Rates of reaction

How fast does a chemical reaction go? Some reactions, like rust forming on an old bicycle, take years. Others, like fireworks exploding, are over in milliseconds. In chemistry, understanding and controlling the speed—or rate—of a reaction is vital for safety, efficiency, and industrial profitability.

What you'll learn:

How to measure the rate of a chemical reaction using different experimental methods.
How to interpret graphs of reactions and calculate rates using tangents (Higher Tier).
How collision theory explains how factors like temperature, concentration, and surface area affect the reaction rate.
What catalysts do to speed up reactions, and how to represent this on reaction profiles.

1. What is the "Rate of Reaction"?

Before we can control a reaction, we must define how to measure its speed. In any chemical reaction, starting materials (reactants) turn into new substances (products).

Definition

Rate of reaction

The rate of reaction is a measure of how quickly reactants are used up, or how quickly products are formed, in a given unit of time.

The mathematical formula to find the rate is:

\text{Rate of reaction} = \frac{\text{Amount of reactant used or product formed}}{\text{Time taken}}

Depending on what we are measuring, the units of rate will vary:

If measuring the mass of a reactant or product in grams, the rate is measured in grams per second ( $g/s\text{g/s}$ or $s−1\text{g s}^{-1}$ ).
If measuring the volume of a gas in cubic centimetres, the rate is measured in cubic centimetres per second ( $cm3/s\text{cm}^3\text{/s}$ or $s−1\text{cm}^3\text{ s}^{-1}$ ).

2. Practical Methods for Measuring Rates

To investigate the rate of a reaction, you must measure how the quantity of a reactant or a product changes over time. Edexcel requires you to know three main practical approaches.

Method A: Measuring the production of a gas (Core Practical 7.1a)

When a reaction produces a gas, you can collect and measure its volume at regular time intervals (e.g., every 10 seconds).

For example, when marble chips (calcium carbonate, $CaCO3(s)\text{CaCO}_3\text{(s)}$ ) react with dilute hydrochloric acid ( $HCl(aq)\text{HCl}\text{(aq)}$ ):

\text{CaCO}_3\text{(s)} + 2\text{HCl}\text{(aq)} \rightarrow \text{CaCl}_2\text{(aq)} + \text{H}_2\text{O}\text{(l)} + \text{CO}_2\text{(g)}

Because carbon dioxide gas ( $CO2\text{CO}_2$ ) is produced, we can collect it using a horizontal gas syringe or an inverted measuring cylinder filled with water.

Setup for measuring gas production

Method B: Change in mass (The balance method)

Alternatively, you can measure the loss in mass of the reaction flask.

Place the reaction flask (containing acid and marble chips) directly on a digital mass balance.
Plug the neck of the flask with a loose ball of cotton wool. This allows the carbon dioxide gas to escape but prevents any acid droplets from spraying out.
As the gas escapes, the total mass of the flask decreases. Record the mass at regular intervals.

Common Mistake

Forgetting the cotton wool

Students often forget to mention the cotton wool in exam questions about the mass-loss method. Without it, acid spit can escape, making the mass decrease look larger than it actually is. This makes your calculated rate inaccurate.

Method C: Observing a colour change (Core Practical 7.1b)

Some reactions produce an insoluble solid called a precipitate, which makes the solution turn cloudy (opaque).

In this core practical, you react clear sodium thiosulfate solution ( $Na2S2O3\text{Na}_2\text{S}_2\text{O}_3$ ) with clear hydrochloric acid ( $HCl\text{HCl}$ ):

\text{Na}_2\text{S}_2\text{O}_3\text{(aq)} + 2\text{HCl}\text{(aq)} \rightarrow 2\text{NaCl}\text{(aq)} + \text{S}\text{(s)} + \text{SO}_2\text{(g)} + \text{H}_2\text{O}\text{(l)}

The reaction produces solid sulfur ( $S(s)\text{S}\text{(s)}$ ), which is a yellow precipitate.

Draw a black cross on a piece of paper and place it under the reaction flask.
Mix the reactants and start a stopwatch.
Look down through the top of the flask. As the yellow sulfur precipitate forms, the liquid becomes increasingly cloudy.
Stop the timer the exact moment you can no longer see the black cross.
The rate is proportional to $1time\frac{1}{\text{time}}$ . A shorter time means a faster reaction.

3. Graphing and Calculating Reaction Rates

If we plot reaction progress on a graph, we put time on the x-axis and the quantity of reactant or product on the y-axis.

How to read a rate graph

At the start: The curve is at its steepest. The reaction is fastest because there is a high concentration of reactant particles, leading to frequent collisions.
As time goes on: The curve becomes less steep (it flatlines/curvatures off). The reaction is slowing down because reactants are being used up, so collisions become less frequent.
At the end: The curve becomes completely horizontal. The reaction has stopped because at least one of the reactants has been completely used up (it is the limiting reactant).

Example

Calculating the mean rate of reaction

A student reacted zinc with excess sulfuric acid. They collected $cm348\text{ cm}^3$ of hydrogen gas over a period of $seconds80\text{ seconds}$ , at which point the reaction completely stopped. Calculate the mean rate of reaction in $cm3/s\text{cm}^3\text{/s}$ .

Identify the equation to use:

\text{Mean rate} = \frac{\text{Total volume of product formed}}{\text{Total time taken}}

Substitute the experimental values into the equation:

\text{Mean rate} = \frac{48\text{ cm}^3}{80\text{ s}}

Calculate the final value and state the units:

\text{Mean rate} = 0.6\text{ cm}^3\text{/s}

Calculating the rate at a specific point in time (Higher Tier only)

The mean rate only gives us an average over a long period. To find the exact rate of reaction at a specific second (e.g., exactly at $seconds30\text{ seconds}$ ), you must calculate the gradient of a tangent drawn on the graph at that point.

Example

Using a tangent to find rate at a specific time

Below is a step-by-step process of how to calculate the rate at $seconds30\text{ seconds}$ from a graph of gas volume against time.

Locate the point on the curve: Find $seconds30\text{ seconds}$ on the x-axis and move vertically up to locate the corresponding point on the curve.
Draw a tangent line: Use a ruler to draw a straight line that touches the curve only at that single point, matching the slope of the curve at that exact position. Extend the line out to the edges of the grid.
Set up gradient coordinates: Choose two widely separated points on your tangent line to construct a right-angled triangle. Let's say your points are $cm3)(10\text{ s}, 15\text{ cm}^3)$ and $cm3)(50\text{ s}, 43\text{ cm}^3)$ .
Calculate change in y ( $Δy\Delta y$ ) and change in x ( $Δx\Delta x$ ):

\Delta y = 43\text{ cm}^3 - 15\text{ cm}^3 = 28\text{ cm}^3

\Delta x = 50\text{ s} - 10\text{ s} = 40\text{ s}

Compute the gradient ( $ΔyΔx\frac{\Delta y}{\Delta x}$ ):

\text{Rate} = \frac{28\text{ cm}^3}{40\text{ s}} = 0.7\text{ cm}^3\text{/s}

4. Collision Theory

To understand how to change the rate of a reaction, we must look at what is happening on a microscopic level. This is described by collision theory.

Key Idea

The two rules of collision theory

For a chemical reaction to occur, reactant particles must:

Collide with each other.
Collide with a minimum amount of energy, known as the activation energy ( $E_a$ ).

If a collision meets both criteria, it is called a successful collision (or effective collision), and products are formed. If particles collide with less than the activation energy, they simply bounce off each other unchanged.

Definition

Activation energy

Activation energy ( $E_a$ ) is the minimum amount of energy that reacting particles must possess when they collide in order to react.

5. Factors Affecting the Rate of Reaction

We can speed up a reaction by increasing the frequency of successful collisions. This can be achieved by either making collisions happen more often (increasing total collision frequency) or by making the collisions more energetic.

Let’s look at the four factors you need to explain:

A. Temperature

The effect: Increasing temperature increases the rate of reaction.
The collision explanation: When you heat particles, they gain kinetic energy and move faster.
1. This means they collide more frequently because they cover ground quicker.
2. Crucially, a much higher proportion of the colliding particles have energy equal to or greater than the activation energy ( $\ge E_a$ ). Consequently, a higher percentage of the collisions are successful.

B. Concentration (and Pressure in gases)

The effect: Increasing concentration (in solutions) or increasing pressure (in gases) increases the rate of reaction.
The collision explanation:
- Higher concentration means there are more reactant particles in the same volume.
- Higher pressure compresses gas particles closer together, meaning more particles in a given space.
- Because the particles are more crowded, they collide more frequently.

C. Surface Area to Volume Ratio (SA:V)

The effect: Breaking a solid reactant into smaller pieces (or a powder) increases the rate of reaction.
The collision explanation:
- For a given volume of solid, smaller pieces have a much larger surface area to volume ratio (SA:V).
- This exposes more reactant particles to the surrounding fluid.
- As a result, there are more places for collisions to occur, leading to a higher frequency of collisions.

Tip

How to phrase exam answers

When explaining these factors, always use the word frequency (e.g., "frequency of collisions increases") rather than just saying "there are more collisions". The examiner wants to know that there are more collisions per second.

6. Catalysts

A catalyst is a chemical helper that can dramatically increase reaction speeds without being used up.

Definition

Catalyst

A catalyst is a substance that speeds up the rate of a chemical reaction without altering the products of the reaction, and is itself chemically and physically unchanged in mass at the end of the reaction.

How do catalysts work?

Catalysts do not provide energy to the particles. Instead, they provide an alternative reaction pathway that has a lower activation energy.

Reaction profile with and without a catalyst

Because the activation energy barrier is lower, a much higher proportion of reactant particles possess enough energy to react when they collide. This increases the frequency of successful collisions.

Enzymes as biological catalysts

Enzymes are proteins that act as biological catalysts inside living organisms.
They are highly specific, meaning each enzyme typically only catalyses one particular reaction.
In industry, yeast (which contains enzymes) is used in the production of alcoholic drinks. Yeast converts sugars into ethanol and carbon dioxide through the process of anaerobic respiration (fermentation).

Exam technique

In the exam

Distinguish between total and successful collisions: When discussing concentration, pressure, and surface area, write that they increase the frequency of collisions. When discussing temperature, explicitly state it increases both the frequency of collisions AND the proportion of collisions that are successful because particles have more energy.
Tangent Drawing: Bring a sharp pencil and a clear ruler to your exam. When drawing a tangent, ensure your line is perfectly balanced on the curve and extends far enough to make reading coordinates easy.
Catalysts are not reactants: Remember that because catalysts are not used up, they do not appear in the overall balanced chemical equation for the reaction.

Self review

Check yourself

A student reacts marble chips with acid. Why does using powdered marble chips instead of large lumps make the reaction faster? Use collision theory in your answer.
Why does a catalyst remain unchanged in mass at the end of a reaction?
Explain why raising the temperature of a reaction mixture has a much larger effect on the rate of reaction than simply increasing the concentration of reactants.

Time / minutes	Cumulative volume of gas / $cm3\text{cm}^3$
0	0
1	38
2	58
3	70
4	76
5	78

Time interval	0 to 1 minute	1 to 2 minutes	2 to 3 minutes	3 to 4 minutes	4 to 5 minutes
Rate of reaction in $min−1\text{cm}^3\text{ min}^{-1}$	38	20	?	6	2

What you'll learn:

How to measure the rate of a chemical reaction using different experimental methods.
How to interpret graphs of reactions and calculate rates using tangents (Higher Tier).
How collision theory explains how factors like temperature, concentration, and surface area affect the reaction rate.
What catalysts do to speed up reactions, and how to represent this on reaction profiles.

1. What is the "Rate of Reaction"?

Before we can control a reaction, we must define how to measure its speed. In any chemical reaction, starting materials (reactants) turn into new substances (products).

Definition

Rate of reaction

The rate of reaction is a measure of how quickly reactants are used up, or how quickly products are formed, in a given unit of time.

The mathematical formula to find the rate is:

\text{Rate of reaction} = \frac{\text{Amount of reactant used or product formed}}{\text{Time taken}}

Depending on what we are measuring, the units of rate will vary:

If measuring the mass of a reactant or product in grams, the rate is measured in grams per second ( $g/s\text{g/s}$ or $s−1\text{g s}^{-1}$ ).
If measuring the volume of a gas in cubic centimetres, the rate is measured in cubic centimetres per second ( $cm3/s\text{cm}^3\text{/s}$ or $s−1\text{cm}^3\text{ s}^{-1}$ ).

2. Practical Methods for Measuring Rates

To investigate the rate of a reaction, you must measure how the quantity of a reactant or a product changes over time. Edexcel requires you to know three main practical approaches.

Method A: Measuring the production of a gas (Core Practical 7.1a)

When a reaction produces a gas, you can collect and measure its volume at regular time intervals (e.g., every 10 seconds).

For example, when marble chips (calcium carbonate, $CaCO3(s)\text{CaCO}_3\text{(s)}$ ) react with dilute hydrochloric acid ( $HCl(aq)\text{HCl}\text{(aq)}$ ):

\text{CaCO}_3\text{(s)} + 2\text{HCl}\text{(aq)} \rightarrow \text{CaCl}_2\text{(aq)} + \text{H}_2\text{O}\text{(l)} + \text{CO}_2\text{(g)}

Because carbon dioxide gas ( $CO2\text{CO}_2$ ) is produced, we can collect it using a horizontal gas syringe or an inverted measuring cylinder filled with water.

Setup for measuring gas production

Method B: Change in mass (The balance method)

Alternatively, you can measure the loss in mass of the reaction flask.

Place the reaction flask (containing acid and marble chips) directly on a digital mass balance.
Plug the neck of the flask with a loose ball of cotton wool. This allows the carbon dioxide gas to escape but prevents any acid droplets from spraying out.
As the gas escapes, the total mass of the flask decreases. Record the mass at regular intervals.

Common Mistake

Forgetting the cotton wool

Method C: Observing a colour change (Core Practical 7.1b)

Some reactions produce an insoluble solid called a precipitate, which makes the solution turn cloudy (opaque).

In this core practical, you react clear sodium thiosulfate solution ( $Na2S2O3\text{Na}_2\text{S}_2\text{O}_3$ ) with clear hydrochloric acid ( $HCl\text{HCl}$ ):

\text{Na}_2\text{S}_2\text{O}_3\text{(aq)} + 2\text{HCl}\text{(aq)} \rightarrow 2\text{NaCl}\text{(aq)} + \text{S}\text{(s)} + \text{SO}_2\text{(g)} + \text{H}_2\text{O}\text{(l)}

The reaction produces solid sulfur ( $S(s)\text{S}\text{(s)}$ ), which is a yellow precipitate.

Draw a black cross on a piece of paper and place it under the reaction flask.
Mix the reactants and start a stopwatch.
Look down through the top of the flask. As the yellow sulfur precipitate forms, the liquid becomes increasingly cloudy.
Stop the timer the exact moment you can no longer see the black cross.
The rate is proportional to $1time\frac{1}{\text{time}}$ . A shorter time means a faster reaction.

3. Graphing and Calculating Reaction Rates

If we plot reaction progress on a graph, we put time on the x-axis and the quantity of reactant or product on the y-axis.

How to read a rate graph

At the start: The curve is at its steepest. The reaction is fastest because there is a high concentration of reactant particles, leading to frequent collisions.
As time goes on: The curve becomes less steep (it flatlines/curvatures off). The reaction is slowing down because reactants are being used up, so collisions become less frequent.
At the end: The curve becomes completely horizontal. The reaction has stopped because at least one of the reactants has been completely used up (it is the limiting reactant).

Example

Calculating the mean rate of reaction

Identify the equation to use:

\text{Mean rate} = \frac{\text{Total volume of product formed}}{\text{Total time taken}}

Substitute the experimental values into the equation:

\text{Mean rate} = \frac{48\text{ cm}^3}{80\text{ s}}

Calculate the final value and state the units:

\text{Mean rate} = 0.6\text{ cm}^3\text{/s}

Calculating the rate at a specific point in time (Higher Tier only)

Example

Using a tangent to find rate at a specific time

Below is a step-by-step process of how to calculate the rate at $seconds30\text{ seconds}$ from a graph of gas volume against time.

Locate the point on the curve: Find $seconds30\text{ seconds}$ on the x-axis and move vertically up to locate the corresponding point on the curve.
Draw a tangent line: Use a ruler to draw a straight line that touches the curve only at that single point, matching the slope of the curve at that exact position. Extend the line out to the edges of the grid.
Set up gradient coordinates: Choose two widely separated points on your tangent line to construct a right-angled triangle. Let's say your points are $cm3)(10\text{ s}, 15\text{ cm}^3)$ and $cm3)(50\text{ s}, 43\text{ cm}^3)$ .
Calculate change in y ( $Δy\Delta y$ ) and change in x ( $Δx\Delta x$ ):

\Delta y = 43\text{ cm}^3 - 15\text{ cm}^3 = 28\text{ cm}^3

\Delta x = 50\text{ s} - 10\text{ s} = 40\text{ s}

Compute the gradient ( $ΔyΔx\frac{\Delta y}{\Delta x}$ ):

\text{Rate} = \frac{28\text{ cm}^3}{40\text{ s}} = 0.7\text{ cm}^3\text{/s}

4. Collision Theory

To understand how to change the rate of a reaction, we must look at what is happening on a microscopic level. This is described by collision theory.

Key Idea

The two rules of collision theory

For a chemical reaction to occur, reactant particles must:

Collide with each other.
Collide with a minimum amount of energy, known as the activation energy ( $E_a$ ).

Definition

Activation energy

Activation energy ( $E_a$ ) is the minimum amount of energy that reacting particles must possess when they collide in order to react.

5. Factors Affecting the Rate of Reaction

Let’s look at the four factors you need to explain:

A. Temperature

The effect: Increasing temperature increases the rate of reaction.
The collision explanation: When you heat particles, they gain kinetic energy and move faster.
1. This means they collide more frequently because they cover ground quicker.
2. Crucially, a much higher proportion of the colliding particles have energy equal to or greater than the activation energy ( $\ge E_a$ ). Consequently, a higher percentage of the collisions are successful.

B. Concentration (and Pressure in gases)

The effect: Increasing concentration (in solutions) or increasing pressure (in gases) increases the rate of reaction.
The collision explanation:
- Higher concentration means there are more reactant particles in the same volume.
- Higher pressure compresses gas particles closer together, meaning more particles in a given space.
- Because the particles are more crowded, they collide more frequently.

C. Surface Area to Volume Ratio (SA:V)

The effect: Breaking a solid reactant into smaller pieces (or a powder) increases the rate of reaction.
The collision explanation:
- For a given volume of solid, smaller pieces have a much larger surface area to volume ratio (SA:V).
- This exposes more reactant particles to the surrounding fluid.
- As a result, there are more places for collisions to occur, leading to a higher frequency of collisions.

Tip

How to phrase exam answers

6. Catalysts

A catalyst is a chemical helper that can dramatically increase reaction speeds without being used up.

Definition

Catalyst

How do catalysts work?

Catalysts do not provide energy to the particles. Instead, they provide an alternative reaction pathway that has a lower activation energy.

Reaction profile with and without a catalyst

Because the activation energy barrier is lower, a much higher proportion of reactant particles possess enough energy to react when they collide. This increases the frequency of successful collisions.

Enzymes as biological catalysts

Enzymes are proteins that act as biological catalysts inside living organisms.
They are highly specific, meaning each enzyme typically only catalyses one particular reaction.
In industry, yeast (which contains enzymes) is used in the production of alcoholic drinks. Yeast converts sugars into ethanol and carbon dioxide through the process of anaerobic respiration (fermentation).

Exam technique

In the exam

Distinguish between total and successful collisions: When discussing concentration, pressure, and surface area, write that they increase the frequency of collisions. When discussing temperature, explicitly state it increases both the frequency of collisions AND the proportion of collisions that are successful because particles have more energy.
Tangent Drawing: Bring a sharp pencil and a clear ruler to your exam. When drawing a tangent, ensure your line is perfectly balanced on the curve and extends far enough to make reading coordinates easy.
Catalysts are not reactants: Remember that because catalysts are not used up, they do not appear in the overall balanced chemical equation for the reaction.

Self review

Check yourself

A student reacts marble chips with acid. Why does using powdered marble chips instead of large lumps make the reaction faster? Use collision theory in your answer.
Why does a catalyst remain unchanged in mass at the end of a reaction?
Explain why raising the temperature of a reaction mixture has a much larger effect on the rate of reaction than simply increasing the concentration of reactants.

What you'll learn:

1. What is the "Rate of Reaction"?

2. Practical Methods for Measuring Rates

Method A: Measuring the production of a gas (Core Practical 7.1a)

Method B: Change in mass (The balance method)

Method C: Observing a colour change (Core Practical 7.1b)

3. Graphing and Calculating Reaction Rates

How to read a rate graph

Calculating the rate at a specific point in time (Higher Tier only)

4. Collision Theory

5. Factors Affecting the Rate of Reaction

A. Temperature

B. Concentration (and Pressure in gases)

C. Surface Area to Volume Ratio (SA:V)

6. Catalysts

How do catalysts work?

Enzymes as biological catalysts

Recap questions

Review Rates of reaction by teaching Genie

Recap your knowledge with an interactive lesson

Put it into practice with exam-style questions

Table 1: Cumulative volume of gas collected

Table 2: Rates of reaction for each 1-minute interval

Remember key concepts with flashcards

Rates of reaction Revision Guide

Rates of reaction

What you'll learn:

1. What is the "Rate of Reaction"?

2. Practical Methods for Measuring Rates

Method A: Measuring the production of a gas (Core Practical 7.1a)

Method B: Change in mass (The balance method)

Method C: Observing a colour change (Core Practical 7.1b)

3. Graphing and Calculating Reaction Rates

How to read a rate graph

Calculating the rate at a specific point in time (Higher Tier only)

4. Collision Theory

5. Factors Affecting the Rate of Reaction

A. Temperature

B. Concentration (and Pressure in gases)

C. Surface Area to Volume Ratio (SA:V)

6. Catalysts

How do catalysts work?

Enzymes as biological catalysts

Recap questions

Review Rates of reaction by teaching Genie

Recap your knowledge with an interactive lesson

Put it into practice with exam-style questions

Table 1: Cumulative volume of gas collected

Table 2: Rates of reaction for each 1-minute interval

Remember key concepts with flashcards

Rates of reaction Revision Guide