Catalysts speed up reactions by providing an alternative pathway for the reaction with a lower activation energy, Ea, and the catalyst is either unchanged during the reactions that occur (as we saw with heterogeneous catalysts in a catalytic converter for example) or the catalyst is regenerated through a series of reactions.
We can visualise the effect of a catalyst on a reaction with a Maxwell-Boltzmann distribution curve …
A homogeneous catalyst is in the same physical state as the reactants e.g. both are gases or both are aqueous ions
The alternative pathway a homogeneous catalyst provides for a reaction involves forming some kind of intermediate.
In many reactions the intermediate is a molecule formed from the reaction of one of the reactants and the catalyst. The intermediate then breaks down to form the products, regenerating the catalyst.
E.g. Cl radicals catalyse the reaction of ozone with oxygen radicals in the stratosphere, resulting in a reaction that is now 1500 times faster. In addition, the Cl radical is regenerated and can go on to catalyse up to a million more ozone depleting reactions.
The enthalpy profile for this reaction is shown below:
Key points:
- Each step in the catalysed reaction has its own activation energy barrier to overcome, but importantly the combined activation energy for each step of the catalysed reaction (reactants ⇾ intermediate and intermediate ⇾ products) is lower than the activation energy for the uncatalysed reaction.
- We need to be clear on the difference between a transition state and an intermediate. Intermediates are reactive and short-lived molecules although they are neither reactants nor products. A transition state is best thought of as an arrangement of atoms the reaction passes through on route from reactant to product (such as a change in geometry).
Transition metal ions are often employed as homogeneous catalysts – they enable the transfer of electrons between reactants with the metal ions being oxidised in the first step of the alternative pathway and reduced in the second (or vice versa).
E.g. the reaction between iodide ions, I–(aq), and peroxodisulfate ions, S2O82-(aq), is very slow because it involves a collision between two negatively charges ions.
S2O82-(aq) + 2I–(aq) ⇾ I2(aq) + 2SO42-(aq)
We can add iron(II) ions, Fe2+(aq) to catalyse the reaction.
Step 1: 2Fe2+(aq) + S2O82-(aq) ⇾ 2Fe3+(aq) + 2SO42-(aq)
Step 2: 2Fe3+(aq) + 2I–(aq) ⇾ 2Fe2+(aq) + I2(aq)
Fe2+ is oxidised to Fe3+ in step 1 and Fe3+ is reduced back to Fe2+ in step 2. The Fe2+ catalyst has been regenerated. If we combine both steps, the catalyst and the intermediate Fe3+ cancel out to give us the overall equation for the reaction between the iodide ions and the peroxodisulfate ions.
We can sketch an energy profile for this reaction just as before:
Practice questions
Bromine radicals, Br, act as catalysts to break down ozone in the stratosphere via a two-step mechanism:
Step 1: Br + O3 ⇾ BrO + O2
Step 2: BrO + O ⇾
(a) Complete the second step and then combine both steps to write an overall equation for the reaction.
(b) Explain why the bromine radical is described as a homogeneous catalyst.
(c) The activation energy for the catalysed reaction must be lower than that of the uncatalysed reaction. Complete the diagrams below to show this, labelling the activation energy in each.
(d) Give two other factors that will speed up the rate of the overall reaction from your answer to part (a), explaining their effect in terms of collision theory.
Answers
(a) Step 2: BrO + O ⇾ Br + O2; overall reaction: O3 + O ⇾ 2O2
(b) Br is a catalyst because it increases the rate of the reaction and is regenerated / not used up in the overall reaction. It is a homogenous catalyst because the catalyst and reactants are both in the gaseous / same physical state.
(c)
(d) Increasing the temperature – the molecules have more energy / move faster so there are more collisions that are successful because their energy exceed activation energy
Increasing the concentration / pressure of ozone – there are more collisions between the O3 molecules and the O radicals per unit time
Increasing the intensity of the uv radiation – causes more O3 molecules to photodissociate per unit time which increases the concentration of O radicals leading to more collisions