Gas solubility in liquids forms liquid-gas solutions, such as soft drinks, where carbon dioxide is dissolved in water, and the ocean, where the solubility of oxygen and carbon dioxide supports marine life. The ability of oceans to dissolve gases impacts weather conditions in the troposphere.
However, gas-liquid interactions vary. For instance, hydrogen chloride gas is highly soluble in water, while oxygen's solubility is much lower. Because these solutions are non-ideal, Raoult’s law, which relates vapor pressure to composition, doesn't apply accurately across all mole fractions.
At low mole fractions, vapor pressure is proportional to the mole fraction, a relationship defined by Henry's law. The proportionality constant in Henry’s law, known as the Henry’s law constant (Ki), depends on the components and temperature. Unlike Raoult’s law, which defines the proportionality constant as the pure component's vapor pressure, Henry’s law uses an experimentally determined value.
Gases that are sparingly soluble in a liquid typically have a low concentration in the solution, making it ideally dilute. In such conditions, Henry's law is applicable. For example, when plotting the mole fraction of dissolved N2 (or H2) in water at 50°C against its partial pressure, the plot for N2 up to 100 atm obeys Henry’s law and is essentially linear. Above 100 atm, the plot shows increasing deviations due to non-ideal gas behavior and pressure dependence on Henry's law constant (Ki). However, H2 follows Henry’s law up to 200 atm.
At low solute concentrations where Henry’s law applies, the solute’s molality and molar concentration are each essentially proportional to its mole fraction. Hence, Henry's law can use molalities or concentrations instead of mole fractions. The larger the Ki value, the smaller the solubility of the gas.
The solubility of most nonpolar gases in water generally decreases as temperature increases at constant pressure, because dissolution is typically exothermic. Near the critical temperature of water (374 ℃), the concept of gas solubility in liquid water is no longer meaningful because liquid water ceases to exist as a distinct phase.
However, Henry’s law does not apply to all solutions, such as a dilute aqueous HCl, since it dissociates completely in water. Even at high dilution, the behavior of the electrolytes is non-ideal due to ion-ion interactions and deviates from the assumptions deriving from Henry’s law.
Gases can dissolve in liquids, forming liquid-gas solutions like soft drinks containing carbon dioxide gas dissolved in water.
But, all gas-liquid interactions are not the same. For instance, HCl gas readily dissolves in water, forming hydrochloric acid, while oxygen is less soluble.
The vapor pressure of volatile components in a solution is governed by Henry's law and Raoult's law, which state that the vapor pressure is proportional to its mole fraction.
The difference is that Raoult's law uses the vapor pressure of the pure component as the proportionality constant, while Henry's law relies on an experimentally determined value.
On plotting vapor pressure against mole fraction, Raoult's law predicts pressures at high mole fractions.
Henry's law applies at low mole fractions, giving an approximately straight line.
Henry’s law relates to sparingly soluble gases in various solvents, such as water, where the dissolved gas concentration is usually low enough for the solution to be ideally dilute.
The solubility of most nonpolar gases in water decreases initially as temperature increases. However, depending on the gas and conditions, it increases sharply as the critical temperature of water is approached.
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