Cameron Miranda
Henry's Law is a gas law in physical chemistry 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. This constant is highly temperature-dependent because vapour pressure and solubility are both temperature-dependent. The Van 't Hoff equation describes the temperature dependence of equilibrium constants and can be used to adjust Henry's law constants for different temperatures. The equation takes into account the standard state, temperature, and enthalpy of solvation for each gas. The maximum temperature of the Henry's law constant also varies with the size of the gas molecule and its solubility in water.
| Characteristics | Values |
|---|---|
| Henry's Law Constant Temperature Dependence | The constant changes with temperature |
| Temperature Dependence Equation | Van 't Hoff equation: \(k(T)=k(T^{\circ})exp\left[-C\left(\frac{1}{T}-\frac{1}{T\circ}\right)\right]\) |
| Solubility and Temperature | Solubility of gases generally decreases with increasing temperature |
| Aqueous Solutions | The solubility constant goes through a minimum for many species |
| Gas Molecule Size | Smaller gas molecules have a lower temperature maximum |
| Example Minimums | Helium: 30 °C, Argon/Nitrogen/Oxygen: 92-93 °C, Xenon: 114 °C |
| Pressure and Solubility | Solubility increases with greater depth and pressure |
| Decompression Sickness | Divers can experience this due to gas accumulation in body tissues at greater depths |
| Carbonated Drinks | Henry's Law applies to the release of CO2 when a bottle is opened |
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What You'll Learn
The Van 't Hoff equation
> {\displaystyle H(T)=H^{\circ }\exp \left [ {\frac {-\Delta _{\text{sol}}H}{R}}\left ({\frac {1}{T}}-{\frac {1}{T^{\circ }}}\right)\right] }.
> {\displaystyle {\frac {d\ln K_{\mathrm {eq} }}{d{\frac {1}{T}}}}=-{\frac {\Delta _{r}H^{\ominus }}{R}}.
This equation is used to estimate a new equilibrium constant at a new absolute temperature, assuming a constant standard enthalpy change over the temperature range. The Van 't Hoff equation can be derived from thermodynamics, and it is important to note that it is an approximation that is strictly true only for dilute ideal solutions. The equation depicts the relationship between the solubility of a specific solute and the temperature of the solution, taking into account the solvent effect.
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Temperature dependence
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 only valid when the molecules in the system are in equilibrium and it does not apply when the gas and solution are involved in a chemical reaction.
Henry's law constant is highly temperature-dependent because vapour pressure and solubility are both temperature-dependent. The temperature dependence of equilibrium constants can be described using the Van 't Hoff equation:
> {\displaystyle H(T)=H^{\circ }\exp \left [{\frac {-\Delta _{\text{sol}}H}{R}}\left ({\frac {1}{T}}-{\frac {1}{T^{\circ}}}\right)\right].}
The Van 't Hoff equation is only valid for a limited temperature range where {\displaystyle \Delta _{\text{sol}}H} does not change much with temperature (around 20K of variation). The solubility of permanent gases usually decreases with increasing temperature at around room temperature. However, for aqueous solutions, the Henry's law solubility constant for many species goes through a minimum. For most permanent gases, this minimum is below 120 °C.
The size of the gas molecule also affects the temperature of the maximum of the Henry's law constant. Smaller gas molecules tend to have lower solubility in water, and lower temperatures for the maximum of the constant. For example, the maximum is at about 30 °C for helium, 92 to 93 °C for argon, nitrogen and oxygen, and 114 °C for xenon.
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Gas molecule size
Henry's law is a gas law formulated by English chemist William Henry in 1803. It states that at a constant temperature, the amount of a gas that dissolves in a liquid is directly proportional to the partial pressure of that gas in equilibrium above the liquid. The constant of proportionality in this relationship is known as Henry's law constant, often denoted by 'kH'.
Henry's law constants are highly temperature-dependent because vapour pressure and solubility are both temperature-dependent. When the temperature of a system changes, the Henry's law constant also changes. The Van 't Hoff equation can be used to describe the temperature dependence of the equilibrium constant. The solubility of permanent gases usually decreases with increasing temperature at around room temperature.
The size of a gas molecule influences the maximum temperature at which the Henry's law constant is observed. Generally, the smaller the gas molecule, the lower the temperature at which the maximum of the Henry's law constant occurs. For instance, the maximum is observed at about 30 °C for helium, a gas with small molecules, while for larger molecules like xenon, the maximum occurs at a much higher temperature of 114 °C.
It is important to note that Henry's law has certain limitations. It is only applicable when the molecules in the system are in a state of equilibrium and at constant temperature. The law does not hold when gases are subjected to extremely high pressures or when the gas and solution undergo chemical reactions with each other.
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Vapour pressure
Henry's Law 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 known as Henry's Law Constant. The constant changes with temperature, and this relationship can be described using the Van 't Hoff equation. The Van 't Hoff equation is only valid for a limited temperature range, around 20K of variation, and the solubility of gases generally decreases with increasing temperature.
The vapour pressure of a liquid is influenced by the strength of intermolecular forces between its molecules. Liquids with strong intermolecular forces tend to have smaller vapour pressures because the attractive interactions between the molecules become less significant as the temperature increases, making it easier for the molecules to escape into the gas phase. Conversely, liquids with weaker intermolecular forces tend to have higher vapour pressures.
The vapour pressure of a substance can be measured by injecting a small amount of the liquid into a closed flask connected to a manometer. This allows for the determination of the equilibrium pressure of the vapour above the liquid or solid sample.
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Solubility
The solubility of a gas in a liquid can be influenced by various factors, including temperature and pressure. When it comes to temperature, Henry's Law Constant is highly temperature-dependent. This is because vapour pressure and solubility are both influenced by temperature. As temperature increases, the solubility of permanent gases tends to decrease at around room temperature. However, for aqueous solutions, the Henry's Law solubility constant often goes through a minimum, with the minimum occurring below 120 °C for most permanent gases. This minimum can vary depending on the specific gas, with helium reaching its maximum at around 30 °C, argon, nitrogen, and oxygen at 92 to 93 °C, and xenon at 114 °C.
The Van 't Hoff equation can be used to describe the temperature dependence of Henry's Law constants:
{\displaystyle H(T)=H^{\circ }\exp \left[{\frac {-\Delta _{\text{sol}}H}{R}}\left({\frac {1}{T}}-{\frac {1}{T^{\circ }}}\right)\right].}
This equation takes into account the change in enthalpy of solvation (represented by ) and the ideal gas constant . It is important to note that this equation is only valid within a limited temperature range where does not vary significantly with temperature, typically around 20K of variation.
Additionally, the solubility of gases is influenced by pressure. According to Henry's Law, as the pressure of a gas above a liquid increases, the solubility of the gas in the liquid also increases. This is particularly relevant in understanding the behaviour of gases dissolved in liquids at different depths, such as in underwater diving. At greater depths, the solubility of gases in body tissues increases, and during decompression, the solubility decreases.
To summarise, solubility is a critical aspect of Henry's Law, and the solubility of gases can be adjusted by varying temperature and pressure. The Van 't Hoff equation provides a way to account for temperature changes in Henry's Law constants, while the law itself describes the relationship between solubility and pressure.
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Frequently asked questions
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 its partial pressure above the liquid at equilibrium.
Henry's Law constants are highly temperature-dependent because vapour pressure and solubility vary with temperature. The Van 't Hoff equation describes this temperature dependence:
H(T)=H^° exp [(-ΔsolH/R)(1/T - 1/T°)]
The Henry's Law constant is typically expressed in units of mol L–1 bar–1.
No, Henry's Law constants vary for different gases and solvents. It also does not apply when gases are under extremely high pressure or when the gas and solvent undergo chemical reactions with each other.
Carbonated drinks, like a bottle of Pepsi, provide a common example of Henry's Law. When the bottle is opened, the pressurised carbon dioxide (CO2) above the liquid escapes, reducing the partial pressure of CO2. This decrease in partial pressure causes the dissolved CO2 in the drink to escape as tiny bubbles, resulting in a flat drink if left open for long enough.
