CHEMICAL BONDINGLEWIS THEORY OF BONDING

Ionic Bonding

Resonance Structures

VSEPRBasic Shapes3-D NotationHybridization (Lab)

M olecular G eom etries

Octet Rule Polar M olecules

Lew is Structures Covalent Bonding

T ypes of Bonds

• Results from the transfer of electrons from a metal to a non-metal.

• A chemical bond between oppositely charged ions

• Held together by electrostatic attraction

IONIC BONDING

• Formed when an orbital from 2 different atoms overlap

• Electrons must have opposite spins

COVALENT BONDING

CHEMICAL BONDS

Bond Type Single Double Triple

# of e’s 2 4 6

Notation — =

Bond order 1 2 3

Bond strength

Increases from Single to Triple

Bond length Decreases from Single to Triple

TYPES OF BONDING CONDITIONS BETWEEN ELEMENTS

Low Electronegativity and low Ionization energy (Metals)

High electronegativity and High Ionization energy (Non-metals)

Low Electronegativity and low Ionization energy (Metals)

Metallic bonding Ionic bonding (transferring of electrons between atoms)

High electronegativity and High Ionization energy (Non-metals)

Ionic bonding Covalent bonding (sharing of electrons between atoms)

Electronegativity

LEWIS DIAGRAMS OR STRUCTURESA convention developed to “show” the relationship between atoms when they form bonds.

Why is it necessary?

• predict where the electrons are in a molecule• needed to predict the shape of a molecule

LEWIS DIAGRAMS FOR IONIC COMPOUNDS

Identify the number of valence shell electrons and determine the charge on the ion using the “stable octet rule”.

Write the elemental symbol, place dots to represent the electrons in the valence shell, enclose in square brackets and write the ionic charge as a superscript.[Na]+ or [ Cl ]-

STRUCTURAL DIAGRAMS FOR COVALENT COMPOUNDS

Draw the Lewis Diagram for nitrogen trifluoride (NF3).

Step 1. Count the valence electrons

N = 5

F = 7

5 + 3( 7) = 26 valence electrons

STRUCTURAL FORMULA FOR COVALENT COMPOUNDS

Step 2. Write a skeletal structure. Use the least electronegative atom in the centre

Electronegativity: N = 3.0 & F = 4.0

FF

F

N = a pair of e-

(a single bond)

Step 3. Complete the octets for each terminal atom (except H)

FF

F

N

Step 4. Assign any additional electrons as lone pairs on the central atom

FF

F

N

Example 2. COCl2 (24 electrons)

C ClCl

O

Step 5. Make multiple bonds where necessary to complete the octets.

C ClCl

O

C ClCl

O

Example 3. Chlorate ion, ClO3

-

((1 x 7) + (3 x 6) + 1) = 26

Cl OO

O

:

:

In some covalent compounds, the bonds between atoms occur because one atom has donated both electrons to the covalent bond. This is called a coordinate covalent bond.

N :

H

H

H H++ N

H

H

H H

+

Nitrogen supplies the two lone pair electrons to this N-H bond. The H+ ion has no electrons.

To determine the number of coordinate covalent bonds – subtract the bonding capacity (lone valence electrons) from the number of bonds the atom has.

N

H

H

H H

+Nitrogen

Bonds 4Bonding capacity 3Coordinate bonds 4-3=1

In some compounds the SCH3U guidelines may not “work”.

On occasion, both elements have the same electronegativity or there may be two or more possible Lewis Structures.

• e.g. CS2

• (both electronegativities = 2.5)• is it S=C=S or C=S=S ?

Exceptions to the Octet RuleExceptions to the Octet Rule

In such situations, one determines the Formal Charge. The option with the lowest formal charge has the most stable and viable structure.

The Formal Charge for an atom is the number of valence electrons in the free neutral atom minus the number of valence electrons assigned to the atom in the Lewis structure.

Formal Charge = (# valence electrons)-(# of bonds)-(# of unshared e-)

C S S

Valence electrons 4 6 6

Electrons assigned 6 4 6

Formal Charge -2 2 0

C=S=S

S C S

Valence electrons 6 4 6

Electrons assigned 6 4 6

Formal Charge 0 0 0

S=C=S

In some structures the Lewis structure does not represent the true structure of the compound.

Bond order is the number of shared pairs of electrons between two atoms. (i.e. – the number of bonds between two atoms)

As the bond order increases. . . • The length of the bond decreases.• The energy associated with breaking the bond

increases.

The number of shared electrons in a bond affects its length and energy.

Bond Type Bond OrderBond Length (pm) (10-12m)

Bond Energy (kJ/mol)

C-O 1 143 351

C=O 2 121 745

C-C 1 154 348

C=C 2 134 615

C C 3 120 812

C-N 1 143 276

C=N 2 138 615

C N 3 116 891

CHO2- is a polyatomic ion with the following

Lewis structure.

The C-O bond lengths are experimentally determined to be between C-O and C=O. The bond order is neither 1 or 2, but considered to be somewhere in between (i.e.-1.5).

The Lewis structure does not support the experimental data.

C

H

O O

-1

The actual structure is a resonance hybrid of the two resonance structures.

C

H

O O

-1

C

H

O O

-1

The resonance structure does not “flip-flop” back and forth between the two. It is a hybrid form of the two.

The bond order for NO3- is,

(1+1+2)/3=1.33.The resonance structure is . . .

N

O

O O

-1

N

O

O O

-1

N

O

O O

-1

C C

C

CC

C

H H

H

H H

H

C C

C

CC

C

H H

H

H H

H

Benzene has a bond order of 1.5.

(1+2+1+2+1+2)/6=1.5

The work on quantum theory in conjunction with the success of Lewis structures resulted in the inevitable connections between the two areas of study.

Linus Pauling , a friend of Gilbert Lewis, connected the two with the valence bond theory.

PRACTICE

COMPLETE THE LEWIS DIAGRAM AND CHEMICAL BONDING WORKSHEETS

e-pairs Notation Name of VSEPR shape Examples

2 AX2 Linear HgCl2 , ZnI2 , CS2 , CO2

3 AX3 Trigonal planar BF3 , GaI3

AX2E Non-linear (Bent) SO2 , SnCl2

4 AX4 Tetrahedral CCl4 , CH4 , BF4-

AX3E (Trigonal) Pyramidal NH3 , OH3-

AX2E2 Non-Linear (Bent) H2O , SeCl2

5 AX5 Trigonal bipyramidal PCl5 , PF5

AX4E Distorted tetrahedral

(see-sawed)

TeCl4 , SF4

AX3E2 T-Shaped ClF3 , BrF3

AX2E3 Linear I3- , ICl2

-

6 AX6 Octahedral SF6 , PF6-

AX5E Square Pyramidal IF5 , BrF5

AX4E2 Square Planar ICl4- , BrF4

-