Chapter 9Molecular Geometry

Introduction

1. Lewis Structures help us understand the compositions of molecules & their covalent bonds, but not their overall shapes.

2. The properties of a substance largely depend on the shape & size of its molecules, together with the strength & polarity of its bonds.

3. Examples: Taxol, smell, vision

VSEPR TheoryValence-shell electron-pair repulsion 1. The overall shape of a molecule is determined by its

bond angles.

2. The VSEPR Model is used to predict Molecular Shapes or Molecular Geometries.

3. The VSEPR Theory assumes that each atom in a molecule will be positioned so that there is minimal repulsion between the valence electrons of that atom.

Five Basic Shapes1. Linear

2. Trigonal Planar

3. Tetrahedral

4. Trigonal Bipyramidal

5. Octahedral

Chart to Memorize for Basic ShapesBond Angle(s) #Electron-Pairs Hybridization

180 2 sp

120 3 sp2

109.5 4 sp3

90 & 120 5 sp3d

90 6 sp3d2

Using VSEPR Model to Predict Shapes

1. Draw the Lewis dot structure to determine the total # of electron pairs around the central atom.

2. Multiple bonds (double and triple) count as one.

3. The total number of bonding and nonbonding electron pairs determines the geometry of the electron pairs (one of the 5 basic shapes).

4. Then use bonding electron pairs only to determine the molecular geometry (actual shape of molecule).

Special Rules1. Lone pairs effect geometry more than bonding pairs. 2. NH3 has one lone pair: reduces angle from 109.5 to 1073. H2O with two lone pairs: reduces angle from 109.5 to

105.4. Multiple bonds affect geometry more than single bonds 5. H2C=O (116 instead of 120) 6. H2C=CH2 (117 instead of 120)7. Lone pairs occupy the axial (middle) position for

trigonal bipyramidal structures.

Practice

Determine the geometry of the following:

BeCl2 CO2 BF3 O3 SO2

CH4 PCl3 H2O PCl5 SF4 ClF3 XeF2 SF6 IF5 XeF4

Consult the next 3 pages to help you.

Linear 180o

BeCl2valence e- = 2 + (2 x 7) = 16e-

Cl....

..BeCl....

..

Two Electron Pairs = Linear Molecule

linear 180o

CO2valence e- = 4 + (2 x 6) = 16e-

CO....

.. O....

.. CO..

O..

.. ..

valence pairs on C ignore double bondstwo

single and double bonds same

linearmolecular shape

molecular geometry linear

trigonal planar 120o

SO2

valence e- = 6+ (2 x 6) = 18e-

valence pairs on Sthree

one lone pairmolecular geometry

molecular shape bent

trigonal

S O....

..O....

:

SO....

.. O....

..:

SO...... O

..

..

:

two bonding pairs

< 120o

tetrahedral 109.5o

CH4

valence e- = 4+ (4 x 1) = 8e-

valence pairs on CfourC HH

H

H109.5o

molecular geometry

molecular shape tetrahedral

tetrahedral

tetrahedral 109.5o

NH3

valence e- = 5+ (3 x 1) = 8e-

valence pairs on Nfour N HH

H

:

< 109.5o

molecular geometry

molecular shape trigonal pyramid

tetrahedral

one lone pair

three bonding pairs

bipyramidal 120o and 1800

PCl5

valence e- = 5+ (5 x 7) = 40e-

valence pairs on Pfive

molecular geometry

molecular shape bipyramidal

bipyramidal

P

Cl....

..

Cl....

..

Cl....

..Cl....

..

Cl ......

90o

120o

180o

bipyramidal 120o and 1800

SF4

valence e- = 6+ (4 x 7) = 34e-

valence pairs on Sfive

molecular geometry

molecular shape seesaw

bipyramidal

one lone pair

four bonding pairsS

..F..

..

..F..

..

..F..

....F..

..

:

< 180o

bipyramidal 120o and 1800

ClF3

valence e- = 7+ (3 x 7) = 28e-

valence pairs on Clfive

molecular geometry

molecular shape T

bipyramidal

two lone pair

three bonding pairs

Cl

..F..

..

..F..

....F..

..

::

180o

90o

bipyramidal 120o and 1800

ICl2-

valence e- = 7+ (2 x 7) + e-

valence pairs on Ifive

molecular geometry

molecular shape linear

bipyramidal

three lone pair on I

two bonding pairs

= 22e-

I..Cl..

..

::

..Cl..

..

:

octahedral 90o

BrF5

valence e- = 7+ (5 x 7)

valence pairs on Brsix

molecular geometry

molecular shape square pyramidal

octahedral

= 42e-

one lone pair

five bonding pairs

Br

F....

..F....

..

F....

..

F....

..

F.... ..:

octahedral 90o

XeF4

valence e- = 8+ (4 x 7)

valence pairs on Xesix

molecular geometry

molecular shape square planar

octahedral

= 36e-

two lone pair

four bonding pairs

Xe

F....

..F....

..F....

..

F....

..

:

:

Cl....

..BeCl....

..

SO....

.. O....

..

S O....

..O....

SO...... O

..

..

S

O O

:

CO....

.. O....

.. CO..

O..

.. ..

B

: :

F

:: : F

:::

F: ::

C HH

H

H

N HH

H

:

O HH

::

P

Cl....

..

Cl....

..

Cl....

..

Cl....

..

Cl ......

S

..F..

..

..F..

..

..F..

..

..F..

..

:

Cl

..F..

....F..

....F..

..

::

I..Cl..

..

::

..Cl..

..

:

F.... ..S

F....

..

F....

..

F....

..F....

..

F.... ..

Br

F....

..

F....

..

F....

..F....

..

F.... ..:

Xe

F....

..

F....

..

F....

..

F....

..

:

:

Valence Bond Theory

1. Valence Bond Theory explains why molecules have the shapes they do based on a concept called hybridization.

2. Hybridization - A mixture of two or more atomic orbitals.

http://www.youtube.com/watch?v=PrNbhuB9W44&feature=related

10 Minutes

Valence Bond Theory & Hybridization

1. In 1931, Linus Pauling, proposed that the outermost orbitals of an atom could be combined to form hybrid atomic orbitals. – Sigma bond (s) - The end-to-end overlapping of two orbitals. – Pi bond (p) - The side-to-side overlapping of two p orbitals.

• Single bonds are made up of one sigma bond. • Double bonds are made up of one sigma bond and one

pi bond. • Triple bonds are made up of one sigma bond and two pi

bonds.

Sigma Bonds

The electron density is concentrated on the internuclear line.

More Sigma Bonds

Pi Bonds

The electron density is concentrated above & below the internuclear line.

Molecular Orbital Theory

MO Theory is used to predict: 1) whether or not a molecule exists 2) its bond type (single, double, or triple) 3) its bond strength (triple are strongest) and

4) some of its properties (paramagnetic or diamagnetic).

MO Diagrams & Magnetic Properties

Paramagnetism – Molecules with one or more UNPAIRED electrons in their MO diagram are attracted into a magnetic field.

Diamagnetism – Molecules with PAIRED electrons in their MO diagram are weakly repelled from a magnetic field

Molecular Orbitals

1. Molecular orbitals, like atomic orbitals, have a definite shape and hold up to 2 electrons of opposite spin.

2. When 2 atomic orbitals combine & overlap, they form 2 molecular orbitals.

3. One of these molecular orbitals in a BONDING orbital and the other is an ANTIBONDING orbital.

Bonding molecular orbital - Electrons in this orbital spend most of their time in the region directly between the two nuclei.

Antibonding molecular orbital - Electrons placed in this orbital spend most of their time away from the region between the two nuclei.

Shown below are a sigma (s) bonding molecular orbital and a sigma antibonding ( s *), molecular orbital.

* = Antibond

(a) Shows sigma MOs(b) & (c) Shows pi MOs

Molecular Orbital DiagramSmall 2s-2p Interaction

Use for Oxygen, Fluorine, & Neon

(Reverse order for o 2p and p 2p for all other others)

Bond Order

Bond order Type of Bond

0 Molecule doesn’t exist

1 Single bond

2 Double bond

3 Triple bond

Molecular Orbital Diagram

Use with the Bond Order Formula to determine if a molecule exists, its type of bond, & whether it is paramagnetic or diamagnetic.

(Small 2s-2p Interaction) (Large 2s-2p Interaction) Use for Oxygen, Fluorine, & Neon Use for Boron, Carbon, Nitrogen

Molecular Orbital Diagrams - Steps to determine bond type1) Sum valence electrons for both atoms2) Place arrows (represent electrons) in center of the diagram3) Determine Bond Order = (# bonding e- - # antibonding e-)/24) ID Bond Type (1 = single, 2 = double, 3 = triple)5) ID as paramagnetic (unpaired e- = attract magnet) or diamagnetic (all paired e- = repel magnet)

Use for Oxygen, Fluorine, & Neon Use for Boron, Carbon, Nitrogen

• http://www1.teachertube.com/viewVideo.php?video_id=57124

• A Little MO Theory • 14 minutes