MediumChemistryHigh School (11–12)Chemical Kinetics › Temperature, the Arrhenius Equation and Collision TheoryNumerical

A catalyst lowers the activation energy of a reaction from 85.0 kJ/mol to 60.0 kJ/mol…

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A catalyst lowers the activation energy of a reaction from $85.0\ \text{kJ/mol}$ to $60.0\ \text{kJ/mol}$ . The pre-exponential factor stays the same. Find the power of ten by which the rate constant is multiplied at $320\ \text{K}$ . A step takes $60$ minutes without the catalyst. How long does it take with it?

Need a reminder? The key ideaShow

A catalyst lowers the barrier by $\Delta E_a$ in both directions. If $A$ stays the same, $\dfrac{k_{cat}}{k} = e^{\Delta E_a/RT}$ . $\Delta H$ and the equilibrium constant do not change. In a mechanism, a catalyst is used in one step and made again in a later one.

The worked solution

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Work it out on paper first. It is fine to get stuck or make a mistake: that is how we learn. When you are ready, go through the solution one step at a time, and check each step against your own.

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← For a reaction in solution, the collision factor is calculated as Z = 2.5 × 10¹¹ L mol⁻¹ s⁻¹…A reaction has a forward activation energy of 94 kJ/mol and Δ H = -38 kJ/mol… →

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