Valence Bond Theory – Chemical Bonding

00:02 In the previous lecture, we talked about electrons and orbitals—specifically, atomic orbitals—so, where electrons reside in their elemental form. And we touched briefly upon what happens in covalent bond formation—that is to say, electrons which want to achieve a low-energy state by forming a complete outer shell can share electrons. Well, what actually happens in that process? It isn't just a simple dot-and-cross diagram that perhaps you were familiar with when you were doing very early qualifications. There's actually more to it than that, and that is molecular bond formation. So let us take the simplest example: the molecule hydrogen. As I said to you before, many atoms don't exist in atomic form in nature; they exist in molecular or ionic form. And here we have the gas H2, or hydrogen. You'll be familiar with H2 potentially, but you won't be aware, necessarily, that what's happening in this scenario is that the electrons which are being shared are being shared as part of a molecular orbital. And that molecular orbital is represented here, in the center of the board, where you can see one hydrogen atom with an electron configuration of 1s1 sharing its electrons with another hydrogen atom, which also has an atom configuration of 1s1. By sharing electrons, the hydrogen now has, in both cases, a full outer shell—specifically, now 1s2. So, like helium, it has now become more stable. To make this happen, the sigma orbitals of… sorry, the s orbitals of the hydrogen overlap to form a sigma bond. And this changes the shape of our atomic sigma orbital... atomic s orbitals into our covalent sigma orbitals.

02:01 The sigma orbital, as you can see here, is shown when one s orbital from a hydrogen overlaps with another s orbital from the hydrogen, and this forms, as you can see on the board, a molecular orbital, which is denoted sigma. It is a molecular, rather than an atomic, orbital but still contains electrons with opposite signs. Remember what we said about the Aufbau and Hund's rule? The fact is that molecular orbitals, like atomic orbitals, tend to be occupied as single-spin electrons before pairing takes place. This orbital holds the atoms together in a bond: a covalent bond. And in this case, when you have two atoms which are identical, the electron density is fixed in the center of that molecular orbital.

02:56 The better the overlap of the orbitals, the stronger the bond, and also, therefore, the closer the atoms can get to each other and the shorter the bond. Both of the orbitals must have the same phase. Now, if you go back to when I talked to you about quantum mechanics, you'll be aware that there is a phase, a wave–particle duality, associated with electrons which are involved in the bonding. We can't treat them as particles. In the case of the sigma bond, though, being formed from two s orbitals, there is no phase issue, because if you recall, they have the same phase in an s orbital. But as you'll come to see when we talk about p orbitals, this phase issue becomes important.