aminapolutan Posted February 12, 2017 Report Share Posted February 12, 2017 Hi, I have a P1 question that I don't know how to solve. Any help would be welcome. Why do gases deviate from the ideal gas law at high pressure? A. Molecules have finite volume. B. Cohesive forces increase the volume from the ideal. C. Increasing pressure increases the temperature of the gas. D. Collisions between molecules occur more frequently as pressure increases. Reply Link to post Share on other sites More sharing options...
kw0573 Posted February 12, 2017 Report Share Posted February 12, 2017 I am taking a thermodynamics course and we were just talking about this. The 2 main assumptions of the ideal gas law is 1) molecules have negligible volume 2) there is no attraction/repulsion between molecules (or no potential energy). If the volume of gas is a significant fraction of the total volume (which is at tow temperature and high pressure), then you have to "subtract" the volume of molecules from the total volume to find the volume the gas occupy. When molecules are just a bit closer from ideal behaviour, the attractive forces (london dispersion forces) make the volume just a bit smaller than ideal, but when they get too close, electrostatic repulsion (electron clouds repel each other) takes over and make volume larger than ideal. If you are interested for more detail, take a look at the van der Waals equation of gas. I would choose A. B is false. (cohesive forces decrease volume, repelling forces increase volume) C. is true (but not explanatory) assuming constant volume but I just said volume changes and D is true in both ideal and non-ideal gas, and is not of concern. If anything, collisions between molecules and walls of container are more pertinent as those increases the pressure. Reply Link to post Share on other sites More sharing options...
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