Real Gases have lower pressure but the same kinetic energy?

[sniderwes]

So if we have a gas at a temperature T, these things are supposed to be true:

1) Regardless of if it is real/ideal, the KE is supposed to be the same at T
2) If the gas were real there would be lower pressure because of intermolecular forces, reducing collision with the walls of the container

But don't these two concepts sort of contract each other? With 2 I think of the velocity of the real gas being decreased due to intermolecular forces pulling the molecules towards each other, thus causing reduced pressure. But wouldn't this cause decreased KE for the real gas as well? How can KE stay the same with pressure reduced?

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[DrknoSDN]

In a real gas the molecules that strike the walls of the container are slowed down due to the attractive forces between the molecules prior to striking the container. This lowers the pressure compared to an ideal gas.
Real gas molecules also occupy space and have a volume where ideal gases do not. Real gas molecules repel one another when they are packed closely together due to normal repulsion between atoms.

If you look at the real gas equation there is a B (bouncy -repulsive) factor, and an A (attractive pulled together) factor that modifies the expression.
From a conceptual standpoint, A and B are what allows temperature to still be related to average kinetic energy for real and ideal gases.
It's also possible to explain it mathematically but the question seemed to be conceptual.

Possibly useful website:
http://www.chemguide.co.uk/physical/kt/realgases.html