COVALENT BONDING
Chapter 8
Section Overview
• 8.1: Molecular Compounds
• 8.2: The Nature of Covalent Compounds
• 8.3: Bonding Theories
• 8.4: Polar Bonds and Molecules
MOLECULAR
COMPOUNDS
Section 8.1
Molecules and Molecular Compounds
• The atoms held together by sharing electrons are joined
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by a covalent bond which is like a “tug of war” for the
electrons.
A molecule is a neutral group of atoms joined together by
covalent bonds (ex. Oxygen).
A diatomic molecule is a molecule consisting of two atoms
(ex. Oxygen).
A compound composed of molecules is called a molecular
compound (ex. Water).
Molecular compounds tend to have relatively lower
melting and boiling points than ionic compounds and are
gases or liquid at room temperature.
Molecular Formulas
• A molecular formula is a chemical formula of a molecular
formula.
• A molecular formula shows how many atoms of each
element a molecule contains (ex. water = H20).
• However, a molecular formula doesn’t tell you about a
molecule’s structure or how the atoms are arranged in the
compound.
THE NATURE OF
COVALENT
COMPOUNDS
Section 8.2
The Octet Rule in Covalent Bonding
• In forming covalent bonds, electron sharing usually occurs
so that atoms attain the electron configuration of noble
gases (eight valence electrons).
• Groups 4A, 5A,6A, and 7A are likely to form covalent
bonds.
Single Covalent Bonds
• Two toms held together by sharing a pair of electrons are
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joined by a single covalent bond.
An electron dot structure can represent the shared pair of
electrons of the covalent bond by two dots.
A structural formula represents covalent bonds by dashes
and shows the arrangement of atoms.
Examples:
Hydrogen Gas =
Fluorine Molecule =
Single Covalent Bonds
• A pair of valence electrons that is not shared between
atoms is called an unshared pair or a lone pair.
• Water, for example, contains three atoms but only two
covalent bonds.
• Other examples are ammonia and methane.
Single Covalent Bonds
• Example Problem: Draw an electron dot structure for
hydrochloric acid (HCl)
Single Covalent Bonds
• Example Problem: Draw an electron dot structure for
hydrochloric acid (HCl)
• Solution:
First, draw separate electron dot structures.
We see that hydrogen has one valence and chlorine has
seven. So, hydrogen will share its one with chlorine to form
an octet.
Double and Triple Covalent Bonds
• Atoms form double or triple covalent bonds if they can
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attain a noble gas structure by sharing two pairs or three
pairs of electrons.
A bond that involves two shared pairs of electrons is a
double covalent bond.
A bond that involves three shared pairs of electrons is a
triple covalent bond.
An example of an element whose molecules contain triple
bonds in nitrogen which has five valence electrons.
Each nitrogen atom in the nitrogen molecule must gain
three electrons to obey to octet rule.
Another example is carbon dioxide, which has two double
bonds.
Double and Triple Covalent Bonds
Coordinate Covalent Bonds
• A coordinate covalent bond is a covalent bond in which
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one atom contributes both bonding electrons.
An example is carbon monoxide.
A carbon atom needs to gain four electrons and an
oxygen atom needs two electrons.
A double covalent bond is formed when these atoms
come together.
However, carbon is unstable in that configuration.
So, oxygen donates a pair of unshared electrons.
Coordinate Covalent Bonds
• Other examples can be found in polyatomic ions.
• A polyatomic ion is a tightly bound group of atoms that
has a positive or negative charge and behaves as a unit.
• An ammonium ion is an example.
• The ammonium ion forms when a positively charged
hydrogen ion attaches to the unshared electron pair of an
ammonia molecule.
Coordinate Covalent Bonds
• Example Problem: Draw the electron dot structure for the
hydronium ion (H3O+).
Coordinate Covalent Bonds
• Example Problem: Draw the electron dot structure for the
hydronium ion (H3O+).
• Solution:
Draw the electron dot structure of a water molecule.
Draw the electron dot structure for a single hydrogen ion.
Put them together into one structure.
Bond Dissociation Energies
• The energy required to break the bond between two
covalently bonded atoms is known as the bond
dissociation energy.
• This is expressed as the energy needed to break one
mole of the bonds.
• A large bond dissociation energy corresponds to a strong
covalent bond and thus a stable compound.
Resonance
• A resonance structure is a structure that occurs when it is
possible to draw two or more valid electron dot structure that
have the same number of electron pairs for a molecule or ion.
• Ozone is an example
• Earlier scientists thought that the electron pairs rapidly flipped
back and forth between the different electron dot structures so
a double headed arrow is used.
• However, today, we know that they do not actually resonate
back and forth.
• The actual bonding of oxygen atoms in ozone is a hybrid, or
mixture, of the extremes represented by the resonance forms.
Exceptions to the Octet Rule
• The octet rule cannot be satisfied in molecules whose
total number of valence electrons is an odd number.
• There are also molecules which an atom has fewer, or
more, than a complete octet of valence electrons.
• Some examples include nitrogen dioxide (NO2), chlorine
dioxide (ClO2), nitric oxide (NO), boron triflouride (BF3),
phosphorus trichloride (PCl3), and phosphorus
pentachloride (PCl5).
BONDING THEORIES
Section 8.3
Molecular Orbitals
• This model assumes that when two atoms combine, their
atomic orbitals overlap and produce molecular orbitals, or
orbitals that apply to the entire molecule.
• Just as an atomic orbital belongs to a particular atom, a
molecular orbital belongs to a molecule as a whole.
• Two electrons are required to fill a molecular orbital.
• A molecular orbital that can be occupied by two electrons
of a covalent bond is called a bonding orbital.
Molecular Orbitals
• Sigma Bonds: When two atomic orbitals combine to form
a molecular orbital that is symmetrical around the axis
connecting the two atomic nuclei.
• Pi Bonds: The bonding electrons are most likely to be
found in sausage-shaped region above and below the
bond axis of bonded atoms.
VSPER Theory
• According to the VSPER theory, the repulsion between
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electrons pairs causes molecular shapes to adjust so that
the valence-electron pairs stay as far apart as possible
which affects their shape.
Unpaired electrons are also important in determining
shape.
The unshared pair strongly repels the bonding pairs which
pushes them closer together.
For example, in methane, all electrons are shared and the
angle between the hydrogens and the central carbon is
109.5.
In ammonia, one of the electron pairs is unshared and the
angles between the hydrogen and the central nitrogen is
107.
VSPER Theory
VSPER Theory
Hybrid Orbitals
• Orbital hybridization provides information on both
molecular bonding and shape.
• In hybridization, several atomic orbitals mix to form the
same total number of equivalent hybrid orbitals.
• Hybridization single bonds: Carbon’s outer electron
configuration is 2s22p2, but one 2s is promoted to 2p to
form 2s12p3, allowing it to bond to four hydrogens in
methane. These orbitals mix together to form four sp3
orbitals. In methane, each of the four sp3 hybrid orbitals
overlaps with a 1s orbital of hydrogen.
Hybrid Orbitals
• Hybridization double bonds: Ethene has one carbon-
carbon double bond and four carbon-carbon single bonds.
In ethene, sp2 hybrid orbitals form from the combination of
one 2s and two 2p atomic orbitals of carbon. Two sp2
hybrid orbitals of each carbon from sigma-bonding with
the four available hydrogen 1s molecules. The third sp2
orbitals of each of the two carbons overlap to form a
carbon-carbon sigma-bonding orbital. The non-hybridized
2p carbon orbitals overlap side-by-side to form a pibonding orbital.
Hybrid Orbitals
• Hybridization triple bonds: Ethyne (aka acetylene) is a
linear molecule. The best hybrid orbital description is
obtained if a 2s atomic orbital of carbon mixes with only
one of three 2p atomic orbitals. The result is two sp hybrid
orbitals for each carbon. One sp hybrid orbital from each
carbon overlaps with a 1s orbital or hydrogen to form a
sigma bond. The other sp hybrid orbital of each carbon
overlaps to form a carbon-carbon sigma bond. The two p
atomic orbitals from each carbon also overlap.
Hybrid Orbitals: Methane
Hybrid Orbitals: Ethene
Hybrid Orbitals: Ethyne
POLAR BONDS AND
MOLECULES
Section 8.4
Bond Polarity
• The bonding pairs of electrons in covalent bonds are
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pulled, kind of like in a tug of war, between the nuclei of
the atoms sharing the electrons.
When the atoms in the bond pull equally, the bonding
electrons are shared equally, and the bond is called a
nonpolar covalent bond.
A polar covalent bond, or just polar bond, is a covalent
bond between atoms in which the electrons are shared
unequally.
The more electronegative atom attracts electrons more
strongly and gains a slightly negative charge.
The less electronegative atom has a slightly positive
charge.
Bond Polarity
• For example, HCl is made of hydrogen and chlorine, to
form hydrogen chloride.
• Hydrogen has an electronegativity of 2.1 and chlorine has
an electronegativity of 3.0.
• These values are significantly different, so the covalent
bond in hydrogen chloride is polar.
• So, chlorine gains a slight negative charge and hydrogen
gains a slight positive charge.
Bond Polarity
Bond Polarity
• Example Problem: Which type of bond will form between
each of the following pairs of atoms?
a. N and H
b. F and F
c. Ca and Cl
d. Al and Cl
Bond Polarity
• Example Problem: Which type of bond will form between
each of the following pairs of atoms?
a. N and H
• Solution:
Find individual electronegativity values
Determine difference
Use chart to determine bond type
N (3.0) and H (2.1)
Difference = 0.9
Bond = covalent
Bond Polarity
• Example Problem: Which type of bond will form between
each of the following pairs of atoms?
b. F and F
c. Ca and Cl
d. Al and Cl
• Solution:
b. F (4.0) and F (4.0); 0.0; nonpolar covalent
c. Ca (1.0) and Cl (3.0); 2.0; ionic
d. Al (1.5), Cl (3.0); 1.5; polar covalent
Polar Molecules
• In a polar molecule, one end of the molecule is slightly
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more negative and the other end is slightly positive.
A molecule that has two poles is called a dipolar
molecule, or dipole.
When polar molecules are placed between oppositely
charged plates, they tend to become oriented with respect
to the positive and negative poles.
The effect of polar bonds on the polarity of an entire
molecule depends on the shape and the orientation of the
polar bonds.
For example, CO2 has two polar bonds, but CO2 is linear
and the bonds are in opposite directions making it overall
nonpolar. Water also has two polar bonds, but it is bent,
making it polar.
Attractions Between Molecules
• Intermolecular attractions are weaker than either ionic or
covalent bonds.
• However, these attractions are responsible for
determining whether a molecular compounds is a gas,
liquid, or solid at a given temperature.
• Two types of attractions are van der Waals Forces and
hydrogen bonds.
Attractions Between Molecules
• Van der Waals Forces: The two weakest attractions
between molecules.
• Dipole interactions occur when polar molecules are
attracted to one another. The attraction involved occurs
between the oppositely charged regions of polar
molecules. The slightly negative region is attracted to the
slightly positive region.
• Dispersion forces are caused by the motion of electrons
and are the weakest of all intermolecular interactions.
When the moving electrons happen to be momentarily
more on the side of a molecule closest to a neighboring
molecule, their electric force influences the neighboring
molecule’s electrons to be momentarily more on the
opposite side.
Attractions Between Molecules
• Hydrogen Bonds: Attractive forces in which a hydrogen
covalently bonded to a very electronegative atom is also
weakly bonded to an unshared electron pair of another
electronegative atom. This other atom can be in the same
molecule or in a nearby molecule.
• Hydrogen bonds are present in water.
• One hydrogen from a water
molecule, is attracted to an
oxygen of another water molecule.
Intermolecular Attractions and Molecular
Properties
• At room temperature, some compounds are gases while
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others are liquids or solids.
The physical properties of a compound depend on the
type of bonding it displays.
A greater range occurs among covalent compounds due
to the varying intermolecular attractions.
In most solids formed from molecules, only the weak
interactions between molecules need to be broken in
order for it to melt.
However, there are network solids, in which all of the
atoms are covalently bonded to each other and can only
melt when all covalent bonds throughout are broken (ex.
Diamond and silicon carbide).
Characteristics of Ionic and Covalent
Compounds
Characteristic
Ionic Compound
Covalent Compound
Representative unit
Formula unit
Molecule
Bond formation
Transfer of one or
more electrons
between atoms
Sharing of electron
pairs between atoms
Type of elements
Metallic and
nonmetallic
Nonmetallic
Physical state
Solid
Solid, liquid, or gas
Melting point
High
Low
Solubility in water
Usually high
High to low
Electrical
conductivity
Good conductor
Poor to
nonconducting