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Valence Bond Theory – Chemical Bonding

by Adam Le Gresley, PhD
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    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. Let's have a look at another example of a


    About the Lecture

    The lecture Valence Bond Theory – Chemical Bonding by Adam Le Gresley, PhD is from the course Chemistry: Introduction.


    Included Quiz Questions

    1. To achieve the lower energy states and stabilities by achieving noble gas like configuration in their outer shells
    2. To get rid of excess negative charge in their outermost shells
    3. To get rid of excess positive charge in the nucleus relative to number of electrons
    4. To achieve high reactivities by having excess of electrons in their outermost shells
    5. To achieve the higher energy states by achieving halogen like configuration in their outer shells
    1. The electrons present in their molecular orbitals
    2. the electrons present in their atomic orbitals
    3. The electrons present in the core atomic orbitals
    4. The electrons present in the middle atomic orbitals
    5. The electrons having anti-spins only
    1. At the center of molecular orbitals
    2. At the center of atomic orbitals
    3. At the corners of atomic orbitals
    4. At the corners of the molecular orbitals
    5. Near the nucleus of a negatively charged oxygen atoms
    1. Better overlapping of molecular orbitals with each other
    2. Atomic mass of bigger atom participating in covalent bonding
    3. Atomic mass of smaller atom involved in covalent bonding
    4. Difference in the atomic masses of the participating atoms
    5. Number of electrons participating in covalent bonding

    Author of lecture Valence Bond Theory – Chemical Bonding

     Adam Le Gresley, PhD

    Adam Le Gresley, PhD


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