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

K12 Standards

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NC K-12 Science Standard Course of Study

9-12

Chemistry

Energy

PS.Chm.7.3

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

Une bouteille de parfum fermée contient une forte concentration de molécules aromatiques gazeuses qui se déplacent constamment et entrent en collision au hasard. Pendant ce temps, l'air à l'extérieur de la bouteille contient essentiellement aucune de ces molécules. Lors de l'ouverture de la bouteille, un gradient de concentration est établi entre ces régions de concentration élevée et faible.Les molécules continuent de se déplacer de façon aléatoire, avec un mouvement global de la région à forte concentration vers la région à faible concentration. Le mélange spontané et la propagation de liquides ou de gaz en réponse à un gradient de concentration est appelée diffusion moléculaire. La diffusion est un processus lent.Même si les particules de gaz se déplacent à grande vitesse, les nombreuses collisions provoquent des changements fréquents de vitesse et de direction. La distance moyenne parcourue par une particule entre les collisions est connue sous le nom de chemin libre moyen. Pour une particule de gaz, son trajet libre moyen est influencé par la densité des particules, ce qui affecte également la pression.Plus la densité des particules augmente, plus la fréquence des collisions augmente. Ainsi, leur chemin libre moyen est plus court. De même, à mesure que la densité des particules diminue, la fréquence de collision diminue, ce qui conduit à un trajet libre moyen plus long.Différents gaz diffusent à des vitesses différentes, en fonction de la vitesse des particules de gaz. Étant donné que la vitesse moyenne quadratique, ou RMS, et la masse molaire d'un gaz sont inversement liées, les gaz plus légers diffusent plus rapidement que les gaz plus lourds. Considérons un tube de verre entre les réservoirs de quantités égales d'ammoniac et de chlorure d'hydrogène gazeux.Lorsque les gaz diffusants se rencontrent, ils réagissent pour former un anneau de chlorure d'ammonium. L'anneau est plus proche de l'extrémité chlorure d'hydrogène du tube parce que les molécules d'ammoniac plus légères ont voyagé plus loin dans le tube que les molécules de chlorure d'hydrogène plus lourdes dans le même laps de temps. L'épanchement est un autre processus qui implique le mouvement des molécules de gaz.C'est la capacité des molécules de gaz à traverser un trou dont le diamètre est beaucoup plus petit que la voie libre moyenne du gaz lui-même en réponse à une différence de pression. C'est pourquoi les ballons d'hélium finissent par se dégonfler-l'hélium efface progressivement à travers de minuscules pores dans le matériau du ballon. Comme la diffusion, le taux d'épanchement dépend de la vitesse RMS et de la masse molaire du gaz.Plus précisément, le taux d'épanchement est inversement proportionnel à la racine carrée de la masse molaire du gaz. Par conséquent, les gaz plus lourds s'écoulent plus lentement que les gaz plus légers. Pour deux gaz quelconques, le rapport de leurs taux d'épanchement est la racine carrée du rapport inverse de leurs masses molaires.C'est ce qu'on appelle la loi d'épanchement de Graham. Considérons deux ballons gonflés à la même pression-l'un rempli d'hélium et l'autre d'oxygène. L'hélium a une masse molaire inférieure à l'oxygène, comme le montre la flottabilité du ballon d'hélium dans l'air.L'application de la loi de Graham à l'hélium et à l'oxygène suggère que l'hélium s'efface 2, 8 fois plus vite que l'oxygène. Ainsi, le ballon d'hélium dégonfle plus rapidement que le ballon d'oxygène.

effusion et diffusion

Bien que les molécules de gaz se déplacent à des vitesses énormes (des centaines de mètres par seconde), elles entrent en collision avec d'autres molécules de gaz et se déplacent dans de nombreuses directions différentes avant d'atteindre la cible souhaitée. À température ambiante, une molécule gazeuse subit des milliards de collisions par seconde. Le libre parcourt moyen est la distance moyenne qu'une molécule parcourt entre les collisions. Le libre parcourt moyen est d'autant plus grand que la pression diminue ; en général, le libre parcourt moyen pour une molécule gazeuse est des centaines de fois le diamètre de la molécule 
 
 En général, lorsqu'un échantillon de gaz est introduit dans une partie d'un récipient clos, ses molécules se dispersent très rapidement dans tout le récipient ; ce processus par lequel les molécules se dispersent dans l'espace en réponse aux différences de concentration est appelé diffusion. Les atomes ou molécules de gaz n'ont évidemment pas conscience d'un gradient de concentration ; elles se déplacent simplement au hasard — les zones avec des concentration plus élevées ont plus de particules que les zones avec des concentration plus faibles, et donc un mouvement net d'espèces des zones à forte concentration vers les zones à faible concentration a lieu. Dans un environnement fermé, la diffusion se traduira en fin de compte par des concentrations égales de gaz de part et d'autres. Les atomes et molécules de gaz continuent de se déplacer, mais comme leurs concentrations sont les mêmes dans les deux bulbes, les vitesses d'échange entre les bulbes sont égales (aucun échange net de molécules ne se produit). La quantité de gaz passant à travers une zone par unité de temps est la vitesse de diffusion. 
 
 <img alt="Eq1" src="https://www.jovecore.com/CDNSource/lessons/1044/1044_Equations-01.png" style="height: 40px;" />
 
 La vitesse de diffusion dépend de plusieurs facteurs : le gradient de concentration (l'augmentation ou la diminution de la concentration d'un point à l'autre), la quantité de surface disponible pour la diffusion et la distance que les particules de gaz doivent parcourir. 
 
 Un processus impliquant le mouvement d'espèces gazeuses semblables à la diffusion est l'effusion, des molécules de gaz s'échappant par un trou minuscule, tel qu'un trou d'épingle dans un ballon sous vide. Bien que les vitesses de diffusion et d'effusion dépendent toutes deux de la masse molaire du gaz impliqué, leurs vitesses ne sont pas égales ; cependant, les rapports de leurs vitesses sont les mêmes. 
 
 Si un mélange de gaz est placé dans un récipient avec des parois poreuses, les gaz s'effusent par les petites ouvertures dans les parois. Les gaz plus légers passent à travers les petits orifices plus rapidement (à une vitesse plus élevée) que les gaz plus lourds. En 1832, Thomas Graham a étudié les vitesses d'effusion de différents gaz et formulé la loi de l'effusion de Graham : la vitesse d’effusion d’un gaz est inversement proportionnelle à la racine carrée de la masse de ses particules : 
 
 <img alt="Eq2" src="https://www.jovecore.com/CDNSource/lessons/1044/1044_Equations-02.png" style="height: 40px;" />
 
 Cela signifie que si deux gaz, A et B, sont à la même température et à la même pression, le rapport de leurs vitesses d'effusion est inversement proportionnel au rapport des racines carrées des masses de leurs particules : 
 
 <img alt="Eq3" src="https://www.jovecore.com/CDNSource/lessons/1044/1044_Equations-03.png" style="height: 40px;" />
 
 La relation indique que le gaz plus léger a une vitesse d'effusion plus élevée. 
 
 Par exemple, un ballon en caoutchouc rempli d'hélium se dégonfle plus rapidement qu'un ballon rempli d'air parce que la vitesse de l'effusion à travers les pores du caoutchouc est plus rapide pour les atomes d'hélium plus légers que pour les molécules d'air. 
 
 Ce texte est adapté de <a href="https://openstax.org/books/chemistry-2e/pages/9-4-effusion-and-diffusion-of-gases">Openstax, Chimie 2e, Section 9.4 : Effusion et diffusion des gaz.</a>

Effusion et diffusion

Bien que les molécules gazeuses se déplacent à des vitesses énormes (des centaines de mètres par seconde), elles entrent en collision avec d’autres molécules gazeuses et se déplacent dans de nombreux ...

Education

Enseignement

JoVE Core

Chimie

Gases

KEY TERMS AND CONCEPTS

TERMES ET CONCEPTS CLÉS

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

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

definition-measurement-pressure-atmospheric-pressure-barometer

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

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

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

applications-of-the-ideal-gas-law-molar-mass-density-and-volume

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

mixtures-of-gases-dalton-s-law-of-partial-pressures-and-mole-fractions

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

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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...

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

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 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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Ohio’s Learning Standards and Model Curriculum

Interactions Of Matter

Gas laws

Virginia  Standards of Learning

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CH.6

CH.6.b

Texas Essential Knowledge and Skills

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CHEM.10

CHEM.10.A

Alabama Course of Study Science

Matter and Its Interactions

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9-12.C.4.7

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Florida Department of Education

Physical Science

SC.912.P.10

SC.912.P.10.5

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New York State Science Learning Standards

HS. Structure and Properties of Matter

HS-PS1-9.

Classifying Matter by State

Chemistry is the study of matter and the changes it undergoes. Matter is anything that has mass and occupies space. Matter is all around us; the air, ...

Classifying Matter by Composition

Matter: Pure Substances and Mixtures
According to its composition, the matter can be classified into two broad categories &mdash; pure substances and ...

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Physical and Chemical Properties of Matter

The characteristics that enable us to distinguish one substance from another are called properties.

Physical Properties of Matter

A physical property is ...

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Molecular Comparison of Gases, Liquids, and Solids

Particles in a solid are tightly packed together (fixed shape) and often arranged in a regular pattern; in a liquid, they are close together with no ...

Intermolecular vs Intramolecular Forces

Intermolecular forces (IMF) are electrostatic attractions arising from charge-charge interactions between molecules. The strength of the intermolecular ...

Intermolecular Forces

Atoms and molecules interact through bonds (or forces): intramolecular and intermolecular. The forces are electrostatic as they arise from interactions ...

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

Intermolecular forces are attractive forces that exist between molecules. They dictate several bulk properties, such as melting points, boiling points, ...

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Surface Tension, Capillary Action, and Viscosity

Surface Tension
The various IMFs between identical molecules of a substance are examples of cohesive forces. The molecules within a liquid are surrounded ...

States of Matter and Phase Changes

The internal energy of a substance&#8212;the total kinetic energy of all its molecules and the potential energy of their associated forces&#8212;depends ...

Phase Transitions

Whether solid, liquid, or gas, a substance&#39;s state depends on the order and arrangement of its particles (atoms, molecules, or ions). Particles in the ...

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Phase Transitions: Vaporization and Condensation

The physical form of a substance changes on changing its temperature. For example, raising the temperature of a liquid causes the liquid to vaporize ...

Vapor Pressure

When a liquid vaporizes in a closed container, gas molecules cannot escape. As these gas phase molecules move randomly about, they will occasionally ...

Clausius-Clapeyron Equation

The equilibrium between a liquid and its vapor depends on the temperature of the system; a rise in temperature causes a corresponding rise in the vapor ...

Phase Transitions: Melting and Freezing

Heating a crystalline solid increases the average energy of its atoms, molecules, or ions, and the solid gets hotter. At some point, the added energy ...

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Phase Transitions: Sublimation and Deposition

Some solids can transition directly into the gaseous state, bypassing the liquid state, via a process known as sublimation. At room temperature and ...

Heating and Cooling Curves

When a substance&#8212;isolated from its environment&#8212;is subjected to heat changes, corresponding changes in temperature and phase of the substance ...

Phase Diagrams

A phase diagram combines plots of pressure versus temperature for the liquid-gas, solid-liquid, and solid-gas phase-transition equilibria of a substance. ...

Use models to explain how changes in energy affect the arrangement and movement of the particles in solids, liquids, and gases, as well as the relative strengths of their intermolecular forces.

PS.Chm.7.1

Energy Basics

Chemical reactions, such as those that occur when you light a match, involve changes in energy as well as matter.

Chemical changes and their accompanying ...

First Law of Thermodynamics

Energy Conservation

Energy can be converted from one form into another, but all of the energy present before a change occurs always exists in some form ...

Internal Energy

The total of all possible kinds of energy present in a substance is called the internal energy (U), sometimes symbolized as E. Suppose a system with ...

Quantifying Heat

Thermal Energy&#160;

Microscopically, thermal energy is the kinetic energy associated with the random motion of atoms and molecules. Temperature is a ...

Quantifying Work

As a system undergoes a change, its internal energy can change, and energy can be transferred from the system to the surroundings, or from the ...

Enthalpy

Chemists ordinarily use a property known as enthalpy (H) to describe the thermodynamics of chemical and physical processes. Enthalpy is defined as the sum ...

Thermochemical Equations

For a chemical reaction (the system) carried out at constant pressure &ndash; with the only work done caused by expansion or contraction &ndash; the ...

Constant Pressure Calorimetry

Calorimetry is a technique used to measure the amount of heat involved in a chemical or physical process or to measure the heat transferred to or from a ...

Constant Volume Calorimetry

Calorimeters are useful to determine the heat released or absorbed by a chemical reaction. Coffee cup calorimeters are designed to operate at constant ...

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Using Differential Scanning Calorimetry to Measure Changes in Enthalpy

Source: Laboratory of Dr. Terry Tritt &#8212; Clemson University
Differential Scanning Calorimetry (DSC) is a method of thermodynamic analysis based on ...

Hess's Law

There are two ways to determine the amount of heat involved in a chemical change: measure it experimentally, or calculate it from other experimentally ...

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.

PS.Chm.7.2

Definition and Measurement of Pressure: Atmospheric Pressure, Barometer, and Manometer

Gas pressure is caused by force exerted by gas molecules colliding with the surfaces of objects. Although the force of each collision is very small, any ...

Gas Laws: Boyle's, Gay-Lussac, Charles', Avogadro's, and Ideal Gas Law

Through experiments, scientists established the mathematical relationships between pairs of variables, such as pressure and temperature, pressure and ...

Applications of the Ideal Gas Law: Molar Mass, Density, and Volume

The volume occupied by one mole of a substance is its molar volume. The ideal gas law, PV = nRT, &nbsp;suggests that the volume of a given quantity of gas ...

Mixtures of Gases: Dalton's Law of Partial Pressures and Mole Fractions

Unless individual gases chemically react with each other, the individual gases in a mixture of gases do not affect each other&rsquo;s pressure. Each gas ...

Chemical Stoichiometry and Gases: Using Ideal Gas Law to Determine Moles

Chemical stoichiometry describes the quantitative relationships between reactants and products in chemical reactions.

In addition to measuring quantities ...

Basic Postulates of Kinetic Molecular Theory: Particle Size, Energy, and Collision

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

Kinetic Molecular Theory and Gas Laws Explain Properties of Gas Molecules

The test of the kinetic molecular theory (KMT) and its postulates is its ability to explain and describe the behavior of a gas. The various gas laws ...

Kinetic Molecular Theory: Molecular Velocities, Temperature, and Kinetic Energy

The kinetic molecular theory qualitatively explains the behaviors described by the various gas laws. The postulates of this theory may be applied in a ...

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

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

Thus far, the ideal gas law, PV = nRT, has been applied to a variety of different types of problems, ranging from reaction stoichiometry and empirical and ...

Ideal Gas Law

Source: Laboratory of Dr. Andreas Z&#252;ttel - Swiss Federal Laboratories for Materials Science and Technology
The ideal gas law describes the behavior ...

Derivation of the Ideal Gas Law
Gases are a fundamental state of matter. A gas is a collection of molecules that have a significant distance between their ...

Source: Smaa Koraym at Johns Hopkins University, MD, USA

	Preparation of the LaboratoryExpand
	Here, we show the laboratory preparation for 10 students ...

Source: Smaa Koraym at Johns Hopkins University, MD, USA

	Determining the Mass of the Unknown LiquidExpand
	According to the ideal gas law for a given ...

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.

Gas Particle Motion, Diffusion, and Effusion

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