The crucial conditions for an ideal gas are as follows. First of all, the gas must have a relatively low pressure. This is because the molecules are quite far apart and only occasionally meet. Since molecules interact only occasionally, their interactions can usually be ignored. Being able to ignore these interactions is part of what contributes to the creation of a gas ideal. You can`t ignore these high-pressure or high-density interactions. The ideal gas law works quite well, but it is not perfect. It is based on molecules without interaction. When molecules interact, the whole thing collapses. While the ideal gas law is easy to understand, remember and apply, it has an obvious limitation. It describes an ideal gas. Gases are not ideal.

This is not to say that gases are terribly different from ideal. The law of ideal gases, also called the general equation of gases, is the equation of state of a hypothetical ideal gas. This is a good approximation of how many gases behave under many conditions, although it has several limitations. It was first formulated in 1834 by Benoît Paul Émile Clapeyron as a combination of Boyles` empirical law, Charlemagne`s law, Avogadro`s law and Gay-Lussac`s law. [1] The law of perfect gases is often written in empirical form: an ideal gas occupies a tank volume of 400 ft3 at a pressure of 200 psig and a temperature of 100 °F. ^ b. In an isenthalpic process, the thalpy (H) system is constant. In the case of free expansion for an ideal gas, there are no molecular interactions and the temperature remains constant. In real gases, molecules interact by attraction or repulsion depending on temperature and pressure, and heating or cooling occurs. This is called the Joule-Thomson effect. For reference, the Joule-Thomson coefficient μJT for air at room temperature and sea level is 0.22 °C/bar. [7] Let q = (qx, qy, qz) and p = (px, py, pz) be the position vector or momentum vector of a particle of an ideal gas.

F denotes the net force on this particle. Next, the average kinetic energy over time of the particle is: at relatively low pressures, gas molecules have virtually no attraction to each other because they are (on average) so far apart and behave almost like particles of an ideal gas. At higher pressures, however, the attraction is no longer insignificant. This force brings the molecules a little closer, slightly decreasing the pressure (if the volume is constant) or decreasing the volume (at constant pressure) ([link]). This change is more pronounced at low temperatures because molecules have a lower KE relative to attractive forces and can therefore overcome these attractive forces less effectively after a collision. where the heat capacity ratio (CP/CV) is 5/3 for a monatomic gas such as helium. Because times and lengths can be measured with great precision, measuring sound speeds by detecting successive resonances in a cylindrical cavity (constant frequency length variation) seems to offer an ideal way to measure temperature. However, this is not quite correct, because the limiting effects (wall and edge) that affect the speed of sound are important even for the simplest case, where there is only one mode in the cavity (frequencies of a few kilohertz). Unfortunately, these effects become greater when the pressure is reduced. An excellent theory relates damping in gas to these changes in speed, but the situation is very complex and satisfactory results are only possible with all the attention to detail.

An alternative configuration uses a spherical resonator, where the acoustic movement of the gas is perpendicular to the wall, eliminating viscosity boundary layer effects. The most reliable recent determination of the gas constant R is based on very careful measurements of the speed of sound in argon as a function of pressure at 273.16 K using a spherical resonator. Please note that the „ideal gas law“ is „ideal“ because it only works if you assume that the conditions are „ideal“. And now all gases behave ideally in conditions of high temperature and low pressure. At low temperatures, there are fewer gas molecules in a given volume. And at high temperatures, interactions between gas molecules are minimized, because the energy of the molecules is relatively higher than the intermolecular forces (among molecules). Under low temperature and high pressure, intermolecular forces and molecular size become important and are no longer negligible, so the law of perfect gases does not work. Under which of the following conditions does a real gas behave most closely with an ideal gas, and under what conditions is a real gas expected to deviate from ideal behavior? Explain.