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testimento
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At very low temps water suck ammonia like nothing. It can easily hold something like 40-50% solutions. 
SO2 needs to be well refrigerated, CaCl2+ice resembles mostly pumpable ice, but if you stuck a tank of that into a freezer with refluxer and lead the
SO2 there, it'll condense. When you get it liquefied and dry, you can store it at room temp in a steel cylinder. I prefer getting an ordinary pressure
cylinder with needle or ball valve and pour or directly condense the SO2 into there and you can store it forever.
-10C ain't enough for SO2, but you can use it for your good, because it holds the CaCl2 ice cooler a lot longer than 30C summer heats.
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AJKOER
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When cold, there could be snow outside!
You may have noticed the measurable rise in uv glare. Suggestion: Do any experiment that requires photolysis.
For example, see my thread "Better/Advanced Chlorate Formation" at http://www.sciencemadness.org/talk/viewthread.php?tid=27013 where I obtained a chlorate by way of a few days of uv exposure. Here is an extract:
Quote: Originally posted by AJKOER  | SUCCESS!
I appear to have relatively easily created KClO3 in good yield without nearly any sophisticated equipment employing some recent published work to
expedite chlorate formation. The method, however, is based on photolysis and takes days in good uv conditions (like snow, for example).
The process is based on simply acidifying a 8.25% Chlorine bleach with Acetic acid (actually used vinegar), adding NaCl and placing the solution in a
glass vase sealed with a thin layer of clear plastic foil (to allow uv rays to react with any escaping Chlorine from the solution that would otherwise
be blocked by the thick glass). The volume ratio employed for the 8.25% bleach was 2.07 parts to one part vinegar. Speculation on the reaction chain:
7 NaOCl + 7 HAc --> 7 NaAc + 7 HOCl
2 HOCl --uv--> 2 HCl + O2
HCl + HOCl = Cl2 + H2O
HCl + NaOCl --> NaCl + HOCl
Net:
8 NaOCl + 7 HAc --> 7 NaAc + NaCl + H2O + O2 + Cl2 + 5 HOCl
And, upon adding NaCl also to expedite the chlorate formation:
8 NaOCl + 7 HAc + 5 NaCl --> 7 NaAc + 6 NaCl + H2O + O2 + Cl2 + 5 HOCl
Further photolysis most likely proceeds along the following paths involving the species ·Cl, ·OH and ·ClO:
Cl2 + hv --> 2 ·Cl
2 ·Cl + 2 HOCl --> 2 ·OH + 2 Cl2 (g)
2 ·OH + 2 HOCl --> 2 H2O + 2 ·ClO (forming, at most, one Cl2O2 as there are poisoning reaction paths)
Cl2O2 + HOCl --> HClO3 + Cl2(g)
forming at most one mole of chlorate for each eight moles of NaOCl given adequate sunlight, pH, ionic strength and solution concentrations.
Source: See for example "Photolysis of free chlorine species (HOCl and OCl- ) with 254 nm ultraviolet light" page 281 attached, and also Table 2, page
797 at http://www.geosci-model-dev-discuss.net/3/769/2010/gmdd-3-76... .
----------------------------------------------------
Now, some references on recent chemistry on paths to chlorate.... |
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cyanureeves
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so can an ultraviolet light work in a closed tool shed if say we only have one day of snow?there is an ultraviolet lamp on ebay for 25 bucks and one
of your pictures show a pink ultravioletlike glow peering through the glass container.interesting!
[Edited on 2-6-2014 by cyanureeves]
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AJKOER
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Cyanureeves:
Depending on what you are trying to accomplish, the photolysis can actually require many days, But remember, patience is a virtue (well, at least, so
I have been told).
The advantage of using sunlight, other than not impacting your utility bill (possibly significant if scaling up) or requiring an investment in new
equipment, is that sunlight affords a wide spectrum band exposure. On the other hand, a low pressure Hg lamp usually has a narrower band, which may or
may not work for your particular application. Safety issues (like eye damage for onlookers) another factor along with the fact that the bulb does burn
out and even breaks. On the positive side, one could always invest in a solar panel to offset the energy load and make money to pay for any medical
bills.
A special advantage in the case of KClO3 formation is that in the presence of KCl when the sunsets in the winter, it can get cold enough so that the
chlorate precipitates out (so no need for that power hungry freezer either, just sell it on ebay).
[Edited on 6-2-2014 by AJKOER]
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Zyklon-A
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@AJKOER, is your method really cheaper than electrolysis of KCl?
| Quote: |
forming at most one mole of chlorate for each eight moles of NaOCl |
That seems to be rather waistful, considering that KCl is cheaper than NaOCl, and you get I mole of KCLO3 for every mole of KCl.
Ok, I just did the math...
Your process:
1 gal NaOCl 6%(aq)= $2.00, 6%=0.87 mol/L, x3.78L=3.2886 mol/$2.00.
3.2886 mole ÷ 8(estimated yield/mole)=0.411075 mole KClO3.
0.411075 mole KClO3=50.15115 grams.
Electrolysis process:
KCl=$2.00/lbs, 1lbs=6.0805 mole, -->6.0805 mole KClO3.
6.0805 mole KClO3 =741.871 grams KClO3.
Your method: $1.00-->25.075575 grams.
Electrolysis method: $1.00--> 370.9355 grams.
Of course with both methods you will inevitably lose some due to solubility issues and what-not.
third edit: Wow! I'm really surprised at how cheap making KClO3 really is. I buy mine at Skylighter for $11.94 per lb.
But with $1.00 I can make 370 grams, that's crazy.
[Edited on 6-2-2014 by Zyklonb]
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Metacelsus
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AJKOER: If you are going to start from NaClO, you should disproportionate it with heat, then precipitate potassium chlorate with potassium chloride.
Per Wikipedia (https://en.wikipedia.org/wiki/Sodium_hypochlorite#Reactions):
3 NaClO → NaClO3 + 2 NaCl
This method is well-known among amateurs, does not require any other reagents, and is more efficient than your proposed pathway. Your "speculation" on
the reaction chain is just that: I find a photolytic mechanism highly implausible. The reference you cite is irrelevant: it is studying reactions of
chlorine species at low concentrations in the gas phase, and nowhere mentions chlorate.
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Zyklon-A
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Yes I've done that and yields were pretty good, but clearly electrolysis of KCl(aq) is the cheapest for large scale production. (The anode can be
quite expensive).
Disproportion of 1 gal 6% NaOCl(aq)-->133.7 grams KClO3-Which isn't too bad, 66.85 grams/$1.00
[Edited on 6-2-2014 by Zyklonb]
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Metacelsus
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You are right; electrolysis is the way to go for any large-scale preparation.
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AJKOER
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In my opinion, there appears to be some misunderstanding on the actual theory underlying the chlorate formation per electrolysis. Here is one source
(link: http://utahpyro.org/compositions/PreparingChlorates.pdf) to quote, in part:
"2.1 theory
Mechanism of chlorate formation
The reactions taking place in chlorate cells are not fully understood even today. A
summarized description of the process will be given here, and the interested reader is
referred to the literature listed below for a more extensive description.
The theory of Foerster and Mueller regarding the reactions in chlorate cells, developed
about 80 years ago, is the most accepted. The following reactions are said to take place at
the electrodes:
At the anode:
2Cl- --> Cl2(aq) + 2 electrons
At the cathode:
2H2O + 2 electrons --> H2 + 2OH-
The dissolved chlorine gas can then react with water to give hypochlorous acid(HClO):
Cl2(aq) + H2O = HClO + H+ + Cl-
From this reaction it can be seen that if the chlorine does not dissolve but escapes to the
atmosphere, no H+ will be generated to neutralize the OH- formed at the cathode and the
pH of the electrolyte will increase.
The hypochlorous acid thus formed will react in acid-base equilibrium reactions with
water to give hypochlorite ions and chlorine gas (dissolved). The exact concentrations of
dissolved Cl2, Hypochlorite (ClO- )and HClO depend on the pH, temperature and
pressure among other things. In the solution, chlorate will be formed (mainly) by the
following reactions:
2HClO + ClO- --> ClO3- + H+ + 2Cl-
and
2HClO + ClO- +2OH- ---> ClO3- + 2Cl- + H2O
These reactions take place at a rather slow rate. Since this reaction pathway is the most
efficient one as we will shortly come to see, the conditions in the cell are usually
optimized to increase their reaction rate. The pH is kept within a range where HClO and
ClO- are simultaneously at their maximum concentration (which is at around pH=6). The
temperature is kept between 60 and 80 degrees centigrade, which is a good compromise
between the temperatures required for a high reaction rate, low anode and cell body
corrosion and high chlorine solubility (remember the chlorine evolved at the anode has to
dissolve in the solution to start with). Many cells also have a large storage tank for
electrolyte in which the electrolyte is kept for a while to give these reactions some time to
take place.
Alternatively, chlorate may also be formed by oxidation of hypochlorite at the anode as
follows:
6HClO + 3 H2O --> 3/2 O2 + 4Cl- + 2ClO3- + 12H+ + 6 electrons
Oxygen is evolved in this reaction, which means a loss of current efficiency (the energy
used for oxidizing the oxygen in water to the free element is lost when the oxygen
escapes to the atmosphere). When the reaction routes are worked out, it turns out that
following this path 9 faradays of charge are required to produce 1 mole of chlorate,
whereas only 6 faradays are required to do that following the route mentioned earlier.
Therefore, optimizing the conditions for that route improves current efficiency.
To prevent the products from being reduced at the cathode again, a membrane around the
cathode was employed in the past. Today, small amounts of chromate’s or dichromate’s
are added. A layer of hydrated oxides of chromium will then form around the cathode
effectively preventing hypochlorite and chlorate ions from reaching the cathode surface.
Finally, it should be mentioned that the reactions forming perchlorates do not take place
until the chloride concentration has dropped to below about 10%. Therefore, cells can be
constructed and operated in such a way that chlorate is produced almost exclusively. The
chlorate can then be purified and fed into a perchlorate cell. Depending on the type of
anodes used in the chlorate cell, the purification step may also be skipped and the
electrolysis continued until all chloride has been converted into perchlorate. Although
slightly less efficient (and therefore not used a lot in industrial setups), this is much less
laborious and therefore probably the preferred method for home setups."
-------------------------------
So the chemistry is pretty much the same, except that I am starting with hypochlorite, and the electrolysis starts at an earlier stage with chlorine
formation.
Now, my positive comments with respect to electrolysis is that it is, by nature, a hot, concentrated and highly ionic solution with the benefits of
significant chloride presence. The latter is important as per a 2000 paper, for example, see http://pubs.acs.org/doi/abs/10.1021/ic991486r , and also per the source, "Effect of Chloride Ion on the Kinetics and Mechanism of the Reaction
between Chlorite Ion and Hypochlorous Acid", at http://www.researchgate.net/publication/23141635_Effect_of_c... .
A less favorably argument for electrolysis, in the same source, to quote:
"Moreover, they found that acetate ion accelerates the formation of · ClO2 enormously. It was interpreted by a steady-state formation and further
reactions of acetyl hypochlorite. "
In other words, the presence of organics (like Acetic acid and Sodium acetate) can apparently accelerate the creation of chlorate in a
non-electrolysis setting.
In addition, per the quoted source above, maintaining the proper pH is necessary for efficiency. The allowed pH range (3.5 to 6.5 for HOCl, or 5.5 to
6.5 for highest joint HOCl and ClO- concentrations) is significant as Chlorate formation is favored in near neutral to mildly acidic conditions as
this allows free chlorine present in the solution to exist primarily as HOCl (fostering chlorate formation) as opposed to Cl2 gas (in lower pH
ranges). High pH can also form chlorate, albeit, not as efficiently. Source: see the chart on 'Available Chlorine Present as HOCl' versus pH on the
last page of this link: http://aquaox.wordpress.com/category/hypochlorous-acid-sodiu... .
Temperature is also a concern, as noted above, to promote low anode corrosion and high chlorine solubility, as otherwise chlorate yield could suffer.
In addition, the author notes the issue of any escaping Chlorine gas, which is needed to promote the reaction chain.
Now, as to whether these concerns are easily managed in a home chlorate cell situation is, in my opinion, debatable. Note, my point here relates not
so much as to quantity that can be conveniently produced, but as to the efficiency of the underlying mechanism.
[EDIT] Now, depending on ones location, KCl electrolysis is still most likely a convenient path to chlorate, but I am now seeming
accessibility/visibility issues with Potassium chloride based water softeners themselves. For example, my local Walmart no longer stocks it, but I can
order it on a credit card for shipment to the local store (with no added charged for placing the buyer, in areas without hard water issues, on a watch
list, I suspect).
[Edited on 13-2-2014 by AJKOER]
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PeeWee2000
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Hey everyone I just had a (maybe not so) great idea as to a reaction well suited to cold weather and as AJKOER pointed out the increased UV index.
Phosgene! It can be made by the reaction of Chloroform and UV light and liquifies aroud 8°C so it would take advantage of both the UV light and
excess of snow!
“Everything is relative in this world, where change alone endures.”
― Leon Trotsky
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testimento
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If you want to do UV experiments, find some shops that sell stuff for home reptiles and snakes and shit and find those UV lamps. If you know what
you're doing(know to use graded UV filter goggles), you will want to find water sterilizing lamps which will give a lot of UVC radiation, though
you'll need quartz glass, because borosilicate and soda glass will mostly filter UVB and UVC off. These are orders of magnitude more efficient for
chlorine photocatalysis.
Chlorine dissolves about 9 grams per liter of 5C water. If you'd install UVC lamp in the middle of the vessel and bubble chlorine into water, you
should be able to make decent concentration of HCl.
If you want decent amount of chlorates fast, get microwave transformer, some copper wire(as thick as you can fit into the tran for 6-7 coils),
powerful enough rectifiers and some lead or graphite and iron slabs. Connect the iron slabs directly to the trans wiring. Install the rectifiers on
the other wire as a positive side, and this will be your anode, and in this you'll plug your graphite or lead slabs. Be sure that you will have
powerful enough rectifiers and large enough connections: for every mm2 of copper you will conduct about 6 amps and for carbon, iron and lead you'll
push about 500mA per square mm. Stainless steels rank much lower, around 50-100mA/mm2. Between transformer and the mains you will get dimmer
controller for few bucks on ebay so you can control the power. With 1kw trannie you can easily push up to 100-150A into the cell. I have tested
welding transformer which vaporized one 500A rectifier within a minute at 6V current and worked well with series of 1000A rectifiers and this all cost
about 200 bucks so it's your best bet. Oh, and you will want to add powerful fan near the transformer, dimmer and rectifiers, they generate a lot of
heat. No computer fans, but just some powerful blower.
Chlorates will need about 3-6 volts of DC and perchlorates about 5-8 volts DC. Only lead(PbO2 will form on it) and platinum group metals and some
exotics will make perchlorates, but if you're making chlorates, your best bet is graphite. Saturate a water solution with KCl and adjust the pH around
6, temperature between 60-80 and the volume of the cell to about 2A per 100ml. For graphite you want enough slabs to get current between 30-40mA/cm2
and chloride concentration over 50g per liter or they will corrode away quickly. You may need to insulate or cool the cell depending on the
conditions, the electrolysis causes some heating but it may not be enough if your cell is larger.
[Edited on 15-2-2014 by testimento]
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AJKOER
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Here is a 'cool' experiment for winter when a nice fluffy snow is available. Try to prepare Chlorine hydrate by filling a vessel with Cl2 gas, some
NaCl and adding some fresh snow followed by shaking. The Chlorine hydrate may (I had not yet performed this experiment) separate out as almost white
crystals.
I suspect that I once formed the Chlorine hydrate by accident when a large amount of dilute HOCl (from adding chlorine bleach, NaOCl, to vinegar), in
a vessel with a large mouth covered loosely with a thin clear plastic wrap, was left standing in sunlight surrounded by snow. Some of the hypochlorous
acid is decomposed by light as follows:
2 HOCl --> 2 HCl + O2
Followed by the formation of dilute Chlorine water:
HCl + HOCl = Cl2 + H2O
However, with prolonged standing in light, the acidic nature of the solution, I suspect from the increased smell of Chlorine, increased (either via
the formation of HClO3 or Chloroacetic acid). This could move the above reaction to the right (more free Chlorine) permitting the formation of
Chlorine hydrate with freezing temperatures. There is a supporting reference for this claim quoting from a source ("A review of the oldest known
gas-hydrate - The chemistry of chlorine hydrate", link provided below) to quote:
"Thenard [4] described chlorine hydrate as a solid formed during the cooling of the aqueous solution of chlorine."
and also the author's discussion around equation [23] with the in situ formation of the hydrate from the action between cold HCl and hypochlorite.
My product separated out (actually floated to the top) in the form of large feathered crystals. There is also a source for this observation (same as
above) to quote:
"Crystals left below 0 C for a few days sublimed with the formation of large brilliant crystals which are delicate prismatic needles extending from
1.25 cm to 5 cm."
Note, on standing in strong sunlight, extracted crystals will melt giving off a strong Chlorine smell.
[EDIT] To reduce waiting for the solution to become more acidic, adding a limited amount of chloride (for example, NaCl or best ZnCl2, an acidic
chloride) can also move the equilibrium to more free Chlorine encouraging the possible formation of the Chlorine hydrate, albeit, at a corresponding
lower temperature as the chloride concentration increases (adding chloride is actually cited in a Patent for preparing the Chlorine hydrate, discussed
below).
-------------------------------------------------------------------------
Other than for display, the Chlorine hydrate may be of use as either a storage device or for direct use in chlorination. It remains a solid in a open
tube until 9.6 C, above which point it decomposes into Cl2 gas and Chlorine water with effervescence. It will remain as a solid in a closed tube below
28.7 C (or 83.7 F), see http://books.google.com/books?id=h4_zgw9DZ1cC&pg=PA103&a... .
It is said to be prepared variously when chlorine gas (or liquid Cl2, see Patent 4,678,656 at http://www.google.com/patents/US4678656 ) is combined with water (or water containing ZnCl2,..) cooled in ice, or by mixing Cl2 gas and steam and
cooling (see Patent 3,908,001 at http://www.google.com/patents/US3908001 ).
The cited formula for the hydrate is Cl2.yH2O where y is between 6.01 and 7.63 depending on its preparation path (see "VARIABLE COMPOSITION OF
CHLORINE HYDRATESSYNTHESIZED UNDER VARIOUS CONDITIONS", link: http://www.google.com/url?sa=t&rct=j&q=&esrc=s&a... ).
A good discussion of Chlorine hydrate can be found in "A review of the oldest known gas-hydrate - The chemistry of chlorine hydrate", link: http://www.google.com/url?sa=t&rct=j&q=&esrc=s&a...
A cautionary note, Chlorine hydrate will react with NH3 and ammonia salts to form N2, HCl and small amounts of the yellow oily and deadly explosive
NCl3, with less forming with ammonia than its acidic salts.
-------------------------------------------------------------------
A cool part of working with Chlorine hydrate is that upon melting:
Cl2. 7 63/100 H2O --Heat--> Cl2 + H2O + 6.63 H2O = HCl + HOCl + 6.63 H2O
implying that the by weight Hypochlorous acid strength may approach 30.5%. At this concentration, it is unstable readily forming HClO3 via
disproportionation. As such, I will try mixing the dry Chlorine hydrate and say, Na2CO3 (or NaOH) plus KCl, and any needed warming, may be a
convenient path to chlorate production with KClO3 separating out on cooling.
[Edited on 25-2-2014 by AJKOER]
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cyanureeves
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i wanted to distilled nitric acid in my backyard using my retort and exploit this cold weather and boy was that a mistake i got an itchy throat now
and ran here and there moving all my set up to different locations.everywhere i looked there was smoke,smoke on the siding,smoke on the trash cans and
trees even,everything was steaming.the nitric dripped fine but the sulfuric stuck to every wet surface and we are having sleet weather right now so,
Never again! this would still happen under a fume hood wouldn't it? or am i missing something about fume hoods?do they employ filters of some sort?
carbon?wont a filter impede exhaust flow?i felt so guilty doing this whenever things go bad i always feel like a drug cook.
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