1. Chapter 7, 8, and 9 “Ionic, Covalent, and Metallic Bonding”
Chemistry
Mt. Hebron High School

2. Section Ions
OBJECTIVES:
Determine the number of valence electrons in an atom of a representative element.

3. Section Ions
OBJECTIVES:
Explain how the octet rule applies to atoms of metallic and nonmetallic elements.

4. Describe how cations form.
Section Ions
OBJECTIVES:
Describe how cations form.

5. Explain how anions form.
Section Ions
OBJECTIVES:
Explain how anions form.

6. In this chapter, you will learn about three types of bonding:
Ionic Bonding
Covalent Bonding
Metallic Bonding

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

8. Keeping Track of Electrons
Atoms in the same column...
Have the same outer electron configuration.
Have the same valence electrons.
The number of valence electrons are easily determined: the group number for a representative element
Group 2A: Be, Mg, Ca, etc.
have 2 valence electrons

9. Electron Dot diagrams are…
A way of showing & keeping track of valence electrons.
How to write them?
Write the symbol - it 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

10. 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.

11. The Octet Rule
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
The Octet Rule: in forming compounds, atoms tend to achieve a noble gas configuration; 8 in the outer level is stable
Each noble gas (except He) has 8 electrons in the outer level

12. Formation of Cations
Metals lose electrons to attain a noble gas configuration.
They make positive ions (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.

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

14. Electron Dots For Cations
Metals will have few valence electrons
Metals will lose the valence electrons Ca

15. Electron Dots For Cations
Metals will have few valence electrons
Metals will lose the valence electrons
Forming positive ions
Ca2+
This is named the calcium ion.
NO DOTS are now shown for the cation.

16. 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 are ions from chlorine or other halogens that gain electrons

17. Electron Dots For Anions
Nonmetals will have many valence electrons (usually 5 or more)
They will gain electrons to fill outer shell. P
P3-
This is called the phosphide ion

18. 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 = already stable!
This is the octet rule (8 in the outer level is particularly stable). Ar

19. Section 7.2 Ionic Bonds and Ionic Compounds
OBJECTIVES:
Explain the electrical charge of an ionic compound.

20. Section 7.2 Ionic Bonds and Ionic Compounds
OBJECTIVES:
Describe three properties of ionic compounds.

21. Ionic Bonding
Anions and cations are held together by opposite charges.
Ionic compounds are called salts.
Simplest ratio of elements in an ionic compound is called the formula unit.
The bond is formed through the transfer of electrons.
Electrons are transferred to achieve noble gas configuration.

22. Ionic Bonding Na Cl
The metal (sodium) tends to lose its one electron from the outer level.
The nonmetal (chlorine) needs to gain one more to fill its outer level, and will accept the one electron that sodium is going to lose.

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

24. Ionic Bonding
Lets do an example by combining calcium and phosphorus: Ca P
All the electrons must be accounted for, and each atom will have a noble gas configuration (which is stable).

25. Ionic Bonding Ca P

26. Ionic Bonding
Ca2+ P

27. Ionic Bonding
Ca2+ P Ca

28. Ionic Bonding
Ca2+
P 3- Ca

29. Ionic Bonding
Ca2+
P 3- Ca P

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

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

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

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

34. = Ca3P2 Ionic Bonding Formula Unit
This is a chemical formula, which shows the kinds and numbers of atoms in the smallest representative particle of the substance.
For an ionic compound, the smallest representative particle is called a: Formula Unit

35. Properties of Ionic Compounds
Crystalline solids - a regular repeating arrangement of ions in the solid
Ions are strongly bonded together.
Structure is rigid.
High melting points

36. 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

37. 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 around.
Melted ionic compounds conduct.
NaCl: must get to about 800 ºC.
Dissolved in water, they also conduct (free to move in aqueous solutions)

38. - 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.

39. Section 7.3 Bonding in Metals
OBJECTIVES:
Model the valence electrons of metal atoms.

40. Section 7.3 Bonding in Metals
OBJECTIVES:
Describe the arrangement of atoms in a metal.

41. Section 7.3 Bonding in Metals
OBJECTIVES:
Explain the importance of alloys.

42. Metallic Bonds are… How metal atoms are held together in the solid.
Metals hold on to their valence electrons very weakly.
Think of them as positive ions (cations) floating in a sea of electrons.

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

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

45. 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

46. Malleable +
Force

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

48. Ionic solids are brittle
Force + -

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

50. Crystalline structure of metal
If made of one kind of atom, metals are among the simplest crystals; very compact & orderly
1. Body-centered cubic:
every atom (except those on the surface) has 8 neighbors
Na, K, Fe, Cr, W

51. Crystalline structure of metal
2. Face-centered cubic:
every atom has 12 neighbors
Cu, Ag, Au, Al, Pb
3. Hexagonal close-packed
every atom also has 12 neighbors
different pattern due to hexagonal
Mg, Zn, Cd

52. Alloys We use lots of metals every day, but few are pure metals
Alloys are mixtures of 2 or more elements, at least 1 is a metal
made by melting a mixture of the ingredients, then cooling
Brass: an alloy of Cu and Zn
Bronze: Cu and Sn

53. 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

54. Chapter 8 “Covalent Bonding”
Chemistry
Mt. Hebron High School

55. Section 8.1 Molecular Compounds
OBJECTIVES:
Distinguish between the melting points and boiling points of molecular compounds and ionic compounds.

56. Section 8.2 Molecular Compounds
OBJECTIVES:
Describe the information provided by a molecular formula.

57. Bonds Ionic bonds – transfer of electrons (gained or lost)
Forces that hold groups of atoms together and make them function as a unit:
Ionic bonds – transfer of electrons (gained or lost)
Covalent bonds – sharing of electrons. The resulting particle is called a “molecule”.

58. Molecules Many elements found in nature are in the form of molecules:
a neutral group of atoms joined together by covalent bonds.
For example, air contains oxygen molecules, consisting of two oxygen atoms joined covalently
Called a “diatomic molecule”

59. How does H2 form?
The nuclei repel each other, since they both have a positive charge, and like charges repel. + +

60. How does H2 form? But, the nuclei are attracted to the electrons
They share the electrons; this is called a “covalent bond”, and involves only NONMETALS! + +

61. Covalent bonds Nonmetals hold on to their valence electrons.
They can’t give away electrons to bond.
Still want noble gas configuration.
Get it by sharing valence electrons with each other = covalent bonding
By sharing, both atoms get to count the electrons toward a noble gas configuration.

62. Covalent bonding
Fluorine has seven valence electrons F

63. F F Covalent bonding Fluorine has seven valence electrons
A second atom also has seven F F

64. F F Covalent bonding Fluorine has seven valence electrons
A second atom also has seven
By sharing electrons… F F

65. F F Covalent bonding Fluorine has seven valence electrons
A second atom also has seven
By sharing electrons… F F

66. F F Covalent bonding Fluorine has seven valence electrons
A second atom also has seven
By sharing electrons… F F

67. F F Covalent bonding Fluorine has seven valence electrons
A second atom also has seven
By sharing electrons… F F

68. F F Covalent bonding Fluorine has seven valence electrons
A second atom also has seven
By sharing electrons… F F

69. F F Covalent bonding …both end with full orbitals
Fluorine has seven valence electrons
A second atom also has seven
By sharing electrons…
…both end with full orbitals F F

70. F F Covalent bonding Fluorine has seven valence electrons
A second atom also has seven
By sharing electrons…
…both end with full orbitals F F
8 Valence electrons

71. F F Covalent bonding Fluorine has seven valence electrons
A second atom also has seven
By sharing electrons…
…both end with full orbitals F F
8 Valence electrons

72. Molecular Compounds
Compounds that are bonded covalently (like water and carbon dioxide) are called molecular compounds
Molecular compounds tend to have relatively lower melting and boiling points than ionic compounds.

73. Molecular Compounds
Thus, molecular compounds tend to be gases or liquids at room temperature
Ionic compounds were solids
A molecular compound consists of a molecular formula:
Shows how many atoms of each element a molecule contains

74. Molecular Compounds The formula for water is written as H2O
The subscript 2 behind hydrogen means there are 2 atoms of hydrogen; if there is only one atom, the subscript 1 is omitted
Molecular formulas do not tell any information about the structure (the arrangement of the various atoms).

75. - Page 215
These are some of the different ways to represent ammonia:
3. The ball and stick model is the BEST, because it shows a 3-dimensional arrangement.
1. The molecular formula shows how many atoms of each element are present
2. The structural formula ALSO shows the arrangement of these atoms!

76. Section 8.2 The Nature of Covalent Bonding
OBJECTIVES:
Describe how electrons are shared to form covalent bonds, and identify exceptions to the octet rule.

77. Section 8.2 The Nature of Covalent Bonding
OBJECTIVES:
Demonstrate how electron dot structures represent shared electrons.

78. Section 8.2 The Nature of Covalent Bonding
OBJECTIVES:
Describe how atoms form double or triple covalent bonds.

79. A Single Covalent Bond is...
A sharing of two valence electrons.
Only nonmetals and hydrogen.
Different from an ionic bond because they actually form molecules.
Two specific atoms are joined.
In an ionic solid, you can’t tell which atom the electrons moved from or to

80. Sodium Chloride Crystal Lattice
Ionic compounds organize in a characteristic crystal lattice of alternating positive and negative ions.

81. How to show the formation…
It’s like a jigsaw puzzle.
You put the pieces together to end up with the right formula.
Carbon is a special example - can it really share 4 electrons: 1s22s22p2?
Yes, due to electron promotion!
Another example: lets show how water is formed with covalent bonds, by using an electron dot diagram

82. H O Water Each hydrogen has 1 valence electron
Each hydrogen wants 1 more
The oxygen has 6 valence electrons
The oxygen wants 2 more
They share to make each other complete H O

83. H O Water Put the pieces together The first hydrogen is happy
The oxygen still wants one more H O

84. H O H Water A second hydrogen attaches
Every atom has full energy levels
Note the two “unshared” pairs of electrons H O H

85. Multiple Bonds
Sometimes atoms share more than one pair of valence electrons.
A double bond is when atoms share two pairs of electrons (4 total)
A triple bond is when atoms share three pairs of electrons (6 total)
Table 8.1, p Know these 7 elements as diatomic:
Br2 I2 N2 Cl2 H2 O2 F2
What’s the deal with the oxygen dot diagram?

86. Dot diagram for Carbon dioxide
CO2 - Carbon is central atom ( more metallic )
Carbon has 4 valence electrons
Wants 4 more
Oxygen has 6 valence electrons
Wants 2 more C O

87. Carbon dioxide
Attaching 1 oxygen leaves the oxygen 1 short, and the carbon 3 short C O

88. Carbon dioxide
Attaching the second oxygen leaves both oxygen 1 short, and the carbon 2 short O C O

89. Carbon dioxide
The only solution is to share more O C O

90. Carbon dioxide
The only solution is to share more O C O

91. Carbon dioxide
The only solution is to share more O C O

92. Carbon dioxide
The only solution is to share more O C O

93. Carbon dioxide
The only solution is to share more O C O

94. Carbon dioxide
The only solution is to share more O C O

95. O C O Carbon dioxide The only solution is to share more
Requires two double bonds
Each atom can count all the electrons in the bond O C O

96. O C O Carbon dioxide The only solution is to share more
Requires two double bonds
Each atom can count all the electrons in the bond
8 valence electrons O C O

97. O C O Carbon dioxide The only solution is to share more
Requires two double bonds
Each atom can count all the electrons in the bond
8 valence electrons O C O

98. O C O Carbon dioxide The only solution is to share more
Requires two double bonds
Each atom can count all the electrons in the bond
8 valence electrons O C O

99. Another way of indicating bonds
Often use a line to indicate a bond
Called a structural formula
Each line is 2 valence electrons H O H H O H =

100. Other Structural Examples
H C N H
C O H

101. A Coordinate Covalent Bond...
When one atom donates both electrons in a covalent bond.
Carbon monoxide (CO) is a good example:
Both the carbon and oxygen give another single electron to share O C

102. Coordinate Covalent Bond
When one atom donates both electrons in a covalent bond.
Carbon monoxide (CO) is a good example:
Oxygen gives both of these electrons, since it has no more singles to share.
This carbon electron moves to make a pair with the other single. C O

103. Coordinate Covalent Bond
When one atom donates both electrons in a covalent bond.
Carbon monoxide (CO)
The coordinate covalent bond is shown with an arrow as: C O
C O

104. Naming Ionic and Covalent Compounds and Writing Ionic Formulas
Chapter 9

105. Writing Ionic Formulas
Criss-Cross Method
Instead of drawing Ionic bonds to determine the chemical formula, we can use our knowledge of ionic charges to determine a compound’s formula.
You will need the use of an ion chart.

106. Criss-Cross Method Rules
1. Write each element (metal on the left and non-metal on the right).
2. Determine the charge on each element (use your ion chart).
3. Exchange charges with each element and write them as subscripts next to the new element.
4. Make sure you do not change the original element symbol. If a compound is present, parenthesis must be used first.

107. Sodium and Chlorine
Sodium Na +1
Chlorine Cl -1
Na1Cl1

108. Aluminum and Iodine
Aluminum Al+3
Iodine I-1
Al1I3

109. Calcium and Nitrogen
Calcium Ca+2
Nitrogen N-3
Ca3N2

110. Ammonium and Oxygen
Ammonium NH4+1
Oxygen O-2
(NH4)2O1

111. Ionic Compounds
There are some basic rules to follow when naming an Ionic Compound.
Make sure you have your Ion Chart and your periodic table.

112. Ionic Naming Rules
1. Identify the Metal and the Non-Metal in the compound.
2. Name the metal using the positive side of your ion chart.
3. Name the Non-Metal using the negative side of your ion chart.
4. Combine both halves together.
5. If there is more than one name for the metal, you must uncriss-cross the charges to determine the original charge of the metal ion.

113. Ionic Naming Examples 1. NaCl Na = Sodium Cl = Chloride
Sodium Chloride

114. 2. MgCl2
Mg = Magnesium
Cl2 = Chloride
Magnesium Chloride

115. AgNO2
CoCl2
Ni(OH)3
Fe(ClO3)2
Silver Nitrate
Cobalt (II) Chloride
Nickel (III) Hydroxide
Iron (II) Chlorate

116. End of Chapter 7,
8, and 9