Chemical reactions occur when particles collide with enough energy and the correct orientation to form new bonds. Activation energy is the minimum energy needed to reach the transition state, where old bonds are partly broken and new bonds are partly formed. This idea matters because a reaction may be thermodynamically favorable but still slow if the activation energy is high.
The Arrhenius equation connects this energy barrier to the observed reaction rate.
Key Facts
- Activation energy, Ea, is the minimum energy needed for reactants to reach the transition state.
- Arrhenius equation: k = Ae^(-Ea/RT).
- Taking logs gives: ln k = ln A - Ea/(RT).
- Two-temperature form: ln(k2/k1) = -Ea/R(1/T2 - 1/T1).
- A catalyst lowers Ea by providing an alternate reaction pathway, but it does not change ΔH or the equilibrium constant.
- On an energy profile, ΔH = Eproducts - Ereactants and Ea = Etransition state - Ereactants for the forward reaction.
Vocabulary
- Activation energy
- The minimum energy that colliding particles must have to form the transition state and continue to products.
- Arrhenius equation
- An equation, k = Ae^(-Ea/RT), that relates a reaction rate constant to activation energy and temperature.
- Transition state
- A short-lived, high-energy arrangement of atoms at the top of the energy barrier during a reaction.
- Catalyst
- A substance that increases reaction rate by lowering the activation energy through an alternate pathway without being consumed overall.
- Rate constant
- The proportionality constant k in a rate law that shows how fast a reaction proceeds under specific conditions.
Common Mistakes to Avoid
- Confusing activation energy with enthalpy change is wrong because Ea is the height from reactants to the transition state, while ΔH is the energy difference between products and reactants.
- Using Celsius in the Arrhenius equation is wrong because temperature must be in kelvins for k = Ae^(-Ea/RT) to work correctly.
- Thinking a catalyst makes products more stable is wrong because a catalyst lowers the pathway barrier but does not change reactant energy, product energy, or ΔH.
- Assuming every collision causes a reaction is wrong because particles must collide with enough energy and the correct orientation to reach the transition state.
Practice Questions
- 1 A reaction has Ea = 55.0 kJ/mol and A = 2.0 x 10^12 s^-1. Calculate k at 300 K using R = 8.314 J/(mol K).
- 2 For a reaction with Ea = 75.0 kJ/mol, k1 = 0.018 s^-1 at 298 K. Use ln(k2/k1) = -Ea/R(1/T2 - 1/T1) to find k2 at 318 K.
- 3 A catalyst is added to an exothermic reaction. Explain how the energy profile changes and identify which quantities stay the same: Ea, ΔH, product energy, and equilibrium position.