Percent Yield: Videos & Practice Problems

Stoichiometry allows us to calculate the theoretical yield of a chemical reaction, while percent yield measures the actual success of the reaction. The formula for percent yield is $\frac{\mathrm{<span\; class="iROOqd">actual\; yield</span>}}{\mathrm{theoretical\; yield}}100$. Percent yields can be categorized as excellent (≥90%), very good (≥80%), good (≥70%), or poor (<40%). Actual yield is the pure product obtained, which is always less than theoretical yield due to inefficiencies in reactions, such as loss of material or external factors affecting the process.

Percent Yield in Reactions

Stoichiometry is a crucial concept in chemistry that allows us to predict the amount of product that can be generated from a given quantity of reactants. However, the true measure of success in a chemical reaction is determined by the percent yield, which reflects the efficiency of the reaction. A higher percent yield indicates a more successful synthesis of the desired product, similar to aiming for high scores on an exam.

Percent yield can be categorized into different levels of success: an excellent yield is 90% or greater, a very good yield ranges from 80% to 89%, a good yield is between 70% and 79%, and a poor yield is anything below 40%. Understanding these classifications helps in evaluating the effectiveness of a chemical process.

The formula for calculating percent yield is given by:

\[\text{Percent Yield} = \left(\frac{\text{Actual Yield}}{\text{Theoretical Yield}}\right) \times 100\]

This formula is essential to memorize, as it is often not provided on formula sheets. In this context, the actual yield refers to the amount of pure product obtained from the experiment, while the theoretical yield is the calculated amount based on stoichiometric principles.

It is important to note that the units of actual yield and theoretical yield must match when using the formula. For instance, if the actual yield is measured in grams, the theoretical yield must also be converted to grams from moles if necessary. This ensures accurate calculations.

No chemical reaction can achieve a 100% yield due to various factors such as loss of material during handling, incomplete reactions, or side reactions. Consequently, the actual yield is always less than the theoretical yield. By analyzing percent yield, actual yield, and theoretical yield, chemists can gain valuable insights into the efficiency and effectiveness of their chemical reactions.

Percent Yield Example

In a balanced chemical reaction involving benzene (C₆H₆) and oxygen, the stoichiometry reveals that 2 moles of benzene react with 15 moles of oxygen to produce 12 moles of carbon dioxide and 6 moles of water. To determine the percent yield of water produced from a 2.6 gram sample of benzene that yielded 1.25 grams of water, we first need to calculate the theoretical yield using stoichiometry.

The percent yield can be calculated using the formula:

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

Here, the actual yield of water is given as 1.25 grams. To find the theoretical yield, we start by converting the mass of benzene to moles. The molar mass of benzene is calculated as follows:

For C₆H₆: (6 × 12.011 g/mol) + (6 × 1.008 g/mol) = 78.108 g/mol

Next, we convert the grams of benzene to moles:

Moles of C₆H₆ = 2.6 g / 78.108 g/mol ≈ 0.0333 moles

Using the stoichiometric coefficients from the balanced equation, we find the moles of water produced. The ratio of moles of C₆H₆ to moles of water is 2:6, or 1:3. Therefore:

Moles of water = 0.0333 moles C₆H₆ × (6 moles H₂O / 2 moles C₆H₆) = 0.0999 moles H₂O

Now, we convert moles of water to grams. The molar mass of water (H₂O) is:

For H₂O: (2 × 1.008 g/mol) + (16.00 g/mol) = 18.016 g/mol

Thus, the theoretical yield in grams is:

Theoretical Yield = 0.0999 moles H₂O × 18.016 g/mol ≈ 1.799 grams

Finally, we can calculate the percent yield:

Percent Yield = (1.25 g / 1.799 g) × 100 ≈ 69.48%

This calculation illustrates the application of stoichiometry in determining the efficiency of a chemical reaction, highlighting the relationship between actual and theoretical yields in the context of percent yield.

For the following chemical reaction 53.1 g HBrO₄ is mixed with 25.8 g Li₂CO₃. Determine the percent yield if 7.17 g CO₂ are produced.
2 HBrO₄ (aq)  +  Li₂CO₃ (aq)  → H₂O (l)  +  CO₂ (g)  +  2 LiBrO₄ (aq)

• A. 73.1%
• B. 88.9%
• C. 90.7%

Answer: (B)

Ammonia, NH₃, can be created from the combining of H₂ and N₂ molecules. How many grams of ammonia are isolated when 10.0 g H₂ reacts with excess N₂ with a chemical reaction that has a 79.3% yield?
3 H₂  +  N₂  →  2 NH₃

• A. 34.4 g
• B. 53.0 g
• C. 44.7 g
• D. 14.5 g

Answer: (C)

The reduction of iron (III) oxide creates the following reaction: 
Fe₂O₃ (s) + 3 H₂ (g) → 2 Fe (s) + 3 H₂O (g) 
If the above reaction only went to 75% completion, how many moles of Fe₂O₃ were require to produce 0.850 moles of Fe? 

• A. 0.963 mol
• B. 2.27 mol
• C. 0.425 mol
• D. 0.567 mol
• E. 1.70 mol

Answer: (D)

Here’s what students ask on this topic:

What is the formula for percent yield in chemistry?

The formula for percent yield in chemistry is:

$\frac{\mathrm{Actual Yield}}{\mathrm{Theoretical Yield}}\times 100$

Here, the actual yield is the amount of pure product obtained from the experiment, and the theoretical yield is the maximum amount of product predicted by stoichiometric calculations. This formula helps measure the efficiency of a chemical reaction, with higher percent yields indicating greater success in obtaining the desired product.

Why is the actual yield always less than the theoretical yield?

The actual yield is always less than the theoretical yield because no chemical reaction is 100% efficient. During experiments, various factors such as incomplete reactions, side reactions, loss of product during transfer, or measurement errors can reduce the amount of product obtained. Additionally, external factors like impurities or environmental conditions may also contribute to the discrepancy. These practical limitations ensure that the actual yield is invariably lower than the theoretical yield.

How do you calculate the theoretical yield in a chemical reaction?

To calculate the theoretical yield, follow these steps:

1. Write the balanced chemical equation for the reaction.
2. Determine the limiting reactant by comparing the mole ratios of the reactants.
3. Use stoichiometry to calculate the maximum amount of product that can be formed from the limiting reactant.
4. Convert the moles of product to grams (if needed) using the molar mass of the product.

The result is the theoretical yield, which represents the maximum amount of product that can be formed under ideal conditions.

What are the categories of percent yield, and what do they indicate?

Percent yield values are categorized as follows:

• Excellent: ≥90% - Indicates very high efficiency in the reaction.
• Very Good: ≥80% - Indicates good efficiency with minimal losses.
• Good: ≥70% - Indicates moderate efficiency with some losses.
• Poor: <40% - Indicates significant inefficiencies or issues in the reaction.

These categories help evaluate the success of a chemical reaction and identify areas for improvement in experimental procedures.

What units should be used when calculating percent yield?

The units for actual yield and theoretical yield must match when calculating percent yield. For example, if the actual yield is measured in grams, the theoretical yield must also be converted to grams. Similarly, if the actual yield is in moles, the theoretical yield should be in moles. Consistent units ensure accurate calculation of the percent yield.

Can percent yield ever exceed 100%? Why or why not?

Percent yield should not exceed 100% because it represents the efficiency of a reaction, and no reaction is perfectly efficient. However, if the percent yield exceeds 100%, it usually indicates an error, such as contamination of the product, incorrect measurements, or incomplete removal of impurities. Proper experimental techniques and careful measurements are essential to avoid such errors.