Group 2 reactivity and reducing power, reactions of Group 2 elements and their oxides and hydroxides, the halogens as oxidising agents, halide displacement, disproportionation of chlorine, and tests for halide ions.

What this topic is asking

OCR specification point 3.1.2 wants you to describe the reactions and reactivity trends of the Group 2 elements (as reducing agents) and of the halogens (as oxidising agents). You must cover Group 2 reactions with water and oxygen, the use of their oxides and hydroxides, halogen-halide displacement, the disproportionation of chlorine, and the silver nitrate test for halide ions. These trends all follow from how easily electrons are lost or gained.

Group 2: reactivity as reducing agents

The metals react with water to form a hydroxide and hydrogen, the reaction becoming more vigorous down the group:

Group 2 oxides and hydroxides

The oxides react with water to form alkaline hydroxide solutions, and hydroxide solubility increases down the group, so the solutions become more strongly alkaline. Group 2 compounds have practical uses as bases: $\text{Mg(OH)}_2$ in antacids and $\text{Ca(OH)}_2$ (lime) to raise the pH of acidic soils and to treat flue gases.

The halogens: reactivity as oxidising agents

A more reactive halogen displaces the ions of a less reactive one from solution. For example, chlorine displaces bromine:

Disproportionation of chlorine

Testing for halide ions

Adding dilute nitric acid then silver nitrate solution gives a precipitate whose colour identifies the halide: chloride white, bromide cream, iodide yellow. The colours are confirmed by solubility in ammonia: $\text{AgCl}$ dissolves in dilute ammonia, $\text{AgBr}$ only in concentrated ammonia, and $\text{AgI}$ in neither. The nitric acid first removes carbonate ions, which would otherwise give a false positive.

Examples in context

Example 1. Water treatment. The disproportionation of chlorine in water produces chloric(I) acid, the active disinfectant, balancing the benefit of clean water against the toxicity of chlorine itself, a classic application of redox.

Example 2. Agricultural lime. Farmers spread calcium hydroxide to neutralise acidic soils; because hydroxide solubility increases down Group 2, the choice of base is tuned to how much alkalinity is needed.

Try this

Q1. Write an equation for the reaction of strontium with water. [1 mark]

• Cue. $\text{Sr} + 2\text{H}_2\text{O} \rightarrow \text{Sr(OH)}_2 + \text{H}_2$.

Q2. Explain why chlorine displaces iodide ions but iodine does not displace chloride ions. [2 marks]

• Cue. Chlorine is a stronger oxidising agent than iodine, so it can take electrons from iodide, but iodine is too weak an oxidiser to remove electrons from chloride.