Orbital Hybridization "Cheat Sheet"?

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I am studying for my AP Chem exam, and have come across this "Cheat Sheet." I have no clue what it means, or how to use it, except for the Electron Geometry (VESPR) Please explain what the sheet does and how to use it. Thanks!

1 Answer
Dec 26, 2017


See below.


Let's go column-by-column.


The first column gives the hybridization. This refers to the idea that atomic orbitals mix, or “hybridize,” to maximize distances between repulsive electron pairs.

  • A quick way to figure out the hybridization of a specific atom in a molecule ("central" atoms) is by the number of bonds that atom has. You begin with #"sp"# hybridization for two bonds. At three bonds, you get #"sp"^2# hybridization, and at four you get #"sp"^3# hybridization. We then move up to #"sp"^3d# and #"sp" ^3 "d,"^2# etc.

  • For example, the carbon in a methane #("CH"_4)# molecule has four (single) bonds, and has hybridization #"sp"^3#.

  • Note that it does not matter if the bonds are single, double, or triple. A triple bond still counts as one bond.


Orbitals Involved

This more or less just reiterates what is stated in the first column; it tells you which orbitals are involved in making the hybrid orbitals. So, for example, an #"sp"^3# hybrid atomic orbital involves 1 s orbital and 3 p orbitals. As mentioned above, #"sp"^3# hybridization implies four bonds.

Orbitals left unhybridized

Not all orbitals which are present are involved in the making of hybrid orbitals. This column tells you which orbitals are left over after the hybrid orbitals are formed.

  • For example, with #"sp"^3# hybridization, we have 1s and 3p orbitals involved in the hybrid orbital and none left over.

  • With #"sp"# hybridization, we have the overlap of an s and p orbital to form the hybrid #"sp"# orbital, but we have 2p orbitals left over, unhybridized.

The electron geometry column is as always, and it was indicated that this was understood.


This column tells you which types of bonds are present. A #pi# bond is the side-to-side overlap of atomic orbitals, whereas a #sigma# bond is the end-to-end overlap of atomic orbitals.

Here is an illustration of the side-to-side overlap of #"p"# orbitals to form a #pi# bond:

UCLA Chemistry

  • A single bond consists of just one #sigma# bond.

  • A double bond consists of one #pi# bond and one #sigma# bond.

  • A triple bond consists of one #pi# bond and two #sigma# bonds.

Here is a visualization of a double bond:


Hope this helps!