Why do gases behave in a less ideal manner when in LOW TEMP. AND HIGH PRESSURE?
Dont real gas molecules occupy volume, and have attractive forces? Do real gases become liquids? Why? Ideal gases cant do this.
I don't know what you just said, but i know that when an ideal gas is at a lower pressure, the molecules do not collide with the container as frequently, and that at a high temperature the molecules do collide frequently. What does this mean, does it even help me?
No. Reread what I asked. I gave you the answer. If you have no idea what I am writing of, you quickly need to reread your text on the failure of the assumptions of the ideal gas vs real gases.
Ok, i understand that for real gases you have to account for volume of particles, and that they have attractive forces that lower the pressure, that make the ideal gas law invalid, but what does that have to do with low temp and high pressure?
Ok, i understand that for real gases you have to account for volume of particles, and that they have attractive forces that lower the pressure, that make the ideal gas law invalid, but what does that have to do with low temp and high pressure?
Is it because at high pressure, molecules have volume, and at low temperature gases move slowly and act like a liqud so they have attractive forces that decrease pressure?
Gases tend to deviate from ideal behavior at low temperatures and high pressures due to several reasons. One of the main reasons is the intermolecular forces between gas molecules.
At low temperatures, the kinetic energy of gas molecules decreases. As a result, the molecular motion becomes slower, and the attractive forces between gas molecules become more significant. In an ideal gas, these intermolecular forces are assumed to be negligible. However, at low temperatures, when the molecular motion is restricted, these forces play a more significant role, causing the gas to behave less ideally.
At high pressures, the gas molecules are close together, leading to an increase in intermolecular interactions. These interactions have a constraining effect on the motion of molecules, causing them to deviate from ideal behavior. The gas molecules may experience stronger attractive forces, which can lead to a decrease in the average distance between molecules and an increase in the volume occupied by the gas particles. This effect is known as the van der Waals forces, and it becomes more pronounced at high pressures.
To quantitatively describe the behavior of gases under non-ideal conditions, scientists use equations of state, such as the van der Waals equation, which include correction factors to account for the effects of intermolecular forces and molecular size. These equations help predict the behavior of real gases at various temperatures and pressures, accounting for deviations from ideal gas behavior.
To study the non-ideal behavior of gases experimentally, one can measure the behavior of the gas under different pressures and temperatures and compare the results to the predictions of the equations of state. Additionally, techniques such as the compression factor, which compares the actual volume of a gas to its volume predicted by the ideal gas law, can be used to determine if a gas behaves ideally.
In summary, gases deviate from ideal behavior at low temperatures and high pressures due to the increasing prominence of intermolecular forces. These deviations can be quantitatively described using equations of state, and experimental techniques can be employed to study and measure these non-ideal behaviors.