Edexcel International A Level Chemistry Energetics Exam Questions

About Edexcel International A Level Chemistry Energetics

Energetics in Unit 2 of Edexcel International A Level Chemistry is about measuring and calculating the heat energy transferred in chemical reactions. You define standard enthalpy changes under standard conditions, carry out calorimetry experiments and use Hess's law to find enthalpy changes that cannot be measured directly. Mean bond enthalpies give another route to estimate enthalpy changes. The topic is extended in Unit 4 through Lattice energy and Entropy. Examiners often set calorimetry calculations with an evaluation of experimental errors, Hess cycles using formation or combustion data, and questions asking why bond enthalpy values differ from data book values.

What you need to know

• Enthalpy change Enthalpy change, delta H, is the heat energy change at constant pressure. Exothermic reactions have negative delta H; endothermic have positive delta H. Standard conditions are 100 kPa and a stated temperature, usually 298 K.
• Standard enthalpy definitions Learn exact definitions of standard enthalpy change of reaction, formation, combustion, neutralisation and atomisation. For example, enthalpy of formation is for one mole of a compound formed from its elements in their standard states.
• Calorimetry Energy transferred Q = mc delta T, where m is the mass of solution (usually taken as the volume in cm3), c is 4.18 J g-1 K-1 and delta T is the temperature change. Divide by moles of the limiting reagent and add the sign.
• Experimental errors Heat loss to the surroundings is the main error. Improve accuracy with an insulated polystyrene cup and lid, and by extrapolating a temperature against time graph back to the time of mixing.
• Hess's law The enthalpy change for a reaction is independent of the route taken. Use formation data (products minus reactants) or combustion data (reactants minus products) in cycles.
• Mean bond enthalpies Delta H = sum of bonds broken minus sum of bonds formed. Results differ from Hess's law values because mean bond enthalpies are averages over many compounds and assume all species are gaseous.

Common mistakes in Energetics

• Forgetting the negative sign for an exothermic reaction when converting Q into delta H.
• Using the mass of the solid added in Q = mc delta T instead of the mass of the solution.
• Leaving Q in J but quoting delta H in kJ mol-1 without dividing by 1000.
• Getting bond enthalpy calculations the wrong way round. It is bonds broken minus bonds made.
• Writing an enthalpy of formation equation for more than one mole of product, or with elements not in their standard states, for example using O instead of O2.

Try three Energetics exam-style questions