1. Understanding the Terms
* Empirical Formula: The simplest whole-number ratio of atoms in a compound. It tells you the relative proportions of each element present.
* Molecular Formula: The actual number of atoms of each element present in a molecule. It shows the exact composition of the molecule.
2. Determining the Empirical Formula
Steps:
1. Convert percentages to grams: If you're given percentages by mass, assume a 100 g sample. This means the percentages directly translate to grams.
2. Convert grams to moles: Divide the mass of each element by its molar mass (found on the periodic table) to get the number of moles.
3. Find the simplest whole-number ratio: Divide each of the mole values by the smallest mole value. Round these results to the nearest whole number. These whole numbers represent the subscripts in the empirical formula.
Example: A compound is found to be 85.7% Carbon and 14.3% Hydrogen by mass.
* Assume 100 g: 85.7 g C and 14.3 g H
* Moles:
* C: 85.7 g / 12.01 g/mol = 7.14 mol C
* H: 14.3 g / 1.01 g/mol = 14.2 mol H
* Ratio:
* C: 7.14 mol / 7.14 mol = 1
* H: 14.2 mol / 7.14 mol = 2
* Empirical Formula: CH₂
3. Determining the Molecular Formula
Steps:
1. Calculate the empirical formula mass: Add the atomic masses of the atoms in the empirical formula.
2. Determine the molecular mass: You'll usually be given the molecular mass. If not, you can use experimental techniques like mass spectrometry.
3. Find the ratio between the molecular mass and the empirical formula mass: Divide the molecular mass by the empirical formula mass.
4. Multiply the subscripts in the empirical formula by the ratio found in step 3: This gives you the molecular formula.
Example: Let's say the molecular mass of the compound in the previous example is 56 g/mol.
* Empirical Formula Mass: CH₂ = 12.01 g/mol + (2 * 1.01 g/mol) = 14.03 g/mol
* Ratio: 56 g/mol / 14.03 g/mol ≈ 4
* Molecular Formula: CH₂ * 4 = C₄H₈
Key Points:
* Empirical formulas provide the simplest representation of a compound's composition.
* Molecular formulas give the actual number of atoms in a molecule.
* If the empirical formula and molecular formula are the same, the compound has a simple, whole-number ratio of atoms.
Let me know if you'd like to work through another example!
