Chemistry · Ch 6 — Chemical Kinetics

Effect of a catalyst on the rate of reaction

6.8

Effect of a catalyst on the rate of reaction

A catalyst is a substance added to the reactants that increases the rate of the reaction without itself being consumed in the reaction. Consider

\mathrm{2\,KClO_3(s) \xrightarrow{\;MnO_2\;} 2\,KCl(s) + 3\,O_2(g)}

Here $\mathrm{MnO_2}$ is the catalyst. It has been observed that the decomposition rate increases with the addition of the catalyst.

Note

The book prints this equation as "2 KClO₃(s)+ 3O₂(g) $\xrightarrow{\;MnO_2\;}$ 2 KCl(s)" — with the oxygen on the reactant side and missing from the products. The decomposition of potassium chlorate in fact produces the oxygen, as the corrected equation above shows.

A catalyst provides an alternative pathway associated with lower activation energy.

Fig. 6.10 compares the potential energy barriers for the catalysed and uncatalysed reactions. The barrier for the uncatalysed reaction, $(E_a)_1$ , is larger than that for the same reaction in the presence of a catalyst, $(E_a)_2$ .

Figure 6.10Potential energy diagram comparing the tall activation energy barrier of an uncatalysed reaction with the lower barrier of the same reaction on the catalysed pathway, both curves starting at the same reactant level and ending at the same product level.

Fig. 6.10 — Potential energy diagram comparing the tall activation energy barrier of an uncatalysed reaction with the lower barrier of the same reaction on the catalysed pathway, both curves starting at the same reactant level and ending at the same product level.

Drawn by us to help you understand the concept clearly, and verified to make sure it's accurate. For exams, practice from your textbook's own diagram.

What this figure shows. Two energy pathways connect the same reactant level to the same product level along the reaction coordinate. The taller curve is the uncatalysed reaction, with barrier height $(E_a)_1$ ; the shorter curve is the catalysed pathway the catalyst provides, with the smaller barrier $(E_a)_2$ . Because $(E_a)_1 > (E_a)_2$ , far more collisions can clear the catalysed barrier — the catalyst speeds the …

Consider the decomposition of $\mathrm{H_2O_2}$ in aqueous solution, catalysed by $\mathrm{I^-}$ ions:

\mathrm{2H_2O_2({\it l}) \xrightarrow{\;I^-\;} 2\,H_2O({\it l}) + O_2(g)}

At room temperature the rate of the reaction is slower in the absence of the catalyst, with its activation energy being 76 kJ mol⁻¹. In the presence of the iodide ion catalyst $\mathrm{I^-}$ the reaction is faster, since the activation energy decreases to 57 kJ mol⁻¹. …

Figure 6.11Energy distribution curve of molecules with the two activation-energy thresholds marked: the lower catalysed threshold admits a visibly larger shaded fraction of activated molecules than the higher uncatalysed threshold.

$Fig. 6.11 — Energy distribution curve of molecules with the two activation-energy thresholds marked: the lower catalysed threshold admits a visibly larger shaded fraction of activated molecules than the higher uncatalysed threshold.$

Drawn by us to help you understand the concept clearly, and verified to make sure it's accurate. For exams, practice from your textbook's own diagram.

What this figure shows. A single distribution curve ( $\mathrm{d}n/N$ , the fraction of molecules) against Energy, with the two threshold energies on the axis: $(E_a)_2$ — the activation energy in the presence of the catalyst — sits to the LEFT of $(E_a)_1$ , the activation energy before use of the catalyst. The shaded area beyond each threshold is the fraction of molecules energetic enough to react: lowering the threshold from $(E_a)_1$ to $(E_a)_2$ sweeps a much larger f …

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