Energetics

Energy is on the move

Every chemical reaction involves a change in energy. When bonds rearrange, energy is either released to the surroundings or taken in from them. The surroundings usually means the water, air or apparatus around the reaction — and the easiest way to detect the energy change is with a thermometer.

Reactions fall into two camps: exothermic (energy out, things get hotter) and endothermic (energy in, things get colder).

Key terms

Exothermic reaction — transfers energy to the surroundings; the temperature of the surroundings rises.

Endothermic reaction — takes in energy from the surroundings; the temperature of the surroundings falls.

Surroundings — everything around the reacting chemicals (often the solvent water).

Exothermic reactions

In an exothermic reaction, energy is given out. If you carry it out in solution, the thermometer reading goes up.

Common exothermic changes:

Combustion — burning fuels such as methane releases a lot of heat.
Neutralisation — acid + alkali always warms the mixture.
Many displacement and oxidation reactions, including respiration.

Endothermic reactions

In an endothermic reaction, energy is taken in, so the thermometer reading goes down.

Common endothermic changes:

Thermal decomposition — e.g. heating calcium carbonate to break it into calcium oxide and carbon dioxide. It needs a constant supply of energy.
Dissolving some salts — e.g. ammonium nitrate or ammonium chloride in water makes the solution feel cold. (Sports injury "cold packs" use this.)

Real world

Self-heating coffee cans use an exothermic reaction (often quicklime + water) to warm the drink. Instant cold packs use the endothermic dissolving of ammonium nitrate to chill a sprain. Same chemistry, opposite direction.

Reaction profiles (energy level diagrams)

A reaction profile shows how the energy stored in the chemicals changes as reactants turn into products. The vertical axis is energy; the horizontal axis is the progress of the reaction.

Two quantities always appear:

Activation energy ($E_a$) — the minimum energy reactants need to collide successfully and start the reaction. It is the size of the "hill" from reactants up to the peak.
Overall energy change ($\Delta H$) — the difference in energy between products and reactants.

For an exothermic reaction, products sit lower than reactants, so energy has been released: $\Delta H$ is negative.

For an endothermic reaction, products sit higher than reactants, so energy has been absorbed: $\Delta H$ is positive.

Exam tip

When you draw a profile, label four things or lose marks: the reactant level, the product level, the activation energy (peak above reactants), and an arrow for $\Delta H$ between the two levels. Always include the curved "hump".

Bonds, breaking and making

Energy changes come from the bonds in the chemicals.

Breaking bonds is endothermic — it takes in energy to pull atoms apart.
Making bonds is exothermic — it gives out energy when new bonds form.

Whether the overall reaction is exothermic or endothermic depends on the balance:

If more energy is released making new bonds than was taken in breaking old bonds → reaction is exothermic ($\Delta H$ negative).
If less energy is released than was taken in → reaction is endothermic ($\Delta H$ positive).

Watch out

A common slip is to write "breaking bonds gives out energy". It does not. Breaking takes energy in; making gives energy out. Remember: Breaking = Bank takes your energy.

Calculating ΔH from bond energies

A bond energy is the energy needed to break one mole of a particular bond. The rule is:

$$\Delta H = \text{(sum of bonds broken)} - \text{(sum of bonds made)}$$

A negative answer = exothermic; a positive answer = endothermic.

Worked example

Find $\Delta H$ for the combustion of hydrogen: $2H_2 + O_2 \rightarrow 2H_2O$

Measuring an energy change: simple calorimetry

You can measure the heat released by a reaction by transferring it to water and recording the temperature change.

A basic experiment to compare the energy released by burning fuels:

1. Measure a known volume (e.g. 100 cm³) of water into a metal can and record its starting temperature.
2. Weigh a small spirit burner of fuel.
3. Light the burner under the can and stir gently.
4. After a set temperature rise (e.g. 20 °C), put out the flame and record the final temperature.
5. Reweigh the burner to find the mass of fuel burned.

The temperature rise tells you how much energy went into the water. More energy released = a bigger rise.

Exam tip

Calorimetry results are always lower than the true value because heat is lost to the air and to the apparatus, and some fuel may not burn completely. To improve it: shield the apparatus from draughts, use a lid, and place the flame close to the can.

Quick recap

Exothermic = energy out, temperature rises, $\Delta H$ negative, products lower on the profile.
Endothermic = energy in, temperature falls, $\Delta H$ positive, products higher on the profile.
Breaking bonds takes energy in; making bonds gives energy out.
$\Delta H = \text{bonds broken} - \text{bonds made}$.