View source for Kinetic theory of gases

[[Image:Translational motion.gif|thumb|upright=1.4|The [[temperature]] of an ideal [[monatomic]] [[gas]] is proportional to the average [[kinetic energy]] of its atoms. The [[Bohr radius|size]] of [[helium]] atoms relative to their spacing is shown to scale under 1950 [[Atmosphere (unit)|atmospheres]] of pressure. The atoms have a certain, average speed, slowed down here two [[1000000000000 (number)|trillion]] fold from room temperature.]]

The '''kinetic theory of gases''' describes a [[gas]] as a large number of submicroscopic [[particle]]s ([[atom]]s or [[molecule]]s), all of which are in constant, rapid, [[randomness|random]] [[motion (physics)|motion]]. The randomness arises from the particles' many collisions with each other and with the walls of the container.

Kinetic theory of gases explains the [[macroscopic scale|macroscopic]] properties of gases, such as pressure, temperature, viscosity, [[thermal conductivity]], and volume, by considering their molecular composition and motion. The theory posits that gas pressure results from particles' collisions with the walls of a container at different velocities.

Kinetic molecular theory defines temperature in its own way, in contrast with the [[temperature|thermodynamic definition]].<ref name="C & C"/>

Under an [[optical microscope]], the molecules making up a liquid are too small to be visible. However, the jittery motion of pollen grains or dust particles in liquid are visible. Known as [[Brownian motion]], the motion of the pollen or dust results from their collisions with the liquid's molecules.

==History==

In approximately 50 [[before common era|BCE]], the Roman philosopher [[Lucretius]] proposed that apparently static macroscopic bodies were composed on a small scale of rapidly moving atoms all bouncing off each other.<ref>{{Cite journal|last1=Maxwell|first1=J. C.|year=1867|title=On the Dynamical Theory of Gases|journal=Philosophical Transactions of the Royal Society of London|volume=157|pages=49–88|doi=10.1098/rstl.1867.0004|pmc=|pmid=}}</ref> This [[Epicureanism|Epicurean]] atomistic point of view was rarely considered in the subsequent centuries, when [[Aristotle]]an ideas were dominant.

[[Image:HYDRODYNAMICA, Danielis Bernoulli.png|thumb|upright|Hydrodynamica front cover]]

In 1738 [[Daniel Bernoulli]] published ''[[Hydrodynamica]]'', which laid the basis for the kinetic theory of gases.  In this work, Bernoulli posited the argument, still used to this day, that gases consist of great numbers of molecules moving in all directions, that their impact on a surface causes the gas pressure that we feel, and that what we experience as [[heat]] is simply the [[kinetic energy]] of their motion. 
Bernoulli also surmised that temperature was the effect of the kinetic energy of the molecules, and thus correlated with the ideal gas law.<ref name=bern>{{cite web|url=https://www.terpconnect.umd.edu/~lvrmr/History/EarlyTheories.html|title=Early Theories of Gases}}</ref>

The theory was not immediately accepted, in part because [[conservation of energy]] had not yet been established, and it was not obvious to physicists how the collisions between molecules could be perfectly elastic.<ref name="PonomarevKurchatov1993">{{cite book|title=The Quantum Dice|author1=L.I Ponomarev|author2=I.V Kurchatov|date=1 January 1993|publisher=CRC Press|isbn=978-0-7503-0251-7}}</ref>{{rp|36–37}}
A competing theory favored by Newton was the repulsion theory, in which heat was a calorific fluid that repulsed molecules in proportion to its quantity (i.e. heat) and the inverse square of the distances between molecules.<ref name=bern/>

Other pioneers of the kinetic theory (which were neglected by their contemporaries) were [[Mikhail Lomonosov]] (1747),<ref>Lomonosov 1758</ref> [[Georges-Louis Le Sage]] (ca. 1780, published 1818),<ref>Le Sage 1780/1818</ref> [[John Herapath]] (1816)<ref>Herapath 1816, 1821</ref> and [[John James Waterston]] (1843),<ref>Waterston 1843</ref> which connected their research with the development of [[mechanical explanations of gravitation]]. In 1856 [[August Krönig]] (probably after reading a paper of Waterston) created a simple gas-kinetic model, which only considered the translational motion of the particles.<ref>Krönig 1856</ref>

In 1857 [[Rudolf Clausius]], according to his own words independently of Krönig, developed a similar, but much more sophisticated version of the theory which included translational and contrary to Krönig also rotational and vibrational molecular motions. In this same work he introduced the concept of [[mean free path]] of a particle.
<ref>Clausius 1857</ref> In 1859, after reading a paper on the diffusion of molecules by [[Rudolf Clausius]], Scottish physicist [[James Clerk Maxwell]] formulated the [[Maxwell distribution]] of molecular velocities, which gave the proportion of molecules having a certain velocity in a specific range.<ref>See:

* Maxwell, J.C. (1860) [https://books.google.com/books?id=-YU7AQAAMAAJ&pg=PA19#v=onepage&q&f=false "Illustrations of the dynamical theory of gases. Part I.  On the motions and collisions of perfectly elastic spheres,"] ''Philosophical Magazine'', 4th series, '''19''' :  19–32.
* Maxwell, J.C. (1860) [https://books.google.com/books?id=DIc7AQAAMAAJ&pg=PA21#v=onepage&q&f=false "Illustrations of the dynamical theory of gases. Part II.  On the process of diffusion of two or more kinds of moving particles among one another,"] ''Philosophical Magazine'', 4th series, '''20''' :  21–37.</ref> This was the first-ever statistical law in physics.<ref>{{cite book|title=The Man Who Changed Everything – the Life of James Clerk Maxwell|author=Mahon, Basil|publisher=Wiley|year=2003|isbn=978-0-470-86171-4|location=Hoboken, NJ|oclc=52358254}}</ref> Maxwell also gave the first mechanical argument that molecular collisions entail an equalization of temperatures and hence a tendency towards equilibrium.<ref>{{Cite journal|last1=Gyenis|first1=Balazs|year=2017|title=Maxwell and the normal distribution: A colored story of probability, independence, and tendency towards equilibrium|journal=Studies in History and Philosophy of Modern Physics|volume=57|pages=53–65|arxiv=1702.01411|bibcode=2017SHPMP..57...53G|doi=10.1016/j.shpsb.2017.01.001}}</ref> In his 1873 thirteen page article 'Molecules', Maxwell states: "we are told that an 'atom' is a material point, invested and surrounded by 'potential forces' and that when 'flying molecules' strike against a solid body in constant succession it causes what is called [[pressure]] of air and other gases."<ref>Maxwell 1875</ref>
In 1871, [[Ludwig Boltzmann]] generalized Maxwell's achievement and formulated the [[Maxwell–Boltzmann distribution]]. Also the [[logarithm]]ic connection between [[entropy]] and [[probability]] was first stated by him.

In the beginning of the twentieth century, however, atoms were considered by many physicists to be purely hypothetical constructs, rather than real objects. An important turning point was [[Albert Einstein]]'s (1905)<ref>Einstein 1905</ref> and [[Marian Smoluchowski]]'s (1906)<ref>Smoluchowski 1906</ref>
papers on [[Brownian motion]], which succeeded in making certain accurate quantitative predictions based on the kinetic theory.

==Assumptions==
The theory for ideal gases makes the following assumptions:

* The gas consists of very small particles known as molecules. This smallness of their size is such that the total [[volume]] of the individual gas molecules added up is negligible compared to the volume of the smallest open ball containing all the molecules. This is equivalent to stating that the average distance separating the gas particles is large compared to their [[atomic radius|size]].
* These particles have the same [[mass]].
* The number of molecules is so large that statistical treatment can be applied.

* The rapidly moving particles constantly collide among themselves and with the walls of the container. Collisions between particles and the wall of the container are non elastic whereas collisions between the particles are elastic. This means the molecules are considered to be perfectly spherical in shape and elastic in nature.
* Except during collisions, the [[interaction]]s among molecules are negligible. (That is, they exert no [[force]]s on one another.) 
:This implies:
::1. [[Special relativity|Relativistic]] effects are negligible.
::2. [[Quantum mechanics|Quantum-mechanical]] effects are negligible. This means that the [[mean inter-particle distance|inter-particle distance]] is much larger than the [[thermal de Broglie wavelength]] and the molecules are treated as [[classical mechanics|classical]] [[physical body|objects]].
::3. Because of the above two, their dynamics can be treated classically. This means that the equations of motion of the molecules are time-reversible.

* The average kinetic energy of the gas particles depends only on the [[thermodynamic temperature|absolute temperature]] of the [[system]]. The kinetic theory has its own definition of temperature, not identical with the thermodynamic definition.
* The elapsed time of a collision between a molecule and the container's wall is negligible when compared to the time between successive collisions.
* There are negligible gravitational force on molecules.

More modern developments relax these assumptions and are based on the [[Boltzmann equation]]. These can accurately describe the properties of dense gases, because they include the volume of the molecules. The necessary assumptions are the absence of quantum effects, [[molecular chaos]] and small gradients in bulk properties. Expansions to higher orders in the density are known as [[virial expansion]]s.

An important book on kinetic theory is that by [[Sydney Chapman (mathematician)|Chapman]] and [[Thomas Cowling|Cowling]].<ref name="C & C">[[Sydney Chapman (mathematician)|Chapman, S.]], [[Cowling, T.G.]] (1939/1970).</ref> An important approach to the subject is called [[Chapman–Enskog theory]].<ref>Kauzmann, W. (1966). ''Kinetic Theory of Gasses'', W.A. Benjamin, New York, pp. 232–235.</ref> There have been many modern developments and there is an alternative approach developed by Grad based on moment expansions.<ref>Grad 1949</ref>
In the other limit, for extremely rarefied gases, the gradients in bulk properties are not small compared to the mean free paths. This is known as the Knudsen regime and expansions can be performed in the [[Knudsen number]].

==Equilibrium properties==
=== {{anchor|Pressure and Kinetic Energy}}Pressure and kinetic energy===<!-- This section is linked from [[Pressure]] -->
In kinetic model of gases, the [[pressure]] is equal to the force exerted by the atoms hitting and rebounding from a unit area of the gas container surface. Consider a gas of ''N'' molecules, each of mass ''m'', enclosed in a cube of volume ''V'' = ''L''<sup>3</sup>. When a gas molecule collides with the wall of the container perpendicular to the ''x'' axis and bounces off in the opposite direction with the same speed (an [[elastic collision]]), the change in [[momentum]] is given by:

:<math>\Delta p = p_{i,x} - p_{f,x} = p_{i,x} - (-p_{i,x}) = 2 p_{i,x} = 2 mv_x,</math>

where ''p'' is the momentum, ''i'' and ''f'' indicate initial and final momentum (before and after collision), ''x'' indicates that only the ''x'' direction is being considered, and ''v'' is the speed of the [[particle]] (which is the same before and after the collision).

The particle impacts one specific side wall once every

:<math>\Delta t = \frac{2L}{v_x},</math>

where ''L'' is the distance between opposite walls.

The [[force]] due to this particle is

:<math>F = \frac{\Delta p}{\Delta t} = \frac{m v_x^2}{L}.</math>

The total force on the wall is

:<math>F = \frac{Nm\overline{v_x^2}}{L},</math>

where the bar denotes an average over the ''N'' particles.

Since the motion of the particles is random and there is no bias applied in any direction, the average squared speed in each direction is identical:

:<math>\overline{v_x^2} = \overline{v_y^2} = \overline{v_z^2}.</math>

By [[Pythagorean theorem]] in three dimensions the total squared speed ''v'' is given by

:<math>\overline{v^2} = \overline{v_x^2} + \overline{v_y^2} + \overline{v_z^2},</math>

:<math>\overline{v^2} = 3\overline{v_x^2}.</math>

Therefore:

:<math>\overline{v_x^2} = \frac{\overline{v^2}}{3},</math>

and the force can be written as:

:<math>F = \frac{Nm\overline{v^2}}{3L}.</math>

This force is exerted on an area ''L''<sup>2</sup>. Therefore, the pressure of the gas is

:<math>P = \frac{F}{L^2} = \frac{Nm\overline{v^2}}{3V},</math>

where ''V'' = ''L''<sup>3</sup> is the volume of the box.

In terms of the kinetic energy of the gas ''K'':

:<math>PV = \frac{2}{3} \times{K}.</math>

This is a first non-trivial result of the kinetic theory because it relates pressure, a [[macroscopic]] property, to the  (translational) [[kinetic energy]] of the molecules <math>N\frac{1}{2} m\overline{v^2}</math>, which is a [[microscopic]] property.

===Temperature and kinetic energy===

Rewriting the above result for the pressure as <math>PV = {Nm\overline{v^2}\over 3} </math>, 
we may combine it with the [[ideal gas law]]

{{NumBlk|:|<math>\displaystyle PV = N k_B T ,</math>|{{EquationRef|1}}}}

where <math>\displaystyle k_B</math> is the [[Boltzmann constant]] and <math>\displaystyle T</math> the
[[Thermodynamic temperature|absolute]] [[temperature]] defined by the ideal gas law, to obtain

:<math>k_B T  =   {m\overline{v^2}\over 3} </math>,

which leads to simplified expression of the average kinetic energy per molecule,<ref>The average kinetic energy of a fluid is proportional to the [[Root-mean-square speed|root mean-square velocity]], which always exceeds the mean velocity - [http://www.mikeblaber.org/oldwine/chm1045/notes/Gases/Kinetic/Gases08.htm Kinetic Molecular Theory]</ref>
:<math>   \displaystyle     \frac {1} {2} m\overline{v^2} =  \frac {3} {2}  k_B T</math>.

The kinetic energy of the system is N times that of a molecule, namely <math> K= \frac {1} {2} N m \overline{v^2} </math>.
Then the temperature <math>\displaystyle T</math> takes the form

{{NumBlk|:|<math>   \displaystyle    T   =   {m\overline{v^2}\over 3 k_B}</math>|{{EquationRef|2}}}}

which becomes

{{NumBlk|:|<math>   \displaystyle    T   =   \frac   {2}   {3}   \frac   {K}   {N k_B}.</math>|{{EquationRef|3}}}}
Eq.({{EquationNote|3}})
is one important result of the kinetic
theory:
''The average molecular kinetic energy is proportional to the ideal gas law's absolute temperature''.
From Eq.({{EquationNote|1}}) and
Eq.({{EquationNote|3}}),
we have
{{NumBlk|:|<math>
   \displaystyle 
   PV 
   =
   \frac
   {2}
   {3}
   K.
</math>|{{EquationRef|4}}}}
Thus, the product of pressure and
volume per [[Mole (unit)|mole]] is proportional to the average
(translational) molecular kinetic energy.

Eq.({{EquationNote|1}}) and Eq.({{EquationNote|4}})
are called the "classical results",
which could also be derived from [[statistical mechanics]];
for more details, see:<ref>
[http://clesm.mae.ufl.edu/wiki.pub/index.php/Configuration_integral_%28statistical_mechanics%29 Configuration integral (statistical mechanics)] {{webarchive|url=https://web.archive.org/web/20120428193950/http://clesm.mae.ufl.edu/wiki.pub/index.php/Configuration_integral_%28statistical_mechanics%29 |date=2012-04-28 }}
</ref>

Since there are
<math>\displaystyle 3N</math>
[[Degrees of freedom (physics and chemistry)|degrees of freedom]] in a monatomic-gas system with
<math>\displaystyle N</math>
particles,
the kinetic energy per degree of freedom per molecule is
{{NumBlk|:|<math>
   \displaystyle 
   \frac
   {K}
   {3 N}
   =
   \frac
   {k_B T}
   {2}
</math>
|{{EquationRef|5}}}}
In the kinetic energy per degree of freedom,
the constant of proportionality of temperature
is 1/2 times [[Boltzmann constant]] or R/2 per mole. In addition to this, the temperature will decrease when the pressure drops to a certain point.{{why?|date=June 2014}}
This result is related
to the [[equipartition theorem]].

As noted in the article on [[heat capacity]], diatomic
gases should have 7 degrees of freedom, but the lighter diatomic gases act
as if they have only 5. Monatomic gases have 3 degrees of freedom.

Thus the kinetic energy per kelvin (monatomic [[ideal gas]]) is 3 [R/2] = 3R/2:
* per mole: 12.47 J
* per molecule: 20.7 yJ = 129 μeV.

At [[Standard conditions for temperature and pressure|standard temperature]] (273.15 K), we get:
* per mole: 3406 J
* per molecule: 5.65 zJ = 35.2 meV.

===Collisions with container===
The total number and velocity distribution of particles hitting the container wall can be calculated<ref name="OCW">{{cite web |title=5.62 Physical Chemistry II |url=https://ocw.mit.edu/courses/chemistry/5-62-physical-chemistry-ii-spring-2008/lecture-notes/29_562ln08.pdf |website=MIT OpenCourseWare}}</ref> based on kinetic theory, and the result can be used for analyzing [[Effusion|effusion into vacuum]]:

Assume that, in the container, the number density is <math>n</math> and particles obey [[Maxwell–Boltzmann distribution|Maxwell's velocity distribution]]:

<math>f_\text{Maxwell}(v_x,v_y,v_z)\,dv_x\,dv_y\,dv_z=\left(\frac{m}{2 \pi k_BT}\right)^{3/2}\,  e^{- \frac{m|v|^2}{2k_BT}} \,dv_x\,dv_y\,dv_z</math>

The particles hitting a small area <math>dA</math> on the container, with speed <math>v</math> at angle <math>\theta</math>  from the normal, in time interval <math>dt</math>, is contained in a [[parallelepiped]] with base area <math>dA</math> and height <math>v\,dt{\times}\cos(\theta)</math>, hence the total number of these particles is:

<math>dN=n v \cos{\theta} dAdt{\times}\left(\frac{m}{2 \pi k_BT}\right)^{3/2}e^{-\frac{mv^2}{2k_BT}} (v^2\sin{\theta}dv d\theta d\phi)</math> 

Note that only the particles within the following constraint are actually heading to hit the wall:

<math>v>0,\qquad 0<\theta<\pi/2,\qquad 0<\phi<2\pi.</math> 

Integrating over all appropriate velocities within the constraint yields the number of atomic or molecular collisions with a wall of a container:

<math>\begin{align}
\Delta N 
&= 
n dA dt \int_{0}^{+\infin} \left(\frac{m}{2 \pi k_BT}\right)^{3/2} v^3 e^{-\frac{mv^2}{2k_BT}} dv \int_{0}^{\frac{\pi}{2}} \cos{\theta} \sin{\theta} d\theta \int_{0}^{2 \pi} d\phi
\\
&= 
n dA dt \int_{0}^{+\infin} \left(\frac{m}{2 \pi k_BT}\right)^{3/2} v^3 e^{-\frac{mv^2}{2k_BT}} dv \times \frac{1}{2} \times 2\pi
\\
&=n dA dt \int_{0}^{+\infin} \frac{v}{4} \left(\frac{m}{2 \pi k_BT}\right)^{3/2} e^{-\frac{mv^2}{2k_BT}} 4\pi v^2 dv
\\
&=n dA dt \frac{v_\text{avg}}{4}
\end{align}</math>

where <math>v_\text{avg}</math> is the average speed from [[Maxwell–Boltzmann distribution|Maxwell's speed distribution]]. Hence the impingement rate is:

:<math>\begin{align}
J_\text{impingement} 
&= \frac{\Delta N}{dA dt} \\
&= \frac{1}{4}n v_\text{avg} = \frac{1}{4} \left(\frac{N}{V}\right) \sqrt{\frac{8 k_{B} T}{\pi m}} \\
&=\frac{P}{\sqrt{2 \pi m k_{B} T}}
\end{align}
 </math>

The last line makes use of  [[ideal gas law]]. This quantity is also known as the [[impingement rate]] in [[Vacuum technology|vacuum physics]]. 

The velocity distribution of particles hitting this small area is:

<math>\begin{align}
f(v,\theta,\phi)\,dv\,{d\theta}\,d\phi
& \propto n v \cos{\theta} dAdt{\times} e^{- \frac{mv^2}{2k_BT}}{\times}(v^2\sin{\theta}\,dv\,{d\theta}\,d\phi)
\\
& \propto v\cos{\theta}{\times}e^{- \frac{mv^2}{2k_BT}}{\times}(v^2\sin{\theta}\,dv\,{d\theta}\,d\phi)
\\
&=\text{const.}{\times}(v^3e^{- \frac{mv^2}{2k_BT}}dv){\times}(\cos{\theta}\sin{\theta}\,{d\theta}){\times}d\phi
\end{align}</math>

with the constraint <math>v>0,0<\theta<\pi/2,0<\phi<2\pi</math>, and the constant can be determined by normalization condition. 

In Cartesian coordinates, this is:

<math>\begin{align}
f(v,\theta,\phi)\,dv\,{d\theta}\,d\phi
& \propto v_ze^{- \frac{mv^2}{2k_BT}}{\times}(dv_x\,dv_y\,dv_z) \\
&=\text{const.}{\times}(e^{- \frac{mv_x^2}{2k_BT}}dv_x){\times}(e^{- \frac{mv_y^2}{2k_BT}}dv_y)
{\times}(v_ze^{- \frac{mv_z^2}{2k_BT}}dv_z)
\end{align}</math>

with the constraint <math>v_z>0</math>, and the constant can be determined by normalization condition. 

From the above distribution, the average velocity of these impinging particles is:

<math>\begin{align}
\overline{v_x}&=\overline{v_y}=0\\
\overline{v_z}&=\frac{\int_{0}^{+\infin} v_z^2 e^{- \frac{mv_z^2}{2k_BT}}dv_z}
{\int_{0}^{+\infin} v_z e^{- \frac{mv_z^2}{2k_BT}}dv_z}=\frac{<v^2>_{1d}}{<v>_{1d}}=\sqrt{\frac{\pi k_BT}{2m}}
\end{align}</math>

The last equal sign used the result from [[Maxwell–Boltzmann distribution#In n-dimensional space|Maxwell's speed distribution]] in n-dimensional space, where n=1.

When these particles bounce off the container wall, each of them transfers a momentum of <math>2 m \overline{v_z}</math>, hence the average force is:

<math>F= 2 m \overline{v_z} \times J_\text{impingement} \times A = PA</math>

===Speed of molecules===
From the kinetic energy formula it can be shown that

:<math>v_\text{p} = \sqrt{2 \cdot \frac{k_B T}{m}},</math>

:<math> \bar v = \frac {2}{\sqrt{\pi}} v_p = \sqrt{\frac {8}{\pi} \cdot \frac{k_B T}{m}},</math>

:<math>v_\text{rms} = \sqrt{\frac{3}{2}} v_p = \sqrt{{3} \cdot \frac {k_B T}{m}},</math>

where ''v'' is in m/s, ''T'' is in kelvins, and ''m'' is the mass of one molecule of gas. The most probable (or mode) speed <math>v_\text{p}</math> is 81.6% of the rms speed <math>v_\text{rms}</math>, and the mean (arithmetic mean, or average) speed <math>\bar v</math> is 92.1% of the rms speed ([[isotropy|isotropic]] [[Maxwell–Boltzmann distribution#Distribution of speeds|distribution of speeds]]).

See: 
* [[Average]], 
* [[Root-mean-square speed]]
* [[Arithmetic mean]]
* [[Mean]]
* [[Mode (statistics)]]

==Transport properties==
{{See also|Transport phenomena}}

The kinetic theory of gases deals not only with gases in thermodynamic equilibrium, but also very importantly with gases not in thermodynamic equilibrium. This means using kinetic theory to consider what are known as "transport properties", such as [[mass diffusivity]], [[viscosity]] and [[thermal conductivity]]. 
===Diffusion coefficient and diffusion flux===
{{See also|Fick's laws of diffusion}}The diffusion constant is related to viscosity by the [[Einstein relation (kinetic theory)]].

===Viscosity and kinetic momentum===
{{See also|Viscosity#Momentum transport}}

In books on elementary kinetic theory<ref name="Sears1975">{{cite book|last1=Sears|first1=F.W.|last2=Salinger|first2=G.L.|year=1975|title=Thermodynamics, Kinetic Theory, and Statistical Thermodynamics|publisher=Addison-Wesley Publishing Company, Inc.|location= Reading, Massachusetts, USA |edition=3|chapter=10|pages=286–291|isbn=978-0201068948}}</ref> one can find results for dilute gas modeling that have widespread use. 
The derivation of the kinetic model for shear viscosity usually starts by considering a [[Couette flow]], where two parallel plates are separated by a gas layer. The upper plate moves with constant velocity to the right due to a force <math>F</math>. The lower plate is stationary, and an equal and opposite force must therefore be acting on it to keep it at rest. The molecules in the gas layer have a forward velocity component <math>u</math> which increases uniformly with distance <math>y</math> above the lower plate. The non-equilibrium flow is superimposed on a [[Maxwell–Boltzmann distribution|Maxwell–Boltzmann equilibrium distribution]] of molecular motions.

Let <math> \sigma </math> be the collision [[Cross section (physics)#Collision among gas particles|cross section]] of one molecule colliding with another. The number density <math> C </math> is defined as the number of molecules per (extensive) volume <math> C = N/V </math>. The collision cross section per volume or collision cross section density is <math> C \sigma </math>, and it is related to the [[Mean free path#Mean free path in kinetic theory|mean free path]] <math>l</math> by 

<math> \quad l = \frac {1} {\sqrt{2} C \sigma} </math>

Notice that the unit of the collision cross section per volume <math> C \sigma </math> is reciprocal of length. The mean free path is the average distance traveled by a molecule, or a number of molecules per volume, before they make their first collision.

Let <math> u_{0} </math> be the forward velocity of the gas at an imaginary horizontal surface inside the gas layer. On the average, a molecule that crosses the surface makes its last collision before crossing at a distance equal to two-thirds of the mean free path (i.e. <math> 2l/3 </math>) away from the surface. At this distance above and below the surface, the forward momentum of the molecule is respectively

<math> \quad p_{x}^{\pm} = m \left( u_{0} \pm \frac {2}{3} l {d u \over dy} \right) </math>

where m is the molecular mass. The molecular flux  <math> \Phi </math> includes all molecules arriving at one side of an element of the surface within the gas layer. The incoming molecules are coming from all directions at the one side of the surface and with all speeds. This molecular flux (i.e. the number flux) is related to the average molecular speed <math> \bar v </math> by

<math> \quad \Phi = \frac {1}{4} \bar v C </math>

Notice that the forward velocity gradient <math>du/dy</math> can be considered to be constant over a distance of mean free path. Next we multiply by the total flux to get the change of momentum per unit time and per unit area, that is carried by the molecules crossing from either above or below the surface area. This gives the equation

<math> \quad \Phi p_{x}^{\pm} = \frac {1}{4} \bar v C \cdot m \left( u_{0} \pm \frac {2}{3} l {d u \over dy} \right) </math>

The net rate of momentum per unit area that is transported across the imaginary surface is thus

<math> \quad \tau = \Phi p_{x}^{+}  - \Phi p_{x}^{-} = \frac {1}{3} \bar v C m \cdot l {d u \over dy} </math>

The defining equation for the (shear) viscosity <math> \eta </math> of the gas is

<math> \quad \frac {F} {A} = \eta {d u \over dy} </math>

Combining the above kinetic equation with defining equation for (shear) viscosity by <math> \tau = F/A </math> gives the equation for shear viscosity, which is usually denoted <math> \eta_{0} </math> when it is a dilute gas:

<math> \quad \eta_{0} = \frac {1} {3} \bar v C m l </math>

Combining this equation with the equation for mean free path gives

<math> \quad \eta_{0} = \frac {1} {3 \sqrt{2} } \frac {m \cdot \bar v} {\sigma}</math>

From statistical thermodynamics for gases we have equations relating average molecular speed to most likely speed and further to temperature. These statistical results gives the average (equilibrium) molecular speed as

<math> \quad \bar v = \frac{2}{\sqrt{\pi}} v_{p} = 2 \sqrt{\frac{2}{\pi} \cdot \frac {k_{B}T}{m_{}}} </math>

where <math>v_{p}</math> is the most probable speed, <math> k_{B} </math> is the Boltzmann constant. We note that

<math> \quad k_{B} \cdot N_{A} = R \quad \text{and} \quad M = m \cdot N_{A} </math>

and insert the velocity in the viscosity equation above. This gives the well known equation for shear viscosity for dilute gases:

<math> \quad \eta_{0} = \frac {2} {3 \sqrt{\pi} } \cdot \frac {\sqrt{m k_{B} T}} { \sigma } 
= \frac {2} {3 \sqrt{\pi} } \cdot \frac {\sqrt{MRT}} { \sigma \cdot N_{A} } </math>

and <math> M </math> is the molar mass. The equation above presupposes that the gas density is low (i.e. the pressure is low). This implies that the kinetic translational energy dominates over rotational and vibrational molecule energies. The viscosity equation further presupposes that there is only one type of gas molecules, and that the gas molecules are perfect elastic and hard core particles of spherical shape. This assumption of elastic, hard core spherical molecules, like billiard balls, implies that the collision cross section of one molecule can be estimated by

<math> \quad \sigma = \pi \left( 2 r \right)^2 = \pi d^2 </math>

The radius <math>r</math> is called collision cross section radius or kinetic radius, and the diameter <math>d</math> is called collision cross section diameter or [[kinetic diameter]] of a molecule in a monomolecular gas. There are no simple general relation between the collision [[Cross section (physics)#Collision among particles|cross section]] and the hard core size of the (fairly spherical) molecule. The relation depends on shape of the potential energy of the molecule. For a real spherical molecule (i.e. a noble gas atom or a reasonably spherical molecule) the interaction potential is more like the [[Lennard-Jones potential]] or [[Morse potential]] which have a negative part that attracts the other molecule from distances longer than the hard core radius. [[Mean free path#Mean free path in kinetic theory|The radius for zero  Lennard-Jones potential]] is then appropriate to use as estimate for the kinetic radius.

Local nomenclature list:
{{div col|colwidth=30em}}
*<math>A \ \ \ </math> : area of moving boundary in Couette flow experiment [m<sup>2</sup>]
*<math>C \ \ \ </math> : [[concentration|number concentration]] or number density [1/m<sup>3</sup>]
*<math>c \ \ \ \, </math> : [[concentration|molar concentration]] or molar density [mol/cm<sup>3</sup>]
*<math>d \ \ \ \, </math> : kinetic diameter in collision cross section [m]
*<math>F \ \ \ </math> : force that move a boundary in Couette flow experiment [N]
*<math>k_{B} \ \, </math> : Boltzmann constant [JK<sup>−1</sup>]
*<math>l \ \ \ \ \ </math> : mean free path [m]
*<math>M \ \ </math> : [[molar mass]] [g/mol]
*<math>m \ \ \ </math> : molecular mass [Da]
*<math>N_{A} \, </math> : Avogadro constant [mol<sup>−1</sup>]
*<math>P \ \ \ </math> : pressure [Pa or bar or atm]
*<math>P_c \ \ </math> : critical pressure [Pa or bar or atm]
*<math>p_{x} \ \ </math> : linear momentum in x-direction of a molecule [kgm/s]
*<math>R \ \ \ </math> : gas constant [JK<sup>−1</sup>mol<sup>−1</sup>]
*<math>r \ \ \ \, \, </math> : kinetic radius in collision cross section or hard core molecular radius [m]
*<math>T \ \ \, \, </math> : temperature [K]
*<math>T_c \ \, \, </math> : critical temperature [K]
*<math>u \ \ \ \ </math> : macroscopic fluid velocity in x-direction [m/s]
*<math>u_{0} \ \ </math> : macroscopic fluid velocity in x-direction on the imaginary surface [m/s]
*<math>\bar v \ \ \ \ </math> : average molecular equilibrium speed [m/s]
*<math>v_p \ \ \, </math> : most probable molecular equilibrium speed [m/s]
*<math>V \ \ \ </math> : molar volume [cm<sup>3</sup>/mol]
*<math>V_c \ \ </math> : critical molar volume [cm<sup>3</sup>/mol]
*<math>V_{ext}</math> : extensive fluid volume [m<sup>3</sup>]
*<math>y \ \ \ \ </math> : distance from non-moving boundary in direction normal to fluid flow [m]
*<math>\eta \ \ \ \ </math> : viscosity [Pas or μP or cP]
*<math>\eta_{0} \ \, \, </math> : viscosity of dilute gas [Pas or μP or cP]
*<math>\Phi \ \ \, \, </math> : molecular flux across an imaginary surface [m<sup>−2</sup>s<sup>−1</sup>]
*<math>\sigma \ \ \ \ </math> : collision cross section [m<sup>2</sup>]
*<math>\tau \ \ \ \ </math> : shear stress [Nm<sup>−2</sup>]
*<math>K_{r} \ </math> : dummy
{{div col end}}

===Thermal conductivity and heat flux===
{{See also|Thermal conductivity}}

==See also==
{{Statistical mechanics}}
* [[BBGKY hierarchy|Bogoliubov–Born–Green–Kirkwood–Yvon hierarchy of equations]]
* [[Boltzmann equation]]
* [[Collision theory]]
* [[Critical temperature]]
* [[Gas laws]]
* [[Heat]]
* [[Interatomic potential]]
* [[Magnetohydrodynamics]]
* [[Maxwell–Boltzmann distribution]]
* [[Mixmaster dynamics]]
* [[Thermodynamics]]
* [[Vlasov equation]]

==Notes==
{{Reflist|2}}

==References==
* {{Citation
| author=Clausius, R.
| title =Ueber die Art der Bewegung, welche wir Wärme nennen
| journal =Annalen der Physik
| volume =176
| pages =353–379
| year =1857
|url=http://gallica.bnf.fr/ark:/12148/bpt6k15185v/f371.table
|doi=10.1002/andp.18571760302|bibcode = 1857AnP...176..353C
| issue=3 }}

* de Groot, S. R., W. A. van Leeuwen and Ch. G. van Weert (1980), Relativistic Kinetic Theory, North-Holland, Amsterdam.
* {{Citation
| author=Einstein, A.
| title =Über die von der molekularkinetischen Theorie der Wärme geforderte Bewegung von in ruhenden Flüssigkeiten suspendierten Teilchen
| journal =Annalen der Physik
| volume =17
| pages =549–560
| year=1905
|url=http://www.cdvandt.org/Band%2017%20Kap%205.pdf
| doi=10.1002/andp.19053220806|bibcode = 1905AnP...322..549E
| issue=8 }}

* {{Citation
| author=Grad, Harold
| title =On the Kinetic Theory of Rarefied Gases.
| journal =Communications on Pure and Applied Mathematics
| volume =2
| pages =331–407
| year =1949
| doi =10.1002/cpa.3160020403
| issue=4 }}

* {{Citation
| first = J.
| last = Herapath
| authorlink= John Herapath
| title =On the physical properties of gases
| journal =[[Annals of Philosophy]]
| year =1816
| pages= 56–60
| url =https://books.google.com/books?id=dBkAAAAAMAAJ&pg=PA56
| publisher= Robert Baldwin}}

* {{Citation
| author=Herapath, J.
| year= 1821
| title=On the Causes, Laws and Phenomena of Heat, Gases, Gravitation
| journal= Annals of Philosophy
| volume =9
| pages =273–293
| url=https://books.google.com/books?id=nCsAAAAAMAAJ&pg=RA1-PA273
| publisher=Baldwin, Cradock, and Joy }}

* {{Citation
| author=Krönig, A.
| title =Grundzüge einer Theorie der Gase
| journal =Annalen der Physik
| volume =99
| pages =315–322
| year =1856
|url=http://gallica.bnf.fr/ark:/12148/bpt6k15184h/f327.table
|doi=10.1002/andp.18561751008|bibcode = 1856AnP...175..315K
| issue=10 }}

* {{Citation
| author=Le Sage, G.-L.
| year=1818
| chapter=Physique Mécanique des Georges-Louis Le Sage
| editor=Prévost, Pierre
| title=Deux Traites de Physique Mécanique
| place=Geneva & Paris
| publisher=J.J. Paschoud
| pages=1–186
| chapter-url=http://resolver.sub.uni-goettingen.de/purl?PPN521099943}}

* Liboff, R. L. (1990), Kinetic Theory, Prentice-Hall, Englewood Cliffs, N. J.
* {{Citation
| author=Lomonosov, M.
| origyear=1758
| year= 1970
| chapter=On the Relation of the Amount of Material and Weight
| editor= Henry M. Leicester
| title= Mikhail Vasil'evich Lomonosov on the Corpuscular Theory
| place = Cambridge
| publisher=Harvard University Press
| pages =224–233
|chapterurl=https://archive.org/details/mikhailvasilevic017733mbp}}

* {{Citation|author=Mahon, Basil
|title=The Man Who Changed Everything – the Life of James Clerk Maxwell
|place=Hoboken, New Jersey
|publisher=Wiley
|year=2003
|isbn= 978-0-470-86171-4}}

* {{Citation|author=Maxwell, James Clerk |title=Molecules |journal=Nature |volume=417 |year=1873 |doi=10.1038/417903a |url=http://www.thecore.nus.edu.sg/landow/victorian/science/science_texts/molecules.html |pages=903 |format=– <sup>[https://scholar.google.co.uk/scholar?hl=en&lr=&q=intitle%3AMolecules&as_publication=Nature&as_ylo=1873&as_yhi=1873&btnG=Search Scholar search]</sup> |pmid=12087385 |issue=6892 |bibcode=2002Natur.417..903M |url-status=dead |archiveurl=https://web.archive.org/web/20070209001003/http://www.thecore.nus.edu.sg/landow/victorian/science/science_texts/molecules.html |archivedate=February 9, 2007 }}
* {{Citation
| author=Smoluchowski, M.
| title =Zur kinetischen Theorie der Brownschen Molekularbewegung und der Suspensionen
| journal =Annalen der Physik
| volume =21
| pages =756–780
| year=1906
|url=http://gallica.bnf.fr/ark:/12148/bpt6k15328k/f770.chemindefer
| doi=10.1002/andp.19063261405|bibcode = 1906AnP...326..756V
| issue=14 }}

* {{Citation
| author = Waterston, John James
| year = 1843
| title = Thoughts on the Mental Functions }} (reprinted in his ''Papers'', '''3''', 167, 183.)

* {{cite book|author=Williams, M. M. R.|year=1971|title=Mathematical Methods in Particle Transport Theory|url=https://archive.org/details/mathematicalmeth0000will|url-access=registration|location=Butterworths, London|isbn=9780408700696}}

==Further reading==
*[[Sydney Chapman (mathematician)|Sydney Chapman]] and T. G. Cowling (1939/1970).  ''The Mathematical Theory of Non-uniform Gases: An Account of the Kinetic Theory of Viscosity, Thermal Conduction and Diffusion in Gases'', (first edition 1939, second edition 1952), third edition 1970 prepared in co-operation with D. Burnett, Cambridge University Press, London.
*J. O. Hirschfelder, C. F. Curtiss, and R. B. Bird (1964).  ''Molecular Theory of Gases and Liquids'', second edition (Wiley).
*R. L. Liboff (2003).  ''Kinetic Theory: Classical, Quantum, and Relativistic Descriptions'', third edition (Springer).
*B. Rahimi and H. Struchtrup, ''[https://dx.doi.org/10.1017/jfm.2016.604 Macroscopic and kinetic modelling of rarefied polyatomic gases]'', Journal of Fluid Mechanics, 806, 437–505, 2016. DOI: https://dx.doi.org/10.1017/jfm.2016.604

== External links ==
* [http://www.math.umd.edu/~lvrmr/History/EarlyTheories.html Early Theories of Gases]
* [http://www.lightandmatter.com/html_books/0sn/ch05/ch05.html Thermodynamics] - a chapter from an online textbook
* [http://physnet.org/modules/pdfmodules/m156.pdf ''Temperature and Pressure of an Ideal Gas: The Equation of State''] on [http://www.physnet.org Project PHYSNET].
* [https://web.archive.org/web/20141103020057/http://www2.ucdsb.on.ca/tiss/stretton/CHEM1/gases9.html Introduction] to the kinetic molecular theory of gases, from The Upper Canada District School Board
* [http://comp.uark.edu/~jgeabana/mol_dyn/ Java animation] illustrating the kinetic theory from University of Arkansas
* [http://hyperphysics.phy-astr.gsu.edu/hbase/kinetic/ktcon.html Flowchart] linking together kinetic theory concepts, from HyperPhysics
* [https://web.archive.org/web/20050409211009/http://www.ewellcastle.co.uk/science/pages/kinetics.html Interactive Java Applets] allowing high school students to experiment and discover how various factors affect rates of chemical reactions.
* https://www.youtube.com/watch?v=47bF13o8pb8&list=UUXrJjdDeqLgGjJbP1sMnH8A A demonstration apparatus for the thermal agitation in gases.
{{Authority control}}

{{DEFAULTSORT:Kinetic Theory of Gasses}}
[[Category:Gases]]
[[Category:Thermodynamics]]
[[Category:Concepts in physics]]