Empirical & Molecular Formula | Cambridge O Level Chemistry Revision Notes 2023

Calculating Empirical & Molecular Formulae

Calculating Empirical Formula

The empirical formula is the simplest whole number ratio of the atoms of each element present in one molecule or formula unit of the compound
- E.g. the empirical formula of ethanoic acid is CH₂O
Organic molecules often have different empirical and molecular formulae
The formula of an ionic compound is always an empirical formula

Worked example

A compound that contains 10 g of hydrogen and 80 g of oxygen.

What is its empirical formula of this compound?

Answer

	hydrogen	oxygen
Write the mass of each element	10 g	80 g
Divide each mass by the relative atomic mass to find the number of moles	10/1 = 10	80/16 = 5
Find the molar ratio by dividing by the smallest number	10/5 = 2	5/5 = 1

Empirical formula = H₂O

Worked example

Substance X was analysed and found to contain 31.58% carbon, 5.26% hydrogen and 63.16% oxygen by mass.

What is the empirical formula of substance X?

Relative atomic masses, Ar: C = 12; H = 1; O = 16

Answer

	carbon	hydrogen	oxygen
Convert % to g by assuming 100 g of substance is present	31.58 g	5.26 g	63.16 g
Divide each mass by the relative atomic mass to find the number of moles in 100 g	31.58/12 = 2.63	5.26/1 = 5.26	63.16/16 = 3.95
Find the molar ratio by dividing by the smallest number	2.63/2.63 = 1	5.26/2.63 = 2	3.95/2.63 = 1.5
Multiply all by 2 to obtain a whole number ratio	2	4	3

Empirical formula = C₂H₄O₃

Examiner Tip

The molar ratio must be a whole number. If you don't get a whole number when calculating the ratio of atoms in an empirical formula, such as 1.5, multiply that and the other ratios to achieve whole numbers.

Calculating Molecular Formula

Molecular formula gives the actual numbers of atoms of each element present in the formula of the compound
To calculate the molecular formula:
- Step 1: Find the relative formula mass of the empirical formula
- Step 2: Use the following equation:

$\frac{relative formula mass of molecular formula}{relative formula mass of empirical formula}$

- Step 3: Multiply the number of each element present in the empirical formula by the number from step 2 to find the molecular formula

Table showing the Relationship between Empirical and Molecular Formula

Relationship between Empirical _ Molecular Formula table, IGCSE & GCSE Chemistry revision notes

Worked example

The empirical formula of X is C₄H₁₀S₁

The relative formula mass (M_r) of X is 180.

What is the molecular formula of X?

(Relative atomic mass, A_r: Carbon : 12 Hydrogen : 1 Sulfur : 32 )

Answer

Step 1 - Calculate the relative formula mass of the empirical formula

M_r= (12 x 4) + (1 x 10) + (32 x 1) = 90

Step 2 - Divide relative formula mass of X by relative formula mass of empirical formula

180 / 90 = 2

Step 3 - Multiply each number of elements by 2

(C_{4 x 2}) + (H_{10 x 2}) + (S_{1 x 2})

Molecular Formula of X = C₈H₂₀S₂

Deducing formulae of hydrated salts

The formula of hydrated salts can be determined experimentally by weighing a sample of the hydrated salt, heating it until the water of crystallisation has been driven off, then reweighing the now anhydrous salt
From the results, you can determine the mass of anhydrous salt and the mass of the water of crystallisation
Applying a similar approach to deducing empirical formulae, the formula of the hydrated salt can be calculated

Worked example

11.25 g of hydrated copper sulfate, CuSO₄.xH₂O, is heated until it loses all of its water of crystallisation. It is reweighed and its mass is 7.19 g. What is the formula of the hydrated copper(II) sulfate?

Answer

	CuSO₄	H₂O
Deduce the mass of water of crystallisation and anhydrous salt	= mass of salt after heating = 7.19 g	mass of hydrated salt - mass of anhydrous salt: 11.25 - 7.19 = 4.06 g
Divide each mass by the relative formula mass to find the number of moles	$\frac{7.19}{160} = 0.045$	$\frac{4.06}{18} = 0.226$
Find the molar ratio by dividing by the smallest number	$\frac{0.045}{0.045} = 1$	$\frac{0.226}{0.045} = 5$

Therefore the formula of hydrated copper(II) sulfate is CuSO₄.5H₂O

Examiner Tip

The specification is not clear about whether deducing the formula of hydrated salts is required, however, it is an application of deducing empirical formulae so it is worth knowing how to do this.

Empirical & Molecular Formula (Cambridge O Level Chemistry)

Revision Note

Calculating Empirical & Molecular Formulae

Calculating Empirical Formula

Worked example

Worked example

Examiner Tip

Calculating Molecular Formula

Table showing the Relationship between Empirical and Molecular Formula

Worked example

Deducing formulae of hydrated salts

Worked example

Examiner Tip

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Author: Caroline