3: Solubility

Using the concentration form of Henry’s law, p = K_H x, the mole fraction of dissolved gas is x = p/K_H. At the same pressure, increasing K_H therefore decreases x, so less gas dissolves. Thus, a gas with a higher Henry’s law constant has lower solubility. It is not correct to say that its solubility is zero or independent of pressure.

At constant temperature, Henry’s law states that the amount or mole fraction of a gas dissolved in a liquid is directly proportional to its partial pressure: x = p/K_H. If the pressure changes from p to p/2 while K_H remains constant, x also changes to x/2. Therefore, the solubility becomes half, not double or unchanged.

Carbon dioxide is dissolved in a sealed cold drink under pressure higher than atmospheric pressure. On opening the bottle, the pressure above the liquid falls suddenly. According to Henry’s law, lower pressure permits less gas to remain dissolved, so excess carbon dioxide escapes as bubbles. The other options do not describe the governing change.

Hydrostatic pressure increases as depth below the water surface increases. A gas in contact with a liquid dissolves more readily at higher pressure, as expressed by Henry’s law. Therefore, gases from the breathing mixture can dissolve more in a diver’s blood at depth. Boiling, absence of liquid, and pressure independence are scientifically incorrect.

Dissolved oxygen is a gas–liquid solubility example. For most gases, increasing temperature increases the kinetic energy of dissolved molecules and makes their escape from the liquid easier. Consequently, warm water generally holds less oxygen than cold water at comparable pressure. Oxygen does not become solid on heating, and it can dissolve in water.

For a gas dissolved in a liquid, higher pressure favours dissolution because more gas molecules are forced into contact with the liquid. Lower temperature generally favours retention of gas because gas molecules have less tendency to escape. Thus, low temperature and high pressure together give the most favourable condition. Zero pressure would not favour dissolution.

A saturated solution contains the maximum amount of a solute that can dissolve in a specified amount of solvent at a specified temperature. If extra solute is added under the same conditions, it remains undissolved or establishes a dissolution–crystallisation equilibrium. Saturation therefore refers to capacity, not absence of solvent or solute, and it is always temperature-dependent.

An unsaturated solution contains less dissolved solute than the maximum possible amount at that temperature. Adding the same solute in small portions allows it to dissolve until the saturation limit is reached; further addition then remains undissolved. Simply leaving it untouched does not generally change its composition, while removing all solvent destroys the stated solution condition.

The slope of a solubility-versus-temperature graph represents the change in solubility for a given change in temperature. A steep slope therefore means that even a small temperature change produces a comparatively large change in solubility. Such behaviour can be useful in crystallisation by cooling. A flat curve, in contrast, indicates weak temperature dependence.

Dissolution is favoured when the new solute–solvent interactions can compensate for breaking solute–solute and solvent–solvent interactions. Substances with similar polarity or intermolecular-force character often form such favourable interactions; this is expressed as “like dissolves like.” It is a useful guideline, not an absolute rule, because molecular size, hydrogen bonding, and other factors also matter.

Stirring continually brings fresh solvent into contact with the solute surface and removes the layer already enriched with dissolved particles. This reduces transport resistance and usually increases the rate at which the solute dissolves. Once equilibrium is reached at the same temperature, however, stirring does not change the maximum amount that can remain dissolved. It affects rate, not equilibrium solubility.

Atmospheric pressure generally decreases as altitude increases. For an open body of water, the partial pressure of gases above the surface therefore tends to be lower. Henry’s law predicts lower equilibrium dissolution at lower gas pressure, so the water may contain less dissolved atmospheric gas. This is a possible pressure effect; temperature and water movement can also influence the actual amount.

For each gas, Henry’s law in the form p = K_H x gives x = p/K_H. Since both gases are at the same pressure, the gas with the smaller K_H has the larger dissolved mole fraction. Therefore, the first gas dissolves more. A lower constant does not mean zero solubility, and gas solubility still depends on pressure.

Solubility is an equilibrium quantity, and changing temperature can shift the dissolution equilibrium. Many solids become more soluble on heating, while some show little change or even a decrease; gas solubility generally decreases on heating. Therefore, a numerical solubility value is incomplete without its temperature, along with the solvent and units. Colour or a changed substance name is irrelevant.

For a gas in a liquid, Henry’s law states that the dissolved amount or mole fraction is proportional to the gas’s partial pressure at constant temperature. Thus, changing pressure changes the equilibrium solubility. A reported value without pressure is incomplete, especially when the gas mixture composition may alter partial pressure. Pressure is not restricted to solids and does not change the solvent’s name.

Before opening, carbon dioxide is held in the drink at elevated pressure. Opening exposes the liquid to a much lower pressure, so the equilibrium solubility falls. The liquid is temporarily supersaturated with respect to carbon dioxide, and the gas escapes through bubbles until a new equilibrium with the atmosphere is approached. Time allows this degassing to continue, so fizziness decreases.

The ionic product Q is calculated from the present ion concentrations, whereas Ksp is the equilibrium limit for the sparingly soluble salt at that temperature. If Q exceeds Ksp, the solution is supersaturated and precipitation can begin. If Q is below Ksp, precipitation is not favoured; equality represents saturation. Mass, colour, smell, or container height do not provide this equilibrium criterion.

Many aquatic organisms obtain oxygen directly from the water through gills, body surfaces, or other respiratory structures. Dissolved oxygen is used in cellular respiration to release energy from food. If its concentration falls, respiration becomes less efficient, causing stress, reduced activity, or death when the shortage is severe. Oxygen is not important only for metals and does not make water solid or simply colour it.

Oxygen is a gas, and the solubility of most gases in water decreases as temperature rises at comparable pressure. Lower temperature reduces the kinetic tendency of dissolved oxygen molecules to escape into the gas phase, so cold water can retain more oxygen. Boiling drives dissolved gases out very strongly, and oxygen solubility is not zero at every temperature.

At a fixed temperature, a saturated solution already contains the maximum equilibrium concentration of the salt’s ions. Adding more solid does not change Ksp or the saturation concentration; instead, the excess solid remains undissolved while dynamic equilibrium continues between dissolution and crystallisation. Only a change in temperature, solvent, or chemical conditions could alter the equilibrium solubility significantly.

At constant temperature, Henry’s law states that the concentration or amount of a gas dissolved in a liquid is proportional to the gas’s partial pressure. Therefore, changing the pressure from p to 2p makes the dissolved amount approximately 2S, provided the solution remains dilute and the gas behaves ideally. The word “approximately” allows for real-system deviations.

The solubility of most gases in water decreases when temperature rises. Therefore, warm water generally contains less dissolved oxygen than cool water. Fish depend on this dissolved oxygen for respiration, so a reduction can make breathing difficult. Temperature does not make all gases completely insoluble or instantly acidify water.

Gas solubility in a liquid generally increases with pressure because more gas molecules are forced into the liquid. It generally decreases with increasing temperature because dissolved gas tends to escape more readily. Thus, combining high pressure with low temperature gives the condition most favourable for maximum solubility.

For an endothermic dissolution, heat is absorbed and can be treated as a reactant in the equilibrium representation. Raising the temperature supplies more of this required heat and generally shifts the equilibrium toward dissolved particles. Therefore solubility usually increases, although the exact magnitude depends on the substance.

In exothermic dissolution, heat is released and may be regarded as a product. Increasing temperature favours the reverse direction according to equilibrium reasoning, so the amount dissolved generally decreases. The result is not necessarily exactly double or zero, and pressure is not the controlling relation for ordinary solid solutes.