Lesson 13A.3

13A.3 Enthalpy of solution and energy cycles Quiz: Pearson Edexcel Chemistry, Unit 13

20 questions

Written by Revision Ninja · Last reviewed 11 October 2026

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Lesson 13A.3, Enthalpy of solution and energy cycles: 20 multiple choice questions for the Pearson Edexcel Chemistry (9CH0), Unit 13: Energetics II, written with Revision Ninja.

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The 20 questions

  1. What is the enthalpy of solution?

    • The enthalpy change when one mole of solute is vaporised from a concentrated solution at its boiling point
    • The enthalpy change when one mole of an ionic solid dissolves in water to form an infinitely dilute solution
    • The enthalpy change when one mole of water is hydrated by the ions released from one mole of an ionic solid
    • The energy change when one mole of gaseous ions forms a solid lattice, the reverse of lattice dissociation

  2. What is the enthalpy of hydration of an ion?

    • The energy needed to break the lattice of one mole of ionic solid into gaseous ions in the gas phase
    • The enthalpy change when one mole of gaseous ions dissolves in water to form aqueous ions
    • The enthalpy change when one mole of a solid ionic compound dissolves in water to give hydrated ions
    • The enthalpy change when one mole of water is formed from gaseous hydrogen and oxygen ions in solution

  3. What is the sign of an enthalpy of hydration?

    • Always zero for small ions
    • Positive for cations and negative for anions
    • Always exothermic (negative)
    • Always endothermic, because water molecules must be separated

  4. Which ion has the more negative enthalpy of hydration?

    • Mg2+, because it has a higher charge density than Na+
    • Cl-, because anions always hydrate more than cations
    • Na+, because it is larger
    • Ca2+, because it has a larger radius than Mg2+

  5. What factors determine the magnitude of an ion's enthalpy of hydration?

    • Only the temperature at which the solution is made
    • The ionic mass only, because heavier ions are hydrated more strongly
    • The number of electrons in the outer shell only
    • The ionic charge and ionic radius, through the charge density of the ion

  6. Which expression correctly relates the enthalpy of solution to the lattice energy and hydration enthalpies?

    • dsolH = dhydH(cation) - dhydH(anion) - LE
    • dsolH = LE + dhydH(cation) + dhydH(anion)
    • dsolH = LE - dhydH(cation) - dhydH(anion)
    • dsolH = dhydH(cation) + dhydH(anion) - LE

  7. Why does a small, highly charged ion have a more negative enthalpy of hydration?

    • It has a high charge density
    • It has fewer electrons, so it hydrates less readily
    • It only dissolves in ethanol, not in water
    • It has a lower charge, so it attracts fewer water molecules

  8. Using a lattice energy of -700 kJ mol-1 and a total hydration enthalpy of -650 kJ mol-1, what is the enthalpy of solution?

    • +1350 kJ mol-1
    • +650 kJ mol-1
    • +50 kJ mol-1
    • -50 kJ mol-1

  9. Using a lattice energy of -820 kJ mol-1 and a total hydration enthalpy of -850 kJ mol-1, what is the enthalpy of solution?

    • +1670 kJ mol-1
    • -30 kJ mol-1
    • +30 kJ mol-1
    • -1670 kJ mol-1

  10. The enthalpy of solution of a salt is +12 kJ mol-1 and its lattice energy is -760 kJ mol-1. What is the total hydration enthalpy?

    • +772 kJ mol-1
    • -748 kJ mol-1
    • +748 kJ mol-1
    • -772 kJ mol-1

  11. Why can a salt with a positive enthalpy of solution still dissolve?

    • It dissolves because the solution becomes more ordered than the solid
    • It does not dissolve enthalpically
    • It dissolves because the lattice energy is positive
    • It dissolves because the hydration enthalpy must be zero

  12. A salt has dsolH = +80 kJ mol-1. Which statement is true of its dissolution?

    • It dissolves only if the lattice energy is zero, so that no energy is needed to separate the ions
    • It is exothermic, so the solution warms as it dissolves and the enthalpy term favours dissolution strongly
    • It is endothermic, so the enthalpy term opposes dissolution and entropy must drive it
    • It cannot dissolve at any temperature because dsolH is positive, which makes the process always unfavourable

  13. When ammonium nitrate dissolves in water and the solution becomes colder, what is the sign of the enthalpy of solution?

    • Negative, so the process is exothermic
    • Negative, and larger in magnitude than the lattice energy
    • Positive, so the process is endothermic
    • Zero, so no heat is exchanged

  14. Using a lattice energy of -2258 kJ mol-1 and a total hydration enthalpy of -2308 kJ mol-1 for CaCl2, what is the enthalpy of solution?

    • -50 kJ mol-1
    • -4566 kJ mol-1
    • +4566 kJ mol-1
    • +50 kJ mol-1

  15. In the enthalpy of solution cycle, what is the first step applied to the solid?

    • Forming dissolved ions in solution with no energy change
    • Forming hydrated ions directly from the solid
    • Breaking the lattice to form gaseous ions
    • Forming gaseous atoms from the element

  16. What is the enthalpy of solution of KBr, given lattice energy -671 kJ mol-1, hydration enthalpy of K+ = -322 kJ mol-1 and of Br- = -334 kJ mol-1?

    • +15 kJ mol-1
    • -15 kJ mol-1
    • -993 kJ mol-1
    • +993 kJ mol-1

  17. A solid has lattice energy -1000 kJ mol-1 and total hydration enthalpy -900 kJ mol-1. A student says it will dissolve readily because dsolH is negative. Which evaluation is best?

    • Correct, because a negative hydration enthalpy always makes dissolution favourable
    • Incorrect, because dsolH = +100 kJ mol-1 is endothermic, so the enthalpy term opposes dissolution
    • Incorrect, because the lattice energy must be positive for a solid to dissolve
    • Correct, because hydration energy is always larger than the lattice energy

  18. A salt has measured dsolH = -35 kJ mol-1 and lattice energy -1200 kJ mol-1. What is its total hydration enthalpy?

    • -1235 kJ mol-1
    • +1235 kJ mol-1
    • +1165 kJ mol-1
    • -1165 kJ mol-1

  19. Why must an enthalpy of solution cycle use gaseous ions as its intermediate state?

    • Because lattice energies and hydration enthalpies are both defined in terms of gaseous ions
    • Because hydration enthalpies are defined for solids rather than ions
    • Because water cannot dissolve any species that is not gaseous
    • Because gaseous ions are the only species that can be measured in solution

  20. A student says the enthalpy of solution is always more negative than the lattice energy. Which evaluation is best?

    • True, because hydration enthalpies are always positive, so the sum is always more negative than the lattice energy
    • False, because dsolH is always zero for any ionic solid
    • False, because dsolH = -LE + dhydH(total), which can be positive if hydration is less exothermic than the lattice energy
    • False, because dsolH is always zero for any ionic solid when the solution is made at 298 K

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