Energy Changes in Reactions
You'll be able to classify a reaction as exothermic or endothermic from observations.
Why some reactions heat up and others go cold, and how to read energy changes from a reaction profile.
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You've probably used an instant cold pack on an injury at sport, it gets cold almost immediately after you crack it. You've also probably felt how warm it gets near a fire or a hot oven. What do you think is happening with energy in each of these two situations?
If a cold pack gets cold by absorbing heat, where exactly is that heat coming from? And if burning wood releases heat, where was that energy stored before the wood was lit? Think about what energy storage inside chemicals might look like.
● Know
- The definitions of exothermic and endothermic reactions
- Examples of each: combustion, neutralisation (exothermic); thermal decomposition, photosynthesis (endothermic)
- That energy can be thought of as a reactant or product conceptually
● Understand
- Why exothermic reactions release heat to the surroundings
- Why endothermic reactions absorb heat from the surroundings
- How energy changes relate to the breaking and forming of bonds conceptually
● Can do
- Classify a reaction as exothermic or endothermic from observations
- Explain everyday and industrial examples in terms of energy transfer
- Describe the direction of energy flow between reaction and surroundings
Crack an instant cold pack and hold it, the bag turns icy cold within seconds even though you are holding it in your warm hands, because the endothermic dissolution process absorbs heat energy from the surroundings. Chemical reactions involve energy changes: some release heat, others absorb it. Understanding these changes is essential for predicting reaction feasibility, designing safe processes, and harnessing chemical energy.
Exothermic reactions release more energy than they absorb. The products have less chemical energy than the reactants, and the excess is released as heat (and sometimes light). Examples: combustion, neutralisation, rusting, and cellular respiration. The surroundings end up warmer than they started.
Endothermic reactions absorb more energy than they release. The products have more chemical energy than the reactants, and the deficit is supplied by heat from the surroundings. Examples of endothermic reactions include thermal decomposition and photosynthesis. Some physical processes, such as dissolving certain salts in water, are also endothermic. The surroundings end up cooler than they started.
At the molecular level, energy changes arise from bond breaking and forming. Breaking bonds always requires energy input. Forming bonds always releases energy. Whether a reaction is exothermic or endothermic depends on the balance: if more energy is released forming new bonds than was required to break old bonds, the reaction is exothermic.
Text description
- Exothermic (ΔH negative): products are lower in energy than reactants; energy is released to the surroundings. The ΔH arrow points down.
- Endothermic (ΔH positive): products are higher in energy than reactants; energy is absorbed from the surroundings. The ΔH arrow points up.
- Both have an activation energy Ea: the height of the peak above the reactants.
A hot pack contains iron powder, salt, water, and vermiculite. When activated, iron oxidises (rusts) in an exothermic reaction: 4Fe + 3O2 → 2Fe2O3. The reaction releases heat slowly over several hours as oxygen diffuses through the pack. A cold pack contains ammonium nitrate and a separate water pouch. When squeezed, the salt dissolves endothermically: NH₄NO₃(s) → NH₄⁺(aq) + NO₃⁻(aq). The ions become dispersed in water; no new chemical substances are formed. Dissolution absorbs heat from the surroundings, making the pack feel cold. Hot packs and cold packs illustrate energy changes in different chemical or physical processes.
Australian energy chemistry: Australian researchers at CSIRO and universities are developing new energy storage systems based on reversible chemical reactions. One approach uses metal hydrides that release hydrogen endothermically when heated, then reabsorb hydrogen exothermically when cooled. Another uses molten salt batteries where charge and discharge involve redox reactions that release and absorb large amounts of energy. These technologies could store renewable energy more efficiently than conventional batteries.
Exothermic reactions start by themselves; endothermic ones have to be forced. This is false, and it runs together two questions that are not the same. How much energy a reaction gives out is one question. Whether it gets going at all is another, and that is about activation energy. Petrol releases an enormous amount of energy when it burns, and a full tank sits in a car all week doing nothing, because nothing has supplied the push needed to start it. Meanwhile an instant cold pack absorbs heat and goes off the second you squeeze it. So "gives out energy" does not mean "starts easily", and "takes in energy" does not mean "needs forcing" — and neither of them tells you how fast the reaction will be once it is under way. That is a third question again, and it is the one lessons 11 to 14 are about.
Classify each process as exothermic or endothermic.
Some reactions can be thermodynamically favourable yet proceed extremely slowly at room temperature. Petrol vapour and oxygen, for example, do not ignite under ordinary conditions without an ignition source because there is an activation-energy barrier. Activation energy (Ea) is the minimum energy barrier that reacting particles must overcome for a successful reaction.
An energy profile diagram shows how the energy of the system changes as the reaction proceeds. The reactants sit at some energy level. To reach the products, the system must first climb to a peak (the transition state) representing the activation energy barrier. Only molecules with sufficient kinetic energy can surmount this barrier.
Catalysts provide an alternative reaction pathway with lower activation energy. They do not change how much energy the reaction gives out or takes in overall, but they make it faster by allowing more molecules to react at a given temperature. Catalysts are not consumed in the reaction.
Enzymes in living organisms are biological catalysts. They enable reactions that would be impossibly slow at body temperature, including digestion, DNA replication, and cellular respiration.
Hydrogen peroxide decomposes slowly at room temperature: 2H2O2 → 2H2O + O2. The activation energy is high enough that a bottle of 3% hydrogen peroxide lasts months before significant decomposition. Adding a small amount of manganese dioxide dramatically speeds up the reaction - bubbles of oxygen form vigorously. The manganese dioxide is unchanged chemically and can be recovered by filtration. This demonstrates catalysis: same reaction, same energy change, but much faster due to lower activation energy.
Australian catalysis research: The Australian Research Council Centre of Excellence for Electromaterials Science develops catalysts for fuel cells, water splitting, and carbon dioxide reduction. One major project at Monash University involves single-atom catalysts where individual metal atoms dispersed on carbon supports catalyse reactions with extraordinary efficiency. These materials could make hydrogen production from water electrolysis economically competitive with fossil fuels.
Catalysts are used up in reactions and must be replenished. This is false. By definition, catalysts are not consumed. They may be poisoned or deactivated by side reactions, but the catalytic mechanism itself leaves the catalyst unchanged. In industrial practice, catalysts do degrade over time due to contamination, sintering (particles clumping together), or structural changes, so they are periodically replaced or regenerated. But this degradation is not part of the catalytic cycle.
Every chemical reaction does two things, and they pull in opposite directions. Old bonds in the reactants have to be pulled apart, which costs energy. New bonds in the products then snap together, which releases it. Nothing else is going on; the whole of exothermic and endothermic is which of those two is bigger.
Burn methane in a gas heater and you break the bonds inside CH₄ and O₂, which takes energy in. Then carbon dioxide and water form, and the bonds in those products are stronger — they release more than the breaking cost. The leftover has nowhere to go but out, into the room, as heat. Heat a lump of limestone in a kiln and the balance runs the other way: the bonds you have to break are stronger than the ones that form, so the reaction cannot pay for itself and stops the moment you stop supplying heat.
That is the difference between the two words. An exothermic reaction ends up with its atoms held together more tightly than they started, and hands the difference to its surroundings. An endothermic one ends up held together more loosely, and has to be paid for from outside.
The rule that carries: ask whether the products are held together more tightly than the reactants were. If yes, energy comes out and the surroundings warm up. If no, energy has to go in and the surroundings cool down.
Test it on every kind of reaction this unit has taught, because a rule you have not tried on the awkward cases is not yet a rule:
Combustion of methane — products more tightly held, heat out, and your hands feel it. Neutralisation — H+ and OH- lock into water, one of the tightest arrangements available, so the beaker warms. Thermal decomposition of limestone — you are prising a tightly held solid apart into looser pieces, so it takes heat in and the kiln has to keep supplying it. Photosynthesis — the same direction, paid for by sunlight instead of a kiln.
An instant cold pack also illustrates energy changes, through a physical process rather than a chemical reaction. As ammonium nitrate dissolves, separating its ionic lattice takes energy, while hydration of the ions releases energy. For ammonium nitrate, the energy absorbed overall is greater than the energy released, so dissolution is endothermic and the pack cools. This energy accounting explains the temperature change without classifying dissolution as a chemical reaction.
Why the energy balance is worth money: Australian coals, natural gas and biofuels are compared by how much heat each releases per kilogram burned. The measurement is the same idea as a school experiment — burn a known amount under a container of water and see how far the water temperature climbs — done carefully enough to be trusted. Those figures decide which fuels a power station buys and feed the national greenhouse-gas accounts, because how much carbon dioxide you emit for a given amount of heat depends on which fuel you burned to get it.
Breaking bonds releases energy. This is the opposite of the truth. Breaking bonds always requires energy input. Forming bonds releases energy. Confusing these directions is one of the most common errors in thermochemistry. Remember: you must put energy in to break something apart (like tearing paper), and energy comes out when something forms (like gravity pulling objects together).
Click each stage of an exothermic reaction energy profile.
Reactants
Molecules have a certain amount of chemical energy stored in their bonds.
Activation energy
Energy input breaks bonds and forms the transition state.
Products
New bonds form, releasing energy. Products have less energy than reactants.
Heat release
The excess energy is released to the surroundings as heat.
Wrong: "All chemical reactions produce heat." No, many reactions absorb heat. Photosynthesis and thermal decomposition are endothermic chemical reactions; dissolving ammonium nitrate in a cold pack is an endothermic physical process.
Right: Many reactions absorb heat (endothermic) rather than releasing it. Photosynthesis and thermal decomposition are endothermic reactions. Dissolving ammonium nitrate is an endothermic physical process.
Wrong: "Endothermic reactions are rare." No, endothermic reactions are common in nature. Photosynthesis, which powers almost all life on Earth, is endothermic.
Right: Endothermic reactions are very common, photosynthesis and some cooking processes are endothermic reactions. Dissolving many salts, including ammonium nitrate in a cold pack, is an endothermic physical process. These processes absorb heat from their surroundings.
Wrong: "A cold feeling proves that a change is physical, not chemical." Temperature change alone cannot classify a change. A cold pack feels cold because ammonium nitrate dissolves endothermically, but chemical reactions can also be endothermic.
Right: A cold feeling shows that a process is absorbing heat from its surroundings, but it does not show whether the change is chemical. In an ammonium-nitrate cold pack, endothermic dissolution is a physical process: the salt separates into its aqueous ions, with no new substance formed.
Bushfires, Sport and Industry
Australian bushfires are among the most dramatic exothermic events on Earth. When eucalyptus leaves burn, the combustion reaction releases enormous heat, enough to create firestorms that generate their own weather. CSIRO researchers study these reactions to predict fire behaviour and protect communities.
At the other end of the temperature scale, Australian sports teams and medical services use instant cold packs for injuries. These packs contain ammonium nitrate and water separated by a barrier. When the barrier is broken, the ammonium nitrate dissolves endothermically, absorbing heat from the pack and its surroundings. This is a physical process rather than a chemical reaction. From bushfires to sports first-aid packs, chemical reactions and physical processes can transfer energy to or from their surroundings.
✍ Copy Into Your Books
▾Exothermic Reactions
- Release heat to surroundings
- Energy is a product conceptually
- Examples: combustion, neutralisation, hand warmers
- Surroundings get warmer
Endothermic Reactions
- Absorb heat from surroundings
- Energy is a reactant conceptually
- Examples: thermal decomposition and photosynthesis; cold packs use endothermic dissolution
- Surroundings get cooler
Energy Flow
- Exothermic: system → surroundings (heat out)
- Endothermic: surroundings → system (heat in)
- Conservation of energy applies
Classify by Energy Change
Energy in Everyday Life
At the start of this lesson, the hook compared endothermic ammonium-nitrate dissolution in a cold pack with exothermic iron oxidation in a hand warmer.
Now that you understand exothermic and endothermic reactions and energy diagrams, can you explain what is happening at the particle level inside each of those products? How has your original idea about where the "cold" or "heat" comes from changed?
Q1. 1. Define exothermic and endothermic reactions. Give one example of each and explain how you could tell which is which by observation. 4 MARKS
Q2. 2. Combustion and photosynthesis can be described as opposite reactions in terms of energy. Explain this statement with reference to whether energy is a reactant or product in each case. 4 MARKS
Q3. 3. A sports trainer uses an instant cold pack to treat an ankle injury. When the inner pouch is broken, ammonium nitrate dissolves in water and the pack becomes cold. Classify the dissolution as exothermic or endothermic and describe the direction of energy transfer between the process and its surroundings. State whether dissolving the salt forms a new substance.
Revisit Your Thinking
Go back to your Think First answer. Has your understanding changed?
- Can you now explain where the heat in a hand warmer comes from?
- Can you explain where the 'cold' in a cold pack goes?
Model answers (click to reveal)
Answers
▾MCQ 1
B An exothermic reaction releases heat energy to the surroundings.
MCQ 2
C Photosynthesis is endothermic because it absorbs light energy from the sun to convert carbon dioxide and water into glucose.
MCQ 3
B A reaction that feels cold is absorbing heat energy from the surroundings (including your skin). This is characteristic of an endothermic reaction.
MCQ 4
A Both combustion and neutralisation are exothermic reactions. They release heat to the surroundings. Choice C includes endothermic dissolution in a cold pack and endothermic photosynthesis; the first is a physical process and the second is a chemical reaction.
MCQ 5
C In an endothermic reaction, energy is absorbed from the surroundings. Conceptually, energy acts as a reactant that must be supplied for the reaction to proceed.
Short Answer 1
Model answer: An exothermic reaction is a chemical reaction that releases heat energy to the surroundings. For example, combustion of wood is exothermic because the surroundings get hot and light is produced. An endothermic reaction is a chemical reaction that absorbs heat energy from the surroundings. For example, photosynthesis is endothermic because it requires light energy from the sun. You can tell them apart by touch: exothermic reactions make the container feel warm, while endothermic reactions make it feel cold.
Short Answer 2
Model answer: Combustion and photosynthesis are opposite reactions in terms of energy. Combustion is exothermic, meaning energy is released as a product, the reaction gives out heat and light. Conceptually: fuel + oxygen → carbon dioxide + water + energy. Photosynthesis is endothermic, meaning energy is required as a reactant, the reaction absorbs light energy. Conceptually: carbon dioxide + water + energy → glucose + oxygen. Together, they form a cycle: photosynthesis stores solar energy in glucose, and respiration or combustion releases that stored energy.
Short Answer 3
Model answer: The cold pack uses endothermic dissolution: when ammonium nitrate dissolves in water, the solid separates into NH₄⁺(aq) and NO₃⁻(aq) ions that become surrounded by water molecules. This physical process absorbs heat from the pack and its surroundings, so their temperature falls. The ions are the dissolved form of the original salt; the process does not form new substances. Endothermic describes the direction of energy transfer and does not by itself show that a chemical reaction occurred.