Rates of Chemical Reactions
Understand why some reactions are fast and others slow by exploring collision theory, activation energy and the factors that control how quickly reactants become products.
Collision Theory
Collision theory states that for a chemical reaction to occur, reactant particles must collide with each other. However, not all collisions lead to a reaction. A successful collision requires two conditions to be met simultaneously.
Sufficient Energy
Particles must collide with energy equal to or greater than the activation energy (Ea) -- the minimum energy needed to break bonds and start the reaction.
Correct Orientation
Particles must collide with the correct geometric alignment so that bonds can form between the right atoms.
Rate of reaction = the change in concentration of a reactant or product per unit time. It can be measured by the volume of gas produced, mass loss, colour change, or change in pH over time.
Energy Profile Diagrams
An energy profile diagram shows the energy changes during a reaction. The y-axis is energy and the x-axis is the progress of the reaction.
Exothermic Reaction Energy Profile
Energy
Reaction Progress
Reactants
Ea (peak)
Products
In an exothermic reaction, products are at a lower energy than reactants (ΔH is negative).
Catalyst effect: A catalyst provides an alternative reaction pathway with a lower activation energy. It increases the rate of reaction without being consumed. On an energy profile, a catalyst lowers the peak but does not change the energy of reactants or products.
Factors Affecting Reaction Rate
Any change that increases the frequency of successful collisions will increase the rate of reaction.
Temperature ↑
Higher temperature gives particles more kinetic energy. They move faster, collide more frequently, and a greater proportion of collisions exceed Ea.
Concentration ↑
More particles per unit volume means more frequent collisions. For gases, increasing pressure has the same effect.
Surface Area ↑
Smaller particles (e.g. powder vs lump) have more exposed surface for collisions. This is why a dust explosion can occur but a solid block will not explode.
Catalyst
Provides an alternative pathway with lower Ea. More collisions now have sufficient energy to react. The catalyst is regenerated and not used up.
Measuring Reaction Rate in Practice
Gas collection: Measure volume of gas produced at regular time intervals using a gas syringe or inverted measuring cylinder.
Mass loss: Place reaction vessel on a balance and record mass at intervals (useful when gas escapes).
Colour change: Use a colorimeter to track how quickly a colour appears or disappears.
Key Vocabulary
Activation Energy (Ea)
The minimum energy that colliding particles must possess for a reaction to occur.
Catalyst
A substance that increases the rate of reaction by providing an alternative pathway with lower activation energy, without being consumed.
Rate of Reaction
The speed at which reactants are converted into products, measured as change in concentration per unit time.
Collision Theory
The theory that reactions occur when particles collide with sufficient energy and correct orientation.
Worked Examples
Explain why increasing temperature increases reaction rate, using collision theory.
Step 1: Increasing temperature gives particles more kinetic energy, so they move faster.
Step 2: Faster particles collide more frequently.
Step 3: A greater proportion of collisions have energy ≥ Ea, so more collisions are successful. Both effects increase the rate.
Marble chips (CaCO3) react with HCl. Would using powdered CaCO3 change the rate? Explain.
Step 1: Powdered CaCO3 has a much greater surface area than marble chips of the same total mass.
Step 2: More surface area means more reactant particles are exposed to HCl molecules at the surface.
Answer: The reaction rate would increase because there are more collision sites, leading to more frequent successful collisions per unit time.
A catalyst is added to the decomposition of hydrogen peroxide. The same amount of oxygen is produced but faster. Explain why the total volume of gas does not change.
Step 1: A catalyst lowers the activation energy but does not change the overall equation: 2H2O2 → 2H2O + O2.
Step 2: The same moles of reactant produce the same moles of product (stoichiometry is unchanged).
Answer: A catalyst increases the rate but not the yield. The reaction reaches completion faster, but the final volume of O2 is the same.
Knowledge Check
Select the correct answer for each question. Click "Check Answer" to see if you are right.
Question 1
According to collision theory, a successful collision requires:
Question 2
A catalyst increases reaction rate by:
Question 3
Why does increasing concentration increase reaction rate?
Question 4
On an energy profile diagram, activation energy is represented by:
Question 5
Increasing temperature increases reaction rate because:
Key Concepts Summary
- ●Collision theory: reactions need collisions with sufficient energy (≥ Ea) and correct orientation.
- ●Activation energy is the minimum energy for a successful collision; shown as the peak on an energy profile.
- ●Rate increases with higher temperature, concentration, surface area and addition of a catalyst.
- ●A catalyst lowers Ea via an alternative pathway; it increases rate but not yield.
- ●Reaction rate can be measured by tracking gas volume, mass loss, or colour change over time.