Covalent bonds form when atoms share electrons to complete their outer electron shells. Covalent compounds are typically made of nonmetal atoms. Binary covalent compounds contain only two elements. Their names follow specific rules based on element position in the periodic table. Molecular shape is determined by VSEPR theory, which predicts the geometry that minimizes electron pair repulsions. Polar covalent bonds result from unequal electron sharing between atoms of different electronegativity. Molecular polarity depends on both bond polarity and molecular geometry.

1. 2/14/13
CHM 101
Ch 4: Covalent Compounds
Covalent Compounds
Covalent bonds form when atoms share
electrons to complete octets.
Covalent bonds are typically between two
nonmetal atoms.

2. Naming Covalent Compounds
Binary Covalent Compounds – molecule that contains atoms of only 2 elements
CS2, SO3, PCl5, etc.
General Rules for Naming
1. Element with lower group number is
named first.
2. If elements are in the same
group, element with higher period
number is named first.
3. Second element named as root + -ide.
4. Greek prefixes are used to designate
the number of atoms.
You must know these!

3. Naming Covalent Compounds
What is the name of SO3?
1. The first nonmetal is S sulfur.
2. The second nonmetal is O named oxide.
3. The subscript 3 of O is shown as the prefix tri.
SO3  sulfur trioxide
The subscript 1 (for S) or mono is understood.

4. Naming Covalent Compounds
Give names for the following binary compounds:
1. PCl5 phosphorous pentachloride
2. CS2 carbon disulfide
3. P4S3 tetraphosphorous trisulfide
Give molecular formulas for the following compounds:
1. Dinitrogen monoxide N2 O
2. Selenium hexafluoride SeF6
3. Dichlorine heptaoxide Cl2O7

5. Naming Covalent Compounds
Use the table below to quiz yourself. Use the formula, write the name. Use
the name, write the formula.

6. Covalent Bonding
Main group elements
Usually make the number of bonds necessary to have noble gas
configuration
Noble gases (except He) have 8 valence electrons
Octet Rule: Main group elements gain, lose, or share e- to achieve a
stable e- configuration with 8 valence e- (except H and He – only need
2 valence e- for stability)

7. Covalent Bonding
Lewis Dot Symbol - Representation of the number of valence electrons
in an atom. Usually only used for main group elements. Maximum
number is 8.
X
H He
Li Be B C N O F Ne

8. Covalent Bonding
Cl2 molecule
Cl: 7 valence electrons, needs to make 1 bond to have octet
Cl Lewis Dot Symbol
Lewis Dot Structure Lone pair of e- (unshared)
Cl Cl Cl Cl
Bonding pair of e- (shared)
Lewis Dot Structure – shows
how the atoms of a molecule
are connected. Shows lone
pairs of electrons and bonding Cl Cl
pairs of electrons.
Count the “shared” e- for both atoms

9. Covalent Bonding
Multiple Bonds – Double & Triple
In carbon dioxide, CO2, the C atom shares 4 electrons with
each O atom in a double bond.

10. Covalent Bonding
Multiple Bonds – Double & Triple
In nitrogen molecule, N2, each N atom shares 6 electrons
with the other N atom in a triple bond.

11. Covalent Bonding
We can often predict the number of covalent bonds an atom will form
based on the number of valence electrons.
# Valence e- # Bonds for Octet Example
Group 7(A) 7 1 HCl
Group 6(A) 6 2 H2 O
Group 5(A) 5 3 NH3
Group 4(A) 4 4 CH4
Cl H Cl O H O H

12. Covalent Bonding
Rules for Drawing Lewis Dot Structures
1. Count the total number of valence electrons in the molecule
NF3 N: 5 e-
F: 7e- x 3 = 21 e-
26 val e- total for the molecule
2. Arrange atoms next to each other and connect with bonds
Central atom is usually written first in the formula and has lower group
number
(If atoms are in same group, central atom is from higher period)
26 val e-
F N F - 6 bonding e-
F 20 remaining e-

13. Covalent Bonding
Rules for Drawing Lewis Dot Structures
3. Place lone pairs around each atom to satisfy octet rule, starting with
terminal atoms
20 remaining e-
F N F -20 lone e-
F 0 remaining e-
Make sure each atom has an octet. If it doesn’t, check your work.
Practice: Draw Lewis Structures for CH3Br and H2Se.

14. Covalent Bonding
Rules for Drawing Lewis Dot Structures
IF there are electrons remaining…
4. Place leftover e- on the central atom. It is ok to exceed the octet
(have more than 8 electrons on an atom) if the central atom is in
Period 3 or higher.
SF4
If the central atom has < 8 e-…
5. Change single bonds to multiple bonds (double or triple) using lone
pairs from terminal atoms.
H2CO, CH3COOH, HCN

15. Covalent Bonding
Draw the Lewis structure for formaldehyde: H2CO
1. Count the total number of valence electrons in the molecule
H2CO H: 1 e- x 2 = 2 e-
C: 4e- x 1 = 4 e-
O: 6e- x 1 = 6 e-
12 val e- total for the molecule
2. Place atoms relative to each other and connect with bonds
H can only make 1 bond, so it has to be a terminal atom. C is a
very common central atom (lower group # than O).
H C H 12 val e-
- 6 bonding e-
O
6 remaining e-

16. Covalent Bonding
Draw the Lewis structure for formaldehyde: H2CO
3. Place lone pairs around each atom to satisfy octet rule, starting with
terminal atoms
H can only accommodate 2 electrons, which it has with the bonding
electrons. Never put lone pairs on H.
H C H 6 val e-
- 6 bonding e-
O 0 remaining e-
The central atom (carbon) has < 8 e-
4. Change single bonds to multiple bonds (double or triple) using lone
pairs from terminal atoms.
H C H Carbon and oxygen now both have
“octets”.
O

17. Lewis Structures of Ions
For anions – add electrons to total valence equal to charge
NO3– : N: 5e- x 1 = 5
O: 6e- x 3 = 18
+ 1e- (b/c of ion charge)
24 e- total
- 6 e- bonding
O 18 e-
- 18 e- lone pairs
O N O 0 e-
– Does each atom have an octet?
O
How can we give N an octet?
O N O
For ions, always put the Lewis structure in
brackets and write the charge as superscript.

18. Lewis Structures of Ions
For cations – subtract electrons from total valence equal to charge
NH4+ : N: 5e- x 1 = 5
H: 1e- x 4 = 4
- 1e-
8 e- total
- 8 e- bonding
+
H 0 e-
H N H
H
Examples for you to practice: ClO3–, ClF4+

19. Covalent Bonding
Exceptions to the Octet Rule
1. Fewer than 8 electrons: Molecules with Be or B as central atom are
often electron deficient. Be usually only needs 4 electrons for stability
and B only needs 6.
BeCl2 BF3
2. Odd # of valence electrons:
NO2:
Free radicals: atoms or molecules with unpaired electrons, highly reactive

20. Covalent Bonding
Exceptions to the Octet Rule
3. More than 8 valence electrons: “expanded valence shells”
- only allowed for atoms in Period 3 and beyond
H2SO4
O
H O S O H
O
Example for you to practice: PCl5

21. Covalent Bonding
Resonance Structures
Draw the Lewis Structure for ozone, O3:
+ +
O O
- -
O O O O
There are 2 possible structures only differing in the location of
electrons and the double bond.
Experimental data says that both bonds are identical.
The actual structure of O3 is neither of them, but a composite of the
two, called a resonance hybrid. Each Lewis structure is a resonance
structure of O3.

22. Covalent Bonding
Resonance Structures
+ +
O O
- -
O O O O
Use a double-headed arrow
for resonance structures
• Resonance structures differ only in the assignment of e- pair
positions, not in atom positions.
Draw resonance structures for NO3-:

23. Covalent Bonding
Molecular Shape – VSEPR Theory
The properties of a molecule are heavily influenced by molecular shape.
Valence-Shell Electron-Pair Repulsion (VSEPR) Theory
Bonding and lone pair electrons surrounding a central atom repel each
other. To minimize the repulsions, electron groups are oriented as far
apart as possible.
How many electron “groups” are on a central atom? Each of the following
counts as one electron group:
• Lone pair of e-
• Single bond
• Double bond
• Triple bond

24. Covalent Bonding
Molecular Shape – VSEPR Theory

25. Covalent Bonding
Ideal bond angles
Electron pair geometry – Arrangement of e- groups (bonds and lone pairs)
around the central atom.
Molecular geometry – Arrangement of atoms in space, shape of molecule.
This is different than e- group geometry if lone pairs are present

26. Covalent Bonding
Molecular Shape – VSEPR Theory
If there are 2 electron “groups” on the central atom…
Linear electron pair geometry
Linear molecular geometry
2 atoms attached to a central atom
180 bond angle
Can’t have any lone pairs, so electron
pair geometry is same as molecular
geometry
Examples: BeCl2, CO2, CS2, HCN

27. Covalent Bonding
Molecular Shape – VSEPR Theory
If there are 3 electron “groups” on the central atom and 0 lone pairs…
Trigonal Planar electron pair geometry
Trigonal Planar molecular geometry
3 atoms attached to a central atom
120 bond angle
If 0 lone pairs, electron pair geometry is
same as molecular geometry
Examples: BF3, SO3, NO3–, CO32–

28. Covalent Bonding
Molecular Shape – VSEPR Theory
If there are 4 electron “groups” on the central atom & 0 lone pairs…
Tetrahedral electron pair geometry
Tetrahedral molecular geometry
4 atoms attached to a central atom
109.5 bond angle
Electron pair geometry If 0 lone pairs, electron pair geometry is
Molecular geometry same as molecular geometry
Example
Examples: CH4, SiCl4, SO42–, ClO4–

29. Covalent Bonding
Molecular Shape – VSEPR Theory
If there are 4 electron “groups” on the central atom & 1 lone pair…
Tetrahedral electron pair geometry
Trigonal pyramidal molecular geometry
3 atoms & 1 lone pair
attached to a central atom
109.5 bond angle
Electron pair geometry
Molecular geometry
Example
Examples: NH3, PF3, ClO3–, H3O+

30. Covalent Bonding
Molecular Shape – VSEPR Theory
If there are 4 electron “groups” on the central atom & 2 lone pairs…
Tetrahedral electron pair geometry
Bent (Angular) molecular geometry
2 atoms & 2 lone pairs
attached to a central
atom
109.5 bond angle
Electron pair geometry
Molecular geometry
Example
Examples: H2O, SCl2

31. Covalent Bonding
Covalent Bonds & Electronegativity
Electrons are shared in covalent bonds, but they usually are not shared
equally.
One atom usually pulls the electrons more strongly than the other.
The bond between H and Cl
in HCl is covalent.
Cl pulls the shared electrons
more strongly than H.
This creates a “partially
negative” region around Cl
and a “partially positive”
region around H.
The bond between H and Cl is a polar covalent bond.

32. Covalent Bonding
Covalent Bonds & Electronegativity
The electronegativity value for an atom indicates the attraction of that atom
for shared electrons (in covalent bonds).
Electronegativity strength
usually increases as the size of
an atom decreases.
WHY???

33. Covalent Bonding
Nonpolar Covalent Bonds
A nonpolar covalent bond is an equal or almost equal sharing of electrons.
The atoms involved have almost no electronegativity difference (0.0 to 0.4).
Examples:
Electronegativity
Atoms Difference Type of Bond
N-N 3.0 - 3.0 = 0.0 Nonpolar covalent
Cl-Br 3.0 - 2.8 = 0.2 Nonpolar covalent
H-Si 2.1 - 1.8 = 0.3 Nonpolar covalent
H-C ??? ???

34. Covalent Bonding
Nonpolar Covalent Bonds
Diatomic molecules are molecules that
contain 2 atoms of the same element.
They are the natural state for elements
H, O, N, Cl, Br, I, and F.
These are the only truly nonpolar covalent
bonds.

35. Covalent Bonding
Polar Covalent Bonds
A polar covalent bond is an unequal sharing of electrons.
The atoms involved have a moderate electronegativity difference (0.5 to 1.7).
Examples:
Electronegativity
Atoms Difference Type of Bond
O-Cl 3.5 - 3.0 = 0.5 Polar covalent
Cl-C 3.0 - 2.5 = 0.5 Polar covalent
O-S 3.5 - 2.5 = 1.0 Polar covalent

36. Covalent Bonding
Ionic Bonds
An ionic bond occurs between metal and nonmetal ions, and is a result of
electron transfer from the metal to the nonmetal.
There is a large electronegativity difference (1.8 or more) between the
atoms.
Examples:
Electronegativity
Atoms Difference Type of Bond
Cl-K 3.0 – 0.8 = 2.2 Ionic
N-Na 3.0 – 0.9 = 2.1 Ionic
S-Cs 2.5 – 0.7 = 1.8 Ionic

37. Covalent Bonding

38. Covalent Bonding
Molecular Polarity
A covalent bond is polar if it connects atoms with different electronegativity
values, i.e. H-Cl, C-F, C-Cl, etc.
How do you know if a molecule is polar?
- Bond Polarity
- Molecular Shape

39. Covalent Bonding
Molecular Polarity
Determine the molecular polarity of CO2:
O C O
Each C=O bond is polar, but the molecule is linear and the two
dipoles cancel each other. Therefore, the CO2 molecule is nonpolar.
COS (carbonyl sulfide):
O C S
C and S have the same EN. There is only 1 dipole, pointing towards
O. Therefore, the COS molecule is polar.

40. Covalent Bonding
Molecular Polarity
What about molecular shape?
H2O:
H O H
Is the molecule nonpolar?
NO! The H2O molecule is not linear, it is tetrahedral!
O
H H

41. Covalent Bonding
Molecular Polarity
Indicate the dipole moment (if any) for CF4 and CHF3.

42. Covalent Bonding
Molecular Polarity
Indicate the dipole moment (if any) for CH2Cl2.