Ch. 5: Chemical Bonding I

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. Practice Problem 5.1

• 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 5.2

• 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 5.3

• 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 e-

pairs 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 trigonal

planar

2 1 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 Bonds Lone Pairs MG

Trigonal

Bipyramidal

5 0 trigonal

bipyramidal

4 1 see-saw

3 2 T-shaped

2 3 linear

1 4 linear

VII. Octahedral Geometries

EG Bonds Lone Pairs MG

Octahedral 6 0 octahedral

5 1 square

pyramidal

4 2 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 e-

repulsions).

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. Practice Problem 5.4

• 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 5.5

• Determine the molecular geometry of

XeF2 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.