Rate of reaction

Definition

Rate of reaction

The amount of a reactant used up or the amount of a product formed per unit time.

The rate of reaction tells you how quickly reactants are turned into products.
You find a rate by measuring how quickly a reactant is used up or how quickly a product is formed.
The average rate over a period of time is found using $taken\text{average rate} = \dfrac{\text{change in amount of product or reactant}}{\text{time taken}}$ .
The units depend on what you measure, such as $cm3/s\text{cm}^3/\text{s}$ for a volume of gas or $g/s\text{g}/\text{s}$ for a mass.
A reaction is usually fastest at the start, when the most reactant is present.
The reaction slows down as the reactants are used up.
The reaction stops when one of the reactants has been completely used up.
Five factors change the rate of a reaction: concentration, pressure, surface area, temperature and the presence of a catalyst.

Concentration and pressure

Increasing the concentration of a reacting solution increases the rate of reaction.
Magnesium ribbon fizzes much more vigorously in $mol/dm32.0\ \text{mol/dm}^3$ hydrochloric acid than in $mol/dm30.5\ \text{mol/dm}^3$ hydrochloric acid.
Decreasing the concentration, for example by diluting the acid with water, decreases the rate.
Increasing the pressure of a reacting gas increases the rate of reaction.
Raising the pressure squeezes the same amount of gas into a smaller volume, which has the same effect as increasing the concentration of the gas.
Pressure only affects the rate of reactions that involve gases.

Surface area

Increasing the surface area of a solid reactant increases the rate of reaction.
Breaking a solid into smaller pieces increases its surface area, so a powder reacts faster than the same mass of lumps.
Calcium carbonate powder fizzes far more quickly in dilute hydrochloric acid than the same mass of large marble chips.
Very fine powders can react so fast that they explode, which is why flour dust in mills and coal dust in mines are serious hazards.

Temperature

Increasing the temperature increases the rate of reaction.
For many reactions, a rise of about $∘C10\ ^\circ\text{C}$ roughly doubles the rate.
Decreasing the temperature decreases the rate, which is why food stays fresh for longer in a fridge.
To compare temperatures fairly, the reactants are warmed in a water bath before they are mixed, and every other factor is kept the same.

Catalysts and enzymes

Definition

Catalyst

A substance that increases the rate of a reaction and is chemically unchanged at the end of the reaction.

Definition

Enzyme

A protein that acts as a biological catalyst, increasing the rate of reactions in living things.

Adding a catalyst increases the rate of a reaction, and removing the catalyst lets the rate fall back again.
A catalyst is unchanged at the end of the reaction, so it has the same mass and the same chemical composition as at the start.
Only a small amount of catalyst is needed, because it is not used up and can be used again.
Manganese(IV) oxide is a catalyst for the decomposition of hydrogen peroxide into water and oxygen: $2H2O2(aq)→2H2O(l)+O2(g)\mathrm{2H_{2}O_{2}(aq) \rightarrow 2H_{2}O(l) + O_{2}(g)}$ .
A catalyst does not change the amount of product formed, only how quickly it forms.
Enzymes are catalysts made by living things, such as the enzymes in yeast that speed up fermentation.
Enzymes work best within a narrow temperature range, and high temperatures change their structure so that they stop working.

Example

Showing that a catalyst is unchanged

Hydrogen peroxide solution on its own gives off oxygen very slowly.
Adding $g0.50\ \text{g}$ of black manganese(IV) oxide powder makes the solution froth with oxygen within seconds.
Once the fizzing stops, the mixture is filtered and the black solid is washed and dried.
The dried solid still has a mass of $g0.50\ \text{g}$ and it speeds up a fresh sample of hydrogen peroxide just as well.
This shows that manganese(IV) oxide increased the rate but was not used up, so it acted as a catalyst.

Measuring the rate of reaction

To follow a reaction, you measure a quantity that changes as the reaction goes on, at regular time intervals.
Volume of gas produced: the gas is collected in a gas syringe or in an upturned measuring cylinder full of water, and its volume is read every few seconds.
Loss in mass: the reaction is carried out in a flask standing on a balance, and the mass falls as a gas escapes into the air.
Time for a precipitate to form: the time taken for a cloudy precipitate to hide a cross drawn under the flask is measured, and a shorter time means a faster rate.
In a fair test you change only one factor, the independent variable, and keep every other factor the same.
The factor you measure, such as the volume of gas, is the dependent variable.

Practical

Measuring gas volume: magnesium and hydrochloric acid

Aim: to find how the concentration of hydrochloric acid affects the rate of its reaction with magnesium by measuring the volume of hydrogen produced.
Apparatus: conical flask, bung fitted with a delivery tube, $cm3100\ \text{cm}^3$ gas syringe held in a clamp and stand, $cm350\ \text{cm}^3$ measuring cylinder, stopwatch, scissors, ruler, magnesium ribbon, hydrochloric acid at $0.5$ , $1.0$ , $1.5$ and $mol/dm32.0\ \text{mol/dm}^3$ , and eye protection.
Method, setting up:
- Clamp the gas syringe horizontally, push the plunger fully in and check that it slides freely and reads zero.
- Cut four $cm3\ \text{cm}$ lengths of magnesium ribbon with scissors, measuring each one against a ruler.
- Measure $cm350\ \text{cm}^3$ of the $mol/dm30.5\ \text{mol/dm}^3$ acid with the measuring cylinder and pour it into the conical flask.
- Connect the delivery tube from the bung to the nozzle of the gas syringe.
Method, running the reaction:
- Drop one length of magnesium into the acid, push the bung in firmly at once and start the stopwatch.
- Read the volume of gas in the syringe every $s10\ \text{s}$ until two readings in a row are the same.
- Rinse the flask and repeat with each of the other concentrations, keeping the volume of acid, the length of magnesium and the room temperature the same.
Method, processing the results:
- Plot volume of hydrogen / $cm3\text{cm}^3$ on the vertical axis against time / $s\text{s}$ on the horizontal axis, with all four sets of results on the same grid.
- Draw a smooth curve of best fit through each set of points and compare how steep each curve is at the start.
Results: the higher the concentration of the acid, the steeper the curve is at the start and the sooner it levels off.
Reaction: magnesium reacts with hydrochloric acid to form magnesium chloride solution and hydrogen, $Mg(s)+2HCl(aq)→MgCl2(aq)+H2(g)\mathrm{Mg(s) + 2HCl(aq) \rightarrow MgCl_{2}(aq) + H_{2}(g)}$ .
Watch out: the acid is in excess in every run, so each curve levels off at the same final volume because the same length of magnesium is used each time.
Safety: wear eye protection, keep flames well away because hydrogen is flammable, and wipe up any spilled acid straight away because it irritates the skin and eyes.

Practical

Measuring loss in mass: marble chips and hydrochloric acid

Aim: to find how the surface area of calcium carbonate affects the rate of its reaction with hydrochloric acid by measuring the loss in mass as carbon dioxide escapes.
Apparatus: conical flask, cotton wool, top-pan balance reading to $g0.01\ \text{g}$ , stopwatch, $cm350\ \text{cm}^3$ measuring cylinder, weighing boat, large marble chips, small marble chips, calcium carbonate powder, $mol/dm31.0\ \text{mol/dm}^3$ hydrochloric acid and eye protection.
Method, setting up:
- Weigh out $g5.00\ \text{g}$ of large marble chips in a weighing boat.
- Measure $cm350\ \text{cm}^3$ of hydrochloric acid into the conical flask and stand the flask on the balance, with a loose plug of cotton wool and the weighing boat of chips beside it on the pan.
Method, running the reaction:
- Note the total mass, tip the chips into the acid, push the cotton wool loosely into the neck of the flask, put the empty weighing boat back on the pan and start the stopwatch.
- The cotton wool lets carbon dioxide escape but stops droplets of acid spraying out, which would add to the loss in mass.
- Record the reading on the balance every $s30\ \text{s}$ for $5$ minutes.
- Repeat with $g5.00\ \text{g}$ of small marble chips and then $g5.00\ \text{g}$ of powder, using fresh acid each time.
Method, processing the results:
- Work out the loss in mass at each time by subtracting each reading from the starting mass.
- Plot loss in mass / $g\text{g}$ against time / $s\text{s}$ for all three sizes on the same grid.
Results: the powder gives the steepest curve, the small chips give a less steep curve and the large chips give the shallowest curve.
Reaction: $CaCO3(s)+2HCl(aq)→CaCl2(aq)+H2O(l)+CO2(g)\mathrm{CaCO_{3}(s) + 2HCl(aq) \rightarrow CaCl_{2}(aq) + H_{2}O(l) + CO_{2}(g)}$
Watch out: the acid is the reactant that runs out in these amounts, so given enough time all three curves level off at the same total loss in mass of about $g1.1\ \text{g}$ .
Safety: wear eye protection and avoid splashing the acid, which irritates the skin and eyes.

Practical

Timing a precipitate: sodium thiosulfate and hydrochloric acid

Aim: to find how temperature affects the rate of the reaction between sodium thiosulfate solution and hydrochloric acid by timing how long the sulfur formed takes to hide a cross.
Apparatus: $cm3100\ \text{cm}^3$ conical flask, sheet of white paper with a bold black cross drawn on it, $cm350\ \text{cm}^3$ and $cm310\ \text{cm}^3$ measuring cylinders, water bath, thermometer, stopwatch, sodium thiosulfate solution, $mol/dm31.0\ \text{mol/dm}^3$ hydrochloric acid and eye protection.
Method, preparing:
- Measure $cm350\ \text{cm}^3$ of sodium thiosulfate solution into the conical flask and warm it in the water bath until the thermometer reads about $∘C20\ ^\circ\text{C}$ .
- Stand the flask on the cross so that the cross can be seen clearly when you look down through the solution.
Method, timing the reaction:
- Add $cm35\ \text{cm}^3$ of hydrochloric acid, swirl the flask once and start the stopwatch.
- Look down through the mouth of the flask and stop the stopwatch at the moment the pale yellow precipitate of sulfur hides the cross completely.
- Record the temperature of the mixture straight away.
- Repeat at about $30$ , $40$ , $50$ and $∘C60\ ^\circ\text{C}$ , using the same volumes, the same flask and the same cross each time.
Method, processing the results:
- Calculate $1time\dfrac{1}{\text{time}}$ for each temperature, because a shorter time means a faster rate.
- Plot $1time\dfrac{1}{\text{time}}$ / $s−1\text{s}^{-1}$ against temperature / $∘C^\circ\text{C}$ and draw a smooth curve of best fit.
Results: as the temperature rises, the cross disappears more quickly, so the time falls and $1time\dfrac{1}{\text{time}}$ increases.
Reaction: $Na2S2O3(aq)+2HCl(aq)→2NaCl(aq)+S(s)+SO2(g)+H2O(l)\mathrm{Na_{2}S_{2}O_{3}(aq) + 2HCl(aq) \rightarrow 2NaCl(aq) + S(s) + SO_{2}(g) + H_{2}O(l)}$
Watch out: deciding when the cross has vanished is a judgement, so the same person should watch the cross in every run.
Safety: wear eye protection, work in a well-ventilated room because the reaction gives off sulfur dioxide, which is toxic and can trigger asthma, and pour the mixtures away promptly once each run is finished.

Interpreting rate graphs

On a graph of the amount of product against time, the gradient, or steepness, of the line shows the rate of reaction.
A steeper line means a faster rate.
The curve is steepest at the start, because this is when the reaction is fastest.
The curve becomes less steep as the reactants are used up and the reaction slows down.
The line becomes horizontal, or levels off, when the reaction has stopped because one reactant has been used up.
The final height of the curve shows the total amount of product, which depends on the amount of the reactant that runs out, not on the rate.
A faster reaction using the same amounts of reactants gives a steeper curve that levels off sooner, but at the same height.
Halving the amount of the reactant that runs out makes the curve level off at half the height.
For a graph of the mass of the flask against time, the curve slopes downwards and then levels off when the reaction stops.
The rate at one instant is found by drawing a tangent to the curve at that time and calculating its gradient.

Example

Average rate from gas volumes

In the magnesium experiment, $cm336\ \text{cm}^3$ of hydrogen is collected in the first $s20\ \text{s}$ .

\text{average rate} = \dfrac{36\ \text{cm}^3}{20\ \text{s}} = 1.8\ \text{cm}^3/\text{s}

Between $s20\ \text{s}$ and $s40\ \text{s}$ , a further $cm318\ \text{cm}^3$ is collected.

\text{average rate} = \dfrac{18\ \text{cm}^3}{20\ \text{s}} = 0.90\ \text{cm}^3/\text{s}

The rate has halved because some of the reactants have already been used up.

Example

Comparing three curves

Curve A comes from $g0.10\ \text{g}$ of magnesium in excess acid at $∘C25\ ^\circ\text{C}$ .
Curve B uses the same amounts at $∘C35\ ^\circ\text{C}$ , so it is steeper at the start and levels off sooner, at the same final volume as A.
Curve C uses $g0.05\ \text{g}$ of magnesium at $∘C25\ ^\circ\text{C}$ , so it levels off at half the final volume of A.
The starting gradients show the order of rates, while the final heights show the amounts of the limiting reactant.

Exam technique

Answering rate questions

When you describe an effect, give its direction clearly, such as “increasing the temperature increases the rate”.
When you describe a graph, use the words steeper, gradient and levels off rather than vague phrases such as “goes up quicker”.
When you are asked when a reaction finished, read off the time at which the line first becomes horizontal.
When curves end at the same height, explain that the same amount of the reactant that was used up was present in each experiment.
When you describe a method, name the apparatus and the measurement, such as “measure the volume of gas with a gas syringe every $s10\ \text{s}$ ”.
Always give a unit with any rate you calculate.

Common Mistake

Do not write that a catalyst is used up or that it makes more product.
Do not confuse how fast a reaction is with how much product it makes.
Do not describe a more concentrated acid as a stronger acid.

Self review

How does increasing the surface area of a solid reactant affect the rate of reaction?
What is a catalyst?
Why does a curve of gas volume against time become horizontal?
Name two ways of measuring the rate of a reaction that produces a gas.
Why does the cotton wool plug matter in the mass-loss method?

Rate of reaction

Definition

Rate of reaction

The amount of a reactant used up or the amount of a product formed per unit time.

The rate of reaction tells you how quickly reactants are turned into products.
You find a rate by measuring how quickly a reactant is used up or how quickly a product is formed.
The average rate over a period of time is found using $taken\text{average rate} = \dfrac{\text{change in amount of product or reactant}}{\text{time taken}}$ .
The units depend on what you measure, such as $cm3/s\text{cm}^3/\text{s}$ for a volume of gas or $g/s\text{g}/\text{s}$ for a mass.
A reaction is usually fastest at the start, when the most reactant is present.
The reaction slows down as the reactants are used up.
The reaction stops when one of the reactants has been completely used up.
Five factors change the rate of a reaction: concentration, pressure, surface area, temperature and the presence of a catalyst.

Concentration and pressure

Increasing the concentration of a reacting solution increases the rate of reaction.
Magnesium ribbon fizzes much more vigorously in $mol/dm32.0\ \text{mol/dm}^3$ hydrochloric acid than in $mol/dm30.5\ \text{mol/dm}^3$ hydrochloric acid.
Decreasing the concentration, for example by diluting the acid with water, decreases the rate.
Increasing the pressure of a reacting gas increases the rate of reaction.
Raising the pressure squeezes the same amount of gas into a smaller volume, which has the same effect as increasing the concentration of the gas.
Pressure only affects the rate of reactions that involve gases.

Surface area

Increasing the surface area of a solid reactant increases the rate of reaction.
Breaking a solid into smaller pieces increases its surface area, so a powder reacts faster than the same mass of lumps.
Calcium carbonate powder fizzes far more quickly in dilute hydrochloric acid than the same mass of large marble chips.
Very fine powders can react so fast that they explode, which is why flour dust in mills and coal dust in mines are serious hazards.

Temperature

Increasing the temperature increases the rate of reaction.
For many reactions, a rise of about $∘C10\ ^\circ\text{C}$ roughly doubles the rate.
Decreasing the temperature decreases the rate, which is why food stays fresh for longer in a fridge.
To compare temperatures fairly, the reactants are warmed in a water bath before they are mixed, and every other factor is kept the same.

Catalysts and enzymes

Definition

Catalyst

A substance that increases the rate of a reaction and is chemically unchanged at the end of the reaction.

Definition

Enzyme

A protein that acts as a biological catalyst, increasing the rate of reactions in living things.

Adding a catalyst increases the rate of a reaction, and removing the catalyst lets the rate fall back again.
A catalyst is unchanged at the end of the reaction, so it has the same mass and the same chemical composition as at the start.
Only a small amount of catalyst is needed, because it is not used up and can be used again.
Manganese(IV) oxide is a catalyst for the decomposition of hydrogen peroxide into water and oxygen: $2H2O2(aq)→2H2O(l)+O2(g)\mathrm{2H_{2}O_{2}(aq) \rightarrow 2H_{2}O(l) + O_{2}(g)}$ .
A catalyst does not change the amount of product formed, only how quickly it forms.
Enzymes are catalysts made by living things, such as the enzymes in yeast that speed up fermentation.
Enzymes work best within a narrow temperature range, and high temperatures change their structure so that they stop working.

Example

Showing that a catalyst is unchanged

Hydrogen peroxide solution on its own gives off oxygen very slowly.
Adding $g0.50\ \text{g}$ of black manganese(IV) oxide powder makes the solution froth with oxygen within seconds.
Once the fizzing stops, the mixture is filtered and the black solid is washed and dried.
The dried solid still has a mass of $g0.50\ \text{g}$ and it speeds up a fresh sample of hydrogen peroxide just as well.
This shows that manganese(IV) oxide increased the rate but was not used up, so it acted as a catalyst.

Measuring the rate of reaction

To follow a reaction, you measure a quantity that changes as the reaction goes on, at regular time intervals.
Volume of gas produced: the gas is collected in a gas syringe or in an upturned measuring cylinder full of water, and its volume is read every few seconds.
Loss in mass: the reaction is carried out in a flask standing on a balance, and the mass falls as a gas escapes into the air.
Time for a precipitate to form: the time taken for a cloudy precipitate to hide a cross drawn under the flask is measured, and a shorter time means a faster rate.
In a fair test you change only one factor, the independent variable, and keep every other factor the same.
The factor you measure, such as the volume of gas, is the dependent variable.

Practical

Measuring gas volume: magnesium and hydrochloric acid

Aim: to find how the concentration of hydrochloric acid affects the rate of its reaction with magnesium by measuring the volume of hydrogen produced.
Apparatus: conical flask, bung fitted with a delivery tube, $cm3100\ \text{cm}^3$ gas syringe held in a clamp and stand, $cm350\ \text{cm}^3$ measuring cylinder, stopwatch, scissors, ruler, magnesium ribbon, hydrochloric acid at $0.5$ , $1.0$ , $1.5$ and $mol/dm32.0\ \text{mol/dm}^3$ , and eye protection.
Method, setting up:
- Clamp the gas syringe horizontally, push the plunger fully in and check that it slides freely and reads zero.
- Cut four $cm3\ \text{cm}$ lengths of magnesium ribbon with scissors, measuring each one against a ruler.
- Measure $cm350\ \text{cm}^3$ of the $mol/dm30.5\ \text{mol/dm}^3$ acid with the measuring cylinder and pour it into the conical flask.
- Connect the delivery tube from the bung to the nozzle of the gas syringe.
Method, running the reaction:
- Drop one length of magnesium into the acid, push the bung in firmly at once and start the stopwatch.
- Read the volume of gas in the syringe every $s10\ \text{s}$ until two readings in a row are the same.
- Rinse the flask and repeat with each of the other concentrations, keeping the volume of acid, the length of magnesium and the room temperature the same.
Method, processing the results:
- Plot volume of hydrogen / $cm3\text{cm}^3$ on the vertical axis against time / $s\text{s}$ on the horizontal axis, with all four sets of results on the same grid.
- Draw a smooth curve of best fit through each set of points and compare how steep each curve is at the start.
Results: the higher the concentration of the acid, the steeper the curve is at the start and the sooner it levels off.
Reaction: magnesium reacts with hydrochloric acid to form magnesium chloride solution and hydrogen, $Mg(s)+2HCl(aq)→MgCl2(aq)+H2(g)\mathrm{Mg(s) + 2HCl(aq) \rightarrow MgCl_{2}(aq) + H_{2}(g)}$ .
Watch out: the acid is in excess in every run, so each curve levels off at the same final volume because the same length of magnesium is used each time.
Safety: wear eye protection, keep flames well away because hydrogen is flammable, and wipe up any spilled acid straight away because it irritates the skin and eyes.

Practical

Measuring loss in mass: marble chips and hydrochloric acid

Aim: to find how the surface area of calcium carbonate affects the rate of its reaction with hydrochloric acid by measuring the loss in mass as carbon dioxide escapes.
Apparatus: conical flask, cotton wool, top-pan balance reading to $g0.01\ \text{g}$ , stopwatch, $cm350\ \text{cm}^3$ measuring cylinder, weighing boat, large marble chips, small marble chips, calcium carbonate powder, $mol/dm31.0\ \text{mol/dm}^3$ hydrochloric acid and eye protection.
Method, setting up:
- Weigh out $g5.00\ \text{g}$ of large marble chips in a weighing boat.
- Measure $cm350\ \text{cm}^3$ of hydrochloric acid into the conical flask and stand the flask on the balance, with a loose plug of cotton wool and the weighing boat of chips beside it on the pan.
Method, running the reaction:
- Note the total mass, tip the chips into the acid, push the cotton wool loosely into the neck of the flask, put the empty weighing boat back on the pan and start the stopwatch.
- The cotton wool lets carbon dioxide escape but stops droplets of acid spraying out, which would add to the loss in mass.
- Record the reading on the balance every $s30\ \text{s}$ for $5$ minutes.
- Repeat with $g5.00\ \text{g}$ of small marble chips and then $g5.00\ \text{g}$ of powder, using fresh acid each time.
Method, processing the results:
- Work out the loss in mass at each time by subtracting each reading from the starting mass.
- Plot loss in mass / $g\text{g}$ against time / $s\text{s}$ for all three sizes on the same grid.
Results: the powder gives the steepest curve, the small chips give a less steep curve and the large chips give the shallowest curve.
Reaction: $CaCO3(s)+2HCl(aq)→CaCl2(aq)+H2O(l)+CO2(g)\mathrm{CaCO_{3}(s) + 2HCl(aq) \rightarrow CaCl_{2}(aq) + H_{2}O(l) + CO_{2}(g)}$
Watch out: the acid is the reactant that runs out in these amounts, so given enough time all three curves level off at the same total loss in mass of about $g1.1\ \text{g}$ .
Safety: wear eye protection and avoid splashing the acid, which irritates the skin and eyes.

Practical

Timing a precipitate: sodium thiosulfate and hydrochloric acid

Aim: to find how temperature affects the rate of the reaction between sodium thiosulfate solution and hydrochloric acid by timing how long the sulfur formed takes to hide a cross.
Apparatus: $cm3100\ \text{cm}^3$ conical flask, sheet of white paper with a bold black cross drawn on it, $cm350\ \text{cm}^3$ and $cm310\ \text{cm}^3$ measuring cylinders, water bath, thermometer, stopwatch, sodium thiosulfate solution, $mol/dm31.0\ \text{mol/dm}^3$ hydrochloric acid and eye protection.
Method, preparing:
- Measure $cm350\ \text{cm}^3$ of sodium thiosulfate solution into the conical flask and warm it in the water bath until the thermometer reads about $∘C20\ ^\circ\text{C}$ .
- Stand the flask on the cross so that the cross can be seen clearly when you look down through the solution.
Method, timing the reaction:
- Add $cm35\ \text{cm}^3$ of hydrochloric acid, swirl the flask once and start the stopwatch.
- Look down through the mouth of the flask and stop the stopwatch at the moment the pale yellow precipitate of sulfur hides the cross completely.
- Record the temperature of the mixture straight away.
- Repeat at about $30$ , $40$ , $50$ and $∘C60\ ^\circ\text{C}$ , using the same volumes, the same flask and the same cross each time.
Method, processing the results:
- Calculate $1time\dfrac{1}{\text{time}}$ for each temperature, because a shorter time means a faster rate.
- Plot $1time\dfrac{1}{\text{time}}$ / $s−1\text{s}^{-1}$ against temperature / $∘C^\circ\text{C}$ and draw a smooth curve of best fit.
Results: as the temperature rises, the cross disappears more quickly, so the time falls and $1time\dfrac{1}{\text{time}}$ increases.
Reaction: $Na2S2O3(aq)+2HCl(aq)→2NaCl(aq)+S(s)+SO2(g)+H2O(l)\mathrm{Na_{2}S_{2}O_{3}(aq) + 2HCl(aq) \rightarrow 2NaCl(aq) + S(s) + SO_{2}(g) + H_{2}O(l)}$
Watch out: deciding when the cross has vanished is a judgement, so the same person should watch the cross in every run.
Safety: wear eye protection, work in a well-ventilated room because the reaction gives off sulfur dioxide, which is toxic and can trigger asthma, and pour the mixtures away promptly once each run is finished.

Interpreting rate graphs

On a graph of the amount of product against time, the gradient, or steepness, of the line shows the rate of reaction.
A steeper line means a faster rate.
The curve is steepest at the start, because this is when the reaction is fastest.
The curve becomes less steep as the reactants are used up and the reaction slows down.
The line becomes horizontal, or levels off, when the reaction has stopped because one reactant has been used up.
The final height of the curve shows the total amount of product, which depends on the amount of the reactant that runs out, not on the rate.
A faster reaction using the same amounts of reactants gives a steeper curve that levels off sooner, but at the same height.
Halving the amount of the reactant that runs out makes the curve level off at half the height.
For a graph of the mass of the flask against time, the curve slopes downwards and then levels off when the reaction stops.
The rate at one instant is found by drawing a tangent to the curve at that time and calculating its gradient.

Example

Average rate from gas volumes

In the magnesium experiment, $cm336\ \text{cm}^3$ of hydrogen is collected in the first $s20\ \text{s}$ .

\text{average rate} = \dfrac{36\ \text{cm}^3}{20\ \text{s}} = 1.8\ \text{cm}^3/\text{s}

Between $s20\ \text{s}$ and $s40\ \text{s}$ , a further $cm318\ \text{cm}^3$ is collected.

\text{average rate} = \dfrac{18\ \text{cm}^3}{20\ \text{s}} = 0.90\ \text{cm}^3/\text{s}

The rate has halved because some of the reactants have already been used up.

Example

Comparing three curves

Curve A comes from $g0.10\ \text{g}$ of magnesium in excess acid at $∘C25\ ^\circ\text{C}$ .
Curve B uses the same amounts at $∘C35\ ^\circ\text{C}$ , so it is steeper at the start and levels off sooner, at the same final volume as A.
Curve C uses $g0.05\ \text{g}$ of magnesium at $∘C25\ ^\circ\text{C}$ , so it levels off at half the final volume of A.
The starting gradients show the order of rates, while the final heights show the amounts of the limiting reactant.

Exam technique

Answering rate questions

When you describe an effect, give its direction clearly, such as “increasing the temperature increases the rate”.
When you describe a graph, use the words steeper, gradient and levels off rather than vague phrases such as “goes up quicker”.
When you are asked when a reaction finished, read off the time at which the line first becomes horizontal.
When curves end at the same height, explain that the same amount of the reactant that was used up was present in each experiment.
When you describe a method, name the apparatus and the measurement, such as “measure the volume of gas with a gas syringe every $s10\ \text{s}$ ”.
Always give a unit with any rate you calculate.

Common Mistake

Do not write that a catalyst is used up or that it makes more product.
Do not confuse how fast a reaction is with how much product it makes.
Do not describe a more concentrated acid as a stronger acid.

Self review

How does increasing the surface area of a solid reactant affect the rate of reaction?
What is a catalyst?
Why does a curve of gas volume against time become horizontal?
Name two ways of measuring the rate of a reaction that produces a gas.
Why does the cotton wool plug matter in the mass-loss method?

6.2.1 Factors affecting the rate of reaction

Rate of reaction

Concentration and pressure

Surface area

Temperature

Catalysts and enzymes

Showing that a catalyst is unchanged

Measuring the rate of reaction

Measuring gas volume: magnesium and hydrochloric acid

Measuring loss in mass: marble chips and hydrochloric acid

Timing a precipitate: sodium thiosulfate and hydrochloric acid

Interpreting rate graphs

Average rate from gas volumes

Comparing three curves

Answering rate questions

Review 6.2.1 Factors affecting the rate of reaction by teaching Genie

6.2.1 Factors affecting the rate of reaction

Rate of reaction

Concentration and pressure

Surface area

Temperature

Catalysts and enzymes

Showing that a catalyst is unchanged

Measuring the rate of reaction

Measuring gas volume: magnesium and hydrochloric acid

Measuring loss in mass: marble chips and hydrochloric acid

Timing a precipitate: sodium thiosulfate and hydrochloric acid

Interpreting rate graphs

Average rate from gas volumes

Comparing three curves

Answering rate questions

Review 6.2.1 Factors affecting the rate of reaction by teaching Genie

6.2.1 Factors affecting the rate of reaction

Rate of reaction

Concentration and pressure

Surface area

Temperature

Catalysts and enzymes

Showing that a catalyst is unchanged

Measuring the rate of reaction

Measuring gas volume: magnesium and hydrochloric acid

Measuring loss in mass: marble chips and hydrochloric acid

Timing a precipitate: sodium thiosulfate and hydrochloric acid

Interpreting rate graphs

Average rate from gas volumes

Comparing three curves

Answering rate questions

Review 6.2.1 Factors affecting the rate of reaction by teaching Genie

6.2.1 Factors affecting the rate of reaction Revision Guide

6.2.1 Factors affecting the rate of reaction

Rate of reaction

Concentration and pressure

Surface area

Temperature

Catalysts and enzymes

Showing that a catalyst is unchanged

Measuring the rate of reaction

Measuring gas volume: magnesium and hydrochloric acid

Measuring loss in mass: marble chips and hydrochloric acid

Timing a precipitate: sodium thiosulfate and hydrochloric acid

Interpreting rate graphs

Average rate from gas volumes

Comparing three curves

Answering rate questions

Review 6.2.1 Factors affecting the rate of reaction by teaching Genie

6.2.1 Factors affecting the rate of reaction Revision Guide