Reaction Quotient

Quotient Equilibrium Constant

There is also Qc, which is the constant at a particular concentration that is not at equilibrium. You would calculate it the same as Kc but this is a comparison with Kc. It is used to determine which way the reaction will proceed at any one given time.

• If Q_c > K_cShift to the left
• Q_c < K_cShift to the Right
• Q_c = K_c Nothing has changed (still at equilibrium) 

*Same process for Q_p

The value of Q in relation to K serves as an index how the composition of the reaction system compares to that of the equilibrium state, and thus it indicates the direction in which any net reaction must proceed.

For example, if we combine the two reactants A and B at concentrations of 1 mol L^–1 each, the value of Q will be 0÷1=0. The only possible change is the conversion of some of these reactants into products. If instead our mixture consists only of the two products C and D, Q will be indeterminately large (1÷0) and the only possible change will be in the reverse direction.

It is easy to see (by simple application of the le Châtelier principle) that the ratio of Q/K immediately tells us whether, and in which direction, a net reaction will occur as the system moves toward its equilibrium state. A schematic view of this relationship is shown below:

Notice the double-ended arrow in the center. This is to remind you that when the reaction is at equilibrium, the forward and reverse reactions continue as always, but at equal rates, thus cancelling each other to yield a state of no net change. 

It is very important that you be able to work out these relations for yourself, not by memorizing them, but from the definitions of Q and K.