chemistry help please?

Question

The equilibrium constant for the reaction below is Kc = 9.1 10-6 at 298 K.

2 Fe3+(aq) + Hg22+(aq) <--> 2 Fe2+(aq) + 2 Hg2+(aq)

Calculate the change in G when [Fe3+] = 0.29 M, [(Hg2)2+]= 0.029 M, [Fe2+] = 0.028 M, and [Hg2+] = 0.040 M.

In which direction will the reaction proceed to achieve equilibrium? (right or left?)

Answer 1

ΔG0= -RTlnKc= -(1.987*10^-3)*298* ln(9.1*10^-6)= 6.87 Kcal/mole

ΔG=ΔG0 + RTln([Fe+2]^2 *[Hg+2]^2 /([Fe+3]^2 * [Hg22+]))=
6.87 + (1.987*10^-3)*298* ln( (0.028^2)*(0.040^2) /((0.29^2)*0.029) )=-4.48 kcal/mole<0 therefore it will proceed to the right.

Or if you prefer ([Fe+2]^2 *[Hg+2]^2 /([Fe+3]^2 * [Hg22+])= 5.14*10^-4>Kc so it will proceed to the right

You can use the calculator for equilibium reactions at http://www.changbioscience.com/biochem/keq.html
to make your life easier.