1. Unit 2: Chemical Bonding
Chemistry2202

2. Outline Bohr diagrams Lewis Diagrams Types of Bonding Ionic bonding
Covalent bonding (Molecular)
Metallic bonding
Network covalent bonding

3. Types of Bonding (cont’d)
London Dispersion forces
Dipole-Dipole forces
Hydrogen Bonding
VSEPR Theory (Shapes)
Physical Properties

4. 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

5. Lewis Diagrams
LD provide a method for keeping track of electrons in atoms, ions, or molecules
Also called Electron Dot diagrams
the nucleus (p+& n0) and filled energy levels are represented by the element symbol

6. Lewis Diagrams
dots are placed around the element symbol to represent valence electrons

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

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

9. Lewis Diagrams
eg. Lewis Diagram for C • • C • •

10. Lewis Diagrams
eg. Lewis Diagram for P • • • P • •

11. Lewis Diagrams
eg. Lewis Diagram for Na • Na

12. 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

13. 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

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

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

16. 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

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

18. Why is graphite soft enough to write with while diamond is the hardest substance known even though both substances are made of pure carbon?

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

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

21. -30 ºC

22. Bonding
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

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

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

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

26. 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

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

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

29. Ionic Bonding

30. 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.

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

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

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

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

35. Metallic Bonding (p. 171) metals tend to lose valence electrons.
valence electrons are loosely held and frequently lost from metal atoms.
This results in 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

36. Metallic Bonding

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

38. 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

39. 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.

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

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

42. 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

43. Valence Shell Electron Pair Repulsion theory (VSEPR)
The shape of molecules is caused by repulsion between valence electron pairs around the atoms in a compound.
Repulsion between valence electron pairs force them to move as far away from each other as possible.

44. Valence Shell Electron Pair Repulsion theory (VSEPR)
To determine molecular shapes, count the # of bonds and # of lone pairs on the central atom(s).
We will examine 5 molecular shapes

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

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

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

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

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

50. For each molecule below draw the Lewis diagram and the shape diagram.
HOCl H2Se H2O2
NBr3 C2F4 C2H6
CHCl3 CH3OH PBr3
I2 SiH4 HCN
SiH2O C2H2

51. Electronegativity & Covalent Bonds
Electronegativity - EN - p. 174
EN measures 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

52. Increases

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

54. Electronegativity & Covalent Bonds
nonpolar covalent bond
a bond between atoms with the same EN
the bonding pair is shared equally between the atoms
Complete: #’s 7 – 9 on p.178

55. Electronegativity & Covalent Bonds
polar molecule
- a molecule in which the bond dipoles do not cancel each other
- a polar molecule has a molecular dipole that points toward the more electronegative end of the molecule.
eg. HCN

56. Electronegativity & Covalent Bonds
nonpolar molecule
- a molecule in which the bond dipoles cancel each other OR
- there are no bond dipoles
eg. CO2 PH3

57. Electronegativity & Covalent Bonds
To determine whether a molecule is polar:
- draw the Lewis diagram & shape diagram
- draw the bond dipoles & determine whether they cancel

58. Intermolecular Forces

59. 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

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

61. Intermolecular Forces (p. 202)

62. Intermolecular Forces
Covalent compounds have low mp and bp because attractive forces between molecules are very weak.
Intermolecular forces were studied extensively by the Dutch physicist Johannes van der Waals
In his honor, two types of intermolecular force are called Van der Waals forces.

63. Intermolecular Forces
Intermolecular forces can be used to account for the physical properties of covalent compounds.

64. Intermolecular Forces

65. 1. London Dispersion Forces
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.

66. 1. London Dispersion Forces
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.

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

68. 1. London Dispersion Forces
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.

69. 2. Dipole-dipole Forces - 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 ?

70. p. 202

71. 3. Hydrogen Bonds
- a special type of dipole-dipole force (about 10 times stronger)
- only occurs between molecules that contain a H atom which is directly bonded to F, O, or N
ie. the molecule contains at least one H-F, H-O, or H-N covalent bond.

72. 3. Hydrogen Bonds
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

73. p. 206

74. 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)

75. Alchem worksheet pp. G32, 33

76. p. 226 #13 Omit parts g), j) – o), q), u), & v)
- Answers on p. 815 for #13
- Incorrect answers
c), d), & s)

77. p. 210

78. 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.

79. 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

80. 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 other compounds.

81. p #’s 13 & 14

82. 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

83. 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 the EN of nonmetals is relatively high

84. Electronegativity and Ionic Bonds
When ions form, the resulting electrostatic force is an ionic bond

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

86. Electronegativity and Covalent Bonds
Atoms that have different EN attract the shared pair of valence electrons at different strengths
The atom with the higher EN exerts a stronger attraction on the shared electron pair
eg. H2O

87. Electronegativity and Covalent Bonds
Since the oxygen atom has a higher EN the bonding electrons will be pulled closer to the oxygen atom
This results in 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).

88. Electronegativity and Covalent Bonds
The region around the oxygen atom will be slightly negative, and around the hydrogens will be slightly positive

89. Electronegativity and Covalent Bonds
The symbol, δ+ represents a partial positive charge (less than +1) and δ− represents a partial negative charge (less than −1).
Since the bond is polarized into a positive area and a negative area the bond has a “bond dipole”.

90. Electronegativity and Covalent Bonds
The arrow points to the atom with the higher EN.
p.178

91. Electronegativity and Covalent Bonds
Covalent bonds resulting from unequal (electronegativities) sharing of bonding electron pairs are called Polar Covalent Bonds

92. Electronegativity Homework
#’s 7, 8, & p. 178
#’s 1, 2, & p. 180

93. 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?

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

95. 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.)