3: Solubility

At a fixed temperature, a saturated solution already contains the maximum amount of solute that can remain dissolved under those conditions. Adding more solute cannot permanently increase that amount; after any brief dissolution, the excess remains as solid and dynamic equilibrium is restored. The solvent does not become a gas.

Solubility is a quantitative property: it is the maximum amount of a solute that dissolves in a specified amount of solvent at a specified temperature, usually forming a saturated solution. The temperature and the reference amount of solvent are essential. Colour, shape, and total volume are not the definition.

Henry's law gives p = K_H x, or x = p/K_H. At constant temperature, K_H does not change. If the original pressure is p and the new pressure is 2p, the new mole fraction is x' = 2p/K_H = 2x. Thus the mole fraction doubles, so option A is correct. Four times would require pressure to become four times, while C reverses the proportionality.

Solids and liquids have very small volume changes under ordinary pressure changes because they are nearly incompressible. Consequently, pressure usually has a negligible effect on the solubility of a solid in a liquid. Gas solubility is much more pressure-sensitive, so the strong pressure rule should not be applied indiscriminately to solids.

An unsaturated solution contains less dissolved solute than the maximum amount possible at the stated temperature and solvent quantity. Because capacity remains, additional solute can still dissolve. Adding more than the limit cannot normally produce a stable unsaturated solution, while evaporation reduces solvent and may instead promote saturation.

For a fair comparison, temperature and the amount of solvent must be the same. The solid that can dissolve in a greater amount before saturation is reached has the greater solubility. Colour, particle size, and settling behaviour do not define equilibrium solubility; particle size mainly affects how fast dissolution occurs.

In contact with excess undissolved solute, a saturated solution reaches dynamic equilibrium: the rate of dissolution equals the rate of crystallisation. The concentrations remain constant even though particles continue moving. An unsaturated solution has no such equilibrium with excess solid, and “ideal” or “volatile” does not describe this solubility condition.

From Henry's law, p = K_H x, so x = p/K_H. Let the initial mole fraction be x = p/K_H. If pressure changes to 3p while K_H remains constant, the new mole fraction is x' = 3p/K_H = 3x. Therefore it becomes three times, making A correct. One third gives the wrong inverse trend, nine times uses an unjustified square, and unchanged ignores the pressure effect.

Henry's law specifically describes the solubility of a gas in a liquid: at constant temperature, the dissolved gas concentration or mole fraction is proportional to the gas's partial pressure above the solution. Therefore Henry's law is the appropriate choice. Conservation of mass is a general accounting principle, the law of multiple proportions concerns compound formation, and the periodic law concerns trends among elements.

Whether more solute can dissolve depends on how much is already present compared with the solubility limit at the actual temperature and solvent amount. If the present amount is below that limit, the solution is unsaturated and more may dissolve. Colour, container shape, and trade name do not determine the equilibrium limit.

Polar water can form strong ion–dipole interactions with ions and dipole–dipole or hydrogen-bond interactions with suitable polar molecules. These interactions can compensate for the forces holding the solute together. This is the basis of “like dissolves like,” although actual solubility also depends on molecular structure and energy changes.

Water is strongly polar and forms an extensive hydrogen-bond network. Most oils consist mainly of non-polar molecules, which cannot make comparably favourable interactions with water. Separating water’s hydrogen bonds to accommodate oil is therefore energetically unfavourable, so the two liquids have very low mutual solubility and usually form separate layers.

Henry’s law relates the partial pressure of a gas above a solution to its mole fraction in the liquid, commonly as p = K_H x. Thus it provides the relevant framework for discussing gas solubility, pressure, and dissolved composition at a fixed temperature. Hess, Faraday, and Newton’s cooling laws describe different chemical or physical relationships.

The maximum amount that dissolves is an equilibrium quantity and usually depends on temperature. A value stated without temperature may be incomplete because heating or cooling can change the solubility and therefore the saturation amount. Temperature does not remove pressure, rename the solute, or have one universal value in all solutions.

For most gases, increasing pressure raises solubility according to Henry’s law, while lowering temperature generally favours dissolution because gas dissolution is usually exothermic. Since both changes act in the same direction, the amount of dissolved gas increases. The exact increase depends on the gas, solvent, and ranges involved, but it is not forced to zero.

A solubility curve plots solubility against temperature. A steep upward slope means that a small rise in temperature produces a relatively large increase in the amount that can dissolve. It does not by itself establish that pressure has no effect, and it certainly does not mean the solute is insoluble. The slope describes temperature sensitivity.

A downward slope means that the plotted solubility value becomes smaller as temperature increases. This behaviour is consistent with an exothermic dissolution and is common for gases dissolved in liquids, although the exact curve depends on the system. It does not imply that saturation is impossible or that the solvent has no role.