Chemistry 2e · Stoichiometry of Chemical Reactions

Reaction Yields

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In 30 seconds
Why this matters
The college version
Eli explains
Worked example
Key takeaway
Check yourself
Study tools
Sources & references

In 30 seconds

A balanced chemical equation is a recipe: it states the exact mole ratios in which reactants combine and products form. But chemistry in a beaker is not chemistry on paper. Real reactions lose material to side reactions, incomplete conversion, evaporation, and transfers between vessels, so the mass of product recovered is almost always less than the equation predicts. Reaction yields provides the vocabulary and the arithmetic for quantifying that gap.

Three quantities anchor this topic:

— the maximum mass of product the can produce, calculated entirely from stoichiometry. It answers the question: if everything went perfectly, how much product would we get?
— the mass of product actually recovered and weighed in the laboratory.
— the ratio of the two, reported as a percentage:

$percent yield = actual yield theoretical yield \times 100 %$

A percent yield near 100% means the reaction ran almost perfectly. Lower values mean material was lost or the reaction did not go to completion. Values above 100% are possible when a product is wet, impure, or contaminated — a red flag that needs investigation, not a sign of a super-efficient reaction.

Why this matters

Percent yield is not an exam-only abstraction; it is an economic and environmental number. A pharmaceutical process that runs at 60% yield wastes 40% of every batch of expensive starting material, and the discarded material still must be handled and disposed of. Industrial chemists tune catalysts, temperatures, and purification steps specifically to push yields upward, because yield improvements translate directly into cost savings and less waste. For students, yield calculations tie together everything from this chapter: mole conversions, molar mass, limiting reactants, and mole ratios. On exams, yield problems are a favorite way to test whether you can run stoichiometry forward (reactants to products) and check your answer against the balanced equation.

The college version

Core Concepts

Theoretical yield comes from the limiting reactant

Before any yield can be computed, the limiting reactant must be identified — the reactant that runs out first and therefore caps product formation. A reliable method: convert each reactant's mass to moles, divide by its coefficient in the balanced equation, and the smallest quotient identifies the limiting reactant. The theoretical yield is then calculated from that reactant alone, using the balanced equation's mole ratio. The other reactant is present in excess and cannot produce additional product once the limiting reactant is gone.

Why actual yield falls short

Actual yield is smaller than theoretical for several common reasons:

Incomplete reactions — reversible reactions may reach equilibrium before all reactants are consumed.
Side reactions — reactants may react with one another or with air and moisture to form unwanted by-products.
Losses during isolation — product sticks to flask walls and filter paper and is lost during transfers, washing, and drying.
Solubility — in precipitation-based purification, some product simply stays dissolved in the liquid phase.

None of these mean the chemist made an error; they mean real processes are inefficient.

Percent yield: putting the numbers together

Percent yield expresses efficiency on a 0–100% scale, and the calculation always follows the same two steps: (1) find the theoretical yield from the limiting reactant, and (2) divide the actual yield by it and multiply by 100. Because theoretical yield is computed from moles, molar masses must be used correctly in both directions — from reactant mass to moles, and from product moles back to mass. Checking that units cancel (grams over grams, mol over mol) is a fast way to catch mistakes.

Percent yield is not atom economy

Percent yield describes how much of a specific product you recovered. describes what fraction of the atoms in the reactants end up in the desired product, regardless of recovery. A reaction can have high atom economy (all reactant atoms incorporated into the product) yet low percent yield (poor recovery), or vice versa. Both metrics matter in green-chemistry assessments of whether a process is sustainable.

Common Confusions

Do Not Confuse	With	Difference
Theoretical yield	Actual yield	One is calculated from stoichiometry; the other is measured in the lab.
Percent yield	Percent error	Percent yield compares product obtained to product possible; percent error compares a measured value to an accepted value.
Yield above 100%	A sign of excellent chemistry	It signals wet, impure, or contaminated product — investigate before celebrating.
Limiting reactant	Reactant with the smallest mass	The limiting reactant has the smallest moles-per-coefficient, not the smallest mass.
Atom economy	Percent yield	Atom economy is about atom incorporation into product; percent yield is about recovery of that product.

Eli explains

The same idea, in plain words

Explain it like I’m 10

Imagine a cookie recipe that says the dough makes 24 cookies. The 24 cookies are the theoretical yield. When you actually bake, you drop one on the floor and burn another, so you end up with 21 — that is your actual yield. Percent yield is like a baking report card: 21 out of 24 is about 88% — pretty good baking!

Worked example

Example 1: Calculating percent yield

Potassium chlorate decomposes on heating:

$2 KClO 3 \to 2 KCl + 3 O 2$

A student heats 12.6 g of KClO₃ (molar mass 122.55 g/mol) and collects 4.20 g of O₂ gas (molar mass 32.00 g/mol). What is the percent yield?

Step 1 — moles of reactant:

$n(KClO 3) = 12.6 g 122.55 g/mol = 0.1028 mol$

Step 2 — theoretical moles and mass of O₂ from the 3:2 mole ratio:

$n(O 2) = 0.1028 mol KClO 3 \times 3 mol O 2 2 mol KClO 3 = 0.1542 mol$

$m(O 2) = 0.1542 mol \times 32.00 g/mol = 4.93 g$

Step 3 — percent yield:

$percent yield = 4.20 g 4.93 g \times 100 % = 85.2 %$

The grams cancel top and bottom, confirming the calculation is dimensionally sound.

Example 2: Identifying the limiting reactant first

Ammonia is synthesized from nitrogen and hydrogen:

$N 2 + 3 H 2 \to 2 NH 3$

A reactor is charged with 28.0 g of N₂ (molar mass 28.02 g/mol) and 10.0 g of H₂ (molar mass 2.016 g/mol).

Moles of each reactant:

$n(N 2) = 28.0 g 28.02 g/mol = 0.999 mol n(H 2) = 10.0 g 2.016 g/mol = 4.96 mol$

Per coefficient: N₂ gives 0.999 ÷ 1 = 0.999; H₂ gives 4.96 ÷ 3 = 1.65. N₂ is limiting. Theoretical yield:

$n(NH 3) = 0.999 mol N 2 \times 2 mol NH 3 1 mol N 2 = 1.998 mol$

$m(NH 3) = 1.998 mol \times 17.03 g/mol = 34.0 g$

If the plant recovers 27.2 g of NH₃, the percent yield is:

$percent yield = 27.2 g 34.0 g \times 100 % = 80.0 %$

Even though H₂ was present in excess, the N₂ supply capped ammonia production at 34.0 g.

Key takeaways

Percent yield = (actual ÷ theoretical) × 100%; good chemistry alone never justifies a yield above 100%.
The limiting reactant sets the theoretical yield — never compute yield from the reactant you have in excess.
Convert masses to moles before applying mole ratios; the molar mass must match the species in the balanced equation.
Actual yield is a measured mass from the lab; theoretical yield is a calculated mass from stoichiometry.
Yields above 100% usually mean wet, impure, or contaminated product (or a weighing error).
Percent yield measures product recovery; atom economy measures atom efficiency — different metrics, both used in green chemistry.

Check yourself

5 review questions from the chapter. Try each one, then open the answer.

Define theoretical yield and actual yield in one sentence each.
Show answer
Theoretical yield is the maximum product mass predicted by stoichiometry from the limiting reactant; actual yield is the mass of product actually recovered.
A reaction has a theoretical yield of 12.5 g and an actual yield of 9.85 g. What is the percent yield?
Show answer
(9.85 g ÷ 12.5 g) × 100% = 78.8%.
Why must you identify the limiting reactant before calculating theoretical yield?
Show answer
The limiting reactant runs out first, so it alone determines the maximum possible product; using the excess reactant would overestimate the yield.
List two reasons a real reaction might produce less than its theoretical yield.
Show answer
Incomplete reactions (equilibrium or reversibility) and losses during isolation (transfers, filtration, evaporation); side reactions are another common cause.
You obtain 104% yield of a solid product. What are two likely explanations?
Show answer
The product is probably wet (solvent not fully removed) or contaminated with an impurity that adds mass — re-dry and re-weigh before trusting the number.

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Study tools & related lessonsKey vocabulary · Related

Key vocabulary

Theoretical yield: Maximum mass of product predicted by stoichiometry from the limiting reactant.
Actual yield: Mass of product actually obtained and weighed in the lab.
Percent yield: (Actual yield ÷ theoretical yield) × 100%.
Limiting reactant: The reactant that runs out first and stops the reaction.
Excess reactant: Reactant left over after the limiting reactant is consumed.
Atom economy: Fraction of reactant atoms that end up in the desired product.

Sources & references

openstax.org — Chemistry 2e

This lesson was adapted from the open educational references above; their licenses and attributions are preserved. See Copyright & Licensing.

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