Name: kinetic molecular theory and gas behavior
Uploaded: 2020-09-03T13:38:26.000Z
Duration: 2 min 43 s
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Basic Postulates of Kinetic Molecular Theory: Particle Size, Energy, and Collision

Terwijl de gaswetten de relaties tussen verschillende eigenschappen van ideale gassen samenvatten, verklaart de kinetische moleculaire theorie waarom gassen de wetten volgen. De theorie is gebaseerd op een aantal aannames of postulaten. De eerste aanname is dat gasdeeltjes verwaarloosbaar klein zijn.Een gas is meestal lege ruimte die bestaat uit kleine deeltjes die worden gescheiden op afstanden die veel groter zijn dan hun eigen afmetingen. Hun gecombineerde volume is verwaarloosbaar in verhouding tot het totale volume waarin het gas zich bevindt. In tegenstelling tot vaste stoffen en vloeistoffen, die onsamendrukbaar zijn vanwege hun nauwe onderlinge afstand tussen de deeltjes, zijn gassen zeer samendrukbaar.Gasdeeltjes zijn in een constante staat van beweging langs rechte lijnen in willekeurige richtingen. Hun pad verandert alleen als ze in botsing komen met andere deeltjes of met de wanden van hun container. De tweede veronderstelling is dat gasdeeltjes perfect elastische botsingen hebben.Ze botsen en stuiteren op elkaar zonder aan elkaar te kleven. Dit is te vergelijken met de botsingen tussen biljartballen tijdens een potje biljart. Wanneer gasdeeltjes botsen, wisselen ze energie met elkaar uit, maar er is geen netto energieverlies.Met andere woorden, de totale energie van het systeem blijft constant. Gasdeeltjes zijn constant in beweging;daarom bezitten ze kinetische energie. De derde aanname stelt dus dat de gemiddelde kinetische energie van een gas evenredig is met de absolute temperatuur in kelvin.Dit betekent dat kinetische energie toeneemt met de temperatuur, en bijgevolg bewegen de deeltjes sneller. Bij hogere temperaturen neemt hun snelheid toe. Omgekeerd, als de temperatuur afneemt, neemt ook de kinetische energie van de deeltjes toe, en bewegen ze langzamer.Bij een bepaalde temperatuur hebben alle gassen, ongeacht hun molecuulmassa, dezelfde gemiddelde kinetische energie. Kinetische energie is gelijk aan massa maal snelheid in het kwadraat. Om verschillende gassen dezelfde gemiddelde kinetische energie te laten hebben, moeten hun gasdeeltjes dus met verschillende gemiddelde snelheden voortbewegen.Daarom hebben zwaardere gassen lagere gemiddelde snelheden, terwijl lichtere gassen hogere gemiddelde snelheden hebben. Helium en neon hebben bijvoorbeeld dezelfde gemiddelde kinetische energie als ze dezelfde temperatuur hebben. Door het verschil in massa bewegen de neonatomen echter veel langzamer dan de heliumatomen.

basispostulaten van de kinetische moleculaire theorie: deeltjesgrootte, energie en botsing

De ideaalgasvergelijking, die empirisch is, beschrijft het gedrag van gassen door relaties tussen hun macroscopische eigenschappen vast te stellen. De wet van Charles stelt bijvoorbeeld dat volume en temperatuur rechtstreeks verband houden. Gassen zetten daarom uit als ze bij constante druk worden verwarmd. Hoewel de gaswetten verklaren hoe de macroscopische eigenschappen ten opzichte van elkaar veranderen, verklaart zij niet de grondgedachte erachter. De kinetische moleculaire theorie is een microscopisch model dat helpt begrijpen wat er met gasdeeltjes op moleculair of atomair niveau gebeurt wanneer omstandigheden zoals druk of temperatuur veranderen. In 1857 publiceerde Rudolf Clausius een volledige en bevredigende vorm van de theorie, die de verschillende gaswetten effectief verklaart door middel van de postulaten die werden ontwikkeld op basis van honderden experimentele observaties van het gedrag van gassen. De meest opvallende kenmerken van deze theorie zijn: <ol> <li>Gassen zijn samengesteld uit deeltjes (atomen of moleculen) die voortdurend in beweging zijn, zich in rechte lijnen voortbewegen en alleen van richting veranderen wanneer ze botsen met andere moleculen of met de wanden van een container. Onderzoek een monster argongas bij standaardtemperatuur en -druk. Hieruit blijkt dat slechts 0,01% van het volume wordt ingenomen door atomen met een gemiddelde afstand van 3,3 nm (de atoomstraal van argon is 0,097 nm) tussen twee argonatomen. De afstand is veel groter dan zijn eigen dimensie.</li> <li>De moleculen waaruit het gas bestaat, zijn verwaarloosbaar klein vergeleken met de afstanden ertussen. Daarom is het gecombineerde volume van alle gasdeeltjes verwaarloosbaar in verhouding tot het totale volume van de container. De deeltjes worden beschouwd als “punten” met massa maar een verwaarloosbaar volume.</li> <li>De druk die door een gas in een container wordt uitgeoefend, is het gevolg van botsingen tussen de gasmoleculen en de containerwanden.</li> <li>Gasmoleculen oefenen geen aantrekkende of afstotende krachten uit op elkaar of op de containerwanden; daarom zijn hun botsingen elastisch (er gaat geen energie verloren). Bij elastische botsingen wordt energie overgedragen tussen de botsende deeltjes. De gemiddelde kinetische energie van de deeltjes blijft daarom constant en verandert niet in de tijd.</li> <li>De gemiddelde kinetische energie van de gasmoleculen is evenredig met de Kelvin-temperatuur van het gas. Alle gassen, ongeacht hun moleculaire massa, hebben dezelfde gemiddelde kinetische energie bij dezelfde temperatuur.</li> </ol> Deze tekst is aangepast van <a href="https://openstax.org/books/chemistry-2e/pages/9-5-the-kinetic-molecular-theory">Openstax, Chemistry 2e, Hoofdstuk 9.5 De Kinetisch-Moleculaire Theorie</a>.

Basispostulaten van de kinetische moleculaire theorie: deeltjesgrootte, energie en botsing

The ideal-gas equation, which is empirical, describes the behavior of gases by establishing relationships between their macroscopic properties. For ...

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kinetic molecular theory and gas behavior

Kinetic molecular theory explains gas behavior at the particle level. It gives a microscopic model for what happens when pressure or temperature changes. Gas laws describe how macroscopic properties change, but they do not explain why those changes happen.

The theory was put into a complete form by Rudolf Clausius in 1857. It was built from hundreds of experimental observations about gases. Those observations helped explain the different gas laws through a few basic postulates.

One key idea is that gases are made of particles, either atoms or molecules, that move constantly. They travel in straight lines until they collide with other particles or with the walls of a container. A sample of argon gas at standard temperature and pressure shows that the particles take up only a tiny part of the total volume. The average distance between two argon atoms is much larger than the size of the atoms themselves. Because of this, gas particles are treated as points with mass but negligible volume.

Gas pressure comes from collisions between gas molecules and the walls of the container. The particles do not attract or repel each other, and they do not interact with the container walls through such forces. For that reason, their collisions are elastic, which means no energy is lost in the collision. Energy can still be transferred between colliding particles, but the average kinetic energy stays constant over time.

The average kinetic energy of gas particles depends on the kelvin temperature of the gas. It is proportional to temperature. At the same temperature, gases with different molecular masses have the same average kinetic energy.

Kinetic Molecular Theory and Gas Behavior

kinetic-molecular-theory-and-gas-behavior

definition and measurement of pressure: atmospheric pressure, barometer, and manometer

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Pressure Units and Tools for Gas Measurement

Pressure is the force applied to a given area, and gas pressure comes from gas molecules colliding with surfaces. Each collision is tiny, but many collisions over a small area can create high pressure. Pressure increases when force increases or when the area gets smaller. It decreases when force drops or when the area gets larger. Pressure is measured in pascals, or Pa, in the SI system. One pascal equals 1 newton per square meter, and a newton is a unit of force equal to 1 kg·m/s2. A pascal...

gas laws: boyle's, gay-lussac, charles', avogadro's, and ideal gas law

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How Gas Pressure, Volume, and Temperature Change

Gas pressure, volume, temperature, and amount are linked by several gas laws. Scientists found these relationships by studying how a confined gas changes when one variable is held constant and another changes. These patterns apply to an ideal gas, which is a useful model for gases under low pressure and high temperature. Gay-Lussac’s law, also called Amontons’s law, connects pressure and temperature. If gas is trapped in a rigid container, the volume and number of moles stay constant. When the...

applications of the ideal gas law: molar mass, density, and volume

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Ideal Gas Law at STP and Gas Density

The ideal gas law links pressure, volume, number of moles, and temperature. It helps explain how gas volume changes when pressure or temperature changes. At standard temperature and pressure, or STP, the values are 273.15 K and 1 atm. At STP, one mole of an ideal gas occupies about 22.4 L. This is called the standard molar volume. The gas can be hydrogen, oxygen, argon, or carbon dioxide, and the volume is still about 22.4 liters. Half a mole occupies 11.2 L, and 2 moles occupy 44.8 L. The...

mixtures of gases: dalton's law of partial pressures and mole fractions

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Gas Pressure in Mixtures: Partial Pressures

Gas pressure in mixtures is described by partial pressure and mole fraction. When gases do not chemically react, each gas in the mixture acts as if it were alone in the container. The pressure from one gas is called its partial pressure. For a mixture of three gases, A, B, and C, the total pressure equals the sum of their partial pressures: PA, PB, and PC. This is Dalton's law of partial pressures. It states that the total pressure of a mixture of ideal gases is the sum of the pressures from...

chemical stoichiometry and gases: using ideal gas law to determine moles

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Stoichiometry of Gases with the Ideal Gas Law

Chemical stoichiometry links the amounts of reactants and products in a chemical reaction. For gases, those amounts can also be measured by volume. If a gas’s volume, pressure, and temperature are known, the ideal gas equation can be used to find the number of moles present. The same idea works in reverse. If the moles of a gas are known, its volume can be found at any temperature and pressure. The transcript gives an example using 8.88 g of gallium reacting with excess hydrochloric acid.

basic postulates of kinetic molecular theory: particle size, energy, and collision

Kinetic molecular theory explains gas behavior at the particle level. It gives a microscopic model for what happens when pressure or temperature changes. Gas laws describe how macroscopic properties change, but they do not explain why those changes happen. The theory was put into a complete form by Rudolf Clausius in 1857. It was built from hundreds of experimental observations about gases. Those observations helped explain the different gas laws through a few basic postulates. One key idea...

kinetic molecular theory and gas laws explain properties of gas molecules

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Gas Pressure Patterns in Kinetic Molecular Theory

Kinetic molecular theory explains gas pressure by linking it to the motion of gas molecules. Gas pressure comes from rapidly moving molecules hitting the walls of a container. It depends on how many molecules strike a unit area in a given time. The theory also helps explain the main gas laws. In Gay-Lussac’s law, also called Amontons’s law, raising temperature increases the average speed and kinetic energy of gas molecules. If volume stays constant, the molecules collide with the walls more...

kinetic molecular theory: molecular velocities, temperature, and kinetic energy

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Gas Particle Speed Distributions and Temperature

Gas particle speed distributions and temperature are key ideas in the kinetic molecular theory. This theory explains the gas laws by describing how gas molecules move, collide, and share energy. Gas molecules do not all move at the same speed. A gas sample has an average speed and an average kinetic energy, but individual molecules move at different speeds. Their collisions are elastic, which means momentum is conserved. Because so many molecules collide, the overall speed distribution stays...

behavior of gas molecules: molecular diffusion, mean free path, and effusion

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Gas Particle Motion, Diffusion, and Effusion

Gas particle motion, diffusion, and effusion explain how gases spread and escape in everyday settings. Gas molecules move at very high speeds, but they collide with other gas molecules and change direction many times before reaching a target. At room temperature, a gas molecule can experience billions of collisions each second. The mean free path is the average distance a molecule travels between collisions. This distance becomes larger when pressure decreases. In general, the mean free path...

real gases: effects of intermolecular forces and molecular volume deriving van der waals equation

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Real Gases: When the Ideal Gas Law Breaks Down

Real gases are gases that do not always follow the ideal gas law, PV = nRT, especially when pressure is high or temperature is low. The ideal gas model assumes gas particles have no attractive or repulsive forces and take up negligible space compared with the container. At room temperature and 1 atm or less, many gases behave close to ideally. When conditions change, the measured relationships between pressure, volume, and temperature can drift away from the ideal prediction. One way to see...

key terms and concepts

Explore how kinetic molecular theory explains gas pressure, particle motion, elastic collisions, and why average kinetic energy depends on kelvin temperature.

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Comparing Intermolecular Forces: Melting Point, Boiling Point, and Miscibility

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Phase Transitions: Vaporization and Condensation

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Clausius-Clapeyron Equation

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Using Differential Scanning Calorimetry to Measure Changes in Enthalpy

Source: Laboratory of Dr. Terry Tritt &#8212; Clemson University
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Hess's Law

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Standard Enthalpy of Formation

Enthalpy changes are typically tabulated for reactions in which both the reactants and products are at the same conditions. A standard state is a commonly ...

Enthalpies of Reaction

Hess&rsquo;s law can be used to determine the enthalpy change of any reaction if the corresponding enthalpies of formation of the reactants and products ...

Use mathematics and computational thinking to execute simple calorimetric calculations based on the Law of Conservation of Energy.

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Behavior of Gas Molecules: Molecular Diffusion, Mean Free Path, and Effusion

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Real Gases: Effects of Intermolecular Forces and Molecular Volume Deriving Van der Waals Equation

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Ideal Gas Law

Source: Laboratory of Dr. Andreas Z&#252;ttel - Swiss Federal Laboratories for Materials Science and Technology
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	Determining the Mass of the Unknown LiquidExpand
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Use mathematics and computational thinking to explain the relationships among pressure, temperature, volume, and quantity of gas, both qualitatively and quantitatively.

Understand the relationship among pressure, temperature, volume, and phase.

Kinetic Molecular Theory and Gas Behavior

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