Chapter 7
Covalent Bonds and Molecular
Structure
The Covalent Bond
• Covalent bond – formed by the sharing of
electrons between two nonmetal atoms
– Forces involved in the bond
• Electrostatic attraction between proton and
electron
• Electrostatic repulsion between electron and
electron
– When will a bond form?
Strengths of Bonds
• Bond dissociation energy – the amount of
energy required to break a bond
– Energy increases as the length of the bond gets
shorter
Bond length and covalent radius.
Internuclear distance
(bond length)
Internuclear distance
(bond length)
Covalent
radius
Covalent
radius
Internuclear distance
(bond length)
Internuclear distance
(bond length)
Covalent
radius
Covalent
radius
Strengths of Bonds
– Single bonds > double bonds > triple bonds
Problem
• Arrange the following bonds in order of
increasing bond strength.
– A.
– B.
– C.
– D.
– E.
C-I < C-Br < C-Cl < C-F
C-F < C-Cl < C-Br < C-I
C-Br < C-I < C-Cl < C-F
C-I < C-Br < C-F< C-Cl
none of these orders is correct
Problem
• Select the strongest bond in the following
group.
– A.
– B.
– C.
– D.
– E.
C-S
C-O
C=C
C≡N
C-F
Electron Dot Structures
• Electron dot symbols
– Aid in understanding the formation of bonds between atomic nuclei
– Elemental symbol represents the type of element and all core electrons;
the valence electrons are represented by dots around the symbol
Electron Dot Structures
• A metal in an ionic loses its electrons to achieve
an octet or pseudo-octet (transition elements) in
its outermost shell
• A nonmetal in an ionic compound gains
electrons to achieve an octet in its outermost
shell
• Period 1 and 2 elements of a covalent
compound share enough electrons to achieve
an octet
Electron-Dot Structures
• Ionic:
• Covalent:
Electron Dot Structures
• Using electron dot symbols build
– H2, H2O, CH4, O2, N2, HCN, CO2
• Least electronegative atom is often central (except
H)
Naming Binary Molecular
Compounds
• Electronegativity – indicates how well an
elements nuclei attract the electrons in a
covalent bond
The Periodic Table and
Electronegativity
Electron Dot Structures
• Using electron dot symbols build
– H2, H2O, CH4, O2, N2, HCN, CO2
• Least electronegative atom is often central (except
H)
Electron Dot Structures
• Single bond: A covalent bond formed by
sharing one electron pair.
• Double bond: A covalent bond formed by
sharing two electron pairs.
• Triple bond: A covalent bond formed by sharing
three electron pairs.
• Single bonds are longer (weaker) than double bonds
• Double bonds are longer (weaker) than triple bonds
Electron-dot Structures
• Step 1: Count the total valence electrons.
• Step 2: Identify the central atom
- Often least electronegative
• Step 3: Place all other atoms around the central
atom
• Step 4: Draw a single bond between each
external atom and the central atom
subtracting 2 electrons for each bond
drawn from the total valence electrons.
Electron-dot Structures
• Step 5: Distribute remaining valence
electrons around the external
atoms giving the external atoms
an octet
• Step 6: If valence electrons still remain,
place them on the central atom in pairs
• Step 7: Verify that each atom has an octet
» Hydrogen needs only 2 electrons
» Boron needs only 6 electrons
Electron-Dot Structures
• Step 8: If the central atom does not have an
octet, form a multiple bond by bringing a
pair of electrons in from the external
atom
• Step 9: Calculate formal charge and minimize
the formal charge if acceptable
» Period 3 elements and greater can have expanded octets if
one is necessary to minimize formal charge
» Formal charge = # valence electrons for the atom – 1 for
every dot on the atom – 1 for every line around the atom
Problems
•
•
•
•
•
•
•
BF3
PF3
C2H6
I3+
NH4+
SO42KClO3
Electron-dot Structures and
Resonance
• How is the double bond formed in O3?
Move lone pair from
this oxygen?
O
O
O
O
O
or
O
Or from this
oxygen?
O
O
O
• The correct answer is that both are correct,
but neither is correct by itself.
Electron-Dot Structures and
Resonance
• When multiple structures can be drawn, the actual
structure is an average of all possibilities.
• The average is called a resonance hybrid. A
straight double-headed arrow indicates resonance.
O
O
O
O
O
O
Problem
•
•
•
•
S3
PO43CO32NO2
Molecular Shapes: The VSEPR
Theory
• The approximate shape of molecules
is given by Valence-Shell Electron-Pair
Repulsion (VSEPR).
Molecular Shapes: The VSEPR
Theory
Molecular
formula
Step 1
Lewis
structure
Step 2
Electron-group
arrangement
Count all e- groups around central
atom (A) – Single, Double and Triple
bonds are all counted as 1 e- group
Step 3
Bond
angles
Note lone pairs and
double bonds
Count bonding
Step 4 and nonbonding egroups separately.
Molecular
shape
(AXmEn)
Problem
• Determine the shape of the molecules for
which Lewis Structures have been
developed.
Valence Bond Theory
• If, in order for a bond to form, a pair of
electrons must be shared, then how does
C form molecules with 4 bonds?
• Valence Bond Theory – hybrid orbitals
Valence Bond Theory
Basic Principle
A covalent bond forms when the orbtials of two atoms overlap
and are occupied by a pair of electrons that have the highest
probability of being located between the nuclei.
Themes
A set of overlapping orbitals has a maximum of two electrons
that must have opposite spins.
The greater the orbital overlap, the stronger (more stable) the
bond.
The valence atomic orbitals in a molecule are different from
those in isolated atoms.
Valence Bond Theory
Key Points
The number of hybrid orbitals obtained equals the number of
atomic orbitals mixed.
The type of hybrid orbitals obtained varies with the types of
atomic orbitals mixed.
Types of Hybrid Orbitals
sp
sp2
sp3
sp3d
sp3d2
Valence Bond Theory
The conceptual steps from molecular formula to the hybrid orbitals
used in bonding.
Step 1
Molecular
formula
Step 2
Lewis
structure
Step 3
Molecular shape
and e- group
arrangement
Hybrid
orbitals
Problems
• Carbon uses ______ hybrid orbitals in
ClCN.
– A.
– B.
– C.
– D.
– E.
sp
sp2
sp3
sp3d
sp3d2
The s bonds in ethane.
both C are sp3 hybridized
s-sp3 overlaps to s bonds
sp3-sp3 overlap to form a s bond
relatively even
distribution of electron
density over all s
bonds
The s and p bonds in ethylene (C2H4)
overlap in one position - s
p overlap - p
electron density
The s and p bonds in acetylene (C2H2)
overlap in one position - s
p overlap - p
Polar Covalent Bonds:
Electronegativity
• Electronegativity – represents the ability of an
atom to attract a shared pair of electrons
• Higher the EN – the more the electrons in a
bond will be pulled toward the atom
– Most electronegative atom is F
• EN ↓ down a group
• EN↑ across a period from left to right w/ few
exceptions
Polar Covalent Bonds:
Electronegativity
Problem
• Which of the following elements is the
most electronegative?
– A.
– B.
– C.
– D.
– E.
S
Ru
Si
Te
Cs
Problem
• Arrange calcium, rubidium, sulfur, and
arsenic in order of decreasing
electronegativity.
– A.
– B.
– C.
– D.
– E.
S > As > Rb > Ca
S > As > Ca > Rb
As > S > Rb > Ca
As > S > Ca > Rb
None of these orders is correct.
Polar Covalent Bonds:
Electronegativity
• % Ionic Character: As a general rule for two
atoms in a bond, we can calculate an
electronegativity difference (∆EN ): ∆EN =
EN(Y) – EN(X) for X–Y bond.

If ∆EN < 0.5 the bond is covalent.

If ∆EN 0.5 - < 2.0 the bond is polar covalent.

If ∆EN > 2.0 the bond is ionic.
Problem
• Select the most polar bond amongst the
following.
– A.
– B.
– C.
– D.
– E.
C-O
Si-F
Cl-F
C-F
C-I
Molecular Orbital Theory
• The molecular orbital (MO) model provides a
better explanation of chemical and physical
properties than the valence bond (VB) model.
– Atomic Orbital: Probability of finding the electron
within a given region of space in an atom.
– Molecular Orbital: Probability of finding the
electron within a given region of space in a
molecule.
Molecular Orbital Theory
• Additive combination of orbitals (s) is
lower in energy than two isolated 1s orbitals
and is called a bonding molecular orbital.
Molecular Orbital Theory
• Subtractive combination of orbitals (s*) is
higher in energy than two isolated 1s
orbitals and is called an antibonding
molecular orbital.
Molecular Orbital Theory
• Molecular Orbital Diagram for H2:
Molecular Orbital Theory
• Molecular Orbital Diagrams for H2– and
He2:
Molecular Orbital Theory
• Additive and subtractive combination of p
orbitals leads to the formation of both sigma
and pi orbitals.
Molecular Orbital Theory
• Second-Row MO Energy Level Diagrams:
Molecular Orbital Theory
• Bond Order is the number of electron pairs
shared between atoms.
• Bond Order is obtained by subtracting the
number of antibonding electrons from the
number of bonding electrons and dividing by 2.