Unit 5
CHEMICAL QUANTITIES
CHAPTER 10 – THE MOLE
How much sand?
3 Ways to Measure Matter
By COUNT – 1 million grains of sand
By MASS – 1,000 grams of sand
By VOLUME – 100 liters of sand
ATOMIC MASS:
What is the atomic mass of Hydrogen?
1.01 a.m.u.
What is the atomic mass of Oxygen?
15.999  16.0 a.m.u.
=
SO3
+
1 S atom
3 O atoms
You can calculate the mass of a molecule by
adding the atomic masses of the atoms making up
the molecule. (32.1 + 16 + 16 + 16 = 80.1 amu)
What is the atomic mass of Water
(H2O)?
H – 2 (1.0) = 2.0
a.m.u.
O – 1(16.0) =16.0 a.m.u.
18.0 a.m.u
What is the atomic mass of CO2?
C – 1(12.0) = 12.0 a.m.u
O – 2 (16.0) = 32.0 a.m.u.
 mass of CO2 = 44.0 a.m.u.
What is the atomic mass of Al2(SO4)3?
What is a MOLE?
A mole is a quantity equal to
Avogadro’s Number (6.02 x 1023)
6.02 x 1023 particles (atoms or
molecules) – depending on what you are
looking at.
A mole of anything contains the same
number of “things” as a mole of anything
else.
one mole is set by defining one mole of
carbon 12 atoms to have a mass of exactly
12 grams.
Create a Mole “Quantity” Poster
Poster states: “A mole of
contains 6.02 x 1023
.”
Surrounding that statement is a bunch of
pictures depicting your object.
On the back it also states:
“The weight of 1 mole depends on
the object”
“If it were a gas, 1 mole would
occupy 22.4 liters”
Grade based on completeness,
creativity, effort, and overall effect.
“Definition” of the molar mass:
Molar Mass = Mass of 1 mole
Weight, in GRAMS, is numerically
equal to what its “atomic” weight was in
a.m.u.’s
What is the molar mass of Water (H2O)?
18.0 a.m.u.’s
18.0 grams / mole
What is the molar mass of CO2?
44.0 grams / mole
Use
Molar Mass
Use 22.4 L
Use 6.02 x 1023
Practice Problem:
How many moles are in 5.0 g of copper?
1 mole Cu = 63.5 g
5.0 g of Cu
Should we have more or less than 1 mole?
5.0 g Cu x 1 mole Cu =
63.5 g Cu
5.0g Cu =
.079 moles Cu
63.5 g/mol
How many moles are in 26.0 g of
NH3?
1 mole NH3 = 14.0 + 3.0 = 17.0 g
More or less than 1 mole?
 26.0 g NH3 / 17.0 g/mol =
1.53 moles NH3
Going the other direction:
How many grams are in 1.5 moles of
H2O?
1 mole H2O = 18.0 g
1.5 mol H2O x 18.0g H2O =
1 mol H2O
1.5 mol H2O x 18.0 g/mol =
27 g H2O
Use
Molar Mass
Use 22.4 L
Use 6.02 x 1023
Volume of a Mole of Gas
The volume of gas varies with
temperature and pressure
At Standard temperature and pressure
(STP), a mole of ANY gas occupies
22.4L (molar volume)
Standard Temp. = 0 oC
Standard Pressure = 101.3 kpa
Practice Problem:
Sulfur Dioxide (SO2) is a gas
produced by burning coal. It is a
pollutant that causes acid rain.
Determine the volume, in liters, of
0.60 moles of SO2 gas at STP.
At STP, CO2 occupies 36.5L. How
many moles of CO2 gas is in the
sample?
Use
Molar Mass
Use 22.4 L
Use 6.02 x 1023
Scientific Notation Refresher:
Multiplication:
Multiply Coefficients
Add Exponents
Division:
Divide Coefficients
Subtract Exponents
Practice Problems:
How many molecules are in 1.5
moles of H2O?
9.03 x 10 23
Going the other direction:
How many moles are in 1.806 x 1024
atoms of Oxygen?
3 moles
Use
Molar Mass
Use 22.4 L
Use 6.02 x 1023
Tying it all TOGETHER:
How many molecules are in 300g of
Na2SO4?
How many grams are in 4.56 x 1023
atoms of Gold (Au)?
What volume, in Liters, would 40g of
O2 gas occupy at STP?
Percent Composition
and
Chemical Formulas
Percent Composition =
the percent by mass of each
element in a compound.
% mass of element =
Mass of element
Mass of compound
x 100%
Percent Composition
Calculate the formula mass of
C2H3O2?
What percent of this compound is
oxygen by mass?
% mass of element =
Mass of element
Mass of compound
32.0g x 100% = 54.2%
59.0g
x 100%
9.03 g Mg combine completely with
3.48 g N to form a compound. What is
the percent composition of this
compound? (9.03 g + 3.48 g = 12.51 g)
% Mg =
%N=
mass of Mg x 100 = 9.03 g x 100 = 72.2% Mg
mass of compound
12.51 g
mass of N x 100 = 3.48 g x 100 = 27.8 % N
mass of compound
12.51 g
Practice Problem:
Propane, C3H8, is commonly used in
gas grills. Calculate the percent
composition of this compound.
Calculating Empirical Formulas
Empirical Formula = the lowest
whole-number ratio of atoms of the
elements in a compound.
may or may not be the same as a
molecular formula.
H2O2 = molecular formula
HO = empirical formula
CO2 = molecular formula
CO2 = empirical formula
Calculating Molecular Formulas:
The molecular formula of a
compound can be determined if
you know its empirical formula and
its molar mass.
Example:
Calculate the molecular formula of the
compound whose molar mass is 60.0 g and
empirical formula is CH4N.
First calculate the empirical formula
mass.
Then divide the empirical formula mass
into the molar mass.
Multiply formula subscripts by this value
to get the molecular formula.