Collision theory explains why chemical reactions happen at different speeds by focusing on the tiny collisions between reacting particles. For a reaction to occur, particles must collide with enough energy and the correct orientation. This matters because reaction rate affects everything from digestion and rusting to industrial chemical production.
The theory gives a particle-level way to predict how changes in conditions will speed up or slow down a reaction.
Most collisions do not form products because the particles either hit too gently or line up in the wrong way. The minimum energy needed for a successful collision is called activation energy, and only particles with enough kinetic energy can overcome it. Temperature, concentration, pressure, surface area, and catalysts change the number or effectiveness of collisions.
A catalyst speeds up a reaction by providing a pathway with lower activation energy, so a larger fraction of collisions becomes successful.
Key Facts
- A reaction occurs only when particles collide with enough energy and the correct orientation.
- Activation energy, Ea, is the minimum energy needed for reactants to form products.
- Higher temperature increases average kinetic energy and increases the fraction of particles with E >= Ea.
- Higher concentration or pressure increases collision frequency, which usually increases reaction rate.
- Rate is proportional to the number of effective collisions per unit time.
- Arrhenius equation: k = Ae^(-Ea/RT), where k is the rate constant.
Vocabulary
- Collision theory
- Collision theory is the model that explains reaction rate by the frequency and success of collisions between reacting particles.
- Effective collision
- An effective collision is a collision that has enough energy and the correct orientation to produce a chemical reaction.
- Activation energy
- Activation energy is the minimum energy that colliding particles must have to begin forming products.
- Orientation factor
- The orientation factor describes how the alignment of particles during a collision affects whether bonds can break and form correctly.
- Catalyst
- A catalyst is a substance that increases reaction rate by lowering the activation energy without being consumed overall.
Common Mistakes to Avoid
- Thinking every collision causes a reaction is wrong because most collisions lack enough energy, the correct orientation, or both.
- Saying temperature only makes particles collide more often is incomplete because the bigger effect is that more particles have enough energy to exceed Ea.
- Assuming a catalyst gives particles more energy is wrong because a catalyst lowers the activation energy by providing an alternate reaction pathway.
- Ignoring molecular orientation is wrong because even high-energy collisions can fail if the reacting parts of the molecules do not meet correctly.
Practice Questions
- 1 In a gas reaction, particles collide 2.0 x 10^8 times per second, but only 0.050% of collisions are effective. How many effective collisions occur each second?
- 2 A reaction has activation energy Ea = 50,000 J/mol. At T = 300 K, calculate the value of e^(-Ea/RT) using R = 8.314 J/(mol K).
- 3 Two samples of the same reactants are tested. One is warmed from 25 degrees Celsius to 45 degrees Celsius, and the other is given a catalyst at 25 degrees Celsius. Explain how each change increases the reaction rate using collision theory.