Chemistry · 6 Chemical reactions

6.2 Rate of reaction

Before you start

Particles are in constant random motion, and move faster when hotter.
Concentration in mol/dm³, and that a more concentrated solution has more solute particles in the same volume.
How to read a gradient from a graph.

Syllabus objectives covered (8)

Describe the effect on the rate of reaction of: (a) changing the concentration of solutions (b) changing the pressure of gases (c) changing the surface area of solids (d) changing the temperature (e) adding or removing a catalyst, including enzymes
State that a catalyst increases the rate of a reaction and is unchanged at the end of a reaction
Describe practical methods for investigating the rate of a reaction including change in mass of a reactant or a product and the formation of a gas
Interpret data, including graphs, from rate of reaction experiments
Describe collision theory in terms of: (a) number of particles per unit volume (b) frequency of collisions between particles (c) kinetic energy of particles (d) activation energy, Ea
Describe and explain the effect on the rate of reaction of: (a) changing the concentration of solutions (b) changing the pressure of gases (c) changing the surface area of solids (d) changing the temperature (e) adding or removing a catalyst, including enzymes using collision theory
State that a catalyst decreases the activation energy, Ea, of a reaction
Evaluate practical methods for investigating the rate of a reaction including change in mass of a reactant or a product and the formation of a gas

Part 1 of 6

What 'rate' actually means here

Rate of reaction is how much product forms, or how much reactant is used up, in a given time. Fast reaction, steep change; slow reaction, gentle change.
Five things change it, and you need all five: concentration of solutions, pressure of gases, surface area of solids, temperature, and the presence of a catalyst.
Four of them speed the reaction up when you increase them. Increasing concentration, pressure, surface area or temperature all raise the rate; a catalyst raises it too.
Note which factor applies to which state. Concentration applies to solutions, pressure to gases, surface area to solids. Writing 'increasing the pressure of the acid' shows you have not sorted this out.
A reaction gets slower as it proceeds, because reactants are being used up, and stops when a reactant runs out. That is why every rate graph flattens off.

Part 2 of 6

Collision theory — the reason behind all five

Particles can only react when they collide. But not every collision works — most bounce apart unchanged.
A collision only causes a reaction if the particles collide with at least a minimum amount of energy. That minimum is the activation energy, $E_{a}$ .
So the rate depends on two things: how often particles collide (the frequency of collisions) and what fraction of those collisions have energy of at least $E_{a}$ .
The number of particles per unit volume controls the first. Pack more particles into the same space and they meet more often.
The kinetic energy of the particles controls the second. Faster particles are more likely to clear the $E_{a}$ bar when they do meet.
Every explanation in this subtopic is one of these two sentences applied to a specific factor. If you can say which one a factor changes, you can build the answer yourself.

Part 3 of 6

Explaining each factor

Concentration. A more concentrated solution has more particles per unit volume, so collisions are more frequent, so the rate increases.
Pressure of a gas. Increasing the pressure squeezes the same particles into a smaller volume — again more particles per unit volume, more frequent collisions, faster rate. It is the same explanation as concentration, applied to gases.
Surface area. Breaking a solid into smaller pieces exposes more of its particles to the other reactant, so there are more frequent collisions at the surface. Powder reacts faster than lumps.
Note that surface area does not change the number of particles per unit volume in the solution — it changes how many solid particles are available to be hit. Say 'more surface exposed', not 'more concentrated'.
Temperature. This one is special, because it does two things at once. The particles have more kinetic energy, so they move faster and collide more frequently — and, more importantly, a greater proportion of collisions have energy of at least $E_{a}$ .
For full marks on temperature, say both, and put the energy point first. It is the dominant effect: a modest temperature rise changes collision frequency a little and the successful fraction a lot.
Catalyst. A catalyst provides an alternative pathway with a lower activation energy, so a greater proportion of collisions are successful and the rate increases.
A catalyst does not increase the number or the energy of collisions. It lowers the bar, rather than raising the jump — a distinction examiners test directly.

Part 4 of 6

Catalysts and enzymes

A catalyst increases the rate of a reaction and is unchanged in mass and chemical composition at the end. That is the definition to write down.
Because it is unchanged, a small amount catalyses a large amount of reaction, over and over.
In energy terms: a catalyst decreases the activation energy, $E_{a}$ . It does not change how much energy the reaction gives out or takes in overall.
Same start, same finish, lower hill. The overall energy change is untouched.
Enzymes are biological catalysts — proteins that catalyse reactions in living things. They work best in a narrow temperature and pH range, and are destroyed (denatured) outside it.
This gives enzymes an unusual rate graph: rate rises with temperature as usual, peaks at the optimum, then falls sharply as the enzyme denatures. An ordinary catalyst shows no such peak.

Part 5 of 6

Measuring a rate in the lab

You need something measurable that changes as the reaction proceeds, and a way to record it against time.
Method 1 — collecting gas. Fit the flask with a gas syringe, or collect the gas over water in an inverted measuring cylinder. Record the volume every 10 or 15 seconds.
Method 2 — loss of mass. Stand the flask on a balance and record the mass at fixed intervals. As gas escapes, the mass falls.
The methods have different weaknesses, and that is what 'evaluate' is asking about. A gas syringe can stick or leak, and it has a fixed maximum volume — overshoot it and you lose the end of the experiment.
Collecting over water is cheap but unreliable for soluble gases: carbon dioxide dissolves appreciably, so you under-read. For carbon dioxide, prefer the syringe or the balance.
The mass-loss method is poor when the gas is hydrogen, because hydrogen is so light that the mass change is too small to measure reliably on a school balance.
Whichever you choose, control the variables that are not being tested — same volume and concentration of solution, same mass and form of solid, same temperature — and start timing the instant the reactants meet.
Gas syringe for volume, balance for mass loss — the cotton wool lets gas out but keeps spray in.

Part 6 of 6

Reading rate graphs

Plot volume of gas (or mass lost) on the vertical axis against time on the horizontal. The shape is always the same: steep at first, gradually flattening, then horizontal.
The gradient is the rate. Steepest at $t = 0$ , because that is when concentrations are highest, and zero once the line is flat.
A flat line means the reaction has finished, not that it has been stopped. The limiting reactant has run out.
Comparing two curves is the standard exam task. A steeper initial gradient means a faster reaction; the same final height means the same total amount of product.
That combination — steeper but finishing at the same height — is exactly what a catalyst, a higher temperature or a larger surface area produces. The reaction goes faster, but the same amount of reactant makes the same amount of product.
A curve finishing at a lower height means less product, which means less reactant — a smaller mass or a smaller volume, not a rate factor at all. Distinguishing these two situations is where most marks are won or lost.

Worked examples

Example 1. Explain, using collision theory, why increasing the temperature increases the rate of a reaction.

Show the working

Start with what temperature does to the particles: they gain kinetic energy and move faster.
Now take the two consequences separately. Faster particles collide more frequently.
More importantly, a greater proportion of the collisions now have energy greater than or equal to the activation energy, $E_{a}$ .
Those are the successful collisions, so more successful collisions per second means a higher rate.
Write both points. An answer that mentions only 'they collide more often' typically scores one mark out of three.

Answer: The particles gain kinetic energy and move faster, so collisions are more frequent. More significantly, a greater proportion of collisions have energy at least equal to the activation energy, so a greater proportion are successful and the rate increases.

Example 2. Excess marble chips are added to $50 {cm}^{3}$ of $1.0 {mol/dm}^{3}$ hydrochloric acid and the volume of gas is recorded. The experiment is repeated with the same volume of $0.5 {mol/dm}^{3}$ acid. Describe how the two curves differ and explain why.

Show the working

Deal with the shape first. The more concentrated acid gives the steeper initial gradient, because there are more acid particles per unit volume and so more frequent collisions.
Now the finishing point, which is the part people get wrong. The acid is the limiting reactant in both runs — the marble is in excess.
Halving the concentration halves the amount of acid, so it produces half as much carbon dioxide. The second curve levels off at half the height.
Both curves flatten, but the dilute one takes longer to get there and stops lower.

Answer: The $1.0 {mol/dm}^{3}$ curve is steeper at the start, because more acid particles per unit volume give more frequent collisions. It also levels off at twice the height, because it contains twice as much acid and the acid is the limiting reactant.

Example 3. A student investigates the reaction between magnesium and hydrochloric acid by placing the flask on a balance and recording the loss of mass. Evaluate this method.

Show the working

First say what the method measures: the gas produced escapes, so the falling mass tracks the progress of the reaction.
Identify the gas, because the evaluation depends on it. Magnesium and hydrochloric acid produce hydrogen.
Hydrogen has a very low mass per unit volume, so even a substantial volume of gas causes only a tiny mass change — often too small to measure reliably on a two-decimal-place balance.
State the better alternative and why: collect the hydrogen in a gas syringe, which measures volume directly and is not limited by the balance's resolution.
An evaluation needs a judgement, not just a list. Say which method you would choose and why.

Answer: The method works in principle, but is unsuitable here: hydrogen is so light that the mass loss is very small and comparable to the balance's uncertainty. A gas syringe measuring the volume of hydrogen would give far more reliable data.

Where people go wrong

Mistake: Saying a catalyst 'gives the particles more energy' or 'increases the number of collisions'.

Instead: A catalyst provides an alternative pathway with a lower activation energy.

It changes the energy needed, not the energy available. This is the single most-tested distinction in the subtopic.

Mistake: Explaining temperature with 'the particles collide more often' alone.

Instead: Add the important half: a greater proportion of collisions have energy of at least $E_{a}$ .

The energy effect dominates. An answer with only the frequency point cannot reach full marks.

Mistake: Saying a catalyst or higher temperature produces more product.

Instead: It produces the same amount of product, faster. The curve is steeper but levels off at the same height.

Rate factors change how quickly reactants are used, not how much product they can make.

Mistake: Applying pressure to solutions or concentration to solids.

Instead: Concentration is for solutions, pressure for gases, surface area for solids.

The factors are not interchangeable, and using the wrong one shows the underlying idea has not been understood.

Mistake: Saying the reaction 'stops because it gets too slow'.

Instead: It stops because one reactant has been completely used up.

The flat section of a rate graph is about running out of the limiting reactant, not about slowing to a halt.

Check yourself

List the five factors that affect the rate of a reaction, and state which state of matter each applies to.
Answer
Concentration (solutions), pressure (gases), surface area (solids), temperature (all), and a catalyst (all).
Define activation energy.
Answer
The minimum energy that colliding particles must have for a reaction to occur.
Explain in terms of collision theory why powdered zinc reacts faster with acid than a single lump of the same mass.
Answer
Powder has a larger surface area, so more zinc particles are exposed to the acid, giving more frequent collisions and a faster rate.
State two things a catalyst does and one thing it does not do.
Answer
It increases the rate and lowers the activation energy; it is unchanged at the end. It does not change the overall energy change of the reaction, nor the amount of product formed.
Name two practical methods for following the rate of a reaction that produces a gas.
Answer
Measure the volume of gas collected in a gas syringe at fixed times, or measure the loss of mass of the flask on a balance.
Why is collecting carbon dioxide over water a poor way to measure its volume?
Answer
Carbon dioxide is appreciably soluble in water, so some dissolves and the measured volume is lower than the volume actually produced.
On a graph of volume of gas against time, what does the gradient represent, and what does a horizontal section mean?
Answer
The gradient is the rate of reaction. A horizontal section means the reaction has finished because a reactant has been used up.

Exam questions

Try each part first. Open the mark scheme after you have written your answer.

Question 1. A student adds $2.0$ g of calcium carbonate chips to $50 {cm}^{3}$ of dilute hydrochloric acid and measures the volume of carbon dioxide produced every 20 seconds. The reaction is repeated at a higher temperature, using the same quantities.

(a) Describe how the student could measure the volume of gas produced. 2 marks No calculator

Mark scheme

B1 connect the flask by a delivery tube to a gas syringe (or measuring cylinder inverted over water)
B1 record the volume at regular timed intervals, starting a stopwatch as the reactants are mixed

(b) Sketch, on the same axes, the curve you would expect at the higher temperature. 2 marks No calculator

Mark scheme

B1 steeper initial gradient than the original curve
B1 levels off at the same final volume

(c) Explain, in terms of collision theory, why the rate is greater at the higher temperature. 3 marks No calculator

Mark scheme

B1 particles have more kinetic energy / move faster
B1 collisions are more frequent
B1 a greater proportion of collisions have energy greater than or equal to the activation energy

(d) Explain why both curves become horizontal at the same volume. 2 marks No calculator

Mark scheme

B1 the same amount of limiting reactant is used in both experiments
B1 so the same amount of carbon dioxide is produced; temperature changes the rate, not the yield, owtte

Question 2. Hydrogen peroxide decomposes slowly to water and oxygen. Adding manganese(IV) oxide powder makes the reaction rapid. The manganese(IV) oxide is recovered, dried and weighed at the end.

(a) State what would be observed about the mass of the manganese(IV) oxide, and what this shows. 2 marks No calculator

Mark scheme

B1 the mass is unchanged
B1 showing it acts as a catalyst / is not used up in the reaction

(b) Explain, in terms of activation energy, how the manganese(IV) oxide increases the rate. 2 marks No calculator

Mark scheme

B1 it provides an alternative pathway (route) of lower activation energy
B1 so a greater proportion of collisions are successful

(c) Predict the effect of the catalyst on the total volume of oxygen produced, and justify your prediction. 2 marks No calculator

Mark scheme

B1 no effect / the total volume is unchanged
B1 the same amount of hydrogen peroxide decomposes, so the same amount of oxygen is formed; only the rate changes

Question 3. Two students investigate the reaction between magnesium and dilute sulfuric acid. Student A measures the loss of mass on a balance; student B collects the gas in a gas syringe.

(a) Evaluate student A's method for this particular reaction. 3 marks No calculator

Mark scheme

B1 the gas produced is hydrogen
B1 hydrogen has a very low mass per unit volume, so the mass change is very small
B1 the change may be too small to measure accurately on a typical balance, so the method is unsuitable here

(b) Suggest one practical problem with student B's method. 1 mark No calculator

Mark scheme

B1 any one from: gas may escape before the bung is fitted / the syringe may stick or leak / the volume may exceed the capacity of the syringe

Words to know

Rate of reaction: The amount of reactant used up, or product formed, per unit time.
Collision theory: The model that particles must collide with energy at least equal to the activation energy for a reaction to occur.
Activation energy, $E_{a}$: The minimum energy colliding particles must have for a reaction to take place.
Catalyst: A substance that increases the rate of a reaction by providing a pathway of lower activation energy, and is unchanged at the end.
Enzyme: A biological catalyst, a protein, which works within a narrow range of temperature and pH.
Limiting reactant: The reactant that is completely used up, and which therefore determines how much product forms.

All Chemistry lessons Next: 6.3 Reversible reactions and equilibrium