Core ideas

• Chemical reactions and physical changes involve energy transfer between the system and the surroundings.

• Total energy is conserved: energy is not created or destroyed, only transferred.

• System = the chemicals/process being studied. Surroundings = everything outside the system.

• Heat is energy transferred because of a temperature difference.

• Temperature is a measure of the average kinetic energy of particles.

• Do not confuse heat with temperature: a substance can gain/lose heat energy and show a temperature change, but they are not the same quantity.

Endothermic vs exothermic

• Endothermic reaction: energy is transferred from the surroundings to the system.

• In an endothermic process, the temperature of the surroundings decreases.

• Exothermic reaction: energy is transferred from the system to the surroundings.

• In an exothermic process, the temperature of the surroundings increases.

• Exam link: if the measured temperature rises, the reaction is usually exothermic; if the measured temperature falls, it is usually endothermic.

• Sign convention:

  • Exothermic: $\Delta H < 0$

  • Endothermic: $\Delta H > 0$

This diagram shows a coffee-cup calorimeter, the standard school setup for measuring enthalpy changes at constant pressure. It helps you visualize which parts reduce heat loss and where the temperature change is measured. Source

Relative stability and enthalpy change

• Whether a reaction is endothermic or exothermic depends on the relative stability of reactants and products.

• More stable species have lower potential energy.

• If products are lower in energy than reactants, energy is released and the reaction is exothermic.

• If products are higher in energy than reactants, energy is absorbed and the reaction is endothermic.

• On an energy profile:

  • y-axis = potential energy

  • x-axis = reaction coordinate

Energy profile diagrams

• Exothermic profile:

  • Reactants start at higher potential energy than products.

  • $\Delta H$ is negative.

• Endothermic profile:

  • Products end at higher potential energy than reactants.

  • $\Delta H$ is positive.

• Always check:

  • relative heights of reactants and products

  • sign of $\Delta H$

  • correct axis labels: reaction coordinate and potential energy

• A common exam error is mixing up activation energy with enthalpy change — in this topic, focus on overall energy difference between reactants and products.

This image compares exothermic and endothermic energy profiles side by side. It is useful for spotting the direction of $\Delta H$ and comparing the relative energies of reactants and products. Source

Standard enthalpy change and conditions

• Standard enthalpy change, $\Delta H^\circ$, is the heat transferred at constant pressure under standard conditions and standard states.

• In this topic, you calculate enthalpy change from the temperature change of a pure substance or solution during an experiment.

• Unit of $\Delta H^\circ$: kJ mol$^{-1}$.

• In calorimetry questions, the measured temperature change is usually used to find $Q$, then converted to enthalpy change per mole.

Calorimetry and measuring enthalpy changes

• Calorimetry measures heat transferred by recording a temperature change.

• Typical school method: measure temperature before and after a reaction in an insulated container.

• Assume the solution behaves like water unless told otherwise.

• Key idea:

  • if the water/solution gains heat, the reaction loses heat

  • if the water/solution loses heat, the reaction gains heat

• This is why the equation for reaction enthalpy includes a negative sign.

These diagrams show how heat flows in exothermic and endothermic processes inside a calorimeter. They are helpful for linking temperature change, sign of $q$, and the direction of energy transfer between system and surroundings. Source

Required equations

• $Q = mc\Delta T$

• $\Delta H = -\dfrac{Q}{n}$

• Meanings:

  • $Q$ = heat energy transferred

  • $m$ = mass of substance heated/cooled

  • $c$ = specific heat capacity

  • $\Delta T$ = temperature change

  • $n$ = amount of reacting substance in mol

• For water, $c$ is given in the data booklet.

• Unit check:

  • $Q$ usually first found in J

  • convert to kJ before finding $\Delta H$ in kJ mol$^{-1}$ if needed

How to calculate enthalpy change in exam questions

• Step 1: Find $\Delta T = T_{final} - T_{initial}$.

• Step 2: Calculate $Q$ using $Q = mc\Delta T$.

• Step 3: Decide which substance the mass refers to, usually the water/solution.

• Step 4: Convert $Q$ from J to kJ if needed.

• Step 5: Find moles, $n$, of the substance asked for in the question.

• Step 6: Use $\Delta H = -\dfrac{Q}{n}$.

• Step 7: Add the correct sign and units: kJ mol$^{-1}$.

Sign conventions and exam traps

• If temperature increases, then $\Delta T$ is positive, $Q$ for the water/solution is positive, so $\Delta H$ for the reaction is negative.

• If temperature decreases, then $\Delta T$ is negative, $Q$ for the water/solution is negative, so $\Delta H$ for the reaction is positive.

• Always be clear whether $Q$ refers to the surroundings or the reaction.

• The reaction enthalpy is usually the opposite sign to the heat change of the water/solution.

• Use the amount in moles of the limiting/reacting substance named in the question, not automatically the mass of solution.

• Do not forget unit conversion from cm$^3$ to g if density is assumed to be 1.0 g cm$^{-3}$ for aqueous solutions.

Practical method and sources of error

• In school calorimetry, the measured temperature change is often smaller than the theoretical value.

• Main reasons:

  • heat lost to the surroundings

  • heat absorbed by the apparatus

  • incomplete combustion or incomplete reaction

  • evaporation or delayed temperature reading

• Therefore, experimental values for enthalpy change are often less exothermic or less endothermic than theoretical values.

• Practical improvements:

  • use a lid

  • improve insulation

  • stir thoroughly

  • measure temperature quickly and carefully

  • use accurate masses/volumes

This diagram shows multiple enthalpy steps adding up to an overall exothermic enthalpy change. It is useful for understanding that enthalpy diagrams can represent energy stages and not just a single up/down arrow. Source