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C h e m i s t r y 1 A : C h a p t e r 1 0 P a g e | 1

Chapter 10: Chemical Bonding II: Molecular Shapes; VSEPR,

Valence Bond and Molecular Orbital Theories

Homework:

Read Chapter 10: Work out sample/practice exercises.

Check for the MasteringChemistry.com assignment and complete before due date

Molecular Shapes:

Properties of molecular substances depend on its 3D structure Bonding neighbors, what is next to what (skeleton arrangement)

Type of bonding; polar, nonpolar, ionic

Shape and Polarity; overall do dipoles cancel or is there an overall dipole

moment

Limitations in Lewis Structures:

• Lewis theory predicts the number of electron regions (lone pair or any bond; single

double, triple), but does not determine actual bond angles.

• Lewis theory predicts trends in properties, but does not give good numerical

predictions of bond strength and bond length

• Lewis theory cannot write one correct structure for molecules where resonance is

important

• Lewis theory often does not correctly predict magnetic behavior of molecules.

Oxygen, O2, is paramagnetic, though the Lewis structure predicts it is diamagnetic

Valence Shell Electron Pair Repulsion (VSEPR) Theory: Three-dimensional

• Electron groups (all negatively charged) around the central atom are most stable

when they are as far apart as possible –valence shell electron pair repulsion theory.

Use all the information gained in the Lewis Dot Structure and convert it to a three

dimensional model to predict electronic and molecular shapes, angles, and polarity

of the molecule.

VSEPR Guidelines:

Start with information from a Lewis Dot Structure

Electronic and Molecular 3D shapes

Bonds angles: When electron groups attach to different size atoms the ideal

bond angles are affected. Lone pairs (nonbonding) use more space.

Polarity of whole substance (ionic, ion, nonpolar, polar molecule)

Electronic and Molecular Geometry:


Count the electron regions. Electron regions will give an electronic shape

while the number of bonded versus nonbonded regions will give the molecular

shape.

# Electron

regions

2 3 4 5 6

Electronic

geometry

five basic

shapes

Linear

180˚

trigonal planar

120˚

tetrahedral

109.5˚

trig. bipyramidal

90˚, 120˚, 180˚

octahedral

90˚, 180˚

molecular

geometry

Linear Trig planar,

bent

Tetrahedral,

Trig. pyramidal,

bent

trig.bipyramidal,

see saw,

T-shaped,

linear

octahedral,

square pyramidal,

square planar

Samples

Imperfect Geometry:

When electron groups attach to different size atoms the ideal bond

angles are affected

CH2O ideally should be trigonal planar with angles of 120° each.

In reality the angle between the smaller H atoms is smaller.

Lone pairs (nonbonding electrons)

use more space. Ideally four

regions should spread out to angles

of 109.5°. Notice how the bond

angles around the atoms are forced

closer together as the unseen

nonbonding electrons take more

space.

Website to try: ChemEdDL.org Click on molecules 360. This website shows the 3D

structure of many chemicals and allows you to rotate in three dimensions, showing bonding,

bond length, dipole arrows, dipole moment, etc.

Writing 3D shapes on paper: May use lines and wedges.


Multiple Central Atoms:

Describe the shape around each central atom

atom separately.

Polarity of the Molecule:

Polar: must have polar bonds (electronegativity difference between the

neighbor atoms with a measureable bond dipole moment) and an

unsymmetrical shape (lone pairs or varying atom neighbors)

Polarity affects properties: boiling points, solubilities (like dissolves like)

HCl and H2O are both polar

CO2 is nonpolar


Valence Bond (VB) Theory: Three-dimensional

The Valence Bond theory is a quantum mechanical model that expands the previous

two theories to describe the electronic nature of covalent bonds.

• Valence bond theory applies principles of quantum mechanics to molecules

• A chemical bond between atoms occurs when atomic orbitals and hybridized

atomic orbitals interact with those in another atom to form a new molecular

orbital with two electrons.

If orbitals align along the axis between the nuclei, sigma bonds which

directly overlap will form ( bonds). It is possible to rotate a sigma bond

If orbitals align outside the axis, pi bonds form, which indirectly overlap

above and below ( bonds). Unable to rotate without breaking bonds. This

causes cis and trans structural isomers.

VB Guidelines:

Use all the information from a Lewis Dot Structure

Hybridizing some orbitals allow for more bonds and more stability

Visualize orbital picture using atomic (s, p, d, f) and hybridized (sp, sp2, sp3,

sp3d, and sp3d2) orbitals

Direct overlap orbitals, sigma () bonds

Indirect overlap orbitals, pi () bonds

All types of bonds have only one bond. Double bonds have 1 and 1

and triple bonds have 1 and 2 bonds

Valence Bond (Bubble) Pictures draw the orbitals in balloon type pictures

Delocalized bonding occurs in substances with resonance


Chemical bonds between atoms occur when atomic orbitals interact with those in

another atom to form a new molecular orbital with two electrons. Sigma bonds

(direct overlap) are stronger than pi bonds (indirect overlap).


Double bond:

CH2O

C

3 sp2 hybridized orbitals and

1 p unhybridized orbital

H

1 s orbital on each

O

1 s unhybridized orbital and

3 p unhybridized orbitals

Triple bond:

C2H2

C

2 sp hybridized

orbitals and

2 p unhybridized

orbitals

H

1 s orbital on each

Limitations in Valence Bond Theory:

Valence Bond theory predicts bond strengths, bond lengths, and bond rigidity better

than Lewis theory.

• Other properties, such as the magnetic behavior of O2, of molecules are not

predicted well.

• VB theory views electrons as localized in overlapping atomic orbitals and it

doesn’t account for delocalization


Molecular Orbital (MO) Theory: The Molecular Orbital Theory is separate from the first three. This theory explains

the paramagnetic behavior found in O2 gas molecules.

• In MO theory, Schrödinger’s wave equation is applied to the molecule to

calculate a set of molecular orbitals

• Electrons and orbitals belong to the whole molecule – Delocalization

• A Bonding Molecular Orbital forms when wave functions combine

constructively, resulting in a molecular orbital with lower energy than the

original atomic orbitals. Most of the electron density is between the nuclei.

Lower energy-stabilizing

• The Antibonding* Molecular Orbital forms when wave functions combine

destructively, resulting in a molecular orbital with more energy than the

original atomic orbitals. Most of the electron density is outside the nuclei

creating nodes between nuclei. Higher energy-unstable

Sigma () 1s molecular orbitals (2s looks the same, but a bit bigger)

Sigma () px molecular orbitals


Pi () py or pz molecular orbitals

MO Guidelines:

Electrons belong to the molecule, not the individual atoms

For this class, limit most of the discussion and examples to diatomic species

such as: H2, O2, CN-1, HF.

Occasionally this gives a more accurate electronic structure than VB

Combination of two atomic orbitals makes a molecular orbital

Bonding orbitals are sigma or pi orbitals. Sigma orbitals directly overlap

and pi orbitals indirectly overlap

Antibonding* sigma or pi orbitals create a node between the atoms with no

overlap

Two atomic s orbitals combine to form a lower energy bonding and a

higher energy * antibonding* orbital

six atomic p orbitals combine to form lower energy bonding orbitals, and

2 degenerate orbitals and higher energy antibonding* orbitals,and 2

degenerate orbitals

Predicts paramagnetic or diamagnetic behavior

Predicts bond order

Compares bond lengths and bond strengths

For diatomic molecules with fewer than 15 total electrons like N2, energy

increases as follows: s, 1s*, 2s, 2s*, 2p,2p, 2p, 2p*, 2p*, 2p*


For diatomic molecules with 15 or more total electrons like O2, energy

increases as follows: s, 1s*, 2s, 2s*, 2p,2p,2p, 2p*, 2p*, 2p*

Magnetic behavior of O2

Diatomic oxygen is attracted to

a magnetic field, indicating

paramagnetic behavior, so it

has unpaired electron(s)


Heteronuclear Diatomic Elements and Ions:

• The more electronegative an atom is, the

lower in energy are its orbitals

• Lower energy atomic orbitals contribute

more to the bonding MOs

• Higher energy atomic orbitals contribute

more to the antibonding MOs

• Nonbonding MOs remain localized on the

atom donating its atomic orbitals


Polyatomic Molecular Orbitals:

• Atomic orbitals of all the atoms in a molecule, even those with 3 or more atoms,

combine to make a set of molecular orbitals, delocalized over the entire molecule

• Predictions made using molecular orbital theory, (especially resonance molecules and

predicting magnetic properties), match the real molecule properties better than either

Lewis or Valence bond theories.

Ozone, O3:

MO theory predicts equivalent bond lengths due to the

delocalized electrons.


Molecular Shapes, Handedness and Drugs:

The shapes of molecules can dramatically change its characteristics. Mirror images

have different biological properties due to the specific shapes of receptor sites in the

body. For a molecule to exhibit handedness it needs four different groups attached to

a carbon.

Identify the electronic and molecular geometries, angles, and VB hybridization

a) h)

b) i)

c) j)

d) k)

e) l)

f)

g) m)


Fill in the following tables: First page follows octet and duet rules, second page has extended octets.

#of electron

regions and

VB hybrid

number of

bonded

atoms

electronic

geometry

name

molecular

geometry

name

bond angles rough

3-D

sketch

an example

molecule or ion

any

1

linear

linear

(180)

O−−O

H2

CO

HF

N2

CN-1

CO2

3

120

3

sp2

bent

or angular

4

109.5

trigonal

pyramidal

H2O


#of electron

regions and

VB hybrid

number of

bonded

atoms

electronic

geometry

name

molecular

geometry

name

bond angles rough

3-D

sketch

an example

molecule or ion

5

trigonal

bipyramidal

see-saw

3

180

90

(120)

5

sp3d

2

6

octahedral

BrF5

6

sp3d2

square

planar

Chapter 10: Chemical Bonding II: Molecular Shapes. Valence Bond and Molecular Orbital Theories P a g e | 15

Examples:

1. The valence bond hybrid atomic orbitals sp3 are used by both C in CH4 and O in H2O. Yet, the

bond angles between atoms in H2O are less than in CH4. Explain.

2. Describe completely the main features of each of the following and explain what useful information

we gain from each.

a) Lewis Structures

b) Valence Shell Electron Pair Repulsion (VSEPR) theory

c) Valence Bond (VB) theory

d) Molecular Orbital (MO) theory

3. a) Draw all possible resonance Lewis structures for NO3-1. Include formal charges and the correct

angles.

b) Draw the "realistic" hybrid resonance structure with appropriate angles that takes and

average of the Lewis structures in part a. Include formal charges (fractions) and bond

orders (fractions). Include nonbonding electrons on central atom but not on terminal

atoms.

c) Sketch the valence bond (bubble) probability picture of one of the NO3-1 resonances.

Identify and label the hybridized orbitals. Identify sigma and pi bonds.

4. Draw and identify the cis and trans isomers for 1,2-dichloroethene, C2H2Cl2

5. For each of the following: B2, Ne2, O2

a) Give the molecular orbital (MO) energy diagram for each.

b) Write the MO configurations for O2 starting with ( 1s)2

c) Give the bond order of each B2, Ne2, O2

d) List the species in decreasing order of bond energy and stability

e) Identify each as diamagnetic or paramagnetic?

f) Using the bond order information, which is least expected to exist. Explain why.

g) Which would have the shortest bond length? Explain.


6. Complete the following table for the indicated substances.

Electronegativities: Na = 0.9, N = 3.0, O = 3.5, F = 4.0, Cl = 3.0, Br = 2.8, I = 2.5

substance SO2 C2H4O2 ICl5 NaBrO3

a) Draw the best

Lewis

structure(s),

resonances, and

structural isomers

if any with octet

b) Include formal

charges if they

are not zero

c) Indicate polar

bonds with dipole

arrows toward

the more

electronegative

name electronic

geometry around

central atom

give hybrid

orbital for center

name molecular

geometry around

central atom

show 3-D sketch

with atoms &

bonds in it

give all bond

angles

how many sigma

bonds? how

many pi bonds?

is it an ionic

compound, polar

or nonpolar

molecule or an

ion?

Draw the VB

hybrid resonance

(bubble) picture


7. Complete the following table for the indicated substances. substance SCN-1 I3

-1 SF6 K2SO3

a)Draw the best

Lewis structure(s),

resonances, and

structural isomers

if any with octet

b) Include formal

charges if they are

not zero

c) Indicate polar

bonds with dipole

arrows toward the

more

electronegative

Answer questions

below for SO3-2

name electronic

geometry around

central atom

give hybrid orbital

for center

name molecular

geometry around

central atom

show 3-D sketch

with atoms &

bonds in it

give all bond

angles

how many sigma

bonds? how many

pi bonds?

is it an ionic

compound, polar

or nonpolar

molecule or an

ion?

Draw the VB

hybrid resonance

(bubble) picture

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