1. Unit 6: Chemical Bonding and Intermolecular Forces
Chemistry I

2. 7.1 Ions: Goals 1-5 REVIEW! Valence Electrons
Electrons that occupy the highest energy level of an atom.
The electrons in the outermost orbitals.
Determining the Number of Valence Electrons
All representative elements have a potential for 8 valence electrons
2 in the s
6 in the p orbitals
The number of valence electrons is the Group Number.
Octet Rule
All representative elements will gain or lose electrons to form an octet (8) of valence electrons.
To achieve the electron configuration of a Noble Gas

3. Electron Dot Structures “Lewis Structures”
Show the VALENCE electrons
Table 7.1
H He
Li Be B C N O F Ne
Na Mg Al Si P S Cl Ar

4. Formation of Cations Cation: Atom Ion Noble Gas Positive Charged Ion
Produced when 1 or more valence electrons are lost
Atom Ion Noble Gas
Electron Dot Na
Electron
Configuration
Electron Dot Mg

5. Formation of Anions Anion: Atom Ion Noble Gas Negative Charged Ion
Produced when 1 or more valence electrons are gained
Atom Ion Noble Gas
Electron Dot Cl
Electron
Configuration
Electron Dot O

6. 7.2 Ionic Bonds & Ionic Compounds
The combination of Anions and Cations.
Attracted to each other by electrical forces (opposites attract).
Although they are made of ions, ionic compounds are electrically neutral
Total (+ Charge) = Total (- Charge)

7. Ionic Bond Attraction by electrical forces (opposites attract).
Attraction of Cations to Anions
Hold Ionic Compounds together
Form a Repeating 3-Dimensional Structure
Form Crystals

8. Sample Problem 7.1 Magnesium & Nitrogen K O → Mg N Mg →
Potassium & Oxygen
Magnesium & Nitrogen
K O →
Mg N Mg →

9. Example Problems (pg 203)
10.
11.
Use electron dot structures to determine formulas of the ionic compounds formed when:
potassium reacts with iodine
b. aluminum reacts with oxygen
What is the formula of the ionic compound composed of calcium cations and chloride anions?

10. Properties of Ionic Compounds
Form solid crystal structures based on number and size of Anions to Cations
Crystal Shape?
Generally have HIGH melting points
Can conduct an electric current when melted or dissolved in water.

11. Metallic Bonding Metallic Bonds: Animation
Force of attraction between free-floating valence electrons and the positively charged metal ions
Metal Cations with positive charge
Valence Electrons act as Anions
Animation

12. Properties of Metals
Malleable and Ductile because of the Metallic Bonds.
The valence electrons and Cations can move about each other
Allow the overall structure to bend or give
Malleable: Can be hammered or pressed into shapes - BENT
Ductile: Can be drawn into wires - STRETCHED
Conduct Electricity because of the Metallic Bonds
Valence Electrons are able to move around or flow
Alloys
Metallic Bonding with more than one type of metal cation.
Examples:
Bronze: Copper & Tin
Brass: Copper & Zinc
Steel: Iron & Carbon

13. Chapter 8: Goals 6-13 8.1 Molecules and Molecular Compounds
Covalent Bond:
A sharing of electrons
Usually occurs between 2 or more NON-METALS
Molecule
A group of atoms joined together by covalent bonds
2 types:
Molecular Compounds: 2 or more Elements Covalently bonded
Example: Water H2O
Diatomic Molecule: Element in Molecule form
Examples: Oxygen O2
Hydrogen
Nitrogen
Halogens

14. Ionic vs. Molecular Compounds
lower melting points
Often Liquids or gases
Molecules
Non-metals together
Electrons Shared
COVALENT BONDS
IONIC
High melting points
Crystal Solids
Formula unit
Metal cation + non-metal anion
Electrons given or taken
IONIC BONDS

15. 8.2 Nature of Covalent Bonding Octet rule in Covalent Bonding
shared electrons count for each atoms “octet”
Hydrogen is an exception (only 2 e-)
Single Covalent Bond
one pair (2 e-) are shared
Ex) H2
Double Covalent Bond
two pairs (4 e-) are shared: Considered 1 bond
Ex) O2
Triple Covalent Bond
three pairs (6 e-) are shared: Considered 1 bond
Ex) N2

16. Sample Problem 8.1
Draw the electron dot structure for HCl
H Cl →

17. Example Problems 7. Draw electron dot structures for each molecule.
a. chlorine
b. bromine
c. iodine
9. The following molecules have single covalent bonds. Draw an electron dot structure for each.
a. H2O2
b. PCl3

18. Coordinate Covalent Bonds
Regular Covalent Bond:
Each atom contributes 1 (or more) e- to be shared
Coordinate Covalent Bond:
One atom contributes 2 e- to be shared
The other atom contributes NO e-
Examples of Coordinate Covalent Bonds:
Carbon monoxide
Sulfate (polyatomic ion)

19. Lewis Structure Practice
Polyatomic Ions:
Covalent Molecules with a charge
NH4+ CO3-2
SO HCO3-
Lewis Structure Practice
RESONANCE!
2 or more valid electron dot structures.

20. VSEPR Theory: Molecule Shape
Valence Shell Electron Pair Repulsion
The shapes of molecules depends on:
the number of covalent bonds
the types of covalent bonds (single, double, etc.),
the number of unshared electrons in the central atom.
Bonded Pairs and Lone Pairs REPELL each other!
Demonstration

21. Lone Pairs (non-bonding)
Total # of
Groups of e-
Approximate
Bond Angle
# of
Covalent Bonds
(# of X)
Lone Pairs (non-bonding)
(# of E)
Geometry Name
(VSEPR class)
Shape
Examples 2
180o
linear
(AX2)
CO2 3
120o
trigonal planar
(AX3)
CO32-
NO3- 1
bent triatomic
(AX2E)
SO2 4
109.5o
tetrahedral
(AX4)
CH4
trigonal pyramidal
(AX3E)
NH3
bent
(AX2E2)
H2O

22. VSEPR Shape Information
Double and Triple bonds count as just 1 bond each
Names correspond to a “central” atom
Some molecules have multiple “central” atoms.
Example: H2O2
Double bent

23. Distorted Tetrahedron
Shape
Expanded Octets
# Bonds Central Atom
# unshared e- pairs
Octahedral 6
Trigonal
Bipyramidal 5
Square Pyramidal 1
Distorted Tetrahedron 4
Square Planar 2
T-shaped 3
Bonus

24. Goal 14: Polar bonds – Bond Polarity
Nonpolar Covalent Bonds:
When electrons are shared equally between atoms.
Elements must have very similar electronegativity values
Examples:
O-O bond: (any diatomic element)
N-O bond
Polar Covalent Bonds:
When electrons are shared unequally between atoms.
Elements must have quite different electronegativity values
Examples: H-O bond
Electronegativity Chart: pg 181

26. Bond Classification Probable Bond Type Electronegativity difference
Nonpolar Covalent
Moderately polar Covalent
Very polar Covalent
Ionic >2.0

27. Sample Problem 8.3
Which type of bond will form between each of the following pairs of atoms?
a. N and H c. Ca and Cl
b. F and F d. Al and Cl
Example Problems
29. Identify the bond type 30. Place the following in order of least to most polar
a. H and Br d. Cl and F a. H – Cl c. H – S
b. K and Cl e. Li and O b. H – Br d. H – C
c. C and O f. Br and Br

28. Polar Molecules Not all polar bonds result in Polar Molecules
The effect of polar bond(s)
Part of the molecule has an effective (-) charge
Part of the molecule has an effective (+) charge
Shape determines effect
Example: Water
Not all polar bonds result in Polar Molecules
Polar bonds can cancel out
Symmetrical shapes
Example: Carbon dioxide

29. Attractions Between Molecules
Dispersion Forces:
weak force caused by electron motion
Only force between non-polar molecules
Dipole interaction:
Slightly stronger force due to charges
force between polar molecules
Hydrogen Bonding:
Much stronger Dipole Interaction
Must be polar molecule due to H and (N, O, or F)
WATER! Special Properties due to H. Bonding
High Surface Tension
High Boiling Point