States of matter

  • 0_1576003435894_SAVE_20191210_173611.jpg Ans is 3 explain please

  • @ankita-maurya

    An ideal gas is the assumption that:

    you have a large number of particles in the gas with a Boltzmann distribution
    the size of the particles is negligible, and
    there are no forces acting between the particles except for elastic collisions
    In a gas with very weak forces, such as hydrogen which has only van der Waals forces, these assumptions are quite reasonable. In a gas with strong intermolecular forces, such as ammonia which can form hydrogen-bonds as well as reglar dipole-dipole interactions, these assumptions no longer make as much sense and the gas deviates from 'ideal' behaviour.

    Molecules of CO2 do not have an overall dipole so it only has van der Waals forces acting on it, but it is much larger and has more electrons than hydrogen molecules so the van der Waals forces are stronger. When you begin to compress the gas the molecules are compressed within range where these forces attract the molecules closer together and the volume is smaller than expected (your negative deviation from ideal gas).

    Eventually there comes a point where this extra attraction is overcome by repulsive forces, because the particles of all gases do have size and can't be compressed infinitely. At the point the gas occupies a larger volume than you would expect for the pressure (the positive deviation from the ideal gas shown by all gases at high pressure).

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