Ch. 6: Chemical Bonding I:
Drawing Lewis Structures and
Determining Molecular Shapes
Dr. Namphol Sinkaset
Chem 200: General Chemistry I
I. Chapter Outline
I. Introduction
II. Electronegativity
III. Lewis Structures
IV. Resonance
V. Exceptions
VI. Bond Energies and Bond Lengths
VII. VSEPR Theory
VIII. Molecular Polarity
I. Bonding Theories
• Chemistry revolves around compounds,
so how these are held together is an
important topic.
• How they are bonded predicts many of
their properties.
• We will cover 3 bonding theories.
• In this chapter, we expand on Lewis
theory.
I. Importance of Shape
• In condensed phases (liquids/solids),
molecules are in close proximity, so they
interact constantly.
• The 3-D shape of a molecule determines
many of its physical properties.
• We want to be able to predict 3-D shape
starting from just a formula of a covalent
compound.
I. Binding Sites
II. Lewis Theory
• Simple interpretation of Lewis theory
implies that e-’s are equally shared.
II. Reality Shows Otherwise
II. Electronegativity
• Atoms don’t share e-’s equally.
• Electronegativity is the relative ability of
a bonded atom to attract shared e-.
 It can be thought of as how greedy an
atom is for e- when it is sharing them.
II. Unequal Sharing of e• More electronegative atoms will pull
shared e- towards them.
• This results in a partial charge
separation which can be indicated in
one of two ways.
This is known as a polar covalent bond.
II. Electronegativity Values
II. Using ΔEN
• Differences in electronegativity can be
used to determine the bond type.
II. Ionic Character of Polar Bonds
III. Lewis Structures
• The first step to getting the 3-D shape of
a molecule is getting the correct 2-D
structure.
• The 2-D structure will be the basis of
our 3-D shape assignment.
• We outline the general steps for
drawing Lewis structures.
III. Steps for Drawing Lewis Structures
1) Determine total # of valence e-.
2) Place atom w/ lower Group # (lower
electronegativity) as the central atom.
3) Attach other atoms to central atom with
single bonds.
4) Fill octet of outer atoms. (Why?)
5) Count # of e- used so far. Place
remaining e- on central atom in pairs.
6) If necessary, form higher order bonds to
satisfy octet rule of central atom.
7) Allow expanded octet for central atoms
from Period 3 or lower.
III. Lewis Structure Practice
• Draw correct Lewis structures for NF3,
CO2, SeCl2, PI5, IF2-, IF6+, and H2CO.
IV. Multiple Valid Lewis
Structures
• Sometimes more than one Lewis
structure can be drawn for the same
molecule.
• For example, ozone (O3).
IV. Resonance Structures
• Resonance structures are also known as
resonance forms.
• A resonance structure is one of two or more
Lewis structures that have the same skeletal
structure (atoms in same place), but different
electron arrangements.
IV. Resonance Hybrid
• Neither resonance form is a true picture of
the molecule.
• The molecule exists as a resonance hybrid,
which is an average of all resonance forms.
• In a resonance hybrid, e- are delocalized over
the entire molecule.
IV. Sample Problem
• Draw the resonance structures of the
carbonate anion.
IV. Important Resonance Forms
• If all resonance forms have the same
surrounding atoms, then each
contributes equally to the resonance
hybrid.
• If this is not the case, then one or more
resonance forms will dominate the
resonance hybrid.
• How can we determine which forms will
dominate?
IV. Formal Charge
• formal charge: the charge an atom
would have if bonding e- were shared
equally
formal charge = (# valence e-) – (unshared e- + ½ shared e-)
IV. Formal Charges in O3
• We calculate formal charge for each atom in
the molecule.
• For oxygen atom A (on the left), there are 6
valence e-, 4 unshared e-, and 4 shared e-.
The formal charge for this O atom is 0.
• NOTE: sum of all formal charges must equal
the overall charge of the molecule!
IV. Using Formal Charges
•
Formal charges help us decide the most
important resonance forms when we
consider to the following guidelines:
1) Small f.c.’s are better than larger f.c.’s.
2) Same sign f.c.’s on adjacent atoms is
undesirable.
3) Electronegative atoms should carry higher
negative f.c.’s.
IV. Sample Problem
• Find the dominant resonance structures
for the sulfate anion.
V. Exceptions to the Octet Rule
• We’ve already discussed expanded
valence cases, but there are other
exceptions as well.
 Compounds w/ odd # of e-’s: free radicals.
Examples include NO and NO2.
 Incomplete octets: e- deficient atoms like
Be and B, e.g. BeCl2 and BF3.
 Expanded octets – when d orbitals are
used to accommodate more than an octet.
VI. Bonding and Energy
• Lewis theory shows a bond as sharing
two electrons, but not all bonds are
identical.
• Bonds can vary in their strength and in
their length.
• Bond energy is the energy needed to
break 1 mole of the bond in the gas
phase.
VI. Average Bond Energies
VI. Bond Length
• Bond length is the distance between
bonded atoms.
• In general, as the bond weakens, the
bond length increases.
• As with bond energies, we can list
average bond lengths.
VI. Average Bond Lengths
VII. VSEPR Theory
• From a correct Lewis structure, we can
get to the 3-D shape using this theory.
• VSEPR stands for valence shell
electron pair repulsion.
• The theory is based on the idea that epairs want to get as far away from each
other as possible!
VII. VSEPR Categories
• There are 5 electron geometries from which
all molecular shapes derive.
VII. Drawing w/ Perspective
• We use the conventions below to depict a 3-D
object on a 2-D surface.
VII. Determining 3-D Shape
• The 5 electron geometries (EG) are a
starting point.
• To determine the molecular geometry
(MG), we consider the # of atoms and
the # of e- pairs that are associated w/
the central atom.
• All the possibilities for molecular
geometry can be listed in a
classification chart.
VII. Linear/Trigonal Planar
Geometries
• First, we have the linear and trigonal
planar EG’s.
EG
Bonds
Lone Pairs
MG
Linear
2
0
linear
Trigonal
planar
3
0
2
1
trigonal
planar
bent
VII. Tetrahedral Geometries
EG
Bonds
Lone Pairs
MG
Tetrahedral
4
0
tetrahedral
3
1
pyramidal
2
2
bent
1
3
linear
VII. Trigonal Bipyramidal
Geometries
EG
Trigonal
Bipyramidal
Bonds Lone Pairs
MG
5
0
4
3
1
2
trigonal
bipyramidal
see-saw
T-shaped
2
1
3
4
linear
linear
VII. Octahedral Geometries
EG
Octahedral
Bonds Lone Pairs
MG
6
0
octahedral
5
1
4
2
square
pyramidal
square planar
3
3
T-shaped
2
4
linear
1
5
linear
VII. Steps to Determine
Molecular Geometry
1) Draw Lewis structure.
2) Count # of bonds and lone pair e-’s on
the central atom.
3) Select electronic geometry.
4) Place e-’s and atoms that lead to most
stable arrangement (minimize erepulsions).
5) Determine molecular geometry.
VII. Trig Bipy is Special
• In other EG’s, all
positions are
equivalent.
• In trig bipy, lone
pairs always choose
to go equatorial first.
• Why?
VII. Lone Pairs Take Up Space
• Lone pair e-’s don’t have another
nucleus to “anchor” them.
VII. Distortion of Angles
• Lone pair e-’s take up a lot of room, and they
distort the optimum angles seen in the EG’s.
VII. Some Practice
• Draw the molecular geometries for SF4,
BeCl2, ClO2-, TeF5-, ClF3, and NF3.
VII. Larger Molecules
VIII. Molecular Polarity
• Individual bonds tend to
be polar, but that
doesn’t mean that a
molecule will be polar
overall.
• To determine molecular
polarity, you need to
consider the 3-D shape
and see if polarity
arrows cancel or not.
VIII. Sample Problem
• Determine the molecular geometry of
IF2- and state whether it is polar or
nonpolar.
VIII. Polarity and Properties
• Polarity is the result
of a compound’s
composition and
structure.
• Knowing that a
compound is
polar/nonpolar
allows us to explain
its properties.