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

Zamknięta butelka perfum zawiera wysokie stężenie gazowych cząsteczek aromatycznych, które stale się poruszają i losowo kolidują. Tymczasem powietrze na zewnątrz butelki nie zawiera w zasadzie żadnej z tych cząsteczek.

Po otwarciu butelki ustala się gradient stężeń między tymi obszarami o wysokim i niskim stężeniu. Cząsteczki nadal poruszają się losowo, z ogólnym ruchem z obszaru o wysokim stężeniu do obszaru o niskim stężeniu. 

Spontaniczne mieszanie i rozprzestrzenianie się cieczy lub gazów w odpowiedzi na gradient stężeń nazywa się dyfuzją molekularną.

Dyfuzja jest procesem powolnym. Mimo że cząsteczki gazu poruszają się z dużą prędkością, liczne zderzenia powodują częste zmiany prędkości i kierunku. 

Średnia odległość, jaką cząstka pokonuje między zderzeniami, jest znana jako jej średnia droga swobodna. W przypadku cząstki gazu na jej średnią swobodną drogę ma wpływ gęstość cząstek, która również wpływa na ciśnienie. 

Wraz ze wzrostem gęstości cząstek wzrasta częstotliwość zderzeń. W związku z tym ich średnia swobodna ścieżka jest krótsza. Podobnie, wraz ze spadkiem gęstości cząstek, zmniejsza się również częstotliwość zderzeń, co prowadzi do wydłużenia średniej ścieżki swobodnej. 

Różne gazy dyfundują z różną szybkością, w zależności od prędkości cząstek gazu. Ponieważ średnia kwadratowa (RMS) prędkość i masa molowa gazu są odwrotnie powiązane, lżejsze gazy dyfundują szybciej niż cięższe gazy.

Rozważ szklaną rurkę między zbiornikami równych ilości amoniaku i gazowego chlorowodoru. 

Kiedy dyfundujące gazy spotykają się, reagują, tworząc pierścień chlorku amonu. Pierścień znajduje się bliżej końca rurki z chlorowodorem, ponieważ lżejsze cząsteczki amoniaku przemieszczały się dalej w dół rurki niż cięższe cząsteczki chlorowodoru w tym samym czasie.  

Efuzja to kolejny proces, który polega na ruchu cząsteczek gazu. Jest to zdolność cząsteczek gazu do przemieszczania się przez otwór, którego średnica jest znacznie mniejsza niż średnia swobodna ścieżka samego gazu w odpowiedzi na różnicę ciśnień.  

To dlatego balony z helem w końcu opróżniają się - hel stopniowo wydostaje się przez maleńkie pory w materiale balonu.

Podobnie jak dyfuzja, szybkość wysięku zależy od prędkości RMS i masy molowej gazu. W szczególności szybkość wysięku jest odwrotnie proporcjonalna do pierwiastka kwadratowego z masy molowej gazu. Dlatego cięższe gazy ulatniają się wolniej niż lżejsze.

Dla dowolnych dwóch gazów stosunek ich szybkości wysięku jest pierwiastkiem kwadratowym z odwrotnego stosunku ich mas molowych. Nazywa się to prawem wysięku Grahama. 

Wyobraźmy sobie dwa balony napompowane do tego samego ciśnienia – jeden wypełniony helem, a drugi tlenem. Hel ma mniejszą masę molową niż tlen, o czym świadczy wyporność balonu z helem w powietrzu.

Zastosowanie prawa Grahama do helu i tlenu sugeruje, że hel ulatnia się 2,8 razy szybciej niż tlen. W ten sposób balon z helem opróżnia się szybciej niż balon tlenowy.

zachowanie cząsteczek gazu: dyfuzja molekularna, średnia droga swobodna i wysięk

Chociaż cząsteczki gazu poruszają się z ogromną prędkością (setki metrów na sekundę), podczas swojej drogi zderzają się z innymi cząsteczkami gazu i podróżują w wielu różnych kierunkach, zanim dotrą do pożądanego celu. W temperaturze pokojowej cząsteczka gazu będzie doświadczać miliardów zderzeń na sekundę. Średnia droga swobodna to średnia odległość, jaką pokonuje cząsteczka pomiędzy zderzeniami. Średnia swobodna droga wzrasta wraz ze spadkiem ciśnienia. W praktyce, średnia swobodna droga cząsteczki gazowej będzie setki razy większa od średnicy cząsteczki

Ogólnie rzecz biorąc, gdy próbkę gazu wprowadza się do jednej części zamkniętego pojemnika, jej cząsteczki bardzo szybko rozpraszają się w całym pojemniku; ten proces, w wyniku którego cząsteczki rozpraszają się w przestrzeni w odpowiedzi na różnice w stężeniu, nazywany jest dyfuzją. Atomy lub cząsteczki gazowe są oczywiście nieświadome jakiegokolwiek gradientu stężeń; po prostu poruszają się losowo — obszary o wyższym stężeniu zawierają więcej cząstek niż obszary o niższym stężeniu, w związku z czym następuje ruch netto cząsteczek z obszarów o wysokim stężeniu do obszarów o niskim stężeniu. W zamkniętym środowisku dyfuzja ostatecznie doprowadza do jednakowego stężenia gazu w całym pomieszczeniu. Atomy i cząsteczki gazowe w dalszym ciągu się poruszają, ale ponieważ ich stężenia są takie same w obu bańkach, szybkości przenoszenia między bańkami są równe (nie następuje transfer netto cząsteczek). Ilość gazu przechodzącego przez dany obszar w jednostce czasu to szybkość dyfuzji.

<img alt="Eq1" src="https://www.jove.com/files/core/CDNSource/lessons/1044/1044_Equations-01.png" style="height: 40px;">


Szybkość dyfuzji zależy od kilku czynników: gradientu stężenia (wzrost lub spadek stężenia z jednego punktu do drugiego), wielkości powierzchni dostępnej do dyfuzji oraz odległości, jaką muszą pokonać cząsteczki gazu.

Procesem obejmującym ruch substancji gazowych podobny do dyfuzji jest wysięk, czyli ucieczka cząsteczek gazu przez maleńki otwór, taki jak dziurka w balonie, do próżni. Chociaż szybkości dyfuzji i wysięku zależą od masy molowej gazu, wartości te nie są równe, ale ich proporcje tak.

Jeśli mieszaninę gazów umieści się w pojemniku o porowatych ściankach, gazy ulatniają się przez małe otwory w ścianach. Lżejsze gazy przechodzą przez małe otwory szybciej (z większą szybkością) niż cięższe. W 1832 roku Thomas Graham zbadał szybkość wysięku różnych gazów i sformułował prawo Grahama wysięku: Szybkość wysięku gazu jest odwrotnie proporcjonalna do pierwiastka kwadratowego masy jego cząstek:

<img alt="Eq2" src="https://www.jove.com/files/core/CDNSource/lessons/1044/1044_Equations-02.png" style="height: 40px;">


Oznacza to, że jeśli dwa gazy, A i B, mają tę samą temperaturę i ciśnienie, stosunek ich szybkości wysięku jest odwrotnie proporcjonalny do stosunku pierwiastków kwadratowych z mas ich cząstek:

<img alt="Eq3" src="https://www.jove.com/files/core/CDNSource/lessons/1044/1044_Equations-03.png" style="height: 40px;">

Zależność wskazuje, że lżejszy gaz ma większą szybkość wysięku.

Na przykład gumowy balon wypełniony helem opróżnia się szybciej niż balon wypełniony powietrzem, ponieważ szybkość wypływu przez pory gumy jest większa w przypadku lżejszych atomów helu niż w przypadku cząsteczek powietrza.

Zachowanie cząsteczek gazu: dyfuzja molekularna, średnia droga swobodna i wysięk

Chociaż cząsteczki gazowe poruszają się z ogromnymi prędkościami (setki metrów na sekundę), zderzają się z innymi cząsteczkami gazowymi i przemieszczają się w wielu ...

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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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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

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

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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...

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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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New York State Science Learning Standards

HS. Structure and Properties of Matter

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Classifying Matter by State

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Classifying Matter by Composition

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Physical and Chemical Properties of Matter

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Physical Properties of Matter

A physical property is ...

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Intermolecular vs Intramolecular Forces

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

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Surface Tension, Capillary Action, and Viscosity

Surface Tension
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States of Matter and Phase Changes

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Phase Transitions

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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

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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

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Phase Transitions: Sublimation and Deposition

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Heating and Cooling Curves

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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

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Quantifying Heat

Thermal Energy&#160;

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Quantifying Work

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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

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Constant Pressure Calorimetry

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Constant Volume Calorimetry

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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

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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

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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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