3.2Physical Chemistry
Rates of Reaction — Edexcel Linear - OL (Chemistry) Revision Notes & Topical Questions
10 min read 13 parts Lesson 5 of 9 Unit quiz
Rates of Reaction: Measuring Rate and Effect of Concentration, Surface Area and Temperature — EDEXCEL LINEAR – OL (CHEMISTRY) notes: Plan and carry out experiments to measure rate and explain the effect of concentration, surface area and temperature using collision theory.The rat…
- rate of a reaction
- Worked example 1
- Worked example 2
- Worked example 3
- Worked example 4
- Rate of reaction
Section 1 of 13Defining and measuring rate
The rate of a reaction tells us how fast reactants are used up or products are formed. Rate = change in amount ÷ time. Reactions range from very slow (rusting, weathering of rocks) to very fast (explosions, precipitation of silver chloride). To measure a rate, choose a property that changes steadily during the reaction.
| Technique | Measured | Example | Limitation |
|---|---|---|---|
| Gas collection (syringe or inverted cylinder over water) | Volume of gas at timed intervals | Magnesium + hydrochloric acid | Gas must not be very soluble |
| Mass loss | Decrease in mass as gas escapes | Marble chips + acid | Light gases such as hydrogen give small changes |
| Disappearing cross | Time for cross to vanish | Sodium thiosulfate + acid | Subjective endpoint; only for precipitates |
| Colour change / colorimeter | Absorbance of coloured solution | Iodine reactions | Needs a coloured species |
Section 2 of 13Method: concentration and surface area
- Use different concentrations made by diluting with water (for example 50, 40, 30, 20 and 10 cm3 of thiosulfate topped up to 50 cm3).
For surface area, react a fixed mass of calcium carbonate as large chips, small chips and powder with the same volume and concentration of acid, and measure gas volume every 10 s. Smaller pieces give a steeper curve because more particles are exposed at once. Equal masses are used so that the final volume of gas is the same.
Section 3 of 13Explaining the effects with collision theory
Particles can only react if they collide with at least the activation energy.
- so the rate rises steeply.
Temperature experimentwarm the thiosulfate solution to 20, 30, 40, 50 and 60 °C in a water bath, add acid and time the cross.
Section 4 of 13Fair testing and safety
Variables
- the independent variable is the one you change;
- the dependent variable is what you measure;
- control variables are kept the same.
- so work in a well-ventilated area, use small volumes, and dispose of waste promptly.
- and food is stored in refrigerators to slow the rate of reactions caused by bacteria.
Section 5 of 13Evaluating the experiments
Evaluation questions ask how reliable and valid the results are.
- so the same person should do all runs, in the same lighting, looking from above.
- so the starting temperature of each solution should be measured immediately before mixing.
Plotting data
- place the independent variable on the x-axis, label axes with units, choose a scale that uses more than half the grid, plot points accurately and draw a best-fit straight line or smooth curve.
- State the pattern in words: "as concentration doubles, rate doubles", or "as temperature increases, rate increases at an increasing gradient".
Section 6 of 13Worked examples
A student times the thiosulfate cross experiment at 20 °C (80 s) and 40 °C (21 s). Calculate rate (1/time) for each and state the effect. [3]3 marks
At 20 °C: 1 ÷ 80 = 0.0125 s-1; at 40 °C: 1 ÷ 21 = 0.048 s-1. The rate increases by about four times for a 20 °C rise.
Explain why powdered zinc reacts faster with acid than zinc granules. [2]2 marks
Powder has a larger surface area, so more zinc particles are exposed to acid particles and collisions are more frequent.
State three variables to control in a concentration experiment. [3]3 marks
Temperature, volume of each solution, and the mass/size of any solid (or the same flask and cross).
Describe one advantage of collecting gas in a syringe instead of measuring mass loss. [1]1 mark
It is more accurate for light gases like hydrogen, and volume can be read continuously.
Section 7 of 13Catalysts and activation energy
A catalyst increases the rate of a chemical reaction but is not used up and does not change the products.
- so a greater proportion of collisions are successful at the same temperature.
- so it does not alter the equilibrium yield.
| Reaction | Catalyst |
|---|---|
| Decomposition of hydrogen peroxide: 2H2O2 → 2H2O + O2 | Manganese(IV) oxide |
| Haber process (making ammonia) | Iron |
| Contact process (making sulfuric acid) | Vanadium(V) oxide |
| Hydrogenation of alkenes (margarine) | Nickel |
| Cracking of alkanes | Zeolite / aluminium oxide |
| Catalytic converters | Platinum, palladium, rhodium |
Enzymes are biological catalysts, for example in fermentation of glucose to ethanol by yeast. Catalysts are reused, so small amounts are needed, and they allow lower temperatures, saving energy and cost. Catalysts can be poisoned by impurities, and expensive metals must be recovered.
Section 8 of 13Energy profile with and without a catalyst
On an energy profile the uncatalysed route has a high hump; the catalysed route has a lower hump but starts and ends at the same energy levels.
Section 9 of 13Reading rate graphs and calculating rates
A graph of product volume against time starts steep and flattens as reactants are used up.
For hydrogen peroxide decomposition, use a gas syringe and compare rates with different catalysts (manganese(IV) oxide, copper(II) oxide, potato).
Section 10 of 13Practical: comparing catalysts and a worked comparison
A fair comparison of catalysts for hydrogen peroxide uses 25 cm3 of 10 volume peroxide each time, 0.5 g of each catalyst, and a gas syringe to record the oxygen at 10 s intervals.
- which shows that catalysts do not change the amount of product.
When drawing an energy profile with and without a catalyst, remember that the catalyst gives a different route through a different transition state, not a lower starting point. Industrial catalysts are often in the form of pellets, gauze or honeycomb to maximise surface area; their activity can drop if they are coated with impurities (poisoning), for example sulfur compounds poisoning the iron in the Haber process, so reactants must be purified first.
Section 11 of 13Worked examples
Explain how a catalyst increases the rate of reaction. [2]2 marks
It provides an alternative route with a lower activation energy, so a larger proportion of collisions have enough energy to react.
A tangent to a gas volume curve goes from (0, 10) to (50, 60). Calculate the rate. [2]2 marks
(60 − 10) ÷ 50 = 1.0 cm3/s.
Describe how to show that manganese(IV) oxide is not used up in the decomposition of hydrogen peroxide. [3]3 marks
Weigh the catalyst before, then after the reaction by filtering, washing and drying it; the mass is the same.
Why does the final volume of gas not depend on the catalyst? [1]1 mark
The amount of product is determined by the amount of limiting reactant, not by the rate.
Section 12 of 13Key terms
| Term | Meaning |
|---|---|
| Rate of reaction | How quickly reactants are used or products formed |
| Collision frequency | Number of collisions per second |
| Surface area | Exposed area of a solid |
| Concentration | Amount of solute per unit volume |
| Control variable | A factor kept constant in an experiment |
| Precipitate | Insoluble solid formed in a solution |
| Term | Meaning |
|---|---|
| Catalyst | Substance that speeds up a reaction and is not used up |
| Activation energy | Minimum energy needed for a successful collision |
| Enzyme | Biological catalyst |
| Gradient | Change in y divided by change in x |
| Tangent | Straight line touching a curve at one point |
| Limiting reactant | Reactant that runs out first |
Section 13 of 13Exam technique
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