REVISION
Period trend in atomic radius: Atomic radius decreases from
left to right across the horizontal rows (i.e. periods) in the
periodic table.
Explanation: Moving across a period from left to right the
nuclear charge increases but the number of screening
electrons in core orbitals remains constant. Hence the
effective nuclear charge increases from left to right and the
outer (i.e. valence) electrons are drawn closer to the
nucleus.
Group trend in atomic radius: Atomic radius increases from
top to bottom of the vertical columns (i.e. groups) in the
periodic table.
Explanation: Each move down a group corresponds to the
completion of a shell and the start of a new one. Screening
by the completed shell exactly balances the increase in
nuclear charge so that the effective nuclear charge remains
constant. Hence atomic radius increases, as expected, with
increasing value of n, the principal quantum number.
Ionisation Energy - the energy required to remove the most
loosely-bound electron to infinity from a neutral isolated
atom, i.e. the energy for the process:
E(g)
E+(g) + e-
Group Trend in ionisation energy: Ionisation Energy
decreases from top to bottom of a group.
Explanation: Charge/size effects – constant ENC but
increasing atomic radius down group – outer electrons
increasingly weakly held by the +ve charge on the nucleus.
Period Trend: Ionisation Energy increases from left to right
across a Period.
Explanation:
1. Overall trend - Charge/size effects – From left to right
across a period ENC increases and atomic radius
decreases – both factors results in outer electrons being
more strongly held.
2. Detailed trend – Electron configuration effects. Atoms
with stable electronic configurations (filled subshell,
half-filled subshell, filled shell) have IE values higher
than expected from charge/size effects alone. Atoms
with one electron more than these configurations have IE
values lower than expected from charge/size effects
alone.
Electron Affinity: The energy change associated with the
addition of an electron to an isolated neutral atom.
Element(g) + e
Element(g)
Group Trend: Electron Affinity becomes less negative (i.e.
less favourable) from top to bottom of a group.
Period Trend: Electron Affinity becomes more negative
(i.e. more favourable) going from left to right across a
period but the detailed trend is irregular.
Trends are explained as for IE, that is charge/size effects
(group trend) and/or a combination of charge/size and
electron configuration effects (period trend).
ATOMIC STRUCTURE AND BONDING
What happens when atoms react with each other?
When atoms interact via their valence electrons two
different kinds of attractive force – called chemical
bonding – can result.
(1) Electron transfer: Ionic bonding – if one of the
reacting atoms has a low Ionisation Energy (e.g. Li) while
the other has high Electron Affinity (e.g. F) electrons will
be transferred between them to form ions. The charged
atoms are held together by an electrostatic attractive force:
Li(g) + F(g)  Li+ + F (= LiF)
Why is just one electron transferred in this reaction?
Li {He} 2s1

Li+ {He} + e
F {He} 2s2 2p5 + e  F {Ne}
The 'Noble Gas Rule' for electron transfer: In forming
ions, atoms gain or lose electrons in such a way as to
attain the stable, filled (closed) valence-shell electronic
structure of the nearest Noble gas.
Another example
- reaction of magnesium metal with
nitrogen:
Mg {Ne} 3s2  Mg2+ {Ne} + 2e
N {He} 2s2 2p3 + 3e  N3 {Ne}
Note that the stoichiometry, i.e. the combining proportions
of the reactants, is determined by the fact that the total
number of positive and negative charges must balance out
- the total number of electrons lost by the metal must equal
the total number of electrons gained by the non-metal.
Hence:
3 Mg + 2 N  Mg3N2
Electron Sharing: Covalent Bonding - Sharing of
electrons between atoms occurs when neither has a strong
tendency to gain or lose electrons. Shared electrons
produce a wave-mechanical attractive force between the
atoms:
H• + •H  H•—•H
What does ‘sharing electrons’ actually mean?
H
+
H
H 
H
Formation of the H2 molecule from two hydrogen atoms.
Two hydrogen 1s1 aromic orbitals overlap to form a new
orbital which encloses both nuclei and is called a
molecular orbital (MO).
The two electrons from the two H atoms fill the molecular
orbital. The same filling rules apply to molecular orbitals
as to atomic orbitals, i.e. no more than two electrons – with
paired (i.e. opposite) spins – can occupy an individual
molecular orbital.
Covalent bonding involves the sharing of one or more
pairs of electrons between the bonded atoms. The shared
pair can be symbolised by a short line joining the two
bonded atoms:
Br + Br
BrBr i.e. Br2
One shared pair: a single bond.
C + O2
O=C=O
Two shared pairs of electrons : a double bond.
N + N
NN i.e. N2
Three shared pairs of electrons : a triple bond.
What determines the number of shared pairs of electrons?
The 'Noble Gas Rule' for electron sharing: - In forming
covalent bonds atoms will share electron pairs in such a
way that both atoms attain the Noble Gas electron
configuration in the newly formed molecule.
A shorthand notation for valence shell electronic
configuration: Lewis Structures.
Valence shell electron pairs are indicated as pairs of dots
adjacent to the symbol for the element. If the valence shell
contains a single electron in an orbital that is indicated by
a single dot.
Examples:
Chlorine, Cl : Z = 17  17 electrons: 1s2 2s22p6 3s23p5:
7 valence electrons (1 electron missing for full shell):
..
.
:Cl
..
Chlorine molecule:
..
..
. + :Cl. 
:Cl
..
..
..
..
:Cl
.. •—• Cl
.. :
The reaction of chlorine with carbon:
2
2
2
C: Z = 6, i.e. 6 electrons - 1s 2s 2p
Energy
1s2 2s1 2p3
.
4 valence electrons : . C. .
..
:Cl:
.
. C. . +
..
.
4 :Cl
..
..
:Cl
..
C
..
Cl
.. :
:Cl
.. :
Carbon tetrachloride
.
.C. +
.
..
.
2.O
..
..
O
..
C
..
O
..
Carbon dioxide
Energy kJ mol-1
The Energetics of Formation of the Hydrogen Molecule:
H2
H
H
0
-200
-458 kJ mol-1
-400
r0
0.1
0.2
Internuclear distance, r (nm)
r0 = 0.074 nm = the bond length of H2 molecule
The H2 molecule is 458 kJ mol-1 more stable than the two
isolated H atoms, i.e. the bond strength of the H2 molecule
is 458 kJ mol-1
The Wave-mechanical Nature of Covalent Bonding –
Molecular Orbital (MO) Theory.
1. Atomic orbitals interact to form new orbitals which
surround both nuclei and hence are called molecular
orbitals.
2. Combination of any two atomic orbitals produces two
molecular orbitals (MOs).
a + b
+
a
b
1s
bonding
1s*
anti-bonding
a - b
One MO (, sigma) corresponds to the sum of the atomic
wave functions, the other MO (*, sigma star)
corresponds to the difference of the atomic wave
functions.
a + b
+
a
1s
bonding
1s*
anti-bonding
b
 a - b
A
B
+
+
A
+ B
+ +
+
The sum of the wave functions produces an orbital with
increased  hence increased 2,
 hence increased
electron density between the nuclei.
This prevents the positively charged nuclei from
repelling each other while both nuclei are held together by
their attraction to the electron density between them.
a + b
+
a
1s
bonding
1s*
anti-bonding
b
a - b
This MO is called a bonding molecular orbital (BMO)
because electrons in this orbital hold the two nuclei
together. The BMO is of lower energy (i.e. more stable)
than the two isolated atomic orbitals.
1s
bonding
a + b
+
a
1s*
anti-bonding
b
 a - b
A
+
B
A
A
+
+
-
B
A -  B
B
The difference of the atomic wave functions produces zero
, hence zero 2, hence no electron density
between the nuclei. This MO is called an anti-bonding
molecular orbital (ABMO). It is higher in energy (i.e.
less stable) than the individual atomic orbitals. Electrons in
this orbital cause the nuclei to repel each other.
Combination of:
two s AO’s
or
end-on combination of one s AO + one p AO
or
end-on combination of two p AOs
generates one  (sigma) and one * (sigma star) MO.
Sigma () MOs are bonding MOs. Sigma star (*) MOs
are anti-bonding MOs.
 bond
s-s
s-p
p-p
In  bonds the electron density is concentrated along the
bond axis, that is along the imaginary line joining the two
atomic nuclei.
Sideways-on combination of two p AOs produces a  (pi,
bonding) and a * (pi star, anti-bonding) MO.
 bond
 bonds are found
in double bonds and
triple bonds.
Unlike  bonds, in  bonds the electron density is
concentrated above and below the bond axis, that is
above and below the imaginary line joining the two atomic
nuclei.
Double bonds involve one  and one  bonding MO.
Triple bonds involve one  and two  bonding Mos.
In molecules electrons occupy MOs in order of increasing
energy (Aufbau Principle). The Pauli Principle (an MO
can contain a maximum of
2 electrons with opposite
spins) and Hund’s Rule (orbitals of equal energy fill singly
before filling in pairs) also apply.