1. Unit 2: Chemical Bonding
Chemistry2202

2. Outline Bohr diagrams & Lewis Diagrams Types of Bonding
Ionic
Covalent (molecular)
Metallic
Network covalent bonding
London Dispersion forces
Dipole-Dipole forces
Hydrogen Bonding
VSEPR Theory (Shapes)
Physical Properties

3. Bohr Diagrams (Review)
How do we draw a Bohr Diagram for
The F atom?
The F ion?
Draw Bohr diagrams for the atom and the ion for the following:
Al S C l Be

4. Lewis Diagrams
Lewis Diagrams provide a method for keeping track of electrons in atoms, ions, or molecules.
AKA: Electron Dot diagrams
the nucleus (p+ & no), and filled energy levels are represented by the symbol
dots are placed around the element symbol to represent valence electrons

5. F • • • • • • • Lewis Diagrams lone pair bonding electron lone pair
eg. Lewis Diagram for F
lone pair • •
bonding electron • F
lone pair • • • •
lone pair

6. Lewis Diagrams
lone pair – a pair of electrons not available for bonding
bonding electron – a single electron that may be shared with another atom

7. P C Na • • • • • • • • • • Lewis Diagrams eg. Draw Lewis Diagrams for:
carbon
phosphorus • • • • • P C • • • •
sodium • Na

8. Li Be Al Si Mg N B O Lewis Diagrams
For each atom draw the Lewis diagram and state the number of lone pairs and number of bonding electrons
Li Be Al Si
Mg N B O

9. Lewis Diagrams for Compounds
draw the LD for each atom in the compound
The atom with the most bonding electrons is the central atom
Connect the other atoms using single bonds (1 pair of shared electrons)
In some cases there may be double bonds or triple bonds

10. Lewis Diagrams for Compounds
eg. Draw the LD for:
PH3
CF4
Cl2O
C2H6
C2H4
C2H2

11. Lewis Diagrams for Compounds
eg. Draw the LD for:
NH3 SiCl4 N2H4 HCN
SI2 CO2 N2H2 CH2O
POI CH3OH
N2 H2 O2

12. Lewis Diagrams for Compounds
A structural formula shows how the atoms are connected in a molecule.
To draw a structural formula:
replace the bonded pairs of electrons with short lines
omit the lone pairs of electrons

13. Why is propane (C3H8) a gas at STP while kerosene (C10H22) a liquid?

14. Why is graphite soft enough to write with while diamond is the hardest substance? (both are C)

15. Graphite
Diamond

16. Other forms of carbon
Buckyball Carbon
Nanotube Carbon

17. Why can you tell if it is ‘real gold’ or just ‘fool’s gold’ (pyrite) by hitting it with a rock?

18. ‘As Slow As Cold Molasses’
‘All Because of Bonding’

19. Viscosity of liquids

20. -30 ºC

21. Malleability, Ductility and Conductivity
A single gram of gold can be
stretched into a wire 3.2km long.
 A gram of gold can be flattened into a sheet with an area of 6.7 sq ft.
 Silver has the highest electrical conductivity of any element and the highest thermal conductivity of any metal. 

22. Melting and Boiling PointsF2
Br2
Cl2 I2

23. dihydrogen monoxide
pepper demo
dd & HB
teacher tube - IF

24. Bonding
Mar. 10
Bonding between atoms, ions and molecules determines the physical and chemical properties of substances.
Bonding can be divided into two categories:
- Intramolecular forces
- Intermolecular forces

25. Bonding
Intramolecular forces are forces of attraction between atoms or ions.
Intramolecular forces include:
ionic bonding
covalent bonding
metallic bonding
network covalent bonding

26. Bonding
Intermolecular forces are forces of attraction between molecules.
Intermolecular forces include:
London Dispersion Forces
Dipole-Dipole forces
Hydrogen Bonding

27. Ionic and Covalent Bonding
ThoughtLab p. 161
Identify #’s 1 - 6

28. Ionic Bonding
Occurs between cations and anions – usually metals and non-metals.
An ionic bond is the force of attraction between positive and negative ions.
Properties:
conduct electricity as liquids and in solution
hard crystalline solids
high melting points and boiling points
brittle

29. Ionic Bonding In an ionic crystal the ions pack tightly together.
Mar. 14
In an ionic crystal the ions pack tightly together.
The repeating 3-D distribution of cations and anions is called an ionic crystal lattice.

30. Ionic Bonding
Each anion can be attracted to six or more cations at once.
The same is true for the individual cations.

31. Ionic Bonding

32. Covalent Bonding Occurs between non-metals in molecular compounds.
Atoms share bonding electrons to become more stable (noble gas structure).
A covalent bond is a simultaneous attraction by two atoms for a common pair of valence electrons.

33. Covalent Bonding
Molecular compounds have low melting and boiling points.
Exist as distinct molecules.

34. Covalent Bonding
Molecular compounds do not conduct electric current in any form

35. Property Ionic Molecular
Type of elements
Metals and nonmetals
Non-Metals
Force of Attraction
Positive ions attract negative ions
Atoms attract a shared electron pair
Electron movement
Electrons move from the metal to the nonmetal
Electrons are shared between atoms
State at room temperature
Always solids
Solids, liquids, or gas

36. Property Ionic Molecular
Solubility
Soluble or low solubility
Soluble or insoluble
Conductivity in solid state
None
Conductivity in liquid state
Conducts
Conductivity in solution

37. Metallic Bonding (p. 171) Na • Na • Na • Na • Na • Na • Na • Na •

38. Na+ Na+ Na+ Na+ Na+ Na+ Na+ Na+ Metallic Bonding (p. 171) • • • • • •

39. Metallic Bonding (p. 171) metals tend to lose valence electrons.
valence electrons are loosely held and frequently lost from metal atoms.
the electron loss produces metal ions surrounded by freely moving valence electrons.
metallic bonding is the force of attraction between the positive metal ions and the mobile or delocalised valence electrons

40. Metallic Bonding

41. Metallic Bonding
This theory of metallic bonding is called the ‘Sea of Electrons’ Model or ‘Free Electron’ Model

42. Metallic Bonding This theory accounts for properties of metals
electrical conductivity
- electric current is the flow of electrons
- metals are the only solids in which electrons are free to move
solids
Attractive forces between positive cations and negative electrons are very strong

43. Metallic Bonding malleability and ductility
metals can be hammered into thin sheets(malleable) or drawn into thin wires(ductile).
metallic bonding is non-directional such that layers of metal atoms slide past each other under pressure.

44. Network Covalent Bonding (p. 199)
occurs in 3 compounds (memorize these)
diamond – Cn
carborundum – SiC
quartz – SiO2
large molecules with covalent bonding in 3D
each atom is held in place in 3D by a network of other atoms

45. Network Covalent bonding
Properties:
the highest melting and boiling points
the hardest substances
brittle
do not conduct electric current in any form

46. 1. Network Covalent (Cn ,SiO2 , SiC)
MP & BP decreases
Strongest
1. Network Covalent (Cn ,SiO2 , SiC)
2. Ionic bonding(metal & nonmetal)
3. Metallic bonding (metals)
4. Molecular (nonmetals)
Weakest

47. Valence Shell Electron Pair Repulsion theory (VSEPR)
The shape of molecules is determined by the arrangement of valence electron pairs around the atoms in a compound.
The shapes are the result of REPULSION between pairs of valence electrons.
Valence electron pairs move as far away from each other as possible.

48. Valence Shell Electron Pair Repulsion theory (VSEPR)
There are 5 shapes that can be determined by the # of bonds and # of lone pairs on the central atom.

49. 1. Tetrahedral (4 bonds; 0 lone pairs)

50. 2. Pyramidal (3 bonds; 1 lone pair)

51. 3. V-shaped (2 bonds; 2 lone pairs)

52. 4. Trigonal Planar (3 bonds; 0 lone pairs)

53. 5. Linear (2 bonds; 0 lone pairs)

54. For each molecule below draw the Lewis diagram and the shape diagram.
1 central atom
HOCl H2Se H2SiO
NBr3 CHCl3 SiH4
PBr3 HCN I2
2 central atoms
C2F4 C2H6 CH3OH
C2H2 H2O2

55. Electronegativity (EN - p. 174)
EN is a measure of the attraction that an atom has for shared electrons.
A higher EN means a stronger attraction or electrostatic pull on valence electrons
EN values increase as you move:
from left to right in a period
up in a group or family

56. Increases

57. Electronegativity & Covalent Bonds
polar covalent bond
a bond between atoms with different EN
the shared electron pair is attracted more strongly to the atom with the higher EN δ− δ+ Cl H

58. Electronegativity & Covalent Bonds
nonpolar covalent bond
occurs between atoms with same EN
the shared electron pair is attracted more strongly to the atom with the higher EN
the separation of charge or bond dipole is shown using an arrow pointing toward the more electronegative atom.
the Greek letter delta (δ) indicates ‘partial’ charges
Complete: #’s 7 – 9 on p.178

59. Electronegativity and Ionic Bonds
Because the EN of metals is so low, metals lose electrons to form cations
Nonmetals gain electrons to form anions because their EN is relatively high
When ions form, the resulting electrostatic force is an ionic bond

60. Electronegativity and Covalent Bonds
Atoms in covalent compounds can either have the same EN
eg. Cl2 , PH3, NCl3
OR different EN
eg. HCl

61. Electronegativity and Covalent Bonds
Atoms with the same EN have the same attraction for shared valence electrons.
Covalent bonds resulting from equal sharing of the bonding electron pairs are called Nonpolar Covalent Bonds
Atoms with different EN attract the shared valence electron pair at different strengths.
(higher EN has a stronger attraction
for the shared electron pair)

62. Electronegativity and Covalent Bonds
eg. HCl
Cl has a higher EN
the bonding electron pair is pulled closer to the chlorine atom
this produces slight positive and negative charges within the bond
these charges are referred to as “partial charges” and are denoted with the Greek letter delta (δ).

63. Electronegativity and Covalent Bonds
The region around the chlorine atom will be slightly negative
The region around the hydrogen will be slightly positive.

64. Electronegativity and Covalent Bonds
Because the bond is polarized into a positive area and a negative area the bond has a “bond dipole”.
an arrow points to the atom with the higher EN.
Covalent bonds resulting from unequal sharing of bonded electron pairs are Polar Covalent Bonds.

65. Electronegativity and Covalent Bonds
eg. H2O

66. Electronegativity and Covalent Bonds
eg. HF

67. Electronegativity Homework
p #’s 7, 8, & 9
p #’s 1, 2, & 3

68. Bond Energy (pp )
1. Describe the forces of attraction and repulsion present in all bonds.
2. What is bond length?
3. Define bond energy.
4. Which type of bond has the most energy?
5. How can bond energy be used to predict whether a reaction is endothermic or exothermic?

69. Test Outline Bohr Diagrams (atoms & ions)
Lewis Diagrams (Electron Dot)
Ion Formation
Ionic Bonding, Structures & Properties
Covalent Bonding, Structures & Properties

70. Test Outline Metallic Bonding Theory& Properties
Network Covalent Bonding & Properties
Electronegativity
Bond Dipoles & Polar Molecules
VSEPR Theory
LD, DD, & H-bonding
Predicting properties (bp, mp, etc.)

71. Molecular Dipoles
The vector sum of all the bond dipoles in a molecule is a Molecular Dipole
A Polar Molecule has a molecular dipole that points toward the more electronegative end of the molecule.
eg. H2O

72. Molecular Dipoles Nonpolar molecules DO NOT have molecular dipoles.
This occurs when:
- bond dipoles cancel
- there are no bond dipoles
To determine whether a molecule is polar:
- draw the LD and the shape diagram
- draw the bond dipoles and determine whether they cancel
eg. CO2
eg. PH3

73. Molecular Dipoles
See Handout #1

74. Intermolecular Forces
Mar. 31
Intermolecular Forces

75. Strongest bonds; Highest mp and bp 1. Network Covalent (Cn SiO2 SiC)
Mar. 31
Strongest bonds; Highest mp and bp
1. Network Covalent (Cn SiO2 SiC)
2. Ionic bonding(metal & nonmetal)
3. Metallic bonding (metals)
4. Molecular (nonmetals)
Weakest bonds; Lowest mp and bp
- Intermolecular forces present

76. To compare mp and bp in covalent compounds you must use:
Mar. 31
To compare mp and bp in covalent compounds you must use:
- London Dispersion forces (p. 204)
(all molecules)
- Dipole-Dipole forces (pp. 202, 203)
(polar molecules)
- Hydrogen Bonding (pp. 205, 206)
(H bonded to N, O, or F)

77. Intermolecular Forces (p. 202)
Mar. 31

78. Intermolecular Forces
Mar. 31
Covalent compounds have low mp and bp because forces between molecules in covalent compounds are very weak.
Intermolecular forces were studied by the Dutch physicist Johannes van der Waals
In his honor, two types of intermolecular force are called Van der Waals forces.
Intermolecular forces can be used to explain physical properties of covalent compounds.

79. 1. London Dispersion Forces
Apr. 1
LD forces exist in ALL molecular elements & compounds.
The positive charges in one molecule attract the negative charges in a second molecule.
The temporary dipoles caused by electron movement in one molecule attract the temporary dipoles of another molecule.

80. 1. London Dispersion Forces
Apr. 1
The strength of these forces depends on:
the number of electrons
more electrons produce stronger LD forces that result in higher mp and bp
eg. CH4 is a gas at room temperature.
C8H18 is a liquid at room temperature.
C25H52 is a solid at room temperature.
Account for the difference.

81. 1. London Dispersion Forces
Apr. 1
Two molecules that have the same number of electrons are isoelectronic
eg. C2H6 and CH3F

82. 1. London Dispersion Forces
Apr. 1
b) shape of the molecule
molecules that “fit together” better will experience stronger LD forces
eg. Cl2 vaporizes at -35 ºC while C4H10 vaporizes at -1 ºC. Use bonding to account for the difference.

83. 2. Dipole-dipole Forces occur between polar molecules
Apr. 5
occur between polar molecules
the δ+ end of one polar molecule is attracted to the δ- end of another polar molecule (& vice-versa)
eg. Which has the higher boiling point; CH3F or C2H6 ?

84. Apr. 5
p. 202

85. 3. Hydrogen Bonds a special type of dipole-dipole force
Apr. 5
a special type of dipole-dipole force
(about 10 times stronger)
- only occurs BETWEEN MOLECULES that contain H directly bonded to F, O, or N
ie. the molecule contains at least one H-F, H-O, or H-N covalent bond.

86. 3. Hydrogen Bonds
Apr. 5
the hydrogen bond occurs between the H atom of one molecule and the N, O, or F of a second molecule.
eg. Arrange these from highest to lowest boiling point;
C3H8 C2H5OH C2H5F

87. Apr. 5
p. 206

88. NOTE: To compare mp and bp in covalent compounds you must use:
- London Dispersion forces
(all molecules)
- Dipole-Dipole forces
(polar molecules)
- Hydrogen Bonding
(H bonded to N, O, or F)

89. WorkSheet: Bonding #4 p. 226 #13 & 14
[answers on p. 815 except c), d), & s)]

90. Intermolecular Forces
1. Use intermolecular forces to explain the following:
a) Ar boils at -186 °C and F2 boils at -188 °C .
b) Kr boils at -152 °C and HBr boils at -67 °C.
c) Cl2 boils at -35 °C and C2H5Cl boils at 13 °C .
2. Examine the graph on p. 210:
a) Account for the increase in boiling point for the hydrogen compounds of the Group IV elements.
b) Why is the trend different for the hydrogen compounds of the Group V, VI, and VII elements?
c) Why are the boiling points of the Group IVA compounds consistently lower than the others.

91. 3. Which substance in each pair has the higher boiling point
3. Which substance in each pair has the higher boiling point. Justify your answers.
(a) SiC or KCl
(b) RbBr or C6H12O6
(c) C3H8 or C2H5OH
(d) C4H10 or C2H5Cl

92. 4. Examine the graph on p. 210:
a) Account for the increase in boiling point for the hydrogen compounds of the Group IV elements.
b) Why is the trend different for the hydrogen compounds of the Group V, VI, and VII elements?
c) Why are the boiling points of the Group IVA compounds consistently lower than the other compounds.

93. p. 210

94. Dipole-Dipole Forces
In the liquid state, polar molecules (dipoles) orient themselves so that oppositely charged ends of the molecules are near to one another.

95. Summary
The types of bonding/forces ranked from strongest to weakest are:
Strongest - Network Covalent
- Ionic
- Metallic
Weakest - Covalent

96. NOTE: To compare covalent compounds you must use:
- London Dispersion forces
(all molecules)
- Dipole-Dipole forces
(polar molecules)
- Hydrogen Bonding
(H bonded to N, O, or F)

97. p #’s 13 & 14

98. Dipole-Dipole Forces
The electrostatic attractions between these oppositely charged ends of the polar molecules are called dipole-dipole forces.

99. Dipole-Dipole Forces Results of dipole-dipole attractions:
polar molecules will tend to attract one another more at room temperature than similarly sized non-polar molecules
energy needed to separate polar molecules is therefore higher than for non-polar molecules of similar molar mass

100. Dipole-Dipole Forces Results of dipole-dipole attractions:
The melting points and boiling points of substances made of polar molecules are higher than for substances made of non-polar molecules.

101. Ion-Dipole Forces
An ion-dipole force is the force of attraction between an ion and a polar molecule (a dipole).

102. Ion-Dipole Forces
NaCl dissolves in water because the attractions between the Na+ and Cl- ions and the partial charges on the H2O molecules are strong enough to overcome the forces that bind the ions together.

103. Induced Intermolecular Forces

104. Induced Intermolecular Forces
Induction of electric charge occurs when a charge on one object causes a change in the distribution of charge on a nearby object. (for example, the balloon)

105. Induced Intermolecular Forces
There are two types of charge -induced dipole forces:
1. An ion-induced dipole force results when an ion in close proximity to a non-polar molecule distorts the electron density of the non-polar molecule

106. Induced Intermolecular Forces
The molecule then becomes momentarily polarized, and the two species are attracted to each other. (ie. hemoglobin)
2. In a dipole-induced dipole force the charge on a polar molecule is responsible for inducing the charge on the non-polar molecule.

107. Dispersion (London) Forces
Bond vibrations, which are part of the normal condition of a non-polar molecule, cause momentary, uneven distribution of charge;
a non-polar becomes slightly polar for an instant, and continues to do so in a random but constant basis.

108. Dispersion (London) Forces
At the instant that one non-polar molecule is in a slightly polar condition, it is capable of inducing a dipole in a nearby molecule
This force of attraction is called a dispersion force.

109. Dispersion (London) Forces
Two factors affecting the magnitude of dispersion forces are:
The number of electrons in the molecule:
Vibrations within larger molecules that have more electrons than smaller molecules can easily cause an uneven distribution of charge.

110. Dispersion (London) Forces
The dispersion forces between these larger molecules are thus stronger, which has the effect of raising the boiling point for larger molecules.
The shape of the molecule:
A molecule with a spherical shape has a smaller surface area than a straight chain molecule that has the same number of electrons

111. Dispersion (London) Forces
Therefore, the substance with molecules that have a more spherical shape will have weaker dispersion forces and a lower boiling point.
London dispersion forces are responsible for the formation and stabilization of the biological membranes surrounding every living cell.

112. Hydrogen Bonding
In order to form a hydrogen bond, a hydrogen atom must be bonded to a highly electronegative atom such as oxygen, nitrogen, or fluorine.

113. Hydrogen Bonding
These bonds are very polar, and since hydrogen has no other electrons, the positive proton, H+, is exposed and can become strongly attracted to the negative end of another dipole nearby

114. … Hydrogen Bonding H H H H O O δ− H δ−
A hydrogen bond is an electrostatic attraction between the nucleus of a hydrogen atom, bonded to fluorine, oxygen, or nitrogen and the negative end of a dipole nearby. H δ+ H δ+ H … H O δ+ δ+ O δ− H δ−

115. Hydrogen Bonding
In biological systems, these polar bonds are often parts of much larger molecules (ie. N H bonds and C O bonds found in biological molecules)

116. Hydrogen Bonding in Water
Hydrogen bonds between the hydrogen atoms in one water molecule and the oxygen atom in another account for many unique properties of water. H δ+ H δ+ H … H O δ+ δ+ O δ− H δ−

117. Hydrogen Bonding in Water
In liquid water, each water molecule is hydrogen bonded to at least four other water molecules.
The large number of bonds between water molecules makes the net attractive force quite strong

118. Hydrogen Bonding in Water
the strong attractive forces are responsible for the relatively high boiling point of water.
The water molecules are farther apart in ice then they are in liquid water making ice less dense than liquid water.

119. Hydrogen Bonding in Water
Hydrogen bonds force water molecules into the special hexagonal, crystalline structure of ice when the temperature is below 4 degrees celcius.

120. Assignment # 4

121. Electronegativity
Electronegativity is a result of the space between the nucleus and the electrons
As the number of protons in the nucleus increases, the attractive force on the electrons increases, pulling them closer to the nucleus