Last updated: 9 October 2026
A hand warmer becomes hot, while an instant cold pack becomes cold. Both involve changes in which energy is transferred, but the direction of that transfer is different. Chemical energetics gives us a way to follow this movement of energy and explain why some reactions warm their surroundings while others cool them.
The ideas in this chapter are closely connected. Exothermic and endothermic reactions describe the overall direction of energy transfer. Energy profile diagrams show that transfer visually, while bond breaking and bond making explain where the energy change comes from at the particle level.
Try these five questions before reading the chapter. They test the ideas you will need when learning about energy changes in reactions.
- 5 quick questions
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- Explanations for every option
1. Chemical reactions involve energy transfer
Every chemical reaction involves an energy change. To understand that change, it helps to distinguish between the reacting system and the surroundings.
The reacting system is the collection of chemicals taking part in the reaction. The surroundings include everything outside the reacting system that can gain or lose energy, such as the solution, container and surrounding air.
Follow the energy
There are two possible overall directions of energy transfer:
- Energy can move from the reacting system to the surroundings.
- Energy can move from the surroundings into the reacting system.
This simple distinction allows us to classify reactions as exothermic or endothermic.
2. Exothermic and endothermic reactions
Exothermic reactions release energy to the surroundings
An exothermic reaction transfers energy from the reacting system to the surroundings.
As the surroundings receive energy, their temperature usually rises. If the reaction takes place in a solution, for example, the solution and container may become warmer.
Examples include:
- combustion;
- neutralisation between an acid and an alkali;
- many reactions between metals and acids;
- respiration.
For example, methane burns in oxygen:
CH₄(g) + 2O₂(g) → CO₂(g) + 2H₂O(l)
Energy is released overall to the surroundings. The products therefore finish at a lower energy level than the reactants.
Endothermic reactions absorb energy from the surroundings
An endothermic reaction takes in energy from the surroundings.
Because the surroundings lose energy, their temperature may decrease.
Thermal decomposition is a common example. Calcium carbonate decomposes when it is heated strongly:
CaCO₃(s) → CaO(s) + CO₂(g)
Energy must be supplied for this reaction to occur. The products finish at a higher energy level than the reactant because energy has been absorbed by the reacting system.
Hot does not automatically mean exothermic
This is a common source of confusion. Calcium carbonate is heated strongly during thermal decomposition, so the test tube becomes extremely hot. Yet the reaction itself is endothermic.
Why? The heat comes from the external flame. Energy is moving from the flame and surroundings into the reacting chemicals.
Therefore, do not classify a reaction merely by asking whether the apparatus is hot or cold. Ask which direction the energy is being transferred.
| Feature | Exothermic | Endothermic |
|---|---|---|
| Direction of energy transfer | From reacting system to surroundings | From surroundings into reacting system |
| Usual effect on surroundings | Temperature rises | Temperature falls |
| Energy level of products | Lower than reactants | Higher than reactants |
| Sign of ΔH | Negative | Positive |
| Typical examples | Combustion, neutralisation | Thermal decomposition |
3. Enthalpy change, ΔH
The overall energy change during a chemical reaction is represented by the enthalpy change, ΔH.
At O-Level, a useful way to understand it is:
ΔH = energy of products − energy of reactants
The sign of ΔH tells us whether the reaction is exothermic or endothermic.
Why is ΔH negative for an exothermic reaction?
In an exothermic reaction, the products contain less energy than the reactants.
So:
lower energy − higher energy = negative value
Therefore:
ΔH < 0
The decrease in energy of the reacting system corresponds to energy that has been transferred to the surroundings.
Why is ΔH positive for an endothermic reaction?
In an endothermic reaction, the products contain more energy than the reactants.
So:
higher energy − lower energy = positive value
Therefore:
ΔH > 0
The extra energy stored in the products came from the surroundings.
Worked reasoning
Suppose a reaction has a ΔH value of −85 kJ mol⁻¹.
- The negative sign tells us that the products are lower in energy than the reactants.
- The reacting system has therefore lost energy overall.
- That energy has been transferred to the surroundings.
- The reaction is exothermic.
The minus sign does not mean that the reaction has “negative energy”. It tells us the direction of the overall energy change.
4. Energy profile diagrams
An energy profile diagram is a graph showing how the energy of a reacting system changes as reactants are converted into products.
The axes should be labelled:
- vertical axis: Energy;
- horizontal axis: Progress of reaction.
The horizontal axis is not normally labelled “time”. The diagram shows the route from reactants to products rather than how many seconds the reaction takes.
Energy profile for an exothermic reaction
In an exothermic energy profile:
- the reactants begin at a higher energy level;
- the curve rises to a maximum;
- the curve then falls;
- the products finish below the reactants;
- ΔH is negative.
The products are lower in energy because energy has been transferred from the reacting system to the surroundings.
Energy profile for an endothermic reaction
In an endothermic energy profile:
- the reactants begin at a lower energy level;
- the curve rises to a maximum;
- the products finish above the reactants;
- ΔH is positive.
The products are higher in energy because the reacting system has absorbed energy from the surroundings.
Where should ΔH be drawn?
ΔH represents the energy difference between the reactants and products.
Therefore, the ΔH arrow should connect:
- the reactant energy level; and
- the product energy level.
For an exothermic reaction, the arrow points downwards. For an endothermic reaction, it points upwards.
ΔH and Eₐ are not the same thing
Students sometimes draw ΔH from the reactant energy level all the way to the top of the curve. That is incorrect.
ΔH is measured between the reactant and product energy levels.
The arrow from the reactants to the peak represents activation energy, Eₐ.
5. Activation energy, Eₐ
Even when a reaction is exothermic overall, it does not necessarily begin immediately.
Reacting particles must first possess enough energy for a successful reaction to occur.
The activation energy, Eₐ, is the minimum energy that reacting particles must possess for a successful reaction to occur.
On an energy profile diagram, Eₐ is represented by the vertical energy difference between:
- the reactant energy level; and
- the top of the energy profile curve.
Why does an exothermic reaction sometimes need a flame?
Consider a fuel. Burning the fuel releases energy overall, but the fuel normally does not burst into flames spontaneously at room temperature.
A flame or spark supplies the initial energy needed for some reacting particles to overcome the activation-energy barrier.
Once the reaction begins, energy is released overall.
This is why saying “an exothermic reaction releases energy” does not mean that no energy is required to start it.
6. Bond breaking and bond making
We can now explain exothermic and endothermic reactions in terms of what happens to chemical bonds.
During a reaction:
- bonds in the reactants are broken;
- the atoms rearrange;
- new bonds are made to form the products.
Bond breaking and bond making have opposite energy effects.
Breaking bonds requires energy
A covalent bond exists because bonded atoms attract one another.
To separate those atoms, the attraction between them must be overcome. That requires an input of energy.
Therefore:
bond breaking is endothermic.
A useful analogy is pulling apart two strong magnets. You have to do work to separate them. Similarly, energy must be supplied to separate bonded atoms.
Making bonds releases energy
When atoms form a bond, they move into a more stable, lower-energy arrangement.
Energy is released during this process.
Therefore:
bond making is exothermic.
| Process | Energy change | Why? |
|---|---|---|
| Bond breaking | Energy absorbed | Energy is required to overcome the attraction between bonded atoms. |
| Bond making | Energy released | Atoms form a more stable, lower-energy arrangement. |
The overall reaction depends on the balance
Every reaction involves both energy absorption and energy release.
Energy is absorbed when reactant bonds are broken, while energy is released when product bonds are made.
The overall energy change depends on which amount is greater.
Exothermic versus endothermic at the bond level
If forming the new bonds releases more energy than breaking the old bonds absorbs, there is a net release of energy. The reaction is exothermic.
If breaking the old bonds absorbs more energy than forming the new bonds releases, there is a net absorption of energy. The reaction is endothermic.
Worked reasoning: combustion of hydrogen
Consider:
2H₂(g) + O₂(g) → 2H₂O(l)
First, energy is absorbed to break the H–H bonds in H₂ and the O=O bond in O₂.
Next, energy is released as O–H bonds are formed in the water molecules.
In this reaction, making the product bonds releases more energy than breaking the reactant bonds absorbs.
There is therefore a net release of energy, so the combustion of hydrogen is exothermic.
Do not decide by simply counting bonds
It is not enough to say, “More bonds are made than broken, so the reaction must be exothermic.”
Different covalent bonds involve different amounts of energy. What matters is the total energy absorbed during bond breaking compared with the total energy released during bond making.
7. Common mistakes in Chemical Energetics
| Common mistake | Why it is wrong | Correct idea |
|---|---|---|
| “Breaking bonds releases energy.” | Energy is required to overcome the attraction holding bonded atoms together. | Bond breaking absorbs energy. |
| “Making bonds needs energy.” | Formation of a bond places atoms in a more stable, lower-energy arrangement. | Bond making releases energy. |
| “An exothermic reaction does not need energy to start.” | Reacting particles may still need to overcome an activation-energy barrier. | An exothermic reaction can require an initial flame or spark. |
| “ΔH is the arrow from the reactants to the peak.” | The reactant-to-peak energy difference represents Eₐ. | ΔH is the difference between reactant and product energy levels. |
| “The horizontal axis of an energy profile is time.” | An energy profile does not show how long the reaction takes. | Use Progress of reaction. |
| “A hot reaction must be exothermic.” | The heat may have been supplied externally. | Identify the direction of energy transfer instead. |
8. Chemical Energetics checklist
Classifying reactions
- Ask which direction energy is transferred.
- Exothermic means energy leaves the reacting system.
- Endothermic means energy enters the reacting system.
- Use the sign of ΔH correctly.
Reading energy profiles
- Vertical axis: Energy.
- Horizontal axis: Progress of reaction.
- Products below reactants means exothermic.
- Products above reactants means endothermic.
Activation energy
- Eₐ is the minimum energy required for reaction.
- Measure Eₐ from the reactant level to the peak.
- Do not confuse Eₐ with ΔH.
Bond reasoning
- Breaking bonds absorbs energy.
- Making bonds releases energy.
- Compare the two energy amounts.
- The balance determines the overall reaction type.
Now use the ideas from the chapter to interpret energy changes, energy profiles and bond changes in unfamiliar situations.
- 10 application questions
- Exam-style reasoning
- Explanations for every option
Chemical Energetics FAQ
The main ideas are simple once the direction of energy transfer is clear, but a few questions often cause confusion.
Are heat and temperature the same thing?
No. Temperature tells us how hot or cold something is and is related to the average kinetic energy of its particles. Heat refers to energy being transferred because of a temperature difference.
During an exothermic reaction, energy transferred to the surroundings often causes their temperature to rise.
Can an exothermic reaction still need heating?
Yes. The reacting particles may need an initial supply of energy to overcome the activation-energy barrier.
Once the reaction proceeds, it can release more energy overall than was needed to start it.
Why does an exothermic energy profile rise before falling?
The rise represents the energy barrier that must first be overcome. Reacting particles need sufficient energy for a successful reaction to occur.
After the peak, formation of product bonds releases energy, and the products eventually finish at a lower energy level than the reactants.
Why is ΔH negative when an exothermic reaction releases energy?
The products of an exothermic reaction contain less energy than the reactants.
Since ΔH compares products with reactants, subtracting the higher reactant energy from the lower product energy gives a negative value.
Does breaking a molecule apart release energy?
Breaking the covalent bonds itself requires energy. The complete chemical reaction may still release energy overall if formation of the new product bonds releases a greater amount of energy.
This is why it is important to distinguish bond breaking from the overall energy change of the reaction.
Can an endothermic reaction happen at a high temperature?
Yes. Thermal decomposition reactions are often carried out at high temperatures because energy must continually be supplied to the reacting chemicals.
The fact that the reaction mixture is hot does not mean that the reaction is releasing energy.
Why is the horizontal axis of an energy profile not time?
An energy profile shows the energetic pathway followed as reactants become products. It does not tell us how long each part of that pathway takes.
The correct label is therefore Progress of reaction.
What is the quickest way to distinguish Eₐ from ΔH on a diagram?
Look at where the arrow finishes.
Eₐ runs from the reactant energy level to the peak.
ΔH runs between the reactant and product energy levels.
Chemical energetics is really one connected energy story. Reactant bonds must be broken, which absorbs energy, and new product bonds are formed, which releases energy. The balance between those changes gives the overall ΔH, while the activation energy explains why a reaction may still need an initial push. Once you learn to follow where the energy starts, where it ends and where it is transferred, exothermic and endothermic reactions become ideas you can reason through rather than definitions you have to memorise.