Form 4 · Chemistry · KCSE

Reaction Rates and Reversible Reactions

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Slide 1 · Learning Outcomes

What you will be able to do

Some reactions are over in a flash (an explosion); others take years (rusting). In this topic you learn what controls how fast a reaction goes, and what happens when a reaction can go both ways.

  • Define rate of reaction and explain it using collision theory and activation energy.
  • Explain how concentration, temperature, surface area, catalysts, pressure and light affect the rate.
  • Describe experiments to measure rate and interpret rate graphs.
  • Calculate average rate from experimental data.
  • Explain reversible reactions and dynamic equilibrium.
  • Apply Le Chatelier's principle to predict the effect of changing conditions on equilibrium.
  • Explain the choice of conditions in the Haber process.
Rate of reaction = change in amount (or concentration) of a reactant or product per unit time. Units: e.g. cm³ s⁻¹, g s⁻¹, mol dm⁻³ s⁻¹.
Slide 2 · Collision theory and activation energy

Particles must collide hard enough

For a reaction to happen, reactant particles must collide, and they must collide with at least a minimum amount of energy called the activation energy (Ea). Collisions with less energy just bounce apart.

Energy profile

A catalyst speeds up a reaction by providing an alternative pathway with a lower activation energy. It is not used up and does not change ΔH or the amount of product. Examples: MnO₂ in decomposing H₂O₂; iron in the Haber process; V₂O₅ in the Contact process.
Slide 3 · Concentration and temperature

Particle simulator

Change the conditions and watch the collisions. Flashes show successful collisions (energy ≥ Ea).

16
25°C
All collisions /s0
Successful /s0
Higher concentration (or higher gas pressure): more particles in the same volume → more frequent collisions → faster rate.
Higher temperature: particles move faster → more collisions, and a bigger fraction have energy ≥ Ea → many more successful collisions. A 10 °C rise roughly doubles many rates.
Slide 4 · Measuring rate

Magnesium + hydrochloric acid

Mg(s) + 2HCl(aq) → MgCl2(aq) + H2(g)

Collect the hydrogen in a gas syringe and record the volume every 30 s. The steeper the curve, the faster the reaction. The curve flattens when one reactant is used up.

Grey line: standard (0.1 g Mg ribbon, excess 1 M HCl, 25 °C). Tap a change to compare.

Tap a change.
Other ways to follow rate: loss in mass as CO₂ escapes (CaCO₃ + HCl on a balance); time for a cross to disappear under cloudy sulphur (sodium thiosulphate + HCl).
Slide 5 · The disappearing cross

Sodium thiosulphate + hydrochloric acid

Na2S2O3(aq) + 2HCl(aq) → 2NaCl(aq) + SO2(g) + H2O(l) + S(s)

The yellow sulphur makes the solution cloudy. We time how long it takes for a cross drawn under the flask to disappear. Rate ∝ 1/time.

0.0 s

Thiosulphate (cm³)Time (s)1/t (s⁻¹)
Run the experiment for each volume to fill the table.
Slide 6 · Reversible reactions and equilibrium

Reactions that go both ways

A reversible reaction can go forward and backward (⇌). In a closed system it reaches dynamic equilibrium: the forward and backward reactions continue at the same rate, so the amounts of reactants and products stay constant.

Activity: chromate ⇌ dichromate

2CrO42−(aq) + 2H+(aq) ⇌ Cr2O72−(aq) + H2O(l)
yellow orange
Yellow
Start: potassium chromate solution (yellow).
Le Chatelier's principle: if a change is made to a system at equilibrium, the equilibrium shifts in the direction that opposes the change.
Other examples: heating hydrated copper(II) sulphate (blue ⇌ white) and heating ammonium chloride (NH₄Cl ⇌ NH₃ + HCl).
Slide 7 · Pressure, temperature and catalysts

Predict the shift

2NO2(g) ⇌ N2O4(g) ΔH = −58 kJ mol⁻¹
brown colourless

Activity: what happens if you…

Brown-ish mixture at equilibrium
Tap a change.
Pressure favours the side with fewer gas moles (it reduces pressure). No effect if gas moles are equal on both sides.
Temperature: raising it favours the endothermic direction; lowering it favours the exothermic direction. A catalyst only makes equilibrium arrive sooner.
Slide 8 · Industrial application

The Haber process

N2(g) + 3H2(g) ⇌ 2NH3(g) ΔH = −92 kJ mol⁻¹

4 moles of gas → 2 moles, and the forward reaction is exothermic. So high pressure and low temperature give the most ammonia. But low temperature is too slow, and very high pressure is expensive and dangerous.

Activity: choose the conditions

Grid: approximate % NH₃ at equilibrium.

Industrial conditions: about 450 °C, 200 atm, finely divided iron catalyst. Unreacted N₂ and H₂ are recycled; NH₃ is removed by cooling it into a liquid.
Slide 9 · Practice

Worked example and quick checks

Worked example: average rate from data

Time (s)0306090120150180
Volume H₂ (cm³)0244051576060

(a) Calculate the average rate in the first 60 s. (b) Why does the rate decrease with time?

Step 1: Volume produced in first 60 s = 40 − 0 = 40 cm³
Step 2: Average rate = volume ÷ time = 40 ÷ 60
Answer (a): 0.67 cm³ s⁻¹
Answer (b): As the reaction goes on, the concentration of acid (and the mass of Mg) falls, so there are fewer collisions per second. The reaction stops at 150 s when the magnesium is used up.

Your turn (use the table above)

1. Average rate between 60 s and 120 s = cm³ s⁻¹ (2 d.p.)
2. Total volume of hydrogen produced = cm³.
3. A catalyst lowers the of a reaction.
4. In 2SO₂ + O₂ ⇌ 2SO₃, increasing pressure shifts the equilibrium to the .
Slide 10 · Quiz

Check your understanding

Part A: Fill in the blanks (auto-marked, 1 mark each)

1. The minimum energy colliding particles need in order to react is the .
2. Powdered calcium carbonate reacts faster than chips because it has a larger surface .
3. A substance that speeds up a reaction without being used up is a .
4. In the decomposition of hydrogen peroxide, the catalyst used is manganese(IV) .
5. In the thiosulphate experiment, rate is taken as 1 divided by .
6. At dynamic equilibrium, the rates of the forward and backward reactions are .
7. For an exothermic forward reaction, increasing temperature shifts the equilibrium to the .
8. Adding acid to yellow chromate solution turns it .
9. The catalyst used in the Haber process is .
10. Increasing pressure favours the side of an equilibrium with moles of gas.

Part B: Short answers (smart-marked against key points)

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