Valence Bond Theory, Hybrid Orbitals, and Molecular Orbital Theory - Hey, it's Professor Dave. Let's talk about molecular orbitals. [MUSIC PLAYING] BAND: (SINGING) He knows a lot about the science stuff. "Professor Dave Explains." - So we know about atomic orbitals. And we learned that an atom can take its atomic orbitals and hybridize them into molecular orbitals. These hybridized orbitals have shapes that resemble combinations of the atomic orbitals. And these are the orbitals that are used to make covalent bonds, which are simply the overlapping of such orbitals. Let's start with something simple, a molecule of H2. Each hydrogen has one electron that sits in a 1s orbital. So to form a sigma bond, the atomic orbitals overlap and form one molecular orbital, which both electrons will reside in. The hydrogen molecule will sit at a lower energy than the two lone hydrogen atoms because of the additional electrostatic interaction between each electron and the opposite nucleus. This is why hydrogen spontaneously forms diatomic molecules. But the number of orbitals must be conserved. So when two atomic orbitals come together to form molecular orbitals, they will form one bonding orbital and one antibonding orbital. The electrons in the covalent bond occupy the lower energy bonding orbital since any orbital atomic or molecular can hold two electrons. But if one electron were to become excited, it could jump from the sigma or bonding orbital to the sigma star or antibonding orbital. This is where hybridization comes into play. If an atom is participating in multiple bonds, it will utilize hybrid orbitals to form them. For example, if carbon is making four bonds as it tends to do, it will take its 2s orbital and all of its 2p orbitals to create four sp3 hybridized molecular orbitals. In order for carbon to hybridize its orbitals, it must first promote one of the two s electrons up to the vacant 2p orbital. Then it can hybridize them, leaving four molecular orbitals that are identical in energy or degenerate. Each of them has one electron, leaving room for another that will be provided by the other atom participating in each covalent bond. Every energy level has one s orbital, three p's, and five d's. And the number of bonds an atom can participate in determines the number of atomic orbitals that will need to be hybridized to form the necessary molecular orbitals to generate the bonds. An sp3 hybridized carbon is typically participating in four sigma bonds. But carbon can also be sp2 or sp hybridized. Let's look at an sp2 carbon. If it is sp2 hybridized, it has just three electron domains, which can happen if it has a double bond to something. Take C2H4 for example. Each of these carbons is sp2 hybridized because of the double bond present. What this means for molecular orbitals is that the carbon will promote one electron and then take its 2s orbital and just two of the 2p orbitals to generate the sp2 hybridized orbitals. That leaves one p orbital unhybridized. It is the overlap of these unhybridized p orbitals each containing one electron that generates a pi bond, which is what the second bond in this double bond is. The p orbitals extend in perpendicular fashion from the plane of the molecule. And they overlap to make the pi bond. The same goes for C2H2, which has a triple bond between the carbons. Each carbon has two electron domains, which means it will be sp hybridized. After promoting one electron, it needs to only utilize the 2s orbital and one 2p orbital to get the sp hybridized orbitals. That leaves two unhybridized p orbitals each with one electron. In this case, they will generate two pi bonds-- one from the overlap of these p orbitals and another one from the other set of p orbitals that will extend perpendicularly from the first pair. What you end up with is a linear molecule with this propellor-like orientation of the unhybridized p orbitals, which form two pi bonds beyond the initial sigma bond, which is the overlap of the sp orbitals triple bond. Sometimes, we have to generate orbital diagrams for a molecule. To do this, we show each atom on the sides with its respective atomic orbitals. And the molecular orbitals go in between. Here, each hydrogen has one electron in the 1s orbital. So they both go to fill up the bonding orbital, generating a molecule of H2. We can use these diagrams to calculate the bond order of the molecule, which is given by the number of electrons in bonding orbitals minus the number of electrons in antibonding orbitals divided by 2. Here, that gives us a bond order of one, which is why H2 has a single covalent bond. Now look at O2. Each atom displays its electron configuration. And every electron will go into a molecular orbital. There's the 1s atomic orbitals, which are both full, that go to make the 1s sigma and 1s sigma star, which are also both full, then the 2s orbitals, also both full, which go to make the 2s sigma and 2s sigma star, again, both full. Next, we have the 2p orbitals. Each oxygen atom has four 2p electrons. There are six 2p orbitals total, three per atom. So we need six 2p molecular orbitals. Those will be, in order of increasing energy, the sigma, pi, pi star, and sigma star. Remember, the starred ones are antibonding orbitals. As we take the four electrons from each oxygen atom and place them in the molecular orbitals bottom to top, we end up with full sigma and pi orbitals as well as half full pi stars. Calculating bond order, we have 10 minus 6 over 2, which gives us 2. That's why atoms in an oxygen molecule have a double bond between them. We can predict the covalent bonding behavior of any two atoms using orbital diagrams like this one. Let's look at nitrogen. Here, the bond order calculation turns out to be 3, which is why atoms in a nitrogen molecule have a triple bond between them. Let's check comprehension. [MUSIC PLAYING] Thanks for watching, guys. Subscribe to my channel for more tutorials. And as always, feel free to email me, professordaveexplains@gmail.com. [MUSIC PLAYING]