Partial pressure - wikipedia Partial pressure In a mixture of gases, each gas has a partial pressure which is the hypothetical pressure of that gas 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. It relies on the following isotherm relation : V x V t o t = p x p t o t = n x n t o t ( displaystyle ( \frac ( V_ ( x ) ) ( V_ ( tot ) ) ) = ( \frac ( p_ ( x ) ) ( p_ ( tot ) ) ) = ( \frac ( n_ ( x ) ) ( n_ ( tot ) ) ) ) V is the partial volume of any individual gas component ( X ) V is the total volume of the gas mixture p is the partial pressure of gas X p is the total pressure of the gas mixture n is the amount of substance of gas ( X ) n is the total amount of substance in gas mixture The partial pressure of a gas is a measure of thermodynamic activity of the gas 's molecules. Gases dissolve, diffuse, and react according to their partial pressures, and not according to their concentrations in gas mixtures or liquids. This general property of gases is also true in chemical reactions of gases in biology. For example, the necessary amount of oxygen for human respiration, and the amount that is toxic, is set by the partial pressure of oxygen alone. This is true across a very wide range of different concentrations of oxygen present in various inhaled breathing gases or dissolved in blood. Contents 1 Symbol 2 Dalton 's law of partial pressures 3 Ideal gas mixtures 4 Partial volume ( Amagat 's law of additive volume ) 5 Vapor pressure 6 Equilibrium constants of reactions involving gas mixtures 7 Henry 's law and the solubility of gases 8 In diving breathing gases 9 In medicine 10 See also 11 References Symbol The symbol for pressure is usually P or p which may use a subscript to identify the pressure, and gas species are also referred to by subscript. When combined these subscripts are applied recursively. Examples : P 1 ( displaystyle ( P_ ( 1 ) ) ) or p 1 ( displaystyle ( p_ ( 1 ) ) ) = pressure at time 1 P H 2 ( displaystyle ( P_ ( H_ ( 2 ) ) ) ) or p H 2 ( displaystyle ( p_ ( H_ ( 2 ) ) ) ) = partial pressure of hydrogen P v O 2 ( displaystyle ( P_ ( v_ ( O_ ( 2 ) ) ) ) ) or p v O 2 ( displaystyle ( p_ ( v_ ( O_ ( 2 ) ) ) ) ) = venous partial pressure of oxygen Dalton 's law of partial pressures Main article : Dalton 's law Dalton 's law expresses the fact that the total pressure of a mixture of gases is equal to the sum of the partial pressures of the individual gases in the mixture. This equality arises from the fact that in an ideal gas the molecules are so far apart that they do not interact with each other. Most actual real - world gases come very close to this ideal. For example, given an ideal gas mixture of nitrogen ( N ), hydrogen ( H ) and ammonia ( NH ) : p = p N 2 + p H 2 + p NH 3 ( displaystyle p = p_ ( ce ( N2 ) ) + p_ ( ce ( H2 ) ) + p_ ( ce ( NH3 ) ) ) where : p ( displaystyle p , ) = total pressure of the gas mixture p N 2 ( displaystyle p_ ( ce ( N2 ) ) ) = partial pressure of nitrogen ( N ) p H 2 ( displaystyle p_ ( ce ( H2 ) ) ) = partial pressure of hydrogen ( H ) p NH 3 ( displaystyle p_ ( ce ( NH3 ) ) ) = partial pressure of ammonia ( NH ) Ideal gas mixtures Ideally the ratio of partial pressures equals the ratio of the number of molecules. That is, 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 : x i = p i p = n i n ( displaystyle x_ ( mathrm ( i ) ) = ( \frac ( p_ ( mathrm ( i ) ) ) ( p ) ) = ( \frac ( n_ ( mathrm ( i ) ) ) ( n ) ) ) and the partial pressure of an individual gas component in an ideal gas can be obtained using this expression : p i = x i ⋅ p ( displaystyle p_ ( mathrm ( i ) ) = x_ ( mathrm ( i ) ) cdot p ) where : x i ( displaystyle x_ ( mathrm ( i ) ) ) = mole fraction of any individual gas component in a gas mixture p i ( displaystyle p_ ( mathrm ( i ) ) ) = partial pressure of any individual gas component in a gas mixture n i ( displaystyle n_ ( mathrm ( i ) ) ) = moles of any individual gas component in a gas mixture n ( displaystyle n ) = total moles of the gas mixture p ( displaystyle p ) = total 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. Partial volume ( amagat 's law of additive volume ) The partial volume of a particular gas in a mixture is the volume of one component of the gas mixture. It is useful in gas mixtures, e.g. air, to focus on one particular gas component, e.g. oxygen. It can be approximated both from partial pressure and molar fraction : V x = V t o t × p x p t o t = V t o t × n x n t o t ( displaystyle V_ ( x ) = V_ ( tot ) \times ( \frac ( p_ ( x ) ) ( p_ ( tot ) ) ) = V_ ( tot ) \times ( \frac ( n_ ( x ) ) ( n_ ( tot ) ) ) ) V is the partial volume of an individual gas component X in the mixture V is the total volume of the gas mixture p is the partial pressure of gas X p is the total pressure of the gas mixture n is the amount of substance of gas X n is the total amount of substance in the gas mixture Vapor pressure Main article : Vapor pressure A log - lin vapor pressure chart for various liquids Vapor pressure is the pressure of a vapor in equilibrium with its non-vapor 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 atmospheric pressure boiling point corresponds to the temperature at which its vapor pressure is equal to the surrounding atmospheric pressure and it is often called the normal boiling point. 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. As can be seen in the chart, the liquids with the highest vapor pressures have the lowest normal boiling points. 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. Note that 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 ). 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 : a A + b B ↔ c C + d D ( displaystyle a , A + b , B leftrightarrow c , C + d , D ) the equilibrium constant of the reaction would be : K p = p C c p D d p A a p B b ( displaystyle K_ ( p ) = ( \frac ( p_ ( C ) ^ ( c ) , p_ ( D ) ^ ( d ) ) ( p_ ( A ) ^ ( a ) , p_ ( B ) ^ ( b ) ) ) ) where : K p ( displaystyle K_ ( p ) ) = the equilibrium constant of the reaction a ( displaystyle a ) = coefficient of reactant A ( displaystyle A ) b ( displaystyle b ) = coefficient of reactant B ( displaystyle B ) c ( displaystyle c ) = coefficient of product C ( displaystyle C ) d ( displaystyle d ) = coefficient of product D ( displaystyle D ) p C c ( displaystyle p_ ( C ) ^ ( c ) ) = the partial pressure of C ( displaystyle C ) raised to the power of c ( displaystyle c ) p D d ( displaystyle p_ ( D ) ^ ( d ) ) = the partial pressure of D ( displaystyle D ) raised to the power of d ( displaystyle d ) p A a ( displaystyle p_ ( A ) ^ ( a ) ) = the partial pressure of A ( displaystyle A ) raised to the power of a ( displaystyle a ) p B b ( displaystyle p_ ( B ) ^ ( b ) ) = the partial pressure of B ( displaystyle B ) raised to the power of b ( displaystyle b ) 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 Main article : Henry 's law 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 ). The equilibrium constant for that equilibrium is : ( 1 ) k = p x C x ( displaystyle k = ( \frac ( p_ ( x ) ) ( C_ ( x ) ) ) ) where : k ( displaystyle k ) = the equilibrium constant for the solvation process p x ( displaystyle p_ ( x ) ) = partial pressure of gas x ( displaystyle x ) in equilibrium with a solution containing some of the gas C x ( displaystyle C_ ( x ) ) = the concentration of gas x ( displaystyle x ) 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 k ( displaystyle k ) is quite often referred to as the Henry 's law constant. Henry 's law is sometimes written as : ( 2 ) k ′ = C x p x ( displaystyle k ' = ( \frac ( C_ ( x ) ) ( p_ ( x ) ) ) ) where k ′ ( displaystyle k ' ) is also referred to as the Henry 's law constant. As can be seen by comparing equations ( 1 ) and ( 2 ) above, k ′ ( displaystyle k ' ) is the reciprocal of k ( displaystyle k ). 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. 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 ) × ( volume fraction of gas component ) For the component gas `` i '' : pp = P × F For example, at 50 metres ( 164 ft ), 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 : ppN = 6 bar × 0.79 = 4.7 bar absolute ppO = 6 bar × 0.21 = 1.3 bar absolute where : pp = partial pressure of gas component i = P i ( displaystyle P_ ( mathrm ( i ) ) ) in the terms used in this article = total pressure = P ( displaystyle P ) in the terms used in this article = volume fraction of gas component i = mole fraction, x i ( displaystyle x_ ( mathrm ( i ) ) ), in the terms used in this article ppN = partial pressure of nitrogen = P N 2 ( displaystyle P_ ( ( mathrm ( N ) ) _ ( 2 ) ) ) in the terms used in this article ppO = partial pressure of oxygen = P O 2 ( displaystyle P_ ( ( mathrm ( O ) ) _ ( 2 ) ) ) 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 bars ( 16 kPa ) absolute. Hypoxia and sudden unconsciousness becomes a problem with an oxygen partial pressure of less than 0.16 bar absolute. Oxygen toxicity, involving convulsions, becomes a problem when oxygen partial pressure is too high. 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 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 when breathing gases at high pressure. Typically, the maximum total partial pressure of narcotic gases used when planning for technical diving is 4.5 bar absolute, based on an equivalent narcotic depth of 35 metres ( 115 ft ). In medicine The partial pressures of particularly oxygen ( p O 2 ( displaystyle p_ ( ( mathrm ( O ) ) _ ( 2 ) ) ) ) and carbon dioxide ( p C O 2 ( displaystyle p_ ( ( mathrm ( CO ) ) _ ( 2 ) ) ) ) are important parameters in tests of arterial blood gases, but can also be measured in, for example, cerebrospinal fluid. Reference ranges for p O 2 ( displaystyle p_ ( ( mathrm ( O ) ) _ ( 2 ) ) ) and p C O 2 ( displaystyle p_ ( ( mathrm ( CO ) ) _ ( 2 ) ) ) Unit Arterial blood gas Venous blood gas Cerebrospinal fluid Alveolar pulmonary gas pressures p O 2 ( displaystyle p_ ( ( mathrm ( O ) ) _ ( 2 ) ) ) kPa 11 -- 13 4.0 -- 5.3 5.3 -- 5.9 14.2 mmHg 75 -- 100 30 -- 40 40 -- 44 107 p C O 2 ( displaystyle p_ ( ( mathrm ( CO ) ) _ ( 2 ) ) ) kPa 4.7 -- 6.0 5.5 -- 6.8 5.9 -- 6.7 4.8 mmHg 35 -- 45 41 -- 51 44 -- 50 36 See also Underwater diving portal Breathing gas Henry 's law Ideal gas and Ideal gas law Mole fraction and Mole ( unit ) Vapor References Jump up ^ Charles Henrickson ( 2005 ). Chemistry. Cliffs Notes. ISBN 0 - 7645 - 7419 - 1. Jump up ^ Staff. `` Symbols and Units '' ( PDF ). Respiratory Physiology & Neurobiology : Guide for Authors. Elsevier. p. 1. Retrieved 3 June 2017. All symbols referring to gas species are in subscript, Jump up ^ IUPAC. Compendium of Chemical Terminology, 2nd ed. ( the `` Gold Book '' ). Compiled by A.D. McNaught and A. Wilkinson. Blackwell Scientific Publications, Oxford ( 1997 ). XML on - line corrected version : http://goldbook.iupac.org ( 2006 - ) created by M. Nic, J. Jirat, B. Kosata ; updates compiled by A. Jenkins. ISBN 0 - 9678550 - 9 - 8. https://doi.org/10.1351/goldbook. Last update : 2014 - 02 - 24 ; version : 2.3. 3. DOI of this term : https://doi.org/10.1351/goldbook.P04819. Jump up ^ Dalton 's Law of Partial Pressures Jump up ^ Frostberg State University 's `` General Chemistry Online '' Jump up ^ Page 200 in : Medical biophysics. Flemming Cornelius. 6th Edition, 2008. Jump up ^ Perry, R.H. and Green, D.W. ( Editors ) ( 1997 ). Perry 's Chemical Engineers ' Handbook ( 7th ed. ). McGraw - Hill. ISBN 0 - 07 - 049841 - 5. CS1 maint : Multiple names : authors list ( link ) CS1 maint : Extra text : authors list ( link ) ^ Jump up to : An extensive list of Henry 's law constants, and a conversion tool Jump up ^ Francis L. Smith & Allan H. Harvey ( September 2007 ). `` Avoid Common Pitfalls When Using Henry 's Law ''. CEP ( Chemical Engineering Progress ). ISSN 0360 - 7275. Jump up ^ Introductory University Chemistry, Henry 's Law and the Solubility of Gases ^ Jump up to : University of Arizona chemistry class notes ^ Jump up to : Derived from mmHg values using 0.133322 kPa / mmHg ^ Jump up to : Normal Reference Range Table from The University of Texas Southwestern Medical Center at Dallas. Used in Interactive Case Study Companion to Pathologic basis of disease. ^ Jump up to : The Medical Education Division of the Brookside Associates -- > ABG ( Arterial Blood Gas ) Retrieved on Dec 6, 2009 ^ Jump up to : PATHOLOGY 425 CEREBROSPINAL FLUID ( CSF ) at the Department of Pathology and Laboratory Medicine at the University of British Columbia. By Dr. G.P. Bondy. 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Question: partial pressure is the term used to describe the amount of gas that is
Answer: hypothetical pressure