9701 · 23.2
Enthalpies of solution and hydration flashcards
Revision flashcards for Cambridge 9701 Enthalpies of solution and hydration (syllabus 23.2). Flip, recall, then mark a real past-paper question.
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What is the standard enthalpy change of solution, $\Delta H_{\text{sol}}^{\ominus}$?
The enthalpy change when one mole of an ionic solid dissolves in a sufficient amount of solvent to form an infinitely dilute solution, under standard conditions (298 K and 100 kPa).
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What is the standard enthalpy change of hydration, $\Delta H_{\text{hyd}}^{\ominus}$?
The enthalpy change when one mole of gaseous ions is completely hydrated by a sufficient amount of water, under standard conditions.
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Is the enthalpy of hydration, $\Delta H_{\text{hyd}}^{\ominus}$, endothermic or exothermic?
Always exothermic (negative value). Energy is released when ion-dipole attractions form between the gaseous ions and water molecules.
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What process does the standard lattice enthalpy, $\Delta H_{\text{latt}}^{\ominus}$, represent?
The enthalpy change when one mole of an ionic solid is formed from its gaseous ions under standard conditions. It is always exothermic.
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What is the key equation linking $\Delta H_{\text{sol}}^{\ominus}$, $\Delta H_{\text{latt}}^{\ominus}$, and $\Delta H_{\text{hyd}}^{\ominus}$?
$\Delta H_{\text{sol}}^{\ominus} = -\Delta H_{\text{latt}}^{\ominus} + \sum \Delta H_{\text{hyd}}^{\ominus}$. Note the negative sign for lattice enthalpy, as we are breaking the lattice (endothermic process).
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Why is the process of breaking the lattice, represented by $-\Delta H_{\text{latt}}^{\ominus}$, always endothermic?
Energy must be supplied to overcome the strong electrostatic forces of attraction between the oppositely charged ions in the crystal lattice.
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How does ionic charge affect the enthalpy of hydration?
A greater ionic charge leads to a stronger electrostatic attraction to water dipoles, resulting in a more exothermic (more negative) enthalpy of hydration. E.g., Mg²⁺ has a much more negative $\Delta H_{\text{hyd}}^{\ominus}$ than Na⁺.
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How does ionic radius affect the enthalpy of hydration?
A smaller ionic radius leads to a higher charge density, resulting in a stronger attraction to water dipoles and a more exothermic (more negative) enthalpy of hydration. E.g., F⁻ has a more negative $\Delta H_{\text{hyd}}^{\ominus}$ than Cl⁻.
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If $\Delta H_{\text{sol}}^{\ominus}$ is positive (endothermic), does this mean the salt will not dissolve?
Not necessarily. Spontaneity of dissolving also depends on the entropy change. A large positive entropy change can overcome a small positive enthalpy change, making the process feasible.
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What is a common mistake when using lattice enthalpy in a dissolution cycle?
Forgetting to reverse the sign. Standard lattice enthalpy ($ΔH_{\text{latt}}^{\ominus}$) is for the formation of the lattice (exothermic). The cycle for solution involves breaking the lattice, which is an endothermic process with a value of $-ΔH_{\text{latt}}^{\ominus}$.