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An ideal gas is made up of minute particles that are randomly moving and colliding with one another in elastic collisions. Real gases are those that do not behave in accordance with the ideal relations of gas law. The departure of real gas behaviour from ideal gas behaviour happens as a result of the assumption that, as pressure increases, volume falls, as pressure increases. The volume will approach zero, but it will never be zero since the molecules will occupy some space that cannot be squeezed any farther as a result of the compression.
The physical observation of gases conforms exactly to the theoretical model of the same gas. The challenge emerges when we try to determine how well the connection pV = nRT, the ideal gas equation, corresponds to the actual pressure-volume-temperature relationship of gases in real life situations. For the sake of verifying this claim, let us plot a graph of the pV vs V. When the temperature remains constant, the pV relationship will remain constant. The graph of pV vs p will be a straight line parallel to the x axis, as seen below.
The graph in the preceding graphic was created using actual data for a variety of gases at 273 degrees Celsius.
When looking at the graph, it can be seen that for real gases, the pV versus p plot does not follow a straight line at constant temperature. There is a large deviance from the optimum state of behaviour. In the case of hydrogen and helium, if the value of p increases, the value of pV increases as well, and vice versa. In some circumstances, such as methane and carbon dioxide, there is an initial negative divergence from the ideal behaviour; but, as pressure increases, the value of pV declines and eventually achieves a minimum value. After reaching the minimum point, the pV value begins to rise and eventually crosses the line for the ideal gas, resulting in a continual positive deviation from the ideal gas.
This means that real gases do not behave in accordance with the ideal gas equation at any temperature or pressure.
When a pressure vs volume graph is plotted, it is possible to see how real gas behaves differently from ideal gas behaviour. The graph of pressure vs volume should be the same for both the experimental data, which represents the real gas, and the theoretical data, which is computed using Boyle’s law as the starting point.
On occasion, it has been noted that the measured amount of gas is more than the estimated volume when operating at high pressure. However, at low pressure, the estimated and measured volumes are getting closer to one other. As a result, it may be concluded that real gases do not exactly obey the Charles law, Boyle’s law, and Avogadro law under all temperature and pressure conditions.
Real Gas
According to the definition of real gas, it is a gas that does not obey gas laws at all conventional pressure and temperature conditions. It begins to depart from its optimum behaviour when the gas grows in size and volume and becomes enormous. True gases have three characteristics: velocity, mass, and volume. They liquefy when chilled to the point when they reach boiling point.When compared to the overall volume of gas, the amount of space occupied by gas is substantial.
The Gas Equation (both ideal and real)
An ideal gas is defined as a gas that, at all pressure and temperature conditions, obeys the rules of physics. Ideal gases have both velocity and mass, and this is true of all gases. They are deafeningly quiet. When compared to the total volume of the gas, the amount of space taken up by the gas is insignificant. Because it does not condense, there is no such thing as a triple-point in this case. The ideal gas law, often known as the universal gas equation, is a mathematical equation that describes the state of a hypothetical ideal gas. It is a reasonable approximation of the behaviour of numerous gases under many conditions, but it has a number of drawbacks and limits. According to Benoît Paul Émile Clapeyron’s original description, it is a modification of the empirical laws of Boyle, Charles, Avogadro, and Gay-Lussac, with the empirical law of Boyle serving as the basis. The ideal gas law can also be expressed in an empirical version, as follows:
pV=nRT.
Compressibility factor
When a real gas’s thermodynamic properties differ from those of an ideal gas, the compressibility factor, also known as gas deviation factor or compression factor, is used as a measure of the difference between the two. We may sum up the concept by saying that it is a correction factor that best captures the amount of gas that deviates from its ideal behaviour at a given temperature and pressure. In most cases, it is utilised to alter the law of ideal gas theory.
According to another definition, the Compressibility Factor can be described as “the relationship between the molar volume of a gas and the volume of an ideal gas.” It is denoted by the equation Z = pV / RT.
Conclusion
Real gases are those that do not behave in accordance with the ideal relations of gas law. The departure of real gas behaviour from ideal gas behaviour happens as a result of the assumption that, as pressure increases, volume falls, as pressure increases.When a real gas’s thermodynamic properties differ from those of an ideal gas, the compressibility factor, also known as gas deviation factor or compression factor, is used as a measure of the difference between the two. We may sum up the concept by saying that it is a correction factor that best captures the amount of gas that deviates from its ideal behaviour at a given temperature and pressure.
