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

K12 Standards

north-carolina-k-12-science-standard-course-of-study

NC K-12 Science Standard Course of Study

9-12

Chemistry

Energy

PS.Chm.7.3

behavior-gas-molecules-molecular-diffusion-mean-free-path

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

Una bottiglia di profumo chiusa contiene un'alta concentrazione di molecole aromatiche gassose, che si muovono costantemente e si scontrano casualmente. Nel frattempo, l'aria all'esterno della bottiglia non contiene essenzialmente nessuna di queste molecole. All'apertura della bottiglia, viene stabilito un gradiente di concentrazione tra queste regioni ad alta e bassa concentrazione.

Le molecole continuano a muoversi in modo casuale, con un movimento diretto dalla regione ad alta concentrazione a quella a bassa concentrazione. La miscelazione e la diffusione spontanea di liquidi o gas in risposta ad un gradiente di concentrazione è detta diffusione molecolare. La diffusione è un processo lento.

Anche se le particelle di gas viaggiano ad alta velocità, le numerose collisioni causano frequenti cambiamenti di velocità e di direzione. La distanza media percorsa da una particella tra le collisioni è indicata come percorso libero medio. Per una particella di gas, il percorso libero medio è influenzato dalla densità delle particelle, che influenza anche la pressione.

Con l'aumentare della densità delle particelle, aumenta la frequenza di collisione, pertanto, il loro percorso libero medio è più breve. Allo stesso modo, quando la densità delle particelle diminuisce, così fa la frequenza di collisione, portando ad un percorso libero medio più lungo. Gas diversi si diffondono a velocità diverse, a seconda della velocità delle particelle di gas.

Poiché la velocità quadratica media, o RMS, e la massa molare di un gas sono inversamente correlate, i gas più leggeri si diffondono più velocemente dei gas più pesanti. Si consideri un tubo di vetro fra i serbatoi di uguali quantità di ammoniaca e acido cloridrico gassoso. Quando i gas diffondenti si incontrano, reagiscono formando un anello di cloruro di ammonio.

L'anello è più vicino all'estremità di acido cloridrico del tubo, poichè le molecole di ammoniaca più leggere hanno viaggiato più lontano lungo il tubo rispetto alle molecole di acido cloridrico più pesanti nella stessa quantità di tempo. L'effusione è un altro processo che coinvolge il movimento delle molecole di gas. È la capacità delle molecole di gas di viaggiare attraverso un foro il cui diametro è molto più piccolo del percorso libero medio del gas stesso, in risposta ad una differenza di pressione.

Questo è il motivo per cui i palloncini di elio alla fine si sgonfiano:l'elio si riversa gradualmente attraverso minuscoli pori nel materiale del palloncino. Come la diffusione, la velocità di effusione dipende dalla velocità RMS e dalla massa molare del gas. Nel dettaglio, la velocità di effusione è inversamente proporzionale alla radice quadrata della massa molare del gas.

Pertanto, i gas più pesanti si riversano più lentamente di quelli più leggeri. Per due gas qualsiasi, il rapporto delle velocità di effusione è la radice quadrata del rapporto inverso delle loro masse molari. Questa è chiamata legge dell'effusione di Graham.

Considerate due palloncini gonfiati alla stessa pressione:uno pieno di elio, l'altro di ossigeno. L'elio ha una massa molare inferiore all'ossigeno, come mostrato dalla galleggiabilità del palloncino di elio in aria. L'applicazione della legge di Graham all'elio e all'ossigeno suggerisce che l'elio fuoriesce 2, 8 volte più velocemente dell'ossigeno.

Pertanto, il palloncino di elio si sgonfia più velocemente del palloncino di ossigeno.

comportamento delle molecole di gas: diffusione molecolare, cammino libero medio ed effusione

Sebbene le molecole gassose viaggino a velocità enormi (centinaia di metri al secondo), si scontrano con altre molecole gassose e viaggiano in molte direzioni diverse prima di raggiungere il bersaglio desiderato. A temperatura ambiente, una molecola gassosa subirà miliardi di collisioni al secondo. Il percorso libero medio è la distanza media percorsa da una molecola tra le collisioni. Il cammino libero medio aumenta al diminuire della pressione; in generale, il percorso libero medio di una molecola gassosa sarà centinaia di volte il diametro della molecola

In generale, quando un campione di gas viene introdotto in una parte di un contenitore chiuso, le sue molecole si disperdono molto rapidamente all'interno del contenitore; questo processo mediante il quale le molecole si disperdono nello spazio in risposta a differenze di concentrazione è chiamato diffusione. Gli atomi o le molecole gassose, ovviamente, non sono consapevoli di alcun gradiente di concentrazione; si muovono semplicemente in modo casuale: le regioni a concentrazione più elevata hanno più particelle rispetto alle regioni a concentrazione più bassa, e quindi avviene un movimento netto di specie dalle aree ad alta concentrazione a quelle a bassa concentrazione. In un ambiente chiuso, la diffusione risulterà infine in uguali concentrazioni di gas ovunque. Gli atomi e le molecole gassosi continuano a muoversi, ma poiché le loro concentrazioni sono le stesse in entrambi i bulbi, le velocità di trasferimento tra i bulbi sono uguali (non si verifica alcun trasferimento netto di molecole). La quantità di gas che passa attraverso una certa area per unità di tempo è la velocità di diffusione.

<img alt="Eq1" src="https://www.jove.com/files/core/CDNSource/lessons/1044/1044_Equations-01.png" style="height: 40px;" />

La velocità di diffusione dipende da diversi fattori: il gradiente di concentrazione (l'aumento o la diminuzione della concentrazione da un punto a un altro), la quantità di superficie disponibile per la diffusione e la distanza che le particelle di gas devono percorrere.

Un processo che coinvolge il movimento di specie gassose simile alla diffusione è l'effusione, la fuga di molecole di gas attraverso un minuscolo foro, come uno stenopeico in un palloncino, nel vuoto. Sebbene le velocità di diffusione e di effusione dipendano entrambe dalla massa molare del gas coinvolto, le loro velocità non sono uguali; tuttavia, i rapporti delle loro tariffe sono gli stessi.

Se una miscela di gas viene posta in un contenitore con pareti porose, i gas fuoriescono attraverso le piccole aperture nelle pareti. I gas più leggeri passano attraverso le piccole aperture più rapidamente (ad una velocità maggiore) rispetto a quelli più pesanti. Nel 1832, Thomas Graham studiò la velocità di effusione di diversi gas e formulò la legge di effusione di Graham: la velocità di effusione di un gas è inversamente proporzionale alla radice quadrata della massa delle sue particelle:

<img alt="Eq2" src="https://www.jove.com/files/core/CDNSource/lessons/1044/1044_Equations-02.png" style="height: 40px;" />

Ciò significa che se due gas, A e B, si trovano alla stessa temperatura e pressione, il rapporto tra le loro velocità di effusione è inversamente proporzionale al rapporto tra le radici quadrate delle masse delle loro particelle:

<img alt="Eq3" src="https://www.jove.com/files/core/CDNSource/lessons/1044/1044_Equations-03.png" style="height: 40px;" />

La relazione indica che il gas più leggero ha un tasso di effusione più elevato.

Ad esempio, un palloncino di gomma riempito di elio si sgonfia più velocemente di uno pieno di aria perché la velocità di effusione attraverso i pori della gomma è più rapida per gli atomi di elio più leggeri che per le molecole di aria.

Questo testo è adattato da <a href="https://openstax.org/books/chemistry-2e/pages/9-4-effusion-and-diffusion-of-gases">Openstax, Chemistry 2e, Section 9.4: Effusion and Diffusion of Gases. </a>

Comportamento delle molecole di gas: diffusione molecolare, cammino libero medio ed effusione

Sebbene le molecole gassose viaggino a velocità enormi (centinaia di metri al secondo), si scontrano con altre molecole gassose e viaggiano in molti ...

Education

Didattica

JoVE Core

Chimica

Gases

KEY TERMS AND CONCEPTS

TERMINI E CONCETTI CHIAVE

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

chemical-stoichiometry-gases-using-ideal-gas-law-to-determine

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

basic-postulates-kinetic-molecular-theory-particle-size-energy

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

kinetic-molecular-theory-gas-laws-explain-properties-gas

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

kinetic-molecular-theory-molecular-velocities-temperature-kinetic

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

real-gases-effects-intermolecular-forces-molecular-volume-deriving

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

chemistry

JoVE Core: Chimica illustra concetti chiave di chimica attraverso video animazioni concise e di immediata comprensione. Inoltre, i video della sezione "scienziati in azione" mostrano procedure metodologiche realizzate da ricercatori provenienti dai migliori laboratori internazionali.

ohios-learning-standards-and-model-curriculum

Ohio’s Learning Standards and Model Curriculum

Interactions Of Matter

Gas laws

Virginia  Standards of Learning

HIgh School

CH.6

CH.6.b

Texas Essential Knowledge and Skills

High School

CHEM.10

CHEM.10.A

Alabama Course of Study Science

Matter and Its Interactions

9-12.C.1.4.e

9-12.C.4.7

9-12.C.4.7.a

Florida Department of Education

Physical Science

SC.912.P.10

SC.912.P.10.5

new-york-state-science-learning-standards

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 ...

physical-and-chemical-properties-of-matter

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 ...

molecular-comparison-of-gases-liquids-and-solids

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 ...

comparing-intermolecular-forces-melting-point-boiling-point

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, ...

surface-tension-capillary-action-and-viscosity

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 ...

phase-transitions-vaporization-and-condensation

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 ...

phase-transitions-sublimation-and-deposition

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 ...

using-differential-scanning-calorimetry-to-measure-changes-in-enthalpy

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

Description

Transcript

Tags

More videos by Standards