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Bugsy23
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Why does the ideal gas state equation fail at high pressures for real gases? I think it has something to do with the forces between the molecules in the gas, but I'm not sure the exact reason
The ideal gas state equation, PV = nRT, assumes that gas molecules have no volume and do not interact with each other. However, at high pressures, the volume of gas molecules becomes significant and the attractive forces between molecules start to affect their behavior. This leads to deviations from the ideal gas behavior and the failure of the ideal gas state equation.
Intermolecular forces, such as London dispersion forces and dipole-dipole interactions, become more significant at high pressures. These forces cause gas molecules to attract each other and take up space, leading to a decrease in the volume available for the gas molecules to move around. This results in a decrease in pressure, which is not accounted for in the ideal gas state equation.
No, the ideal gas state equation is only valid at low pressures and high temperatures, where the volume and intermolecular forces of gas molecules are negligible. At high pressures, the equation fails to accurately predict the behavior of real gases and should not be used.
At high pressures, real gases deviate from ideal gas behavior by exhibiting lower volumes and lower pressures than predicted by the ideal gas state equation. This is due to the effects of intermolecular forces and the finite size of gas molecules, which are not taken into account by the ideal gas state equation.
In addition to intermolecular forces, real gases can also deviate from ideal gas behavior at high pressures due to non-ideal behavior, such as chemical reactions between gas molecules or phase changes. These factors can also affect the volume and pressure of gas molecules and lead to deviations from the ideal gas state equation.