[Araconan]

I am currently attempting the following problem:

At a certain temperature K = 1.1 x 10³ L/mol for the following reaction:

Fe³⁺(aq) + SCN^-(aq) <-------> FeSCN²⁺(aq)

Calculate the concentrations of Fe³⁺, SCN^-, and FeSCN²⁺ at equilibrium if 0.02 moles of Fe(NO₃)₃ is added to 1L of 0.01M KSCN. (Neglect any volume change)

My attempt:

With the provided amounts of compounds, once placed in an aqueous solution, there will be a concentration of 0.02M Fe³⁺, and 0.1M SCN^-.

Constructing an ICe table:
                           Fe³⁺(aq) + SCN^-(aq) <-------> FeSCN²⁺(aq)
Initial:                 0.02M         0.1M                      0M

Change:              -x               -x                           +x

Equilibrium:        0.02 - x       0.1 - x                      x

The quadratic equation would then be:

x/(0.02-x)(0.1-x) = 1.1*10³

Simplifying to:

1100x² - 113.2x + 2.2 = 0

Solving for x, the two values would be

x1 = 0.0769
x2 = 0.026

However, both values can't be correct since I only have 0.02M of Fe³⁺. Even if I round 0.026 to 0.02, I would get 0M, which isn't what the answer states. (Answer being 2*10^-4)

So where did I go wrong?

Thank you in advance!

[Borek]

Equilibrium:        0.02 - x       0.1 - x                      x

Check your math.

[Araconan]

Ahh, that awkward moment when after double checking my work, I still make an error 
I see where I went wrong now, equation should've been 1100x² - 133x + 2.2 = 0

Thank you!