Title: Lesson 3 Lewis Structures
Learning Objectives:
• Know how to draw and interpret Lewis structures
• Describe what a co-ordinate (dative) bond is
• Explain why some elements don’t follow the octet rule
Lewis structures

Show the position of outer-shell electrons in a covalent compound

Various types: all show the same thing, any is fine
dots and crosses crosses only
dots only
lines
Blue Circles: These are the bonding pairs of electrons – the ones involved in the
bonds.
Red Circles: These are non-bonding or lone pairs of electrons. They are very
important, but students often forget about them!
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Lone pairs
What is a lone pair?
Lone pairs occur in elements from group
5, 6 and 7
Lone pair
How many lone pairs
does Oxygen have?
Lone pairs
Lone pairs affect
the shape of the
molecule
Working out a Lewis structure

Example: diazene, N2H2
Don’t worry about
the shape…more
on that later!
Step 1: Write the number of
Nitrogen: 5 electrons, 3 bonds
electrons in each atom and the
Hydrogen: 1 electron, 1 bond
number of bonds each atom can form
Step 2: Draw the structure using
lines for bonds
There will be 2 N-H bonds and 1 N=N bond
Step 3: Add in the lone pairs
The N started with 5 electrons, and 3 are in
bonds, so that leaves 2 remaining…each N
will have one lone pair
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Alternative Steps to working out Lewis
Structure
First
method is
probably
easier but
you decide
which suits
you best!
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Time to practice…

Draw Lewis structures for the following, bearing in mind the previous two
slides
1.
H2
6.
NH3
2.
O2
7.
CO2
3.
N2
8.
HCN
4.
H2O
9.
C2H4
5.
HCl
10.
C2H2
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Lewis Structures for Ions


Calculate the valence electrons as above and then add one electron for
each negative charge and subtract one for each positive charge
Put Lewis structure in a square bracket with the charge shown outside.
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DATIVE COVALENT (CO-ORDINATE) BONDING
A dative covalent bond differs from covalent bond only in its formation
Both electrons of the shared pair are provided by one species (donor) and it
shares the electrons with the acceptor
Donor species will have lone pairs in their outer shells
Acceptor species will be short of their “octet” or maximum.
Lewis base
Lewis acid
a lone pair donor
a lone pair acceptor
Ammonium ion, NH4+
The lone pair on N is used to share
with the hydrogen ion which needs
two electrons to fill its outer shell.
The N now has a +ive charge as
- it is now sharing rather than
owning two electrons.
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Co-ordinate bonding
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Boron trifluoride-ammonia NH3BF3
Boron has an incomplete shell in BF3 and can accept a share of a pair of
electrons donated by ammonia. The B becomes -ive as it is now shares a
pair of electrons (i.e. it is up one electron) it didn’t have before.
Stopwatch Graph Home
The octet rule is not always followed...
Small atoms such as Beryllium and Boron form stable molecules in which the central
atom has fewer than eight electrons in its valence shell. This is an incomplete octet.
Incomplete octets are electron deficient and will accept an electron pair from a
molecule with a lone pair. This leads to the formation of a co-ordinate bond.
The expanded octet

In this example, the Lewis structure of
SO3 shows it with 12 electrons in the
outer shell

This is because sulphur can make use of
its empty d-orbitals (the 3d ones)

This is called an expanded octet

Period 2 elements can’t do this as they have
no d-orbitals
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Solutions
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