Energetics and Thermodynamics
Energetics and thermodynamics is the study of enthalpy changes that accompany chemical reactions, including standard enthalpies of formation, combustion and reaction, and how these are calculated using bond enthalpies, Hess's law cycles and experimental calorimetry. It also introduces entropy as a measure of disorder, and is examined throughout A Level Chemistry.
Method
- Identify which type of standard enthalpy change is described (formation: 1 mole of compound from elements; combustion: 1 mole of substance burned completely in oxygen; reaction: as written in the equation) and check all substances are in their standard states at standard conditions.
- For calorimetry calculations, use q = m x c x delta T to find the heat energy transferred to the water or solution, then divide by the amount (in mol) of substance reacted to find the enthalpy change per mole.
- For bond enthalpy calculations, add up the bond enthalpies of all bonds broken in the reactants, add up the bond enthalpies of all bonds formed in the products, then calculate delta H = (energy to break bonds) - (energy released forming bonds).
- Draw a Hess's law cycle linking the enthalpy change you want to find with enthalpy changes you know (such as combustion or formation data), keeping every arrow direction consistent, then add or subtract enthalpy values around the cycle to find the unknown value.
- Remember that reversing the direction of an enthalpy change reverses its sign, and that Hess's law states the total enthalpy change for a reaction is independent of the route taken, provided the initial and final conditions are the same.
- To predict the sign of an entropy change, compare the number of moles of gas (and the amount of disorder) on each side of the equation; entropy increases when moles of gas increase, or a solid/liquid becomes a gas.
Worked example
Standard enthalpies of combustion (kJ/mol): delta Hc[C(s)] = -393, delta Hc[H2(g)] = -286, delta Hc[C2H6(g)] = -1560. Use a Hess's law cycle to calculate the standard enthalpy of formation of ethane, C2H6.
- Write the balanced formation equation: 2C(s) + 3H2(g) -> C2H6(g).
- Set up the Hess's law relationship using combustion data: delta Hf(C2H6) = [2 x delta Hc(C) + 3 x delta Hc(H2)] - delta Hc(C2H6).
- Substitute the values: delta Hf(C2H6) = [2 x (-393) + 3 x (-286)] - (-1560).
- Evaluate the bracket: 2 x (-393) = -786, and 3 x (-286) = -858, giving -786 + (-858) = -1644.
- Complete the calculation: -1644 - (-1560) = -1644 + 1560 = -84.
- Final answer: delta Hf(C2H6) = -84 kJ/mol
Practice questions
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Q1State two standard conditions that must apply when quoting a standard enthalpy change.Show answer
Answer: Pressure of 100 kPa, and a stated temperature (usually 298 K); solutions, if present, at 1 mol/dm3
Q2Define standard enthalpy of combustion.Show answer
Answer: The enthalpy change when 1 mole of a substance is completely burned in oxygen, with all reactants and products in their standard states under standard conditions
Q3A reaction has delta H = +65 kJ/mol. State, with a reason, whether it is exothermic or endothermic.Show answer
Answer: Endothermic, because delta H is positive (energy is absorbed from the surroundings)
Q4Mean bond enthalpies (kJ/mol): H-H = +436, Br-Br = +193, H-Br = +366. Calculate delta H for H2(g) + Br2(g) -> 2HBr(g).Show answer
Answer: -103 kJ/mol (bonds broken = 436 + 193 = 629; bonds formed = 2 x 366 = 732; delta H = 629 - 732 = -103)
Q5Calculate the energy transferred, q, when 200 g of water is heated from 21.0 degrees C to 35.0 degrees C. Specific heat capacity of water, c = 4.18 J/(g K).Show answer
Answer: 11.7 kJ (q = 200 x 4.18 x 14.0 = 11704 J = 11.7 kJ)
Q6Lattice dissociation enthalpy of KCl(s) = +711 kJ/mol. Enthalpy of hydration of K+(g) = -322 kJ/mol. Enthalpy of hydration of Cl-(g) = -363 kJ/mol. Calculate the enthalpy of solution of KCl(s).Show answer
Answer: +26 kJ/mol (enthalpy of solution = lattice dissociation enthalpy + sum of hydration enthalpies = 711 + (-322) + (-363) = 711 - 685 = +26)
Exam-style questions
Written in the style of a A Level Science exam paper, with a full mark scheme.
State Hess's law, and explain what it allows you to do when the enthalpy change of a reaction cannot be measured directly.
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A student burns 0.68 g of ethanol (Mr = 46.0) to heat 100 g of water, raising its temperature from 19.5 degrees C to 42.3 degrees C. Calculate the experimental enthalpy of combustion of ethanol, in kJ/mol, giving your answer to 3 significant figures. Specific heat capacity of water, c = 4.18 J/(g K).
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Explain, in terms of bond breaking and bond making, why the reaction between hydrogen and fluorine, H2(g) + F2(g) -> 2HF(g), is highly exothermic, given mean bond enthalpies (kJ/mol): H-H = +436, F-F = +158, H-F = +567.
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