I'm going to repost this just to clear up the difference between K and Q:
"If I'm wrong about any of this I'm sorry, but I'll give it a go. Hopefully someone will correct me if I'm wrong. K basically tells you the ratio of products over reactants when a system is at equilibrium. Q tells you the ratio of products over reactants whether or not the system is at eq. So if Q = K then the reaction is at equilibrium. If K > Q the there will be a rightward shift in the reaction rate (reaction favor production of products). A way to remember this is: K --> Q. See how you can make an arrow? Lastly, if K < Q there will be a leftward shift.
Now tying this into Gibb's we have delta G = delta G* + RTlnQ and delta G* = -RTlnK. Let's start with delta G*. This basically tells you the value of delta G* when the reaction is at equilibrium. Remember, although the * sign means standard conditions, we can have std conditions at any temp, not just 298k. With the second equation, that basically relates the delta G of the rxn at equilibrium to Q, or the ratio of products over reactants of the given. If the Q = K, then delta G = 0 which tells us that the rxn is at equilibrium. Another way to look at it is delta G = -RTlnK + RT lnQ. So when Q = K, the two will cancel each other out and you'll get zero. Remember though the delta G* tells you the rxn is spontaneous at equilibrium, that doesn't mean the delta G value will be negative or spontaneous as well. The spontaneity of a rxn depends on the starting concentrations represented by the Q.
Well I hope that helps some. This is was my amount of studying today since I took today off. It felt good to look up stuff and teach myself, while helping someone else. 🙂
Edit: Oh yeah, you can think of temperature in delta G equation as giving a greater priority to the starting concentrations."
About the cath/an question, I'll get to that sometime tomorrow if no one has responded.
Edit 2: Did you still need that explanation or have you found the information?
