Sorry, there was something wrong with the original post - the question I cut & paste did not appear. Here is the question in full, I hope.

Find the equilibrium constant KC for O2 (g) + O(g) -> O3 (g) given that NO2 (g) -> NO(g) + O(g) has KC = 6.8E-49 mol/L and O3 (g) + NO(g) -> NO2 (g) + O2 (g) has KC = 5.8E-34. If the overall reaction is initially at equilibrium with [O2 (0)] = 8 mol m−3 and [O3 (0)] = 1.25 mol m−3, and .25 mol m−3 of O3 is added at constant temperature, what are the new final concentrations of O3 , O2 and O? Are these in accord with Le Chatelierâ€™s principle?

*NOTE: All the arrows -> mean that the reaction can proceed forward or backwards... I just don't know how to do a double-headed arrow.

I find it difficult to follow your question because of the layout. I think a solution may be clearer if you adopt the following layout and write the equations in the following order

NO

_{2} <-> NO + O K

_{c1 }= . some expression using the data you have ...

O

_{2} + O <-> O

_{3} K

_{c2} = . some expression using the data you have ...

NO + O

_{3} <-> NO

_{2} + O2 K

_{c3} = . some expression using the data you have ...

Then build an ICE table for

**each** and start by saying at equlibrium the change in equation 1 is due to a loss of "x" amount of NO

_{2}. What does equation 1 give you by applying ICE as suggested above?

Then do the same for equation 2 but try to link it to the change in equation 1.

Then do the same for equation 3 but try to link it to the change in equation 1 (and 2?).

Now you should have equations in "x". Try and eliminate things and find "x". I haven't actually waded through this but it

**might** work (no guarantees express or implied ... )

Interesting that you are given concentrations to work with when all the reactants are in the gas phase. I thought K

_{p} is more appropriate in such a situation but I

**guess** concentrations and Kc will work ...

Clive