The aim of kinetic theory is to account for the properties of gases in terms of the forces between the molecules, assuming that their motions are described by the laws of mechanics (usually classical Newtonian mechanics, although quantum mechanics is needed in some cases). The present discussion focuses on dilute ideal gases, in which molecular collisions of at most two bodies are of primary importance. Only the simplest theories are treated here in order to avoid obscuring the fundamental physics with complex mathematics.
The ideal gas equation of state can be deduced by calculating the pressure as caused by molecular impacts on a container wall. The internal energy and John Dalton’s law of partial pressures also emerge from this calculation, along with some free-molecule phenomena. The calculation is significant because it is basically the same one used to explain all dilute-gas phenomena.
Sir Isaac Newton’s second law of motion can be stated in not-so-familiar form as impulse equals change in momentum, where impulse is force multiplied by the time during which it acts. A molecule experiences a change in momentum when it collides with a container wall. During the collision an impulse is imparted by the wall to the molecule that is equal and opposite to the impulse imparted by the molecule to the wall. This is required by Newton’s third law. The sum of the impulses imparted by all the molecules to the wall is, in effect, the pressure.
Consider a system of molecules of mass m traveling with a velocity v in an enclosed container. In order to arrive at an expression for the pressure, a calculation will be made of the impulse imparted to one of the walls by a single impact, followed by a calculation of how many impacts occur on that wall during a time t. Although the molecules are moving in all directions, only those with a component of velocity toward the wall can collide with it; call this component vz, where z represents the direction directly toward the wall. Not all molecules have the same vz, of course; perhaps only Nz out of a total of N molecules do. To find the total pressure, the contributions from molecules with all different values of vz must be summed. A molecule approaches the wall with an initial momentum mvz, and after impact it moves away from the wall with an equal momentum in the opposite direction, -mvz. Thus, the total change in momentum is mvz - (-mvz) = 2mvz, which is equal to the total impulse imparted to the wall.
The number of impacts on a small area A of the wall in time t is equal to the number of molecules that reach the wall in time t. Since the molecules are traveling at speed vz, only those within a distance vz t and moving toward the wall will reach it in that time. Thus, the molecules that are traveling toward the wall and are within a volume Avz t will strike the area A of the wall in time t. On the average, half of the molecules in this volume will be moving toward the wall. If Nz molecules with speed component vz are present in the total volume V, then (1/2)(Nz/V)(A)(vzt) molecules in the collision volume will hit, and each one contributes an impulse of 2mvz. The total impulse in time t is therefore (1/2)(Nz/V) (A)(vzt)(2mvz) = (Nz//V)(mvz2)(At), which is equal to Ft, where F is the force on the wall due to the impacts. Equating these two expressions, the time factor t cancels out. Since pressure is defined as the force per unit area (F/A), it follows that the contribution to the pressure from the molecules with speed vz is thus (Nz/V)mvz2. Because there are different values of vz for different molecules, the average value, denoted vz2, is used to take into account the contributions from all the molecules. The pressure is thus given as p = (N/V)mvz2.
Since the molecules are in random motion, this result is independent of the choice of axis. For any choice of (x, y, z) axes, the magnitude of the velocity is v2 = vx2 + vy2 + vz2 (which is just the Pythagorean theorem in three dimensions), and taking the average gives v2 = vx2 + vy2 + vz2. The gas is in equilibrium, so it must appear the same in any direction, and the average velocities are therefore the same in all directions—i.e., vx2 = vy2 = vz2; thus v2 = 3vz2. When the value (1/3)v2 is substituted for vz2 in the expression for pressure, the following equation is obtained:
To rewrite this in molar units, N is set equal to nN0—i.e., the product of the number of moles n and Avogadro’s number N0—to give
where M = N0m is the molecular weight of the gas and v is the volume per mole (V/n). Since the ideal gas equation of state relates pressure, molar volume, and temperature as pv = RT, the temperature T must be related to the average kinetic energy of the molecules as
This expression is often written in molecular (rather than molar) terms as (1/2)[mv2] = (3/2)kT, where k = R/N0 is called Boltzmann’s constant, named after Ludwig Boltzmann, a 19th-century Austrian physicist who substantially contributed to the foundation and development of statistical mechanics. If the gas is a mixture, the foregoing calculation shows that the impacts of the different species are simply added separately, and Dalton’s law of partial pressures follows directly.
The energy law given as equation (16) also follows from equation (19): the kinetic energy of translational motion per mole is (3/2)RT. Any energy residing in the internal motions of the individual molecules is simply carried separately without contributing to the pressure.
Average molecular speeds can be calculated from the results of kinetic theory in terms of the so-called root-mean-square speed vrms. The vrms is the square root of the average of the squares of the speeds of the molecules: (v2)1/2. From equation (19) the vrms is (3RT/M)l/2. At 20 °C (68 °F) the value for air (M = 29) is 502 metres (1,647 feet) per second, a result very close to the rough estimate of 500 metres per second given above.
Molecule-molecule collisions were not considered in the calculation of the expression for pressure even though many such collisions occur. Such collisions could be ignored because they are elastic; i.e., linear momentum is conserved in the collision, provided that no external forces act. Two molecules therefore continue to carry the same momentum to the wall even if they collide with one another before striking it. The ideal gas equation of state remains valid as the density is decreased, even holding for a free-molecule gas. The equation eventually fails as the density is increased, however, because other molecules exert forces and change the rate of collisions with the walls.
It was not until the mid- to late 19th century that kinetic theory was successfully applied to such calculations as gas pressure. Such notable scientists as Newton and Dalton had believed that gas pressure was caused by repulsions between molecules that pushed them against the container walls. For many reasons, the kinetic theory had overshadowed such static theories (and others such as vortex theories) by about 1860. It was not until 1875, however, that James Clerk Maxwell actually proved that a static theory was in conflict with experiment.
Consider the system described above in the calculation of gas pressure, but with the area A in the container wall replaced with a small hole. The number of molecules that escape through the hole in time t is equal to (1/2)(N/V)vz (At). In this case, collisions between molecules are significant, and the result holds only for tiny holes in very thin walls (as compared to the mean free path), so that a molecule that approaches near the hole will get through without colliding with another molecule and being deflected away. The relationship between vz and the average speed 
is rather straightforward: vz = (1/2).
If the rates for two different gases effusing through the same hole are compared, starting with the same gas density each time, it is found that much more light gas escapes than heavy gas and that more gas escapes at a high temperature than at a low temperature, other things being equal. In particular,
The last step follows from the energy formula, (1/2)mv2 = (3/2)kT, where (v2)1/2 is approximated to be v, even though v2 and ()2 actually differ by a numerical factor near unity (namely, 3π/8). This result was discovered experimentally in 1846 by Graham for the case of constant temperature and is known as Graham’s law of effusion. It can be used to measure molecular weights, to measure the vapour pressure of a material with a low vapour pressure, or to calculate the rate of evaporation of molecules from a liquid or solid surface.
Suppose that two containers of the same gas but at different temperatures are connected by a tiny hole and that the gas is brought to a steady state. If the hole is small enough and the gas density is low enough that only effusion occurs, the equilibrium pressure will be greater on the high-temperature side. But, if the initial pressures on both sides are equal, gas will flow from the low-temperature side to the high-temperature side to cause the high-temperature pressure to increase. The latter situation is called thermal transpiration, and the steady-state result is called the thermomolecular pressure difference. These results follow simply from the effusion formula if the ideal gas law is used to replace N/V with p/T;
When a steady state is reached, the effusion rates are equal, and thus
This phenomenon was first investigated experimentally by Osborne Reynolds in 1879 in Manchester, England. Errors can result if a gas pressure is measured in a vessel at very low or very high temperature by connecting it via a fine tube to a manometer at room temperature. A continuous circulation of gas can be produced by connecting the two containers with another tube whose diameter is large compared with the mean free path. The pressure difference drives gas through this tube by viscous flow. A heat engine based on this circulating flow unfortunately has a low efficiency.
The kinetic-theory explanation of viscosity can be simplified by examining it in qualitative terms. Viscosity is caused by the transfer of momentum between two planes sliding parallel to one another but at different rates, and this momentum is transferred by molecules moving between the planes. Molecules from the faster plane move to the slower plane and tend to speed it up, while molecules from the slower plane travel to the faster plane and tend to slow it down. This is the mechanism by which one plane experiences the drag of the other. A simple analogy is two mail trains passing each other, with workers throwing mailbags between the trains. Every time a mailbag from the fast-moving train lands on the slow one, it imparts its momentum to the slow train, speeding it up a little; likewise each mailbag from the slow train that lands on the fast one slows it down a bit.
If the trains are too far apart, the mailbags cannot be passed between them. Similarly, the planes of a gas must be only about a mean free path apart in order for molecules to pass between them without being deflected by collisions. If one uses this approach, a simple calculation can be carried out, much as in the case of the gas pressure, with the result that
where a is a numerical constant of order unity, the term (N/V)l is a measure of the number of molecules contained in a small counting cylinder, and the mass m is a measure of the momentum carried between the sliding planes. The cross-sectional area of the counting cylinder and the relative speed of the sliding planes do not appear in the equation because they cancel one another when the drag force is divided by the area and speed of the planes in order to find η.
It can now be seen why η is independent of gas density or pressure. The term (N/V) in equation (23) is the number of carriers of momentum, but l measures the number of collisions that interfere with these carriers and is inversely proportional to (N/V). The two effects exactly cancel each other. Viscosity increases with temperature because the average velocity does; that is, momentum is carried more quickly when the molecules move faster. Although increases as T1/2, η increases somewhat faster because the mean free path also increases with temperature, since it is harder to deflect a fast molecule than a slow one. This feature depends explicitly on the forces between the molecules and is difficult to calculate accurately, as is the value of the constant a, which turns out to be close to 1/2.
The behaviour of the viscosity of a mixture can also be explained by the foregoing calculation. In a mixture of a light gas and a viscous heavy gas, both types of molecules have the same average energy. However, most of the momentum is carried by the heavy molecules, which are therefore the main contributors to the viscosity. The light molecules are rather ineffective in deflecting the heavy molecules, so that the latter continue to carry virtually as much momentum as they would in the absence of light molecules. The addition of a light gas to a heavy gas therefore does not reduce the viscosity substantially and may in fact increase it because of the small extra momentum carried by the light molecules. The viscosity will eventually decrease when there are only a few heavy molecules remaining in a large sea of light molecules.
The main dependence of η on the molecular mass is through the product m in equation (23), which varies as m½ since varies as 1/m1/2. Owing to this effect, heavy gases tend to be more viscous than light gases, but this tendency is compensated for to some degree by the behaviour of l, which tends to be smaller for heavy molecules because they are usually larger than light molecules and therefore more likely to collide. The often confusing connection between viscosity and molecular weight can thus be accounted for by equation (23).
Finally, in a free-molecule gas there are no collisions with other molecules to impede the transport of momentum, and the viscosity thus increases linearly with pressure or density until the number of collisions becomes great enough so that the viscosity assumes the constant value given by equation (23). The nonideal behaviour of the gas that accompanies further increases in density eventually leads to an increase in viscosity, and the viscosity of an extremely dense gas becomes much like that of a liquid.
The kinetic-theory explanation of heat conduction is similar to that for viscosity, but in this case the molecules carry net energy from a region of higher energy (i.e., temperature) to one of lower energy (temperature). Internal molecular motions must be accounted for because, though they do not transport momentum, they do transport energy. Monatomic gases, which carry only their kinetic energy of translational motion, are the simplest case. The resulting expression for thermal conductivity is
which has the same basic form as equation (23) for viscosity, with (3k/2) replacing m. The (3k/2) is the heat capacity per molecule and is the conversion factor from an energy difference to a temperature difference.
It can be shown from equation (24) that the independence of density and the increase with temperature is the same for thermal conductivity as it is for viscosity. The dependence on molecular mass is different, however, with λ varying as 1/m1/2 owing to the factor . Thus, light gases tend to be better conductors of heat than are heavy gases, and this tendency is usually augmented by the behaviour of l.
The behaviour of the thermal conductivity of mixtures may be qualitatively explained. Adding heavy gas to light gas reduces the thermal conductivity because the heavy molecules carry less energy and also interfere with the energy transport of the light molecules.
The similar behaviour of λ and η suggests that their ratio might provide information about the constants a and a′. The ratio of a′/a is given as
Although simple theory suggests that this ratio should be about one, both experiment and more refined theory give a value close to 5/2. This means that molecules do not “forget” their past history in every collision, but some persistence of their precollision velocities occurs. Molecules transport both energy and momentum from a somewhat greater distance than just one mean free path, but this distance is greater for energy than for momentum. This is plausible, for molecules with higher kinetic energies might be expected to have greater persistences.
Attempts to calculate the constants a and a′ by tracing collision histories to find the “persistence of velocities” have not met with much success. The molecular “memory” fades slowly, too many previous collisions have to be traced, and the calculations become almost hopelessly complicated. A different theoretical approach is needed, which was supplied about 1916–17 independently by David Enskog and Sydney Chapman. Their theory also shows that the same value of l applies to both η and λ, a fact that is not obvious in the simple theory described here.
The thermal conductivity of polyatomic molecules is accounted for by simply adding on a contribution for the energy carried by the internal molecular motions:
where cint is the contribution of the internal motions to the heat capacity (per molecule) and is easily found by subtracting (3k/2) from the total measured heat capacity. As might be expected, the constant a’ is only about half as large as a′.
The pressure or density dependence of λ must be similar to that of η—an initial linear increase in the free-molecule region, followed by a constant value in the dilute-gas region and finally an increase in the dense-fluid region.
Both of these properties present difficulties for the simple mean free path version of kinetic theory. In the case of diffusion it must be argued that collisions of the molecules of species 1 with other species 1 molecules do not inhibit the interdiffusion of species 1 and 2, and similarly for 2–2 collisions. If this is not assumed, the calculated value of the diffusion coefficient for the 1–2 gas pair, D12, depends strongly on the mixture composition instead of being virtually independent of it, as is shown by experiment. The neglect of 1–1 and 2–2 collisions can be rationalized by noting that the flow of momentum is not disturbed by such like-molecule collisions owing to the conservation of momentum, but it can be contended that the argument was simply invented to make the theory agree with experiment. A more charitable view is that the experimental results demonstrate that collisions between like molecules have little affect on D12. It is one of the triumphs of the accurate kinetic theory of Enskog and Chapman that this result clearly emerges.
If 1–1 and 2–2 collisions are ignored, a simple calculation gives a result much like those for η and λ:
where a12 is a numerical constant, v12 is an average relative speed for 1–2 collisions given by v122 = (½)(v12 + v22), and l12 is a mean free path for 1–2 collisions that is inversely proportional to the total molecular number density, (N1 + N2)/V. Thus, D12 is inversely proportional to gas density or pressure, unlike η and λ, but the concentration difference is proportional to pressure, with the two effects canceling one another, as pointed out previously. The actual transport of molecules is therefore independent of pressure. The numerical value of a12, as obtained by refined calculations, is close to 3/5.
The pressure dependence of pD12 should be qualitatively similar to that of η and λ—an initial linear increase in the free-molecule region, a constant value in the dilute-gas region, and finally an increase in the dense-fluid region.
Thermal diffusion presents special difficulties for kinetic theory. The transport coefficients η, λ, and D12 are always positive regardless of the nature of the intermolecular forces that produce the collisions—the mere existence of collisions suffices to account for their important features. The transport coefficient that describes thermal diffusion, however, depends critically on the nature of the intermolecular forces and the collisions and can be positive, negative, or zero. Its dependence on composition is also rather complicated. There have been a number of attempts to explain thermal diffusion with a simple mean free path model, but none has been satisfactory. No simple physical explanation of thermal diffusion has been devised, and recourse to the accurate, but complicated, kinetic theory is necessary.
The simple mean free path description of gas transport coefficients accounts for the major observed phenomena, but it is quantitatively unsatisfactory with respect to two major points: the values of numerical constants such as a, a′, a”, and a12 and the description of the molecular collisions that define a mean free path. Indeed, collisions remain a somewhat vague concept except when they are considered to take place between molecules modeled as hard spheres. Improvement has required a different, somewhat indirect, and more mathematical approach through a quantity called the velocity distribution function. This function describes how molecular velocities are distributed on the average: a few very slow molecules, a few very fast ones, and most near some average value—namely, vrms = (v2)1/2 = (3kT/2)1/2. If this function is known, all gas properties can be calculated by using it to obtain various averages. For example, the average momentum carried in a certain direction would give the viscosity.
The velocity distribution for a gas at equilibrium was suggested by Maxwell in 1859 and is represented by the familiar bell-shaped curve that describes the normal, or Gaussian, distribution of random variables in large populations. Attempts to support more definitively this result and to extend it to nonequilibrium gases led to the formulation of the Boltzmann equation, which describes how collisions and external forces cause the velocity distribution to change. This equation is difficult to solve in any general sense, but some progress can be made by assuming that the deviations from the equilibrium distribution are small and are proportional to the external influences that cause the deviations, such as temperature, pressure, and composition differences.
Even the resulting simpler equations remained unsolved for nearly 50 years until the work of Enskog and Chapman, with a single notable exception. The one case that was solvable dealt with molecules that interact with forces that fall off as the fifth power of their separation (i.e., as 1/r5), for which Maxwell found an exact solution. Unfortunately, thermal diffusion happens to be exactly zero for molecules subject to this force law, so that phenomenon was missed.
It was later discovered that it is possible to use the solutions for the 1/r5 Maxwell model as a starting point and then calculate successive corrections for more general interactions. Although the calculations quickly increase in complexity, the improvement in accuracy is rapid, unlike the persistence-of-velocities corrections applied in mean free path theory. This refined version of kinetic theory is now highly developed, but it is quite mathematical and is not described here.
Deviations from ideal gas behaviour occur both at low densities, where molecule-surface collisions become important, and at high densities, where a description in terms of only two-body collisions becomes inadequate. The low-density case can be handled in principle by including both molecule-surface and molecule-molecule collisions in the Boltzmann equation. Since this branch of the subject is now quite advanced and mathematical in character, only the high-density case will be discussed here.
To a first approximation, molecule-molecule collisions do not affect the ideal gas equation of state, pv = RT, but real gases at nonzero densities show deviations from this equation that are due to interactions among the molecules. Ever since the great advance made by van der Waals in 1873, an accurate universal formula relating p, v, and T has been sought. No completely satisfactory equation of state has been found, though important advances occurred in the 1970s and ’80s. The only rigorous theoretical result available is an infinite-series expansion in powers of 1/v, known as the virial equation of state:
where B(T), C(T), … are called the second, third, … virial coefficients and depend only on the temperature and the particular gas. The virtue of this equation is that there is a rigorous connection between the virial coefficients and intermolecular forces, and experimental values of B(T) were an early source (and still a useful one) of quantitative information on intermolecular forces. The drawback of the virial equation of state is that it is an infinite series and becomes essentially useless at high densities, which in practice are those greater than about the critical density. Also, the equation is wanting in that it does not predict condensation.
The most practical approaches to the equation of state for real fluids remain the versions of the principle of corresponding states first proposed by van der Waals.
Despite many attempts, there is still no satisfactory theory of the transport properties of dense fluids. Even the extension of the Boltzmann equation to include collisions of more than two bodies is not entirely clear. An important advance was made in 1921 by Enskog, but it is restricted to hard spheres and has not been extended to real molecules except in an empirical way to fit experimental measurements.
Attempts to develop a virial type of expansion in 1/v for the transport coefficients have failed in a surprising way. A formal theory was formulated, but, when the virial coefficients were evaluated for the tractable case of hard spheres, an infinite result was obtained for the coefficient of the 1/v2 term. This is a signal that a virial expansion is not accurate in a mathematical sense, and subsequent research showed that the error arose from a neglected term of the form (1/v2)ln(1/v). It remains unknown how many similar problematic mathematical terms exist in the theory. Transport coefficients of dense fluids are usually described by some empirical extension of the Enskog hard-sphere theory or more commonly by some version of a principle of corresponding states. Much work clearly remains to be done.