How Chemical Equilibrium Is Established

A reverse reaction becomes possible as products accumulate

Imagine a closed vessel initially containing only reactants A and B. They form products C and D. Once products exist, collisions between them can regenerate the reactants. A reversible reaction is represented by a double arrow because both chemical directions are possible under the conditions being considered. The net change depends on the difference between their rates.

At first, there may be little or no reverse reaction because little product is present. As products accumulate and reactants are consumed, the forward rate commonly falls and the reverse rate rises. Eventually they are equal. The concentrations then remain constant because the amount consumed per unit time is replenished at the same rate. Constancy does not mean that the reactants have all disappeared.

Some equilibria leave mostly reactants; others leave mostly products; still others contain substantial amounts of both. These are different extents of reaction, not different definitions of equilibrium. Industrial conditions are chosen partly to favour the useful product, but a good final yield and a fast reaction are separate requirements.

A water-transfer model makes the balance visible

Use two measuring cylinders and two narrow tubes. Begin with coloured water in the first cylinder and none in the second. Transfer a tube-full in each direction during each cycle. The volume carried by a tube depends on how deeply it is immersed, so the higher water level initially produces a larger transfer. Repeating the cycle gradually establishes levels at which the two opposite transfers are equal.

At that stage, continue transferring water. The levels remain steady even though water is moving both ways. If the tubes have different diameters, the final levels can differ because different column heights can give equal transferred volumes. This is an analogy for equal reaction rates with unequal concentrations. It is a model, so it does not claim that actual molecular reaction rates obey the same tube geometry.

Isotopes reveal reactions that concentration measurements hide

Ammonia synthesis provides direct chemical evidence. One mixture reaches equilibrium with ordinary hydrogen, H2, nitrogen and ammonia. Another reaches equilibrium using deuterium, D2, producing deuterated ammonia, ND3. Deuterium is an isotope of hydrogen: it has the same element identity but a different nuclear mass. The mixtures can be followed by mass spectrometry because molecules containing different numbers of D atoms have different masses.

After the equilibrated mixtures are combined under suitable conditions, mixed species such as NH2D, NHD2 and HD appear. The observed composition can remain approximately steady while atoms exchange among molecules. If every reaction had stopped at equilibrium, these mixed molecules could not have formed. The isotope experiment supports continuing forward and reverse reactions. Isotope substitution can slightly change quantitative equilibrium values; the evidence here is the exchange itself.

Equilibrium can be approached from either direction

For hydrogen reacting with iodine, a mixture initially containing H2 and I2 forms HI until the concentrations level off. Starting instead with HI allows decomposition to form H2 and I2. At the same temperature and volume, with the same total numbers of hydrogen and iodine atoms and the same accessible phases, both preparations approach the same equilibrium composition. Merely saying “same temperature” is insufficient if the atom inventories are different.

On a concentration–time graph, equilibrium is shown by horizontal concentration curves. Their heights may differ. The curve for each species must reflect the balanced equation: forming two moles of HI consumes one mole each of H2 and I2. A concentration graph describes amounts per volume; a rate graph describes how fast they change. Confusing their vertical axes causes many errors.

Reversible binding is also important in living systems. Oxygen binds to haemoglobin in the lungs and can be released where oxygen is needed. Carbon monoxide competes strongly for haemoglobin binding, disrupting oxygen transport. These applications show why changing concentrations can have important consequences in an equilibrium system; they do not mean every biological process is itself at equilibrium.