Question

How do I calculate the theoretical yield of the product in grams and mol?

Answer 1

Consider this example:

10 grams of hydrogen gas are combusted in the presence of excess oxygen gas to produce water. How much water is produced?

The reaction where hydrogen gas combines with oxygen gas to produce water is:

H ₂ (g) + O ₂ (g) ? H ₂ O(l)

Step 1

Make sure your chemical equations are balanced equations.

The equation above is not balanced. After balancing , the equation becomes:

2 H ₂ (g) + O ₂ (g) ? 2 H ₂ O(l)

Step 2

Determine the mole ratios between the reactants and the product.

This value is the bridge between the reactant and the product.

The mole ratio is the stoichiometric ratio between the amount of one compound and the amount of another compound in a reaction. For this reaction, for every two moles of hydrogen gas used, two moles of water are produced. The mole ratio between H ₂ and H ₂ O is 1 mol H ₂ /1 mol H ₂ O.

Step 3

Calculate the theoretical yield of the reaction.

There is now enough information to determine the theoretical yield . Use the strategy:

1. Use molar mass of reactant to convert grams of reactant to moles of reactant
2. Use the mole ratio between reactant and product to convert moles reactant to moles product
3. Use the molar mass of the product to convert moles product to grams of product.

In equation form:

grams product = grams reactant x (1 mol reactant/molar mass of reactant) x (mole ratio product/reactant) x (molar mass of product/1 mol product)

The theoretical yield of our reaction is calculated using:

molar mass of H ₂ gas = 2 grams
molar mass of H ₂ O = 18 grams

grams H ₂ O = grams H ₂ x (1 mol H ² /2 grams H ₂ ) x (1 mol H ₂ O/1 mol H ₂ ) x (18 grams H ₂ O/1 mol H ₂ O)

We had 10 grams of H ₂ gas, so

grams H ₂ O = 10 g H ₂ x (1 mol H ₂ /2 g H ₂ ) x (1 mol H ₂ O/1 mol H ₂ ) x (18 g H ₂ O/1 mol H ₂ O)

All the units except grams H ₂ O cancel out, leaving

grams H ₂ O = (10 x 1/2 x 1 x 18) grams H ₂ O
grams H ₂ O = 90 grams H ₂ O

Ten grams of hydrogen gas with excess oxygen will theoretically produce 90 grams of water.

This strategy can be slightly modified to calculate the amount of reactants needed to produce a set amount of product. Let's change our example slightly: How many grams of hydrogen gas and oxygen gas are needed to produce 90 grams of water?

We know the amount of hydrogen needed by the first example , but to do the calculation:

grams reactant = grams product x (1 mol product/molar mass product) x (mole ratio reactant/product) x (grams reactant/molar mass reactant)

For hydrogen gas:

grams H ₂ = 90 grams H ₂ O x (1 mol H ₂ O/18 g) x (1 mol H ₂ /1 mol H ₂ O) x (2 g H ₂ /1 mol H ₂ )

grams H ₂ = (90 x 1/18 x 1 x 2) grams H ₂ grams H ₂ = 10 grams H ₂

This agrees with the first example. To determine the amount of oxygen needed, the mole ratio of oxygen to water is needed. For every mole of oxygen gas used, 2 moles of water are produced. The mole ratio between oxygen gas and water is 1 mol O ² /2 mol H ₂ O.

The equation for grams O ₂ becomes:

grams O ₂ = 90 grams H ₂ O x (1 mol H ₂ O/18 g) x (1 mol O ₂ /2 mol H ₂ O) x (32 g O ₂ /1 mol H ₂ )

grams O ₂ = (90 x 1/18 x 1/2 x 32) grams O ₂
grams O ₂ = 80 grams O ₂

To produce 90 grams of water, 10 grams of hydrogen gas and 80 grams of oxygen gas are needed.

Theoretical yield calculations are straightforward as long as you have balanced equations to find the mole ratios needed to bridge the reactants and the product.