What Is Electrolysis? Using Electricity to Drive Chemical Change

An electrolytic cell contains an electrolyte and two electrodes connected to an external power source. The electrolyte may be a molten ionic compound or a solution containing ions that can move. The power source pushes electrons through the external circuit and creates conditions that drive chemical reactions at the electrode surfaces. The electrical energy is converted into chemical change.

This is the core of electrolysis. Positive ions, called cations, are attracted toward the cathode, where reduction occurs because particles gain electrons. Negative ions, called anions, move toward the anode, where oxidation occurs because particles lose electrons. Remembering reduction at the cathode and oxidation at the anode helps organise many electrolysis questions.

The electrodes provide surfaces where electron transfer can happen, but they are not always passive. Inert electrodes such as graphite or platinum are chosen when the aim is to avoid the electrode material reacting significantly. Active electrodes can take part in the chemistry. The products therefore depend on the electrolyte, the electrode materials, and whether the substances are molten or dissolved in water.

In a molten ionic compound, the ions from that compound are free to move. For molten sodium chloride, sodium ions can be reduced at the cathode and chloride ions can be oxidised at the anode. In a solid crystal, the ions are locked in place, so the material does not conduct in the same way and electrolysis cannot proceed through mobile ions.

Aqueous solutions are more complicated because water contributes additional possible reactions. If several ions could be discharged at an electrode, their relative tendency to gain or lose electrons matters, along with concentration and electrode material. This is why electrolysis in solution cannot always be predicted by simply choosing the positive ion at the cathode and the negative ion at the anode.

Electrolysis has many practical uses. It can extract reactive metals from molten compounds, purify metals such as copper, coat objects by electroplating, and produce industrial chemicals. In each case, the useful product depends on controlling which oxidation and reduction reactions occur. The process can require substantial electrical energy, so energy cost and electricity source matter in large-scale applications. A good way to solve electrolysis problems is to separate the two electrodes. Write the ions present, decide which species can be reduced at the cathode, and decide which can be oxidised at the anode. Then write half-equations with electrons. This step-by-step approach makes electrolysis a pair of electron-transfer processes rather than one confusing reaction happening everywhere at once. Checking charge balance is a useful final step. The electrons lost in oxidation and gained in reduction must fit the half-equations, and the products should make sense for the ions and electrodes present. This keeps electrolysis connected to the conservation of charge as well as to chemical change.

Concentrated sodium chloride solution, called brine, is one of the most important electrolytes in industry. The solution contains sodium ions and chloride ions from the salt, along with hydrogen ions and hydroxide ions from water. At the cathode, sodium is too reactive to be discharged from solution, so hydrogen is reduced instead: 2H+ + 2e- gives H2. At the anode, the high concentration of chloride ions means chlorine is usually produced: 2Cl- gives Cl2 + 2e-. Bubbles of both gases appear at their electrodes.

What remains in solution is sodium ions and hydroxide ions, so the liquid gradually becomes sodium hydroxide solution. One process therefore makes three useful products: hydrogen, chlorine for disinfecting water and making plastics such as PVC, and sodium hydroxide for soap, paper and many other industries. The example also shows why concentration matters. In very dilute salt solution, more oxygen and less chlorine forms at the anode, because the chloride ions are no longer so dominant at the electrode surface.

Electrolysis reshaped chemistry soon after batteries were invented around 1800. In 1807 Humphry Davy passed current through molten caustic potash and saw tiny globules of metal form, some bursting into flame. He had isolated potassium, and he soon isolated sodium the same way. No chemical reaction then available was strong enough to pull these reactive metals out of their compounds, but electricity could. His assistant Michael Faraday later worked out how the amount of substance produced depends on the quantity of electricity, and, with the scholar William Whewell, introduced words such as electrode, anode, cathode, ion and electrolyte.

Aluminium shows the economic impact. It is the most abundant metal in Earth's crust, yet in the mid-nineteenth century it was costly to produce and was treated almost as a precious metal; a small aluminium pyramid was placed on top of the Washington Monument in 1884. In 1886 Charles Martin Hall in the United States and Paul Heroult in France independently found a practical method. Aluminium oxide melts at over 2,000 degrees Celsius, but it dissolves in molten cryolite at a much lower temperature, and that solution can be electrolysed. Aluminium soon became cheap enough for everyday use.

Faraday's discovery makes electrolysis quantitative. The charge passed equals current multiplied by time, and one mole of electrons carries about 96,500 coulombs, a quantity called the Faraday constant. Suppose you electroplate a key with copper using a current of 2.0 amps for 30 minutes. The charge is 2.0 times 1,800 seconds, which is 3,600 coulombs, or 3,600 divided by 96,500, about 0.0373 moles of electrons. Each copper ion needs two electrons, since Cu2+ + 2e- gives Cu, so about 0.0187 moles of copper are deposited. With a molar mass of 63.5 grams per mole, that is roughly 1.2 grams.

The same counting explains gas volumes. In the electrolysis of water, making one molecule of hydrogen takes two electrons, but making one molecule of oxygen takes four, because 2H2O gives O2 + 4H+ + 4e-. For the same charge, therefore, twice as many hydrogen molecules form as oxygen molecules. Since equal numbers of gas molecules occupy equal volumes at the same temperature and pressure, the hydrogen tube fills twice as fast. In practice the measured ratio can drift slightly, partly because oxygen is a little more soluble in water than hydrogen and some can react with the electrodes.

Copper for electrical wiring must be extremely pure, because even small amounts of impurity noticeably reduce how well it conducts. Refineries use electrolysis. A slab of impure copper is made the anode and a thin sheet of pure copper, or a steel blank, the cathode, both dipping into copper sulfate solution with sulfuric acid. Copper atoms at the anode lose electrons and enter the solution as ions, then gain electrons at the cathode and build up as copper of about 99.99 percent purity. More reactive impurities such as iron and zinc stay dissolved, while unreactive metals such as silver and gold fall to the bottom as anode sludge.

That sludge is valuable, and selling the silver, gold and other precious metals it contains helps pay for the refining. Anodising runs electrolysis the other way round on purpose. An aluminium object is made the anode in dilute sulfuric acid, so oxygen forms at its surface and combines with the metal, thickening its thin natural oxide layer from a few nanometres to many micrometres. The new layer is hard and protective, and because it is full of tiny pores it can soak up dyes before being sealed in hot water. That is how coloured aluminium laptops, bicycle parts and climbing gear get their tough, bright finishes.

Electrolysis is the use of electrical energy to drive chemical reactions in an electrolyte. Cations move toward the cathode, where reduction occurs, while anions move toward the anode, where oxidation occurs. Keep the two electrodes separate, identify the possible particles, and track electrons carefully. That turns electrolysis into a clear application of redox chemistry.