While silicon dioxide is chemically described as SiO2, you never actually encounter any in quite that form. You always encounter huge conglomerations of SiO2 molecules, linked together in extensive chemical 3-dimensional matrixes. This kind of linkage, for some simple single molecule like SiO2, is known as "polymerized". Each individual molecule is called a "monomer"; ordinary quartz (pure SiO2, remember) is a polymer. All substances known as "silicates" are also polymers, thanks to seemingly endless linkages of silicon-to-oxygen-to-silicon-to-oxygen-to-silicon-...
Okay, it is widely known that carbon and silicon are "related" atoms with similar chemical properties. Almost all the more widely-known polymers (polyethylene, polyester, polyvinyl chloride or PVC, polyurethane, etc.) are in fact based on chains of carbon atoms. Silicones, in fact, are silicon polymers that were designed BECAUSE of the fact that carbon molecules form polymeric chains so easily -- and silicon is chemically similar to carbon.
So...if polymerized silicon dioxide is so common, why don't we ever hear of polymerized carbon dioxide? When I first thought of this idea, a bunch of years ago, I could hardly believe that no one else had thought of it, but I couldn't find any references (which of course doesn't mean that there aren't any). I suppose, even today, that the stuff, if any was ever made, would prove to be unstable. Carbon dioxide may simply "prefer", energetically speaking, to exist in the form of individual molecules. (This sort of thing is well-known to be true for nitrogen gas molecules, N2. Most chemical explosives depend partly on the fact that if given half a chance, nitrogen will break loose from certain molecules, such as potassium nitrate, in order to recreate N2 molecules. That initial break-up leads to the release of other molecules, such as oxygen, which can fan the flames of the main explosive event.)
On the other hand, perhaps polymerized CO2 is stable, but requires such special conditions to create that you never find any in Nature. If this is true, what might its physical properties be? Well, since we are talking about a perfect analogue of quartz, it should be able to exist in a crystalline form. This is a remarkable form, because it is the same form that pure crystalline carbon takes -- otherwise known as diamond. And we all know that diamond is a remarkable material.
Pure crystalline silicon will also possess the same structure as diamond, but it is nowhere near as hard a substance as diamond. Why? Two reasons: One, the silicon-silicon bond is much weaker than the carbon-carbon bond; Two, carbon is a significantly smaller-than-average atom, and there are vastly more carbon-carbon bonds per unit of volume in a diamond, than than bonds-per-volume in any other substance.
Now consider quartz: It is harder than crystalline silicon, but not as hard as a diamond. This is because the silicon-oxygen bond is quite a strong chemical bond (stronger, if I recall right, than even the carbon-carbon bond). However, there are still a lot fewer chemical bonds per unit of volume in quartz than in diamond, so diamond remains the champion in hardness.
There is another substance, rather harder than quartz, known as "silicon carbide", which also possesses the diamond structure, and is among the hardest of all substances -- except for diamond. Due to the size of the silicon atom, there are again fewer carbon-silicon bonds per unit of volume than carbon-carbon bonds in diamond. (They are also not quite as strong as the carbon-carbon bond, if I recall right.)
With the background in place, we can now consider the suggested polymerized carbon dioxide. The carbon-oxygen bond is a VERY strong bond, significantly stronger than the carbon-carbon bond in diamond. However, oxygen atoms are larger than carbon atoms, and so there will again be fewer bonds per unit of volume than in diamond. Will the bond-strength make up the difference? Will poly-CO2 be as hard as, or even harder than diamond? THAT might explain why we never see any of these crystals in Nature! We all know what Nature has to do to make a diamond, and they are pretty rare....
Not to mention that ordinary CO2 is a gas, while ordinary pure carbon is a solid (graphite, for example). HOW in Nature will it collect, pressurize, and polymerize carbon dioxide???
However, we humans collect and pressurize pure CO2 everywhere we make soda-pop. One might think that if we added a bunch of heat, and significantly more pressure (and maybe some electrical discharges or blasts of laser light at the right frequency), we might find ourselves possessed of poly-CO2.
Will it be stable? I dunno. A friend of mine at work told me that he once saw something on the Web about some other dude who had made some, and it decomposed and destroyed his laboratory, but I know that this friend often exaggerates for entertainment value. He might have made the story up. Certainly I haven't been able to find any reference to the alleged incident.
IF it is stable, though, then simply because it will probably be relatively easier to make than diamond, we could find a lot of uses for it. For example, it is known that ordinary glass gets STRONGER under pressure (up to the limit where it collapses, of course); some submersibles are simply two glass hemispheres and a rubber seal. Poly-CO2 should be stronger yet. This is, in fact, the stuff of which vacuum balloons might be made....-- Vernon, Sep 25 2001 Crystalline Structure of Carbon Dioxide Seen for the First Time http://www.ars.usda...r/1998/980223.b.htm [egnor, Sep 25 2001, last modified Oct 04 2004] Nitrogen crystal http://www.sciencen...s/20040717/fob4.aspI wish they had compared the energy stored in this stuff to other explosives. [bungston, Oct 04 2004] The many faces of carbon dioxide http://www.llnl.gov/str/Yoo.html"a team led by physicist Choong-Shik Yoo has verified two forms of solid CO2 never before seen in the laboratory", a polymer form. [nomel, Oct 04 2004] Silicon based life http://www.daviddar.../S/siliconlife.htmlmentions co2 polymers. [nomel, Oct 04 2004] Carbon Dioxide phase diagram http://www.science....c123/phasesdgm.htmlAbout half way down the page. [shapu, Oct 04 2004] Apparently Now Discovered! http://news.bbc.co..../nature/5083222.stmGlass-like, though, not a pure crystal, so to precisely match this Idea, they have a ways to go. [Vernon, Jun 19 2006] Molecular orbital theory http://www.chm.davi...Orbitals/index.htmlFor [Vernon] - how bonding actually works once you get past the lies of school and college [K o R, Sep 17 2006] Shared e˜ pairs? Consider the similarity between Si and S (similarly amorphous).-- reensure, Sep 25 2001 PeterSealy, I am quite certain that even though it is a crystal, SiO2 is ALSO a polymer. In most crystals the molecules merely fit together neatly, and are held in place by van der Waals forces. In SiO2, silicon carbide, and diamond, there are actual chemical bonds connecting to the neighbors. If you could actually make and accumulate pure monomolecular SiO2, it would probably be a fairly ordinary liquid at room temperature, and possibly even a gas.
egnor, hmmm, it does look more like a zero than an oh...(fixed).-- Vernon, Sep 25 2001 UnaBubba, yes, that is true. However, whenever you cool down that gas back to ordinary human temperatures, it spontaneously re-polymerizes. This happens because AS the molecules cool, they collide with less energy...but still hard enough to cause the chemical bonds to break/flow/merge. To obtain monomolecular SiO2 at room temperature, that process has to be prevented somehow. Perhaps extremely rapid cooling with liquid helium spray...to cool them before they can collide with each other, so that when they DO collide, they have too little energy to merge.
egnor, that link to ars.usda... seems to be defective.-- Vernon, Sep 25 2001 It ([egnor]'s link) works for me.-- angel, Sep 26 2001 Wahey! Vernon's back!-- lewisgirl, Sep 26 2001 Well, today egnor's link worked for me, too. Previously my browser kept giving me error messages (two different browsers on two different computers).
Anyway, the destination of the link pretty clearly describes ordinary frozen carbon dioxide, in which the molecules are (as I described in prior annotation) merely neatly fitting together, and held by van der Waals forces. Not like the proposed Poly-CO2 at all, which would have chemical bonds holding all the atoms together, into a single regular structure (just like quartz).-- Vernon, Sep 26 2001 Thank you, RonHamel!
However, it appears that I was still first in thinking up the idea. AND shortly after I first wrote it down, years ago, I mailed off copies to various places, and have saved their reply-letters. So perhaps the idea eventually percolated to LLL (and, of course, perhaps someone else independently thought of it).
More details to follow, after I get at my records. My original document specifies a whole list of things to try to polymerize, MUCH longer than that list I saw at the link-site.
Note regarding the stability issue: "Once formed, the quartzlike CO2 remains stable at room temperature at pressures above 1 GPa and the researchers hope to able to isolate it at ambient pressures in the near future. One can expect that this new material has high thermal conductivity, just like diamond, and is also a very good candidate for a superhard material like diamond..." ---from the link-site
1 GPa is a GigaPascal of pressure. That is a LOT of pressure, just to maintain the stability of poly-CO2. I will remain doubtful that it will stay stable at room pressure (ESPECIALLY in the presence of humidity/water-vapor -- have they forgotton the reaction CO2+H20->H2C03, creating the tangy carbonic acid in soda pop?). But I will also remain hopeful -- vacuum balloons would be SUCH a nice technology.... (I suspect the best way to stabilize it will be to bond it with a coating of Teflon. Then there will be no "dangling" chemical bonds at the surface of the solid, for anything to react with and promote destabilization.)
OK, I now have the original document, and it is dated April 16, 1978. It is three pages, and starts off with "Let us consider the condensation of silicon dioxide gas into the molten state, which takes place at ordinary pressure when the temperature is approximately 2300 degrees Celsius, and falling."
Well, I sort of already described that in a prior Annotation: the molecules spontaneously polymerize. The essay goes on to wonder about obtaining monomolecular (or "triatomic") silicon dioxide (there are 3 atoms in the molecule, after all), using liquid argon instead of the helium I mentioned previously here.
Then: "Would triatomic silicon dioxide be a gas, a liquid, or a solid under room conditions? (This depends upon the strength of the van der Waals forces between its molecules.) Can it be used as a refrigerant? Or as a catalyst? How easily can polymerization be initiated? The questions and possibilities are endless.
"Silicon dioxide is not the only natural polymeric substance that we might be able to break into individual small molecules. An obvious alternative is germanium dioxide, since its molecular geometry is so similar to that of silicon dioxide. But there are less obvious possibilities, as well....
"Sulfur naturally occurs in ring-shaped molecules of eight atoms each, while oxygen, a close chemical relative, comes in molecules of only two atoms each. What would diatomic sulfur be like at room temperature and pressure?
"Is alumina (dialuminum trioxide) a polymeric substance? The various iron oxides? How about their sulfides? The oxides and sulfides of most of the transition elements can be examined with the question of polymericity in mind.
"Now for the other side of the coin: If there are natural polymers which may be broken up into their ultimate monomer constituents, then perhaps there are natural compounds which may be seen as potential monomers --- they may be polymerized.
"The whole plastics industry is proof enough of this speculation, but let's go one step further. Consider again the difference between the naturally-occurring molecules of sulfur and oxygen. If we can a diatomic allotrope of sulfur, can we not at least contemplate the creation of octatomic oxygen molecules? It would take a considerable amount of temperature and pressure, but it shouldn't be impossible. (However, we might expect the resulting oxygen allotrope to be highly unstable --- a super oxidizer for use in rockets.)
"When liquid sulfur is poured into water, the rapid cooling that results causes the sulfur atoms to link into long chains containing many thousands of them, and this the allotrope known as 'rubber sulfur' is formed. Can oxygen be made to do the same? (What an oxidizer THAT would be!)
"Consider the difference between carbon dioxide gas and silicon dioxide polymer...."
OK, I snipped that last paragraph because everything in it is already posted here. But the essay concludes:
"An examination of the Chemical Rubber Company's Handbook of Chemistry and Physics locates the following inorganic compounds which may easily be seen as potential monomers, substances which perhaps can be polymerized: carbon monoxide; nitric oxide; nitrosyl fluoride; nitrosyl chloride; carbonyl difluoride; carbonyl chloride fluoride; carbonyl dichloride; silicyl difluoride; silicyl chloride fluoride; silicyl dichloride; nitrogen dioxide; carbon disulfide; sulfur dioxide; carbon oxysulfide; nitryl fluoride; nitryl chloride; sulfuryl difluoride; sulfuryl chloride fluoride; sulfuryl dichloride; and sulfur trioxide. There are certainly others, all of which may be mixed before the polymerization reaction is initiated...."
End of original essay. I will assume you don't need me to drag out the reply-letters mentioned above.-- Vernon, Sep 26 2001, last modified Sep 27 2001 Carbon dioxide does exist in solid form at very low temperatures and at very high pressures. "Dry Ice" is used in ice-cream-truck freezers to keep popsicles cold during hot summer months; at deep ocean depths, carbon dioxide forms "clathrates," solid, nonbuoyant chunks of material. I don't know if anyone has ever done extensive materials analysis at those depths. Perhaps it does crystallize, or polymerize.-- whlanteigne, Oct 05 2002 Would you settle for a CO2 copolymer? I think CO2 can be made to copolymerize with certain epoxides....-- mhh5, May 22 2003 (neutral) ..... But there is very little around on this because its a no-brainer that CO2 can form such structures..... so can H20.... seen a snowflake up close? C02 can form a solid.... ice... which if made with a certain method can be made to have startling similarities to quartz.-- venomx, May 22 2003 whlanteigne and venomx, you are not paying attention to the fact that a solid object is not necessarily made of polymerized molecules. In a snowflake, for example, the water molecules are linked together by "hydrogen bonds", not full-fledged chemical bonds. A hydrogen bond is just an electrical attraction between two atoms, but is not a sharing of electrons. A polymer is always constructed using full-fledged chemical bonds.-- Vernon, Sep 10 2003 <Vernon>, I'm not a chemist, so these questions are sincere.
Just looking at the first dictionary.com deffenition for polymer, can a polymer be hard? Wouldn't it, somewhat by deffenition, be extremely flexible?
If all the bonds for all the atoms are filled in CO2, how could a chemical bond to other CO2 atoms be made? From what I understand, a polymer is a chain of similar molecules. If you wanted a CO2 polymer, then wouldn't the CO2 have to be chained together, and only contain CO2? If this is the case, how can CO2 bond if it's bonds are filled?
Or, would a CO2 polymer consist of mostly CO2?
What about these? (didn't add this one to link on purpose admin ;)) http://broken.pc.cz/gallery/halfbakery/CO2_polymers.htm Both could be extended. Would they be considered a polymer? Would they work?
Seems like both would be very weak/unstable, but very flexible.
Would a carbon monoxide polymer be easier to make?-- nomel, Jul 29 2004 Weeelll... What plants do with the CO2 they get from the air is to polymerize it into carbohydrates, losing molecular oxygen... Somehow I think this is not exactly what you are thinking of...-- GutPunchLullabies, Jul 29 2004 nomel, (FYI) a "monomer" is the starting chemical from which a polymer is made. Sometimes more than one monomer is involved. ALL monomers have one simple thing in common: A chemical double-bond between two of the atoms in the monomer (and sometimes a triple bond).
Consider an imaginary molecule consisting of atoms A and B, we might portray it as A=B to show the chemical double bond. Polymerization breaks those bonds and allows separate molecules to join to make a chain: ...A-B-A-B-A-B... Carbon monoxide is such an A=B molecule, but how easily it can be polymerized remains to be seen.
Flexibility is related to the length of the chains (longer is usually more flexible). Monomers that have more than one double-bond can offer additional linkage points, so that the polymer might have as many connections as a casual examination of a chain-link fence (a grid of squares). These sorts of polymers are generally not very flexible.
Finally, in CO2, both oxygens have double-bonds with the carbon. Polymerized, and because of the orientations of those links from a carbon atom, Poly-CO2 will have a fully 3-dimensional linkage structure, the same as quartz, pure crystalline silicon, and diamond (the least flexible thing known).-- Vernon, Aug 04 2004 [Vernon], Would those shown on the page I posted in my last comment work? Would they be considered co2 polymers?
This idea seems baked. I finally searched for some material since it's pretty obviouse once you realise the similarities between silicon and carbon.
I have added the link. From the paper, it can be done! So, good job vernon! ;-)
And from another site, "[C and Si] form long chains, or polymers, in which they alternate with oxygen. In the simplest case, carbon-oxygen chains yield poly-acetal, a plastic used in synthetic fibers, while from a backbone of alternating atoms of silicon and oxygen come polymeric silicones."
I give you a crossaint anyways.-- nomel, Aug 12 2004 Doesnt really make sense to me. Silicon and Carbon arent in the (what was the word family?, the columns are different) meaning the properties are different, so it would behave differently.-- gnaeuspompeiusmagnus, Aug 12 2004 nomel, thanks, what you described is indeed a kind of poly-CO2. However, I gather that only one of the two double-bonds is used for the linkages, since you have described a single chain. The type of poly-CO2 that I had im nind was the fully-crosslinked 3D variant, analogous to quartz crystal. A year or two ago somebody linked to a place that indicated that it had been done, but the link since broke, so hard to be sure.
gnaeuspompeiusmagnus, you are mistaken; carbon and silicon ARE in the same column of the Periodic Table of Elements, and they do have a number of similar properties. For example, all that pure crystalline silicon used in microchip manufacturing is structurally identical to diamond (just silicon atoms instead of carbon atoms).-- Vernon, Aug 12 2004 [Vernon], did you read the link I posted. "The many faces of carbon dioxide"?
and, those in my pictures use all bonds available.-- nomel, Aug 17 2004 nomel, sorry, I had not noticed the link before, and was just responding to your descriptions. It does look like the CO2-V that they describe there might be the Poly-CO2 that I had in mind all along. I also note that they are wondering if it can exist at ordinary pressures...just like I've previously posted here.
Thanks!-- Vernon, Aug 17 2004 Well, the answer you're looking for would be in the phase diagram for carbon dioxide. I've found one for you and linkinatified it.
However, note that crystallizing a substance does lead to added stability, and would shift the phases in the diagram down (lower pressures) and to the right (higher temperatures). I doubt, though, just based on the readiness with which C02 is utilized in the natural world, that crystalline C02 would ever exist at STP, which I guess is where you really want it.
Now, egnor's link did mention that the C02 crystals were held at 320 below 0 F, or -195 C. That's really not all that cold - it's still 78 degrees from absolute zero. There is a possibility that at 1 atm, this structure could exist in the natural world in exceedingly cold conditions, but I doubt you'd ever find anything that cold IRL.-- shapu, Aug 17 2004 My ignorance in chemistry will surely blind everyone, but if you could get the pressure to make it, could you then once it was made let it go and be in aluminum reflective sheathed existence out in space? I was thinking of something like this for use as an unmanned self constructing scaffolding in space. I just don't have the knowledge that is displayed here to push much farther with idea without extensive study.-- Zimmy, Jul 11 2005 [Zimmy], the question being asked here is not really about the pressure needed to create some Poly-CO2, but what happens after the pressure is released. There are reasonable doubts that it will be stable at ordinary temperature/pressure conditions (to say nothing of heat/cold/vacuum of space). The stuff may decompose, and perhaps may decompose explosively. Have to make some to find out!
The best-case scenario I can imagine is that it could be stable provided no water molecules are in the vicinity. So, any freshly-made Poly-CO2 should be given a surface treatment with fluorine gas, to seal off all the chemical bonds that are "dangling" and which would react decompositionally with water (even with water vapor in the air). Again, though, the experiments have to be done before we will know the realities regarding this speculation.-- Vernon, Jul 11 2005 Funny how I came across this idea somehow in searching for crystallization of gases in extreme cold, though. I started with Hydrogen & worked my way through various common gasses. I'm no master of the internet & wish I were. It seems that too often I get inundated with articles on dry ice when that's not really what I'm looking for.-- Zimmy, Jul 11 2005 Vernon's new link churned this up, so I figured I might comment...
Under immense heat and pressure, poly-CO2 was produced. It's stable at room temperature once manufactured - as long as 400,000 bar is maintained.-- david_scothern, Jun 19 2006 I wonder if the folks pumping CO2 into deep earth reservoirs have considered pumping (dropping) it in as clathrates, for better stability?-- bungston, Sep 11 2006 The reason you get molecular CO2 rather than polymeric CO2 is to do with the energies of the bonds. A C=O double bond requires more energy to break than the total of two C-O single bonds. In silica the opposite is true, more energy is required for two Si-O single bonds than an Si=O double bond.-- K o R, Sep 16 2006 [K o R], thanks! Can I assume that's related to the sizes of the carbon and silicon atoms, relative to oxygen's size? That is, the two available bonds for oxyen "come out" at an angle, and this is true for two of the four bonds for either carbon or silicon -- but the distance between two bonds of carbon is a better fit with oxygen, than the distance between two silicon bonds?-- Vernon, Sep 16 2006 Welcome to Molecular Orbital Theory, [Vernon]!
Added a link for you, this is how bonding actually happens - many interesting topics.
Size will factor into it, since C and O are similar in size, the symmetry of the overlapping P-orbitals will be better between C and O giving a stronger Pi bond and hence a stronger C=O bond. Si is bigger than O so its Pi bond is weaker and hence the two single bonds are more energetically beneficial to the molecule.-- K o R, Sep 17 2006 Okay, then if we were talking about polymerizing carbon disulfide, instead of carbon dioxide, we should get something stable at room temperature, right?-- Vernon, Sep 18 2006 Possible, but not really a good idea as CS2 is a neurotoxin...-- K o R, Sep 18 2006 according to the first link CO2 crystals evaporate at extremely low temperatures. However, you may be able to encase it in a thick layer of some other material to keep it cool- It wouldn't last long against the impact of solid objects, however it would still be able to resist evenly distributed pressure. It could be used in deep sea exploration, in which case you could put a refrigeration system on the inside of the craft and some heat-proof material on the outside. You could put refrigerating systems on either side of a thin layer of the substance , but then it could only resist impact from a small range of angles.-- apocalyps956, Sep 21 2006 random, halfbakery