Partial pressure

The partial pressure is the pressure which is associated in an (ideal) gas mixture of a single component. The partial pressure corresponds to the pressure which the individual gas component would exert on the sole in the presence of the relevant volume.

Distinction

In meteorology, the term vapor pressure is used as a synonym for the partial pressure of water in the air. In a gas mixture such as air, the boiling or condensation temperature of a gas component (e.g., water vapor) is the one that is associated with the partial pressure of the component, not the total pressure. It is referred to as dew point in the field of meteorology.

Application

In biology and medicine, are primarily of the oxygen and carbon dioxide partial pressure is very important. Herein, the term is also applied to the concentrations of these gases in solution, for example in the blood or in water. This is expressed as partial pressure of that pressure of the gas which communicates with the respective concentration in the solution ( on an imaginary or real interface of gas and liquid) in a diffusion equilibrium. Henry's law describes macroscopically the ratio of the partial pressures that occur in the gas phase, depending on the respective concentrations in the liquid phase. The partial pressure is always used instead of the mass concentration when the diffusion behavior of the dissolved gas will be considered. Typical topics for this are the respiratory exchanges in the lungs, the risk of gas embolism in diving and aviation medicine and the formation of gas bubble disease of fish. Because of this, the calculation of partial pressures in technical scuba diving or nitrox basis of the associated courses.

Dalton's law

The Dalton's law ( Dalton 's law, law of partial pressures ), formulated in 1805 by John Dalton, stating that the sum of all partial pressures in ideal gases is equal to the total pressure of the mixture. For components results

And followed by conversion

If an area over a gas mixture is not practical or not, the resulting over gas space under the two formulas is filled. Here is the partial pressure of a component of the ruling at this temperature boiling pressure of the pure component.

It follows from this that the partial pressure of th gas is equal to the product of mole fraction of the gas times the total pressure of the mixture (eg, air pressure). While this is an idealized representation of the event that the particles of the gas phase out of the mechanical mutual interaction have not ( ideal gas ), but this can be ignored in most cases.

The ratio of the number of particles (molar ) of a component to the total number of particles of the mixture corresponds to the mole fraction, and thus the partial pressure of the component to the total pressure of the mixture. It therefore applies:

Example: dry air at sea level

The following table shows the composition of completely dry air at sea level (normal condition) so at an air pressure of 1013.25 hPa

One can equate ( mole fraction ) the volume fraction of the particle fraction, as it is nearly ideal gases. Other parts of the air can be neglected due to their small share.

Airworthiness

For the issue of an airworthiness of passengers expected at altitude oxygen partial pressure should be above 50-55 mmHg.