1. Chapter 6 Bonding

2. Objectives
Write the electron configuration of ions of representative elements.
Define ionic bonds in terms of the difference in electronegativity of the atoms.
Write the Lewis electron dot structures of atoms, ions and ionic compounds.
Write the name of an ionic compound (binary, polyatomic and transition metal) given its formula, or its formula given its name.
Describe the properties of ionic, metallic and covalent molecules and identify the forces holding them together.

3. Valence Electrons
Mendeleev ordered his periodic table according to chemical behavior.
A pattern he saw was that elements in the same group behaved the same chemically.
This is due to its valence electrons.
Valence electrons are the outermost electrons and are involved in bonding.

4. Calculating valence electrons
Groups on the periodic table can help!
What’s left out on the periodic table?

5. Transition Metals
Transition metals can have different amounts of valence electrons depending on how they bond with nonmetals

6. Decide how many valence electrons each one has?Li F Kr Mg
Group 13
Carbon
Phosphorus
1 valence electron
7 valence electron
8 valence electron
2 valence electron
3 valence electron
4 valence electron
5 valence electron

7. Gilbert Lewis
Lewis developed the idea of the octet and coined the term Lewis dot structures
He was nominated 35 times for the Nobel Prize in chemistry, but never won.

8. Electron Dot structures
A diagram that shows only the valence electrons around the atom as dots
Some rules:
Each side can only have 2 dots, for a maximum amount of 8 dots
An Exception to group 18 is Helium, it will only have 2

9. Electron Dot structures
Draw the correct Lewis dot structure for each given:
1 V.E 2 V.E 3 V.E V.E 5 V.E
6 V.E 7 V.E 8 V.E

10. Try these on your own:

11. Electron dot structures

12. Think back to Nobel Gases
Nobel gases are unreactive…
What do you notice about their electron dot structures??

13. Octet Rule
Atoms want to achieve stability or noble gas configuration by attaining 8 valence electrons.
Maximum number in the s and p orbitals.
Atoms do not want to gain or lose more than 3 electrons when bonding!

14. Which is the only group that has full octets?
NOBLE GASES!

15. Ions

16. Ions Ions are formed when an atom loses or gains electrons
Ions can gain or lose just one or multiple electrons to bond and achieve stability
Ions do not want to gain or lose more than 3 electrons though when bonding

17. Cations
Cations are formed when an atom loses electrons and becomes positive
Metals form cations
Examples: magnesium, potassium, aluminum

18. Anions
Anions are formed when an atom gains electrons and becomes negative
Nonmetals form anions
Examples: fluorine, oxygen, phosphorus

19. How do we know if atoms lose or gain electrons?
Let’s look at their Valence electrons
Every atom wants to be at 8 valence electrons
Sodium has 1 VE, so its easier to lose 1 and have a +1
Chlorine has 7 VE, so its easier to gain 1 and would have a – 1

20. Charge or Oxidation state
The charge corresponds with the amount of valence electrons:

21. Did it lose or gain? Oxide ion Cesium ion Aluminum ion Bromide ion
Phosphide ion
O-2: gained 2 electrons
Cs+1: lost 1
Al+3 lost 3 electrons
Br-1 gained 1 electrons
P-3 gained 3 electrons

22. Valence electrons: Gain or lose, and how many?
Strontium
Hydrogen
Sulfur
Xenon
Iodine
Lose 2
Lose 1
Gain 2
Neither, neutral
Gain 1

23. Polyatomic Ions

24. Polyatomic Ions

25. Ions
Review

26. Predicting Ionic Charges
Group 1:
Lose 1 electron to form 1+ ions H+
Li+
Na+ K+

27. Predicting Ionic Charges
Group 2:
Loses 2 electrons to form 2+ ions
Be2+
Mg2+
Ca2+
Sr2+
Ba2+

28. Predicting Ionic Charges
Group 13:
B3+
Al3+
Ga3+
Loses 3
electrons to form
3+ ions

29. Predicting Ionic Charges
Nitride
Group 15:
P3-
Phosphide
As3-
Arsenide
Gains 3
electrons to form
3- ions

30. Predicting Ionic Charges
Oxide
Group 16:
S2-
Sulfide
Se2-
Selenide
Gains 2
electrons to form
2- ions

31. Predicting Ionic Charges
F1-
Fluoride
Br1-
Bromide
Group 17:
Cl1-
Chloride
I1-
Iodide
Gains 1
electron to form
1- ions

32. Predicting Ionic Charges
Group 18:
Stable Noble gases do not form ions!

33. Predicting Ionic Charges
Groups :
Many transition elements
have more than one possible oxidation state.
Iron(II) = Fe2+
Iron(III) = Fe3+

34. Predicting Ionic Charges
Groups :
Some transition elements
have only one possible oxidation state.
Zinc = Zn2+
Silver = Ag+

35. Chapter 6 Types of Bonding

36. Atoms seldom exist in nature as independent particles.
Nearly all substances are made up of a combination of atoms that are held together by chemical bonds.
Chemical Bond – a mutual electrical attraction between the nuclei and valence electrons of different atoms that bind the atoms together.

37. When atoms bond, their valence electrons are redistributed in ways that make the atoms more stable.
The ways in which the electrons are redistributed determines the type of bonding.
Metals tend to lose electrons to form positive ions, or cations, and nonmetals tend to gain electrons to form negative ions, or anions.

38. Chemical bonding that results from the electrical attraction between cations and anions is called ionic bonding.
Chemical bonding that results from the sharing of electrons between atoms is covalent bonding.
In covalent bonds the shared electrons are owned equally by the two bonded atoms.

40. You can estimate the type of bonding (ionic or covalent) between elements by the difference in electronegativities (Chapter 5).

41. Electronegativity 0.0-0.3 Non polar covalent 0.3-1.7 polar covalent
This chart will help you determine if it is polar or nonpolar
Non polar covalent
polar covalent
>/= 1.8 Ionic

42. Polar Covalent vs. Non-polar Covalent
Blue shading represents electron density

43. Sample Problem
Page 163
Use electronegativity differences and Figure 6-2 to classify bonding between sulfur, S, and the following elements: hydrogen, H; cesium, Cs; and chlorine, Cl as either ionic, polar covalent or non-polar covalent.

44. electronegativity
difference bond type
S and H = polar
S and Cs = ionic
S and Cl = polar

45. Ionic Bonds

46. Opposites attract!
An ionic compound is composed of positive and negative ions that are combined so that the charges are equal.
Cations will combine with anions to form ionic compounds or salts. (metal and nonmetal)
This electrostatic force that holds them together is called an ionic bond!

47. Na+ + Cl- NaCl Example: sodium chloride consists of a
positive ion (Na) with a +1 charge and
negative ion (Cl) with a -1 charge.
Na+ + Cl NaCl
Draw electron configurations and Lewis Dot Structures.

48. Ionic Bonds Cation + Anion  Ionic bond (Ionic compound)
Salt (NaCl) Ionic bond between a metal and non-metal
Look at the electron dot structures.
Note how the sodium donates its electron to chlorine, now chlorine has an octet and a negative charge

49. Structure of Sodium Chloride Crystal

50. Structure of Sodium Chloride Crystal

51. Ionic bonds
Show the ionic bonds of the following using the electron dot structure method.
Magnesium and oxygen
(MgO)
2. Potassium and Sulfur
(K2S)
3. Calcium and chloride
(CaCl2)

52. Properties of ionic compounds

53. Properties of ionic compounds
Electrically neutral
Hard, but brittle
Most form crystal lattices at room temperature
Generally, have high melting points and Boiling Points
Cinnabar (HgS)
Fluorite (CaF2)

54. Properties of ionic compounds
Conduct electricity when dissolved in water or as a liquid. 
Solids do not conduct electricity.

55. Ionic Compounds in Water
When ionic compounds are placed in water, they will dissociate.
This means the anions and cations will split apart and form weak bonds with the water molecules.

56. Sodium chloride vs sugar
C12H22O11
NaCl
Sodium is a metal
Chloride is a nonmetal
Melting point ~800oC
C, H, and O are nonmetals
Melting point ~185oC
Which one (or both) is/are an ionic compound(s)?

57. Crystal Lattice
Ionic compounds form in repeating patterns of anions and cations.
Their crystal structure will be the same for a particular compound.

58. Ionic crystals
Ionic compounds that are crystals are made out of small pieces called unit cells
A unit cell is the simplest repeating unit in a crystal
NaCl’s unit cell:

59. Lesson check: True or False: Ionic compounds have low melting points
Nitrogen and oxygen form an ionic compound
Ionic Compounds are normally liquids
Ionic Compounds conduct electricity when dissolved
False
True

60. Metallic Bonding

61. What are some properties of metals?
Think back to Chapter 1

62. Properties of metals Good conductors of electricity
Ability to be drawn into a wire (Ductile)
Ability to be hammered, without breaking (Malleable)
Not brittle

63. Metallic bonds
Metals are composed of closely packed cations held together by their outer electrons.
These outer electrons can be referred to as a sea of electrons. The electrons move freely between metal atoms.
This “sea of electrons” is what allows metals to be molded, hammered, or stretched.

64. Metallic bonding
Metallic Bonds are the forces or attraction between those free floating outside electrons and the positively charged metal ions

65. Metal alloys
An alloy is a mixture of 2 or more elements (one must be a metal)
These are uniform throughout, so a homogeneous mixture
Examples: Brass (copper and zinc); Sterling silver (silver and copper); Bronze (copper and tin)

66. Why have alloys?
Alloys are important because they are combining properties and are often superior compared to the pure elements
Typically, more inexpensive than the pure element:
Sterling silver vs pure silver
$0.95 vs $1.68

67. Think about this:
A bronze statue is beginning to turn green; bronze is an alloy made of copper and tin
Which element is causing it to become green?
Hint: Think statue of liberty

68. Covalent Molecules

69. Molecules and Molecular Compounds
Ionic bonds are a + and -, but what about CO2?
What is type of elements are C and O?
When nonmetals bond together a covalent bond is created and we call them molecules or molecular compounds!

70. Molecules
Molecules are neutral atoms that are joined together by covalent bonds
Molecular Compound another way a saying molecule
Molecular formula shows you how many atoms of each element is in a substance
Example: CO2 , NH4

71. Octet Rule and Covalent Bonding
An octet is 8 valence electrons and want to achieve noble gas configuration!
Molecules want the same thing, but they share their valence electrons

72. Sharing electrons Recall that ionic bonds give and take electrons…
Molecules share their electrons between the 2 atoms.
When they share their valence electrons, a covalent bond is made

73. Single covalent bonds
When atoms share one pair of electrons they form a single covalent bond
Example: H2
Let’s draw it:

74. Show these diatomics:
Cl2
Br2 I2 F2
What about H2O?

75. Structural Formula
Electron dot structure represents bonds as 2 dots coming together:
A structural formula represents covalent bonds as dashes

76. What did we call those 2 dots next to one another?
Lone pairs or unshared pairs
They do NOT participate in bonding, but you must show them!

77. Try these on your own:
NH3
CH4
H2O2
PCl3

78. Double Bonds
Atoms that share two pairs of electrons
Example: CO2

79. Double Bonds
Draw O2:

80. Triple Bonds
Atoms that share three pairs of electrons:
Example: N2

81. Properties of Covalent Molecules

82. Properties of covalent molecules
Made out of nonmetals
Can be a solid, liquid, or gas at room temperature
Low melting point and boiling points
Poor to non-conductors of heat and electricity

83. H2O vs NaCl
Liquid water
Solid water

84. Strengths of covalent bonds vs. ionic bonds

86. Bonding Theories

87. How do we decide where to find electrons?
The modern atomic theory tells us that they are most probable at certain locations

88. Molecular Orbitals
When two atoms combine, their atomic orbitals combine and overlap to produce molecular orbitals
A molecular orbital belongs to the whole of the molecule
Each orbital can only contain 2 electrons
When an orbital overlaps and participates in a covalent bond, it can be classified as a bonding orbital

89. Sigma bonding (σ)
The first bond between sharing atoms is classified as sigma bonds

90. Pi bonding (π) The second type of bonding can be a pi
Remember the first is a sigma and the rest can be pi bonds

91. Molecular Geometry
VSEPR Theory

92. VSEPR Theory Valence-Shell-Electron-Pair-Repulsion theory
This theory helps us understand the 3D structure of molecules and their properties.
Bonding and unshared pairs of valence electrons become very important to us within VSEPR theory!
The shapes of molecules are determined because electron pairs want to be far apart from each other (repulsion).

93. AXE – Method to represent compounds
A represents the central atom X represents the bonding atoms E represents the lone pairs the central atom has

94. Compounds with
no lone pairs

95. Bonding pairs are far apart from each other
Draw or build CO2
Meaning 1 central atom, 2 bonded atoms
It has a linear shape
No lone pairs
A bond angle of 180o
Bonding pairs are far apart from each other

96. This has a trigonal planar
Draw or build BF3
This has a trigonal planar
Meaning 1 central, 3 bonded
Bond angle: 120o
Bonds pointing to the corners of a triangle

97. Draw or build CH4 This has a tetrahedral AX4 Bond angle: 109.5o
Meaning 1 central, 4 bonded
Bond angle: 109.5o

98. Problem Use VSEPR theory to predict the shape of:
aluminum trichloride, AlCl3
hydrogen iodide, HI
carbon tetrabromide, CBr4
dichloromethane, CH2Cl2

99. Compounds with lone pairs
Lone pairs occupy space, but only bonded atoms determine the name

100. Draw or build H2O
This has a bent shape
AX2E2
Meaning 1 central, 2 bonded, 2 lone pairs
Bond angle: 105o
Similar angles to the
tetrahedral bond angles

101. This has a trigonal pyramidal AX3E
Draw or build NH3
This has a trigonal pyramidal
AX3E
Meaning 1 central, 3 bonded, 1 lone pair
Bond angle: 107o
Similar angles to the
tetrahedral bond angles

102. What shape is each one? Linear BeCl2 Bent OF2 Trigonal Planar AlCl3
Trigonal pyramidal
Tetrahedral
BeCl2
OF2
AlCl3
PCl3
CF4

103. Switch presentations – slide 80
Bond Polarity
Switch presentations – slide 80

104. Bond polarity Since, atoms are sharing within a covalent bond…
If they share equally they are a nonpolar covalent bond
Examples: Diatomic atoms are nonpolar because they pull on each other evenly

105. Bond polarity Since, atoms are sharing within a covalent bond…
If they share unequally they are a polar covalent bond
Examples: HCl, H2O

106. Bond polarity
Since, polar bonds are unequally sharing we will have dipoles.
But how will we decide polarity??
Electronegativity!!

107. Dipoles
The more electronegative will have the arrow point towards it and have a slightly negative charge
The less electronegative will have a slightly positive charge

108. Electronegativity 0.0-0.4 Non polar covalent
This chart will help you determine if it is polar or nonpolar
Non polar covalent
Slightly polar covalent
Very polar covalent
>/= 2.0 Ionic

109. Decide the polarity of the following:
N-H
F-F
Ca- Cl
Al- Cl
0.9 slightly polar
0 Nonpolar
2.0 ionic
1.5 very polar