Trystan Casey
Henry's law, formulated by English chemist William Henry in 1803, is a gas law that illustrates the relationship between the solubility of a gas and its partial pressure. The law is expressed through the formula: solubility of gas = Henry's Law Constant x Partial Pressure of the Gas. The solubility of a gas in a liquid is directly proportional to the partial pressure of that gas above the liquid. The law is only applicable when the molecules are in equilibrium, and it does not apply to gases at high pressures. The value of the Henry's Law Constant varies with temperature and solvent and can be calculated by determining the ratio of the concentration of a gas in a liquid to the partial pressure of the gas above the liquid when the system is at equilibrium.
| Characteristics | Values |
|---|---|
| What is Henry's Law? | A gas law formulated by William Henry in 1803 that states: "At a constant temperature, the amount of a given gas that dissolves in a given type and volume of liquid is directly proportional to the partial pressure of that gas in equilibrium with that liquid." |
| When is it applicable? | When the molecules are in equilibrium. It does not apply to gases at high pressures or when the solution and gas are involved in a chemical reaction with each other. |
| How to calculate it | The formula to calculate Henry's Law constant (kH) is: solubility of gas = kH * partial pressure of gas. The units of Henry's Law constant will depend on the units used for concentration and pressure. It's important to maintain consistency in the units when applying the formula. |
| Factors it depends on | Temperature and solvent. |
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What You'll Learn
Calculating Henry's law constant
Henry's law, formulated by English chemist William Henry in the early 19th century, is a gas law that states that the amount of dissolved gas in a liquid is directly proportional to its partial pressure above the liquid when the molecules are in equilibrium. The proportionality factor is called Henry's law constant.
Henry's law constant is expressed in mol L–1 bar–1. The value of the constant varies with temperature and solvent and is highly temperature-dependent because vapour pressure and solubility are both temperature-dependent. The constant is also influenced by the units used for concentration and pressure, which must be consistent. The units for pressure can be in atmospheres, torrs, or millimetres of mercury.
The formula to calculate Henry's Law constant (kH) is:
> kH = C/P
Where:
- C is the concentration of the dissolved gas in the liquid (solubility)
- P is the partial pressure of the gas above the liquid
For example, to calculate the solubility of carbon dioxide (CO2) when its Henry's law constant is 8.20 * 10^2 molarities per atmosphere at a pressure of 3.29 atmospheres, you would use the formula:
> S gas, CO2 = Henry's constant x Pressure of gas
> S gas, CO2 = 8.20 * 10^2 x 3.29 atmospheres
Henry's law is applicable in various scenarios, such as the change in the dissolution of oxygen and nitrogen in the blood of underwater divers during decompression and the release of pressurised carbon dioxide when opening a carbonated beverage.
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Henry's law and respiration
Henry's law, named after the English physician William Henry, defines the relationship between the partial pressure of gases overlying a solution and the gases' ability to dissolve in that solution. In other words, it states that the amount of a gas that dissolves in a liquid is directly proportional to the partial pressure of that gas above the liquid. This law is particularly relevant in respiratory physiology, where it helps predict how gases dissolve in the alveoli and bloodstream during gas exchange.
During respiration, the primary gases involved are oxygen and carbon dioxide. Oxygen has a larger partial pressure gradient, allowing it to diffuse into the bloodstream more readily. The partial pressure of oxygen is higher in alveolar air than in deoxygenated blood, making oxygen more likely to dissolve into the latter. Conversely, carbon dioxide has a higher partial pressure in deoxygenated blood, causing it to diffuse out of the solution and back into its gaseous form.
The solubility of a gas in a liquid is influenced by its partial pressure. Gases with higher partial pressures have more molecules and are more likely to dissolve. This principle can be observed in everyday examples like carbonated soft drinks. Before opening the container, the gas above the drink is primarily carbon dioxide, which is under pressure. When the container is opened, the pressure drops, causing some of the dissolved carbon dioxide to escape as bubbles. Over time, the concentration of carbon dioxide in the drink equalizes with the surrounding air, and the drink becomes flat.
Understanding Henry's Law is crucial in certain medical specialties, such as aerospace medicine. As individuals ascend from sea level, the pressure and oxygen content decrease while altitude increases. This can lead to various physiological changes, and understanding Henry's Law helps predict the behaviour of gases within the body during these changes. By applying Henry's Law, medical professionals can better safeguard the health of pilots, aircrew, and passengers in dynamic atmospheric conditions.
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Henry's law and underwater diving
Henry's Law is a gas law that states that the amount of dissolved gas in a liquid is directly proportional at equilibrium to its partial pressure above the liquid. The proportionality factor is called Henry's Law Constant. The law is expressed through the formula:
> Sgas = K × Pgas
Where:
- Sgas is the solubility of the gas
- K is Henry's Law Constant
- Pgas is the partial pressure of the gas
The concentration, also referred to as the solubility of a dissolved gas, can be determined from its Henry's Law Constant and partial pressure. Henry's Law Constant can be calculated by determining the ratio of the concentration of a gas in a liquid to the partial pressure of the gas above the liquid when the system is at equilibrium.
Henry's Law is highly relevant to underwater diving. Gas is breathed at ambient pressure, which increases with depth due to hydrostatic pressure. According to Henry's Law, the solubility of gases increases with greater depth (greater pressure), so the body tissues of a diver take on more gas over time in greater depths of water. When ascending, the diver is decompressed and the solubility of the gases dissolved in the tissues decreases, so the excess gas is carried away by the blood and released into the lung gas.
If a diver ascends too quickly, the supersaturation may be too great, causing gas bubbles to form and grow. These bubbles can cause blockages in capillaries or distortion in the more solid tissues, resulting in damage known as decompression sickness. To avoid this injury, the diver must ascend slowly.
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Henry's law and carbonated drinks
Henry's law is a gas law formulated by English chemist William Henry in the early 19th century. It states that the amount of dissolved gas in a liquid is directly proportional to the partial pressure of that gas above the liquid, provided the temperature is kept constant. The constant of proportionality is called Henry's law constant, often denoted by 'kH'. The law is expressed through the formula:
> Sgas = Henry's constant x Pgas
Where Sgas is the solubility of the gas and Pgas is the partial pressure of the gas, usually expressed in atmospheres, atm, or Pascals, Pa.
Henry's law is applicable when the molecules are in equilibrium. It does not apply to gases at high pressures or when the gas and solution chemically react with each other. The law is highly temperature-dependent because vapour pressure and solubility are temperature-dependent.
Carbonated drinks provide a common example of Henry's law in action. Before opening a carbonated drink, the gas above the liquid is almost pure carbon dioxide, maintained at a pressure slightly higher than atmospheric pressure. As per Henry's law, the solubility of carbon dioxide in the unopened drink is high. When the bottle is opened, the pressurised CO2 escapes into the atmosphere, causing a rapid decrease in the partial pressure of CO2 above the liquid. Consequently, the solubility of carbon dioxide in the drink also decreases, leading to the formation of tiny bubbles as the dissolved CO2 escapes into the atmosphere. This process is known as degassing.
If the carbonated drink is left open for long enough, the concentration of carbon dioxide in the drink will reach an equilibrium with the concentration of carbon dioxide in the atmosphere (around 0.05%). At this point, the drink will go flat as the concentration of carbon dioxide in the drink matches the atmosphere, and no further degassing occurs.
Henry's law also explains why beer served directly from a tap (gravity) is less carbonated than beer served via a hand pump. Beer served via a hand pump is pressurised on its way to the point of service, causing carbon dioxide to dissolve in the beer and increasing its carbonation.
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Henry's law and temperature
Henry's law is a gas law formulated in the early 19th century by English chemist William Henry. The law states that the amount of gas that is dissolved in a liquid is directly proportional to the partial pressure of that gas above the liquid when the temperature is kept constant. The constant of proportionality for this relationship is called Henry's law constant (often denoted by 'kH'). The law is expressed through the formula:
> Sgas = kH * Pgas
Where Sgas is the solubility of the gas and Pgas is the partial pressure of the gas, usually in atmospheres, torrs, or millimetres of mercury. It's important to maintain consistency in units when applying the formula.
Henry's law is only applicable when the molecules are in equilibrium. It does not apply to gases at high pressures or when the gas and solution undergo a chemical reaction. The law's constants are highly temperature-dependent because vapour pressure and solubility are both temperature-dependent.
An example of Henry's law in action is the opening of a carbonated drink. When the bottle is unopened, the gas above the drink is usually pure carbon dioxide, kept at a pressure slightly above atmospheric pressure. As a result, the solubility of carbon dioxide in the drink is high. When the bottle is opened, the pressurised CO2 escapes into the atmosphere, decreasing the partial pressure of CO2 above the drink and, in turn, decreasing the solubility of the gas in the drink.
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Frequently asked questions
Henry's Law is a gas law formulated in the early 19th century by English chemist William Henry. It states that the amount of dissolved gas in a liquid is directly proportional to the partial pressure of that gas above the liquid when the temperature is constant.
An unopened carbonated drink contains gas, usually carbon dioxide, at a pressure slightly above atmospheric pressure. Due to Henry's Law, the solubility of the gas is high. When the drink is opened, the pressure decreases, and so does the solubility, causing the gas to escape into the atmosphere.
As a diver descends, the ambient pressure and solubility of gases increase. This causes the body tissues to absorb more gas over time. When ascending, the diver must do so slowly to allow the excess dissolved gas to be carried away by the blood and released through the lungs, thus avoiding decompression sickness.
The formula for calculating Henry's Law Constant (kH) is: kH = P / C, where P is the partial pressure of the gas above the liquid, and C is the concentration of the gas in the liquid. The units for P can be atmospheres, torrs, or millimeters of mercury, while the units for C can be molarity or mass per volume.
Henry's Law constants are highly temperature-dependent because vapour pressure and solubility vary with temperature. As temperature increases, solubility decreases, and vice versa. Therefore, the Henry's Law constant is inversely proportional to temperature.
