10-1 Notes
The Mole – A measurement of matter
Measuring Matter
 Chemistry is a quantitative science
 We will be analyzing the composition of
samples of matter and perform chemical
calculations that relate quantities of reactants
in a chemical reaction to the quantities of
products
 Ex: How many grams of H2 and N2 must be
combined to make 200 grams of fertilizer (NH3)?
 1st we must find a way to measure the
amount of matter we have.
 By counting how much matter, finding the
mass, or finding the volume of matter
What is a Mole?
 Review: Matter is composed of atoms,
molecules, and ions
 Particles that are much to small to count
 We use a unit called a mole (mol) to count
particles of matter
 A mole of a substance is 6.02 x 1023
representative particles of that substance
 Known as Avogadro’s Constant
see table 10.1
 Ex. 1 mole of Na = 6.02 x 1023 atoms of Na
Mole Conversions
 How many moles of Mg are in 1.25 x 1023
atoms of Mg?
 Use this conversion: 1 mol = 6.02 x 1023
 Answer: 0.208 mol Mg
 Practice Problem # 3
 How many atoms are in 2.12 mol of
propane?
 Conversion – moles  molecules  atoms
 Answer: 1.40 x 1025 atoms
 Practice Problem #5
The Mass of a Mole
 Review: The atomic mass of an element is
expressed in atomic mass units (amu).
 The atomic masses are relative values based
on the most common isotope of carbon.
 The atomic mass of C is 12 amu which is 12
times heavier then H (mass of 1 amu)
 Results in a constant 12:1 ratio
Therefore, 12 grams of C and 1 gram of H
must contain the same number of atoms
The Mass of a Mole
 The atomic mass of an element expressed
in grams is the mass of a mole of the
element
Ex. The molar mass of Carbon is 12
grams
 Because 1 mole of any element contains
6.02 x 1023 atoms, the number of atoms
in 12 grams of Carbon = the number of
atoms in 16 grams of Oxygen
 Molar Mass is the mass of 1 mol of atoms
of any element
The Mass of a Mole
Step 1: Determine the formula for the
compound
Step 2: Find the number of grams of each
element in one mole of the compound
Step 3: Add the masses of the elements in
the compound
This method is used for any type of
compound
The Mass of a Mole
Find the mass of 1 mol of Sulfur trioxide
SO3 = 1 sulfur atom and 3 oxygen atoms
1 atom sulfur x 32.1 amu = 32.1 amu
3 atoms oxygen x 16.0 amu (per oxygen) =
48.0 amu
32.1 amu + 48.0 amu = 80.1 amu for Sulfur
trioxide
 Substitute grams for amu.
1 mol of SO3 has a mass of 80.1 grams
 Practice Problem # 7
10-2 Notes
Mole-Mass and Mole-Volume Relationships
Mole-Mass Relationship
Use the molar mass of an element or
compound to convert between the mass
of a substance and the moles of that
substance.
Mass (grams) = number of moles x
(mass / 1 mole)
Ex. The molar mass of NaCl is 58.5 grams.
What is the molar mass of 3 moles of NaCl?
3.00 mol NaCl x 58.5g/1 mol = 176 g
Sample problem 10.5
Mole-Mass Relationship
 Suppose in the lab you were able to obtain
10.0 grams of Na2SO4.
 How many moles is this?
 moles = mass x (1mol/mass)
 moles of Na2SO4 = 10.0 grams x (1 mol/142.1 g)
 moles of Na2SO4 = 7.04 x 10-2
 Sample problem 10.6
Mole-Volume Relationship
 The volumes of one mole of different solid
and liquid substances are not the same
 fig. 10.7 on page 295
 The volumes of gases are much more
predictable
 In 1811 Avogadro stated that equal
volumes of gases at the same temperature
and pressure contain equal number of
particles
 Even though the particles that make up
different gases are not the same size, they
take up relatively the same space
Mole-Volume Relationship
 The volume of a gas will vary depending on the
temperature or the pressure
 Increase in pressure = decrease in volume
 The volume of a gas is usually measured at
standard temperature and pressure (STP)
 STP is at O oC and at a pressure of 101.3 kPa
or 1 atmosphere (atm)
 At STP 1 mol of any gas occupies a volume of
22.4 L
 This is referred to as molar volume
Mole-Volume Relationship
 To calculate volume of a gas at STP:
 Volume of gas = moles of gas x (22.4 L/1
mol)
 Determine the volume (Liters) of 0.60 mol SO2 at
STP
 You can also calculate molar mass from
Density.
 The density of a gas at STP is 22.4L/mol
 Molar mass = density @ STP x molar
volume @ STP
 Determine the mass of a compound that has
a density of 1.964 g/L at STP
10-3 Notes
Percent Composition
and Chemical Formulas
Percent Composition
 The % composition of a compound consists of a
% value for each different element in the
compound
 % composition from mass data
 The % mass of an element is the # of grams of the
element divided by the mass in grams of the
compound, multiplied by 100
 % mass of element =
mass of element
x 100
mass of compound
 Sample problem 10.9
Percent Composition
 % Composition from the Chemical Formula
 % mass =
Mass of element in 1 mol compound
Molar mass of compound
 Sample Problem 10.10
x 100
% Composition as a Conversion
Factor
 You can use % comp. to calculate the
number of grams of any element in a specific
mass of a compound
 Multiply the mass of the compound by a
conversion factor based on the % comp of
the element in the compound.
 Sample problem:
 In 10.10 we found in propane, the % comp of
Carbon was 81.8%. What would the mass of
Carbon be in an 82.0 g sample?
Empirical Formulas
 The % comp. of any compound is the data we use to
calculate the basic ratio of elements contained in a
compound
 The basic ratio that gives the lowest whole-number
ratio of the atoms of the elements in a compound is
called the empirical formula
 An empirical formula may or may not be the same as
a molecular formula
 Example: The molecular formula for hydrogen peroxide
is H2O2 the empirical formula is HO
 Sample problem 10.11
Molecular Formulas
 The molecular formula of a compound is
either the same as its experimentally
determined empirical formula, or it is a
simple whole-number multiple of its
empirical formula
 To find the molecular formula you must
first know the molar mass of the
compound.
 Sample Problem 10.12