Everything about Partial Pressure totally explained
In a mixture of
ideal gases, each gas has a
partial pressure which is the pressure which the gas would have if it alone occupied the volume. The total
pressure of a gas mixture is the sum of the partial pressures of each individual gas in the mixture.
In
chemistry, the partial pressure of a
gas in a mixture of gases is defined as above. The partial pressure of a gas dissolved in a liquid is the partial pressure of that gas which would be generated in a gas phase in equilibrium with the liquid at the same temperature. The partial pressure of a gas is a measure of thermodynamic activity of the gas's
molecules. Gases will always flow from a region of higher partial pressure to one of lower pressure; the larger this difference, the faster the flow. Gases dissolve, diffuse, and react according to their partial pressures, and not necessarily according to their
concentrations in a gas mixture.
Dalton's law of partial pressures
The partial pressure of an
ideal gas in a mixture is equal to the pressure it would exert if it occupied the same volume alone at the same temperature. This is because ideal gas molecules are so far apart that they don't interfere with each other at all. Actual real-world gases come very close to this ideal.
A consequence of this is that the total pressure of a mixture of ideal gases is equal to the sum of the partial pressures of the individual gases in the mixture as stated by
Dalton's law. For example, given an ideal gas mixture of
nitrogen (N
2),
hydrogen (H
2) and
ammonia (NH
3):
»
where:
|
|
| | = total pressure of the gas mixture
|
| | = partial pressure of nitrogen (N2)
|
| | = partial pressure of hydrogen (H2)
|
| | = partial pressure of ammonia (NH3)
|
Ideal gas mixtures
The
mole fraction of an individual gas component in an ideal gas mixture can be expressed in terms of the component's partial pressure or the
moles of the component:
» in the terms used in this article
