Chemistry

Unit 5: Advanced Theories of Bonding

Now that we've considered the names and shapes of molecules, it's time to go deeper into the formation of atomic bonds.

Thus far, we've focused on the fact that atoms can trade or share electrons. However, the ways that these electrons get shared will change the type of bond we see.

The video by Professor Dave gives a general overview of these advanced theories, and the subsections go into greater detail.

Professor Dave Explains | "Valence Bond Theory, Hybrid Orbitals, and Molecular Orbital Theory"

Valence Bond Theory

In a simplistic view of bonding, atoms simply share electron orbitals. However, this doesn't explain the molecular shapes we see in nature. Instead, we see that electron orbitals change their shape depending upon the number of bonds. These orbitals affect the shape of the atom:

Sigma Bonds (σ): in single bonds, two atomic orbitals merge to create a dense cloud of electrons between the atomic nuclei.
Pi Bonds (π): in double and triple bonds, not every electron can go in the cloud between the nuclei. In the pi (π) bonds, some electrons will be above and below the nuclei. Basically, they avoid overlapping with the sigma (σ) bonds that are right between the atoms.

Najam Academy | "Pi and Sigma Bonds"

Hybrid Orbitals and Multiple Bonds

Hybrid Orbitals

When we look at the shapes of orbitals, the sigma and pi bonds discussed previously describe how electrons are shared. But before this sharing occurs, the electrons around the individual atoms might change shape.

Crash Chemistry Academy | "Hybrid Orbitals Explained - Valence Bond Theory"

Review: Orbitals Are Regions of Probability

As quick review, remember that atomic orbitals come in different shapes depending upon their distance from the nucleus and the energy. The shape of an orbital is not the exact path an electron travels — instead, the shape shows where the electron is most likely to be found when you observe the atom.

Crash Chemistry Academy | "Orbitals, the Basics: Atomic Online Tutorial — probability, shapes, energy"

Molecular Orbital Theory

Just when you thought chemical bonds might almost make sense . . . they get even more complicated! This video's a bit longer so it can go into more detail. Here are some key considerations:

Not all electron orbital combinations lead to a chemical bond. Electron orbitals can cancel each other out if they're out of phase, and this will prevent a bond.
Bonding orbitals have a lower energy than the original atomic orbitals. This leads to molecular bonds.
Antibonding orbitals have a higher energy because the electrons are out of phase. This prevents molecular bonds.

Molecular Orbitals in Solids

In solids, the molecular orbitals stretch across the entire solid, and they can be very close together in energy. Because of this, there are so many orbitals that we refer to the collection bonding and antibonding orbitals as bands.

Electrons can jump from bonding orbitals to antibonding orbitals. The energy needed for this jump is the band gap. The size of the band gap affects how easily electrons can move from those bonding to antibonding orbitals — and electrons are electricity.

Conductors have no band gap, so electrons can easily move across the material.
Insulators have a large band gap, so it's very difficult for electrons to travel across the material.
Semiconductors have a moderate band gap, so electrons need some extra energy to get to the conduction band.

Professor Dave Explains | "Conductivity and Semiconductors"

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