1

Ch 8 notes part 2

2

3

4

5

6

7

8

9

10

11

12

13

14

15

16

17

18

19

20

21

22

23

Attachments

8_CovalentBonding.ppsx

CovalentBonding

Bonding models for methane, CH4. Models are NOT reality. Each has its own strengths and limitations.

1

Polar-Covalent bonds

Nonpolar-Covalent bonds

Covalent Bonds

Electrons are unequally shared

Electronegativity difference between .3

and 1.7

Electrons are equally shared

Electronegativity difference of

0 to 0.3

2

Covalent Bonding Forces

Electron – electron repulsive forces

Proton – proton repulsive forces

Electron – proton attractive forces

3

Bond Length and Energy

BondBond typeBond length (pm)Bond Energy(kJ/mol)

C - CSingle154347

C = CDouble134614

C CTriple120839

C - OSingle143358

C = ODouble123745

C - NSingle143305

C = N Double138615

C NTriple116891

Bonds between elements become shorter and stronger as multiplicity increases.

4

Bond Energy and Enthalpy

D = Bond energy per mole of bonds

Energy required

Energy released

Breaking bonds always requires energy

Breaking = endothermic

Forming bonds always releases energy

Forming = exothermic

5

The Octet Rule

Combinations of elements tend to form so that each atom, by gaining, losing, or sharing electrons, has an octet of electrons in its highest occupied energy level.

Monatomic chlorine

Diatomic chlorine

6

The Octet Rule and Covalent Compounds

Covalent compounds tend to form so that each atom, by sharing electrons, has an octet of electrons in its highest occupied energy level.

Covalent compounds involve atoms of nonmetals only.

The term “molecule” is used exclusively for covalent bonding

7

The Octet Rule: The Diatomic Fluorine Molecule

F

F

1s

1s

2s

2s

2p

2p

Each has seven valence electrons

F

F

8

The Octet Rule: The Diatomic Oxygen Molecule

O

O

1s

1s

2s

2s

2p

2p

Each has six valence electrons

O

O

9

The Octet Rule: The Diatomic Nitrogen Molecule

N

N

1s

1s

2s

2s

2p

2p

Each has five valence electrons

N

N

10

Lewis structures show how valence electrons are arranged among atoms in a molecule.

Lewis structures Reflect the central idea that stability of a compound relates to noble gas electron configuration.

Shared electrons pairs are covalent bonds and can be represented by two dots (:) or by a single line ( - )

Lewis Structures

11

Comments About the Octet Rule

2nd row elements C, N, O, F observe the octet rule (HONC rule as well).

2nd row elements B and Be often have fewer than 8 electrons around themselves - they are very reactive.

3rd row and heavier elements CAN exceed the octet rule using empty valence d orbitals.

When writing Lewis structures, satisfy octets first, then place electrons around elements having available d orbitals.

12

Show how valence electrons are arranged among atoms in a molecule.

Reflect the central idea that stability of a compound relates to noble gas electron configuration.

Lewis Structures

13

The HONC Rule

Hydrogen (and Halogens) form one covalent bond

Oxygen (and sulfur) form two covalent bonds

One double bond, or two single bonds

Nitrogen (and phosphorus) form three covalent bonds

One triple bond, or three single bonds, or one double bond and a single bond

Carbon (and silicon) form four covalent bonds.

Two double bonds, or four single bonds, or a triple and a single, or a double and two singles

14

C

H

H

H

Cl

..

..

..

..

Completing a Lewis Structure -CH3Cl

Add up available valence electrons:

C = 4, H = (3)(1), Cl = 7 Total = 14

Join peripheral atoms

to the central atom

with electron pairs.

Complete octets on

atoms other than

hydrogen with remaining

electrons

Make the atom wanting the most bonds central

..

..

..

15

Multiple Covalent Bonds:Double bonds

Two pairs of shared electrons

Ethene

16

Multiple Covalent Bonds:Triple bonds

Three pairs of shared electrons

Ethyne

17

Acetic Acid

Two electrons (one bond) per hydrogen

Eight electrons (four bonds) per carbon

Eight electrons (two bonds, two unshared pairs) per oxygen

18

bondsbrokenbondsformed

HDD

D=-

åå

C

C

H

H

H

H

C

C

H

H

H

H

C

C

H

H

C

C

H

H

H

O

O

H

H

H

C

C

SMART Notebook

Page 1Page 2Page 3Page 4Page 5Page 6Page 7Page 8Page 9Page 10Page 11Page 12Page 13Page 14Page 15Page 16Page 17Page 18Page 19Page 20Page 21Page 22Page 23Attachments Page 1