COVALENT BONDS

1. Introduction and revision

In Grade 10, we saw how atoms can bond together to form molecules. You might want to revise this topic before going on. The theory of covalent bond formation was originally developed by G. N. Lewis. It described how certain elements could fill their outer electron shells by sharing electrons with the atoms of other elements. Covalent bonds normally occur between non-metals, or between hydrogen and non-metals. Thus, in the example below hydrogen shares its single 1s electron with a 3p electron of chlorine, to form a covalent bond between the H and Cl atoms:

We commonly represent a covalent bond in either the Lewis or Couper notations, shown here on the right:

In the Lewis notation, only the outer, valence electrons are shown, while in the Couper notation, a covalent bond is represented as a solid line. We call such bonds SINGLE BONDS. Remember, a single solid line joining two atoms represents TWO ELECTRONS!


2. Multiple covalent bonds

Atoms can form covalent bonds by sharing more than one pair of electrons between them. Thus, oxygen, with an electronic structure 1s2 2s2 2p4, can link two atoms to form the oxygen molecule, O2, by sharing two pairs of electrons. This arrangement is called a DOUBLE BOND. Nitrogen, with the electronic structure 1s2 2s2 2p5, gives rise to the nitrogen molecule, N2 by sharing three pairs of electrons. This is called a TRIPLE BOND.

3. Bond energies

The hydrogen molecule is very stable, and at room temperature, there is little or no tendency to split the hydrogen molecule into atoms. At high temperatures however, this does happen, suggesting that to break a covalent bond, energy must be supplied to the molecule. The strength of the H-H covalent bond may be discussed in terms of the energy required to break that covalent bond and form individual atoms, each with an unpaired electron (Note for educators). This energy is called the BOND DISSOCIATION ENERGY, ΔH for the H-H bond, which is found experimentally to be 436 kJ·mol-1 ().

One would reasonably expect that double and triple bonds would have higher bond dissociation energies than single bonds, and this is in practice found to be the case, as shown in the table on the right.

Bond ΔH
kJ·mol-1
H-H
O=O
N≡N
436
494
941

4. Bond lengths

The length of a covalent bond can be measured by various physical techniques, and it is expressed as the BOND LENGTH or BOND DISTANCE, which is the distance between the centres of the two atoms joined by the covalent bond.

Multiple bonds are shorter than single bonds.

Bond Length
pm
O-O
O=O
N-N
N=N
N≡N
148
121
145
125
110

The bond length for a covalent bond is approximately equal to the sum of the COVALENT RADII of the atoms that are linked. The covalent radius of an atom is the radius of the sphere representing that atom when it is covalently bonded to another atom. It is indicated in the diagram on the left by the symbol r:


5. Electronegativity of atoms

When a covalent bond is formed between two identical atoms, such as H-H or Cl-Cl, the pair of electrons which joins the atoms is evenly shared between the two atoms. However, when two different atoms are joined together by a covalent bond, the sharing of the electron pair is not even, and the pair of electrons is shifted towards one of the atoms:

The covalent bond is said to have been POLARIZED, and one refers to such a bond as a POLAR COVALENT BOND.

The ability to polarize a covalent bond differs from one atom to another, and is known as the ELECTRONEGATIVITY of the atom.

Electronegativities are measured on an arbitrary scale ranging from 0 (He) to 4.1 (F).

The rare gases (He, Ne, Ar, Kr, Xe, Rn) have zero electronegativities, that is, they only form covalent bonds, and this only in exceptional cases (), and have no tendency to attract electrons of that bond.

The halogens (F, Cl, Br, I, At) appear at peak values of electronegativities. Within a period, they have the highest tendency to polarize covalent bonds. Note that within the group, the electronegativity decreases.

On the whole, electronegativities INCREASE from left to right along a period (see Li to F) and DECREASE from top to bottom within a group (see F to At).

6. Additional questions


Note for educators

The bond dissociation energy is, strictly speaking, the change in ENTHALPY, ΔH, that occurs when one mole of the bond undergoes complete dissociation. The energy change, ΔE is equal to ΔH - PΔV where ΔV is the volume change occurring at a constant pressure P.