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I read about a hypothetical compound called hexazine on Wikenigma. It's a 6-membered ring all of whose atoms are nitrogen and they form alternate single and double bonds just similar to benzene. But it also says that this would be highly unstable. How can it be so much unstable when it's structure resembles so much to that of benzene which is exceptionally stable. Has it something to do with nitrogen's compact size?

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    $\begingroup$ Related: chemistry.stackexchange.com/questions/53864/… $\endgroup$ Mar 4, 2023 at 4:19
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    $\begingroup$ The decomposition to 3 molecules of nitrogen gas would be highly exothermic $\endgroup$
    – Waylander
    Mar 4, 2023 at 7:06
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    $\begingroup$ Well, chemistry is like that: all elements are different, even if they are neighbors. Size is certainly not an explanation: one N to another N is the same as one C to another C. $\endgroup$ Mar 4, 2023 at 7:21
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    $\begingroup$ It's hard to beat that NN triple bond. You can make N clusters but they are not rings: arxiv.org/ftp/arxiv/papers/1805/1805.00835.pdf $\endgroup$
    – Buck Thorn
    Mar 4, 2023 at 7:51
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    $\begingroup$ Notice relative bonding energies in simple, double and triple bonds of carbon and nitrogen. Get the benzene stabilization energy due aromaticity. Evaluate and conclude. $\endgroup$
    – Poutnik
    Mar 4, 2023 at 17:09

2 Answers 2

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How can it be so much unstable when it's structure resembles so much to that of benzene which is exceptionally stable.

Benzene is not exceptionally stable to start with. As you substitute more and more CH groups with N in benzene, the enthalpy of formation (already positive for benzene) becomes larger and larger (e.g. 1,3,5 triazene). A positive enthalpy of formation tells you how much energy is released when the molecule turns back into elements. (It does not tell you about the activation energy needed for the process, so a molecule can exist even if a positive enthalpy of formation).

Playing around with the different unsubstituted species on MolCalc, it seems that the predicted enthapies of formation are higher when two nitrogen atoms are directly next to each other. In hexazine, there would be 6 of those pairs.

There are no examples of pentazine or hexazine. There are different tetrazines, but 1,2,3,4 tetrazine only exists fused to benzene (in the picture below, compounds (1), (4), and (5) have not been observed yet).

enter image description here Source: https://www.sciencedirect.com/science/article/pii/B9780080965185001386

Like the comments of the question already say, bonding in elemental carbon is very different from that in elemental nitrogen. Similarly, substituting N for CH in ethane yields hydrazine, a rocket fuel. So while the Lewis structure look the same, many properties change dramatically when you substitute N for CH.

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    $\begingroup$ Hydrazine isn't explosive... unless you mix it with oxidant, but so is ethane. There are plenty of nitrogen rich explosives, though, like tetrazene explosive $\endgroup$
    – Mithoron
    May 1, 2023 at 15:57
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    $\begingroup$ @Mithoron Point well taken, I removed that bit. Ethane has a negative enthalpy of formation, while hydrazine has a positive enthalpy of formation. Thus, the latter can be used as a monopropellant in thrusters operating in outer space, e.g. here. $\endgroup$
    – Karsten
    May 1, 2023 at 21:33
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An aromatic five-nitrogen ring is known as the pentazolate ion, $\ce{N_5^-}$. Among several documentations of this anion, which is isoelectronic with the cyclopentadienyl anion, the Wikipedia article identifies an isolated salt with moderately good thermal stability:

In 2017, white cubic crystals of the pentazolate salt, $\ce{(N5)6(H3O)3(NH4)4Cl}$ [the article and cited reference render the anion first in the formula] were announced. In this salt, the $\ce{N_5^-}$ rings are planar. The bond lengths in the ring are 1.309 Å, 1.310 Å, 1.310 Å, 1.324 Å, and 1.324 Å.[1] When heated, the salt is stable up to 117 °C, and over this temperature it decomposes to ammonium azide.[1]

Crystal packing presumably stabilizes the anion in the presence of protic species that would surely react with it in solution.

Cited Reference

  1. Zhang, Chong; Sun, Chengguo; Hu, Bingcheng; Yu, Chuanming; Lu, Ming (26 January 2017). "Synthesis and characterization of the pentazolate anion cyclo-N5ˉ in (N5)6(H3O)3(NH4)4Cl". Science. 355 (6323): 374–376. Bibcode:2017Sci...355..374Z. doi:10.1126/science.aah3840. PMID 28126812. S2CID 206651670.
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    $\begingroup$ Cool reference! From the editorial introduction: "The flip side of the robust stability of N2 is the instability of any larger molecules composed exclusively of nitrogen. These molecules nonetheless remain enticing targets for explosive and propellant applications." $\endgroup$
    – Karsten
    Mar 4, 2023 at 20:05
  • $\begingroup$ $\ce{N3- and N5+}$ also exist, see here. They aren't aromatic, though (not even a ring). $\endgroup$
    – Karsten
    Mar 4, 2023 at 20:09
  • $\begingroup$ @karsten the WP article says chemists trued to synthesize $\ce{N5^+N5^-}$ for use as rocket fuel. That apparently failed, but we do have $\ce{LiN5}$ synthesized wt high pressure. Might that work with an appropriate pentazenium compound? $\endgroup$ Mar 4, 2023 at 20:11
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    $\begingroup$ I can't tell, I'm not a rocket scientist. ;-) $\endgroup$
    – Karsten
    Mar 4, 2023 at 21:50

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