# Electrolytic Solutions

Ionic compounds dissolve in water if the energy released by hydration is similar in magnitude to the energy holding the ionic crystal together.

An electrolyte is a compound that does not just dissolve into individual molecules but dissociates into separate ions in water. Electrolytes are so named because the dissociated ions can conduct electricity. Glucose, C6H12O6, is a nonpolar molecule that dissolves in water but does not dissociate, and a solution of glucose in water with no impurities will not conduct electricity. Sodium chloride, NaCl, dissociates into Na+ and Cl ions and can conduct electricity. An ionic compound that dissociates completely into its component ions in solution is a strong electrolyte. An ionic compound that only partially dissociates into its component ions in solution is a weak electrolyte.

Ionic compounds have a large exothermic hydration enthalpy. This is because water molecules are polar and therefore have large attractive forces to both ions in the compound: the positive side (hydrogen atoms) of the water molecule attracts the anion, and the negative side (oxygen atoms) attracts the cation. The oxygen atom is more electronegative than the hydrogen atoms, creating a dipole (molecule with partial positive and negative charges). Hydrating ions is an exothermic process. The electrostatic interaction between ions and polar ends of molecules with permanent dipoles is called an ion-dipole interaction.

#### Hydrated Ions

Because the energy needed to dissociate an ionic compound into its composite ions is equal to the energy needed for a solid compound to transition into gaseous ions, the forming of an ionic solution can be imagined as taking place in the following three steps. M+ represents a metal cation, A represents an anion, and MA is the ionic compound formed by them.

1. Ionic compound dissociates: ${\rm{M}}_m{\rm{A}}_a(s)\rightarrow m{\rm{M}}^+(g)+a{\rm{A}}^-(g),\;\;\;\;\;\;\;\;\Delta H_1>0$

2. Cation becomes hydrated: ${\rm{M}}^+(g)+{\rm{H}}_2{\rm{O}}\rightarrow {\rm{M}}^+(aq),\;\;\;\;\;\;\;\;\Delta H_2<0$

3. Anion becomes hydrated: ${\rm{A}}^-(g)+{\rm{H}}_2{\rm{O}}\rightarrow {\rm{A}}^-(aq),\;\;\;\;\;\;\;\;\Delta H_3<0$

An input of energy is required to break ionic bonds, so $\Delta H_1>0$. However, the hydration energy of ionic compounds releases energy, so $\Delta H_2<0$ and $\Delta H_3<0$. $\Delta H_{\rm{soln}}$ is the sum of the enthalpies of the three steps.
$\Delta H_{\rm{soln}}=\Delta H_1+\Delta H_2+\Delta H_3$
If the total hydration energy for an ionic compound, $\Delta H_2+\Delta H_3$, were equal in magnitude to the energy needed to dissociate the ions, $\Delta H_1$, the hydration would be exothermic, and $\Delta H_{\rm{soln}}$ would be zero.
\begin{aligned}\Delta H_{\rm{soln}}&=\Delta H_1+(\Delta H_2+\Delta H_3)\\&\approx\Delta H_1+(-\Delta H_1)\\&=0\end{aligned}
This is $\Delta H_{\rm{soln}}$ of an ideal solution. In reality the hydration enthalpies of most soluble ionic compounds are slightly larger than the dissociation energy, so the net $\Delta H_{\rm{soln}}$ for ionic solutions is small but endothermic.