How Bonding Type Predicts Solubility

Stir a spoonful of table salt into water and it vanishes. Stir the same salt into cooking oil and it sits at the bottom forever. Swap the salt for candle wax and the result flips: the wax ignores the water but softens and disperses in warm oil. Sand ignores both. A lump of copper ignores both too, although it will happily dissolve into molten tin.

These everyday results are not random. Each one follows from the type of bonding in the solid and the type of forces in the liquid. Once you can sort a substance into one of four structure types, you can predict its solubility with surprising confidence.

Dissolving is a trade

Think of dissolving as a trade in which particles swap partners. Three things have to happen:

1. The particles of the solid must be pulled apart from each other. This costs energy.
2. The particles of the solvent must be pushed apart to make room. This also costs energy.
3. Solid particles and solvent particles attract each other in their new arrangement. This releases energy.

Dissolving goes well when step 3 pays back most or all of the cost of steps 1 and 2. (Entropy, the natural tendency to spread out and mix, also helps, which is why a small energy shortfall can still be overcome.) If the new attractions are weak compared with what was broken, the substance stays undissolved.

A useful analogy is a dance hall. Dancers already have partners. They will only swap if the new partner is at least as good as the old one. Water molecules are strongly attached to each other through hydrogen bonds; they will only let a newcomer in if the newcomer can offer a similar grip.

This is the physical reasoning behind the old slogan “like dissolves like”. Polar and ionic substances offer water the strong attractions it expects. Non-polar substances offer only weak ones, and do better in non-polar solvents whose own attractions are equally weak.

Type 1: Ionic compounds

In an ionic solid such as sodium chloride, positive and negative ions are held in a lattice by strong electrostatic attraction. Pulling the lattice apart costs a lot of energy; the size of that cost is the lattice energy.

Water can pay the bill because its molecules are polar. The slightly negative oxygen ends cluster around each cation, and the slightly positive hydrogen ends cluster around each anion. Each ion ends up wearing a shell of water molecules, a process called hydration, and the energy released is large.

So many ionic compounds dissolve in water, and essentially none dissolve in non-polar solvents like hexane or oil. A hexane molecule has no charged or partially charged ends, so it can offer an ion almost nothing in return for leaving its lattice.

But “ionic means soluble in water” is only a tendency. Calcium carbonate (limestone), silver chloride and barium sulfate are all ionic and all barely soluble. In these cases the lattice holds together more strongly than hydration can repay, often because the ions are small or highly charged. The detailed patterns are collected in the solubility rules, which exist precisely because bonding type alone cannot settle every ionic case.

Type 2: Simple molecular substances

Simple molecular substances are made of separate molecules with strong covalent bonds inside them and weak forces between them. When they dissolve, the covalent bonds stay intact. Only the intermolecular forces are rearranged.

Here the question is: which forces does the molecule offer?

• Molecules that can hydrogen bond (with O–H or N–H groups) usually dissolve well in water if they are small. Ethanol mixes with water in any proportion. Sugar dissolves in large amounts because each sucrose molecule is covered in O–H groups. See hydrogen bonding for why these attractions are unusually strong.
• Polar molecules without hydrogen bonding are moderately soluble, depending on size and shape.
• Non-polar molecules such as methane, iodine or the long hydrocarbon chains in wax and oil offer only weak London dispersion forces. Water will not give up its hydrogen bonds for them, so they barely dissolve in water. They do dissolve in non-polar solvents, where dispersion forces are all anyone has. Iodine, for instance, is only slightly soluble in water but dissolves readily in hexane, giving a violet solution.

Size matters as well as polarity. Methanol, ethanol and propanol mix completely with water. As the carbon chain grows, the non-polar part of the molecule takes over, and by the time you reach alcohols with long chains the O–H group can no longer carry the rest of the molecule into solution. A molecule is a mixture of polar and non-polar regions, and the bigger region usually wins.

A special case: some polar molecules react with water rather than simply dissolving. Hydrogen chloride gas dissolves enormously in water because it ionises to form hydrochloric acid. That is a chemical change as well as a physical one.

Type 3: Giant covalent structures

Diamond, graphite and silicon dioxide (quartz, sand) are single networks of atoms held together by covalent bonds in every direction. To dissolve them, you would have to break those strong covalent bonds, and no ordinary solvent offers anything comparable in return.

So the rule is simple: giant covalent substances are insoluble in all common solvents, polar or not. This is why sand sits unchanged at the bottom of the sea and why diamonds survive in river gravels for millions of years. (Silicon dioxide can be attacked by hydrofluoric acid or hot concentrated alkali, but that is a chemical reaction that breaks bonds, not dissolving in the ordinary sense.) See giant covalent structures for how these networks are built.

Type 4: Metals

Metals consist of positive ions in a sea of delocalised electrons. Neither water nor oil can offer anything that matches metallic bonding, so metals do not dissolve in ordinary solvents.

Two things can look like dissolving but are not:

• Reaction. Sodium in water disappears, but it is reacting to form sodium hydroxide and hydrogen. The sodium ends up as Na⁺ ions, a different substance.
• Acids. Zinc “dissolves” in hydrochloric acid only by being oxidised to Zn²⁺.

Metals do genuinely dissolve in other metals. Molten metals mix to form alloys, and mercury dissolves many metals at room temperature to form amalgams. Like dissolves like, once again: metallic bonding is happy to share its electron sea with another metal.

A quick prediction table

Structure type	Example	Water (polar)	Hexane or oil (non-polar)
Ionic	NaCl	Often soluble (check solubility rules)	Insoluble
Molecular, hydrogen bonding	Ethanol, sucrose	Soluble if small or with many O–H groups	Usually poor
Molecular, non-polar	I₂, CH₄, wax	Very slightly soluble	Soluble
Giant covalent	Diamond, SiO₂	Insoluble	Insoluble
Metallic	Cu, Fe	Insoluble (some react)	Insoluble

Common mistakes

“Dissolving breaks the bonds in the molecule.” For simple molecular substances it does not. When sugar dissolves, sucrose molecules are separated from each other, but every C–H, C–O and O–H bond stays intact. Evaporate the water and you get sugar back.

“All ionic compounds dissolve in water.” Many do not. Calcium carbonate, barium sulfate and silver chloride are classic insoluble salts. Bonding type tells you the solvent that has a chance, not that success is guaranteed.

“Polar molecules always dissolve in water.” Big molecules with one small polar group behave mostly as non-polar. Long-chain alcohols and fats are good examples.

“Metals dissolve in acid.” In exam language, say they react with acid. The metal is oxidised to ions and a new compound forms.

“Non-polar means insoluble.” Non-polar substances are insoluble in water. In non-polar solvents they may be very soluble. Always ask “soluble in what?”

Confusing dissolving with melting. Sugar in tea dissolves; it does not melt. Melting needs heat and involves one substance; dissolving involves a solvent.

Key takeaways

• Dissolving is a trade: attractions broken in the solid and solvent must be repaid by new solute–solvent attractions.
• Ionic compounds can dissolve in water through hydration, but many are insoluble; they never dissolve in non-polar solvents.
• Simple molecular substances dissolve without breaking covalent bonds; polar and hydrogen-bonding molecules suit water, non-polar molecules suit non-polar solvents, and molecular size shifts the balance.
• Giant covalent networks are insoluble in every common solvent because dissolving would mean breaking covalent bonds.
• Metals do not dissolve in water or oil; they may react with water or acids, and they dissolve in other metals.
• “Like dissolves like” is a summary of the energy trade, not a law; always say what the solvent is. For the matching story for heat, see how bonding type predicts melting point.