Multiple choiceQuestion 1 · Recall
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By the end you can describe the trend in the strength of the halogen–halogen bond down Group 17, explain it in terms of bond length and orbital overlap, and account for the unusual strength of the F–F bond.
The strength of a covalent bond is measured by its bond enthalpy: the energy needed to break one mole of that bond in the gaseous state.
For a halogen molecule, , there is only one bond to break, so the value is exact.
| bond | bond length / nm | bond enthalpy / kJ mol⁻¹ |
|---|---|---|
| – | 0.142 | 158 |
| – | 0.199 | 242 |
| – | 0.228 | 193 |
| – | 0.267 | 151 |
Read the table carefully. From chlorine to iodine, the bond gets longer and weaker. Fluorine breaks that pattern: its bond is the shortest of all, yet it is not the strongest.
Down Group 17, each atom has its bonding electron in a shell further from the nucleus: chlorine uses 3p, bromine 4p and iodine 5p.
That has two consequences for the – bond:
Both effects weaken the attraction between the nuclei and the shared pair, so the bond enthalpy decreases: – (242) > – (193) > – (151).
The overlap of the two orbitals also gets poorer as the atoms get larger and more diffuse. Less effective overlap means a weaker bond.
Imagine two magnets holding a rubber band stretched between them. The pull is strong when the magnets are close.
Move them further apart, and the same band is held less tightly. The shared pair in a bond is like the band: the further it is from the nuclei, the weaker their grip.
Where the analogy breaks: a covalent bond is not a physical band, and the electrons are shared quantum-mechanically. The useful part is the distance idea — longer bond, weaker attraction.