5.2.1 Measuring and interpreting rate of reaction

Mean rate is change in amount over time

Definition

Rate of reaction

How quickly a reactant is used up or a product is formed, measured as the change in amount divided by the time taken.

The mean rate over a time interval is: $taken\text{mean rate} = \frac{\text{change in amount of reactant or product}}{\text{time taken}}$
The units of rate follow from what is measured, such as $cm3/s\text{cm}^3/\text{s}$ for a gas volume or $g/s\text{g/s}$ for a mass.
To follow a reaction, measure a quantity that changes as the reaction goes on, at regular time intervals.
When the time for a fixed change is measured instead, a shorter time means a greater rate.
1. For a fixed change, a relative rate can be worked out as $1time\dfrac{1}{\text{time}}$ , in $s−1\text{s}^{-1}$ .

Example

A reaction gives off $cm348\,\text{cm}^3$ of gas in the first $s60\,\text{s}$ ; calculate the mean rate over this time.
Mean rate $\dfrac{48\,\text{cm}^3}{60\,\text{s}}$ .
Mean rate $0.80\,\text{cm}^3/\text{s}$ .

Following a reaction by measuring gas volume

A reaction that gives off a gas, such as marble chips with hydrochloric acid, can be followed with a gas syringe.
Instead, the gas can be collected over water in an upturned measuring cylinder.
Start the stop clock as soon as the reactants are mixed and the bung is replaced.
Record the volume of gas at regular intervals, such as every 10 seconds, until no more gas is given off.
The gas collection method needs a closed apparatus, so the bung must fit tightly and the tube must not leak.

Practical

Marble chips and hydrochloric acid

Aim: follow how the rate changes during the reaction of marble chips with hydrochloric acid.
Apparatus: a conical flask with a bung and delivery tube, a $cm3100\,\text{cm}^3$ gas syringe, a clamp and stand, a $cm350\,\text{cm}^3$ measuring cylinder, a digital balance, a stop clock, marble chips of different sizes, dilute hydrochloric acid of different concentrations and cotton wool.

Method

Measure $cm350\,\text{cm}^3$ of dilute hydrochloric acid into the conical flask.
Weigh out a set mass of marble chips of one size.
Connect the bung and delivery tube to the gas syringe, checking that the plunger is pushed fully in.
Add the chips to the acid, replace the bung quickly and start the stop clock.
Record the volume of gas every 10 seconds until it stops changing.
Repeat with a different acid concentration or chip size, keeping the volume of acid, the mass of chips and the temperature the same.
Plot volume of gas against time for each run on the same axes, and compare how steep the lines are at the start.

Accuracy and safety

Replace the bung straight away, or some carbon dioxide escapes before it is measured.
Another method is to stand the flask on a balance with a cotton wool plug in the neck and record the loss in mass as carbon dioxide escapes.
The cotton wool lets the gas out but stops acid spray escaping, which would also lower the mass.
Wear eye protection, because hydrochloric acid is an irritant.

Following reactions by mass loss or visible change

If a gas escapes, put the flask on a balance and record the mass at regular intervals.
1. A cotton wool plug in the neck lets the gas out but stops acid spray escaping.
This method suits heavy gases such as carbon dioxide, but hydrogen is so light that the mass change is too small to measure well.
A reaction that forms a precipitate can be timed until a cross drawn under the flask disappears from view.
1. Sodium thiosulfate reacts with hydrochloric acid to form a yellow sulfur precipitate: $Na2S2O3(aq)+2HCl(aq)→2NaCl(aq)+S(s)+SO2(g)+H2O(l)\text{Na}_2\text{S}_2\text{O}_3\text{(aq)} + 2\text{HCl(aq)} \rightarrow 2\text{NaCl(aq)} + \text{S(s)} + \text{SO}_2\text{(g)} + \text{H}_2\text{O(l)}$
A reaction that changes colour can be timed until the colour change reaches the same stage each time.

Note

Choose the method from what the reaction produces: a gas suggests gas volume or mass loss, while a precipitate or colour change suggests timing a visible change.

The gradient of the graph shows the rate

Definition

Gradient

The steepness of a line on a graph, found by dividing the change in the vertical value by the change in the horizontal value.

On a graph of gas volume against time, a steeper curve means gas is given off faster, so the rate is greater.
The curve is steepest at the start, because the reactants are at their most concentrated.
The curve gets less steep as the reactants are used up, so the rate decreases.
A horizontal line means the reaction has stopped, because one of the reactants has been used up.
A graph of mass against time falls as gas escapes, and it too levels off when the reaction stops.

Common Mistake

Do not judge which reaction was faster from the height of a curve, because height shows how much product has formed, while the gradient shows the rate.

Comparing curves and drawing tangents shows how rates differ

When two curves are drawn on the same axes, the one that is steeper at the start shows the faster reaction.
If both curves level off at the same height, the same amount of product formed, so the amount of the used-up reactant was the same.
1. A curve that levels off at half the height shows that half as much of that reactant was used.
The rate at an instant is the gradient of a tangent to the curve at that time.
1. Draw a straight line that touches the curve at that point, then divide its rise by its run using two points far apart on the line.

Example

A tangent drawn at $s20\,\text{s}$ on a gas volume curve passes through $cm3)(0\,\text{s}, 10\,\text{cm}^3)$ and $cm3)(40\,\text{s}, 50\,\text{cm}^3)$ ; find the rate at $s20\,\text{s}$ .
Gradient $\dfrac{50 - 10}{40 - 0} = \dfrac{40\,\text{cm}^3}{40\,\text{s}}$ .
Rate at $cm3/s20\,\text{s} = 1.0\,\text{cm}^3/\text{s}$ .

Self review

How is the mean rate of a reaction calculated?
Name two pieces of apparatus that can collect and measure the volume of a gas.
Why is the mass-loss method unsuitable for a reaction that gives off hydrogen?
What does a horizontal section of a gas volume curve show?
How do you find the rate of a reaction at a particular time from a curve?

5.2.2 Factors affecting rate and collision theory

Colliding particles need at least the activation energy

Definition

Activation energy

The minimum energy that colliding particles must have for a reaction to happen.

Definition

Successful collision

A collision between reactant particles that has at least the activation energy, so the particles react.

Reactant particles must collide with each other before they can react.
Most collisions do not lead to a reaction, because the particles have too little energy.
Only collisions with energy equal to or greater than the activation energy are successful.
The rate of reaction depends on the number of successful collisions per second.
So anything that makes collisions more frequent or more energetic increases the rate.

Key Idea

Every rate explanation comes back to how often particles collide and what fraction of those collisions have enough energy to react.

Higher temperature gives more frequent, more energetic collisions

Definition

Kinetic energy

The energy an object or particle has because it is moving.

Raising the temperature increases the rate of a reaction, and for many reactions a $∘C10\ ^{\circ}\text{C}$ rise roughly doubles it.
The particles gain kinetic energy and move faster.
1. Faster particles collide more frequently, so there are more collisions per second.
More of the collisions have energy equal to or greater than the activation energy.
1. This effect is the larger of the two, so a much greater fraction of collisions is successful.
Lowering the temperature has the opposite effect, which is why food keeps longer in a fridge.

Practical

The effect of temperature on rate

Aim: find how temperature affects the rate of the reaction between magnesium and hydrochloric acid.
Apparatus: a conical flask with a bung and delivery tube, a $cm350\,\text{cm}^3$ measuring cylinder, a $cm3100\,\text{cm}^3$ measuring cylinder to collect the gas, a trough of water, a water bath, a thermometer, a stop clock, a ruler, scissors, magnesium ribbon and dilute hydrochloric acid.

Method

Cut magnesium ribbon into $cm4\,\text{cm}$ strips of equal length.
Measure $cm350\,\text{cm}^3$ of dilute hydrochloric acid into the conical flask and warm it in the water bath until it reaches the chosen temperature.
Fill the $cm3100\,\text{cm}^3$ measuring cylinder with water and turn it upside down in the trough, with the delivery tube from the flask leading into it.
Add a strip of magnesium to the acid, replace the bung at once and start the stop clock.
Stop the clock when $cm330\,\text{cm}^3$ of hydrogen has collected.
Repeat at four or five other temperatures, using the same volume and concentration of acid and the same length of ribbon.
Work out the mean rate at each temperature by dividing the volume of gas by the time, then plot rate against temperature.

Accuracy and safety

Measure the temperature of the acid immediately before adding the magnesium, because it starts to cool once out of the water bath.
Repeat each temperature and find a mean time, so that any anomalous result can be spotted.
Keep flames away, because hydrogen is flammable.
Wear eye protection when handling the acid.

Higher concentration or pressure makes collisions more frequent

Increasing the concentration of a solution puts more reactant particles in the same volume.
1. The particles are closer together, so they collide more frequently and the rate increases.
Increasing the pressure of a gas squeezes the same number of particles into a smaller volume.
1. This also gives more frequent collisions, so pressure affects the rate only when a reactant is a gas.
Neither change gives the particles more energy, so the fraction of collisions that succeed stays the same.

Example

Magnesium ribbon fizzes much more vigorously in concentrated hydrochloric acid than in dilute acid, because acid particles hit the surface of the metal more frequently.

Smaller pieces react faster through larger surface area

Definition

Surface area to volume ratio

The surface area of a particle divided by its volume.

Only the particles on the surface of a solid can collide with particles in the solution or gas around it.
Cutting a solid into smaller pieces keeps its volume the same but increases its surface area, so its surface area to volume ratio increases.
1. A greater fraction of the solid's particles is then exposed to collisions.
Collisions become more frequent, so the rate increases, although the energy of each collision is unchanged.
Powdered calcium carbonate reacts with acid much faster than the same mass of marble chips.
1. Very fine powders such as flour dust can burn so fast in air that they explode.

Example

Compare the surface area to volume ratio of a $cm2\,\text{cm}$ cube with that of the same cube cut into eight $cm1\,\text{cm}$ cubes.
For the large cube, surface area $\times 2^2 = 24\,\text{cm}^2$ and volume $2^3 = 8\,\text{cm}^3$ , so the ratio is $24 : 8 = 3 : 1$ .
For the eight small cubes, surface area $\times 6 \times 1^2 = 48\,\text{cm}^2$ and volume $8\,\text{cm}^3$ , so the ratio is $48 : 8 = 6 : 1$ .
Cutting the cube doubles the surface area to volume ratio, so more particles are exposed to collisions.

Explaining rate changes through collision frequency and energy

For temperature, link it to faster particles, more frequent collisions and more collisions with at least the activation energy.
For concentration or pressure, link it to more particles in the same volume and more frequent collisions.
For size of pieces, link it to the surface area to volume ratio, more exposed particles and more frequent collisions.
For a decrease in any factor, reverse the chain to give fewer successful collisions per second.
Say that the reaction has a greater rate or takes less time, rather than calling it a faster reaction time.

Exam technique

When a question asks you to explain a change in rate, name the factor, say what happens to the particles and finish with the effect on successful collisions per second.

Self review

What is a successful collision?
Give the two reasons why a higher temperature increases the rate of reaction.
Why does increasing the concentration of an acid increase the rate of its reaction with magnesium?
Why does pressure affect only reactions involving gases?
Why does powdered calcium carbonate react faster than marble chips of the same mass?

5.2.3 Catalysts, activation energy and enzymes

A catalyst increases rate and ends chemically unchanged

Definition

Catalyst

A substance that increases the rate of a reaction by providing an alternative pathway with a lower activation energy without being used up.

A catalyst takes part in the reaction but is regenerated, so the same mass is left at the end.
Because it is not used up, only a small amount of catalyst is needed and it can be used again.
A catalyst does not change the amount of product formed, only how quickly it forms.
1. A catalysed reaction reaches the same final volume of gas as an uncatalysed one, but sooner.
Catalysts are usually specific, so a catalyst that speeds up one reaction may have no effect on another.

Common Mistake

Do not write that a catalyst does not take part in the reaction, because it does take part and is then re-formed.

Identifying a catalyst and writing it over the arrow

In an equation a catalyst is written over the arrow, not as a reactant or product.
In $2H2O2(aq)→MnO22H2O(l)+O2(g)2\text{H}_2\text{O}_2\text{(aq)} \xrightarrow{\text{MnO}_2} 2\text{H}_2\text{O(l)} + \text{O}_2\text{(g)}$ , manganese(IV) oxide is the catalyst.
In an experiment, a substance is a catalyst if the rate increases and the same mass of it can be recovered afterwards.
1. Filter off the solid after the reaction, then wash, dry and reweigh it.
Transition metals and their compounds are common industrial catalysts, such as iron in making ammonia and platinum in car catalytic converters.

Practical

Testing catalysts for hydrogen peroxide

Aim: find out which of several solids catalyse the decomposition of hydrogen peroxide.
Apparatus: a conical flask with a bung and delivery tube, a $cm3100\,\text{cm}^3$ gas syringe, a clamp and stand, a $cm325\,\text{cm}^3$ measuring cylinder, a digital balance, a stop clock, a filter funnel and filter paper, hydrogen peroxide solution and solids to test, such as manganese(IV) oxide, copper(II) oxide and zinc oxide.

Method

Measure $cm325\,\text{cm}^3$ of hydrogen peroxide solution into the conical flask.
Weigh $g0.5\,\text{g}$ of the first solid, such as manganese(IV) oxide.
Add the solid to the flask, replace the bung at once and start the stop clock.
Record the volume of oxygen in the gas syringe every 10 seconds.
Repeat with no solid and with each other solid, keeping the volume, concentration and temperature of the hydrogen peroxide the same.
Filter off each solid at the end, then wash, dry and reweigh it.

Accuracy and safety

The run with no solid is a control, which shows how fast hydrogen peroxide decomposes on its own.
The solid that gives the steepest line at the start of the graph is the most effective catalyst.
A catalyst has the same mass at the end as at the start, because it is not used up.
Wear eye protection, because hydrogen peroxide is an irritant.

A catalyst gives a pathway with lower activation energy

Definition

Activation energy

The minimum energy that colliding particles must have for a reaction to happen.

On a reaction profile, the activation energy appears as the height of the peak above the energy of the reactants.
With a catalyst, the reaction follows a different route whose peak is lower.
At the same temperature, more of the collisions now have energy equal to or greater than the lower activation energy.
1. So there are more successful collisions per second, and the rate increases.
The catalyst does not give the particles more energy, and it does not change the overall energy change of the reaction.

Note

On a reaction profile with and without a catalyst, the reactant and product levels stay the same, and only the height of the peak between them is lower for the catalysed pathway.

Enzymes are the catalysts in living things

Definition

Enzyme

A protein that acts as a catalyst in living organisms.

Enzymes catalyse the reactions in living cells, which would otherwise be far too slow at body temperature.
Like other catalysts, enzymes are not used up and work by lowering the activation energy.
Each enzyme is specific, so it catalyses only one reaction or a small group of similar reactions.
Catalase, found in liver and potato, catalyses the breakdown of hydrogen peroxide into water and oxygen.
1. Amylase in saliva catalyses the breakdown of starch into sugars.
Enzymes in yeast catalyse fermentation, which turns sugar into ethanol and carbon dioxide.

Example

Dropping a small piece of raw liver into hydrogen peroxide solution makes it froth with oxygen almost at once, because the catalase in the liver catalyses the decomposition.

Self review

What happens to the mass of a catalyst during a reaction?
Where is a catalyst written in a chemical equation?
How would you show that manganese(IV) oxide is a catalyst for the decomposition of hydrogen peroxide?
How does a catalyst increase the rate of a reaction?
What are enzymes?

5.2.1 Measuring and interpreting rate of reaction

Mean rate is change in amount over time

Definition

Rate of reaction

How quickly a reactant is used up or a product is formed, measured as the change in amount divided by the time taken.

The mean rate over a time interval is: $taken\text{mean rate} = \frac{\text{change in amount of reactant or product}}{\text{time taken}}$
The units of rate follow from what is measured, such as $cm3/s\text{cm}^3/\text{s}$ for a gas volume or $g/s\text{g/s}$ for a mass.
To follow a reaction, measure a quantity that changes as the reaction goes on, at regular time intervals.
When the time for a fixed change is measured instead, a shorter time means a greater rate.
1. For a fixed change, a relative rate can be worked out as $1time\dfrac{1}{\text{time}}$ , in $s−1\text{s}^{-1}$ .

Example

A reaction gives off $cm348\,\text{cm}^3$ of gas in the first $s60\,\text{s}$ ; calculate the mean rate over this time.
Mean rate $\dfrac{48\,\text{cm}^3}{60\,\text{s}}$ .
Mean rate $0.80\,\text{cm}^3/\text{s}$ .

Following a reaction by measuring gas volume

A reaction that gives off a gas, such as marble chips with hydrochloric acid, can be followed with a gas syringe.
Instead, the gas can be collected over water in an upturned measuring cylinder.
Start the stop clock as soon as the reactants are mixed and the bung is replaced.
Record the volume of gas at regular intervals, such as every 10 seconds, until no more gas is given off.
The gas collection method needs a closed apparatus, so the bung must fit tightly and the tube must not leak.

Practical

Marble chips and hydrochloric acid

Aim: follow how the rate changes during the reaction of marble chips with hydrochloric acid.
Apparatus: a conical flask with a bung and delivery tube, a $cm3100\,\text{cm}^3$ gas syringe, a clamp and stand, a $cm350\,\text{cm}^3$ measuring cylinder, a digital balance, a stop clock, marble chips of different sizes, dilute hydrochloric acid of different concentrations and cotton wool.

Method

Measure $cm350\,\text{cm}^3$ of dilute hydrochloric acid into the conical flask.
Weigh out a set mass of marble chips of one size.
Connect the bung and delivery tube to the gas syringe, checking that the plunger is pushed fully in.
Add the chips to the acid, replace the bung quickly and start the stop clock.
Record the volume of gas every 10 seconds until it stops changing.
Repeat with a different acid concentration or chip size, keeping the volume of acid, the mass of chips and the temperature the same.
Plot volume of gas against time for each run on the same axes, and compare how steep the lines are at the start.

Accuracy and safety

Replace the bung straight away, or some carbon dioxide escapes before it is measured.
Another method is to stand the flask on a balance with a cotton wool plug in the neck and record the loss in mass as carbon dioxide escapes.
The cotton wool lets the gas out but stops acid spray escaping, which would also lower the mass.
Wear eye protection, because hydrochloric acid is an irritant.

Following reactions by mass loss or visible change

If a gas escapes, put the flask on a balance and record the mass at regular intervals.
1. A cotton wool plug in the neck lets the gas out but stops acid spray escaping.
This method suits heavy gases such as carbon dioxide, but hydrogen is so light that the mass change is too small to measure well.
A reaction that forms a precipitate can be timed until a cross drawn under the flask disappears from view.
1. Sodium thiosulfate reacts with hydrochloric acid to form a yellow sulfur precipitate: $Na2S2O3(aq)+2HCl(aq)→2NaCl(aq)+S(s)+SO2(g)+H2O(l)\text{Na}_2\text{S}_2\text{O}_3\text{(aq)} + 2\text{HCl(aq)} \rightarrow 2\text{NaCl(aq)} + \text{S(s)} + \text{SO}_2\text{(g)} + \text{H}_2\text{O(l)}$
A reaction that changes colour can be timed until the colour change reaches the same stage each time.

Note

Choose the method from what the reaction produces: a gas suggests gas volume or mass loss, while a precipitate or colour change suggests timing a visible change.

The gradient of the graph shows the rate

Definition

Gradient

The steepness of a line on a graph, found by dividing the change in the vertical value by the change in the horizontal value.

On a graph of gas volume against time, a steeper curve means gas is given off faster, so the rate is greater.
The curve is steepest at the start, because the reactants are at their most concentrated.
The curve gets less steep as the reactants are used up, so the rate decreases.
A horizontal line means the reaction has stopped, because one of the reactants has been used up.
A graph of mass against time falls as gas escapes, and it too levels off when the reaction stops.

Common Mistake

Do not judge which reaction was faster from the height of a curve, because height shows how much product has formed, while the gradient shows the rate.

Comparing curves and drawing tangents shows how rates differ

When two curves are drawn on the same axes, the one that is steeper at the start shows the faster reaction.
If both curves level off at the same height, the same amount of product formed, so the amount of the used-up reactant was the same.
1. A curve that levels off at half the height shows that half as much of that reactant was used.
The rate at an instant is the gradient of a tangent to the curve at that time.
1. Draw a straight line that touches the curve at that point, then divide its rise by its run using two points far apart on the line.

Example

A tangent drawn at $s20\,\text{s}$ on a gas volume curve passes through $cm3)(0\,\text{s}, 10\,\text{cm}^3)$ and $cm3)(40\,\text{s}, 50\,\text{cm}^3)$ ; find the rate at $s20\,\text{s}$ .
Gradient $\dfrac{50 - 10}{40 - 0} = \dfrac{40\,\text{cm}^3}{40\,\text{s}}$ .
Rate at $cm3/s20\,\text{s} = 1.0\,\text{cm}^3/\text{s}$ .

Self review

How is the mean rate of a reaction calculated?
Name two pieces of apparatus that can collect and measure the volume of a gas.
Why is the mass-loss method unsuitable for a reaction that gives off hydrogen?
What does a horizontal section of a gas volume curve show?
How do you find the rate of a reaction at a particular time from a curve?

5.2.2 Factors affecting rate and collision theory

Colliding particles need at least the activation energy

Definition

Activation energy

The minimum energy that colliding particles must have for a reaction to happen.

Definition

Successful collision

A collision between reactant particles that has at least the activation energy, so the particles react.

Reactant particles must collide with each other before they can react.
Most collisions do not lead to a reaction, because the particles have too little energy.
Only collisions with energy equal to or greater than the activation energy are successful.
The rate of reaction depends on the number of successful collisions per second.
So anything that makes collisions more frequent or more energetic increases the rate.

Key Idea

Every rate explanation comes back to how often particles collide and what fraction of those collisions have enough energy to react.

Higher temperature gives more frequent, more energetic collisions

Definition

Kinetic energy

The energy an object or particle has because it is moving.

Raising the temperature increases the rate of a reaction, and for many reactions a $∘C10\ ^{\circ}\text{C}$ rise roughly doubles it.
The particles gain kinetic energy and move faster.
1. Faster particles collide more frequently, so there are more collisions per second.
More of the collisions have energy equal to or greater than the activation energy.
1. This effect is the larger of the two, so a much greater fraction of collisions is successful.
Lowering the temperature has the opposite effect, which is why food keeps longer in a fridge.

Practical

The effect of temperature on rate

Aim: find how temperature affects the rate of the reaction between magnesium and hydrochloric acid.
Apparatus: a conical flask with a bung and delivery tube, a $cm350\,\text{cm}^3$ measuring cylinder, a $cm3100\,\text{cm}^3$ measuring cylinder to collect the gas, a trough of water, a water bath, a thermometer, a stop clock, a ruler, scissors, magnesium ribbon and dilute hydrochloric acid.

Method

Cut magnesium ribbon into $cm4\,\text{cm}$ strips of equal length.
Measure $cm350\,\text{cm}^3$ of dilute hydrochloric acid into the conical flask and warm it in the water bath until it reaches the chosen temperature.
Fill the $cm3100\,\text{cm}^3$ measuring cylinder with water and turn it upside down in the trough, with the delivery tube from the flask leading into it.
Add a strip of magnesium to the acid, replace the bung at once and start the stop clock.
Stop the clock when $cm330\,\text{cm}^3$ of hydrogen has collected.
Repeat at four or five other temperatures, using the same volume and concentration of acid and the same length of ribbon.
Work out the mean rate at each temperature by dividing the volume of gas by the time, then plot rate against temperature.

Accuracy and safety

Measure the temperature of the acid immediately before adding the magnesium, because it starts to cool once out of the water bath.
Repeat each temperature and find a mean time, so that any anomalous result can be spotted.
Keep flames away, because hydrogen is flammable.
Wear eye protection when handling the acid.

Higher concentration or pressure makes collisions more frequent

Increasing the concentration of a solution puts more reactant particles in the same volume.
1. The particles are closer together, so they collide more frequently and the rate increases.
Increasing the pressure of a gas squeezes the same number of particles into a smaller volume.
1. This also gives more frequent collisions, so pressure affects the rate only when a reactant is a gas.
Neither change gives the particles more energy, so the fraction of collisions that succeed stays the same.

Example

Magnesium ribbon fizzes much more vigorously in concentrated hydrochloric acid than in dilute acid, because acid particles hit the surface of the metal more frequently.

Smaller pieces react faster through larger surface area

Definition

Surface area to volume ratio

The surface area of a particle divided by its volume.

Only the particles on the surface of a solid can collide with particles in the solution or gas around it.
Cutting a solid into smaller pieces keeps its volume the same but increases its surface area, so its surface area to volume ratio increases.
1. A greater fraction of the solid's particles is then exposed to collisions.
Collisions become more frequent, so the rate increases, although the energy of each collision is unchanged.
Powdered calcium carbonate reacts with acid much faster than the same mass of marble chips.
1. Very fine powders such as flour dust can burn so fast in air that they explode.

Example

Compare the surface area to volume ratio of a $cm2\,\text{cm}$ cube with that of the same cube cut into eight $cm1\,\text{cm}$ cubes.
For the large cube, surface area $\times 2^2 = 24\,\text{cm}^2$ and volume $2^3 = 8\,\text{cm}^3$ , so the ratio is $24 : 8 = 3 : 1$ .
For the eight small cubes, surface area $\times 6 \times 1^2 = 48\,\text{cm}^2$ and volume $8\,\text{cm}^3$ , so the ratio is $48 : 8 = 6 : 1$ .
Cutting the cube doubles the surface area to volume ratio, so more particles are exposed to collisions.

Explaining rate changes through collision frequency and energy

For temperature, link it to faster particles, more frequent collisions and more collisions with at least the activation energy.
For concentration or pressure, link it to more particles in the same volume and more frequent collisions.
For size of pieces, link it to the surface area to volume ratio, more exposed particles and more frequent collisions.
For a decrease in any factor, reverse the chain to give fewer successful collisions per second.
Say that the reaction has a greater rate or takes less time, rather than calling it a faster reaction time.

Exam technique

When a question asks you to explain a change in rate, name the factor, say what happens to the particles and finish with the effect on successful collisions per second.

Self review

What is a successful collision?
Give the two reasons why a higher temperature increases the rate of reaction.
Why does increasing the concentration of an acid increase the rate of its reaction with magnesium?
Why does pressure affect only reactions involving gases?
Why does powdered calcium carbonate react faster than marble chips of the same mass?

5.2.3 Catalysts, activation energy and enzymes

A catalyst increases rate and ends chemically unchanged

Definition

Catalyst

A substance that increases the rate of a reaction by providing an alternative pathway with a lower activation energy without being used up.

A catalyst takes part in the reaction but is regenerated, so the same mass is left at the end.
Because it is not used up, only a small amount of catalyst is needed and it can be used again.
A catalyst does not change the amount of product formed, only how quickly it forms.
1. A catalysed reaction reaches the same final volume of gas as an uncatalysed one, but sooner.
Catalysts are usually specific, so a catalyst that speeds up one reaction may have no effect on another.

Common Mistake

Do not write that a catalyst does not take part in the reaction, because it does take part and is then re-formed.

Identifying a catalyst and writing it over the arrow

In an equation a catalyst is written over the arrow, not as a reactant or product.
In $2H2O2(aq)→MnO22H2O(l)+O2(g)2\text{H}_2\text{O}_2\text{(aq)} \xrightarrow{\text{MnO}_2} 2\text{H}_2\text{O(l)} + \text{O}_2\text{(g)}$ , manganese(IV) oxide is the catalyst.
In an experiment, a substance is a catalyst if the rate increases and the same mass of it can be recovered afterwards.
1. Filter off the solid after the reaction, then wash, dry and reweigh it.
Transition metals and their compounds are common industrial catalysts, such as iron in making ammonia and platinum in car catalytic converters.

Practical

Testing catalysts for hydrogen peroxide

Aim: find out which of several solids catalyse the decomposition of hydrogen peroxide.
Apparatus: a conical flask with a bung and delivery tube, a $cm3100\,\text{cm}^3$ gas syringe, a clamp and stand, a $cm325\,\text{cm}^3$ measuring cylinder, a digital balance, a stop clock, a filter funnel and filter paper, hydrogen peroxide solution and solids to test, such as manganese(IV) oxide, copper(II) oxide and zinc oxide.

Method

Measure $cm325\,\text{cm}^3$ of hydrogen peroxide solution into the conical flask.
Weigh $g0.5\,\text{g}$ of the first solid, such as manganese(IV) oxide.
Add the solid to the flask, replace the bung at once and start the stop clock.
Record the volume of oxygen in the gas syringe every 10 seconds.
Repeat with no solid and with each other solid, keeping the volume, concentration and temperature of the hydrogen peroxide the same.
Filter off each solid at the end, then wash, dry and reweigh it.

Accuracy and safety

The run with no solid is a control, which shows how fast hydrogen peroxide decomposes on its own.
The solid that gives the steepest line at the start of the graph is the most effective catalyst.
A catalyst has the same mass at the end as at the start, because it is not used up.
Wear eye protection, because hydrogen peroxide is an irritant.

A catalyst gives a pathway with lower activation energy

Definition

Activation energy

The minimum energy that colliding particles must have for a reaction to happen.

On a reaction profile, the activation energy appears as the height of the peak above the energy of the reactants.
With a catalyst, the reaction follows a different route whose peak is lower.
At the same temperature, more of the collisions now have energy equal to or greater than the lower activation energy.
1. So there are more successful collisions per second, and the rate increases.
The catalyst does not give the particles more energy, and it does not change the overall energy change of the reaction.

Note

On a reaction profile with and without a catalyst, the reactant and product levels stay the same, and only the height of the peak between them is lower for the catalysed pathway.

Enzymes are the catalysts in living things

Definition

Enzyme

A protein that acts as a catalyst in living organisms.

Enzymes catalyse the reactions in living cells, which would otherwise be far too slow at body temperature.
Like other catalysts, enzymes are not used up and work by lowering the activation energy.
Each enzyme is specific, so it catalyses only one reaction or a small group of similar reactions.
Catalase, found in liver and potato, catalyses the breakdown of hydrogen peroxide into water and oxygen.
1. Amylase in saliva catalyses the breakdown of starch into sugars.
Enzymes in yeast catalyse fermentation, which turns sugar into ethanol and carbon dioxide.

Example

Dropping a small piece of raw liver into hydrogen peroxide solution makes it froth with oxygen almost at once, because the catalase in the liver catalyses the decomposition.

Self review

What happens to the mass of a catalyst during a reaction?
Where is a catalyst written in a chemical equation?
How would you show that manganese(IV) oxide is a catalyst for the decomposition of hydrogen peroxide?
How does a catalyst increase the rate of a reaction?
What are enzymes?

5.2 Controlling reactions

5.2.1 Measuring and interpreting rate of reaction

Mean rate is change in amount over time

Following a reaction by measuring gas volume

Marble chips and hydrochloric acid

Method

Accuracy and safety

Following reactions by mass loss or visible change

The gradient of the graph shows the rate

Comparing curves and drawing tangents shows how rates differ

5.2.2 Factors affecting rate and collision theory

Colliding particles need at least the activation energy

Higher temperature gives more frequent, more energetic collisions

The effect of temperature on rate

Method

Accuracy and safety

Higher concentration or pressure makes collisions more frequent

Smaller pieces react faster through larger surface area

Explaining rate changes through collision frequency and energy

5.2.3 Catalysts, activation energy and enzymes

A catalyst increases rate and ends chemically unchanged

Identifying a catalyst and writing it over the arrow

Testing catalysts for hydrogen peroxide

Method

Accuracy and safety

A catalyst gives a pathway with lower activation energy

Enzymes are the catalysts in living things

Review 5.2 Controlling reactions by teaching Genie

Recap your knowledge with an interactive lesson

Put it into practice with exam-style questions

Remember key concepts with flashcards

5.2 Controlling reactions Revision Guide

5.2 Controlling reactions

5.2.1 Measuring and interpreting rate of reaction

Mean rate is change in amount over time

Following a reaction by measuring gas volume

Marble chips and hydrochloric acid

Method

Accuracy and safety

Following reactions by mass loss or visible change

The gradient of the graph shows the rate

Comparing curves and drawing tangents shows how rates differ

5.2.2 Factors affecting rate and collision theory

Colliding particles need at least the activation energy

Higher temperature gives more frequent, more energetic collisions

The effect of temperature on rate

Method

Accuracy and safety

Higher concentration or pressure makes collisions more frequent

Smaller pieces react faster through larger surface area

Explaining rate changes through collision frequency and energy

5.2.3 Catalysts, activation energy and enzymes

A catalyst increases rate and ends chemically unchanged

Identifying a catalyst and writing it over the arrow

Testing catalysts for hydrogen peroxide

Method

Accuracy and safety

A catalyst gives a pathway with lower activation energy

Enzymes are the catalysts in living things

Review 5.2 Controlling reactions by teaching Genie

Recap your knowledge with an interactive lesson

Put it into practice with exam-style questions

Remember key concepts with flashcards

5.2 Controlling reactions Revision Guide