Solubility — Solid in Liquid, Gas in Liquid

What Solubility Means

Solubility of a substance is its maximum amount that can be dissolved in a specified amount of solvent at a specified temperature. The qualifier "at a specified temperature" is not decorative: because solubility is a state of balance, it carries a definite value only once the conditions are pinned down. NCERT lists four controlling variables, namely the nature of the solute, the nature of the solvent, the temperature and the pressure, and the whole of Section 1.3 is an inventory of how each of these acts on a solid solute and on a gaseous solute.

The two cases behave very differently, and the differences are physical rather than arbitrary. A solid and a liquid are nearly incompressible, so squeezing the system changes little; a gas occupies a large, compressible volume, so pressure becomes a lever. This is the structural reason behind almost every objective question on the topic. It also helps to keep solubility distinct from a concentration unit: solubility is a property of the solute–solvent pair under fixed conditions, even though it is reported using a unit such as molarity or molality. Keeping the definition strict — maximum amount, fixed solvent, fixed temperature — is what lets the four controlling variables be discussed cleanly.

Like Dissolves Like

Not every solid dissolves in a given liquid. Sodium chloride and sugar dissolve readily in water, whereas naphthalene and anthracene do not; reverse the solvent to benzene and the behaviour inverts, with naphthalene and anthracene dissolving freely while sodium chloride and sugar refuse. The pattern is captured by the observation that polar solutes dissolve in polar solvents and non-polar solutes dissolve in non-polar solvents. In general, a solute dissolves in a solvent when the intermolecular interactions in the two are similar, summarised in the maxim "like dissolves like."

The physical sense behind the maxim is energetic. For a solute to disperse, the solute–solute and solvent–solvent attractions must be replaced by solute–solvent attractions of comparable strength. When the two species interact in the same way — both by hydrogen bonding, both by dipole forces, or both by weak dispersion forces — the new contacts are about as favourable as those they replace and mixing proceeds; when they are mismatched, as for an ionic solid in a non-polar solvent, the solvent cannot pay the cost of breaking the lattice and the solid stays undissolved. This is why the nature of solute and solvent is listed first among the variables: it decides whether dissolution happens at all, before temperature and pressure decide how much.

Saturation and Dynamic Equilibrium

When a solid solute is added to a solvent, some of it dissolves and its concentration in solution rises; this process is dissolution. Simultaneously, solute particles already in solution collide with the solid and separate out again, a process called crystallisation. A stage is reached at which these two opposing processes occur at the same rate, so the number of particles entering solution equals the number leaving it, and a state of dynamic equilibrium is established.

At equilibrium the concentration of dissolved solute stays constant under the given temperature and pressure. This equilibrium is written as a reversible change, and the same description applies when a gas dissolves in a liquid solvent:

$\ce{Solute + Solvent <=> Solution}$

A solution in which no more solute can be dissolved at the same temperature and pressure is a saturated solution; an unsaturated solution is one that can still take up more solute under the same conditions. The saturated solution, being in dynamic equilibrium with undissolved solute, holds the maximum amount of solute for that solvent — and that maximum concentration is precisely the solubility.

Two points about this equilibrium repay attention. First, "dynamic" means the dissolving and crystallising have not stopped; they continue at equal and opposite rates, so the macroscopic concentration is steady while individual particles keep crossing the boundary. Second, because the equilibrium is genuine, it responds to disturbances in exactly the way Le Chatelier's Principle predicts — and that single fact governs the temperature behaviour of both solids and gases discussed below. The same dissolution-and-crystallisation picture is followed when a gas dissolves in a liquid, except that the "undissolved solute" is now the gas phase sitting above the solution.

Solubility of a Solid in a Liquid

The solubility of a solid in a liquid is significantly affected by temperature. Because the dissolution equilibrium $\ce{Solute + Solvent <=> Solution}$ is a dynamic equilibrium, it must obey Le Chatelier's Principle. The direction in which solubility moves with temperature therefore depends entirely on the sign of the enthalpy of dissolution, $\Delta_{sol}H$.

If, in a nearly saturated solution, the dissolution process is endothermic, with $\Delta_{sol}H > 0$, then supplying heat drives the equilibrium forward and the solubility increases with rise in temperature. If the dissolution is exothermic, with $\Delta_{sol}H < 0$, then heating pushes the equilibrium backward and the solubility decreases as temperature rises. Both trends are confirmed experimentally, so no single rule can be quoted for "solid in liquid" — the question stem will supply the sign of $\Delta_{sol}H$ or name a familiar substance, and that sign is the switch: positive climbs with temperature, negative drops.

Pressure, by contrast, has no significant effect on the solubility of solids in liquids. The reason is structural: solids and liquids are highly incompressible and remain practically unaffected by changes in pressure, so the dissolution equilibrium does not shift. This single sentence answers a recurring class of objective questions, and it is the cleanest point of contrast with the gas case that follows.

Solubility of a Gas in a Liquid

Many gases dissolve in water. Oxygen dissolves only to a small extent, yet it is this dissolved oxygen that sustains all aquatic life; hydrogen chloride gas, on the other hand, is highly soluble in water. Unlike the solid case, the solubility of a gas in a liquid is greatly affected by both pressure and temperature, and the two act in opposite directions of intuition.

Effect of pressure

Consider a closed system whose lower part is a solution and whose upper part is gas at pressure $p$ and temperature $T$, in dynamic equilibrium so that gas particles enter and leave the solution at equal rates. If the gas is now compressed to a smaller volume, the number of gas particles per unit volume above the solution rises, and so does the rate at which they strike the surface and enter it. The solubility increases until a new equilibrium is established at the higher pressure. In short, the solubility of gases increases with increase of pressure.

Effect of temperature

The solubility of gases in liquids decreases with rise in temperature. Once dissolved, the gas molecules are in the liquid phase, so dissolution resembles condensation and heat is evolved — the process is exothermic. Applying Le Chatelier's Principle to this exothermic dissolution equilibrium, raising the temperature shifts it backward, expelling gas from solution. This is exactly why aquatic species are more comfortable in cold waters than in warm ones, and why warming a glass of an aerated drink drives the dissolved gas out.

The Henry's Law Link

Henry was the first to give a quantitative relation between pressure and the solubility of a gas in a solvent. At a constant temperature, the solubility of a gas in a liquid is directly proportional to the partial pressure of the gas present above the surface of the liquid. Using the mole fraction $x$ of the gas in solution as the measure of solubility, the most commonly used statement is that the partial pressure of the gas in the vapour phase is proportional to its mole fraction in solution:

$p = K_H\, x$

Here $K_H$ is the Henry's law constant. A plot of partial pressure against mole fraction is the straight line of Figure 2. Different gases have different $K_H$ values at the same temperature, which shows that $K_H$ is a function of the nature of the gas. Rearranging to $x = p / K_H$ makes a key consequence explicit: at a given partial pressure, the higher the value of $K_H$, the lower the solubility of the gas.

The temperature trend is encoded in the same table when read across NCERT's data: $K_H$ for both $\ce{N2}$ and $\ce{O2}$ increases as temperature rises (for example $\ce{O2}$ moves from $34.86$ to $46.82$ kbar), which means their solubility falls with warming — the gas-temperature rule again, now in numbers.

Three applications follow directly, each a favourite source of objective questions. To increase the solubility of $\ce{CO2}$ in soft drinks and soda water, the bottle is sealed under high pressure, so more gas stays dissolved; opening it drops the pressure and the gas fizzes out. Scuba divers breathing air at high pressure dissolve more nitrogen in the blood, and on rapid ascent the falling pressure releases it as bubbles that block capillaries — the dangerous condition called the bends. At high altitude the partial pressure of oxygen is low, so little oxygen dissolves in the blood, producing the weakness and impaired thinking of anoxia.

Henry's law holds well only within limits worth memorising: the pressure must not be too high and the temperature not too low, and the gas must not chemically react with, associate in, or dissociate in the solvent. Within these limits the law reliably describes how a gas partitions between vapour and liquid.

Factors Affecting Solubility — Summary

The two tables below collect every controlling factor for the two cases, in the form NEET tests them. They are deliberately parallel so the contrasts stand out at a glance.