AQA A Level Chemistry
AQA A Level Chemistry 3.1.4 Energetics Exam Questions
Free AQA A Level Chemistry exam questions for 3.1.4 Energetics, with a mark scheme for every pack. Covers enthalpy changes, standard conditions, calorimetry with q = mcΔT, Hess's law cycles and mean bond enthalpy calculations.
3.1.4 Energetics exam questions
1 question pack of real past exam questions, each with its mark scheme. Do not stop after a few: answer as many as you can without notes, then mark every answer with the mark scheme. The more questions you do, the more quickly you will spot how examiners ask about each idea.
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About AQA A Level Chemistry 3.1.4 Energetics
Energetics deals with the heat energy transferred during chemical reactions. On the AQA A Level course you need to define standard enthalpy changes, measure enthalpy changes by calorimetry using q = mcΔT (Required practical 2), use Hess's law to find enthalpy changes that cannot be measured directly, and calculate enthalpy changes from mean bond enthalpies. The topic leads directly into 3.1.8 Thermodynamics, where Born-Haber cycles and Gibbs free energy extend the same ideas. Examiners frequently test definitions word for word, calorimetry calculations with sign and unit traps, Hess cycles from formation or combustion data, and evaluation of why experimental values differ from data book values.
What you need to know
- Enthalpy change ΔH is the heat energy change measured at constant pressure. Exothermic reactions have negative ΔH; endothermic reactions have positive ΔH. Standard conditions are 100 kPa and a stated temperature, usually 298 K.
- Standard enthalpy definitions Standard enthalpy of combustion: the enthalpy change when one mole of a substance is completely burned in oxygen with all reactants and products in their standard states under standard conditions. Standard enthalpy of formation: the enthalpy change when one mole of a compound is formed from its elements, all in standard states under standard conditions.
- Calorimetry q = mcΔT, where m is the mass of the water or solution heated (not the fuel), c is 4.18 J g-1 K-1 and ΔT is the temperature change. Divide q by the moles reacting to get ΔH in kJ mol-1 and add the correct sign.
- Required practical 2 Measure an enthalpy change, for example of a neutralisation or displacement reaction in a polystyrene cup. Plotting temperature against time and extrapolating back to the time of mixing corrects for heat loss.
- Hess's law The enthalpy change for a reaction is independent of the route taken. With formation data, ΔH = sum of ΔfH(products) minus sum of ΔfH(reactants). With combustion data, ΔH = sum of ΔcH(reactants) minus sum of ΔcH(products).
- Mean bond enthalpies ΔH = sum of bond enthalpies broken minus sum of bond enthalpies made. Values are averages over many compounds and assume gaseous species, so answers differ from accurate data book values.
Common mistakes in 3.1.4 Energetics
- Using the mass of fuel or solid in q = mcΔT. The mass is the water or solution whose temperature changes.
- Forgetting the sign or the unit conversion. Convert J to kJ and give exothermic values as negative.
- Reversing bond enthalpy calculations. It is bonds broken minus bonds made.
- Defining enthalpy of formation without saying one mole of compound, from its elements, in standard states. Each part of the definition is a marking point.
- Forgetting that the standard enthalpy of formation of an element in its standard state is zero.
Try three 3.1.4 Energetics exam-style questions
1. Define the standard enthalpy of formation. (2 marks)
Show the mark scheme answer
The enthalpy change when one mole of a compound is formed from its elements; with all reactants and products in their standard states under standard conditions.
2. Burning 0.150 g of ethanol (Mr 46.0) raised the temperature of 50.0 g of water by 18.5 K. Calculate the enthalpy of combustion of ethanol in kJ mol-1. c = 4.18 J g-1 K-1. (3 marks)
Show the mark scheme answer
q = 50.0 x 4.18 x 18.5 = 3867 J (3.87 kJ); amount of ethanol = 0.150 / 46.0 = 3.26 x 10^-3 mol; ΔcH = -3.87 / 3.26 x 10^-3 = -1190 kJ mol-1 (negative sign required).
3. Use the mean bond enthalpies C-H 413, O=O 498, C=O 805 and O-H 464 kJ mol-1 to calculate ΔH for CH4(g) + 2O2(g) → CO2(g) + 2H2O(g). (3 marks)
Show the mark scheme answer
Bonds broken = (4 x 413) + (2 x 498) = 2648 kJ mol-1; bonds made = (2 x 805) + (4 x 464) = 3466 kJ mol-1; ΔH = 2648 - 3466 = -818 kJ mol-1.
3.1.4 Energetics FAQs
What does AQA A Level Chemistry 3.1.4 Energetics cover?
It covers enthalpy changes and their definitions, calorimetry with q = mcΔT, Hess's law and mean bond enthalpy calculations. It includes Required practical 2 on measuring an enthalpy change.
Which paper is 3.1.4 Energetics examined in?
It can be tested in both Paper 1 and Paper 2, and Paper 3 can draw on it too, often through practical calorimetry data.
Which definitions should I learn word for word?
Learn standard enthalpy of combustion, standard enthalpy of formation and mean bond enthalpy, plus standard conditions. Examiners award separate marks for one mole, standard states and standard conditions, so missing a phrase costs a mark.
