Kinetic theory 64204 217760763 2008-06-07T15:20:33Z D.H 1523162 /* History */ Correcting link '''Kinetic theory''' (or '''kinetic theory of gases''') attempts to explain [[macroscopic]] properties of [[gas]]es, such as pressure, temperature, or volume, by considering their [[molecule|molecular]] composition and [[motion (physics)|motion]]. Essentially, the theory posits that pressure is due not to static repulsion between molecules, as was [[Isaac Newton]]'s conjecture, but due to [[collision]]s between molecules moving at different velocities. Kinetic theory is also known as the '''kinetic-molecular theory''' or the '''[[collision theory]]'''. ==History== 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), <ref>{{Citation | author=Lomonosow, M. | author-link =Mikhail Lomonosov| year= 1970 | title=[http://www.archive.org/details/mikhailvasilevic017733mbp On the Relation of the Amount of Material and Weight (1758)] | Herausgeber= Henry M. Leicester | journal= Mikhail Vasil'evich Lomonosov on the Corpuscular Theory | place = Cambridge | publisher=Harvard University Press | pages =224-233 }}</ref> [[Georges-Louis Le Sage]] (ca. 1780, published 1818), <ref>{{Citation | author=Le Sage, G.-L. | author-link=Georges-Louis Le Sage | year=1818 | chapter=Physique Mécanique des Georges-Louis Le Sage | editor=Prévost, Pierre | editor-link=Pierre Prévost | title=Deux Traites de Physique Mécanique | place=Geneva & Paris | publisher=J.J. Paschoud | pages=1-186 | chapter-url=http://dz1.gdz-cms.de/index.php?id=img&no_cache=1&IDDOC=304083}}</ref> [[John Herapath]] (1816) <ref>{{Citation | author= Herapath, J. | title =On the physical properties of gases | journal =Annals of Philosophy | year =1816 | pages= 56-60| url =http://books.google.com/books?id=dBkAAAAAMAAJ&pg=PA56}}<br /> {{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=http://books.google.com/books?id=nCsAAAAAMAAJ&pg=RA1-PA273 }}</ref> and [[John James Waterston]] (1843), <ref>{{cite book | last = Waterston | first = JJ | authorlink = John James Waterston | year = 1843 | title = Thoughts on the Mental Functions }} (reprinted in his ''Papers'', '''3''', 167, 183.)</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>{{Citation | author=Krönig, A. | title =[http://gallica.bnf.fr/ark:/12148/bpt6k15184h/f327.table Grundzüge einer Theorie der Gase] | journal =Annalen der Physik | volume =99 | pages =315-322 | year =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>{{Citation | author=Clausius, R. | title =[http://gallica.bnf.fr/ark:/12148/bpt6k15185v/f371.table Über die Art der Bewegung, die wir Wärme nennen] | journal =Annalen der Physik | volume =100 | pages =353-379 | year =1857 }}</ref> 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.<ref>{{cite book|author=Mahon, Basil |title=The Man Who Changed Everything – the Life of James Clerk Maxwell|location=Hoboken, NJ | publisher=Wiley|year=2003|id=ISBN 0-470-86171-1}}</ref> In his 1875 thirteen page article 'Molecules', published in the September issue of ''Nature'', 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, James Clerk, "[http://www.thecore.nus.edu.sg/landow/victorian/science/science_texts/molecules.html Molecules]". ''Nature'', September, 1873.</ref> 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) <ref>{{Citation | author=Einstein, A. | title =[http://www3.interscience.wiley.com/homepages/5006612/549_560.pdf 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}}</ref> and [[Marian Smoluchowski]]'s (1906) <ref>{{Citation | author=Smoluchowski, M. | title =[http://gallica.bnf.fr/ark:/12148/bpt6k15328k/f770.chemindefer Zur kinetischen Theorie der Brownschen Molekularbewegung und der Suspensionen] | journal =Annalen der Physik | volume =21 | pages =756-780 | year=1906}}</ref> papers on [[Brownian motion]], which succeeded in making certain accurate quantitative predictions based on the kinetic theory. == Postulates == 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, [[randomness|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 [[interaction]]s among molecules are [[negligible]]. They exert no [[force]]s 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 [[dimension|size]]. * The molecules are perfectly spherical in shape, and elastic in nature. * The average [[kinetic energy]] of the gas particles depends only on the [[thermodynamic temperature|temperature]] of the [[system]]. * [[Special relativity|Relativistic]] effects are negligible. * [[Quantum mechanics|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 mechanics|classical]] [[physical body|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}} <!-- Please give the full titles and other relevant information about these books in the form of a reference citation. --> 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 ==<!-- This section is linked from [[Pressure]] --> [[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: :<math>\Delta p_x = p_i - p_f = 2 m v_x\,</math> where ''v<sub>x</sub>'' is the ''x''-component of the initial velocity of the particle. The particle impacts the wall once every 2''l/v<sub>x</sub>'' time units (where ''l'' is the length of the container). Although the particle impacts a side wall once every 1''l/v<sub>x</sub>'' 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/v<sub>x</sub>'' time units. :<math>\Delta t = \frac{2l}{v_x}</math> The [[force]] due to this particle is: :<math>F = \frac{\Delta p}{\Delta t} = \frac{2 m v_x}{\frac{2l}{v_x}} = \frac{m v_x^2}{l}</math> The total force acting on the wall is: :<math>F = \frac{m\sum_j v_{jx}^2}{l}</math> where the summation is over all the gas molecules in the container. The magnitude of the velocity for each particle will follow: :<math> v^2 = v_x^2 + v_y^2 + v_z^2 </math> Now considering the total force acting on all six walls, adding the contributions from each direction we have: :<math>\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}</math> 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: :<math>F = \frac{1}{6} \left(2 \cdot \frac{m \sum_j v_{j}^2}{l}\right) = \frac{m \sum_j v_{j}^2}{3l}</math> The quantity <math>\sum_j v_{j}^2</math> can be written as <math>{N} \overline{v^2}</math>, where the bar denotes an average, in this case an average over all particles. This quantity is also denoted by <math>v_{rms}^2</math> where <math>v_{rms}</math> is the [[root mean square|root-mean-square]] velocity of the collection of particles. Thus the force can be written as: :<math>F = \frac{Nmv_{rms}^2}{3l}</math> Pressure, which is force per unit area, of the gas can then be written as: :<math>P = \frac{F}{A} = \frac{Nmv_{rms}^2}{3Al}</math> 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 :<math>P = {Nmv_{rms}^2 \over 3V} </math> where ''V'' is the volume. Also, as ''Nm'' is the total mass of the gas, and mass divided by volume is density :<math> P = {1 \over 3} \rho\ v_{rms}^2</math> 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''mv<sub>rms</sub>''<sup>2</sup>), 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]], :{| style="width:100%" border="0" |- | style="width:95%" | <math> \displaystyle PV = N k_B T </math> | style= | (1) |} where <math>\displaystyle k_B</math> is the [[Boltzmann constant]], and <math>\displaystyle T</math> the [[Thermodynamic temperature|absolute]] [[temperature]], it follows from the above result that the temperature <math>\displaystyle T</math> takes the form :{| style="width:100%" border="0" |- | style="width:95%" | <math> \displaystyle PV = N k_B T = \frac {N m v_{rms}^2} {3} \Longrightarrow T = \frac {m v_{rms}^2} {3 k_B} </math> | style= | (2) |} and the kinetic energy <math>\displaystyle K</math> of the system can now be written as :{| style="width:100%" border="0" |- | style="width:95%" | <math> \displaystyle K = \frac {1} {2} N m v_{rms}^2 = \frac {3} {2} N k_B T \ {\rm and} \ T = \frac {2} {3} \frac {K} {N k_B} </math> | style= | (3) |} Eq.(3)<sub>1</sub> is one important result of the kinetic theory: <i>The average molecular kinetic energy is proportional to the absolute temperature</i>. From Eq.(1) and Eq.(3)<sub>1</sub>, we have :{| style="width:100%" border="0" |- | style="width:95%" | <math> \displaystyle PV = \frac {2} {3} K </math> | style= | (4) |} Thus, the product of pressure and volume per [[Mole (unit)|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 <ref> [http://clesm.mae.ufl.edu/wiki.pub/index.php/Configuration_integral_%28statistical_mechanics%29 Configuration integral (statistical mechanics)] </ref>. Since there are <math>\displaystyle 3N</math> [[Degrees of freedom (physics and chemistry)|degrees of freedom]] (dofs) in a monoatomic-gas system with <math>\displaystyle N</math> particles, the kinetic energy per dof is :{| style="width:100%" border="0" |- | style="width:95%" | <math> \displaystyle \frac {K} {3 N} = \frac {k_B T} {2} </math> | style= | (5) |} In the kinetic energy per dof, the constant of proportionality of temperature is 1/2 times [[Boltzmann constant]]. 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 gases act as 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 conditions for temperature and pressure|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<ref>[http://www.chem.arizona.edu/~salzmanr/480a/480ants/collsurf/collsurf.html Collisions With a Surface<!-- Bot generated title -->]</ref> results in an equation for total number of collisions per unit time per area: ::<math>A = \frac{1}{4}\frac{N}{V} v_{avg} = \frac{\rho}{4} \sqrt{\frac{8 k T}{\pi m}} \frac{1}{m} \,</math> == RMS speeds of molecules == From the kinetic energy formula it can be shown that :<math>v_{rms}^2 = \frac{3RT}{\mbox{molar mass}}</math> 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% ([[Maxwell-Boltzmann distribution#Distribution of speeds|distribution of speeds]]). == See also == * [[Gas laws]] * [[Heat]] * [[Maxwell-Boltzmann distribution]] * [[Thermodynamics]] * [[Collision theory]] * [[Critical temperature]] ==References== {{reflist}} The Mathematical Theory of Non-uniform Gases : An Account of the Kinetic Theory of Viscosity, Thermal Conduction and Diffusion in Gases Sydney Chapman, T. G. Cowling == 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]. * [http://www.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 * [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. * [http://www.bustertests.co.uk/answer/molecular-kinetic-theory/ Molecular kinetic theory fundamentals] [[Category:Gases]] [[Category:Thermodynamics]] [[Category:Fundamental physics concepts]] [[ca:Teoria cinètica molecular]] [[cs:Kinetická teorie látek]] [[de:Kinetische Gastheorie]] [[el:Κινητική θεωρία]] [[es:Teoría cinética]] [[fa:تئوری سینتیک مولکولی]] [[fr:Théorie cinétique des gaz]] [[gl:Teoría Cinética]] [[it:Teoria cinetica dei gas]] [[he:התאוריה הקינטית של הגזים]] [[lv:Molekulāri kinētiskā teorija]] [[nl:Kinetische gastheorie]] [[ja:気体分子運動論]] [[nn:Kinetisk teori]] [[ru:Молекулярно-кинетическая теория]] [[simple:Kinetic theory]] [[sl:Kinetična teorija plinov]] [[fi:Kineettinen kaasuteoria]] [[sv:Kinetiska gasteorin]] [[th:ทฤษฎีจลน์ของแก๊ส]] [[tr:Kinetik teori]] [[zh:分子运动论]]