Reaction quotient, Q

Q = [C]c[D]d
        [A]a[B]b

The reaction quotient, Q, for the general reaction is given as follows:
aA (g) + bB (g)  cC (aq)  + dD (g)

The reaction quotient has the same form as the equilibrium constant, but it involves specific values that are not necessarily equilibrium concentrations. If they are equilibrium concentrations, then Q = Kc. The concept of the reaction quotient is very useful. We can compare the magnitude of Q with that of Kc for a reaction under given conditions to decide whether the forward or the reverse reaction must occur to a greater extent to establish equilibrium.

If at any time Q < Kc, the forward reaction must occur to a greater extent than the reverse reaction for equilibrium to be established. This is because when Q < Kc, the numerator of Q is too small and the denominator is too large. To increase the numerator and to reduce the denominator, A and B must react to produce C and D. Conversely, if Q>  Kc, the reverse reaction must occur to a greater extent than the forward reaction for equilibrium to be reached. When the value of Q reaches the value of Kc, the system is at equilibrium, so no further net reaction occurs.

At a very high temperature, Kc = 65.0 for the following reaction.
  			2HI (g)  H2 (g) + I2 (g)
The following concentrations were detected in a mixture. Is the system at equilibrium? If not, in which direction must the reaction proceed for equilibrium to be established?
[HI] =  0.500 mol / L,	 [H2] = 2.80 mol / 		[I2] = 3.40 mol / L

Q =     [H2][I2]      =       (2.80) (3.40)   =  38.08
            [HI]2                      (0.500)2

But Kc = 65.0, so Q <  Kc. The system is not at equilibrium. For equilibrium to be established, the value of Q must increase until it equals Kc. This can occur only if the numerator increases and the denominator decreases. Thus, the forward (left-to-right) reaction must occur to a greater extent than the reverse reaction; that is, some HI must react to form more H2 and I2 to reach
equilibrium.