1. Determining the Empirical Formula for a Compound
The empirical formula for a compound is the simplest __________ number __________ of the atoms in the compound.
Examples: H2O is the empirical formula for water. (can’t be reduced)
Glucose, C6H12O6, has the empirical formula ____________.
Step 1: Change % to grams.
Step 2: Convert grams to moles using the mole chart.
Step 3: Divide each of these answers by the smallest number.
Step 4: Your answers will be the subscripts for the empirical formula.
whole
ratio
C1H2O1
Helpful Rhyme: % to mass, mass to mole, divide by small, times ’til whole.

2. Assume 100g since % , just change to % sign to g.
Practice Problems:
Calculate the empirical formula if a compound is found to be composed of 7.8% Carbon and 92.2% Cl
Assume 100g since % , just change to % sign to g.
Find moles of each element
7.8g C mole =
92.2g Cl 1 mole =
(% to mass)
(mass to moles)
0.65 moles C
12.0 g C
2.63 moles Cl
35.0 g Cl

3. Divide each mole found by lowest mole. (Round to whole #’s)
0.65 mole C = mole Cl =
4. Final answer, putting numbers in as subscripts. 1 4
(÷ by small)
C1Cl4
CCl4

4. Assume 100g since % , just change to % sign to g.
Practice Problems:
Calculate the empirical formula if a compound is found to be composed of 42.9% C and 57.1% O
Assume 100g since % , just change to % sign to g.
Find moles of each element
42.9g C mole =
57.1g O 1 mole =
(% to mass)
(mass to moles)
3.57 moles C
12.0 g C
3.57 moles O
16.0 g O

5. Divide each mole found by lowest mole. (Round to whole #’s)
3.575 mole C = mole O =
4. Final answer, putting numbers in as subscripts. 1 1
(÷ by small)
C1O1 CO

6. Determining the Molecular Formula for a Compound
The molecular formula for a compound is either the same as the empirical formula ratio or it is a “_________ _________ of this ratio. It represents the true # of atoms in the molecule.
Examples: 1) H2O is the empirical & molecular formula for water ) CH2O is the empirical formula for sugar, ethanoic acid, and methanol. The molecular formula for glucose is C6H12O6, (___times the empirical ratio!)
Step 1: Determine the empirical formula for the compound. (See the previous steps in the notes.)
Step 2: Determine the “whole # multiple” by dividing the molecular formula mass (given in the problem) by the empirical formula molar mass. Multiply each of the empirical ratios by this whole number.
whole # multiple 6

7. H5.9O5.9 H1O1 HO 5.9 5.9 H2O2 Practice Problems:
1) An unknown compound is composed of 5.9% hydrogen and 94.1% oxygen. The molecular formula mass is 34 g. Determine the molecular formula for the compound.
H = 5.9% = 5.9 g
(mass to moles)
5.9 g H ÷ 1.0 = 5.9 moles H
O = 94.1% = 94.1 g
94.1 g O ÷ 16.0 = 5.9 moles O
(÷ by small)
H5.9O5.9
H1O1 HO
5.9
5.9
H2O2
[H1O1 ]x 2 =
Molar mass of H1O1 = 17 g
(34 ÷ 17 = 2)
Molecular formula mass divided by molar mass

8. C3.3H6.6O3.3 C1H2O1 CH2O 3.3 3.3 3.3 C4H8O4 Practice Problems:
2) An unknown compound is composed of 40% carbon, 6.6% hydrogen, and 53.4% oxygen. Determine the molecular formula for the compound if the mass of one mole of the compound is 120 g.
C = 40% = 40 g
(mass to moles)
40 g C ÷ 12.0 = 3.3 moles C
H = 6.6% = 6.6 g
6.6 g H ÷ 1.0 = 6.6 moles H
O = 53.4% = 53.4 g
53.4 g O ÷ 16.0 = 3.3 moles O
(÷ by small)
C3.3H6.6O3.3
C1H2O1
CH2O
3.3
3.3
3.3
(Compare)
C4H8O4
[C1H2O1 ]x 4 =
Our formula mass = 30 g
(120 ÷ 30 = 4)