Anders Hoveland - 27-6-2010 at 13:51
I read that there is an equilibrium between Cl2O7 and anhydrous HF.
Cl2O7 + HF <--> ClO3F + HClO4
If this is so, than would it be expected that NaF and Cl2O7 would react to make NaClO4 and ClO3F? Would this reaction be endothermic? If it was, it
might absorb heat as perchloryl fluoride gas was given off, further shifting the equilibrium to the right.
It was in "Advances in inorganic chemistry and radiochemistry" in google books but now I cannot find exactly where in the book it was.
Similary, one would expect NO2+ nitronium ions to react with fluoride ions to make NO2F, which is covalent. Perhaps a nitrate, concentrated sulfuric
acid, and NaF slowly added to the mixture, would give off NO2F. Do you think NO2+ would attack CaF2, or would the extreme insolubility prevent the
equilibrium from shifting to a gas? If CaF2 did not react with NO2+, then it is logical to infer that NO2F would react with
Ca(NO3)2 to make N2O5.
This could be a way for the home chemist to make extreme fluorinating compounds, without having to first make F2 electrolytically.
Fleaker - 27-6-2010 at 15:42
Good lord man, you're mad.
UnintentionalChaos - 27-6-2010 at 15:58
I don't think we've actually pulled out a ban in years (except maybe for spammers, I wouldn't know). This is seriously bordering on spam.
Anders Hoveland - 27-6-2010 at 16:07
Yes, but if you look at the description under my name, you can infer that I theorize, urging others to do the suicidal reactions of my design. This
forum IS called "scienceMADness" after all. If you really want to see some treacherous reaction, check out my topic "Reactions to Easily make new
Precursors".
By the way, if you ever accidently take a lethal breath of HF fumes, you can quickly inject yourself with calcium acetate solution, and this will
double your body's tolerance against HF poisoning. [Harwood-Nuss' Clinical Practice of Emergency Medicine By Allan B Wolfson]
JohnWW - 27-6-2010 at 16:11
Nitronium fluoride mentioned above, NO2F, also called nitryl fluoride, which is used as an oxidant in rocket propulsion and a fluorinating agent, if
it contains pentavalent N through being made from a nitrate without being reduced, CANNOT be ENTIRELY covalent, because the N atom has only four 2s
and 2p orbitals available for bonding, usually as sp2 or sp3 hybrids. For trivalent N to become pentavalent, it must either lose an electron or add a
cation of some sort, to become positively charged. The vacant 3s and 3p orbitals are at much too high an energy level to be utilized. The electronic
structure for it shown on http://en.wikipedia.org/wiki/Nitryl_fluoride shows a positive charge on the N, a N-F single bond, and two equivalent O atoms bonded to the N with
average bond order 1½ with a negative charge alternating by resonance between the O atoms, and thus with a total of 4 covalent bonds. In spite of its
polarity and molecular weight, its being a single covalent molecule with no tendency to ionize to NO2+ and F- results in a boiling-point of only
-72ºC See also http://www.chemindustry.com/chemicals/0906111.html
It is also theoretically possible for an isomeric NO2F with trivalent N to exist, with the two O atoms not being equivalent due to an O-F bond:
O=N-O-F , but it would not be particularly stable.
I may also mention that, for the same reasons as above, and also on account of steric crowding, noting that F cannot form bonds with bond-order
greater than 1 unlike O, all attempts (even by Karl Christe, noted for his seemingly impossible hypervalent F compounds) to synthesize NF5, as either
a covalent compound or [NF4]+F-, have failed, including attempts to trap it at very low temperatures and very high pressures, only NF3 and F2 being
recovered. Of course, the cation NF4+ has been known since 1966, quite stable in salts of the strongest complex oxy- and fluoro-acids, e.g. with
ClO4-, NO3-, BF4-, PF6-, SbF6-, AsF6-, etc.; but it is readily hydrolyzed by H2O to NF3 and O2 ( http://en.wikipedia.org/wiki/Tetrafluoroammonium ).
See:
“Nitrogen Pentafluoride: Covalent NF5 versus Ionic NF4+F- and Studies on the Instability of the Latter," J. Am. Chem. Soc., 114, 9934 (1992), K
Christe with W. W. Wilson abstracted in http://pubs.acs.org/doi/abs/10.1021/ja00051a027
http://www-rcf.usc.edu/~kchriste/christe.htm
http://www.osti.gov/energycitations/product.biblio.jsp?osti_...
[Edited on 28-6-10 by JohnWW]
Anders Hoveland - 27-6-2010 at 16:18
Yes, and the same thing for ClF6+ and F-.
ClF7 cannot be made. I these particularly interesting reactions.
A parallel exists with the reaction of permanganate and acidified H2O2.
I have posted several speculative reactions using NF4+ as a reactant.
ScienceSquirrel - 27-6-2010 at 16:29
When it comes to NF4+, it does nor exist, nor can it exist.....
Zzzzzzzz :-(
Anders Hoveland - 27-6-2010 at 16:36
tetrafluorammonium salts definitely exist.
http://pubs.acs.org/doi/abs/10.1021/ic00287a016
From wikipedia:
"The tetrafluoroammonium ion forms salts with a large variety of fluorine-bearing anions. These include the bifluoride anion (HF2-),
tetrafluorobromate (BrF4-), Ge, Sn, or Ti pentafluorides; P, As, Ni, Sb, Bi, or Pt hexafluorides, XeF7-), XeF8-, various oxyfluorides (XF5O- where
X is W or U; FSO3-, BrF4O-), and perchlorate. The nitrate salt, NF4NO3, has also been made."
^ Hoge, B. (2001). "On the stability of NF+4NO−3 and a new synthesis of fluorine nitrate". Journal of Fluorine Chemistry 110: 87–88.
doi:10.1016/S0022-1139(01)00415-8.
I wonder if NF4HF2 would react with Cl2O7 to decompose into NF3, ClO4F, and ClO3F ?
NF4HF2 + Cl2O7 --> NF4ClO4 + ClO3F + HF
or...
NF4HF2 + Cl2O7 --> NF4(+),F(-) + ClO3F + HF
NF4(+),F(-) --> NF3 + F2
HClO4 + F2 --> HF + ClO4F
the last two of these reactions have been well documented.
[Edited on 28-6-2010 by Anders Hoveland]
woelen - 27-6-2010 at 22:56
Anders, what chemicals do you have access to? You have posted many theoretical reactions, but if you have access to e.g. anhydrous HClO4, HF and other
interesting (but for most home chemists considered too dangerous) chemicals, then I would love to see you write about a real experiment.
The problem is that virtually none of us will ever have access to the chemicals you are mentioning. Sometimes you make a remark about making such
things (an example you mentioned as response to my experiment is making borane and a borane hydrazine adduct). Making _and isolating_ that kind of
chemicals is a lot of a difference with e.g. making copper chloride, ferric oxide and so on.
I do like theorizing about strange and uncommon compounds, but the theory must be accompanied with practical results and insights and there must be
some chance that people actually can test the theory.
Alexein - 28-6-2010 at 07:35
By all means Anders, go for it and tell us your results. I would like to see some good characterization of the products though.
And remember to take all the safety precautions you can, most of those chemicals are nasty.
blogfast25 - 28-6-2010 at 08:34
... and good luck with the Cl2O7, just for starters! 
Anders Hoveland - 1-7-2010 at 19:55
I have some SO3 that I distilled. Making it is not too difficult.
When I add some NaF into SO3, nothing noticeable happens, but the fluoride IS soluble. Probably NaOSO2F is made. Adding nitrate to this, nothing
happens.
I think the problem is that NO2(+) (-)SO3F is forming. The SO3 probably holds the NO2F very strongly, so apparently I will need to make pure N2O5
without any trace of SO3, and react the N2O5 with HF.
By the way, I silver over my glassware, and then put a nickel plate over the silver to handle the HF, since I do not have teflon. Both metals can be
plated chemically, but I could not obtain the hydrophosphite necessary to chemically plate the nickel, so must resort to electric plating.
I can easily make Cl2O7, with KClO4 and SO3, but it is contaminated with SO3, so I would probably not be able to make ClO3F, since the SO3 would hold
it as O3ClOSO2F
[Edited on 2-7-2010 by Anders Hoveland]
woelen - 1-7-2010 at 23:15
How did you distill the SO3? From what? SO3 is not an easy chemical to obtain, not at all!
SO3 is a solid at room temperature, did you mix solid SO3 with solid NaF?
Could you make pictures of your silvered and nickel plated glassware? It would be really cool to have that kind of stuff for handling solutions of HF.
I also made silver mirrors, but they are extremely vulnerable, not something robust enough to work with.
Anders Hoveland - 1-7-2010 at 23:42
My SO3 is liquid. It probably absorbs too much water vapor from the air. Apparently solid SO3 is almost completely miscible in H2S2O7.
I am using sodium pyrosulfate and boric oxide in my distillations. I put it into a metal pipe and put the pipe over a propane burner. The pipe can
only be used once. I do not get high yields (that is a drastic understatement), but the boric oxide seems to help me do it at a lower temperature.
I have also tried burning SO2 in Cl2 with an activated carbon catalyst. A little bit of SO2Cl2 can be made this way, which would react with water to
make pure H2SO4. I once reacted sulfuryl chloride with NaNO3; it gave off chlorine. Not sure what the reaction was. That was years ago, have not done
anything with SO2Cl2 since. It was really messy and the Cl2 fumed too much.
I might try reacting SO2 with N2O5. The N2O5 I made from nitric acid and P2O5 suspended above the liquid, in a sealed container, and left for a month.
SO2 + 2N2O5 --> 2NO2 + (NO2)2SO4
I forgot to mention, before you silver your glass, let the glassware soak in a bath of NH4F solution (in a plastic bucket) until it gets all frosty.
This makes the silver adhere much better to the glass. After you nickel plate it, the plating becomes much stronger, but still definitely not
impervious to scratches or dents. It is not an ideal solution, and pieces of the plating keep falling off after repeated use.
http://www.brownells.com/1/1/27045-electroless-nickel-platin...
I just reacted superconcentrated perchloric acid (mostly HClO4, but containing some dissolved Cl2O7) with NaF. I used just a drop of the liquid on a
speck of NaF. It bubbled a little. This is obviously HF being given off. When I used excess NaF, not all of it dissolved, and my HF was unable to
spontaneously ignite a paper strip suspended above it. I think pure Cl2O7 is needed for the reaction, but it is really hard to dehydrate HClO4 all the
way to pure Cl2O7 without dangerous distillation.
[Edited on 2-7-2010 by Anders Hoveland]
Anders Hoveland - 6-7-2010 at 00:01
I made barium perchlorate, heated it to drive off moisture, and (tried to) dissolved it in pure HClO4, but interestingly it did not seem very soluble.
Adding a drop of SO3 caused the solution to become cloudly, meaning BaSO4 was precipitating. So now the solution was a mix of anhydrous HClO4 and a
little Cl2O7. Unfortuneately, reacting SO3 with dry solid barium perchlorate would not work very to make Cl2O7, because of the protective BaSO4 that
would inhibit further reaction. Normally, I might use a pestle and mortar to keep grinding at the powder and force a reaction, but that would probably
be, in this case, unwise in the extreme.
woelen - 6-7-2010 at 01:18
If you have pure anhydrous HClO4, then why not try adding some P4O10 to this? I can imagine that this gives Cl2O7 and H3PO4 (or HPO3) and you might be
able to use the Cl2O7 as is, without the need to isolate it from the phosphoric acid.
P4O10 is much easier to obtain than SO3 and also much safer to handle. But still, this kind of experiments is very dangerous. I see no way of
obtaining pure Cl2O7 without dangerous distillation. I myself certainly will not attempt that kind of operation, the chance of getting an accident
simply is too large.
Anders Hoveland - 6-7-2010 at 08:41
The distillation requires low temperatures, and it goes without saying, reduced pressure. I probably could do the distillation through warm
evaporation, but this would be very dangerous.
"The compound is obtained by cautiously regulated action of phosphoric oxide on perchloric acid"
Michael and Conn, Am. Journ. Chem. ,1900.
“Perchloric acid is commonly obtained as an aqueous solu- tion, although the pure anhydrous compound can be prepared by vacuum distil- lation as a
colorless liquid, which freezes at minus 128C and boils at 168C at 2.4 kPa (18 mm Hg) without decomposition. The pure acid cannot be distilled at
ordinary pressures and explodes at 908C after standing at room temperature for 10–30 days. The aqueous solution can be concentrated by boiling at
101 kPa (1 atm) at 203C, at which point an azeotropic solution is attained which contains 72.4% HClO4. For purification by distillation, reduced
pressure is needed below 200 mm to avoid partial decomposition to chlorine, chlorine oxides, and oxygen (19–24). Commercial 72% perchloric acid
contains only slightly more water than the dihydrate”
“Cl2O7 is obtained as a colorless oily liquid by dehydration of perchloric acid using a strong dehydrating agent such as phosphorus pentoxide.
The Cl2O7 decomposes spontaneously on standing for a few days. The acid dehydration reaction requires a day for completion at minus 10C and explosions
can occur. Upon ozonation of chlorine or gaseous ClO2 at 30C, Cl2O7 is formed (13). >>comment: chlorine catalyzes the decomposition of ozone, so
the reaction would waste most of the ozone. Oxidation of ClO2 with O3, I read, can also afford Cl2O6, presumably when less O3 is used<<
Chlorine heptoxide is more stable than either chlorine monoxide or chlorine
dioxide; however, the anhydride detonates when heated or subjected to shock. It melts at minus 91.5 C, boils at 80C . It decomposes to chlorine and
oxygen at low (0.2–10.7 kPa (1.5–80 mm Hg)) pressures and in a tempera- ture range of 100–120C . It is soluble in benzene, slowly attacking the
solvent with water to form perchloric acid; it also reacts with iodine to form iodine pentoxide.”
Interestingly, the site mentions an alternate reaction, one that indirectly generates chlorine, and uses it to oxidize the ammonia away from the
ammonium salt. I had this idea before, but was told by someone here that they had tried using ammonium sulfate and chlorine and it did not work.
“reaction of ammonium perchlorate with nitric and hydrochloric acids, and then concentration at 198–200C to eliminate the unreacted acids by
vacuum distillation:
34 NH4ClO4 +36 HNO3 + 8 HCl --> 34 HClO4 +4 Cl2 + 35 N2O + 73 H2O “
“Perchloric acid is commercially manufactured by reacting saturated solution of sodium per- chlorate with hydrochloric acid. Precipitated sodium
chloride is separated from the dilute solution (32% by weight HClO4) by filtration, and the solution is concentrated to 70% by weight via vacuum
distillation.”
19. A. A. Shilt,Perchloric Acid and Perchlorates, G. F. Smith Chemical Co., Columbus,
Ohio, 1979.
20. F. C. Mathers,J. Am. Chem. Soc 32, 66 (1910).
21. H. H. Willard,J. Am. Chem. Soc 34, 1480 (1912).
22. D. Vorlander and R. von Schilling, Liebig’s Ann 310, 369 (1900).
23. A. Michael and W. T. Conn,Am. Chem. J 23, 445 (1900).
24. T. W. Richards and H. H. Willard,J. Am. Chem. Soc 32, 4 (1910
http://www.scribd.com/doc/30127718/Perchloric-Acid-and-Perch...
[Edited on 6-7-2010 by Anders Hoveland]
Formatik - 8-7-2010 at 21:18
Cl2O7 has been made by letting anhydrous HClO4 in contact for one day with P2O5 below -10 C, then it is slowly heated to 85 C. Practically pure Cl2O7
distills over. By this preparation, severe explosions can occur. This freshly prepared Cl2O7 can be distilled once more without hazard under
atmospheric pressure - A. Michael, Conn (Am. chem. J. 23 [1900] 445, 25 [1901] 92). More "harmless" method is from KClO4 and chlorosulfonic acid
though product is only 98-99% pure, F.Meyer, Keszler (Ber. 54 [1921] 567).
Perchloryl fluoride (FClO3) is very stable, poisonous and reactive (Bp. -46.7C, Mp. -147.7C). Electrolysis of satd. NaClO4 in anhydrous HF yields the
compound. Another way in 85-90% yield, is to warm a mixture of KClO4, HF and SbF5 at 40-50 C (Kirk Othmer). FClO3 is also stable up to 400 C, and
hydrolyzes slowly. Grease and rubber tubing has caused explosions, for more reactivity see Brethericks.
The German wikipedia claims alkali fluorides reacting with Cl2O7 does yield FClO3, though no exact reference is given for this.
Hantzsch claimed to have made fluoronium perchlorate [FH2]ClO4 by reacting anhydrous HClO4 with anhydrous liquid HF, which under strong heat evolution
was said to yield solid [FH2]ClO4 (Ber. 60 [1927] 1946), and which compound he said reacts explosively with H2O (Ber. 63 [1930] 97). Brauer and
Distler (Z. anorg. u. allgem. Chem. 275 [1954], 157) tried to make this compound, but could not repeat preparation despite mixing in various ratios
and temperatures.
FClO3 is also made by reacting fluorine with KClO3 at -20 C in SbF5: KClO3 + F2 = KF + FClO3 Or by reacting KClO4 with HSO3F: KClO4 + HSO3F = FClO3 +
KHSO4 (From: Lehrbuch der anorganischen Chemie by A.F. Holleman, E.Wiberg, N.Wiberg). On the last one, no more decent details given by Holleman et
al., it could react right away or may need some warming, time to hit the more serious lit.
With ammonia, FClO3 forms a perchlorylamide: FClO3 + NH3 = ClO3(NH2) + HF. This has acidic protons and they are replaceable by metal ions: K[ClO3(NH)]
and K2[ClO3N] these are colorless, up to 300 C stable compounds, which explode by impact.
Electrophilic substitution in presence of Lewis acids of FClO3 on aromatics has yielded some interesting compounds via introduction of -ClO3 groups.
Perchlorylbenzene, nitroperchlorylbenzene, etc. preparation is in US3067211, US3937627. Nitrogen heterocyclics also, US3332955 describes
N-perchlorylpiperidine from ClO3F, this compound has been known to explode on storage. The same have been made from Cl2O7 and cyclic nitrogen amines
in an inert solvent. Perchloryl aromatics are usually shock sensitive. 3-Nitroperchlorylbenzene is about as shock sensitive as lead azide, and said to
have a very high detonation rate. Not exactly inert either, even to one of the chemicals used to make them: a perchlorylbenzene/AlCl3 mixture does
nothing, then explodes after some time. With FClO3, alcohols said to turn into extremely explosive alkyl perchlorates, etc.
[Edited on 9-7-2010 by Formatik]
Formatik - 11-7-2010 at 01:03
Caution on mixing anhydrous HClO4 with SO3:
From Brethericks: "Interaction of the anhydrous acid and sulfur trioxide is violent and highly exothermic, even in presence of chloroform as diluent,
and explosions are frequent."
It seems the same is also the case for P2O5 (and so the cooling mentioned above), since a solution of HClO4 in CHCl3 explodes on contact with P2O5
(Vorl., v.Sch., Lieb. Ann. 310 [1900] 374).