Empirical and Molecular Formulas

Empirical & Molecular Formulas

Definitions

• Empirical Formula: The lowest whole number ratio of atoms in a compound.
• Molecular Formula: The actual formula of a molecular compound.
• Formulas for ionic compounds are always empirical.

Determining Empirical Formula

1. Determine the number of moles for each element.
   • If given percent, assume 100 g.
   • If given mass (g), convert grams to moles.
2. Divide each number of moles by the smallest number of moles.
3. If necessary, multiply each number by an integer to obtain all whole numbers.

Molecular Formula

• The molecular formula is a whole number multiple of the empirical formula.

Steps to find Molecular Formula

1. Determine the mass of the empirical formula.
2. Divide the molecular molar mass by the empirical formula mass to find the multiplier $n$.
3. Multiply the subscripts in the empirical formula by $n$ to obtain the molecular formula: $n(empirical formula) = molecular formula$

Examples:

• Empirical Formula: $CH_2$
  • Molecular Formulas: $C<em>2H</em>4$, $C<em>3H</em>6$, $C<em>4H</em>8$

Classifying Formulas

• Empirical: Simplest whole number ratio (e.g., $NaCl$).
• Molecular: Actual number of atoms in a molecule (e.g., $C<em>6H</em>{12}O_6$).
• Some compounds can have both empirical and molecular formulas (e.g., $H_2O$).

Example Formulas

• Acetylene ($C<em>2H</em>2$) and benzene ($C<em>6H</em>6$) have the same empirical formula: $CH$