Formation of Covalent Bonds
Senior 4 · Chemistry
Formation of Covalent Bonds
Lesson 1 of 5
25 min
A biogas digester fed with cow dung gives a fuel gas that is mostly methane, . Carbon and hydrogen have similar pulls on electrons; neither can take electrons from the other, so no ions form. Instead the atoms share electrons. Most of the substances around you, from water and the gases of the air to sugar, plastics and DNA, are held together this way. Unit 3 explained bonding by electron transfer. This unit explains sharing: how bonds form from overlapping orbitals, how to draw them, and how they decide the shape of a molecule and how molecules attract each other.
Definition
Covalent bond: a pair of electrons shared between two atoms. Both nuclei attract the shared pair, and this holds the atoms together.
Bond length: the distance between the two nuclei when the atoms are bonded. Bond energy: the energy needed to break one mole of the bond in gaseous molecules.
Bring two hydrogen atoms together. When they are far apart they do not interact. As they approach, the 1s orbital of each atom begins to overlap the other, and the two electrons (with opposite spins) occupy the region between the nuclei, where both nuclei attract them. The energy falls. If the atoms come too close, the two positive nuclei repel each other and the energy rises steeply. The bond forms at the distance of lowest energy (Figure 1.1).
So the bond length of is 74 pm, and 436 kJ is released when one mole of H–H bonds forms; the same energy is needed to break them.
Definition
Sigma (σ) bond: a covalent bond formed by head-on overlap of two orbitals along the line joining the two nuclei (the internuclear axis). The electron density is greatest on that axis. It can come from s–s, s–p or end-on p–p overlap.
Pi (π) bond: a covalent bond formed by sideways overlap of two parallel p orbitals. The electron density lies in two regions, above and below the internuclear axis, and none on the axis itself.
The first bond between two atoms is always a σ bond. A double bond is one σ bond and one π bond; a triple bond is one σ bond and two π bonds, at right angles to each other. Sideways overlap is less effective than head-on overlap, so a π bond is weaker than a σ bond, and the π electrons, lying away from the axis, are easier for other reagents to reach. That is why alkenes, with a C=C bond, react much more readily than alkanes.
Table 1.1: Carbon–carbon bonds (mean values)
| Bond | Made of | Bond length / pm | Bond energy / |
|---|---|---|---|
| C–C | 1 σ | 154 | 347 |
| C=C | 1 σ + 1 π | 134 | 612 |
| C≡C | 1 σ + 2 π | 120 | 838 |
Worked example 1: counting σ and π bonds
How many σ and π bonds are there in (a) ethene, ; (b) hydrogen cyanide, ; (c) carbon dioxide, ? Use Table 1.1 to estimate the strength of the π bond in ethene.
Solution. Every single bond is σ; each double bond adds one π; each triple bond adds two π.
π bond in C=C ≈ , less than the 347 of a σ bond, as expected for weaker sideways overlap.
A Lewis structure (or dot-and-cross diagram) shows how the outer electrons are arranged in a molecule: shared pairs between atoms, and lone pairs (non-bonding pairs) on single atoms. In most stable molecules of Period 2 elements, each atom ends up with eight outer electrons (an octet), hydrogen with two. To draw one:
In a displayed formula each shared pair becomes a line: H–O–H, O=C=O, N≡N.
Worked example 2: the Lewis structure of methanal,
Solution. (1) Outer electrons: C 4 + H 2 × 1 + O 6 = 12 (6 pairs). (2) Carbon in the centre with single bonds to H, H and O: 3 pairs used. (3) The remaining 3 pairs go on oxygen as lone pairs, giving oxygen an octet, but carbon now has only 6 electrons. (4) Move one oxygen lone pair into the C–O bond to make C=O.
Answer: : carbon has 4 bonding pairs (an octet), oxygen has 2 bonding pairs and 2 lone pairs. The molecule has 3 σ bonds and 1 π bond.
Warning
Common mistakes. (1) Counting a double bond as two σ bonds: it is one σ and one π. (2) Forgetting lone pairs: water has two on oxygen, and they matter for its shape and properties. (3) Drawing a π bond on the line between the nuclei: π density is above and below the axis. (4) Thinking a double bond is twice as strong as a single bond: 612 is less than .
Tip
Check every Lewis structure by counting: total electrons drawn = total outer electrons; each H has 2; each C, N, O and F has 8. Then count σ (one per bonded pair of atoms) and π (the extra lines).
Chemistry Student: Sigma and Pi Bonds EXPLAINED! (8 min). Watch how head-on overlap gives a σ bond and sideways overlap of p orbitals gives a π bond, and why the π bond is the weaker of the two.
Practice
Key points
Sources and credits
Quick check
How many σ and π bonds are there in one molecule of ethyne, H–C≡C–H?
