• To complete the valence or outer energy level
• Octet Rule – 8 is great!
• A bond forms when 2 atoms attract the same electron or electron pair

• Attractive Forces pull atoms together:
- (+ ) Nucleus of one atom attracts (–) electrons of another atom
- Opposite charges attract = ELECTROSTATIC ATTRACTION
• Bond Energy
- When bonds form Energy is released = Exotherimc
- When bonds break Energy is absorbed = Endothermic
- When 2 atoms are held together by a chemical bond they are at a LOWER energy condition than when they are separate!

• Bond Stability
- Compounds with LOW energy are STABLE
- Compounds with HIGH energy are UNSTABLE
- Chemical Change = the rearrangement of atoms to form low energy and stable compounds
• Bond Strength
- Weak Bond – when a small amount of energy is released
- Strong Bond – when a large amount of energy is released

• Exothermic reactions are favored (bond formation)
UNSTABLE
REACTIVE
HIGH ENERGY
WEAK BOND




STABLE
UNREACTIVE
LOW ENERGY
STRONG BOND

• An atom’s ability to attract electrons
• Highest EN value = 4.0 = Fluorine (non-metal)
• Lowest EN value = 0.7 = Francium (metal)
• The difference in EN values of two bonded atoms can tell you the bond type

ELECTRONEGATIVITY DIFFERENCE AND BOND TYPE
If the EN Difference is between:
4.0 – 1.7
Ionic Bond
Examples:
BaO
1.6 – 0.2
Polar Covalent
CH 4 Cl 2
0.1 – 0.0
Non-Polar Covalent

• Ionic
• Covalent:
- Polar Covalent
- Non-Polar Covalent
- Coordinate Covalent
• Metallic

• Electrons are TRANSFERRED from one atom to another
• Transfer from a metal to a nonmetal
• Metals lose electrons to the nonmetals (gain e-)
• High EN Diff. – the greater the difference, the stronger the ionic character
• Form ionic solids called Ionic Crystals (crystal lattice)
• Example: Which of the following has the strongest ionic bond?
MgO NaCl LiF

• High Melting Point
• High Boiling Point
• Conducts electricity when dissolved in solutions
• Does not conduct electricity in the solid phase!
• Soluble in water
• Forms crystals in solid phase

• Are ALWAYS IONIC
• Occurs when Hydrogen bonds with a group 1 or 2 element
• Examples: LiH NaH KH
• H behaves as a nonmetal and is written second in the formula

• When 2 atoms SHARE electrons in the same orbital
• Possible to have single, double, triple bonds (1, 2, or 3 shared e- pairs)
• Occurs between 2 nonmetals
• Includes the Diatomic Molecules (B O FINCH)
• 2 types of Covalent Bonds
1. POLAR Covalent – sharing of e- is uneven
2. NONPOLAR Covalent – sharing of e- is even

• There is an uneven distribution of + and – charges due an uneven sharing of electrons in the bond
• EN Difference is between 0.2 – 1.6
• Examples: HCl H2O

• Even distribution of + and – charges due to even sharing of electrons
• EN Diff = 0.0 – 0.1
• ALL Diatomic molecules have nonpolar covalent bonds

• Occurs when only one atom contributes BOTH electrons in forming a bond
• Occurs with the polyatomic ions (see Reference Table E)
• Polyatomic Ions – 2 or more atoms that act as a single unit and have one charge on the entire group
• Ex. NH
4
+ H
3
O + NO
3
SO
4
-

• Can exist as solids, liquids, or gases
• SOFT in the solid phase
• Poor conductors of heat and electricity *Good insulators
• Low melting and boiling points
• Does not dissolve in water
• Soluble in Benzene or Ether (organic molecules)

• In molecular solids, the the covalently bonded atoms are linked into a
GIANT NETWORK throughout the entire solid
• MACROMOLECULES
• EXTENSIVE WEB OF COVALENT BONDS (3-D FISHNET)
• Examples: Diamond, Silicon Dioxide, Asbestos
• PROPERTIES:
- HIGH melting points
- HARD solids
- Poor conductors of heat and electricity

• Bonds between metal atoms
• Valence electrons drift from atom to atom
• “SEA OF MOBILE ELECTRONS”
• Properties:
- Very malleable
- Good Conductors of Heat and electricity
- Has Luster
- Tenacity

Ionic Bonds
• Metal – Nonmetal
• Forms Crystal Compound
• Exists as solids
• EN Diff = 1.7 – 4.0
• High Melting/Boiling Pt
• Conducts electricity when in solution
• Soluble in water
Covalent Bonds
• Nonmetal – Nonmetal
• Forms Molecule
• Exists as solids, liquids, gases
• EN Diff = 0.2 – 1.6 Polar
0.1 – 0.0 NonPolar
• Lo Melting/Boiling pt
• Does not conduct electricity in solutions
• Not soluble in water (dissolves in
Benzene)

Metallic Bonds
• Metal – Metal
• Sea of Mobile Electrons
• Good Conductors of Heat & Elec
• Malleable, Luster
• High melting/boiling points
Network Solids
• Covalently bonded clusters of atoms
• Macromolecule
• Poor Conductors of Heat & Elect
• Hard Solids
• High melting point

• 2 TYPES OF MOLECULES
1. Polar Molecules – molecule has a (+) end and a (-) end
- uneven or asymmetrical distributions of + -
- Bond angle is less than 180 degrees
- also called DIPOLES Ex. H
2
0, NH
3
, HCl
2. NonPolar Molecules – symmetrical, bond angle = 180
- All of the DIATOMIC MOLECULES
- B O FINCH

• All molecules with polar covalent bonds are dipoles EXCEPT:
CF4 CCl4 CO2
• A molecule can have polar covalent bonds but be a NONPOLAR molecule
• Example: Draw Lewis Dot Diagram of CCl4 (Carbon Tetrafluoride), CO2, H2O

• Dipole-Dipole Attraction
- When + end of one molecule is attracted to the – end of another molecule
- Very strong attraction gives a high boiling point due to energy needed to break the force of + - attraction
• Hydrogen Bond
• NOT A BOND
• Occurs when H is covalently bonded to an atom which has a
- Small atomic radius
- Hi Electronegativity
- H is bonded to N, O, F (NH3, H2O, HF)

• WEAK attractive forces
• Van der Waals Forces OR London Dispersion Forces
• Occurs between: Diatomic Molecules (B O FINCH)
NonPolar Molecules
Noble Gases
*Force increases as they get closed together:
Distance btwn molecules decreases OR
Size of molecules increases

• EXAMPLE: NaCl dissolves in H2O
• Hydration of the ions
• Draw diagram