3.2.2 (c,d) Catalysts and Rate

Syllabus

(i) in increasing reaction rate without being used up by the overall reaction

(ii) in allowing a reaction to proceed via a different route with lower activation energy, as shown by enthalpy profile diagrams

{Details of processes are not required.}

(d) (i) explanation of the terms homogeneous and heterogeneous catalysts

(ii) explanation of catalysts' economic importance & benefits for sustainability by lowering temperatures & reducing energy demand from fossil fuel combustion with resulting reduction in CO₂ emissions

{Benefits to the environment of improved sustainability weighed against toxicity of some catalysts.}

What Does this Mean?

A catalyst is still a substance that speeds up a reaction without being used up -which is the GCSE definition.

The only addition to the definition at A level is a substance that speeds up a reaction without being used up by providing an alternative reaction-route with a lower Activation Energy.

What this means in practice is that the catalyst changes the way that the reaction happens - often by allowing it to pass through one or more intermediate stages.

A High Activation Energy slows a reaction because no matter how high the collision frequency the vast majority of collisions will be unsuccessful (unless the temperature is very high)

But a Lower Activation Energy allows the collisions to be successful much more frequently - even at lower temperatures.

So for any given temperature the catalysed reaction is bound to be faster than the uncatalysed reaction.

Examiners like to ask for reasons to use catalysts.

The most obvious one is to allow the reaction to happen quickly at low temperatures - thereby saving energy and lowering the need to burn fossil fuels, emit CO₂ etc.

However, they often want more than one reason and the other usual answers are that catalysed reactions are often more specific so they reduce side-reactions and the amount of waste.

Which is one way to improve atom-economy and avoid producing possibly toxic and unwanted products.

Although some catalysts can themselves be toxic this is less of a problem since the catalyst should not need replacing very frequently.

Also, if you recall catalytic converters from GCSE - many of the unwanted emissions would react in the atmosphere if released eventually.

But don't have time to do so in an exhaust-pipe unless they pass over a catalyst.

In the graph you can see the best catalysts almost eliminate un-burned fuel (hydrocarbons) as well as Nitrogen Oxides, and doing a good job on the Carbon Monoxide.

Heterogeneous Catalysts

Hetero- means different.

A Heterogeneous catalyst is a catalyst in a different phase to the reactants.

This generally means a solid catalyst that allows gases or solutions to react faster.

You shouldn't be asked how this works but it is the same explanation you may have been given about the catalytic converter in previous years.

One of the reactants adsorbs (sticks to) onto the surface of the catalyst.

This lowers the strength of its bonds so that when the other reactant collides the Activation Energy is lower - since this is often said to be the energy needed to break all the bonds in the reactants.

When the reaction has happened the products desorbs - drifts away - leaving space for more reactant to adsorb.

Heterogeneous catalysts need to be very finely divided or in fine meshes to provide a big enough surface area for the reactants to stick to.

Although, sometimes heterogeneous catalysts sometimes react with a reactant to make an intermediate that then reacts again to form the desired product and remake the catalyst.

For example, the Contact process involves the reaction:

SO₂ + 1/2 O₂ → SO₃ - which is very difficult to achieve a decent yield for.

But with a vanadium oxide catalyst the SO₂ has another way to gain Oxygen with a lower Activation Energy.

SO₂ + V₂O₅ → SO₃ + V₂O₄

The V₂O₄ can then be reoxidised by the Oxygen

1/2 O₂ + V₂O₅ → V₂O₅

So the overall reaction is still SO₂ + 1/2 O₂ → SO₃and the catalyst has not been destroyed.

Homogeneous Catalysts

Homogeneous catalysts may also bind temporarily to one of the reactants.

But more usually they react with the reactants to form a totally new intermediate product or more than one.

Because a homogeneous catalyst is in the same phase it mixes without the need for any meshes.

But this means that they have to be separated from the products which may add expense to a process and mean that the catalyst is destroyed in the process.

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Exam-style Questions

1. The gas-phase reaction between Hydrogen and Chlorine is very slow at room temperature.

H_2(g) + Cl_2(g)→ 2HCl_(g)

(a) Define the term activation energy.

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(b) Give one reason why the reaction between Hydrogen and Chlorine is very slow at room temperature.

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.........................................................................................................................................................................(1)

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(d) Explain why small increases in temperature lead to large increases in rate of reaction between Hydrogen & Chlorine.

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(e) Give the meaning of the term catalyst.

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(f) Suggest one reason why a solid catalyst for a gas-phase reaction is often in the form of a powder.

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(Total 9 marks)

Answers

1. (a) minimum energy 1

to start a reaction/ for a reaction to occur/ for a successful collision 1

(b) activation energy is high / few molecules/particles have sufficient

energy to react/few molecules/particles have the required activation energy 1

(or breaking bonds needs much energy)

therefore collide more often 1

(d) many 1

more molecules have energy greater than activation energy (QoL) 1

(e) speeds up a reaction but is chemically unchanged at the end 1

(f) increases the surface area 1

[9]

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