Exothermic and Endothermic Reactions

Every chemical reaction shifts energy between the chemicals and their surroundings, which is why a reaction mixture can turn noticeably hotter or colder as it goes. Reactions that give out energy are exothermic; reactions that take in energy are endothermic. Being able to tell which is which — and to back it up with a temperature change you have measured, or with a reaction profile — is a core skill in Topic C3 of OCR Gateway Combined Science. This lesson explains the two types, works through the temperature-change required practical, and then uses reaction profiles to show the energy of the chemicals as a reaction proceeds, including the activation energy and how a catalyst helps.

By the end of this lesson you should be able to define exothermic and endothermic reactions, classify a reaction from a temperature change, give everyday examples of each, describe the temperature-change required practical, and draw and interpret reaction profiles including activation energy and the effect of a catalyst.

This lesson builds AO1 recall of the exothermic/endothermic definitions, AO2 application through the temperature-change required practical and (Higher) the bond-energy calculation, and AO3 interpretation when you classify a reaction from data and read a reaction profile.

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Energy and the Surroundings

When chemists think about energy in a reaction, they split the world into 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 single temperature reading tells you the type: temperature up = exothermic; temperature down = endothermic. The energy is usually transferred as heat.

Exam Tip: Define each type by the direction of energy transfer first, then by the temperature change. "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 that difference is released to the surroundings as heat (so the mixture warms up). Key examples to remember:

• 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 of all.
• 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 noticeably.

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

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Endothermic Reactions

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

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

A practical use of an endothermic process is an instant cold pack for a sports injury. Squeezing the pack mixes a salt with water; as the salt dissolves it takes in energy, and the pack turns 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$25 cm3 of sodium hydroxide solution) into an insulated polystyrene cup standing in a beaker for support.
2. Measure and record the starting temperature with a thermometer.
3. Add a measured volume of acid, stir, and record the highest temperature reached.
4. Repeat with different volumes of acid, keeping everything else the same.
5. Work out the temperature change ($\Delta T$ΔT = highest temperature − starting temperature) for each, and plot a graph.

The polystyrene cup (usually with a lid) is used because it is a good insulator, cutting 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 grows 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 alkali left to neutralise. (These are illustrative figures, 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 means more energy was released. Comparing the rows, the temperature change keeps rising up to $15 \text{ cm}^3$15 cm3 and then holds 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.

Answer: $\Delta T = 7.0 \text{ °C}$ΔT=7.0 °C; the rise shows the reaction is exothermic.

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Why the Temperature Changes: A Particle View

It is worth knowing why a reaction gives out or takes in energy. During a reaction, the bonds in the reactants have to be broken and new bonds have to be made in the products. Breaking bonds always takes in energy, and making bonds always releases energy.

• If more energy is released making the new bonds than is used breaking the old ones, there is a net release of energy to the surroundings — the reaction is exothermic and the temperature rises.
• If more energy is taken in breaking the old bonds than is released making the new ones, there is a net intake of energy from the surroundings — the reaction is endothermic and the temperature falls.

So the temperature change you measure in the cup is just the balance between the energy needed to break bonds and the energy given out making them. This picture explains why combustion — which forms strong new bonds in carbon dioxide and water — is so strongly exothermic.

Exam Tip: A tidy one-line explanation: a reaction is exothermic when more energy is released making bonds than is used breaking bonds, and endothermic when the reverse is true.

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Reaction Profiles

A reaction profile (energy profile) turns these energy ideas into a diagram. It is a graph with energy on the y-axis and the progress of the reaction on the x-axis, and it shows three things:

• the energy of the reactants (a level on the left);
• the energy of the products (a level on the right);
• the activation energy — the "hump" the chemicals must climb over to react.

The overall energy change of the reaction is the difference in height between the reactant level and the product level.

Profile of an Exothermic Reaction

In an exothermic reaction the products end up at a lower energy than the reactants, so energy is released to the surroundings. The product level sits below the reactant level.