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Look at the molecular structure of benzene:

enter image description here

It's a perfect hexagon. Why aren't there any molecules arranged in a triangular fashion with bonds forming the edges and the molecules the vertices? Or, for that matter, square molecules? And what about other n-gonal molecules? octagons?

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    $\begingroup$ There are triangular molecules: en.wikipedia.org/wiki/Cyclopropane $\endgroup$
    – Philipp
    Commented Aug 28, 2013 at 5:39
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    $\begingroup$ Molecules whose shapes are connected to squares or higher polygons than a hexagon are usually subject to deformations due to ring strain etc., so that they deviate from the perfect polygon geometry. $\endgroup$
    – Philipp
    Commented Aug 28, 2013 at 5:47

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There exists triangular molecule, eg. cyclopropane, but they are highly reactive due to ring strain.

Ring strain results from a combination of angle strain, conformational strain or Pitzer strain, and transannular strain or van der Waals strain.

C-C-C bond angle in rings should be approximately 109.5° degrees which in case of small rings like cyclopropane, bond angles are 60° whereas tetrahedral 109.5° bond angles are expected. The intense angle strain leads to nonlinear orbital overlap of its sp3 orbitals. Because of the bond's instability, cyclopropane is more reactive than other alkanes. Since any three points make a plane and cyclopropane has only three carbons, cyclopropane is planar.

Due to this strain they exists in higher energy state and therefore Because of their high strain, the heat of combustion for these small rings is elevated.

Kindly see Ring Strain, it will help you to understand about these kinds of molecules.

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    $\begingroup$ You might want to comment on the actual underlying phenomenon behind ring strain, i.e. electron-electron repulsion. $\endgroup$
    – Aesin
    Commented Aug 28, 2013 at 20:33
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    $\begingroup$ Ring strain is an awful concept. It is purely built on phenomenology as in "stuff that one could be expecting from hybrid orbital angles." It has only little more scientific justification than the VSEPR model. $\endgroup$
    – Martin - マーチン ♦
    Commented Sep 8, 2014 at 11:00
  • $\begingroup$ @Martin : But Ring Strain is the only reason that was explained to me in 12th grade. $\endgroup$
    – ashu
    Commented Sep 8, 2014 at 11:49
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    $\begingroup$ Unfortunately topics like this will always be taught as something that appears to be a law, without questioning the reasons, that underlie this concept. Most of steric repulsion, attraction and other effects can be described and explained by more advanced theories, i.e. electronic structure theory. For someone in school it might be appropriate to conclude with models like these, but if the underlying theory is not too far off it should be included as well. $\endgroup$
    – Martin - マーチン ♦
    Commented Sep 8, 2014 at 12:38
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    $\begingroup$ While ring strain may be phenomenological, I think it should be taught. The interactions of atoms and orbitals in tight rings does lead to a strained high energy structure, and opening the ring releases this energy, relieving the strain. It would be complicated to approach all those interactions from the ground up, the concept of ring strain wraps that up nicely. And with a good molecular model kit, you can "experience" ring strain by jamming a cyclopropane together. $\endgroup$
    – user137
    Commented Sep 9, 2014 at 6:14
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white phosphorous allotrope (P_4) is a deltahedron (tetrahedron) consisting of four triangles fused by mean of each side. It doesn't bear that much angular strain, cause the measurements done on 2nd-row elements-based compounds (i.g. hydrocarbons and related) does not necessarily include 3rd-row and heavier elements. This also holds in non cyclic compounds. While the H-O-H in water is close to theoretically predictable 109,5°, the H-S-H angle in hydrogen sulfide is very close to a right angle, showing little or no hybridization to have occurred. Moreover the rules with closed shell molecules and non-electron deficient ones, appear to be inadequate to descrive electron deficient compounds, spanning from closo-boranes (very often deltahedranes) to metal clusters P.S. sorry for my english :)

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The best example for relatively stable 3 membered rings are the epoxides. They are very reactive, yes, but they can be isolated and stored, and then used in synthetic routes, like those sold by sigma aldrich:

http://www.sigmaaldrich.com/materials-science/material-science-products.html?TablePage=20203736

You probably wouldn't have guessed that all of those guys above did exist, would you? :-D

You can also understand its stability trough the molecular orbital theory. It turns out that molecular orbitals of these 3 membered rings, like cyclopropanes or epoxides has quite similar molecular orbitals compared to those of alkenes, so their reactivity is somehow similar.

Also, square, penta-, hepta-, octa-....gonal molecules do exist! You would be surprised about some macrocyclic compounds being stable like those over here:

http://www.google.com/patents/US6200254

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