Introduction

Objectives

After studying this chapter, you should be able to:

• identify the dynamic nature of equilibrium in physical processes (solid–liquid, liquid–vapour, solid–vapour, and dissolution of solids/gases in liquids);
• explain the law of chemical equilibrium and write the equilibrium constant expression ($K_c$, $K_p$) for a reaction, and relate the two through $K_p = K_c(RT)^{\Delta n}$;
• use the reaction quotient $Q$ to predict the direction in which a reaction not yet at equilibrium will proceed;
• apply Le Chatelier's principle to predict how concentration, pressure, volume, and temperature changes shift an equilibrium;
• classify electrolytes as strong or weak, and describe ionic equilibrium in solution;
• state and compare the Arrhenius, Brønsted–Lowry, and Lewis concepts of acids and bases;
• calculate pH, pOH, ionization constants ($K_a$, $K_b$), and degree of ionization for weak acids and bases;
• explain how a buffer solution resists changes in pH, and calculate buffer pH using the Henderson–Hasselbalch relation; and
• apply the solubility product constant ($K_{sp}$) to calculate the solubility of a sparingly soluble salt and to predict whether a precipitate will form.

Equilibrium isn't a phenomenon confined to a chemistry lab — it runs quietly through biology

and the environment too. The oxygen you breathe in is carried around your body because of an

equilibrium between O2 molecules and the haemoglobin in your blood, which picks up oxygen in

the lungs and releases it where tissues need it. Carbon monoxide is dangerous for exactly the

same reason: it competes for the same haemoglobin equilibrium, binding more strongly than

oxygen does and blocking it out.

You can see the same basic idea — two opposing processes running at equal rates — in

something far simpler: a liquid sealed inside a closed container. Molecules with enough

kinetic energy keep escaping the liquid surface into the space above it, while vapour

molecules keep striking the surface and rejoining the liquid. Once the rate of escape

(evaporation) exactly balances the rate of return (condensation), the vapour pressure inside

the container stops changing, e.g. $H_2O(l) \rightleftharpoons H_2O(g)$. Nothing has actually

stopped moving — the system has simply reached a dynamic equilibrium, and the resulting

mix of liquid and vapour is called an equilibrium mixture.

The same reasoning extends to genuine chemical reactions run in a closed vessel. As a

reaction proceeds, reactant concentrations fall and product concentrations rise, but only for

a while — eventually the forward and reverse rates become equal and every concentration in

the vessel stops changing, even though both reactions keep running underneath. How far a

reaction gets before settling into this balance varies a great deal: some reactions run

almost to completion, leaving only a trace of unreacted starting material; some barely

proceed at all, leaving mostly reactants; and others settle somewhere in between, with

comparable amounts of reactant and product both present at equilibrium. Because that extent

depends on conditions such as concentration and temperature, learning to predict and control

it is exactly what makes equilibrium such a practically important idea — in industrial

reactors just as much as in a lab.

The chapter builds this picture in three stages. It starts with equilibrium in physical processes — a liquid boiling in a sealed container, a solid dissolving in water, a gas

dissolving under pressure — since these are the simplest cases of a system settling into a

constant, measurable state with no chemical change involved at all. It then moves to

equilibrium in chemical reactions, where the law of chemical equilibrium gives a precise

relationship (the equilibrium constant, $K_c$ or $K_p$) between reactant and product

concentrations once equilibrium is reached, and Le Chatelier's principle predicts how that

balance shifts when you add a reactant, remove a product, or change the pressure or

temperature. Finally, the chapter turns to ionic equilibrium in solution — the equilibria

behind acids, bases, salts, pH, buffers, and the solubility of sparingly soluble compounds —

which underlies everything from antacid tablets to blood pH regulation to water purification.

A theme worth keeping in mind throughout: equilibrium is never about a reaction "stopping." It

is about two opposing processes racing at exactly the same speed, so nothing appears to

change even though, at the molecular level, both directions stay fully active. Every

experiment and demonstration in this chapter is really just a different way of making that

dynamic balance visible.