Cycles are used to calculate enthalpies of solution for ionic compounds…

The idea

Dissolving an ionic solid is modelled as two steps. First the lattice is pulled apart into gaseous ions, which costs the lattice dissociation enthalpy and is strongly endothermic. Then each gaseous ion is surrounded by water molecules attracted to its charge, releasing the enthalpy of hydration, which is negative and larger in magnitude for small, highly charged ions. The enthalpy of solution is the sum: ∆solH = ∆latH(dissociation) + Σ∆hydH. For calcium chloride, +2237 for the lattice, −1650 for the calcium ion and two lots of −364 for the chloride ions give −141 kJ mol⁻¹, so the solution warms. Remember to count one hydration enthalpy for every ion in the formula, and to use the dissociation value, not the formation value, in this cycle.

What earns the marks

• A cycle or expression combining the lattice dissociation enthalpy with the hydration enthalpies.
• Multiplying the hydration enthalpy by the number of that ion in the formula.
• Using the correct sign for each term, positive for dissociation and negative for hydration.
• The final enthalpy of solution with sign and units, and for deducing the temperature change it implies.

Common mistakes and how to put them right

• Using the lattice formation enthalpy without reversing its sign for the dissolving step.

• Counting only one chloride hydration enthalpy for a compound of formula CaCl₂.

• Assuming dissolving must be endothermic because the lattice has to be broken.

• Predicting the temperature falls when the calculated enthalpy of solution is negative.

• Defining enthalpy of hydration without specifying that the ions start as gaseous.

Exam tips

• Write the cycle with gaseous ions at the top, solid at one corner and aqueous ions at the other, then follow the arrows.
• Count the ions from the formula before substituting; the doubling mark is the one most often lost.
• A negative enthalpy of solution means heat is released, so the temperature of the water rises.