Electrochemistry

Electrolysis: principles

Electrolysis is chemistry driven by an external voltage. A cell has three parts you must be able to name — two electrodes and one electrolyte — and every product depends on which ion ends up at which electrode.

Why solids don’t work

An ionic compound in the solid state has its ions locked in a lattice; they cannot move, so no current flows. Melt it or dissolve it and the ions are free to drift, and electrolysis becomes possible. Covalent compounds have no ions at all — they cannot be electrolysed regardless of state.

Building the cell

Two electrodes dip into the electrolyte and connect to a d.c. power supply. The electrode wired to the negative terminal becomes the cathode; the one wired to the positive terminal becomes the anode. Once the current flows, the two electrodes pull ions of opposite charge toward themselves.

Remembering which is which

The mnemonic that keeps the signs straight is PANIC: Positive is Anode, Negative Is Cathode. Then match ions to the opposite sign — cations go to the cathode, anions go to the anode.

Electrolysis of molten compounds

Molten electrolysis is the clean case. Only two ions in the electrolyte — one metal cation, one non-metal anion — and each one heads for the opposite electrode. No water, no competition.

The rule

The cation, always a metal, is attracted to the cathode; there it gains electrons and forms the neutral metal, which either plates onto the electrode or falls as a solid. The anion, a non-metal, is attracted to the anode; it loses electrons and forms the neutral element — usually a gas.

Worked example: molten lead(II) bromide

Predict the products at each electrode when molten PbBr₂ is electrolysed with graphite electrodes. The ions present are Pb²⁺ and Br⁻. Pb²⁺ is the cation, so it migrates to the cathode and gains 2 electrons to form molten lead metal, which sinks below the electrolyte. Br⁻ is the anion, so it migrates to the anode and loses an electron; two bromine atoms pair up as Br₂, which bubbles off as a red-brown gas. The graphite electrodes are inert — the reactions happen on their surface but the electrodes themselves do not react. Molten KCl works the same way: potassium metal at the cathode, chlorine gas at the anode.

Electrolysis of aqueous solutions

Once water enters the picture there are four ions in the electrolyte, not two — the ones from the dissolved compound, plus H⁺ and OH⁻ from water itself. Every product depends on which ion wins the race to each electrode.

Concentrated aqueous NaCl (brine)

Brine contains Na⁺, Cl⁻, H⁺ and OH⁻. At the cathode, H⁺ is discharged in preference to Na⁺ (sodium is far too reactive), and hydrogen gas bubbles off. At the anode, Cl⁻ is discharged in preference to OH⁻, and chlorine gas is released. The Na⁺ and OH⁻ are left behind — the solution around the cell becomes sodium hydroxide.

Dilute sulfuric acid

Dilute H₂SO₄ contains H⁺, SO₄²⁻ and OH⁻. At the cathode, H⁺ is discharged as hydrogen. At the anode, OH⁻ is discharged as oxygen — SO₄²⁻ is too stable to discharge. In a Hofmann voltameter you can measure the ratio: twice as much hydrogen is collected as oxygen, matching the 2 : 1 ratio in H₂O. Overall, water is being split.

Testing the gases

ExtendedPredicting any aqueous case

At the cathode, the less reactive cation wins. Compare the metal against hydrogen: if the metal is above H in the reactivity series (K, Na, Ca, Mg, Al, Zn, Fe…), hydrogen is discharged; if below (Cu, Ag, Au), the metal is discharged. At the anode, halides beat hydroxide if concentrated enough — concentrated halide solutions release the halogen, dilute halide solutions release oxygen from OH⁻. Sulfates and nitrates never discharge, so oxygen is always the anode product with them.

ExtendedCopper(II) sulfate — two setups, two outcomes

With inert graphite electrodes, Cu²⁺ is discharged at the cathode (Cu is below H) and coats it pink-brown; OH⁻ is discharged at the anode and oxygen bubbles off; the blue colour fades as Cu²⁺ is used up. With active copper electrodes, Cu²⁺ still deposits at the cathode — but at the anode the copper electrode itself dissolves instead of oxygen forming. Copper atoms lose electrons to become Cu²⁺, so the anode loses mass while the cathode gains it, and the blue colour stays constant. This is the basis of copper refining.

Electroplating

Electroplating is the same idea as active-electrode CuSO₄ electrolysis, put to work: a thin coat of one metal deposited on the surface of another. The object gets the coating; the coating metal supplies it.

The setup

The object to be plated is wired as the cathode. A bar of the plating metal is wired as the anode. The electrolyte is an aqueous solution containing ions of the plating metal — typically a soluble salt such as a sulfate or nitrate. When the current flows, ions of the plating metal migrate to the cathode and deposit as a thin metal layer, while the anode dissolves — atoms of the plating metal give up electrons and enter the solution as ions. The concentration of the electrolyte stays constant, so plating can carry on until the anode is used up.

Worked example: tin-plating an iron strip

A strip of iron is electroplated with tin using tin(II) chloride solution. Wire the iron strip as the cathode and a bar of pure tin as the anode; the electrolyte is aqueous SnCl₂. At the anode, tin atoms lose 2 electrons each, entering the solution as Sn²⁺, so the anode loses mass. At the cathode, Sn²⁺ ions gain 2 electrons and deposit as tin atoms on the iron surface, so the iron gains mass and appears silvery. The Sn²⁺ concentration is preserved — every ion deposited at the cathode is replaced by one dissolving at the anode.

Charge transfer & half-equations

Everything so far has been described in words. Here the same events are written as equations — one for the anode, one for the cathode, both showing the electrons explicitly.

ExtendedCharge carriers in the circuit

Two different particles carry the charge, in two different parts of the circuit. In the external circuit (the wires), electrons flow from the negative terminal of the supply to the cathode, then back from the anode to the positive terminal. In the electrolyte, ions do the carrying — cations drift toward the cathode, anions toward the anode. Electrons do not pass through the solution.

ExtendedWriting half-equations

At the cathode (reduction), positive ions gain electrons, and the number of electrons matches the charge on the ion: Mⁿ⁺ + ne⁻ → M. So Li⁺ + e⁻ → Li, Cu²⁺ + 2e⁻ → Cu, Al³⁺ + 3e⁻ → Al. For hydrogen, two ions pair up to make one molecule: 2H⁺ + 2e⁻ → H₂.

At the anode (oxidation), negative ions lose electrons. Halides pair up into diatomic halogens: 2X⁻ → X₂ + 2e⁻ (X = Cl, Br, I). Hydroxide is the tricky one — it produces oxygen and water: 4OH⁻ → O₂ + 2H₂O + 4e⁻.

Hydrogen–oxygen fuel cells

A fuel cell is electrolysis running in reverse. Instead of pushing current through a compound to break it apart, you let two elements combine — and harvest the electron flow as electricity.

The reactions

Hydrogen is fed to one electrode, oxygen (from the air) to the other. At the hydrogen electrode, H₂ is oxidised, losing electrons to form H⁺: H₂ → 2H⁺ + 2e⁻ (anode). Those electrons flow around the external circuit, doing work. At the oxygen electrode, O₂ is reduced, gaining electrons and combining with H⁺ to make water: O₂ + 4H⁺ + 4e⁻ → 2H₂O (cathode). Overall: 2H₂ + O₂ → 2H₂O.

ExtendedFuel cells vs. petrol engines