O2 769.23 1.3 E-3 4.259 E4 3.180 E-2
H2 1282.05 7.8 E-4 7.099 E4 1.907 E-2
CO2 29.41 3.4 E-2 0.163 E4 0.8317
N2 1639.34 6.1 E-4 9.077 E4 1.492 E-2
He 2702.7 3.7 E-4 14.97 E4 9.051 E-3
Ne 2222.22 4.5 E-4 12.30 E4 1.101 E-2
Ar 714.28 1.4 E-3 3.955 E4 3.425 E-2
CO 1052.63 9.5 E-4 5.828 E4 2.324 E-2
where:
= moles of gas per liter of solution = liters of solution
= partial pressure of gas above the solution, in atmospheres of absolute pressure
= mole fraction of gas in solution = moles of gas per total moles ≈ moles of gas per mole of water
= atmospheres of absolute pressure.
As can be seen by comparing the equations in the above table, the Henry's Law constant kH,pc is simply the inverse of the constant kH,cp. Since all kH may be referred to as the Henry's Law constant, readers of the technical literature must be quite careful to note which version of the Henry's Law equation is being used.[2]
It should also be noted the Henry's Law is a limiting law that only applies for dilute enough solutions. The range of concentrations in which it applies becomes narrower the more the system diverges from non-ideal behavior. Roughly speaking, that is the more chemically different the solute is from the solvent.
It also only applies for solutions where the solvent does not react chemically with the gas being dissolved. A common example of a gas that does react with the solvent is
carbon dioxide, which rapidly forms hydrated carbon dioxide and then carbonic acid (H2CO3) with water.
[edit] Temperature dependence of the Henry constant
When the temperature of a system changes, the Henry constant will also change.[2] This is why some people prefer to name it Henry coefficient. There are multiple equations assessing the effect of temperature on the constant. This form of the van 't Hoff equation is one example:[4] where
k for a given temperature is the Henry's Law constant (as defined in the first section of this article), identical with kH,pc defined in Table 1, T is in Kelvin,
the index Θ (theta) refers to the standard temperature (298 K).
The above equation is an approximation only and should be used only when no better experimentally derived formula for a given gas exists.
The following table lists some values for constant C (dimension of kelvins) in the equation above:
Table 2: Values of C Gas O2 H2 CO2 N2 He Ne Ar CO C 1700 500 2400 1300 230 490 1300 1300 Because solubility of gases is decreasing with increasing temperature, the partial pressure a given gas concentration has in liquid must increase. While heating water (saturated with nitrogen) from 25 °C to 95 °C the solubility will decrease to about 43% of its initial value. This can be verified when heating water in a pot. Small
bubbles evolve and rise, long before the water reaches boiling temperature. Similarly, carbon dioxide from a carbonated drink escapes much faster when the drink is not cooled because of the increased partial pressure of CO2 in higher temperatures. Partial pressure of CO2 in seawater doubles with every 16 K increase in temperature.[5] The constant C may be regarded as: where
is the enthalpy of solution R is the gas constant.
[edit] Henry's law in geophysics
In geophysics a version of Henry's law applies to the solubility of a noble gas in contact with silicate melt. One equation used is where:
subscript m = melt subscript g = gas phase
ρ = the number densities of the solute gas in the melt and gas phase β = 1 / kBT an inverse temperature scale kB = the Boltzmann constant
μex,m and μex,g = the excess chemical potential of the solute in the two phases.
[edit] Henry's law versus Raoult's law
Both Henry's law and Raoult's law state that the vapor pressure of a component, p, is proportional to its concentration.
Henry's law: Raoult's law:
where:
is the mole fraction of the component;
is the Henry constant; (Note that the numerical value and dimensions of this constant change when mole fractions are used rather than molarity, as seen in Table 1.) is the equilibrium vapor pressure of the pure component.
If the solution is ideal, both components follow Raoult's law over the entire composition range, but Henry noticed that at low concentrations of non-ideal solutions, the constant of proportionality is not p*. Therefore Henry's law uses an empirically-derived constant, k, based on an infinitely-dilute solution, i.e. x = 0, that is specific to the components in the mixture and the temperature.
In most systems, the laws can only be applied over very limited concentrations at the extreme ends of the mole-fraction range. Raoult's law, which uses the vapor pressure of the pure component, is best used for the major component (solvent) and in
mixtures of similar components. Henry's law applies to the minor component (solute) in dilute solutions.
In ideal-dilute solutions, the minor component follows Henry's law, while the solvent obeys Raoult's law. This is proved by the Gibbs-Duhem equation.
Dalton's law
In chemistry and physics, Dalton's law (also called Dalton's law of partial
pressures) states that the total pressure exerted by a gaseous mixture is equal to the sum of the partial pressures of each individual component in a gas mixture. This empirical law was observed by John Dalton in 1801 and is related to the ideal gas laws.
Mathematically, the pressure of a mixture of gases can be defined as the summation
P total =P1+P2+…+Pn
where P1, P2, Pn represent the partial pressure of each component. It is assumed that the gases do not react with each other.
Pi=P total Xi
where Xi = the mole fraction of the i-th component in the total mixture of m components .
The relationship below provides a way to determine the volume based concentration of any individual gaseous component. Pi=P total Ci/1000000
Where, Ci is the concentration of the i-th component expressed in ppm.
Dalton's law is not exactly followed by real gases. Those deviations are considerably large at high pressures. In such conditions, the volume occupied by the molecules can become significant compared to the free space between them. Moreover, the short average distances between molecules raises the intensity of intermolecular forces between gas molecules enough to substantially change the pressure exerted by them. Neither of those effects are considered by the ideal gas model.