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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Quick-Fire Definitions

Rate of reaction
How quickly reactants are used up or products are formed, measured per unit time.
Activation energy (Eₐ)
The minimum energy particles must have when they collide for a reaction to occur.
Catalyst
A substance that increases reaction rate by providing an alternative pathway with lower activation energy. Not used up.
Dynamic equilibrium
When the forward and reverse rates are equal in a closed system - concentrations stay constant but both reactions continue.
Le Chatelier’s principle
If a system at equilibrium is disturbed, the position of equilibrium shifts to oppose the change (Higher Tier).
Closed system
A system where no substances can enter or leave - required for equilibrium to be established.

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.

Successful Collision BAM! Fast-moving particles Energy ≥ Activation Energy (Eₐ) Reaction occurs Unsuccessful Collision Slow-moving particles Energy < Activation Energy (Eₐ) Particles just bounce off

Left: Particles collide with sufficient energy (≥ Eₐ) to react. Right: Particles collide with insufficient energy (< Eₐ) and simply bounce apart.

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.

Large Solid Block Low Surface Area (Particles trapped inside cannot react) Broken into Powders High Surface Area (Many more exposed particles to collide)

Left: A large piece of reactant only exposes its outer faces to collisions. Right: Breaking it down exposes new surfaces, increasing the frequency of successful 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.
Progress of Reaction → Energy → Reactants Products Eₐ without catalyst Eₐ with catalyst Uncatalysed Pathway Catalysed Pathway

A catalyst provides an alternative reaction pathway with a lower activation energy, meaning more particles have sufficient energy to react.

Biological catalysts are called enzymes. They are proteins that catalyse reactions in living organisms - for example, amylase breaks down starch.

Industrial Catalysts

  • Iron - used in the Haber process (N₂ + 3H₂ ⇌ 2NH₃) to manufacture ammonia.
  • Vanadium(V) oxide (V₂O₅) - used in the Contact process to make sulfuric acid.
  • Manganese dioxide (MnO₂) - catalyses the decomposition of hydrogen peroxide.
Catalysts reduce costs in industry because they can be reused, reduce the energy needed (lower activation energy), and speed up production. However, they must be kept clean - impurities can "poison" a catalyst and stop it working.

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.

Method 1: Gas Collection Gas Syringe Measure gas volume at regular intervals Method 2: Mass Loss 120.45g Cotton Wool Gas escapes, mass decreases Method 3: Disappearing Cross Precipitate forms, cross becomes invisible

Three common methods for measuring the rate of a reaction depending on the state of the products.

Required Practical: Rate of Reaction Required Practical

Investigate how changing the concentration of sodium thiosulfate affects the rate of reaction with hydrochloric acid.

Na₂S₂O₃(aq) + 2HCl(aq) → 2NaCl(aq) + H₂O(l) + SO₂(g) + S(s)
  1. Draw a cross on white paper. Place a conical flask on top.
  2. Measure 10 cm³ of HCl into the flask.
  3. Measure 40 cm³ of Na₂S₂O₃ solution and pour it into the flask. Start the timer.
  4. Look down through the solution. Stop the timer when the cross can no longer be seen.
  5. Repeat with different concentrations of Na₂S₂O₃ (diluting with water to keep total volume constant).
The sulfur precipitate makes the solution go cloudy. A shorter time = faster rate. Rate is proportional to 1/time.

Calculating mean rate of reaction

In an experiment, 48 cm³ of gas was collected in 60 seconds. Calculate the mean rate.

Formula: mean rate = volume of gas ÷ time

Calculation: 48 ÷ 60 = 0.80 cm³/s

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.
Time → Amount of Product → (or reactant used) Final amount of product (reaction is complete) Reaction A finishes Reaction B finishes Steeper initial gradient = Faster rate (More frequent successful collisions) Higher Temp / Conc. (Faster) Lower Temp / Conc. (Slower)

Rate graphs show how the amount of product changes over time. The steeper the line, the faster the reaction.

Interpreting two rate curves

Two experiments use the same mass of marble chips and volume of HCl. Experiment A is at 20°C; Experiment B is at 40°C. Both curves level off at the same volume of CO₂. Explain the differences.

Gradient: Curve B is steeper because at 40°C particles have more kinetic energy, so collisions are more frequent and more energetic → more successful collisions per second.

Final volume: Both level off at the same volume because the same amount of reactant was used - temperature does not change the total product, only how fast it forms.

Time to finish: Curve B levels off sooner because the reaction is faster at higher temperature.

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. If you are asked to calculate rate from a graph, draw a tangent at the required time and calculate gradient = Δy/Δx.

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.
Reactants Products Forward Rate Reverse Rate Rates are EQUAL Constant Level Constant Level

In dynamic equilibrium, the forward and reverse reactions happen at the exact same rate. The amount of reactants and products remains constant because they are being formed as fast as they are used up.

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.

Applying Le Chatelier’s Principle: The Haber Process

N₂(g) + 3H₂(g) ⇌ 2NH₃(g)     (forward reaction is exothermic)

The Haber process is used to manufacture ammonia. Conditions are chosen as a compromise between yield and rate:

Worked Example: Predicting equilibrium shifts

For N₂ + 3H₂ ⇌ 2NH₃ (exothermic forward), predict what happens when:

1. Temperature is increased: Equilibrium shifts LEFT (endothermic direction) to absorb extra heat → less NH₃, lower yield. However, rate is faster.

2. Pressure is increased: Left side has 4 moles of gas (1 + 3), right side has 2 moles. Equilibrium shifts RIGHT (fewer moles) → more NH₃, higher yield.

3. More N₂ is added: Equilibrium shifts RIGHT to use up the extra N₂ → more NH₃ produced.

Compromise conditions: 450°C (moderate - low temp gives better yield but reaction is too slow), 200 atm (high pressure for better yield), iron catalyst (speeds up reaction without affecting yield). These give about 15% yield per pass - unreacted gases are recycled.

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