Halogen reactions and reactivity

For example, sodium reacts with chlorine to form sodium chloride:

$$2\mathrm{Na} + \mathrm{Cl}_2 \rightarrow 2\mathrm{NaCl}$$

Each sodium atom transfers one electron to a chlorine atom. The resulting Na⁺ and Cl⁻ ions combine in a 1:1 ratio because their charges balance.

A metal ion with a charge of +2 needs two halide ions. Magnesium bromide therefore has the formula MgBr₂: one Mg²⁺ ion balances two Br⁻ ions. Similarly, calcium reacts with bromine to form calcium bromide:

$$\mathrm{Ca} + \mathrm{Br}_2 \rightarrow \mathrm{CaBr}_2$$

Iodine also forms metal halides, called iodides. For example, sodium and iodine form sodium iodide, NaI.

Reactivity decreases down Group 7. With a comparable metal under comparable conditions, chlorine reacts more vigorously than bromine, and iodine reacts less vigorously still. The shared reaction pattern lets us predict that other halogens also form metal halides: fluorine forms fluorides, while astatine is predicted to form astatides. A less vigorous reaction does not mean that no reaction can occur.

Reactions with hydrogen form acidic solutions

Halogens also react with hydrogen, but the products are covalent molecular compounds called hydrogen halides:

hydrogen + halogen → hydrogen halide

Chlorine forms hydrogen chloride, HCl; bromine forms hydrogen bromide, HBr; and iodine forms hydrogen iodide, HI. For chlorine:

$$\mathrm{H}_2 + \mathrm{Cl}_2 \rightarrow 2\mathrm{HCl}$$

These hydrogen halides dissolve in water to form acidic solutions containing H⁺ ions. Hydrogen chloride gas dissolved in water forms hydrochloric acid. Hydrogen bromide and hydrogen iodide likewise form acidic solutions.

The reactions with hydrogen become less vigorous down the group. Chlorine can react with hydrogen in sunlight, whereas bromine requires a higher temperature. Iodine reacts less vigorously than chlorine.

The common pattern predicts that other halogens also form hydrogen halides which give acidic solutions in water. For example, fluorine forms hydrogen fluoride. Predicting an acidic solution does not mean that every hydrogen halide produces an acid of the same strength.

Displacement reactions reveal the reactivity order

A more reactive halogen displaces a less reactive halogen from an aqueous solution of its halide. Here, aqueous means dissolved in water.

For example, adding chlorine solution to potassium bromide solution produces potassium chloride and bromine:

chlorine + potassium bromide → potassium chloride + bromine

$$\mathrm{Cl}_2 + 2\mathrm{KBr} \rightarrow 2\mathrm{KCl} + \mathrm{Br}_2$$

The solution becomes orange as bromine forms. Chlorine also displaces iodine from potassium iodide solution, producing a brown solution containing iodine:

$$\mathrm{Cl}_2 + 2\mathrm{KI} \rightarrow 2\mathrm{KCl} + \mathrm{I}_2$$

Bromine displaces iodine from potassium iodide, but cannot displace chlorine from a chloride:

$$\mathrm{Br}_2 + 2\mathrm{KI} \rightarrow 2\mathrm{KBr} + \mathrm{I}_2$$

The outcomes can be summarised as follows. Each row names the halogen added; each column names the ions in the dissolved salt.

These reactions establish the order chlorine > bromine > iodine in reactivity. Astatine lies below iodine, so it is predicted to be less reactive than all three. Chlorine, bromine and iodine should therefore displace astatine from an aqueous astatide solution; astatine should not displace any of them from their halide solutions.

Why reactivity decreases down Group 7

Halogens react by gaining an electron. Their relative reactivity depends on how strongly their nuclei attract an incoming electron.

Fluorine has the electronic configuration 2,7, while chlorine has 2,8,7. Both have seven outer electrons, but chlorine has an extra occupied shell. Moving further down Group 7 adds more occupied shells.

Down the group, the outer shell is further from the positively charged nucleus, and there are more inner electrons to shield an incoming electron from its attraction. Together, the greater distance and increased shielding weaken the attraction for an electron entering the outer shell. An electron is gained less easily, so reactivity decreases.

This explains both the less vigorous reactions down the group and why chlorine can displace bromine and iodine, rather than the reverse.

Higher tier: displacement involves electron transfer

Halogen displacement reactions are redox reactions because oxidation and reduction happen together. Oxidation is loss of electrons; reduction is gain of electrons.

Consider chlorine displacing bromine from potassium bromide. The potassium ions remain unchanged, so they are spectator ions. Leaving them out gives the ionic equation:

$$\mathrm{Cl}_2 + 2\mathrm{Br}^- \rightarrow 2\mathrm{Cl}^- + \mathrm{Br}_2$$

Chlorine gains electrons to become chloride ions, so chlorine is reduced:

$$\mathrm{Cl}_2 + 2\mathrm{e}^- \rightarrow 2\mathrm{Cl}^-$$

Bromide ions lose electrons to become bromine, so bromide ions are oxidised:

$$2\mathrm{Br}^- \rightarrow \mathrm{Br}_2 + 2\mathrm{e}^-$$

The two electrons lost by the bromide ions are gained by chlorine. The same reasoning applies to the other displacement reactions: the added, more reactive halogen is reduced, while the halide ions of the less reactive halogen are oxidised.

Chlorine > bromine > iodine > astatine in predicted reactivity.

• Chlorine displaces bromine and iodine from their aqueous halides.
• Bromine displaces iodine, but not chlorine.
• Iodine displaces neither chlorine nor bromine.
• Astatine is predicted not to displace any of these; all three should displace astatine from astatides.

Electronic explanation

Fluorine: 2,7; chlorine: 2,8,7. Both have seven outer electrons.

Down Group 7: more occupied shells → greater distance and shielding → weaker attraction for an incoming electron → electron gained less easily → lower reactivity.

Higher tier: redox

OIL RIG: oxidation is loss; reduction is gain of electrons.

$${\mathrm{C}\mathrm{l}}_2+2{\mathrm{B}\mathrm{r}}^{-}\to 2{\mathrm{C}\mathrm{l}}^{-}+{\mathrm{B}}_{}$$