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CHEM1901/3 Worksheet 6: Molecular Geometry
Model 1: Oxidation Numbers
Oxidation numbers are a useful “accountancy” tool to help keep track of electrons in
compounds and reactions. This is particularly important in redox reactions where some atoms
lose (are oxidised) and others gain (are reduced) electrons. It is the positive or negative charge
the atom would have if the molecule was completely ionic.
Rules for working out oxidation numbers
The rules should be used in this order – the higher the rule, the higher its priority.
1.
An atom in its elemental form (e.g. Fe, Cl2, graphite etc) has oxidation number = 0
2.
The sum of the oxidation number of all the atoms in a molecule equals zero.
3.
The sum of the oxidation number of all the atoms in an ion equals the charge of the ion.
4.
The oxidation number of fluorine is -1 (except in F2 where it is 0 [rule 1]).
5.
The oxidation number of group 1 elements is +1.
6.
The oxidation number of group 2 elements is +2
7.
The oxidation number of oxygen is -2, except peroxides where it is -1.
8.
The oxidation number of halogens is usually -1.
9.
The oxidation number of hydrogen is +1 when bonded to non-metals and -1 when bonded
to metals.
Examples
(a)
ClF3 F has oxidation number = -1 (rule 4)
Cl must have oxidation number = +3 so that overall charge is zero: (+3) + (3×-1) = 0
(b)
ClF4+ F has oxidation number = -1 (rule 4)
Ion has charge of +1.
Cl must have oxidation number = +5 so that overall charge is +1: (+5) + (4×-1) = +1
(c)
ClF4- F has oxidation number = -1 (rule 4)
Ion has charge of -1.
Cl must have oxidation number = +3 so that overall charge is -1: (+3) + (4×-1) = -1
Critical thinking questions
1.
Work out the oxidation number of the N atom in the molecules and ions below:
System Oxidation number System Oxidation number HNO3 N2O NO3 -­‐
N2 NO2 NH2OH N2O4 N2H4 HNO2 NH3 NO2 -­‐
NH2 -­‐
NO NH4+ Model 2: Lewis Structures
Lewis structures can give us information about the electron arrangement in molecules.
The steps for drawing Lewis structures are:
Step 1. Count the total number of valence electrons on the molecule
Step 2. Draw single bonds between the atoms
Step 3. Count how many electrons are left and assign them as lone pairs to fill the
valence shells
Step 4: If there are unfilled valence orbitals, share the lone pairs to fill the valence
orbitals
Step 5: If there is more than one possible plausible structure, then we consider both
(all) to be resonance structures
Step 6: Minimise the formal charges on all atoms
Critical thinking questions
1.
For each of the following molecules and ions, draw the Lewis structure on the
underlined atom.
System Lewis structure System Lewis structure XeF2 XeF4 XeF3+ XeO2F2 SF2 SF4 SF6 SOF4 Model 3: Molecular Shape
The ‘Valence Shell Electron Pair Repulsion’ (VSEPR) model assumes that the because of the
repulsion between electrons, the bonds and lone pairs surrounding an atom try to get as far
from each other as possible. For molecules and ions of the type XYn in which the central atom
X makes bonded to n atoms and has m lone pairs, the shape is dictated by the total number
of bonded atoms and lone pairs: n + m.
The 3D-arrangements of bonds and lone pair which maximize the distance between these
bonds and lone pairs for common values of (n + m) are shown in the table below.
The molecular shape or geometry is the arrangement of the bonds. If lone pairs are present, they
help to dictate what this arrangement is but are not included when the geometry of the molecule is
described. The table below shows the geometries possible when n + m = 4 - 6.
n + m 180o
120o trigonal planar bent 109.5o
4 X
120 o
Y
Y
tetrahedral X
Y
Y
bent or V shaped Y
Y
Y
X
X
Y
X
X
Y
trigonal pyramidal Y
90o m = 3 X
3 5 m = 2 X
2 m = 1 m = 0 Y
trigonal bipyramidal Y
see-­‐saw X
Y
T-­‐shaped Y
Linear Y
6 90o
Y
X
Y
octahedral X
Y
Y
Y
Y
square-­‐based pyramid X
Y
Y
square planar Note that when n = 5, the lone pairs occupy the sites in the equatorial plane.
T-­‐shaped Critical thinking questions
1.
2.
3.
Complete the table below. The central element is underlined.
System n m n + m Arrangement SF2 SF4 SF6 SOF4 XeF2 XeF4 XeF3+ XeO2F2 Sketch and describe the molecular geometries of the following systems.
(a) SF2
(b) SF4
(c) SF6
(d) SOF4
(e) XeF2
(f) XeF4
(g) XeF3+
(h) XeO2F2
Which of these molecules or ions is polar?