Chapter 10 Notes

Chemical Bonding II

Molecular Shape,

Valence Bond Theory, and

Molecular Orbital Theory

Sections 10.1 – 10.8

I. Artificial Sweeteners: Fooled by Molecular Shape

Artificial sweeteners such as aspartame (Nutrasweet) taste sweet like

sugar but do not have the calories of sugar. This is because taste and

nutritional value are independent properties. The caloric value of

food depends on the amount of energy released when the food is

metabolized in the body. Many artificial sweeteners are not even

metabolized by the body- they just pass straight though.

The taste of a food begins with specialized cells on the tongue that

detect different molecules of food. These cells are so sensitive that

the tongue can detect one molecule of sugar out of thousands of

different molecules in a bite of food. We experience certain tastes

when molecules of food fit into a special part of a taste receptor on

our tongue, much in the same way a key fits into a lock. Artificial

sweeteners “trick” our tongues into tasting sweetness because these

molecules mimic the molecular shape of the sugar molecule and fit

snuggly into our taste receptors.

- used to determine the 3- D shape of a molecule

- based on the idea that electrons groups (bonds & lone pair e─)

repel one another to varying degrees

- the combination of repulsions on the central atom of a molecule

determines its 3-D shape

• VSEPR THEORY (VALENCE SHELL ELECTRON PAIR REPULSION):

II. VSEPR Theory: The Five Basic Shapes

*The basic idea behind VSEPR is that repulsions between electron

groups determine molecular geometry. The preferred geometry is

the one in which the electron groups have the maximum separation

(and therefore the minimum energy) possible.

Preview

We will first look at molecular geometries where there are two to

six electron groups around a central atom and all of the electron

groups are bonding groups (single or double bonds). We will then

look at what happens to the shapes when some of the electron

groups become lone pair electrons.

A. TWO ELECTRONS GROUPS: LINEAR GEOMETRY

Bond Angles: 180°

Other examples:

SiO2

CS2

Si O O

C S S

16

16

B. THREE ELECTRONS GROUPS: TRIGONAL PLANAR

GEOMETRY

Bond Angles: 120° Other examples:

BH3 SO3

B

H

H H

S

O

O O

In CH2O, the bond angles deviate slightly from

120° because there is more electron density in a

double bond than in a single bond.

Different electron groups repel each other in

slightly different ways.

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C. FOUR ELECTRONS GROUPS: TETRAHEDRAL GEOMETRY

Bond Angles: 109.5°

Other examples:

SiCl4 CF4

Si

Cl

Cl Cl

Cl

C

F

F F

F

32 32

D. FIVE ELECTRONS GROUPS: TRIGONAL BIPYRAMIDAL

GEOMETRY

Bond Angles: 90°, 120°

Other examples:

SOF4 ClO2F3

S

O

F F

F

F

40 40

Cl

O

F O

F

F

E. SIX ELECTRONS GROUPS: OCTAHEDRAL GEOMETRY

Bond Angles: 90°

Other examples:

SCl6

S

Cl

Cl Cl

Cl

Cl

Cl

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III. VSEPR Theory: The Effect of Lone Pairs

Preview

We will now look at molecular geometries where there are four to

six electron groups around a central atom and some of the electron

groups are lone pairs. Keep in mind that these shapes are all

variations of the geometries from the last section, but now one or

more bonding groups have been replaced with lone pairs. The

difference in shapes results from the fact that lone pair electrons

repel other lone pair and bonding electrons to a greater extent than

bonding electrons repel one another.

THE ORDER OF ELECTRON PAIR REPULSION:

lone pair - lone pair > lone pair- bonding pair > bonding pair- bonding pair

Most repulsive Least repulsive

* Want to be as far away

from each other as

possible

• In this section, don’t get confused between electron and molecular

geometries.

• In the last section, the electron and molecular geometries were the same

(same name for each).

• In this section, they are different.

• There are only 5 electron geometries, but there are 11 molecular

geometries.

THE DIFFERENCE BETWEEN:

Electron Geometry:

Molecular Geometry:

The geometrical arrangement of electron groups

The geometrical arrangement of the atoms

A. FOUR ELECTRON GROUPS WITH LONE PAIRS

(TETRAHEDRAL ELECTRON GEOMETRY)

Bond Angles: 107°

TRIGONAL PYRAMIDAL (MOLEC. GEO.): ONE LONE PAIR

Other examples:

PF3 NCl3

P

F F F

N

Cl Cl Cl

26 26

Bond Angles: 104.5°

BENT (MOLEC. GEO.): TWO LONE PAIRS

Other examples:

SCl2 H2S

A. FOUR ELECTRON GROUPS WITH LONE PAIRS

(TETRAHEDRAL ELECTRON GEOMETRY)

S Cl Cl

S H H

18 8

Summarizing Tetrahedral Electron Geometries

B. FIVE ELECTRON GROUPS WITH LONE PAIRS

(BIPYRAMIDAL ELECTRON GEOMETRY)

Bond Angles: 90°, 120°

SEESAW (MOLEC. GEO.): ONE LONE PAIR

Other examples:

SeCl4 IOF3

Se

Cl

Cl

Cl

Cl

I

O

F

F

F

34

Lone pair must go equatorial

34

B. FIVE ELECTRON GROUPS WITH LONE PAIRS

(BIPYRAMIDAL ELECTRON GEOMETRY)

Bond Angles: 90°, 120°

T-SHAPED OR LINEAR (MOLEC. GEO.): TWO OR THREE

LONE PAIRS

Other examples:

ClF3 I3

-

Cl

F

F

F I

I

I

28 28

Lone pairs must go equatorial

C. SIX ELECTRON GROUPS WITH LONE PAIRS

(OCTAHEDRAL ELECTRON GEOMETRY)

Bond Angles: 90°

SQUARE PYRAMIDAL (MOLEC. GEO.): ONE LONE PAIR

Other examples:

XeOF4 BrF4-

Xe

O

F

F

F

F

Br F

F

F

F

SQUARE PLANAR (MOLEC. GEO.): TWO LONE PAIRS

Lone pairs must go axial

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