Empirical Formula
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Transcript of Empirical Formula
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