Download Chapter 11.1

Chapter 11.1
The Mole
Mole
SI base unit to measure the amount
of a substance
 Equal to Avogadro's number (NA)
(6.02 x 1023) of things i.e. atoms,
pennies, pencils, etc.

The Mole
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A counting unit
Similar to a dozen, except instead
of 12, it’s 602 billion trillion
602,000,000,000,000,000,000,000
6.02 X 1023 (in scientific notation)
This number is named in honor of
Amedeo Avogadro (1776 – 1856),
who studied quantities of gases
and discovered that no matter what
the gas was, there were the same
number of molecules present
Just How Big is a Mole?
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Enough pop cans to cover the
surface of the earth to a depth of
over 200 miles.
If you had Avogadro's number of
unpopped popcorn kernels, and
spread them across the United
States of America, the country would
be covered in popcorn to a depth of
over 9 miles.
If we were able to count atoms at the
rate of 10 million per second, it
would take about 2 billion years to
count the atoms in one mole.
Converting Moles to Particles
# of moles x 6.02 x 1023 particles
1 mole
Practice Problems p. 311
Assignment
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Determine the number of representative particles for each of the
following:
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11.5 mol Ag
18 mol water
0.15 mol NaCl
3.75 mol Fe
12.5 mol CaCO3
How many moles of CaCl2 contains 1.26 x 1024 formula units of
CaCl2
How many moles of Ag contains 4.59 x 1025 formula units of Ag
Chapter 11.2
Molar Mass
Molar Mass

The Mass of 1 mole (in grams)

Equal to the numerical value of the average
atomic mass (get from periodic table)
1 mole of C atoms = 12.0 g
1 mole of Mg atoms =
24.3 g
Some people used to call this a (gram
formula mass) when used with ionic
compounds
Find the molar mass
(round to the tenths place)
A.1 mole of Br atoms = 79.9 g/mole
B.1 mole of Sn atoms= 118.7 g/mole
Molar Mass of Molecules and
Compounds
Mass in grams of 1 mole equal
numerically to the sum of the atomic
masses
1 mole of CaCl2
= 111.1 g/mol
1 mole of N2O4 = 92.0 g/mol
Converting Moles to Grams

grams = moles X molecular mass
2
moles of CuSO4 = _____g
 G = 2 x 159.6 = 319.2 g
Converting Grams to Moles
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moles =
 779
grams
molecular mass
g CuSO4 = _____ mol
 Mol = 779/159.6 = 4.88 moles
Atoms/Molecules and Grams
Since 6.02 X 1023 particles = 1 mole
AND
1 mole = molar mass (grams)
 You can convert atoms/molecules to
moles and then moles to grams!
(Two step process)

Atoms/Molecules and Grams
You can’t go directly from atoms to
grams!!!! You MUST go thru MOLES.
 That’s like asking 2 dozen cookies
weigh how many ounces if 1 cookie
weighs 4 oz? You have to convert to
dozen first!
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Converting Grams to Atoms
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# grams/molar mass x NA
Atoms/Molecules and Grams
How many atoms of Cu are present in
35.4 g of Cu?
35.4 g Cu
1 mol Cu
63.5 g Cu
6.02 X 1023 atoms Cu
1 mol Cu
= 3.4 X 1023 atoms Cu
How many atoms of K are present in
78.4 g of K?
78.4 g K
1 mol K
39.1 g K
6.02 X 1023 atoms K
1 mol K
= 1.20 X 1024 atoms K
Converting Atoms to Grams
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# atoms/NA x Molar Mass
 1.2
x 1025 atoms of N = _____ g
x 1025 x 14 = 279.1
6.02 x 1023
 1.2
Assignment
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Get a book
Questions 11-19
Chapter 11.3
Molar
Relationships
Molar Mass
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 (p.321 20-24)
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Chapter 11.4
Percent
Composition and
Formulas
Percent Composition
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Percent by mass of each
element in a compound
Mass of Element x100=%mass
Mass of Compound
Percent Composition
Fe2O3 = 160 g
Fe= 56 x 2 = 112 g 112/160x100 = 70%
48/160x100 = 30%
O = 16 x 3 = 48 g
CaCl2
H3PO4
Chemical Formulas of Compounds

Formulas give the relative
numbers of atoms or moles of
each element in a formula unit always a whole number ratio (the
law of definite proportions).
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
Chemical Formulas of
Compounds
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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.
Empirical Formula
The formula of a compound that
expresses the smallest whole
number ratio of the atoms
present.
 Ionic formula are always
empirical formula
 May or may not be the same as
the molecular formula

To obtain an Empirical Formula
1. Determine the mass in grams of
each element present, if
necessary.
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.
To obtain an Empirical
Formula
3. Divide each 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
*
Be careful! Do not round off
numbers prematurely
Assignment
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A blue solid is 36.8% nitrogen and 63.2% oxygen what is the
empirical formula?
Determine the empirical formula for a compound that contains
36% aluminum and 64% Sulfur
Propane is 81.8% carbon and 18.2% hydrogen, what is the
empirical formula for propane?
Aspirin is 60% carbon, 4.4% hydrogen, and 35.6% oxygen, what
is the empirical formula of aspirin?
Determine the empirical formula for a compound that contains
10.9% magnesium, 31.8% chlorine, and 57.3% oxygen
Molecular Formula
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The formula that states the actual
number of each kind of atom
found in the compound
Necessary since multiple
compounds can have the same
empirical formula
Molecular Formula =
Empirical Formula x factor by which
everything is multiplied
Determining Molecular
Formula
1.
2.
3.
4.
Calculate the empirical formula
Calculate the empirical formula mass
by multiplying number of moles by
molar mass
Divide the given molar mass by the
mass of the empirical formula
Multiply subscripts in the empirical
formula by the answer in #3
Assignment
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Page 335 &337, #’s 51-62
Chapter 11.5
Hydrates
Hydrate
Solids that contain water
molecules
 Each hydrate has a specific
number of water molecules
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Anhydrous – Water has been
removed from a hydrate
Uses for Hydrates
Used to store solar energy
 Anhydrous materials are used
as desiccants
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Naming Hydrates
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# of water molecules follows
the solid formula and a dot
MgSO4 • 4H2O
Prefixes for water are on
page 338
Mass of water must be
included in all calculations
Determining the Formula of
a Hydrate
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Write the formula for the solid
Determine the mass of the water
(Original mass – anhydrous mass)
Calculate the moles of water
(mass/18g)
Calculate the moles of compound
Determine the amount of water
(mol H2O/mol compound)
Rewrite formula
Determining the Formula of
a Hydrate
If 2.50 g of copper
sulfate reduces
to1.59 g after
heating what is
the formula for
the copper sulfate
Hydrate?
1.
2.
3.
4.
5.
6.
CuSO4 • _ H2O
2.50 g – 1.59 g
= 0.91 g
0.91g/18g =
0.05 mols
1.59g/160g =
0.01 mols
0.05/0.01 = 5:1
CuSO4 • 5 H2O
Assignment
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Find the formula for the
following hydrates:
48.8% MgSO4 and 51.2% H2O
An 11.75 g sample of cobalt
(II) chloride hydrate is heated,
9.25 g of anhydrous cobalt
chloride remains.
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What is a hydrate?
Describe the experimental
procedure for finding the
formula of a hydrate.
Name the compound having
the formula SrCl26H2O
Arrange these hydrates in
order of increasing percent
water content:
MgSO47H2O,
Ba(OH)28H2O,
CoCl26H2O