In an equilibrium, why are partial pressures not affected by the coefficients of the balanced equation? It's probably easiest for me to explain in an example. For the following reaction: $$\ce{2NOBr(g) &lt;=&gt; 2NO(g) + Br2(g)}$$
At equilibrium, if the partial pressure of $\ce{NOBr}$ was 6 atm, the partial pressure of $\ce{NO}$ was 4 atm, and the partial pressure of $\ce{Br2}$ was 3 atm. Why is the equilibrium constant: $$Kp = \frac{3\cdot 4^2}{6^2}$$ instead of the individual pressures multiplied by their coefficients like in molar solubility: $$Kp = \frac{3\cdot \left(2\cdot 4\right)^2}{\left(2\cdot 6\right)^2}$$
One way of thinking about this is writing the original equation
$$\ce{2NOBr(g) &lt;=&gt; 2NO(g) + Br2(g)}$$
as
$$\ce{NOBr(g) + NOBr(g)  &lt;=&gt; NO(g) +NO(g) +  Br2(g)}$$
Now it should be clear you have to multiply, not add concentrations.