The GCSE rates of reaction required practical uses two classic experiments: marble chips (calcium carbonate) reacting with hydrochloric acid — tracked by loss of mass — and sodium thiosulfate reacting with hydrochloric acid — tracked by the time taken for a cross to disappear. Both test how changing concentration, surface area, or temperature affects reaction rate.
What is the purpose of the required practical?
The rates of reaction required practical develops three core skills:
- Planning: identifying independent, dependent, and controlled variables
- Collecting data: measuring rate quantitatively (mass loss, volume of gas, or cloudiness)
- Analysing results: drawing graphs, finding rates from tangents, and explaining trends using collision theory
GCSE examiners test understanding of both the practical technique and the underlying theory. You must be able to describe what you would do, predict what results you would see, explain the results using collision theory, and evaluate sources of error.
Experiment 1: Marble chips and hydrochloric acid (loss of mass method)
The reaction
CaCO₃(s) + 2HCl(aq) → CaCl₂(aq) + H₂O(l) + CO₂(g)
Carbon dioxide gas is produced and escapes from the flask, causing the total mass to decrease over time. The rate of mass loss indicates the rate of reaction.
Equipment
- Conical flask on a top-pan balance
- Marble chips (calcium carbonate) — available as large lumps or small chips or powder
- Hydrochloric acid (various concentrations)
- Cotton wool plug (to prevent acid spray escaping while allowing CO₂ to leave)
- Stopwatch
Method
- Place the conical flask (with cotton wool in the neck) on the balance and zero (tare) it.
- Add a measured mass of marble chips.
- Add a measured volume of hydrochloric acid of known concentration.
- Record the mass every 30 seconds (or every minute) until the reaction is complete (mass stops changing).
- Calculate mass of CO₂ lost at each time point = initial mass − mass at that time.
- Plot mass of CO₂ produced (y-axis) against time (x-axis).
Typical results table
| Time (s) | Mass of flask (g) | Mass of CO₂ lost (g) |
|---|---|---|
| 0 | 150.00 | 0.00 |
| 30 | 149.76 | 0.24 |
| 60 | 149.57 | 0.43 |
| 90 | 149.45 | 0.55 |
| 120 | 149.39 | 0.61 |
| 150 | 149.37 | 0.63 |
| 180 | 149.37 | 0.63 |
The mass stops changing when a reactant is used up (the reaction is complete).
Variables
- Independent variable: concentration of HCl (or surface area of marble, or temperature)
- Dependent variable: mass of CO₂ produced per unit time
- Controlled variables: mass of marble, volume of acid, temperature (keep the same for all repeats except when temperature is the variable being tested)
Experiment 2: Sodium thiosulfate and hydrochloric acid (disappearing cross method)
The reaction
Na₂S₂O₃(aq) + 2HCl(aq) → 2NaCl(aq) + S(s) + SO₂(g) + H₂O(l)
A yellow precipitate of sulfur (S) forms, making the mixture progressively cloudier (more opaque). The rate of cloudiness is measured by timing how long it takes for a cross drawn on paper under the flask to become invisible when viewed from above.
Equipment
- Conical flask placed on a piece of paper with a bold cross drawn on it
- Sodium thiosulfate solution (various concentrations)
- Hydrochloric acid
- Stopwatch
- Thermometer (if investigating temperature)
Method
- Draw a cross on a piece of white paper.
- Place the conical flask directly over the cross.
- Add a measured volume of sodium thiosulfate solution to the flask.
- Add a measured volume of hydrochloric acid and start the stopwatch immediately.
- Look down through the flask from above. Record the time at which the cross is no longer visible.
- Repeat with different concentrations (or temperatures), keeping other variables constant.
Calculating rate
The rate is proportional to 1/time:
$$\text{rate} \propto \frac{1}{t}$$
A shorter time means a faster rate. When you plot rate (1/t) against concentration, you expect a directly proportional relationship — a straight line through the origin.
Typical results
| Concentration of Na₂S₂O₃ (mol/dm³) | Time for cross to disappear (s) | Rate (1/t, s⁻¹) |
|---|---|---|
| 0.20 | 24 | 0.042 |
| 0.16 | 30 | 0.033 |
| 0.12 | 40 | 0.025 |
| 0.08 | 61 | 0.016 |
| 0.04 | 121 | 0.008 |
A graph of 1/t against concentration should give a straight line through the origin, confirming that rate is directly proportional to concentration.
How do you find the rate from a graph?
Gradient of a straight line (constant rate)
If the graph of gas volume (or mass lost) against time has a straight-line section, the rate in that section = gradient = Δy / Δx.
Tangent method (for a curve)
When the graph is a curve (which is typical because rate slows as reactants are used up), draw a tangent to the curve at the time of interest:
- Use a ruler to draw a straight line touching the curve at exactly one point (the tangent point)
- Extend the tangent line to form a right-angled triangle on the graph
- Rate = rise / run = (change in y) / (change in x) for the triangle
The steeper the tangent at t = 0 (initial rate), the faster the reaction was at the start.
How does collision theory explain the results?
| Change | Effect on collisions | Effect on rate |
|---|---|---|
| Increase concentration (Na₂S₂O₃) | More particles per unit volume → more frequent collisions | Rate increases |
| Increase temperature | Particles have more kinetic energy → more frequent collisions AND more exceed activation energy | Rate increases (steeply) |
| Increase surface area (marble chips → powder) | More exposed surface → more collisions per second | Rate increases |
| Add a catalyst | Provides alternative pathway with lower activation energy → more collisions exceed Eₐ | Rate increases without changing reactant concentration |
Common sources of error and how to reduce them
| Source of error | How to reduce it |
|---|---|
| Judging when the cross disappears (subjective) | Use the same observer throughout; define "invisible" precisely |
| Starting the clock late | Start timing immediately when acid is added; have materials ready |
| Heat produced by the reaction changing temperature | Carry out reaction in a water bath; monitor temperature |
| Variation in chip size (surface area) | Sieve marble chips to get uniform size; record mass not number |
| Losing CO₂ before reading (in experiment 1) | Tare the balance with flask and cotton wool in place before adding reactants |
Frequently asked questions
Why is cotton wool used in the marble chips experiment?
The cotton wool plug prevents acid droplets from being carried out of the flask by the escaping CO₂ gas. Without it, you would lose mass from acid spray as well as from CO₂, making your results inaccurate. Cotton wool allows CO₂ to escape freely (so you measure the mass loss from gas production only) while retaining liquid droplets.
Why do you plot 1/time rather than time itself on the rate graph?
Time is inversely related to rate — a slower reaction takes longer. If you plotted time on the y-axis against concentration on the x-axis, you would get a curve (hyperbola) falling as concentration increases, which is harder to interpret. Plotting rate (= 1/t) against concentration gives a directly proportional straight-line relationship, making the pattern clear and making it easier to draw a best-fit line.
How do you make the sodium thiosulfate experiment a fair test when investigating temperature?
Keep the concentrations of both sodium thiosulfate and hydrochloric acid the same in every trial. Change only the temperature, by warming the sodium thiosulfate solution in a water bath to the target temperature before adding acid. Measure the temperature of the mixture at the start of each trial. Use the same conical flask (same cross beneath) and the same observer. Repeat each trial at least twice and calculate a mean time.
Why does the rate of reaction slow down over time even at constant temperature?
As the reaction proceeds, reactant particles are consumed and their concentration falls. Fewer particles per unit volume means fewer collisions per second, so the rate decreases. On a graph of gas produced (or mass lost) against time, this appears as a curve that becomes less steep over time, eventually flattening out completely when a reactant is exhausted. The steepest gradient is always at the very start, when reactant concentration is highest.
For Socratic GCSE chemistry with Professor Curie — connecting particle-level collision theory to what you measure in a flask — visit aitutors.me.