Effect of a catalyst and the collision theory
A catalyst gives a new path with lower Ea and is not used up. Collision theory says molecules react only in effective collisions: enough energy and the right orientation.
What a catalyst does
Potassium chlorate gives oxygen slowly on heating. A little manganese dioxide makes it much faster: 2KClO₃ → 2KCl + 3O₂ (heat, MnO₂ catalyst). At the end the MnO₂ is unchanged. Such a substance is a catalyst. By the intermediate complex theory, the catalyst forms temporary bonds with the reactants to give an intermediate complex. This breaks down to give the products and gives back the catalyst. The new path has a lower activation energy, so more molecules can react at the same temperature. A substance that lowers the rate is called an inhibitor, not a catalyst.
The catalysed path
Because the catalyst is given back, a small amount of it can change a large amount of reactants.
How to draw the energy profile with and without a catalyst
- Label the y-axis Potential energy and the x-axis Reaction coordinate.
- Draw a reactant level on the left and a lower product level on the right (exothermic).
- Join them with a tall hump labelled "uncatalysed" and a shorter hump labelled "catalysed".
- Mark Ea (uncatalysed) and Ea′ (catalysed) from the reactant level to each peak. Ea′ is smaller.
- Mark ΔH, the gap between reactants and products: it is the same for both paths.
Key points about catalysts
- A catalyst does not change ΔG of a reaction, so it can speed up only reactions that are already spontaneous. It cannot make a non-spontaneous reaction happen.
- It does not change the equilibrium constant. It speeds up the forward and backward reactions equally, so equilibrium is reached sooner.
- A small amount can catalyse a large amount of reactants.
- It is recovered chemically unchanged at the end, though its physical form may change.
Collision theory of chemical reactions
Collision theory was developed by Max Trautz and William Lewis (1916 to 1918). It treats molecules as hard spheres and says a reaction happens only when reactant molecules collide. The number of collisions per second per unit volume is the collision frequency (Z). For A + B → products, rate = Z_AB e^(−Ea/RT). Not every collision works. An effective collision needs two things: (i) energy at least equal to the threshold energy (activation energy + energy the molecules already have), and (ii) proper orientation. So the rate is written rate = P Z_AB e^(−Ea/RT), where P is the probability or steric factor. One weakness: the theory treats molecules as hard spheres and ignores their structure.
Effective vs ineffective collisions
The catalytic converter in petrol cars and two-wheelers in India uses platinum-group metals to change harmful CO and nitrogen oxides into CO₂ and N₂. Enzymes in our saliva and gut are natural catalysts that digest food at body temperature.
2 marks: “How does a catalyst increase the rate of a reaction? Draw the energy diagram to show this.” It provides an alternative path with lower activation energy via an intermediate complex. Diagram: two humps between the same reactant and product levels; catalysed hump lower; same ΔH.
3 marks: “State the two conditions for an effective collision. Write the collision theory rate expression and explain the term P.” (i) Energy at least equal to the threshold energy; (ii) proper orientation. Rate = P Z_AB e^(−Ea/RT). P is the probability or steric factor, which allows for the fact that only properly oriented collisions give products.
2 marks: “Give two things a catalyst does not change in a reversible reaction.” ΔG of the reaction and the equilibrium constant (it only helps equilibrium to be reached sooner).
A catalyst cuts the hill, not the height difference: lower Ea, same ΔH, same ΔG, same K.