Chemical Formulas and the Mole
Introduction
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84 potatoes
63 carrots
15 onions
3 heads of garlic
27 turnips
42 pieces of celery
9 cans of green beans
6 cans of tomato puree
6 cans of diced peppers
9 cans of corn
6 cans of lima beans
Serves 81
• If only 1 head of garlic is
available, how many of each
ingredient will be needed to
adjust the recipe in order to
make the stew correctly and
how many people will it
feed?
• If you wanted to feed 324
people, how much of each
ingredient would you need?
How are formulas represented?
• Calculate the % composition of the formula you were
given.
• % composition comes from experimental analysis and
provides information about the make up of the
compound.
• Compare your calculations with the students around
you having a different colored card.
• What did you find regarding your % compositions?
• Although the formulas are different, are there any
similarities among your formulas that could account for
the similarities of your calculations?
Empirical Formulas
• The molecular formula of a compound gives the actual
number of atoms of each element making up a compound.
• (The molecular formula is the formula that is written on your
card)
• The empirical formula of a compound is the formula with the
smallest whole number ratio of the elements making up the
compound.
• What is the empirical formula for the compound written on
your card?
• The empirical formula may or may not be the same as the
actual molecular formula.
• If the two formulas are different, the molecular formula will
always be a simple multiple of the empirical formula.
Practice Problems
• Write the empirical formula for each of the
following molecular formulas.
1. C6H12O6
2. H2O2
3. C4H10
4. H2O
5. N2O4
6. C3H6O2
Determining Empirical Formulas
• If % composition can be determined from
the formula, then the formula can be
determined from the % composition of
the compound.
Rules for determining Empirical Formulas
1. Assume a 100 g sample, so % becomes “g”.
2. Convert grams to moles by dividing by the
atomic mass of each element.
3. Divide each result (mole) by the smallest
result present (mole ratio).
4. Look for whole number ratios.
5. The whole number ratios become the
subscripts for the formula.
Rhyme for Remembering Rules
“Percent to mass
Mass to mole
Divide by small
Multiply 'til whole“
• A Simple Rhyme for a Simple Formula
by Joel S. Thompson
Journal of Chemical Education
Vol. 65, No. 8; August 1988, p. 704
Sample Problems
• Determine the empirical formula for a
compound that is 27.3% carbon and 72.7%
oxygen.
• Determine the empirical formula for a
compound that is 56.6% K, 8.7% C, and
34.7% O.
• Determine the empirical formula for a
compound that is 69.9% Fe and 30.1% O.
Multiples for use when ratio is not
initially a whole number
Decimal
Multiplier
0.5
X2
0.30-.35
X3
0.63 - .67
X3
0.22-0.25
X4
0.72-0.75
X4
Molecular Formulas
• Molecular formulas are multiples of empirical
formulas.
Compound
Empirical
Formula
Molar Mass
Molecular
Formula
Formaldehyde
CH2O
30
CH2O
Acetic acid
CH2O
60
C2H4O2
Glucose
CH2O
180
C6H12O6
Sample Problems
• The empirical formula for a compound
containing phosphorus and oxygen was found
to be P2O5. Experiments show that the molar
mass of the compound is 284 g/mol. What is
the molecular formula and name of the
compound?
• Determine the molecular formula of a
compound having an empirical formula of CH
and a molar mass of 78.11 g/mol.
Putting it All Together
• A compound with a molar mass of 92 g/mol
contains 0.608 g of nitrogen and 1.388 g of
oxygen. What is the empirical and molecular
formula of the compound?
• A compound with a molar mass of 86 g/mol
contains 83.62% C and 16.38% H. What is the
molecular formula of the compound?
Hydrates
• Hydrates are solid ionic compounds in which
water molecules are trapped.
• Some products, such as electronic equipment,
are boxed with small packets labeled
dessicant. These packets control moisture by
absorbing water. Some contain ionic
compounds called hydrates.
• Each hydrate has a specific number of water
molecules bound to its atoms.
Example
• An opal is hydrated silicon dioxide (SiO2)
»
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The presence of water and
various mineral impurities
accounts for the variety of
colors.
Naming Hydrates
• In the formula for a hydrate, the number of water
molecules associated with each formula unit of the
compound is written following a dot.
• For example: Na2CO3 ∙ 10 H2O
• This compound is called sodium carbonate decahydrate.
• The prefix deca- means ten and the root word hydrate
refers to water.
• When naming hydrates use the same prefixes for the
water as used with covalent molecules.
• A decahydrate has ten water molecules associated with
each formula unit of compound.
• The number of water molecules associated with hydrates
varies widely.
Examples of Hydrates
Anhydrous (without water) CoCl2
The hydrate CoCl2 ∙ 6 H2O
Hydrating CuSO4
Hydrated CuSO4 ∙ 5H2O
Anhydrous CuSO4
Calculating the Formula Mass of Hydrates
Na2CO3 ∙ 10 H2O
Na: 2 x 23 = 46
C: 1 x 12 = 12
O: 3 x 16 = 48
H2O: 10 x 18 = 180
(H: 20 x 1 = 20; O: 10 x 16 = 160; 20 + 160 = 180)
Total = 286 g/mol
**Reminder-the “dot” in the formula means to
add the mass of the water (not multiply as in math)
Practice Problem
Calculate the formula mass of the following:
1. CuSO4 ∙ 5 H2O
Answer : 249.5 g/mol
2. BaCl2 ∙ 2 H2O
Answer : 244 g/mol
Determining the Formula of
Hydrates from Experimental Data
• The composition of a hydrate is determined to
contain 48.8% MgSO4 and 51.2% H2O.
• What is the formula of the hydrate?
48.8/120 (mass of MgSO4) = 0.407 moles MgSO4
51.2/18 (mass of H2O) = 2.84 moles H2O
0.407/0.407 = 1
2.84/0.407 = 7
Formula is MgSO4 ∙ 7 H2O
Let’s Try Another
• A 1.628 g sample of a sample of a hydrate of
magnesium iodide (MgI2) heated until its mass
is reduced to 1.072 g and all water has been
removed. What is the formula of the hydrate?
Answer:
1.072/278=0.00386 moles MgI2
1.628-1.072 = 0.556 g water/18 = 0.0309 moles H2O
0.00386/0.00386 = 1
0.0309/0.00386 = 8
Formula: MgI2 ∙ 8 H2O