The difference between an exothermic and an endothermic reaction is the direction in which thermal energy is transferred. An exothermic reaction transfers heat from the reacting system to the surroundings, whereas an endothermic reaction absorbs heat from the surroundings into the system.
This difference determines the sign of the enthalpy change, the relative enthalpies of reactants and products, and the temperature change normally observed in a calorimetry experiment. In the current IB Chemistry course, these ideas form part of Reactivity 1.1: Measuring enthalpy changes and provide the foundation for later work on Hess's law, bond enthalpies, fuels, entropy, and spontaneity.
Exothermic vs endothermic reactions at a glance
FeatureExothermic reactionEndothermic reactionDirection of heat transferSystem to surroundingsSurroundings to systemEffect on surroundingsSurroundings usually become warmerSurroundings usually become coolerSign of enthalpy changeΔH < 0ΔH > 0Relative enthalpiesProducts have lower enthalpy than reactantsProducts have higher enthalpy than reactantsEnergy-profile directionOverall downwardOverall upwardBond-energy balanceBond formation releases more energy than bond breaking absorbsBond breaking absorbs more energy than bond formation releasesCommon examplesCombustion, acid-base neutralization, respirationThermal decomposition, photosynthesis, some dissolving processes
The central memory rule is:
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Exothermic: energy exits the system, so ΔH is negative.
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Endothermic: energy enters the system, so ΔH is positive.
For broader topic coverage, see IB Chemistry Energetics Explained for Exams. This article concentrates specifically on the exothermic-endothermic distinction rather than duplicating calorimetry, Hess cycles, fuels, and Gibbs energy in full.
What do system and surroundings mean?
The system is the chemical reaction or physical process being studied. The surroundings are everything outside that system that can exchange energy with it, including the solution, reaction vessel, thermometer, air, and calorimeter.
This distinction matters because students sometimes describe energy transfer without stating where the energy goes. Saying that an exothermic reaction “produces heat” is less precise than saying that it transfers thermal energy from the system to the surroundings.
Heat is not a substance stored inside the reactants. Heat is energy transferred because of a temperature difference. Enthalpy, symbolized by H, is a thermodynamic state function, and the measurable quantity in a reaction is normally the change in enthalpy:
ΔH = H(products) − H(reactants)
At constant pressure, the enthalpy change corresponds to the heat transferred to the system. The IUPAC definition of enthalpy states that enthalpy is internal energy plus the pressure-volume product, and that its change equals heat brought into the system at constant pressure.
What happens in an exothermic reaction?
An exothermic reaction transfers thermal energy from the system to its surroundings. The system therefore loses enthalpy, giving a negative enthalpy change:
H(products) < H(reactants)
ΔH < 0
The negative sign does not mean that the amount of energy is physically negative. It records the direction of transfer according to the thermodynamic sign convention: energy has left the system.
The surroundings usually become warmer because they absorb the energy released by the reaction. In a school calorimetry experiment, the solution temperature therefore normally rises. This observed temperature increase is evidence of an exothermic process, but the formal definition concerns energy transfer, not temperature alone.
Real-world examples of exothermic reactions
Combustion is exothermic. When methane burns, energy is transferred to the surroundings as heat and light:
CH₄(g) + 2O₂(g) → CO₂(g) + 2H₂O(l)
The products have lower enthalpy than the reactants, so ΔH is negative. Fuels are useful precisely because their oxidation can transfer substantial energy to the surroundings.
Acid-base neutralization is another important example. When hydrogen ions and hydroxide ions form water, the solution normally warms:
H⁺(aq) + OH⁻(aq) → H₂O(l)
Cellular respiration is also overall exothermic, although cells capture part of the released energy through coupled biochemical processes rather than allowing all of it to appear immediately as heat. For a focused explanation, see Exothermic Reactions Explained for IB Chemistry.
What happens in an endothermic reaction?
An endothermic reaction absorbs thermal energy from its surroundings. The system gains enthalpy, so the products have higher enthalpy than the reactants:
H(products) > H(reactants)
ΔH > 0
Because energy enters the system, the sign of ΔH is positive. In a simple calorimetry experiment, the solution or other surroundings lose energy and their temperature normally decreases.
Real-world examples of endothermic reactions
The thermal decomposition of calcium carbonate requires a continuous energy input:
CaCO₃(s) → CaO(s) + CO₂(g)
This reaction is used in lime and cement production. Heating supplies the energy required to convert calcium carbonate into calcium oxide and carbon dioxide.
Photosynthesis is an overall energy-absorbing process powered by light:
6CO₂(g) + 6H₂O(l) → C₆H₁₂O₆(s) + 6O₂(g)
Some instant cold packs use the endothermic dissolution of a salt such as ammonium nitrate in water. Strictly, dissolution is a physical or solution process, not necessarily a chemical reaction because no new covalent substance must be formed. It is still correctly described as endothermic when the overall process absorbs heat from the surroundings.
This distinction is useful in IB answers: exothermic and endothermic can describe both chemical reactions and physical processes. Melting and vaporization, for example, are endothermic processes even though they are phase changes rather than reactions. See Endothermic Reactions Explained for further examples.
Why do reactions release or absorb energy?
Chemical reactions involve changes in bonding and intermolecular interactions. Two principles help explain the overall enthalpy change:
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Breaking bonds requires energy. It is an endothermic step.
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Forming bonds releases energy. It is an exothermic step.
A reaction is exothermic when forming the new bonds and interactions releases more energy than is required to break the original ones. It is endothermic when breaking the original bonds and interactions requires more energy than is released during product formation.
For reactions involving covalent substances in the gas phase, this idea can be represented approximately using mean bond enthalpies:
ΔH ≈ Σ(bond enthalpies of bonds broken) − Σ(bond enthalpies of bonds formed)
If the result is negative, the reaction is exothermic. If it is positive, the reaction is endothermic.
This explanation must be phrased carefully. It is incorrect to say that “breaking bonds releases energy.” Breaking an attractive bond always requires an energy input. Energy is released when new bonds form, and the overall balance determines the sign of ΔH. The RevisionDojo explanation of why reactions release different amounts of energy develops this comparison further.
How energy-profile diagrams show the difference
An energy-profile diagram plots the enthalpy of the reacting system against reaction progress. The vertical positions of the reactants and products show the overall enthalpy change, while the peak represents the high-energy transition region that must be reached for reaction to occur.
Diagram featureExothermic profileEndothermic profileReactant levelHigher than productsLower than productsProduct levelLower than reactantsHigher than reactantsΔH arrowPoints downwardPoints upwardSign of ΔHNegativePositiveForward activation energyReactant level to peakReactant level to peak
The activation energy, Eₐ, is not the same as ΔH. Activation energy is the minimum energy barrier associated with starting the reaction, whereas ΔH is the difference between the enthalpies of products and reactants.
Both exothermic and endothermic reactions need an activation energy. Combustion is exothermic, for example, but a fuel may still require a spark before it burns. A catalyst lowers the activation energy by providing an alternative pathway, but it does not change ΔH or turn an endothermic reaction into an exothermic one.
When sketching an IB energy profile, label:
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the vertical axis as enthalpy or potential energy
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the horizontal axis as reaction progress or reaction coordinate
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the reactants and products
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the forward activation energy, Eₐ
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the enthalpy change, ΔH, with the correct direction
How calorimetry identifies energy transfer
In solution calorimetry, the reaction is treated as the system and the solution is treated as part of the surroundings. The heat absorbed or released by the solution is calculated from:
q = mcΔT
where:
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q is the heat absorbed by the measured substance in J
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m is its mass in g
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c is its specific heat capacity in J g⁻¹ K⁻¹
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ΔT = T(final) − T(initial)
Energy conservation gives:
q(reaction) = −q(surroundings)
The negative sign is essential. If the solution warms, its ΔT and calculated q are positive because the solution gained energy. The reaction therefore has a negative q and is exothermic. The University of Calgary calorimetry explanation confirms this system-surroundings relationship.
Worked example
Suppose a reaction heats 100.0 g of solution from 20.0°C to 25.0°C. Assume the solution has a specific heat capacity of 4.18 J g⁻¹ K⁻¹.
ΔT = 25.0 − 20.0 = +5.0 K
q(solution) = 100.0 × 4.18 × 5.0 = +2090 J
The solution gained 2090 J, so:
q(reaction) = −2090 J = −2.09 kJ
If 0.0250 mol reacted, the molar enthalpy change would be:
ΔH = −2.09 ÷ 0.0250 = −83.6 kJ mol⁻¹
The negative sign identifies the reaction as exothermic. Further explanation of the notation is available in Enthalpy Change Explained for IB Chemistry.
Common IB Chemistry misconceptions
Temperature is evidence, not the definition
A temperature rise in the surroundings usually indicates an exothermic reaction, while a fall usually indicates an endothermic reaction. However, poor insulation, incomplete reaction, evaporation, slow energy transfer, or an initially hot reactant can complicate the observation. Define the reaction using the direction of heat transfer and use temperature change as supporting evidence.
Exothermic does not mean fast
Thermodynamics and kinetics answer different questions. The sign of ΔH describes the net energy change, while reaction rate depends on factors such as activation energy, temperature, concentration, pressure, surface area, and catalysts. An exothermic reaction can be extremely slow.
Exothermic does not automatically mean spontaneous
At IB Higher Level, spontaneity is assessed using Gibbs energy:
ΔG = ΔH − TΔS
A negative ΔH can favor thermodynamic spontaneity, but entropy and temperature also matter. Conversely, an endothermic process can be spontaneous if the entropy contribution is sufficiently favorable.
Reversing a reaction reverses the sign
If a reaction has ΔH = −50 kJ mol⁻¹ in the forward direction, the reverse reaction has ΔH = +50 kJ mol⁻¹. The same energy transferred out in one direction must be transferred in when the process is reversed under equivalent conditions.
Physical states matter
Enthalpy depends on physical state. Forming H₂O(l) does not have the same enthalpy change as forming H₂O(g), because condensation or vaporization involves an additional energy change. Always retain state symbols in thermochemical equations.
How to answer exothermic vs endothermic exam questions
A complete explanation should normally connect observation, energy flow, enthalpy levels, and sign. For example:
The temperature of the solution increases, showing that the solution has absorbed thermal energy. Energy is therefore transferred from the reacting system to the surroundings. The products have lower enthalpy than the reactants, so ΔH is negative and the reaction is exothermic.
For an endothermic process, reverse the direction and sign:
The temperature of the surroundings decreases because energy is transferred from the surroundings to the reacting system. The products have higher enthalpy than the reactants, so ΔH is positive and the reaction is endothermic.
Avoid one-word conclusions when the command term asks you to explain. Practice identifying the sign and interpreting profiles with the IB Chemistry energy-transfer resources, then apply the ideas through the Energetics and Thermochemistry Questionbank. The energetics flashcards are useful for recalling definitions, but exam preparation should also include calculations and written explanations.
Conclusion
The difference between exothermic and endothermic reactions is the direction of thermal energy transfer. Exothermic reactions transfer energy to the surroundings and have negative ΔH, while endothermic reactions absorb energy from the surroundings and have positive ΔH.
On an energy profile, exothermic products lie below the reactants and endothermic products lie above them. In calorimetry, the observed temperature change belongs to the measured surroundings, so the reaction's heat has the opposite sign. RevisionDojo study notes, Questionbank practice, flashcards, and Jojo AI can help you test this distinction through energy profiles, explanations, and calculation questions.
