1. BONDING Bond types bond energies

2. Types of Chemical Bonds (what holds atoms together):
Covalent – sharing electrons between non metal atoms to form molecules.
Nonpolar – equal sharing
Polar – unequal sharing
Ionic – giving or taking electrons between atoms to form compounds.
Metallic – a “sea of electrons” among metal atoms to form the metal solid.

3. Types of Bonds IONIC COVALENT
Bond Formation
e- are transferred from metal to nonmetal
e- are shared between two nonmetals
Type of Structure
crystal lattice
true molecules
Physical
State
solid
liquid or gas
Melting
Point
high
low
Solubility in
Water
yes
usually not
Electrical Conductivity
yes (solution or liquid) no
Other
Properties
odorous

4. METALLIC e- are delocalized among metal atoms “electron sea” solid
Bond Formation
e- are delocalized among metal atoms
Type of Structure
“electron sea”
Physical
State
solid
Melting
Point
very high
Solubility in
Water no
yes (any form)
Electrical Conductivity
Other
Properties
malleable, ductile, lustrous

5. Ionic Bonding - Crystal Lattice
Types of Bonds
Ionic Bonding - Crystal Lattice

6. Covalent Bonding - True Molecules
Types of Bonds
Covalent Bonding - True Molecules
Diatomic Molecule

7. Metallic Bonding - “Electron Sea”
Types of Bonds
Metallic Bonding - “Electron Sea”

8. NaCl CO2 CHEMICAL FORMULA IONIC COVALENT Formula Unit Molecular

9. Chemical Bond
attractive force between atoms or ions that binds them together as a unit
bonds form in order to…
decrease potential energy (PE)
increase stability

10. Bond Polarity
Most bonds are a blend of ionic and covalent characteristics.
Difference in electronegativity determines bond type.

11. Bond Polarity Electronegativity
Attraction an atom has for a shared pair of electrons.
higher e-neg atom  -
lower e-neg atom +

12. Bond Polarity
Electronegativity Trend
Increases up and to the right.

13. Bond Polarity Nonpolar Covalent Bond e- are shared equally
symmetrical e- density
usually identical atoms

14. + - Bond Polarity Polar Covalent Bond e- are shared unequally
asymmetrical e- density
results in partial charges (dipole) + -

15. Bond Polarity
Nonpolar (Pure)
Polar
Ionic
View Bonding Animations.

16. Bond Polarity 3.0-3.0=0.0 Nonpolar 3.0-2.1=0.9 Polar 3.0-0.9=2.1 Ionic
Examples:
Cl2
HCl
NaCl
=0.0
Nonpolar
=0.9
Polar
=2.1
Ionic

17. MOLECULAR POLARITY When does a MOLECULE become polar??
Write some examples:

18. ATTRACTIONS BETWEEN ATOMS
INTRAMOLECULAR FORCES
Covalent Bonds
Ionic (Electrostatic) Bonds
Metallic Bonds
Examples:

19. 8.4 in TEXTBOOK: Attractions between Molecules
INTERMOLECULAR FORCES
Van der Waals Forces – weakest attractions
Result of dipole interactions – when one polar molecule “lines up” to another molecule
Dispersion forces – occur between ALL molecules (polar or not) and are momentary
Hydrogen Bonds – strongest type of intermolecular force
Always involves hydrogen
An attraction to a hydrogen atom ALREADY bonded to another strongly electronegative atom
Hydrogen becomes strongly polarized with a bond and will therefore try to compensate when close to another molecule

22. BOND ENERGIES

23. Energy of Bonding CD + 100kJ  C + D Reaction 1(endo): Endothermic –
heat is a reactant
heat is absorbed
breaking a bond
Reaction 1(endo):
CD + 100kJ  C + D

24. Reaction 2 (exo): A + B  AB + 100kJ
Energy of Bonding
Exothermic –
heat is a product
heat is released
making a bond
less potential energy
bond is more stable
Reaction 2 (exo): A + B  AB + 100kJ
Reaction 3: E + B  EB + 100kJ
Reaction 4: F + G  FG + 400kJ

25. Which bond is more stable, EB or FG?
Reaction 1(endo): CD + 100kJ  C + D
Reaction 2 (exo): A + B  AB + 100kJ
Reaction 3: E + B  EB + 100kJ
Reaction 4: F + G  FG + 400kJ
Which bond is more stable, EB or FG?
FG, because it takes more energy (300kj more) to break the bond than EB.

26. Energy of Bond Formation
Bond Energy
Short bond = high bond energy