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How can the versatility of non-metals be explained?

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

A focused VCE Chemistry Unit 1 answer on solubility. Covers the dissolution of ionic compounds in water (hydration shells and hydration enthalpy), why polar solvents dissolve polar solutes and non-polar solvents dissolve non-polar solutes (like dissolves like), and the difference between saturated, unsaturated and supersaturated solutions.

Generated by Claude OpusReviewed by Better Tuition Academy8 min answerUpdated 2026-05-19

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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:

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.
The combined pull of many water dipoles is strong enough to peel ions off the lattice.
Each separated ion becomes surrounded by a hydration shell of oriented water molecules and is now an aqueous ion.

NaCl(s) \rightarrow Na^{+}(aq) + Cl^{-}(aq)

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

\Delta H_{sol} = \Delta H_{lat} + \Delta H_{hyd}

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:

IMATH_18 : ionic, both ions are well hydrated. Yes, dissolves.
IMATH_19 : non-polar, dispersion only. No, essentially insoluble.
IMATH_20 (methanol): polar, hydrogen bonds with water. Yes, miscible.
IMATH_21 : 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.

Past exam questions, worked

Real questions from past VCAA papers on this dot point, with our answer explainer.

2024 VCE SAC-style5 marksSodium chloride dissolves readily in water but not in hexane, while iodine ($I_2$) dissolves in hexane but barely in water. (a) Explain in terms of bonding and intermolecular forces why sodium chloride dissolves in water. (b) Explain why iodine prefers hexane over water. (c) State the general rule these two observations illustrate.

Show worked answer →

A 5-mark answer needs the dissolution mechanism for each solute, the role of solvent polarity, and the general rule.

(a) NaCl is an ionic compound: a 3D lattice of $Na^{+}$ and $Cl^{-}$ ions held by strong electrostatic attraction. Water is a polar molecule with a partially negative O and partially positive H. When NaCl is added, the partially negative O of water molecules orients towards the $Na^{+}$ ions and the partially positive H towards the $Cl^{-}$ ions, forming ion-dipole attractions. These ion-dipole interactions release energy (the hydration enthalpy); when the energy released on hydration exceeds the lattice enthalpy needed to separate the ions, the salt dissolves. The ions become surrounded by hydration shells of water molecules and dissociate into solution as aqueous ions: $NaCl(s) \rightarrow Na^{+}(aq) + Cl^{-}(aq)$ .

(b) $I_2$ is a non-polar covalent molecular substance held in the solid by dispersion forces. Water cannot solvate it well: water-water hydrogen bonds are far stronger than any water- $I_2$ attraction, so to dissolve $I_2$ would require breaking hydrogen bonds and replacing them with much weaker dispersion forces, which is unfavourable. Hexane is also non-polar; its molecules attract one another and $I_2$ only by dispersion. Mixing $I_2$ with hexane swaps similar-strength dispersion forces for similar-strength dispersion forces, with the entropy increase of mixing making the process favourable. So $I_2$ dissolves in hexane.

(c) Like dissolves like. Polar and ionic solutes dissolve in polar solvents; non-polar solutes dissolve in non-polar solvents.