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AJKOER
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A SAFE TEST RUN ON THE CHLORINATION OF (NH4)2SO4 TO MAKE DILUTE H2SO4 (WITH Na2SO4 IN THIS TEST)
I thought of a somewhat safe procedure for a small test run on the chlorination of Ammonium sulfate to make H2SO4. This procedure is only intended to
gain insights (or suggest improvements) as to the original cited synthesis. Note, this particular preparation is intended to quickly and somewhat
safely create dilute H2SO4 containing Na2SO4 in this instance. The process was:
1. Prepare (NH4)2SO4 per the reaction of NH4OH on a slight excess of Epom Salt (MgSO4.7H2O). Filter out the Mg(OH)2 to obtain a clear solution of
(NH4)2SO4.
2. Add NaHSO4 to NaClO/NaCl (6% Bleach) to create only the required amount of Chlorine. As a departure from the original synthesis, one immediately
add this solution mix to (NH4)2SO4 from Step 1, which now also includes a small amount of 3% H2O2 (note, precise quantities are listed for this test
run below) in a large (3 liter) clear plastic bottle.
Some of the reaction equations:
NaHSO4 + NaOCl --> Na2SO4 + HOCl
NaHSO4 + NaCl --> Na2SO4 + HCl
Cl2 + H2O <---> HCl + HOCl
and, per Watt's for dilute Chlorine water solutions only:
HCl + HOCl + H2O2 --> 2 HOCl + H2O
where an excess of H2O2 is to be avoided (as it would further reduced the HOCl to HCl and O2). Reference: "Watts' dictionary of chemistry", Volume 2,
page 16, link: http://books.google.com/books?id=ijnPAAAAMAAJ&q=HOCL#v=o...
The key speculated reaction in this synthesis is:
(NH4)2SO4 + 2 HOCl + 2 H2O --> 2NH2Cl + 4 H2O + H2SO4
After the reaction is completed, one can add more H2O2 to remove unwanted NH2Cl (this is a suggested improvement to the original synthesis in
converting HCl to HOCl, and also possibly decomposing Chloroamines):
2 NH2Cl + 2 H2O2 --> N2 + Cl2 + 2 H2O
Note, no heat was applied to the solution to reduce the possibilty of the NCl3 formation.
--------------------------------
More precisely, here are the quantites employed:
MgSO4 11.1 grams 6.6 ml
NaOCl/NaCl 110 grams 100 ml
NH4OH 30.7 grams 33 ml
H2O2 2.76 grams 3 ml
NaHSO4 10.85 grams 3.96 ml
93.2% NaHSO4 11.64 grams 4 ml
However, preparation of the (NH4)2SO4 via MgSO4 (which also includes the addition of 10 ml H2O to MgSO4 and an about 3 ml of H2O2) upon filtering
suffered a loss in the amount of Ammonium sulfate available (the solution went from to a total of 46 ml to 35 ml after filtering). I proceeded anyway
with the synthesis. The reaction upon mixing the (NH4)2SO4 and Chlorine generating solution at first displayed white smoke upon pouring into the 3
liter vessel. Letting noxious gases escape for about a minute (do outdoors, these were Chlorine and I would guess Chloroamines and perhaps Nitrogen),
I compressed the plastic container (to test for continuing gas expansion) and sealed it. A cloud was formed that dissipated after several minutes of
shaking, but a greenish tinct was observed. After an hour, a more intense noticeable greenish color was observed, which even with shaking, remained.
I decided at this point to prepare and add more (NH4)2SO4 to make up for the original shortfall. Immediately, upon adding 11 ml of the new
(NH4)2SO4/H2O2 mix, all visual evidence of Cl2 disappeared, however I did not witness a significant gas generation as previously. The whole reaction I
found surprisingly less violent than I anticipated (but, for a more concentrated preparation, this may not be the case). The final product was a clear
solution with a strong Chlorine smell with seemingly greater viscosity (most likely testing acidic as well given the chlorination). I plan on testing
it further starting by freezing the solution.
[Edited on 21-2-2012 by AJKOER]
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Kola
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I don't know much about lead Chamber process but I know H2SO3 can be safely prepared in a laboratory by dissolving SO2 in water. SO2 can be easily
prepared by heating sulphur &Oxygen in the presence of a catalyst eg Pt. The H2SO3 can be oxidized with H2O2 to yield H2SO4...any problems with
this Scheme?
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White Yeti
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Quote: Originally posted by Kola  | I don't know much about lead Chamber process but I know H2SO3 can be safely prepared in a laboratory by dissolving SO2 in water. SO2 can be easily
prepared by heating sulphur &Oxygen in the presence of a catalyst eg Pt. The H2SO3 can be oxidized with H2O2 to yield H2SO4...any problems with
this Scheme?
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Expensive, the sulphur is not too expensive, but the hydrogen peroxide is definitely expensive when working on small scales.
"Ja, Kalzium, das ist alles!" -Otto Loewi
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Kola
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| Quote: Originally posted by White Yeti |
Expensive, the sulphur is not too expensive, but the hydrogen peroxide is definitely expensive when working on small scales. |
expensive but less risky, besides I figure O2 can do the oxidation but the resulting solution will be dilute cos of SO3 evolution..and perhaps the
reaction wont go to completion
Nothing gets higher than where God dwells. If there is an upper limit to the universe, call it the abode of God
Chemistry is the central science!
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White Yeti
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Hydrogen peroxide is dangerous mind you. The 3% is safe, but once you go over 30% (the lowest concentration for this to be useful), things get
dangerous.
Of course, 18M sulphuric acid is not by any stretch of the imagination a compound to be handled without care.
Also, sulphur dioxide is dangerous. Do not treat it as an ordinary gas, it will choke you if you do not employ proper ventilation. If you have asthma,
this gas will not have mercy, so be careful.
"Ja, Kalzium, das ist alles!" -Otto Loewi
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AJKOER
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UPDATE ON THE TEST RUN (CHLORINATION OF AMMONIUM SULFATE)
First, I previously stated that the key speculated reaction in this synthesis of dilute H2SO4 was:
(NH4)2SO4 + 2 HOCl + 2 H2O --> 2NH2Cl + 4 H2O + H2SO4 [1]
I have since found a reference stating that stock solutions for the purpose of testing Chloramine levels are prepared by adding NaOCl (Bleach) to
(NH4)2SO4. This supports the speculated key reaction if one replaces Na with H (but not as to speed as NaOCl is clearly more ionic than HOCl). See
page 52, top left, at:
http://www.fwrj.com/techarticles/0608%20tech%205.pdf
However, I also found an interesting equation that suggests a portion of the H2SO4 could be consumed by liberated NH3:
"The overall chloramine decomposition reaction, based on acid-catalyzed disproportionation reaction of NH2Cl (Eq 2.43):
3 NH2Cl + H+ --> N2 + NH4+ + 3 H(+) + 3 Cl(-)" [2]
Source: "White's Handbook of Chlorination and Alternative Disinfectants" by Black & Veatch Corporation, page 150 (middle of page):
Link:
http://books.google.com/books?id=mGVbIoW2lNAC&pg=PA150&a...
If we rescale Equation [2] to correspond to Equation [1] and add H2SO4:
2 NH2Cl + 2/6 H2SO4 + 4/6 H2SO4 --> 2/3 N2 + 2/6 (NH4)2SO4 + 2 HCl + 4/6 H2SO4
which suggests that 1/3 of the expected H2SO4 yield is reduced by the presence of Monochloramine. However, the author also cites the following
reaction (also on page 150):
"NHCl2 + NH2Cl --> N2 + 3 H(+) + 3 Cl(-)"
So, if we added more HOCl, then per the equation:
HOCl + NH2Cl ===> NHCl2 + H2O
(as a reference of the above and many other associated reactions, see: http://www.h2o4u.org/chloramination/chemistry.shtml )
then, some NHCl2 could be created, which correspondingly reduces the NH2Cl. This would increase the H2SO4 yield as upon adding more HOCl, we have
NH2Cl + (HOCl + NH2Cl) + 4 H2O + H2SO4
= NH2Cl + (NHCl2 + H2O) + 4 H2O + H2SO4
Or:
(NH2Cl + NHCl2) + 5 H2O + H2SO4 = N2 + 3 HCl + 5 H2O + H2SO4
So the new target reaction is:
(NH4)2SO4 + 3 HOCl + 2 H2O --> N2 + 3 HCl + 5 H2O + H2SO4
and in the direct application of HOCl synthesis (as opposed to chlorination) the best one could hope for is apparently dilute H2SO4 in HCl.
Also, the original targeting of the production of NH2Cl is not advised as the Monochloramine is more stable than NHCl2 which, in contrast, is much
more volatile and easier removed. In fact, prolonged boiling is required for NH2Cl's decomposition combined with aeration only slowly decomposing it,
and only very advanced/costly filtering being able to remove it from solution.
In genral, one should also be aware that the chemistry of Chloramine is complex with the species of chloramine rapidly and constantly shifting as a
function of temperature, pH, turbulence, and Cl2/NH3 ratio. See: http://www.chloramine.org/literature_pdf/chloramine_facts_06...
------------------------------------
Update Frozen Solution
An interesting salt has separated (tiny long thin crystals which in solution shimmer and resembles cotton fibers). Some bubbling also in a perfectly
clear solution with a strong chlorine-like smell.
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AirCowPeaCock
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These guys do NOT like Chloramines(;
BOLD
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AJKOER
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Actually, I do not like Chloramines either. If in California or (or Massachusetts) and you see a green color in the local swimming pool, avoid getting
the water (actually Chloramine treated) in your eyes, and certainly don't let a pet drink it (toxic to dogs and fish). Apparently, Campden pills which
releases SO2 in water are able to remove Chloramines in about a minute.
------------------------------
I thought it was interesting to note if we take the new target equation scaled by two and net the H2O:
2 (NH4)2SO4 + 6 HOCl --> 2 N2 + 6 HCl + 6 H2O + 2 H2SO4
and noting that:
6 HOCl <--> 3 Cl2O + 3 H2O
we have:
2 (NH4)2SO4 + 3 Cl2O --> 2 N2 + 6 HCl + 3 H2O + 2 H2SO4
Note, the treatment of dry Ammonium Sulfate with any strong oxidizer (as DiChlorine Mono-oxide) is strongly not advised (at worst, explosive), and it
is intended that a concentrated aqueous (NH4)2SO4 solution be treated.
Also, upon adding more Cl2O:
2 (NH4)2SO4 + 6 Cl2O --> 2 N2 + 6 HCl + 6 HOCl + 2 H2SO4
but:
6 HCl + 6 HOCl <==> 3 Cl2 + 3 H2O
So:
2 (NH4)2SO4 + 6 Cl2O --> 2 N2 + 3 Cl2 + 3 H2O + 2 H2SO4
So assuming this very vigorous reaction can even be performed  , a massive
treatment of a concentrated (NH4)2SO4 solution with Cl2O gets one at best 40% H2SO4 (and a lot of Chlorine). Caution, DiChlorine Mono-oxide is many
times more poisonous than Cl2 and has explosive properties as well.
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Kola
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| Quote: Originally posted by AJKOER | Actually, I do not like Chloramines either. If in California or (or Massachusetts) and you see a green color in the local swimming pool, avoid getting
the water (actually Chloramine treated) in your eyes, and certainly don't let a pet drink it (toxic to dogs and fish). Apparently, Campden pills which
releases SO2 in water are able to remove Chloramines in about a minute.
------------------------------
I thought it was interesting to note if we take the new target equation scaled by two and net the H2O:
2 (NH4)2SO4 + 6 HOCl --> 2 N2 + 6 HCl + 6 H2O + 2 H2SO4
and noting that:
6 HOCl <--> 3 Cl2O + 3 H2O
we have:
2 (NH4)2SO4 + 3 Cl2O --> 2 N2 + 6 HCl + 3 H2O + 2 H2SO4
Note, the treatment of dry Ammonium Sulfate with any strong oxidizer (as DiChlorine Mono-oxide) is strongly not advised (at worst, explosive), and it
is intended that a concentrated aqueous (NH4)2SO4 solution be treated.
Also, upon adding more Cl2O:
2 (NH4)2SO4 + 6 Cl2O --> 2 N2 + 6 HCl + 6 HOCl + 2 H2SO4
but:
6 HCl + 6 HOCl <==> 3 Cl2 + 3 H2O
So:
2 (NH4)2SO4 + 6 Cl2O --> 2 N2 + 3 Cl2 + 3 H2O + 2 H2SO4
So assuming this very vigorous reaction can even be performed , a massive
treatment of a concentrated (NH4)2SO4 solution with Cl2O gets one at best 40% H2SO4 (and a lot of Chlorine). Caution, DiChlorine Mono-oxide is many
times more poisonous than Cl2 and has explosive properties as well.
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this scheme is too complex...i don't think its logical for any laboratory chemist. Maybe can there be a simpler pathway?
Axhandle's synthesis is much more realistic and it can be used on a large scale.
Have you experimented your theory? What were your results?
Nothing gets higher than where God dwells. If there is an upper limit to the universe, call it the abode of God
Chemistry is the central science!
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AJKOER
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Kola: I agree with your comment. The most practical approach is also the simplest, mentioned in the original synthesis, just over chlorinate (NH4)2SO4
forming HCl and H2SO4. To be more precise, per the equation:
2 (NH4)2SO4 + 6 HOCl --> 2 N2 + 6 HCl + 6 H2O + 2 H2SO4
upon adding 6 HCl to each side:
2 (NH4)2SO4 + 6 HOCl + 6 HCl --> 2 N2 + 12 HCl + 6 H2O + 2 H2SO4
and as:
6 Cl2 + 6 H2O <---> 6 HOCl + 6 HCl
2 (NH4)2SO4 + 6 Cl2 + 6 H2O --> 2 N2 + 12 HCl + 6 H2O + 2 H2SO4
or:
2 (NH4)2SO4 + 6 Cl2 --> 2 N2 + 12 HCl + 2 H2SO4
or, more precisely working with aqueous Ammonium sulfate:
(NH4)2SO4 + x H2O + 3 Cl2 --> N2 + 6 HCl + H2SO4 + x H2O [CAUTION: No heat to avoid the creation of an oily yellow NCl3 explosive]
at which point, removal of the HCl (without heating) via dilute H2O2, aeration,.. becomes the concern, without significantly diluting the H2SO4 or
excessive expense. Note, adding FeSO4 to bleach has been suggested previously on a Sciencemadness thread as a means, without acid, to generate Cl2.
Nevertheless, the synthesis is far from perfect with smelly toxic fumes and stills provides much more HCl (to be addressed) than H2SO4.
[Edited on 1-3-2012 by AJKOER]
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Formatik
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| Quote: Originally posted by AJKOER | (NH4)2SO4 + x H2O + 3 Cl2 --> N2 + 6 HCl + H2SO4 + x H2O [CAUTION: No heat to avoid the creation of an oily yellow NCl3 explosive]
at which point, removal of the HCl (without heating) via dilute H2O2, aeration,.. becomes the concern, without significantly diluting the H2SO4 or
excessive expense. Note, adding FeSO4 to bleach has been suggested previously on a Sciencemadness thread as a means, without acid, to generate Cl2.
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It's full well possible that the nitrogen trichloride might not be skipped as an intermediate and accumulates regardless of overchlorination, in which
case it would have to be decomposed by standing for 24 hours, maybe even less. It would make it an immense explosion hazard. Cold temperatures also
favors nitrogen trichloride formation. And then if the reaction goes as thought you may be able to isolate some sulfuric acid.
For generation chlorine, oxidizing hydrochloric acid has been the preferable chemical route.
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AJKOER
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Formatik:
In my test run, I did avoid an excess of Chlorine (via stoichiometric calculations) and without heating, observed no yellow oily liquid. Given that
the formation of the deadly NCl3 is highly endothermic, these measures I would think, are necessary first safety measures. Note, strong light may
serve as a heat substitute (as well as a trigger), and should also be avoided.
Now, NCl3 reputedly has limited solubility in water, but with continuous solution turbulence (dangerous, I would otherwise believe) and aeration
(commercially employed to remove NCl3 in so called Breakpoint Chlorination of water) may speed up the hydrolysis/removal of any Nitrogen trichloride
that has formed.
Still, I think your comment is to some extent appropriate ("immense explosion hazard"), and hence my reluctance to make especially the chlorination
path the recommended synthesis route.
As a source on dealing with unwanted NCl3 see: "White's Handbook of Chlorination and Alternative Disinfectants" by Black & Veatch Corporation,
page 23, where the authors mention breakpoint chlorination followed by aeration, and an alternate method of pre-treating the chlorine gas with UV
light in the spectrum 3600-4400 (my speculation on how this works is as NCl3 is know to explosively decompose with UV light that this treated Cl2
decomposes the NCl3 on formation). Link:
http://books.google.com/books?id=mGVbIoW2lNAC&pg=PA23&am...
See also, for example, Patent 4435291, "Breakpoint chlorination control system", which does mention the use of mechanical agitation.
http://www.patentgenius.com/patent/4435291.html
[Edited on 2-3-2012 by AJKOER]
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Formatik
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What I was thinking when saying nitrogen trichloride formation is favored in cold temperatures was in reference to hydrolysis of nitrogen trichloride
which increases with temperature, whereas when it is kept cool then nitrogen trichloride better accumulates. I had phrased it badly, cold does not
promote its formation but slows its hydrolysis. Ideal temperatures for forming nitrogen trichloride are cool to warm, but not freezing.
AJOKER, it's a good sign you observed no yellow oily liquid. However, nitrogen trichloride is an elusive compound not only because of decomposability,
but because it needs to be let sit to see it (typically after 20 minutes), because it can be in finely divided form where especially also milky
appearance of solutions makes it harder to see.
The reaction between ammonium sulfate solution and chlorine could be:
6 Cl2 + (NH4)2SO4 --> 2 NCl3 + H2SO4 + 6 HCl
Then NCl3 could possibly react with an excess ammonium sulfate per this equation:
(NH4)2SO4 + 2 NCl3 = 2 N2 + 6 HCl + H2SO4
Note: similarly that if NH4Cl is present in large excess the following reaction has been observed to occur slowly with ammonium chloride: NCl3 + NH4Cl
= N2 + 4 HCl (Bray, Dowell, J. Am. Soc. 39 [1917] 905).
Thus we could eventually have something like the:
(NH4)2SO4 + 3 Cl2 = N2 + 6 HCl + H2SO4
Dilute H2SO4 decomposes NCl3 rapidly, but dilute HCl is said not to decompose NCl3 (according to Davy, Phil. Trans. 103 [1813] 1). However both acids
when concentrated decompose the trichloride rapidly under N2 evolution according to the same reference (these might also act as trigger initiators to
the pure compound when added to it). So there is the possibility that as nitrogen trichloride is formed, can be decomposed by dilute H2SO4 before it
has the chance to accumulate (speculation on assuming it forms in an atmosphere corresponding to dilute sulfuric acid). Though Chapman and Vodden (J.
chem. Soc. 95 [1909] 138) describe NCl3 as solubilizing more rapidly in aq. HCl than in water or aq. H2SO4, because HClO which is formed in
equilibrium reaction of water hydrolysis, is destroyed by HCl but not water or sulfuric acid. The presence of both acids seems to have some
promotional decomposing effect.
Note in some points it may look conflicting but it may really be more complicated with nitrogen trichloride. I've formed it in acidic ammonium sulfate
solution and it did not decompose right away (probably too dilute of an acid concentration), in fact it could be isolated.
Addendum: we may be interested in destroying any NCl3, and there are substances which are reported to prevent the formation of NCl3 mentioned in Gilb.
Ann. 47 (1814) 58 coal powder, CO2, sulfur powder, etc. But if the nitrogen trichloride is needed for the reaction to work it could be in that sense,
counter-productive to destroy it.
[Edited on 3-3-2012 by Formatik]
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AJKOER
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Formatik:
Some interesting stuff (at least to me).
OK, I found a source with interesting comments on the decomposition of NCl3.
First, the decomposition reaction:
NCl3 + H2O → NHCl2 + HOCl
is first order in time, so your comment on letting NCl3 sit for hours is supported.
Next, the source cites work by Kumar that NCl3, in the presence of excess HOCl, will either decompose into HOCl or be subject to reduction/oxidation.
Noting that Chlorine water is:
Cl2 + H2O <---> HOCl + H(+) + Cl (-)
I would argue that upon adding a small amount of NH3 (via the hydrolysis of (NH4)2SO4), the reaction is first moved to the right forming more HOCl, as
much more rapidly:
H(+) + Cl(-) + NH3 --> NH4Cl
which would thus add to the decomposition of NCl3 per the above author's assertion on the effect of increasing HOCl (see also H2O2 comment below).
In addition, the author's assertion on the ability of HOCl to decompose NCl3 adds support to my equation, to quote:
| Quote: Originally posted by AJKOER | UPDATE ON THE TEST RUN (CHLORINATION OF AMMONIUM SULFATE)
............
So the new target reaction is:
(NH4)2SO4 + 3 HOCl + 2 H2O --> N2 + 3 HCl + 5 H2O + H2SO4
and in the direct application of HOCl synthesis (as opposed to chlorination) the best one could hope for is apparently dilute H2SO4 in HCl.
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However, on old Sciencemadness thread (http://www.sciencemadness.org/talk/viewthread.php?tid=3347 ), the implication is clearly that one can directly formulate NCl3 with a sufficient
amount of HOCl (so carefully monitor/restrict the quantity of HOCl to keep all your body parts!). To quote:
| Quote: Originally posted by kazaa81 | Hi,
in roguesci's forum I found a different synthesis of NCl3, with acetic acid, sodium hypochlorite and ammonium nitrate (probably any ammonium salt
would be right). I've reformatted this and post it now:
==
Nitrogen chloride
3 NaOCl + 2 CH3COOH + NH4NO3 -----> NCl3 + NaNO3 + 2 CH3COONa + 3 H2O
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although an excess of HOCl may disrupt the created Nitrogen trichloride. Note, Acetic acid plus NaOCl yields HOCl, so this is a reworking of the
synthesis. Also, replacing a sulfate with a nitrate in a pure HOCl reaction could produce dilute HNO3.
---------------------------
Finally, your comment on accumulation of NCl3 is in agreement with the author's statement that NCl3 can be quite stable in the presence of excess Cl2.
However, excess Cl2/H2O equals HCl/HOCl, and in the presence 3% H2O2, per my updated synthesis, this is mostly HOCl and may hinder NCl3 without even
assuming that H2O2 directly attacks the Nitrogen trichloride.
Source: Per page 14 at link:
http://scholar.lib.vt.edu/theses/available/etd-042299-143911...
[Edited on 4-3-2012 by AJKOER]
[Edited on 4-3-2012 by AJKOER]
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AJKOER
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Apparently Chloramines can be decomposed by super chlorination, potassium peroxymonosulfate (2KHSO5·KHSO4·K2SO4), and ozone. The latter two
processes support my contention that dilute H2O2 may be effective on NCl3. To quote:
"Chloramines can be removed from the water by the following three methods:
By adding a high dose of chlorine and raising the levels to 10 times the level of combined chlorines (5 to 10 ppm) for a minimum of 4 hours. This is
called super chlorination. To remove chloramines, the ratio of chlorine to chloramines must be at least 7.6 to 1. If this ratio is not obtained more
chloramines will be produced.
By adding a non-chlorine shock to the water. The most common chemical used for this is potassium peroxymonosulfate (MPS). This "shocking" requires the
addition of 1 oz. per 625 gallons of water.
By adding ozone to the water. If an ozone generator is installed and wired so that it comes on each time the pump comes on, then oxidation of the
ammonia and nitrogen compounds will take place on a continuous basis. This reduces, and can even eliminates the need for shocking. Each time ammonia
and nitrogen enter the ozonated water, they are oxidized by the ozone."
Link:
http://www.rhtubs.com/chlorine.htm
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ScienceSquirrel
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Thread Pruned
13-3-2012 at 05:25 |
ScienceSquirrel
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Thread Pruned
13-3-2012 at 05:26 |
Cemtek
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SO3 generator.
I'm in need of a small SO3 generator to produce a known quatity on SO3 (between 1ppm & 100ppm) from SO2 & H2O for testing & calibrating a
laser based SO3 analyzer.
Does any one know of a company that makes something like that?
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elementcollector1
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I bought a bottle of Rooto's H2SO4 at McLendon's yesterday, and after measuring 100mL on a scale, it weighed 168.9 g. This translates to a density of
1.69, or 78%.
However, MSDS says 93.2%.
Which one is the correct value?
(Does my scale need calibration?)
Elements Collected:52/87
Latest Acquired: Cl
Next in Line: Nd
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ScienceSquirrel
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| Quote: Originally posted by elementcollector1 | I bought a bottle of Rooto's H2SO4 at McLendon's yesterday, and after measuring 100mL on a scale, it weighed 168.9 g. This translates to a density of
1.69, or 78%.
However, MSDS says 93.2%.
Which one is the correct value?
(Does my scale need calibration?) |
I would check the scale by weighing 100g and 200g of water. I have done this with my 0.1 and 0.01g scales using a 100ml pipette and they are spot on.
If you use a measuring cylinder it will be nowhere near as accurate but it should do.
It is drain opener and i doubt most end users care if it is ca 80% or 90%, it will still open drains so I would take the MSDS with a pinch of salt!
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Poppy
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Not that this equation should be proposed at all. After all, the way up here has shown uses of bleach, so let come the exotric equations!!!
I've been interested on making tripped equations out of standard electrode solutions, such an equation should provide means of a different contact for
implanting the damn oxygen to sulfur dioxide, the problems found is, first, to surpass the thumb-rule 0,6V that should give aprecciable reaction
rates:
SO4(2-) + 4H+ + 2e- --> SO2 + 2H2O +0.17 (reversing it)
SO2 + 2H2O --> SO4(2-) + 4H+ + 2e- -0.17
2Fe3+ + 2e- --> 2Fe2+ +0.77V
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SO2 + 2H2O + 2Fe3+ --> SO4(2-) + 4H+ + 2Fe2+ +0.5V
furthermore quite acidic solutions are already needed to bear Fe3+, so the equilibrium would shift, which could be compensated with pressurized SO2
until the overvoltage is achieved.
Seriously, why is the direct reaction so slow?
K = 1.08
E° = 1,03.10^(-3)V
(From harsh calculations) I think the harsh Fe3+ enviroment would increase chances of success.
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White Yeti
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Using Fe+3 is really overkill. Why not use oxygen?
O2(g) + 2 H2O + 4 e− ----> 4 OH−(aq)
"Ja, Kalzium, das ist alles!" -Otto Loewi
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Poppy
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Fe3+ enchances surface contact. A 1 molar solution is in a steric proportion of 1 to 55 iron/ H2O
oxygen bubbles would hardly beat this
Indeed, the Fe3+ feedback is provided by bubbling oxygen, but since Fe2+ oxydises faster in the presence of oxygen than does SO2, iron as a
catalyst...
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White Yeti
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I see the reasoning, but your previous post made it seem as though you were using ironIII as the oxidizer, not as a catalyst, otherwise you would have
included oxygen in the overall reaction as well.
Try it out and see how it goes.
"Ja, Kalzium, das ist alles!" -Otto Loewi
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Poppy
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| Quote: Originally posted by White Yeti | I see the reasoning, but your previous post made it seem as though you were using ironIII as the oxidizer, not as a catalyst, otherwise you would have
included oxygen in the overall reaction as well.
Try it out and see how it goes. |
Sure, then pick in the apparatus to be given a try to this.
Direct oxygen doesn't work because it would flush, I mean, just vent the SO2 away.
It must proceed alligator style, step by step (wait, did this make any sense?)
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testimento
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What are methods of producing sulfur dioxide the most straightforward way from OTC sources? OTC means for ex. sodium sulfate, potassium sulfate, etc.
One idea was to flush the salt with very hot air(1000-1500C) to cause it to pyrolyze into oxide and sulfur dioxide. Another idea was to roast it with
carbon to form sulfide, react with acid to form H2S and burn it to form SO2.
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Jay
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The most straightforward way to get sulphuric acid is to BUY IT AS BATTERY ACID. The price is pretty good. Do you guys not have it available in your
area?
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