Percent Yield Formula and Examples

The most common mistake in a first chemistry lab is treating the theoretical yield as the number you should see on the balance. It is not. The theoretical yield is a calculation, never a measurement. It tells you the maximum mass of product possible if every molecule of the limiting reactant converted perfectly, nothing spilled, no side reaction happened, and no product was lost during isolation. The number you actually weigh is the actual yield, and it is always smaller. The gap between them is the percent yield.

Percent yield = (actual yield ÷ theoretical yield) × 100%. That is the only formula from OpenStax Chemistry 2e section 4.4 that matters for interpreting your experiment. If you calculate 2.5 g of theoretical product but recovered only 1.8 g, your percent yield is (1.8 / 2.5) × 100% = 72%. The reaction ran at 72% of its ideal capacity.

Percent yield is calculated by comparing actual yield to theoretical yield, can be used to find an unknown actual or theoretical yield, is nearly always below 100%, can exceed 100% for specific reasons, and has different benchmarks for what constitutes a good yield depending on the reaction type.

The percent yield formula is a single division and a multiplication:

Percent yield = (actual yield ÷ theoretical yield) × 100%

Both yields must be in the same unit, usually grams. If you have a theoretical yield of 4.20 g and recover 3.15 g, the calculation is (3.15 ÷ 4.20) × 100% = 75.0%. The theoretical yield is always calculated from the limiting reactant using the balanced chemical equation. Without the theoretical yield, percent yield cannot be computed.

This formula is universal. It applies whether you are making sodium chloride in a beaker or isolating a pharmaceutical intermediate in a pilot plant. The IUPAC Gold Book defines it the same way: yield divided by theoretical yield multiplied by 100.

Each example follows the same sequence: write the balanced equation, identify the limiting reactant, calculate the theoretical yield, measure or state the actual yield, then apply the formula.

Example 1: Sodium Chloride from Sodium Hydroxide

React 5.00 g of sodium hydroxide (NaOH, molar mass 40.00 g/mol) with excess hydrochloric acid (HCl) to produce sodium chloride (NaCl, molar mass 58.44 g/mol) and water.

Balanced equation: NaOH + HCl → NaCl + H₂O. The mole ratio is 1:1.

Moles of NaOH = 5.00 g ÷ 40.00 g/mol = 0.1250 mol. Theoretical yield of NaCl = 0.1250 mol × 58.44 g/mol = 7.305 g. If you collect 6.20 g, percent yield = (6.20 ÷ 7.305) × 100% = 84.9%.

The reaction was efficient. Losses of about 15% are typical from transfer and incomplete recovery of the solid salt.

Example 2: Decomposition of Potassium Chlorate

Heat 2.45 g of potassium chlorate (KClO₃, molar mass 122.55 g/mol) to produce potassium chloride (KCl, molar mass 74.55 g/mol) and oxygen gas.

Balanced equation: 2 KClO₃ → 2 KCl + 3 O₂. The mole ratio of KClO₃ to KCl is 2:2, or 1:1.

Moles of KClO₃ = 2.45 g ÷ 122.55 g/mol = 0.0200 mol. Theoretical yield of KCl = 0.0200 mol × 74.55 g/mol = 1.491 g. If you recover 1.12 g, percent yield = (1.12 ÷ 1.491) × 100% = 75.1%.

The loss is typical. Decomposition reactions can lose product as fine powder during heating or if the reaction is incomplete.

Example 3: Gas-Phase Combustion of Methane

Burn 4.00 g of methane (CH₄, molar mass 16.04 g/mol) in excess oxygen to produce carbon dioxide (CO₂, molar mass 44.01 g/mol) and water.

Balanced equation: CH₄ + 2 O₂ → CO₂ + 2 H₂O. The mole ratio of CH₄ to CO₂ is 1:1.

Moles of CH₄ = 4.00 g ÷ 16.04 g/mol = 0.2494 mol. Theoretical yield of CO₂ = 0.2494 mol × 44.01 g/mol = 10.98 g. If you capture 8.50 g of CO₂, percent yield = (8.50 ÷ 10.98) × 100% = 77.4%.

Gas-phase yields are often lower because of incomplete combustion or leakage of gaseous product.

The percent yield formula can be rearranged to solve for either the actual yield or the theoretical yield when the other two values are known.

To find actual yield: actual yield = (percent yield ÷ 100) × theoretical yield.
If the theoretical yield is 5.00 g and the percent yield is 85%, actual yield = (85 ÷ 100) × 5.00 g = 4.25 g.

To find theoretical yield: theoretical yield = (actual yield ÷ percent yield) × 100.
If you recovered 3.20 g of product with a percent yield of 64%, theoretical yield = (3.20 ÷ 64) × 100 = 5.00 g.

This rearrangement is essential when planning a synthesis. If you need 10 g of final product and expect a 75% yield per step, you know the theoretical yield must be at least 13.3 g, and the starting materials must be scaled accordingly.

A percent yield under 100% is normal. The three main causes are:

• Incomplete reaction. Not every molecule of limiting reactant converts to product. Equilibrium reactions, slow kinetics, or insufficient reaction time all leave starting material unreacted.
• Side reactions. The limiting reactant can form unwanted byproducts instead of the desired product. This is especially common in organic chemistry where multiple reaction pathways exist.
• Physical losses. Product adheres to glassware, is lost during filtration, evaporates, or is left behind during transfer. These losses are unavoidable even with careful technique.

These losses compound in a multistep synthesis. If each step has an 80% yield, a three-step synthesis gives an overall yield of 0.80 × 0.80 × 0.80 = 0.512, or 51.2%. That is typical for many organic preparations.

A percent yield above 100% is physically impossible if the product is pure. The theoretical yield is the absolute maximum. When your actual yield exceeds it, the problem is always measurement error or contamination. The three most common causes are:

• Wet product. Solvent or water remains trapped in the product. Drying the sample thoroughly brings the mass down.
• Impure product. Unreacted starting material, byproducts, or catalyst residue add mass. Purification by recrystallization or chromatography fixes this.
• Weighing errors. A balance that is not calibrated, a wet weighing boat, or misreading the mass all inflate the number. Recheck the measurement on a tared, dry surface.

If you consistently get over 100%, dry the product in a desiccator or oven (if it is thermally stable) and reweigh. If the number remains above 100%, run a purity assay like melting point or NMR.

What counts as a good percent yield depends on the reaction type. The ranges below come from organic chemistry journal guidelines and standard textbooks.

Yield Range	Classification	Typical Context
>90%	Excellent	Single-step, optimized, high-purity reactions. Rare in organic synthesis on first attempt.
70-90%	Good	Most well-run laboratory syntheses. Acceptable for publication.
50-70%	Fair to moderate	Common for multistep or difficult reactions. Often requires optimization.
<50%	Poor	Indicates significant problems. Check stoichiometry, conditions, and technique.

A 50% yield in a three-step synthesis (each step around 80%) is a good result. A 50% yield for a simple single-step reaction like a salt metathesis is poor. Compare your result to literature values for the same reaction under similar conditions. If your yield is far below the reported range, check your procedure and starting materials.

When your yield is unexpectedly low or high, use this table to identify the most likely cause.

Symptom	Likely Cause	What to Check
Percent yield <20%	Incomplete reaction or wrong limiting reactant	Verify stoichiometry. Run the reaction longer. Heat or add catalyst.
Percent yield 20-60%	Significant side reactions or product loss	Check for byproducts by TLC or GC. Improve isolation technique.
Percent yield 60-90%	Minor losses typical for the method	Fine-tune conditions if desired. Accept for most purposes.
Percent yield >100%	Wet or impure product	Dry thoroughly. Purify and reweigh. Calibrate balance.
Consistently low yield despite good technique	Reagent purity or stoichiometry error	Check reagent purity. Recalculate the limiting reactant using actual reagent purity.

Percent yield measures your experimental efficiency. Atom economy, defined by Trost in Science 1991, measures the theoretical waste inherent in the reaction itself: the percentage of reactant atoms that end up in the desired product. A reaction can have 100% atom economy (all atoms incorporated into product) but a low percent yield (if the reaction does not go to completion or product is lost). Conversely, a reaction with poor atom economy can still give a high percent yield if you recover most of the product. Both numbers matter for green chemistry, but they answer different questions. Percent yield tells you how well you ran the reaction; atom economy tells you how well the reaction was designed.