1. 4.2 The Nature of the Chemical Bond
Unit 1: Structure and Properties

2. Valence Bond Theory Developed by Linus Pauling
A covalent bond is the overlap of two atomic orbitals to form a new combined orbital containing two electrons of opposite spin.
Example: 1s 1s

3. Valence Bond Theory
Electron density becomes concentrated between nuclei (shared pair of electrons)
Example: hydrogen fluoride

4. Problems with Valence Bond Theory
NOT capable of explaining BOND ANGLES
Example: H2O
*predicted bond angle using valence bond theory is 90o
*measured angle is about 105o

5. The real magic of electrons are their ability to change shapes depending on the needs for bonding. For example, on the left is nitrogen's 3 lobe-shaped p orbital electrons that sit along the x, y, and z axis. The s orbital is the large orange translucent sphere. In this configuration, p orbital electrons are at 90° angles. Depending on how many other atoms it needs to connect to, the electrons can rearrange themselves into different shapes to accommodate the number of bonds it needs to make. Those can have different angles. Roll cursor over image to see these 4 electrons change their shapes. Note: The p orbitals are drawn thinner than normal so you can see their orientations better.

6. Bonding Carbon and Hydrogen

7. Up until now we have talked about the s, p, d, and f orbitals
Up until now we have talked about the s, p, d, and f orbitals. That alone gives atoms all kinds of flexibility in connecting with other atoms. This tutorial is about how these orbitals can blend with each other to form new shapes, which gives even more flexibility in bonding. This blending is called hybridization. It's a good word because the final orbital is a hybrid of the original orbitals.
In my example on the left, I show an s orbital and a p orbital at the top. When they blend they make an sp orbital. You can see some of the traits of both s and p orbitals in the sp orbital.

8. Since organic chemistry and biochemistry rely on the carbon atom, that's the element that we put our attention on first. The diversity of carbon to make complex molecules is only possible because of the hybridization that its electrons undergo. On the left are 3 carbon atoms with their electrons in their ground state (lowest energy level). Carbon's two 2p orbital electrons (2px and 2py) are the more available for bonding than the s orbital electrons. However, those p orbital bonds would be at right angles to each other, which would make the atoms crowded. It also would only allow carbon to connect to 2 other atoms. So carbon's electrons do something that allows for more spacing and for connection to more atoms.

9. Here is the ground state of carbon
Here is the ground state of carbon. The 1s electrons are small and not shown. They aren't involved in bonding. The two paired 2s electrons are spherical and are symbolized with two arrows. [↑↓]. There are 2 unpaired p orbital electrons that align on the x and y axes.
As the carbon atom nears another atom, such as hydrogen, one of the electrons in the 2s orbital gets pulled into a higher energy orbital. In this case, it goes into the unfilled 2pz orbital (see next panel).

10. Again, one of the electrons in the 2s orbital gets pulled into the unfilled 2pz orbital. Notice the green 2pz orbital is now present. The orange dotted arrow shows it came from the 2s orbital. The 2s orbital still has 1 electron left. So the orange sphere is still there.
What happens next is that these four electrons blend with each other (hybridize), and they all become the same shape. The shape is two lobes with one much larger than the other. In the right image I'm showing the four larger lobes and not the small ones.
The shape is named after the number of the orbitals involved. The name "sp3" means there was one s orbital and three p orbitals that got hybridized. Warning: In the electron configuration symbols of "2s22p3", the exponents are the number of electrons. In "sp3" the exponent is number of orbitals. There can be 4 to 8 electrons in the "sp3" orbitals. Notice in the image the 4 electrons in the sp3 orbital area unpaired. If all were paired, you would have 8 electrons (the octet).

11. The sp3 orbital forms a tetrahedron (a pyramid with a triangular base and all equal sides) The angle formed between the bonds is 109.5°. When carbon combines with 4 other atoms, this is the shape of the molecule.

12. Hybridization Theory
proposes that atomic orbitals combine to form new hybrid orbitals
hybrid orbitals then overlap with orbitals of other atoms to form covalent bonds
Example: combining a “p” orbital and an “s” orbital
s orbital + p orbital = two sp hybridized orbitals
Note: these orbitals are hybridized only when bonding occurs to form a molecule. They do not exist in an isolated atom.

13. SP2 Connections

14. When carbon is surrounded by 2 hydrogen atoms and an oxygen atom, carbon's electrons will do something different. Remember oxygen needs 2 electrons to make an octet, so carbon needs to share two electrons with oxygen by forming a double-bond with oxygen. At the top I show the electron configuration of carbon after one of the 2s orbital electrons have been promoted to the 2pz orbital.
Three electrons are going to get hybridized. From the name of "sp2" you can deduce that it will be one s orbital electron and two p orbital electrons.
This image shows the result of 3 of carbon's electrons hybridizing into sp2 electrons. Notice these three sp2 electrons spread out in a triangle with 120° between them. In the earlier tutorial on VSEPR (Valence Shell Election Pair Repulsion) we saw that when a central atom bonds with 3 other atoms and there's no lone pair on the central atom, the shape of the molecule is called planar triangular. This hybridizing is how that happens.
The one 2py electron (yellow lobes) did not get hybridized which is how carbon forms a double bond with oxygen shown next.

15. The image on the left is of carbon combining with 2 hydrogen atoms and one oxygen atom to make formaldehyde. The top diagram is the Lewis Dot Structure. The bottom illustration shows the orbitals as they prepare for bonding. Notice that the oxygens electrons underwent the same hybridization as did carbon's electrons. The presence of carbon triggered oxygen to also do hybridization. The difference is that oxygen has 2 more electrons than carbon, which explains why two of the sp2 orbitals are full. That's indicated by the 2 black dots, which represent the 2 paired electrons. Since these are not used for bonding, these are called lone pairs.
Now we will see how the bonded molecule looks...

16. Now we will see how the bonded molecule looks...
What happens is the two sp2 electrons that were pointing towards each other overlap and form a bond. The two vertical 2py electrons from both carbon and oxygen bend toward each other and overlap. That also forms a bond.
By the way, bonds that overlap directly between 2 atoms are called sigma bonds. The overlapped region is spherically shaped, much like the s orbital. The Greek letter for the "s" sound is σ (sigma). The sigma bonds are the strongest of the bonds.
The bond formed by the overlapping p orbital electrons is called a pi (π) bond. Note the overlapping at the top and bottom of the p orbital electrons, but it still counts as just one pi bond. In the Lewis Structure at the top, the double bond symbol (=) shows 2 lines. One is the sigma bond and the other line represents the pi bond.

17. This is the same molecule as above but seen from the top
This is the same molecule as above but seen from the top. I left off the pi bond so you can see the sp2 bonds better. These bonds sit flat (planar) and form 120° angles.

18. To show its flexibility even more, carbon can connect to an element that needs 3 electrons to reach its octet. The molecule of hydrogen cyanide has carbon in the center of hydrogen and nitrogen. A triple bond is needed with nitrogen.
Like before, carbon promotes one of its 2s electrons to the 2pz orbital. Then the other 2s electron plus one of the p orbital electrons gets hybridized into an sp orbital. The two sp orbital electrons point in the opposite direction. One will bond with hydrogen and the other with nitrogen. The two 2px and 2py electrons will form the other two bonds. (Shown in next panel.)

19. Here the hydrogen, carbon, and nitrogen atoms are set to bond
Here the hydrogen, carbon, and nitrogen atoms are set to bond. The pink sp electrons point directly at the atom they will bond with. Notice that the electrons in nitrogen did the same hybridization as carbon. The only difference is that nitrogen has 1 more electron than carbon which explains why one of nitrogen's sp electrons is full (2 dots) and forms a lone pair (non-bonding pair).
The yellow lobes are the unchanged p orbital electrons as are the green lobes. They are going to bend towards each other and overlap.
Here the sigma bond between the sp electrons have overlapped and those bonds are in place. The 2 pairs of p electrons are moving towards each other. In the next panel I show overlap of the p electrons
The yellow lobes form a single pi bond above and below the sigma bond. The green lobes form a single pi bond to the front and back of the sigma bond. So in the Lewis structure that shows the triple bond, one line represent the sigma bond and the other 2 lines represent the 2 pi bonds. The lone pair of electrons on the nitrogen atom are the 2 dots in the non-bonding sp orbital. The hydrogen atom also forms a sigma bond with carbon's second sp orbital electron.
The shape of the molecule is linear because the sp electrons are on opposite sides of carbon and nitrogen and there are no lone pairs on carbon to repel these bonds.

20. What about Phosphorus Pentachloride?

21. Phosphorus is in the 3rd row on the Periodic Table
Phosphorus is in the 3rd row on the Periodic Table. So its principle quantum number is n=3. At n=3, the l values available (if you remember) are 0, 1, and 2. That means s, p, and d orbitals are available. Carbon is in row 2, so it doesn't have any d orbitals available. Because d orbitals are available for phosphorus, it can promote one of its 3s electrons to a d orbital.
What you see here is that 1 of the 3s electrons got promoted to one of the d orbitals. That causes the appearance of four turquoise egg shape lobes. Those 4 lobes are just one d orbital electron that came from the 3s orbital. I picked the dxy orbital because it's easier to draw.
The beauty of this move is that phosphorus now has five electrons that are in 3 different orbitals. That allows it to hybridize all of them into 5 identically shaped orbitals. These are called sp3d orbitals because the one s orbital, three p orbitals, and one d orbital were hybridized
After the five electrons are hybridized, they are five electrons of equal size and shape. They are five sp3d orbital electrons. Three form a triangle (120° separation) in the middle. One bond is above and one is below those 3. Here is the geometry of how it will bond with the 5 chlorine atoms.

22. End Result
See link on website

23. Predicting Hybridization
Draw the Lewis Diagram of the molecule/ion
Count the number of bonds on the central atom (NOTE: Double and triple bonds count as ONE)
Add any non-bonding electron pairs
The total number that you get corresponds to the number of orbitals 2 3 4 5 6 sp
sp2
sp3
sp3d
sp3d2

24. Homework Read p 219-225 Answer p 223 #20-22; p 227 #9, 13
Lewis Structures & Hybridization Worksheet

25. Draw Lewis structures for each of the following molecules, draw the energy level diagrams for both ground and promoted states, and state it’s hybridization
(a) TeF4 (b) ClF5 (c) XeF2
(d) XeF4 (e) ClO4- (f) AsH3
(g) AsCl4+ (h) PCl6- (i) FCl2+
(j) AsF5 (k) AsF3 (l) SeO2
(m) IO4- (n) ICl4- (o) TeF6
(p) ICl2- (q) SbH3 (r) PCl4+
(s) ClO3- (t) SiO44- (u) SO3
(v) OF2 (w) SbCl6- (x) BrCl3
(y) XeF4