Energy Changes in Reactions
In 2023, NSW Ambulance used chemical instant cold packs at over 850,000 callouts, each one drops to −5°C in 10 seconds using endothermic chemistry.
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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 chemical reaction inside is pulling heat energy out of 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 enthalpy change (delta H) is negative.
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: thermal decomposition, photosynthesis, and dissolving some salts in water. The enthalpy change (delta H) is positive.
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.
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 water dissolves the ammonium nitrate in an endothermic process: NH4NO3(s) -> NH4+(aq) + NO3-(aq). The dissolution absorbs heat from the surroundings, making the pack feel cold. Both packs exploit the same principle: energy changes accompanying 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 with large enthalpy changes. These technologies could store renewable energy more efficiently than conventional batteries.
Exothermic reactions happen spontaneously; endothermic reactions do not. This is false. Spontaneity depends on Gibbs free energy (delta G), not just enthalpy. Endothermic reactions can be spontaneous if the entropy increase is large enough. Dissolving ammonium nitrate in water is spontaneous but endothermic. Many spontaneous processes in nature are endothermic, including evaporation and photosynthesis. Enthalpy is one factor in spontaneity, but not the only one.
Classify each process as exothermic or endothermic.
Even exothermic reactions do not occur spontaneously at room temperature. Petrol and oxygen can coexist indefinitely without reacting unless ignited. This is because reactions require activation energy (Ea) - the minimum energy needed to break bonds in reactants and initiate the 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 the overall energy change (delta H) or the equilibrium position, but they make the reaction 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.
Quantitative understanding of reaction energetics uses bond energies - the energy required to break one mole of a specific bond. Bond energies are always positive (energy is required to break bonds). When bonds form, the same amount of energy is released.
Common bond energies: H-H = 436 kJ/mol; O=O = 498 kJ/mol; C-H = 413 kJ/mol; O-H = 463 kJ/mol; C=O = 745 kJ/mol.
To estimate the enthalpy change of a reaction:
delta H = (sum of bond energies broken) - (sum of bond energies formed)
If more energy is released forming bonds than was required to break bonds, delta H is negative (exothermic). If less energy is released, delta H is positive (endothermic).
Calorimetry measures actual heat changes experimentally. A simple calorimeter is an insulated container with water. The reaction occurs in or near the water, and the temperature change is measured. Using Q = mcT, the heat absorbed or released by the water is calculated, which equals the heat change of the reaction (assuming no heat loss).
Estimate the enthalpy change for hydrogen combustion: 2H2 + O2 -> 2H2O.
Bonds broken: 2 H-H (2 * 436 = 872 kJ) + 1 O=O (498 kJ) = 1,370 kJ.
Bonds formed: 4 O-H (4 * 463 = 1,852 kJ).
delta H = 1,370 - 1,852 = -482 kJ per 2 moles of H2O, or -241 kJ/mol.
The experimental value is -286 kJ/mol. The estimate is close but not exact because bond energies are average values that vary slightly depending on molecular environment. This calculation illustrates the power and limitations of bond energy estimation.
Australian thermochemistry: The Combustion and Energy Research Centre at the University of Adelaide measures precise enthalpies of combustion for Australian coals, natural gas, and biofuels. These data inform power station efficiency calculations and greenhouse gas accounting. Accurate calorimetry is essential for Australia national greenhouse gas inventory, which reports emissions to the United Nations Framework Convention on Climate Change.
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, thermal decomposition and cold packs are all endothermic.
Right: Many reactions absorb heat (endothermic) rather than releasing it. Photosynthesis, dissolving ammonium nitrate, and thermal decomposition are all everyday examples of endothermic reactions.
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, cooking, dissolving many salts, and cold packs are all endothermic. They are simply less dramatic than exothermic reactions because they don't produce visible heat or light.
Wrong: "If a reaction feels cold, it is not a chemical reaction." No � a cold feeling simply means the reaction is absorbing heat from your skin. It is still a chemical reaction with new substances forming.
Right: A cold feeling is simply evidence of an endothermic reaction absorbing heat from the surroundings (including your hand). New substances are still forming, the reaction is chemical, not physical.
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, drawing heat from the injured area and reducing swelling. From the Outback to the operating theatre, exothermic and endothermic reactions are part of Australian life.
✍ 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, photosynthesis, cold packs
- 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 described an instant cold pack dropping to –5°C in seconds without any electricity, and a hand warmer running the opposite reaction to keep rescuers warm in the Snowy Mountains.
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, two chemicals mix and the pack becomes cold. Explain whether this is an exothermic or endothermic reaction, and describe the energy transfer between the reaction and the surroundings. 4 MARKS
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
BAn exothermic reaction releases heat energy to the surroundings.
MCQ 2
CPhotosynthesis is endothermic because it absorbs light energy from the sun to convert carbon dioxide and water into glucose.
MCQ 3
BA reaction that feels cold is absorbing heat energy from the surroundings (including your skin). This is characteristic of an endothermic reaction.
MCQ 4
ABoth combustion and neutralisation are exothermic reactions. They release heat to the surroundings.
MCQ 5
CIn 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 instant cold pack involves an endothermic reaction. When the chemicals mix, the reaction absorbs heat energy from the surroundings. This energy transfer goes from the surroundings (the injured ankle, the air, the pack itself) into the chemical system. As heat is removed from the surroundings, the temperature drops and the pack feels cold. This reduces blood flow and swelling at the injury site. The reaction is endothermic because energy must be supplied from outside for the reaction to proceed.