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Because when you decrease the pressure, the volume increases, which means that the molar concentrations of the reactants and products decrease. Since the molar concentrations of gases and aqueous solutions are inversely related to volume, increasing the volume decreases the molar concentration of the gases and aqueous solutions. The side with more moles of gas and aqueous solutions will experience a greater decrease, which will shift the equilibrium toward the side towards it.
A decrease in pressure corresponds to a higher volume (as pressure and volume are inversely proportional according to the ideal gas law, PV=nRT). So the reaction side with more moles of GAS is favored, as the higher moles of gas can occupy the larger value to return the reaction to equilibrium.
In retrospect by decreasing pressure you are increasing volume since P and V are inversely related to each other. By increasing the volume you essentially have more space for gas to move about, therefore, favoring the side with more moles of gas.
To summarize, you can think of this in two ways: thinking about it in terms of concentration as Victoria Zheng--2F explained above, or the quicker way of gases occupying areas naturally as CameronDis2K and Natalie Benitez 1E have said. Whatever floats your boat!
We did an example in class with actual numbers to show how decreasing the volume, increased concentration (n/V) and made equilibrium shift to the side with fewer gas moles. You can use the same numbers to show how increasing the volume makes equilibrium shift to the side with more gas moles.
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