As
we have already learned, covalent bonding occurs between
two or more non-metals. We now need to describe exactly how
this
type of bonding occurs, and represent covalent bonding
using dot and cross diagrams.
1. Single
Covalent Bonds.
In covalent
bonding, pairs of electrons are shared between the two non-metals.
A shared pair of electrons is called a single covalent bond, and each
atom contributes
one electron to the bond.
A single
covalent bond is shown using a single line e.g. H2 is written as H-H
Both the
non-metals now have full outer shells of electrons. This arrangement
of electrons will be the same as in the noble gases, and this is a
very stable arrangement.
Checkpoint
1.
Draw
dot and cross diagrams for the following molecules.
a.
C2
Click
here to see the answe
b.
H2O
Click
here to see the answer
c.
NH3
Click
here to see the answer
d.
CH4
Click
here to see the answer
2. Multiple
Covalent Bonds
Sometimes
molecules share more than one pair of electrons in order to achieve
full shells of electron. If two pairs of electrons are shared then
a double bond is formed, and if three pairs are shared then a triple
bond is formed. It is not possible to form a "quadruple" bond!
Oxygen
(O2) molecules contain a double bond, as shown in the diagram below.
The double bond is written as O=O.
Checkpoint 2.
Draw dot and cross
diagrams for the following molecules. (Hint: they all contain a least one double
or triple bond, and all the atoms have full outer shells of electrons)
a. Nitrogen molecules
(N2)
Click
here to see the answer
b. Ethene
(C2 H4)
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here to see the answer
c. Carbon dioxide
(CO2)
Click
here to see the answer
Further
Work.
You
should now take a look at the following
website. Read through the notes, and answer
the questions.
You can then try the test at the end!
BBC
Chemistry Resources
- 
3. "Unusual" Covalent
Compounds.
So
far all the compounds we have looked at have had complete outer
shells of electron - stable noble gas configurations. There
are however a few covalent compounds that break this rule -
some have less than the usual eight electrons in the outer
shell, and some have more!
An
example of a compound with an incomplete outer shell is boron
trifluoride - BF3.
Notice
that although the fluorine atoms have full outer shells of
electrons, the boron atom only has six electrons around it.
An
example of a compound with an expanded outer shell is sulphur
hexafluoride - SF6.
Notice
that although the fluorine atoms have full outer shells of
electrons, the sulphur atom has twelve electrons around it.
The reasons why this can happen are not required at GCSE, but
the sulphur atom has other energy levels that the extra electrons
can go into.
Checkpoint 3.
Draw
dot and cross diagrams for the molecules below. The first two
have incomplete outer shells of electrons, and the last two
have expanded outer shells.
a.
BCl3
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here to see the answer
b.
BeCl2
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here to see the answer
c.
PF5
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here to see the answer
d.
BrF3
Click
here to see the answer
4. Lone
Pairs and Bond Pairs of Electrons.
If
we examine the structure of methane (CH4), we can see that
there are four pairs of electrons involved in covalent bonds.
We say there are four bond pairs of electrons around carbon.
If we now
examine the structure of water (H2O), we can see that there
are two bond pairs of electrons around oxygen, but there
are also two pairs of electrons around oxygen that are not involved
in covalent bonds. These are known as lone pairs of
electrons.
In water
there are two lone pairs of electrons, and two bond pairs
of electrons.
Checkpoint
4.
Using
the dot and cross diagrams you have already drawn in this section,
count up the bond pairs and lone pairs around the central atom
in the molecule. Check your answers below.
a. H2
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here to see the answer
b. BeCl2
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here to see the answer
c. BF3
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here to see the answer
d. PF5
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here to see the answer
e. SF6
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here to see the answer
f. NH3
Click
here to see the answer
g. H2S
(you will need to draw a dot and cross diagram for this one)
Click
here to see the answer
5. Dative Covalent Bonds.
A
covalent bond contains a pair of electron, with each atom contributing
one electron to the bond. In a dative covalent bond, both of
the shared electrons in the bond have come from one atom. Dative
covalent bonds are also known as coordinate bonds. This type
of bonding can occur between molecules with a pair of electrons
to donate i.e. molecules with a lone pair, and molecules that
have room for a pair of electrons i.e. those with an incomplete
outer shell.
The
ammonium ion (NH4+) is an example of a compound containing
a dative covalent bond. It is formed when ammonia and a hydrogen
ion react together.
Notice
how all the atoms now have full outer shells of electrons
- remember that there are only two electrons in the first
shell.
The
dative covalent bond is shown using an arrow, with the arrow
pointing to the atom accepting the electrons.
One important
thing to remember is that there is now way to tell apart
a covalent and dative covalent bond - both contain a shared
pair of electrons and there is no way to tell where these
electrons came from!
Checkpoint 5.
Draw
dot and cross diagrams for the following compounds that all
contain a dative covalent bond. It will help if you draw dot
and cross diagrams for the starting materials first! All atoms
should have full outer shells of electrons.
a. BF3NH3 This compound is formed when BF3 and NH3 react together.
Click
here to see the answer
b. H3O+ This compound is formed when H2O and H+ react together.
Click
here to see the answer
c. CO This is very difficult - this molecule is called carbon monoxide.
It has a multiple bond as well as a dative covalent bond.
Click
here to see the answer
