Ideal gas law

Branches

Equilibrium / Non-equilibrium

State
Equation of state Ideal gas Real gas State of matter Equilibrium Control volume Instruments
Processes
Isobaric Isochoric Isothermal Adiabatic Isentropic Isenthalpic Quasistatic Polytropic Free expansion Reversibility Irreversibility Endoreversibility
Cycles
Heat engines Heat pumps Thermal efficiency

System properties

Note: Conjugate variables in italics

Functions of state
Property diagrams Intensive and extensive properties
Temperature / Entropy (introduction) Pressure / Volume Chemical potential / Particle number Vapor quality Reduced properties
Process functions
Work Heat

Material properties

Property databases

Specific heat capacity

c=

$T$	$\partial S$
$N$	$\partial T$

Compressibility

\beta =-

$1$	$\partial V$
$V$	$\partial p$

Thermal expansion

\alpha =

$1$	$\partial V$
$V$	$\partial T$

Equations

Potentials

Internal energy
$U(S,V)$
Enthalpy
$H(S,p)=U+pV$
Helmholtz free energy
$A(T,V)=U-TS$
Gibbs free energy
$G(T,p)=H-TS$

History
Culture

History
General Heat Entropy Gas laws "Perpetual motion" machines
Philosophy
Entropy and time Entropy and life Brownian ratchet Maxwell's demon Heat death paradox Loschmidt's paradox Synergetics
Theories
Caloric theory Theory of heat Vis viva ("living force") Mechanical equivalent of heat Motive power
Key publications
"An Experimental Enquiry Concerning ... Heat" "On the Equilibrium of Heterogeneous Substances" "Reflections on the Motive Power of Fire"
Timelines
Thermodynamics Heat engines
Art Education
Maxwell's thermodynamic surface Entropy as energy dispersal

Scientists

Book:Thermodynamics

Isotherms of an ideal gas. The curved lines represent the relationship between pressure (on the vertical, y-axis) and volume (on the horizontal, x-axis) for an ideal gas at different temperatures: lines that are farther away from the origin (that is, lines that are nearer to the top right-hand 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 of the behavior of many gases under many conditions, although it has several limitations. It was first stated by Émile Clapeyron in 1834 as a combination of the empirical Boyle's law, Charles' law and Avogadro's Law.^[1] The ideal gas law is often written as

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 (in moles),
$R$ is the ideal, or universal, gas constant, equal to the product of the Boltzmann constant and the Avogadro constant,
$T$ is the absolute 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: the appropriate SI unit is the 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 absolute 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 kelvins (the Kelvin scale is a shifted Celsius scale, where 0.00 K = −273.15 °C, the lowest possible temperature). R has the value 8.314 J/(K·mol) ≈ 2 cal/(K·mol), or 0.08206 L·atm/(mol·K).

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_{\text{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_{\text{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 ${\bar {R}}$ 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:

PV=Nk_{\text{B}}T,

where P is the absolute pressure of the gas, N is the number of molecules in the given volume V (the number density is given by the ratio n = N/V), 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_Ak_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. In extreme conditions the principles of statistical mechanics may break down as some of the assumptions relating a real life example to an ideal gas become untrue.

From this we notice that for a gas with an average particle mass of μ times the atomic mass constant m_u (i.e., the mass is μ u) the number of molecules will be given by

N={\frac {m}{\mu m_{\text{u}}}},

and since ρ = m/V = nμm_u, we find that the ideal gas law can be rewritten as

P={\frac {1}{V}}{\frac {m}{\mu m_{\text{u}}}}k_{B}T={\frac {k_{B}}{\mu m_{\text{u}}}}\rho T.

In SI units, P is measured in pascals, V in cubic metres, Y is a dimensionless number, and T in measured kelvins. k_B has the value 1.38·10⁻²³ J/K 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 (which are 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	P₂	V₂	T₂
Isobaric process	Pressure	V₂/V₁	P₂ = P₁	V₂ = V₁(V₂/V₁)	T₂ = T₁(V₂/V₁)
Isobaric process	Pressure	T₂/T₁	P₂ = P₁	V₂ = V₁(T₂/T₁)	T₂ = T₁(T₂/T₁)
Isochoric process (Isovolumetric process) (Isometric process)	Volume	P₂/P₁	P₂ = P₁(P₂/P₁)	V₂ = V₁	T₂ = T₁(P₂/P₁)
	Volume	T₂/T₁	P₂ = P₁(T₂/T₁)	V₂ = V₁	T₂ = T₁(T₂/T₁)
Isothermal process	Temperature	P₂/P₁	P₂ = P₁(P₂/P₁)	V₂ = V₁/(P₂/P₁)	T₂ = T₁
Isothermal process	Temperature	V₂/V₁	P₂ = P₁/(V₂/V₁)	V₂ = V₁(V₂/V₁)	T₂ = T₁
Isentropic process (Reversible adiabatic process)	Entropy^[a]	P₂/P₁	P₂ = P₁(P₂/P₁)	V₂ = V₁(P₂/P₁)^(−1/γ)	T₂ = T₁(P₂/P₁)^{(γ − 1)/γ}
		V₂/V₁	P₂ = P₁(V₂/V₁)^−γ	V₂ = V₁(V₂/V₁)	T₂ = T₁(V₂/V₁)^{(1 − γ)}
		T₂/T₁	P₂ = P₁(T₂/T₁)^{γ/(γ − 1)}	V₂ = V₁(T₂/T₁)^{1/(1 − γ)}	T₂ = T₁(T₂/T₁)
Polytropic process	P Vⁿ	P₂/P₁	P₂ = P₁(P₂/P₁)	V₂ = V₁(P₂/P₁)^(-1/n)	T₂ = T₁(P₂/P₁)^{(n - 1)/n}
		V₂/V₁	P₂ = P₁(V₂/V₁)⁻ⁿ	V₂ = V₁(V₂/V₁)	T₂ = T₁(V₂/V₁)⁽¹⁻ⁿ⁾
		T₂/T₁	P₂ = P₁(T₂/T₁)^{n/(n − 1)}	V₂ = V₁(T₂/T₁)^{1/(1 − n)}	T₂ = T₁(T₂/T₁)

^{^} a. In an isentropic process, system entropy (S) is constant. Under these conditions, P₁ V₁^γ = P₂ V₂^γ, where γ is defined as the heat capacity ratio, which is constant for a calorifically perfect gas. The value used for γ is typically 1.4 for diatomic gases like nitrogen (N₂) and oxygen (O₂), (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 that 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 is

PV=nRT.

Theoretical

Kinetic theory

Main article: Kinetic theory of gases

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

Main article: 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 time-averaged potential energy of the particle is:

{\begin{aligned}\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{aligned}}

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_Ak_B is the gas constant.

References

↑ Clapeyron, E (1834). "Mémoire sur la puissance motrice de la chaleur". Journal de l'École Polytechnique (in French). XIV: 153–90. Facsimile at the Bibliothèque nationale de France (pp. 153–90).
↑ Krönig, A. (1856). "Grundzüge einer Theorie der Gase". Annalen der Physik und Chemie (in German). 99 (10): 315–22. Bibcode:1856AnP...175..315K. doi:10.1002/andp.18561751008. Facsimile at the Bibliothèque nationale de France (pp. 315–22).
↑ Clausius, R. (1857). "Ueber die Art der Bewegung, welche wir Wärme nennen". Annalen der Physik und Chemie (in German). 176 (3): 353–79. Bibcode:1857AnP...176..353C. doi:10.1002/andp.18571760302. Facsimile at the Bibliothèque nationale de France (pp. 353–79).
↑ "Equation of State".
↑ Moran and Shapiro, Fundamentals of Engineering Thermodynamics, Wiley, 4th Ed, 2000.

External links

Configuration integral (statistical mechanics) where an alternative statistical mechanics derivation of the ideal-gas law, using the relationship between the Helmholtz free energy and the partition function, but without using the equipartition theorem, is provided. Vu-Quoc, L., Configuration integral (statistical mechanics), 2008. this wiki site is down; see this article in the web archive on 2012 April 28.

Mole concepts

Physical quantities	Mass concentration Molar concentration Molality Mass Volume Density Mole fraction Mass fraction Amount of substance Molar mass Atomic mass Particle number Avogadro constant Pressure Kelvin temperature Gas constant Molar volume Specific volume

Laws	Charles's law Boyle's law Gay-Lussac's law Combined gas law Ideal gas law

Statistical mechanics

Statistical ensembles	Microcanonical Canonical Grand canonical Isothermal–isobaric Isoenthalpic–isobaric ensemble

Statistical thermodynamics	Characteristic state functions

Partition functions	Translational Vibrational Rotational

Equations of state	Dieterici Van der Waals Ideal gas law Birch–Murnaghan

Entropy	Sackur–Tetrode equation Tsallis entropy Von Neumann entropy

Particle statistics	Maxwell–Boltzmann statistics Fermi–Dirac statistics Bose–Einstein statistics

Statistical field theory	Conformal field theory Osterwalder–Schrader axioms

Quantum statistical mechanics	Density matrix Gibbs measure Partition function Phase space formulation of quantum mechanics Slater determinant

See also	Probability distribution Elementary particles

Diving medicine, physiology, physics and environment

Diving medicine:

Injuries
and
Disorders

Pressure

Oxygen	Deep water blackout Hyperoxia Oxygen toxicity Shallow water blackout

Inert gases	Atrial septal defect Avascular necrosis Decompression sickness Dysbaric osteonecrosis High-pressure nervous syndrome Hydrogen narcosis Isobaric counterdiffusion Nitrogen narcosis Taravana Uncontrolled decompression

Carbon dioxide	Hypercapnia

Immersion

Treatments

Diving physiology

Diving physics

Diving environment

Researchers in
diving medicine,
physiology and physics

Diving medical
research
organisations

Aerospace Medical Association

Divers Alert Network

Diving Diseases Research Centre

Fitness to dive

List of diving hazards and precautions

National Board of Diving and Hyperbaric Medical Technology

Naval Submarine Medical Research Laboratory

Royal Australian Navy School of Underwater Medicine

Rubicon Foundation

South Pacific Underwater Medicine Society

Undersea and Hyperbaric Medical Society

United States Navy Experimental Diving Unit

Categories: Diving medicine
Underwater diving physiology
Underwater diving physics
Physical oceanography
Glossary
Portal

This article is issued from Wikipedia - version of the 12/2/2016. The text is available under the Creative Commons Attribution/Share Alike but additional terms may apply for the media files.

Ideal gas law

Equation

Common form

Molar form

Statistical mechanics

Applications to thermodynamic processes

Deviations from ideal behavior of real gases

Derivations

Empirical

Theoretical

Kinetic theory

Statistical mechanics

See also

References

Further reading

External links