Percent Yield Calculator
Calculate percent yield, theoretical yield, and actual yield instantly with full step-by-step working. Solve stoichiometry problems, predict experimental yield, and analyze multi-step reactions — all in one free percent yield calculator.
Using This Number to Predict the Experimental Yield
Chemists calculate the theoretical yield first using stoichiometry, then apply a realistic percent yield — based on experience with the reaction type — to predict how much product they will actually collect. For example, if a reaction typically achieves 75% yield and the theoretical yield is 45 g, the predicted experimental yield is 0.75 × 45 = 33.75 g. This prediction helps labs prepare the correct amount of starting materials and set realistic expectations before running the experiment.
Using this number to predict the experimental yield is especially important in industrial chemistry, where even a 5% improvement in yield can translate to significant cost savings at scale. The slider above lets you explore how the predicted actual yield changes as efficiency assumptions change, giving an intuitive visual feel for the relationship between theoretical and experimental quantities.
Multi-Step Formula: Overall Yield = (Step 1 % × Step 2 % × Step 3 % × …) ÷ 100(n−1)
What is Percent Yield?
Percent yield is a measure of how efficient a chemical reaction is in practice. It compares how much product you actually collected in the lab (the actual or experimental yield) to the maximum amount you theoretically could have produced based on stoichiometric calculations (the theoretical yield). Expressed as a percentage, it tells chemists how close their reaction ran to perfect efficiency.
Percent Yield Formula:
Percent Yield = (Actual Yield ÷ Theoretical Yield) × 100
In reality, reactions almost never achieve 100% yield. Losses occur due to:
- Side reactions — competing reactions consume the reactants to form unwanted products
- Incomplete reactions — not all limiting reagent is consumed, especially in equilibrium reactions
- Product loss during transfer — some product sticks to glassware or is lost during filtration, recrystallization, or distillation
- Measurement imprecision — weighing errors introduce small discrepancies
- Volatile products — some product evaporates before it can be collected and weighed
Theoretical Yield vs Actual Yield vs Experimental Yield
These three terms are closely related but refer to different quantities. Understanding the distinction is essential for answering any percent yield problem correctly.
| Term | Definition | How to Calculate / Measure | Example |
|---|---|---|---|
| Theoretical Yield | The maximum possible amount of product if the reaction were 100% efficient | Calculated from stoichiometry using the limiting reagent | 36 g H₂O from 4 g H₂ |
| Actual Yield | The amount of product actually collected after the experiment | Measured by weighing the purified product in the lab | 30 g H₂O collected |
| Experimental Yield | The same as actual yield — the measured quantity from the experiment | Measured directly; sometimes corrected for purity | 30 g H₂O (same as actual) |
Note: Experimental yield and actual yield mean exactly the same thing in most chemistry courses. Both refer to the amount of product you physically collect from the reaction. Some textbooks use them interchangeably.
How to Calculate Percent Yield — Step by Step
The following worked example shows the complete procedure from a balanced equation through to final percent yield.
Reaction: 2H₂ + O₂ → 2H₂O
Given: 4 g of H₂ reacts with excess O₂. Actual yield of water collected = 30 g.
Step 1 — Find Moles of Limiting Reagent (H₂)
Molar mass of H₂ = 2 g/mol
Step 2 — Use Stoichiometry to Find Theoretical Yield of H₂O
From the balanced equation: 2 mol H₂ produces 2 mol H₂O (1:1 ratio). Molar mass of H₂O = 18 g/mol.
Theoretical Yield = 2 mol × 18 g/mol = 36 g H₂O
Step 3 — Calculate Percent Yield
% Yield = (30 ÷ 36) × 100 = 83.3%
This is a good yield — typical for many organic chemistry lab reactions.
Worked Examples
1 — How Do You Figure Out Percent Yield?
Divide the actual yield by the theoretical yield, then multiply by 100. The formula is: % Yield = (Actual Yield ÷ Theoretical Yield) × 100. For example, if you collected 25 g of product and the theoretical yield was 32 g: % Yield = (25 ÷ 32) × 100 = 78.1%. A percent yield above 70% is generally considered a successful reaction in undergraduate lab settings.
2 — How to Find Percent Yield When Given Grams
When both actual yield and theoretical yield are given in grams, apply the formula directly — no unit conversion is needed. Simply divide actual by theoretical and multiply by 100. If actual = 18.5 g and theoretical = 22.0 g: % Yield = (18.5 ÷ 22.0) × 100 = 84.1%. Always make sure both yields are in the same unit before dividing.
3 — How to Compute Percent Yield from a Reaction
Start with the balanced chemical equation to find the stoichiometric ratio. Calculate the theoretical yield from the moles of limiting reagent. Weigh the actual product collected. Then compute: % Yield = (Actual ÷ Theoretical) × 100. For the reaction A + B → C, if 3 mol A gives a theoretical yield of 60 g C, but you only collected 48 g: % Yield = (48 ÷ 60) × 100 = 80%.
4 — How to Find Theoretical Yield from Moles
Once you have the moles of product from stoichiometry, multiply by the molar mass of the product. For example, if you calculate 2.5 mol of NaCl will form, and molar mass of NaCl = 58.44 g/mol: Theoretical yield = 2.5 × 58.44 = 146.1 g. This is the maximum mass of NaCl that could be produced under perfect conditions.
5 — How to Find Experimental Yield in Chemistry
The experimental yield is simply what you collect and weigh at the end of the reaction after purification. It is measured directly using a balance — after filtering, drying, or any other purification step. If your reaction produces a white crystalline solid and after drying you have 12.3 g on the filter paper, then your experimental (actual) yield = 12.3 g. This is what you use in the percent yield formula.
6 — What is a Good Percent Yield in Chemistry?
In industrial chemistry, yields above 90% are considered excellent and are economically important. In organic synthesis labs (undergraduate level), yields of 70–85% are typically considered good, and anything above 85% is excellent. Some complex multi-step syntheses may have overall yields as low as 5–20% — not because the chemist made errors, but because each individual step compounds the losses. Yields greater than 100% indicate an error such as incomplete drying of the product (water adds mass).
7 — Why is Percent Yield Less Than 100%?
Percent yield is less than 100% because no real reaction is perfectly efficient. Causes include: incomplete reactions (especially reversible equilibrium reactions), side reactions consuming reactants, product lost during purification steps like filtration and recrystallization, product sticking to glassware, volatile products evaporating, and measurement errors. The theoretical yield represents an ideal maximum that is physically impossible to achieve in practice.
8 — How to Calculate Overall Yield for a Multi-Step Reaction
Multiply each individual step's percent yield together, converting each to a decimal first, to get the overall fraction. Then multiply by 100 for the overall percent yield. For three steps at 85%, 72%, and 90%: Overall = (0.85 × 0.72 × 0.90) × 100 = 0.5508 × 100 = 55.1%. You can also write this as: (85 × 72 × 90) ÷ 100² = 55.08%.
9 — How to Predict Experimental Yield from Theoretical Yield
Multiply the theoretical yield by the expected percent yield (as a decimal). If the theoretical yield is 50 g and you expect 75% efficiency based on past experience: Predicted actual yield = 50 × 0.75 = 37.5 g. This prediction is used before running the experiment to ensure enough starting material is available and to set procurement expectations. Use the Predictor tool above to explore this relationship interactively.
10 — How to Find Percentage Yield When Actual and Theoretical Yield Are Given
This is the direct application of the percent yield formula. Simply divide the actual yield by the theoretical yield and multiply by 100. Always double-check that both values are in the same unit (both in grams, or both in moles). If actual yield = 14.8 g and theoretical yield = 17.5 g: % Yield = (14.8 ÷ 17.5) × 100 = 84.6%. Use Mode A in the calculator above to compute this instantly.
Percent Yield in Real Chemistry — Reference Table
Percent yield varies enormously between reaction types. Industrial processes are optimized for efficiency; lab syntheses depend on the reaction mechanism, purification method, and skill of the chemist.
| Reaction | Typical % Yield | Reason for Yield Loss |
|---|---|---|
| Haber process (N₂ + H₂ → NH₃) | ~15% per pass | Equilibrium-limited; gases recycled continuously |
| Esterification (acid + alcohol → ester) | ~60–70% | Reversible reaction; water must be removed |
| Aspirin synthesis (lab) | ~70–80% | Product loss during recrystallization and filtration |
| NaCl precipitation in solution | ~95% | Small amount remains in solution; fast, clean reaction |
| Combustion of methane (CH₄ + O₂) | ~99% | Near-complete reaction; only trace incomplete combustion |
| Fermentation (glucose → ethanol) | ~80–90% | Enzyme efficiency; by-products (CO₂ loss accounted for) |
| Grignard reaction (organic synthesis) | ~50–75% | Moisture sensitivity, side reactions, multiple steps |
Industrial processes like the Haber process achieve low per-pass yields but are economically viable because unreacted gases are recycled. Lab reactions like aspirin synthesis lose product during purification. Precipitation reactions are among the most efficient because the product immediately leaves solution as a solid.
Frequently Asked Questions
Related Calculators
| Reaction | % Yield |
|---|---|
| Haber (NH₃) | ~15% |
| Esterification | 60–70% |
| Aspirin (lab) | 70–80% |
| NaCl precipit. | ~95% |
| Combustion CH₄ | ~99% |
| Fermentation | 80–90% |