The ideal gas law, also called the general gas equation, 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 Benoît Paul Émile Clapeyron in 1834 as a combination of the empirical Boyle's law, Charles's law, Avogadro's law, and Gay-Lussac's law.^{[1]} The ideal gas law is often written in an empirical form:
where , and are the pressure, volume and temperature; is the amount of substance; and is the ideal gas constant. It is the same for all gases. 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]}
Note that this law makes no comment as to whether a gas heats or cools during compression or expansion. An ideal gas may not change temperature, but most gases like air are not ideal and follow the Joule-Thomson effect.^{[dubious – discuss]}
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]}
The most frequently introduced forms are:
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 0.0821 l·atm/(mol·K).
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 total mass of the gas (m) (in kilograms) divided by the molar mass (M) (in kilograms per mole):
By replacing n with m/M and subsequently introducing density ? = m/V, we get:
Defining the specific gas constant R_{specific(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 and meteorological 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 or 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]}
In statistical mechanics the following molecular equation is derived from first p
where P is the absolute pressure of the gas, n is the number density of the molecules (given by the ratio n = N/V, in contrast to the previous formulation in which n is the number of moles), T is the absolute temperature, and k_{B} is the Boltzmann constant relating temperature and energy, given by:
where N_{A} is the Avogadro constant.
From this we notice that for a gas of mass m, with an average particle mass of ? times the atomic mass constant, m_{u}, (i.e., the mass is ? u) the number of molecules will be given by
and since ? = m/V = n?m_{u}, we find that the ideal gas law can be rewritten as
In SI units, P is measured in pascals, V in cubic metres, T in kelvins, and k_{B} = in SI units.
Combining the laws of Charles, Boyle and Gay-Lussac gives the combined gas law, which takes the same functional form as the ideal gas law save that the number of moles is unspecified, and the ratio of to is simply taken as a constant:^{[6]}
where is the pressure of the gas, is the volume of the gas, is the absolute temperature of the gas, and is a constant. When comparing the same substance under two different sets of conditions, the law can be written as
According to assumptions of the kinetic theory of ideal gases, we assume that there are no intermolecular attractions between the molecules of an ideal gas. In other words, its potential energy is zero. Hence, all the energy possessed by the gas is in the kinetic energy of the molecules of the gas.
This is the kinetic energy of n moles of a monatomic gas having 3 degrees of freedom; x, y, z.
Energy of gas | Mathematical expression |
---|---|
energy associated with one mole of a monatomic gas | |
energy associated with one gram of a monatomic gas | |
energy associated with one molecule (or atom) of a monatomic gas |
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 | p_{2} | V_{2} | T_{2} |
---|---|---|---|---|---|
Isobaric process | p_{2} = p_{1} | V_{2} = V_{1}(V_{2}/V_{1}) | T_{2} = T_{1}(V_{2}/V_{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) |
p_{2} = p_{1}(p_{2}/p_{1}) | V_{2} = V_{1} | T_{2} = T_{1}(p_{2}/p_{1}) | ||
p_{2} = p_{1}(T_{2}/T_{1}) | V_{2} = V_{1} | T_{2} = T_{1}(T_{2}/T_{1}) | |||
Isothermal process | p_{2} = p_{1}(p_{2}/p_{1}) | V_{2} = V_{1}/(p_{2}/p_{1}) | T_{2} = T_{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) |
p_{2} = p_{1}(p_{2}/p_{1}) | V_{2} = V_{1}(p_{2}/p_{1})^{(-1/?)} | T_{2} = T_{1}(p_{2}/p_{1})^{(? - 1)/?} | ||
p_{2} = p_{1}(V_{2}/V_{1})^{-?} | V_{2} = V_{1}(V_{2}/V_{1}) | T_{2} = T_{1}(V_{2}/V_{1})^{(1 - ?)} | |||
p_{2} = p_{1}(T_{2}/T_{1})^{?/(? - 1)} | V_{2} = V_{1}(T_{2}/T_{1})^{1/(1 - ?) } | T_{2} = T_{1}(T_{2}/T_{1}) | |||
Polytropic process | 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} | ||
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)} | |||
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}) | |||
Isenthalpic process (Irreversible adiabatic process) |
p_{2} = p_{1} + (p_{2} - p_{1}) | T_{2} = T_{1} + ?_{JT}(p_{2} - p_{1}) | |||
p_{2} = p_{1} + (T_{2} - T_{1})/?_{JT} | 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}^{?} = p_{2}V_{2}^{?}, where ? is defined as the heat capacity ratio, which is constant for a calorifically perfect gas. The value used for ? is typically 1.4 for diatomic gases like nitrogen (N_{2}) and oxygen (O_{2}), (and air, which is 99% diatomic). Also ? is typically 1.6 for mono atomic gases like the noble gases helium (He), and argon (Ar). In internal combustion engines ? 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 ?_{JT} for air at room temperature and sea level is 0.22 °C/bar.^{[7]}
The equation of state given here (PV=nRT) 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.
The empirical laws that led to the derivation of the ideal gas law were discovered with experiments that changed only 2 state variables of the gas and kept every other one constant.
All the possible gas laws that could have been discovered with this kind of setup are:
where "P" stands for pressure, "V" for volume, "N" for number of particles in the gas and "T" for temperature; where are not actual constants but are in this context because of each equation requiring only the parameters explicitly noted in it changing.
To derive the ideal gas law one does not need to know all 6 formulas, one can just know 3 and with those derive the rest or just one more to be able to get the ideal gas law, which needs 4.
Since each formula only holds when only the state variables involved in said formula change while the others remain constant, we cannot simply use algebra and directly combine them all. I.e. Boyle did his experiments while keeping N and T constant and this must be taken into account.
Keeping this in mind, to carry the derivation on correctly, one must imagine the gas being altered by one process at a time. The derivation using 4 formulas can look like this:
at first the gas has parameters
Say, starting to change only pressure and volume, according to Boyle's law, then:
Using then Eq. (5) to change the number of particles in the gas and the temperature,
Using then Eq. (6) to change the pressure and the number of particles,
Using then Charles's law to change the volume and temperature of the gas,
Using simple algebra on equations (7), (8), (9) and (10) yields the result:
Another equivalent result, using the fact that , where "n" is the number of moles in the gas and "R" is the universal gas constant, is:
If you know or have found with an experiment 3 of the 6 formulas, you can easily derive the rest using the same method explained above; but due to the properties of said equations, namely that they only have 2 variables in them, they can't be any 3 formulas. For example, if you were to have Eqs. (1), (2) and (4) you would not be able to get any more because combining any two of them will give you the third; But if you had Eqs. (1), (2) and (3) you would be able to get all 6 Equations without having to do the rest of the experiments because combining (1) and (2) will yield (4), then (1) and (3) will yield (6), then (4) and (6) will yield (5), as well as would the combination of (2) and (3) as is visually explained in the following visual relation:
where the numbers represent the gas laws numbered above.
If you were to use the same method used above on 2 of the 3 laws on the vertices of one triangle that has a "O" inside it, you would get the third.
For example:
Change only pressure and volume first: (1´)
then only volume and temperature: (2´)
then as we can choose any value for , if we set , Eq. (2´) becomes: (3´)
combining equations (1´) and (3´) yields , which is Eq. (4), of which we had no prior knowledge until this derivation.
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.
The fundamental assumptions of the kinetic theory of gases imply that
Using the Maxwell-Boltzmann distribution, the fraction of molecules that have a speed in the range to is , where
and denotes the Boltzmann constant. The root-mean-square speed can be calculated by
Using the integration formula
it follows that
from which we get the ideal gas law:
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 kinetic 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/N_{A} is the number of moles of gas and R = N_{A}k_{B} is the gas constant.
For a d-dimensional system, the ideal gas pressure is:^{[8]}
where is the volume of the d-dimensional domain in which the gas exists. Note that the dimensions of the pressure changes with dimensionality.