1. Chapters 3 & 4
Chemical Bonding

2. Valence Electrons Outermost electrons
s and p electrons for main group elements
Responsible for chemical properties of atoms
Participate in chemical reactions
Valence Electron
Core Electrons

4. Problems
Write out the electron configurations for the following elements and identify how many core and valence electrons each has. Mg S Br Kr

5. Lewis Dot Structures
LDS: a representation of an atom using its chemical symbol surrounded by dots that signify valence electrons

6. Problems Write the Lewis Dot Structures for the following atoms Li BeBr C N Ne

7. Li: [He]1s1
Na: [Ne]2s1
K: [Ar]3s1

8. Octet Rule
Octet Rule: the tendency for atoms to seek 8 electrons in their outer shells
Natural electron configuration of the Noble Gases
Done by gaining, losing, or sharing electrons
Increases stability
H and He seek a “Duet”

9. Ionic Bonding
Ions: atoms that have a charge due to gain or loss of electrons
Anion: (-) charged atom
Cation: (+) charged atom
Ionic Bond: a bond formed through the transfer of one or more electrons from one atom or group of atoms to another atom or group of atoms

10. Formula Unit

13. Polyatomic Ion + Non-metal Polyatomic Ion + Polyatomic Ion
Ionic Compounds: compounds composed of oppositely charged ions that are held together by their attraction to each other
Metal + Non-metal
NaCl
Metal + Polyatomic Ion
NaNO3
Polyatomic Ion + Non-metal
NH4Cl
Polyatomic Ion + Polyatomic Ion
NH4NO3
Net charge on compound equal to zero

15. Oxyanions SO42- Sulfate SO32- Sulfite PO43- Phosphate PO33- Phosphite
NO3-
Nitrate
NO2-
Nitrite
ClO4-
Perchlorate
ClO3-
Chlorate
ClO2-
Chlorite
ClO-
Hypochlorite

16. Rules For Naming Ionic Compounds
Name the cation by its elemental/polyatomic name
If the metal is a transition metal with a variable charge, indicate its charge with a Roman Numeral in parentheses
Next, name the anion and change its ending to “-ide”
If the anion is polyatomic, do not change the ending to “-ide”
Do NOT use prefixes (mono, di, tri etc.) to indicate how many of each atom are present

17. Problems Write the name for the following compounds: KI MgBr2 Al2O3
FeCl2
CaSO4
Ba(NO2)2
Cu(NO3)2

18. Write the Formula for the following ionic compounds:
Sodium Fluoride
Calcium Sulfite
Calcium Chloride
Iron (III) Oxide
Cobalt (II) Hydroxide
Ammonium Bromide
Ammonium Carbonate
Aluminum Carbonate

19. Iron (II) Chloride
Iron (III) Chloride

20. Covalent Compounds
Covalent Compounds: compounds composed of atoms bonded to each other through the sharing of electrons
Electrons NOT transferred
No + or – charges on atoms
Non-metal + Non-metal
Also called “molecules”
Examples:
H2O
CO2
Cl2
CH4

21. or H-H
Duet or

22. Naming Covalent Compounds
Name the first non-metal by its elemental name
Add a prefix to indicate how many 1 6 2 7 3 8 4 9 5 10
Name the 2nd non-metal and change its ending to “-ide”
Add a prefix to indicate how many

23. Problems Write the name of the following compounds: CO NI3 N2O SF6
B2O3

24. Write the formula for the following compounds:
Phosphorous Pentachloride
Nitrogen Monoxide
Dinitrogen Tetroxide
Tetraphosphorous Decoxide

25. Problems
KCl
Na2S
H2O
SO2
K3PO4
FeCl3
(NH4)2SO4
SCl2
Cu(OH)2
P2O5

26. Sodium Iodide
Aluminum Sulfate
Phosphorous Pentabromide
Magnesium Nitride

27. Naming Acids Acids that do not contain oxygen HCl HF
Begin the name with “hydro”
Name the anion, but change the ending to “-ic”
Add “acid” on the end
HCl HF

28. Acids that contain oxygen
Do not put “hydro” at the beginning
Begin the name with the anion
If the anion has the ending “-ate,” change this to “-ic acid”
If the anion has the ending “-ite,” change this to “-ous acid”
HClO4
HClO3
HClO2
HClO

29. Problems Name the following HBr(g) HBr(aq) HNO2(aq) HNO3(aq) HI (aq)
HI (g)
H2CO3 (aq)
H3PO4 (aq)
H3PO3 (aq)
HCN (aq)

30. Molecular Structures

31. Ball & Stick Models
Space-Filling Models
Water
Methane

32. Ethanol

34. Lewis Dot Structures
Count the total number of valence electrons in the molecule. Ex: PCl3
Use atomic symbols to draw a proposed structure with shared pairs of electrons.
Atoms don’t tend to bond to other atoms of the same element when they can avoid it
Exception: Carbon

35. Place any leftover electrons on the central atom
Place lone pair electrons around each (except H) to satisfy the octet rule, beginning with the terminal atoms
Place any leftover electrons on the central atom
If the number of electrons around the central atom is less than 8, change single bonds to the central atom to multiple bonds (double or triple).
Ex: CH2O

36. Problems
Draw the LDS’s for the following molecules:
Cl2O
C2H4
C2H6O

37. What Things Like To Do Halogens Like to be terminal
Like to have one single bond and 3 lone pairs (non-bonding electrons)
Carbon
Likes to have 4 single bonds and no lone pairs
A double bond counts as two singles
A triple bond counts as three singles
Likes to be central
Likes to bond to other carbons

38. Silicon
Likes to do what carbon does
Oxygen
Likes to have two single bonds and 2 lone pairs
Sulfur
Likes to do what oxygen does
May expand its octet
Nitrogen
Likes to have 3 single bonds and one lone pair

39. Phosphorous
Likes to do what nitrogen does
May expand its octet
Hydrogen
Likes to be terminal with only one single bond
No lone pairs!
Boron
Likes 3 bonds and no lone pairs (sextet)

40. Problems Draw the Lewis Structures for the following molecules: SH2
C3H8
Si2H6
PI3
CH3OH
C2H2
BF3
CCl2O

41. N2H4
CH2OS
C2H6O CO
BrHO

42. Electronegativity
The measure of the ability of an atom to attract electrons to itself
Increases across period (left to right) and
Decreases down group (top to bottom)
fluorine is the most electronegative element
francium is the least electronegative element

44. Electronegativity Scale

45. Types of Bonding Non-Polar Covalent Bond:
Difference in electronegativity values of atoms is 0.0 – 0.4
Electrons in molecule are equally shared
Examples: Cl2, H2, CH4
ENCl = 3.0
= 0
Pure Covalent

46. Polar Covalent Bond:
Difference in electronegativity values of atoms is 0.4 – 2.0
Electrons in the molecule are not equally shared
The atom with the higher EN value pulls the electron cloud towards itself
Partial charges
Examples: HCl, ClF, NO
ENCl = 3.0
ENH = 2.1
3.0 – 2.1 = 0.9
Polar Covalent

47. Ionic Bond: Difference in EN above 2.0
Complete transfer of electron(s)
Whole charges
ENCl = 3.0
ENNa = 1.0
3.0 – 0.9 = 2.1
Ionic

49. Problems
Predict the type of bonding in the following compounds using differences in EN values of the atoms. Indicate the direction of the dipole moment if applicable
KBr HF
BrI FI

50. Valence Shell Electron Pair Repulsion Theory
VSEPR theory:
Electrons repel each other
Electrons arrange in a molecule themselves so as to be as far apart as possible
Minimize repulsion
Determines molecular geometry

56. Defining Molecular Shape
Electron pair geometry: the geometrical arrangement of electron groups around a central atom
Look at all bonding and non-bonding e-’s
Molecular Geometry: the geometrical arrangement of atoms around a central atom
Ignore lone pair electrons

59. 2 e- groups surrounding the central atom
e- pair geometry: linear
MG: linear
AXE designation: AX2E0
A: Central Atom
X: Bonding pairs
E: Non-bonding pairs
Example: BeCl2

60. 3 e- groups 3 Bonds, 0 Lone Pairs 2 Bonds, 1 Lone Pair
e- PG: Trigonal Planar (Triangular planar)
MG: Trigonal Planar
AX3E0
BF3
2 Bonds, 1 Lone Pair
e- PG: Trigonal Planar (Triangular planar)
MG: Bent/angular
AX2E1
GeCl2

61. 4 e- groups 4 bonds, 0 Lone Pairs 3 bonds, 1 Lone Pair
e- PG: Tetrahedral
MG: Tetrahedral
AX4E0
CH4
3 bonds, 1 Lone Pair
e- PG: Tetrahedral
MG: Triangular Pyramidal
AX3E1
NH3

62. 2 bonds, 2 Lone Pairs
e- PG: Tetrahedral
MG: Bent/Angular
AX2E2
H2O

63. Drawing LDS With Correct Geometry

64. Molecular Polarity

66. Problems BF3 CH2O CBr4 CHCl3 CH2Cl2
Draw the 3D Lewis Dot Structures, using wedges and dashes when applicable, for the following molecules and then identify the net dipole, if any.
BF3
CH2O
CBr4
CHCl3
CH2Cl2