Empirical Formula Fundamentals
An empirical formula shows the simplest whole-number ratio of atoms in a compound. It may differ from the molecular formula, which shows actual numbers. For example, glucose (C?H??O?) has empirical formula CH?O. To find empirical formula from percent composition: convert percentages to grams (assume 100 g sample), convert grams to moles using molar mass, divide all mole values by the smallest, and round to whole numbers. If needed, multiply by a common factor to get whole numbers.
Step-by-Step Process
Example: A compound is 40.0% C, 6.7% H, 53.3% O. Assume 100 g sample: 40.0 g C, 6.7 g H, 53.3 g O. Convert to moles: C = 40.0/12.01 = 3.33 mol, H = 6.7/1.008 = 6.65 mol, O = 53.3/16.00 = 3.33 mol. Divide by smallest (3.33): C = 1, H = 2, O = 1. Empirical formula is CH?O. If ratios aren't whole numbers, multiply all by a common factor (e.g., if you get 1:1.5:2, multiply by 2 to get 2:3:4).
From Empirical to Molecular Formula
The molecular formula is a whole-number multiple of the empirical formula. If empirical formula is CH?O (mass = 30 g/mol) and actual molar mass is 180 g/mol, the multiplier is 180/30 = 6, giving molecular formula C?H??O?. Combustion analysis determines empirical formulas of organic compounds by burning samples and measuring CO? and H?O produced. This technique is fundamental in organic chemistry for identifying unknown compounds.
Quick Tips
- Always verify units are consistent
- Use scientific notation for very large/small numbers
- Results are approximations — real conditions may vary
Frequently Asked Questions
Empirical formula is the simplest whole-number ratio of atoms. Molecular formula shows actual numbers of atoms and is a whole-number multiple of empirical formula. H?O? (molecular) has empirical formula HO.
Yes, when the molecular formula is already in simplest form. Water (H?O), methane (CH?), and sodium chloride (NaCl) have identical empirical and molecular formulas.
Assuming 100 g makes percentages equal to grams directly (40% becomes 40 g), simplifying calculations. Any sample size works-the mole ratio will be the same-but 100 g is most convenient.
Multiply all ratios by the smallest integer that gives whole numbers. Ratios like 1:1.5:2 become 2:3:4 when multiplied by 2. Common multipliers are 2, 3, 4, or 5.
Divide the compound's actual molar mass by empirical formula mass to get the multiplier. Multiply all subscripts in the empirical formula by this number. If multiplier is 1, molecular = empirical.