Identify the limiting reactant in a reaction, calculate the theoretical yield it allows, and calculate percent yield by comparing theoretical yield to actual yield.
Real reactions rarely start with the exact amounts of each reactant needed to use everything up evenly. The limiting reactant is whichever reactant runs out first, and once it is gone, the reaction simply stops -- no matter how much of the other reactant (the excess reactant) is left over unused. The limiting reactant determines the maximum possible amount of product, called the theoretical yield.
To find the limiting reactant, calculate how much product COULD form from each reactant separately (as if it were the only one that mattered), using mass-to-mass stoichiometry for each. Whichever reactant produces the SMALLER amount of product is the limiting reactant, since it runs out first and caps the reaction there.
In an actual lab, you rarely recover 100% of the theoretical yield -- some product is lost to spills, incomplete reactions, or side reactions. Percent yield measures how efficient the real reaction was: percent yield = (actual yield / theoretical yield) x 100. A percent yield close to 100% signals a clean, efficient lab process; a low percent yield signals product was lost somewhere along the way.
Mastery looks like: They correctly identify the limiting reactant by comparing possible product from each reactant, and correctly calculate percent yield from theoretical and actual yield.
Common stumbles: Assuming the reactant with the smaller starting amount is automatically limiting without doing the stoichiometry (the mole ratio can flip that assumption); dividing theoretical by actual instead of actual by theoretical when finding percent yield.