Activation Energy

• Activation energy is defined as the minimum energy required for a reaction to occur.

• It acts as a barrier that prevents molecules from readily reacting with each other. Only those molecules possessing energy equal to or greater than the activation energy can engage in a reaction.

• It’s vital to note that not all collisions between reactant molecules result in a reaction. The particles must collide with the correct orientation and with sufficient energy.

• The Arrhenius Equation demonstrates the relationship between the rate constant (k) of a reaction and the activation energy. The equation is typically expressed as k = Ae^(-Ea/RT), where A is the pre-exponential factor, Ea is the activation energy, R is the gas constant, and T is temperature (in Kelvin).

• A catalyst works by providing an alternative reaction pathway with a lower activation energy, hence increasing the reaction rate. However, it’s critical to understand that while a catalyst increases the rate of reaction, it does not impact the position of the equilibrium nor the enthalpy change of the reaction.

• Maxwell Boltzmann Distribution curves illustrate the distribution of molecular energies in a sample. The area under the curve represents the total number of molecules. The line often shifts to the right with increasing temperature, indicating more molecules with energy exceeding the activation energy, and thus a higher rate of reaction.

• A higher activation energy generally results in slower chemical reactions as fewer molecules possess the necessary energy to react. Lowering the activation energy through catalysis or increase in temperature can increase the rate of reaction.

• The concept of activation energy also finds usage in explaining why certain reactions require an initial input of heat (an energy source) to get started.

• The Transition State Theory proposes an ‘activated complex’ at the peak of the energy profile, where old bonds are breaking and new ones forming. This state of highest energy is fleeting, after which the reaction can proceed to products without a further energy input.

• The enthalpy profile of a reaction can provide valuable visual aid in understanding activation energy: the peak of the graph indicates activation energy for the forward reaction whereas the difference in energy between peak and energy of products gives the activation energy for the backward reaction.