- Online Text
- 1. Introduction
- 2. Valence Electron Patterns and Lewis Structures
- 3. Ionic Bonds
- 4. Covalent Bonds and the Octet Rule
- 5. Polarity and Basic Lewis Structures
- 6. Advanced Lewis Structures
- 7. VSEPR Theory
- 8. Hybrid Orbitals
- 9. Intermolecular Forces
- 10. Physical Properties of Molecules
- 11. Conclusion
- 12. Further Reading
- Unit Guide (PDF)
Section 8: Hybrid Orbitals
So far, we've talked about sharing electrons in valence shells and we've seen pictures illustrating the bonds connecting the various components of molecules. But there is a more precise and fundamental way to discuss the manner in which atoms combine to form molecules: We can describe the combining of the "electron clouds" that surround the atomic nuclei. It is possible to determine the shapes and charges of these combined clouds, which tells us a lot about the behavior of the molecule. To describe these clouds we have to mathematically merge the atomic orbitals together to create new orbitals. When an electron is shared between two different atoms in a covalent bond, the two electron clouds that correspond to each of the atomic orbitals merge to become a larger cloud. This combination of atomic orbitals is called a "molecular orbital."
σ (sigma) and π (pi) Bonds
Between any two atoms, there is at least a single covalent bond, and there need to be molecular orbitals to make this connection. The molecular orbitals that form these single covalent bonds are called "σ (sigma) bonds"; this comes from the fact that they are derived from the valence s orbitals and s in Greek is a sigma (σ). Single covalent bonds have a cylindrically symmetric molecular orbital; the electrons move at identical distances around the axis between the two atomic nuclei (Figure 5-16). Sigma bonds are flexible in that the atoms in them can spin relative to one another while maintaining their strength. So, if a sigma bond is holding two atoms together, the molecule can rotate freely about that bond.
FIgure 5-16. σ (sigma) Bonds
This two-carbon molecule is shown here with its single bond. Due to the shape of the sigma bond, one tetrahedral end of the molecule can rotate relative to the other end.
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A double bond consists of a sigma bond and a second bond whose molecular orbital is not symmetric around the nuclear axis (Figure 5-17 (top)). This second bond is called a "π (pi) bond " and it is shaped like a hot dog bun; a pi bond is called a pi bond because it is derived from p orbitals, which in Greek would be the letter π (pi). Recall that p orbitals have two portions on opposite sides of the nucleus, which is why a pi bond has two parts. Unlike sigma bonds, pi bonds are rigid; they cannot twist, because if the molecule were to twist, the pi bond would break. Molecules that are held together by pi bonds are unable to rotate and create a rigid planar section of a molecule. A double bond is stronger than a single bond, but it is not twice as strong because pi bonds are not as strong as sigma bonds.
Figure 5-17. σ (sigma) and π (pi) Bonds
The two-carbon molecules are shown here with the two possibilities of what can connect them together with multiple bonds: a double bond or a triple bond. While a single σ (sigma) bond is free to rotate, double and triple bonds are rigid because of their π (pi) bonds.
© Science Media Group.
Accordingly, a triple bond consists of a sigma bond and two pi bonds (Figure 5-17 (bottom)). A triple bond will always be found in a geometrically linear portion of the molecule. While a triple bond is the strongest type of bond, it is not three times stronger than a single bond because pi bonds are weaker than sigma bonds.