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

2. Covalent Bonds F F F F F F A. Types of Covalent Bonds
Forms when 2 atoms share a pair of valence e-
A. Types of Covalent Bonds
1. Single Covalent Bond – two atoms share one pair of electrons
Ex: F2
Unshared pair – e- not shared between atoms ● ● ● ● ● ● ● ● ● ● ● ● F F F F ● ● ● ● ● ● F F ● ● ● ● ● ● ● ● or ● ● ● ● ● ● ● ● ● ● ● ● ● ●
What makes this bonding work?
Atoms have 8 e- in their outer level to make them stable

3. Covalent Bonds (cont.) H H H H H H Ex: H2● H H H ● ● or ●
Why does H2 only need 2 e- to be stable?
first energy level only contains 2 e-

4. Covalent Bonds (cont.) O O O O O O
2. Double Covalent Bond – 2 pairs of electrons are shared between atoms
Ex: O2 ● ● ● ● ● ● ● ● O O O O O O ● ● ● ● ● ● or ● ● ● ● ● ● ● ● ● ●

5. Covalent Bonds (cont.) N N N N N N
3. Triple Covalent Bond – 3 pairs of electrons are shared between atoms
Ex: N2 ● ● ● ● N N N ● ● N N N ● ● ● ● ● ● ● ● ● ● or ● ● ● ● ● ● ● ●

6. 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.

7. 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 ● ● ● ● ● ● ● ● ● ● ● ● ● ● ● ● ● ●

8. N N H H H H H H O C O C O O 2. NH3 3. CO2 ● ● ● ● ● ● ● ● ● ● ● ● ● ●

9. C C Cl Cl Cl Cl Cl Cl Cl Cl H H O O H H 4. CCl4 5. H2O ● ● ● ● ● ● ● ●

10. Covalent Bond Practice Problems:
1. CH OF2
2. H CHI3
3. PH CO2

11. 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.

12. Shape Formula Bond Angle Electrons Linear AX2 180o 4 shared 0 unshared
Bent
105o
2 shared
2 unshared
Trigonal Pyramidal
AX3
107o
3 shared
1 unshared
Tetrahedral
AX4
109.5o
Trigonal Planar
120o
Contains a double bond

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

14. Bond Polarity (cont.)
Nonpolar Covalent Bond – when 2 atoms are joined by a covalent bond and the bonding electrons are shared equally

15. Differences between polar, nonpolar, and ionic bonds

16. How do you determine if a bond is polar, nonpolar, or ionic?
Subtract the electronegativities of the bonding atoms (p. 405 in textbook)

17. Electronegativity Differences & Bond Type
Type of Bond
Electronegativity Difference Range
Nonpolar Covalent Bond
0.0 – 0.4
Polar Covalent Bond
0.5 – 2.0
Ionic Bond
greater than 2.0

18. 3. Potassium and Chlorine 4. Fluorine and Fluorine
Tell if the bonds between the following atoms are polar, nonpolar, or ionic:
1. Hydrogen and Carbon
2. Oxygen and Carbon
3. Potassium and Chlorine
4. Fluorine and Fluorine
5. Nitrogen and Oxygen
H 2.1
C 2.5
0.4 Nonpolar
O 3.5
C 2.5
1.0 Polar
K 0.8
Cl 3.0
2.2 Ionic
F 4.0
0.0 Nonpolar
N 3.0
O 3.5
0.5 Polar

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

20. 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.
Carbon
Oxygen
Oxygen

21. Practice:
Write the dot structure of the following molecules – then predict the shape and polarity I2
PCl3
H2S
CHI3
SiO2
CH2O