3.2 Determining Empirical and Molecular Formulas

Chemical formulas are like a secret code for molecules. They tell us what elements make up a compound and how many atoms of each are present. Empirical formulas show the simplest ratio, while molecular formulas reveal the actual number of atoms.

Understanding these formulas is crucial for chemists. They help us calculate quantities in reactions, predict properties, and identify unknown substances. Mastering how to derive and interpret these formulas opens up a world of chemical knowledge and problem-solving skills.

Empirical and Molecular Formulas

Percent composition of compounds

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• Represents the mass percentage of each element in a compound calculated by dividing the mass of each element by the total mass of the compound and multiplying by 100
• Formula for percent composition: \text{[Percent composition](https://www.fiveableKeyTerm:Percent_Composition) of element} = \frac{\text{Mass of element}}{\text{Total mass of compound}} \times 100\%
• Steps to calculate percent composition:
  1. Determine the mass of each element in the compound using atomic masses (e.g., C = 12.01 g/mol, H = 1.01 g/mol, O = 16.00 g/mol)
  2. Calculate the total mass of the compound by adding the masses of all elements
  3. Divide the mass of each element by the total mass and multiply by 100 to obtain percent composition (e.g., glucose (C₆H₁₂O₆): C = 40.00%, H = 6.72%, O = 53.28%)

Empirical formulas from elemental data

• Represents the simplest whole-number ratio of atoms in a compound derived from elemental composition data
• Steps to derive empirical formula:
  1. Convert the percent composition or mass of each element to moles using the molar mass with the formula: \text{Moles of element} = \frac{\text{Mass of element}}{\text{[Molar mass](https://www.fiveableKeyTerm:Molar_Mass) of element}}
  2. Divide each mole value by the smallest mole value to obtain the simplest whole-number ratio (e.g., C:H:O = 1:2:1 for glucose)
  3. If necessary, multiply the ratio by an integer to obtain whole numbers (e.g., empirical formula of glucose: CH₂O)
• The empirical formula may or may not be the same as the molecular formula depending on the actual number of atoms in the compound
• Empirical formulas are essential in stoichiometry calculations and writing balanced chemical equations

Molecular formulas using empirical formulas

• Represents the actual number of atoms of each element in a compound determined using empirical formulas and molar mass information
• Steps to determine molecular formula:
  1. Calculate the empirical formula mass by adding the atomic masses of each element in the empirical formula (e.g., CH₂O = 12.01 + 2(1.01) + 16.00 = 30.03 g/mol)
  2. Divide the molar mass of the compound by the empirical formula mass to obtain a ratio using the formula: $\text{Ratio} = \frac{\text{Molar mass of compound}}{\text{Empirical formula mass}}$ (e.g., glucose molar mass = 180.16 g/mol, ratio = 180.16 / 30.03 = 6)
  3. Multiply the subscripts in the empirical formula by the ratio to obtain the molecular formula (e.g., glucose molecular formula: (CH₂O)₆ = C₆H₁₂O₆)
• The molecular formula will always be a whole-number multiple of the empirical formula (e.g., benzene: empirical formula = CH, molecular formula = C₆H₆)
• Subscripts in molecular formulas indicate the number of atoms of each element present in a single molecule

Additional Methods for Determining Formulas

• Combustion analysis: A technique used to determine the empirical formula of organic compounds by completely burning the sample and measuring the masses of CO₂ and H₂O produced
• Mass spectrometry: An analytical technique that can determine the molecular mass of a compound, which is crucial for confirming molecular formulas