What's up with Polyyne?

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Jorpho
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What's up with Polyyne?

Postby Jorpho » Wed Dec 03, 2008 3:18 am UTC

Another item in the list-of-cool-science-things-I've-heard-of-that-seem-to-have-gone-nowhere is polyynes. Remember those? They're long chains of carbon atoms with alternating single and triple bonds, and they're supposed to have all kinds of nifty structural and conductive properties.

Is this stuff just more nano-hype, or is there really something to them?
http://en.wikipedia.org/wiki/Polyyne

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Re: What's up with Polyyne?

Postby Carnildo » Wed Dec 03, 2008 4:23 am UTC

Jorpho wrote:Another item in the list-of-cool-science-things-I've-heard-of-that-seem-to-have-gone-nowhere is polyynes. Remember those? They're long chains of carbon atoms with alternating single and triple bonds, and they're supposed to have all kinds of nifty structural and conductive properties.

Is this stuff just more nano-hype, or is there really something to them?
http://en.wikipedia.org/wiki/Polyyne


Nifty, yes, but long-chain polyynes are also explosively unstable.

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Jorpho
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Re: What's up with Polyyne?

Postby Jorpho » Wed Dec 03, 2008 4:35 am UTC

Aha. There's one little factor that never came up. Why exactly is that so?

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danpilon54
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Re: What's up with Polyyne?

Postby danpilon54 » Wed Dec 03, 2008 4:38 am UTC

why is it explosive? Probably because you have lots of carbons with very little oxygens. Carbon likes oxygen. Same reason gas burns.
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Re: What's up with Polyyne?

Postby Carnildo » Wed Dec 03, 2008 5:22 am UTC

Jorpho wrote:Aha. There's one little factor that never came up. Why exactly is that so?


The triple bonds. There's a great deal of potential energy bound up in a carbon-carbon triple bond, and breaking it has a fairly low activation energy. Once the triple bonds start breaking, it's not going to stop. It's what makes acetylene hazardous to handle, and the longer-chain versions are even worse.

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Re: What's up with Polyyne?

Postby tantalum » Wed Dec 03, 2008 5:51 am UTC

Are these molecules actually unstable? Wikipedia doesn't mention anything about instability, and wikipedia is god.

In a weird way, I think it's possible that a hunk of polyyne would probably be very rigid because of the occasional 2+2+2 cycloaddition reactions fusing chains together. In that way it would start to resemble graphite, although nowhere near as stable.

In terms of conducting organic polymers, we already have some good stuff. See the Jensen lab at MIT for some interesting polymers (polyphenylenevinylenes)

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Re: What's up with Polyyne?

Postby Jorpho » Wed Dec 03, 2008 6:08 am UTC

danpilon54 wrote:why is it explosive? Probably because you have lots of carbons with very little oxygens. Carbon likes oxygen. Same reason gas burns.
But by that reasoning graphite and diamond ought to be explosive too.

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Re: What's up with Polyyne?

Postby Mr_Rose » Wed Dec 03, 2008 9:19 am UTC

Jorpho wrote:
danpilon54 wrote:why is it explosive? Probably because you have lots of carbons with very little oxygens. Carbon likes oxygen. Same reason gas burns.
But by that reasoning graphite and diamond ought to be explosive too.

Carbon-carbon single bonds are much more stable than triple bonds, which is why single-bond isomorphs like graphite and diamond only burn rather than exploding. Same reaction, just slower....
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Re: What's up with Polyyne?

Postby BeerBottle » Wed Dec 03, 2008 11:36 am UTC

The fun thing to do is attach a transition metal complex to each end of a polyyne: you can get all sorts of cool structures depending on whether you have an even or an odd number of carbons, and how exactly you attach them to the metals. Say you place exactly the same metal complex on each end of a symmetrical polyyne. Now, the redox potential of the two metal complexes (essentially the energy required to add/remove an electron to/from the metal) will be different, i.e. one end will be oxidised first and this change will be communicated through the chain to the other end, making oxidation there slightly harder. Now this effect falls off with increasing chain length, but the good part is that even at very long chain lengths (extrapolating to infinity), there is still a residual effect - i.e. the metal centres can communicate over long (in molecular terms) distances. Cue hype about molecular electronics.

But even if molecular electronics never becomes a reality, it's still interesting science.

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Re: What's up with Polyyne?

Postby meat.paste » Wed Dec 03, 2008 3:24 pm UTC

I would have guessed that polyyne would get some stability from pi bond delocalization. Each carbon has 2 p orbitals for bonding in this molecule. Would a state where each carbon is double bonded to its neighbor be feasible? It's been too long since I looked at molecular orbital theory.

In any case, you have a molecule that has a large, negative delta-G when converted to oxide, and a low activation energy (breaking a delocalized pi bond?).
Huh? What?

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Re: What's up with Polyyne?

Postby Jorpho » Wed Dec 03, 2008 7:59 pm UTC

meat.paste wrote:Would a state where each carbon is double bonded to its neighbor be feasible? It's been too long since I looked at molecular orbital theory.
Apparently that's cumulene.

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Re: What's up with Polyyne?

Postby AFedchuck » Wed Dec 03, 2008 9:26 pm UTC

meat.paste wrote:I would have guessed that polyyne would get some stability from pi bond delocalization. Each carbon has 2 p orbitals for bonding in this molecule. Would a state where each carbon is double bonded to its neighbor be feasible? It's been too long since I looked at molecular orbital theory.

In any case, you have a molecule that has a large, negative delta-G when converted to oxide, and a low activation energy (breaking a delocalized pi bond?).

Delocalisation is disrupted because it's 1D. Look up Peierls distortion

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Re: What's up with Polyyne?

Postby meat.paste » Wed Dec 03, 2008 10:30 pm UTC

AFedchuck wrote:Delocalisation is disrupted because it's 1D. Look up Peierls distortion


I can check off "learn something" from my to-do list today. Thanks.
Huh? What?

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Re: What's up with Polyyne?

Postby Sir_Elderberry » Thu Dec 04, 2008 12:24 am UTC

Mr_Rose wrote:
Jorpho wrote:
danpilon54 wrote:why is it explosive? Probably because you have lots of carbons with very little oxygens. Carbon likes oxygen. Same reason gas burns.
But by that reasoning graphite and diamond ought to be explosive too.

Carbon-carbon single bonds are much more stable than triple bonds, which is why single-bond isomorphs like graphite and diamond only burn rather than exploding. Same reaction, just slower....


Wait, diamond burns? I'd feel terrible/awesome/wasteful/rich burning a diamond. Sounds like a fun experiment.
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Re: What's up with Polyyne?

Postby ducksan » Thu Dec 04, 2008 12:27 am UTC

Carbon subnitride is fun. C4N2, also known as dicyanoacetylene or butynedinitrile. The heat of formation is about 500 kJ/mol.
It may be the hottest-burning compound known.
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Re: What's up with Polyyne?

Postby Mr_Rose » Thu Dec 04, 2008 9:58 am UTC

Sir_Elderberry wrote:
Mr_Rose wrote:
Jorpho wrote:
danpilon54 wrote:why is it explosive? Probably because you have lots of carbons with very little oxygens. Carbon likes oxygen. Same reason gas burns.
But by that reasoning graphite and diamond ought to be explosive too.

Carbon-carbon single bonds are much more stable than triple bonds, which is why single-bond isomorphs like graphite and diamond only burn rather than exploding. Same reaction, just slower....

Wait, diamond burns? I'd feel terrible/awesome/wasteful/rich burning a diamond. Sounds like a fun experiment.

Yup. Soak it in oxygen and light it up and it will burn like really expensive coal. IIRC, this was demonstrated to the public a few times at some of the popular exhibitions of the nineteenth century.
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Re: What's up with Polyyne?

Postby Diadem » Thu Dec 04, 2008 12:38 pm UTC

Mr_Rose wrote:
Sir_Elderberry wrote:
Mr_Rose wrote:
Jorpho wrote:
danpilon54 wrote:why is it explosive? Probably because you have lots of carbons with very little oxygens. Carbon likes oxygen. Same reason gas burns.
But by that reasoning graphite and diamond ought to be explosive too.

Carbon-carbon single bonds are much more stable than triple bonds, which is why single-bond isomorphs like graphite and diamond only burn rather than exploding. Same reaction, just slower....

Wait, diamond burns? I'd feel terrible/awesome/wasteful/rich burning a diamond. Sounds like a fun experiment.

Yup. Soak it in oxygen and light it up and it will burn like really expensive coal. IIRC, this was demonstrated to the public a few times at some of the popular exhibitions of the nineteenth century.

Wow i never knew that. That is so cool.

*Wants to get rich just so he can burn diamonds*
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Re: What's up with Polyyne?

Postby ducksan » Sun Dec 07, 2008 3:16 am UTC

To elaborate on alkyne instability: The carbon-carbon single bond is about 350 kJ/mol, compared with 850 for the triple bond.
We all know nitrogen gas to be very stable based on its triple bond (nitrogen-containing explosives get a lot of their power from the energy released when N2 is formed.) The difference between nitrogen and carbon is that only the latter burns. Activation energy is about 1050-1200 kJ/mol to break three carbon-carbon or carbon-hydrogen bonds; this is SIGNIFICANTLY more than the 850 you'd need to break the
C=C attraction. The pi bonds only contribute a rather weak ~250 kJ/mol each. (Adding even small amounts, e.g. 10 kJ/mol to the required activation energy at constant temperature, by comparison, will really fuck with the reaction rate. Adding that same amount to the delta-G value has a very powerful impact on equilibrium.)
Less energy needed to break the bonds + same amount of energy released during combustion for each carbon/hydrogen burned = alkynes (and alkenes, somewhat) are more energetic compounds than you might expect. Most flammable gases can be used as improvised explosives when blended with oxygen, but acetylene really takes the cake.
Though generally regarded as stable, the nitriles (cyanides) must release a good deal of energy when combusted as well. Cf. cyanogen and carbon subnitride as exotic examples.

Polyynes might be some of the most reactive hydrocarbons I know of. Very high ring strain or steric hindrance are other major causes of reactivity. (The triphenylmethyl radical is stable. And yellow. The dimer, which isn't hexaphenylethane, has an almost nonexistent C-C bond due to steric effects, though the BE value I've seen might've been the Ph3C-CPh3 value in Ph6Et, which would explain a lot. It coexists with its monomer - nitrogen dioxide, another radical, is somewhat resistant to dimerization for related reasons.)

Finally, in regards to the above comment about metal complexes stuck onto polyynes - the effect that makes one harder to oxidize is analogous to dicarboxylic acids. Oxalic is almost a strong acid for the first proton because the carboxylate groups are close to each other. Something like decanedioic acid has a pKa1 closer to 4-5.
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Re: What's up with Polyyne?

Postby Jorpho » Sun Dec 07, 2008 4:47 am UTC

How very enlightening. I never quite understood why unsaturated fats were supposed to be more digestible than saturated fats when the double bond was supposed to be stronger. But now I do. (I also never new you could represent a triple bond with an underlined equal sign!)

FYI, the major reference I first encountered on the subject was McMurry's "Organic Chemistry", 4th edition (1996), which references a 1995 report by Richard Lagow at Rice. (I'm guessing it might not appear in subsequent editions simply because it was topical at the time.)

EDIT:
Very high ring strain or steric hindrance are other major causes of reactivity.
Why exactly would there be ring strain or steric hinderance?


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