Molecule - Chemical Paradigms

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4.2.3

Drawing Lewis (electron dot) Structures for Molecules

GN Lewis (1875-1946) a professor and Dean of Chemistry at Berkeley University proposed that chemical bonds resulted when two atoms shared a pair of electrons. The Lewis concept allowed for a kind of "electron dot bookkeeping" to show how atoms could share electrons to achieve their quota as a noble gas or an octet of electrons.

Some general rules

No. of valence electrons of an atom = Group number

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

Octet Rule - there is a total of 4 bonding pairs and lone pairs around the central atom.

Those atoms that have more or less than 4 are non octet.

The central atom is the one with lowest electron affinity. It is never H. Electron affinity is defined as the energy released when a gaseous atom gains a electron. Electron affinity increases across a period and down a group.

Some common shapes of molecules (more about how these are determined later)

Formula Lewis (electron dot) structure

Structural formula

(Use molymod models make the structure. Draw in the correct shape)

Cl

2

Chlorine

HF

Hydrogen

Fluoride

NH

3

Ammonia

H

2

O

Water

1

O

2

Oxygen

CO

2

Carbon dioxide

N

2 nitrogen

CH

4 methane

C

2

H

6 ethane

CH

3

Cl chloromethane

NF

3

Nitrogen trifluoride

C

2

H

5

OH ethanol

C

2

H

4 ethene

C

2

H

2 ethyne

2

CO

Carbon monoxide

H

2

S

Hydrogen sulphide

CH

3

CHO ethanal

Some covalent molecules are non octet, they do not obey the octet rule. These molecules either have more or less than four pairs of electrons around the central atom or can have an odd number of electrons. E.g. BF

3

, NO

2

BF

3

Boron trifluroide

NO

2

Nitrogen dioxide

SF

4

Sulfur tetrafuoride

When covalent bonds form we usually think of both atoms in the bond contributing an equal number of electrons to the bond. However, it is not uncommon for one atom to provide both electrons for bond formation. Such donor atoms tend to be nitrogen, oxygen and chlorine and these bonds are called coordinate or dative covalent bonds. E.g. SO

2

, O

3

, Al

2

Cl

6

SO

2

Sulfur dioxide can exist as two possible

Lewis Structures that follow the octet rule.

These two equivalent structures are called

resonance

structures. The true structure is a hybrid of the two. There is also a non-octet version.

3

O

3 ozone

Al

2

Cl

6

Two AlCl

3

molecules join to form the dimer, Al

2

Cl

6

. Two dative covalent bonds are formed when a non-bonding electron pair on a chlorine atom of one

AlCl

3

molecule is shared with the aluminium atom of the second AlCl

3 molecule in the dimer.

More advanced Lewis Structures

HCOOH

Methanoic acid

CH

3

OCH

3

Methoxy methane

CH

3

COCH

3 propanone

Two possible structures

N

3

H

Hydrozoic acid

4

N

2

H

4

Dinitrogen tetrahydride

N

2

H

2

Dinitrogen dihydride

CH

3

N

2

Element

H

Table of Electron Affinities of common Elements

Electron Affinity (kJ/mol)

72.8

Electron Configuration

1s

1

He <0 1s

2

Li

Be

B

C

N

O

F

Ne

59.8

<0

27

122.3

<0

141.1

328.0

<0

[He] 2s

1

[He] 2s

2

[He] 2s

2

2p

1

[He] 2s

2

2p

2

[He] 2s

2

2p

3

[He] 2s

2

2p

4

[He] 2s

2

2p

5

[He] 2s

2

2p

6

5

Octet obeying structure of SO

2

. These two equivalent structures are called resonance

structures. The true structure is a hybrid of the two.

6

7

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