What Percent Yield Measures

Percent yield tells you how much product you actually made in a chemical reaction compared to the maximum amount you could theoretically make. It is expressed as a percentage. If a reaction has 100 percent yield, you got every bit of product the math predicted. If it has 50 percent yield, you got half. Real reactions almost always fall short because some reactants are lost, some energy escapes as heat, or the reaction does not go to completion.

You need three pieces of information to calculate percent yield: the theoretical yield (what the balanced equation predicts), the actual yield (what you measured after the reaction), and a straightforward formula. The calculation takes less than a minute once you have those numbers.

Key Takeaways

  • Percent yield equals actual yield divided by theoretical yield, then multiplied by 100.
  • Theoretical yield comes from stoichiometry—the balanced chemical equation and the amounts of reactants you started with.
  • Actual yield is the mass or moles of product you physically measured after the reaction finished.
  • Percent yields below 100 percent are normal and expected in real laboratory work.

Finding Your Theoretical Yield

The theoretical yield is what stoichiometry predicts. Start with a balanced chemical equation. For example, if you are burning methane: CH₄ + 2O₂ → CO₂ + 2H₂O. The coefficients (1, 2, 1, 2) tell you the mole ratios.

Next, convert the mass of your limiting reactant to moles. Divide the mass in grams by the molar mass of that substance. If you started with 16 grams of methane (molar mass 16 g/mol), you have 16 ÷ 16 = 1 mole of methane. Then use the mole ratio from the balanced equation to find how many moles of product should form. In this case, 1 mole of CH₄ produces 1 mole of CO₂. Finally, convert moles of product back to grams by multiplying by its molar mass. One mole of CO₂ (molar mass 44 g/mol) equals 44 grams. That 44 grams is your theoretical yield.

The most common mistake here is using the wrong reactant. If you have two reactants, calculate the theoretical yield for both, then use the smaller number. That smaller amount comes from the limiting reactant—the one that runs out first and controls how much product you can make.

Measuring Your Actual Yield

The actual yield is the mass or moles of product you physically obtained after the reaction. You measure this by weighing the product on a balance or collecting it in a graduated cylinder, depending on whether it is a solid, liquid, or gas. Make sure your measurement includes only the pure product, not any leftover reactants or byproducts.

If the product is a solid, let it dry completely before weighing. If it is a liquid, pour it into a clean, dry container and measure its mass. If it is a gas, you may need to collect it in a balloon or graduated cylinder and convert the volume to moles using the ideal gas law or molar volume (22.4 L/mol at standard temperature and pressure). Record this number in grams or moles—use the same unit you used for the theoretical yield so the division works cleanly.

The Percent Yield Formula

Once you have both numbers, the formula is straightforward:

Percent Yield = (Actual Yield ÷ Theoretical Yield) × 100

Divide the actual yield by the theoretical yield. Then multiply the result by 100 to convert it to a percentage. If your actual yield was 35 grams and your theoretical yield was 44 grams, the calculation is (35 ÷ 44) × 100 = 79.5 percent yield.

Always make sure both numbers are in the same units before you divide. If theoretical yield is in grams, actual yield must also be in grams. If one is in moles and the other is in grams, convert one of them first.

A Complete Worked Example

Suppose you react 10 grams of sodium (Na) with excess chlorine gas (Cl₂) to make sodium chloride (NaCl). The balanced equation is 2Na + Cl₂ → 2NaCl.

First, find the theoretical yield. Sodium has a molar mass of 23 g/mol, so 10 grams ÷ 23 g/mol = 0.435 moles of Na. The mole ratio from the equation is 2 moles Na to 2 moles NaCl, which simplifies to 1:1. So 0.435 moles of Na should produce 0.435 moles of NaCl. Sodium chloride has a molar mass of 58.5 g/mol, so 0.435 moles × 58.5 g/mol = 25.4 grams. That is your theoretical yield.

After the reaction, you collect and dry the product and weigh it. Suppose you get 20 grams of NaCl. That is your actual yield. Now calculate: (20 ÷ 25.4) × 100 = 78.7 percent yield. This means you recovered about 79 percent of the product the equation predicted.

Why Percent Yield Is Always Less Than 100

In a perfect world, every atom of reactant would turn into product with no waste. In reality, several things prevent that. Some product sticks to the walls of the container or is lost when you transfer it. Some reactant does not fully convert because the reaction reaches equilibrium before completion. Heat escapes, side reactions occur, or impurities in your starting materials interfere. Evaporation, incomplete precipitation, and measurement error also reduce the actual yield.

A percent yield between 70 and 90 percent is typical for most laboratory reactions. Yields below 50 percent suggest something went wrong—check your measurements, your balanced equation, or whether you identified the limiting reactant correctly. Yields above 95 percent are rare and usually mean either the reaction was exceptionally well-controlled or the theoretical calculation included an error.

Common Mistakes to Avoid

The most frequent error is forgetting to multiply by 100 at the end. If you divide 35 by 44 and get 0.795, that is not your answer—multiply by 100 to get 79.5 percent. Another common mistake is mixing units: if you calculated theoretical yield in moles but measured actual yield in grams, the division will give you nonsense. Always convert both to the same unit first.

A third mistake is using the wrong reactant for stoichiometry. If you have two reactants and neither is labeled "excess," calculate the theoretical yield for both and use the smaller one. Using the larger number will make your percent yield appear lower than it actually is. Finally, do not forget to include all the product in your actual yield measurement. If some product remains in the reaction vessel or filter paper, weigh those too and add them to your total.

Frequently Asked Questions

Can percent yield be more than 100 percent?

In theory, no—you cannot make more product than the equation predicts. If you calculate a yield above 100 percent, something is wrong. You may have measured the actual yield incorrectly (included impurities or water), miscalculated the theoretical yield, or used the wrong molar masses. Recheck your math and your measurements.

What if I do not know the molar mass?

Look it up on the periodic table or in a chemistry reference. Add the atomic masses of all atoms in the molecule. For example, H₂O is 2(1) + 16 = 18 g/mol. Your textbook, a chemistry website, or your instructor can provide a periodic table if you do not have one.

Do I need to use grams, or can I use moles?

You can use either, as long as both your theoretical and actual yields are in the same unit. If you measure product in moles directly (for example, by gas volume), convert theoretical yield to moles too. The percent yield will be the same either way.

What does a low percent yield tell me?

It tells you that you recovered less product than the balanced equation predicted. This is normal. If your yield is unusually low (below 50 percent), check whether you identified the limiting reactant correctly, whether your product dried completely before weighing, or whether some product was lost during transfer or collection.

How do I find the limiting reactant?

Calculate the theoretical yield for each reactant separately using stoichiometry. The reactant that produces the smallest amount of product is the limiting reactant—use that theoretical yield in your percent yield calculation. The other reactant is in excess and does not limit how much product forms.