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4.1 Chemical Bonds
A chemical bond is an attractive force
that holds two atoms together in a
more complex unit.
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4–1
4.1 Chemical Bonds
Ionic Bonds are formed through the
transfer of one or more electrons
from on atom or group of atoms to
another.
2 Na + Cl2
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
2 Na 1+ + 2 Cl1
4–2
4.1 Chemical Bonds
Ionic Bonds

E Na __ __
Na ___ ___
N
E

R
G
Y
Cl __ __
Cl __ __

2 Na + Cl2

2 Na1+ + 2 Cl1–
Electrons transferred from higher energy orbitals to lower energy orbitals
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4–3
4.1 Chemical Bonds
Covalent Bonds are formed through
the sharing of one or more pairs
of electrons between two atoms
H + F
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 HF
4–4
4.1 Chemical Bonds
Covalent Bonds

E
N
E
R
G
Y

____
H
H
__
___

__
F
___
F
__
Atomic orbitals combine to form molecular orbitals.
Electrons are shared in the lowest energy molecular orbitals.
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4–5
4.2 Valence Electrons and
Lewis Structures
Valence electrons are the electrons in the
outermost electron shell.
Valence electrons are transferred or shared
when chemical bonds form.
Lewis structures of representative elements
consist of the element’s symbol and one
dot for each valence electron.
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4–6
Figure 4.1 Lewis structures for selected
representative and noble-gas elements.
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4–7
4.3 The Octet Rule
In compound formation, an atom of an
element loses, gains, or shares electrons in such a way that its electron
configuration becomes identical to that
of the noble gas nearest to it on the
periodic table.
Atoms other than hydrogen obtain eight
electrons (an octet) in their outermost
filled shells.
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4–8
4.4 The Ionic Bond Model
An ion is an atom or group of atoms that
is electrically charged as a result of the
loss or gain of electrons.
Cations positive (+) charge
lost electrons
Anion
negative (–) charge
gained electrons
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4–9
4.4 The Ionic Bond Model
Cations:
Na1+
Anions:
P3–
Mg2+
S2–
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Al3+
Cl1–
NH41+
CO32–
4–10
Figure 4.3
A cation is smaller than the corresponding atom.
An anion is larger than the corresponding atom.
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4–11
4.4 The Ionic Bond Model
Ionic bonds, which are found in ionic
compounds, are electrostatic forces
between particles of unlike charge.
The higher the charge, and the closer the
ions approach, the stronger the bond.
Ionic compounds are generally formed
between metals and nonmetals.
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4–12
4.5 Ions and the Octet Rule
An atom tends to gain or lose electrons
until it obtains the same electron configuration as the nearest noble gas.
Its outermost shell of electrons contains
eight electrons (an octet).
Na  Na1+ + e1
(like Ne)
Cl + e1  Cl1
(like Ar)
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4–13
4.6 Ionic Compound Formation
Ionic Bonds

E Na __ __
Na ___ ___
N
E

R
G
Y
Cl __ __
Cl __ __

2 Na + Cl2

2 Na1+ + 2 Cl1–
Electrons transferred from higher energy orbitals to lower energy orbitals
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4–14
4.6 Ionic Compound Formation
Draw Lewis structure to show formation of
the following compounds:
NaCl
Na2O
CaCl2
AlCl3
Al2O3
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4–15
4.7 Formulas for Ionic
Compounds
The ratio in which positive and negative ions
combine is the ratio that achieves charge
neutrality for the resulting compound.
This can be used in lieu of Lewis Structures
to determine formulas of ionic compounds.
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4–16
4.7 Formulas for Ionic
Compounds
The charge on ions of representative elements
can be determined from their positions on
the Periodic Table.
Cations: Charge = Group #
Charge on Rb, Sr, Ga
Anions: Charge = Group # – 8
Charge on N, O, F
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4–17
4.7 Formulas for Ionic
Compounds
Determine formulas for compounds containing:
Rb and F
Sr and F
Ga and F
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Rb and O
Sr and O
Ga and O
Rb and N
Sr and N
Ga and N
4–18
4.8 The Structure of
Ionic Compounds
Ionic bonds are nondirectional; ions
arrange themselves to maximize
interactions between ions of opposite
charge.
Ionic compounds do not form molecules.
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4–19
Figure 4.5
Two-dimensional cross section of an ionic
solid (NaCl). No molecule can be
distinguished. Instead, a basic formula
unit is repeated indefinitely.
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4–20
Chemistry at a Glance:
Ionic Bonds and Ionic Compounds
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4–21
4.9 Recognizing and Naming
Binary Ionic Compounds
A binary compound contains two elements.
A binary ionic compound contains a metallic
element and a nonmetallic element.
The metallic element gives the cation.
The nonmetallic element gives the anion.
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4–22
4.9 Recognizing and Naming
Binary Ionic Compounds
The full name of the metallic element is given
first, followed by a separate word that contains the stem of the nonmetallic element
name and the suffix -ide.
NaCl
Aluminum sulfide
MgO
Zinc iodide
AgBr
Calcium fluoride
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4–23
4.9 Recognizing and Naming
Binary Ionic Compounds
Most transition metals can form more than
one ion. Know these terms:
Fe2+ ferrous
Fe3+ ferric
Cu1+ cuprous
Cu2+ cupric
Hg22+ mercurous Hg2+ mercuric
Sn2+ stannous
Sn4+ stannic
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4–24
4.9 Recognizing and Naming
Binary Ionic Compounds
IUPAC (International Union of Pure and Applied Chemistry) rules for naming compounds in which the cation can have more
than one charge:
Name the metal, give its charge in Roman
numerals, in parentheses, then name the
anion. Give charge only when necessary.
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4–25
4.9 Recognizing and Naming
Binary Ionic Compounds
Give names:
Cu2O
CuO
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FeCl2
FeCl3
4–26
Figure 4.6
(a) Copper (II) oxide (CuO) is black;
copper (I) oxide (Cu2O) is brown.
(b) Iron (II) chloride (FeCl2) is green;
iron (III) chloride (FeCl3) is yellow.
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4–27
Figure 4.7 A periodic table in which the metallic
elements that exhibit a fixed ionic charge
are highlighted.
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4–28
4.10 The Covalent Bond Model
• Covalent bonds, which are found in
molecular compounds, are directional
bonds that result from sharing electrons
• Covalent bonds generally form between
two nonmetallic atoms
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4–29
Figure 4.8
Electron sharing can occur only when
electron orbitals from two different atoms
overlap, forming molecular orbitals.
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4–30
4.10 Covalent Bond Model
Covalent Bonds

E
N
E
R
G
Y

____
H
H
__
___

__
F
___
F
__
Atomic orbitals combine to form molecular orbitals.
Electrons are shared in the lowest energy molecular orbitals.
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4–31
4.11 Lewis Structures for
Molecular Compounds
Draw Lewis Structures for the following
molecules:
H2
HF
F2
H2O
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4–32
4.12 Single, Double, and Triple
Covalent Bonds
Draw Lewis Structures for the following
molecules:
N2
O2
CO2
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4–33
4.13 A Systematic Method for
Drawing Lewis Structures
Calculate N  A = S
Identify Central Atom
Draw Structure
Check Structure
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4–34
4.13 A Systematic Method for
Drawing Lewis Structures
Calculate N  A = S
N = Needs = # of electrons to give every
atom in formula an octet
(H needs 2)
A = Available = # of valence electrons
for all atoms in formula
S = Shared = N  A = # of electrons shared
to give each atom an octet
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4–35
4.13 A Systematic Method for
Drawing Lewis Structures
Identify Central Atom
Has fewest electrons available
Is NEVER hydrogen
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4–36
4.13 A Systematic Method for
Drawing Lewis Structures
Draw Structure
Central atom in center
Others around it, on all four sides
Pairs of electrons between atoms
Share more than one pair if
S > # of outside atoms
Extra electrons around atoms
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4–37
4.13 A Systematic Method for
Drawing Lewis Structures
Check Structure
Central atom in center?
# of Shared electrons = S?
# of total electrons = A?
All atoms surrounded by 8 electrons?
(except H, shares 2 electrons)
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4–38
Practice Structures
Draw:
PF3
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HCN
SO2
4–39
Resonance in Lewis Structures
In some structures, like SO2, it is possible to
put a second pair of electrons in several
equivalent places.
The resulting structures are called resonance
structures. In molecules with resonance,
the extra pair of electrons is shared among
the atoms that could accommodate it.
How many resonance structures are there for
SO3?
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4–40
Figure 4.10
The phosphorus trifluoride
(PF3) molecule.
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4–41
Figure 4.11
The hydrogen cyanide (HCN) molecule.
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4–42
Figure 4.9
The sulfur dioxide (SO2) molecule.
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4–43
4.14 The Shape of Molecules:
Molecular Chemistry
The shape of a molecule has a major impact
on its properties.
The shapes of molecules can be predicted.
Use:
Lewis Structure
VSEPR (Valence Shell Electron
Pair Repulsion Theory)
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4–44
4.14 The Shape of Molecules:
Molecular Chemistry
Valence Shell Electron Pair Repulsion Theory
Procedures for predicting 3-D shapes of
molecules, using Lewis Structures
Electron groups minimize their interactions
by getting as far apart as possible in a
molecule.
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4–45
4.14 The Shape of Molecules:
Molecular Chemistry
VSEPR Electron "Pairs" (electron groups):
Nonbonding electron pairs
Any set of bonding electron groups
1 electron pair
2 electron pairs
3 electron pairs
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(single bond)
(double bond)
(triple bond)
4–46
Figure 4.12
Arrangements of valence electron pairs
about a central atom that minimize
repulsions between the pairs.
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4–47
VSEPR Pairs;
Electronic and Molecular Geometry
# of VSEPR
Groups
Electronic
Geometry
Bonding and Nonbonding Pairs
2
Linear
2 Bonding
Linear
3
Trigonal Planar
3 Bonding
Trigonal Planar
2 Bonding
1 Nonbonding
4
Tetrahedral
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4 Bonding
Molecular
Geometry
Angular or Bent
120 Bond Angle
Tetrahedral
3 Bonding
1 Nonbonding
Trigonal
Pyramidal
2 Bonding
1 Nonbonding
Angular or Bent
109 Bond Angle
4–48
CO2, a Molecule with Linear Electronic
and Molecular Geometry
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4–49
Molecules with Trigonal Planar Electronic Geometry
and Trigonal Planar or Angular Molecular Geometry
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4–50
Molecules with Tetrahedral Electronic Geometry;
Angular, Trigonal Pyramidal, or Tetrahedral
Molecular Geometry
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4–51
Chemistry at a
Glance:
The Shape
(Geometry) of
Molecules
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4–52
4.15 Electronegativity
Electronegativity is a measure of the relative
attraction that atoms have for the shared
electrons in a bond.
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4–53
Figure 4.14
Abbreviated periodic table showing
Pauling electronegativity values for
selected representative elements.
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4–54
4.16 Bond Polarity
Ionic and Covalent Bonds are ends
of a continuum of bond types
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4–55
Figure 4.15
(a) In the nonpolar covalent bond in H2,
equal sharing of electrons occurs.
(b) In the polar covalent bond present
of HCl, unequal sharing of electrons
occurs. The electronegativity of Cl
is greater than that of H
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4–56
4.16 Bond Polarity
Ionic Bonds:
Electronegativity difference between
atoms is 2.0 or more.
Polar Covalent Bonds:
Electronegativity difference between
atoms is between 1.9 and 0.5.
Nonpolar Covalent Bonds:
Electronegativity difference between
atoms is 0.4 or less.
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4–57
Chemistry at a Glance:
Covalent Bonds and Molecular Compounds
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4–58
4.17 Molecular Polarity
Molecular polarity is a measure of the
degree of symmetry of electron
distribution in a molecule.
A nonpolar molecule has a symmetric
distribution of electronic charge.
A polar molecule has an assymmetric
distribution of electronic charge.
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4–59
4.17 Molecular Polarity
Molecular polarity is determined by
two factors:
1. Bond Polarities
2. Molecular Geometries
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4–60
Figure 4.16
(a) Methane (CH4) is a nonpolar tetrahedral molecule.
(b) Methyl chloride (CH3Cl) is a polar tetrahedral molecule.
Bond polarities cancel in CH4, but not CH3Cl.
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4–61
4.17 Molecular Polarity
Molecular polarity is determined by
two factors:
1. Bond Polarities
2. Molecular Geometries
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4–62
4.17 Molecular Polarity
Rule:
If all the electron regions around the central
atom in a simple structure are bonds,
and
all the substituent (surrounding) atoms are
the same:
the molecule or ion is not polar.
In any other case, it is polar.
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4–63
4.17 Molecular Polarity
Examples:
Which of these molecules are polar?
N2
CO2
CH4
SO3
CHCl3
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CH2O
NH3
SO2
H 2O
4–64
4.18 Naming Binary
Molecular Compounds
Give full name of less electronegative
element
Give stem of more electronegative element,
followed by "ide"
Numerical prefixes precede names of both
nonmetals, unless only one atom of the
first element is present in the formula.
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4–65
Table 4.2
Prefixes for 1 through 10.
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4–66
4.18 Naming Binary
Molecular Compounds
Examples:
Name: CO2, CO, SiO2, N2O, N2O4
Formula:
Sulfur dioxide
Nitrogen monoxide
Xenon tetrafluoride
Diphosphorous pentoxide
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4–67
Table 4.3
Selected Binary Molecular Compounds
that Have Common Names.
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4–68
Figure 4.17
Models of
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(a) a sulfate ion (SO42-)
(b) a nitrate ion (NO31-)
4–69
Table 4.4
Formulas and Names of Some Common
Polyatomic Ions.
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4–70
4.19 Polyatomic Ions
A polyatomic ion is a charged group of
covalently bound atoms.
OH1-
hydroxide
CN1-
NH41+
ammonium
NO31- nitrate
SO42-
sulfate
PO43- phospate
CO32-
carbonate
C2H3O21- = CH3COO1Copyright © Houghton Mifflin Company. All rights reserved.
cyanide
acetate
4–71
4.19 Polyatomic Ions
Naming compound with polyatomic anions:
Name the metal, using Roman numerals to
indicate charge if necessary
Name the anion
Naming compounds with ammonium cation:
Name the cation (ammonium)
Name the anion
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4–72
4.19 Polyatomic Ions
Examples:
Name:
NaOH Na2CO3 CuSO4
Formula:
(NH4)2SO4
Potassium cyanide
Sodium phosphate
Lead (II) acetate
Ammonium chloride
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4–73
4.19 Polyatomic Ions
Mixed salts are compounds that contain
Na1+ or K1+, hydrogen, and an anion with
a charge of -2 or -3.
Name:
NaHCO3
KH2PO4
KHSO4
K2HPO4
Formula: Sodium hydrogen sulfide
Sodium dihydrogen phosphate
Disodium hydrogen phosphate
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4–74
Nomeclature Practice
Name:
CO2
CaSO4
SiO2
MnO2
TiO2
NH4NO3 Li2CO3
Formula: Phosphorus trichloride
Copper (II) sulfate
Silver bromide
Lead (IV) oxide
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4–75