Topic 8 · page 25 of 29

Halogens with metals, and chlorine with water

Halogens oxidise Group 1 and 2 metals to ionic halides, and the reactions are most vigorous for chlorine. Chlorine reacts with water by disproportionation to give HCl and chloric(I) acid, which is why it is used to treat water.

12 min read8.268.26(i)8.26(ii)

Read first: Colours of the halogens and displacement reactions

Oxidation of Group 1 and 2 metals 8.268.26(i)

The halogens are strong oxidising agents: they take electrons from metals. A Group 1 metal gives an ionic halide MX, and a Group 2 metal gives MX₂, where X is a halogen.

2 M(s)+XX2→2 MX(s)M(s)+XX2→MXX2(s) \ce{2M(s) + X2 -> 2MX(s)} \qquad \ce{M(s) + X2 -> MX2(s)}
Reaction Oxidation number changes
2 Na(s)+ClX2(g)→2 NaCl(s)\ce{2Na(s) + Cl2(g) -> 2NaCl(s)} Na 0 → +1 (oxidised); Cl 0 → −1 (reduced)
2 K(s)+BrX2(l)→2 KBr(s)\ce{2K(s) + Br2(l) -> 2KBr(s)} K 0 → +1; Br 0 → −1
Mg(s)+IX2(s)→MgIX2(s)\ce{Mg(s) + I2(s) -> MgI2(s)} Mg 0 → +2; I 0 → −1
Ca(s)+ClX2(g)→CaClX2(s)\ce{Ca(s) + Cl2(g) -> CaCl2(s)} Ca 0 → +2; Cl 0 → −1

In each, the metal is oxidised (its oxidation number increases) and is the reducing agent, and the halogen is reduced and is the oxidising agent. The products are white ionic solids.

The reaction is most vigorous with chlorine and least with iodine, in line with the oxidising power of the halogens: Cl₂ > Br₂ > I₂. Hot sodium burns brightly in chlorine to give white sodium chloride.

Worked exampleWriting an equation by pattern

Write the equation for magnesium with bromine, and state the oxidation number changes.

  1. Magnesium is Group 2, so Mg²⁺. Bromide is Br⁻. Two Br⁻ for each Mg²⁺: MgBr₂.
  2. Mg(s)+BrX2(l)→MgBrX2(s)\ce{Mg(s) + Br2(l) -> MgBr2(s)}
  3. Mg: 0 → +2 (oxidised). Br: 0 → −1 (reduced).

Chlorine with water 8.268.26(ii)

Chlorine dissolves in water, and some of it reacts. The reaction is reversible.

ClX2(aq)+HX2O(l)⇌HCl(aq)+HClO(aq) \ce{Cl2(aq) + H2O(l) <=> HCl(aq) + HClO(aq)}

In ionic form: ClX2(aq)+HX2O(l)⇌2 HX+(aq)+ClX−(aq)+ClOX−(aq)\ce{Cl2(aq) + H2O(l) <=> 2H+(aq) + Cl-(aq) + ClO-(aq)}, with the weak acid HClO only partly ionised.

The oxidation numbers show the reaction type:

Species Oxidation number of Cl Change
Cl₂ 0 starting point
HCl (Cl⁻) −1 reduced
HClO (ClO⁻) +1 oxidised

One element in one species, chlorine in Cl₂, is both oxidised and reduced: disproportionation. HClO is named chloric(I) acid and its ion, ClO⁻, is chlorate(I).

Chlorine in water treatment 8.268.26(ii)

Chlorate(I) ions and chloric(I) acid kill bacteria. Chlorine is added to drinking water and to swimming pools so that disease-causing bacteria are killed and the water is safe.

Point
Benefit Kills bacteria: prevents water-borne diseases such as cholera and typhoid. Cheap and effective. Some stays in the water to keep it safe afterwards.
Risks Chlorine gas is toxic: it is hazardous to store and handle. It reacts with organic matter in the water to form chlorinated organic compounds, some of which are harmful. It can give water an unpleasant taste and smell.

Check yourself

  1. Write equations for (a) potassium with chlorine, (b) calcium with bromine, (c) magnesium with iodine.
  2. For the reaction in (a), state the oxidation numbers and name the oxidising agent.
  3. Write the equation for chlorine with water, and give the oxidation numbers of chlorine in each chlorine-containing species.
  4. Explain why the reaction of chlorine with water is a disproportionation.
  5. Name the species in water that kills bacteria, and give one risk and one benefit of using chlorine to treat drinking water.
  6. Explain why sodium reacts more vigorously with chlorine than with iodine.

Answers cover these until you have tried the questions

  1. (a) 2 K(s)+ClX2(g)→2 KCl(s)\ce{2K(s) + Cl2(g) -> 2KCl(s)} (b) Ca(s)+BrX2(l)→CaBrX2(s)\ce{Ca(s) + Br2(l) -> CaBr2(s)} (c) Mg(s)+IX2(s)→MgIX2(s)\ce{Mg(s) + I2(s) -> MgI2(s)}
  2. K: 0 → +1 (oxidised); Cl: 0 → −1 (reduced). The oxidising agent is chlorine (Cl₂).
  3. ClX2+HX2O⇌HCl+HClO\ce{Cl2 + H2O <=> HCl + HClO}. Cl in Cl₂ = 0; in HCl = −1; in HClO = +1 (H +1, O −2).
  4. Chlorine, from one species (Cl₂), goes both down (0 → −1, reduced) and up (0 → +1, oxidised) at the same time.
  5. HClO (chloric(I) acid) or ClO⁻ (chlorate(I)). Benefit: kills disease-causing bacteria. Risk: chlorine is toxic, or can form chlorinated organic compounds.
  6. Chlorine is a stronger oxidising agent than iodine. Its atoms attract an incoming electron more strongly, because the electron goes into a shell that is closer to the nucleus and less shielded, so it takes electrons from sodium more readily.