Hydrogen gas has a solubility in water of \(0.00157 \mathrm{~g} \mathrm{~L}^{-1}\) under 1.00 atm of \(H_{2}\) pressure at \(25^{\circ} \mathrm{C}\). At 0.85 atm and \(25^{\circ} \mathrm{C}\), its solubility is \(0.00133 \mathrm{~g} \mathrm{~L}^{-1}\). Does hydrogen gas obey Henry's law?

Understanding the solubility of gases in liquids is crucial for comprehending a range of natural and industrial processes, including how fish breathe underwater and how carbonated beverages are made. At its core, the solubility of a gas refers to how much of that gas can dissolve in a liquid at a given temperature and pressure.

The solubility is primarily influenced by factors such as the nature of the gas and liquid, temperature, and pressure. Generally, as temperature increases, the solubility of gases in liquids decreases because most gases are less soluble in hotter liquids. Conversely, an increase in pressure typically enhances the amount of gas that can be dissolved, which is precisely what Henry's Law explores.

Real-World Examples

• In carbonated drinks, carbon dioxide is dissolved in the liquid under high pressure, which is why bubbles form when the pressure is released as you open the bottle.
• Scuba divers must consider the solubility of gases such as nitrogen in their blood to prevent decompression sickness, often known as 'the bends'.

Partial pressure is a key concept when discussing gaseous systems such as our atmosphere, where various gases are present, each exerting its own pressure independent of the others. The total pressure of the system is the sum of all these individual pressures. According to Dalton's Law of Partial Pressures, this sum equals the total pressure exerted by the gas mixture.

It's useful to think of partial pressure as the contribution of each gas in a mixture to the overall pressure. For example, in a container with hydrogen and oxygen gases, the total pressure is a combination of the pressure from hydrogen molecules bouncing against the container's walls and the pressure from the oxygen molecules doing the same.

Everyday Applications

• Medical applications such as determining the correct mixtures of gases for anesthesia
• Predicting the behavior of gases at high altitudes, which is valuable for climbers and pilots

Gas-liquid equilibrium is a state where a liquid and a gas phase are in balance with each other, with no net transfer of molecules from the gas to the liquid or vice versa. At this point, the rate of gas molecules entering the liquid is equal to the rate of gas molecules escaping. Understanding this concept allows us to predict the behavior of gases in liquids under various conditions.

The equilibrium is dynamic, meaning that although the overall quantities of gas and liquid stay constant, individual molecules continually switch between the liquid and gas phases.

Factors Influencing Equilibrium

• Temperature: Typically, increasing temperature shifts the equilibrium, favoring gas formation because higher kinetic energy makes it easier for molecules to escape the liquid.
• Pressure: The equilibrium can also be shifted by changes in pressure. Henry's Law provides a useful guide to understanding how changes in pressure can affect the solubility of a gas in a liquid at equilibrium.