Empirical and Molecular Formulas
% Composition
 The percent by mass of each element in a compound
42. ) 36.11% Ca; 63.89% Cl
43. ) 32.37% Na, 22.58%S, 45.05% O
44. ) H2SO3
45. ) 3.08% H, 31.61% P, 65.31% O
Empirical Formula
 The formula with the lowest whole-number mole
ratio of the elements
EX: CH2O is the empirical formula glucose
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.
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.
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”.
Determining the Molecular
Formula
To determine the molecular formula of a compound,
1. Calculate the following:
Experimentally determined molar mass
mass of the empirical formula
2. Multiply that answer through the empirical formula
molecular formula = (empirical formula)n
Determining the Formula for a
Hydrate
1.
2.
3.
4.
5.
6.
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 11-14 on page 340.