Partial pressure

In a mixture of

The partial pressure of a gas is a measure of thermodynamic activity of the gas's

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.[4]^[5]

Examples:

• ${\displaystyle P_{1}}$ or ${\displaystyle p_{1}}$ = pressure at time 1
• ${\displaystyle P_{{\ce {H2}}}}$ or ${\displaystyle p_{{\ce {H2}}}}$ = partial pressure of hydrogen
• ${\displaystyle P_{a_{{\ce {O2}}}}}$ or ${\displaystyle p_{a_{{\ce {O2}}}}}$ or P_aO₂ = arterial partial pressure of oxygen
• ${\displaystyle P_{v_{{\ce {O2}}}}}$ or ${\displaystyle p_{v_{{\ce {O2}}}}}$ or P_vO₂ = venous partial pressure of oxygen

Dalton's law of partial pressures

Dalton's law expresses the fact 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.^[6] 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 (N₂), hydrogen (H₂) and ammonia (NH₃):

$${\displaystyle p=p_{{\ce {N2}}}+p_{{\ce {H2}}}+p_{{\ce {NH3}}}}$$

• ${\displaystyle p}$ = total pressure of the gas mixture
• ${\displaystyle p_{{\ce {N2}}}}$ = partial pressure of nitrogen (N₂)
• ${\displaystyle p_{{\ce {H2}}}}$ = partial pressure of hydrogen (H₂)
• ${\displaystyle p_{{\ce {NH3}}}}$ = partial pressure of ammonia (NH₃)

Ideal gas mixtures

Ideally the ratio of partial pressures equals the ratio of the number of molecules. That is, the mole fraction ${\displaystyle x_{\mathrm {i} }}$ 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:

$${\displaystyle x_{\mathrm {i} }={\frac {p_{\mathrm {i} }}{p}}={\frac {n_{\mathrm {i} }}{n}}}$$

and the partial pressure of an individual gas component in an ideal gas can be obtained using this expression:

$${\displaystyle p_{\mathrm {i} }=x_{\mathrm {i} }\cdot p}$$

where:
${\displaystyle x_{\mathrm {i} }}$	= mole fraction of any individual gas component in a gas mixture
${\displaystyle p_{\mathrm {i} }}$	= partial pressure of any individual gas component in a gas mixture
${\displaystyle n_{\mathrm {i} }}$	= moles of any individual gas component in a gas mixture
${\displaystyle n}$	= total moles of the gas mixture
${\displaystyle p}$	= 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.^[7]

The ratio of partial pressures relies on the following isotherm relation:

$${\displaystyle {\frac {V_{\rm {X}}}{V_{\rm {tot}}}}={\frac {p_{\rm {X}}}{p_{\rm {tot}}}}={\frac {n_{\rm {X}}}{n_{\rm {tot}}}}}$$

• V_X is the partial volume of any individual gas component (X)
• V_tot is the total volume of the gas mixture
• p_X is the partial pressure of gas X
• p_tot is the total pressure of the gas mixture
• n_X is the amount of substance of gas (X)
• n_tot is the total amount of substance in 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:^[8]

$${\displaystyle V_{\rm {X}}=V_{\rm {tot}}\times {\frac {p_{\rm {X}}}{p_{\rm {tot}}}}=V_{\rm {tot}}\times {\frac {n_{\rm {X}}}{n_{\rm {tot}}}}}$$

• V_X is the partial volume of an individual gas component X in the mixture
• V_tot is the total volume of the gas mixture
• p_X is the partial pressure of gas X
• p_tot is the total pressure of the gas mixture
• n_X is the amount of substance of gas X
• n_tot is the total amount of substance in the gas mixture

Vapor pressure

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.^[9] 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,

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:

$${\displaystyle {\ce {{{\mathit {a}}A}+{{\mathit {b}}B}<=>{{\mathit {c}}C}+{{\mathit {d}}D}}}}$$

the equilibrium constant of the reaction would be:

$${\displaystyle K_{\mathrm {p} }={\frac {p_{C}^{c}\,p_{D}^{d}}{p_{A}^{a}\,p_{B}^{b}}}}$$

where:
${\displaystyle K_{p}}$	=  the equilibrium constant of the reaction
${\displaystyle a}$	=  coefficient of reactant ${\displaystyle A}$
${\displaystyle b}$	=  coefficient of reactant ${\displaystyle B}$
${\displaystyle c}$	=  coefficient of product ${\displaystyle C}$
${\displaystyle d}$	=  coefficient of product ${\displaystyle D}$
${\displaystyle p_{C}^{c}}$	=  the partial pressure of ${\displaystyle C}$ raised to the power of ${\displaystyle c}$
${\displaystyle p_{D}^{d}}$	=  the partial pressure of ${\displaystyle D}$ raised to the power of ${\displaystyle d}$
${\displaystyle p_{A}^{a}}$	=  the partial pressure of ${\displaystyle A}$ raised to the power of ${\displaystyle a}$
${\displaystyle p_{B}^{b}}$	=  the partial pressure of ${\displaystyle B}$ raised to the power of ${\displaystyle b}$

For reversible reactions, changes in the total pressure, temperature or reactant concentrations will shift the

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).^[10] The equilibrium constant for that equilibrium is:

$${\displaystyle k={\frac {p_{x}}{C_{x}}}}$$

(1)

where:

• ${\displaystyle k}$ =  the equilibrium constant for the solvation process
• ${\displaystyle p_{x}}$ =  partial pressure of gas ${\displaystyle x}$ in equilibrium with a solution containing some of the gas
• ${\displaystyle C_{x}}$ =  the concentration of gas ${\displaystyle x}$ in the liquid solution

The form of the equilibrium constant shows that the concentration of a

and the equilibrium constant ${\displaystyle k}$ is quite often referred to as the Henry's law constant.

Henry's law is sometimes written as:^[13]

$${\displaystyle k'={\frac {C_{x}}{p_{x}}}}$$

(2)

where ${\displaystyle k'}$ is also referred to as the Henry's law constant.^[13] As can be seen by comparing equations (1) and (2) above, ${\displaystyle k'}$ is the reciprocal of ${\displaystyle k}$. 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 underwater diving the physiological effects of individual component gases of breathing gases are a function of partial pressure.^[14]

Using diving terms, partial pressure is calculated as:

partial pressure = (total absolute pressure) × (volume fraction of gas component)^[14]

For the component gas "i":

p_i = P × F_i^[14]

For example, at 50 metres (164 ft) underwater, the total absolute pressure is 6 bar (600 kPa) (i.e., 1 bar of

approximately 79% by volume are:

pN₂ = 6 bar × 0.79 = 4.7 bar absolute

pO₂ = 6 bar × 0.21 = 1.3 bar absolute

where:
pi	= partial pressure of gas component i  = ${\displaystyle P_{\mathrm {i} }}$ in the terms used in this article
P	= total pressure = ${\displaystyle P}$ in the terms used in this article
Fi	= volume fraction of gas component i  =  mole fraction, ${\displaystyle x_{\mathrm {i} }}$, in the terms used in this article
pN2	= partial pressure of nitrogen  = ${\displaystyle P_{\mathrm {N_{2}} }}$ in the terms used in this article
pO2	= partial pressure of oxygen  = ${\displaystyle P_{\mathrm {O_{2}} }}$ in the terms used in this article

The minimum safe lower limit for the partial pressures of oxygen in a breathing gas mixture for diving is 0.16 bars (16 kPa) 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 35 metres (115 ft).

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.^[14]

In medicine

The partial pressures of particularly oxygen (${\displaystyle p_{\mathrm {O_{2}} }}$) and carbon dioxide (${\displaystyle p_{\mathrm {CO_{2}} }}$) are important parameters in tests of

]

Reference ranges for ${\displaystyle p_{\mathrm {O_{2}} }}$ and ${\displaystyle p_{\mathrm {CO_{2}} }}$

	Unit	Arterial blood gas	Venous blood gas	Cerebrospinal fluid	Alveolar pulmonary gas pressures
${\displaystyle p_{\mathrm {O_{2}} }}$	kPa	11–13[16]	4.0–5.3[16]	5.3–5.9[16]	14.2
	mmHg	75–100[17]	30–40[18]	40–44[19]	107
${\displaystyle p_{\mathrm {CO_{2}} }}$	kPa	4.7–6.0[16]	5.5–6.8[16]	5.9–6.7[16]	4.8
	mmHg	35–45[17]	41–51[18]	44–50[19]	36

See also

• Blood gas tension – Partial pressure of blood gases
• Breathing gas – Gas used for human respiration
• Henry's law – Gas law regarding proportionality of dissolved gas
• Ideal gas – Mathematical model which approximates the behavior of real gases
  • Ideal gas law – Equation of the state of a hypothetical ideal gas
• Mole fraction – Proportion of a constituent in a mixture
  • Mole (unit) – SI unit of amount of substance
• Vapor – Substances in the gas phase at a temperature lower than its critical point

References