Reading Notes for Chapter 8


These are Dr. Bodwin's reading notes for Chapter 8 of "Chemistry 2e" from OpenStax. I am using a local .pdf copy that was downloaded in May 2020.

Chapter Summary:

As with many topics, there are a variety of different ways we can envision and explain chemical bonding. Many of these "competing" theories exist and are valid within their limitations. In general, always use the "simplest" theory to explain whatever you are trying to explain; just because a theory or model is more complex, doesn't mean it does a better job of explaining an observation.

Valence Bond Theory:

When orbitals on adjoining atoms overlap, they allow sharing of the electrons between them and form a bond.
VB theory treats orbitals on atoms as distinct and fully "belonging" to one atom, although the electrons that are shared can bounce back and forth between the atoms.
"Stronger" bonds have more "effective" overlap.
The energy well graph in Figure 8.2 is a very important concept in chemistry. It is widely used to describe energy-based interactions between charged particles.
There are a number of different ways that orbitals can overlap simply based upon their shape and orientation. This terminology is used common in chemistry courses beyond General Chemistry, so it's important to specify here:
  1. Sigma bond/interaction - these are the simplest overlaps where there is one point of interactions. They're called "sigma" because this is the only interaction that two "s" orbitals can have, as shown in Figure 8.4a. All orbital types can form sigma interactions. In a molecule or polyatomic ion, the first interaction between any 2 atoms is almost always a sigma interaction.
  2. Pi bond/interaction - This is a two-point overlap as shown in Figure 8.5. This type of overlap is a "side-by-side" interaction that can take place between two "p" orbitals, or a "p" and a "d" orbital, or two "d" orbitals. When we form double or triple bonds in molecules and polyatomic ions, the second and third interactions between any 2 atoms is almost always a pi interaction.
  3. Delta bond/interaction - These aren't discussed in your textbook, but as long as we're looking at interactions, we might as well include this one. A delta overlap is a "face-to-face" interaction that can take place between two "d" orbitals. Since both orbitals have to be "d" orbitals, this type of interaction only occurs between transition metals, so it is a less common interaction when we are looking at "simple" molecules. For some good pictures and further description of delata interactions, check out:
  4. https://www.quora.com/What-is-a-delta-bond-and-how-do-d-orbitals-form-a-delta-bond
What is an "orbital"? As we get deeper into our discussion of bonding and bonding theories, let's not forget what an orbital is. An orbital is a visual representation of a mathematical probability surface where we are most likely to find an electron or electron "pair". The red/blue shading in your textbook represents the "sign" of each lobe of an orbital, and effective overlap can only occur between lobes of the same sign. That's why the pi interaction is specifically shown with the blue lobes of the p-orbitals lined up and the red lobes lined up. This is an important concept that we'll get back to soon...

Hybrid Orbitals:

The cartesian geometry of atomic orbitals causes some problems when we look at the geometry of actual molecules. We could explain that away by just saying the orbitals overlap at an angle, but it's hard to come up with a good reason for that. Since orbitals are just visualizations of a mathematical relationship, it seems reasonable that we can combine those orbitals to form new orbitals that retain some character of the old ones. These are hybrid orbitals and can be used to explain a much wider range of molecular geometries than the standard "s", "p", "d", and "f" orbital set we have already developed.
Forming hybrids - We can think in simple visual math terms. Orbitals are added and subtracted to give a new set of shapes. The number of atomic orbitals that go in to making the hybrids must equal the number of hybrids we get out. Look at the left side of Figure 8.8. Remember, we said that the red and blue color represent the +/- "sign" of the orbital. If we superimpose the "s" and "p" orbital shown in the picture, where blue overlaps with blue, the blue part gets bigger; where blue overlaps with red, the red part gets smaller. That's the first hybrid orbital shown on the right side of the picture. Now, if we "subtract" these instead of "adding" them, change the color of the s-orbital to red (the "negative" of blue) and repeat the process. That gives the 2nd hybrid orbital on the right side of the figure. Two "standard" orbitals go in, two hybrid orbitals come out!
Hybrid Orbital Types:
  1. "sp" - one s and one p orbital form 2 "sp" hybrid orbitals. These orbitals are 180° apart and result in linear electronic geometry.
  2. "sp2" - one s and two p orbitals form 3 "sp2" hybrid orbitals. These orbitals are 120° apart and result in trigonal planar electronic geometry.
  3. "sp3" - one s and three p orbitals form 4 "sp3" hybrid orbitals. These orbitals are 109.5° apart and result in tetrahedral geometry.
  4. "sp3d" - one s, three p, and one d orbital form 5 "sp3d" hybrid orbitals. These orbitals are 180°/120°/90° apart and result in trigonal bipyramidal electronic geometry.
  5. "sp3d2" - one s, three p, and two d orbitals form 6 "sp3d2" hybrid orbitals. These orbitals are 180°/90° apart and result in octahedral electronic geometry.
Those bond angles are the ideal angles in a "perfect" shape... bond type, bond length, bond strength, atom type, and other things can cause deviations.
Hybrid orbitals work the same way as "standard" s/p/d/f orbitals, they just allow different bond angles and interactions.

Molecular Orbital Theory:

Linear Combination of Atomic Orbitals (LCAO) - in some resources, Molecular Orbital Theory ("MO Theory" of just "MO") is explicitly called "LCAO-MO" because we get molecular orbitals by taking linear combinations of atomic orbitals. It's similar to the method we used to "make" hybrid orbitals, but now it's done on the molecule as a whole.
Your textbook shows some nice pictures of simple MOs. Start with the simple ones because some MOs can get pretty complex!
The pictures can sometimes get a little ponderous, so it's also important to start to think about these as energy level diagrams. You've seen some simple electron-filling diagrams before, a MO energy level diagram if just a little extension of that. Check out Figure 8.34 for a nice example.
These energy level diagrams can also get pretty complex, but the simple ones aren't too ponderous and give us a good foundation for how they work.

MO diagrams are extremely powerful tools, but again, don't use a more complicated tool if you don't need to. You don't need to take a Ferrari or a semi tractor-trailer to the grocery store for a loaf of bread, use something simpler. Same thin here. If all you want to do is look at the bonding an carbon dioxide, Valence Bond theory works great! If you need to understand why molecule oxygen behaves as if it has unpaired electrons, you might need to step up to MO theory.





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