1. Chapter 8 Covalent Bonding
Key: C O
H N
Ball-and-stick model of epinephrine (adrenaline)

2. 8.1 Molecular Compounds OBJECTIVES:
Distinguish between the melting points and boiling points of molecular compounds and ionic compounds
Describe the information provided by a molecular formula

3. Bonds are…
Forces that hold groups of atoms together and make them function as a unit. Two types:
Ionic bonds – transfer of electrons results in oppositely charged particles that attract one another.
Ionic solids are crystals with regular arrangements of +/- ions
In solutions, cations and anions exist separately – there are NO molecules!
Covalent bonds – sharing of electrons. The resulting particle is called a molecule = individual units of a covalent compound

4. Sodium Chloride Crystal Lattice
Ionic compounds organize in a characteristic crystal lattice of alternating positive and negative ions, repeated over and over.

5. 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” (O2)

6. Formation of a Covalent Bond
(Video: Formation of H2)

7. Covalent Bonds Nonmetals hold on to their valence electrons.
They acquire a stable, noble gas configuration by sharing valence electrons with each other = covalent bonding
By sharing, both atoms “count” the shared electrons toward a noble gas configuration

8. F Covalent Bonding Fluorine has seven valence electrons
(but would more stable with 8) F

9. F F Covalent bonding Fluorine has seven valence electrons
A second F atom with seven… F F

10. F F Covalent bonding Fluorine has seven valence electrons
A second atom also has seven
If they share a pair of electrons… F F

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

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

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

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

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

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

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

18. Molecular Compounds
Compounds that are bonded covalently (like H2O, CO2) are called molecular compounds
Molecular compounds tend to have relatively lower melting points than ionic compounds
Ions are held together in a crystal solid by relatively strong ionic bonds
Covalent bonds are very strong but this is an intramolecular force
It’s the forces between molecules that holds a solid together (intermolecular forces)
and these are often rather weak

19. Molecular Compounds
Thus, molecular compounds are often gases or liquids at room temperature
Ionic compounds are solids
A molecular compound has a molecular formula:
Shows how many atoms of each element a molecule contains

20. Molecular Compounds The formula for water is H2O
The subscript “2” after 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).

21. - 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!

22. 8.2 Nature of Covalent Bonding
OBJECTIVES
Describe how electrons are shared to form covalent bonds, and identify exceptions to the octet rule
Demonstrate how electron dot structures represent shared electrons
Describe how atoms form double or triple covalent bonds

23. 8.2 Nature of Covalent Bonding
OBJECTIVES:
Distinguish between a covalent bond and a coordinate covalent bond, and describe how the strength of a covalent bond is related to its bond dissociation energy
Describe how oxygen atoms are bonded in ozone

24. A Single Covalent Bond is...
A sharing of two valence electrons.
Only nonmetals (including H).
Different from ionic bonds because discrete molecules are formed
Two atoms are joined

25. Sodium Chloride Crystal Lattice
Ionic compounds organize in a characteristic crystal lattice of alternating positive and negative ions, repeated over and over.

26. 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 – it can share 4 electrons: 1s22s22p2?
Let’s show how water is formed with covalent bonds, by using an electron dot diagram…

27. H O Water Each hydrogen has 1 valence electron
- Each hydrogen would be stable if it had 1 more
The oxygen has 6 valence electrons
- Each would be stable with 2 more electrons
They share to give each other complete octets H O

28. H O Water Put the pieces together The first hydrogen is stable
The oxygen still needs one more H O

29. H O H Water So, a second hydrogen attaches
Every atom has full energy levels
Note the “unshared” pairs of electrons (also called lone pairs) H O H

30. Examples:
Conceptual Problem 8.1 on page 220
Do PCl3

31. 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?

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

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

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

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

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

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

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

39. O C O Carbon dioxide The only solution is to share more
Requires two double bonds
all the electrons in each double bond “count” for each atom O C O

40. 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

41. 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

42. 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

43. Covalent Bonding
Note how the formation of covalent bonds fills the ½ full s and p orbitals of the bonding atoms.

44. How to draw Lewis Structures?
Add up all the valence electrons.
Count up the total number of electrons needed to give all the atoms a noble gas configuration.
Subtract #2 - 1; then Divide by 2
Tells you how many bonds to draw
Fill in the rest of the valence electrons to “fill up” the atoms’ valence shells.

45. N H Example NH3, which is ammonia
N – central atom; has 5 valence electrons, needs 8
H - has 1 (x3) valence electrons, needs 2 (x3)
NH3 has 5+3 = 8
NH3 needs 8+6 = 14
(14-8)/2= 3 bonds
4 atoms with 3 bonds N H

46. H H N H Examples Draw in the bonds; start with singles
All 8 electrons are accounted for
Everything is full – done with this one. H H N H

47. Example: HCN HCN: C is central atom
N - has 5 valence electrons, needs 8
C - has 4 valence electrons, needs 8
H - has 1 valence electron, needs 2
HCN has = 10
HCN needs = 18
(18-10)/2= 4 bonds
3 atoms with 4 bonds – this will require multiple bonds - not to H though

48. H C N HCN Put single bond between each atom Need to add 2 more bonds
Must go between C and N (Hydrogen is full) H C N

49. H C N HCN Put in single bonds Must have 2 more bonds
Must go between C and N, not the H
Uses 8 electrons – need 2 more to equal the 10 it has H C N

50. H C N HCN Put in single bonds Need 2 more bonds
Must go between C and N
Uses 8 electrons - 2 more to add
Must go on the N to fill its octet H C N

51. 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 =

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

53. 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

54. 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. C O
These two electrons make a lone pair.

55. 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

56. Coordinate covalent bond
Most polyatomic cations and anions contain covalent and coordinate covalent bonds
Table 8.2, p.224
Sample Problem 8.2, p.225
The ammonium ion (NH41+) can be shown as another example

57. Bond Dissociation Energies...
The total energy required to break the bond between 2 covalently bonded atoms
High dissociation energy usually means the chemical is relatively unreactive, because it takes a lot of energy to break it down.

58. Resonance is... When more than one valid dot diagram is possible.
Consider the two ways to draw ozone (O3)
Which one is it? Does it go back and forth?
It is a hybrid of both, like a mule; and shown by a double-headed arrow
found in double-bond structures!

59. Resonance in Ozone
Note the different location of the double bond
Neither structure is correct, it is actually a hybrid of the two. To show it, draw all varieties possible, and join them with a double-headed arrow.

60. Resonance
Occurs when more than one valid Lewis structure can be written for a particular molecule (due to position of double bond)
These are resonance structures of benzene.
The actual structure is an average (or hybrid) of these structures.

61. Polyatomic ions – note the different positions of the double bond.
Resonance in a carbonate ion (CO32-):
Resonance in an acetate ion (C2H3O21-):

62. The 3 Exceptions to Octet rule
For some molecules, it is impossible to satisfy the octet rule
#1. usually when there is an odd number of valence electrons
NO2 has 17 valence electrons, because the N has 5, and each O contributes 6. Note “N” page 228
It is impossible to satisfy octet rule, yet the stable molecule does exist

63. Exceptions to Octet rule
Another exception: Boron
Page 228 shows boron trifluoride, and note that one of the fluorides might be able to make a coordinate covalent bond to fulfill the boron
#2 -But fluorine has a high electronegativity (it is greedy), so this coordinate bond does not form
#3 -Top page 229 examples exist because they are in period 3 or beyond

64. 8.4 Polar Bonds & Molecules
OBJECTIVES:
Describe how electronegativity values determine the distribution of charge in a polar molecule
Describe what happens to polar molecules when they are placed between oppositely charged metal plates
Evaluate the strength of intermolecular attractions compared with the strength of ionic and covalent bonds
Identify the reason why network solids have high melting points

65. Bond Polarity Covalent bonding means shared electrons
but, do they share equally?
Electrons are pulled, as in a tug-of-war, between the atoms’ nuclei
In equal sharing (such as diatomic molecules), the bond that results is called a nonpolar covalent bond

66. Bond Polarity
When two different atoms bond covalently, there is almost always an unequal sharing of the bonding e-
the more electronegative atom has a stronger attraction for the bonding e-
Thus the bonding pair of e- spend most time near the more electronegative atom  acquire a partial (−) charge
called a polar covalent bond, or simply polar bond.

67. Electronegativity?
The ability of an atom in a molecule to attract shared electrons to itself.
Linus Pauling

68. Table of Electronegativities
Higher electronegativity

69. Bond Polarity Consider HCl H: electroneg = 2.1 Cl: electroneg = 3.0
the bond is polar
the chlorine acquires a slight negative charge, and the hydrogen a slight positive charge

70. Bond Polarity d+ d- d+ and d-
Only partial charges, less than a true 1+ or 1− as in ionic bond
Written as:
H Cl
the positive and minus signs (with the lower case delta: ) denote partial charges.
d+ d-
d+ and d-

71. Bond Polarity H Cl Can also be shown:
*arrow points to more electronegative atom
H Cl

72. Polar molecules Sample Problem 8.3, p.239
A polar bond may make the entire molecule “polar”
Depends on shape of molecule
HCl has polar bonds, and one side is partially +, the other partially –, so the whole molecule is polar
A molecule like HCl that has two poles is said to have a dipole

73. Practice Problems

74. Polar molecules
The effect of polar bonds on the polarity of the entire molecule depends on the molecule shape
carbon dioxide has two polar bonds, and is linear = nonpolar molecule!

75. Polar molecules
The effect of polar bonds on the polarity of the entire molecule depends on the molecule shape
water has two polar bonds and a bent shape; the highly electronegative oxygen pulls the e- away from H’s = very polar!

76. Polar molecules
When polar molecules are placed between oppositely charged plates, they tend to become oriented with respect to the positive and negative plates.
With current off With current on

77. Intermolecular Forces
= attractions between molecules
They are what determine if a compound will exist as a solid liquid or gas at room temperature
The weakest are called van der Waal’s forces - there are two kinds:

78. Van der Waal’s Forces #1. Dispersion forces
Weakest, caused by random motions of e-
Strength increases as # e- increases (size of atom/molecule)
Exist in all compounds
halogens F and Cl are gases; bromine is liquid; iodine is solid – (Group 7A)

79. Dispersion Forces Animation
Link to Dispersion Forces Animation

80. Dipole-dipole attraction (weak)
Van der Waal’s Forces
#2 Dipole Interactions
Occur when polar molecules attract each other
Usually, slightly stronger than dispersion forces
Opposites attract, but these are only partial charges (), so attractions NOT as strong as ionic bonds + −
Dipole-dipole attraction (weak)

81. #3. Hydrogen bonding
…is the attractive force between a strongly + H atom (polar bonded to a N, O, F), and a lone pair of electrons on a N, O, or F atom of another molecule.
N, O, F, are very electronegative (so strong – charge)
Strongest type of dipole attraction
The H atom “shares” a lone pair of electrons with an O, N, or F atom in another molecule
On average, the strongest of the 3 IM forces.

82. #3. Hydrogen bonding defined:
When a hydrogen atom is:
a) covalently bonded to a highly electronegative atom (O, N, or F), AND
b) is also weakly bonded to an unshared electron pair of a nearby highly electronegative atom
The hydrogen is very electron deficient (its lone electron is being shared in a covalent bond with a highly electronegative atom)
so it is strongly attracted to electrons on electronegative atoms on other molecules
Hydrogen atom has no shielding for its nucleus, so that proton is strongly attracted to any nearby electrons

83. Hydrogen Bonding H O d+ d- H O d+ d-
This H is bonded covalently to the O in the same molecule, and H-bonded to a nearby unshared pair on the O atom of another molecule.

84. H-bonding causes H2O to be a liquid at room conditions
H-bonds form between H2O and NH3 molecules

85. Order of Strength: IM Forces
On average:
Dispersion forces are weakest
Dipole interactions a little stronger
Hydrogen bonding strongest
All 3 are weaker than ionic or covalent bonds
All are variable: not all dispersion forces alike, not all H-bonds the same strength
For example: dispersion forces are weakest, but in large molecules, they can be significantly stronger than the other IM forces of smaller molecules

86. Attractions and properties
Why are some chemicals gases, some liquids, some solids at room temperature?
Depends on the type of bonding!
Characteristic
Ionic Compounds
Covalent Compounds
Representative Unit
Formula unit
Molecule
Bond Formation
Electrical attraction between oppositely charged ions
Attraction between atoms for shared electrons
Type of Elements
Metals and Non-metals
Nonmetals
Usual Physical State
Solid
Solid, liquid or gas
Forces Holding Particles Together in a Solid (IM Forces)
Ionic Bonds
Dispersion forces, dipole-dipole, H-bond, covalent bond
Melting Point
High (usually > 300ºC)
Low (usually < 300ºC)

87. Attractions and properties
Network covalent solids – solids in which all the atoms are covalently bonded to each other
The force holding the particles together in a solid (IM force) is the strongest type of chemical bond!

88. Attractions and properties
Network covalent solids melt at very high temperatures, or not at all (decompose instead)
Diamond does not really melt, but vaporizes to a gas at 3500oC
SiC, used in grinding, has a melting point of about 2700oC

89. Covalent Network Compounds
Some covalently bonded substances DO NOT form discrete molecules.
Diamond, a network of covalently bonded carbon atoms
Graphite, a network of covalently bonded sheets of carbon atoms

90. 8.4 Polar Bonds & Molecules
OBJECTIVES (Accomplished??)
Describe how electronegativity values determine the distribution of charge in a polar molecule
Describe what happens to polar molecules when they are placed between oppositely charged metal plates
Evaluate the strength of intermolecular attractions compared with the strength of ionic and covalent bonds
Identify the reason why network solids have high melting points