The Mole
Section 10.1 Measuring Matter
Section 10.2 Mass and the Mole
Section 10.3 Moles of Compounds
Section 10.4 Empirical and
Molecular Formulas
Section 10.5 Formulas of Hydrates
Click a hyperlink or folder tab to view
the corresponding slides.
Exit
Section 10.1 Measuring Matter
• Explain how a mole is
used to indirectly count
the number of particles of
matter.
molecule: two or more
atoms that covalently
bond together to form a
unit
• Relate the mole to a
common everyday
counting unit.
mole
• Convert between moles
and number of
representative particles.
Avogadro’s number
Chemists use the mole to count atoms,
molecules, ions, and formula units.
Counting Particles
• Chemists need a convenient method for
accurately counting the number of atoms,
molecules, or formula units of a substance.
• The mole is the SI base unit used to measure
the amount of a substance.
• 1 mole is the amount of atoms in 12 g of pure
carbon-12, or 6.02  1023 atoms.
• The number is called Avogadro’s number.
Converting Between Moles and Particles
• Conversion factors must be used.
• Moles to particles
Number of molecules in 3.50 mol of sucrose
Converting Between Moles and Particles (cont.)
• Particles to moles
• Use the inverse of Avogadro’s number as the
conversion factor.
Section 10.1 Assessment
What does the mole measure?
A. mass of a substance
B. amount of a substance
C. volume of a gas
D
C
A
0%
B
D. density of a gas
A. A
B. B
C. C
0%
0%
0%
D. D
Section 10.1 Assessment
What is the conversion factor for
determining the number of moles of a
substance from a known number of
particles?
A
D. 1 mol  6.02  1023 particles
0%
D
C. 1 particle  6.02  1023
C
B.
A. A
B. B
C. C
0%
0%
0%
D. D
B
A.
Section 10.2 Mass and the Mole
• Relate the mass of an atom conversion factor: a
to the mass of a mole of
ratio of equivalent
atoms.
values used to express
the same quantity in
• Convert between number
different units
of moles and the mass of
an element.
• Convert between number
of moles and number of
atoms of an element.
molar mass
A mole always contains the same number
of particles; however, moles of different
substances have different masses.
The Mass of a Mole
• 1 mol of copper and 1 mol of carbon have
different masses.
• One copper atom has a different mass than 1
carbon atom.
The Mass of a Mole (cont.)
• Molar mass is the mass in grams of one
mole of any pure substance.
• The molar mass of any element is
numerically equivalent to its atomic mass and
has the units g/mol.
Using Molar Mass
• Moles to mass
3.00 moles of copper has a mass of 191 g.
Using Molar Mass (cont.)
• Convert mass to moles with the inverse
molar mass conversion factor.
• Convert moles to atoms with Avogadro’s
number as the conversion factor.
Using Molar Mass (cont.)
• This figure shows the steps to complete
conversions between mass and atoms.
Section 10.2 Assessment
The mass in grams of 1 mol of any pure
substance is:
A. molar mass
B. Avogadro’s number
D
A
0%
C
D. 1 g/mol
A. A
B. B
C. C
0%
0%
0%
D. D
B
C. atomic mass
Section 10.2 Assessment
Molar mass is used to convert what?
A. mass to moles
B. moles to mass
C. atomic weight
D
C
A
0%
B
D. particles
A. A
B. B
C. C
0%
0%
0%
D. D
Section 10.3 Moles of Compounds
• Recognize the mole relationships shown by a
chemical formula.
• Calculate the molar mass of a compound.
• Convert between the number of moles and mass of
a compound.
• Apply conversion factors to determine the number of
atoms or ions in a known mass of a compound.
representative particle: an atom, molecule, formula
unit, or ion
Section 10.3 Moles of Compounds
(cont.)
The molar mass of a compound can be
calculated from its chemical formula
and can be used to convert from mass
to moles of that compound.
Chemical Formulas and the Mole
• Chemical formulas indicate the numbers
and types of atoms contained in one unit of
the compound.
• One mole of CCl2F2 contains one mole of C
atoms, two moles of Cl atoms, and two moles
of F atoms.
The Molar Mass of Compounds
• The molar mass of a compound equals the
molar mass of each element, multiplied by
the moles of that element in the chemical
formula, added together.
• The molar mass of a compound
demonstrates the law of conservation of
mass.
Converting Moles of a Compound to Mass
• For elements, the conversion factor is the
molar mass of the compound.
• The procedure is the same for compounds,
except that you must first calculate the molar
mass of the compound.
Converting the Mass of a Compound to Moles
• The conversion factor is the inverse of the
molar mass of the compound.
Converting the Mass of a Compound to
Number of Particles
• Convert mass to moles of compound with
the inverse of molar mass.
• Convert moles to particles with Avogadro’s
number.
Converting the Mass of a Compound to
Number of Particles (cont.)
• This figure summarizes the conversions
between mass, moles, and particles.
Section 10.3 Assessment
How many moles of OH— ions are in 2.50
moles of Ca(OH)2?
A. 2.00
B. 2.50
D
A
0%
C
D. 5.00
A. A
B. B
C. C
0%
0%
0%
D. D
B
C. 4.00
Section 10.3 Assessment
How many particles of Mg are in 10 moles
of MgBr2?
A. 6.02  1023
B. 6.02  1024
D
A
0%
C
D. 1.20  1025
A. A
B. B
C. C
0%
0%
0%
D. D
B
C. 1.20  1024
Section 10.4 Empirical and Molecular Formulas
• Explain what is meant by
the percent composition
of a compound.
• Determine the
empirical and molecular
formulas for a
compound from mass
percent and actual
mass data.
percent by mass: the
ratio of the mass of each
element to the total
mass of the compound
expressed as a percent
percent composition
empirical formula
molecular formula
A molecular formula of a compound is
a whole-number multiple of its
empirical formula.
Percent Composition
• The percent by mass of any element in a
compound can be found by dividing the
mass of the element by the mass of the
compound and multiplying by 100.
Percent Composition (cont.)
• The percent by mass of each element in a
compound is the percent composition of
a compound.
• Percent composition of a compound can also
be determined from its chemical formula.
Empirical Formula
• The empirical formula for a compound is
the smallest whole-number mole ratio of
the elements.
• You can calculate the empirical formula from
percent by mass by assuming you have
100.00 g of the compound. Then, convert the
mass of each element to moles.
• The empirical formula may or may not be the
same as the molecular formula.
Molecular formula of hydrogen peroxide = H2O2
Empirical formula of hydrogen peroxide = HO
Molecular Formula
• The molecular formula specifies the
actual number of atoms of each element in
one molecule or formula unit of the
substance.
• Molecular formula is always a whole-number
multiple of the empirical formula.
Molecular Formula (cont.)
Section 10.4 Assessment
What is the empirical formula for the
compound C6H12O6?
A. CHO
B. C2H3O2
D
A
0%
C
D. CH3O
A. A
B. B
C. C
0%
0%
0%
D. D
B
C. CH2O
Section 10.4 Assessment
Which is the empirical formula for
hydrogen peroxide?
A. H2O2
B. H2O
D
A
0%
C
D. none of the above
A. A
B. B
C. C
0%
0%
0%
D. D
B
C. HO
Section 10.5 Formulas of Hydrates
• Explain what a hydrate
is and relate the name of
the hydrate to its
composition.
• Determine the formula
of a hydrate from
laboratory data.
crystal lattice: a threedimensional geometric
arrangement of particles
hydrate
Hydrates are solid ionic compounds in
which water molecules are trapped.
Naming Hydrates
• A hydrate is a compound that has a
specific number of water molecules bound
to its atoms.
• The number of water molecules associated
with each formula unit of the compound is
written following a dot.
• Sodium carbonate decahydrate =
Na2CO3 • 10H2O
Naming Hydrates
(cont.)
Analyzing a Hydrate
• When heated, water molecules are
released from a hydrate leaving an
anhydrous compound.
• To determine the formula of a hydrate, find
the number of moles of water associated with
1 mole of hydrate.
Analyzing a Hydrate
(cont.)
• Weigh hydrate.
• Heat to drive off the water.
• Weigh the anhydrous compound.
• Subtract and convert the difference to moles.
• The ratio of moles of water to moles of
anhydrous compound is the coefficient for
water in the hydrate.
Use of Hydrates
• Anhydrous forms of hydrates are often
used to absorb water, particularly during
shipment of electronic and optical
equipment.
• In chemistry labs, anhydrous forms of
hydrates are used to remove moisture from
the air and keep other substances dry.
Section 10.5 Assessment
Heating a hydrate causes what to
happen?
A. Water is driven from the hydrate.
B. The hydrate melts.
D
A
0%
C
D. There is no change in the
hydrate.
A. A
B. B
C. C
0%
0%
0%
D. D
B
C. The hydrate conducts
electricity.
Section 10.5 Assessment
A hydrate that has been heated and the
water driven off is called:
A. dehydrated compound
B. antihydrated compound
D
A
0%
C
D. hydrous compound
A. A
B. B
C. C
0%
0%
0%
D. D
B
C. anhydrous compound
Chemistry Online
Study Guide
Chapter Assessment
Standardized Test Practice
Image Bank
Concepts in Motion
Section 10.1 Measuring Matter
Key Concepts
• The mole is a unit used to count particles of matter
indirectly. One mole of a pure substance contains
Avogadro’s number of particles.
• Representative particles include atoms, ions,
molecules, formula units, electrons, and other similar
particles.
• One mole of carbon-12 atoms has a mass of exactly
12 g.
• Conversion factors written from Avogadro’s relationship
can be used to convert between moles and number of
representative particles.
Section 10.2 Mass and the Mole
Key Concepts
• The mass in grams of 1 mol of any pure substance is
called its molar mass.
• The molar mass of an element is numerically equal to
its atomic mass.
• The molar mass of any substance is the mass in grams
of Avogadro’s number of representative particles of the
substance.
• Molar mass is used to convert from moles to mass. The
inverse of molar mass is used to convert from mass to
moles.
Section 10.3 Moles of Compounds
Key Concepts
• Subscripts in a chemical formula indicate how many
moles of each element are present in 1 mol of the
compound.
• The molar mass of a compound is calculated from the
molar masses of all of the elements in the compound.
• Conversion factors based on a compound’s molar
mass are used to convert between moles and mass of
a compound.
Section 10.4 Empirical and
Molecular Formulas
Key Concepts
• The percent by mass of an element in a compound
gives the percentage of the compound’s total mass
due to that element.
• The subscripts in an empirical formula give the smallest
whole-number ratio of moles of elements in the
compound.
• The molecular formula gives the actual number of
atoms of each element in a molecule or formula unit of
a substance.
• The molecular formula is a whole-number
multiple of the empirical formula.
Section 10.5 Formulas of Hydrates
Key Concepts
• The formula of a hydrate consists of the formula of
the ionic compound and the number of water
molecules associated with one formula unit.
• The name of a hydrate consists of the compound name
and the word hydrate with a prefix indicating the
number of water molecules in 1 mol of the compound.
• Anhydrous compounds are formed when hydrates are
heated.
What does Avogadro’s number represent?
A. the number of atoms in 1 mol of
an element
B. the number of molecules in 1 mol of
a compound
D
C
B
A
A. A
C. the number of Na+ ions in 1 mol of
B. B
NaCl (aq)
C. C
D. all of the above
0%
0%
0%
0%
D. D
The molar mass of an element is
numerically equivalent to what?
A. 1 amu
B. 1 mole
D
A
0%
C
D. its atomic number
A. A
B. B
C. C
0%
0%
0%
D. D
B
C. its atomic mass
How many moles of hydrogen atoms are
in one mole of H2O2?
A. 1
B. 2
D
A
0%
C
D. 0.5
A. A
B. B
C. C
0%
0%
0%
D. D
B
C. 3
What is the empirical formula of Al2Br3?
A. AlBr
B. AlBr3
D
A
0%
C
D. Al2Br3
A. A
B. B
C. C
0%
0%
0%
D. D
B
C. Al2Br
What is an ionic solid with trapped water
molecules called?
A. aqueous solution
B. anhydrous compound
D
A
0%
C
D. solute
A. A
B. B
C. C
0%
0%
0%
D. D
B
C. hydrate
Two substances have the same percent by
mass composition, but very different
properties. They must have the same ____.
A. density
A
0%
D
D. molar mass
C
C. molecular formula
A. A
B. B
C. C
0%
0%
0%
D. D
B
B. empirical formula
How many moles of Al are in 2.0 mol of
Al2Br3?
A. 2
B. 4
D
A
0%
C
D. 1
A. A
B. B
C. C
0%
0%
0%
D. D
B
C. 6
How many water molecules are
associated with 3.0 mol of CoCl2 • 6H2O?
A. 18
B. 1.1  1025
D
A
0%
C
D. 1.8  1024
A. A
B. B
C. C
0%
0%
0%
D. D
B
C. 3.6  1024
How many atoms of hydrogen are in
3.5 mol of H2S?
A. 7.0  1023
B. 2.1  1023
D
A
0%
C
D. 4.2  1024
A. A
B. B
C. C
0%
0%
0%
D. D
B
C. 6.0  1023
Which is not the correct formula for an
ionic compound?
A. CO2
B. NaCl
D
A
0%
C
D. LiBr2
A. A
B. B
C. C
0%
0%
0%
D. D
B
C. Na2SO4
Click on an image to enlarge.
Figure 10.6 Molar Mass
Table 10.1 Formulas of Hydrates
Click any of the background top tabs
to display the respective folder.
Within the Chapter Outline, clicking a section
tab on the right side of the screen will bring you
to the first slide in each respective section.
Simple navigation buttons will allow you to
progress to the next slide or the previous slide.
The Chapter Resources Menu will allow you to
access chapter specific resources from the Chapter
Menu or any Chapter Outline slide. From within any
feature, click the Resources tab to return to this slide.
The “Return” button will allow you to return to the
slide that you were viewing when you clicked either
the Resources or Help tab.
To exit the presentation, click the Exit button on the Chapter Menu slide or hit
Escape [Esc] on your keyboards while viewing any Chapter Outline slide.
This slide is intentionally blank.