General Chemistry
posted by marthy on .
At a certain temperature, the equilibrium constant for the following chemical equation is 2.90. At this temperature, calculate the number of moles of NO2(g) that must be added to 2.53 mol of SO2(g) in order to form 1.10 mol of SO3(g) at equilibrium.
SO2(g)+ NO2(g) <> SO3(g) + NO(g)
I tried doing the ICE chart but there seems to be something wrong with my calculation. _

Try this
..........SO2 + NO2 ==> SO3 + NO
I.........2.53...0.......0.....0
C..........x....x.......1.1...1.1
E.........2.53x..x......1.1....1.1
Total mols = 2.53x_x_1,1+1.1 = 4.73x
XSO2 = (2.53x)/(4.73x)
XNO2 = x/(4.73x)
XSO3 = 1.1/(4.73x)
XNO = 1.1(4.73x)
Then partial pressures are
pSO2 = (2.53x)/(4.73x)]*Ptotal
pNO2 = [x/(4.73x)]*Ptotal
pSO3 = [1.1/(4.73x)]*Ptotal
pNO = [1.1/4.73x)]*Ptotal
I won't finish but here is what you do.
Substitute partial pressures from above into K expression. Solve for x = mols NO2 and you have it. Don't worry about all of this extraneous stuff; it cancels and you are left with a simple quadratic to solve.