1 chapter 9 the shapes of molecules - mission...

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© 2006 Brooks/Cole - Thomson

Chapter 9 The Shapes of Molecules

Cocaine

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10.1 Depicting Molecules & Ions with Lewis Structures

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The number of covalent bonds can be determined from the number of electrons needed to complete an octet.

Number of Covalent Bonds

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Molecules with Single Bonds

Ammonia, NH3

1. Decide on the central atom; never H.

Central atom is atom of lowest affinity for electrons.

Therefore, N is central

2. Count total valence electrons

H = 1 and N = 5

Total = (3 x 1) + 5

= 8 electrons / 4 pairs

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3. Form a single bond between the central atom and each surrounding atom

H H

H

N

Drawing Lewis Structure

4. Remaining electrons form

LONE PAIRS to complete octet as

needed.

3 BOND PAIRS and 1 LONE PAIR.

Note that N has a share in 3 pairs (6 electrons), while H shares 1 pair.

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Sulfite ion, SO32-

Remaining pairs become lone pairs, first on outside atoms and then on central atom.

Each atom is surrounded by an octet of electrons.

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Molecules with Multiple Bonds

Step 4: Place lone pairs on outer atoms.

Step 5: To fulfill the octet on C, we form

DOUBLE BONDS between C and O.

Formaldehyde H2CO

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Double and even triple bonds are commonly observed for C, N, P, O, and S

H2CO

SO3

C2F4

N2

CO2

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RESONANCE: Delocalized Electron-Pair Bonding

These equivalent structures are called

RESONANCE STRUCTURES. The true electronic

structure is a HYBRID of the two.

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Electron Delocalization

Lewis structures depict electrons as localized either on an

individual atom (lone pairs) or in a bond between two atoms

(shared pair).

In a resonance hybrid, electrons are delocalized: their

density is “spread” over a few adjacent atoms.

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Fractional Bond Orders

Resonance hybrids often have fractional bond orders due

to partial bonding.

3 electron pairs

2 bonded-atom pairs For O3, bond order = = 1½

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Nitrate ion, NO3-

Each atom is surrounded by an octet of electrons.

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Formal Charge: Selecting the More

Important Resonance Structure

Formal charge is the charge an atom would have if all

electrons were shared equally.

Formal charge of atom =

# of valence e- - (# of unshared valence e-

+ ½ # of shared valence e-)

For OA in resonance form I, the formal charge is given by

6 valence e- - (4 unshared e- + ½(4 shared e-) = 6 – 4 – 2 = 0

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Determine Formal Charge

OA [6 – 4 – ½(4)] = 0

OB [6 – 2 – ½(6)] = +1

OC [6 – 6 – ½(2)] = -1

OA [6 – 6 – ½(2)] = -1

OB [6 – 2 – ½(6)] = +1

OC [6 – 4 – ½(4)] = 0

- Formal charges must sum to the actual charge on

the species for all resonance forms.

For both these resonance forms the formal charges sum

to zero, since O3 is a neutral molecule. - Smaller formal charges (positive or negative) are

preferable to larger ones.

- The same nonzero formal charges on adjacent atoms

are not preferred.

Avoid like charges on adjacent atoms.

- A more negative formal charge should reside on a

more electronegative atom.

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Example: NCO− has 3 possible resonance forms: +2 0 -1 -1 0 0 0 0 -1

Resonance forms with smaller formal charges are

preferred. Resonance form I is therefore not an important

contributor. A negative formal charge should be placed on a

more electronegative atoms. Resonance form III is

preferred to resonance form II.

The overall structure of the NCO- ion is still an average of all three

forms, but resonance form III contributes most to the

average.

Choosing the More Important Resonance Form

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Formal Charge Versus Oxidation Number

For a formal charge, bonding electrons are shared

equally by the atoms. The formal charge of an atom may change between resonance

forms.

For an oxidation number, bonding electrons are

transferred to the more electronegative atom. The oxidation number of an atom is the same in all resonance

forms.

+2 0 -1 -1 0 0 0 0 -1

Formal charges

-3 +4 -2 -3 +4 -2 -3 +4 -2

Oxidation numbers

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Exceptions to the Octet Rule Molecules with Electron-Deficient Atoms

Odd-Electron Species

A molecule with an odd number of electrons is

called a free radical.

B and Be are commonly

electron-deficient.

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Exceptions to the Octet Rule

Expanded Valence Shells

An expanded valence shell is only possible for nonmetals

from Period 3 or higher because these elements have

available d orbitals.

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Valence Shell Electron Pair Repulsion theory.

• Most important factor in determining geometry is relative repulsion between electron pairs. In general, the relative repulsion strength are in the order:

LP-LP > LP-BP > BP-BP

Molecule adopts

the shape that

minimizes the

electron pair

repulsions.

10.2 The VSEPR Model

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Electron Group Arrangements & Molecular Shapes

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Figure 10.3

Molecular shape with TWO Electron Groups

(Linear Arrangement)

This key refers to Figures 10.3 through 10.8.

Examples:

CS2, HCN, BeF2

AX2 Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.

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Figure 10.4

Molecular Shapes with THREE Electron

Groups (Trigonal Planar Arrangement)

Examples:

SO3, BF3, NO3–, CO3

2−

AX2E

AX3 Examples:

SO2, O3, PbCl2, SnBr2

Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.

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Factors Affecting Bond Angles

A lone pair repels bonding pairs more

strongly than bonding pairs repel each

other. This decreases the angle between

the bonding pairs.

Nonbonding (Lone) Pairs

Double Bonds

A double bond has greater electron

density than a single bond, and repels the

single bond electrons more than they

repel each other.

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Figure 10.5

Molecular Shapes with FOUR Electron

Groups (Tetrahedral Arrangement)

Examples:

CH4, SiCl4,

SO42–, ClO4

–

AX4

AX3E

Examples:

NH3, PF3

ClO3–, H3O

+

Examples:

H2O, OF2, SCl2

AX2E2


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Figure 10.6

Lewis structures do not indicate molecular shape.

twist to the right twist to the left

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Figure 10.7

Molecular Shapes with FIVE Electron Groups

(Trigonal Bipyramidal Arrangement)

AX5 Examples:

PF5, AsF5, SOF4

Examples:

SF4, XeO2F2

IF4+, IO2F2

–

AX4E

AX3E2 Examples:

ClF3, BrF3

Examples:

XeF2, I3–, IF2

–

AX2E3


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Axial and Equatorial Positions

Where possible, lone pairs in a five electron-group

system occupy equatorial positions.

A five electron-group system has two different

positions for electron groups, and two ideal bond

angles. Equatorial-equatorial

repulsions are weaker than

axial-equatorial repulsions.

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Figure 10.8

Molecular Shapes with SIX Electron

Groups (Octahedral Arrangement)

Examples:

SF6, IOF5

AX6

AX5E

Examples:

BrF5, TeF5–,

XeOF4

AX4E2

Examples:

XeF4, ICl4–

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Figure 10.9

Molecular shapes for central atoms

in Period 2 and in higher periods

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© 2006 Brooks/Cole - Thomson Figure 10.11

The four steps in converting

a molecular formula to a molecular shape

Molecular

Formula

Count all e- groups

around central atom (A). Step 2

Electron-

group

arrangement

Molecular shape

(AXmEn)

Step 4 Count bonding and

nonbonding e- groups

separately.

Lewis

structure

Draw Lewis structure. Step 1

Note positions of any lone

pairs and double bonds. Step 3

Bond

angles

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Structure Determination by VSEPR

Ammonia, NH3

1. Draw electron dot structure

2. Count BP’s and LP’s = 4

H

H

H

lone pair of electronsin tetrahedral position

N

3. The 4 electron pairs are at the

corners of a tetrahedron.

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Structure Determination by VSEPR Formaldehyde, H2CO

1. Draw electron dot structure

The electron pair geometry

is TRIGONAL PLANAR with

120o bond angles.

•

•

C H H

O •

•

•

•

C

H H

O •

•

2. Count BP’s and LP’s at C

3. There are 3 electron “lumps” around C at

the corners of a planar triangle.

The molecular geometry is also trigonal planar.

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Figure 10.12

Molecular Shapes with More than One Central Atom

ethane

CH3CH3

ethanol

CH3CH2OH


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H-C-H = 109o

C-O-H = 109o

In both cases the atom is surrounded by 4 electron groups.


Methanol, CH3OH

Determine H-C-H and C-O-H

bond angles

H

H

••

•• H—C—O—H

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Acetonitrile, CH3CN ••

H

H

H—C—C N

180˚ 109˚ H-C-H = 109o

C-C-N = 180o

One C is surrounded by 4 electron groups

and the other C by 2 electron groups

Determine bond angles

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10.3 Molecular Shape

& Molecular Polarity

A molecule is polar if

- it contains one or more polar bonds and

- the individual bond dipoles do not cancel.

Overall molecular polarity depends on

both shape and bond polarity.

The polarity of a molecule is measured by its dipole

moment (μ), which is given in the unit debye (D).

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Figure 10.13

The orientation of polar molecules in an electric field.

Molecules are randomly oriented. Molecules become oriented when

the field is turned on.


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Bond Polarity, Bond Angle, and Dipole Moment

Compare CO2 and H2O. Which one is polar?

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CH4 … CCl4 Polar or Not?

• Only CH4 and CCl4 are NOT polar. They are the two molecules that are “symmetrical.”

Bond Polarity, Bond Angle,

and Dipole Moment

Molecules with the same shape may have

different polarities.

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Polar or Nonpolar?

• Consider AB3 molecules: BF3, Cl2CO, and NH3.

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Effect of Molecular Polarity on Behavior

Example: cis and trans isomers of 1,2-dichloroethylene, C2H2Cl2

The cis isomer is polar while the trans isomer is not. The

boiling point of the cis isomer boils is 13°C higher than that

of the trans isomer.

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