Energy of Activation
Definition and meaning of Energy of Activation in chemistry.
Energy of activation is the minimum energy that colliding reactant particles must possess for a chemical reaction to proceed, corresponding to the energy barrier between the reactants and the transition state.
In more detail
This barrier exists because existing bonds must partially stretch or break before new bonds can form, passing through a high-energy, unstable arrangement of atoms called the transition state or activated complex. At a given temperature, only a fraction of molecules collide with enough energy and proper orientation to clear this barrier, so a higher activation energy means a slower reaction rate. Catalysts speed up reactions by offering an alternative pathway with a lower activation energy, without being consumed themselves. The relationship between activation energy, temperature, and rate constant is captured quantitatively by the Arrhenius equation.
Key facts
| Symbol | Ea |
|---|---|
| Typical units | kJ/mol or kcal/mol |
| Key equation | k = Ae^(−Ea/RT) (Arrhenius equation) |
| Field | Physical Chemistry |
The decomposition of hydrogen peroxide, 2H2O2(aq) → 2H2O(l) + O2(g), has an activation energy of about 75 kJ/mol uncatalyzed, but drops to roughly 8 kJ/mol when the enzyme catalase catalyzes it, dramatically speeding up the reaction.
Frequently asked questions
How does a catalyst affect activation energy?
A catalyst lowers the activation energy by providing an alternative reaction mechanism with a lower-energy transition state, increasing the reaction rate without being permanently consumed.
Why does raising temperature increase reaction rate?
Raising temperature increases the fraction of molecules whose kinetic energy exceeds the activation energy, following the Boltzmann distribution, so more collisions succeed in forming product per unit time.