DONATSciences and Maths
Chemistry · Practice book · Senior 5 · Unit 14 of 15

Enthalpy change of reactions

Senior 5 Chemistry · 14 questions with answers. Try each question on paper, then tap it to see the answer.

Unit 14 Enthalpy change of reactions

Key unit competence: Use experimental data to determine the enthalpy changes of reactions and apply Hess's law, Born–Haber cycles and bond enthalpies.

Key ideas

  • Enthalpy change ΔH is the heat change at constant pressure. Exothermic: ΔH negative; endothermic: ΔH positive. Standard conditions: 298 K, 100 kPa (1 atm), 1 mol dm−3 solutions, substances in standard states.
  • Standard enthalpies: of formation ΔH°f (1 mol formed from its elements; zero for elements), of combustion ΔH°c (1 mol burnt completely in oxygen), of neutralisation (1 mol water formed; −57.3 kJ mol−1 for strong acid + strong base), of atomisation, of solution.
  • Calorimetry: q = mcΔT (c of water = 4.18 J g−1 K−1), then ΔH = −q ÷ n. Experimental ΔHc values are less exothermic than data-book values because of heat loss, incomplete combustion and evaporation.
  • Hess's law: the enthalpy change of a reaction depends only on the initial and final states, not the route. ΔHr = ΣΔH°f(products) − ΣΔH°f(reactants) = ΣΔH°c(reactants) − ΣΔH°c(products).
  • Lattice enthalpy: enthalpy change when 1 mol of an ionic solid forms from its gaseous ions (exothermic). It is more negative for smaller ions and higher charges (MgO much more than NaCl). A Born–Haber cycle finds it: ΔHf = ΔHat(metal) + IE + ΔHat(non-metal) + EA + ΔHlatt.
  • Bond enthalpy: energy to break 1 mol of a bond in gaseous molecules (average values). ΔH = Σ(bonds broken) − Σ(bonds formed). Results are approximate because average bond enthalpies are used and substances may not be gaseous.

Common misconceptions

  • "Breaking bonds releases energy." No: Breaking bonds always needs (absorbs) energy; forming bonds releases it.
  • "A negative ΔH means the surroundings get colder." No: Negative ΔH is exothermic: heat goes out and the surroundings get warmer.
  • "The enthalpy of formation of oxygen gas is large because it is reactive." No: ΔH°f of any element in its standard state is zero.

Worked examples

Example 1. Describe how to measure the enthalpy of neutralisation of HCl with NaOH.

Solution

Measure 50.0 cm3 of 1.00 mol dm−3 HCl into an insulated polystyrene cup; record its steady temperature.

Measure 50.0 cm3 of 1.00 mol dm−3 NaOH at the same temperature; add quickly, stir, record the highest temperature.

q = (100 g) × 4.18 × ΔT; n(H2O) = 0.0500 mol.

ΔH = −q ÷ n (in kJ mol−1). Use a lid and correct for cooling by plotting temperature against time.

Example 2. Draw (describe) a Born–Haber cycle for magnesium oxide and say how the lattice enthalpy is found.

Solution

Start: Mg(s) + ½O2(g) → MgO(s) is ΔHf.

Steps up: atomise Mg; first and second ionisation energies of Mg; atomise ½O2; first EA of O (exothermic) and second EA (endothermic).

Gaseous Mg2+ and O2− then form MgO(s): lattice enthalpy.

By Hess's law: ΔHlatt = ΔHf − (sum of all the other steps).

Practice questions

  1. Define standard enthalpy of formation.

    Show answer

    The enthalpy change when 1 mol of a compound is formed from its elements in their standard states under standard conditions.

    Common mistake: Forgetting 'one mole' or 'standard states'.

  2. State Hess's law.

    Show answer

    The total enthalpy change of a reaction is the same whatever route is taken, provided the initial and final states are the same.

    Common mistake: Saying the energy depends on the catalyst.

  3. Why is a polystyrene cup used in calorimetry?

    Show answer

    It is a good insulator with a low heat capacity, so less heat is lost.

    Common mistake: Saying it reacts with the acid.

  4. Explain why the enthalpy of neutralisation is the same for HCl and HNO3 with NaOH.

    Show answer

    Both are strong acids, fully ionised; the reaction is always H+ + OH− → H2O.

    Common mistake: Saying the anions react.

  5. Why is the second electron affinity of oxygen endothermic?

    Show answer

    Adding an electron to the negative O− ion needs energy to overcome repulsion.

    Common mistake: Saying all electron affinities are exothermic.

  6. Explain why MgO has a much higher melting point than NaCl.

    Show answer

    Its lattice enthalpy is much more exothermic: smaller ions with 2+ and 2− charges attract more strongly.

    Common mistake: Saying MgO is covalent.

  7. Calculate q when 100 g of water warms by 4.0 K.

    Show answer

    q = 100 × 4.18 × 4.0 = 1672 J = 1.67 kJ.

    Common mistake: Using the mass of the solid instead of the water.

  8. Give two sources of error in measuring ΔHc with a spirit burner.

    Show answer

    Heat loss to the air and can; incomplete combustion (soot); evaporation of fuel; heat absorbed by the can (any two).

    Common mistake: Saying the thermometer causes heat loss.

  9. Write the equation for the standard enthalpy of formation of water.

    Show answer

    H2(g) + 1/2O2(g) → H2O(l)

    Common mistake: Writing 2 mol of water.

  10. Is the enthalpy of atomisation endothermic or exothermic? Explain.

    Show answer

    Endothermic: bonds or metallic bonding must be broken to form gaseous atoms.

    Common mistake: Saying exothermic.

Matching and drag-and-drop

  1. Matching

    Match each item in column A with the correct answer in column B. Write the letter next to each number.

    Column AColumn B
    1. The standard enthalpy of combustion refers to burning:a. exothermic
    2. An energy profile with products lower than reactants shows a reaction that is:b. 1 mol of a substance completely in oxygen
    3. Average bond enthalpies give approximate answers because:c. graphite
    4. The standard state of carbon used for enthalpies of formation is:d. a bond's strength depends slightly on its environment
    Show answer

    1 → b (1 mol of a substance completely in oxygen) · 2 → a (exothermic) · 3 → d (a bond's strength depends slightly on its environment) · 4 → c (graphite)

    • Complete combustion of one mole.
    • Energy is released.
    • e.g. C–H differs a little between molecules.
    • Graphite is the most stable form.

  2. Drag and drop

    Choose the correct word or phrase from the box for each gap. One word in the box is not needed.

    298 K  ·  negative  ·  zero  ·  positive
    1. The standard enthalpy of formation of an element in its standard state is .
    2. An exothermic reaction has ΔH that is .
    3. Standard conditions for enthalpy changes include a temperature of .
    Show answer

    1: zero · 2: negative · 3: 298 K

    • No change is needed to form it from itself.
    • Heat is given out to the surroundings.
    • 25 °C and 100 kPa.

  3. Matching

    Match each item in column A with the correct answer in column B. Write the letter next to each number.

    Column AColumn B
    1. Which change is endothermic?a. endothermic
    2. ΔH for a reaction from bond enthalpies is:b. kJ mol−1
    3. Breaking chemical bonds is:c. bonds broken − bonds formed
    4. The SI unit of molar enthalpy change is:d. Melting ice
    Show answer

    1 → d (Melting ice) · 2 → c (bonds broken − bonds formed) · 3 → a (endothermic) · 4 → b (kJ mol−1)

    • Heat is absorbed to separate the molecules.
    • Endothermic part minus exothermic part.
    • Energy must be supplied.
    • Energy per mole.

  4. Drag and drop

    Choose the correct word or phrase from the box for each gap. One word in the box is not needed.

    Na(s) → Na(g)  ·  is independent of the route taken  ·  is always negative  ·  −57 kJ mol−1
    1. In a Born–Haber cycle, the enthalpy of atomisation of sodium is for .
    2. Hess's law states that the enthalpy change of a reaction .
    3. The enthalpy of neutralisation of a strong acid with a strong base is about .
    Show answer

    1: Na(s) → Na(g) · 2: is independent of the route taken · 3: −57 kJ mol−1

    • Solid to gaseous atoms.
    • Energy is conserved.
    • Both are fully ionised: H+ + OH− → H2O.

Take the Unit 14 quiz → · Back to contents