Ideal gas law

The most frequently introduced forms are:

${\displaystyle pV=nRT=nk_{\text{B}}N_{\text{A}}T,}$

where:

• ${\displaystyle p}$ is the pressure of the gas,
• ${\displaystyle V}$ is the volume of the gas,
• ${\displaystyle n}$ is the amount of substance of gas (also known as number of moles),
• ${\displaystyle R}$ is the ideal, or universal, gas constant, equal to the product of the Boltzmann constant and the Avogadro constant,
• ${\displaystyle k_{\text{B}}}$ is the Boltzmann constant
• ${\displaystyle N_{A}}$ is the Avogadro constant
• ${\displaystyle T}$ 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 0.0821 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 (kg) (in kilograms) divided by the molar mass (M) (in kilograms per mole):

${\displaystyle n={\frac {m}{M}}.}$

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

${\displaystyle pV={\frac {m}{M}}RT}$

${\displaystyle p={\frac {m}{V}}{\frac {RT}{M}}}$

${\displaystyle p=\rho {\frac {R}{M}}T}$

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

${\displaystyle p=\rho R_{\text{specific}}T}$

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

${\displaystyle pv=R_{\text{specific}}T.}$

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 ${\displaystyle {\bar {R}}}$ or ${\displaystyle R^{*}}$ 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]

In statistical mechanics the following molecular equation is derived from first principles

${\displaystyle P=nk_{\text{B}}T,}$

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:

${\displaystyle k_{\text{B}}={\frac {R}{N_{\text{A}}}}}$

where N_A is the Avogadro constant.

From this we notice that for a gas of mass m, 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

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

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

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

In SI units, P is measured in pascals, V in cubic metres, and T in measured kelvins. k_B has the value 1.38·10⁻²³ J/K 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 ${\displaystyle PV}$ to ${\displaystyle T}$ is simply taken as a constant:[6]

${\displaystyle {\frac {PV}{T}}=k,}$

where ${\displaystyle p}$ is the pressure of the gas, ${\displaystyle V}$ is the volume of the gas, ${\displaystyle T}$ is the absolute temperature of the gas, and ${\displaystyle k}$ is a constant. When comparing the same substance under two different sets of conditions, the law can be written as

${\displaystyle {\frac {P_{1}V_{1}}{T_{1}}}={\frac {P_{2}V_{2}}{T_{2}}}.}$

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:

${\displaystyle {PV}=C_{1}}$ or ${\displaystyle P_{1}V_{1}=P_{2}V_{2}}$ (1) known as Boyle's law

${\displaystyle {\frac {V}{T}}=C_{2}}$ or ${\displaystyle {\frac {V_{1}}{T_{1}}}={\frac {V_{2}}{T_{2}}}}$ (2) known as Charles's law

${\displaystyle {\frac {V}{N}}=C_{3}}$ or ${\displaystyle {\frac {V_{1}}{N_{1}}}={\frac {V_{2}}{N_{2}}}}$ (3) known as Avogadro's law

${\displaystyle {\frac {P}{T}}=C_{4}}$or ${\displaystyle {\frac {P_{1}}{T_{1}}}={\frac {P_{2}}{T_{2}}}}$ (4) known as Gay-Lussac's law

${\displaystyle NT=C_{5}}$or ${\displaystyle N_{1}T_{1}=N_{2}T_{2}}$ (5)

${\displaystyle {\frac {P}{N}}=C_{6}}$ or ${\displaystyle {\frac {P_{1}}{N_{1}}}={\frac {P_{2}}{N_{2}}}}$ (6)

Relationships between Boyle's, Charles's, Gay-Lussac's, Avogadro's, combined and ideal gas laws, with the Boltzmann constant k_B = R/N_A = n R/N  (in each law, properties circled are constant and properties not circled are variable)

where "P" stands for pressure, "V" for volume, "N" for number of particles in the gas and "T" for temperature; Where ${\displaystyle C_{1},C_{2},C_{3},C_{4},C_{5},C_{6}}$ 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 ${\displaystyle P_{1},V_{1},N_{1},T_{1}}$

Say, starting to change only pressure and volume, according to Boyle's law, then:

${\displaystyle P_{1}V_{1}=P_{2}V_{2}}$ (7) After this process, the gas has parameters ${\displaystyle P_{2},V_{2},N_{1},T_{1}}$

Using then Eq. (5) to change the number of particles in the gas and the temperature,

${\displaystyle N_{1}T_{1}=N_{2}T_{2}}$ (8) After this process, the gas has parameters ${\displaystyle P_{2},V_{2},N_{2},T_{2}}$

Using then Eq. (6) to change the pressure and the number of particles,

${\displaystyle {\frac {P_{2}}{N_{2}}}={\frac {P_{3}}{N_{3}}}}$ (9) After this process, the gas has parameters ${\displaystyle P_{3},V_{2},N_{3},T_{2}}$

Using then Charles's law to change the volume and temperature of the gas,

${\displaystyle {\frac {V_{2}}{T_{2}}}={\frac {V_{3}}{T_{3}}}}$ (10) After this process, the gas has parameters ${\displaystyle P_{3},V_{3},N_{3},T_{3}}$

Using simple algebra on equations (7), (8), (9) and (10) yields the result:

${\displaystyle {\frac {P_{1}V_{1}}{N_{1}T_{1}}}={\frac {P_{3}V_{3}}{N_{3}T_{3}}}}$ or ${\displaystyle {\frac {PV}{NT}}=K_{B}}$, Where ${\displaystyle K_{B}}$ stands for Boltzmann's constant.

Another equivalent result, using the fact that ${\displaystyle nR=NK_{B}}$ ,where "n" is the number of moles in the gas and "R" is the universal gas constant, is:

${\displaystyle PV=nRT,}$ which is known as the ideal gas law.

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:

Relationship between the 6 gas laws

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: ${\displaystyle P_{1}V_{1}=P_{2}V_{2}}$ (1´)

then only volume and temperature: ${\displaystyle {\frac {V_{2}}{T_{1}}}={\frac {V_{3}}{T_{2}}}}$ (2´)

then as we can choose any value for ${\displaystyle V_{3}}$, if we set ${\displaystyle V_{1}=V_{3}}$, Eq. (2´) becomes: ${\displaystyle {\frac {V_{2}}{T_{1}}}={\frac {V_{1}}{T_{2}}}}$(3´)

combining equations (1´) and (3´) yields ${\displaystyle {\frac {P_{1}}{T_{1}}}={\frac {P_{2}}{T_{2}}}}$ , 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

${\displaystyle PV={\frac {1}{3}}Nmv_{\text{rms}}^{2}.}$

Using the Maxwell–Boltzmann distribution, the fraction of molecules that have a speed in the range ${\displaystyle v}$ to ${\displaystyle v+dv}$ is ${\displaystyle f(v)\,dv}$, where

${\displaystyle f(v)=4\pi \left({\frac {m}{2\pi kT}}\right)^{\frac {3}{2}}v^{2}e^{-{\frac {mv^{2}}{2kT}}}}$

and ${\displaystyle k}$ denotes the Boltzmann constant. The root-mean-square speed can be calculated by

${\displaystyle v_{\text{rms}}^{2}=\int _{0}^{\infty }v^{2}f(v)\,dv=4\pi \left({\frac {m}{2\pi kT}}\right)^{\frac {3}{2}}\int _{0}^{\infty }v^{4}e^{-{\frac {mv^{2}}{2kT}}}\,dv.}$

Using the integration formula

${\displaystyle \int _{0}^{\infty }x^{2n}e^{-{\frac {x^{2}}{a^{2}}}}\,dx={\sqrt {\pi }}\,{\frac {(2n)!}{n!}}\left({\frac {a}{2}}\right)^{2n+1},}$

it follows that

${\displaystyle v_{\text{rms}}^{2}=4\pi \left({\frac {m}{2\pi kT}}\right)^{\frac {3}{2}}{\sqrt {\pi }}\,{\frac {4!}{2!}}\left({\frac {\sqrt {\frac {2kT}{m}}}{2}}\right)^{5}={\frac {3kT}{m}},}$

from which we get the ideal gas law:

${\displaystyle PV={\frac {1}{3}}Nm\left({\frac {3kT}{m}}\right)=NkT.}$

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:

${\displaystyle {\begin{aligned}\langle \mathbf {q} \cdot \mathbf {F} \rangle &={\Bigl \langle }q_{x}{\frac {dp_{x}}{dt}}{\Bigr \rangle }+{\Bigl \langle }q_{y}{\frac {dp_{y}}{dt}}{\Bigr \rangle }+{\Bigl \langle }q_{z}{\frac {dp_{z}}{dt}}{\Bigr \rangle }\\&=-{\Bigl \langle }q_{x}{\frac {\partial H}{\partial q_{x}}}{\Bigr \rangle }-{\Bigl \langle }q_{y}{\frac {\partial H}{\partial q_{y}}}{\Bigr \rangle }-{\Bigl \langle }q_{z}{\frac {\partial H}{\partial q_{z}}}{\Bigr \rangle }=-3k_{B}T,\end{aligned}}}$

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

${\displaystyle 3Nk_{B}T=-{\biggl \langle }\sum _{k=1}^{N}\mathbf {q} _{k}\cdot \mathbf {F} _{k}{\biggr \rangle }.}$

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

${\displaystyle -{\biggl \langle }\sum _{k=1}^{N}\mathbf {q} _{k}\cdot \mathbf {F} _{k}{\biggr \rangle }=P\oint _{\mathrm {surface} }\mathbf {q} \cdot d\mathbf {S} ,}$

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

${\displaystyle \nabla \cdot \mathbf {q} ={\frac {\partial q_{x}}{\partial q_{x}}}+{\frac {\partial q_{y}}{\partial q_{y}}}+{\frac {\partial q_{z}}{\partial q_{z}}}=3,}$

the divergence theorem implies that

${\displaystyle P\oint _{\mathrm {surface} }\mathbf {q} \cdot d\mathbf {S} =P\int _{\mathrm {volume} }\left(\nabla \cdot \mathbf {q} \right)dV=3PV,}$

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

Putting these equalities together yields

${\displaystyle 3Nk_{B}T=-{\biggl \langle }\sum _{k=1}^{N}\mathbf {q} _{k}\cdot \mathbf {F} _{k}{\biggr \rangle }=3PV,}$

which immediately implies the ideal gas law for N particles:

${\displaystyle PV=Nk_{B}T=nRT,\,}$

where n = N/N_A is the number of moles of gas and R = N_Ak_B is the gas constant.