Percent Yield vs Theoretical Yield: The Difference
You do the stoichiometry, the equation says you should get 28.0 g, and the balance says 24.5 g. Nothing has gone wrong — that's just chemistry in a real flask. These two yields are how chemists describe that gap.
The short answer: theoretical yield is the maximum amount of product a balanced equation predicts if everything reacts perfectly. Percent yield compares what you actually collected to that maximum: percent yield = (actual yield ÷ theoretical yield) × 100%.
Quick comparison at a glance
| Feature | Theoretical yield | Percent yield |
|---|---|---|
| What it is | The predicted maximum product | How efficient the reaction was |
| Where it comes from | Calculation from the balanced equation | Comparing lab result to prediction |
| Units | Grams or moles | Percent — no units |
| Needs lab work? | No, it's pure calculation | Yes, you need the actual yield |
| Based on | The limiting reactant | Actual ÷ theoretical |
| Typical value | A fixed number, e.g. 28.0 g | Usually 70–90%; never above 100% |
There's a third term hiding in that table: actual yield — simply what you weighed at the end. Theoretical is the prediction, actual is the measurement, percent is the comparison.
What is theoretical yield?
Theoretical yield is the answer to "if this reaction went perfectly, with nothing lost and nothing left over, how much product could I possibly get?"
You calculate it with ordinary stoichiometry: convert the limiting reactant to moles, apply the mole ratio, convert to grams. That last point matters — theoretical yield is always based on the reactant that runs out first, never on one that's in excess.
It's a ceiling, not a forecast. Nobody expects to hit it.
What is percent yield?
Percent yield measures how close you got:
percent yield = (actual ÷ theoretical) × 100%
A percent yield of 87.5% means you recovered 87.5% of the maximum possible. It's the number chemists actually report, because it's comparable across reactions of any size.
Why real yields fall short
- The reaction doesn't finish. Many reactions are reversible and settle at equilibrium with reactant still present.
- Side reactions. Some reactant goes off and makes something else entirely.
- Physical losses. Product clings to the filter paper, stays in solution, or is lost in transfers.
- Impure reactants. If your starting material isn't pure, you had less of it than you weighed.
Worked example
Heating 50.0 g of calcium carbonate gives 24.5 g of calcium oxide. What is the percent yield?
CaCO₃ → CaO + CO₂
Step 1 — moles of reactant. Molar mass CaCO₃ = 40.078 + 12.011 + 3(15.999) = 100.09 g/mol. 50.0 ÷ 100.09 = 0.4996 mol CaCO₃
Step 2 — mole ratio. CaO : CaCO₃ is 1 : 1, so 0.4996 mol CaO.
Step 3 — theoretical yield. Molar mass CaO = 40.078 + 15.999 = 56.08 g/mol. 0.4996 × 56.08 = 28.0 g — the theoretical yield.
Step 4 — percent yield. (24.5 ÷ 28.0) × 100 = 87.5%
More worked examples
- Ammonia. Theoretical yield 34.1 g of NH₃, actual yield 28.0 g. Percent yield = (28.0 ÷ 34.1) × 100 = 82.1%.
- Finding actual yield. A reaction has a theoretical yield of 15.0 g and runs at 76.0%. Actual = 0.760 × 15.0 = 11.4 g.
- Finding theoretical yield. You collected 9.60 g at 80.0% yield. Theoretical = 9.60 ÷ 0.800 = 12.0 g.
- A yield over 100%? If you calculate 104%, the product is almost certainly still wet or contaminated — dry it and weigh again. Genuinely exceeding the theoretical yield is impossible; atoms aren't created.
Common mistakes to avoid
- Basing theoretical yield on the wrong reactant. It must come from the limiting reactant. Using one that's in excess inflates the prediction and deflates your percent yield.
- Inverting the fraction. Actual goes on top, theoretical on the bottom. If your answer exceeds 100%, check this before blaming the experiment.
- Forgetting to balance the equation. The mole ratio comes from the coefficients, so an unbalanced equation makes the theoretical yield wrong from the start.
FAQ
What is the difference between percent yield and theoretical yield?
Theoretical yield is the maximum product predicted by the balanced equation, measured in grams or moles. Percent yield is the actual yield divided by the theoretical yield, times 100 — a measure of efficiency.
How do you calculate percent yield?
Divide the actual yield by the theoretical yield and multiply by 100. For example, 24.5 g actual out of 28.0 g theoretical gives 87.5%.
Can percent yield be more than 100%?
Not genuinely. A result above 100% means the product was impure or not fully dried, or the theoretical yield was miscalculated.
Why is the actual yield always lower than the theoretical yield?
Because of incomplete or reversible reactions, competing side reactions, impure reactants, and product lost during filtering and transfers.
The takeaway
Theoretical yield is what the equation promises; actual yield is what the balance tells you; percent yield is the honest ratio between them. Calculate the theoretical from the limiting reactant, divide the actual by it, multiply by 100 — and you've turned a disappointing-looking flask into a number that means something.
Needed first → [What Is Stoichiometry?] and [What Is Molar Mass?] (sibling posts). Read next → [What Is a Limiting Reactant?] — the reactant every theoretical yield depends on. Background → What Is the Mole? Avogadro's Number Made Simple.
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