CHAPTER 8 COVALENT BONDS
8.1 Molecular Compounds
8.2 The Nature of Covalent Bonding
8.3 Bonding Theories
8.4 Polar Bonds and Molecules
8.1 Molecular Compounds
Key Concepts:
1. How are melting points and boiling points of molecular
compounds different from ionic compounds?
2. What information does a molecular formula provide?
Molecules and Molecular Compounds
1. Covalent bond – occurs when two or more atoms share
valence electrons.
2. Molecule – is a neutral group of atoms joined together by
covalent bonds.
3. Diatomic molecule – is a molecule consisting of two atoms.
4. Compound – a substance that contains two or more elements
chemically combined in a fixed proportion
5. Molecular compounds – a compound composed of molecules
a. Have low melting points
b. Have low boiling points
c. Most are gas or liquid at room temperature
d. Composed of two or more non-metals
6. Using page 214 illustrate some differences between ionic and
covalent compounds.
Molecular Formula:
1. Molecular Formula – the chemical formula of a molecular
compound
a. Describes how many of each atom a molecule contains
b. Subscripts are used after the element’s symbol to indicate
the number of atoms of each element in the molecule.
c. Reflects the actual number of atoms in each molecule and
are not necessarily the lowest whole-number ratios.
d. Can describe molecules consisting of one element.
e. Does not tell you about the molecule’s structure
2. Using page 215 state the molecular formula for Ammonia and
describe the types of diagrams and models used to represent
Ammonia.
8.2 The Nature of Covalent Bonding
Key Concepts
1. How does electron sharing occur in forming covalent bonds?
2. How do electron dot structures represent shared electrons?
3. How do atoms form double or triple covalent bonds?
4. How are coordinate covalent bonds different other covalent
bonds?
5. How is the strength of a covalent bond related to its bond
dissociation energy?
6. How are oxygen atoms bonded in ozone?
7. What are some exceptions to the octet rule?
The Octet Rule in Covalent Bonding
1. In forming covalent bonds, electron sharing
usually occurs so that atoms attain the
electron configurations of noble gases.
2. That is to say the valence electrons arrange
themselves so that each atom sees an octet.
3. Hydrogen has a noble gas configuration with 2
electrons
4. Groups four to seven are likely to form
covalent bonds
Single Covalent Bonds
1. Single Covalent Bond – Two atoms held together by sharing a
pair of electrons.
2. Hydrogen is an example.
3. An electron dot structure can be used to show the shared
pair of electrons of the covalent bond.
4. Using page 218 use electron dots to combine two Fluorine
atoms then show the electron configuration for each atom.
5. Structural Formula – represents the covalent bonds by using
dashes, each dash represents one electron pair.
6. Unshared Pair – are electrons not shared between atoms –
also called lone pair, nonbonding pair.
7. Draw the electron dot structure for ammonia (NH3) show the
unpaired bonds and the shared pairs properly.
8. Now draw the structure for methane
9. Draw the electron configuration for Carbon then using
p220 of the text explain why Carbon usually forms four
bonds.
Double and Triple Covalent Bonds
1. Atoms sometime bond by sharing more than one pair of
electrons.
2. Double Covalent Bond – Shares two pair of electrons
3. Triple Covalent Bond – Shares three pairs of electrons
4. Try showing bonding Carbon Dioxide
Coordinate Covalent Bonds
1. Is a covalent bond in which one atom contributes both bonding
electrons
2. In a structural formula, you can show coordinate covalent bonds as
arrows that point from the atom donating the pair of electrons.
3. Once formed, a coordinate covalent bond is like any other covalent
bond.
4. Most polyatomic cations and anions contain both covalent and
coordinate covalent bonds.
5. Compounds containing polyatomic ions include both ionic and
covalent bonding.
6. Polyatomic ions have charge in order to satisfy the octet rule for
each atom present in the group.
7. Show the coordinate covalent bond of Carbon Monoxide.
8. Show the formation of the Ammonium ion.
9. Show the formation of Sulfate.
10. Using page 224 of your text, show the chemical and structural
formula for the following Molecular Compounds.
a. Nitrous Oxide
b. Sulfur Trioxide
c. Hydrogen Fluoride
d. Nitric Oxide
e. Hydrogen Peroxide
f. Nitrogen Dioxide
g. Hydrogen Cyanide
h. Hydrogen Chloride
i. Sulfur Dioxide
11. The electron dot structure for a neutral molecule contains the
same number of electrons as the total number of valence
electrons in the combining atoms.
12. The negative charge of a polyatomic ion shows the number of
electrons in addition to the valence electrons.
13. Because a negatively charged polyatomic ion is part of an ionic
compound, the positive charge of the cation of the compound
balances these additional electrons.
Bond Dissociation Energies
1. The energy required to break the bond between two
covalently bonded atoms.
2. Usually expressed as the energy needed to break one mole
of bonds.
3. A large bond dissociation energy corresponds to a strong
covalent bond.
4. High dissociation energies tend to create very stable
compounds that tend to be chemically unreactive.
5. Units are measured in kJ/mo1
6. A mol is a chemical quantity of an element or compound in
which there are 6.02x1023 atoms or molecules present.
Link
Resonance
1. A structure that occurs when it is possible to draw two or more
valid electron dot structures that have the same number of
electron pairs for a molecule or ion.
2. Although no back-and-forth changes occur, double –headed
arrows are used to connect resonance structures.
3. Show the structural formation of ozone.
Exceptions to the Octet Rule
1. The octet rule cannot be satisfied in molecules whose total
number of valence electrons is an odd number. There are also
molecules in which an atom has fewer, or more, than a complete
octet of valence electrons.
2. Draw two resonance structures for Nitrogen Dioxide
3. Other odd number electron molecules are Chlorine Dioxide and
Nitric Oxide.
4. Several molecules with an even number of electrons, such as
some compounds of Boron, also fail to follow the octet rule.
5. Draw the structure for Boron Trifluoride and show the
significance of it reacting with ammonia.
6. A few atoms, Phosphorus and Sulfur, can have ten or twelve
electrons instead of eight
7. Draw the structure for Phosphorus Pentachloride and Sulfur
Hexafluoride
Exceptions to the Octet Rule
There are three classes of exceptions to the octet rule
1) Molecules with an odd number of electrons;
2) Molecules in which one atom has less than an octet;
3) Molecules in which one atom has more than an octet.
Odd Number of Electrons
Few examples. Generally molecules such as ClO2, NO,
and NO2 have an odd number of electrons.
Less than an Octet Less
Molecules with less than an octet are typical for
compounds of Groups 1A, 2A, and 3A.
Most typical example is BF3.
More electrons than an Octet
This is the largest class of exceptions. Atoms from
the 3rd period onwards can accommodate
more than an octet. Beyond the third period, the dorbitals are low enough in energy to participate in bonding and
accept the extra electron density.
8.3 Bonding Theories
1. How are atomic and molecular orbitals related?
2. How does VESPR theory help predict the shapes of molecules?
3. In what ways is orbital hybridization useful in describing molecules?
Hybridization of atomic orbitals
Quantum mechanical approaches by combining the wave
functions to give new wavefunctions are called
hybridization of atomic orbitals. Hybridization has a
sound mathematical fundation, but it is a little too
complicated to show the details here. We can say that
an imaginary mixing process converts a set of atomic
orbitals to a new set of hybrid orbitals that are a
combination of the two overlaping orbitals.
At this level, we consider the following hybrid orbitals:
sp
sp2
sp3
Molecular Orbitals
1. Molecular orbitals are created when two atoms combine by the
overlap of each atoms atomic orbital creating an orbital that
applies to the entire molecule.
2. Each atomic orbital is full when it contains two electrons.
3. Bonding Orbitals – in covalent bonds two electrons are also
required to fill a molecular orbital.
4. Sigma Bonds – are created when two atomic orbitals combine to
form a molecular orbital that is symmetrical around the axis
connecting two atomic nuclei
hybridization of single bonds
Two examples of sigma bonds
Are H2 and F2
5. Pi Bonds – Are created by the side by side overlap of p orbitals
the bonding electrons are most likely to be found in sausage-shaped
regions above and below the bond axis. Hybridization of double bonds
Atomic orbitals of pi bonding
overlap less than in sigma bonding
therefore, pi bonds tend
to be weaker than sigma bonds.
Example = Ethene
Simply – sigma are single bonds
pi are double bonds
6. Hybridization of triple bonds –
Example = acetylene
double bonds =1 pi bond + 1 sigma.
triple bond=2 pi bonds +1 sigma
single bonds = 1 sigma.
VSEPR – Valence Shell Electron-Pair Repulsion Theory
The repulsion theory between electron pairs causes molecular
shapes to adjust so that the valence –electron pairs stay as far apart
as possible creating three dimensional structures
Therefore, VSEPR diagrams are characterized by the number of
lone pair electrons (unshared electron pairs) and the angles between
the shared pairs of electrons
The AXE system
American* general chemistry textbooks adopt the excellent
AXmEn system, where A is the central atom, m the number
of ligands X, and n the number of nonbonded lone-pairs of
electrons, E, about the central atom.
methane, CH4, is AX4
ammonia, H3N:, is AX3E1
water, H2O, is AX2E2
Note that different AXmEn designations can give rise to the
same overall geometry or shape:
AX2E1 and AX2E2 both give rise to bent or angular geometries
AX2 and AX2E3 both give rise to linear geometries
The AXE system gives rise to a pattern, from which
the various atomic geometric shapes can be
determined/assigned:
A Couple of More Advanced Examples:
a
r
Hybrid Orbitals
-Provides information about both
T 3 smolecular bonding and molecular
r just
p deals with molecular shape.
shape unlike VSEPR theory that
Valence Electron
Pair Geometry
Linear
i
g
o
Numbernof
Orbitals
a
2 l
2
Hybrid Orbitals
sp
Trigonal Planar
3
sp2
Tetrahedral
4 P
sp3
5
sp3d
Trigonal
Bipyramidal
Octahedral
l
a
n
6
a
r
T4 s
e p
t 3
r
a
h
e
sp3d2