Chemists make first-ever ring of pure carbon
nature.com
nature.com
The molecule (an all-carbon cycle of 18 atoms) was prepared in a very unusual way - by directly manipulating the atoms using an atomic force microscope.
In other words, each molecule is made individually. This is not the way that chemists typically work, and will not result in quantities of material that can be seen.
The abstract says nothing about chemical stability, but I suspect C-18 quite unstable and may never be prepared in gram quantities.
Higher cycles containing more carbons may be more stable, but this is likely to remain a curiosity for some time.
Still, this is a new kind of "allotrope" of carbon. The cyclic, relatively rigid nature of the structure and the potential for electrons to circulate under applied fields could lead to some unusual applications.
I have been a chemist, and also played with explosives as a child. So it's my experience that C≡C bonds are very unstable. I mean, liquid HC≡CH (acetylene) must be diluted with acetone, and stored in tanks packed with diatomaceous earth, to keep it from falling apart explosively. A lot like nitroglycerine, really.
And AgC≡CAg (silver acetylide, with the Ag-C bonds being almost ionic) is a damn fine primary detonator. They used to use it in party poppers and cigarette/cigar "loads". But you could use it instead of lead azide or mercury fulminate, and it's much easier to make. But it's also much more expensive.
So anyway, I can't imagine that you'd want gram quantities of this stuff. Even if you could prepare it.
Edit: A gram of silver acetylide would punch a hole through 15-20 gauge mild steel. Uncontained. Just sitting there, in a little pile.
Depends how many graduate students they can hire.
Is there not merit in doing it via this 'hard way' first before optimizing stuff and finding a production pathway? (IANAchemist)
A milllimole (10^20 moleculess) is considered small scale by many chemists. Working with a single molecule as described in the artice is just short of science fiction, and hardly any chemists have ever done it. It requires one of the most expensive instruments in the world.
Wrong url?
(For anyone else wondering, parent quoted just the paragraph >"Gawel... says Scriven", then added their thoughts.)
The molecules where the bonds have wrong angles usually have more energy and they end to decompose in other molecules where the bonds have the correct angles or they have other bonds. [Oversimplifying warning, Chemistry is more complicated.]
Some nice graphics: https://www.sciencedirect.com/topics/chemistry/triple-bond Perhaps the text is too technical, but the graphics are nice.
But upon further consideration , the potential energy surface may be more like a saddle point, with each atom having a stable local minimum in-plane but unstable local maxima perpendicular to the plane. (I’m not a chemist).
I studied a little of Chemistry, but this is out of the scope that I know well. IIUC correctly from https://en.wikipedia.org/wiki/Cubane the stability of cubane is not due to symmetry. The stability comes from the fact that locally it looks like a alkane, i.e. a molecule with Carbon and Hydrogen that only have single bonds. They are quite stable, the most common example is gasoline that is mostly alkanes.
Double bonds are more reactive. For example vegetable oils (that have also Oxygen, not only Carbon and Hydrogen) may have or not have double bonds. The one with many double bonds become rancid easier. https://en.wikipedia.org/wiki/Rancidification
I have a gut feeling that in this case it's better if the number of carbons is like 4n+2 (i.e. even, but not a multiple of 4, just like 18 :) ). This rule is important when you have a similar structure with double bonds and Hydrogen. https://en.wikipedia.org/wiki/Aromaticity I'm not sure if it translates to triple bonds.
(Some handwaving: Here the pi bonds in the plane act like the H in the aromatic compounds, and the pi bonds perpendicular to the plane form an aromatic system.) (Note that handwaving is never a replacement of knowledge, so this many be very bad.)
Edit: Expanded a bit
Nit: Article confusingly uses name carbon-18 to denote a molecule with 18 atoms of carbon. Carbon-18 suggests otherwise an isotope of carbon.
”Carbon (6C) has 15 known isotopes, from ⁸C to ²²C, of which ¹²C and ¹³C are stable”
⇒ ¹⁸C exists (albeit with a half-life of about 92ms)
I'm not very familiar with chemistry. Are there other instances where other molecules can interlink in a way that, while not having bonds between them cannot be separated without breaking the bonds of one or the other?
At the very least it's fun to imagine a form of ringmail made of these.
For some interesting reading in self-assembling structures see: Hao Li et al. Quantitative self-assembly of a purely organic three-dimensional catenane in water, Nature Chemistry (2015). DOI: 10.1038/nchem.2392
In order to be stable, they have to be capped with a less-reactive group than hydrogen. If you break a ring, and get a C18 with naked electrons instead of protons on the ends, that's going to react with something, quickly and probably energetically.
That bond-bending energy has to go somewhere, too. If you break the ring, you're probably going to break the molecule near the antipode, too, giving you two highly reactive C9s, or a C10 and a C8. Like if you cut a bowstring at full draw, you're probably going to break one of the limbs of the bow.
> ...showed that the 18-carbon rings had alternating triple and single bonds.
Stable alternating single/triple bonds? As in the bonds don't flicker around? How is that even possible?
1. Surely the bonds can't stay in the same place
2. Surely they would 'nebulise'/spread out/whatever the right term is, over the ring
3. if 2. doesn't apply, why on earth is it more energetically favourable to have a very high energy (ok, my assumption) triple bond next to a much lower energy (ditto) single bond, and for it not to immediately snap into a pair of double bonds?
Given my n00b level of teh chemistry, what is going on here?
Resonance is not possible with this configuration because the orbitals are not symmetric. Benzene is often drawn with alternating alkane and -ene bonds, ⌬. It's not obvious, but benzene is radially symmetric in its electronic configuration. Each carbon has p-orbitals pointing vertically from the disc, forming a cloud, rather than discrete pi bonds between neighbors, and the electrons can "skate around the ring".
The first bond between carbons is always sigma. In alkenes, there is one sigma bond, in -ynes, there are two pi's at 90 degrees to each other. I'm a bit fuzzy on my o-chem, but iirc there is an energy gap between this configuration and cumulene C=C=C, which makes spontaneous transition unlikely.
In this c18 cycloalkyne, the p-orbitals are further strained into a bent cross shape if looking toroidally. This makes for an even bigger difference in energy levels, which means it can't flip to alkene or even swap positions.
The role of spdf and sigma/pi electrons, I'm aware of these things but don't have a mental model for them. And "p-orbitals pointing vertically from the disc"??.
That was brutal but you've given me a lot to aim for so I'll get to reading this afternoon - thanks!
Aromaticity is basically alternating single and double bonds. So when one pi-bond flips to the next bond down the chain (preferably into the same slot, but not necessarily--molecules can twist and flex), you get two single bonds in a row on one side, and two double bonds in a row on the other. Ordinarily, this would be a strong incentive to flip back to the way it was before, but if a double bond further down the chain can be flipped, it might propagate further down. In a ring, like benzene or a phenyl group, that wave of bond-flipping can just chase its own tail forever. In a polymer, it can go all the way down to the end of the (long) molecule before bouncing back.
When you have alternating single and triple bonds, that triple bond uses up both mutually-perpendicular slots for the side bonds. No other pi-bond can flip over next to the adjacent sigma-bond, because the bonds on the other side just bounce it right back, like having Manute Bol and Hakeem Olajuwon on either side of the paint when the second forward of your high school junior varsity basketball team tries to go around to one side or the other to score a new pi-bond. It's just not going to happen, unless you pump Jeremy Cohen full of so much cocaine, amphetamines, and sugar that he can charge into a gigantic center and somehow bounce the defender into the top row of the bleachers.
I'll have a goggle but if you have any introductory resources for this, that would be very helpful. Books, sites, appropriate terms to search for I may not know, anything for a beginner just to get started, I'd appreciate it. Thanks.
The spherical mode (s) and the dumbbell-shaped modes (px, py, pz) can also hybridize when each has a single electron in it. Those hybrid modes can point lobes out in different directions. A stable mode can hold up to 2 electrons, each spinning in opposite directions.
When an electron probability function of one atom overlaps with that of an adjacent atom, their functions can further hybridize into a bonding function. (They can also hybridize into anti-bonding functions, but ignore that for now.)
When the bonding function puts the electron pair mostly along the line connecting the centers of the bound atoms, that's a sigma bond. It's the lowest-energy kind of bond, so it always (aka exceedingly rare exceptions) happens before any other kind of bond.
In a carbon triple bond, the spherical orbital (s) will hybridize with just one of the dumbbell orbitals (p), forming one function (spx+) that mostly occupies one lobe of the former dumbbell, and one that mostly occupies the other (spx-). These two will then form direct sigma bonds with atoms on either side, separated by 180 degrees. That leaves the other two dumbbell modes (py, pz), mutually perpendicular to the two hybrids. Those can align with the parallel dumbbell modes of the adjacent atom and create bonding functions where the electron positions sort of average out to between the two bound atoms, but mostly occupy places that do not overlap with the sigma bond. Those are pi bonds. They happen between two aligned p orbitals. (There is another kind of indirect bond, a delta bond, that happens between two aligned d orbitals, which have more complex shapes. Those would be the fourth bond in a quadruple bond.)
When those carbons are triple-bonded, their p-orbital atoms are leaning toward the adjacent bound atom. It's much harder for the atom on the opposite side to get their attention to maybe participate in a pi bond on that side.
If you look at the bond through a microscope, you're just going to see a big blob of electron probability between the nuclei, with 6 electrons somewhere in it most of the time, where the triple bond is. It's not organized into discrete spaces for each separate bond. It only has separate slots when you look at the mathematical model.
I don't have any specific resources, other than any modern chemistry textbook you can find. Possible search terms "electron orbitals", "bonding orbitals"
Can someone 'in the know' on semiconductor device fabrication expand on this more?
When the Superconducting Super Collider was proposed, the newspapers said that research done at the SSC would lead to breakthroughs in computer graphics and flat panel displays.
> Initial studies of the properties of the molecule, called a cyclocarbon, suggest that it acts as a semiconductor.
One of the amazing things about carbon is that it can lots of interesting things (conductor, insulator, semi-conductor, Etc) with no added elements, just different bonding and structures. Further, it is an excellent thermal conductor. As a result it has the potential to be the material that replaces silicon in computer chips on a very large scale.
...for the purposes of scrubbing CO2 from the air? There's absolutely no reason we'd use this as the final product. Graphite or diamond would do the trick just as well, and either is cheap to make.
The practical reality is that each transformation step has imperfect yeilds and I would guess that the equipment required to do the molecule by molecule transformation in the last step would require significant energy to run.
So to attempt to this today, you would end up producing more CO2 into the atmosphere than you removed.
Also, alternating single and double bonds leave a bonding site on each carbon for a functional group. Alternating single and triple bonds use up all four (easy) carbon bonds, and you can't cram another bond in there without an insanely powerful acid.
Maybe 18 carbons is the minimum (C≡C-C)ₙ, because smaller rings would be too strained. But larger rings would probably be more stable.