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.