In a mixture of , each constituent gas has a partial pressure which is the notional pressure of that constituent gas as if it alone occupied the entire volume of the original mixture at the same temperature. The total pressure of an ideal gas mixture is the sum of the partial pressures of the gases in the mixture (Dalton's Law).
In respiratory physiology, the partial pressure of a dissolved gas in liquid (such as oxygen in arterial blood) is also defined as the partial pressure of that gas as it would be undissolved in gas phase yet in equilibrium with the liquid. This concept is also known as blood gas tension. In this sense, the diffusion of a gas liquid is said to be driven by differences in partial pressure (not concentration). In chemistry and thermodynamics, this concept is generalized to non-ideal gases and instead called fugacity. The partial pressure of a gas is a measure of its thermodynamic activity. Gases dissolve, diffuse, and react according to their partial pressures and not according to their Concentration in a gas mixture or as a solute in solution.Collman, J. P.; Brauman, J. I.; Halbert, T. R.; Suslick, K. S. (1976). “ Nature of O2 and CO binding to metalloporphyrins and heme proteins”. Proceedings of the National Academy of Sciences. 73 (10): 3333-3337. This general property of gases is also true in chemical reactions of gases in biology.
Examples:
where:
and the partial pressure of an individual gas component in an ideal gas can be obtained using this expression:
where: | |
The mole fraction of a gas component in a gas mixture is equal to the volumetric fraction of that component in a gas mixture. Frostberg State University's "General Chemistry Online"
The ratio of partial pressures relies on the following isotherm relation:
It can be approximated both from partial pressure and molar fraction:Page 200 in: Medical biophysics. Flemming Cornelius. 6th Edition, 2008.
The higher the vapor pressure of a liquid at a given temperature, the lower the normal boiling point of the liquid.
The vapor pressure chart displayed has graphs of the vapor pressures versus temperatures for a variety of liquids.
For example, at any given temperature, methyl chloride has the highest vapor pressure of any of the liquids in the chart. It also has the lowest normal boiling point (−24.2 °C), which is where the vapor pressure curve of methyl chloride (the blue line) intersects the horizontal pressure line of one atmosphere (atm) of absolute vapor pressure. At higher altitudes, the atmospheric pressure is less than that at sea level, so boiling points of liquids are reduced. At the top of Mount Everest, the atmospheric pressure is approximately 0.333 atm, so by using the graph, the boiling point of diethyl ether would be approximately 7.5 °C versus 34.6 °C at sea level (1 atm).
the equilibrium constant of the reaction would be:
where: | |
For reversible reactions, changes in the total pressure, temperature or reactant concentrations will shift the equilibrium so as to favor either the right or left side of the reaction in accordance with Le Chatelier's Principle. However, the reaction kinetics may either oppose or enhance the equilibrium shift. In some cases, the reaction kinetics may be the overriding factor to consider.
where:
The form of the equilibrium constant shows that the concentration of a solute gas in a solution is directly proportional to the partial pressure of that gas above the solution. This statement is known as Henry's law and the equilibrium constant is quite often referred to as the Henry's law constant. Introductory University Chemistry, Henry's Law and the Solubility of Gases
Henry's law is sometimes written as:
where is also referred to as the Henry's law constant. As can be seen by comparing equations () and () above, is the reciprocal of . Since both 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.
Henry's law is an approximation that only applies for dilute, ideal solutions and for solutions where the liquid solvent does not react chemically with the gas being dissolved.
Using diving terms, partial pressure is calculated as:
For the component gas "i":
For example, at underwater, the total absolute pressure is (i.e., 1 bar of atmospheric pressure + 5 bar of water pressure) and the partial pressures of the main components of air, oxygen 21% by volume and nitrogen approximately 79% by volume are:
The minimum safe lower limit for the partial pressures of oxygen in a breathing gas mixture for diving is absolute. Hypoxia and sudden unconsciousness can become a problem with an oxygen partial pressure of less than 0.16 bar absolute.
Narcosis is a problem when breathing gases at high pressure. Typically, the maximum total partial pressure of narcotic gases used when planning for technical diving may be around 4.5 bar absolute, based on an equivalent narcotic depth of .
The effect of a toxic contaminant such as carbon monoxide in breathing gas is also related to the partial pressure when breathed. A mixture which may be relatively safe at the surface could be dangerously toxic at the maximum depth of a dive, or a tolerable level of carbon dioxide in the breathing loop of a diving rebreather may become intolerable within seconds during descent when the partial pressure rapidly increases, and could lead to panic or incapacitation of the diver.
In diving breathing gases
where:
In medicine
+ for and
! !! Unit !! Arterial blood gas !! vein blood gas !! Cerebrospinal fluid !! Alveolar pulmonary
gas pressures14.2 107 4.8 36
See also
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