Solubility

Stir some salt into a glass of water and it vanishes almost instantly. Try the same with a lump of wax and nothing happens, no matter how long you stir. Why does one substance dissolve while another refuses? And is there a limit to how much can dissolve? These are the questions that the concept of solubility answers.

What Solubility Means

Solubility (the maximum amount of a substance that can dissolve in a specified quantity of solvent at a given temperature) tells you the dissolving capacity of a particular solute-solvent pair. It is not a single fixed number for a substance; it changes depending on four factors:

• Nature of the solute and nature of the solvent (what you are dissolving and what you are dissolving it in)
• Temperature
• Pressure

The sections below explore how each of these controls the solubility of a solid in a liquid.

Why Some Solids Dissolve and Others Do Not: Like Dissolves Like

Not every solid cooperates with every liquid. $NaCl$ (table salt) and sugar dissolve quickly in water, yet naphthalene and anthracene (both non-polar organic solids) will not dissolve in water no matter how hard you try. Flip the solvent to benzene (a non-polar liquid), and the picture reverses: naphthalene and anthracene dissolve easily, while salt and sugar stay behind as undissolved solids.

The pattern behind this behaviour is straightforward: a solute tends to dissolve in a solvent when their intermolecular interactions are similar. This is often summarised as the like dissolves like rule.

• Polar solutes (substances whose molecules carry partial charges, such as $NaCl$, sugar) dissolve well in polar solvents (like water) because both have strong electrostatic or hydrogen-bonding interactions.
• Non-polar solutes (substances whose molecules have no significant charge separation, such as naphthalene, anthracene) dissolve well in non-polar solvents (like benzene) because both rely on weak London dispersion forces.

When you try to dissolve a polar solute in a non-polar solvent (or vice versa), the mismatch in intermolecular forces means the solvent cannot pull solute particles apart effectively, and the solute stays undissolved.

Dissolution, Crystallisation, and Dynamic Equilibrium

What actually happens at the molecular level when a solid meets a compatible solvent?

Step 1: Dissolution begins. When you add a solid solute to a suitable solvent, solvent molecules start pulling solute particles away from the solid surface and into the surrounding liquid. This process is called dissolution (the movement of solute particles from the solid phase into the solution). As more and more particles enter the solution, the concentration of dissolved solute rises steadily.

Step 2: Crystallisation starts competing. As the solution grows more concentrated, some dissolved solute particles collide with the remaining solid and rejoin it. This reverse process is called crystallisation (the return of dissolved particles from solution back onto the solid surface).

Step 3: Equilibrium is reached. Eventually, the rate of dissolution and the rate of crystallisation become equal. From this point on, the number of solute particles entering the solution every second is the same as the number leaving it. The concentration stops changing. This is a state of dynamic equilibrium, represented by:

$\text{Solute + Solvent} \rightleftharpoons \text{Solution} \qquad \text{(1.10)}$

The word “dynamic” is important here. Both processes are still happening; neither has stopped. But because they balance each other perfectly, the overall composition of the solution remains constant.

Saturated and Unsaturated Solutions

Once this dynamic equilibrium is established at a given temperature and pressure:

• The solution is called a saturated solution (a solution that has dissolved the maximum possible amount of solute and is in dynamic equilibrium with any undissolved solute). No additional solute can dissolve at that temperature and pressure.
• The concentration of solute in a saturated solution is exactly equal to the solubility of that solute in that solvent at that temperature.

If the solution has not yet reached this limit, meaning you could still dissolve more solute into it under the same conditions, it is called an unsaturated solution.

Think of it this way: a saturated solution is “full” of solute. An unsaturated solution still has room for more.

How Temperature Affects Solubility

Temperature has a significant effect on how much solid can dissolve in a liquid. The dissolution equilibrium (Equation 1.10) is a dynamic equilibrium, and like all dynamic equilibria, it follows Le Chatelier’s principle (when a system at equilibrium is disturbed by a change in conditions, the system shifts in the direction that partially counteracts that change).

The key question is: does the dissolution process absorb heat or release it?

• If dissolution is endothermic ($\Delta_{\text{sol}}H > 0$, meaning the process absorbs heat from the surroundings): raising the temperature provides more thermal energy, which favours the forward process (dissolution). The equilibrium shifts to the right, and solubility increases with rising temperature.

• If dissolution is exothermic ($\Delta_{\text{sol}}H < 0$, meaning the process releases heat into the surroundings): raising the temperature pushes the equilibrium backward, away from dissolution. The equilibrium shifts to the left, and solubility decreases with rising temperature.

Both of these predictions match what is observed experimentally for nearly saturated solutions. Most common solids (like sugar, $KNO_3$) dissolve endothermically in water, which is why you can dissolve more of them in hot water than in cold water.

Why Pressure Barely Matters for Solids in Liquids

Unlike temperature, pressure has no significant effect on the solubility of a solid in a liquid. The reason is simple: both solids and liquids are highly incompressible (their volumes change very little even under large pressure changes). Since pressure influences a system by changing volumes, and neither the solid solute nor the liquid solvent changes volume appreciably, altering the pressure does virtually nothing to the dissolution equilibrium.

This is quite different from gases dissolved in liquids, where pressure plays a major role, but that is a topic for a later section.