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November 10

Benzene ring

Many, many moons ago (I would have been about 15) our chemistry teacher told us about the benzene ring. But because my mathematics was abysmal, bordering on non-existent, I soon came to the unhappy conclusion that I was never going to realise my boyhood dream: to be some sort of chemist, preferably of the organic type. Even the concept of a mole did my head in. About a decade later (having pursued a somewhat different career path) I came across Kekulé's fabled constitutional wander, and began to understand how utterly important his insight was.

Can anyone explain in plain language why the benzine ring appears to be so fundamental to org. chem.? Almost every diagram of any complex molecule I come across on WP seems to include it almost as a matter of course. What makes it so bindable, as it were? Is there anything else quite so widespread?

Our teacher also suggested The Peptide Link™ as the name of a funky jazz beat combo, but it seems never to have taken off. MinorProphet (talk) 02:24, 10 November 2022 (UTC)[reply]

Articles: cyclic compound, aromatic compound, aromaticity. Very technically, a "benzene ring" is only C₆H₆. There are lots of compounds containing a ring structure that isn't a "benzene ring", although they can be modeled as starting from a benzene ring and then substituting atoms of the ring with other things. Aromatic rings, of which the benzene ring is an example, are very stable chemically because the electrons of the atoms in the ring "overlap" and are delocalized throughout the ring, in a great demonstration of quantum chemistry and the wave–particle duality of electrons. This confers increased stability, which is often a desirable property. Not all "common" molecules contain ring structures. For instance carbohydrates often can adopt a cyclic form, but can freely convert between this and a linear form. Fatty acids are linear molecules, and only some of the proteinogenic amino acids contain rings. --47.147.118.55 (talk) 05:22, 10 November 2022 (UTC)[reply]

Interestingly, the part of the article on the history of benzene's structure was improved greatly within the last month thanks to Leyo, who took the trouble to upload to Commons the original diagrams as published in the late 1800s. Our Talk Page discussion is WT:WikiProject Chemistry#Historic benzene structures. One reason organic chemists focus on benzene is because its structure is used to teach the general principles both of aromaticity and reaction mechanism, where many fall in love with explanations based on the curly arrow. The use of benzene as a part-structure in drugs and agrochemicals largely stems from intermediates containing it being readily available and cheap as a byproduct of petrochemical production. Mike Turnbull (talk) 12:18, 10 November 2022 (UTC)[reply]

The importance of benzene and its structure is the presence of the delocalized pi electrons; it was the first structure where such a thing was shown to exist; it explains a LOT of the properties of benzene, which are paradoxical. Benzene is highly unsaturated, meaning that it has excess electrons and thus in theory should easily make bonds to other atoms readily. Usually, unsaturated hydrocarbons are more reactive than saturated ones, because there is a local area of electron density in the molecule; basically things that have double bonds in them have a concentrated area of electron density that electrophiles can easily form bonds to. And yet benzene is essentially as unreactive as a saturated alkane. How can one have an unreactive molecule that is so unsaturated? The reason is that the "extra electrons" in benzene are delocalized and spread evenly across the ring structure, meaning there is no local area of high electron density; so nothing for electrophiles to "grab on to". Recognizing that this could happen was the key to opening up entire new areas of organic chemistry, we call this aromaticity, and such compounds are called aromatic compounds (the etymology of the terminology escapes me, but the use of the word aromatic in organic chemistry has basically nothing to do with odor). If you need me to go more into what "pi electrons" are I can, but really as long as you know that the "extra" electrons in the highly unsaturated benzene ring are not concentrated between any particular atoms in the ring, but are instead delocalized and spread evenly all over the molecule is the key to the situation. I hope that all works for you. --Jayron 32 13:15, 10 November 2022 (UTC)[reply]
The word "aromatic" does have its origin in the odor of such compounds. See Aromaticity#The term "aromatic". It was used in that sense by August Wilhelm von Hofmann in 1855, before the relevant structures were confirmed: and in some cases what he called aromatic turned out not to contain what we today recognise the word to imply. Mike Turnbull (talk) 13:51, 10 November 2022 (UTC)[reply]
Oh, it has its origin in that way. It has no current such meaning. The fact that something meant something in the past has no bearing on what it means in the present. See Etymological fallacy. When an organic chemist uses the term "aromatic" in the context of ring compounds with delocalized pi systems, they are not saying anything about their scent. --Jayron 32 15:08, 10 November 2022 (UTC)[reply]

To appreciate the importance of the benzene ring you should consider somewhat similar compound - cyclohexane. In both six carbon atoms make a ring but their properties are very different. Ruslik_Zero 19:19, 10 November 2022 (UTC)[reply]

Wow, what a lot of clear, helpful answers! I'm now much more aware of its importance, and why; and lots more to be reading. Thanks all for your expertise. MinorProphet (talk) 23:35, 10 November 2022 (UTC)[reply]

Jeez, this is wonderfully complex... It was finding File:Benzene-resonance-structures.svg which allowed me to conceptually grasp (I think) the dual quantum states of benzene and how the ring model is a sort of approximation of the states; and how quantum chemistry can explain it.

Having read and attempted to comprehend the articles recommended above, I've got a load more questions. Some of them may be plain wrong, but I would be grateful for any further straightforward answers.

How does the benzine resonance diagram above relate (if at all) to File:StationaryStatesAnimation.gif and the Schrödinger equation?
Which type of molecular orbital does benzene have?
Do all ring structures have delocalised electrons?
What importance does molecular symmetry have in forming pi bonds in benzene?
Is it only pi bonds that have de-localised electrons?
What effect do the conformations of cyclohexane have in the real world? Do other compounds exhibit this behaviour?

@MinorProphet: These are the sorts of questions that chemists on graduate courses need about a year of study to master! Try reading Woodward–Hoffmann rules next, which covers many of these topics but does get quite difficult to follow without the sort of background best found by reading a decent organic chemistry textbook. Mike Turnbull (talk) 18:11, 13 November 2022 (UTC)[reply]

To answer a few of these good questions, benzene has both sigma and pi molecular orbitals. As one approximation, the sigmas represent the "first" bond along each edge and are "in the plane" (in-line between the atomic nuclei), whereas the pi are the additional bonding electrons "above and below the plane"; overall, it's two halves of a bagel (pi) with an onion-ring (nuclei and sigma) as the sandwich filling. File:Benzene MO diagram.png illustrates how the p atomic orbitals of the carbon atoms align and combine in various ways to form the pi molecular orbitals.

The aromaticity is a pi phenomenon here, and resonance/delocalization in general is usually a pi and rarely a sigma situation. It's the alignment of more than two atomic orbitals that allows delocalization, not the molecular shape (ring) itself: allyl cation has delocalized electrons but cyclohexane does not. Aromaticity is a more complex situation that relates to delocalization that is around a ring: benzene has it, but cyclohexa-1,3-diene has delocalization but not aromaticity and cyclohexa-1,4-diene does not even have delocalization.

Conformations of cyclohexane are just one notable application of the ideas of conformational analysis, where pairs of bonded atoms can rotate relative to each other, leading to more- or less-stable geometries in those parts of the structure. Many biological functions are dependent upon molecular shape (the lock and key model, for example), so the ways a molecule can be shaped at all (flexibility) and their relative stability in different shapes are both important when structures interact with each other. DMacks (talk) 19:44, 13 November 2022 (UTC)[reply]

Thanks all for your great comments and explanations. Reading the above links makes me realise how much I would have liked to have understood the maths in order to study this sort of thing at uni. It makes so much sense of what I formerly had a mere vague grasp of. I did actually buy a 2nd-hand c. 1960s organic chemistry textbook after some interesting experiences with psilocybin, but again the maths and algebra defeated me. Ah well, it's been satisfying in other ways, less concerned with the scientific method and the essential matter of the universe.

So, although I have probably reached the limits of my understanding beyond these clear explanations, it's been very enlightening finding out how far I could get. Wouldn't it be possible to gather all the above info into a sub-section of benzene, headed ==Importance/significance of benzene==?

So, as a final unanswered question from above, please: How does the File:Benzene-resonance-structures.svg diagram relate (if at all) to File:StationaryStatesAnimation.gif and the Schrödinger equation? Many thanks, MinorProphet (talk) 17:48, 14 November 2022 (UTC)[reply]

The answer to that question is "not really at all". What we call "resonance structures" is an artifact of Lewis diagrams as a model, and does not represent anything all that profound about molecules. Molecules are just collections of nuclei and electrons arranged in a relatively stable geometry. When you think of a single atom, you can envision a nucleus at the center of an electron cloud, that represents the sort of "average" of all of the quantum states of the electrons spread out over the space around the nucleus. A molecule is similar, in that case you have multiple nuclei floating in a more complex-shaped electron cloud, the geometry of which is controlled by the geometry of molecular orbitals. Lewis diagrams are useful for simple molecules that feature simple bonding, where "bonds" of two electrons each are represented by lines. However, the molecular electron cloud is more complex than this; what a line represents is not a rigid, stationary pair of electrons staying still between the atoms, but rather represents about 2 electrons worth of negative charge acting to hold the respective nuclei together; given that electrons are essentially not standing still in that space, that 2-electrons-worth-of-charge is an average of all of the motions of those electrons (ish... it's a bit messier than that, but it works well enough for this explanation). Since it's just an average anyways, there's nothing that says it has to be exactly 2 (or a multiple of 2) electrons. You can have 1 electron's worth of charge, or even a fraction of an electrons worth of charge holding the nuclei together in a molecule. There is nothing wrong with that from a physics point of view. The problem is that Lewis diagrams are a rather blunt tool. How do you represent a bond consisting of 3 electrons worth of charge with them? What about 2.67 electrons worth of charge? I mean, all I have are dots and lines to do that. What we call "resonance structures" are just a kludge to make Lewis diagrams work with more complex electronic situations. Simply put: resonances doesn't necessarily represent anything profound about molecules; which are happily doing their own thing with no particular problems; they're just a way so that our diagrams and models can "fit" what is actually going on. One of the impetuses of molecular orbital theory was to abandon the overly simplistic Lewis theory and develop a more robust model; the model often lacked a nice little picture we can write with letters and lines, but in abandoning that we developed a more predictive and robust model of molecules. We still use Lewis diagrams today because they are convenient, but they kind of fall apart when we get to anything beyond simple 2-electron bonds. --Jayron 32 13:33, 15 November 2022 (UTC)[reply]

@Jayron32: Your explanations seem so clear and free of jargon, thank you. I think I get the idea. The more you know, the deeper it gets... I wonder if you (or someone) could please answer some tangential questions: <What, more?>

What are the solid black thin triangular pointers in eg Arrow pushing#E1 eliminations?
If people are still talking about valences as a valid way of understanding atomic theory, how much of all this theory on bonds might Kekulé, working on Edward Frankland's idea, have actually understood?
What needed to happen to mathematics/physics between Kekulé's 1865 paper and, say, Einstein's of 1905, to to come up with quantum chemistry? How far is Aleph null, my favourite infinity, involved?
Who came up with the concept of bond lengths and angles between bonds?

I'm very grateful for everyone's interest and expertise in explaining unbelievably complex ideas in straightforward terms. MinorProphet (talk) 16:04, 15 November 2022 (UTC)[reply]

Sure thing! Glad to be of service! Let me do my best.

For the triangular bits, see skeletal formula for more details. In organic chemistry, a highly condensed form of Lewis diagrams is shown, omitting carbons and hydrogens from drawings. In an organic skeletal formula, each end or vertex is considered to be a carbon atom (plus enough H atoms to give it 4 bonds). In order to represent three dimensions on the page, we use the "wedge and dash" method. Any bond that comes "towards" the viewer ("out of the page") is written as a wedge (what you call "solid black thin triangular pointers") and any bond going "away" from the viewer ("into the page") is written as a series of dashes. See also here for more details.
August Kekulé was really on the "tip of the spear" for developing what we know more formally call valence bond theory, which Gilbert N. Lewis put together in his landmark textbook on the subject in 1916, but Lewis didn't really invent it out of whole cloth; these ideas had been kicking around Chemistry for decades and decades, it really can't be credited to any one person, August Kekulé both developed some of the early ideas on valence and also expanded on ideas developed earlier. Really, valence can be traced all the way back to things like the Law of definite proportions and the Law of multiple proportions and now we're getting back to Proust, Dalton, and Lavoisier; back to the real founding fathers of Chemistry. The notion of valence had already existed when fellow Russian Dmitri Mendeleev incorporated it into his periodic law, working concurrently with Kekulé. Valence (chemistry) discusses some of this in the "Historical development" section.
So, the thing about Einstein is (and I am NOT saying this to diminish him, he's really deserving of all the credit he gets!), is that he's just got a better PR machine than other equally (and in some fields more) important scientists working at the time. Einstein's contributions to quantum theory are really limited; in his "Annus Mirabilis" papers of 1905, it was his application of quantum theory to the photoelectric effect. In terms of other contributions, it was mostly that he was a pretty good foil in trying to poke holes in quantum theory (his famous "God does not play dice" criticism of QM for example, and the EPR Paradox as another), except as far as we can tell, he was usually wrong in his feelings about quantum mechanics. If we really want to trace QM from its roots, it starts with Max Planck, who invented QM as a literal mathematical kludge to fix the ultraviolet catastrophe, a deviation from experiment caused by the Rayleigh–Jeans law in trying to predict the results of blackbody radiation. I don't even think that Planck himself at first believed it was true, he merely created it as a mathematical tool to help fix the problem. From a chemist's point of view, the real genesis of QM in chemistry starts with the Bohr model of the atom developed by Neils Bohr (building on the Rydberg formula and several earlier models explained in the intro to the Bohr model article). Modern quantum theory really starts with Erwin Schrödinger and the application of wave dynamics to quantum theory; the article Old quantum theory will answer your question of how we got there from it's earliest beginnings.
I'm not sure what you mean "came up with". They are actually experimentally measurable things. Frequently, things like bond lengths and angles can be obtained via a process known as X-ray crystallography, so if we want to blame anyone for how we came up with this knowledge, it was the two Williams Bragg, the father-son team that came up with Bragg's law, which allows us to extract crystal shapes (and thus bond lengths and bond angles) from the scattering of X-rays off of solid crystals. Later techniques such as Raman spectroscopy and Gas electron diffraction have been used, among many others. In terms of predicting bond angles and lengths, that comes down to VSEPR theory, which provided the predictions that experimental techniques either confirmed or refined. So, if we have to blame anyone specific for first developing notions of bond length and bond angles and the like, it would be Ronald Gillespie and Ronald Sydney Nyholm for developing VSEPR theory, and William Henry Bragg and William Lawrence Bragg for Bragg's law and X-ray crystallography. And the experimental techniques predated the theoretical by some decades. --Jayron 32 16:42, 15 November 2022 (UTC)[reply]

I think your description of K as being "on the tip of the spear" is entirely appropriate. It's a pity that WP articles tend to shy away from assigning extraordinary greatness to individuals - there appears to be little to choose between the biography of, say, an Estonian footballer and eg Lavoisier (whose portrait appears at the head of Annales de chimie - Ext. links section partly mea culpa). The exquisite fertility of the minds of the founders of chemistry appears to confound our petty existences - most of us can only follow in their wake and murmur, "Yes, boss." Even more inviting reading material - many thanks once again. MinorProphet (talk) 21:13, 15 November 2022 (UTC)[reply]