Name: gas pressure patterns in kinetic molecular theory
Uploaded: 2020-09-03T13:38:27.000Z
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NC K-12 Science Standard Course of Study

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Chemistry

Energy

PS.Chm.7.3

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Kinetic Molecular Theory and Gas Laws Explain Properties of Gas Molecules

Las observaciones de las diferentes propiedades de los gases, como lo expresan las diversas leyes de gases establecidas por Boyle, Charles, Gay-Lussac y Avogadro, siguen conceptualmente a la teoría de la cinética molecular. La presión ejercida por un gas resulta del impacto de partículas que están en constante movimiento contra las paredes de su contenedor. Disminuir el volumen del contenedor mientras el número de moles y la temperatura se mantienen constantes acerca a las partículas de gas, reduciendo el espacio interparticular.En este volumen más pequeño, aumenta la densidad del gas y la frecuencia de colisión, o sea, la frecuencia de colisiones entre las moléculas y las paredes. Por lo tanto, la presión ejercida por el gas incrementa. La relación inversa entre la presión y el volumen está establecida en la ley de Boyle.Añadir más moles de gas al contenedor a una temperatura constante incrementa la densidad del gas y, por lo tanto, la frecuencia de colisión. Para mantener la presión inicial, el volumen debe expandirse. Esta relación directa entre volumen y moles está establecida en la ley de Avogadro.Ahora, consideremos que el número de moles se mantiene constante y se aumenta la temperatura. Como la energía cinética media de las partículas de gas incrementa proporcionalmente con la temperatura, las partículas colisionan con más frecuencia y con mayor fuerza. Si el volumen se mantiene constante mientras se incrementa la temperatura, la densidad del gas y la frecuencia de colisión aumentan, y, por lo tanto, la presión también aumentará.La relación directa entre presión y temperatura está establecida en la ley de Gay-Lussac. Si, por otro lado, la presión debe mantenerse constante junto con un número constante de moles, entonces, se debe acompañar con un aumento en la temperatura con un incremento del volumen para dispersar las colisiones sobre la superficie en un área más grande. Esta relación directa entre volumen y temperatura está establecida en la ley de Charles.Finalmente, de acuerdo con la teoría cinética molecular, las partículas de gas no se atraen ni se repelen entre ellas. En una mezcla de diferentes gases, los componentes actúan de forma independiente y sus presiones individuales no se ven afectadas por la presencia de otro gas. La presión total de una mezcla es, por lo tanto, la suma de las presiones parciales individuales.Esta es la ley de Dalton.

teoría cinético molécular y las leyes de los gases

La prueba de la teoría cinético molecular (TCM) y sus postulados es su capacidad para explicar y describir el comportamiento de un gas. Las diversas leyes de los gases (las leyes de Boyle, Charles, Gay-Lussac, Avogadro y Dalton) pueden derivarse de las suposiciones de la TCM, que han llevado a los químicos a creer que las suposiciones de la teoría representan con precisión las propiedades de las moléculas de gas.
 
 <h4>La Teoría Cinético Molecular explica el comportamiento de los gases</h4>
 
 Recordando que la presión de gas es ejercida por moléculas de gas que se mueven rápidamente y depende directamente del número de moléculas que golpean una unidad de área de la pared por unidad de tiempo, la TCM explica conceptualmente el comportamiento de un gas de la siguiente manera:
 
 <ul>
 <li>Ley de Gay-Lussac: Si la temperatura aumenta, la velocidad media y la energía cinética de las moléculas de gas aumentan. Si el volumen se mantiene constante, el aumento de la velocidad de las moléculas de gas provoca colisiones más frecuentes y más fuertes con las paredes del contenedor, lo que aumenta la presión. Esto también se denomina ley de Amontons.</li>
 <li>Ley de Charles: Si la temperatura de un gas aumenta, la presión constante sólo puede mantenerse si aumenta el volumen ocupado por el gas. Esto dará como resultado mayores distancias medias recorridas por las moléculas para alcanzar las paredes del contenedor, así como un área mayor de la superficie de la pared. Estas condiciones disminuirán tanto la frecuencia de las colisiones molécula-pared como el número de colisiones por unidad de área, cuyos efectos combinados equilibran el efecto del aumento de las fuerzas de colisión debido a la mayor energía cinética a la temperatura más alta.</li>
 <li>Ley de Boyle: Si el volumen de gas disminuye, el área de la pared del contenedor disminuye y la frecuencia de colisión molécula-pared aumenta, lo que aumenta la presión ejercida por el gas.</li>
 <li>Ley de Avogadro: A presión y temperatura constantes, la frecuencia y la fuerza de las colisiones molécula-pared son constantes. En tales condiciones, el aumento del número de moléculas gaseosas requerirá un aumento proporcional en el volumen del contenedor para producir una disminución en el número de colisiones por unidad de área y compensar el aumento de la frecuencia de las colisiones.</li>
 <li>Ley de Dalton: Debido a las grandes distancias entre ellas, las moléculas de un gas en una mezcla bombardean las paredes del contenedor con la misma frecuencia, independientemente de si otros gases están presentes o no, y la presión total de una mezcla de gases es igual a la suma de las presiones (parciales) de los gases individuales.</li>
 </ul>
 
 Este texto ha sido adaptado de <a href="https://openstax.org/books/chemistry-2e/pages/9-5-the-kinetic-molecular-theory">Openstax, Química 2e, Sección 9.5: La Teoría Cinético-Molecular.</a>

Teoría Cinético Molécular y las Leyes de los Gases

La prueba de la teoría cinética molecular (KMT) y sus postulados es su capacidad para explicar y describir el comportamiento de un gas. Las diversas leyes de los gases ...

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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 often and more forcefully, so pressure increases.

Charles’s law shows what happens when temperature rises at constant pressure. The gas must expand to keep pressure steady. That larger volume means molecules travel farther before reaching the walls, and the wall area increases too. These changes lower collision frequency and the number of collisions per unit area, which balances the stronger collisions caused by the higher kinetic energy.

Boyle’s law describes the effect of smaller volume. When gas volume decreases, the container wall area also decreases. Molecule-wall collisions become more frequent, and the pressure rises. Avogadro’s law says that at constant pressure and temperature, the collision rate and collision force stay constant. If the number of gas molecules increases, the volume must increase in proportion so the collisions per unit area can drop enough to offset the added molecules.

Dalton’s law applies to gas mixtures. Because gas molecules are far apart, one gas in a mixture strikes the container walls with the same frequency whether other gases are present or not. The total pressure of the mixture equals the sum of the partial pressures of the individual gases. These gas laws show how the assumptions of kinetic molecular theory fit the behavior of gases.

Gas Pressure Patterns in Kinetic Molecular Theory

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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...

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

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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

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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

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...

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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, temperature, volume changes, and the gas laws of Boyle, Charles, Gay-Lussac, Avogadro, and Dalton.

Gas Pressure Patterns in Kinetic Molecular Theory | JoVE K12

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Quantifying Work

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Standard Enthalpy of Formation

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Gas Pressure Patterns in Kinetic Molecular Theory

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