Yes, as I will explore the chemistry, but apparently your mother's common household salt should be banned along with your Dad's seltzer water (he can
keep the booze for now), that window/floor cleaner with that dangerous ammonia (just let the dirt alone) and worst of all, that reputedly safe
disinfectant for all those minor cuts and scrapes (just let it get infected). There actually may be some merit, after all, on those warning labels
about mixing household cleaners. I do all of this, of course, in the light of knowledge to keep all of us safe.
The unspeakable alternative, of course, is to stop banning all chemicals period as a fruitless venture, given that apparently even the most common and
innocent chemicals among us can lead to such dangers.
Now, onto the chemistry for the disbelievers, or those who still want to hold onto their salt:
First, the Solvay process to turn that dangerous salt (NaCl is know to increase blood pressure, which can lead to a stroke, a common killer among
adults) along with that evil Selzter water (H2CO3, employed in alcoholic and soda beverages no less) along with that powerful and dangerous ammonia
cleaner (NH3):
NaCl + H2CO3 + NH3 --> NaHCO3 + NH4Cl
Now, comes the final step, add H2O2 and place into a closed container, wait and watch out. The lesser known reaction, I claim with time (a day or
more) is that the ammonium ion is oxidized to nitrite by the H2O2, or more rapidly, upon mild heating (to 60 C at most), where the basic Sodium
bicarbonate serves to stabilize the creation of the dangerous Ammonium nitrite, a high explosive and very toxic compound (if ingested). This statement
is also true even in aqueous solutions if it is concentrated, and/or has less than a neutral pH.
So the basic baking soda (NaHCO3) fosters the creation of the Ammonium nitrite, and the increasing acidity (from the Hydrochloric acid, HCl) makes the
solution a lot less stable (safe).
The good news is that with luck, the NH4NO2 may simply and safely decompose forming Nitrogen (N2) gas, but in a sealed container, there still could be
a gas explosion.
So, hold onto your salt while you can.
(This really is potentially a very dangerous synthesis, and I do not recommend that anyone performs it at home, especially around children)
[Edited on 5-4-2013 by AJKOER]AJKOER - 8-5-2013 at 16:20
I performed a variation on this experiment.
First, I formed (NH4)2SO4 by combining aqueous ammonia and a solution of MgSO4 (Epsom salt).
MgSO4 + 2 NH3 + 2 H2O --> Mg(OH)2 (s) + (NH4)2SO4
Filtered out the white Mg(OH)2 suspension, which I am keeping to make MgO (which is reported to be able to decompose alcohols).
Next, added CaCl2:
(NH4)2SO4 + CaCl2 <---> 2 NH4Cl + CaSO4
that is, only a very dilute milk colored solution is formed as apparently CaSO4 is soluble in NH4Cl (and, upon research, other ammonium salts as well,
and especially soluble in Ammonium acetate, citrate,..).
Next added, Na2CO3.H2O2, and witnesses a a very thick and bright white salt (CaCO3 or CaSO4):
2 NH4Cl + CaSO4 + Na2CO3.H2O2 --> (NH4)2CO3 + CaCO3 (s) + H2O2 --> 2 NH3 + H2O + CO2 + CaCO3 (s) + H2O2
or:
2 NH4Cl + Na2CO3.H2O2 + CaSO4 (aq) --> (NH4)2CO3 + 2 NaCl + CaSO4 (s) + H2O2 --> 2 NH3 + H2O + CO2 + CaSO4 (s) + H2O2
The reaction occurred in a sealed and partially compressed plastic container. I proceeded to add more aqueous ammonia, warmed the solution in a hot
water bath and let sit. After a day, a significant gas expansion occurred with the probable formation and decompositon of NH4NO2. I tossed the vessel
and on landing it bursted as a result of the gas pressure.
Some related research to quote:
Per Wikipedia on NH4NO2:
"Ammonium nitrite solution is stable at higher pH and lower temperature. If there is any decrease in pH lower than 7.0, It may lead to explosion. It
is desirable to maintain pH by adding ammonia solution. The mole ratio of Ammonium Nitrite to Ammonia must be above 10% mole ratio.
NH4NO2 → N2 + 2 H2O "
also:
"Ammonium nitrite may explode at a temperature of 60–70 °C.[2] It decomposes more quickly when a concentrated solution than when it is a dry
crystal."
An interesting reaction from the literature of Marcellin Berthelot, dry ammonia gas reacts with the nitrogen dioxide and nitric oxide, at room
temperature,
2 NO2 + 2 NO + 4 NH3 --> 2 NH4NO2 + 2 H2O + 2 N2
where it is reported that solid ammonium nitrite inside a tube explodes if heated on a water bath to between 60-70 C. Also, the substance gradually
decomposes at room temperature, slower if cold, or faster in aqueous solutions, forming nitrogen gas.
Also:
"Strong solutions of H2O2 with a few drops of NH4OH or solutions of ammonium carbonate (with or without NaOH or Na2CO3) can be let to stand 24 hours
without any nitrite formation occurring. But upon longer standing, even with a small amount of hydroxide then nitrite forms. Nitrite also forms when a
dilute solution of H2O2 is mixed with NH4OH and a little Na2CO3 and is evaporated over pure conc. H2SO4 with a bell jar. H2O2 forms (even in very
dilute solutions) nitrite very rapidly, if the H2O2 solution is mixed with a few drops of NH4OH and a little NaOH or Na2CO3, and this then boiled in a
retort to a very small volume. They suggest this nitrite formation as a demonstration experiment because it is very quick to do, and then after
acidification of the colorless liquid with H2SO4, the HNO2 can be nicely be proven to be present." Hoppe-Seyler woelen - 9-5-2013 at 13:05
I do not think that this is really interesing. You just have an explosion due to pressure buildup from formed gases. H2O2 always leads to formation of
oxygen. Combinations of nitrite and ammonia/ammonium lead to formation of N2. So, when put in a sealed bottle, it is just a matter of waiting to see
it go BOOM.kristofvagyok - 9-5-2013 at 15:58
Ammonium nitrite is not just decomposing forming nitrogen and water... I have made an oxidation on a large scale where the oxidant was ammonium
nitrate and one product was ammonium nitrite. It mainly decomposed to nitrogen and probably water, but a not nitrogen monoxide also evolved. It's
usually not described, but it always present, especially if a larger scale is made.
And it was on a fairly acidic pH.... I'm still alive. AJKOER - 12-5-2013 at 18:07
OK, for those who believe that only a relatively safe pressure reaction is the main outcome here always, please note the following comments in this
old (1905) report (see page 242 at http://books.google.com/books?pg=PA242&lpg=PA242&dq=... from Journal Chemical Society, London, Volume 88, Part 2), to quote:
"Electrolytic Oxidation of Ammonia to Nitrites. Erich Muller and Fritz Spitzer (Ber., 1905, 38, 778—782. Compare Traube and Biltz, Abstr., 1904, ii,
727).—In the presence of a small amount of sodium hydroxide, ammonia may be oxidised electrolytically to nitrite even in the absence of copper
compounds.
In the presence of copper hydroxide and sufficient alkali, the oxidation of ammonia to nitrite does not cease suddenly when the nitrite concentration
has reached a certain value, but appears to proceed quite independently of the nitrite concentration. In these experiments, the oxidation was allowed
to proceed for a comparatively short time only, so that the amount of alkali present was not greatly reduced. The formation of nitrite is intimately
connected with the amount of alkali present, and when no sodium 'hydroxide is present, but only ammonia, nitrite, and copper hydroxide, it is found
that the nitrite is transformed into nitrate more rapidly than the ammonia into nitrite, and thus the concentration of the nitrite tends to decrease.
Nitrogen is also formed during the oxidation. J. J. S."
Note the interesting statement "In these experiments, the oxidation was allowed to proceed for a comparatively short time only, so that the amount of
alkali present was not greatly reduced." So the author performing scientific research in an aqueous environment, felt it was advisable to cut the
research short. Clearly, this was with respect for (or fear of) the associated product and its properties, this being in line with my Wikipedia
comments above, to quote again:
"Ammonium nitrite solution is stable at higher pH and lower temperature. If there is any decrease in pH lower than 7.0, It may lead to explosion. It
is desirable to maintain pH by adding ammonia solution."
[Edited on 13-5-2013 by AJKOER]ScienceSquirrel - 14-5-2013 at 04:55
You are not going to get an explosion for the simple reason that even if the chemistry does work, and I think that it is a little dubious, you are
unlikely to have a high enough concentration of reagents.
I think it is time for AJKOER to put down his pipe, climb out of his armchair and go down the mall to buy a bottle of soda, some hydrogen peroxide and
some ammonia solution.
Open the soda, tip out some of it, add the rest of the reagents and screw the lid on tightly.
If you use a plastic soda bottle, it will not shatter but split safely down the side.
Even glass beer bottles, and I have plenty of experience of those, do not explode when filled with beer and over pressurised due to secondary
fermentation because liquids hold very little energy.
Connecting a bottle to a gas cylinder and pressurising it is highly dangerous however as gases do hold a lot of energy and the bottle will burst
sending fragments everywhere.
This is why steam boilers are tested with water before being used for live steam.AJKOER - 14-5-2013 at 05:46
You are not going to get an explosion for the simple reason that even if the chemistry does work, and I think that it is a little dubious, you are
unlikely to have a high enough concentration of reagents......
ScienceSquirrel:
Thanks for reading and responding. With respect to your concentration comment I quote from above:
"In the presence of copper hydroxide and sufficient alkali, the oxidation of ammonia to nitrite does not cease suddenly when the nitrite concentration
has reached a certain value, but appears to proceed quite independently of the nitrite concentration."
But, of course, how are my starting concentrations? Per my revised preparation (yes, there could be concentration issues with the 1st suggested
route), we have (NH)4SO4 (plus, I would recommend, aqueous ammonia) as the more than ample ammonia source. For H2O2 and alkaline we have 2Na2CO3.3H2O2
(Sodium percarbonate) which, in commercial bleaching products, also have added Na2CO3. These reactants should all be sufficiently strong for the
formation of NH4NO2.
I also assure you that my last modified experiment run is real and became a bit scary for me given the starting quantities involved. I am sorry (or
perhaps not) that I added extra ammonia which should have reduced gaseous decompostion and promote aqueous NH4NO2 formation. Eventually, I decided to
abandon ship around 30 hours into the run. Perhaps some would say I came to my senses upon witnessing the gaseous expansion, and converting it into a
relatively safe pressure event, a course I would recommend for all replicating this experiment although even handling the solution at his stage could,
nevertheless, trigger an event. The reason most likely being the associated increase in concentration of NH4NO2 and lowering of the pH (promoting
instability). I would personally describe the whole process (perhaps a bit exaggerated) as constantly checking dynamite, and then it begins to sweat
(gaseous decomposition in this case).
I am willing to take the word of an authority when they call NH4NO2 a high explosive. I have presented above a report displaying the respect (fear?)
granted to aqueous solutions of Ammonium nitrite by researchers.
But if you think my focus on NH4NO2 is misplaced, dare I address the ease of accidental formation and properties of NCl3 from household chemicals?
Perhaps it is better to talk about NH4NO2 afterall, and hope that I have, in fact, exaggerated without ever having moved from the safety of my
armchair.
------------------------------------------
Some extracts from "Bretherick’s Handbook of Reactive Chemical Hazards", Sixth Edition, Volume 1, to quote:
"Lawrence, G. M. , Plant/Oper. Progr., 1989, 8(1), 33
An explosion in the vent of an ammonia combustion plant was attributed to deposition
of ammonium nitrite/nitrate crystals. It is considered that the very unstable
nitrite acts as a sensitiser to the nitrate, and that explosion is triggered by contact
with acid."
"Heating a mixture of an ammonium salt with a nitrite salt causes a violent explosion
on melting [1], owing to formation and decomposition of ammonium nitrite. Salts
of other nitrogenous bases behave similarly. Mixtures of ammonium chloride and
sodium nitrite are used as commercial explosives [2]. Accidental contact of traces
of ammonium nitrate with sodium nitrite residues caused wooden decking on a
truck to ignite [3]."
Interestingly, these reports apply to the reputedly more stable dry NH4NO2. However, as this reference addresses largely reported commercial
incidents, I understand the omission of smaller laboratory aqueous accidents.
[Edited on 14-5-2013 by AJKOER]ScienceSquirrel - 14-5-2013 at 09:35
Ammonium nitrite in dilute aqueous solution is never going to explode, just like dilute aqueous solutions of picric acid will not explode.
9% hydrogen peroxide is a pussy cat, it becomes slightly warm on adding a catalyst that causes it to break down, 35% peroxide gets seriously hot but
you need high test peroxide of about 70 - 80 % + before it turns to steam and oxygen.
Water has a high specific heat and it is a very effective buffer.AJKOER - 14-5-2013 at 10:15
Ammonium nitrite in dilute aqueous solution is never going to explode, just like dilute aqueous solutions of picric acid will not explode.....
ScienceSquirrel:
Thanks for bringing up this point, which apparent has a somewhat unique, and unexpected property, for those with knowledge of energetic materials.
Here is a reference: "The Effect of Chromium (VI) Compounds on the Stability of Concentrated Aqueous Solution of Ammonium Nitrite", November 1973, by
ARMY FOREIGN SCIENCE AND TECHNOLOGY CENTER CHARLOTTESVILLE VA, link: http://oai.dtic.mil/oai/oai?verb=getRecord&metadataPrefi... ), other than Wikipedia, which I gave previously. To quote:
"Abstract: Ammonium nitrite is highly unstable, both in a dry state and in concentrated aqueous solutions. When such solutions are used industrially
they are normally stored at a pH of 8 - 9 and a temperature of 5C. or less. The effect of additions of chromium compounds on the decomposition of
ammonium nitrite solutions was studied, under isothermic and non-isothermic conditions. It was found that such compounds catalyze the decomposition of
ammonium nitrite more intensely than HCl. This led to the conclusion that contamination by even small amounts of such compounds should be carefully
avoided in industrial work with ammonium nitrite solutions in order to minimize the danger of explosive decomposition."
"Ammonium nitrite is a potentially explosive solid, and so the aqueous ammonium nitrite solution is made by adding ammonium chloride and sodium
nitrite, both stable compounds, to water. Even so, the solution must be heated carefully."
The aqueous stability issue is general known among those working with the compound and is generally described as, for example: (see http://www.google.com/url?sa=t&rct=j&q=ammonium%20ni... ) "Since ammonium nitrite is know to decompose, spontaneously, into nitrogen and
water".
[Edited on 14-5-2013 by AJKOER]kristofvagyok - 14-5-2013 at 10:35
In this flask, there was a total 1mol of ammonium nitrate dissolved in 150cm3 water with 1% of copper-acetate, used for an oxidation.
The ammonium nitrate was the oxidant in the system and ammonium-nitrite was generated. And it was heated on a water bath to 50 Celsius.
A lot nitrogen evolved, with some nitrogen-monoxide and be surprised: I'm still alive. AJKOER - 18-5-2013 at 08:57
Kristofvagyok:
I did notice that your target temperature was 50 Celsius, which is a safe temperature for aqueous NH4NO2 (below 60 C). pH is also a major concern and
your use of aqueous NH4NO3 with a 1% copper-acetate, if used soon after mixing, is most likely OK as well.
Interestingly, HNO2 is reported to be sensitive to agitation as well as concentration and temperature. Hence, you may see terms like 'carefully
heating' (and not vigorous boiling) in various preparation involving nitrites to produce nitrogen. Also, lesser known is that HNO2 decomposition is
sensitive to surface area. In my reaction above, I created a large surface with the formation of a very fine calcium salt, which is probably not on
the recommended to do list.
[Edited on 18-5-2013 by AJKOER]APO - 23-5-2013 at 02:49
Kristofvagyok, what camera did you take that picture with? That is some sexy depth of field.kristofvagyok - 23-5-2013 at 13:56
Kristofvagyok, what camera did you take that picture with? That is some sexy depth of field.
It is a Nikon D5000. The dept of field is not a camera selective thing, it comes from the lens what was used. In this case I used a Nikkor 50mm f1.4
AI lens at full aperture.
