3: Solubility

Water molecules are strongly polar and form extensive hydrogen-bond interactions with one another. Most oils consist mainly of non-polar hydrocarbon molecules, which cannot form equally favourable interactions with water. Consequently, mixing is energetically unfavourable and separate phases form. The separation is not caused by colour, by oil converting into water, or by the absence of molecules in water.

For a gas dissolved in a liquid, increasing pressure generally increases solubility, as expressed by Henry’s law. Increasing temperature usually decreases gas solubility because escape of gas from the liquid becomes more favourable. Therefore, lowering pressure or raising temperature can reduce the dissolved amount. Keeping a container closed is not by itself a universal change in solubility, and increasing partial pressure has the opposite effect.

Solubility is normally specified for a definite amount of solvent, such as grams of solute per 100 grams of solvent, at a stated temperature. The same solute mass may be unsaturated in a large amount of solvent but exceed the solubility limit in a small amount. Therefore, solvent amount and temperature are needed to judge saturation or make a fair comparison.

At high temperature, a large amount of the solid can dissolve in the solvent. When the hot saturated solution is cooled, its solubility falls substantially, so the solution can no longer retain all the dissolved material. The excess separates as crystals, especially if cooling is controlled. Pressure is not the required condition, and the statement says nothing about whether the solid dissolves in water at all.

Begin by identifying the solute and solvent and deciding whether they are ionic, polar, or non-polar. Then inspect the temperature and pressure conditions, because temperature strongly affects many solids and gases, while pressure is especially important for gases. Finally check the stated amounts and units. This systematic order prevents applying a gas rule to a solid or ignoring a required condition.

Henry’s law states that, at constant temperature and for a dilute solution, the concentration or mole fraction of a dissolved gas is proportional to its partial pressure above the liquid. Thus, if that partial pressure changes from p to 3p, the dissolved amount changes approximately from x to 3x, provided the solution remains in the range where Henry’s law applies. The answer is therefore nearly three times.

Using the form p = K_H x, where p is the gas partial pressure and x is its mole fraction in the solution, a fixed p gives x = p/K_H. Therefore, a smaller Henry’s law constant produces a larger dissolved-gas mole fraction and hence greater solubility. Colour and molecular mass alone do not determine the comparison, although molecular properties can influence the value of K_H.

A hot saturated solution contains the maximum amount of solute that can dissolve at the higher temperature. If the solid’s solubility decreases on cooling, the lower-temperature solution cannot retain all of that solute. The excess leaves the solution and arranges into an ordered crystal lattice. Slow cooling often allows larger, better-formed crystals, whereas the other options describe no relevant chemical process.

For the convention p = K_H x, the dissolved-gas mole fraction is x = p/K_H. At the same pressure, increasing K_H therefore decreases x, so the gas is less soluble. Henry’s law constants can be defined in other reciprocal conventions in some texts, so the equation or convention must be checked; with the standard relation used here, option B is correct.

In a hot saturated solution, the higher temperature may permit a relatively large amount of solid solute to remain dissolved. On cooling, the solubility of many solids decreases, so the lower-temperature solution cannot hold all of the original dissolved amount. The excess solute separates and forms crystals. This depends on the solubility–temperature relationship, not on formation of a new element or unavoidable zero pressure.

A supersaturated solution contains more dissolved solute than is normally stable at the given temperature, but it may remain clear if no nucleation site is present. A seed crystal supplies an ordered surface on which solute particles can attach, so crystallization can begin and excess solute separates until a stable saturated state is approached. The solvent does not disappear and solubility does not become infinite.

Water is a polar solvent and ionic salt is stabilized by ion–dipole interactions and hydration in water. Oil is mainly non-polar, so it does not form comparably favourable interactions with water and remains as a separate phase. This illustrates the general “like dissolves like” idea. Density, colour, and whether a substance is a gas do not explain the observed solubility difference.

At a fixed temperature, Henry’s law states that the solubility, or dissolved concentration, of a gas in a liquid is directly proportional to the partial pressure of that gas above the liquid. If the pressure is reduced to one-half, the dissolved amount also becomes approximately one-half, assuming other conditions remain unchanged. Therefore, option A is correct; doubling pressure would instead tend to double solubility.

With the convention p = K_H x, the mole fraction of dissolved gas is x = p/K_H. At the same pressure, a smaller K_H therefore gives a larger x, meaning a greater proportion of gas is dissolved and the solubility is higher. This conclusion depends on the stated convention; under the reciprocal convention the numerical constant is defined differently, so the equation should always be checked.

A hot saturated solution may contain more dissolved solid than the solvent can retain at a lower temperature. When cooling reduces the solubility, the solution becomes supersaturated temporarily and the excess solute separates from it. With suitable nucleation, the separated particles arrange into crystals. The solvent does not normally vanish on cooling, and the solute does not generally become a gas under this process.

A supersaturated solution contains more dissolved solute than the stable equilibrium amount at that temperature. It can remain in this metastable state until a disturbance, impurity, scratch, or seed provides a nucleation site. Solute particles then begin to assemble into a crystal, lowering the dissolved concentration toward saturation. The process is not caused by absence of solute or pressure.

Henry’s law states that at constant temperature the mole fraction or concentration of a dissolved gas is directly proportional to its partial pressure above the solution. If the partial pressure changes from p to 4p, the dissolved amount changes approximately from x to 4x, assuming the solution remains dilute and the law remains applicable. Therefore, the solubility becomes nearly four times.

At great depth, the surrounding pressure is high, so a larger amount of gases such as nitrogen can dissolve in body fluids. If a diver rises suddenly, pressure falls rapidly and the dissolved gas can come out as bubbles in blood and tissues. These bubbles may cause decompression sickness. A slow ascent allows the gas to leave gradually and safely.

A supersaturated solution contains more dissolved solute than is normally stable at that temperature and is therefore metastable. A seed crystal supplies an ordered surface and lowers the barrier to nucleation. Excess solute deposits on it until the solution approaches the ordinary saturated state.

For a gas in a liquid, increasing the gas pressure above the liquid generally increases the number of gas molecules entering and remaining dissolved, as expressed by Henry’s law. Increasing temperature usually gives dissolved gas molecules more energy to escape, so gas solubility decreases. Therefore, option A correctly combines the usual pressure and temperature effects. The reverse trend in option B is incorrect.

Water is polar: its oxygen end is partially negative and its hydrogen ends are partially positive. These poles attract cations and anions, surround them through hydration, and help separate them from the ionic crystal lattice. Whether dissolution occurs also depends on lattice energy, but ion–water attraction is the relevant reason among these options.

At a high altitude, the partial pressure of a gas in contact with water is generally lower. Henry’s law states that, at constant temperature, dissolved mole fraction is proportional to partial pressure. Therefore equilibrium may shift toward gas leaving the water, so dissolved gas amounts can decrease, although the exact effect depends on conditions.

Using Henry’s law p = K_H x, the constant can be written as K_H = p/x. When the pressure is the same for both gases, the gas that dissolves less has the smaller mole fraction x. Dividing the same pressure by this smaller x gives a larger K_H. Therefore, low solubility corresponds to a high Henry’s law constant, while high solubility corresponds to a lower one.

Represent dissolution as gas + liquid ⇌ dissolved gas + heat. When temperature is raised, the system tends to consume the added heat by favouring the reverse direction, in which dissolved gas escapes. Consequently the equilibrium concentration of dissolved gas decreases. This is a general explanation; the exact temperature effect depends on the system.

For an endothermic dissolution, heat acts like a reactant. Raising temperature supplies heat and generally shifts equilibrium toward dissolved particles, so solubility may increase. The word “may” is important because real solid solubility also depends on lattice and hydration effects; temperature trends should be confirmed for the particular substance.