Isotherms of an
ideal gas. The curved lines represent the relationship between
pressure (on the vertical,
yaxis) and
volume (on the horizontal,
xaxis) for an ideal gas at different
temperatures: lines which are further away from the
origin (that is, lines that are nearer to the top righthand corner of the diagram) represent higher temperatures.
The ideal gas law is the equation of state of a hypothetical ideal gas. It is a good approximation to the behaviour of many gases under many conditions, although it has several limitations. It was first stated by Émile Clapeyron in 1834 as a combination of Boyle's law and Charles's law.^{[1]} The ideal gas law is often introduced in its common form:

PV=nRT\,
where P is the pressure of the gas, V is the volume of the gas, n is the amount of substance of gas (measured in moles), R is the ideal, or universal, gas constant, and T is the temperature of the gas.
It can also be derived microscopically from kinetic theory, as was achieved (apparently independently) by August Krönig in 1856^{[2]} and Rudolf Clausius in 1857.^{[3]}
Equation
The state of an amount of gas is determined by its pressure, volume, and temperature. The modern form of the equation relates these simply in two main forms. The temperature used in the equation of state is an absolute temperature: in the SI system of units, Kelvin.^{[4]}
Common form
The most frequently introduced form is

PV=nRT\,
where:
P is the pressure of the gas
V is the volume of the gas
n is the amount of substance of gas (also known as number of moles)
R is the ideal, or universal, gas constant, equal to the product of the Boltzmann constant and the Avogadro constant.
T is the temperature of the gas
In SI units, P is measured in pascals, V is measured in cubic metres, n is measured in moles, and T in kelvin (273.15 kelvin = 0.00 degrees Celsius). R has the value 8.314 J·K^{−1}·mol^{−1} or 0.08206 L·atm·mol^{−1}·K^{−1}or ≈2 calories if using pressure in standard atmospheres (atm) instead of pascals, and volume in liters instead of cubic metres.
Molar form
How much gas is present could be specified by giving the mass instead of the chemical amount of gas. Therefore, an alternative form of the ideal gas law may be useful. The chemical amount (n) (in moles) is equal to the mass (m) (in grams) divided by the molar mass (M) (in grams per mole):

n = {\frac{m}{M}}
By replacing n with m / M, and subsequently introducing density ρ = m/V, we get:

\ PV = \frac{m}{M}RT

\ P = \rho \frac{R}{M}T
Defining the specific gas constant R_{specific} as the ratio R/M,

\ P = \rho R_{\rm specific}T
This form of the ideal gas law is very useful because it links pressure, density, and temperature in a unique formula independent of the quantity of the considered gas. Alternatively, the law may be written in terms of the specific volume v, the reciprocal of density, as

\ Pv = R_{\rm specific}T
It is common, especially in engineering applications, to represent the specific gas constant by the symbol R. In such cases, the universal gas constant is usually given a different symbol such as R to distinguish it. In any case, the context and/or units of the gas constant should make it clear as to whether the universal or specific gas constant is being referred to.^{[5]}
Statistical mechanics
In statistical mechanics the following molecular equation is derived from first principles:

\ P = nk_BT
where P is the absolute pressure of the gas measured in Pascals; n is the number density in the gas measured in 1/(meters cubed); k_{B} is the Boltzmann constant relating temperature and energy; and T is the absolute temperature.
The number density contrasts to the other formulation, which uses N, the number of moles and V, the volume. This relation implies that R=N_{A}k_{B} where N_{A} is Avogadro's constant, and the consistency of this result with experiment is a good check on the principles of statistical mechanics.
From this we can notice that for an average particle mass of μ times the atomic mass constant m_{u} (i.e., the mass is μ u)

N = \frac{m}{\mu m_\mathrm{u}}
and since ρ = mn, we find that the ideal gas law can be rewritten as:

P = \frac{1}{V}\frac{m}{\mu m_\mathrm{u}} kT = \frac{k}{\mu m_\mathrm{u}} \rho T .
In SI units, P is measured in pascals; V in cubic metres; N is a dimensionless number; and T in Kelvin. k has the value 1.38·10^{−23} J·K^{−1} in SI units.
Applications to thermodynamic processes
The table below essentially simplifies the ideal gas equation for a particular processes, thus making this equation easier to solve using numerical methods.
A thermodynamic process is defined as a system that moves from state 1 to state 2, where the state number is denoted by subscript. As shown in the first column of the table, basic thermodynamic processes are defined such that one of the gas properties (P, V, T, or S) is constant throughout the process.
For a given thermodynamics process, in order to specify the extent of a particular process, one of the properties ratios (listed under the column labeled "known ratio") must be specified (either directly or indirectly). Also, the property for which the ratio is known must be distinct from the property held constant in the previous column (otherwise the ratio would be unity, and not enough information would be available to simplify the gas law equation).
In the final three columns, the properties (P, V, or T) at state 2 can be calculated from the properties at state 1 using the equations listed.
Process

Constant

Known ratio

P_{2}

V_{2}

T_{2}

Isobaric process

Pressure

V_{2}/V_{1}

P_{2} = P_{1}

V_{2} = V_{1}(V_{2}/V_{1})

T_{2} = T_{1}(V_{2}/V_{1})

T_{2}/T_{1}

P_{2} = P_{1}

V_{2} = V_{1}(T_{2}/T_{1})

T_{2} = T_{1}(T_{2}/T_{1})

Isochoric process
(Isovolumetric process)
(Isometric process)

Volume

P_{2}/P_{1}

P_{2} = P_{1}(P_{2}/P_{1})

V_{2} = V_{1}

T_{2} = T_{1}(P_{2}/P_{1})

T_{2}/T_{1}

P_{2} = P_{1}(T_{2}/T_{1})

V_{2} = V_{1}

T_{2} = T_{1}(T_{2}/T_{1})

Isothermal process

Temperature

P_{2}/P_{1}

P_{2} = P_{1}(P_{2}/P_{1})

V_{2} = V_{1}/(P_{2}/P_{1})

T_{2} = T_{1}

V_{2}/V_{1}

P_{2} = P_{1}/(V_{2}/V_{1})

V_{2} = V_{1}(V_{2}/V_{1})

T_{2} = T_{1}

Isentropic process
(Reversible adiabatic process)

Entropy^{[a]}

P_{2}/P_{1}

P_{2} = P_{1}(P_{2}/P_{1})

V_{2} = V_{1}(P_{2}/P_{1})^{(−1/γ)}

T_{2} = T_{1}(P_{2}/P_{1})^{(γ − 1)/γ}

V_{2}/V_{1}

P_{2} = P_{1}(V_{2}/V_{1})^{−γ}

V_{2} = V_{1}(V_{2}/V_{1})

T_{2} = T_{1}(V_{2}/V_{1})^{(1 − γ)}

T_{2}/T_{1}

P_{2} = P_{1}(T_{2}/T_{1})^{γ/(γ − 1)}

V_{2} = V_{1}(T_{2}/T_{1})^{1/(1 − γ)}

T_{2} = T_{1}(T_{2}/T_{1})

Polytropic process

P V^{n}

P_{2}/P_{1}

P_{2} = P_{1}(P_{2}/P_{1})

V_{2} = V_{1}(P_{2}/P_{1})^{(1/n)}

T_{2} = T_{1}(P_{2}/P_{1})^{(n  1)/n}

V_{2}/V_{1}

P_{2} = P_{1}(V_{2}/V_{1})^{−n}

V_{2} = V_{1}(V_{2}/V_{1})

T_{2} = T_{1}(V_{2}/V_{1})^{(1−n)}

T_{2}/T_{1}

P_{2} = P_{1}(T_{2}/T_{1})^{n/(n − 1)}

V_{2} = V_{1}(T_{2}/T_{1})^{1/(1 − n)}

T_{2} = T_{1}(T_{2}/T_{1})

^{^} a. In an isentropic process, system entropy (S) is constant. Under these conditions, P_{1} V_{1}^{γ} = P_{2} V_{2}^{γ}, where γ is defined as the heat capacity ratio, which is constant for an ideal gas. The value used for γ is typically 1.4 for diatomic gases like nitrogen (N_{2}) and oxygen (O_{2}), (and air, which is 99% diatomic). Also γ is typically 1.6 for monatomic gases like the noble gases helium (He), and argon (Ar). In internal combustion engines γ varies between 1.35 and 1.15, depending on constitution gases and temperature.
Deviations from ideal behavior of real gases
The equation of state given here applies only to an ideal gas, or as an approximation to a real gas that behaves sufficiently like an ideal gas. There are in fact many different forms of the equation of state. Since the ideal gas law neglects both molecular size and intermolecular attractions, it is most accurate for monatomic gases at high temperatures and low pressures. The neglect of molecular size becomes less important for lower densities, i.e. for larger volumes at lower pressures, because the average distance between adjacent molecules becomes much larger than the molecular size. The relative importance of intermolecular attractions diminishes with increasing thermal kinetic energy, i.e., with increasing temperatures. More detailed equations of state, such as the van der Waals equation, account for deviations from ideality caused by molecular size and intermolecular forces.
A residual property is defined as the difference between a real gas property and an ideal gas property, both considered at the same pressure, temperature, and composition.
Derivations
Empirical
The ideal gas law can be derived from combining two empirical gas laws: the combined gas law and Avogadro's law. The combined gas law states that

\frac{PV}{T}= C
where C is a constant which is directly proportional to the amount of gas, n (Avogadro's law). The proportionality factor is the universal gas constant, R, i.e. C = nR.
Hence the ideal gas law

PV = nRT \,
Theoretical
Kinetic theory
The ideal gas law can also be derived from first principles using the kinetic theory of gases, in which several simplifying assumptions are made, chief among which are that the molecules, or atoms, of the gas are point masses, possessing mass but no significant volume, and undergo only elastic collisions with each other and the sides of the container in which both linear momentum and kinetic energy are conserved.
Statistical mechanics
Let q = (q_{x}, q_{y}, q_{z}) and p = (p_{x}, p_{y}, p_{z}) denote the position vector and momentum vector of a particle of an ideal gas, respectively. Let F denote the net force on that particle. Then the timeaveraged potential energy of the particle is:

\begin{align} \langle \mathbf{q} \cdot \mathbf{F} \rangle &= \Bigl\langle q_{x} \frac{dp_{x}}{dt} \Bigr\rangle + \Bigl\langle q_{y} \frac{dp_{y}}{dt} \Bigr\rangle + \Bigl\langle q_{z} \frac{dp_{z}}{dt} \Bigr\rangle\\ &=\Bigl\langle q_{x} \frac{\partial H}{\partial q_x} \Bigr\rangle  \Bigl\langle q_{y} \frac{\partial H}{\partial q_y} \Bigr\rangle  \Bigl\langle q_{z} \frac{\partial H}{\partial q_z} \Bigr\rangle = 3k_{B} T, \end{align}
where the first equality is Newton's second law, and the second line uses Hamilton's equations and the equipartition theorem. Summing over a system of N particles yields

3Nk_{B} T =  \biggl\langle \sum_{k=1}^{N} \mathbf{q}_{k} \cdot \mathbf{F}_{k} \biggr\rangle.
By Newton's third law and the ideal gas assumption, the net force of the system is the force applied by the walls of the container, and this force is given by the pressure P of the gas. Hence

\biggl\langle\sum_{k=1}^{N} \mathbf{q}_{k} \cdot \mathbf{F}_{k}\biggr\rangle = P \oint_{\mathrm{surface}} \mathbf{q} \cdot d\mathbf{S},
where dS is the infinitesimal area element along the walls of the container. Since the divergence of the position vector q is

\nabla \cdot \mathbf{q} = \frac{\partial q_{x}}{\partial q_{x}} + \frac{\partial q_{y}}{\partial q_{y}} + \frac{\partial q_{z}}{\partial q_{z}} = 3,
the divergence theorem implies that

P \oint_{\mathrm{surface}} \mathbf{q} \cdot d\mathbf{S} = P \int_{\mathrm{volume}} \left( \nabla \cdot \mathbf{q} \right) dV = 3PV,
where dV is an infinitesimal volume within the container and V is the total volume of the container.
Putting these equalities together yields

3Nk_{B} T = \biggl\langle \sum_{k=1}^{N} \mathbf{q}_{k} \cdot \mathbf{F}_{k} \biggr\rangle = 3PV,
which immediately implies the ideal gas law for N particles:

PV = Nk_{B} T = nRT,\,
where n = N/N_{A} is the number of moles of gas and R = N_{A}k_{B} is the gas constant.
See also
References
Further reading

Davis and Masten Principles of Environmental Engineering and Science, McGrawHill Companies, Inc. New York (2002) ISBN 0072350539

Website giving credit to Benoît Paul Émile Clapeyron, (1799–1864) in 1834
External links

Configuration integral (statistical mechanics) where an alternative statistical mechanics derivation of the idealgas law, using the relationship between the Helmholtz free energy and the partition function, but without using the equipartition theorem, is provided.
Diving medicine, physiology, physics and environment


Diving medicine:


Pressure

Oxygen



Inert gases



Carbon dioxide




Immersion




Treatments





Diving physiology



Diving physics



Diving environment



Researchers in
diving medicine,
physiology and physics



Diving medical
research
organisations





This article was sourced from Creative Commons AttributionShareAlike License; additional terms may apply. World Heritage Encyclopedia content is assembled from numerous content providers, Open Access Publishing, and in compliance with The Fair Access to Science and Technology Research Act (FASTR), Wikimedia Foundation, Inc., Public Library of Science, The Encyclopedia of Life, Open Book Publishers (OBP), PubMed, U.S. National Library of Medicine, National Center for Biotechnology Information, U.S. National Library of Medicine, National Institutes of Health (NIH), U.S. Department of Health & Human Services, and USA.gov, which sources content from all federal, state, local, tribal, and territorial government publication portals (.gov, .mil, .edu). Funding for USA.gov and content contributors is made possible from the U.S. Congress, EGovernment Act of 2002.
Crowd sourced content that is contributed to World Heritage Encyclopedia is peer reviewed and edited by our editorial staff to ensure quality scholarly research articles.
By using this site, you agree to the Terms of Use and Privacy Policy. World Heritage Encyclopedia™ is a registered trademark of the World Public Library Association, a nonprofit organization.