1. Chapter 7 “Ionic and Metallic Bonding”

2. 7.1 - Ions

3. Valence Electrons are…?
Valence electrons – The s and p electrons in the outer energy level
the highest occupied energy level
Core electrons – electrons in the energy levels below.
The electrons responsible for the chemical properties of atoms, and are those in the outer energy level.

4. Keeping Track of Electrons
Atoms in the same column...
Have the same outer electron configuration.
Have the same valence electrons.
Electron configuration of Group 1A?
𝒔 𝟏
Number of valence electrons in Group 1A? 1

5. The number of valence electrons are easily determined
The number of valence electrons are easily determined. It is the group number for a representative element
Group 2A: Be, Mg, Ca, etc.
have 2 valence electrons

6. Electron Dot diagrams are…
A way of showing valence electrons.
How to write them?
Write the symbol – represents the nucleus and inner (core) electrons
Put one dot for each valence electron (8 maximum)
They don’t pair up until they have to (Hund’s rule) X

7. The Electron Dot diagram for Nitrogen
Nitrogen has 5 valence electrons to show.
First we write the symbol. N
Then add 1 electron at a time to each side.
Now they are forced to pair up.
We have now written the electron dot diagram for Nitrogen.

8. Practice with e- dot structureLi N Be O B F C Ne

9. The Octet Rule
The Octet Rule: in forming compounds, atoms tend to achieve a noble gas configuration; 8 in the outer level
Each noble gas (except He, which has 2) has 8 electrons in the outer level
In Chapter 6, we learned that noble gases are unreactive in chemical reactions
In 1916, Gilbert Lewis used this fact to explain why atoms form certain kinds of ions and molecules

10. Formation of Cations
Metals lose electrons to attain a noble gas configuration.
Where are metals located? Left
Make positive ions (cations)

11. Formation of Cations
If we look at the electron configuration, it makes sense to lose electrons:
Na 1s22s22p63s1 1 valence electron
Na1+ 1s22s22p6 This is a noble gas configuration with 8 electrons in the outer level.

12. Electron Dots For Cations
Metals will have few valence electrons (usually 3 or less); calcium has only 2 valence electrons Ca

13. Electron Dots For Cations
Metals will lose all valence electrons Ca

14. Electron Dots For Cations
Metals will lose the valence electrons
Form positive ions
Ca2+
This is the “calcium ion”.
NO DOTS are now shown for the cation.

15. Practice Cation formed Na = 1 𝑁𝑎 +1 Mg = 2 𝑀𝑔 +2 Al = 3 𝐴𝑙 +3
# of valence electron
Na = 1
Mg = 2
Al = 3
Cation formed
𝑁𝑎 +1
𝑀𝑔 +2
𝐴𝑙 +3

16. The transition metals get funky…

17. Electron Dots For Cations
Let’s do Scandium, #21
The electron configuration is: 1s22s22p63s23p64s23d1
Thus, it can lose 2e- (making it 2+), or lose 3e- (making 3+)
Sc = Sc2+ Sc
= Sc3+
Scandium (II) ion
Scandium (III) ion

18. Electron Dots For Cations
Let’s do Silver, element #47
Predicted configuration is: 1s22s22p63s23p64s23d104p65s24d9
Actual configuration is: 1s22s22p63s23p64s23d104p65s14d10
Ag = Ag1+ (can’t lose any more, charges of 3+ or greater are uncommon)

19. Electron Dots For Cations
Silver did the best job it could, but it did not achieve a true Noble Gas configuration
Instead, it is called a “pseudo-noble gas configuration”

20. Electron Configurations: Anions
Nonmetals gain electrons to attain noble gas configuration.
They make negative ions (anions)
S = 1s22s22p63s23p4 = 6 valence electrons
S2- = 1s22s22p63s23p6 = noble gas configuration.
Halide ions - ions from chlorine or other halogens that gain electrons

21. Electron Dots For Anions
Nonmetals will have many valence electrons (usually 5 or more)
They will gain electrons to fill outer shell. 3- P
(This is called the “phosphide ion”, and shows dots)

22. Stable Electron Configurations
All atoms react to try and achieve a noble gas configuration.
Noble gases have 2 s and 6 p electrons.
8 valence electrons = stable
This is the octet rule (8 in the outer level is particularly stable). Ar

23. Review How many valence electrons do the following elements have?
Ga = 3
K = 1
P = 5
Li = 1
Cl = 7
He = 2
S = 6
F = 7
Mg = 2
Ne = 8

24. Review Draw the electron dot structure of the following elements.Ga K P He S F Mg Ne

25. Review What ions will the following elements form?
𝐺𝑎 +3
𝑀𝑔 +2
𝑃 −3
𝐹 −1
𝐾 +1
𝑆 −2

26. 7.2 - Ionic Bonds and Ionic Compounds

27. Ionic Bonding
Anions and cations are held together by opposite charges (+ and -)
Ionic compounds are called salts.
Formula unit – Simplest ratio of elements in an ionic compound.

28. Practice with Formula Unit
What is the formula unit for the following compounds?
Formula Unit
NaCl :1
𝐻 2 𝑂 2:1
𝐶𝑎𝐶𝑂 :1:3

29. Ionic Bonds Formed through the transfer of electrons (lose and gain)
Electrons are transferred to achieve noble gas configuration.

30. Ionic Compounds Also called SALTS Made from a CATION and ANION
A metal and a nonmetal)

31. Ionic Bonding Na Cl
The metal loses its one electron from the outer level.
The nonmetal gains one more to fill its outer level, and will accept the one electron that the metal is going to lose.

32. Ionic Bonding
Na+
Cl -
Note: Remember that NO DOTS are now shown for the cation!

33. Ca P Ionic Bonding Example- combine Ca and P:
All electrons must be accounted for, and each atom will have a noble gas configuration (which is stable).

34. Ionic Bonding Ca P

35. Ionic Bonding
Ca2+ P

36. Ionic Bonding
Ca2+ P Ca

37. Ionic Bonding
Ca2+
P 3- Ca

38. Ionic Bonding
Ca2+
P 3- Ca P

39. Ionic Bonding
Ca2+
P 3-
Ca2+ P

40. Ionic Bonding Ca
Ca2+
P 3-
Ca2+ P

41. Ionic Bonding Ca
Ca2+
P 3-
Ca2+ P

42. Ionic Bonding
Ca2+
Ca2+
P 3-
Ca2+
P 3-

43. Ionic Bonding
= Ca3P2
Chemical formula- shows the kinds and numbers of atoms in the smallest representative particle of the substance.

44. Balancing Ionic formulas
Crisscross method
The numerical value of the charge of each ion is crossed over and becomes the subscript for the other ion.

45. Practice What will the chemical formula look like?
Elements
Ca+2 F-
Al+3 O-2
Ca+2 O-2
Formula
𝐶𝑎𝐹 2
𝐴𝑙 2 𝑂 3
𝐶𝑎𝑂

46. Properties of Ionic Compounds
Crystalline solids – regular repeating arrangement of ions
Strongly bonded together.
Structure is rigid.
High melting points
Coordination number- number of ions of opposite charge surrounding it

47. NaCl CsCl TiO2 - Page 198 Coordination Numbers:
Both the sodium and chlorine have 6
NaCl
Both the cesium and chlorine have 8
CsCl
Each titanium has 6, and each oxygen has 3
TiO2

48. Do they Conduct?
Conducting electricity means allowing charges to move.
In a solid, the ions are locked in place.
Ionic solids are insulators.
When melted, the ions can move.
Melted ionic compounds conduct.
NaCl: must get to about 800 ºC.
Dissolved in water, they also conduct (free to move in aqueous solutions)

49. - Page 198
The ions are free to move when they are molten (or in aqueous solution), and thus they are able to conduct the electric current.

50. Review What is the formula unit for the following compounds?
KCl 1:1
𝐹𝑒 2 𝑂 :3
𝐹𝑒𝑆 2 1:2
𝐵𝑎𝑆𝑂 :1:4
HgS 1:1
𝑁𝑎 2 𝑂 2:1

51. Practice What will the chemical formula look like?
Elements
Na+ Cl-
Mg+2 F-
Ca+2 N-3
Al+3 N-3
K+ S-2
Formula
𝑁𝑎𝐶𝑙
𝑀𝑔𝐹 2
𝐶𝑎 3 𝑁 2
𝐴𝑙𝑁
𝐾 2 𝑆

52. 7.3 – Bonding in Metals

53. Metallic Bonds are… How do we get sheets of Aluminum (or any metal)?
Metals hold on to their valence electrons very weakly.
Think of them as positive ions (cations) floating in a sea of electrons

54. Sea of Electrons + Electrons are free to move through the solid.
Metals conduct electricity. +

55. Metals are Malleable Hammered into shape (bend).
Ductile - drawn into wires.
Both malleability and ductility explained in terms of the mobility of the valence electrons

56. Due to the mobility of the valence electrons, metals have:
- Page 201
Due to the mobility of the valence electrons, metals have:
Notice that the ionic crystal breaks due to ion repulsion!
1) Ductility
2) Malleability
and

57. Malleable +
Force

58. Malleable + + + + Force + + + + + + + +
Mobile electrons allow atoms to slide by, sort of like ball bearings in oil. + + + +
Force + + + + + + + +

59. Ionic solids are brittle
Force + -

60. Ionic solids are brittle
Strong Repulsion breaks a crystal apart, due to similar ions being next to each other. + -
Force + - + - + -

61. Crystalline structure of metal
Metals are arranged in very compact and orderly patterns.

62. Alloys Alloys - mixtures of 2 or more elements, at least 1 is a metal
Made by melting a mixture of elements, then cooling
Brass: an alloy of Cu and Zn
Bronze: Cu and Sn

63. Why use alloys? Properties are often superior to the pure element
Sterling silver (92.5% Ag, 7.5% Cu) is harder and more durable than pure Ag, but still soft enough to make jewelry and tableware
Steels are very important alloys
corrosion resistant, ductility, hardness, toughness, cost

64. The
END

65. Predicting Ionic Charges
Group B elements:
Many transition elements
have more than one possible oxidation state.
Note the use of Roman numerals to show charges
Iron (II) = Fe2+
Iron (III) = Fe3+

66. Naming cations
Two methods can clarify when more than one charge is possible:
Stock system – uses roman numerals in parenthesis to indicate the numerical value
Classical method – uses root word with suffixes (-ous, -ic)
Does not give true value

67. Naming cations We will use the Stock system.
Cation - if the charge is always the same (like in the Group A metals) just write the name of the metal.
Transition metals can have more than one type of charge.
Indicate their charge as a roman numeral in parenthesis after the name of the metal (Table 9.2, p.255)

68. Predicting Ionic Charges
Some of the post-transition elements also
have more than one possible oxidation state.
Tin (II) = Sn2+
Lead (II) = Pb2+
Tin (IV) = Sn4+
Lead (IV) = Pb 4+

69. Predicting Ionic Charges
Group B elements:
Some transition elements
have only one possible oxidation state, such as these three:
Silver = Ag1+
Zinc = Zn2+
Cadmium = Cd2+

70. Do not need to use roman numerals for these:
Exceptions:
Some of the transition metals have only one ionic charge:
Do not need to use roman numerals for these:
Silver is always 1+ (Ag1+)
Cadmium and Zinc are always 2+ (Cd2+ and Zn2+)

71. Practice by naming these:
Ca2+
Al3+
Fe3+
Fe2+
Pb2+
Li1+

72. Write symbols for these:
Potassium ion
Magnesium ion
Copper (II) ion
Chromium (VI) ion
Barium ion
Mercury (II) ion

73. Anions are always the same charge
Naming Anions
Anions are always the same charge
Change the monatomic element ending to – ide
F1- a Fluorine atom will become a Fluoride ion.

74. Practice by naming these:
Cl1-
N3-
Br1-
O2-
Ga3+

75. Write symbols for these:
Sulfide ion
Iodide ion
Phosphide ion
Strontium ion

76. Polyatomic ions are…
Groups of atoms that stay together and have an overall charge, and one name.
Usually end in –ate or -ite
Acetate: C2H3O21-
Nitrate: NO31-
Nitrite: NO21-
Permanganate: MnO41-
Hydroxide: OH1- and Cyanide: CN1-?

77. (One of the few positive polyatomic ions)
Know Table 9.3 on page 257
Phosphate: PO43-
Phosphite: PO33-
Ammonium: NH41+
Sulfate: SO42-
Sulfite: SO32-
Carbonate: CO32-
Chromate: CrO42-
Dichromate: Cr2O72-
(One of the few positive polyatomic ions)
If the polyatomic ion begins with H, then combine the word hydrogen with the other polyatomic ion present: H CO → HCO hydrogen + carbonate → hydrogen carbonate ion

78. Writing Ionic Compound Formulas
Example: Barium nitrate (note the 2 word name)
1. Write the formulas for the cation and anion, including CHARGES!
( )
Ba2+
NO3- 2
2. Check to see if charges are balanced.
Now balanced.
Not balanced!
= Ba(NO3)2
3. Balance charges , if necessary, using subscripts. Use parentheses if you need more than one of a polyatomic ion. Use the criss-cross method to balance subscripts.

79. Writing Ionic Compound Formulas
Example: Ammonium sulfate (note the 2 word name)
( )
1. Write the formulas for the cation and anion, including CHARGES!
NH4+
SO42- 2
Now balanced.
2. Check to see if charges are balanced.
Not balanced!
= (NH4)2SO4
3. Balance charges , if necessary, using subscripts. Use parentheses if you need more than one of a polyatomic ion. Use the criss-cross method to balance the subscripts.

80. Writing Ionic Compound Formulas
Example: Iron (III) chloride (note the 2 word name)
1. Write the formulas for the cation and anion, including CHARGES!
Fe3+
Cl- 3
Now balanced.
2. Check to see if charges are balanced.
Not balanced!
= FeCl3
3. Balance charges , if necessary, using subscripts. Use parentheses if you need more than one of a polyatomic ion. Use the criss-cross method to balance the subscripts.

81. Writing Ionic Compound Formulas
Example: Aluminum sulfide (note the 2 word name)
1. Write the formulas for the cation and anion, including CHARGES!
Al3+
S2- 2 3
2. Check to see if charges are balanced.
Now balanced.
Not balanced!
= Al2S3
3. Balance charges , if necessary, using subscripts. Use parentheses if you need more than one of a polyatomic ion. Use the criss-cross method to balance the subscripts.

82. Writing Ionic Compound Formulas
Example: Magnesium carbonate (note the 2 word name)
1. Write the formulas for the cation and anion, including CHARGES!
Mg2+
CO32-
2. Check to see if charges are balanced.
They are balanced!
= MgCO3

83. Writing Ionic Compound Formulas
Example: Zinc hydroxide (note the 2 word name)
1. Write the formulas for the cation and anion, including CHARGES!
( )
Zn2+
OH- 2
Now balanced.
2. Check to see if charges are balanced.
Not balanced!
= Zn(OH)2
3. Balance charges , if necessary, using subscripts. Use parentheses if you need more than one of a polyatomic ion. Use the criss-cross method to balance the subscripts.

84. Writing Ionic Compound Formulas
Example: Aluminum phosphate (note the 2 word name)
1. Write the formulas for the cation and anion, including CHARGES!
Al3+
PO43-
2. Check to see if charges are balanced.
They ARE balanced!
= AlPO4

85. Naming Ionic Compounds
1. Name the cation first, then anion
2. Monatomic cation = name of the element
Ca2+ = calcium ion
3. Monatomic anion = root + -ide
Cl- = chloride
CaCl2 = calcium chloride

86. Naming Ionic Compounds
(Metals with multiple oxidation states)
some metals can form more than one charge (usually the transition metals)
use a Roman numeral in their name:
PbCl2 – use the anion to find the charge on the cation (chloride is always 1-)
Pb2+ is the lead (II) cation
PbCl2 = lead (II) chloride

87. Things to look for:
If cations have ( ), the number in parenthesis is their charge.
If anions end in -ide they are probably off the periodic table (Monoatomic)
If anion ends in -ate or –ite, then it is polyatomic

88. Practice by writing the formula or name as required…
Iron (II) Phosphate
Stannous Fluoride
Potassium Sulfide
Ammonium Chromate
MgSO4
FeCl3