Topic I: Rate of Reaction

Reactions occur at widely different speeds, from explosions to rusting. This part defines reaction rate and shows how it is measured, then explains how concentration, surface area, temperature and catalysts change it. Collision theory ties everything together by explaining rate in terms of the frequency and energy of particle collisions.

Defining and measuring rate

The rate of a reaction is the change in amount of a reactant or product per unit time. It is usually fastest at the start, when reactant concentrations are highest, and slows as reactants are used up, becoming zero when the reaction stops. Rate can be followed by any property that changes measurably: the volume of gas produced collected in a syringe, the loss of mass as gas escapes from a flask on a balance, the time for a cross to disappear behind a precipitate (turbidity), or changes in colour, pH or conductivity. A graph of the measured quantity against time gives a curve whose gradient is the rate; the steeper the curve, the faster the reaction at that moment.

Collision theory

Collision theory explains rate at the particle level. For a reaction to occur, particles must collide, and the collision must satisfy two conditions: the particles must have at least the activation energy, the minimum energy needed to react, and they must collide with the correct orientation. Only a fraction of collisions meet both conditions, and these are called successful or effective collisions. The rate of reaction depends on the frequency of successful collisions per second. Anything that increases either the number of collisions per second or the fraction that have enough energy will speed up the reaction. This single idea underlies the explanation of every factor that affects rate, so it must be applied carefully in answers.

Concentration and pressure

Increasing the concentration of a dissolved reactant, or the pressure of a gas, packs the particles more closely together. This means the particles collide more often, so the frequency of successful collisions per second rises and the reaction speeds up. For gases, raising the pressure has the same effect as raising concentration because it reduces the volume and crowds the particles. As a reaction proceeds and reactants are consumed, their concentration falls, which is why the rate decreases over time. In experiments, doubling the concentration of a reactant often roughly doubles the initial rate, a relationship that can be investigated by measuring the rate at several different starting concentrations.

Surface area

When a solid reacts with a liquid or gas, only the particles on the surface are exposed to collisions. Breaking the solid into smaller pieces, or grinding it into a powder, greatly increases the total surface area available. This exposes more particles to the other reactant, so collisions are more frequent and the rate increases. A lump of marble chips reacts slowly with acid, but the same mass of powdered marble reacts much faster. This is why powdered or finely divided reactants are used in industry to speed up production, and why fine dusts of flammable materials such as flour or coal can be dangerously explosive when suspended in air.

Temperature

Raising the temperature increases reaction rate for two linked reasons. First, the particles gain kinetic energy and move faster, so they collide more frequently. More importantly, a greater fraction of the particles now have energy equal to or above the activation energy, so a much larger proportion of collisions are successful. Because the second effect is the stronger one, even a modest rise in temperature can produce a large increase in rate; as a rough guide, the rate of many reactions roughly doubles for every 10 degrees Celsius rise. The reverse is also true: cooling slows reactions down, which is why food is refrigerated to slow the chemical changes that cause spoiling.

Catalysts

A catalyst is a substance that speeds up a reaction without being used up, so it can be recovered unchanged at the end. It works by providing an alternative reaction pathway with a lower activation energy, so a greater fraction of collisions have enough energy to be successful and the rate rises. A catalyst does not change the products or the overall enthalpy change, only the speed. Small amounts can catalyse large quantities of reactant. Manganese(IV) oxide catalyses the decomposition of hydrogen peroxide, and transition metals and their compounds are common industrial catalysts, such as iron in the Haber process. Enzymes are biological catalysts that work under mild conditions inside living things.