Name: how real gases deviate from ideal behavior
Uploaded: 2025-05-22T09:05:49.000Z
Duration: 1 min 10 s
Channel: JoVE

La ecuación del gas ideal tiene dos inconvenientes significativos: en primer lugar, no considera el volumen de las moléculas de gas y, en segundo lugar, no tiene en cuenta las fuerzas intermoleculares atractivas entre las moléculas de gas. 
Si el gas contiene n número de moles y cada uno de ellos ocupa el volumen b, entonces el volumen total ocupado por las moléculas de gas es nb. Por lo tanto, el volumen disponible para que se muevan las moléculas de gas será el volumen total menos nb.
Las fuerzas intermoleculares de atracción reducen la presión proporcional al cuadrado de la densidad molar. En ausencia de fuerzas intermoleculares atractivas, se requiere una presión más alta para alcanzar la misma nRT.
Esta es la ecuación de estado de van der Waals que trata de abordar los inconvenientes de la ecuación del gas ideal. Aquí, a y b son constantes empíricas que dependen del tipo de gas bajo estudio.
La ecuación de van der Waals predice adecuadamente el comportamiento del gas real, como la transición de líquido a vapor y el efecto Joule-Thomson.

<ol>
	<li>Young, H.D. y Freedman, R.A. (2012). Física Universitaria con Física Moderna. San Francisco, CA: Pearson; Sección 18.1; páginas 595 y 617.</li>
	<li>OpenStax. (2019). Física Universitaria Vol. 2. [Versión web]. Obtenido de <a href="https://openstax.org/details/books/university-physics-volume-2">https://openstax.org/details/books/university-physics-volume-2</a>; Sección 2.1; Página 75.</li>
</ol>

la ecuación de van der waals

Van der Waals Equation

La ley del gas ideal es una aproximación que funciona bien a altas temperaturas y bajas presiones. La ecuación de estado de van der Waals (llamada así por el físico holandés Johannes van der Waals, 1837-1923) mejora esta aproximación al considerar dos factores. 

 En primer lugar, se consideran las fuerzas atractivas entre moléculas, que son más fuertes a densidades más altas y reducen la presión, al agregar a la presión un término igual al cuadrado de la densidad molar multiplicado por un coeficiente positivo a. En segundo lugar, el volumen de las moléculas se representa mediante una constante positiva b, que puede considerarse como el volumen de un mol de moléculas. Esto se resta del volumen total para obtener el volumen restante en el que las moléculas pueden moverse. Las constantes a y b se determinan experimentalmente para cada gas. La ecuación resultante es 
 <img alt="Equation1" src="/files/ftp_upload/13668/13668_Equation_1.svg" style="text-align: center; height: 30px;"> 

 Para el gas dióxido de carbono con la ecuación de van der Waals, la constante a es de 0.364 J·m3/mol2 y la constante b es de 4.27 x 10−5 m3/mol. Si 1 mol de este gas se confina en un volumen de 300 cm3 a 300 K, entonces la presión del gas se puede calcular utilizando la ecuación de van der Waals. Reorganizando la ecuación de van der Waals para la presión 
 <img alt="Equation2" src="/files/ftp_upload/13668/13668_Equation_2.svg" style="text-align: center; height: 30px;"> 

 y sustituyendo las magnitudes conocidas en ella 
 <img alt="Equation3" src="/files/ftp_upload/13668/13668_Equation_3.svg" style="text-align: center; height: 32px;"> 

 se obtiene la presión del gas dióxido de carbono 
 <img alt="Equation4" src="/files/ftp_upload/13668/13668_Equation_4.svg" style="text-align: center; height: 16px;"> 

 En el límite de baja densidad (pequeño n), los términos a y b son despreciables, y la ecuación de van der Waals se reduce a la ley del gas ideal. Por otro lado, si el segundo término de la ecuación de van der Waals es pequeño, lo que significa que las moléculas están muy cerca una de otra, entonces la presión debe ser mayor para dar el mismo nRT, como se esperaría en la situación de un gas altamente comprimido. Sin embargo, el incremento en la presión es menor de lo que ese argumento sugeriría debido a que, a altas densidades, el término de corrección de presión de la ecuación de van der Waals es significativo. Dado que el término de corrección de presión es positivo, se requiere una presión más baja para dar el mismo nRT. La ecuación de estado de van der Waals funciona bien para la mayoría de los gases en diversas condiciones, como para predecir transiciones de fase líquido-gas.

La ecuación de van der Waals

La ley de los gases ideal es una aproximación que funcione bien a altas temperaturas y bajas presiones. La ecuación de estado de van der Waals (llamada así por el ...

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how real gases deviate from ideal behavior

Real gases often deviate from the ideal gas law, especially at low temperatures and high pressures. The van der Waals equation improves the ideal gas model by adding two corrections that account for these real-world effects.

One correction accounts for attractive forces between molecules. These forces become stronger at higher densities and lower the measured pressure, so the equation adds a pressure term based on the square of the molar density and a positive constant a.

The second correction accounts for the space occupied by the molecules themselves. A positive constant b represents the volume of the molecules, or the volume of a mole of molecules, and this amount is subtracted from the total volume to leave the space the gas can actually move in. The values of a and b are found experimentally for each gas.

For carbon dioxide, the van der Waals constants are a = 0.364 J·m3/mol2 and b = 4.27 x 10−5 m3/mol. If 1 mole of carbon dioxide is confined to 300 cm3 at 300 K, the pressure can be calculated by rearranging the van der Waals equation for pressure and substituting the known values.

At low density, the a and b terms are small, so the van der Waals equation reduces to the ideal gas law. At very high density, the molecules are close together, so a higher pressure is needed to reach the same nRT. Even then, the pressure increase is smaller than a simple compression argument suggests because the pressure correction term becomes significant. This makes the van der Waals equation useful for many gases, including cases that involve liquid-gas phase transitions.

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How Real Gases Deviate from Ideal Behavior | JoVE K12

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How Real Gases Deviate from Ideal Behavior

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