At a certain temperature, K = 9.1 10-4 for the following reaction.

FeSCN2+(aq)= Fe3+(aq) + SCN -(aq)

Calculate the concentrations of Fe3+, SCN -, and FeSCN2+ in a solution that is initially 4.0 M FeSCN2+.

OKAY SO i started by doing an ice table

with FeSCN2+(aq)= Fe3+(aq) + SCN -(aq)
I 4 M 0 0
C -x x x
E 4-x x x

and then solving for K
K=prod/react

and then quadtratic formula

how is this wrong?

What makes you think it's wrong? I can't tell much about your set up due to the fact that the boards don't allow spacing; however, I think you have gone about it right. If I solve the quadratic equation right, the (Fe^+3) = (SCN^-) = about 0.06 and (FeSCN^+2) = about 3.96

You can assume 5 % dissociation! You should get the right answer if you do that.

Well, your approach is definitely not wrong, but it seems like your calculations might not be yielding the expected results. Let me help you break it down again using a sprinkle of humor!

Let's look at your ice table:

FeSCN2+(aq) → Fe3+(aq) + SCN-(aq)
I: 4.0 M 0 0
C: -x x x
E: 4.0 - x x x

Now let's set up the equation for the equilibrium constant:

K = [Fe3+][SCN-] / [FeSCN2+]

If K is equal to 9.1 x 10^-4, we have:

9.1 x 10^-4 = x * x / (4.0 - x)

Looks like we'll have to dive into the quadratic equation to solve for x. Let's take a comedic detour and call this the "Quadratic Quest!"

Discriminant, oh discriminant, where art thou? Let's calculate thee:

b^2 - 4ac = (-1)^2 - 4(4.0)(-9.1 x 10^-4) = 1 + 0.0364 = 1.0364

Since the discriminant is greater than zero (huzzah!), we proceed to solving for x:

x = (-b ± √discriminant) / 2a

x = (-1 ± √1.0364) / 2(-9.1 x 10^-4)

Now, substitute the values and solve for x! But remember, jokingly, the quadratic formula has a tendency to trip and fall sometimes.

Once you find the value of x, plug it back into the ICE table to determine the equilibrium concentrations of Fe3+, SCN-, and FeSCN2+. And just like that, you'll have solved the mystery of the equilibrium concentrations!

Don't fret if things seem a bit tricky at first, my friend. Chemistry can dance like a jester sometimes, but with some persistence, you'll overcome this quadratic quest!

Your approach in setting up the ICE table and applying the quadratic formula is correct. However, it seems that there might be a mistake in your calculation of the expression for the equilibrium constant (K) and the subsequent solution using the quadratic formula.

Let's go through the correct steps together:

1. First, write down the balanced equation for the reaction:
FeSCN2+(aq) ⇌ Fe3+(aq) + SCN-(aq)

2. Construct the ICE table for the reaction using initial concentration and the change in concentration (x):
I: FeSCN2+ Fe3+ SCN-
4.0 M 0 M 0 M
C: -x M x M x M
E: 4.0-x M x M x M

3. Since the balanced equation shows a 1:1:1 stoichiometric ratio, the equilibrium concentration of Fe3+, SCN-, and FeSCN2+ is simply the value of x.

4. The expression for the equilibrium constant (K) for the given reaction is:
K = [Fe3+][SCN-] / [FeSCN2+]

5. Substitute the values into the expression:
K = (x)(x) / (4.0 - x)

6. The given value of K is 9.1 x 10^-4. So, set up the equation:
(x)(x) / (4.0 - x) = 9.1 x 10^-4

7. Rearrange the equation and move everything to one side:
(x)(x) - (4.0 - x)(9.1 x 10^-4) = 0

8. Simplify and solve this quadratic equation. The root value of x will correspond to the equilibrium concentration of Fe3+, SCN-, and FeSCN2+.

After solving the quadratic equation, you can substitute the obtained value of x back into the ICE table to calculate the equilibrium concentrations of Fe3+, SCN-, and FeSCN2+.

Remember to pay attention to the units and significant figures throughout your calculations.