We’re being asked to calculate the equilibrium constant of HA if a 0.0538 M solution is 3.57% ionized.
Recall that the percent ionization is given by:
The dissociation is as follows:
N2O4(g) ⇌ 2NO2(g)
From this, we can construct an ICE table. Remember that liquids are ignored in the ICE table and Ka expression.
The Kc expression for is:
Note that each concentration is raised by the stoichiometric coefficient: [NO2] is raised to 2 and [N2O4] is raised to 1.
Exactly 1.0 mol N2O4 is placed in an empty 1.0-L container and allowed to reach equilibrium described by the equation
N2O4(g) ⇌ 2NO2(g).
If at equilibrium N2O4 is 28.0% dissociated, what is the value of the equilibrium constant, Kc, for the reaction under these conditions?