There are actually two ways of approaching this problem.
You're dealing with a redox reaction in which ammonia,
Take a look at the
Keep this in mind, it will become very useful in a short while.
Now, you know that an 8.5-g sample of ammonia reacts to produce 4.5 g of nitric oxide. Before doing anything else, check to see whether or not the aforementioned mole ratio is respected.
Use ammonia and nitric oxide's molar masses to determine how many moles of each reacted/were produced
The number of moles of ammonia and nitric oxide are not equal. Your reaction produced less moles of nitric oxide than the number of moles of ammonia that reacted, which means that one of two things happened
- not all the ammonia reacted, i.e. oxygen acted as a limiting reagent;
- the percent yield of the reaction is not 100%.
Let's take the first scenario. If not all ammonia reacted, then oxygen must have acted as a limiting reagent. If you work backwards from the number of moles of nitric oxide produced, you have
This means that, out of the 8.5 g sample, only
of ammonia actually reacted.
In the second scenario, you assume that all the ammonia reacted, but that the reaction doesn't have a 100% yield.
If all the ammonia actually reacted, then the reaction should have produced
The mass of nitric oxide should have been
This is the reaction's theoretical yield. The actual yield is 4.5 g of nitric oxide produced. This means that the percent yield was