Module 3: Periodic Table and Energy · Year 12
3.2.2 Reaction rates
Use collision theory, gradients and Maxwell–Boltzmann distributions to explain rates, then connect catalysts to activation energy and practical rate measurements.
What you need to know
Open a line for a quick recap. If it feels obvious, move straight to the linked practice.
3.2.2 a Explain how concentration and gas pressure affect rate by changing collision frequency. Quick revision
If you increase the concentration of a solution, there are more reactant particles in a given volume, so collisions happen more frequently. If you compress reacting gases, the particles are closer together and again collide more frequently.
At the same temperature, you do not need to invent an energy change here. The main reason the rate increases is that there are more collisions per unit time, so there are more successful collisions per unit time.
3.2.2 b Calculate reaction rate from gradients of graphs such as volume-time, mass-time or concentration-time. Quick revision
Average rate comes from a change over a time interval. For the rate at one instant, draw a tangent to the curve at that point and calculate its gradient using two well-separated points on the tangent.
Pay attention to what is on the vertical axis. If mass is falling, your raw gradient is negative; a question asking for the rate of reaction will often want the positive magnitude. Carry the graph units through your gradient.
3.2.2 c Explain catalysis as an alternative reaction route with a lower activation energy; the catalyst is regenerated overall. Quick revision
When you add a catalyst, the reaction gets an alternative route with a lower activation energy. At the same temperature, a larger fraction of collisions now have enough energy to react, so the rate increases.
The catalyst can take part in intermediate steps, but it is regenerated overall. On an enthalpy profile, show a lower catalysed peak while keeping the same reactant level, product level and ΔH.
3.2.2 d Explain homogeneous and heterogeneous catalysts and discuss their economic and environmental importance. Quick revision
If the catalyst is in the same phase as the reactants, you call it homogeneous; if it is in a different phase, it is heterogeneous. For a solid heterogeneous catalyst with gaseous reactants, the chemistry happens at the surface, so available surface and active sites matter.
Industrially, you can use a catalyst to get a useful rate at a lower temperature or pressure, cutting energy use and cost. It can also improve process efficiency and reduce unwanted emissions, although catalysts themselves may be expensive or rely on scarce metals.
3.2.2 e Describe practical methods for investigating reaction rates using mass, gas volume and time. Quick revision
Choose a measurement that changes as the reaction proceeds. If a gas escapes, you can follow the loss of mass on a balance; if you want the gas directly, collect it in a gas syringe and record volume against time.
Keep the variable you are testing separate from the control variables. Take enough readings to define the curve, start timing consistently, and choose apparatus with suitable resolution. If the reaction is very fast, the delay while mixing and starting the timer can become a significant source of uncertainty.
3.2.2 f Connect the Maxwell–Boltzmann energy distribution with activation energy and the fraction of molecules able to react. Quick revision
A Maxwell–Boltzmann distribution shows how molecular energies are spread at a particular temperature. The curve starts at the origin, rises to a most-probable energy, then tails off at high energy; the total area under it represents the total number of molecules.
Draw Eₐ as a vertical line. Molecules to the right have at least the activation energy, so that area tells you the fraction capable of reacting when they collide with a suitable orientation.
3.2.2 g Use Boltzmann distributions to explain the effect of temperature and catalysts on rate. Quick revision
When you raise the temperature, the distribution becomes broader and lower, and its peak moves to higher energy. The useful thing to focus on is the area to the right of Eₐ: it becomes much larger, so a greater fraction of molecules can react.
A catalyst is different. You do not change the temperature distribution; you lower Eₐ, so draw the activation-energy line further left. Again, a larger fraction of molecules lies beyond the threshold.