UNIT 7
Chemical Bonding & Molecular Geometry
Chemical Bonding
• Valence electrons are the electrons
in the outer shell (highest energy
level) of an atom.
• A chemical bond is a mutual electrical attraction
between the nuclei and valence electrons of
different atoms that binds the atoms together.
• A chemical bond forms when the potential
energy of the bonded atoms is less than the
potential energy of the separate atoms.
The Three Major Types of Chemical Bonding
• Metallic Bonding results from the attraction between
metal cations and surrounding delocalized electrons.
• Ionic Bonding results from the electrical attraction
between oppositely-charged ions.
• Covalent Bonding
results from the
sharing of electron
pairs between
two atoms.
Types of Bonding Summarized
Bond
Types of Atoms Type of Bond
Characteristic
Electrons
Metal to Metal
Metallic
Pooled
Metal to
Electrons
Ionic
Nonmetal
Transferred
Nonmetal to
Electrons
Covalent
Nonmetal
Shared
The Metallic Bond
• In metals, overlapping orbitals allow the outer
electrons of the atoms to roam freely
throughout the entire metal.
• These mobile electrons form
a sea of electrons around the
metal atoms, which are packed
together in a crystal lattice.
• A metallic bond results from the attraction
between metal cations and the surrounding
sea of electrons.
Properties of Metals
• The characteristics of metallic bonding give
metals their unique properties, listed below.
 electrical conductivity
 thermal (heat) conductivity
 malleability (can be
hammered into thin sheets)
 ductility (can be pulled
or extruded into wires)
 luster (shiny appearance)
Visual Concept
The Octet Rule
• When atoms bond, they tend to
gain, lose, or share electrons to
result in 8 valence electrons.
• Atoms are the most stable when
they have completely full valence
shells (like the noble gases.)
• The Octet Rule – Compounds tend to form so
that each atom has an octet (group of eight)
electrons in its highest energy level.
Visual Concept
Lewis Symbols
• Lewis Symbols
(also called electron
dot symbols) are
indicated by dots
placed around the
element’s symbol.
Only the valence
electrons are shown.
Inner-shell electrons
are not shown.
Lewis Theory and Ionic Bonding
• Lewis symbols can be used to represent the
transfer of electrons from metal atom to
nonmetal atom, resulting in ions that are
attracted to each other and therefore bond.

Li

+


F



Li+1


+ 


F





1
Ionic Bonding
• Cations (+) and anions (-)
are attracted to each
other because of their
opposite electrical charges.
• An ionic bond is a bond
that forms between
oppositely-charged ions
because of their mutual
electrical attraction.
Visual Concept
Ionic Bonding and the Crystal Lattice
• In an ionic crystal, ions minimize their potential
energy by combining in an orderly arrangement
known as a crystal lattice.
• A formula unit is the smallest repeating unit of an
ionic compound.
Sodium Chloride crystal lattice (many Na and Cl atoms)
Formula Unit = NaCl
Lattice Energy
• Lattice energy is the energy released when the
solid crystal forms.
 Lattice energy depends directly on size of
charges and inversely on distance between ions.
Lattice Energy vs. Ion Size
• larger ion = weaker attraction
= smaller lattice energy.
Lattice Energy
Metal Chloride
(kJ/mol)
LiCl
-834
NaCl
-787
KCl
-701
CsCl
-657
Lattice Energy vs. Ion Charge
• larger charge = stronger attraction
= larger lattice energy.
 Of the two factors that affect lattice
energy, ion charge is generally
more important.
Lattice Energy =
-910 kJ/mol
Lattice Energy =
-3414 kJ/mol
Properties of Ionic Compounds
• Ionic compounds have very high melting and
boiling points (generally > 300oC.)



All ionic compounds are solids at room temp.
Because the attractions between ions are
strong, breaking down the crystal requires
a lot of energy.
Lattice
Melting
Compound Energy
Point ( C)
the stronger the attraction
(kJ/mol)
KBr
-671
734
(the larger the lattice
NaCl
-787
801
energy), the higher the
SrO
-3217
2530
melting point.
CaO
-3414
2613
o
Properties of Ionic Compounds (cont’d)
• Ionic compounds are hard, but brittle.

displacing the ions from their positions results in
like charges close to each other and the repulsive
forces take over.
Properties of Ionic Compounds (cont’d)
• Ionic compounds cannot conduct electricity in the
solid phase, but as dissolved ions they can.
in NaCl(s), the
ions are stuck in
position and not
allowed to move
to the charged
rods
in NaCl(aq), the
ions are
separated and
allowed to move
to the charged
rods
Covalent Bonds
• A covalent bond is formed from shared pairs of
electrons.
• A molecule is
a neutral group
of atoms held
together by
covalent bonds.
Visual Concept
Properties of Covalent Compounds
• Molecular compounds have low melting points
and boiling points (generally < 300oC).

Melting and boiling involve breaking the attractions
between the molecules (intermolecular forces), but
not the covalent bonds within the molecule.
Properties of Covalent Compounds (cont’d)
• Covalent compounds generally do not conduct
electricity as solids or in aqueous solution (no
ions to carry a charge.)

Exception - molecular acids (and bases) are
ionized when dissolved in water and therefore
have the ability to conduct electricity.
Why Do Covalent Bonds Form?
• When two atoms form a covalent bond, their
shared electrons form overlapping orbitals.
• This gives both atoms stable noble-gas configurations.
Multiple Covalent Bonds
• In a single covalent bond, one pair of
electrons is shared between 2 atoms.
• A double bond is a covalent bond in
which two pairs of electrons are
shared between 2 atoms.
• A triple bond is a covalent bond in
which three pairs of electrons are
shared between 2 atoms.
• Multiple bonds are often found
in molecules containing carbon,
nitrogen, and oxygen.
Single Bond
Double Bond
Triple Bond
Covalent Bond Length & Strength
• In general, bonds get weaker as they get longer.
• Double bonds
are shorter and
stronger than
single bonds.
• Triple bonds are
shorter and stronger
than double bonds.
Polar Covalent Bonds
• Bonding is usually somewhere between ionic
and covalent, depending on the electronegativity
difference between the two atoms.
• In polar covalent bonds, the bonded atoms have
an unequal attraction for the shared electron.
Trends in Electronegativity
Ionic or Covalent?
• If the electronegativity difference is less than or
equal to 0.4, the bond is nonpolar covalent.
• If the electronegativity difference is 0.5 to 1.9, the
bond is polar covalent.
• If the electronegativity difference is larger than or
equal to 2.0, the bond is ionic.
0
0.4
2.0
3.3
Ionic or Covalent?
Sample Problem
Use electronegativity values to classify bonding
between bromine and the following elements:
hydrogen, cesium, and chlorine.
In each pair, which atom will be more negative?
Solution:
Bonding
between
Br and:
hydrogen
cesium
chlorine
Electroneg.
difference
2.8 – 2.1 = 0.7
2.8 – 0.7 = 2.1
3.0 – 2.8 = 0.2
More
negative
Bond type
atom
polar-covalent
bromine
ionic
bromine
nonpolar-covalent N/A
Lewis Structures
• Lewis symbols of two or more atoms can be
combined to represent molecules.
• Unpaired electrons will pair up to form a
shared pair or covalent bond.
Lewis Structures (continued)
• The pair of dots representing
the shared pair of electrons in
a covalent bond is often replaced
by a long dash.
• An unshared pair, also called a
lone pair, is a pair of electrons
that is not involved in bonding Shared pair
(covalent bond)
and that belongs exclusively
to one atom.
•
•
Lone
pair
How to Draw Lewis Structures
1. Count valence electrons to determine the total
number of electrons in the compound.
2. Write the correct skeletal structure for the molecule,
putting the least electronegative atom(s) in the
center (except H.)
3. Subtract 2 electrons per bond from your total
number of electrons. Distribute remaining electrons
as lone pairs, first on outside atoms then on the
central atom.
4. Make sure all atoms (except H) have octets. If
necessary, introduce multiple bonds.
Lewis Structures
Sample Problem A
Draw the Lewis structure of iodomethane, CH3I.
C
:
4
e
Step 1 - Count valence electrons.
3H: 3(1 e-)= 3 eStep 2 - Draw correct structure.
I:
+ 7 eStep 3 - Subtract e- used in bonding.
Total:
14
e
Distribute remaining e- as lone pairs.
Used: - 8 eStep 4 - Make sure all atoms
Remaining: 6 e-
H
(except H) have octets.
••

If not, introduce
H C I ••
multiple bonds.
••
H
Lewis Structures
Sample Problem B
Draw the Lewis structure of nitric acid, HNO3.
H:
1 eN:
5 e(In oxyacid, H outside attached to O) 3O: 3(6 e-)= 18 e+
Step 3 - Subtract e used in bonding.
Total: 24 eDistribute remaining e- as lone pairs.
Used: - 8 eStep 4 - Make sure all atoms
Remaining:
16
e
••
Step 1 - Count valence electrons.
Step 2 - Draw correct structure.
(except H) have octets.
If not, introduce
multiple bonds.
 ••
•• O ••
••
H O N O ••
••
••
Polyatomic Ions
• A charged group of covalently bonded atoms
is known as a polyatomic ion.
• Draw a Lewis structure for a polyatomic ion
with brackets around it and the charge in the
upper right corner.
hydroxide ion, OH-
ammonium ion, NH4+
Formal Charges
• during bonding, atoms may wind up with more or
less electrons in order to fulfill octets - this results in
atoms having a formal charge.
FC = valence e- - nonbonding e- - ½ bonding eExample:
••
••
••
O S O ••
••
••
left O
S
right O
FC = 6 - 4 - ½ (4) = 0
FC = 6 - 2 - ½ (6) = +1
FC = 6 - 6 - ½ (2) = -1
Formal Charges (cont’d)
• Formal charges can help distinguish between two
possible skeletal structures, according to these rules:
The sum of all the formal charges in a molecule = 0.
 In an ion, the total equals the charge on the ion.
 Small (or 0) formal charges are better than large ones.
 When formal charge cannot be avoided, negative formal
charge should reside on the most electronegative atom.
Example: Which is the correct structure for hydrogen cyanide?

H
0
C N ••
0
0
OR
H
N C ••
0
+1
-1
Resonance
• when there is more than one Lewis structure for a
molecule that differ only in the position of the
electrons, they are called resonance structures.
• the actual molecule is a combination of the
resonance forms – a resonance hybrid.
Isomers
• Not to be confused with resonance structures,
Isomers are compounds with the same molecular
formula but different structural formulas.


Changes in structure result in changes in physical
and chemical properties.
Example: butane (C4H10)
n-butane
bp: 0oC
isobutane
bp: -12oC
Resonance and Formal Charge
Sample Problem
Assign formal charges to each atom in the following
resonance forms of the cyanate ion (OCN-). Which
resonance form is likely to contribute most to the
correct structure of OCN-?
A
B
C
• •
••O
••
-1
C N••
-
••
••O
••
C N••
-
••O
••
••
C N
••
0
-1
0
0
0
+1
0 -2
A and B have less formal charge, so they are
preferred over C.
A is preferable to B because it has the negative
formal charge on the more electronegative atom.
-
Octet Rule Exception:
Odd-Electron Species
• Molecules and ions with an odd number of
valence electrons are called free radicals.
Ex: Nitrogen monoxide (NO), a pollutant from car exhaust
•
•• N



• ••
O•
11 valence eThese molecules break the
rules of Lewis Theory, but
they do exist in nature.
They tend to be unstable and reactive.
Free radicals can damage living cells, causing aging,
degenerative conditions and disease.
Octet Rule Exception:
Incomplete Octets
• H, Li, Be and B are exceptions to the octet rule,
since they will attain electron configurations like
He (2 valence electrons.)
Example: BF3
 Li loses its1valence electron.
 H can gain, share, or lose its
•• ••
••
F
1electron.
••
• ••
 Be loses 2 electrons to become Be2+ •
F
B
F
•
••
• It can also share its 2 electrons
• ••

FC’s all 0
in covalent bonds, resulting in
4 valence electrons.
24 valence eB loses 3 electrons to become B3+
• though it commonly shares its 3 electrons in
covalent bonds, resulting in 6 valence electrons .
Octet Rule Exception:
Expanded Octets
• Some atoms in row 3 or lower will form
expanded octets, including S, P, As, Cl, I and Xe.

Ex: SO3 can minimize formal charges by expansion.
•• • -1
• ••
•O
O
•
•
-1 ••O
••

• •
S
• • • -1
O
• ••
•
•
••O
•
S
• ••
O•
+3
All 0
And sometimes there’s no other way to draw the
• ••
structure, such as for XeF2 or AsF5.
F•
•
•
•
•
•
••O Xe•
•• ••
• ••
O
• ••
22 valence e-
• •
•• •F• As •F• •
•
•• ••
•
•
• ••
•
F
40 valence e • F
• • • ••
VSEPR Theory
• The abbreviation VSEPR (say it
“VES-pur”) stands for “valence-shell
electron-pair repulsion.”
• VSEPR theory – repulsion between
pairs of valence electrons around
an atom causes the electron pairs to
be oriented as far apart as possible.
• Treat double and triple bonds the same as
single bonds.
Visual Concept
Linear Geometry
• 2 electron groups around the central atom will
occupy positions opposite each other.
• This results in linear geometry.
• The bond angle is 180°.
•
•• •• •Cl
••
••
Be
•• •
••
Cl
•
••
••
•• O
••
••
C
••
••
O
••
Trigonal Planar Geometry
• 3 electron groups around the central atom will take
positions in the shape of a triangle.
• This results in trigonal planar geometry.
• The bond angle is 120°.
•
•
•• O
••
-
•• •
O•
N
• ••
O
•••
•
•
•• F
••
•• ••
F ••
B
• ••
F
• ••
Tetrahedral Geometry
• 4 electron groups around the central atom will
take positions in the shape of a tetrahedron.
• This results in tetrahedral geometry.
• The bond angle is 109.5°.
H
H C H
H
Trigonal Bipyramidal Geometry
• 5 electron groups around the central atom results
in trigonal bipyramidal geometry.
• The axial positions are above
and below the central atom.
• The equatorial positions are in the
same base plane as the central atom.
• Bond angle between equatorial
positions is 120°.
• Bond angle between axial and equatorial
positions is 90°.
Trigonal Bipyramidal Geometry cont’d
••
••Cl ••
•• • •
Cl
•• ••
•• Cl
••
P
••
••
••
Cl
••
••
Cl
••
Octahedral Geometry
• 6 electron groups around the central atom will take
positions in the shape of a octahedron
• This results in octahedral geometry.
• The bond angle is 90°.
•• •
•• ••
•
•
F••F•
•• •F• S •F• •
•
•• ••
•
•
•• F •F• ••
•• ••
Effect of Lone Pairs
• Lone pairs of electrons also take up space.
• But lone pair groups occupy more space on the
central atom than bonding pairs do.

Their electron density is solely on the central atom
rather than being shared.
• Molecules with lone pairs will have distorted
bond angles.
Bent (Angular) Shape
• 2 bonding pairs and 1 lone
pair of electrons will result
in a bent (angular) shape
with a bond angle of <120o
• 2 bonding pairs and 2
lone pairs of electrons
will result
in a bent
shape with a
bond angle
of <109.5o
••
••
••
O S O ••
••
••
•• ••
H
O
H
Trigonal Pyramidal Shape
• 3 bonding pairs and 1lone pair of electrons
will result in a trigonal pyramidal shape with
a bond angle of <109.5o
••
H N H
H
Equatorial Positioning of Lone Pairs
• When there are 5 electron groups around the
central atom, and some are lone pairs, they will
occupy the
equatorial
positions
because there
is more room
for them there.
Seesaw (Sawhorse) Shape
• When there are 4 bonding pairs
of electrons around the central
atom, and 1 lone pair, the result
is called see-saw shape (aka
sawhorse or teeter-totter.)
• Bond angles are
•• ••
••
F
•• •
<120o (equatorial)
F• ••
•
o
and <90 (between • S
• ••
F
axial & equatorial.)
•
•
•
•• F ••
••
T-shaped
• 3 bonding pairs and 2 lone pairs of electrons
results in a T-shaped molecule with bond
angles of <90o.
•• ••
F ••
••
•• Br ••
F
••
••
•• F ••
••
Linear Shape
• 2 bonding pairs and 3 lone pairs of electrons
results in a Linear
molecule with
bond angles
of 180o.
•
•
•
•
•
••O Xe•
•• ••
• ••
O
• ••
Square Pyramidal Shape
• When there are 5 bonding pairs of electrons
around the central atom and 1 lone pair, the
resultis called a square pyramidal shape.
• Bond angles
are <90o.
•• •
•• ••
•
•
F••F•
•• Br •F• ••
•
•
•• •F• •F• ••
•• ••
Square Planar Shape
• 4 bonding pairs and 2 lone pairs of electrons
results in a Square Planar molecule with bond
angles of 90o.
•• ••
F ••
• ••
•• •F• ••
F
Xe
•• •
• ••
•
•• F ••
••
Multiple Central Atoms
• Many molecules have larger structures with
many interior atoms.
• For these, describe the shape around each
central atom in sequence.
Methanol: CH3OH
Glycine: C2H5NO2
Molecular Polarity
• Molecular Polarity depends on 2 things:
Bond polarity (i.e. electronegativity difference)
2. Molecular shape (symmetrical or not?)
1.
• For a molecule to be polar it must have both
polar bonds & a non-symmetrical shape.
Polar or Nonpolar?
• If dipole moments
cancel, the molecule
is nonpolar, even if
the bonds are polar.
Effects of Molecular Polarity
• Determines the type of
intermolecular forces (polar
molecules are attracted to
other polar molecules.)
• Determines solubility
(like dissolves like.)
• some molecules have both polar
and nonpolar parts (i.e. soap.)
Types of Intermolecular Forces
But Wait a Minute…Quantum
Theory Doesn’t Match Our Models?!
• Look at bonding between C and H, for example.
• According to Quantum Theory, C has 2 half-filled
orbitals, so should form CH2 (with 90o bond angles.)
• Actually forms CH4 (with 109.5o bond angles.)
Valence Bond Theory
• Linus Pauling and others applied the principles of
quantum mechanics to molecules.
• According to their Valence Bond Theory, as two
atoms approached, valence atomic orbitals on the
atoms would interact to form molecular orbitals.
• It was postulated that the valence atomic orbitals
would hybridize before bonding took place.
• The number of standard atomic orbitals combined
= the number of hybrid orbitals formed.
3
sp
Hybridization of Carbon
• The one s orbital and three p orbitals
of carbon hybridize to form four
sp3 orbitals.
• All four of the hybrid orbitals
have the same energy.
Lone Pairs and Hybrid Orbitals
• Hybrid orbitals can accommodate lone pairs as well
as bonding pairs of electrons.
• Example:
ammonia (NH3)
exhibits sp3
hybridization.
Hybridization & Electron Geometry
• The electron
geometry of the
molecule allows
us to predict the
type of hybridization.
• The number of
standard orbitals
added together
always equals the
number of hybrid
orbitals formed.
2
Linear
sp
3
Trigonal
planar
sp2
4
Tetrahedral
sp3
5
Trigonal
bipyramidal
sp3d
6
Octahedral
sp3d2
Why Don’t Multiple Bonds Get a
Hybridized Orbital?
unhybridized
carbon orbitals
sp hybridized

 
2p
2s
 
2sp
 
2p
sp2 hybridized

 
2sp2
sp3 hybridized



2sp3

2p

68
Sigma and Pi Bonds
• a sigma (s) bond results when the bonding
orbitals overlap end to end.
• a pi (p) bond results when the bonding orbitals
overlap side to side.
Multiple Bonds and Hybridization
• One – and only
one - sigma (s)
bond forms
between any
two atoms.
Additional
bonds must
be pi (p)
bonds.
Valence Bond Theory
Practice Problem
Predict the hybridization and bonding scheme
for CH3CHO.
18 valence eC1 – tetrahedral
••
•
•
3
O
•
•
sp hybridization
H
4 sigma bonds
H C1 C2 H
C2 – trigonal planar
sp2 hybridization
H
3 sigma bonds & 1 pi bond
ABE Notation
• ABE notation clearly shows the number of bonding
and non-bonding electron groups. (useful to know
since that determines molecular geometry.)
• A = the central atom
B = the bonding electron groups
x y
E = the non-bonding electron groups
Subscripts (x & y) = the number of each.
• “Electron groups" can mean double
or triple bonds as well as single.
AB E
ABE Notation (cont’d.)
AB4
AB3E
AB2E2
AB2
Bonding Groups: 4
Nonbonding
Groups: 0
Bonding Groups: 3
Nonbonding
Groups: 1
Bonding Groups: 2
Nonbonding
Groups: 2
Bonding Groups: 2
Nonbonding
Groups: 0
CH4
NH3
H2O
CO2
Methane
Ammonia
Water
Carbon Dioxide
ABE Notation (cont’d.)
• ABE notation provides a convenient way of
remembering the hybridization at the central atom.
• Total number of substituents (bonding plus non-bonding
groups) equals the number of hybridized orbitals.
Molecule
ABE Notation
# of Substituents
Hybridization
CH4
AB4
4
sp3
NH3
AB3E
4
sp3
H2O
AB2E2
4
sp3
CO2
AB2
2
sp
SF6
AB6
6
sp3d2
(I3)- ion
AB2E3
5
sp3d