Hello and welcome to a video about hybrid orbitals, often called valence bond theory. Developed in the 1930s by the great chemist Linus Pauling as a model of bonding to understand the three-dimensional placement of atoms in a molecule, and that is critical to our understanding of the properties that molecules have. In this video we will look at methane, ethene, ethyne, ammonia, and water as our models for hybridization. There is only a small group of atoms in the second period that the model really works for, but among those are carbon nitrogen and oxygen, which make up the vast majority of molecules that exist on earth. So the model applies to a limited number of elements but it applies by far to the majority of molecules. Let's take a look at the word hybrid: it is a blending of two varieties. If you get a horse and donkey together you get a mule, a blending or hybrid of a horse and a donkey. Let's get to hybrid orbitals using carbon as our model. As a single atom not bonded to anything, carbon has two electrons in 1s, two electrons and 2s, and two electrons in 2p. This electron configuration is the energy arrangement of carbons electrons. However carbon rarely exists in nature as an individual atom except momentarily while undergoing chemical reactions. Carbon exists with its valence electrons bonded to other atoms. When carbon bonds to four other atoms, carbon's four bonds are experimentally seen to be equivalent. And so when the carbon atoms finds itself in a bonding situation, its bonding electrons themselves exist at equivalent energies, which requires that they hybridize to an energy that is intermediate between the 2s energy and the 2p energy. Or you can think of it as a blending or hybridization of the two energies, the s and p energies. And since the energies of these electrons have now changed, the shape of the orbitals they occupy are different as well, which we will see momentarily, and those are called hybrid orbitals. They are named 2sp3 hybrid orbitals. The naming often confuses students so before we go any further, let's take a look at where the 2sp3 name comes from: The 2 comes from the second principal energy level that the valence orbitals are in. The s comes from the 2s orbital contributing to the hybridization, and the p comes from the 2p orbitals contributing to the hybridization, and the 3 comes from the number of 2p orbitals used in hybridization. Once hybridized, the 2s and 2p orbitals no longer exist, and so we have four 2sp3 hybrid orbitals. Four 2sp3 hybrid orbitals derived from combining the energies of one 2s orbital and three 2p orbitals, which gives a total of four 2sp3 orbitals. Before we look at the shape of hybrid orbitals it would be helpful to briefly review the atomic orbitals. The 1s orbital is a sphere, the 2s orbital is a larger sphere surrounding 1s, and here we will get rid of 1s since we are only concerned with the valence electrons. Each 2s orbital is a two lobed shape converging at the nucleus. So there are the three 2p orbitals. However when hybridization occurs the s and p orbitals cease to exist, and the 2sp3 orbitals have an entirely different shape. We can see that orbital hybridization explains the VSEPR placement of carbon's four valence electrons since all four 2sp3 orbitals are equivalent, each 2sp3 orbital repels the others with equal force, resulting in identical bond angles. The carbon atom only hybridizes when it is in a bonding situation. Here, four hydrogen atoms bond to carbon by overlapping their orbitals with carbon's hybrid orbitals. So what would be the reason for this? If we go back and see that both carbon and hydrogen have unpaired electrons, the overlap allows the electrons to pair and thus go to a lower potential energy. The illustration here contains the valence electrons of both carbon and hydrogen, and since everyone likes to visualize atoms as spheres we can do the same: here is our carbon atom, and here are the hydrogens, with the overlapping spheres, indicating the overlapping orbitals that constitute the bond. The bonds are more readily discernible in a ball and stick model, which also makes the bond angle more visible. Since all four sp3 orbitals are equivalent, each bonding orbital repels the others with equal force resulting in identical bond angles. The bonds in hybridization also have their own nomenclature. The overlapping orbitals are called sigma bonds which represents the single bond occupied by a single pair of electrons. What about double bonds? How does the hybridization model explain double bonds? We will use ethene, C2H4, to see what happens in a double bond. The single bonds we know are sigma bonds, and the double bond also has a sigma bond, but the second bond of a double bond is a pi bond. Let's see how hybridization and orbital overlap can explain a double bond. The two carbon atoms in ethene are equivalent, so let's look at one of the carbon atoms first. A pi bond comes from the overlap of unhybridized p orbitals, and so the atom hybridizes only three orbitals, leaving a p orbital unhybridized for the pi bond. The hybrid orbital is called 2sp2 the superscript 2 denoting that only two 2p orbitals have contributed to the hybridization. The 2sp2 hybrid orbitals exist in a plane perpendicular to the unhybridized 2p orbital. Let's remove the 2p orbital for now to more easily see that. The 2sp2 hybrid orbitals are spread out at a 120 degree angle, which means that they exist in a plane, and the plane is perpendicular to the unhybridized 2p orbital. So this is what both carbon atoms do when bonding occurs in ethene. Each carbon atom is sp2 hybridized. The sigma bond occurs with 2sp2 orbital overlap. What about the pi bond? The second bond of the double bond. Previously we said that it comes from the unhybridized p orbitals, which we see here from both carbon atoms. The top and bottom lobes of the 2p orbitals overlap above and below the axis of the sigma bond forming a single pi bond. The space in which the now paired electrons move around. The ball and stick model shows this double bond with two dashes. To summarize, sigma bonds occur along the axis between nuclei. The pi bond occurs above and below the Sigma axis where the p orbital lobes have overlapped. The ethene molecule also bonds to four hydrogen atoms by overlapping with both carbon's other 2sp2 hybrid orbitals, creating four more sigma bonds. In the ball-and-stick model we can readily see that each carbon has a trigonal planar geometry, and thus the whole molecule exists in a plane with the single pi bond above and below that plane. Now let's look at how hybridization can be a model for the triple bond using ethyne, C2H2. The carbon-hydrogen bonds are sigma bonds, and the triple bond is one sigma bond and two pi bonds. Let's see how hybridization can accommodate this. Since pi bonds come from p orbitals, and we need two pi bonds, then two 2p orbitals have to remain unhybridized, and so the remaining single 2s orbital and a single 2p orbital will hybridize to two 2sp hybrid orbitals. And there is also the violet 2px and the blue 2py orbitals. Here each green lobe is a single orbital, and so they each have an electron, while both violet lobes constitute the single 2px orbital with a single electron. And both blue lobes constitute the single 2py orbital with a single electron. The other carbon in C2H2 also has that same triple bond, and so it has the same hybridization. Let's see what happens during bonding. sp orbitals from both carbons overlap, forming a sigma bond. The upper lobes of the blue 2p orbitals overlap, as do the lower two lobes, creating the first of the 2 pi bonds. Can you guess where the second of the two pi bonds comes from? Yes that's right! It is the overlapping of the violet 2p lobes. Let's get rid of the sigma bond for a moment to take a look at something interesting. Each pi bond lies on a separate plane and the two planes are perpendicular to each other, and so the two pi bonds are perpendicular to each other. Finally, two hydrogens will overlap with remaining sp hybrid orbitals, creating the C2H2 molecule. The overlap of the space-filling model reflects the overlapping orbitals, which is also represented by the ball-and-stick model. In the remainder of the video we will look at hybridization of nitrogen and oxygen using NH3, ammonia, as our model for nitrogen hybridization, and H2O, water, for our oxygen model. In NH3, nitrogen has three sigma bonds and a lone pair, so how does hybridization account for this? The hybridization is 2sp3, and nitrogen has 5 valence electrons so one of the four 2sp3 hybrid orbitals has a pair of electrons. The three sigma bonds come from the sp3 orbitals with a single electron, so they can pair up, and so the remaining electron pair is a lone pair, an unbonded pair of electrons. As with sp3 hybridization in carbon, nitrogen hybrid orbitals spread out in a tetrahedral shape. And lastly water. Here oxygen has two sigma bonds and two lone pairs. In water oxygen is also 2sp3 hybridized, but with six valence electrons: Two of the sp3 orbitals have paired electrons. You can probably guess that the sp3 orbitals with a single electron will overlap with hydrogen, and the remaining two pairs are unbonded, they are lone pairs. Again oxygen's hybrid orbitals spread out in a tetrahedral shape. That's it for hybridization, the product of a mad scientist! SEEYA!