Ideal gas law

Isotherms of an ideal gas. The curved lines represent the relationship between pressure (on the vertical axis) and volume (on the horizontal 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's law and Avogadro's Law.[1] The ideal gas law is often written as

where:

It can also be derived from the microscopic 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

where:

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 total mass of the gas (m) (in grams) divided by the molar mass (M) (in grams per mole):

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

Defining the specific gas constant Rspecific(r) as the ratio R/M,

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

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 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

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, in contrasts to the previous formulation in which n is the number of moles), T is the absolute temperature, and kB is the Boltzmann constant relating temperature and energy, given by:

where NA is Avogadro's constant.

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

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

In SI units, P is measured in pascals, V in cubic metres, and T in measured kelvins. kB has the value 1.38·10−23 J/K in SI units.

Energy associated with a gas

According to the assumptions of the kinetic theory of gases, we assumed that there are no inter molecular attractions between the molecules of an ideal gas its potential energy is zero. hence, all the energy possessed by the gas is kinetic energy.

This is the kinetic energy of one mole of a gas.

Energy of gas Mathematical formula
energy associated with one mole of a gas
energy associated with one gram of a gas
energy associated with one molecule of a gas

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, S, or H) 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 or delta P2 V2 T2
Isobaric process
Pressure
V2/V1
 P2 = P1 V2 = V1(V2/V1) T2 = T1(V2/V1)
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)
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
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)
Isenthalpic process
(Irreversible adiabatic process)
Enthalpy[b]
P2 − P1
 P2 = P1 + (P2 − P1) T2 = T1 + μJT(P2 − P1)
T2 − T1
 P2 = P1 + (T2 − T1)/μJT T2 = T1 + (T2 − T1)

^ a. In an isentropic process, system entropy (S) is constant. Under these conditions, P1 V1γ = P2 V2γ, 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 mono atomic 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.

^ b. In an isenthalpic process, system enthalpy (H) is constant. In the case of free expansion for an ideal gas, there are no molecular interactions, and the temperature remains constant. For real gasses, the molecules do interact via attraction or repulsion depending on temperature and pressure, and heating or cooling does occur. This is known as the Joule–Thomson effect. For reference, the Joule–Thomson coefficient μJT for air at room temperature and sea level is 0.22 °C/bar.[6]

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 inter molecular 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

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

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 = (qx, qy, qz) and p = (px, py, pz) 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:

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

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

where dS is the infinitesimal area element along the walls of the container. Since the divergence of the position vector q is

the divergence theorem implies that

where dV is an infinitesimal volume within the container and V is the total volume of the container.

Putting these equalities together yields

which immediately implies the ideal gas law for N particles:

where n = N/NA is the number of moles of gas and R = NAkB is the gas constant.

See also

References

  1. ↑ 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).
  2. ↑ 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).
  3. ↑ 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).
  4. ↑ "Equation of State".
  5. ↑ Moran and Shapiro, Fundamentals of Engineering Thermodynamics, Wiley, 4th Ed, 2000.
  6. ↑ J. R. Roebuck (1926). "The Joule-Thomson Effect in Air". Proceedings of the National Academy of Sciences of the United States of America. 12 (1): 55–58. PMC 1084398 Freely accessible.

Further reading

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