Hope you don't mind, I fixed a superscript problem in your post. Thank you, BTW, for using superscripts and subscripts in your post. Most people don't use them and it makes it a pain to read posts without them.

OK, the first thing to do is take a look at the chemical reaction that is in equilibrium here:

Fe^{+3} + SCN^{-} --> FeSCN^{+2}

Now, I'm assuming that you've learned about the ICE charts. They are charts used to describe the initial concentration, the change in concentration, and the equilibrium concentration of all species present in solution. You fill in the appropriate data. From what you have given in your problem, the chart would look like this:

Fe^{+3} + SCN^{-} --> FeSCN^{+2}

I: 1.0 x 10^{-5} 2.0 x 10^{-6} ?

C:

E: 0.17

That's the data you've given. Now, are you sure your calibration curve isn't in units of 10^{-5} M or something like this? Because there's no way to get 0.17 moles from 0.00005 moles to begin with. I used to TA a lab that was almost the exact same thing, and people got hung up on that all the time.

Now, the change part is where you add variables. You know that the iron III and the thiocyanate are going to come together to form the iron thiocyanate, but you're not sure to what extent. And you know that for every mol of iron and thiocyanate that react, you get one mol of iron thiocyanate (just stoichiometry). This means that for every mol of iron that disappears on the left side of the equation, a mol of iron thiocyanate will appear on the right side. The same goes for thiocyanate. We arbitarily say that this change in mols is the variable X. So, we will in our ICE chart some more (and correct our concentration of FeSCN^{+2}:

Fe^{+3} + SCN^{-} --> FeSCN^{+2}

I: 1.0 x 10^{-5} 2.0 x 10^{-6} ?

C: -X -X X

E: 1.7 x 10^{-6}

Why did we put negative X on the left? That's because you start with only reactants, and the reaction must shift to the right in order to obtain equilibrium. Remember, this is the change that must occur to reach equilibrium. So you dump the two reactants together, and their concentration goes down as the concentration of the products goes up. Does that make sense?

OK, so, now you're asking, 'How do we find out the equilibrium concentrations of the iron and the thiocyanate?' Well, you can write them in terms of X as well.

Fe^{+3} + SCN^{-} --> FeSCN^{+2}

I: 1.0 x 10^{-5} 2.0 x 10^{-6} ?

C: -X -X X

E: 1.0 x 10^{-5} - X 2.0 x x 10^{-6} - X 1.7 x 10^{-6}

Now you can write an expression for K. K is given as concentration of products over concentration of reactants at equilibrium, and so it looks like this:

[ FeSCN^{+2} ]

K_{c} = [ Fe^{+3} ][ SCN^{-} ]

From your table, you can fill these values in:

[ 1.7 x 10 ^{-6} ]

K_{c} = [ 1.0 x 10^{-5} - X ][ 2.0 x 10^{-6} - X ]

Take a look at the ICE chart. Is there anywhere for the FeSCN^{+2} value to have come from other than the X? Nope. This means that the equilibrium mols of the iron thiocyanate *is* X. Now that you know that, this whole problem just becomes math. You can calculate all that you asked for, plus the equilibrium constant.

The only other thing to watch out for is that problem is asking for equilibrium concentrations, so once you get your # of mols, you'll have to divide by the number of milliliters of solution that you have in order to calculate the molarity of the chemicals. Take a stab at it and see how it turns out. Come on back here if you have any futher questions.