Question #d78b6
1 Answer
Explanation:
You know that a certain temperature, you have
#color(red)(2)"NO"_ (2(g)) rightleftharpoons "N"_ 2"O"_ (4(g))#
with an equilibrium constant of
#K_(eq) = 1.15#
Now, by definition, the equilibrium constant for this reaction takes this form--don't forget that the expression of the equilibrium constant takes into account the stoichiometric coefficients of the chemical species involved in the reaction!
#K_(eq) = (["N"_2"O"_4])/(["NO"_2]^color(red)(2))#
This means that you have
#["N"_ 2"O"_ 4] = K_(eq) * ["NO"_ 2]^color(red)(2)#
Plug in your values to find--since you didn't get units for
#["N"_ 2"O"_ 4] = 1.15 * (0.05)^color(red)(2) = color(darkgreen)(ul(color(black)("0.003 mol L"^(-1))))#
The answer is rounded to one significant figure, the number of sig figs you have for the equilibrium concentration of nitrogen dioxide.