Find the percent yield, actual yield or theoretical yield of any reaction, enter the two values you have and the third is calculated as you type.
Percent yield
Step by step breakdown
Percent yield equals actual yield divided by theoretical yield, multiplied by 100. The calculator above applies that formula, converts your units, and rates the result against the standard quality bands. A reaction that returns 4.0 g of product where the equation allowed 5.0 g ran at 80%.
01. Definition
Percent yield is the ratio of the actual yield of a chemical reaction to its theoretical yield, expressed as a percentage. It answers a single question: of all the product the balanced equation said you could make, how much did you actually isolate?
Chemists abbreviate it percent yield (Yp). A Yp of 85% means 85 of every 100 possible grams of product ended up on the balance, while 15 were lost to incomplete conversion, competing side reactions, or the mechanical losses of filtration and transfer. The actual yield is also called the experimental yield, and the two terms are interchangeable. That is the whole percent yield definition.
Because it compresses everything that happened in the flask into a single number, the figure is the standard measure of reaction efficiency, a reaction success rate that captures both how completely the chemistry proceeded and how much product recovery survived the workup. It is the one number chemists reach for when comparing chemical efficiency between routes, because it normalises reaction output against the ceiling the equation set. A chemical yield of 0% means nothing usable was isolated. 100% is the stoichiometric limit that no real procedure reaches.
The figure is dimensionless. Because it divides one mass by another mass, or one mole count by another mole count, the units cancel and only a percentage survives. That is why a reaction yield can be compared directly between a milligram-scale medicinal chemistry step and a tonne-scale industrial process.
Percent yield is not the same as percent purity, which measures how much of an isolated sample is genuinely the target compound, nor percent recovery, which measures how much material survives a purification step. Both distinctions are covered further down this page.
02. Formula
In words: the percent yield equals the actual yield divided by the theoretical yield, multiplied by one hundred. That single sentence is the whole method, and every other calculation on this page is a rearrangement of the same percent yield equation.
Both quantities must share the same unit before you divide: grams with grams, or moles with moles. Mixing a mass with an amount of substance means converting one of them first, which needs the molar mass of the product. The grams to moles converter moves between the two. The calculator at the top of this page performs that conversion for you, and only asks for a molar mass when your two units genuinely span both dimensions.
Rearranged for the other two unknowns, the same relationship gives actual = (Yp ÷ 100) × theoretical and theoretical = actual ÷ (Yp ÷ 100). Switching the tab on the calculator applies whichever form you need.
03. Method
Write the reaction and balance it, because every mole ratio you use afterwards comes from those coefficients. See the equation balancer.
Convert each reagent to moles and find the one that runs out first, since it alone sets the ceiling. The limiting reactant calculator does this comparison.
Convert the limiting reactant to moles, apply the mole ratio, then multiply by the product's molar mass to return to grams. Full worked method on the theoretical yield calculator.
Isolate, purify and dry your product to constant mass, then weigh it, because any trapped solvent will corrupt the result.
Divide the actual yield by the theoretical yield and multiply by 100 to get the percent yield.
The order matters. Skipping straight to step five with an unbalanced equation, or calculating the theoretical yield from a reagent that was present in excess, produces a number that looks reasonable but is meaningless. Steps one to three are where almost every incorrect percent yield actually originates.
Step four is where lab technique accuracy decides the result. Any experimental error analysis of a yield calculation starts at the balance: a tare that drifted, a wet filter cake, or product left in the reaction mixture will all shift the numerator without any chemistry having gone wrong. Weigh on a calibrated balance, dry to constant mass, and record the figure before the sample can pick up atmospheric moisture.
04. Comparison
The two inputs to the formula come from completely different places: one is derived on paper, the other is read off a balance. Confusing them is the most common source of a wrong percent yield.
Calculated
The maximum mass of product a reaction could possibly produce. It is derived entirely on paper from the balanced equation and the limiting reactant: convert the limiting reactant to moles, apply the mole ratio, convert back to mass. It assumes every molecule reacts and nothing is lost, so it is a ceiling rather than a prediction, and no experiment will ever beat it.
Measured
The mass of product you genuinely end up with, also called the experimental yield. It can only be obtained experimentally: run the reaction, isolate the product, apply purification techniques such as recrystallisation, filtration or column chromatography, dry it to constant mass, and weigh it. Because it reflects everything that really happened in the flask, it is always the smaller of the two figures.
In practice the actual yield is expected to be less than or equal to the theoretical yield, never more. The theoretical yield describes a perfect reaction with perfect recovery, and no real procedure achieves both, so any measured value that exceeds it is a signal that the sample still contains solvent or impurities rather than evidence of an unusually good reaction.
05. Benchmarks
A percent yield above 90% is considered excellent. Below that, the conventional bands used in preparative chemistry run from very good down to poor, as set out in the table.
| Yield range | Rating |
|---|---|
| ≥ 90% | Excellent |
| 80 – 89% | Very good |
| 70 – 79% | Good |
| 50 – 69% | Fair |
| < 50% | Poor |
As a working rule of thumb, anything at or above 70% is generally treated as an acceptable yield for teaching and preparative work, which is why 70% appears so often as a target in lab manuals.
These bands are a general guide rather than a universal standard. What counts as good depends heavily on the difficulty of the reaction: a single-step precipitation that returns 75% would be disappointing, while the same figure across a demanding multi-step synthesis with a sensitive intermediate would be a strong outcome. Always judge a yield against comparable literature procedures.
06. Causes
Almost every reaction lands under 100%, and the handful that appear to exceed it are measurement artefacts. The causes split cleanly into two groups.
Where the mass goes
Detailed diagnosis and fixes for each: common lab errors.
Why the balance lies
Both have the same remedy: dry the sample to constant mass, weighing repeatedly until the reading stops falling, and then re-weigh. A true percent yield will always settle below 100%. A worked case: a product weighing 5.31 g against a 4.35 g theoretical yield returns 122.1%. After proper drying the same sample came to 4.00 g, a real yield of 92.0%.
Four variables account for most of the spread between two runs of the same procedure. Temperature and pressure set the reaction kinetics and, for a reversible reaction, the equilibrium position, so pushing them can raise conversion or, for an exothermic equilibrium, quietly lower the ceiling. Reaction time decides whether the mixture actually reached completion. Concentration and solvent choice govern how fast the reactants meet and how much product stays dissolved. And a catalyst changes the rate and the selectivity of the pathway without moving the equilibrium at all. Getting the reaction conditions right is what separates a reproducible 70% from a run-to-run scatter.
Yield optimization works on two fronts: form more product, then lose less of it. Driving the reaction further means using an excess of the cheaper reagent, extending time or temperature within the limits of stability, removing a product to pull a reversible reaction forward, and selecting a catalyst that favours your pathway over the side reactions. Losing less means fewer transfers, washing with a solvent your product is barely soluble in, cooling a recrystallisation slowly, and applying proper drying methods to constant mass before the final weighing. Diagnosing which of the two is costing you is the subject of our common lab errors guide.
07. Worked Examples
Four reactions, each taken from the given values through to a rated result. Tap a card to load its numbers into the calculator.
NaCN + (CH₃)₂CO → (CH₃)₂C(OH)CN
Very good Sodium cyanide adds to acetone to give acetone cyanohydrin, a compound also indexed as hydroxyacetonitrile. At 85.3% this sits in the very good band, and the shortfall is typical of an aqueous workup.
CaCO₃ + 2CH₃COOH → (CH₃COO)₂Ca + H₂O + CO₂
Excellent A clean acid and carbonate reaction that goes close to completion. Above 90%, so the procedure is effectively optimised.
Same reaction, poorly executed workup
Poor Recovering under a quarter of the ceiling almost always means product was lost in the filtrate rather than never formed. Check solubility before blaming the reaction.
Glucose, C₆H₁₂O₆, M = 180.16 g/mol
Good Convert first, divide second. Dividing 3.50 g by 0.0250 mol before the conversion gives a meaningless number, because the two figures must share a dimension at the moment you divide, not afterwards. Rounding the mole value to 0.0194 before dividing returns 77.6%, which is why that figure is often quoted, but carrying the full value through gives 77.71%.
The arithmetic does not change with the chemistry. A Friedel-Crafts methylation taking benzene to toluene, or a recrystallisation from methanol, is handled with exactly the same ratio, and only the molar masses differ. Five further worked problems, including limiting-reactant and mole-ratio variations, are set out in the worked examples library.
08. Distinction
Percent yield measures the output of a synthesis reaction. Percent recovery measures how much material survives a purification step such as recrystallisation or filtration.
The two share the same formula shape, a measured mass over a reference mass times 100, but the denominator means something different in each. For yield it is the stoichiometric maximum the equation allows. For recovery it is simply the mass of material you put into the purification, none of which was chemically transformed. That is the whole percent recovery formula: recovered ÷ starting × 100. Recrystallising 2.00 g of crude solid and collecting 1.60 g of pure crystals is an 80.0% recovery, and no balanced equation is involved anywhere in it.
To calculate a recovery figure, use the dedicated percent recovery calculator. A related measure, percent error, compares a measured value against an accepted one and answers a different question again.
09. Upstream
Percent yield depends entirely on a number you have to calculate first. The theoretical yield is a product of stoichiometry, and stoichiometry begins with the limiting reactant.
The limiting reactant, also called the limiting reagent, is the reagent that is fully consumed first, and it alone sets the ceiling on how much product can form. Every other reagent is an excess reactant and has no influence on the theoretical yield, which is why calculating from the wrong reagent silently inflates the denominator and depresses your percentage.
Three quantities carry the stoichiometric calculations from that reactant to a mass of product. The mole ratio comes from the coefficients of the balanced equation and tells you how many moles of product each mole of reactant can give. The molar mass converts between grams and moles at both ends of that conversion. And the balanced equation itself is the source of the ratio, so it must be correct before anything downstream can be.
The mole itself is fixed to a count, Avogadro's number, 6.022 × 10²³ units of whatever you're counting, which is what lets a mole ratio move between amounts of different substances at all; converting a measured mass into that count is covered on the grams-to-moles calculator.
Run the whole chain in one pass with the stoichiometry calculator, or identify the constraint first with the limiting reactant calculator.
The two are routinely confused. Conversion rate is the fraction of limiting reactant that was consumed. Percent yield is the fraction of the theoretical product you actually isolated. They only agree when the reaction is perfectly selective. A synthesis can reach full conversion, with no starting material left in the reaction mixture, and still return a poor yield, because the reactant went into by-products rather than the compound you wanted. Reading the two figures together is what tells you whether to work on the chemistry or on the workup.
10. Reference Chart
What counts as a normal yield varies enormously by the kind of chemistry involved. These are the ranges usually reported for each class of process.
| Reaction type | Typical yield |
|---|---|
| Organic synthesis | 60 – 85% |
| Inorganic precipitation | 85 – 98% |
| Enzymatic reactions | 40 – 90% |
| Industrial processes | 85 – 99% |
| Academic lab synthesis | 50 – 80% |
11. Questions
Divide the actual yield by the theoretical yield, then multiply by 100. The actual yield is the mass you isolated and weighed, and the theoretical yield is the maximum the balanced equation allows. For an actual yield of 4.0 g against a theoretical yield of 5.0 g, the percent yield is 4.0 ÷ 5.0 × 100 = 80%.
Percent yield in chemistry measures how efficiently a reaction converted its limiting reactant into isolated product, written as a percentage. It compares what you actually recovered against the theoretical maximum. A percent yield of 92.0% means you obtained 92 of every 100 possible grams, losing the remaining 8 to incomplete conversion, side reactions or handling.
None of its own. Percent yield is a ratio of two masses or two amounts, so the units cancel and the answer is a bare percentage. The inputs do need units, and they must match in dimension before dividing. Grams over grams, or moles over moles, both give the same percentage.
Percent loss is the fraction of the theoretical yield you did not recover, expressed as a percentage. Calculate it as (theoretical − actual) ÷ theoretical × 100, or simply 100 − percent yield whenever the actual yield sits below the theoretical. A 92.0% yield therefore carries an 8.0% loss.
Rearrange the formula so that actual yield = (percent yield ÷ 100) × theoretical yield. With a percent yield of 70% and a theoretical yield of 5 g, the actual yield is 0.70 × 5 g = 3.5 g. Switch the calculator to Find Actual Yield and it applies the rearrangement for you.
The actual yield is 10.5 g. Use actual yield = (percent yield ÷ 100) × theoretical yield, which gives (70 ÷ 100) × 15 g = 0.70 × 15 g = 10.5 g. The remaining 4.5 g is the percent loss, equal to 30% of the theoretical yield.
Yes. Percent yield is a ratio, so it is dimensionless as long as both figures share the same dimension. Two masses work and two mole amounts work equally well. What does not work is dividing a mass by an amount, so comparing 3.50 g against 0.0250 mol needs the molar mass to convert one of them first.
Stoichiometry supplies the theoretical yield that percent yield divides by. It converts the limiting reactant's mass to moles, applies the mole ratio from the balanced equation, then converts back to a product mass. Without that step you have no denominator, which is why an unbalanced equation makes every percent yield after it wrong.
The limiting reactant is the reagent that runs out first, and it matters because it alone sets the theoretical yield. Divide each reagent's moles by its coefficient, and the smallest quotient identifies it. Reagents in excess never enter the calculation, so using one by mistake inflates the theoretical yield and understates the percentage.
Yes arithmetically, but it is never a real chemical result. A figure above 100% means the mass you weighed contains something besides pure product, almost always residual solvent. Dry the sample to constant mass and re-weigh: a reading of 5.31 g against a 4.35 g ceiling fell to 4.00 g once dry, turning 122.1% into 92.0%.
Theoretical yield is calculated and actual yield is measured. The theoretical yield is the maximum mass the balanced equation permits from the limiting reactant, assuming complete conversion and no losses. The actual yield is what you isolate, dry and weigh. The first is a ceiling, the second is what the bench gave you, and it is expected to be lower.
No, not in practice. Some limiting reactant always fails to react, competing reactions consume material, and product is always left on glassware or in the filtrate. An exact 100% on a preparative scale should be read as a measurement problem, usually a sample that is not yet dry, rather than as a perfect reaction.
Above 90% is excellent and 70% is the usual working target. Between 80% and 89% is very good, 70% to 79% is good, 50% to 69% is fair, and below 50% is poor. Judge it against the reaction though, because 60% on a crowded multi-step synthesis is a stronger result than 90% on a simple precipitation.
Five causes account for most of it: incomplete conversion of the limiting reactant, competing side reactions, product lost during transfer and filtration, measurement error, and an impure reagent that lowers the real ceiling. On small scales transfer losses dominate, because a fixed few milligrams left on glassware is a larger share of a small batch.
Percent yield measures a synthesis and percent recovery measures a purification. Percent yield divides by a theoretical yield calculated from the balanced equation. Percent recovery divides by the mass you started the purification with, so no equation is involved. Recrystallising 2.00 g of crude solid and recovering 1.60 g is an 80.0% recovery.
Percent yield tells you how much product you got, and percent purity tells you how much of that mass is actually the compound. They are independent, so a 90% yield that is 80% pure gives less usable product than a clean 75% yield. Low purity is also a common reason a yield reads above 100%.
Conversion rate is the fraction of limiting reactant consumed. Percent yield is the fraction of theoretical product actually isolated. They agree only when a reaction is perfectly selective. A synthesis can reach full conversion, with no starting material left, and still return a poor yield, because the reactant went into by-products instead.
Yes, through two separate routes. Temperature changes the reaction kinetics, so a warmer mixture usually reaches completion faster. For a reversible reaction it also shifts the equilibrium position, and for an exothermic equilibrium a higher temperature lowers the ceiling. Pushing past a compound's stability limit destroys product and cuts the yield instead.
A catalyst raises the yield you actually obtain, but not the theoretical maximum. It speeds the reaction and can favour your pathway over competing side reactions, which leaves more product and fewer by-products. It does not move the equilibrium position of a reversible reaction, so it cannot push a yield past the stoichiometric limit.
Usually by driving the reaction further and losing less on the way out. Adding an excess of the cheaper reagent, extending reaction time, controlling temperature and removing a product to shift an equilibrium all raise conversion. Reducing transfers, washing with cold solvent and drying to constant mass recover material that would otherwise be discarded.
12. Related Tools
Yield tools
Stoichiometry tools
Reactant tools