The kinetic theory of gases assumes that all collisions between gas molecules are completely elastic. So kinetic energy is conserved in collisions between molecules. Thus the average value of velocity remains constant for the gas. Pressure is caused by the change in momentum associated with collisions of gas molecules against the walls of the container. So as the average value of velocity remains constant it is safe to ignore the effects of the collisions between the molecules themselves, when calculating the pressure of the system. Is this reasoning correct?
If it's correct, PV=nRT successfully calculates the true pressure of the gas. Consider my thought experiment. Two ideals gases are sealed in a container. There would be some temperature and total pressure associated with the system. Now, if we can successfully ignore the effects of the collisions between the gas molecules themselves, then this system is equivalent to having the two gases separate, in similar containers ( they just add up, that's all) If so, The individual pressures of the gases are going to be equal to their partial pressure, which is a measure of how much a given gas contributes to the total pressure.
But the gases are in equilibrium are they not? So their pressures must equal the total pressure. Where does the additional pressure come from? Is there is something wrong in the above reasoning? Can somebody point out where?
Thanks for any help offered