Kinetics
Collision Theory
Reactions occur when molecules collide with sufficient energy and right orientation
Collision theory states that chemical reactions occur when molecules collide with: (1) sufficient kinetic energy (≥ activation energy), and (2) correct geometric orientation. Most collisions don't result in reaction. Rate = (collision frequency) × (fraction with energy ≥ Ea) × (fraction with correct orientation). Explains: why rate depends on concentration (more particles → more collisions), temperature (higher T → more energetic collisions), and surface area (more contact → more collisions). Foundation of chemical kinetics.
- Conditions(1) Energy ≥ Ea, (2) correct orientation
- Most collisionsDon't result in reaction
- Concentration effectMore particles → more collisions
- Temperature effectHigher T → more particles with E ≥ Ea
- Orientation factorSteric factor (~10⁻¹ to 10⁻⁵)
- Foundation1916-1920 development
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Why collision theory matters
- Kinetics foundation. Why concentration and T matter.
- Catalysis. Lower Ea → more effective collisions.
- Reaction mechanisms. Visualize what happens.
- Physical chemistry. Quantitative kinetics.
- Atmospheric. Reactions in upper atmosphere.
- Combustion. Why heat speeds reactions.
- Biology. Enzyme efficiency vs uncatalyzed.
Common misconceptions
- Every collision reacts. Most don't.
- Energy enough. Orientation also required.
- Temperature only effect on collisions. Mostly on energy distribution.
- Collision theory exact. Approximation; transition state theory more accurate.
- Catalyst increases collisions. Lowers Ea instead.
- Solid reactions don't apply. Surface area gives collision sites.
Frequently asked questions
How does collision theory work?
Reactions require: (1) molecules to collide. (2) Collision energy ≥ activation energy (so bonds can break). (3) Correct orientation (so bonds can form correctly). Rate proportional to: collision frequency × probability of sufficient energy × probability of correct orientation. Most collisions ineffective; only fraction lead to reaction.
How is collision frequency calculated?
For gases: Z = N(A)N(B) × σ × sqrt(8RT/πμ), where σ is collision cross-section, μ is reduced mass. Collisions ~10²⁹-10³⁰ per second for typical gas. Most collisions don't react. Even for reactive system, only small fraction lead to product.
What's the steric factor?
Probability of correct orientation. Some reactions: nearly any orientation works (e.g., simple atom transfer; steric factor ≈ 1). Others: very specific orientation needed (e.g., enzyme-substrate; complex molecules; steric factor ≈ 10⁻³ to 10⁻⁵). Known as "p-factor"; multiplies rate constant. Reflects: geometric requirements of transition state.
How does temperature affect collision theory?
Two effects. (1) Slightly more collisions per second (k_T ~ T^0.5). (2) Much larger fraction with energy ≥ Ea (Boltzmann distribution → exp(-Ea/RT)). Total rate increase depends on both, but exponential factor dominates. Doubling T can change rate by 10× or more (depending on Ea).
What's the activation energy in collision theory?
Minimum collision energy (kinetic energy of approach) for reaction. Below: no reaction. Above: reaction possible (subject to orientation). Ea reflects the energy "barrier" between reactants and products. Visualized in reaction coordinate diagram: peak between reactant valley and product valley.
Why are most collisions ineffective?
Most don't have enough energy. Only those above Ea react. Fraction with E ≥ Ea = exp(-Ea/RT). For Ea = 50 kJ/mol at room T: ~10⁻⁹ of collisions have enough energy. Plus: orientation factor reduces it more. So: only ~10⁻⁹ to 10⁻¹⁴ of collisions actually react.
How does it explain catalysis?
Catalyst lowers Ea. More collisions have sufficient energy (exp(-Ea/RT) larger). Total rate higher. Doesn't change collision frequency or orientation. Just gives access to lower-energy pathway. Same overall reaction; different transition state.