Limiting Reactant and Yield

Limiting Reactant

In experimental chemistry, one or more reactants are often used in excess so that the other, more expensive or important, reactant is completely consumed. The reactant that is completely used up first and thus controls the amount of product formed is called the Limiting Reactant. Once it is consumed, the reaction stops and no additional product can form. The reactant that remains unreacted is called the Excess Reactant.
Purpose of excess: An excess of one reactant ensures the other reactant is fully consumed, avoiding wastage of expensive materials.
Reaction speed: Sometimes an excess reactant (such as oxygen) is used deliberately to make the reaction proceed faster.
Product limitation: The quantity of every product in the reaction is determined by the limiting reactant — no amount of excess reactant can produce more product once the limiting reactant is gone.
To identify the Limiting Reactant in a reaction where the quantities of two or more reactants are given, follow three systematic steps using the Stoichiometry of the balanced equation.
Step 1 — Convert to moles: Calculate the number of moles of each reactant from the given mass using where is the given mass and is the molar mass.
Step 2 — Predict product moles: Using the mole ratio from the balanced chemical equation, calculate how many moles of product each reactant can produce.
Step 3 — Identify limiting reactant: The reactant that produces the smaller (least) number of moles of product is the limiting reactant.
Once the Limiting Reactant is identified, you can calculate how much of the Excess Reactant remains unreacted by using the stoichiometric ratio from the balanced equation.
Find moles needed: From the balanced equation, determine how many moles of the excess reactant are required to fully react with all of the limiting reactant.
Subtract from initial: The difference between the initial moles of the excess reactant and the moles that actually reacted gives the moles left over.
Convert back to mass: Multiply the leftover moles by the molar mass of the excess reactant to get the mass remaining unreacted.

Yield — Theoretical, Actual and Percentage

The amount of product calculated from a balanced chemical equation, assuming perfect conditions, is called the Theoretical Yield. It represents the maximum possible amount of product. The amount of product actually obtained from a real experiment is called the Actual Yield. The ratio of actual yield to theoretical yield, expressed as a percentage, is the Percentage Yield, which measures the efficiency of the reaction.
Percentage yield expresses how efficient a chemical reaction is by comparing the product actually obtained to the maximum possible product from stoichiometry.
=The mass of product obtained from the experiment(g or kg)
=The maximum mass of product predicted by stoichiometric calculation(g or kg)
=Efficiency of the reaction as a percentage(%)
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Percentage yield = 100% (ideal, rarely achieved)
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Percentage yield < 100% (the typical case)
Theoretical yield: Always calculated from the limiting reactant using the balanced equation — it assumes no losses.
Actual yield: Always less than theoretical yield in practice due to experimental imperfections.
Percentage yield: A value between 0% and 100% that tells you how efficient the reaction was. Higher percentage means less wastage.

Reasons Actual Yield is Less Than Theoretical Yield

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Incomplete reactions — some reactant may not fully convert to product
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Side reactions — reactants may form undesired by-products
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Mechanical losses during workup — filtration, washing, transfer, and crystallization can lose material
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Impure reactants — starting materials that are not 100% pure give less product
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Practical handling errors — inexperienced technique by the operator