Using Orbital Hybridization and Valence Bond Theory to Predict Molecular Shape
- 0:01 Hybridization
- 1:37 Orbital Hybridization Theory
- 3:00 Sigma and Pi
- 3:24 Number of Orbitals
You'll learn how to explain how shapes of molecules can be predicted using valence bond theory and hybridization. When finished, you'll understand the difference between sigma and pi bonds and how the VSEPR theory, along with the hybridization theory, helps predict the shape of a molecule.
When you think of hybrid, you may think of a car, like the Toyota Prius, or an animal, like the zebroid or the liger. But in this lesson, we are using the term 'hybrid' to refer to the orbitals of electrons in atoms.
Hybridization is the mixing of two or more atomic orbitals to form new orbitals that describe the covalent bonding in molecules. Orbital hybridization shows the relationships between the geometry of a molecule and the orbitals of the bonding electrons. Hybridization works because the net energy of the hybridized bonding electron orbitals is reduced compared to the un-hybridized orbitals.
Orbital Hybridization Theory
In order to figure out the hybridization, we need to know the valence electrons in the participating atoms and the valance bond theory. The valence bond theory essentially says that all bonds are made by an atom donating a valence electron to another atom to complete its octet. The theory, combined with knowledge of valence electrons, tells us how many bonds there are between two atoms in a molecule. The VSEPR theory, as you've learned previously, helps predict the shape of a molecule based on the repulsion of the electrons in the orbitals. The VSEPR theory fails to explain all of the interactions scientists see in molecules, though. So they've developed the concept of orbital hybridization.
Take methane as an example (CH4). The carbon atom has four valence electrons: two in the 2S orbital and two in the 2P orbital. With what you've learned, it might not make sense that these four electrons form bonds in a tetrahedral shape, because two electrons in the S orbital are already paired up compared to the two single P orbital electrons. You may think it would only have three orbitals, not four.
You also know that the S orbital is spherical and the P orbital is dumbbell-shaped. To explain the known bonding of the carbon atom, you have to assume that the 2S and 2P orbits get combined and rearranged to make four orbitals. In other words, they get hybridized.
A good analogy for understanding hybridization is colored water. Start with one beaker with 50mL red water that represents the 2S orbital and three beakers, each with 50mL blue water, that represent three 2P orbitals. Mix all four beakers and get 200mL of purple water. Divide the purple water into four beakers with 50mL in each. These are the hybrid orbitals. Just as each of the beakers is made of a mixture of red and blue water, each hybrid orbital is made of a mixture of 2S and 2P orbitals.
Sigma and Pi
There are two types of covalent bonds: the very strong sigma and the not-as-strong pi. The sigma bond is when two orbitals directly overlap but there is only one bonding interaction. A pi bond is weaker than the sigma bond. This overlap occurs when two orbitals overlap and there are two bonding interactions. It looks like two dumbbells put side-by-side and overlapped.
Number of Orbitals
The number of hybrid orbitals that are made equals the number of orbitals that have combined. So in the case of methane, there is one S orbitals and three P orbitals for a total of four. There are three P orbitals, even though there are only electrons in two of them. This is written as sp3. The superscript of three shows that three P orbitals were included in the hybridization.
The four orbits on the carbon are located 109.5 degrees apart. This allows the orbitals to be as far apart from each other as possible, which leads to a tetrahedral shape. Each of these orbitals has one electron from the carbon and can overlap, or merge, with another atom that has room in its electron orbital. This forms a covalent bond because they are sharing their electrons between two atoms.
If an atom has only two electrons to share, as in oxygen, the electron energy shell levels look like this. You can see that it has one S orbital and two P orbitals (the third p orbital is full), so it is designated as sp2. When you draw an orbital picture of it, it has three total orbits, which are angled 120 degrees apart and are planar, or flat.
Hybridization rules can be summarized:
Hybridization isn't for single atoms. It is a model used to explain covalent bonding between two or more atoms.
- Hybridization is the mixing of two atomic orbitals, such as S and P orbitals.
- A hybrid orbital is not a pure orbital, so hybrid orbitals and pure atomic orbitals have different shapes.
- The number of hybrid orbitals is the same as the number of atomic orbitals that participate in the process.
- Although hybridization requires energy, the system has a net loss of energy because of the lower bond energy of each bond.
- Covalent bonds are formed by the overlap of hybrid orbitals.
Chapters in Chemistry 101: General Chemistry
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