Textbook Chemistry Form 4 KSSM (DLP), pages 218–251
What this chapter covers
Classify fast and slow reactions and explain the meaning of the rate of reaction (7.1)
Identify the changes that can be observed and measured during a reaction (7.1)
Calculate the average rate and the instantaneous rate of reaction from data and graphs (7.1)
Investigate how the size of reactants, concentration, temperature and a catalyst affect the rate of reaction (7.2)
Explain how the factors that affect the rate of reaction are applied in daily life (7.3)
Use the collision theory and activation energy to explain the rate of reaction (7.4)
Interpret energy profile diagrams of exothermic and endothermic reactions, with and without a catalyst (7.4)
7.1Determining rate of reaction
The rate of reaction is the change in the quantity of a reactant, or of a product, per unit time.
Some reactions are fast (ignition of matches, fireworks, combustion of gases, electric cells). Others are slow (rusting, food decay, photosynthesis, fermentation, erosion of rock).
Rate of reaction = change in the quantity of reactant or product ÷ time taken for the change.
Quantity measured
Common units of rate
Mass of a solid
g s−1 or g minute−1
Volume of a gas
cm3 s−1 or cm3 minute−1
Concentration of a solution
mol dm−3 s−1 or mol dm−3 minute−1
Changes you can observe and measure
Change
Example
How it is measured
Precipitate forms
AgNO3 + NaCl gives AgCl
Time for a "X" mark to disappear; mass of precipitate
Mass of reactants decreases
HNO3 + CaCO3 releases CO2
Electronic balance
Volume of gas increases
HCl + Mg releases H2
Gas syringe or burette over water
Pressure changes
Reactions of gases
Pressure meter
Electrical conductivity changes
Reactions with mobile ions
Ammeter
pH changes
Acids or bases in solution
pH meter
Average rate and instantaneous rate
Average rate The rate over a period of time. Overall average rate = total volume of gas ÷ time taken for the reaction to finish. For a period t1 to t2: (V2 − V1) ÷ (t2 − t1).
Instantaneous rate The rate at one moment. Draw a tangent to the curve at that time. The rate is the gradient of the tangent, ΔV ÷ Δt. A bigger triangle gives a more accurate gradient.
Analogy: a car that covers 400 km in 4 hours has an average speed of 100 km h−1. A speed camera that reads 140 km h−1 gives the speed at that instant.
Worked example
Excess MgCO3 reacts with sulphuric acid. The volume of CO2 is recorded every minute (Table 7.2 in the textbook).
Time (minute)
0
1
2
3
4
5
6
7
8
9
10
Volume (cm3)
0
25
40
51
58
63
68
70
70
70
70
Volume of gas against time. The curve is steepest at the start and flat once the reaction stops.
Overall: the volume stops rising at 7 minutes, so 70 ÷ 7 = 10 cm3 minute−1.
For the overall average rate, divide by the time the reaction ended (7 minutes here), not the end of the recording (10 minutes).
A reaction releases 48 cm3 of gas in 4 minutes and then stops. Calculate the overall average rate.
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48 ÷ 4 = 12 cm3 minute−1.
7.2Factors affecting the rate of reaction
A catalyst is a substance that changes the rate of a reaction without itself being chemically changed at the end. It does not change the quantity of product.
Factor
Change
Effect on rate
Experiment in the textbook
Size of reactant
Smaller pieces (larger total surface area)
Increases
Marble chips with HCl, volume of CO2 against time
Concentration
Higher concentration
Increases
Na2S2O3 with H2SO4, time for the "X" mark to disappear
Temperature
Higher temperature
Increases
Same reaction from room temperature to 55 °C
Catalyst
Catalyst present
Increases
H2O2 with MnO2, glowing splinter relights
Remember the four factors with CSTP: Concentration, Size of solid reactant, Temperature, Presence of catalyst. For gases, higher pressure also raises the rate because it raises the concentration.
For a fixed mass, a powder has a larger total surface area than a lump. Cutting a 2 cm cube into smaller pieces raises the total surface area (24 cm2 to 32 cm2 to 40 cm2 in the textbook figure).
In the sodium thiosulphate experiment, rate = 1 ÷ time. Graph the rate against concentration. At room temperature, a rise of 10 °C roughly doubles the rate.
A catalyst may change physically (a lump becomes powder) but its mass and chemical properties stay the same.
A change in rate does not change the total amount of product. If the reactants are the same, the final volume of gas is the same. Only the time to get there changes.
Worked example
Experiment I uses 1.0 mol dm−3 HCl and takes 120 s. Experiment II uses 2.0 mol dm−3 HCl and takes 60 s. Same antacid tablet, same temperature.
Rate = 1 ÷ time: I = 1 ÷ 120 = 0.0083 s−1; II = 1 ÷ 60 = 0.017 s−1. The higher concentration doubles the rate.
Name the manipulated and the responding variable in the experiment on the effect of temperature on the reaction of sodium thiosulphate with sulphuric acid.
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Manipulated: temperature of the sodium thiosulphate solution. Responding: time for the "X" mark to disappear (rate of reaction).
7.3Factors in daily life
Factor
Application
Why it works
Size
Chewing an antacid tablet; slicing potatoes thinly before frying
Larger total surface area exposed
Concentration
Iron corrodes faster near industry; petrol vapour compressed in a car engine
More acid rain pollutant; more particles per unit volume
Temperature
Washing with hot water and detergent; frying in oil above 100 °C; keeping food in a refrigerator
Higher temperature speeds reactions up; low temperature slows bacterial decay
Catalyst
Platinum catalytic converter; yeast enzymes make ethanol at 37 °C
Pollutants (CO, NOx, hydrocarbons) become CO2, N2, H2O
Why does coal burn faster as small chips than as a large lump?
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Small chips have a larger total surface area exposed to oxygen, so collisions with oxygen are more frequent.
7.4Collision theory
The collision theory says that reactant particles must collide for a reaction to occur, and the rate depends on the frequency of effective collisions. A collision is effective only if the particles have energy ≥ the activation energy, Ea, and collide in the correct orientation.
Factor increased
What happens to the particles
Concentration or pressure
More particles per unit volume, so collisions are more frequent, so effective collisions are more frequent
Total surface area
More particles exposed to collision, so effective collisions are more frequent
Temperature
Particles gain kinetic energy, more of them reach Ea, effective collisions are more frequent
Catalyst
An alternative pathway with a lower activation energy Ea′, so more particles reach it
Energy profile diagrams
Exothermic Reactants have more energy than products. Heat is released to the surroundings. ΔH is negative.
Endothermic Products have more energy than reactants. Heat is absorbed from the surroundings. ΔH is positive.
A catalyst lowers the peak (Ea′ < Ea). ΔH stays the same.
A catalyst lowers the activation energy but does not change the energy of reactants or products, so ΔH is unchanged. Temperature does not lower Ea: it gives more particles energy ≥ Ea.
State two conditions for an effective collision.
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The particles have energy equal to or more than the activation energy, and they collide in the correct orientation.
Exothermic
Endothermic