What percent yield measures
Percent yield is the fraction of the stoichiometric maximum you actually got out of the flask. It is the single number chemists use to judge how well a preparation ran, and it appears in every experimental write-up and every process report.
It has two parts, and only one of them is arithmetic. The denominator, the theoretical yield, is a calculation from the limiting reactant and the balanced equation. The numerator is a weighing, and it is only meaningful if the product is pure and dry. Most bad percent yields are bad numerators — crude material weighed before recrystallisation, or a product still holding solvent.
What the number does not do is diagnose. A yield of 62% tells you 38% of the possible product did not reach your balance; it does not say whether that is because the reaction stalled at equilibrium, because a side reaction consumed the substrate, or because you left the product in the mother liquor. Separating those requires more than one weighing, which is why careful workers record the crude mass as well as the purified one.
The ratio, and the conditions under which it is valid
Divide actual by theoretical and multiply by 100. The two masses must refer to the same product and be in the same units, and beyond that the arithmetic is trivial. Everything interesting is in the conditions.
The comparison works equally in mass or in moles. Because both figures convert through the same molar mass, the ratio is unchanged: 22.00 g of carbon dioxide against 44.01 g is 50%, and 0.500 mol against 1.000 mol is also 50%. Reporting in moles becomes useful when you are comparing steps of a synthesis whose products have very different molecular sizes.
The denominator must come from the limiting reactant. Computing a theoretical yield from a reagent that was in excess produces a ceiling the reaction could never reach, and the percent yield that follows is meaninglessly low. If you are not certain which reactant limited, settle it with the limiting reagent calculator first.
Above 100% is impossible. The theoretical yield is a ceiling derived from conservation of atoms, so exceeding it means one of the inputs is wrong. In descending order of likelihood: the product is not dry, the product is not pure, the equation is unbalanced, or the wrong limiting reactant was used. Dry the sample to constant mass before you look any further.
Worked example: an aspirin preparation
Suppose the stoichiometry allows 5.00 g of product and your dried, recrystallised solid weighs 4.20 g.
- Divide. 4.20 ÷ 5.00 = 0.840.
- Multiply by 100. The percent yield is 84.0%.
- Mass not recovered. 5.00 − 4.20 = 0.80 g, which is 16.0% of the ceiling.
Now interpret the 0.80 g. If the crude solid weighed 4.85 g before recrystallisation and 4.20 g after, then 0.65 g of the 0.80 g was lost in purification and only 0.15 g never formed. The reaction ran at roughly 97% of the ceiling; the purification cost 13 percentage points. Those are entirely different problems, and the single figure of 84% conceals which one you have.
Convert to moles if you want to compare across steps. Aspirin, C₉H₈O₄, has a molar mass of 180.159 g/mol, so 4.20 g is 0.02331 mol against a ceiling of 0.02776 mol — the same 84.0%, now in a form you can chain. A three-step sequence at 84%, 76% and 91% gives an overall yield of 0.84 × 0.76 × 0.91 = 58%, and that multiplication is only legitimate on mole-basis yields.
What counts as a good yield
There is no universal benchmark, because what is excellent for one reaction class is poor for another. The rough expectations practitioners work to are these: a clean inorganic precipitation or a simple acid–base salt formation should exceed 90%; a well-behaved single-step organic transformation typically lands between 70% and 90%; a difficult ring closure, a sterically hindered coupling or an enzymatic step may be counted a success at 40%. Treat these as rules of thumb, not standards — the honest benchmark for your reaction is the yield reported in the procedure you are following.
Pay attention to the compounding across steps. A linear five-step synthesis at 80% per step delivers 0.80⁵ = 32.8% overall. That arithmetic is why process chemists will spend months raising a single step from 80% to 92%, and why convergent routes — combining two branches late — beat linear ones for anything long.
Finally, separate yield from purity. A 95% yield of material that is 80% pure is worse than an 80% yield of material that is 99% pure, and the percent yield figure alone cannot tell them apart. Report both, and state how purity was established.
Reference: yields and what they imply
| Percent yield | Fraction lost | Typical reading | Overall after 3 such steps |
|---|---|---|---|
| 100% | 0% | Quantitative — verify the product is dry before believing it | 100.0% |
| 95% | 5% | Excellent; typical ceiling for a clean precipitation | 85.7% |
| 90% | 10% | Very good for a synthetic step | 72.9% |
| 80% | 20% | Good; a common working figure in organic synthesis | 51.2% |
| 70% | 30% | Acceptable; look at workup losses | 34.3% |
| 50% | 50% | Poor for a simple step, normal for a hard one | 12.5% |
| 30% | 70% | Investigate: side reactions or an unfavourable equilibrium | 2.7% |
| 10% | 90% | Usually a failed or misidentified reaction | 0.1% |
The last column is the yield cubed, and it is the reason step yields matter so much more in a long sequence than in a single preparation.
Where the missing mass usually goes
- Product left in the mother liquor. Nothing is perfectly insoluble. Recrystallisation always discards some product in the filtrate, and a second crop can recover a useful fraction of it.
- Incomplete reaction. An equilibrium stops short of consuming the limiting reactant. Longer reaction times, more heat, or removing a product can push it further.
- Side reactions. Substrate consumed to make something you did not want, which lowers selectivity without lowering conversion.
- Transfer losses. Material left on glassware, in the filter cake, or on the filter paper. Rinsing quantitatively costs a minute and often recovers several percent.
- Wrong theoretical yield. An unbalanced equation or a non-limiting reactant in the denominator makes the percentage wrong even when the chemistry went perfectly.
- Product still wet. This one inflates the yield rather than lowering it, and it is the single commonest cause of an impossible result above 100%.
Yield, conversion, selectivity and atom economy
Four measures describe the efficiency of a reaction and they are routinely confused. Conversion is the fraction of the limiting reactant that reacted. Selectivity is the fraction of the reacted material that became your product. Yield is their product — which is why a reaction can have 99% selectivity and still deliver 30% yield if only a third of the substrate reacted. Atom economy is different in kind: it is the fraction of all reactant mass that appears in the desired product, fixed by the reaction you chose rather than by how well you ran it.
In a research laboratory, yield is the number reported. In a plant, conversion and selectivity are often more useful, because unreacted feed can be recycled while material lost to side products cannot. A process deliberately run at 40% conversion with 98% selectivity and full recycle beats one at 95% conversion with 80% selectivity, despite the second having a much better single-pass yield.
A related term worth keeping distinct is percent recovery, used in purification rather than synthesis. It compares the mass of purified material with the mass of crude you started the purification from. It answers "how much did the recrystallisation cost me", not "how well did the reaction go", and quoting one for the other is a common slip in write-ups.
To build the denominator this page needs, start from the limiting reagent calculator and then the theoretical yield calculator. If you need to convert between species rather than judge a run, the mole-ratio calculator is the right tool.