Rates of Reaction
What is Rate of Reaction?
The rate of reaction is the amount of reactant used up or product formed per unit time:
Rate of reaction = amount of reactant used / time Rate of reaction = amount of product formed / time
Amount can be measured as mass (g), volume (cm3) or moles. Units are therefore g/s, cm3/s or mol/s.
Factors Affecting Rate of Reaction
Four factors affect the rate of a reaction:
- Surface area of a solid reactant
- Concentration of a solution (or pressure of a gas)
- Temperature
- Catalyst
Increasing any of these (except a catalyst, which is added or not) increases the rate of reaction.
Effect of Surface Area
A solid with a larger surface area (smaller particle size) reacts faster.
Increasing the surface area means that more particles of the solid are exposed to the other reactant, leading to more frequent and successful collisions.
Graph: A reaction with smaller particles (larger surface area) has a steeper initial gradient and reaches completion sooner, but produces the same total amount of product.
Effect of Concentration (and Pressure)
Increasing concentration (for solutions) or increasing pressure (for gases) increases the rate of reaction.
More particles are present in the same volume, so collisions occur more frequently and there are more successful collisions per second.
Graph: Higher concentration gives a steeper initial gradient and reaches completion sooner, producing the same total amount of product.
Effect of Temperature
Increasing temperature increases the rate of reaction.
At higher temperatures, particles have more kinetic energy. They move faster, so collisions are more frequent and more energetic. A greater proportion of particles have energy greater than or equal to the activation energy, so more collisions are successful.
As a rough guide, for every 10 degrees Celsius increase in temperature, the rate of reaction approximately doubles.
Graph: Higher temperature gives a steeper initial gradient and reaches completion sooner.
Effect of a Catalyst
Adding a catalyst increases the rate of reaction. The catalyst provides an alternative reaction pathway with a lower activation energy, so more particles have sufficient energy to react.
A catalyst is not used up in the reaction; it remains chemically unchanged at the end. Only small amounts are needed.
Collision Theory
Collision theory explains why the above factors affect rate of reaction. For a reaction to occur, particles must:
- Collide with each other
- Have sufficient energy (at least the activation energy) when they collide
Any factor that increases the frequency of collisions OR increases the proportion of collisions with sufficient energy will increase the rate of reaction.
| Factor | How it increases collision frequency or energy |
|---|---|
| Increase concentration | More particles per unit volume; more frequent collisions |
| Increase pressure | Same number of particles in a smaller volume; more frequent collisions |
| Increase surface area | More exposed particles; more frequent collisions with other reactant |
| Increase temperature | Particles have more kinetic energy; more frequent AND more energetic collisions; more particles exceed activation energy |
| Add catalyst | Provides lower activation energy pathway; more particles have sufficient energy to react |
Activation Energy
The activation energy (Ea) is the minimum energy that colliding particles must have for a reaction to occur.
Reactions with a low activation energy occur readily, as many particles have sufficient energy. Reactions with a high activation energy are slow at room temperature.
Catalysts and Activation Energy
A catalyst provides an alternative reaction pathway with a lower activation energy. Because more particles now have energy equal to or greater than this lower activation energy, a greater proportion of collisions are successful. The rate of reaction therefore increases.
The catalyst itself is unchanged at the end of the reaction.
Reaction Profile Diagrams
Higher Tier
Spec point 3.14C: draw and explain reaction profile diagrams showing DH and activation energy.
Reaction profile diagrams (also called reaction profiles) show how the energy of the reacting system changes during the course of a reaction.
Key features of a reaction profile:
- The x-axis shows the progress of the reaction (from reactants to products)
- The y-axis shows the energy
- The curve rises to a peak (the transition state) before falling to the products
- The difference in height between reactants and products = the enthalpy change (DH)
- The height from the reactants to the peak = the activation energy (Ea)
For an exothermic reaction: products are lower energy than reactants; DH is negative; the arrow pointing from reactants to products points downward.
For an endothermic reaction: products are higher energy than reactants; DH is positive; the arrow points upward.
Effect of a catalyst on a reaction profile: A catalyst lowers the activation energy. The peak in the profile is at a lower energy level, but the energies of reactants and products (and therefore DH) remain unchanged.
Practical: Effect of Surface Area on Rate of Reaction
Aim (spec 3.15): Investigate the effect of changing the surface area of marble chips on the rate of reaction with dilute hydrochloric acid.
Reaction: CaCO3(s) + 2HCl(aq) → CaCl2(aq) + H2O(l) + CO2(g)
Method:
- Add a fixed volume of dilute hydrochloric acid to a conical flask.
- Connect the flask to an inverted measuring cylinder in a water trough (to collect gas).
- Add marble chips and close the bung.
- Measure the volume of carbon dioxide produced at regular time intervals.
- Repeat using marble chips of different sizes (but the same total mass), keeping the acid volume and concentration constant.
Conclusion: Smaller marble chips (larger surface area) produce gas more quickly (steeper rate curve) because more carbonate particles are exposed to the acid, increasing collision frequency.
Also investigate the effect of changing concentration of hydrochloric acid (keeping surface area constant): higher concentration gives a faster rate.
Practical: Effect of Catalysts on Rate of Reaction
Aim (spec 3.16): Investigate the effect of different solids on the catalytic decomposition of hydrogen peroxide solution.
Reaction: 2H2O2(aq) → 2H2O(l) + O2(g)
Method:
- Add hydrogen peroxide solution to a conical flask.
- Connect the flask to an inverted measuring cylinder in a water trough.
- Add the chosen solid catalyst and close the bung.
- Measure the volume of oxygen produced at regular time intervals.
- Repeat with different catalysts (e.g. manganese(IV) oxide, lead(II) oxide, iron(III) oxide, copper(II) oxide).
- Compare results by plotting volume of gas produced over time for each catalyst.
Conclusion: Different catalysts give different rates; manganese(IV) oxide is particularly effective. The catalyst providing the highest rate of gas production in a given time is the most effective catalyst.
Exam Questions
Q: Explain, using collision theory, why increasing the temperature increases the rate of reaction.
At a higher temperature, particles have more kinetic energy, so they move faster. This means collisions occur more frequently. Additionally, a greater proportion of particles have energy equal to or greater than the activation energy. Therefore more collisions result in successful reactions, and the rate of reaction increases.
Q: Explain why a catalyst increases the rate of reaction but is not used up.
A catalyst provides an alternative reaction pathway with a lower activation energy. A greater proportion of reactant particles now have sufficient energy to react, so the rate increases. The catalyst takes part in the reaction but is regenerated at the end, so it remains chemically unchanged.
Q: A student investigates the reaction between marble chips and dilute hydrochloric acid. Describe how surface area can be changed while keeping other variables constant.
Use marble chips of different sizes (e.g. large lumps versus small chips versus powder), but use the same total mass of marble each time. Keep the volume and concentration of acid constant and carry out the reaction at the same temperature.
Q: In an experiment, 50 cm3 of gas was collected in 30 seconds. Calculate the rate of reaction.
Rate = amount of product formed / time = 50 cm3 / 30 s = 1.67 cm3/s.
Exam Tip
When explaining the effect of any factor on rate, always refer to collision theory: state whether the factor increases the frequency of collisions, the energy of collisions, or both, and explain why. Answers that simply state “more collisions” without explaining why, or that do not link to the activation energy, will not earn full marks for explanation questions.