Partial pressure
WikiDoc Resources for Partial pressure 
Articles 

Most recent articles on Partial pressure Most cited articles on Partial pressure 
Media 
Powerpoint slides on Partial pressure 
Evidence Based Medicine 
Clinical Trials 
Ongoing Trials on Partial pressure at Clinical Trials.gov Trial results on Partial pressure Clinical Trials on Partial pressure at Google

Guidelines / Policies / Govt 
US National Guidelines Clearinghouse on Partial pressure NICE Guidance on Partial pressure

Books 
News 
Commentary 
Definitions 
Patient Resources / Community 
Patient resources on Partial pressure Discussion groups on Partial pressure Patient Handouts on Partial pressure Directions to Hospitals Treating Partial pressure Risk calculators and risk factors for Partial pressure

Healthcare Provider Resources 
Causes & Risk Factors for Partial pressure 
Continuing Medical Education (CME) 
International 

Business 
Experimental / Informatics 
EditorInChief: C. Michael Gibson, M.S., M.D. [1]
Partial Pressure
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.^{[1]} 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.
Vapor pressure is the pressure of a vapor in equilibrium with its nonvapor phases (i.e., liquid or solid). Most often the term is used to describe a liquid's tendency to evaporate. It is a measure of the tendency of molecules and atoms to escape from a liquid or a solid. A liquid's boiling point corresponds to the point where its vapor pressure is equal to the surrounding atmospheric pressure.
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 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 realworld 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.^{[2]} For example, given an ideal gas mixture of nitrogen (N_{2}), hydrogen (H_{2}) and ammonia (NH_{3}):
 <math>P = P_{{\mathrm{N}}_2} + P_{{\mathrm{H}}_2} + P_{{\mathrm{NH}}_3}</math>
where:  
<math>P \,</math>  = total pressure of the gas mixture 

<math>P_{{\mathrm{N}}_2}</math>  = partial pressure of nitrogen (N_{2}) 
<math>P_{{\mathrm{H}}_2}</math>  = partial pressure of hydrogen (H_{2}) 
<math>P_{{\mathrm{NH}}_3}</math>  = partial pressure of ammonia (NH_{3}) 
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:
 <math>x_{\mathrm{i}} = \frac{P_{\mathrm{i}}}{P} = \frac{n_{\mathrm{i}}}{n}</math>
and the partial pressure of an individual gas component in an ideal gas can be obtained using this expression:
 <math>P_{\mathrm{i}} = x_{\mathrm{i}} \cdot P</math>
where:  
<math>x_{\mathrm{i}}</math>  = mole fraction of any individual gas component in a gas mixture 

<math>P_{\mathrm{i}}</math>  = partial pressure of any individual gas component in a gas mixture 
<math>n_{\mathrm{i}}</math>  = moles of any individual gas component in a gas mixture 
<math>n</math>  = total moles of the gas mixture 
<math>P</math>  = pressure of the gas mixture 
The mole fraction of a gas component in a gas mixture is equal to the volumetric fraction of that component in a gas mixture.^{[3]}
Equilibrium constants of reactions involving gas mixtures
It is possible to work out the equilibrium constant for a chemical reaction involving a mixture of gases given the partial pressure of each gas and the overall reaction formula. For a reversible reaction involving gas reactants and gas products, such as:
 <math>a\,A + b\,B \leftrightarrow c\,C + d\,D</math>
the equilibrium constant of the reaction would be:
 <math>K_P = \frac{P_C^c\, P_D^d} {P_A^a\, P_B^b}</math>
where:  
<math>K_P</math>  = the equilibrium constant of the reaction 

<math>a</math>  = coefficient of reactant <math>A</math> 
<math>b</math>  = coefficient of reactant <math>B</math> 
<math>c</math>  = coefficient of product <math>C</math> 
<math>d</math>  = coefficient of product <math>D</math> 
<math>P_C^c</math>  = the partial pressure of <math>C</math> raised to the power of <math>c</math> 
<math>P_D^d</math>  = the partial pressure of <math>D</math> raised to the power of <math>d</math> 
<math>P_A^a</math>  = the partial pressure of <math>A</math> raised to the power of <math>a</math> 
<math>P_B^b</math>  = the partial pressure of <math>B</math> raised to the power of <math>b</math> 
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.
Henry's Law and the solubility of gases
Gases will dissolve in liquids to an extent that is determined by the equilibrium between the undissolved gas and the gas that has dissolved in the liquid (called the solvent).^{[4]} The equilibrium constant for that equilibrium is:
 (1) <math>k = \frac {P_X}{C_X}</math>
where: <math>k</math> = the equilibrium constant for the solvation process <math>P_X</math> = partial pressure of gas <math>X</math> in equilibrium with a solution containing some of the gas <math>C_X</math> = the concentration of gas <math>X</math> in the liquid solution
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 <math>k</math> is quite often referred to as the Henry's Law constant.^{[4]}^{[5]}^{[6]}
Henry's Law is sometimes written as:^{[7]}
 (2) <math>k' = \frac {C_X}{P_X}</math>
where <math>k'</math> is also referred to as the Henry's Law constant.^{[7]} As can be seen by comparing equations (1) and (2) above, <math>k'</math> is the reciprocal of <math>k</math>. 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.
Partial pressure in diving breathing gases
In recreational diving and professional diving the richness of individual component gases of breathing gases is expressed by partial pressure.
Using diving terms, partial pressure is calculated as:
 partial pressure = total absolute pressure x volume fraction of gas component
For the component gas "i":
 ppi = P x Fi
For example, at 50 metres (165 feet), the total absolute pressure is 6 bar (600 kPa) (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 79% by volume are:
 ppN2 = 6 bar x 0.79 = 4.7 bar absolute
 ppO2 = 6 bar x 0.21 = 1.3 bar absolute
where:  
ppi  = partial pressure of gas component i = <math>P_{\mathrm{i}}</math> in the terms used in this article 

P  = total pressure = <math>P</math> in the terms used in this article 
Fi  = volume fraction of gas component i = mole fraction, <math>x_{\mathrm{i}}</math>, in the terms used in this article 
ppN2  = partial pressure of nitrogen = <math>P_{{\mathrm{N}}_2}</math> in the terms used in this article 
ppO2  = partial pressure of oxygen = <math>P_{{\mathrm{O}}_2}</math> in the terms used in this article 
The minimum safe lower limit for the partial pressures of oxygen in a gas mixture is 0.16 bar (16 kPa) absolute. Hypoxia and sudden unconsciousness becomes a problem with an oxygen partial pressure of less than 0.16 bar absolute. The NOAA Diving Manual recommends a maximum single exposure of 45 minutes at 1.6 bar absolute, of 120 minutes at 1.5 bar absolute, of 150 minutes at 1.4 bar absolute, of 180 minutes at 1.3 bar absolute and of 210 minutes at 1.2 bar absolute. Oxygen toxicity, involving convulsions, becomes a risk when these oxygen partial pressures and exposures are exceeded. The partial pressure of oxygen determines the maximum operating depth of a gas mixture.
Nitrogen narcosis is a problem with gas mixes containing nitrogen. A typical planned maximum partial pressure of nitrogen for technical diving is 3.5 bar absolute, based on an equivalent air depth of 35 metres (115 feet).
Related Chapters
 Vapor and Vapor pressure
 Gas, Ideal gas and Ideal gas law
 Mole fraction and Mole (unit)
 Dalton's law
 Henry's law
 Breathing gas
References
 ↑ University of Illinois class notes on thermodynamics
 ↑ Chemistry notes at Ohio Sate University
 ↑ Pittsburgh University chemical engineering class notes
 ↑ ^{4.0} ^{4.1} Intute University Introductory Chemistry
 ↑ University of Delware physical chemistry lecture
 ↑ Rice University chemistry class notes
 ↑ ^{7.0} ^{7.1} University of Arizona chemistry class notes
da:Partialtryk de:Partialdruck eo:Parta premo it:Pressione totale e parziale in un miscuglio di gas he:לחץ חלקי nl:Partiële druk sl:Delni tlak sv:Partialtryck uk:Парціальний тиск