6.3 - Rate Factors & Collision Theory

Collision theory

• A reaction can only happen when reacting particles collide with each other.
• Collision theory explains reaction rates in terms of these collisions.
• A collision that results in the particles reacting is a successful collision.

Two factors determine the rate of a reaction

1. Collision frequency - How often reacting particles collide. More collisions per second means more chances to react, giving a faster reaction.
2. Energy of the collision - Particles must collide with enough energy to react. The minimum energy needed is the activation energy; collisions with less energy than this do not produce a reaction.

Factors that affect reaction rate

Temperature

Increasing the temperature increases the rate for two reasons:

1. The particles move faster, so they collide more often.
2. The particles have more energy, so a greater proportion of collisions have the activation energy.

Concentration and pressure

• Increasing the concentration of a reactant in solution puts more particles into the same volume, so they collide more often.
• Increasing the pressure of a reacting gas squeezes the particles into a smaller volume, so they collide more often.
• For many reactions, rate is directly proportional to concentration (rate ∝ concentration).

Surface area

• Breaking a solid into smaller pieces increases its surface area to volume ratio.
• This exposes more of the solid to the other reactant, so collisions happen more often.

Catalyst

• Adding a catalyst increases the rate of reaction.
• A catalyst speeds a reaction up without being used up in the process.

Catalysts

• A catalyst is a substance that speeds up a chemical reaction without being used up in the process.
• It remains chemically unchanged, so its mass is the same at the start and end of the reaction.
• A catalyst doesn't appear in the balanced equation for the reaction.
• Enzymes are biological catalysts.

How a catalyst works

• A catalyst provides an alternative reaction pathway with a lower activation energy.
• A greater proportion of colliding particles have enough energy to react, so there are more successful collisions per second.

Catalysts in industry

• Catalysts allow reactions to run at lower temperatures and pressures.
• This reduces energy costs and burns less fossil fuel, making processes cheaper and more sustainable.

Example - Decomposing hydrogen peroxide

Hydrogen peroxide decomposes very slowly on its own. Adding solid manganese(IV) oxide speeds up the reaction:

hydrogen peroxide → water + oxygen

2H₂O₂(aq) → 2H₂O(l) + O₂(g)

Enzymes as biological catalysts

• Enzymes are biological catalysts - they speed up chemical reactions in living organisms without being used up.
• Enzymes are specific to the reactions they catalyse.

Optimum conditions

• Enzymes work fastest at an optimum temperature and optimum pH.
• Away from these optimum conditions, enzyme activity decreases.

Uses of enzymes

• Inside living cells - In processes such as respiration and photosynthesis.
• In industry - Yeast enzymes convert sugar into alcohol during fermentation.

Reaction profiles with a catalyst

• A reaction profile can compare an uncatalysed reaction with a catalysed reaction on the same axes.
• Both curves start at the same reactant energy level and end at the same product energy level, so the overall energy change is unchanged.
• The catalysed reaction has a lower activation energy peak, because the catalyst provides an alternative pathway with a lower activation energy.