Energy changes — Eduqas A-Level Chemistry

Energy changes in chemistry, also known as thermochemistry, is the study of heat energy transferred during chemical reactions. This topic is fundamental because it explains why reactions occur—whether they release heat (exothermic) or absorb heat (endothermic)—and how we can quantify these changes using enthalpy (H). In the WJEC A-Level Chemistry specification, you'll explore concepts like enthalpy change (ΔH), standard conditions, and Hess's law, which allows you to calculate enthalpy changes indirectly. Understanding energy changes is crucial for predicting reaction feasibility, designing industrial processes (e.g., in the Haber process), and grasping the thermodynamics behind everyday phenomena like combustion and respiration.

This topic builds on GCSE ideas of exothermic and endothermic reactions but introduces rigorous calculations and definitions. You'll learn to draw and interpret enthalpy profile diagrams, use calorimetry experiments to measure temperature changes, and apply bond enthalpy values to estimate reaction enthalpies. Mastery of energy changes is essential for later topics like entropy and free energy, which determine whether reactions are spontaneous. In the WJEC exam, you can expect questions on calculating ΔH using Hess's law, interpreting energy cycles, and explaining the difference between activation energy and enthalpy change.

Why does this matter? Energy changes underpin everything from the efficiency of fuels to the metabolic pathways in your body. For example, the combustion of petrol in a car engine is an exothermic reaction that releases energy to power the vehicle, while photosynthesis is an endothermic process that stores energy from sunlight. By studying this topic, you'll gain a deeper appreciation of how energy flows in chemical systems—a key concept for fields like materials science, environmental chemistry, and biochemistry.

• →Enthalpy change (ΔH): The heat energy transferred at constant pressure, measured in kJ mol⁻¹. Exothermic reactions have negative ΔH (heat released), endothermic have positive ΔH (heat absorbed).
• →Standard enthalpy changes: Defined under standard conditions (100 kPa, 298 K, 1 mol dm⁻³ for solutions). Key types include enthalpy of combustion (ΔHc), formation (ΔHf), and neutralisation (ΔHneut).
• →Hess's law: The total enthalpy change for a reaction is independent of the route taken. This allows calculation of ΔH for reactions that are difficult to measure directly, using known enthalpy values.
• →Calorimetry: Experimental technique to measure ΔH by tracking temperature change in a known mass of water. Requires calculations using q = mcΔT, then dividing by moles to get ΔH in kJ mol⁻¹.
• →Bond enthalpy: The energy required to break one mole of a specific covalent bond in gaseous molecules. Average bond enthalpies can be used to estimate ΔH for reactions: ΔH = Σ(bonds broken) – Σ(bonds formed).

• Definition of enthalpy changes of atomisation, lattice formation, lattice breaking, hydration, and solution
• Construction and interpretation of Born-Haber cycles
• Relationship between lattice breaking enthalpy, hydration enthalpies, and enthalpy of solution
• Definition of entropy as a measure of particle freedom
• Calculation of entropy change using absolute entropy values
• Application of the Gibbs free energy equation (ΔG = ΔH - TΔS)
• Interpretation of ΔG values regarding reaction spontaneity
• Explanation of spontaneous endothermic processes using entropy changes

• 💡Always check units for entropy (J K-1 mol-1) and enthalpy (kJ mol-1) before using the Gibbs equation
• 💡Draw clear energy cycles to avoid sign errors in Hess's law or Born-Haber calculations
• 💡Remember that ΔG must be negative for a reaction to be feasible
• 💡Use the correct state symbols when defining enthalpy changes
• 💡Be prepared to explain why a reaction becomes feasible at higher temperatures based on the TΔS term
• 💡Always include the correct sign for ΔH: negative for exothermic, positive for endothermic. In Hess's law calculations, carefully track the direction of arrows in enthalpy cycles—reversing a reaction flips the sign of ΔH.
• 💡In calorimetry questions, remember to convert temperature change from °C to K (the numerical change is the same) and ensure units are consistent: q in J, then convert to kJ. Also, account for heat loss to the surroundings by using a polystyrene cup and lid, and stirring to ensure even temperature distribution.
• 💡When using bond enthalpies, note that these are average values from different compounds, so calculated ΔH is an estimate. For more accurate results, use standard enthalpy of formation data. In exam questions, state whether your answer is an estimate if using bond enthalpies.

• Confusing lattice formation enthalpy with lattice breaking enthalpy
• Incorrectly assigning signs to enthalpy values in Born-Haber cycles
• Failing to convert units (e.g., J to kJ) when calculating Gibbs free energy
• Misinterpreting the effect of temperature on the feasibility of reactions
• Assuming all spontaneous reactions must be exothermic
• Misconception: 'Exothermic reactions always feel hot, endothermic always feel cold.' Correction: While many exothermic reactions release heat to the surroundings, some may not feel hot if the reaction is slow or heat is dissipated. Similarly, endothermic reactions absorb heat, but the surroundings can feel cold only if the reaction is rapid enough.
• Misconception: 'ΔH is the same as activation energy.' Correction: ΔH is the overall energy change between reactants and products, while activation energy (Ea) is the minimum energy needed to start the reaction. A reaction can have a large ΔH but a low Ea, or vice versa.
• Misconception: 'Hess's law only works for reactions that don't actually happen.' Correction: Hess's law is valid for any reaction, whether it occurs directly or not, because enthalpy is a state function. It's commonly used for reactions that are difficult to measure experimentally, like the formation of organic compounds.