1. Covalent Bonds
Other Bonds
James Treasury Sen. Kit
Bond Bond Bond

2. Bonding
A metal and a nonmetal form an ionic bond by the transfer of e-.
Nonmetal atoms form covalent bonds by sharing e-.

3. Bonding
A metal and a nonmetal form an ionic bond by the transfer of e-.
Nonmetal atoms form covalent bonds by sharing e-.
…that’s not the whole story.

4. Consider electronegativity

5. (We can also define % ionic character)
Call a bond with electronegativity difference of:
nonpolar covalent bond
polar covalent bond
1.9 and above -- ionic bond
(We can also define % ionic character)

6. What kind of bond forms between chlorine and phosphorus atoms?

7. What kind of bond forms between chlorine and phosphorus atoms?
Cl P
=.9
This is a polar covalent bond
(use an absolute value for the difference)

8. What kind of bond forms between these atoms?
H and Cl Cl and C
Cl and F F and F
Na and O Mg and N
Mg and Mg N and O

9. The Lewis Diagram Bar represents a covalent bond, 2 shared electrons
Unshared pairs fill out the octets.
Double and triple bars represent double and triple bonds.

10. Lewis diagrams Step 1: Count the total valence electrons available
--use the columns of the periodic chart
--negative ions have extra electrons,
--positive ions are missing electrons
Step 2: Count the total valence electrons needed
--duet rule for hydrogen, or the
--octet rule for everything else
Step 3: Number of bonds = (electrons needed-electrons available) / 2 electrons per bond

11. Lewis diagrams
Step 4: Choose the central atom (almost always the unique one), surround it with the others.
Step 5: Connect with one bond to each outer atom. (PS: Recheck your formula!)
Step 6: Fill in enough multiple bonds to satisfy step 3
Step 7: Draw in unshared pairs to fill valence levels.

12. Don’t
Don’t try to figure out whose electrons are whose. Electrons are identical.
Don’t string the atoms along. Put one atom in the center, unless you have 6 or more atoms.
Don’t EVER put two bonds or an unshared pair on H.

13. Lewis diagrams
Draw a Lewis diagram of hydrogen cyanide, HCN

14. Lewis diagrams
Draw a Lewis diagram of hydrogen cyanide, HCN
H C N

15. Try an ion.
Draw a Lewis diagram for the nitrite, NO2- , ion

16. Lewis diagrams
Draw a Lewis diagram for the nitrite, NO2- , ion
[ O N O ]-

17. Resonance
Which one is preferable?
[ O N O ]-

18. (the double-headed arrow signifies resonance)
Each is valid. The multiple bond exists in both locations. This is called resonance.
[ O N O ]-
(the double-headed arrow signifies resonance)

19. Resonance
Draw three resonance structures for carbon dioxide.

20. Formal Charge (#valence e-) - (#bonds/2 + unshared e-)
Try to minimize each. The more electronegative atom gets a more negative formal charge.

21. Formal Charge
Draw two valid arrangements each for the atoms in:
H2CO
H2O2
N2O
(count bonds, put them in, fill in lone pairs)

22. Formal Charge
Which is preferable? H H O C H O C H

23. Formal Charge Which is preferable? H H O C H O C H

24. Formal Charge Choose this one! Which is preferable? H H O C H O C H

25. Formal Charge
Which is preferable? H H O O H O O H

26. Formal Charge Which is preferable? H H O O H O O H

27. Formal Charge Choose this one! Which is preferable? H H O O H O O H

28. Formal Charge
Which is preferable? O N N N O N

29. Formal Charge
Which is preferable? O N N N O N or

30. Formal Charge Choose this one! Which is preferable? O N N N O N

31. Formal Charge
Which is preferable? O N N O N N

32. Formal Charge
Which is preferable? O N N O N N or

33. Formal Charge Choose this one! Which is preferable? O N N O N N

34. Coordinate covalent bonds
How did this
become this? C O C O

35. Coordinate covalent bonds
How did this
become this?
Carbon monoxide really does have the third bond. The oxygen donates both electrons to share. This is a coordinate covalent bond C O C O C O

36. Coordinate covalent bonds
Draw a Lewis diagram of the ozone (O3) molecule. Count the formal charge for each atom and mark a coordinate covalent bond

37. Coordinate covalent bonds
Draw a Lewis diagram of the ozone (O3) molecule. Count the formal charge for each atom and mark a coordinate covalent bond O O O

38. Exceptions to the octet rule
Draw a Lewis diagram for the triiodide ion, I3-

39. Exceptions to the octet rule
Draw a Lewis diagram for the triiodide ion, I3-
Gadzooks!
When you try to find the number of bonds, (24-22)/2=1 bond.
That’s not enough to tie the ion together.

40. Exceptions to the octet rule
When that happens—go old school. Circle your electrons
[ I I I ]-

41. Exceptions to the octet rule
When that happens—go old school. Circle your electrons
[ I I I ]-

42. Exceptions to the octet rule
When that happens—go old school. Circle your electrons
[ I I I ]-
Two single bonds will satisfy the outer two iodine atoms, the middle one breaks the octet rule (with 10 electrons).

43. Exceptions to the octet rule
Draw a Lewis diagram for XeF4
(The point here is to find out how many unshared pairs are on the central atom)

44. Exceptions to the octet rule
Draw a Lewis diagram for XeF4
(The point here is to find out how many unshared pairs are on the central atom) F F Xe F
Start here day 2 purple F

45. Polar bonds We use a symbol to show a polar covalent bond.
The arrow points toward the more electronegative atom, the (+) end is less electronegative O H H H

46. Polar bonds Or, mark the molecule’s (+) and (-) parts
The d is the small Greek delta. It indicates a small change. In this case, a partial charge d- O d+ H d+ H H

47. Three properties of polar bonds:
The less electronegative end of a polar bond: d+ d-
H Cl
--is more positive
--cannot attract the electrons as well
--is farther from the shared pair of electrons
Start here white day

48. Molecular Shapes
Most molecules have a central atom that follows the octet rule. This allows the following shapes.
Tetrahedral
Trigonal pyramid (trigonal=having three
Bent corners)
Linear and
Trigonal planar

49. Molecular Shapes
Four bonds in four directions makes a tetrahedral shape

50. Molecular Shapes
Three bonds and one lone pair in four directions makes a trigonal pyramid shape

51. Molecular Shapes
Two bonds and two lone pairs in four directions makes a bent shape

52. Molecular Shapes
A double bond holds two electron pairs in the same direction. With no lone pairs, this makes a trigonal planar molecule

53. Molecular Shapes
One lone pair, with a single and a double bond gives a bent molecule.

54. Molecular Shapes
Two double bonds, or a single and a triple makes a linear molecule

55. Molecular Shapes
Two atoms are always in a straight line, a linear molecule.

56. Look for double bonds and unshared pairs
If A=central atom, B=atoms bonded to it, E=e- pairs:
AB4—tetrahedral
(no double bonds)
AB3E-trigonal pyramid
AB2E2 –bent
ABE3 –linear
AB3— trigonal planar
(one double bond)
AB2E –bent
ABE2—linear
AB2—linear
(2 doubles or 1 triple)
ABE—linear (“)
Look for double bonds and unshared pairs

57. Determine the shape of each molecule and ion on the lab that has a single central atom.

58. Polarity of molecules
When polar bonds are not cancelled by symmetry, you get a polar molecule. A polar molecule has (+) and (-) parts.
POLARITY is the first property to look for when analyzing a molecule !

59. Polarity CH4 has no polar bonds. It is symmetric
PH3 has no polar bonds It is not symmetric
CO2 has polar bonds. It is symmetric
H2O has polar bonds. It is not symmetric

60. Polarity CH4 has no polar bonds. It is symmetric Not polar!
PH3 has no polar bonds It is not symmetric
CO2 has polar bonds. It is symmetric
H2O has polar bonds. It is not symmetric
Polar!

61. Mark each molecule on the lab that is polar.
For those that are not polar—why not?
(PS—don’t even look at the ions. If it has a whole charge, ignore the partial charges)

62. (Hydrogen is the only exception)
Hybridization
Atomic orbitals combine to form hybrid orbitals before bonding
(Hydrogen is the only exception)

63. Before bonding The first step is a hybridization of the valence level
H H H H
forms…
p orbitals
s orbitals

64. Hybridization The first step is a hybridization of the valence level C
H H H H
The s and p orbitals hybridize to form sp3 orbitals. The sp3 designation shows one s orbital and 3 p orbitals make the new ones
sp3 orbitals

65. Hybridization The first step is a hybridization of the valence level C
H H H H
The number of orbitals is preserved
(4 in  4 out)
sp3 orbitals

66. Hybridization C
H H H H
All four bonds are identical. Methane is a symmetrical molecule.

67. sp2 Hybridization
When one p orbital is left out of the hybridization, it is used to make a double bond C
H H O
…forms….
p orbitals
s orbitals

68. sp2 Hybridization
When one p orbital is left out of the hybridization, it is used to make a double bond C
H H O
sp2 orbitals
Unused p orbitals—will form the second bond between C and O

69. sp2 Hybridization
When one p orbital is left out of the hybridization, it is used to make a double bond C
H H O
Makes the double bond!

70. sp2 Hybridization H C O::
Carbon & oxygen share electrons in unused p orbitals H
C O::
Carbon shares electrons in sp2 orbitals

71. sp2 Hybridization p bond s bonds H C O::
Carbon & oxygen share electrons in unused p orbitals H
C O::
Carbon shares electrons in sp2 orbitals
p bond
s bonds

72. sp Hybridization
When two p orbitals are left out of the hybridization, it is used to make two double bonds, or a triple bond C
O O
…forms….

73. sp Hybridization
When two p orbitals are left out of the hybridization, it is used to make two double bonds, or a triple bond C
O O
sp orbitals
Unused p orbitals
sp2 orbitals

74. sp Hybridization
When two p orbitals are left out of the hybridization, it is used to make two double bonds, or a triple bond C
O O

75. sp Hybridization p bonds
When two p orbitals are left out of the hybridization, it is used to make two double bonds, or a triple bond C
O O

76. sp Hybridization ::O C O::
Carbon & oxygen share electrons in unused p orbitals
::O C O::
Carbon shares electrons in sp orbitals

77. Look for multiple bonds!
#of Multiples Bonding patterns Hybridization
None AB4, AB3E, sp3
AB2E2, ABE3
One AB3, AB2E sp2
ABE2
Two AB2, ABE sp

78. Look for multiple bonds!
#of Multiples Bonding patterns Hybridization
None sp3
One sp2
Two sp

79. What is the hybridization of the carbon atoms in…
CCl4
H2CO
C2H6
C2H4
C2H2 CO
CH3OH
HCOOH

80. Molecular Orbital Theory (MOT)
Overlapping s orbitals, or hybridized orbitals makes a s (sigma) bond
The electron density is on the SAME line as the nuclei s* s s s

81. Molecular orbitals Overlapping p orbitals, makes a p (pi) bond
The electron density is on a PARALLEL line to the line of the nuclei p p p + + p*

82. Molecular orbitals A single bond is a s bond
A double bond is a s bond, and a p bond above and below the s
A triple bond is a s bond, with two p bonds– above/below and front/back

83. For every bonding molecular orbital (s or p) an antibonding orbital is formed (s* or p*)
A bond is formed when there are more bonding than antibonding electrons

84. VSEPR
Valence Shell Electron Pair Repulsion
Theory

85. VSEPR Valence Shell Electron Pair Repulsion Theory
--pronounced “Vesper”
Electron pairs repel each other. Just as it says.

86. VSEPR is used to predict bond angles
VSEPR is used to predict bond angles. The pairs will space themselves out as far as possible.
A lone pair will take as much room as a bond AND MORE!
Consider sp3 hybridization

87. AB4—like methane. Tetrahedral 109.5o
AB3E—like ammonia. Pyramidal 107o
AB2E2—like water. Bent o angles
--the unshared pairs force the bonds closer together—bond angles decrease

88. With sp2 hybridization:
AB3—like carbonate. Trigonal planar: 120o
AB2E—like nitrite. Bent: less than 120o
ABE2—like O2(2 atoms, has to be linear)

89. With sp hybridization:
AB2—like carbon dioxide. Linear: 180o
ABE—like carbon monoxide. Linear: 180o

90. …but that’s just if you always follow the rules…
– like the octet rule.

91. With dsp3 hybridization:
AB5—trigonal bipyramid
AB4E—seesaw
AB3E2—t-shaped
AB2E3—linear
ABE4—linear

92. With d2sp3 hybridization:
AB6— octahedral
AB5E—square pyramid
AB4E2—square planar
AB3E3—t-shaped
AB2E4 , ABE5—linear

93. What is the shape of… All of the molecules and ions on the lab?
I3-, SF6, XeF4, PCl5, IF5?

94. Count the s and p bonds in the following molecule
Count the s and p bonds in the following molecule. Label each bond as s or p H H H C C C C C H H H O

95. 11 3
Count the s and p bonds in the following molecule. Label each bond as s or p H H H C C C C C H H H O

96. Determine the hybridization of the carbons and the oxygen atom

97. Determine the hybridization of the carbons and the oxygen atom
sp sp sp3 sp2 sp3
sp2 H H H C C C C C H H H O

98. The molecular aufbau diagram

99. The molecular aufbau order
s1s2s*1s2s2s2s*2s2s2px2p2py,z4p*2py,z4s*2px2….
For example:
O2 has 16 electrons. Its electron configuration is:
O2 s1s2s*1s2s2s2s*2s2s2px2p2py,z4p*2py,z2

100. The molecular aufbau order
What is the electron configuration of… N2 NO
Ne2
Remember: we couldn’t do a Lewis diagram with an odd number of electrons!

101. The molecular aufbau diagram

102. The molecular aufbau diagram
These two can switch places—no effect on bonding, but it causes magnetic effects we can measure

103. The molecular aufbau diagram
These two can switch places—no effect on bonding, but it causes magnetic effects we can measure

104. Bond order
The order of a bond in a diatomic molecule is half the number of shared electrons not cancelled by antibonding electrons.
Or:
(number of bonding e- in the atoms-antibonding e-)/2

105. Bond order What is the bond order of… N2 NO Ne2
Remember: we couldn’t do a Lewis diagram with an odd number of electrons!

106. You will be responsible for:
Writing the molecular orbital electron configuration and
Calculating the bond order…
…of any pair of atoms from the second period as they attempt to form a diatomic molecule.

107. Bond Energies
The energy it takes to break a bond is the amount of energy released as the bond is formed.
--measured in kJ/mol
--can be used to estimate DHrxn
--can be absorbed or emitted as light.

108. What is the DHf of NH3?

109. What is the DHf of NH3?
Write the reaction
N2 + 3H2 2NH3

110. What is the DHf of NH3? Count the bonds made and broken N2 + 3H2 2NH3
1 NN triple bond, 3 HH single bonds broken
6 NH single bonds made

111. What is the DHf of NH3? Look up bond energies, and find a total
N2 + 3H2 2NH3
1 molx941kJ/mol+3 molx436kJ/mol= 2249kJ used
6 molx393 kJ/mol=2358 kJ released

112. What is the DHf of NH3? Find the difference, express as kJ/mol
N2 + 3H2 2NH3
2358 kJ-2249kJ=109 kJ more is released, as 2mol NH3 is produced, DHf=-109kJ/2mol=-55kJ/mol

113. It’s an estimate.
My book claims -46 kJ/mole.

114. What is the heat of reaction for:
CH4 + O2 H2O +CO2

115. What is the heat of reaction for:
CH4 +2O2 2H2O +CO2

116. What is the heat of reaction for:
CH4 +2O2 2H2O +CO2
Break 2 O=O and 4 C-H
Form 4 H-O and 2 C=O

117. What is the heat of reaction for:
CH4 +2O2 2H2O +CO2
Break 2 O=O and 4 C-H
kJ +
Form 4 H-O and 2 C=O kJ

118. What is the heat of reaction for:
CH4 +2O2 2H2O +CO2
Break 2 O=O and 4 C-H
kJ +
Form 4 H-O and 2 C=O kJ
-882kJ/mol

119. What is the heat of reaction for:
2H2O2 2H2O +O2

120. Tasks

121. On your test, you will be asked to:
Describe how and why ionization of metals and non-metals occurs
Write EC’s for atoms and ions
Show formation of ionic and covalent bonds by electron dot diagrams
Describe metallic bonding
Define and identify electrolytes

122. On your test, you will be asked to:
Identify particles and types of substances by bonding
Draw Lewis diagrams
Identify shapes of molecules and ions
Identify types of bonds between atoms
Describe polarity

123. On your test, you will be asked to:
Identify polar and nonpolar molecules
Identify hybridizations
Describe single and double bonds by MOT
Estimate DHrxn by bond energies
Use VSEPR to predict molecular shapes and bond angles.
Calculate and justify bond orders for diatomics from the second period