Download Chapter 11: The Mole1 In what quantity do you purchase the

Chapter 11: The Mole1
In what quantity do you purchase the
following?
Eggs?
Shoes?
Paper?
Chemists have created their own counting
unit to measure atoms, molecules, or formula
units in a sample.
The mole is the unit used to measure the
amount of a substance.
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The mass of a mole of one element will be
different from the mass of a mole of
another element because each element
has a different composition.
Remember the atomic mass unit?
An atom of carbon-12 has a mass of 12
amu.
Roses?
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Mass and the Mole
Abbreviated mol
SI base unit to measure the amount
of a substance
Number of representative particles
of carbon-12 in 12 g of C-12
6.02x1023 particles; 602 billion
trillion
known as Avogadro’s number, in
honor of the scientist who
determined the volume of one mole
of gas (1811). The number he
determined is the same for all gases.
How big is a mole?
1. Enough 12-oz. aluminum cans to cover
the surface of the earth to a depth of
over 200 miles.
2. Enough of unpopped popcorn kernels to
spread over the U.S. to a depth of more
than 9 miles.
3. Enough marbles to cover the surface of
the earth to a depth of more than 6 Km.
Mass related to the mass of a mole
The mass in grams of one mole of any
pure substance is called its molar mass.
Units for molar mass is g/mol.
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Equal to the numerical value of the
average atomic mass (from PT)
o 1 mole of C atoms = 12.0 g
o 1 mole of Mg atoms = 24.3 g
o Called a gram formula mass
when used with ionic compounds
Find the molar mass (round to the tenths
place) for:
1 mole of Br atoms = __________________
1 mole of Sn atoms = __________________
Molar mass of molecules and compounds
Mass in grams of 1 mole of a compound
equals the sum of the atomic masses of
the elements involved.
1 mole of CaCl2 = ____________________
1 mole of N2O4 = ____________________
1 mole of CH4 = _____________________
Converting Moles to Particles and Vice-versa
Converting Moles to Grams
# of moles x 6.02 x 1023 particles = particles
1 mole
Grams = moles x molecular mass (in g)
1 mole
# particles x
2 moles of CuSO4 = __________________g
1 mole________ = moles
6.02x1023 particles
3 moles of NaCl = ___________________g
Chapter 11: The Mole2
Converting Grams to Moles
Converting from mass to atoms
Moles = ____grams_____
molecular mass
You cannot make a direction conversion
from the mass of a substance to the number
of representative particles (atoms) in that
substance.
779 g CuSO4 = _______________________mol
525 g ZnCl2 = _______________________mol
Atoms to mass conversion
# Atoms x
1 mole___ x molar mass = __g
6.02x1023 atoms
1 mole
1.2 x 1025 atoms of N = __________g
You have to convert to moles first, then
moles to atoms.
Step 1: Mass x _____mole_____ = # moles
molecular mass
Step 2: #moles x 6.02x1023 atoms = #atoms
mole
25.0 g Au = _?__ atoms
35.4 g Cu = _?__ atoms
5.5 x 1022 atoms He = ______g
Wrap up!
Mass must always be converted to
moles before being converted to atoms.
Atoms must always be converted to
moles before calculating their mass.
Chapter 11: The Mole3
Chemical Compounds in Moles
Molar Mass of Compounds
You can also have a mole of different
compounds.
To determine the molar mass of a
compound, use the periodic table to
determine the mass of one mole of each
element in the compound.
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Each compound will have the sum of
the molar masses of the component
elements.
In one mole of a compound, the ratio of
moles of each element is the same as
for one molecule.
o Ex. In a molecule of freon (CCl2F2),
you have 1 atom of C, 2 atoms of Cl,
and 2 atoms of F.
In a mole of Freon, you would have
1 mole of C, 2 moles of Cl, and 2
moles of F.
You have 1.25 moles of aluminum oxide
(Al2O3). How many moles of aluminum ions
do you have?
Ex. Potassium chromate (K2CrO4)
2 mol K x 39.1 g =
1 mol K
78.2 g
1 mol Cr x 52 g =
1 mol Cr
52 g
4 mol O x 16 g =
1 mol O
64 g
Molar mass of K2CrO4 =
194.2 g
Determine the molar mass of methane (CH4)
Converting Moles to Mass
Converting Mass to Moles
To convert moles to mass in grams, you
first determine the molar mass of the
compound. Then, you multiply by the
number of moles given.
To convert mass of a compound to moles,
determine the compound’s molar mass
and set up your conversion factor.
What is the mass of 2.50 moles of allyl
sulfide, (C3H5)2S?
How many moles are in 325 g of Ca(OH)2?
Chapter 11: The Mole4
Converting Mass of a Compound to Particles
Recap:
Remember, you can’t directly convert
mass to particles; you must go through
moles, first.
1. Molar mass (# of grams/1 mol) and
the inverse (1 mol/# of grams) are
the conversion factors between the
mass of a substance and the number
of moles of the substance.
How many particles of Al3+ and Cl1- are in
35.6 g of AlCl3?
What is the first step:
Find molar mass of AlCl3
Second step:
Find # of moles of AlCl3 in 35.6 g
Third step:
Find the ratio of Al in AlCl3
Fourth step:
Find the ratio of Cl in AlCl3
How would you find the mass in g of one
formula unit of AlCl3?
Start with the molar mass of AlCl3, and use
the inverse of Avogadro’s number as a
conversion factor.
Empirical and Molecular Formulas
2. Avogadro’s number and its inverse
are the conversion factors between
the moles of a substance and the
number of representative particles.
3. To convert between the number of
moles of a compound and the
number of moles of atoms or ions
contained in the compound, you
need the ratio of moles of atoms (or
ions) to 1 mole of compound (or its
inverse).
a. The ratios are shown in the
subscripts in the chemical
formula.
Percent means parts
per 100, so the
percents by mass of all
the elements of a
compound must always
add up to 100.
Percent Composition
Chemical Formulas of Compounds
Percent by mass of each element in a
compound
Formulas give the relative numbers of
atoms or moles of each element in a
formula unit.
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Molar mass of element x 100 = % mass
Molar mass of compound
Find the percent composition of
Fe2O3 = 160 g
Fe = 56 x 2 = 112 g
112 x 100 = 70%
160
CaCl2
H3PO4
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O = 16 x 3 = 48 g
48 x 100 = 30%
160
Always a whole number ratio (law
of definite proportions)
Ex. NO2: 2 atoms of O for every
atom of N
1 mole of NO2 = 2 moles of O
atoms to every 1 mole of N atoms
If we know, or can determine, the
relative number of moles of each
element in a compound, we can
determine a formula for the compound.
Chapter 11: The Mole5
Empirical Formula
Writing an Empirical Formula
The empirical formula of a compound that
expresses the smallest whole number mole
ratio of the elements in the compound.
1. Determine the mass in grams of
each element present, if necessary.
a. If the problem gives you a
percent composition, assume
a 100 g sample.
2. Calculate the number of moles of
each element from the masses.
3. Divide each of the elements by the
smallest number of moles to obtain
the simplest whole number ratio.
4. If whole numbers are not obtained
in step 3, multiply through by the
smallest number that will give all
whole numbers.
a. Note: do not round off
numbers prematurely!
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Ratio provides the subscripts in the
empirical formula
Ionic formulas are always empirical
formulas
The empirical formula may (or may
not) be the same as the molecular
formula)
o If different, the molecular
formula will always be a simple
multiple of the empirical
formula
Molecular Formula
The molecular formula states the actual
number of each kind of atom found in the
compound.
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More than one compound can have the
same empirical formula
Molecular formula = empirical formula
multiplied by a factor by which
everything is multiplied.
Practice time!
Determine the empirical formula for a
compound that has 40.05% S and 59.95%
O.
To determine the molecular formula:
1. Calculate the empirical formula
2. Calculate the empirical formula mass by
multiplying number of moles by the molar
mass
3. Divide the given molar mass by the mass
of the empirical formula
4. Multiply subscripts in the empirical
formula by the answer in #3.
Molecular formula = (empirical formula)n
Practice time!
Succinic acid is composed of 40.68% carbon,
5.08% hydrogen, and 54.24% oxygen and has a
molar mass of 118.1 g/mol. Determine the
empirical and molecular formulas for succinic
acid.
Work space!
Chapter 11: The Mole6
Chapter 11: The Mole7
Hydrates
Determining the Formula of a Hydrate
Hydrates are solids in which water
molecules are trapped
1. Write the formula for the solid
2. Determine the mass of the water
(original mass ― anhydrous mass)
3. Calculate the moles of water
(Mass/18g)
4. Calculate the moles of compound
5. Determine the amount of water
(mol H2O/mol compound)
6. Rewrite the formula
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A compound that has a specific
number of water molecules bound
to its atoms
Ex. Opal – colors is the result of the
presence of water in the mineral
Anhydrous – means water has been
removed from a hydrate
Example
You have a 5.0-g sample of a hydrate of
barium chloride.
1. Formula: BaCl2·xH2O [you don’t know
x]
Uses for Hydrates
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Used to store solar energy
Anhydrous materials are used as
desiccants (drying agents)
Naming Hydrates
When hydrates are named, the number of
water molecules associated with each
formula unit of the compound is written
following a dot:
MgSO4 · 4H2O (magnesium sulfate
tetrahydrate)
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Prefixes are the same as used when
naming covalent compounds
1
mono2
di3
tri4
tetra5
penta6
hexa7
hepta8
octa9
nona10
decaMass of water must be included in
all calculations
After heating, the mass of the anhydrous
BaCl2 has a mass of 4.26 g.
2. To find x, subtract the anhydrous mass
from the original mass.
5.0g – 4.26g = 0.74 g H2O
3. Calculate the moles of water
.74 g x 1 mol = 0.041 mol H2O
18 g
4. Calculate the moles of the compound
4.26 g x 1 mol = 0.02 mol BaCl2
208.3 g
5. Mol of H2O = .041 mol H2O = 2
Mol of BaCl2
.02 mol BaCl2
6. Rewrite the formula: BaCl2·2H2O
Practice
2.50 g of hydrated copper sulfate
(CuSO4·xH2O) reduces to 1.59 g after heating.
What is the formula for the copper sulfate
hydrate?