Half-reaction

A half reaction (also called a half-cell reaction) is one of two parts of a redox reaction: either the oxidation piece or the reduction piece. It is written by looking at how the oxidation states of individual substances change during the reaction. Half reactions are commonly used to explain what happens at each electrode in an electrochemical cell, like a galvanic cell battery. They are also a tool for balancing redox reactions.

In metal plating and stripping, the same half reaction can describe both the metal being oxidized (at the anode) and the metal being reduced (at the cathode). For balancing redox reactions in acidic conditions, after atoms and oxidation numbers are balanced, H⁺ ions are added to balance hydrogen in the half reaction. In basic conditions, the reaction is first treated as if it were acidic, then OH⁻ ions are added to neutralize the H⁺ ions, producing H₂O.

**Example: Zn and Cu galvanic cell** A galvanic cell can be built with a piece of zinc in zinc sulfate solution and a piece of copper in copper(II) sulfate solution. The overall reaction is: Zn(s) + CuSO₄(aq) → ZnSO₄(aq) + Cu(s) At the zinc anode, oxidation occurs—the metal loses electrons. The oxidation half reaction is: Zn(s) → Zn²⁺ + 2 e⁻ At the copper cathode, reduction occurs—electrons are gained. The reduction half reaction is: Cu²⁺ + 2 e⁻ → Cu(s)

**Example: oxidation of magnesium** When magnesium ribbon burns in air, it reacts with oxygen to form magnesium oxide: 2 Mg(s) + O₂(g) → 2 MgO(s) Magnesium oxide is ionic, containing Mg²⁺ and O²⁻ ions, while the starting elements have no charge. Magnesium goes from 0 to +2, losing two electrons per atom. With two magnesium atoms, four electrons are lost in the oxidation half reaction: 2 Mg(s) → 2 Mg²⁺ + 4 e⁻ Oxygen goes from 0 to –2, gaining four electrons in the reduction half reaction: O₂(g) + 4 e⁻ → 2 O²⁻ Adding the two half reactions gives: 2 Mg(s) + O₂(g) + 4 e⁻ → 2 Mg²⁺ + 2 O²⁻ + 4 e⁻ Electrons on both sides cancel, leaving: 2 Mg(s) + O₂(g) → 2 Mg²⁺ + 2 O²⁻ The positive and negative ions then combine to form MgO due to electrostatic attraction. Every redox reaction has two half reactions: one oxidation and one reduction.

**Half-reaction balancing method** In the reaction between iron and chlorine: Cl₂ + 2 Fe²⁺ → 2 Cl⁻ + 2 Fe³⁺ Iron changes from +2 to +3, and chlorine from 0 to –1. The half reactions are: Fe²⁺ → Fe³⁺ + e⁻ Cl₂ + 2 e⁻ → 2 Cl⁻ By balancing the electrons (multiplying the iron half reaction by 2) and adding, the full reaction is obtained: 2 Fe²⁺ → 2 Fe³⁺ + 2e⁻ plus Cl₂ + 2 e⁻ → 2 Cl⁻ gives Cl₂ + 2 Fe²⁺ → 2 Cl⁻ + 2 Fe³⁺ Breaking a reaction into half reactions helps understand electron transfer and simplifies balancing—each half reaction can be balanced for atoms and charge separately.

**Basic and acidic conditions** Sometimes a half reaction must be balanced for acidic or basic conditions because the electrolyte is acidic or basic. To balance atoms and charges, H₂O, OH⁻, e⁻, and/or H⁺ are added to either side, assuming the reaction occurs in water. For the half reaction PbO₂ → PbO in basic conditions: 2 e⁻ + H₂O + PbO₂ → PbO + 2 OH⁻ For the same half reaction in acidic conditions: 2 e⁻ + 2 H⁺ + PbO₂ → PbO + H₂O Both versions are balanced for atoms and charge (overall neutral). In practice, both H⁺ and OH⁻ may be present, but they react to form water.