Chapter 6 — Equilibrium
Many reactions never run to completion: they settle into a balance in which the forward and backward reactions continue at equal rates. Equilibrium constants, the reaction quotient, Le Chatelier's principle, acids, bases and pH, buffers, salt hydrolysis and solubility products.
In 1909 Fritz Haber passed nitrogen and hydrogen over hot osmium at about 200 atmospheres and collected a few drops of liquid ammonia an hour. The reaction he was fighting, $\mathrm{N_{2}(g)+3H_{2}(g)\rightleftharpoons2NH_{3}(g)}$, never goes to completion: however long the gases stay in the reactor, a mixture remains in which ammonia, nitrogen and hydrogen are all present in fixed proportions. Haber's achievement — and Carl Bosch's, who scaled it up into a factory — was to understand how temperature and pressure change those proportions, and to choose the conditions that give enough ammonia, fast enough, to be worth making. Today more than 150 million tonnes of ammonia a year are made this way, and the fertilisers made from it feed roughly half the people on Earth.
A mixture that has stopped changing is in chemical equilibrium, and this chapter is about the rules that govern it. It begins with equilibria you can see — ice in water, a liquid and its vapour, a gas dissolved in a sealed bottle — and shows that equilibrium is dynamic: the opposing processes never stop, they just balance. The law of chemical equilibrium then gives every reversible reaction a number, the equilibrium constant $K$, that predicts how far the reaction goes and in which direction a mixture will move. Le Chatelier's principle explains how an equilibrium responds when it is disturbed — the key to the Haber process. The second half applies the same ideas to ions in water: acids and bases, the pH scale, weak acids and bases, buffers that keep the pH of your blood close to $7.4$, the pH of salt solutions, and the solubility of sparingly soluble salts. $K_{c}$, $K_{p}$, pH and $K_{sp}$ calculations appear in every Board, JEE and NEET paper.
- describe equilibrium in physical processes (melting, vaporisation, sublimation, dissolution, gases in liquids) and explain why chemical equilibrium is dynamic;
- state the law of chemical equilibrium, write $K_{c}$ and $K_{p}$ for homogeneous and heterogeneous equilibria, give their units and convert between them with $K_{p}=K_{c}(RT)^{\Delta n}$;
- use the size of $K$ and the reaction quotient $Q$ to predict the extent and direction of a reaction, and calculate equilibrium concentrations with an ICE table;
- relate $K$, $Q$ and $\Delta G$, and use Le Chatelier's principle to predict the effect of concentration, pressure, inert gas, temperature and a catalyst;
- explain the Arrhenius, Brønsted–Lowry and Lewis concepts of acids and bases and identify conjugate acid–base pairs;
- use $K_{w}$ and the pH scale, and calculate the pH of strong and weak acids and bases with $K_{a}$, $K_{b}$ and the degree of ionisation;
- explain the common-ion effect, salt hydrolysis and buffer action, and use the Henderson equation;
- relate solubility and $K_{sp}$ and predict whether a precipitate forms.
Sections in this chapter
Chemical Equilibrium and the Equilibrium Constant
Equilibrium in physical processes and Henry's law; the dynamic nature of chemical equilibrium; the law of chemical equilibrium; $K_{c}$ and $K_{p}$ for homogeneous and heterogeneous equilibria.
Using K, Gibbs Energy and Le Chatelier's Principle
The extent of a reaction from the size of $K$; the reaction quotient and the direction of change; ICE-table calculations; $K$, $Q$ and $\Delta G$; Le Chatelier's principle and the Haber process.
Ionic Equilibrium: Acids, Bases and pH
Strong and weak electrolytes; Arrhenius, Brønsted–Lowry and Lewis acids and bases; the ionic product of water and the pH scale; $K_{a}$, $K_{b}$ and the degree of ionisation; polybasic acids; acid strength; the common-ion effect.
Buffers, Salt Hydrolysis and Solubility Product
The pH of salt solutions; buffer solutions and the Henderson equation; solubility equilibria, $K_{sp}$ and molar solubility; the common-ion effect on solubility; predicting precipitation.