Covalent Bonding
8.1 Molecules & Molecular
Compounds
• Molecule: a neutral group of atoms joined
by covalent bonds
• Diatomic Molecule: two atoms joined by a
covalent bond
• Examples: H2, Cl2, O2, NO, CO
– Diatomic elements: Dr. Brinclhof
• Molecular Compounds: Compounds
composed of molecules (covalent bonds)
Comparison of Molecular & Ionic
Compounds
Molecular
Ionic
Bonding
Covalent
Ionic
Melting point
Lower
Higher
Electrolyte
Weak or non
Strong
Physical state
@ room temp
(s), (l), (g)
(s)
Molecular Formulas
• Show number & type of atoms in a
molecule
• CH4, H2S
• HNO3
• C6H6
• C3H7OH
• (NH4)3PO4
Structural Formulas
• Show the arrangement of atoms in a molecule
8.2 Nature of Covalent Bonding
• Octet rule is a guide
• Electrons are shared to form a covalent
bond
Formation of a Single Covalent
Bond
• Formed when two atoms share one pair of
electrons
Why do some elements form
diatomic molecules?
Single Covalent Bonds
The hydrogen and oxygen atoms attain noble-gas
configurations by sharing electrons.
Ammonia, NH3
Drawing Electron Dot (Lewis)
Structures
Lewis structure is a type of structural formula that
depicts all the valence electrons in the
molecule or ion
See Tutorial
1. Determine the total # ve
2. Connect atoms in such a way that all have a
noble gas configuration (octet rule)
3. Carbon is often a central atom
4. Check
Draw Lewis Structures for these
Molecular Compounds
•
•
•
•
•
•
HCl
Cl2
I2
H2O2
PCl3
CH4
hydrogen chloride
chlorine
iodine
hydrogen peroxide
phosphorous trichloride
methane
Single, Double and Triple
Covalent Bonds
• Sometimes atoms share more than one
pair of ve’s
• A bond that involves on shared pair of e-s
is a single covalent bond
• Two shared pairs of electrons is a double
covalent bond.
• Three shared pairs of electrons is a triple
covalent bond.
Acetylene
•
•
•
1.
2.
3.
4.
A gas used in cutting steel
Molecular formula is C2H2
Draw the Lewis structure for acetylene
Connect the atoms
Calculate ve’s
Form single covalent bonds between atoms
Complete octets until remainder of ve’s are
used
5. Form double or triple bonds if needed to
complete octets.
Polyatomic Ions
• Same process except…
• Add or subtract e-s to account for the
charge of the ion, for example
• [NH4]+
• [SO4]2-
Coordinate Covalent Bonds
• Bonds in which one of the shared pair comes
completely from one of the bonding atoms
• Carbon Monoxide
Bond Energies
•
•
•
•
Energy required to break a chemical bond
Energy released when a bond is formed
Is a measure of the strength of the bond
Large bond energies = strong bonds
Type of bond
Bond Energy
(kJ/mol)
C─C
347
C=C
657
C≡C
908
Resonance Structures
• A resonance structure is 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.
• Actual bonding is a hybrid of all the
possible resonance structures
Ozone
•
•
•
•
•
Is an allotropic form of oxygen
Molecular formula is O3
Is a pollutant (smog)
Protects earth by absorbing UV radiation
Draw the resonant Lewis structures for
ozone
Nitrogen Dioxide
•
•
•
•
•
Formed by lightning strikes
Molecular formula NO2
Also a pollutant in automobile exhaust
Draw the resonance structures for NO2
Why is this an exception to the octet rule?
Exceptions to Octet Rule
• When there is an odd number of ve,
NO2
• Less than an octet:
– Boron
BF3
• More than an octet:
– Phosphorous
PCl5
– Sulfur
SF6
– Unfilled d-shells accept additional
electrons, creating an “expanded” octet
8.3 Bonding Theories
• Molecular orbitals
• When covalent bonds form, atomic orbitals
merge to form molecular orbitals
Sigma and Pi Bonds
• Sigma bonds result atomic orbitals merge
along the axis between nuclei (internuclear
axis)
• Pi bonds result when atomic orbitals
merge to around the internuclear axis
Sigma Bonds
σ bonds are present in single covalent bonds.
Pi Bonds
π bonds are present in double and triple
covalent bonds
Sigma and Pi Bonds
C2 H 2
VSEPR Theory
• Valence Shell Electron Pair Repulsion
Theory
• The big idea:
• Because covalent bonds and non-bonding
pairs of electrons are areas of negative
charge, they repel one another
• Covalent bonds and non-bonding
electrons are called “electron domains”
VSEPR Predicts the shape of small
molecules
According to VSEPR theory, the repulsion
between electron pairs causes molecular
shapes to adjust so that the valence-electron
pairs stay as far apart as possible.
How to predict the shape of the following
molecules:
1. Draw the Lewis structure
2. Count the electron domains
3. Determine the geometry of the molecule (the
way the atoms are arranged
Methane, CH4
Tetrahedron, bond angles of 109.5°
Ammonia, NH3
Trigonal pyramid, 107°
Why is this not trigonal planar?
Why is the H-N-H bond angle not 109.5 °?
Water, H2O
•
•
•
•
•
Draw the Lewis structure
Determine the total domains
Determine the bonding domains
Determine the shape of the molecule
Why is water a bend molecule and not a
linear one?
Hybrid Orbitals
• When covalent bonds form, atomic orbitals
mix together to form hybrid orbitals
• Atomic orbitals involved in bonding often
contain a single unpaired electron
• When the orbitals hybridize, a pair of
electrons is shared
• These hybrid orbitals are equal in number
to the atomic orbitals which made them
Covalent Bond formation in CH4
In order for carbon’s 4 ve
to be used in bonding, one
2s2 electron is promoted to
2p.
This results in 4 unpaired
ve, which can then bond
with unpaired e’s of other
atoms.
In order to accomplish this,
the atomic orbitals of C
containing these ve
hybridize.
One s and three p orbitals
hybridize to form four
equivalent orbitals, called
sp3 orbitals
Covalent bonding in CH4
• The s (one) and p (three) orbitals in the
valence shell of C hybridize (merge) to
form four equivalent sp3 orbitals.
• They are called sp3 orbitals because they
are formed from one s orbital and three p
orbitals
Formation of Hybrid Orbitals
• http://www.mhhe.com/physsci/chemistry/e
ssentialchemistry/flash/hybrv18.swf
Hybrid Orbitals
–Hybridization Involving Single Bonds
Hybrid Orbitals
–Hybridization Involving Double Bonds
Hybrid Orbitals
–Hybridization Involving Triple Bonds
How to Determine Hybridization
about an Atom
• The principle: the number of hybrid orbitals
must equal the number of atomic orbitals
hybridized
• Count the number of covalent bonds about an
atom
• This must equal the number of hybridized
orbitals
• Beginning with s, continue to add orbitals until
the total equals the number of covalent bonds
about the atom
Hybridization Chart
# bonds
Hybridization
2
sp
3
sp2
4
sp3
5
??
6
??
Predicting Hybridization
• What hybridzation would be found about
carbon in the following molecules?
• HC≡CH
• sp
• H2C=CH2
• sp2
• H3C-CH3
• sp3
8.4 Polar Bonds and Molecules
• Electrons in a covalent bond are attracted to
the nuclei of both atoms. Why?
Unequal Sharing of Bonding Electrons
• When covalently bonded to another atom,
some atoms attract electrons more
strongly than others
• These atoms have greater
“electronegativity”
• When bonded atoms differ in
electronegativity, they do not share the
bonding electrons equally
Bonding Electrons in HCl
• Bonding e’s spend
more time near Cl
than H
• What does this imply
about Cl?
• What does this imply
about the distribution
of electrical charge in
HCl?
Polar Covalent Bonds
• When bonded atoms are sufficiently different in
electronegativity, the bond develops negative (-)
and positive (+) ends
• Why? Because the bonding e’s spend more
time around the more electronegative element
• This unequal distribution of (-) charge is called a
dipole
• The bond is called a polar covalent bond
Bond Character
• Describes the type of charge distribution in a
chemical bond
• Based upon differences in electronegativity
Differences in Electronegativity
and Bond Character
Polar Molecules
• Molecules containing polar bonds may
have an net dipole
• The molecule may have a (+) and (-) side
• Depends upon two factors
– Presence of polar bonds
– Geometry (shape) of molecule
Intermolecular Forces
• Types of intermolecular forces account for differences
between ionic and molecular substances.
Polar Molecules
Intermolecular Forces of Attraction
• Not chemical bonds
• Much weaker than covalent or ionic bonds
• Van der Waals Forces
dipole-dipole interactions
London dispersion forces
• Hydrogen Bonds
very important
Hydrogen Bonds
• Hydrogen bonds
– Attraction between a hydrogen covalently
bonded to a very electronegative atom to an
unshared electron pair of another
electronegative atom
– often involve different molecules
• Hydrogen bonding accounts for the
unusual properties of water.
Hydrogen Bonding in Water