Electrochemistry: electrolysis and cells made clear

Electrochemistry looks intimidating because it has two halves that seem to work in opposite directions, electrolysis pushes electricity into a chemical to break it apart, while a chemical cell lets a reaction produce electricity. Once you see that both are simply redox reactions with the electrons routed through wires, the whole topic becomes one clear picture. This explainer draws that picture from the ground up.

Electrolysis: forcing a reaction with electricity

Electrolysis is the decomposition of an ionic compound, in molten or aqueous form, by passing electricity through it. The liquid that conducts is the electrolyte, and it conducts because its ions are now free to move. Two rods called electrodes dip into it: the cathode is connected to the negative terminal of the power supply, and the anode to the positive terminal. This is the heart of the redox equilibrium chapter, and it runs on one simple rule of attraction.

Opposite charges attract, so:

• Cations (positive ions) move to the cathode, gain electrons, and are reduced.
• Anions (negative ions) move to the anode, lose electrons, and are oxidised.

That is genuinely all the mechanism you need. Everything else is deciding which ion reacts when there is a choice.

Molten compounds: the simple case

When the electrolyte is a molten compound, there is no water and therefore no competition, only the compound’s own ions are present. Molten lead(II) bromide is the classic example. Lead(II) ions go to the cathode and bromide ions to the anode:

• Cathode: Pb²⁺ + 2e⁻ → Pb (a grey bead of molten lead forms)
• Anode: 2Br⁻ → Br₂ + 2e⁻ (brown bromine vapour is released)

The overall change is PbBr₂ → Pb + Br₂. Notice this is a redox reaction split across two electrodes: reduction at the cathode, oxidation at the anode, every time.

Aqueous solutions: why “selective discharge” appears

Water changes everything. In an aqueous solution the water itself supplies a small amount of H⁺ and OH⁻ ions, so at each electrode there are now two candidates competing to be discharged. Which one wins is decided by three factors you must weigh in order:

1. Position in the electrochemical series. The ion of the less reactive element is discharged more readily. At the cathode this means ions like Cu²⁺ or H⁺ are favoured over reactive ones like Na⁺; at the anode, OH⁻ is discharged in preference to sulfate.
2. Concentration of ions. A high concentration of a particular ion can override the series. This is why concentrated sodium chloride gives chlorine at the anode even though OH⁻ would otherwise win.
3. Type of electrode. Usually the electrodes are inert (carbon or platinum) and take no part. But a reactive electrode such as copper can itself dissolve at the anode instead of an ion being discharged.

Work through the standard examples under electrolysis of aqueous solutions and the pattern soon becomes second nature. For dilute sulfuric acid, for instance, you effectively electrolyse water: hydrogen at the cathode, oxygen at the anode, in a 2 : 1 ratio by volume.

Copper electrodes: a special, useful twist

Electrolysing copper(II) sulfate with copper electrodes shows factor three in action and explains electroplating and copper purification. The reactive copper anode dissolves, Cu → Cu²⁺ + 2e⁻, while at the cathode copper is deposited, Cu²⁺ + 2e⁻ → Cu. Copper is effectively transferred from the anode to the cathode, which is exactly how impure copper is refined and how objects are electroplated with a thin, even copper layer.

Chemical cells: the reaction running in reverse

A chemical (voltaic) cell does the opposite of electrolysis: instead of using electricity to force a reaction, it uses a spontaneous redox reaction to produce electricity. Dip two different metals into an electrolyte and connect them with a wire, and a current flows. The reason is reactivity difference.

The more reactive metal loses electrons more readily, so it is oxidised and becomes the negative terminal. Those electrons travel through the external wire to the less reactive metal, the positive terminal, doing electrical work on the way. In a zinc–copper cell, zinc is the negative terminal and slowly dissolves, while copper is the positive terminal. You can build one yourself following constructing a simple voltaic cell.

There is a neat, examinable consequence: the further apart the two metals are in the reactivity series, the larger the voltage the cell produces. A magnesium–copper cell gives a higher reading than a zinc–copper cell because magnesium and copper are further apart.

Holding the two ideas together

The cleanest way to keep electrolysis and cells straight is to remember their direction of energy. Electrolysis: electrical energy → chemical change, driven by an outside supply, used to break compounds down or coat metals. Chemical cell: chemical change → electrical energy, spontaneous, used to generate current. Both are redox at their core, oxidation at one electrode, reduction at the other, only the source of the push is different.

Making it click

Do not memorise long lists of electrode products. Instead, for any electrolysis question, name the ions present, send cations to the cathode and anions to the anode, then apply the three selective-discharge factors in order. For any cell, find the more reactive metal and call it negative. Practise five or six examples this way and the topic settles into place. If the selective-discharge rules still feel slippery, a focused online one-to-one lesson can walk you through them with worked cases, our teachers teach in English from RM50 an hour, with a paid one-hour trial to begin.