1. Chapter 6
Sections 6.1 – 6.4

2. 6.1 – Chemical Bonds Chemical Bond = A link between atoms
Why does it occur?
The nucleus of one atom is attracted to the electrons of another.

3. Types of Bonds (an overview) (You will see all of these again later in the chapter!)
Ionic Bond
Ion = Atom which has gained or lost electron(s)
Metal =
-LEFT side of Periodic Table
-Weak nucleus / Low Electronegativity
-LOSERS of electrons
-Become + charged ions

4. Types of Bonds (an overview)
Ionic Bond
Nonmetal =
-RIGHT side of Periodic Table
-Strong nucleus / High Electronegativity
-GRABBERS of electrons
-Become - charged ions

5. Types of Bonds Ionic Bonds
Atoms gain or lose valence electrons to become a NOBLE GAS CONFIGURATION
Examples:
Na =
Na ION =
Cl =
Cl ION =

6. Types of Bonds Ionic Bonds
Ionic bond = A chemical bond between a cation (+) and an anion (-). Caused by a TRANSFER of electron(s).
Usually a metal + a nonmetal

7. Types of Bonds (an overview) (You will see all of these again later in the chapter!)
2. Covalent Bond = A bond caused by a SHARING of electrons
Usually a nonmetal + a nonmetal
Nonpolar Covalent = Equal sharing of the electrons. Atoms are close in strength
Polar Covalent = Unequal sharing of the electrons. One atom is a little bit stronger than the other
World of Chemistry; #8 Chemical Bonds; End at 16:25 – formation of Hydrogen molecule

8. Types of Bonds How do you tell which type of bond it is?
-By ELECTRONEGATIVITY
A chart of electronegativity will be provided to you:
-The greater the difference in electronegativity – the more ionic the bond.
-Electrons spend more time closer to the element with higher electronegativity.

9. Types of Bonds
If the ABSOLUTE VALUE of the electronegativity difference is:
GREATER THAN 1.7 = IONIC Bond
LESS THAN 0.3 = NONPOLAR COVALENT Bond
0.3 – 1.7 = POLAR COVALENT Bond
Examples:

10. Types of Bonds (an overview)
3. Metallic bond
Usually metals only
-The metal gives up valence electrons.
-Electrons are free to move about.
Atom
Electron Sea

11. More Detail on the Bond Types 6.2 – Covalent Bonds
Covalent Bond = A sharing of electrons
Molecule = A group of atoms held by covalent bonds (ex – water)
Diatomic Molecule = Molecule with only 2 atoms
Molecular Compound = Compound made of molecules
Molecular Formula = The type and number of atoms in a molecule (ex – H2O)

12. Formation of Covalent Bonds
Sharing electrons in a covalent bond makes the atoms more stable and decreases the energy of the atoms. Energy is released when a bond is FORMED.
Overlapping of Orbitals – Example H2: + H H H2

13. The Octet Rule
Atoms in a compound obtain the electron configuration of a NOBLE GAS to gain stability

14. Drawing Lewis Structures
-A picture of the covalent bonds in a molecule
-Connect valence electrons with LINES
-For Academic classes, atoms follow the octet rule unless stated
Examples:

17. Single Bond = 1 pair of electrons (2 e-s total) shared between two atoms
Double Bond = 2 pairs of electrons (4 total e-s) shared between two atoms
Triple Bond = 3 pairs of electrons (6 total e-s) shared between two atoms
*Single bonds are the LONGEST in length; Triple are the SHORTEST
*Single bonds have the LOWEST bond energy; Triple have the HIGHEST

18. More Detail on the Bond Types 6.3 – Ionic Bonds
Ionic Bond = Bond formed by the attraction of a cation (lost electrons) to an anion (gained electrons)
Crystal Lattice = 3-Dimensional network of ions
Formula Unit = Simplest ratio of ions
NaCl

19. Dot structures for Ionic Compounds:
-Will reach noble gas configuration
-Draw an ARROW to show the transfer of e-
-Draw as many of each ion as needed
Examples:

20. Comparison of Ionic and Molecular
Molecular Compounds
IONIC Compounds
Bond Type
Covalent Bonds
Ionic
Structure
Individual Molecules
Crystal Lattice
Strength of Bond
Strong Bonds
Very Strong Bonds
mp/bp
Low mp / bp
High mp / bp
Drawing
Connect Dots
Arrows with Charges
Other
Conduct electricity when melted or in water

21. Drawing the Pictures – when you’re not told the TYPE of substance
Do electronegativity difference first!!
Examples:

22. 6. 4- Metallic Bonding
Metals have LOW electronegativity – Will LOSE electrons
The steps:
-Donate valence electrons to electron sea
-Electrons free to move about
-All electrons in sea are shared by all atoms

23. 6.4 – Metallic Bonding Properties of Metals:
Good conductors of heat – e- sea shakes
Good conductors of electricity – e- in sea can move
Malleable – atoms can be pushed closer
Ductile – atoms can be pushed closer
Luster – light bounces off e- sea

24. Chapter 6
Section 6.5

25. Putting Partial Charges on Molecules

26. Properties of MOLECULAR Compounds
VSEPR = Valence Shell Electron Pair Repulsion Theory
Valence electrons move as far away from each other as possible
Draw Lewis Structure
Look at Central Atom
Count electron areas (bonds + lone pairs)
Use chart info
*Academic will be given a WORD BANK and the option to use Model Kits*

27. VSEPR http://cheminfo.chem.ou.edu/~mra/jmol/jmol.php Areas Bonds
Lone pairs
Shape
Bond Angles
Structure
Example 2
Linear
180o
BeCl2;CO2 3
Trigonal Planar
120o
BF3 1
Bent
<120o
SO2 4
Tetrahedral
109.5o
CH4
Trigonal Pyramidal
<109.5o
NH3
H2O

28. VSEPR
Examples:

29. VSEPR
Examples:

30. VSEPR
Examples:

31. Hybridization
Combination of equal energy orbitals to form new orbitals which all have the same shape and energy
Carbon:
BECOMES C C
1s22s22p2
four sp3 hybrid

32. Types of Molecules
FIRST – Draw Lewis Structure & Include Partial Charges!
1. Dipole = Molecule with overall charge
2. NonPolar With Polar Sites = Molecules with area of charge which cancel out
3. Nonpolar = Molecule with no areas of charge

33. Types of Molecules How do you tell the difference?
-Ask yourself these questions…
1. Is there charge on the molecule?
Yes
No = Nonpolar
2. Can it be sliced?
YES = Dipole
NO = NPWPS

34. Examples

35. Intermolecular Forces
AKA – EXTERNAL BONDS
The attraction BETWEEN Molecules
Types of External Bonds:
Dipole-Dipole Interactions
-Occur due to attraction between partial charges
-Occur between two dipoles (Fix notes) – the strongest external bond
Hydrogen Bond = External bond that involves a hydrogen atom
d -
d +

36. Intermolecular Forces
London Force
-Occurs between Nonpolar (or Nonpolar With Polar Sites) molecules – CHANGE THIS IN THE NOTES!!
-Very weak connection (nonpolar to nonpolar is the weakest)
The Steps: (only needed for Honors)
Electrons in one molecule shift instantaneously to one side
Instantaneous charge results
Electrons in another molecule are repelled
Very weak attraction results

37. Properties Based on Number / Strength of External Bonds
1. State of Matter
s>l>g
2. Evaporation (*volatility)
slow>fast
3. Thickness (*viscosity)
thick>thin
4. Wetness (*adhesion)
To feel wet the substance must bond to your skin (to the Na+Cl-)
5. Dissolving
LIKE DISSOLVES LIKE

38. Demonstrations