Contents: 6 sections
Syllabus points
- Explain reaction rate in terms of collision theory and activation energy.
- Explain the effect of concentration, pressure, surface area and temperature on rate.
- Describe and interpret the Boltzmann distribution of molecular energies.
- Explain how a catalyst increases rate by providing an alternative route of lower activation energy.
- Distinguish between homogeneous and heterogeneous catalysis.
Collision theory
For a reaction to happen, particles must collide, and a collision is only successful if two conditions are met:
- The particles collide with energy equal to or greater than the activation energy.
- They collide in the correct orientation.
Activation energy (Eₐ) is the minimum energy that colliding particles must have for a reaction to occur.
Most collisions achieve nothing. Rate depends on the frequency of successful collisions, and every factor below works by changing either how often particles meet or what fraction of those meetings have enough energy.
The factors
| Factor | Effect on rate | Why |
|---|---|---|
| Higher concentration | Faster | More particles per unit volume, so collisions are more frequent |
| Higher pressure (gases) | Faster | Same reason: the particles are closer together |
| Larger surface area | Faster | More of the solid is exposed, so more collisions per second |
| Higher temperature | Much faster | Particles move faster, and a much larger fraction exceed Eₐ |
| Catalyst | Faster | Alternative route with a lower Eₐ |
Temperature is different in kind from the first three, and the difference is the point of the topic. Concentration, pressure and surface area only change the frequency of collisions. Temperature changes the frequency and the proportion of collisions that are energetic enough, and the second effect is far the larger. A rise of 10 °C often roughly doubles the rate, which a modest increase in collision frequency could never explain.
The Boltzmann distribution
The Boltzmann distribution shows how molecular energies are spread in a sample. Its shape matters as much as its position:
- It starts at the origin, because no molecule has zero energy.
- It rises to a peak at the most probable energy.
- It has a long tail to the right that never touches the axis, because there is no upper limit on energy.
- The total area under the curve is the total number of molecules.
The activation energy is marked as a vertical line well to the right of the peak. Only the molecules in the shaded area beyond it can react.
Raising the temperature
At a higher temperature the curve becomes lower and flatter, and its peak shifts to the right. The area under the curve is unchanged, because the number of molecules has not changed.
The shaded area beyond Eₐ becomes much larger. That is the explanation the mark scheme wants: a small rise in temperature produces a large increase in the proportion of molecules with energy greater than the activation energy.
Two details that are frequently drawn wrongly: the new curve must start at the origin like the old one, and the two curves must cross, since the area is conserved.
Adding a catalyst
A catalyst does not change the distribution at all. The curve stays exactly where it was. What moves is the activation energy line, which shifts to the left, so a greater proportion of the same molecules now have enough energy.
That distinction is worth being precise about: temperature changes the curve, a catalyst changes the line.
Catalysis
A catalyst increases the rate of a reaction by providing an alternative reaction pathway of lower activation energy, and is chemically unchanged at the end.
A catalyst is not consumed, which is why a small quantity is enough for a large amount of reaction. It is also worth remembering from the equilibria topic that a catalyst does not shift the position of equilibrium: it lowers Eₐ for the forward and reverse reactions equally.
Homogeneous catalysis has the catalyst in the same phase as the reactants. The oxidation of sulfur dioxide by NO₂ in the atmosphere is a gas-phase example.
Heterogeneous catalysis has the catalyst in a different phase, usually a solid with gaseous or aqueous reactants. Iron in the Haber process and platinum in a catalytic converter are the standard examples. The reaction happens on the surface, which is why these catalysts are used as fine powders, meshes or thin coatings on a honeycomb: it maximises the surface area available.
On an energy profile, a catalysed reaction has a lower hump, and often two smaller humps if the route goes through an intermediate. The reactant and product levels are unchanged, so ΔH is the same. Drawing a catalyst as lowering the products is a serious error.
Common mistakes
- Saying a catalyst lowers the energy of the reaction. It lowers the activation energy only, and ΔH is unaffected.
- Explaining the temperature effect purely by faster movement. The dominant reason is the larger proportion of molecules exceeding Eₐ.
- Drawing a Boltzmann curve that starts above the origin, or one that does not cross the original curve.
- Saying a catalyst shifts the Boltzmann curve. It moves the activation energy line instead.
- Saying a catalyst increases the yield at equilibrium. It only gets there faster.
- Forgetting orientation. Energy alone does not make a collision successful.