There are two ways that energy can change in a reaction: exothermic and endothermic. An exothermic reaction releases energy to its surroundings, which causes the temperature to rise. An endothermic reaction absorbs energy from its surroundings, lowering the temperature. We see these chemical reactions in combustion and instant cold packs. It is linked to the sign of the enthalpy change.
This article explains both chemical reactions, including how to recognise them from temperature change and simple equations, and provides detailed examples of their use in the real world. It is suitable for GCSE Chemistry students revising this topic.
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Endothermic and exothermic reactions
A chemistry experiment consists of two parts: the system, which includes the reacting chemicals, and the surroundings, such as the beaker, the air, and your hand. The surroundings will heat up when energy flows from the system (exothermic). When the energy moves from the surroundings to the system, the surroundings will get colder (endothermic).
This flow of energy shows the enthalpy change. Enthalpy is the total heat in the system, so the enthalpy change is the heat absorbed or released by a system during a process. Exothermic reactions release heat into their surroundings, and the products store less chemical energy than the reactants. Enothermic reactions absorb heat from their surroundings because the products store more energy than the reactants, which is released during the change.
Put simply, breaking bonds absorbs energy, while making bonds releases energy. An exothermic reaction is when more energy is released in bond-making than is absorbed during bond-breaking. The opposite is true of endothermic reactions.
The chemical equations for both reactions are below:
- Exothermic change | CH4 + 202 -> C02 + 2H20 + energy
- Endothermic change | Energy + CaC03 -> Ca0 + C02
The best indicator for the reaction type is temperature. A rise in your thermometer means an exothermic process, while a fall suggests an endothermic one. The reaction itself finishes lower than it started for exothermic reactions, and it finishes higher for endothermic reactions. Don't forget that both need energy to begin the process, and catalysts lower this activation energy without changing the enthalpy change.
These are a few examples of both reaction types:
- Exothermic examples | Combustion of fuels (e.g. a candle or the gas hob on your cooker), neutralisation reactions (e.g. hydrochloric acid with sodium hydroxide), and the oxidation used in iron-powder hand warmers.
- Endothermic examples | Thermal decomposition of limestone (CaC03 -> Ca0 + C02), the dissolution of ammonium nitrate used in instant cold packs, and the cooling effect of evaporation from the skin.
1
A solution is mixed in a polystyrene cup. The temperature falls from 22 °C to 16 °C. What does this mean?
Examples of Exothermic reactions
An exothermic reaction releases energy to the surroundings. When you're working in the lab, a rise in temperature happens when the solution, beaker and surrounding air gain energy. When new bonds are formed, more energy is released than when the old bonds are broken. This means the excess energy flows out as heat, often in the form of light, such as flames:
CH4 + 202 -> C02 + 2H20
Another example is neutralisation, which is a reaction between an acid and a base to form a salt and water. When you mix hydrochloric acid with sodium hydroxide, the solution heats up. This is because creating water from H+ and OH- releases energy:
HC1(aq) + Na0H(aq) -> NaC1(aq) + H20(1)
In the real world, you see this reaction with iron-powder hand warmers, which you may take with you when camping or hiking. The iron in the warmers reacts with oxygen to create iron(III) oxide. The oxidation releases energy, and the pouch heats your hands. The same concept is found with self-heating food cans - when you open the can, an exothermic reaction triggers in a separate chamber. The heat created is transferred into the meal.
If you are experimenting, you should be looking to see if your surroundings warm up from the energy that leaves the reactants. A simple test is to use a polystyrene cup and mix an exothermic solution to see if the temperature rises by a few degrees.
2
A student adds 50 cm³ of 1.0 mol dm-³ HCl to 50 cm³ of 1.0 mol dm-³ NaOH in a polystyrene cup. The temperature increases by 6 °C. Which statement is correct?
Examples of Endothermic reactions
An endothermic reaction absorbs energy from the surroundings. Essentially, the reading on the thermometer drops because the solution and container lose energy in the reaction. The products sit higher than the reactants, and the rise between them is the positive enthalpy change. More energy is needed to break the bonds than when new bonds are formed. This means the system pulls in energy.
A clear classroom example is thermal decomposition. This is when a compound is broken down into simpler substances when it's heated. For instance, heating limestone breaks the bonds in calcium carbonate. The heat is absorbed by the reaction and enables the decomposition:
CaC03(s) + energy -> Ca0(s) + C02(g)
Another example is dissolving ammonium nitrate in water. When the crystals disperse and ions become hydrated, energy is absorbed into the process. This cools the beaker used for the experiment.
A real-world example is instant cold packs, such as if you have an injury. When you squeeze the pack, the seal is broken and ammonium nitrate mixes with water. The energy is absorbed by the reaction, and the pack becomes cold.
With all of these examples, energy passes from the surroundings into the system, which cools down the surroundings. As a final laboratory example, you can add ammonium nitrate to water via a cup calorimetry setup. There will be a clear drop in temperature after a few minutes, showing the direction of the endothermic reaction.
3
A student dissolves 5.0 g of ammonium nitrate in 50 cm³ of water. The temperature falls from 20 °C to 12 °C. Which statement is accurate?
Measuring the temperature change
The easiest way to identify the type of reaction is by tracking the temperature change in the surroundings. Follow the steps below for this simple experiment:
- Use a polystyrene cup in a beaker for support. Add a thermometer. Record a starting temperature over a few minutes to ensure it is steady.
- Mix the reagents and stir gently.
- The reaction is exothermic if the reading climbs and levels off. This means the surroundings have gained energy. The reaction is endothermic if the temperature drops before stabilising. This means energy has been drawn into the system.
- The size of the temperature change shows how much energy was transferred into or out of the surroundings.
Be sure to note the chemicals, starting temperature, the highest/lowest temperature recorded, and the time the experiment took. The cup needs to be covered to avoid losing heat naturally to the room.
Neutralisation using acid and alkali normally shows a temperature rise. Dissolving ammonium nitrate usually drops the reading on the thermometer.
4
A student dissolves salt in water and records these temperatures: 21.0 °C (start), 18.4 °C (minimum), 18.6 °C (steady). What is the right conclusion?
Why the chemical reaction happens
The change in energy is due to the bonds breaking in the reactants and new bonds forming in the products. When a bond breaks, the particles are pulled apart, which absorbs energy. Creating a bond releases energy because the particles move to a stable arrangement.
An exothermic reaction means the total energy from forming bonds is higher than the energy created from breaking bonds. The excess energy flows into the surroundings.
An endothermic reaction is when the energy needed to break bonds is greater than bond-making gives back. The system has to draw energy in, which means the enthalpy change is positive.
We can see this balance on an energy diagram, also called a reaction profile. The reactants climb a "hill" called the activation energy. This is the energy needed for particles to collide, which leads to products. When they have peaked, the curve falls back to the energy level of the products.
If the product is below the reactants, the vertical drop equals a negative enthalpy change (exothermic). If the product sits above the reactants, the enthalpy change is positive (endothermic).
You can use a catalyst for an alternative route with a lower activation energy. This makes it easier for the particles to react, without changing the starting and finishing levels.
Combustion releases lots of energy by forming multiple strong bonds in CO2 and H2O. This creates more energy than the energy absorbed to break C–H and O=O bonds.
5
Why is methane combustion exothermic?
Further examples
We have listed some additional examples below to compare the different chemical reactions.
Dissolving salts
When ammonium nitrate dissolves in water, energy is required to separate ions from the crystal lattice and organise water molecules around them. Hydration releases energy, but not enough to compensate. Instead, the mixture absorbs energy and the beaker cools. This is the same endothermic change we see in instant cold packs.
By contrast, the hydration in the dissolving of calcium chloride needs more energy than that required to separate the ions. The extra energy flows into the surroundings and warms the solution. This is an exothermic process.
Crystallisation and melting
You can form a crystal from a supersaturated solution. This creates new bonds in the crystal and releases energy. A real-world example is crystallising sodium acetate in a reusable hand warmer. The pouch heats up, which means the crystallisation is exothermic.
The opposite is melting a solid. Energy is absorbed by the solid to break apart the structure. Meanwhile, the surroundings will cool unless heat is applied, making the process of melting endothermic.
Neutralisation and thermal decomposition
Water is formed when we mix a strong acid with a strong alkali. The new O-H bonds release energy, which makes the temperature rise. This is an exothermic reaction.
When we heat calcium carbonate, calcium oxide and carbon dioxide are created. This requires a constant source of energy to split the carbonate. This energy is absorbed, and the products sit at a higher enthalpy (endothermic reaction).
6
A reusable hand warmer is clicked. A supersaturated sodium acetate solution then crystalises, which heats up the pack. What statement is correct?
Conclusion - Endothermic and exothermic reactions
An exothermic reaction releases energy to the surroundings, which increases the temperature. Endothermic reactions absorb energy from the surroundings, which makes the temperature fall.
Think of these chemical reactions in terms of bond-breaking and bond-making, which explain both contrasting processes. It also helps with examples of each reaction, such as combustion and neutralisation (exothermic), and thermal decomposition or dissolving ammonium nitrate (endothermic).
For further reading, you can watch this video experimenting with exothermic reactions using everyday items you find around your house. You can also test yourself with past paper questions on exothermic and endothermic reactions by ExamPapersPractice.
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