Chapter 9
Covalent Bonding
I. The Covalent Bond
A. Why do atoms bond?
• When two atoms need to gain electrons,
they can share electrons to acquire a noblegas configuration
B. What is a covalent bond?
• In a covalent bond, shared electrons are
considered to be part of the complete outer
energy level of both atoms involved
• Generally occurs when elements are
relatively close to each other on the periodic
table
• A majority form between nonmetallic
elements
• A molecule is formed
when two or more
atoms bond covalently
Examples of covalent
bond – carbohydrates,
sugars, proteins, fats,
and DNA
1. Formation of a covalent bond
• There are seven elements that naturally occur
as diatomic molecules (H2, N2, O2, Fl2, Cl2,
Br2, and I2)
• Opposite charges of the protons of one atom
and electrons of another atom are attracted to
each other
• Like charges of the protons of each atom repel
each other
• As a result, the most stable arrangement of
atoms exists at the point of maximum
attraction
C. Single Covalent Bonds
• When a single pair of electrons is shared, a
single covalent bond forms
• Group 7A elements have seven valence
electrons and can form one covalent bond
with another element
• Group 6A element shave six valence electrons
and can form two covalent bond and so on.
1. The sigma bond
• Single covalent bonds are also called sigma
bonds, σ
• Occurs when the electron pair is shared in an
area centered between the two atoms
• The valence atomic orbital of one atom
overlaps or merges with the valence atomic
orbital of another atom
• Sigma bonds can
form from the
overlap between
s orbitals,
between an s and
p orbital, or
between 2 p
orbitals
D. Multiple Covalent Bond
• Some atoms attain a noble-gas configuration
by sharing more than one pair of electrons
between two atoms
• Carbon, nitrogen, oxygen, and sulfur most
often form multiple bonds
• A double covalent bond occurs when two pairs
of electrons are shared, O2
• A triple covalent bond occurs when three pairs
of electrons are shared, N2
1. The pi bond (π)
• Formed when parallel orbitals overlap to share
electrons
• Electron pairs occupy the space above and
below the line that represents where the two
atoms are joined together
• Always accompanies a sigma bond when
forming double and triple bonds
• One pi bond in a double bond
• Two pi bonds in a triple bond
E. Strength of Covalent Bonds
• Depends on how much distance separates
bonded nuclei
• Bond length is the distance between the two
nuclei at the position of maximum attraction
- Determined by atomic radius and how many
electron pairs are shared
• As the number of shared electron pairs
increases, bond length decreases
• Single bonds are weaker than double bonds
which are weaker than triple bonds
• Energy is released when a bond forms
(exothermic)
• Energy is absorbed when a bond breaks
(endothermic)
• The amount of energy required to break a
specific covalent bond is tis bond dissociation
energy
• Bond energy is directly related to bond length
II. Naming Molecules
A. Naming Binary Molecular Compounds
• Many binary molecular compounds were
discovered and given common names long
before the modern naming system was
developed
• Binary molecular compounds are composed
of two different nonmetals and do not
contain metals or ions
1. The first element in the formula is always
named first, using the entire element name
2. The second element in the formula is named
using the root of the element and adding the
suffix –ide
3. Prefixes are used to indicate the number of
atoms of each type that are present in the
compound
- The first element never uses the prefix mono-
B. Naming Acids
1. Naming binary acids
• Contains hydrogen and one other element
• Use the prefix hydro- to name the hydrogen
part of the compound
• The rest of the name consists of a form of
the root of the second element plus the
suffix –ic, followed by the word acid
2. Naming oxyacids
• Any acid that contains hydrogen and an
oxyanion
• If the oxyanion suffix is –ate, it is replaced
with the suffix –ic
• If the oxyanion suffix is –ite, it is replaced
with –ous
• They hydrogen in an oxyacid is not part of the
name
III. Molecular Structures
A. Structural Formulas
• Uses letter symbols and bonds to show
relative positions of atoms
1. Predict the location of atoms
- Hydrogen is always a terminal atom
because it can share only one pair of
electrons
2. Find the total number of electrons available
for bonding
3. Determine the number of bonding pairs by
dividing the number of electrons available for
bonding by two
4. Place one bonding pair between the central
atom and each of the terminal atoms
5. Place lone pairs around each terminal atom
bonded to the central atom to satisfy the
octet rule
6. If the central atom is not surrounded by four
electron pairs, it does not have an octet
- You must convert one or two of the lone pairs
on the terminal atoms
B. Resonance Structures
• It is possible to have more than one correct
Lewis structure when a molecule or
polyatomic ion has both a double bond and a
single bond
• Resonance is a condition that occurs when
more than one valid Lewis structure can be
written
• Resonance structures differ only in the position
of the electron pairs
• Molecules or ions behaves as if it has only one
structure
- Bond lengths are shorter than single bonds
but longer than double bonds
C. Exceptions to the Octet Rule
1. Some molecules have an odd number of
valence electrons and cannot form an octet
around each atom
2. Some molecules have less than eight valence
electrons and form a coordinate covalent bond
with an atom with a lone pair of electrons
3. Some molecules have a center atom with more
than eight valence electrons and forms an
expanded octet
IV. Molecular Shape
A. VSEPR Model
• The shape of the molecule determines
whether or not molecules can get close
enough to reach
• Based on an arrangement that minimizes
the repulsion of shared and unshared pairs
• Repulsions among electron pairs result in
atoms exist at fixed angles, or bond angles
- Supported by experimental evidence
• Lone pairs occupy a slightly larger orbital than
shared electrons
• Shared bonding orbitals are pushed together
slightly by lone pairs
1. Linear 180°
• 2 shared pairs
2. Trigonal planar 120°
• 3 shared pairs
3. Tetrahedral 109.5°
• 4 shared pairs
4. Trigonal pyramidal 107.3 °
• 3 shared pairs, 1 lone pair
5. Bent 104.5°
• 2 shared pairs, 2 lone pairs
6. Trigonal bipyramidal 90/120°
• 5 shared pairs
7. Octahedral 90°
• 6 shared pairs
C. Hybridization
• A hybrid results from combining two of the
same type of object, and it has characteristics
of both
• Atomic orbitals undergo hybridization during
bonding to form new hybrid orbitals
V. Electronegativity and Polarity
A. Electronegativity Difference and Bond
Character
• Electronegativity indicates the relative
ability of an atom to attract electrons in a
chemical bond
• Chemical bonds can be predicted using the
electronegativity difference of the elements
in the bond
• For identical atoms, with an electronegativity
difference of zero, the electrons in the bond
are equally shared, or covalent
• Chemical bonds between different atoms are
never completely ionic or covalent
- Depends on how strongly atoms are attracted
to electrons
• Unequal sharing of electrons between different
elements result in a polar covalent bond
• Large differences indicate an electron was
transferred from one atom to another,
resulting in ionic bonds
• An electronegativity difference of 1.70 is
considered 50% covalent and 50% ionic
• Ionic bonds generally have electronegativities
above 1.70
B. Polar Covalent Bonds
• In a polar covalent bond,
electrons spend more time
around one atom than
another
• Partial changes occur at
the ends of the bond, δ+
and δ• Referred to as a dipole
(two poles)
1. Molecular polarity
• Nonpolar molecules are not attracted by an
electric field
• Polar molecules tend to align with an electric
field because of their dipole
2. Polar molecule or not?
• Symmetric molecules are usually nonpolar
• Asymmetrical molecules are usually polar
3. Solubility of polar molecules
• The ability of a substance to dissolve in
another substance
• Polar molecules and ionic compounds are
usually soluble in polar substances
• Nonpolar molecules dissolve only in nonpolar
substances
C. Properties of Covalent Compounds
• Differences in properties are a result of
differences in attractive forces
• Weak forces of attractions are known as
intermolecular forces, or van der Waal forces
• In nonpolar substances, weak attractions are
called dispersion forces
• The force in polar molecules is stronger and is
called a dipole-dipole force
• A hydrogen bond is a strong intermolecular
force that is formed between a hydrogen and a
dipole end of another atom
• Physical properties such as melting point,
boiling point, hardness, etc. are determined by
these intermolecular forces