ajkoer

Some more history: This is a case of a water treatment tank exploding. To quote: "Redding, California – June 17, 2011 Fred Crumb, 54, was killed Thursday when a large water tank exploded at the Clear Creek Wastewater Treatment Plant. The California Occupational Safety and Health Administration report that Mr. Crumb was working at the sewage treatment plant when a steel 4,500-gallon tank exploded. The tank was 33 years old. According to the Shasta County Coroner’s Office, Mr. Crumb was killed as a result of critical injuries he sustained to his head in the explosion. . Redding Public Works and the California Safety and Health Administration are still investigating the explosion." This could be a case of pre-aeration (H2S accumulation) or active chlorination occurring (NCl3 formation/detonation). The presence of an old steel tank, most likely corroded, may indicate some FeCl3. Ferric chloride can increase the so called 'activity level' of acids in highly salt rich solutions, so dilute HCl and HOCl (from Cl2 hydrolysis) and even H2CO3 (from CO2) are effectively stronger acids. This would lower pH and favor NHCl2, NCl3 and even Fe(HCO3)2 formation. However, is also likely, given that the steel was 33 years old, it was galvanized steel (meaning coated with Zn). As such, an electrochemical oxidation (that similarly occurs in the presence of Copper) of the NH2Cl and NH3 to nitrogen via nitrites, implying a possible rapid N2 gas decomposition, may be responsible as well

While this rare event of massive sized water tanks rupturing occurs rarely, it is oftened accompanied by some loss of life and large property loss. I have put together some chemical based theories on how some of the evidently pressure detonations could be occurring. Those more familar with these large water tanks may be able to supply more insights (like what metal is employed, ventilation/pressure release valves, is the water stored pre or post aeration and/or chlorination,.etc). If a coherent list of possible causes can be produced, I will forward to OSHA, which apparently review these incidents, for considerations. As to why I am presenting on this forum is a net search reveals little insight and I can understand why a very respected scientist may feel reluctant to be associated with a topic on water tank pressure explosions. First, some history of the events to ascertain some possible patterns. Here is a report of a large explosion from Fox News reported on April 07, 2011 "Two Killed From 300,000 Gallon Water Tank Explosion" (see http://www.foxnews.com/us/2011/04/07/killed-300000-gallon-wa... ). To quote: "Two men died Thursday when a 300,000 gallon water tank exploded in Florida, unleashing a flood and causing an adjacent building to collapse. The victims were in the midst of repairing a pump that filled the tank inside an adjacent concrete block building. The force of the water from the explosion caused the building to collapse, MyFoxTampaBay.com reported." Here is another incident in Tomball, Texas where a worker was killed after a water storage tank exploded (see http://www.khou.com/news/Worker-injured-in-church-water-tank... ) to quote: "The man was cutting on the top of the tank to provide ventilation," said Lieutenant Chad Shaw with the Harris County Fire Marshal’s Office. "The tank was about three quarters full of water but there was a build-up of combustible vapors above the water. Sparks or a flame caused by the cutting ignited the vapors, causing the explosion." Here is a report of yet another large water tank explosion in Galax, Virginia (see http://www.thecarrollnews.com/view/full_story/22210488/artic... ), and also in Chester, New York (see http://chroniclenewspaper.com/apps/pbcs.dll/article?AID=/201... ) where to quote: "Internal pressure blew the end of the tank off and through the attached treatment building, completely demolishing the building,” police said. This is also report at http://www.jstor.org/discover/10.2307/41232342?uid=3739808&a... called "Investigation of a Water Pressure Tank Explosion" that occurred in 1938 following a water pressure tank explosion in the muncipal water supply of Bricelyn, Minnesota. Apparently, some eight months prior the tank was drained and received a Zinc lining. Less credible, but perhaps a valuable clue to the chemistry involved is even a report of an exploding fish tank (see http://www.ratemyfishtank.com/phpBB3/topic1309.html). But this may be just someone's nightmare, well perhaps not, as here is a report of a 33-ton Shark tank in a Shanghai shopping center lobby with 6 inch thick glass walls that cracked in just 2 seconds flat (see the video at http://thestir.cafemom.com/in_the_news/148711/shark_tank_exp... ). Source: New York Daily News, Dec 27, 2012 and also ABC News. ------------------------------------------------------------------------------------ Hypothesis: For those water tank events not related to a pressure eruption, my first suggestion for this class is most likely a flammable gas, Hydrogen sulfide (H2S ), which per Wikipedia (http://en.wikipedia.org/wiki/Hydrogen_sulfide ) is both flammable and explosive. It can be formed by the action of bacteria in sulfur rich water with deficient oxygen content, to quote: "Hydrogen sulfide often results from the bacterial breakdown of organic matter in the absence of oxygen, such as in swamps and sewers; this process is commonly known as anaerobic digestion. H2S also occurs in volcanic gases, natural gas, and some well waters." So the chemistry (or biochemistry) here would be water in a tank loses O2 on warming and in the presence of organics fosters the creation of some H2S gas, which being very heavy, could form dangerous explosive accumulation. -------------------------- Next, hypothesis for this class is the formation of explosives Chloramine (NH2Cl) vapors from chlorine in water (via chlorination) producing Hypochlorous acid (HOCl), which forms Chloramine in the presence of ammonia (from decaying matter): Cl2 + H2O <--> HCl + HOCl NH3 + HOCl <--> NH2Cl + H2O --------------------------- The last hypothesis for the explosive gas formation is perhaps the least likely cause. It is formation of Nitrogen trichloride or trichloramine (NCl3), a yellow oily liquid that floats on water and only slowly undergoes hydrolysis, which is explosively sensitive to heat, shock and the presence of organic matter. Now, a source for its creation is per Wikipedia (see http://en.wikipedia.org/wiki/NCl3 ) to quote: "Nitrogen trichloride can form in small amounts when public water supplies are disinfected with monochloramine, and in swimming pools by disinfecting chlorine reacting with urea in urine from bathers" So, a large water tank may provide a collection vessel for the formation of explosive NCl3 and its vapors. Some chemistry: NHCl2 + HOCl <--> NCl3 + H2O --------------------------------------------------------------------------------- The following hypotheses relate to the class of Water Tank Pressure Explosions, which I view as more complex with respect to chemistry: First path is per the source provided below, the decomposition of NH2Cl itself which forms many products including N2 gas leading to a possible pressure explosion/rapid decomposition reaction: "As shown in Table 1, chloramine loss by auto-decomposition is a relatively complex process. However, the overall rate of chloramine loss for neutral pH values and above is primarily limited by the rate of formation of dichloramine (Jafvert and Valentine, 1992). Dichloramine formation occurs through both monochloramine hydrolysis (reactions 1.2 and 1.3) and by a general acid catalyzed monochloramine disproportionation reaction (reaction 1.5). The relative importance of these pathways on the formation of dichloramine is dependent on factors like pH, ionic strength, temperature, and alkalinity. Once dichloramine forms it decomposes via a series of rapid redox reactions. The products of these reactions are primarily ammonia, chloride, and nitrogen gas, however, nitrate also forms under some conditions (Vikesland et al., 1998)." where the dichloramine is formed variously including: NH2Cl + NH2Cl --> NHCl2 + NH3 NH2Cl + H2O <--> HOCl + NH3 NH2Cl + HOCl <--> NHCl2 + H2O Link: http://www.researchgate.net/publication/12006087_Monochloram... ------------------------------------------------ A second path is a more simple model that does not require Chlorine or Hypochlorous acid. Just air (actually O2 and CO2), ammonia (from decaying organic matter) and the appropriate metal (Copper or Zinc). Per this source, "Kinetics and Mechanism of Copper Dissolution In Aqueous Ammonia", fully available after signing on to ones Facebook account at http://www.academia.edu/292096/Kinetics_and_Mechanism_of_Cop... ) Copper, for example, is capable of reacting slowly (or rapidly depending on concentrations) with ammonia and air to form a soluble cupric salt. A side product is the formation of nitrites. Upon acidification (with CO2), nitrites (like NH4NO2) can produce a rapid gaseous decomposition yielding N2. Some of the underlying reactions cited by this source include: 2 Cu + 4 NH3 + 1/2 O2 + H2O --> 2 [Cu(NH3)2]OH 2 [Cu(NH3)2]OH + 4 NH3 (aq) + 1/2 O2 + H2O --> 2 [Cu(NH3)4](OH)2 Cu + [Cu(NH3)4](OH)2 <---> 2 [Cu(NH3)2]OH And, with respect to this thread, an important side reaction forming a nitrite: 2 NH3 (aq) + 3 O2 + [Cu(NH3)4](OH)2 --> [Cu(NH3)4](NO2)2 + 4 H2O Another author cites the following reaction, see "Copper-Mediated Non-Enzymatic Formation of Nitrite from Ammonia and Hydrogen peroxide at Alkaline pH", that is pertinent relating to nitrite formation noted above (please see http://www.google.com/url?sa=t&rct=j&q=reaction%20of...Fchem%2Fchempdf%2FCT%3D23(646-656)AJ11.pdf&ei=iS-mUfCNN4nr0gGYw4D4BA&usg=AFQjCNFaObAi5_3NNOdt8e1DiRoiHzg9bg&bvm=bv.47008514,d.dmQ , which have some important subtle differences: Cu2O + 4 NH3 + H2O → 2 [Cu(NH3)2]OH Diamminecopper(I), then generated from reduction of the copper(II) salt or added exogenously, then facilitates the oxidation of ammonia: ...[Catalyst Role]..Cu(NH3)2]OH......................... NH3 + 3 H2O2 -----------------> HNO2 + 4 H2O With additional ammonia, the reaction with nitrous acid proceeds as follows: HNO2 + NH3•H2O --> NH4NO2 + H2O The important subtle difference here being only a trace amount of Copper need be present with Nitrous acid and Ammonium nitrite formation being the cause of a possible significant and sudden Nitrogen gas evolution. This reaction could occur over time, as a function of the rapidity of water turnover in the tank (the lack of would could occur following a large scale evacuation for, say, a hurricane event, which is interestingly one of the reported events above in Chester, New York with hurricane Irene). Now, I actually performed the above reaction replacing atmospheric oxygen with some dilute H2O2 to speed things up. To my surprise, Copper pennies (my Cu source) became readily covered with O2 in agreement with a cathodic reduction reaction of oxygen at the copper's surface per the author's electrochemical dissolution model. ------------------------------------------------------ The last model has relatively simple chemistry requiring an oxygen source (like air or Hypochlorous acid from the action of Chlorine and water), CO2 and a significant presence of Iron bicarbonate. The reaction is, for example: 2 Fe(HCO3)2 + HOCl + H2O --> 2 Fe(OH)3↓ + 4 CO2↑ + HCl where one mole of Hypochlorous acid (or half a mole of O2) liberates 4 moles of CO2 gas. --------------------------------------------------------- Side Notes: The decomposition of NH2Cl is also known to be expedited in the presence of nitrites (http://www.researchgate.net/publication/12006087_Monochloram... ) and also cupric salts .These latter comments may be significant when working with fish tanks fed by ones internal copper plumbing or with exposure to zinc. Municipal drinking water is frequently aerated for various reasons, I suspect, including purification, taste and to check the formation such gases. As a source see http://www.gewater.com/handbook/ext_treatment/ch_4_aeration.... to quote: "Aeration as a water treatment practice is used for the following operations: • carbon dioxide reduction (decarbonation) • oxidation of iron and manganese found in many well waters (oxidation tower) • ammonia and hydrogen sulfide reduction (stripping) Aeration is also an effective method of bacteria control" What is interesting about the above is sufficient aeration could remove nearly all the suggested paths to either a flammable gas, or ammonia and iron that could form H2S, N2 or CO2 gases. However, once Chloramine, NH2Cl, has been formed and given its reported relative stability (being one of the reason cite for its employment over Chlorine), aeration is not one of the more effective means for its removal. In summary, the incidents relating to explosive vapors could be attributed to H2S, NH2Cl vapors (which are reportedly explosive) or the NCl3 hypothesis (flame or shock initiated explosion), but the fish tank and other obvious pressure ruptures lends support to paths forming nitrogen gas (from the auto-decomposition of NH2Cl, or from ammonia forming nitrites catalyzed by Cu, or perhaps even Zn) or CO2 gas emission (from O2 or HOCl on Iron rich water).

Came across an interesting patent "Method for the production of anhydrous potassium tert.butoxide US 4577045", link: http://www.google.com/patents/US4577045 Some extracts of some possible important points as potassium tert.butoxide is presumed formed and decomposed to Potassium in my reaction chain: "a) using cyclohexane or hexane as withdrawing agent, (b) using the tert.butyl alcohol in such an excess with respect to the aqueous potash lye and the withdrawing agent that in the bottom of the column a 10 to 18 wt.-% solution of potassium tert.butoxide is present, and the content of tert.butyl alcohol in the gas mixture situated in the center of the column is between 50 and 90 wt.-%, and © distilling out a mixture of withdrawing agent, tert.butyl alcohol and water at temperatures between 65" and 75. "Preferably, however, the tert.butyl alcohol is to have a water content of less than 0.1% by weight". So having too much water is problematic to the formation of potassium tert.butoxide." "Lyes having KOH contents of about 50% by weight can be used; the KOH content can be even lower, but then correspondingly larger amounts of water have to be distilled out. For this reason the use of potash lyes of KOH contents under 30 weight-percent is not recommended." "In the column, the reaction product that forms therein and is dissolved in the tert.butyl alcohol/hexane or cyclohexane mixture is washed into the bottom of the column, which is kept at the boiling temperature throughout the reaction. The potassium tert.butoxide is then in the bottom in the form of a 10 to 18% solution in pure, anhydrous tert.butyl alcohol." "The tert.butyl alcohol is not only reacting agent, it serves simultaneously as solvent for the potassium butoxide obtained, up to 90% by weight of total tert.butyl alcohol amount. Less than 1% by weight of total tert. butyl alcohol is part of the aqueous phase. Therefore the amount of the tert.butyl alcohol needed for the whole process should be sufficiently great that a 10 to 18% by weight, preferably a 10 to 15% by weight, solution of the potassium salt will be present in the bottom of the distillation column." Now this last comment is interesting because without sufficient excess tert.butyl alcohol, the potassium salt, lying at the base, may be capable of thermal decompositon (to K?) due to uneven heating. In general, note the relatively low temperature (under 100 C) use to prepare potassium tert.butoxide, and the possible consequences of a higher temperature (around 200 C) suggested in the Mg/KOH/tert.butyl alcohol path to Potassium. Also, in this preparation of potassium tert.butoxide, hexane acts as withdrawing agent to remove water via distillation in place of Mg. So, could one use hexane together with MgO (and possible H2) at higher temperatures to this preparation to form metallic Potassium (Sodium)? A speculation for would be experimentors.

John Cuthber: Yes, I agree that it is a dangerous preparation, but having started a synthesis thread using Oxalic acid on various salts on Sciencemadness, the actual salt that did prove problematic was not the chlorate, but would you believe, the sulfate. Apparently, heating H2C2O4 and a sulfate gets to a point where the H2SO4 is so concentrated it literally shouts out of the test tube. Many chemists, including those with experience, are generally unaware of the nasty behavior of some common acids that are apparently very dangerous when highly concentrated (as they are not, for good cause, available).

As an alternative to the above, consider the following preparation. Per Wikipedia (http://en.wikipedia.org/wiki/Chloric_acid ): "It [Chloric acid] is stable in cold aqueous solution up to a concentration of approximately 30%, and solution of up to 40% can be prepared by careful evaporation under reduced pressure. Above these concentrations, and on warming, chloric acid solutions decompose to give a variety of products, for example: 8HClO3 → 4HClO4 + 2H2O + 2Cl2 + 3 O2 3HClO3 → HClO4 + H2O + 2 ClO2 The decomposition is controlled by kinetic factors: indeed, chloric acid is never thermodynamically stable with respect to disproportionation." So upon mild heating of solid NaClO3 and Oxalic acid dihydrate: 2 NaClO3 + H2C2O4.2H2O --> 2 HClO3 (g) + 2 H2O (g) + Na2C2O4 (s) and condensing the vapors, one could theoretically have highly conc HClO3, and with careful evaporation under reduced pressure, the disproportionation to HClO4, with possibly around 33% to 50% of the Chloric acid converted into HClO4. However, the yield is most likely lower for several reasons including the need for an excess of NaClO3 as otherwise, Oxalic acid may reduce the HClO3 to explosive ClO2. I would still, however, consider this preparation as potentially dangerous, and exercise appropriate safety precautions, especially limiting the quantities involved. Also, be aware that although ClO2 has a more pleasing smell than Chlorine, it is many times more poisonous as measured by recommended exposure limits.

Since Bordeaux mixture is prepared (see Wikipedia) from mixing solutions of copper sulfate and Ca(OH)2 or CaO, I suspect that what we have is a mixture of Cu(OH)2 and CaSO4. Now, apparently CaSO4 is very insoluble and cannot even be dissolved in HCl. So add HCl and recover a solution of CuCl2 leaving CaSO4 behind. Now add NaOH (try to avoid exposure to air and CO2): CuCl2 + 2 NaoH --> NaCl2 + Cu(OH)2 (s) Good luck. As a side note, I would also avoid the use of any carbonates as your final product would be more likely a basic copper carbonate (Cu2(OH)2CO3, for example). Dissolving the Bordeaux mixture in ammonia water will also work by dissolving the Cu(OH)2, but an issue may surface on removing the ammonia from the copper complex. For example, adding H2O2 may react with NH3 forming a mixed nitrite salt.Source, per Wikipedia (http://en.wikipedia.org/wiki/Ammonium_nitrite ) to quote: "It can also be prepared by oxidizing ammonia with ozone or hydrogen peroxide" and "Ammonium nitrite may explode at a temperature of 60–70 °C.[2] It decomposes more quickly when a concentrated solution than when it is a dry crystal." As such, I would avoid ammonium nitrite and related mixed Copper salt for fear of potential extreme instability even in aqueous solutions. Using NaClO or HOCl on the copper ammonium salt to separate out a copper salt is equally problematic, very toxic and explosive side products (NH2Cl, NHCl2, NCl3, or N2H4) with the likely formation of a basic copper oxychloride. However, treating the copper ammonium complex with CaO may produce a copper salt precipitate (see http://www.jstor.org/discover/10.2307/25040124?uid=3739808&uid=2129&uid=2&uid=70&uid=4&uid=3739256&sid=21102222166897 ), but this may be CuO and only slowly formed, hence my recommended route.

I did not mention the lowering of the pH upon boiling H2SO4 because the question asks what 'chemical tests' to employ, and this requires the use of somewhat imprecise litmus paper. If the H2SO4 is very dilute, upon boiling it may be difficult to discern a lowering of the pH via litmus paper, in my opinion.

Actually, I recommend 'John Cuthber' classic method also. However, if one is interested in a less conventional, more accessible for some, but more expensive (but perhaps safer in avoiding the storage of strong acids for the home chemists) route to Chlorine (Cl2) which occurs as a reported side product in this proposed method. The chief product of interest, cited in multiple patents (for example, see United States Patent 4380533 at http://www.freepatentsonline.com/4380533.html or, a non-patent textbook reference http://books.google.com/books?id=yZ786vEild0C&pg=PA74&lpg=PA74&dq=dibasic+magnesium+hypochlorite&source=bl&ots=WsG8bHQyX8&sig=XnsYZC6fir8XW61NDm3hvEc_iW8&hl=en&sa=X&ei=jal-UeqrCNLh0wG024GQDw&ved=0CDoQ6AEwBg#v=onepage&q=dibasic%20magnesium%20hypochlorite&f=false ) is dibasic Magnesium hypochlorite Mg(OCl)2.2Mg(OH)2 (used in bleaching and sanitizing applications) in which Cl2 is also apparently generated. The bleaching salt is reported to separate out easily. So, apparently one does not need HCl (impure or otherwise) just NaCl, NaOCl and MgSO4 hydrate (Epsom salt). The process for liberating Cl2 and in more dilute solutions (or, upon addition of H2O to a closed vessel and shaking) would result in chlorine water. The cited net ionic reaction per the above patent: 3 Mg+ + 4 ClO- + 2 Cl- + 2 H2O ➝ Mg(OCl)2.2Mg(OH)2 + 2 Cl2 Or, upon employing NaCl, NaOCl and MgSO4.7H2O, the net reaction could be stated as: 3 MgSO4.7H2O + 4 NaClO + 2 NaCl ---> Mg(OCl)2.2Mg(OH)2 + 2 Cl2 + 3 Na2SO4 + 19 H2O As NaOCl found in commericial chlorine bleach is basic per the addition of a little NaOH and, at times, Na2CO3, adding acetic acid (or ascorbic acid,..) to adjust the pH to be between 3 to 7 may be required (or the use of a more acidic chloride salt like CaCl2 in place NaCl). Dilution here may not actually be a bad thing as the goal, after all, is chlorine water, however, I suspect that the reaction actually requires a concentrated solution. My logic: the normally weak acid hydolysis of sulfate salt produces, in highly ionic conditons, a higher 'activity level' for the acid (or a lower one for H2O), resulting in the liberation of chlorine. Note, in earlier versions of this patent, a solid hydrated salt to solid reaction is recommended, which is in line with my ionic characterization. Some may find more interesting a claim from an old sciencemadness thread that Cl2 is generated by the action of bleach (actually a mixture of NaOCl and NaCl due to its creation by the action of chlorine on NaOH) on FeSO4 (see comments by Proteios at https://www.sciencemadness.org/whisper/viewthread.php?tid=1305 ). If confirmed, this reaction bares a strong parallel to the one cited above and is, in line, with my change in 'activity level' argument of the acid salt. [Edit] Some confirmation from the cited thread, page 2, to quote Proteios "from memory the reaction goes best when the FeSO4/ bleach mix is a thick goopy brown mess." Nevertheless, the general lack of discussion in the literature of this apparently observed reaction suggests that the mechanism of the FeSO4/hypochlorite reaction is not generally well understood (except, of course, now by us, or, at least, so we have an understandable and tentative reaction path which appears to more broadly concur with observations).

If using litmus paper is considered a chemical test, 1st test the solutions pH with litmus paper. Boil and re-test. If the solution has become more basic, it is H2SO3. Reaction logic: H2SO3 ---Heat--> SO2 (g) + H2O and the pH approaches that of water. Boiling H2SO4 will not raise its pH. -------------------------------------------------------- Another test, add a large amount of zinc foil. If the Zn dissolves and no fine precipitate, it is ZnSO4 and the acid is H2SO4. If the Zn is only very slowly attacked, and forms a fine precipitate, it is ZnSO3 and the acid is H2SO3.

As I noted previously Mg powder is observed to be inferior, but large metallic Mg surface is needed. So one possible path is a temperature dependent reduction process directly as a result of say MgO on some metal surface. For example, see the full paper "Theoretical study of the decomposition of HCOOH on an MgO(100) surface" at http://www.qcri.or.jp/lab/wp-content/uploads/2011/07/p236.pdf to quote: "It is well known that metal oxides are good catalysts for a variety of chemical processes [1– 26]. For example, methanol, formaldehyde and formic acid readily decompose on MgO catalysts [7,8,10–12]." See also related full paper "Chemical reactivity of oxygen vacancies on the MgO surface: Reactions with CO2, NO2 and metals at: http://www.captura.uchile.cl/bitstream/handle/2250/7121/Florez_Elizabeth.pdf?sequence=1 where the author notes, to quote: "It was proposed that the interaction is dominated by an electrostatic mechanism of electron transfer and that this is strictly connected to the oxygen vacancies." relating to impurities on the MgO surface. See also comments at "Acid-base reactions on model MgO surfaces" at http://link.springer.com/article/10.1007%2FBF00767207#page-2. So, in the current context, it may be that a particular KOR may be reduced forming potassium in a metallic phase principally due to MgO and heating based on my repeated below Wikipedia comment (link: http://en.wikipedia....AlkoxideSection ) to quote: "Many metal alkoxides thermally decompose in the range ~100–300 °C. Depending on process conditions, this thermolysis can afford nanosized powders of oxide or metallic phases. This approach is a basis of processes of fabrication of functional materials intended for aircraft, space, electronic fields, and chemical industry..." Also possible is via chemisportion occurring on the metallic Mg (but not strongly supported in the literature). Now, there is apparently much literature focus, which is great news, for a nano MgO surface reactions (which can be introduced by other cheap routes, perhaps eliminating Mg metal altogether, but replacing with something to remove water and provide support for the MgO, perhaps Al foil and a drop of Iodine, Zn,...). Also, my limited research on the reaction mechanism occurring on the MgO surface is complex via the formation of some intermediaries and I would not be surprised if we actually saw some 'novel' chemistry on the MgO surface itself (as I have see in atmospheric gas/solid reactions at very ambient temperatures, for example, see full paper "Water Chemisorption and Reconstruction of the MgO Surface" at http://arxiv.org/pdf/mtrl-th/9508001.pdf). To quote: "The presence of surface hydroxyl groups on MgO powders exposed to H2O has been demonstrated by infra-red spectroscopy7,9,10. Hydroxyls are clearly distinguishable from physisorbed molecular water by the HOH bending mode which disappears above 100◦C, while the OH stretching mode persists even above 500◦C. Furthermore there is complete monolayer coverage of the surface by hydroxyls, as shown by microgravimetry measurements7 Despite these observations, the most reliable theoretical calculations predict that water molecules do not dissociate on the (001) surface." and: "In summary, water demonstrably chemisorbs onto MgO but trustworthy calculations show that H2O molecules should not dissociate on the only known stable surface." Also, see: http://www.chem.tamu.edu/rgroup/goodman/pdf%20files/343_jcpb_103_99_3391_jp983729r.pdf ---------------------------------------------------------------------------------------------------- Now, with respect to your proposed use of Calcium, here is an interesting reference comparing MgO and CaO catalyst in "Natural Gas Conversion VI" edited by T.H. Fleisch, J.J. Spivey, Enrique Iglesia, page 213. Link: http://books.google.com/books?id=I0SCzBkz7IoC&pg=PA213&lpg=PA213&dq=reduction+of+CH3OH+by+MgO&source=bl&ots=bhoppFUF9m&sig=oFkDYcYkIe2kAlmz8Q8qFJy3IMk&hl=en&sa=X&ei=vKl6UdHiAsfj4AO80YHYCA&ved=0CDwQ6AEwAw#v=onepage&q=reduction%20of%20CH3OH%20by%20MgO&f=false --------------------------------------------------------------------------------------------------- With respect to cutting the relative Mg use in half, I am not sure if that is the best test, but here is the logic: Non-aqueous reaction cutting relative Mg input: 2 KOH + 2 ROH --> 2 ROK + 2H2O Mg + 2 H2O --KOH--> MgO + H2O + H2* (g) 2 ROK + H2* + Heat --MgO--> 2 ROH + 2 K (s) Net: 2 KOH + Mg ---> MgO + 2 K + H2O + H2* --> Potassium decomposition/reduced yield on water contact. So this equation's indicated amount of Mg is 1/2 of the available # of moles of KOH, or actually ROH if it is less than KOH per the 1st equation in the chain. This amount of Mg is then added to the amount of Mg required, per below, to remove all water from the aqueous first phase: Mg + H2O --KOH--> MgO + H2* (g) where I would assume 20% of the KOH, for example, is actually water. The hypothesis is that this total amount of Mg is insufficient for a successful Potassium production run. Interestingly, note that in the usual preparation of Potassium having too much water initially with sufficient Mg around is not necessarily a bad thing as it just produces more MgO (which I claim may have a positive role in the non-aqueous phase).

Actually and unfortunately, perhaps what I suggested isn't a story after all, more of a re-wording (or a re-write as it were). Stating (as per Wikipedia) ""Many metal alkoxides thermally decompose in the range ~100–300 °C. Depending on process conditions", means to me wrong temperature and process conditions and your synthesis of KOR is toast. And to further confuse the hapless would be chemist, he is presented with some potassium (or, perhaps the oxide). Solution: just re-write the failed preparation as a patentably path to K (or whatever) instead. The only major issue, of course, is that you are most likely clueless as to exactly the how and why, and thus are purposefully vague as to reaction mechanics, or worse, make-up a replacement reaction as a potential explanation. Sound familar? Funny, but only until someone really takes the reaction mechanism seriously.

Elementcollector1: There is no 'nascent hydrogen' mentioned here. Activated H2 via chemisportion, and its many associated recents references (I only gave two) are easy found in the same very highly regarded journals I cited (note the copyright from the American Chemical Society). So forget all the non-truths to support a baseless position (asserted absolutely, no less, despite numerous possible chain reaction failings and, even more problematic, that proposition cannot explain a single experimental observed particularity). Now, the open question in this Na forum, is not the formation of Sodium without electrolysis (that has been done), but can it be performed safely. Now, assume you have a reaction preparation that does form K at ambient (or safe temperatures). Does adding NaCl, mean that you now have a safe path to Sodium? Or, can you just use NaOH in place of KOH? My understanding of the reaction path for K, still being explored and apparently debated, indicate paths for which this would not be favored. So, back to path exploration, here are the results of more research, and I would appreciate some honest feedback (but fear of that moderator may inhibit your honesty, I understand): First, per Wikipedia (link: http://en.wikipedia.org/wiki/AlkoxideSection ), under heading "Thermal stability" of metal alkoxides, to quote: "Many metal alkoxides thermally decompose in the range ~100–300 °C. Depending on process conditions, this thermolysis can afford nanosized powders of oxide or metallic phases. This approach is a basis of processes of fabrication of functional materials intended for aircraft, space, electronic fields, and chemical industry: individual oxides, their solid solutions, complex oxides, powders of metals and alloys active towards sintering." Now, this is applicable to both Na and K metal alkoxides. Also, in a recent thread at ScienceMadness, Nicodem warned Blogfast about stability issues asociated with alkoxides per his personal experience. However, the possible decomposition products cited here by Wikipedia are of particular interest, including a metallic phase (the liberation of metallic Potassium?) and/or nanosized K2O. This is in further support ot the reverse formation proposition of KOR to K that I have proposed previously. Next, in the particular case of nanosized Potassium oxide formation, K2O (from the decomposition of KOR) could be attacked with H2 via chemisorption (on nanostructured MgO, see, for example, "H2 chemisorption and consecutive UV stimulated surface reactions on nanostructured MgO", in Phys. Chem. Chem. Phys., 1999,1, 713-721., see also "Theoretical aspects of H2 and CO chemisorption on MgO surfaces", Surface Science (May 1982), 117 (1-3), pg. 571-580.) This may not be the case for Na2O as there is no supporting literature on chemisorption. But, even disgarding this path, there is still a possible hydrogenation reaction, although occurring rarely (per Wikipedia, link: http://en.wikipedia.org/wiki/Hydrogenation) for reactions below 480 °C between H2 and organic compounds in the absence of metal catalysts. However, with respect to rare exceptions, Wikipedia also states under the topic "Metal-free Hydrogenation", that to quote: "Hydrogenation can, however, proceed from some hydrogen donors without catalysts, illustrative hydrogen donors being diimide and aluminium isopropoxide. Some metal-free catalytic systems have been investigated in academic research. One such system for reduction of ketones consists of tert-butanol and potassium tert-butoxide and very high temperatures.[24]". I would observe that perhaps this is related to the formation of K2O, which is apparently employed in mixed oxide catalyst for hydrogenation (see, for example, "Promotion effect of K2O and MnO additives on the selective production of light alkenes via syngas over Fe/silicalite-2 catalysts"). This path is particular to K and, as such, excludes the formation of Sodium again. In any event, my opinion is that the simple thermal decompostion of potassium tert-butoxide (simple and perhaps now the best explanation, and also for the Na salt), by itself, or in the presence of activated H2 per chemisportion on MgO, or via hydrogenation in the presence of K2O, may be some alternate explanations with peer reviewed references. Now, I want to be fair, some more research on the direct reaction of Mg with KOR (or NaOR), and there is nothing supporting the direct replacement reaction by Mg as proposed liberating K (or Na). For example, see "Reactions of Magnesium and Titanium Alkoxides. Preparation and Characterization of Alkoxy-Derived Magnesium Titanate Powders and Ceramics". To quote the entire abstract: "Abstract The interaction between magnesium and titanium alkoxides is studied in order to chose the best precursors for synthesis of MgTiO3. No reaction between magnesium and titanium methoxides and isopropoxides occurs. The solubility diagrams for Mg(OR)2-Ti(OR)4-ROH, R = Et,-Bu at 20°C are studied. Magnesium ethoxotitanates of variable composition MgnTi4-n (OEt)16-2nċ2nEtOH (n=2.0-0) which are structural analogs of Ti4(OR)16 (R = Me, Et) are isolated. This is a quite unusual example of statistical distribution of heteroatoms in molecular structures of metal alkoxides. Among the systems of metal alkoxides with simple aliphatic radicals only Mg(OBu)2-Ti(OBu)4-BuOH gives a convenient precursor for the synthesis of MgTiO3. A simple scheme of preparation of magnesium titanate from the alkoxide solutions is suggested. The phase purity of MgTiO3 is to a considerable extent dependent on the hydrolysis conditions. The alkoxy-derived magnesium titanate is obtained in the form of a uniform fine powder, it can be sintered into dense ceramics in the temperature range of 1140–1220°C which is 150–200°C lower in comparison with the conventional powders" Link: http://link.springer.com/article/10.1023%2FA%3A1008616329847... Interestingly here, when products are produced, mixed salts (and not deposits of titanium) occurred. Other searches on Springer's articles also only produced references to oxide formation reactions. With time, and observations, my basic argument appears only to get more supporting paths and, so far, a total lack of rationale for the other side at the very ambient temperature recommended for this reaction. But, if I have missed something, please cited it, and unlike the others pushing their position, I welcome open supported discussion, which so far your comments are lacking, both in accuracy (no nascent H2 here, just well researched paths to activated hydrogen in journals of physical chemistry), and no apparent attempt to educate yourself by researching and citing references. Remember, good science always wins in the end. -------------------------------------------------------------------------------------------- [EDIT] I have recently come to a somewhat uncomplimentary view, based on my research, as to how this Potassium reaction was discovered. It owes it genesis, I suspect, to a failed preparation of KOR. The KOR is formed, and subsequently, perhaps unexpectingly, decomposes, liberating K on occasion. Is this translatable to Na? Wow, what a revelation if I am correct!

OK, burning a Mg ribbon in O2 (caution, very hot reaction temperature and blinding light) is an example of a highly exothermic reaction with a very stable compound Magnesium oxide, MgO, forming. In air, mention that Mg3N2 is also formed, which liberates NH3 with water (a possible fuel source). 3 Mg + N2 + 461 kJ/mol --> Mg3N2 Mg + 1/2 O2 + 601 kJ/mol --> MgO Mg3N2 + 6 H2O ---> 3 Mg(OH)2 (s) + 2 NH3 (g) Reacting purple Iodine vapors (from gently heating I2 at one end of a flask) with NH3 fumes, should procedure a red smoke in a flask. This is an example of an endothermic reaction (show condensation). It forms a very unstable and highly sensitive novelty explosive 8 NI3 · NH3 (only demo in very small amounts). 24 I2 (s) + 9 NH3 (g) + Heat → (8NI3. NH3) (s) + 24 HI (g) 2 NI3 (s) → N2 (g) + 3 I2 (g) (–290 kJ/mol) Dry salt deconposing: 8 NI3 · NH3 → 5 N2 + 6 NH4I + 9 I2 If the flash, bang, all the equations and talk of green fuel does get the judges all excited, I don't know show business.

..... -------------------------------------------------------------------------- Quote from Woelen: "Try the following to make your FeCl3 work: Use a very concentrated solution, or add a lot of NaCl also. The chloride concentration needs to be really high. The complex formation is one of the main driving forces behind the reaction and that complex only is formed at very high concentration of Cl(-). If your solution is very turbid, then you probably have impure FeCl3 with a lot of basic Fe(3+) compounds, e.g. FeCl2(OH), FeCl(OH)2 and Fe(OH)3 and all kinds of hydrated forms of that. If this is the case, then add a few drops of HCl, until the solution is clear." -------------------------------------------------------------------------- My comment is that having a very concentrated solution of FeCl3 with added NaCl, means to me that in the hydrolysis of FeCl3: 2 FeCl3 + 6 H2O <---> 6 HCl + 2 Fe(OH)3 one has changed the 'activity level' of the normally dilute (and weak) HCl to a much stronger acting acid. Hence the witnessed increase in etching behavior.

Before one separates the H2O2 out of Na2CO3.H2O2, sometimes the Na2CO3 can act like a base catalyst (for example, in the oxidation of ammonia to nitrite), so don't bother to remove it, use as is and keep it in its more stable form. Also, l suspect given the sizzle (when one adds water to commercial oxygen bleaches), one really cannot extract the conc H2O2, as I would guess it is preset to decompose, and release O2 on contact with water, per its bleaching role.

To quote from a reputable authority: "In the past, quite a few great chemists were severely poisoned, not by HNO3, H2SO4 or HCl, but by HF. Some to name: Davies, Moissan (first person who made elemental fluorine). Apparently, some of these great chemists have suffered great pains from fluoride poisoning." On dimethyl mercury, Wikipedia has an article about a chemist who died from just a few drops on her hand a couple of months after exposure. Wow! Apparently, Karen Wetterhahn died a horrific death from dimethyl mercury. Always research chemicals before you use or produce them (start with Wikipedia and MSDSs). Evaluate the risk of your exposure and/or to a third party. If possible, ask the opinion of a professional chemists (they know that some chemicals are just so dangerous they should never be made or employed - seek substitutes).

For those who still believe in the reactions posted previously on the road to K and Na: KOR + Mg ---> 2 K + Mg(OR)2 Mg(OR)2 + H2O --> MgO + 2 ROH Please explain to me why Mg(OR)2 gets to react with water before Potassium, and even if, it is a one pot reaction, there still should still be something observed. In other words, this accepted reaction is ignored entirely: 2 K + H2O --> 2 KOH + H2 (g) because, if you accepted this reaction occurs at all, you would have more Hydrogen than observed, less K and some Mg(OR)2 lying around, but not reported. And why is the reverse reaction not occurring, namely: 2 K + Mg(OR)2 --> 2 KOR + Mg More bad news is a search for verification uncovered the following,"Reactions of Magnesium and Titanium Alkoxides. Preparation and Characterization of Alkoxy-Derived Magnesium Titanate Powders and Ceramics". To quote the complete abstract: "The interaction between magnesium and titanium alkoxides is studied in order to chose the best precursors for synthesis of MgTiO3. No reaction between magnesium and titanium methoxides and isopropoxides occurs. The solubility diagrams for Mg(OR)2-Ti(OR)4-ROH, R = Et,-Bu at 20°C are studied. Magnesium ethoxotitanates of variable composition MgnTi4-n (OEt)16-2nċ2nEtOH (n=2.0-0) which are structural analogs of Ti4(OR)16 (R = Me, Et) are isolated. This is a quite unusual example of statistical distribution of heteroatoms in molecular structures of metal alkoxides. Among the systems of metal alkoxides with simple aliphatic radicals only Mg(OBu)2-Ti(OBu)4-BuOH gives a convenient precursor for the synthesis of MgTiO3. A simple scheme of preparation of magnesium titanate from the alkoxide solutions is suggested. The phase purity of MgTiO3 is to a considerable extent dependent on the hydrolysis conditions. The alkoxy-derived magnesium titanate is obtained in the form of a uniform fine powder, it can be sintered into dense ceramics in the temperature range of 1140–1220°C which is 150–200°C lower in comparison with the conventional powders." Note, the second sentence, 'No reaction between magnesium and titanium methoxides and isopropoxides occurs", and further, when they do react, double salts and no titanium metal precipitation. So barring a reaction with a higher alcohol with a very weak bond, the prospective of the professed reaction really needs a good source. The bottom line is, you just can't make up bad stuff, it just gets worse for you. ---------------------------------------------------------------------------------------------------------- The good news is that if you reject this increasing unlikely reaction, and even partially accept my path, there is hope in replacing Mg in a ambient temperature synthesis of K and/or Na. In fact, perhaps, Aluminum foil (contains Si and Fe impurities which are good here perhap activated with a drop of Iodine) together with dry MgO from say: MgSO4 + 2 NH3 + 2 H2O --> Mg(OH)2 + (NH4)2SO4 Mg(OH)2 --Heat--> MgO + H2O (g) may even work! While the particular ROH employed is still a problem, but more research may suggests a more accessible and cheaper substitute. Power to Chemistry!

If you have metallic K, then just add NaCl making Sodium in situ (of course, still doing this safely, is the issue). But, assuming one could address this problem, how would one safely make K? Someone mentioned Sciencemadness which is currently running a very long thread on making Potassium at much safer ambient temperatures via Mg turnings and KOH and (for this forum as to not violate rules with respect to recipes for dangerous substances like metallic K) called the alcohol ROH. This could also give insight as to a low temperature Sodium preparation. Since the Potassium reaction temperature is around 300 C (correct me, it is a 53 page thread), my take on the reaction mechanism: 1. Aqueous phase since as much as 27% of KOH could be water: Mg + 2 KOH + H2O --> [K2MgO2 + H2O] + H2* (g) K2MgO2 + H2O <---> 2 KOH + MgO Net: Mg + H2O --KOH--> MgO + H2* (g) 2. Non-aqueous reaction: KOH + ROH --> ROK + H2O Mg + H2O --KOH--> MgO + H2* (g) ROK + 1/2 H2* <---> ROH + K (s) Net: KOH + Mg ---> MgO + K (s) + 1/2 H2* ------------------------------------------------------ Comments: Now, Wikipedia actually cites the laboratory preparation for KOR from the un-named alcohol as follows: K + ROH → ROK + 1/2 H2 (g) + Heat where ROK species is, itself, noted as being a strong, non-nucleophilic base in organic chemistry. So, what I am postulating here is that activated hydrogen is formed (most likely via chemisorption, discussed more below) and that with excess active H2* (from the initial dehydration), pressure (balloon employment), and heat applied to the reaction chamber, that the ROK formation reaction, cited above, is reversed to some extent releasing K (in agreement with Le Chatelier's principle to remove stresses relating to temperature, pressure and/ or concentrations): ROK + 1/2 H2* (in excess) + Heat (applied) ---> ROH + K (s) Now, more on this chemisorbed hydrogen most likely created via the presence of MgO. Here is an abstract (source: "H2 chemisorption and consecutive UV stimulated surface reactions on nanostructured MgO", in Phys. Chem. Chem. Phys., 1999,1, 713-721. To quote from the abstract: "MgO nanoparticles obtained by chemical vapour deposition (CVD) were exposed to H2 and subsequently to UV irradiation and/or molecular oxygen at room temperature. A combined IR/EPR study reveals the role of low coordinated surface sites and anion vacancies in the diverse reaction steps. The hydride groups emerging from the initial H2 chemisorption processes (heterolytic splitting) play an active role in the consecutive reactions. They provide the electrons which are required for the UV induced formation of surface colour centres and for the production of superoxide anions (redox reaction). Both the colour centres and the superoxide anions are EPR active. The hydroxy groups resulting from H2 chemisorption do not actively participate in the consecutive reactions. Together with the OH groups formed in the course of colour centre formation they rather play the role of an observer. They undergo specific electronic interactions with both the colour centre and the superoxide anion which are IR inactive (or IR inaccessible) surface species. They may, however, be observed by IR spectroscopy via the specifically influenced OH stretching vibrations. This proves the intimate interplay between IR and EPR spectroscopy as applied to the surface processes under investigation. As a result, two paths were found for the three consecutive surface reaction steps: H2 chemisorption, colour centre formation and superoxide anion formation. In the first one a single, well defined surface area element is involved, namely a low coordinated ion pair, the cation of which is a constituent of an anion vacancy. In the second path a diffusion controlled intermediate step has to be adopted in which the electron required for the colour centre is transported by an H atom travelling from a hydride group to a remote anion vacancy. In either case there is clear experimental evidence that the finally resulting superoxide anions are complexed by the colour centre cations." See also "Theoretical aspects of H2 and CO chemisorption on MgO surfaces", Surface Science (May 1982), 117 (1-3), pg. 571-580, to quote from the abstract: "Preliminary ab initio calculations at the SCF level and beyond are reported for the chemisorption of H2 and CO at the (001) surface of MgO. It is concluded that the dissociative chemisorption of H2 requires the presence of defects and that at anion vacancies, V− centres and self-trapped holes the overall process is exothermic in each case. It is predicted to be non-activated at anion vacancies and possibly the same at the other two defects. Binding energies are calculated for the interaction of CO with a non-defective (001) surface of MgO and at impurity ions therein. They range from 2.5 kcal/mole at Al3+ to 20.8 kcal/mole at Cu2+ and are shown to be highly sensitive to lattice relaxation of the defective surface." where there is an interesting reference to the role of impurities in the MgO. Now, there are also so studies citing the reaction of between hydrogen and magnesium, but mostly as fine Mg powder (or nano). However, Mg surface attacked by KOH, may be more amiable to H2. Probably, an important speculation is that absence MgO, no or reduced chemisorbed H2 formation, no reduction reaction and no K is produced! Also, less than completely pure Mg, KOH or KOH, could produce detective surfaces on the MgO, increasing yield. I would speculate that MgO dust on Mg turnings may provide a good contact point for gaseous H2 and, with infrequent stirring to add ROK, help to form potassium. Neither frequent or very infrequent (in agreement with the patent instructions) would be advisable. Now why is Mg dust no good for this reaction? My speculations, first, the reaction rate would be too fast (also more heat) and limit gaseous contact. Also, absence the Mg turnings, less support for formed MgO thereby limiting the H2 contact. In closing for disclosure, the K creation is based on a patent and a questionable (at least in my opinon, but not Sciencemadness's moderator) reaction chain, depicted very differently from my opinion: Mg + 2 KOH + H2O --> [K2MgO2 + H2O] + H2 (g) K2MgO2 + H2O <---> 2 KOH + MgO Net Aqueous phase (same reactions): Mg + H2O ---KOH---> MgO + H2 (g) Non-Aqueous: 2 KOH + 2 ROH ---> 2 KOR + 2 H2O Mg + H2O ---KOH---> MgO + H2 (g) 2 KOR + Mg ---> 2 K + Mg(OR)2 Mg(OR)2 + H2O --> MgO + 2 ROH Net reaction: 2 KOH + 2 Mg ---> 2 MgO + 2 K(s) + H2 or: KOH + Mg ---> MgO + K (s) + 1/2 H2 where one of the main disagreement is that Magnesium shavings replaces Potassium in KOR: 2 KOR + Mg ---> 2 K + Mg(OR)2 occurring, no less, at ambient temperature (under 350 C, with K boiling at twice this) with no apparent rationale based on Le Chatelier's principle, as far as I can see. Observations suggest this is not, in effect, in the reaction chain as it has been observed that fine Mg powder does in fact produce a much lower yield (which should support, not detract, from this reaction's effectiveness), while the same observation as to the best Mg size supports the concept of surface formation on Mg for MgO to increase yield. This same observation generally suggests a more complex chemical and physical role for the Mg metal itself. With respect to documentation of this reaction, other than a patent's quasi unbalanced reaction, there is no, not even one, reputable source for this questioned reaction anyway on the internet or elsewhere. What do you think? [EDIT] For the solid to solid reaction, I conceive that relative lattice energies and that at high tempeatures (over 1,000 C), the more volatile nature of either Na or K versus Mg, may, per Le Chatelier's principle explain Mg replacing Potassium (or Sodium), but this is not as strong an argument at a lower temperature. Also, my reaction dynamics fits many of the observed particularies (like upon changing the size of Mg employed), negative influence of faster reaction rate (limiting H2 contact), role of temperature in the reaction, best stirring frequency, and why the preparation is generally problematic (seen in select hydrogenation reactions). Also, this should not be described solely as a hydrogenation reaction, or solely as a Magnesium (or MgO) reduction reaction, but as a dissociative chemisorption of H2 activated through a Mg/Mg complex formation on a MgO surface.

If the H2SO4 is pure and concentrated, react directly with the NaNO3 and distill off to get pure HNO3: H2SO4 + 2 NaNO3 --> Na2SO4 + 2 HNO3 ------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------ If the H2SO4 is impure or not concentrated, start by making FeSO4 by dissolving some iron in H2SO4. Caution: Hydrogen gas is generated. H2SO4 + Fe --> FeSO4 + H2 Now add to this solution NaNO3 and more H2SO4 per the equation: 6 FeSO4 + 2 NaNO3 + 4 H2SO4---->4 H2O + 2 NO + Na2SO4 + 3 Fe2(SO4)3 This is the so called ring test for Nitric acid. Now that we have NO, there are many paths to HNO3. Examples: 1. Capture the NO gas and mix with air to form NO2: NO + 1/2 O2 --> NO2 2. React the NO2 in a series of water towers (this is the commercial process): 3 NO2 + H2O --> 2 HNO3 + NO Alternate process, dissolve the NO2 in H2O2: 2 NO2 + H2O2 -->2 HNO3 Alternate process: The action of the strong oxidizer HClO on NO: 3 NO + 3 HClO + H2O --> 2 HNO3 +3 HCl + NO Good Luck.

Burn Aluminum and dissolve in pure NH4OH (impure would be to use household ammonia, so heat household ammonia and send the NH3 into distilled H2O to make pure NH4OH). Avoid using an excess of NH4OH as the solution could become clear with an Ammonia Aluminum complex. Otherwise, you should see whitish gelatinous Al(OH)3 and some undissolved residue from the burned Al (most likely silicon impurities). Extract only the pure gelatinous Al(OH)3 (more easily said than done) form the Si residue To this extract add distilled water and CO2. Heat until all the fluid is evaporated. Repeat adding distilled water and CO2. If there is a new residue, this could be either Iron, Magnesium or Manganese Carbonate. In Reynold's wrap Aluminum foil, for example, it is 98% Al and the rest is Si and Fe. Other sources of Al, like Al cans, have Mg and Mn impurities. Hopefully this process of carbonation will separate out these other impurities. Note, Aluminum carbonate is highly unstable decomposing into CO2 and Al(OH)3 upon heating. Heat the final product to yield pure Al2O3 (or dehydrated the Al(OH)3 gel using ethanol). Caution, the process of dissolving of the Al2O3 (and.or the Al) in the NH4OH produces hydrogen gas. Note, the Al2O3 produced from burning Al in air (no not matter how thoroughly you heat it) contains some unreactive Al and Aluminum Nitride. Also, AlN slowly dissolves in H2O to produce NH3. This is why you should dissolve the burned Al in pure NH4OH to separate and purify it. Using Al only also works, but is dissolves more slowly due to possible acrylic & oil coating, and processing techniques to make the protective Al2O3 layer more effective (why?, in my opinion the Al foil manufacturers want to reduce the possible leaching of Al into our food given the developing understanding of the magnitude/or lack thereof of Al toxicity and the potential detrimental, if any, effect on their industry). However, per one Al coating maker, vinegar penetrates the defenses and (I have done this) soaking Al foil in acetic acid for several hours, even though it appears unreactive, does make the Al foil dissolve more rapidly in NH4OH. I probably missed some points in the preparation, so please comment.

A recipe that I have not tried so I cannot state that you will make something useful is as follows: 1. Prepare fresh HClO. One way to prepare is by adding vinegar to bleach (NaClO). Boil this dilute solution of HClO (and sodium acetate) and capture vapor and gas (HClO and Cl2O, the gaseous anhydride of HClO) in distilled water. Only boil off half of the solution as this will capture 5/6 of the total available HClO in solution. Note, with time concentrated HClO disproportionates into HCl and HClO3, and also on heating at 70 C or higher. 2. Gentle heat Sulfur in HClO (even dilute HClO should work per Watt's Dictionary of Chemistry Vol 2). This will make H2SO4 and HCl. You should be able to buy sulfur in a plant and garden store as "Flowers of Sulfur", which should be washed in distilled water prior to use to remove any organic wetting agents. If solution is too dilute for your purpose, carefully distill to concentrate. Note, substituting phosphorous for sulfur, per Watt's, makes phosphoric acid. Using Iodine makes Iodic acid, etc., as HClO is a powerful oxidizer. Good luck.

My speculation is that someone added excess ammonia to silver nitrate. This produces ammonium nitrate and a soluble silver hydroxide ammonium complex. The problem is that the latter salt can decompose with age to form silver nitride. This compound is so sensitive that crystals in solution can set off an explosion, and has a history of causing injuries. The evacuation was likely just a precaution against a known potential hazard that could have produced a large liability suit against the university if it failed to act in a prudent manner and, in my opinion, it was not likely that a significant part of the university would have been consumed in an explosion, just a financial disaster.

An interesting idea is to replace CaCO3 with Na2CO3, and the reaction directly producing Na2O: 4 H2 + Na2CO3 --> CH4 + Na2O + 2 H2O The reaction temperature is also possibly in the range of 500C as well. As a word of caution, the auto-ignition temperature for H2 is varyingly reported as between 500C to 585C (other values 536C, 560C and 565.5C) with a heated glass vessel yielding the lower value, jet streams and heated wire somewhat higher. Also, the lower flammability level for H2 is unusually low at 4% (so 4 parts of H2 in 96 parts of air may auto-ignite at around 500C depending on the nature of the heat source) and the upper flammability limit is 75. The bottom line, the proximity of the reaction temperature to H2 auto-ignition threshold and the wide range in H2 flammability limits are dangers to be addressed in attempting this experiment.

Per Wikipedia (link below), the following reaction occurs at only 500C and may provide a more accessible way for some to make CaO from CaCO3 for use in the preparation of NaOH: 4 H2 + CaCO3 --> CH4 + CaO + 2 H2O That is, heat Calcium carbonate to 500C in an atmosphere of hydrogen. Note, in the same Wikipedia article, the cited temperature decomposition of CaCO3 is 1,500C!!! Caco3 --> CO2 + CaO One reason given on the web for the wide range in temperature ranges for the decomposition of CaCO3 is that there is a significant difference between the "onset" temperature and the "practical" temperature where a significant yield can be obtained. http://en.wikipedia.org/wiki/Abiogenic_petroleum_origin