Title: Lesson 9 Expanded Octet Structures
and Formal Charge
Learning Objectives:
• Explain the exceptions to the octet rule
• Explain using formal charges the stability of comparable resonance
structures
Exceptions to the Octet Rule

Small atoms like Boron and Beryllium form stable molecules with fewer
than an octet of electrons
Atoms in period 3 or below may expand their octet by using the d orbitals
in their valence shell
- This is because the d orbitals available in the valence shell of these atoms
have energy values relatively close to those of the p orbitals
- Promotion of electros from 3p to empty 3d orbitals will allow additional
electron pairs to form.
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Elements such as phosporous and sulphur expand their octets, forming
species with five or six electron domains.
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Species with five domains

If all five electron domains are bonding electrons, the
shape will be triangular bipyramidal (90o, 120o, and
180o)

Four bonding pairs and one lone pair will give an
unsymmetrical tetrahedron or see-saw shape (90o,
117o, and 180o)
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Three bonding pairs and two lone pairs will give a Tshaped structure (90o and 180o)

Two bonding and three lone pairs will give a linear
shape (180o)
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Species with six domains

All six bonding pairs will give an
octahedral shape (90o)

Five bonding pairs and one lone pair will
give a square pyramidal shape (90o)

Four bonding pairs and two lone pairs will
give a square planar shape (180o)
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Formal Charges can be used to compare
Lewis Structures

Sometimes we can come up with more than one possible structure. This
can be the case with molecules that can form an expanded octet.

E.g. Sulphur dioxide can be represented as:
(i)
Standard structures
(ii)

Expanded octet
These structures only differ in the distribution of the electrons not the
skeletal structure. How can we tell which one is the most stable?
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Formal charge can help us decide on the
most stable structure...

Treats all covalent bonds as purely covalent with equal distribution – so
forgets electronegativity!

Idea is to count how many electrons ‘belong’ to each atom and compare
this with the number of valence electrons in the non bonded atom.
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Valence (V) is the element’s group number
 Number of electrons assigned to atom is calculated by
(a) Each atom having an equal share of atoms (even co-ordinate
bonds) (1/2 B)
(b) Lone pairs are own by it’s atom (L)

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Low formal charges mean that...

Less charge transfer has taken place in forming a structure
from it’s atoms  most stable and preferred structure
Standard structures
Expanded octet
TASK: Apply the formula for formal charge to the
different Lewis structures for SO2 and work out the FC
for each atom in each structure.
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Solution
Mirrored
structure = same
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Things to note...

All resonance structures contribute to the electronic
structure of the molecule, and the real or observed
structure is a combination of them.

The more stable the structure, the more it contributes to
the ‘real or observed’ structure.

Formal charges are used to compare resonance
structures that have different numbers of single, double or
triple bonds.

The sum of the FCs for a species must be zero for a
neutral molecule or equal to the charge of the ion.
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The most stable of the resonance Lewis
structures are:


Structures with the lowest formal charges on atoms and...
Negative values for formal charge on the more
electronegative atoms...
E.g.
Two structures with formal structures for N2O:
 Both have the same difference for formal charge  same
stability?
 But (i) is preferable because the more electronegative O
has the negative formal charge...
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Ozone

3 domains around the central atom  triangular planar,
but only 2 bonding  v shaped/bent

The double bond is composed of a sigma (σ) bond and a
pi ( ) bond.
Electrons in the pi bond are held less tightly so because
delocalised  forming resonance structures
True structure is resonance hybrid (intermediate bond
length) Bond order is 1.5 (3 pairs of electrons in 2
bonding positions)
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Ozone is a polar molecule...


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Surprising, since bonding is only between Oxygen atoms...
By checking the FC on each atom shows an uneven
distribution of electrons.
This gives a net dipole across the molecule making it
polar...
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Ozone is important...
The atmosphere is crucial for life on Earth...
 Traps heat
 Blocks dangerous radiation
 Contains Oxygen
 Lower part of stratosphere is the ozone layer
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Ozone levels are maintained in the ozone
cycle...

Oxygen dissociation
represents a free radical
(unpaired electron and
highly reactive)

Ozone dissociation

NOTE: Oxygen is broken down with a shorter wavelength of
light. Stronger O3 require higher energy radiation to break...
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Ozone protects us...



Absorbs radiation of
wavelengths 200nm315nm
UV-B & UV-C (these
can damage living
tissue)
Ozone absorbing UV
radiation is major
source of heat
(stratosphere gets
hotter as you get
higher)
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Ozone reacts easily with compounds
released in the atmosphere...

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Nitrogen oxides NOx
Chlorofluorocarbons CFCs
Nitrogen monoxide – produced in vehicle engines
Nitrogen dioxide – produced by the oxidation of Nitrogen
oxide
NO.(g) is a catalyst because is has been regenerated
Net change is breakdown of ozone
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CFCs...



Used in aerosols, refrigerants, solvents, plastics etc...
Unreactive/low toxicity in the troposphere...
In stratosphere, UV radiation breaks them down
(photochemical decomposition) Cl.

Weaker C-Cl bond breaks (not C-F bond  more polar)

Cl.(g) is a catalyst and regenerated and net reaction of
breakdown of ozone is:
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Solutions
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