An electrovalent compound is insoluble in water when its lattice enthalpy is smaller in ma… — JAMB UTME

An electrovalent compound is insoluble in water when its lattice enthalpy is smaller in magnitude than the total hydration enthalpy of its constituent gaseous ions.

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False. An electrovalent compound is soluble in water when the magnitude of its hydration enthalpy exceeds its lattice enthalpy, allowing ion-water electrostatic attractions to overcome ionic crystal lattice forces.

The statement is false because for an electrovalent compound to dissolve in water, the hydration energy released when ions interact with water molecules must overcome the lattice energy holding the crystal together. If hydration enthalpy is greater in magnitude than lattice enthalpy, the compound is soluble rather than insoluble.

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Identify the enthalpy changes during the dissolution of an ionic solid.

Dissolution depends on two key thermodynamic quantities: lattice enthalpy (energy required to separate solid ions into gaseous ions) and hydration enthalpy (energy released when gaseous ions are solvated by water).

The overall enthalpy of solution is approximated by

\Delta H_{\text{solution}} = \Delta H_{\text{lattice}} + \Delta H_{\text{hydration}}

Evaluate the condition where

|\Delta H_{\text{lattice}}| < |\Delta H_{\text{hydration}}|

The energy released during ion hydration is greater than the energy required to break the ionic lattice.

This leads to an exothermic dissolution process (

\Delta H_{\text{solution}} < 0

), which strongly favors the compound dissolving in water.

Determine the truth value of the statement.

The statement asserts that such a compound is insoluble, which contradicts chemical thermodynamic principles.

Therefore, the given statement is false.

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Lattice Enthalpy vs Hydration Enthalpy in Ionic Compound Solubility

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