\[Percentage Yield = \frac{ Actual Yield }{ Theoretical Yield } \times 100\]
or
\[Percentage Yield = \frac{ Actual Yield \times 100 }{ Theoretical Yield }\]
Rearranging, we can get:
\[Actual Yield \times 100 = (Percentage Yield)(Theoretical Yield)\]
which means that
\[Actual Yield = \frac{ (Percentage Yield)(Theoretical Yield) }{ 100 }\]
From this rearranged equation, let's consider what we know and what we don't know based on the information given to us for the reaction:
- The percentage yield is 84.1% (or simply 84.1 in this case).
- The theoretical yield is unknown.
However, based on the balanced chemical equation for the reaction:
\[4P + 5O _{2} \rightarrow 2P _{2}O _{5}\]
we can work out the number of moles, and thus the number of grams of P2O5 we would EXPECT to form (i.e. our theoretical yield), given that we have reacted 3.54 g of phosphorus. The phosphorus is the limiting reagent whilst the other reactant, oxygen gas, is in excess, meaning that the amount of phosphorus we start with, and its molar ratio with the product we are forming, we be the determining factor into how much P2O5 can possibly be synthesised.
Here's a tutorial @taramgrant0543664 put together on how to go about finding the theoretical yield for a product from the balanced equation. One you have found it (in grams remember!), simply plug it in alongside our percentage yield value into the rearranged equation above to find the actual yield. Hope that helped you out a bit! :)

http://openstudy.com/study#/updates/559c49b4e4b0f93dd7c278ad