1.2.2 Empirical & Molecular Formula

Understanding the empirical and molecular formulas of compounds is essential for determining the composition and structure of molecules.

Empirical Formula

The empirical formula of a compound represents the simplest whole-number ratio of atoms of each element present. It does not necessarily show the actual number of atoms in a molecule but provides the simplest ratio.

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Example: For glucose $C_6H_{12}O_6$ , the empirical formula is $CH_2O$ , which shows the simplest 1:2:1 ratio of carbon, hydrogen, and oxygen atoms.

Molecular Formula

The molecular formula provides the actual number of atoms of each element in a molecule. It is often a whole-number multiple of the empirical formula.

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Example: For glucose $C_6H_{12}O_6$ , the molecular formula shows the actual number of atoms in the molecule. In this case, the molecular formula is 6 times the empirical formula $CH_2O$ .

Relationship Between Empirical and Molecular Formula

The molecular formula is always a whole-number multiple of the empirical formula. This relationship can be expressed as:

\text{Molecular Formula} = (\text{Empirical Formula}) \times n

Where $n$ is the ratio of the molar mass of the compound to the molar mass of the empirical formula.

n = \frac{M_r (\text{Molecular})}{M_r (\text{Empirical})}

Calculating Empirical Formula

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Example: How to calculate the empirical formula from given data

Step 1: Convert masses to moles:

Use the formula:

\text{Moles} = \frac{\text{Mass}}{A_r}

Where $A_r$ is the relative atomic mass of each element.

Step 2: Divide each mole value by the smallest number of moles:

This gives the simplest ratio.

Step 3: Adjust the ratio to whole numbers:

If the resulting ratio isn't a whole number, multiply all values by a suitable factor to obtain whole numbers (e.g., multiply by 2 if one value is 1.5).

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Example: A compound contains 2.4 g of carbon and 0.4 g of hydrogen.

Calculate the empirical formula.

Step 1: Convert masses to moles:

\text{Moles of C} = \frac{2.4}{12} = 0.2

\quad \text{Moles of H} = \frac{0.4}{1} = 0.4

Step 2: Divide each mole value by the smallest number of moles:

\frac{0.2}{0.2} = 1

\quad \frac{0.4}{0.2} = 2

Step 3: Adjust the ratio to whole numbers:

The ratio is 1:2, so the empirical formula is CH₂.

Calculating Molecular Formula

To calculate the molecular formula, you need:

The empirical formula.
The molar mass ( $M_r$ ) of the compound.

Steps:

Find the molar mass of the empirical formula.
Divide the molar mass of the molecular formula by the empirical formula mass to find $n$ .
Multiply the empirical formula by $n$ to obtain the molecular formula.

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Example: The empirical formula of a compound is $CH_2O$ and its molar mass is 180 g/mol

Determine the molecular formula.

Step 1: Find the molar mass of the empirical formula.

Molar mass of

CH_2O = 12 + 2 + 16 = 30 g/mol.

Step 2: Divide the molar mass of the molecular formula by the empirical formula mass to find $n$ .

Find $n$ :

n = \frac{180}{30} = 6

Step 3: Multiply the empirical formula by $n$ to obtain the molecular formula.

Multiply the empirical formula by 6:

\text{Molecular formula} = C_6H_{12}O_6

Empirical & Molecular Formula Simplified Revision Notes for A-Level AQA Chemistry

1.2.2 Empirical & Molecular Formula

Empirical Formula

Molecular Formula

Relationship Between Empirical and Molecular Formula

Calculating Empirical Formula

Example: How to calculate the empirical formula from given data

Calculating Molecular Formula

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