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

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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!

Lone pairs

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 affect the shape of the molecule

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Working out a Lewis structure Example: diazene,

N2H2Step 1: Write the number of electrons in each atom and the number of bonds each atom can form

Nitrogen: 5 electrons, 3 bondsHydrogen: 1 electron, 1 bond

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

Don’t worry about the

shape…more on that later!

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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

2. O2

3. N2

4. H2O

5. HCl

6. NH3

7. CO2

8. HCN

9. C2H4

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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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 a lone pair donor

Lewis acid a lone pair acceptor

DATIVE COVALENT (CO-ORDINATE) BONDING

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.

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.

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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