Topic 6 of 10

Rate & Extent of Chemical Change

Learn what controls reaction speed — from collision theory and catalysts to reversible reactions and dynamic equilibrium.

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Rate of Reaction

The rate of a chemical reaction is a measure of how quickly reactants are used up or how quickly products are formed.

rate = amount of product formed ÷ time

Some reactions are very fast (explosions), while others are very slow (rusting of iron).

Collision Theory

For a chemical reaction to occur, particles must:

  1. Collide with each other.
  2. Collide with sufficient energy — at least the activation energy (Eₐ).

Collisions that have enough energy to react are called successful collisions.

Not every collision leads to a reaction. Only those with energy ≥ Eₐ (the activation energy) are successful.

Factors Affecting Rate

Temperature

Increasing temperature increases rate. Particles move faster (more kinetic energy), so collisions are more frequent AND more energetic. A greater proportion of collisions exceed the activation energy.

Concentration (or Pressure for gases)

Increasing concentration increases rate. There are more particles in the same volume, so collisions are more frequent.

Surface Area

Increasing surface area increases rate. Using smaller pieces (or powders) exposes more reactant particles on the surface, so there are more opportunities for collisions.

When explaining a rate factor, always use collision theory: (1) state what happens to the frequency or energy of collisions, (2) explain how this affects the number of successful collisions.

Catalysts

A catalyst is a substance that increases the rate of a reaction without being used up. It works by providing an alternative reaction pathway with a lower activation energy.

  • Catalysts are not used up — they can be reused.
  • They do not change the amount of product — they only make the reaction faster.
  • They are often specific to a particular reaction.
Biological catalysts are called enzymes. They are proteins that catalyse reactions in living organisms — for example, amylase breaks down starch.

Measuring Rate of Reaction

Method 1: Gas Collection

Collect the gas produced in a gas syringe. Measure the volume of gas at regular time intervals.

Method 2: Mass Loss

Place the reaction on a balance. As gas escapes, the mass decreases. Record the mass at regular time intervals.

Method 3: Disappearing Cross

For reactions that produce a precipitate (e.g., sodium thiosulfate + hydrochloric acid). Time how long it takes for a cross underneath the flask to become invisible as the solution turns cloudy.

Rate Graphs

The steeper the graph's gradient, the faster the rate. The gradient decreases over time as reactants are used up. The graph eventually levels off when the reaction is complete.

Interpreting Changes on Graphs

  • Higher temperature or concentration: Steeper initial gradient, but same final amount of product (same amount of reactant).
  • Using a catalyst: Steeper gradient, same final product.
  • Using more reactant: Steeper gradient AND more total product.
The graph always starts steep and flattens out. The flat part shows the reaction has finished. A catalyst makes the graph steeper but doesn't change where it levels off.

Reversible Reactions

A reversible reaction is one that can proceed in both directions — products can re-form the reactants.

A + B ⇌ C + D

The ⇌ symbol indicates a reversible reaction.

Energy in Reversible Reactions

If the forward reaction is exothermic, the reverse reaction is endothermic — and they involve the exact same amount of energy.

Hydration of anhydrous copper sulfate:

CuSO₄ + 5H₂O ⇌ CuSO₄·5H₂O

Forward: white → blue (exothermic). Reverse: blue → white (endothermic, by heating).

Dynamic Equilibrium

In a closed system (nothing can enter or leave), a reversible reaction reaches dynamic equilibrium. At equilibrium:

  • The rate of the forward reaction equals the rate of the reverse reaction.
  • The concentrations of reactants and products remain constant (but are not necessarily equal).
At equilibrium, both reactions are still happening — it's "dynamic." But the overall concentrations don't change because the rates are balanced.

Le Chatelier's Principle (HT)

If a system at equilibrium is subjected to a change in conditions, the position of equilibrium will shift to oppose the change.

Effect of Temperature

  • Increase temperature: Equilibrium shifts in the endothermic direction (to absorb the extra heat).
  • Decrease temperature: Equilibrium shifts in the exothermic direction.

Effect of Pressure (for gas reactions)

  • Increase pressure: Equilibrium shifts to the side with fewer moles of gas.
  • Decrease pressure: Equilibrium shifts to the side with more moles of gas.

Effect of Concentration

  • Increase concentration of a reactant: Equilibrium shifts to the right (forward), producing more product.
  • Increase concentration of a product: Equilibrium shifts to the left (backward).
A catalyst does NOT affect the position of equilibrium — it speeds up both forward and reverse reactions equally. It only makes equilibrium reached faster.