4 ms·

These alkyne (C≡C) bonds are indeed strained and unstable, but there is an energy barrier to jumping to alkane (C-C) or alkene (C=C). Otherwise, polyynes like I

by kortex 7y ago

These alkyne (C≡C) bonds are indeed strained and unstable, but there is an energy barrier to jumping to alkane (C-C) or alkene (C=C). Otherwise, polyynes like Ichthyothereol [1] would not be stable. 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. https://en.m.wikipedia.org/wiki/Ichthyothereol https://en.m.wikipedia.org/wiki/Ichthyothereol

tempguy9999 7y ago
OMG which part of "I Is A N00b" did I not make clear? :-) 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!
- logfromblammo 7y ago
  In order for a pi-bond to easily flip parallel to a different sigma-bond, there has to be some factor making the other end of the new bond more positive, or the other end of the existing bond more negative. And pi-bonds basically occupy two side slots. You can think of them as left&right for one slot and top&bottom for the other. 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.
  - tempguy9999 7y ago
    Oh my. I had no idea of "pi-bonds basically occupy two side slots" or "uses up both mutually-perpendicular slots for the side bonds". This geometry aspect is completely new. 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.
    - logfromblammo 7y ago
      Electrons around an atom occupy certain probability functions that correspond to the electron's energy level. The most basic and lowest energy of these at a given tier of orbital distance is nested spherical shells (s). Next is 3 different 2-lobed dumbbell shapes (p), oriented mutually perpendicularly (px, py, pz). 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"

by kortex 7y ago

tempguy9999 7y ago

OMG which part of "I Is A N00b" did I not make clear? :-) 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!

logfromblammo 7y ago
In order for a pi-bond to easily flip parallel to a different sigma-bond, there has to be some factor making the other end of the new bond more positive, or the other end of the existing bond more negative. And pi-bonds basically occupy two side slots. You can think of them as left&right for one slot and top&bottom for the other. 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.
- tempguy9999 7y ago
  Oh my. I had no idea of "pi-bonds basically occupy two side slots" or "uses up both mutually-perpendicular slots for the side bonds". This geometry aspect is completely new. 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.
  - logfromblammo 7y ago
    Electrons around an atom occupy certain probability functions that correspond to the electron's energy level. The most basic and lowest energy of these at a given tier of orbital distance is nested spherical shells (s). Next is 3 different 2-lobed dumbbell shapes (p), oriented mutually perpendicularly (px, py, pz). 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"