Name: kinetic molecular theory and gas behavior
Uploaded: 2020-09-03T13:38:26.000Z
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NC K-12 Science Standard Course of Study

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Energy

PS.Chm.7.3

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

Mientras que las leyes de los gases resumen las relaciones entre las distintas propiedades de los gases ideales, la teoría cinética molecular explica por qué los gases siguen dichas leyes. Esta teoría se basa en algunos supuestos o postulados. El primer supuesto es que el tamaño de las partículas de gas es insignificante.Un gas es sobre todo espacio vacío compuesto por pequeñas partículas que están separadas por distancias muchísimo más grandes que sus propias dimensiones. Su volumen combinado es insignificante en comparación con el volumen total donde el gas está contenido. A diferencia de los sólidos y líquidos, que son incompresibles debido a sus espacios intermoleculares estrechos, los gases son muy compresibles.Las partículas de gas están en constante movimiento, y van en líneas rectas y en direcciones aleatorias. Sus rutas solo cambian cuando colisionan con otras partículas o con las paredes de su contenedor. El segundo supuesto es que entre las partículas de gas se producen colisiones perfectamente elásticas.Colisionan y rebotan sin quedarse pegadas entre ellas. Esto se puede comparar a la colisión de las bolas en un juego de billar. Cuando las partículas de gas chocan, intercambian energía entre ellas, pero no hay pérdida de energía.En otras palabras, la energía total del sistema se mantiene constante. Las partículas de gas se mueven constantemente, por lo que poseen energía cinética. Por consiguiente, el tercer supuesto postula que la energía cinética promedio de un gas es proporcional a su temperatura absoluta en Kelvin.Esto significa que la energía cinética aumenta con la temperatura y, por eso, las partículas se mueven más rápido. A temperaturas más altas, su velocidad incrementa. En cambio, cuando la temperatura disminuye, la energía cinética de las partículas también lo hace, por lo que se mueven más lento.A una temperatura determinada, todos los gases, independientemente de su masa molecular, tienen la misma energía cinética promedio. La energía cinética es igual a la masa por la velocidad al cuadrado. Por lo tanto, para que distintos gases tengan la misma energía cinética promedio, sus partículas de gas deben viajar a diferentes velocidades promedio.Es por esto que los gases más pesados tienen velocidades promedio más bajas, mientras que los gases más livianos tienen velocidades promedio más altas. Por ejemplo, cuando el helio y el neón están a la misma temperatura, tienen la misma energía cinética promedio. Sin embargo, dada la diferencia en sus masas, los átomos de neón se mueven mucho más lento que los del helio.

teoría cinético molecular: postulados básicos

La ecuación de los gases ideales, que es empírica, describe el comportamiento de los gases estableciendo relaciones entre sus propiedades macroscópicas. Por ejemplo, la ley de Charles establece que el volumen y la temperatura están directamente relacionados. Los gases, por lo tanto, se expanden cuando se calientan a presión constante. Aunque las leyes de los gases explican cómo las propiedades macroscópicas cambian en relación unas con otras, no explican la razón que hay detrás de ello.
 
 La teoría cinético molecular es un modelo microscópico que ayuda a entender lo que sucede con las partículas de gas a nivel molecular o atómico cuando condiciones como la presión o la temperatura cambian. En 1857 Rudolf Clausius publicó una versión completa y satisfactoria de la teoría, que explica efectivamente las diferentes leyes de los gases a través de los postulados que fueron desarrollados sobre la base de cientos de observaciones experimentales del comportamiento de los gases.
 
 Las características más destacadas de esta teoría son:
 
 <ol>
 <li>Los gases se componen de partículas (átomos o moléculas) que están en movimiento continuo, viajando en líneas rectas y cambiando de dirección sólo cuando chocan con otras moléculas o con las paredes de un contenedor. 
 Examine una muestra de gas argón a temperatura y presión estándar. Muestra que sólo el 0,01% del volumen es ocupado por átomos con una distancia promedio de 3,3 nm (el radio atómico del argón es 0,097 nm) entre dos átomos de argón. La distancia es mucho mayor que su propia dimensión.</li>
 <li>Las moléculas que componen el gas son insignificantemente pequeñas en comparación con las distancias entre ellas. Por lo tanto, el volumen combinado de todas las partículas de gas es insignificante en relación con el volumen total del contenedor. Las partículas se consideran “puntos” que tienen masa pero volumen insignificante.</li>
 <li>La presión ejercida por un gas en un contenedor es el resultado de colisiones entre las moléculas de gas y las paredes del contenedor.</li>
 <li>Las moléculas de gas no ejercen ninguna fuerza atractiva o repulsiva sobre las paredes de los contenedores; por lo tanto, sus colisiones son elásticas (no implican una pérdida de energía). 
 Durante las colisiones elásticas, la energía se transfiere entre las partículas en colisión. La energía cinética media de las partículas, por lo tanto, permanece constante y no cambia con el tiempo.</li>
 <li>La energía cinética media de las moléculas de gas es proporcional a la temperatura kelvin del gas. 
 Todos los gases, independientemente de su masa molecular, tienen la misma energía cinética media a la misma temperatura.</li>
 </ol>
 
 Este texto ha sido adaptado de <a href="https://openstax.org/books/chemistry-2e/pages/9-5-the-kinetic-molecular-theory">Openstax, Química 2e, Capítulo 9.5 La Teoría Cinético-Molecular</a>.

Teoría Cinético Molecular: Postulados Básicos

La ecuación del gas ideal, que es empírica, describe el comportamiento de los gases estableciendo relaciones entre sus propiedades macroscópicas. Para...

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kinetic molecular theory and gas behavior

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

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

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

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

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

Kinetic Molecular Theory and Gas Behavior

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

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Pressure Units and Tools for Gas Measurement

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

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How Gas Pressure, Volume, and Temperature Change

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

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

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Ideal Gas Law at STP and Gas Density

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

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

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Gas Pressure in Mixtures: Partial Pressures

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

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

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Stoichiometry of Gases with the Ideal Gas Law

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

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

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

kinetic molecular theory and gas laws explain properties of gas molecules

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Gas Pressure Patterns in Kinetic Molecular Theory

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

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

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Gas Particle Speed Distributions and Temperature

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

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

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Gas Particle Motion, Diffusion, and Effusion

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

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Real Gases: When the Ideal Gas Law Breaks Down

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

key terms and concepts

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

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Kinetic Molecular Theory and Gas Behavior

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