1. Today is Tuesday, March 24th, 2015
In This Lesson:
Covalent Bonding, VSEPR Theory, and Intermolecular Forces
(Lesson 4 of 4)
Stuff You Need:
See Below
Today is Tuesday, March 24th, 2015
Pre-Class:
What kind of bond joins the two oxygen atoms in oxygen gas – O2?
What kind of bond joins the carbon and oxygen atoms in carbon dioxide – CO2?
Please take a video worksheet from the Turn-In Box and get a paper towel. Also find your bonding overview worksheets (the big grid one).

2. Today’s Agenda Covalent/Molecular Bonding Bond/Molecular Polarity
VSEPR Theory
Intermolecular Forces
Where is this in my book?
P. 213 and following…

3. By the end of this lesson…
You should be able to describe the nature of a molecular bond.
You should be able to describe the shape of a molecule based on its constituent atoms.
You should be able to identify forces between molecules based on the nature of the participating molecules.
You should be able to draw resonance structures of molecules.

4. What you need to know…
…is on this checklist.

5. Reminder
This is an atom H O
This is a molecule

6. Video Introduction
Find your Bonding Overview sheets and have your video worksheets ready (on the Covalent Bonding side).
It’s another Australian video!

7. Covalent Bonds: The Important Stuff
Covalent bonds are sometimes called molecular bonds.
Covalent bonds generally occur between nonmetals and nonmetals (or metalloids).
Electron pairs are shared (or fought-over).
Bonds can be polar or non-polar.
The term “molecule” is used only for covalent compounds.
Ionic compounds sometimes get the term formula unit.

8. Characteristics of Molecular Compounds
Poor conductors of electricity.
Low melting points.
Polar compounds can dissolve in water.
Non-polar compounds cannot dissolve in water.
Many are gases at room temperature.

9. Ionic or Covalent? But wait, you say.
How can you determine whether a compound is ionic or covalent?
There are two ways.
The “official” way:
By electronegativity.
The “unofficial” way:
By location on the periodic table.

10. The Official Way
Remember that electronegativity is the ability of an atom to attract electrons in a covalent compound.
Why are noble gases not on here?
They don’t form compounds!

11. Determining Bond Type Use the file Periodic Table – Electronegativity.
Compare the electronegativity values of the two atoms.
Ionic: Difference ≥ 1.67
Sometimes people use numbers as high as 2.0.
Polar Covalent: 1.66 ≥ Difference ≥ 0.4
This means that one atom (the more electronegative one) “hogs” the shared electrons.
Non-Polar Covalent: 0.3 ≥ Difference
The atoms play nice.

12. The Unofficial Way
Because ionic compounds occur between cations and anions, they often occur between elements on opposite sides of the periodic table.
Consider both ways when making the distinction between bonds.

13. Identify the Bond Types
NaCl
Ionic (electronegativity difference of 2.1) F2
Non-Polar Covalent (electronegativity difference of 0)
MgCl2
Ionic (electronegativity difference of 1.8)
BaO
Ionic (electronegativity difference of 2.6)
CH4
Polar Covalent (electronegativity difference of 0.4) O2

14. Aside: Bond Type Exceptions?
You will not need to know them for this class, but keep in mind that there are exceptions to this bond type rule.
A notable one is compounds involving Beryllium (Be) – they tend to be covalent despite Beryllium not achieving an octet.
BeCl2, for example, is a covalent compound as confirmed by experimental evidence involving boiling point and behavior.
However, it still likes to form things that look like crystal lattices when it’s a solid. Weird.

15. The Octet Rule
In ionic compounds, electrons are transferred such that atoms achieve the octet, either by losing or gaining.
In covalent compounds, electrons are shared (or fought over) so that each atom has eight.
The key?
Whichever type of compound it is, everyone either needs to lose all their valence electrons or share/gain to eight.

16. Diatomic Molecules
Recall that seven elements on the periodic table naturally bond to themselves.
Elements that exist diatomically:
Br-I-N-Cl-H-O-F
Br2, I2, N2, Cl2, H2, O2, F2
These elements are all bonded covalently to themselves.

17. The HONC Rule* *Can be broken on occasion.
The HONC Rule is an easy way to show how many bonds form for an element.
Hydrogen and Halogens form one covalent bond.
Oxygen (and sulfur) form two covalent bonds.
Either two single bonds or one double bond.
Nitrogen (and phosphorus) form three covalent bonds.
Either a triple bond, a double and single bond, or three single bonds.
Carbon (and silicon) form four covalent bonds.
Either a triple and a single bond, two double bonds, a double and two single bonds, or four single bonds.

18. Facebook?
Back in the day, Facebook had a “visualizer” for your friends.
Basically, it made a web of friend connections.
Those with more connections tended to be in the inner part of the web, those with fewer were on the outside.

19. Facebook

20. Central Atoms
In a molecule, the central atom is the one that forms the most bonds.
According to the HONC Rule, Carbon and Silicon (and others in their group) are most likely to be the central atom.
Similarly, hydrogens and halogens are rarely the central atoms.

21. Now, onto the bonding…
As you know, Lewis Structures (dot notation? Remember?) show the arrangement of valence electrons in an atom.
In covalent bonds, shared/fought over electrons can be shown by two dots ( : ) or a single line ( - ).
IMPORTANT: Each single line (-) connects two dots, so you should count them as two electrons.

22. Methods of Visualization
If you’re using strictly dots, you’re drawing the Lewis Dot Structure.
If you’re using lines, you’re drawing the Structural Formula.
Throughout this lesson, you’ll learn many ways to accomplish one goal. Use what works for you and/or the problem.

23. Single Bonds Single bonds involve two shared electrons (or one pair).
Single bonds are sometimes drawn with a single line.

24. F F F F Single Bond Example F F Each has seven valence electrons.
One pair of electrons is shared. F F

25. Think of it another way…
Shared
Electrons
Lewis Structure

26. Think of it another way…
Shared
Electrons F
Structural Formula

27. Covalent Bonding Practice
Try this example: Hydrochloric Acid
HCl H Cl H Cl

28. Covalent Bonding Practice
Try this example: Iodine I2 I I I I

29. Covalent Bonding Practice
Try this example: Dihydrogen Monoxide
H2O O H O H H H

30. Notice something…
In each of the examples we did, I drew in the non-bonding pairs of electrons.
As we saw in the video, these electrons have a strong influence on the shape of the molecule even without bonding to another atom.
More to come…

31. Covalent Bonding Practice
Now it’s your turn.
Covalent Bonding Worksheet (first page) #2, 4, 5, 7, 9, 10

32. There are two unshared pairs of electrons.
In our all examples (but let’s return to H2O), there are pairs of electrons that are not shared.
There are known as unshared pairs, lone pairs, or non-bonding pairs.
IMPORTANT: Unshared pairs only count on the central atom (the one with the most bonds).
You should still draw them, though. O H H
There are two unshared pairs of electrons.

33. Unshared Pairs Ammonia (NH3) has one unshared pair of electrons.
The electrons are on the nitrogen atom.

34. Unshared Pairs Practice
How many unshared pairs are on the methane (CH4) molecule?
None.

35. Unshared Pairs Practice
How many unshared pairs are on the carbon tetrachloride molecule?
None (carbon is the central atom). Cl Cl C Cl Cl

36. Double Bonds
Double bonds involve four shared electrons (or two pairs).
Double bonds are sometimes drawn with double lines.

37. O O O O Double Bond Example O O Each has six valence electrons.
Two pairs of electrons are shared. O O

38. Covalent Bonding Practice
Try this one: Carbon Dioxide
CO2 O C O O C O

39. Covalent Bonding Practice
Now it’s your turn.
Covalent Bonding Worksheet (first page) #3, 12

40. Triple Bonds
Triple bonds involve six shared electrons (or three pairs).
Triple bonds are sometimes drawn with triple lines.

41. N N  N N Triple Bond Example N N Each has five valence electrons.
Three pairs of electrons are shared.  N N

42. Covalent Bonding Practice
Now it’s your turn.
Covalent Bonding Worksheet #1
Covalent Bonding Worksheet (Reverse)
WARNING: I have done something a little sneaky on Page 2.

43. Drawing Lewis Structures
Bigger molecules can make for complicated Lewis Structure drawings.
Imagine drawing the Lewis Structure for CH3Cl.
Thankfully, there are two systematic ways to do it.

44. Covalent Bonding Tips and Tricks Jot it like you mean it!
Hydrogen and halogens will only ever make one bond. Connect them with single bonds first.
Make sure you don’t use too many electrons! You can only use as many as the atoms bring.
Follow the HONC Rule when possible.
Single bonds count as 2 electrons, double count as 4, triple count as 6.
DO NOT MAKE A QUADRUPLE BOND EVER!

45. CH3Cl: Method #1
Make carbon the central atom (it wants four bonds).
Add up the available valence electrons:
C = 4, H = 1 (3), Cl = 7
TOTAL = 14.
Join peripheral atoms to the central atoms with electron pairs.
Complete octets on atoms other than hydrogen with remaining electrons. H Cl C H H

46. CH3Cl: Method #2
Make carbon the central atom (it wants four bonds).
Add up the total electrons needed:
C = 8, H = 2 (3), Cl = 8
TOTAL = 22.
Subtract available electrons:
C = 4, H = 1 (3), Cl = 7
TOTAL = 14.
= 8.
Divide by two. (8/2 = 4)
Draw in 4 total bonds between atoms, then complete octets.
Remember, double bonds count as two and triple as three. H Cl C H H

47. Practice Molecule: HCN H C N

48. Practice
Covalent Bonding and VSEPR Theory Worksheet
#1-3
#4-5

49. VSEPR Theory
So dot structures are just peachy, but they leave out something important.
Something like a third dimension.
The thing is, thinking about things in three dimensions can be difficult.
So…let’s introduce the next topic with a video:
TED: George Zaidan and Charles Morton – What is the Shape of a Molecule?

50. VSEPR Theory
VSEPR Theory (pronounced “vesper”) stands for Valence Shell Electron Pair Repulsion Theory.
You can think of it as a theory that helps scientists predict the shape of a molecule.
Because electrons involved in bonding pairs or lone pairs are both negatively charged, they repel one another and get as far apart as possible.

51. VSEPR Theory Chart Treat this VSEPR Theory Chart like gold.
If you need another, it’s on my website.
This is invaluable and is basically your lifesaver.
For reals.
It summarizes a lot of information that can’t be put into a PowerPoint and I’ll be referring to it frequently.
Unfortunately, it will not be available to you on the test.

52. VSEPR Theory To use the VSEPR Chart I gave you, do this:
Draw the dot structure.
Find:
The central atom.
The number of atoms bonded to the central atom.
The number of lone pairs on the central atom.
Then use the VSEPR Chart to find its shape.

53. VSEPR Chart Example: H2O
We know H2O looks like this:
According to the HONC rule, O forms two bonds and H only forms one, so O is the central atom.
Oxygen has two bonded atoms and two unshared pairs of electrons.
The chart tells us that this is molecule takes on the “bent” geometry. O H H

54. VSEPR Geometry: Linear
1 Bonded Atom or 2 Bonded Atoms with 0 Lone Pairs.
Example: CO2 O C O

55. VSEPR Geometry: Trigonal Planar
3 Bonded Atoms, 0 Lone Pairs.
Example: BBr3 Br B Br Br

56. Wait, what? Al Br Try drawing the dot structure for AlBr3.
Notice anything?
As you probably could have guessed, there are exceptions to the octet rule occasionally.
The most common exceptions are compounds that have only single bonds formed with Group 3A cations.
These compounds do not obey the octet rule and take on trigonal planar geometry.

57. Covalent Bonding Worksheet
Take a look at #9 on the reverse of your Covalent Bonding Worksheet.
If this were Boron (and thus not an ionic compound), this would be a molecular compound and would exhibit the same geometry.
The reason it’s ionic is because of the electronegativity difference between Al and F. ;) F B F F

58. VSEPR Geometry: Trigonal Pyramidal
3 Bonded Atoms, 1 Lone Pair.
Example: PH3 P H H H

59. O H H VSEPR Geometry: Bent 2 Bonded Atoms, 1 or 2 Lone Pairs.
Example: H2O O H H

60. VSEPR Geometry: Tetrahedral
4 Bonded Atoms, 0 Lone Pairs.
Example: CH4 H C H H H

61. Other Crazy Shapes T-Shaped Octahedral Trigonal Bipyramidal
Square Planar
Square Pyramidal
Seesaw

62. VSEPR Summary Bonded Atoms Lone Pairs Shape 1 0, 1, 2, or 3 Linear 2
1 or 2
Bent 3
Trigonal Planar
Trigonal Pyramidal 4
Tetrahedral

63. Practice
Find your Covalent Bonding Worksheet and write in the geometry of each covalently-bonded molecule.
Example: #1 (N2) is linear because it has one bonded atom and one lone pair.
Now try Covalent Bonding and VSEPR Theory Worksheet. #6
Now add the geometry to #4 and 5 from Covalent Bonding and VSEPR Theory.

64. Practice Now try: Covalent Bonding and VSEPR Theory Worksheet
Reverse – Bottom section

65. Practice
Lab – Molecules I

66. Aside: Empty Space
Remember earlier in the year when I told you that we never really touch anything?
As in, repulsive forces between atoms keep us from actually coming in contact with them so we just get really really close?
Well, this VSEPR stuff, the repulsion of electrons, is what’s causing that.
VSEPR Theory is part of the reason you can’t walk through a wall right now.
You actually have plenty of space to sneak your atoms’ nuclei through the wall’s – if it weren’t for the repulsive forces.

67. Review
Quick – what’s the shape?

68. Practice Now try: Covalent Bonding and VSEPR Theory Worksheet
Reverse – Top section

69. Lastly… Finally, let’s do Lewis Structures for covalent ions.
Like polyatomic ions from waaaay back when…

70. Polyatomic Ion Lewis Structures Or other really tricky ones…
Procedure:
Determine total valence electrons.
Draw single bonds between all atoms.
Give bonded atoms octets (except H).
Satisfy the octet rule for the central atom(s).
Any leftover electrons go to the central atom.
If the central atom has six electrons, it double bonds to a bonded atom with a lone pair.
If the central atom has four electrons, it triple bonds to a bonded atom with a lone pair, or makes two double bonds.
BIG HUGE IMPORTANT NOTE!
The HONC rule can be broken, and will always be for ions but never for hydrogen or halogens.

71. Example: Oxalate (C2O42-)
Electrons Left: 24
Electrons Left: 34
Electrons Left: 0
Example: Oxalate (C2O42-) 2-
Draw a single bond between all atoms.
Give bonded atoms octets.
If a central atom has only six total electrons, it double bonds to a bonded atom with a lone pair.
Find the total e-:
Each carbon has 4 and each oxygen has 6, so 32 total.
The overall charge is 2-, so there are 2 more electrons.
Grand total 34.
Start filling in… O O C C O O
Then put it all in brackets with a charge indicated.

72. BIG IMPORTANT NOTE DEUX
The method just described works well for challenging dot structures.
In case of emergency, use this technique, but don’t forget to count your electrons.

73. Practice
Covalent Bonding and VSEPR Theory Worksheet (Reverse, Middle Section)
Letters D, E, F, and G are challenging. Leave F for last!

74. Covalent Bonding
Remember that there are two types of covalent bonding:
Polar Covalent
Electrons are not shared equally.
Electrons are pulled closer to one atom than another because one atom is more electronegative.
Non-Polar Covalent
Electrons are shared equally.
Pulls from both atoms are roughly equal strength.

75. Water Has Polar Bonds
Let’s think of it this way. You tell me which of these two guys is gonna hold onto his electrons more… (draw them too)
OMG OXYGEN!
3.5 PAULINGS ELECTRONEGATIVITY!!!1
Meet Hydrogen – Poor guy’s only 2.1 Paulings in electronegativity.

76. - - - + + + + O O O H H H In a diagram… - - - - - - - -
There are 8 total valence electrons in the molecule.
Meet water: (say hi)
This is the source of cohesion/surface tension, and why it hurts to do a belly-flop.
…which causes neighboring water molecules to form hydrogen bonds with the first one.
…and all that’s left on the other end are the hydrogen protons, causing that end of the molecule to take on a positive charge…
Having all that negative charge built up around oxygen causes that end of the molecule to take on a negative charge…
But…since oxygen is such a strong atom, it doesn’t share electrons nicely. + H O -
Covalent Bond + H O - + + H O
Hydrogen Bond - - - - - - - - -

77. Polar Covalent Bonds
As a result of that big oxygen dude hoggin’ all the negative electrons (and being more electronegative), his end of the water molecule takes on a slightly negative charge.
We use the (lower case) delta symbol to indicate a negative charge: -
Because poor ol’ Hydrogen’s doesn’t see much of the electron action (and being less electronegative), his end gets a slightly positive charge (due to the hydrogen protons).
We use the (lower case) delta symbol to indicate a positive charge: +

78. Polar Covalent Bonds
Because water molecules have poles, the positive end of one water molecule (hydrogen) can stick to the negative end of another (oxygen).
This is what allows water to undergo adhesion and cohesion.

79. Non-Polar Covalent Bonds
Electrons shared (fought-over?) equally.
Example: H2 (or an H-H bond)

80. Bonds and Molecules Now we separate the molecules from the bonds.
As we have learned, individual bonds can be polar.
However, as in the lab, sometimes molecules made only of polar bonds can be non-polar overall.
Here’s what I mean…

81. This is a polar molecule, because the pulls do not cancel each other.
This is a non-polar molecule, because the pulls do cancel each other.
A figurative look at it?
Think of when a child pulls on his/her parent’s arm to lead him/her somewhere:
This is a polar bond, because “Parent” is likely stronger than “Kid.”
But if “Other Kid” is ripped, the pulls may not cancel out.
Add another kid with equal strength and you have another polar bond.
Kid
Parent
Other Kid

82. Polar Bonds, Non-Polar Molecules
Carbon tetrachloride (CCl4) has 4 polar bonds.
Electronegativity difference of 0.5 in each direction.
However, because all chlorine atoms are pulling equally on the electrons, the overall molecule remains symmetrical, and thus non-polar. Cl - Cl C Cl + - - Cl -

83. Identifying Polar Molecules
If the “pulls” in a molecule are symmetrical and cancel each other, it’s non-polar.
Even if the bonds are polar.
If they are asymmetrical, it is polar.
Example: PH3
Electronegativity difference: 2.1 (P) – 2.1 (H) = 0.
So all bonds are non-polar.
BUT! There is an unshared pair on P, making the molecule assymetrical (trigonal pyramidal geometry), leading to a polar molecule*. P H H H
*Up for debate.

84. H2O – Polar or Non-Polar? - + Here’s another model of water.
We know water’s polar, but it appears to be symmetrical, right?
However, when you look at the direction of the pulls, it is clearly not pulling equally in all directions.
So, to determine molecular polarity, draw the “pulls” from polar bonds and see if they cancel out. - +

85. Another Look [sketch me]
Let’s look at the “pulling” going on in a water molecule again:
2.1 + +
2.1 H O
…making O slightly negative…
O pulls downward on electrons more strongly than H pulls upward…
…and H slightly positive.
KEY: To show this, we add a “plus” to the arrow.
3.5 -

86. Polar or Non-Polar? Carbon dioxide (CO2) Ammonia (NH3) Chlorine (Cl2)
Non-polar; linear molecule
Ammonia (NH3)
Polar; trigonal pyramidal molecule
Chlorine (Cl2)

87. Practice Lab – Molecules I Front: #2, 4
Back: “Molecule Polarity” Column

88. Aside: Microwave Ovens
Microwaves work by heating water molecules in food.
Since the plate usually doesn’t have water, it doesn’t get warm.
How do they heat the water? With polarity!
The oven induces an electromagnetic field that causes the polar water molecules to orient in a certain direction.
Then the microwave changes the orientation of the field and the molecules need to change their orientation too.
The shifting back and forth and resulting friction creates heat and deliciousness.

89. Intermolecular Forces
Water in Space!

90. Intermolecular Forces
Time for another Australian video!
Find your video worksheets and turn to the reverse side for Intermolecular Forces.

91. Intermolecular Forces
Intermolecular forces are interactive forces between molecules. Here are three, in order of increasing strength:
WEAKEST: London Dispersion Forces
Occur: Always
Part of Van der Waals Forces
Dipole-Dipole Forces
Occur: With two polar molecules.
STRONGEST: Hydrogen Bonding Forces
Occur: When H is bonded to O, N, or F on one molecule and there’s an O, N, or F on another molecule.

92. London Dispersion Forces Demo
I need a nucleus or two!

93. London Dispersion Forces
The weakest type of intermolecular force.
It results from the motion of electrons around the nucleus.
When, momentarily, electrons in one atom happen to be on the same side of the atom, a transient dipole forms.
It’s like a temporary polar molecule.
As a result, the molecule has a polar end and non-polar end, and nearby molecules can also gain transient dipoles, causing them to attract each other.

94. London Dispersion Forces
Part of Van der Waals Forces:

95. Dipole-Dipole Forces
Also part of Van der Waals forces, dipole-dipole forces occur when two polar molecules come together.
The positive end of one polar molecule is attracted to the negative end of another.
Polar molecules are called dipoles, after all!

96. Van der Waals Forces Where do we see van der Waals forces in nature?
Geckos, for one!

97. Geckos

98. Hydrogen Bonding The strongest force between neutral molecules.
Also known as “when Hydrogen cheats on its molecule.”
When is there H-bonding? (NEED BOTH CONDITIONS)
On one molecule, H is bonded to O, N, or F.
Another molecule with O, N, or F is nearby.
Mostly because H only has two electrons.

99. Hydrogen Bonding Explained
When is there H-bonding (long answer)?
When H, small little atom that it is, bonds to a strongly electronegative atom like oxygen, nitrogen, or fluorine, none of the electrons hang out near it at all, making it positive.
Similarly, the O, N, or F on another atom usually has all the electrons hanging out near it, making it negative.
As a result, the H on one molecule and the O, N, or F on another are attracted to one another. Like a specific dipole-dipole interaction.
Or some kind of extra-molecular couple.

100. Intermolecular Forces Practice
Draw the pair and identify:
The shape of each.
Whether each is polar or non-polar.
The intermolecular forces between the two.
H2O and BF3
H2O Bent, BF3 Trigonal Planar
H2O Polar, BF3 Non-Polar
London Dispersion, H-Bonds
HCN and CH3F
HCN Linear, CH3F Tetrahedral
HCN Polar, CH3F Polar
London Dispersion, Dipole-Dipole

101. Important Note
Generally speaking, the more intermolecular forces between two molecules, the higher the boiling and melting points.
This is because the molecules hang onto each other more strongly.

102. Intermolecular Forces Practice
Intermolecular Forces Worksheet

103. Bond Dissociation Energies
For chemical reactions to occur, bonds must be broken.
Bond dissociation energy is the amount of energy needed to break a covalent bond.
Of one mole of the substance – more to come.
To find the dissociation energy of the whole molecule, add the bond dissociation energies of each bond in the molecule.

104. Bond Dissociation Energies (kJ/mol)
H—H
H—Br
C—Cl
F—F
O=O
436
366
330
158
498
H—C
H—I
C—S
Cl—Cl
C≡C
393
298
272
243
839
H—N
C—C
C—Br
Br—Br
N≡N
391
347
288
192
945
H—O
C—O
C—I
I—I
464
360
216
151
H—F
C—N
N—N
C=C
568
308
170
657
H—Cl
C—F
O—O
C=O
432
488
145
736

105. Bond Dissociation Examples
CH4 I2
393 * 4 = 1572 kJ/mol
302 kJ/mol
CO2 O2
736 * 2 = 1472 kJ/mol
498 kJ/mol
H2O
O—O
464 * 2 = 928 kJ/mol
145 kJ/mol
Br2
NH3
192 kJ/mol
391 * 3 = 1173 kJ/mol

106. Polystyrene Dissolution Demo
Styrofoam (non-polar)
Acetone (non-polar)
Permanent marker (to draw the wicked witch or something)

107. Bond Dissociation Energies
Styrofoam (called Polystyrene)
Chemical Formula: C8H8
Calculate the total bond dissociation energy for polystyrene.
7 C-H Bonds:
393 * 7 = 2751 kJ/mol
4 C-C Bonds:
347 * 4 = 1388 kJ/mol
4 C=C Bonds:
657 * 4 = 2628 kJ/mol
Total
6767 kJ/mol

108. O N O O Resonance Example Let’s draw the nitrate ion:
Add up total electrons available:
24.
Put single bonds everywhere.
Complete octets on bonded atoms.
Satisfy the octet rule for the central atom.
Did you put your double bond somewhere else? O N O O

109. Resonance in Nitrate – NO3-

110. Resonance
Resonance occurs when more than one valid Lewis structure can be drawn for a particular molecule.
The actual structure (or bond order) is an average of resonance structures.
Additional bonds are shown as dotted lines, as in this image. O N O O

111. Resonance in Benzene (C6H6)

112. Aside: Coordinate Covalent Bonds
On your Covalent Bonding Worksheets, #6 on the front (SO3) is an example of a substance with a coordinate covalent bond.
This is when there is a “donation” of an electron pair from one atom to another.
Both electrons are donated from one atom.
Here’s how to draw it: O
 Double Bond S
Coordinate  Bond
 Coordinate Bond O O

113. Aside: Coordinate Covalent Bonds
Note that coordinate covalent bonds (also known as dipolar bonds or a dative covalent bond) frequently occur in resonance structures.
A good sign of this is when an atom has all its valence electrons shown as lone pairs, yet is still forming a covalent bond.
Usually oxygen, it seems like, but could be others.
Let’s redraw nitrate’s resonance structures to show coordinate covalent bonds.

114. Resonance in Nitrate – NO3-

115. What’s the Big Idea?
Coordinate covalent bonds do not function any differently from the kind of covalent bonding you already know how to draw.
Drawing a coordinate bond simply allows you to specify the origin of the electrons that are being shared.
You do not need to show coordinate bonds for this class, but it’s good to know about them for those of you with AP Chem in your futures.

116. As a single resonance structure: With coordinate bonds:
And the answer to #6?
As a single resonance structure:
With coordinate bonds: S O S O

117. Closure Wow. We’ve done a lot of stuff throughout this lesson.
Questions?

118. Okay, one last thing…
Should you decide to take AP Chemistry, you’ll also cover a little bit more of VSEPR Theory called hybridization.
Hybridization occurs in covalent bonding, as electron repulsion changes the shape of orbitals (generally just the s and p ones, but sometimes d).
When hybridization occurs, the orbitals re-shape to be more like a strange combination of both of them.
When those orbitals form bonds with each other, the bond type is renamed as well.
Here’s a short summary…

119. Orbital Hybridization
Groups Attached to Central Atom
Geometry
Hybridization 2
Linear sp 3
Trigonal Planar
sp2 4
Tetrahedral
sp3 5
Trigonal Bipyramidal
sp3d 6
Octahedral
sp3d2

120. Hybrid Bonds Bond Type Hybrid Bond Types Single 1 Sigma Bond* Double
1 Pi Bond**
Triple
2 Pi Bonds**
*Overlapping of two sp hybrid orbitals.
**Overlapping of two non-hybridized p orbitals.