The Activation Energy of Catalysed Reactions

We now enquire how it is that a catalyst is able to accelerate the rate of a reaction. We may start with the concept proposed by Svante Arrhenius to describe the effect of temperature on a homogeneous (i.e. non-catalysed) gas-phase reaction he stated that reaction rate r depended on the fraction of colliding molecules that between them had more than a critical amount of energy, which he called the activation energy E. This fraction increased exponentially with temperature in line with the Boltzmann distribution fraction, so that 

It must do this by creating a new and energetically more favourable reaction path, and we can visualise this by recalling that the activation energy can also be represented as the potential energy barrier that exists between reactants and products. This is the barrier that has to be broken down, so 

If we display the temperature dependences of a reaction proceeding both homogeneously and heterogeneously catalysed as Arrhenius plots, using the equation in the form 

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