In the early days of thermodynamics, scientists were working on the basic principles of gases. One of these scientists was Johannes Petrus Loschmidt who described what we now know as the Ideal Gas Law in 1856. However, this law is not actually accurate when it comes to describing how gases behave in the real world. In this blog post, we will explore the differences between ideal and real gases, and discuss how this affects our understanding of thermodynamics. Stay tuned!
What is Ideal Gas?
Ideal gas is an imaginary gas that consists of perfectly elastic collisions between atoms or molecules. The molecules are assumed to be point particles with no size or intermolecular interactions. Ideal gas is also assumed to have constant temperature and volume. Ideal gas is a useful concept because it allows us to simplify and understand the behavior of real gases. For example, the Ideal Gas Law can be used to predict the behavior of a real gas under different conditions. However, it is important to remember that ideal gas is not a reality, and real gases do not always behave as predicted by the Ideal Gas Law.
What is Real Gas?
Real gas is a gas that closely obeys the ideal gas law over a wide range of temperatures and pressures. The term is used in contrast to an ideal gas, which is a hypothetical gas that exactly obeys the ideal gas law. Real gases differ from ideal gases due to the presence of intermolecular forces, which cause the molecules of a real gas to deviate from perfect random motion. As a result, real gases deviate from the ideal gas law at high densities and low temperatures, where the effects of intermolecular forces are more pronounced.
Real gases also undergo phase transitions, such as condensation and freezing, which cannot be predicted by the ideal gas law. Despite these deviations, the behavior of real gases can be accurately described by the equations of state for ideal gases with modified constants. These modified constants, known as virial coefficients, depend on the type of intermolecular forces present in the gas. By taking into account the virial coefficients of a real gas, engineers and scientists are able to make accurate predictions about the behavior of gases under a wide range of conditions.
Difference between Ideal Gas and Real Gas
Ideal gas and real gas are two types of gases that have some similarities and some differences.
- They both consist of molecules that are in constant motion and exert forces on their surroundings.
- The main difference between ideal gas and real gas is that ideal gas is a hypothetical gas, while real gas is a physical gas.
- Ideal gas obeys the laws of thermodynamics perfectly, while real gas doesn’t obey those laws perfectly.
- Real gas molecules are attracted to each other and they take up space, while ideal gas molecules don’t have those attractions and don’t take up space.
- The Ideal Gas Law uses the assumptions that ideal gases make to calculate properties of gases, such as pressure, volume, temperature, and moles. The Ideal Gas Law is: PV = nRT, where P is pressure, V is volume, n is moles, R is the Ideal Gas Constant, and T is temperature.
- Real gases deviate from Ideal gases at high pressures or low temperatures because their molecules are attracted to each other. The van der Waals equation corrects for these attractions: (P + an^2/V^2)(V – b) = RT. Where a and b are constants that depend on the type of molecule.
The Ideal Gas Law is a good approximation for most gases under most conditions.
Conclusion
In conclusion, the Ideal Gas Law is a mathematical model that helps to describe the physical properties of gases. However, it does not always perfectly reflect the behavior of real gases. The Real Gas Law takes into account the effects of molecular interactions and thus provides a more accurate description of gas behavior.
