3: Solubility

Let Q be the ionic product and Ksp the solubility product. When Q > Ksp, the concentrations of ions are higher than permitted at equilibrium. The solution is therefore supersaturated with respect to the salt, and ions combine to form a solid precipitate until the ionic product falls back to Ksp. Q < Ksp would instead indicate an unsaturated solution.

When Q, the ionic product, is less than Ksp, the current ion concentrations are below the equilibrium precipitation limit. The solution is unsaturated with respect to that salt and can dissolve more of it, provided the solid is available and other conditions remain fixed. Precipitation is expected only when Q exceeds Ksp; equality represents saturation equilibrium.

The condition Q = Ksp means that the ion concentrations have reached the equilibrium value for the salt. The solution is saturated: it is in equilibrium with the solid phase if solid is present, and there is no net tendency for further precipitation or dissolution. Equality does not mean that the solution must visibly contain a precipitate; it describes the equilibrium concentration condition.

For a salt, precipitation begins when the ionic product Q exceeds Ksp. Therefore, to ensure that no precipitation occurs and the solution remains unsaturated, Q should be kept below Ksp. At Q = Ksp the solution is just saturated, so any further increase could cause precipitation. Q greater than Ksp definitely favours formation of a solid phase.

Dissolving an ionic solid requires separating its ions from the crystal lattice and solvating them in water. Strong ion–dipole attractions between the ions and polar water molecules stabilise the separated ions and can make dissolution favourable. If ions remain away from water or solute particles never separate, dissolution cannot proceed. Water’s polarity is essential rather than absent.

For the same mass, smaller particles expose a larger total surface area to the solvent, so solvent molecules can attack them more rapidly and the dissolving rate increases. However, at a fixed temperature and equilibrium, the final solubility is determined mainly by the substance, solvent, and conditions, not by ordinary particle size. Size changes kinetics, not the equilibrium limit.

Vigorous shaking creates nucleation sites and increases the liquid–air contact area. In an open container, the gas can escape freely, and the disturbance helps dissolved gas come out as bubbles until the liquid approaches equilibrium with the lower external pressure. Shaking does not make solubility infinite or necessarily raise pressure; the liquid also does not become solid merely because it is shaken.

Henry’s law constant is not a universal number for a gas; it depends on temperature, solvent, and the chosen convention for defining the constant. To compare two gases fairly using their constants, temperature and solvent should be the same and the same convention should be used. Pressure should also be compared on the same basis. Different solvents or temperatures could change the constants themselves.

Under the usual Class 12 convention, Henry’s law is p = K_H x. At a fixed pressure, x is therefore inversely proportional to K_H. If K_H increases as temperature rises, the equilibrium mole fraction of dissolved gas decreases, so gas solubility decreases. This agrees with the common observation that heating makes many dissolved gases escape more readily.

At a fixed temperature, solubility specifies the maximum amount of solute per given amount of solvent. If solvent evaporates, the amount of solvent decreases while the dissolved solute may initially remain the same. The solution can then contain more solute than the new solvent amount can hold, so excess solute crystallises or precipitates until saturation is restored. The intrinsic solubility does not automatically increase.

Suppose a solution is saturated at the original temperature, so it contains the maximum amount then possible. If heating increases the solubility of that solid, the new maximum is larger while the amount already dissolved may remain unchanged. It is therefore below the new limit and becomes unsaturated. If solubility decreases on heating, solid would instead tend to separate.

Dissolution separates solute particles and surrounds them with solvent molecules. If the solvent forms strong, favourable interactions with the separated particles, it lowers their tendency to recombine and stabilises the dissolved state. This increases the likelihood and extent of dissolution, provided the energy needed to break the original interactions is also manageable. Stabilisation does not destroy the particles or make solubility automatically infinite.

Cooling produces crystals only when the solubility decreases enough that the solution can no longer hold all the dissolved solute. If the solubility curve is nearly flat, the cold solution can still hold almost the same amount as the hot solution, so little or no solid separates. A large increase in solubility on heating would usually favour crystallisation on cooling, not prevent it.

At constant temperature, Henry’s law gives a direct increase of gas solubility with pressure. For most gases, increasing temperature weakens retention in the liquid and increases escape, so solubility generally decreases. The two effects should be considered separately, and “generally” is important because particular gases or chemical reactions can show exceptions. The other choices reverse or ignore these trends.

In the mole-fraction form of Henry’s law, p = K_H x, where K_H is constant at fixed temperature and solvent. Thus x = p/K_H, so the dissolved gas mole fraction is directly proportional to pressure. Replacing p by 2p gives x' = 2x. This proportional result assumes the dilute-solution and ideal Henry-law conditions stated in the question.

For a salt such as AB, dissolution produces A⁺ and B⁻. If one of these ions is already present, the concentration of a product is higher from the beginning. Le Chatelier’s principle shifts the equilibrium toward undissolved AB, reducing the amount that dissolves. This is the common ion effect. It does not turn the salt into a gas or remove the solvent’s polarity.

In a saturated solution containing excess solid, dissolution continues as particles leave the solid and enter the liquid, while dissolved ions simultaneously return to the solid by crystallisation. At equilibrium, the forward and reverse rates are equal, so the concentrations remain constant even though microscopic exchange continues. This is dynamic equilibrium, not complete insolubility or destruction of the salt.

A sealed carbonated drink contains dissolved carbon dioxide under pressure. When the bottle is opened, the pressure above the liquid falls suddenly. According to Henry’s law, lower pressure reduces gas solubility, so dissolved carbon dioxide escapes as bubbles. The water mass and physical state of the solvent do not suddenly change.

Using the convention p = K_H x, where p is the gas pressure and x is its mole fraction in solution, we get x = p/K_H. At fixed pressure and temperature, increasing K_H therefore makes x smaller, meaning less gas dissolves. Thus option A is correct. Option B reverses the relationship, while C and D ignore the role of pressure, solvent, and temperature.

A supersaturated solution contains more dissolved solute than is stable at the given temperature. It can remain temporarily clear because crystal nuclei have not formed. A disturbance supplies a nucleation site, allowing excess solute to arrange into an ordered crystal phase and separate until a stable saturated state is approached.

Henry's law, in the convention commonly used in this chapter, states that the partial pressure p of a gas above the solution is proportional to its mole fraction x in the solution. The proportionality constant is Henry's constant K_H, so p = K_H x. Option B incorrectly divides by K_H, while C and D do not express the required direct proportionality.

For a gas obeying Henry's law, p = K_H x, so K_H = p/x. When the pressure is the same, a more soluble gas has a larger mole fraction x in water. Since p is fixed and x is larger, p/x is smaller; therefore its Henry's constant is lower. Hence option A is correct, not B, C, or D.

Water pressure increases with depth, and increased pressure raises the amount of breathing gases that can dissolve in body fluids. If a diver ascends too quickly, the pressure falls and the dissolved gas may come out as bubbles. This is why controlled ascent is important; blood does not solidify and gas mass does not disappear.

Henry's law is p = K_H x. At constant temperature, K_H is treated as constant for the given gas and solvent. Rearranging gives x = p/K_H, so increasing the partial pressure increases the mole fraction in direct proportion. Therefore option A is correct. It does not decrease or remain fixed; option D is physically and mathematically unrelated to Henry's law.

Saturation is defined at a particular temperature. If heating permits additional solid to dissolve, the solubility has increased with temperature. This behaviour is consistent with an endothermic dissolution, because added heat favours the dissolving direction. It does not prove that every dissolution is endothermic, but it supports that possibility for this system.