Solubility

Solubility is defined as the maximum amount of solute that can dissolve in a given quantity of solvent at a specified temperature and pressure to form a stable, homogeneous solution. When no more solute can dissolve, the solution is said to be saturated. Understanding solubility is fundamental to chemistry because it governs everything from how medicines are absorbed in the body to how carbonated beverages retain their fizz. Two of the most important categories studied at this level are solids dissolving in liquids and gases dissolving in liquids. For solids dissolving in liquids, solubility is primarily influenced by temperature and the nature of the solute–solvent interaction. The guiding principle is 'like dissolves like': polar and ionic solids tend to dissolve in polar solvents such as water, while nonpolar solids dissolve more readily in nonpolar solvents such as hexane. The dissolution of an ionic solid like sodium chloride (NaCl) in water involves the breaking of ionic lattice bonds and the formation of ion–dipole interactions between Na⁺ or Cl⁻ ions and surrounding water molecules, a process called hydration. For gases dissolving in liquids, pressure becomes the dominant variable, and this relationship is quantified by Henry's Law. Everyday examples include the dissolution of oxygen in blood (critical for respiration), carbon dioxide in soft drinks, and the dangerous buildup of nitrogen in a diver's blood during deep-sea diving. Mastering both solid and gas solubility allows students to predict, explain, and control solution behaviour across a wide range of real-world and laboratory contexts.

Key takeaways
  • Solid in Liquid: The Dissolution Process
  • Effect of Temperature on Solid Solubility
  • Gas in Liquid: Henry's Law
  • Effect of Temperature on Gas Solubility
  • Real-World Applications of Henry's Law

Solid in Liquid: The Dissolution Process

When an ionic solid such as NaCl is placed in water, polar water molecules orient themselves around the surface ions. The slightly negative oxygen end of water surrounds Na⁺ cations, and the slightly positive hydrogen ends surround Cl⁻ anions. This ion–dipole attraction supplies enough energy to overcome the ionic lattice energy, pulling ions free into solution. The overall enthalpy of dissolution (ΔH_soln) is the sum of the lattice enthalpy (endothermic, energy required to separate ions) and the hydration enthalpy (exothermic, energy released when ions are surrounded by water): ΔH_soln = ΔH_lattice + ΔH_hydration. If ΔH_soln is negative (exothermic), solubility often increases with temperature only moderately or may even decrease; if positive (endothermic), solubility typically increases significantly with temperature, as seen with KNO₃.

Effect of Temperature on Solid Solubility

For most solid solutes, solubility increases with increasing temperature because the dissolution process is endothermic — adding heat shifts the equilibrium toward the dissolved state (Le Chatelier's Principle). For example, the solubility of potassium nitrate (KNO₃) increases from approximately 13 g/100 mL at 0 °C to 247 g/100 mL at 100 °C. A notable exception is calcium sulfate (CaSO₄), whose solubility decreases with increasing temperature because its dissolution is exothermic. Pressure has negligible effect on the solubility of solids in liquids.

Gas in Liquid: Henry's Law

Unlike solids, the solubility of a gas in a liquid increases with increasing partial pressure of that gas above the solution. This is described by Henry's Law: C = k_H × P, where C is the concentration of dissolved gas (mol/L), k_H is Henry's Law constant (unique for each gas–solvent pair at a given temperature), and P is the partial pressure of the gas above the solution (atm or Pa). Physically, higher pressure forces more gas molecules into contact with the liquid surface, increasing the rate of dissolution until a new equilibrium is established. Henry's Law assumes ideal behaviour and applies only to dilute solutions.

Effect of Temperature on Gas Solubility

Dissolution of most gases in water is exothermic: \ce{Gas(g) -> Gas(aq) + heat}. By Le Chatelier's Principle, increasing temperature shifts equilibrium to the left, decreasing gas solubility. This is why warm soda goes flat faster than cold soda — CO₂ is less soluble at higher temperatures. It also explains thermal pollution in rivers: warmer water holds less dissolved O₂, harming aquatic life. The temperature dependence of Henry's constant k_H can be approximated using a van't Hoff-type equation: k_H(T) = k_H° × exp[C × (1/T − 1/T°)], where C is a solute-specific constant.

Real-World Applications of Henry's Law

Carbonated beverages are bottled under high CO₂ pressure (~2–4 atm). When the bottle is opened, pressure drops to atmospheric (~0.0004 atm partial CO₂), and dissolved CO₂ rapidly comes out of solution as bubbles. In medicine, decompression sickness ('the bends') occurs when a scuba diver ascends too quickly: nitrogen gas that dissolved in blood and tissues at high underwater pressure (obeying Henry's Law) comes out of solution as bubbles, blocking blood vessels. Treatment involves a hyperbaric chamber to slowly re-dissolve the nitrogen. Oxygen therapy also relies on Henry's Law — breathing pure O₂ at elevated pressure increases dissolved O₂ in plasma beyond what haemoglobin alone can carry.

The Key Equation

C=kH×PC = k_H \times P
  • C — Concentration of dissolved gas (mol/L (M)): The molar concentration of the gas that has dissolved in the liquid phase at equilibrium.
  • k_H — Henry's Law constant (mol/(L·atm)): A proportionality constant specific to each gas–solvent pair at a given temperature. It decreases as temperature rises, reflecting lower gas solubility at higher temperatures.
  • P — Partial pressure of the gas (atm (or Pa)): The pressure exerted by the gas of interest above the solution surface. Higher partial pressure drives more gas into solution.

Worked Example

Problem. A bottle of sparkling water is sealed under a CO₂ partial pressure of 3.50 atm at 25 °C. The Henry's Law constant for CO₂ in water at 25 °C is k_H = 3.4 × 10⁻² mol/(L·atm). (a) Calculate the concentration of dissolved CO₂ in the sealed bottle. (b) When the bottle is opened, the partial pressure of CO₂ above the liquid drops to 0.000397 atm (its atmospheric partial pressure). Calculate the new equilibrium concentration of dissolved CO₂.

Csealed=0.12 mol/L;Copen=1.35×10−5 mol/LC_{\text{sealed}} = 0.12\ \text{mol/L}; \quad C_{\text{open}} = 1.35 \times 10^{-5}\ \text{mol/L}

Answer. Sealed: 0.12 mol/L; Open (at equilibrium with atmosphere): 1.35 × 10⁻⁵ mol/L

Figure
Diagram
Henry's Law: Effect of Pressure on Gas SolubilityLow Pressure (P₁)High Pressure (P₂ = 3P₁)gasgasliquidliquid (more dissolved gas)Few gas molecules dissolvedMore gas molecules dissolvedC₁ = k_H × P₁C₂ = k_H × 3P₁ = 3C₁Gas moleculeDissolved gas
Origins · 1 minWilliam Henry and how this was discovered
WH
William Henry
1774–1836 · British

William Henry was an English chemist who, in 1803, published the quantitative relationship now known as Henry's Law. Through careful experiments measuring the absorption of gases such as CO₂, N₂, and H₂ by water at different pressures, he established that the mass of gas dissolved by a given volume of liquid is directly proportional to the pressure of the gas in contact with the liquid. This was a landmark contribution to physical chemistry and the study of solutions, forming the basis for understanding gas solubility in industrial, environmental, and biological contexts. Henry was a close friend of the chemist John Dalton (famous for atomic theory), and the two frequently collaborated and exchanged ideas about the behaviour of gas mixtures, which informed both Henry's Law and Dalton's Law of Partial Pressures.