What is the difference between theoretical yield and actual yield?

Theoretical yield is the maximum amount of product you could possibly get if the reaction went perfectly and every bit of the limiting reactant were converted — it's calculated purely from stoichiometry. Actual yield is what you actually recover in the lab, which is almost always less because of side reactions, incomplete reactions, and product lost during transfer and purification. Percent yield = actual / theoretical × 100.

Why is percent yield almost never 100%?

Real reactions lose product to several unavoidable causes: some reactant is consumed by competing side reactions, the reaction may not go fully to completion (it can reach equilibrium before all reactant converts), and material is inevitably lost sticking to glassware, during filtration, or in purification steps like recrystallization. A percent yield above 100% almost always means the product is impure or still wet (contains solvent or unreacted material adding extra mass), not that you beat stoichiometry.

Why do I use the limiting reactant and not the other reactant?

The limiting reactant is the one that runs out first, so it caps how much product can form — any excess of the other reactant just sits there unreacted. Theoretical yield is always calculated from the limiting reactant. If you're not sure which reactant is limiting, compute the moles of product each reactant could make on its own; whichever gives the smaller amount is the limiting one.

Can percent yield be greater than 100%?

Not really, in the sense that you can't create more product than stoichiometry allows. If your calculation gives over 100%, the actual-yield mass is too high — usually because the product still contains impurities, leftover solvent, or water it hasn't been fully dried of. Drying and purifying the product to constant mass before weighing fixes this.

Percent Yield Calculator — Theoretical & Actual, Free

Limiting reactant

Mass of limiting reactant (g)

Its molar mass (g/mol)

Mole ratio (limiting reactant : product) from the balanced equation

Reactant coefficient

Product coefficient

Product molar mass (g/mol)

Actual yield you obtained (g)

Theoretical yield

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Percent yield

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How percent yield is built, step by step

Theoretical yield is the most product the reaction could possibly make, calculated purely from the limiting reactant and the balanced equation — it assumes every bit of limiting reactant converts perfectly. You get it in three moves: convert the limiting reactant's mass to moles (divide by its molar mass), convert those to moles of product using the mole ratio from the coefficients, then convert to grams of product (multiply by the product's molar mass). Percent yield is then just how much you actually recovered as a fraction of that maximum: actual ÷ theoretical × 100.

Worked example: the Haber process (the default)

Reaction: N2 + 3 H2 → 2 NH3 (N2 is the limiting reactant) Step 1 — moles of N2: moles = mass / molar mass = 28 / 28.014 = 0.9995 mol Step 2 — moles of NH3 (mole ratio N2:NH3 = 1:2): moles NH3 = 0.9995 × (2/1) = 1.999 mol Step 3 — theoretical yield of NH3: mass = 1.999 × 17.031 = 34.04 g Step 4 — percent yield (actual = 30 g): % yield = 30 / 34.04 × 100 = 88.1%

That 88.1% is a realistic lab result — a bit of ammonia is always lost to the reverse reaction (the Haber process is a classic equilibrium that never goes fully to completion) and to handling. Change any input above and every step updates.

Practice problems

1. 5.40 g of aluminum reacts with excess chlorine: 2 Al + 3 Cl₂ → 2 AlCl₃. If you recover 24.0 g of AlCl₃ (molar mass 133.34 g/mol, Al molar mass 26.98), what is the percent yield?

Show answer

Moles Al = 5.40 / 26.98 = 0.2001 mol Mole ratio Al:AlCl3 = 2:2 = 1:1, so moles AlCl3 = 0.2001 mol Theoretical = 0.2001 × 133.34 = 26.69 g % yield = 24.0 / 26.69 × 100 = 89.9%

2. A student calculates a theoretical yield of 12.0 g but weighs 12.6 g of product and reports a 105% yield. What went wrong, and what should they do?

Show answer

A real percent yield cannot exceed 100% — you can't make more product than stoichiometry allows. A value above 100% means the weighed product is heavier than pure product should be, almost always because it still contains impurities, leftover solvent, or water it wasn't fully dried of. Fix: dry the product to constant mass (and/or purify it, e.g. by recrystallization) before weighing, then recompute. The true yield will come out at or below 100%.

FAQ

Theoretical vs. actual yield?
Theoretical is the maximum possible from stoichiometry (assumes perfect conversion of the limiting reactant); actual is what you recover in the lab, always less due to side reactions, incomplete conversion, and losses in transfer/purification.

Why is it almost never 100%?
Competing side reactions, reactions that reach equilibrium before completing, and product lost on glassware or during filtration/purification all cut into the recovered amount.

Why the limiting reactant, not the other one?
The limiting reactant runs out first and caps how much product can form; excess of the other reactant just sits unreacted. Theoretical yield is always figured from the limiting reactant.

Where does the mole ratio come from?
Straight from the balanced equation's coefficients — for N₂ + 3H₂ → 2NH₃, N₂:NH₃ is 1:2. Balance the equation first, or the ratio (and the yield) will be wrong.

Can it be over 100%?
Not genuinely. Over 100% means your actual-yield mass is inflated by impurities or leftover solvent/water — dry and purify to constant mass, then recompute.