Combustion reactions
I'm having a problem with an A.P. Chemistry problem. It doesn't make sense because the equation starts with an unknown amount of Carbon, Hydrogen, and Nitrogen. This compoud combusts and forms Carbon dioxide and water, but the Nitrogen is gone. I'm not sure why this happens (or even it's supposed to, and this is just a typo), so if any one can help me I would love you forever. The problem is:
A compound contains only C, H, and N. Combustion of 35.0 mg of the compund produces 33.5 mg of CO2 and 41.1 mg of H2O. What is the empirical formula of this compound?
First, I balnced the equation as far as I could: ?CHN+?O2 yields ?CO2+?H2O.
Second, I obtained the mass of O2 by subtracting the mass of the products by the mass of the known reactant: (41.1mg+33.5mg)-35.0mg=39.6mg
Third, I changed the milligrams into grams, ending up with 0.0335g CO2, 0.0411g H2O, 0.0396g O2, and 0.0350g ?CHN.
Fourth, I converted to moles and got: 0.0335g CO2/44.01=7.61E-4 mol CO2.
0.0411g H2O/18.02=0.00228 mol H2O.
0.0396g O2/32.00=0.0012375 mol O2.
Fifth, I divided each amount of moles by the smallest amount.
7.61E-4 mol CO2/7.61E-4 mol CO2=1CO2
0.00228 mol H2O/7.61E-4 mol CO2=3H20
0.0012375 mol O2/7.61E-4 mol CO2=1.62615O2
So, I multiplied each answer by 8 to get 1.62615 to a whole number.
?CHN+13O2+8CO2+24H2O
But, as you can tell, that doesn't balance and N is still missing. SO if anyone has any suggestions, please help. Thanks in advance.