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 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. 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.