Unit 2 Chp 9 notes B and L.ppt

UNIT 2

Molecular Geometries and

Bonding Theories

Localized Electron Model

Localized Electron Model

Lewis structures are an application of the

“

Localized Electron Model

”

L.E.M.

says: Electron pairs can be

thought of as

“

belonging

”

to pairs of

atoms when bonding

VSEPR Model

VSEPR model uses the

Coulombic repulsion

between electrons as a

basis for predicting the

arrangement (shapes) of

electron pairs around a

central atom.

VSEPR – Valence Shell Electron Pair

VSEPR – Valence Shell Electron Pair

Repulsion

Repulsion

X + E Overall StructureOverall Structure FormsForms

2 Linear AX₂

3 Trigonal Planar AX₃, AX₂E

4 Tetrahedral AX₄, AX₃E, AX₂E₂

5 Trigonal bipyramidal _{AX5, AX4E, AX3E2, AX2E3}

6 Octahedral AX₆, AX₅E, AX₄E₂

A = central atom

A = central atom

X =

X =

atoms bonded to A

atoms bonded to A

E = nonbonding electron pairs on A

How to Predict Shapes

Shapes depend on two things:

How many atoms bonded to center?

How many lone pairs of electrons around

center?

Must ask yourself –

How many electron

domains around center atom?

VSEPR: Linear

VSEPR: Linear

AX

VSEPR: Trigonal Planar (electronic

VSEPR: Trigonal Planar (electronic

geometries)

geometries)

AX

AX

AX

AX

E

E

BF

BF

SnCl

SnCl

Trigonal Planar

VSEPR: Tetrahedral

VSEPR: Tetrahedral

AX

AX

AX

AX

E

E

AX

AX

E

E

CCl

CCl

PCl

PCl

Cl

Cl

O

O

Tetrahedral

Trigonal Pyrimidal

VSEPR: Trigonal Bi-pyramidal

VSEPR: Trigonal Bi-pyramidal

AX

AX

AX

AX

E

E

AX

AX

E

E

AX

AX

E

E

PCl

PCl

SF

SF

ClF

ClF

I

I

-Trigonal Bipyrimidal

Distorted Tetrahedral

T-Shaped

VSEPR: Octahedral

VSEPR: Octahedral

AX

AX

AX

AX

E

E

AX

AX

E

E

SF

SF

ICl

ICl

-

-BrF

BrF

Octahedral

Square Pyrimidal

Predicting Molecular Geometry

1. Draw Lewis structure for molecule.

2. Count number of lone pairs on the central atom and number of atoms bonded to the central atom.

3. Use VSEPR to predict the geometry of the molecule.

What are the molecular geometries of SO₂ and SF₄?

S

O O

AB₂E bent

S

F F

F F

AB₄E

Distorted Tetrahedral

10.1

AP Practice Question

According to the VSEPR model, the progressive

decrease in the bond angles in the series of

molecules CH

, NH

, and H

O is best accounted for

by the

A.increasing strength of the bonds

B.decreasing size of the central atom

AP Practice Question

Which of the following complete Lewis

diagrams represents a molecule containing

a bond angle that is closest to 120°?

AP Practice Question

Refer to the following diatomic species (A) Li₂

(B) B₂ (C) N₂ (D) O₂ (E) F₂

1. Has the largest bond-dissociation energy?

N₂

2. Has a bond order of 2?

AP Practice Question

In the following diagrams, elements are

represented by X and Z, which form molecular

compounds with one another. Which diagram

represents a molecule that has a bent molecular

geometry?

AP Practice Question

Which of the following Lewis electron-dot

diagrams represents the molecule that

contains the smallest bond angle?

H F H F

Polar covalent bond

or

polar bond

is a covalent

bond with greater electron density around one of

the two atoms

electron rich region electron poor

region _e- _{poor e}- _rich

-Electronegativity Classification

• Non-polar: equal sharing of electrons

• Polar: unequal sharing of electrons

• Ionic: no sharing of electrons

*The greater the difference in electronegativity, the more

polar the bond!

Subtraction Value

Bond Classification

0.0 - 0.3

Non-polar covalent (NPC)

0.4 – 1.9

Polar covalent (PC)

Can measure Polarity using Dipole

Moment, μ

μ = Q r

μ is Dipole moment, measured in Debyes, D

Q = charge

r = bond distance

*equation not very important nor

Sample Question

Molecular Polarity

1. Polarity

: results from

having one positive and

one negative part in a

molecule (Polar

Covalents).

2. More electronegative

atom is partially

negative, less

Polarity in Molecule

a. Polar bonds can be overall polar molecules if

the molecule has a net direction of charge.

b. Assign vectors along the bond to assess the

molecule.

c. Examples:

Dipole Moments and Polar Molecules

H F

electron rich region

electron poor region

Which of the following molecules have a dipole moment? H₂O, CO₂, SO₂, and CH₄

O _H H dipole moment polar molecule S O O C O O

no dipole moment nonpolar molecule dipole moment polar molecule C H H H H

AP Practice Question

Refer to the following gaseous molecules:

(A) BeCl

(B) SO

(C) N

(D) O

(E) F

1

Is best represented by two or more resonance forms?

2. Is a polar molecule?

2

Lewis diagrams of molecules of three different

hydrocarbons are shown above. Which of the following

claims about the molecules is best supported by the

diagrams?

A.All the atoms in molecule 1 lie in one plane.

B.All the molecules have the same empirical formula.

C.The C-C-C bond angle in molecule 2 is close to 180°

AP Practice Question

Which of the following molecules contains polar

covalent bonds but is a nonpolar molecule?

A. CH

Cl

B. CH

Cl

C. NH

Helpful Polarity Website

www.chemeddl.org

Hybridization Explanation Video:

Hybridization -

Hybridization -

The Blending of Orbitals

The Blending of Orbitals

AKA Valence Bond Theory

AKA Valence Bond Theory

Poodle

+

+ Cocker Spaniel

= = = = + +

s orbital p orbital

Cockapoo

We have studied electron configuration notation and the sharing of electrons in the formation of covalent bonds.

Methane is a simple natural gas. Its molecule has a carbon atom at the center with four hydrogen

atoms covalently bonded around it.

What Proof Exists for Hybridization?

What Proof Exists for Hybridization?

What is the expected orbital notation of carbon in its ground state?

(Hint: How many unpaired electrons does this carbon atom have available for bonding?)

Can you see a problem with this?

You should conclude that You should conclude that

carbon only has

carbon only has TWOTWO electrons available for electrons available for

bonding. That is not bonding. That is not

enough! enough!

How does carbon overcome this problem so that How does carbon overcome this problem so that

it may form four bonds? it may form four bonds?

Carbon

The first thought that

The first thought that

chemists had was that

chemists had was that

carbon promotes one of

carbon promotes one of

its

its

2s

2s

electrons…

electrons…

Carbon

Do you see a problem with such an arrangement…

Three of the carbon-hydrogen bonds would involve an electron pair in which the carbon electron was a 2p, matched with the lone 1s electron from a

hydrogen atom.

A Problem Arises…

This would mean that three of the bonds in a methane molecule would be identical, because they would involve electron pairs of equal energy.

The fourth bond is between a 2s electron from the carbon and the lone 1s hydrogen electron.

Such a bond would have slightly less energy than the other bonds in a methane molecule.

This bond would be slightly different in character than

the other three bonds in methane.

This difference would be measurable to a chemist by determining the bond length and bond energy.

The simple answer is,

“No”.

Chemists have proposed an explanation – they call it

Hybridization.

Hybridization is the combining of two or more orbitals of nearly equal energy within the same atom into

orbitals of equal energy. Measurements show that all four bonds in methane are equal. Thus, we need a new explanation for the bonding in methane.

In the case of methane, they call the hybridization

sp3, meaning that an s orbital is combined with three

p orbitals to create four equal hybrid orbitals.

These new orbitals have slightly MORE energy than the 2s orbital…

… and slightly LESS energy than the 2p orbitals.

Here is

another way to look at the sp3

hybridization and energy profile…

Another hybrid is the sp2, which combines two orbitals

from a p sublevel with one orbital from an s sublevel.

One p orbital

remains unchanged.

While sp3 is the hybridization observed in methane,

there are other types of hybridization that atoms undergo.

These include sp

hybridization, in which one s orbital combines with a

single p orbital.

This produces two hybrid orbitals, while leaving two normal p orbitals

Hybridization and Molecular Geometry

Hybridization and Molecular Geometry

Forms

Forms Overall StructureOverall Structure (Electronic)

(Electronic)

Hybridization

Hybridization

of

of ““A_A”” Angle_AngleBond Bond

AX₂ Linear sp 180°

AX₃, AX₂E Trigonal Planar sp2 _120°

AX₄, AX₃E, AX₂E₂ Tetrahedral sp3 _109.5°

A = central atom

A = central atom

X =

X =

atoms bonded to A

atoms bonded to A

E = nonbonding electron pairs on A

Examples:

Predict the hybridization of the following

molecules:

1.NH

2.C

H

3.CO

*You must first start from a Lewis Structure,

Predict the VSEPR shape, remember the

sp

2

hybridization of C

Sigma and Pi Bonds

Sigma and Pi Bonds

Sigma () bonds exist in the region directly between two bonded atoms.

Pi () bonds exist in the region above and below a line drawn between two bonded atoms.

Single bond 1 sigma bond Double Bond 1 sigma, 1 pi bond

Another way to show pi (



) bonds

Sigma (



) and Pi Bonds (



)

How many



and



bonds are in the acetic acid

(vinegar) molecule CH

COOH?

C H

H

C H

O

O H  bonds = 6 + 1 = 7

 bonds = 1

The De-Localized Electron Model

The De-Localized Electron Model

H H H H H H H H H H H H

Generalization:

1. Bond formation is overlap between orbitals.

2. In multiple bonds, overlap leads to formation of

sigma and pi bonds.

3. Overlap stronger in sigma bonds (sigma bonds

have larger bond energy).

4. Pi bonds prevent rotation of molecules

5. Pi bonds can be delocalized, such as in

benzene.

AP Practice Question

Pi (π) bonding occurs in each of the following

species EXCEPT

A. CO

B. C

H

C. CN

AP Practice Question

Which of the following molecules contains exactly

three sigma (σ) bonds and two pi (π) bonds?

In the reaction represented above, what is the

hybridization of the C atoms before and after the

reaction occurs?

Before After

sp sp2

Before After

sp sp3

Before After

sp2 _sp

Before After