Equilibrium in Physical Processes

Equilibrium is a state in which two opposing changes occur at exactly equal rates, so that no net

change in the system is visible even though, at the microscopic level, matter continues to move

between the two sides. Several everyday physical processes reach exactly this kind of balance, and

studying them first builds the intuition needed for the chemical equilibria that follow.

Solid-liquid equilibrium. Consider ice and water in a perfectly insulated vessel at $0\,^\circ \text{C}$ and $1\ \text{atm}$. Molecules at the surface of the ice continually break free and join

the liquid (melting), while liquid molecules continually lose kinetic energy and rejoin the solid

lattice (freezing). At $0\,^\circ\text{C}$ these two rates are equal, so the mass of ice and the mass

of water both stay constant with time — this is why a mixture of ice and water, left alone, holds

steady at $0\,^\circ\text{C}$ for as long as both phases remain present.

Liquid-vapour equilibrium. In a closed vessel containing a liquid and empty space above it,

molecules continually escape from the liquid surface into the vapour (evaporation) while vapour

molecules continually strike the surface and are recaptured (condensation). Initially evaporation

dominates, but as vapour accumulates, the rate of condensation rises until it exactly equals the rate

of evaporation. The pressure exerted by the vapour at this point is called the vapour pressure of

the liquid at that temperature, and it is a genuine equilibrium property — constant, and

reproducible, for a given liquid at a given temperature, regardless of how much liquid or empty space

the vessel contains.

Solid-vapour equilibrium. Certain solids, such as camphor, ammonium chloride and iodine, pass

directly into the vapour phase without an intervening liquid — a process called sublimation. In a

closed vessel, this too settles into a dynamic balance between sublimation and the reverse process,

deposition, characterized by its own equilibrium vapour pressure.

Equilibrium involving dissolution of a solid or a gas in a liquid. When excess solid solute is

stirred into a solvent, dissolution and crystallization occur simultaneously; once the solution is

saturated, the two rates become equal and the concentration of dissolved solute (the solubility)

remains constant, even as individual solute particles continue to dissolve and re-deposit. The same

logic applies to a gas dissolved in a liquid, such as $\text{CO}_2$ dissolved in a sealed bottle of

a carbonated drink: the rate at which gas molecules enter solution equals the rate at which dissolved

molecules escape back into the gas phase, and the amount dissolved at equilibrium (at a fixed

temperature) is proportional to the pressure of the gas above the liquid, a relationship formalized

later as Henry's law.

General characteristics common to all physical equilibria. Four features recur in every case

above: (i) equilibrium can be attained only in a closed system, from which matter is not allowed to

escape or enter freely; (ii) all measurable, macroscopic properties of the system (mass, pressure,

concentration) become constant; (iii) the equilibrium is dynamic — the two opposing processes never

stop, they merely proceed at equal rates; and (iv) equilibrium can be approached from either

direction and, at a given temperature, exhibits a definite, reproducible, characteristic value (a

particular vapour pressure, a particular solubility) irrespective of the exact quantities of the

phases present.