Thermodynamics and Electrochemistry

This is Unit Eleven of Chemistry, Part Four: batteries, corrosion and electrolysis. Part Three built the galvanic cell and everything needed to describe one — the anode where oxidation happens and the cathode where reduction happens, the standard cell potential computed from tabulated values, and Faraday's constant, the charge carried by one mole of electrons. Two of those come back constantly here. This part takes that cell out of the laboratory and puts it in three places: inside a battery, where it has been packaged for portable use; on the hull of a ship, where the same mechanism runs uninvited and is called corrosion; and in reverse, where an external power source drives the reaction against its own natural direction. The point running through all three is that none of them is new chemistry. Each is the galvanic cell of Part Three, arranged toward a different end. One thing to carry in from Part Three specifically: a galvanic cell's anode is negative and its cathode positive, and in an electrolytic cell those signs reverse, while oxidation at the anode holds throughout. By the end you will be able to say what a battery actually stores; say what makes one rechargeable; say why wet iron rusts where dry iron does not; say how a sacrificial anode protects the metal it is attached to; and say why electrolysis does not make a reaction spontaneous.

Three entries packaging the galvanic cell. A battery is one or more of those cells sealed into a device, and the first question worth settling is what is actually inside it — because the intuitive answer, a store of electrons, is wrong in a way that makes everything after it confusing. Then the distinction between the batteries you discard and the ones you recharge, which turns out to rest on a property of the reaction rather than of the product. Then a device that removes the fixed supply altogether.

Cells in a box.

Chem 2210 Battery as a Multi-Cell Galvanic Device

A battery is a self-contained device built from one or more galvanic cells, delivering a stated voltage and current. A car battery links several lead-acid cells in series, each contributing about two volts, to reach its rating. What it holds is not a store of electrons waiting to be released — it holds chemical reactants, and the current appears when those reactants are allowed to react. That distinction is why a battery has a shelf life and why its voltage is set by its chemistry rather than by its size.

Can it run backwards?

Chem 2213 Primary vs. Secondary Battery

A primary battery's redox reaction is not practically reversible, so once its reactants are spent the device is discarded. A secondary battery's reaction is reversible, so driving it backward restores the reactants and it can be recharged. A single-use alkaline cell is primary; a lithium-ion cell is secondary. Rechargeable describes the reversibility of a reaction — it is a fact about the chemistry inside, not about the brand on the label or the price paid for it.

Keep feeding it.

Chem 2216 Fuel Cell Mechanism

A fuel cell generates current from a redox reaction between reactants supplied continuously from outside, so it runs as long as fuel and oxidant keep arriving — unlike a battery, which carries a fixed supply sealed in. A hydrogen fuel cell consumes hydrogen and oxygen fed in steadily, producing water and current. It is not a battery with unlimited charge; remove the supply and it stops, because it was never storing anything.

Two entries on the galvanic cell nobody built. Corrosion is the same mechanism from Part Three occurring spontaneously on a metal surface, with the environment supplying what a designed cell supplies deliberately — which is why the defence against it is not a coating but a second, deliberately worse metal.

Rust is electrochemistry.

Chem 2219 Corrosion as Unwanted Spontaneous Oxidation

Corrosion is the unwanted spontaneous oxidation of a metal by its environment — the same galvanic mechanism already built, running without anyone designing a circuit for it. Iron in moist air oxidises to rust through an electrochemical process rather than by reacting with oxygen directly. Water is doing real work there, carrying ions between the anodic and cathodic regions on the metal's own surface, which is why dry iron does not rust nearly as readily as wet iron does.

Give it something worse.

Chem 2222 Cathodic Protection: Sacrificial Anode Principle

Cathodic protection connects the metal you want to keep to a metal that oxidises more easily, so the second corrodes in its place as a sacrificial anode while the protected metal is held as the cathode. Zinc blocks bolted to a steel hull corrode instead of the steel, because zinc's reduction potential is the more negative of the two. The sacrificial metal has to oxidise more readily than what it protects — attach one that does not and it is consumed for nothing.

Three entries on forcing the issue. Everything so far ran because it wanted to; here an external power supply drives a reaction against its own direction. The registered twin in this part names what that device is — not a different sort of chemistry, but the galvanic cell reversed — and this is where the charge signs from Part Three swap over. Then the calculation that turns a current and a clock into a quantity of product.

Push it uphill.

Chem 2225 Electrolysis: Nonspontaneous Redox Driven by Current

Electrolysis forces a nonspontaneous redox reaction to proceed by supplying electrical energy from outside, running the reaction against its natural direction. Passing current through molten sodium chloride drives its decomposition into sodium metal and chlorine gas. It does not make the reaction spontaneous — the cell potential is still negative and the reaction still has no tendency of its own to run. External energy overpowers that, which is a different thing from removing it.

The same cell, reversed.

Chem 2228 Twin: Electrolytic Cell as the Forced-Reverse Galvanic Cell

An electrolytic cell runs the same redox chemistry as a galvanic cell, forced nonspontaneous by external power rather than generated from an already favourable reaction. Recharging a secondary battery is exactly this: the cell that discharged spontaneously as a galvanic cell is driven backward by the charger. Galvanic and electrolytic are not two different chemistries — one rechargeable cell is literally both, in turn, and its electrode charge signs swap as it changes role.

Current, time, product.

Chem 2231 Calculating Product Quantity from Current and Time

Multiply current by time to get total charge in coulombs, divide by Faraday's constant for moles of electrons, use the half-reaction's electron count to convert to moles of product, and finish with molar mass if a mass is wanted. Five amperes for one hour delivers eighteen thousand coulombs, which is about zero point one eight seven moles of electrons; at two electrons per copper atom, that deposits zero point zero nine three moles of copper, close to five point nine grams. The slip to avoid is dividing the current by Faraday's constant directly, which treats amperes as though they were already coulombs and skips the clock entirely.

The unit closes on a capstone that makes its argument in one move. A single zinc-copper pair is run three ways — as a generator, as a packaged battery, and as its own reversal under an external supply — and nothing about the chemistry changes between them. Only the arrangement and the direction do.

One pair, three settings.

Chem 2235 Capstone: One Redox Couple, Three Cell Types

Zinc with copper two plus, built from the two tabulated half-reactions, gives a standard cell potential of one point one zero volts as a galvanic cell — spontaneous, with zinc oxidising at the anode and copper depositing at the cathode. Wire several such cells together and seal them, and you have a battery delivering a multiple of that voltage; the mechanism inside is unchanged, only the packaging and the combined voltage differ. Now take the same cell fully discharged and push three amperes through it for two hours: twenty-one thousand six hundred coulombs, about zero point two two four moles of electrons, which strips zero point one one two moles of copper back into solution and redeposits the same quantity of zinc. That is electrolysis, working against the minus one point one zero volts the reversed direction carries, with the electrode charge signs swapped exactly as Part Three described. One redox pair, three settings. A dead battery has not undergone some different irreversible chemistry — its reaction has run forward to near completion, and recharging is ordinary electrolysis.

Five things to carry out of this part, and out of the unit. A battery stores chemical reactants, not electrons — the current appears when those reactants are permitted to react, which is why a battery has a shelf life and why its voltage is set by its chemistry rather than its size. What makes one rechargeable is the reversibility of that reaction: a secondary cell's chemistry can be driven backward to restore the reactants, a primary cell's cannot practically be, and the word describes the reaction rather than the product's brand or price. Wet iron rusts where dry iron does not because corrosion is electrochemical rather than a direct reaction with oxygen — water carries ions between anodic and cathodic regions on the metal's own surface, completing a circuit nobody designed. A sacrificial anode protects a metal by being easier to oxidise than it is: bolted to the protected metal, the sacrifice corrodes preferentially while the protected metal is held as cathode, and a sacrificial metal that oxidises less readily than what it guards is consumed for nothing. And electrolysis does not make a reaction spontaneous. The cell potential stays negative and the reaction retains no tendency of its own to run; an external supply overpowers that, which is a different claim entirely — and it is why a galvanic cell and an electrolytic cell are one chemistry in two directions, with a rechargeable battery being literally both in turn. That closes Unit Eleven. Unit Twelve takes up descriptive, organic and nuclear chemistry.