In an effort to understand molecular structure in terms of atomic orbitals, chemists developed valence-bond theory in the 1920s. While VB theory doesn't explain the properties of all molecules, it's a good model that is especially useful in organic chemistry. The concept of hybrid orbitals is important in VB theory. Before you start, you should already have a good working knowledge of Lewis dot structures and VSEPR (valence shell electron pair repulsion).
- Skill level:
- Challenging
Instructions
Things you need
- Periodic table of the elements
-
1
Draw the most plausible Lewis dot structure for the molecule you want to analyze. Recall that an atomic orbital is the wave function of an electron, which is obtained as a solution of the Schrodinger equation. The simplest way to think about an atomic orbital is as a region of high probability density, meaning the region within which the electron is most likely to be found. VB theory describes bonds between atoms as regions of orbital overlap, where orbitals from both atoms overlap or fuse to form a single region of high probability density. In the hydrogen molecule, for example, the s-orbital of each hydrogen atom would overlap with the other to form a continuous region of high probability density around both atoms. Bonding between atoms will often necessitate hybrid orbitals, which are combinations of the orbitals in an atom (combinations of an s- and p- type orbital, for example).
-
2
Count the number of regions of electron density around an atom in the Lewis dot structure to determine its hybridization. A lone pair counts as one region; a bond, whether it be single, double or triple, also counts as one region. The oxygen atom in water, for example, has four regions of high electron density (two lone pairs and two bonds), while either of the two carbon atoms in ethene (C2H4) has three regions of high electron density (the double bond counts as only one region).
-
3
Determine how many hybrid orbitals the atom needs. The atom will have one hybrid orbital for each region of electron density you counted on the Lewis dot structure. The oxygen in water, for example, will have four hybrid orbitals, whereas either of the carbons in ethene will have three hybrid orbitals. A hybrid orbital is formed by combining an s-orbital with one or more p- and d-orbitals to form a hybrid that looks somewhat like a cross between both orbital types. It takes N atomic orbitals to form N hybrid orbitals, so the number of hybrid orbitals is the same as the number of orbitals you combine to make them.
-
4
Determine the atom's hybridization scheme based on the number of hybrid orbitals you need. An sp-hybrid has two hybrid orbitals, whereas an sp2 hybrid has three and an sp3 hybrid has four. Some elements from period 3 and below can also use d-orbitals to form additional hybrid orbitals -- sp3d has four hybrid orbitals, for example, while sp3d2 has five. In organic chemistry you'll typically be working with carbon compounds, so sp3 should be the most you'll see.
-
5
Describe the molecule's geometry by inferring it from the hybridization scheme. Owing to repulsion between electrons, the regions of electron density will separate out as far as possible, so sp-hybrid orbitals are linear, while sp2 are trigonal planar and sp3 are tetrahedral. Double and triple bonds are formed by unhybridized overlapping p-orbitals. The oxygen atom in water, for example, is sp3-hybridized, so its hybrid orbitals will have a tetrahedral arrangement.
Tips and warnings
- Note that there is an exception to these rules. If an atom is surrounded by two or more adjacent p-orbitals and has a lone pair, the lone pair can move into one of the p-orbitals and the atom will have one less hybrid orbital than we would expect based solely on the Lewis dot structure. The most common example is an amide group (a carbonyl group bonded to a nitrogen bonded to a hydrogen). We would expect, based solely on the Lewis dot structure, that the nitrogen in an amide would be an sp3 hybrid. But because of the adjacent p-orbitals, it is actually an sp2-hybrid and thus forms bonds arranged in a trigonal planar fashion.
