Notice that the bond angles around each carbon are all 120°. Each carbon is bonded to three other atoms in the same kind of plane trigonal configuration that we saw in the case of boron trifluoride, where the same kind of hybridization occurs. We can explain this trivalence by supposing that the orbital hybridization in carbon is in this case not sp 3, but is sp 2 instead in other words, only two of the three p orbitals of carbon mix with the 2 s orbital to form hybrids the remaining p-orbital, which we will call the i orbital, remains unhybridized. Here, we can regard carbon as being trivalent. Figure 9.11. An sp 3 hybridized atomic orbital.Īccording to VSEPR theory, the four degenerate orbitals will arrange as far apart from each other as possible, giving a tetrahedral geometry with each orbital 109.5 o apart (Figure 9.12 “ A carbon atom’s four tetrahedral sp 3 hybridized orbitals”).Trigonal hybridization in carbon: the double bondĬarbon and hydrogen can also form a compound ethylene (ethene) in which each carbon atom is linked to only three other atoms. This larger lobe is typically used for orbital overlap in covalent bonding. The sp 3 orbitals, being a combination of a spherical s orbital and propeller- (or peanut-) shaped p orbital, give an unsymmetrical propeller shape where one lobe of the orbital is larger (fatter) than the other (Figure 9.11 “ An sp 3 hybridized atomic orbital”). Hybridization of carbon to generate sp 3 orbitals. Note that in hybridization, the number of atomic orbitals hybridized is equal to the number of hybrid orbitals generated. The 2 s and three 2 p orbitals are averaged mathematically through hybridization to produce four degenerate sp 3 hybrid orbitals (Figure 9.10 “ Hybridization of carbon to generate sp 3 orbitals”). In 1931, Linus Pauling (Figure 9.9) proposed a mathematical mixing of atomic orbitals known as hybridization. Pi bonds result from the sideways overlap of p orbitals, placing electron density on opposite sides of the internuclear axis (Figure 9.7 “Pi bond diagram showing sideways overlapping of p orbitals”).įigure 9.9. A diagram showing the overlap of s orbitals of two hydrogen atoms to form H 2.įor molecules that contain double or triple bonds, one of these bonds is a sigma bond, and the remaining multiple bonds are a different type of bond known as a pi bond (π bond). The optimal distance between atoms, which maximizes the attractive forces and minimizes the repulsive forces, gives the H-H sigma bond a length of 74 pm. Repulsion forces between the two nuclei and between the two electrons are also present. The 1 s orbitals of the two hydrogens approach each other and overlap to form a bond that has cylindrical symmetry known as a sigma bond (σ bond). Let’s examine the simplest case of atomic overlap resulting in a covalent bond, the formation of H 2 from two hydrogen atoms (Figure 9.6 “ A diagram showing the overlap of s orbitals of two hydrogen atoms to form H 2“). Since these electrons are simultaneously attracted to both nuclei, the electron pair holds the two atoms together. This creates an area of electron pair density between the two atoms. The valence bond theory states that atoms in a covalent bond share electron density through the overlapping of their valence atomic orbitals. Representations of s and p atomic orbitals.
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