Sigma bonds come in six varieties: Pi bonds come in one

by James

in Alkanes, Alkenes, Chemical Bonds, Organic Chemistry 1, Where Electrons Are

You may recall from Gen chem (and no doubt your first week of o-chem as well), that orbitals on carbon come in two flavors: s and p.

s orbitals should be familiar as the spherical-shaped orbitals. The electrons of hydrogen, for instance, are in a 1s orbital.

p orbitals are shaped like figure-eights, or loops. The electrons in p orbitals are slightly farther away from the nucleus than those in s orbitals, so they are a little bit higher in energy. The p orbitals therefore fill with electrons only after the s orbitals are filled.

Hybridization is a concept that hasn’t been addressed here yet, but there’s an excellent video discussion of it here.

What’s observed from analyzing the structure of molecules such as CH4 is that the shapes cannot result from the electrons being in s or p orbitals alone, but instead are a consequence of the electrons in s and p orbitals mixing to form hybrid orbitals. If you draw an analogy to how we could make a “hybrid” soft drink by mixing different proportions of Sprite and Pepsi, these new orbitals aren’t fully s or fully p, but are a combination of both. The “flavor” of each bond depends on the relative proportions of s orbital and p orbital content:

sp3 = 25% s character, 75% p character

sp2 = 33% s character, 66% p character

sp = 50% s character, 50% p character.

It is these hybrid orbitals that form sigma bonds (σ bonds). Sigma bonds are created from the head-on overlap of orbitals. Those orbitals will be some combination of s orbitals and p orbitals.

[What’s happened to the missing p orbitals for sp2 and sp hybridized carbon? Well, these p orbitals aren’t involved in hybridization or in sigma bonding – instead, they maintain 100% p character, are available to form π bonds, which are created by the side-by-side overlap of orbitals. π bonding involves p orbitals exclusively.*]

Since carbon can exist in one of these three hybridization states, we can therefore have six varieties of carbon-carbon sigma bonds:

Now, recall that for any given quantum number, s orbitals are lower in energy than p orbitals. The electrons in s orbitals are held more closely to the nucleus than electrons in p orbitals. Take a look at the relative bond lengths and strengths for each of these six situations.

General principle – the more s character the bond has, the more tightly held the bond will be.

[Note that this is also why the protons in acetylene are more acidic than those in ethene and ethane – the electrons in the conjugate base are more tightly held, making them more stable. If it seems contradictory that the bond is both stronger and yet more easily broken, remember that bond energies measure homolytic cleavage]

Now, the number of π bonds that can form will be dependent on the number of unhybridized p orbitals available – 1 for sp2 hybridized carbons, 2 for sp hybridized carbons (the two π bonds will be at right angles to each other in the latter case).

In contrast to sigma orbitals, there is only one way to form a C-C π bond – from the overlap of two p orbitals.

Question: Relative to sigma bonds, do you think p-p bonds (π bonds) would be stronger or weaker?

1) What’s their % s-character?

2) Energy required to break C-C bond in ethane:

Energy required to break apart the C=C bond in ethene:

*green text = not the whole truth! ( but close enough for your course)

Thank you, that was really helpful. Could you explain the diplole moment, its calculation, bent, linear shape, trigonal, and tetrahedral shape effect on dipole moment. Also, could please an example lets say why does Acetonitrile is more polar than Methanol? From electronegativity perspective, Oxygen has a higher electronegativity than Nitrogen but why in case of Acentonitrile it’s having higher electronegativity than Methanol? Does the triple bond has an effect due shortness of its bond compared to single bond?

Thank you ver much

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