Give two exceptions to Henry's law.

Under everyday circumstances, Henry's Law provides us with a reliable relation between the pressure of a gas and its solubility in liquids. When applying a low pressure to a gas, we find that solubility is directly proportional to this pressure. However, this principle becomes less dependable under high pressure conditions.

When the pressure exceeds a certain threshold, the behavior of gases significantly changes. At these high pressures, the molecular volume and the interactions between gas molecules can no longer be overlooked. These factors contribute to a deviation from the ideal gas behavior that Henry's Law presumes. As a result, you may observe a plateau or even a decrease in solubility, which diverges from the linear relationship that Henry's Law predicts. This phenomenon is particularly important to consider in industrial processes, where gases are often subjected to exceedingly high pressures.

Henry's Law stands on the assumption that gases behave ideally, meaning their molecules do not attract or repel each other and they occupy no space. Real gases, however, seldom comply with this idealized version, especially under certain conditions.

As pressure increases, the interactions among gas molecules become significant, and the occupied space by these molecules can't be ignored. This results in non-ideal gas behavior, which is described by real-world equations, such as the Van der Waals equation, that take these interactions into account. This non-ideal behavior plays a crucial role in understanding how gases will actually dissolve in liquids at varying pressures, and it explains why at high pressures, the predictions of Henry's Law may fail.

Temperature is a critical factor that impacts the solubility of gases in liquids, thereby influencing the applicability of Henry's Law. Typically, Henry's Law is relevant at a standard temperature, commonly around room temperature.

As the temperature increases, gas molecules gain kinetic energy, which tends to make them escape from the liquid, leading to decreased solubility. Conversely, a decrease in temperature usually increases solubility as gas molecules have less energy to leave the liquid phase. Consequently, at extreme temperatures, the solubility of gases does not vary linearly with pressure, challenging the precision of Henry's Law. This temperature-solubility relationship is key in various applications, for example, when managing dissolved oxygen levels in aquatic environments or carbonation in the beverage industry.