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1/17/16
Chapter 3
Molecules, Compounds and Chemical Equations
Chapter 3
Molecules, Compounds and Chemical Equations
3.1
3.2
3.3
3.4
3.5
3.6
3.7
3.8
3.9
3.10
3.11
3.12
Hydrogen, Oxygen and Water
Chemical Bonds
Representing Compounds: Chemical Formulas and Molecular
Models
An Atomic Level View of Elements and Compounds
Ionic Compounds: Formulas and Names
Molecular Compounds: Formulas and Names
Summary of Inorganic Nomenclature
Formula Mass and the Mole Concept for Compounds
Composition of Compounds
Determining a Chemical Formula from Experimental Data
Writing and Balancing Chemical Equations
Organic Compounds
2
Section 3.1
Hydrogen, Oxygen and Water
Mixtures vs Compounds
•  When two or more elements combine to form a compound
an entirely new substance results.
–  Balloon contains a mixture of hydrogen (H2) and oxygen gas (O2)
–  Glass contains a compound composed of hydrogen and oxygen
(H2O)
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Section 3.1
Hydrogen, Oxygen and Water
Properties of Hydrogen, Oxygen and Water
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Section 3.1
Hydrogen, Oxygen and Water
Sodium Chloride NaCl
When two or more elements combine to form a compound
an entirely new substance results.
Sodium (Na) = Soft reactive metal
Chlorine (Cl) = toxic gas
Sodium Chloride (NaCl) =
Tasty condiment/preservative
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Section 3.1
Hydrogen, Oxygen and Water
Ionic vs Covalent Bonds
•  Compounds form through Bond formation
•  2H2 + O2 → 2H2O
covalent bond
•  2Na + Cl2 → 2NaCl
ionic bond
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Section 3.2
Chemical Bonds
Ionic Bonds
• 
Ionic bonds form when one atom loses
electrons and another atom gains electrons.
–  Metal atoms tend to lose electrons and
become positive cations
–  Nonmetal atoms tend to gain electrons and
become negative anions
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Section 3.2
Chemical Bonds
Ionic Bonds
• 
• 
The oppositely charged ions attract one
another by electrostatic forces and form an
ionic bond
The result is an ionic compound, which
(in the solid phase) is composed of a
lattice – a regular three dimensional array
of cations and anions
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Section 3.2
Chemical Bonds
Covalent Bonds
•  Covalent bonds form when two atoms share
electrons
–  Form between non-metals – neither are willing to give
up electrons so in order to form a bond – they share.
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Section 3.2
Chemical Bonds
Covalent Bonds
•  Water (H2O), ammonia (NH3), and Methane
(CH4) all have covalent bonds. A covalent
bond is a shared pair of electrons (represented
here by a line). Notice covalent compounds
are individual units (molecules)
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Section 3.2
Chemical Bonds
Why do bonds form?
•  Bonds form to lower the potential energy of the
atoms
–  In the process each atom ends up with an octet of
electrons (8 electrons in outer shell)
–  This is the result
–  Not the driving force
•  A major driving force in nature is to increase
stability (lower potential energy).
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Section 3.3
Representing Compounds: Chemical Formulas and Molecular Models
Types of Chemical Formulas
•  Empirical formula
–  Simplest – tells what types of atoms are present in a
compound
–  Empirical formula for hydrogen peroxide is HO
•  Molecular Formula
–  Give the actual number of atoms in a compound
–  Molecular formula for hydrogen peroxide is H2O2
•  Structural Formula
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Section 3.3
Representing Compounds: Chemical Formulas and Molecular Models
Types of Chemical Formulas
•  Structural Formula
–  Can be simple (L)
–  Can indicate geometry of the compound (R)
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Section 3.3
Representing Compounds: Chemical Formulas and Molecular Models
Molecular Models
•  Ball and Stick Models
–  Atoms are balls and bonds are sticks
•  Space filling Models
–  Atoms fill the spaces between each
other
–  Probably the most “accurate” way to
imagine a molecule if it was blown up
to a size we could see
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Section 3.4
An Atomic-Level View of Elements and Compounds
Elements vs Compounds
• 
Pure substances can be elements or
compounds (Ch 1)
H2, Ca, Ag
H2O, CaCl2
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Section 3.4
An Atomic-Level View of Elements and Compounds
Elements vs Compounds
• 
Elements and Compounds can be further categorized
based on the types of units that compose them
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Section 3.4
An Atomic-Level View of Elements and Compounds
Elements vs Compounds
• 
Elements can be atomic (exist as individual
atoms) or molecular (oxygen is a diatomic
molecule). There are even polyatomic
elements (next slide).
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Section 3.4
An Atomic-Level View of Elements and Compounds
Molecular Elements
•  H2, N2, O2,
F2, Cl2, Br2
and I2
•  P4
•  S8, Se8
•  You should
learn the 7
diatomic
elements
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Section 3.4
An Atomic-Level View of Elements and Compounds
Elements vs Compounds
• 
• 
• 
Compounds can be
molecular (H2O) or
ionic (NaCl)
The smallest unit of a
molecular compound
is the molecule
The smallest unit of
an ionic compound is
a formula unit
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Section 3.4
An Atomic-Level View of Elements and Compounds
Molecules vs Formula Units
•  A water molecule is a discrete (individual) stable
unit that consists of three atoms tightly bound
together.
–  Covalent bonds are very strong
•  A sodium chloride formula unit is the smallest
unit of a sample of NaCl.
–  Does not exist on its own as a stable discrete
unit.
–  Single ionic bond is very weak
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Section 3.5 and 3.6
Ionic and Molecular Compounds
Types of Compounds
• 
Binary Ionic Compounds
– 
Metal—nonmetal
NaCl, CaBr2
•  Ionic Compounds with Polyatomic Ions
−  Ions with more than one atom
•  Hydrated Ionic Compounds
• 
NaNO3, NH3Cl
MgSO47 H2O
Binary Covalent Compounds
– 
Nonmetal—nonmetal
CO2, H2O, NH3
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Section 3.5
Ionic Compounds: Formulas and Names
Writing Formulas for Ionic Compounds
•  Binary Ionic Compounds form between a metal
and a nonmetal
–  Metal forms a cation
–  Non metal forms an anion
•  All compounds are charge neutral
•  For these two reasons it is fairly simple to reason
out the formula of most ionic compounds just by
knowing the identity of the component elements.
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Section 3.5
Ionic Compounds: Formulas and Names
Writing Formulas for Ionic Compounds
•  If you have sodium (Na) and chlorine (Cl)
–  Na forms a Na+ cation
–  Cl forms a Cl– anion
–  They combine 1:1 to form NaCl
•  If you have calcium (Ca) and chlorine (Cl)
–  Ca forms a Ca2+ cation
–  Cl forms a Cl– anion
–  They combine 1:2 to form CaCl2
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Section 3.5
Ionic Compounds: Formulas and Names
Learning Check
•  Write formulas for ionic compounds formed from
the following elements
1. Sodium and Oxygen
2.  Calcium and bromine
3.  Aluminum and Sulfur
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Section 3.5
Ionic Compounds: Formulas and Names
Naming Ionic Compounds
•  Ionic compounds have common names and
systematic names
•  NaCl is sodium chloride or table salt (common
name)
•  NaHCO3 is sodium bicarbonate or baking soda
(common name)
•  We will learn the systematic naming system
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Section 3.5
Ionic Compounds: Formulas and Names
Naming Ionic Compounds
•  Binary Ionic compounds (metals and non metals)
•  Metals are of two types
–  Metals that only make one type of cation
•  Groups I -3
–  Metals that can from more than one type of cation
•  Transition Group Metals and Main Groups 4 and above
(basically everything else)
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Section 3.5
Ionic Compounds: Formulas and Names
Naming Binary Ionic Compounds Containing
a Metal that Forms Only One Type of Cation
1. The cation is always named first and the anion
second.
2. Cation takes its name from the name of the
parent element.
3. Anion is named by taking the root of the
element name and adding –ide.
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Section 3.5
Ionic Compounds: Formulas and Names
Naming Binary Ionic Compounds Containing
a Metal that Forms Only One Type of Cation
KCl
Potassium chloride
MgBr2
Magnesium bromide
CaO
Calcium oxide
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Section 3.5
Ionic Compounds: Formulas and Names
Concept Check
Write the names of the following compounds.
A. MgO
B. Al2S3
C. MgF2
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Section 3.5
Ionic Compounds: Formulas and Names
Naming Binary Ionic Compounds Containing
a Metal that Form More than One Type of
Cation
• 
• 
• 
• 
Metals in these compounds form more than
one type of positive ion
Charge on the metal ion must be specified.
Roman numeral indicates the charge of the
metal cation.
Transition metals and large metals (period IV
and higher usually require a Roman numeral
• 
Exceptions: Ag+, Cd2+, Zn2+ (form only a single ion)
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Section 3.5
Ionic Compounds: Formulas and Names
Naming Binary Ionic Compounds Containing
a Metal that Form More than One Type of
Cation
1. The cation is always named first and the anion
second.
2.  A monatomic cation takes its name from the
name of the parent element.
3.  Charge of cation in roman numerals (in
parenthesis)
3. A monatomic anion is named by taking the root
of the element name and adding –ide.
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Section 3.5
Ionic Compounds: Formulas and Names
Naming Binary Ionic Compounds Containing
a Metal that Form More than One Type of
Cation
CuBr
Copper (I) bromide
FeS
Iron (II) sulfide
PbO2
Lead (IV) oxide
Figure out the charge on the cation based on the
charge of the anion.
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Section 3.5
Ionic Compounds: Formulas and Names
Learning Check
Name the following ionic compounds containing
metals that form two kinds of positive ions:
A.  Fe2O3
B.  SnCl2
C.  PbI4
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Section 3.5
Ionic Compounds: Formulas and Names
Naming Ionic Compounds Containing
Polyatomic Ions
•  Naming these compounds is the same only now
we use the name of the polyatomic ion wherever
it occurs
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Section 3.5
Ionic Compounds: Formulas and Names
Naming Ionic Compounds Containing
Polyatomic Ions
• 
• 
Must be memorized (polyatomic ion handout).
Examples of compounds containing polyatomic
ions:
NaOH
Sodium hydroxide
Pb(NO3)2
lead (II) nitrate
(NH4)2SO4
Ammonium sulfate
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Section 3.5
Ionic Compounds: Formulas and Names
Learning Check
Write the names of the following compounds.
A. NaC2H3O2
B. Ca3(PO4)2
C.  CuNO2
D. K2Cr2O7
E. Mg(ClO2)2
F.  NH4CN
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Section 3.5
Ionic Compounds: Formulas and Names
Hydrated Ionic Compounds
•  Some ionic compounds are hydrates
•  This means they have a specific number of
water molecules associated with each formula
unit in the lattice structure
•  An example is magnesium sulfate heptahydrate
•  MgSO47H2O
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Section 3.5
Ionic Compounds: Formulas and Names
Hydrated Ionic Compounds
•  Cobalt chloride hexahydrate before and after
heating to drive off the water
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Section 3.5
Ionic Compounds: Formulas and Names
Common hydrate prefixes
• 
• 
• 
• 
• 
• 
• 
• 
• 
Hemi = ½
Mono = 1
Di = 2
Tri = 3
Tetra = 4
Penta = 5
Hexa = 6
Hepta = 7
Octa = 8
These prefixes are also used in
naming molecular compounds
For example: carbon dioxide
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Section 3.6
Molecular Compounds: Formulas and Names
Molecular compounds vs Ionic Compounds
•  Ionic compounds form between metal and non metal
–  The ions in ionic compounds can only combine one way
because compounds are charge neutral
–  Na+ and Cl– can only combine one way
•  Molecular compounds form between two nonmetals
–  The atoms in molecular compounds don’t form ions – they
share electrons. For this reason the same combination of
elements can form a number of different molecular
compounds.
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Section 3.6
Molecular Compounds: Formulas and Names
Naming Molecular Compounds
1. The first element in the formula is named first,
using the full element name.
2. The second element is named as if it were an
anion.
3. Prefixes are used to denote the numbers of
atoms present.
4. The prefix mono- is never used for naming the
first element.
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Section 3.6
Molecular Compounds: Formulas and Names
Naming Molecular Compounds
Prefixes Used to
Indicate Number in
Chemical Names
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Section 3.6
Molecular Compounds: Formulas and Names
Naming Molecular Compounds
CO2
Carbon dioxide
SF6
Sulfur hexafluoride
N2O4
Dinitrogen tetroxide
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Section 3.6
Molecular Compounds: Formulas and Names
Learning Check
Write the name of each covalent compound.
A. SeO
B. NO2
C. PF3
D. CBr4
E. P2O5
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Section 3.6
Molecular Compounds: Formulas and Names
Conceptual Connection
•  The compound NCl3 is nitrogen trichloride , but
AlCl3 is just aluminum chloride. Why?
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Section 3.6
Molecular Compounds: Formulas and Names
Acids
•  Acids can be defined as molecular compounds
than release hydrogen ions (H+) when dissolved
in water.
•  Acids are composed of hydrogen, usually written
first in the formula, followed by one or more non
metals.
•  Two categories of acids for naming
–  Binary acids
–  Oxy Acids
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Section 3.6
Molecular Compounds: Formulas and Names
Naming Binary Acids
• 
• 
Hydrogen plus a non metal, the acid is named
with the prefix hydro– followed by the root of the
anion and the suffix –ic.
Examples:
HCl
Hydro chlor ic acid
HBr
Hydro brom ic acid
H 2S
Hydro sulfur ic acid
Notice sulfur is kind of a weirdo sulfur instead of sulf
47
Section 3.6
Molecular Compounds: Formulas and Names
Naming OxyAcids
• 
Hydrogen plus an oxyanion (anion containing a
nonmetal and oxygen)
 
 
The suffix –ic is added to the root name if anion name ends in –ate.:
HNO3
Nitrate
Nitric acid
H2SO4
Sulfate
Sulfuric acid
HC2H3O2
Acetate
Acetic acid
The suffix –ous is added to the root name if anion name ends in –ite.
HNO2
Nitrite
Nitrous acid
H2SO3
Sulfite
Sulfurous acid
HClO2
Chlorite
Chlorous acid
Notice sulfate is kind of a weirdo sulfur instead of sulf
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Section 3.6
Molecular Compounds: Formulas and Names
Flow Chart for Naming Acids
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Section 3.6
Molecular Compounds: Formulas and Names
Learning Check
Name the following acids
a)  HF (aq)
b)  H2CO3 (aq)
c)  HI (aq)
d)  HClO2 (aq)
e)  H2CrO4 (aq)
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Section 3.6
Molecular Compounds: Formulas and Names
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Section 3.6
Molecular Compounds: Formulas and Names
Learning Check
Name the following compounds
a)  KNO3
b)  TiO2
c)  Sn(OH)4
d)  PBr5
e)  H2SO3 (aq)
52
Section 3.8
Formula Mass and the Mole Concept for Compounds
Formula Mass
•  In Chapter 2 we defined the average mass of an
atom as its atomic mass
–  Measured in atomic mass units (amu)
•  The atomic mass of carbon is 12.01 amu
•  Now that we have been introduced to ionic and
molecular compounds we can build on this
definition to something called a Formula Mass
53
Section 3.8
Formula Mass and the Mole Concept for Compounds
Formula Mass
•  Formula mass is the average mass of a
molecule (or formula unit)
–  Molecular mass and molecular weight are also
common (and interchangeable terms)
•  What does it mean though?
•  Formula mass is the sum of all the atomic
masses in the chemical formula
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Section 3.8
Formula Mass and the Mole Concept for Compounds
Formula Mass
Formula Mass of H2O
(2 × 1.008 amu) + 16.00 amu = 18.02 amu
Formula Mass of Ba(NO3)2
137.33 g + (2 × 14.01 g) + (6 × 16.00 g) = 261.35
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Section 3.8
Formula Mass and the Mole Concept for Compounds
Molar Mass of a Compound
•  Also in chapter 2 we saw that the atomic mass in
amu/atom was equal to the molar mass in g/mole.
–  This is because the definition for amu and mole are
related to each other.
•  The same is true for a compound.
•  The formula mass in amu/formula unit = molar
mass in g/mole
56
Section 3.8
Formula Mass and the Mole Concept for Compounds
Molar Mass of a Compound
•  So if we know a formula mass for water = 18.02
amu/molecule
–  18.02 amu in one molecule
•  We also know that the molar mass for water =
18.02 g/mole.
–  18.02 grams in one mole
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Section 3.8
Formula Mass and the Mole Concept for Compounds
Using Molar Mass to Count by Weighing
•  The fact that amu/molecule is related to g/mole
allows us to count molecules by weighing
them.
•  This is the same thing we did when we counted
atoms by weighing them.
•  Molar mass – g/mole – is a conversion factor
between mass and numbers of molecules.
58
Section 3.8
Formula Mass and the Mole Concept for Compounds
Example
•  Glucose is one of the end products of photosynthesis,
the process that converts CO2 and H2O to complex
carbohydrates. The formula for glucose is C6H12O6.
Determine the molar mass of glucose.
Determine the number of moles in 50.0 g of glucose.
Determine the number of molecules in 50.0 g of glucose.
59
Section 3.8
Formula Mass and the Mole Concept for Compounds
Example
•  Glucose is one of the end products of photosynthesis,
the process that converts CO2 and H2O to complex
carbohydrates. The formula for glucose is C6H12O6.
Determine the molar mass of glucose.
(6 x 12.01) + (12 x 1.008) + (6 x 16.00) = 180.16 g/mol
Determine the number of moles in 50.0 g of glucose.
Determine the number of molecules in 50.0 g of glucose.
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Section 3.8
Formula Mass and the Mole Concept for Compounds
Learning Check
Ethanol is produced from sugars by yeast in the process
called fermentation. The formula for ethanol is C2H5OH.
Determine the molar mass of ethanol.
Determine the number of moles in 525 g of ethanol.
Determine the number of molecules in 525 g of ethanol.
61
Section 3.9
Composition of Compounds
Mass Percent
•  We use something called Mass Percent to
determine the relative amounts of elements in a
compound.
•  Mass percent takes the different masses of the
different elements into consideration in
calculating relative amounts.
62
Section 3.9
Composition of Compounds
Mass Percent
•  Mass Percent of hydrogen in H2O
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Section 3.9
Composition of Compounds
Mass Percent (Mass Percent Composition)
•  General Equation
64
Section 3.9
Composition of Compounds
How is Mass Percent Composition Useful
•  A lot of times we need to know how much of one
specific element is present in a compound
•  A toxin
•  An active ingredient
•  Mass Percent (mass percent composition)
allows us to calculate this.
65
Section 3.9
Composition of Compounds
Learning Check
•  Determine the mass percent of each element in
calcium carbonate (CaCO3)
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Section 3.9
Composition of Compounds
Mass Percent Composition as a Conversion
Factor
•  Mass percent can also be used as a conversion
factor
•  From the last problem
•  Mass Percent of Oxygen = 47.96%
67
Section 3.9
Composition of Compounds
Mass Percent Composition as a Conversion
Factor
•  How do we use mass percent as a conversion
factor?
•  To calculate the mass of an element in a given
mass of a compound.
68
Section 3.9
Composition of Compounds
Example
•  Sodium chloride (table salt) is 39% sodium by
mass. How much sodium chloride is allowed per
the RDA (recommended daily allowance) if a
person is allowed to consume 2.4 g of sodium
per day.
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Section 3.9
Composition of Compounds
Example
•  Sodium chloride (table salt) is 39% sodium by
mass. How much sodium chloride is allowed per
the RDA (recommended daily allowance) if a
person is allowed to consume 2.4 g of sodium
per day.
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Section 3.9
Composition of Compounds
Learning Check
•  Calcium carbonate (a common calcium
supplement) is 40.04% calcium by mass. How
much calcium is present in one tablet of calcium
carbonate that has a mass of 500. mg?
71
Section 3.9
Composition of Compounds
Conversion Factors from Chemical Formulas
•  We just determined the amount of calcium in a
given mass of calcium carbonate using mass
percent as a conversion factor.
•  There is actually another way to do this
calculation
•  Using the Mole Relationships in the Chemical
Formula
•  What does this mean and how do we use it?
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Section 3.9
Composition of Compounds
Mole Relationships
•  Chemical formulas tell us the number of atoms in
a compound
•  For a single formula unit of calcium carbonate
(CaCO3)
•  But what if we have a mole of CaCO3?
•  Mole Relationships for CaCO3
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Section 3.9
Composition of Compounds
Mole Relationships
•  Mole Relationships for CaCO3
•  There is a 1:1 relationship between moles of
calcium or carbon atoms and moles of
compound
•  There is a 3:1 relationship between moles of
oxygen atoms and moles of compound
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Section 3.9
Composition of Compounds
Conversion Factors from Chemical Formulas
•  How do we use mole relationships to determine
composition of a compound?
•  The mole relationship is a conversion factor.
•  So lets say we want to know how many grams of
calcium are in a 500. mg sample of CaCO3
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Section 3.9
Composition of Compounds
Conversion Factors from Chemical Formulas
•  How many grams of calcium are in a 500. mg
sample of CaCO3
•  We cannot convert directly from grams to grams.
•  The relationship between atoms in a compound
is moles to moles
•  How do we use the mole:mole conversion to do
this.
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Section 3.9
Composition of Compounds
Conversion Factors from Chemical Formulas
•  How many grams of calcium are in a 500. mg
sample of CaCO3
•  Grams of CaCO3 to moles CaCO3
•  Moles CaCO3 to moles of Ca
•  Moles of Ca to grams of Ca
•  GMMG
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Section 3.9
Composition of Compounds
Learning Check
•  The book uses the compound CCl2F2
(chlorofluorocarbons) for several examples
•  Write out the mole relationships for this
compound
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Section 3.9
Composition of Compounds
Learning Check
•  Use the mole relationships from the last problem
to determine how many grams of fluorine (F) are
in 1500 g of the compound CCl2F2.
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Section 3.10
Determining a Chemical Formula from Experimental Data
Decomposition Analysis
•  How do we determine the formula of a
compound that is newly isolated?
•  This is commonly done in the lab when new
chemicals/drugs are synthesized
•  Also common when new compounds are
isolated
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Section 3.10
Determining a Chemical Formula from Experimental Data
Decomposition Analysis
•  First step is to decompose the compound into its
constituent parts
•  Weigh them
•  But this gives us a ratio of masses (in grams)
•  Does not really tell us anything about the
relationship of the number of atoms in the
compound
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Section 3.10
Determining a Chemical Formula from Experimental Data
Decomposition Analysis
• 
• 
• 
• 
If we decompose 7.717 g of water (H2O) we get
6.868 g of oxygen
0.857 g of hydrogen
If we divide these it gives an 8:1 ratio of oxygen
to water (well we know that’s not right)
•  To get the ratio of atoms in the compound we
have to correct for the different in their
atomic weight
•  We have to convert to moles
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Section 3.10
Determining a Chemical Formula from Experimental Data
Decomposition Analysis
•  We have to convert to moles
•  Now this looks more like it
•  H0.850 O0.429
•  Divide by the smaller number = H1.98O = H2O
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Section 3.10
Determining a Chemical Formula from Experimental Data
Empirical vs Molecular Formula
•  A decomposition analysis breaks a complex
compound down to its elements to determine the
relationship between them
•  This relationship is called an empirical formula
•  The simplest whole number ratio of atoms in the
compound
•  The form of the compound before it was broken
down might actually be a multiple of this
empirical formula
•  This is called the molecular formula
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Section 3.10
Determining a Chemical Formula from Experimental Data
Empirical vs Molecular Formula
•  To determine the molecular formula you need
the empirical formula and the molecular weight
of the original compound.
•  For water the empirical formula and molecular
formula are the same
•  Not true for other compounds
•  Molecular formula = empirical formula x n
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Section 3.10
Determining a Chemical Formula from Experimental Data
Example
•  The empirical formula for a compound by
decomposition analysis is CH2O and the molar
mass is determined to be 180.2. What is the
molecular formula?
•  Molecular formula = (CH2O) x n
•  Molecular formula = C6H12O6
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Section 3.10
Determining a Chemical Formula from Experimental Data
Learning Check
•  A compound with the following percent
composition has a molar mass of 60.10 g/mol.
Determine its molecular formula. Assume 100 g
of the compound.
•  C, 39.97%
•  H, 13.41%
•  N, 46.62%
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Section 3.10
Determining a Chemical Formula from Experimental Data
Combustion Analysis
•  One way to determine hydrogen and carbon
content of a compound is to react it with oxygen
then collect the H2O and CO2 produced
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Section 3.10
Determining a Chemical Formula from Experimental Data
Combustion Analysis
• 
• 
• 
How can this information be used to determine a formula?
If a substance has been isolated that contains only C, H and N.
If 0.1156 g of this sample is reacted with oxygen and 0.1638 g of CO2
and 0.1676 g of H2O are collected.
• 
• 
The nitrogen is what is left over
0.1156 – (0.04470 + 0.01875) = 0.05215 g N
• 
To determine formulas though we need numbers of atoms and grams
don’t really help us with that
89
Section 3.10
Determining a Chemical Formula from Experimental Data
Combustion Analysis
• 
• 
Ratio of moles is the same as ratio of atoms (Section 3.3 The
Mole)
Divide by the smallest value
90
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Section 3.10
Determining a Chemical Formula from Experimental Data
Combustion Analysis
• 
So we need to convert grams to moles
• 
Empirical Formula = C1H5N1
91
Section 3.11
Writing and Balancing Chemical Equations
How to Write Balanced Chemical Equations
•  Chemical equations include the states of the
reactants and products
•  CH4 (g) + O2 (g) → CO2 (g) + H2O (g)
•  All of the reactants and products in this
reaction are gases
92
Section 3.11
Writing and Balancing Chemical Equations
States of reactants and products
•  Notice the
difference
between (l)
liquid and (aq)
aqueous
•  Aqueous is
something
dissolved in
water.
93
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Section 3.11
Writing and Balancing Chemical Equations
How to Write Balanced Chemical Equations
•  Balanced equation
•  CH4 (g) + 2O2 (g) → CO2 (g) + 2H2O (g)
94
Section 3.11
Writing and Balancing Chemical Equations
Tips for balancing
•  Balance polyatomic ions as a group
•  Always balance oxygen last
•  Can used fractions to balance – then multiply
them out at the end
•  Practice, practice, and practice.
95
Section 3.11
Writing and Balancing Chemical Equations
Example
Write a balanced equation for the reaction of
aqueous lead (II) nitrate with aqueous sodium
phosphate to produce solid lead (II) phosphate and
aqueous sodium nitrate.
Pb(NO3)2 (aq) + Na3PO4 (aq) → Pb3(PO4)2 (s) + NaNO3 (aq)
3 Pb(NO3)2 (aq) + 2 Na3PO4 (aq) → Pb3(PO4)2 (s) + 6 NaNO3
(aq)
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Section 3.11
Writing and Balancing Chemical Equations
Learning Check
Write a balanced equation for the reaction of
ammonia (NH3) with the oxygen to produce
nitrogen monoxide and water
Section 3.11
Writing and Balancing Chemical Equations
Learning Check
Write a balanced equation for the reaction of
aqueous cobalt (I) nitrate with sulfuric acid to produce
solid cobalt (I) sulfate and nitric acid. Both acids are
aqueous solutions
Section 3.11
Writing and Balancing Chemical Equations
Conceptual Connection
•  Which quantity or quantities must always be the
same on both sides of a chemical equation?
•  (a) the number of atoms of each kind
•  (b) the number of molecules of each kind
•  (c) the number of moles of each kind of
molecule
•  (d) the sum of the masses of all the substances
involved
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Section 3.12
Organic Compounds
Organic vs Inorganic Compounds
•  Early chemists divided compounds into two
types
•  Organic - Originated from living things
–  sugars
•  Inorganic - Originated from the earth
–  salts
100
Section 3.12
Organic Compounds
Organic vs Inorganic Compounds
•  Organic and inorganic compounds have different
chemical properties
•  Organic –  Easy to decompose
–  Very hard to synthesize – very complex
•  Inorganic –
–  Hard to decompose
–  Easier to synthesize
101
Section 3.12
Organic Compounds
What are organic compounds
•  Major components of living organisms
–  Proteins, fats, DNA
• 
• 
• 
• 
Smells, tastes, fuels
Drugs
Food
Pretty much everything we consume
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Section 3.12
Organic Compounds
What are organic compounds made of?
•  Organic compounds are composed primarily of
carbon and hydrogen with a few other elements
–  Nitrogen, oxygen and sulfur
•  Key element in an organic compound is carbon
103
Section 3.12
Organic Compounds
What are organic compounds made of?
•  Carbon makes 4 bonds and often bonds to itself
to make chain, branched and ring structures.
104
Section 3.12
Organic Compounds
Hydrocarbons
•  Organic compounds that contain only carbon
and hydrogen.
•  These are all the common fuels that we use
•  Gasoline, propane, natural gas
•  Can have single, double or triple bonds
–  Alkane – single bonds
–  Alkenes – double bonds
–  Alkynes – triple bonds
105
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Section 3.12
Organic Compounds
Naming Hydrocarbons
•  Base name determined by a prefix that indicates
the number of carbons
–  Meth = 1, eth = 2, prop = 3, but = 4, pent = 5
–  hex = 6, hept = 7, oct = 8, non = 9, dec = 10
•  Suffix indicates the presence of a multiple bond
–  Single bonds - ane
–  Double bond - ene
–  Triple bond - yne
106
Section 3.12
Organic Compounds
Naming Hydrocarbons
•  Butane CH3CH2CH2CH3
•  Butene CH3CH2CH=CH2
107
Section 3.12
Organic Compounds
Common Hydrocarbons
108
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Section 3.12
Organic Compounds
Learning Check
•  Write the name for the following hydrocarbons
•  CH3CH2CH2CH2CH2CH3
109
Section 3.12
Organic Compounds
Functionalized Hydrocarbons
•  A functionalized hydrocarbon is a carbon
hydrogen chain that has a functional group
added to it.
•  What is a functional group?
•  A functional group is a group of atoms that
change the characteristics of the hydrocarbon.
110
Section 3.12
Organic Compounds
Functionalized Hydrocarbons
•  Example of a functionalized hydrocarbon.
•  Ethane
•  Ethanol
111
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Section 3.12
Organic Compounds
Functional Groups
112
Section 3.12
Organic Compounds
Functional Groups
113
Section 3.12
Organic Compounds
Learning Check
•  For the following compounds, give the family of
the functional group (try to name the compound!)
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Section 3.12
Organic Compounds
Suggested Homework
•  Chapter 3
•  Review Questions 1 – 5, 8, 14 – 22
•  Odd numbered problems 23 – 118 (skip 25, 63,
69, 73, 79, 87, 89, 97)
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