Collision theory and effects of conditions on reaction rate

The rate and extent of chemical change • Rate of reaction

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What role does collision orientation play in reaction rate?

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Correct orientation during collisions is necessary for bond breaking/forming; many collisions remain ineffective due to incorrect orientation.

Key concepts

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Collision theory - definition

Collision theory states that particles must collide to react. Only collisions with sufficient energy and correct orientation produce products. Reaction rate depends on the number of collisions per unit time and the proportion of those collisions that are effective.

Activation energy and effective collisions

Activation energy is the minimum energy required for a collision to be effective. Only collisions with energy equal to or greater than the activation energy cause reaction. The Maxwell–Boltzmann distribution describes particle energy spread; temperature shifts this distribution so more particles exceed the activation energy at higher temperatures.

Simple proportionality in collision frequency

Collision frequency changes in proportion to particle concentration (in solutions) or number density (in gases). Doubling concentration of a reactant approximately doubles the collision frequency between reacting particles, which tends to double the rate if other factors remain constant.

Effect of concentration on rate

Higher concentration of reactants in solution increases the number of particles per unit volume. Cause: more particles per unit volume → effect: more collisions per second → higher rate. Limiting factor: if collisions remain ineffective due to insufficient energy or orientation, increasing concentration has reduced effect.

Effect of pressure on rate (gases)

Increasing pressure for reacting gases reduces the volume and increases number density of gas molecules. Cause: more molecules per unit volume → effect: more frequent collisions → higher rate. Limiting factor: pressure changes do not alter activation energy; if few collisions reach activation energy, rate increase is limited.

Effect of temperature on rate

Raising temperature increases average kinetic energy of particles and shifts the energy distribution so a larger fraction of collisions exceed the activation energy. Cause: more collisions have required energy and collisions occur slightly more often → effect: large increase in rate. Rule of thumb: moderate temperature increases often cause substantial rate increases (commonly rates double for every 10 °C rise for many reactions), though exact change depends on activation energy.

Key notes

Important points to keep in mind

Collision frequency is roughly proportional to concentration (solutions) or number density (gases).

Only collisions with energy ≥ activation energy and correct orientation are effective.

Temperature increases both average kinetic energy and the fraction of effective collisions; rate often rises strongly with temperature.

Pressure changes affect gas collision frequency but not activation energy.

Doubling concentration tends to double rate for simple collision-limited reactions; proportionality depends on reaction order.

Catalysts increase rate by lowering activation energy, not by increasing collision frequency.

Surface saturation limits the effect of further increases in concentration for heterogeneous reactions.