Reaction rates — OCR A-Level Chemistry

Simple collision theory explains why reaction rates vary. Particles react only when they collide with at least the activation energy and with correct orientation. Most collisions are unsuccessful because the particles lack sufficient energy. Raising concentration or pressure increases the number of particles per unit volume, so collisions become more frequent and the rate rises. Raising temperature increases both collision frequency and the proportion of particles with energy at or above the activation energy, so the effect is usually larger. A catalyst provides an alternative route with lower activation energy. The theory links the observable rate to the frequency of successful collisions.

Increasing the concentration of a reactant in solution, or the pressure of a gaseous reactant, packs more particles into a given volume. With more particles per unit volume, collisions between reactant particles happen more often, so the frequency of successful collisions rises and the rate of reaction increases. For gases, raising pressure compresses the same amount of gas into a smaller volume, giving an equivalent increase in particle concentration. The key link is always between concentration or pressure and collision frequency, not activation energy, which is unchanged. Rate therefore increases with concentration or pressure provided other factors stay constant.

Reaction rate is found from the gradient of a graph of a physical quantity against time. The quantity may be concentration, mass, gas volume or absorbance; the method is the same. Draw a tangent to the curve at the chosen time, then divide the change in the quantity by the change in time: rate = Δy/Δt. The gradient's sign shows whether the quantity rises or falls, so a decreasing concentration gives a negative gradient; the rate of reaction is the magnitude, so report the positive value. Units come from the axes, for example mol dm⁻³ s⁻¹ for concentration against time, or cm³ s⁻¹ for gas volume against time. At the start the tangent is steepest, so the initial rate is greatest; as reactants are used up the curve flattens and the rate falls. Choose two widely separated points on the tangent to reduce reading error.

A catalyst increases the rate of a reaction by providing an alternative reaction pathway of lower activation energy. It is not used up overall and is regenerated, so a small amount can catalyse a large amount of reaction. Because the activation energy is lower, a greater proportion of collisions has energy at or above the activation energy, so more frequent successful collisions occur and the rate rises. A catalyst does not change the position of equilibrium or the enthalpy change of the reaction; it only speeds the approach to equilibrium, helping both forward and reverse reactions equally. A homogeneous catalyst is in the same phase as the reactants, for example aqueous acid; a heterogeneous catalyst is in a different phase, usually a solid with gaseous or liquid reactants, and works by adsorption onto its surface.

A catalyst provides an alternative reaction pathway of lower activation energy, so a greater proportion of collisions between particles have at least the activation energy and the rate of reaction increases. The catalyst is not consumed by the overall reaction: it takes part in one step and is regenerated in a later step, so its mass and chemical identity are restored at the end. For example, a solid catalyst adsorbs reactant molecules, weakens their bonds and holds them in a favourable orientation, then releases the products and is left unchanged. The catalyst does not change the position of equilibrium or the overall enthalpy change, because it speeds the forward and reverse reactions equally. Explain the role by linking lower activation energy to more successful collisions per unit time.

A catalyst increases the rate of a reaction while remaining chemically unchanged at the end of the overall reaction. It does this by taking part in one step of the mechanism and being regenerated in a later step, so its mass and identity are restored. Because it is not consumed, a small amount of catalyst can speed up a large amount of reaction. For example, manganese(IV) oxide catalyses the decomposition of hydrogen peroxide and is recovered unchanged afterwards. The catalyst does not alter the equilibrium position or the overall enthalpy change, because it accelerates the forward and reverse reactions equally. Explain the increase in rate through an alternative pathway of lower activation energy, and explain the 'not used up' point through regeneration in the mechanism.

A catalyst speeds a reaction by providing an alternative pathway whose activation energy is lower than that of the uncatalysed route. On an enthalpy profile diagram the reactants and products sit at the same enthalpy levels whether or not a catalyst is present, because a catalyst does not change the overall enthalpy change of the reaction. What changes is the height of the energy barrier between reactants and products: the catalysed curve peaks lower, so a greater fraction of collisions at a given temperature has enough energy to succeed. The catalysed route may pass through one or more intermediate stages, shown as extra humps, but each step has a lower barrier than the single uncatalysed peak. A catalyst is not used up overall and does not shift the position of equilibrium; it simply lets the system reach equilibrium faster.

Catalysts are classified by the phase relationship between the catalyst and the reactants. A homogeneous catalyst is in the same physical state as the reactants, so it is usually dissolved in the same solution or present in the same gas mixture; an example is acid catalysis in aqueous solution, where hydrogen ions in solution catalyse a reaction between species also in solution. A heterogeneous catalyst is in a different physical state from the reactants, typically a solid catalyst with gaseous or liquid reactants; the reaction occurs at the catalyst surface, where reactant molecules are adsorbed, react and then desorb. Recognising the phase relationship lets you explain why heterogeneous catalysis depends on surface area while homogeneous catalysis depends on the catalyst being intimately mixed with the reactants.

A catalyst provides an alternative reaction pathway of lower activation energy, so a given rate is achieved at a lower temperature and with less heating. This matters economically because energy is a major cost in industrial processes, and a lower operating temperature can reduce fuel consumption and equipment demands. It matters for sustainability because much industrial heating relies on burning fossil fuels; using less fuel lowers energy demand and therefore reduces the CO₂ emissions associated with that combustion. A strong answer links the chain explicitly: catalyst lowers activation energy, so a lower temperature gives an acceptable rate, so less fossil fuel is burned, so less CO₂ is released. Keep the link causal rather than listing benefits separately.

Reaction rates can be followed by measuring a quantity that changes as the reaction proceeds and recording it against time. When a gas is produced, the mass of the reaction vessel can be measured on a balance at intervals; the loss in mass corresponds to gas escaping, so a steeper initial fall means a faster rate. Alternatively, the volume of gas collected can be measured with a gas syringe at intervals; a steeper initial rise in volume means a faster rate. In both methods time is recorded with a stopwatch, readings are taken at regular intervals, and the rate is found from the gradient or from the change in quantity per unit time. The initial rate is usually taken from the steepest part of the graph near the start.

The Boltzmann distribution shows how molecular energies are spread at a given temperature. The curve starts at the origin (0,0), rises to a peak at the most probable energy, and tails off at higher energies without touching the energy axis. The area under the curve represents the total number of molecules, so it stays constant when temperature changes. Raising the temperature flattens and broadens the curve: the peak moves to higher energy and the tail extends further right, so a greater proportion of molecules have energy at or above the activation energy, Ea. This explains why reaction rate increases with temperature. The distribution applies to gases and to the energies of molecules in a sample; it does not show individual molecules changing energy over time.

A qualitative explanation of the Boltzmann distribution describes, without calculation, how molecular kinetic energies are spread at a given temperature and why only some collisions lead to reaction. Molecules in a gas or liquid have a range of energies; the distribution peaks at a moderate energy and tails off at high energy. The activation energy, Ea, is the minimum energy colliding particles need for a successful reaction, shown as a vertical line on the energy axis. Only the fraction of molecules with energy at or above Ea can react, and this fraction is small at low temperature. Raising the temperature broadens the distribution and increases the number of molecules beyond Ea, so more successful collisions occur per second and the rate rises. A catalyst lowers Ea, moving the line left so a larger fraction exceeds it.

A Boltzmann distribution plots the number of molecules against kinetic energy at a fixed temperature. The curve starts near the origin, peaks at moderate energy and tails off at high energy; the area under it represents all the molecules. Only molecules with energy at or above the activation energy, Ea, can react, so the proportion able to react is the area under the curve to the right of Ea. Raising the temperature shifts the peak to higher energy and broadens the curve, so a much larger fraction of molecules lies beyond Ea; the rate therefore rises. Lowering the temperature moves the peak to lower energy and narrows the curve, so fewer molecules exceed Ea and the rate falls. The explanation is qualitative: describe the direction of the shift and its effect on the proportion exceeding Ea, not exact numbers.

Temperature is one of the factors affecting reaction rate. Raising the temperature increases the average kinetic energy of the particles, so a larger proportion of molecules has energy at or above the activation energy. More frequent successful collisions result, so the rate increases. Lowering the temperature reduces the proportion exceeding the activation energy and slows the reaction. In a multiple-choice question, select the option that correctly links the temperature change to the proportion of molecules exceeding the activation energy and to the rate, and reject options that wrongly say the activation energy itself changes or that all molecules gain equal energy.

A catalyst raises the rate of a reaction by providing an alternative route of lower activation energy, and it is not consumed overall. In homogeneous catalysis the catalyst is in the same phase as the reactants, for example Fe²⁺/Fe³⁺ ions in the peroxodisulfate–iodide reaction; in heterogeneous catalysis the catalyst is in a different phase, typically a solid surface such as iron in the Haber process, where reactant molecules adsorb, bonds weaken, reaction occurs and products desorb. The catalyst lowers the activation energy for both the forward and reverse reactions equally, so it speeds the approach to equilibrium but does not change the position of equilibrium or the yield at equilibrium. Enzymes act as biological catalysts with high specificity. A catalyst may be poisoned by impurities that block active sites, which reduces its effectiveness.