Ch. 16 Covalent Bonding VSEPR Theory, Polarity, and using Electronegativity

Ch. 16 Covalent Bonding

VSEPR Theory, Polarity, and using Electronegativity

Covalent Bonds– Forms when 2 atoms share a pair of valence e-

A. Types of Covalent Bonds1. Single Covalent Bond – two atoms share one pair of electrons

Ex: F2

F F● ●

●●

●●

● ●●●●

●● ●F F● ●

●●

●●

● ●●●●

●● ●or F F

● ●●●

●

●●●●

●● ●

Unshared pair – e- not shared between atoms

What makes this bonding work?Atoms have 8 e- in their outer level to make them stable

Ex: H2

H● H● H ● H● or H H

Why does H2 only need 2 e- to be stable?first energy level only contains 2 e-

Covalent Bonds (cont.)

Covalent Bonds (cont.)

2. Double Covalent Bond – 2 pairs of electrons are shared between atomsEx: O2

O●●

●●●

●●O

●●●●●

●O● ●

● O or●●●

● ● ●

● ●O● ● O

● ●

● ●●

●

Covalent Bonds (cont.)

3. Triple Covalent Bond – 3 pairs of electrons are shared between atomsEx: N2

NN●●

●●

●N●●

●●

●N●

● N●● or N●

●●●● ●● ●● ●

Covalent Lewis Dot Structures

1. Determine the # of valence e- in each atom in the molecule

(# valence e- = roman numeral for group A atoms)

2. The central atom is often the first atom written & is usually the atom with the least # of e-. (Exception – H can’t be the central atom). This is going to be the atom that needs to share the most electrons.

Lewis Dot Structures for Compounds

3. Place the electrons around the atoms so each is stable (8 around it, except H – only 2)

Examples:1. Br2

Br● ●

● ●●

●

●Br● ●

●●

●

●

●Br● ●

●●

●

●Br● ●

●

●●

●

2. NH3

N●

● ●

●

●

H●

H●

H●

N● ●

HH

H

3. CO2

C●

●

●

●

O

O● ●

● ●●●

●●●

●

●●

CO O● ●

●●

● ●

●●

4. CCl4

C ●

●

●

Cl●

Cl●

Cl●C ClCl

Cl

5. H2O

H●

O● ●●

●●

●

●

Cl

● ●

● ●● ●

● ●

● ●● ●

● ●

●●

●●

●●

●●●

● ●

Cl● ●

● ●● ●

●●

● ●

● ●

●●

● ●

●●

●●

●●

●●

H●

O

H

H● ● ●

●

Covalent Bond Practice Problems:

1. CH4 4. OF2

2. H2 5. CHI3

3. PH3 6. CO2

VSEPR Theory

• Explains the shapes of molecules. • The VSEPR theory states: b/c electrons repel each

other, molecules adjust their shapes so that the valence e- pairs are as far apart from each other as possible.

Shape Formula Bond Angle ElectronsLinear AX2 180o 4 shared

0 unshared

Linear AX 180o 1 shared3 unshared

Bent AX2 105o 2 shared2 unshared

Trigonal Pyramidal

AX3 107o 3 shared1 unshared

Tetrahedral AX4 109.5o 4 shared0 unshared

Trigonal Planar AX3 120o 4 shared0 unshared

Contains a double bond

Bond Polarity

Polar Covalent Bond – when 2 atoms are joined by a covalent bond and the bonding electrons are not shared equally

Nonpolar Covalent Bond – when 2 atoms are joined by a covalent bond and the bonding electrons are shared equally

Bond Polarity (cont.)

Differences between polar, nonpolar, and ionic bonds

How do you determine if a bond is polar, nonpolar, or ionic?

Subtract the electronegativities of the bonding atoms (p. 405 in textbook)

Electronegativity Differences & Bond Type

Type of BondElectronegativity Difference Range

Nonpolar Covalent Bond 0.0 – 0.4

Polar Covalent Bond 0.5 – 2.0

Ionic Bond greater than 2.0

Tell if the bonds between the following atoms are polar, nonpolar, or ionic:

H 2.1C 2.5 0.4 Nonpolar

1. Hydrogen and Carbon

2. Oxygen and Carbon

3. Potassium and Chlorine

4. Fluorine and Fluorine

5. Nitrogen and Oxygen

O 3.5C 2.5 1.0 Polar

K 0.8Cl 3.0 2.2 Ionic

F 4.0F 4.0 0.0 NonpolarN 3.0O 3.5 0.5 Polar

Polar Molecule – a molecule with a positive and negative end. Polar bonds must be present.

Polarity of Molecule

Polarity of Molecule (cont.)

It is possible to have polar bonds but not a polar molecule!• Carbon dioxide has 2 polar bonds and is linear.• Bond polarities cancel out b/c they are in opposite directions.

CarbonOxygen Oxygen

Practice:

Write the dot structure of the following molecules – then predict the shape and polarity

1. I2

2. PCl3

3. H2S

4. CHI3

5. SiO2

6. CH2O