Question

Consider the following:

A student mixes 5.00 mL 2.00 × 10−3 M Fe(NO3)3 with 3.00 mL 2.00 × 10−3 M KSCN. She finds that

in the equilibrium mixture the concentration of FeSCN2+ is 1.28 × 10−4 M. Find Kc for the reaction

Fe3+(aq) + SCN−(aq) ↔ FeSCN2+(aq).

a. What is the initial concentration of Fe3+ in the reaction mixture?

[Fe3+] = ___ x 10-3 M

b. What is the initial concentration of SCN- in the reaction mixture?

[SCN-] = ___ x 10-4 M

c. We can construct an ICE table to help us organize the concentrations of the reactants and products and help us to calculate the equilibrium constant. Based on the first two calculations and the equilibrium concentration of FeSCN2+ ([FeSCN2+] = 0.000128 M), the ICE looks as follows:

	[Fe3+]	[SCN-]	[FeSCN2+]
Initial	0.00125 M	0.00075 M	0
Change	-x	-x	+x
Equilibrium	0.00125 M - x	0.00075 M - x	0.000128 M

Based on this, what is the value of "x"?

x = ___ M

d.   What is the equilibrium concentration of Fe3+ And SCN-??

[Fe3+] = ___ M [SCN-] = ___ M

e.   What is the value of the equilibrium constant for the reaction?

Kc = ___ M-1

Answer 1