The ideal gas equation (or the van der Waals equation, for this matter) states that holding volume and temperature constant, increasing the number of moles of particles in the container will increase the total pressure exerted by the gas mixture.
That’s all well and good, but when it comes to looking at chemical reactions and their equilibria, Le Châtelier’s Principle states that when a gas mixture is pressurized, the equilibrium shifts to the side of the reaction that produces a lesser total number of moles. This would mean less particles, and therefore less pressure, as I see from the ideal gas equation.
What I do not understand is why the mass of each molecule does not affect the pressure exerted by the gas. If you have heavier molecules, they would each have more momentum, and would exert more force on the container’s walls—i.e. more pressure. So is this not a factor when considering Le Châtelier’s Principle? The equilibrium shifts toward the formation of more particles, but they are each lighter and exert less force. Why is this favorable over having less but heavier particles?