Exothermic and Endothermic Reactions

Every chemical reaction involves a transfer of energy to or from its surroundings, which is why a reaction mixture can grow noticeably hotter or colder as it proceeds. Reactions that give out energy are called exothermic; reactions that take in energy are called endothermic. Recognising which is which — and being able to back it up with a temperature change you have measured — is a core skill in Topic C3 of OCR Gateway Science A. This lesson explains the two types, looks at everyday examples, and works through the required practical in which you measure a temperature change.

By the end of this lesson you should be able to define exothermic and endothermic reactions, classify reactions from a temperature change, give everyday examples of each, and describe the temperature-change required practical.

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Energy and the Surroundings

When chemists talk about energy in reactions, they think of two parts: the reacting chemicals (the system) and everything around them — the solution, the container, the air — which is the surroundings.

• In an exothermic reaction, energy is transferred from the chemicals to the surroundings, so the temperature of the surroundings rises. (Exo = out.)
• In an endothermic reaction, energy is taken in from the surroundings by the chemicals, so the temperature of the surroundings falls. (Endo = in.)

So a simple temperature reading tells you the type: temperature up = exothermic; temperature down = endothermic. The energy involved is usually transferred as heat.

Exam Tip: Define the two types by the direction of energy transfer, then by the temperature change of the surroundings. "Exothermic transfers energy to the surroundings, so the temperature rises" is a complete, mark-worthy statement.

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Exothermic Reactions

In an exothermic reaction the products store less energy than the reactants, and the difference is released to the surroundings as heat (so the mixture warms up). Important examples to remember are:

• Combustion (burning a fuel), such as $\text{CH}_4 + 2\text{O}_2 \rightarrow \text{CO}_2 + 2\text{H}_2\text{O}$CH4​+2O2​→CO2​+2H2​O — this releases a great deal of energy and is the most familiar exothermic reaction.
• Neutralisation of an acid by an alkali, $\text{H}^+ + \text{OH}^- \rightarrow \text{H}_2\text{O}$H++OH−→H2​O, which warms the solution.
• Oxidation reactions, including the slow oxidation of iron and the reaction of reactive metals with oxygen.
• Many displacement and metal–acid reactions, which warm the mixture.

A practical use of an exothermic reaction is a hand warmer. One common type uses the slow oxidation of iron (iron + oxygen + water → hydrated iron oxide); the reaction releases heat steadily, warming your hands. Self-heating cans of food and drink work on the same principle.

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Endothermic Reactions

In an endothermic reaction the products store more energy than the reactants, so energy must be taken in from the surroundings (and the mixture cools down). Examples to remember are:

• Thermal decomposition, such as heating a metal carbonate: $\text{CaCO}_3 \rightarrow \text{CaO} + \text{CO}_2$CaCO3​→CaO+CO2​. Energy must be supplied continuously, which is why the substance only decomposes while it is being heated.
• The reaction between citric acid and sodium hydrogencarbonate (the fizzing reaction in some bath bombs and sherbet), which makes the mixture noticeably colder.
• Dissolving certain salts, such as ammonium nitrate, in water.

A practical use of an endothermic reaction (or process) is an instant cold pack for sports injuries. Squeezing the pack mixes a salt with water; as the salt dissolves it takes in energy, and the pack becomes cold enough to reduce swelling without needing a freezer.

	Exothermic	Endothermic
Energy transfer	Released to the surroundings	Taken in from the surroundings
Temperature of surroundings	Rises	Falls
Energy stored in products	Less than in reactants	More than in reactants
Examples	Combustion, neutralisation, oxidation	Thermal decomposition, citric acid + sodium hydrogencarbonate
Everyday use	Hand warmers, self-heating cans	Instant cold packs

Exam Tip: "Thermal decomposition is endothermic" is a fact examiners reward — a substance only decomposes while it is being heated because the reaction needs a continuous supply of energy.

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Required Practical: Measuring a Temperature Change

A key required practical for C3 is to investigate the temperature change in a reaction — for example, how the volume of acid added to an alkali (or the type of metal added to acid) affects the temperature change of the mixture.

Variables (using acid added to alkali as the example):

Variable	What it is
Independent (you change)	Volume of acid added to the alkali
Dependent (you measure)	Maximum temperature change of the mixture
Control (kept the same)	Concentration of acid and alkali, starting temperature, volume of alkali, same insulated cup

Method (numbered):

1. Use a measuring cylinder to put a fixed volume of alkali (for example $25 \text{ cm}^3$25 cm3 of sodium hydroxide solution) into an insulated polystyrene cup standing in a beaker for support.
2. Measure and record the starting temperature with a thermometer.
3. Add a measured volume of acid, stir, and record the highest temperature reached.
4. Repeat with different volumes of acid, keeping everything else the same.
5. Calculate the temperature change ($\Delta T$ΔT = highest temperature − starting temperature) for each, and plot a graph.

The polystyrene cup (often with a lid) is used because it is a good insulator, reducing heat loss to the surroundings so the temperature change you measure is closer to the true value.

Example readings (illustrative):

Volume of acid / $\text{cm}^3$cm3	Start temp / °C	Highest temp / °C	$\Delta T$ΔT / °C
5	20.0	23.5	3.5
10	20.0	27.0	7.0
15	20.0	28.0	8.0
20	20.0	28.0	8.0

Here the temperature change rises as more acid is added until all the alkali has reacted; after that, adding more acid does not raise the temperature further because there is no more alkali left to neutralise. (These figures are example readings, not data from a specific experiment.)

Exam Tip: The reason for the insulated cup and lid is to reduce heat loss to the surroundings. If asked to improve the method, suggesting a lid or repeating and taking a mean are reliable marks.

Worked Example — Reading a temperature change

From the table above, calculate the temperature change when $10 \text{ cm}^3$10 cm3 of acid is added, and state what it tells you.

Step 1 — read the values: start temperature $= 20.0 \text{ °C}$=20.0 °C; highest temperature $= 27.0 \text{ °C}$=27.0 °C.

Step 2 — calculate $\Delta T$ΔT (highest − start):

$\Delta T = 27.0 - 20.0 = 7.0 \text{ °C}$ΔT=27.0−20.0=7.0 °C

Step 3 — interpret: the temperature rose, so the reaction is exothermic. A larger $\Delta T$ΔT would mean more energy was released. Comparing rows, the temperature change keeps rising up to $15 \text{ cm}^3$15 cm3 and then stays at $8.0 \text{ °C}$8.0 °C, which shows that by $15 \text{ cm}^3$15 cm3 all the alkali has reacted, so adding more acid releases no further energy.