# hybridization – Why can there not be more than one sigma bond in a set of bonds?

## The Question :

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A question on an exam asked why there is exactly one sigma bond in double and triple covalent bonds. I looked in my text and online after the exam, but couldn’t find an anawer to the question.

Why can there not be more than one sigma bond in a set of covalent bonds?

• What atomic orbitals overlap to form a sigma bond?
• @LordStryker, the s and p orbitals overlap to form the orbitals involved in sigma bonds, if that’s what you mean.
• “Why can there not be more than one sigma bond in a set of covalent bonds?” Actually that is quite a profound question. Your thinking is far ahead of your test. The answer to your question is that while introductory texts often display double and triple bonds as one sigma bond and the rest pi bonds, they can also be equivalently described as 2 or 3 “bent” sigma bonds. So double bonds and triple bonds can be described using only sigma bonds or as a mixture of pi and sigma bonds.
• Strictly speaking, you can get two $\sigma$ bonds between the same two atoms, though it is rare. One example is in the gaseous dimolybdenum molecule $\ce{Mo2}$ with its sextuple bond, which has both a $s\! -\! s$ $\sigma$ sigma bond and a $d\! -\! d$ $\sigma$ sigma bond, as well as two $d\! -\! d$ $\pi$ pi bonds and two $d\! -\! d$ $\delta$ delta bonds. I’ve never heard of a double sigma bond for $\ce{C=C}$, though in some interpretations dicarbon $\ce{C2}$ has two $\pi$ bonds with no $\sigma$ bond.
• This is where chemistry really gets interesting — when you try to pin a concept down and, after some wild times down the rabbit hole, find out that things are far more subtle and more weird than you ever thought possible. Nearly every concept is initially taught at a (relatively) comprehensible, modest-to-extreme level of approximation. As learning proceeds, the approximations are gradually stripped away until one finally starts to bump up against the limits of human knowledge.

There can be, even in simple carbon compounds. Bent bonds, tau bonds or banana bonds; whatever you might like to call them were proposed by Linus Pauling; Erich Hückel proposed the alternative $\sigma – \pi$ bonding formalism. Hückel’s description is the one commonly seen in introductory texts, but both methods produce equivalent descriptions of the electron distribution in a molecule.