# Empirical: based on observation and experiment

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

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

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

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?

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
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

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

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

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

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

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

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

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