Unless otherwise stated, all images in this file have been reproduced from:
Blackman, Bottle, Schmid, Mocerino and Wille,
Chemistry, 2012 (John Wiley & Sons)
ISBN: 9 78047081 0866
1
Lewis structures
Lecture 15
Ionic bonds
Ionic solids
Lecture 16
Covalent bonds
Lewis structures
Lecture 17
VSEPR
Polyatomic molecules
Earlier we described
bonding in diatomic
molecules in terms of the
molecular orbitals.
HF
H2O
σ*
For polyatomic molecules,
the MOs become much
more complicated and
often involve many atoms.
n
We need a simpler model
to describe bonding.
σ
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Lewis model of bonding
• Simple model to allow us to determine the connectivity in and structure
of molecules
• Phrased as a set of rules or algorithm for structure-writing
• A chemical bond is thought of as a pair of electrons shared between
two atoms; atoms share bonds to achieve a full valence shell (8 e−)
• Developed in 1916
• before the nuclear structure of the atom was understood
• before the wave nature of the electron was discovered
• before molecular orbital theory was developed
• Still widely used by chemists to represent molecules and bonds
• Extremely useful but incomplete model → need to use it judiciously and
understand its limitations as well as its strengths
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Lewis structures of simple molecules
A simple strategy for drawing molecular structures:
Step 1:
Count the total number of valence electrons in the molecule.
Step 2:
Draw single bonds between the atoms. (Assume 2 e− per bond.)
Step 3:
Count how many electrons are left and assign them as lone
pairs to fill valence shells (usually 8 e−). Octet rule. Start with
outer atoms.
Step 4:
If any octets are incomplete, share lone pairs to form multiple
bonds.
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Lewis structures of simple molecules
Lewis structure of hypochlorous acid (HOCl):
O has two lone pairs.
Cl, typical of halogens, has 3.
In polyatomic molecules, the centre atom is the one that needs to share the
greatest number of electrons.
Lewis structure of Oxygen molecule (O2):
O2 has double bond.
Unfilled
valence
shell
Bond order = 2
Filled valence shells
Each O atom has two lone pairs
Worksheet questions 1 and 2
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Resonance structures and delocalisation of electrons
Sometimes the Lewis model allows several valid structures, e.g. O3:
This structure satisfies the Lewis rules, but is
usually disregarded because three-membered
rings are not very stable (but they are known).
These resonance structures are equivalent representations of bonds that are
“in-between” single and double bonds.
We use a double headed arrow, not equilibrium arrows ( ) for resonance.
In ozone, this means that both O-O bonds are equivalent,
and we have an average bond order of 3/2 or 1½.
Remember that electrons are really delocalised between
more than two nuclei in molecular orbitals. Resonance is
simply a way to describe this in terms of Lewis structures.
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Lewis structures of simple molecules
A simple strategy for drawing molecular structures:
Step 1:
Count the total number of valence electrons in the molecule.
Step 2:
Draw single bonds between the atoms. (Assume 2 e− per bond.)
Step 3:
Count how many electrons are left and assign them as lone
pairs to fill the valence shells (usually 8 e−). Octet rule. Start
with outer atoms.
Step 4:
If any octets are incomplete, share lone pairs to form multiple
bonds.
Step 5:
If there is more than one plausible structure, then we consider
both (all) to be resonance structures.
Worksheet question 3
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Formal charges (i)
Formal charge is the charge assigned to an atom in a molecule, assuming that electrons are
shared equally between atoms.
→ It is the difference between the number of valence electrons of each atom, and the
actual number of electrons associated with it.
For example, an oxygen atom has 6 valence electrons. In the ozone molecule:
2 lone pair electrons
6 shared electrons
FC = 6−2−½(6) = +1
6 lone pair electrons
2 shared electrons
FC = 6−6−½(2) = −1
4 lone pair electrons
4 shared electrons
FC = 6−4−½(4) = 0
Formal
Valence electrons
Lone pairs
½ (shared
=
−
−
Charge
of free atom
of electrons
electrons)
Worksheet question 4
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Formal charges (ii)
Formal charge can be used to decide the most likely Lewis structure from a number of possibilities.
1
S: FC = 6−6−½(2) = −1
C: FC = 4−0−½(8) = 0
N: FC = 5−2−½(6) = 0
2
S: FC = 6−4−½(4) = 0
C: FC = 4−0−½(8) = 0
N: FC = 5−4−½(4) = −1
3
S: FC = 6−2−½(6) = +1
C: FC = 4−0−½(8) = 0
N: FC = 5−6−½(2) = −2
Step 6: Minimise the formal charges
on all atoms.
• Structure 3 has larger formal charges
→ unlikely
• Structures 1 and 2 have the same
formal charges (just in different
places) → compare
electronegativities
• N is more electronegative than S, so
structure 2 is more likely than
structure 1
• But, both 1 and 2 are seen in nature
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Lewis structures of simple molecules
A simple strategy for drawing molecular structures:
Step 1:
Step 2:
Step 3:
Step 4:
Step 5:
Step 6:
Count the total number of valence electrons in the molecule.
Draw single bonds between the atoms. (Assume 2 e− per bond.)
Count how many electrons are left and assign them as lone
pairs to fill the valence shells (usually 8 e−). Octet rule. Start
with outer atoms.
If any octets are incomplete, share lone pairs to form multiple
bonds.
If there is more than one plausible structure, then we consider
both (all) to be resonance structures.
Worksheet question 5
Minimise the formal charges on all atoms.
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Breaking the octet rule? (i)
Atoms can have a maximum of 4 electron pairs (8 e−) surrounding them
(H can only have 1 electron pair, 2 e−).
But, B and Be may have fewer than 8 valence electrons, e.g. BF3.
Elements below the second period may bind more than 4 other atoms.
If you follow the rules for these molecules, it is looks like you have to break the octet:
Group 15 (P, As, Sb)
up to 5 atoms
Group 16 (S, Se, Te)
up to 6 atoms
Group 17 (Cl, Br, I)
up to 7 atoms
10 valence e− for P?
12 valence e− for S?
12 valence e− for Br?
In reality, these molecules do obey the octet rule → the links here comprise < 2 e−
(e.g. SF6 has only 4 bonding MOs to cover the 6 links → each BO = 2/3.)
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Hypercoordinate molecules
1. Count the total number of valence electrons
2. Form bonds as usual, but make sure you don’t
exceed the octet for any atom
3. Assign remaining electrons as lone pairs
4. If all octets are full, un-share bonding pairs
to make space for left over electrons
5. This gives structures which appear to have
electrostatic interactions between charged
ions (partial ionic bonding)
6. But… all possible resonance structures must
be considered
7. Overall, two bonds are shared over three
linkages → on average, bond order is 2/3
7 + 7  3 = 28
Consistent with MO
theory – only two
true bonding MOs
Worksheet Question 6
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Breaking the octet rule? (ii)
For SO2, you might come across either of the following structures:
Analogous structure to O3
CO2-like structure
O: 6 valence electrons
6 lone pair electrons
2 shared electrons
formal charge = -1
O: 6 valence electrons
4 lone pair electrons O: 6 valence electrons
S: 6 valence electrons
4 shared electrons
2 lone pair electrons
4 lone pair electrons
formal charge = 0
8 shared electrons
4 shared electrons
formal charge = 0
formal charge = 0
S: 6 valence electrons
2 lone pair electrons
• Assumes breaking the octet is ok for S
6 shared electrons
formal charge = +1
• Uses this to eliminate formal charges
• This is a much better approximation to
the MO description of this molecule
• This is what you will see in most text
books and in a lot of scientific writing
Worksheet question 7
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Summary of the basic principles of the Lewis model
Lewis structures allow us to draw out plausible or reasonable structures for connectivity in
molecules and ions and to get some insight into bond orders and electron distributions.
Octet rule: atoms share electrons in order to have filled valence shells, i.e. adopt the
electron configuration of the following noble gas.
• A filled valence shell for H contains 2 electrons.
• A filled valence shell for all other main block
elements contains 8 electrons.
• Elements in the 3rd, 4th etc. rows may have extra
linkages, but they still obey the octet rule.
• Elements at the beginning of a row may not
achieve a full valence shell.
Note: Bonding electrons are counted towards the valence shells of both bonding partners.
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Summary of rules
Step 1:
Count the total number of valence electrons in the molecule.
Step 2:
Draw single bonds between atoms. Assume 2 e− per bond
– do not exceed octets!
Step 3:
Count how many electrons are left and assign them as lone pairs to
fill valence shells (8 e− except for H, Be, B). Start with outer atoms.
Step 4a: If any octets are incomplete, share lone pairs → multiple bonds.
Step 4b: If electrons are left over, un-share bonding pairs → lone pairs.
Step 5:
If there is more than one plausible structure, consider both (all)
resonance structures.
Step 6:
Minimise the formal charges on all atoms. (But don’t exceed octets!)
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Learning outcomes
After this lecture you should be able to:
• Draw out plausible Lewis structures for simple polyatomic molecules.
• Assign bond orders based on sharing of electrons, resonance
structures and formal charges.
• Explain the relationship between resonance and electron
delocalisation in molecular orbitals.
Further Reading: Blackman 5.3 (but ignore everything they say about
expansion of the octet)
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