The sorting you did at the start of this unit had an awkward result. Salt dissolves in water and not in oil. Sugar dissolves in water. Wax dissolves in oil and not in water. Iodine barely touches water and colours oil deeply. Nothing about a substance’s appearance predicts which column it lands in — and “water dissolves everything” turns out to be one of the most confidently held wrong ideas in chemistry.
The pattern is real and it has a single explanation, and the explanation starts with the shape of one very small molecule.
Water is bent, and that changes everything
Two facts about , both from Lewis Structures and Models:
Oxygen is much more electronegative than hydrogen, so each bond is strongly polar — the shared pair spends more time near the oxygen, leaving the oxygen and each hydrogen .
And the molecule is bent, at roughly 105°, because the oxygen also carries two lone pairs that take up room.
Put those together and the two bond dipoles do not cancel. Water has a definite negative end and a definite positive end: it is a polar molecule. Compare it with carbon dioxide, which also has two strongly polar bonds — but is linear, the two pulls point in exactly opposite directions, and they sum to nothing. Same kind of bonds, opposite result, purely because of shape.
Being polar lets water do something further. When a hydrogen is bonded to oxygen (or to nitrogen or fluorine — small, very electronegative atoms), the hydrogen’s own electron is pulled so far off that its nucleus is left almost bare on the outside of the molecule. That exposed proton is strongly attracted to a lone pair on a neighbouring molecule. The result is a hydrogen bond: not a covalent bond, but by far the strongest kind of attraction between molecules.
Water is unusual in having two hydrogens to donate and two lone pairs to accept, so each molecule can hold four neighbours at once, building a network rather than pairs. Almost every strange property of water comes from that network:
- It boils at 100 °C. A molecule that small, without hydrogen bonding, would be a gas well below room temperature — hydrogen sulfide, which is heavier, boils at about −60 °C.
- Ice floats. The network locks into an open hexagonal arrangement that holds the molecules further apart than in the liquid, so the solid is less dense — which is why lakes freeze from the top and things survive underneath.
- It takes a lot of energy to warm up, which is why large lakes moderate the climate of the land beside them.
- It pulls itself into drops and supports insects on its surface.
How something dissolves
Dissolving is not disappearing. It is a competition: the attractions holding the solute together, against the attractions the solvent can offer instead.
An ionic solid in water. The lattice is held by strong attractions between ions, so this should not work — and often it does not. What makes it possible is that water molecules can surround each ion with their appropriate end pointing inwards: negative oxygen ends towards a cation, positive hydrogen ends towards an anion. Those ion–dipole attractions release energy, and if enough is released to compensate for pulling the lattice apart, the solid dissolves and every ion ends up wrapped in a shell of water molecules. If not enough is released, it does not. That is precisely why some ionic compounds dissolve freely and others sit at the bottom of the beaker, which is the pattern catalogued in Solubility Rules and exploited in Precipitation and Net Ionic Equations.
A polar molecular solid in water. Sugar has many groups, so it makes hydrogen bonds with water directly. Ethanol does too, which is why it mixes with water in any proportion. No ions are produced, which is why sugar solution does not conduct electricity and salt solution does — a test you can do in ten seconds and which settles what kind of substance you are holding.
Non-polar in non-polar. Grease dissolves in oil. Both are held together only by weak dispersion forces, and swapping one weak attraction for another costs nothing.
Non-polar in water: no. This is the case worth explaining properly, because the usual account is backwards. Oil does not dissolve in water not because water molecules dislike oil, but because water molecules like each other so much. To make room for an oil molecule, the hydrogen-bond network has to be broken open, and the oil has nothing to offer in return but feeble dispersion forces. The network closes back up and squeezes the oil out.
So like dissolves like is a fair summary, as long as you can say what “like” means: polar dissolves polar and ionic, non-polar dissolves non-polar, and the reason is always which attractions are being traded for which.
Oxygen gas is the useful exception to keep in mind. is completely non-polar, yet a small amount does dissolve in water — enough for every fish in the lake, and no more. Small non-polar molecules can be accommodated in the network without breaking too much of it.
Temperature and pressure
Solubility is not a fixed property. It has conditions attached, and the two that matter behave differently for solids and gases.
| Solute | Raise the temperature | Raise the pressure above the solution |
|---|---|---|
| Most solids | more dissolves | no measurable effect |
| Gases | less dissolves | more dissolves, in proportion |
Heating helps a solid dissolve because the added energy helps break the lattice apart, and because faster-moving solvent molecules carry the pieces away more effectively. This is why you make a hot drink with hot water and why a supersaturated solution can be made by cooling a hot saturated one carefully.
“Most” is doing honest work in that table. Sodium chloride is barely more soluble at 100 °C than at 20 °C, and a few compounds — cerium(III) sulfate is the standard example — become less soluble as temperature rises. Read the curve rather than assuming it.
Gases run the other way, and the reason follows from what dissolving is for a gas: a gas molecule in solution has been captured by the solvent, and heating gives it the energy to escape. Two consequences you can observe:
- Bubbles appear in a pan of water long before it boils. That is dissolved air coming out of solution as the temperature rises.
- A lake warmed by a hot summer, or by water discharged from an industrial process, holds less dissolved oxygen — at exactly the time when the fish in it are more active and need more. Thermal pollution kills without adding a single molecule of anything toxic. That is a thread worth following into The Water Report.
Pressure barely touches solids and liquids, because they are already about as close-packed as they get. For gases it is direct: double the pressure of the gas above the liquid and roughly twice as much dissolves. Sealing a carbonated drink under pressurised carbon dioxide is that principle; opening the can, and watching it fizz, is that principle running backwards.
Saturated, and what a solubility curve tells you
A solution is saturated when it holds as much solute as it can at that temperature, with any excess sitting undissolved at the bottom. Below that it is unsaturated. A supersaturated solution holds more than it should — made by cooling a saturated solution slowly and undisturbed — and it will dump the excess suddenly if you drop a crystal into it.
Because solubility depends on temperature, it is published as a curve rather than a number: grams of solute per 100 g of water, against temperature. Reading one is a skill of its own, and Reading a Data Table is worth reviewing before you are asked to.
Next: once you know something dissolves, the question becomes how much of it is in there. That is Concentration, and it is where this unit becomes quantitative.
Curriculum connection
E3.1
describe the properties of water (e.g., polarity, hydrogen bonding), and explain why these properties make water such a good solvent
Link to original
E3.2
explain the process of formation for solutions that are produced by dissolving ionic and molecular compounds (e.g., salt, oxygen) in water, and for solutions that are produced by dissolving non-polar solutes in non-polar solvents (e.g., grease in vegetable oil)
Link to original
E3.3
explain the effects of changes in temperature and pressure on the solubility of solids, liquids, and gases (e.g., explain how a change in temperature or atmospheric pressure affects the solubility of oxygen in lake water)
Link to original