3.1 Energetics Notes

Edexcel IGCSE Chemistry: Revision notes

Key facts

Exothermic reactions release heat (surroundings warm up); endothermic reactions absorb heat (surroundings cool).
Heat energy change: $Q = mc\Delta T$ , with $c$ = 4.2 J/g °C for water.
Molar enthalpy change: $\Delta H = Q$ / moles; negative if exothermic, positive if endothermic.
Exothermic: products lower in energy than reactants. Endothermic: products higher.
Bond breaking is endothermic; bond making is exothermic.

Exothermic and endothermic

Exothermic reactions give out heat; endothermic reactions take it in.

An exothermic reaction gives out heat to the surroundings, so the temperature increases. Examples: combustion, neutralisation and most oxidation. An endothermic reaction takes in heat, so the temperature decreases. Examples: thermal decomposition and dissolving some salts such as ammonium nitrate.

Illustrative energy profiles: exothermic products end lower than the reactants (heat given out); endothermic products end higher (heat taken in).

Exothermic

Gives out heat
Surroundings warm up
Combustion, neutralisation

Endothermic

Takes in heat
Surroundings cool down
Thermal decomposition, dissolving ammonium nitrate

A reaction makes the temperature of the surroundings fall. What type of reaction is it?

Measuring heat change

Measure the temperature change of a known mass of water, then use Q = mcΔT.

Calorimetry works for combustion, displacement, dissolving and neutralisation, using an insulated container and a thermometer. For neutralisation: measure a known volume of acid into an insulated cup and record the temperature, add a known volume of alkali, stir, record the highest (or lowest) temperature, and find the change. $m$ is the mass heated (g), $c$ is 4.2 J/g °C, $\Delta T$ is in °C and $Q$ is in joules.

1
Measure the acid
known volume into an insulated cup; record the temperature
2
Add the alkali
known volume; stir
3
Record the extreme
highest (or lowest) temperature reached
4
Calculate
Q = m × c × ΔT

Measuring the heat change of a neutralisation.

Heat energy change, $Q$ $m \times c \times \Delta T$

Worked example

50 g of water rises in temperature from 20.0 °C to 27.0 °C in a neutralisation experiment. Calculate $Q$ .

1
$\Delta T = 27.0 - 20.0 = 7.0$ °C.
2
$Q = 50 \times 4.2 \times 7.0$ .

Which quantities are needed to calculate $Q$ in a calorimetry experiment?

Molar enthalpy change

Divide the heat energy change by the number of moles and convert to kJ/mol.

The molar enthalpy change $\Delta H$ is the heat energy change per mole of a stated substance reacting. Divide by 1000 to convert J/mol to kJ/mol. $\Delta H$ is negative for exothermic reactions (energy released) and positive for endothermic reactions (energy absorbed).

1
Q = m × c × ΔT
heat energy change in joules
2
Moles
of the stated substance that reacted
3
ΔH = Q ÷ moles
in J/mol
4
Convert and sign
÷ 1000 for kJ/mol; negative if exothermic, positive if endothermic

From a temperature change to a molar enthalpy change.

$\Delta H$ $\dfrac{Q}{\text{moles}}$
Exothermic $\Delta H$ negative
Endothermic $\Delta H$ positive

Worked example

In a neutralisation experiment, 0.05 mol of a substance releases 4200 J of heat. Calculate $\Delta H$ in kJ/mol.

1
It releases heat, so it is exothermic and $\Delta H$ is negative.
2
$4200 / 0.05 = 84\,000$ J/mol.
3
Divide by 1000.

A reaction absorbs 2500 J when 0.10 mol reacts. What is $\Delta H$ in kJ/mol?

Energy level diagrams

Exothermic profiles end lower than they start; endothermic profiles end higher.

An energy level diagram shows the energy of reactants and products. Exothermic: products lower, $\Delta H$ negative. Endothermic: products higher, $\Delta H$ positive. Breaking bonds needs energy (endothermic) and making bonds releases energy (exothermic). If more energy is released making bonds than used breaking them, the reaction is exothermic overall.

Energy level diagram for an exothermic reaction (illustrative)

$\Delta H$ energy to break bonds − energy released making bonds

Worked example

Use bond energies to find $\Delta H$ for $\ce{H2 + Cl2 -> 2HCl}$ . Bond energies (kJ/mol): H–H 436, Cl–Cl 243, H–Cl 432.

1
Energy to break bonds: 436 + 243 = 679 kJ.
2
Energy released making bonds: 2 x 432 = 864 kJ.
3
$\Delta H = 679 - 864$ .

In an endothermic reaction, which statement is correct?

Try an exam question

In an experiment, 25 g of water is heated from 18.0 °C to 30.0 °C by burning 0.010 mol of ethanol. Calculate the molar enthalpy change in kJ/mol. (c = 4.2 J/g °C)

[4 marks]

That's the notes covered.

Carry on to the next subtopic.