Percent Yield Calculator

Calculate the percent yield of a reaction. Understand the percent yield of a reaction.

Percent yield
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Chemistry is as much about precision as it is about discovery. When conducting chemical reactions, one of the most important calculations you'll perform is determining the percent yield. This metric helps chemists understand how efficient their reactions are and identify potential areas for improvement.

What is percent yield?

Percent yield compares the amount of product actually obtained from a chemical reaction (actual yield) to the maximum amount that could theoretically be produced (theoretical yield). It's expressed as a percentage:

Percent Yield=Actual YieldTheoretical Yield×100%\text{Percent Yield} = \frac{\text{Actual Yield}}{\text{Theoretical Yield}} \times 100\%

For example, if a reaction should theoretically produce 50 grams of product, but you only obtain 43 grams, the percent yield would be:

Percent Yield=43 g50 g×100%=86%\text{Percent Yield} = \frac{43 \text{ g}}{50 \text{ g}} \times 100\% = 86\%

Why is percent yield important?

Understanding percent yield helps chemists:

  1. Evaluate reaction efficiency
  2. Identify possible sources of product loss
  3. Optimize reaction conditions
  4. Determine economic feasibility of industrial processes
  5. Compare different synthetic routes

Calculating theoretical yield

To determine percent yield, you must first calculate the theoretical yield using stoichiometry:

  1. Write a balanced chemical equation
  2. Identify the limiting reagent (the reactant that will be completely consumed)
  3. Calculate the moles of limiting reagent
  4. Use the mole ratio from the balanced equation to calculate moles of product
  5. Convert moles of product to mass (or volume for gases)
Theoretical Yield (g)=Moles of Product×Molar Mass (g/mol)\text{Theoretical Yield (g)} = \text{Moles of Product} \times \text{Molar Mass (g/mol)}

Common reasons for less than 100% yield

It's rare to achieve a 100% yield in real-world reactions due to various factors:

  • Side reactions producing unwanted byproducts
  • Incomplete reactions (not reaching equilibrium)
  • Product losses during purification steps
  • Measurement errors
  • Product remaining in solution or adsorbed on surfaces

Example calculation

Let's walk through a complete example:

In the reaction between sodium hydroxide and hydrochloric acid:

NaOH (aq)+HCl (aq)→NaCl (aq)+H2O (l)\text{NaOH (aq)} + \text{HCl (aq)} \rightarrow \text{NaCl (aq)} + \text{H}_2\text{O (l)}

If we react 20.0 g of NaOH with excess HCl and collect 29.5 g of NaCl after evaporation, what is the percent yield?

Step 1: Calculate moles of NaOH

Moles of NaOH=20.0 g40.0 g/mol=0.500 mol\text{Moles of NaOH} = \frac{20.0 \text{ g}}{40.0 \text{ g/mol}} = 0.500 \text{ mol}

Step 2: Calculate theoretical yield of NaCl

Moles of NaCl=0.500 mol NaOH×1 mol NaCl1 mol NaOH=0.500 mol\text{Moles of NaCl} = 0.500 \text{ mol NaOH} \times \frac{1 \text{ mol NaCl}}{1 \text{ mol NaOH}} = 0.500 \text{ mol} Theoretical yield=0.500 mol×58.5 g/mol=29.25 g\text{Theoretical yield} = 0.500 \text{ mol} \times 58.5 \text{ g/mol} = 29.25 \text{ g}

Step 3: Calculate percent yield

Percent yield=29.5 g29.25 g×100%=100.9%\text{Percent yield} = \frac{29.5 \text{ g}}{29.25 \text{ g}} \times 100\% = 100.9\%

Note: A yield slightly above 100% suggests experimental error, possibly due to impurities or measurement inaccuracies.

Interpreting your results

  • 90-100%: Excellent yield, near-perfect reaction conditions
  • 70-90%: Good yield, typical for many laboratory syntheses
  • 50-70%: Moderate yield, may benefit from optimization
  • Below 50%: Low yield, consider reviewing reaction conditions or purification methods
  • Above 100%: Indicates measurement error or impurities in the product

Remember that percent yield is just one metric for evaluating reaction success. Consider other factors like product purity, reaction time, and reagent costs when optimizing chemical processes.