AP Chemistry Unit 5 Study Notes

AP Chemistry 5.4: Collision Theory and Activation Energy

Explain successful collisions, temperature effects, and the Arrhenius equation.

Aligned to Kinetics from the current College Board AP Chemistry course outline. Exam weighting for this unit: 7%-9% of the multiple-choice score range listed by College Board.

Study these notes

Start with each main idea, then follow the indented explanations and worked examples. Try the next calculation before reading its answer.

Organized from the provided Unit 5 study document. Further study: Khan Academy.

Activation Energy and Reaction Rate

  • Chemical reactions require particles to collide, but not every collision produces a reaction.

    • According to collision theory, reacting particles generally must collide:
    • with enough energy, and
    • with an appropriate orientation
    • for bonds to rearrange successfully.
    • This explains why a container can contain countless molecular collisions without every collision causing a chemical reaction.
Activation Energy

  • The activation energy, written Eₐ, is the minimum energy barrier that reacting particles must overcome for a particular reaction pathway.

  • Even an exothermic reaction can require activation energy.

    • Imagine a ball sitting in a valley separated from another valley by a hill. Even if the second valley is lower, the ball must first be pushed uphill before it can roll down.
    • Chemical reactions behave similarly.

  • Reactants must reach a high-energy arrangement called the transition state before becoming products.

Collision Theory

  • A successful collision requires enough energy to reach the transition state.

    • If two particles collide with too little kinetic energy, they simply separate again.

  • Particles may also need the correct orientation.

    • For example, if one part of molecule A needs to collide with a particular part of molecule B, hitting the wrong ends together may not produce a reaction even if the collision energy is high enough.
    • Therefore:
    • successful collision = sufficient energy + suitable orientation
Temperature and Reaction Rate

  • Increasing temperature causes reaction rates to rise significantly.

    • There are two important reasons:

  • Particles generally move faster, which can increase collision frequency.

  • More importantly, a larger fraction of particles have kinetic energies greater than or equal to Eₐ.

    • The second effect is usually the more important explanation for why temperature strongly affects reaction rate.
Maxwell–Boltzmann Distribution

  • A Maxwell–Boltzmann distribution shows the range of kinetic energies present in a collection of particles at a particular temperature.

  • At a higher temperature, the distribution becomes broader and shifts so that more particles occupy the high-energy region.

  • If a vertical line represents Eₐ, the area of the curve to the right of Eₐ represents particles with enough energy to overcome the activation barrier.

    • When temperature increases, that area becomes larger.
    • Therefore:
    • higher T → larger fraction with E ≥ Eₐ → more successful collisions → faster reaction

  • Do not say that increasing temperature lowers activation energy. It normally does not.

  • Temperature changes the distribution of particle energies.

  • A catalyst lowers the activation-energy pathway.

    • Those are different effects.
The Arrhenius Equation

  • The relationship between rate constant, activation energy, and temperature is described by the Arrhenius equation:

    • k = Ae^(−Eₐ/RT)
    • where:
    • k = rate constant
    • A = frequency/pre-exponential factor
    • Eₐ = activation energy
    • R = gas constant
    • T = Kelvin temperature
    • The exponential term represents the strong effect activation energy and temperature have on k.

  • If temperature increases, the negative exponent becomes less negative, so k increases.

  • If activation energy is smaller, k is larger, assuming other factors are comparable.

    • This is why even a moderate increase in temperature can cause a major increase in reaction rate.
Two-Temperature Arrhenius Form

  • A useful rearrangement is:

    • ln(k₂/k₁) = −Eₐ/R(1/T₂ − 1/T₁)
    • This can be used when comparing rate constants at two temperatures.

  • Make sure temperature is in Kelvin.

  • If Eₐ is given in kJ/mol but R is in J/(mol·K), convert Eₐ to J/mol before substituting.

Intermediates on an Energy Diagram

Common Mistakes

  • Do not say all collisions produce reactions.

  • Do not say increasing temperature lowers Eₐ.

  • Higher temperature increases the fraction of particles that can overcome Eₐ.

  • Do not confuse activation energy with the overall energy change of the reaction.

  • A reaction can be exothermic and still have a large activation energy.

  • Do not confuse intermediates with transition states on an energy diagram.

Intermediates on an Energy Diagram

Remember This

  • A reaction needs particles to get over an energy hill.

    • Higher temperature gives more particles enough energy to get over that hill.
    • A catalyst gives them a lower hill.