1. Introduction to Chemical Bonding

2. Why do bonds form?
Atoms are lazy! They want to have the lowest possible potential energy.
As independent particles, atoms have relatively high potential energy.
By bonding, atoms lower their potential energy and become more stable.

3. Types of Chemical Bonding
Chemical bond-a mutual electrical attraction between the nuclei & valence electrons of different atoms that binds the atoms together.
Major types of bonding:
Covalent bonding-sharing of electron pairs
Ionic bonding-electrical attraction between cations and anions
Metallic bonding-sharing of electrons in an “electron sea”

4. Vocab
Molecule-a neutral group of atoms that are held together by covalent bonds
Molecular compound-a compound whose simplest units are molecules
Example-water
Molecular formula-shows the types and numbers of atoms combined in a single molecule
Example-H2O
Diatomic molecule-a molecule containing only two atoms

5. What type of bond will form?
Use electronegativity values to determine
Remember that electronegativity is a measure of an atom’s ability to attract electrons.
Bond Type
EN Difference
Nonpolar covalent
0-0.4
Polar covalent
Ionic
>1.7

6. What type of bond will form?
Nonpolar covalent-a covalent bond in which the bonding electrons are shared equally by the bonded atoms, resulting in a balanced distribution of electrical charge
Polar covalent-a covalent bond in which the bonded atoms have an unequal attraction for the shared electrons
Ionic bond-a bond in which electron(s) are completely transferred from one atom to another, forming ions.

7. Three Possible Types of Bonds
Covalent
Polar covalent
Ionic

8. Comparing Electron Density in a Nonpolar & Polar Covalent Bond

9. What type of bond will form? examples
Bonding between sulfur and
EN difference
Bond type
More-negative atom
Hydrogen
0.4
Polar covalent
Sulfur
Cesium
1.8
Ionic
Chlorine
0.5

10. What type of bond will form? examples
Bonding between chlorine and
EN difference
Bond type
More-negative atom
Calcium
2.0
Ionic
Chlorine
Oxygen
0.5
Polar covalent
Bromine
0.2
Nonpolar covalent

11. 11/26 WELCOME BACK  (19 more school days until Christmas
Pick up the papers from the side table
Nothing to turn in
Take out a piece of paper and something to write with
We will review naming ionic compounds and then learn to name covalent compounds
HW: naming covalent compounds
FYI: Polyatomic Memorization test is 12/13
FYI: Test corrections are AFTERSCHOOL Tuesday and Thursday
RETEST IS NEXT TUESDAY-THURSDAY

12. On the BACK OF THE NOMENCLATURE REVIEW ANSWER THE FOLLOWING:
1. What is an ionic compound?
2. What is a covalent compounds?
3. What are the properties of covalent and ionic compounds?
4. Why do atoms bond?
5. Where are the nonmetals located on the periodic table?
6. ON THE FRONT OF THE NOMENCLATURE REVIEW:
1. label each compound as ionic or covalent NEXT to the number (you can put an “I” or a “C”
2. NAME THE IONIC COMPOUNDS ONLY

13. NAMING Covalent Compounds

14. Vocab
Molecule-a neutral group of atoms that are held together by covalent bonds
Molecular compound-a compound whose simplest units are molecules
Example-water
Molecular formula-shows the types and numbers of atoms combined in a single molecule
Example-H2O
Diatomic molecule-a molecule containing only two atoms

15. Naming Binary Molecular Compounds
Named using a prefix system to indicate how many atoms of each element are present.
Prefixes for Naming Covalent 1
Mono- 2
Di- 3
Tri- 4
Tetra- 5
Penta- 6
Hexa- 7
Hepta 8
Octa- 9
Nona- 10
Deca-

16. Rules for Naming Binary Molecular Compounds
The less-electronegative element is given first. It is given a prefix only if it contributes more than one atom to a molecule. (Names will never start with mono-.)
Order of nonmetals: C, P, N, H, S, I, Br, Cl, O, F
The second element is named by combining a prefix indicating the number of atoms, the root of the name of the element, and the ending –ide.
You may drop an o or a at the end of the prefix if the name of the element starts with a vowel. Example: pentoxide, not pentaoxide.

17. Naming Binary Molecular Compounds Example
Name the following compound: P4O10.
P4O10
Prefix indicating number of atoms contributed by more-electronegative element
Root name of more-electronegative element + -ide. +
Prefix needed if less electronegative element contributes more than one atom
Name of less-electronegative element +
tetraphosphorus
decoxide

18. Naming Binary Molecular Compounds Practice
Name the following compounds:
As2O5
Diarsenic pentoxide
XeF4
Xenon tetrafluoride
CCl4
Carbon tetrachloride

19. Naming Binary Molecular Compounds Practice
Write formulas for the following compounds:
Carbon diselenide
CSe2
Sulfur pentachloride
SCl5
Dihydrogen monoxide
H2O

20. 11/27
Turn in your naming HW
Pick up the Lewis Structure paper and a calculator
Take out something to answer a warm-up on
We will quickly review naming/writing formulas for covalent compounds and then learn how to draw a Lewis structure for covalent compounds.
HW: Lewis structures
FYI: test corrections TODAY and THURSDAY
FYI: retest TUESDAY-THURSDAY of next week
FYI: Covalent test 12/5 & 12/6
FYI: Polyatomic test 12/13

21. Name or the write the formula for the following:
1.As2O5
2. XeF4’
3. MgI2
4. CCl4
5. N3Br6
6. NaOH
7. Carbon diselenide
8. Sulfur pentachloride
9. Dihydrogen dioxide
10. Cesium nitrate
11. Lithium sulfide
12. DO YOU REMEMBER:
Draw the transfer of electrons for lithium and sulfur

22. LEWIS STRUCTURES FOR Covalent Compounds

23. SOME Vocab for Covalent Compounds:
Molecule-a neutral group of atoms that are held together by covalent bonds
Molecular compound-a compound whose simplest units are molecules
Example-water
Molecular formula-shows the types and numbers of atoms combined in a single molecule
Example-H2O
Diatomic molecule-a molecule containing only two atoms. The atom does not exist by itself, it will bond to itself or another atom to form a molecule)
(Br, I, N, Cl, H, O, F)  MEMORIZE THESE

24. Octet Rule All atoms want to be a noble gas. Why?
Noble gases have a full outer energy level. This gives great stability!
When atoms bond, they gain stability by filling their outer energy level through sharing or exchanging electrons.
Octet rule-chemical compounds tend to form so that each atom, but gaining, losing, or sharing electrons, has an octet of electrons in it highest occupied energy level.

25. Exceptions to the Octet Rule
Hydrogen (H)-2
Beryllium (Be)-4
Boron (B)-6
Some atoms want more than 8 electrons. This is called expanded valence, and it involves bonding in the d orbitals.

26. Bonding electron pair in overlapping orbitals
How do atoms achieve the octet rule in covalent compounds: Orbitals Overlap 1s 2s
2px
2py
2pz F
Fluorine atoms 1s 2s
2px
2py
2pz F
Fluorine molecule
Bonding electron pair in overlapping orbitals

27. F F F Lewis Structures Used to represent molecules.
Combine electron dot notations to show sharing of electrons. F F F
Fluorine atoms
Fluorine molecule

28. Parts of a Lewis Structure
Unshared or lone pair-electrons not involved in bonding that belong only to one atom
A line represents two electrons that are shared, bonding the 2 atoms together

29. NAS Method to drawing Lewis Structures
Needed, Available, Shared Method
Determine how many electrons are available by adding up the total number of valence electrons each atom in the molecule has.
Determine how many electrons are needed to satisfy each atoms’ octet. Don’t forget exceptions!
Subtract available from needed. This will tell you how many electrons have to be shared.
Divide the shared electrons by 2 to figure out how many bonds you will need.

30. Helpful Hints
ALWAYS count the electrons that you used. You can’t use more or less than you had available!
If you have a carbon (C) atom, this will always be in the center. Otherwise, usually the least electronegative element will be the central atom.
Hydrogen is NEVER the central atom since it will only bond two atoms!

31. NAS EXAMPLE #1 Draw the Lewis structure for water, H2O.
Always count how many e- you used &make sure it’s not more than you had available!
How many electrons are available (A)?
How many electrons are needed (N)?
How many electrons must be shared (S)?
How many bonds will this molecule have?
Draw it!
H-2x1=2
O-1x6=6
Total=8
H-2x2=4
O-1x8=8
Total=12
Bond=2 e- N A S 12 H O H 8 4
2=2 bonds
Don’t forget the lone pairs!

32. 11/28
Nothing to pick up
Take out your Lewis Structure HW, something to write with, and a calculator if you think you need one
We will continue practicing Lewis structures today and then quickly take notes over VSEPR determining shapes the second part of the period (first period will do this on Thursday)
HW: Lewis Structures
FYI: Test corrections TOMORROW and retest next week Tuesday-Thursday
FYI: Unit 6 Test next Wednesday/Thursday
FYI: Polyatomic test is 12/13
FYI: Mole due the Monday of midterm week

33. Helpful Hints
ALWAYS count the electrons that you used. You can’t use more or less than you had in the calculated available part of the NAS!
If you have multiple atoms, place the ATOM WITH THE FEWEST AMOUNT IN THE MIDDLE of your Lewis Structure
If you have a carbon (C) atom, this will always be in the center. Otherwise, usually the least electronegative element will be the central atom.
Hydrogen is NEVER the central atom since it will only bond two atoms AND IT NEVER HAS LONE PAIRS!

35. NAS Example 2 Draw the Lewis structure for CH3I.
How many electrons are available (A)?
How many electrons are needed (N)?
How many electrons must be shared (S)?
How many bonds will this molecule have?
Draw it! H C H
C-1x4=4
H-3x1=3
I-1x7=7
Total=14
C-1x8=8
H-3x2=6
I-1x8=8
Total=22 H N A S 22 14 8
2=4 bonds

36. Multiple Bonds X-X X=X X≡X
A single line connecting 2 atoms represents 2 electrons being shared and is called a single bond.
Atoms can share 2, 4, or 6 electrons.
2=single bond
4=double bond
6=triple bond
X-X
X=X
X≡X

37. Multiple Bond Example #1
Draw the Lewis structure for ethene, CH2O.
Use NAS to determine how many bonds the molecule will have.
What atom will be the central atom?
Draw it!

38. O H C H Multiple Bond Example 1 N A S 20 12 8 2=4 bonds
Carbon in center, with other atoms arrange around it.
Single bond atoms together. Have you used all of your bonds?
Where can you put the last bond?
Remember octet exceptions! O N A S 20 H C H 12 8
2=4 bonds

39. N N Multiple Bond Example 2 N A S 16 10 6 2=3 bonds
Draw the Lewis structure for nitrogen, N2. N A S 16 N N 10 6
2=3 bonds

40. [ ] O S Resonance Structures Lewis structures for SO2
Which one is correct?
Neither is exactly right. True structure is actually “in between” the two.
This is called resonance, and is indicated by placing brackets around the structures and a double headed arrow between them.
Resonance only occurs when there are multiple bonds! [ ] O S

41. Ex. CO32-
# of valence electrons: = 24

42. LEWIS STRUCTURE REVIEW USING NAS METHOD
LEWIS STRUCTURE REVIEW USING NAS METHOD. SHOW ALL RESONATING STRUCTURES:
Carbon dioxide
HCN
BeH2
BeF3
PF3
Oxygen gas
CF4

43. Molecular Geometry: AKA VSEPR

44. Molecular Structure: The VSEPR Model
Stands for Valence Shell Electron Pair Repulsion
allows us to use electron dot structures to determine molecular shapes
the structure around a given atom is determined primarily by minimizing electron repulsions
bonding and nonbonding pairs of electrons around an atom position themselves as far apart as possible
Steps:
Draw Lewis structure
Count effective electron pairs on central atom (double and triple bonds count as one)
Arrange the electron pairs as far apart as possible

45. VSEPR Theory Bonding pairs - form bonds Lone pairs - nonbonding e-
Types of electron Pairs
Bonding pairs - form bonds
Lone pairs - nonbonding e-

46. Steps to Determine Molecular Shape
Draw the Lewis Diagram.
Tally up e- pairs & atoms bonded to central atom.
Shape is determined by the # of atoms and e- pairs bonded to central atom.
Know the 5 common shapes
& their bond angles!

47. Determining Molecular Shape
A=central atom
B=any atoms bonded to central atom
E=e- pairs belonging to central atom

48. BeH2 Bond angle:180° Formula: AB2 #1: LINEAR
2 atoms bonded to central atom
No lone pairs ON CENTRAL ATOM
Bond angle:180°
Formula: AB2
BeH2

49. H2O #2: BENT Or AB2E BOND ANGLE<120° FORMULA: AB2E2
2 atoms bonded to central atom
1 or 2 lone pairs ON CENTRAL ATOM
H2O
BOND ANGLE<120°
FORMULA: AB2E2
Or AB2E

50. BF3 FORMULA: AB3 #3: TRIGONAL PLANAR BOND ANGLE:120°
3 atoms bonded to central atom
No lone pair ON CENTRAL ATOM
BF3
BOND ANGLE:120°
FORMULA: AB3

51. NH3 FORMULA: AB3E #4: TRIGONAL PYRAMIDAL
3 atoms bonded to central atom
1 lone pair ON CENTRAL ATOM
NH3
BOND ANGLE: 107°
FORMULA: AB3E

52. CH4 #5: TETRAHEDRAL 4 atoms bonded to central atom
No lone pairs ON CENTRAL ATOM
CH4
BOND ANGLE:109.5°
FORMULA: AB4

53. Molecular Geometry Shape Atoms bonded to central atom Lone pairs
Type of molecule
Linear 2
AB2
Bent/angular 1
AB2E
Trigonal-planar 3
AB3
Tetrahedral 4
AB4
Trigonal-pyramidal
AB3E
AB2E2

54. 11/30 HAPPY FRIDAY! Turn in your molecular geometry homework
Pick up the paper from the side table
You will perform an activity with gum drops and the molecular model kits.
Monday you will learn to determine hybridization and dipole moments. If you want to get ahead, look through the Unit 6: Covalent notes online.
FYI: NO HOMEWORK UNLESS YOU DO NOT FINISH THE ACTIVITY TODAY
FYI: Covalent Test next 12/5 & 12/6
FYI: Polyatomic test 12/13
FYI: mole due 12/17

55. F P F F 107° EXAMPLE #1 TRIGONAL PYRAMIDAL PF3
Atoms bonded to central atom? 3
Lone pairs? 1
AB3E
F P F F
PF3
TRIGONAL PYRAMIDAL
107°

56. O C O 180° Example #2 LINEAR CO2 Atoms bonded to central atom?
Lone pairs?
AB2
LINEAR
180°

57. + - H Cl Dipole Moment Direction of the polar bond in a molecule
Arrow points toward the more e-neg atom
H Cl + -

58. Determining Molecular Polarity
Depends on:
dipole moments
molecular shape

59. Determining Molecular Polarity
Nonpolar Molecules
Dipole moments are symmetrical and cancel out.
BF3 F B

60. Determining Molecular Polarity
Polar Molecules
Dipole moments are asymmetrical and don’t cancel .
H2O H O
net
dipole
moment

61. Determining Molecular Polarity
Therefore, polar molecules have...
asymmetrical shape (lone pairs) or
asymmetrical atoms
CHCl3 H Cl
net
dipole
moment

62. “Intermolecular Bonding”
IMF’S
“Intermolecular Bonding”

63. Intermolecular Forces
Intermolecular forces-forces of attraction between molecules
Vary in strength, but generally weaker than bonds
3 types:
Dipole-dipole forces
Hydrogen bonding
London Dispersion forces

64. Br-F F-F Dipole-dipole forces
Forces of attraction between 2 polar molecules
Short-range, act only between nearby molecules
Compound
Boiling point
Bromine fluoride
-20°C
Fluorine
-188°C
Br-F
F-F

65. Hydrogen bonding Very strong type of dipole-dipole force
Found in compounds that contain H-F, H-O, or H-N
Strong electronegativity difference makes bonds highly polar
Partially positive hydrogen is attracted to an unshared pair of electrons on an adjacent molecule.
Hydrogen bonding-intermolecular force in which a hydrogen atom that is bonded to a highly electronegative atom is attracted to an unshared pair of electrons of an electronegative atom in a nearby molecule

66. Hydrogen Bonding in Water

67. London Dispersion Forces
AKA “van der Waals” forces
LDF-the intermolecular attraction resulting from the constant motion of electrons and the creation of instantaneous dipoles
Electrons are always moving around, which sometimes causes uneven electron distribution. This creates temporary “instantaneous” dipoles.