OneTaskBy Tim King · Updated
Percent Yield Calculator
Theoretical Yield From a Limiting Reactant and Its Mole Ratios
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How the Percent Yield Calculator works
This percent yield calculator compares how much product you actually made with the most you could have made. Enter the actual and theoretical yield, or let it work out the theoretical yield from the mass of the limiting reactant, the two formulas and the coefficients from the balanced equation. It shows the working and how much product was lost.
Real reactions rarely give every gram the equation promises. Some product sticks to glassware, some stays dissolved, side reactions use up starting material and a few reactions never go to completion. Percent yield measures how close a practical run came to the ideal.
How the percent yield calculator works
The formula is percent yield = actual yield ÷ theoretical yield × 100. The theoretical yield comes from the limiting reactant: moles of reactant = grams ÷ molar mass, moles of product = moles of reactant × product coefficient ÷ reactant coefficient, and grams of product = moles × the product's molar mass.
First example. A reaction should give 18 g but you collect 12.5 g. Percent yield = 12.5 ÷ 18 × 100 = 69.44%, and 5.5 g was lost or never formed.
Second example. Making aspirin, C9H8O4, from 13.8 g of salicylic acid, C7H6O3, in a 1 : 1 ratio: 13.8 ÷ 138.12 = 0.0999 mol, times 180.16 g/mol = 18.0 g theoretical. Collecting 15.2 g is an 84.44% yield.
How many grams should I expect at a 75% yield? Multiply the theoretical yield by 0.75: 18 g theoretical gives 13.5 g.
Percent yield for an 18 g theoretical yield
| Percent yield | Actual yield (g) | Lost (g) |
|---|---|---|
| 5% | 0.9 | 17.1 |
| 10% | 1.8 | 16.2 |
| 20% | 3.6 | 14.4 |
| 25% | 4.5 | 13.5 |
| 30% | 5.4 | 12.6 |
| 40% | 7.2 | 10.8 |
| 50% | 9 | 9 |
| 55% | 9.9 | 8.1 |
| 60% | 10.8 | 7.2 |
| 65% | 11.7 | 6.3 |
| 70% | 12.6 | 5.4 |
| 75% | 13.5 | 4.5 |
| 80% | 14.4 | 3.6 |
| 85% | 15.3 | 2.7 |
| 90% | 16.2 | 1.8 |
| 92% | 16.56 | 1.44 |
| 94% | 16.92 | 1.08 |
| 95% | 17.1 | 0.9 |
| 96% | 17.28 | 0.72 |
| 97% | 17.46 | 0.54 |
| 98% | 17.64 | 0.36 |
| 99% | 17.82 | 0.18 |
| 100% | 18 | 0 |
Tips and pitfalls
- Balance the equation first. The coefficients set the mole ratio; the equation balancer on this site finds them.
- Find the limiting reactant. Use the reactant that makes the least product, not the one with the smallest mass.
- Over 100% means a problem. Usually the product is still wet or contains impurities, or the theoretical yield was worked out wrongly.
- Dry before weighing. Weigh the product to a constant mass so trapped solvent is not counted.
FAQ
How do you calculate percent yield?
Divide the actual yield, the mass of product you collected, by the theoretical yield, the most the balanced equation allows, and multiply by 100. Collecting 8 g when 10 g was possible gives an 80% yield. Both masses must be in the same unit.
How do you find the theoretical yield?
Convert the mass of the limiting reactant to moles, use the mole ratio from the balanced equation to get moles of product, then multiply by the product's molar mass. Choose the theoretical yield mode and enter both formulas and the coefficients to have this done for you.
What is a good percent yield?
It depends on the reaction. Simple precipitations can exceed 95%, while a multi-step organic synthesis may be pleased with 50% per step. In school labs, 60 to 90% is typical for preparations such as aspirin, where some product is lost during recrystallisation.
Why is my percent yield over 100%?
A yield above 100% is not physically possible for pure, dry product. It almost always means the product still holds water or solvent, contains unreacted starting material or by-products, or the theoretical yield used the wrong reactant or mole ratio.
What is the difference between actual and theoretical yield?
Theoretical yield is the maximum mass of product calculated from the balanced equation if every bit of the limiting reactant reacted perfectly. Actual yield is what you weigh at the end of the experiment, which is almost always less.
See also: all Chemistry tools, Physics, Electronics.
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