Water as a solvent

A solvent is the substance present in larger amount that does the dissolving; the solute is what gets dissolved; together they form a solution. Water is an exceptional solvent because each $H_2O$ molecule is a tiny electric dipole. The oxygen atom pulls shared electrons toward itself, leaving oxygen slightly negative ($\delta^-$) and each hydrogen slightly positive ($\delta^+$). When an ionic solid such as common salt is dropped into water, the negative oxygen ends cluster around the solid's positive ions and the positive hydrogen ends cluster around its negative ions. This pulling-apart action, repeated at every ion on the crystal surface, peels the lattice apart ion by ion until the whole solid has dispersed — the process of hydration. Polar molecular substances such as sugar dissolve by a similar attraction between water's dipole and the solute molecule's own polar groups, without any ions being involved at all. Not every mixture with water behaves the same way, because the categories depend on how large the dispersed particles are.

• True solution — particles smaller than about $1$ nanometre. They cannot be seen even under a microscope, never scatter a light beam, and never settle out. Salt water and sugar water are true solutions.
• Colloid — particles roughly $1$ to $1000$ nanometres. Too small to settle, but large enough to scatter light passing through them, an effect called the Tyndall effect. Milk and starch solution are colloids.
• Suspension — particles large enough to see with the naked eye, which settle out under gravity if left undisturbed. Chalk powder shaken into water is a suspension.

For a solid solute in water, the amount already dissolved relative to the maximum possible at that temperature defines three states: a saturated solution holds the maximum solute the solvent can take at that temperature, so any extra solute added stays undissolved; an unsaturated solution holds less than that maximum, so more solute can still go in; and a supersaturated solution holds more than the normal saturation point, an unstable state made by carefully cooling a hot saturated solution, which crystallises out rapidly the moment it is disturbed.

Solubility curves put numbers on this idea. The solubility of a substance at a given temperature is defined as the mass of solute that saturates a fixed mass of water:

where: $S$ is the solubility (grams of solute per $100$ g of water), $m_{\text{solute}}$ is the mass of solute needed to just saturate that water, and the temperature must always be stated because solubility changes with it. Plotting $S$ against temperature gives a solubility curve. Most solids, such as potassium nitrate ($KNO_3$) and sodium chloride ($NaCl$), show rising curves — $KNO_3$ rises steeply, $NaCl$ only gently. Calcium sulphate ($CaSO_4$) is the standard exception: its solubility is low to begin with and drifts slightly downward as temperature rises. Reading these curves tells you, at a chosen temperature, how much of each solute will dissolve and which of two solutes is more soluble — no calculation beyond reading a graph is expected at this level.

Water's ability to dissolve so many substances traces back entirely to its polarity, and it is worth walking through why that single property has such wide reach.

Ionic solids are pulled apart electrically. A crystal like sodium chloride is held together by the attraction between $Na^+$ and $Cl^-$ ions. When the crystal meets water, the negative oxygen ends of nearby water molecules are drawn to each $Na^+$ ion and the positive hydrogen ends are drawn to each $Cl^-$ ion. Enough water molecules cluster around a surface ion that their combined pull exceeds the pull holding that ion in the lattice, and the ion is drawn away into the liquid, surrounded by a shell of water molecules. Repeated over every ion, the crystal dissolves from the outside in.

Polar covalent solutes dissolve without ions. Sugar has no charged particles, but its molecule carries polar $-OH$ groups. Water's dipole forms similar attractions with these groups, so sugar molecules are pulled away from each other and surrounded by water molecules just as ions are — the mechanism is the same electrical attraction, only the solute contributes a polar group instead of a full charge.

Nonpolar substances are left out. Oil and other nonpolar molecules have no charge separation for water to grip. Water molecules would rather hold onto each other (through hydrogen bonds) than surround an oil molecule that offers nothing back, so oil and water do not mix — the general rule chemists summarise as "like dissolves like." This is also why gases with polar or reactive molecules, such as ammonia or hydrogen chloride, dissolve readily in water while a nonpolar gas dissolves only sparingly.

Temperature changes how much fits. Raising the temperature usually gives solute particles and water molecules more kinetic energy, which for most solids helps more solute separate and disperse, raising solubility — the rising branch of the $KNO_3$ and $NaCl$ curves. $CaSO_4$'s slightly falling curve is the exception the syllabus flags rather than explains in detail; the general trend, and the existence of exceptions, is what you need to know.

Question: A student keeps adding sugar to a glass of water and stirring. Beyond a point, extra sugar grains stop dissolving and settle at the bottom no matter how long the stirring continues. (a) What term describes this solution? (b) The student gently warms the glass. What happens to the undissolved sugar, and why? (c) The student then repeats the experiment with sand instead of sugar, stirring for the same length of time. Sand never disappears and always settles when stirring stops — classify muddy water of this kind, and explain how it differs from the sugar solution in terms of particle size.

Step 1. Once the water can hold no more dissolved sugar and excess solid remains undissolved at the bottom, the solution has reached its maximum solute capacity for that temperature — this is a saturated solution by definition.

Step 2. For most solids, including sugar, solubility rises with temperature — the rising solubility curve behaviour described above. Warming lets the water hold more dissolved sugar, so some or all of the remaining grains dissolve, moving the solution from saturated back toward unsaturated.

Step 3. Sand particles are far larger than dissolved sugar molecules — large enough to be seen and to settle under gravity once stirring stops. This is a suspension, not a solution: sugar particles are below about $1$ nanometre and stay dispersed indefinitely, while sand particles are large, visible, and always settle given time. Answer: (a) saturated solution; (b) some or all of the sugar dissolves because solubility increases with temperature; (c) muddy water is a suspension, distinguished from the true sugar solution by its much larger, visible, settling particles.

Competitive stretch

If the same glass of muddy water were left to stand and, instead of settling, the water stayed permanently and uniformly cloudy while scattering a visible torch beam through it, the mixture would instead be classified as a colloid — the Tyndall effect (light scattering without settling) is the deciding test, separating a colloid from both a true solution (no scattering, no settling) and a suspension (no scattering test needed, since it visibly settles).

• Confusing a colloid with a suspension: colloids do not settle on standing, while suspensions do — the Tyndall effect, not settling, is the defining colloid test.
• Assuming solubility always rises with temperature for every solute — calcium sulphate is the standard counter-example, with solubility falling slightly as temperature rises.
• Believing a saturated solution can never dissolve any more solute under any condition — raising the temperature can let it dissolve more, for solutes whose solubility increases with heat.
• Treating "concentrated" and "saturated" as the same word — a solution can be concentrated (a large amount of solute) without being saturated (unable to dissolve any more); the two describe different things.

Ask the tutor: want to walk through a few more real-life examples of colloids, suspensions and true solutions, or see how to read an unfamiliar solubility curve?