Name: gas particle motion, diffusion, and effusion
Uploaded: 2020-09-03T13:38:27.000Z
Duration: 3 min 48 s
Channel: JoVE

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NC K-12 Science Standard Course of Study

9-12

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PS.Chm.7.3

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

Een gesloten flesje parfum bevat een hoge concentratie gasvormige aromatische moleculen die constant in beweging zijn en willekeurig botsen. Ondertussen bevat de lucht buiten de fles in wezen geen van deze moleculen. Bij het openen van de fles wordt een concentratiegradiënt ingesteld tussen deze hoge en lage concentratie-gebieden.De moleculen blijven willekeurig bewegen, met algemene beweging van het gebied met hoge concentratie naar het gebied met lage concentratie. Het spontaan mengen en verspreiden van vloeistoffen of gassen in reactie op een concentratiegradiënt wordt moleculaire diffusie genoemd. Diffusie is een langzaam proces.Hoewel gasdeeltjes zich met hoge snelheden verplaatsen, veroorzaken de talrijke botsingen frequente veranderingen in snelheid en richting. De gemiddelde afstand die een deeltje aflegt tussen botsingen staat bekend als het gemiddelde vrije weglengte. Voor een gasdeeltje wordt zijn gemiddelde vrije weglengte beïnvloed door de deeltjesdichtheid, die ook de druk beïnvloedt.Naarmate de deeltjesdichtheid toeneemt, neemt ook de botsingsfrequentie toe. Hun gemiddelde vrije weglengte is dus korter. Evenzo, naarmate de deeltjesdichtheid afneemt, neemt ook de botsingsfrequentie toe, wat leidt tot een langer gemiddeld vrije weglengte.Verschillende gassen verspreiden zich met verschillende snelheden, afhankelijk van de snelheid van de gasdeeltjes. Omdat kwadratisch gemiddelde, of RMS, snelheid en de molaire massa van een gas omgekeerd evenredig zijn, verspreiden lichtere gassen zich sneller dan zwaardere gassen. Denk aan een glazen buis tussen reservoirs met gelijke hoeveelheden ammoniak en waterstofchloridegas.Wanneer de diffunderende gassen elkaar ontmoeten, reageren ze om een ring van ammoniumchloride te vormen. De ring bevindt zich dichter bij het waterstofchloride-uiteinde van de buis omdat de lichtere ammoniakmoleculen in dezelfde tijd verder door de buis gingen dan de zwaardere waterstofchloridemoleculen. Effusie is een ander proces waarbij gasmoleculen worden verplaatst.Het is het vermogen van gasmoleculen om door een gat te verplaatsen waarvan de diameter veel kleiner is dan het gemiddelde vrije pad van het gas zelf als reactie op een drukverschil. Dit is de reden waarom heliumballonnen uiteindelijk leeglopen het helium stroomt geleidelijk door kleine poriën in het ballonmateriaal. Net als diffusie is de effusiesnelheid afhankelijk van de RMS snelheid en de molaire massa van het gas.In het bijzonder is de uitstroomsnelheid omgekeerd evenredig met de wortel van de molaire massa van het gas. Daarom spreiden zwaardere gassen langzamer uit dan lichtere. Voor twee gassen is de verhouding van hun effusiesnelheden de wortel van de omgekeerde ratio van hun molaire massa.Dit wordt Graham's wet van effusie genoemd. Denk aan twee ballonnen die tot dezelfde druk zijn opgeblazen de ene gevuld met helium en de andere met zuurstof. Helium heeft een lagere molaire massa dan zuurstof, zoals blijkt uit de opwaartse kracht van de heliumballon in lucht.Het toepassen van de wet van Graham op helium en zuurstof suggereert dat helium 2, 8 keer sneller uitstroomt dan zuurstof. De heliumballon loopt dus sneller leeg dan de zuurstofballon.

gedrag van gasmoleculen: moleculaire diffusie, gemiddeld vrij pad en effusie

Hoewel gasvormige moleculen zich met enorme snelheden (honderden meter per seconde) voortbewegen, botsen ze met andere gasvormige moleculen en reizen ze in veel verschillende richtingen voordat ze het gewenste doel bereiken. Bij kamertemperatuur zal een gasvormig molecuul miljarden botsingen per seconde ervaren. Het gemiddelde vrije pad is de gemiddelde afstand die een molecuul aflegt tussen botsingen. Het gemiddelde vrije pad neemt toe met afnemende druk; in het algemeen zal het gemiddelde vrije pad voor een gasvormig molecuul honderden keren de diameter van het molecuul zijn Wanneer een gasmonster in een deel van een gesloten container wordt gebracht, verspreiden de moleculen ervan zich over het algemeen zeer snel door de container; dit proces waarbij moleculen zich in de ruimte verspreiden als reactie op concentratieverschillen wordt diffusie genoemd. De gasvormige atomen of moleculen zijn zich uiteraard niet bewust van enige concentratiegradiënt; ze verplaatsen zich eenvoudigweg willekeurig: regio's met een hogere concentratie hebben meer deeltjes dan regio's met een lagere concentratie, en dus vindt er netto een verplaatsing van soorten plaats van gebieden met een hoge naar een lage concentratie. In een gesloten omgeving zal diffusie uiteindelijk resulteren in overal gelijke gasconcentraties. De gasvormige atomen en moleculen blijven bewegen, maar aangezien hun concentraties in beide bollen hetzelfde zijn, zijn de overdrachtssnelheden tussen de bollen gelijk (er vindt geen netto overdracht van moleculen plaats). De hoeveelheid gas die per tijdseenheid door een bepaald gebied stroomt, is de diffusiesnelheid. <img alt="Eq1" src="https://www.jove.com/files/core/CDNSource/lessons/1044/1044_Equations-01.png" style="height: 40px;"> De diffusiesnelheid is afhankelijk van verschillende factoren: de concentratiegradiënt (de toename of afname van de concentratie van het ene punt naar het andere), de hoeveelheid oppervlak die beschikbaar is voor diffusie en de afstand die de gasdeeltjes moeten afleggen. Een proces dat de beweging van gasvormige soorten omvat, vergelijkbaar met diffusie, is effusie, het ontsnappen van gasmoleculen door een klein gaatje, zoals een gaatje in een ballon, naar een vacuüm. Hoewel diffusie- en effusiesnelheden beide afhankelijk zijn van de molaire massa van het betrokken gas, zijn hun snelheden niet gelijk; de verhoudingen van hun snelheden zijn echter hetzelfde. Als een mengsel van gassen in een container met poreuze wanden wordt geplaatst, stromen de gassen door de kleine openingen in de wanden. De lichtere gassen passeren sneller (met een hogere snelheid) door de kleine openingen dan de zwaardere. In 1832 bestudeerde Thomas Graham de effusiesnelheden van verschillende gassen en formuleerde de effusiewet van Graham: De effusiesnelheid van een gas is omgekeerd evenredig met de vierkantswortel van de massa van zijn deeltjes: <img alt="Eq2" src="https://www.jove.com/files/core/CDNSource/lessons/1044/1044_Equations-02.png" style="height: 40px;"> Dit betekent dat als twee gassen, A en B, dezelfde temperatuur en druk hebben, de verhouding van hun uitstromingssnelheden omgekeerd evenredig is met de verhouding van de vierkantswortels van de massa van hun deeltjes: <img alt="Eq3" src="https://www.jove.com/files/core/CDNSource/lessons/1044/1044_Equations-03.png" style="height: 40px;"> De relatie geeft aan dat het lichtere gas een hogere effusiesnelheid heeft. Een met helium gevulde rubberen ballon loopt bijvoorbeeld sneller leeg dan een met lucht gevulde ballon, omdat de effusiesnelheid door de poriën van het rubber sneller is voor de lichtere heliumatomen dan voor de luchtmoleculen. Deze tekst is aangepast van <a href="https://openstax.org/books/chemistry-2e/pages/9-4-effusion-and-diffusion-of-gases">Openstax, Chemistry 2e, Sectie 9.4: Effusie en diffusie van gassen.</a>

Gedrag van gasmoleculen: moleculaire diffusie, gemiddeld vrij pad en effusie

Although gaseous molecules travel at tremendous speeds (hundreds of meters per second), they collide with other gaseous molecules and travel in many ...

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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 of a gas molecule is hundreds of times greater than the molecule’s diameter.

Diffusion is the spread of gas particles through space because of differences in concentration. When a gas is added to one part of a closed container, its molecules quickly disperse throughout the container. The molecules move randomly, but the net movement is from higher concentration to lower concentration. In a closed environment, diffusion eventually leads to equal concentrations throughout. The gas particles still move, but there is no net transfer between regions when the concentrations are the same.

The rate of diffusion is the amount of gas that passes through an area in a given time. It depends on the concentration gradient, the available surface area, and the distance the particles must travel. A related process is effusion, which is the escape of gas molecules through a tiny opening, such as a pinhole in a balloon into a vacuum. Diffusion and effusion rates both depend on molar mass, but the rates are not the same.

If a mixture of gases is placed in a container with porous walls, the gases effuse through the small openings. The lighter gases pass through more quickly than the heavier gases. Thomas Graham studied these rates in 1832 and formed Graham’s law of effusion. This law states that the rate of effusion is inversely proportional to the square root of the mass of the particles. For two gases at the same temperature and pressure, the ratio of their effusion rates follows the same relationship. A helium-filled rubber balloon, for example, deflates faster than an air-filled balloon because helium atoms effuse faster through the rubber pores than air molecules.

Gas Particle Motion, Diffusion, and Effusion

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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...

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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...

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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...

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Gas Pressure in Mixtures: Partial Pressures

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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.

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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...

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Gas Pressure Patterns in Kinetic Molecular Theory

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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

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...

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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...

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Explore how gas molecules move, diffuse, and effuse, plus why lighter gases pass through tiny openings faster than heavier gases.

Gas Particle Motion, Diffusion, and Effusion | JoVE K12

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PS.Chm.7.2

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

Gas Particle Motion, Diffusion, and Effusion

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