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I understand that reaction quotient (Q) measures the relative amounts of products and reactants at any time in the reaction at any particular point and equilibrium constant (K) describes the reaction at equilibrium. I also know that K and Q can be compared so that we can tell which way the reaction is shifting. However, in terms of the Gibbs Free Energy equation, why do we use K and Q? Can someone explain this three-way relationship for these three variables conceptually?
Perhaps this question can be revisited after we have covered this portion of the Gibb's Free Energy topic, but Dr. Lavelle actually hasn't covered those versions of the Gibb's Free Energy equation yet. The only one he has gone over so far has been the deltaG= deltaH - T*deltaS.
When the reaction is not at equilibrium, deltaG = deltaG(naught) + RTlnQ. There is a difference between deltaG and deltaG(naught). When the system is at equilibrium then delta G would equal 0 and the equation would be deltaG(naught)=-RTlnK. So depending on what state the system is, depends on what equation would be used.
In equilibrium, the activities constant Q is equivalent to the equilibrium constant K, that is why, when a reaction is at equilibrium, when can state that Q=K and replace Q in the equation delta G= -RTlnQ with K: delta G=-RTlnK. I hope this helps!
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