# Percentage Yield

### Unit 7; Learning Target 3

## What is Percent Yield?

**Background:**Up until now, all the reactions we have discussed have taken place under ideal conditions. However, ideal conditions don't exist in real life. Reactants may be impure, reactions may not go to completion, or given reactions may have to compete with several smaller side reactions. In fact, in the laboratory, if you get 60% of the expected amount of product, that is considered very good.

Terms to Know:Terms to Know:

The calculated or expected amount of product is called the

*. The theoretical yield is the one that is calculated based on*

**theoretical yield***mass-mass calculations*, as we have been doing so far.

Typically, what is produced in an experiment is less then predicted by the theoretical yield. The amount of product actually produced is called the

*. The actual yield is what is formed in the*

**actual yield***experiment*--you cannot calculate this; you must conduct an experiment for this amount.

When you divide actual yield by theoretical yield you get a decimal percentage known as the

**of a reaction. Formula for percent yield is:**

*percent yield*## Steps to Solve Percent Yield Problems

- Write a balanced chemical reaction.
- Look at the problem; there should be an amount of product formed given. This will be your actual yield.
- Look at the problem again; there should be an amount of a reactant given. This will be your "given" for your mass-mass calculation.
- Set up a mass-mass calculation, using this reactant given, and solving for the product mentioned by the actual yield. You must solve for the same product as is formed in the actual yield. The answer to the mass-mass calculation will be your theoretical yield.
- Once you have solved for the theoretical yield, solve for percent yield using the given formula.

## Percent Yield Example 1

Calcium carbonate decomposes. What is the percent yield if 60.0 grams of CaCO3 is heated to form 15.0 grams of CaO?

2.

15.0 grams: this is actual yield because it has to be an amount of PRODUCT formed.

3.

60.0 g: this is the "given" because it must be an amount of REACTANT used.

4.

**Write a balanced equation:**

2.

**Find the actual yield:**15.0 grams: this is actual yield because it has to be an amount of PRODUCT formed.

3.

**Find "given" for mass-mass calculation of theoretical yield:**60.0 g: this is the "given" because it must be an amount of REACTANT used.

4.

**Set up a mass-mass calculation:**

**Solve for percent yield using formula.**

Actual Yield = 15.0 grams

Theoretical Yield = 33.62 grams

Percent Yield = 15 g / 33.62 g (x 100)

= 0.4462 (x 100)

**= 44.62 % = Percent Yield**

(this is your final answer).

(this is your final answer).

## Practice Problems

Solve the following practice problems on your paper.

What is the percentage yield in the following reaction? 5.50 grams of hydrogen gas reacts with nitrogen gas to form 20.40 grams of ammonia (NH3)?

In the reaction of sodium hydroxide and carbon dioxide to produce sodium carbonate and water,0.40 g of sodium hydroxide produces 0.51 g of sodium carbonate. What is the percentage yield for this reaction?

*Practice Problem 1)*What is the percentage yield in the following reaction? 5.50 grams of hydrogen gas reacts with nitrogen gas to form 20.40 grams of ammonia (NH3)?

*Practice Problem 2)*In the reaction of sodium hydroxide and carbon dioxide to produce sodium carbonate and water,0.40 g of sodium hydroxide produces 0.51 g of sodium carbonate. What is the percentage yield for this reaction?

Click here when finished with both practice problems

You must click here when you are finished and answer the questions on the form. This counts as a unit 7 assignment.