BNG_Valence Bond Theory Lesson

The Scientific Method

Lewis structures and the VSEPR theory are extremely useful; however, they do not tell us specifically how the electrons are shared. Two theories attempt to explain how electrons are shared in covalent bonds. They are called the valence bond theory (VB theory) and the molecular orbital theory (MO theory). We will focus on the VB theory in this course. We have already been using the VB theory to explain bonding throughout this module, so it should sound consistent with what you have already learned.

The VB theory explains that covalent bonds are formed between two atoms by the overlap of half-filled valence atomic orbitals. The theory assumes that electrons occupy atomic orbitals of individual atoms within a molecule, and that the electrons of one atom are attracted to the nucleus of another atom. A bond between two atoms is formed when a pair of electrons with their spins paired is shared by two overlapping atomic orbitals. This shared area of electron density (or area where the electrons are likely to be found) is what forms the covalent bond. Let's use the VB theory to explain the formation of the hydrogen bond, H₂. First, you must recall the electron configuration (focusing on the valence electrons) of hydrogen,1s1 . Each of these s electrons is shown below, represented by a sphere. They overlap to form the bond.   You can also see by the orbital diagram how these electrons can fit together.

Now, let's explain the formation of the HF molecule using the VB theory. To begin, write the electron configuration of each atom. Remember, that we are only concerned with the valence electrons, so noble gas configurations will be sufficient.

H             1s¹        

F               [He] 2s², 2p⁵

In order to figure out which specific electrons are forming the bond, we can draw an orbital diagram. Remember to use Hund's Rule to distribute the electrons (with each orbital half full before any is full).

Now we can see that the s electron from H will overlap with the p electron from F.

This can be shown in a drawing like this:

 

or in a combined orbital diagram like this:

Hybrid Orbitals

Sometimes theories need to be modified to explain observations. This is exactly what happened in the VB theory. Read the following slides to see what observation was made that required modification to the VB theory and how the theory was changed.

As explained above, electron orbitals can mix to form new orbitals known as hybrid atomic orbitals. These new hybrid orbitals are equal to each other in shape and energy. When one s and one p orbital mix, two new  sp hybrid orbitals  are formed. When one s and two p orbitals mix, three new sp² hybrid orbitals are formed. When one s and three p orbitals mix, four new sp³ hybrid orbitals are formed. You will see reference to sp3d, sp3d2, etc. hybrid orbitals. However, current evidence suggests that hybridization involving d orbitals does not exist. So, disregard any references to such orbitals. Links to an external site.

These new orbitals have new shapes and new directional properties, and they can be overlapped to give structures that have realistic bond angles. It's important to realize that hybrid orbitals are part of the valence bond theory. They can't be directly observed in an experiment. We use them to describe molecular structures that have been determined experimentally. Hybrid orbitals have orientations around the central atom that correspond to the electron-domain geometry predicted by the VSEPR Theory as shown in the chart below.

Remember to work on the module practice problems as you complete each section of content.   

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