(b) Rates of reaction
- Syllabus
- 2024
- Topic
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- Level
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Learning objectives
3.9Experiments on factors affecting reaction rateDescribe experiments to investigate the effects of changes in surface area of a solid, concentration of a solution, temperature and the use of a catalyst on the rate of a reaction3.10Effects of factors on reaction rateDescribe the effects of changes in surface area of a solid, concentration of a solution, pressure of a gas, temperature and the use of a catalyst on the rate of a reaction3.11Collision theory and reaction rateExplain the effects of changes in surface area of a solid, concentration of a solution, pressure of a gas and temperature on the rate of a reaction in terms of particle collision theory3.12Catalyst definitionKnow that a catalyst is a substance that increases the rate of a reaction, but is chemically unchanged at the end of the reaction3.13Catalysts and activation energyKnow that a catalyst works by providing an alternative pathway with lower activation energy314C Reaction profile diagramsDraw and explain reaction profile diagrams showing ΔH and activation energy3.15Surface area and concentration rate practicalPractical: investigate the effect of changing the surface area of marble chips and of changing the concentration of hydrochloric acid on the rate of reaction between marble chips and dilute hydrochloric acid3.16Catalytic decomposition practicalPractical: investigate the effect of different solids on the catalytic decomposition of hydrogen peroxide solution
Investigate four factors that change reaction rate
To investigate a rate factor, change only that factor, measure how quickly a fixed amount of product forms or reactant disappears, and keep every other condition constant.
| Factor changed | Suitable comparison | Rate measurement | Essential controls |
|---|---|---|---|
| surface area | equal masses of marble chips with different chip sizes | carbon dioxide volume or mass loss against time | acid volume and concentration, temperature, marble mass |
| concentration | different hydrochloric acid concentrations with equal marble samples | gas volume, mass loss or time to a fixed endpoint | total acid volume, marble size and mass, temperature |
| temperature | sodium thiosulfate and acid at several measured temperatures | time until a cross is obscured; rate may be compared using 1/t | solution volumes and concentrations, flask, cross and observer position |
| catalyst | equal hydrogen peroxide samples with equal amounts of different solids | oxygen volume in a fixed time or time to a fixed volume | peroxide volume and concentration, catalyst mass and surface area, temperature |
Repeat each condition and calculate a mean after checking anomalies. Begin timing at mixing, collect readings at regular intervals and use the initial gradient when comparing continuous rate graphs.
Do not change two variables together. A shorter completion time means a faster rate, while a steeper product–time or mass-loss–time gradient means a faster rate; the final amount may remain unchanged when only rate changes.
Predict how five factors affect reaction rate
Reaction rate increases when particles can collide successfully more often. The required factors alter collision frequency or the energy barrier without necessarily changing the final quantity of product.
| Change | Effect on rate |
|---|---|
| increase the surface area of a solid | increases |
| increase solution concentration | increases |
| increase gas pressure | increases |
| increase temperature | increases |
| add a suitable catalyst | increases |
On a product–time graph, a faster reaction has a steeper initial gradient and reaches its plateau sooner. If the reacting amounts are unchanged and only rate changes, both experiments reach the same plateau.
Rate describes change per unit time, not final yield. Greater pressure is a factor for gases; it is not used as the explanation for particles in a liquid solution.
Explain rate changes with collision theory
A reaction occurs only when reactant particles collide with enough energy to react. Rate therefore depends on the number of successful collisions per unit time.
| Factor increased | Particle-level change | Why successful collisions become more frequent |
|---|---|---|
| solid surface area | more reactant particles are exposed at the surface | more collisions can occur at the solid surface each second |
| solution concentration | more reactant particles occupy the same volume | collision frequency increases |
| gas pressure | gas particles are closer together; there are more particles per unit volume | collision frequency increases |
| temperature | mean kinetic energy increases; particles move faster | collisions are more frequent and a larger fraction have energy at least equal to the activation energy |
As acid reacts with excess marble, acid concentration falls. Fewer acid particles remain per unit volume, so successful collisions become less frequent and the graph becomes less steep before levelling off when the acid is used up.
Higher concentration or pressure does not make individual particles move faster or give them more energy. That energy explanation belongs specifically to increasing temperature.
Recognise what a catalyst does—and does not do
A catalyst increases the rate of a reaction but is chemically unchanged at the end. It participates in reaction steps, yet is regenerated rather than consumed overall.
| Test | Evidence expected if the solid is a catalyst |
|---|---|
| compare rate with and without the solid | the reaction is faster when the solid is present |
| recover the solid after completion | filter, wash if needed, dry and reweigh it |
| compare before and after | the same substance and approximately the same mass remain |
For manganese(IV) oxide in hydrogen peroxide decomposition, oxygen is produced faster, but the manganese(IV) oxide can be recovered after the peroxide has reacted.
A catalyst is not a reactant and is not included as a consumed substance in the overall equation. 'Chemically unchanged' does not mean it performs no role during the reaction.
Link catalysts to a lower activation-energy pathway
A catalyst provides an alternative reaction pathway with a lower activation energy, Ea.
At the same temperature, particle energies are not raised by the catalyst. Lowering the energy barrier means a larger fraction of collisions already have enough energy to react, so successful collisions occur more often and rate increases.
| Changed by a catalyst | Not changed by a catalyst |
|---|---|
| reaction pathway and activation energy | reactant and product energy levels |
| reaction rate and time to reach completion or equilibrium | ΔH and the final equilibrium position |
A catalyst does not supply energy to particles, increase their speed or make ΔH more negative. It changes the route, not the starting and finishing energy levels.
Construct and explain a reaction profile
A reaction profile plots energy vertically against progress of reaction horizontally. It shows reactant and product energy levels, the activation-energy barrier and the enthalpy change.
| Feature | How to draw or label it |
|---|---|
| reactants and products | labelled horizontal levels at the start and end |
| pathway | a curve rising to one peak and falling to the product level |
| activation energy, Ea | upward arrow from reactant level to the top of the peak |
| enthalpy change, ΔH | arrow from reactant level to product level; downward/negative for exothermic, upward/positive for endothermic |
| catalyst pathway | second curve with a lower peak but the same start and end levels |
The peak represents the minimum energy barrier along that pathway. A catalyst lowers this peak; because the endpoints remain fixed, it does not alter ΔH.
Do not measure Ea from zero or from the product level for the forward reaction. Do not move product energy when adding a catalyst, and keep the activation-energy arrow pointing upward to the peak.
Compare marble reaction rates fairly
Investigate marble and dilute hydrochloric acid by following carbon dioxide production while changing either marble surface area or acid concentration.
\ce{CaCO3(s) + 2HCl(aq) -> CaCl2(aq) + H2O(l) + CO2(g)}
| Stage | Method |
|---|---|
| 1 | place a measured acid volume and concentration in a conical flask |
| 2 | add a measured mass of marble, immediately fit a gas syringe or place the flask on a balance with a cotton-wool plug, and start timing |
| 3 | record gas volume or mass at regular times until it stops changing |
| 4 | repeat with equal-mass smaller chips while keeping acid constant; separately repeat with different acid concentrations while keeping marble mass and size constant |
| 5 | repeat trials and compare initial gradients or time to a fixed gas volume/mass loss |
Control temperature, total acid volume, apparatus and mixing. A cotton-wool plug prevents acid spray leaving while allowing carbon dioxide to escape, so mass loss is attributable to gas.
Never change chip size and acid concentration in the same comparison. Smaller chips may make a steeper curve without changing the final gas amount when reactant quantities and limiting reactant are unchanged.
Compare catalysts for hydrogen peroxide decomposition
Compare different solid catalysts by measuring how quickly equal hydrogen peroxide samples produce oxygen under otherwise identical conditions.
\ce{2H2O2(aq) -> 2H2O(l) + O2(g)}
| Stage | Method and control |
|---|---|
| 1 | place the same volume and concentration of hydrogen peroxide in a conical flask connected to a gas syringe |
| 2 | add the same mass and comparable particle size of the first solid, seal immediately and start the timer |
| 3 | record oxygen volume at regular intervals or the time to a fixed volume |
| 4 | clean the apparatus and repeat at the same temperature with each solid; repeat each catalyst trial |
| 5 | compare initial gradients or mean times; the steepest gradient or shortest time identifies the most effective catalyst |
Wear eye protection, use dilute hydrogen peroxide and small quantities, and keep the apparatus unblocked so oxygen cannot build dangerous pressure.
Equal catalyst mass alone is insufficient if particle sizes differ greatly, because surface area would also change. Catalyst effectiveness is judged from rate, not from final oxygen volume when the peroxide amount is fixed.