Electrochemistry

Every electrochemical change is really a transfer of electrons from one species to another. The atom that hands electrons over is oxidised; the atom that takes them in is reduced. The two always happen together — one cannot give unless another takes — so the combined change is called a redox reaction.

An oxidation number is the charge an atom would carry if every bond in the molecule were fully ionic. It is a piece of electron bookkeeping that lets us see, at a glance, which atom was oxidised (its number goes up) and which was reduced (its number goes down).

• A free element is 0 (e.g. Zn, O₂, Cu).
• A simple ion equals its charge (Na⁺ = +1, Cl⁻ = −1).
• O is usually −2 (peroxides −1); H is usually +1 (metal hydrides −1).
• The sum of all oxidation numbers equals the overall charge.

On the right the contributions are tallied like a shopkeeper's receipt — the only "unknown" entry is forced to the value that makes the total balance.

• Oxidising agent — accepts electrons (and is itself reduced). e.g. O₂, KMnO₄, Cl₂.
• Reducing agent — donates electrons (and is itself oxidised). e.g. metals, H₂, C.

The clearest everyday example is a rusting iron gate. The iron is the reducing agent, slowly handing its electrons to the oxidising agent — the oxygen in damp air — and turning into reddish-brown iron oxide. It is the same electron transfer as a battery, only spread out over months instead of seconds.

• Galvanic cell — converts chemical energy → electrical energy through a spontaneous redox reaction. Every battery is one.

The classic Daniell cell is a zinc rod in zinc sulphate and a copper rod in copper sulphate, joined by a wire. Zinc is oxidised at the anode (−), copper ions are reduced at the cathode (+), and the electrons released by the zinc travel through the wire to the copper — and on the way they can light a bulb or run your torch.

• Salt bridge — a tube of inert electrolyte (e.g. KNO₃) that completes the circuit and keeps both half-cells electrically neutral.

Think of the two beakers as two riverbanks. As zinc dissolves, its bank fills with positive charge; as copper deposits, its bank turns negative. The salt bridge is a pedestrian bridge that lets ions stroll across — negative ions toward the zinc, positive ions toward the copper — keeping both banks balanced so the current can keep flowing.

• Standard electrode potential (E°) — the voltage of a half-cell measured against the standard hydrogen electrode (E° = 0.00 V) at 298 K, 1 M, 1 atm.

Picture the metals as ledges at different heights on a waterfall. A reactive metal like zinc sits high up — its electrons have lots of potential energy and tumble away easily. An unreactive metal like copper sits low down. Electrons flow like water from the high ledge to the low one, and the height of the drop is the cell voltage.

You can build a cell at home. Push a galvanised (zinc) nail and a copper coin into a lemon — the citric acid is the electrolyte — and wire them to an LED. The zinc is the anode, the copper the cathode, and the lemon's acid carries the ions. The voltage you get is just the difference of the two electrode potentials.

A positive E°cell means the reaction is spontaneous — the cell really does deliver a voltage and light the LED.

• Electrolytic cell — uses electrical energy → chemical change: an external supply drives a non-spontaneous reaction. Here the anode is positive.

To electroplate a spoon with silver, the steel spoon is the cathode and a bar of pure silver is the anode, both dipped in a silver-salt solution. The battery pulls electrons off the silver bar (oxidation) and pushes silver ions onto the spoon (reduction), laying down a bright, even coat.

• First law: the mass deposited ∝ the charge passed (Q = I × t).
• Second law: for the same charge, mass deposited ∝ the equivalent mass.