Periodic Trends 1
Elements show trends in their physical and chemical properties across periods and
INORGANIC CHEMISTRY
down groups.
s block
2 Group
1 1
Period
2 Li
18
p block
13 14 15 16 17 He
H
Be
d block
B
C
N
O
F
Ne
Al
Si
P
3 Na Mg
3
4
5
6
S
Cl
Ar
4 K
Sc
Ti
V
Cr Mn Fe Co Ni Cu Zn Ga Ge As Se
Br
Kr
Y
Zr Nb Mo Tc Ru Rh Pd Ag Cd
Ca
5 Rb Sr
f block
6 Cs Ba La Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu Hf Ta
7 Fr
Ra Ac Th Pa
U
Np Pu Am Cm Bk Cf
Figure 10.2 A modern form of the Periodic
Table. The elements shown in red are good
conductors of electricity, and the ones shown
in blue are poor conductors. The pink shading
indicates that the principal form of the element
is a semiconductor. As only a few atoms of the
elements at the end of the Periodic Table have
10_02 Cam/Chem AS&A2
been made, these have been left white.
Es Fm Md No Lr
7
8
9
W Re Os
r
10 11
12
n
Pt Au Hg TI
Sn Sb Te
Pb
Bi
Xe
Po At Rn
Rf Db Sg Bh Hs Mt Ds Rg Cn 113 Fl 115 Lv 117 118
in rows or periods across the table. The elements in a period have different physical
and chemical properties, but trends become apparent in these properties as we move
across a period.
The periodic table is divided into groups, periods and blocks. The International Union
of Pure and Applied Chemistry, IUPAC, now recommends that the groups in the
The s, p, d and f blocks
periodic table should be numbered
1 togroups
18. form the s block (Groups 1 and 2). These have the outer
The first two
1
electronic configurations ns1 and ns2, respectively, where n is the number of the
shell. The last six groups form the p block (Groups 13 to 18), in which the p
sub-shell is being progressively filled. These elements have the outer electronic
configurations ns2 np1 to ns2 np6.
In between the s and p blocks is the d block (Groups 3–12), in which the d
sub-shell is being progressively filled. Our study of the d block is largely restricted
to the elements scandium to zinc, and these elements have the outer electronic
configurations 4s2 3d1 to 4s2 3d10, as we shall see in Topic 24.
About a quarter of the known elements belong to the f block. The first row, the
4f, includes all the elements from cerium (Ce) to lutetium (Lu) inclusive. In this block,
the f sub-shell in the fourth principal shell is being progressively filled. In the past,
these elements were called the ‘rare earths’. This was a misnomer, as their abundance
in the Earth’s crust is fairly large – for some it is comparable to that of lead. However,
it is true that they are mostly quite thinly spread, so mining them is quite expensive.
They are now called the lanthanoids, as their properties are similar to those of the
element lanthanum (La) that precedes them in the Periodic Table. Although they are
not particularly rare, the lanthanoids are difficult to purify from one another because
their chemical properties are nearly identical. The elements of the second row of the
f block, the 5f, are called the actinoids. All the actinoids are radioactive, and most
have to be made artificially.
Groups
Dog Art1 and 2 remain the same as before – classified as the s-block.
•Barking
• The transition elements now become Group 3 to 12 – classified as the d-block.
• Groups 3 to 7 now become Groups 13 to 17 and the noble gases become Group
18 – classified as the p-block.
All the elements in a group contain the same number of electrons in their outer shell.
For example, all the elements in group 1 have one electron in their outer shell, which is
in an s-orbital. All the elements in group 17 have seven electrons in their outer orbit,
arranged s2p5.
The horizontal rows are called periods. All elements in the same period have the
same number of shells containing
electrons.
For
example,
all
the
elements
in
the
third
10.3 Periodic trends in the elements of
period, Na to Ar, have electrons
in
the
first,
second
and
third
shells.
the third period (sodium to argon)
1
Key
atomic number
symbol
relative atomic mass
1
H
In this section we shall look at the trends in the properties of the elements and their
18
compounds in the third period of the Periodic Table.
2
He
Appearance
1.0
2
3
4
Li
Be
6.9
9.0
11
12
Na
Mg
23.0
24.3
3
4
5
6
19
20
21
22
23
24
K
Ca
Sc
Ti
V
Cr
39.1
40.1
45.0
47.9
50.9
52.0
37
38
39
40
41
42
Rb
Sr
Y
Zr
Nb
Mo
85.5
87.6
88.9
91.2
92.9
95.9
55
56
72
73
74
Cs
Ba
Hf
Ta
W
178.5
180.9
183.8
186.2
190.2
192.2
195.1
197.0
200.6
104
105
106
107
108
109
110
111
112
114
116
Rf
Db
Sg
Bh
Hs
Mt
Ds
Rg
Cn
Fl
Lv
188
132.9
137.3
87
88
Fr
Ra
57-71
89–103
13
14
15
16
17
5
6
7
8
9
204.4
207.2
209.0
4.0
10
The elements on the left of the Periodic Table have low values of ionisation energy
B
C
N
O
F
Ne
and electronegativity, and so they show the properties
associated
with
metallic
10.8
12.0
14.0
16.0
19.0
20.2
bonding (see section 4.11), for example, they13are shiny
conduct16 electricity.
In18
14 and 15
17
Al higher
Si values
P of ionisation
S
ClenergyAr
the middle of the Periodic Table, elements with
27.0
31.0
35.5
7
8
11
12
9
and
electronegativity
are10semiconductors:
they
have 28.1
a dull shine
to32.1
them and
are 39.9
25
26
27
30
28
29
31
32
33
34
35
36
−12
poor
of electricity
(typically
times that
metal).SeThe elements
Mn conductors
Fe
Co
Zn10 Ga
Ni
Cu
Ge of aAs
Br
Kr
on
the right
the highest
values
of74.9
ionisation
and
72.6
79.9
54.9
55.8 of the
58.9 Periodic
58.7 Table,
65.4
63.5 with
69.7
79.0 energy
83.8
43
44
45are dull
46 in appearance
47
48 and are
49 such
50poor conductors
51
52
54
electronegativity,
that53they are
Tc as electrical
Ru
Rhinsulators
Pd (their
Ag conductivities
Cd
In are only
Sn about
Sb 10−18
Tetimes Ithat ofXe
used
102.9
106.4
101.1
107.9
112.4
114.8
118.7
121.8
127.6
126.9
131.3
a 75
metal). 76
86
79
80
83
84
85
77
78
81
82
Ir
Au
Re
Os
Pt
Hg
Po
At
Rn
Bi
Tl
Pb
333_10_AS_Chem_BP_186-201.indd 188
18/09/14 1
57
58
59
60
61
62
63
64
65
66
67
68
69
70
71
La
Ce
Pr
Nd
Pm
Sm
Eu
Gd
Tb
Dy
Ho
Er
Tm
Yb
Lu
138.9
140.1
140.9
144.2
144.9
150.4
152.0
157.2
158.9
162.5
164.9
167.3
168.9
173.0
175.0
89
90
91
92
93
94
95
96
97
98
99
100
101
102
103
Ac
Th
Pa
U
Np
Pu
Am
Cm
Bk
Cf
Es
Fm
Md
No
Lr
232.0
238.1
Figure 8.1 The periodic table.
The International Union of Pure and Applied Chemistry, IUPAC, now
recommends that the groups in the periodic table should be numbered
1 to 18.
●
Groups 1 and 2 remain the same as before – classified as the s-block.
●
The transition elements now become Group 3 to 12 – classified as
the d-block.
●
Groups 3 to 7 now become Groups 13 to 17 and the noble gases
Groups
3
to
7
now
become
Groups
13
to
17
and
the
noble
gases
Periodic Trends 2
become Group 18 – classified as the p-block.
●
The table is also divided into blocks:
The blocks are shown in Figure 8.2.
pblock
sblock
dblock
1
Figure 8.2 Blocks of the periodic table.
The s-block consists of the elements in groups 1 and 2. An s-block element has its
The
vertical
columns
are
called
groups
and
the
horizontal
rows
are
highest-energy electron in an s-orbital, i.e. the last electron added goes into an scalled
periods.
Trends
in
the
groups
(vertical
columns)
and
periods
orbital.
(horizontal rows) reflect the structures of the atoms of the elements
within
them,in and
these in
affect
the These
chemical
properties
the
The
elements
the s-block
areturn
reactive
metals.
metals
(includingof
potassium,
elements.
These
properties
sodium,
calcium
andrepeating
magnesium)
are all highdemonstrate
in the activity periodicity.
(electrochemical) series.
They have lower densities, lower melting points and lower boiling points than most
other metals and they form stable, involatile ionic compounds.
Group 1: Also called the Alkali Metals. Soft, highly reactive group of metals. ns1.
30/03/15
Group 2. Also called the Alkaline Earth Metals. Fairly reactive metals. ns2
The d-block elements occupy space in the periodic table between Group 2 and
Group 13. The d-block contains the elements scandium to zinc in period 4 and
those elements below them.
A d-block element cannot be defined in terms of energy because the energy level of
the d-orbitals is altered by the presence of electrons in the outer s-orbital. It can only
be defined in aufbau terms, i.e. the last electron added to a d-block element goes into
a d-orbital.
These metals (including chromium, iron, copper, zinc and silver) are much less
reactive than the metals in Groups 1 and 2.
Periodic Trends 3
Although the periodic table does not classify elements as metals and
The p-block contains
elements
in groups
13 tobetween
17 and group
18 (the
noble
non-metals,
there is athe
fairly
obvious
division
the two
(‘fairly
obvious’,
but,
notA ‘clear
cut’).
Separating
the elements into
either
or i.e.
non-metals
gases).
p-block
element
has its highest-energy
electron
in ametals
p-orbital,
the last is
rather
like
trying
to
separate
all
the
shades
of
grey
into
either
black
or
white.
electron added goes into a p-orbital.
The fairly obvious division between metals and non-metals is shown by a thick
stepped
line
in
Figure
9.5.
The
20
or
so
non-metals
are
packed
into
the
top
Some of the elements in the p-block are metals, like tin, lead and bismuth. They are
right-hand corner above the thick stepped line. Some of the elements next to
usually low in the activity (electrochemical) series and they have some resemblances
the thick steps, such as germanium, arsenic and antimony, have similarities to
to
non-metals.
Other
elements
in
the
p-block
are
non-mental.
both metals and non-metals and it is difficult to place these in one class or the
other. Chemists sometimes use the name metalloid for these elements which
2 np5
Group
17.
Also
called
the
Halogens.
Very
reactive
group
of
non-metals.
ns
are difficult to classify one way or the other.
8
Periodicity
8.1 The
Periodic
Table
1
Figure
9.8Also
shows
a classification
of elements
as metals,
metalloids
and nonGroup 18.
called
the Noble Gases
because they
do not react.
Form few
metals
on
the
basis
of
their
electrical
conductivity.
In
this
classification:
2
6
compounds. ns np
Metals
Metalloids
Non-metals
The Periodic Table is a list of all the elements in order of increasing atomic
number. You can predict the properties of an element from its position in
Li Be
B C N O F
Ne
the table. You can use it to explain the similarities of certain elements and
Na Mg
Al Si P S Cl
Ar
the trends in their properties, in terms of their electronic arrangements.
Metals
K Ca
Ga Ge As Se Br
Rb Sr
In Sn
Sb Te
The structure
ofd-block
the and
Periodic
Table
Cs Ba
f-block metals
I
Kr
Non-metals
Xe
Tl Pb Bi Po At
Rn
The Periodic Table has been written in many forms including pyramids
Ra one shown below is one common layout. Some areas
and spirals.FrThe
ofFigthe
Table
are asgiven
are shown
in Figure
9.8 Periodic
Classifying the
elements
metals,names.
metalloidsThese
and non-metals
on the basis
of their 1.
1
2
alkali metals
alkaline earth metals
electrical conductivity. In this figure, metalloids are coloured grey
3
4
5
6
7
0
noble (rare or inert) gases
4
halogens
1 Metals are good conductors of electricity with atomic electrical conductivity*
1 greater than 10−3 ohm−1 cm−4.
2 Metalloids are poor conductors of electricity with atomic electrical conductivity
2
−3
−5
−1
−4
usually less than 10 but greater than 10 ohm cm .
metalloids
3 3 Non-metals are virtually non-conductors
(insulators). Their atomic electrical
−10
−1
−4
conductivity is usually less than 10 ohm cm .
Notice in Figure 9.8 that the cell for carbon is shaded less heavily than those for
transition metals
5
the other metalloids. This is because carbon exists as three different allotropes –
6 graphite, a poor conductor classed as a metalloid, plus diamond and fullerenes,
insulators classed as non-metals.
7
In spite of problems such as this, the classification of elements into metals,
metalloids and non-metals is useful and convenient.
Lanthanides
lanthanides
Actinides
actinides
▲ Figure 1 Named areas of the Periodic Table
Metals and non-metals
The red stepped line in Figure 1 (the ‘staircase line’) divides metals
(on its left) from non-metals (on its right). Elements that touch this
line, such as silicon, have a combination of metallic and non-metallic
properties. They are called metalloids or semi-metals. Silicon,
for example, is a non-metal but it looks quite shiny and conducts
electricity, although not as well as a metal.
t
The
Periodic
Table
reveals
patterns
in
the
properties
of
e
from Na table
→ Mg omits
→ Al. the t
period number. This so-called ‘short form’ of the periodic
Periodic Trends 4
example, every time you go across a period you go from
left to non-metals
on
right. 15This is16 an example
1
2
13 the 14
17
18 of pe
On the other hand, metalloids and non-metals with molecular structures
Period
have
no mobile
electrons.
This means
that their electrical
conductivity regularly.
is
1
word
periodic
means
recurring
2
almost nil.
Metals, such as sodium, magnesium and aluminium, on the left of each period
with mobile, delocalised electrons in their structure are good conductors of
electricity. When a battery is attached to them, electrons flow out of the metals
into the positive terminal of the battery. At the same time, electrons flow into
Group
the metal
from the negative terminal of the battery.
H
1
He
1
2
3
4 and
5
6
7
8
9
10
Periodicity
properties
of
elements
in
The
periodicity
of
atomic
properties
29.7
Li
Be
B
C
N
O
F
Ne
2,1 is explained
2,2
2,3
2,5
2,6 arrangements
2,7
2,8
Periodicity
by2,4the electron
o
Table 9.3 shows the electronic (shell and sub-shell) structures for the elements
11
12
13
14
15
in Periods 2 and 3.
In this section you will learn to:
16
17
18
• Give the electronic structure
of Period 2 and 3 elements
and compare
the number
of
2,8,6
2,8,7
2,8,8
electrons in the last sub-shell
• Describe and explain changes in
atomic and ionic radii as you go
across a period
Na
Mg
Al
Si
P
S
Cl
Ar
•
The
elements
in
Groups
1,
2,
and
3
(sodium,
magne
Elements in the same group of the periodic table have similar electron
2,8,1
2,8,2I elements
2,8,3
2,8,5
configurations. For
example, Group
(Li to Fr) 2,8,4
have one electron
in their aluminium)
outer shells (ns ). Group II elements
(Be
to Ra) have two
electrons in
are
metals.
They
have
giant
structures.
19) and Group20VII elements (F to At) have seven outer shell
their outer shells (ns
electrons
(ns
these electron configurations
recur in a periodic
pattern,
4 np ). As
outer
electrons
to
form
ionic
compounds.
K
Ca
it would not be surprising to find that atomic properties, such as atomic radii,
3
1
1
2
2
5
2,8,8,1 energies,
2,8,8,2
ionic radii and ionisation
show similar periodicity.
• Describe and explain other
trends in atomic properties such
as first ionisation energies and
electronegativity
•
Silicon
in
Group
4
has
four
electrons
in
its
outer
she
Atomic radii
Based
on
the
electron
arrangements
of
the
elements
(Chapter
12),
the
pe
The
atomic
radii
of
atoms
can
be
obtained
from
X-ray
analysis
and
electron
it
forms
four
covalent
bonds.
The
element
has
some
density maps. Using these techniques, it is possible to measure the distance
divided
intoof four
blocks
of elements
(Figure
3.7):
between
the nuclei
atoms and
then estimate
the radius of individual
atoms.
properties and is classed as a semi-metal.
KEY POINT
The atomic
radii of metals are
obtained
by measuring
the distance between the
The
periodicity
of
atomic
properties
■
s-block
elements
nuclei of neighbouring atoms in metal crystals (Figure 9.10 (a)). The atomic
•thisisThe
elements
indistance.
Groups
5, and
6, and
radius
simply half
of the
In
section
we
ll inter-nuclear
look at the
trends
in atomic
ionic 7 (phosphorus, sulfu
■ p-block elements
and electronegativity
as you goThey
across aeither
period. accept electrons to form io
are non-metals.
■ d-block elements
Table 9.3 The electronic structures of the elements in Periods 2 and 3
or share their outer electrons to form covalent compou
■
f-block
Electron shell structure
Period 2
Elements in the same group of the
periodic
table have similar
electron
radii,
first ionisation
energies
configurations. This results in
similar atomic properties.
elements.
2, 1
2, 2
Li
Be
B
C
N
O
F
Ne
2, 3
2, 4
2, 5
2, 6
2, 7
2, 8
• Argon in Group 0 is a noble gas – it has a full outer s
unreactive.
Electron sub-shell structure
1s22s1
1s22s2
1s22s22p1
1s22s22p2
1s22s22p3
1s22s22p4
1s22s22p5
1s22s22p6
Period 3
Na
Mg
Al
Si
P
S
Cl
Ar
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
1s22s22p6
3s1
1s22s22p6
3s2
1s22s22p6
3s23p1
1s22s22p6
3s23p2
1s22s22p6
3s23p3
1s22s22p6
3s23p4
1s22s22p6
3s23p5
1s22s22p6
3s23p6
Table 1 shows some trends across Period 3 (see Figure 1
trends are found in other periods.
185
1 2
3 4
5 6 7
0
1
2
3 Na Mg
Al Si P S Cl Ar
4
5
6
7
▲
Figure
1
The
Periodic
Table
with
Period
3
highlighted
To understand how such atomic properties vary we have to understand all the forces
at play inside an atom.
3
luminium
4
silicon
5
phosphorus
6
sulfur
2
1
Ne] 3s 3p
2
2
[Ne] 3s 3p
2
3
[Ne] 3s 3p
2
4
[Ne] 3s 3p
7
chlorine
2
[Ne] 3s 3
Forces that act on electrons in an
Periodic Trends 5
Forces that act on electrons in an atom
tom
The
forces
that
act
on
electrons
in
an
atom
are
electromagn
The forces that act on electrons in an atom are electromagnetic. They obey
Coulomb’s
law,
which
states
that
the
magnitude
of
the
force
betw
Coulomb’s law, which states that the magnitude of the force between two charged
objects
is
directly
proportional
to
the
product
of
their
charg
objects
is
directly
proportional
to
the
product
of
their
charges
and
inversely
. Theyproportional
obey
to
the
square
of
the
distance
between
their
centres.
proportional to the square of the distance between their centres.
two charged
mathematically as:
q+q−
and inversely
force ∝ 2
r
s is expressed
where q+ and q− are the charges on the two objects and r is the
theirto:centres.
It simplifies
•
1
Thecharges
rules are:
opposite
attract; like charges repel
the bigger
the charges,
the stronger
is the
● opposite
charges
attract;
likeforce
charges repel
•
ance between
● the bigger the charges, the stronger is the force
• the further apart the particles, the weaker is the force
● the further apart the particles, the weaker is the force
This means that electrons are attracted towards the nucleus (F
greater the atomic number, the stronger is the force of attractio
are further away from the nucleus are attracted less than those clos
In addition, because they have –the same
charge, electrons repel ea
Repulsion
they are more densely packed, the inner electrons repel outer elec
Attraction
–
+ repel each other.
than the other outer electrons
re 2.18). The
– Repulsion
Electrons
that
Shielding
and the effective nuclear charge
o the nucleus.
Effective nuclear charge across period 1
In a hydrogen atom, the nucleus has a charge of +1 and there is
ther. Because
1s-orbital. There are no forces of repulsion, so the electron feels
Figure
2.18
The
forces
acting
ns much
more
attraction
of
a
+1
charge.
This means that electrons are attracted towards the nucleus. The greater the atomic
on
an
electron
in
an
atom
number,In
theastronger
is
the
force
of
attraction.
Electrons
that
are
further
away
from
helium atom, the nucleus has a charge of +2 and there
the nucleus
are
attracted
less
than
those
closer
to
the
nucleus.
the 1s-orbital. These electrons repel each other slightly. The
are
res
force
of
attraction
between
the
nucleus
and
each
electron
is
sligh
In addition, because they have the same charge, electrons repel each other. Because
between
a
+2
charge
and
one
electron.
Therefore,
the
helium
nucl
they are more densely packed, the inner electrons repel outer electrons much more
an
eff
ective
nuclear
charge
of
slightly
less
than
2.
than the other outer electrons repel each other.
electron in a
Effective nuclear charge across period 2
e full force
of
The situation becomes more complicated for lithium and the remai
lithium, the nucleus has a charge of +3 and the outer 2s electron is
o electrons in
s that the net
less than that
is said to have
Key term
The effective nuclear
charge is the net charge
07404_C02_Edexcel_GF_Chem_009-036.indd
on the27nucleus, after
Periodic Trends 6
7
Calculate the
−34 J s × 3.00 × 10 8 m s−1)
(6.63
×
10
hc
= outer
=
wavelength
(in m)
Factors that effect
the force
of attractionE on
electrons
−9 m
λ
600
×
10
of electromagnetic
Nuclear charge
−17 J
4.0
×
10
radiation with
= 1.2 × 102 p
When the nuclear charge
becomes
of additional
a frequency
of more positive (due to the
3.31presence
× 10 −19 J/photon
protons), its attraction1368
on allkHz.
theDeduce
electrons increases.
Calculate the number of photons with wavelength
which part of the
electromagnetic
−34 J s × 3.
Atomic Radius
(6.63
×
10
hc
spectrum it belongs
For one photon: E =
=
λ attraction of 4 × 10 −9
As the distance of theto.
outer electrons from the nucleus increases, the
1
the positive nucleus
for
the
negatively
charged
electrons
falls.
8 Calculate the
= 2.01 × 1016
So for one joule:
−17
4.9725
×
10
frequency of
yellow light with
Shielding Effect
a wavelength of
The outer or valence electrons are
repelled by all the other electrons in the atom in
−8
cm.
5800
×
10
74 2 Atomic structure
addition to being attracted by the positively charged nucleus. The outer electrons are
9 The laser used to
Ionization
energy
shielded from the attraction
of
the
nucleus
by
the
shielding
effect.
read information
The third shell can hold a
The first ionization energy
is
the
minimum
from a compact disc
However, when there are e
has a wavelength of
one
mole
of
isolated
gaseous
atoms
to
form
o
shell
there
is
a
degree
of
st
–
–
780 nm. Calculate
thermodynamic conditions.
example,
the
electronsFor
enter
the fourth
the energy
required
to
bring
about
the
reaction:
metals
beyond
calcium
the
associated with This electron does not
–
the
third
shell
until
it
con
feel
the
full
effect
of
the
positive
– one photon of this
+
−
Cl(g)
→
Cl
(g)
+
e
electrons.
In addition, the
charge
of
the
nucleus
–
radiation.
1
Electron configuration
Explaining periodic patterns
Atomic rad
+
■
shells are divided into a nu
The electron is removed
from
the
outer sub-s
–
(other
than
hydrogen)
also
–
is,
a
3p
electron).
Table
12.1
gives
some
exam
before
they
can
form
chemi
–
energy,
which
is
the
enthalpy
change
for
the
This
process
is
called
–
important concept in
the
IB the
Chemistry
dataAn
booklet.
These electrons
shield
outer
3 and also in Chapter 12
electron from the nucleus
–
radii
(Figure 2.50).Atomic
The electro
Table 12.1 Selected
Element
Ionization equation
experience
different
attra
ionization energies
+ (g) + e −
Oxygen
O(g)
→
O
Figure 2.50 Electron shielding
presence of other electron
+ (g) + e −
Sulfur
S(g)
→
S
experience the most shiel
Na
ToK Link 10–
Cu(g) → Cu+(g) + e −
–
Copper
1
Figure 8.3 Periodici
Jacob Bronowski: ‘OneFactors
aim of the physical
sciences has
been to give anenerg
exact p
that affect
ionization
One achievement… has been to prove that this aim is unattainable’.
are t
Across a What
period,
Values
of
ionization
energies
depend
on
the
f
for the aspirations
of
natural
sciences
in
particular
and
for
knowledge
in
genera
11+
one
the
size of the
atom (orFrom
ion)
This claim is probably related
to ‘modern’
(approximately
100
years
old) atom
physics suo
special relativity and quantum physics, including the Uncertainty Principle. An exa
the
nuclear
charge
●
material or physical world is impossible.
Quantum
mechanics
means
that
the
wor
the
charge
of
shielding
molecular level, and there are
limits
of
experimental
precision
dictated
by
the
Unc
the
shielding
effect.
electrons in
the atomic level is ‘schizophrenic’ due to its wave–particle
duality.
The
material
w
●
the shielding r
full shells
of scientific models, all of which are limited and incomplete as descriptions of phy
Atomic
radius
There is 8.4
no absolute
knowledge
in science
method
th
Across
theassumes
period
Figure
Shielding
effect of
inner and the scientific
of
objects
and
phenomena
existing
out
there
that
is
independent
of
the
observ
As
the
distance
shell electrons reduces the pull of the
increases
and
the
some physicists might question the assertion that the material world is indepen
the
attraction
of
nucleus
in
the
outer
shell.
Schrödinger’s cat is both alive and dead until the observation is made, i.e. the b
10 Find out about the
thought experiment
‘Schrödinger’s Cat’
Down
falls. Tt
function is collapsed, and one of the eventualities – alive
orelectrons
dead a
– isgroup,
manifeste
–
ionization
energ
Niels Bohr wrote, ‘It is wrong to think that the task of physics is to find out how
what we can say about nature.’
3+
nuclear
pull
From one atom o
Nuclear charg
● the charge of
When the nucle
ToK Link
of addii
● presence
the shielding
–
Heisenberg’s
Uncertainty Principle states that there is a theoretical limit to the p
increases.
This
c
can know the momentum and the position of a particle. What are the implicatio
–
Down
the
group
human knowledge?
Shielding
effec
decreases
and
th
In 1927 the German physicist Heisenberg stated the Uncertainty Principle, which
ity
The
outer
or
val
dual behaviour of matter and radiation
(de
Broglie’s
hypothesis).
It
states
that
it
repulsion from
simultaneously the exact position and exact
(orin
velocity)
of an in
elect
the atom
a
innermomentum
shell of
particle) along a given direction.
electrons (’shielding’)
Mathematically,
it can beforces
described
by the equation:
Figure
12.13 Electrostatic
operating
on the
h
nucleus. The ou
the nucleus by t
Periodic Trends 7
nickel
Shielding and the effective nuclear charge
The effective nuclear charge is the net charge on the nucleus, after allowing for the
electrons in orbit around the nucleus shielding its full charge.
■ Electron arrangement and the periodic table
Effective
nuclear
charge
across
period
1
It is the electrons in the outer or valence shell that determine the chemical a
In
a
hydrogen
atom,
the
nucleus
has
a
charge
of
+1
and
there
is
one
electron
in
a
1sproperties of the chemical element. The position of a chemical element in th
orbital.
There
are
no
forces
of
repulsion,
so
the
electron
feels
the
full
force
of
attraction
related to its electron arrangement. The period number indicates the numbe
of
a
+1
charge.
atom of the element. All chemical elements in the same period have the sam
In groups
1
and
2,
the
number
of
valence
electrons
is
equal
to
the
group
num
In a helium atom, the nucleus has a charge of +2 and there are two electrons in the
18, 1s-orbital.
the number
of
valence
electrons
is
equal
to
the
group
number
minus
10.
These electrons repel each other slightly. The result is that the net force of
Figure
3.6
shows
how
the
electron
arrangement
of
a
chemical
element
is
re
attraction between the nucleus and each electron is slightly less than that between a
period
number.
This
so-called
‘short
form’
of
the
periodic
table
omits
the
trans
+2 charge and one electron. Therefore, the helium nucleus is said to have an
1
effective nuclear charge of slightly less than 2.
Group
Period
1
2
13
14
15
16
17
1
2
Effective
nuclear
charge
across
period
2
1
H
He
1
2
3
5
2,2
2,3
2,4
2,5
11
12
13
14
15
Mg
Al
C
7
Be
Na
B
6
Li
2,1
3
4
Si
N
P
18
2
8
O
9
10
2,6
F
2,7
Ne
16
17
18
S
Cl
2,8
Ar
The situation
complicated
lithium and
elements. In
2,8,1 becomes
2,8,2 more
2,8,3
2,8,4 for2,8,5
2,8,6the remaining
2,8,7
2,8,8
lithium, the
nucleus
has
a
charge
of
+3
and
the
outer
2s
electron
is
strongly
repelled
19
20
4 two
by the
electrons. The nucleus is shielded by the inner electrons and the
Kinner 1sCa
2,8,8,1
2,8,8,2
effective
nuclear
charge is approximately +1. This is the +3 nuclear charge, minus
the effect of two negatively charged screening electrons.
Based on the electron arrangements of the elements (Chapter 12), the period
divided
intoelement
four blocks
of elements
(Figure
The next
is beryllium.
The nucleus
has a3.7):
charge of +4, there are two 1s
shielding the nucleus and two 2s electrons that also repel each other
■ electrons
s-block elements
slightly. Therefore, the effective nuclear charge is not exactly +2 (+4 nuclear charge
■ p-block elements
minus the effect of the two negative inner electrons) — it is slightly less than +2
■ because
d-blockofelements
the extra repulsion by the two electrons in the outer orbit.
■ f-block elements.
The situation is slightly more complicated with the next element, boron. The atomic
number of boron is five (a nuclear charge of +5). There are two 1s electrons that
shield the outer electrons from the nucleus. The two 2s electrons are closer to the
nucleus than the single 2p electron and they repel it. Therefore, the effective
nuclear charge is significantly less than the +3 value predicted by the simplified
idea that effective nuclear charge is equal to the atomic number of the element minus
the number of inner-shell electrons.
1.
Periodic Trends 8
Electron arrangement and the periodic table
he electrons
in the outer or valence shell that determine the chem
shielding
force
erties of the chemical
element. The position of a chemical element
Li
Be
B TheCperiod
N number
O
F
Ne the n
ed to its electron arrangement.
indicates
of the element. All chemical elements in the same period have th
oups 1 and 2, the number of valence electrons is equal to the group
he number of valence electrons is equal to the group number minu
electrostatic
igure 3.6 shows
how the electron arrangement of a chemical elemen
attraction
towards
d number. This
so-called ‘short form’ of the periodic table omits the
Group
d
1
positive
nucleus
1
1
2
13
14
15
16
17
Figure 12.15 A diagram illustrating how the balance
18
2
Similar arguments apply to other periods — the effective nuclear charge increases
1
between shielding and nuclear charge changes across
across a period, but does not increase by as much as +1 between successive
1 period 2
2
electron
elements.
attracted
3
4
5
6
7
8
9
10
by an
2
effective
H
He
inner
elect
shiel
valen
elect
from
nucle
Be of attraction
B
C
N electrons,
O so they
F areNe
creasesLithe force
on all the
2,1
2,2
2,3
2,4
2,5 charge
2,6
2,711+ 2,8
gly. Each additional electron across a periodofenters
the
same
+1
11
12
13
14
15
16
17
18
increase
in
shielding
is
minimal
(Figure
12.15).
3
Na
Mg
Al
Si
P
S
Cl
Ar
for the
ionization
energy
to
increase
across
the
period,
there
2,8,1
2,8,2
2,8,3
2,8,4
2,8,5
2,8,6
2,8,7
2,8,8
energy19across 20
periods 2 and 3 (Figure 12.14). These dipssodium
can
core
charge
=
11
–
10
=
+1
The
core
charge
increase
model
of
electronic
structure.
4 electron
K
Ca
inner
this
electron
attracted
2,8,8,1
2,8,8,2
electrons
of chlorine ther
od is the result of a change inelectrons
the sub-shell (sub-level) from
experiences
by an
shield the
than
does
the
outer-shell
hangeeffective
in electron shielding. These
have
a
greater
effect
than
the
a
stronger
valence
attraction
d
on
the
electron
arrangements
of
the
elements
(Chapter
12),
the
p
charge
electron
11+
crease in atomic radius.
In period 2, this firstthan
decrease
occurs
17+
Atomic
radius
decrease
the
of +1
from the
2
2
ed
into
four
blocks
of
elements
(Figure
3.7):
nd boron. When it is ionized,nucleus
the berylliumelectron
atom
(1s
2s
)
loses
attraction
experienced
by
in sodium
2
2
1
m
(1s 2selements
2p ) loses a 2p electron (Figure 12.16).
More
energy
is
block
are all in the third electr
m the lower energy 2s orbital in beryllium than
from
the
higher
increases,
the
electrostat
-block elements sodium
chlorine
ugh the 2score
andcharge
2pissub-levels
areto in
the from
sameanthe
shell,
the
energy
electron
more
remove
argon
atom.
Theincreases.
argon atom Th
=
11 –diffi
10 cult
= +1
nucleus
core charge = 17 – 10 = +7
-block
elements
is 2)
alsothat
smaller
the sodium
atom and, therefore,
the
outer
electron
is
l (Chapter
thethan
energy
gap between
shells
and
sub-levels
this electron
the nucleus and making
experiences
closer
to
the
nucleus
and
more
strongly
held.
Figurein
3.2.8
Core2pcharge
can be used to ex
in shell
number.
In
addition,
a
single
electron
the
subblock
elements.
a stronger
The
trend
in
ionic
radii
is
2
trends
within
the
periodic
table.
attraction
y the inner electrons
than
the
2s
electrons
(Figure
12.17).
17+
For the metals (sodium t
across the period. Silicon
4−
(Si
)
ion.
For
the
non-me
2p
3− • atomic radius decrease
(P ) to
the
chloride
ion
(
• electronegativity increa
chlorine
radius
than
the
original
•
first
ionization
energy
core charge = 17 – 10 = +7
•
ionic
radius
decreases
negative ion has a larger
Figure 3.2.8 Core charge can be used to explain
of extra negative charges
trends within the periodic table.
ions have a larger radius
electron
Figure 3.2.9 Trends in properties across a p
The incr
•
atomic
radius
decreases
The atomic radius is basically used to describe the size of an atom.from
The left
than the
electron
in sodium
Periodic Trends 9
Figure 8.3 Periodicity of atomic radii in the periodic table.
Effective nuclear charge in a group
The nuclear charge and the number of inner shielding electrons increase by the
Across a period, the atomic radius decreases from left to right (Figure 8.3).
same amount in a group. This leads to an assumption that the effective nuclear
From acting
one atom
an element
theelements
next across
period:
charge
on theofouter
electronsto
of the
in theasame
group of the
periodic
hardly
varies,
but thisincreases
is a simplification of a more complex situation.
● the table
charge
of the
nucleus
● instance,
the shielding
remains
the same
For
sodium
has 11 protons
and, therefore, a nuclear charge of +11. It has
two
electrons
in
the
first
shell
and
eight
in
the
second
shell.
These
ten
electrons
Across the period the effect of the nuclear charge on the outer electrons
shield
the
outer,
third-shell
electron
very
efficiently
and
the
effective
nuclear
increases and the atomic radii decreases.
charge is close to +1 (+11 − 2 − 8 = +1).
Down a group, the atomic radius increases.
1
Potassium has 19 protons and, therefore, a nuclear charge of +19. The outer, fourthFrom one atom of an element to the next down a group:
shell electron is shielded from the nucleus by two electrons in the first shell, eight
● the charge
of theshell
nucleus
increases
electrons
in the second
and eight
electrons in the third shell, making a total of 18
inner
shielding
electrons.
From sodium to potassium, the number of protons has
● the
shielding
increases
increased by eight and the number of shielding electrons has also increased by eight.
Down
the
group
the
effect
of
the
nuclear
charge
on
the
outer
electrons
Therefore, potassium also has a similar effective nuclear charge close to +1.
decreases and the atomic radii increases.
Li
Na
K
outer electron
shielded by 2
inner electrons
outer electron
shielded by 10
inner electrons
outer electron
shielded by 18
inner electrons
Figure 8.5 Atoms of the Group 1 elements lithium, sodium and potassium.
Test yourself
1 Put the following elements in order of increasing atomic radius. Justify
your answers:
a) Mg, S, Si
b) Mg, K, Al
c) Si, Cl, K
30/03/15
ments
of
Period
3
Periodic Trends 10
Atomic radius
Some
key
of atoms,
such as
and between
ionisation
energy,
are
The
radius
of properties
an atom is found
by measuring
thesize
distance
the
nuclei
of
two
−12
1 pm = 1 × 10 m
periodic,
thatand
is, there
are similar
trends as you go across each period
touching
atoms,
then halving
that distance.
in the Periodic Table.
In general the atomic radius of an atom is determined by the balance between two
opposing factors:
Atomic radii
•
These
tell us effect
aboutbythe
of atoms.
Youshell(s)
cannot
measure
theatomic
the shielding
thesizes
electrons
of the inner
– this
makes the
Although
it
is
not
possible
to
measure
radius
of
an
isolated
atom
because
there
is
no
clear
point
at
which
radius larger. The shielding effect is the result of repulsion between the electrons
an
atomic
radius
for
Ar,
it
is
possible
to
the
cloud
around
it valence
drops to
zero. Instead half the
in electron
the inner shell
anddensity
those in the
outer or
shell.
measure
a
value
for
the
van
der
Waals’
−12
distance between
the
of a pair ofradius
atoms
is
used,
see
Figure
1.
1 pm = 1 × 10 centres
m
of this element.
1
•
the nuclear charge (due to the protons) – this is an attractive force that pulls all the
The atomic radius of an element can differ as it is a general term.
electrons closer to the nucleus. With an increase in nuclear charge, the atomic
It
depends
on
the
type
of
bond
that
it
is
forming
–
covalent,
ionic,
omic radius
decreases
acrosssmaller.
a period is basically the
radius
becomes
metallic,
van der
Waals,
and
so on. The covalent radius is most
egativity
and ionisation
energy
increase:
an increase
measure
of the size of the atom. Figure 2 shows a
oss thecommonly
period but noused
signifias
canta increase
in shielding.
ne have
theof
same
number of
inner shells
of electrons
plot
covalent
radius
against
atomic
number.
Atomic
radius
decreases
across
a
period
Although
it
is
not
possible
to
measure
nt of shielding is similar); however, chlorine has
an
atomic
radius
for
Ar,
it
is
possible
to
Remember
that
effective
nuclear
charge
increases
across
the
period.
This
means
(Even
metals
can
form
covalent
molecules
such
as
Na
in
the
gas
7+ whereas sodium has a nuclear charge of only
2
measure
a
value
for
the
van
der
Waals’
that
the
outer-shell
electrons
ofdo
chlorine
therefore
experience
a greater
the outer
electrons
arenoble
pulled
in
more
strongly
phase.
Since
gases
not in
bond
covalently
with
one attraction
another,to
radius of this element.
um and
the
atomic
radius
is smaller.
the
nucleus
than
does
the outer-shell
sodium.
they
do not
have
covalent
radiielectron
and soofthey
are often left out of
comparisons of atomic sizes.)
cally the
measure
of graph
the size shows
of an ion.that:
The
ncrease
shielding.
•
atomic
radius
is
a
periodic
property
because
it
decreases
across
each
nic
radii
of
positive
ions
are
smaller
than
electrons
period
and
there
is a jump
when
starting the next period
dii,
and
the
ionic
radii
of
negative
ions
are
has
eir
atomic
radii.
•
atoms
get larger down any group.
only
ngly in
.10 (overleaf
) shows
a comparison
of thedecrease
atomic and across a period
Why the
radii
of atoms
or theYou
alkalican
metals.
Each 1+this
ion istrend
smallerby
than
the
explain
looking
at the electronic structures of
formed
(by
loss
of
an
electron).
the
elements
in
a
period,
for
example,
sodium
to
chlorine
in
Period
3,
+
Na as it has one extra shell of electrons – the
as shown in Figure 3. +
on of Na is 2, 8, 1, whereas that of Na is 2, 8. Also,
Atomic
radius
decreases
from
left
to
right
across
period
3
due
to
the
increasing
me nuclear
charge
pulling
in
the
electrons
(11+),
but
As you move from sodium to chlorine you are adding protons to the
r than
attraction
experienced
by the
unt of
electron–electron
repulsion
in outer-shell
Na, as there electrons. These outer-shell electrons are
ns are
nucleus and electrons
to the outer main shell, which is the third shell.
+
electron
is however, as the effective nuclear charge
paredallwith
only
10 in
Na . The
in the
third
electron
shell
of thecloud
atoms;
The
charge
on
the
nucleus
increases
from
+11
to
+17.
This
increased
+
there
are
more
electrons
repelling
thanincreases,
in Na , asthe
electrostatic
attraction
between
the
outer-shell
electrons
and
the
charge
pulls
the
electrons
in
closer
to
the
nucleus.
There
are
no
charge pulling the electrons in.
tomic
and increases. This has the effect of pulling these outer-shell electrons in closer
nucleus
additional
electron
shells
tocan
provide
ive ions
are larger than
their parent
atoms
be seen more shielding. So the size of the
han to
thethe nucleus and making the atom smaller.
es of halogen
atoms with as
their
ionsgo
(1−)
in Figure
atom decreases
you
across
the period.
larger than Cl, because it has more electrons for the
he
and, therefore,
greater repulsion
between
electrons.
Na
Mg
Al
P
S
Cl
Si
atom
, 8. Also,
−
nd 17 protons in the nucleus. Cl also has 17 protons
2,8,6
2,8,7
2,8,5
(11+), but
has 18size
electrons.
The repulsion between 18 electrons
of
s there
en 17 electrons,
so the electron cloud expands as an
atom
oud is
−
d to a Cl atom to make Cl .
epelling
2,8,2
2,8,4
2,8,5
2,8,6
2,8,7
2,8,1
2,8,3
atomic
(covalent)
an be seen
radius
/
nm
Figure
0.156
0.136
0.125
0.117
0.110
0.104
0.099
ns for the
nuclear
ectrons.
16+
17+
11+
12+
13+
14+
15+
charge
7 protons
electrons
▲ Figure 3 The sizes and electronic structures of the elements sodium to chlorine
nds as an
Periodic Trends 11
1
Summary: Across a period, the atomic radius decreases from left to right. From one
atom of an element to the next across a period:
•
the
shielding
remains
the
same
•
the charge of the nucleus increases
Across the period the effect of the nuclear charge on the outer electrons increases
and the atomic radii decreases.
Atomic radius
The radius of an atom is found by measuring the distance between the nuclei of two
touching atoms, and then halving that distance.
Trends across a period
Atomic radius/pm
Even though extra electrons are being added, the atoms get smaller going across
a period from left to right. From lithium to fluorine, the outer electrons are all
2
in the 2nd shell, being screened by the 1s electrons. The increasing number of
protons in the nucleus as you go across the period pulls the electrons in more
tightly. This is slightly offset by the increased repulsion of the electrons from each
other. However, the net effect is a decrease in radius (Figure 2.19). This pattern is
repeated in the 3rd period.
250
200 reason that atomic radius decreases across a period is basically the
The
same reason electronegativity and ionisation energy increase: an increase
150
in nuclear charge across the period but no significant increase in shielding.
Sodium
and chlorine have the same number of inner shells of electrons
100
(and hence the amount of shielding is similar); however, chlorine has
50
sodium
nuclear
of only
a nuclear charge
Li
Be of
B 17+
C whereas
N
O
F
Na has
Mg a Al
Si
Pcharge
S
Cl
K
Ca
2nd
period
3rd
period
11+. This means that the outer electrons are pulled in more strongly in
Figure 2.19 Trends in atomic radius
chlorine than in sodium and the atomic radius is smaller.
Trends in a group
Down a group, the trend is for atomic radii to increase steadily because of the
244 12
Periodic Trends
262
Atomic radius
increases down a group0
10
20
30
0
Down a group from top to bottom, atomic radii increase. In each new period the
270
At
outer-shell electrons enter a new energy level so are located further away from the
nucleus. This has a greater effect than the increasing nuclear charge because of
ion of atomic radii in group 1 ■ Figure 3.20 Bar chart showing the va
shielding by the inner-shell electrons (effective nuclear charge stays nearly the
same.)
300
mber Atomic radius/pm
Atomic radius/pm
1
58
99
200
Cl
100
Br
F
114
Summary: Down
a
group,
the
atomic
radius
increases.
From
one
atom
of
an
element
133
0
to the next down a group:
10
20
30
0
140
At
•
The variation of ionic radius across a period is not a clear-cut trend, as
■
Figure
3.21
Bar
chart
showing
the
va
on ofthe
atomic
radii
in
group
17
shielding
effect
increases
(effective
nuclear
charge
stays
the
same)
•the type of ion changes from one side to the other. Thus positive ions are
formed
on
the
left-hand
side
of
the
period
and
negative
ions
on
the
rightincrease
in
the
number
of
shells
•hand
Moving
down
a
group,
both
the
nuclear
charge
and
the
shiel
side.
the charge of the nucleus increases
outer
electrons
enter
new
shells.
So,
although
the
nucleus
gai
Down
the
group
the
effect
of
the
nuclear
charge
on
the
outer
electrons
decreases
For positive ions there is a decrease in ionic radius as the charge on
and
the
radii increases.
further
away,
butforalso
moreions
effectively
screened
an additio
the
ionatomic
increases,
but
negative
the size increases
as thebycharge
increases (Figure 4.12).
+
2+
:
both
ions
have
the
same
electronic
Let us consider Na and Mg
atomic
radii
decrease
2+
configuration, but Mg has one more proton in the nucleus (Figure
4.13). Because there is the same number of electrons in both ions, the
amount of electron–electron repulsion is the same; however, the higher
2+
nuclear charge in Mg means that the electrons are pulled in more
atomic radii increase
strongly and so the ionic radius is smaller.
Ionic radii for ions of the same charge also increase down a
and 3.7). Ionic radii are the radii for ions in a crystalline ion
Periodic Trends 13
How is atomic radius measured
We know that electrons in atoms are in located in atomic orbitals, which are regions of
space where there is a high probability of finding an electron.
This meansproperties
that the
Periodic
position of the electron is not fixed, so we cannot measure the radius of the atom in
the same way as we measure the radius of a circle. One way of overcoming this
problem and finding the radius of an atom is to measure the distance between the
Figure
7.7
!
nuclei of two closest atoms of the element and dividing it by 2.
Periodicity in the first ionisation energies
The atomic radii of atoms can be obtained
from X-ray analysis and electron density
of the elements.
maps. Using these techniques, it is possible to measure the distance between the
nuclei of atoms and then estimate the radius of individual atoms.
1
The atomic radius of an element can differ as it is a general term. It depends on the
typeArof bond that it is forming – covalent, ionic, metallic, van der Waals, and so on.
group 0
The atomic radii of metals are obtained by measuring the distance between the
Cl nuclei of neighbouring atoms in metal crystals. The atomic radius is simply half of the
inter-nuclear distance. This is sometimes referred to as metallic radius.
S
Ca
K
group 2
group 1
8.3
More
trends
in
the
proper
20
the elements of Period 3
Learning objectives:
Some key properties of atoms, such as
periodic,
that
is,
there
are
similar
trend
➔ Describe the trends in atomic
in
the
Periodic
Table.
atomic
radiiand
of non-metals
radius
first ionisationare obtained from the distance between the nuclei
anyThe
pattern
energyjoined
of theby
elements
in bond. So, for non-metals, the atomic radius is
of and
similar3.atoms
a covalent
ds 2
Atomic
radii
Period
3.
half
of
the
covalent
bond
length.
Because
of
this
link
to
covalent
bonds, the atomic
s decrease
Data
These
tell
us
about
the
sizes
of
atoms.
Y
➔
Explain
these
trends.
radii
of
non-metals
are
sometimes
called
covalent
radii.
?
radius
of
an
isolated
atom
because
ther
Specification reference: 3.2.1
the electron cloud density around it dr
distance
between the centres of a pair o
atomic
radius
The atomic radius of an element can di
It depends on the type of bond that it i
metallic, van der Waals, and so on. The
commonly used as a measure of the siz
plot of covalent radius against atomic n
molecules. Chemists use
distance betweenr the
be defined precisely
number of bonds.
(Even metals can form covalent molecu
non-metallic
molecule
nces between atoms in
phase. Since noble gases do not bond c
Figurecalculated
1 Atomic radii are taken to be
they
do
not
have
covalent
radii
and
so
n-metals▲are
atomic radius
half the distance between the centres of
comparisons
of
atomic
sizes.)
ules (covalent
radii).
a pair of atoms
The graph shows that:
d. Across the period Na
99 nm forHint
chlorine. From
• atomic radius is a periodic property
Periodic Trends 14
How is atomic radius measured
In the simple molecular structures of non-metals, one molecule touches the next and
it is sometimes useful to compare the distances between neighbouring atoms which
are not chemically bonded. This distance in non-metal crystals is called the van der
Waals radius.
1
Consider a group of gaseous argon atoms. When two argon atoms collide with one
another there is very little penetration of their electron cloud densities. Argon does not
form a diatomic species. If argon is frozen in the solid phase the atoms would touch
3.2 PE
each other but would not be chemically bonded. In this case the distance between
the argon atoms could be measured and hence non-bonding atomic radius. The
atoms could be measured and hence Rnb could be found (figure 2). The
non-bonding atomic radius is often termed the van der Waals’ radius.
non-bonding atomic radius is often termed the van der Waals’ radius
d = 2Rnb
▲ Figure 2 Atoms of argon in the solid phase. The atoms are touching but not chemically
bonded. The non-bonding atomic radius of argon Rnb is 188 pm (d = 376 pm)
Section 9 of the Data booklet provides data for the covalent atomic radii
of the elements. The general term “atomic radius” is used to represent
the mean bonding atomic radius obtained from experimental data
over a wide range of elements and compounds. Note that the bonding
atomic radius is always smaller than the non-bonding atomic radius. Th
approximate bond length between two elements can also be estimated
from their atomic radii.
For example, for the interhalogen compound BrF:
atomic radius of bromine = 117 pm
atomic radius of fluorine = 60 pm
bond length of Br-F = 177 pm
Compare this with the experimental bond length of Br-F in the gas
F
9
h
chemical properties across periods and down grou
Periodic Trends 15
First ionization energy
The Data booklet provides data for the atomic radii of the elements. The general term
The
first
ionization
energy
is
the
minimum
energy
req
“atomic radius” is used to represent the mean bonding atomic radius obtained from
from one mole of gaseous atoms (under standard therm
experimental data. Note that the bonding atomic radius (metallic and covalent) is
+
−
(g) + e
In general:
always smaller than the non-bonding
atomic radiusX(g)
(van →
derX
Waals).
For example, the first ionization energy of hydrogen is g
0.20
H(g) →
Na
+
H (g)
+
−
e
∆H =
−1
+1310 kJ mol
Ar
The amount of energy required to carry out this proces
1310
kilojoules.
Mg
Period 3
Atoms of each element have different values of first
atomic radius/nm
Al
1
Electronegativity
Si
The electronegativityPof an atom
is
the
ability
or
powe
S
Cl
shared pairs of electrons to itself. The greater the electr
ability to attract shared pairs of electrons to itself.
Electronegativity values are usually based on the Paul
the most electronegative atom. The least electronegative
value of 0.7. The values for all the other elements lie betw
electronegativity values are pure numbers with no units.
0.10
0
11
12
13
14in the
15 properties
16
17 of the
18
■ Trends
elem
group 17
proton number
Phosphorus, sulfur and chlorine all form small covalent molecules, P4, S8 and
Trends
in
atomic
and
ionic
radii
Cl2 respectively. When these substances melt, it is only necessary to break weak
a
intermolecular
bonds
and notAt
strong
melting
points radius is d
the interatomic
right of theattractions.
periodicThe
table,
the atomic
10_04 Cam/Chem
AS&A2
decrease
phosphorus
> covalently
chlorine (seebonded
Figure 10.5),
the(Figure 3.19). Fo
r in the order sulfur >nuclei
of
two
atoms
Barking Dog Art
intermolecular bonds becoming
weaker
as
the
molecules
become
smaller.
molecule (the distance between two chlorine nuclei) is
r
The melting point of argon is low, as the attraction between the argon atoms is
−1
of chorine is ½ × 199 = 99 pm (1 picometre [pm] = 10
very small.
metallic radius
b
r
r
Si
covalent radius
c
Mg
r
left of the periodic table, the atomic radius is that of th
radius). For the noble gases the atomic radius is that of
radius).
In general the atomic radius of an atom is determin
factors:
■ the shielding effect by the electrons of the inner sh
Cr shielding effect is the result of repulsion
larger.V The
Ti those in the outer or valence shell.
and
Fe Co Ni
Sc
■ the nuclear charge (due to the protons) – this is an
Cu
Ge
Mn
closer to the nucleus. With an increase in nuclear c
Ca
However, when moving downZna group in the periodic t
radius
as
the
nuclear
charge
increases
(Tables
3.4
and
3
S
Se
Na
P
K
result
of
two
factors:
Br
Cl
As
Ga
Ar
Kr
van der Waals’ radius
■ the increase in the number of complete electron sh
Ne
(for group 18)
and
the
nucleus
10 11
12 13 14
15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36
■ Figure
3.19
proton number
■ the increase in the shielding effect of the outer elec
Atomic radius
Al
Electrical conductivity
Sodium, magnesium and aluminium are metals. They have delocalised electrons that
are free to move in the lattice of cations (see section 4.11). Silicon is a semiconductor.
The other elements in the third period form covalent bonds with no free electrons,
negative ion is always larger than its neu
Periodic Trends in negative
ionic radiusion, an atom must gain one or more
The radii of cations andthe
anions
vary from
the parent
atoms from
they of
are attracti
same
atomic
number,
sowhich
forces
formed in the following is
way.
extra repulsion due to the increased numbe
causes
the
ion
to
expand,
moving
the
elect
The radii of cations are smaller than those of their parent atoms. The reason
again,
between
for this is that there are once
more protons
thanthere
electronsisinathebalance
cation so the
valence the f
electrons are more strongly
attracted to the nucleus.
repulsion.
Periodic Trends 16A
1
Ion
Atomic number
Li+
Na
Ionic radius/pm
3
Ion
+
Na
Atomic number
Cl
Ionic radius/pm
9
133
F–
68
Electron affinity (EA)
–
+
Cl
11
98
Na
If an atom loses all its outer electrons, the radius of the resulting ion17is much smaller181
Br–
35
148
I–
353
219
At–
85
No data
+
133
K
than the atomic19radius. This is because:
196
•
can be represented by the eq
there
are
fewer
electrons
in
the
positive
ion
than
in
the
atom,
so
the
electron–
•
−
−
A(g) + e → A (g)
Rb+
37
Cs
55
167
Fr+
87
No data
there
is one fewer shell
of
electrons
+
Electron affinity
■
■ Table 3.7 The variation of ionic radii in group 17
electron
repulsion
is
less,
causing
a
further
reduction
in
the
radius.
Table 3.6 The variation of ionic radii in group 1
Li
Li
+
The negatively charged electron being adde
charged nucleus. There is a force of attracti
energy is released when the two are brought c
Na
Na
+
K
K
+
The first electron affinity of oxygen is −14
per mole, for the process:
−
O(g)
+
e
Trends in first ionization energy
→
−
O (g)
On moving down a group, the atomic radius increases as additional electron shells are added. Th
The
second
electron
affi
nity
can
be
represe
causes the shielding effect to increase. The further the outer or valence shell is from the nucleus
the smaller the attractive force exerted by the protons in the nucleus. Hence, the more easily an
−
−
2−
e the→ionization
A (g)
Aand(g)
outer electron can be removed
the +
lower
energy. So, within each group, the
first ionization energies decrease down the group. This is shown in Table 3.8 and Figure 3.24.
494
418
402
376
on of first ionization
First ionization energy/kJ mol
519
–1
mber
First ionization
energy/kJ mol –1
The second
electron
affi
nity
of
an
element
is
a
Li
Na
500
energy is required to add an electron to an
K
Rb negative ion. The
electron
is
repelled
by
the
400
Cs
bring the ion and the electron together.
300
200
100
0
0
10
20
30
40
Atomic number
50
60
Periodic Trends 17
charge
increases
from
11
to
12
to
13.
The
force
of
attraction
between
+
2+
3+
For ions with the same electron configuration (e.g. Na , Mg and Al ), the ion
f the
aluminium
ion
(13
protons)
and
its
ten
electrons
is,
therefore,
with the greatest charge will have the smallest radius.
hat between the nucleus of the other ions and their ten electrons. This
3+ ion
Theto
nuclear
increases radius.
from 11 to 12 to 13. The force of attraction between the
havecharge
the smallest
nucleus of the aluminium ion (13 protons) and its ten electrons is, therefore, greater
than that between the nucleus of the other ions and their ten electrons. This causes
the Al3+ ion to have the smallest radius.
egative ions
on is always larger than its neutral atom (Figure 2.20). To form a
The
radii
of
anions
are
larger
than
those
of
their
parent
atoms.
A
negative
ion
is
an atom must gain one or more electrons. The atom and the ion have
always
larger
than
its
neutral
atom.
To
form
a
negative
ion,
an
atom
must
gain
one
or
mic number, so forces of attraction remain the same. However, there
more electrons. The atom and the ion have the same atomic number, so forces of
sion due to the increased number of electrons in the same shell. This
attraction remain the same. However, there is extra repulsion due to the increased
n to expand, moving the electrons further from the nucleus until,
number of electrons in the same shell. This causes the ion to expand, moving the
here is a balance between the forces of attraction and the forces of
1
electrons further from the nucleus.
+
Na
–
Cl
Cl
n affinity (EA)
ity can be represented by the equation:
→
−
A (g)
ly charged electron being added is brought towards the positively
eus. There is a force of attraction between the two and, therefore,
ased when the two are brought closer together.
ctron affinity of oxygen is
the process:
−1
−142 kJ mol .
This is the energy change,
The variation of ion
−
→ O (g)
the type of ion change
formed on the left-ha
electron affinity can be represented by the equation:
hand side.
− → A 2−(g)
Table 2.6 Atomic and ionic radii
For positive ions th
lectron affinity of
an
element
is
always
positive
(endothermic)
because
Group 1 Group 2 Group 3 Group 4 Group
5 ion
Groupincreases,
6 Group 7 but f
the
uiredAtomic
to add anNaelectron
to
an
already
negative
ion.
The
incoming
Mg
Al
C
N increases
O
F
(Figure
4.12
pelledRadius/pm
by the negative
ion.
Therefore,
energy
has
to
be
supplied
to
+
186
160
143
70
65
60
50
Let us consider Na
and the electron together.
2−
Mg2+
Al3+
N3− confiOguration,
F− but Mg
Ionic
Na+
Radius/pm
95
65
50
171
140
136
4.13). Because there is
amount
of
electron–el
The nuclear charge increases from 11 to 12 to 13. The force of attraction between Io
the nucleus of the aluminium ion (13 protons) and its tennuclear
electrons is,
therefore,
charge
in Mg
s
Periodic Trends 18
Increasing electronegativity and first ionization energy
1
Increasing electronegativity and first ionization energy
1310
2370
H
He
2.1
519
799
900
1090
1400
1310
1680
–
(O +844)
2080
Li
Be
1.0
494
1.5
736
2.0
577
2.5
786
Na
Mg
Al
Si
0.9
418
1.2
590
1.5
577
1.8
762
2.1
966
2.5
941
3.0
1140
B
632
661
648
Ca
Sc
Ti
0.8
402
1.0
548
1.3
636
1.5
669
Rb
Sr
0.8
376
1.0
502
Cs
0.7
181
K
653
762
757
736
745
908
N
3.0
1060
P
Ne
F
O
3.5
4.0
1000
1260
–
(S +532)
Mn
Fe
Co
Ni
Cu
Zn
Ga
Ge
As
Se
Br
1.6
653
1.6
694
1.5
699
1.8
724
1.8
745
1.8
803
1.9
732
1.6
866
1.6
556
1.8
707
2.0
833
2.4
870
2.8
1010
Zr
Nb
Mo
Tc
Ru
Rh
Pd
Ag
Cd
In
Sn
Sb
Te
1.2
540
1.4
531
1.6
760
1.8
770
1.9
762
2.2
841
2.2
887
2.2
866
1.9
891
1.7
1010
1.7
590
1.8
716
1.9
703
2.1
812
2.5
920
Ba
La
Hf
Ta
W
Re
Os
Ir
Pt
Au
Hg
Tl
Pb
Bi
Po
At
0.9
1.1
1.3
1.5
1.7
1.9
2.2
2.2
2.2
2.4
1.9
1.8
1.8
1.9
2.0
2.2
510
1520
Ar
Cl
S
Cr
Y
V
716
C
1350
Kr
1170
Xe
I
1040
Rn
669
Fr
Ra
Ac
0.7
0.9
1.1
First ionization
–1
energy (kJ mol )
Element
Electronegativity
Figure 3.2.1 Electronegativity and first ionization energy values, showing trends within the periodic table. A similar table can be found in the IB Data Booklet © IBO 2007.
Decreasing atomic and ionic radii
30
H
He
Increasing atomic and ionic radii
154 (1–)
152
Li
68 (1+)
186
Na
98 (1+)
231
K
133 (1+)
244
Rb
148 (1+)
262
Cs
167 (1+)
270
Fr
88
112
Be
B
30 (2+)
16 (3+)
143
160
Mg
65 (2+)
160
Ca
94 (2+)
110 (2+)
217
Ba
34 (2+)
220
Ra
131
125
129
126
125
124
128
133
260 (4–)
N
171 (3–)
117
66
O
146 (2–)
110
58
F
133 (1–)
104
Ne
99
Si
P
S
Cl
Ar
42 (4+)
212 (3–) 190 (2–) 181 (1–)
271 (4–)
141
122
121
117
114
Sc
Ti
V
Cr
Mn
Fe
Co
Ni
Cu
Zn
Ga
Ge
As
Se
Br
Kr
Y
Zr
Nb
Mo
Tc
Ru
Rh
Pd
Ag
Cd
In
Sn
Sb
Te
I
Xe
Tl
Pb
Bi
Po
At
Rn
81 (3+)
90 (2+)
88 (2+)
63 (3+)
80 (2+)
76 (2+)
74 (2+)
72 (2+)
96 (1+)
74 (2+)
62 (3+)
53 (4+)
222 (3–) 202 (2–) 196 (1–)
68 (4+)
59 (5+)
60 (4+)
64 (3+)
63 (3+)
69 (2+)
272 (4–)
180
157
141
136
135
133
134
138
144
149
166
162
141
137
133
215
Sr
146
C
70
Al
45 (3+)
197
77
93 (3+)
188
La
115 (3+)
80 (4+)
157
Hf
81 (4+)
70 (5+)
143
Ta
73 (5+)
68 (4+)
137
W
68 (4+)
137
Re
65 (4+)
86 (2+)
134
Os
67 (4+)
135
Ir
66 (4+)
126 (1+)
138
Pt
144
Au
137 (1+)
85 (3+)
97 (2+)
152
Hg
127 (1+)
110 (2+)
81 (3+)
112 (2+) 245 (3–) 222 (2–) 219 (1–)
71 (4+)
171
175
170
140
140
95 (3+)
120 (2+)
84 (4+)
120 (3+)
200
Ac
Atomic
radius
(10–12 m)
Element
Ionic
radius
(10–12 m)
Figure 3.2.2 Atomic and ionic radii values showing trends within the periodic table. A similar table can be found in the IB Data Booklet © IBO 2007.
Periodic Trends 19
Skill Check 1
ne
>, =, <
First particle
a
Second particle
−
Chlorine atom (Cl)
Chloride ion (Cl )
3+
b
Aluminium ion (Al )
Aluminium atom (Al)
c
d
Calcium atom (Ca)
Sulfur atom (S)
e
f
+
−
Sodium ion (Na )
Fluoride ion (F )
2+
Magnesium ion (Mg )
Calcium ion (Ca )
+
2−
Sulfide ion (S )
2+
1
Potassium ion (K )
EMENTS AND THEIR OXIDES
iodicity within groups and periods,
ents and some of their compounds.
n arrangement of the elements.
group
eactive nature. Their tendency to
that they must be stored under
in secondary school laboratories
olent reaction with water.
er increases down group 1. While
ter and reacts slowly, producing
ently, whizzing around on the
gas in what is sometimes described
3.3.1
Discuss the similarities and
differences in the chemical
properties of elements in the
same group. © IBO 2007
g its progress on the surface of the
.1). The white smoke in figure 3.3.1
. Potassium burns spontaneously
olution when they react with water.
en in figure 3.3.1. Phenolphthalein
e the reaction with sodium and has
ydroxide.
by the decrease in electrostatic
and the positive nucleus of the
Figure 3.3.1 Sodium will burn with a yellow flame
if its motion on the surface of water is stilled.
factor is Trends
the increase
Periodic
20 in the atomic radius, making the outermost electro
from the nucleus and so easier to remove.
Periodic trends in ionisation energy
The general trend across a period is for the first ionisation energy to
The
first
ionisation
energies
of
the
elements
show
periodicity.
The
pattern
from
However, there are a number of slight variations from this trend.
1st ionisation energy/kJ mol–1
lithium to neon is repeated exactly with the elements sodium to argon. Apart from the
Figure
2.6
shows
that
there
are
maxima
at
each
noble
gas
and
minima
at
e
insertion of the d-block elements, this pattern is seen again from potassium to
1
metal.
There
are
dips
after
the
second
and
fi
fth
elements
in
both
period
rubidium.
2400
He
2200the radii of atoms
Ne
Why
increase down a group
2000 down a group in the Periodic Table, the atoms of each ele
Going
have
one
extra
complete
main
shell
of
electrons
compared
with
1800
F
one1600
before. So, for example, inArGroup 1 the outer electron in sod
is in main shell N3, whereas in potassium it is in main shell
4.
So
Kr
1400
down the group, the outer electron main shell is further from th
O
Cl
1200 H
Br
nucleus and the atomic radii
increase.
P
1
1000
Be
C
Zn
S
Se
Si
First ionisation
energy
B
Mg
800
As
Ge
600
The first ionisation energy
is
the
energy
required
to
convert
a
m
Al
Li
Ga
Na
400
Rb
isolated gaseous atoms into a mole
of
singly
positively
charged
g
K
200
ions,
that is, to remove one electron from each atom.
0
E(g)
➝
0
+
–
E
(g)
+
e
5
10
(g)15
where
E
stands
for
any
elemen
25
30
35
40
20
Atomic number
The
first
ionisation
energies
also
have
periodic
patterns.
These
ar
Figure 2.6 Variation of the first ionisation energy with atomic number
shown in Figure 4.
2500
He
first ionisation energy / kJ mol
–1
Ne
2000
F
1500
Ar
N
02_Edexcel_GF_Chem_009-036.indd 17
O
H
1000
Cl
Be
P
C
Mg
B
500
0
S
Ca
Al
Li
0
Si
Na
5
10
atomic number
▲ Figure 4 The periodicity of first ionisation energies
K
15
20
01
Periodic Trends 21
of sodium,
magnesium
aluminium
isoelectronic
species
(Table to
3.15).
The nuclear
The Ions
general
trend
across and
a period
is forare
the
first ionisation
energy
increase.
2
2
6
2p ,
In Figure
notice
the ion
small
between
phosphorus
(1s , pulls
2s ,all
charge
increases2,
from
the sodium
to thedrop
aluminium
ion. The
higher nuclear charge
2
3
2
2
6
2
4
the
electron
shells
closer
to
the
nucleus.
Hence,
the
ionic
radii
decrease.
3s
,
3p
)
and
sulfur
(1s
,
2s
,
2p
,
3s
,
3p
).
In phosphorus,
As the effective nuclear charge increases from left to right
across a periodeach of the
Similarly, the nuclear charge increases from the phosphide ion to the chloride ion. The
three
3p
orbitals
contains
just
one
electron,
while
in
sulfur,
one
of
the
(nuclear
charge
increases
and
the
shielding
effectively
remains
the
same),
the
8.3
More nuclear
trends incharge
the properties
of the
elements
Period
3
higher
causes the
electron
shellsofto
be pulled
closer to the nucleus. Again, the
3p
orbitals
must
contain
two
electrons.
The
repulsion
between
these
ionic
radii
decrease
(Table
3.16).
valence electrons are pulled closer to the nucleus, so the attraction between the
•
paired
electrons
makes
it
easier
to
remove
one
of
them,
despite
the
+
2+
3+
3–
2–
–
electrons and the nucleus
increases.
This
makes
it
more
difficult
to
remove
an
Species
Na
Mg
Al
Species
P
S
Cl
The first ionisation energy generally increases across a period
increase
in
nuclear
charge,
see
Figure
4.
Nuclear
charge
+11
+12
+13
Nuclear
charge
+15 and lithium,
+16
(see
Figure
4),
alkali
metals
like sodium, Na,
Li,+17
have the
electron from the atom.
Number of electrons
lowest values
and Number
the nobleofgases
(helium,18He, neon,18Ne, and18
argon,
10
10
electrons
10
•
Ar) have the highest values.
Ionic
radius/pm
98 across
65 a period
45
Ionic radius/pm
190the
Atomic
radii decrease
– because
the distance212
between
181
2p ionisation energy decreases going3pdown any group. The
The first
the nucleus
it3sGroup
becomes
more
difficult
■ valence
Table 3.15 electrons
Atomic data
for sodium,
magnesium
Table
3.16 Atomic
data for
phosphide,
sulfide
2s and
1s
trends
fordecreases,
Group■1 and
0 are shown
dotted
in to
redremove
and green,
and aluminium ions
andgraph.
chloride ions
respectively
on
the
an electron from the atom.
1
The large increase in size from theYou
aluminium
ion these
to thepatterns
phosphide
ion is due
to the presence
can explain
by looking
at electronic
arrangements
of an additional
shell.inThis
causes5).
a large increase in the shielding effect and as a result
orbitalselectron
(sub-shells)
phosphorus
(Figure
However, there are a number of slight variations from this trend.
easier
the ionic radius increases.
Outer electrons are harder to remove as nuclear charge
2pincreases
Trends in 1s
first ionization
energy
2s
3p
to lose
3s
The first ionization energies of the elements in period 3 are listed in Table 3.17. The general
11+
12+
13+
14+
15+
16+
17+
18+
trend is an increase in first ionization energy across the periodic table. When moving across
a period from
left to
right the nuclear
orbitals
(sub-shells)
in sulfurcharge increases but the shielding effect only increases
2 3p2
2 3p3
2 3p4
2 3p5
2 3p6
[Ne]3s1
[Ne]3s2 3p1
[Ne]3s2
[Ne]3s
[Ne]3s
[Ne]3s
[Ne]3s
[Ne]3s
slightly (since electrons enter the same shell). Consequently, the electron shells are pulled
closer to
the nucleus
and as sodium
a resultto first
▲progressively
Figure 5 The electronic
structures
of the elements
argonionization energies increase.
▲ Figure 4 Electron arrangements of phosphorus and sulfur
Element
Sodium
Magnesium
Aluminium
Phosphorus
Why
the first
ionisationSilicon
energy
increases Sulfur
across aChlorine
period
As you
a period from
right, the number
of1260
protons
736go across 577
786 left to 1060
1000
Both these cases, whichin go
against
the expected
trend,
aresame
evidence
the nucleus
increases
but the electrons
enter the
main shell,
see Figure
increased
charge on theThese
nucleus were
means that it gets
that confirms the existence
of 5.s-The
and
p-sub-shells.
increasingly
to remove
anthere
electron.
However, the increase in first ionization
energydifficult
is not uniform
and
are two decreases –
First ionization energy/
kJ mol –1
494
Down
a group
predicted
by
quantum
theory
and
the
Schrödinger
equation.
between
magnesium
aluminium
phosphorusenergy
and sulfur.
decreases can
The
trend
down aand
group
is for and
thebetween
first ionisation
toThese
decrease.
Down
a
group,
the
ionisation
energy
decreases.
Why
the
first
ionisation
energy decreases going down a group
only
be
explained
by
reference
to
sub-shells
and
orbitals.
Hint
Although the number of protons
increases,
so
does
the
number
of down
shielding
The
number
of
filled
inner
shells
increases
the
group.
This
The first ionization energy of aluminium is lower
than
that
of
magnesium,
even
though
–1
Trends
in
ionisation
energies
down
a
group
in
the
results
in
an
increase
in
shielding.
Also,
the
electron
to
be
removed
is
Filled
inner
electron
shells
are
said
)
of
the
elements
in
Group
1
of
the
The
first
ionisation
energies
(kJ
mol
electrons,
keeping
theradius.
effective
nuclearin
charge
constant.
Thefrom
main
factor is
aluminiumhence
has a smaller
atomic
The decrease
first ionization
energy
magnesium
at
an6 increasing
distance from the nucleus and is therefore held less
to shield
electrons
in the outer
2
2
6
2
2
2
2
1
periodic
are
shown
in
Table
8.6.
(1sPeriodic
2s 2p table
3s in
) tothe
aluminium
(1s
2s
2p
3s
3p
)
occurs
because
the
electrons
in
the
filled
3s
Table
the
increase
atomic
radius,
making
the
outermost
electron
further
from
the
strongly. Thus the outer electrons get easier to remove going down a
shell from the nuclear charge.
orbital are more effective at shielding
the electron
in the
orbitalaway
thanfrom
they the
are nucleus.
at shielding each
group
because they
are3pfurther
nucleus
and
so
easier
to
remove.
Table
8.6
First
ionisation
energies
ofis
the
1 elements.
Figure
5
shows
that
thereto
a Group
general
decrease
in tofirst
ionisation
other. Therefore less energy is
needed
remove
a single
3p
electron than
remove
a paired 3s
Why there is a drop in ionisation energy from one period
electron.
energy
going
down
Group
2,
and
the
same
pattern
is
seen
in
other
Li
Na
K
Rb
Cs
Element
2
2
6
2
2
1
1
to
the
next
The first ionization energy of sulfur (1s 2s 2p 3s 3p 3p 3p ) is less than that of phosphorus
groups.
This
is
because
the
outer
electron
is
in
a
main
shell
that
gets
Moving
from
neon
in
Period
0
(far
right)
with
electron
arrangement
2
2
6
2
1
1
1
4
–1 energy is required to remove an electron from the 3p orbitals
(1s
2s
2p
3s
3p
3p
3p
)
because
less
496
419
403
376
First ionisation energy/kJ mol2,8 to520
sodium,
2,8,1
(Period 1, far left) there is a sharp drop in the
further
from
the
nucleus
in
each
case.
of sulfur than from the half-filled 3p orbitals of phosphorus. The presence of a spin pair of electrons
first ionisation energy. This is because at sodium a new main shell
results in greater electron repulsion compared
to
two
unpaired
electrons
in
separate
orbitals.
starts and so there is an increase in atomic radius, the outer electron is
Factor 1 Atomic radius further
– increases
down
the
group
and
the
outer
from the nucleus, less strongly attracted and easier to remove.
■ Trends electron
in electronegativity
values
is
further
from
the
nucleus.
This
tends
to
decrease
900
Be
The electronegativities
of
the
elements
in
period
3
are
listed
in
Table
3.18.
The
general
trend
the ionisation energy
as
we
go
down
the
group.
Summary questions
–1
is an increase
in
first
ionization
energy
across
the
periodic
table.
When
moving
across
a
period
850
from left to
the nuclear
charge increases
but thedown
shielding
effect
only increases
slightly
Factor
2 right
Nuclear
charge
–
increases
the
group.
This
tends
toacross a
1 What happens to the size of atoms as you go from left to right
(since electrons enter the same shell). Consequently,
the
electron
shells
are
pulled
progressively
period? Choose
fromas
increase,
decrease,
no change.
800 increase the ionisation
energy
we
go
down
the
group.
closer to the nucleus and as a result electronegativity values increase.
2 What happens to the first ionisation energy as you go from left to right
Mg
750
across a period? Choose from increase, decrease, no change.
first ionisation energy / kJ mol
nts
3.2 Periodic trends 101
Factor
3
Shielding
effect
–
increases
down
the
group.
This
reduces
Element
Sodium Magnesium Aluminium
Silicon
Phosphorus
Sulfur
Chlorine
3 nucleus.
What happens
to the
nuclearto
charge
of the atomsthe
as you go left to right
the
effect
of
the
This
tends
decrease
Electronegativity
0.9
1.3
1.6
1.9
2.2
2.6
3.2
700
across a period?
ionisation energy4 as
we
go
down
the
group.
Why do the noble gases have the highest first ionisation energy of all
Generally, the electronegativity values of chemical elements increase across a period and
650
the
elements
in
their
period?
decrease down
a group
(Figure 3.28).
This observation
can be used
to compare
the relative
Factors
1 and
3 outweigh
factor
2 as ionisation
energy
decreases
down
electronegativity values of two elements
in the periodic table. To do this, find the positions of
Ca
a group.
600
154
Sr
Across550
a period
Across a period, ionisation energy increases. This is illustrated in
Ba Table 8.7.
500
0 ionisation
10 energies20
30in Period 2.
40
50
60
Table 8.7 First
of elements
B
C
N
atomic number
O
F
▲ Figure
5
The
first
ionisation
energies
of
the
elements
of
Group
2
801
1086
1402
1314
1681
Ne
2081
18/0
8.4 A closer look at ionis
Periodic Trends 22
Why there is a drop in ionisation energy from one period to the next
Moving from neon in Period 2 Group 18 (far right) with electron arrangement 2,8 to
sodium, 2,8,1 (Period 3 Group 1 far left) there is a sharp drop in the first ionisation
This
chapter
revisits
the
trends
in
ionisation
energies
first
dealt
wi
energy. This is because at sodium a new main shell starts and so there is an increase
in
Topicradius,
1.6, the
in outer
the electron
context
of periodicity.
Theless
graph
of attracted
first ionisa
in atomic
is further
from the nucleus,
strongly
energy
and easieragainst
to remove.atomic number across a period is not smooth. Figu
below shows the plot for Period 3.
A closer look at ionisation energies of Period 3
first ionisation energy / kJ mol
–1
1750
Group 0
Ar
1
1500
1250
Group 5
P
1000
500
Group 6
S
Group 2
Mg
750
Group 1
Na
Group 7
Cl
Group 4
Si
Group 3
Al
250
0
10
11
12
13
14
15
atomic number
16
17
18
▲ Figure 1 Graph of first ionisation energy against atomic number for
Theelements
graph of first
the
of ionisation
Period 3 energy against atomic number across a period is not
smooth. It shows that:
It
shows
that:
the first ionisation energy actually drops between Group 2 and Group 13, so that
has aionisation
lower ionisation
energy actually
than magnesium
•aluminium
the first
energy
drops between Group 2 and
Group 3, so that aluminium has a lower ionisation energy than
the ionisation energy drops again slightly between Group 15 (phosphorus) and
magnesium
Group 16 (sulfur).
• the ionisation energy drops again slightly between Group 5
Similar
patterns
occur
in
other
periods.
You
can
explain
this
if
you
look
at
the
electron
(phosphorus) and Group 6 (sulfur).
arrangements of these elements.
Similar patterns occur in other periods. You can explain this if you
look at the electron arrangements of these elements.
The drop in first ionisation energy between
Groups 2 and 3
For the first ionisation energy:
• magnesium,
• aluminium,
2
1s
2
1s
2
2s
2
2s
6
2p
6
2p
2
3s ,
2
3s
loses a 3s electron
1
3p ,
loses the 3p electron.
Specification reference: 3.2.1
Periodic Trends 23
3p
3s
energy
The drop in first ionisation energy between Groups 2 and 3
actually
betweenenergy:
Group 2 and
For thedrops
first ionisation
m has Hint
a lower ionisation energy than
2s
magnesium, 1s2 2s2 2p6 3s2 , loses a 3s electron
Ionisation
energies
are
sometimes
ps again slightly between Group 5
1s
2
2
6
2
1
aluminium,
1s
2s
2p
3s
3p
,
loses
the
3p
electron
called
ionisation
enthalpies.
6 (sulfur).
2p
•
•
2 2s2 2p6 3s2
magnesium
1s
periods.
You can
explaininthis
if you
The p-electron
is already
a higher
energy level than the s-electron, so it takes less
er
ments
of these
elements.
energy
to remove
it
ation energy between
complete removal
1
1st IE
y:
1st IE
3p
3p
loses a 3s
3s electron
1
3p ,
energy
energy
2,
2
complete removal
loses the 2p
3p electron.
2s
energy
3s
2p
2s
higher
level
than
the
s-electron,
rgies
ove it, see Figure 2.
1s
1s
2
2
6
2
2
2
6
2
1
The drop
in
first
ionisation
energy
between
magnesium 1s 2s 2p 3s
aluminium 1s 2s 2p 3s 3p
Groups 5 and 6
▲ Figure 2 The first ionisation
energies
of
magnesium
and
aluminium
The
drop
in
first
ionisation
energy
between
Groups
15
and
16
An electron in a pair will be easier to remove that one in an orbital
complete removal
(not
to
scale)
An
in abecause
pair will be
to remove
thatrepelled
one in an by
orbital
its own
because
onelectron
its own
it easier
is already
being
theonother
electron.
1st
shownbeing
in IEFigure
itAs
is already
repelled3:
by the other electron.
•
2
1s
2
2s
6
2p
2
3s
155
3
3p ,
• phosphorus, 1s 2s2 2p6 3s2 3p3, as no paired
has electrons
no paired
electrons
in
a
in a p-orbital because
p-orbital
because
each
p-electron
is
in
a
different
orbital
3sp-electron is in a different orbital
each
energy
3p 2
phosphorus,
2
1s
2
2s
6
2p
2
3s
4
3p ,
• sulfur, 2p
has two of its p-electrons paired in a
2 2s2 2p6 3s2 3p4, has two of its p-electrons paired in a p-orbital so one of
sulfur,
1s
p-orbital
so
one
of
these
will
be
easier
to
remove
than
an
unpaired
2s
these
beto
easier
remove than
unpaired
due to in
thethe
repulsion
the
onewill
due
thetorepulsion
ofan
the
other one
electron
sameoforbital.
•
other electron in the same orbital.
1s
2s
1s aluminium
2
1s
2
2s
6
2p
2p
2
3s
1
3p
3s
3p
▲ Figure 2 The first ionisation
energies of magnesium and aluminium
orbitals
(sub-shells)
in
phosphorus
(not to scale)
1s
2s
2p
155
3s
3p
easier
to lose
orbitals (sub-shells) in sulfur
▲ Figure 3 Electron arrangements of phosphorus and sulfur
Successive ionisation energies
If you remove the electrons from an atom one at a time, each one is
harder to remove than the one before. Figure 4 is a graph of ionisation
hence
there
is
a
greater
attraction
(higher
effective
nuclear
charge)
fo
Periodic Trends 24
Ionization energy
He
filling up the
p sub-shell
H
Ne
filling up the
p sub-shell Ar
half-filled
p sub-shell
Li
s sub-shell
s sub-shell
s sub-shell
1
erstand by the term atomic orbital?
show the shape of:
half-filled
p sub-shell
Na
[1
K
Atomic number
2
2
6
2
6
8
2
[1
[1
electronic configuration 1s 2s 2p 3s 3p 3d 4s .
TheCheck
‘drops’
groups
2 and to?
13 occur because an electron is remove[1
theSkill
Periodic
Table2between
does element
X belong
mumThe
number
of energy
electrons
inofaseveral
dansub-shell.
higher
in
than
s orbital
is easier
remove.
There is[1
1st ionisation
energies
elementsand
with hence
consecutive
atomicto
numbers
are shown
2+
ms aninion
of type
X .Thebetween
‘drops’
groups
15 and
16elements.
occurs because there are pa
the The
graph
below.
letters are not
the symbols
of the
2
[1
ectronic
confienergy
gurationisfor
this ion using
1s notation.
so
less
required
to
remove
one.
This
is
due
to
the
increased
a. Which of the elements A to I belong to Group I in the Periodic Table? Explain your answer.
ol for theThere
sub-shell
begins
afterend
the of
3d each
and 4speriod
are completely
is which
a large
droptoatfillthe
becausefull.
an electro[1
2 2s2 2p6 3s2?Total =
b.which
Which ofresults
the elements
A
to
I
could
have
the
electronic
configuration
1s
in a large increase in shielding. Hence there is less attra
nucleus.
c.the
Explain
the rise in 1st ionisation energy between element E and element G.
ergies of several elements with consecutive atomic numbers are shown in the graph
For
the
sand
p-blocks
the
increase
in
nuclear
charge
across
a
per
Estimate
the of
1stthe
ionisation
energy of element J.
notd.the
symbols
elements.
on the outer shell (valence) electrons because the inner shielding onl
–1
First ionisation energy / kJ mol
2000
1600
1200
1800
400
0
A
B
C
D E F
Element
G H
I
J
ents A to I belong to Group I in the Periodic Table? Explain your answer.
2
2
6
2
ents A to I could have the electronic configuration 1s 2s 2p 3s ?
1st ionisation energy between element E and element G.
nisation energy of element J.
[3
[1
[4
[2
80
80
−
Periodic Trends B
25 ( Br− Br)
79
81
−
C ( Br− Br)
Skill Check 3
79
81
+
( Br−
The first ionisation of D
elements
sodiumBr)
to argon is shown below.
d?
(1)
(2)
B Mark on the graph where the value for potassium would be.
The
firstforionisation
of elements
tolarger than
C Explain3why
the value
the second ionisation
of sodiumsodium
is very much
that of its first
ionisation.
argon
is shown below.
First ionisation energy
(1)
(1)
d)
State
and
outline
one
modern
use
of
mass
A Explain why the general trend from sodium to argon is upwards but why the value for
(3)
sulfur is less thanspectrometry.
that for phosphorus.
(Total 9 marks)
1
Na
Mg
Al
Si
P
S
Cl
Ar
K
a) Explain why the general trend from sodium
to
argon
is
upwards
but
why
the
value
for
Skill Check 4
sulfur
is
less
than
that
for
phosphorus.
(5)
(1) For each of the following pairs, state which element has the higher first ionisation energy
b)
Mark
on
the
graph
where
the
value
for
ks) and explain your answer:
potassium
would
be.
(1)
A Mg and Al
c) Explain why the value for the second
he
ionisation of sodium is very much larger
than that of its first ionisation.
(2)
e B Mg and Ca
(Total 8 marks)
(1)
(1)
C Ne and Na
(1)
ic 1)