Name: _______________________________________________ Date: _______________ Period: _____
Chapter 6 – Covalent Bonding Notes
REVIEW
Two Types of Bonds

Ionic:

Covalent:
o Non-polar covalent:
o Polar covalent:
Electronegativity:
Which element can attract electrons the most?
____________________________: covalent bond with greater electron density around 1 of the 2 atoms.
Electronegativity Difference
Bond Type
increasing difference in electronegativity
Examples:
1. CsCl
2. H2S
3. N2
Do you notice a pattern for the combination of elements that are ionic vs. covalent?

Ionic bonds form between:

Covalent bonds form between:
Example: Identify the following as ionic, covalent, or both:
CaCl2
BaSO4
CO2
AlPO4
SO3
H2O
Properties of Covalent Compounds:
1.
2.
3.
4.
5.
6.
Properties of Ionic Compounds:
1.
2.
3.
4.
5.
6.
7.
NAMING COMPOUNDS
Nonmetal – Nonmetal
USE PREFIXES!
1. Change the ending of the second word to _____________
2. No _____________ on the first word
3. Drop any double vowels
Number of
Atoms
Prefixes
Examples:
1.
CO
2.
CO2
2
3.
SO2
3
4.
SO3
4
5.
N2H4
6.
N2O3
1.
disilicon hexafluoride
2.
tricarbon octachloride
3.
phosphorus pentabromide
8
4.
nitrogen monoxide
9
5.
selenium difluoride
10
6.
dihydrogen monoxide
1
5
6
7
EMPIRICAL AND MOLECULAR FORMULAS
__________________________________: a chemical formula that gives the simplest whole-number
ratio of the elements in the formula.
Example: Which of the following is an empirical formula?
CO2
C2O4
N2H4
NH2
_________________________________: a chemical formula that gives the actual number of the
elements in the molecular compound.
Molecular
Empirical
C2H4
C6H12O6
C9H21O6N3
LEWIS STRUCTURES
Octet rule: ________ electrons in the valence shell (filling s and p orbitals) make an atom ____________
Bond formation follows the octet rule… Chemical compounds tend to form so that each atom:
Lewis Dot Diagrams:
 an electron-configuration notation with only ________________________ electrons of an
element shown, indicated by __________________ placed around the element’s symbol.

Tracks the number of ______________________________ electrons

The ________________________________ electrons are not shown
Examples:
Lewis Structures for Compounds
- The pair of dots between two symbols represents a ________________________________.
How many shared pairs does each fluorine have in F2?
-
An unshared pair, also called a _____________________________, is a pair of electrons that is
not involved in bonding and belongs exclusively to one atom.
How many lone pairs does each fluorine have in F2?
-
The pair of dots representing a _____________________________ of electrons in a covalent
bond is often replaced by a __________________________.
* Each dash represents ________ electrons!
Why do atoms share electrons?
Multiple Covalent Bonds:
- Double covalent bond or double bond: covalent bond in which ________________________ of
electrons are shared between two atoms.
o Shown by two side-by-side pairs of dots or by two parallel dashes.
Examples:
-
Triple covalent bond or triple bond: covalent bond in which __________________________ of
electrons are shared between two atoms.
Examples:
Bond Lengths
Bond Length and Bond Energy
 As atomic size _____________________, bond length ____________________, and as a result
bond energy _____________________.

As you increase the number of bonds between two atoms, energy ___________________, while
bond length ____________________.
Examples:
1. Which bond is greater in length: Br2 or F2
2. The HF bond is 570 pm, the H2 bond is 436 pm, which bond requires more energy to break?
3. Which bond would require more energy to break C-C single bond or C=C double bond?
Which bond is longer?
STEPS FOR WRITING LEWIS DOT STRUCTURES:
1. Draw skeletal structure of compound showing what atoms are bonded to each other. Put least
electronegative element in the center.
2. Count total number of valence e-. Add 1 for each negative charge. Subtract 1 for each positive
charge.
3. Complete an octet for all atoms except hydrogen
4. If structure contains too many electrons, form double and triple bonds on central atom as
needed.
Examples:
NF3
CO32-
Resonance Structures:
Example: Review the picture of carbonate above, is there another way to draw the double bonds?
Exceptions to the Octet Rule
VSEPR THEORY
- Lewis dot diagrams are 2D but we live in a 3D world.
- How are molecules actually arranged?
o Follows ______________________________________________________________
(VSEPR)
*SEE ATTACHED VSEPR MOLECULAR GEOMETRY HANDOUT!
Predicting Molecular Geometry
1. Draw Lewis structure for molecule.
2. Count number of lone pairs on the central atom and number of atoms bonded to the central
atom.
3. Use VSEPR to predict the geometry of the molecule.
Example: What are the molecular geometries of SO2 and SF4?
INTERMOLECULAR FORCES
Intermolecular forces:
Intramolecular forces (bonds):
intramolecular forces are much ____________________ than intermolecular forces
What is a dipole?


Dipole – Dipole Forces
_____________________________________________________________________________________
Dipole – Dipole Forces: Hydrogen Bond
Hydrogen bond:
 Special type of dipole-dipole
 Attraction between:
______________________________________________________________________________
Dipole – Induced Dipole Forces
______________________________________________________________________________
London Dispersion Forces
______________________________________________________________________________
Strength of Intermolecular Forces
Hydrogen Bond
Dipole – Dipole
Dipole – Induced Dipole
London Dispersion Forces
POLAR MOLECULES
Ex: Which of the following molecules have a dipole moment? H2O, CO2, SO2, and CH4
Ex: What type(s) of intermolecular forces exist between each of the following molecules? HBr, CH 4, SO2
What does intermolecular forces effect?
Viscosity:

Surface tension:

Boiling Point:

Cohesion:

Adhesion:
