Every stoichiometry calculation you've done so far predicts a theoretical yield — the maximum possible amount of product, assuming the reaction goes perfectly to completion with none of the material lost along the way. Real reactions rarely hit that number. When a chemist actually runs the reaction, weighs the real product, and gets an actual yield, it's almost always somewhat less than the theoretical prediction. Percentage yield is simply a way of expressing how close the real result came to the ideal calculation. Think of baking: a recipe that "makes 24 cookies" is a theoretical yield. If some dough sticks to the bowl and two cookies burn, you end up with 21, an actual yield of . Nothing about the recipe was wrong; real processes simply lose material. In a chemistry lab, a careful student synthesis often lands somewhere between about and , and the percent yield is a direct measure of both the reaction and the technique.
Percentage Yield Formula
- Theoretical yield is calculated from the balanced equation, always based on the limiting reagent (see the previous topic).
- Actual yield is a measured, experimental quantity — it comes from weighing the real product on a balance, not from a calculation.
The two yields must be in the same unit (both in grams, or both in moles) so the units cancel and leave a pure percentage. The theoretical yield comes from the full stoichiometry chain you already know: identify the limiting reagent, convert it to moles, apply the mole ratio, and convert to mass of product. Percent yield matters enormously in industry, because every percent lost is raw material paid for and thrown away. It matters even more for products made in several steps, such as medicines, because the losses multiply: the overall yield is the product of the individual step yields. Three steps at each deliver barely half the theoretical product, which is why chemists work hard to raise the yield of every single step.
Why Actual Yield Is Normally Lower
- Side reactions: some reactant is consumed by a competing reaction that doesn't produce the desired product.
- Incomplete reaction: the reaction doesn't fully finish reacting before it's stopped or measured.
- Loss during transfer or purification: some product sticks to glassware, is lost while filtering, or is removed while purifying the sample.
- Impure reactants: if the starting material is not 100% pure, less of the limiting reagent is really present than its mass suggests, so less product can form.
- Calculated on paper from the balanced equation
- Always based on the limiting reagent
- The maximum possible, assuming perfect conversion and no losses
- Unaffected by anything done in the lab
- Measured on a balance after the experiment
- Reduced by side reactions, incomplete reaction and handling losses
- Inflated by water or impurities left in the product
- Its sig figs come from the balance reading
| What happened in the lab | Effect on the measured product | Effect on percent yield |
|---|---|---|
| Product weighed before it was fully dry | extra water mass | too high (can exceed 100%) |
| Unreacted solid reactant mixed into the product | impurity adds mass | too high |
| Some product spilled or stuck to the filter paper | mass lost | too low |
| Reaction stopped before it was complete | less product formed | too low |
| A side reaction made a different product | less of the wanted product | too low |
Working Backward: Planning with Percent Yield
Once a process's typical percent yield is known, you can use it to predict or plan. Predicting is multiplication: a reaction with a yield and a theoretical yield of should give about . Planning is division, and it is where students most often slip. If you need of product and the yield is only , you must aim for a larger theoretical yield, , and then work back through the stoichiometry to find how much reactant to use. A quick reasonableness test catches the error: the planned theoretical yield must always be bigger than the amount you actually want, never smaller.