At 25 degrees C, the delta G naught for a certain reaction A < - -> B + 2C is 0. If the concentration of A, B, and C in the cell at 25 degrees C are all 10 mM, how does the delta G compare to the measurement taken at 1 M concentrations?

A. Delta G is greater than delta G naught, thus the reaction is spontaneous

B. Delta G is less than delta G naught, thus the reaction is spontaneous

C. Delta G is greater than delta G naught, thus the reaction is nonspontaneous

D. Delta G is less than delta G naught, thus the reaction is nonspontaneous

The answer is B and you can plug in the numbers into the equation, delta G = delta G naught + RT lnQ to get a negative delta G, which means it's spontaneous.

My question is two parts:

Part 1) from a conceptual perspective, why is delta G negative? If I change the concentration of both the reactants and products by the same amount shouldn't the reaction favor the products, since there is more products than reactants (3 products, which are B and 2C versus 1 reactant, A). I understand that this would be true for Keq<Q, which does not apply here, since we aren't given Keq (i.e. you don't know where equilibrium is).

__It might be helpful to see how someone goes through this problem conceptually without any equation and plugging in numbers.__

Part 2) How do we account for the change in delta G if we were to use 20 M (a non-decimal number) instead of 10 mM?