Chapter 8
Bonding:
General Concepts
Section 8.1
Chemical
Bonds Bonds
Types
of Chemical
 Forces that hold groups of atoms together and make them
function as a unit.
 A bond will form if the energy of the compound formed is
lower than that of the separated atoms.
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2
Section 8.1
Types of Chemical Bonds
All chemical reactions involve making
and breaking chemical bonds.
 Bond formation always releases energy
 Bond breaking always requires energy
Section 8.1
I won the Nobel Prize
in Chemistry and the
Types of Chemical Bonds
Nobel Peace Prize
Electronegativity:
 The ability of an atom in a molecule
to attract shared electrons to itself.
 The most widely accepted method
for determining values of
electronegativity was derived by
Linus Pauling.
 The range of electronegativity values
is from 4.0 for fluorine (the most
electronegative) to 0.7 for cesium
(the least electronegative).
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4
Section 8.1
Types of Chemical Bonds
The electronegativity increases across a period:
 Same energy level, but the number of protons
increases which increases the effective nuclear
charge.
The electronegativity decreases going down a group
 because the more energy levels are added and
the size of the atoms increases. The effective
nuclear charge is not as great due to shielding.
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5
Section 8.1
Types of Chemical Bonds
CONCEPT CHECK!
If lithium and fluorine react, which has more attraction
for an electron? Why?
 Fluorine has a greater nuclear pull. Both are in the
same energy level, but F has more protons. More
protons, greater nuclear attraction
In a bond between fluorine and iodine, which has more
attraction for an electron? Why?
 Fluorine has a greater nuclear pull. Although iodine
has a greater nuclear charge, the charge is being
‘shielded’ by the increase in energy levels.
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6
Section 8.1
Types of Chemical Bonds
As the different in electronegativity increases:
-bond polarity increases
-ionic character increases
Section 8.1
Types of Chemical Bonds
Three major type of intermolecular bonds can be described
by the electronegativity difference
1. Covalent (non-polar covalent): Bond difference 0-0.4
 electrons are shared equally
2. Polar Covalent: Bond difference 0.4  1.9
 Bond becomes increasingly polar with higher values;
 electrons shared unequally
3. Ionic: Bond difference greater than 1.9
 Large difference; Bond becomes more ionic in nature ;
electron transfer can occur.
Section 8.1
Types of Chemical Bonds
Put the following bonds in order of increasing ionic
character. Predict the type of bond each will incur.
H-Cl H-F
H-H
H-O
H-S
NP
Cov
NP
Cov
P
Cov
P
Cov
Ionic
Section 8.1
Types of Chemical Bonds
EXERCISE!
Arrange the following bonds from most to least polar:
a) N–F
b) C–F
c) Cl–Cl
a) C–F,
b) Si–F,
c) B–Cl,
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O–F
N–O
B–Cl
N–F,
C–F,
S–Cl,
C–F
Si–F
S–Cl
O–F
N–O
Cl–Cl
10
Section 8.1
Types of Chemical Bonds
CONCEPT CHECK!
Which of the following bonds would be the least
polar yet still be considered polar covalent?
Mg–O C–O O–O Si–O N–O
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Section 8.1
Types of Chemical Bonds
CONCEPT CHECK!
Which of the following bonds would be the most
polar without being considered ionic?
Mg–O C–O O–O Si–O N–O
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12
Section 8.1
TypesofofChemical
ChemicalBonds
Bonds
Types
1. Ionic Bonding
 Electrons are transferred
 Metal and nonmetal react: ions form so that both ions
achieve noble gas configurations (usually) resulting in an
electrostatic attraction of oppositely charged ions
 Any compound that conducts an electric current when
melted will be classified as ionic.
Section 8.1
TypesofofChemical
ChemicalBonds
Bonds
Types
Metals lose electrons to achieve noble gas configurations.
Non-metals gain electrons to achieve noble gas
configurations.
Section 8.1
Types of Chemical Bonds
Ex: Explain how atoms of sodium and oxygen form an ionic
compound using their electron configurations
Na: 1s22s22p63s1
O: 1s22s22p4
Oxygen has a higher electronegativity(the ability to attract an
electron), so electrons are transferred from sodium to
oxygen. Oxygen needs 2 electrons, so 2 sodium atoms are
needed. Na2O is formed.
2Na  2Na+ + 2e-
O + 2e-  O2-
Na+: 1s22s22p6
O2-: 1s22s22p6
Section 8.1
Types of Chemical Bonds
Lewis Structure (Electron Dot Diagrams)
 A simplistic way of showing the reactive electrons (the
valence electrons) of an atom.
 Reflects central idea that stability of a compound relates
to noble gas electron configuration.
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16
Section 8.1
dog structures
structures
Electron
dot
Types
of Chemical
Bonds
Draw the Lewis structure for each of the following:
1. O
2. Li
3. S
4. O25. Ca2+
Section 8.1
Types of Chemical Bonds
Use the electron dot structure to show how Na and O
can become an ion
Na

Na+
+
sodium Ion
Sodium atom
e1 electron
2-
O
Oxygen atom
+
2e2 electrons

O
Oxide Ion
Section 8.1
Showofthe
formation
of Na2O using Lewis structures
Types
Chemical
Bonds
[ ]
[ ]
Na2O
Section 8.1
Types of Chemical Bonds
 Show the formation of (A) calcium fluoride and
(B) aluminum oxide using electron-dot structures
Section 8.1
Lattice
Types
ofenergy
Chemical Bonds
Lattice Energy
 Lattice energy is the energy release that occurs when
separated gaseous ions are packed together to form an
ionic solid (lattice crystal)
Xx+(g) + Yy-(g)  XyYx(s) + energy
Ex: Na+(g) + Cl-(g) → NaCl(s) + 7787.3 kJ
Ho= -787.3 kJ/mol
 Lattice energy indicates how strongly the ions attract
each other in the solid state.
21
Section 8.1
Coulomb’s
Law Bonds
Types
of Chemical
We can calculate the lattice energy of an ionic compound by
using Coulomb’s Law:
 Q1Q2 
Lattice energy = k 

 r 
E = (2.31 x 10-19 J·nm) x Q1Q2
r
k = proportionality constant =2.31 x 10-19 J·nm
Q1 and Q2 = charges on the ions
r = distance between the ion centers in nm
A negative sign indicates an attractive force!
Section 8.1
Lattice
trend
Types
ofenergy
Chemical
Bonds
 Charge on ions: The greater the charge on the
ions, the stronger the ionic bond
 Radius or size of ion: If the charges are equal, the
smaller ion will have the stronger ionic bond.
Section 8.1
Types of Chemical Bonds
Example: Which would you predict has the most
exothermic lattice energy: NaCl or KCl? Explain.
 Q1Q2 
Lattice energy = k 

 r 
Answer: NaCl
 Both NaCl and KCl all have +1 or -1 charges.
 Since Na+ is smaller than K+, the distance, r, is
less, making the energy more.
 As this energy increases, the strength and stability
of the crystal lattice increases.
Section 8.1
Types of Chemical Bonds
Example: In solid sodium chloride, the distance between
the centers of the Na+ and Cl- ions is 2.76Å. Determine
the ionic energy per pair of ions. (1m = 1010Å).
E = (2.31 x 10-19 J·nm) x Q1Q2
r
E = (2.31 x 10-19 J·nm) x (+1)(-1)
0.276nm
= -8.37 x 10-19 J
The negative sign indicates an attractive force. The
ion pair has LOWER energy than the separated ions.
Section 8.1
Types of Chemical Bonds
 Coulomb’s law can also be used to calculate the
repulsive energy when two like charged ions are
brought together. This energy will have a positive
sign.
Example: The interaction of two hydrogen atoms
 The hydrogen atoms will position themselves so
that the system will achieve the lowest possible
energy.
 The distance where energy is minimal is called the
bond length.
Section 8.1
Types of Chemical Bonds
The Interaction of Two Hydrogen Atoms: the molecule is
more stable than two separated hydrogen atoms by a
certain quantity of energy.
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27
Section 8.1
Types of Chemical Bonds
2. Covalent Bonding
 Electrons are shared by nuclei
 Usually formed by bonding
nonmetals
 Non-polar covalent –
electrons are shared equally
 Polar covalent – electrons
are shared unequally
 Results in a partial
positive and partial
negative charge.
Section 8.1
Types of Chemical Bonds
Ionic and Covalent bonding animation
http://www.youtube.com/watch?v=QqjcCvzWwww
Section 8.2
Electronegativity
 Hydrogen forms stable molecules where it shares two
electrons. (duet rule)
Section 8.2
Electronegativity
Octet Rule
 Almost all other elements form stable molecules when
surrounded by eight electrons.
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Section 8.2
Electronegativity
 Single covalent bond: two atoms share one pair
of electrons.
H–H
 Example: All halogens form single covalent
bonds
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32
All th
Section 8.2
Electronegativity
 A double covalent bond:
two atoms share two pairs
of electrons.
Example: O=O
 A triple covalent bond:
two atoms share three
pairs of electrons.
Example: N  N
Section 8.2
Electronegativity
Bond Dissociation Energy
 Energy required to break a covalent bond
 Large dissociation energy means strong bond
 The more bonds located between 2 atoms, the
shorter the bond.
 The shorter the bond, the stronger the bond.
H–H Single bond
Weaker
O=O Double bond Stronger
N≡N Triple Bond
Strongest
Section 8.2
Bond Energies
Electronegativity
 To break bonds, energy must be added to the system
(endothermic, energy term carries a positive sign).
 To form bonds, energy is released (exothermic, energy
term carries a negative sign).
H = D(bonds broken) – D(bonds formed)
D represents the bond energy per mole of
bonds (always has a positive sign).
35
Section 8.2
Electronegativity
CONCEPT CHECK!
Predict ∆H for the following reaction:
CH3N  C(g )  CH3C  N(g )
Given the following information:
Bond Energy (kJ/mol)
C–H
413
C–N
305
C–C
347
891
CN
[3(413) + 305 + 891] – [3(413) + 347 + 891] = –42 kJ
∆H = –42 kJ
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36
Figure 10.3
Using bond energies to calculate DH0comb. of Methane, CH4
BOND BREAKAGE
4BE(C-H)= +1652kJ
2BE(O2)= + 996kJ
DH0(bond breaking) = +2648kJ
BOND FORMATION
Enthalpy,H
2[-BE(C O)]= -1598kJ
4[-BE(O-H)]= -1868kJ
DH0(bond forming) = -3466kJ
DH0rxn= -818 kJ/mol
Section 8.2
How to draw a Lewis structure
Electronegativity
1. Find the total number of valence electrons. For anions,
add electrons; for cations, subtract.
2. Choose the central atom. The atom in front is often the
central. The central atom is usually the least
electronegative, or the atom with the highest valence.
3. Draw a skeletal structure – connect the atoms. No ring
around the rosy!
4. Arrange the remaining electrons to satisfy the octet
rule. If there are not enough electrons, form double or
triple bonds.
Section 8.2
Electronegativity
Example: Draw a Lewis structure for water.
Notice that a pair of
valence electrons are not
shared. They are ‘lone’
pairs.
Draw a Lewis structure for each of the following
molecules:
1. NH3
5. NH4+
2. HCN
6. CO2
3. CCl4
7. SO3
4. SO428. NO2-
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39
Section 8.2
Exceptions to the octet rule
Electronegativity
Some molecules fail to follow the octet
rule, yet are stable molecules.
 The total number of valence electrons is
odd. (free radicals)
 Ex: NO2
 Incomplete octets need less than 8
valence electrons.
 Ex: B and Be often do not obey the octet
rule
H
BH3
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H B H
40
Section 8.2
Electronegativity
 Expanded octets exceed the octet rule. Extra electrons
are placed on the central atom. NOT on AP exam!
 Example:
SF4 = 34e–
F
F S
F
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F
AsBr5 = 40e–
Br Br
Br As Br
Br
41
Section 8.2
Resonance
Electronegativity
Resonance occurs when more than one valid Lewis
structure can be written for a particular molecule.
Example: NO3–

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
42
Section 8.2
Electronegativity
 Actual structure is an average or hybrid of the resonance
structures.
 Electrons are really delocalized – they can move around
the entire molecule.
 Its as if the structures were changing infinitely fast, so
that electrons and charges were everywhere at once.
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43
Section 8.2
Electronegativity
CONCEPT CHECK!
Draw a Lewis structure for each of the following
molecules:
CO
CH3OH
C2H6
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CO2
OCN–
C2H4
44
Section 8.13
Molecular Structure: The VSEPR Model
VSEPR: Valence Shell Electron-Pair Repulsion.
 Used to predict a 3-dimensional shape of a
molecule
 Based on the premise that electrons repel each
other. To minimize energy, electrons try to get
as far apart as possible.
 Works in many, but not all cases.
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45
Section 8.13
Molecular Structure: The VSEPR Model
VSEPR Theory is based upon how many electron
‘groups’ are bonded to a central atom.
 Steric Number (or substituents)
 The steric number is equal to the number of
atoms attached to the central atom plus the
number of lone pairs of valence electrons on
the central atom.
 Examples: How many steric numbers are on
the central atom?
CH4
H20
CO2
Section 8.13
2 atomsStructure:
joined The
together
Molecular
VSEPR Model
 Shape: Linear
Bond Angle: 180°
Ex: H2, HCl
Steric8.13
Number: 2
Section
Molecular
Structure: sp
The VSEPR Model
Hybridization:
 Shape: Linear
Bond Angle: 180°
Number of lone pairs: 0
Ex: BeH2, CO2
Section 8.13
VSEPR: Two Electron Pairs
Molecular Structure: The VSEPR Model
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49
Steric8.13
Number: 3
Section
Molecular
Structure: sp
The2 VSEPR Model
Hybridization:
 Shape: trigonal planar
Bond Angle: 120°
Number of lone pairs: 0
Ex: SO3, CO32-
 Shape: bent
Bond angle: 120°
Number of lone pairs: 1
Ex: SO2, NO2-
Section 8.13
VSEPR: Three
Electron
Pairs
Molecular
Structure:
The
VSEPR Model
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51
Section
Steric8.13
Number: 4
Molecular
Structure: The
Hybridization:
sp3VSEPR Model
 Shape: tetrahedral
Bond angle: 109.5°
Number of lone pairs: 0
Examples: SiH4, CH2Cl2
 Shape: trigonal pyramidal
Bond angle: 107°
Number of lone pairs: 1
Examples: PH3, NH3
 Molecular Geometry: bent
Bond angle: 105°
Number of lone pairs: 2
Example: H2O
Section 8.13
VSEPR: Four Electron Pairs
Molecular Structure: The VSEPR Model
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53
Section 8.13
VSEPR: Review
of allThe
shapes
Molecular
Structure:
VSEPR Model
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54
Section 8.13
Molecular Structure: The VSEPR Model
CONCEPT CHECK!
Determine the shape for each of the following
molecules, and include bond angles:
HCN
PH3
O3
HCN – linear, 180o
PH3 – trigonal pyramid, 109.5o (107o)
O3 – bent, 120o
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55
Section 8.13
Hybridization
Molecular
Structure: The VSEPR Model
 Covalent bonds are formed by the sharing of
electrons; orbitals overlap to allow for this
sharing.
 The mixing of two or more atomic orbitals of an
atom forms hybrid orbitals. This process is called
hybridization.
 Predicting hybridization is easy – just count the
total number of steric numbers.
Section 8.13
Molecular Structure: The VSEPR Model
 The total number of steric numbers (also known as
substituents - bonding plus non-bonding groups) is
equal to the number of atomic orbitals that
participate in the hybrid orbital.
# of substituents
(steric numbers)
Hybridization
Example Molecule
2
sp
CO2
3
sp2
CH3
4
sp3
CH4, NH3, H2O
5
sp3d
PCl5, I3-
6
sp3d2
SF6
Section 8.13
The sp hybrid
orbitalThe VSEPR Model
Molecular
Structure:
 In sp hybridization, only one p orbital is mixed with the s
orbital
 Example: BeF2 - Steric Number: 2
 Electron configuration of Be: 1s22s2
 Electron configuration of F:1s22s22p5
Section 8.13
Molecular Structure: The VSEPR Model
In the ground state, Be has no unpaired electrons – so how
can the Be atom form a covalent bond with a fluorine?
Be
F
Be obtains an unpaired electron by moving one electron
from the 2s orbital to the 2p orbital resulting in two
unpaired electrons, one in a 2s orbital and another in a
2p orbital
Section 8.13
Molecular Structure: The VSEPR Model
 The Be atom can now form two covalent bonds with
fluorine atoms
Be
F
 Although we would not expect these bonds to be
identical (one is in a 2s electron orbital, the other is
in a 2p electron orbital), the structure of BeF2 is
linear and the bond lengths are identical
 The 2s and 2p electrons produced a "hybrid" orbital
for both electrons
Section 28.13
The sp hybrid
orbital
Molecular
Structure:
The VSEPR Model
 In sp2 hybridization,
two p orbitals are mixed
with the s orbital to
generate three new
hybrids
 Example: BF3
Section 38.13
The sp hybrid
orbital
Molecular
Structure:
The VSEPR Model
 In sp3 hybridization, all three p orbitals are mixed
with the s orbital to generate four new hybrids
 Example: CH4
Section 8.13
Sigma (s)
and Pi Bonds
(p) Model
Molecular
Structure:
The VSEPR
1 sigma bond
Single bond
Double bond
1 sigma bond and 1 pi bond
Triple bond
1 sigma bond and 2 pi bonds
How many s and p bonds are in the acetic acid
Molecule, CH3COOH?
H
C
H
O
H
C
O
H
s bonds = 6 + 1 = 7
p bonds = 1
More fun stuff about
POLAR BEARS
BONDS
and
MOLECULES
What is a Polar Molecule?
• In a polar molecule,
one end is slightly
positive, the other
slightly negative.
• Dipole: A molecule
that has two poles
+
-
I know how to predict if a bond will
be polar, but how can I predict if
the molecule will be polar?
In order to determine this, you must
first know:
1 - bond polarity
2 - molecular shape (VSEPR)
POLAR MOLECULES
1.Must have at least 1 polar bond.
2.Most molecules with lone pairs
on the central atom are polar!
ex: bent & pyramidal
3.Tetrahedral shape must be
unsymmetrical (Different atoms
bonded to the central atom)
ex: CH2Cl2
Nonpolar Molecules
1. May have polar bonds.
2. Molecules with no lone pairs on the
central atom are usually nonpolar.
ex: trigonal planar & triatomic linear
3. Tetrahedral shape must be
symmetrical (Same atoms bonded to
the central atom)
ex: CH4, CF4
Examples . . .
• HCl
• Cl2
Examples . . .
• HCN
• CO2
Examples . . .
• H2O
• SI2
Examples . . .
• BF3
• CH2O
Examples . . .
• NH3
• PH3
Examples . . .
• CH4
• CH2Cl2
The End . . .
Attraction between molecules
Hey baby . . .
Wanna bond,
sugar-pie?
Nice hydrogen's,
sweetie, . . . Sure
would like to put
my oxygen on
them . . .
Intramolecular vs
Intermolecular bonding . . .
• Intramolecular forces are the forces
of attraction which hold an individual
molecule together
– Non-polar covalent bonds
– Polar covalent bonds
– Ionic bonds
• Intermolecular attractions are
attractions between one molecule and a
neighboring molecule.
All molecules experience
intermolecular attractions
This is the force that enables molecules
to change state to form solids &
liquids
– Solid: Intermolecular forces are strong
– Liquid: Intermolecular forces are weak
– Gas: Intermolecular forces are
especially weak
Types of intermolecular forces:
1. Dispersion
forces
2. Dipole-dipole
attraction
3. Hydrogen
bonds
Collectively, these
are called “van der
Waals forces” after
me, Johannes van der
Waals
Dispersion Forces
• Dispersion Forces are
weak forces caused
by an instantaneous
dipole, in which
electron distribution
becomes
asymmetrical.
• These are the ONLY
forces of attraction
that exist among
noble gas atoms and
nonpolar molecules.
Dispersion forces are
also called London
Dispersion Forces, or
just London Forces after
me, Fritz London.
Dispersion forces are temporary
. . . They are just caused by the
motion of electrons!
London Dispersion Forces
Dispersion forces increase in
strength as the number of
electrons increases & the molecular
weight increases.
The more electrons
you have and the
bigger you are,
the bigger the
possible temporary
dipoles.
Would you expect
neon or xenon to have
a higher boiling point?
Why?
Because of the greater temporary dipoles,
xenon molecules are "stickier" than neon
molecules. Neon molecules will break
away from each other at much lower
temperatures than xenon molecules
Compare the boiling points of the noble gases:
helium
-269°C
neon
-246°C
argon
-186°C
krypton
-152°C
xenon
-108°C
radon
-62°C
Dipole-Dipole Forces
• Dipole-Dipole attraction:
occurs in molecules that have a
permanent dipole.
• Molecules with dipoles orient
themselves so that “+” and “-”
ends of the dipoles are close to
each other.
• The strength of the attraction
increases as the polarity
increases.
• These occur in addition to
dispersion forces
Dipole
Forces
A special type of Dipole-Dipole. . .
Hydrogen Bonding
• Hydrogen bond: A dipole-dipole attraction
in which hydrogen is bound to a highly
electronegative atom.
• Hydrogen bonding is the strongest of the
intermolecular forces.
What types of molecules will hydrogen bond?
• The molecule must have a hydrogen bonded
to a very electronegative atom (F, O, N)
• This hydrogen can bond to another
electronegative atom with a lone pair.
• For example, all of these molecules can
hydrogen bond . . .NH3, H2O, HF. Why??
• If an F, O, or N has at least one lone
pair, and it is attached to a H, then it can
hydrogen bond!
Hydrogen Bonding in Water . . .
The perfect example
Review:
Is the molecule polar?
No
Yes
Is there a H attached to F, O, or N
with at least 1 lone pair
No
Dispersion
Dipole-Dipole
Yes
H-Bond
Properties
Ionic Compounds:
•At room temperature they are a
crystalline solid.
•Hard & Brittle, High melting point
Covalent (Molecular Compounds):
•At room temperature they can be a
solid, liquid or gas.
•Usually soft & low melting
point
Properties
Metallic Solids
•Soft to very hard solids
Why?
Strength increases with an increase in #
of e- available for bonding.
•Low to very high melting points
•Ductile:Metal can be drawn into thin wires
•Malleable: Metal can be hammered into
different shapes
Compound Arrangement
Ionic Compound
•Three dimensional pattern
How is the 3D pattern created?
•Oppositely charged ions form
closely packed layers.
•Layers stack together to make a
regular repeating 3D pattern called a:
Unit cell
•Closely packed layers will keep
stacking to form a: Crystal
Coordination
Number of
the Ion
Is the # of
ions of
opposite
charge that
surround the
ion in a
crystal
How many
Na+ surround
Cl-?
Compound Arrangement
Molecular Compound
•Consist of
atoms
(nonmetals)
which form
molecules that
are held
together by
Weak
weak forces. forces
Water molecules
Compound Arrangement
Metallic Solid
•Consist of closely packed cations.
How are metallic bonds formed?
•It is formed from the attraction of
free-moving valence electrons
•These electrons bind the atoms
together to create a metallic solid.
Electron
cloud
e- eForm a
bond
Type of Bond Force & Strength
Electrostatic forces
Very Strong
Ionic Compounds:
Covalent (Molecular Compounds):
Intermolecular forces
Weak
Dispersion forces, Dipole-dipole, H-bonding
Metallic Solids: A “sea” of Valence Electrons
Strong
Conductors
Ionic Compounds:
Good conductors of electricity
Two ways ionic compounds
conduct electricity:
• Conduct an electric current in the
melted state.
• Conduct electricity when the
compound is dissolved in water.
Conductors
How?
•Bonds break or dissociate.
•The ions separate.
•Ions are free to move around.
•When a voltage is applied the
cations move to one electrode & the
anions to the other.
The ion movement produces a flow
of electricity
Conductors
Molecular Compounds: Poor conductors
Why?
• Bonds do not break easily.
• Compound does not easily dissolve in
water.
• Compound does not contain ions that
conduct electricity.
Conductors
Metallic Solids:
Good conductors of electricity
Why?
• Free-moving valence electrons
Good conductors of heat
Why?
• Free-moving valence
electrons
Good conductors of heat
& electricity.
Network Solids
•Consist of atoms held together by large
chains of covalent bonds.
•Each atom is covalently bonded in a
large chain or network.
•There are NO molecules
in a network solid, only
atoms bonded together.
•Bonds are very strong.
Example: Diamond
Metallic solids can arrange its bonds to
form 3 orderly patterns:
Body-centered cubic
Face-centered cubic
+
Hexagonal close- packed
Why are metallic solids the
simplest crystalline solid?
Section 8.13
Molecular Structure: The VSEPR Model
Section 8.13
Molecular Structure: The VSEPR Model
CONCEPT CHECK!
True or false:
A molecule that has polar bonds will always be polar.
-If true, explain why.
-If false, provide a counter-example.
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108
Section 8.13
Molecular Structure: The VSEPR Model
Let’s Think About It
 Draw the Lewis structure for CO2.
 Does CO2 contain polar bonds?
 Is the molecule polar or nonpolar overall? Why?
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Section 8.13
Molecular Structure: The VSEPR Model
CONCEPT CHECK!
True or false:
Lone pairs make a molecule polar.
-If true, explain why.
-If false, provide a counter-example.
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110
Section 8.13
Molecular Structure: The VSEPR Model
PUT IN CH 7 NOTES Isoelectronic Series
 A series of ions/atoms containing the same number of
electrons.
O2-, F-, Ne, Na+, Mg2+, and Al3+
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Section 8.13
Molecular Structure: The VSEPR Model
Ionic Radii
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112
Section 8.13
Molecular Structure: The VSEPR Model
CONCEPT CHECK!
Choose an alkali metal, an alkaline earth metal, a noble gas, and
a halogen so that they constitute an isoelectronic series when
the metals and halogen are written as their most stable ions.



What is the electron configuration for each species?
Determine the number of electrons for each species.
Determine the number of protons for each species.
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Section 8.13
Molecular Structure: The VSEPR Model
Periodic Table Allows Us to Predict Many Properties
 Trends for:
 Atomic size, ion radius, ionization energy,
electronegativity
 Electron configurations
 Formula prediction for ionic compounds
 Covalent bond polarity ranking
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Section 8.13
Molecular Structure: The VSEPR Model
isoelectronic
Section 8.13
Molecular Structure: The VSEPR Model
Metallic bonding
The valence electrons of a metal can be modeled
as a "see of electrons."
The valence electrons are free to move.
 The metallic bond consists of the attraction of
free-floating valence electrons for the positivelycharged metal ions.
This model explains many of the metallic
properties: conductivity, malleability, and ductility
Section 8.13
Molecular Structure: The VSEPR Model
Bond Polarity and Dipole Moments
 Dipolar Molecules: Molecules with a somewhat negative
end and a somewhat positive end (a dipole moment)
 Use an arrow to represent a dipole moment.
 Point to the negative charge center with the tail of the
arrow indicating the positive center of charge.
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117
Section 8.13
Molecular Structure: The VSEPR Model
Dipole Moment
118
Section 8.13
Molecular Structure: The VSEPR Model
No Net Dipole Moment (Dipoles Cancel)
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119
Section 8.13
Molecular Structure: The VSEPR Model
 Molecules that have a polar bond may or may not
show a dipole moment – the shape of the
molecule must be considered.
Section 8.13
Molecular Structure: The VSEPR Model
Localized Electron Model
 A molecule is composed of atoms that are bound
together by sharing pairs of electrons using the atomic
orbitals of the bound atoms.
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Section 8.13
Molecular Structure: The VSEPR Model
Localized Electron Model
 Electron pairs are assumed to be localized on a
particular atom or in the space between two atoms:
 Lone pairs – pairs of electrons localized on an atom
 Bonding pairs – pairs of electrons found in the space
between the atoms
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122
Section 8.13
Molecular Structure: The VSEPR Model
Localized Electron Model
1. Description of valence electron arrangement (Lewis
structure).
2. Prediction of geometry (VSEPR model).
3. Description of atomic orbital types used by atoms to
share electrons or hold lone pairs.
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Section 8.13
Molecular Structure: The VSEPR Model
 https://www.youtub
e.com/watch?v=7Djs
D7Hcd9U
Section 8.13
Molecular Structure: The VSEPR Model