1. Empirical: based on observation and experiment
Empirical Formula
Empirical: based on observation and experiment

2. Empirical Formula
The lowest, whole number ratio of the atoms in a compound
The empirical formula of a compound does not always equal the molecular formula
Example: Hydrogen Peroxide
Molecular Formula = H2O2
Empirical Formula = HO

3. Ionic Formula Ionic formula always equals empirical formula
Ionic compounds are always simple, whole-number ratios of elements
Examples:
FeS
Ammonium Phosphate
CaCO3

4. Determining Empirical Formula
Example: A compound has a percent composition of 27.29% carbon and 72.71% oxygen. What is the compound’s empirical formula?

5. STEP ONE: Assume sample size is 100g
STEP TWO: Determine how many grams of each element are present using percent composition
27.29g C
72.71g O
STEP THREE: Determine the number of moles of each element in the sample

6. Moles carbon = 27.29 g x 1 mol C = 2.27 moles C
Moles oxygen = g x 1 mol O = 4.54 moles O g
STEP FOUR: Convert the ratio of moles to the lowest whole number ratio by dividing each number by the lowest number of moles present

7. Therefore, the empirical formula of this compound = CO2
C = 2.27 mol = O= 4.54 mol = 2
2.27 mol mol
Therefore, the empirical formula of this compound = CO2

8. Example #2
If 2.5 g of Al is heated with 5.28g of F, what is the EF of the resulting compound?
2Al F2 2AlF3

9. Empirical Formula 2Al + 3F2 2AlF3 Change into grams 2.50g 5.28g
Law of Conservation of Mass = g Al and 5.28g F
Total mass of the compound = 7.78g
Al = 2.50g/7.78g x 100% = 32.1%
F = 5.28g/7.78g x 100% = 67.9%
Change into grams

10. Al % F %
32.1g g
Determine how many moles of each you have Al F

11. Molecular Formula
Either the same as empirical formula or a simple, whole number multiple of its empirical formula
Example: Benzene
Empirical = CH
Molecular = C6H6
Example: Methanol
Empirical = CH4O
Molecular = CH4O

12. Determining Molecular Formula
From empirical formula, empirical formula mass (efm) can be determined
Example: HO = 17.0 g/mol
Molar mass is determined experimentally
Example: g/mol
Number of empirical formula units can be determined by these two values
Molar Mass = Empirical Formula Multiplier
efm

13. Example: HO 34.0 g/mol = 2 17.0 g/mol
Therefore, the empirical formula of HO needs to be multiplied by two in order to find the molecular formula:
(HO)x2= H2O2