4.4
Gas Solubility
Gas solubility in liquids forms liquid-gas solutions, such as soft drinks, where carbon dioxide is dissolved in water, and the ocean, where the solubi…
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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Q1: Why does HCl gas dissolve more readily in water than oxygen gas?
HCl gas is highly soluble in water because it undergoes chemical interaction, forming hydrochloric acid through dissociation. Oxygen, a nonpolar gas, exhibits much lower solubility due to weak physical interactions with water molecules. Gas-liquid interactions vary significantly depending on the chemical nature of both components.
Q2: What is the difference between Henry's law and Raoult's law?
Both laws relate vapor pressure to mole fraction, but differ in their proportionality constants. Raoult's law uses the pure component's vapor pressure and applies at high mole fractions. Henry's law uses an experimentally determined constant and applies at low mole fractions where solutions are ideally dilute, such as sparingly soluble gases in water.
Q3: When does Henry's law apply to gas-liquid solutions?
Henry's law applies when dissolved gas concentration is low enough for the solution to be ideally dilute, typically at low mole fractions. For example, nitrogen and hydrogen in water follow Henry's law up to specific pressure limits. However, Henry's law does not apply to electrolytes like aqueous HCl, which dissociate completely and exhibit non-ideal behavior due to ion-ion interactions.
Q4: How does temperature affect the solubility of nonpolar gases in water?
Most nonpolar gases show decreased solubility as temperature increases at constant pressure because dissolution is typically exothermic. However, near water's critical temperature of 374°C, solubility behavior becomes undefined since liquid water ceases to exist as a distinct phase, making the concept of gas solubility in liquid water no longer meaningful.
Q5: What does Henry's law constant tell you about gas solubility?
Henry's law constant (Ki) is an experimentally determined proportionality value that depends on the gas, solvent, and temperature. A larger Ki value indicates smaller gas solubility, meaning the gas dissolves less readily. This constant allows prediction of dissolved gas concentration at given partial pressures in ideally dilute solutions.
Q6: Why do non-ideal solutions deviate from Henry's law predictions?
Non-ideal solutions deviate from Henry's law due to strong molecular interactions or chemical reactions. Electrolytes like HCl dissociate completely in water, causing ion-ion interactions that violate Henry's law assumptions. Additionally, at high pressures, real gases exhibit non-ideal behavior, and Henry's law constant becomes pressure-dependent, causing deviations from linear relationships.
Q7: How can Henry's law be expressed using molality or molar concentration instead of mole fraction?
At low solute concentrations where Henry's law applies, molality and molar concentration are each proportional to mole fraction. This allows Henry's law to be reformulated using these alternative concentration units instead of mole fractions, making calculations more practical for dilute solutions while maintaining the same proportionality relationship with partial pressure.