Real Gas and Ideal Gas
An ideal gas obeys the gas laws regardless of temperature or pressure.an ideal gas follows kinetic molecular theory. The gas particles must have zero volume and no attraction forces between them. But there is no such thing as a perfect gas since none of the conditions can be fulfilled. On the other hand, a real gas does not act according to the assumptions of the kinetic-molecular theory. Real gases, fortunately, behave similar to ideal gases at the temperature and pressure conditions observed in a laboratory. Therefore, we can say that a real gas deviates from the rules of an ideal gas at low temperatures and high pressures.
In this article, we will discuss the real and ideal gases in detail and the differences between them and also what is a perfect gas.
Ideal Gases
An ideal gas is made up of several randomly moving point particles with no interparticle interactions. The ideal gas obeys the ideal gas law. At higher temperatures and lower pressures, a gas typically behaves like an ideal gas because the potential energy from intermolecular forces is less important compared to the kinetic energy of the particles and the size of the molecules is less important compared to the void space between them. No attraction or repellence exists between the molecules of ideal gases. The sole interaction between molecules of an ideal gas would be an elastic collision when they collided or an elastic collision with the container walls. The molecules of an ideal gas have no volume. The molecules of the gas take up volume because they spread out over a broad area of space. Although there are many gases that are close to ideal, there are few perfect gases, making the idea of an ideal gas a very useful approximation for many.
Real gas
Real gases are non-ideal gases with molecules that occupy a specific amount of space and can interact with one another. Real gases do not always follow the gas laws, regardless of temperature or pressure. Real gases have mass, volume, and velocity. When cooled down to their boiling point, they get liquified. The volume filled by the gas is not small when compared to the overall volume of the gas.
A number of things need to be taken into account in order to understand how real gases behave. The following is a list that must be taken into account when working with actual gases.
Impacts of compressibility
Particular heat capacity that varies;
Forces of van der Waals;
Thermodynamic non-equilibrium effects
Molecular dissociation problems and elementary reactions with varying composition
Difference Between Ideal Gas and Real Gas
Perfect Gas and Ideal Gas
In the view of gas laws, ideal gases are similar to perfect gases.
Ideal gas, sometimes known as perfect gas, is a gas that physically complies with the ideal, or general, gas law, a particular idealised relationship between pressure, volume, and temperature. The only difference is that temperature affects specific heat. Since intermolecular interaction is minimal at large distances between molecules, actual gases frequently behave like ideal gases at high temperatures or low pressures.
Ideal and perfect gases both contain internal energy, and entropy changes with temperature.
Conclusion
Ideal gases obey all gas laws regardless of pressure or temperature. In contrast to real gases, which have restricted quantities, they have no volume. Particles in an ideal gas have no forces of attraction and collide elastically. Real gases produce few attractive forces. An ideal gas has far greater pressure than a real gas because its particles lack the forces of attraction that keep particles apart when colliding. The difference might be more obvious when the gas particles are larger and have strong attraction forces.
FAQs on Real Gas and Ideal Gas|Difference Between Real and Ideal Gas for JEE
1. Why do real gases behave differently than ideal gases?
The volume filled by gas molecules is negligibly small compared to the volume occupied by the gas. In the case of low temperatures and high pressures, the first assumption holds true, when the volume filled by gas molecules is less than the total volume of the gas.
However, because molecules are compressible at low temperatures and high pressures, their volumes are no longer insignificant in contrast to the total volume of the gas. The forces of attraction are negligible between the gas molecules. The second assumption is erroneous when the pressure is high and the temperature is low.
2. Why does the Earth's atmosphere act as a perfect gas?
The Earth's atmosphere comprises mainly nitrogen and oxygen, both of which behave quite well at the temperatures and pressures found there. Nitrogen, the most common gas in the atmosphere, is especially close to exhibiting perfect behaviour. At 20 atmospheres at normal temperature, gaseous oxygen deviates from ideal by around 3%, with deviations from ideal decreasing more or less linearly when pressure is dropped.
H2O is the only component whose behaviour is clearly non-ideal. Water vapour is a trace gas that accounts for only a few percent of the atmosphere (and that extreme occurs only in very humid tropical regions).
3. Does the ideal gas law apply to liquids?
The gas rules are only applicable to ideal gases. Ideal gases are those whose behaviour can be predicted. Listed below are some distinctive features of ideal gases:
They take up very little space.
There are no intermolecular interactions between the molecules.
The gas particles move at random.
Elastic collisions happen at random.
Because solids and liquids have intermolecular interaction, which gives them their particular structure and characteristics, the ideal gas rules do not apply to them. So it is understood that the ideal gas law may not be applicable to all liquids.