1. Valence electrons- electrons in the highest energy level (only s & p)
Ch7- Ionic Bonds
Valence electrons- electrons in the highest energy level
(only s & p)
- # of valence electrons corresponds to the group.
H He
Li Be B C N O F Ne
Na Mg Al Si P S Cl Ar
lose electrons to gain electrons to form
form positive ions negative ions called
called cations anions.
(A Negative ION)
All atoms are trying to satisfy the octet rule
- get 8 electrons in outer energy level.

2. Ex1) Write the electron configurations and dot diagrams for:
Na Cl
Write the e.c. & dot diagram for the ions:
Na+ : Cl- :
What is the chemical formula for sodium chloride?

3. Ex1) Write the electron configurations and dot diagrams for:
Na Cl
[Ne] 3s1 [Ne] 3s2 3p5
Write the e.c. & dot diagram for the ions:
Na+ : [Ne] or [Na]+ Cl- : [Ne] 3s2 3p6 or [Cl]-
What is the chemical formula for sodium chloride? NaCl

4. Ex 2) Write the electron configure & dot diagram for:
Al Br
Write the e.c. & d.d. for the ions:
Al Br-1
What is the chemical formula for aluminum bromide?

5. Ex 2) Write the electron configure & dot diagram for:. . . . . . Br . Al . . .
[Ne] 3s23p1 [Ar] 4s23d104p5
Write the e.c. & d.d. for the ions:
Al+3 = [Ne]+3 Br-1 = [Ar] 4s23d104p6 or [Kr]-1
What is the chemical formula for aluminum bromide?
AlBr3

6. Ex 3) Same stuff for:
Atoms: Mg N
Ions:
Formula:

7. Ex 3) Same stuff for:
Atoms:
Ions:
Formula: N . . . . . . . Mg
[Ne]3s [He] 2s22p3
Mg+2 = [Ne]+2 N-3 = [He] 2s22p6 = [Ne]-3 . Mg . . .
Mg3N2 . N . . . Mg . . . . . N
Properties of Ionic Compounds
- solids at room temp
- don’t conduct electricity in solid state
- will conduct electricity if
1 molten (liquid state)
2 dissolved in water
in both these cases the ions are free to move. . . Mg
Ch7 HW#1 .

8. K+1 C+4 Mg+2 O-2 C-4 Ch7 HW#1 1) How many valence electrons for:
a) Potassium b) Carbon c) Magnesium d) Oxygen
2) Dot structures & electron configurations:
K C Mg O
3) Gain or Lose?
K  C  Mg  O 
4) E.C. for ions
K C Mg O-2
C-4
5) Why do nonmetals form anions when reacting?

9. 6) Dot formula to determine formulas:
a) potassium & iodine b) Ca & S
K I Ca S
c) Al & O d) Na & P
Al Na
Na P Na O O Al O
7) Name:

10. - + Ch7.2 – Metallic Compounds - have free floating valence electrons - +
- have free floating valence electrons
- good conductors of heat and electricity
- if electrons enter one end, others will exit other end.
(basis of electric circuits)
- ductility and malleability caused by electrons shielding cations
from each other, even when the metal is smashed or bent.
(ionic solids shatter because like charges get pressed
together, then repel.)

11. Metallic Structures: (common, simple ones)
Body-centered cubes (BCC)
- every atom has 8 neighbors
- Na,K,Fe,Cr,W
Face-centered cubic (FCC)
- every atom has 12 neighbors
- Cu,Ag,Au,Al,Pb
Hexagonal Closest Packing (HCP)
-12 neighbors, but diff shape
- Mg,Zn,Cd
Quiz tomorrow:
1.Balance equation:
2.Mass-mass: 10 grams of ____ reacts with excess ____.
How much ____ produced?
3.Density: Givens: grams and cm3
4.Temp conversion: ˚C = ___K
CH7 HW#2 8-12

12. Ch7 HW#2 8 – 12
8) Why do metals tend to form cations?
9) Electron configurations:
a) Copper (I) Cu+1 Cu: [Ar] 4s2Cd [Ar] 4s13d10
b) Gold (I) Au Au: [Xe] 6s24f145d9 6s14f145d10
c) Cadmium Cd Cd: [Kr] 5s24d10
d) Mercury (II) Hg+2 Hg: [Xe] 6s24f145d10
10) Ductile -
Malleable -
11) Electron configurations for +3 charged:
a) Chromium Cr: [Ar] 4s23d4
b) Manganese Mn [Ar] 4s23d5
c) Iron Fe: [Ar] 4s23d6

13. 12) Use dots to combine:
a) Cu(I) & Cl:
Cu Cl
b) Cu(II) & Cl: Cl
c) Fe (II) & O
Fe O
d) Fe (III) & O Fe O Fe O O

14. Ch7.3 Formulas of Ionic Compounds
1. Elements exchange electrons in ionic bonds.
2. Cations form from metals by the loss of valence electrons.
3. Anions form from nonmetals by the gain of electrons.
4. Ionic bonds form as the result of oppositely charged ions
attracting one another. An ionic compound always contains at least
one positive ion (cation) and one negative ion (anion.) These ions must
combine in such a way as to produce a neutral compound.
5. The formula unit of an ionic compound represents the smallest sample
of an ionic compound that has the composition of that compound.
6. This formula unit will reflect the balance of charges of the compound’s ions.
7. The use of electron dot formulas is a helpful tool in predicting formulas
of ionic compounds.

15. Ch7 HW#3 1 – 12
1. Use the periodic table to find the number of valence electrons in an atom.
a. sodium b. carbon c. phosphorus
2. Draw the electron dot formulas of these representative elements:
K AI O CI
3. Describe the formation of a cation from an atom of a metallic element,
using the octet rule and the importance of noble-gas electron configurations.
Describe the formation of the sodium ion using an electron dot structure.

16. 4. Describe the formation of an atom of a non metallic element.
Describe the formation of the chloride ion using an electron dot structure.
5. List the characteristics of an ionic bond.

17. 6. Explain the electrical conductivity of melted and of aqueous solutions of
ionic compounds, using the characteristics of ionic compounds.
7.Explain the physical properties of metals, using theory of metallic bonding.

18. 8. Write electron dot structures for the following atoms:
a. silicon b. rubidium c. barium d. tin e. iodine f. arsenic
Si Rb Ba Sn I As
9. Complete the following table.
outer electron outer electron formula of ion type of ion
config of atom config of ion Se K Ca Br N

19. 10. Use electron dot formulas to determine chemical formulas of the ionic
compound formed when the following elements combine,
and then name them.
a. strontium and fluorine
b. magnesium and chlorine
c. sodium and oxygen
d. aluminum and nitrogen

20. 11. Determine the formula of the ionic compound formed when barium
and phosphorus combine.
12. What is the formula for the compound formed when astatine and
strontium combine?

21. Lab7.1 – Models
- due in 2 days
- Ch7 Rev WS due at beginning of period

22. . . : Ex1) H + H H H H-H Ex2) F2 Ex3) H2O Ch8.1 – Covalent Bonds
Single Covalent Bonds – one pair of electrons shared between 2 atoms.
Ex1) H + H H H H-H
one shared pair Structural formula
Each dash represents
one shared pair of electrons. . . :
Ex2) F2
Ex3) H2O

23. Ch8.1 – Covalent Bonds
Single Covalent Bonds – one pair of electrons shared between 2 atoms.
Ex1) H + H H H H-H
one shared pair Structural formula
Each dash represents
one shared pair of electrons. . . : .. .. .. .. . . : : : :
Ex2) F2 F F F–F
Unshared pairs of electrons
(nonbonding pairs)
Take up more space then the shared pairs. .. .. .. ..
Ex3) H2O

24. Double Covalent Bonds – 2 shared pairs of electrons.
Triple Covalent Bonds – 3 shared pairs of electron.
Ex4) O2
Ex5) N2

25. Covalent Compounds
Ex6) NH3
Ex7) CH4
Ex8) CO2
Ch8 HW # 1

26. Ch8 HW#1 1 – 5
1) Dot structures for diatomic molecules
a) Chlorine, Cl2 b) Bromine, Br2 c) Iodine, I2
Cl Cl Br Br I I
2) How many unshared pairs are in each halogen molecule
3) Why necessary to form double & triple bond sometimes?
4) How many electrons does each atom contribute in:
Double Bond?
Triple Bond?

27. 5) Dot structures
a) H2S H S H
b) PH3 P H
c) ClF Cl F

28. Ch8.2 – More Covalent Bonds
Ex) Write electron dot structures for:
a) carbon monoxide, CO
Coordinate covalent bond - one atom contributes a pair
of electrons to the bond.
b) Hydroxide, OH–

29. Resonant Structures - double or triple bond that can jump around.
c) SO3–2
d) NH4+
e) H30+
f) O3
Resonant Structures - double or triple bond that can jump around.
Ch8 HW#2 6 – 8

30. Lab8.1 – Conductivity
- due tomorrow
- Ch8 HW#2 due at beginning of period

31. S O O O O C O O O Ch8 HW #2 6 – 8 6) Dot structures for SO4-2 & CO3-2
7) Resonant structures for CO3-2

32. 8) Dot structures:
a) BF3 B F F F
b) O2 O O
c) NO2-1 N O O
d) F2 F F

33. Ch8.3 – VESPR Theory
Ex 1) Methane, CH4
C H H H H

34. Ch8.3 – VESPR Theory C H H H H Ex 1) Methane, CH4 H H H C H or H-C-H
- Electrons pairs seek maximum separations in 3 dimensions
Instead of a 2-D cross, they get greater separation
in a 3-D tetrahedral shape. H
C bond angle Tetrahedral bond angles = 109.5˚
C H H H H

35. Ex2) Ammonia, NH3
N H H H

36. N H H H Pyramidal shape Bond angles = 107˚ Ex2) Ammonia, NH3
- Unshared pairs of electrons take up more space,
forces shared pairs closer together.
Ex3) Water, H2O
H H O

37. N H H H Pyramidal shape Bond angles = 107˚ Ex2) Ammonia, NH3
- Unshared pairs of electrons take up more space,
forces shared pairs closer together.
Ex3) Water, H2O
H H O
Bent shape
Bond angle = 180˚

38. C O O B F F F Ch8 HW#3 9 – 11 Ex4) Carbon Dioxide, CO2
Ex5) Boron trifluoride, BF3
B F F F
Ch8 HW#3 9 – 11

39. Ch8 HW#3 9 – 11
9. BF3 is trigonal planar. Add a fluoride ion, F–, in a coordinate
covalent bond, what is its shape?
B F F F F–
10. Use VSEPR to draw:
a. CCl b. PCl3
C Cl Cl Cl Cl P Cl Cl Cl
c. SeCl2
Se Cl Cl

40. 11. Draw
a. CO b. SiCl4
C O O Si Cl Cl Cl Cl
c. SO d. SCl2
S O O O S Cl Cl
e. CO f. I3+
C O I I I+

41. Ch8 HW#4
1. Predict the electron dot structure, shape and bond angles
a. silicon dioxide b. PH3
Si O O P H H H
c. sulfur dioxide d. N2O (1 N as central atom)
S O O N N O
e. CH2O f. Dinitrogen tetroxide
C H H O N N O O
O O
g. hydrogen peroxide
H H O O

42. Ch8.4 – Hybridization
- orbitals blend together to form the same #
of equivalent orbitals.
- most notably happens with carbon
Find the BP, EC, and OFD for carbon:
then bond it in methane

43. Experimentally determined that in methane, all the bond lengths
Ch8.4 – Hybridization
- orbitals blend together to form the same #
of equivalent orbitals.
- most notably happens with carbon
Find the BP, EC, and OFD for carbon:
C: [He] 2s2 2p2 H
then bond it in methane: H C H
Experimentally determined that in methane, all the bond lengths
are the same!

44. Atomic Blender
- mix 1 part s with 1 part p

45. Atomic Blender
- mix 1 part s with 1 part p
2 sp leftover p’s
- mix 1 part s with 2 parts p

46. 3 sp2 1 leftover p Atomic Blender - mix 1 part s with 1 part p
2 sp leftover p’s
- mix 1 part s with 2 parts p
3 sp leftover p
- mix 1 part s with 3 parts p

47. Atomic Blender
- mix 1 part s with 1 part p sp leftover p’s
- mix 1 part s with 2 parts p sp leftover p
- mix 1 part s with 3 parts p sp no leftover p
The first bond between 2 elements is always a direct overlap of these
hybridized orbitals. It is calles a sigma bond, σ.
The 2nd and 3rd bonds are sideways overlaps, called pi bonds, π.
These occur with the leftover p orbitals.

48. Ex1) Methane, CH4
- how many elements bonded to carbon? ___
- which recipe provides for that? ___
- any leftover p’s? ___
a) Draw and label the σ bonds.
b) Draw and label the π bonds.
c) Shape?

49. Ex2) Ethane, C2H6
- how many elements bonded to carbon? ___
- which recipe provides for that? ___
- any leftover p’s? ___
a) Draw and label the σ bonds.
b) Draw and label the π bonds.
c) Shape

50. Ex3) Ethene, C2H4
- how many elements bonded to carbon? ___
- which recipe provides for that? ___
- any leftover p’s? ___
a) Draw and label the σ bonds.
b) Draw and label the π bonds.
c) Shape
Ch8 HW#5 13 – 15

51. Ch8 HW#5 13 – 15
13. What shape would you expect a simple carbon-containing compound
to have when the carbon atom has the following hybridization?
a. sp2 b. sp3 c. sp

52. 14. What types of hybrid orbitals are involved in bonding of each atom
in the following molecules?
a. CO
b. C2H2
c. H2C=O
d. N C–C N

53. 15. The molecule CO2 has two carbon-oxygen double bonds.
Draw the bonding in the CO2 molecule using the hybridized orbitals
for carbon and oxygen.

54. Ch8.5 – Polarity
Nonpolar Bonds - electrons are shared equally between 2 atoms
Ex 1) Cl2

55. Nonpolar Bonds - electrons are shared equally between 2 atoms
Ex 1) Cl2 Cl Cl Cl – Cl
All of HNOFClBrI is nonpolar. But sometimes the random motion of the
shared electrons causes temporary polarity. This induces polarity in the
neighboring molecules.
Cl – Cl Cl – Cl

56. Nonpolar Bonds - electrons are shared equally between 2 atoms
Ex 1) Cl2 Cl Cl Cl – Cl
All of HNOFClBrI is nonpolar. But sometimes the random motion of the
shared electrons causes temporary polarity. This induces polarity in the
neighboring molecules.
δ– Cl – Cl δ+ δ– Cl – Cl δ+
This type of bond between molecules
is called Dispersion Forces. It is the
weakest of all bonds.
(Extremely low melting & boiling points)

57. Polar Bonds – Electrons are not shared equally between 2 atoms.
Ex2) HCl

58. Polar Bonds – Electrons are not shared equally between 2 atoms.
Ex2) HCl H Cl + H – Cl –
This arrangement is called a Dipole.
The Dipole – Dipole Bond between molecules is a little stronger.
+ H – Cl - + H – Cl -
higher melting & boiling points

59. Ex3) H2O

60. Ex3) H2O
Water contains 2 polar bonds, and overall is a polar molecule.
Its polarity attracts other water molecules.
When hydrogen is involved in polar bonds, it is called
Hydrogen Bonding. (Helps explain higher melting and boiling points.)
Ex4) CO2

61. Ex3) H2O
Water contains 2 polar bonds, and overall is a polar molecule.
Its polarity attracts other water molecules.
When hydrogen is involved in polar bonds, it is called
Hydrogen Bonding. (Helps explain higher melting and boiling points.)
Ex4) CO2
Contains 2 polar bonds, but the symmetry of the molecule cancels out
the polarity, making this molecule nonpolar.

62. Ex5) BF3

63. Ex5) BF3
Linear, trig planar, and tetrahedral molecules will be nonpolar IF all
the atoms attached to the central atom are the same.
Bent and pyramidal are always nonpolar.
Ex6) NaCl
Na Cl
Ch8 HW#6 16 – 20

64. Ex5) BF3
Linear, trig planar, and tetrahedral molecules will be nonpolar IF all
the atoms attached to the central atom are the same.
Bent and pyramidal are always nonpolar.
Ex6) NaCl
Na Cl Na+ Cl–
Ionic!
Ch8 HW#6 16 – 20

65. Lab8.2 – Models
- due in 3 days
- Ch8 HW#6 due at beginning of period.

66. CH8 HW#6 16 – 20
16) ID bonds a) H and Br b) K and CI
(ionic, polar cov,
nonpolar cov)
c) C and O d) CI and F e) Li and O
f) Br and Br g) F and F
17)draw with details a)HF b)HOOH c) BrCl d) H2O
18) Not every molecule with polar bonds
is nonpolar. Why is CCI4 nonpolar? CI
CI C CI CI

67. O–H H H CH8 HW#6 16 – 20 16) ID bonds a) H and Br b) K and CI
(ionic, polar cov,
nonpolar cov)
c) C and O d) CI and F e) Li and O
f) Br and Br g) F and F
17)draw with details a)HF b)HOOH c) BrCl d) H2O
a) H–F b) H–O
O–H
c) Br–CI d) O
H H
18) Not every molecule with polar bonds
is nonpolar. Why is CCI4 nonpolar? CI
CI C CI CI

68. 19) Hydrogen bonding between 2 NH3’s & between NH3 & H2O
20)Rank Forces
strongest middle weakest

69. Ch7,8 Rev
1. Define valence electron.
2. Electron configs for N and N-3 .
3. Dot structures for correct bonding between Na and O
Ionic or covalent?
4. Describe diff between ionic and covalent bonds.
5. Dot structure shape and polarity for OF2.
6. Resonant structures for SO2.
7. Why do compounds with strong intermolecular forces have higher
boiling points than compounds with weak intermolecular forces?

70. 8. What is the hybridization of each carbon in propene, aka propylene, C3H6
List σ bonds, π bonds, and shape around each carbon.