Key Takeaways
- The empirical formula shows the simplest whole number ratio of atoms in a compound
- It's derived by converting mass or percent composition to moles, then finding the smallest ratio
- The molecular formula is always a whole number multiple of the empirical formula
- Example: Glucose (C6H12O6) has the empirical formula CH2O
What Is an Empirical Formula?
An empirical formula represents the simplest whole number ratio of atoms of each element in a chemical compound. Unlike the molecular formula, which shows the actual number of atoms, the empirical formula shows only the relative proportions.
For example, hydrogen peroxide (H2O2) has the molecular formula showing 2 hydrogen and 2 oxygen atoms, but its empirical formula is HO, showing a 1:1 ratio. Glucose (C6H12O6) has the empirical formula CH2O, representing a 1:2:1 ratio.
Common Examples
| Compound | Molecular Formula | Empirical Formula |
|---|---|---|
| Water | H2O | H2O |
| Hydrogen Peroxide | H2O2 | HO |
| Glucose | C6H12O6 | CH2O |
| Benzene | C6H6 | CH |
| Acetic Acid | C2H4O2 | CH2O |
How to Calculate Empirical Formula
Step-by-Step Process
Convert to Grams (if given percentages)
If given percent composition, assume 100g sample so percentages equal grams directly. For example, 40% carbon = 40g carbon.
Convert Grams to Moles
Divide the mass of each element by its molar mass (atomic weight from periodic table). Moles = Mass / Molar Mass
Find the Mole Ratio
Divide all mole values by the smallest number of moles. This gives the simplest ratio.
Round to Whole Numbers
If ratios are close to whole numbers (within 0.1), round directly. If not, multiply all ratios by a common factor (usually 2, 3, or 4).
Worked Example: Compound with 40% C, 6.7% H, 53.3% O
Step 1: Assume 100g sample: 40g C, 6.7g H, 53.3g O
Step 2: Convert to moles:
- C: 40g / 12.01 g/mol = 3.33 mol
- H: 6.7g / 1.008 g/mol = 6.65 mol
- O: 53.3g / 16.00 g/mol = 3.33 mol
Step 3: Divide by smallest (3.33):
- C: 3.33/3.33 = 1
- H: 6.65/3.33 = 2
- O: 3.33/3.33 = 1
Result: Empirical formula = CH2O
Empirical Formula vs. Molecular Formula
The molecular formula shows the actual number of atoms in one molecule, while the empirical formula shows only the simplest ratio. The molecular formula is always a whole number multiple of the empirical formula.
Molecular Formula = (Empirical Formula)n
To find the molecular formula from the empirical formula, you need the molar mass of the compound. Divide the molar mass by the empirical formula mass to find n, then multiply all subscripts by n.
Pro Tip: Handling Tricky Ratios
If your mole ratios don't round to nice whole numbers, look for common patterns: ratios near 0.25, 0.33, 0.5, 0.67, or 0.75 suggest multiplying by 4, 3, 2, 3, or 4 respectively. For example, a ratio of 1.5 means multiply all ratios by 2.
Applications of Empirical Formulas
- Analytical Chemistry: Determining the composition of unknown compounds
- Pharmaceutical Research: Identifying drug compounds and their structures
- Environmental Science: Analyzing pollutants and their compositions
- Food Science: Determining nutritional composition of foods
- Materials Science: Characterizing new materials and alloys
Frequently Asked Questions
The empirical formula shows the simplest whole number ratio of atoms (e.g., CH2O), while the molecular formula shows the actual number of atoms in one molecule (e.g., C6H12O6 for glucose). The molecular formula is always a whole number multiple of the empirical formula.
Check that: (1) all subscripts are whole numbers with no common factor greater than 1, (2) the percent composition calculated from your formula matches the original data, and (3) the subscripts represent the smallest possible integers.
If ratios are close to whole numbers (within 0.1), you can round. Otherwise, multiply all ratios by the same factor to get whole numbers. Common multipliers are 2 (for X.5), 3 (for X.33 or X.67), or 4 (for X.25 or X.75).
Yes! Compounds with the same empirical formula are called "isomers" if they share the same molecular formula, or they may have different molecular formulas. For example, glucose (C6H12O6) and acetic acid (C2H4O2) both have the empirical formula CH2O.
You need the molar mass of the compound. Calculate the mass of the empirical formula, then divide the molar mass by the empirical formula mass to get n. Multiply all subscripts in the empirical formula by n to get the molecular formula.