Electrochemistry: Electrolysis, Batteries, Fuel Cells and Corrosion

Dip two copper strips into copper sulphate solution and apply an external DC voltage, and Cu2+ ions discharge at the cathode, depositing as copper metal, while at the anode, copper metal dissolves back into Cu2+ ions, a reaction that would never proceed on its own without that external push. This exact setup is the basis for an entire industrial process: making impure copper the anode, so it dissolves under applied current, while pure copper deposits at the cathode, refining copper to remarkably high purity purely through electrochemistry. The same basic principle scales up to extracting metals that have no available chemical reducing agent strong enough to isolate them any other way; sodium and magnesium are produced by electrolysing their molten (fused) chlorides, and aluminium by electrolysing aluminium oxide dissolved in molten cryolite, all genuinely non-spontaneous reactions made to happen purely by supplying electrical energy from outside.

Michael Faraday worked out the precise, quantitative relationship between electricity and electrolysis in 1833-34, in two laws. The first: the amount of chemical change at an electrode is directly proportional to the quantity of electricity passed through, Q = It (charge equals current times time, in coulombs). The second: for a fixed quantity of electricity, different substances liberated are proportional to their chemical equivalent weight, atomic mass divided by the number of electrons needed to reduce that particular ion. The charge on exactly one mole of electrons is called a Faraday, approximately 96487 C/mol (rounded to 96500 for most calculations), and this lets a chemist calculate exact masses directly: electrolysing CuSO4 solution with 1.5 A for 10 minutes passes 900 coulombs of charge, and since depositing one mole (63 g) of copper needs 2 Faradays (since Cu2+ + 2e- to Cu), the mass deposited works out to (63 x 900)/(2 x 96487), roughly 0.29 grams, calculated purely from current and time.

Electrolysing molten NaCl is simple: only Na+ can be reduced at the cathode and only Cl- can be oxidised at the anode, so sodium metal and chlorine gas are the only possible products. Electrolysing aqueous NaCl is genuinely more complicated, since water itself contributes H+ and OH- ions alongside Na+ and Cl-, creating real competition at each electrode. At the cathode, Na+ to Na (E° = -2.71 V) competes against H+ to 1/2H2 (E° = 0.00 V), and since the reaction with the higher E° is preferred, hydrogen wins, water is reduced instead of sodium. At the anode, water oxidising to oxygen (E° = 1.23 V) is theoretically preferred over chloride oxidising to chlorine (E° = 1.36 V), since lower E° is preferred for oxidation, yet in practice chlorine gas forms instead, because oxygen evolution suffers from a significant overpotential, an extra voltage penalty beyond the simple E° prediction, making chlorine the reaction that actually, practically occurs.

A primary battery undergoes its redox reaction only once, becoming permanently dead once the reactants are consumed, and it cannot be recharged. The familiar dry cell (Leclanche cell) uses a zinc container as its own anode, a central carbon rod cathode surrounded by manganese dioxide and carbon powder, and a moist ammonium chloride/zinc chloride paste as electrolyte, delivering roughly 1.5 V; zinc oxidises while manganese is reduced from +4 to +3 at the cathode. The mercury cell, better suited to low-current devices like hearing aids and watches, uses a zinc-mercury amalgam anode against a mercuric oxide/carbon cathode, giving a notably useful property: since its overall reaction, Zn(Hg) + HgO(s) to ZnO(s) + Hg(l), involves no dissolved ion whose concentration can change during use, its voltage, about 1.35 V, stays essentially constant throughout its entire working life, unlike the dry cell's voltage, which gradually drifts as it discharges.

A secondary battery can be recharged, undergoing its discharge reaction repeatedly over its lifetime, since passing current through it backward reverses the chemistry entirely. The lead storage battery, standard in automobiles, uses a lead anode and a lead dioxide (PbO2) cathode, both immersed in roughly 38% sulphuric acid, with discharge reactions Pb(s) + SO42- to PbSO4(s) + 2e- at the anode and PbO2(s) + SO42- + 4H+ + 2e- to PbSO4(s) + 2H2O at the cathode, both electrodes converting to lead sulphate as the battery discharges. Charging simply drives both of these reactions in reverse, converting the lead sulphate back into lead metal at one electrode and lead dioxide at the other, restoring the battery to its original, ready-to-discharge state. This same reversible-reaction principle underlies other rechargeable batteries too, such as the longer-lived but more expensive nickel-cadmium cell.

A fuel cell is a galvanic cell with a genuinely different operating principle: reactants are fed in continuously and products removed continuously, rather than being sealed inside from the start, so it runs for as long as fuel keeps arriving, never needing to recharge at all. The most successful design reacts hydrogen with oxygen, bubbled through porous carbon electrodes into concentrated aqueous NaOH, with platinum or palladium catalysts speeding the electrode reactions along; famously, this exact cell powered the Apollo space missions, with its only byproduct, water vapour, condensed and added directly to the astronauts' own drinking water supply. Fuel cells convert chemical energy to electricity at roughly 70% efficiency, compared to a conventional thermal power plant's roughly 40%, since thermal plants waste energy converting heat to steam to mechanical turbine motion to electricity, a longer, lossier chain than a fuel cell's single, direct chemical-to-electrical conversion step.

Corrosion, most familiar as iron rusting, is fundamentally an electrochemical process, a tiny galvanic cell forming spontaneously right on a metal's own surface. At one spot on an iron object, oxidation occurs, Fe(s) to Fe2+(aq) + 2e-, and that spot behaves exactly like an anode. The electrons released travel through the metal itself to a different spot, where they reduce oxygen in the presence of H+ (supplied by dissolved atmospheric CO2 forming carbonic acid), O2(g) + 4H+(aq) + 4e- to 2H2O(l), and that spot behaves exactly like a cathode. The overall reaction, 2Fe(s) + O2(g) + 4H+(aq) to 2Fe2+(aq) + 2H2O(l), has a standard cell potential of +1.67 V, genuinely spontaneous, and the resulting Fe2+ ions are further oxidised by atmospheric oxygen into rust, hydrated ferric oxide, releasing yet more H+ to keep the whole cycle going.

Since rusting requires a metal surface to contact both water and oxygen simultaneously, the most direct prevention strategy is simply keeping that contact from happening at all: painting a surface, or coating it with a protective chemical layer, physically blocks the metal from air and moisture. A second approach coats the metal with a different, less reactive metal instead, tin-plating or zinc-plating being common examples, again forming a physical barrier. A genuinely elegant third approach, purely electrochemical, uses a sacrificial electrode: attaching a more reactive metal, like magnesium or zinc, directly to the object needing protection. Because that attached metal sits lower on the activity series, it becomes the anode instead of the iron, corroding away itself while the iron, now effectively forced into the cathode role, stays protected, a strategy genuinely used to protect buried pipelines and ship hulls, sacrificing a replaceable block of metal to save the much larger structure.