1. Empirical and Molecular Formulas
Ch. 11, Sec. 4
Empirical and Molecular Formulas

2. % Composition The percent by mass of each element in a compound
) % Ca; 63.89% Cl
) % Na, 22.58%S, 45.05% O
) H2SO3
) 3.08% H, 31.61% P, 65.31% O

3. Empirical Formula
The formula with the lowest whole-number mole ratio of the elements
EX: CH2O is the empirical formula glucose

4. Molecular Formula
Formula that specifies the actual # of atoms of each element in one molecule or formula unit of the substance
EX: C6H12O6 is the molecular formula for glucose.

5. Steps to Determine an Empirical Formula
1.) Find the # of moles of each element. (Divide the amount of each element by its molar mass.) 2.) Divide each mole value by the smallest answer. 3.) Round to the nearest whole #, if it is close. This is the mole ratio of the elements and is represented by subscripts in the empirical formula. **If you get ___.5 at this step, multiply EACH by 2. **If you get ___.25 at this step, multiply EACH by 4. **If you get ___.33 or ___.66, multiply EACH by 3.

6. If you are given “%” of each element instead of a mass in grams, just replace the “%” with “g” and calculate as you normally do with “g”.

7. Determining the Molecular Formula
To determine the molecular formula of a compound,
Calculate the following:
Experimentally determined molar mass
mass of the empirical formula
Multiply that answer through the empirical formula
molecular formula = (empirical formula)n

8. Determining the Formula for a Hydrate
Write down ALL info., including the mass of both the hydrated compound AND the anhydrate
Calculate: hydrated cpd. - anhydrate = g H2O
Find the moles of the anhydrate and the moles of H2O
Get the mole to mole ratio (by dividing by the smaller number from #3 above)
Write the formula, using a “ “ (dot) to attach the H2O
Write the name of the hydrate
Look at Example Problem on page 340.