Colourful Solutions > The covalent model > Multiple covalent bonds

The Mad Science Lab

Standard level

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 ref: 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.

Guidance

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.

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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

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Worked examples

Q215-01 When the Lewis structure for HCOOCH3 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, C4H10, propanal, C3H6O and propan-1-ol, C3H8O 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 N3H

Answer
Lewis structure 1 Lewis structure 2
Q215-05 Draw Lewis structures of each of the following species, NO2- and NO2+

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 BF3 and NF3

Answer
BF3 NF3
Q215-07 Write two Lewis electron dot structures for the methanoate ion HCOO-

Answer
Structure 1 Structure 2
Q215-08 Write Lewis electron dot structures for H2NNH2 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 CO2 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 NCl3. Predict giving a reason the Cl-N-Cl bond angle in NCl3

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
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