In solid `NaCl`, the `Na^(+)` and `Cl^(-)` ions are held together by strong positive-negative electrostatic forces, but when a small crystal of NaCl dissolves in water, the three-dimensional network of ions breaks into its individual units. The separated `Na^(+)` and `Cl^(-)` ions are stabilized in solution by their interaction with water molecules. These ions are said to be hydrated.
Here water plays a role similar to that of a good electrical insulator. Water molecules shield the ions `(Na^(+)` and `Cl^(-))` from each other and effectively reduce the electrostatic attraction that held them together in the solid state.
The heat of solution is defined by the following process:
`NaCl(s)overset(H_(2)O)(rarr) Na^(+)(aq.)+Cl^(-)(aq.)`
Dissolving an ionic compound such as `NaCl` in water involves complex interactions among the solute and solvent species. However, for the sake of analysis, we can imagine the solution process takes place in two separate steps.
First, the `Na^(+)` and `Cl^(-)` ions in the solid crystal are separated from each other in a gaseous state:
`"Energy" + NaCl(s) rarr Na^(+)(g) + Cl^(-)(g)`
The energy required for this process is called lattice enthalphy (U), the energy required to completely separate one mole of a solid ionic compound into gaseous ions. The lattice enthalpy of `NaCl` is `788 kJ mol^(-1)`. In oyher words, we would need to supply `788 kJ` of energy to break 1 mole of solid `NaCl` into 1 mole of `Na^(+)` ions and 1 mol of `Cl^(-)` ions in gas phase. Here, we have define lattice enthalpy as a positive quantity. Next, the "gaseous" `Na^(+)` and `Cl^(-)` ions enter the water and become hydrated:
`Na^(+)(g)+Cl^(-)(g)overset(H_(2)O)(rarr)Na^(+)(aq.)+Cl^(-)(aq.)+"energy"`
The enthalpy change associated with the hydration process is called the enthalpy of hydration `(DeltaH_(hyd))`. It is a negative quantity for cations and anions.
Applying Hess's law, it is possible to consider `DeltaH_("soln")` as the sum of two related quanitities, lattice enthalpy (U) and anthalpy of hydration `(DeltaH_(hyd))`:
`NaCl(s) rarr Na^(+)(g)+Cl^(-)(g) U=788 kJ`
`Na^(+)(g)+Cl^(-)(g)overset(H_(2)O)(rarr)Na^(+)(aq.)+Cl^(-)(aq.)DeltaH_(hyd)= -784 kJ`
`NaCl(s)overset(H_(2)O)(rarr) Na^(+)(aq.)+Cl^(-)(aq.)DeltaH_("soln")=4 kJ`
Therefore, when `NaCl` dissolves in water, `4 kJ` of heat will be absorbed from the immediate surroundings. We would observe this effect by noting that the beaker containing the solution becomes slightly colder. Depending on the nature of the cation and anion involved, `DeltaH_("soln")` for an ionic compound may be positive (endothermic) or negative (exothermic).
In general. we predict that ionic compounds should be much more soluble in polar solvents such as water, liquid `NH_(3)`, and liquid HF, than in nonpolar solvents, such as benzene and `CCl_(4)`. Since the molecules of nonpolar solvents lack a dipole moment, they cannot effectively solvate the `Na^(+)` and `Cl^(-)` ions. (Solvation is the process in which an ion or a molecule is surrounded by solvent molecules arranged in a specific manner. When the solvent is water, the process is called hydration.)
The predominant intermolecular interaction between ions and nonpolar molecules is ion-induced dipole interaction, which is much weaker than ion-dipole interaction. Consequently, ionic compounds usually have extremely law solubility in nonpolar solvents.
