Kinetic theory

Kinetic theory

temperature of an ideal monatomic gas is a measure related to the average kinetic energy of its atoms as they move. In this animation, the size of helium atoms relative to their spacing is shown to scale under 1950 atmospheres of pressure. These room-temperature atoms have a certain, average speed (slowed down here two trillion fold).] Kinetic theory (or kinetic theory of gases) attempts to explain macroscopic properties of gases, such as pressure, temperature, or volume, by considering their molecular composition and motion. Essentially, the theory posits that pressure is due not to static repulsion between molecules, as was Isaac Newton's conjecture, but due to collisions between molecules moving at different velocities. Kinetic theory is also known as the kinetic-molecular theory or the collision theory.


The theory for ideal gases makes the following assumptions:
* The gas consists of very small particles, each of which has a mass or weight in SI units, kilograms.
* The number of molecules is large such that statistical treatment can be applied.
* These molecules are in constant, random motion. The rapidly moving particles constantly collide with each other and with the walls of the container.
* The collisions of gas particles with the walls of the container holding them are perfectly elastic.
* The interactions among molecules are negligible. They exert no forces on one another except during collisions.
* The total volume of the individual gas molecules added up is negligible compared to the volume of the container. This is equivalent to stating that the average distance separating the gas particles is relatively large compared to their size.
* The molecules are perfectly spherical in shape, and elastic in nature.
* The average kinetic energy of the gas particles depends only on the temperature of the system.
* Relativistic effects are negligible.
* Quantum-mechanical effects are negligible. This means that the inter-particle distance is much larger than the thermal de Broglie wavelength and the molecules can be treated as classical objects.
* The time during collision of molecule with the container's wall is negligible as comparable to the time between successive collisions.
* The equations of motion of the molecules are time-reversible.In addition, if the gas is in a container, the collisions with the walls are assumed to be instantaneous and elastic.

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 expansions. The definitive work is the book by Chapman and Enskog but there have been many modern developments and there is an alternative approach developed by Grad based on moment expansions.Fact|date=November 2007 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.

The kinetic theory has also been extended to include inelastic collisions in granular matter by Jenkins and others.Fact|date=November 2007


Pressure is explained by kinetic theory as arising from the force exerted by gas molecules impacting on the walls of the container. Consider a gas of "N" molecules, each of mass "m", enclosed in a cuboidal container of volume "V". When a gas molecule collides with the wall of the container perpendicular to the "x" coordinate axis and bounces off in the opposite direction with the same speed (an elastic collision), then the momentum lost by the particle and gained by the wall is:

:Delta p_x = p_i - p_f = 2 m v_x,

where "vx" is the "x"-component of the initial velocity of the particle.

The particle impacts the wall once every 2"l/vx" time units (where "l" is the length of the container). Although the particle impacts a side wall once every 1"l/vx" time units, only the momentum change on one wall is considered so that the particle produces a momentum change on a particular wall once every 2"l/vx" time units.:Delta t = frac{2l}{v_x}

The force due to this particle is:

:F = frac{Delta p}{Delta t} = frac{2 m v_x}{frac{2l}{v_x = frac{m v_x^2}{l}

The total force acting on the wall is:

:F = frac{msum_j v_{jx}^2}{l}

where the summation is over all the gas molecules in the container.

The magnitude of the velocity for each particle will follow:

: v^2 = v_x^2 + v_y^2 + v_z^2

Now considering the total force acting on all six walls, adding the contributions from each direction we have:

:mbox{Total Force} = 2 cdot frac{m}{l}(sum_j v_{jx}^2 + sum_j v_{jy}^2 + sum_j v_{jz}^2) = 2 cdot frac{m}{l} sum_j (v_{jx}^2 + v_{jy}^2 + v_{jz}^2) = 2 cdot frac{m sum_j v_{j}^2}{l}

where the factor of two arises from now considering both walls in a given direction.

Assuming there are a large number of particles moving sufficiently randomly, the force on each of the walls will be approximately the same and now considering the force on only one wall we have:

:F = frac{1}{6} left(2 cdot frac{m sum_j v_{j}^2}{l} ight) = frac{m sum_j v_{j}^2}{3l}

The quantity sum_j v_{j}^2 can be written as {N} overline{v^2}, where the bar denotes an average, in this case an average over all particles. This quantity is also denoted by v_{rms}^2 where v_{rms} is the root-mean-square velocity of the collection of particles.

Thus the force can be written as:

:F = frac{Nmv_{rms}^2}{3l}

Pressure, which is force per unit area, of the gas can then be written as:

:P = frac{F}{A} = frac{Nmv_{rms}^2}{3Al}

where "A" is the area of the wall of which the force exerted on is considered.

Thus, as cross-sectional area multiplied by length is equal to volume, we have the following expression for the pressure

:P = {Nmv_{rms}^2 over 3V}

where "V" is the volume. Also, as "Nm" is the total mass of the gas, and mass divided by volume is density

: P = {1 over 3} ho v_{rms}^2

where ρ is the density of the gas.

This result is interesting and significant, because it relates pressure, a macroscopic property, to the average (translational) kinetic energy per molecule (1/2"mvrms"2), which is a microscopic property. Note that the product of pressure and volume is simply two thirds of the total kinetic energy.

Temperature and kinetic energy

From the ideal gas law,:Eq.(3)1is one important result of the kinetictheory: The average molecular kinetic energy is proportional tothe absolute temperature.

From Eq.(1) and Eq.(3)1,we have:Thus, the product of pressure andvolume per mole is proportional to the average(translational) molecular kinetic energy.

Eq.(1) and Eq.(4)are called the "classical results", which could also be derived from statistical mechanics; for more details, see [ [ Configuration integral (statistical mechanics)] ] .

Since there are displaystyle 3N
degrees of freedom(dofs)in a monoatomic-gas system with displaystyle Nparticles,the kinetic energy per dof is:In the kinetic energy per dof,the constant of proportionality of temperature is 1/2 times
Boltzmann constant. This result is relatedto the equipartition theorem.

As noted in the article on heat capacity, diatomicgases should have 7 degrees of freedom, but the lighter gases actas if they have only 5.

Thus the kinetic energy per kelvin (monatomic ideal gas) is:
* per mole: 12.47 J
* per molecule: 20.7 yJ = 129 μeV

At standard temperature (273.15 K), we get:
* per mole: 3406 J
* per molecule: 5.65 zJ = 35.2 meV

Number of collisions with wall

One can calculate the number of atomic or molecular collisions with a wall of a container per unit area per unit time.

Assuming an ideal gas, a derivation [ [ Collisions With a Surface ] ] results in an equation for total number of collisions per unit time per area:

::A = frac{1}{4}frac{N}{V} v_{avg} = frac{ ho}{4} sqrt{frac{8 k T}{pi m frac{1}{m} ,

RMS speeds of molecules

From the kinetic energy formula it can be shown that

:v_{rms}^2 = frac{3RT}{mbox{molar mass

with "v" in m/s, "T" in kelvins, and "R" is the gas constant. The molar mass is given as kg/mol. The most probable speed is 81.6% of the rms speed, and the mean speeds 92.1% (distribution of speeds).


In 1740 Daniel Bernoulli published "Hydrodynamica", which laid the basis for the kinetic theory of gases. In this work, Bernoulli positioned 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. 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.

Other pioneers of the kinetic theory (which were neglected by their contemporaries) were Mikhail Lomonosov (1747), [Lomovosov 1758] Georges-Louis Le Sage (ca. 1780, published 1818), [Le Sage 1780/1818] John Herapath (1816) [Herapath 1816, 1821] and John James Waterston (1843), [Waterston 1843] 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. [Krönig 1856]

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. [Clausius 1857] In 1859, after reading a paper by Clausius, James Clerk Maxwell formulated the Maxwell distribution of molecular velocities, which gave the proportion of molecules having a certain velocity in a specific range. This was the first-ever statistical law in physics. [Mahon 2003] 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.” [Maxwell 1875] In 1871, Ludwig Boltzmann generalized Maxwell's achievement and formulated the Maxwell–Boltzmann distribution. Also the logarithmic connection between entropy and probability was first stated by him.

In the beginning of 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) [Einstein 1905] and Marian Smoluchowski's (1906) [Smoluchowski 1906] papers on Brownian motion, which succeeded in making certain accurate quantitative predictions based on the kinetic theory.

See also

* Gas laws
* Heat
* Maxwell-Boltzmann distribution
* Thermodynamics
* Collision theory
* Critical temperature


author=Clausius, R.
title =Ueber die Art der Bewegung, welche wir Wärme nennen
journal =Annalen der Physik
volume =100
pages =353–379
year =1857

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

author= Herapath, J.
title =On the physical properties of gases
journal =Annals of Philosophy
year =1816
pages= 56–60
url =

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

author=Krönig, A.
title =Grundzüge einer Theorie der Gase
journal =Annalen der Physik
volume =99
pages =315–322
year =1856

author=Le Sage, G.-L.
chapter=Physique Mécanique des Georges-Louis Le Sage
editor=Prévost, Pierre
title=Deux Traites de Physique Mécanique
place=Geneva & Paris
publisher=J.J. Paschoud

author=Lomonosow, M.
year= 1758/1970
chapter=On the Relation of the Amount of Material and Weight
editor= Henry M. Leicester
journal= Mikhail Vasil'evich Lomonosov on the Corpuscular Theory
place = Cambridge
publisher=Harvard University Press
pages =224–233

*Citation|author=Mahon, Basil
title=The Man Who Changed Everything – the Life of James Clerk Maxwell
place=Hoboken, NJ
isbn= 0-470-86171-1

author=Maxwell, James Clerk
title =Molecules
journal =Nature
volume =417

author=Smoluchowski, M.
title =Zur kinetischen Theorie der Brownschen Molekularbewegung und der Suspensionen
journal =Annalen der Physik
volume =21
pages =756–780

author = Waterston, John James
year = 1843
title = Thoughts on the Mental Functions
(reprinted in his "Papers", 3, 167, 183.)


The Mathematical Theory of Non-uniform Gases : An Account of the Kinetic Theory of Viscosity, Thermal Conduction and Diffusion in GasesSydney Chapman, T. G. Cowling

External links

* [ Early Theories of Gases]
* [ Thermodynamics] - a chapter from an online textbook
* [ "Temperature and Pressure of an Ideal Gas: The Equation of State"] on [ Project PHYSNET] .
* [ Introduction] to the kinetic molecular theory of gases, from The Upper Canada District School Board
* [ Java animation] illustrating the kinetic theory from University of Arkansas
* [ Flowchart] linking together kinetic theory concepts, from HyperPhysics
* [ Interactive Java Applets] allowing high school students to experiment and discover how various factors affect rates of chemical reactions.
* [ Molecular kinetic theory fundamentals]

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