How do you measure matter?
 You may count how many you have.
 _____________________________
 Determine a substances mass and weight.
 _____________________________
 Determine a substances volume.
 _____________________________
 But how can you relate these three types of
measurements to one another?
How do you measure matter?
 Knowing how to relate the count, mass and volume
relate to a dozen apples allows you to convert between
units.
 In chemistry we will use the mole as our common
factor.
 The ___________________is the SI unit that measure the
amount of a substance.
The Number of Particles in a Mole
 One way to measure the amount of a substance is to
count the number of particles in the substance.
 What problem arises because atoms, molecules, and
ions are extremely small?
 The number of individual particles in a sample of any
substance would be _____________________________
The Number of Particles in a Mole
 How can you solve this problem?
 Just as a dozen equals 12 apples, a mole represents a certain
number of particles.
 1 mole of a substance equals 6.02 x 1023 particles of that
substance.
 This number is called _________________number = 6.02 x
1023
 Representative particle refers to the type present in a
substance:
__________________________________________________
__________________________________________________
 Usually the _____________ for most elements
The Number of Particles in a Mole
Substance
Representative
Particle
Chemical Formula
Representative
particles in 1 mol
Atomic nitrogen
Atom
N
6.02 x 1023
Nitrogen gas
Molecule
N2
6.02 x 1023
Water
Molecule
H2O
6.02 x 1023
Calcium ion
Ion
Ca2+
6.02 x 1023
Calcium fluoride
Formula unit
CaF2
6.02 x 1023
Sucrose
Molecule
C12H22O11
6.02 x 1023
How to determine how many atoms
are in a mole of a compound?
First you must know how many atoms are in a representative
particle of the compound.
The number is determined from the chemical formula.





A molecule of CO2 is composed of ________________.
A mole of CO2 contains Avogadro’s number of CO2 molecules.
Thus a mole of CO2 contains 3 times Avogadro’s number of
atoms.
To figure out the number of atoms in a mole of a
compound, you 1st determine the number of atoms in a
representative particle of that compound and then
multiply that number by Avogadro’s number.
Review
 1 mole = 6.02 x 1023 molecules
 Number of atoms = 1 molecule=6.02 x 1023
Problems
 How many moles is 2.80 x 1024 atoms of silicon?
1) Analyze
Known:
Number of atoms
 1 mole of Si= 6.02 x 1023
 Conversion= atoms to moles
Unknown: moles = ___________ moles of Si

2) Calculate
3) Evaluate
Make sure unit cancel out correctly.
Problems
 How many molecules is 0.360 mol of water?
1) Analyze
Known:
Number of moles
 1 mole of Water= 6.02 x 1023 molecules of water
 Conversion= moles to molecules
Unknown: molecules of water

2) Calculate
3) Evaluate
Make sure unit cancel out correctly.
1) How many moles are equal to 2.41 x 1024
formula units of sodium chloride (NaCl)?
2) How many moles are equal to 9.03 x 1024
atoms of mercury (Hg)?
3) How many atoms are equal to 4.5 moles
of copper (Cu)?
4) How many molecules are equal to 100.0
moles of carbon dioxide (CO2)?
5) How many atoms are in 1.00
moles of sucrose C12H22O11?
6) How many atoms of C are in 2.0
moles of C12H22O11?
7) How many atoms of H are in 2.0
moles of C12H22O11?
8) How many atoms of O are in
3.65 moles of C12H22O11?
 The gram atomic masses of any 2 elements contains the same
number of atoms, because the atomic masses of the elements
are relative values.
 In other words….
 Suppose that the mass of an atom of element X is twice as
great as the mass of an atom of element Y. Now suppose
that you have 10 grams of element X and 10 grams of
element Y. Would you expect both samples to contain the
same number of atoms? Why or Why not?
 ________________________________________________
________________________________________________
________________________________________________
________________________________________________
________________________________________________
________________________________________________
Review Continue
 What would you have to do to get the same number of
atoms in both samples?
 ____________________________________________________
 Although the mass of a single atom can be expressed in
atomic mass units (______) it isn’t realistic to work with
single atoms.
 Chemists work with large numbers of atoms (moles of
atoms) for which the mass can be determined in grams.
 The _______________________(gam) of an element is the
mass of an mole of atoms of that element expressed in
grams.
The Mass of a Mole of an Element
 By checking the atomic masses in the periodic table
you can find the gram atomic mass of the element.
 The gram atomic masses of any 2 elements must
contain the same number of atoms.
 How many atoms are contained in the gram atomic
mass of an element?
 _____________________________
What is the mass of a mole of a
compound?
To answer you need to know:
1.
_____________________________(remember this
tell you the number of atoms of each element in a
representative particle of the compound)
2. Then you can calculate the mass of a molecule by
adding the atomic masses of the atoms making up
the molecule. To give you the
_____________________________.
3. If you substitute the unit grams for atomic mass
units you will have the gram molecular mass.
Mass of a Mole of a compound
 The _____________________________(gmm) of any
molecular compound is the mass of 1 mole of that
compound.
 The gmm equals the molecular mass expressed in
grams.
Gram molecular mass may be
calculated directly from gam.
 For each element in a compound find the number of
grams of that element per mole of the compound.
 Then sum the masses of the elements in the
compound.
Practice
 Find the gram molecular mass of each of the following
compounds:
 C2H6
 N2O5
 C3H7OH
Practice
 What is the mass of 1.00 mol of each of the following
substances?
 Carbon tetrabromide
 Silicon dioxide
 Chlorine
Practice
 Determine the mass of 1 mole of the following
compounds
1.
2.
3.
4.
5.
6.
7.
8.
CO2
SO3
Br2
H2
N2
NaOH
Al2
Ba(NO3)2
 You do not calculate the gmm of





_____________________________.
The representative particle of an ionic compound is a
_____________________________, not a molecule.
The mass of one mole of an ionic compound is the
gram formula mass (gfm)
gfm= the formula mass expressed in grams.
Gfm is calculated the same way as the gmm by simply
taking the sum of the atomic masses of the ions in the
formula of the compound.
Out of gam, gmm, and gfm gfm is the most inclusive.
The term gfm can be used to refer to a mole of any
substance.
Review for 7.1
Explain what a mole of a substance represents.
1.

__________________________________________________
Demonstrate how to convert the number of atoms or
molecules of a substance to moles.
2.

___________________________________________________
Define atomic mass unit, and gram atomic mass, gram
molecular mass and gram formula mass
3.





___________________________________________________
___________________________________________________
___________________________________________________
Determine the molar mass of As2(CO3)3
___________________________________________________
Molar Mass of a Substance
 gam, gmm, and gfm are used to represent a mole of a




particular kind of substance.
Gam contains a mole of atoms
Gmm of a molecular compound contains a mole of
molecules
Gfm of an ionic compound contains a mole of formula
units.
We can use molar mass to refer to a mole of an
element, a molecular compound, or an ionic
compound because it is the mass in grams of one mole
of the substance.
Review
 Name the units that determine the mass of a mole
 ________________________________________________
 Which units should be used to describe the volume of a
mole?
 _______________________________________________
 Remember that unlike solids or liquids the molar
volume of gases is more predictable and is affected by
temperature and pressure.
Volume of a Mole of Gas
 How does temperature affect the volume of a gas?
 ________________________________________________
________________________________________________
________________________________________________
________________________________________________
 How does pressure affect the volume of a gas?
 ________________________________________________
________________________________________________
________________________________________________
________________________________________________
Volume of a Mole of gas
 When comparing the molar volumes of gases, it is




necessary to have the gases at the same conditions of
temperature and pressure.
If the gasses are at 0oC and 101.3 kPa
(_____________________), they are a standard
conditions of temperature and pressure (____).
At STP, 1 mole of any gas occupies __________.
The ___________________of any gas at STP is 22.4
L/mole.
Molar volume contains 6.02 x 1023 particles of that gas
Volume of a Mole of Gas
 The density of a gas is usually measured in the
_______________________________________________
_______________________________________________
Review
 Density = Mass/ Volume
 What are densities units?
 ________________________________________________
 If you had a mole of gas at STP, how could you
calculate the density?
 _______________________________________________
 What information do you need to calculate the molar
mass of a gas?
 _______________________________________________
The Mole Road Map
Check Prior Knowledge
 What is 73% of 150?
 ________________
 What percent of 6.5 is 3.1?
 ________________
 Look at figure 7.15 which compound is the better
source of potassium?
 ________________
Calculating the % Composition of a
Compound
 The relative amounts of each element in a compound are
expressed as the percent composition, or the percent by
mass of each element in a compound.
 The percent by mass of an element in a compound is the
number of grams of the element divided by the number of
grams of the compound multiplied by 100%.
% mass of element E = (grams of element E)/(grams of compound) 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?
A: Known mass of Mg, Mass of N, mass of compound
(9.03 + 3.48 = 12.52g)
Unknown: % Mg, % N
C:
% Mg = mass of Mg / grams of compound x 100% =
72.2%
% N = mass of N/grams of compound x 100% = 27.8%
Calculating % Composition
 Of a known compound
 ________________________________________________
________________________________________________
________________________________________________
 ________________________________________________
________________________________________________
________________________________________________
% mass = (grams of element in 1 mole of
compound) / molar mass of compound x 100%
Calculate the percent composition of
these compounds: Ethane (C2H6)
Ethane
% C= mass of C / grams of compound x 100%
% H= mass of H / grams of compound x 100%
Mass of C = 2 * 12= 24
Mass of H = 6 *1= 6
Mass of compound = 30
%C = 24/30 x 100% = 80%
% H = 6/30 x 100%= 20%
% Composition Practice
1) Fe2O3
2) HgO
3) Ag2O
4) Na2O
Using % as a Conversion Factor
 You can use the percent composition to calculate the
number of grams of an element contained in a specific
amount of a compound.
 _____________________________________________________
_____________________________________________________
_____________________________________________________
____________________________________________________
 For example if you know that the %C in the compound is
81.8% and you have a 82.0 g sample
82.0g of compound x 81.8grams C/ 100 g of compound= 67.1 g C
Calculating Empirical Formulas
 What is an Empirical Formula? Why do we need it?
 Suppose you are given a sample of a substance that
contains hydrogen and oxygen. You are told that the
ratio of moles of hydrogen to moles of oxygen is 1:1.
What is the formula of the substance?
Calculating Empirical Formula
 % composition can be used to calculate the empirical
formula of a compound.
 The empirical formula gives the lowest whole-number
ratio f the atoms of the elements in a compound.
Calculate the empirical formula of
94.1% O, 5.9% H
A: Known- % O, % H Unknown- empirical formula
C: In 100.0 g of the compound, there are 94.1g of O and
5.9 g H. These values are used to convert to moles.
94.1 g O x 1 mol O / 16 g O = 5.88 mol O
5.9 g H x 1 mol H/ 1 g H = 5.9 mol H
Next you divide each by the smaller mole number.
5.88/5.88= 1.00 5.9/5.88= 1.00
OH is the empirical formula
25.9% N and 74.1 % O what is the
empirical formula?
A: known % N and % O Unknown empirical formula
C: 25.9 g N x 1 mol N/ 14 g N = 1.85 mol N
74.1 g O x 1 mol O / 16 g O = 4.63 mol O
Then divide each by the smaller mole number
1.85/1.85 = 1 mol N
4.63/1.85= 2.5 mol O
However we can not have NO2.5 So we multiply both by
2 to get N2O5
Calculating Molecular Formulas
 Look at table 7.2 on page 194
 Both ethyne and benzene have the same empirical
formula of CH. But each of these have different molar
masses.
 The molecular formula is either the same as the
empirical formula or a simple whole-number multiple
of it.
Calculating Molecular Formula
 You can determine the molecular formula of a
compound if you know the
_______________________and its
______________________.
 From the empirical formula you calculate the
empirical formula mass (efm). Or the molar mass of
the empirical formula.
 Then the known molar mass is divided by the efm.
Calculate the molecular formula of the
compound whose molar mass is 60.0g and
empirical formula is CH4N.
A: Known empirical formula; molar mass
Unknown molecular formula
C:
Empirical formula
efm
CH4N
30
molar mass/efm
60/30=2
molecular formula
C2H8N2
When you known the %
composition and gram molecular
mass, you must first use the %
composition to calculate the
empirical formula. Then you can
calculate the empirical formula
mass, and compare it to the gram
molecular mass to determine the
molecular formula.