Empirical FormulaEmpirical Formula

Empirical: based on observation and Empirical: based on observation and experimentexperiment

Empirical FormulaEmpirical 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

Ionic FormulaIonic Formula

• Ionic formula always equals empirical formula

• Ionic compounds are always simple, whole-number ratios of elements

• Examples:

– FeS

– Ammonium Phosphate

– CaCO3

Determining Empirical FormulaDetermining Empirical Formula

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

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

Moles carbon = 27.29 g x 1 mol C = 2.27 moles C

1 12.0 g

Moles oxygen = 72.71 g x 1 mol O = 4.54 moles O

1 16.0 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

C = 2.27 mol = 1 O= 4.54 mol = 2

2.27 mol 2.27 mol

Therefore, the empirical formula of this compound = CO2

Example #2

If 2.5 g of Al is heated with 5.28g of F, what is the EF of the resulting compound?

2Al + 3F2 2AlF3

Empirical Formula

2Al + 3F2 2AlF3

2.50g 5.28g Law of Conservation of Mass = 2.50g 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

Al 32.1% F 67.9%

32.1g 67.9g

Determine how many moles of each you have

Al

F

Molecular FormulaMolecular 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

• From empirical formula, empirical formula mass (efm) can be determined

• Example: HO = 17.0 g/mol

• Molar mass is determined experimentally• Example: 34.0 g/mol

• Number of empirical formula units can be determined by these two values

Molar Mass = Empirical Formula Multiplier efm

Determining Molecular FormulaDetermining Molecular Formula

• 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