3: Solubility

The solubility of most gases in liquids decreases when temperature rises. Heating gives dissolved gas molecules more kinetic energy, so they escape from the liquid more readily; gas dissolution is also commonly exothermic. The liquid does not become solid, gases do not form only in cold containers, and heating does not make pressure automatically zero. Therefore, option A explains the observation.

If heating increases the solubility of a solid, heat acts as a factor that favors dissolution. By Le Chatelier’s principle, adding heat shifts an equilibrium in the direction that absorbs heat; this indicates that the dissolution process is generally endothermic. This is a useful inference rather than an absolute rule for every solid. The other choices do not explain the temperature dependence.

When heating decreases the solubility of a solid, the dissolution equilibrium is generally shifted away from the dissolved state. This indicates that dissolution releases heat and is therefore usually exothermic. The observation does not imply that the solvent is a gas or that the solute must be a metal; those options are unrelated.

A saturated solution contains the maximum amount of solute that can remain dissolved at the stated temperature and conditions. If more solute is added without changing those conditions, it cannot dissolve appreciably and remains as undissolved solid. The solution does not automatically become dilute, and colour does not determine its saturation limit.

At a specified temperature, an unsaturated solution contains less dissolved solute than the maximum amount possible at equilibrium. Consequently, some additional solute can still dissolve, provided the solvent and temperature remain suitable. A solution need not be solute-free, and it must contain a solvent. Unsaturated solutions can occur in many kinds of systems, not only in gases, so option A is the correct definition.

A supersaturated solution contains more dissolved solute than is stable at the current temperature. It is metastable, so a seed crystal supplies a surface and an ordered starting point for crystallisation. The excess solute separates as crystals until the normal equilibrium solubility is approached. The solvent is not destroyed and solubility does not become infinite.

A hot saturated solution contains a large amount of solute because solubility is often higher at high temperature. On slow cooling, solubility decreases and the excess solute leaves the solution as well-formed crystals. Slow growth tends to exclude impurities from the crystal lattice, whereas rapid cooling can trap impurities and produce small crystals.

Pressure has a strong effect on phases whose volume changes greatly, especially gases. Solids and liquids are condensed phases and are only slightly compressible, so their volumes and chemical potentials change little over ordinary pressure changes. Consequently, the solubility of a solid in a liquid is usually nearly pressure-independent, unlike gas solubility.

Atmospheric pressure decreases as altitude increases. For an open water surface, the partial pressure of gases above the liquid also generally decreases, so Henry's law predicts a lower equilibrium concentration of dissolved gases. Water's polarity is a solvent property, not the altitude-related cause, and container shape or colour does not explain the main effect.

As a diver descends, the pressure exerted by the surrounding water increases. At nearly constant temperature, Henry’s law indicates that the amount of a gas dissolved in a liquid increases with its partial pressure. Consequently, gases such as nitrogen can dissolve in greater amounts in body fluids under depth. High temperature is not the reason, blood contains water, and gases are pressure-sensitive.

Henry’s law is commonly written as p = K_H x, where p is the partial pressure of the gas, x is its mole fraction in the solution, and K_H is Henry’s law constant at a specified temperature. If temperature and the gas–solvent system remain fixed, K_H is constant, so increasing p increases x in direct proportion. Thus, option A is correct.

At constant temperature, Henry’s law gives a direct proportionality between gas solubility and the gas pressure above the solution. If the initial pressure is P and solubility is S, reducing the pressure to P/2 changes the solubility to S/2. Therefore, it becomes half. It does not double or become four times, and it would remain unchanged only if pressure had no effect.

At constant temperature, gas solubility S is directly proportional to pressure P, so S/P remains constant for the same gas–solvent system. If the desired solubility is 3S, the pressure must be changed from P to 3P. Reducing pressure would reduce solubility, keeping it unchanged would not produce the required increase, and zero pressure cannot produce three times the original dissolution.

Gas solubility in a liquid depends on several conditions, especially pressure and temperature. At constant temperature, increasing the gas pressure generally increases its solubility according to Henry’s law. For most gases, increasing temperature decreases solubility because dissolved gas escapes more readily. Liquid colour and container height alone are not governing general factors, so option A is the scientifically appropriate statement.

A polar solute has separated or partially separated charges, and a polar solvent can interact with those charges through dipole-dipole forces, hydrogen bonding, or ion-dipole attraction. If solute-solvent attractions are sufficiently strong compared with the original interactions, mixing is favoured. This is the basis of the rule “like dissolves like,” though it is a general guide, not an absolute law.

Dissolution requires solute particles to leave their original arrangement and become accommodated in the solvent. Strong and favourable solute-solvent interactions stabilize the separated particles, lowering the tendency for them to return to the pure solute phase. Hence solubility generally increases. Stabilization does not destroy the solute or make solubility zero.

For salts having the same stoichiometric type, a smaller Ksp generally corresponds to smaller equilibrium ion concentrations and hence lower molar solubility. The comparison must be made carefully because different formulas produce different powers and numerical Ksp values. A small Ksp does not mean that no ions dissolve; it indicates a small equilibrium amount.

At a fixed temperature, the saturated concentration is limited by the amount of solvent and its solubility. Evaporation reduces the solvent quantity, so the same amount of dissolved solute may exceed the new saturation limit. The excess therefore separates as crystals or precipitate. The solution does not become unsaturated, and evaporation does not automatically increase intrinsic solubility.

A saturated solution is defined relative to a particular temperature. If heating increases the solubility of the solute, the amount already dissolved may be below the new, larger saturation limit. The solution then becomes unsaturated and can dissolve more solute. If heating decreases solubility, some solute would instead crystallise, so option B gives the opposite outcome.

Aquatic organisms obtain oxygen for cellular respiration from the oxygen dissolved in water, not directly from the surrounding air. Because water can hold only a limited amount of oxygen, changes in temperature, pressure, turbulence, and biological activity can affect the available supply. Low dissolved oxygen can stress or kill aquatic organisms; oxygen does not colour or solidify water.

Oxygen is a gas, and the solubility of most gases in water decreases as temperature increases. Therefore, cold water can retain more dissolved oxygen than warm water, assuming comparable pressure and water composition. Boiling drives dissolved gases out, while zero pressure would not favor gas dissolution. In natural waters, oxygen content can also depend on agitation, salinity, and biological activity, but among these choices low temperature is correct.

Vigorous shaking increases the liquid-gas interfacial area and can create bubbles. In an open container, these bubbles provide a rapid path for dissolved gas to leave the liquid and approach equilibrium with the lower external pressure. Shaking does not make Ksp relevant, does not necessarily increase pressure, and does not convert the gas into a solid. The actual outcome also depends on whether the container is sealed.

Before opening, the carbonated drink is equilibrated with carbon dioxide at a pressure greater than atmospheric pressure. Opening the bottle lowers the gas pressure above the liquid, so the equilibrium solubility of CO2 decreases. The excess dissolved gas escapes as bubbles until a new equilibrium is approached. Time allows this escape; gas is not continuously generated by the liquid.

Using the concentration form of Henry’s law, p = K_H x, the mole fraction is x = p/K_H. At the same pressure and temperature, a smaller K_H gives a larger value of x, meaning greater solubility in the liquid. Therefore, the first gas dissolves more. The conclusion assumes the same convention for Henry’s constant and comparable solvent conditions; a lower constant does not mean that no gas dissolves.

Solubility is an equilibrium property, and changing temperature can change the balance between dissolved and undissolved phases. Many solids become more soluble on heating, while some become less soluble; gas solubility usually decreases on heating. Therefore, a numerical solubility value is incomplete without its temperature. Pressure is important especially for gases, but it is not the only variable for all substances.