the solubility of ionic compounds and covalent molecular substances in water and in non-polar solvents, explained in terms of bond polarity, intermolecular forces and the energy changes (including hydration enthalpy) associated with dissolving, and the formation of saturated and unsaturated solutions

What this dot point is asking

VCAA wants you to explain why some substances dissolve in water and others do not, using the polarity of the solvent, the bonding inside the solute, and the energy changes of dissolving (lattice enthalpy breaking ionic lattices, hydration enthalpy released when ions are surrounded by water). You also need to know the like dissolves like rule and the difference between saturated, unsaturated and supersaturated solutions.

The answer

Solute, solvent, solution

A solution is a homogeneous mixture. The solvent is the component in larger amount (often the liquid); the solute is what is dissolved. For most of Unit 1, the solvent is water and we are asking whether the solute will dissolve in it.

Why water dissolves ionic compounds

Water is a bent, polar molecule. Each O has two lone pairs and is partially negative ($\delta^{-}$); each H is partially positive ($\delta^{+}$).

When an ionic solid like NaCl is added:

1. Water molecules cluster around each surface ion, with $\delta^{-}$ O atoms pointing at $Na^{+}$ cations and $\delta^{+}$ H atoms pointing at $Cl^{-}$ anions. These are ion-dipole attractions.
2. The combined pull of many water dipoles is strong enough to peel ions off the lattice.
3. Each separated ion becomes surrounded by a hydration shell of oriented water molecules and is now an aqueous ion.

The energy released as ion-dipole bonds form is the hydration enthalpy ($\Delta H_{hyd}$, negative). The energy needed to break apart the ionic lattice is the lattice enthalpy ($\Delta H_{lat}$, positive). The overall enthalpy of dissolving is

If $\Delta H_{hyd}$ is large enough (more negative) compared to $\Delta H_{lat}$, dissolving is energetically favourable. For NaCl the two are similar in magnitude and the overall enthalpy of dissolving is close to zero, but the entropy gain from spreading ions through water makes the process spontaneous.

Why some ionic compounds are insoluble

Not all ionic compounds dissolve. Compounds with very large lattice enthalpies relative to their hydration enthalpies (silver chloride AgCl, barium sulfate $BaSO_4$, calcium carbonate $CaCO_3$) do not dissolve significantly in water; the cost of breaking the lattice exceeds the energy returned by hydration.

Why water dissolves polar molecular substances

Polar molecules like ethanol, glucose and ammonia have $\delta^{+}$ and $\delta^{-}$ regions of their own. Water can form dipole-dipole or hydrogen-bond interactions with them, similar in strength to the water-water hydrogen bonds they replace. So they mix freely.

Why water does not dissolve non-polar substances

Non-polar substances (oils, $I_2$, hexane, $O_2$) have only dispersion forces. To dissolve in water, the strong water-water hydrogen-bond network would have to be broken and replaced with much weaker water-solute dispersion. That is energetically very unfavourable, so non-polar substances stay in their own non-polar phase.

Non-polar solvents

A non-polar solvent (hexane, toluene, cyclohexane) interacts with its solutes only by dispersion. Adding a non-polar solute swaps dispersion for dispersion and dissolves; adding a polar or ionic solute would require breaking strong solute-solute forces and replacing them with much weaker dispersion, so they do not dissolve. This is why oil-and-water emulsions separate but oil and petrol mix.

Like dissolves like

A short summary you can quote in any solubility question:

• Polar solvent + polar/ionic solute: dissolves.
• Non-polar solvent + non-polar solute: dissolves.
• Polar solvent + non-polar solute: does not dissolve.
• Non-polar solvent + ionic/polar solute: does not dissolve.

Saturated, unsaturated, supersaturated

At a given temperature, a fixed mass of solvent has a maximum amount of solute it will hold. That maximum is the solubility (often in g per 100 g water).

• Unsaturated: less solute than the solubility limit. More solute would still dissolve.
• Saturated: solute equals the solubility limit. Any extra solute remains undissolved at the bottom of the container, in dynamic equilibrium with dissolved solute.
• Supersaturated: more solute than the solubility limit, achieved by cooling a hot saturated solution gently. Highly unstable; a seed crystal causes the excess to crystallise out.

Solubility of most solids in water increases with temperature; solubility of most gases decreases with temperature.

Worked example

Predict whether each of the following will dissolve in water and justify:

• $KNO_3$: ionic, both ions are well hydrated. Yes, dissolves.
• $CH_4$: non-polar, dispersion only. No, essentially insoluble.
• $CH_3OH$ (methanol): polar, hydrogen bonds with water. Yes, miscible.
• $CaCO_3$: ionic but lattice enthalpy outweighs hydration enthalpy. No, only very slightly soluble.

Common traps

Saying NaCl "breaks the covalent bonds in water"

Dissolving NaCl breaks the ionic bonds in the NaCl lattice and forms new ion-dipole attractions; the O-H bonds in water are untouched.

Calling NaCl(aq) a covalent solution

$Na^{+}(aq)$ and $Cl^{-}(aq)$ are dissociated free ions surrounded by hydration shells, not NaCl molecules.

Forgetting entropy

Some dissolutions are slightly endothermic (e.g. $NH_4NO_3$, used in instant cold packs) and still happen because the entropy of mixing outweighs the small unfavourable enthalpy.

Using "miscible" for solids

Miscible means two liquids mix in any proportion (ethanol and water). Use "soluble" for solids.

Mixing up saturated and concentrated

Saturated means at the solubility limit. Concentrated just means a lot of solute relative to solvent; a small-solubility salt can have a saturated solution that is still very dilute.

In one sentence

A solute dissolves when the new solute-solvent interactions (ion-dipole for ionic solutes in water, dipole-dipole or hydrogen bonding for polar molecular solutes, dispersion for non-polar solutes in non-polar solvents) return enough energy and entropy to make up for breaking the original solute-solute and solvent-solvent forces, which is the molecular basis of the like dissolves like rule.

Examples in context

Example 1. Salinity in the Murray-Darling Basin. The Murray-Darling Basin Authority monitors river salinity, which often exceeds the $800 \, \mu\text{S/cm}$ threshold below Mildura in dry years. The salts (mostly $\text{NaCl}$ but with $\text{CaSO}_4$ and $\text{MgSO}_4$) dissolve from ancient marine sediments. Sodium chloride is fully soluble (around $360 \, \text{g/L}$ at $25^{\circ}\text{C}$) because the hydration enthalpies of $\text{Na}^{+}$ and $\text{Cl}^{-}$ approximately offset the lattice enthalpy of $787 \, \text{kJ}$ per mole. Calcium sulfate (gypsum) is much less soluble (about $2.4 \, \text{g/L}$) because $\text{Ca}^{2+}$ and $\text{SO}_4^{2-}$ have high charges and strong lattice attractions that water cannot fully overcome. Below the gypsum saturation limit, irrigators see white scaling on drip lines.

Example 2. Oil-spill response in Port Phillip Bay. The Australian Maritime Safety Authority deploys dispersants such as Corexit when a ship's bunker fuel spills near Williamstown. Crude-oil hydrocarbons are non-polar molecules held by dispersion forces only and are essentially insoluble in water (solubility under $1 \, \text{mg/L}$). The dispersant contains an amphiphilic surfactant: its hydrocarbon tail dissolves in oil ($\text{like dissolves like}$), while the polar sulfonate head dissolves in water. The surfactant breaks oil films into micron-sized droplets that bacteria can digest. Without surfactants, the oil floats indefinitely because solute-solute dispersion forces between oil molecules vastly exceed weak oil-water interactions.

Try this

Q1. Explain why iodine ($\text{I}_2$) is more soluble in hexane than in water. [2 marks]

• Cue. $\text{I}_2$ is non-polar; only dispersion forces. Hexane non-polar matches. Water polar; energy cost to break water hydrogen bonds is not repaid. Like dissolves like.

Q2. A solution is prepared by dissolving $11.7 \, \text{g}$ of $\text{NaCl}$ in $250 \, \text{mL}$ of water. Calculate (a) the molar concentration and (b) the concentration of $\text{Na}^+$ ions in $\text{mol/L}$. [3 marks]

• Cue. (a) $n = 11.7 / 58.5 = 0.200$ mol; $c = 0.200 / 0.250 = 0.800 \, \text{mol/L}$. (b) NaCl dissociates fully; $[\text{Na}^+] = 0.800 \, \text{mol/L}$.

Q3. Compare the solubility behaviour of glucose, sodium chloride and naphthalene in water. (a) Identify the dominant solute-solvent interaction in each case. (b) Rank the three by approximate solubility in water. (c) Explain the ranking. [2+2+2 marks]

• Cue. (a) Glucose: H-bonding through OH groups. NaCl: ion-dipole. Naphthalene: only weak dispersion with induced dipole. (b) NaCl > glucose > naphthalene. (c) NaCl fully dissociates with strong hydration; glucose very soluble via H-bonds; naphthalene almost insoluble (no polar groups).