Topic G: Redox Reactions, Chemical Cells and Electrolysis

This part treats reactions as transfers of electrons. It defines oxidation and reduction, assigns oxidation numbers to track them, and shows how redox underlies displacement reactions, the chemical cells that generate electricity, and electrolysis, which uses electricity to drive reactions and to deposit metals by electroplating.

Oxidation, reduction and electron transfer

In modern terms, oxidation is the loss of electrons and reduction is the gain of electrons, summarised by the mnemonic OIL RIG (Oxidation Is Loss, Reduction Is Gain). The two always occur together, since electrons lost by one species are gained by another, making a redox reaction. The substance that causes oxidation by accepting electrons is the oxidising agent and is itself reduced; the substance that causes reduction by donating electrons is the reducing agent and is itself oxidised. For example, when magnesium reacts with oxygen, the magnesium loses electrons (oxidised) and the oxygen gains them (reduced); magnesium is the reducing agent and oxygen the oxidising agent. Identifying these roles is fundamental to all of electrochemistry.

Oxidation numbers

Oxidation numbers are bookkeeping values that track electron transfer, assigned by rules: an uncombined element is 0, a simple ion equals its charge, hydrogen is usually +1 and oxygen usually -2, and the values in a neutral compound sum to zero while in an ion they sum to the ionic charge. A rise in oxidation number signals oxidation, and a fall signals reduction. For example, in the change from Fe2+ to Fe3+ the oxidation number rises by one, so iron is oxidised. Oxidation numbers are especially useful for reactions with no obvious oxygen or hydrogen transfer, and for naming compounds such as iron(II) and iron(III), where the number states the oxidation state of the metal.

Redox in displacement and tests

Many familiar reactions are redox. In metal displacement, a more reactive metal donates electrons to the ion of a less reactive metal: Zn + Cu2+ -> Zn2+ + Cu, where zinc is oxidised and copper(II) reduced. Halogen displacement works the same way, with a more reactive halogen displacing a less reactive one from solution, for example Cl2 + 2 KBr -> 2 KCl + Br2. Useful tests rely on colour changes: acidified potassium permanganate (purple) is decolourised as it is reduced by a reducing agent, and acidified potassium dichromate turns from orange to green when reduced. These colour changes let you detect oxidising and reducing agents and follow the direction of electron transfer.

Simple and galvanic chemical cells

A chemical cell converts the energy of a spontaneous redox reaction into electrical energy. In a simple cell, two different metals dip into an electrolyte; the more reactive metal becomes the negative electrode, losing electrons that flow through the external circuit to the less reactive positive electrode. The bigger the reactivity difference, the larger the voltage. A galvanic cell separates the two half-reactions into half-cells joined by a salt bridge, which completes the circuit and keeps each solution electrically neutral. Oxidation occurs at the negative anode and reduction at the positive cathode, with electrons flowing through the wire from anode to cathode. The salt bridge prevents the build-up of charge that would otherwise stop the reaction.

Electrolysis of molten and aqueous compounds

Electrolysis uses electrical energy to drive a non-spontaneous redox reaction, decomposing a molten or dissolved ionic compound. The compound is the electrolyte; positive ions (cations) move to the negative cathode and are reduced, while negative ions (anions) move to the positive anode and are oxidised. Electrolysing molten lead(II) bromide gives lead at the cathode and bromine at the anode. In aqueous solutions, water can also be discharged, so the products depend on the relative ease of discharge of the ions present: at the cathode, less reactive metals or hydrogen are released; at the anode, halides or oxygen. Electrolysis of brine, for instance, gives hydrogen, chlorine and sodium hydroxide.

Electroplating and applications

Electroplating uses electrolysis to coat one metal with a thin layer of another, for protection or appearance. The object to be plated is made the cathode, the plating metal is the anode, and the electrolyte is a solution containing ions of the plating metal. As current flows, metal ions are reduced and deposited onto the cathode, while the anode dissolves to replenish the ions, so silver-plating uses a silver anode and a silver salt solution. Electrolysis is also used industrially to extract reactive metals such as aluminium, to purify copper, and to produce chlorine and sodium hydroxide from brine. These applications show how controlling electron transfer turns electrical energy into valuable chemical products.