equilibrium constant problem. Please Help.?

Question

Using the value of Ksp for Ag2S, Ka1 and Ka2 for H2s, and Kf=1.1x10 to the fifth for AgCl2, calculate the equilibrium constant for the following reaction:

Ag2S+4Cl+2H<-----> 2 AgCl2+H2S

Can anyone help me with this?

Answer 1

First of all you are not providing the values
Second, the reaction is not properly written: AgCl2 doesn't exist; Ag has +1charge. So I guess you mean that a complex is formed of the type AgCl2(-). Also is Ag2S solid or in solution? If it is the former which is the more likely case, then

Ag2S (s) + 4Cl-(aq)+ 2H+ (aq) <=> 2AgCl2(-)(aq) + H2S (aq)

In the K expression solids are left out, thus
K= [AgCl2(-)]^2 * [H2S] / [H+]^2 *[Cl-]^4 =>
K= ( [H2S]/[H+]^2 ) * ([AgCl2(-)]/[Cl-]^2)^2 (1)

For H2S:
H2S <=> HS- + H+ with Ka1=[HS-][H+] / [H2S]
HS- <=> S(-2) + H+ with Ka2 =[H+][S(-2)] / [HS-]

Thus Ka1*Ka2= [S(-2)]*[H+]^2 / [H2S] =>
[H2S]/[H+]^2 = [S(-2)] / Ka1Ka2 (2)

For the complex formation
Ag+ +2Cl- <=> AgCl2(-) with Kf= [AgCl2(-)]/[Cl-]^2[Ag+] =>
[AgCl2(-)] / [Cl-]^2 =Kf*[Ag+] (3)

Substitute the ratios in (1) using (2) and (3)

K= ( [S(-2)] / Ka1Ka2 ) *(Kf*[Ag+] )^2 =>
K= ( Kf^2 / (Ka1Ka2)) *[Ag+]^2*[S(-2)]

but KspAg2S= [Ag+]^2[S(-2)] so the equation becomes

K= Ksp*Kf^2 / (Ka1*Ka2)

substitute the values and calculate K

Answer 2

This is too much to ask! You haven't given a properly balanced equation, no Ksp, Ka values, and there is no such thing as AgCl2 it's AgCl.

Ag2S + 2HCl --> 2AgCl + H2S

maybe that will help you get started.