2 Atomic structure
2.3 Electrons: energy levels, atomic orbitals, ionisation energy
alt
Electronic Structure
Electronic Structure
Bilal Hameed
1
Electronic Structure
Cambridge International AS and A Level Chemistry 9701 syllabus Syllabus content
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
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 The
topicnucleus
describes
the type,a)number
andthe
distribution
of the
particles to
which
make
up an
2.2
of the
describe
contribution
of fundamental
protons and neutrons
atomic
nuclei
in
atomatom
and the impact of this on some
atomic
properties.
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:
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
relative charges and relative masses
a) describe the number and relative energies of the s, p and d orbitals for
b) the
deduce
the behaviour
of beams1,
of2protons,
neutrons
in
principal
quantum numbers
and 3 and
also theand
4s electrons
and 4p orbitals
electric fields
b) describe and sketch the shapes of s and p orbitals
c) describe the distribution of mass and charge within an atom
c) state the electronic configuration of atoms and ions given the proton
2 electrons
d) number
deduce the
of protons,
neutrons 1s
and
andnumbers
charge, using
the convention
2s22p6 , etc.present in both
atoms and ions given proton and nucleon numbers and charge
d) (i) explain and use the term ionisation energy
alt
2.1 Particles in the atom
2.3 Electrons: energy
levels, atomic
orbitals, ionisation
energy, electron
affinity
2.2 The nucleus of the
atom
2.3 Electrons: energy
levels, atomic
orbitals, ionisation
energy, electron
affinity
(ii) explain the factors influencing the ionisation energies of elements
a) describe the contribution of protons and neutrons to atomic nuclei in
(iii) explain the trends in ionisation energies across a Period and down a
terms of proton number and nucleon number
Group of the Periodic Table (see also Section 9.1)
b) distinguish between isotopes on the basis of different numbers of
e) deduce the electronic configurations of elements from successive
neutrons present
ionisation energy data
c) recognise and use the symbolism xy A for isotopes, where x is the nucleon
f) interpret successive ionisation energy data of an element in terms of the
number and y is the proton number
position of that
element within the Periodic Table
g) explain and use the term 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)
18
www.cie.org.uk/alevel
18
www.cie.org.uk/alevel
Electronic Structure
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
Back to contents page
Back to contents page
2
Bilal Hameed
18
Relative abundance
70
When sunlight (which contains
41 all wavelengths of visible light) passes through
a prism, the different wavelengths are bent (or refracted) through different
29
angles so that the light is broken into its components, producing a continuous
spectrum of colours. A similar effect is seen when sunlight passes through
raindrops to produce a rainbow.
40
1
Lear
As you have seen in Topic 1.4, in a simple modelFindingofthe visible
the
atom the
spectrum
20
15
electrons
are thought of as being arranged in shells around the
The Bohr model of the atom
ELECTRONIC STRUCTURE 1
nucleus.
The
shells
canBohrhold
increasing
numbers
of electrons as they
In 1913, Danish
physicist Niels
proposed
that the electrons move
around
0
the nucleus in fixed energy levels called shells. He proposed that each atom
N
10
30
40
50
60
70
20
M
When chemical reactions
take
electrons
the
outer
parts
of atoms
are
hasplace,
a series
of these shells.
shells
close to
nucleus
are of low
energy
get further
from
theinThe
nucleus
–thethe
pattern
is 2, 8, 18,432and so
on.
L
m/z one atom to another or they may
redistributed. The electrons may
be transferred
the nucleus
(1, 2, 3 …). Theyfrom
are also identified by letters (K, L, M …).
Electrons move around the nucleus in these shells in pathways called orbits.
be shared between the reacting
atoms
in
a
different
way. Protons and neutrons, in the
According to Bohr’s model of the atom, different shells (or energy levels) hold
different
numbers
of
electrons.
There
is a maximum number of electrons that
nuclei of atoms, take no part in chemical reactions.
2
and those further out are of higher energy. Shells are numbered outwards from
nucleus
K
1
alt
Energy levels
in order from greatest energy (E )
can fit in any shell. This maximum number is 2n , where n = the shell number.
to least energy: E > E > E > E
Electrons
in different shells have differing amounts
of energy. They
For example, sodium (Z = 11) has 11 electrons. Two electrons are in the first
Figure 1.3.4 Bohr’s model of the atom. The
shell, eight
are inatom
the second
and one electron
is in the
thirdare
shell.outside
This
According to Bohr’s can
atomic
model,
each
hasshell
electrons
in orbits
that
energy of an electron depends
largely upon
therefore
be
represented
on
an
energy
level
diagram.
The
shells
is called the electron arrangement or electron configuration of sodium and
its distance from the nucleus.
You
may and
be familiar
with
aenergy
picture
of
atoms
based
on
model
is written
as 2,8,1.
electrons
incan
the outer
shell can
also a
be called
valence
the
nucleus
are represent
of increasing
size.
TheThe
first
shell
accommodate
two
electrons,
levels
they
are
labelled
1,is2, 3, and so on (Figure 1).
electrons. As all of the
electronsand
are as close
to the
nucleus
as possible, this
suggested
by electrons,
the Danish
scientist,
Each
has
the third
lowest
energy
of a sodium
atom.
The
lowest atom
energy
of an electrons
atomin
the
second eight
the
shell state
18,Niels
and
soBohr.
on.
The
number
ofstate
electrons
Each main
canstate.
hold up to a maximum number of electrons given
is knownshell
as the ground
in orbits
are outside
the nucleus
and are of increasing size. The first
n is the number
of the
shell.
each
shell =that
2n2 , where
2, where
by the formula
2n
n is the
number ARRANGEMENTS
of the main shell. So, you
1.3.1electrons,
ELECTRONS IN SHELLS
TABLE
1.3.2 ELECTRON
shell can accommodateTABLE
two
the second
eight
electrons, and OF SOME ELEMENTS
Shell
number
Maximum number
can have
two
electrons
in the first
main
shell,
eight in
the
next, 18 in the
Element
name
Atomic number
Electron
arrangement
in this
so on. Such diagrams show(n)what ofiselectrons
known
as theNitrogen
ground state of
the
7
2,5
shell
(2n
)
next,
on. are in positions
8
Oxygen
2,6
atom. This means
thatand
its so
electrons
that are as close
1
3
2
1
ATOMIC STRUCTURE
1
4
➔ Il
c
io
p
Speci
2
2
10
Neon
2,8
8
as allowable to Apart
the nucleus
atom.
These
diagrams
are useful
as are divided
from23 of
thethe
first
shell,
these
main
energy
levels
into
17
Chlorine
2,8,7
18
a basic representation
of atomic
structure and help us to understand
4called s, p,32d, and f, which have slightly different energies
sub-shells,
the nature of chemical
bonding
Chapter
7). However,
their
use is
(Figure
2). Shell(see
2ground
has
an
s-sub-shell
and
a p-sub-shell.
Shell 3 an
Such
diagrams
show what
is relatively
known as the
state of the
atom. This
means
that
limited
to atoms
with
few
electrons.
Electrons
can
however
s-sub-shell,
and
a d-sub-shell.
its
electrons
areone
in positions
thatanother
are a
asp-sub-shell,
close
as allowable
to the
of the atom.
move
from
orbit to
within
an atom.
If nucleus
theUSEelectrons
arePROGRAMME
in a STANDARD LEVEL 15
CHEMISTRY:
FOR
WITH THE IB DIPLOMA
position of higher energy they are said to be excited.
CHAPTER 1
to give the
ration of an
ould give the
that element.
Figure 1.3.3 White light produces a continuous spectrum (ROYGBIV) when passed through a prism.
60
Blue
Indigo
Violet
3
2
1
up to 18 electrons
up to 8 electrons
Ground states of some elements
energy
3
3
energy
of an atom
aturally exists
ns in their lowest
.
in atoms
80
d
p
s
up to 2 electrons
The simplest atom is hydrogen.1.4
With
atomic number 1, an atom of
p
The arrangement of the electrons
2
2
s
hydrogen contains one proton and hence one electron. The next atom,
helium, has atomic number 2. Therefore, there are two protons in its
nucleus. The first shell can accommodate two electrons only and it is 1
s
Electron shells 1
Study
therefore completely filled. An atom
oftip
lithium has three protons in its
Themain
firstshells
shell, which issub-shells
closest to the nucleus
electron shells
nucleus. The third electron
is in aYounew,
shell.
mustlarger
remember
the number of
second, and so on. The number of electrons i
n is the number of the shell, so:
Table 2.2 shows▲the
arrangement
the electrons for the first
three
Figure
1 Electronof
shells
▲ Figure
2 Main shells and sub-shells
These diagrams are useful as a basic representation of atomic structure and help us
elements.
• the first shell holds up to two electrons
1.4 The arrangement
of the
electrons
to understand
the nature of chemical
bonding.
Table 2.2 Electron arrangement for the ground state of the first three elements.• the second shell holds up to eight electron
electrons in each shell.
Quantum mechanics
• the third shell holds up to 18 electrons.
a more
complete
of the electrons
in atoms a theory called
1 description
2
3
Atomic number/For
number
of protons
C which was developed
Electron
diagrams
quantum mechanics
during
the 1920s. This
H is used,
He
Li
Electron
shells
If
you
know
the
number of protons in an ato
Study tip describes the atom mathematically with an equation (the Schrödinger
Arrangement of electrons
number
of
electrons
it has. This
because
th
The
first
shell, which
is closest
to the
fillsis first,
then
equation).
The solutions
to this
equation
give the
probability
ofnucleus,
finding an
You must remember
the number
of
can
therefore
draw
an
electron
diagram
for
carbon
(2,4) and so on. The number of electrons in each shell =a2
second,
electron
in
a
given
volume
of
space
called an atomic
orbital.
electrons in each shell.
▲ Figure 1 Electron
diagram
of carbon
n is the
number
of thecarbon
shell,has
so:six electrons. The four electrons in
Symbol
Atomic orbitals
usually drawn spaced out around the atom (F
• the first shell holdsSulfur
up tohas
two
electronsIt has six electrons in it
16 electrons.
whenup
drawing
diagrams
to space out t
• considered
the secondto
shell
to
eight
electrons
The electron is no longer
be aholds
particle
but
abonding
cloud
of
and then add the next two electrons to form pa
• thefills
third
shell holds
up to called
18 electrons.
negative charge. An electron
a volume
in space
its atomic
You
can
also
draw
orbital. The concept of theS main shells and the sub-shells iselectron
then diagrams of ions,
C
30/03/15
PM
Electron
diagrams
the number
of 2:42
electrons.
For example, a sodi
included in the following way.
electrons,
but its ion
hasatom,
10, so ityou
has also
a positiv
If you know the number
of protons
in an
kn
• Different atomic orbitals
have
different
energies.
Each
orbital
has
number of electrons itAn
has.
Thisatom
is because
atom is
neutr
oxygen
has eightthe
electrons,
but
its io
2– (Figure
a
number
that
tells
us
the
main
energy
shell
that
it
corresponds
can
therefore
draw
an
electron
diagram
for
any
element.
Fo
negative
charge,
O
4).
carbon (2,4)
sulfur
(2,8,6)
to: 1, 2, 3, and so on.carbon has six electrons.
four electron
electrons
in theinouter
she
You The
can write
diagrams
shorthand
▲ Figure 1 Electron diagram of carbon
▲
Figure
2
Electron
diagram
of
sulfur
usually
drawn
spaced
out
around
the
atom
(Figure
1).
• The atomic orbitals of each main shell have
different shapes, which
• write the number of electrons in each
she
▲ Figu
in turn have slightly Sulfur
different
energies.
These
are
the
sub-shells.
shell
and
working
outwards
has 16 electrons. It has six electrons in its outer shell.
It
d-orbit
They are described by
the letters
s, bonding
p, d, and
f.
• diagrams
separate
each
number
a comma.
+
when
drawing
to space
outbythe
first four (as
Bilal Hameed
Bilal Hameed
Marginalizer
you
write
2,4;Structure
for sulfur
2,8,6;
for
Electronic
and3then add the next For
twocarbon
electrons
to form
pairs
(Figure
2).
Na
S
You can also draw electron diagrams of ions, as long as you
The maximum number of electrons each shell (main energy level) can hold is given by the
expression 2n2. Hence the first, second, third and fourth shells can hold up to a maximum of 2,
C
8, 18 and 32 electrons.
Chemists often use a shorthand notation to describe the arrangement of electrons in shells.
It indicates the number of electrons in each shell without drawing the shells. It is known as the
2
electron
arrangement.
Hydrogenof
hasthe
an electron
arrangement of 1; lithium has an electron
1.4 The
arrangement
electrons
arrangement 2,1 or 2.1 and sodium has an electron arrangementcarbon
of 2,8,1 or
2.8.1. Table 2.4 lists
(2,4)
electron arrangements for the first 20 chemical elements; Figure 2.49 shows the shell structures
ELECTRONIC
STRUCTURE 2
for selected elements.
▲ Figure 1 Electron diagram of carbon
■ Table 2.4 Electron
arrangements for
the first 20 chemical
Study Atomic
tip
Electron
shells
Energy shell
Atomic
Energy shell
Electron dia
If you know th
number of ele
can therefore
carbon has six
usually drawn
has 16 fi
e
1st 2nd 3rd 4th
numbershell,
The first
which is closest toSulfur
the nucleus,
when
drawing
You must remember
the number of Sodiumsecond,
Hydrogen
1
1
11
elements
and2 so8 on.1 The number of
electrons
in
Helium
2 shell.
Magnesium
12
2
8
2
a 2 each
b
electrons in
and
then
add t
CHEM COMPLEMENT
n
is
the
number
of
the
shell,
so:
Lithium
3
2
1
Aluminium
13
2
8
3
Element
number 1st 2nd 3rd
4th
Element
alt
Why K, L, M, and not
A, B, C?
8
4
can also d
• the14first2shell
holds
up to twoYou
electrons
Beryllium
4
2
2
Silicon
Boron
5
2
3
Phosphorus
Charles G. Barkla was a
Carbon
spectroscopist who studied the
Nitrogen
X-rays emitted by atoms and
Oxygen
found that there appeared to be
Fluorine
two types, which he originally
6
2
4
Sulfur
7
2
5
Chlorine
8
2
6
Argon
9
2
7
Potassium
named A and B. Later, he Neon
renamed them K and L, to leave
room
for the
■
Figure
2.49possibility that the
K type was
not the highest
Electron
arrangements
energy
X-ray an
atom can emit. H
of
hydrogen,
lithium,
We now know that this is the
sodium, argon and
hydrogen
highest energy X-ray, produced
potassium, shown as
when an electron in the
shell structures
innermost shell is knocked out
and then recaptured. The
innermost shell is therefore
called the K shell. Barkla won
the 1971 Nobel Prize for Physics.
10
15
S2
8
5
the number o
• the16second
2
8shell
6 holds up to eight electrons
electrons, but
8
7
• the17third2 shell
holds
up to 18 electrons.
18
2
8
8
An oxygen ato
negative charg
Calcium
20
2
8
8
2
phosphorus
If you know the number of protons in an atom
sulfur (2,8,6)
proton (+)
neutron (no charge)
electron (–)
number
of electrons it has. This is
because
the a
You
can write
▲
Figure
2
Electron
diagram
of
sulfur
can
therefore
draw
an
electron
diagram
for
any
carbon (2,4)
•
write
the
Figure 1.3.5
Li Bohr atomic model diagrams of (a) carbon and (b) phosphorus.
carbon
has six electrons. The four electrons innt
Na carbon
Ar
▲ Figure 1 Electron diagram of
shell
and
w
usually drawn spacedK out around the
atom
(Fig
lithium
Atoms may also be represented diagrammatically. The Bohr model of the atom
• separate ea
+
can be shown in full detail with numbers of
protons,has
neutrons
and electrons
Sulfur
16 electrons.
It has six electrons in its o
sodium
argon
fully labelled.
carbon
when drawing bonding
diagrams For
to space
out yo
the
The electron arrangement of an ion will be different from that of potassium
the atom from
and
then
add
the
next
two
electrons
to
form
pair
Na
which it was formed, because an ion is an atom that has lost or gained electrons.
C
8
1
Electron diagrams
1
2
8
carbon
19
2
8
Positive ions are atoms that have lost electrons and negative ions are atoms
When an element
forms
an ion,
the electrons are also
accommodated
in theelectron diagrams of ions, as
You
can also draw
that have
gained
electrons.
S
appropriate shells.
the number of electrons. For example, a sodium
Summary
TABLE 1.3.3 ELECTRON ARRANGEMENTS OF SOME ELEMENTS AND THEIR IONS
+
Element and
ion name
829055_02_IB_Chemistry_052-084.indd 73
Symbol of ion
Atomic number
Nitride ion
N3−
7
Oxide ion
O2−
electrons, but its ion has 10, so it has a positive
Na
sodium
ion
1 Draw the e
Charge
on ion
Electron
arrangement
11
protons,
10
electrons
An oxygen atom
has eight electrons,
but its ion
18/05/15 9:26 amfollowing n
(2,8) 3−
2,8
2–
negative charge, O (Figure 4).
a 3
2−
Electron
diagram2,8
of a
diagrams in
Na
Sodium ion
11
2,8
1+ write electron
2 shorthand:
State, in sh
sodium You
ion can
▲
Figure
2
Electron
diagram
of
sulfur
2+
Ca
Calcium ion
20
2,8,8
2+
• write the number of electrons in each
a 4shell,
elect
shell and working outwards
3 Identify wh
2.3.2
Evidence for the Bohr
model: line spectra
• separate each2–number by a comma.
+
Distinguish between a
ions. Give t
continuous spectrum and a line
Experimental evidence for Bohr’s model came from studies of the emission
For
carbon
you
write
2,4;
for
sulfur
2,8,6;
forTa
N
spectrum. © IBO 2007
spectra of atoms. These spectra are the emissions of light from atoms that
Periodic
8 Figure 3
▲
sulfur (2,8,6)
+
2.3.3
Explain how the lines in the
emission spectrum of hydrogen
are related to electron energy
levels. © IBO 2007
16
have beenNa
provided with energy such as heat, light or O
electricity. The bright
×
colours of fireworks are the result of such emissions.
Summary questions
N
Bohr explained emission spectra by suggesting that if atoms are subjected to
A
large amounts of energy from heat, light or electricity, the electrons can change
× from the nucleus
energy levels. The electrons jump to energy levels further
B
than
they would usually occupy. The atom is said to be in an excited state
+ sodium
2– Draw
Na
ion
1
the
electron
arrangement
diagrams
of ato
O
oxygen
ion
when this happens. When the electrons return to the ground state this extra
C
11 protons, 10 electrons
8 protons,
energy is released in the form of light. The electrons
make specifi
c jumps,
following
numbers
of
electrons:
(2,8)
depending on the energy levels involved, therefore
the light(2,8)
released has a
10 electrons
D
3
b 9 looksc 14
specific wavelength. The emitted light, a line (oraemission)
spectrum,
▲ Figurelike
3 Electron
diagram
of
a
a series of coloured lines on ▲
a black
background.
Somediagram
of the emissions
E
Figure
4 Electron
of an
PRAC 1.2
State,
in naked
shorthand,
sodium ion
may be radiation of a wavelength that is not2visible
to the
eye. Thethe electron arrangements
Flame tests and emission spectra
oxygen
ioncalled emission spectroscopy.
study of this light emitted from the
atom is
2–
O
×
×
Marginalizer
Electronic Structure
a 4 electrons
b 13 electrons
c
22 3 Identify which of the following are atoms, positi
ions. Give the size of the charge on each ion, inc
Periodic Table to identify the elements A–E.
O2– oxygen ion
8 protons,
10 electrons (2,8)
4
▲ Figure 4 Electron diagram of an
oxygen ion
A
B
C
D
E
Number of protons Number of electro
12
10
2
2
17
18
Bilal Hameed
10
10
Bilal Hameed
3
2
3
ELECTRONIC STRUCTURE 3
Ionisation energy
It is possible to obtain information about the arrangement of electrons in atoms by
studying the ease with which atoms lose electrons.
alt
Electrons can be removed from atoms and the energy it takes to remove them can be
measured. This is called ionisation energy because as the electrons are removed, the
atoms become positive ions.
The energy needed to remove one electron from each atom in a mole of gaseous
atoms is known as the first ionisation energy.
X (g)
1➝ X (g)
+
+ e—
First ionisation energy is the energy required to remove a mole of electrons from a
mole of atoms in the gaseous state, and is measured in kJ mol —1.
All ionisation energy values will be positive as they are endothermic as energy is
required to remove an electron from the attractive power of the nucleus. Energy has
to be absorbed and work done so that the negatively charged electron can be
removed from the influence of the positively charged nucleus.
Ionisation energies
electron removed
–
energy needed =
ionisation energy
–
–
–
–
–
The first ionisation energy is the energy required to
electrons from one mole of gaseous atoms to form one
ions.
–
+
–
Successive ionisations give the first, second, third, fou
energies. Only one mole of electrons is removed with
–
–
–
–
For example for sodium
–
Figure 1.21 What element is being
The first ionisation energy
of sodium
is the energy required for:
ionised
in this diagram?
Na (g)
➝
Na+ (g)
+ e—
➝
O+ (g)
+ e—
The first ionisation of sodium is represented by the eq
Na(g) → Na+(g) + e−
∆H = +494 kJ mol-1
The second ionisation of sodium is represented by the
For oxygen, it is the energy change per mole for:
O (g)
The electrons in atoms and ions are attracted to the po
is required to overcome this attraction and remove ele
removing electrons from atoms and ions is called ioni
energy is the energy required to remove electrons.
Na+(g) → Na2+(g) + e−
-1
∆H = +1314 kJ molThe
third ionisation of sodium is represented by the e
This means that a mole of gaseous positive ions is formed, regardless of whether
Na2+(g) the
→ Na3+(g) + e−
element is a metal or a non-metal.
The outer-shell electron is the most easily removed, and so first ionisation energy is a
TIP
measure of how tightly the outer-shell electrons are held in an atom.
Bilal Hameed
Bilal Hameed
STRUCTURE
Equations for ionisation energies are often asked for.
state symbols. The atoms and ions must be in the ga
Marginalizer
5
Electronic Structure
corresponds to the first ionization energy.
(6.63 × 10 −34 J s × 3.00 × 108 m s−1) periods account
for the existence of
= 4.9725 × 10 −17 J
For one photon:Trends
E = hc =in first ionization energy across
λ
× 10 −9
m
main energy levels4and
sub-levels
in atoms.
1
16 photons
Successive
ionization
energy
data for an element give information that shows
= 2.01 × 10
So for one joule:
4.9725 × 10 −17
4
relations to electron configurations.
electromagnetic
spectrum it belongs
to.
8
Calculate the
frequency of
yellow light with
a wavelength of
5800 × 10 −8 cm.
9
The laser used to
read information
from a compact disc
has a wavelength of
780 nm. Calculate
the energy
associated with
one photon of this
radiation.
ELECTRONIC
4 energy
Ionization
ElectronSTRUCTURE
configuration
Ionization energye−
−
e
The first ionization energy
is the minimum energy per molee−required to remove electrons from
one mole of isolated gaseous atoms to
form one
mole of gaseous unipositive ions under standard
valence
electron
+3
+3 energy of chlorine is the energy
thermodynamic conditions.
For example, lost
the first ionization
e−
required to bring about −the reaction:
alt
e
Cl(g) → Cl+gaseous
(g) + e− lithium
gaseous lithium
ion, Li+(g)
atom, Li(g)
The electron is removed from the outer sub-shell (energy sub-level) of the chlorine atom (that
Figure
of firstofionization
is, a 3p electron).
Table12.1
12.1The
givesconcept
some examples
ionizations,energy
and in each case the ionization
In
general,
electrons
in
metals
are
easily
removed,
so
metals
have
low
energy, which is the enthalpy change for the equation. Ionization ionisation
energies are listed Table 8 of
Ionization
energies strongly,
may be and
measured
inside a mass spectrometer,
energies.
Non-metals
electrons
so haveexperimentally
high ionisation energies.
the IB Chemistry
datahold
booklet.
Table 12.1 Selected
ionization energies
Element
Oxygen
Sulfur
Copper
1
which vaporizes substances and then fires high-speed electrons at them to cause
ionization.
Ionization equation
First ionization energy/kJ mol−1
O(g) → O+(g) + e −
1314
Factors that influence ionization energy
S(g) → S+(g) + e −
Cu(g) → Cu+(g) + e −
1000
745
The size of the nuclear charge
Factors thationisation
affect ionization
energy
Successive
energies
As the atomic
number (number of protons) increases, the nuclear charge
Values
of ionization
energies
on
the following
factors: of energy required
the
successive
ionisation
energies
of anthe
element
are the
amounts
increases.
Thedepend
larger
positive
charge,
the greater the attractive electrostatic
the size
the
atom
(or
ion)
to remove
allof
the
electrons
from
one
mole
of
an
element
in
the
gaseous
state,
one amount of energy is
force between the nucleus and all the electrons. So,
a larger
molethe
of electrons
at
a
time.
nuclearneeded
charge to overcome these attractive forces if an electron is to be removed (during
the shielding
effect. As the proton number increases across a row of the periodic table the
ionization).
For example for beryllium:
ionization energy tends to increase.
+ (g) +
Atomic
radius
e—
I1 = +900 kJ mol-1
Be (g) ➝ Be
As2+of
the
distance
of the
from nucleus
the nucleus increases,
from
Be+ Distance
(g) ➝ Be
(g) outer
+ e— electrons
I1 = outer
+1760electrons
kJ mol-1 the
the attraction of the positive nucleus for the negatively charged
–
The
electrostatic
between
positive
and negative charges
2+ (g)force
3+ (g) +falls.
-1
Be
➝of Be
e—attraction
= +14
800 ionization
kJ mol
electrons
ThisI1 causes
the
energy to decrease. Hence,
decreases rapidly
as theenergy
distance
between
them
increases.
Hence,
ionization
decreases
as the
atomic
or ionic
radius electrons
increases.
-1
Be
(g) ➝
Be4+energy
(g) + levels)
e— further
I1 = +21
000 from
kJ molthe
in3+shells
(main
away
nucleus are more weakly
3+
–
Nuclear
attracted to the
nucleuscharge
than those closer to the nucleus. The further the outer
nuclear
When
the
nuclear
charge
positiveenergy.
(due to Ionization
the
electron
shell
is
from
the
nucleus,
the becomes
lower themore
ionization
pull
Electrons in energy levels
that are
furtheroffrom
the nucleus
have higher
energy than
presence
additional
protons),
its
attraction
on
all
the
electrons
energies tend to decrease down a group of the periodic table.
those that are closer to the nucleus;
therefore,
is these
electrons
thatenergy
can beto increase.
increases.
This itcauses
the
ionization
–
removed more easily by the addition of energy.
Shielding
effecteffect
Shielding
The first electrons
to be
the
of energy
will be
those
thateach
already
Thebyouter
or valence
electrons
are
repelled
by
allother.
the other
electrons
Since
allremoved
electrons
are addition
negatively
charged,
they
repel
Electrons
in
repulsion
from
have the highest
energy—those
in
the
valence
electron
shell.
This
can
also
be
in
the
atom
in
addition
to
being
attracted
by
the
positively
charged
inner
shell
of
full inner shells repel electrons in outer shells. The full inner shells of electrons
electrons
(’shielding’)
nucleus.
The
outer electrons
are shielded
of This
explained
in terms
of thethe
attraction
of the
electrons
to the nucleus.
As electrons
theattraction
prevent
full
nuclear
charge
being
experienced
byfrom
theinthe
outer
electrons.
Figure 12.13 Electrostatic
forces
operating
on
the
the
nucleus
by
the
shielding
effect
(an
effect
of
electron–electron
valence shell are
from the nucleus,
are The
not attracted
it as
strongly of
as outer electrons by
is furthest
called shielding
(Figurethey
12.2).
greater to
the
shielding
outer or valence electron in a lithium atom
repulsion)
(Figure
12.13).
electrons in other
shells, andshells,
so theythe
willlower
be thethe
ones
that are removed
first. forces between the
theelectron
inner electron
electrostatic
attractive
829055_12_IB_Chemistry_435-450.indd 442
The first electron to be removed is one that already has a high energy and is least
strongly attracted to the nucleus; that is, one in the valence shell. This will require the
lowest ionisation energy. If there are more electrons in this outer shell, then they will
be removed next, with the ionisation energy gradually increasing.
The successive ionisation energies of an element get bigger and bigger. This is not
899713_12_IBDip Chem_RG_164-170.indd 164
surprising because, having removed one electron, it is more difficult to remove a
second electron from the positive ion formed.
Marginalizer
Electronic Structure
Bilal Hameed
6
Bilal Hameed
18/05/15 10
The first electron to be removed is one that
5 already has a high energy and is
least strongly attracted to the nucleus; that is, one in the valence shell. This
will require the lowest ionization energy. If there are more electrons in this
STRUCTURE
5
outerELECTRONIC
shell, then they
will be removed
next, with the ionization energy
gradually
increasing.
outermost
electron
shelltoisbeempty,
next
When the
outermost When
electronthe
shell
is empty, the
next electron
lost willthe
come
electron
to
be
lost
will
come
from
the
next
closest
shell
to
the
nucleus.
from the next closest shell to the nucleus. But since this shell is full, and thereforeBut
since this shell is full, and therefore stable, a great deal more energy will be
stable, and closer to the nucleus, a great deal more energy will be required to remove
required to remove one electron from this shell than was needed to remove the
one electron from this shell than was needed to remove the previous electron.
previous electron.
alt
The The
pattern
formed
successiveionisation
ionization
energies
of anprovides
atom provides
pattern
formedby
by the
the successive
energies
of an atom
evidence
for for
thethe
existence
levels
around
the nucleus
andus
allows
evidence
existence of
of energy
energy levels
around
the nucleus
and allows
to workus to
workout
out
electron
configuration
thethe
electron
configuration
of an atom..of an atom.
Consider the successive ionization energies for nitrogen and magnesium below:
Successive ionisation energies for nitrogen
The first two successive ionization energies of magnesium
Consider the successive ionisation energies for nitrogen:
1
TABLE 1.1.1 SUCCESSIVE IONIZATION ENERGIES
to 1451 kJ mol 1. The next ionization energy
736 kJFOR
molNITROGEN
1
7732.6
kJ(kJ
mol
that the 3rd electron is being
Ionization
Ionization
energy
mol,–1indicating
)
next
level. This energy level is full and closer to th
I1 energy
1400
N(g)
N (g) e
greater
amount
of energy is required for the 3rd electron
I2
2856
N (g)
N2 (g) e
for
the
2nd
electron.
The successive ionization energies inc
I3
4578
N3 (g) e
N2 (g)
electrons are removed from the second electron shell. The
I4
7475
N4 (g) e
N3 (g)
jump
from
35 461 kJ mol 1 for the 10th ionization energy
4
5
Figure 1.1.2 An atom of nitrogen
I5 the 11th
9440
N (g) e
N (g)
for
ionization energy, indicating
that this electro
has 5 electrons in its 2nd electron
5
6
I6 the53
266 energy level, the inner
N (g) e
N (g)
from
next
electron
shell and 2 electrons inshell.
the This le
6
7
1st electron shell.
configuration
I7
64 358 of 2,8,2 for magnesium.
N (g) e
N (g)
1
The pattern formed by successive ionization energies can
see that
the first1.1.1
five ionisation
gradually
increase
from 1400
to
graph
ofenergies
electron
being
removed against ionization energ
You You
willwill
notice
in table
that theenergies
first five
ionization
gradually
—1. The sixth requires
1
1kJ mol—1. This large increase
about
53000
9440
kJ
mol
Figure
1.1.3
shows energy
the successive
ionization energies of n
the 6th
ionization
is
increase from 1400 kJ mol to 9440 kJ mol , then
1
indicates
thatkJ
themol
6th electron
beingincrease
removed form
a shell closer
to the
nucleus.
A to energies for the first 5 electron
. Thisislarge
indicates
that
the
6th
electron
equal
to 53 266
figure
1.1.3,
the
ionization
much greater
amountfrom
of energy
needed
to remove
one
electron
from
this
shell
than
be removed
is coming
the is
next
energy
level,that
which
is full
and
closer
to the in the outer shell, then there
there
are
five
electrons
nucleus.
A much
greater
amount
energy are
is needed
to
remove
one
electron
from the
previous
shell because
theofelectrons
experiencing
a
significantly
larger
the ionization energies for the 6th and 7th electrons, whi
fromattraction
this new
level than from the previous energy level because the
to energy
the nucleus.
electron shell.
electrons are experiencing a significantly larger attraction to the nucleus. This
to electron
the configuration
of 2,7. of 2,5 for nitrogen.
leadsThis
us leads
to the
configuration
6.00
70000
60000
TABLE 1.1.2
SUCCESSIVE IONIZATION ENERGIES FOR MAGNESIUM
Mg3 (g)
Mg2 (g) 10000
0
Ionization energy (kJ mol
I1
736
2nd
energy
I level 1451
2
61st
Ionization
7
(g)
Mgenergy
Mg (g)
e
3.00
I7
21 703
Mg7 (g)
Mg8 (g) 2.00
e
I8
25 656
I
31 642
I
35 461
I
189 363
level
e
I3
7733
Mg8 (g)
Mg9 (g) 1.00
e
e
I4
10 540
Mg9 (g)
Mg10 (g)
0.00
Mg4 (g)
Mg4 (g)
Mg5 (g)
1
2
3
4
5
6
7
10
number
of
the
electron
being
removed
I5
13 630
e
Mg11 (g)
Mg (g)
Mg5 (g)
Mg6 (g)
e
I
17 995
Figure 1.1.3 The successive ionization energies6 of nitrogen.
Mg11 (g)
Mg12 (g)
e
e
2nd
Ionization energy (kJ mol
4.00
Mg3 (g)
0
3rd energy level
5.00
10
50000
Ionization
Mg(g) 40000
Mg (g) e
30000
Mg2 (g) e
Mg (g)
20000
log (ionization energy)
–1
ionization energy (kJ mol )
A similar trend can be seen for the successive ionization energies of magnesium.
9
4th energy level
0
2
4 10 6
8 10 1
number
of
the
b
I11 169 electron
987
12
Figure 1.1.4e The successive ionization
energies of potass
log10(ionization energy), to make the increases more ob
CHEMISTRY: FOR USE WITH THE IB DIPLOMA PROGRAMME
Bilal Hameed
Bilal Hameed
Evidence for the four electron shells of potassium can be
this graph there are three sharp increases in ionization e
Marginalizer
after 1 electron has
been removed and the second and thi
8 electrons
have been removed.Electronic
This leads
to the electron
7
Structure
2,8,8,1 for potassium.
6
ELECTRONIC STRUCTURE 6
Successive Ionisation enegeries are more endothermic
The second ionisation energy is always higher than the first, and this can be
explained in two ways:
alt
1. Once an electron has been removed from an atom, a positive ion is
created.A positive ion attracts a negatively charged electron more strongly than a
neutral atom does. More energy is therefore required to remove the electron from
a positive ion.
2. 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). If they are closer to the nucleus, they are more strongly attracted and
more difficult to remove.
1
The second ionisation ener
this can be explained in two w
1 Once an electron has been r
created. A positive ion attrac
strongly than a neutral atom
remove the electron from a p
2 Once an electron has been r
repulsion between the remai
in closer to the nucleus (Figu
they are more strongly attrac
This graph shows the energy re
from a gaseous potassium atom
Marginalizer
Electronic Structure
A log scale is being used here to allow all
the data to be plotted on one graph, but
Bilal
although on one level this has made
theHameed
data
easier8to interpret and supported the explanations
that
Bilal Hameed
have been given, it has also distorted the data. The
difference between the first and second ionisation
−1
data are presen
their theories?
scale on a grap
particular tren
or is it a matte
The argument
the reaction so the energy change is given a positive sign.
canelectron
be removed
from atoms and the energy it takes to remov
oms. According
to
ionisation
energies
a)suggests
down that theyElectrons
evidence
for➔
electron
sub-shells.
evidence
for
thedefinition
data.
Thisof
are very close
to
the sub-shells.
State the
nisation
energies
them
can beusing
measured.
This is called
each electron in an
Scientists
can
also
determine
energies
a spectroscope
to ionisation energy because as th
aionisation
groupnucleus.
andenergy.
b) across
a ionisation
Th
ere must
be a very great
force by
of attraction
electrons
are
removed,
the atoms become positive ions.
te
energy.
When
cessive
ionisation
the light
given
by atoms
whenelectrons
heated inand
a flame
(as
period
in terms
of out
electron
Ionisation➔study
energy
Ionisation
energy
between
the
nucleus
and these
there
arein a flame test).
Describe
the
trend
in
e energy, the
electronsfor Theconfigurations.
electrons
removed
spectroscope
shows up to
a series of
bright
lines
(Figure
1.18).
therequired to remove a mole of electron
7them.
•Electrons
energy
the from
energy
no inner
Th
e large
inis Heating
Electrons
can be ionisation
removed
fromelectrons
atoms
andshield
the energy
itIonisation
takes
can
toincrease
remove
be removed
atoms and the energy it takes to remov
energies
a) down
ergy
level to
another.
atoms gives them
energy
which makes some of the electrons jump to higher
the base
10)
is used
from
a be
mole
atomsThis
in the
gaseous
state, andenergy
is measured
in kJasmo
ionisation
energy
between theenergy
9th
and
10th
electrons
them can be ➔
measured.
This
is across
called
them
because
can
measured.
asofthe
is called
ionisation
because
th
Explain
how
trends
in ionisation
a group
and b)
energy
levels.
Each
line ain the spectrum arises from the energy given out as
ionisationelectrons
energies have
confi
rms
that
the
10th
electron
is
in
a
shell
closer
to
the
areELECTRONIC
removed,
the
atoms
become
positive
ions.
electrons
are
removed,
the
atoms
become
positive
ions.
•
Ionisation
energy
has
the
abbreviation
IE.
ionisation
energies
STRUCTURE
in terms
electron
the period
electrons
drop of
back
from7a higher energy level to a lower level.
nucleus
than
the
9th electron.
provide
evidence
for
the
configurations.
• Ionisation
energy
is the
energy required
tofor
remove
• aIonisation
mole of electrons
energy
is the energy
required
ionisation
energies
sodium
about sodium
from Successive
Removing
the
electrons
one by
one to remove a mole of electron
existence
of
electron
shells
–1.
from
a
mole
of
atoms
in
the
gaseous
state,
and
is
measured
from
a
mole
in
kJ
mol
of
atoms
in
the
gaseous
state,
and is
measured
kJ mo
Explain
trends
trons at fixed or
A➔
similar
trend how
can be
seen in
for the successive ionisation
of the
sodium.
You canenergies
measure
energies required to
remove
the
electronsinone
by
and
sub-shells.
sometimes
called
•
Ionisation
energy
has
the
abbreviation
IE.
•
Ionisation
energy
has
the
abbreviation
IE.
ionisation
energies
one
from
an
atom,
starting
from
the
outer
electrons
and
working
inwar
s a low 1st ionisation
Specification
reference:
The of
word
provide evidence
for3.1.1
the
es.rest
the‘quantum’
data. It is very
1 electron;
•Removing
The first
electron
needs the
least
Removing
the
electrons
one
byshells
one
the electrons
one
by energy
one to remove it because it is
be
something
related
existence
of
electron
. It is therefore likely to be
very
easily
being
removed
from
a neutral
atom.
is the
IE. one by
You can measure the energies required to remove the
You
electrons
can measure
one bythe
energies
required
to This
remove
thefirst
electrons
or awell
fixedshielded
level.
removed
nd
by inner and sub-shells.
2 electrons;
alt
1
log IE
log IE
one from an atom, starting from the outer electrons •one
andThe
from
working
an atom,
inwards.
starting
frommore
the outer
electrons
andfirst
working
inwar
second
electron
needs
energy
than the
because
it i
difficult electrons
Specificationvery
reference:
3.1.1 in 11+
8
electrons;
being
removed
from
a
+1
ion.
This
is
the
second
IE.
Figure 1.18
The
line
ofshell
hydrogen
in the
region
of the
electromagnetic
to remove
• The first electron
needs
thespectrum
least
energy
• visible
The
it because
first less
electron
it is
needs the least energy to remove it because it is
main
1 to remove
easily
•
The
third
electron
needs
even more
spectrum.
being removed from aelectrons
neutralin atom. This is the first
being
IE. removed
a neutral
atom.energy
This is to
theremove
first IE.it because
removed from
is
being
removed
from
a
+2
ion.
This
is
the
third
IE.
main shell
2 energy than the
• The second electron needs
more
• first
The because
second electron
it is
needs more energy than the first because it i
Using from
data from
spectra,
measure
the needs
energy
required
to This
n
electrons
inpossible
•to IE.
The
fourth
yet
and
so on.
being removed
a +1 ion.
Thisitisisthe
second
being
removed
from
amore,
+1 ion.
is the second IE.
nucleus
1
main
remove electrons from
ions shell
with1increasing charges. A succession of ionisation
• The third electron
evenFor
more
energy to remove
• Theare
third
it called
because
electron
it needs even
more energy
to remove it because
These
successive
ionisation
energies.
electronsisneeds
inobtained.
energies
example:
electrons in
is being removed
from
a
+2
ion.
This
is
the
third
IE.
is
being
removed
from
a
+2
ion.
This
is
the
third
IE.
main shell
3
Figuremain
3.5
The
of electrons For
in an atom
of sodium
can be −1
shellarrangement
2−
+(g)
example,
sodium:
+
e
first
ionisation
energy
=
+496
kJ
mol
Na(g)
→
Na
values
of
successive
ionisation
energies.
1 2 deduced
3 4more,
5from
6 the
7
8 so
9 10
11
• The fourth 0needs
yet
and
on.
• The
fourth needs yet more, and so on.
+
2+
−
+(g)−1 + e−
total
number
of
electrons
removed
Na (g) → Na (g) + e second ionisation energy
= +4563 Na
kJ mol
Na(g)
first IE
= + 496 kJ mol−1
These are calledelectrons
successive
ionisation energies.These are
called ➝
successive
ionisation
energies.
in
2+(g)
▲ Figure
1(g)The
successive ionisation
−1 + e−
Na+=
(g)+ 6913
second IE = + 4563 kJ mol−1
➝ kJNa
Na2+shell
→
mol
main
3 Na 3+(g) + e− third ionisation energy
For example,energies
sodium:
For
example,
sodium:
of sodium against number of
1
0 1 2 3 Check-up
5 6 7 8 9 10 11
Na2+(g) ➝ Na3+(g) + e− third IE
= + 6913 kJ mol−1
There
are
114 electrons
a sodium
electrons
removed.
Note thatinthe
log of theatom so there are 11 successive ionisation
−1to support
+(g) + e−
d
total
of +
electrons
removed
successive
ionisation
for
sodium
evidence
Na(g)The
Na
Na(g)
e−eenergies
first
= +provide
496
mol
first IE
= + 496 kJ mol−1
➝ number
➝ Na
energies
for+3(g)
this
ionisation
energy
has
been
plotted
inIE
order
and
sokJ
on,
see
1. the
a element.
Th
successive
ionisation
energies
of Table
2+(g)
− atom
+mol
−1
2+(g) the
theory
that
electrons
invalues
are
arranged
in+a4563
seriesNa
of levels
or shells
around
n
▲fit
Figure
1Na
The
successive
ionisation
Na+(g)
+
ean
second
IEfor
(g)
Na
+This
e− second IE = + 4563 kJ mol−1
➝
to
large
range
of
on
the
scale
boron
are
shown
in=an
Table
3.3.kJget
Thethe
successive
ionisation
energies
element
bigger➝and
bigger.
f
energies of sodium
against
number of
nucleus.
−
Na2+(g)is not
+ e− third
IEremoved
= + 6913
Na
kJ 2+
mol
(g)−1
Na3+
(g) + eto
third IE
= + 6913 kJ mol−1
➝ Na3+(g) because,
having
one electron,
it is➝more
difficult
▼
Tablesurprising
1removed.
Successive
sodium
of the
electrons
Noteionisation
that the
logenergies
of
the of3rd
Ionisation
1st
2nd
4th
5th
remove a second electron from the positive ion formed.
order
ionisation
has been plotted in order
and so on, see
Table energy
1.
and so on,
Table
3rd
4th see
5th 1. 6th
7th
8th
9th
10th
11t
Electron
removed
2420 3660
25 000
32 800
Ionisation1st 799 2nd
scale
toThe
fit the
large range
of values
the scale evidence to support the theory that the
graph
in Figure
1.19onprovides
energy
Ionisation energy
/ / are arranged in a series of levels or shells around the
electrons
in–1an atom
4563
6913
9544 13 352 16 611 20 115 25 491 28 934 141 367 159 0
−1 496
6 7 of8 sodium
9 10 11 ▼ Table
nergies
1
Successive
ionisation energies of sodium
kJ
mol
kJ mol
nucleus.
ectrons removed
2nd
3rd
4th Electron
5th removed
6th 3.3 Successive
1st
7th ionisation
2nd
8th energies
3rd
9thof boron.
10th
4th
11th
5th
6th
7th
8th
9th
10th
11t
Table
Ionisation
energy /
Tip
Notice
that
the159
second
is not 20
the115
energy
change
continued
13 352 16
611 20
496
115 25
4563
491 28
6913
934 141
9544
367
13
352
079 IE
16 611
25 491
28 for
934 141 367 159 0
Study
tip
kJ mol–1
Logarithms reduce the range of numbers that vary over several
orders
of
magnitude.
2+
−
Na(g) ➝ Na (g) + 2e
The
shape
thelogarithms
graph in which
Figurework
1 like this: log 10 = 1, log 100 = 2, log 1000 = 3
Figure
1.19of
uses
for this
process
would be (first IE + second IE).
andare
on.electrons
A calculator
be used
to so
findthere
theThe
values
the logarithms
(log) of other
There
11
in a can
sodium
atom,
areenergy
11 of
successive
ionisation
has
tosobe
thought
about
carefully.
Notice that the second IE is not the energy changeNotice
for
that the second IE is not the energy change for
numbers.
energies
forelectron
this element.
The
first
removed is in the
Study
tip
If you plot a graph2+
of the values shown in Table 1 you get Figure 1.
2+(g) + 2e−
Na(g) ➝
Na
Na(g) ➝ Na (g) + 2e−
outer main shell and the 10th and
1
Theisshape
of the graph
in Figure
1
There
a big difference
between
ionisation
energies.
For sodiumisthe
first
Notice
that
one electron
relatively
easy to remove, then comes a
11th
electrons
removed
aresome
in thesuccessive
The
energy
for
this
process
would
be
(first
second
The energies.
energy
IE). for this process would be (first IE + second IE).
y.
to be thought
about carefully.
2nd+ionisation
bighas
difference
occurs between
the 1st andIE
group
of
eight
that
are
more
difficult
to remove, and finally two that
innermost main shell.
he structure
The first
mic
and the Periodic
Tableelectron removed is in the
are
very
difficult
to
remove.
If you
plot aThere
graph
of
the
values
shown
in
Table
1
If
you
you
get
plot
Figure
a
graph
1.
of
the
values
shown
in Table 1 you get Figure 1.
is a big jump in the value of the second ionisation energy. This suggests that the
d
outer main shell and the 10th and
electron isisinrelatively
a shell closer to theto
nucleus than
the
first
electron.
Notice that second
one
electron
Notice
then
that
comes
one aelectron is relatively easy to remove, then comes a
he
11th
electrons removed areeasy
in the remove,
26
group of eight
that
are
more
difficult
to
remove,
and
group
finally
of
eight
two
that
areelectron
more difficult
to remove, and finally two that
st and
nd ionisation
Taken
together,
the
1
2
energies
suggest
that
sodium
has one
in
innermost main shell.
are very difficult
to
remove.
are
very
difficult
to
remove.
24
13/04/19
10:15
PM
its outer shell.
f563
ionisation6913
energy of sodium
9544
d.
From the second to the ninth electrons removed there is only a gradual change in
26
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.
Bilal Hameed
Bilal Hameed
Marginalizer
9
Electronic Structure
Na2+(g) → Na3+(g) + e–
Definitions
The first ionisation energy of an
element is the energy needed to
remove one electron from each atom
in one mole of gaseous atoms.
third ionisation energy, ∆Hi3 = +6913 kJ mol–1
There are 11 electrons in a sodium atom, so there are 11 successive ionisation
energies for this element.
Amountionisation
of substance
The successive
energies 8
of an element are all endothermic and
they get bigger and bigger. This is not surprising because, having removed one
A successive ionisation energy of
electron, it is more difficult to remove a second electron from the positive ion
an element measures the energy
formed.
needed to remove a second, third or ELECTRONIC STRUCTURE 8
The graph
in Figure
5.2 shows a logarithmic plot of the successive
mount fourth
of substance
Amount
of substance
electron and so on from one
ionisation
energies
of
sodium
against
thethe
number
of electrons
removed.
This
The
graph
below
shows
a
logarithmic
plot of
successive
ionisation
energies
of
mole of gaseous ions of the element
provides
evidence
to
support
the
theory
that
electrons
in
an
atom
are
with the appropriate positive charge. sodium against the number of electrons removed. The logarithmic plot allows an
arranged in a series of levels or shells around the nucleus. The logarithmic
extremely wide range of ionisation energies – from 496 kJ mol—1 to 159 080 kJ mo–1—
plot allows an extremely wide range of ionisation energies – from 496 kJ mol
1 – to be shown on the–1same graph.
Learning
to 159 080objectives:
kJ mol – to be shown on theThe
samepatterns
graph. in first ionisation energies across a perio
2
1.6 Electron arrangements and
ionisation energy
1.6 Electron
Electron arrangements
and arrangements and
ionisation energy ionisation
energy
evidence for electron sub-shells.
2
alt
State the definition of
ionisation energy.
Ionisation
energy
bjectives:
Learning
objectives:
Describe
the trend
in
first
ionisation
energies
across a period
The patterns
provide in first ionisation energies across a perio
Electrons
canelectron
be removed
from atoms and the energy it
ionisation
energies of
a) down
evidence for➔
electron
sub-shells.
evidence for
sub-shells.
efinition of
State the
definition
them
can
be
measured.
This
is called ionisation energ
aionisation
group andenergy.
b) across a
nergy.
electrons
are removed,
period in terms of electron
Ionisation➔energy
Ionisation
energythe atoms become positive ions.
e trend in
Describe the trend in
•Electrons
energy
is the from
energy
required
to remove
Electrons can be configurations.
removed energies
from atoms
and the energy
itIonisation
takescan
to remove
be removed
atoms
and the
energy ai
nergies a) down
ionisation
a) down
from
a
mole
of
atoms
in
the
gaseous
state,
and
is mea
them can be ➔
measured.
This
is across
called
energy
them because
can be measured.
as the
This is called ionisation energ
Explain
trends
in ionisation
b) across a
a group how
and b)
a
electrons are removed,
the
atoms
become positive ions.
are removed,
thethe
atoms
become positive
ions.
•electrons
Ionisation
energy has
abbreviation
IE.
ionisation
energies
rms of Note
electron
period in
terms
of electron
provide
evidence
the
ons. The shells of electrons •at fixed
configurations.
Ionisation
energy
is the
energyfor
required
to remove
• aIonisation
mole of electrons
energy
is the energy
required
or
Removing
the
electrons
one by
one to remove a
existence
of
electron
shells
–1.
specific levels are sometimes
called
from
a
mole
of
atoms
in
the
gaseous
state,
and
is
measured
from
a
mole
in
kJ
mol
of
atoms
in
the
gaseous
state,
and is
mea
w trendsquantum
in
➔ Explain how trends in
You can measure the energies required to
remove
the
el
shells, because the word
and sub-shells.
•
Ionisation
energy
has
the
abbreviation
IE.
•
Ionisation
energy
has
the
abbreviation
IE.
nergies
ionisation
energies
one
from
an
atom,
starting
from
the
outer
electrons
and
‘quantum’ is used to describe
reference:
something
dence for
the related to a fixed amount Specification
provide evidence
for3.1.1
the
•Removing
The first electron
needs the
least
Removing
the
electrons
one
by one
the electrons
one
by energy
one to remove
or fixed level.
f electron shells
existence of electron shells
being
removed
from
a
neutral
atom.
is the
You can measure the energies
required to remove
the
You
electrons
can measure
one10
bythe energies required to This
remove
thefir
el
0
5
ells.
and sub-shells.
Highest
energy
one
from
an atom, starting from the outer
electrons
one
and
from
working
an
atom,
inwards.
starting
from
the
outer
electrons
and
•
The
second
electron
needs
more
energy
than
the
fi
Number of electrons removed
level – electron Specification reference:electrons
ference: 3.1.1
3.1.1 in
being removed from a +1 ion. This is the second IE
Log ionisation energy
➔
Figure 5.2 !
Log (ionisation energy) plotted against
the number of electrons removed for
The patterns➔
in
sodium.
removed
•easily
The
first electron needs the least
energy
• The
it because
first electron
it is needs the least energy to remove
main
shell 1 to remove
in Figure
the bigthe
jumps
in
value
between
theelectron
first
and second
• first
The
third
needs
even more
TheNotice
bigfrom
jumps
inneutral
value5.2
between
firstis
and
second
ionisation
energies
and
being removed
a
atom.
This
the
being
IE.
removed
from
aagain
neutral
atom.energy
This is to
therem
fir
Intermediate
electrons and
in again between the ninth and tenth ionisation
ionisation
energies
energy level – between the ninth and tenth ionisation energies are
is
being
removed
from
a
+2
ion.
This
is
the
third
IE
very
visible
in
the
graph
above.
mainsuggests
shell
2 that
energies. This
sodium
atoms
electron
•electrons
The harder
second electron
needs
more
energy
than
the
• have
first
Theone
because
second
electron
itinisan outer
needs more energy than the fi
electrons
in second
• IE.
The
fourth
yet
more,
and
so on.
to remove
shellfrom
or energy
the nucleus.
This
outer needs
electron
is aeasily
being removed
a +1level
ion.furthest
This is from
the
being
removed
from
+1 ion.
This
is the second IE
log IE
electrons in
main shell 1
1
Thisremoved
suggestsbecause
that sodium
atoms
have
an shielded
outer shell
or energy
shell
1 one
it ismain
furthest
from
theelectron
nucleusinand
from
the fulllevel
log IE
•Lowest
Theenergy
thirdfurthest
electron
needs
even
moreouter
energy
to inner
remove
• easily
The
third
it called
because
electron
it needs
even
more energy
to rem
These
are
successive
ionisation
energies.
electrons
inthe
from
nucleus.
electron
is
removed
because
it is furthest
attraction
of
the
positive
by 10
electrons.
ectrons in
electrons
in Thisnucleus
levelis– being
electronsremoved
from
a
+2
ion.
This
is
the
third
IE.
is
being
removed
from
a
+2
ion.
This
is
the
third
IE
main
shell
3
Below
this
outer
electron, sodium atoms seem to have eight electrons
from the
nucleus.
ain shell 2
main
shellsingle
2
For example, sodium:
hardest to remove
in
a
second
shell
–
all
at
roughly
the
same
energy
level.
These
eight
electrons
1 2 yet
3 4more,
5 6 7
8 so
9 10
• The fourth 0needs
and
on.11
• The fourth needs yet more, and so on.
are
closer
to single
theelectrons
nucleus
than
the single
and only
have two
Figure 5.3 "
Below
this
outer
electron,
sodium
atomsouter
seemelectron
to
have eight
electrons
in a + e−
total
number
of
removed
Na(g)
Na+(g)
first IE
= + 49
➝
areatom.
called
successive
ionisation
energies.
called
successive
ionisation
energies.
inner
electrons
shielding them
from the These
positiveare
nucleus.
Energy levels of electrons These
in a sodium
electrons
in
second
shell
–
all
at
roughly
the
same
energy
level.
These
eight
electrons
are
closer −
+
2+
60
▲ Figure
1
The
successive
ionisation
Na (g) ➝ Na (g) + e
second IE = + 45
main shell 3
tosodium:
the nucleus
thanagainst
the single
outerofelectron. For example,
For example,energies
sodium: 3+
of sodium
number
2+
−
5 6 7 8 9 10 11
0 1 2 3 4 5 6 7 8 9 10 11
Na (g) ➝ Na (g) + e
third IE
= + 69
electrons removed. Note that the log of the
+
−
−1
+
−
of electrons removed
total
number
of
electrons
removed
Finally,
sodium
atoms
have
two
inner
electrons
in
a
shell
or
level
closest
to
the
Na(g)ionisation
Na(g)
(g)has+been
e plotted
first inIEorder = +and
496sokJ
mol
first IE
= + 49
➝ Na
➝ Na
energy
on,
see Table
1. (g) + e
+
2+
−
+
−1
2+
−
nucleus.
These
two
electrons
feel
the
full
attraction
of
the
positive
nucleus
and
are
uccessive ionisation
▲
Figure
1
The
successive
ionisation
Na (g)to fit the
Na range
(g) +
e
second
IE = + 4563 Na
kJ mol
(g) ➝ Na (g) + e
second IE = + 45
➝ large
of values
on the scale
m against number of
energies
of remove.
sodium
against
number
of endothermic ionisation
hardest to
They− have
the most
energies. 3+
2+
3+
2+
−1
−
Na (g) ➝ Na (g) + e
third IE
= + 6913 Na
kJ mol
(g) ➝ Na (g) + e
third IE
= + 69
▼
Table 1removed.
Successive
d. Note that the log of the
electrons
Noteionisation
that the logenergies
of the of sodium
has been plotted in order
ionisation
has been1st
plotted in2nd
order
and so on, see
Table energy
1.
and so on,
3rd
4th see Table
5th 1. 6th
7th
8th
9th
Electron
removed
ge of values on the scale
to fit the large range of values on the scale
Ionisation energy /
496
4563
6913
9544 13 352 16 611 20 115 25 491 28 934
–1
ssive ionisation energies of sodium
▼ Table
Successive
ionisation energies of sodium
kJ1mol
5th removed
6th
1st
7th
2nd
8th
3rd
9th
10th
4th
11th
5th
6th
7th
8th
9th
Electron
Ionisation energy /
Notice
that
the159
second
is not 20
the115
energy
change
496
4563
6913
9544 13 352 16
115 25
4563
491 28
6913
934 141
9544
367
13
352
079 IE
16 611
25 491
28 for
934
Study
tip–1 611 20496
kJ mol
2+
−
Na(g) ➝ Na (g) + 2e
The shape of the graph in Figure 1
The energy for this process would be (first IE + second
has to be thought about carefully.
Notice that the second IE is not the energy changeNotice
for
that the second IE is not the energy change for
The
first electron
Study
tip removed is in the
If
you
plot
a graph of the values shown in Table 1 you
−
Na(g) ➝
Na2+
Na(g) ➝ Na2+(g) + 2e−
(g) shell
+ 2eand
outer
main
the 10th and
e graph in Figure 1
The shape of the graph in Figure 1
Notice that one electron is relatively easy to remove,
electrons
removed
are(first
in the
The energy for11th
this
process
would
The energy
IE). for this process would be (first IE + secon
ght about carefully.
has
to be
thought
about be
carefully.IE + second
group of eight that are more difficult to remove, and
innermost main shell.
on removed is in the
The first
removed
is inin
theTable 1are
very
difficult
remove.
If you plot
a graph
ofelectron
the values
shown
Ifyou
you
get
plot
Figure
a graph
1.to of
theBilal
values
shown in Table 1 you
Marginalizer
Hameed
ll and the 10th and
outer main shell and the 10th and
Notice that one
electron
is removed
relatively
easy
Notice
thenthat
comes
one aelectron is
relatively
easy to remove,
removed are
in the
11th
electrons
are
in theto remove,
Electronic
Structure
10
Bilal
Hameed
26
group of eightinnermost
that are main
moreshell.
difficult to remove, and
group
finally
of eight
two that are more difficult to remove, and
n shell.
are very difficult to remove.
are very difficult to remove.
ed
gy /
1st
2nd
3rd
4th
of the nucleus. The higher the value, the more energy is required to
remove
of electrons. The first ionisation energy for magnesium is
12th 1 mole
13th
−1. 738 kJ of energy are required to convert 1 mole of Mg(g)
+738
kJ
mol
200 000 222 000
1 mole of electrons. There are trends in
to 1 mole of Mg+(g) by removing
9
ionisation energy values in the Periodic Table that provide evidence for the
existence
of electron
arrangement in energy levels and in sub-shells.
● The data
in Table
ELECTRONIC
STRUCTURE
9 2.3 show that the fi rst three electrons are considerably easier
to remove than the fourth, as there is a big jump from the third to the fourth
Ionisation
energies as evidence for energy levels
ionisation energies. These first three electrons come from the outer shell.
● The last two electrons are very much harder to remove than the preceding eight,
General increase in successive ionisation
6 as there is a huge jump between the eleventh and twelfth ionisation energies.
energy values due to increase in effective
nuclear
charge,come
i.e. ratio
of protons
to shell.
These
electrons
from
the inner
electrons
increasing
5
● As the shells or orbits of the electrons get further from the nucleus, the energy level
rises, so less energy is required to remove an electron from that shell. The values
4 in Table 2.3, with the jumps after the third and eleventh ionisation energies, mean
Large gap between ninth and
that an aluminium atom
haselectrons
three electrons
tenth
as tenth in its outer orbit, eight nearer to the
3 nucleus in an inner orbit
electron
is removed
from very
an
and two
electrons
close to the nucleus.
electrons is
each ionisation,
e ions formed may
xample Cl+(g) or
Log(ionisation energy)
alt
energy level closer to the nucleus
The
jumps
providefirst
evidence for quantum shells. Aluminium has two electrons
2 big
Large
gap between
and
second
electrons
in the first quantum shell, eight in the second and three in the third.
og(ionisation energy)
er of electrons
odium atom.
19
1
For Al, th
occurs b
and four
energies
outer ele
group 3.
as second electron is
1
removed
an energy
The
group
infrom
which
an element is found can be worked out by looking at where the
level closer to the nucleus
first big jump in the successive ionisation energy occurs. If it occurs between the
0
fourth
and
fi fth3ionisation
element
2
4
5 energies,
6
0
1
7 then
8
9the 10
11 has four outer electrons and is
in group 4. Number of electrons removed
Log IE
Another way of presenting the data is in graphical form. As the variation between
The
in Figure
1.23 shows
the successive
ionisation
energies
a values as the
thegraph
first and
last ionisation
energies
is so great,
it is usual
to plotofthe
sodium
atom,
becomes
clearenergy.
that there
is ais distinct
of energy
logarithm
of itthe
ionisation
This
shown set
in Figure
2.8levels.
for the element
The electron
arrangement
in
sodium
is
usually
written
as
2,8,1.
The
diagram
shows
the
log(ionisation
energy)
against
the
number
of
sodium (Z = 11).
electrons removed from a sodium atom.
This met
the grou
element
not appl
element
all have
their out
from chr
which ha
electron.
6
The log of1 the ionisation energy
is used to condense the diagram as the
8
electron
electrons
ionisation
energies vary inacross
a wide range of values.
5 in outer
middle
shell
shell
2
The existence of energy levels is proven by the
large
gaps in the successive
electrons
4
inner
ionisation
energies as these correspond to theinremoval
of electrons from
shell
19
energy levels closer to the nucleus and so more energy is required to
3
remove
the electron.
The 2general increase in successive ionisation energies is caused by the
0
1
2
3
4
5
6
7
8
9 10 11
increase in the ratio of protons
to electrons
asremoved
successive electrons are
Number
of electrons
removed.
is often
called effective
Figure 2.8This
Successive
ionisations
of sodiumnuclear charge.
As the successive ionisation energies of an element increase, there is a big jump in
value each
startthat
to bethere
removed
the next
shellbetween
nearer the nucleus.
It time
canelectrons
be seen
is from
a big
jump
the first
and second ionisation
energies and another big jump between the ninth and tenth. This means that sodium
By noting where the first big jump comes in the successive ionisation energies, it is
has one electron in its outer orbit, eight in the next inner orbit and two in the orbit
possible to predict the group to which an element belongs. For example, the first big
nearest to the nucleus. Thus, the electronic structure of sodium is 2,8,1.
jump in the successive ionisation energies for sodium comes after the first electron is
removed. This suggests that sodium has just one electron in its outermost shell, so it
must be in Group 1.
Bilal Hameed
Ionis
Marginalizer
11
Remem
electron
comes
3/27/19 1:54 PM
807404_C02_Edexcel_GF_Chem_009-036.indd 19
Bilal Hameed
Tip
Electronic Structure
0000
0000
aluminium atom.
alt
NOTE
Logarithms to base 10
Logarithmic scales are a means
of bringing a very wide range of
numbers onto the same scale. For
a number, x, log10x is the power to
which 10 must be raised to equal
x. So,
If x = 10,
If x = 100,
If x = 106,
If x = 0.1,
log1010 = 1
log10100 = 2
log 10106 = 6
log100.1 = –1
The logarithmic plot in Figure 2.16 allows an extremely
energies to be shown on the same graph, from 494 kJ m
Notice also in Figure 2.16 how there are two big jumps
the first and second ionisation energies and the ninth an
energies. See if you can now deduce the electronic struc
answering question 14.
1
Figure 2.17 (a) shows a sketch graph of log10 for the succe
of potassium. This graph has been used to construct an en
(Figure 2.17 (b)) for the electrons in a potassium atom. Th
three big jumps in the ionisation energies between the fir
the ninth and tenth and between the seventeenth and eig
These jumps suggest that the electronic structure of potass
the number of electrons in each quantum shell as we mov
Successive ionisation energies for potassium
Energy
log10 ionisation energy
level
If 5a graph of ionisation energy (rather than log10 ionisation energy) is
plotted for the removal of the first few electrons from a silicon
atom, more n = 4
2 electrons
The first two successive ionization energies of magnesium increase from
close
features can be seen on the graph (Figure 2.36). It can bevery
seen
that there
736 kJ mol 1 to 1451 kJ mol 1. The next ionization energy increases greatly to to nucleus,
n=3
1
is a larger
jumpthat
in the
the3rd
ionisation
between
and third
, indicating
electron isenergy
being removed
fromthe
the second
7732.6 kJ mol
n = 1 shell
next energy level.
This energy level is full and closer to the nucleus, so a much
4
ionisation
energies.
n=2
greater amount of energy is required for the 3rd electron to be removed than
2 2 6 2 2
2s
2p
3s
3p
.
The
fi
rst
The
full
electronic
confi
guration
for
silicon
is
1s
for the 2nd electron. The successive ionization energies increase
gradually
as
8 electrons
closer to
electrons are
fromare
the removed
second electron
shell.
There
is another
large
nucleus,
n = 2 shell
n=1
tworemoved
electrons
from
the
3p
sub-level
(subshell),
whereas the
jump from 35 461 kJ mol 1 for the 10th ionization energy to 169 987 kJ mol 1
third
electron
is removed
fromthis
theelectron
3s sub-level
(Figure 2.37). The 3p subfor the 11th
ionization
energy,
indicating that
is being removed
(b)
from the next
energy
level,
the
inner
electron
shell.
This
leads
us
to
the
3 is higher in energy than the 3s sub-level, andelectron
level
therefore less energy is
configuration of 2,8,2 for magnesium.
required to remove
the electron.
This provides evidence for the existence
8 electrons
far
The pattern formed by successive
ionization
energies can be seen clearly in a
from
nucleus,
n
=
3
shell
of subbeing
energy
levelsagainst
(subshells)
in an
atom.
graph of electron
removed
ionization
energy
(kJ mol 1).
Figure 1.1.3 shows the successive ionization energies of nitrogen (Z 7). In
1 electron
veryfor
farthe
from
figure 1.1.3, the ionization
energies
first 5 electrons are low, suggesting
2
nucleus, n = 4 shell
that there are five electrons in the outer shell, then there is a sharp increase in
15 Thefortable
shows
theelectrons,
successive ionisation
of
some elements.
the ionization energies
the15
first16 17 18 19
1
2 3 the4 6th
5 and
6 7th
7 8 9 10 which
11 12are
13in 14
Deduce which group in the periodic table each element is in.
electron(a)
shell.
Number of the ionisation
Fig 2.17 (a) A sketch graph of log10 of the successive ionisation energies for potassium
(b) An energy level diagram for the electrons in a potassium
atom level
1st energy
6.00
32
1st
energy
level
2nd
are usingenergy
reasoning
level
0000
to deduce the existence
0000
of energy levels in an
0
0
1
2that energy
3
4
5
6
. Do we
know
number of the electron being removed
exist?
0000We
1.3 The successive ionization energies of nitrogen.
10
0000
ionisation when all the electrons are successively re
Successive ionisation energies for silicon
log (ionization energy)
0000
Graphs of successive ionisation energyelectrons
give us information
in each shell.about
(The electronic structure
how many electrons are in a particular
energy
level.
Consider
the
graph
15 The first six ionisation energies
of an element in kJ
for silicon shown in Figure 2.35. There
is a large
2920,
4960,jump
6270 in
andthe
21ionisation
270.
How
electrons
are there
in the outer shell
10and athe
energy graph between the fourth
fifthmany
ionisation
energies,
which
element?
suggests that these electrons are removed from different main energy
Which group in the periodic table does the elem
levels. It can STRUCTURE
therefore be deduced
thatb silicon
has four electrons in its
ELECTRONIC
10
16 Sketch a graph of log10 ionisation energy against th
outer main energy level (shell) and is in group 4 of the periodic table.
2nd energy level
4.00
3.00
2.00
1.00
0.00
7
3rd energy level
5.00
4th energy level
0
2
4
6
8 10 12 14 16 18 20
number of the electron being removed
Figure 1.1.4 The successive ionization energies of potassium. This graph is plotted as
log10(ionization energy), to make the increases more obvious.
Marginalizer
Evidence
for the four electron shells of potassium can be seen in figure 1.1.4. In Bilal Hameed
this graph there are three sharp increases in ionization energy. The first occurs
Electronic Structure
12the second and third each after another Bilal Hameed
after 1 electron has been removed and
8 electrons have been removed. This leads to the electron configuration of
2,8,8,1 for potassium.
ELECTRONIC STRUCTURE 11
Successive ionisation energies for argon
12.1 Electrons in atoms 447
alt
Ionization energy/kJ mol–1
and 7 is due to a change in sub-shell from 3p to 3s. The slight
increase in ionization energy for electrons 4, 5 and 6 relative to
electrons 1, 2 and 3 can be accounted for by enhanced electron–
electron repulsion.
11 Use a reliable source on the internet to find twelve successive
ionization energies for calcium and vanadium. Plot logarithms (to
the base 10) of the values. Relate the oxidation states and detailed
electron configuration of each element to the graphs.
1
Electron-in-a-box model
The electron-in-a-box model (Figure 12.23) is a very simple
quantum mechanical model that describes an electron in a covalent
bond (Chapter 14). This simple model imagines an electron trapped
between two infinitely high potential ‘walls’. The potential energy
of the electron is zero inside the box but is infinite at the walls. The
2
3
4
5energy
6
7 (rather
8
9 than
10 11
is trapped
in a potential
If a graph 1of ionisation
log10electron
ionisation
energy)
is well and cannot escape.
Number of electrons being removed
This model assumes that the electron behaves as a standing
plotted for the removal of the first few electrons from a silicon atom, more
wave (wave–particle duality) and is subject to boundary conditions
Figure 12.22 Eleven successive ionization energies of
features
can be seen on the graph (Figure 2.36).similar
It canto be
seen
thatto there
those
applied
the tension waves of a violin string
an argon atom
Successive
ionisation
energies
for
period
3
elements
1
1.1.3 VALUES
IONIZATION
ENERGIES,between
In, (kJ molfixed
) FOR
ELEMENTS
OF PERIOD
3
both
ends.and
The
standing
waves have nodes (regions of
is a TABLE
larger
jump OF
inSUCCESSIVE
the ionisation
energy
theatTHEsecond
third
Element
I
I2
I3
I4
I5
I6 vibration
I7 or zero
I8 electron
I9 density) and antinodes (regions of
no
ionisation
energies.
Na
494
4562
6912
9543
13 353
16 610
20 114
25 490
28 933
maximum
vibration
and maximum
electron density).
Mg
736
1451
7733
10 540
13 630
17 995 2 21
703 6
25 6562
31 642
3s227
3p
. The
first
The
full electronic
confi2745
guration
for silicon
is 1s 2s2322p
Al
577
1817
11 575
14 839
18 376
293
457
31 857
Si
786
1577
3231
4356
16
091
19
784
23
786
29
252
33
876
two electrons are removed from the 3p sub-level (subshell),
the
V =whereas
∞
P
1060
1903
2912
4956
6273
21 268
25 397
29 854
35 867
n=3
third
electron
is
removed
from
the
3s
sub-level
(Figure
2.37).
The
3p
subS
1000
2251
3361
4564
7013
8495
27 106
31 669
36 578
2297
3826 the 3s
5158sub-level,
6540
9362 therefore
11 020
33
610 energy
38 600 is
levelCl is higher1260
in energy
than
and
less
Ar
1520
2666
3928
5770
7238
8811
12 021
13 844
40 759
∞
V(x)
box
required to remove the electron. This provides evidence for the existence
First ionization
of sub
energyenergies
levels (subshells)
in anV =atom.
V= ∞
V= ∞
0
n=2
V= ∞
ψ3
energy
Recall that the first ionization energy of an element is the amount of energy
required to remove one mole of electrons from one mole of atoms of an element
in the gaseous state. The state of the element is important because if the
x
element was solid or liquid, the input of energy would change its state before
a
0
electrons could be removed. A graph of first ionization energies of the first 54
Deduce
what groups fo Elements X,electron
Z and Q belong to.
elements shows distinct patterns that lead us to a greater understanding of
the electron structure of an atom.
Figure 12.23 The electron-in-a-box
ψ2
Skill Check 1
n=1
15 The table shows the successive ionisation of some elements.
Deduce which group in the periodic table each element
is in.x
V=0
0
first ionization energy (kJ mol–1)
model (V represents the potential
energy of the electron)
2500
He
WORKSHEET 1.1
Ionization energies
ψ1
0
L
0
L
0
x
L
L
energy levels
wave functions
Figure 12.24 Electron standing waves
Ne
2000
The walls of the box must always be nodes, so these standing waves must have an integer
(whole) number of half wavelengths within the ‘box’. This results in quantization (Figure 12.24)
of the energy levels available to the electron, just as the boundary conditions for a violin string
500
(see Chapter 24 on the accompanying website) produce first, second and third harmonics etc.
0
The electron-in-a-box
is important because it shows how discrete energy levels arise
0
10
20
30
40
50model 60
atomic number (Z )
when
an electron is confined to a tiny region of space. The model can be made quantitative and
canto be
Figure 1.1.5 The first ionization energies of hydrogen
xenon.used to calculate the energy levels and to show that that as the ‘box’ is made larger, the
wavelengths increase and the energy decreases.
1500
1000
Ar
N
H
O
Be
C
Mg
Si
B
Li
P
Na
Al
Cl
S
Kr
Xe
I
Cd Sb
Pd
Mn Co
Se
Te
Tc
Nb
Ti
Y
Ca
Sr
Fe Ni Cu Ge
Sn
Ru Rh Ag
Zr Mo
Sc V Cr
Ga
In
K
Rb
Zn
As
Br
ATOMIC STRUCTURE
F
The most obvious feature of this graph is the periodic series of peaks
corresponding to the first ionization energy of the noble gases (He, Ne, Ar, Kr,
Xe). These elements have high first ionization energies because they have a
Each
in an atom
canThe
be uniquely described by a set of four quantum numbers
full electron shell and an associated
high electron
degree of energetic
stability.
next most obvious feature of the graph
is the
lowest
point
of each periodic
(Table
12.3):
the
principal
quantum number (n), the angular momentum quantum number
series. These troughs correspond to the group 1 elements (Li, Na, K, Rb). These
quantum
number
(m) and the spin quantum number (s).
elements have only one electron inmagnetic
the outer shell
and so the
first ionization
energy is small, as little energy is required to remove this electron from the
atom. Recall that the attraction between the valence electrons and the nucleus
is not great due to the low core charge (Chemistry: For use with the IB Diploma
Programme Standard Level, p. 85).
Quantum numbers
(l), the
CHAPTER 1
easoning
existence
s in an
t energy
for silicon shown in Figure 2.35. There is a large jump in the ionisation
energy graph between the fourth and the fifth ionisation energies, which
suggests that these electrons are removed from different main energy
11 silicon has four electrons in its
levels. It can therefore be deduced that
outer main energy level (shell) and is in group 4 of the periodic table.
CHEMISTRY: FOR USE WITH THE IB DIPLOMA PROGRAMME
Bilal Hameed
Bilal Hameed
829055_12_IB_Chemistry_435-450.indd 447
5
Marginalizer
13
Electronic Structure18/05/15
10:35 a
12
ELECTRONIC STRUCTURE 12
Skill Check 2
(a) Write an equation to represent the 5th ionisation energy of Fluorine.
alt
(b) 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.
1
Skill Check 3
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
Skill Check 4
The successive ionisation energies of beryllium are 900, 1757, 14,849 and 21,007 kJ
mol-1.
(a) Why do successive ionisation energies of beryllium always get more
endothermic?
(b) To which group of the Periodic Table does this element belong?
Marginalizer
Electronic Structure
Bilal Hameed
14
Bilal Hameed
estions are
etting you to
n a particular
Using successive ionisation energies
13
Successive ionisation energies are an indicator of the group to which an
element belongs.
ELECTRONIC STRUCTURE 13
Skill Check 5
EXAMPLE 12
The first five successive ionisation energies for four different elements are given in
The first five successive ionisation energies for four different elements
the table.
alt
are given in the table.
Element
Ionisation energy (kJ mol−1)
First
Second
Third
Fourth
Fifth
W
+496
+4562
+6912
+9543
+13353
X
+1087
+2353
+4621
+6223
+37831
Y
+578
+1817
+2745
+11577
+14842
Z
+738
+1451
+7733
+10543
+13630
1
1 State which element belongs to Group 4.
(a)
State which
element
belongs
4. ion with a charge of 2+?
2 Which
element
would
formtoaGroup
simple
(b)
Which
element
would
form
a
simple
ion
with ain
charge
of 2+?
3 Which element would have one electron
its outer
energy level?
(c)
Which element
have
onean
electron
its outer
energy level?
4 Which
elementwould
would
form
oxide in
with
the formula
M2O3 where M
thewould
element?
(d)represents
Which element
form an oxide with the formula M2O3 where M represents
the element?
Answers
1 Element
Skill
CheckX6has a large increase in ionisation energy after the fourth
electron has been removed (+6223 to remove the fourth electron and
The+37831
successive
ionisation
mol−1This
of element
X are listed
below.
Identifyin
to remove
theenergies/kJ
fifth electron).
would suggest
four
electrons
thethe
group
in
the
periodic
table
in
which
X
occurs.
Ionisation
energies
of
X:
outer energy level and the fifth electron in an energy level closer to the
1stnucleus.
950; 2nd
3rdX2700;
4thelectrons
4800; 5th
6000;
12 300;
7th
15
So,1800;
element
has four
in the
outer6th
energy
level,
which
is characteristic of an element in Group 4. Element X is actually carbon.
000
2 An element in Group 2 would form a simple ion with a charge of 2+.
Element Z has a large increase in ionisation energy after the second
electron
Skill
Checkhas
7 been removed (+1451 to remove the second electron and
+7733 to remove the third electron). This would suggest there are two
Which equation represents the second ionisation energy of an element X?
electrons in the outer energy level and the third electron is in an energy
closer
to +the
So, element Z has two electrons in its outer
level
X2+(g)
2enucleus.
A X(g) ➝
energy level, which is characteristic of an element in Group 2.
magnesium.
B Element
X+(g) ➝ XZ2+is(g)actually
+ e3 One electron in the outer energy level would suggest a Group 1
C element.
X(g) + 2e-Element
➝ X2-(g) W has a large increase in ionisation after the
first electron has been removed (+496 to remove the first electron
to remove
D and
X-(g)+4562
+ e- ➝
X2-(g) the second electron). This would suggest one
electron in the outer energy level and the second electron in an energy
level closer to the nucleus. So, element W has one electron in the
outer energy level, which is characteristic of an element in Group 1.
Element
Skill
CheckW8is actually sodium.
4 An element which forms an oxide with the formula M2O3 would suggest
(a) Write equations for the following ionisations including state symbols:
an element in Group 3. Element Y has a large increase in ionisation
energy
afterofthe
third electron has been removed (+2745 to remove
(i)first
ionisation
silicon
the third electron and +11577 to remove the fourth electron). This
electrons in the outer energy level and the fourth
(ii) would
secondsuggest
ionisationthree
of potassium
electron in an energy level closer to the nucleus. So, element Y has
electrons
the outer energy level, which is characteristic of an
(iii)three
third ionisation
of in
carbon
element in Group 3. Element Y is actually aluminium.
(b) Why do successive ionisation energies always get more endothermic?
Bilal Hameed
Marginalizer
3/27/19 1:54 PM
Bilal Hameed
15
Electronic Structure
radiation with
a frequency of
1368 kHz. Deduce
which part of the
electromagnetic
spectrum it belongs
to.
3.31 × 10 −19 J/photon
= 1.2 × 102 photons
Calculate the number of photons with wavelength 4.00 nm that can prov
(6.63 × 10 −34 J s × 3.00 × 108 m s−1)
= 4.972
For one photon: E = hc =
λ
4 × 10 −9 m
14
1
Calculate the
= 2.01 × 1016 photons
So for one joule:
4.9725 × 10 −17
ELECTRONIC STRUCTURE
frequency of 14
yellow light with
Evidence for sub-shells
of electrons
a wavelength
of
−8 cm.
5800 × 10
By studying the first ionisation
energies
of successive elements in the periodic table,
8
Electron configuration
9 easy
The laser
to
we can compare how
it is used
to remove
anIonization
electron from energy
the highest energy level in
read information
different atoms. This provides us with evidence
for
the
arrangement
of electrons
in
The first ionization energy
is the minimum
energy per mole req
from a compact disc
sub-shells.
has a wavelength of
one mole of isolated gaseous atoms to form one mole of gaseous
alt
780 nm. Calculate
thermodynamic conditions. For example,
the energy energies
Factors affecting ionisation
required to bring about the reaction:
associated with
The ionisation energy of
an
atom
is
influenced
by three atomic properties:
one photon of this
Cl(g) → Cl+(g) + e−
radiation.
the first ionization ene
1. Distance of the outermost electron from the nucleus: As the distance from
The electron is removed from the outer sub-shell (energy sub-lev
the nucleus increases, the attraction of the
nucleusTable
for the12.1
negative
is, positive
a 3p electron).
gives some examples of ionizations,
electron decreases, and this tends to reduce
the which
ionisation
energy.
energy,
is the
enthalpy change for the equation. Ionization
the IB Chemistry data booklet.
2. Size of the positive nuclear charge: As the positive nuclear charge increases
Table
Selected
with atomic number,
its 12.1
attraction
for outermost
electrons increases,
andequation
this
Element
Ionization
First ionization
ionization
energies
+
tends to increase the ionisation energy. Oxygen
O(g) → O (g) + e −
13
1
S(g) → S+(g) + e −
Sulfur
10
3. Shielding effect of inner shells of electrons:
inner→
shells
exert
Copper Electrons inCu(g)
Cu+(g)
+ e −a
7
repelling effect on electrons in the outermost shell of an atom. This effect shields
the outermost electrons from the attractive
force ofthat
the nucleus
reduces its energy
Factors
affectand
ionization
pull on them. This shielding means that the
‘effective
nuclear charge’
Values
of ionization
energiesattracting
depend on the following factors:
electrons in the outer shell is much less thanthe
the size
full positive
charge
of
the
of the atom (or ion)
nucleus. As expected, the shielding effect increases
significantly
the nuclear charge as the number
of inner shells increases.
the shielding effect.
Electrons in the same shell exert a relatively small shielding effect on each other.
Atomic radius
As the distance of the outer electron
the attraction of the positive nucleu
electrons falls. This causes the ioniz
ionization energy decreases as the at
–
3+
–
nuclear
pull
–
Nuclear charge
When the nuclear charge becomes m
presence of additional protons), its a
increases. This causes the ionization
Shielding effect
repulsion from
inner shell of
electrons (’shielding’)
Figure 12.13 Electrostatic forces operating on the
outer or valence electron in a lithium atom
The outer or valence electrons are r
in the atom in addition to being att
nucleus. The outer electrons are shie
the nucleus by the shielding effect (
repulsion) (Figure 12.13).
829055_12_IB_Chemistry_435-450.indd 442
Marginalizer
Electronic Structure
Bilal Hameed
16
Bilal Hameed
12.1 Electron
15
In general, the shielding effect is most effective if the electrons are close to the nuc
ELECTRONIC STRUCTURE
15
Consequently,
electrons in the first shell (energy level), where there is high electron de
have
a
stronger
shielding
effectin than
in the second shell, which in turn have
From the study of ionisation energies, we know that
the electrons
atomselectrons
are
shielding
effect
than
electrons
in
the
third
shell.
grouped together in shells or energy levels. Principal quantum numbers 1, 2, 3 and so Electrons in the same shell exert a re
small shielding effect on each other.
on are used to denote these shells working out from the nucleus.
Figure 12.14 shows the first ionization energies for the chemical elements of periods 1, 2
The
generalenergies
increase
energy
across
The graph below shows the first
ionisation
forin
theionization
elements of
periods
1, 2 each period is due to the increase in nucle
This occurs because across the period each chemical element has one additional proton, w
and 3.
increases the nuclear charge by +1.
alt
2500
He
First ionization energy/kJ mol–1
Ne
shielding
force
1
2000
1500
Li
Ar
Be
B
C
N
O
H
1000
500
0
Li
1
Na
5
electrostatic
attraction
towards
positive
nucleus
K
10
15
Atomic number, Z
20
Figure 12.14 First ionization energies for periods 1, 2 and 3
Figure 12.15 A diagram illustrating how the b
between shielding and nuclear charge changes
period 2
The most obvious feature of this graph is the periodic series of peaks corresponding
to the
firstatom,
ionisation
energy of the
gasesin
(He,
Ne, Ar).charge
These elements
have
Thenoble
increase
nuclear
increases
the force of attraction on all the electrons, so
boron
B
high1sfirst ionisation
energies
because
they have
full electron
and anEach additional electron across a period enters th
held closer
andahence
moreshell
strongly.
2s
2p
associated high degree of energetic
shell stability.
(energy level) and hence the increase in shielding is minimal (Figure 12.15).
Although the general trend is for the ionization energy to increase across the perio
+
boron
B hold only a limited number of electrons. If all the shells in the atoms of
Each
shellion,
can
are two distinct dips in ionization energy across periods 2 and 3 (Figure 12.14). These d
1s
2s
2p
an element are full, it will be veryonly
stable
a highly endothermic
first ionisation
bewith
explained
using an orbital
model of electronic structure.
energy.
The first decrease in each period is the result of a change in the sub-shell (sub-level)
which the electron is lost and a change in electron shielding. These have a greater effec
atom, Be
Theberyllium
first quantum
shell nearest the nucleus (n = 1) is full and stable when it contains
increase in nuclear charge and decrease in atomic radius. In period 2, this first decrease
2s
2pis the case in helium atoms, which are very stable and
just 1s
two electrons.
This
between the elements beryllium and boron. When it is ionized, the beryllium atom (1s2
unreactive with a higher first ionisation energy than neighbouring elements in
the
a 2s electron, whereas a boron atom (1s22s22p1) loses a 2p electron (Figure 12.16). More e
+
periodic
table.
beryllium ion, Be
required to remove an electron from the lower energy 2s orbital in beryllium than from
energy
2pn orbital
inand
boron.
Although
the 2s and 2p sub-levels are in the same shell, the
The second quantum shell (energy
level,
= 2) is full
stable
when it contains
difference
is
relatively
large.
Recall
(Chapter
2) that the energy gap between shells and
eight electrons. This is the situation in neon atoms. Neon has a filled first shell with 2
becomes smaller with an increase in shell number. In addition, a single electron in the 2
Figure
12.16
Orbital
electrons
and
a filled
second shell with 8 electrons. Its electronic structure is 2,8 and
2
notations for boron and level is more effectively shielded by the inner electrons than the 2s electrons (Figure 12
1s
2s
2p
neon has a higher first ionisation energy than its neighbours in the periodic table like
beryllium
atoms
helium.
The high
firstand
ionisation energies of helium and neon show that their
their
unipositive
electronic structuresions
are very stable 2s
and explain why the elements are so unreactive.
Bilal Hameed
Bilal Hameed
2p
Figure 12.17 Electron density clouds of the 2s and 2p orbitals (only one lobe shown). The d
Marginalizer
shows the extent of the 1s orbital; the 2s electron
can partially penetrate the 1s orbital, incre
stability
17
Electronic Structure
16
ELECTRONIC STRUCTURE 16
Ingraph
general,
shielding
is most effective if the electrons ar
The next most obvious feature of the
is thethe
lowest
point ofeffect
each periodic
Consequently,
inNa,
theK).first
shell
(energy level), where there i
series. These troughs correspond
to the group 1 electrons
elements (Li,
These
elements
have
a
stronger
shielding
effect
than
electrons
have only one electron in the outer shell and so the first ionisation energy is small, as in the second shell, wh
shielding
effect
than
electrons in the third shell. Electrons in the sam
little energy is required to remove
this electron
from
the atom.
alt
small shielding effect on each other.
Figureacross
12.14 each
showsperiod
the first
ionization
energies
The general increase in ionisation energy
is due
to the increase
in for the chemical eleme
The
general
increase
in
ionization
energy
across
nuclear charge. This occurs because across the period each chemical element has each period is due to the
This occurs
because
across
one additional proton, which increases
the nuclear
charge
by the
+1. period each chemical element has one ad
increases the nuclear charge by +1.
2500
He
1
First ionization energy/kJ mol–1
Ne
2000
shielding
force
Ar
1500
Li
Be
B
H
1000
500
0
Li
1
Na
5
10
15
Atomic number, Z
electrostatic
attraction
towards
positive
nucleus
K
20
Figure 12.14 First ionization energies for periods 1, 2 and 3
Figure 12.15 A diagram illu
between shielding and nucle
period 2
The points between one noble gas and the next in the graph above can be divided
intoboron
sub-sections.
providein
evidence
sub-shells
(sub-levels)
The increase
nuclearforcharge
increases
the force of attraction on al
atom, B These sub-sections
of electrons.
held closer and hence more strongly. Each additional electron across
1s
2s
2p
shell (energy level) and hence the increase in shielding is minimal (F
After both He and Ne there are deep
troughs followed
by small
intermediate
peaks
at
Although
the general
trend
is for the
ionization
energy to increas
ion,These
B+ are subsections with just two points.
Be boron
and Mg.
are two distinct dips in ionization energy across periods 2 and 3 (Figu
1s
2s
2p
only be explained using an orbital model of electronic structure.
Immediately after Be and Mg there are similar sub-sections of six points (B to Ne and
The first decrease in each period is the result of a change in the su
Al to Ar) or two sub-sections of three points.
which the electron is lost and a change in electron shielding. These h
radius. In period 2,
between the elements beryllium and boron. When it is ionized, the be
The n = 1 shell can hold 2 electrons
in the same
sub-shell.
a 2s electron,
whereas
a boron atom (1s22s22p1) loses a 2p electron (Fig
+
beryllium ion, Be
required to remove an electron from the lower energy 2s orbital in ber
1s n = 2s
The
2 shell can2p
hold 8 electrons:
2 inorbital
one sub-shell
and Although
6 in a slightlythe 2s and 2p sub-levels are in th
energy 2p
in boron.
higher sub-shell.
difference is relatively large. Recall (Chapter 2) that the energy gap be
becomes smaller with an increase in shell number. In addition, a singl
Figure
Orbital
The
n = 312.16
shell can
hold 18 electrons: 2 in one sub-shell, 6 in a slightly
2
notations for boron and level is more effectively shielded by the inner electrons than the 2s el
beryllium atom, Be
inchemists
nuclearhave
charge
and decrease
in atomic
By studying
energiesincrease
in this way,
deduced
the following:
1s
2sionisation2p
•
•
•
•
higher sub-shell and 10 electrons in a sub-shell that is slightly higher still.
beryllium atoms and
their
ions
The nunipositive
= 4 shell can
hold 32 electrons, with sub-shells containing 2, 6, 10 and 14
2s
electrons.
Marginalizer
Electronic Structure
2p
Bilal Hameed
2s and 2p orbitals (only one l
18 Figure 12.17 Electron density clouds of the
Bilal
Hameed
shows the extent of the 1s orbital; the 2s electron can partially penetrate t
stability
17
ELECTRONIC STRUCTURE 17
Sub-shells
•The sub-shells that make up the main shells are given names:
•the sub-shells that can hold up to 2 electrons are called s sub-shells
•those that can hold up to 6 electrons are called p sub-shells
•those that can hold up to 10 electrons are called d sub-shells and
•those that can hold up to 14 electrons are called f sub-shells.
alt
sub-shell
4f
3p
4s
n=3
3s
2p
2s
1s
n=4
s of electron sub-shells
s
s
2
2
2
p
6
8
Sub-shell
4p
3d
Maximum number
of electrons in the
shell (energy level)
n=1
n=2
4d
1
Number of
electrons in
sub-shell
Shell (energy
level)
s
2
p
6
d
10
s
2
p
6
d
10
f
14
18
32
■ Table 2.7 Structure of sub-shells (not hydrogen)
Potential energy
4d
4p
4s
shell
sub-shell
3d
3p
n=3
n=4
3s
1s
n=1
4f
n=4
Shell (energy
level)
n=1
n=2
4d
4p
2p
2s
n=3
Potential energy
Potential energy
The electron structure of an atom can be described in terms of the shells occupied
Figure 2.53 shows the energy levels of the atomic orbitals (except hydrogen). Note that the
by electrons.
In terms
of the
shells,
electron
of lithiumfill
is 2,1
4s sub-shell (sub-level) has
a lower energy
than
3d the
sub-shell
andstructure
hence electrons
theand that of
sodium
is
2,8,1.
It
is
also
possible
to
describe
the
electron
structure
an atom more
4s sub-shell before they occupy the 3d sub-shell. This sub-shell overlap (Figure 12.54) first of
occurs
in terms(Chapter
of sub-shells.
with the first row of theprecisely
d-block metals
13). However, the 3d sub-shell is then stabilized
across the first row of the
d-block
metals.(sub-levels) are being filled, electrons always occupy the lowest
When
sub-shells
available
energy level first. The figure below shows the relative energy levels of the
76 2 Atomic
structure
4f
various sub-shells in the first four quantum shells.
3d
n=2
3p
4s
n=3
3s
2p
n=2
2s
3d
n=1
4s
1s
■ Figure 2.52 Energy levels of electron sub-shells
4s
3d orbitals are
empty (Ca)
Bilal Hameed
Bilal Hameed
Energy levels■inFigure
hydrogen
2.53 Orbital
n=4
3d
Sub-shell
Number of
electrons in
sub-shell
Maximum number
of electrons in the
shell (energy level)
s
s
2
2
2
p
6
8
s
2
p
6
d
10
s
2
p
6
d
10
f
14
18
32
■ Table 2.7 Structure of sub-shells (not hydrogen)
Figure 2.53 shows the energy levels of the atomic orbitals (except hydrogen). Note that the
4s sub-shell (sub-level) has a lower energy than the 3d sub-shell and hence electrons fill the
3d orbitals occupied (Sc)
4s sub-shell before they occupy the 3d sub-shell. This sub-shell overlap (Figure 12.54) first occurs
Marginalizer
with the first row of the d-block metals (Chapter 13).
However, the 3d sub-shell is then stabilized
across the first row of the d-block
metals.
19
Electronic Structure
In a hydrogen atom
the sub-shells
all have the same energy. For example, the 2s and 2p
structure
of atoms
4f
an orbital and a subshell (sub
energy level)
18
evidence, such as ion
simple treatment of c
energy levels is a use
ELECTRONIC STRUCTURE 18
The principal energy levels (shells) get closer together as you get further from the
nucleus. This results in an overlap of sub- levels.
Each main energy lev
(subshells). The first m
second main energy
The sub-levels in eac
alt
1
Within any main
(subshells)
is always s
From it we can deduce that the order in which the sub-shells are filled in the first
four
principal quantum shells is: 1s, 2s then 2p, 3s then 3p, 4s then 3d, then 4p orders between sub-l
So, the single electron in a hydrogen atom goes in the 1s sub-shell, and the electronic
the subshells are show
structure of hydrogen can be written in sub-shell notation as 1s1. The electronic
structure of helium is 1s2, then lithium is 1s2 2s1 and so on.
The Aufbau principl
out the electronic co
Marginalizer
Electronic Structure
Bilal Hameed
20
Bilal Hameed
shell
1s
n=1
19
First quantum
shell
So, the single electron in a hydrogen atom goes in the 1s sub-shell, and the
ELECTRONIC
STRUCTURE
19
electronic
structure of hydrogen
can be written in sub-shell notation as 1s1.
Following on, the electronic structure of helium is 1s2, then lithium is 1s22s1
The electron and
shellsoand
on. sub-shell structures of the first 20 elements in the periodic
electron shell and sub-shell structures of the first 20 elements in the
table are shownThe
below.
periodic table are shown in Figure 5.6.
alt
Figure 5.6 !
The electron shell and
sub-shell structure of the
first 20 elements in the
periodic table.
Period 1
H
He
Atomic no.
1
2
Electron
shell structure
1
2
Electron
sub-shell
structure
1s1
1s2
Period 2
Li
Be
B
Atomic no.
3
4
5
Electron
shell structure
2, 1
2, 2
2, 3
Electron
sub-shell
structure
1s2
2s1
1s2
2s2
Period 3
Na
Atomic no.
1
C
N
O
F
Ne
6
7
8
9
10
2, 4
2, 5
2, 6
2, 7
2, 8
1s2
2s22p1
1s2
2s22p2
1s2
2s22p3
1s2
2s22p4
1s2
2s22p5
1s2
2s22p6
Mg
AI
Si
P
S
Cl
Ar
11
12
13
14
15
16
17
18
Electron
shell structure
2, 8, 1
2, 8, 2
2, 8, 3
2, 8, 4
2, 8, 5
2, 8, 6
2, 8, 7
2, 8, 8
Electron
sub-shell
structure
1s2
2s2 2p6
3s1
1s2
2s2 2p6
3s2
1s2
2s2 2p6
3s2 3p1
1s2
2s2 2p6
3s2 3p2
1s2
2s2 2p6
3s2 3p3
1s2
2s2 2p6
3s2 3p4
1s2
2s2 2p6
3s2 3p5
1s2
2s2 2p6
3s2 3p6
Period 4
K
Ca
Atomic no.
19
20
Electron
shell structure
2, 8, 8, 1
2, 8, 8, 2
Electron
sub-shell
structure
1s2
2s2 2p6
3s2 3p6
4s1
1s2
2s2 2p6
3s2 3p6
4s2
Understanding
the pattern in ionisation energies
Notice that the electron structure of calcium is 2,8,8,2. In sub-shell notation, this
becomes 1s2 2s2 2p6 3s2 3p6 4s2 .
8,8,2. In sub-
Calcium
including 4s
covers the
ic table.
m, is therefore
ure 5.7).
m, this would be
. (Hint: for
Shell
structure
2,
8,
8,
2
Sub-shell 1s2, 2s2 2p6, 3s2 3p6, 4s2
structure
Iron
For calcium (atomic number Z = 20), all sub-shells up to and including 4s are filled.
Shell
2,
8,
14,
2
structure
Sub-shell 1s2, 2s2 2p6, 3s23p63d6, 4s2
structure
Figure 5.7 !
The relationship between the shell and
sub-shell structures of calcium and iron.
Bilal Hameed
Bilal Hameed
Marginalizer
21
Electronic Structure
Understanding the pattern in ionisation energies
20
Electrons fi
the lowest energy level
upwards –
f calcium is 2,8,8,2.
Inll sub-levels
sub- from
Calcium
ELECTRONIC STRUCTURE
20
this gives the lowest possible (potential) energy.
Shell
2,
8, into the
8, 3d sub-shell,
2
The
next 10 electrons
go
which covers the elements from
structure
hells up to and
including
4s configuration
Thus the
full electronic
of
sodium
(11
electrons)
can
be
scandium (Z = 21) to zinc (Z = 30) in the periodic table.
b-shell, which
covers
the
built up
as follows:
2
2 2:6this sub-level
2
6
The
fi
rst
two
the
1s sub-level
is2 places after calcium, is therefore 2,8,14,2,
Sub-shell
) in the periodic table.electrons go into
The
electron
structure
of iron,
1s
, 2s 1s
2p
, 3swhich
3p , is4ssix
now full.
structure
ces after calcium, is therefore and its sub-shell structure2 is 1s2 2s2 2p6 3s2 3p6 3d6 4s2
The next two electrons go into the 2s sub-level
2s : this sub-level is
3p63d64s2 (Figure
now full.5.7).
Iron
6
The next six electrons go into
the
2p
sub-level
2p
:
this
sub-level is
Shell
2,
8,
14,
2
now full.
structure
1
The last electron goes into the 3s sub level
3s : the full electronic
configuration of sodium is thus 1s22s22p63s1. This can also be
(Hint: for sodium,
this would be
6
Sub-shell
the electronic
abbreviated to [Ne]3s1, where
1s2,confi
2s2 guration
2p6, 3s2of
3pthe
3d6, 4s2
previous noble gas is assumedstructure
and everything after that is given in full.
The full electronic configuration of iron (26 electrons) is:
2 2 6 2 6 2 6
Figure
!sub-level is lower in
2p 3supwards
3p 4s 3d–. Note that,
because5.7
the 4s
1s
the lowest energy2slevel
The
relationship
shell and
energy
than
the
3d
sub-level
it
is
fi
lled
fi
rst.
In
otherbetween
words, twothe
electrons
(potential) energy.
go into the fourth main energysub-shell
level beforestructures
the third main
energy
level
is iron.
of calcium and
2 6
3d
.
fi
lled.
This
can
also
be
written
as
[Ar]4s
ion of sodium (11 electrons) can be
alt
1
ms in Question 4. (Hint: for
6 2
is2sometimes
written
the 1s sub-levelThis 1s
: this sub-level
is as [Ar]3d 4s to keep the sub-levels in order
of the main energy levels.
ns.
2s2: this sub-level is
The full electronic configuration of tin (50 electrons) is:
2 2 6 6 2 6 2 10 6 2 10 2
2p: this
3s 3p
4s 3d 4p
he 2p sub-level1s 2s2p
sub-level
is 5s 4d 5p . Or, in abbreviated form:
[Kr]5s24d105p2.
in which the sub-levels are filled can be remembered most
3s sub level
3s1The
: the order
full electronic
2 2 6 1 easily from the periodic table (Figure 2.21).
1s 2s 2p 3s . This can also be
he
electronic
confioutermost
guration of the
ctures
in their
shells.
nd everything after that is given in full.
Skill Check 9
n of iron (26 electrons) is:
Write the electronic sub-shell structures of the following atoms.
because the 4s sub-level is lower in
filled first. In other words, two electrons
(a) Oxygen
vel before the third main energy level is
2 6
(b) Neon
Ar]4s 3d .
the 2s sub-level
d64s2 to keep the sub-levels in order
n of tin (50 electrons) is:
elop
ideas about
p2. Or,their
in abbreviated
form:
the
sses and charges of protons
are fiprobability
lled can be remembered
most
eelsthe
of finding
ure 2.21).
ations have led chemists to
an electron or a pair of
an atom. These regions are
n at millions of nanosecond
ng the electron ‘smeared out’
se smeared-out pictures are
lectron density maps. The
ely to be and lightest where
(c) Silicon
(d) Potassium
(e) Titanium (Z = 22)
Skill Check 10
Definition
Write the electronic sub-shell structures of the following ions.
An orbital is a region around the
(a)nucleus
Be2+ of an atom that can hold up
+ electrons with opposite spins.
two
(b)toNa
(c) Cl—
(d) Ca2+
(e) Br—
electron density plots by
re is a 95% chance of finding
Marginalizer
ins.
Electronic
Structure
described
as a spherical
ring
The p sub-shells contain
ped like dumbbells and
Bilal Hameed
22
Bilal Hameed
21
ELECTRONIC STRUCTURE 21
Skill Check 11
Identify the elements with the following electron structures in their outermost shells.
alt
(a) 1s2
(b) 2s2 2p2
(c) 3s2
(d) 3s2 3p4
(e) 3d3 4s2
(f) 4p3
Electrons and Orbitals
1
Chemists have used complex mathematics to develop their ideas about the
arrangement of electrons in sub-shells. These calculations have led chemists to
believe that electrons do not occupy fixed positions within an atom, nor do they follow
orbits in the shells. Electrons occupy volumes or regions of space called orbitals.
There is a high probability of finding an electron or a pair of electrons in certain regions
around the nucleus of an atom.These regions are called orbitals.
By pinpointing the likely position of an electron at millions of nanosecond intervals, it is
possible to build up a picture showing the electron ‘smeared out’ over its orbital as a
negatively charged cloud. These smeared-out pictures are sometimes described as
electron density plots or electron density maps. The plots are darkest where the
electrons are more likely to be and lightest where the electrons are less likely to be.
The idea of an electron as a particle is replaced by a description based on the electron
as a cloud of negative charge. This charge is spread like a mist around the nucleus of
the element. The mist varies in its density at different distances from the nucleus.
These regions of ‘electron mist’ are called orbitals and correspond to the energy
levels used by an atom. Each orbital can hold one or two electrons and has a density
of charge that depends on the energy level being described. If the orbital has two
electrons the two electrons will differ in the direction in which they spin.
Bilal Hameed
Bilal Hameed
Marginalizer
23
Electronic Structure
of fin
22
An orbital can contain a
s-orbitals
two
electrons.
The maximum
simplest type of orbital toof
describe
is the orbital
corresponding to an ‘s’ sub-shell
ELECTRONIC STRUCTURE 22
(sub-energy level). It consists of a spherical volume of negative charge with the
nucleus at its centre. All s-orbitals have this shape.
alt
There a
The fi
and this
2.24a).
The 1
is movin
probabil
The dar
that poi
The e
nucleus
certain d
The s
made up
consists
orbitals.
bigger t
p orb
up the 2
appropr
x axis. T
as dege
For example, Selenium (Se) is in p
Note: all atoms in the same group
therefore the last part of the electron
(vertical column) in the periodic
electronic configuration can be wor
Atomic and electronic structure
table have the same outer shell
H He
1s2
electronic configuration. For
Li Be
2s2
example, all the elements in group
Na Mg
3s2
6 (like Se)
have theinouter
K Ca
4s2
Sc Zn
The shape of a 1s orbital on the left, and the electron
density
a 1sshell
orbital
on
the
electronic configuration ns2np4.
(remember
to
go
down
1 in the d
right. The difference between a 1s-orbital andnais2s-orbital
is the distance from the
the period number.
Therefore the electronic configur
1
2.12
Electrons and orbitals
nucleus to where the major density of charge is concentrated. In the 1s-orbital, the Figure 2.23 shows an alternative w
In this section you will learn to:
So far, we have described the electrons of atoms in shells and sub-shells at
which sub-levels are filled.
electron density is closer to the nucleus. As the
levels (principal
quantum
numbers)
increasing
distances from
the nucleus.
• Outline the concept of orbitals
and theat
probability
of finding
increase, the radius
which the
chargeanis most
densehave
becomes
further
away
from to calculate the probability
Chemists
also used
complex
mathematics
electron in a particular area of
of
finding
an
electron
at
any
point
in
an
atom. Their calculations have led
the nucleus. For
example, a 4s-orbital is larger than a 3s because the point at which
an atom
13 Give
theelectron
full electronic
confi
chemists to believe that there is a high probability
that an
or a pair
o
the 4s-orbital
has the
its greatest
density of negative
charge
further
fromregions,
the nucleus
a Naround the nucleus
b Ar of
• State
‘Aufbau principle’
electrons
will is
occupy
certain
called orbitals,
an atom.
d Sr
e Te
Draw
‘electrons-in-boxes’
than it is for •the
3s-orbital.
diagrams to show electronic
arrangement of some common
elements
DEFINITION
An orbital is a region in space
around the nucleus of an atom
occupied by an electron or a pair
of electrons with opposite spins.
By calculating the probability of finding an electron in different regions of the
orbital, it is possible to compose a density map showing how the electron is
distributed throughout its orbital. Figure 2.22 shows one such electron densit
map for the 1s and 2s electrons in a lithium atom. The density maps are darke
where the electrons are most likely to be and lightest where the electrons are
Electrons in atoms occupy atomic o
least likely. The charged cloud for the two 1s electrons and that for the single
electron are both spherical in shape.
An orbital is a region of space i
The general shapes of orbitals are deduced from electron density plots by
of finding
ana electron.
It repres
determining the boundary of the region in which
there is
95% chance
of
finding an electron or a pair of electrons. As a result of these studies, chemists
believe that all s sub-shells contain one orbital,
as a types
spherical
Therebest
are described
four different
of ato
An orbital can contain a
annulus – like extra thick peel on an orange. The
On the
p orbitals
ar
firstother
shell hand,
(maximum
number
electrons. ‘dumb-bell’ shaped with the nucleus located
notmaximum
spherical, of
buttwo
approximately
and this makes up the entire 1s subbetween the two halves of the dumb-bell.
2.24a).
In fact, each p sub-shell has three separate p orbitals,
each of
The 1s orbital
is which
centredcan
on hold
the n
maximum of two electrons. This makes a total
of six all
electrons
a filled
is moving
the timeinand
the inten
p sub-shell.
probability of finding the electron at
1s
Figure 2.23 shows the shapes of the three orbitals
in athe
p sub-shell.
are th
The darker
colour theThey
greater
identical except for their axes of symmetrythat
which,
like
therepresents
axes of a the
threepoint.
This
electro
dimensional co-ordinate system, are mutuallyThe
at right
angles.
it isanywh
electron
can Thus,
be found
convenient to distinguish between the p orbitals by labelling them px, py and
nucleus – at the centre of the orbital
pz. Electrons will, of course, occupy the three p orbitals singly at first because
certainisdistance
from
nucleus.one
their mutual repulsions. When a fourth electron
added to
a p the
sub-shell,
The second main energy level (m
of the three orbitals will contain a pair of electrons.
made up of the 2s sub-level and the
consists of thez2s orbital, whereas the
orbitals. The 2s orbital (like all other
bigger than the 1s orbital (Figure 2.2
p orbitals have a ‘dumb-bell’ shape
x
x
x
up the y2p sub-level. These point at 9
y
y
px
py appropriately aspzpx, py, pz (Figure 2.2
x axis. The three 2p orbitals all have
The shapes and relative positions of the three p orbitals in a p sub-shell
as degenerate.
z
2s
Fig 2.22 An electron density map for the
charge clouds of electrons in a lithium atom
Fig 2.23
z
All s orbitals can be represented as spheres. They differ only in size and energy. The
So, the first quantum shell contains just one sub-shell (1s) with one s orbital.
3s orbital is larger than the 2s orbital, which is The
larger
thanquantum
the 1s orbital.
The larger
second
shell contains
two sub-shells (2s and 2p) with a total of
four orbitals, one 2s orbital and three 2p orbitals (2px, 2py and 2pz).
orbitals are described as being more diffuse since
the electron density is less.
38
Marginalizer
Electronic Structure
Bilal Hameed
24
Bilal Hameed
The generalhave
increase
in ionization
each period
is due
to the
increase
in turn
nuclear
a stronger
shieldingenergy
effect across
than electrons
in the
second
shell,
which in
havecharge.
a stronger
This occurs shielding
because across
the period
each
has oneinadditional
proton,
which
effect than
electrons
inchemical
the third element
shell. Electrons
the same shell
exert
a relatively
small
shielding
effect
increases the
nuclear
charge
byon
+1.each other.
23 energies for the chemical elements of periods 1, 2 and 3.
Figure 12.14 shows the first ionization
The
general
increase
in
ionization
energy
across each
d as spheres. They differ only in size and energy. The 3s orbital
is period is due to the increase in nuclear charge.
This
occurs
because
across
the
period
each
chemical
h is larger than theNe1sELECTRONIC
orbital. TheSTRUCTURE
larger orbitals
as element has one additional proton, which
23 are described
shielding
increases the nuclear charge by +1.
ctron density is less.
force
p
orbitals
2500
e
First ionization energy/kJ mol–1
es forming He
a ‘dumb-bell’
shape and have
different orientations Li p . The
Be p-orbitals
B
C moreN
Ar three orbitals labelled px, py and
The p sub-shells contain
are
z
Ne
ight angles to each other
and
are
labelled
p
,
p
and
p
to
reflect
x They
y have an
z elongated
complex and harder to visualise.
dumbbell shape and a
shielding
2000
rbitals
all have the same
energy
–
the
orbitals
are
said
to
be
force
variable charge density. The dumbbell has two lobes and between
the two lobes
e the same shape as there
the 2p
butthe
areprobability
larger.
Li
Be is not
B
C
is a orbitals
node at which
of
finding
an
electron
is
zero.
This
Ar
electrostatic
1500
2.2
Electron
configuration
77
as the
between
the twoguide,
lobes contains
the 3d orbitals is notsurprising,
required
by region
the IB
Chemistry
but the nucleus.
attraction
H
3.)
towards
For each principal quantum number there are three p-orbitals. They are identical and
Li
Na shells, which are involved in thepositive
sShapes
are1000
thoseof
in orbitals
the outer
electrostatic
K
have the same energy, differing
only in their orientation
in space. They are arranged
nucleus
attraction
ovalent
bonds
arefixed
formed
when
atomic
orbitals
overlap
and
trons
do not
occupy
positions
within
an
atom,
nor
do
they
follow
orbits
in the shells.
at right angles to each other. They are labelled ‘x’, ‘y’ and
‘z’ to correspond
to the
towards
500
trons
occupy
volumes
orthree
regions
of
space
called
orbitals
(Figure
2.55).
The
four
types
of
s (Chapter
14).
5
10
15
20
a py orbital is aligned with
principal
Li
Naaxes. (px orbital is aligned with the x-axis,positive
O
F
Ne
alt
N
O
F
Ne
1
Figure 12.15 A diagram illustrating how the balance
tals, s, p, d and f,Atomic
all have
different
shapes. (The shapes Kand energies
of atomic
nucleus orbitals are
number,
z-axis. ) shielding
the
y-axisZand a pz orbital is aligned with the
between
and nuclear charge changes across
fined
an bys 0orbital
the pwave
,
p
and
p
orbitals
solving theand
Schrödinger
equation.)
x
y
z
1
5
10
15 3
4 First ionization
energies
for periods
1, 2 and
20 period 2
herical with a nucleus located
at the centre. The radii of the Figure 12.15 A diagram illustrating how the balance
Atomic number, Z
between shielding and nuclear charge changes across
al quantum The
number
(n). Figure
2.56charge
showsincreases
the 1s and
2sforce
orbitals.
increase
in
nuclear
the
of attraction
on all the electrons, so they are
Figure 12.14 First ionization energies for periods 1, 2 and 3
period
2
r and2phenceheld
the electron
is lessstrongly.
than theEach
1s orbital.
closer anddensity
hence more
additional electron across a period enters the same
bell
shape. The three pThe
orbitals
arehence
alongincreases
the
x, ythe force
increase
inarranged
nuclear
attraction
on all12.15).
the electrons, so they are
level)
and
thecharge
increase
in shielding
is of
minimal
(Figure
boron atom, Bshell (energy
shows
three
p orbitals
of
an
atom.
The
sizes
of
the
p
orbitals
are
held
closer
and
hence
more
strongly.
Each
additional
electron
across
a
enters the
same
1s
2s
2p
Although the general trend is for the ionization energy to increase acrossperiod
the period,
there
shell
(energy
level)
and
hence
the
increase
in
shielding
is
minimal
(Figure
12.15).
bital.2p
are two distinct dips in ionization energy across periods 2 and 3 (Figure 12.14). These dips can
Although the general trend is for the ionization energy to increase across the period, there
can
only
be
explained
using
an
orbital
model
of
electronic
structure.
er than the 2swhich
orbital,the
which
is larger
thanand
the
orbital.in
The
larger orbitals
are These
described
as a greater effect than the
py a1schange
electron
is lost
electron
shielding.
have
Thedensity
first decrease
in each period is the result of a change in the sub-shell (sub-level) from
,gBe
more diffuse
since
the
electron
is
less.
increase in nuclear
charge
andis decrease
atomic
radius.
In
period
2,
first
occurs
z
Figure
2.28
shows
orbitals that
make
up the
2sthis
and have
2p
sub-levels
in
zchange
which
the
electron
lost and
apin
inthe
electron
shielding.
These
adecrease
greater
effect
than the
The
p2porbitals
have
lobes
forming
a ‘dumb-bell’
shape
and
have
orientations
22s2) loses
beryllium
atom,between
Be twothe
the
second
main
energydifferent
level.
elements
beryllium
and
boron.
When
it
is
ionized,
the
beryllium
atom
(1s
increase
in
nuclear
charge
and
decrease
in
atomic
radius.
In
period
2,
this
first
decrease
occurs
1s They
2s are arranged
2p
pace.
at right angles to each other and
py and pz to reflect
22plabelled
1) loses pax,2p
2 2 is
electron
(Figure 12.16). More energy
a 2s electron,
whereas
aelements
boron atom
(1s22sare
between
the
beryllium
and
boron.
When
is ionized,
r orientation. The three p orbitals all have the same energy – the2 orbitals
areitsaid
to be the beryllium atom (1s 2s ) loses
2
1
Be+
required
to remove
an electron
theatom
lower
energy
orbital
in beryllium
than
from
theenergy
higheris
a 2s electron,
whereasfrom
a boron
(1s
2sare
2p larger.
)2s
loses
a 2p electron
(Figure
12.16).
More
enerate.
3p
but
x +orbitals have the same shape as the 2p orbitals
2p The
pxand
beryllium
ion, Be
required
to
remove
an
electron
from
the
lower
energy
2s
orbital
in
beryllium
than
from
the
higher
energy
2p
orbital
in
boron.
Although
the
2s
2p
sub-levels
are
in
the
same
shell,
the
energy
(Knowledge
shapes of the 3d orbitals is not required by the IB Chemistry guide, but
1s
2s of the2p
energy
2p
orbital
in
boron.
Although
the
2s
and
2p
sub-levels
are
in
the
same
shell,
the
energy
relatively large. Recall (Chapter 2) that the energy gap between shells and sub-levels
are describeddifference
in Chapteris13.)
difference
isan
relatively
large.
Recall
(Chapter
that in
thethe
energy
betweeninshells
andsubsub-levels
becomes
smaller
in
shell
number.
In2)addition,
a singlegap
electron
the 2p
The
most important orbitals arewith
those
in increase
the outer
shells,
which
are involved
6 Orbital
becomes
smaller
with
an
increase
in
shell
number.
In
addition,
a
single
electron
in
the
2p
2
Figure
12.16
Orbital
(Figure 12.17). sublevel
is moreCovalent
effectively
shielded
by the
inner
electrons
than
the 2s
mation
of chemical
bonds.
bonds
are formed
when
atomic
orbitals
overlap
andelectrons
boron
and
2 electrons (Figure 12.17).
level
is
more
effectively
shielded
by
the
inner
electrons
than
the
2s
boron and
genotations
to form for
molecular
orbitals (Chapter 14).
boron ion, B+ only be explained using an orbital model of electronic structure.
two
distinct
dipsdiffer
in ionization
energy
across
periods
and 3 (Figure
12.14). These dips
1s
2scan be represented
2p firstare
stalorbitals
as
spheres.
only
in size
andof
energy.
Thein
3s2the
orbital
is
The
decrease
in They
each
period
is the
result
a change
sub-shell
(sub-level)
from
ms and
ve
ions
ognizing
atoms and
ndberyllium
2s orbitals
■ Figure 2.57 Three p orbitals
theions
shape
of an s orbital and theThepxthird
, pshell
pz orbitals
(maximum
182p
electrons) consists of the 3s, 3p and 3d
y and
their unipositive
2s
2s
2p the 3p sub-level consists
sub-levels.at
The
3s sub-level
just the
3s orbital;
s atomic orbital is always spherical
with a nucleus located
the
centre.isThe
radii
of the
of three 3p orbitals; and the 3d sub-level is made up of five 3d orbitals.
bitals increases with principal quantum number (n).One
Figure
2.56
shows
the
1s
and
2s
of the five 3d orbitals is shown in Figureorbitals.
2.29.
model
e that the 2s orbital is larger and hence the electron density
less(maximum
than the
1s orbital.
The fourthisshell
32 electrons)
consists of one 4s, three 4p,
ve 4d and are
sevenarranged
4f orbitals.along
The seven
orbitals
The
have model
a dumb-bell
shape.
three function
p fiorbitals
the4f x,
y make up the 4f subel
is pa orbitals
probability
(which
usesThe
a wave
level. One ofto
the describe
f orbitals is shown in Figure 2.30.
z axes in space. Figure 2.57 shows three p orbitals of an atom. The sizes of the p orbitals are
the electron). The orbitals described previously are drawn as
< 3p < 4p. There is no 1p orbital.
Within any sub-shell all the orbitals have the same energy (they
ns spend 95% of their time. A more accurate description
ofe.g. the three 2p orbitals are degenerate and the
are degenerate),
five 3d orbitals are degenerate.
2s orbital
y
py 2s
Figure 12.17 Electron density
cloudsofof
the
2p orbitals
lobe
The number
orbitals
in and
each energy
level is (only
shown one
in Table
2.4.shown). The dotted line
Figure 12.17 Electron density clouds of the 2s and 2p orbitals (only one lobe shown). The dotted line
z
Theshows
most the
important
those inthe
the2s
outer
shells,can
which
in thethe 1s orbital, increasing its
ppartially
extentorbitals
of the are
1s orbital;
electron
penetrate
z are involved
shows the extent of the 1s orbital; the 2s electron
can partially
penetrate
the 1s orbital,
increasing
The diagrams
of atomic
orbitals
thatorbitals
for more complex
atoms. its
What implications does this
formation
of chemical bonds. Covalent bonds
are formed
when
atomic
stability
orbital
y
z
we have seen here are derived from
stability overlap and merge to form molecular orbitals.
x
x
Bilal
B_Chemistry_435-450.indd 443
Hameed
Figure 2.56 Shapes of the 1s and 2s orbitals
0.indd 443
Bilal Hameed
antum mechanical model
mathematical functions that are solutions
to the Schrödinger equation. Exact solutions
of the
px
Schrödinger equation are only possible for a system
involving one electron, i.e. the hydrogen atom. It is
not possible to derive exact mathematical solutions
have for the limit of scientific knowledge? When we
describe more complex atoms in terms of orbitals,
we are actually just extending the results from the
hydrogen atom and gaining an approximate view of
the properties of electrons in atoms.
18/05/15 9:26Marginalizer
am
■ Figure 2.57 Three p orbitals
25
18/05/15 10:35
18/05/15
Electronic Structure
10:35 am
Figure 2.12 A 1s orbital.
z
to an ‘s’ energy level. It consists of a spherical volume of
negative charge with the nucleus at its centre. All s-orbitals have
this shape. The difference between a 1s-orbital and a 2s-orbital is
the distance from the nucleus to where the major density of charge
24
is concentrated. In the 1s-orbital,
the electron density is closer to
the nucleus. As the levels (principal quantum numbers) increase, the
radius atSTRUCTURE
which the charge
ELECTRONIC
24is most dense becomes further away from
the nucleus. For example, a 4s-orbital is larger than a 3s because
Orbitalsthe
and
sub-shells
point
at which the 4s-orbital has its greatest density of negative
charge
is
further
from energy
the nucleus
than
it is for the 3s-orbital.
A group of orbitals with
the same
is called
a subshell.
z
There is only one s orbital in an s subshell,
alt
2.12 E
p-orbitals
there xareThe
threep-orbitals
p orbitals in
a pmore
subshell,
are
complex and harder to visualise
(Figure
Theyand
have an elongated dumbbell shape and a
five d orbitals
in a2.14).
d subshell
x
variable charge density, with the area of greatest concentration
seven f orbitals
in an with
f subshell.
increasing
the NUMBERS
distance from
the nucleusINasORBITALS
the principal
TABLE
1.3.1 MAXIMUM
OF ELECTRONS
quantum
increases.
For each
y Main
energynumber
Subshells
Orbitals
in
principal
quantum
number
Maximum
Maximum
there
They
are identical
and
the same
levelare three p-orbitals.the
subshell
number
of have
number
of
electrons
in They
electrons
energy, differing only in their orientation
in space.
are in
the subshell energy level
labelled ‘x’, ‘y’ and ‘z’ to correspond to the three principal axes.
y
Electron configurations
1s-orbital
2s-orbital
Figure 2.13 The shape of 1s- and 2s-orbitals.
1
1s
z2
2s
1
2p
3
Tip
y
x
4
Orbitals that have the same
energy are said to be degenerate.
px-orbital
1
2
2
z1
2
z 8
3
6
3s
1
3p
3
3d
5
4s
1
4p
3
y
4f
Figure 2.14 The shape of 2p-orbitals.
18
y
10
x
2
x
32
6
py-orbital
5
4d
2
6
pz-orbital
10
7
14
n
2
2
Representing the electron configuration of the
elements
Tip
Fig 2.24
Suppose
magnified
an atom
one million
times toan
the‘electrons
size of thein
2010
World
From
these
ideasweabout
orbitals,
chemists
havemillion
developed
boxes’
Cup Stadium in Johannesburg. The nucleus would be the size of a pea at the centre of the pitch
For each principal quantum
and the outermostfor
electrons
would bestructures
moving around
at the edges
of therepresents
stadium an orbital
representation
the electronic
of atoms.
Each box
number of 3 and above there are
Electron
energy
levels
can
be
represented
in
‘box’
diagrams,
as shown in
and each orbital can contain a maximum of two electrons.
five d-orbitals.
It is not
necessary
CHEM
COMPLEMENT
Figuretheir
2.15.knowledge about sub-shells and orbitals, chemists have developed
From
to know their shapes at this stage.
an ‘electrons-in-boxes’ notation for the electronic structures of atoms. Using
Electron spin
this notation, each box represents an orbital. Using a set of three rules called
4d
Detailed study of the line spectra
of ‘Aufbau
many-electron
atoms it is spin.
This gives
two possible
values
for a fourth
quantum of all
the
principle’,
possible
to work
out the
electronic
structure
5s
(i.e. atoms other than hydrogen)atoms.
showedIn
that
each
line
was
number,the
spin
magnetic
quantum
number,
ms.rules
German,
the
word
Aufbau
means
‘build
up’.
The
three
in the
Tip
4p
Note
that
the
actually a pair of very closely spaced
lines.principle’
This was explained
This quantum number completes the description of an
‘Aufbau
are:
5s-orbital
is
at
3d
Energy
twothe
Dutch
physicists
electron and allows the Pauli exclusionaprinciple
to be
The reasonsbywhy
orbitals
areGeorge Uhlenbeck and Samuel
lower energy
4sNo first.
1 Otto
Electrons
enter
the orbital
of lowest
energy
Goudsmit after an experiment by
Stern and
Walther
completely
satisfied.
two electronslevel
can have
the
than
thesame
4d set
filled in this Gerlach
order are
complex
Note thattwo
the
3p one
in 1922 in which electrons
were deflected
in a either
2 Orbitals
can hold
electron
spins.
of four
quantumor
numbers.electrons
This resultswith
in theopposite
need for two
and you willmagnetic
not be asked
tofigure 1.3.2).3 Electrons occupy3sorbitalselectrons
4s-orbital is at
field (see
at the same
sub-level
singly
they pair
up.
occupying
the same
orbital before
to have opposite
spins.
provide any Electrons
explanation.
were postulated to have an intrinsic property, called
a lower energy
These are represented in orbital diagrams (see figure 1.3.1)
level than
the 3d notation and also the subFigure 2.25 shows
electrons-in-boxes
(orbitals)
2p asthe
by arrows pointing in opposite directions.
electron spin, which causes each electron to behave
if it
shell
(s, p,
d, f)generates
notationa for the electronic structures of beryllium, carbon,
were spinning about an axis. The
spinning
charge
2s
oxygen.
magnetic field whose direction nitrogen
depends onand
the direction
of These representations of the electronic structures of
atoms1sare often called electronic configurations.
Figure 2.15 Energy-level diagram.
S
Electron sub-shell
1s
slit
OCR_A_Level_Chemistry.indb 22
beam of
atoms
2s
–
beryllium
magnet
2p
s,p,d,f electron
notation
30/03/15
1s2 2s2
collector plate
N
nitrogen
Figure 1.3.2 Atoms in which the electron spin quantum number ofoxygen
the unpaired
1s2 2s2 2p3
2
2
4
Figure 1.3.3 The direction in which an electron1sspins 2s 2p
electron is 21 are deflected in one direction and those in which it is 21 are
determines the direction of its magnetic field.
Fig 2.25 An electrons-in-boxes notation and a sub-shell notation for the electronic configurations
deflected in the other.
of beryllium, carbon, nitrogen and oxygen
14
Marginalizer
Electronic Structure
QUESTIONS
21 In the electron
notation of ele
any two arrow
box always po
directions. Wh
22 Write the ‘elec
structure for t
elements:
a lithium
b fluorine
2:43 PM
c sulfur
d calcium
1s2 2s2 2p2
carbon
beam
KEY POINT
In a stable atom,
the lowest availab
(sub-shells). Whe
half-filled, furthe
to pair up.
23 For the shell o
quantum num
down:
a the number
b the total nu
c the maximu
electrons in
Bilal Hameed
26
Bilal Hameed
1s
2s
It can be written as 1s2. The two electrons must fo
The lithium atom (atomic number 3) has three
(as a spin pair). The 1s orbital is now full; so the t
orbital with the lowest energy). This is in accorda
configuration is:
Some examples of electron configurations using ‘electrons in box’
diagrams are given in Figure 2.16. It25
has been mentioned that where an
orbital contains two electrons they differ in their direction of spin. To
distinguish between the two electrons they are shown with an upward
ELECTRONIC
STRUCTURE
25 ( ).
arrow ( ) and
a downward arrow
x
y z
2p
x
y z
2p
1s
Energy
Energy
2s
It can be written as 1s1. The large number represe
number), the letter represents the sub-shell and th
of electrons in the sub-shell.
The detailed electron configuration of the heli
1s
Lithium
2s
1s
Boron
Carbon
■ Figure 2.61
Electron spin
Energy
alt
2s
1s
Figure 2.16 ‘Electrons in box’ diagrams for the electron configurations of lithium, boron
carbon.
Fillingand
atomic
orbitals
1s
diagrams of this kind you can
so represent the electrons
ing half arrows.
1
The electrons
are arranged
in atomic
orbitals
certain
principles:
Box diagrams
are a useful
visual
way ofaccording
showing to
how
electrons
are
S
b anticlockwise
The detailed electron configuration of the hydrog
distributed but it is often more helpful to identify the orbitals using the
Each orbital
hold The
up tosystem
a maximum
of two
This, in simplified form, is the
orbitalcan
names.
for doing
thiselectrons.
is as follows.
Pauli exclusion principle.
● An atom of hydrogen has one electron, which in its ground state is in
Electrons the
enter
and occupy
an empty atomic
with the
This is
(You would
saylowest
this asenergy.
‘one s one’)
1s-orbital.
It is represented
as 1s1.orbital
2.2 Electron configuration 79
known as the Aufbau principle.
● An atom of helium has two electrons that occupy the 1s-orbital. Its
■ Within a sub-shell, electrons experience repulsion a
S
2
. (You
would
say
this
assame energy. This is known as Hund’s rule.
electron configuration
is represented
as 1s
Within
a
sub-shell,
electrons
experience
repulsion
and
hence
enter
two
different
■ Within
sub-shell, electrons experience repulsion and hence
enter two different orbitals of the
‘one satwo’)
orbitals
of the
same energy.
This
is known
as
Hund’srule.
rule.
■ Electrons behave as particles and hence possess a c
same
energy.
This
is
known
as
Hund’s
● An atom of lithium has three electrons and its electron structure iselectron can spin in two different directions: clock
2s1and
.particles
represented
as 1s2as
■ Electrons
and hence
possess
a classical
property
as spin.inAn
Electrons
behave behave
as particles
hence possess
a classical
property
knownsymbols
as
spin. orknown
. Two electrons
the same orbital m
N
2
2
electron
can
spin
in
two
different
directions:
clockwise
and
anticlockwise,
shown
by
the
An electron
can
spin instructure
two different
directions:
and anticlockwise,i.e.
shownand not .
● The
electron
of beryllium
is clockwise
1s 2s .
N
a clockwise
N
■ Electron configurations of ato
When a charged particle spins on its axis, a magnet
have magnetic properties. Protons show similar beh
field. This property is exploited in the technique of
S
Electron energy levels and orbitals
Each box represents an orbital and each orbital can contain a maximum
of two electrons.
■ Within a sub-shell, electrons experience repulsi
same energy. This is known as Hund’s rule.
■ Electrons behave as particles and hence posses
electron can spin in two different directions: cl
symbols or . Two electrons in the same orbit
i.e.
and not .
■ Single electrons in the same sub-shell must ha
i.e.
and not
.
a clockwise
or .. Twohas
electrons
in the same
orbital must have
opposite
spins,
2.2have
Electth
■ Single
electrons
in the same sub-shell must
by thesymbols
symbols
●
An atomorof boron
five electrons and the next highest energy
2p,not
is used. .The electron configuration is represented as i.e.
and not
.
i.e. orbital,
and
2
2
1
2s
2p
,
although
the
subscript
x
is
not
really
necessary
because
1s
■
Within
a
sub-shell,
electrons
experience
repulsion
and
hence
enter two differ
a clockwise
x
S same sub-shell mustNhave the same
When
a charged spin,
particle spins on its axis, a magnetic fi
■ Single
electrons
the
any of
the three in
2p-orbitals
could be usedsame
as they
all have
same(parallel)
energy.
This the
is known
as Hund’s rule.
2.2
have magnetic properties. Protons show similar behavio
label not
x has no direct .meaning
until the behave
y- and z-orbitals
2.2 El
i.e. energy. The and
■ Electrons
as particles
and
hence
possess
a
classical
property
known
field. This property is exploited in the technique of nuc
are also occupied. Labelling is only necessary
to distinguish
between
electron
can spina sub-shell,
in
two different
directions:
clockwise
and hence
anticlockwise,
■ Within
electrons
experience
repulsion
enter twos
N
When orbitals
a charged
particle energy.
spins onSits axis,
a
magnetic
field
is
produced
2.61).
Thus,and
electrons
of equivalent
■ Within
a sub-shell,
electrons
experience
repulsion
and
hence enter two di
S symbols same
or
.energy.
Two electrons
in
the(Figure
orbital
This is known
assame
Hund’s
rule.must have opposite spins,
same energy.
This
is known
Hund’s
rule.
Electron
of atoms
haveEach
magnetic
properties.
Protons
similar
behaviour
to■electrons
andasconfigurations
also
produce
a magnetic
orbital of
equivalent N
energy is show
occupied
by one
electron
before
i.e.
not
.
■Sand
Electrons
behave
as
particles
and
hence
possess
a
classical
property2.2
knE
■ ElectronsThe
behave
as particles
and hence
possess a classical
kno
detailed
electron
configuration
of the property
hydrogen
a
field.the
This
property
isa isclockwise
exploited
the technique
of
nuclear
magnetic
resonance
(NMR).
second
electron
added. Theinreason
for
this is that
two electrons
b anticlockwise
electron
can
spin
in
two
different
directions:
clockwise
and
anticlockw
■ Single electrons
in the
same
mustdirections:
have the clockwise
same (parallel)
spin,
electron
can
spin sub-shell
in two different
and anticlockwi
■ Within
2.2 a
Electron
configuration
79
1s a sub-shell, electrons experience repulsion and hence enter two d
within the same orbital experience
degree
of
repulsion
the
S that makes
or . Two
electrons
thethesame
orbital
haveopposite
oppositespin
sp
. Two
in
orbital
must have
symbols
■ Figure 2.61symbols
energy.
Thisin
is known
assame
Hund’s
rule.must
i.e.
and
not orsame
. electrons
pairing
of electrons
slightly
less
favourable.
N
N
■
Within
a
sub-shell,
electrons
experience
repulsion
and
hence
enter
two
different
orbitals
of
the
Electron spin i.e. i.e.and not
■ not
Electrons
and
.. . behave as particles and hence possess a classical property kn
S
Two
in therule.
same
orbitalNmust havea opposite
spins,
i.e. and
not
■Thiselectrons
configurations
of
atoms
1. Thedirections:
same energy.
isElectron
known as Hund’s
electron
can
spin inaas
two
clockwiserepresents
and
anticlockw
When
charged
particle
spins
on
axis,
magnetic
field
is number
produced
(Figure
2.6
It can
be
written
1sdifferent
large
t
clockwiseclockwise
● Carbon, atomic number 6,atherefore
has
the aground
state
■
Single
electrons
inits
the
same
sub-shell
must
have
the
same
(parallel)
spin
S
■
Single
electrons
in
the
same
sub-shell
must
have
the
same
(parallel)
sp
symbols
or
.
Two
electrons
in the same
orbital
must have
opposite spin
■ Electrons behave as particles
and
hence
possess
a
classical
property
known
as
spin.
An
2
2
1
1
have
magnetic
properties.
Protons
show
similar
behaviour
to
electrons
and
also
p
number),
the
letter
represents
the
sub-shell
and
the
su
The
detailed
electron
configuration
of
the
hydrogen
atom
(atomic
number
1)
is:
2s
2p
2p
.
1s
and
electron can spin in two different directions:
clockwise and anticlockwise, shown by the
b anticlockwise
x
y
i.e.
and not . .
i.e.N i.e.is exploited
and
notinnot
field.
Thisa property
the in
technique
of nuclear magnetic resonance (
of electrons
the sub-shell.
symbols or . Two electrons in the same orbital must have opposite spins,
clockwise
1s
2
2
1
1
1
N
■
Single
electrons
must have
same (parallel)
spi
● Nitrogen, atomic number 7, has the
charged
particle
spinsinonthe
itssame
axis,sub-shell
a magnetic
field the
is produced
(Figure
1s When
2s 2pxa 2p
2pz The
. spins
N ground state
i.e.
and not .
y particle
FigureN2.61
detailed
electron
configuration
the helium
When a charged
on
its
axis,
a .magnetic
field isof
produced
(Figu
i.e.
and
not
have
magnetic
properties.
Protons
show
similar
behaviour
to
electrons
and
als
a clockwise
Singlein electrons
in the same
sub-shell
must have the same
(parallel) spin,
i.e.
■ Single electrons
the same sub-shell
the
same (parallel)
Protons
similar
behaviour
electrons
and
After
nitrogen,must
thehave
electrons
in thespin,
p-orbitals
pairhave
up: magnetic
1s charged
ectron spin
N configurations
field.
This properties.
property
exploited
the
technique
nuclearto
magnetic
resonan
■ Electron
ofinshow
atoms
When a is
particle
spins
on its axis, aofmagnetic
field
is produced
(Figu
i.e.
andnot
not
..
1
and
field.
This
property
is
exploited
in
the
technique
of
nuclear
magnetic
resona
S
It
can
be
written
as
1s
.
The
large
number
represents
the
shell
number
(principal
quantum
have
magnetic
properties.
Protons
show
similar
behaviour
to
electrons
and a
2
2
1
1
N
● spins
oxygen
is 1s
2sb22p
2p
2p
The
detailed
electron
configuration
of
the
hydrogen
atom
(atomic
number
1) is
When a charged particle
on its axis,
a magnetic
field
is
produced
(Figure
2.61).
Thus,
electrons
x
y
z
anticlockwise
field. This property is exploited2in the technique of nuclear magnetic resona
have magnetic number),
properties. Protons
showletter
similar behaviour
to electronsthe
and also
produce a magnetic
■
Electron
configurations
of
atoms
the
represents
sub-shell
and
the
superscript
number
represents
the
number
It
can
be
written
as
1s
.
The
two
electrons
must
form
1s
2
2
S
● neon
field. This property is exploited
in the
of
nuclear
magnetic
is technique
1s22s■22p
2py22.61
2p
. resonance (NMR).
Figure
■ Electron
configurations
x
z
The detailed■
electron
configuration
theatoms
hydrogen
atom3)(atomic
number
The lithium
atomofof
(atomic
has three
ele
23
of electrons in the
sub-shell.
Electron
configurations
ofnumber
atoms
b anticlockwise
S
S
Electron spin
1s
The
detailed
electron
configuration
of
the
hydrogen
atom
(atomic
numbe
(as
a
spin
pair).
The
1s
orbital
is
now
full;
so
the
third
The
detailed
electron
configuration
of
the
hydrogen
atom
(atomic
number
■ Electron configurations
of
atoms
b
anticlockwise
anticlockwise
■ FigureIt2.61
The detailed electron configuration
of the
helium
(atomic
number
2) is: the shell number (princip
can
beb written
as atom
1s1. The
large number
represents
S
1s
1s
The detailed electron configuration of the hydrogen atom (atomic number
1) is: spin
orbital
with the lowest energy). This is in accordance
Electron
b anticlockwise
■ Figure 2.61
■
Figure
2.61
number),
the
letter
represents
the
sub-shell
and
the superscript number represe
1s
1s
It can be written
as 1s1. The large
Electron spin
■ Figure 2.61
configuration
is:1 number represents the shell number (prin
Electron spin
It
can
be
written
as
1s
.
The
large
number
represents
the shell
numberrepr
(pri
of
electrons
in
the
sub-shell.
Electron spin
number), the letter
the
and the
superscript
number
_A_Level_Chemistry.indb
23
2:43 sub-shell
PM
1s1.represents
2s30/03/15
It can beelectron
writtennumber),
as
1s
The
large
shell number
number
(pr
the
letter
represents
the represents
sub-shell
and the
the
superscript
number2)
rep
It can be written as 1s1. The large number represents the shell number (principal quantum
The detailed
configuration
ofnumber
the helium
atom
(atomic
of electrons in
the sub-shell.
number), the letter represents the sub-shell and the superscript
number represents the numbernumber), the letter
of electrons
in thethe
sub-shell.
represents
sub-shell
and
the
superscript
number
re
2
The
electron
of the of
helium
atomatom
(atomic
number
be written as 1s . The two electrons
a detailed
spin The
pair.
1s must form
of electrons in It
the can
sub-shell.
detailedconfiguration
electron configuration
the helium
(atomic
numb
of
electrons
in
the
sub-shell.
2
1
The detailed electron configuration of the helium atom (atomic number 2) is:
1s
It can
be written
as 1senter
2s . the 1s orbital
1sTwo
The lithium atom (atomic number 3) has three
electrons
The electrons.
detailed electron
configuration
of (atomic
the helium
atom 4)
(atomic
numb
1s
The
beryllium
atom
number
has the
det
2. The two
(as a spin pair). The 1s orbital is now full;
electron
enters
the
2s
orbital
next
It canso
bethe
written
as 1s
electrons
must
form
a (the
spin
pair.
2. The
1sthird
2
It
can
be
written
as
1s
two
electrons
must
form
a
spin
pair.
It can be written1s
as 1s2s
. The two electrons must form a spin pair.
It can be written as 1s2. The two electrons must form a spin pair.
Thenumber
lithium
atom
(atomic
number
3) has electron
electrons.
electron
lithium atom
(atomic
3)number
has
electrons.
Two
electrons
en
with
the
energy).
This
isthein1sThe
accordance
with
Hund’s
rule.
Thethree
detailed
The
lithium
atom
(atomic
3) has
three three
electrons.
TwoTwo
electrons
The lithiumorbital
atom (atomic
number
3) haslowest
three electrons.
Two electrons
enter
orbital
(as a spin pair). The 1s orbital is now full; so the third electron enters the 2
(as
a
spin
pair).
The
1s
orbital
is
now
full;
so
the
third
electron
enters
the
2s
2
(as a spin pair).configuration
The 1s orbital is now full;
so
the
third
electron
enters
the
2s
orbital
(the
next
spin pair).
1s the
orbital
iselectrons
now full;
soisthe
thirdaelectron
entersrule.
theorb
2s
It can(as
beawritten
as The
1s with
. The
two
spin
is:
orbital
lowest
energy).
This
in form
accordance
withpair.
Hund’s
Th
22s
2must
orbital with the lowest energy). This is in accordance with Hund’s rule. The detailed
electron
It
can
be
written
as
1s
.
orbital
with
the
lowest
energy).
This
is
in
accordance
with
Hund’s
rule.
The
orbital
with The
the
lowest
energy).
This
is
in
accordance
with
Hund’s
rule.
The
de
configuration
is: number 3) has three electrons. Two electro
lithium
atom
(atomic
configuration is:
1s 2s
configuration is:1sThe
2s boron atom (atomic number 5) has five electro
configuration
is:
1s 2s
(as a spin
pair). The 1s orbital
is now full; so the third electron enters the
1s 2s 1s and 2s orbitals. The fifth electron occupies the 2p o
1s 2s
orbital
with
the
lowest
energy).
This is in accordance with Hund’s rule. T
It can be written as 1s22s1.
It can be written as 1s22s1.
is:
22s1.
configuration is: configuration
The beryllium
atom (atomic number 4) has the detailed electron confi
It
can
be
written
as
The beryllium
atom (atomic
number 4) has
the1s
detailed
electron configuration shown below.
2
1
It can be writtenMarginalizer
as 1s 2s .
Bilal Hameed
2p
22s1. 1s1s 2s 2s
1s 2s
1s 2s
It can
written
as 1sberyllium
The beryllium atom (atomic number
4)be
has
the
detailed
electron
configuration
below.
The
atom (atomic
number 4) has shown
the detailed
electron config
Bilal Hameed
27beryllium1satom
Electronic
Structure
The
number 4)
has
the detailed
electron configurat
2s (atomic
22s2.
22s2.
It can be written as 1s1s
It
can
be
written
as
1s
2s
22s22p1.
be1. written
as 1snumber
It can be writtenItascan
1s22s
The boron atom (atomic number 5) has five electrons. The first four electrons occupy
1s the2s
The
boron
atom (atomic
5) has five electrons. The first four el
1s and 2s orbitals. The fifth electron occupies the 2p orbital. The correct detailed electron
The
carbon
atom
(atomic
number
has
six correct
electr
1s atom
and
2s1s
orbitals.
fifth electron
occupies
the 2p6)
orbital.
The
22s2. The
The
beryllium
(atomic
number
4) has
the
detailed
electron
confd
It
can
be
written
as
configuration is:
configuration is: is:
configuration
1s 2s
2 2
The boron
atom (atomic number 5) has five electrons. The first four ele
S
N
NN
a
clockwise
a clockwise
■
■
NN
i.e.
and not
not ..
i.e.
and
Single electrons
electrons in
in the
the same
same sub-shell
sub-shell must
must have
have the
the same
same (parallel)
Single
(parallel) spin,
spin,
i.e.
and
not
.
i.e.
and not
.
When aa charged
charged particle
particle spins
spins on
on its
its axis,
axis, aa magnetic
magnetic field
is produced
produced (Figure
When
(Figure 2.61).
2.61). Thus,
Thus, electrons
electrons
26 field is
have
magnetic
properties.
Protons
show
similar
behaviour
to
electrons
and
also
produce
have magnetic properties. Protons show similar behaviour to electrons and also produce aa magnetic
magnetic
field. This
This property
property isis exploited
exploited in
in the
the technique
technique of
of nuclear
nuclear magnetic
magnetic resonance
field.
resonance (NMR).
(NMR).
ELECTRONIC
STRUCTURE 26
■ Electron
Electron configurations
configurations
of atoms
atoms
■
of
S
S
b anticlockwise
b anticlockwise
■ Figure 2.61
■ Figure 2.61
Electron spin
Electron spin
The detailed
detailed electron
electron configuration
configuration of
of the
the hydrogen
hydrogen atom
atom (atomic
(atomic number
The
number 1)
1) is:
is:
1s
1s
alt
can be
be written
written as
as 1s
1s11.. The
The large
large number
number represents
represents the
the shell
shell number
ItIt can
number (principal
(principal quantum
quantum
number),
the
letter
represents
the sub-shell
sub-shell and
and the
the superscript
superscript number
number), the letter represents the
number represents
represents the
the number
number
of
electrons
in
the
sub-shell.
of electrons in the sub-shell.
The detailed electron configuration of the helium atom (atomic number 2) is:
The
detailed electron configuration of the helium atom (atomic number 2) is:
1s
1s
It can be written as 1s2. The two electrons must form a spin pair.
It can be written as 1s2. The two electrons must form a spin pair.
The lithium atom (atomic number 3) has three electrons. Two electrons enter the 1s orbital
The lithium atom (atomic number 3) has three electrons. Two electrons enter the 1s orbital
(as a spin pair). The 1s orbital is now full; so the third electron enters the 2s orbital (the next
(as a spin pair). The 1s orbital is now full; so the third electron enters the 2s orbital (the next
orbital with the lowest energy). This is in accordance with Hund’s rule. The detailed electron
orbital with the lowest energy). This is in accordance with Hund’s rule. The detailed electron
configuration is:
configuration is:
1s
1s
1
2s
2s
It can be written as 1s22s1.
It can be written as 1s22s1.
The beryllium atom (atomic number 4) has the detailed electron configuration shown below.
The beryllium atom (atomic number 4) has the detailed electron configuration shown below.
1s
1s
2s
2s
It can be written as 1s22s2.
It can be written as 1s22s2.
The boron atom (atomic number 5) has five electrons. The first four electrons occupy the
The boron atom (atomic number 5) has five electrons. The first four electrons occupy the
1s and 2s orbitals. The fifth electron occupies the 2p orbital. The correct detailed electron
1s and 2s orbitals. The fifth electron occupies the 2p orbital. The correct detailed electron
configuration is:
configuration is:
1s
1s
2s
2s
2p
2p
It can be written as 1s222s222p11.
It can
becarbon
writtenatom
as 1s(atomic
2s 2p . number 6) has six electrons and the correct detailed electron
The
The
carbon
atom
(atomic
number 6) has six electrons and the correct detailed electron
configuration is:
configuration is:
1s
1s
80
2 Atomic structure
2s
2s
2p
2p
It can be written as 1s222s222p22.
It can
be written
1s 2s 2pdetailed
.
Note
that theasfollowing
electron configurations are not allowed (i.e. are forbidden)
thatatom
the following
detailed
for Note
a carbon
in its ground
state.electron configurations are not allowed (i.e. are forbidden)
for a carbon atom in its ground state.
Detailed electron configuration
829055_02_IB_Chemistry_052-084.indd 79
829055_02_IB_Chemistry_052-084.indd 79
Reason for error (principle violated)
1s
2s
2p
The 2p electrons should occupy different orbitals.
Hund’s rule has been violated.
1s
2s
2p
The single electrons in the same sub-shell should have the same spin.
1s
2s
2p
The 2s orbital can accept one more electron, so it should contain two
electrons.
The Aufbau principle has been violated.
18/05/15 9:26 am
18/05/15 9:26 am
■ Division of the periodic table into blocks
The long form of the periodic table is divided into four blocks: the s-, p-, d- and f-blocks (Chapter 3).
This division reflects the filling of the outermost orbitals with electrons (Figure 2.62).
The detailed electron configurations of the first 37 elements are shown in Table 2.9. These
are the ground state (lowest energy) configurations of the atoms.
Two elements in the first row of the d-block have unexpected electron configurations
(highlighted in Table 2.9) that do not obey the Aufbau principle. The outer electron
configuration of the chromium atom is 4s13d5 and not 4s23d4 as expected. The outer
configuration of the copper atom is 4s13d10 and not 4s23d9. A simplified explanation for these
observations is that a half-filled or completely filled 3d sub-shell is a particularly stable electron
configuration. The outer electron configurations for copper and chromium atoms can also be
written as 3d54s1 and 3d104s1 so that the 3d sub-shell is placed into the third shell.
Group
Period
1
2
3
4
5
6
s-block
1
H
2
Li
Electronic Structure
7
8
9
10
11
12
13
14
d-block
15
16
17
18
p-block
He
Marginalizer
Be
3
Na
Mg
4
K
Ca
B
28
Sc
Ti
V
Cr
Mn
Fe
Co
Ni
Cu
Zn
C
N
O
Bilal Hameed
F
Ne
Bilal Hameed
Al
Si
P
S
Cl
Ar
Ga
Ge
As
Se
Br
Kr
■ Filling atomic orbitals
The electrons are arranged in atomic orbitals according to certain principles:
27electrons. This, in simplified form, is the
■ Each orbital can hold up to a maximum of two
Pauli exclusion principle.
■ Electrons enter and occupy an empty atomic orbital with the lowest energy. This is known as
the Aufbau principle (see Figures 2.59 and 2.60).
1s
ELECTRONIC STRUCTURE 27
2s
2p
3s
3p
3d
4s
4p
4d
4f
5s
5p
5d
5f
3d
3s
alt
6d
2p
Energy
6p
3p
3s
Energy
6s
3d
3p
2s
2p
2s
hydrogen, 1s1
■ Figure 2.59
The Aufbau principle
carbon, 1s2 2s2 2p2
1s
1s
for filling atomic
orbitals with electrons
There are three 3p
rules for allocating electrons to atomic orbitals: 3p
1
Study tip
orbitals to show the 2
application of the
Aufbau or building-up 3
principle
2
2 fill
6 singly
1
Atomic orbitals of the same
before pairing starts.
sodium,energy
1s 2s 2p
3s
This is the
because
electrons repel
each other.
Once
2p-orbitals
have
been filled, the 3s-orbital and the 3p-orbitals
1s
No atomic
orbital can hold more than two electrons. 1s
are occupied in a similar way so that, for example, the ground state of
The electron diagrams for the elements
hydrogen
2
2p63sto23px13py1 and that of chlorine with 17
silicon (14 electrons) is 1s22s
sodium are shown in Figure 6.
2
2
6
2
2
electrons is 1s 2s 2p 3s 3px 3py23pz1.
3p
3s
3p
3s
Note that once the 3p-orbitals have been filled, the next orbital to be
occupied
is the 4s (not the 3d). Therefore, the ground state of potassium
18/05/15 9:26 am
He
H
(19 electrons) is 1s22s22p63s23p64s1.
2s
2p
2s
1s
3p
3p
3s
3s
3p
3p
3s
3s
3p
3s
3p
3s
3p
3s
The 3d-orbitals are filled in the atoms of scandium (21 electrons) to zinc
(30 electrons). These elements have a number of properties in common
Li
Be
B
C
N
O
F
Ne
and are called d-block elements. For example, scandium would be
written as 1s22s22p63s23p64s23d1 or 1s2, 2s2, 2p6, 3s2, 3p6, 3d1, 4s2.
2p
1s
2s
2p
2s
2p
2s
1s
1s
2p
2s
2p
1s
1s
2s
2p
2s
1s
1s
2p
2s
2p
1s
3p
3s
2s
2p
‘Electrons in box’ representations of the electron ground states of some
Na
elements
are given in Figure 2.17.
1s
Synoptic link
▲ Figure 6 The electron arrangements for the elements hydrogen to sodium – note
how they obey the rule above
Number of
electrons structures
1s 2s
2p
Writing
Element electronic
3s
3p
3d
AHydrogen
shorthand way1of writing electronic structures is as follows, for
example, for sodium which has 11 electrons:
Helium
tom has a large
electrons, its
often abbreviated
e symbol of a group
to provide its inner
nfiguration. It is then
ary to provide details
most electrons.
e scandium can be
to [Ar] 4s2, 3d1 or
o [Ar] 4s2, 3d3.
2p
1s
3p
3s
2s
tional to write
n configuration as
2
3p63d14s2 (i.e. with
3d energy levels
r all elements with
3d electrons.
Practise working out the shorthand
electronic structure of all the
elements
at6 least
up
to krypton
2
2
2
4
sulfur,(atomic
1s 2s number
2p 3s 3p
36).
Energy
Energy
3s
1 Atomic
orbitals of lower energy are filled first – so the3slower main
■ Figure 2.60 Electrons shell is filled2p
first and, within this shell, sub-shells of lower energy
2p
are filled
first.
in energy levels or
2s
2s
corresponds to
on the ion that is
hen the atom loses
ctrons. So sodium is
n829055_02_IB_Chemistry_052-084.indd
becomes O2− and 78
n NO3−.
1
Atomic structure
3d
3d
2
1s2
2s2 2p6
3s1
Lithium
3
2
8
1
Carbon
6
Note
how this matches
the simpler 2,8,1 you used at GCSE.
Neon with 20
10electrons would be:
Calcium,
Sodium
1s2
2s211
2p6
3s2 3p6
4s2
which matches 2,8,8,2
Sulfurhow the16
Notice
4s orbital is filled before the 3d orbital because it is of
lower
energy. 18
Argon
After
calcium, electrons
begin to fill the 3d orbitals, so vanadium with
Potassium
19
23 electrons is: 1s2 2s2 2p6 3s2 3p6 3d3 4s2
Scandium 21
Iron
26
Sometimes
it simplifies
things to use the previous noble gas symbol.
Zinc
30
So the electron arrangement of calcium, Ca, could be written [Ar] 4s2
Bromine
35 [1s2 2s2 2p6 3s2 3p6] 4s2 because 1s2 2s2 2p6 3s2 3p6
as
a shorthand for
Krypton with 36 electrons is: 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p6
isKrypton
the electron arrangement
of argon.
36
Strontium
You
can use the38
same notation for ions. So a sodium ion, Na+, would
have the electron arrangement 1s2 2s2 2p6, one less than a sodium
atom, 1s2 2s2 2p6 3s1.
You will learn how electron
arrangements
4s affect
4p the properties
5s
of the transition metals in
Topic 23.1, The general properties
of transition metals.
Summary questions
1 a Give the full electron
arrangement for
phosphorus.
b Give the electron
arrangement for phosphorus
using an inert gas symbol as
a shorthand.
2 a Give the full electron
arrangements of:
i Ca2+ and ii F−
b Give their electron
arrangements using an
inert gas symbol as a
shorthand.
Figure 2.17 The electronic ground states of some elements.
25
Test yourself
8 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 =
Bilal Hameed
Marginalizer
c) Number of different types of orbital =
Bilal Hameed
29 =
Electronic Structure
d) Maximum number of electrons in the shell
9 Give the electron orbital configuration for the ground state of the
following atoms or ions:
28
ELECTRONIC STRUCTURE 28
Section 1.3 Exercises
Skill Check 12
1 State the maximum number of electrons that can be found in:
Section 1.3 Exercises
a any px orbital
thethe
first
energy level
1 b
State
maximum
number of electrons that can be found in:
ca any
d
subshell
any px orbital
d
energy
level
b the
the second
first energy
level
ec any
any sd orbital
subshell
alt
d thean
second
energy
levelfor each of the following elements.
Draw
orbital
diagram
e
any
s
orbital
Skill
Check 13
a Magnesium
Silicon
2 b
Draw
an orbital diagram for each of the following elements.
Magnesium
ca Chlorine
b Siliconhow three electrons would be distributed in a p subshell that was
3 Describe
c Chlorineempty.
previously
2
1
43
Describe
three
electrons
would
p subshell
that was
Name
thehow
block
of the
periodic
tablebeindistributed
which eachinofa the
following
previously
Skill
Check
14
elements
isempty.
found.
and
complete
for theeach
quantum
shell
with principal
aName
Sodium
4 Copy
the
block ofthe
thefollowing
periodicinformation
table in which
of the
following
elements
is
found.
quantum
number 3.
b Carbon
a
Sodium
c Iron
a)b Total
number of sub-shells = __________
Carbon
d Chlorine
c Uranium
Iron
b)e Total
number of orbitals = __________
d
Chlorine
f Silver
Uranium
c)e Number
of different types of orbital = __________
5 For each of the following elements, state the highest energy subshell that
f Silver
is being filled.
d) Maximum number of electrons in the shell = __________
End of c
5 a
ForOxygen
each of the following elements, state the highest energy subshell that
is being filled.
b Chromium
a Oxygen
8 The
accurate relative isotopic masses of five isotopes
15 a List three factors which influence the
c Strontium
Skill
Check 15
are
shown below:
ionisation energy of an element.
b Chromium
End of chapter
1d Aluminium
2
12 structure for
14 the following
16
Draw
the
‘electrons
in
boxes’
elements:
b
The first ionisation energies of the ele
H
H
(D)
C
N
O
1c Strontium
1
6
7
8
e Arsenic
Group I are shown below.
d Aluminium
1.0078
2.0141
12.0000
14.0031
15.9949
8 a)
The
accurate
relative isotopic
masses
of five isotopes
15 a List three factors which influence the size o
f boron
Rhodium
e Calculate
Arsenicbelow:
aare
the accurate relative molecular masses for:
shown
ionisation
anionisation
element.
Elementenergy of
First
1
12
16
6 b)
State
configuration
i
N2theii21electron
CO
iv C214
Hof
v Cof
Rhodium
/kJ mol
HDCN
(D) iii
b The first ionisationenergy
energies
of –1the elements
4 each
2D8the
2O following elements.
1fHfluorine
6C
7N
Group
Nitrogen
ba The
relative molecular
mass of
a certain15.9949
gas in a
Li I are shown below.
520
12.0000
14.0031
6 1.0078
State the 2.0141
electron configuration
of each of the following elements.
high-resolution
mass
spectrometer
was
28.0171.
c)
phosphorus
b
Chlorine
Na
500
aaCalculate
the accurate relative molecular masses for:
Element
First ionisation
Nitrogen
What
gas
is probably
under
observation?
–1
i
N
ii
DCN
iii
CO
iv
C
H
v
C
D
energy /kJ mol420
K
2
2 4
2 2
cb Calcium
Chlorine
d)
potassium.
The relative
molecular
mass of a certain gas in a
9 bDraw
the shape
of
Li Rb
520 400
d
c Nickel
Calcium
mass spectrometer
was 28.0171.
Na Cs
500 380
ae high-resolution
an
s
orbital.
b
a
p
orbital.
Selenium
Skill
Check
gas is 16
probably under observation?
dWhat
Nickel
K
420
10 List
the orbitals below in order of increasing energy.
fe Tin
Selenium
9 Draw
the shape of
Rb Explain the change
400 in the first ionisat
1s 3p 3d 4s 3s 2s 2p 4p
proton
number
increases.
f an
Tin
7 aState
electron configuration
of each of the following negative
ions.
s the
orbital.
b a p orbital.
Cs
380
11 Write down the chemical2symbols of the elements
that
3
c Explain, with examples, how ionisati
aState
Fthethe
b inTe
c ofP
d Br
List
orbitals
below
order of increasing
energy.
710 have
electron
configuration
ofgiven
each
the following negative
ions.
the
electronic
configurations
below.
Explain
the change
in for
the the
firstarrangemen
ionisation en
provide
evidence
2
3
1s
3p
3d
4s
3s
2s
2p
4p
2
2
3
Skill
Check
17
a 1sF , the
Te
c Pof the following
d positive
Brproton
number increases.
8 aState
configuration
of each
2s 2pelectron b
in ions.
shells.
11 Write2 down
of the elements that
2 2 6the 1chemical symbols
2
Explain,
with examples, how ionisation ene
aState
Mg
b Al3
c ofRb
dc16Zn
1s
, the
2s
2pelectron
, 3s
8 b
configuration
of each
the following positive
Theions.
graph
in Figure
2.26 shows the first
have
the
electronic
configurations
given
below.
provide
evidence
for the arrangement of el
2 22 6
2 5
3
2
caa 1s
1sMg
2 , 2s2 2p3 , 3s 3p b Al
ionisation
energies
of the elemen
c Rb
d Zninsecond
, 2s 2p
shells.
to calcium.
2 the
2 electronic
6
1
12 Write
configurations
of
the
atoms
and
b 1s , 2s 2p , 3s
16 The graph in Figure 2.26 shows the first and
2
2
a Why is there a large decrease in the f
6
ions
c 1s2given
, 2s22pbelow,
, 3s23p5in the form 1s , 2s …
second ionisation energies of the elements nit
energy after neon and after argon?
a Li
b O
c Cl
to calcium.
12 Write the electronic configurations
of the –atoms and
+
b
is the
first ionisation
m
d
K given below, inetheLiform 1s2, 2s2…
f Cl
a WhyWhy
is there
a large
decrease inenergy
the firstofion
ions
greater
than
that
for
aluminium?
2–
+
energy after neon and after argon?
g
O
a Li
bh OK
c Cl
c Why
is first
the second
ionisation
+
–
b
Why
is the
ionisation
energy ofenergy
magne
d K following table
e shows
Li
f Cl
13 The
the ionisation
energies
element
greater
than the correspondi
greater
than
that
for
aluminium?
2–
+
–1
g OkJ mol ) of five
h elements
K
Marginalizer
Bilal Hameed
(in
lettered A, B, C, D and E.
energy?
c Whyionisation
is the second
ionisation energy of each
13
The
following
table
shows
the
ionisation
energies
d Whygreater
do the
maxima
for the two grap
Element –11st
2nd 30
3rd
4th
element
than
the corresponding
firs
Electronic Structure
Bilal
Hameed
(in kJ mol ionisation
) of five elements
lettered
A, B, C, D
and E.
ionisation
ionisation
ionisation
ionisation
energy?
different
proton numbers?
Element
A
energy
1st
500
ionisation
energy
2nd
4600
ionisation
energy
3rd
6900
ionisation
energy
4th
9500
ionisation
d Why do the maxima for the two graphs occ
different proton numbers?
of gaseous isolated atoms in the ground
22p 6
Ne
[He]2s
22
ms in10
bold are discussed
on page
80).
1
11
Na
[Ne]3s
of gaseous isolated atoms in the ground 23
[Ne]3s2
Mg
ms in 12
bold are discussed
on page 80).
figurations of excited species
36
Kr
37 Ti
V
2
2
[Ar]4s
Rb 3d [Kr]5s1
[Ar]4s23d3
34
Se
[Ar]4s23d104p4
35
Br
[Ar]4s23d104p5
[Ar]4s13d5
36
Kr
[Ar]4s23d104p6
37
Rb
[Kr]5s1
24
Cr
29
[Ar]4s 3d 4p
■ electrons
Table 2.9 Detailed
electron
configurations
of gaseous
isolated
in the groundinto higher
re
absorb
thermal
or electrical
energy,
theyatoms
are promoted
figurations
excited
state (the electronof
configurations
ofspecies
the atoms in bold are discussed on page 80).
ELECTRONIC
STRUCTURE
The atoms and electrons
are in an excited
state. 29
re electrons absorb thermal or electrical energy, they are promoted into higher
Electron
configurations
ofinexcited
species
le
an
excited
sodium
atom
shown
below
Figure
2.63. The
return of the
Electron
configurations
of excited
species
Additional
heofatoms
and electrons
are inisan
excited
state.
When
one
or
more
electrons
absorb
thermal
or
electrical
energy,
they arethey
promoted
o thePerspectives
ground state will
give
rise
to
emission
of
electromagnetic
radiation
When one or more electrons absorb thermal or electrical energy,
are promoted into higher
into
higher
energy
orbitals.
The
atoms
and
electrons
are
in
an
excited
state.
lea of
an
excited
sodium
atom
is
shown
below
in
Figure
2.63.
The
return
of
the
orbitals.
The atoms
and electrons
specific line in the energy
emission
spectrum
of sodium
atoms.are in an excited state.
o the ground state willAgive
riseexample
to emission
of electromagnetic
radiation
specific
of excited
an excited
is shown
in Figure
A specific
of
sodium
atomatom
is shown
below.below
The return
of the2.63. The return of the
1s
2s
2p example
3san
3p sodium
a specific line in the emission
spectrum
sodium
excited electron
to theof
ground
stateatoms.
will give rise to emission of electromagnetic radiation
d state:
alt
12.1
excited electron to the ground state will give rise to emission of energy.
In general, the shielding effect is most
corresponding
to3s
a specific line
2p
3p in the emission spectrum of sodium atoms.
effective if the electrons are close t
Consequently, electrons in the first shell (energy level), where there is high e
1s
2spromoted
2p
3s
3p
■ Figure 2.63 Orbital
d state:
electron
have
a stronger shielding effect than electrons in the second shell, which in
sodium
atom,
excited
state:
notation for sodium
1s
2s
2p
3s
3pshielding effect than electrons in the third shell. Electrons in the same shell
atoms in ground and
electron
promoted
d state:
small
shielding effect on each
other.
electron
promoted
excited states
1s
2s
2p
3s
3p
shows
ionization energies for the chemical elements of p
1s
2s Figure 12.14
2p
3s the first3p
d state:
sodium atom, ground state:
The general increase in ionization energy across each period is due to the increas
This occurs because across the period each chemical element has one additional
c configuration of ions
increases the nuclear charge by +1.
Electronic configuration of ions
1s
2s
1
First ionization energy/kJ mol–1
Pauli exclusion principle
the Aufbau
principle also apply
when extra
■ and
Electronic
configuration
of ions
Hund’s
rule,
the Pauli
2500
cedconfiguration
of
ions
Heexclusion principle and the Aufbau principle also apply when
to form negative ions
(anions).
The
fluoride
ion
(Figure
2.64)
is
formed
when
Hund’s
rule, the
Pauli
exclusion
principle
the Aufbau
principle
apply when extra
extra
electrons
are added
to form negative
ionsand
(anions).
The fluoride
ion is also
formed
2
2
5
Pauli
exclusion
principle
and
the
Aufbau
principle
also
apply
when
extra
Ne
1s 2s 2p ) gains an additional
electron.
electrons
are
added
to
form
negative
ions
(anions).
The
fluoride
ion
(Figure
2.64) shielding
is formed when
2
2
5
2p ) gains an additional electron.
when a fluorine atom 2(1s2 2s2
2000
ed
to form configuration
negative ions
(anions).
fluoride
ionan(Figure
2.64)
is formed when
a fluorine
atomThe
(1s 2s
2p5) gains
additional
electron.
e electron
force
22s22p5) gains an additional
the 1s
electron2sconfiguration
1s
electron.
2p
cations),
electrons are To deduce
Li
Be
B
C
Ar
1s 2 2s
2p
of positive
ions (cations), electrons are
2 2p6
configuration
1500
or
simply
1s
2s
F – ion:
seelectron
order (that
is, the last
or simply 1s2 2s2 2p6
F – ion:
removed in reverse order (that is, the last
2p
ations),
are electron is removed
H 1sfirst). 2s
ed
first).electrons
(An exception
(An exception
2
6
■
2.64 Orbital
notation and■ detailed
electron
ion:
or
simply
1s
2s2 2pnotation
F – Figure
e order
(thattransition
is, the last
Figure
2.64
Orbital
and detailed electron
rs
with the
1000
todeduce
this ‘rule’
the transition
To
the occurs
electronwith
configuration
of positive ions
(cations),
electrons
are
removed
–configuration for a fluoride ion, F – electrostatic
configuration
for
a
fluoride
ion,
F
ed
see Chapter
13.)last electron is removed first). (An exception to this ‘rule’
pterfirst).
13.) (An exception
inmetals
reverse–order
(that is, the
attraction
+(g)
■
Figure
2.64
Orbital
notation
detailed electron
+
For
example,
the
O
ion to
is formed
by and
rs
with
the
transition
he O (g) ion is formedoccurs
by with500
towards
the transition metals,
be discussed
ahead.) For example, the O+(g) ion
–
1s
2s
2p
the
removal
of
one
electron
from
an
oxygen
configuration
for
a
fluoride
ion,
F
positive
Li
1s
2s
2p
2
2
4
pter
13.)
Na
e electron from an oxygen
).
is formed 2by 2the4removal of one electron from an oxygen
atom (1s 2s 2p
K 3
or simply
1s2 2s2 2p3
nucleus
atom (1s 2s+ 2p ) (Figure 2.65). This ionization O+ ion:
2
2
+
or simply 1s 2s 2p
O ion:
he O (g)
ionThis
is formed
by
Figure
2.65).
ionization
process can 0be made to occur inside a mass
1s
2s
2p
electron
from
an oxygen
■ Figure
and detailed electron
1 electron
5removed
10
152.65 Orbital notation
20
de
to occur
inside
a mass
spectrometer.
The
is the
Figure 12.15 A diagram illustrating
+
2 an
2 O+(g)
3 ion
■
Figure
2.65
Orbital
notation
and
detailed
electron
configuration
for
ion:
or
simply
1s
2s
2p
O
igure
2.65).
This ionization
electron
removed
is thelast electron from the 2p sub-shell. Atomic number, Z
between shielding and nuclear charg
configuration for an O+(g) ion444 12 Atomic structure
de
inside a mass
theto2poccur
sub-shell.
period 2
Figure 12.14 First ionization energies for periods 1, 2 and 3
■ Figure 2.65 Orbital notation and detailed electron
electron removed is the
configuration for an O+(g) ion
A similar
explanation
also
fo
the 2p sub-shell.
The increase in
nuclear
charge
the force
of attraction
onaccounts
all the ele
boron atom, B
nitrogen
atom,
N increases
1s
2s
2p
+
boron ion, B
1s
2s
2p
_Chemistry_052-084.indd 81
beryllium atom, Be
1s
2s
2p
+
beryllium ion, Be
1s
2s
2p
Figure 12.16 Orbital
notations for boron and
beryllium atoms and
their unipositive ions
Bilal Hameed
Bilal Hameed
magnesium
andelectron
aluminium.
decre
held closer and1shence
Each
additional
acrossThe
a period
2smore strongly.
2p
2
2
6
2
1
to aluminium
(1sis2sminimal
2p 3s 3p(Figure
) arises12l
shell (energy level) and hence the increase
in shielding
effective
at
shielding
the
electron
in th
Although the general trend
is
for
the
ionization
energy
to
increase
across
nitrogen ion, N+
The second
in first 12.14
ioni
are two distinct
in ionization
periodsdecrease
2 and 3 (Figure
1sdips2s
2p energy 2across
22p 12p 12p 1) and oxygen (1s22
(1sof2s
only be explained using an orbital model
electronic
structure.
x
y
z
nitrogen
arein18/05/15
inthe
three
separ
The first decrease in each period is the result
of a atom
change
sub-shell
9:26 am
which
states
that
every
orbital
in
which the electron
is
lost
and
a
change
in
electron
shielding.
These
have
aagrs
oxygen atom, O
1s charge
2s
orbitalradius.
is doubly
occupied.
Howe
increase in nuclear
and2pdecreaseone
in atomic
In period
2, this
firs
18/05/15
9:26 amThe
2p orbital.
two electrons
in the
between the elements beryllium and boron.
When
it is ionized,
the beryllium
repulsion
makes (Figure
it easier12.1
to
2s22p1) loses
a 2p electron
a 2s electron, whereas a boron
atom (1s2electron
oxygen ion, O+
required to remove
an
electron
from
the
lower
energy
2s
orbital
in
beryllium
t
electron
from
a
half-filled
2p
orbital
18/05/15 9:26 am
z
1s
2s
2p
energy 2p orbital in boron. Although the
2s and and
2p sub-levels
in to
thethe
same
nitrogen
oxygen isaredue
ad
difference is relatively large. Recall (Chapter
2) atom
that the
energy
gap between s
oxygen
(Figure
12.18).
becomes smallerFigure
with an
increase
addition, a single
electro
A similarInexplanation
accounts
for
12.18
Orbitalin shell number.
2 electrons
level is more effectively
shielded
by
the
inner
electrons
than
the
2s
phosphorus and sulfur in period 3. Th
notation for nitrogen
is less than that of phosphorus (1s22s22
and oxygen atoms and
an electron from the 3p4 orbitals of su
their unipositive ions
Marginalizer
2s
2pfirst ionization ener
The patterns of
the corresponding
31
Electronicionization
Structure energies f
removed are in a second shell closer to
elements. The outer electrons in period
30
ELECTRONIC STRUCTURE 30
octet rule
According to the octet rule, atoms usually form stable ions by losing or gaining
electrons to attain an octet. For example, the nitrogen atom gains three electrons to
attain the stable electron arrangement of neon, the nearest noble gas. The calcium
atom loses two electrons to attain the electron arrangement of the noble gas, argon.
Lithium and beryllium atoms lose electrons to attain the electronic arrangement of a
helium atom, with two electrons. The lithium atom loses one electron to form the
lithium ion, Li+.
1s
2s
2p
alt
ron arrangements of noble gases
ely
stable
and their
atoms do
The
octet
rule
electrons
to form ions.
er gain
electron
arrangements
of noble gases
nitrogen
82 atom:
2 Atomic structure
noble
gases,
with
the
exception
relatively stable and their atoms do
1s
2s 2p
2p
1s
2s
haveoreight
lose
gainelectrons
electronsintotheir
form ions. 3–
atom): atom:
N ion (or neon nitrogen
■ The octet rule
ls. This
arrangement
is known
oms
of noble
gases, with
the exception
t.
arrangements of noble gases
1s
2s The electron
2p
helium,
have eight electrons in their
octet
3–
ding
to
the
octet
rule,
atoms
are relatively stable and their atoms do
N ion (or neon atom):
er shells. This arrangement is known
■ Figure 2.66 Orbital notation for a nitrogen
m
stable
ions
by
losing
or
not lose or gain electrons to form ions.
n octet.
atom and the nitride ion, N3 −
octet
ectrons
to
attain
an
octet
Atoms
of noble gases, with the exception
According to the octet rule, atoms
4).
example,
the nitrogen
■ Figure 2.66 Orbital notation for aofnitrogen
helium, have eight electrons in their
allyFor
form
stable ions
by losing or
3−
N3
s
three
electrons
to
attain
the
atom
and
the
nitride
ion,
N
outer shells.
3s
3p
4s This arrangement is known
ning electrons to attain an octet
tron arrangement
of neon,
as an octet.
calcium atom: [Ne]
hapter
4). For example,
the nitrogen
(Figure
2.66).
According to the octet rule, atoms
mt noble
gains gas
three
electrons
toThe
attain the
3s 3p
3p 4s
4s
3s
■
om
loses
two
electrons
to
usually
form stable ions by losing or
ble electron arrangement of neon, Ca2+ ion (or argon atom):
calcium [Ne]
atom: [Ne]
ato
gaining electrons to attain an octet
electronnoble
arrangement
of the
nearest
gas (Figure
2.66). The
example, the nitrogen
3s (Chapter
3p 4). For 4s
argon
(Figure
Lithium to
octet
cium
atom
loses2.67).
two electrons
atom
gains
three
electrons to attain the
Ca2+ ion (or argon atom): [Ne]
ium
atoms
lose
electrons
to
in the electron arrangement of the■ Figure 2.67 Orbital notation for a calcium atom and
stable electron arrangement of neon,
electronic
arrangement
of Lithium
a
le
gas, argon
(Figure 2.67).
the calcium
ion, Ca2+of ions
1s
2s
octet
Electron
configurations
the
nearest noble gas (Figure 2.66). The
om,
with two
electrons.
The
beryllium
atoms
lose electrons
to
lithium
atom:or notation
An
ion
is
formed
when
an
atom
loses
gains
electrons.
While
an
will
■
Figure
2.67
Orbital
for
a
calcium
atom loses
and two electrons to
calcium
atom
om
loses
one electron
to form of ahave the same atomic number as an atom of the same element, it ion
in the
electronic
arrangement
will have a
Ca
2+
+
the calcium
ion,
attainconfiguration.
the electron arrangement of the
1s a different
2s electron
mum
ion,
Li
(Figure
2.68).
different number
of electrons
andCa
so will have
atom, with two electrons. TheThe driving force behind formation of+ ions is the gaining of the same stability
noble gas, argon (Figure 2.67). Lithium
lithium ion Li :
as the nearest noble gas, so ions of main group elements have the same
ium atom loses one electron to form
WORKSHEET 1.2
and beryllium atoms lose
electrons to
electron configuration
as the nearest
noble gas.exclusion
Electron configurations
glithium
of the
Aufbau
principle,
Hund’s
rule and
the Pauli
■
ion,
Li+ (Figure
2.68).
■ Figure 2.68 Orbital
attain the electronic arrangement of a
e to write electron configurations
for
atoms
and
ions
up
to
Z
=
36
the
TABLE 1.3.4 ELECTRON CONFIGURATIONS OF IONS OF SOME MAIN GROUP ELEMENTS
for a lithium
COMPARED TO THATnotation
OF THE NEAREST
NOBLE GAS helium atom, with two electrons. The
ron
configuration
of any atom
or ion
(Z <Hund’s
36) can rule
be
predicted
by applying
the Aufbau
plying
of the Aufbau
principle,
and the
PauliSymbol
exclusion
atom
lithium
lithium
atom configuration
loses one electron to form
Element
Atomicand the
Group
Electron
number
(Z)
Hund’s rule
and theelectron
Pauli exclusion
principle. Thefor
two
exceptions
are
copper
and
+
nciple
to write
configurations
atoms
and
ions
up
to
Z
=
36
ion, Li
the lithium
ion, Li+ (Figure 2.68).
Neon
10
0
Ne
1s2 2s2 2p6
.
2
2
e electron configuration of any atomMagnesium
or ion (Z < 36) 12
can be predicted
by Mg
applying
the
2
1s
2s2Aufbau
2p6
2
2
6
3
Nitrogen principle. The
7
5
1s 2sand
2p
N
nciple, Hund’s rule and the Pauli exclusion
two exceptions
are
copper
Applying
of
the Aufbau principle, Hu
2
2
3+
2+
3−
Argon for the following
18
0, Cu, P , Cl
Arand Ga.1s 2s 2p6 3s2 3p6
ull
electron
configurations
and
orbital
diagrams
Ti
,
Cr
omium.
principle
to6 write
electron configurati
2
2
2
6
1
e all atoms are gaseous and in the groundPotassium
state.
19
1
K
16
6
S2
1s 2s 2p 3s 3p
The electron
of any atom or ion (
1s2 2s2 2p6configuration
3s2 3p6
3+
2+
3−
Write
full
electron
configurations
and
orbital
diagrams
for
the
following
Ti
,
Cr
,
Cu,
P
,
Cl
and
Ga.
principle,
Hund’s
rule
and
the
Pauli exclusion p
h of the following, decide whether the full or condensed electronic configuration
shown represents
Note
in the table
above that Ne, Mg2 and N3 ions have the same electron
Assume
all
atoms
are
gaseous
and
in
the
ground
state.
tom, a positive ion (cation) or a negativeconfiguration,
ion (anion) and
of the
shown,
chromium.These
Ar, Kelement
and S2 have
the same electron configuration.
Sulfur
groups of ions are referred to as isoelectronic—having the same number
e ground state or an excited state.
of electrons.
For
the full or condensed electronic configuration shown represents
1s2each
2p1 of the following, decide whether
configurations of metaland
ions orbital diagra
11 Write full electronElectron
configurations
2
a1san atom, a positive ion (cation) or a negative
ionmetal
(anion)
Transition
ionsof the element shown,
Assume all atoms are gaseous and in the ground st
b[He]
in the
2s2 ground
2p6 3s2 state
3p4 or an excited state.
The majority of the 3d transition metals have a full 4s subshell, although as
2
1
1s 3 2p
1si2 2sLi1 2p
discussed above, chromium and copper have only one electron in a 4s subshell.
25
12filled
For 3d
each
of the following, decide whether the full o
2
1
A full 4s subshell has a slightly higher energy than a partially
subshell
ii
H
1s
[He] 2s 2p 3s
a
an
atom,
a positive ion (cation) or a negative ion
2
6
2
4
(see
chapter
3),
so
when
a
3d
transition
metal
loses
electrons,
the
electrons
are
iii S [He] 2s 2p 3s 3p
initially
lost electrons
from the 4s in
subshell.
If more than 2 electrons are lost,
then
the
b
in
the
ground
state or an excited state.
2
1
3
t about
the
use
of
quantum
numbers
to
describe
atoms.
iv N 1s 2s 2p
subsequent electrons are taken from the 3d subshell. Similarly, electrons are
2 2p1
i
Li
1s
2
5
1
v F [He] 2s 2p 3s
taken from the 5s subshell before the 4d subshell.
ii H 1s2
The electron configurations of some transition metal ions are shown below.
6 3s2Hameed
Find out about the use of quantumMarginalizer
numberswill
to describe
electrons
in atoms.
Bilal
iii S [He] 2s2 2p
3p4
be discussed
in more detail
in chapter 3.
Q2 AThese
chemical
element
with the
symbol X has the
2
1
3
iv N 1s 2s 2p
Electronic Structure
32
Bilal Hameed
electron
2,8,6. Which
chemical
TABLE 1.3.5arrangement
ELECTRON CONFIGURATIONS
OF SOME
TRANSITION
v METAL
F [He]IONS
2s2 2p5 3s1
questions –
species
element
likely
form?
Element is this chemical
Atomic
Ion symbol most
Number
of toElectron
configuration
Q2
A chemical
element
with the
symbol
X has
13
outthe
about the use of quantum numbers to des
number
(Z)
electrons
in Find
3+
A the ion X
the ion
TURE
on questions –
ng
Electron configurations of elements with more than 18 electrons
In figure 1.2.7, the energies of all the subshells up to 5f are shown. Notice that
the energy of the 4s subshell is less than
31that of the 3d subshell. This means
that the 4s subshell fills before the 3d subshell. As a consequence of this filling
order, the majority of the elements of the first transition series (V to Zn) have a
ELECTRONIC STRUCTURE 31
full 4s subshell.
Periodic
Table
When the
number
of electrons becomes great ( 20), even the shorter (subscript)
the modern
table,configuration
elements are arranged
in order
of atomic
number.
The
form ofInwriting
theperiodic
electron
becomes
tedious.
For
note-taking
horizontal
rows inshorthand
the table are form
called periods
each
these ends
a noble
purposes,
a further
can beand
used
inofwhich
thewith
electron
gas.
configuration
of the previous noble gas is represented by the symbol for that
element in square brackets (e.g. [Ar]) and is followed by the rest of the electron
The vertical
in the table
are calledUsing
groups this
and these
can be divided
into four
configuration
forcolumns
the element
required.
shorthand
the electron
2f-block on the basis of the electron
blocks:
the
s-block,
p-block,
d-block
and
configuration of calcium would be [Ar]4s .
alt
structures of the elements.
Note that this shorthand form should not be used when answering examination
questions.
Our modern arrangement of elements in the periodic table reflects the pattern in
1
electronic
structures
of the
atoms,
while the
more sophisticated
modelelectron
of electron
The periodic
table
can be
used
to great
advantage
in writing
structure
in
terms
of
orbitals
allows
chemists
to
explain
the
properties
of
configurations. The periodic table can be divided up into blockselements
that reflect
more effectively.
fourfilled.
blocks in the periodic table are shown in different colours
the subshell
that is The
being
below.
s-block
p-block
period
1 1s
2
2s
3
3s
d-block
3p
4
4s
3d
4p
5
5s
4d
5p
6
6s
5d
6p
7
7s
6d
Lanthanides
Actinides
1s
2p
f-block
4f
5f
TheThe
s-block
consists
hydrogen,
and groups
1 (alkali
Figure 1.3.6
periodic table
can beofdivided
up intohelium
blocks reflecting
the subshell
thatmetals)
is beingand
filled.2. All the
1
s-block elements have a half-filled s orbital (s ) or a completely filled s orbital (s2) in
the outermost shell.
This view of the periodic table can be most useful in determining the electron
The p-block
consists
of groups
to 18. The s-an
andelectron
p-blocks are
collectively called
configuration
of any
element.
To13
determine
configuration:
the main group elements. Each p-block element has an outer electron configuration
that varies from s2p1 (group 13), s2p2 (group 14) through to s2p6 (the noble gases in
Step 2: group
Determine
the name of the block that the element is in (s, p, d, f)
18).
Step 1: Locate the element on the periodic table.
Step 3: Count how many groups (vertical columns) from the left of that block
The d-block consists of three series of metals. Each series of d-block metals
the element is in.
contains ten metals with outer electron configurations ranging from d1s2 to d10s2.
Step 4: Determine which period the element is in by counting down from the
firstare
period,
which
consists
H and
He.
There
two series
of metals
at the of
bottom
of the
periodic table known as
f-block
metalsand
because
they contain
that
are being
filled. than
The two
rows
Step 5: the
Fill
all shells
subshells
thatf orbitals
have a
lower
energy
this
ofsubshell
the f-block(see
series,
known1.2.7).
as the lanthanoids and actinoids, each contain 14
figure
elements.
Bilal Hameed
Bilal Hameed
Marginalizer
33
Electronic Structure
2
2
3
4
3
e any s orbital
c any d subshell
Draw
orbital
diagram
d
thean
second
energy
levelfor each of the following elements.
a any
Magnesium
e
s orbital
32
b Silicon
Draw an orbital diagram for each of the following elements.
c Magnesium
Chlorine
a
ELECTRONIC
Describe
three electrons32
would be distributed in a p subshell that was
b
Siliconhow STRUCTURE
previously
empty.
c Chlorine
Skill
Check 18
Name thehow
block
of the
periodic
tablebeindistributed
which eachinofathe
following
Describe
three
electrons
would
p subshell
that was
elements
is
found.
previously empty.
a Sodium
Name the block of the periodic table in which each of the following
b Carbonis found.
elements
c Sodium
Iron
a
d Carbon
Chlorine
b
Uranium
ce Iron
alt
4
5
5
f Chlorine
Silver
d
eForUranium
each of the following elements, state the highest energy subshell that
is being
filled.
f
Silver
Skill Check
19
a
Oxygen
For each of the following elements, state the highest energy subshell that
b being
Chromium
is
filled.
c Oxygen
Strontium
a
d Chromium
Aluminium
b
Arsenic
ce Strontium
f Aluminium
Rhodium
d
1
6
eState
Arsenic
the electron configuration of each of the following elements.
fa Rhodium
Nitrogen
6
b Chlorine
State
the electron configuration of each of the following elements.
c
Calcium
a
Nitrogen
Skill Check 20
b
d Chlorine
Nickel
Which
of the following elements have atoms that contain only one unpaired p-orbital
ce Calcium
Selenium
electron?
d
f Nickel
Tin
(a)
phosphorus,
eState
Selenium 15P
the electron configuration of each of the following negative ions.
(b)
35Br
fa bromine,
Tin
F
b Te2
c P3
d Br
(c) aluminium, 13Al
State the electron configuration of each of the following negative ions.
State the electron configuration
of each of3 the following positive ions.
a F 2
b Te23
c P
d Br 2
a Mg
b Al
c Rb
d Zn
State the electron configuration of each of the following positive ions.
a Mg2
b Al3
c Rb
d Zn2
7
7
8
8
Marginalizer
Electronic Structure
Bilal Hameed
34
Bilal Hameed
33
ELECTRONIC STRUCTURE 33
Electron configurations of the d-block elements
As the shells of electrons around the nuclei of atoms get further from the nucleus,
they become closer in energy. So the difference in energy between the second and
third shells is less than that between the first and second. When the fourth shell is
reached, there is, in fact, an overlap between the orbitals of highest energy in the third
shell (the 3d orbitals) and thatTutorial
of lowest energy in the fourth shell (the 4s orbital).
alt
4p
Orbitals in
the 4th shell
16.1
Electron configurations
As the shells of electrons around the nuclei of atoms get further from the
nucleus, they become closer in energy. So the difference in energy between
3d
4s
the second and third shells is less than that between the first and second.
When the fourth shell is reached, there is, in fact, an overlap between the
orbitals of highest energy in the third shell (the 3d orbitals) and that of lowest
3p
energy in the fourth shell (the 4s orbital) (Figure 16.2).
Orbitals in
The 3d sub-shell is on average closer to the nucleus than the 4s sub-shell,
the 3rd shell
but at a higher energy level. So, once the 3s and 3p sub-shells are filled, the
next electrons go into the 4s sub-shell because it occupies a lower energy level
3s
than the 3d sub-shell.
Figure 16.2"
This means that potassium and calcium have the electron structure [Ar]4s1
Relative
energy levelsisofon
orbitals
in the
2 respectively
220 The
3d sub-shell
average
closer to
nucleus
than the(Table
4s sub-shell,
andthe
[Ar]4s
16.1). but at a
1
third and fourth shells.
higher energy level. So, once the 3s and 3p sub-shells are filled, the next electrons go
into the 4s sub-shell because it occupies a lower energy level than the 3d sub-shell.
Electron configuration
EXAMPLE
6
This means that potassium and calcium have the electron structure [Ar]4s1 and
Write the
electron
configuration of an oxygen atom.
2 respectively.
[Ar]4s
Answer
2p
configuration
s2 2p 4 .
The last element before the first member of the d-block is calcium, whose atom has
The atomic number of oxygen is 8.
the detailed electron configuration 1s2 2s2 2p6 3s2 3p6 4s2. However, with the next
The 1s and 2s sub-shells are full and this leaves 4 electrons to place
element, scandium, the additional electron is placed in a 3d sub-level which was
in the 2p sub-shell. There is one orbital with 2 electrons (spinning in
unoccupied
in the
opposite
directions(empty)
shown as
↑↓)calcium
and theatom.
other two 2p orbitals in this
sub-shell have one electron (both shown as ↑).
The 3d sub-level has five orbitals into which successive electrons are placed
Oxygen atoms have only two unpaired electrons as one of the 2p orbitals
according to the Aufbau or building up principle. In particular:
has a pair of electrons which spin in opposite directions to minimise their
repulsion for each other.
Electrons are, if possible, placed in 3d orbitals without being paired up, unless there
are no more empty orbitals.
4p
nfiguration of a vanadium atom.
↑↓
f vanadium is 23. For 23
the 1s2 2s2 2p6 3s2 3p6 subctrons in total). This leaves five
he 4s and 3d sub-shells. The 4s
ectrons and then the remaining
aced in the 3d (all spinning in
different orbitals).
Vanadium:
↑↓
↑
4s
↑↓
↑↓
↑↓
3p
↑↓
↑↓
↑↓
2p
↑
3d
3s
↑↓
2s
↑↓
1s
Figure 1.18 The electron configuration of a vanadium atom
6 3s 2 3p 6 3d 3 4s 2 .
1s2 2s2up
2pin
If electrons areispaired
the same 3d orbital, then a spin pair results. The electron
configurations of the first-row d-block metals are given on the next page.
4p
nfiguration of a chromium atom.
f chromium is 24.
↑
↑
Bilal Hameed
Bilal
6 areHameed
full as
e 1s2 2s2 2p6 3s2 3p
total). This leaves six electrons
3d sub-shells. By moving one
↑↓
↑↓
↑
4s
↑↓
↑↓
↑↓
3p
↑↓
↑↓
↑↓
2p
3s
2s
↑
↑
↑
↑
3d
Marginalizer
35
Electronic Structure
than the 3d sub-shell.
34
Table 15.1 Electron configurations from potassium to zinc in Period 4 of the Periodic
Write the electron configuration of an oxygen atom.
Table. ([Ar] represents the electronic configuration of argon.) Note the way that the
ELECTRONIC
STRUCTURE
for chromium34
and copper atoms do not fit the general pattern.
Answer electron configurations
EXAMPLE 6
Element of oxygen
Symbol
The atomic number
is 8.
Electronic structure
s,p,d,f notation
↑
2p
Electrons-in-boxes
The 1s and 2s sub-shells are full and this leaves 4 electrons
to place
notation
in the 2p sub-shell.
There
is
one
orbital
with
2
electrons
(spinning
in
1
Potassium
K
[Ar]4s
[Ar]
opposite directions shown as ↑↓) and the other
two
2p
orbitals
in
this
[Ar]4s2
Calcium
Ca
[Ar]
sub-shell have one electron (both shown as ↑).
Electron configuration
This means that potassium and calcium have the electron structure [Ar]4s1
and [Ar]4s2 respectively (Table 15.1).
alt
ectron configuration
n 1s2 2s2 2p 4 .
↑
↑
↑
↑
Scandium
Sc
[Ar]3d14s2
[Ar]
Oxygen atoms have only two unpaired electrons as one of the↑2p orbitals ↑ ↑
2
[Ar]3d24sdirections
↑ ↑
has a pair ofTitanium
electrons whichTispin in opposite
to[Ar]
minimise
their ↑
↑
3
2
repulsion forVanadium
each other.
[Ar] ↑ ↑ ↑
V
[Ar]3d 4s
↑
Cr
[Ar]3d54s1
[Ar] ↑ ↑ ↑ ↑ ↑
↑
Manganese
Mn
[Ar]3d54s2
[Ar] ↑ ↑ ↑ ↑ ↑
↑
Iron
Fe
[Ar] ↑ ↑ ↑ ↑ ↑
Co
[Ar]3d64s2
4p
[Ar]3d74s2
4s Ni
[Ar]3d84s2
Chromium
on configuration of a vanadium atom.Cobalt
Nickel
↑↓
Copper
Cu ↑↓
1
3p4s1
↑↓ [Ar]3d
↑↓ 10
↑
↑ ↑
↑
↑
↑
↑
↑
↑
[Ar] ↑ ↑ ↑ ↑ ↑ ↑
3d ↑
↑ ↑ ↑ ↑
[Ar] ↑ ↑ ↑ ↑ ↑ ↑
[Ar]
↑ ↑ ↑ ↑ ↑
↑ ↑ ↑ ↑ ↑ ↑
mber of vanadium is 23. For 23
↑ ↑ ↑ ↑ ↑ ↑
10 4s2
[Ar]3d
Zinc
Zn
[Ar]
↑ ↑ ↑ ↑ ↑ ↑
↑↓ 3s
again the 1s2 2s2 2p6 3s2 3p6 sub18 electrons in total). This leaves five
↑↓ ↑↓ ↑↓ 2p
ce in the 4s and 3d sub-shells. The 4s
Look are,
carefully
Table
15.1. In orPeriod
4, the
d-block
elements run
There
however,
unexpected
anomalous
electron
configurations
that from
break
↑↓ 2sat two
two electrons and then the remaining
22s22p63s23p63d14s2) to zinc (1s22s22p63s23p63d104s2). But, notice
scandium
(1s
the Aufbau principle, namely, those of chromium and copper. A simple explanation to
are placed in the 3d (all spinning in that the electronic configurations of chromium and copper do not fit the
explain the existence of these electronic arrangements is to suggest that half-filled
ion in different orbitals).
↑↓ 1s The explanation of these irregularities lies in the stability
general pattern.
and filled 3d sub-levels are both particularly stable electron configurations.
associated with half-filled and filled sub-shells. So, the electronic structure
Figure
1.18 5The
configuration
of a vanadium
atom
4s1, electron
with
half-filled
sub-shells
an equal
distribution
of chromium,
[Ar]3d
54s1, with and
half-filled
sub-shells
and
So,
the electronic
structure
chromium,
2 2s
2 2p 6 3sof
2 3p
6 3d 3 4s 2[Ar]3d
is
1s
.
of
charge
around
the
nucleus,
is
more
stable
than
the
electronic
structure
an equal4 distribution
of charge around the nucleus, is more stable than the electronic
[Ar]3d 4s2. 4 2
structure [Ar]3d 4s .
Similarly, the electronic structure of copper, [Ar]3d10 4s1, with a filled 3d
9 2
sub-shell and a half-filled 4s sub-shell
4p is more stable than [Ar]3d 4s .
on configuration of a chromium atom.
Along the series of d-block elements from
3dnumber
↑ scandium
↑
↑ to↑ zinc,
↑ the
of protons↑ in the
4s nucleus increases by one from one element to the next.
However, the added ↑↓
electrons
go into an inner d sub-shell, but the outer
↑↓ ↑↓ 3p
mber of chromium is 24.
electrons are always in the 4s sub-shell. This means that there are clear
↑↓ 3s
ans the 1s2 2s2 2p6 3s2 3p6 are full assimilarities amongst the transition elements. Changes in their chemical
↑↓ 2pless marked than the big changes across
↑↓ ↑↓
are much
rons in total). This leaves six electronsproperties across the series
a
series
of
p-block
elements
such
as aluminium to argon.
↑↓ 2s
s and 3d sub-shells. By moving one
he 4s to the 3d the chromium atom can
d 3d sub-shell (3d5). A half-filled or
is more stable so this is a more stable
uration for the chromium atom.
5
↑↓
1s
15
Figure 1.19 The electron configuration of a chromium atom
is 1s2 2s2 2p 6 3s2 3p 6 3d5 4s1.
15.1 The atoms and ions of transition e
469983_15_Chem_Y1-2_418-449.indd 419
Marginalizer
Electronic Structure
Bilal Hameed
36
Bilal Hameed
3/27/19 1:54 PM
The electron configuration of vanadium can be written as 1
Note that the three 3d electrons are in different orbitals and
35
Tip
3p
3s
Vanadium’s
electronic 35
ELECTRONIC
STRUCTURE
2p
configuration
sometimes
Similarly,
the electronic is
structure
of copper, [Ar]3d104s1 , with a filled 3d sub-shell
2s
2 2s2 2p6 3s2
written 1s
and a half-filled
4s sub-shell
is more stable than [Ar]3d94s2.
4s
3p6 are full (18
rons to place in
mium, copper
as one electron
giving a more
↑↓
↑↓
↑↓
4s
↑↓
3d
3p
alt
↑
↑↓
Energy
opper atom.
3p6 4s2 3d3, which is in
order of increasing energy.
4p
4p
3s
1s
↑↓
↑↓
↑↓
2p
3d
2s
↑↓
↑↓
↑↓
3p
↑↓
↑↓
↑↓
2p
3s
2s
1
1s
1s
(a) magnesium
(b) vanadium
Figure 2.15 The electron configuration of (a) a magnesium atom and (
A full list of the electron configurations of the first 36 eleme
Figure 1.20 The electron configuration of a copper atom is
A short
waybeing
of writing
an electron configuration is to give the
For all
the d-block metals the 3d and 4s sub-levels,
despite
from different
1s 2 2s2 2p 6 3s2 3p 6 3d10 4s1.
shells, are relatively close in energy. The low energy
means that
3d detail
and
nobledifference
gas followed
bythethe
of the electrons added subseq
4s electrons can both be regarded as valence electrons and involved in bonding.
electron configuration of mang
The atoms of chromium and copper have an unusual electron configuration
because of the stability of filled and half-filled sub-shells. Chromium is
1s2 2s2 2p6 3s2 3p6 3d5 4s1 instead of 3d4 4s2. This is because one electron
in each of the orbitals of the 3d sub-shell makes it more stable – they are
symmetrical around the nucleus. Copper is 1s2 2s2 2p6 3s2 3p6 3d104p4s1
3d
in
instead of 3d9 4s2 as would have been expected. Again the ten electrons
the 3d sub-shell make it more stable due to symmetry around the4snucleus.
Electron configuration of simple ions
that of bromine is [Ar] 3d10 4s2 4
2s2 2p6 3s2 3p6.
5s
Note that the d-orbitals fill ‘la
the next orbit (or shell) has rece
there are two slight variations i
3p
When atoms form simple ions they can either lose or 3s
gain electrons.
Ions formed from metal atoms
The order of filling orbitals is
4p, 5s, 4d, 5p, 6s… This is sho
Figure 2.16.
2p
●
2s
●
Metal atoms tend to lose electrons
to become positive ions.
1s
The number of electrons they
lose is thend same as therd positive charge
on
1st
2
3
4th
5th
the ion.
period
period
period
period
period
The name of a positive
ion is
theThe
same
as inthe
atom,
e.g. Na
a sodium
Figure
2.16
order
which
orbitals
are is
filled
+
atom and Na is a sodium ion; Al is an aluminium atom and Al3+ is an
aluminium ion.
When metal atoms lose electrons they lose them from the outermost level
except atoms of d block elements.
24
2 Atomic structure and the periodic table (Topic 1)
Atoms of d block elements lose
their 4s electrons first then their 3d.
chromium is [Ar] 3d5 4s1 no
copper is [Ar] 3d10 4s1 not [A
This is because an atom is mo
filled or filled set of 3d-orbital
the 4s-orbital, rather than four
two in the 4s-orbital.
Ions formed from non-metal atoms
Non-metal atoms tend to gain electrons to become negative ions.
The number of electrons they gain is the same as the negative charge on
807404_C02_Edexcel_GF_Chem_009-036.indd 24
the ion.
The name of a negative ion is the atom stem with ‘-ide’ on the end, e.g.
O is an oxygen atom and O2− is an oxide ion; Br is a bromine atom and
Bilalion.
Hameed
Br− is a bromide
Hydrogen
Bilal Hameed
A hydrogen atom has only one 1s electron.
37
Marginalizer
Electronic Structure
Potassium
19
Sulfur
16
1
K
1s2 2s2 2p6 3s2 3p6
6
2
1s2 2s2 2p6 3s2 3p6
S
Note in the table above that Ne, Mg2 and N3 36
ions have the same electron
configuration, and Ar, K and S2 have the same electron configuration. These
groups of ions are referred to as isoelectronic—having the same number
of electrons.
ELECTRONIC STRUCTURE 36
Electron configurations of metal ions
Transition metal ions
Ions
of theofd-block
elements
The majority
the 3d transition
metals have a full 4s subshell, although as
discussed
above, chromium
and copper
have
only one
electron
in the
a 4s first
subshell.
When
a d-block
metal ionizes
to form
a simple
positive
ion,
electrons to be
A full 4s subshell has a slightly higher energy than a partially filled 3d subshell
lost
are
the
4s
electrons,
followed
by
the
3d
electrons.
In
other
words,
when a d-block
(see chapter 3), so when a 3d transition metal loses electrons, the electrons are
0
n
initially
lost from
the 4s ions
subshell.
If more which
than 2 electrons
the configurations.
3d then
electron
metal
ionizes,
positive
are formed
possess are
4s lost,
alt
subsequent electrons are taken from the 3d subshell. Similarly, electrons are
taken from the 5s subshell before the 4d subshell.
For example, if the iron(ii) ion is formed, only the two 4s electrons are lost, but if the
The electron configurations of some transition metal ions are shown below.
iron(iii)
ion is formed an additional electron is lost from the spin pair of the 3d subThese will be discussed in more detail in chapter 3.
level. Some examples of common d-block simple ions are shown below.
Ion symbol
Number of
electrons in
the ion
26
Fe2
24
1s2 2s2 2p6 3s2 3p6 3d 6
26
Fe
3
23
1s2 2s2 2p6 3s2 3p6 3d 5
22
2
20
1s2 2s2 2p6 3s2 3p6 3d 2
3
21
1s2 2s2 2p6 3s2 3p6 3d 3
Iron
Iron
Vanadium
V
1
Electron configuration
Chromium
24
Cr
Copper
29
Cu
28
1s2 2s2 2p6 3s2 3p6 3d 10
29
454 13 The periodic table – Copper
the transition metals
2
27
1s2 2s2 2p6 3s2 3p6 3d 9
■ Table 13.3 Selected
examples of the more
stable simple ions from
the first row of the
d-block
d-block metal
Scandium
Titanium
Vanadium
2
Explain why
vanadium is a
d-block metal and
transition element.
Give the full
detailed electron
configuration for
the copper(III),
cobalt(III) and
chromium(II) ions.
Nature of Science
■ Table 13.4
Comparison of a
typical transition
element with calcium
Chromium
Manganese
Iron
Cobalt
Nickel
Copper
Zinc
Simple
ion
Sc3+
Ti3+
Ti4+
V2+
V3+
Cr3+
Mn2+
Mn4+
Fe2+
Fe3+
Co2+
Ni2+
Cu+
Cu2+
Zn2+
A transition element is defined as a d-block
Detailed
metal that forms at least one stable cation with
outer electron
configuration
anFOR
incomplete
3d sub-level.
the elements
CHEMISTRY:
USE WITH THE
IB DIPLOMA All
PROGRAMME
3d04s0
in Table 13.3 conform to this except zinc which
3d14s0
is therefore not a transition element. Copper
3d04s0
is regarded as a transition element since it
3d34s0
forms the stable copper(ii) ion, which has
3d24s0
an incomplete d sub-level. Scandium is also
3d34s0
regarded as a transition element since it can
3d54s0
form Sc+ ([Ar]3d14s1) and Sc2+ ([Ar]3d14s 0) in
3
0
3d 4s
a limited number of compounds. For example
3d64s0
the compound CsScCl 3 [Cs+ Sc2+ 3Cl−] has
5
0
3d 4s
scandium in oxidation state +2.
3d74s0
Ions with a half-filled 3d sub-level (3d5) or a
8
0
3d 4s
filled 3d sub-level (3d10) are usually relatively stable,
3d104s0
but a number of factors are involved in determining
3d94s0
the stability of transition metal compounds in the
3d104s0
solid state (including lattice enthalpies).
Looking for trends and discrepancies – the anomalous behaviour of zinc, chromium
and copper
The elements from scandium to zinc form the first row of the d-block (3d-block). The first row
of this block contains ten elements, because the 3d sub-level contains five 3d orbitals, each
able to accommodate two electrons (a spin pair). All of the elements from scandium to copper
are in the d-block because the 3d sub-level is being progressively filled. Originally the d-block
was known as the ‘transition metals’ because some of their properties show a gradual change
between the reactive metal calcium in group 2 to the much less reactive metal gallium in
group 13. The term ‘transition metal’ is now reserved for those metals in the d-block that show
properties characteristically different from those in the s- and p-blocks (the representative
elements). Table 13.4 compares the properties of a typical transition element with calcium.
Property
Transition element
Calcium
Melting point
Very high (>1000 °C)
Lower than transition elements (850 °C)
Density
Very high
Lower than transition elements
Atomic radius
Smaller than calcium
Larger than transition elements
Ionic radius (M2+ ion)
Smaller than calcium
Larger than transition elements
First ionization energy
Marginalizer
Electronic Structure
Cu
Electrical conductivity
Larger than calcium
Smaller than transition elements
Bilal
Good, but poorer than
38
calcium
Very good – better than most transition elements
Hameed
Bilal Hameed
Zinc is a d-block element that is excluded from being classified as a transition element.
CHAPTER 1
Atomic
number (Z)
ATOMIC STRUCTURE
TABLE 1.3.5 ELECTRON CONFIGURATIONS OF SOME TRANSITION METAL IONS
Element
19