Rate of Reaction — Cambridge O Level Chemistry 5070 Revision Notes & Topical Questions

Section 1 of 21Collision theory

• Requirements for reaction
  • particles must collide, and the collision must have enough energy to break bonds.
• Activation energy
  • the minimum energy colliding particles need in order to react.
  • Successful collision
    • a collision with energy at or above the activation energy that leads to a reaction.
  • Unsuccessful collision
    • particles bounce off each other because they have less than the activation energy.

Section 2 of 21What controls the number of successful collisions

• Particles per unit volume
  • more particles in the same volume collide more often.
• Frequency of collisions
  • more collisions per second give more successful collisions per second.
• Kinetic energy
  • faster particles collide more often and a greater proportion exceed the activation energy.
• Activation energy
  • a higher value means fewer collisions are successful.
• Rate
  • the rate of reaction depends on the number of successful collisions per unit time.
  • Industry
    • a higher rate means faster production and a more efficient, sustainable process.

Section 3 of 21Factors that change the rate

Factor	Effect on rate	Reason using collisions
Higher concentration or gas pressure	Increases	More particles per unit volume, so more frequent collisions
Higher temperature	Increases	Particles have more kinetic energy, collide more often and a larger proportion exceed the activation energy
Larger surface area (smaller lumps)	Increases	More surface exposed, so more collisions per second
Catalyst	Increases	Alternative pathway with lower activation energy

Section 4 of 21Temperature in detail

• Graph
  • the line is steeper at the start and levels off sooner at a higher temperature.
• Reason
  • more kinetic energy means more frequent collisions and more particles above the activation energy.
  • Size of effect
    • a small rise in temperature causes a large rise in rate.
  • Rule of thumb
    • in aqueous and gaseous systems the rate roughly doubles for every 10 °C rise.

Section 5 of 21Concentration, pressure and surface area in detail

• Concentration
  • the line is steeper and levels sooner at higher concentration.
• Explanation
  • more particles per unit volume, so more frequent successful collisions.
  • Proportionality
    • doubling the number of particles doubles the collisions per second.
• Gas pressure
  • raising pressure squeezes the same particles into a smaller space, like raising concentration.
• Surface area
  • powdered solid reacts faster than lumps because more particles are exposed to the other reactant.
• Doubling
  • doubling the surface area doubles the collisions per second.

Section 6 of 21Catalysts

• Definition
  • a catalyst speeds up a reaction and is unchanged at the end.
• Mass
  • it is the same before and after and does not appear in the equation.
  • Mechanism
    • it provides an alternative pathway with lower activation energy.
  • Result
    • a larger proportion of particles have enough energy, so more collisions are successful.
• Examples
  • iron for the Haber process, used as beads to increase surface area.
• Enzymes
  • biological catalysts that work best at a specific temperature and pH.
  • Amount
    • only small amounts of catalyst are usually needed.

Section 7 of 21How rate is measured

• Principle
  • measure how fast reactants are used up or products form, using a property that changes with concentration.
• Properties
  • colour, mass or gas volume.
  • Timing
    • fast reactions are easier to measure over the whole reaction, while slower ones can be followed as they go, giving more data.
• Common techniques
  • mass loss on a balance, volume of gas produced, and time for a colour change.

Section 8 of 21Experiment: surface area

1. Put dilute hydrochloric acid into a conical flask.
2. Connect the flask by a delivery tube to an upturned measuring cylinder over water (downward displacement).
3. Add calcium carbonate chips and quickly replace the bung.
4. Measure the gas volume collected in a fixed time.
5. Repeat with lumps, crushed pieces and powder.

• Result
  • smaller pieces give a faster rate because more surface is exposed to the acid.

Section 9 of 21Experiment: concentration (sodium thiosulfate)

1. Measure 50 cm³ of sodium thiosulfate solution into a flask.
2. Measure 5 cm³ of dilute hydrochloric acid in a measuring cylinder.
3. Place the flask over a cross drawn on paper.
4. Add the acid and start the stopwatch at once.
5. Look down from above and stop the clock when the cross can no longer be seen.
6. Repeat with different concentrations made by diluting with water.

• Result
  • higher concentration means more particles per unit volume and a shorter time.

Section 10 of 21Experiments: temperature and catalyst

1. Heat dilute hydrochloric acid in a water bath to a set temperature.
2. Pour it into a conical flask, add a magnesium strip and start the stopwatch.
3. Stop the clock when the magnesium disappears.
4. Repeat at different temperatures.

• Temperature result
  • a higher temperature shortens the time because particles have more energy.
• Catalyst experiment
  • add hydrogen peroxide to a flask connected to an upturned measuring cylinder over water, add manganese(IV) oxide,
  • and measure the oxygen volume in a fixed time.
• Control
  • repeat without the catalyst and compare.
  • Result
    • the catalyst gives more gas in the same time because of lower activation energy.

Section 11 of 21Worked examples

1. Explain, using collision theory, why powdered marble reacts faster with acid than marble chips. [3]3 marks

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2. State and explain the effect of raising the temperature by 10 °C on the rate. [3]3 marks

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3. Explain how a catalyst increases the rate of a reaction. [2]2 marks

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4. In the sodium thiosulfate experiment, why is the same volume of mixture used each time? [1]1 mark

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Section 12 of 21Monitoring change in mass

• Method
  • perform the reaction in an open flask on a balance and record the loss in mass as gas escapes.
• Cotton wool
  • a plug in the neck lets gas out but stops liquid spraying out in vigorous reactions.
  • Limit
    • it is not suitable for hydrogen or other gases of very small Mr because the mass loss is too small to measure.
  • Alternative
    • collect the gas in a gas syringe or measuring cylinder and record volume against time.

Section 13 of 21Evaluating methods

Method	Advantage	Disadvantage
Gas collected over water or in a syringe	Gives readings during the reaction	Gas may dissolve or leak, bung must be fitted quickly
Loss of mass on a balance	Simple and gives continuous readings	Not suitable for light gases such as hydrogen
Time for colour change or cross to disappear	Easy equipment	Judging the end point is subjective and gives one reading

• Good planning
  • compare the available methods and choose the one suited to the reaction.
• You may be asked
  • to devise and evaluate your own method.

Section 14 of 21Rate graphs

• Plotting
  • data from a rate study are plotted to show how rate changes with time.
• Shape
  • the line is steepest at the start, becomes less steep as reactants are used up and is finally horizontal when the reaction ends.
  • Meaning of steepness
    • the steeper the curve, the faster the rate.
  • Several runs
    • plotting different conditions on one graph makes comparison easy.
• Why the rate falls
  • reactant concentration decreases as the reaction goes on.
• Finish
  • when one reactant is used up the rate is zero.
  • Mass or volume
    • volume and mass of product are proportional, so a mass graph has the same shape.
  • Reactant graph
    • if reactant is being monitored the graph falls steeply and levels off.

Section 15 of 21Amount of product

• Limiting reactant
  • determines the total product formed.
• More limiting reactant
  • more product.
  • More excess reactant
    • no change in the amount of product.
• Rate and amount
  • a faster rate gives a steeper line but the same final amount if the limiting reactant is unchanged.

Section 16 of 21Worked pattern: changing concentration

1. 0.2 g MnO₂ with 25 cm³ of 0.1 mol/dm³ hydrogen peroxide gives a first curve of oxygen volume against time.
2. Repeat with 0.2 mol/dm³ hydrogen peroxide at the same temperature and catalyst mass.
3. The solution is twice as concentrated, so the initial gradient is steeper.
4. Twice as many peroxide particles are present, so the final volume of oxygen is double.
5. Sketch a steeper curve levelling off at twice the height.

Section 17 of 21Calculating rate at a point

1. Draw a tangent to the curve at the chosen time.
2. Complete a right-angled triangle on the tangent.
3. Read off the change in y and the change in x.
4. Rate = change in y ÷ change in x.

• Example
  • iodine and methanoic acid, volume of CO₂ against time.
• Tangent at 20 s
  • rises 24 cm³ over 40 s.
  • Result
    • 24 ÷ 40 = 0.60 cm³/s.

Section 18 of 21Average rate

• Whole reaction
  • average rate = total product ÷ total time.
    • (or reactant used)
• Between two times
  • use the difference in readings divided by the difference in time.
  • Units
    • cm³/s for gas volume, g/s for mass.
• Example
  • 60 cm³ of gas in 120 s gives 0.50 cm³/s.

Section 19 of 21Worked examples

1. The mass of a flask and contents falls by 0.44 g in 110 s. Calculate the average rate. [2]2 marks

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2. Explain why the curve of gas volume against time becomes less steep. [2]2 marks

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3. Why is loss of mass unsuitable for measuring the rate of the reaction of magnesium with acid? [2]2 marks

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4. A tangent to a curve at 30 s rises from 10 to 40 cm³ over 60 s. Find the rate. [2]2 marks

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Section 20 of 21Key terms

Term	Meaning
Collision theory	Particles react only when they collide with energy at or above the activation energy.
Activation energy	The minimum energy colliding particles need to react.
Catalyst	A substance that increases reaction rate and is unchanged at the end.
Surface area	The area of a solid exposed to the other reactant.
Rate of reaction	How fast reactants are used up or products formed.
Enzyme	A biological catalyst.

Term	Meaning
Tangent	A straight line touching a curve at one point, used to find the gradient.
Gradient	Change in y divided by change in x.
Limiting reactant	The reactant that is used up first.
Average rate	Total change divided by total time.
Downward displacement	Collecting a gas in an upturned measuring cylinder over water.

Section 21 of 21Exam technique

1. Saying a catalyst gives particles more energy — it lowers the activation energy needed.
2. Saying higher temperature only increases collision frequency — the proportion with enough energy also rises.
3. Saying more concentration means bigger particles — it means more particles per unit volume.
4. Forgetting to repeat with the same conditions apart from the one variable being tested.

1. Saying the rate is constant through the reaction — it decreases as reactants are used.
2. Reading a tangent badly — draw it at the point and use a large triangle.
3. Saying a faster rate means more product — only more limiting reactant gives more product.
4. Using the loss-of-mass method for hydrogen — its mass is too small.