Chapter 8
Covalent Bonding
8.1
Molecular Compounds
Molecules and Molecular Compounds
Covalent Bonds
Define: A bond formed when two atoms share a pair
of electrons
Molecule
Define: A neutral group of atoms joined together by
covalent bonds.
Diatomic Molecule:
Define: A molecule consisting of two like elements.
Ex: O2
Molecular Compound:
·
a compound composed of molecules. (nonmetals
only, including H)
· tend to have relatively low boiling/melting
points ( 300o)
Molecular Formulas
· is the chemical formula of a molecular compound
· depicts the number of atoms of each atom
present in a molecule
8.2 The Nature of Covalent Bonding
The Octet Rule in Covalent Bonding
In forming covalent bonds, the octet rule is still
obeyed. Electron sharing usually occurs so that atoms
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attain the electron configuration of the noble gases.
Single Covalent Bonds
Define: Two atoms held together by a sharing of a
single pair of electrons.
The simplest compound with covalent bonds: H2
Orbital Picture:
Lewis Dot Picture:
Each hydrogen gets the stable electron configuration
of helium.
Covalent bonds occur between elements which want
electrons:
Another example: F
A pair of shared electrons may also be represented by
a dash ( ─ ).
These are called structural formulas.
Another example: H2O
2
Nitrogen trichloride:
NCl3
Double and Triple Covalent Bonds
Example: the oxygen molecule, O2
In order to write a Lewis Structure, following the
octet rule, oxygen must share 4 electrons.
Define:
Double Covalent Bond a bond which involves the
sharing of two pairs of electrons
also:
Triple Covalent Bond a bond which involves the
sharing of three pairs of electrons
Example: the nitrogen molecule, N2
Coordinate Covalent Bonds
· a special type of covalent bond
Defined:
A covalent bond in which one atom contributes both
bonding electrons; the shared pair comes from one of
the bonding atoms. Once formed, it acts as a normal
covalent bond.
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Example: carbon monoxide, CO
To indicate a coordinate covalent bond, an arrow is
used in the structural formula, showing which atom
contributed both electrons.
Lewis Structures for Polyatomic Ions
Polyatomic ions contain covalent bonds within
themselves, even though they are part of an ionic
compound
Example NH4+
Many polyatomic ions are negatively charged meaning
electrons must be added
Example: OH-
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Example: SO32-
Bond Dissociation Energies
Atoms form bonds to become stable (lower energy)
Energy is required to break atoms apart
Bond Dissociation Energy:
Defined:
The energy required to break a chemical
bond
Example:
H-H
+
435kJ

H· +
H·
See Table 15.4
Note: Double bonds have higher bond energies than
single bonds
C─O 356k1
C═O 736kJ
CO 1074kJ
Bond Energies may be used to predict ΔH of reactions
Example:
If H─Cl bond energy = 431 kJ,
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calculate AH for: H2 + Cl2  2HCl
Resonance
Defined:
A phenomenon that occurs when two or more equally
valid electron dot structures can be written for a
molecule or ion.
Example: ozone, O3
·
·
·
both structures are different and valid
the bonding changes while positions of the atoms do
not
the actual structure is considered to be an average
or hybrid of both
Generally, the more resonance forms that can be drawn,
the more stable the molecule or ion.
Exceptions to the Octet Rule
Sometimes, molecules exist which do not follow the
octet rule. If there is an odd number of electrons or
more/less than 8 valence electrons, the octet rule
cannot be followed.
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Example: NO2
Note: Both structures have an unpaired electron
Substances with one or more unpaired electrons are
said to be paramagnetic
· are attracted by an external magnetic field
Substances with all paired electrons are said to be
diamagnetic
· not attracted by an external magnetic field
Oxygen, O2 is found to be paramagnetic
most likely, oxygen is a resonance hybrid of:
Other exceptions to the Octet Rule:
BF3
PCl5
BeCl2
8.3 Bonding Theories
VSEPR Theory
SF6
VSEPR Theory is used to determine the shapes of
covalent molecules from their Lewis Structures.
Consider 5 cases:
Case #1 AB2
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Dot formula:
Case #2 AB3
Dot formula:
Picture:
Case #3 AB4
Dot formula:
Picture:
Case #4 AB2
Dot formula:
Picture:
Case #5 AB3
Dot formula:
Picture:
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8.4 Polar Bonds and Molecules
Bond Polarity
Not all covalent bonds are the same
· the character of the bond depends on the atoms
which are joined
· the character of the molecule depends on the bonds
When atoms share electrons, they may not share them
equally
· some atoms have a stronger pull on the electrons
Two types of Covalent Bonds
 Nonpolar Covalent Bond: a covalent bond in which
the electrons are equally shared
Examples: H2, O2, N2...
· atoms have an equal pull on the electrons
 Polar Covalent Bond: a covalent bond in which the
electrons are unequally shared
· one atom has a greater attraction for
electrons
Electronegativity
· the measure of the attraction an atom has for
electrons within a bond (see table )
Example:
HCl
To indicate this:
Polar Molecules
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Polar molecules: one end is more electronegative than
the other.
When polar molecules are placed between oppositely
charged plates, they tend to orient themselves with
respect to the positive and negative plates. P.239
Fig. 8.24
H2O
Other molecules:
BeH2
PCl3
CH4
The amount of polarity depends on the differences in
electronegativities.
Four levels of Polarity
Δ EN
Type of Bond
Example
0.0 - 0.4
Covalent (nonpolar)
H─H
0.4 - 1.0
Covalent(moderately
polar)
H─Cl
1.0 - 2.0
Covalent (very polar)
H─F
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2.0
Ionic
NaCl
Attractions between Molecules-Intermolecular
Attractions
Ionic substances are held together by strong ionic
bonds
· most are solids
Covalent (molecular) substances are individual
molecules
· many are gases
· some are liquids or solids
Molecular substances are held together by
Intermolecular Forces (forces between molecules).
These forces are weaker than bonds.
Two Types:
 van der Waals Forces
a. Dipole interactions - the attraction
between polar molecules
b. Dispersion Forces - due to interactions of
electrons and nuclei of neighboring
molecules
Example:
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Intermolecular forces help determine melting points,
boiling points, solubilities, etc.
 Hydrogen "Bonds"
· are the result of attractive forces between
a hydrogen on one molecule and a highly
electronegative atom (N, O, F) on another
Example: H2O
Intermolecular Attractions and Molecular Properties
Molecular substances:
· most are gases, liquids, or low-melting solids
(below 300 oC)
· do not conduct electricity
· solubility in water varies (depends on
intermolecular forces)
network solids: solids that are covalently bonded and
form a network: high BP/MP (diamond, SiC)
See Table 8.4 p. 244
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HW:
1,2,3,4,5,6,7,8,9,10,11,12,15,16,17,18,19,21,22,24,30,31,32,34,36,37,38,3
9,40,41,43,44,45,46,49,50,51,52,54,57,58,59,60,61,63,64,65,68,69,70,73,
75,76,79,82
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