Download Section 2 Covalent Bonding and Molecular Compounds Chapter 6

Chapter 6
Chemical Bonding
Table of Contents
Section 1 Introduction to Chemical Bonding
Section 2 Covalent Bonding and Molecular Compounds
Section 3 Ionic Bonding and Ionic Compounds
Section 4 Metallic Bonding
Section 5 Molecular Geometry
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Chapter 6
Section 1 Introduction to
Chemical Bonding
Objectives
• Define chemical bond.
• Explain why most atoms form chemical bonds.
• Describe ionic and covalent bonding.
• Explain why most chemical bonding is neither purely
ionic nor purely covalent.
• Classify bonding type according to
electronegativity differences.
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Chemical Bond
An attractive force that holds atoms or ions together
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Types of Chemical Bonds
Ionic
-electrical attraction between positive and negative ions
-transfer of electrons
-usually a metal and a nonmetal
Cation- positive ion
Anion- negative ion
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Types of Chemical Bonds, cont’d
Covalent
-sharing of electrons between two atoms to make a bond
-can be:
nonpolar covalent (equal sharing)
polar covalent (unequal sharing)
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Types of Chemical Bonds, cont’d
Determining the Type of Bond
-bonds are not 100% ionic or covalent
-use the table of electronegativities (pg.159)
3.3 – 1.8 ionic bonds
1.7- 0.3 polar covalent
0.2- 0 nonpolar covalent
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Chapter 6
Section 1 Introduction to
Chemical Bonding
Chemical Bonding, continued
Sample Problem A
Use electronegativity values listed on page 159, and
Figure 2 in your book, on page 172, to classify bonding
between sulfur, S, and the following elements:
hydrogen, H; cesium, Cs; and chlorine, Cl. In each pair,
which atom will be more negative?
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Chapter 6
Section 2 Covalent Bonding and
Molecular Compounds
Objectives
• Define molecule and molecular formula.
• Explain the relationships among potential energy,
distance between approaching atoms, bond length,
and bond energy.
• State the octet rule.
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Chapter 6
Section 2 Covalent Bonding and
Molecular Compounds
Molecular Compounds
• A molecule is a neutral group of atoms that are held
together by covalent bonds.
• A chemical compound whose simplest units are
molecules is called a molecular compound.
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Chapter 6
Visual Concepts
Molecule
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Molecular Compounds
One type of molecular compound is called a
Diatomic molecule
examples- O2, this is composed of two atoms of oxygen
others:
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Chapter 6
Section 2 Covalent Bonding and
Molecular Compounds
Formation of a Covalent Bond
• Atoms have lower potential energy when they are
bonded to other atoms
• The figure below shows potential energy changes
during the formation of a hydrogen-hydrogen bond.
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Chapter 6
Section 2 Covalent Bonding and
Molecular Compounds
Characteristics of the Covalent Bond
• The distance between two bonded atoms at their
most stable energy state is the bond length.
• In forming bond, atoms release energy. The same
amount of energy must be added to separate the
bonded atoms. This is called Bond energy .
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Chapter 6
Section 2 Covalent Bonding and
Molecular Compounds
Bond Energies & Bond Lengths for Single
Bonds- Inverse relationship
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Characteristics of the Covalent Bond
When two atoms form a covalent bond, their outer orbitals
overlap.
This achieves a noble-gas configuration.
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Chapter 6
Section 2 Covalent Bonding and
Molecular Compounds
The Octet Rule
• Bond formation follows the octet rule:
• Compounds tend to form so that each atom, by
gaining, losing, or sharing electrons, has an
octet of electrons in its highest energy level.
• Become like noble gases.
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Chapter 6
Section 2 Covalent Bonding and
Molecular Compounds
Exceptions to the Octet Rule
• Atoms have more or less than 8 electron in outer
energy level
• Hydrogen - only two electrons.
• Boron – only six electrons.
• Main-group elements in Periods 3-7 can form bonds
with expanded valence, involving more than eight
electrons.
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Electron-Dot Notation
Bonding involves the outer s & p electrons (valence
electrons).
Using a dot notation shows the outer electrons each atom
has for bonding.
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Chapter 6
Section 2 Covalent Bonding and
Molecular Compounds
Electron-Dot Notation, continued
Sample Problem B
a. Write the electron-dot notation for hydrogen.
b. Write the electron-dot notation for nitrogen.
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Chapter 6
Section 2 Covalent Bonding and
Molecular Compounds
Objectives, continued
• Use the basic steps to write Lewis structures.
• Determine Lewis structures for molecules containing
single bonds, multiple bonds, or both.
• Explain why scientists use resonance structures to
represent some molecules.
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Chapter 6
Section 2 Covalent Bonding and
Molecular Compounds
Lewis Structures
• Electron-dot notation can also be used to represent
molecules.
• The pair of dots between the two symbols represents
the shared electron pair
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Chapter 6
Section 2 Covalent Bonding and
Molecular Compounds
Lewis Structures
• The pair of dots between the two symbols represents
the shared pair of a covalent bond.
• In addition, each atom is surrounded by three pairs of
electrons that are not shared in bonds.
• An unshared pair, also called a lone pair, is a pair of
electrons that is not involved in bonding and that
belongs exclusively to one atom.
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Section 2 Covalent Bonding and
Molecular Compounds
Chapter 6
Lewis Structures
• The pair of dots representing a shared pair of
electrons in a covalent bond is often replaced by a
long dash.
• example:
H
H
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Rules for Drawing Lewis Structures
1. Determine the total valence electrons for all atoms.
2. Draw a skeleton. Central atom is C or least
electronegative. Never H.
3. Give each atom an octet. (some exceptions)
4. Double or triple bond if you don’t have enough
electrons. (COPSN)
5. Extra electrons? Put on the central atom.
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Chapter 6
Section 2 Covalent Bonding and
Molecular Compounds
Lewis Structures, continued
Examples:
Draw the Lewis structure of
a. iodomethane, CH3I.
b. Ammonia, NH3
c. Silane, SiH4
d. Phosphorous trifluoride, PF3
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Chapter 6
Section 2 Covalent Bonding and
Molecular Compounds
Multiple Covalent Bonds
• A double bond- two pairs of electrons are shared
between two atoms.
• A triple bond-three pairs of electrons are shared
between two atoms.
• In general, double bonds have greater bond energies and
are shorter than single bonds.
• Triple bonds are even stronger and shorter than
double bonds.
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Lewis Structures with Multiple Bonds
Examples:
a.
b.
c.
d.
Formaldehyde, CH2O
Carbon dioxide, CO2
Hydrogen cyanide, HCN
Ethyne, C2H2
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Chapter 6
Section 3 Ionic Bonding and
Ionic Compounds
Polyatomic Ions
• A charged group of covalently bonded atoms is
known as a polyatomic ion.
• Like other ions, polyatomic ions have a charge that
results from either a shortage or excess of electrons.
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Chapter 6
Section 3 Ionic Bonding and
Ionic Compounds
Polyatomic Ions, continued
• Some examples of Lewis structures of polyatomic
ions
• nitrate NO3-1
• ammonium, NH4 +1
• sulfate, SO4-2
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Chapter 6
Section 3 Ionic Bonding and
Ionic Compounds
Objectives
• Compare a chemical formula for a molecular
compounds with one for an ionic compound.
• Discuss the arrangements of ions in crystals.
• Define lattice energy and explain its significance.
• List and compare the distinctive properties of ionic
and molecular compounds.
• Write the Lewis structure for a polyatomic ion given
the identity of the atoms combined and other
appropriate information.
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Chapter 6
Section 3 Ionic Bonding and
Ionic Compounds
Ionic Compounds
• Most of the rocks and minerals that make up Earth’s
crust consist of positive and negative ions held
together by ionic bonding.
• example: table salt, NaCl,
• An ionic compound is composed of positive and
negative ions that are combined so that the numbers
of positive and negative charges are equal.
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Chapter 6
Section 3 Ionic Bonding and
Ionic Compounds
Ionic Compounds
• Most ionic compounds exist as crystalline solids.
• In contrast to a molecular compound, an ionic
compound is not composed of independent, neutral
units that can be isolated.
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Chapter 6
Section 3 Ionic Bonding and
Ionic Compounds
Ionic Compounds, continued
• The chemical formula of an ionic compound
represents the simplest ratio of the compound’s
ions.
• A formula unit is the simplest collection of atoms
from which an ionic compound’s formula can be
established.
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Chapter 6
Section 3 Ionic Bonding and
Ionic Compounds
Ionic Vs. Covalent Bonding
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Chapter 6
Section 3 Ionic Bonding and
Ionic Compounds
Formation of Ionic Compounds
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Chapter 6
Section 3 Ionic Bonding and
Ionic Compounds
Formation of Ionic Compounds, continued
• In an ionic crystal, ions minimize their potential energy
by combining in an orderly arrangement known as a
crystal lattice.
• The combined attractive and repulsive forces within a
crystal lattice determine:
• the distances between ions.
• the pattern of the ions’ arrangement in the crystal
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Chapter 6
Section 3 Ionic Bonding and
Ionic Compounds
NaCl and CsCl Crystal Lattices
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Chapter 6
Section 4 Metallic Bonding
Objectives
• Describe the electron-sea model of metallic bonding,
and explain why metals are good electrical
conductors.
• Explain why metal surfaces are shiny.
• Explain why metals are malleable and ductile but
ionic-crystalline compound are not.
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Chapter 6
Section 4 Metallic Bonding
Metallic Bonding
• Chemical bonding is different in metals than it is in
ionic, molecular, or covalent-network compounds.
• The unique characteristics of metallic bonding gives
metals their characteristic properties, listed below.
• electrical conductivity
• thermal conductivity
• malleability
• ductility
• shiny appearance
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Chapter 6
Section 4 Metallic Bonding
Properties of Substances with Metallic, Ionic, and
Covalent Bonds
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Chapter 6
Section 4 Metallic Bonding
The Metallic-Bond Model
• In a metal, the vacant orbitals in the atoms’ outer
energy levels overlap.
• This overlapping of orbitals allows the outer electrons
of the atoms to roam freely throughout the entire metal.
• The electrons are delocalized, which means that they
do not belong to any one atom but move freely about
the metal’s network of empty atomic orbitals.
• These mobile electrons form a sea of electrons around
the metal atoms, which are packed together
in a crystal lattice.
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Chapter 6
Section 4 Metallic Bonding
Metallic Bonding, continued
• Mobile electrons allow charge/heat to pass and
allow the metal to be hammered because there is
not a rigid structure.
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Chapter 6
Section 5 Molecular Geometry
Objectives
• Explain VSEPR theory.
• Predict the shapes of molecules or polyatomic ions
using VSEPR theory.
• Explain the shapes of molecules or polyatomic ions
using VSEPR theory.
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Chapter 6
Section 5 Molecular Geometry
Molecular Geometry
• The properties of molecules depend not only on the
bonding of atoms but also on molecular geometry: the
three-dimensional arrangement of a molecule’s atoms.
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Chapter 6
Section 5 Molecular Geometry
VSEPR Theory
• As shown at right, diatomic
molecules, like those of
(a) hydrogen, H2, and
(b) hydrogen chloride, HCl,
can only be linear because
they consist of only two
atoms.
• To predict the geometries of more-complicated
molecules, one must consider the locations of all
electron pairs surrounding the bonding atoms. This
is the basis of VSEPR theory.
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Chapter 6
Section 5 Molecular Geometry
VSEPR Theory
• VSEPR stands for
• “valence-shell electron-pair repulsion.”
• VSEPR theory states that repulsion between the sets
of valence-level electrons surrounding an atom
causes these sets to be oriented as far apart as
possible.
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Chapter 6
Section 5 Molecular Geometry
VSEPR Theory
• Find the VSEPR formula
•A=central atom
•B=bonded atoms
•E=lone pairs of e- on the central atom
•Use the chart on pg.200 to determine the molecular shape
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Chapter 6
Section 5 Molecular Geometry
VSEPR Theory, continued
Sample Problems
Use VSEPR theory to predict the molecular geometry of
a. BCl3
b. NH3
c. CH4
d. NO3-1
e. CO2
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Chapter 6
Section 5 Molecular Geometry
VSEPR Theory, continued
• Unshared electron pairs repel other electron pairs
more strongly than bonding pairs do.
• This is why the bond angles in ammonia and water are
somewhat less than the 109.5° bond angles of a perfectly
tetrahedral molecule.
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Chapter 6
Section 5 Molecular Geometry
VSEPR Theory, continued
• Treat double and triple bonds the same way as
single bonds.
• Treat polyatomic ions similarly to molecules.
• The next slide shows several more examples of
molecular geometries determined by VSEPR
theory.
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Chapter 6
Section 5 Molecular Geometry
VSEPR and Molecular Geometry
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Chapter 6
Section 5 Molecular Geometry
VSEPR and Molecular Geometry
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Chapter 6
Section 5 Molecular Geometry
Intermolecular Forces a.k.a. van der Waals Forces
• The forces of attraction between molecules are
known as intermolecular forces.
• The boiling point is a good measure of the intermolecular
forces between its molecules:
•the higher the boiling point, the stronger the forces
between the molecules.
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Chapter 6
Section 5 Molecular Geometry
Intermolecular Forces, continued
• The strongest intermolecular forces exist between
polar molecules.
•Identical atoms attached = nonpolar
•Different atoms attached = polar
•Polar bonds are the result of an asymmetrical
distribution of electrons- causes a dipole
•Nonpolar bonds- equal distribution of electrons
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Chapter 6
Section 5 Molecular Geometry
Intermolecular Forces, continued
• A dipole is represented by an arrow with its head
pointing toward the negative pole and a crossed tail at
the positive pole. The dipole created by a hydrogen
chloride molecule is indicated as follows:
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Chapter 6
Section 5 Molecular Geometry
Intermolecular Forces, continued
• Determine if each of the molecules is polar or
nonpolar:
• CH4
• Cl2
• CCl3H
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Intermolecular Forces
a.k.a-van der Waals forces
Three types
1. Hydrogen Bonding
2. Dipole-Dipole
3. London Dispersion Forces
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Hydrogen Bonding
-Strongest
-occurs between molecules that contain O-H, N-H or H-F
atoms bonding.
-example:
H2O vs.
H2S
Boiling point
100 C
-61 C
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Dipole-Dipole
-2nd in strength
- occurs in polar molecules
- will see an increase in boiling point as the mass
increases
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London Dispersion Forces
-3rd in strength
-created from “momentary” dipoles that occur between
two nonpolar molecules
-also-increase in strength with increase in mass
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Intermolecular Forces
Let’s try some:
1. Draw Lewis Structure
2. Determine if the molecule is polar or nonpolar
3. Pick based on polarity or if you see an HF, NH or OH
you will know.
1. HCl
2. NH3
3. CH3Cl
4. SiF4
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