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Why are sigma bonds stronger than pi bonds?

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  • $\begingroup$ Simply, it is due the fact that the extent of overlap in sigma bonds is greater than that in pi bonds. It is also interesting to note that because of this difference in strength, a double bond is not twice as strong as a single bond (which is sigma) because one of them will be a pi bond. $\endgroup$ – Nick Sep 19 '13 at 16:44
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The reason behind this is the orientation of the overlapped orbitals. Sigma bonds result from head-on(co-axial) overlapping while pi bonds are outcome of lateral(para-axial) overlapping. Here is a pictorial representation of ethene(sp2 hybridized C atoms) :

enter image description here

The greater the extent of overlapping, the higher the probability of finding the valence electrons in between the nuclei and hence the bond will be stronger & shorter.

In MOT, this can be explained using Overlap Integral. This is how Atkins depicts it :

enter image description here

In simple terms, after forming a sigma-bond (a pre-requisite for pi-bonds), the two atoms get locked along the inter-nuclear axis. As a result, the orbitals available for pi-bonding can only partially overlap, thus forming a weaker bond.

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    $\begingroup$ Bond strength and bond length do not necessarily correlate. With a higher overlap, there is also more inter electron repulsion (derived from the Pauli principle). It is also untrue, that a $\sigma$ bond is a necessity for a $\pi$ bond. The reason why $\pi$ bonds are weaker is the same as why p-orbitals have higher energy than s orbitals. $\endgroup$ – Martin - マーチン Jun 19 '14 at 3:00
  • $\begingroup$ @Martin : I presented as far my understanding. I appreciate that you point out the anomalies. You can post an answer or suggest an edit(doesn't seem viable due to lots of changes), if you care to ! $\endgroup$ – blackSmith Jul 2 '14 at 12:52
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As stated previously, it is due to the head-on overlap of sigma bonds and the lateral overlap of pi-bonds.

The smaller overlap of pi bonds also explains why double and triple bonds basically exist only for 2nd row elements (C,N,O especially) and not for higher row elements. A C=C bond has a length of 133 pm. A Si-Si bond length is at around 186 pm, therefore the contribution of the pi-integral overlap is almost negligible.

The inability of Silicium to form strong pi bonds is also one of the answer why life on earth is carbon based and not silicium based, since the richness of organic chemistry is in part due to the ability of carbon to form strong double and triple bonds.

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Unsaturated compounds (due to the existence of a double bond) are more reactive than saturated compounds, because the electron density at the internuclear axis is zero in a pi bond (in contrast to the sigma bond).

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Pi bonds involve sideways overlap while sigma bond involves head on or axial overlap. Axial overlaps have higher degree of overlapping than sideways overlap.

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    $\begingroup$ How does this answer differ from the accepted one, apart from it is shorter and has no explanation? $\endgroup$ – Martin - マーチン Jun 19 '14 at 3:03

protected by jonsca Jun 19 '14 at 13:00

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