What is the difference between real gas and ideal gas? This is a fundamental question in the field of thermodynamics and chemistry. To understand the distinction, we need to delve into the characteristics and behaviors of both types of gases.
An ideal gas is a theoretical concept that assumes certain properties to simplify the study of gas behavior. It is defined by the following postulates: 1) Gas particles have no volume, 2) There are no intermolecular forces between gas particles, 3) Collisions between gas particles are perfectly elastic, and 4) The average kinetic energy of gas particles is directly proportional to the absolute temperature. Ideal gases can be described using the ideal gas law, which states that the pressure, volume, and temperature of a gas are related by the equation PV = nRT, where P is pressure, V is volume, n is the number of moles, R is the ideal gas constant, and T is temperature.
On the other hand, a real gas is a gas that does not perfectly adhere to the assumptions of the ideal gas model. Real gases have particles with finite volume and intermolecular forces, and their collisions are not perfectly elastic. As a result, real gases exhibit deviations from the ideal gas behavior. The most significant differences between real gases and ideal gases are as follows:
1. Volume: In an ideal gas, particles are assumed to have no volume, which means that the volume of the gas is solely determined by the container. However, real gas particles have a finite volume, which affects the overall volume of the gas.
2. Intermolecular forces: Ideal gases have no intermolecular forces between particles, while real gases do. These forces can cause deviations from the ideal gas behavior, particularly at low temperatures and high pressures.
3. Elastic collisions: In an ideal gas, collisions between particles are perfectly elastic, meaning that no kinetic energy is lost during the collision. In real gases, collisions are not perfectly elastic, and some kinetic energy is lost, which can lead to deviations from the ideal gas law.
4. Deviations from the ideal gas law: At high pressures and low temperatures, real gases deviate from the ideal gas behavior. This is because the assumptions of the ideal gas model break down under these conditions. To account for these deviations, modifications to the ideal gas law, such as the Van der Waals equation, are used.
In conclusion, the main difference between real gases and ideal gases lies in the assumptions made about the behavior of gas particles. Ideal gases are a theoretical concept that simplifies the study of gas behavior, while real gases represent the actual behavior of gases under various conditions. Understanding these differences is crucial for accurately predicting and describing the properties of gases in real-world applications.
