What assumption is made for an ideal gas, and what gives us the right to make that assumption?

An ideal gas is a theoretical gas composed of a large number of randomly moving, non-interacting point particles. The behavior of an ideal gas can be described by the ideal gas law, which is given by the equation: \[PV = nRT\] Where P is the pressure, V is the volume, n is the number of moles of the gas, R is the ideal gas constant, and T is the temperature.

The main assumption made for an ideal gas is that the volume occupied by the individual gas molecules is negligible compared to the volume of the container. In other words, the size of gas particles is much smaller than the distances between them, and the gas particles do not occupy any significant volume within the container. Another important assumption is that there are no intermolecular forces between the gas particles, meaning there is no attraction or repulsion between the particles. Because of this, the only forces acting on a gas particle are due to collisions with the container walls.

The assumption of an ideal gas can be made in many real-life situations where the gas particles are sufficiently far apart so that their volume and the intermolecular forces can be ignored. In most cases, this assumption is valid at relatively low pressures and high temperatures. At low pressures, the gas molecules are far apart, occupying a large volume, and their individual volumes become insignificant. At high temperatures, the kinetic energy of the gas particles is high, and the influence of intermolecular forces is reduced. In conclusion, we can make the assumption of an ideal gas when the volume occupied by the gas particles is negligible compared to the volume of the container and there are no significant intermolecular forces between the particles. This assumption is often valid for many real-life situations, especially at low pressures and high temperatures.