Everything from the 3 lessons of this topic on a few pages: the key idea of each lesson, the diagram to remember, the points to learn and the traps examiners set. Read it, cover it, test yourself.
The rate of reaction is the change in amount or concentration of a reactant or product per unit time, often in mol dm−3 s−1. It is found from the gradient of a concentration-time (or volume-time) graph; the gradient of the tangent at t = 0 is the initial rate. By collision theory, particles react only if they collide with at least the activation energy, Ea, and in the right orientation. Raising the concentration, the pressure of a gas or the surface area of a solid makes collisions more frequent, so successful collisions happen more often.
Increasing the concentration does not give the particles more energy; it only increases the frequency of collisions, so the number of successful collisions per second rises. Say “more frequent successful collisions”, not just “more collisions”. Temperature is the factor that changes the energy of the particles, covered in the next lesson.
The Boltzmann distribution shows the number of molecules with each energy. It starts at the origin (no molecule has zero energy), rises to a peak (the most probable energy) and then falls, approaching but never touching the energy axis, because there is no maximum energy. The area under the curve equals the total number of molecules. Only molecules to the right of Ea can react. At a higher temperature the curve is lower and broader, with its peak moved to the right and the same area; the area beyond Ea increases greatly, so a much larger proportion of collisions succeed.
On the higher-temperature curve the peak is lower and further right, and the area is the same, because the number of molecules has not changed. And temperature does not change Ea; it changes how many molecules have enough energy to overcome it. The main reason for the faster rate is the larger proportion of molecules with E ≥ Ea, not the increase in collision frequency.
A catalyst increases the rate of a reaction by providing an alternative route with a lower activation energy; it is chemically unchanged at the end. On the Boltzmann distribution, Ea moves left, so a larger proportion of molecules can react. ΔH is unchanged, and because forward and reverse rates rise equally, a catalyst does not move the position of equilibrium. A homogeneous catalyst is in the same phase as the reactants and usually forms an intermediate; a heterogeneous catalyst is in a different phase and works by adsorbing reactants onto its surface.
A catalyst lowers the activation energy of the forward and reverse reactions, so it does not change ΔH, the position of equilibrium or the yield. On a Boltzmann diagram for a catalysed reaction, do not redraw the curve; just move the Ea line to the left.