CONTENTS
1
2
3
4
5
6
ATOMIC STRUCTURE
9
ATOMIC STRUCTURE WS 1
29
ATOMIC STRUCTURE WS 2
42
ELECTRONIC CONFIGURATION
65
ELECTRONIC CONFIGURATION WS 1
115
ELECTRONIC CONFIGURATION WS 2
117
ELECTRONIC CONFIGURATION WS 3
128
PERIODIC TRENDS
135
PERIODIC TRENDS WS 1
157
PERIODIC TRENDS WS 2
167
COVALENT BONDING
189
COVALENT BONDING WS 1
211
COVALENT BONDING WS 2
217
SHAPES OF MOLECULES
219
SHAPES OF MOLECULES WS 1
237
INTERMOLECULAR FORCES
247
INTERMOLECULAR FORCES WS 1
269
INTERMOLECULAR FORCES WS 2
278
DATA BOOKLET
285
3
HOMEWORK
DATE
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PARTICULARS
4
DATE
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PARTICULARS
5
NOTES
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6
NOTES
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NOTES
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NOTES
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1 Atoms, molecules and stoichiometry
This topic illustrates how quantitative relationships can be established when different
substances react. (The term relative formula mass or Mr will be used for all
compounds including ionic.)
1.1 Relative masses of atoms and molecules
1.3 The determination of relative atomic masses, Ar
2 Atomic structure
2.1 Particles in the atom
2.2 The nucleus of the atom
ATOMIC STRUCTURE
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ATOMIC STRUCTURE
10
Physical chemistry
1 Atoms, molecules and stoichiometry
This topic illustrates how quantitative relationships can be established when different substances react.
(The term relative formula mass or Mr will be used for all compounds including ionic.)
Learning outcomes
Candidates should be able to:
1.1 Relative masses
of atoms and
molecules
a) define and use the terms relative atomic, isotopic, molecular and
formula masses, based on the 12C scale
1.2 The mole and the
Avogadro constant
a) define and use the term mole in terms of the Avogadro constant
1.3 The determination
of relative atomic
masses, Ar
a) analyse mass spectra in terms of isotopic abundances
(knowledge of the working of the mass spectrometer is not required)
1.4 The calculation
of empirical and
molecular formulae
a) define and use the terms empirical and molecular formula
1.5 Reacting masses
and volumes (of
solutions and gases)
a) write and construct balanced equations
b) calculate the relative atomic mass of an element given the relative
abundances of its isotopes, or its mass spectrum
b) calculate empirical and molecular formulae, using combustion data or
composition by mass
b) perform calculations, including use of the mole concept, involving:
(i) reacting masses (from formulae and equations)
(ii) volumes of gases (e.g. in the burning of hydrocarbons)
(iii) volumes and concentrations of solutions
When performing calculations, candidates’ answers should reflect the
number of significant figures given or asked for in the question. When
rounding up or down, candidates should ensure that significant figures
are neither lost unnecessarily nor used beyond what is justified (see also
Practical Assessment, Paper 3, Display of calculation and reasoning on
page 52)
c) deduce stoichiometric relationships from calculations such as those in
1.5(b)
Back to contents page
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ATOMIC STRUCTURE
17
Cambridge International AS and A Level Chemistry 9701 syllabus Syllabus content
11
2 Atomic structure
This topic describes the type, number and distribution of the fundamental particles which make up an
atom and the impact of this on some atomic properties.
Learning outcomes
Candidates should be able to:
2.1 Particles in the atom
a) identify and describe protons, neutrons and electrons in terms of their
relative charges and relative masses
b) deduce the behaviour of beams of protons, neutrons and electrons in
electric fields
c) describe the distribution of mass and charge within an atom
d) deduce the numbers of protons, neutrons and electrons present in both
atoms and ions given proton and nucleon numbers and charge
2.2 The nucleus of the
atom
a) describe the contribution of protons and neutrons to atomic nuclei in
terms of proton number and nucleon number
b) distinguish between isotopes on the basis of different numbers of
neutrons present
c) recognise and use the symbolism xy A for isotopes, where x is the nucleon
number and y is the proton number
2.3 Electrons: energy
levels, atomic
orbitals, ionisation
energy, electron
affinity
a) describe the number and relative energies of the s, p and d orbitals for
the principal quantum numbers 1, 2 and 3 and also the 4s and 4p orbitals
b) describe and sketch the shapes of s and p orbitals
c) state the electronic configuration of atoms and ions given the proton
number and charge, using the convention 1s22s22p6 , etc.
d) (i) explain and use the term ionisation energy
(ii) explain the factors influencing the ionisation energies of elements
(iii) explain the trends in ionisation energies across a Period and down a
Group of the Periodic Table (see also Section 9.1)
e) deduce the electronic configurations of elements from successive
ionisation energy data
f)
interpret successive ionisation energy data of an element in terms of the
position of that element within the Periodic Table
g) explain and use the term electron affinity
18
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Back to contents page
ATOMIC STRUCTURE
estimated that the diameter of the nucleus could not be greater than 0.000 01 times that
of the atom itself. Eventually, Rutherford was able to chip away from this nucleus small
positively charged particles. He showed that
12these were also identical to each other, no
matter which element they came from. This positive particle is called the proton. It is
much heavier than the electron, having nearly the mass of the hydrogen atom.
ht) and
Manchester
seen with
etect and
clei of
WHAT WE KNOW ABOUT ATOMS
Atoms are mostly empty space!
Atoms have 2 charged particles:
• Positively charged protons.
• Negatively charged electrons.
Atoms have 1 uncharged particle called a neutron.
The total number of protons, neutrons and electrons in any one atom
determines its properties.
1
It was another 20 years before the last of the three sub-atomic particles, the neutron,
was discovered. Although its existence was first suspected in 1919, it was not until 1932
that James Chadwick eventually pinned it down. As its name suggests, the neutron is
electrically neutral, but it is relatively heavy, having about the same mass as a proton.
Scientists had therefore to change the earlier picture of the atom. In a sense the
picture had become more complicated, showing that atoms had an internal structure,
and were made up of other, smaller particles. But looked at in another way it had
become simpler – the 90 or so different types of atoms that are needed to make up
the various elements
had been
by just
three sub-atomic
particles. It turns
WHAT
WEreplaced
KNOW
ABOUT
ATOMS
out that these, in different amounts, make up the atoms of all the different elements.
m in the
most of the volume
of the atom is occupied
by the electrons
the nucleus is very small;
it contains the protons
and the neutrons
21
02_02 Cam/Chem AS&A2
Barking Dog Art
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21/10/14 9:22 AM
ATOMIC STRUCTURE
13
THE STRUCTURE OF ATOMS
Mass / g
Charge / C
Relative mass
Relative charge
PROTON
1.672 x 10-24
1.602 x 10-19
1
+1
NEUTRON
1.675 x 10-24
0
1
0
ELECTRON
9.109 x 10-28
1.602 x 10-19
1
1836
-1
3
EXAMPLE
Calculate the mass of a carbon-12 atom.
A neutral 12C atom has 6 protons, 6 neutrons, and 6 electrons.
Therefore, its mass is:
(6 x 1.672 x 10-24) + (6 x 1.675 x 10-24) + (6 x 9.109 x 10-28)
= 2.0089 x 10-23 g
4
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ATOMIC STRUCTURE
Table 2.1 lists some of the properties of the
Table 2.1 The properties of the sub-atomic
particles. Note that the masses in the last row
114
are given relative to 12 the mass of an atom
of carbon-12. These masses are often quoted
relative to the mass of the proton instead,
when the relative mass of the electron is 1 ,
1836
and the relative masses of the proton and
neutron are both 1.
Property
Electro
electrical charge/coulombs
−1.6 ×
charge (relative to that of the proton)
−1
mass/g
9.11 ×
mass/amu (see section 1.4)
5.485
BEHAVIOUR OF SUB ATOMIC PARTICLES
p+
n
e–
–
+
Because of their relative masses and charge
an electric field, as shown in Figure 2.3. Ne
neutral. Protons are attracted towards the n
positive pole. If their initial velocities are th
greater extent than protons because they ar
The picture of the atom assembled from
● Atoms are small, spherical structures with
beam of particles
5
Figure 2.3 The behaviour
of protons,
neutrons and electrons in an electric field
02_03 Cam/Chem AS&A2
Barking Dog Art
ISOTOPES
Isotopes are atoms with:
3 × 10−10 m.
● The particles that contribute to the atom’s
within a very small central nucleus that ha
● The electrons occupy the region around the
inside the atom but outside the nucleus, wh
● All the atoms of a particular element cont
equals the number of electrons within tho
● The atoms of all elements except hydroge
nucleus along with the protons. Almost th
of the atom is to increase its mass.
2.3 Isotopes
At the same time as Rutherford and his team
the nucleus, it was discovered that some ele
masses, but identical chemical properties.
These atoms were given the name isotop
place (topos) in the Periodic Table. The first
of the unstable radioactive element thorium
Periodic Table.) In 1913, however, Thomson
obtained from liquid air contained atoms w
as those with the usual relative atomic mass
stable, unlike the thorium isotopes. Many o
which are listed in Table 2.2.
the same number of protons and electrons but different number of
Table 2.2 The relative abundance of
neutrons.
Isotope
some isotopes
the same atomic number but different mass number boron-10
boron-11
Mass relative
10.0
11.0
neon-20
Chemical properties of isotopes are identical, whereas,
physical
neon-22
properties (such as density) can differ.
20.0
magnesium-24
24.0
magnesium-25
25.0
magnesium-26
26.0
The isotopes of an element differ in their
composition in only one respect – although
they all contain the same
numbers of
6
electrons and protons, they have different
numbers of neutrons.
Most of the naturally occurring isotopes are
and also many artificially made ones, are un
radioactive isotopes.
22
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ATOMIC STRUCTURE
181333_02_AS_Chem_BP_020-043.indd 22
22.0
e proton and
1836
5.485 × 10−4
mass/amu (see section 1.4)
1.007
1.009
Because of their relative masses and charges, the three particles behave differently in
15
an electric field, as shown in Figure 2.3. Neutrons are undeflected, being electrically
neutral. Protons are attracted towards the negative pole, and electrons towards the
positive pole. If their initial velocities are the same, electrons are deflected to a
greater extent than protons because they are much lighter.
ISOTOPES
The picture of the atom assembled
from these observations is as follows.
e–
+
−10
● Atoms are small, spherical structures with diameters ranging from 1 × 10
m to
tritium
protium
deuterium
3 × 10 m.
● electron
The particles that contribute to
the atom’s mass (protons and neutrons) are contained
neutron
within a very small central nucleus that has a diameter of about 1 × 10−15 m.
● proton
The electrons occupy the region around the nucleus. They are to be found in the space
inside the atom but outside the nucleus, which is almost the whole of the atom.
● All the atoms of a particular element contain the same number of protons. This also
equals the number of electrons within those atoms.
protons
1
1
1
● The atoms of all elements except hydrogen also contain neutrons. These are in the
neutrons
0 the protons. Almost the only
1 effect they have on the properties
2
nucleus along with
3
1 increase its mass.
2
of the symbol
atom is to
isotopic
1
1
1
−10
cles
of protons,
electric field
H
H
H
2.3 Isotopes
gure 2.8 The atomic structure and isotopic symbols for the three isotopes of hydrogen.
7
At the same time as Rutherford and his team
were finding out about the structure of
nucleus, it was discovered that some elements contained atoms that have different
he nucleon numberthe
after
the name. For example,
The chloride ion has a single negative charge becaus
masses, but identical chemical properties.
he isotopes of hydrogen
can be called hydrogen-1,
there are 17 protons (+) and 18 electrons (−
).
These atoms were given the name isotopes, since they occupy the same (iso)
ydrogen-2 and hydrogen-3.
place (topos) in the Periodic Table. The first isotopes to be discovered were those
Mgin2+the +
2e−
of the unstable radioactive element thorium. (ThoriumMg
is element→
number 90
magnesium
2 electron
Table.) In 1913, however, Thomson was magnesium
able to show that a sample
of neon
Check-up Periodic
atom
ionwell
removed
obtained from liquid air contained atoms with a relative
atomic mass of 22 as
as those with the usual relative atomic mass of 20.12
These
heavier neon atoms
were
protons
12 protons
3 Use the Periodic
Table
on
page
497
to
ISOTOPES
stable, unlike the thorium isotopes. Many other elements
contain isotopes,
some of
12 electrons
10 electrons
help you. which are listed in Table 2.2.
The magnesium ion has a charge of 2+ because it ha
a Write isotopic symbols for the following
Isotope
Mass relative to hydrogen
abundance
protons (+)Relative
but only
10 electrons (−
).
neutral atoms:
The isotopic
boron-10
10.0
20%symbol for an ion derived from sulfu
i bromine-81
33 2−
is sulfide ion has 16 protons, 17 neutr
33 is 16S . Th
boron-11
11.0
80%
ii calcium-44
(because 33 −
16 = 17) and 18 electrons (because 16 + 2
20.0
91%
iii iron-58neon-20
neon-22
22.0
9%
iv palladium-110
79%
b What is themagnesium-24
number of protons and 24.0
25.0
10%
neutrons inmagnesium-25
each of these atoms?
Check-up
ndance of
nt differ in their
espect – although
numbers of
ey have different
magnesium-26
26.0
4 Deduce the number of electrons in each of
Most of the naturally occurring isotopes are stable, but some, like those of uranium
ions: These are called
and also many artificially made ones, are unstable andthese
emit radiation.
+
radioactive isotopes.
a 40
19K
2.4 How many protons,
neutrons and electrons?
8
n a neutral atom the number of positively charged
rotons in the nucleus equals the number of negatively
harged electrons outside the nucleus. When an atom
ains or loses electrons, ions are formed which are
ectrically charged.
For
example:
BILAL
HAMEED
Cl
11%
+
e−
→
Cl−
b 157N3−
c
18
8
O2−
3+
d 71
31Ga
15/10/14 12:19 PM
ATOMIC STRUCTURE
16
RELATIVE MASSES
The 12C (carbon-12) isotope is chosen as a standard and given a relative
mass of exactly 12.
Important relative masses measured against this standard are:
Relative Atomic Mass (Ar)
The relative atomic mass (Ar) of an element is the weighted average mass
of the naturally occurring isotopes of the element relative to one-twelfth
the mass of the Carbon-12 isotope.
9
RELATIVE MASSES
Relative Isotopic Mass
The relative isotopic mass of an element is the mass of an atom of the
isotope of the element relative to one twelfth of the mass of an atom of
the isotope carbon-12.
Relative Molecular Mass (Mr)
The relative molecular mass (Mr) of a compound is the mass of a
molecule of that compound relative to one-twelfth the mass of the
Carbon-12 isotope.
10
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ATOMIC STRUCTURE
17
RELATIVE MASSES
Relative Formula Mass
Used for ionic compounds and any formula of a species or ion e.g. NaCl
or OH— etc…
11
SKILL CHECK 1
A How many protons, electrons and neutrons are present in a sulphide
ion, S2-? Sulphur has atomic number 16 and mass number 32.
B How many protons, neutrons and electrons are present in a potassium
ion, K+? Potassium has atomic number 19 and mass number 39.
12
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ATOMIC STRUCTURE
18
MASS SPECTROMETER
A mass spectrometer is used to calculate the relative atomic mass of an
atom by comparing it with the mass of a 12C atom.
You can also use a mass spectrometer to calculate the relative
abundance of different isotopes of any element.
2 .1 T H E N U C L E A R ATO M
●
13
The relative atomic mass, Ar , is the ratio of the average mass of the
atom to the unified atomic mass unit.
As mentioned in sub-topic 1.2, the average mass of the atom is a weighted
average of the atomic masses of its isotopes and their relative abundances.
Study tip
Relative atomic ma
so it does not have u
The mass spectrometer
The mass spectrometer is an instrument used to determine the relative
atomic mass of an element. It can also show its isotopic composition.
HOW DOES A MASS SPECTROMETER WORK?
positive ions are
accelerated in the electric
field (stage 3)
heating filament to vaporize
sample (stage 1)
inlet to inject
sample
lighest particles
(deflected most)
detector
(stage 5)
magnet (stage 4)
heaviest particles
(deflected least)
N
electron beam to
ionize sample (stage 2)
S
14
Figure 10 Schematic diagram of a mass spectrometer
There are five stages in this process:
Stage 1 (vaporization): The sample is injected into the instrument
where it is heated and vaporized, producing gaseous atoms or
BILAL HAMEED
ATOMIC STRUCTURE
molecules.
●
●
Stage 2 (ionization): The gaseous atoms are bombarded by high-
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HOW DOES A MASS
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edh^i^kZ
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BILAL HAMEED
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ATOMIC STRUCTURE
20
HOW DOES A MASS SPECTROMETER WORK?
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Note: particles are defl
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17
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18
BILAL HAMEED
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Figure 2.10. Ar = (78.6 × 24) + (10.1 ×
1
The&&#(relative atomic mass can b
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40
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M(g)STRUCTURE
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ATOMIC
field in a mass spectrometer.
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21
MASS SPECTRA
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19
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5 Chromium has four naturally occurring isotopes,
and their masses and natural abundances are shown
in the table below.
>hdideZ
SKILL
CHECK 2
CVijgVaVWjcYVcXZ
The graph shows
*%
8g the mass spectrum of Neon.
)#(*
*'
8g
-(#,.
1. How many isotopes
does Neon
*(
8g
.#*%
contain?
*)
8g
'#(+
2. What are the relative isotopic
Calculate
the relative
atomic mass of chromium to
masses of the
isotopes
of Neon?
two decimal places.
3. What are the relative amounts of the
6 Silicon has three naturally occurring isotopes and
isotopes oftheir
Neon?
details are given in the table below.
>hdideZ
'-
H^
'.
H^
(%
H^
7 a Indium has two naturally
indium-113 and indium-1
atomic mass of indium is
natural abundance of each
b Gallium has two naturally
gallium-69 and gallium-7
mass of gallium is 69.723.
abundance of each isotope
8 Arrange the following ions i
they will be deflected in a m
one that will be deflected m
84
2+
81
+
1
2 3Kr
35Br
CVijgVaVWjcYVcXZ
.'#'
)#,
20
(#&
Calculate the relative atomic mass of silicon to two
decimal places.
BILAL HAMEED
ATOMIC STRUCTURE
2 .1 T H E N U C L E A R ATO M
tio of the average mass of the
22
Ar of neon
Study
(20.0 ×tip
90.9) (21.0 × 0.3) (22.0 × 8.8)
= 20.2
=
100is a ratio
Relative atomic mass
Fact file
A high-resolution mass spectrometer can gi
16
relative isotopic masses. For example O =
32
S = 31.972. Because of this, chemists can d
which is consistent with the data given.
molecules such as SO2 and S2 which appear
Calculate the average relative atomic mass of neon using the information
relative molecular mass.
mass of the atom is a weighted
it does
not have
units.
that
this answer
is given
to 3 significant figures,
nd their relative abundances. Noteso
SKILL
CHECK
3
The mole and the Avogadro con
60
40
20
0
heaviest particles
(deflected least)
8.8 %
detector
(stage 5)
1.4 Amount of substan
80
0.3 %
Relative abundance / %
lighest particles
(deflected most)
100
90.9 %
below:
used to determine the relative
its isotopic composition.
19
20
21
22
Mass/charge (m/e) ratio
23
21
Figure 1.5 The mass spectrum of neon, Ne.
Check-up
40
36.7 %
2 Look at the mass spectrum of germanium, Ge.
The formula of a compound shows us
atoms of each element present in one fo
molecule of the compound. In water w
atoms of hydrogen (Ar = 1.0) combine
of oxygen (Ar = 16.0). So the ratio of m
atoms to oxygen atoms in a water mole
matter how many molecules of water w
will always be the same. But the mass o
is far too small to be weighed. We have
more than this to get an amount of sub
easy to weigh.
The relative atomic mass or relative m
a substance in grams is called a mole of
So a mole of sodium (Ar = 23.0) weigh
abbreviation for a mole is mol. We defi
terms of the standard carbon-12 isotop
SKILL CHECK 4
re deflected by a magnetic
gree of deflection depends on
The species with the smallest
be deflected the most.
d in the magnetic field.
s species of a particular m/z
ectrical signal is generated.4
7.6 %
100
0
Relative abundance
are attracted to negatively
ctric field.
10
7.7 %
following peaks in its mass spectrum:
are bombarded by highharged species:
27.4 %
30
20.6 %
Abundance / %
One mole of a substance is the amou
substance which has the same numb
particles
Calculate the average relative atomic mass of an element producing
the (atoms, molecules or ions) a
20
atoms in exactly 12 g of the carbon-1
jected into the instrument
cing gaseous atoms or
80.1 (area under peak)
70
75
Mass/charge (m/e) ratio
80
Figure501.6 The mass spectrum of germanium.
a Write the isotopic
forpeak)
the heaviest
19.9 formula
(area under
isotope of germanium.
b 0Use the % abundance of each isotope
2
4 relative
6 atomic
8
10 of12
to0 calculate
the
mass
m/z
germanium.
22 of boron. The two
Figure 11 Mass spectrum
to two isotopes
1peaks
Molescorrespond
and equations
We often refer to the mass of a mole of
molar mass (abbreviation M). The uni
are g mol−1.
The number of atoms in a mole of at
6.02 × 1023 atoms. This number is calle
constant (or Avogadro number). The s
Avogadro constant is L. The Avogadro
to atoms, molecules, ions and electrons
sodium there are 6.02 × 1023 sodium at
of sodium chloride (NaCl) there are 6.0
ions and 6.02 × 1023 chloride ions.
positive ions of a single
n depend only on the mass.
ative abundance (of each
The height
of each
peak
BILAL
HAMEED
ective isotope.
ATOMIC STRUCTURE
23
SKILL CHECK 5
Boron has two naturally occurring isotopes with the natural
abundances shown. Calculate Boron’s relative atomic mass.
ISOTOPE
NATURAL ABUNDANCE (%)
10 B
19.9
11 B
80.1
23
SKILL CHECK 6
Calculate the average relative atomic mass of lead using the information
below:
Isotopic mass
Relative abundance/%
204
2
206
24
207
22
208
52
24
BILAL HAMEED
ATOMIC STRUCTURE
atoms to oxygen atoms in a water mo
19
20
21
22
23
matter how many molecules of water
Mass/charge (m/e) ratio
will always be the same. But the mas
24
is far too small to be weighed. We ha
Figure 1.5 The mass spectrum of neon, Ne.
more than this to get an amount of s
easy to weigh.
The relative atomic mass or relative
SKILL
CHECK
7
a
substance
in grams is called a mole
Check-up
So a mole of sodium (Ar = 23.0) weig
Calculate the average
relative
atomic
mass
of
Germanium
using
the
abbreviation
for a mole is mol. We d
2 Look at the mass spectrum of germanium, Ge.
terms of the standard carbon-12 isot
information below:
36.7 %
0
10
0
7.6 %
7.7 %
20
One mole of a substance is the am
substance which has the same num
particles (atoms, molecules or ions
atoms in exactly 12 g of the carbon
27.4 %
30
20.6 %
Abundance / %
40
70
75
Mass/charge (m/e) ratio
80
Figure 1.6 The mass spectrum 25
of germanium.
a Write the isotopic formula for the heaviest
isotope of germanium.
b Use the % abundance of each isotope
to calculate the relative atomic mass of
germanium.
4
We often refer to the mass of a mole
molar mass (abbreviation M). The u
are g mol−1.
The number of atoms in a mole of
6.02 × 1023 atoms. This number is cal
constant (or Avogadro number). Th
Avogadro constant is L. The Avogadr
to atoms, molecules, ions and electro
sodium there are 6.02 × 1023 sodium
of sodium chloride (NaCl) there are
ions and 6.02 × 1023 chloride ions.
1 Moles and equations
MASS SPECTRA: CALCULATING ABUNDANCE
Naturally occurring potassium consists of potassium-39 and
potassium-41. Calculate the percentage of each isotope present if the
average is 39.1:
Assume there are x nuclei of 39K in every 100, so there will be (100 – x)
of 41K
This means that:
39x + 41 (100 – x)
100
= 39.1
26
BILAL HAMEED
ATOMIC STRUCTURE
25
MASS SPECTRA: CALCULATING ABUNDANCE
therefore
39x + 4100 - 41x
thus
- 2x = - 190
= 3910
x = 95
There will be 95% 39K and
5% 41K!
27
SKILL CHECK 8
Silver has two isotopes. 51.35% of the atoms are Silver-107 and
48.65% of the atoms are Silver-109. Calculate the relative atomic mass
of Silver.
28
BILAL HAMEED
ATOMIC STRUCTURE
26
SKILL CHECK 9
The relative atomic mass of copper is 63.5. Calculate the
relative abundance of the two copper isotopes with relative isotopic
masses of 63.0 and 65.0
29
SKILL CHECK 10
Silver has two isotopes. 51.35% of the atoms are Silver-107 and
48.65% of the atoms are Silver-109. Calculate the relative atomic mass
of Silver.
30
BILAL HAMEED
ATOMIC STRUCTURE
27
SKILL CHECK 11
A sample of element X contains 69% of 63X and 31% of 65X. What is the
relative atomic mass of X in this sample?
31
SKILL CHECK 12
A sample of element X contains 69% of 63X and 31% of 65X. What is the
relative atomic mass of X in this sample?
32
BILAL HAMEED
ATOMIC STRUCTURE
28
MASS SPECTRA: MOLECULAR IONS
Molecules also give peaks in their mass spectrum.
For example the mass spectrum of chlorine gaseous molecule includes
both peaks of its isotopes and peaks of its molecular form (molecular
ions):
33
MASS SPECTRA: MOLECULAR IONS
Cl2+ molecular ion produces three peaks:
peak at m/e @ 70: 35Cl - 35Cl
Species responsible: 70Cl2+
peak at m/e @ 72: 35Cl - 37Cl
Species responsible: 72Cl2+
peak at m/e @ 74: 37Cl - 37Cl
Species responsible: 74Cl2+
34
BILAL HAMEED
ATOMIC STRUCTURE
29
ATOMIC STRUCTURE WS 1
SECTION A
1 Which one the following has more neutrons than electrons and more electrons than
protons?
A 19F--
C 9Be
B 37Cl--
D 9Be2+
2 Chlorine exists as two isotopes 35Cl with an abundance of 75.5% and 37Cl with an abundance
of 24.5%.
Phosphorus has only one isotope, 31P. The mass spectrum of PCl3 has four lines at m/z = 136,
138, 140 and 142.
Which one of these lines will have the smallest height?
A 136
C 138
B 140
D 142
3 Antimony has two isotopes 121Sb and 123Sb. The relative atomic mass of a naturally occurring
sample of antimony is measured as 121.75.
Which one of the following is the best approximate estimate of the percentage of 121Sb
present in the naturally occurring sample?
A 20%
C 40%
B 60%
D 80%
4 When sulfur, 32S is bombarded with neutrons 1n, two particles are formed. One of them is a
hydrogen atom, 1H and the other is an element, X.
32S
+
1n
1H
+ X
Which one of the following correctly represents X?
5
A 32S
C 33S
B 32P
D 33P
What is the number of protons, electrons, and neutrons in boron-11?
A 5 protons, 5 electrons and 11 neutrons
C 5 protons, 5 electrons and 6 neutrons
B 5 protons, 5 electrons and 10.8
neutrons
D 11 protons, 11 electrons and 5 neutrons
CEDAR COLLEGE
ATOMIC STRUCTURE WS 1
30
6
7
What is the number of protons, electrons, and neutrons in 34S2--?
A 18 protons, 16 electrons and 18 neutrons
C 16 protons, 18 electrons and 18 neutrons
B 16 protons, 18 electrons and 18 neutrons
D 16 protons, 16 electrons and 18 neutrons
Which statements about the isotopes of chlorine, 35Cl and 37Cl are correct?
I. They have the same chemical properties.
II. The have the same atomic number.
III. The have the same physical properties.
8
9
A I and II only
C II an III only
B I and III only
D I, II and III
A sample of element X contains 69% of 63X and 31% of 65X. What is the relative atomic mass
of X in this sample?
A 63.0
C 65.0
B 63.6
D 69.0
What is the relative atomic mass of an element with the mass spectrum shown below?
A 24
C 26
B 25
D 27
CEDAR COLLEGE
ATOMIC STRUCTURE WS 1
A
electron
C
neutron
Section A
2
For each
there are four possible answers, A, B, C, and D. Choose the one you consider to be
B question
ion
correct.
31 A
Section
For each question there are four possible answers, A, B, C, and D. Choose the one you consider to be
10
1correct.
InDthe proton
radioactive decay of an isotope of lead to an isotope of bismuth, a particle –10 X is emitted.
12
Which particle is –10 X?
As use
a simplification,
an adult
human
can
be considered to have a daily diet of 1.80 kg of
The
of the Data Booklet
is relevant
to this
question.
O).
carbohydrate
(empirical
formula
CH
A electron
2
What is the number of molecules in 500 cm3 of oxygen under room conditions?
BWhich
ion mass of carbon dioxide does a person produce each day if all the carbohydrate eaten is
Adigested
1.25 xand
1022oxidised?
C
neutron
B proton
1.34 x 1022
D
A 0.267 kg
C
B
3.0 x 1022
C
0.800 kg
D
1.32 kg
2.64 kg
[S’02 Q1]
2
11
3
D a3.0
x 1026
As
simplification,
an adult human can be considered to have a daily diet of 1.80 kg of
The diagram shows the mass spectrum of a sample of naturally-occurring copper.
carbohydrate (empirical formula CH2O).
2
In the preparation
of softdioxide
margarine,
trieleostearate
Which
mass of carbon
doesglyceryl
a person
produce each day if all the carbohydrate eaten is
7
digested and oxidised?
6
CH3(CH2)3CH
CHCH 5 CHCH
CH(CH2)7CO2CH2
A 0.267 kg
B 0.800 kg
D 2.64 kg
relative C 1.32 kg
4
abundance
CH3(CH2)3CH
CHCH
CHCH
CH(CH2)7CO2CH
3
The diagram shows
mass spectrum
of2a sample
naturally-occurring
copper.
CHCH
CHCH of CH(CH
CHthe
3(CH2)3CH
2)7CO2CH2
1
3
is suitably hydrogenated so that, on 7average, one of its side-chains is converted into the
61 62 63 64 65 66 67 68
CH3(CH2)4CH=CHCH2CH=CH(CH2)7CO
6 2 residue and two side-chains are converted into the
m/e
CH3(CH2)7CH=CH(CH2)7CO2 residue. 5
relative
4
abundance
How
many
moles
of hydrogen
are required
convert one mole of glyceryl trieleostearate into the
What
is the
relative
atomic mass
of 3thisto
copper?
soft margarine?
2
A
63.3
B
63.5
C D 63.6
D 64.0
A 4
B 5
C 6 1
9
[S’02 Q3]
61 62 63 64 65 66 67 68
3
12
Which isotope of an element in the third period of them/e
Periodic Table contains the same number of
32
neutrons as 16 S?
What23is the relative atomic mass of this copper?
A
11 Na
A
B
24
63.3
12 Mg
C
28
14 Si
31
15 P
D
B
C
63.5
D
63.6
64.0
3
[S’03 Q3]
–1 9701/1/M/J/02
4
The successive ionisation energies, in kJ mol , of an element X are given below.
13 Unnilpentium is an artificial element. One of its isotopes is 262Unp.
4
105
870
1800
3000
3600
5800
7000
13 200
Which of the following statements is correct?
What is X?
A 262
105Unp has a nucleon number of 105.
B 40Zr
C 52
Te
D 53I
A 33As
262
B The atom 260
9701/1/M/J/02
105X is an isotope of 105Unp.
C
There are 262 neutrons in 262
105Unp.
D
The proton number of 262
105Unp is 262.
9701/1/M/J/03
[W’03 Q4]
5
The table gives the successive ionisation energies for an element X.
CEDAR COLLEGE
ionisation energy / kJ mol–1
1st
2nd
3rd
WS 1
4th ATOMIC
5th STRUCTURE
6th
950
1800
2700
4800
6000
12300
B
2
3
mixture
contains
304cm of carbon
dioxide and 20 cm of unreacted oxygen. All gaseous volumes
were measured under identical conditions.
C
3
4
3½
32
What
D is the
3 formula4of the hydrocarbon?
3
A
14
2
2
C 2H 6
B
C3H6
C
C3H8
D
C4H10
The diagram shows the mass spectrum of a sample of zinc. Use the data to calculate the relative
atomic mass of the sample.
On collision, airbags in cars inflate rapidly due to the production of nitrogen.
100
The nitrogen is formed according to the following equations.
2NaN3 → 2Na + 3N80
2
relative
10Na
+ 2KNO3 → K60
2O + 5Na2O + N2
abundance / %
How many moles of nitrogen gas are produced
from 1 mol of sodium azide, NaN3?
40
A
1.5
B
1.6
C
20
3.2
D
4.0
0
The first six ionisation energies of four elements,
A to D, are given.
66 67
64 65
Ar
Which element is most likely to be in Group IV of the Periodic Table?
3
A 65
B 65.25
C 65.5
D 65.66
ionisation
2
[W’04 Q2]
4 Use of the Data−1Booklet
question.4th
1stis relevant
2nd to this3rd
5th
6th
energy / kJ mol
3 The
thatthat
skunks
spray
is due Section
toexists
a number
of thiols,
one ofthe
which
is methanethiol,
A several
15
It is foul
nowAsmell
thought
where
an element
as
isotopes,
stable
ones usually
494
4560
6940
9540
13400 16600
CH
SH,
which
burns
as
follows.
3
contain a ‘magic number’ of neutrons. One of these magic numbers is 126.
For each question
there are
A,
B, C, and
D. Choose
B
736four possible
1450 answers,
7740
10500
13600
18000the one you consider to
CH
SH
+
3O
→
CO
+
SO
+
2H
O
be correct.
3
2
2
2
2
Which isotope is unstable?
C
1090
2350
4610
6220
37800 47000
3
3
209
208
208
AA sample
was exploded
with
60
cm
Bi of 10 cm ofBmethanethiol
Pb
C 210
Po
Dof oxygen.
Tl
D
1400
2860
4590 2 7480
9400
1 The petrol additive tetraethyl-lead(IV), Pb(C2H5)4, is now banned53200
in many countries. When
it is
[W’04 Q4]
What
wouldburned
be the final
ofoxide,
the resultant
mixture
of gases
when cooled to room
and
H
O
are
formed.
completely
in air,volume
lead(II)
CO
2
2
Section A
temperature?
5
An
atom
has
eight
electrons.
16
4 How
In which
species
are
the
numbers
of electrons
and one
neutrons
equal?
many
moles
ofare
oxygen
are
to burn
mole
Pb(C2H5)3the
4? one you consider to
3 required
3 C,
For each
question
there
possible
answers,
and of
D.
A 20
cm3
B four
30 cm
C 23
50A,
cmB,
D Choose
70 cm
18
+
2–
9
19
WhichBe
diagram shows
electronic configuration
in itsOlowest energy state?
be correct.
A
BB the11
C C 1113.5
Na of this atom
D D
4
9F
8 27
A
9.5
3
[S’05 Q4]
A2O4 is a poisonous gas. It can be disposed of safely by reaction with sodium hydroxide.
N
Which ion has more electrons than protons and more protons than neutrons?
N2O4(g) + 2NaOH(aq) → NaNO3(aq) + NaNO2(aq) + H2O(l)
[H = 11 H ; D = 12 H ; O = 168 O]
B
What
is the minimum volume of 0.5 mol dm–3 NaOH(aq) needed to dispose of 0.02 mol of N2O4?
B H3O+
C OD–
D OH–
A D–
B 12.5 cm3
C 40 cm3
D 80 cm3
A 8 cm3
© UCLES 2004
1
17
2
9701/01/O/N/04
[W’05 Q2]
© UCLES 2005
3
18
2
9701/01/M/J/05
C
What
is theofelectronic
of an element
withnumbers
a second35ionisation
higher than
A sample
chlorine configuration
containing isotopes
of mass
and 37 energy
was analysed
in a
that
of
each
of
its
neighbours
in
the
Periodic
Table?
mass-spectrometer.
A
1s22s22p63s2
D
How many peaks corresponding to Cl 2+ were recorded?
B 1s22s22p63s23p1
A 2 2 2 6 2 2B 3
C 4
D 5
2s 2pa 3s
3p lizard, can climb up a smooth glass window. The gecko has millions of
6 C
The 1s
gecko,
small
[S’06 Q2]
microscopic hairs on its toes and each hair has thousands of pads at its tip. The result is that the
D 1s22s22p63s23p3
molecules
in theGaN,
pads could
are extremely
closethe
to the
glassofsurface
which
the because
gecko is climbing.
3 Gallium
nitride,
revolutionise
design
electriconlight
bulbs
only a small
length used as a filament gives excellent light at low cost.
What is the attraction between the gecko’s toe pads and the glass surface?
4 Which compound has a boiling point which is influenced
by hydrogen bonding?
Gallium nitride is an ionic compound containing the Ga3+ ion.
A co-ordinate bonds
A CH3CHO
What
is the electron
B covalent
bonds arrangement of the nitrogen ion in gallium nitride?
B CH3OCH3
2
2
2sbonds
A
C 1s
ionic
C HCO2H
2
2
3
2s
2p
B
D 1s
van
der
Waals’
forces
D HCO2CH3
CEDAR
ATOMIC STRUCTURE WS 1
C COLLEGE
1s2 2s2 2p4
7
5
2
2
6
What
angles in the PH3 molecule likely to be?
2sthe
2pbond
D
1sare
Which gas is likely to deviate most from ideal gas behaviour?
B
1s 2s 2p
2
4
For each
question
are four possible answers, A, B, C, and D. Choose the one you consider to
C 1s
2s2 2pthere
be correct.
D 1s2 2s2 2p6
33
1 The isotopic composition of an element is indicated below.
19
4 A radioactive isotope of thallium, 201
81 Tl, is used to assess damage in heart muscles after a heart
5
attack.
4
Which statement about 201
81 Tl is correct?
relative 3
A This isotope
has
a nucleon number of 120. 2
abundance
2
B The number of electrons in one atom of this isotope is 81.
Section A
1
C The number of neutrons in one atom of this isotope is 201.
0 possible answers, A, B, C, and D. Choose the one you consider to
201
For each
question there are four
D 201
82 X is an isotope of 81 Tl.
9
10
11
12
be correct.
mass number
[S’06 Q4]
20
the relative
atomic mass
the element?
1 What
The is
isotopic
composition
of anofelement
is indicated below.
A
10.2
B
10.5
5
C
10.8
D
11.0
4
Use2006
of the Data Booklet is relevant to this 9701/01/M/J/06
question.
2© UCLES
Oxides of nitrogen arerelative
pollutant 3gases which are emitted from car exhausts.
abundance
2 one kilometre, it releases 0.23 g of an oxide of nitrogen N O ,
In urban traffic, when a car travels
x y
3
which occupies 120 cm .
1
What are the values of x and y?
(Assume 1 mol of gas molecules0 occupies 24.0 dm3.)
9
A
x = 1, y = 1
B
x = 1, y = 2
10
11
12
mass number
What is the relative atomic mass of the element?
C
x = 2, y = 1
A
10.2
D
x = 2, y = 4
B
10.5
C
10.8
D
11.0
[S’07 Q1]
Use Dalton’s
of the Data
Booklet
relevant to
this question.
21
32 John
atomic
theory,ispublished
in 1808,
contained four predictions about atoms.
Oxides
areispollutant
gases to
which
are emitted from car exhausts.
Which
of of
hisnitrogen
predictions
still considered
be correct?
AIn urban
Atoms traffic,
are very
smallaincar
size.
when
travels one kilometre, it releases 0.23 g of an oxide of nitrogen NxOy,
which occupies 120 cm3.
B
No atom can be split into simpler parts.
the values
of x and element
y?
CWhat
All are
the atoms
of a particular
have the same mass.
(Assume 1 mol of gas molecules occupies 24.0 dm3.)
D
All the atoms of one element are different in mass from all the atoms of other elements.
A
x = 1, y = 1
B
x = 1, y = 2
© UCLES 2007
3
C
x = 2, y = 1
D
x = 2, y = 4
[S’07 Q3]
9701/01/M/J/07
John Dalton’s atomic theory, published in 1808, contained four predictions about atoms.
Which of his predictions is still considered to be correct?
A
Atoms are very small in size.
CEDARBCOLLEGE
No atom can be split into simpler parts.
C
All the atoms of a particular element have the same mass.
ATOMIC STRUCTURE WS 1
Section A
1 ForIneach
the question
Basic Oxygen
steel-making
P4 O
impurity
is removed
there are
four possibleprocess
answers,the
A, B,
C,10 and
D. Choose
the oneby
youreacting
considerittowith
34
3
calcium
oxide.
The
only
product
of
this
reaction
is
the
salt
calcium
phosphate,
Ca
(PO
be correct.
3
4 )2 .
3
4 this
The
diagrams
the
possible
paths oxide
of subatomic
particles
moving
In
reaction,
howshow
many
moles
of calcium
react with
one mole
of P4 in
O10an
? electric field in a
The
diagrams
the ispossible
subatomic particles moving in an electric field in a
Usevacuum.
of the Datashow
Booklet
relevantpaths
to thisofquestion.
Avacuum.
1
B 1.5
C 3
D 6
+ oxide produced is used in +the
Titanium(IV) oxide, TiO2, +is brilliantly white and much of the
+
+
manufacture of paint. +
2 Use
of the Data Booklet is relevant to
this question.
electrons
neutrons
protons
obtainable from 19.0 tonnes protons
of the ore ilmenite, FeTiO3?
What
is
the
maximum
amount
of
TiO
2
electrons
neutrons
A typical solid fertiliser for use with household plants and shrubs contains the elements N, P, and
– gtonnes
10.0
tonnes
14.0 tonnes
D– 17.7 tonnes
KAin the
ratio
of 15 g :B30 g12.7
:15
per 100 g Cof fertiliser.
The recommended
usage of fertiliser is–14 g
3
–
–
–
2
3
of fertiliser per 5 dm 1of water.
1
2
3
2 What
Carbon
disulphide
vapour
in oxygen
to the following equation.
Which
diagrams
are correct?
is the
concentration
ofburns
nitrogen
atomsaccording
in this solution?
Which diagrams are correct?
CS2(g) + 3O2(g) → CO2(g) + 2SO2(g)
A 1mol
anddm
2 –3
only
A 0.03
A 1 and 2 only–3
B 1mol
and
3 only
BA sample
0.05
cm3 of carbon disulphide was burned in 50 cm3 of oxygen. After measuring the
ofdm
10
B
1 and
3gas
onlyremaining,
volume
ofand
the product was treated with an excess of aqueous sodium hydroxide
C 2mol
3 –3
only
Cand0.42
dm
the volume
of gas measured again. All measurements were made at the same temperature
C 2pressure,
and 3 only
such conditions that carbon disulphide was gaseous.
–33
D 1,
2 and
Dand0.75
mol
dmunder
41
22
D 1, 2 and 3
[S’07 Q4]
What were the measured volumes?
5 The CN– ion is widely used in the synthesis of organic compounds.
32
3 Skin cancer
can be treated using a radioactive isotope of phosphorus, 15
P . A compound
23
volume
of
gas
after
volume
gas in the32synthesis
widelyofused
of
organic
compounds.
5 The CN– ion is
3−
What is after
the
electron
pairs, in
thisNaOH(aq)
ion?in a plastic
3
wrapped
containing
thepattern
phosphide
ion
15 P adding
/ cm3 sheet, is strapped to the affected
burningof/ cm
area. is the pattern of electron pairs in this ion?
What
A
0 pairs on
bonding 30
pairs of
lone pairs on
lone
32 3 −
electrons
carbon
atom
nitrogen
atom
What
the composition
the pairs
phosphide
ion,
P ?on
15pairs
pairs
on
lone
B isbonding
30 of oflone
20
electrons
carbon atom
nitrogen atom
250
1
C A
20 1
protons
neutrons
electrons
A B
2 2
1 2
1 1
D
50
40
152 3
17
BA C
2 1 18
1 1
3
64
24
B D
153 3
17
32
1 electron
1 in2 an s orbital in their ground states?
InCwhich
pair do
both atoms have1 one
only
173
15
17
DC
1
2
A Ca, Sc
B Cu, Be
C H, He
D Li, Cr
6 DPlastic bottles
drinks’ are made
32 for ‘fizzy 17
15 from a polymer with the following structure.
[S’08 Q3]
Plastic
forBooklet
‘fizzy drinks’
are made
a polymer with the following structure.
Use of bottles
the Data
is relevant
to thisfrom
question.
H
X
C
Hard water contains calcium ions and hydrogencarbonate
ions arising from dissolved calcium
H
X
hydrogencarbonate, Ca(HCO3)2.
C
C
H
H n anion?
How many electrons are present in the hydrogencarbonate
C
H 32
A The
30 ability of theBpolymer
31
D 33 through the wall of the bottle
to preventCescape
ofH carbon dioxide
n
Q4]
depends on the ability of the group X to form hydrogen bonds with the carbon dioxide in [W’08
the drink.
The ability of the polymer to prevent escape of carbon dioxide through the wall of the bottle
Whichon
group
X bestofprevents
loss
dioxide?bonds with the carbon dioxide in the drink.
depends
the ability
the group
X of
to carbon
form hydrogen
© UCLES 2008
A
9701/01/M/J/08
Cl
B CN
C CO2 CH3
Which group X best prevents loss of carbon dioxide?
A 2008
Cl
© UCLES
© UCLES 2007
© UCLES 2007
CEDAR COLLEGE
B
CN
C
CO2 CH3
9701/01/O/N/08
9701/01/M/J/07
9701/01/M/J/07
D
D
OH
OH
[Turn over
[Turn over
ATOMIC STRUCTURE WS 1
Section A
be correct.
For each question there are four possible answers, A, B, C, and D. Choose the one you consider to
1 correct.
The molecular energy distribution curve represents
the variation in energy of the molecules of a
be
35
gas at room temperature.
25
1
An element X consists of four isotopes. The mass spectrum of X is shown in the diagram.
2
proportion
3
of
80
molecules
In the ideal gas equation, pV = nRT, what are the units of n and T ?
100
relative 60
0T
abundance
0
40 °C
no units%
n
A
energy
B
no units
K
Which curve applies for the20same gas at a lower temperature?
C
mol
°C
A
0 K
D
mol
90
91
2
B
93
92
94
m/e
proportion
3 The reaction between acidified dichromate(VI) proportion
ions, Cr2O72–, and aqueous Fe2+ ions results in the
Section
A
of dichromate(VI) ions being reduced to Cr3+ ions.of
What is the relative atomic mass of X?
molecules
molecules
For What
each isquestion
there
are
four
possible
answers,
A, B, C, and D. Choose the one you consider to
the correct equation for this reaction?
A 91.00
be correct.
91.30
+
C
3+
91.75
D
92.00
3+
[W’09 1 Q4]
Cr
0 2O7 + Fe + 14H → 2Cr + Fe + 7H2O 0
0 2–
0
2+
3+
B
Cr2mol
O7He,
+energy
14H+element
→
2Cr3+in +the
2Fe
+ 7H
of+isa2Fe
hydrocarbon
undergo
complete
combustion
to giveenergy
35.2 g of carbon dioxide and
0.200
2O
Helium,
the
second
Periodic
Table.
A
21
26
B
2+
2–
water as the only+products.
14.4 g of 2–
→ 2Cr33+ + 3Fe3+ + 7H2O
Cr2O7 + 3Fe2+ + 14H
C
D
Tritium is the isotope of hydrogen H.
2+
+
3+ hydrocarbon?
3+
2– molecular
What
is
the
formula
of
the
D Cr2O7 + 6Fe + 14H → 2Cr + 6Fe + 7H2O
What is the same in an atom of 4He and an atom of 3H?
B C2 H6
C C 4H 4
D C 4H 8
A C 2 H4
proportion
proportion
4 Sodium
azide,
NaN
is
an
explosive
used
to
inflate
airbags
in
cars
when
they crash. It consists of
A the number of3 electrons
C
of positive sodium ions and negative azide ions. of
molecules
molecules
B the number of neutrons
What are the numbers of electrons in the sodium ion and the azide ion?
C
D
the number of protons
0
A
27
2
2
sodium ion
10
0
energy
B
10 correctly shows
22
Which
diagram
the bonding in the ammonium ion, NH4+?
In which species are the numbers of protons, neutrons and electrons all different?
C
12
20
−
+
key
B 2+
B 199 F A
C 23
D 24
A 115B
12 Mg
11Na
D
12
22
N electron
28
5
0
azide ion
energy
20
0 relative atomic mass
the
+
H
H electron
+
H
[S’11 2 Q1]
[S’12 2 Q2]
it undergoes a H
natural
process to give
The 68Ge isotope is medically
H Nuseful because
H
N radioactive
H
an isotope of a different element, 68X, which can be used to detect tumours. This transformation
of 68Ge occurs when an electron enters the
nucleus and changes a proton into a neutron.
© UCLES 2012
9701/12/M/J/12
H
H
68
Which statement about the composition of an atom of X is correct?
A
It has 4 electrons in its outer p orbitals.
B
It has 13 electrons in its outer shell.
C
It has 37 neutrons.
D
Its proton number is 32.
C
© UCLES 2009
H
H
N
H
D
+
9701/11/O/N/09
H
CEDAR
© UCLESCOLLEGE
2016
+
H
H
N
H
[S’12 1 Q2]
H
9701/11/O/N/16
ATOMIC STRUCTURE
WS 1
[Turn over
What is the order of their first ionisation energies?
reactants
enthalpy
C
+5
+3
0
–2
least
+5
+3
endothermic
D
29
2
+6 most
36
endothermic
+4
products
A
Fr
Ra
Rn
Use of the Data Booklet is relevant to this question.
progress of reaction
B
Fr
Rn
Ra
In which species are the numbers of protons,
2 neutrons and electrons all different?
C
Ra change could
Fr the diagramRn
Which
enthalpy
not apply to?
B 23
A 199 F
11 Na
D
Rn
A enthalpy of atomisationRa
−
31
+
C 15 P
Section
Fr A
D
32 2 −
16 S
[S’10 1 Q1]
For each
question of
there
are four possible answers, A, B, C, and D. Choose the one you consider to
B enthalpy
combustion
3
The
first
six
ionisation
energies of four elements are given.
be
correct.
30
5 In which species are the numbers of protons, neutrons and electrons all different?
1
C enthalpy of formation
−to be in Group 32
+
2 − the Periodic Table?
27 element is most likely
Which
IVSof
B 35
C 16
D 39
A
13 Al
17 Cl
19 K
D
enthalpy
of
neutralisation
The diagram below represents, for a given temperature, the Boltzmann distribution of the kinetic
[S’13 1 Q5]
−1 react slowly together.
energy of the molecules inionisation
a mixtureenergy
of two gases
/ kJ molthat
6
31
4
An experiment is set up to measure the rate of hydrolysis of ethyl ethanoate.
2nd
3rd to
5th
6th
is4th
marked.
The
energy
for the
reaction,
Ea,this
Use activation
of the1st
Data
Booklet
is relevant
question.
H + H29
O 540 CH
C2600
H5OH
CH3CO26C2940
3CO
2H + 16
A most
494
560
13
400
When
the reaction
is4ion-molecule
catalysed,
the 5reaction
rate of reaction
little.
The
common
in gas increases
clouds
ofathe
Universe is as shown.
The
is found
to be 7slow
in neutral
solution
B hydrolysis
1 450of
10 500aqueous
13 600
18 but
000it proceeds at a measurable
the catalysed
reaction?
What
will be736
the position
Ea for740
H2+(g) → acid.
H(g) + H3+(g)
H2(g)
rate when the solution is acidified
with +hydrochloric
C
1 090
2 350
4 610
6 220
37 800
47 000
+
What is
the function
of the hydrochloric
acid?
What
be the 2composition
of an H
ion?
37 480
D could
1 400
860
4 590
9 400
53 200
A
to dissolve the ethyl ethanoate
protons number
neutrons
electrons
of
to ensure thatmolecules
the reaction reaches equilibrium
B
A
C
2
B
D
1
1
to increase the reaction rate by catalytic action
2
1
2
to suppress ionisation of the ethanoic acid formed
kinetic
C
3
0
1
energy
D
3
0
2
A
32
2
5
Ea
B
C
D
[S’14 3 Q4]
The electrolysis of brine using the diaphragm cell is an important industrial process.
Use of the Data Booklet is relevant to this question.
What happens at the anode?
In some types of spectroscopy, it is important2 to know if ions are isoelectronic. This means that
they contain equal numbers of electrons. 9701/13/M/J/13
© UCLES 2013
A
Chloride ions are oxidised.
Which
ion is notgas
isoelectronic
with K+?
B Hydrogen
is produced.
Section A
For each question
there are four– possible answers,2–A, B, C, and D. Choose
the one you consider to
2+
C S
D Ti3+
A
C Ca
Hydroxide ionsBareCl
formed.
be correct.
[W’14 3 Q2]
–
D
3
33
1
–
The electrode reaction is 2Cl (aq) + 2e → Cl 2(g).
Which row shows properties of a ceramic material?
Use of the Data Booklet is relevant to this question.
point
conductivity
melting
In which
optionpoint
do all boiling
three particles
have the same electronic configuration and the same
K
/
K
of
solid
/
number
of
neutrons?
© UCLES 2013
9701/11/M/J/13
[Turn over
A A 15N3–
16
156
O2–
19 –
B B 18O2–
19
922–
C C 19F–
D D 22Ne
352
none
20
Ne 1380
good
20
2130
Ne
23
Na+ 2943
good
23
3125
Na
24
F
© UCLES 2014
[Turn over
9701/13/M/J/14
F
Mg2+3873
none
[S’15 2 Q1]
2
The shell of a chicken’s egg makes up 5% of the mass of an average egg. An average egg has a
mass of 50 g.
Assume the egg shell is pure calcium carbonate.
© UCLES
9701/13/O/N/14
How2014
many complete chicken’s egg shells
would be needed to neutralise 50 cm3 of 2.0 mol dm–3
CEDAR
COLLEGE
ATOMIC STRUCTURE WS 1
ethanoic acid?
A
1
B
2
C
3
D
4
[H = 11 H; D = 12 H; O = 168 O]
D–
A
B
H3O+
–
OD
37
C
D
OH–
D
He+
34
4 Which species contains the smallest number of electrons?
B3+
A
B
Be2+
H–
3
C
[M’16 2 Q4]
4 When nuclear reactions take place, the elements produced are different from the elements that
35
reacted. Nuclear equations, such as the one below, are used to represent the changes that
occur.
235
92 U
+
1
0n
→ 144
56 Ba +
89
36 Kr
+ 3 01 n
The nucleon (mass) number total is constant at 236 and the proton number total is constant
at 92.
In another nuclear reaction, uranium-238 is reacted with deuterium atoms, 21 H. An isotope of a
new element, J, is formed as well as two neutrons.
238
92 U
2
1H
+
→ J + 2 01 n
What is isotope J?
238
238
Np
A
B
240
Pu
C
240
Np
D
Pu
[S’16 1 Q4]
3
5 Dicarbon monoxide, C2O, is found in dust clouds in space. The structure of this molecule is
36
4 C=C=O.
The relative
atomic mass
of copper
is 63.5.electrons.
The molecule
contains
no unpaired
Which
chart
is apairs
correct
mass spectrum
thatinwould
lead toofthis
value?
How
many
lone
of electrons
are present
a molecule
C2O?
© UCLES 2016
A
1
B
9701/12/F/M/16
2 A
C
3
D
B
4
100%
6
abundance
abundance
A white powder is known to be a mixture of magnesium oxide and aluminium oxide.
50%
33%
100 cm3 of 2 mol dm–3 NaOH(aq) is just sufficient to cause the aluminium
17% oxide in x grams of the
mixture to dissolve.
61 62 63 64 65 66
61 62 63 64 65 66
–
–
+
3H
O
→
2Al
(OH)
.
The reaction occurring is Al 2O3 + 2OH
2
4
m/e
m/e
800 cm3 of 2 mol dm–3 HCl (aq) is just sufficient to cause all of the oxide in x grams of the mixture
to dissolve.
C
D
The reactions occurring are Al 2O3 + 6H+ → 2Al 3+ + 3H2O
abundance
abundance
74%
and MgO + 2H+ → Mg2+ + H2O. 33%
26%
33%
22%
12%
How many moles of each oxide are present in x grams of the mixture?
61
aluminium
oxide
62magnesium
63 64 65 66
oxide
m/e
61
62
63
64
65 66
m/e
[S’16 3 Q4]
A
5
0.05
0.25
Which
isolated
B
0.05 gaseous atom
0.50 has a total of five electrons occupying spherically shaped orbitals?
AC boron 0.10
0.25
0.10
BD fluorine
0.50
C
sodium
D
potassium
CEDAR COLLEGE
© UCLES 2016
6
9701/11/M/J/16
Carbon and silicon have the same outer electronic structure.
ATOMIC STRUCTURE WS 1
[Turn over
2
37
C
N–H covalent bonds are stronger than P–H covalent bonds.
D
There is one lone pair in each ammonia molecule but no lone pair in each phosphine
molecule.
38
Neutrons are passed through an electric field. The mass of one neutron relative to 121 the mass of
a 12C atom and any deflection in the electric field is recorded.
Which row is correct?
mass of
neutron
behaviour of beam of
neutrons in an electric field
A
0
deflected
B
1
deflected
C
0
not deflected
D
1
not deflected
[S’18 3 Q2]
3
The table refers to the electron distribution in the second shell of an atom with eight protons.
Which row is correct for this atom?
orbital shape
orbital shape
orbital type
number of
electrons
orbital type
number of
electrons
A
p
2
s
4
B
p
4
s
2
C
s
2
p
4
D
s
4
p
2
© UCLES 2018
CEDAR COLLEGE
9701/13/M/J/18
ATOMIC STRUCTURE WS 1
11
39
14
Section
B
SECTION
Section BB
For each of the questions in this section, one or more of the three numbered statements 1 to 3 may
be correct.
For each of the questions in this section, one or more of the three numbered statements 1 to 3 may
be correct.
Decide whether each of the statements is or is not correct (you may find it helpful to put a tick against
the Decide
statements
thateach
youofconsider
to be correct).
whether
the statements
is or is not correct (you may find it helpful to put a tick against
the statements that you consider to be correct).
The responses A to D should be selected on the basis of
The responses A to D should be selected on the basis of
11
A
B
C
D
D
Section C
B
1 only
2 and 3
11and
2
1, 21,and
3
1 only
2 and 3
and 2
2 and 3
is
only
are
are
areare
For each of the
questions in thisonly
section,
one
or
more
of
the
three
numbered
1 to 3 may
is statements
only are
only are
correct
correct
correct
be correct. correct
correct
correct
correct
correct
A
B
Decide whether each of the statements is or is not correct (you may find it helpful to put a tick against
No
other
combination
statements
is used
used
as
No
combination
ofof
statements
is
asaacorrect
correctresponse.
response.
the other
statements
that you
consider
to be
correct).
The responses A to D should be selected on the basis of
31 The diagram illustrates the energy changes of a set of reactions.
1 The isotope cobalt-60 ( 60 Co ) is used to destroy cancer cells in the human body.
31
27
_
A
C
D
HB= _134 kJ12
mol 1
Which statements aboutRan atom of cobalt-60 are Scorrect?
1 only
1 and 2 Section B 2 and 3
1, 2 and 3
1 It contains
is
only are
are 33 neutrons.only are
correct
correct
correct
For each of the
questions in this section,
one or more of thecorrect
three numbered statements
1 to 3 may
2 Its nucleus has a relative charge of 27+.
be correct.
1
H = +92
J molisotopes of cobalt.
3 It has a different number of neutrons from the atoms
of other
Decide
whether
each of of
thestatements
statements is
is not
(youresponse.
may find it helpful to put a tick against
No
other
combination
is or
used
ascorrect
a correct
the statements that you consider to be correct).
[S’04 Q31]
32
The conversion
graphite
has only
small
positive
value of !H.
The
A to
D of
should
be
selected
on athe
basis
of
31
The relative
molecular
mass
of a molecule
of_ 1 chlorine
is 72.
2 responses
_
H = 75 kJ mol
U
T
C (graphite)
" C (diamond)
!H = +2.1 kJ mol –1
A
B in this molecule C
D
Which properties
of the atoms
are the same?
Which of
theproduction
following statements
arediamonds
correct? using this reaction is very difficult.
However,
the
of
synthetic
1, 2 and
3
1 and
2
2 and 3
1 only
1 radius
are enthalpy change
only
onlyUare
is –1 .
The
forare
the transformation
⎯→ R is + 42 kJ mol
correct
correct
The enthalpy change for the transformation T ⎯→ S is endothermic.
1
Which
statements
help to correct
explain this?
2 nucleon
number
correct
2
13
relative
isotopicenergy
mass of the reaction is large.
The
activation
The enthalpy
change foristhe
transformation
⎯→ T is – 33 kJ mol–1 .
No other3combination
of statements
used
as a correct R
response.
2 An equilibrium exists between diamond and graphite.
32 3Which
molecules
are reactions
planar? can be made to occur readily.
Only
exothermic
[S’05 Q31]
3 Use of the Data Booklet is relevant to this question.
31
1
BCl3
The technetium–99 isotope (99Tc) is radioactive and has been found in lobsters and seaweed
to nuclear fuel
reprocessing
plants.of a catalyst are correct?
33 adjacent
Which
about
the properties
2 NHstatements
3
Which
statements
are correct
an atom
of 99Tc?
PH
13 A
catalyst
increases
theabout
average
kinetic
energy of the reacting particles.
3
1
2
ItAhas
13 more
neutronsthe
than
protons.
catalyst
increases
rate
of the reverse reaction.
o
ItAhas
protons.
33 23
Boron
is 43
a non-metallic
element
is placed
in Group III of the Periodic
the reaction.
catalyst
has no effect
on thewhich
enthalpy
changeabove
!H ofaluminium
Table. It forms a compound with nitrogen known as boron nitride which has a graphite structure.
3
It has 99 nucleons.
Which of the following conclusions can be
drawn from this information?
9701/01/O/N/04
[S’07 Q31]
© UCLES 2004
32 Which of the following solids contain more than one type of chemical bond?
1
The empirical formula of boron nitride is BN.
1
brass (an alloy of copper and zinc)
2
The boron and nitride atoms are likely to be arranged alternately in a hexagonal pattern.
9701/01/M/J/04
[Turn over
! UCLES 2004
2 graphite
3 COLLEGE
Boron nitride has a layer structure with van der Waals’ forcesATOMIC
betweenSTRUCTURE
the layers. WS 1
CEDAR
3 ice
correct
A
correct
B
correct
C
correct
D
1, 2 and 3
1 and 2
2 and 3
No other combination
of statements
are
only is
areused as a correct
are
40onlyresponse.
correct
correct
correct
1 only
is
correct
10 is 12 the relative molecular mass of a particular
31
4 On a scale in which the mass of a 12C atom
No other
combination
is used as a correct response.
sample
of chlorineofisstatements
72.
Section B
Which properties of the atoms in this sample are always the same?
31 each
How may
existininthis
compounds?
For
of thenitrogen
questions
section, one or more of the three numbered statements 1 to 3 may
1
radius
be correct.
1 bonded by a triple covalent bond
2 nucleon number
Decide
of the statements is or is not correct (you may find it helpful to put a tick against
2 whether
as part each
of a cation
the statements
that
you consider to be correct).
3 isotopic
mass
3 in an oxidation state of +5
[S’09 1 Q31]
The responses A to D should be selected on the basis of
12 2–
32 An ideal gas obeys the
gas laws under all conditions
of temperature and pressure.
31 3–
32 13
5 The phosphide
P andBsulfide ion 16
S have
32
ion 15
A
C the same number
D of which sub-atomic
Which
of the following are true for an ideal
gas? B
particles?
Section
Section
B3
1, 2 and 3
1 and 2
2 and
1 only
1
The
molecules
have
negligible
volume.
arequestions in thisonly
are one or more
only
For each neutrons
of the
section,
of are
the three numberedisstatements 1 to 3 may
Forcorrect.
each ofcorrect
the questions in this
section, one or more
of the three numbered
correct
correct
correct statements 1 to 3 may
be
electrons
22
There
are no forces of attraction between molecules.
be correct.
Decide
each of the
statements
is kinetic
or is not
correct
(youismay
find it helpful
put a tick against
3 whether
protons
The
molecules
an average
energy
which
proportional
to itstoabsolute
No other
combination
of have
statements
is used
response.
Decide
whether
each
the statements
is as
or aiscorrect
not correct
(you may find it helpful to put a tick against
the
statements
that
youof
consider
to be correct).
temperature.
[S’12 2 Q32]
the statements that you consider to be correct).
The
responses
A to D should
selected
on the basis of
1 substances
33
Which
have
abe
giant
structure?
H3+ ionAwas
characterised
by J.
Thomson
over a century ago. 6Li is a rare isotope of
31
6 The
o basis of
The
responses
to Dfirst
should
selected
onJ.∆H
the
6 +be
represent
both a standard enthalpy change of
33 lithium
For which
reactions
does
the
value
of
which forms the Li ion.
1 calcium
oxidea standard enthalpy
combustion
change of formation?
A and
B
C
D
A
C
D
Which
statements
are correct? B
2
calcium
1, 2 +
and
3 → CO2(g) 1 and 2
2 and 3
1 only
1 C(s)
O2(g)
only
are
only2 are
is 1 only
1, 2are
and
3found in
1 and
2
and 3
13 Both
ions
thecrockery
same
number
of protons.
baked
claycontain
2 2C(s)
+ O2(g) → 2CO(g)correct
correct
correct
correctis
are
only are
only are
2 Both
ions 1contain the same
number of electrons.correct
correct
correct
correct
3 CO(g)
+ 2 O2(g) → CO2(g)
No other
combination
of statements
used as
correct response.
3 Both
ions contain
the sameisnumber
ofaneutrons.
[S’04 3 Q31]
No other combination of statements is used as a correct response.
31
2011
an international
group
of scientists
agreed
to add two
new elements
to temperature.
the Periodic
32
The
diagram
represents the
Boltzmann
distribution
of molecular
energies
at a given
7 In
Table. Both elements had been made artificially and were called ununquadium (Uuq)
and
31 ununhexium
In a solution(Uuh).
that contains both Br2 and Cl 2 , a process takes place that produces BrO3 – ions.
The process is represented by the following equations.
9701/01/M/J/09
Uuq
© UCLES 2009
© UCLES 2012
Uuh
HBrO
equation 1
Brproton
2 + Hnumber
2 O → HBr + 114
116
number
of molecules
equation
2
3HBrO
Cl 2 9701/13/O/N/12
→ 2Cl – 289
+ BrO3 – 292
+ Br2 + 3H+
nucleon+ number
[Turn over
Whichstatements
statementsabout
aboutthese
these
reactions
correct?
Which
elements
areare
correct?
11
One
atom is
ofreduced
Uuh has in
one
more neutron
than one atom of Uuq.
Chlorine
equation
2.
22
has theinsame
of 1electrons
as one2.atom of Uuh.
One
Uuq2–ision
Bromine
oxidised
both number
equation
and energy
equation
ionreduced
has the in
same
number
of electrons
as one2.
Uuq– ion.
33 One
Uuh+ is
Bromine
boththe
equation
equation
Which
of the factors
that affect
rate of 1a and
reaction
can be explained using such a Boltzmann
[S’14 3 Q31]
distribution?
32 P and Q are two liquid compounds with similar Mr values. Molecules of P attract each other by
328 1
Which
statements
are correct when
referring to the isotopes of a single element?
increasing
the concentration
of reactants
hydrogen bonds. Molecules of Q attract each other by van der Waals’ forces only.
2
the have
temperature
1 increasing
The isotopes
different masses.
How do the properties of P and Q differ?
3
the
of have
a catalyst
2
Theaddition
isotopes
different numbers of nucleons.
1 P has higher surface tension than Q.
3
The isotopes have different chemical reactions.
2
P has a higher boiling point than Q.
3
P is less viscous than Q.
© UCLES 2013
CEDAR COLLEGE
[S’14 3 Q32]
9701/13/O/N/13
ATOMIC STRUCTURE WS 1
correct
correct
correct
correct
Decide whether each of the statements is or is not correct (you may find it helpful to put a tick against
the
statements
that you
to be
correct).
No other
combination
of consider
statements
is used
as a correct response.
41
The responses A to D should be selected on the basis of
319 Use of the Data Booklet is relevant to this question.
A
B
C
D
Which statements about the phosphide ion,1431P3–, and the chloride ion, 35Cl –, are correct?
1, 2 and 3
1 and 2
2 and 3
1 only
only are
is
1 Theyare
have the same number
of electrons.
Section B only are
correct
correct
correct
correct
the same
number
of one
neutrons.
10 of the three numbered statements 1 to 3 may
For 2eachThey
of thehave
questions
in this
section,
or more
be correct.
3 They
have theofsame
numberisof
protons.
No other
combination
statements
used
as a correct
response.
Section
B
[S’15 2 Q31]
Decide whether each of the statements is or is not correct (you may find it helpful to put a tick against
the each
statements
you consider
be correct).
For
of the that
questions
in this to
section,
one or more of the three numbered statements 1 to 3 may
10 correct.
, sublime
at the relatively low temperature of 180 °C?
32
Why
aluminium
31
X
is adoes
particle
with 18 chloride,
electronsAland
neutrons.
2Cl 620
be
The responses A to D should be selected on the basis of
1 whether
The
intermolecular
forces
between
2Cl 6 molecules
What
could
be the
of X?
Decide
each
of symbol
the statements
is or isthe
notAlcorrect
(you may are
find weak.
it helpful to put a tick against
the statementsAthat you consider toBbe correct).
C
D
2 The
co-ordinate bonds between aluminium and chlorine are weak.
38
1
18 Ar
1, 2 and
1selected
and 2 on the basis
2 and
3
The responses
A to3 D should
ofchlorine
3 The
covalent
bonds be
between
aluminium and
are weak.1 only
are
only are
only are
is
40
2+
2
20 Ca
correct
correct
correct
correct
A
B
C
D
39 statements
+
33 3Which
are
correct
for
all
exothermic
reactions?
1,K2 and 3 of statements
1 and
2
2 and
3
1 only
No other 19
combination
is used
as a correct
response.
are
only
are
only
are
is
[S’16 1 Q31]
1 ∆H for the reaction is negative.
correct
correct
correct
correct
31 What
of boron
contains diagram
aluminium
the
onlyare
impurity.
spectrum
of the mixture
11
2A sample
On
reaction
pathway
theasproducts
shownAlower
than
the
reactants.
32
area basic
assumptions
of the kinetic
theory
as
applied
tomass
an ideal
gas?
shows three lines corresponding to three ions, X+, Y+ and Z+.
No other
combination
statements
is used as a correct response.
3 The
reactionofwill
happen spontaneously.
1
Gas particles are in continuous random motion.
Y+
Z+
X+
Use of the Data Booklet may be ion
appropriate for some
questions.
2
Gas particles experience no intermolecular forces.
3
The volume of each gas particle is zero.
m/e
10
11
27
31 For complete combustion, 1 percentage
mol of an organic
compound
requires 2.5 mol of O2.
15.52
74.48 X10.00
abundance
compounds could be X?
33 Which
Bromine
reacts with water.
Which statements are correct?
1 C2H5OH
HOBr + HBr
Br2 + H2O
in X+.
1 There are more electrons in Z+ than
2 C2H2
Which
oxidation
states
ofsample
bromine
are present
in the equilibrium
2 The
Ar of boron
in the
is 10.83
to four significant
figures. mixture?
3 CH3CHO
There are more protons in Y+ than in X+.
13 +3
32
2 which
0 pairs do both species have the same number of electrons?
12 In
32 Which
elements can form bonds in their compounds?
Cl and 37Cl
13 35
–1
1 carbon
2 35Cl – and 40Ar
2 oxygen
3 40Ar and 40K+
3 nitrogen
© UCLES 2015
9701/12/M/J/15
[Turn over
[S’18 2 Q32]
33 For which reactions does the value of H o represent both a standard enthalpy change of
33 combustion
For which enthalpy
changesenthalpy
is the value
of Hofalways
negative?
and a standard
change
formation?
combustion
© UCLES
11 2016
C(s)
+ O (g)
2
22
3
3
→ CO2(g)
9701/11/M/J/16
[Turn over
hydration
2C(s)
+ O2(g) → 2CO(g)
solution
CO(g) + 21 O2(g) → CO2(g)
© UCLES 2016
CEDAR COLLEGE
9701/11/O/N/16
ATOMIC STRUCTURE WS 1
42
ATOMIC STRUCTURE WS 2
1 Give the numbers of protons, neutrons and electrons present in each of the following atoms:
a) 40Ar
c) 197Au+
b) 127I
d) 52Cr3+
2 This question concerns the following five species:
16O2--
19F
20Ne
23Na
25Mg2+
a) Which two species have the same number of neutrons? ........................................
b) Which two species have the same ratio of neutrons to protons? ........................................
c) Which two species do not have 10 electrons? ........................................
3 The element Rhenium (Re) has two main isotopes, 185Re with an abundance of 37.1% and
187Re with an abundance of 62.9%.
Calculate the weighted mean atomic mass of rhenium.
4 Antimony has two main isotopes, 121Sb and 123Sb. A forensic scientist was asked to help a
crime investigation by analysing the antimony in a bullet. This was found to contain 57.3% of
121Sb and 42.7% of 123Sb.
a) Calculate the relative atomic mass of the sample of antimony from the bullet. (Write your
answer to three significant figures)
b) State one similarity and one difference between isotopes in terms of subatomic particles.
...............................................................................................................................................................................
...............................................................................................................................................................................
...............................................................................................................................................................................
CEDAR COLLEGE
ATOMIC STRUCTURE WS 2
43
5
Bromine exists as a molecule with two bromine atoms combined together. Bromine has two
isotopes: bromine-79 and bromine-81
a) A molecule of bromine containing two atoms of bromine can be written as 79Br2. Write the
formulae for the two other possible molecules of bromine.
...............................................................................................................................................................................
b) The mass spectrum of molecules of bromine is shown below:
i. Explain why these peaks are observed.
...............................................................................................................................................................................
...............................................................................................................................................................................
...............................................................................................................................................................................
ii. The peaks at 79 and 81 are the same height. What does this tell you about the relative
abundances of the two isotopes?
...............................................................................................................................................................................
...............................................................................................................................................................................
iii. Explain why the peak at 160 is twice the height of the peaks at 158 and 162.
...............................................................................................................................................................................
...............................................................................................................................................................................
...............................................................................................................................................................................
6
a) Explain why the relative atomic mass of copper is not an exact whole number?
...............................................................................................................................................................................
...............................................................................................................................................................................
...............................................................................................................................................................................
CEDAR COLLEGE
ATOMIC STRUCTURE WS 2
44
b) The relative atomic mass of copper is 63.5. Calculate the relative abundance of the two
copper isotopes with the relative isotopic masses of 63.0 and 65.0.
7
Chlorine exists as a molecule with two chlorine atoms combined together. Chlorine has two
isotopes: chlorine-35 and chlorine-37.
The mass spectrum of chlorine os shown below:
a) The peak at 35 is three times as high as the peak at 37. Calculate the relative atomic mass
of chlorine.
...............................................................................................................................................................................
...............................................................................................................................................................................
b) Explain why the peaks are observed at 70, 72 and 74.
...............................................................................................................................................................................
...............................................................................................................................................................................
c) The heights of the peaks at 70, 72 and 74 are in the ratio 9 :6 :1. Explain why the heights
are in this ratio.
...............................................................................................................................................................................
...............................................................................................................................................................................
...............................................................................................................................................................................
CEDAR COLLEGE
ATOMIC STRUCTURE WS 2
copper isotopes with relative isotopic masses
of 63.0 and 65.0.
45
8
3 The data about silicon in the table below were
The data about silicon in obtained
the table below
obtained
from a mass spectrometer.
fromwere
a mass
spectrometer.
10 List the p
increasin
11 Explain w
noble ga
m/z
% abundance
28
92.2
12 The succ
X, are giv
periodic
29
4.7
Ionisatio
30
3.1
1st
2nd
Calculate
relative
Calculate the relative atomic
mass ofthe
silicon
to oneatomic
decimalmass
place.of silicon to
3rd
one decimal place.
2
19
For
For
4th
Examiner’s
Examiner’s
Use
Use
5th
4 Bromine has two isotopes of mass numbers
79 and 81. The mass spectrum of a sample of
bromine is shown below.
..........................................................................................................................................
(a) Define an isotope in terms of its sub-atomic particles.
6th
7th
Relative
..........................................................................................................................................
intensity
[1]
8th
9th
10th
(b) In a mass spectrometer some hydrogen chloride molecules will split into atoms. The
mass spectrum of HCl is given. Chlorine has two isotopes. The hydrogen involved here
is the isotope 11H only.
13 Discuss t
of ESI an
100
80
60
158
160
162
m/z
relative
a) Identify the particles responsible for the
abundance
40
peaks.
14 By consi
and d ele
why the
of fluorin
between
20
0
35
m /e
40
807404_C02_Edexcel_GF_Chem_009-036.indd 35
(i) What particle is responsible for the peak at mass 35? .............................................
(ii) What particle is responsible for the peak at mass 38? .............................................
[2]
(c) Use the relative heights of the peaks to determine the proportions of the two isotopes of
chlorine. Explain simply how you obtained your answer.
[2]
[2]
(d) Use
Use your
your answer
answer to
to (c)
(c) to
to explain
explain why
why chlorine
(d)
chlorine has
has a
a relative
relative atomic
atomic mass
mass of
of 35.5.
35.5.
CEDAR COLLEGE
ATOMIC STRUCTURE WS 2
46
[2]
(d) Use your answer to (c) to explain why chlorine has a relative atomic mass of 35.5.
[1]
2
2
Answer
in the
the space
space provided.
provided.
Answer all
all the
the questions
questions in
[S’03:Q1]
[Total
6]
9701/2/M/J/03
1
1
10
Iron
Table. Iron
Iron has
has three
three main
main naturally
naturally
Iron and
and cobalt
cobalt are
are adjacent
adjacent elements
elements in
in the
the Periodic
Periodic Table.
occurring isotopes, cobalt has one.
For
For
Examiner’s
Examiner’s
Use
Use
(a) Explain the meaning of the term isotope.
..........................................................................................................................................
..........................................................................................................................................
.................................................................................................................................... [2]
56Fe; the only naturally occurring isotope of cobalt
(b) The most common isotope of iron is 56
59
59
is Co.
56
Use the Data Booklet to complete the table below to show the atomic structure of 56
Fe
59
and of 59Co.
number of
isotope
protons
neutrons
electrons
56Fe
59Co
[3]
(c) A sample of iron has the following isotopic composition by mass.
isotope mass
54
56
57
% by mass
5.84
91.68
2.17
(i) Define the term relative atomic mass.
..................................................................................................................................
..................................................................................................................................
..................................................................................................................................
..................................................................................................................................
(ii) By using the data above, calculate the relative atomic mass of iron to three
significant figures.
CEDAR COLLEGE
ATOMIC STRUCTURE WS 2
..................................................................................................................................
47
..................................................................................................................................
(ii) By using the data above, calculate the relative atomic mass of iron to three
significant figures.
[5]
[Total:
[S’05 10]
Q1]
© UCLES 2005
CEDAR COLLEGE
9701/02/M/J/05
ATOMIC STRUCTURE WS 2
2
Answer all the questions
48in the spaces provided.
1
11
For
Examiner’s
Use
In the 19th and 20th centuries, scientists established the atomic theory and showed that
three sub-atomic particles, electron, neutron and proton, exist. The masses and charges of
these three particles were subsequently determined.
When separate beams of electrons, neutrons or protons are passed through an electric field
in the apparatus below, they behave differently.
+
–
beam of particles
(a) (i)
Which of these three particles will be deflected the most by the electric field?
.........................................
(ii)
In which direction will this particle be deflected?
..................................................................................................................................
(iii)
Explain your answer.
..................................................................................................................................
(b) (i)
..................................................................................................................................
[4]
Define the term proton number.
..................................................................................................................................
..................................................................................................................................
(ii)
Why is the proton number of an atom of an element usually different from the
nucleon number of an atom of the element?
..................................................................................................................................
..................................................................................................................................
[2]
© UCLES 2006
CEDAR COLLEGE
9701/02/O/N/06
ATOMIC STRUCTURE WS 2
49
3
(c) Protons and neutrons have been used in nuclear reactions which result in the formation
of artificial elements. In such processes, protons or neutrons are accelerated to high
speeds and then fired like ‘bullets’ at the nucleus of an atom of an element.
For
Examiner’s
Use
Suggest why neutrons are more effective than protons as ‘nuclear bullets’.
..........................................................................................................................................
..................................................................................................................................... [2]
(d) In some cases, when neutrons are fired at atoms of an element, the neutrons become
part of the nucleus of those atoms.
What effect does the presence of an extra neutron have on the chemical properties of
the new atoms formed? Explain your answer.
..........................................................................................................................................
..........................................................................................................................................
..................................................................................................................................... [2]
[Total: 10]
[W’06 Q1]
© UCLES 2006
CEDAR COLLEGE
9701/02/O/N/06
[Turn over
ATOMIC STRUCTURE WS 2
2
50
Answer all the questions in the spaces provided.
1
12
For
Examiner’s
Use
Magnesium, Mg, and radium, Ra, are elements in Group II of the Periodic Table.
Magnesium has three isotopes.
(a) Explain the meaning of the term isotope.
..........................................................................................................................................
..........................................................................................................................................
..................................................................................................................................... [2]
A sample of magnesium has the following isotopic composition by mass.
isotope mass
24
25
26
% by mass
78.60
10.11
11.29
(b) Calculate the relative atomic mass, Ar, of magnesium to four significant figures.
3
Ar = ………………
[2]
Radium, proton number 88, and uranium, proton number 92, are radioactive elements.
For
Examiner’s
Use
The isotope 226Ra is produced by the radioactive decay of the uranium isotope 238U.
(c) Complete the table below to show the atomic structures of the isotopes 226Ra and
238U.
number of
isotopes
protons
neutrons electrons
226Ra
238U
[3]
(d) Radium, like other Group II elements, forms a number of ionic compounds.
(i)
© UCLES 2009
What is the formula of the radium cation?
………………
9701/21/O/N/09
(ii)
Use the Data Booklet to suggest a value for the energy required to form one mole of
the gaseous radium cation you have given in (i) from one mole of gaseous radium
atoms. Explain your answer.
CEDAR COLLEGE
ATOMIC STRUCTURE WS 2
..................................................................................................................................
226Ra
238U
51
[3]
(d) Radium, like other Group II elements, forms a number of ionic compounds.
(i)
What is the formula of the radium cation?
………………
(ii)
Use the Data Booklet to suggest a value for the energy required to form one mole of
the gaseous radium cation you have given in (i) from one mole of gaseous radium
atoms. Explain your answer.
..................................................................................................................................
..................................................................................................................................
.............................................................................................................................
[3]
2
Answer all the questions in the space provided.
13
1
[W’09
Q1]
[Total:110]
The element magnesium, Mg, proton number 12, is a metal which is used in many alloys
which are strong and light.
For
Examiner’s
Use
Magnesium has several naturally occurring isotopes.
(a) What is meant by the term isotope?
..........................................................................................................................................
..........................................................................................................................................
...................................................................................................................................... [2]
(b) Complete the table below for two of the isotopes of magnesium.
isotope
number of
protons
number of
neutrons
number of
electrons
24Mg
26Mg
[2]
© UCLES 2009
[Turn over
9701/21/O/N/09
A sample of magnesium had the following isotopic composition:
24Mg, 78.60%; 25Mg, 10.11%; 26Mg, 11.29%.
(c) Calculate the relative atomic mass, Ar, of magnesium in the sample.
Express your answer to an appropriate number of significant figures.
[2]
CEDAR COLLEGE
ATOMIC STRUCTURE WS 2
..........................................................................................................................................
52
...................................................................................................................................... [2]
(b) Complete the table below for two of the isotopes of magnesium.
isotope
number of
protons
number of
neutrons
number of
electrons
24Mg
26Mg
[2]
A sample of magnesium had the following isotopic composition:
24Mg, 78.60%; 25Mg, 10.11%; 26Mg, 11.29%.
(c) Calculate the relative atomic mass, Ar, of magnesium in the sample.
Express your answer to an appropriate number of significant figures.
2
[2]
Answer all the questions in the space provided.
14
1
[W’10 3 Q1]
Sulfur, S, and polonium, Po, are both elements in Group VI of the Periodic Table.
For
Examiner’s
Use
Sulfur has three isotopes.
(a) Explain the meaning of the term isotope.
..........................................................................................................................................
..........................................................................................................................................
...................................................................................................................................... [2]
(b) A sample of sulfur has the following isotopic composition by mass.
© UCLES 2010
9701/23/O/N/10
isotope mass
32
33
34
% by mass
95.00
0.77
4.23
Calculate the relative atomic mass, Ar, of sulfur to two decimal places.
Ar = ...............
CEDAR COLLEGE
[2]
ATOMIC STRUCTURE WS 2
(c) Isotopes of polonium, proton number 84, are produced by the radioactive decay of
several elements including thorium, Th, proton number 90.
..........................................................................................................................................
..........................................................................................................................................
..........................................................................................................................................
53
......................................................................................................................................
[2]
...................................................................................................................................... [2]
(b)
following isotopic
isotopic composition
composition by
by mass.
mass.
(b) A
A sample
sample of
of sulfur
sulfur has
has the
the following
isotope
mass
isotope mass
32
32
33
33
34
34
% by mass
mass
95.00
95.00
0.77
0.77
4.23
4.23
places.
Calculate the relative atomic mass, A r, of sulfur to two decimal places.
r
Arr = ...............
...............
[2]
[2]
(c) Isotopes of polonium, proton number 84, are produced by the radioactive
radioactive decay
decay of
of
several elements including thorium, Th, proton number 90.
232
The isotope 213Po is produced from the thorium isotope 232
Th.
213Po and 232
Th.
Complete the table below to show the atomic structures of the isotopes
isotopes 213
Po and 232Th.
number of
isotope
protons
neutrons
electrons
213Po
232Th
[3]
[3]
© UCLES 2011
©
CEDAR COLLEGE
9701/23/O/N/11
ATOMIC STRUCTURE WS 2
54
3
Radiochemical reactions, such as nuclear fission and radioactive decay of isotopes, can be
represented by equations in which the nucleon (mass) numbers must balance and the proton
numbers must also balance.
For
Examiner’s
Use
1
For example, the nuclear fission of uranium-235, 235
92 U, by collision with a neutron, 0n,
produces strontium-90, xenon-143 and three neutrons.
1
90
143
38Sr + 54 Xe + 3 0n
235
1
92 U + 0n
In this equation, the nucleon (mass) numbers balance because: 235 + 1 = 90 + 143 + (3x1).
The proton numbers also balance because:
92 + 0 = 38 + 54 + (3x0).
(d) In the first stage of the radioactive decay of
element E and an alpha-particle, 42 He.
(i)
232
90 Th, the products are an isotope of
By considering nucleon and proton numbers only, construct a balanced equation
for the formation of the isotope of E in this reaction.
232
90 Th
................. + 42 He
Show clearly the nucleon number and proton number of the isotope of E.
nucleon number of the isotope of E ...........
proton number of the isotope of E .............
(ii)
Hence state the symbol of the element E.
2
............ ................................................................................................................. [3]
Answer all the questions in the spaces provided.
[Total:
[W’11 310]
Q1]
15
1
(a) Explain what is meant by the term nucleon number.
....................................................................................................................................................
.............................................................................................................................................. [1]
(b) Bromine exists naturally as a mixture of two stable isotopes, 79Br and 81Br, with relative isotopic
masses of 78.92 and 80.92 respectively.
(i) Define the term relative isotopic mass.
.............................................................................................................................................
.............................................................................................................................................
....................................................................................................................................... [2]
(ii) Using the relative atomic mass of bromine, 79.90, calculate the relative isotopic abundances
of 79Br and 81Br.
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ATOMIC STRUCTURE WS 2
.............................................................................................................................................
.............................................................................................................................................
55
2
.......................................................................................................................................
[2]
all the mass
questions
in the 79.90,
spacescalculate
provided.
(ii) Using theAnswer
relative atomic
of bromine,
the relative isotopic abundances
of 79Br and 81Br.
1
(a) Successive ionisation energies for the elements magnesium to barium are given in the table.
element
1st ionisation
2nd ionisation
3rd ionisation
energy / kJ mol–1 energy / kJ mol–1 energy / kJ mol–1
Mg
736
1450
7740
Ca
590
1150
4940
Sr
548
1060
4120
Ba
502
966
3390
[3]
(i) Explain why the first ionisation energies decrease down the group.
(c) Bromine reacts with the element A to form a compound with empirical formula ABr3. The
percentage composition by mass of ABr3 is A, 4.31; Br, 95.69.
.............................................................................................................................................
Calculate the relative atomic mass, Ar, of A.
.............................................................................................................................................
Give your answer to three significant figures.
.............................................................................................................................................
....................................................................................................................................... [3]
(ii) Explain why, for each element, there is a large increase between the 2nd and 3rd ionisation
energies.
.............................................................................................................................................
.............................................................................................................................................
Ar of A = ....................... [3]
....................................................................................................................................... [2]
[S’14 2 Q1]
16
(b) A sample of strontium, atomic number 38, gave the mass spectrum shown. The percentage
© UCLES
2014
abundances
are given above each peak.9701/22/M/J/14
100
82.58
percentage
abundance
0
9.86 7.00
0.56
80
81
82
83
84
85
86
87
88
89
90
atomic mass units
© UCLES 2014
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ATOMIC STRUCTURE WS 2
3
(i) Complete the full electronic configuration of strontium.
56
1s2 2s2 2p6 ................................................................................................................... [1]
(ii) Explain why there are four different peaks in the mass spectrum of strontium.
.............................................................................................................................................
....................................................................................................................................... [1]
(iii) Calculate the atomic mass, Ar, of this sample of strontium.
Give your answer to three significant figures.
Ar = ............................. [2]
[W’14 2 Q1]
(c) A compound of barium, A, is used in fireworks as an oxidising agent and to produce a green
colour.
(i) Explain, in terms of electron transfer, what is meant by the term oxidising agent.
.............................................................................................................................................
....................................................................................................................................... [1]
(ii) A has the following percentage composition by mass: Ba, 45.1; Cl , 23.4; O, 31.5.
Calculate the empirical formula of A.
empirical formula of A ........................................... [3]
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ATOMIC STRUCTURE WS 2
2
Answer all the questions
57in the spaces provided.
17
1
(a) Chemists recognise that atoms are made of three types of particle.
Complete the following table with their names and properties.
name of particle
relative mass
relative charge
0
1/1836
[3]
(b) The relative atomic mass of an element can be determined using data from its mass spectrum.
The mass spectrum of element X is shown, with the percentage abundance of each isotope
labelled.
60
49.61
50
40
percentage
abundance 30
23.77
20
10
0
9.37
7.63
76
77
8.73
0.89
73
74
75
78
m/e
79
80
81
82
83
(i) Define the terms relative atomic mass and isotope.
relative atomic mass ...........................................................................................................
.............................................................................................................................................
.............................................................................................................................................
isotope ................................................................................................................................
.............................................................................................................................................
[3]
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ATOMIC STRUCTURE WS 2
58
3
(ii) Use the data in the mass spectrum to calculate the relative atomic mass, Ar, of X.
Give your answer to two decimal places and suggest the identity of X.
Ar of X ....................................
identity of X ....................................
[2]
(c) The element tellurium, Te, reacts with chlorine to form a single solid product, with a relative
formula mass of 270. The product contains 52.6% chlorine by mass.
(i) Calculate the molecular formula of this chloride.
molecular formula .................................... [3]
(ii) This chloride melts at 224 °C and reacts vigorously with water.
[M’15 1 Q1]
State the type of bonding and structure present in this chloride and explain your reasoning.
.............................................................................................................................................
.............................................................................................................................................
.............................................................................................................................................
....................................................................................................................................... [2]
(iii) Suggest an equation for the reaction of this chloride with water.
....................................................................................................................................... [1]
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ATOMIC STRUCTURE WS 2
2
Answer all the questions
59in the spaces provided.
18
1 (a) Chemists recognise that atoms are made of three types of particle.
Complete the following table with their names and properties.
name of particle
relative mass
relative charge
+1
1/1836
[3]
(b) Most elements exist naturally as a mixture of isotopes, each with their own relative isotopic
mass. The mass spectrum of an element reveals the abundances of these isotopes, which can
be used to calculate the relative atomic mass of the element.
Magnesium has three stable isotopes. Information about two of these isotopes is given.
isotope
relative
isotopic mass
percentage
abundance
24
Mg
24.0
79.0
26
Mg
26.0
11.0
(i) Define the term relative isotopic mass.
.............................................................................................................................................
.............................................................................................................................................
....................................................................................................................................... [2]
(ii) The relative atomic mass of magnesium is 24.3.
Calculate the percentage abundance and hence the relative isotopic mass of the third
isotope of magnesium. Give your answer to three significant figures
percentage abundance = .................................
isotopic mass = .................................
[3]
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ATOMIC STRUCTURE WS 2
.............................................................................................................................................
.......................................................................................................................................
[2]
60
(c) Neon has three stable isotopes.
isotope
mass number
percentage abundance
1
9.25
2
20
90.48
3
21
0.27
(i) Define the term relative atomic mass.
.............................................................................................................................................
....................................................................................................................................... [2]
(ii) Use the relative atomic mass of neon, 20.2, to calculate the mass number of isotope 1.
mass number = ................................. [2]
[S’15 3 Q1]
© UCLES 2015
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ATOMIC STRUCTURE WS 2
2
Answer all the questions
61 in the spaces provided.
20
1
(a) Complete the table to show the composition and identity of some atoms and ions.
name of
element
nucleon
number
atomic
number
lithium
6
3
.............
.............
.............
54
26
26
.............
.............
oxygen
.............
.............
number of
protons
number of
neutrons
number of
electrons
.............
.............
.............
9
10
17
.............
18
overall
charge
+1
.............
24
.............
.............
0
[4]
(b)
eams of protons neutrons and electrons beha e differentl in an electric field due to their
differing properties.
The diagram sho s the path of a beam of electrons in an electric field
dd and label lines to represent the paths of beams of protons and neutrons in the same field
electron beam
[3]
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ATOMIC STRUCTURE WS 2
62
4
(d) A sample of strontium exists as a mixture of four isotopes. Information about three of these
isotopes is given in the table.
mass number
86
87
88
abundance
9.86%
7.00%
82.58%
(i) Calculate the abundance of the fourth isotope.
abundance = ............................. % [1]
(ii) The relative atomic mass of this sample of strontium is 87.71.
Calculate the mass number of the fourth isotope.
mass number = ............................. [2]
[Total:
[S’16 116]
Q1]
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ATOMIC STRUCTURE WS 2
3
63
21
(c) A sample of oxygen exists as a mixture of three isotopes. Information about two of these
isotopes is given in the table.
mass number
16
17
abundance
99.76%
0.04%
(i) Calculate the abundance of the third isotope.
abundance = ............................. % [1]
(ii) The relative atomic mass of this sample of oxygen is 16.0044.
Calculate the mass number of the third isotope. You
st show your working.
mass number = ............................. [2]
[S’16 211]
Q1]
[Total:
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ATOMIC STRUCTURE WS 2
5
64
22 A naturally occurring sample of cerium contains only four isotopes. Data for three of the
(iii)
isotopes are shown in the table.
isotope
136
Ce
138
Ce
140
Ce
142
Ce
relative
isotopic mass
135.907
137.906
139.905
to be
calculated
percentage
abundance
0.185
0.251
88.450
to be
calculated
The Ar of the sample is 140.116.
Use these data to calculate the relative isotopic mass of the fourth isotope in this sample
of cerium.
Give your answer to three decimal places.
relative isotopic mass = .............................. [3]
[M’1717]
Q1]
[Total:
CEDAR
COLLEGE
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ATOMIC STRUCTURE
WS 2
[Turn over
65
2 Atomic structure
2.3 Electrons: energy levels, atomic orbitals, ionisation energy
ELECTRONIC CONFIGURATION
BILAL HAMEED
ELECTRONIC CONFIGURATION
Cambridge International AS and A Level Chemistry
9701 syllabus
Syllabus
content
relative charges
and relative
masses
b) deduce the behaviour of beams of protons, neutrons and electrons in
electric fields
66of mass and charge within an atom
c) describe the distribution
2 Atomic structure
d) deduce the numbers of protons, neutrons and electrons present in both
atoms and ions given proton and nucleon numbers and charge
This topic describes the type, number and distribution of the fundamental particles which make up an
2.2
The
nucleus
of the
the contribution
atom
and
the impact
of thisa)on describe
some atomic
properties.of protons and neutrons to atomic nuclei in
atom
terms of proton number and nucleon number
b) distinguish between isotopes on the basis of different numbers of
Learning
outcomes
neutrons present
Candidates should be able to:
2.1 Particles in the atom
2.3 Electrons: energy
levels, atomic
orbitals, ionisation
energy, electron
affinity
c) recognise and use the symbolism xy A for isotopes, where x is the nucleon
number and y is the proton number
a) identify and describe protons, neutrons and electrons in terms of their
relativethe
charges
and
a) describe
number
andrelative
relativemasses
energies of the s, p and d orbitals for
b)thededuce
the
behaviour
of beams
of protons,
neutrons
and4p
electrons
principal
quantum
numbers
1, 2 and
3 and also
the 4s and
orbitals in
electricand
fields
b) describe
sketch the shapes of s and p orbitals
describe
the distribution
of mass
and charge
within
c) c)state
the electronic
configuration
of atoms
and ions
givenan
theatom
proton
2
2
6
number
and
charge,
using
the
convention
1s
2s
2p
,
etc.
d) deduce the numbers of protons, neutrons and electrons present in both
and
proton
and nucleon
d) (i) atoms
explain
andions
usegiven
the term
ionisation
energy numbers and charge
(ii) explain the factors influencing the ionisation energies of elements
2.2 The nucleus of the
atom
explain the
in ionisation
energies
a Period
and down
a in
a)(iii)describe
thetrends
contribution
of protons
andacross
neutrons
to atomic
nuclei
Group
of
the
Periodic
Table
(see
also
Section
9.1)
terms of proton number and nucleon number
e)b)deduce
the electronic
configurations
successive
distinguish
between
isotopes on of
theelements
basis offrom
different
numbers of
ionisation
energy
data
neutrons present
x
x
f) c)interpret
successive
energy data
of an element in terms
of the
recognise
and useionisation
the symbolism
y A for isotopes, where is the nucleon
position of that element within the Periodic Table
number and y is the proton number
g) explain and use the term electron affinity
2.3 Electrons: energy
levels, atomic
orbitals, ionisation
energy, electron
affinity
a) describe the number and relative energies of the s, p and d orbitals for
the principal quantum numbers 1, 2 and 3 and also the 4s and 4p orbitals
b) describe and sketch the shapes of s and p orbitals
c) state the electronic configuration of atoms and ions given the proton
number and charge, using the convention 1s22s22p6 , etc.
d) (i) explain and use the term ionisation energy
(ii) explain the factors influencing the ionisation energies of elements
(iii) explain the trends in ionisation energies across a Period and down a
Group of the Periodic Table (see also Section 9.1)
e) deduce the electronic configurations of elements from successive
ionisation energy data
18
www.cie.org.uk/alevel
f)
interpret successive ionisation energy data of an element in terms of the
Back to contents page
position of that element within the Periodic Table
g) explain and use the term electron affinity
18
www.cie.org.uk/alevel
BILAL HAMEED
Back to contents page
ELECTRONIC CONFIGURATION
67
DISCLAIMER
A complete discussion of the experimental evidence for the modern
theory of atomic structure is beyond the scope of the CIE A Level
Syllabus.
In this chapter only the results of the theoretical treatment will be
described. These results will have to be memorized as “rules of the
game,” but they will be used so extensively throughout the general
chemistry course that the notation used will soon become familiar.
1
ELECTRON ARRANGEMENT
The electronic configuration describes the arrangement of electrons in
atoms.
An atom’s electrons are arranged outside the nucleus in energy levels (or
shells).
Each shell or energy level holds a certain maximum number of electrons.
The energy of levels becomes greater as they go further from the
nucleus and electrons fill energy levels in order.
2
BILAL HAMEED
ELECTRONIC CONFIGURATION
68
IONISATION ENERGY
Ionisation energy is a measure of the energy needed to remove an
electron from a gaseous atom or ion. It measures how strongly an atom
or ion holds on to its electrons.
Attraction between the nucleus and an electron
+
The greater the pull of the nucleus, the harder it will be to pull an electron
away from an atom.
3
FIRST IONISATION ENERGY
Ionisation energies give evidence for the arrangement of electrons in
atoms in shells and sub-shells.
The first ionisation energy for an element is the energy needed to
remove one mole of electrons from one mole of gaseous atoms.
Na (g) ➝ Na+(g) + e-
∆H i1 = 496 kJ mol-1
Ca (g) ➝ Ca+(g) + e-
∆H i1 = 590 kJ mol-1
F (g) ➝ F+(g) + e-
∆H i1 = 1680 kJ mol-1
4
BILAL HAMEED
ELECTRONIC CONFIGURATION
69
SUCCESSIVE IONISATION ENERGIES
Successive ionisation energies for the same element measure the
energy to remove a second, third, fourth electron and so on.
Na+(g) ➝ Na2+(g) + e- ∆H i2 = 4563 kJ mol-1
Na2+(g) ➝ Na3+(g) + e- ∆H i3 = 6913 kJ mol-1
It is possible to measure energy changes involving ions which do not
normally appear in chemical reactions.
5
SKILL CHECK 1
Which equation represents the second ionisation energy of an element X?
A X(g)
X2+(g) + 2e-
B X+(g)
X2+(g) + e-
C X(g) + 2e-
X2-(g)
D X-(g) + e-
X2-(g)
6
BILAL HAMEED
ELECTRONIC CONFIGURATION
70
SKILL CHECK 2
Write equations to represent the first ionisation of:
(a) potassium
(b) argon
(c) bromine
(d) fluorine
7
SKILL CHECK 3
The successive ionisation energies kJmol-1 of element X are listed
below. Identify the group in the periodic table in which X occurs.
Ionisation energies of X:
1st
2nd
3rd
4th
5th
6th
7th
950
1800
2700
4800
6000
12300
15000
8
BILAL HAMEED
ELECTRONIC CONFIGURATION
71
they repel
pel electrons
ns prevent
uter electrons.
hielding of
s, the lower
nd the
mber of full
ns and the
energies
• The second electron is much more difficult to remove
than the first electron. There is a big jump in the value
IONISATION
ENERGIES
ofSUCCESSIVE
the ionisation energy.
This suggests that
the second
inbeen
a shellremoved
closer tothethe
nucleus
themore
After anelectron
electronishas
rest
of themthan
will be
rst electron.
Taken
together, the 1st and 2nd
stronglyfiattracted
by the
nucleus.
ionisation energies suggest that sodium has one
Hence more energy is required to pull the 2nd electron and thus the 2nd
electron in its outer shell.
I.E. is greater than the 1st I.E.
• From the second to the ninth electrons removed there is
Successive
always greater
thanenergies.
the previous
onlyionisation
a gradual energies
change inare
successive
ionisation
one. This suggests that all these eight electrons are in the
same shell.
• The 10th and 11th electrons have extremely high
ionisation energies, when compared to the rest of
the data. This suggests that they are very close to the
nucleus. There must be a very great force of attraction
between the nucleus and these electrons and there are
no inner electrons to shield them. The large increase in
ionisation energy between the 9th and 10th electrons
confirms that the 10th electron is in a shell closer to the
nucleus than the 9th electron.
9
isation
emoved for
0) is used
energies have
ium from
SUCCESSIVE IONISATION ENERGIES
ionisation
e data. It is very
efore likely to be
ielded by inner
2 electrons;
very difficult
to remove
1 electron;
very easily
removed
11+
8 electrons;
less easily
removed
nucleus
10
Figure 3.5 The arrangement of electrons
in an atom of sodium can be
deduced from the values of successive ionisation energies.
BILAL HAMEED
Check-up
ELECTRONIC CONFIGURATION
3 a The successive ionisation energies of
of an
d to
ch atom
s.
ergy of
ergy
third or
m one
element
charge.
72
EVIDENCE OF ENERGY LEVELS
The arrangement of electrons in an atom of any element can be deduced
from the values of successive ionisation energies.
The successive I.E of sodium illustrate the change clearly.
In neon advertising signs, gaseous neon atoms are continually bombarded by
electron
electrons. This
produces positive
neon
as neon
1
2 ions,3which 4then reform
5
6 atoms.
7
removed
This is what
happens when scientists measure ionisation energies.
Scientists can also determine ionisation energies from the emission spectra
ionisation
of atoms. Using
data from spectra, it is possible to measure the energy needed
500 4600 6900 9500 13400 16600
20100
to remove electrons
energy from ions with increasing positive charges. A succession
of ionisation energies, represented by the symbols ∆Hi1, ∆Hi2, ∆Hi3 and so on,
is obtained. For example:
Na(g) → Na+(g) + e–
Na (g) → Na (g) + e
second ionisation energy, ∆Hi2 =11+4563 kJ mol–1
Na2+(g) → Na3+(g) + e–
third ionisation energy, ∆Hi3 = +6913 kJ mol–1
2+
9
25500
28900
10
11
141000 158000
first ionisation energy, ∆Hi1 = +496 kJ mol–1
–
+
8
There are 11 electrons in a sodium atom, so there are 11 successive ionisation
energies for this element.
The successive ionisation energies of an element are all endothermic and
they get bigger and bigger. This is not surprising because, having removed one
electron, it is more difficult to remove a second electron from the positive ion
formed.
The graph in Figure 5.2 shows a logarithmic plot of the successive
ionisation energies of sodium against the number of electrons removed. This
provides evidence to support the theory that electrons in an atom are
arranged in a series of levels or shells around the nucleus. The logarithmic
plot allows an extremely wide range of ionisation energies – from 496 kJ mol–1
to 159 080 kJ mol–1 – to be shown on the same graph.
IONISATION ENERGIES OF SODIUM
against
ed for
Log ionisation energy
The graph shows shows
successive ionisation energies
against the number of
electrons removed for sodium.
d or
called
e word
e
amount
st energy
electron
removed
0
5
10
Number of electrons removed
12
Notice in Figure 5.2 the big jumps in value between the first and second
ionisation energies and again between the ninth and tenth ionisation
energies. This suggests that sodium atoms have one electron in an outer
shell or energy level furthest from the nucleus. This outer electron is easily
removed because it is furthest from the nucleus and shielded from the full
t energy
attraction of the positive nucleus by 10 inner electrons.
electrons
Below this outer single electron, sodium atoms seem to have eight electrons
st to remove
BILAL
HAMEED
in a second
shell – all at roughly the same energy level. These eight electrons
are closer to the nucleus than the single outer electron and only have two
dium atom. inner electrons shielding them from the positive nucleus.
ediate
y level –
ons harder
ove
ELECTRONIC CONFIGURATION
73
IONISATION ENERGIES OF SODIUM
There is a big difference between some successive ionisation energies.
For sodium the first big difference occurs between the 1st and 2nd
ionisation energies.
These large changes indicate that for the second of these two ionisation
energies, the electron being removed is from a shell closer to the
nucleus.
13
IONISATION ENERGIES OF SODIUM
electron
removed
1
ionisation
energy
500
2
3
4
4600 6900 9500
5
6
7
8
9
13400
16600
20100
25500
28900
10
11
141000 158000
There is a big jump in the value of the second ionisation energy. This
suggests that the second electron is in a shell closer to the nucleus than
the first electron.
Taken together, the 1st and 2nd ionisation energies suggest that sodium
has one electron in its outer shell.
14
BILAL HAMEED
ELECTRONIC CONFIGURATION
74
IONISATION ENERGIES OF SODIUM
electron
removed
1
2
ionisation
energy
500
4600
3
4
5
6900 9500 13400
6
7
8
9
16600
20100
25500
28900
10
11
141000 158000
From the second to the ninth electrons removed there is only a gradual
change in successive ionisation energies. This suggests that all these
eight electrons are in the same shell.
There is a big jump in the value of the 10th ionisation energy. This
suggests that the 10th electron is in a shell closer to the nucleus than the
9th electron.
15
2
The structure of the atom
Ionisation energies are used to probe electronic configurations in two ways:
● successive ionisation energies for the same atom
● first ionisation energies for different atoms.
We shall look at each in turn.
Successive ionisation energies
We can look at an atom of a particular element, and measure the energy required to
remove each of its electrons, one by one:
X(g) → X+(g) + e−
X+(g) → X2+(g) + e−
X2+(g) → X3+(g) + e−
∆ H = IE1
∆ H = IE2
∆ H = IE3
etc.
These successive ionisation energies show clearly the arrangement of electrons
in shells around the nucleus. If we take the magnesium atom as an example, and
measure the energy required to remove successively the first electron, the second, the
third, and so on, we obtain the plot shown in Figure 2.29.
IONISATION ENERGIES OF MAGNESIUM
Figure 2.29 Graph of the twelve ionisation
energies of magnesium against electron
number. The electronic configuration of
magnesium is 1s2 2s2 2p6 3s2.
200 000
first
shell
180 000
160 000
ionisation energy/kJ mol–1
140 000
120 000
100 000
80 000
60 000
40 000
second
shell
20 000
third
shell
0
1
2
3
4
5
6
7
8
9
10
11
12
number of electrons removed
Successive ionisation energies are bound to increase because the remaining electrons
are closer to, and less shielded from, the 16
nucleus. But a larger increase occurs when
the third electron is removed. This is because once the two electrons in the outer
(third) shell have been removed, the next has to be stripped from a shell that is very
muchCam/Chem
nearer to AS&A2
the nucleus (the second shell). A similar, but much more enormous,
02_29
jump in ionisation energy occurs when the eleventh electron is removed. This has
Barking
Art the first, innermost shell, right next to the nucleus. These two large
to comeDog
from
jumps in the series of successive ionisation energies are very good evidence that the
electrons in the magnesium atom exist in three different shells.
37
BILAL HAMEED
181333_02_AS_Chem_BP_020-043.indd 37
ELECTRONIC CONFIGURATION
15/10/14 12:19 PM
75
PHYSICAL CHEMISTRY
The jumps in successive ionisation energies are more apparent if we plot the
logarithm of the ionisation energy against proton number, as in Figure 2.30. (Taking
the logarithm is a scaling device that has the effect of decreasing the differences
between adjacent values for the larger ionisation energies, so the jumps between the
shells become more obvious.)
IONISATION ENERGIES OF MAGNESIUM
Figure 2.30 Graph of logarithms of the
twelve ionisation energies of magnesium
against electron number
5.5
1s
5
2s
4.5
log (IE)
2p
4
3.5
3s
3
2.5
2
1
2
3
4
5
6
7
8
9
10
11
12
number of electrons removed
17
Worked example
02_30 Cam/Chem AS&A2
The
first five successive ionisation energies of element X are 631, 1235, 2389, 7089 and
−1
8844
kJ mol
.
Barking
Dog Art
How many electrons are in the outer shell of element X?
Answer
The differences between the successive ionisation energies are as follows:
Graphs 2of
ionisation
energy give us information about
− 1:successive
1235 − 631 = 604
kJ mol
3 − 2: 2389 − 1235 = 1154 kJ mol
how
many
electrons
are
in
a
particular
energy level. Consider the graph
4 − 3: 7089 − 2389 = 4700 kJ mol
Decide which group element Y is in,
5 − 4: 8844 − 7089 = 1755 kJ mol
based on the following successive
for silicon
shown in Figure 2.35. There is a large jump in the ionisation
The largest jump comes between the third and the fourth ionisation energies, therefore X
ionisation energies:
has
three between
electrons in its outer
graph
theshell.
fourth and the fifth ionisation energies, which
590, 1145, 4912, 6474, 8144 kJenergy
mol
suggests that these electrons are removed from different main energy
First ionisation energies
levels.
It The
cansecond
therefore
be ENERGIES
deduced
that
silicon
has four
in its
way that ionisation
energies show
us the
details ofOF
electronic
confielectrons
guration
IONISATION
SILICON
is to look at how the first ionisation energies of elements vary with proton number.
outer main
energy level (shell) and is in group 4 of the periodic table.
Figure 2.31 is a plot for the first 40 elements.
=A
Now try this
−1
−1
−1
−1
−1
This graph shows us the following.
(''+
38
181333_02_AS_Chem_BP_020-043.indd 38
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*
1 All ionisation energies are strongly endothermic – it takes energy to separate an
electron from an atom.
2 As we go down a particular group, for example from helium to neon to argon, or
from lithium to sodium to potassium, ionisation energies decrease. The larger the
atom, the easier it is to separate an electron from it.
%
&
'
(
)
*
+
,
. &%
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&)
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18
6eea^XVi^dc'
If a graph of ionisation energy (rather than log10 ionisation energy) is
plotted for the removal of the first few electrons from a silicon atom, more
¤
features can be seen on the graph (Figure 2.36). It can be seen that there
is a larger jump in the ionisation energy between the second and third
BILAL HAMEED ionisation energies.
ELECTRONIC CONFIGURATION
The full electronic configuration for silicon is 1s22s22p63s23p2. The first
two electrons are removed from the 3p sub-level (subshell), whereas the
iigVXi^dcWZilZZcZaZXigdc
cYcjXaZjh
2 Once an electron has been removed from an atom, there is less
repulsion between the remaining electrons. They are therefore pulled
in closer to the nucleus (Figure 2.33). If they are closer to the nucleus,
76
they are more strongly attracted and more difficult to remove.
HjXXZhh^kZ^dc^hVi^dcZcZg\^Zhd[ediVhh^jb
hhgZejah^dcWZilZZc
ZXigdch"ZaZXigdchejaaZY
XadhZgidcjXaZjh
This graph shows the energy required to remove each electron in turn
from a gaseous potassium
atom.
IONISATION
ENERGIES
OF POTASSIUM
+#%
'
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-ZaZXigdch
cZaZXigdc^hgZbdkZY
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19
scale is being used here to allow all
ta to be plotted on one graph, but
gh on one level this has made the data
ret and supported the explanations that
n, it has also distorted the data. The
een the first and second ionisation
assium is about 2600 kJ mol−1, but the
een the 18th and 19th ionisations
30 000 kJ mol−1! How can the way
data are presented be used by scientists to support
their theories? Can you find examples where the
scale on a graph has been chosen to exaggerate a
particular trend – is scientific knowledge objective
or is it a matter of interpretation and presentation?
The arguments for and against human-made
climate change are a classic example of where the
interpretation and presentation of data are key in
influencing public opinion.
SKILL CHECK 4
th
to represent the 5 ionisation energy of Fluorine.
of these numbersWrite an equation
The simple electronic configuration of potassium is 2,8,8,1, and this
The first
can be deduced directly from the graph above. The large jumps in the
of potassium is
graph occur between main energy levels (shells).
eas the 19th is
t would be very
ese values on a
20
BILAL HAMEED
ELECTRONIC CONFIGURATION
77
SKILL CHECK 5
What group is X in?
21
SKILL CHECK 6
The first ionisation energies of four consecutive elements in the Periodic table
are:
Sodium = 494 kJ mol-1
Magnesium = 736 kJ mol-1
Aluminium = 577 kJ mol-1
Silicon = 786 kJ mol-1
a Explain the general increase in ionisation energies from Sodium to Silicon.
b Explain why Aluminium has a lower first ionisation energy than Magnesium.
22
BILAL HAMEED
ELECTRONIC CONFIGURATION
78
SKILL CHECK 7
The first six ionisation energies of an element are, or 1090, 2250, 4610,
6220, 37,800, and 47,300kJ mol-1. Which group in the Periodic Table
does this element belong to? Explain your decision.
23
SKILL CHECK 8
The successive ionisation energies ∆Hi, of an element X are shown in the
table below. Which group in the Periodic Table does X belong to?
Number of electrons removed
1
∆Hi / kJ mol-1 1000
2
3
4
5
2260 3390 4540
7010
6
7
8
9
10
8500 27100 31670 36580 43140
24
BILAL HAMEED
ELECTRONIC CONFIGURATION
79
SUCCESSIVE IONISATION ENERGIES
We can use successive ionisation energies in this way to confirm:
• The simple electronic configuration of elements.
• The number of electrons in the outer shell of an element and hence
the group to which the element belongs.
The successive ionisation energies for an element rise and there are big
jumps in value each time electrons start to be removed from the next
shell in towards the nucleus.
25
IONISATION ENERGIES OF OXYGEN
electron removed
1
2
3
4
5
6
7
8
ionisation energy
1310
3390
5320
7450
11000
13300
71300
84100
Large increases can be used to predict the group of any element. The
electron configuration of oxygen is 2,6.
Since the large change is after the removal of 6 electrons, it signifies that
there are 6 electrons in the shell farthest from the nucleus.
Therefore, Oxygen is in Group VI.
26
BILAL HAMEED
ELECTRONIC CONFIGURATION
80
SUCCESSIVE IONISATION ENERGIES
ELECTRONS REMOVED
Element
1
2
3
4
5
6
7
8
9
10
1
H
1310
2
He
2370
5250
3
Li
519
7300
11800
4
Be
900
1760
14850
21000
5
B
799
2420
3660
25000
32800
6
C
1090
2350
4620
6220
37800
47300
7
N
1400
2860
4580
7480
9450
53300
64400
8
O
1310
3390
5320
7450
11000
13300
71300
84100
9
F
1680
3370
6040
8410
11000
15200
17900
92000
106000
10
Ne
2080
3950
6150
9290
12200
15200
20000
23000
117000
131400
11
Na
494
4560
6940
9540
13400
16600
20100
25500
28900
141000
11
158700
27
SKILL CHECK 9
The successive ionisation energies, in kJ mol–1, of different elements are
given below. Which groups are the following elements in?
1
2
3
4
5
6
7
A
799
2420
3660
25000
B
736
1450
7740
10500
C
418
3070
4600
5860
D
870
1800
3000
3600
5800
7000
13200
E
950
1800
2700
4800
6000
12300
8
28
BILAL HAMEED
ELECTRONIC CONFIGURATION
81
SKILL CHECK 10
The successive ionisation energies of beryllium are 900, 1757, 14,849
and 21,007 kJ mol-1.
a What is the atomic number of beryllium?
b Why do successive ionisation energies of beryllium always get
more endothermic?
c To which group of the Periodic Table does this element belong?
29
SHELLS (ENERGY LEVELS)
n=1
The principal energy levels are
designated n = 1, 2, 3, and so forth.
The energy levels are not equally
spaced.
The energy gap between successive
levels gets increasingly smaller as the
levels move further from the nucleus.
n=2
n=3
n=4
30
BILAL HAMEED
ELECTRONIC CONFIGURATION
82
SUB-LEVELS (SUB-SHELLS)
Electron shells are numbered 1,2,3 etc. These numbers are known as the
principle quantum numbers.
Each energy level (shell) consists of a number of sub-levels (sub-shells),
labeled s, p, d, or f.
Energy Level
Number of sub-levels
Name of sub-levels
1
1
s
2
2
s, p
3
3
s, p, d
4
4
s, p, d, f
31
SUB-LEVELS (SUB-SHELLS)
INCREASING ENERGY / DISTANCE FROM NUCLEUS
PRINCIPAL
ENERGY
LEVELS
SUB LEVELS
4
3
2
1
32
BILAL HAMEED
ELECTRONIC CONFIGURATION
Quantum sub-shells
We can use successive ionisation energies in this way to:
• predict or confirm the simple electronic configuration
of elements
• confirm the number of electrons in the outer shell
of an element and hence the group to which the
element belongs.
The principal quantum shells, apart from the first, are
split into sub-shells (sub-levels). Each principal quantum
shell contains a different number of sub-shells. The sub83
shells are distinguished by the letters s, p or d. There are
also f sub-shells for elements with more than 57 electrons.
Figure 3.6 shows the sub-shells for the first four principal
quantum levels. In any principal quantum shell, the
energy of the electrons in the sub-shells increases in the
order s < p < d.
Worked example
SUB-LEVELS (SUB-SHELLS)
4d
5s
n =4
4p
3d
4s
Increasing energy
1 The successive ionisation energies, ∆Hi, of an
element X are shown in Table 3.4. Which group
in the Periodic Table does X belong to?
We look for a large jump in the value of the
ionisation energy. This occurs between the
removal of the 6th and 7th electrons. So, six
electrons have been removed comparatively
easily. The removal of the 7th electron requires
about three times the energy required to remove
the 6th electron. So, there must be six electrons
in the outer shell of X. So, element X must be
in Group VI of the Periodic Table.
n =3
3p
3s
2p
Check-up
n =2
2s
4 a The first six ionisation energies of an
element are 1090, 2350, 4610, 6220,
37 800 and 47 300 kJ mol−1. Which group
in the Periodic Table does this element
belong to? Explain your decision.
n =1
1s
principal quantum
shell
sub-shell
continued
33
Figure 3.6 The sub-shells for the first four principal quantum shells.
∆Hi /
kJ mol−1
Number of electrons removed
1
2
3
4
1000
2260
3390
4540
5
6
7010
7
8500
8
27 100
9
31 670
36 580
10
43 140
Table 3.4 The successive ionisation energies of an element X. For Worked example 1.
3 Electrons in atoms
37
SUB-LEVELS
Each sub-level can hold a certain maximum number of electrons.
Type of sub-level
Maximum # of electrons
s
2
p
6
d
10
f
14
34
BILAL HAMEED
ELECTRONIC CONFIGURATION
84
ORBITALS
An atomic orbital is a region of space around the nucleus of an atom
which can be occupied by one or two electrons only.
Each sub-level contains a fixed number of orbitals that contain electrons.
Type of sub-level
Maximum # of electrons
Number of orbitals
s
2
1
p
6
3
d
10
5
f
14
7
35
ORBITALS
Electrons are viewed as charged clouds and the region, which
encloses almost all the charge cloud, is the orbital.
The region in space where the probability of finding an electron in
maximum is called the orbital.
The boundary surface encloses the region where the probability of
finding an electron is high.
The orbitals are of different three-dimensional shapes and are named s,
p, d, f etc.
36
BILAL HAMEED
ELECTRONIC CONFIGURATION
85
2 . 2 E L E C T R O N C O N F I G U R AT I O N
ORBITALS
could perhaps be at the airport
might even have gone home to South Africa!
he class went looking for the teacher they would most likely start looking in
cluster
of dotsBut
show
the
e most probable locationsThe
closest
to the classroom.
at 8.15
am they do
t know with any degree ofprobability
certainty precisely
where thean
teacher
is.
of finding
electron
at
hree-dimensional graph could
be drawn
with a cluster
of dots
showing
areas
different
distances
from
the
nucleus.
ere there is a high probability of finding him. This is the idea of an orbital. A
undary surface could be drawn around this cluster of dots to define a region
space where there is a 99% chance of finding the teacher. This might be the
hool perimeter, or Quito where he lives.
you were also asked to measure the distance from the classroom to the
act location where the teacher is you could not do this at 8.15 am, as you do
t know his exact location with absolute certainty.
hat aspects of quantum mechanics does this analogy capture?
Figure 8 An orbital is a three-dimensional
graph with a cluster of dots showing the
probability of finding the electron at different
distances from the nucleus
37
e s atomic orbital
s orbital is spherically symmetrical. The sphere represents a
ndary surface, meaning that within the sphere there is a 99%
nce or probability of finding an electron (figure 9).
y
e p atomic orbital
orbital is dumbbell shaped. There are three p atomic orbitals, px, py,
pz, all with boundary surfaces conveying probable electron density
nting in different directions along the three respective Cartesian
s, x, y, and z (figure 10).
SHELLS, SUB SHELLS AND ORBITALS
x
Level
of sub-level
ergy levels, sublevels,Energy
orbitals,
andType
electron
spin Number of orbitals
1
s
Bohr model introduced the idea of a main energy level, described
2
n, which is called the principal quantum
number. sThis can have
tive integer values 1, 2, 3, etc. In the quantum mechanical
model, as n
p
eases, the mean position of an electron is further from the nucleus. The
s
rgies of the orbitals also increase as3n increases. Each main
energy level
hell can hold a maximum number of electrons given pby 2n2. So the
tron capacity for n = 1 is 2, for n = 2 is 8, for n = 3 is 18. That is why
d
have two elements in the first row of the periodic table, eight elements
4
s
he second, etc.
p there are four
energy levels are split up into sublevels, of which
mmon types: s, p, d, and f. Each sublevel contains ad number of
tals, each of which can hold a maximum of 2 electrons (table 2).
ublevel
s
p
d
f
Number of orbitals in
sublevel
1
3
5
7
Maximum number of
electrons in sublevel
2
6
10
14
BILAL HAMEED
ble 2 Sublevels of the main energy levels in the quantum mechanical model
38
Maximum # of electrons
1
2
z
1
2
3
6
Figure
9
The
s
atomic
orbital
is spherically
1
2
symmetrical
3
6
5
10
y
1
y
3
y
2
6
5
x
z
px
10 x
z
pz
x
z
py
Figure 10 The three p atomic orbitals are
dumbbell shaped, aligned along the x, y, and
z axes
ELECTRONIC CONFIGURATION
86
“S” ORBITAL
The s orbital is spherically shaped. There is
only one s orbital for each shell.
The 2s orbital in the second principal
quantum shell has the same shape as
the 1s orbital in the first quantum shell.
They are both spherical, but electrons in the
2s orbital have more energy than electrons
in the 1s orbital.
39
“P” ORBITAL
For each shell (except the first), there are three p orbitals.
p
=A
o
n
ep
en
eo
V
;^\jgZ'#', I]Zi]gZZedgW^iVahi]VibV`ZjeVehjW"aZkZaed^ciVi.%•idZVX]di]Zg#
;^\jgZ'#'+ VI]Zh]Ve
Wi]ZZaZXigdcYZch^in^c
40
Figure 2.28 shows the orbitals that make up the 2s and 2p sub-levels in
the second main energy level.
'h
BILAL HAMEED
ELECTRONIC CONFIGURATION
;^\jgZ'#'. DcZd[i]ZÒk
(YhjW"aZkZa#
bV^cZcZg\naZkZa#HiVgi^c\Vi&h![daadli]Z
Vggdlhid\^kZi]ZdgYZg^c\d[i]ZhjW"aZkZah#
of finding an electron. It represents a discrete energy level.
There are four different types of atomic orbital:
87
eo
s
p
d
The first shell (maximum number of electrons 2) consists of a 1s orb
and this makes up the entire 1s sub-level. This is spherical in shape (Fig
2.24a).
;^\jgZ'#', I]Zi]gZZedgW^iVahi]VibV`ZjeVehjW"aZkZaed^ciVi.%•idZVX]
The 1s orbital is centred on the nucleus (Figure 2.24b). The electro
is moving all the time and the intensity of the colour here represents th
Figure 2.28 shows
the orbitals
thatelectron
makeat up
the distance
2s andfrom
2p the
sub-lev
probability
of finding the
a certain
nucleu
V
W
The darkerlevel.
the colour the greater the probability of the electron being
the
second
main
energy
;^\jgZ'#') VI]Zh]VeZd[V&hdgW^iVa0
that point. This represents the electron density.
Wi]ZZaZXigdcYZch^in^cV&hdgW^iVa#
The electron can be found anywhere in this region of space (except
nucleus – at the centre of the orbital) but it is most likely to be found
'h
certain distance from the nucleus.
'hdgW^iVa
The second main energy level (maximum number of electrons 8) is
made up of the 2s sub-level and the 2p sub-level. The 2s sub-level just
consists of the 2s orbital, whereas the 2p sub-level is made up of three
orbitals. The 2s orbital (like all other s orbitals) is spherical in shape and
bigger than the 1s orbital (Figure 2.25).
&hdgW^iVa
p orbitals have a ‘dumb-bell’ shape (Figure 2.26). Three p orbitals m
up the 2p sub-level. These point at 90° to each other and are named
appropriately as px, py, pz (Figure 2.27). The px orbital points along the
x axis. The three 2p orbitals all have the same energy – they are describ
;^\jgZ'#'* 6XgdhhhZXi^dcd[i]ZZaZXigdc
YZch^ind[i]Z&hVcY'hdgW^iVahid\Zi]Zg#
as degenerate.
An orbital can contain a
maximum of two electrons.
en
THE FIRST TWO SHELLS
;^\jgZ'#'- I]Z'hVcY'ehjW"aZkZah^ci]ZhZXdcYbV^cZcZg\naZkZa#
,%
41
The third shell (maximum 18 electrons) consists of the 3s, 3p and
sub-levels. The 3s sub-level is just the 3s orbital; the 3p sub-level con
of three 3p orbitals; and the 3d sub-level is made up of five 3d orbi
One of the five 3d orbitals is shown in Figure 2.29.
The fourth shell (maximum 32 electrons) consists of one 4s, three
five 4d and seven 4f orbitals. The seven 4f orbitals make up the 4f
level. One of the f orbitals is shown in Figure 2.30.
ELECTRONIC CONFIGURATION
Within any sub-shell all the orbitals have the same energy (th
are degenerate), e.g. the three 2p orbitals are degenerate an
The order in which they
fill up
sub-levels
is governed by stability.
five
3d the
orbitals
are degenerate.
Electrons are distributed in different energy levels in the atom of the element.
When electrons fill up the orbitals having the least energy they attain
maximum stability.The number of orbitals in each energy level is shown in Table 2.4.
There are three principles that describe how electrons fill up in orbitals.
The diagrams of atomic orbitals that
Aufbau Principle: Electrons enter the orbital that is available with the lowest
we have seen here are derived from
energy. The orbitals are arranged
in the order of
increasing
energy
and the
mathematical
functions
that
are solutions
electrons are added until the proper number of electrons for the element
to the Schrödinger equation. Exact solutions of the
have been accommodated
Schrödinger equation are only possible for a system
involving one electron,
i.e. the hydrogen atom. It is
42
not possible to derive exact mathematical solutions
BILAL HAMEED
ELECTRONIC CONFIGURATION
for mor
have fo
describ
we are
hydrog
the pro
88
ELECTRONIC CONFIGURATION
Pauli’s Exclusion Principle: No orbital can accommodate more than two
electrons. If there are two electrons in an orbital, they must have opposite
spin.
Hund’s Rule of Maximum Multiplicity: When there are a number of
orbitals of equal energy, electrons first fill them up individually and then
get paired. By filling up individually, mutual repulsion between electrons
is avoided and thereby maximum stability is achieved.
43
ELECTRON SPIN
Electrons are all identical. The only way of distinguishing them is by
describing how their energies and spatial distributions differ.
Thus an electron in a 1s orbital is different from an electron in a 2s orbital
because it occupies a different region of space closer to the nucleus,
causing it to have less potential energy.
An electron in a 2px orbital differs from an electron in a 2py orbital
because although they have exactly the same potential energy, they
occupy different regions of space.
44
BILAL HAMEED
ELECTRONIC CONFIGURATION
2
89
ATOMIC STRUCTURE
f sublevel (seven boxes representing the seven f orb
Useful resource
The “Orbitron” website,
ELECTRON
SPIN
developed by Professor Mark
For convenience, an “arrow-in-box” notation call
diagram is used to represent the electrons in the
(figure 11). We shall use orbital diagrams to repre
Figure 11 Orbital diagrams are used to represe
configurations.
Winter at the University of
Arrows are drawn in the boxes to represent ele
Sheffield, UK isin
an each
excellentorbital, sand
There can only be two electrons
have
sublevelthey
(one(orbital)
boxmust
representing
an s orbital)
box
resource for exploring the
opposite directions of spin.
shapes of the various atomic
Two electrons in the same orbital hav
orbitals. It also provides
N
S
magnetic quantum number, ms. T
information on the associated
sophisticated mathematical
indicates
orientation
of the
p sublevel
(three boxesthe
representing
the three p orbitals
px, pmagn
y, and pz
wavefunctions.
electron. A pair of electrons in an orb
http://winter.group.shef.ac.uk/
orbitron/
facing in opposite directions and ther
by two arrows in a box (figure 12).
d sublevel (five boxes representing the five d orbitals)
Quantum numbers
S
N
N
S
45
magnet analogy
S
half-arrows representing
N spin
electrons of opposite
in an orbital
N
S
Figure 12 Electron spin is represented by
ELECTRON SPIN
arrows in orbital diagrams
In this mathematical model of the electr
four quantum numbers. The first is the p
f sublevelrepresents
(seven boxes representing
the seven
f orbitals)
the energy
level.
The second
quantum number, l, describes the suble
number, the magnetic quantum number
quantum number, the spin magnetic qu
Figure spatial
11 Orbital orientation
diagrams are used
represent
the electron
of to
the
electron
spin. cQ
Arrows are drawn in the boxes to represent electrons, a maxim
examined in the IB Chemistry Diploma, b
box (orbital)
of energy levels, sublevels, atomic orbit
Two electrons in the same orbital have opposite v
You might think of the four quantum num
magnetic quantum number, ms. The sign of ms
address.
The country
represents
the ene
indicates
the orientation
of the
magnetic field
gen
sublevels,
town the
orbitals,
the
electron.
A pair of the
electrons
in an
orbital and
behaves
facing in
opposite
spin
of the directions
electron. and therefore is com
by two arrows in a box (figure 12).
numbers
For convenience, an “electrons-in-boxes” notation Quantum
is used to
represent
58
In this mathematical model of the electronic structure
the electrons in these atomic orbitals:
four quantum numbers. The first is the principal quan
s sub-level
S
N
N
S
p sub-level
magnet analogy
Each box represents one orbital:
S
half-arrows representing
electrons of opposite spin
in an orbital
N
represents the energy level. The second quantum num
quantum number, l, describes the sublevel, and the th
number, the magnetic quantum number, ml, the atomi
d sub-level
quantum
number, the spin magnetic quantum numbe
spatial orientation of the electron spin. Quantum numb
examined in the IB Chemistry Diploma, but you need t
of energy levels, sublevels, atomic orbitals, and electr
You might think of the four quantum numbers as an ele
address. The country represents the energy level, the p
sublevels, the town the orbitals, and the street numbe
spin of the electron.
Figure
12 Electron
spinelectrons
is represented
an orbital
with
inbyopposite spins
arrows in orbital diagrams
58
BILAL HAMEED
46
ELECTRONIC CONFIGURATION
r Determine the full electronic
6bdgZVYkVcXZY
higjXijgZd[Vidbh
configuration of an atom with
up to 54 electrons
r Understand what is meant by
an orbital and a subshell (sub
energy level)
90
The emission spectra of ato
evidence, such as ionisation
simple treatment of consid
energy levels is a useful firs
HjWZcZg\naZkZah
ORDER OF FILLING ORBITALS
*e
)Y
The principal energy levels (shells)
get closer together as you get
further from the nucleus. This
results in an overlap of sub- levels.
*h
Each main energy level in
(subshells). The first main e
second main energy level i
The sub-levels in each mai
)e
(Y
)h
(h
&
'e
hjW"aZkZa
'h
bV^cZcZg\naZkZa
^cXgZVh^c\ZcZg\n
(e
BV^cZcZg\n
aZkZa
HjW"aZkZah
&h
'
'h
'e
(
(h
(e
)
)h
)e
*
*h
*e
IVWaZ'#( I]ZhjW"aZkZah^cZV
&h
;^\jgZ'#'& I]ZdgYZg^c\d[i]ZZcZg\n
aZkZahVcYhjW"aZkZahl^i]^cVcVidb#I]Z
hjW"aZkZahl^i]^cVbV^cZcZg\naZkZaVgZ
h]dlc^ci]ZhVbZXdadjg#
47
Within any main energy
(subshells) is always s < p <
orders between sub-levels i
the subshells are shown in F
I]Z6j[WVjWj^aY^c\"j
The Aufbau principle is sim
out the electronic configur
+-
The first example of this overlap
occurs when the 4s orbital is
filled before the 3d orbital
INCREASING ENERGY / DISTANCE FROM NUCLEUS
ORDER OF FILLING ORBITALS
4
4p
3d
4s
3
3p
3s
2
1
2p
2s
1s
48
BILAL HAMEED
ELECTRONIC CONFIGURATION
91
INCREASING ENERGY / DISTANCE FROM NUCLEUS
THE ‘AUFBAU’ PRINCIPAL
4
3
The following sequence will show the
‘building up’ of the electronic structures of
the first 36 elements in the periodic table.
4p
3d
4s
3p
3s
2
1
2p
Electrons are shown as half-headed
arrows and can spin in one of two
directions.
Orbitals are color-coded as below:
s orbitals
2s
p orbitals
1s
d orbitals
49
HYDROGEN
INCREASING ENERGY / DISTANCE FROM NUCLEUS
1s1
4
3
Hydrogen atoms have one
electron. This goes into a vacant
orbital in the lowest available
energy level.
4p
3d
4s
3p
3s
2
1
2p
2s
1s
50
BILAL HAMEED
ELECTRONIC CONFIGURATION
92
HELIUM
INCREASING ENERGY / DISTANCE FROM NUCLEUS
1s2
4
3
Every orbital can contain 2
electrons, provided the electrons
are spinning in opposite directions.
4p
3d
4s
3p
The two electrons in a helium atom
both go in the 1s orbital.
3s
2
1
2p
2s
1s
51
LITHIUM
INCREASING ENERGY / DISTANCE FROM NUCLEUS
1s2 2s1
4
3
4p
3d
4s
3p
3s
2
1
2p
2s
1s
1s orbitals can hold a maximum of two
electrons so the third electron in a
lithium atom must go into the next
available orbital of higher energy. This
will be further from the nucleus in the
second principal energy level.
The second principal level has two
types of orbital (s and p). An s orbital
is lower in energy than a p.
52
BILAL HAMEED
ELECTRONIC CONFIGURATION
93
BERYLLIUM
INCREASING ENERGY / DISTANCE FROM NUCLEUS
1s2 2s2
4
3
4p
3d
4s
3p
3s
2
1
Beryllium atoms have four electrons
so the fourth electron pairs up in the
2s orbital. The 2s sub level is now
full.
2p
2s
1s
53
BORON
INCREASING ENERGY / DISTANCE FROM NUCLEUS
1s2 2s2 2p1
4
3
4p
3d
4s
3p
3s
2
1
2p
As the 2s sub level is now full, the
fifth electron goes into one of the
three p orbitals in the 2p sub level.
The 2p orbitals are slightly higher in
energy than the 2s orbital.
2s
1s
54
BILAL HAMEED
ELECTRONIC CONFIGURATION
94
CARBON
INCREASING ENERGY / DISTANCE FROM NUCLEUS
1s2 2s2 2p2
4
3
The next electron in doesn’t pair up
with the one already there. This
would give rise to repulsion
between the similarly charged
species. Instead, it goes into
another p orbital which means less
repulsion, lower energy and more
stability.
4p
3d
4s
3p
3s
2
1
2p
2s
1s
55
NITROGEN
INCREASING ENERGY / DISTANCE FROM NUCLEUS
1s2 2s2 2p3
4
3
The next electron will not pair up
but goes into a vacant p orbital. All
three electrons are now unpaired.
This gives less repulsion, lower
energy and therefore more
stability.
4p
3d
4s
3p
3s
2
1
2p
2s
1s
56
BILAL HAMEED
ELECTRONIC CONFIGURATION
95
OXYGEN
INCREASING ENERGY / DISTANCE FROM NUCLEUS
1s2 2s2 2p4
4
3
With all three orbitals half-filled, the
eighth electron in an oxygen atom
must now pair up with one of the
electrons already there.
4p
3d
4s
3p
3s
2
1
2p
2s
1s
57
FLUORINE
INCREASING ENERGY / DISTANCE FROM NUCLEUS
1s2 2s2 2p5
4
3
The electrons continue to pair up
with those in the half-filled orbitals.
4p
3d
4s
3p
3s
2
1
2p
2s
1s
58
BILAL HAMEED
ELECTRONIC CONFIGURATION
96
NEON
INCREASING ENERGY / DISTANCE FROM NUCLEUS
1s2 2s2 2p6
4
3
The electrons continue to pair up
with those in the half-filled orbitals.
The 2p orbitals are now completely
filled and so is the second principal
energy level.
4p
3d
4s
3p
3s
2
1
2p
In the older system of describing
electronic configurations, this would
have been written as 2,8.
2s
1s
59
SODIUM TO ARGON
Na
1s2 2s2 2p6 3s1
Mg
1s2 2s2 2p6 3s2
Al
1s2 2s2 2p6 3s2 3p1
Si
1s2 2s2 2p6 3s2 3p2
P
1s2 2s2 2p6 3s2 3p3
S
1s2 2s2 2p6 3s2 3p4
Cl
1s2 2s2 2p6 3s2 3p5
Ar
1s2 2s2 2p6 3s2 3p6
60
BILAL HAMEED
ELECTRONIC CONFIGURATION
97
POTASSIUM
INCREASING ENERGY / DISTANCE FROM NUCLEUS
1s2 2s2 2p6 3s2 3p6 4s1
4
3
4p
3d
4s
3p
3s
2
1
2p
2s
Because the principal energy levels
get closer together as you go
further from the nucleus coupled
with the splitting into sub energy
levels, the 4s orbital is of a LOWER
ENERGY than the 3d orbitals so it
gets filled first.
1s
61
CALCIUM
INCREASING ENERGY / DISTANCE FROM NUCLEUS
1s2 2s2 2p6 3s2 3p6 4s2
4
3
As expected, the next electron
4p
3d
4s
3p
pairs up to complete a filled 4s
orbital.
3s
2
1
2p
2s
1s
62
BILAL HAMEED
ELECTRONIC CONFIGURATION
98
ELEMENTS WITH PROTON NUMBERS 1 TO 10
63
ELEMENTS WITH PROTON NUMBERS 11 TO 20
64
BILAL HAMEED
ELECTRONIC CONFIGURATION
99
SKILL CHECK 11
Copy and complete the following information for the quantum shell
with principal quantum number 3.
(a) total number of sub-shells
(b) total number of orbitals
(c) number of different types of orbital
(d) maximum number of electrons in the shell
65
SKILL CHECK 12
An atom has eight electrons. Which diagram shows the electronic configuration of
this atom in its lowest energy state?
66
BILAL HAMEED
ELECTRONIC CONFIGURATION
re
5.
wo
sses
re
to
a) Phosphorus: [Ne]…
b) Cobalt: [Ar]…
100
6 Write the full electron configuration of an
excited sodium atom.
SKILL CHECK 13
7 Fill in the outer electrons of a phosphorus
Fill in atom
the outer
of a phosphorus
inelectrons
the boxes
below. atom in the boxes below.
3s
3p
8 Explain, with an example, the meaning of the
term periodicity.
9 Consider atoms of boron, magnesium and
bromine. In which is the effective nuclear charge
the largest and in which is it the smallest?
10 List the particles Cl, Cl− and K+ in order of
increasing radius.
67
11 Explain why the first ionisation energies of the
noble gases decrease from helium to krypton.
12 The successive ionisation energies of an element,
X, are given in SKILL
the table.
To which
CHECK
14 group of the
periodic table does element X belong?
Give the electron orbital configuration for the ground state of the
following atoms or ions:
Ionisation energy/kJ mol−
(a) N 1st
1 000
(b) O2- 2nd
2 260
(c) Ca2+3rd
3 390
(d) Al3+4th
4 540
(e) P3- 5th
6 990
6th
8 490
7th
27 100
8th
31 700
9th
36 600
10th
43 100
Ionisation
BILAL HAMEED
1
68
ELECTRONIC CONFIGURATION
13 Discuss the relative advantages and disadvantages
101
SKILL CHECK 15
Write the full electronic configuration of:
(a) 75As
(b) 75As3(c) 32S
(d) 32S2-
69
SKILL CHECK 16
Name all element in the third period (row) of the periodic table
with the following:
a three valence electrons
b four 3p electrons
c six 3p electrons
d two 3s electrons and no 3p electrons
70
BILAL HAMEED
ELECTRONIC CONFIGURATION
102
EXCEPTIONS IN TRANSITION ELEMENTS
Though the s orbital is at lower energy level than the d of the penultimate
shell, after the filling of the d sub-level, the order changes.
The d electrons, because of the shape of the d orbital, penetrate into the
region of space between the nucleus and the s orbital and repel the s
electrons and push them to higher energy level.
Before filling: 1s 2s 2p 3s 3p 4s 3d
After filling: 1s2 2s2 2p6 3s2 3p6 3dx 4s2
71
EXCEPTIONS IN TRANSITION ELEMENTS
Thus the electronic configuration of iron is
1s2 2s2 2p6 3s2 3p6 3d6 4s2
and not
1s2 2s2 2p6 3s2 3p6 4s2 3d6
Remember: 3d is higher than 4s in terms of energy levels!
72
BILAL HAMEED
ELECTRONIC CONFIGURATION
103
SCANDIUM
INCREASING ENERGY / DISTANCE FROM NUCLEUS
1s2 2s2 2p6 3s2 3p6 3d1 4s2
4
3
4p
3d
4s
3p
3s
2
1
With the lower energy 4s orbital
filled, the next electrons can now fill
the 3d orbitals. There are five d
orbitals. They are filled according to
Hund’s Rule.
2p
2s
1s
73
TITANIUM
INCREASING ENERGY / DISTANCE FROM NUCLEUS
1s2 2s2 2p6 3s2 3p6 3d2 4s2
4
3
4p
3d
4s
3p
3s
2
1
The 3d orbitals are filled according
to Hund’s rule so the next electron
doesn’t pair up but goes into an
empty orbital in the same sub level.
2p
2s
1s
74
BILAL HAMEED
ELECTRONIC CONFIGURATION
104
VANADIUM
INCREASING ENERGY / DISTANCE FROM NUCLEUS
1s2 2s2 2p6 3s2 3p6 3d3 4s2
4
3
4p
3d
4s
3p
3s
2
1
The 3d orbitals are filled according
to Hund’s rule so the next electron
doesn’t pair up but goes into an
empty orbital in the same sub level.
2p
2s
1s
75
CHROMIUM
INCREASING ENERGY / DISTANCE FROM NUCLEUS
1s2 2s2 2p6 3s2 3p6 3d5 4s1
4
3
4p
3d
4s
3p
3s
2
1
2p
2s
1s
One would expect the configuration
of chromium atoms to end in 3d4
4s2.
To achieve a more stable
arrangement of lower energy, one of
the 4s electrons is promoted into the
3d to give six unpaired electrons
with lower repulsion.
76
BILAL HAMEED
ELECTRONIC CONFIGURATION
105
MANGANESE
INCREASING ENERGY / DISTANCE FROM NUCLEUS
1s2 2s2 2p6 3s2 3p6 3d5 4s2
4
3
4p
3d
4s
3p
The new electron goes into the 4s
to restore its filled state.
3s
2
1
2p
2s
1s
77
IRON
INCREASING ENERGY / DISTANCE FROM NUCLEUS
1s2 2s2 2p6 3s2 3p6 3d6 4s2
4
3
Orbitals are filled according to
Hund’s Rule. They continue to pair
up.
4p
3d
4s
3p
3s
2
1
2p
2s
1s
78
BILAL HAMEED
ELECTRONIC CONFIGURATION
106
COBALT
INCREASING ENERGY / DISTANCE FROM NUCLEUS
1s2 2s2 2p6 3s2 3p6 3d7 4s2
4
3
4p
3d
4s
3p
Orbitals are filled according to
Hund’s Rule. They continue to pair
up.
3s
2
1
2p
2s
1s
79
NICKEL
INCREASING ENERGY / DISTANCE FROM NUCLEUS
1s2 2s2 2p6 3s2 3p6 3d8 4s2
4
3
4p
3d
4s
3p
Orbitals are filled according to
Hund’s Rule. They continue to pair
up.
3s
2
1
2p
2s
1s
80
BILAL HAMEED
ELECTRONIC CONFIGURATION
107
COPPER
INCREASING ENERGY / DISTANCE FROM NUCLEUS
1s2 2s2 2p6 3s2 3p6 3d10 4s1
4
3
4p
3d
4s
3p
3s
2
1
2p
2s
One would expect the configuration
of chromium atoms to end in 3d9
4s2.
To achieve a more stable
arrangement of lower energy, one
of the 4s electrons is promoted into
the 3d.
1s
81
ZINC
INCREASING ENERGY / DISTANCE FROM NUCLEUS
1s2 2s2 2p6 3s2 3p6 3d10 4s2
4
3
4p
3d
4s
3p
The electron goes into the 4s to
restore its filled state and complete
the 3d and 4s orbital filling.
3s
2
1
2p
2s
1s
82
BILAL HAMEED
ELECTRONIC CONFIGURATION
108
GALLIUM TO KRYPTON
The 4p orbitals are filled in exactly the same way as those in the 2p and 3p sub levels:
2
2
6
2
6
10
2
1s 2s 2p 3s 3p 3d 4s
1
Ga
- 4p
Ge
- 4p
As
- 4p
Se
- 4p
Br
- 4p
Kr
- 4p
2
3
4
5
6
83
ELECTRONIC CONFIGURATION OF IONS
Metallic elements, belonging to Group I, II and III form positively charged
ions by losing electrons in their outermost shell, so that the ions may be
iso-electronic with the preceding noble gas.
SODIUM
Na 1s2 2s2 2p6 3s1
Na+ 1s2 2s2 2p6
(1 electron removed from the 3s orbital)
84
BILAL HAMEED
ELECTRONIC CONFIGURATION
109
ELECTRONIC CONFIGURATION OF IONS
Non-metallic elements belonging to Group V, VI and VII form negative
ions by accepting electrons so that they achieve the noble gas
configuration:
CHLORINE Cl 1s2 2s2 2p6 3s2 3p5
Cl¯ 1s2 2s2 2p6 3s2 3p6
(1 electron added to the 3p orbital)
85
TRANSITION METAL IONS
When transition metals form ions, electrons are lost first from the
outermost s orbital and then from the penultimate d sub-level.
Electrons in the 4s orbital are removed before any electrons in the 3d
orbitals.
TITANIUM
Ti
1s2 2s2 2p6 3s2 3p6 3d2 4s2
Ti+
1s2 2s2 2p6 3s2 3p6 3d2 4s1
Ti2+
1s2 2s2 2p6 3s2 3p6 3d2
Ti3+
1s2 2s2 2p6 3s2 3p6 3d1
Ti4+
1s2 2s2 2p6 3s2 3p6`
86
BILAL HAMEED
ELECTRONIC CONFIGURATION
110
SKILL CHECK 17
Write electronic configurations for the following ions:
a
Al3+
d Cu2+
b
O2-
e Cu+
c
Fe3+
87
SKILL CHECK 18
Write the electron configuration for each ion.
a O2b Brc Sr2+
d Co3+
88
BILAL HAMEED
ELECTRONIC CONFIGURATION
111
SKILL CHECK 19
Q. What is the order of increasing energy of the listed orbitals in the
atom of titanium?
A 3s 3p 3d 4s
B 3s 3p 4s 3d
C 3s 4s 3p 3d
D 4s 3s 3p 3d
89
SKILL CHECK 20
A simple ion X+ contains eight protons.
What is the electronic configuration of X+?
A 1s2 2s1 2p6
B 1s2 2s2 2p3
C 1s2 2s2 2p5
D 1s2 2s2 2p7
90
BILAL HAMEED
ELECTRONIC CONFIGURATION
112
SKILL CHECK 21
Write down the electronic configurations of the following ions:
A bromide
B magnesium
C iron(II)
D copper (II)
E iron(III)
91
SKILL CHECK 22
Q. Determine the number of unpaired electrons in each of these
atoms:
A phosphorus
B chromium
C oxygen
92
BILAL HAMEED
ELECTRONIC CONFIGURATION
113
SKILL CHECK 23
Use the periodic table to determine the element corresponding
to each electron configuration.
A [Ar] 4s2 3d10 4p6
B [Ar] 4s2 3d2
C [Kr] 5s2 4d10 5p2
D [Kr] 5s2
93
PERIODIC TABLE
The modern Periodic Table is arranged such that elements with similar
electronic configuration lie in vertical groups.
Thus elements with electronic configuration ending up as ns1 are put in
one group while those that end up as ns2 are put in another group.
The elements that fill up the s orbital of the highest energy level are said
to belong to the s block, comprising of Groups 1 and 2.
94
BILAL HAMEED
ELECTRONIC CONFIGURATION
114
PERIODIC TABLE
The elements that similarly fill up the p orbitals of the highest energy are
said to belong to the p block, which comprises of Groups 13 to 0.
So there are two columns in the s block and six columns in the p block.
The elements that fill up the d sub-level of the penultimate shell are
called the d block elements or the Transition Metal series.
95
SKILL CHECK 24
Write down the last term of the electronic configuration for the elements in the
table, the sequence for Hydrogen has been entered for you. Label the s, p and d
blocks on the table.
1s1
96
BILAL HAMEED
ELECTRONIC CONFIGURATION
12
12
17neutrons
20
35
44
20
18
22
10
m
a) Phosphorus: [Ne]…
16
electrons
peak
at z = 79.
b)
Cobalt:
[Ar]…
18 36
18
6 Write
the full electron
configuration
an
115
5 Complete
the
electronofconfigurations
of atoms
18
17
excited sodium atom.
ntify the two species from A−F that are
of the following elements:
12the same element.
10
7 Fill in the outer
electrons of a phosphorus
opes of
ELECTRONIC
CONFIGURATION
WS 1
a)
Phosphorus:
atom in the boxes below. [Ne]…
20 of A−F are
16neutral atoms.
ntify which
[Ar]…
3s b) Cobalt:
3p
ntify which
44 of A−F are
36cations.
ntify which of A−F are1 anions.
Complete the electronic configurations of atoms of the following elements:
22 of A−F have
18 the same
6 Write the full electron configuration of an
ntify which
8 Explain, with an example,
the meaning
of the
a) Phosphorus: [Ne].................
b) Cobalt:
[Ar].................
excited
sodium
atom.
term periodicity.
olainspecies
from A−F that are
why the relative atomic mass of
9 Consider
of boron,
magnesium
and of a phosphorus
7 Fillatoms
in the
outer
electrons
same
per
is notelement.
an exact whole number.
bromine.
In
which
is
the
effective
nuclear
charge
2of copper
Fill in the
outer electrons of a phosphorus
atom
in thebelow.
boxes below:
relative
atomic
is 63.5.
atom
in
the
boxes
of
A−F
are mass
neutral
atoms.
the largest and in which is it the smallest?
culate the relative abundance of the two
3s
of isotopes
A−F are
10 List the particles
Cl, Cl− and3p
K+ in order of
per
withcations.
relative isotopic masses
increasing radius.
3.0
65.0.
of and
A−F
are anions.
11 Explain why the first ionisation energies of the
of A−F have the same
ta about silicon in the table below were
noble
gases decrease
froman
helium
to krypton.
8 Explain,
with
example,
the meaning of the
uration.
d from a mass spectrometer.
12 The successive
ionisation energies of an element,
termofperiodicity.
3 The successive ionisation energies
an element, X, are given in the table. To which group
X, are given in the table. To which group of the
% abundance
the periodic
table does element X belong?
he relative atomic ofmass
of
periodic
table does element
9 Consider
atoms Xofbelong?
boron, magnesium and
92.2
tron configuration.
n exact whole
number.
4.7
omic mass3.1
of copper is 63.5.
elative
abundance
two
te the relative
atomic massofof the
silicon
to
cimal
place.
s with relative isotopic masses
e has two isotopes of mass numbers
0.
81. The mass spectrum of a sample of
e is shown below.
con in the table below were
ass spectrometer.
% abundance
92.2
−1
bromine. In
which
is themol
effective
nuclear charge
Ionisation
energy/kJ
1st the largest 1
000
and in which is it the smallest?
Ionisation
2nd
2 260
5th
6 990
3rd List the particles
3 390
10
Cl, Cl− and K+ in order of
4th
540
increasing 4radius.
8 490
11 Explain why
the first ionisation energies of the
7th
27 100
noble gases decrease from helium to krypton.
6th
8th
31 700
9th The successive
36 600 ionisation energies of an element,
12
10th
43 100
X, are given
in the table. To which group of the
13 Discuss
the relativetable
advantages
disadvantages
periodic
doesandelement
X belong?
...........................................................
of ESI and TOF mass spectrometers.
14 By considering
the relative effectiveness
of s, p energy/kJ mol−
Ionisation
Ionisation
and d electrons in shielding the nucleus, suggest
3.1
5
Write down the complete electronic
following elements:
1st configurations for the
1 000
why the difference between the atomic radii
158
160
162 m/z
of fluorine
2ndand chlorine is 50 pm whereas
2 260 that
Mn:
ive the
atomic
ofa) silicon
ntify
particlesmass
responsible
for.........................................................................................
the to
between chlorine and bromine is only 15 pm.
3rd
3 390
ks.
b) Mn2+: .........................................................................................
4.7
sotopes of mass numbers
ss spectrum of a sample of
below.
4th
Questions
5th
6 990
6th
8 490
7th
27 100
8th
31 700
9th
36 600
10th
43 100
c) Cu2+: ..........................................................................................
36.indd 35
4 540
1
35
27/02/2015 19:58
13 Discuss the relative advantages and disadvantages
of ESI and TOF mass spectrometers.
CEDAR COLLEGE
160
162
m/z
14 By considering the relative effectiveness of s, p
and d electrons in shielding
theCONFIGURATION
nucleus, suggest
ELECTRONIC
WS 1
why the difference between the atomic radii
of fluorine and chlorine is 50 pm whereas that
116
6
a) Explain what you understand by the word ‘orbital’ as applied to electrons.
...............................................................................................................................................................................
...............................................................................................................................................................................
b) Draw an s- and a p- orbital.
c) How does a 1s orbital differ from a 2s orbital?
...............................................................................................................................................................................
d) How many electrons are there in:
i. a 3d orbital? ..........................
ii. orbitals with a principal quantum number of 3? ..............................
CEDAR COLLEGE
ELECTRONIC CONFIGURATION WS 1
Multiple choice questions 1–10
1 Which one of the following has more neutrons
than electrons and more electrons than protons?
117
A 199F–
B
6 When sulfur, 16S is bom
two particles are form
One of them is a hydr
other is an element, X
S + 10n → 11H + X
32
16
Cl–
37
17
ELECTRONIC CONFIGURATION WS 2
1
1817
2745 enthalpies
11578of an element
14831
2 Successive
ionisation
–1
X in kJ mol are as follows:
578,is1817,
Which one of the following
X? 2745, 11 578, 14 831, 18 378
Which one of the following is X?
A boron
C aluminium
A boron
578
2
A
B
C
D
4
18378
P
B
32
15
C
33
16
D
33
15
S
P
7 Chlorine has two isoto
has two isotopes, 79Br
B carbon
How many lines woul
B carbon
D silicon
C aluminium
the mass spectrum of
D silicon
A 5
C 7is
Which one of the rows3 giving
aboutgiving
the fourth
periodabout
of the periodic table
Whichinformation
one of the rows
information
Use the key below to an
correct?
the fourth period of the periodic table is correct?
Total
number of
orbitals
3
Which one of the follo
A 32
S
15
C 94Be
Successive ionisation energies of an element X in kj/mol are as follows:
D 94Be2+
4
9
9
9
The
number of
different
types of
orbital
2
2
3
3
Maximum
number of
electrons in
the shell
8
18
8
18
A
B
1, 2 & 3 correct 1, 2 correct
8 Which of the followin
contain only one unpa
1 phosphorus, 15P
2 bromine, 35Br
3 aluminium, 13Al
9 Which of the followin
elements in the third
4 Chlorine exists as two isotopes, 35
Cl with an
17
37
1 They all have elect
abundance of 75.5% and 17Cl with an abundance
What is the order of increasing
energy of the orbitals within a single energy level?
different orbitals.
of 24.5%.
31
2 They all have som
Phosphorus has only one
A d>s<f<p
C pisotope,
< s < f <15P.d
electron structures
The mass spectrum of PCl3 has four lines
3
Only two of the el
B s<p<d<f
f < 142.
d<p<s
at m/z = 136, 138, 140Dand
structures with all
Which one these lines will have the smallest height?
pairs of electrons.
A 136
B 138
What is the condensed electronic
C 140 configuration for Co3+?
D 142
10 Which of the followin
121 6
number of neutrons a
5 Antimony has two isotopes
Sb and 123Sb.
C [Ar]3d
equal to 12?
The relative atomic mass of a naturally
occurring
B [Ar]4s23d4
D [Ar]4s13d5
1 three helium atom
sample of antimony is measured as 121.75.
2 two lithium ions (73
Which one of the following is the best
3 six hydrogen mole
approximate estimate of the percentage of 121Sb
Which are the values of the
successive
ionisation
energies
for
an
element
in
Group
14 of the
present in the naturally occurring sample?
11 Give the numbers of p
periodic table?
A 20%
and electrons present
C 60%
A 496, 738, 578, 789, 1012,
1000
C 1086, 2353, 4621, 6223, 37832, 47278
following atoms.
B 40%
a) 40
b
Ar
18
B 578, 1817, 2745, 11578, 14831, 18378
D 1314, 1000, 941, 869, 812
D 80%
Au
d) 197
e)
79
A [Ar]4s23d5
5
9781471827068_OCR_A_Level_Chemistry.indb 25
CEDAR COLLEGE
ELECTRONIC CONFIGURATION WS 2
What does 35Cl • contain?
CH3(CH2)3CH
CHCH
CHCH
CH(CH2)7CO2CH2
118
6
73
3
CH3(CH2)3CH
CHCH
CHCH
CH(CH2)7CO2CH
protons
neutrons
electrons
CHCH energies
CHCH ofCH(CH
CH
Which
are
for the
first16ionisation
consecutive
elements
in the same
2)3CH
2)7CO2CH
2
A
17 the values
183(CH
period?
B
17
18
17
is
so1000
that, on average,
one of
its side-chains
converted
A suitably
496, 738,hydrogenated
578, 789, 1012,
C 1086,
2353,
4621, 6223, is
37832,
47278into the
C (CH 18
17 CH=CH(CH
17 ) CO residue and two side-chains are converted into the
)
CH=CHCH
CH
3
2 4
2
2 7
2
)7CH=CH(CH
)
CO
residue.
CH
BD578,
2745, 11578,
14831,
18378
D 1314, 1000, 941, 869, 812
3(CH1817,
218
2
7
2
17
18
How many moles of hydrogen are required to convert one mole of glyceryl trieleostearate into the
soft margarine?
Use of the Data Booklet is relevant to this question.
A 4
B 5
C 6
D 9
In the gas phase, aluminium and a transition element require the same amount of energy to form
one mole of an ion with a 2+ charge.
Which isotope of an element in the third period of the Periodic Table contains the same number of
32
What is the
element?
neutrons
as transition
S?
16
A
A
23
Co
11 Na
B
B
24
Cr
Mg
C
C
D
D
12
28
Cu
14 Si
31
Ni
15 P
[W’02 3]
8
4
The successive ionisation energies, in kJ mol–1
3 , of an element X are given below.
4
870
1800
3000
3600
5800
7000
13 200
Unnilpentium is an artificial element. One of its isotopes is 262
105Unp.
9701/1/O/N/02
What is X?
Which of the following statements is correct?
95
AA
As
B 40Zr
C 52Te
262
33
105Unp has a nucleon number of 105.
B
262
The atom 260
105X is an isotope of 105Unp.
53I
[S’03 4]
3
9701/1/M/J/03
Which
of the
following
particles
would, on losing an electron, have a half-filled set of p orbitals?
C There
are
262 neutrons
in 262
105Unp.
D
A
1056
D
–
The
proton number
B of
N 262
C
105Unp is 262.C
+
N–
D
O
[S’04 5]
The table gives
theissuccessive
an element
Magnesium
oxide
used to lineionisation
industrialenergies
furnacesfor
because
it hasX.a very high melting point.
Which type of bond needs to be broken for magnesium oxide to melt?
1st
2nd
3rd
4th
5th
A co-ordinate
1800
2700
4800
6000
ionisation energy / kJ mol–1 950
B covalent
6th
12300
C
Whationic
could be the formula of the chloride of X ?
D
A
metallic
XCl
B
XCl2
C
XCl3
D
XCl4
[W’03 5]
7
6
7
Which solid exhibits more than one kind of chemical bonding?
Which molecule contains only six bonding electrons?
A brass
B C2F6
C H2O
D NF3
A C2H4
B copper
C diamond
Chemists have been interested in the properties of hydrogen selenide, H2Se, to compare it with
D
‘badice
egg’ gas hydrogen sulphide, H2S.
Which set of data would the hydrogen selenide molecule be expected to have?
CEDAR
COLLEGE
CONFIGURATION
WS 2
8 The
standard enthalpy changes of formation of iron(II)ELECTRONIC
oxide, FeO(s),
and aluminium oxide,
–1
–1
Al2O3(s), are –266 kJ mol and –1676 kJ mol respectively.
number of lone pairs
bond angle
Calculate the volume of nitrogen this will produce at room temperature.
A
114
4
9.2 dm3
B
13.9 dm3
C
119
3
2
27.7 dm3
72.0 dm3
D
What is the order of increasing energy of the listed orbitals in the atom of titanium?
Use of the Data Booklet is relevant to this question.
Section A
A
3s 3p 3d 4s
It is now thought that where an element exists as several isotopes, the stable ones usually
For each question there are four possible answers, A, B, C, and D. Choose the one you consider to
contain
a ‘magic
B
3s 3p
4s 3dnumber’ of neutrons. One of these magic numbers is 126.
be correct.
Which
is unstable?
C
3sisotope
4s 3p 3d
1
125
209 hydrocarbon is used
A
in bottled gasCfor 210
cooking
and heating.
A pure
D
4sBi3s 3p 3d B 208Pb
Po
D 208 Tl
[S’04 4]
When 10 cm3 of the hydrocarbon is burned in 70 cm3 of oxygen (an excess), the final gaseous
3
3
mixture contains 30 cm of carbon dioxide and 20 cm of unreacted oxygen. All gaseous volumes
An atom
has eight
electrons.
were
measured
under
identical conditions.
Which diagram shows the electronic configuration of this atom in its lowest energy state?
What
is the formula of the hydrocarbon?
! UCLES
A 2004
C 2H 6
B
A
2
C3H6
C 9701/01/M/J/04
C3H8
D
C4H10
On collision, airbags in cars inflate rapidly due to the production of nitrogen.
B
The nitrogen is formed according to the following equations.
2NaN3 → 2Na + 3N2
C
2
10Na + 2KNO3 → K2O + 5Na2O + N2
Section A
How many moles of nitrogen gas are produced from 1 mol of sodium azide, NaN3?
D
For each
question there are
A 1.5
B four
1.6 possible answers,
C 3.2A, B, C, and D.
D Choose
4.0 the one you consider to
be correct.
[W’04 5]
6
133
1
2
7
144
3
The gecko, a small lizard, can climb up a smooth glass window. The gecko has millions of
The
first six ionisation
energies
of each
four elements,
A to D, areofgiven.
microscopic
hairs ontetraethyl-lead(IV),
its
toes and
hair
has
pads at
tip. countries.
The result When
is that itthe
The
petrol additive
Pb(C
)4, is now banned
in its
many
is
2H5thousands
molecules in
the pads
arelead(II)
extremely
close
the
glass
surface
on
which
the
gecko
is
climbing.
and
H
O
are
formed.
completely
burned
in air,
oxide,
COto
2
Which element is most likely to be in Group2IV of the
Periodic Table?
Whatmany
is themoles
attraction
between
gecko’stotoe
pads
theofglass
How
of oxygen
arethe
required
burn
oneand
mole
Pb(Csurface?
2H5)4?
ionisation
A 9.5
co-ordinate bonds
2nd
5th
A
B 1st11
C3rd 13.5 4th
D 27 6th
energy / kJ mol−1
B covalent bonds
A
494
4560
6940
9540
13400
16600
C ionic
Which
ionbonds
has more electrons than protons and more protons than neutrons?
B
736
1450
7740
10500
13600
18000
D van der Waals’ forces
[H = 11 H 4610
; D = 12 H 6220
; O = 168 O]
C
1090
2350
37800
47000
– D
+ 2860
– 7480
–53200
4590
9400
B1400
H3inOthe
C OD
A
WhatDare the bond angles
PH3 molecule
likely to be? D OH
[S’05 3]
A 90 o
B 104 o
C 109 o
D 120 o
In
which
species
are
the
numbers
of
electrons
and
neutrons
equal?
What is the electronic configuration of an element with a second ionisation energy higher than
23
that of
9 each of its neighbours
A
B 19 F in the Periodic
C Table?
Na+
D 18 O2–
Be
4
9
2
2
6
11
8
2
A
1s 2s 2p 3s
B
1s22s22p63s23p1
C
1s22s22p63s23p2
D
1s22s22p63s23p3
[S’05 3]
© UCLES 2004
9701/01/O/N/04
4
Which compound has a boiling point which
is influenced by hydrogen bonding?
© UCLES 2005
9701/01/M/J/05
A CH3CHO
CEDAR COLLEGE
B CH3OCH3
C
HCO2H
[Turn over
ELECTRONIC CONFIGURATION WS 2
N2O4(g) + 2NaOH(aq) → NaNO3(aq) + NaNO2(aq) + H2O(l)
0
0
–3
Al of 0.02
S mol of N2O4?
Mg
Al
S
NaOH(aq) needed toMg
dispose
What is the minimum volume of 0.5 mol dm 120
B 12.5 cm3
C 40 cm3
D 80 cm3
A 8 cm3
1513 In 1999, researchers working in the USA believed that they had made a new element and that it
had the following electronic structure.
3
2 A sample of chlorine containing isotopes of mass numbers 35 and 37 was analysed in a
2
[Rn] 5f14Table
6d107shave
7p6.first ionisation energies which have the
2 mass-spectrometer.
Three successive elements in the Periodic
pattern shown in the diagram.
+
In which
group
of corresponding
the Periodic Table
you expect to find this element?
How
many
peaks
to Clwould
2 were recorded?
A
A
2II
B
B
3IV
C
C
4VI
x
D
D
50
[S’11 2 13]
x
first
14 Gallium
Consecutive
elements
X,
Y,
Z
are
in
period
of the of
Periodic
Element
Y hasonly
the ahighest
163
nitride, GaN, could
revolutionise the 3design
electricTable.
light bulbs
because
small
ionisation
2point.
first ionisation
and
the lowest
melting
length
used asenergy
a filament
gives
excellent
light
at
low
cost.
energy
Section
A Ga3+ ion.
What could
be is
theanidentities
of X, Y and
Z?
Gallium
nitride
ionic compound
containing
the
x
For each
questionmagnesium,
there are four
possible answers, A, B, C, and D. Choose the one you consider to
A sodium,
aluminium
What
is the electron
arrangement
of the nitrogen ion in gallium nitride?
be correct.
B magnesium, aluminium, silicon
2
A 1s2 2s2
C
aluminium,
silicon, phosphorus
atomic Anumber
1 B
Use 1s
of 2the
Section
2s2Data
2p3 Booklet is relevant to this question.
D silicon,
phosphorus,
sulphur
2
4
C
1s
2s2 be
2poxide,
, is possible
brilliantly
white and
of the
oxide produced
is used
in the
Titanium(IV)
TiO
What
could
the first
element
of thisanswers,
sequence?
2
For each
question
there
are
four
A, much
B, C, and
D. Choose
the one you
consider
to
manufacture
of
paint.
be correct.
2
2
6
D
A 1s
C 2s 2p
B N
C F
D Na
What is the maximum amount of TiO2 obtainable from 19.0 tonnes of the ore ilmenite, FeTiO3? [S’06 3]
1 Use of the Data Booklet is relevant201to this question.
A radioactive
10.0
B of 12.7
tonnes81
Cisquestion.
14.0 to
tonnes
D 17.7in
tonnes
used
assess damage
heart muscles after a heart
isotope
thallium,
Use
of thetonnes
Data
Booklet
is
relevant
toTl,
this
1743 A
When a sports medal with a total surface area of 150 cm2 was evenly coated with silver, using
attack.
The
electronic
of calcium,
krypton,
electrolysis,
itsstructures
mass increased
by 0.216
g. phosphorus and an element X are shown.
2 Carbon disulphide vapour201burns in oxygen according to the following equation.
Which statement about 81 Tl is correct?
Which
electronic
thatdeposited
of elementper
X?cm2 on the surface of the medal?
How many
atomsstructure
of silver is
were
© UCLES 2005
9701/01/O/N/05
[Turn over
→ CO2(g) + 2SO2(g)
CS2(g) + 3O2(g)
2 isotope
2
6
2 has
3 a nucleon number of 120.
A
This
183s 3p
2s
2p
A
1s
A 8.0 × 10
3
of 2carbon disulphide was burned in 50 cm3 of oxygen. After measuring the
A sample
of 610 cm
2 number
2
2 of 6electrons
B
The
one atom
this isotope
81. of aqueous sodium hydroxide
2s
2p
3s
3p
4s
1s
19
volume
gas remaining, the in
product
was of
treated
with anisexcess
B
1.8 of
× 10
and the2 volume
of
gas
measured
again.
All
measurements
were
2
2 of 6neutrons
6
2
C The
in one atom of this isotope is 201. made at the same temperature
1s
2snumber
2p2163s
3p 3d
4sconditions
and
pressure,
under
such
that carbon disulphide was gaseous.
C 1.2 × 10
2012
201
2
6
2
6
10
2
6
isotope of 4p
Tl.
D 1s
2s is2pan
82 X
223s 3p 3d 4s 81
D
× 10
What4.1
were
the
measured volumes?
[W’06 3]
182
volume of gas after
volume of gas
3
Use of the after
Databurning
Booklet/ cm
is relevant
to
this question.
adding
NaOH(aq) / cm3
In Aforming ionic compounds,
elements generally
30
0 form an ion with the electronic structure of a
noble gas.
B
30
20
Which
ion
does
not
have
a
noble
gas
electronic
© UCLESC2006
9701/01/M/J/06
50
20 structure?
A D I–
50B
Rb+
C
2+
Sn
40
Sr2+
D
[W’07 2]
1933
The
first pair
stage
the atoms
manufacture
of nitric
acidonly
is the
oxidation
by oxygen.
In which
doinboth
have one
electron
in an
s orbitalofinammonia
their ground
states?
A
Ca, Sc
B
Cu,wNH
Be 3(g) + xO
C2(g)H,→He
D2O(g)
Li, Cr
yNO(g) + zH
[W’08 3]
4
Which values for w, x, y and z are needed to balance the equation?
Use of the Data Booklet is relevant to this question.
w
x
y
z
Hard water contains calcium ions and hydrogencarbonate ions arising from dissolved calcium
hydrogencarbonate,
A
4
5 Ca(HCO
4 3)2 .
6
How
in the
B many4 electrons
6 are present
4
5 hydrogencarbonate anion?
© UCLES 2006
A C 30
[Turn over
9701/01/O/N/06
5
CEDAR COLLEGE
D
6
6B
31 5
4
5
6
4
C
32
D
33
ELECTRONIC CONFIGURATION WS 2
–3
mol dm
For C
each0.875
question
there
are four possible answers, A, B, C, and D. Choose the one you consider to
be correct.
D 1.40 mol dm–3
121
1
3
20
Use of the Data Booklet is relevant to this question.
The first seven ionisation energies of an element between lithium and neon in the Periodic Table
are
as follows.
Lead(IV)
chloride will oxidise bromide ions to bromine. The Pb4+ ions are reduced to Pb2+ ions in
this reaction.
1310
3390
5320
7450
11 000
13 300
kJ mol–1
71 000
If 6.980 g of lead(IV) chloride is added to an excess of sodium bromide solution, what mass of
What
is the
outer
configuration of the element?
bromine
would
beelectronic
produced?
2
2s
0.799 g
A
A
2
2s
2p1g
1.598
B
B
C
C
2
2s
2p4g
3.196
2
D
2p6g
D 2s
6.392
[S’09 1 3]
212
Which element has an equal number of electron pairs and of unpaired electrons within orbitals of
principal quantum number 2?
A
beryllium
B
carbon
C
nitrogen
D
oxygen
[W’11 1 3]
3
elements,
Y andisZ,relevant
have the
physical
properties shown in the table.
223 4 Three
Use of
the Data X,
Booklet
to this
question.
© UCLES 2009
9701/01/M/J/09
Which graph represents the melting
numberpoint
of unpaired
p orbital electrons
boiling point
density for atoms with proton
numbers 13 to 18? element
/ g cm-3
/ °C
/ °C
X
A
Y
–7
59
3.12
98
883
0.97
1107
number of
unpaired
electrons
1.74
Z
649
number of
unpaired
What
could be the identities of X, Y and Z?
electrons
X
Y
Z
A
Br2
Al
Si
B
Na
Mg
C
Br2
0
I2 13
D
I2
14Mg 15 16 Na17
proton number
Si
K
0
18
13
14
B
15 16 17
proton number
C
18
D
number of
unpaired
electrons
number of
unpaired
electrons
© UCLES 2011
9701/13/O/N/11
0
13
14
15 16 17
proton number
18
0
13
14
15 16 17
proton number
18
[S’09 1 4]
5
Which statement explains why the boiling point of methane is higher than that of neon?
[Ar: H, 1; C, 12; Ne, 20]
CEDAR COLLEGE
ELECTRONIC CONFIGURATION WS 2
A
A molecule of methane has a greater mass than a molecule of neon.
B
A
C
4
23
3
cryolite is a base; aluminium oxide is amphoteric
B 1:2
C 1:4
1:1
cryolite reacts with the aluminium oxide to form ions
D
2:1
122
D
aluminium
oxide
alone would
conduct electricity
Use molten
of the Data
Booklet
is relevant
to thisnot
question.
In which set do all species contain the same number of electrons?
A 10 cm3 sample of 0.30 mol dm–3 Tl +NO3– required 20 cm3 of 0.10 mol dm–3 acidified NH4VO3 to
4+
in solution.
Vanadium is the only element reduced in this reaction.
oxidise
it ,toCo
Tl3+3+, Co
A Co2+
–
Cl –
B F–is, Br
What
the,oxidation
number of the vanadium in the reduced form?
+1
Na+, Mg2+, Al 3+ B
A
C
4
+2
C
+3
D
+4
D K2SO4, K2SeO4, K2TeO4
Use of the Data Booklet is relevant to this question.
[W’13 2 4]
Element
X forms X– ions that can 9701/11/O/N/13
be oxidised to element X by acidified potassium
24
© UCLES
2013
manganate(VII).
What could be the values of the first four ionisation energies of X?
1st
2nd
3rd
A
418
3070
4600
D
1010
1840
2040
2 4th
5860
Section
A
B
577
1820
2740
11 600
For each question there are four possible answers, A, B, C, and D. Choose the one you consider to
C
590
1150
4940
6480
be correct.
1
4030
2
Use of the Data Booklet is relevant to this question.
25
Section A
Atoms of element X have six unpaired electrons.
[W’13 3 4]
For each
are X?
four possible answers, A, B, C, and D. Choose the one you consider to
Whatquestion
could bethere
element
be correct.
A carbon
Use of the Data Booklet may be appropriate for some questions.
© UCLES
9701/13/O/N/13
B 2013
chromium
1
C iron
Which
compound contains two different elements with identical oxidation states?
D selenium
A HCl O
B
Mg(OH)2
C
Na2SO4
D
NH4Cl
[S’14 1 1]
2
26
2
Use of the Data Booklet is relevant to this question.
For the element sulfur, which pair of ionisation energies has the largest difference between them?
a black,
non-metallic
AIodine
thirdisand
fourth shiny,
ionisation
energies solid and a member of Group VII. It sublimes easily on
heating to give a purple vapour.
B fourth and fifth ionisation energies
A sample of iodine vapour of mass 6.35 g has a volume of 1.247 dm3 when maintained at
Cconstant
fifth and
sixth ionisation
temperature
and aenergies
pressure of 1.00 × 105 Pa.
D sixth and seventh ionisation energies
If iodine vapour acts as an ideal gas, what is the temperature of the iodine vapour?
[M’16 2 2]
3
B 600 K
C 300 000 K
D 600 000 K
A 300 K
Which ion has both more electrons than protons and more protons than neutrons?
[H = 11 H; D = 12 H; O = 168 O]
3
Enthalpy
changes that are difficult
to measure directly
can often be– determined using Hess’ Law
+
Ato construct
D–
B Hcycle.
C OD–
D OH
3O
an enthalpy
4
Which enthalpy change is indicated by X in the enthalpy cycle shown?
Which species contains the smallest number of electrons?
C(s) + 2H– 2(g) + 2O2(g)
B Be2+
C H
D He+
A B3+
X
CEDAR COLLEGE
CH4(g) + 2O2(g)
ELECTRONIC CONFIGURATION WS 2
CO2(g)
+
2H2O(l)
4
The relative atomic mass of copper is 63.5.
3
123
Which chart is a correct mass spectrum that would
lead to this value?
3
2
The diagram shows the pathways of a reaction, with and without a catalyst.
3
27
A same group of the Periodic Table.
Elements X and Y are in the
B
Which letter represents the overall energy change for the reaction?
100%
The table shows the first six ionisation energiesabundance
of X and Y in kJ mol–1.
abundance
1st
2
3rd 3
2nd
4th
33%
6th
5th
17%
X
800
1600
2400
4300
5400
10 400
B
62 63
In the ideal gas 61
equation,
pV = 64
nRT,65
what66
are the units of n and61
T ? 62 63 64
A6000
Y
1000
1800 m / e 2700
4800
D 12 300
energy
reactants
n
T
What could be the identities of X and Y?
C °C
A
no units
Y
B
no X
units
K
abundance
74%
A
antimony,
arsenic,
C
mol Sb
°C As
B
D
arsenic,
mol As
C
abundance
26%
D
products
33%
33%
22%
extent of reaction
antimony,
K Sb
65 66
m/e
12%
selenium,
61 Se62
tellurium,
63
64 Te
65 66
61 62 63 64 65 66
2+ for m
2–
g of
salt Cr
(sodium
chloride) per day
reasons.
People
are advised
eat
lessdichromate(
than
e VI
/e
m /6.00
D reaction
tellurium,
Te toacidified
selenium,
Se
The
between
) ions,
ions health
results
in the
2O7 , and aqueous Fe
3+
dichromate(VI) ions being reduced to Cr ions.
[S’16 2 3]
C
33
50%
Which mass of sodium is present in 6.00 g of sodium chloride?
5
28
4
4
Which
gaseous
atom
total
of five
occupying in
spherically
orbitals?
What
isisolated
the
equation
thisa reaction?
H12O6, is measured
mmol / l . shaped
In Pakistan,
the
In
China,
thecorrect
concentration
offorhas
blood
glucose,
C6electrons
2.36isg measured C
D 3.88 g
A 0.261 g of bloodBglucose
/ dl .g
concentration
in mg3.64
A
A boron
Cr2O72– + Fe2+ + 14H+ → 2Cr3+ + Fe3+ + 7H2O
3
3
The unit l is
2+dm ). The
+ unit dl is
3+ a decilitre
3+ (0.1 dm ).
2– a litre (1
B
fluorine
B
Cr
O
+
2Fe
+
14H
→
2Cr
+
2Fe
+
7H
O
7
2
When 2copper
reacts with a 50% solution of nitric acid,
nitrogen monoxide is evolved and a blue
2+
3+
3+
2–
l indicates
health problem.
A
blood
glucose
concentration
solution
results.
C
sodium
+ 3Fe
+ 14H+ of
→18.5
2Crmmol
+ /3Fe
+ 7Ha2O
C
Cr2O
7
+ to mg / dl
3+
2– mmol / l2+converted
D
What
is2O
18.5
D
Cr
+ equation
6Fe + 14H
→reaction
2Cr3+ ?+ is6Fe
+ 7H2O
Thepotassium
balanced
for this
shown.
7
[S’16 2 5]
64
30
5
B 178 mg / dl
C 333 mg / dl
D 3330 mg / dl
A 33.3 mg / dl
p Cu + q HNO3 → r Cu(NO3)2 + s H2O + t NO
Carbon
and
silicon
have
the
same
outer
electronic
structure.
Sodium azide, NaN3 is an explosive used to inflate airbags in cars when they crash. It consists of
positive
sodium
ions and
negative
azide
? gaseous.
What are
the
of in
the
integers
p, ions.
q, r,isolated
s and tand
Each
fourvalues
species
this
question
Why isofathe
Si–Si
bond
weaker
than
a C–Care
bond?
What are the numbers of electrons in the sodium ion and the azide ion?
Which
species
is not
A Silicon
have
a larger
patoms
q planar?
r atomics radius than
t carbon atoms.
+
sodium
ion B CH
azide
ion
B ABF
Silicon
nuclear
charge
H
A
3
31
1has a greater
4
2 CthanC2carbon.
24
D
NH3
10
CA
first220
ionisation
BSilicon2has a smaller
6
3 energy2than carbon.
6
Argon
is a gas
usedmetallic
to fill electric
light bulbs.
22 carbon.
D B Silicon
is10
more
than
C
2
8
2
4
4
C which 12
20
Under
conditions of pressure
and temperature will argon behave most like an ideal gas?
D
3
D
12
pressure
8
3
4
2
22
temperature
2
What
electronic configuration
of an isolated Ni2+ ion?
A 68 is thehigh
high
The Ge isotope is medically useful because it undergoes a natural radioactive process to give
68
A to detect tumours. This transformation
anBisotope
of a6 different
element,
X, whichSection
can be used
2
2p 3s23p63d64slow
A 68 1s22s2high
of Ge occurs when an electron enters the nucleus and changes a proton into a neutron.
For each
C question
low there are four
highpossible answers, A, B, C and D. Choose the one you consider to
B 1s22s22p63s23p63d84s2
be correct.
Which statement about the composition of an atom of 68X is correct?
D
low
low
C 1s22s22p63s23p63d104s2
It has
4 electrons
in itsbe
outer
p orbitals.for some questions.
Use A
of the
Data
Booklet may
appropriate
© UCLES 2016 2
9701/13/M/J/16
[Turn over
D 1s 2s22p63s23p63d8
31
5
5
B
32
1
It has 13 electrons in its outer shell.
Which
ion37
has
the same electronic configuration as Cl –?
C It has
neutrons.
© UCLES 2016
D
A
Its
P+
F– proton numberBis 32.
© UCLES 2016
[Turn over
9701/12/M/J/16
3+
C
Sc
9701/12/O/N/16
4+
D
Si
[Turn over
CEDAR
COLLEGE J and K each contain 40% carbon by mass.
ELECTRONIC CONFIGURATION WS 2
2 Compounds
What
© UCLES
2016could J and K be?
9701/11/O/N/16
[Turn over
For each question there are four possible answers, A, B, C and D. Choose the one you consider to
9 correct.
Oxidation numbers should be used to answer this question.
be
+
A the
redox
reaction
takes
between hydroxylammonium
ions, [NH3OH] , and acidified iron(III)
Use of
Data
Booklet
mayplace
be appropriate
for some
124 questions.
ions, Fe3+. The products are iron(II) ions, Fe2+, H+ ions, water and a compound of nitrogen.
1
33
The question
mole ratiorefers
of reacting
hydroxylammonium
This
to isolated
gaseous atoms.ions to reacting iron(III) ions is 1 : 2.
Which
compound
In
whichnitrogen-containing
atom are all electrons
paired?could be formed in the reaction?
NH3
AA Ba
N2O
BB Br
C
C
NO
S
D
NO2
Si
[S’18 1 Q1]
Element
X has a higher
first ionisation
energy
than element
Y.
210 Which
compound
has a boiling
point that
is influenced
by hydrogen
bonding?
34
B CH
C factor
HCO2that
CH3 helps D
HCO2H
ATwoCH
students
they
believe
is one
to explain
this.
3CHO state what
3OCH
3
2
student 1 “X has a higher first ionisation energy than Y because an atom of X has more
3 Which fuel would produce
largest
mass
protonsthe
in its
nucleus
thanofanCO
atom
of Y.”10 kg of the fuel undergo complete
2 when
Section A
combustion?
student 2 “X has a higher first ionisation energy than Y because X has a smaller atomic
For A
eachbiodiesel,
questionCthere
are
four
possible
answers, A, B, C and D. Choose the one you consider to
O
than
Y.”
17H34radius
2
be correct.
BOnlyethanol,
C2Htwo
6O students is correct.
one of the
Use of the Data Booklet may be appropriate for some questions.
C octane, C8H18
What could X and Y be?
D propane, C3H8
1 Which feature is present in both ethene and poly(ethene)?
X
Y
4
A A bond
angles
of 109°
The
diagram
shows
the Boltzmann
carbon
boron distribution of energies in a gas. The gas can take part in a
reaction with an activation energy, Ea. The gas is maintained at a constant temperature.
B B covalent
bonds
magnesium
aluminium
C C covalent
bonds
oxygen
nitrogen
oxygen
D D sp3 orbitals
sulfur
P
proportion of
molecules with
211 The
electronic
of an atomofofesters.
sulfur is 1s22s22p63s23p4.
35
Hydrogen
ions configuration
catalyse the hydrolysis
a given energy
[S’18 1 Q10]
How
many
valence
shell and unpaired electrons are present in one sulfur atom?
Which
statement
is correct?
A
The hydrogen ions act as a heterogeneous
catalyst.
0
valence shell
unpaired
0
Ea
electronsions are electrons
B The hydrogen
in the same phase as the reactants.
molecular energy
CA The hydrogen
ions
are
used
up
in
the
reaction.
2
1
Which statement is correct?
DB The hydrogen
ions have no2effect on the activation energy of the reaction.
4
A If a catalyst is added, peak P will be lower and Ea will move to the left.
C
6
0
B If a catalyst is added, peak P will be lower and Ea will move to the right.
D
6
2
C If a catalyst is added, peak P will be the same and Ea will move to the left.
3
the right.
D
If a catalyst
is added,
peak substance
P will be the
same
and Eboiling
a will move
In which
pair does
the second
have
a lower
pointtothan
the first substance?
A
C2H6 and C2H5Cl
© UCLES
B 2018
CH3OCH3 and C2H5OH
9701/11/M/J/18
C 2018
Ne and Ar
© UCLES
9701/11/M/J/18
D
4
[S’18 2 Q2]
CH3NH2 and C2H6
Compound J burns in excess oxygen to give carbon dioxide and water only. When a 3.00 g
sample of compound J is burnt in excess oxygen, 4.40 g of carbon dioxide and 1.80 g of water are
formed.
What is the empirical formula of J?
CEDAR COLLEGE
A CH
B
CHO
C
CH2
ELECTRONIC CONFIGURATION WS 2
D CH2O
3
36
C
0
not deflected
D
1
not deflected
125
The table refers to the electron distribution in the second shell of an atom with eight protons.
Which row is correct for this atom?
orbital shape
orbital shape
orbital type
number of
electrons
orbital type
number of
electrons
A
p
2
s
4
B
p
4
s
2
C
s
2
p
4
D
s
4
p
2
[S’18 3 Q3]
© UCLES 2018
CEDAR COLLEGE
9701/13/M/J/18
ELECTRONIC CONFIGURATION WS 2
correct
correct
correct
correct
126
No other combination of statements is used as
a correct response.
14
11 B
SECTION
Section
B equation.
31 The gas laws can be summarised in the
ideal gas
Section
For each of the questions in this section, one or
more ofB the three numbered statements 1 to 3 may
pV = nRT
be correct.
For each of the questions in this section, one or more of the three numbered statements 1 to 3 may
where
each
symbol
its usualismeaning.
be correct.
Decide
whether
each
of thehas
statements
or is not correct
11 (you may find it helpful to put a tick against
the statements that you consider to be correct).
Which
statements
arestatements
correct? is or is not correct (you may find it helpful to put a tick against
Decide
whether
each of the
Section B
the
statements
that
you
consider
to be correct).
The responses A to D should be selected
on the basis of
1 One mole of an ideal gas occupies the same volume under the same conditions of
For each of the questions in this section, one or more of the three numbered statements 1 to 3 may
D
2
The density of an ideal gas at constant pressure is inversely proportional to the temperature,
A
B 2
C 3 (you may find
Dit helpful to put a tick against
Decide whether
is or is not
correct
1 only
2 and
1 and
1,
T. 2 and 3each of the statements
is
are
only
are
the statements
to2be correct). only
1, 2are
andthat
3 you consider
1 and
2 and 3
1 only
3 The
volume of a givencorrect
mass of an ideal gas
is doubled if its temperature
is raised from 25 °C
correct
correct
correct
are
only are
only are
is
tocorrect
50 °CAattoconstant
The responses
D shouldpressure.
be selected on the basis
correct
correctof
correct
temperature
and pressure.
The responses
A to D should
be selected on the basis of
be correct. A
B
C
No other combination of statements is used as a correct response.
A
B
C
No
statements
is used
a correct
response.
32 other
Usecombination
of the Data of
Booklet
is relevant
toasthis
question.
D
1, 2 and
3
1 and
2
and 3
1 only
31 The
diagram
illustrates
energy
changes
of a set
of2 reactions.
do boththe
species
have
the same
number
of unpaired p electrons?
1 In which pairs
are
only
are
only
are
is
31 How may nitrogen exist in compounds?
correct+
correct
correct
correct
_
1 O and Cl
H = _134 kJ mol 1
1 bonded by a tripleRcovalent bond
S
+
2
Fcombination
and Ga– of statements is used as a correct response.
No2other
as part of a cation
+
P andlost
Ne 3 electrons to form an anion
33 having
[S’12 3 32]
31 A space shuttle’s upward thrust came from the following reaction between aluminium and
1
ammonium perchlorate.
H = +92 J mol
2 Use of the Data Booklet is relevant to this question.
32
10Al + 6NH4Cl O4 → 4Al 2O3 + 2Al Cl 3 + 12H2O + 3N2
The isotope 99Tc is radioactive and has been found in lobsters and seaweed adjacent to nuclear
fuel reprocessing plants.
Which statements about this reaction
are_correct?
H = _75 kJ mol 1 99
Which statements are Tcorrect about an atom of Tc?
U
1 Aluminium is oxidised.
1 It has 13 more neutrons than protons.
Which
of the following
statements are correct?
2
Chlorine
is reduced.
2 It has 43 protons.
Nitrogen
is change
oxidised.
1 3 The
enthalpy
for the transformation U ⎯→ R is + 42 kJ mol–1 .
3 It has 99 nucleons.
© UCLES
2012 enthalpy change for the transformation
9701/11/M/J/12
2 The
T ⎯→ S is endothermic.
32 3 Use
of enthalpy
the Datachange
Bookletforisthe
relevant
to this question.
The
transformation
R ⎯→ T is – 33 kJ mol–1 .
[Turn over
[S’07 31]
33 Which of these substances have a giant structure?
3
Which statements are correct when referring to the atoms 23Na and 24Mg?
1 silicon(IV) oxide
1
2
They have the same number of full electron orbitals.
baked clay found in crockery
2
They have the same number of neutrons.
phosphorus(V) oxide
3
3
They are both reducing agents.
[S’13 3 32]
© UCLES 2004
9701/01/O/N/04
© UCLES 2012
9701/12/O/N/12
CEDAR COLLEGE
[Turn over
ELECTRONIC CONFIGURATION WS 2
3
39 +
19 K
127
32 Use of the Data Booklet is relevant to this question.
4
Carbon and nitrogen are adjacent in the Periodic Table.
Which properties do they both have?
1
There is an empty 2p orbital in one atom of the element.
2
The principal quantum number of the highest occupied orbital is 2.
3
They form compounds in which their atoms form bonds with four other atoms.
[S’13 1 32]
33 What are necessary properties of a dynamic equilibrium?
1
Equal amounts of reactants and products are present.
2
Concentrations of reactants and products remain constant.
3
The rate of the forward reaction is the same as the rate of the reverse reaction.
© UCLES 2013
CEDAR COLLEGE
9701/11/O/N/13
[Turn over
ELECTRONIC CONFIGURATION WS 2
128
2
Answer all the questions.
ELECTRONIC CONFIGURATION WS 3
11 Electrons are arranged in energy levels.
(a) An orbital is a region in which an electron may be found.
Draw diagrams to show the shape of an s orbital and of a p orbital.
s orbital
p orbital
[2]
(b) Complete the table below to show how many electrons completely fill each of the following.
number of electrons
a d orbital
a p sub-shell
the third shell (n = 3)
[3]
(c) The energy diagram below is for the eight electrons in an oxygen atom. The diagram is
incomplete as it only shows the two electrons in the 1s level.
energy
1s
Complete the diagram for the oxygen atom by:
(i)
adding labels for the other sub-shell levels,
[1]
(ii) adding arrows to show how the other electrons are arranged.
[1]
© OCR 2007
CEDAR COLLEGE
ELECTRONIC CONFIGURATION WS 3
129
3
(d) Successive ionisation energies provide evidence for the arrangement of electrons in atoms.
Table 1.1 shows the eight successive ionisation energies of oxygen.
Table 1.1
ionisation number
1st
2nd
3rd
4th
5th
6th
7th
8th
ionisation
energy / kJ mol–1
1314
3388
5301
7469
10 989
13 327
71 337
84 080
(i)
Define the term first ionisation energy.
...........................................................................................................................................
...........................................................................................................................................
...........................................................................................................................................
..................................................................................................................................... [3]
(ii)
Write an equation, with state symbols, to represent the third ionisation energy of
oxygen.
..................................................................................................................................... [2]
(iii)
Explain how the information in Table 1.1 provides evidence for two electron shells in
oxygen.
...........................................................................................................................................
...........................................................................................................................................
...........................................................................................................................................
..................................................................................................................................... [2]
[Total: 14]
© OCR 2007
CEDAR COLLEGE
[Turn over
ELECTRONIC CONFIGURATION WS 3
2
Answer all the questions
in the spaces provided.
130
1 3 In the 19th and 20th centuries, scientists established the atomic theory and showed that
three sub-atomic particles, electron, neutron and proton, exist. The masses and charges of
these three particles were subsequently determined.
When separate beams of electrons, neutrons or protons are passed through an electric field
in the apparatus below, they behave differently.
+
–
beam of particles
(a) (i)
Which of these three particles will be deflected the most by the electric field?
.........................................
(ii)
In which direction will this particle be deflected?
..................................................................................................................................
(iii)
Explain your answer.
..................................................................................................................................
(b) (i)
..................................................................................................................................
[4]
Define the term proton number.
..................................................................................................................................
..................................................................................................................................
(ii)
Why is the proton number of an atom of an element usually different from the
nucleon number of an atom of the element?
..................................................................................................................................
..................................................................................................................................
[2]
© UCLES 2006
CEDAR COLLEGE
9701/02/O/N/06
ELECTRONIC CONFIGURATION WS 3
For
Examiner’s
Use
131
3
(c) Protons and neutrons have been used in nuclear reactions which result in the formation
of artificial elements. In such processes, protons or neutrons are accelerated to high
speeds and then fired like ‘bullets’ at the nucleus of an atom of an element.
Suggest why neutrons are more effective than protons as ‘nuclear bullets’.
..........................................................................................................................................
..................................................................................................................................... [2]
(d) In some cases, when neutrons are fired at atoms of an element, the neutrons become
part of the nucleus of those atoms.
What effect does the presence of an extra neutron have on the chemical properties of
the new atoms formed? Explain your answer.
..........................................................................................................................................
..........................................................................................................................................
..................................................................................................................................... [2]
[Total: 10]
© UCLES
2006
CEDAR
COLLEGE
9701/02/O/N/06
[Turn over
ELECTRONIC CONFIGURATION
WS 3
For
Examiner’s
Use
132
4
CEDAR COLLEGE
ELECTRONIC CONFIGURATION WS 3
2
Answer all the questions
in the spaces provided.
133
15 (a) Explain what is meant by the term ionisation energy.
....................................................................................................................................................
....................................................................................................................................................
.............................................................................................................................................. [3]
(b) The first seven ionisation energies of an element, A, in kJ mol–1, are
1012
1903
2912
4957
6274
21 269
25 398.
(i) State the group of the Periodic Table to which A is most likely to belong. Explain your
answer.
.............................................................................................................................................
.............................................................................................................................................
.............................................................................................................................................
....................................................................................................................................... [2]
(ii) Complete the electronic configuration of the element in Period 2 that is in the same group
as A.
1s2 ................................................................................................................................. [1]
(c) Another element, Z, in the same period of the Periodic Table as A, reacts with chlorine to
form a compound with empirical formula ZCl 2. The percentage composition by mass of ZCl 2
is Z, 31.13; Cl , 68.87.
(i) Define the term relative atomic mass.
.............................................................................................................................................
.............................................................................................................................................
....................................................................................................................................... [2]
(ii) Calculate the relative atomic mass, Ar, of Z.
Give your answer to three significant figures.
Ar of Z = ....................... [2]
© UCLES 2014
CEDAR COLLEGE
9701/21/M/J/14
ELECTRONIC CONFIGURATION WS 3
134
NOTES
CEDAR COLLEGE
ELECTRONIC CONFIGURATION WS 3
135
9 The Periodic Table: chemical periodicity
This topic illustrates the regular patterns in some physical properties of the elements
in the Periodic Table.
9.1 Periodicity of physical properties of the elements in the third period
PERIODIC TRENDS
BILAL HAMEED
PERIODIC TRENDS
136
Inorganic chemistry
9 The Periodic Table: chemical periodicity
This topic illustrates the regular patterns in chemical and physical properties of the elements in the
Periodic Table.
Learning outcomes
Candidates should be able to:
9.1 Periodicity of
physical properties
of the elements in
the third period
a) describe qualitatively (and indicate the periodicity in) the variations in
atomic radius, ionic radius, melting point and electrical conductivity of
the elements (see the Data Booklet )
b) explain qualitatively the variation in atomic radius and ionic radius
c) interpret the variation in melting point and electrical conductivity in terms
of the presence of simple molecular, giant molecular or metallic bonding
in the elements
d) explain the variation in first ionisation energy (see the Data Booklet)
e) explain the strength, high melting point and electrical insulating
properties of ceramics in terms of their giant structure; to include
magnesium oxide, aluminium oxide and silicon dioxide
9.2 Periodicity of
chemical properties
of the elements in
the third period
a) describe the reactions, if any, of the elements with oxygen (to give
Na2O, MgO, Al 2O3, P4O10, SO2, SO3), chlorine (to give NaCl , MgCl 2,
Al 2Cl 6, SiCl 4, PCl 5) and water (Na and Mg only)
b) state and explain the variation in oxidation number of the oxides (sodium
to sulfur only) and chlorides (sodium to phosphorus only) in terms of
their valence shell electrons
c) describe the reactions of the oxides with water
(treatment of peroxides and superoxides is not required)
d) describe and explain the acid/base behaviour of oxides and hydroxides
including, where relevant, amphoteric behaviour in reaction with acids
and bases (sodium hydroxide only)
e) describe and explain the reactions of the chlorides with water
f)
interpret the variations and trends in 9.2(b), (c), (d) and (e) in terms of
bonding and electronegativity
g) suggest the types of chemical bonding present in chlorides and oxides
from observations of their chemical and physical properties
9.3 Chemical periodicity
of other elements
a) predict the characteristic properties of an element in a given Group by
using knowledge of chemical periodicity
b) deduce the nature, possible position in the Periodic Table and identity of
unknown elements from given information about physical and chemical
properties
Back to contents page
BILAL HAMEED
www.cie.org.uk/alevel
PERIODIC TRENDS
29
137
PERIODICITY
The outer electronic configuration is a periodic function, it repeats ever so often.
Many physical and chemical properties are influenced by the outer shell
configuration of an atom and hence also exhibit periodicity, such as:
atomic radius, ionic radius, ionisation energy, electron affinity, electronegativity,
electrical conductivity, melting point and boiling point
The first two periods in the periodic table are not typical...
Period 1
(H, He) contains only two elements
Period 2 elements at the top of each group have small sizes and high I.E.values
Period 3 (Na-Ar) is the most suitable period for studying trends
1
ATOMIC RADIUS
The atomic radius is basically used to describe the size of an atom.
Larger the atomic radius, larger the atom.
Atomic radius increases down a group.
This is because, as we go down a group in the periodic table the atoms
have increasingly more electron shells.
In each new period the outer-shell electrons enter a new energy level
and so are located further away from the nucleus.
2
BILAL HAMEED
PERIODIC TRENDS
charge. The 18 core electrons partially cancel this
positive charge and the effective nuclear charge is
approximately 1:
4s valence electron
▲ Figure 3 Shielding of the outer valence electron in the
Zeff ≈ Z - S = 19 - 18 = 138
1
potassium atom
Periodic trends in atomic radius
Across a period from left to right, atomic radii decrease. This is because of
the increasing effective nuclear charge, Zeff, going from left to right across
the period. This pulls the valence (outer-shell) electrons closer to the
nucleus, reducing the atomic radius.
ATOMIC RADIUS
Down a group from top to bottom, atomic radii increase. In each new
Atomic radius decreases across a period
period the outer-shell electrons enter a new energy level so are located
further away from the nucleus. This has a greater effect than the
charge, Z,
because of shielding
by the core electrons.
The magnitude of the positive chargeincreasing
of thenuclear
nucleus
increases,
its “pull”
These trends are summarized in figures 4 and 5. Figure 5 shows that the
on all of the electrons increases, hence
the
ofelements
shells
remain
the across a period.
atomic
radii number
of the transition
do not
change greatly
The reason for this is that the number of electrons in the outermost energy
same and the electrons are drawn closer
toprincipal
the nucleus.
level of the
quantum number, n, remains almost constant across
the period. As electrons are added they enter the (n - 1) rather than the
nth energy level. So the number of valence electrons and hence Zeff remain
essentially constant, resulting in little variation in atomic radius.
atomic radii increase down a group
This results in a contraction of the atomic radius and therefore a
decrease in atomic size.
3
atomic radii decrease across a period
▲ Figure 4 Trends in atomic radii. Some people think of these shapes as snowmen – going down
a group the snowman is standing upright, while across a period the snowman is sleeping!
78
3.2 PERIODIC TRENDS
ATOMIC RADIUS
20
300
23
24
Cs 8
Fr 2
21
radius (pm)
K 0
19
2
Ra 1
200
2
Rb15
Ba06
20
Sr 0
19
La 4
Ac 1
1
Na60
Li 0
174
Ca
1
M 40
176
Sc 9
Y
16
164
Hf
15
Ta 8
H32
9
g
Be9
15
Zr 4
148
Ti
1
Nb56
15
W 0
14
V 4
1
Mo46
141
Re
100
0
13
1
2
3
4
5
6
13
Cr 0
138
Tc
13
Os 6
1
Mn29
1
Ru36
12
Fe 4
118
12
84
Co
11
1
B
Al 4
Ni 7 Cu22
13
75
11
Pd 0 A 136 14 Zn120 123
Si 4 1 C
132
71
g C 0 1 Ga
09
N
Ir
120
d
42
130
P
6
104
In
12
14 Ge
Pt
13
O4
Sn 0 140 As 0 11 S
6
100
Au 0 132 144
F0
8
Sb
Hg Ti
10
1
117 Cl
137 Se
Ar 1 Ne62
Pb45 150
B
T
e
116
136
r
Bi
K
I
13
r
14
Xe 6
Po 2 A 148
t
1
Rn46
7
8
9
decr
10
easin
11
g at
omic
radii
1
R 34
h
12
13
14
15
3
He7
ii
ad
cr
mi
o
t
ga
sin
rea
c
n
i
16
17
18
4 found in section 9 of
▲ Figure 5 Values of atomic radii of elements in pm. These data can be
the Data booklet
Periodic trends in ionic radius
The radii of cations and anions vary from the parent atoms from which
they are formed in the following way.
The radii of cations are smaller than those of their parent atoms; for
the atomic radius of K is 200 pm while the ionic radius of
BILAL HAMEED example,
K is 138 pm. The reason for this is that there are more protons than
+
Ions
electrons in the cation so the valence electrons are more strongly
attracted to the nucleus.
An ion is a charged species.
Ions are either cations or
anions:
The radii of anions are larger than those of their parent atoms; for
• A cation is positively
PERIODIC TRENDS
139
ATOMIC RADIUS MEASUREMENTS
The radius of an atom is determined by the distance between the center of the
nucleus and the boundary of the region where the valence electrons have a
probability of being located.
Since the probability of finding an electron in an atom extends up to infinity it is
impossible to determine10
the exact size of the atom.
Hence atomic radius is taken as half the distance between the centre of two
adjacent atoms in a molecule or in a closed packed system.
Na
Mg
Al
The bond between two Melting
adjacent
atoms may be covalent,
metallic
or simply660
van
point/°C
98
649
der Waal’s forces.
736
577
First ionisation energy/kJ mol−1 494
Electronegativity
Electrical conductivity/S cm−1
Atomic radius/nm
5
Si
P
1410
4
786
1
0.9
1.2
1.5
1.8
2
2.1 × 105
2.3 × 105
3.8 × 105
2.5 × 10−6
1
0.186
0.160
0.143
0.117
0
Table 10.1 Some physical properties of the
elements in the third period
Interatomic and interio
The term ‘atomic radius’ does not mea
we can fairly easily measure the distan
element (see Figure 10.3), if we are to
with another, we must be aware that t
two elements. For example, tables of
sodium: its covalent radius is 0.154 nm
der Waals’ radius is 0.230 nm. The onl
radius (0.190
Chemists use X-ray diffraction and other techniques to measure
thenm), since it does not fo
sizes of the atoms in a sensible way, w
distance between the nuclei of atoms.
Comparing the van der Waals’ radii of
atom (0.230 nm) is larger than the argo
The atomic radius of an atom cannot be defined precisely
because it
the nuclear attraction for the outer ele
depends on the type of bonding and on the number of bonds.
across the group.
ATOMIC RADIUS
Metallic radius in metals is half the distance between the inter-nuclear
Figure 10.3 The atomic radius is half the
a
distance of what are internuclear
effectively
ions.
distance.
half this distance
b
r
cation
6
on the left – metallic radius
BILAL HAMEED
on the rig
In
crossing
the third
period of the Peri
10_03
Cam/Chem
AS&A2
PERIODIC
TRENDS
11 to 18. This means that the 3s and 3p
Barking
Dog Artand as a result their dist
to
the nucleus
This is the reason why the atomic radii
irst ionisation energy/kJ mol
494
736
577
lectronegativity
0.9
1.2
is 0.154 nm;
1.5 sodium:
1.8its covalent
2.1 radius 2.5
3.0 its metallic
— radius is 0.186 nm, and its van
786
1060
1000
1260
1520
lectrical conductivity/S cm−1
2.1 × 105
2.3 × 105
is ×0.230
for argon is its van der Waals
~0 radius listed
~0
3.8 ×der
105 Waals’
2.5 × radius
10−6
1.0
10−11 nm.
5.0 ×The
10−18only
tomic radius/nm
0.186
0.160
radius (0.190
nm),0.110
since it does
0.143
0.117140
0.104 not form
0.099compounds.
(0.190) If we are to compare the
sizes of the atoms in a sensible way, we need to compare radii of the same type.
Comparing the van der Waals’ radii of the two elements shows that the sodium
Interatomic
and
interionic
radii
atom
(0.230
nm) is larger
than the argon atom (0.190 nm), as we might expect, sin
The term ‘atomicthe
radius’
does not
mean thefor
same
for electron
every element.
nuclear
attraction
thething
outer
shell Although
increases with proton number
we can fairly easily
measure
the
distance
between
two
adjacent
nuclei
in
an
across the group.
ble 10.1 Some physical properties of the
ments in the third period
ATOMIC
RADIUS MEASUREMENTS
element (see Figure 10.3), if we are to use those distances to compare one element
with another, we must be aware that the interatomic bonding might differ in the
•
radius
is
between
the nuclei
of
atoms
joined
Foradistance
example, tables
of interatomic
radii include
three values
for by a
Figure 10.3 TheCovalent
atomic radius
is two
half elements.
thehalf the
b
sodium: its covalent radius is 0.154 nm; its metallic radius is 0.186 nm, and its van
internuclear distance.
covalent bond.
The values are measured by X-ray or electron diffraction.
c
der Waals’ radius is 0.230 nm. The only radius listed for argon is its van der Waals’
radius (0.190 nm), since it doeshalf
notthis
form
compounds. If we are to compare the
r distance
sizes of the atoms in a sensible way, we need to compare radii of the same type.
r
Comparing the van der Waals’ radii of the two elements shows that the sodium
chlorine moleculeatom (0.230 nm) is larger than the argon atom (0.190 nm), as we might expect, since
the nuclear attraction for the outer electron shell increases with proton number
across the group.
r
molecules (two
• Van der Waal’s radius is half the distancebbetween two monoatomic
c
gure 10.3 The atomic radius is half the
a
ernuclear distance.
atoms held by van der Waal’s forces)
on the left – metallic radius
r
for Group 18 –
van der Waals’ radius
on the right – covalent radius
r
half this distance
r
argon molecule
In
crossing
the third
period of the Periodic Table, the proton number increases from
10_03
Cam/Chem
AS&A2
11 to 18. This means that the 3s and 3p outer electrons become more firmly attracte
Barking
Dog Art
7 and as a result their distance away from the nucleus becomes less.
to
the nucleus
This is the reason why the atomic radii decrease from sodium to chlorine (see Figur
10.4).
the somewhat
larger
for Group
18 – value of 0.190 nm listed for argon
on the left – metallic
radiusAs mentioned
on the right – above,
covalent radius
vanwhich
der Waals’
radiusthe distance between the two
Table 10.1 is the van der Waals’ radius,
is half
nuclei when the atoms touch one another in the solid state.
In
crossing
the third
period of the Periodic Table, the proton number increases from
10_03
Cam/Chem
AS&A2
Sodium, magnesium and aluminium are metals and form Na+, Mg2+ and Al3+ ions.
11 to 18. This means that the 3s and 3p outer electrons become more firmly attracted
6
Barking
Dog Artand
These
ions their
all have
theaway
electronic
1s2 2s2 2p
. As the proton number increas
to
the nucleus
as a result
distance
from thestructure
nucleus becomes
less.
11 atomic
with sodium
to 13 from
withsodium
aluminium,
the (see
attraction
This is the reasonfrom
why the
radii decrease
to chlorine
Figure for the outer electrons
10.4). As mentioned
above, the
larger value
of 0.190
nm listed
for argon
in This means that the ioni
increases
andsomewhat
they become
drawn
in closer
to the
nucleus.
+
2+
3+ two
Table 10.1 is the van
der
Waals’
radius,
which
is
half
the
distance
between
the
radii get smaller (Na 0.095 nm, Mg 0.065 nm, Al 0.050 nm).
nuclei when the atoms touch one another in the solid state.
Sodium, magnesium and aluminium are metals and form Na+, Mg2+ and Al3+ ions.
These ions all have the electronic structure 1s2 2s2 2p6. As the proton number increases
from 11 with sodium
13 with aluminium,
theelement
attraction that
for the
outer
electrons
Thetomelting
point of an
has
a giant
structure is high because many
increases and they
become drawn
in closer
to be
the nucleus.
This
means thattothe
ionicplace. The atomic radii
interatomic
bonds
must
broken
for
melting
take
radii get smaller (Na+ 0.095 nm, Mg2+ 0.065 nm, Al3+ 0.050 nm).
ATOMIC RADIUS
Melting points
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become smaller from sodium to silicon and so the bonding becomes stronger
and
the
melting
points
become
higher,
as
the
bonding
electrons
are
closer
to the
CV Melting points
nuclei.
The melting pointadjacent
of an element
that has a giant structure is high because many
interatomic bonds must be broken for melting to take place. The atomic radii
B\
become smaller from sodium to silicon and so the bonding becomes stronger
and the melting points6>become higher, as the bonding electrons are closer to the
adjacent nuclei.
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18
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The reason that atomic radius decreases across a period is basically the
BILAL HAMEED
same
reason electronegativity and ionisation energy increase: an PERIODIC
increase TRENDS
in nuclear charge across the period but no significant increase in shielding.
141
SKILL CHECK 1
Put the following elements in order of increasing atomic radius. Justify
your answers:
A Mg, S, Si
B Mg, K, Al
C Si, Cl, K
9
SKILL CHECK 2
Explain why the ionic radius of a Group 2 ion is smaller than the
atomic radius of the corresponding Group 2 atom.
10
BILAL HAMEED
PERIODIC TRENDS
142
IONIC RADIUS
Positive ions (cations) are smaller than the parent atom.
The cation has more protons than electrons (an increased nuclear charge).
The excess nuclear charge pulls the remaining electrons closer to the
nucleus.
Also, cation formation often results in the loss of all outer-shell electrons,
resulting in a significant decrease in radius.
11
IONIC RADIUS
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The variation of ionic radius across a period is not a clear-cut trend, as
the type of ion changes from one side to the other. Thus positive ions
are
BILAL HAMEED
PERIODIC TRENDS
formed on the left-hand side of the period and negative ions on the righthand side.
143
IONIC RADIUS
Negative ions (anions) are larger than the parent atom.
The anion has more electrons than protons.
Owing to the excess negative charge, the nuclear “pull” on each
individual electron is reduced. The electrons are held less tightly,
resulting in a larger anion radius in contrast to the neutral atom.
13
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The variation of ionic radius across a period
the type of ion changes from one side to the o
formed on the left-hand side of the period and
hand side.
For positive ions there is a decrease in ionic
the ion increases, but for negative ions the size
increases (Figure 4.12).
Let us consider Na+ and Mg2+: both ions ha
configuration, but Mg2+ has one more proton
PERIODIC TRENDS
4.13). Because there is the same number of ele
amount of electron–electron repulsion is the sa
2+
V#
IONIC RADIUS ACROSS A PERIOD
Across a period the radii of ions with the same electron configuration
decreases as nuclear charge increases.
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hand side.
For positive ions there is a decrease in ionic radius as the charge on
the ion increases, but for negative ions the size increases as the charge
144
increases (Figure 4.12).
Let us consider Na+ and Mg2+: both ions have the same electronic
configuration, but Mg2+ has one more proton in the nucleus (Figure
4.13). Because there is the same number of electrons in both ions, the
amount of electron–electron repulsion is the same; however, the higher
nuclear charge in Mg2+ means that the electrons are pulled in more
strongly and so the ionic radius is smaller.
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7
8
9
10
11
12
13
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10
10
10
10
10
10
10
16
BILAL HAMEED
PERIODIC TRENDS
145
SKILL CHECK 3
Explain why the ionic radii of halides are larger than atomic radii of the
corresponding halogens
17
SKILL CHECK 4
Which diagram represents the change in ionic radius of the elements
across the third period (Na to Cl)?
18
BILAL HAMEED
PERIODIC TRENDS
146
SKILL CHECK 5
The species Ar, K+ and Ca2+ are isoelectronic (have the same number of
electrons).
In what order do their radii increase?
19
WHAT IS IONISATION ENERGY?
Ionisation Energy is the amount of energy needed to remove electrons from atoms.
The ionisation energy of an atom measures how strongly an atom holds its
electrons.
Attraction between the nucleus and an electron
As electrons are negatively charged and protons in the nucleus are positively
charged, there will be an attraction between them. The greater the pull of the
nucleus, the harder it will be to pull an electron away from an atom.
20
BILAL HAMEED
PERIODIC TRENDS
147
WHAT AFFECTS IONISATION ENERGY?
The value of the 1st Ionisation Energy depends on the electronic structure
Lithium
Helium
Hydrogen
2+
+
1310 kJ mol-1
2370 kJ mol-1
3+
519 kJ mol-1
Helium has two protons in the nucleus. The nuclear charge is greater so the
pull on the outer electrons is larger. More energy will be needed to pull an
electron out of the atom.
21
WHAT AFFECTS IONISATION ENERGY?
Lithium
Helium
Hydrogen
2+
+
1310 kJ mol-1
2370 kJ mol-1
3+
519 kJ mol-1
Lithium atoms have 3 protons so you would expect the pull on electrons to be
greater. However, the 1st I.E. of Li is lower than that of He because
1. Filled inner shells exert a shielding effect - lowers the effective nuclear pull and,
2. Electrons are further away from the nucleus- lowers effective nuclear attraction.
22
BILAL HAMEED
PERIODIC TRENDS
148
FACTORS INFLUENCING IONISATION ENERGY
Nuclear Charge
As nuclear charge increases the attractive force would increase and hence
more energy would be needed to overcome the increasing nuclear
attraction.
Atomic Radius
As the atomic radius increases, the outermost electron would be further
away from the nucleus experiencing a weaker attractive force. Hence the
value decreases with increasing size.
Attraction decreases very rapidly with distance. An electron close to the
nucleus will be much more strongly attracted than one further away.
23
FACTORS INFLUENCING IONISATION ENERGY
Shielding effect/Screening
Electrons in full inner shells repel electrons in outer shells. Full inner shells of
electrons prevent the full nuclear charge being felt by the outer electrons.
The inner electron orbitals effectively screen or shield the outermost electrons
from the nucleus, due to which the ionisation energy decreases.
The lessening of the pull of the nucleus by the inner electrons is known as
shielding.
It is due to the shielding effect that the electrons in the outer shell are attracted
by the effective nuclear charge which is less than the full charge on the
nucleus.
24
BILAL HAMEED
PERIODIC TRENDS
&.
of the outer electron is not simply
in an orbit around the nucleus as shown,
are three full shells of electrons between the outermost electron and the
and this is why the effective nuclear charge felt by the outer electron
nucleus, and if this shielding were perfect the effective nuclear charge
:miZch^dc
felt by the outer electron would be +1 (19+ in nucleus, 18 shielding in potassium is greater than 1. There are various ways of estimating or
calculating
the aeffpoint
ectivecharge
nuclear
for a particular electron in an atom
In however,
electrostatics
sphere
of charge
behaves like
at charge
its centre;
electrons). This shielding is not perfect,
and athe
effective
nuclear
(e.g. Slater’s
rules).
Calculations
suggest that the effective nuclear charge felt
to the outer electron,
spheres
of charge
inside
(the
charge felt by the outermost electrontherefore
is higherrelative
than +1.
149
@
the outer
electron
in potassium
about 3.5+.
electron shells) behave as if their by
charge
is at the
nucleus.
The chargeis felt
&.
by the outer electron is thus 19+ + 18− = 1+ acting at the nucleus.
;^\jgZ'#() I]ZdjiZgZaZXigdc^cV
An alternative view of shielding is that The
the outer
electron
is attracted
electrons
do not
form perfect
of electron
charge, and
movement
Oncespheres
the ediVhh^jb^dc^hh]^ZaYZY[gdbi]Z[jaa
first
hasthe
been
removed, the next electron is
by the nucleus but repelled by the inner
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of the
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around the nucleus as shown,
ViigVXi^kZ[dgXZd[i]ZcjXaZjhWni]Z^ccZg
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and this is why the effective nuclear charge h]Zaahd[ZaZXigdchh]VYZY^cWajZ#
felt by the outer electron
and second ionisation energies). This is consistent with the electron being
in potassium is greater than 1. There are various ways of estimating or
removed
from
a new electron
main energy
level (shell). This electron is closer
:miZch^dc
calculating the effective nuclear charge for a particular
in an atom
to the that
nucleus
and
therefore
more
strongly
(e.g. Slater’s rules). Calculations suggest
the eff
ective
nuclear
charge
felt attracted. It is also shielded
In electrostatics a sphere of charge behaves like a point charge at its centre;
byisfewer
electrons
(the ten electrons in the inner main energy
levels), as
@
by
the
outer
electron
in
potassium
about
3.5+.
therefore relative to the outer electron, spheres of charge inside (the electrons in the same shell do not shield each other very well
&. (they do not
@ @
&. &.
electron shells) behave as if their charge is at the nucleus. The charge felt
get between the electron and the nucleus).
by the outer electron is thus 19+ + 18
− = 1+
at the nucleus.
Once
theacting
first electron
has been removed,
the nextenergy
electron
is rises steadily as the electrons are removed
The ionisation
now
The electrons do not form perfect considerably
spheres of charge,
more and
diffithe
cultmovement
to remove
(there
is
a
large
jump
between
firstlevel. There is no significant change
urthest from the nucleus and =A
successively from the same main energy
of the outer electron is not simply in and
an orbit
around
the nucleus
as shown,
second
ionisation
energies).
This
is consistent
with
the electron
being
eus. This electron is thus
in
shielding,
but
as
the
positive
charge
on the ion increases it becomes more
and this is why the effective nuclear charge
feltfrom
by the
outermain
electron
.i]ZaZXigdc
.i]ZaZXigdc
removed
a new
energy level (shell). This electron is closer
ned) from
the full attractive
in potassium
is greater than 1. There are various ways of estimating or difficult to remove a negatively charged electron (less electron–electron
to
the
nucleus
and
therefore
more
strongly
attracted.
It
is
also
shielded
ns in the
atom (Figure
2.34).
repulsion, so the electrons are pulled in closer to the nucleus).
calculating
the eff
ective nuclear charge for a particular electron in an atom
by fewer electrons (the ten electronsThere
in theisinner
main
energy
as
nucleus
andSlater’s
a particular
another
large
jumplevels),
in ionisation
energies between the ninth
(e.g.
rules). Calculations suggest that the effective nuclear charge felt
@ @
@ @
electrons in@ the same shell do notand
shield
each
other
very
well
(they
do
not last to be removed
ucleus for
that
electron.
There
the tenth, as the ninth
electron is the
@
by the outer electron in potassium is about 3.5+.
&. &.
&. &. from the
&. the electron and the nucleus).
get between
&.
outermost electron and the
third main energy level
but the tenth is the first to be removed from the
The ionisation energy now rises
steadily
as the
are removed
e effective nuclear charge
second
level.
Theelectrons
tenth electron
is significantly closer to the nucleus and is
Once 18
theshielding
first electron has been removed,
the next
electron
is main energy level. There is no significant change
successively
from
the same
in nucleus,
less shielded than the ninth electron to be removed.
&%i]ZaZXigdc
&%i]ZaZXigdc
considerably more difficult to removein(there
is a large
between
first on the ion increases it becomes more
shielding,
but asjump
the positive
charge
ever, and the effective nuclear
.i]ZaZXigdc
.i]ZaZXigdc
and second ionisation energies). Thisdiffi
is consistent
with
the
electron
being
cult to remove a negatively charged electron (less electron–electron
her than +1.
'6IDB>8HIGJ8IJG:
removed from a new main energy repulsion,
level (shell).
This
electronareis pulled
closer in closer to the nucleus).
so the
electrons
to the nucleus and therefore more strongly
attracted.
It
is
also
shielded
There is another large jump in ionisation energies between the ninth
;^\jgZ'#() I]ZdjiZgZaZXigdc^cV
outer electron
attracted (the ten electrons
@ @
by feweris electrons
in the
main
energy
as is the last to be removed from the
and
theinner
tenth,
as the
ninthlevels),
electron
&. &.
ctrons. electrons in the same shell doediVhh^jb^dc^hh]^ZaYZY[gdbi]Z[jaa
@
not shield
verylevel
wellbut
(they
not is the first to @be removed
ViigVXi^kZ[dgXZd[i]ZcjXaZjhWni]Z^ccZg
thirdeach
mainother
energy
thedo
tenth
from
the
&. &.
get between the electron andh]Zaahd[ZaZXigdchh]VYZY^cWajZ#
the nucleus).
second level. The tenth electron is significantly closer to the nucleus and is
The ionisation energy now rises steadily
as the electrons
removed
less shielded
than the are
ninth
electron to be removed.
&%i]ZaZXigdc
&%i]ZaZXigdc
successively from the same main energy level. There is no significant change
in shielding,
butitsascentre;
the positive charge on the ion increases it becomes more
ike a point
charge at
,*
'6IDB>8HIGJ8IJG:
.i]ZaZXigdc
.i]ZaZXigdc
difficultinside
to remove
res of charge
(the a negatively charged electron (less electron–electron
25
repulsion,
the electrons
are pulled in closer to the nucleus).
the nucleus.
Thesocharge
felt
There
is another large jump in ionisation energies between the ninth
+ acting at
the nucleus.
@ @
and the
as the ninth electron is the last to be removed from the
of charge,
andtenth,
the movement
&. &.
t around
themain
nucleus
as shown,
third
energy
level but the tenth is the first to be removed from the
elt by the
outerlevel.
electron
second
The tenth electron is significantly closer to the nucleus and is
ous ways
estimating
lessofshielded
thanorthe ninth electron to be removed.
&%i]ZaZXigdc
&%i]ZaZXigdc
particular electron in an atom
he effective nuclear charge felt
,*
'6IDB>8HIGJ8IJG:
@
.5+.
FACTORS INFLUENCING IONISATION ENERGY
&.
he next electron is
is a large jump between first
istent with the electron being
shell). This electron is closer
ttracted. It is also shielded
nner main energy levels), as
h other very well (they do not
the electrons are removed
There is no significant change
ion increases it becomes more
ron (less electron–electron
to the nucleus).
energies between the ninth
t to be removed from the
first to be removed from the
tly closer to the nucleus and is
moved.
FACTORS INFLUENCING IONISATION ENERGY
@ @
&. &.
Nature of sub level
Electrons in the s orbital are closer to the nucleus as compared to the
.i]ZaZXigdc
.i]ZaZXigdc
electrons in the p orbital of the same shell.
@ @
&. &. the s electrons are the more firmly held and require a greater amount
Hence
of energy to be removed as compared to the p electrons which are slightly
&%i]ZaZXigdc
&%i]ZaZXigdc
further.
'6IDB>8HIGJ8IJG:
,*
Also, the p orbital of the same energy level will experience the shielding
effect of the s electrons from the same shell in addition to previous s orbitals.
Note: Only S Filled Orbitals Provide Shielding
26
BILAL HAMEED
PERIODIC TRENDS
,*
150
FACTORS INFLUENCING IONISATION ENERGY
Stability of Certain Configurations
Two electrons in the same orbital experience a bit of repulsion from each
other. This offsets the attraction of the nucleus so that the paired
electrons are removed rather more easily.
Completely filled sub levels are more stable than others requiring large
amounts of energy for their disruption. This additional stability is
associated with half filled sub levels also.
Stability of
completely filled
sub levels
>
Stability of halffilled sub levels
>
Stability of other
configurations
27
IONISATION ENERGY DOWN A GROUP
A^ energy decreases down a group.
Ionisation
;^ghi^dc^hVi^dcZcZg\n$`?bdaÄ&
&+%%
&)%%
CV
ZaZXigdcXadhZg
idi]ZcjXaZjh
Despite an increased nuclear charge the outer
electron is easier to remove. This is due to
&%%%
@
GW and increased
greater distance from the nucleus
-%%
8h
shielding.
A^
(
&'%%
+%%
Down the group, as the number of shells
%
<gdje&
increase, the outer electrons are farther away
;^\jgZ)#* ;^ghi^dc^hVi^dcZcZg\n[dg\gdje&#
from the nucleus.
)%%
@
&.
Potassium has a lower first
ionisation energy than
Down any group in the periodic table the first ionisation
energy
lithium.
decreases.
28
Although the nuclear charge also increases down a group, this is largely
balanced out by an increase in shielding down the group, as there are
more electron energy levels (shells). It is the increase in size that governs
the change in first ionisation energy.
KVg^Vi^dc^cÒ
ghi^dc^hVi^dcZcZg\nVXgdhhVeZg^dY
BILAL
HAMEED
The general trend is that first ionisation energy increases from left
;^\jgZ)#+ EdiVhh^jb]VhVadlZgÒghi
^dc^hVi^dcZcZg\ni]Vca^i]^jb#
PERIODIC
TRENDS
The increase
in first ionisation
energy (Figure 4.7) can also be
explained in terms of the effective
151
IONISATION ENERGY DOWN A GROUP
Also, down the group each successive element contains more electrons
in the shells between the nucleus and the outermost electrons which
increases the shielding effect.
This increased shielding causes the outermost electrons to be held less
tightly to the nucleus.
The increased shielding effect and increased atomic radius/size
outweigh the increase in nuclear charge. Therefore ionisation energy
decreases.
29
IONISATION ENERGY ACROSS A PERIOD
General Trend: Ionisation energy increases across a Period
Across a period, the number of protons and the number of electrons increase by
one each.
The additional proton increases the nuclear charge.
The additional electron is added to the same outer shell in each of the elements.
A higher nuclear charge more strongly attracts the outer electrons in the same
shell, but the electron-shielding effect from inner-level electrons remains the same.
Thus, more energy is required to remove an electron because the attractive force
on them is higher.
30
BILAL HAMEED
PERIODIC TRENDS
The nuclear charge increases from Na (11+) to Ar (18+) as protons are
added to the nucleus. The electrons are all removed from the same main
energy level (third shell) and electrons152
in the same energy level do not
shield each other very well. Therefore the force on the outer electrons due
to the nucleus increases from left to right across the period and the outer
IONISATION ENERGY ACROSS A PERIOD
;^ghi^dc^hVi^dcZcZg\n$`?bda
Ä&
&+%%
6g
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8>
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Level stud
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-%%
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+%%
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outer elec
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i.e. 1+ if s
The effec
felt by the
argon ato
charge) −
electrons)
perfect.
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%
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EZg^dY(
31
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SKILL CHECK 6
For each of the following pairs, state which element has the higher
first ionisation energy and explain your answer:
A Mg and Al
B Mg and Ca
C Ne and Na
32
BILAL HAMEED
PERIODIC TRENDS
153
1st IONISATION ENERGY / kJmol-1
1 ST IONISATION ENERGIES OF PERIOD 3
Along a Period as electrons are being added on to the same shell the
magnitude of the ionisation energy increases due to the increase in the
nuclear charge, a decrease in atomic size and no extra shielding.
3s
3p
P
3s
3s
There is a DROP in the value for Aluminum.
This is because the extra electron has gone
into one of the 3p orbitals. The increased
shielding from the 3s makes the electron
easier to remove. And the 3p is further away
from the nucleus.
3p
Si
Mg
Al
3s
3s
3p
Na
ATOMIC NUMBER
33
1 ST IONISATION ENERGIES OF PERIOD 3
1st IONISATION ENERGY / kJmol-1
3s
3p
Ar
3s
3s
3p
3p
Cl
P
3s
3s
3p
S
3s
Si
Mg
Also, Phosphorus, has a more stable
half-filled 3p configuration.
Al
3s
3s
3p
There is a DROP in the value for
Sulfur. The extra electron has paired
up with one of the electrons already
in one of the 3p orbitals. The
repulsive force between the two
paired-up electrons means that less
energy is required to remove one of
them.
3p
Na
ATOMIC NUMBER
34
BILAL HAMEED
PERIODIC TRENDS
154
Periodic pr
IONISATION ENERGIES OF ELEMENTS
2500
Figure 7.7 !
Periodicity in the first ionisation e
of the elements.
First ionisation energy/kJ mol –1
Ne
2000
F
1500
Ar
N
H
1000
O
C
Be
P
Mg
B
500
0
Cl
S
Si
Ca
Al
Li
Na
1
group 0
5
group 1
K
10
Atomic number
group 2
15
20
35
13 Use data on the CD-ROM to explore whether there is any pattern
in the boiling points of the chlorides of elements in Periods 2 and 3.
14 Why do you think the first ionisation energies of elements decrease
with atomic number in every group of the periodic table?
Data
Periodicity
atomic radius
Atomic radii
Ionisation energies
Atomic radii measure the size of atoms in crystals and molecules. Chemists use
X-ray diffraction and other techniques to measure the distance between the
nuclei of atoms. The atomic radius of an atom cannot be defined precisely
because it depends on the type of bonding and on the number of bonds.
2500
Atomic
radii for metals are calculated from the distances between atoms in
He
metal crystals (metallic radii). The atomic radii for non-metals are calculated
from the lengths of covalent
bonds in crystals or molecules (covalent radii).
Ne
2000
Atomic radii decrease from left to right across a period. Across the period Na
to Ar, atomic radii
fall from 0.191 nm for sodium to 0.099 nm for chlorine. From
F
one element to the next across a period
the charge on the nucleus increases by
Ar
1500
N
one as the number of electrons in the same outer shell increases by
Kr one.
Shielding
by electrons
in the same
shell is limited, so the ‘effective nuclear
Cl
H
O
charge’ increasesC and the electrons
are drawn more tightly to the nucleus.
P
The first ionisation energies generally increase from sodium to argon as the
proton number increases. In two instances, this increase does not take place:
between magnesium and aluminium, and between phosphorus and sulfur (see
Figure 10.6).
IONISATION ENERGIES OF ELEMENTS
first IE/kJ mol–1
first ionisation
r
on and
er melting points
(see Figure 10.5).
g and structure in
He
1000
Mg
Si
B
Al
Li
Na
metal crystal
AtomicKradii decrease
Rb
0
0
2
atomic radius
S
Be
500
non-metallic molecule
4
6
8
10
12
14
16
18
20
22
24
26
28
30
32
34
36
proton number
38
40
Figure 7.8 "
Atomic radii and the internuclear
in a molecule and a crystal.
36
Atomic radii increase
Worked example
Suggest why:
Cam/Chem
AS&A2
a10_06
the ionisation
energy
of aluminium is less than that of magnesium
b the ionisation energy of sulfur is less than that of phosphorus.
Test yourself
15 Arrange these elements in
atomic radius: Al, B, C, K a
a This is because the outer electron in aluminium is in the 3p orbital, and so further from
germanium to be
16
Which atom or ion in each
the nucleus and less strongly held than the outer electron in magnesium, which is in the
conductor or an
3s orbital.
pairs has the larger radius?
7.9 the
" electron that is being removed from sulfur comes from a
b ThisFigure
is because
ng point of
doubly-occupied orbital, in which it suffers inter-electron repulsion from the other electron
a) Cl or Cl–
ber 34) is higher
Periodicity of atomic radii in the periodic table.
BILAL HAMEED
PERIODIC TRENDS
occupying that orbital.
roton number 33).
b) Al or N
Barking Dog Art
Answer
The reactions of the elements with oxygen,
chlorine and water
155
SKILL CHECK 7
Use of the Data Booklet is relevant to this question.
The elements radon (Rn), francium (Fr) and radium (Ra) have consecutive proton numbers
in the Periodic Table.
What is the order of their first ionisation energies?
37
3
a What do you understand by the term atomic orbital?
b Draw diagrams to show the shape of:
i an s orbital
ii a p orbital.
c Element X has the electronic configuration 1s2 2s2 2p6 3s2 3p6 3d8 4s2.
i Which block in the Periodic Table does element X belong to?
ii State the maximum number of electrons in a d sub-shell.
iii Element X forms an ion of type X2+.
Write the full electronic configuration for this ion using 1s2 notation.
iv Write the symbol for the sub-shell which begins to fill after the 3d and 4s are completely full.
SKILL CHECK 8
The 1st ionisation energies of several
4 Thelements
e 1st ionisation energies of several elements with consecutive atomic numbers are shown in the g
a. Which of the elements A to I belong to Group I in
the Periodic Table? Explain your answer.
b. Which of the elements A to I could have the
electronic configuration 1s2 2s2 2p6 3s2?
c. Explain the rise in 1st ionisation energy between
element E and element G.
d. Estimate the 1st ionisation energy of element J.
First ionisation energy / kJ mol–1
below.
The letters
with consecutive atomic numbers are
shown
in are not the symbols of the elements.
the graph below. The letters are not the symbols
2000
of the elements.
1600
1200
1800
400
0
A
B
C
D E F
Element
G H
I
J
a Which of the elements A to I belong to Group I in the Periodic Table? Explain your answer.
b Which of the elements A to I could have the electronic configuration 1s2 2s2 2p6 3s2?
c Explain the38rise in 1st ionisation energy between element E and element G.
d Estimate the 1st ionisation energy of element J.
BILAL HAMEED
PERIODIC TRENDS
B 578, 1817, 2745, 11 578, 14 831, 18 378
two decimal places.
(2)
C 1086, 2353, 4621, 6223, 37 832, 47 278
c) The mass spectrum of bromine has lines at
D 1314, 1000, 941, 869, 812
(1)
m values of 158, 160 and 162, but none at
ii) Which are the values for the first
z
159 or 161. What causes the line at 160?
ionisation energies of consecutive 156
A (80Br−80Br)+
elements in the same period?
B (80Br−80Br)−
A 496, 738, 578, 789, 1012, 1000
C (79Br−81Br)−
B 578, 1817, 2745, 11 578, 14 831,
D (79Br−81Br)+
(1)
18 378
d) State and outline one modern use of mass
C 1086, 2353, 4621, 6223, 37 832,
spectrometry.
(3)
47 278
(Total 9 marks)
D 1314, 1000, 941, 869, 812
(1)
iii) Which are
values for thesodium
first
3 Theisfirst
ionisationbelow.
of elements sodium to
The first ionisation
oftheelements
to argon
shown
ionisation energies of elements in the
argon is shown below.
same group, as the group is descended?
A 496, 738, 578, 789, 1012, 1000
A Explain why
the 1817,
general
sodium
B 578,
2745,trend
11 578,from
14 831,
18 378
to argon is upwards but why the value for
C 1086, 2353, 4621, 6223, 37 832,
sulfur is less than
that for phosphorus.
47 278
D 1314, 1000, 941, 869, 812
(1)
c) on
i) the
Define
the term
electron
affinity.
B Mark
graph
where
the value
for (2)
ii) Which is the equation that relates to
potassium would
be.
the first electron affinity of chlorine?
First ionisation energy
SKILL CHECK 9
Na
Mg
Al
Si
P
S
Cl
Ar
K
A 12 Cl2(g) + e− → Cl−(g)
−(g)
a) Explain
why
the general
→ Cl
Cl(g)
+ e− for
C Explain whyB the
value
the
second ionisation of sodium
is very
much
larger trend
than from
that sodium
of its
−
−
to
argon
is
upwards
but
why
the
value for
→
Cl
(g)
C
Cl(g)
−
e
first ionisation.
sulfur is less than that for phosphorus.
(5)
(1)
D Cl2(g) + 2e− → 2Cl−(g)
b) Mark on the graph where the value for
(Total 8 marks)
potassium would be.
(1)
2 This question is about mass spectrometry.
c) Explain why the value for the second
a) A mass spectrometer can be used to find39
the
ionisation of sodium is very much larger
percentage composition of the isotopes of
than that of its first ionisation.
(2)
an element. Explain how the following are
(Total 8 marks)
achieved in a mass spectrometer:
i) ionisation
(1)
ii) acceleration
(1)
iii) deflection
(1)
36
2 Atomic structure and the periodic table (Topic 1)
SKILL CHECK 10
Which experience a greater effective nuclear charge: the valence
electrons in beryllium or the valence electrons in nitrogen? Why?
807404_C02_Edexcel_GF_Chem_009-036.indd 36
27/02/2015 19:58
40
BILAL HAMEED
PERIODIC TRENDS
A
2
rate
B
0
D
One of the most important chemical species responsible for the removal of ozone from the
stratosphere is a free radical of chlorine, 35Cl •.
157
What does 35Cl • contain?
rate
rate
rate
protons
0
C
A time 17
PERIODIC TRENDS WS 1
neutrons
0
18 0
electrons
16
time
SECTION A
0
0
0
time
0
time
B
17
18
17
1 Which species represented by the following formulae has the largest radius?
12
C
18
17
17
AD
P3– 18
B
Cl –
17
C18 Ar
D
5
K+
11
13
32
Swimming pool water can be kept free of harmful bacteria by adding aqueous sodium chlorate(I),
Which
the following
oxides
is unlikely
to dissolve
in aqueous
sodium
hydroxide?
NaOCl.ofThis
reacts with
water
to produce
HOCl molecules
which
kill bacteria.
Use of the Data Booklet is relevant to this question.
B MgO OCl –C
A Al 2O3
(aq) SO
+ H22O 6 D
OH–SiO
(aq)2 + HOCl (aq)
In the gas phase, aluminium and a transition element require the same amount of energy to form
13 The
chloride
of ion
element
hydrolysed by water to form an acidic solution and its oxide reacts
one
mole
of an
with
aQ2+is– charge.
ion is broken down by ultra-violet light.
In bright
sunshine,
the OCl
acid to form
salt. period (Na to S) is heated in chlorine. The product is purified and then
14 with
An element
of thea third
–
–
added
water.
The resulting
solution
found
What istothe
transition
element?
(aq) +isuv
lightto→beClneutral.
(aq) + ½O2(g)
OCl
What could be the element Q?
2
What
ismethod
the element?
A
Co
Which
would maintain the highest concentration
of HOCl (aq)?
A magnesium
B sodium
Cr
A
Section A
A aluminium
acidify the pool water
B
C aluminium
Cu
B
B
add
a solution
chloride
ions
For C
eachsilicon
question
thereofare
four possible
answers, A, B, C, and D. Choose the one you consider to
D
Ni
be correct.
C silicon
C phosphorus
add a solution of hydroxide ions
D
D phosphorus
D bubble air through the water
1
The petrol additive tetraethyl-lead(IV), Pb(C2H5)4, is now banned in many countries. When it is
3 Which
14
diagram
represents
the change
ionic
radius of the elements across the third period
completely
burned
in air, lead(II)
oxide,inCO
2 and H2O are formed.
(Na
to
Cl)?
15 Which equation represents the reaction that occurs when calcium nitrate is heated strongly?
12 Which solid-line curve most accurately represents the distribution of molecular speeds in a gas at
HowKmany
of oxygen
are
requiredthe
to corresponding
burn one mole distribution
of Pb(C2H5)for
4? the same gas at 300 K?
if the moles
dotted-line
curve
represents
500
) → Ca(NO
A
Ca(NO
2)2 + O2 B
9701/1/O/N/02
A3 2
C
D
9.5
B +N
11O + 2O
C 13.5
D 27
BA Ca(NO
)
→
CaO
A 2
B
3 2
2
2
C Ca(NO3)2 → CaO2 + 2NO2
ion has
moreofelectrons than protons and more
protons
fraction
fraction
of than neutrons?
DWhich
2Ca(NO
3)2 → 2CaO + 4NO2 + O2
molecules
molecules
[H = 11 H ; D = 12 H ; O = 168 O]
Na D–
A
+
Cl B Na
H3Ospeed
Cl
C
Na–
OD
– Na
D ClOHspeed
Cl
15
propellant
used in configuration
the solid rocket
of awith
space
shuttle is
a mixture
of aluminium and
4
3 The
What
is the electronic
of booster
an element
a second
ionisation
C
D energy higher than
compound
X.
Compound
X
contains
chlorine
in
an
oxidation
state
of
+7.
that of each of its neighbours in the Periodic Table?
2
2
6
A 1sof
2s
2p
3s2 of could be compound X?
Which
the
following
fraction
2
2molecules
6
2
1
AB
1s 2s
2p 3s B
3p NH ClO
NH
4Cl
4
3
C
1s22s22p63s23p2
C
fraction of
molecules
NH4ClO4
D N2H5Cl
9701/1/M/J/02
[Turn over
2
2–, HCO– , Cl–, and NO–
speed area may contain Ca2+, Mg2+, CO
speed
16 River
water
in63s
a 2chalky
2s22p
3p3 agricultural
D 1s
3
3
3
ions. In a waterworks, such water is treated by adding a calculated quantity of calcium hydroxide.
5 In which pair is the radius of the second atom greater than that of the first atom?
13
will
be precipitated
the addition
calcium hydroxide?
4 What
Which
compound
has a following
boiling point
which isof
influenced
by hydrogen bonding?
A
A
A
Na,
CH3Mg
CHO
CaCl
2
BB
CaCO
CH3OCH
3
3
CC
MgCO
HCO2H
3
B
Sr, Ca
C
P, N
D
DD Mg(NO
HCO2CH
3)23
CEDAR COLLEGE
5
Which gas is likely to deviate most from ideal gas behaviour?
Cl, Br
PERIODIC TRENDS WS 1
der Waals’ forces.
C
The change from diamond to graphite has a high activation energy.
D
The density of graphite is less than that158
of diamond.
6 The sketch below shows the variation of first ionisation energy with proton number for six
13
elements of consecutive proton numbers between 1 and 18 (H to Ar).
ionisation
energy
X
proton number
What is the identity of the element X?
A
Mg
B
Al
C
D 3P
Si
27 Three successive elements in the Periodic Table have first ionisation energies which have the
14 pattern
The metals
of Group
II react readily with oxygen to form compounds of general formula MO.
shown
in the diagram.
When each of these oxides is added to water, which forms the most alkaline solution?
x
A
MgO
B
CaO
C
D
SrO
BaO
x
6
first
ionisation
15 One mole of each of the following compounds is strongly heated and any gas produced is
11 In
some fireworks
thereenergy
is a and
reaction
between powdered aluminium and powdered barium nitrate
collected
at room temperature
pressure.
in which heat is evolved and an unreactive gas is produced.
From which compound is 24 dm3 of gas likely to be collected?
x
What
is theofequation
for this reaction?
[One mole
any gas occupies
24 dm3 at room temperature and pressure.]
A
A
Al2O33 + BaOC+ 2NO
2Al
+2Ba(NOB
3)2 →
MgCO
Mg(NO3)2
MgCl
B
number
4Al + 4Ba(NO3)2 → 2Al2O3 + 4Ba(NOatomic
2)2 + O 2
D
Mg(OH)2
What
could+ be
the first
sequence?
→ 5Al2of
O3this
+ 3BaO
+ 3N2
C 10Al
3Ba(NO
3)2element
A
D
C
B 3)2N→ 10Al(NO3)3C+ 18BaO
F
10Al + 18Ba(NO
+ 3N2
D
Na
9701/1/O/N/03
38
12
[W’06 Q2]
Use
of the
Data
is is
relevant
to this
question. size?
Which
group
ofBooklet
particles
in order
of increasing
The
A electronic
N
O structures
F of calcium, krypton, phosphorus and an element X are shown.
Which
electronic
is that of element X?
O2– structure
F–
B N3–
13
A
C
2+ 2
6 2+
2pMg
3s23p3Al 3+
1s
Na2s
B
D
2+ 2
6
2pNe
3s23p6F
4s– 2
1s
Na2s
C
1s22s22p63s23p63d64s2
[S’09 1 Q12]
6
2s22p6in
3s2a3pchalky
3d104s2agricultural
4p6
D
1s2water
River
area may contain Ca2+, Mg2+, CO 2− , HCO − , Cl − and NO −
3
3
3
ions. In a waterworks, such water is treated by adding a calculated quantity of calcium hydroxide.
What will be precipitated following the addition of calcium hydroxide?
A
CaCl2
B
CaCO3
C
Ca(NO
CEDAR COLLEGE
PERIODIC TRENDS WS 1
13 Which species has the largest radius?
For each question there are four possible answers, A, B, C, and D. Choose the one you consider to
be correct.
B Cl −
C Ar
D K+
A P3−
159
14
of thespecies
Data Booklet
relevant
to thisatom
question.
1
In which
does is
the
underlined
have an incomplete outer shell?
9 Use
–
The
of one
B the
CHvariation
C physical
F2O or chemical
D property
H3O+ with another for the
A sketch
BF3 graph shows
3
Group II elements.
2
Ammonia is manufactured by the Haber Process, in an exothermic reaction.
Assuming that the amount of catalyst remains constant, which change will not bring about an
y
increase in the rate of the forward
reaction?
A
decreasing the size of the catalyst pieces
B
increasing the pressure
x
C increasing
the temperature
What
are the correct
labels for the axes?
D
removing the ammonia as it is formed
x-axis
y-axis
5
A
number
mass
number
12
is atomic
a white
solid which
was used
to make
the early plastic celluloid. Camphor contains
3 Camphor
The equation
for a reaction
is shown.
theBsame percentage
by mass of hydrogen
and oxygen.
atomic number
melting
point
–1
1
H2O(l) ; ∆H = x kJ mol
H2(g) + 2 Oatomic
2(g)
C is the
firstmolecular
ionisationformula
energy
number
What
of camphor?
DC H
first
ionisation
atomic
A
B energy
C10 His8 Ofully correct
C for
C10radius
H16Oreaction?
D
Which
this
10 pair
6O6of descriptions
C10 H10 O2
[W’07 1 Q14]
[Turn over
type(s)
of enthalpy
change
value ofthan
x the first ionisation energy of silicon?
13 Why is the
first ionisation
energy
of phosphorus greater
10
© UCLES 2007
9701/01/O/N/07
formation
onlyone more proton in its
positive
A AA phosphorus
atom has
nucleus.
negative
B BThe atomicformation
radius of only
a phosphorus atom is greater.
combustion,
positive
C CThe outer
electron information
a phosphorus atom is more
shielded.
combustion,
negative
D DThe outer
electron information
a phosphorus atom is paired.
[W’10 1 Q13]
which
three kJ
gases
are given off?
14
magnesium
nitrate, Mg(NO
3 )2 .7H
2 O, is heated,
11
mol–1
4 When
The value
of the second
ionisation
energy
of calcium
is 1150
.
A
dinitrogen oxide, oxygen, water vapour
Which equation correctly represents this statement?
B
hydrogen, nitrogen, oxygen
A
Ca(g)
C
hydrogen, nitrogen dioxide, oxygen
Ca2+(g) + 2e− ;
Ca2+(g) + e− ;
H o = +1150 kJ mol–1
B
Ca+(g)
D
nitrogen dioxide, oxygen, water vapour
C
Ca+(g) 
Ca2+(g) + e− ;
H o = +1150 kJ mol–1
H o = –1150 kJ mol–1
o
15 Ammonium
in nitrogenous
fertilisers
in kJ
themol
soil–1 can be slowly oxidised by air producing
D Ca(g) sulfate
Ca2+(g)
+ 2e− ;
H
= –1150
sulfuric acid, nitric acid and water.
[S’12 2 Q18]
How many moles of oxygen gas are needed to oxidise completely one mole of ammonium
sulfate?
A
1
B
2
C
3
D
4
as an impurity.
16
Chile
saltpetre, NaNO3 , contains sodium iodide
© UCLES
2012
9701/11/M/J/12
Aqueous silver nitrate is added to an aqueous solution of Chile saltpetre. Concentrated aqueous
ammonia is then added.
Which observations are made?
CEDAR COLLEGE
with acidified silver
nitrate
PERIODIC TRENDS WS 1
with concentrated
aqueous ammonia
C
to reduce its loss during the reaction
D
The magnesium glows and a white solid is produced.
D
to reduce its surface area
18
160
Use of the Data Booklet is relevant to this question.
2
12
The following
equations
the letters
W,sodium
X, Y and
Z allNa
represent
whole numbers.
Sodium
and sulfur
react together
to form
sulfide,
2S.
When
balanced,
equation
requires
one ofwith
letters
W,ofX,sulfur?
Y or Z to be 5?
How
docorrectly
the atomic
radius andwhich
ionic radius
of sodium
compare
those
A
+ XO2 → YCO2ionic
+ ZH
WC3Hatomic
7COOH
2O
radius
radius
→ YCO2 + ZH
B A WC4Hsodium
8 + XO>2 sulfur
2O
sodium
> sulfur
7
HPO4 +< ZH
C B WH3PO
sodium
> sulfur → YNa2sodium
sulfur
4 + XNaOH
2O
15 Which
substance
will
not be a product
the thermal decomposition of hydrated magnesium
sodium
<
sulfur
sodium of
> sulfur
D C WNH
3 + XO2 → YN2 + ZH2O
nitrate?
D
sodium < sulfur
sodium < sulfur
A dinitrogen monoxide
[M’1 Q12]
3
Use of the Data Booklet is relevant to this question.
magnesium
19 B
Which
substanceoxide
does not produce a poisonous gas, when burnt in a limited amount of air?
13
From which particle is the removal of an electron the most difficult?
C
oxygen
A hydrogen
–
A Cl (g)
D
B steam
methane
B
F–(g)
C
K+(g)
D
Na+(g)
[W’11 2 Q3]
C propene
+
4 The
Usespecies
of the Data
Booklet
is 2+
relevant
to this question.
and Mg
are isoelectronic.
This means that they have the same number of
16
Ne, Na
14
D
sulfur
electrons.
560 kg of nitrogen and 120 kg of hydrogen are pressurised, heated and passed over an iron
In
which order
dothe
theirmixture
radii increase?
catalyst.
When
of gases reaches equilibrium, it contains 96 kg of hydrogen.
Which mass
of ammonia does it contain?
smallest
largest
AA 24 kg Ne
+
B Na
68 kg
Mg2+C
136 kg
B
Ne
Mg2+
Na+
C
Mg2+
Ne
Na+
Mg2+
Na+
9701/12/M/J/12
Ne
© UCLES
D2012
D
680 kg
[S’14 3 Q16]
17 Use of the Data Booklet is relevant to this question.
In an experiment, 0.6 mol of chlorine gas, Cl 2, is reacted with an excess of hot aqueous sodium
hydroxide. One of the products is a compound of sodium, oxygen and chlorine.
Which mass of this product is formed?
A
21.3 g
B
44.7 g
© UCLES 2011
C
63.9 g
D
128 g
9701/12/O/N/11
18 Which properties do compounds of aluminium and silicon have in common?
A
Aqueous solutions of their chlorides contain aluminium or silicon cations.
B
Their chlorides have co-ordinate bonding.
C
Their oxides are amphoteric.
D
Their oxides are insoluble in water.
19 When strontium is burnt in oxygen, what colour is the flame?
A
green
B
red
C
white
yellow
CEDARD COLLEGE
PERIODIC TRENDS WS 1
161
8
15
15 The diagram shows the first ionisation energies of 11 consecutive elements.
first
ionisation
energy
/ kJ mol–1
X
Y
atomic number
Which type of elements are labelled X and Y ?
A
Group I metals
B
Group II metals
C
halogens
D
noble gases
[S’11 12 Q15]
8
19
the Data
Bookletoxide
is relevant
to this
16
16 Use
Whyofdoes
aluminium
dissolve
in question.
sodium hydroxide solution?
Which
graph correctly
relativeaselectronegativity
plotted against relative atomic radius for
A Aluminium
oxide shows
can behave
a base.
the elements Na, Mg, Al and Si?
B
Aluminium oxide can behave as an acid.
C
Aluminium oxide has a giant structure.
D
The bonding in aluminium oxide is ionic.
A
B
Si
Al
electronegativity
Al
Mg
electronegativity
Mg
17 Concentrated sulfuric acid can behave both as a strong acid and as an oxidising agent.
Na
Na
Si
With which compound does concentrated sulfuric acid react in this way?
atomic radius
A
ethanol
B
magnesium carbonate
C
propanenitrile
D
sodium bromide
atomic radius
C
D
Si
Na
Mg
Al
electronegativity
Al
electronegativity
Mg
Na
Si
atomic radius
atomic radius
20 Many organic reactions need to be heated before reaction occurs, but some do not require
© UCLES
2011
9701/12/M/J/11
heating.
CEDAR
COLLEGE
PERIODIC TRENDS WS 1
Which reaction occurs quickly at room temperature?
argon
graphite anode
molten strontium
162
7
molten
17
18 The graph below
shows
the variation of the first ionisation energy with the number of protons for
strontium
bromide
some elements.
Q
heat
Y
Why is an atmosphere of argon used around the cathode?
A
A thin film
a compound of strontium andX argon forms on the surface protecting the freshly
firstof
ionisation
formed energy
metal. / kJ mol–1
B
The argon keeps the strontium molten.
V
C
D
W
S
The argon stops the molten strontium rising too high in the tube.
U
Without the argon, strontium oxide would form in the air.
R
13
T
A metal, X, reacts with water to produce a colourless solution which gives a white precipitate
proton number
when mixed with aqueous sulfuric acid.
Which statement is correct?
What
is metal X?
A
Elements Q and Y are in the same period in the Periodic Table.
A
barium
B
The general increase from elements R to Y is due to increasing atomic radius.
B
magnesium
C
The small decrease between elements S and T is due to decreased shielding.
C
potassium
D
The small decrease between elements V and W is due to repulsion between paired
D
sodium
electrons.
[W’13 3 Q18]
19 Which
Use of property
the Data Booklet
is relevant
to going
this question.
18
14
increases
in value
down Group II?
Elements
J and K react together to form compound L. Elements J and K are both in Period 3.
A
electronegativity
Element J has the smallest atomic radius in Period 3. There are only two elements in Period 3
which
have
a lower melting point than element K.
B
ionic
radius
C
maximum
oxidation
Which
compound
could benumber
L?
D
A
second
B energy
MgS
MgCl 2 ionisation
C
Na2S
5
D
PCl 3
[W’13 2 Q14]
12
elements
X,skeletal
Y and formula
Z are inof Period
3 of the Periodic Table. Element Y has the
19
20 Consecutive
Which diagram
gives the
2-bromopentan-4-ol?
highest first ionisation energy and the lowest melting point of these three elements.
© UCLES 2013A
9701/11/O/N/13 C
B
D
What are the identities of X, Y and Z?
OH
Br
Br
Br
A sodium, magnesium, aluminium
B
C
magnesium, aluminium, silicon
OH
aluminium, silicon, phosphorus
D
silicon, phosphorus, sulfur
OH
Br
OH
[M’16 Q12]
© UCLES 2013
9701/13/O/N/13
[Turn
over
13
When dealing with a spillage of metallic sodium
it is important that no toxic or flammable
products
are formed.
Which material should be used if there is a spillage of metallic sodium?
CEDAR COLLEGE
PERIODIC TRENDS WS 1
A dilute hydrochloric acid
B
ethanol
5
C Glucokinase only operates on a narrow range of substrate molecules.
10 When solid ammonium chloride dissociates5at a certain temperature in a 0.500 dm3 container,
ammonia
and hydrogen
chloride
are formed.
D
Glucokinase
provides
an alternative
route
for the reactions it catalyses.
163
11 Enzymes are biological catalysts. Many enzymes show specificity. An example of an enzyme
which shows specificity is glucokinase.
Glucokinase
is involved
in the metabolism of glucose.
NH3(g)
+ HCl (g)
NH4Cl (s)
20
12 Why is the ionic radius of a chloride ion larger than the ionic radius of a sodium ion?
Whatinitial
does specificity
in this context?
The
amount ofmean
ammonium
chloride was 1.00 mol, and when the system had reached
A
A chloride
ionwas
has0.300
one more
occupied
electron
shell than a sodium ion.
mol of
ammonium
chloride.
equilibrium
there
A Glucokinase is most effective as a catalyst over a narrow pH range.
B
Chlorine
has a higher
proton
number
than sodium.
forathis
reaction
these
conditions?
What
is the numerical
value
of Kc as
B Glucokinase
is most
effective
catalyst
overunder
a narrow
range
of temperatures.
CC Ionic
radius increases
regularly across range
the third period.
onlyBoperates
A Glucokinase
0.490
1.63 on a narrow
C 1.96of substrate molecules.
D 3.27
DD Sodium
is a metal,
chlorine
is a non-metal.
Glucokinase
provides
an alternative
route for the reactions it catalyses.
[W’12 1 Q13]
11 Which stage in the free radical substitution of ethane by chlorine has the lowest activation
energy? D and E are both in Period 3. Element D has the smallest atomic radius in Period 3.
13
21
12 Elements
Why is the ionic radius of a chloride ion larger than the ionic radius of a sodium ion?
There are only two elements in Period 3 which have a lower melting point than element E.
2Cl
●
A
Clchloride
2 →
Elements
D and
E has
react
together
to form compound
L. than a sodium ion.
A A
ion
one
more occupied
electron shell
C2H5proton
● + HCl
B
Cl ● + Chas
2H6 a→
B Chlorine
higher
Which
compound
could
be L? number than 5sodium.
C Ionic
C2H5●radius
+ Clincreases
●
the third period.
2 → C2Hregularly
5Cl + Clacross
B excess
MgS of chlorine, C
Na2S compoundDJ. PCl 3
A MgCl
2
12 Silicon
is heated
in an
producing
a 5metal,
chlorine
D Sodium
Cl ● + Cis2H
● → C
2H5Cl is a non-metal.
[S’12 1 Q13]
Excess water is added to the sample of J produced.
14 X and Y are both Group 2 metals.
13 Which
Elements
Dis and
E react
are both
in Period
3. Element
D has the
atomic radius in Period 3.
22
S.
12
Sodium
and
sulfur
together
to form
sodium sulfide,
Na2smallest
row
correct?
only
twohydroxide
elementscompounds,
in Period 3but
which
have
lower
melting
point than
element
more
soluble
in water
than Y(OH)
XThere
and Yare
both
form
X(OH)
2 is a
2. E.
Elements
D
and
E
react
together
to
form
compound
L.
How do the atomic radius and ionic radius of sodium compare with those of sulfur?
Is HCl
produced
If a piece of metal Y is put into cold
water
a very slow reaction occurs, and only a very few, small
structure
of
J
when
water is
Which compound
beseen.
L?
hydrogen
bubblescould
can be
atomic radius
ionicto
radius
added
J?
B
MgS
C
D PCl 3
A
MgCl
2
2S
What
be the <identities
Y? > Na
A could
sodium
sulfur of X and
sodium
A
giant
molecular
no sulfur
B
sodium
< sulfur
sodium
giant
yes < sulfur
Xmolecular
14 X B
and Y are
both
Group 2 metals.Y
C
sodium
> sulfur
sodium
C
simple
molecular
no > sulfur
barium
XA
and Y both
form hydroxidemagnesium
compounds, but X(OH)2 is more soluble in water than Y(OH)2.
D
sodium
> sulfur
sodium < sulfur
D
simple
molecular
B
barium
strontium yes
If a piece of metal Y is put into cold water a very slow reaction occurs, and only a very few,[S’16
small
2 Q12]
hydrogen
bubbles
C
calciumcan be seen.
strontium
13 Which
Solid aluminium
chloride
sublimessmallest
at 178 o C.
13
element has
the second
atomic radius in its group and the third lowest first
23
D
magnesium
calcium
What
could
be
the
identities
of
X
and
Y?
ionisation energy in its period?
Which structure best represents the species in the vapour at this temperature?
A boron
X
Y
A
B
C
D
B A calciumbarium
magnesium
Cl magnesium
Cl
Cl
Cl
Cl
Cl
Al 3+(Cl –)3
Cl
Cl
C
B
barium
strontium
Al
Al calcium
Al
Al
D C sodium
strontiumAl
[S’18 1 Q13]
Cl
magnesium
calcium
ClD
Cl
Cl
Cl
Cl
Cl
9701/11/M/J/16
[Turn over
14 Chlorine reacts with cold aqueous sodium hydroxide to produce sodium chloride, water and
compound X.
© UCLES 2016
Chlorine reacts with hot aqueous sodium hydroxide to produce sodium chloride, water and
compound Y.
What are the oxidation states of chlorine in compound X and compound Y?
X
Y
A
–1
–5
B
–1
+5
C
+1
–5
D
+1
+5
© UCLES
UCLES 2016
2016
©
9701/11/M/J/16
9701/12/M/J/16
CEDAR COLLEGE
15 In which reaction does ammonia behave as a Brønsted-Lowry base?
[Turn
over
[Turn
over
PERIODIC TRENDS WS 1
Which nitrogen-containing compound could be formed in the reaction?
A
NH3
B
N2O
C
NO
164
D
NO2
10 Element X has a higher first ionisation energy than element Y.
24
Two students state what they believe is one factor that helps to explain this.
student 1
“X has a higher first ionisation energy than Y because an atom of X has more
protons in its nucleus than an atom of Y.”
student 2
“X has a higher first ionisation energy than Y because X has a smaller atomic
radius than Y.”
Only one of the two students is correct.
What could X and Y be?
X
Y
A
carbon
boron
B
magnesium
aluminium
C
oxygen
nitrogen
D
oxygen
sulfur
[S’18 1 Q10]
11 Hydrogen ions catalyse the hydrolysis of esters.
Which statement is correct?
A
The hydrogen ions act as a heterogeneous catalyst.
B
The hydrogen ions are in the same phase as the reactants.
C
The hydrogen ions are used up in the reaction.
D
The hydrogen ions have no effect on the activation energy of the reaction.
© UCLES 2018
CEDAR COLLEGE
9701/11/M/J/18
PERIODIC TRENDS WS 1
15
32 Ammonia is produced commercially by the Haber
16514 process in which nitrogen and hydrogen react
as shown.
SectionBB
SECTION
N2(g) + 3H2(g)
2NH3(g) ; ∆H = – 92 kJ mol–1
For each of the questions in this section, one or more of the three numbered statements 1 to 3 may
Which
statements are true of the commercial process?
be
correct.
o
1 A temperature
C is used. is or is not correct (you may find it helpful to put a tick against
Decide
whether eachofof1000
the statements
the statements that you consider to be correct).
2 A pressure of 100 - 200 atm is used.
The
A ammonia
to D should
selected
3 responses
The yield of
is be
less
than 20on
%.the basis of
A
B
C
D
33 What factors can affect the value of the activation energy of a reaction?
1 only
2 and 3
1 and 2
1, 2 and 3
is
only are
are of a catalyst only are
1 the presence
correct
correct
correct
correct
2 changes in temperature
3 other
changes
in concentration
of theis reactants
No
combination
of statements
used as a correct response.
illustrates
theofenergy
changes
of a setinofthe
reactions.
341 31
The The
firstdiagram
ionisation
energies
successive
elements
Periodic Table are represented in the
graph.
_
H = _134 kJ mol 1
2500
R
S
2000
H = +92 J mol
1500
first ionisation
_
energy / kJ mol 1
1000
1
_
T
H = _75 kJ mol 1
U
Which of the following statements are correct?
500
1 The enthalpy change for the transformation U ⎯→ R is + 42 kJ mol–1 .
2
The enthalpy change for the transformation T ⎯→ S is endothermic.
3
0 change for the transformation R ⎯→ T is – 33 kJ mol–1 .
The enthalpy
A B C D E F G H
I
J
K
L M N O P Q R S
T
element
Which of these statements about this graph are correct?
1
Elements B, J and R are in Group 0 of the Periodic Table.
2
Atoms of elements D and L contain 2 electrons in their outer shells.
3
Atoms of elements G and O contain half-filled p orbitals.
© UCLES
2004 2004
© UCLES
CEDAR COLLEGE
9701/01/O/N/04
9701/01/O/N/04
[Turn over
PERIODIC TRENDS WS 1
Decide whether each of the statements is or is not correct (you may find it helpful to put a tick against
2
nucleon number
the statements that you consider to be correct).
E1
E2
3 isotopic mass
The responses A to D should be selected on the 166
basis of
molecular energy
322 Which of the
energy of an atom?
A following influenceBthe size of the ionisation
C
D
Which statements are correct?
1, amount
2 and 3 of shielding 1byand
1 only
1 the
the2inner electrons2 and 3
1 The are
shaded area represents
of molecules
with energies
between E 1 and E 2 at
only arethe proportiononly
are
is
.
temperature
T
2
the
charge
on
the
nucleus
2
correct
correct
correct
correct
No distance
particles between
have zero
energy
at either temperature.
32 the
the
outer electrons
and the nucleus
No other combination of statements is used as a correct response.
3 T2 is a lower temperature than T1 .
33 Which equations can apply to an ideal gas?
3 [p
31
Compound
X is
from Mtwo
elements.
One
second highest
valueconstant,
of first
= pressure,
V made
= volume,
= molar
mass,
ρ =element
density,has
c = the
concentration,
R = gas
O
is
heated,
the
following
reaction
occurs.
34 T
When
KCl
3
ionisation
energy
in
its
group
and
the
other
element
has
the
third
highest
value
of
first
ionisation
= temperature]
12
energy in its group.
ρRT
4KCl O3 →
4 + KCl
1 p=
2
pV = cRT
3 3KCl
pVB=OMRT
Section
M
M
Which compounds
could be compound
X?
Which statements are correct?
For each
of the questions
1 calcium
chloride in this section, one or more of the three numbered statements 1 to 3 may
be correct.
1 The oxidation state of Cl in KCl O3 is +5.
2 magnesium bromide
Decide
each ofstate
the statements
or is not
correct (you
2 whether
The oxidation
of some Clisatoms
decreases
by 6.may find it helpful to put a tick against
3 potassium
sulfide
the statements
that you
consider to be correct).
3 The reaction involves disproportionation.
The responses A to D should be selected on the basis of
32 Water has some unusual physical properties compared to other hydrides of Group 16 elements.
Someis of
these
propertiesenergy
are due
hydrogenless
bonds.
intermolecular forces are much
4 Why
35
the
first ionisation
of to
aluminium
than These
that of magnesium?
A water
D
stronger in
than they areBin H2S, for example. C
1 The
aluminium
atom is2more
nuclear charge.
1, 2 outer
and 3electron in the
1 and
2
and 3shielded from the
1 only
Which statements are correct?
are
only are
only are
is
2 The outer electron in the aluminium atom is in a higher energy orbital.
correct bonds cause correct
correct
correct
1 Hydrogen
the melting point
of ice
to be higher than expected.
© UCLES 2015
9701/12/O/N/15
[Turn over
3 The outer electron in the aluminium atom is further from the nucleus.
2 Hydrogen bonds cause the surface tension of water to be higher than expected.
No other combination of statements is used as a correct response.
3 Hydrogen bonds cause the viscosity of water14
to be higher than expected.
36 Ammonia is a colourless gas that is produced by the Haber process.
5 Compound
31
frombe
two
elements.
Onebasis
element
The
responsesXA isto made
D should
selected
on the
of has the second highest value of first
Which statements
ammonia
areother
correct?
ionisation
energy inabout
its group
and the
element has the third highest value of first ionisation
energy in its group.
A
B bond pairs and one
C lone pair of electrons.
D
1 An ammonia
molecule has three
Which compounds could be compound X?
2 and 3is bubbled into
1 water
and 2 the pH of the solution
2 and 3 will increase. 1 only
2 If 1,
ammonia
13
2
arechloride
only are
only are
is
calcium
Ammonia
by warming ammonium
sulfate with aqueous
hydrochloric acid.
correctgas can be made
correct
correct
correct
magnesium bromide
3 potassium
sulfide
No other
combination
of statements is used9701/13/M/J/16
as a correct response.
[Turn over
© UCLES 2016
[S’16 3 31]
6 Water
32
has
some unusual
33
X is an
element
that hasphysical properties compared to other hydrides of Group 16 elements.
Some of these properties are due to hydrogen bonds. These intermolecular forces are much
example.
stronger in water than they are in H2S, for9701/13/M/J/16
© UCLES 2016
●
its outer electrons in the 4th principal quantum shell,
Which statements
are correct?
● a higher
1st ionisation energy than calcium.
1
Hydrogen bonds cause the melting point of ice to be higher than expected.
What could be the identity of X?
2
Hydrogen bonds cause the surface tension of water to be higher than expected.
1
bromine
3
Hydrogen bonds cause the viscosity of water to be higher than expected.
2
krypton
3
xenon
34 Methanol, CH3OH, can be produced industrially by reacting CO with H2.
CEDAR COLLEGE
CO(g) + 2H2(g)
CH3OH(g)
The process can be carried out at 4 × 103 kPa and 1150 K.
PERIODIC
TRENDS
∆H = –91
kJ mol–1WS 1
167
2
PERIODIC TRENDS WS 2
Answer all the questions in the space provided.
For
Examiner’s
Use
1 1 The first six ionisation energies of an element X are given below.
ionisation energy / kJ mol–1
first
second
third
fourth
fifth
sixth
950
1800
2700
4800
6000
12300
(a) Define the term first ionisation energy.
..........................................................................................................................................
..........................................................................................................................................
..........................................................................................................................................
.................................................................................................................................... [3]
(b) Write an equation, with state symbols, for the second ionisation energy of element X.
.................................................................................................................................... [2]
(c) Use the data given above to deduce in which Group of the Periodic Table element X is
placed. Explain your answer.
Group ...............................................................................................................................
explanation ......................................................................................................................
..........................................................................................................................................
..........................................................................................................................................
.................................................................................................................................... [3]
The first ionisation energies (I.E.) for the elements of Group IV are given below.
element
C
Si
Ge
Sn
Pb
1st I.E. / kJ mol–1
1090
786
762
707
716
(d) Explain the trend shown by these values in terms of the atomic structure of the elements.
..........................................................................................................................................
..........................................................................................................................................
..........................................................................................................................................
..........................................................................................................................................
.................................................................................................................................... [4]
[Total: 12]
© UCLES 2005
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PERIODIC TRENDS WS 2
168
4
2 2 The Periodic Table we currently use is derived directly from that proposed by Mendeleev
in 1869 after he had noticed patterns in the chemical properties of the elements he had
studied.
The diagram below shows the first ionisation energies of the first 18 elements of the Periodic
Table as we know it today.
2500
He
Ne
2000
first
ionisation
energy
/kJ mol-1
1500
1000
500
0
Ar
H
Li
Na
1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18
proton number
(a) Give the equation, including state symbols, for the first ionisation energy of fluorine.
...................................................................................................................................... [2]
(b) Explain why there is a general increase in first ionisation energies from sodium to
argon.
..........................................................................................................................................
..........................................................................................................................................
..........................................................................................................................................
...................................................................................................................................... [3]
(c) (i)
Explain why the first ionisation energy of aluminium is less than that of
magnesium.
..................................................................................................................................
..................................................................................................................................
..................................................................................................................................
CEDAR
© UCLESCOLLEGE
2008
9701/02/M/J/08
PERIODIC TRENDS WS 2
For
Examiner’s
Use
169
5
(ii)
Explain why the first ionisation energy of sulphur is less than that of phosphorus.
..................................................................................................................................
For
Examiner’s
Use
..................................................................................................................................
..................................................................................................................................
[4]
4
3 Magnesium
2The
react
withsodium
chlorine,
oxygen,
orisnitrogen
to give the chloride, or
table belowwill
refers
to on
theheating
elements
to or
sulphur
and
incomplete.
oxide, or nitride respectively. Each of these compounds is ionic and in them magnesium has
the same +2Na
oxidation state.
element
Mg
Al
Si
P
S
For
Examiner’s
Use
(a) (i) Write an equation,
melting point
high with state symbols, for the second ionisation energy of
magnesium.
conductivity
high
..................................................................................................................................
(d) (i) (ii)
Complete
‘melting
point’
row by using
only thechange
words that
‘high’
or ‘low’.
Use thethe
Data
Booklet
to calculate
the enthalpy
occurs
when one mole
2+
of gaseous magnesium ions, Mg , is formed from one mole of gaseous magnesium
(ii) Complete the ‘conductivity’ row by using only the words ‘high’, ‘moderate’ or ‘low’.
atoms.
[5]
Include a sign in your answer.
(e) When Mendeleev published his Periodic Table, the elements helium, neon and argon
were not included.
Suggest a reason for this.
enthalpy change = ……………… kJ mol–1
..........................................................................................................................................
[3]
...................................................................................................................................... [1]
(b) Separate samples of magnesium chloride and magnesium oxide are shaken with water.
In each case, describe what you would see when this is done, and state the approximate
[Total: 15]
pH of the water after the solid has been shaken with it.
4 Elements in the same period of the Periodic Table show trends in physical and chemical
properties.
(i) magnesium chloride
On the grid below, draw a clear sketch to show the variation of the stated property.
Below the grid, briefly explain the variation you have described in your sketch.
observation ...............................................................................................................
You should refer to the important factors that cause the differences in the property you are
describing.
approximate pH of the water …………………
(ii)
magnesium oxide
observation ...............................................................................................................
approximate pH of the water …………………
[4]
CEDAR COLLEGE
© UCLES 2008
PERIODIC TRENDS WS 2
9701/02/M/J/08
[Turn over
On each of these grids, draw a clear sketch to show the variation of the stated property.
Below each grid, briefly explain the variation you have described in your sketch.
For each explanation you should refer to the important factors that cause the differences in
170
the property you are describing.
(a)
atomic
radius of
element
Na
Mg
Al
Si
P
S
Cl
explanation ......................................................................................................................
..........................................................................................................................................
..........................................................................................................................................
..........................................................................................................................................
[3]
(b)
melting
point of
element
Na
Mg
Al
Si
P
S
Cl
explanation ......................................................................................................................
..........................................................................................................................................
..........................................................................................................................................
..........................................................................................................................................
[4]
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PERIODIC TRENDS WS 2
171
8
45 Barium, Ba, was discovered by Davy in 1808. The element gets its name from the Greek ‘barys’
meaning ‘heavy’.
(a) The table below compares some properties of barium with caesium.
(i)
element
Cs
Ba
group
1
2
atomic number
55
56
atomic radius / pm
531
435
Why do caesium and barium have different atomic numbers?
..................................................................................................................................... [1]
(ii)
State the block in the Periodic Table in which caesium and barium are found.
..................................................................................................................................... [1]
(iii)
Explain why the atomic radius of barium is less than the atomic radius of caesium.
...........................................................................................................................................
...........................................................................................................................................
...........................................................................................................................................
..................................................................................................................................... [3]
(iv)
Predict and explain whether a barium ion is larger, smaller or the same size as a
barium atom.
...........................................................................................................................................
...........................................................................................................................................
...........................................................................................................................................
..................................................................................................................................... [2]
© OCR 2007
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PERIODIC TRENDS WS 2
172
More Past Papers at: www.pastpapers.org
6
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PERIODIC TRENDS WS 2
173
6
3 7 This question is about the elements in Group II of the Periodic Table, magnesium to barium.
(a) Complete the table below to show the electronic configuration of calcium atoms
and of strontium ions, Sr2+.
1s
2s
2p
Ca
2
2
6
Sr2+
2
2
6
3s
3p
3d
4s
4p
[2]
(b) Explain the following observations.
(i)
4d
The atomic radii of Group II elements increase down the Group.
..................................................................................................................................
..................................................................................................................................
(ii)
The strontium ion is smaller than the strontium atom.
..................................................................................................................................
..................................................................................................................................
(iii)
The first ionisation energies of the elements of Group II decrease with increasing
proton number.
..................................................................................................................................
..................................................................................................................................
..................................................................................................................................
..................................................................................................................................
[4]
© UCLES
2007
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COLLEGE
9701/02/M/J/07
PERIODIC TRENDS WS 2
For
Examiner’s
Use
174
4
2 8 The alkali metals are a series of six elements in Group I of the Periodic Table. The first
ionisation energy of these elements shows a marked trend as the Group is descended.
(a) Define the term first ionisation energy.
..........................................................................................................................................
..........................................................................................................................................
.................................................................................................................................... [2]
(b) (i)
State and explain the trend in first ionisation energy as Group I is descended.
..................................................................................................................................
..................................................................................................................................
..................................................................................................................................
(ii)
Suggest how this trend helps to explain the increase in the reactivity of the elements
as the Group is descended.
..................................................................................................................................
..................................................................................................................................
[3]
(c) In a redox reaction, 0.83 g of lithium reacted with water to form 0.50 dm3 of aqueous
lithium hydroxide.
2Li(s) + 2H2O(l)
(i)
2LiOH(aq) + H2(g)
Calculate the amount, in moles, of lithium that reacted.
CEDAR
© UCLESCOLLEGE
2010
9701/23/M/J/10
PERIODIC TRENDS WS 2
For
Examiner’s
Use
2
For
Examiner’
Use
Answer all the questions in the spaces provided.
175
19 Although the actual size of an atom cannot be measured exactly, it is possible to measure
the distance between the nuclei of two atoms. For example, the ‘covalent radius’ of the Cl
atom is assumed to be half of the distance between the nuclei in a Cl 2 molecule. Similarly,
the ‘metallic radius’ is half of the distance between two metal atoms in the crystal lattice of a
metal. These two types of radius are generally known as ‘atomic radii’.
The table below contains the resulting atomic radii for the elements of period three of the
Periodic Table, Na to Cl.
element
Na
Mg
Al
Si
P
S
Cl
atomic radius / nm
0.186
0.160
0.143
0.117
0.110
0.104
0.099
(a) (i) Explain qualitatively this variation in atomic radius.
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
(ii) Suggest why it is not possible to use the same type of measurement for argon, Ar.
....................................................................................................................................
....................................................................................................................................
[4]
(b) (i) Use the Data Booklet to complete the following table of radii of the cations and
anions formed by some of the period three elements.
radius of cation / nm
Na+
© UCLES 2012
CEDAR COLLEGE
Mg2+
radius of anion / nm
Al 3+
P3–
S2–
Cl –
9701/23/M/J/12
PERIODIC TRENDS WS 2
176
3
(ii) Explain the differences in size between the cations and the corresponding atoms.
For
Examiner’s
Use
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
(iii) Explain the differences in size between the anions and the corresponding atoms.
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
[5]
(c) Each of the elements Na to Cl forms at least one oxide. Na2O is an ionic oxide, SO2 is a
covalent oxide. Both oxides react with water.
(i) Write an equation for the reaction of each of these oxides with water.
Na2O ..........................................................................................................................
SO2 ............................................................................................................................
(ii) What is the pH of the resulting solution in each case?
Na2O ................
SO2 .....................
(iii) Write an equation for the reaction that occurs between the products of your reactions
in (i).
....................................................................................................................................
[5]
[Total: 14]
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COLLEGE
© UCLES 2012
9701/23/M/J/12
PERIODIC TRENDS WS
2 over
[Turn
1774
210 The alkali metals are a series of six elements in Group I of the Periodic Table. The first
ionisation energy of these elements shows a marked trend as the Group is descended.
(a) Define the term first ionisation energy.
..........................................................................................................................................
..........................................................................................................................................
.................................................................................................................................... [2]
(b) (i)
State and explain the trend in first ionisation energy as Group I is descended.
..................................................................................................................................
..................................................................................................................................
..................................................................................................................................
(ii)
Suggest how this trend helps to explain the increase in the reactivity of the elements
as the Group is descended.
..................................................................................................................................
..................................................................................................................................
[3]
(c) In a redox reaction, 0.83 g of lithium reacted with water to form 0.50 dm3 of aqueous
lithium hydroxide.
2Li(s) + 2H2O(l)
(i)
2LiOH(aq) + H2(g)
Calculate the amount, in moles, of lithium that reacted.
CEDAR
COLLEGE
© UCLES 2010
9701/23/M/J/10
PERIODIC TRENDS WS 2
For
Examiner’s
Use
178 5
(ii)
Calculate the volume of hydrogen produced at room temperature and pressure.
(iii)
Calculate the concentration, in mol dm–3, of the LiOH(aq) formed.
6
For
Examiner’s
Use
[5]
11 (b) The graph below shows the variation of the first ionisation energies across Period 3.
(d) When heated in chlorine, all of the alkali metals react to form the corresponding
chloride.
1600
Describe what
you see when sodium is heated in chlorine and write a balanced equation
1400
for the reaction.
1200
description 1000
first ionisation
..........................................................................................................................................
800
energy / kJ mol–1
600
..........................................................................................................................................
400
..........................................................................................................................................
200
equation
0
Na
Mg
Al
Si
P
S
Cl
Ar
..........................................................................................................................................
(i) Explain why the first ionisation energy of Ar is greater than that of Cl.
[2]
[Total: 12]
.............................................................................................................................................
....................................................................................................................................... [1]
(ii) Explain why the first ionisation energy of Al is less than that of Mg.
.............................................................................................................................................
....................................................................................................................................... [1]
(iii) Explain why the first ionisation energy of S is less than that of P.
.............................................................................................................................................
....................................................................................................................................... [1]
[Total: 10]
© UCLES 2010
CEDAR
COLLEGE
9701/23/M/J/10
PERIODIC TRENDS[Turn
WS 2over
Answer all the questions in the spaces provided.
1
179
Neon is a noble gas.
12 (a) Complete the full electronic configuration of neon.
1s2 ........................................................................................................................................ [1]
(b) (i) Explain what is meant by the term first ionisation energy.
.............................................................................................................................................
.............................................................................................................................................
....................................................................................................................................... [3]
(ii) Explain why the first ionisation energy of neon is greater than that of fluorine.
.............................................................................................................................................
....................................................................................................................................... [2]
(c) Neon has three stable isotopes.
isotope
mass number
percentage abundance
1
9.25
2
20
90.48
3
21
0.27
(i) Define the term relative atomic mass.
.............................................................................................................................................
....................................................................................................................................... [2]
(ii) Use the relative atomic mass of neon, 20.2, to calculate the mass number of isotope 1.
mass number = ................................. [2]
© UCLES 2015
CEDAR COLLEGE
9701/23/M/J/15
PERIODIC TRENDS WS 2
2
180
Answer all the questions
in the spaces provided.
113 (a) Successive ionisation energies for the elements magnesium to barium are given in the table.
element
1st ionisation
2nd ionisation
3rd ionisation
energy / kJ mol–1 energy / kJ mol–1 energy / kJ mol–1
Mg
736
1450
7740
Ca
590
1150
4940
Sr
548
1060
4120
Ba
502
966
3390
(i) Explain why the first ionisation energies decrease down the group.
.............................................................................................................................................
.............................................................................................................................................
.............................................................................................................................................
....................................................................................................................................... [3]
(ii) Explain why, for each element, there is a large increase between the 2nd and 3rd ionisation
energies.
.............................................................................................................................................
.............................................................................................................................................
....................................................................................................................................... [2]
(b) A sample of strontium, atomic number 38, gave the mass spectrum shown. The percentage
abundances are given above each peak.
100
82.58
percentage
abundance
0
9.86 7.00
0.56
80
81
82
83
84
85
86
87
88
89
90
atomic mass units
© UCLES 2014
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PERIODIC TRENDS WS 2
181
3
(i) Complete the full electronic configuration of strontium.
1s2 2s2 2p6 ................................................................................................................... [1]
(ii) Explain why there are four different peaks in the mass spectrum of strontium.
.............................................................................................................................................
....................................................................................................................................... [1]
(iii) Calculate the atomic mass, Ar, of this sample of strontium.
Give your answer to three significant figures.
Ar = ............................. [2]
(c) A compound of barium, A, is used in fireworks as an oxidising agent and to produce a green
colour.
(i) Explain, in terms of electron transfer, what is meant by the term oxidising agent.
.............................................................................................................................................
....................................................................................................................................... [1]
(ii) A has the following percentage composition by mass: Ba, 45.1; Cl , 23.4; O, 31.5.
Calculate the empirical formula of A.
empirical formula of A ........................................... [3]
© UCLES 2014
CEDAR COLLEGE
9701/21/O/N/14
[Turn over
PERIODIC TRENDS WS 2
2
Answer all the questions in the spaces provided.
182
1 14 (a) Successive ionisation energies for the elements fluorine, F, to bromine, Br, are shown on the
graph.
12 000
F
10 000
8000
Cl
ionisation
6000
energy / kJ mol–1
Br
4000
2000
0
1
2
3
electrons removed
4
5
(i) Explain why the first ionisation energies decrease down the group.
.............................................................................................................................................
.............................................................................................................................................
.............................................................................................................................................
....................................................................................................................................... [3]
(ii) Explain why there is an increase in the successive ionisation energies of fluorine.
.............................................................................................................................................
.............................................................................................................................................
....................................................................................................................................... [2]
© UCLES 2014
CEDAR COLLEGE
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PERIODIC TRENDS WS 2
183
3
15
(c) The fifth to eighth ionisation energies of three elements in the third period of the Periodic Table
are given. The symbols used for reference are not the actual symbols of the elements.
ionisation energies, kJ mol–1
fifth
sixth
seventh
eighth
X
6274
21 269
25 398
29 855
Y
7012
8496
27 107
31 671
Z
6542
9362
11 018
33 606
(i) State and explain the group number of element Y.
group number ...............................
explanation .........................................................................................................................
.............................................................................................................................................
[1]
(ii) State and explain the general trend in first ionisation energies across the third period.
.............................................................................................................................................
.............................................................................................................................................
....................................................................................................................................... [2]
(iii) Explain why the first ionisation energy of element Y is less than that of element X.
.............................................................................................................................................
.............................................................................................................................................
.............................................................................................................................................
....................................................................................................................................... [2]
(iv) Complete the electronic configuration of element Z.
1s2 ................................................................................................................................. [1]
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PERIODIC TRENDS WS 2
.............
.............
.............
.............
[4]
184
16 The fifth to eighth ionisation energies of three elements in the third period of the Periodic Table
(b)
are given. The symbols used for reference are not the actual symbols of the elements.
ionisation energies, kJ mol–1
fifth
sixth
seventh
eighth
X
7012
8496
27 107
31 671
Y
6542
9362
11 018
33 606
Z
7238
8781
11 996
13 842
(i) State and explain the group number of element Y.
group number ...............................
explanation .........................................................................................................................
.............................................................................................................................................
[1]
(ii) State and explain the general trend in first ionisation energies across the third period.
.............................................................................................................................................
.............................................................................................................................................
....................................................................................................................................... [2]
(iii) Complete the electronic configuration of element X.
1s2 ................................................................................................................................. [1]
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PERIODIC TRENDS WS 2
185
8
317 The elements in the third period exhibit periodicity in both their chemical and physical properties.
(a) A graph of the atomic and ionic radii across the third period is shown.
0.25
= atomic radius / nm
= ionic radius / nm
0.20
atomic or 0.15
ionic radius
/ nm
0.10
0.05
0.00
(i)
Na Na+ Mg Mg2+ Al
Al 3+ Si Si4+ P
atoms and ions
P3–
S
S2–
Cl
Cl –
Explain the decrease in atomic radius across the third period.
.............................................................................................................................................
.............................................................................................................................................
.............................................................................................................................................
....................................................................................................................................... [2]
(ii)
Explain why, for sodium to silicon, the ionic radii are less than the atomic radii.
.............................................................................................................................................
....................................................................................................................................... [1]
(iii)
Explain why, for phosphorus to chlorine, the ionic radii are greater than the atomic radii.
.............................................................................................................................................
.............................................................................................................................................
....................................................................................................................................... [2]
(b) The first ionisation energies of the elements across the third period show a general increase.
Aluminium and sulfur do not follow this general trend.
(i)
Explain why aluminium has a lower first ionisation energy than magnesium.
.............................................................................................................................................
.............................................................................................................................................
....................................................................................................................................... [2]
[S’18 1 Q3]
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CEDAR COLLEGE
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PERIODIC TRENDS WS 2
186
18
CEDAR COLLEGE
PERIODIC TRENDS WS 2
187
2
119 Sir James Jeans, who was a great populariser of science, once described an atom of carbon
as being like six bees buzzing around a space the size of a football stadium.
(a) (i)
Suggest what were represented by the six bees in this description.
...................................................................................................................................
(ii)
Explain (in terms of an atom of carbon) what stopped the bees from flying away
from the space of the football stadium.
...................................................................................................................................
...................................................................................................................................
(iii)
What is missing from Jeans’ description when applied to an atom of carbon?
...................................................................................................................................
...................................................................................................................................
[3]
(b) The diagram below represents the energy levels of the orbitals in atoms of the second
period, lithium to neon.
Label the energy levels to indicate the principal quantum number and the type of
orbital at each energy level.
energy
(i)
nucleus
(ii)
In the space below, sketch the shapes of the two types of orbital.
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PERIODIC TRENDS WS 2
For
Examiner’s
Use
188
3
(iii)
Complete the electron configurations of nitrogen and oxygen on the energy level
diagrams below, using arrows to represent electrons.
nitrogen
(iv)
For
Examiner’s
Use
oxygen
Explain, with reference to your answer to (iii), the relative values of the first
ionisation energies of nitrogen and oxygen. The values are given in the Data
Booklet and should be quoted in your answer.
...................................................................................................................................
...................................................................................................................................
...................................................................................................................................
...................................................................................................................................
[6]
(c) (i)
State the formulae of the negatively charged ions formed by these elements in
simple binary compounds (nitrides and oxides).
...................................................................................................................................
(ii)
Why do nitrogen and oxygen form negative ions, but not positive ions, in simple
binary compounds?
...................................................................................................................................
...................................................................................................................................
...................................................................................................................................
[2]
[Total : 11]
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PERIODIC TRENDS WS 2
189
3 Chemical bonding
This topic introduces the different ways by which chemical bonding occurs and the
effect this can have on physical properties.
3.2 Covalent bonding and co-ordinate (dative covalent) bonding
COVALENT BONDING
BILAL HAMEED
COVALENT BONDING
Cambridge International AS and A Level Chemistry 9701 syllabus Syllabus content
3 Chemical bonding
190
This topic introduces the different ways by which chemical bonding occurs and the effect this can have
3
bonding
onChemical
physical properties.
This topic introduces the different ways by which chemical bonding occurs and the effect this can have
on physical properties.
Learning outcomes
Candidates should be able to:
3.1 Ionic bonding
Learning outcomes
Candidates
be able to:
a) describeshould
ionic bonding,
as in sodium chloride, magnesium oxide and
calcium fluoride, including the use of ‘dot-and-cross’ diagrams
3.1 Ionic bonding
3.2 Covalent bonding
and co-ordinate
(dative covalent)
3.2 Covalent
bonding
bonding
including
and
co-ordinate
shapes covalent)
of simple
(dative
molecules
bonding
including
shapes of simple
molecules
a) describe ionic bonding, as in sodium chloride, magnesium oxide and
including
use of ‘dot-and-cross’
diagrams
a) calcium
describe,fluoride,
including
the usethe
of ‘dot-and-cross’
diagrams:
(i) covalent bonding, in molecules such as hydrogen, oxygen, chlorine,
a) describe,
including
the carbon
use of ‘dot-and-cross’
diagrams:
hydrogen
chloride,
dioxide, methane,
ethene
(i)
bonding,
molecules
such assuch
hydrogen,
oxygen,
chlorine,
(ii) covalent
co-ordinate
(dativeincovalent)
bonding,
as in the
formation
of the
hydrogen
chloride,
dioxide,
methane,
ethene
ammonium
ion andcarbon
in the Al
Cl
molecule
2
6
co-ordinate
(dative
covalent)
bonding,
such
as in the
formation
of the
b) (ii)
describe
covalent
bonding
in terms
of orbital
overlap,
giving
σ and π
2
3
ammonium
in the of
Alhybridisation
Cl
molecule
bonds,
includingion
theand
concept
to
form
sp,
sp
and
sp
2
6
orbitals (see
also Section
b) describe
covalent
bonding14.3)
in terms of orbital overlap, giving σ and π
theof,
concept
of hybridisation
to form sp,
sp2 and
c) bonds,
explain including
the shapes
and bond
angles in, molecules
by using
thesp3
orbitals
(see
also Section
14.3) repulsion (including lone pairs), using
qualitative
model
of electron-pair
as simple
CO2 in,
(linear),
CH4 (tetrahedral),
3 (trigonal),
c) explain
theexamples:
shapes of,BF
and
bond angles
molecules
by using theNH3
SF6 (octahedral),
PF5 (trigonal
bipyramidal)
(pyramidal),model
H2O (non-linear),
qualitative
of electron-pair
repulsion (including
lone pairs),
using
simple
CO2 in,
(linear),
CH4 (tetrahedral),
NH3
d) as
predict
theexamples:
shapes of,BF
and
bond angles
molecules
and ions analogous
3 (trigonal),
SF6also
(octahedral),
PF5 (trigonal bipyramidal)
(pyramidal),
H2O (non-linear),
to those specified
in 3.2(b) (see
Section 14.3)
3.3 Intermolecular
forces,
electronegativity
3.3 Intermolecular
and bond properties
forces,
electronegativity
and bond properties
d) predict the shapes of, and bond angles in, molecules and ions analogous
those specified
3.2(b) (see
also
Section and
14.3)water as simple
a) to
describe
hydrogeninbonding,
using
ammonia
examples of molecules containing N–H and O–H groups
a)
hydrogen
bonding,
ammonia
water as simple
b) describe
understand,
in simple
terms,using
the concept
of and
electronegativity
and apply
examples
ofthe
molecules
containing
N–H and
O–H
groups
it to explain
properties
of molecules
such
as bond
polarity (see
also Section in
3.3(c)),
the
dipolethe
moments
(3.3(d))and
andapply
the
b) understand,
simple
terms,
conceptof
ofmolecules
electronegativity
behaviour
of
oxides
with
water
(9.2(c))
it to explain the properties of molecules such as bond polarity (see
Section
3.3(c)),
the energy,
dipole moments
of molecules
(3.3(d)) and
c) also
explain
the terms
bond
bond length
and bond polarity
and the
use
behaviour
of oxides
water (9.2(c))
them to compare
thewith
reactivities
of covalent bonds (see also Section
5.1(b)(ii))the terms bond energy, bond length and bond polarity and use
c) explain
d) them
describe
intermolecular
forces (van
der Waals’
forces),
on
to compare
the reactivities
of covalent
bonds
(seebased
also Section
permanent and induced dipoles, as in, for example, CHCl 3(l); Br2(l) and
5.1(b)(ii))
the
liquidintermolecular
Group 18 elements
d) describe
forces (van der Waals’ forces), based on
3.4 Metallic bonding
3.4 Metallic bonding
3.5 Bonding and
physical properties
3.5 Bonding and
physical properties
permanent and induced dipoles, as in, for example, CHCl 3(l); Br2(l) and
the
liquid metallic
Group 18
elements
a) describe
bonding
in terms of a lattice of positive ions
surrounded by delocalised electrons
a) describe metallic bonding in terms of a lattice of positive ions
by delocalised
electrons
a) surrounded
describe, interpret
and predict
the effect of different types of bonding
(ionic bonding, covalent bonding, hydrogen bonding, other intermolecular
interactions,
metallic
bonding)
on the
physical
properties
ofof
substances
a) describe,
interpret
and
predict the
effect
of different
types
bonding
b) (ionic
deduce
the type
of bonding
present
from given
information
bonding,
covalent
bonding,
hydrogen
bonding,
other intermolecular
metallic of
bonding)
on reactions
the physical
properties
of substances
c) interactions,
show understanding
chemical
in terms
of energy
transfers
associated
breakingpresent
and making
chemical
bonds
b) deduce
the with
type the
of bonding
from of
given
information
c) show understanding of chemical reactions in terms of energy transfers
associated with the breaking and making of chemical bonds
Back to contents page
Back to contents page
BILAL HAMEED
www.cie.org.uk/alevel
19
www.cie.org.uk/alevel
19
COVALENT BONDING
191
O + O
CHEMICAL BONDING
O O
O
O
Note that in this Lewis structure of O2 there are a total of f
bonding
pairsthe
of electrons
(the held
lone pairs) and two bon
atoms are
C H E MIC A L BON DING AWhen
N D S T Rtwo
U C T Uor
R Emore atoms form a chemical compound,
pairs of electrons.
●
together in a characteristic arrangement by attractive forces.
of a covalent bond
Covalent bonding
Nitrogen, N2
TheInchemical
bond
is how
the atoms
forcecan
ofeither
attraction
between
atoms
a
ionic bonding
we saw
lose or gain
electrons
in
Nitrogen
is inany
grouptwo
15, so
has fiveinvalence
electrons. Henc
nd is formed by the
order to attain a noble gas electron configuration. Aacquiring
second type
of
three more electrons nitrogen would achieve a n
The
attraction
the repulsion of
ttraction between a shared compound.
chemical bond
exists,
however, is
in the
whichforce
atoms that
share overcomes
electrons
each
electron with
configuration
with a complete octet of electrons.
ns and the positively
in order charged
to attain a noble
gas electron
This type of
the other
positively
nuclei
of the configuration.
two atoms.
If two nitrogen atoms each share three electrons with each oth
●
●
i. According to IUPAC (the
bonding is covalent bonding, and it usually occurs between non-metals.
electron configuration can be achieved and results in the forma
Union of Pure and Applied
In order to look
at this type
of bonding
in detail, itare
is useful
first
to between
Interactions
involving
valence
electrons
responsible
for the
chemical
covalent
bond
the two
nitrogen atoms. This covalent
covalent bond is a region of
introduce the idea of a Lewis symbol, which is a simple
and
convenient
triple
bond
and
the
three
shared
pairs can be represented by
electron density between bond. We shall focus our attention on these electrons and the electron
method of representing the valence (outer shell) electrons of an element.
ses at least partly from
In sub-topic 4.3 we shall develop this further into what weNterm
the
N
N N
arrangement
of atoms both before and after bond+
formation.
electrons and gives rise
Lewis (electron dot) structure of a compound, based on a system
e force and characteristic
devised by the US chemist, Gilbert N. Lewis (1875–1946).
istance.
N N
In a Lewis symbol representation, each element is surrounded by a
●
Noteelectrons
that in this
Lewis structure of N2 there are a total of t
number of dots (or crosses), which represent the valence
of the
1
bonding
pairs
of
electrons (the lone pairs) and three bo
element. Some examples are given in figure 1.
pairs of electrons.
Let us consider the presence of covalent bonding in four different
species, F2, O2, N2, and HF.
Hydrogen fluoride, HF
N
Fluorine is in group 17, so has seven valence electrons. Hence
acquiring one
4 . 2 more
C O V A electron,
L E N T B O Nfluorine
D I N G would attain a noble ga
configuration
with a complete octet of electrons. Hydrogen is
Fluorine is in group 17, so has seven valence electrons.
Hence by
has just
acquiring one more electron, fluorine would attain1,aso
noble
gas one valence electron. Hence by acquiring just on
electron,
hydrogen
would attain the noble gas configuration
electron configuration with a complete octet of electrons.
●
Fluorine, F2
Cl
●
O + O
O O
Note that hydrogen does not acquire an octet (the octet ru
is historical in nature, and the key point to remember here
O O
hydrogen is the formation of a noble gas electron configur
F
●
Note that in this Lewis structure of O2 there are
a
total of four non● the
Theshared
Lewis symbols
forpair
hydrogen
and fluorine are:
When electrons are shared rather than transferred,
electron
is referred
bonding pairs of electrons (the lone pairs) and two bonding
●
If
two
fluorine
atoms
share
one
electron
each
with
each
other,
each
x
bond.
H
pairs of electrons.
ure 1 Lewis symbols of three to as a covalent
fluorine atom gains one more electron to attain a complete octet of
ments. Nitrogen has five valence
electrons,
whichtoresults
the formation
a covalent
betweento gain or lose
Covalent
bondsNtend
form in
between
atomsofwith
similarbond
Ftendencies
ctrons, chlorine has seven
Nitrogen,
the
two
fluorine
atoms.
This
covalent
bond
is
a
single
bond
and the
2
ence electrons, and boron has electrons. The most obvious examples are the diatomic molecules
2, N
2, O2, Fsymbols
2, Cl2, (a cross and a dot) f
●
For
convenience
use
different
Nitrogen
is
in
group
15,
so
has
five
valence
electrons.
Hence
byweH
shared
pair
can
be
represented
by
a
line:
ee valence electrons
electrons
each of
the two Lewis symbols to signify different
three more electrons nitrogen would
achieve in
a noble
gas
Br2, and acquiring
I2.
forelectrons.
the two elements.
electron configuration with a complete octet of
●
The Lewis symbol for fluorine is:
●
COVALENT BONDING
B
●
deduce the number of
rons of an element you can
number from the periodic
nts. For example, sodium
group 1, so has one valence
um (also s-block) is in group
valence electrons. For the
nts you simply drop the ‘1’
umber to find the number of
ons: silicon (p-block) is in
as four valence electrons.
p-block) is in group 17, so
ence electrons, and so on.
●
F + F
O + O
F F
O O
achieve
If two nitrogen atoms each share three electrons ●withToeach
other,noble
this gas configurations, fluorine and hydrogen
share oneofelectron
with each other, forming a covalen
electron configuration can be achieved and results ineach
the formation
a
F two nitrogen
F
O Obond and the shared pair ca
covalent
covalent bond between the
atoms. ThisThis
covalent
bondbond
is a is a single
represented
a Lewis
line. structure of O there are a total of
triple bond and the three shared pairs can be represented
by three
lines:
●
Note that
in by
this
2
Note that in this Lewis structure of F2 there are a total of six nonbonding
pairs
of
electrons
(the
lone
pairs) and two bo
x one
bonding +
pairs
(often called lone pairs) and
xF
N of electrons
N N
N
H
F
pairsHof +
electrons.
bonding pair of electrons.
Oxygen, O2
N
N
Nitrogen, N2
H
F
is in group 15, so has five valence electrons. Hen
Note that
in group
this Lewis
structure
of N2 there
are a Nitrogen
total
of two
●●
Oxygen
is in
16, so
has six valence
electrons.
Hence
by non2
acquiring
three more electrons nitrogen would achieve a
bondingtwo
pairs
of electrons
lonewould
pairs)attain
and three
bonding
acquiring
more
electrons, (the
oxygen
a noble
gas
electron
configuration
with a complete octet of electrons.
pairs ofconfiguration
electrons. with a complete octet of electrons.
electron
●
If two nitrogen atoms each share three electrons with each ot
●
If two oxygen atoms each share two electrons with each other, this
electron configuration can be achieved and results in the form
Hydrogen
fluoride,
HF
electron configuration can be achieved and results in the formation of a
covalent
bond
the two nitrogen atoms. This covalent
●
Fluorinebond
is in group
17,the
so has
valence
Hence
by between
covalent
between
twoseven
oxygen
atoms.electrons.
This covalent
bond
is a
triple
bond
and
the
three
shared pairs can be represented by
acquiring
one and
more
electron,
fluorine
a nobleby
gastwo
electron
double
bond
the
two shared
pairswould
can beattain
represented
lines.
configuration with a complete octet of electrons. Hydrogen is in group
BILAL HAMEED
COVALENT
+ more
N
N N BONDING
N one
1, so has just one valence electron. Hence by acquiring just
electron, hydrogen would attain the noble gas configuration of helium.
●
N
N
192
COVALENT BONDING
attraction
electron configuration of each atom is then like that of neon – the nearest
noble gas (Figure 6.17).
Chemists draw a line between symbols to represent a covalent bond
҆ (Section
1.2), so they write a fluorine molecule as F−F. This is the structural formula,
.
which shows the atoms and bonding. The molecular formula +
of fluorine is F+
2
fluorine atoms
҆
fluorine molecule
shared electrons
Covalent bonding
Definition
A covalent bond involves a shared
pair of electrons between two atoms.
In a normal covalent bond, each atom
protons
contributes one electron to the
shared pair.
AF covalent bondF is the electrostatic force
of attraction
between the positively charged
F
F
nuclei of both atoms and their shared pair(s) of electrons.
Covalent bondi
Figure 6.17 "
\gZViZhiViigVXi^dc
CovalentaZVhiViigVXi^dc
bonding in a fluorine molecule.
ZaZXigdch^cWdcY
Ä
electron
configuration of each atom is then like that of neon – the nearest
Definition
Ä
Ä
noble
gas (Figure 6.17).
Covalent bonds also link the atomsÄ in non-metal compounds andÄ Figure 6.18
Ä
Chemists draw
a
line
between
symbols to represent
(Section
3
3 a covalentÄ bond
shows the covalent bonding in methane.
Ä
Ä
1.2),
so they write a fluorine molecule
as F−F. This is the structural
formula,
cjXaZjh
cjXaZjha
‘Dot-and-cross’ diagrams showing
inand
outer
shells
provide
Ä only the electrons
Ä
H of fluorine is F H.
which
shows the atoms
bonding.
The molecular
formula
2
Ä bonding. Three of these ‘dot-and-cross’
simple way of representing covalent
ig^eaZWdcY
YdjWaZWdcY
h^c\aZWdcY
diagrams are shown in Figure 6.19. Rememberfluorine
from atoms
Section 1.2 that each
fluorine molecule
non-metal usually forms
the sameI]ZbdgZZaZXigdchbV`ZjeVXdkVaZciWdcY!i]Z\gZViZgi]ZViigVXi^dcWZilZZc
number of covalent bonds in3all its
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;^\jgZ(#&-
compounds. This should
help you predict the structures of different molecules
i]ZZaZXigdchVcYi]ZcjXaZ^!VcYi]ZgZ[dgZi]Zhigdc\Zgi]ZWdcY#
(Table 1.1).
F
F
F
A covalent bond involves a shared
pair of electrons between two atoms
In a normal covalent bond, each atom
contributes one electron to the
shared pair.
F
H
H
It should also
H be noted that
H triple bonds are shorter than double
7dcY AZc\i]$cb 7dcYZcZg\n$
than
bonds. This is, again, due to greater
N single
H O are shorter H
H
Cl Cl bonds, which
&
Figure 6.17`?bda
"
molecule, CH4
Covalent bonding in a fluorine molecule
attraction between the bonding electrons and the nuclei when there methane
are
8Ä8
%#&*)
()more electrons in the bond.
Figure 6.18 !
H
H
828
%#&()
+&'
Covalentthe
bonding
in methane.
Covalent
bonds
also link thejust
atoms
in non-metal
and Figure
6.18
In general,
when we
are comparing
single
bonds, the compounds
longer
2
88
%#&'%
-(,
shows
in methane.
is. covalent
Data
two
in the periodic table are shown
Cl
Cl bond theHweaker
O itthe
H for
N bonding
H groups
a %#&)( H
H
8ÄD
(+%
in Table 3.4. ‘Dot-and-cross’ diagrams showing only the electrons in outer shells provide
simple way of ammonia
representing covalent bonding. Three of these ‘dot-and-cross’
water
chlorine
82D
%#&''
,)(
diagrams are shown in Figure 6.19. Remember from Section 1.2 that each
Figure 6.19 !
<gdje) non-metal usually forms the
<gdje,
same number of covalent bonds in all itsIVWaZ(#( I]ZgZaVi^dch]^eWZilZZccjbWZg
C
‘Dot-and-cross’ diagrams show the single covalent bonds in molecules. A simpler
compounds.
This
should
help
you
predict
the
structures
of
different
molecules
d[WdcYhVcYWdcYaZc\i]$higZc\i]#
7dcY
AZc\i]$cb
:cZg\n$
7dcY
AZc\i]$cb
:cZg\n$
way of showing the bonding in molecules is also included. This shows each covalent
bond as a line between two symbols.(Table 1.1).`?bda&
`?bda&
H
H
H ')'
8Ä8
%#&*)
()8aÄ8a H
%#&..
H O
H N H
Cl Cl
Multiple bonds H^ÄH^
%#'(*
''+
7gÄ7g
%#''&.(
Length
‘DOT҆AND҆CROSS’ DIAGRAMS
methane molecule, CH4
One shared pair of electrons
makes%#')&
a single bond.&-Double bonds
and triple
single bonds
<ZÄ<Z
>Ä>
%#'+,
&*&
decreasing
bonds with two or three shared pairs are also possible.
Figure 6.18 !
H
H
double bonds
Covalent bonding
in methane.
IVWaZ(#)
I]ZgZaVi^dch]^eWZilZZcaZc\i]d[WdcYhVcYWdcYhigZc\i]#
There are two covalent
bonds
between the two oxygen atoms in an oxygen
length
triple bonds
molecule and two covalent bonds between both
the
oxygen
atoms
and
the
Cl
Cl
H
O
H
N
H
carbon atom in carbon dioxide
(Figurethe
6.20).
two 4,
electron
pairs
arethat the single bond
If we consider
dataWhen
for group
it can be
seen
chlorine
water
ammonia
involved in the bonding,
therethe
is aelements
region ofgets
high
electron
density
between
Bonding
and structure
between
weaker
as the
bond
gets longer.
This
is because,
the two atoms joined as
bythe
a double
bond.
6.19 the
! electron pair in the covalent bond is further
atomsFigure
get bigger,
‘Dot-and-cross’
show the
covalent
bonds
in molecules.
away from the
nuclei of thediagrams
atoms making
upsingle
the bond.
If the
electron
pair A simpler
H
H
way of showing the bonding
in molecules
is also included. This shows each covalent
Lone pairs of elec
is further O
away
it istwo
Clesssymbols.
Cstrongly attracted and the covalent
C from
O O
Oas athe
bond
linenuclei
between
H
H
N
N
H
C
C
H
In many molecules an
bond is weaker (Figure 3.19). A similar trend can be seen down group 7.
oxygen
carbon dioxide
Multiple bonds
ethene
One shared
electrons
makes
a single
bond.
Double bonds and triple
The trends should
only pair
reallyof
compared
down
a group,
as elements
Hbe
H
N
N
H
C
C
H
three
areshell
also electrons,
possible. and
O bonds
C
Owith
C shared
C ofpairs
in the same
group
have two
the or
same
number
outer
There are two H
covalent bonds between the two oxygen atoms in an oxygen
therefore anymolecule
effects due
to effective nuclearHcharge or shielding
are
Figure
and two covalent bonds between both
the 6.21
oxygen!atoms and the
most
similar.
In
general,
comparisons
such
as
this
are
most
useful
andelectron
molecules
with triple
covalent
Figure 6.20 !
carbon atom in carbon dioxide (Figure 6.20).Two
When
two
pairs are
bonds.
Three molecules with double
covalent
bonds.
valid when
similar
molecules,
bonds
or
compounds
are
considered.
involved in the bonding, there4is a region of high electron density between
O
O
the two atoms joined by a double bond.
V
\gZViZgViigVXi^dc
Ä
W
O O
O C O
oxygen
carbon dioxide
H
Ä
H
C C
H
H
H
H
H
XadhZgidi]ZcjXaZ^i]Vc^cW#
Figure 6.20 !
Three molecules with double covalent bonds.
water and ammoni
77
H
H
aZhhViigVXi^dc
O
O
O
C
O
C
C
BILAL HAMEED ;^\jgZ(#&. I]ZWdcY^cV^hVh]dgiZgWdcY^cl]^X]i]ZWdcY^c\ZaZXigdchVgZ
● affect the shapes of
● form dative covale
● are important in th
H
ethene
Ä
Ä
outer shells which are
call these ‘lone pairs’
O
Figure 6.22 !
Molecules and ions with
COVALENT BONDING
Dative covalent b
193
COVALENT BONDING
The bond is the force of attraction between the pair of electrons and the
two positive nuclei of the atoms.
This increased negative charge in the centre holds the two positively
charged nuclei together, thus forming the bond.
Depending on the number of electrons involved, the bonding is classified
as single double and triple.
H H
H Cl
O O
N N
5
COVALENT COMPOUNDS
H
H C H
H N H
H
H
H O H
6
BILAL HAMEED
COVALENT BONDING
194
SKILL CHECK 1
Draw dot and cross diagrams of PCl3, N2, CS2, C2H4, NH3, H2O2, N2H4,
SO3
magnesium has a higher melting point than sodium. The first o
is that magnesium forms a 2+ ion, whereas sodium forms a 1+
XdkVaZci
WdcY
This means that the electrostatic attraction between the ions an
delocalised electrons is stronger in magnesium. Also, there are t
delocalised electrons per atom in magnesium, so there will be
number of electrostatic attractions between the ions and the de
electrons. Thirdly, the Mg2+ ion (65 pm) is smaller than the Na
(98 pm), and therefore the delocalised electrons are closer to th
of the positive ion in magnesium and more strongly attracted.
Silicon has a very high melting point because it has a giant
H^Vidb
7 structure, and covalent bonds must be broken when it is melte
bonds are very strong, and a lot of energy is required to break
There are various allotropes (structural modifications) of ph
XdkVaZci
WdcY
The data given in Table 4.1 are for white phosphorus, which i
magnesium
has a higher
melting
than
tetrahedra
most common form
of phosphorus.
It consists
of P4point
is that
magnesium
formsPa4 2+
ion, whereas
tetrahedra
also es
van der Waals’ forces
between
them. These
XdkVaZci
WdcY
This means
the Waals’
electrostatic
liquid state, and therefore
only that
van der
forces attraction
are brokenb
delocalised
electrons
is stronger
in and
magnesiu
phosphorus is melted.
Van der Waals’
forces
are weak
little
EVidb
delocalised
electrons
per
atom
in
magnesium
required to break them. The melting point of phosphorus is th
EXAMPLES OFmuch
MOLECULES
number
of electrostatic
attractions
between
lower than that
of silicon,
because, when
phosphorus
is m
2+
ion
(65
pm)
is
s
XdkVaZci
electrons.
Thirdly,
the
Mg
weak van der Waals’ forces are broken rather than covalent bon
WdcY
(98 pm),has
anda therefore
the delocalised
electrw
magnesi
Sulfur, like phosphorus,
covalent molecular
structure,
of the positive
ionAinsulfur
magnesium
and
more
ishowev
that m
der Waals’ forces between
molecules.
molecule,
XdkVaZci
Silicon
has
a
very
high
melting
point
bec
WdcY cantly higher relative
Thismol
me
of eight atoms and, as S8 has a signifi
H^Vidb
structure,
and
covalent
bonds
must
be
broke
S8 mole
than P4, the van der Waals’ forces are stronger between delocali
bonds
are
very
strong,
and
a
lot
of
energy
HVidb
breakist
between P4 molecules. More energy is thus required todelocali
are various
m
XdkVaZciWaals’ forces betweenThere
number
thanallotropes
between (structural
P4 molecule
S8 molecules
WdcY
The
datathan
given
in Table 4.1 are for white
ph
electron
has a higher melting
point
phosphorus.
most
common
of mass
phosphorus.
It consi
(98
pm)
Going fom sulfur
to argon,
the form
realtive
of the molecule
van
der
Waals’
forces
between
them.
These
of (M
ther p2P
the case of argon) decreases. Sulfur exists as S8 molecules
liquid state,
and therefore
onlyasvan
der Silico
Waal3
(Mr 70.90)
and argon
atoms
(A
chlorine as Cl2 molecules
r
phosphorus
is
melted.
Van
der
Waals’
forces
H^Vidb
the decrease in melting point from sulfur to argon can structur
be attria
EVidb
to break
them.
melting
point
bonds
decrease in van derrequired
Waals’ forces
as the
massThe
of the
particles
deao
much lowerXdkVaZci
than that of silicon, because,
wh
Ther
XdkVaZci8
weak van der
Waals’ forces are broken
rather
WdcY
The
dat
WdcY
Sulfur, like phosphorus, has a covalent
most mo
co
der
Waals’
forces
between
molecules.
A
sulfu
van der
AZVgc^c\dW_ZXi^kZh
cantlysth
of eight atoms and, as S8 has a signifiliquid
Understand
that
elements
in
stronge
than P4, the van der Waals’ forces arephospho
r
8]Zb^XVaegdeZgi^Zhd[ZaZbZcih^ci]ZhVb
EVidb
the same group have similarHVidb
BILAL HAMEED
COVALENT
BONDING
is thus
between P4 molecules. More energyrequired
\gdje
chemical properties
Waals’ forces between S8 molecules than
muchbet
lo
Describe
some
reactions
of
The
reactions
of
an
atom
are
determined
by
the
number
of ele
has a higher
melting point than phosphorus
XdkVaZci
r
)#( 8]Zb^XVaegdeZgi^Zhd[ZaZbZci
\gdje&VcY\gdje,
195
SKILL CHECK 2
When barium metal burns in oxygen, the ionic compound barium
peroxide, BaO2, is formed.
Which dot-and-cross diagram represents the electronic structure of the
peroxide anion in BaO2?
9
SKILL CHECK 3
Dicarbon monoxide, C2O, is found in dust clouds in space. Analysis of
it shows that the sequence of atoms in this molecules is C-C-O. All
bonds are double bonds and there are no unpaired electrons.
How many lone pairs of electrons are present in a molecule of C2O?
A 1
B 2
C 3
D 4
10
BILAL HAMEED
COVALENT BONDING
nd
ells
196
COVALENT BONDING
Atoms share electrons to get the nearest noble gas electronic configuration
Some don’t achieve an “octet” as they haven’t got enough electrons
e.g. Al in AlCl3
Others share only some — if they share all they will exceed their “octet”
e.g. NH3 and H2O
Atoms of elements in the 3rd period onwards can exceed their “octet” if they
wish as they are not restricted to eight electrons in their “outer shell”
e.g. PCl5 , SO2 , SO3 and SF6
Only in period 2 are the elements restricted to form an octet. However, in period 3,
more an 8 electrons can be taken in the outermost shell due to the s, p and d
orbitals which take up 2, 6 and 10 electrons respectively.
11
a Boron trifluoride
H
DEFYING OCTET
F
H
+
F trifluoride
B
a 3 Boron
H
H
F +
three
fluorine
atoms (2,7)
B
boron
atom (2,3)
three
boron
3
H
H
C
H
Sulfur hexafluoride
b fluorine
atom (2,3)
HH
H
C
atoms (2,7)
H
H
O
H
H
O
F
F
F
B
F
boron trifluoride
molecule
F
B
F
F
boron trifluoride
molecule
F
F
F
Fsix +
fluorine
atoms (2,7)
S
sulfur
atom
(2,8,6)
S
Fsulfur hexafluoride
F
12
molecule
F
F
six
sulfur
sulfur hexafluoride
S
F
F
F
F
F
B
F
F
S
F
F
F
F
F
F
F
F
S
F
F
F
F
fluorine
atom
molecule
Figure
4.10 Dot-and-cross
diagrams for
a boron trifluoride, BF3, and b
atoms
(2,7)
(2,8,6)
sulfur hexafluoride, SF6.
BILAL HAMEED
H
B
F +
S
hexafluoride
b6 Sulfur
6
H
F
F
COVALENT BONDING
Figure 4.10 Dot-and-cross diagrams for a boron trifluoride, BF3, and b
sulfur hexafluoride, SF6.
elements
Figure 3.10 Dot-and-cross diagrams for
phosphorus trifluoride, PF3, and phosphorus
pentafluoride, PF5
F
F
Unlike elements in the second row of the Periodic Table, those in the third and
P p orbitals.
F
F in bonding, as well as their s and
subsequent rows can useF their197
dP orbitals
They can therefore form more than four covalent bonds to other atoms.
Like nitrogen,
F
2 2
6 2
3
phosphorus (1s 2s 2p 3s 3p ) has
F
F five electrons in its valence shell. But because
it can make use of five orbitals (one 3s, three 3p and one 3d) it can use all five of
its valence-shell electrons in bonding with fluorine. It therefore forms phosphorus
In a similar way, sulfur can use all its valence-shell electrons in six orbit
pentachloride, PF5, as well as phosphorus trifluoride, PF3 (see Figure 3.10).
three 3p and two 3d) to form sulfur hexafluoride, SF6. Chlorine does no
however. The chlorine atom is too small for seven fluorine atoms to asse
03_10 Cam/Chem AS&A2
it.
Chlorine does, though, form FClF3 and ClF5 in addition to ClF.
DEFYING OCTET
3.10 Dot-and-cross diagrams for
rus trifluoride, PF3, and phosphorus
oride, PF5
Barking Dog Art
F
P
Worked example
F
F
P
F
Draw a diagram to showFthe bonding in chlorine pentafluoride, ClF5.
F
F
Answer
In a similar
way,
sulfur can use all its valence-shell electrons in six orbitals (one 3s,
Figure
3.11
F
F
three 3p and two 3d) to form
sulfur hexafl
uoride, SF6. Chlorine does not form ClF7,
however. The chlorine atom is Cl
too small for seven fluorine atoms to assemble around
03_10 Cam/Chem AS&A2
it.
Chlorine does, though,F form ClF3 and
F ClF5 in addition to ClF.
Barking Dog Art
F
Worked example
Draw a diagram to show the bonding in chlorine pentafluoride, ClF5.
13
Now try this
Answer
03_11 Cam/Chem AS&A2
Figure 3.11
F
F
1 Sulfur forms a tetrafluoride, SF4. Draw a diagram to show its bonding.
2 a How many valence-shell electrons does chlorine use for bonding in chlo
Cl
trifluoride, ClF3?
b
So
how many electrons are left in the valence shell?
F
Other ways of representing the bonding in carbon monoxide are shown
or C
C
Oatom
c So how many lone pairs of electrons are there on the chlorine
in c
F
in Figure
3.24.
trifluoride?
F
Barking Dog Art
;^\jgZ(#') Di]Zgl
Dative covalent bonds are important in the bonding in transition
metal complexes.
=A WdcY^c\^cXVgWdcb
MOLECULAR ORBITALS
Now try this
03_11 Cam/Chem AS&A2
3.6 The covalency table
Covalent bonding is brought about by the atomic orbitals of each atom overlapping
1 Sulfur
forms a tetraflAs
uoride,
SF4. Draw
a diagram
to show
8dkVaZciWdcYh^ciZgbhd[dgW^iVah
explained
in sections
3.4 and
3.5, its
thebonding.
number of covalent bonds formed
Barking
Dog Art
with each other,
coalescing
forming
a molecular
2 a
How manyand
valence-shell
electrons
does orbital.
chlorine use for bonding in chlorine
thetwo
number
of electrons
for bonding, and the numbe
A covalent bond isdepends
formed on
when
atomic
orbitals, available
each
trifluoride, ClF3?
Only individually filled
orbitals one
takeshell
partorbitals
in such
bonding,
so
the resultant
it can
use to
putthat
the
electrons
into.toFor many elements, the nu
containing
electron,
overlap.
The
atomic
orbitals
combine
b So how many electrons are left in the valence shell?
bonds
formed
is
a
fi
xed
quantity
and
is
termed
the
covalency of the eleme
form
a molecular
inInelectrons
which
both
up atom
(Figures
molecular orbitalchas
only
two electrons.
the molecular
the
two
electrons
So
how
many
loneorbital
pairs of
areelectrons
thereorbital
on are
thepaired
chlorine
in chlorine
3.25
and
3.26). These
electrons
belong
totwo
bothbonded
atoms
simultaneously.
related
tonuclei
its group
number
in
the Periodic
Table. As shown above, elemen
trifl
uoride?
circulate around and
are
attracted
to two
of the
atoms.
Vidb^XdgW^iVa
48
15 to 17 in the third and subsequent rows of the Periodic Table (Period 3 an
bdaZXjaVgdgW^iVa
use a variable number of electrons in bonding. They can therefore display m
covalency. Table 3.1 shows the most usual covalencies of some common el
Vidb^XdgW^iVa
3.6 The covalency table
As explained in sections 3.4 and 3.5, the number of covalent bonds formed by an atom
depends on the
number of electrons available for bonding, and the number of valencecjXaZjh
shell orbitals it can use to put the electrons into. For many elements, the number of
181333_03_AS_Chem_BP_044-068.indd 48
bonds formed is a fixed quantity and is termed the covalency of the element. It is often
'ZaZXigdcheV^gZYje
related dgW^iVaXdciV^ch
to its group number in the
As shown above, elements in Groups
14 Periodic Table.
^cVbdaZXjaVgdgW^iVa
&ZaZXigdc
15 to 17 in the third and subsequent rows of the Periodic Table (Period 3 and higher) can
;^\jgZ(#'* I]ZdkZgaVed[ildhdgW^iVahid[dgbVXdkVaZciWdcY#
use a variable number of electrons in bonding. They can therefore display more than one
covalency. Table
3.1 shows the most usual covalencies of some common elements.
edgW^iVa
BP_044-068.indd 48
BILAL HAMEED
COVALENT BONDING
;^\jgZ(#'+ I]ZdkZgaVed[ildedgW^iVahid[dgbVXdkVaZciWdcY#
30/09/14 5:5
198
ORBITAL OVERLAP
Covalent bonds are formed when orbitals, each containing one electron, overlap.
This forms a region in space where an electron pair can be found; new molecular
orbitals are formed.
molecular orbital
orbital
containing 1
electron
orbital
containing 1
electron
overlap of orbitals provides a
region in space which can contain
a pair of electrons
The space that these shared electrons move within is called a molecular
orbital. A molecular orbital is made when two atomic orbitals overlap.
15
SIGMA BONDS10/23/04 3:58 PM Page 221
CVM06_ch06_COV_220-221_final
A Sigma bond (σ bond) is formed when orbitals overlap head-on in a
covalent bond.
In a sigma bond, the electron density is concentrated in the orbital
INTERPRETING GRAPHICS
overlap volume between the two nuclei.
Directions (10–13): For each question below, record the correct answer o
A sigma bond can be rotated without
breaking
bond.
separate
sheet ofthe
paper.
Use the diagram below to answer question 10.
The region where the the two orbitals
overlap is the molecular orbital
Molecular orbital
which contains electrons of the sigma bond.
–
–
16
0 The diagram above best represents which type of chemical bond?
F. ionic
G. metallic
BILAL HAMEED
H. nonpolar covalent
I. polar covalent
COVALENT
BONDING
The table below shows the connection
between
electronegativity and bond
strength (kilojoules per mole). Use it to answer questions 11 through 13.
Electronegativity Difference for Hydrogen Halides
199
SIGMA BONDS
Sigma bonds can form by overlap of two s-orbitals, and s-orbital and porbital or two p-orbitals.
1s
1s
H
H
1s
H҆H
3p
H
Cl
H ҆ Cl
3p
3p
Cl
Cl
Cl ҆ Cl
17
PI BONDS
A pi bond (π bond) is formed when orbitals overlap sideways in a
covalent bond. Pi bonds are found in molecules with double and triple
bonds.
2p
2p
O
O
18
BILAL HAMEED
COVALENT BONDING
200
PI BONDS
A single π bond is drawn as two electron clouds, one arising from each
lobe of the p orbitals.
The two clouds of electrons in a π bond represent one bond consisting of
a total of two electrons.
19
PI BONDS
20
BILAL HAMEED
COVALENT BONDING
Mutual repulsion of the electrons in the three σ bonds around each carbon atom
tends to place them as far apart from one another as possible (see page 51). The
predicted angles (H¬C¬H = H¬C¬C =201
120°) are very close to those observed
(H¬C¬H = 117°, H¬C¬C = 121.5°).
The π bond confers various properties on the ethene molecule.
● Being the weaker of the two, the π bond is more reactive than the σ bond. Many
reactions of ethene involve only PI
the breaking
of the π bond, the σ bond remaining intact.
BONDS
● The two areas of overlap in the π bond (both above and below the plane of the
ethene to
be a rigid
molecule, withinno
easy
rotation
of one end
In amolecule)
pi bond, cause
the electron
density
is concentrated
the
orbital
overlap
with respect
the below
other (see
3.41). This
is innuclei.
contrast to ethane, and leads
volumes
abovetoand
theFigure
line joining
the two
to the existence of cis–trans isomerism in alkenes.
A pi bond cannot be rotated with out breaking the bond.
These features are developed further in Topics 12 and 14.
prevents rotation
n double bond.
H
H
H
C
C
H
C
H
H
H
C
H
H
H
no rotation possible
rotation possible around
the single bond
Triple bonds
21
Triple bonds are formed when two p orbitals from each of the bonding atoms overlap
03_41 Cam/Chem AS&A2
sideways, forming two π orbitals (in addition to the σ orbital formed by the end-on
overlap
of the
Barking Dog
Artthird p orbital on each atom). This occurs in the ethyne molecule,
C2H2, and in the nitrogen molecule, N2 (see Figure 3.42).
s in a nitrogen and
a N2
N
N
SKILL CHECK 4
N
N
What is always involved in a carbon-carbon π bond?
py +electrons
pz
py + ppair
A a shared
of
z
/y + /z
H2
Bb aC2sideways
overlap of p orbitals
C delocalized electrons
H
C
C
py + p z
py + pz
H
H
C
C
H
/y + /z
A03_42
single
carbon atom
Cam/Chem
AS&A2can also use two of its p orbitals to form π orbitals with two
separate bonding atoms. This occurs in the
22 carbon dioxide molecule, CO2 (see
Barking Dog Art
Figure 3.43).
bonds in carbon
BILAL HAMEED
O
C
O
O
COVALENT
BONDING
C
O
202
DOUBLE BONDS
A double bond consists of one sigma bond and one pi bond, for example
the C=C double bond in ethene and the O=O double bond in oxygen.
A typical triple bond, for example in nitrogen, consists of one sigma bond
and two pi bonds in two mutually perpendicular planes.
Pi bonds are weaker than sigma bonds.
23
MULTIPLE BONDS - ETHENE
An example of a pi bond is the C ˞ C bond in ethene. This bond consists
of two parts.
One part consists of two p orbitals of carbon overlapping in a sigma
bond.
In the other part of the bond, two p orbitals overlap in a pi bond
24
BILAL HAMEED
COVALENT BONDING
203
MULTIPLE BONDS - ETHENE
25
ETHANE AND ETHENE
26
BILAL HAMEED
COVALENT BONDING
204
TRIPLE BOND - ETHYNE
27
TRIPLE BOND - NITROGEN
28
BILAL HAMEED
COVALENT BONDING
205
TWO DOUBLES - CARBON DIOXIDE
29
SKILL CHECK 5
Which statements about covalent bonds are correct?
A A triple bond consists of one π bond and two σ bonds.
B The electron density is a σ bond is highest along the axis between the
two bonded atoms.
C A π bond restricts rotation about the σ bond axis.
30
BILAL HAMEED
COVALENT BONDING
206
DATIVE COVALENT BONDING
A dative covalent bond (or coordinate covalent bond) is a covalent bond
between two atoms in which the shared electrons are contributed by
only one of the atoms.
e.g. Boron trifluoride & ammonia NH3BF3
Boron has an incomplete shell in BF3 and can accept/share a pair of
electrons donated by ammonia.
H
F
H F
xo
x
o
xo
xo
x
o
xo
xo
H N B xo F
H
F
H F
+
x
o
xo
B F
H N
o
o
o
o
xo
xo
31
DATIVE COVALENT BONDING
H
H N oo
F
B
+
H
F
F
F
H N
B
H
F
F
+
H
H N oo
H
H
+
H+
H N
H
H
H
32
BILAL HAMEED
COVALENT BONDING
W
=
=
=
H3O+ is formed when a lone pair of electrons is donated from the O in
H+:covalent bond has been formed, it is identical to an
HOnce
2O to athe
dative
C
+
+
207
‘ordinary’ covalent
= bond.
D =For example,
= DNH
= 4 can be formed when H
becomes bonded to NH3:
=
=
+
NH3 + H
=
=
+
n NH
4
YVi^kZXdkVaZci
WdcY
2
YVi^kZXdkVaZci
WdcY
=
+
H
does
not
have
any
electrons
with
which
to formprovided.
a covalent bond, but
DATIVE
COVALENT
BONDING
Answer
allcan
thecombine
questions
in the
spaces
]ZVbbdc^jb^dc!Vl^i]
=
= C =
and
BF
to
form
an
adduct
(two
molecules
bonded
NH
a covalent
bond
NH3 3has a lone3 pair of electrons, which can be used to form
Yh]dlcVcYWl^i]cd
=
together):
lZZci]ZineZhd[WdcYh#
(Figure 3.20).
1
Elements and compounds
which have small molecules usually exist as gases or liquids.
kVaZci
YVi^kZXdkVaZci
= bond
; is sometimes =
; as an arrow (Figure 3.21a).
A dative covalent
shown
WdcY
= C formed,
7 a;dative bond
= Cis, however,
7 ;
has been
the same Br
as any
other
ample
of a Lewis
, is aitgas
at room temperature
whereas bromine,
(a) acidChlorine,
Cl 2Once
9Vi^kZXdkVaZciWdcY^c\^cC=
2, is a liquid under the
) #
=
;
=
;
covalent bond. The ammonium ion can be represented as shown in Figure
(see page 314). same conditions.
W
C
=
=
=
8D
..........................................................................................................................................
Normally carbon shares
four electrons toYVi^kZXdkVaZci
form four covalent bonds,
YVi^kZXdkVaZci
]ZhigjXijgZd[XVgWdc=
WdcY
WdcY
and oxygen shares two to form two covalent bonds. If a carbon atom
dÈdgY^cVgnÉXdkVaZciWdcYh
.................................................................................................................................... [2]
I]ZVbbdc^jb^dc!Vl^i]
=
8 D
=
3.21b, in whichYVi^kZXdkVaZci
no distinction is made between
the individual bonds.
YVi^kZXdkVaZci
+
WdcY
WdcY
H3O is formed when a lone pair of electrons is donated from the O in
Explain these observations.
H2O to the H+:
In
BF3 there are only six electrons in the outer shell of the boron;
=
..........................................................................................................................................
therefore there is space
= Dfor the
= boron to
= accept
D = a pair of electrons.
combines
withBF
an3 oxygen
atom with
thean
formation
twomolecules
covalent bonded
can combine
to form
adduct of
(two
NH3 and
cYh]dlcVcYWl^i]cd
bonds,
we get the structure shown
in Figure 3.22. However, in this
together):
33
ZilZZci]ZineZhd[WdcYh#
CO, are isoelectronic, that is they have
(b) The gases nitrogen,
N2, and carbon monoxide,
structure, although
the oxygen atom has a full outer shell (octet), the
the same number of electrons
molecules.
= in their
;
= ;
carbon atom only has six electrons in its outer shell.
=
C
7
;
=
C
7 ;
xample of a Lewis acidBoth atoms can attain an octet if the oxygen atom donates a pair of
= boiling
; point than =
;
CO.
Suggest why N2 has a lower
n 8(see
electrons to carbon in the formation of a dative covalent bond. There
D page 314).
YVi^kZXdkVaZci
is thus a triple bondYVi^kZXdkVaZci
between the two atoms,
made up of two ‘ordinary’
]ZhigjXijgZd[XVgWdc
WdcY
WdcY
..........................................................................................................................................
covalent bonds and one dative covalent bond (Figure 3.23). Both atoms
have
a lone
pairare
ofonly
electrons.
In BF
six electrons in the outer shell of the boron;
3 there
..........................................................................................................................................
therefore there is space for the boron to accept a pair of electrons.
dÈdgY^cVgnÉXdkVaZciWdcYh
XdkVaZciWdcYlZgZ[dgbZY#
8 D
8D DATIVE COVALENT BONDING
....................................................................................................................................
[2]
Normally carbon shares four electrons to form four covalent bonds,
I]ZhigjXijgZd[XVgWdc
and oxygen
shares
to form
two covalent
bonds.below.
If a carbon
(c) A ‘dot-and-cross’
diagram
oftwo
a CO
molecule
is shown
Onlyatom
electrons from outer
ldÈdgY^cVgnÉXdkVaZciWdcYh
#
combines
with
an
oxygen
atom
with
the
formation
of
two
covalent
shells are represented.
bonds, we get the structure shown in Figure 3.22. However, in this
structure, although the oxygen atom has a full outer shell (octet), the
carbon atom only has six electrons in its outer shell.
Both atoms can attain an octet if the oxygen atom donates a pair of
electrons to carbon in the formation of a dative covalent bond. There
is thus a triple bond between the two atoms, made up of two ‘ordinary’
covalent bonds and one dative covalent bond (Figure 3.23). Both atoms
have a lone pair of electrons.
C
8 D
I]ZhigjXijgZd[XVgWdc
ldÈdgY^cVgnÉXdkVaZciWdcYh
kZXdkVaZciWdcYlZgZ[dgbZY#
O
In the table below, there are three copies of this structure.
On the structures, draw a circle round a pair of electrons that is associated with each of
the following.
34
(i) a co-ordinate bond
C
BILAL HAMEED
O
(ii) a covalent bond
C
O
(iii) a lone pair
C
O
COVALENT BONDING
208
AL 2 CL 6
o
o
oo
Cl
o
o
o
o
oo
o
o
Cl
Al
oo
o
o
Cl
oo
Cl
o
o
o
o
Al Cl
+
o
o
o
o
oo
Cl
oo
o
o
o
o
o
o
o
o
o
o
o
o
oo
o
Cl oo Al Cl o
oo
o
o
o
o
o
o
Cl
oo
Al Cl oo
Cl
oo
o
o
Cl
oo
Cl
Al
oo
o
o
Cl
oo
o
o
Cl
Al oo Cl
oo
oo
oo
o
o
oo
Cl oo
o
Cl
o
oo
35
AL 2 CL 6
Cl
Cl
Al
Cl
Cl
Cl
Cl
+
Al
Al
Cl
Cl
Cl
Al
Cl
Cl
Cl
Cl
Al
Cl
Cl
Cl
Al
Cl
Cl
Figure 4.14 A dot-and-cross diagram for an aluminium chloride molecule, Al2Cl6.
Check-up
4 a Draw dot-and-cross diagrams to show the
formation of a co-ordinate bond between
the following:
i boron trifluoride, BF3, and ammonia,
NH3, to form the compound F3BNH3
BILAL HAMEED
ii phosphine, PH3, and a hydrogen ion,
H+, to form the ion PH4+.
b Draw the displayed formulae of the
products formed in part a. Show the co-
36
Bond
C C
C C
C O
C O
Bond energy / kJ mol−1 Bond length / nm
350
0.154
610
0.134
360
0.143
740
0.116
Table 4.1 Examples of values for bond energies and bond lengths.
COVALENT BONDING
Bond strength can influence the reactivity of a
209
DATIVE COVALENT BONDING
A dative covalent bond differs from covalent bond only in its formation
Both electrons of the shared pair are provided by one species (donor) and
it shares the electrons with the acceptor
Donor species will have lone pairs in their outer shells
Acceptor species will be short of their “octet” or maximum.
Chemists call it a dative covalent bond because the word dative means
‘giving’ and one atom gives both the electrons to make the bond. Once
formed, there is no difference between a dative bond and any other
covalent bond.
37
BOND ENERGY
The strength of a covalent bond is measured by the bond energy.
The bond energy is the amount of energy required to break a covalent
bond, per mole of bonds. The greater the bond energy, the stronger the
bond.
Very large bond energies can make molecules unreactive. The nitrogen
molecule (N2) is unreactive because of the very large N≡N bond energy
of 944 KJ mol҆1.
38
BILAL HAMEED
COVALENT BONDING
Ä
ig^eaZWdcY
YdjWaZWdcY
h^c\aZWdcY
210
;^\jgZ(#&- I]ZbdgZZaZXigdchbV`ZjeVXdkVaZciWdcY!i]Z\gZViZgi]ZViigVXi^dcWZilZZc
i]ZZaZXigdchVcYi]ZcjXaZ^!VcYi]ZgZ[dgZi]Zhigdc\Zgi]ZWdcY#
It should also be noted that triple bonds are shorter than double
bonds, which are shorterBOND
than singleENERGY
bonds. This is, again, due to greater
attraction between the bonding electrons and the nuclei when there are
moreforming
electronsreleases
in the bond.
Bond
energy, and is therefore exothermic.
In general, when we are comparing just single bonds, the longer the
Bond
requires
energy,
and
is therefore
endothermic.
the
bondbreaking
the weaker
it is. Data
for two
groups
in the periodic
table areFor
shown
same
two3.4.
atoms, triple bonds are stronger than double bonds which in
in Table
turn are stronger than single bonds.
<gdje)
<gdje,
7dcY
AZc\i]$cb
8Ä8
%#&*)
H^ÄH^
%#'(*
<ZÄ<Z
%#')&
:cZg\n$
C҆C
`?bda&
7dcY
350
AZc\i]$cb
()C˞C
''+
C≡C
&--
8aÄ8a
610
7gÄ7g
>Ä>840
%#&..
')'
%#''-
&.(
%#'+,
&*&
:cZg\n$
`?bda&
IVWaZ(#) I]ZgZaVi^dch]^eWZilZZcaZc\i]d[WdcYhVcYWdcYhigZc\i]#
39
7dcY
AZc\i]$c
8Ä8
%#&*)
828
%#&()
28
8
%#&'%
8ÄD
%#&)(
82D
%#&''
IVWaZ(#( I]ZgZaVi^d
d[WdcYhVcYWdcYaZ
Length
single bonds
double bonds
triple bonds
If we consider the data for group 4, it can be seen that the single bond
between the elements gets weaker as the bond gets longer. This is because,
as the atoms get bigger, the electron pair in the covalent bond is further
away from the nuclei of the atoms making up the bond. If the electron pair
is further away from the nuclei it is less strongly attracted and the covalent
bond is weaker (Figure 3.19). A similar trend can be seen down group 7.
BOND LENGTH
The trends should only really be compared down a group, as elements
Theincovalent
length
distance
between
nuclei
of the and
two atoms
the samebond
group
have is
thethe
same
number
of outerthe
shell
electrons,
therefore
ectscovalent
due to effbonds
ective nuclear charge or shielding are
linked
by oneany
or eff
more
most similar. In general, comparisons such as this are most useful and
For the same two atoms, triple bonds are shorter than double bonds which in
valid when similar molecules, bonds or compounds are considered.
turn are shorter than single bonds.
V
W
\gZViZgViigVXi^dc
Ä
Ä
Ä
Ä
aZhhViigVXi^dc
;^\jgZ(#&. I]ZWdcY^cV^hVh]dgiZgWdcY^cl]^X]i]ZWdcY^c\ZaZXigdchVgZ
XadhZgidi]ZcjXaZ^i]Vc^cW#
40
BILAL HAMEED
COVALENT BONDING
ionisation energy / kJ mol–1
950
1800
2700
4800
6000
12300
211
What could be the formula of the chloride of X ?
A
C XCl3 BONDING
D XCl4 WS 1
XClCOVALENT
2
B
XCl
SECTION A
6 1 Which molecule contains only six bonding electrons?
A
B
C2H4
C
C2F6
D
H2O
NF3
[W’03 Q6]
5
7 Chemists have been interested in the properties of hydrogen selenide, H2Se, to compare it with
5 2 Which diagram describes the formation of a π bond from the overlap of its orbitals?
‘bad egg’ gas hydrogen sulphide, H2S.
3
4
A
Which
set of data would the hydrogen selenide molecule be expected to have?
The diagrams show the possible paths of subatomic particles moving in an electric field in a
vacuum.
number of lone pairs
bond angle
+
+
+
on Se atom
B
A
1
104 °
electrons
neutrons
protons
B
2
104 °
C
C
D
–
2
2
–
109 °
1
–
2
180 °
3
Which diagrams are correct?
A
1 and 2 only
D
B
1 and 3 only
C
2 and 3 only
D
1, 2 and 3
[W’06 Q5]
6 For an ideal gas, the plot of pV against p is a straight line. For a real gas, such a plot shows a
5 3 The CN– ion is widely used in the synthesis of organic compounds.
deviation from ideal behaviour. The plots of pV against p for three real gases are shown below.
[Turn over
9701/1/O/N/03
What is the pattern of electron pairs in this ion?
3
The gases represented are ammonia, hydrogen
and nitrogen.
3
lone
on of lone
pairs on
What isbonding
the mostpairs
likelyofshape
of apairs
molecule
hydrazine,
N2H4?
electrons
carbon atom
A
B
A
H
2
B 90°
1
H
H
2
N
C
N
3
H
D
H
3
HpV N
H
90° 2
nitrogen atom
X
H
1
Y
H
1120°
N1
N1 H
H
D
C
1
N
H
2
Z
idealHgas
H
107°
N
N
H
In whichbottles
species
underlined
atomfrom
haveaan
incomplete
shell? structure.
64 4 Plastic
fordoes
‘fizzythe
drinks’
are made
polymer
with outer
the following
0
+
B CH
C Cl
D H2Cl C•
A Al 2Cl 6
0 3
p 2O
H
X
What are the identities of the gases X, Y and Z?
C
5 Which solid contains more than one kind of bonding?
A
iodine
X
B A silicon
dioxide
ammonia
Y
nitrogen
C
Z
H
H
hydrogen
H
[J’12 1 Q11]
n
C B sodium
chloride
hydrogen
nitrogen
The
ability
of
the polymer
to prevent ammonia
escape of carbon dioxide through the wall of the bottle
CEDAR COLLEGE
BONDS WS 1
depends on the ability of the group X to form hydrogen bonds with theCOVALENT
carbon dioxide
in the drink.
D C zinc nitrogen
ammonia
hydrogen
Which
X best prevents
loss of carbon
dioxide?
D group
nitrogen
hydrogen
ammonia
What is element X?
C increasing the temperature
A
D
B
sodium
removing the ammonia as it is formed
212
magnesium
aluminium
25 C
In which
species does the underlined atom have an incomplete outer shell?
D
A
phosphorus
BF3
B
CH3–
C
F 2O
D
H3O+
[M’1 Q12]
15 Aluminium chloride sublimes at 178 C.
3 6 The value of the second ionisation energy of calcium is 1150 kJ mol–1.
o
Which structure best represents the species in2 the vapour at this temperature?
Which equation correctly represents this statement?
A
Ca(g)
A
Ca2+(g) + 2e− ;
B
Section
A –1
H o = +1150
kJ mol
C
D
3+
+
2+ are
For Cl
eachCa
question
there
B,–1
C and D. Choose
the one you
Cl H o =answers,
Al + 3Cl
Cl kJA,
Cl
Alconsider
(Cl –)3 to
(g)Cl Ca
(g)Cl+ four
e− ; possible
+1150
mol
B
be correct.
Al+(g)  Ca
Al2+(g) + e− ;
H o Al
= –1150 kJAlmol–1
C Ca
Use of the Data Booklet may be appropriate for some questions.
Cl
D Ca(g) Cl Ca2+(g) Cl
+ 2e− ; Cl H o = –1150
Cl kJ mol–1
Cl
[S’06 Q15]
1 Which equation shows the reaction that 7occurs during the standard enthalpy change of
atomisation
of for
bromine?
4
The
equation
a reaction
is shown.
16 Use
of
the Data
Booklet
is relevant
to this question.
14
Which
element
is expected to show the greatest tendency to form some covalent compounds?
7
A Brmass
2(l) → 2Br(g)
–1
What
of solid residue
can
from; ∆H
the=thermal
of 4.10 g of
+ 21be
O2obtained
(g)
H2O(l)
x kJ moldecomposition
H2(g)
A aluminium
anhydrous
nitrate?
B Br2(g) calcium
→ 2Br(g)
B calcium
Which
isgfully correct
thisgreaction? D 2.25 g
1 pair
C for1.40
A
C 0.70
Br2g(l) of→descriptions
Br(g)B 1.00
2
C magnesium
1
Br2(g) →ofBr(g)
D
enthalpy change
value of x
2 type(s)
D sodium
17 What happens when chlorine is bubbled through aqueous potassium iodide?
[S’13 Q17]
A
formation only
positive
A
Chlorine
is
oxidised
to
chloride
ions.
215
theData
correct
number
of bonds
each
type in the Al 2Cl 6 molecule?
Use ofisthe
Booklet
is relevant
to of
this
question.
8 What
B
formation only
negative
B Hydrochloric acid is formed.
The
of fossil
fuels
is a major source of
increasing atmospheric carbon dioxide, with a
C combustion
combustion,
formation
positive
co-ordinate
covalent
rise
global
warming.
Another
significant
contribution to carbon dioxide levels
Cconsequential
Iodide ions
are in
oxidised
to
iodine.
(dative covalent)
comes
from
the thermalformation
decomposition of limestone,
in the manufacture of cement and of lime for
D
combustion,
negative
Dagricultural
Potassium
iodide is reduced to iodine.
purposes.
A
18
6
1
Cement
works6roast 1000 million tonnes
of limestone per year and a further 200 million tonnes is
B
2
The
emissions
a power
roasted
in kilns from
to make
lime. station contain about 14 tonnes of SO2 per hour from the oxidation
C 2 contained
7 in the coal.
0
of FeS
What
annual mass 1output of carbon dioxide (in million tonnes) from these two
D is the total
7
processes?
What
is the most practical way of preventing the SO2 from being released into the atmosphere?
[S’16 Q2]
4
440 the gases and
B the
527SO2 will liquefy
C 9701/13/M/J/12
660 can be removed.
D 880
© UCLES
and
AA 2012
Cool
H5)4, hasisbeen
used
additive.
36 9 Tetraethyl
Pb(C2chloride
When solidlead,
aluminium
heated,
Al 2as
Cl 6aispetrol
formed.
B Dissolve the ionic FeS2 in hexane.
theofpercentage
by
in tetraethyl
lead?
Whichisbonding
is present
in mass
Aland
? carbon
16 What
Properties
chlorine,
iodine
their
compounds
are compared.
2Cl 6of
C Pass the emissions through a bed of calcium oxide.
A 10.2
covalent
co-ordinate
(dative
covalent)
A
B is14.9
29.7
D 32.0
Property
Q forand
chlorine
smaller
than
for C
iodine.
D Pass the gases through concentrated sulphuric acid to dissolve the SO2.
B covalent only
What is property Q?
C ionic and co-ordinate (dative covalent)
A
oxidising ability of the element
© UCLES 2006
9701/01/M/J/06
D ionic only
B solubility of the silver halide in NH3(aq)
7
[W’16 2 Q6]
C
strength
of van
derH–X–H
Waals’bond
forces
between
the molecules of the element
In which
hydride
is the
angle
the smallest?
D
A
thermal
stabilityBof the
halide
CHhydrogen
C C 2H 6
BH3
4
D
NH3
17
Which
reagent, when
mixedofand
heated
ammonium
ammonia?
8
In an experiment,
a sample
a pure
gaswith
is put
into a gas sulphate,
syringe atliberates
a temperature
of 300 K and
pressure of 16 kPa. The gas is compressed until the volume occupied by the gas is halved.
A COLLEGE
aqueous bromine
CEDAR
COVALENT BONDS WS 1
Afterdilute
compression,
the temperature
of the gas in the syringe is 375 K and the pressure is 40 kPa.
B
hydrochloric
acid
C
A
–980 kJ mol–1
B
O
–540 kJ mol–1
360
C
+540 kJ mol–1
213
5
+980 kJ mol–1
D
10 When solid ammonium chloride dissociates at a certain temperature in a 0.500 dm3 container,
6
diagram represents the overlap of two orbitals which will form a π bond?
10 Which
ammonia and hydrogen chloride are formed.
2
NH3(g) + HCl (g)
NH4Cl (s)
A
Section A
The initial amount of ammonium chloride was 1.00 mol, and when the system had reached
molpossible
of ammonium
chloride.
therethere
was are
0.300
For equilibrium
each question
four
answers,
A, B, C and D. Choose the one you consider to
be correct.
What is the numerical value of Kc for this reaction under these conditions?
UseBof the Data Booklet may be appropriate for some questions.
A 0.490
B 1.63
C 1.96
D
3.27
1 Which
Which stage
equation
shows
the reaction
that occurs
during
the standard
change of
11
in the
free radical
substitution
of ethane
by chlorine
has theenthalpy
lowest activation
atomisation
of
bromine?
energy?
Br2(l)→ →
C
2Cl2Br(g)
●
AA Cl
2
Br●2(g)
→H2Br(g)
+ C
BB Cl
2 6 → C2H5● + HCl
1
Br●2(l)+ →
CC C22H
Cl 2Br(g)
→ C2H5Cl + Cl ●
5
D
1
→5● Br(g)
2(g)
DD Cl2 ●Br+
C 2H
→ C2H5Cl
[W’15 1 Q4]
12
sulfur react
together
to form
sodium
sulfide,
Cl 6 molecule?
2 Sodium
What is and
the correct
number
of bonds
of each
type
in theNa
Al22S.
11
How do the atomic radius and ionic radius of sodium compare with those of sulfur?
co-ordinate
covalent
(dative covalent)
atomic radius
ionic radius
A
6
1
A
sodium < sulfur
sodium > sulfur
B
6
2
B
sodium < sulfur
sodium < sulfur
C
7
0
C
sodium > sulfur
sodium > sulfur
D
7
1
D
sodium > sulfur
sodium <9701/13/O/N/12
sulfur
© UCLES 2012
[Turn over
3 Tetraethyl lead, Pb(C2H5)4, has been used as a petrol additive.
13 Solid aluminium chloride sublimes at 178 o C.
12
What is the percentage by mass of carbon in tetraethyl lead?
Which structure best represents the species in the vapour at this temperature?
A
10.2 A
B
Cl
Cl
Cl
Al
Cl
C
B
14.9
Cl
Al
Cl
Cl
Al
Cl
Cl
29.7
D
Cl
Al
Cl
Cl
Cl
32.0
C
D
Cl
Al
Al 3+(Cl –)3
Cl
[S’06 Q15]
© UCLES 2016
9701/12/M/J/16
CEDAR COLLEGE
© UCLES 2016
[Turn over
COVALENT BONDS WS 1
9701/11/M/J/16
C
sodium
D
potassium
214
613 Carbon and silicon have the same outer electronic structure.
Why is a Si–Si bond weaker than a C–C bond?
A
Silicon atoms have a larger atomic radius than carbon atoms.
B
Silicon has a greater nuclear charge than carbon.
C
Silicon has a smaller first ionisation energy than carbon.
D
Silicon is more metallic than carbon.
[S’16 2 6]
© UCLES 2016
CEDAR COLLEGE
9701/13/M/J/16
[Turn over
COVALENT BONDS WS 1
Decide whether each of the statements is or is not
Section
B (you may find it helpful to put a tick against
12correct
the statements that you consider to be correct).
The each
responses
to D shouldinbe
selected
onone
the basis
For
of theAquestions
this
section,
or more
14of of the three numbered statements 1 to 3 may
215
The
responses
A
to
D
should
be
selected
on
the
basis
of
be correct.
Section B C
A
B SECTION
D
B
Decide whether
is or is not correct
A each of the statements
B
C (you may find itDhelpful to put a tick against
1, 2questions
and
arein this
1to
and
only
are
2 and
only
arenumbered
1 only isstatements
correct 1 to 3 may
For each of the
section,
one
or more
of 3the
three
the statements
that
you 3consider
be2correct).
correct
correct
correct
1, 2 and 3
1 and 2
2 and 3
1 only
be correct.
are
only
are
only
are
is
The responses A to D should be selected on the basis of
Decide
whether
of the statements
or aiscorrect
not correct
(you may find itcorrect
helpful to put a tick against
correct each
correct
correct
No other
combination
of statements
is usedisas
response.
the statements that you consider to be correct).
A
B
C
D
No otherresponses
combination
is used as
athe
correct
A toofDstatements
should be diamond
selected
on
basis response.
of
33 The
The conversion
is an
endothermic
kJ mol–1).
1, 2 and 3of graphite into
1 and 2
2 and 3reaction (∆H = +3
1 only
are
only are
only are
→ C(diamond)
A
B
C
31 Which diagrams
ofC(graphite)
a giant molecular
structure?
correct represent part
correct
correct
is
D
correct
2 and 3
1, 2 and 3are correct? 1 and
Which 1statements
2 2
3 1 only
is
onlyresponse.
are
only
are of statements
No other combination
is are
used as a correct
=
C
correct
correct
correct
correct
= change
C
1
The enthalpy
of atomisation of diamond is smaller than that of graphite.
2
The bond energy of the C–C bonds in graphite is greater than that in diamond.
= Na
= Cl
Molecules of P attract each other by
31 No
P and
are two compounds
similar
Mr values.
otherQ
combination
of statements with
is used
as a correct
response.
hydrogen
bonds. Molecules
of Q attractofeach
otherisby
van der
Waals’
only.
3
The enthalpy
change of combustion
diamond
greater
than
that offorces
graphite.
31
The diagram
illustrates
the energy
a set
of reactions.
A sample
of liquid
P is compared
to changes
a sampleofof
liquid
Q.
341 Which of the following statements are correct for the sequence of compounds below considered
from left
right?
How
will to
their
properties differ? H = _134 kJ mol _ 1
R
S
NaF
MgO in water
Al N thanSiC
1 P
is more soluble
Q.
32 2Which
reactions
are melting
redoxdifference
reactions?
has
a higher
point than
Q. the elements in each compound increases.
1 P
The
electronegativity
between
312
32
The
formula-units
of these
are isoelectronic (have the same number of
P
is more
Q.4 + Brcompounds
CaSO
CaBr
2 + 2Hviscous
2SO4 → than
2 + SO2 + 2H2O
1
electrons).
H = +92 J mol
2 CaBr2 + 2H3PO4 → Ca(H2PO4)2 + 2HBr
3
The bonding becomes increasingly covalent.
The reaction
[W’03 1 Q34]
3 CaBr2 + 2AgNO3 → Ca(NO3)2 + 2AgBr
_75
G used
+ Has a food preservative?
E kJ
+ dioxide
F _ 1 is
35 Which statements are reasons why
H =sulphur
mol
U
332 Sodium hydrogensulfide, T
NaSH, is used to remove hair from animal hides.
is
by platinum.
1 catalysed
It is a reducing
agent and therefore an anti-oxidant.
Which
of the following
statements
are correct?
correct?
Which
statements
about the
SH– ion are
2
It prevents
alcohols
sour-tasting
acids.
Which
statements
aboutforming
the properties
of the
catalyst are correct? –1
1
The
enthalpy
change
for
the
transformation
U ⎯→ R is + 42 kJ mol .
1
It
contains
18
electrons.
3
It does not smell and therefore can be used in more than trace quantities.
1 2TheThe
catalyst
haschange
no effect
enthalpy change
ofisthe
reaction.
enthalpy
for on
the the
transformation
T ⎯→ S
endothermic.
2 Three lone pairs of electrons surround the sulfur atom.
2 3TheThe
catalyst
increases
ratetransformation
of the reverse
reaction.
enthalpy
change the
for the
R ⎯→
T is – 33 kJ mol–1 .
36 3
WhySulfur
is thehas
addition
of concentrated
sulphuric acid to solid potassium iodide unsuitable for the
an oxidation
state of +2.
[S’10 3 Q32]
of hydrogen
iodide?
3preparation
The catalyst
increases
the average kinetic energy of the reacting particles.
1
Hydrogen iodide is not displaced by sulphuric acid.
333 Which
elements
have
atomstowhich
2
Iodide
ions are
oxidised
iodine.can form π bonds with atoms of other elements?
13
oxygen
The product is contaminated by sulphur compounds.
© UCLES 2010
2
nitrogen
3
fluorine
9701/11/M/J/10
9701/1/M/J/03
© UCLES 2004
© UCLES 2014
CEDAR COLLEGE
9701/01/O/N/04
9701/12/M/J/14
COVALENT BONDS WS 1
2
2+
(aq) + H2(g)
X(s) is+ 10.83
2H+(aq)
The Ar of boron in the sample
to →
fourX significant
figures.
3
YZ(s)
+ 2Hin+(aq)
→ Y2+(aq) + H2Z(g)
X+. 216
There are more protons in
Y+ than
When 1.0 g of either X or YZ is reacted with an excess of acid, the total volume of gas formed is
4 the
same.
32
Which
elements can form π bonds in their compounds?
1 carbon
Which
statements are correct?
12
oxygen
A
r(X) = Mr(YZ)
23
nitrogen
X
and Y are metals.
[M’16 2 Q31]
3 X and Y must both be in the same Group of the Periodic Table.
33 For which enthalpy changes is the value of ∆H always negative?
1 thecombustion
355 In
gas phase, aluminium chloride exists as the dimer, Al2Cl6.
2 using
hydration
By
this information, which of the following are structural features of the Al2Cl6 molecule?
3
1
solution
Each aluminium atom is surrounded by four chlorine atoms.
2
There are twelve non-bonded electron pairs in the molecule.
3
Each aluminium atom contributes electrons to four covalent bonds.
© UCLES 2016
9701/11/O/N/16
© UCLES 2012
9701/13/O/N/12
CEDAR COLLEGE
COVALENT BONDS WS 1
Aluminium reacts with chlorine.
4
217
217
(b)
how,hydrocarbon
starting from
aluminium
powder,
reaction
be carried out For
2 (i)
The Outline
unsaturated
ethyne
(acetylene),
C2H2, this
is widely
usedcould
in ‘oxy-acetylene
torches’
cuttingor
and
weldinglaboratory
metals. In the
torch,a ethyne
burnedof
in aluminium
oxygen to produce
a Examiner’s
in aforschool
college
to give
small is
sample
chloride.
A
Use
flamediagram
with a temperature
of
3400
K.
COVALENT BONDS WS 2
is not necessary.
COVALENT BONDING WS 2
(a)1 Ethyne
is a linear molecule with a triple bond, C!C, between the two carbon atoms.
..................................................................................................................................
Draw a ‘dot-and-cross’ diagram of an ethyne molecule.
..................................................................................................................................
..................................................................................................................................
(ii)
Describe what you would see during this reaction.
[1]
(b) When
used for cutting or welding, ethyne is transported in cylinders which contain the
..................................................................................................................................
gas under pressure. A typical cylinder has a volume of 76 dm3 and contains ethyne gas
at
1515 kPa pressure at a temperature of 25 °C.
..................................................................................................................................
Use the general gas equation, pV = nRT, to calculate the amount, in moles, of ethyne in
(iii)2 this
At low
temperatures, aluminium chloride vapour has the formula Al2Cl6.
cylinder.
Draw a ‘dot-and-cross’ diagram to show the bonding in Al2Cl6.
Show outer electrons only.
Represent the aluminium electrons by .
Represent the chlorine electrons by x.
[2]
(c) In some countries, ethyne is manufactured from calcium carbide, CaC2, which is
produced by heating quicklime and coke together at 2300 K.
CaO + 3C
CaC2 + CO
When water is added to the CaC2, calcium hydroxide, Ca(OH)2, and ethyne, C2H2, are
produced.
(i)
Construct a balanced equation for the formation of ethyne from calcium carbide.
..................................................................................................................................
(ii)
[6]
Use this equation and your answer to part (b) to calculate the mass of CaC2 which
will react with an excess of water to produce enough ethyne to fill 100 cylinders of
the gas.
© UCLES 2009
9701/21/M/J/09
[3]
© UCLES 2006
BILAL HAMEED
CEDAR COLLEGE
9701/02/M/J/06
COVALENT BONDING WS 2
COVALENT BONDS WS 2
218
NOTES
CEDAR COLLEGE
COVALENT BONDS WS 2
219
3 Chemical bonding
This topic introduces the different ways by which chemical bonding occurs and the
effect this can have on physical properties.
3.2 shapes of simple molecules
SHAPES OF MOLECULES
BILAL HAMEED
SHAPES OF MOLECULES
Cambridge International AS and A Level Chemistry 9701 syllabus Syllabus content
3 Chemical bonding
220
This topic introduces the different ways by which chemical bonding occurs and the effect this can have
3
bonding
onChemical
physical properties.
This topic introduces the different ways by which chemical bonding occurs and the effect this can have
on physical properties.
Learning outcomes
Candidates should be able to:
3.1 Ionic bonding
Learning outcomes
Candidates
be able to:
a) describeshould
ionic bonding,
as in sodium chloride, magnesium oxide and
calcium fluoride, including the use of ‘dot-and-cross’ diagrams
3.1 Ionic bonding
3.2 Covalent bonding
and co-ordinate
(dative covalent)
3.2 Covalent
bonding
bonding
including
and
co-ordinate
shapes covalent)
of simple
(dative
molecules
bonding
including
shapes of simple
molecules
a) describe ionic bonding, as in sodium chloride, magnesium oxide and
including
use of ‘dot-and-cross’
diagrams
a) calcium
describe,fluoride,
including
the usethe
of ‘dot-and-cross’
diagrams:
(i) covalent bonding, in molecules such as hydrogen, oxygen, chlorine,
a) describe,
including
the carbon
use of ‘dot-and-cross’
diagrams:
hydrogen
chloride,
dioxide, methane,
ethene
(i)
bonding,
molecules
such assuch
hydrogen,
oxygen,
chlorine,
(ii) covalent
co-ordinate
(dativeincovalent)
bonding,
as in the
formation
of the
hydrogen
chloride,
dioxide,
methane,
ethene
ammonium
ion andcarbon
in the Al
Cl
molecule
2
6
co-ordinate
(dative
covalent)
bonding,
such
as in the
formation
of the
b) (ii)
describe
covalent
bonding
in terms
of orbital
overlap,
giving
σ and π
2
3
ammonium
in the of
Alhybridisation
Cl
molecule
bonds,
includingion
theand
concept
to
form
sp,
sp
and
sp
2
6
orbitals (see
also Section
b) describe
covalent
bonding14.3)
in terms of orbital overlap, giving σ and π
theof,
concept
of hybridisation
to form sp,
sp2 and
c) bonds,
explain including
the shapes
and bond
angles in, molecules
by using
thesp3
orbitals
(see
also Section
14.3) repulsion (including lone pairs), using
qualitative
model
of electron-pair
as simple
CO2 in,
(linear),
CH4 (tetrahedral),
3 (trigonal),
c) explain
theexamples:
shapes of,BF
and
bond angles
molecules
by using theNH3
SF6 (octahedral),
PF5 (trigonal
bipyramidal)
(pyramidal),model
H2O (non-linear),
qualitative
of electron-pair
repulsion (including
lone pairs),
using
simple
CO2 in,
(linear),
CH4 (tetrahedral),
NH3
d) as
predict
theexamples:
shapes of,BF
and
bond angles
molecules
and ions analogous
3 (trigonal),
SF6also
(octahedral),
PF5 (trigonal bipyramidal)
(pyramidal),
H2O (non-linear),
to those specified
in 3.2(b) (see
Section 14.3)
3.3 Intermolecular
forces,
electronegativity
3.3 Intermolecular
and bond properties
forces,
electronegativity
and bond properties
d) predict the shapes of, and bond angles in, molecules and ions analogous
those specified
3.2(b) (see
also
Section and
14.3)water as simple
a) to
describe
hydrogeninbonding,
using
ammonia
examples of molecules containing N–H and O–H groups
a)
hydrogen
bonding,
ammonia
water as simple
b) describe
understand,
in simple
terms,using
the concept
of and
electronegativity
and apply
examples
ofthe
molecules
containing
N–H and
O–H
groups
it to explain
properties
of molecules
such
as bond
polarity (see
also Section in
3.3(c)),
the
dipolethe
moments
(3.3(d))and
andapply
the
b) understand,
simple
terms,
conceptof
ofmolecules
electronegativity
behaviour
of
oxides
with
water
(9.2(c))
it to explain the properties of molecules such as bond polarity (see
Section
3.3(c)),
the energy,
dipole moments
of molecules
(3.3(d)) and
c) also
explain
the terms
bond
bond length
and bond polarity
and the
use
behaviour
of oxides
water (9.2(c))
them to compare
thewith
reactivities
of covalent bonds (see also Section
5.1(b)(ii))the terms bond energy, bond length and bond polarity and use
c) explain
d) them
describe
intermolecular
forces (van
der Waals’
forces),
on
to compare
the reactivities
of covalent
bonds
(seebased
also Section
permanent and induced dipoles, as in, for example, CHCl 3(l); Br2(l) and
5.1(b)(ii))
the
liquidintermolecular
Group 18 elements
d) describe
forces (van der Waals’ forces), based on
3.4 Metallic bonding
3.4 Metallic bonding
3.5 Bonding and
physical properties
3.5 Bonding and
physical properties
permanent and induced dipoles, as in, for example, CHCl 3(l); Br2(l) and
the
liquid metallic
Group 18
elements
a) describe
bonding
in terms of a lattice of positive ions
surrounded by delocalised electrons
a) describe metallic bonding in terms of a lattice of positive ions
by delocalised
electrons
a) surrounded
describe, interpret
and predict
the effect of different types of bonding
(ionic bonding, covalent bonding, hydrogen bonding, other intermolecular
interactions,
metallic
bonding)
on the
physical
properties
ofof
substances
a) describe,
interpret
and
predict the
effect
of different
types
bonding
b) (ionic
deduce
the type
of bonding
present
from given
information
bonding,
covalent
bonding,
hydrogen
bonding,
other intermolecular
metallic of
bonding)
on reactions
the physical
properties
of substances
c) interactions,
show understanding
chemical
in terms
of energy
transfers
associated
breakingpresent
and making
chemical
bonds
b) deduce
the with
type the
of bonding
from of
given
information
c) show understanding of chemical reactions in terms of energy transfers
associated with the breaking and making of chemical bonds
Back to contents page
Back to contents page
BILAL HAMEED
www.cie.org.uk/alevel
19
www.cie.org.uk/alevel
19
SHAPES OF MOLECULES
221
SHAPES OF MOLECULES
The shape of a molecule plays a large part in determining its properties and
reactivity.
The specific orientation of electron pairs in covalent molecules imparts a
characteristic shape to the molecules.
The shape of a molecule made of only two atoms, such as H2 or CO, is easy to
determine. Only a linear shape is possible when there are two atoms. Determining
the shapes of molecules made of more than two atoms is more complicated.
Using Valence Shell Electron Pair Repulsion (VSEPR) theory one can predict the
shape of a molecule by examining the Lewis structure of the molecule.
1
VSEPR THEORY
The electron pairs around the central atom
of the molecule arrange themselves to
minimise electronic repulsion and so that
they can be as far away as possible from
each other.
Bonds are closer
together so repulsive
forces are greater
Al
Bonds are further
apart so repulsive
forces are less
This fact is used to predict molecular shape.
Molecules contain covalent bonds.
Al
As covalent bonds consist of a pair of
electrons, each bond will repel other bonds.
All bonds are equally
spaced out as far apart
as possible
2
BILAL HAMEED
SHAPES OF MOLECULES
222
VSEPR THEORY
Bonds will therefore push each other as far apart as possible to reduce
the repulsive forces.
Because the repulsions are equal, the bonds will also be equally spaced.
Denniston: General,
Organic and Biochemistry,
ourth Edition
4. Structure and Properties
of Ionic and Covalent
Compounds
Bonds© The
areMcGraw−Hill
closer together so
Companies,
2003 are greater
repulsive
forces
Text
Al
Bonds are further apart so repulsive
forces are less
103
4.4 Drawing Lewis Structures of Molecules and Polyatomic Ions
BeH2
HSBeSH
O
SFS
OSBS
O F
OS
SF
Q
Q
BF3
All bonds are equally spaced out as
far apart as possible
Al
3
H
OSH
HSC
Q
H
CH4
NH3
O
HSN
QSH
H
H2O
HSO
QSH
e electron pairs around the central atom of the molecule arrange themselves
imize electronic repulsion. This means that the electron pairs arrange themso that they can be as far as possible from each other. We may use this fact
dict molecular shape. This approach is termed the v alence shell electron p air
ion (VSEPR) theory.
t’s see how the VSEPR theory can be used to describe the bonding and
ure of each of the preceding molecules.
LINEAR
Beryllium chloride has two shared electron pairs around the beryllium
Only four electrons surround the
atom.
beryllium atom in BeH . Consequently,
saw in Example 4.14, beryllium hydride has two shared electron pairs
d the beryllium atom. These electron pairs have minimum repulsion if they
ated as far apart as possible while still bonding the hydrogen to the central
This condition is met if the electron pairs are located on opposite sides of the
ule, resulting in a linear structure, 180! apart:
2
BeH2 is a stable exception to the octet
rule.
These electron pairs have minimum repulsion if they are located as far
apart
as possible while still bonding the chlorine to the central atom.
HSBeSH
ThisH—Be—H
condition is met if the electron pairs are located on opposite sides of
180!
the molecule,
resulting in a linear structure, 180o apart:
nd angle, the angle between H—Be and Be—H bonds, formed by the two
ng pairs is 180! (Figure 4.4).
180!
BeCl2
Cl
Linear
BeH
2
o
angle: 180
shape: linear
trifluoride has three shared electron pairs around the central atom. Placing
ctron pairs in a plane, forming a triangle, minimizes the electron pair repulthis molecule, as depicted in Figure 4.5 and the following sketches:
S
S
S
S
S
S
S
S
F F
B
O
SQ
FS
or
F
120°
H E
B
F
|
F
BILAL HAMEED
120°
or
F
H E
B
CO2
180°
Figure 4.4
Be
Cl
Bonding and geometry in beryllium O
180°
C
O
hydride, BeH2.
BF3 has4only six electrons around the
central atom, B. It is one of a number of
stable compounds that are exceptions to
the octet rule.
F
|
F
4-27
SHAPES OF MOLECULES
223
TRIGONAL PLANAR
Denniston: General,
Organic and Biochemistry,
Fourth Edition
4. Structure and Properties
Text
© The McGraw−Hill
2003
Ionic and Covalent
NowofCompounds
think
about what happens when the Companies,
central
atom is surrounded by
three shared pairs. Look at BF3.
Boron trifluoride has three shared electron pairs around the central atom.
104
Chapter 4 Structure and Properties of Ionic and Covalent Compounds
Placing the
electron pairs in a plane, forming a triangle, minimizes the
Such a structure is trigonal planar, and each F—B—F bond angle is 120!. We also
find that compounds with central atoms in the same group of the periodic table
electron pair
repulsion
in this molecule.
have similar
geometry. Aluminum,
in the same group as boron, produces compounds such as AlH3, which is also trigonal planar.
F
120°
F
B
120°
120!
F
angle: 120o
120°
shape: trigonal planar
Planar
Figure 4.5
5
Representation of the two-dimensional
structure of boron trifluoride, BF3.
(a) Trigonal planar structure. (b) Ball and
stick model of trigonal planar BF3.
CH4, NH3, and H2O all have eight
electrons around their central atoms; all
obey the octet rule.
(a)
(b)
BF3
CH4
Methane has four shared pairs of electrons. Here, minimum electron repulsion is
achieved by arranging the electrons at the corners of a tetrahedron (Figure 4.6).
Each H—C—H bond angle is 109.5!. Methane has a three-dimensional tetrahedral
structure. Silicon, in the same group as carbon, forms compounds such as SiCl4
and SiH4 that also have tetrahedral structures.
Projecting away
from you, behind the
plane of the paper H
H
H
TETRAHEDRAL
In the plane of
the paper
Methane
has four shared pairs of electrons.
109.5°
C
H
1.09Å C
H
(a)
Figure 4.6
H
H
C
Here, minimum electron
repulsion is achievedHby arranging the electrons
H
H
at theH corners
of a tetrahedron.
Projecting toward
H
you, in front of the
plane of the paper
Each H—C—H (b)
bond angle is 109.5º.
(c)
Methane NH
has a three҆dimensional tetrahedral structure.
Representations of the three-dimensional
structure of methane, CH4. (a) Tetrahedral
methane structure. (b) Ball and stick
model of tetrahedral methane. (c) Threedimensional representation of
structure (b).
3
Ammonia also has four electron pairs about the central atom. In contrast to
H
methane, in which all four pairs are bonding, ammonia
has three pairs of bonding
electrons and one nonbonding lone pair of electrons. We might expect CH4 and
NH3Hto have electron pair arrangements
109.5° that are similar but not identical. The lone
pairxoin ammonia is more negative than the bonding pairs; some of the negative
C
charge
of the hydrogen atoms with
x on the bonding pairs is offset by the presence
Hxo Cpositive
H in ammonia is disoH
their
nuclei. Thus the arrangement of electron pairs
xo
torted. The resulting distribution appears
H as
H
4-28
109.5°
H
angle: 109.5o
shape: tetrahedral
6
BILAL HAMEED
SHAPES OF MOLECULES
224
5 AND 6 BONDS
F
shape: trigonal bipyramidal
90°
F
P
120°
angle: 90o and 120o
F
F
F
F
F
F
S
F
angle: 90o
F
shape: octahedral
F
7
IRREGULAR SHAPES
NH3 also has four electron pairs about the central atom. In contrast to CH4, in
which all four pairs are bonding, ammonia has three pairs of bonding
electrons and one nonbonding lone pair of electrons.
H
oo
xo
H xo Nxo xo H
H xo Cxo xo H
H
H
The lone pair in ammonia is closer to the central atom, N, than the bonding
pairs and has greater repulsion.
Thus the arrangement of electron pairs in ammonia is distorted.
8
BILAL HAMEED
SHAPES OF MOLECULES
has four pairs
of electrons
surrounding
ammonia
and water.
Space-fithe
lling models of these
m. Note that
in
drawing
three-dimensional
molecules are shown in Figure 4.17. Each of these
he triangular
‘wedge’has
is the
molecules
fourbond
pairscoming
of electrons surrounding the225
u and thecentral
dashedatom.
black line
the bond
going three-dimensional
Noteis that
in drawing
you.
diagrams, the triangular ‘wedge’ is the bond coming
towards you and the dashed black line is the bond going
IRREGULAR SHAPES
away from you.
H
C
The H atoms in NH3 are pushed closer together than in CH4. The bond
H
angle is 107o because lone pair–bond pair repulsions are greater than
repulsions
109.5º
aH Methane bond
H equal pair repulsions.
Cpair–bond
H H
H
a
Check-up
N
H
O
H
H
H
H
C
109.5º
H
C
H
H lone pair
repulsions
equal
H
greatest repulsion between lone pairs
5H The table lists bond lengths and bond
greater
halides.
a energies
Ammonia of Nsome hydrogen
repulsion
H
107º
H
O
lone pair
Hydrogen H Bond
Bond energy /
halideH
kJ mol−1
N length
H / nm
greater
N431 repulsion
H Cl
0.127
H
H
0.141 intermediate
366
H Br
107º
H
greatest
H 299
HH I
O 0.161
H repulsion
repulsion
104.5º
What is the relationship
between the
a aWater
H
H
intermediate repulsion
H
H
O
H angle is 104.5°
least repulsion between bonding pairs
9
Figure
4.15 Repulsion between lone and bonding electron pairs in water.
Working out the shapes of molecules
The differences in electron-pair repulsion determine
bond length and the bond energy for
the shape and bond angles in a molecule. Figure 4.16
intermediate
these hydrogen
halides?
he bond angles in a methane,
b ammonia
and c water
compares the shapes and bond angles of methane,
greatest
H
O the Hbond energy values O
H repulsion
ttype of electron-pair
b Suggest
why
repulsion.
repulsion
ammonia and water. Space-filling models of these
decrease in the order HCl > HBr > HI. 104.5º
molecules are shown in Figure 4.17. Each of these
c Suggest a value for the bond length in H
molecules
has four pairs of electrons surrounding the
IRREGULAR
SHAPES
hydrogen fluoride, HF.
central atom. Note that in drawing three-dimensional
Figure 4.16 The bond angles in a methane, b ammonia and c water
the triangular
‘wedge’
the bond coming
If a molecule,
or ion,
has lone pairs on the diagrams,
central atom,
the shapes
areisslightly
depend on the type
of electron-pair
repulsion.
distorted away from the regular shapes. towards you and the dashed black line is the bond going
away from you.
This is because of the extra repulsion caused by the lone pairs.
4.4 Shapes of molecules
Electron-pair repulsion theory
BOND PAIR ҆ BOND PAIR < LONE PAIRa҆Methane
BOND PAIR < LONE PAIR ҆ LONE PAIR
Because all electrons have the same (negative)
charge,
H
H
they repel each other when they are close together. So,
a pair of electrons in the bonds surrounding the central
109.5º
H
H
C
H
C
atom in a molecule will repel other electron
pairs. This
O
O
O
H
repulsion forces the pairs of electrons apart until the
H
H
repulsive forces are minimised.
repulsions
equal
As aangles
resultofofa the
extra repulsion,
tend to be slightly less as the bonds
The shape and bond
covalently
bonded bond angles
a Ammonia
molecule dependare
on:squeezed together.
lone pair
• the number of pairs of electrons around each atom
H
N
H
10
• whether these pairs are lone pairs or bonding pairs.
greater
N
Lone pairs of electrons have a more concentrated
electron charge cloud than bonding pairs of electrons.
Their cloud charges are wider and slightly closer to the
nucleus of the central atom. This results in a different
amount of repulsion between different types of electron
BILAL HAMEED
pairs. The order of repulsion is lone pair–lone pair (most
repulsion) > lone pair–bond pair > bond pair–bond pair
(least repulsion).
H
H
repulsion
H
107º
H
a Water
intermediate
H
O
H
greatest OF MOLECULES
SHAPES
repulsion
repulsion
H
O
104.5º
H
226
AMMONIA
N
• Nitrogen has five electrons in its outer shell
• 3 covalent bonds are formed and a pair of non-
H
H
H
bonded electrons is left
• As the total number of electron pairs is 4, the
N
H
shape is BASED on four bond tetrahedral shape
• Not all the repulsions are the same.
Repulsions: LONE PAIR - BOND PAIR > BOND PAIR - BOND PAIR
• The N-H bonds are pushed closer together
• Lone pairs are not included in the shape
H
angle: 107o
H
shape: trigonal pyramidal
N
H
107°
H
H
11
WATER
• Oxygen has six electrons in its outer shell
• 2 covalent bonds are formed and 2 pairs of non—
O
H
H
bonded electrons are left
• As the total number of electron pairs is 4, the shape
O
is BASED on the four bond tetrahedral shape.
• Not all the repulsions are the same.
Repulsions: LONE PAIR — LONE PAIR > LONE PAIR — BOND PAIR >
H
angle: 104.5o
H
shape: angular / bent
BOND PAIR — BOND PAIR
• The O—H bonds are pushed even closer together
• Lone pairs are not included in the shape
O
H
104.5°
H
12
BILAL HAMEED
SHAPES OF MOLECULES
227
SUMMARY OF SHAPES
bonds
lone pairs
shape
angle
example
2
0
linear
180o
BeCl2 CO2
3
0
trigonal planar
120o
BF3 AlCl3
2
1
bent / angular
117o
SO2
4
0
tetrahedral
109.5o
SiCl4 CH4
3
1
trigonal pyramidal
107o
NH3 PCl3
2
2
bent / angular
104.5o
H2O
5
0
trigonal bipyramidal
90o & 120o
PCl5
6
0
octahedral
90o
SF6
13
CALCULATING SHAPES
The shape of a molecule or a complex ion is calculated by:
1. Calculating the number of electrons in the outer shell of the central
species
2. Pairing up electrons, making sure the outer shell maximum is not
exceeded
3. Calculating the number of bond pairs and lone pairs
4. Using VSEPR to calculate shape and bond angle(s)
14
BILAL HAMEED
SHAPES OF MOLECULES
228
CALCULATING SHAPES
For IONS:
The number of electrons in the outer shell depends on the charge on the ion
•
if the ion is positive you remove as many electrons as there are positive
charges
•
if the ion is negative you add as many electrons as there are negative
charges
e.g.
for PF6҆
add one electron to the outer shell of P
for PCl4+
remove one electron from the outer shell of P
15
EXAMPLES
NH3
NH4+
NH2҆
BOND PAIRS
3 PYRAMIDAL
LONE PAIRS
1 H҆N҆H 107°
A compound with an overall charge may
not be ionic, it can also be covalent.
BOND PAIRS
4 TETRAHEDRAL
e.g: NH4+
LONE PAIRS
0 H҆N҆H 109.5°
BOND PAIRS
2 ANGULAR
LONE PAIRS
2 H҆N҆H 104.5°
Number of valence electrons = 5
But due to “+” it loses 1e҆ ∴ = 4
+
҆
16
BILAL HAMEED
SHAPES OF MOLECULES
229
SKILL CHECK 1
Determine the number, and type, of electron pairs around the central
atom(s) in each of the following. Predict the shape and bond angles of
each. (Hint: it may help to draw ‘dot-and-cross’ diagrams)
(a) phosphine, PH3
(b) sulfur dichloride, SCl2
(c) dichloromethane, CH2Cl2
(d) cobalt(II) chloride, COCl2
(e) xenon tetrafluoride, XeF4
17
MORE
BF3
3 bp 0 lp
120º
trigonal planar
boron pairs up all 3 electrons in its
outer shell
SiCl4
4 bp 0 lp
109.5º
tetrahedral
silicon pairs up all 4 electrons in its
outer shell
PCl4+
4 bp 0 lp
109.5º
tetrahedral
as ion is +, remove an electron in the
outer shell then pair up
PCl6҆
6 bp 0 lp
90º
octahedral
as the ion is ҆ , add one electron to
the 5 in the outer shell then pair up
SiCl62҆
6 bp 0 lp
90º
octahedral
as the ion is 2҆, add two electrons
to the outer shell then pair up
18
BILAL HAMEED
SHAPES OF MOLECULES
230
SKILL CHECK 2
Draw the shapes for the following molecules. The underlined atom is
the central atom
a. NF3
b. CH3OH
c. BeF2
d. CCl2F2
e. HOCl
19
SKILL CHECK 3
Draw the shapes of the following molecules
a. CF4
b. Cl2O
c. PF3
d. NCl3
e. SiCl4
f. SO3
20
BILAL HAMEED
SHAPES OF MOLECULES
231
SKILL CHECK 4
Deduce the shapes of the following ions:
a. NH4+
b. CH3+ c. CH3҆ d. CO32҆
e. SO42҆ f. AlH4҆
g. AlF63҆
21
SKILL CHECK 5
Lycra® is a polyurethane fibre used in the fashion industry. It is a polymer
made from two monomers, one of which has the following formula.
O=C=N—(CH2)n—N=C=O
What is the O-C-N bond angle in this molecule?
A 90°
B 109°
C 120°
D 180°
22
BILAL HAMEED
SHAPES OF MOLECULES
232
SKILL CHECK 6
Organic nitrates in photochemical smog can cause breathing difficulties.
The diagram shows an example of an organic nitrate molecule.
What is the correct order of the bond angles shown in ascending order
(smallest first)?
A1
2
3
B 2
1
3
C 3
1
2
D 3
2
1
23
HYBRIDISATION
Hybridisation is the mixing of atomic orbitals to produce a new set of orbitals (the
same number as originally) that have characteristics of the original orbitals and
are better arranged spatially for covalent bonding.
To form a covalent bond, an orbital containing one electron is required. These
orbitals overlap to form a covalent bond.
Carbon has six electrons and and the outer shell electronic configuration 2s2 2p2:
2p
2s
As it has only two unpaired electrons, carbon should form two covalent bonds.
24
BILAL HAMEED
SHAPES OF MOLECULES
233
HYBRIDISATION
However, it is well known that carbon virtually always forms four covalent bonds.
One of the electrons in the 2s orbital must then be promoted to the 2p sub-shell
to give four unpaired electrons. This requires energy.
2p
2s
However, bond formation releases energy and the formation of four bonds instead of two
more than pays back the energy needed to promote an electron to a higher sub-shell.
Carbon now has four unpaired electrons and can form four covalent bonds, but the
atomic orbitals do not point in the correct direction for bonding.
25
HYBRIDISATION
CH4 is tetrahedral with bond angles of 109.5° but the p orbitals are at 90° to each other.
=
=A
&%.•
8
=
=
=
;^\jgZ(#*. I]ZWdcYVc\aZhd[8=)#
=
=
8
26
=
This requires energy. However, bond forma
the formation of four bonds instead of two mor
energy needed to promote an electron to a high
Carbon now has four unpaired electrons and
bonds, but the atomic orbitals do not point in th
bonding. CH4 is tetrahedral with bond angles of
the p orbitals are at 90° to each other (Figure 3.
When carbon forms methane, the four atomi
s and three p, then mix to give four sp3 hybri
to the vertices of a tetrahedron. This is the proce
(Figure 3.60).
=
;^\jgZ(#+% I]ZedgW^iVah^c8=)#
BILAL HAMEED
Hybridisation produces orbitals that
V
=
SHAPES OF MOLECULES
=
=
(Figure 3.60).
V
234
=
W
SP 3 HYBRIDISATION
8
When=carbon forms methane, the four atomic orbitals on carbon, one s and three p,
then mix to give four sp3 hybrid
= orbitals, which point to the vertices of a tetrahedron.
This is the=process of hybridisation.
;^\jgZ(#+& V;djghe(]nWg^YdgW^iVahed^ciidlVgYhi]ZkZgi^XZhd[ViZigV]ZYgdc
VcYVgZWZiiZghZije[dgWdcY^c\idi]Z=Vidbh0WdcZhe(dgW^iVa#
The four sp3 hybrid
orbitals all have the same energy (they are
The four sp3 hybrid orbitals all have the same energy
degenerate):
(
he
8'=)
27
The Lewis structure for ethene is shown in Figure 3.62.
The shape about each C atom is trigonal planar, and the molecule is
planar overall:
=
8
=
8
=
=
Of the three p orbitals on each C atom, one of them is not in the same
plane as the H atoms or the other
C 3atom.
This p orbital is not involved in
SP
HYBRIDISATION
hybridisation. Mixing the two p orbitals and one s orbital all in the
3 hybridised atom, then the
Should
only three
bondpointing
to the sp
towards
the corners of
same plane
produces
threeatoms
sp2 orbitals
an equilateral
triangle:
molecular
geometry is trigonal pyramidal. Ammonia is an example of this.
he']nWg^Y
dgW^iVa
=
=
8
8
e
=
he'
=
In water the oxygen is sp3 hybridised but because only two atoms have
bonded to the oxygen, the molecular geometry is bent.
28
BILAL HAMEED
SHAPES OF MOLECULES
Carbon now has four unpaired electrons and can form four covalent
bonds, but the atomic orbitals do not point in the correct direction for
bonding. CH4 is tetrahedral with bond angles of 109.5° (Figure 3.59), but
235
the p orbitals are at 90° to each other (Figure 3.60).
When carbon forms methane, the four atomic orbitals on carbon, one
s and three p, then mix to give four sp3 hybrid orbitals, which point
to the vertices of a tetrahedron. This is the process of hybridisation
(Figure 3.60).
SP 2 HYBRIDISATION
In an sp2 hybridisation one 2s and
W two 2p orbitals combine to form a new hybrid shape.
V
=
In this hybridised state the carbon will make two single bonds and one double bond.
Of the three p orbitals on each C atom, one of them is not in the same plane as the H
atoms
8 or the other C atom.
=
This p orbital is not involved in hybridisation.
=
=
;^\jgZ(#+& V;djghe(]nWg^YdgW^iVahed^ciidlVgYhi]ZkZgi^XZhd[ViZigV]ZYgdc
VcYVgZWZiiZghZije[dgWdcY^c\idi]Z=Vidbh0WdcZhe(dgW^iVa#
The four sp3 hybrid orbitals all have the same energy (they are
degenerate):
(
he
29
8'=)
The Lewis structure for ethene is shown in Figure 3.62.
The shape about each C atom is trigonal planar, and the molecule is
planar overall:
=
8
=
8
=
=
SP 2is not
HYBRIDISATION
Of the three p orbitals on each C atom, one of them
in the same
plane as the H atoms
or thethe
other
C atom.
This pand
orbital
involved
in
Mixing
two
p orbitals
oneis snot
orbital
all in
hybridisation. Mixing the two p orbitals and one s orbital
all
in
the
the same plane produces three sp2 orbitals
same plane produces three sp2 orbitals pointing towards the corners of
pointing towards the corners of an equilateral
an equilateral triangle:
triangle:
he']nWg^Y
dgW^iVa
=
=
8
8
e
=
he'
=
This leaves one p orbital, containing one
electron, perpendicular to the single bond ()
framework on each C atom:
30
BILAL HAMEED
SHAPES OF MOLECULES
236
HYBRIDISATION
sp3 hybridisation
sp2 hybridisation
sp hybridisation
Number of Atoms
Bonded to the Central
Atom
4
3
2
Angle between Atoms
Bonded to Central
Atom
109.5°
120°
180°
Molecular Geometry
Tetrahedral with four atoms
bonded. Trigonal pyramidal with
three atoms bonded. Bent with
two atoms bonded.
Trigonal planar with three
atoms bonded.
Linear with two atoms
bonded.
Types of Bonds Found
Four single bonds.
One double bond and two
single bonds.
One single and one triple
bond. (Or) Two double
bonds.
31
BILAL HAMEED
SHAPES OF MOLECULES
D
CH3 CH2CH2CH2CH2OH
237
Which of the following molecules has no permanent dipole?
5
A
OF
1
B SHAPES
CHCl3
C MOLECULES
C2Cl4
D CWS
2H5 Cl
CCl2F2
SECTION A
6 1 Which is the most likely shape of a molecule of hydrazine, N2H4 ?
3
4
A
B
5
H
H
N
N
H
H
Unnilpentium90°
is an artificial element. One of its isotopes is 262
105Unp.
Which of the following statements is correct?
A
262Unp has a nucleon number of 105.
105
B
262
The atom 260
105X is an isotope of 105Unp.
C
There are 262 90°
neutrons in 262
105Unp.
D
The proton number of 262
105Unp is 262.
H
H
H
N
N
H
The table gives the successive ionisation energies for an element X.
H
120°
C
N
H
N
H
H
ionisation energy / kJ mol–1
1st
2nd
3rd
4th
5th
6th
950
1800
2700
4800
6000
12300
What could be the formula of the chloride of X ?
D
A
XCl
N
H
6
7
B
107° H
XCl2
N
H C
D
XCl3
XCl4
H
Which molecule contains only six bonding electrons?
A
B
C2H4
C2F6
C
D
H2O
[W’02 Q6]
NF3
Which of the following exists in the solid state as a giant covalent lattice?
27 A Chemists
ice
have been interested in the properties of hydrogen selenide, H2Se, to compare it with
‘bad egg’ gas hydrogen sulphide, H2S.
B
iodine
set of data would the hydrogen selenide molecule be expected to have?
IV) oxide
C Which
silicon(
D
tin(number
IV) chloride
of lone pairs
bond angle
on Se atom
9701/1/O/N/02
A
1
104 °
B
2
104 °
C
2
109 °
D
2
180 °
[Turn over
[W’03 Q7]
CEDAR COLLEGE
SHAPES OF MOLECULES WS 1
9701/1/O/N/03
[Turn over
C
is insoluble in water
What can be deduced about element X from the graph?
D is
used as
a food-preservative
The
gecko,
a small
lizard, can climb up a smooth glass window. The gecko has millions of
A It is in thehairs
second
period
to Ne)
of hair
the238
Periodic
Table. of pads at its tip. The result is that the
microscopic
on its
toes(Li
and
each
has thousands
molecules
in
the
pads
are
extremely
close
to
the
glass
surface on which the gecko is climbing.
B It is a d-block element.
203 Which molecule is planar?
What
attraction
between
theTable.
gecko’s toe pads and the glass surface?
C It is the
in Group
II of the
Periodic
A NF3
D It
is in Groupbonds
III of the Periodic Table.
A
co-ordinate
B C2Cl 4
B covalent bonds
C C3H6 bonding can occur between molecules of methanal, HCHO, and molecules of liquid Y.
5 Hydrogen
C ionic bonds
D C 3H 8
What could
liquid Y be?
[S’04 Q20]
D van der Waals’ forces
6
A
CH OH
3
3
21 Which of these always applies to a nucleophile?
B CH
3CHO
7 4 What
are
the bond angles in the PH3 molecule likely to be?
5 Use
ofattacks
the Data
Bookletbond.
is relevant to this question.
A It
a double
C CH3COCH3
A 90 o
B 104 o
C 109 o
D 120 o
B
It
has
a
lone
pair
of
electrons.
Nickel
up 20 % of the total mass of a coin. The coin has a mass of 10.0 g.
D CHmakes
3CO2CH3
[W’04 Q7]
C It is a single atom.
How many nickel atoms are in the coin?
6 5 Lycra
a polyurethane
fibre used
in the fashion industry.
It is a polymer24 made from two
D It® isisnegatively
charged.
22
B has
4.30
× following
1022
C 1.03 × 1023
D 1.20 × 10
A
2.05 × 10
monomers,
one of which
the
formula.
–N=C=O
22 Compound P displays cis-transO=C=N–(CH
isomerism2)nand
gives a red-brown precipitate with Fehling’s
6 Which
ion
has
more
electrons
than
protons
and
more protons than neutrons?
solution.
What
bond
H the
; D O–C–N
= 12 H ; O
= 168angle
O] in this molecule?
[H
= 11is
What
A
90is° P?
B 109 °
C 120 °
D 180 °
[W’07 Q6]
A D–
B H3O+
C OD–
D OH–
© UCLES 2004
9701/01/O/N/04
[Turn over
A
B
COCH
CH2CHO
H
H
7 6 Organic nitrates in photochemical
smog can cause breathing difficulties.3
C C
C C
The diagram shows anCH
example of an organic nitrate
CHmolecule.
H
H
3
3
H
H
© UCLES 2007
1
C
H
C
C
2
3
O
O
NO2
[Turn over
9701/01/O/N/07
O
D
CH
What is the
the bond angles shown in 3ascending order
CHof
H correct order
H (smallest first)?
3
A
1→2→C
3
H
! UCLES 2004
CEDAR COLLEGE
CB
2→1→3
CHO
C
3→1→2 C
H
9701/01/M/J/04
DC 3 → 2 → 1
CHO
[W’10 Q7]
[Turn over
SHAPES OF MOLECULES WS 1
pressure.
If the flasks are connected at constant temperature, what is the final pressure?
property
239
3
A 8 kPa
B 9 kPa
C 10 kPa
D 11 kPa
4
7 Methyl isocyanate, CH3NCO, is a toxic liquid which is used in the manufacture of some
pesticides.
8 Use of the Data Booklet is relevant to this question.
In the methyl isocyanate molecule, the sequence of atoms is H3C — N C O.
The enthalpy change of formation, ∆Hf, of hydrated calcium ions is the enthalpy change of the
What is reaction.
the approximate angle between the bonds formed by the N atom?
following
Be
Mg
Ca
Sr
Ba
–
2+
A
B
C
D
Ca(s) + aq – 2e → Ca (aq)
What is this property?
N are
C notOquoted in the Data
C Oenthalpy changes
N Booklet.
C O
TheNfollowing
A ionic radius
H 3C N C O
HB3C ionisation energy
H3C
Ca(s) → Ca(g) H3C ∆Ha = 177 kJ mol–1
C
Ca2+(g) + aq → Ca2+(aq)
neutron
104° / proton ratio
109°
∆Hhyd = –1565 kJ mol–1
120°
180°
[W’11 1 Q4]
D rate
of reaction
waterof formation of hydrated calcium ions?
What
is the
enthalpywith
change
5 At room temperature
and pressure chlorine does not behave as an ideal gas.
A –1388 kJ mol–1
8
4 The antidote molecule
shown can help to prevent liver damage if someone takes too many
–1
At which
temperature
and pressure would the behaviour of chlorine become more ideal?
B
–798 kJ
mol
paracetamol
tablets.
–238pressure
kJ mol–1
temperature
H
H
H
/ kPa–1
x /K
y
D +352 kJ mol
represents a
A
50 H
C
N
S 200C
lone pair
z
B
50
400
9 The following equilibrium is set up in a mixture of concentrated nitric and sulfuric acids.
H
H
H
C
200
200
H2NO3+ + HSO4–
HNO3 + H2SO4
What
order of decreasing
size of the bond angles x, y and z?
D is the 200
400
Which row correctly describes the behaviour of each substance in the equilibrium mixture?
largest
smallest
6 The standard enthalpy change for the reaction
H SO
H2NO3+
HSO4–
HNO3
A
x
y 2 4 z
(g) → 2N(g) +base
6F(g)
is ∆H o = +1668 kJ
2NF3base
A
acid
B
x
z acid
y
C
B
acid
base
What
is the
of the N–F
C
y bond energy
z base
x bond?
acid
–1
AC
D –556base
zkJ mol
y acid
acid
base
BD –278 base
kJ mol–1
acid
base
acid
C
x
[W’09 1 Q4]
+278 kJ mol–1
109 Which
molecule
kJ mol–1or structure does not contain three atoms bonded at an angle between
D +556
109° and 110°?
© UCLES 2009
A
ethanoic acid
B
graphite
C
propane
D
silicon(IV) oxide
9701/11/O/N/09
[Turn over
[’1 Q]
© UCLES 2010
© UCLES 2011
CEDAR COLLEGE
9701/12/O/N/10
[Turn over
9701/11/O/N/11
[Turn over
SHAPES OF MOLECULES WS 1
∆Hn-1 (g) → XYn(g)
CmolY(g) + XY
12 / kJ
Consecutive
elements X, Y and Z are in Period 3/ kJ
of mol
the Periodic Table. Element Y has the
On heating, Mn(NO3)2 decomposes into MnO2 and NO2.
highest first ionisation energy
and the lowest melting point of these three elements.
n
D nXY(g) → nX(g) + Y2 (g)
2
MnO2(s) + 2NO2(g)
240
What are the identities of X, Y Mn(NO
and Z?3)2(s) →
extent
extent of reaction
What
is the magnesium,
value
of of
thereaction
standard enthalpy change of this reaction?
sodium,
aluminium
10 AIn which
12
pair do the molecules have the same shape as each other?
–1
BA magnesium,
silicon
–242 kJ molaluminium,
A H2 O and CO2
–1
CB aluminium,
silicon,
phosphorus
C
D
–209 kJ mol
B H2 O and SCl 2
DC silicon,
phosphorus,
sulfur
mol–1
+209 kJ
C NH3 and BH3
EA
EA
D +242 kJ mol–1
D SCl 2 and BeCl 2
13 When dealing with a spillage of metallic sodium it is important that no toxic or flammable[W’12
products
1 Q12]
enthalpy
enthalpy
are formed.
o
o
–1
∆H
∆H
/ kJ mol–1
11/ kJ
10
Xmol
is an
element
in Period
13 Why
is the
ionic radius
of a2.chloride ion larger than the ionic radius of a sodium ion?
Which material should be used if there is a spillage of metallic sodium?
In which
fluoride
theone
F – X – F angle the largest?
A chloride
ionishas
AA dilute
hydrochloric
acid more occupied electron shell than a sodium ion.
B CF
C NF
D OF2
A Chlorine
BF3
4
3
has
a higher
proton
number than
sodium.
extent
of reaction
extent of reaction
BB ethanol
Ionic radius increases regularly across the third period.
CC sand
11 Which reaction has an enthalpy change equal to the standard enthalpy change of formation of
612 Which
series
shows
molecules
in is
order
of increasing bond angle?
D water
Sodium
is a metal,
chlorine
a non-metal.
Dpropane?
spray
A CH4 → BF3 → NH3
A 3C(g) + 4H2(g) → C3H8(g)
14 Chlorine
What are
the
trendsused
in the
stated
properties water.
as Group II is descended from magnesium to
14
is widely
to treat
contaminated
Bbarium?
H2O gas
→ CO
2 → BF3
B 3C(g) + 8H(g) → C3H8(g)
3
Which
present in water when chlorine gas has been added is responsible for killing
C
NHspecies
3 → CH4 → CO2
C3Hplace,
C 3C(s)
+
4H2(g) →take
bacteria?
8(g) the elements produced are different from the elements that
decomposition
4 When
nuclear
reactions
D
NH
→
CH
→
H
O
4 equations,
2 of the such as
temperature
firstthe
ionisation
energy
reacted.3 Nuclear
one below,
are used to represent
the changes that
–
[M’16 Q6]
3C(s)
4H2(g)
O2– + carbonate
B →ClC
C HCl
D OCl –
AD Cl
3H8(l)
occur.
7 What
of steam
when
of ice is heated
to 323 C at a pressure of
235 produced
144 1.00 g 89
A is the volume
decreases
+ 10 n of
→decreases
Ba + 336element.
Kr + 3 01 nThis is what he wrote down as his
15
Which
row ofinvestigated
the table is correct?
12
A student
the92 U
chloride
a 56
Period
13 101
kPa?
B
record
of whatdecreases
he did and what he saw.increases
3
3
The
nucleon
is constant
and the
A 0.27
dm3 (mass)Bnumber
1.3 dmtotal
C 2.7at
dm236
D proton
48present
dm3number total is constant
bonds
C
increases shape
decreases
at 92.
ammonia
ammonium
ammonia
ammonium
The
compound
crystalline
It dissolved easily
in water to
D
increases was a white
increasessolid.
2
molecule
ion
molecule
ionin
In anothergive
nuclear
reaction,
uranium-238
reacted
deuterium
12. When is
placed
in awith
test-tube
and atoms,
heated
a An isotope of a
a solution
of pH
1 H.
roaring
flame,
thetwo
compound
new element,
J, is Bunsen
formed as
well as
neutrons.melted after several minutes heating.
A
pyramidal
regular tetrahedral
B
pyramidal
238
+ 21H → J + 2 01 n
regular
92 U tetrahedral
[Turn over
© UCLES
2016 can be deduced from this record?
9701/12/F/M/16
What
C
regular tetrahedral
pyramidal
What is isotope J?
A At least one of the recorded observations is incorrect.
D
regular tetrahedral
pyramidal 240
B 238Pu
C
Np
D
A 238Np
B The compound was magnesium chloride,
MgCl 2.
240
Pu
C The compound was phosphorus pentachloride, PCl 5.
514 Dicarbon monoxide, C2O, is found in dust clouds in space. The structure of this molecule is
C=C=O.
moleculewas
contains
no chloride,
unpaired NaCl.
electrons.
D TheThe
compound
sodium
© UCLES 2012
9701/11/O/N/12
[Turn over
How many lone pairs of electrons are present in a molecule of C2O?
A
1
© UCLES 2014
B
2
C
3
D
4
9701/13/M/J/14
6
A white powder is known to be a mixture 9701/12/F/M/16
of magnesium oxide and aluminium oxide.
© UCLES 2016
[S’13 2 Q9]
[Turn over
[Turn over
100 cm3 of 2 mol dm–3 NaOH(aq) is just sufficient to cause the aluminium oxide in x grams of the
mixture to dissolve.
The reaction occurring is Al 2O3 + 2OH– + 3H2O → 2Al (OH)4–.
800 cm3 of 2 mol dm–3 HCl (aq) is just sufficient to cause all of the oxide in x grams of the mixture
to dissolve.
CEDAR
COLLEGE
SHAPES OF MOLECULES WS 1
The reactions occurring are Al 2O3 + 6H+ → 2Al 3+ + 3H2O
and MgO + 2H+ → Mg2+ + H O.
He used the following
values for2nd
the standard
of combustion
1st
3rd enthalpy
4thchanges5th
6th of carbon and
Chydrogen.
–30.0 kJ mol–1
X –1
800
1600
2400
4300
5400
10 400
D –0.150 kJ mol
carbon = –394 kJ mol–1
241
Y
1000
1800
2700
4800
6000
12 300
hydrogen = –286 kJ mol–1
2
615 Al Cl 3 vapour forms molecules with formula Al 2Cl 6 as it is cooled.
What
could be the identities of X and Y?
He calculated the enthalpy change of formation of ethane to be –140 kJ mol–1.
Section
A from Al Cl 3 to Al 2Cl 6?
What happens to the bond angles during the
change
X
What was his experimental
valueYfor the standard enthalpy change of combustion of ethane?
A Some
decrease,
some
remain
the same.
For each
question
there are
four
possible
answers, A, B, C, and D. Choose the one you consider to
–1
be correct.
kJ mol Sb
–2364
antimony,
arsenic,
BAA Some
increase,
some remain
theAs
same.
1
mol–1As
–1506
arsenic,
antimony, Sb
CBB They
allkJdecrease.
In which reaction –1
does an element undergo the largest change in oxidation state?
–1112
mol Se
selenium,
tellurium, Te
DCC They
allkJincrease.
[S’14 1 Q6]
–
2OH
→ OCl – + Cl – + H2O
AD Cl
2 + kJ
mol–1
D –540
tellurium,
Te
selenium, Se
B
3Cl 2 + 6OH– → Cl O3– + 5Cl – + 3H2O
2+
+ different
3+ shapes?
3+
16 Which
has
species
with
Cr2pair
O7 2–the
+ 6Fe
+ 14H
→of2Cr
+ 6Fe
+ 7H
H12O6, is measured in mmol / l . In Pakistan, the
44 C
In China,
concentration
blood
glucose,
C26O
concentration
ofCO
blood
glucose is measured
in mg / dl .
+
–
BeCl 242–and
2→ MnO2 + 2MnO4 + 2H2O
DA 3MnO
+ 4H
3
3
+
The
is aNH
litre
B unit
CH4 land
4 (1 dm ). The unit dl is a decilitre (0.1 dm ).
2
Use
of
Data
is relevant to this question.
and
BFBooklet
NHthe
3glucose
3concentration of 18.5 mmol / l indicates a health problem.
AC blood
D 68SCl
and H2Ois medically useful because it undergoes a natural radioactive process to give a
Ge2 isotope
The
68 / l converted to mg / dl ?
mmol
What
is
18.5
[S’15occurs
3 Q4]
Ga, which can be used to detect tumours. This transformation of 68 Ge
gallium isotope,
when an electron enters the nucleus, changing a proton into a neutron.
mg / dl does
B an178
mg / dl
C largest
333 mg
/ dl in oxidation
D 3330
mg / dl
33.3 reaction
5 A
In which
element
have the
change
number?
© UCLES 2014
9701/11/M/J/14
[Turn over
68
Which statement
about
the composition
of an atom
of the Ga isotope is correct?
2+
+
3+
3+
2–
A Cr2O7 + 6Fe + 14H → 2Cr + 6Fe + 7H2O
517 A
EachIt of
the
four species
this question
are isolated and gaseous.
has
4 electrons
in itsinouter
p subshell.
B 3OCl – → Cl O3– + 2Cl –
B
It has
13 electrons
in its outer shell.
Which
species
is not planar?
C 5Fe2+ + MnO4– + 8H+ → 5Fe3+ + Mn2+ + 4H2O
+
C
It has 37 neutrons.
2+
C2HO4
D NH3
A
3
D BF
PbO
+ B4H+CH
→3 Sn4+ + Pb2+C + 2H
2 + Sn
2
[S’16 2 Q5]
D Its proton number is 32.
66 Argon
a gas used
electric light
bulbs.
Which is
statement
canto
befill
explained
by intermolecular
hydrogen bonding?
318 Sodium borohydride, NaBH4, and boron trifluoride, BF3, are compounds of boron.
Under
whichhas
conditions
pressure
A Ethanol
a higherofboiling
pointand
thantemperature
propane. will argon behave most like an ideal gas?
What are the shapes around boron in the borohydride ion and in boron trifluoride?
B Hydrogen chloride has a higher boiling point than silane, SiH4.
pressure
temperature
C Hydrogen
iodide forms
solution
when dissolved in water.
borohydride
ion an acidic
boron
trifluoride
A
high
high
DA Propanone
a higher boiling pyramidal
point than propane.
squarehas
planar
B
high
low
B
square planar
trigonal planar
C
low
high
C
tetrahedral
pyramidal
D
low
low
D
tetrahedral
trigonal planar
[W’12 2 Q3]
UCLES 2016
2015
©©UCLES
9701/13/M/J/15
9701/12/M/J/16
CEDAR COLLEGE
9701/11/O/N/12
© UCLES 2012
[Turn
overover
[Turn
SHAPES OF MOLECULES WS 1
242
4
619 Histamine is produced in the body to help fight infection. Its shape allows it to fit into receptors
which expand blood vessels.
H
H
H
H
xN
yC
C
H
H
H
N
C
z
C
N
H
C
H
histamine
What are the bond angles x, y and z in histamine, from the smallest to the largest?
5
smallest
largest
10 The enthalpy
change of the neutralisation
given below is –114 kJ mol–1.
bond angle
bond angle
A
x
2NaOH(aq) + H2SO4(aq) → Na2SO4(aq) + 2H2O(l)
y
z
4
ByBusing thisy information, xwhat is the most
likely
value for the enthalpy change of the following
z
neutralisation?
C solid aluminium
y
z is heated,xAl 2Cl 6 is formed.
6
When
chloride
→ BaCl2(aq) + 2H2O(l)
Ba(OH)
D
z
y 2(aq) + 2HCl(aq)
x
Which bonding is present in Al 2Cl 6?
[W’16 1 Q6]
B –76 kJ mol–1
C –114 kJ mol–1 D –228 kJ mol–1
A –57 kJ mol–1
A covalent and co-ordinate (dative covalent)
7 The approximate percentage composition of the atmospheres on four different planets is given in
table below.
Bthe covalent
only
11
20 Which molecule is planar?
CA
ionic
and co-ordinate
(dative
covalent)C
Which
mixture
of gases
density?
B has
C
Hthe greatest
CH
D NF
C Cl
2
D
4
3
6
3
8
3
[S’04 Q20]
ionic only
major gases /
by number
molecules
12 Use of theplanet
Data Booklet is %
relevant
to thisofquestion.
7 21 In which hydride is the H–X–H bond angle the smallest?
CH
4
Why is the ionic radius ofHa2 sulfide ion He
larger than the
ionic
radius of a potassium ion?
B CH4
C C 2H 6
D NH
A BH3
10.2
0.0 number. 3
A A Ionic Jupiter
radius always 89.8
decreases with
increasing atomic
[W’16 2 Q7]
B
Neptune
80.0
19.0
1.0
B Positive ions have smaller radii than negative ions.
8
In an experiment, a sample of a pure gas is put into a gas syringe at a temperature of 300 K and
C
Saturn
96.3is compressed
3.3
0.4
gas
the volume
by ion.
the gas is halved.
pressure
16 kPa. The
C The of
potassium
ion has
more protons inuntil
its nucleus
thanoccupied
the sulfide
D
Uranus
82.5
15.2
2.3
D The
sulfide ionthe
is doubly
charged;
thegas
potassium
ion is singly
K and the pressure is 40 kPa.
After
compression,
temperature
of the
in the syringe
is 375 charged.
Which statement is correct?
8 Mohr’s
An important
in thecrystalline
manufacture
of which
nitric acid
is the catalytic
of ammonia.
13
salt is reaction
a pale green
solid
is soluble
in water.oxidation
It contains
two cations, one
2+
2–
Aof which
Intermolecular
forces
between
theisgas
are significant.
one anion
which
SOmolecules
is Fe and
4 .
4NH3(g) + 5O2(g)
4NO(g) + 6H2O(g)
B It is possible to calculate the number of moles of gas present using these data alone.
The identity of the second cation was determined by heating solid Mohr’s salt with solid sodium
For every mole of O2 that reacts in this way, 181.8 kJ of energy are released.
hydroxide
and
a colourless
gas was evolved. The gas readily dissolved in water giving an alkaline
C The gas
is behaving
ideally.
solution.
5
makes 2.50
× 10
NOforevery
DA factory
The pressures
used
aremol
too of
high
idealday.
gas behaviour.
A grey-green solid residue was also formed which was insoluble in water.
How much energy, in kJ, is released every day?
9
of athe
fuelsolid
are residue?
burnt. 45.0% of the energy 8released is absorbed
In
a calorimetric
experiment
What
are the identities
of the1.60
gasgand
A 3.64 × 107
B 4.55 × 107
C 5.68 × 107
D 2.27 × 10
by 200 g of water. The temperature of the water rises from 18.0 °C to 66.0 °C.
gas
residue
© UCLES
2016
What
is the total energy released per gram9701/11/O/N/16
of fuel burnt (to 3 significant figures)?
A
NH3
25 100 J
A
B
B
NH3
CEDAR COLLEGE
C
SO2
D
SO2
Fe(OH)2
55 700 J
Na2SO4
C
89 200 J
D
143 000 J
SHAPES OF MOLECULES WS 1
Fe(OH)2
Na2SO4
C
More electrons in the molecules results in stronger van der Waals’ forces.
D
More hydrogen atoms in the molecules results in stronger hydrogen bonding.
243
522 The characteristic smell of garlic is due to alliin.
H
H
xC
C
O
y
CH2
S
CH2
H
alliin
NH2
z
C H
CO2H
What are the approximate bond angles x, y and z in a molecule of alliin?
x
y
z
A
90o
90o
109o
B
120o
109o
90o
2
Section A
o
For each
possible
C question
120o there are
120four
109o answers, A, B, C and D. Choose the one you consider to
be correct.
D
180o
109o
109o
Use of the Data Booklet may be appropriate for some questions.
[M’17 Q5]
6 Which gas sample contains the fewest molecules?
123 Which feature is present in both ethene and poly(ethene)?
A 1.00 dm3 of carbon dioxide at 27 °C and 2.0 kPa
A bond angles of 109°
B 1.00 dm3 of hydrogen at 100 °C and 2.0 kPa
B π covalent bonds
C 1.00 dm3 of nitrogen at 300 °C and 4.0 kPa
C σ covalent bonds
D 1.00 dm3 of oxygen at 250 °C and 3.0 kPa
D sp3 orbitals
3
[S’18 2 Q1]
424 Which statement describes the bond between carbon2and
hydrogen in an ethene molecule?
2 The electronic configuration of an atom of sulfur is 1s 2s22p63s23p4.
A a π bond between an s orbital and an sp2 orbital
How many valence shell and unpaired electrons are present in one sulfur atom?
B a π bond between an s orbital and an sp3 orbital
valence shell
unpaired
C a σ bond
between an selectrons
orbital and an sp2 orbital
electrons
D
5
A
a σ bond between an s orbital and an sp3 orbital
2
1
[S’18 3 Q4]
B
4
2
9701/12/F/M/17
[Turn over
Aspirin, C9H8O4, Mr = 180.0, can be made by a reaction between 2-hydroxybenzoic acid, C7H6O3,
6 ethanoic anhydride,
0
MrC= 138.0, and
C4H6O3, Mr = 102.0. The balanced equation for the reaction
is D
shown.
6
2
© UCLES 2017
C7H6O3 + C4H6O3
3
In which pair does the second substance have a lower boiling point than the first substance?
If a reaction mixture consists of 10.0 g of each of the two reactants, what is the maximum mass of
aspirin
that
canCbe
and
H produced?
Cl
A CH
2
A
B
6
C9H8O4 + C2H4O2
6
2
5
10.0 g
5.7
CH3gOCH3 and CB2H5OH
C
13.0 g
D
17.6 g
C Ne and Ar
Which diagram correctly describes the behaviour of a fixed mass of an ideal gas? (T is measured
D K.)CH3NH2 and C2H6
in
C
B
D
A
Compound J burns in excess oxygen to give carbon dioxide and water only. When a 3.00 g
sample
of compound
J is burnt
in excess
of carbon
g of water
constant
p
constant
T oxygen, 4.40 gconstant
T dioxide and 1.80
constant
T are
formed.
CEDAR
COLLEGE
SHAPES OF MOLECULES WS 1
4
VWhat is the empirical formula
p
of J?
pV
pV
second oxygen atom
electron from
barium atom
244
25
5 In this question, the methyl group, CH3, is represented by Me.
Trimethylamine, Me3N, reacts with boron trifluoride, BF3, to form a compound of formula
Me3N.BF3.
How may this reaction be written in terms of the shapes of the reactants and products?
!
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[S’08 Q5]
© UCLES 2008
CEDAR COLLEGE
9701/01/M/J/08
[Turn over
SHAPES OF MOLECULES WS 1
11
245
14
12
Section B
SECTION
Section
B
Section
BB
For each of the questions in this section, one or more of the three numbered statements 1 to 3 may
be
For each
of the
questions
in this
section,
or more
of the
three
numbered
statements
may
Forcorrect.
each
of the
questions
in this
section,
oneone
or more
of the
three
numbered
statements
11
toto3 3may
be
correct.
be correct.
Decide whether each of the statements is or is not correct (you may find it helpful to put a tick against
the
statements
that
you
to be correct).
Decide
whether
each
of the
statements
is is
ornot
is not
correct
(you
may
find
it helpful
put
tickagainst
against
Decide
whether
each
of consider
the
statements
is or
correct
(you
may
find
it helpful
to to
put
aa
tick
the
statements
that
you
consider
to
be
correct).
the statements that you consider to be correct).
The responses A to D should be selected on the basis of
The responses A to D should be selected on the basis of
The responses A to D should be selected on the basis of
11
A
B
C
D
A
B
C
D
Section B C
A
B
D
1 only
2 and 3
1 and 2
1, 2 and 3
1 only
2 are
and 3
1are
and 2
1, 2 and 3
is
only
2 and
3are numbered1statements
1 and
2are
and
3
For each of1,the
questions
in this only
section,
one or more only
of
the
1 to 3 may be
is
onlythree
only2 are
are
correct
correct
correct
correct
14
is
only
are
only
are
are
correct.
correct
correct
correct
correct
correct
correct
correct
correct
Decide
whether
each
the
statements
is oron
isthe
not
correct
The
responses
A
to Dof
be selected
basis
of (you may find it helpful to put a tick against
No other
combination
ofshould
statements
is used
as
a
correct
response.
the statements
that you consider
to be correct).
No other combination
of statements
is used as a correct response.
No other combination of statements is used as a correct response.
A
B
C
D
The
to D should
be selected
on theofbasis
of is 72.
31 responses
The relativeA molecular
mass
of a molecule
chlorine
31 The diagram
the
energy
changes of a2 set
of 3reactions.
1 only
and
andfollowing
3 illustrates
1 and
2 ions
311 Which1,of2 the
molecules
and
have a regular
trigonal planar shape?
Which properties
of
the
atoms
in
this
molecule
are
the
same?
only
are
only
are
are
A
B_
C
D is
_
1
correct
correct
correct
H = 134 kJ mol correct
1 Al Cl
3
1 radius 1, 2 and 3 are R 1 and 2 only are 2 and 3S only are 1 only is correct
+
2 CH3
correct
correct
correct
2 nucleon
number
No other
combination
of statements is used as a correct response.
3 PH3
3 relative isotopic mass
[S’05 Q31]
No other combination of statements is used as a correct response.
34 Ammonia and chlorine react in the gas phase.
1
H = +92 J mol
32 A quantity of solid Y was placed in a previously evacuated vessel and the apparatus was then
322 Which
molecules
planar?
held atpairs
a series
ofare
different
temperatures.
At
temperature,
31 Which
of compounds
have
the 3same
empirical
8NH
+ 3Cl
→ N
+ 6NH4Clthe mass of Y in the vapour state
2each
2formula?
was
calculated
from
pressure
measurements.
The
results
are shown below.
1 BCl3
1
ethane
and ethene
Which
statements
are correct?
_
2 NH3
H = _75 kJ mol 1
2
ethenenitrogen
and cyclohexane
mass
U
1 Each
atom
isTof
oxidised.
vapour
3 PH3
3
cyclohexane and oct-1-ene
[W’05 Q32]
2
Each chlorine
atom isstatements
reduced.
Which
of the following
are correct?
m
–1
behaves
as
a base.
1Ammonia
enthalpy
change
for the
transformation
U ⎯→ aluminium
R is + 42 kJinmol
. III of the Periodic
33 3
Boron
isThe
a non-metallic
element
which
is placed above
Group
323 In which sequences are the molecules quoted in order of increasing bond angle within the
Table. It forms a compound with nitrogen known as boron nitride which has a graphite structure.
molecule?
2 The enthalpy change for the transformation T ⎯→ S is endothermic.
35 Which
statements
about
calcium and
strontium
Which of
the following
conclusions
can
be drawn compounds
from
this T
information?
transformation
R ⎯→
isare
– 33correct?
kJ mol–1 .
NH3 change
CH4for the T
1 3H2OThe enthalpy
temperature
1 When
calciumformula
oxide and
strontium
are added to water they both produce alkalis.
The empirical
of boron
nitrideoxide
is BN.
2
H2O
SF6
BF3
What can
be deduced
from the
diagram?
22 Calcium
hydroxide
isSF
more soluble
than
strontium
hydroxide.
The
and nitride
atoms
are likely
to be
arranged
alternately in a hexagonal pattern.
3
CH4 boron
CO
2
6
1 The mass of Y used in the experiment was m.
[S’02 Q32]
Boron nitride
hasisaless
layersoluble
structure
with
van der Waals’
3 Calcium
sulfate
than
strontium
sulfate.forces between the layers.
2 The pressure of the vapour was constant for all temperatures above temperature T.
33 The concepts of bond energy, bond length and bond polarity are useful when comparing the
ofappeared
similar molecules,
e.g. thermal
stability.
3 Liquid
temperature
T. ion are
36 4 behaviour
Which
descriptions
ofatthe
ammonium
correct?
For
it could
be said
1 example,
It contains
ten electrons.
2
It has awith
bond
of 109.5°. the bond ……….X…………. of the HI molecule is
“Compared
theangle
HCl molecule,
………..Y……….. .”
3 It has only three bonding pairs of electrons.
Which
pairs of words correctly complete the 9701/01/O/N/04
above sentence?
© UCLES
© UCLES
2005 2004
9701/01/O/N/05
[Turn over
37
Compound
Q is obtained by adding H2O9701/01/M/J/05
across the double bond in compound P.
© UCLES
2005
X
Y
1
energy greater
CEDAR COLLEGE
2
length
greater
3
polarity
less
OH
SHAPES OF MOLECULES WS 1
Ar of
boron
in theissample
is to
10.83
to four significant figures.
31 2UseThe
of the
Data
Booklet
relevant
this question.
3 There are99more protons in Y+ than in X+.
The isotope Tc is radioactive and has been
246 found in lobsters and seaweed adjacent to nuclear
fuel reprocessing plants.
325 Which elements can form π bonds in their compounds?
12 of 99Tc?
Which statements are correct about an atom
1 carbon
Section B
1 It has 13 more neutrons than protons.
2 oxygen
For each
of
the questions
in this section, one or more of the three numbered statements 1 to 3 may
2
It has
43 protons.
be correct.
3 nitrogen
3 It has 99 nucleons.
Decide whether each of the statements is or is not correct (you may find it helpful to put a tick against
the statements
you consider
correct).
33
For which that
enthalpy
changestoisbethe
value of ∆H always negative?
326 Urea is a product of animal metabolism. It can also be used as a fertiliser.
The responses
A to D should be selected on the basis of
1 combustion
H
2 hydration
A
B
C
D
N
x 2 and 3
1 only
1 and 2
1, 2 and 3
3 solution
H
O
Conly are
is
only are
are
correct
correct H
correct
correct
N
No other combination of statements is used as a correct response.
H
Use of the Data Booklet may be appropriate for 9701/11/O/N/16
someurea
questions.
© UCLES 2016
The diagram shows angle x in this molecule.
31 One mole of sulfuric acid is used to make an aqueous solution. The solution contains H2SO4
2–
molecules,
H+ ions, SO
ions
HSO4–ofions.
4
Which statements
about
the and
structure
urea are correct?
Which
statements
are correct? 120°.
1 Angle
x is approximately
12
The
contains
6.02π ×bonds.
1023 sulfur atoms.
Thesolution
molecule
has two
23 The
contains
an exactly
equalpairs
number
of H+ ions and HSO4– ions.
Thesolution
molecule
has only
three lone
of electrons.
3
One mole of SO42– ions contains two moles of electrons.
32
Which
statements are correct?
©7UCLES
2015
[S’15 1 Q32]
9701/11/M/J/15
1
The hydrogen bonds in ice are more regularly arranged than in water.
2
The solidification of water to form ice is exothermic.
3
Pure water is less dense than ice.
[S’18 1 Q32]
33 Calcium reacts with water to form calcium hydroxide and hydrogen.
Ca(s) + 2H2O(l)
Ca(OH)2(s) + H2(g)
The standard enthalpy change for this reaction is – 414 kJ mol–1.
What further information is needed in order to calculate the standard enthalpy change of
formation of calcium hydroxide,
Ca(OH)2(s)?
1
for H2O(l)
2
for H2(g)
3
first and second ionisation energies of Ca
© UCLES 2018
CEDAR COLLEGE
9701/11/M/J/18
SHAPES OF MOLECULES WS 1
247
3 Chemical bonding
This topic introduces the different ways by which chemical bonding occurs and the
effect this can have on physical properties.
3.3 Intermolecular forces, electronegativity and bond properties
INTERMOLECULAR FORCES
BILAL HAMEED
INTERMOLECULAR FORCES
and co-ordinate
(dative covalent)
bonding including
shapes of simple
molecules
Cambridge
International
AS andsuch
A Level
Chemistry 9701
syllabus
Syllabus content
(i) covalent
bonding,
in molecules
as hydrogen,
oxygen,
chlorine,
hydrogen chloride, carbon dioxide, methane, ethene
(ii) co-ordinate (dative covalent) bonding, such as in the formation of the
248
ammonium ion
and in the Al 2Cl 6 molecule
b) describe covalent bonding in terms of orbital overlap, giving σ and π
bonds, including the concept of hybridisation to form sp, sp2 and sp3
3 Chemical bonding
orbitals (see also Section 14.3)
This topic introduces the different ways by which chemical bonding occurs and the effect this can have
c) explain the shapes of, and bond angles in, molecules by using the
on physical properties.
qualitative model of electron-pair repulsion (including lone pairs), using
as simple examples: BF3 (trigonal), CO2 (linear), CH4 (tetrahedral), NH3
Learning
outcomes
(pyramidal),
H2O (non-linear), SF6 (octahedral), PF5 (trigonal bipyramidal)
Candidates
should
be able
to: bond angles in, molecules and ions analogous
d) predict the shapes
of, and
3.1 Ionic bonding
3.3 Intermolecular
forces,
electronegativity
3.2 Covalent
bonding
andco-ordinate
bond properties
and
(dative covalent)
bonding including
shapes of simple
molecules
3.4 Metallic bonding
to those specified in 3.2(b) (see also Section 14.3)
a) describe ionic bonding, as in sodium chloride, magnesium oxide and
including
theusing
use of
‘dot-and-cross’
diagrams
a) calcium
describefluoride,
hydrogen
bonding,
ammonia
and water
as simple
examples of molecules containing N–H and O–H groups
a)
including
the terms,
use of the
‘dot-and-cross’
diagrams:
b) describe,
understand,
in simple
concept of electronegativity
and apply
it
to
explain
the
properties
of
molecules
such
as
bond
polarity
(i) covalent bonding, in molecules such as hydrogen, oxygen, (see
chlorine,
alsohydrogen
Section 3.3(c)),
thecarbon
dipoledioxide,
moments
of molecules
(3.3(d)) and the
chloride,
methane,
ethene
behaviour
of oxides
with
water (9.2(c))
(ii)
co-ordinate
(dative
covalent)
bonding, such as in the formation of the
c) explain
the terms
energy,
length and bond polarity and use
ammonium
ionbond
and in
the Al bond
2Cl 6 molecule
them to compare the reactivities of covalent bonds (see also Section
b) describe covalent bonding in terms of orbital overlap, giving σ and π
5.1(b)(ii))
bonds, including the concept of hybridisation to form sp, sp2 and sp3
d) orbitals
describe
intermolecular
(see
also Sectionforces
14.3) (van der Waals’ forces), based on
permanent and induced dipoles, as in, for example, CHCl 3(l); Br2(l) and
c) explain the shapes of, and bond angles in, molecules by using the
the liquid Group 18 elements
qualitative model of electron-pair repulsion (including lone pairs), using
as simple examples: BF3 (trigonal), CO2 (linear), CH4 (tetrahedral), NH3
a) (pyramidal),
describe metallic
bonding in terms
of a lattice PF
of 5positive
SF6 (octahedral),
(trigonalions
bipyramidal)
H2O (non-linear),
surrounded
by
delocalised
electrons
d) predict the shapes of, and bond angles in, molecules and ions analogous
to those specified in 3.2(b) (see also Section 14.3)
3.5 Bonding and
physical properties
3.3 Intermolecular
forces,
electronegativity
and bond properties
Back to contents page
a) describe, interpret and predict the effect of different types of bonding
(ionic bonding,
covalent
bonding,
other
intermolecular
a) describe
hydrogen
bonding,
usinghydrogen
ammonia bonding,
and water
as simple
interactions,
metallic
bonding)
on
the
physical
properties
of
substances
examples of molecules containing N–H and O–H groups
b) understand,
deduce the type
of bonding
from of
given
information and apply
b)
in simple
terms,present
the concept
electronegativity
c) itshow
understanding
of chemical
reactions
in terms
ofpolarity
energy (see
transfers
to explain
the properties
of molecules
such
as bond
associated
the the
breaking
makingof
ofmolecules
chemical bonds
also
Sectionwith
3.3(c)),
dipoleand
moments
(3.3(d)) and the
behaviour of oxides with water (9.2(c))
c) explain the terms bond energy, bond length and bond polarity and use
them to compare the reactivities of covalent bonds (see also Section
5.1(b)(ii))
www.cie.org.uk/alevel
19
d) describe intermolecular forces (van der Waals’ forces), based on
permanent and induced dipoles, as in, for example, CHCl 3(l); Br2(l) and
the liquid Group 18 elements
3.4 Metallic bonding
a) describe metallic bonding in terms of a lattice of positive ions
surrounded by delocalised electrons
3.5 Bonding and
physical properties
a) describe, interpret and predict the effect of different types of bonding
(ionic bonding, covalent bonding, hydrogen bonding, other intermolecular
interactions, metallic bonding) on the physical properties of substances
b) deduce the type of bonding present from given information
c) show understanding of chemical reactions in terms of energy transfers
associated with the breaking and making of chemical bonds
Back to contents page
BILAL HAMEED
www.cie.org.uk/alevel
INTERMOLECULAR FORCES
19
(#, EdaVg^in
249
:aZXigdcZ\Vi^k^in
BOND POLARITIES
In a covalent bond between two different atoms, the atoms do not
attract
thetwo
electron
pairatoms
in thebond
bond
equally.
How strongly
theareelectrons
The example
in which
hydrogen
is simple
because
both atoms
the same. are attracted depends on the size of the individual atoms and their
nuclear
charge.
Also, each one
has a single
proton and a single electron, so the attractions are easy
molecule,
the two F atoms attract the electrons in the bond
In
the
F
2
to identify.
equally and thus the electrons lie symmetrically (Figure 3.87). This
However, many covalent bonds form between two different atoms.
molecule is non-polar.
These atoms often have different attractions for shared electrons.
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However, in HF, F is more electronegative than H and attracts the
electrons in the H–F bond more strongly than the H atom does. The
electrons in the bond thus lie closer to the F than to the H (Figure 3.8
H–F is a polar molecule. The unsymmetrical distribution of electron
density results in small charges on the atoms. F is E− because the electro
E
EÄ
However,
in HF,has
F isbeen
morepulled
electron
in the bond lie closer to F, whereas electron
density
aw
=
;
electrons in the H–F bond more stron
from H, so it is E+.
in the bond thus lie closer to
BOND POLARITIESelectrons
V
W
H–F is a polar molecule. The unsymm
Now consider a diatomic molecule composed
=
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density results
=
;
E
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of two different elements; HF is a common
in ZaZXigdcha^Z!dcVkZgV\Z!
the bond lie closer to F, whereas el
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However, in
tes a small negativeelectrons in the H—F bond are not equally
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Z
EVja^c\ZaZXigdcZ\Vi^k^i^Zh
electrons
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th
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H–F istoa realise
polar
This is because fluorine
is
a
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;^\jgZ(#--
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that, although they are derived from
physical quantities
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E
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electronegative
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than
hydrogen.
energies),
the numbers themselves are not physical
they
havelie
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fromcommonly
H, so it i
charge. units and must only be used in a comparative way. The most
EVja^c\ZaZXigdcZ\Vi^k^i^Zh
2
used scale of electronegativity is that developed by Linus Pauling. The
There areare
various
ofV electronega
electronegativity values for some elements
shownscales
in Table
3.11.
that, although they are derived from p
=
energies),
the
numbers
themselves
are
Noble gases do not have electronegativity values as they do not
als have higher
units and must only be used in a com
form compounds.
HAMEED
INTERMOLECULAR FORCES
gativities thanBILAL
metals.
used scale of electronegativity is that d
V
electronegativity values ;^\jgZ(#--
for some elem
bond enthalpy reflects the strength of the covalent bond. As we
ve from single to double to triple bonds the number of electrons in
250
bond increases resulting in an increase in electrostatic
forces and a
tening of the bond length. This trend can be seen in the carbon–
on bond of the homologous series of alkanes, alkenes, and alkynes.
nd polarity
151
I–I
Table 1 Average bond enthalpies
at 298 K
POLAR BONDS
polarity of a bond can be described by the difference in
Polar covalent bond is the preferred term for a bond made up of
tronegativity of the bonded atoms (tables 3 and 4).
m
r
unequally shared electron pairs.
Electronegativity
Bond 4(Electronegativity) Bond enthalpy/
One end of the bond (in this case, the F atom)
–1 more electron rich
kJ molis
2.2
(higher electron density), hence,
more negative.
1.8
567
H–F
4.0
3.2
3.0
The other end
(in this case, the431
H atom) is less electron rich
H–Clof the bond1.0
δ(lower electron density), hence,
0.8 more positive.
366
H–Br
These two ends,
one somewhat positive and the other
Table 4 Bond polarity and bond enthalpy at 298 K
le 3 Electronegativity values
somewhat negative may be described as electronic
orine has the highest electronegativity value of any element.
poles,
hence
the term
bonds.
polarity of the
H–F bond
results
in a polar
partialcovalent
charge on
each
m. The bond is said to have ionic character. The partial charges
3
act one another, increasing the strength of the bond (figure 1).
δ+
H
F
Figure 1 The polar H–F bond
153
POLAR BONDS
In a polar covalent bond, the shared electrons, which are in a molecular orbital,
are more likely to be found nearer to the atom whose electronegativity is higher.
This unequal distribution of charge makes the bond polar covalent.
To emphasize the dipole nature of the HF molecule, the formula can be written
as Hδ+ Fδ−. The symbol δ means partial.
With polar molecules, such as HF, the symbol δ+ is used to show a partial
positive charge on one end of the molecule.
Likewise, the symbol δ− is used to show a partial negative charge on the other
end.
4
BILAL HAMEED
INTERMOLECULAR FORCES
251
ELECTRONEGATIVITY
Electronegativity is a measure of the ability of an atom to attract electrons
in a chemical bond. Elements with high electronegativity have a greater
ability to attract electrons than do elements with low electronegativity.
The four most electronegative elements are F, O, N, Cl.
increases
electronegativity increases
Electronegativity increases across a period and decreases down a group
5
BOND POLARITIES
Non-polar bond
Polar bond
• Similar atoms have the same
• Different atoms have different
electronegativity
• They will both pull on the
electrons to the same extent
• The electrons will be equally
shared
electronegativities
• One will pull the electron pair closer to its end
• It will be slightly more negative than average, δ• The other will be slightly less negative, or more
positive, δ+
• A dipole is formed and the bond is said to be
polar
• Greater electronegativity difference = greater
polarity
6
BILAL HAMEED
INTERMOLECULAR FORCES
252
MOLECULAR
Linus Pauling
Linus Pauling
related
Linus the
related
Pauling
electronegativity
the
related
electronegativity
the POLARITY
electronegativity
difference
diffbetween
erence
diffbetween
erence between
Exam
two atoms
two
thetwo
ionic
toatoms
the
character
ionic
to
the
character
of
ionic
a between
bond.
character
of aHe
bond.
suggested
of
He
a bond.
suggested
that
Heansuggested
that an that an
Thetoatoms
electronegativity
difference
two
atoms
covalently
bonded
electronegativity
electronegativity
difference
erence
1.7diff
corresponded
oference
1.7 more
corresponded
of 1.7
tocorresponded
50%one
to
ionic
50%
character
ionic
to 50%
character
in
ionic character
in If aske
in
togetherelectronegativity
results
in diff
theof
electrons
lying
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atom
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a bond and
aother.
bond
reasoned
and
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bond
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and
greater
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a electronegativity
greater
that aelectronegativity
greater electronegativity
difference
diffthan
erence
this
diffthan
erence
thisthan this
We
such
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corresponded
corresponded
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a structure
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that
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was
ionic
more
covalent,
than
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than covalent,
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However, whether an overall molecule is polar also depends on the
differe
if the diff
ifshape
erence
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erence
the
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than
iserence
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the
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iscovalent
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than ionic.
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ofifis
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This is aThis
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it must
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The polarity of molecules is distinct from the polarity of individual bonds;
less th
KI (electronegativity
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(electronegativity
KI (electronegativity
difference
diff1.7)
erence
would
diff1.7)
erence
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of this discussion
a non-polar molecule may have polar bonds.
as having
as 50%
having
ionic
as50%
having
and
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50%
50%
and
covalent
ionic
50%and
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character,
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(diffNaI
erence
and
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erence
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would
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be covalent,
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covalent,
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whereas
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bothasbehave as
predominantly
predominantly
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predominantly
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Pauling,
compounds.
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in his original
Pauling,
in his original
discussion
in his original
discussion
of discussion
of
of
this, wasthis,
actually
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was referring
actually
to diatomic
referring
to diatomic
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to diatomic
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and not
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and
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not and
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not to macroscopic
compounds.
compounds.
compounds.
7
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The electronegativity
The electronegativity
The electronegativity
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diffbetween
erence
two atoms
between
twocovalently
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atoms
bonded
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results in
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results
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the electrons
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lying
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POLAR
polar also
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depends
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be
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must
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have
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NH
andand
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For
instance,
HCl,
NH
and
HNH
are
allNH
polar.
3 HCl,
2O
3 and
32O
23O
2O are all
polar (Figure
polar
(Figure
3.89).
polar
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3.89).
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molecules
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Y^edaZbdb
Y
These molecules all have an overall dipole moment, and the arrow
and the and
arrow
the
indicates
arrow
and
the
indicates
the
arrow
direction
indicates
the
direction
of
the
direction
dipole
of
the
moment.
dipole
of
the
moment.
dipole
moment.
edh^i^kZ cZ
ed
indicates the direction of the dipole moment.
D
=
D Ä D D D Ä DD Ä D
DÄ DÄ
D D
DÄ
DDÄ DD
C = 8a
8a = = 8a =
= D C = = D C === D
=D
=D
DÄ
D
DÄ
DD
==
D
D
=
=D
;^\jgZ(#-.
;^\jgZ(#-.
I]ZhZbdaZXjaZhVgZVaaedaVg/dcZZcYd[i]ZbdaZXjaZ^hha^\]ian
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edh^i^kZXdbeVgZYl^i]i]Zdi]ZgZcY#I]ZVggdl^cY^XViZhi]ZY^edaZbdbZci#
edh^i^kZXdbeVgZYl^i]i]Zdi]ZgZcY#I]ZVggdl^cY^XViZhi]ZY^edaZbdbZci#
edh^i^kZXdbeVgZYl^i]i]Zdi]ZgZcY#I]ZVggdl^cY^XViZhi]ZY^edaZbdbZci#
8
Saying
momen
Although
Although
individual
Although
individual
bonds
individual
may
bonds
bemay
polar,
bonds
bethe
may
polar,
overall
bethe
polar,
overall
molecule
the molecule
overall molecule
saying t
may bemay
non-polar
be may
non-polar
if,beowing
non-polar
if, to
owing
theif,symmetry
to
owing
the symmetry
to of
thethe
symmetry
molecule,
of the molecule,
ofthe
the molecule,
the
the
BILAL HAMEED
INTERMOLECULAR
FORCES
dipole
moments
dipole moments
dipole
of themoments
individual
of the individual
of bonds
the individual
cancel
bonds out.
cancel
bonds
out.
cancel out.
that the
bonding between
a bond and reasoned that a greater electronegativity difference
this
a bond than
and reasoned
that a greater
electronegativity
diffe
two
elements
is
ionic
if
the
corresponded to a structure that was more ionic than covalent,
whereas
corresponded
to a structure that was more ionic than cov
diffthe
erence
is more
than cov
1.7
f the difference is less than that the bonding is more covalent
than
ionic.is less than that
if253
the diff
erence
bonding
is more
andbecovalent
if the
diffcautio
eren
This is a useful idea, but it must be used with great caution.
ThisFor
is a instance,
useful idea, but it must
used with
great
less1.7)
than
1.7. come out of
KI (electronegativity difference 1.7) would come out ofKI
this(electronegativity
discussion
difference
would
having
50% ionic and 50% covalent character, and NaI
as having 50% ionic and 50% covalent character, and NaIas (diff
erence
of behave
1.6) would
of 1.6) would appear to be mostly covalent,POLAR
whereas both
as appear to be mostly covalent, whereas both
MOLECULES
predominantly ionic compounds. Pauling, in his originalpredominantly
discussion of ionic compounds. Pauling, in his original
this,
actually referring to diatomic molecules and not
his, was actually referring to diatomic molecules and not
to was
macroscopic
δ+
compounds.
compounds.
CHCl3
EdaVgbdaZXjaZh
Cl
NH3
δ+
δ+
δ-
EdaVgbdaZXjaZhO
H
δ-
δ-
C
N
C
H
Cl
The electronegativity
difference between
two atoms cova
Cl
The electronegativity difference
between
two atoms Hcovalently
bonded
δδ+
δ+
Cl
H
Cl
δtogether
results
lying
more towards one a
δδ+
δ+
ogether results in the electrons
lying
more
towards one
atomδ+than
the in theδ-electrons
COCl
δ+
δcall suchis a bond polar. However, whether an o
other. We call such a bond polar. However, whether an other.
overallWe
molecule
polar also depends on the shape of the molecule.
polar also depends on the shape of the molecule.
a molecule to be polar it must have a positive en
For a molecule to be polar it must have a positive end For
to the
and a negative end. For instance, HCl, NH3
molecule and a negative end. For instance, HCl, NHmolecule
3 and H2O are all
polar (Figure 3.89). These molecules all have an overall d
polar (Figure 3.89). These molecules all have an overall dipole moment,
and the arrow indicates theY^edaZbdbZci
direction of the dipole mome
and the arrow indicates the direction of the dipole moment.
edh^i^kZ cZ\Vi^kZ
2
9
D
=
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DÄ
DÄ
D
C
=
=D
D
=
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D
D
=
=
DÄ
8a
DÄ
DÄ
D
C
=
D
=
=D
D
D
=
=D
=D
;^\jgZ(#-. I]ZhZbdaZXjaZhVgZVaaedaVg/dcZZcYd[i]ZbdaZXj
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edh^i^kZXdbeVgZYl^i]i]Zdi]ZgZcY#I]ZVggdl^cY^XViZhi]ZY^e
edh^i^kZXdbeVgZYl^i]i]Zdi]ZgZcY#I]ZVggdl^cY^XViZhi]ZY^edaZbdbZci#
Saying that a molecule has a
individual bonds
may be
polar,
the overa
NON-POLAR Although
MOLECULES
moment
is just
another
way
Although individual bonds may be polar, the overallmay
molecule
be non-polar if, owing
tothat
the itsymmetry
saying
is polar. of the
Although
individual
may beofpolar,
the
overall molecule
may be non-polar
if, owing
to thebonds
symmetry
the dipole
molecule,
the
moments of the individual bonds cancel ou
if, owing
the symmetry
dipole momentsmay
of be
thenon-polar
individual
bondstocancel
out. of the molecule,
the dipole moments of the individual bonds cancel out.
CO2 is a non-polar molecule. Each C–O bond is pola
molecule.
Each
C–O
bond
is
polar,
because
oxygen
CO2 is a non-polar
is bond
more
than
CO2 is a non-polar molecule. Each C–O
iselectronegative
polar,
DEÄ carbon,
8E but overall
DEÄ the dipol
s more electronegative
thanoxygen
carbon,isbut
overall
the dipoles
cancel
so
thatbutdipole moment and the molecule is no
there
is no
overall
because
more
electronegative
than
carbon,
Y^edaZhXVcXZa
here is no overall overall
dipole moment
andcancel
the molecule
is non-polar.
is also non-polar.
Again,
each individual bond is p
the dipoles
so that there
isBF
no3 overall
dipole
Again, each individual bond is polar but the
BF3 is also non-polar.
moment and the molecule is non-polar.dipoles cancel.
dipoles cancel.
BF3 is also non-polar. Again, each individual bond is polar but
;EÄcancel. ;EÄ
the dipoles
7
E
;EÄ
BILAL HAMEED
;EÄ
7
E
;EÄ
;EÄ
10
INTERMOLECULAR FORCES
(7
means that the b
That we simp
intermolecular fo
can be seen if we
254
NON-POLAR MOLECULES
In contrast, a molecule containing polar bonds may be either polar or non- 8dbedjcY
polar depending on the relative arrangement of the bonds and any lone pairs
but CHCl
is polar
(Figure
3.90).
4 is non-polar,
3but
CCl
non-polar,
CHCl
polar
(Figure 3.90).
4 isCCl
3 isCHCl
ofCCl
electrons,
e.g.
4 is non-polar, but
3 is polar.
7g'
:miZch^dc
dkZgVaaY^edaZ
dkZgVaaY^edaZ
p
W
V
W
V
why as to why
nswer
D
D
Ä
8>
=
D =
8> D Ä
gh ‘although
olar:
DÄ
DÄ
olar
8D DÄ 8D
8D DÄ 8D
l bond is polar
8>
8>
D
Ä
8>
8>8>D Ä
8>
8> D Ä
8>D Ä
erence in
Ä
DÄ
D
8>
8>
8>D Ä
8>D Ä
oms,
ty of the atoms,
Y^edaZhXVcXZa Y^edaZhXVcXZa
of
the
symmetry
of the;^\jgZ(#.% V88a ^hcdc"edaVgWZXVjhZi]Z^cY^k^YjVaY^edaZhXVcXZa#W8=8a ^hedaVg
)
(
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(^hedaVg
el’.
ipoles cancel’. WZXVjhZi]ZY^edaZhYdcdiXVcXZa0i]ZgZ^hVedh^i^kZZcYidi]ZbdaZXjaZVcYV
WZXVjhZi]ZY^edaZhYdcdiXVcXZa0i]ZgZ^hVedh^i^kZZcYidi]ZbdaZXjaZVcYV
Explain why BrF has a higher boiling
point than ClBr, despite ClBr having
the greater relative molecular mass.
cZ\Vi^kZZcY#6ai]dj\]i]Z8^c8=8a
!^i^hcdiVhedh^i^kZVhi]Z=Vh
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!^i^hcdiVhedh^i^kZVhi]Z=Vh
(^hh]dlcVhE
8^hbdgZZaZXigdcZ\Vi^kZi]Vc=0i]ZgZ[dgZ!i]Z8^hha^\]iancZ\Vi^kZXdbeVgZYl^i]
8^hbdgZZaZXigdcZ\Vi^kZi]Vc=0i]ZgZ[dgZ!i]Z8^hha^\]iancZ\Vi^kZXdbeVgZYl^i]
i]Z=!Vai]dj\]^i^hedh^i^kZdkZgVaa^ci]ZbdaZXjaZ#
i]Z=!Vai]dj\]^i^hedh^i^kZdkZgVaa^ci]ZbdaZXjaZ#
11
Some polarSome
and non-polar
shown in
3.12.
polar and molecules
non-polar are
molecules
areTable
shown
in Table 3.12.
uct
of product
the
s the
of the
EdaVg
Cdc"edaVg Cdc"edaVg
EdaVg
ween
stancethe
between the
eis (D).
!8=8a
8D''8a
!8'!'='!8'8D
8a)'!7;
!MZ;
the debye (D).=8a!='D!C==8a!=
(!HD''
'! (!8=
+!E;
*! )!H;+!E;*!
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!HD''8a!8=8a
!8'(=!MZ;
!MZ;
'!8''8a
)!7;)(!H;
'!MZ;
MZD)!HD(
I]ZhZbdaZXjaZhVgZcdc"edaVg!WZXVjhZ
I]ZhZbdaZXjaZhVgZcdc"edaVg!WZXVjhZ
i]ZhnbbZignd[i]ZbdaZXjaZhXVjhZh
i]ZhnbbZignd[i]ZbdaZXjaZhXVjhZh
i]Z^cY^k^YjVaY^edaZbdbZcihd[i]Z
i]Z^cY^k^YjVaY^edaZbdbZcihd[i]Z
WdcYhidXVcXZadji#
WdcYhidXVcXZadji#
This will also be discussed on pages
487–491. DIPOLE MOMENTS
IVWaZ(#&' HdbZedaVgVcYcdc"edaVgbdaZXjaZh#
IVWaZ(#&' HdbZedaVgVcYcdc"edaVgbdaZXjaZh#
jghZa[
g atoms
in order
of in order of=A 15 Sort=Athe
polar and
he
following
atoms
15following
Sort themolecules
following into
molecules
into polar and
allest
fi
rst):
non-polar
molecules.
For
the
polar
molecule
gativity (smallest first):
non-polar
=
= molecules.
= For the polar molecule
8aEÄ
Br H NaBr
draw diagrams
dipoles.the dipoles.
Na
drawshowing
diagramstheshowing
SF
PCl
SF
E 6
EXeF
E2 5
4 4
XeF
SFE45
PCl
SF2
XeF46
ecules
the following
polar from
molecules
from the following8XeF
8
8
8
ClF
BrF
SOCl
SF
6
ClF35
BrF3 2
SOCl2
SF65
lecules,
draw diagrams
he
e polar molecules,
draw diagrams
i^kZh
EÄ
CF
SCl4 2
C
2Cl32
BCl
8a
CF4
C2Cl2
EÄ
EÄ
8a
Y_i"&!'"Y^X]adgdZi]ZcZ
edaVg
8a
12
>ciZgbdaZXjaVg[dgXZh
>ciZgbdaZXjaVg[dgXZh
=
jhWdi"&!'"Y^X]adgdZi]ZcZ
cdc"edaVg
;^\jgZ(#., 9_i"VcYjhWdi"&!'"
(#- >ciZgbdaZXjaVg[dgXZh
Y^X]adgdZi]ZcZ#
(#-
>ciZgbdaZXjaVg[dgXZh
lecular
ow
w intermolecular
7g;
8a7g
Br2 is the only
boiling point as
therefore, stronge
=A 8dbeVg^hdcd[
8=(8a(!H8a'!MZD
MZD
(
)!HD((
8=(8a(!E8a
!MZD'(!ED8a
!E8a(!HD8a
(!H8a('!HD8a
'!ED8a
he dipoles.
SCl
PH3
HCN
PH23
O
BCl
OCl32
O3 3
CH2Cl2
8a;
Cis-1,2-dichloro
isomers – they d
C=C (Figure 3.9
8dbedjcY
Y_i"&!'"Y^X]adgdZi
jhWdi"&!'"Y^X]adgd
Both molecul
except that one
the boiling poin
interactions in th
=nYgd\ZcW
Br2 is a liquid
temperature.
It consists of
Br2atis room
a liquid
at room temperature.
It discrete
consists molecules
of discrete in
molecules in
ow
w physical
which
the
two
Br
atoms
are
joined
by
a
covalent
bond
(an
intramolecular
which the two Br atoms are joined by a covalent bond (an intramolecular
lecular molecular
ovalent
BILALforce).
HAMEED
INTERMOLECULAR
FORCES
to if
beBr
a 2liquid,
there
must there
be some
forces
between
But
if Br2 isBut
is to be
a liquid,
must
be some
forces
between
force).
end on the
molecules holding
them
in thethem
liquid
(otherwise
it would beitawould
gas). be a gas).
molecules
holding
in state
the liquid
state (otherwise
forces
These forces
are intermolecular
forces (Figure
3.91).
These
forces are intermolecular
forces
(Figure 3.91).
ints
of points of
boiling
Hydrogen bondi
responsible for ic
255
SKILL CHECK 1
Decide whether the following bonds are polar or non-polar. if the
bond is polar, state which is the δ+ atom, and explain whether or not
the molecule is polar:
A C-O as in CO2
B C-I as in CH3I
13
SKILL CHECK 2
Predict which of the following bonds are polar, and, if polar, in which
direction the electrons are pulled:
a. O—S
b. Cl—Cl
c. C—N
d. I—Cl
Predict whether each of the following molecules is polar:
a. BCl3
b. HCl
c. NH3
d. SiCl4
14
BILAL HAMEED
INTERMOLECULAR FORCES
256
SKILL CHECK 3
Predict whether each of the following molecules is polar:
a. CO2
b. BrCl
c. SCl2
d. CS2
Using the shapes drawn on slide 48 to 51, predict whether each of the
following molecules is polar:
a. CF4
b. Cl2O
c. PF3
d. NCl3
e. SiCl4
f. SO3
15
INTERMOLECULAR FORCES
Intermolecular forces are weak attractive forces between molecules.
These forces determine such properties as the solubility of one substance in
another and the freezing and boiling points of liquids.
Without intermolecular forces there could be no molecular liquids or solids.
16
BILAL HAMEED
INTERMOLECULAR FORCES
257
INTERMOLECULAR FORCES
Weak intermolecular forces arise from electrostatic attractions between
dipoles, including attractions between:
• Molecules with permanent dipoles such as hydrogen chloride
• A permanent dipole in one molecule and a dipole induced in a
neighbouring molecule, such as the attraction between iodine and
water
• Temporary dipoles are created fleetingly in non-polar atoms or
molecules
17
VAN DER WAALS’ FORCES DUE TO PERMANENT DIPOLES
Van der Waal’s forces due to permanent dipoles are interactions between
polar molecules - the positive end of one molecule attracts the negative
end of a neighbouring molecule.
H Cl
H Cl
H Cl
δ+
H Cl
δ+
δ+
δ+
δ-
δ-
C
δ+
δ-
δ-
Cl
H
δ-
δ-
Cl
Cl δCl
δ-
C
H
δ+
Cl
Cl δCl
C
H
δ+
δ-
δ-
van der Waal’s forces
Cl δCl
δ-
18
BILAL HAMEED
INTERMOLECULAR FORCES
258
VAN DER WAALS’ FORCES DUE TO PERMANENT DIPOLES
DÄ
8>
DÄ
=
DÄ
=
8>
= DÄ D
8> =
8>
D
D
D
8>
8>
D
l(l) but,
molecular
–
ttractions.
ar
der Waals’
n der
8>
DÄ
8>
DÄ
=
DÄ
=
8>
D
=
DÄ
=
8>
D
he H–Cl
DÄ D
DÄ D
D
8> =
=
=
DÄ
D
;^\jgZ(#.+ EZgbVcZciY^edaZÄeZgbVcZci
Y^edaZ^ciZgVXi^dchZm^hiWZilZZcbdaZXjaZh#
I]ZhZVgZh]dlc^cejgeaZ#
19
All other things being equal:
this basically means that we should
VAN DER WAALS’
FORCES
IN NON-POLAR
MOLECULES
compare
compounds
with relative
ole
molecular
as molecules.
similar Van
as der Waal’s forces
also existmasses
in non-polar
ling pointsIntermolecular forces
due to induced dipoles are the intermolecular attraction resulting from the
possible.
uneven distribution of electrons and the creation of temporary dipoles.
Because electrons move quickly in
ecular mass
their position is constantly
be similar.orbitals,
changing; at If
anymolecules
given instant they
with similar relative
ractions ascould be anywhere in an atom.
molecular masses are compared,
moleculesThe possibility will exist that one side
polar molecules have higher
es and will have more electrons than the other.
melting and boiling points than
This will give rise to a dipole.
ust be
non-polar molecules.
20
ci$•8
'
BILAL HAMEED
INTERMOLECULAR FORCES
259
VAN DER WAALS’ FORCES DUE TO INDUCED DIPOLES
Consider liquid argon. The electrons in an atom are in constant motion, and at any one
time the electrons will not be symmetrically distributed about the nucleus.
This results in a temporary (instantaneous) dipole in the atom, which will induce an
opposite dipole in a neighbouring atom.
These dipoles will attract each other so that there is an attractive force between atoms.
Although the dipoles are constantly disappearing and
reappearing, the overall force between the argon
atoms is always attractive, because a dipole always
induces an opposite one.
21
er Waals’ forces
he relative
ncreases.
molecular mass also increases, and we normally talk about ‘an increase in the
VAN
DER WAALS’ FORCES DUE TO INDUCED DIPOLES
strength of van der Waals’ forces as the relative molecular mass increases’.
A clear correlation between boiling point and relative molecular/
In general, van der Waals’ forces get stronger as the number of electrons in a
atomic mass can be seen down group 7 (blue line) and down group 0 (red
molecule
increases.
line) in Figure
3.94. As the number of electrons increases, the relative molecular
mass also increases, resulting in an increase in the strength of van der Waals’ forces.
'*%
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22
;^\jgZ(#.) KVg^Vi^dc^cWd^a^c\ed^cihd[ZaZbZcih^c\gdjeh,VcY%#
Consider group 7: fluorine (Mr 38.00) is a gas at room temperature,
but iodine (Mr 253.80) is a solid. This is because there are more electrons
in an iodine molecule, and the atoms making up the molecule are larger.
More electrons means that the temporary dipoles will be larger with INTERMOLECULAR
more
BILAL HAMEED
FORCES
electrons moving around. The larger atoms in the molecule means that
the outer electrons will be less strongly held, and therefore the induced
In addition to the relatively weak forces described above, ionic and covalent bonds
(discussed in detail in Chapter 5) are strong forces that greatly affect the melting point
of a compound:
260 a major source of
! In ionic solids, ionic bonds (electrostatic interactions) provide
attraction between particles. These types of solids tend to be hard but brittle (the
ionic lattice can crack) and have very high melting points. Lattice energy is a measure
of the strength of the interactions between ions in the lattice of an ionic solid. The
larger the lattice energy, the stronger the ion-ion interactions.
! In covalent solids, particles are bound to each other within strong networks of covalent
bonds. These solids are often exceptionally hard and have very high melting points.
VAN DER WAALS’ FORCES DUE TO INDUCED DIPOLES
e 11-2: a)
b)
Attraction
−
Inter+
δ+ −
ecular
eδe− δ+
δ−
ctions
2+
2+
δ+
δ+
−
−
nclude
e
e
+
−
dipolevan der Waal’s forces
−
inter+
dipoles
due to induced dipoles
induced
ctions,
ondon
−
+
ersion
+
−
s, and
drogen
Attraction
δ− δ+
δ− δ+
bonds.
Repulsion
c)
δ+
δ+
δ−
δ+
δ−
Electrons
b.p. / oC
δ+
NOBLE GASES
hydrogen bond
He
2
-269
Ne
10
-246
Ar
18
-186
Kr
36
-152
ALKANES
CH4
10
-161
C2H6
18
-88
C3H8
26
-42
oving Through States with Phase Diagrams
23
The previous sections described how microscopic interactions between particles can lead to
large-scale differences in the properties of a sample, especially by causing the sample to be
in different states (solid, liquid, or gas). This section describes some tools you can use to
track the state of a sample as it moves through different regions of temperature, pressure, or
added heat energy.
Phases and phase diagrams
Each state (solid, liquid, gas) is called a phase. When matter moves from one phase to another
due to changes in temperature and/or pressure, that matter is said to undergo a phase transition. The way a particular substance moves through states as temperature and pressure vary is
summarized by a phase diagram. Phase diagrams usually display pressure on the vertical axis
and temperature on the horizontal axis. Lines drawn within the temperature-pressure field of
the diagram represent the equilibrium boundaries between phases. A representative phase diaBOILING
HYDRIDES
OF
14
gram is shown
in FigurePOINTS
11-3. ReferOF
back
to this figure
as GROUP
you read through
the section.
EXAMPLES OF VAN DER WAAL’S
BOILING POINT / C°
Moving from liquid 100
to gas is called boiling, and the temperature at which boiling occurs is
called the boiling point. The normal boiling point is when this transition occurs at 1 atmosphere of pressure. Moving from
solid topoints
liquid is
melting,increase
and the temperature
at which
The boiling
ofcalled
the hydrides
with increasing
number of
melting occurs is called the melting point. The melting point temperature is the same as the
electrons. CH4 has the lowest boiling point as it is the smallest molecule.
freezing point temperature, but freezing implies matter moving from liquid to solid phase.
The melting point a substance has at 1atm pressure is called the normal melting point. Just
as freezing and melting points are the same, condensation points and boiling points are the
0
Mr
same temperature. For this reason published
tables are of100
freezing points
140 and boiling points.
50
PbH4
SiH4
GeH4 Larger molecules have more electrons and
therefore have greater intermolecular forces and
therefore higher boiling points.
-160
CH4
24
BILAL HAMEED
INTERMOLECULAR FORCES
261
I]Zdg^\^cd[]nYgd\ZcWdcY^c\
<gdje+]n
Let us compare the boiling points of the hydrides of group 6 elements:
HYDRIDES OF GROUP 16
='D
='H
&'%
&%%
7d^a^c\ed^ci$•8
='HZ
='D
='IZ
-%
+%
)%
D
'%
=
%
'%
Ä'%
)%
+%
-%
&%%
GZaVi^kZbdaZXjaVgbVhh
Ä)%
Ä+%
Ä-%
&'%
='IZ &)%
='HZ
='H
%
25
The boiling point increases from H2S to H2Te, as would be expected
from the increase in the strength of van der Waals’ forces as the relative
molecular mass increases.
H2O has, however, a much higher boiling point than would be
expected from its relative molecular mass. This suggests that there must
be intermolecular forces other than van der Waals’ forces between H2O
molecules and that these intermolecular forces are stronger than van der
HYDROGEN
Waals’ forces. These extra
intermolecularBONDING
forces between H2O molecules
are called hydrogen bonds. A hydrogen bond between two water molecules
HO
is shown in Figure 3.98.
100
2
BOILING POINT / C°
The higher than
It is important to realise that, although hydrogen bonding is the
expected boiling
strongest of the intermolecularHF forces, it is much weaker than
points
of NHbonding.
3, H2O
covalent
GROUP VI
GROUP V
and HF are due to
HI
NH3
HS
GROUP VII
intermolecular
PbH4
HCl
Hydrogen bonding occurs between molecules
when
a very
HBr
HYDROGEN
IV
electronegative atom (N, O, F) is joined to a hydrogen atomGROUP
in the
PH
BONDING
GeH
molecule. The electronegative atom withdraws electron
density from the
SiH
H, polarising the bond such that there is a strong interaction between the
E+ H and the E− atom (N, O,CHF) on the
other molecule.
26
The hydrogen bonding between ammonia molecules and hydrogen
fluoride molecules is shown in Figure 3.99.
0
100
50
140
2
3
4
4
-160
4
DÄ
D
=
D
=
C
=
D
DÄ
BILAL HAMEED
]nYgd\Zc
WdcY
Mr
;^\jgZ(#.- 6
lViZgbdaZXj
YVh]ZYejgeaZ
d[VEdmn\Z
]nYgd\Zcdc
:miZch^dc
There is mo
than just an
between dip
to a very ele
the electron
strongly from
an interactio
density of th
in one mole
in a differen
a directiona
bonds. If the
dipole–dipo
electronegat
expected to
bonding, bu
be explained
more diffuse
interacting e
INTERMOLECULAR FORCES
D
=
;
DÄ D
=
;
DÄ
The requ
are that
expected from its relative molecular mass. This suggests that there must
be intermolecular forces other than van der Waals’ forces between H2O
molecules and that these intermolecular forces are stronger than van der
Waals’ forces.
These extra intermolecular forces between H2O molecules
262
are called hydrogen bonds. A hydrogen bond between two water molecules
is shown in Figure 3.98.
It is important to realise that, although hydrogen bonding is the
HYDROGEN
BONDING
strongest of the
intermolecular forces, it is much weaker than
covalent bonding.
Strong hydrogen bonds can form with a hydrogen atom that is covalently bonded to
very electronegative atoms in the upper-right part of the periodic table: nitrogen,
Hydrogen bonding occurs between molecules when a very
oxygen, and fluorine.
electronegative atom (N, O, F) is joined to a hydrogen atom in the
elements:
7d^a^c\ed^ci$•8
molecule. The electronegative atom withdraws electron density from the
The hydrogen bonding between ammonia molecules and hydrogen
='HZ
)&
The
partially positive hydrogen
atom
polar molecules
be 3.99.
attracted to the
fluorideofmolecules
is shown incan
Figure
='IZ
)
unshared pairs of electrons of neighbouring molecules.
DÄ
D
=
D
DÄ
D
C
=
=D
D
=
DÄ
]nYgd\Zc
WdcY
D
=
D
=
;
DÄ D
=
;
DÄ
=D
='IZ &)%
=
]nYgd\Zc
WdcY
C D ÄD
=
&'%
D
D
=
D
=
D
;^\jgZ(#.- 6]nYgd\ZcWdcYWZilZZcild
;^\jgZ(#.. =nYgd\ZcWdcY^c\WZilZZcVbbdc^VVcY]nYgd\ZcÓ
jdg^YZbdaZXjaZh#
27
lViZgbdaZXjaZh#I]Z]nYgd\ZcWdcY^hi]Z
YVh]ZYejgeaZa^cZWZilZZci]ZadcZeV^g
d[VEdmn\ZcdcdcZbdaZXjaZVcYi]ZE ]nYgd\Zcdci]Zdi]ZgbdaZXjaZ#
expected
he relative
ld be
here must
ween H2O
han van der
O molecules
ater molecules
is the
han
y
in the
ity from the
between the
hydrogen
;
<gdje+]nYg^YZ
When
bonds to
an atom of N, O, or F, the hydrogen atom has a
='D a hydrogen atom&%%
H, polarising the bond such that there is a strong interaction between the
large,
E+ H and the E− atom (N, O, F) on the other molecule.
='H partially positive charge.
+%
:miZch^dc
There is more to hydrogen bonding
than just an electrostatic interaction
between dipoles. When H is attached
to a very electronegative atom,
the electron density is pulled away
strongly
from thethat
H. There
is then bonds are such strong forces is because the
One
reason
hydrogen
an interaction between the electron
hydrogen
ison
small
density of theatom
lone pair
N, O and
or F has only one electron.
in one molecule and the H nucleus
When
thatmolecule.
electron
is pulled
in a different
There
is thus away by a highly electronegative atom, there
a directional component to hydrogen
are
no more electrons under it. Thus, the single proton of the hydrogen
bonds. If the interaction were purely
dipole–dipole,
Cl, with theexposed.
same
nucleus
is partially
electronegativity as N, would also be
expected
to participate
in hydrogen
As
a result,
hydrogen’s
proton is strongly attracted to the lone pair of
bonding, but it does not. This could
electrons
other
molecules.
be explained of
by the
higher
energy and The combination of the large
more diffuse lone pair on the Cl not
electronegative
difference (high polarity) and hydrogen’s small size
interacting effectively with the H.
HYDROGEN BONDING
accounts for the strength of the hydrogen bond.
DÄ
g^YZbdaZXjaZh#
28
The requirements for H bonding
are that the H atom is attached to
a very electronegative atom – N,
O, F – which possesses at least
one lone pair of electrons.
BILAL HAMEED
(7DC9>C<
&',
INTERMOLECULAR FORCES
There is mo
than just an
between dip
to a very ele
the electron
strongly fro
an interactio
density of th
in one mole
in a differen
a directiona
bonds. If th
dipole–dipo
electronega
expected to
bonding, bu
be explaine
more diffus
interacting
The requ
are that
a very el
O, F – w
one lone
There is no hydrogen bonding between molecules of CH3CH2F, as the
H is not joined directly to an F atom:
=
=
=
=
8
8
=
=
263
;
CH3OCH3 is an ether with the structure shown in Figure 3.100. It is a
polar molecule, but there is no hydrogen bonding between molecules.
The hydrogen bonding between some molecules of ethanol is shown in
=
Figure
3.101.
=
The cis form has a lower melting point,8=
because
as well as intermolecular
^cigVbdaZXjaVg
=A
'8=(
]nYgd\Zc
DÄ
hydrogen bonding it is also able to participate
in
intramolecular
hydrogen
DÄ
D
WdcY
D D
DÄ
higjXijgZ[dg
D
D
=
bonding (Figure 3.102). This means that =
there is =
less hydrogen bonding
D DÄ
D
8='8=(
DÄ
between molecules andethanol
that the intermolecular forces are weaker
than in
D
D
D Ä :i]Vcda^hhdajWaZ^clViZg
D
Substanceshydrogen
that are able to
8
8 DÄ
=
the trans form. The trans form participates=only in intermolecular
8='8=
DÄ
n influence
Ethanol (C2HD
(
5OH) is very soluble in water, because ethan
participate
in
hydrogen
bonding
bonding, as the COOH groups are further
8 8
D D away from each other; the
ances in
hydrogen
bond
to the water (Figure 3.105). The hydrogen
will generally be soluble in water,
intermolecular forces are therefore
8=(8=' stronger
= in the trans form.
=
=
cules that are
between
water
and
ethanol molecules in the solution rele
as they are able to hydrogen bond
;^\jgZ(#&%& =nYgd\ZcWdcY^c\^cZi]Vcda#
;^\jgZ(#&%' >cigVbdaZXjaVg]nYgd\Zc
d are soluble
pays
back
the
energy
to break theacid
hydrogen bonds in pure
to the water.
but-2-ene-1,4-dioic
BZai^c\VcYWd^a^c\ed^cih
WdcY^c\^cY_i"Wji"'"ZcZ"&!)"Y^d^XVX^Y#
discussed
ethanol.
=
EXAMPLES OF HYDROGEN BONDING
D
=A =nYgd\ZcWdcY^c\XVcVahddXXjgl^i]^cbdaZXjaZh
EÄ molecular
Only
intermolecular forces are broken when covalent
D notE broken. E
Consider
cis andortrans
forms
butenedioic
E acid:
substances
arethe
melted
boiled
– of
covalent
bonds
are
=
E
Adc\ZgX]V^cVaXd]dahWZXdbZegd\gZhh^kZanaZhh
lViZg
=
=
EÄ
Octan-1-ol is insoluble in water. Although there is some h
D
ethanol
inuence
water
jhWdi"Wji"'"ZcZ"&!)"Y^d^XVX^Y
We haveY_i"Wji"'"ZcZ"&!)"Y^d^XVX^Y
looked at several factors
that infl
the strength of E
The stronger
the intermolecular
bonding
between the O–H group of the alcohol and the
=
8=
8=
'
(
bVaZ^XVX^Y
[jbVg^XVX^Y
intermolecular
forces and hence the melting and
boiling points
of
EÄ
forces, the more
energy
must be chain prevents water molecules on
the
long
hydrocarbon
D E
covalent substances.
D
D Now let
D us consider all the aboveE points together
supplied to hydrogen
break them
and theto each other (Figure 3.106). Energy is
bonding
=
=
with some examples.
higher the boiling
point.
break the
hydrogen bonds between the water molecules, b
=
8
=
;^\jgZ(#&%*
=nYgd\ZcWdcY^c\WZilZZc
8
8
D
paid
back
as only van der Waals’ forces form between the
lViZgVcYZi]Vcda#
8
=
=
=
=
29
Intermolecular
forces increase
in strength
as:
D
D < hydrogen bonding
van der Waals’ forces < permanent dipole–dipole
WEAKEST
~~~~~~~~~~~~~~~n
STRONGEST
bZai^c\ed^ci/&(%•8
bZai^c\ed^ci/'-,•8
DÄ
=
D
=
D
D
D
DÄ
D
=
^dc"Y^edaZ
Compare the
boiling
points of
of sulfur,
chlorine
and argon.
The
amount
energy
required
to breakD Ä the van der Waals’ forces in pure
hexane and
^ciZgVXi^dc
D
D
=
DÄ
D
=
D
=
=
DÄ
D
=
=
DÄ
D D=
Generally a substance will dissolve in a solvent if the intermolecular
D Ä forces in the solute
D in water.
D
D not
and solvent are similar. E.g. pentane is
but
CVreadily soluble in hexane
Ldg`ZYZmVbeaZh
^dc"Y^edaZ
^ciZgVXi^dc
D
=
DÄ
=
=
D
D
D
=
IVWaZ(#&( 7d^a^c\ed^cih^chjWhiVcXZhl^i]Y^[[ZgZci^ciZgbdaZXjaVg[dgXZh#
D
D Ä
D
=
D
D
D
=
DÄ
= =
= =
= =
= =
D
SOLUBILITY
8
D
=
DÄ
kVcYZgLVVahÉ[dgXZh
iZbedgVgnY^edaZh
D
=
=
eZgbVcZciY^edaZÄY^edaZ
D
]nYgd\ZcWdcY^c\
D
D
cdc"edaVghjWhiVcXZh
=
DÄ
edaVghjWhiVcXZh
Higdc\Zhi^ciZgbdaZXjaVg[dgXZ
YZXgZVh^c\
Wd^a^c\ed^ci
hjWhiVcXZhXdciV^c^c\=WdcYZY
idD!C!;
D
D
D
Relative molecular mass –
=
generally substances with higher D Ä
D
D
D
relative molecular masses have
=
DÄ =
D
higher melting points and
boiling
= D
= = = =
D
= = = =
D Ä van der
points due to stronger
=
D D=
8 8 8
Waals’ forces. D
=
8 8 8 8
=
IneZd[hjWhiVcXZ
DÄ
D
If relative molecular masses are approximately equal, then the boiling
points usually go in the order shown in Table 3.13.
and the hydrocarbon part of the molecule.
=
=
D
D
D
8
D
8
D
8
D
8
D Ä =
D
=
D
=
=
=
EÄ
D
D
D
D
pure pentane is paid back when van
der
Waals’
forces are formedÄbetween the
;^\jgZ(#&%+ I]Z]nYgdXVgWdcX]V^c^cdXiVc"&"daegZkZcih]nYgd\
8>
WZilZZclViZgbdaZXjaZhdcZ^i]Zgh^YZ#
= therefore=the strongest intermolecular
Sulfur, chlorine and argon are all non-polar substances, and
forces
are van
der Waals’ forces.
Any difference
betweenand
thesepentane
substances is thus due to the strength of van der Waals’ forces,
molecules
of hexane
W
V
9^hhdak^c\^dc^XhjWhiVcXZh
with a
which is affected by relative molecular mass (relative atomic mass for argon). Sulfur forms S8 molecules
DÄ
atomic
of
relative molecular mass of 256.48; Cl2 has a relative molecular mass of 70.90; and Ar has a relative
Consider
themass
dissolving
of sodium chloride:
D
D
39.95. Therefore the boilingCV
point of sulfur would be expected
toDbe highest,
as S8 has a greater relative molecular
=
=
‘+ water’
^dc"Y^edaZ
is thus required to
break ~
the
mass and therefore stronger van der Waals’
forces than Cl2 and Ar. More energy
NaCl(s)
~
~n NaCl(aq)
^dc"Y^edaZ
^ciZgVXi^dc
DÄ
^ciZgVXi^dc
intermolecular forces in sulfur than to break the intermolecular forces in chlorine
and the interatomic forces in Ar.
The aqueous solution contains aqueous ions (Na+(aq) and
D
D
Dto have
Chlorine would be expected
a higher boiling point than 8>
Ar,
Ä again due to the greater mass and the stronger
=
=
Ion–dipole interactions form between the water and the i
van der Waals’ forces.
3.107). Energy is required to break the electrostatic forces
W
V
lattice but energy is released when hydrated ions are form
;^\jgZ(#&%, >dcÄY^edaZ^ciZgVXi^dch
released when
the hydrated&'.
ions are formed is comparable
VWZilZZclViZgVcYhdY^jb^dch0VcY
(7DC9>C<
WWZilZZclViZgVcYX]adg^YZ^dch#
required to break apart the lattice then the substance is ge
30
&('
BILAL HAMEED
INTERMOLECULAR FORCES
264
SOLUBILITY
Pentane does not dissolve in water because there is hydrogen bonding
between water molecules.
If pentane were to dissolve in water there would be van der Waals’ forces
between water molecules and pentane. The energy released if van der Waals’
forces were to form between water molecules and pentane molecules would not
pay back the energy required to break the hydrogen bonds between water
molecules, as hydrogen bonds are stronger than van der Waals’ forces
31
SOLUBILITY
Ethanol (C2H5OH) is very soluble in water, because ethanol is able to form
hydrogen bonds with the water.
:i]Vcda^hhdajWaZ^clViZg
Substances that are able to
(C2H5OH) is very soluble in water, beca
participate
hydrogen bonding
The hydrogen bonding between water
andinethanol
moleculesEthanol
in the
hydrogen
bond to the water (Figure 3.105). The
will generally be soluble in water,
solution releases energy and pays back
the
energy
to
break
the
hydrogen
between
water
and ethanol molecules in the solu
as they are able to hydrogen bond
pays
back
the
energy
to break the hydrogen bond
to the water.
bonds in pure water and pure ethanol.
ethanol.
E
=
;^\jgZ(#&%* =nYgd\ZcWdcY^c\WZilZZc
lViZgVcYZi]Vcda#
32
DÄ
D
D
BILAL HAMEED
=
DÄ
D
=
D
=
D
=
D
DÄ
D
D
=
D
=
DÄ
8='8=(
Octan-1-ol is insoluble in water. Although there
bonding between the O–H group of the alcohol
the long hydrocarbon chain prevents water molec
hydrogen bonding to each other (Figure 3.106).
break the hydrogen bonds between the water mo
paid back as only van der Waals’ forces form betw
and the hydrocarbon part of the molecule.
D
E
=
=
=
D
Adc\ZgX]V^cVaXd]dahWZXdbZegd\gZhh^kZ
lViZg
D
EÄ
D
=
E
=
E
EÄ
D
D
ethanol in water
=
E
DÄ
=
EÄ
E
=
EÄ
D
E
= =
D
= = = =
=
INTERMOLECULAR
D D= FORCES
D
8
=
8 8 8 8
D 8
=
= = = = =
DÄ
EÄ
=
E
E
=
EÄ
=
EÄ
E
=
Adc\ZgX]V^cVaXd]dahWZXdbZegd\gZhh^kZanaZhhhdajWaZ^c
lViZg
E
=
Octan-1-ol is insoluble in water. Although there is some hydrogen
bonding265
between the O–H group of the alcohol and the water molecules,
the long hydrocarbon chain prevents water molecules on either side from
hydrogen bonding to each other (Figure 3.106). Energy is needed to
break the hydrogen bonds between the water molecules, but this is not
paid back as only van der Waals’ forces form between the water molecules
and the hydrocarbon part of the molecule.
D
E
D
=
D
8='8=(
E
=
;^\jgZ(#&%* =nYgd\ZcWdcY^c\WZilZZc
lViZgVcYZi]Vcda#
SOLUBILITY
DÄ
D
D =
Octan-1-ol is insoluble in water. Although
D
=
there is some hydrogen bonding between
DÄ
D
D D
the O–H group of the alcohol and the
=
DÄ =
D
D
=
= = = = cdc"edaVg
water molecules, the long hydrocarbon
]nYgdXVgWdc
D
= =
=
=
DÄ
=
X]V^c
D
chain prevents water molecules on either D D =
8 8 8 =
=
8 8 8 8
side from hydrogen bonding to each
D 8
=
other.
= = = = = =
D
DÄ
=
D
D Ä =
D
=
D
D
=
D
D
D
DÄ
=
D
D
=
=
D
DÄ
=
D
D
=
=
D
DÄ
D
DÄ
D
=
=
D
D
=
DÄ
D
D =
DÄ
=
D
=
D
D
D
D
^dc"Y^edaZ
^ciZgVXi^dc
D
=
9^hhdak^c\^dc^XhjWhiVcXZh
Consider the dissolving of sodium chloride:
D
^dc"Y^edaZ
^ciZgVXi^dc
8>Ä
=
DÄ
D
=
=
D
D
DÄ
D
=
^dc"Y^edaZ
^ciZgVXi^dc
D
D
=
=
DÄ
W
DÄ
D
=
D
D
D
=
;^\jgZ(#&%+ I]Z]nYgdXVgWdcX]V^c^cdXiVc"&"daegZkZcih]nYgd\ZcWdcY^c\
8>Ä 33
WZilZZclViZgbdaZXjaZhdcZ^i]Zgh^YZ#
D
=
V
CV
D
DÄ
D
D D=
D
D
=
= =
DÄ
=
Energy is needed to break the hydrogen
bonds between the water molecules, but
DÄ
this is not paid back as only van der Waals’
D
D
CV
forces form between
the water molecules
=
^dc"Y^edaZ
^ciZgVXi^dc
and the hydrocarbon
part of the molecule.
DÄ
D
D
DÄ
=
D
=
D
=
=
D
V
EÄ
D
E
W
;^\jgZ(#&%, >dcÄY^edaZ^ciZgVXi^dch
VWZilZZclViZgVcYhdY^jb^dch0VcY
WWZilZZclViZgVcYX]adg^YZ^dch#
‘+ water’
NaCl(s) ~
~
~n NaCl(aq)
The aqueous solution contains aqueous ions (Na+(aq) and Cl−(aq)).
Ion–dipole interactions form between the water and the ions (Figure
3.107). Energy is required to break the electrostatic forces in the ionic
lattice but energy is released when hydrated ions are formed. If the energy
released when the hydrated ions are formed is comparable to the energy
required to break apart the lattice then the substance is generally soluble.
SKILL CHECK 4
&(' Explain why, in comparison with the other group 6 hydrides, water has
an anomalous boiling temperature.
34
BILAL HAMEED
INTERMOLECULAR FORCES
266
SKILL CHECK 5
State all the forces operating between molecules of:
(a) ammonia, NH3
(b) methane, CH4
(c) oxygen fluoride, OF2
35
HYDROGEN BONDING IN ICE
• Each water molecule is hydrogen-bonded to 4
others in a tetrahedral formation
• Ice has a “diamond-like” structure
• Volume is larger than the liquid making it less
dense
• When ice melts, the structure collapses slightly
and the molecules come closer; they then move a
little further apart as they get more energy as they
warm up
This is why water has a maximum density at 4°C and
ice floats.
Hydrogen bonding
Lone pair
36
BILAL HAMEED
INTERMOLECULAR FORCES
PH3
–100
GeH4
H2O
Barking
Dog Art
SiH4
Group 14
100
Group 16
–150
267
2
3
50
4
boiling point/°C
CH4
–200
5
period
gure 3.54 Thymine, adenine, cytosine
nd guanine are bases in DNA. The hydrogen
onding between them allows the DNA
olecule to replicate (make new copies of
self) and also allows the genetic code to be
xpressed in the synthesis of proteins.
C
C
C
N
C
N
N
C
C
HYDROGEN BONDING IN DNA
thymine
eVgZci
higVcY
Group 15
C
IC
<
N
8
I
to DNA
chain
C
N6
C
SiH4
Group 14
N
<
I
<
<
6
<
8
8
I
;^\jgZ7(. I]ZYdjWaZ"]Za^mhigjXijgZ
d[9C6#
Remember that proteins are
made up of a specific sequence
of amino acids (called the
primary structure). This sequence
gives the protein its shape and
therefore its function, and it is
the sequence of nucleotide
bases within each gene that
dictates the primary structure
of the protein produced. In
other words, the information
for the structure of each protein
is carried in the sequence of
nucleotide bases within the
particular gene.
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DNA
higVcY
chain
I
8
N
C
C
I
<
6
<
C
C
<
<
6
<
and guanine are bases in DNA. The hydrogen
to DNA
bonding
between them allows the DNA
chainto replicate (make new copies of
molecule
itself) and also allows the genetic code to be
expressed in the synthesis of proteins.
N
6
6
8
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C
I
6
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8
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<
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cytosine
C
C
N
C
C
N
C
N
C
to DNA
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adenine
66
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protein. It is these proteins produced by the cell that carry out the
thousands of biochemical processes that are responsible for life.
For a cell to produce a protein, the gene must first undergo a process
called transcription. This occurs in the nucleus, and the first step in this
process is the unwinding and separation of the two strands of DNA on
which the gene is situated. In a similar manner to DNA replication, this
exposes the nucleotide bases of the gene and allows the bases on one of
the DNA strands to be used as a template.
Transcription differs to DNA replication, however, in that only one
strand of DNA is used as a template and the complementary strand
produced is a ribonucleic acid called messenger RNA (mRNA). The
complementary strand of mRNA is built up through complementary
RNA nucleotides (called ribonucleotides) forming base pairs with
the exposed bases of the DNA template. Guanine pairs with cytosine
and adenine pairs with uracil (not thymine, as in DNA). As each
ribonucleotide comes in and forms a base pair, it is joined covalently to
the growing mRNA chain by a phosphodiester link. This results in the
production of a strand of mRNA that has the complementary sequence
of bases to the gene of the DNA template strand. This mRNA then leaves
the nucleus and enters the cytoplasm, where it takes part in the second
process to produce a protein, known as translation.
SKILL CHECK 6
What is involved when a hydrogen bond is formed between two
molecules?
Translation is the process of
protein synthesis in which the
code held in the sequence of
bases of the mRNA is translated
into the sequence of amino acids
(primary structure) of the protein.
A a hydrogen atom bonded to an atom less electronegative than itself
B a lone pair of electrons
C an electrostatic attraction
between opposite charges
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38
BILAL HAMEED
N
to DNA
chain
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3
C
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O
to DNA
chain
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6
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2
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I
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–200
<
O
I
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6
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HC
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8
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66
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HF
INTERMOLECULAR FORCES
cytos
268
SKILL CHECK 7
Which types of intermolecular forces can exist between adjacent urea
molecules?
A hydrogen bonding
B permanent dipole-dipole forces
C temporary induced dipole-dipole forces
39
BILAL HAMEED
INTERMOLECULAR FORCES
269
INTERMOLECULAR
FORCES
3
SECTION
A
3
4
A slow stream of water from a tap can be deflected by an electrostatically charged plastic rod
because
is a from
polar amolecule.
4 1 A slow
streamwater
of water
tap can be deflected by an electrostatically charged plastic rod
because water is a polar molecule.
liquid
liquid
burette
burette
negatively charged rod
negatively charged rod
Why is a water molecule polar?
Why is a water molecule polar?
A
B
C
D
A Molecules are bonded together by hydrogen bonds.
Molecules are bonded together by hydrogen bonds.
B The oxygen and hydrogen atoms have different electronegativities.
The oxygen and hydrogen atoms have different electronegativities.
C The oxygen atom has two lone pairs of electrons.
The oxygen atom has two lone pairs of electrons.
D
is to
able
to dissociate
into ions.
WaterWater
is able
dissociate
into ions.
[S’02 Q4]
4
Why
does
copper
wire conduct
electricity
when
a potential
difference
is applied?
57 2 5
Why
does
copper
wire
conduct
electricity
when a
potential
difference
is applied?
When
heated,
solid
iodine
readily
forms iodine
vapour.
A Bonding
electrons
in
the crystal
lattice move.
AWhat
Bonding
electrons
in thesuggest
crystal
lattice
does this
information
about move.
the nature of the particles in these two physical states
of iodine?
II)move
ions move
the cathode.
B Copper(
II) ions
to the to
cathode.
B Copper(
C
C solid
The atoms
of copper
become
ionised.
vapour
The
atoms
of copper
become
ionised.
D ionic
The crystal
lattice
breaks
atomic
DA The
crystal
lattice
breaks
down.down.
B
6
8
9
ionic
molecular
3 of helium at 2 kPa pressure and flask Y contains
31of
dm3 of
6
Flask
X contains
dm
helium
at 2 kPa pressure and flask Y contains 2 dm3 of2 neon
at neon
1 kPaat 1 kPa
Flask
X contains
1 dm
C molecular
atomic
pressure.
pressure.
D molecular
molecular
[S’02 Q7]
If theIfflasks
are connected
at constant
temperature,
what iswhat
the final
pressure?
the flasks
are connected
at constant
temperature,
is the
final pressure?
Which statement about the standard enthalpy change of formation of carbon dioxide is correct?
$" kPa1 $" kPa D 2 kPa
A 1A!" kPa
D 2 kPa
1 !" kPa B 1B!# kPa1 !# kPa C 1 C
A It is equal to the standard enthalpy change of combustion of carbon.
B
It is equal to twice the bond energy of the C=O bond.
C
It is the energy released when one mole of carbon dioxide is formed from carbon at the
temperature of combustion of the carbon.
D
It is the same for carbon dioxide produced from graphite and from diamond.
Use of the Data Booklet is relevant to this question.
Hydrazine was used as a fuel for the Messerschmidt 163 rocket fighter in World War II and for the
CEDAR
COLLEGE
INTERMOLECULAR FORCES WS 1
American
Gemini and Apollo spacecraft. It has the following formula.
9701/1/M/J/02
9701/1/M/J/02
H
H
[Turn over
[Turn over
Which diagram shows the electronic configuration of this atom in its lowest energy state?
270
3
A
4 3 The African weaver ant defends its territory by spraying an intruder with a mixture of compounds.
The ease by which these compounds are detected by other ants depends upon the volatility,
which decreases as the strength of the intermolecular forces in the compound increases.
B
Which compound would be the most volatile?
A
C
B
C
D
D
CH3 CH2CH2CH2CH2CH3
CH3 CH2CH2CH2CH2CHO
CH3 CH2CH2CH2CH2NH2
CH3 CH2CH2CH2CH2OH
[W’02 Q4]
The gecko,
small lizard,
can has
climbnoup
a smoothdipole?
glass window. The gecko has millions of
56 4 Which
of the afollowing
molecules
permanent
microscopic hairs on its toes and each hair has thousands of pads at its tip. The result is that the
molecules in the pads are extremely close to the glass surface on which the gecko is climbing.
B CHCl3
C C2Cl4
D C2H5 Cl
A CCl2F2
What is the attraction between the gecko’s toe pads and the glass surface?
6
Which
is the most
likely shape of a molecule of hydrazine, N2H4 ?
A co-ordinate
bonds
B
C
A
D
covalent bonds
H
H
ionic bonds
90°
N
N
van der Waals’ forces
H
H
[W’04 Q6]
43
In which
which
reaction
does
theincarbon-containing
product
have a smaller bond angle than the organic
What
areprocess
the bondare
angles
the PH
likely
to be?
3 molecule
576 5 In
hydrogen
bonds
broken?
reactant?
H
H o
o
A 90
C 109 o
D 120 o
90° 104
A
H2(l) → H2(g) B
ethanolic sodium hydroxide
N with H
H
N refluxed
BA bromoethane
B NH3(l) → NH3(g)
B complete combustion of methane in air
C 2HI(g) → H2(g) + I2(g)
C methane and an excess of chlorine under ultraviolet light
H → C(g)120°
+ 4H(g)H
D CH4(g)
D polymerisation of ethene
[S’06 Q5]
N
N
C
6 Which ofHthe following leastHresembles an ideal gas?
7 6 A crystal of iodine produces a purple vapour when gently heated.
A
ammonia
Which
this process?
© UCLES
2004 pair of statements correctly describes
9701/01/O/N/04
helium
107° H
type of bond broken
C hydrogen
N
N H
D
covalent
H
D A trichloromethane
H
B
covalent
[Turn over
B
7
7
8
formula of purple species
I
I2
C diagram
induced
dipole-dipole
I2 of solids X and Y. [In X,
The
shows
part of the lattice structures
Which of the following exists in the solid state as a giant covalent lattice?
particles of different elements.]
D
permanent dipole-dipole
I2
A ice
represent
[S’09 1 Q7]
B iodine
Hydrogen peroxide slowly decomposes into water and oxygen. The enthalpy change of reaction
be calculated
using standard enthalpies of formation.
IV) oxide
Ccan silicon(
D
tin(IV) chloride
(hydrogen peroxide(l)) = –187.8 kJ mol–1
[Turn over
9701/1/O/N/02
(water(l))
= –285.8 kJ mol–1
CEDAR COLLEGE
Using a Hess cycle, what is the enthalpy change of reaction INTERMOLECULAR
for this decomposition?
FORCES WS 1
X 2H2O2(l) → 2H2O(l) + O2(g)
Y
be correct.
0
0
Use of the Data Booklet
be appropriate
13 14may15
16 17 for
18some questions.
271
proton number
13
14
15 16 17
proton number
18
1 7 Why is the boiling point of ammonia, NH3, higher than the boiling point of phosphine, PH3?
A
C polar; phosphine molecules are not.
Ammonia molecules are
D
B
Ammonia molecules have significant hydrogen bonding; phosphine molecules do not.
C N–H covalent bonds are stronger than P–H covalent bonds.
number of
number of
D There is one lone pair in each ammonia molecule but no lone pair in each phosphine
unpaired
unpaired
molecule.
4
electrons
electrons
[S’18 3 Q1]
5 8 Which molecule has the largest overall dipole?
2 Neutrons are passed through an electric field. The mass of one neutron relative to 121 the mass of
B electric field is recorded.
C
D
A and any deflection in the
a 12C atom
H
Cl
Cl
Which row
0 is correct?
17 18
O
C
C
C13 14 15 16
mass of
behaviour
proton
numberof beam of
Cl neutron
Cl
H
neutrons in an electric
field
O
Cl
0
C 13
14
Cl
15 O 16 C 17 O 18
proton number
[W’09 1 Q5]
A statement
0
5 9 Which
explains why deflected
the boiling point of methane is higher than that of neon?
6 [A
The
first
stage
in
the20]
industrial production of nitric acid from ammonia can be represented by the
: H, 1; C, 112; Ne,
rB
deflected
following equation.
A C A molecule
greater mass than a molecule of neon.
0 of methane has
not adeflected
4NO(g) + 6H2O(g)
4NH3(g) + 5O2(g)
1 of methane form
not hydrogen
deflected bonds, but those of neon do not.
B D Molecules
Using the following standard enthalpy change of formation data, what is the value of the standard
Centhalpy
Molecules
of methane
have
stronger intermolecular forces than those of neon.
reaction?
change,
∆Ho , for this
3 D
The The
tablemolecule
refers to of
themethane
electron is
distribution
the
shell
polar, butinthat
neon
is
not.of–1 an atom with eight protons.
3 ofsecond
o
[S’09 1 Q5]
compound
∆Hf / kJ mol
Which row is correct for this atom?
410 The boiling points of methane, ethane, propane and butane are given.
NH3(g)
–46.1
orbital shape
compound
orbitalCH
shape
NO(g)
3CH3
CH4
+90.3
CH3CH2CH3
CH3CH2CH2CH3
112of H2O(g)
185number–241.8
231
273
number
of
orbital type
electrons
electrons
–1
kJ mol
A
Which+905.2
statement
explains the increase in boiling point from methane to butane?
p
2
s 9701/01/M/J/09
4
© UCLESA2009
[Turn over
B
–105.4packing
kJ mol–1of molecules results in stronger van der Waals’ forces.
A Closer
B
p
4
s
2
C
–905.2
kJ mol–1
B More
covalent
bonds are present and therefore more energy is required to break the bonds.
C
s
2
p
4
–1
D
–1274.0
kJ
mol
C More electrons in the molecules results in stronger van der Waals’ forces.
D
s
4
p
2
D More hydrogen atoms in the molecules results in stronger hydrogen bonding.
[M’17 2 Q4]
7
Which conversion involves a reduction of chromium?
boiling point / K
orbital type
5
2−
The CrO
characteristic
smell of garlic is due to alliin.
A
4 → CrO3
H
CrO24 − → Cr2O72 −
B
© UCLES 2018
C
CrO2Cl2 → CrO24 −
D
CrO2Cl2 → Cr2O3
xC
H
C
O
y
9701/13/M/J/18
CH2
S
H
alliin
CH2
NH2
z
H
C
CO2H
What are the approximate bond angles x, y and z in a molecule of alliin?
x
y
z
A
90o
CEDAR COLLEGEo
B
120
90o
109o
109o
90o
120o
120o
109o
C
INTERMOLECULAR FORCES WS 1
Which statement about this reaction is incorrect?
14 Two students, P and Q,
were asked to draw bar charts
to represent
of the
substance
R
S how some properties
T
halogens
and
their
compounds
differ
in
magnitude.
Their
diagrams
are
shown.
For each
question
there
are
four
possible
answers,
A,
B,
C,
and
D.
Choose
the
one
you
consider
to
A Increased pressure gives a higher yield of SO3.
be correct.
mp / oC
801
2852
3550
student gives
P
student Q
B Increased temperature
a higher 272
yield of SO3.
o
bp / C
1413
3600
4827
C ability
In the of
forward
reaction
the oxidation
ofelectrons
sulfurochanges
+4 to
111 The
an atom
in a covalent
bond tostate
attract
to itselffrom
is called
its+6.
electronegativity.
electrical conductivity of solid
poor ∆H
poor
good
I asthe
Br
D greater
Vanadium(
) oxide isbetween
used
a catalyst.
The
the V
difference
electronegativities0 of theCltwo atoms
inI the bond, the more
Whatiscould
be the identities
of R, S and T ?
polar
the bond.
Br
solubility
8
Clthe most
When pair
making
sparkler
fireworks,
a mixture
barium the
nitrate
powder with aluminium powder,
Which
willRform
polar
bond
atoms?
S covalent
T ofbetween
water and glue is coated onto wires and allowed to dry. At this stage, the following exothermic
A
and
bromine NaCl
reaction
may
occur.
Achlorine
C [graphite]
0MgO
B Bchlorine MgO
and iodine
NaCl
+3(aq)
36H2O SiO
→ 23Ba(OH)
(OH)3 + 6NH3
16Al +of3Ba(NO
3)2NH
2 +H 16Al
solubility
AgX(s)
in
∆H o for
2(g) + X2(g) → 2HX(g)
C Cfluorine NaCl
and chlorine
MgO
C [graphite]
Which of
conditions
would
be bestare
tocorrect?
reduce the rate of this reaction during the drying process, and
Which
theand
student’s
diagrams
D
would
also keep
the aluminium
nitrate
unchanged?
Dfluorine
NaCliodine
MgOand barium SiO
2
[W’10 2 Q1]
A both P and Q
/ K represents
pH the Boltzmann distribution of molecular energies at two
temperature
2 B
Which
diagram
correctly
P only
11
In
which
change
would
only
van
have
12 temperatures T1 and T2, where T1 =der
300Waals’
K and forces
T2 = 310
K? to be overcome?
A
298
7
C Q only
A evaporation of ethanol A
C2H5OH(l) → C2H5OH(g)
B
14
D B neither P 298
nor Q
B melting of ice
H2O(s) → H2O(l)
T1
C
398
7
T1
T2
CD melting
solid
carbon
dioxide
CO
2(s) → CO2(l)
T2
15 When
iodineproportion
isof
heated,
a
vapour
is
produced.
proportion
398
14
of molecules
molecules
D solidification
of butane
(l) →ofCin
4H10
4Hthe
10(s)
Which
row of the table
correctly identifiesCthe
species
vapour and its colour?
913 Which molecular structure will have the smallest overall dipole?
species
0
0 colour
molecular
molecular
A
B
C
energy
energy
A
I(g)
brown
H3 C
© UCLES
B 2011
I(g)
C
CI2(g)
C
Cl
D H2 C
Cl
I2(g) H
HC
purple3
C
brown C
H3 C T2
purple
T1
Cl9701/11/M/J/11 Cl
C
Cl
C
Cl
H3 C
D
Cl
D
C
CH
T23
CH3 over
[Turn
C
H3 C
C
Cl
T1
[W’09 2 Q4]
proportion
proportion
of strengths
moleculesof the covalent bonds within
of molecules and the van der Waals’ forces
14 How do the
16
+ I2(g)
2HI(g), is 60 at 450 °C.
10
The equilibrium
constant, Kc, for the reaction H2(g) molecules,
between molecules, vary going down Group VII from chlorine to bromine to iodine?
What is the number of moles of hydrogen iodide in equilibrium with 2 mol of hydrogen and 0.3 mol
0
0
strength
molecular strength of
molecular
°C? of
of iodine at 450
covalent bonds energy
van der Waals’ forces
energy
1
1
AA
B
C 6
D 36
decrease
decrease
100
10
B
decrease
increase
C
increase
decrease
D
increase
increase
[S’13 3 Q16]
© UCLES 2014
© UCLES 2010
© UCLES 2013
CEDAR COLLEGE
9701/11/M/J/14
9701/12/O/N/10
9701/13/M/J/13
INTERMOLECULAR FORCES WS 1
1st
ionisation energy / kJ mol–1
2nd
3rd
4th
5th
1800
2700
4800
6000
273
4
What could be the formula of a chloride of X?
7 15 At room temperature and pressure, H2O is a liquid and H2S is a gas.
A XCl
B XCl 2
C XCl 3
D XCl 4
What is the reason for this difference?
2
950
6th
12 300
A O has
higher
first and
second
ionisation
energies
than S.
Which
set of
conditions
gives
the highest
yield
of ammonia
at equilibrium?
B
The covalent bond between O and H is stronger than the covalent
bond between S and H.
2NH3 (g)
∆H o = –92 kJ mol–1
N2 (g) + 3H2 (g)
There is significant hydrogen bonding between H2O molecules but not between H2S
molecules.
catalyst
pressure
temperature
C
D A The instantaneous
dipole-induced
dipole low
forces between H2O molecules are stronger than
absent
high
the instantaneous dipole-induced dipole forces
between H2S molecules.
3
[S’16 1 Q7]
B
absent
low
high
516 The presence of dipoles helps to explain why the element Br2 and the compound CHCl 3 exist as
C
present
high can be decomposed
high
8
Gaseous
phosphorus
pentachloride
into gaseous phosphorus trichloride and
liquids at room
temperature.
chlorine
by
heating.
The
table
gives
the
bond
energies.
D
present
low
low
Which types of dipole are involved?
bond
bond energy / kJ mol–1
Br2 is relevant to this question. CHCl 3
3 Use of the Data Booklet
P–Cl (in both chlorides)
330
induced dipoles and permanent induced dipoles and permanent
A compound S O is hydrolysed by water to produce sulfuric acid and oxygen only.
The
2 dipoles
7
dipoles
Cl –Cl
242
induced dipoles and permanent
induced dipoles
only
B
Which
of oxygen,
measured
at room temperature
pressure,
is evolved when 0.352 g
dipoles
What isvolume
the enthalpy
change
for the decomposition
of PCl 5 and
to PCl
3 and Cl 2?
of S2 O7 is hydrolysed?
induced dipoles and permanent
C
induced
dipoles only
–1
dipoles
A –418 kJ
3 mol
B 24 cm3
C 48 cm3
D 96 cm3
A 12 cm
–1
D
induced
dipoles
only
induced
dipoles
only
B –88 kJ mol
[W’11 2 Q5]
–1
C +88 kJNmol
417 Nitrogen,
2 , and carbon monoxide, CO, both have Mr = 28.
6 Three compounds
have the physical properties shown in the table.
D +418 kJ mol–1
The boiling point of N2 is 77 K.
compound
P
Q
R
K.
The
boiling point
of COwas
is 82prepared
9
An aqueous
solution
containing a mixture of 1.0 mol of AgNO3 and 1.0 mol of
melting point / °C
2852
993
–119
FeSO4 in 1.00 dm3 of water. When equilibrium was established, there was 0.44 mol of Ag+(aq) in
What
could
be
responsible
for
this
difference
in
boiling
points?
the mixture.
boiling point / °C
3600
1695
39
are not polar.
CO molecules
have a permanent
dipole,
+
2 molecules
conductivity
poorthe NAg(s)
poor
(aq) + Fe2+(aq)
+ Fe3+(aq) poor
Ag(solid)
B N2 has σ and
π bonding,
CO has σ bonding
conductivity
(liquid)
good only. good
poor
What is the numerical value of Kc?
strong N≡N (aqueous)
bond, CO hasinsoluble
a C=O bond. good
C N2 has a conductivity
insoluble
A 0.62
B 1.40
C 1.62
D 2.89
D The CO molecule has more electrons than the N2 molecule.
A
[W’15 1 Q4]
What might be the identities of P, Q and R?
10 The equation for the reaction between carbon monoxide and hydrogen is shown.
© UCLES 2015
9701/11/O/N/15
P
Q
CO(g) +R 3H2(g)
A
MgO
KCl
NH3
What are the units of Kp for this reaction?
B
MgO
NaF
C2H5 Br
C
A kPa
B kPa–1
C
SiO2
KCl
C2H5 Br
D
7
SiO2
NaF
kPa–2
2SO3 (g), what will change the value of Kp ?
9701/11/M/J/16
adding more O2
C COLLEGE
increasing the pressure
CEDAR
D
D
adding a catalyst
© UCLES 2016
B
kPa2
HCl
For the equilibrium 2SO2(g) + O2(g)
A
CH4(g) + H2O(g)
increasing the temperature
INTERMOLECULAR FORCES WS 1
4
Which pair has species with different shapes?
A
BeCl 2 and CO2
B
CH4 and NH4+
274
3
BF3 of methane, ethane, propane and butane are given.
3 and
418 C
The NH
boiling
points
D
5
SCl 2 and H2O
compound
CH4
CH3CH3
CH3CH2CH3
CH3CH2CH2CH3
boiling point / K
112
185
231
273
In which reaction does an element have the largest change in oxidation number?
Which statement
explains
the +increase in
from methane to butane?
2+
3+ boiling point
3+
2–
A
Cr2O7
A
Closer– packing of– molecules
results in stronger van der Waals’ forces.
–
+ 6Fe
+ 14H → 2Cr
+ 6Fe
+ 7H2O
B
3OCl
B
More
covalent bonds
are
present 3+
and therefore
more energy is required to break the bonds.
2+
+
2+
–
C
5Fe
C
More electrons in the molecules results in stronger van der Waals’ forces.
D
PbO2 + Sn2+ + 4H+ → Sn4+ + Pb2+ + 2H2O
D
More hydrogen atoms in the molecules results in stronger hydrogen bonding.
→ Cl O3 + 2Cl
+ MnO4 + 8H → 5Fe
+ Mn
+ 4H2O
[M’17 2 Q4]
619 Which statement can be explained by intermolecular hydrogen bonding?
5
The characteristic smell of garlic is due to alliin.
A
Ethanol has a higher boiling point than propane.
B
Hydrogen chloride has a higheryboiling point than silane,
SiH4.
z
C
Hydrogen iodide forms an acidic solution when dissolved in water.
D
Propanone has a higher boiling point than
3 propane.
H
xC
H
O
C
CH2
S
CH2
H
NH2
C
H
CO2H
alliin
[W’12 3 Q2]
420 The
points
of methane,
are given.
Whatboiling
are the
approximate
bondethane,
angles propane
x, y and and
z in butane
a molecule
of alliin?
compound
x
boiling pointo/ K
A
90
y CH4
112
90o
z
CH3CH3
CH3CH2CH3
CH3CH2CH2CH3
109o
185
231
273
B statement
120o explains
109o the increase
90o in boiling point from methane to butane?
Which
o
o
o
120
© UCLES
A C2015
Closer
packing of120
molecules 109
results in 9701/13/M/J/15
stronger van der Waals’ forces.
o
o
[Turn over
o
109 are present
109 and therefore more energy is required to break the bonds.
B D More 180
covalent bonds
6
C More electrons in the molecules results in stronger van der Waals’ forces.
Which gas sample contains the fewest molecules?
D More hydrogen atoms in the molecules results in stronger hydrogen bonding.
5
[M’17 2 Q4]
1.00 dm3 of carbon dioxide at 27 °C and 2.0 kPa
A
The 1.00
characteristic
smell ofat
garlic
is due
alliin.
dm3 of hydrogen
100 °C
andto2.0
kPa
B
1.00 dm3 of nitrogen at 300
H 4.0 kPa O
H °C and
y
kPa S
1.00 dm3 of oxygen at 250
x C°C and
C 3.0
CH
2
C
D
H
alliin
CH2
NH2
z
H
C
CO2H
What are the approximate bond angles x, y and z in a molecule of alliin?
x
y
z
A
90o
90o
109o
B
120o
109o
90o
C
120o
120o
109o
D
180o
109o
109o
© UCLES 2017
9701/12/F/M/17
CEDAR COLLEGE
6
Which gas sample contains the fewest molecules?
[Turn over
INTERMOLECULAR FORCES WS 1
1, 2 and 3
are
correct
2 and 3
only are
correct
1 and 2
only are
correct
1 only
is
correct
27514
No other combination of statements is used as a correct response.
SECTION
SectionB
B
31 For
Theeach
relative
molecular
mass
of asection,
molecule
72.three numbered statements 1 to 3 may
of the
questions
in this
oneoforchlorine
more ofisthe
be correct.
Which properties of the atoms in this molecule are the same?
Decide whether each of the statements is or is 14
not correct (you may find it helpful to put a tick against
1 statements
radius
the
that you consider to be correct).
Section
12 B
2 responses
nucleon number
The
A to D should be selected on the basis of
For
of the Aquestions
in this
section, on
onethe
or basis
more of
of the three numbered statements 1 to 3 may
3responses
relative
isotopic
mass
Theeach
to D should
be selected
be correct.
A
B
C
D
A each
B
C and
D
Decide
whether
statements
is 2or is not correct
(you3 may find it helpful
to put a tick against
1 only
2
1 and
1,
2 andof
3 the
32 Which
molecules
are
planar?
the statements that
be correct).
is
only are
are
areyou consider to only
1, 2 and
3
1 and
2
2 and
3
1 only
1 BCl
correct
correct
correct
correct
3
are
only
are
only
are
is
The responses A to D should be selected on the basis of
correct
correct
correct
2 NHcorrect
3
No other combination
of statements
B is used as a correct
C response.
D
3 PH3 A
No other combination of statements is used as a correct response.
1, 2 and 3
1 and 2
2 and 3
1 only
31 The diagram
illustrates the
a setare
of reactions.
are
onlyenergy
are changes ofonly
is
331 Boron is a non-metallic element which is placed above aluminium in Group III of the Periodic
correct
correct
correct
correct
32 Table.
The diagram
the with
Boltzmann
distribution
molecular
ataagraphite
given temperature.
It formsrepresents
a compound
nitrogen
known as
boron
nitride energies
which has
structure.
_ of
H = _134 kJ mol 1
R
S
No other
combination
of statements
is used
as drawn
a correct
response.
Which
of the following
conclusions
can be
from
this information?
1 The empirical formula of boron nitride is BN.
31 A monomer undergoes addition polymerisation. A 1 mol sample of the monomer is completely
2 The boron and nitride atoms are likely to be arranged alternately in a hexagonal pattern.
polymerised.
proportion
1
= +92between
J mol the layers.
3 Boron nitride
has a layer structure with van13
der Waals’Hforces
of molecules
How many moles of polymer might, theoretically, be formed?
[W’05 Q33]
342 1
What1 is involved when a hydrogen bond is formed between two molecules?
_
–6
_75 kJ
1
a hydrogen
atom bonded to H
an=atom
less
than itself
10
molelectronegative
U
T
1
3
2 a lone pair
of electrons
energy
6 .02 × 1023
Which of the following statements are correct?
3Which
an of
electrostatic
attraction between opposite charges
the factors that affect the rate of a reaction can be explained using
such a Boltzmann
[S’11 2 Q34]
1
The
enthalpy
change for the transformation U ⎯→ R is + 42 kJ mol–1 .
distribution?
323 Which physical properties are due to hydrogen bonding between water molecules?
© UCLES 2005
9701/01/O/N/05
[Turn over
2 the
The
enthalpy
change
transformation
T ⎯→
S is endothermic.
is the
stirred
into aqueous
sodium
hydroxide, the reaction that occurs
35 When
yellow
NCl 3for
1 increasing
theliquid
concentration
of
reactants
1
Water
has a higher the
boiling
point than
H2 S.
can 3
be represented
following
equation.
The enthalpyby
change
for the transformation
R ⎯→ T is – 33 kJ mol–1 .
2 increasing the temperature
2 Ice floats on water.
+ 6NaOH(aq) → N2(g) + 3NaCl (aq) + 3NaOCl (aq) + 3H2O(l)
2NClof3(l)
3 the addition
a catalyst
3 The H−O−H bond angle in water is approximately 104°.
What will be the result of this reaction?
[W’09 Q32]
1
2
334 Which types of intermolecular forces can exist between adjacent urea molecules?
1 The
nitrogeninundergoes
a redox are
reaction.
33 Which
equilibria,
which all species
gaseous, would have equilibrium constants, Kp, with no
units?
O
2 A bleaching solution remains after the reaction.
1 sulfur dioxide and oxygen in equilibrium with sulfur trioxide
3 The final solution gives a precipitate withCacidified silver nitrate.
NH2iodide
2 hydrogen and iodine in equilibriumHwith
2N hydrogen
urea
monoxide
andoxide
steam
equilibrium
withnon-metallic
carbon dioxide
and hydrogen
36 3
In a carbon
car engine
pollutant
Y,inwhich
contains
element
X, is formed.
1 hydrogen bonding
Further oxidation of Y to Z occurs in the atmosphere. In this further oxidation, 1 mol of Y reacts
© UCLES 2004
9701/01/O/N/04
oxygen.
with
0.5 mol of gaseous
2 permanent
dipole-dipole
forces
3 could
temporary
induced
dipole-dipole
X
be either
nitrogen
or sulfur. forces
Which statements about X, Y and Z can be correct?
CEDAR COLLEGE
1
The oxidation number of X increases by two from Y to Z.
© UCLES 2009
9701/11/O/N/09
[W’10 1 Q33]
INTERMOLECULAR FORCES WS 1
2 3 H2O
CH4
SF6CO2
BF3 SF6
3
CO2
SF6
CH4
11
Section
276 B
33 The concepts of bond energy, bond length and bond polarity are useful when comparing the
For each
of the questions inmolecules,
this section,
one
or more
of the three numbered statements 1 to 3 may
behaviour
e.g.
thermal
335 The
conceptsofofsimilar
bond energy, bond
length
andstability.
bond polarity are useful when comparing the
be correct.
behaviour of similar molecules, e.g. thermal stability.
For example, it could be said
12 correct (you may find it helpful to put a tick against
Decide whether each of the statements is or is not
For example, it could be said
the statements
thatwith
you the
consider
to be correct).
“Compared
HCl molecule,
the bond ……….X…………. of the HI molecule is
Section B
………..Y………..
.” molecule, the bond
“Compared
with the HCl
……….X…………. of the HI molecule is
The responses A to D should be selected on the basis of
………..Y……….. .”
For each of the questions in this section, one or more of the three numbered statements 1 to 3 may
Which pairs of words correctly complete the above sentence?
be correct.
Which pairs
the above sentence?
A of words correctly complete
B
C
D
12
Decide whether X
each of theYstatements is or is not correct (you may find it helpful to put a tick against
1, X2 and
1 and
2
2 and 3
1 only
the statements
that 3
youYconsider to
be correct).
Section
B
are
only
are
only
are
is
1
energy greater
correct
correct
correct
correct
1
energy
greater
The responses
A to D
should be selected on the basis of
For each
questions
in this section, one or more of the three numbered statements 1 to 3 may
2 of the
length
greater
2
length
greater
be correct.
A
B is used as a correct
C response.
D
No other
of less
statements
3 combination
polarity
3
polarity
less
Decide whether each of the statements is or is not correct (you may find it helpful to put a tick against
[S’02 Q33]
1, 2 and
2 and 3
1 only
the statements
that3you consider1toand
be 2correct).
are
only are
only are
is
316 Kevlar has the structure below.
correct
correct
The responses
A to D should be correct
selected on the basiscorrect
of
[Turn over
9701/1/M/J/02
O
O
O
[Turn over
9701/1/M/J/02O
No other combination
of statementsBis used as a correct C
response.
A
D
C
C
N
N
C
C
N
N
1, 2 and 3
1 and 2
2 and 3
1 only
areX is made from
are has the second
is
31 Compound
twoare
elements. H
Oneonly
element
Honly
H highest value of
H first
correct
correct
correct
correct
ionisation energy in its group and the other element has the third highest value of first ionisation
energy in its group.
Compared to a steel rope of similar dimensions, a Kevlar rope is both lighter and stronger.
No other combination of statements is used as a correct response.
Which compounds could be compound X?
Which properties of Kevlar help to explain these facts?
1 calcium chloride
31 The
chemical
terms
can be illustrated
1 definitions
The fibresofofmany
Kevlar
align due
to hydrogen
bonding.by chemical equations.
2 magnesium bromide
2 The
mass
length by
is less
in a Kevlar
than
a steel rope.
Which
terms
canper
beunit
illustrated
an equation
that rope
shows
theinformation
of a positive ion?
3 potassium sulfide
Theionisation
Kevlar molecule
13 first
energy has no permanent dipole.
[W’08 Q31]
2 heterolytic fission
327 Water has some unusual physical properties compared to other hydrides of Group 16 elements.
32 Some
Whichofof these
the following
canare
act due
as a to
Bronsted-Lowry
acid?
properties
hydrogen bonds.
These intermolecular forces are much
3 enthalpy
change
of atomisation
S,
for
example.
stronger in
water
than
they
are
in
H
2
1 H 3O+
statements
are correct?
+
32 Which
Why
NH aluminium chloride, Al2Cl6, sublime at the relatively low temperature of 180 °C?
2 does
4
11
3
22
Hydrogen
bonds cause
thebetween
melting the
point
ice to be higher than expected.
The intermolecular
forces
Al of
2Cl 6 molecules are weak.
H 2O
Hydrogen
bondsbonds
causebetween
the surface
tension and
of water
to be
higher
The co-ordinate
aluminium
chlorine
are
weak.than expected.
Hydrogen
bonds
cause
thegoverns
viscosity
of water
toabe
higher
than expected.
Thegiven
covalent
bonds
between
aluminium
and of
chlorine
are weak.
33 33Under
conditions,
what
the
rate
forward
reaction?
[S’16 3 Q32]
1 the activation energy of the reaction
338 The three statements that follow are all true.
2 the enthalpy change of the reaction
Which of these can be explained, at least in part, by reference to hydrogen bonding?
3 the equilibrium constant of the reaction
1 At 0 °C ice floats on water.
2
The boiling point of propan-2-ol is 82 °C. The boiling point of propanone is 56 °C.
3 2008
At 20 °C propanone and propanal mix completely.
© UCLES
9701/01/O/N/08
[Turn over
[W’11 3 Q32]
CEDAR COLLEGE
© UCLES 2016
INTERMOLECULAR FORCES WS 1
9701/13/M/J/16
correct
correct
correct
correct
For each of the questions in this section, one or more of the three numbered statements 1 to 3 may
be other
correct.
No
combination of statements is used as a correct response.
277
Decide whether each of the statements is or is not correct (you may find it helpful to put a tick against
the
that you forces
consider
to be correct).
9 statements
36
The intermolecular
between
iodine molecules are instantaneous dipole-induced dipole
forces.
The responses A to D should be selected on the basis of
Which statements explain why iodine has these intermolecular forces?
A
B
C
D
1 An iodine molecule is polar and experiences an attraction from a lone pair of electrons on an
adjacent
1, 2 andmolecule.
3
1 and 2
2 and 3
1 only
are
only are
only are
is
2 An iodine molecule has a fluctuating dipole because the electrons in a molecule are more
correct
correct
correct
correct
mobile than the nuclei.
become unsymmetrical and may then
3 The
electron charge
cloud within
an as
I2 molecule
No other
combination
of statements
is used
a correct may
response.
attract other I2 molecules.
[S’14 3 Q36]
10 Which molecules have an overall dipole moment?
31
37 Which reactions must be warmed to form a solid product?
1
1
2
2
33
carbon monoxide, CO
CH3 CH2CHO + 2,4-dinitrophenylhydrazine reagent
phosphine, PH
CH3 CH2CHO + 3 Fehling’s reagent
carbon
dioxide,
CO2
CH CHO
+ Tollens’
reagent
CH
3
2
32 When
The diagram
illustrates
the enthalpy
changes
a set of reactions
38
hydrolysed
with dilute
sulfuric acid,
whichof
compounds
produce propanoic acid?
1
CH3 CH2CO2CH3
2
CH3 CH2CH2CN
3
CH3 CH2CH2Cl
R
∆H = –134 kJ mol–1
S
∆H = +92 kJ mol–1
T
∆H = –75 kJ mol–1
U
Which statements are correct?
1
The enthalpy change for the transformation U → R is + 42 kJ mol–1.
2
The enthalpy change for the transformation T → S is endothermic.
3
The enthalpy change for the transformation R → T is – 33 kJ mol–1.
© UCLES 2014
© UCLES 2016
CEDAR COLLEGE
9701/13/M/J/14
9701/12/F/M/16
[Turn over
INTERMOLECULAR FORCES WS 1
278
279
2
1
INTERMOLECULAR FORCES WS 2
INTERMOLECULAR
FORCES WS 2
SECTION A
Answer all the questions in the spaces provided.
11
Ethene, C2H4, and hydrazine, N2H4, are hydrides of elements which are adjacent in the
Periodic Table. Data about ethene and hydrazine are given in the table below.
C2H4
N2H4
melting
point/°C
–169
+2
boiling
point/°C
–104
+114
solubility in
water
insoluble
high
solubility in
ethanol
high
high
For
Examiner’s
Use
(a) Ethene and hydrazine have a similar arrangement of atoms but differently shaped
molecules.
(i)
What is the H-C-H bond angle in ethene?
..................................................................................................................................
(ii)
Draw a ‘dot-and-cross’ diagram for hydrazine.
(iii)
What is the H-N-H bond angle in hydrazine?
..................................................................................................................................
[4]
(b) The melting and boiling points of hydrazine are much higher than those of ethene.
Suggest reasons for these differences in terms of the intermolecular forces each
compound possesses.
..........................................................................................................................................
..........................................................................................................................................
..........................................................................................................................................
..........................................................................................................................................
......................................................................................................................................[3]
CEDAR COLLEGE
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INTERMOLECULAR FORCES WS 2
INTERMOLECULAR FORCES WS 2
2
283
279 in the spaces provided.
Answer all the questions
1
15
The structural formulae of water, methanol and methoxymethane, CH3 OCH3 , are given
below.
O
O
H
(a) (i)
H3 C
H
For
Examiner’s
Use
O
H
H3 C
CH3
How many lone pairs of electrons are there around the oxygen atom in
methoxymethane?
..................................................................................................................................
(ii)
Suggest the size of the C–O–C bond angle in methoxymethane.
..................................................................................................................................
[2]
The physical properties of a covalent compound, such as its melting point, boiling point,
vapour pressure, or solubility, are related to the strength of attractive forces between the
molecules of that compound.
These relatively weak attractive forces are called intermolecular forces. They differ in their
strength and include the following.
A
interactions involving permanent dipoles
B
interactions involving temporary or induced dipoles
C
hydrogen bonds
(b) By using the letters A, B, or C, state the strongest intermolecular force present in each
of the following compounds.
For each compound, write the answer on the dotted line.
ethanal
CH3 CHO
..............
ethanol
CH3 CH2OH
..............
methoxymethane
CH3 OCH3
..............
2-methylpropane
(CH3 )2CHCH3
..............
© UCLES 2008
BILAL
HAMEED
CEDAR
COLLEGE
[4]
9701/02/M/J/08
INTERMOLECULAR
INTERMOLECULARFORCES
FORCESWS
WS22
(c) Methanol and water are completely soluble in each other.
(i) Calculate the average Mr of the mixture.
(i) Which intermolecular force exists between
methanol molecules and water molecules
284
280
3
that makes these two liquids soluble in each other?
1
(c) Methanol and water are completely soluble in each other.
..................................................................................................................................
(ii)
(i) Which intermolecular force exists between methanol molecules and water molecules
Draw
diagram
thattwo
clearly
shows
thisinintermolecular
thatamakes
these
liquids
soluble
each other? force. Your diagram should
show any lone pairs or dipoles present on either molecule that you consider to be
important.
..................................................................................................................................
average Mr = ................................. [2]
For
Examiner’s
Use
For
Examiner’s
Use
(ii) Draw a diagram that clearly shows this intermolecular force. Your diagram should
(ii) Use your answer to (i) to calculate the percentage of neon in the mixture.
show any lone pairs or dipoles present on either molecule that you consider to be
Give your answer to three significant figures.
important.
[4]
(d) When equal volumes of ethoxyethane, C2H5OC2H5, and water are mixed, shaken, and
then allowed to stand, two layers are formed.
percentage of neon = ................................. % [1]
[4]
Suggest
why
ethoxyethane
does
not
fully
dissolve
in
water.
Explain
your
answer.
6
(e) Neon and argon can both be obtained by fractional distillation of liquid air as they have different
(d) When
equal volumes of ethoxyethane, C2H5OC2H5, and water are mixed, shaken, and
boiling
points.
then allowed to stand, two layers are formed.
..........................................................................................................................................
Neon has a boiling point of 27.3 K. The boiling point of argon is 87.4 K.
Suggest why ethoxyethane does not fully dissolve in water. Explain your answer.
..........................................................................................................................................
(i) Name the force that has to be overcome in order to boil neon or argon and explain what
..........................................................................................................................................
..........................................................................................................................................
causes
it.
..........................................................................................................................................
......................................................................................................................................
[2]
.............................................................................................................................................
..........................................................................................................................................
[Total: 12]
.............................................................................................................................................
...................................................................................................................................... [2]
.......................................................................................................................................
[3]
(ii) Explain why argon has a higher boiling point than neon.
[Total: 12]
.............................................................................................................................................
.............................................................................................................................................
....................................................................................................................................... [2]
[Total: 18]
© UCLES 2015
© UCLES 2008
© UCLES 2008
BILAL
HAMEED
CEDAR
COLLEGE
9701/23/M/J/15
9701/02/M/J/08
9701/02/M/J/08
[Turn over
[Turn over
[Turn over
INTERMOLECULAR
INTERMOLECULARFORCES
FORCESWS
WS22
285
2814
1
2 7 Carbon disulphide, CS2, is a volatile, stinking liquid which is used to manufacture viscose
rayon and cellophane.
4
(a) The carbon atom is in the centre of the CS2 molecule.
2
Sulphur and its compounds are found in volcanoes, in organic matter and in minerals.
Draw a ‘dot-and-cross’ diagram of the carbon disulphide molecule.
Sulphuric acid, an important industrial chemical, is manufactured from sulphur by the
Contact
process.
There are
three consecutive reactions in the Contact process which are
Show
outer electrons
only.
essential.
(a) Write a balanced equation (using
in the correct sequence.
For
Examiner’s
Use
For
Examiner’s
Use
where appropriate) for each of these reactions
1 .......................................................................................................................................
[2]
2 .......................................................................................................................................
(b) Suggest the shape of the molecule and give its bond angle.
3 ................................................................................................................................. [4]
shapecatalyst
.........................................................
(b) What
is used?
bond
angle .................................................
[2]
....................................................................................................................................
[1]
(c) Explain
the termHstandard
enthalpy change of formation, !H f .
Hydrogen
sulphide,
2S, is a foul-smelling compound found in the gases from volcanoes.
Hydrogen sulphide is covalent, melting at –85 °C and boiling at –60 °C.
..........................................................................................................................................
(c)
(i) Draw a ‘dot-and-cross’ diagram to show the structure of the H2S molecule.
..........................................................................................................................................
.................................................................................................................................... [3]
(d) Calculate the standard enthalpy change of formation of CS2 from the following data.
standard enthalpy change of formation of SO2
= –298 kJ mol–1
standard enthalpy change of formation of CO2
(ii) Predict the shape of the H2S molecule.
= –395 kJ mol–1
standard enthalpy change of combustion of CS2 = –1110 kJ mol–1
.............................................................
(iii) Oxygen and sulphur are both in Group VI of the Periodic Table.
Suggest why the melting and boiling points of water, H2O, are much higher than
those of H2S.
..................................................................................................................................
..................................................................................................................................
[3]
............................................................................................................................ [4]
© UCLES 2005
BILAL
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9701/02/O/N/05
INTERMOLECULAR
INTERMOLECULARFORCES
FORCESWS
WS2 2
2
Answer all the questions
in the spaces provided.
286
282
1
18 (a) Fill the gaps in the table for each of the given particles.
name
of isotope
type
of particle
charge
symbol
carbon-13
electron
configuration
1s22s22p2
–1
sulfur-34
atom
iron-54
cation
37
–
17 Cl
0
1s22s22p63s23p63d6
[5]
(b) One of the factors that determines the type of bonding present between the particles of a
substance is the relative electronegativities of the bonded particles.
(i) Explain the meaning of the term electronegativity.
.............................................................................................................................................
.............................................................................................................................................
....................................................................................................................................... [2]
(ii) Name and describe the type of bonding you would expect to find between particles with
equal electronegativities.
.............................................................................................................................................
.............................................................................................................................................
....................................................................................................................................... [2]
(iii) Name and describe the type of bonding you would expect to find between particles with
very different electronegativities.
.............................................................................................................................................
.............................................................................................................................................
....................................................................................................................................... [2]
© UCLES 2015
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FORCESWS
WS22
287
283
1
3
(c) The boiling points of some molecules with equal numbers of electrons are given.
substance
fluorine
argon
hydrogen
chloride
methanol
formula
F2
Ar
HCl
CH3OH
boiling point / K
85
87
188
338
(i) Explain why the boiling points of fluorine and argon are so similar.
.............................................................................................................................................
.............................................................................................................................................
....................................................................................................................................... [2]
(ii) Explain why the boiling point of hydrogen chloride is higher than that of fluorine.
.............................................................................................................................................
.............................................................................................................................................
....................................................................................................................................... [2]
(iii) Explain why methanol has the highest boiling point of all these molecules.
.............................................................................................................................................
.............................................................................................................................................
....................................................................................................................................... [2]
[Total: 17]
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[Turn over
INTERMOLECULAR
FORCES
WS
22
INTERMOLECULAR
FORCES
WS
2
288
284
1
Answer all the questions in the spaces provided.
19 Elements and compounds which have small molecules usually exist as gases or liquids.
(a) Chlorine, Cl2, is a gas at room temperature whereas bromine, Br2, is a liquid under the same
conditions.
Explain these observations.
...................................................................................................................................................
...................................................................................................................................................
.............................................................................................................................................. [2]
(b) The gases nitrogen, N2, and carbon monoxide, CO, are isoelectronic, that is they have the
same number of electrons in their molecules.
Suggest why N2 has a lower boiling point than CO.
...................................................................................................................................................
...................................................................................................................................................
.............................................................................................................................................. [2]
(c) A ‘dot-and-cross’ diagram of a CO molecule is shown below. Only electrons from outer shells
are represented.
C
O
In the table below, there are three copies of this structure.
On the structures, draw a circle around a pair of electrons that is associated with each of the
following.
a co-ordinate bond
C
O
a covalent bond
C
O
a lone pair
C
O
[3]
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285
Cambridge International AS and A Level Chemistry 9701 syllabus General syllabus requirements and information
DATA BOOKLET
1 Important values, constants and standards
molar gas constant
R = 8.31 J K–1 mol–1
the Faraday constant
F = 9.65 × 104 C mol–1
the Avogadro constant
L = 6.02 x 1023 mol–1
the Planck constant
h = 6.63 × 10–34 J s
speed of light in a vacuum
c = 3.00 × 108 m s–1
rest mass of proton, 11H
mp = 1.67 × 10–27 kg
rest mass of neutron, 01 n
mn = 1.67 × 10–27 kg
rest mass of electron, -10e
me = 9.11 × 10–31 kg
electronic charge
e = –1.60 × 10–19 C
molar volume of gas
Vm = 22.4 dm3 mol–1 at s.t.p
Vm = 24.0 dm3 mol–1 under room conditions
(where s.t.p. is expressed as 101 kPa, approximately, and 273 K (0 °C))
ionic product of water
Kw = 1.00 × 10–14 mol2 dm–6
(at 298 K (25 °C))
specific heat capacity of water
= 4.18 kJ kg–1 K–1
(= 4.18 J g–1 K–1)
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2 Ionisation energies (1st, 2nd, 3rd and 4th) of selected elements,
in kJ mol–1
78
Proton
number
First
Second
Third
Fourth
H
1
1310
–
–
–
He
2
2370
5250
–
–
Li
3
519
7300
11800
–
Be
4
900
1760
14800
21000
B
5
799
2420
3660
25000
C
6
1090
2350
4610
6220
N
7
1400
2860
4590
7480
O
8
1310
3390
5320
7450
F
9
1680
3370
6040
8410
Ne
10
2080
3950
6150
9290
Na
11
494
4560
6940
9540
Mg
12
736
1450
7740
10500
Al
13
577
1820
2740
11600
Si
14
786
1580
3230
4360
P
15
1060
1900
2920
4960
S
16
1000
2260
3390
4540
Cl
17
1260
2300
3850
5150
Ar
18
1520
2660
3950
5770
K
19
418
3070
4600
5860
Ca
20
590
1150
4940
6480
Sc
21
632
1240
2390
7110
Ti
22
661
1310
2720
4170
V
23
648
1370
2870
4600
Cr
24
653
1590
2990
4770
Mn
25
716
1510
3250
5190
Fe
26
762
1560
2960
5400
Co
27
757
1640
3230
5100
Ni
28
736
1750
3390
5400
Cu
29
745
1960
3350
5690
Zn
30
908
1730
3828
5980
Ga
31
577
1980
2960
6190
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Proton
number
First
Second
Third
Fourth
Br
35
1140
2080
3460
4850
Rb
37
403
2632
3900
5080
Sr
38
548
1060
4120
5440
Ag
47
731
2074
3361
–
I
53
1010
1840
2040
4030
Cs
55
376
2420
3300
–
Ba
56
502
966
3390
–
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3 Bond energies
3(a) Bond energies in diatomic molecules (these are exact values)
Homonuclear
80
Heteronuclear
Bond
Energy/kJ mol
H−H
–1
Bond
Energy/kJ mol–1
436
H−F
562
D−D
442
H−Cl
431
N≡N
944
H−Br
366
O=O
496
H−I
299
P≡P
485
C≡O
1077
S=S
425
F−F
158
Cl −Cl
242
Br−Br
193
I−I
151
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3(b) Bond energies in polyatomic molecules (these are average values)
Homonuclear
Heteronuclear
Bond
Energy/kJ mol–1
Bond
Energy/kJ mol–1
C−C
350
C−H
410
C=C
610
C−Cl
340
C≡C
840
C−Br
280
C
520
C−I
240
N−N
160
C−N
305
N=N
410
C=N
610
O−O
150
C≡N
890
Si−Si
222
C−O
360
P−P
200
C=O
740
S−S
264
C=O in CO2
805
N−H
390
N−Cl
310
O−H
460
Si−Cl
359
Si−H
320
Si−O (in SiO2(s))
460
Si=O (in SiO2(g))
640
P−H
320
P−Cl
330
P−O
340
P=O
540
S−H
347
S−Cl
250
S−O
360
S=O
500
C (benzene)
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4 Standard electrode potential and redox potentials, E⦵ at 298 K (25 °C)
For ease of reference, two tables are given:
(a) an extended list in alphabetical order
(b) a shorter list in decreasing order of magnitude, i.e. a redox series.
(a) E in alphabetical order
⦵
Electrode reaction
E /V
Ag+ + e–
⇌
Ag
+0.80
Al 3+ + 3e–
⇌
Al
–1.66
Ba2+ + 2e–
⇌
Ba
–
⇌
2Br
Ca2+ + 2e–
⇌
Ca
–
⇌
2Cl
2HOCl + 2H+ + 2e–
⇌
Cl2 + 2H2O
+1.64
ClO– + H2O + 2e–
⇌
Cl – + 2OH–
+0.89
Co + 2e
–
⇌
Co
–0.28
3+
–
⇌
Co
2+
–
⇌
Co + 6NH3
–0.43
2+
–
⇌
Cr
–0.91
3+
Cr + 3e
–
⇌
Cr
–0.74
3+
–
⇌
+
–
⇌
2Cr + 7H2O
+1.33
+
Cu + e
–
⇌
Cu
+0.52
2+
Cu + 2e
–
⇌
Cu
2+
–
⇌
Cu
2+
[Cu(NH3)4] + 2e
–
⇌
Cu + 4NH3
F2 + 2e
–
Fe2+ + 2e–
Br2 + 2e
Cl 2 + 2e
2+
Co + e
[Co(NH3)6] + 2e
Cr + 2e
Cr + e
2–
Cr2O7 + 14H + 6e
Cu + e
–2.90
–
+1.07
–2.87
–
+1.36
2+
+1.82
2+
Cr
–0.41
3+
+0.34
+
+0.15
–0.05
⇌
–
2F
+2.87
⇌
Fe
–0.44
–
⇌
Fe
–0.04
Fe3+ + e–
⇌
Fe2+
3+
Fe + 3e
+0.77
–
⇌
[Fe(CN)6]
+0.36
Fe(OH)3 + e–
⇌
Fe(OH)2 + OH–
–0.56
+
–
⇌
H2
0.00
2H2O + 2e–
⇌
H2 + 2OH–
–
⇌
3–
[Fe(CN)6] + e
2H + 2e
I2 + 2e
82
⦵
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4–
2I
–
–0.83
+0.54
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Electrode reaction
E /V
⦵
⇌
K
–2.92
–
⇌
Li
–3.04
Mg2+ + 2e–
K + + e–
+
Li + e
⇌
Mg
–2.38
–
⇌
Mn
–1.18
Mn3+ + e–
⇌
Mn2+
+1.49
–
⇌
2+
Mn + 2H2O
+1.23
MnO4– + e–
⇌
MnO42–
+0.56
+
–
⇌
MnO2 + 2H2O
+1.67
MnO4– + 8H+ + 5e–
⇌
Mn2+ + 4H2O
+1.52
–
⇌
NO2 + H2O
+0.81
NO3– + 3H+ + 2e–
⇌
HNO2 + H2O
+0.94
2+
Mn + 2e
+
MnO2 + 4H + 2e
–
MnO4 + 4H + 3e
–
+
NO3 + 2H + e
–
⇌
NH4 + 3H2O
+0.87
Na+ + e–
⇌
Na
–2.71
2+
–
⇌
Ni
–0.25
[Ni(NH3)6]2+ + 2e–
–
+
NO3 + 10H + 8e
Ni + 2e
+
⇌
Ni + 6NH3
–0.51
–
⇌
2H2O
+1.77
HO2– + H2O + 2e–
⇌
3OH–
+0.88
–
⇌
2H2O
+1.23
O2 + 2H2O + 4e–
⇌
4OH–
+0.40
–
⇌
H2 O 2
+0.68
O2 + H2O + 2e–
⇌
HO2– + OH–
–0.08
2+
–
⇌
Pb
–0.13
Pb4+ + 2e–
⇌
Pb2+
+1.69
–
⇌
2+
Pb + 2H2O
+1.47
SO42– + 4H+ + 2e–
+
H2O2 + 2H + 2e
+
O2 + 4H + 4e
+
O2 + 2H + 2e
Pb + 2e
+
PbO2 + 4H + 2e
⇌
SO2 + 2H2O
+0.17
–
⇌
2SO4
2–
+2.01
S4O62–+ 2e–
⇌
2S2O32–
+0.09
–
⇌
Sn
–0.14
Sn4+ + 2e–
⇌
Sn2+
+0.15
–
⇌
V
–1.20
V3+ + e–
⇌
V2+
–0.26
–
⇌
3+
V + H2 O
+0.34
VO2+ + 2H+ + e–
⇌
VO2+ + H2O
+1.00
2–
S2O8 + 2e
2+
Sn + 2e
2+
V + 2e
2+
+
VO + 2H + e
–
–
⇌
VO + 2H2O
+1.00
Zn2+ + 2e–
⇌
Zn
–0.76
+
VO3 + 4H + e
2+
All ionic states refer to aqueous ions but other state symbols have been omitted.
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(b) E in decreasing order of oxidising power
⦵
(a selection only – see also the extended alphabetical list on the previous pages)
Electrode reaction
E /V
F2 + 2e–
⇌
2F–
+2.87
S2O82–+ 2e–
⇌
2SO42–
+2.01
+
–
⇌
2H2O
+1.77
MnO4– + 8H+ + 5e–
⇌
Mn2+ + 4H2O
+1.52
H2O2 + 2H + 2e
⇌
2+
Pb + 2H2O
+1.47
⇌
2Cl –
+1.36
–
⇌
3+
2Cr + 7H2O
+1.33
O2 + 4H+ + 4e–
⇌
2H2O
+1.23
–
⇌
2Br
ClO– + H2O + 2e–
⇌
Cl – + 2OH–
+
–
Cl 2 + 2e–
PbO2 + 4H + 2e
2–
+
Cr2O7 + 14H + 6e
Br2 + 2e
–
+1.07
+0.89
–
⇌
NH4 + 3H2O
+0.87
NO3– + 2H+ + e–
–
+
NO3 + 10H + 8e
+
⇌
NO2 + H2O
+0.81
–
⇌
Ag
+0.80
Fe3+ + e–
⇌
Fe2+
+0.77
I2 + 2e
–
⇌
–
2I
+0.54
O2 + 2H2O + 4e–
⇌
4OH–
+0.40
+
Ag + e
2–
2+
Cu + 2e
–
⇌
Cu
+0.34
+
–
⇌
SO2 + 2H2O
+0.17
4+
Sn + 2e
–
⇌
2–
S4O6 + 2e
–
⇌
2S2O3
+0.09
+
–
⇌
H2
0.00
2+
–
⇌
Pb
–0.13
2+
–
⇌
Sn
–0.14
Fe2+ + 2e–
SO4 + 4H + 2e
2H + 2e
Pb + 2e
Sn + 2e
Sn
2+
+0.15
2–
⇌
Fe
–0.44
–
⇌
Zn
–0.76
2H2O + 2e–
⇌
H2 + 2OH–
–0.83
–
⇌
V
–1.20
Mg2+ + 2e–
2+
Zn + 2e
2+
V + 2e
⇌
Mg
–2.38
–
⇌
Ca
–2.87
K + + e–
⇌
K
–2.92
2+
Ca + 2e
84
⦵
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5 Atomic and ionic radii
(a) Period 1
atomic/nm
ionic/nm
single covalent
H
0.037
H+
van der Waals
He
0.140
Li
0.152
0.208
(b) Period 2
metallic
single covalent
van der Waals
Li+
0.060
2+
0.031
Be
0.112
Be
B
0.080
B3+
0.020
C
0.077
4+
0.015
N
C4–
0.260
0.074
N3–
0.171
O
0.073
O
2–
0.140
F
0.072
F–
0.136
Ne
0.160
Na
0.186
C
(c) Period 3
metallic
single covalent
van der Waals
Na+
0.095
2+
0.065
Mg
0.160
Mg
Al
0.143
Al 3+
0.050
Si
0.117
Si4+
0.041
P
0.110
P3–
0.212
S
0.104
S2–
0.184
Cl
0.099
–
0.181
Ar
0.190
Be
0.112
Be2+
0.031
Mg
0.160
Mg
2+
0.065
Ca
0.197
Ca2+
0.099
Sr
0.215
Sr
2+
0.113
Ba
0.217
Ba2+
0.135
0.220
2+
0.140
Cl
(d) Group 2
metallic
Ra
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(e) Group 14
single covalent
metallic
atomic/nm
ionic/nm
C
0.077
Si
0.117
Si4+
0.041
Ge
0.122
Ge2+
0.093
Sn
0.162
Sn2+
0.112
Pb
0.175
Pb
2+
0.120
F
0.072
F–
Cl
0.099
(f) Group 17
single covalent
0.136
Cl
–
0.181
–
0.195
Br
0.114
Br
I
0.133
I–
At
0.140
Sc
0.164
0.216
(g) First row transition elements
metallic
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0.081
3+
0.067
Ti
0.146
Ti
0.090
Ti
V
0.135
V2+
0.079
V3+
0.064
Cr
0.129
2+
Cr
0.073
Cr
3+
0.062
Mn
0.132
Mn2+
0.067
Mn3+
0.062
Fe
0.126
2+
Fe
0.061
Fe
3+
0.055
Co
0.125
Co2+
0.078
Co3+
0.053
Ni
0.124
2+
Ni
0.070
3+
0.056
Cu
0.128
Cu2+
0.073
0.135
2+
0.075
Zn
86
2+
Sc3+
Zn
Ni
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6 Typical proton (1H) chemical shift values (δ) relative to TMS = 0
Type of
proton
Environment of proton
Example structures
Chemical shift
range (δ)
–CH3, –CH2–, >CH–
0.9–1.7
CH3–C=O, –CH2–C=O, >CH–C=O
2.2–3.0
alkyl next to aromatic ring
CH3–Ar, –CH2–Ar, > CH–Ar
2.3–3.0
alkyl next to
electronegative atom
CH3–O, –CH2–O, –CH2–Cl,
>CH–Br
3.2–4.0
attached to alkyne
≡C−H
1.8–3.1
attached to alkene
=CH2, =CH–
4.5–6.0
attached to aromatic ring
–H
6.0–9.0
alkane
alkyl next to C=O
C–H
O
aldehyde
R
C
9.3–10.5
H
alcohol
O–H
(see note
below)
RO–H
0.5–6.0
–OH
phenol
4.5–7.0
O
carboxylic acid
R
9.0–13.0
C
O
alkyl amine
R–NH–
1.0–5.0
–NH2
aryl amine
N–H
(see note
below)
H
3.0–6.0
O
amide
R
C
5.0–12.0
N
H
Note: δ values for –O–H and –N–H protons can vary depending on solvent and concentration.
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296
7 Typical carbon (13C) chemical shift values (δ) relative to TMS = 0
88
Hybridisation of
the carbon atom
Environment of carbon
atom
Example structures
sp3
alkyl
CH3–, CH2–, –CH<
0–50
sp3
next to alkene/arene
−CH2−C=C, −CH2−
10–40
sp3
next to carbonyl/carboxyl
−CH2−COR, −CH2−CO2R
25–50
sp3
next to nitrogen
−CH2−NH2, −CH2−NR2, −CH2−NHCO
30–65
sp3
next to chlorine
(−CH2−Br and −CH2−I are in
the same range
as alkyl)
−CH2−Cl
30–60
sp3
next to oxygen
−CH2−OH, −CH2−O−CO−
50–70
sp2
alkene or arene
>C=C<, C
C
C
C
C
C
Chemical
shift range
( δ)
110–160
sp2
carboxyl
R−CO2H, R−CO2R
160–185
sp2
carbonyl
R−CHO, R−CO−R
190–220
sp
alkyne
R−C≡C−
65–85
sp
nitrile
R−C≡N
100–125
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8 Characteristic infra-red absorption frequencies for some selected
bonds
Functional groups containing
the bond
Bond
Absorption range (in
wavenumbers)/cm–1
Appearance of peak
(s = strong, w = weak)
C−O
alcohols, ethers, esters
1040–1300
s
C=C
aromatic compounds, alkenes
1500–1680
w unless conjugated
C=O
amides
ketones and aldehydes
esters
1640–1690
1670–1740
1710–1750
s
s
s
C≡C
alkynes
2150–2250
w unless conjugated
C≡N
nitriles
2200–2250
w
C−H
alkanes, CH2−H
alkenes/arenes, =C−H
2850–2950
3000–3100
s
w
N−H
amines, amides
3300–3500
w
O−H
carboxylic acids, RCO2−H
H-bonded alcohol, RO−H
free alcohol, RO−H
2500–3000
3200–3600
3580–3650
s and very broad
s
s and sharp
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9 The orientating effect of groups in aromatic substitution reactions.
The position of the incoming group, Y, is determined by the nature of the group, X, already bonded to the
ring, and not by the nature of the incoming group Y.
X
X
2
+Y+
+
+H
3
4
90
Y
X– groups that direct the incoming
Y group to the 2– or 4– positions
X– groups that direct the incoming
Y group to the 3– position
−NH2, −NHR or −NR2
−NO2
−OH or −OR
−NH3
−NHCOR
−CN
−CH2, −alkyl
−CHO, –COR
−Cl
−CO2H, −CO2R
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299
10 Names, structures and abbreviations of some amino acids
structure of side chain R– in
Name
3-letter
abbreviation
NH2
1-letter symbol
CH
R
CO2H
alanine
Ala
A
CH3−
aspartic acid
Asp
D
HO2CCH2−
cysteine
Cys
C
HSCH2−
glutamic acid
Glu
E
HO2CCH2CH2−
glycine
Gly
G
H−
lysine
Lys
K
H2NCH2CH2CH2CH2−
phenylalanine
Phe
F
CH2
serine
Ser
S
HOCH2−
tyrosine
Tyr
Y
HO
CH2
CH3
valine
Val
V
CH
CH3
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45.0
39
Y
40.1
38
39.1
37
88
Ra
radium
–
87
Fr
francium
–
actinoids
lanthanoids
137.3
132.9
Hf
Ba
barium
Cs
lanthanoids
56
caesium
72
57–71
87.6
55
Th
thorium
232.0
Ac
–
140.1
90
138.9
89
actinium
Ce
cerium
La
58
–
lanthanum
57
Rf
actinoids
rutherfordium
104
178.5
hafnium
91.2
89–103
88.9
Zr
zirconium
85.5
yttrium
Sr
strontium
Rb
rubidium
47.9
40
titanium
Ti
Sc
scandium
Ca
calcium
K
22
21
potassium
4
3
23.0
19
7
Mn
Tc
Mo
Nd
60
–
seaborgium
Sg
106
183.8
tungsten
W
74
95.9
231.0
protactinium
Pa
91
140.9
238.0
uranium
U
144.4
92
–
neptunium
Np
–
93
promethium
Pm
61
–
bohrium
Bh
107
186.2
rhenium
Re
75
–
molybdenum technetium
54.9
43
52.0
42
manganese
Cr
chromium
25
24
6
praseodymium neodymium
Pr
59
–
dubnium
Db
105
180.9
tantalum
Ta
73
92.9
niobium
Nb
50.9
41
vanadium
V
23
5
relative atomic mass
24.3
20
Mg
magnesium
Na
sodium
9.0
12
9
10
11
12
name
6.9
11
B
boron
atomic symbol
Be
beryllium
Li
lithium
–
plutonium
Pu
150.4
94
samarium
Sm
62
–
hassium
Hs
108
190.2
osmium
Os
76
101.1
ruthenium
Ru
55.8
44
iron
Fe
26
8
1.0
Ds
110
195.1
platinum
Pt
78
106.4
palladium
Pd
58.7
46
nickel
Ni
28
Rg
111
197.0
gold
Au
79
107.9
silver
Ag
63.5
47
copper
Cu
29
Cn
112
200.6
mercury
Hg
80
112.4
cadmium
Cd
65.4
48
zinc
Zn
30
–
americium
Am
152.0
95
europium
Eu
63
–
–
curium
Cm
157.3
96
gadolinium
Gd
64
–
–
berkelium
Bk
158.9
97
terbium
Tb
65
–
–
californium
Cf
162.5
98
dysprosium
Dy
66
–
meitnerium darmstadtium roentgenium copernicium
Mt
192.2
109
iridium
Ir
77
102.9
rhodium
Rh
58.9
45
cobalt
Co
27
–
einsteinium
Es
164.9
99
holmium
Ho
67
204.4
thallium
Tl
114.8
81
indium
In
69.7
49
gallium
Ga
27.0
31
aluminium
Al
10.8
13
5
atomic number
–
fermium
Fm
167.3
100
erbium
Er
68
–
flerovium
Fl
114
207.2
lead
Pb
82
118.7
tin
Sn
72.6
50
germanium
Ge
28.1
32
silicon
Si
12.0
14
carbon
C
6
–
mendelevium
Md
168.9
101
thulium
Tm
69
209.0
bismuth
Bi
83
121.8
antimony
Sb
74.9
51
arsenic
As
31.0
33
phosphorus
P
14.0
15
nitrogen
N
7
–
nobelium
No
173.1
102
ytterbium
Yb
70
–
livermorium
Lv
116
–
polonium
Po
84
127.6
tellurium
Te
79.0
52
selenium
Se
32.1
34
sulfur
S
16.0
16
oxygen
O
8
–
lawrencium
Lr
175.0
103
lutetium
Lu
71
–
astatine
At
126.9
85
iodine
I
79.9
53
bromine
Br
35.5
35
chlorine
Cl
19.0
17
fluorine
F
9
18
–
radon
Rn
86
131.3
xenon
Xe
83.8
54
krypton
Kr
39.9
36
argon
Ar
20.2
18
neon
Ne
4.0
10
helium
17
hydrogen
16
2
15
He
14
H
4
13
1
3
Key
Group
2
1
The Periodic Table of Elements
300
Cambridge International AS and A Level Chemistry 9701 syllabus
General syllabus requirements and information
DATA
BOOKLET
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