if the pressure was increased in this reaction :
2NO(g)+ O2 =====> 2NO2
<====
wouldnt the product side be favored?
Please tell me if I'm right or wrong
see above.
To determine if the product side would be favored when the pressure is increased in a reaction, we need to analyze the stoichiometry of the reaction and consider Le Chatelier's principle.
In the balanced equation you provided:
2NO(g) + O2 ⇌ 2NO2
Increasing the pressure means increasing the concentration of the reactants or products in the reaction vessel. In this case, since there are no coefficients denoting the number of moles of gases, we cannot definitively conclude how the pressure change would affect the equilibrium.
However, we can apply Le Chatelier's principle to make an educated prediction. According to Le Chatelier's principle, if the pressure on a system at equilibrium is increased, the system will shift in the direction that reduces the total number of moles of gas. Thus, if the reaction produced fewer moles of gas on the product side, the product side would be favored and the equilibrium position would shift to the right.
In this case, since the reaction produces 2 moles of NO2 gas from 3 moles of gaseous reactants (2NO + O2), the forward reaction (2NO(g) + O2 ⇌ 2NO2(g)) actually increases the total number of moles of gas. Therefore, increasing the pressure would not favor the product side of the reaction.
So, in summary, you are incorrect in your assumption. Increasing the pressure would not favor the product side in this specific reaction.