MCQ Detailed View

The solubility of gases in liquids is an important concept in physical chemistry. It is explained by Henry’s Law, which states:

C=kH⋅PC = k_H \cdot PC=kH⋅P

Here, C is the concentration of the gas in the liquid, P is the partial pressure of the gas above the liquid, and k_H is Henry’s Law constant, which depends on the nature of the gas, the liquid, and temperature. This law shows that if the partial pressure of a gas increases, the amount of gas that dissolves in the liquid also increases proportionally.

This relationship is widely seen in everyday life and in industrial applications. For example:

• 
  

  Carbonated beverages like soda are bottled under high pressure of carbon dioxide. When the bottle is sealed, the high pressure keeps a large amount of CO₂ dissolved in the liquid. Once opened, the pressure is released, and gas escapes, forming bubbles.

  
  


• 
  

  Scuba diving provides another application. At greater depths, the partial pressure of gases such as nitrogen increases, leading to more nitrogen dissolving in the diver’s blood. If the diver ascends too quickly, the dissolved nitrogen comes out of solution rapidly, causing decompression sickness (the bends).

  
  


• 
  

  In respiration physiology, the exchange of gases like oxygen and carbon dioxide between blood and alveolar air also follows the principle of Henry’s Law.

  
  

Comparing with other laws:

• 
  

  Dalton’s Law deals with total pressure of gas mixtures, not solubility.

  
  


• 
  

  Gay-Lussac’s Law explains pressure–temperature relationships for gases.

  
  


• 
  

  Raoult’s Law describes vapor pressure lowering in solutions.

  
  

Only Henry’s Law directly explains the proportional relationship between the partial pressure of a gas and its solubility in a liquid. This principle is fundamental in solution chemistry, chemical engineering, and biological systems.