Please explain how to solve for equilibrium concentrations? From week 1 homework

[805751553]

The reaction

N2O4↽−−⇀2NO2

is allowed to reach equilibrium in a chloroform solution at 25 ∘C . The equilibrium concentrations are 0.424 mol/L N2O4 and 2.18 mol/L NO2 .
Kc=11.2


Calculate the equilibrium concentrations of N2O4 and NO2 after the extra 1.00 mol NO2 is added to 1.00 L of solution.

I'm not quite sure how to proceed with the second part of the problem can someone please explain the steps to proceed?

[Nicole Li 2B]

We would solve this problem using the ICE table method. For the initial, we would put the given equilibrium concentrations of N2O4 and NO2. In the "change" row of the ICE table, you would add 1M of NO2. Since the concentration of products is being increased, we know that there must also be an increase in reactants to keep the equilibrium constant the same. Thus, we would write "+x" in the "change" row for N2O4.

You can then put the values of the "E" row into the equilibrium constant expression, equate the expression to Kc (11.2), and solve for x.

I hope this helps!