Reactant A can participate in both of these reactions: 1) A + B --> C + D Grxn = –15.6 kJ 2) A + E --> F + G Grxn = –20.5 kJ Reaction 1 has a much higher rate than Reaction 2 under the same reaction conditions. Explain this observation in terms of activation energy

Reaction 1 has a lower activation energy than reaction 2

Explanation:

A close look at the both reactions show that the change in free energy for both reactions are negative showing that both are spontaneous reactions as shown.

However, according to Arrhenius equation, the rate of reaction, which is given by the rate constant depends on the activation energy of the reaction and the reaction temperature.

Since it was explicitly stated in the question that the both reactions were allowed to proceed under the same conditions, that means that the temperature of the both processes are the same.

The difference in rate of reaction now stems from difference in activation energy. The higher the activation energy the lower the rate of reaction. Hence reaction 1 has a lower activation energy than reaction 2 as shown by the information in the question about the relative rates of both reactions.