Molecular Geometry Made Easy: VSEPR Theory and How to Determine the Shape of a Molecule — Transcript

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1. 0:00Molecular Geometry Made Easy brought to you by Ketzbook
2. 0:03Welcome back and thanks to all my subscribers and patrons out there for making this video possible.
3. 0:08This video is sponsored by Genius Lab Gear, the creator of “The Pocket Chemist.”
4. 0:12In this video we are going to learn how to determine the shape of a molecule, what the
5. 0:15bond angles are in different molecules, and the VSEPR theory to understand why molecules
6. 0:20have different shapes and different bond angles.
7. 0:23However, we will not look at molecules with an “expanded octet” or coordination complexes.
8. 0:27I will cover those in subsequent videos.
9. 0:30But first, I would like to ask a fascinating question.
10. 0:33Why are you able to hold something in your hand?
11. 0:35Why doesn’t it just fall through your hand? Remember that atoms are mostly
12. 0:39empty space, and by “mostly” I mean that an atom is more than 99.99999999999% empty space.
13. 0:48It is highly unlikely that the protons, neutrons, or electrons of the atoms in different objects will actually touch.
14. 0:55So if they aren’t touching, what gives an object its sense of boundaries?
15. 0:59Electron-Electron Repulsion.
16. 1:01When an atom approaches another atom, the negatively-charged electrons of the two atoms repel each other.
17. 1:07So, the size and boundaries of an atom are determined by electrons repelling each other.
18. 1:11That same principle also determines the shape of a molecule. Let’s see how this works
19. 1:15in the example of sulfur dioxide. The first step in determining the shape of a molecule
20. 1:20is to draw the Lewis diagram, and if you don’t know the Lewis diagram already, you’ll have
21. 1:24to count the valence electrons. The one sulfur has 6 valence electrons, and the two oxygens
22. 1:28have 6 times 2 or 12 valence electrons, which gives sulfur dioxide a total of 18 valence electrons.
23. 1:35Because there is only one sulfur, we put sulfur in the middle, surrounded by two oxygens with
24. 1:39single bonds to the sulfur. Each bond has two electrons, so we have used up four electrons
25. 1:44so far, which means we have 14 electrons remaining for lone pairs. 2, 4, 6, 8, 10, 12, and 14.
26. 1:53Notice that the sulfur only has 6 valence electrons around it, so we should use two
27. 1:57of oxygen’s electrons to form a bond between it and sulfur. The resulting Lewis diagram
28. 2:03is complete with each atom having 8 valence electrons.
29. 2:07After you draw the Lewis diagram, you need to count the number of “things” on the central atom.
30. 2:11By “things” I mean bonded atoms or lone pairs.
31. 2:15Technically speaking, these are called “electron domains.”
32. 2:18The double bonded oxygen is a thing, as is the single-bonded oxygen and, the lone pair on the sulfur.
33. 2:24Remember that only lone pairs on the central atom count.
34. 2:27So, there are 3 electron domains, or things on sulfur dioxide that determine its shape.
35. 2:32Those 3 things will repel each other and get as far apart as possible.
36. 2:37That repulsion with the lone pair of electrons will cause the two oxygens to bend down,
37. 2:42so the resulting shape of SO2 is called BENT.
38. 2:46Since there are three things attached to the central atom,
39. 2:48we would expect the bond angle to be 360 divided by 3 or 120°.
40. 2:53In fact, the real bond angle is slightly less than that;
41. 2:57it is 119° because lone pairs occupy more space close to the central atom,
42. 3:03while electrons in bonds are stretched out between the two bonding atoms.
43. 3:07Okay, now let’s try to apply these principles to determine the shape of several other molecules.
44. 3:12Let’s start with carbon dioxide. You may wish to pause the video now and try to draw
45. 3:16the Lewis diagram and determine the shape of carbon dioxide on your own.
46. 3:21For the Lewis diagram, carbon is in the middle,
47. 3:24and it forms double bonds to each oxygen, which both have two lone pairs.
48. 3:28Next, we need to count the number of things on the central atom.
49. 3:31In this case, there are only two things on carbon.
50. 3:34Although there are lone pairs in the molecule,
51. 3:37the lone pairs on oxygen do not affect the shape because they are not on the central atom.
52. 3:42Since there are only two things, you can see in our drawing that those two things are already
53. 3:46as far apart as possible. Because the three atoms are all in a line, we say that the shape
54. 3:51of CO2 is linear, and the bond angle is a linear 180°.
55. 3:56As we take a step back and look at SO2 and CO2, we can see that for a 3 atom molecule
56. 4:02the only two possible shapes are linear or bent, and the deciding factor is whether or
57. 4:07not there is a lone pair on the central atom. Bent molecules have one or more lone pairs
58. 4:13on the central atom, while linear molecules have no lone pairs on the central atom.
59. 4:19Before we move on to more complex molecules, let’s consider two atom molecules, such as N2.
60. 4:24Two atom molecules do not have a central atom, which means they cannot have any bond angles.
61. 4:30So, the only possible shape for two atoms is linear.
62. 4:33N2 is linear, and all two atom molecules are linear.
63. 4:37Next, let’s look at some 4 atom molecules, like boron trihydride.
64. 4:41When drawing the Lewis diagram of BH3, boron is in the middle,
65. 4:45and it makes 3 single bonds to the 3 hydrogen atoms.
66. 4:49Remember that because boron starts with only 3 valence electrons,
67. 4:52and each hydrogen starts with 1 valence electron,
68. 4:55there are no more electrons remaining for lone pairs, leaving boron without an octet.
69. 5:00Boron has three things around it, and each of those things will repel each other to get
70. 5:04as far away from each other as possible. That repulsion will lead to a symmetrical structure
71. 5:09with all the bond angles being equal. That is, all the bond angles are 120°, and the
72. 5:15molecule is perfectly flat without any puckering. Because the molecule looks like a flat triangle,
73. 5:20the name of its shape is trigonal planar.
74. 5:23Let’s try another 4 atom molecule, CH2O, that is, formaldehyde. Because carbon prefers
75. 5:29to make 4 bonds, we put carbon in the center. Each hydrogen only forms one bond, and oxygen
76. 5:34prefers to form 2 bonds. That leaves 4 electrons for lone pairs, which both go on the oxygen
77. 5:39to give every atom an octet or duet.
78. 5:42There are 3 things on the central atom, and those three things repel each other.
79. 5:46Once again, we can see that the molecule looks like a flat triangle, so the shape is trigonal planar.
80. 5:51However, the bond angles are not exactly 120°. The H–C=O bond angles are both 122°;
81. 5:59they are a little bit bigger than 120 because double bonds occupy more space than single bonds.
82. 6:04The H–C–H bond is consequently only 116°. Remember that a double bond has
83. 6:104 electrons, and a single bond has two electrons, so the double bond would repel more and need
84. 6:15more space than a single bond. But, the shape is still the same.
85. 6:19Next, let’s check out a 5 atom molecule, like CH4, that is, methane. Carbon is in the
86. 6:25middle with 4 single bonds to hydrogens, and there are no lone pairs. Now there are 4 things
87. 6:30on the central atom, and each of those things will repel each other to be as far apart as possible.
88. 6:35This time it appears as though they already are as far apart as possible, and
89. 6:39it appears as though the bond angles should all be 90°.
90. 6:43However, looks can be deceiving when you are trying to visualize a 3-dimensional molecule
91. 6:48on a 2-dimensional surface. The hydrogens can actually move away from you or towards
92. 6:52you in 3-dimensional space. The resulting shape is not flat but is a 3-dimensional shape
93. 6:57called tetrahedral, and the bond angles are consequently larger than 90°.
94. 7:02The bond angles are all 109.5°
95. 7:05It can be difficult to visualize this shape, so I recommend that you build methane and
96. 7:09similar molecules yourself with a physical molecular model or an application like MolView.
97. 7:14Here is a ball and stick 3d model of methane.
98. 7:16Notice the bond angles and the symmetry of the molecule.
99. 7:19To give it a bit of a 3-dimensional perspective, methane is sometimes drawn like this, or like this.
100. 7:25The dashed bond represents a bond that is going away from you,
101. 7:29while a solid wedge bond represents a bond that is going toward you.
102. 7:33Next, let’s try NH3, which is called ammonia. Unlike BH3, which we said was trigonal planar,
103. 7:39NH3 has a lone pair of electrons on the central atom. In some ways, ammonia is similar to
104. 7:44methane because it also has 4 things attached to a central atom. In general, molecules that
105. 7:49have the same number of things on the central atom are said to have the same electron-domain
106. 7:53geometry and tend to have similar bond angles.
107. 7:56However, the actual shape of these molecules will differ because lone pairs of electrons
108. 8:00occupy space but are not observed as part of the shape.
109. 8:04As for ammonia, those four things attached to nitrogen will all repel each other, and
110. 8:08because there are four things attached to a central atom—just like in methane—we
111. 8:12expect that the bond angles will be about 109.5°. We can represent this 3-dimensional
112. 8:17shape the same way we did methane, only this time we put the lone pair on top.
113. 8:22Here is a ball and stick model of ammonia. Its shape is called trigonal pyramidal because it essentially
114. 8:27looks like a pyramid with a triangular base. You can think of trigonal pyramidal as tetrahedral
115. 8:33with one atom removed, or you could think of trigonal pyramidal as trigonal planar with
116. 8:37the central atom pulled up some.
117. 8:40For ammonia, the actual bond angle is very close to 109.5°, but it is a little bit smaller
118. 8:45because lone pairs occupy more space close to the central atom causing the hydrogen atoms
119. 8:49to squeeze closer together. The actual bond angle of ammonia is 106.7°
120. 8:55Okay, we have already looked at the five basic shapes of molecules along with an example
121. 9:00for each one and the guiding principle of electron-electron repulsion.
122. 9:04Let’s take some time to summarize all of this before we go through one last example molecule.
123. 9:09Previously we said that 2 atom molecules like N2 are always linear.
124. 9:13If a molecule has 3 atoms, there are two possible shapes.
125. 9:16If there are no lone pairs of electrons on the central atom, the shape is linear, like CO2;
126. 9:21but if there are lone pair(s), the shape is bent, like SO2.
127. 9:25For 4 atom molecules, the shape also depends on whether or not there is a lone pair of
128. 9:29electrons on the central atom. If there is no lone pair on the central atom, then the
129. 9:33molecule is trigonal planar, like BH3 and formaldehyde. If there is a lone pair on the
130. 9:39central atom, then the molecule is trigonal pyramidal, like ammonia.
131. 9:42Lastly, 5 atom molecules are typically tetrahedral, like methane. Although there are some exceptions,
132. 9:48those exceptions are non-octet molecules, which I will discuss in a different video.
133. 9:53Another way to summarize these shapes is by counting the number of things attached to the central atom.
134. 9:58For example, in a 3-atom molecule, there are 2 atoms attached to the central atom.
135. 10:02If there are zero lone pairs of electrons, the shape is linear, and the bond angle is 180°
136. 10:08If there is one lone pair, the shape is bent, and the bond angle is approximately 120°.
137. 10:13If there are two lone pairs, the shape is also bent, but the bond angle is closer to 109.5°.
138. 10:20In a 4-atom molecule, there are 3 atoms attached to the central atom. If there are zero lone
139. 10:24pairs on the central atom, the shape is trigonal planar, and the bond angle is about 120°.
140. 10:29If there is one lone pair, then the shape is trigonal pyramidal, and the bond angle is about 109.5°.
141. 10:35Finally, if there are 4 atoms, the shape is tetrahedral, and the bond angle is also 109.5°.
142. 10:41Notice that if the total number of things attached to the central atom is the same,
143. 10:45then the bond angle will also be the same. A total of 2 things leads to a 180° bond angle,
144. 10:51a total of 3 things leads to 120° bond angle,
145. 10:54and a total of 4 things on the central atom leads to 109.5° bond angle.
146. 10:59Now, it’s time for you to try a sample problem. What is the shape of H2O, and what is the
147. 11:04approximate bond angle? Pause the video, and try to figure it out yourself.
148. 11:09The first thing that you should have done is draw the Lewis diagram. Oxygen is in the middle,
149. 11:13and there are two single bonds to the hydrogens. This leaves 4 valence electrons,
150. 11:18which we put as lone pairs on oxygen, giving it an octet. This structure has 2 bonded atoms
151. 11:24and 2 lone pairs attached to the central atom, for a total of four things that all repel each other.
152. 11:29Just like in methane, these things will move as far apart as possible in 3-dimensional space.
153. 11:34The hydrogens can move away from you and the lone pairs can move toward you, and
154. 11:38the bond angle will be greater than 90°. In fact, because there are four things on
155. 11:43oxygen, the bond angle should be about 109.5°.
156. 11:47Any 3-atom molecule is either linear or bent, and water is bent because of the two lone pairs.
157. 11:52We can redraw water in a way that emphasizes its bent shape. As for the bond angle,
158. 11:58we know that it should be less than 109.5 because lone pairs take up more space close to the
159. 12:03central atom than single bonds. Because there are two lone pairs, we expect the bond angle
160. 12:07to be even smaller than in ammonia, which was 106.7. These predictions are correct,
161. 12:12and the actual bond angle in water is 104.5°.
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174. 13:08Thanks for watching! If you found this video helpful, please like, subscribe, or check
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176. 13:15below, and have a wonderful day!