As all halogens are capable of accepting electrons and being reduced they are all oxidising agents:

• F₂(g) + 2e ? 2F^–(aq)
• Cl₂(g) + 2e ? 2Cl^–(aq)
• Br₂(l) + 2e ? 2Br^–(aq)
• I₂(aq) + 2e ? 2I^–(aq)

As the number of shells in an atom decreases so the distance between the electrons and nucleus also decreases and the electrons are less shielded from the nucleus. As the attractions between the electrons and the nucleus is stronger the chances that the atom will accept electrons is higher.

The halides – reducing agents

As halide ions can lose electrons and get oxidised they are known as reducing agents:

• 2Fl^–(aq) ? F₂(g) + 2e
• 2Cl^–(aq) ? Cl₂(g) + 2e
• 2Br^–(aq) ? Br₂(aq) + 2e
• 2I^–(aq) ? I₂(aq) + 2e

As the number of shells in the ion rises so does the shielding of the nucleus meaning that the outer electrons are not held as tightly. This means that these electrons can be lose more easily making the halide more easily oxidised.

Displacement reactions between halogens and halides

Displacement reactions between halogens and halides illustrate their oxidising and reducing properties in an aqueous solution.

The halogens which are the strongest oxidising agents will displace the halides which are the strongest reducing agents from solutions of their ions:

• chlorine will displace bromide ions producing an orange colour
• chlorine will displace iodide ions producing a brown colour
• bromine will displace iodide ions producing a brown colour
• bromine will not chloride ions
• iodine will not displace bromide or chloride ions

Reactions of halide ions with concentrated sulphuric acid

Mixing sodium halides with concentrated sulphuric acid illustrates the variation in reducing halide strength.

Concentrated sulphuric acid is a strong acid. It is able to convert sodium salts of halides into hydrogen halides:

• H₂SO₄(l) + NaX(s) ? NaHSO₄(s) + HX(g)
• or H2SO₄(l) + X- ? HSO₄^– + HX(g)

Instead of being oxidised the halide ions act as bases.

Concentrated sulphuric acid is also an oxidising agent and can be reduce to SO2, S or H₂S:

• H₂SO₄ + 2H⁺ + 2e ? SO₂ + 2H₂O (reduction of S from +6 to 4)
• H₂SO₄ + 6H⁺ + 6e ? S + 4H₂O (reduction of S from +6 to 0)
• H₂SO₄ + 8H⁺ + 8e ? H₂S + 4H₂O (reduction of A from +6 to -2)

As Cl- is not a strong reducing agent H₂SO₄ does not oxidise it. Instead an acid-base reaction occurs in which HCl gas is produced:

• H₂SO₄ + Cl^– ? HSO₄^– + HCl

HCl is visible as white fumes which turn blue litmus paper red.

Br- is a better reducing agent and is oxidised. However, the sulphur in the H₂SO₄ is reduced only from +6 to +4 (SO₂). It is also possible for an acid-base reaction to occur:

• Acid-base reaction: H₂SO₄ + Br^– ? HSO + HBr
• Redox reaction: H₂SO₄ + 2H⁺ + 2Br^– ? SO₂ + Br₂ + 2H₂O

Another good reducing agent is I- which is also oxidised. It reduces the sulphur in H₂SO₄ from +6 to +4 (SO₂), 0 (S) or -2 (H₂S). It is also possible for an acid-base reaction to occur:

• Acid-base reaction: H₂SO₄ + I^– ? HSO₄^– + HI
• Redox reaction: H₂SO₄ + 6H⁺ + 6I^– ? S + 3I₂ + 4H₂O
• Redox reaction: H₂SO₄ + 8H⁺ + 8I^– ? H₂S + 4I₂ + 4H₂O

White fumes emerge which change blue litmus paper red as well as the purple colour of iodine vapour and the rotten egg smell of H₂S.