Stoichiometry and Limiting Reactants: Finding the Maximum Product

A reaction can stop producing product while one reactant remains. The limiting reactant determines the maximum amount of product. Comparing the starting masses alone will not identify it, because the equation requires specific mole ratios.

Stoichiometry uses a balanced chemical equation to relate amounts of reactants and products. The limiting reactant permits the smallest theoretical product amount under the stated reaction model. Theoretical yield is that calculated maximum. Percent yield compares the measured amount of desired product with the theoretical yield on the same basis.

Which ratio comes from the equation?

For

\[ \mathrm{N_2+3H_2\rightarrow2NH_3}, \]

one mole of nitrogen requires three moles of hydrogen and can form two moles of ammonia. Convert any starting masses into moles before applying those ratios.

A useful path is mass of starting substance, moles of that substance, moles of product, then mass of product. Write each unit so you can see the cancellations.

Worked example: which reactant limits?

Suppose 2.00 mol of nitrogen and 3.00 mol of hydrogen are available. Assume the stated reaction proceeds to the stoichiometric limit, with no competing reactions.

Nitrogen could support

\[ 2.00\,\mathrm{mol\ N_2}\times\frac{2\,\mathrm{mol\ NH_3}}{1\,\mathrm{mol\ N_2}}=4.00\,\mathrm{mol\ NH_3}. \]

Hydrogen could support

\[ 3.00\,\mathrm{mol\ H_2}\times\frac{2\,\mathrm{mol\ NH_3}}{3\,\mathrm{mol\ H_2}}=2.00\,\mathrm{mol\ NH_3}. \]

Hydrogen limits the product to 2.00 mol of ammonia. That consumes 1.00 mol of nitrogen, leaving 1.00 mol of nitrogen. Real ammonia production is equilibrium-limited, so this calculation is an ideal stoichiometric maximum.

How do you calculate percent yield?

If a different reaction has a theoretical yield of 12.5 g and gives 10.0 g of dry, purified product, then

\[ \text{percent yield}=\frac{10.0}{12.5}\times100\%=80.0\%. \]

A value above 100% suggests that the measured mass may include solvent or impurities, or that the calculation or measurements need checking. It does not mean extra atoms were created.

What should the final check include?

1. Confirm the equation is balanced.
2. Calculate the same product amount from every potentially limiting reactant.
3. Use the smallest supported product amount.
4. Calculate leftovers from the amount actually consumed.
5. Keep actual and theoretical yields separate.

Can you apply the idea?

1. Can the reactant with the smallest mass be assumed to limit?

   Check your answer

   No. Compare available amounts using the balanced mole ratios.

2. For \(\mathrm{2H_2+O_2\rightarrow2H_2O}\), how much water can 1.00 mol of oxygen support with excess hydrogen?

   Check your answer

   \(2.00\,\mathrm{mol}\) of water.

3. In the ammonia example, why is nitrogen in excess?

   Check your answer

   Only 1.00 of the available 2.00 mol is needed to react with all 3.00 mol of hydrogen.

4. A theoretical yield is 20.0 g and actual yield is 15.0 g. Find percent yield.

   Check your answer

   75.0%, from actual divided by theoretical times 100%.

5. Why should product be dry before a mass-based yield is interpreted?

   Check your answer

   Retained solvent adds mass that does not belong to the desired product.

6. What does theoretical yield assume about competing reactions?

   Check your answer

   The calculation attributes the limiting reactant to the stated product-forming reaction. Side reactions can reduce the actual desired yield.

Watch the idea explained

Introduction to Limiting Reactant and Excess Reactant — Tyler DeWitt.

Open this video on YouTube.

Where does this fit?

Use the chemistry learning hub to choose a lesson or practice test. Connect this topic with reaction types and redox, calorimetry and heat, balancing chemical equations, reaction rates.

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