Energetics

What you'll learn

How to tell whether a reaction is exothermic or endothermic.
How simple calorimetry experiments measure temperature changes.
How to calculate heat energy using $mc\Delta T$ and molar enthalpy change, $ΔH\Delta H$ .
How energy level diagrams and bond energies explain reaction energy changes (Paper 2 only).

Energy changes in chemical reactions

Chemical reactions involve energy transfers. Often we notice this as a temperature change in the surroundings: the reaction mixture, thermometer, cup, air, or water being heated.

Definition

Exothermic and endothermic reactions

An exothermic reaction gives out heat energy to the surroundings, usually causing a temperature rise. An endothermic reaction takes in heat energy from the surroundings, usually causing a temperature fall.

Common exothermic reactions include combustion, many neutralisation reactions, and many displacement reactions. For example:

$CH4(g)+2O2(g)→CO2(g)+2H2O(l)CH_4(g) + 2O_2(g) \rightarrow CO_2(g) + 2H_2O(l)$

Common endothermic changes include some salts dissolving in water, such as ammonium nitrate:

$NH4NO3(s)→NH4+(aq)+NO3−(aq)NH_4NO_3(s) \rightarrow NH_4^+(aq) + NO_3^-(aq)$

Key Idea

Temperature change tells you the direction of heat transfer

If the reaction mixture gets warmer, heat has been released by the reaction: it is exothermic. If it gets colder, heat has been absorbed by the reaction: it is endothermic.

Example

Using temperature change to identify the energy transfer

A student mixes two solutions. The temperature changes from 19.5 °C to 31.0 °C.

Compare the final temperature with the initial temperature: 31.0 °C is higher than 19.5 °C, so the mixture warmed up.
A temperature rise means heat energy was transferred from the reacting chemicals to the surroundings.
Therefore, the reaction is exothermic.

Calorimetry: measuring heat changes

Definition

Calorimetry

Calorimetry is the experimental measurement of heat energy changes, usually by measuring the temperature change of water or a solution.

For reactions in solution, you usually use an insulated polystyrene cup. The cup reduces heat transfer to the room, so the measured temperature change is closer to the true value.

Polystyrene cup calorimetry apparatus with thermometer, lid, stirrer and measured reactant volume

Practical method: reactions in solution

This method can be used for salts dissolving, neutralisation reactions, and displacement reactions.

Measure a known volume of solution or water into a polystyrene cup.
Record the initial temperature.
Add the second reactant, such as a salt, an alkali, or a metal powder.
Quickly replace the lid, stir gently, and record the highest or lowest temperature reached.
Calculate the temperature change, $ΔT\Delta T$ .

Examples you might meet:

Salt dissolving: add a measured mass of salt to water. Some salts cause a temperature fall; others cause a temperature rise.
Neutralisation: mix an acid and alkali, such as $H C l (a q)$ and $N a O H (a q)$ . These usually give a temperature rise.
Displacement: add a metal to a salt solution, such as magnesium to copper(II) sulfate solution:
$CuSO_4(aq) \rightarrow MgSO_4(aq) + Cu(s)$

Practical method: combustion

In combustion calorimetry, a fuel is burned to heat water. The temperature rise of the water is used to estimate the energy released by the fuel.

Combustion calorimetry apparatus showing a spirit burner heating water in a copper calorimeter

A typical method is:

Put a known mass or volume of water in a copper calorimeter.
Record the initial temperature of the water.
Weigh the spirit burner and fuel.
Burn the fuel to heat the water, stirring gently.
Record the final temperature of the water.
Reweigh the burner to find the mass of fuel burned.

Tip

Key variables in calorimetry practicals

Keep the volume of solution, concentration, starting temperature, insulation, stirring method, and apparatus the same when comparing reactions. For combustion, keep the mass of water and the distance between flame and calorimeter constant.

Common Mistake

Heat loss to the surroundings

Calorimetry experiments are often inaccurate because heat is lost to the air, cup, thermometer, or bench. In combustion, incomplete combustion and heat escaping around the calorimeter make the calculated energy released too small.

Calculating heat energy: $mc\Delta T$

The heat energy transferred to or from water or solution is calculated using:

mc\Delta T

Where:

$Q$ is the heat energy change in joules, J.
$m$ is the mass of water or solution in grams, g.
$c$ is the specific heat capacity. For water and dilute solutions, use 4.18 J/g °C.
$ΔT\Delta T$ is the temperature change in °C.

For dilute aqueous solutions, you usually assume the density is 1 g/cm³. So 50 cm³ of solution has a mass of about 50 g.

Definition

Specific heat capacity

Specific heat capacity is the energy needed to raise the temperature of 1 g of a substance by 1 °C.

Example

Calculating heat energy from a temperature change

50 cm³ of acid is mixed with 50 cm³ of alkali. The temperature rises from 20.0 °C to 26.8 °C. Calculate the heat energy transferred to the solution.

Find the total mass of solution. The total volume is 100 cm³, so the mass is approximately 100 g.
Calculate the temperature change:
$∘C\Delta T = 26.8 - 20.0 = 6.8\ ^\circ\text{C}$
Substitute into $mc\Delta T$ :
$\times 4.18 \times 6.8 = 2842.4\ \text{J}$
Convert to kilojoules if needed:
$kJ2842.4\ \text{J} = 2.84\ \text{kJ}$

Common Mistake

Using the wrong mass

Use the mass of the water or solution being heated, not the mass of the solid reactant or the mass of the fuel. For aqueous solutions, volume in cm³ is usually taken as mass in g.

Molar enthalpy change, $ΔH\Delta H$

Definition

Molar enthalpy change

The molar enthalpy change, $ΔH\Delta H$ , is the heat energy change per mole of substance reacting or formed. It is usually measured in kJ/mol.

For a reaction mixture that warms up, the solution gains heat, so the reaction has released heat. That means the reaction’s $ΔH\Delta H$ is negative.

A useful relationship is:

ΔH=−Qn\Delta H = -\frac{Q}{n}

for an exothermic reaction where $Q$ is the heat gained by the solution and $n$ is the number of moles of the chosen reactant.

Example

Calculating molar enthalpy of neutralisation

50.0 cm³ of 1.00 mol/dm³ hydrochloric acid reacts with 50.0 cm³ of 1.00 mol/dm³ sodium hydroxide. The temperature rises by 6.8 °C. Calculate $ΔH\Delta H$ for the neutralisation.

$\rightarrow NaCl(aq) + H_2O(l)$

Calculate the heat transferred to the solution. From the previous method:
$2.84\ \text{kJ}$
Calculate moles of acid used. Convert 50.0 cm³ to 0.0500 dm³:
$\times 0.0500 = 0.0500\ \text{mol}$
Divide heat energy by moles and include the sign. The temperature rose, so the reaction is exothermic:
$kJ/mol\Delta H = -\frac{2.84}{0.0500} = -56.8\ \text{kJ/mol}$

Tip

Sign of enthalpy change

Temperature rise means $ΔH\Delta H$ is negative. Temperature fall means $ΔH\Delta H$ is positive.

Energy level diagrams (Paper 2 only)

Energy level diagrams show the energy of reactants and products during a reaction.

Energy level diagrams for exothermic and endothermic reactions showing activation energy and enthalpy change

Definition

Activation energy

Activation energy is the minimum energy particles need for a successful reaction to occur.

In an exothermic energy level diagram:

Reactants are higher in energy than products.
The arrow for $ΔH\Delta H$ points down.
$ΔH\Delta H$ is negative.

In an endothermic energy level diagram:

Reactants are lower in energy than products.
The arrow for $ΔH\Delta H$ points up.
$ΔH\Delta H$ is positive.

Key Idea

Reading energy diagrams

The vertical gap between reactants and products is the enthalpy change, $ΔH\Delta H$ . The hump shows the activation energy needed before the reaction can happen.

Bond energies (Paper 2 only)

Chemical reactions involve breaking bonds in the reactants and making new bonds in the products.

Key Idea

Breaking and making bonds

Bond-breaking is endothermic because energy is needed to break attractions between atoms. Bond-making is exothermic because energy is released when new attractions form.

To estimate the enthalpy change from bond energies:

bonds\Delta H = \text{energy to break bonds} - \text{energy released making bonds}

So:

If more energy is released making bonds than is taken in breaking bonds, $ΔH\Delta H$ is negative: exothermic.
If more energy is taken in breaking bonds than is released making bonds, $ΔH\Delta H$ is positive: endothermic.

Example

Using bond energies to calculate enthalpy change

Use these bond energies to calculate $ΔH\Delta H$ for:

$H2(g)+Cl2(g)→2HCl(g)H_2(g) + Cl_2(g) \rightarrow 2HCl(g)$

Bond energies: H–H = 436 kJ/mol, Cl–Cl = 242 kJ/mol, H–Cl = 431 kJ/mol.

Identify bonds broken in the reactants: one H–H bond and one Cl–Cl bond.
Energy absorbed = $678\ \text{kJ/mol}$
Identify bonds made in the products: two H–Cl bonds.
Energy released = $\times 431 = 862\ \text{kJ/mol}$
Calculate the enthalpy change:
$kJ/mol\Delta H = 678 - 862 = -184\ \text{kJ/mol}$
Interpret the sign: $ΔH\Delta H$ is negative, so the reaction is exothermic.

Common Mistake

Reversing the bond energy calculation

Do not do “made minus broken”. For bond energy calculations, use broken minus made: energy taken in to break bonds minus energy released when bonds form.

Exam technique

In the exam

For calorimetry calculations, write down $mc\Delta T$ , use mass in g, and convert J to kJ before calculating $ΔH\Delta H$ .
Always link a temperature rise to exothermic and a temperature fall to endothermic.
For energy diagrams, check whether products are above or below reactants before deciding the sign of $ΔH\Delta H$ .
For bond energies, count every bond carefully, then use broken minus made.

Self review

Check yourself

Why does a reaction mixture get colder during an endothermic reaction?
In a calorimetry experiment, why is a polystyrene cup used instead of a glass beaker?
For $CH4+2O2→CO2+2H2OCH_4 + 2O_2 \rightarrow CO_2 + 2H_2O$ , which bonds are broken and which bonds are made?

What you'll learn

How to tell whether a reaction is exothermic or endothermic.
How simple calorimetry experiments measure temperature changes.
How to calculate heat energy using $mc\Delta T$ and molar enthalpy change, $ΔH\Delta H$ .
How energy level diagrams and bond energies explain reaction energy changes (Paper 2 only).

Energy changes in chemical reactions

Chemical reactions involve energy transfers. Often we notice this as a temperature change in the surroundings: the reaction mixture, thermometer, cup, air, or water being heated.

Definition

Exothermic and endothermic reactions

Common exothermic reactions include combustion, many neutralisation reactions, and many displacement reactions. For example:

$CH4(g)+2O2(g)→CO2(g)+2H2O(l)CH_4(g) + 2O_2(g) \rightarrow CO_2(g) + 2H_2O(l)$

Common endothermic changes include some salts dissolving in water, such as ammonium nitrate:

$NH4NO3(s)→NH4+(aq)+NO3−(aq)NH_4NO_3(s) \rightarrow NH_4^+(aq) + NO_3^-(aq)$

Key Idea

Temperature change tells you the direction of heat transfer

If the reaction mixture gets warmer, heat has been released by the reaction: it is exothermic. If it gets colder, heat has been absorbed by the reaction: it is endothermic.

Example

Using temperature change to identify the energy transfer

A student mixes two solutions. The temperature changes from 19.5 °C to 31.0 °C.

Compare the final temperature with the initial temperature: 31.0 °C is higher than 19.5 °C, so the mixture warmed up.
A temperature rise means heat energy was transferred from the reacting chemicals to the surroundings.
Therefore, the reaction is exothermic.

Calorimetry: measuring heat changes

Definition

Calorimetry

Calorimetry is the experimental measurement of heat energy changes, usually by measuring the temperature change of water or a solution.

For reactions in solution, you usually use an insulated polystyrene cup. The cup reduces heat transfer to the room, so the measured temperature change is closer to the true value.

Polystyrene cup calorimetry apparatus with thermometer, lid, stirrer and measured reactant volume

Practical method: reactions in solution

This method can be used for salts dissolving, neutralisation reactions, and displacement reactions.

Measure a known volume of solution or water into a polystyrene cup.
Record the initial temperature.
Add the second reactant, such as a salt, an alkali, or a metal powder.
Quickly replace the lid, stir gently, and record the highest or lowest temperature reached.
Calculate the temperature change, $ΔT\Delta T$ .

Examples you might meet:

Salt dissolving: add a measured mass of salt to water. Some salts cause a temperature fall; others cause a temperature rise.
Neutralisation: mix an acid and alkali, such as $H C l (a q)$ and $N a O H (a q)$ . These usually give a temperature rise.
Displacement: add a metal to a salt solution, such as magnesium to copper(II) sulfate solution:
$CuSO_4(aq) \rightarrow MgSO_4(aq) + Cu(s)$

Practical method: combustion

In combustion calorimetry, a fuel is burned to heat water. The temperature rise of the water is used to estimate the energy released by the fuel.

Combustion calorimetry apparatus showing a spirit burner heating water in a copper calorimeter

A typical method is:

Put a known mass or volume of water in a copper calorimeter.
Record the initial temperature of the water.
Weigh the spirit burner and fuel.
Burn the fuel to heat the water, stirring gently.
Record the final temperature of the water.
Reweigh the burner to find the mass of fuel burned.

Tip

Key variables in calorimetry practicals

Common Mistake

Heat loss to the surroundings

Calculating heat energy: $mc\Delta T$

The heat energy transferred to or from water or solution is calculated using:

mc\Delta T

Where:

$Q$ is the heat energy change in joules, J.
$m$ is the mass of water or solution in grams, g.
$c$ is the specific heat capacity. For water and dilute solutions, use 4.18 J/g °C.
$ΔT\Delta T$ is the temperature change in °C.

For dilute aqueous solutions, you usually assume the density is 1 g/cm³. So 50 cm³ of solution has a mass of about 50 g.

Definition

Specific heat capacity

Specific heat capacity is the energy needed to raise the temperature of 1 g of a substance by 1 °C.

Example

Calculating heat energy from a temperature change

50 cm³ of acid is mixed with 50 cm³ of alkali. The temperature rises from 20.0 °C to 26.8 °C. Calculate the heat energy transferred to the solution.

Find the total mass of solution. The total volume is 100 cm³, so the mass is approximately 100 g.
Calculate the temperature change:
$∘C\Delta T = 26.8 - 20.0 = 6.8\ ^\circ\text{C}$
Substitute into $mc\Delta T$ :
$\times 4.18 \times 6.8 = 2842.4\ \text{J}$
Convert to kilojoules if needed:
$kJ2842.4\ \text{J} = 2.84\ \text{kJ}$

Common Mistake

Using the wrong mass

Use the mass of the water or solution being heated, not the mass of the solid reactant or the mass of the fuel. For aqueous solutions, volume in cm³ is usually taken as mass in g.

Molar enthalpy change, $ΔH\Delta H$

Definition

Molar enthalpy change

The molar enthalpy change, $ΔH\Delta H$ , is the heat energy change per mole of substance reacting or formed. It is usually measured in kJ/mol.

For a reaction mixture that warms up, the solution gains heat, so the reaction has released heat. That means the reaction’s $ΔH\Delta H$ is negative.

A useful relationship is:

ΔH=−Qn\Delta H = -\frac{Q}{n}

for an exothermic reaction where $Q$ is the heat gained by the solution and $n$ is the number of moles of the chosen reactant.

Example

Calculating molar enthalpy of neutralisation

50.0 cm³ of 1.00 mol/dm³ hydrochloric acid reacts with 50.0 cm³ of 1.00 mol/dm³ sodium hydroxide. The temperature rises by 6.8 °C. Calculate $ΔH\Delta H$ for the neutralisation.

$\rightarrow NaCl(aq) + H_2O(l)$

Calculate the heat transferred to the solution. From the previous method:
$2.84\ \text{kJ}$
Calculate moles of acid used. Convert 50.0 cm³ to 0.0500 dm³:
$\times 0.0500 = 0.0500\ \text{mol}$
Divide heat energy by moles and include the sign. The temperature rose, so the reaction is exothermic:
$kJ/mol\Delta H = -\frac{2.84}{0.0500} = -56.8\ \text{kJ/mol}$

Tip

Sign of enthalpy change

Temperature rise means $ΔH\Delta H$ is negative. Temperature fall means $ΔH\Delta H$ is positive.

Energy level diagrams (Paper 2 only)

Energy level diagrams show the energy of reactants and products during a reaction.

Energy level diagrams for exothermic and endothermic reactions showing activation energy and enthalpy change

Definition

Activation energy

Activation energy is the minimum energy particles need for a successful reaction to occur.

In an exothermic energy level diagram:

Reactants are higher in energy than products.
The arrow for $ΔH\Delta H$ points down.
$ΔH\Delta H$ is negative.

In an endothermic energy level diagram:

Reactants are lower in energy than products.
The arrow for $ΔH\Delta H$ points up.
$ΔH\Delta H$ is positive.

Key Idea

Reading energy diagrams

The vertical gap between reactants and products is the enthalpy change, $ΔH\Delta H$ . The hump shows the activation energy needed before the reaction can happen.

Bond energies (Paper 2 only)

Chemical reactions involve breaking bonds in the reactants and making new bonds in the products.

Key Idea

Breaking and making bonds

Bond-breaking is endothermic because energy is needed to break attractions between atoms. Bond-making is exothermic because energy is released when new attractions form.

To estimate the enthalpy change from bond energies:

bonds\Delta H = \text{energy to break bonds} - \text{energy released making bonds}

So:

If more energy is released making bonds than is taken in breaking bonds, $ΔH\Delta H$ is negative: exothermic.
If more energy is taken in breaking bonds than is released making bonds, $ΔH\Delta H$ is positive: endothermic.

Example

Using bond energies to calculate enthalpy change

Use these bond energies to calculate $ΔH\Delta H$ for:

$H2(g)+Cl2(g)→2HCl(g)H_2(g) + Cl_2(g) \rightarrow 2HCl(g)$

Bond energies: H–H = 436 kJ/mol, Cl–Cl = 242 kJ/mol, H–Cl = 431 kJ/mol.

Identify bonds broken in the reactants: one H–H bond and one Cl–Cl bond.
Energy absorbed = $678\ \text{kJ/mol}$
Identify bonds made in the products: two H–Cl bonds.
Energy released = $\times 431 = 862\ \text{kJ/mol}$
Calculate the enthalpy change:
$kJ/mol\Delta H = 678 - 862 = -184\ \text{kJ/mol}$
Interpret the sign: $ΔH\Delta H$ is negative, so the reaction is exothermic.

Common Mistake

Reversing the bond energy calculation

Do not do “made minus broken”. For bond energy calculations, use broken minus made: energy taken in to break bonds minus energy released when bonds form.

Exam technique

In the exam

For calorimetry calculations, write down $mc\Delta T$ , use mass in g, and convert J to kJ before calculating $ΔH\Delta H$ .
Always link a temperature rise to exothermic and a temperature fall to endothermic.
For energy diagrams, check whether products are above or below reactants before deciding the sign of $ΔH\Delta H$ .
For bond energies, count every bond carefully, then use broken minus made.

Self review

Check yourself

Why does a reaction mixture get colder during an endothermic reaction?
In a calorimetry experiment, why is a polystyrene cup used instead of a glass beaker?
For $CH4+2O2→CO2+2H2OCH_4 + 2O_2 \rightarrow CO_2 + 2H_2O$ , which bonds are broken and which bonds are made?

What you'll learn

Energy changes in chemical reactions

Calorimetry: measuring heat changes

Practical method: reactions in solution

Practical method: combustion

Calculating heat energy: $mc\Delta T$

Molar enthalpy change, $ΔH\Delta H$

Energy level diagrams (Paper 2 only)

Bond energies (Paper 2 only)

Review Energetics by teaching Genie

Remember key concepts with flashcards

Energetics

What you'll learn

Energy changes in chemical reactions

Calorimetry: measuring heat changes

Practical method: reactions in solution

Practical method: combustion

Calculating heat energy: $mc\Delta T$

Molar enthalpy change, $ΔH\Delta H$

Energy level diagrams (Paper 2 only)

Bond energies (Paper 2 only)

Review Energetics by teaching Genie

Remember key concepts with flashcards

Energetics

What you'll learn

Energy changes in chemical reactions

Calorimetry: measuring heat changes

Practical method: reactions in solution

Practical method: combustion

Calculating heat energy: Q=mcΔTQ = mc\Delta TQ=mcΔT

Molar enthalpy change, ΔH\Delta HΔH

Energy level diagrams (Paper 2 only)

Bond energies (Paper 2 only)

Review Energetics by teaching Genie

Remember key concepts with flashcards

Energetics Revision Guide

Energetics

What you'll learn

Energy changes in chemical reactions

Calorimetry: measuring heat changes

Practical method: reactions in solution

Practical method: combustion

Calculating heat energy: Q=mcΔTQ = mc\Delta TQ=mcΔT

Molar enthalpy change, ΔH\Delta HΔH

Energy level diagrams (Paper 2 only)

Bond energies (Paper 2 only)

Review Energetics by teaching Genie

Remember key concepts with flashcards

Energetics Revision Guide

Calculating heat energy: $mc\Delta T$

Molar enthalpy change, $ΔH\Delta H$

Calculating heat energy: $mc\Delta T$

Molar enthalpy change, $ΔH\Delta H$