Halogens: Disproportionation and Halide Ions

Learning Objectives (OCR Spec 3.1.3)

By the end of this lesson you should be able to:

• Define disproportionation and recognise it from oxidation numbers
• Write equations for the disproportionation of chlorine in water and chlorine in cold dilute NaOH
• Describe the use of chlorine in water treatment and discuss the ethics/benefits
• Describe and explain the trend in reducing ability of halide ions (F- → I-)
• Link the trend in halide reducing ability to atomic radius and shielding

Disproportionation: Definition

Disproportionation is a redox reaction in which the same element is simultaneously oxidised and reduced. One species is split into two products with different oxidation numbers - one higher and one lower than the starting number.

Example: Hydrogen Peroxide

2H2O2 → 2H2O + O2 (disproportionation)

Oxygen starts at -1 (in peroxide); ends as -2 in H2O (reduced) and 0 in O2 (oxidised). Same element, same reaction, both oxidised and reduced.

Chlorine + Water (Water Treatment)

Chlorine dissolves slightly in water and undergoes disproportionation:

Cl2(aq) + H2O(l) → HClO(aq) + HCl(aq)

Or in ionic form:

Cl2(aq) + H2O(l) → HClO(aq) + H+(aq) + Cl-(aq)

Oxidation numbers of chlorine:

• Cl2: 0
• HClO (chloric(I) acid): +1 ← oxidised
• HCl / Cl-: -1 ← reduced

Half of the chlorine is oxidised (0 → +1) and half is reduced (0 → -1). This is disproportionation.

Why This Matters: Killing Pathogens

HClO (chloric(I) acid) is a powerful oxidising agent that kills bacteria and viruses by oxidising essential cellular components. Adding a tiny amount of chlorine to drinking water and swimming pools provides continuous sterilisation. The equilibrium maintains a small concentration of HClO, which is effective even at ~1 ppm.

Risks and Benefits

Benefit	Risk
Kills waterborne pathogens (cholera, typhoid)	Cl2 is toxic at high concentration
Easy to produce and store as a liquid	Can react with organic matter to form chlorinated hydrocarbons (some are carcinogenic)
Long-lasting effect in pipes	Taste/smell can be unpleasant
Millions of lives saved worldwide	Historical use as a chemical weapon (WWI)

The consensus is that the benefits massively outweigh the risks for public health.

Chlorine + Cold Dilute NaOH (Bleach Production)

At room temperature with cold, dilute sodium hydroxide:

Cl2(aq) + 2NaOH(aq) → NaCl(aq) + NaClO(aq) + H2O(l)

Oxidation numbers of chlorine:

• Cl2: 0
• NaCl: -1 (reduced)
• NaClO (sodium chlorate(I)): +1 (oxidised)

Again this is disproportionation. The product mixture is household bleach - a solution of NaClO (the active ingredient) and NaCl. NaClO is used as a disinfectant and bleaching agent.

Conditions matter: with hot concentrated NaOH, chlorine disproportionates further to chlorate(V):

3Cl2 + 6NaOH(hot conc) → 5NaCl + NaClO3 + 3H2O

But this is not required for OCR AS-level - focus on the cold dilute reaction.

Reaction Tree for Chlorine

graph TD
  A[Cl2 gas] --> B[+ H2O]
  A --> C[+ cold dilute NaOH]
  A --> D[+ Na metal]
  B --> B1[HClO + HCl disproportionation, water treatment]
  C --> C1[NaCl + NaClO + H2O bleach, disproportionation]
  D --> D1[2NaCl simple redox, no disproportionation]

Reducing Ability of Halide Ions: Opposite Trend

We have seen that oxidising ability decreases down Group 7 (F2 > Cl2 > Br2 > I2). The mirror-image trend is the reducing ability of the halide ions:

Reducing ability of halide ions increases down the group: I- > Br- > Cl- > F-

A reducing agent loses electrons; halide ions lose an electron to become a halogen atom:

2X-(aq) → X2 + 2e- (oxidation)