Covalent bonds may be single, double or triple. This section looks at the effect of bond length on bond energy, as well as the impact of multiple bonds.
Syllabus reference S2.2.2
Structure 2.2.2 - Single, double and triple bonds involve one, two and three shared pairs of electrons respectively.
- Explain the relationship between the number of bonds, bond length and bond strength.
Tools and links
- Reactivity 2.2 - How does the presence of double and triple bonds in molecules influence their reactivity?
Double and triple bonds
A double bond is shown as two shared pairs of electrons, each of the bonded
atoms provide two electrons for the bond. In the oxygen molecule at the left,
their are two shared pairs of electrons giving a stable octet (eight) of electrons
around each oxygen atom. 
In triple bonds there are three pairs of electrons holding the two atoms
together. The molecule shown at the right is the ethyne (acetylene) molecule.
Each
of the carbon atoms (group IV) originally started with 4 outer electrons.
In the ethyne molecule at the right, both of the carbon atoms now have a full
octet (eight) of electrons and the two hydrogen atoms have a full first shell,
with two electrons. There are no non-bonding pairs.
^ top
Length and strength of covalent bonds
A covalent bond is a shared pair of electrons in which the electron charge density lies along the inter-nuclear axis. This attracts the nucleus of each bonded atom holding them together.
The bond is a result of the two electrostatic forces, nucleus-electrons and electrons-nucleus.
The greater the electron density between the two nuclear centres, the greater the electrostatic force and the stronger the bond.
Hence, double bonds are shorter and stronger than single bonds.
The size of the atoms is also a factor in bond strength, as smaller atoms have the bonding electrons closer to the nuclei where they can exert a stronger electrostatic force. In general, short bonds are stronger bonds.
Covalent bond strength
Bond type
Length (x 10-12 m)
Strength /kJ mol-1
C-C
154
346
C=C
134
614
C≡C
120
839
^ top
Worked examples
Q215-01 When the Lewis structure
for HCOOCH
3 is drawn how many bonding and how many lone pairs of
electrons are present?
| |
Bonding pairs |
Lone pairs |
| A. |
8 |
4 |
| B. |
7 |
5 |
| C. |
7 |
4 |
| D. |
5 |
5 |
Answer
|
The best way is to draw out the structure:
Counting up there are 4 lone pairs and
8 bonding pairs.
|
Q215-02 What is the Lewis
(electron dot) structure of sulfur dioxide?
Answer
|
Sulfur dioxide has the formula SO2. Sulfur has six valence
(outer shell) electrons and each oxygen also has six valence (outer
shell) electrons, giving a total of 18. By double bonding one of the
oxygens to the sulfur, the oxygen and sulfur now have a full octets.
The sulfur can then donate a pair of electrons into the second oxygen's
outer shell to leave all the atoms with full octets
|
Q215-03 Butane, C
4H
10,
propanal, C
3H
6O and propan-1-ol, C
3H
8O
have very similar molar masses (59±1). Draw the Lewis structures of each
of these molecules.
Answer
|
|
|
|
|
Click on the compound name
for the Lewis structure
|
|
Q215-04 Hydrazoic acid can
be represented by two possible Lewis structures, in which the atoms can be arranged
as NNNH. Draw the two possible Lewis structures of N
3H
Answer
| Lewis structure 1 |
Lewis structure 2 |
|
|
|
|
Q215-05 Draw Lewis structures
of each of the following species, NO
2- and NO
2+
Answer
|
With the negative ion, NO2-, you have to add
one electron to the structure. So N (5 valence electrons ) and 2 x
oxygen (six valence electrons) = 17 electrons + 1 for the negative
charge = 18 electrons.
With the positive ion, NO2+, you have to subtract
one electron from the structure. So N (5 valence electrons ) and 2
x oxygen (six valence electrons) = 17 electrons - 1 for the negative
charge = 16 electrons.
| NO2- |
NO2+ |
|
|
|
Q215-06 Draw Lewis structures
to represent BF
3 and NF
3
Answer
| BF3 |
NF3 |
|
|
|
Q215-07 Write two Lewis electron
dot structures for the methanoate ion HCOO
-
Answer
| Structure 1 |
Structure 2 |
|
.gif)
|
|
Q215-08 Write Lewis electron
dot structures for H
2NNH
2 and HNNH. What bond angle is
expected for the H-N-N atoms is each molecule?
Answer
| Structure 1 |
Structure 2 |
|
|
| Both nitrogen atoms are electronically
tetrahedral. The tetrahedral bond angle is 109º 28', but
there is a lone pair squeezing the angle together. The expected
bond angle is 107º. |
Both nitrogen atoms are electronically trigonal
planar (three regions of electrical charge). The trigonal bond
angle is 120º', but there is a lone pair squeezing the angle
together. The expected bond angle is a little less than 120º. |
|
Q215-09 Draw Lewis (electron
dot) structures for CO
2 showing all valence electrons.
Answer
|
Count up the available valence electrons:
- Carbon = 4, 2 x oxygen = 12
- Total valence electrons = 16 (8 pairs)
Connect both of the oxygen atoms to the central carbon by a double
bond. O=C=O
This uses up four pairs of electrons, leaving four pairs to place
on the structure, so that all atoms have a full octet.
|
Q215-10 Draw the Lewis structure
of NCl
3. Predict giving a reason the Cl-N-Cl bond angle in NCl
3
Answer
 |
Count up the valence electrons.
- Nitrogen from group V = 5 valence electrons
- Each chlorine provides 1 electron = 3 electrons
- Total = 8 electrons = 4 pairs
The molecule is electronically tetrahedral, but only 3 pairs
of electrons are used in bonding. Thus it has a lone pair.
The expected bond angle for a tetrahedral electronic arrangement
is 109º 28'. However, the repulsion expected between the
lone pair and the bonding pairs squeezes the Cl-N-Cl bonds together
by about 2.5º.
Therefore, the expected bond angle
is 107º
|
|
