1. The problem statement, all variables and given/known data When the supply of oxygen is limited, iron metal reacts with oxygen to produce a mixture of FeO and Fe2O3. In a certain experiment, 20.00 g iron metal was reacted with 11.20 g oxygen gas. After the experiment, the iron was totally consumed, and 3.24 g oxygen gas remained. Calculate the amounts of FeO and Fe2O3 formed in this experiment. 2. Relevant equations 3. The attempt at a solution 3 Fe + 2 O2 ------> FeO + Fe2O3 20g 0,12mol 0,12mol 0,36mol As each 3 mol Fe produces 1 mol FeO and Fe2O3 , just by multplying moles to their molar masses it gives: 8,64g FeO and 19,2g Fe2O3. In my solution no need for O2 in calculations. But, solution manual disagrees :( Firts of all it states each reaction separately (WHY?) Then it calculates O2's mole by subtracting 3.24 from 11.2 and by dividing it to its molar mass and then: " Let’s assume x moles of Fe reacts to form x moles of FeO. Then 0.3581 – x, the remaining moles of Fe, reacts to form Fe2O3. Balancing the two equations in terms of x:... " etc etc But why?? What's wrong with my calculation? I assumed when it says "3.24 g oxygen gas remained" as O2 is surplus, excess. Did I assume wrong?