Exothermic and Endothermic Reactions
This topic covers the fundamental classification of chemical reactions based on energy flow. Understanding whether a reaction releases heat (exothermic) or absorbs it (endothermic) is key to predicting reaction outcomes and interpreting energy data.
Part of the ESAT Chemistry syllabus — revision for the Engineering and Science Admissions Test (ESAT), the UAT-UK admissions test for Cambridge, Imperial, Oxford and UCL.
Key points
- Exothermic reactions release heat energy into the surroundings, causing the temperature of the surroundings to increase.
- Endothermic reactions absorb heat energy from the surroundings, causing the temperature of the surroundings to decrease.
- The enthalpy change, ΔH, is the measure of this heat exchange at constant pressure.
- For exothermic reactions, the products have less stored chemical energy than the reactants, so ΔH is negative.
- For endothermic reactions, the products have more stored chemical energy than the reactants, so ΔH is positive.
- Common examples include combustion (exothermic) and melting ice (an endothermic physical process).
Diagram
Formulae
Exothermic: ΔH < 0 Used to represent any reaction where heat is released to the surroundings. The products are more stable (lower energy) than the reactants.
Endothermic: ΔH > 0 Used to represent any reaction where heat is absorbed from the surroundings. The products are less stable (higher energy) than the reactants.
Definitions
- Exothermic Reaction
- A process that releases energy to its surroundings, usually in the form of heat. The system's enthalpy decreases, so ΔH is negative.
- Endothermic Reaction
- A process that absorbs energy from its surroundings, usually in the form of heat. The system's enthalpy increases, so ΔH is positive.
- Enthalpy Change (ΔH)
- The overall heat energy change of a system during a reaction at constant pressure. Its sign indicates the direction of heat flow.
Worked example
The thermal decomposition of 0.2 moles of calcium carbonate absorbs 35.6 kJ of energy. What is the standard enthalpy change of reaction, ΔH, in kJ mol⁻¹? A) -178 kJ mol⁻¹ B) +178 kJ mol⁻¹ C) -7.12 kJ mol⁻¹ D) +35.6 kJ mol⁻¹
- 1
Step 1:
Identify the direction of energy flow.
The reaction 'absorbs' energy, which means it is an endothermic process.
- 2
Step 2:
Determine the sign of ΔH.
For an endothermic reaction, the enthalpy change ΔH must be positive.
This eliminates options A and C.
- 3
Step 3:
Calculate the enthalpy change per mole.
The energy given (35.6 kJ) corresponds to 0.2 moles of reactant.
- 4
Step 4:
Scale the value to one mole.
To find the energy per mole, divide the energy by the number of moles:
ΔH = 35.6 kJ / 0.2 mol - 5
Step 5:
Perform the calculation without a calculator.
Dividing by 0.2 is the same as dividing by 1/5, which is equivalent to multiplying by 5.
So, 35.6 × 5 = (35 × 5) + (0.6 × 5) = 175 + 3 = 178.
- 6
Step 6:
State the final answer with the correct sign and units.
The enthalpy change is +178 kJ mol⁻¹.
Answer: B
Common mistakes
- ×Confusing the sign convention: Remember that heat *exiting* the system results in a *negative* ΔH because the system's energy has decreased.
- ×Forgetting to scale values to a per-mole basis. Enthalpy changes are typically reported in kJ mol⁻¹, so if the data is for a different amount (e.g., 0.5 moles or 2 moles), you must adjust your final answer accordingly.
- ×Mixing up the system and the surroundings. An exothermic reaction feels hot because the *surroundings* are gaining heat, but the *system* (the reacting chemicals) is losing it.
No-calculator tips
- ✓Use mnemonics: 'EXOthermic' means heat 'EXIts'. 'ENDOthermic' means heat 'ENters'.
- ✓To divide by a decimal without a calculator, convert it to a fraction. For example, dividing by 0.2 is the same as multiplying by 5; dividing by 0.5 is the same as multiplying by 2; dividing by 0.25 is the same as multiplying by 4.