ajkoer

To quote from Wikipedia (http://en.m.wikipedia.org/wiki/Water_of_crystallisation#undefined ): "Classically, "water of crystallization" refers to water that is found in the crystalline framework of a metal complex or a salt, which is not directly bonded to the metal cation." So per Wikipedia, the water molecule in the formed crystal structure that are not directly bonded to metal cation (usually containing a +2 and +3 cations as well as a −2 anions) are usually termed "water of crystallization". Now, in the case of copper(II) sulfate, it is more actually represented as [Cu(H2O)4]SO4·H2O. So, only one molecular of water is technically a "water of crystallization", the other four are coordinated to the Copper cation. For Glauber's salt, Na2SO4(H2O)10, there is technically no water of crystallization. For Epsom salts, I recall seeing it express as MgSO4(H2O)6.H2O implying a single water of crystallization. In fact, for reason of accessibility most likely, if you need to extract water from say isopropyl alcohol, to make it drier, using MgSO4 is often recommended (I have done it with some success). If I am correct here, the recovered hydrated salt after being dried and heated, could liberate some alcohol per the single water of crystallization, which I would describe as in effect, solution water (a mix of water and alcohol, in the current case).

Here is an idea based on the reaction of the hydroxyl radical with graphite. Per this source "The reaction of hydroxyl radicals with carbon at 298 K" by M.F.R. Mulcahy, B.C. Young, to quote from the abstract: "The reaction of free OH radicals with graphite was studied in a flow system by mass spectrometry, the OH being produced by the reaction H + NO2 → OH + NO. The OH radicals react rapidly at 298 K to produce approximately equal amounts of CO and CO2." Link: http://www.sciencedirect.com/science/article/pii/0008622375902687 The reactions, involving radical species, appear to proceed as follows: C + OH → CO + H (Source: see http://en.m.wikipedia.org/wiki/Hydroxyl_radical CO + OH → CO2 + H (Source: see http://www.sciencedirect.com/science/article/pii/S0082078475803432 ) H + OH + M → H2O + M (Source: see https://www.google.com/url?sa=t&source=web&rct=j&ei=aGAwVPW-OsrlsAS-rII4&url=http://www.nist.gov/data/PDFfiles/jpcrd9.pdf&cd=6&ved=0CDMQFjAF&usg=AFQjCNFsRKYlwzIasvags7FxSiEgsZdq2w&sig2=1grL8k3UWaTU9-wIclokfQ ) I would, in your case, create the OH radicals via the photolysis of an aqueous nitrate (as a source, see, for example, the abstract at http://www.ncbi.nlm.nih.gov/pubmed/22819875 ), like KNO3, in the presence of lime water where the formation of CO2 produces a precipitate of CaCO3. With time and uv exposure, a gas built-up could occur (CO, CO2). Cool and shake the reaction vessel and you may be able to witness a cloudy suspension of CaCO3 being formed and a reduction in volume. The magnitude of CaCO3 produced is the indicator test for the presence of Carbon. To test for the presence of Carbon monoxide, open the vessel and extract the remaining gas, add O2 and heat. Repeat test for CO2 per above. The hydroxyl radical appears, in practice, to be a powerful agent in the enviromental remediation of toxic organic compounds, which puts a green bent on your science project.

The starting materials are a copper alloy, air, moisture and NH4NO3. An electrochemical reaction proceeds with the reduction of the NH4NO3 to NH4NO2 (which is very poisonous and unstable usually forming N2, but the dry salt can explode, especially in the presence of impurities) and then further reduce to NH3. Interestingly, you did actually detect the smell of ammonia. Copper is a known sensitizer for ammonium nitrate (that is, NH4NO3 is now transformed into a sensitive explosive, probably due to the creation of NH4NO2 and at higher temperatures, other unstable intermediaries, so be very very wary). I would suspect a copper alloy is even more problematic due to the possible mixed metal electrochemical based REDOX reaction alluded to above. I have actually experience with the REDOX reaction of aqueous KNO3 with Aluminum foil, where a similar electrochemical reaction ensues. I avoided using an excess of the Al as I wanted just KNO2, but I removed some of the nitrite solution for testing. As a consequence of the now relative increase in Al, the next day I was surprised when greeted by a rush of escaping NH3 gas upon opening the vessel with the remainder of the once aqueous KNO3. [Edit] If a mixed salt is formed with ammonia, it could be, indeed, Cu(NH3)4(NO3)2, the later having been actually reported as forming in nitrate ammunitions upon warming in contact with corroding Copper (this could be galvanic corrosion from the action of NH3, H2O and O2 on Cu producing cupric ammonium hydroxide). The cupric ammonium nitrate could then be formed via the dissociation of NH4NO3 into NH3 and HNO3 upon mild heating (similar to what happens on heating NH4Cl) acting on the Cu(NH3)4(OH)2.

John: I agree on your comment on the reaction cited and have edited my comments to be more clear that the author is, in my opinion, referring to an electrochemical reaction using non-standard notation for that field, which certainly can be confusing and viewed, to quote you, as "utter garbage". But, your comment on the important of concentration is similarly misplaced, to the best of my understanding, in an electrochemical setting. That is why an insignificant amount of HOCl, found normally in Bleach, is still significant. The other less obvious point is that a low concentration, but still significant reaction, over time can be, in the long term, damaging/corrosive to an anode. Another point I should make is that in the references on these electrochemical reactions I have provided above, there is a large number of possible side reactions in addition to the major electrochemical reaction. My nitrate reduction example is, in my opinion, but one possible minor reaction scenario with potentially disastrous consequences. Finally, the degrading of my analysis on this mystery to the category of speculation is NOT appropriate in a legal sense, in my opinion. My opening comments alluded to an authoritative report (a government commissioned investigation) on one of the exploding water tank occurrences. The first page of that report clearly alludes to the re-finishing of the water tank with a zinc coating as potential causative factor. While it has taken me some time, the commission's comments to me now suggest, more likely than not, an electrochemical based explanation. As such, my latest path is clearly within this legally endorsed scenario with suggested (and documented) electrochemical paths. More interesting in another field, and pure speculation by a non-lawyer myself, but could a derogatory comment to another espousing a legally endorsed position (by a respected government authority no less) be considered libelous? My argument being there may be an existing basis for legal validity with respect to a general framework (electrochemical), which therefore could provide a frame for establishing the untruth of a contradictory statement which dismisses, in its entirety, said viewpoint. Add a pinch of ridicule and we are off to the courts, ... just joking, but I would like a real lawyer's comment anyway, as I find this stuff fascinating, and useful on how to behave on forums and the like in the internet age.

John: Note my quote from the author: "Even though there is relatively little HOCl in bleach" the reaction with HOCl apparently proceeds nevertheless because of its greater electrial potential. Now, so called 'free-chlorine' (defined as Hypochlorous acid and hypochlorite ion) is in all of our drinking water unless you water is made safe via ozone or chloramine. Interestingly, employing higher chlorination levels is one way of addressing issues with declining water purity, so my suggested reaction path may be observed even more frequently in Aluminum alloy tanks with exposure to copper. It also makes the water a better electrolye for such electrochemical reactions. Unfortunately, if my hypothesiis is correct, it is supportive for the use of NH2Cl in place of Cl2 or NaOCl or ClO2. However, it also argues for ozone if there are a large number of these Aluminum alloy water tanks. --------------------------------------------------------------------------------------------------------------------- Fortunately for those fans of my prior speculations into the possible contributory nature of nitrates in our water to the problem of erupting/exploding water tanks, there is ,good news. Here is some of the chemical reactions associated with Aluminum (and Al tanks) after the protective Al2O3 has been penetrated (see http://en.wikibooks.org/wiki/Inorganic_Chemistry/Qualitative_Analysis/Tests_for_anions ) to quote: "1. 3NO3- + 2Al + 3H2O → 3NO-2 + 2Al(OH)3 2. NO-2 + 2Al + 3H2O → NH3 + 2Al(OH)3 + OH- 3. 2NO-2 + 2Al + 4H2O → N2 + 2Al(OH)3 +2OH- Nitrate reduction was found to be pH dependant. At pH values less than eight no nitrate reduction took place. Above pH 10.5 nitrate was reduced upon addition of the aluminium powder. Aluminium powder has been suggested for the denitrification of sodium-based nuclear wastes, employing the nitrate to ammonia and ceramic (NAC) process (Mattus et al., 1993, 1994)." Here are encouraging comments from another source (see http://en.wikibooks.org/wiki/Inorganic_Chemistry/Qualitative_Analysis/Tests_for_anions ): "The Nitrate ion can easily be reduced to ammonia with either Devardas Alloy or Aluminium Foil. The aluminium is a very powerful reducing agent, and this combined with heating causes the nitrate ions to form ammonia gas. This can be tested for by holding a piece of damp red litmus paper over the end of the test tube. The ammonia will form alkaline ammonium ions in the water and turn the paper blue. 4NO3-(aq) + 6H20(l) -> 4NH3(g) + 9O2(g) Aluminium powder is not shown as it merely catalyses the reaction." In my opinion, the author's last quoted reaction above must be more clearly presented and interpreted in an electrochemical setting. Interestingly, Devarda's alloy is an alloy of aluminium (44% – 46%), copper (49% – 51%) and zinc. Also, again the Aluminum, acting as the anode, is corroded to Al(OH)3 in an electrochemical reaction.

OK, lets assume there is not ever a high concentration of any compounds to form a gaseous eruption. Also, your proposal of the action of Zn and water to form H2 is not likely given the water temperature. However, I believe your idea is close. I was recently exploring (seemingly unrelated to this topic) a hypochlorite based Aluminum and Copper battery. Here is the chemistry per my research. To quote (see http://www.exo.net/~pauld/saltwater/ and http://sci-toys.com/scitoys/scitoys/echem/batteries/batteries.html ): "In the bleach battery, sodium hypochorite (NaOCl), the major constituent in bleach, and hypochorous acid (HOCl), a minor constituent, are reduced, according to equations 6 and 7: (6) ClO- + H2O + 2 e- --> Cl- + 2 OH- Eo = 0.89V (7) HOCl + H+ + e --> 1/2 Cl2(g) + H2O Eo = 1.3V The possible reactions involving aluminum are given by equations 8 and 9: (8) 3 ClO- + 2 Al + 2 OH- + 3 H2O --> 2 Al(OH)4- + 3 Cl- Eo net = 3.21 V (9) 3 HOCl + Al --> Al(OH)3 (s) + 3/2 Cl2 (g) Eo net = 3.93 V " The author also notes that "Even though there is relatively little HOCl in bleach, the latter reaction is more favored because of its large potential of 3.93 volts. Over time, you will see both green cupric hydroxide (Cu(OH)2) and black cupric oxide (CuO). Black CuO is formed from green Cu(OH)2 by the loss of water, which happens over time." The implied reactions at anode: Al + 3OH- ⇒ Al(OH)3 + 3e- given reaction at cathode (copper); 3 HOCl + 3 H3O+ + 3 e- --> 3/2 Cl2(g) + 3 H2O for an implied net of: 3 HOCl + Al --> Al(OH)3 (s) + 3/2 Cl2 (g) Eo net = 3.93 V " More, fully, I would express the reactions with respect to HOCl powered cell as follows: H2O <--> H3O+ + OH- with the implied reactions at the anode: Al + 3OH- ⇒ Al(OH)3 + 3e- and the reaction at cathode (copper); 3 HOCl + 3 H3O+ + 3 e- --> 3/2 Cl2(g) + 3 H2O for an implied net reaction of: 3 HOCl + Al --> Al(OH)3 (s) + 3/2 Cl2 (g) where Eo net = 3.93 V A very important comment by the author above is "Even though there is relatively little HOCl in bleach", implying that the above reactions proceed even at low concentrations. --------------------------------------------------------------------------------------------- So, how does this apply potentially to water tank problems (ruptures, explosions, etc.)? Consider an Aluminum alloy tank with at least one copper pipe in contact with an ionic (mineral rich) water containing hypochlorite or HOCl (from the reaction of H2CO3 on NaOCl added during purification, for example). Bottom line, with time the Aluminum tank will corrode (forming Al(OH)3 more precisely) causing it to weaken. In addition, the protective Al2O3 will be removed allowing the side reaction: 2 Al + 3 H2O --> Al(OH)3 + 3 H2 (g) to proceed. As such, an explosive combination of H2, O2 and Cl2 may potentially form over time in the tank. The actual air concentration of Cl2, given its solubility, may vary with water temperature and the presence of dissolved CO2. Note, this explanation requires no high concentration of any compound, just a persistent use of chlorination in an Aluminum tank with continuous (or periodic with change in water levels) exposure to copper fittings. Also, in the case of a Zinc plated tank, similar chemistry follows. I have performed the above reaction as follows: Combined bleach (NaOCl) and vinegar (which contains Acetic acid HAc) in the volume ratio 1.4 parts of 5% vinegar to one part of 8.25% extra strength chlorine bleach. Then, add a piece of copper metal which will function as the cathode and an Aluminum source to act as the anode (finely cut up Al foil, for example, will do). Finally, add a touch of sea salt (better than NaCl) to act as the electrolyte to get things started. An interesting aspect of the reaction, performed in a closed vessel with shaking, is how long its takes to completely dissolve the Aluminum (a couple of hours) and a warming of the reaction mixture. No heating is required. As expected, using an excess of Aluminum slowly over a course of days, does form hydrogen gas. I tested it by exposing it to a flame and did observe a loud retort (caution: it is know that H2 and Cl2, and/or O2, can produce a kinetically powerful explosion).

John: Thanks for raising questions. I think we can forget the point on whether any more gas (meaning O2) is generated via iron bacteria as I do not feel it alone is the answer, but would explain, on testing the gas mixture, any increased in the amount of oxygen. The important point is that Fe(HCO3)2, all by itself, on standing with exposure to air breaks down releasing CO2 and deposits Fe2O3.xH2O and consumes O2. Absence oxygen, Iron bacteria could attack ferrous salts (including Fe(HCO3)2) depositing more iron oxide and,in the case of this bicarbonate, releasing CO2, Now, the hypothesis is that depending on several factors (like temperature, pH, surface area, water air flux, dissolved gases, carbonates,..) this may lead/contribute to a pressure event. Similarly, there are denitrifying bacteria that will breakdown nitrates in O2 deficient water releasing N2. Now, for those who still believe that nitrates in drinking water is not a growing problem exacerbated by droughts (global warming a possible contributing factor) with associated change in water tables, non-optimal farming practices and septic tank leakage,... please explain why typing the following into google "reduce nitrates in drinking water" returns so many US states issuing nitrate related public health material. Strange for a non-existing problem. Absent bacteria, I have presented some chemistry on how certain heavy metals (like Fe coated with Zn or Sn) can act to break down ammonia, nitrites and nitrates ultimately to N2. One does not have to accept the formation of either HNO2 or NH4NO2, just the final product Nitrogen gas which is infrequently released in laboratory situations in an explosive manner. So called laboratory preparation of Nitrogen gas employing NH4NO2 from NaNO2 + NH4Cl, at the proper pH and concentration level, requires 'careful' boiling or 'gentle heating' (for example, see page 2 at http://www.edudigm.in/downloads/Grp%2015.pdf ) as perhaps a rapid heating can precipitated an unwanted explosion. Interestingly, the dry mixture of NaNO2 + NH4Cl + a Stabilizer is employed commercially for 'safe' blasting, to quote one source (see page 240 at https://www.google.com/url?sa=t&rct=j&q=&esrc=s&source=web&cd=29&ved=0CHAQFjAIOBQ&url=http%3A%2F%2Fciteseerx.ist.psu.edu%2Fviewdoc%2Fdownload%3Fdoi%3D10.1.1.137.1104%26rep%3Drep1%26type%3Dpdf&ei=i_jKUc-kA-7e4APUyIGoCQ&usg=AFQjCNH_R1tjaiZ34fYeyHWF9gX-BbQqBw&sig2=Jy1020D3J7a1_xsgIeZn6A ): "The dry reaction [ NaNO2 + NH4Cl] proceeds briskly with 0.42 kcal heat developed per gram of mixture and has been used with a stabilizing additive in a mixture called "Hydrox" for safe blasting in coal mines, the explosive effect being caused only by pressure from expansion of the gases.[6] Such nitrogen is, however, not pure and contains some ammonia and oxides of nitrogen." You mention of Copper is interesting as, per the chemistry I have documented, the latest research suggests it is actively engaged in the reaction mechanism by forming a cuprous and then a cupric salt which, in the presence of Copper, reforms the cuprous salt. The Zinc (and Tin) metal may behave in a similar fashion.

Yes, I agree if the water is pure it shouldn't make gas. Or, one could state, if it makes gas, the water is impure. So, are there rare instances in varying locales where the water could be impure? I claim the reported incidents indicate a possible yes as well as potentially newly implemented solutions in some countries (like adding CuSO4 to the water supply). Now, on the solubility of the CO2 in the context of the specifics of the water tank, there are several parameters to examine. First, solubility is a function of temperature. Also the presence of other dissolved gases (like O2), and I would suspect, in the case for the formation of Carbonic acid, the pH of the water is a factor (see http://ion.chem.usu.edu/~sbialkow/Classes/3650/CO2%20Solubility/DissolvedCO2.html and more advanced discussion at http://www.pwtag.org/researchdocs/Used%20Ref%20docs/52%20Carbondioxide%20in%20water%20equilibrium.pdf ). Next, the available surface area at the top of the tank (to permit a rapid re-absorption if needed to avoid a pressure situation) and the flux between CO2 and the water (see http://en.wikipedia.org/wiki/Solubility_pump ). On the biochemistry, my reading is that the reaction is normally reversible in nature as decaying vegetation consumes oxygen. In the current context, with O2 accumulating at the top of the tank, no vegetation, limited air/surface contact, and the Fe2O3 at the bottom of a large water tank, there may be the possibility of different dynamics. Finally, I do not claim rust as structural failure mechanism to explain a pressure explosion. I do agree that in the presence of certain metals (Zn, Cu,..) a galvanic corrosion reaction is possibility as this leads to both corrosion and, more importantly, in my opinion, the possibility of large gas generation.

I like the concept of why it was felt necessary to add a cupric salt to the water storage tank to kill (Copper is highly toxic to lower organisms) micro-organisms. If they are present, some micro-organisms are reputedly capable of breaking the otherwise stable nitrates into nitrites (see, for example, page 527 at http://books.google.com/books?id=8aw4ZWLABQkC&pg=PA527&lpg=PA527&dq=bacteria+convert+gaseous+nitrogen+into+ammonia+nitrates+and+nitrites&source=bl&ots=SQWZ3JcRrc&sig=Of3QJShC3vIQrrFR_1Ah4-URVSY&hl=en&sa=X&ei=T-PIUYf-IeTx0wHvxIHYAQ&ved=0CDQQ6AEwAjgK#v=onepage&q=bacteria%20convert%20gaseous%20nitrogen%20into%20ammonia%20nitrates%20and%20nitrites&f=false and the role of denitrifying bacteria), and the associated problems caused therein that I had speculated on (namely, possible massive rapid gas evolution events). I can undertstand the logic, albeit drastic, in my opinion. I would list this as another possible route as it avoids the necessity of the presence of a significant amount of a heavy metal (like Zn, Cu, Pb,...), even if acting in the role of a catalysis. However, it still requires one to accept that nitrate levels may be rising and impacting water quality. Another path, for those firmly believing in the ability and willingness of goverments to keep water standards high, is to focus on Iron consuming bacteria. Apparently (see http://en.wikipedia.org/wiki/Iron_bacteria ), it is possible for certain bacteria to attack ferrous Iron salts (like Iron bicarbonate) and gradually (or rapidly?) release massives amount of CO2. Another related path, Chlorine treated Iron rich water could form FeCl2. On standing in a water tank, upon warming and in the presence of O2: 4 FeCl2 + O2 + H2O ---> 2 Fe2O3 (s) + 8 HCl Or, Iron bicarbonate on standing with exposure to oxygen or Iron bacteria: 4 Fe(HCO3)2 + O2 + 2 H2O --> 4 Fe(OH)3↓ + 8 CO2↑ so that Iron oxide accumulates at the bottom of the tank and in the presence of the right microbes, could further generate gas. Per Wikipedia (http://en.wikipedia.org/wiki/Iron_bacteria ) a reaction in low oxygen conditions, to quote: " H2O + Fe2O3 → 2 Fe(OH)2 + O2 " which is, unfortunately, not balanced with respect to Hydrogen, but does gives an idea of what is occurring. A parallel argument could go for water orginally rich in H2S and treated with Chlorine: H2S + Cl2 --> 2 HCl + S (s) where the free Sulfur could also accumulate in the water tanks. In the presence of the right bacteria, metals, pH and water temperature, gas generation is also possible. For example, Sulfur can be converted into sulfites and sulfates by either bacteria or available hypochlorous acid: Cl2 + H2O <--> HOCl + HCl 2 HOCl + S --> 2 HCl + SO2 SO2 + H2O <--> H2SO3 H2SO3 + HOCl --> H2SO4 + HCl and to quote Wikipedia source above: "Corrosion of of pipes is another source of soluble iron for the first reaction above and the sulfurous smell of rot or decay results from enzymatic conversion of soil sulfates to volatile hydrogen sulfide as an alternative source of oxygen in anaerobic environments.[5]" So the bottom line in this microbe assisted scenario, is that bacteria may serve a catalytic role in a changing oxygen content environment in the water tank, along with the presence of any nitrates, Iron or Sulfur compounds, to form problematic gaseous accumulations.

Now, to quote Wikipedia again: "In anything other than very dilute, cold solutions, nitrous acid rapidly decomposes into nitrogen dioxide, nitric oxide, and water: 2 HNO2 → NO2 + NO + H2O " So, assuming some very dilute aqueous HNO2 (or NH4NO2) is formed, it is stable, until the water temperature rises, or the solution becomes acidic. So, correct me, but dilution itself, is not a negative as it contributes to Nitrous acid's stability, and more interestingly, per this source, HNO2 has limited solubility (see http://pubs.acs.org/doi/abs/10.1021/j100333a025 ). The severity of pressure eruption is most likely a function of the quantity of nitrite present, and the change in conditions that trigger its decomposition (temperature, concentration and pH).

John: One of the pictures I posted clearly shows that the top of the tank has been blow off as if by a pressure eruption. This observation is not consisent of a failure by corrosion (there is no solution contact on the top of the tank). Also, the corrosion argument can be used to support a pressure reaction before a mechanical failure. Galvanic corrosion, like for example, with the Zn-Cu couple, does consume Zn forming some Zn(OH)2 which can eventually lead to a failure. But the other product is a whole lot of H2 gas: Zn + 2 H2O ---Cu & Heat--> Zn(OH)2 + H2 in fact, for each mole of Zn(OH)2 formed by evenly dissolving of the lining forms 22.4 liters of Hydroen gas forms! ----------------------------------------- Now, with respect to Nitric oxide formation per the decomposition of HNO2, assuming it hasn't formed the very problematic NH4NO2, it presence would be insignificant in comparison to amount of possible N2 formation from the metal induced decomposition of nitrates and nitrities in significant concentrations. Read my quote above, this isn't 'obscure chemistry', very common in testing for the presence of bacteria for nitrate determination (many google hits), but the field is perhaps more biochemistry. Also, my dated reference is perhaps a 100 years old that notes the breakdown of nitrates by various metals in neutral and other conditons forming nitrogen as one of the gaseous products. The fact that it is new to many (including chemists) may be related to the fact that only in the last 50 years has the underlying chemistry been examined and explained. In essence, bad water with high nitrates and nitrities could behave (meaning explode, erupt, burns,..) more and more like chemical plants waste water for which such reactions are common place. In my opinion, it may all be a question of water quality. Mechanical failure (like of the aquarium) by miscalculating the required thickness of the glass (no corrosion there) for the expected weight of water, fish and the like just does not happen as much in the computer age. In fact, the likelihood of someone trying to break the glass may have been anticipiated by requiring the glass to be even thicker than needed.

If one accepts the presence of nitrates in the drinking water, then in the presence of the following metals Fe and Pb in neutral solutions, and Zn, Cd, Cu, Mg and Al also appear active (see "American journal of science", Volume 112, page 188 at http://books.google.com/books?id=MvcQAAAAIAAJ&pg=PA188&lpg=PA188&dq=zinc+reduces+nitrates+to+nitrites&source=bl&ots=Dq9GZDHiTU&sig=sVyXwq5mDYxfUJzj51enKIVdpjg&hl=en&sa=X&ei=K1K4UcXCGOji4AOJ0oDYBA&ved=0CFIQ6AEwCDge#v=onepage&q=zinc%20reduces%20nitrates%20to%20nitrites&f=false in reducing nitrates to nitrites. In fact, Zinc powder is widely used in test for nitrate reducing bacteria (see, for example, http://www.mesacc.edu/~johnson/labtools/Dbiochem/nit.html via a sensitive nitrite based test. Bottomline, if there is a water quality issue associated with nitrates and nitrites in the presence of select metals, reduction reaction can occur forming N2 gas. In fact, this is pretty much common knowledge as to quote from the last source: "However, it is possible that the nitrate was reduced to nitrite but has been further reduced to ammonia or nitrogen gas."

John: I actually came up with something that supports your zinc-copper couple hydrogen based model (opened minded in spite of my critics). The argument goes, assume there is only a small amount of copper salt around (this has been one of main points of concern), but a double replacement reaction with Zn will deposit the copper on the tank's lining over time. So, with sufficient time and some copper presence, it could accumulate. Then, upon sufficient warming of the water, the reaction: Zn + 2 H2O ---Copper & Heat--> Zn(OH)2 + H2 (g) and more rapid generation of hydrogen could occur. I am still not confident, however, why pressure release valves could not address (relieve) this pressure buildup over time. ---------------------------------------------------------------------------------------------- On the nitrite argument, I was reading that ammonia contamination of the water table is a problem is increasing common in agriculture settings from ammonia based fertilizers and a movement in the water table attributed to a drought. Also, septic tanks can be a contributing factor. Here is one source (http://www.google.com/url?url=http://scholar.google.com/scholar_url%3Fhl%3Den%26q%3Dhttps://info.ngwa.org/GWOL/pdf/750600769.PDF%26sa%3DX%26scisig%3DAAGBfm3aEthCEvtudr9v10x6q9dduq6EuA%26oi%3Dscholarr&rct=j&sa=X&ei=1pW3UcHCCZXA4APa7oDYDQ&ved=0CDYQgAMoAjAA&q=ammonia+contamination+in+groundwater&usg=AFQjCNHAxw6vBvSYKGG2jZTNycURei1eUA ) where the nitrate presence based on 230 water samples was 250 mg per liter of water. One word, incredible, but even more shocking is the range from less than 1 mg/Liter to 3,100 mg/Liter. Not surprising, some cattle actually were documented as having died of anoxia (excess nitrate poisoning). Perhaps my nitrite hypothesis isn't a mystery after all, it is just an unpleasant truth reflective of a much larger problem that governments have been unwilling or unable to address, as I have been hearing more talk (case in point, Bloomberg News June 10, 2013 ) of a worldwide water problem. The first symptoms, of course, would be water quality, and an exploding tank may simply be an obvious indication of that reality. Someone tell me I am wrong, I will sleep better.

To quote a source (see http://cgmp.blauplanet.com/adv/nomol.html ): "At room temperature and at atmospheric pressure Nitric oxide is a colorless gas with low solubility in water". See also Wikipedia solubility table results at (http://en.wikipedia.org/wiki/Solubility_table#N ), low solubility indeed, .0056 g/100g water at 20 C. Now as NO molar mass is 30 g which would occupy 22.4 liters, the dissolved amount of gas of .0056 g equates to 4.2 ml. Now, add O2 to the water and things change as, to quote Wikipedia (http://en.wikipedia.org/wiki/Nitric_oxide ): "In water, NO reacts with oxygen and water to form HNO2 or nitrous acid. The reaction is thought to proceed via the following stoichiometry: 4 NO + O2 + 2 H2O → 4 HNO2 So, in oxygen poor water, Nitric oxide can accumulate per low solubility, but undergoes a chemical reaction in fresh oxygen rich water. An important point is that one accept the formation (and decomposition) of HNO2 and NH4NO2. The gaseous decomposition/explosive nature of NH4NO2 alone could account for the tank ruptures.

In my opinion, Points 3 and 4 are did not necessarily contradict. Timing is important, as Chlorine, per a source cited above, to quote: " Chloramine does not dissipate easily compared to chlorine." My take, although Cl2 can kill nearly all kinds of bugs, it dissipiates and the water can be subsequently contaminated with organic matter and bacteria (forming NH3). Hence, the argument for Chloramine, although a much weaker disinfectant, having greater longevity. You may be aware of some arguments about chlorine cleansers, that they can create super bugs in your home, well, at least, that is the assertion. No one has mentioned drinking water, but perhaps a biochemist will confirm with some words of comfort. On point 7, I would add NO pending exposure to sufficient oxygen (form newly added fresh water) at which point the highly soluble NO2 is formed and HNO2 transported from an earlier tank composition. While temporary, this may actually play a role at times (my speculation) in a nitrite concentration mechanism as an answer to valid dilution arguments. Also, NH3 and HNO2 can co-exist, at least until they react, as my reference provided previously cites the creation of ammonium nitrite by the action of ammonia on Nitrous acid: HNO2 + NH3•H2O --> NH4NO2 + H2O The other comment is more directed toward active followers of that other chemical forum.

John: Lets list what you do accept (please correct me) on these large water storage tanks. Lets restrict the water to either drinking water or existing in a somewhat vented acquarium setting. 1. Explosion of a volatile gas from sparks, or open flame, or possibly on heating in a closed vessel from the sun. 2. Gas pressure eruptions are apparently occurring. 3. The water at some prior point has been treated (Cl2, NH2Cl or aeration). 4. The water has varying mineral content, Sulfur, H2S, CH4, CO2 level, bacteria levels, free NH2Cl or Chlorine levels, organic matter (per filtering levels) and associated ammonia from decay and the like. 5. The level of gases, organic material, ... could vary based on temperature (time of year). 6. The structural composition of the tank and its age may be significant. For example, zinc plated iron or non-metallic. 7. The only applicable chemistry relates to the one of very dilute solutions and limited change in temperature, or do you accept that chemically reactive gases can accumulate producing local concentrations varying from time to time. Now the last point is important as to expanding the horizon as to the chemistry that may be at play here. By the way, there is no 'my chemistry' here, as I have given, for the most part, peer reviewed published work. Of course, the applicability is an open question as the actual conditions are not precisely known and apparently vary with the seasons and local water minerals, water treatment processes, type of storage vessel access to venting, etc., and more work on showing how the reactions come about probably has to be done. Further complicating the topic is the, at times, advanced nature of some of the peered reviewed work that may be applicable, the borderline quackery of the topic (close to exploding water), and my prior history of maintaining the purity of the science (over making friends, brown nosing, etc) and worst of all, being right on occasion. As far as my errors at SC, what was rejected previously was my assertions on the role of metals, or more precisely Cu, and as I proposed in this thread, its applicability also to Zinc. Believe me, I miss the fact that a prior investigation into an incident actually does cited Zinc lining in a refurbished water tank as a factor, else I would have made the connection originally with the cited research on Copper oxidation of ammonia. So, Elementcollector, why don't you go back to Sciencesquirrel and tell him the truth, there a 1938 documented investigation producing corroborating evidence to part of my theory, involving the ongoing loss of human lives and loss of property? Better idea yet, you re-introduce it? I need no recognition, I already have 2 advanced degrees (MS, to be precise with enough credits for a pH in one), or someone telling me I was right. There is no apparent answer out there, or some of the media would be citing something, and not the caption 'mystery'. I would gladly combine, citing source, any suggested paths for investigation in a report to the OSHA. By the way, just writing a report to the OSHA is probably not enough to get any actual funding for a study to effect preventive measures. The media may well be needed to shine the spotlight on rare event incidents.

"Ammonia won't be released unless the conditions are alkaline. HNO2 won't be produced unless the conditions are acid" In alkaline to neutral, free ammonia exists and upon warming, is expelled, there is no problem with my assertions here. HNO2 exists in highly dilute solutions, again near neutral pH. I recall someone stating on how dilute these solutions are. Remember, this is a closed vessel, both reactions did not have to occur simultaneously, but within a range of the neutral conditions. I agree this is a difficult nut to crack for several reasons. Sciencemadness (in spite of its name) refuses to acknowledge that there is a rare event, in spite of reported pressure eruptions in the media. So far, I have appreciated the discussion has it has generated some new ideas and questions to proposed paths as any theory presented must survive arguments as to how and why not. If anyone has new theories, please present them, but claiming that all the incidents and investigations thereof, are all nonsense is really closed minded, One investigation in 1938 link the mechanism to the presence of a zinc lining, which I have embraced and discussed possible associated chemistry and one other has used it to present a whole new ammonia free pathway. I am attempting to adress a serious issue involving significant property damage and loss of life. Personally attacking me, based on how I have, on many occasions embarassed a super moderator at SC (latest example, apparently not having seen/read a preparation for NCl3, he claims its sinks because of its density, but it reportedly floats in an NH4Cl solution, an important point I noted earlier) who thinks, because he has a longer career as a practicing organic (not inorganic) chemist, he is a de facto authority. I wish him the best nevertheless.

Yes, I agree not a path in for nitrite accumulation in an aquarium setting. ----------------------------------------------------------- I also just envisioned another manner employing HNO2 that is more direct, may facilitate accumulation and more frightening. As before, on warming HNO2 decomposes: 2 HNO2 --> NO + NO2 + H2O Simultaneously, NH3 gas would be released also as a result of a rise in temperature. The frightening reaction is that it is known that ammonia fumes react with NO and NO2 gas mixture forming a white smoke of ammonium nitrite, NH4NO2: 2 NH3 + NO + NO2 + H2O --> 2 NH4NO2 Here is a quote from Wikipedia (http://en.wikipedia.org/wiki/NH4NO2 ): "Ammonium nitrite forms naturally in the air and can be prepared by the absorption of equal parts nitrogen dioxide and nitric oxide upon aqueous ammonia.[2]" This powder could accumulate on the sides of the water tank. Being an unstable and explosive compound, this may present an issue. Fortunately, Ammonium nitrite slowly decomposes at room temperature liberating nitrogen gas with the dry salt being reputedly more stable. ----------------------------------------------------------------------------- "BTW, do you realise that the reactions you have cited 2 NO + O2 --> 2 NO2 2 NO2 + H2O --> HNO3 + HNO2 give rise to a reduction in gas volume and pressure?" The reaction chain of interest is first, Nitric oxide gas is accumulated via various ammonia oxidation paths. Second, NO is exposed to new O2 and per the reaction: 4 NO + O2 + 2 H2O --> 4 HNO2 in sufficient quantity to reform aqueous HNO2, but the solution concentration (or, at least, the local concentration at the top of the tank) may have increased. Third, it is decomposed again (heat, pH,..) to NO gas and the chain is repeated. Well, at least, until (if ever) there is a sufficient high HNO2 or NH4NO2 concentration to form a massive gas decomposition event.

As I cited in the reported investigation of the pressure eruption above, the water tank is equipped with a pressure release valve. So, a slow release of H2 from a trace amount of Copper catalyst, does seem to me to be the most plausible explanation. I do admit, your concentration argument may be problematic, even if a large amount of solution is present. However, changing focus from NH4NO2 to HNO2 and NO may be appropriate. Source, per Wikipedia (http://en.wikipedia.org/wiki/HNO2 ) to quote: "In anything other than very dilute, cold solutions, nitrous acid rapidly decomposes into nitrogen dioxide, nitric oxide and water: 2 HNO2 --> NO + NO2 + H2O So, upon change in temperature, or pH, the dilute Nitrous acid could decompose releasing Nitric oxide gas (NO). This gas could accumulate, cause pressure issues or act as a path to the sudden formation of nitrites upon future O2 exposure. For example: 2 NO + O2 --> 2 NO2 2 NO2 + H2O --> HNO3 + HNO2 or, with excess NO relative to oxygen: 4 NO + O2 + 2 H2O --> 4 HNO2 (see Wkipedia http://en.wikipedia.org/wiki/Nitric_oxide ) So NO gas may act as a nitrite accumulation/storage vehicle, as well as a potential cause of gas pressure issues.

John: I will consider adding Hydrogen to the pressure based tank ruptures after some experimenting. I have witness the nitrite based nitrogen gas formation reaction. Using accelerators like H2O2, overnight a pressure detonation (use a plastic container for safety) is easily obtainable. If the H2 from warming water (to accelerator) with a Zn-Cu couple can produce a gas pressure explosion, it is a good candidate, in my opinion, requiring only warm water, zinc and a small amount of a soluble copper salt. The argument for adding ammonia to the input list appears reasonable as in high NH3/Urea environment (an aquarium, for example) there is, at least, one unexplained gas pressure rupture. The rarity of the events also suggest the need for an unusual concentration level of a common impurity. The chief factor favoring the nitrite theory is the known violent decomposition property of the reputedly formed NH4NO2. An interesting aspect of the nitrite decomposition model is that per this reference (http://www.google.com/url?sa=t&rct=j&q=rapid%20decomposition%20of%20ammonium%20nitrite&source=web&cd=9&ved=0CFoQFjAI&url=https%3A%2F%2Fsrac.tamu.edu%2Findex.cfm%2Fevent%2FgetFactSheet%2Fwhichfactsheet%2F169%2F&ei=fz6zUandA-nT0wGI54CoDA&usg=AFQjCNFRsBbL2istZBuWQrhrHLAM-wK7MA ) to quote: "Bacteria oxidize ammonia in a two-step process, first to nitrite (NO2-) and then to nitrate (NO3-). The main factors that affect nitrification rate are ammonia concentration, temperature and dissolved oxygen concentration. During summer, ammonia concentration is very low and so nitrification rates are also very low. During winter, low temperature suppresses microbial activity. During spring and fall, ammonia concentration and temperature are intermediate, conditions that favor maximum nitrification rates. Spring and fall peaks of nitrite concentration are commonly seen in fish ponds." So an interesting statistical correlation, but by no means definitive, based on seasonality and a larger sample, may be observed for municipal water tanks with respect to nitrite formation and reported incidents (interestingly, 2 instances of all the events I listed relating to pressure ruptures occurred in the month of April). I would also expect for flammable/explosion events, that summer months would be slightly more prevalent.

Since we are talking about the existence of ammonia in drinking water, it is perhaps appropriate to cite some sources (see http://www.atsdr.cdc.gov/phs/phs.asp?id=9&tid=2 ) to quote: "It [ammonia] is found in water, soil, and air, and is a source of much needed nitrogen for plants and animals. Most of the ammonia in the environment comes from the natural breakdown of manure and dead plants and animals." Also, per this reference (http://www.hc-sc.gc.ca/ewh-semt/consult/_2012/ammonia-ammoniac/draft-ebauche-eng.php ):: "Levels of ammonia, either naturally present in the source water or added as part of a disinfection strategy, can affect water quality in the distribution system (e.g., nitrification) and should be monitored." and: "The concentration of free ammonia entering the distribution system can lead to nitrification and the potential increase of nitrate and nitrite in drinking water." So the bottom line is a possible cause for municipal water tank pressure related explosions could be an unusual increase in nitrite concentration, especially from the oxidation of ammonia in the presence of certain metals precipitating a dangerous and rapid nitrogen gas evolution depending on conditions (including pH, concentration and water temperature). -------------------------------------------------------------------------------------------- " have also on several occasions had first hand observations of light galv. steel electrical raceways with a painted surface sustain an electrolysis processes that reduced it to crumbs in less than a year." Galvanic corrosion is certainly real. Here is a comment from a Wikipedia article on this subject (http://en.wikipedia.org/wiki/Galvanic_corrosion ) to quote: "A common example of galvanic corrosion is the rusting of corrugated iron sheet, which becomes widespread when the protective zinc coating is broken and the underlying steel is attacked. The zinc is attacked preferentially because it is less noble, but once it has been consumed, rusting of the base metal can occur in earnest." Also, some important factors listed include "the conductivity of the water within the system" and its pH. In the case of the sign, exposure to acidic exhaust could lower pH. Dust (containing minerals) and moisture form the electrolyte. "Could electrolysis provide the means for structural failure and/or means to create hydrogen gas within these tanks?" Yes, corrosion could increase pitting and possible structural issues. In my opinion, hydrogen gas generation should not be such an issue (assuming it is slow), assuming the presence of functioning pressure release valves, as to precipiate a gas pressure tank eruption, but I could be wrong.

OK John, a zinc-copper couple will form Hydrogen with water, but only if hot water (see http://www.ucc.ie/academic/chem/dolchem/html/elem001.html ). Not clear if the larger water tanks can be heated to measurably form Hydrogen. But I admit, H2S and H2 are not safe gases around sparks and could be responsible for some of the maintenance related explosions. "Are you aware that they chlorinate water supplies and that destroys ammonia quite well?" Yes, but the job isn't done until you drink the water. The problem is keeping the bacteria and decaying matter from contaminating the once clean chlorinated water. One of the argument for the use of Chloramine is its increased stability, so its is around to disinfectant after all the Cl2 is spent. Chlorine, forming HOCl, is much more effective, especially in killing viruses, if still present in the water. Per the summary source I provided the link for above, some relevant points to quote: "• Chloramine is a less effective disinfectant than chlorine. The World Health Organization (WHO, PDF 950 KB) says that "monochloramine is about 2,000 and 100,000 times less effective than free chlorine for the inactivation of E. Coli and rotaviruses, respectively." • Chloramine does not dissipate easily compared to chlorine. • Chloramine stays in the water distribution system longer than chlorine." One of the reason I introduce the aquarium example is because it is an extreme case of where the NH3 (and/or Urea) content can get high. In other words, normally the ammonia content is low, but, on rare instances, it may spike as it occurs in fish tanks. Note, aeration helps to remove NH3 and limits bacteria growth. But, for whatever reason, a temporary loss in aeration can raise ammonia levels. Also, yes, H2O2 is not used commercially to purify water. The citation of H2O2 is per the chemistry supplied in reference paper noted above. The reaction using even dilute Hydrogen peroxide is many times faster than waiting for atmospheric oxygen to act on ammonia, but otherwise not that different.

JohnCuthber: I will not try to defend the NCl3 path as I have already stated it is a weak hypothesis among better alternatives. I also will further characterize it as very unlikely. I also accept and will note the essential dilution associated with NH2Cl presence that you stated. ----------------------------------------------------------------------------------------------------------------------------------------------------- My claim that Zinc is a potential issue as well as copper has mixed, but insightful, reviews in the literature. For example, this old reference from the Proceedings of the Chemical Society (Great Britain), Vol 22-24, pages 39-40, at http://books.google.com/books?id=LNEfAQAAMAAJ&pg=PA40&lpg=PA40&dq=oxidation+of+ammonia+by+H2O2+in+the+presence+of+Zinc&source=bl&ots=t8qGMmn74s&sig=EjFv4ljd7Y2S2xU-Q_ONK2K3too&hl=en&sa=X&ei=ITGxUe6uMZCi4APbtIGgDA&ved=0CCkQ6AEwADge#v=onepage&q=oxidation%20of%20ammonia%20by%20H2O2%20in%20the%20presence%20of%20Zinc&f=false characterizes the oxidation of ammonia in the presence of air and zinc as sometimes leading to positive results and sometimes negative. This suggested to me a missing piece, which occurred to me is a trace amount of Copper or soluble Copper salt. In essence, a so called copper-zinc couple. The chemistry, for example, in the presence of NH3 and O2 (previously given) for Copper: 2 Cu + 4 NH3 + 1/2 O2 + H2O --> 2 [Cu(NH3)2]OH ...[Catalyst Role]..Cu(NH3)2]OH......................... NH3 + 3 H2O2 -----------------> HNO2 + 4 H2O HNO2 + NH3•H2O ---------------> NH4NO2 + H2O 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 But now, in the presence of a Zinc plate, a significant reaction could be: Zn + [Cu(NH3)4](OH)2 <---> Cu + [Zn(NH3)4](OH)2 so the Copper would be regenerated at the expense of Zinc, to form potentially more nitrite. Even more profound is that this reaction is not limited to Zinc, but potentially could occur in the presence of any metal (or alloy thereof) that is chemically more reactive than Copper that forms a complex with ammonia. So, some metal container could be at risk, for example, an iron container coated in Zn, Sn, Cr,...

John Cuthber: Thanks for your comment. My first comment is that it is generally believed that Zn does not react with water (see for example, http://www.webelements.com/zinc/chemistry.html ). However, in the presence of a salt solution, oxygen and NH3, I can accept an electrochemical oxidation reaction forming a Zinc salt. In acidic water, yes, some Hydrogen. However, my main focus is on the unexplained pressure buildup, most likely sudden, that is rupturing these tanks. I did state I am skeptical on the NCl3 hypothesis, but there are known explosions that accidentally occur during water chlorination attributed to NCl3. Maintaining the proper water turbulence is usually employed to reduce the possible creation of NCl3 by promoting the hydrolysis reaction which normally occurs in hot water (see Wikipedia at http://en.wikipedia.org/wiki/Nitrogen_trichloride ) : NCl3 + 3 H2O → NH3 + 3 HOCl See also this summary source (http://www.chloramine.org/literature_pdf/chloramine_facts_060911.pdf ), to quote: "The three species of chloramine constantly and rapidly shift from one form to another. The species that predominates is dependent on pH, temperature, turbulence, and the chlorine to ammonia ratio." Also, here is a small part of a preparation for NCl3 that is relevant, to quote: "The chlorine is absorbed and oily drops of the trichloride float on the surface of the solution" where, I would speculate, NCl3 droplets could coalesce absence any turbulence. Now, NCl3 by itself is not a candidate, in my opinion, on any of pressure based explosions, which you did not comment on. Now, this link (http://www.jstor.org/discover/10.2307/41232342?uid=3739808&uid=2129&uid=2&uid=70&uid=4&uid=3739256&sid=21102372208537 ) supplied previously per a 1938 investigation, in the opening paragraph of the investigation, does point to the Zinc lining as a contributing factor to the water pressure tank explosion. Some of the chemistry I have presented, based on more recent chemistry (for example, the Cu catalysted oxidation of ammonia in air is 1962 and the second author 2011) indicate the formation of nitrites noted for their notorious N2 decomposition properties (namely rapid & violent, pH and concentration sensitive). I do think the term mystery is appropriate as per this article " Tank Explosion Still Mystifies Investigators?". To quote: "Officials Are Hunting For The Cause Of A Water Tank Accident That Killed A Minneola Worker. One Said There Were No Safety Violations. March 15, 1995, by Terri Coole Sentinel Correspondent MINNEOLA — Officials are still baffled about what could have caused a water tank explosion two weeks ago that killed a city worker who was crushed against a chain-link fence by a 3,000-gallon blast of water." The picture below is, I believe, an example of a pressure based rupture on a water tank.

CaptainPanic: Thanks, my transfer from a word document appears to have broken some of the links. I will repost the whole document below and, I may be able to insert (with edits) an abstract to improve the flow. I sincerely believe this thread serves a public need and thank you again for your input. ---------------------------------------------------------- REPOST (for the purpose of a report to OSHA with the help of the members of Science Forum and its administrators): ABSTRACT While the rupturing and/or explosion of massive sized water tanks occurs rarely, it is oftened accompanied by some loss of life and large property loss. With the intent of forwarding a report to the OSHA, this resport assembles some a chemical based theories differentiating by class for explosive vapors versus clearly pressure eruptions. A brief history of water tank incidents along with an interesting aquarium incident are documented. Hypotheses are delineated for each class of type of explosion. Chemistry for pressure explosions, the chief interest for this work, suggest auto-decomposition of Chloramine (NH2Cl) forming N2 gas, together with possible rapid decomposition of nitrites in the presence of select metals to N2 as a potential cause, as well as even the possible liberation of CO2 from Iron rich water. HISTORY 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-water-tank-explosion/ ). 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-explosion-127869308.html ) 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/article-Water-tank-explodes-in-Galax , and also in Chester, New York (see http://chroniclenewspaper.com/apps/pbcs.dll/article?AID=/20121003/NEWS01/121009993/Water-tank-explodes-in-Chester- ) where to quote: "Internal pressure blew the end of the tank off and through the attached treatment building, completely demolishing the building,” police said. There is also a report at http://www.jstor.org/discover/10.2307/41232342?uid=3739808&uid=2&uid=4&uid=3739256&sid=21102371396667 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_explodes_all_over ). 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 Normally, more acidic conditions (low pH) promote the formation of Nitrogen trichloride. However, where the water has high H2S content, to de-odorize the water, more heavy chlorination can be used. The reaction are: Cl2 + H2O <--> HOCl + HCl H2S + HOCl --> H2O + S (s) + HCl S + 2 HOCl + H2O --> H2SO3 + 2 HCl H2SO3 + HOCl --> H2SO4 + HCl all of which consumes Chlorine (actually HOCl from Chlorine hydrolysis) requiring more Cl2 to address the H2S problem. This makes the water more acidic with the creation of HCl, H2SO3 and H2SO4 contributing to the problematic pH senstive nitrite decomposition reaction discussed below. --------------------------------------------------------------------------------- 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: To quote: "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_Monochloramine_decay_in_model_and_distribution_system_waters ------------------------------------------------ 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_Copper_Dissolution_In_Aqueous_Ammonia 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=copper-mediated%20non-enzymatic%20formation%20of%20nitrite&source=web&cd=1&sqi=2&ved=0CCoQFjAA&url=http%3A%2F%2Fsphinxsai.com%2Fvol3.no2%2Fchem%2Fchempdf%2FCT%3D23(646-656)AJ11.pdf&ei=NLWwUdOMJJWu4APZy4DwAQ&usg=AFQjCNFaObAi5_3NNOdt8e1DiRoiHzg9bg&bvm=bv.47534661,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 notorious for significant and sudden Nitrogen gas evolution. This reaction could occur slowly over time, as a function of the rapidity of water turnover in storage tank. However, 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), the water turnover could decline. 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_Monochloramine_decay_in_model_and_distribution_system_waters ) 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.jsp 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 problematic 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 suggested paths for future investigation include explosive vapors attributed to H2S, NH2Cl vapors 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).