Question

A solution is prepared that contains 0.0200 M AgNO# and 0.150 M KSCN. What is the equilibrium concentration of Ag+ ion in this solution? The Kf for Ag(SCN)4^3+ is 1.21x10^10

I'm not even sure how to set up this problem? Would I need to do an ice box?

Ag(SCN) --> Ag+ + SCN

Answer 1

An ICE chart makes it easier, I think, to see. You have written the equation for the dissociation constant for the complex, the number given is a formation constant; therefore, I will rewrite it as (and note that the charge on the complex ion is 3- and not 3+.) The easier way to solve this problem is to assume that the reaction proceeds far to the right; i.e., to completion. That goes like this (and with a large K like this that isn't a bad assumption).
.........Ag^+ + 4SCN^- ==> Ag[(SCN)4]^3-
I......0.02....0.150M........0
C......0.02....-4*0.02.....0.02
E.......0.......0.07.......0.02
--------------------------------
Then use the E line for a new ICE chart with the E values as the initial (I) values and to the reverse; i.e., Ag[(SCN)]^3- dissociating to Ag^+ and SCN. That will look like this (using the same set up as above but thinking in reverse).
I........0.......0.07.........0.02
C........x........4x...........-x
E........x.....0.07+4x........0.02-x

Then set up the Kf expression, substitute the E line into the expression and solve for x = (Ag^+). For the first go-round I would assume 0.07+4x = 0.07 and 0.02-x = 0.02.
Kf = [Ag[(SCN)4]^3-/(Ag^+)(SCN^-)^4
Post your work if you get stuck.