yield percentage

Chemical reactions in the real world do not always proceed exactly as predicted on paper. Over the course of an experiment, many things will contribute to less product being formed than predicted. In addition to experimental errors, there are often losses due to incomplete reaction, unwanted side reactions, etc. The percent yield is a measure that indicates the degree of success of a reaction , and it is essential for chemical work in the laboratory.

Definition

The yield percentage is the ratio between the actual yield and the theoretical yield expressed as a percentage . The  actual yield  is the result that the chemical reaction actually produces with all the factors found in the environment. The actual percentage can be close to 100%, but never quite there.

To calculate percent yield, one must first determine the amount of product that should be formed based on stoichiometry, called the ” theoretical yield ,” that is, the maximum amount of product that could be formed from the given amounts of reactants.

If the actual and theoretical yield are equal, the percent yield is 100%. Typically, the percentage return is less than 100% because the actual return is often less than the theoretical value. Reasons for this may include incomplete or competitive reactions and loss of sample during recovery. It is also possible for the percent yield to be greater than 100%, which means that more sample was recovered from a reaction than expected. Finally, the percentage yield is always a positive value.

Theoretical and experimental performance, why do they differ?

Whenever we mix two or more reactants to carry out a chemical reaction, we can calculate by simple stoichiometry the amount of product that we should obtain from the known amounts of the reactants that we add. Since this amount of product (referred to as the yield) is calculated from the stoichiometric ratios of the chemical reaction, it is referred to as the theoretical yield.

On the other hand, the amount of product that we actually obtain when we mix the amounts of reactants and carry out the chemical reaction, is what is known as experimental yield, practical yield, or actual yield .

In the ideal case, we would obtain exactly the same quantity of product as the one calculated by stoichiometry. In this case, the percentage yield would be 100%. However, there is a wide variety of factors that make the experimental performance is never equal to the theoretical. Some of these factors are:

• Experimental measurement errors both in the amounts of mixed reagents and in the weighing or determination of the amount of product obtained.
• The presence of impurities in the reagents.
• The presence of chemical equilibria that prevent the reaction from progressing to completion because part of the products are converted back into reactants.
• The speed of reaction. If the reaction is very slow and we stop it prematurely, we will obtain less product than expected.
• Losses of reactants and products during the processes of transferring substances from one container to another.
• The occurrence of parallel chemical reactions that compromise part of the reagents, among others.

Many of these factors can be controlled to some degree, but most will always be present.

Percent Yield Formula

The equation for percentage yield is: Percentage yield = (actual yield / theoretical yield) x 100% where:

• The actual yield is the amount of product obtained from a chemical reaction.
• Theoretical yield is the amount of product that is obtained from the stoichiometric equation using the limiting reagent to determine the product.
• Units for actual and theoretical yield must be the same (moles or grams).

examples

The decomposition of magnesium carbonate forms 15 grams of magnesium oxide in one experiment, the theoretical yield is known to be 19 grams. 

1. What is the percentage yield of magnesium oxide?

MgCO ₃  → MgO + CO ₂

The calculation is simple if you know the real and theoretical returns, the next thing is to enter the values ​​in the formula:

1. Percentage yield = actual yield / theoretical yield x 100%
2. Percentage yield= 15g / 19g x 100%
3. Percentage yield= 79%

2. Percentage yield (%R) of the following chemical reaction:

2N ₂  + 5O ₂  → 2N ₂ O ₅

1. Calculate the molecular mass of the substances that are part of the chemical reaction:

N ₂ = 28 g/mol

O ₂ = 32 g/mol

N ₂ O ₅  = 108.01 g/mol

• Calculate the limiting reactant, comparing the reactants:

first relationship

2N2 _→ 5O2 _{_}

2mol x 28g/mol → 5mol x 32g/mol

40g→x

56g → 160g

40g → x

X= 114.29 g of O ₂

second relationship

2N2 _→ 5O2 _{_}

2mol x 28g/mol → 5mol x 32g/mol

x →55g

56g → 160g

x→ 55g

x= 19.25 g of N ₂

Generally we must calculate the theoretical yield based on the balanced equation. In this equation, the reactant and the product have a 1:1 molar ratio, so if you know the amount of reactant, we know that the theoretical yield is the same value in moles.

To get the amount in grams we must take the amount of grams of reactant, later convert it to moles and then use this same amount to find out how many grams of product to expect. 

References

• Brown, T. (2021). Chemistry: The Central Science (11th ed.). London, England: Pearson Education.
• Chang, R., Manzo, Á. R., Lopez, PS, & Herranz, ZR (2020). Chemistry (10th ed.). New York City, NY: MCGRAW-HILL.
• Flowers, P., Neth, EJ, Robinson, WR, Theopold, K., & Langley, R. (2019). Chemistry: Atoms First 2e . Retrieved from https://openstax.org/books/chemistry-atoms-first-2e/pages/1-introduction
• Flowers, P., Theopold, K., Langley, R., & Robinson, WR (2019b). Chemistry 2e . Retrieved from https://openstax.org/books/chemistry-2e/pages/1-1-chemistry-in-context