3: Solubility

Insoluble impurities remain undissolved even when the solvent is hot. The hot solution should therefore be filtered before cooling, so these impurities are retained by the filter while the desired compound passes through in solution. Cooling first could cause desired crystals to form on the filter and reduce recovery. Hence hot filtration is correct.

At a specified temperature, a saturated solution contains the maximum equilibrium amount of solute that can remain dissolved. A supersaturated solution contains more than this normal limit, often after careful cooling, and is therefore metastable. A small disturbance or seed crystal can cause the excess solute to crystallize. Thus B is unambiguous.

The mole fraction of a component is defined as its number of moles divided by the total number of moles of all components in the solution. For a dissolved gas, its mole fraction therefore indicates the relative amount of gas present in the liquid solution. It is a dimensionless ratio, not a mass or a visual property.

Solubility depends on the physical state and nature of the solute and solvent, and also on temperature and pressure. These factors do not all act in the same way for solids, liquids and gases. Separating the information first prevents applying Henry's law to a solid or using a temperature rule without checking the process. Therefore, A is the sound strategy.

At constant temperature, Henry's law states that the dissolved amount or mole fraction of a gas is directly proportional to its partial pressure above the solution, within the law's applicable range. Increasing pressure therefore drives more gas into the liquid and increases its solubility. Opening a carbonated drink reverses this effect by lowering pressure.

An exothermic dissolution releases heat, so heat behaves like a product of the process. Raising the temperature shifts equilibrium in the direction that consumes heat, which is generally toward gas leaving the liquid. Consequently, the amount of dissolved gas decreases. The exact change depends on the system, but the usual qualitative conclusion is a decrease, making A correct.

Lower temperature usually favours retention of a gas in the liquid because gas dissolution is commonly exothermic. Higher pressure increases the number of gas molecules entering the liquid and, at constant temperature, increases dissolved concentration according to Henry's law. Combining both favourable changes gives low temperature and high pressure, option A.

A saturated solution already contains the maximum equilibrium amount of solute that can dissolve at the stated temperature. Adding more solute does not increase that limit. The added excess may remain as solid at the bottom, while the dissolved portion and solution composition stay at saturation. Therefore, the extra solute does not dissolve completely, making A correct.

At constant temperature, Henry’s law gives p = K_H x, so the dissolved gas concentration or mole fraction x is directly proportional to the gas pressure p when K_H is fixed. If x must become twice its original value, p must also become twice its original value: p₂ = 2p₁. Halving pressure would halve the solubility, while fourfold pressure would generally produce fourfold, not twofold, solubility in this ideal comparison.

Carbon dioxide is bottled under pressure, which keeps more gas dissolved in the liquid than would be present at atmospheric pressure. Opening the bottle suddenly lowers the pressure above the drink. The equilibrium solubility of CO₂ then falls, so excess dissolved gas leaves the liquid as bubbles until a new equilibrium is approached. The bubbling is not caused by freezing or solid formation.

Water pressure increases with depth. Under the higher pressure experienced by a diver, gases in the breathing mixture have higher partial pressures and can dissolve in blood and tissues in greater amounts. If the diver ascends too rapidly, pressure falls and these gases may come out as bubbles, which explains the need for controlled decompression.

Atmospheric pressure generally decreases as altitude increases. The partial pressure of gases above the water therefore tends to be lower, and Henry's law predicts a lower dissolved concentration at lower pressure, assuming temperature and other conditions are comparable. The gas does not lose its mass; it simply has less pressure-driven dissolution in the water.

An unsaturated solution contains less dissolved solute than the maximum amount possible at that temperature, so it has remaining dissolving capacity. Adding the same solute gradually increases the dissolved amount. Eventually the saturation limit is reached; further added solid would remain undissolved. This procedure directly converts the solution to saturation without requiring a temperature change.

At a specified temperature, a supersaturated solution contains more dissolved solute than the normal equilibrium saturation limit. This state is metastable rather than fully stable. A small disturbance, scratching of the vessel, or addition of a seed crystal can provide a nucleation site, causing the excess solute to crystallize until the solution returns toward saturation.

A steep temperature dependence means the solid dissolves abundantly in the hot solvent but much less in the cold solvent. Preparing a hot saturated solution and cooling it slowly allows the excess solute to separate as crystals. Slow cooling generally promotes better crystal growth and can improve purity, whereas keeping the solution hot prevents crystallization.

For most gases, increasing temperature decreases solubility because dissolved gas tends to escape, while decreasing pressure reduces the driving force for gas molecules to enter the liquid. High temperature and low pressure therefore act together to produce the least solubility. Low temperature or high pressure would instead favour retention of gas in solution.

Oxygen is a gas and its dissolution in water is limited by an equilibrium between molecules entering and leaving the liquid. Heating generally favours escape of dissolved gas, so the equilibrium concentration of oxygen becomes lower. This is why warm water may support less dissolved oxygen for aquatic organisms. The water remains a solvent, and pressure is not necessarily higher.

At constant temperature, Henry's law gives a direct proportionality between gas pressure and dissolved concentration or mole fraction: x ∝ p. Therefore, if the pressure changes from p to 3p, the dissolved amount changes from x to 3x, assuming the solution remains within the law's valid range. The relation is linear, not inverse or squared.

A saturated solution contains the maximum amount of dissolved solute at its current temperature. If the solubility of the solid increases on heating, the solvent can hold more solute at the higher temperature. The amount already dissolved remains the same initially, but it becomes less than the new maximum capacity; therefore, the solution becomes unsaturated. Cooling or evaporating would not produce this stated result.

Pressure is a major factor for gases, but temperature also changes gas solubility, usually lowering it when temperature rises. The chemical nature of the gas and solvent controls their intermolecular interactions, so different gases can have different solubilities at the same pressure and temperature. Container colour, glass thickness and viewing angle do not determine equilibrium solubility.

Dissolution becomes favourable when interactions between solute and solvent particles can compensate for the interactions present in the pure solute and pure solvent. A polar solvent can strongly attract a polar or ionic solute through dipole, ion–dipole or hydrogen-bonding interactions. Consequently, suitable solute–solvent attraction promotes greater solubility; container dimensions and naming do not.

Henry’s law states that, at constant temperature, the concentration or mole fraction of a gas dissolved in a liquid is proportional to its partial pressure above the liquid: x = p/K_H. Thus, increasing pressure generally forces more gas into solution, and reducing pressure allows some gas to escape. The relation is not inverse or independent, and container shape alone does not determine gas solubility.

Henry's law is commonly written as p = K_H x. At a fixed pressure, a larger K_H gives a smaller mole fraction x of the gas in solution. Therefore, a gas with a large Henry's law constant has a lower tendency to dissolve. Infinite solubility and complete independence from pressure are not valid conclusions.

At constant temperature, Henry’s law states that the solubility of a gas in a liquid is directly proportional to the partial pressure of that gas above the liquid. Therefore, if the pressure changes from P to 4P, the solubility changes in the same ratio, from S to 4S. It does not become double, half, or remain unchanged under the ideal conditions stated.

For most gases, solubility in a liquid increases when pressure increases and decreases when temperature increases. Thus, high pressure forces more gas molecules into the liquid, while low temperature helps prevent them from escaping. Carbonated drinks are therefore bottled under pressure and kept cool. Low pressure or high temperature would favor gas release rather than retention.