Percentage yield

A reaction rarely gives you all the product the equation says it could. The percentage yield compares the mass you actually collect with that maximum, the theoretical yield. Yields fall below 100% because reactions may not finish, product is lost when it is separated, and other reactions can happen.

Key facts

  • Percentage yield = mass of product made ÷ maximum (theoretical) mass × 100.
  • The theoretical yield is the maximum mass of product, worked out from the balanced equation and the mass of reactant used.
  • Yields are below 100% because the reaction may not finish, some product is lost when it is separated, and other reactions can happen.

Worked examples

Example 1

Heating some calcium carbonate could make 28 g of calcium oxide, but only 22.4 g is collected. Find the percentage yield.

  1. Percentage yield = mass made ÷ maximum mass × 100.
  2. Put the numbers in. 22.4 ÷ 28 × 100 = 80%

Answer: 80%

Example 2

A reaction should make 12.5 g of copper sulfate crystals, but only 10.0 g is collected. Find the percentage yield.

  1. Percentage yield = mass made ÷ maximum possible mass × 100. 10.0 ÷ 12.5 × 100 = 80%

Answer: 80%

Example 3

A reaction could make 27.5 g of product, but only 17.6 g is collected. What is the percentage yield?

  1. Percentage yield = (mass made ÷ maximum possible mass) × 100.
  2. 17.6 ÷ 27.5 × 100 = 64%
  3. So the yield is 64%.

Answer: 64%

Example 4

Give three reasons why the percentage yield of a reaction is less than 100%.

  1. Not all the reactants turn into product: a reversible reaction never goes to completion.
  2. Some product is lost: left on filter paper or in the beaker, or spilt when transferring.
  3. Some reactants may react in other ways (side reactions), making unwanted products.

Answer: the reaction may be reversible and not go to completion; some product is lost when it is separated or moved from one container to another; other reactions may happen and make different products

Example 5

Nitrogen reacts with hydrogen to make ammonia: N2 + 3H2 ⇌ 2NH3. 14 g of nitrogen could make 17 g of ammonia, but only 2.55 g forms. Find the percentage yield, and give one reason it is so low.

  1. Divide the actual mass by the maximum mass. 2.55 ÷ 17 × 100 = 15%
  2. The ⇌ sign shows the reaction is reversible: ammonia breaks back down into nitrogen and hydrogen, so most of the reactants never become product.

Answer: 15%; the reaction is reversible, so it does not go to completion

Example 6

C + O2 → CO2. 6 g of chem C reacts and 15.4 g of chem CO2 is collected. What is the percentage yield, to 1 decimal place? (Aᵣ: C = 12, O = 16)

  1. Moles of chem C: 6 ÷ 12 = 0.5
  2. Moles of chem CO2: 0.5 × 1 ÷ 1 = 0.5; maximum mass = 0.5 × 44 = 22 g
  3. Percentage yield = 15.4 ÷ 22 × 100 = 70.0%

Answer: 70%

Example 7

Mg + 2HCl → MgCl2 + H2. 12 g of chem Mg reacts and 36.1 g of chem MgCl2 is collected. What is the percentage yield, to 1 decimal place? (Aᵣ: Mg = 24, H = 1, Cl = 35.5)

  1. Moles of chem Mg: 12 ÷ 24 = 0.5
  2. Moles of chem MgCl2: 0.5 × 1 ÷ 1 = 0.5; maximum mass = 0.5 × 95 = 47.5 g
  3. Percentage yield = 36.1 ÷ 47.5 × 100 = 76.0%

Answer: 76%