ch.9 - molecular geometry & bonding...

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CHEMISTRY - BROWN 13E

CH.9 - MOLECULAR GEOMETRY & BONDING THEORIES

CONCEPT: ELECTRONIC GEOMETRY

When drawing a compound you have to take into account two different systems of geometrical shape.

• The simpler system known as electronic geometry or __________ shape treats lone pairs (nonbonding electrons)

and surrounding elements as the same.

Key: A = Central Element

X = Lone Pairs and Surrounding Elements

O C O H C N

AX2 = Linear (2 Groups)

AX3 = Trigonal Planar (3 Groups)

F

BF F

SnF F

AX4 = Tetrahedral (4 Groups)

NH

HH C

Cl

ClCl

Cl

Cl

PCl

ClCl

Cl

AX5 = Trigonal Bipyramidal (5 Groups)

Xe FF

Cl

SCl Cl

Cl Cl

AX6 = Octahedral (6 Groups)

Cl

XeH

H

H

H

!

AX

XX

!



PRACTICE: ELECTRONIC GEOMETRY

EXAMPLE: Draw each of the following compounds and determine their electronic geometries.

PH3 BeCl2

PRACTICE 1: Draw the following compound and determine its electronic geometry.

SBr4


IF3


H2S


PO43-



CONCEPT: MOLECULAR GEOMETRY

When drawing a compound you have to take into account two different systems of geometrical shape.

• With the molecular geometry you treat lone pairs (nonbonding electrons) and

surrounding elements as different.

Key: A = Central Element X = Surrounding Element E = Lone Pair

O C O H C N

3 Groups2 Groups F

BF F

SnF F

AX3 - Trigonal Planar AX2E1 - Bent, Angular or V-ShapedAX2 - Linear

4 Groups NH

HHC

Cl

ClClCl H

OH

AX4 - Tetrahedral AX2E2 - Bent, Angular or V-ShapedAX3E1 - Trigonal Pyramidal

5 Groups

Xe FF

AX5 - Trigonal Bipyramidal

ClP

ClCl

Cl Cl

F Cl F

FAX4E1 - Seesaw AX2E3 - Linear

FS FF

F

AX3E2 - T-Shaped

6 Groups ClSCl Cl

Cl ClClXe

H

H

H

HAX6 - Octahedral AX4E2 - Square Planar

F

SFF F

F

AX5E1 - Square Pyramidal

AX

XX



PRACTICE: MOLECULAR GEOMETRY

EXAMPLE: Draw each of the following compounds and determine their molecular geometries.

PH2 – XeCl2

PRACTICE 1: Draw the following compound and determine its molecular geometry.

OBr2

PRACTICE 2: Draw the following compound and determine its molecular geometry.

SO42-



CONCEPT: IDEALIZED BOND ANGLES

According to the ___________________________________________ (VSEPR) model bond and lone electron pairs will

position themselves around the central element so that they are as far apart as possible.



PRACTICE: IDEALIZED BOND ANGLES

EXAMPLE: Determine the bond angles of each of the following compounds.

CO2 BrF4+

PRACTICE 1: Determine the bond angle of the following compound.

AsCl5

PRACTICE 2: Determine the bond angle of the following compound.

IF3



CONCEPT: POLARITY

Molecules that have _______________ sharing of electrons contain a molecular polarity.

• For these molecules, both _________________ and _________________ determine the molecular polarity. POLARITY RULES TO BEING NON-POLAR:

1) If central element has NO lone pair(s): a. Central element must be connected to the ___________ elements.

b. Central element must be ______________ electronegative than the surrounding elements.

PRACTICE 1: Determine if carbon dioxide, CO2, is polar or nonpolar.

2) If central element has lone pair(s): a. Central element must be connected to the ___________ elements. b. Central element must be ______________ electronegative than the surrounding elements. c. Use dipole arrows to point to the ________ electronegative element. These dipole arrows must cancel out. d. Dipole arrows extend _____________________ lone pairs. These lone pair dipole arrows must cancel out.

PRACTICE 2: Determine if xenon tetrafluoride, XeF4, is polar or nonpolar.



CONCEPT: POLARITY PT.2

Molecules that have unequal sharing of electrons contain a molecular polarity.

• Some molecular shapes are seen as perfect and will always lead to a non-polar molecule overall. POLARITY RULES TO BEING NON-POLAR:

1) A perfect shape will be non-polar as long as: a. The central element is connected to the same elements.

b. The central element is less electronegative than the surrounding elements.

O C O H C N

3 Groups2 Groups F

BF F

SnF F

AX3 - Trigonal Planar AX2E1 - Bent, Angular or V-ShapedAX2 - Linear

4 Groups NH

HHC

Cl

ClClCl H

OH

AX4 - Tetrahedral AX2E2 - Bent, Angular or V-ShapedAX3E1 - Trigonal Pyramidal

5 Groups

Xe FF

AX5 - Trigonal Bipyramidal

ClP

ClCl

Cl Cl

F Cl F

FAX4E1 - Seesaw AX2E3 - Linear

FS FF

F

AX3E2 - T-Shaped

6 Groups ClSCl Cl

Cl ClClXe

H

H

H

HAX6 - Octahedral AX4E2 - Square Planar

F

SFF F

F

AX5E1 - Square Pyramidal

EXAMPLE 1: Determine if silicon tetrachloride, SiCl4, is polar or nonpolar. EXAMPLE 2: Determine if phosphorus trihydride, PH3, is polar or nonpolar.



PRACTICE: POLARITY PRACTICE 1: Determine if the following compound is polar or nonpolar.

a. SiBr42- PRACTICE 2: Determine if the following compound is polar or nonpolar.

a. H2S PRACTICE 3: Determine if the following compound is polar or nonpolar.

a. PCl2F3 PRACTICE 4: Determine if the following compound is polar or nonpolar.

a. IF2 –



CONCEPT: HYBRIDIZATION

Covalent bonds are formed when atomic orbitals on different atoms overlap and electrons are shared.

Cl Cl

H H H H

ClCl

But what happens when we need to form covalent bonds with different atomic oribitals, for example BeCl2?

Be Cl

[He]2s2 [Ne]3s23p5

2

2p 2s 2p2s

Promotion

& Hybridization2p sp 2p2s

Be



PRACTICE: HYBRIDIZATION

EXAMPLE: For each of the given covalent compounds draw out the Lewis Structure and answer the questions

CH2Cl2 Hybridization: XeCl5+ Hybridization:

Unhybridized Orbitals: Unhybridized Orbitals: Bonding orbitals (C – H): Bonding orbitals (Xe – Cl):

PRACTICE 1: For the given covalent compound draw out the Lewis Structure and answer the questions.

IF2– Hybridization:

Unhybridized Orbitals: Bonding orbitals (I – F):

PRACTICE 2: For the given covalent compound draw out the Lewis Structure and answer the questions.

CH3+ Hybridization:

Unhybridized Orbitals:

Bonding orbitals (C – H):



CONCEPT: MO THEORY In the molecular orbital theory electrons are seen as being _______________, or spread out over a molecule instead of concentrated in a covalent bond.

• A(n) _______________ orbital is the region of high electron density between nuclei where a bond forms.

• A(n) _______________ orbital is the region that has zero electron density (a node) between the nuclei where a

bond can’t form.

EXAMPLE: Use a MO diagram to write the electron configuration of each of the following:

a. C22-

b. F2+



PRACTICE: MO THEORY

The MO diagram can be connected to the MO bond order:

Bond Order = 12

(# of e – in bonding MO – # of e – in anti-bonding MO)

A bond order __________________________ zero means that the compound is stable and exists.

A bond order __________________________ zero means the compound is unstable and does not exist.

• In general, the _______________ the bond order, the _______________ the bond.

PRACTICE: Use a MO diagram to determine if the following compound exists or not.

a. O22-

b. B2-



CONCEPT: HETERONUCLEAR DIATOMIC MOLECULES Molecular Orbital Theory can also be applied to heteronuclear diatomic molecules, which are composed of two different elements covalently bonded together. The key differences it has from a homonuclear diatomic molecule include:

The _________ electronegative element determines the particular MO diagram used.

The more electronegative element will possess atomic orbitals that are ________ in energy.

The _________ energy orbital contributes more to the bonding molecular orbital, while the

__________ energy orbital contributes more to the antibonding molecular orbital.

s2s

s*2s

p2p

s2p

p*2p

s*2p

Atomic Orbitals

Atomic Orbitals

Molecular Orbitals

s2s

s*2s

p2p

s2p

p*2p

s*2p

Atomic Orbitals

Atomic Orbitals

Molecular Orbitals

2s2s

2p

2p

Hydrogen Nitrogen Oxygen Neon

2s2s

2p

2p

EXAMPLE: Using your knowledge of molecular orbital diagrams, determine the bond order of the NO– ion.



ch.9 - molecular geometry & bonding...

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