Chemical Bonding

Lewis Theory of Bonding

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 Structures

A 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 ]-

Lewis 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

Lewis Diagrams 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)

Lewis Diagrams for Covalent Compounds

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

FF

F

N

Lewis Diagrams for Covalent Compounds

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

FF

F

N

Lewis Diagrams for Covalent Compounds

Example 2. COCl2 (24 electrons)

C ClCl

O

Lewis Diagrams for Covalent Compounds

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

C ClCl

O

C ClCl

O

Lewis Diagrams for Covalent Compounds

Example 3. Chlorate ion, ClO3-

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

Cl OO

O

:

:

Co-ordinate covalent bonds

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.

Co-ordinate covalent bonds

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 ?

Formal Charge

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

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

Resonance Structures

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.

Bond orderThe 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

Resonance structure

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

Resonance structure

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.

Resonance structure

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

Nature of the chemical bond

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.

Valence Bond Theory

The half filled orbital of one atom overlaps a half filled orbital of a second atom to form a full orbital with two electrons spinning in opposite directions.

The bonding atoms arrange themselves in order to maximize the overlap of the half-filled orbitals.

Maximum overlapping of the orbitals creates a bonding orbital with a lower energy and increased stability.

Valence Bond Theory

The combination of electrostatic repulsion and opposing magnetic fields (due to the electron’s spin) creates the stability associated with a bonding orbital.

Negative charge

Electrostatic repulsion

Negative charge

Electron

Spin

Electron

Spin

“North” magnetic field

“South” magnetic field

Magnetic attraction

Valence Bond Theory During this process, two atoms approach

each other and allow their half filled orbitals to overlap and form the stability of a filled bonding orbital.

Hybrid Orbitals

In some situations a more advantageous bonding scenario can be established by promoting electrons from a full orbital to a similar empty orbital to create two half filled orbitals that are available for bonding.

The resulting orbital is a mixture of the two original orbitals and is called a hybrid orbital.

Hybrid Orbitals

Electrons are promoted from the full “s-orbital” into an empty “p-orbital” which results in hybrid orbitals that have one electron per orbital and characteristics unique to the newly formed orbitals.

The new hybrid orbitals are free to become involved in bonds by overlapping with other half filled valence orbitals.

Hybridization

1s 2s 2p

C (z = 6)

1s 2s 2p

C (z = 6)

1s sp3

C (z = 6)

sp3 hybridization

1s 2s 2p

C (z = 6)

1s 2s 2p

C (z = 6)

1s sp2

C (z = 6)

sp2 hybridization

p

1s sp

C (z = 6)

sp hybridization

p p

sp2 Hybrid atomic orbitals

sp3 Hybrid atomic orbitals

Double & Triple Bonds Sigma () bonds – The “end-to-end” overlapping

of half filled orbitals to make a full bonding orbital of lower energy level (i.e. – more stable) They occur between “s”, “p” and hybrid orbitals (“sp”,“sp2” & “sp3”) to make single covalent bonds.

Double & Triple Bonds Pi () bonds – The “side-to-side” overlapping of

half filled “p” orbitals to make more stable filled bonding orbitals.

Double & Triple Bonds A combination of “” and “” bonds makes

double and triple bonds.

Single bonds Sigma only ()

Double bonds 1 Sigma () and 1 Pi ()

Double & Triple Bonds A combination of “” and “” bonds makes

double and triple bonds.

Triple bonds 1 Sigma () and 2 Pi ()

Pi (bond

Pi (bond

Sigma (bond

Molecular shapes

Valence shell electron pair repulsion (VSEPR) theory

Only valence shell electrons of the central atom are important in the molecular shape.

Valence shell electrons will repel to maximize the distance between the pairs

Bond pairs and lone pairs behave similarly

Molecular shapes

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

-