Difference Between Ideal Gas and Real Gas with FAQs

Q: What is the Z compression factor?

The ratio of the observed molar volume V m to the ideal molar volume V m ideal ( = RT/p) as the function of pressure at constant temperature and this ratio is called the compression factor Z.

Edited By Team Careers360 | Updated on Jul 02, 2025 04:55 PM IST

In this article, we will discuss the difference between real gas and ideal gas, non ideal gas, ideal and real gas, ideal and real gas examples.

Differentiate between ideal gas and real gas:

Ideal Gas	Real gas
Ideal gas has no definite volume	Real gas has definite volume
There are elastic collisions of particles in ideal gas.	There is non - elastic collisions between particles in non-ideal gas.
There is absence of intermolecular attraction force in ideal gas.	There is presence of intermolecular attraction force in real gas or non-ideal gas.
It does not exist in reality that is in the environment and is a hypothetical gas.	It exists in the environment
The pressure is high	The pressure is less when compared to other ideal /perfect gas.
It is independent	It interacts with other gases.
It obeys gas laws such as pV=nRT	It obeys p + ((n² a)/V²)(V - nb) = nRT

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The above table gives the difference between real gas and ideal gas. So, we have assumed that all gases follow the gas laws under all conditions of temperature and pressure, however, this is not true for real gas. Difference between real gases and ideal gasis that real gas obeys the gas laws under limited conditions of low pressure and high temperature.

Theideal gas and real gas have major difference thatreal gases exhibit deviations from gaseous laws and the deviations increase when temperature and pressure are near to conditions in which the gas can condense into liquid, thus the Boyle’s and Charles’s law derived ideal gas equation is only applicable at relatively low pressure and moderately high temperatures.

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Ideal gas and real gas difference are that ideal gas is a hypothetical gas that obeys the gas law exactly under all conditions of temperature and pressure which is perfect gas or ideal gas. The universal gas constant R, is an experimentally derived quantity which has the same value for all ideal gases.

Now, as we know that we can distinguish between ideal gas and real gas is that ideal gas equation pV= nRT, is not applicable to real gas, hence the evaluation of universal gas constant R cannot be done directly by using pressure, volume, and temperature data of real gas.

In ideal gas and real gas, we know through various experiments it is known that gases approach ideal behaviour as the pressure is decreased. Thus, the extrapolation method (p → 0) on the data of real gases can be utilized to determine the corresponding properties of ideal gas and hence the data gained from this method, after extrapolation, should be independent of the characteristic of the real gas used for experiment.

Here, ideal and real gas can be differentiated after calculating the volumes of one mole of a real gas at different pressures and constant temperature, a graph between pV and p can be drawn, then on extrapolating the graph to zero pressure and for departure from ideal behaviour it is now possible to determine the value of pV which is expected to be appropriate to one mole of ideal gas. By following this it is found that the value of (pV)p→0 at 273.15 K is 22.711 dm³ bar. Thus, if p=1 bar, then V=22.711 dm³. The volume occupied by one mole of an ideal gas at standard temperature 273.15 K and 1 bar pressure is 22.711 dm³.

In ideal gas and real gas, the difference between real and ideal gas is that real gas is the gases which do not obey the ideal gas laws exactly under all conditions of temperature and pressure and is also called as non-ideal gas. According to experiments at low pressure and moderately high temperature, the gases obey the laws of Boyle’s Charles and Avogadro approximately. But as the pressure rises or the temperature is decreased a visible departure from ideal behaviour is seen also a difference between ideal and non-ideal gas is seen. We can see the example of the type of deviation that occurs in Boyle’s law for H₂ at room temperature.

The deviations or the difference between ideal gas and real gas can be displayed more distinctively by plotting the graph with the ratio of the observed molar volume V_m to the ideal molar volume V_m _ideal( = RT/p) as the function of pressure at constant temperature and this ratio is called the compression factor Z and this can be expressed as

Z = V_m/ V_m _ideal = pV_m/RT

The major difference between ideal gas and real gas is that for an ideal gas, Z = 1 and is independent of pressure and temperature. For example a real gas, Z = f(T,p) a function of both temperature and pressure as shown in figure below a graph between Z and p for some gases at 273.15 K, the inference we can conclude is that

File:Factor Z vs.png - Wikimedia Commons

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Difference between ideal gas and real gas is that in real gas Z the compression factor is always greater than 1 for H₂ because the attractive interactions are so weak that the repulsive interactions dominate even at low pressures.
For CH₄, at low pressure, Z is less than 1. That is, their molar volumes of the methane gas are smaller than that of a perfect gas, showing that the molecules are drawn together slightly and also, we can conclude that for these molecules and these conditions the attractive interactions are more dominant and this is the differentiate between real gas and ideal gas.

In the below figure, it gives the impression that the nature of deviations depends upon the nature of the gas. The more important determining factor is the temperature relative to the critical temperature of a selected/ particular gas which is near the critical temperature in real gas.

File:Realgasfaktor Stickstoff.jpg - Wikimedia Commons

Given the pressure is of the order of 1 bar or less, and the temperature is not too near the point of liquefaction, the observed deviations from the ideal gas laws are not more than a few % . Thus under these conditions, the equation pV =nRT and related expressions may be used. These are some of the examples of ideal gas and real gas.

The ideal gas differentiate from real gas or the difference between ideal gas and real gasis because ideal gas deviates from its ideal behaviour and this is because the ideal gas can be derived from the kinetic theory of gases which is based on the following two assumptions:

The volume occupied by the molecules is negligible in comparison to the total volume of the gas in consideration.
The molecules exert no forces of attraction on one another molecule.

As we know neither of these assumptions can be regarded as applicable to real gases that the latter show departure from the ideal gas behaviour.

The person who was the first to systematically introduce the correction terms due to the above two invalid assumptions in the ideal gas equation p_iV_i = nRT was Van der Waals.

His corrections are

Correction in Volume V_i = V – nb where b is excluded or co- volume.
Correction for forces of attraction.

The equation which is applicable to all real gas is and is known as van der waals equation.

(p + n²a / V²) ( V- nb) = nRT

Thus, due to this, there is major difference between real and ideal gas.

Also, students can refer,

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Frequently Asked Questions (FAQs)

1. What is non-ideal gas?

Non ideal gases are gases which do not obey the ideal gas laws exactly under all conditions of temperature and pressure.

2. What is ideal gas?

A hypothetical gas that obeys the gas law exactly under all conditions of temperature and pressure is perfect gas or ideal gas.

3. What is a real molecule?

The real molecules possess finite size and do not overlap and thus has a strong, short – ranged repulsion between them.

4. What is the major difference between ideal and non ideal gas?

Ideal gas and real gas have major difference of: A hypothetical gas that obeys the gas law exactly under all conditions of temperature and pressure is ideal gas and real gas is a gas that does not follow the assumptions made according to kinetic – molecular theory.

5. What is the Z compression factor?

The ratio of the observed molar volume V_m to the ideal molar volume V_m _ideal( = RT/p) as the function of pressure at constant temperature and this ratio is called the compression factor Z.

6. What is the main difference between an ideal gas and a real gas?

The main difference is that ideal gases are theoretical models that perfectly follow gas laws, while real gases are actual gases that deviate from these laws due to intermolecular forces and finite particle size. Ideal gases assume no intermolecular attractions and negligible particle volume, whereas real gases have these properties, leading to deviations from ideal behavior.

7. Under what conditions do real gases behave most like ideal gases?

Real gases behave most like ideal gases at high temperatures and low pressures. These conditions minimize intermolecular forces and the effect of particle volume, allowing real gases to more closely follow ideal gas laws. As temperature increases or pressure decreases, gas molecules have more space and energy, reducing interactions and behaving more ideally.

8. What is the van der Waals equation, and how does it relate to real gases?

The van der Waals equation is a modified version of the ideal gas law that accounts for the behavior of real gases. It introduces two constants: 'a' to account for intermolecular attractions, and 'b' to account for the volume of gas particles. This equation provides a more accurate description of real gas behavior, especially at high pressures and low temperatures where deviations from ideal behavior are more significant.

9. How do intermolecular forces affect the behavior of real gases?

Intermolecular forces cause real gas molecules to attract each other, leading to deviations from ideal gas behavior. These attractions can cause real gases to have a lower pressure than predicted by the ideal gas law at the same volume and temperature. The effect is more pronounced at lower temperatures and higher pressures, where molecules are closer together and interactions are stronger.

10. Why do real gases deviate more from ideal behavior at high pressures?

At high pressures, real gas molecules are forced closer together, increasing the significance of both intermolecular forces and the volume occupied by the gas particles. This leads to greater deviations from ideal gas behavior. The assumptions of negligible particle volume and no intermolecular interactions in the ideal gas model become less valid under these conditions.

11. Why do we study ideal gases if they don't exist in reality?

We study ideal gases because they provide a simplified model that helps us understand gas behavior under many conditions. Ideal gas laws are easier to work with mathematically and give good approximations for real gases at normal temperatures and pressures. This model serves as a foundation for understanding more complex gas behaviors.

12. What is compressibility factor (Z), and how does it indicate deviation from ideal gas behavior?

The compressibility factor (Z) is the ratio of the actual volume of a real gas to the volume predicted by the ideal gas law under the same conditions. For an ideal gas, Z = 1. When Z < 1, the gas is more compressible than an ideal gas due to attractive forces. When Z > 1, the gas is less compressible due to repulsive forces at high pressures. Z helps quantify how much a real gas deviates from ideal behavior.

13. How does the molecular size of a gas affect its deviation from ideal behavior?

Larger gas molecules tend to deviate more from ideal behavior because they occupy more space and have stronger intermolecular forces. The assumption of negligible particle volume in the ideal gas model becomes less accurate for larger molecules. Gases with smaller molecules, like helium, behave more ideally over a wider range of conditions compared to gases with larger molecules.

14. What is the Boyle temperature, and why is it significant for real gases?

The Boyle temperature is the temperature at which a real gas behaves most like an ideal gas over a wide range of pressures. At this temperature, the effects of intermolecular attractions and repulsions approximately cancel out. The Boyle temperature is significant because it represents a point where a real gas closely follows the ideal gas law, making calculations and predictions more accurate.

15. How does the ideal gas law need to be modified to describe real gas behavior?

To describe real gas behavior, the ideal gas law (PV = nRT) needs to be modified to account for intermolecular forces and the volume of gas particles. This is typically done by introducing correction factors or using more complex equations like the van der Waals equation. These modifications adjust the pressure and volume terms to better represent the actual behavior of real gases under various conditions.

16. What is the Lennard-Jones potential, and how does it relate to real gas behavior?

The Lennard-Jones potential is a mathematical model that describes the interaction between two neutral atoms or molecules. It accounts for both attractive and repulsive forces between particles. This potential is crucial in understanding real gas behavior as it provides a more realistic representation of intermolecular interactions. It helps explain phenomena like gas condensation and deviations from ideal gas behavior at high pressures and low temperatures.

17. What is the virial equation of state, and how does it improve upon the ideal gas law?

The virial equation of state is an expansion of the ideal gas law that includes additional terms to account for real gas behavior. It's expressed as a power series in terms of density or pressure. Each term in the series (virial coefficients) represents the effect of interactions between increasing numbers of molecules. This equation provides a more accurate description of real gas behavior over a wider range of conditions than the ideal gas law.

18. What is the principle of corresponding states, and how does it apply to real gases?

The principle of corresponding states suggests that all gases behave similarly when compared at the same reduced temperature and pressure (relative to their critical points). This principle allows for the prediction of real gas behavior using generalized charts or equations, even when specific data for a particular gas is not available. It's based on the idea that the deviation from ideal behavior is similar for different gases under corresponding conditions.

19. How does the presence of polar molecules in a gas affect its deviation from ideal behavior?

Gases composed of polar molecules generally deviate more from ideal behavior than non-polar gases. Polar molecules have stronger intermolecular attractions (dipole-dipole forces) in addition to van der Waals forces. These stronger attractions lead to greater deviations from ideal gas laws, especially at lower temperatures and higher pressures. Polar gases may also exhibit more complex behavior in terms of compressibility and phase transitions.

20. What is the significance of the van der Waals radius in understanding real gas behavior?

The van der Waals radius represents the size of an atom or molecule in a real gas, accounting for the fact that particles cannot overlap. This concept is crucial in understanding real gas behavior because it relates to the 'b' constant in the van der Waals equation, which corrects for the finite volume of gas particles. The van der Waals radius helps explain deviations from ideal gas behavior, especially at high pressures where molecular volume becomes significant.

21. What is the Peng-Robinson equation of state, and how does it improve upon the van der Waals equation for real gases?

The Peng-Robinson equation of state is an improved model for predicting the behavior of real gases, especially in the oil and gas industry. It builds upon the van der Waals equation by introducing more complex temperature-dependent terms. This equation provides better accuracy near the critical point and for predicting liquid densities. It's particularly useful for calculating vapor-liquid equilibria and other thermodynamic properties of hydrocarbon mixtures.

22. How does the concept of fugacity coefficient help in understanding real gas behavior?

The fugacity coefficient is the ratio of a gas's fugacity to its pressure. It quantifies how much a real gas deviates from ideal behavior. A fugacity coefficient of 1 indicates ideal behavior, while values above or below 1 indicate positive or negative deviations, respectively. This concept is crucial in thermodynamic calculations involving real gases, as it allows for more accurate predictions of chemical equilibria and phase behavior than using pressure alone.

23. What is the significance of the Redlich-Kwong equation of state in real gas calculations?

The Redlich-Kwong equation of state is an empirical modification of the van der Waals equation that provides improved accuracy for real gases. It introduces a temperature-dependent term in the attractive forces part of the equation. This modification makes it more accurate over a wider range of temperatures and pressures, especially for non-polar gases. The Redlich-Kwong equation is widely used in chemical engineering for process design and optimization involving real gases.

24. What is the concept of compressibility charts, and how are they used for real gases?

Compressibility charts are graphical tools used to determine the compressibility factor (Z) for real gases under various conditions. These charts plot Z against reduced pressure and temperature (relative to critical values). They are based on the principle of corresponding states and allow for quick estimation of real gas behavior without complex calculations. Compressibility charts are particularly useful in engineering applications where rapid, approximate solutions are needed.

25. How does the presence of impurities affect the behavior of real gases compared to pure substances?

Impurities in a gas can significantly alter its behavior compared to a pure substance. They can change the intermolecular forces, affecting properties like compressibility, critical point, and phase behavior. Even small amounts of impurities can lead to deviations from expected real gas behavior based on pure substance models. This is particularly important in industrial applications where gas purity can vary and affect process efficiency and safety.

26. What is the significance of the acentric factor in describing real gas behavior?

The acentric factor is a measure of the non-sphericity (or acentricity) of a molecule and its deviation from simple spherical molecule behavior. It's used in equations of state to improve predictions of real gas and liquid behavior. A higher acentric factor indicates greater deviation from ideal gas behavior. This concept is particularly useful in the petroleum industry for predicting the properties of complex hydrocarbon mixtures.

27. What is the Joule-Thomson effect, and how does it differ between ideal and real gases?

The Joule-Thomson effect is the temperature change that occurs when a gas expands at constant enthalpy. In an ideal gas, this effect is zero – the temperature doesn't change during expansion. However, real gases can experience either cooling or heating during this process, depending on their initial temperature and pressure. This difference arises from the intermolecular forces present in real gases but absent in ideal gases.

28. How does the critical point of a gas relate to its deviation from ideal behavior?

The critical point is the temperature and pressure at which the liquid and gas phases of a substance become indistinguishable. Near the critical point, real gases deviate significantly from ideal behavior due to strong intermolecular forces and the breakdown of distinct phase boundaries. Understanding the critical point is crucial for predicting real gas behavior under extreme conditions.

29. Why do some gases, like hydrogen and helium, behave more ideally than others?

Hydrogen and helium behave more ideally because they have very weak intermolecular forces and small molecular sizes. These properties align closely with the assumptions of the ideal gas model. Their weak attractions and small volume mean they deviate less from ideal behavior over a wider range of temperatures and pressures compared to gases with stronger intermolecular forces or larger molecules.

30. How does the concept of fugacity relate to real gas behavior?

Fugacity is a measure of the tendency of a substance to escape from a phase. For real gases, fugacity is used instead of pressure to more accurately describe their behavior in thermodynamic calculations. The ratio of fugacity to pressure (fugacity coefficient) indicates how much a real gas deviates from ideal behavior. For an ideal gas, fugacity equals pressure, but for real gases, it differs, especially at high pressures.

31. What is the significance of the second virial coefficient in describing real gas behavior?

The second virial coefficient is a measure of the deviation of a real gas from ideal behavior due to two-particle interactions. It appears in the virial equation of state, which is an expansion of the ideal gas law. A positive second virial coefficient indicates that repulsive forces dominate, while a negative value suggests attractive forces are more significant. This coefficient helps in quantifying and understanding real gas behavior under different conditions.

32. How does the internal energy of a real gas differ from that of an ideal gas?

The internal energy of a real gas includes contributions from both kinetic energy and potential energy due to intermolecular forces. In contrast, the internal energy of an ideal gas depends only on temperature (kinetic energy) because it assumes no intermolecular interactions. This difference leads to variations in thermodynamic properties and behavior between real and ideal gases, especially at low temperatures or high pressures.

33. How do real gases behave differently from ideal gases during adiabatic processes?

During adiabatic processes (no heat exchange with surroundings), real gases can exhibit different temperature changes compared to ideal gases. This is due to the Joule-Thomson effect and the presence of intermolecular forces. While ideal gases follow a simple relationship between pressure and volume during adiabatic processes, real gases may deviate from this, especially at extreme conditions, leading to different final states and work done.

34. What is the significance of the Boyle point in understanding real gas behavior?

The Boyle point is the temperature at which the second virial coefficient of a gas becomes zero. At this point, the gas behaves most like an ideal gas over a range of pressures. The Boyle point is significant because it represents a condition where the attractive and repulsive forces between gas molecules effectively cancel out, allowing the gas to follow the ideal gas law more closely. It's a useful reference point for understanding real gas behavior.

35. How does the speed of sound in a real gas differ from that in an ideal gas?

The speed of sound in a real gas can differ from that in an ideal gas due to the effects of intermolecular forces and finite particle size. In an ideal gas, the speed of sound depends only on temperature and the gas's molecular weight. In real gases, it can also be affected by pressure and the gas's compressibility. This difference becomes more pronounced at high pressures or low temperatures where real gas effects are more significant.

36. What is the Maxwell-Boltzmann distribution, and how does it apply differently to ideal and real gases?

The Maxwell-Boltzmann distribution describes the distribution of molecular speeds in a gas at thermal equilibrium. For ideal gases, this distribution depends only on temperature and molecular mass. In real gases, intermolecular forces can affect the distribution, especially at high pressures or low temperatures. This leads to slight deviations in the actual speed distribution compared to what the ideal Maxwell-Boltzmann distribution predicts.

37. How do real gases behave differently from ideal gases during Joule expansion?

In Joule expansion, a gas expands into a vacuum. For an ideal gas, there's no temperature change during this process. However, real gases can experience temperature changes due to intermolecular forces. Most real gases cool slightly during Joule expansion (except at very high temperatures), while ideal gases maintain constant temperature. This difference highlights the impact of intermolecular attractions in real gases.

38. What is the significance of the Boyle-Mariotte law for real gases?

The Boyle-Mariotte law states that for an ideal gas at constant temperature, pressure and volume are inversely proportional. For real gases, this law holds approximately at low pressures and high temperatures. Deviations from this law at high pressures or low temperatures indicate the presence of intermolecular forces and finite molecular volume in real gases. Understanding these deviations is crucial for accurate predictions of real gas behavior.

39. How does the heat capacity of a real gas differ from that of an ideal gas?

The heat capacity of a real gas can differ from that of an ideal gas due to intermolecular forces. In an ideal gas, heat capacity depends only on the gas's degrees of freedom. In real gases, additional energy can be stored in intermolecular potentials, leading to a heat capacity that varies with temperature and pressure. This difference is more pronounced at low temperatures or high pressures where intermolecular interactions are more significant.

40. How do real gases behave differently from ideal gases during isothermal compression?

During isothermal compression, real gases deviate from the behavior predicted by Boyle's law for ideal gases. At high pressures, real gases are generally less compressible than ideal gases due to the finite volume of molecules and repulsive forces. At moderate pressures, some real gases may be more compressible due to attractive forces. These deviations become more pronounced as the gas approaches its critical point or liquefaction.

41. How do real gases behave differently from ideal gases in terms of effusion and diffusion rates?

While Graham's law of effusion and diffusion applies to both ideal and real gases, real gases can show slight deviations. In ideal gases, these rates depend solely on molecular mass. For real gases, intermolecular forces and molecular size can affect these rates, especially at high pressures or low temperatures. Larger molecules or those with stronger intermolecular attractions may diffuse or effuse more slowly than predicted by Graham's law.

42. How do real gases behave differently from ideal gases in terms of their Joule-Thomson coefficient?

The Joule-Thomson coefficient describes how the temperature of a gas changes during an isenthalpic expansion. For an ideal gas, this coefficient is zero, meaning no temperature change occurs. Real gases, however, can have positive or negative Joule-Thomson coefficients, leading to cooling or heating during expansion. This behavior is due to intermolecular forces and varies with temperature and pressure, playing a crucial role in processes like gas liquefaction.