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Lattice Enthalpy, Hydration Enthalpy And Enthalpy Of Solution - Practice Questions & MCQ

Edited By admin | Updated on Sep 18, 2023 18:35 AM | #JEE Main

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Lattice Enthalpy, Hydration Enthalpy And Enthalpy Of Solution is considered one of the most asked concept.
11 Questions around this concept.

Solve by difficulty

Given below are two statements : one is labelled as Assertion (A) and the other is labelled as Reason (R).

Assertion (A): $\mathrm{Cu}^{2+}$ in water is more stable than $\mathrm{Cu}^{+}$ .
Reason (R) : Enthalpy of hydration for $\mathrm{Cu}^{2+}$ is much less than that of $\mathrm{Cu}^{+}$ .

In the light of the above statements, choose the correct answer from the options given below :

Both (A) and (R) are correct and (R) is the correct explanation of (A)

(A) is not correct but (R) is correct

(A) is correct but (R) is not correct

Both (A) and (R) are correct but (R) is not the correct explanation of (A)

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Concepts Covered - 1

Lattice Enthalpy, Hydration Enthalpy And Enthalpy Of Solution

Lattice Enthalpy

The lattice enthalpy of an ionic compound is the enthalpy change which occurs when one mole of an ionic compound is formed from its ions in gaseous state.

$\mathrm{Na}^{+}(\mathrm{g})+\mathrm{Cl}^{-}(\mathrm{g}) \longrightarrow \mathrm{NaCl}(\mathrm{s}), \Delta \mathrm{H}=-788 \mathrm{kJmole}^{-1}$

Heat of Hydration

The enthalpy change during hydration of one mole of any gaseous ion is called heat of hydration.

$\begin{aligned} & \mathrm{Na}^{+}(\mathrm{g}) \longrightarrow \mathrm{Na}^{+}(\mathrm{aq}), \Delta \mathrm{H}=\Delta \mathrm{H}_{\mathrm{hyd}_{\mathrm{Na}^{+}}} \\ & \mathrm{Cl}^{-}(\mathrm{g}) \longrightarrow \mathrm{Cl}^{-}(\mathrm{aq}), \Delta \mathrm{H}=\Delta \mathrm{H}_{\mathrm{hyd}_{\mathrm{Cl}-}}\end{aligned}$

Heat of Solution

It is change in enthalpy when one mole of a solid solute is dissolved in excess of solvent.

$\mathrm{NaCl}(\mathrm{s}) \longrightarrow \mathrm{Na}^{+}(\mathrm{aq})+\mathrm{Cl}^{-}(\mathrm{aq}), \Delta \mathrm{H}=\Delta \mathrm{H}_{\mathrm{sol}_{\mathrm{NaCl}}}$

Solubility of an Ionic Compound in water

When an ionic compound dissolves in a solvent, the ions leave their ordered positions on the crystal lattice. These are now more free in solution. But solvation of these ions (hydration in case solvent is water) also occurs at the same time. The enthalpy of solution of any ionic solid, in water is, therefore, determined by the selective values of the lattice enthalpy and enthalpy of hydration of ions. This will be more clear with the help of the diagram given below:

Thus, the enthalpy of solution, enthalpy of hydration and the lattice energy can be related as

$\Delta_{\text {sol }} \mathrm{H}^0=\Delta_{\text {lattice }} \mathrm{H}^0+\Delta_{\text {hyd }} \mathrm{H}^0$

For most of the ionic compounds, $\Delta_{\text {sol }} \mathrm{H}^0$ is positive and the dissociation process is endothermic. Therefore the solubility of most salts in water increases with rise of temperature. If the lattice enthalpy is very high, the dissolution of the compound may not take place at all.