When air contacts a liquid, gases naturally dissolve into it. This process, where a gas enters a liquid phase, is fundamental to natural and industrial systems, such as the oxygen fish breathe or the fizz in a soft drink. Henry’s Law, formulated by William Henry in the early 19th century, governs this phenomenon. The law establishes a quantitative relationship between the pressure of a gas above a liquid and the concentration of that gas dissolved within the liquid. The Henry’s Law Constant (\(K_H\)) is the specific numerical value used to calculate this distribution.
Defining Henry’s Law and the Constant (\(K_H\))
Henry’s Law states that at a constant temperature, the amount of a gas dissolved in a liquid is directly proportional to the partial pressure of that gas above the liquid surface. For example, if the gas pressure is doubled, the amount of dissolved gas at equilibrium will also double. The law applies primarily to dilute solutions where the gas does not chemically react with the solvent.
The Henry’s Law Constant (\(K_H\)) is the constant of proportionality linking pressure and dissolved concentration. Mathematically, the law uses two primary, reciprocal forms, leading to different definitions for \(K_H\).
One common representation defines the partial pressure (\(P\)) as proportional to the concentration (\(C\)): \(P = K_H \cdot C\). In this form, \(K_H\) is the Henry’s Law volatility constant, typically carrying units like atmospheres per molarity (atm/M).
The alternate form defines the concentration as proportional to the pressure: \(C = K_H \cdot P\). This version uses the Henry’s Law solubility constant, where the units are inverted, such as moles per liter per atmosphere (M/atm). Both forms describe the same physical reality: the equilibrium partitioning of a gas between the two phases. Since the numerical value of \(K_H\) depends heavily on the unit convention used, it is important to specify which definition is being applied.
Interpreting the Constant: Solubility and Volatility
The numerical value of \(K_H\) measures a gas’s preference for the liquid or gas phase under specific conditions. When using the volatility constant definition (\(P = K_H \cdot C\)), a large \(K_H\) value signifies low solubility, meaning a high partial pressure is required to dissolve a small amount of gas. This indicates the gas prefers to remain in the atmosphere and is considered highly volatile from the solution.
Conversely, a small \(K_H\) value indicates high solubility because a low pressure achieves a high concentration in the liquid. For example, at 25°C, oxygen has a \(K_H\) in water of approximately 4.4 \(\times\) \(10^4\) atm (using mole fraction), while carbon dioxide has a smaller constant of about 1.6 \(\times\) \(10^3\) atm. This difference shows why carbon dioxide is much more soluble in water than oxygen, readily partitioning into the liquid phase.
The constant reflects how strongly gas molecules interact with the solvent compared to their self-interaction in the gas phase. Gases like helium and nitrogen are relatively insoluble in water due to weak attractive forces with water molecules, resulting in large \(K_H\) values. Highly soluble gases, such as carbon dioxide, exhibit a lower constant because they interact more favorably with the water.
Key Factors Influencing the Henry’s Law Constant
Although termed a “constant,” \(K_H\) is only invariant for a specific gas in a specific solvent at a fixed temperature. The most significant factor influencing \(K_H\) is temperature, as gas solubility in liquids is highly dependent on it. For most gases in water, solubility decreases as temperature increases, meaning the \(K_H\) value (when expressed as the volatility constant, P/C) increases with rising temperature.
This inverse relationship occurs because higher temperatures increase the kinetic energy of gas molecules in the liquid phase. This increased motion provides the dissolved gas molecules with the energy needed to overcome solvent attraction and escape back into the gas phase. This temperature dependence can be mathematically described by a conceptual relationship similar to the van’t Hoff equation, which links the constant’s change to the enthalpy of dissolution.
The nature of the solvent also affects the constant. For example, solubility in pure water differs from solubility in saltwater because the presence of dissolved salts alters the attractive forces within the liquid. This “salting out” effect typically reduces gas solubility, which increases the numerical value of \(K_H\). \(K_H\) is therefore a unique parameter for every gas-solvent-temperature combination.
Real-World Applications of Henry’s Law
Henry’s Law is a foundational principle with numerous applications across environmental science and industry.
Environmental Modeling
In environmental chemistry, the law models the exchange of gases between large bodies of water, such as oceans and lakes, and the atmosphere. Calculating the \(K_H\) for carbon dioxide allows scientists to estimate how much atmospheric \(\text{CO}_2\) is absorbed by the world’s oceans, a process that significantly influences global climate.
Chemical Engineering
The law is employed in chemical engineering for processes like gas stripping and aeration, which are designed to remove volatile contaminants from water supplies. By manipulating the partial pressure of the contaminant gas above the water, engineers can efficiently drive the unwanted substance out of the liquid phase. This principle is important for both industrial pollution control and the purification of drinking water.
Food and Beverage Industry
Henry’s Law dictates the level of carbonation in soft drinks and beer. Manufacturers dissolve carbon dioxide under high pressure to achieve the desired fizz. When the container is opened, the pressure above the liquid drops, causing the \(\text{CO}_2\) to escape as bubbles. A low \(K_H\) for \(\text{CO}_2\) ensures that significant amounts of the gas remain dissolved while the container is sealed.
Scuba Diving Physiology
The law is critical in understanding the physiology of scuba diving, particularly concerning nitrogen gas. As a diver descends, increasing ambient pressure causes more nitrogen from the breathing air to dissolve into the diver’s blood and tissues. A slow, controlled ascent is necessary to allow the dissolved nitrogen to escape harmlessly through the lungs. This prevents decompression sickness, commonly known as “the bends,” which occurs if the gas comes out of solution too rapidly.