| Pages:
1
2 |
Neil
National Hazard
Posts: 556
Registered: 19-3-2008
Member Is Offline
Mood: No Mood
|
|
Not likely, burnt charcoal losses most of its density and has its pore structure destroyed my marauding oxygen. Just pop some wood into a tin can,
cover the top of the can with tin foil and have a small hole poked through the foil.
Set it next to a fire or set burning charcoal around it, as soon as it stops venting cover the hole with a bit of ash and let it cool; instant
charcoal.
If you soak the wood in an activating agent before you destructively distil it; you get a better performance out of the end charcoal.
I've used home-made charcoal to soak the yellow lemon scented satan spit out of cleaning ammonia with great success.
|
|
|
Chemistry Alchemist
Hazard to Others
Posts: 403
Registered: 2-8-2011
Location: Australia
Member Is Offline
Mood: No Mood
|
|
in other words activated carbon would be the simple explanation? when some of the water evaporates, a black precipitate forms... i dont know what it
is tho...
|
|
|
bbartlog
International Hazard
Posts: 1139
Registered: 27-8-2009
Location: Unmoored in time
Member Is Offline
Mood: No Mood
|
|
I decided to have a try at producing potassium chlorate from commercial bleach (sodium hypochlorite).
First, I took 1240g of 6% bleach (containing one mole of sodium hypochlorite). I added 75g of sodium bicarbonate, reasoning that this would lower the
pH somewhat by turning the NaOH used to stabilize the bleach into Na2CO3 (even after all of the HCO3 was driven out of solution). This I put into a
glass coffee pot, covered with a watch glass (to slow evaporation, initially), and heated on a hot plate so that the temperature stayed in the range
of 88-95C. I left it this way for 18 hours.
At that point the solution had lost any trace of yellowish color. I then removed the watch glass and turned up the heat to produce a rolling boil.
After another five hours, the liquid volume had been reduced to about 400ml. I let the solution cool to -5C (it gets cold in my shed at night). A
large mass of crystals precipitated - two different types, presumably Na2CO3 hydrate and NaCl, were apparent.
The liquid was decanted from the crystals. Only about 150ml of solution (sp gr 1.33) was obtained.
To this I added 100ml of distilled water and 25g (about 1/3 mole) of KCl. This was heated in a 500ml beaker and stirred to speed dissolution of the
KCl. Once a clear solution resulted (at about 50C) I set the beaker aside, covered, to cool.
Once again this was left to cool overnightat about -5C. In the morning I decanted the liquid from the crystals that had formed, and dried them between
paper towels (probably a stupid choice for KClO3, come to think of it). 20g of crystals were obtained (49% of theory). Losses are probably due to
incomplete decomposition, liquid trapped in the initial mass of precipitated crystals, and KClO3 still dissolved in the solution.
No tests for purity have yet been done.
If I do this again I will probably try to use HCl to adjust the pH. I shied away from it this time, having chlorinated myself previously, but the
Na2CO3 complicates things needlessly here.
|
|
|
AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
First, recall that
NaClO + NaHCO3 --> Na2CO3 + HClO
or adding a weak acid (Boric acid has been mentioned) to NaClO will also liberate HClO.
Now, much of the discussion on pH control, in my opinion, could be viewed as relating to HClO production. The reason is based on the following
important reactions:
NaClO + HClO --> NaClO2 + HCl
NaClO2 + HClO --> NaClO3 + HCl
Also, an important catalyst to create Chlorate is the concentration of the solution. Hence, only incidental to the act of concentrating is the
prolonged boiling although not optimal from a temperature perspective.
----------------------
As a final comment, on the disputed disproportionation of concentrated HClO into HClO3 yields an equilibrium involving HClO + HClO3 (on the right) and
some ClO2 (on the left), the reaction I cited is best expressed as:
3 HClO --> 2 HCl (g) + HClO3
See Wikipedia on Chloric acid "Another method is the heating of hypochlorous acid, of which products include chloric acid and hydrogen chloride:
3HClO → HClO3 + 2 HCl "
where the reference is most likely only correct (thanks) as to the products upon distillation of HClO (but I would add HClO and Cl2O to the
distillation products).
In the context of aqueous solution, I agree there can be some ClO2 as:
6 ClO2 + 3 H2O <==> 5 HClO3 + HCl
where the presence of even a small amount of HCl (relative to HClO3) appears to form an equilibrium creating ClO2.
[Edited on 3-12-2011 by AJKOER]
[Edited on 3-12-2011 by AJKOER]
[Edited on 3-12-2011 by AJKOER]
|
|
|
bbartlog
International Hazard
Posts: 1139
Registered: 27-8-2009
Location: Unmoored in time
Member Is Offline
Mood: No Mood
|
|
| Quote: | | NaClO + NaHCO3 --> Na2CO3 + HClO |
Not really... this being a solution we have to think about ion concentrations rather than compounds. And once you heat the solution, excess CO2
(HCO3-) will be driven off as CO2. I added the NaHCO3 reasoning that by turning all of the NaOH into Na2CO3 I would drop the pH to just a little over
11, where before it was likely around 12. pH 7 would of course be better.
| Quote: | | Hence, only incidental to the act of concentrating is the prolonged boiling |
I don't agree. Both concentration and temperature increase the rate, but temperature does so more dramatically. See here: http://www.omegachem.com.au/docs/mega_handbook.pdf (they are trying to *avoid* chlorate formation but have some very nice tables and references).
You can see in the chart they give for rate constants of decomposition that increasing concentration from 5% to 16% speeds up the reaction 4-5x, while
increasing the temperature from 15C to 55C (to say nothing of 90C) speeds it up about 150x.
|
|
|
AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
bbartlog: I believe we agree as I cited (per my first post) the ideal temperature (per one source from recollection) at around 70 C.
My concern is that increasing the temperature to boiling, while concentrating, could lose HClO necessary to chlorate formation.
|
|
|
AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
While my discussion has focused on heating HClO solutions, I believe some may find very interesting what is reported in the literature with respect to
very dilute solutions and the power of diffused sunlight and my actual observations last winter.
From "A treatise on chemistry", Volume 1 By Henry Enfield Roscoe, Carl Schorlemmer, page 192:
"Saturated chlorine water gives off chlorine freely on exposure to the air, and bleaches organic colouring matters. When exposed to direct sunlight it
is, if sufficiently dilute, gradually converted into hydrochloric acid with evolution of oxygen:
Cl2 + 2H20 = 4HCl + 02.
It has been proposed to employ this reaction in measuring the chemical action of light, but the decomposition is not sufficiently regular for this
purpose ; thus Pedler (1) has shown that a solution containing 1 molecule of chlorine to 64 of water undergoes no appreciable alteration during two
months' exposure to tropical sunlight, whilst more dilute solutions undergo more or less decomposition, as shown in the following table
Mols. H20 for 1 mol. Cl2 / Percentage of Cl2 acting on water.
64 no action
88 29%
130 46%
140 29%
412 78%
In the case of more dilute solutions, the reaction in sunlight appears to take place almost completely in accordance with the above equation, except
in so far as small quantities of chloric acid are formed. In diffused daylight, however, a considerable quantity of the latter acid is obtained, so
that in this case the reactions are probably those put forward by Popper (2):
Cl2 + H20 = HCl + HClO
8HC10 = 2HCl03 + 6HCl + 02
Under certain conditions, however, sunlight brings about the reverse change, causing the formation of free chlorine from a mixture of hydrogen
chloride and oxygen (see p. 200).
1 Journ. Chem. Soc. 1890, 57, 613. 1 Annalen, 1885, 227, 161 "
Now, last winter I left out in the sun (partially open to the air) some fresh dilute HClO in a thick transparent glass flower vase (diffused light?),
which I further re-diluted (actually intending to discard hence the diluting to save the pumbling!). After two weeks, I noticed that the solution
developed a much stronger chlorine like smell (so much for my effort to dilute!). Passing NH3 near the top produced a cloud of NH4Cl. Upon discarding
the solution down the drain in an old shower, I noticed that where I splashed some on the shower floor, an intense bleaching action occurred (HClO3?).
I now suspect that the diffused sunlight on the dilute HClO produced HCl and HClO3, the latter acid being so strong as to account for the smell and
bleaching action. The cold temperature and dilution may have assisted in keeping gases dissolved in the solution and preserving the HClO to be acted
upon by the sunlight.
Note, as the vase was covered, but not sealed, some CO2 may have dissolved into the solution. The above authors also report that when Cl2 is mixed
with other gases (including CO2), the volume of Cl2 that is dissolved increases. Thus, this also is a possible factor in producing the results. I will
more rigorously repeat the whole experiment this winter to confirm my observations.
[Edited on 5-12-2011 by AJKOER]
[Edited on 6-12-2011 by AJKOER]
|
|
|
woelen
Super Administrator
Posts: 7977
Registered: 20-8-2005
Location: Netherlands
Member Is Offline
Mood: interested
|
|
Keep in mind that formation of HCl from decomposition of HClO causes formation of Cl2. In solutions, containing only HOCl the only way of forming Cl2
is by means of a complicated disproportionation reaction. In the presence of HCl, the formation of Cl2, however, is fast and easy:
HCl + HOCl --> Cl2 + H2O
This reaction is used for making Cl2 by adding hydrochloric acid to hypochlorites.
I once added dilute HNO3 to pure Ca(ClO)2. When this is done, then a pale green solution is obtained and hardly any gas is produced. When dilute HCl
of similar concentration is added, then immediately there is bubbling and copious amounts of Cl2 are produced.
So, your observation can perfectly be explained. HOCl decomposes to HCl and O2. The resulting HCl then further reacts with HOCl to form Cl2. This
explains the much stronger smell of Cl2.
HClO3 hardly is involved in the bleaching reactions and smell. Solutions of HClO3 are odorless and have no appreciable bleaching capabilities. They
also are much more acidic than solutions of HOCl.
|
|
|
AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
Quote: Originally posted by AndersHoveland  |
Actually, a solution of hypochlorous acid will not disproportionate into chloric acid. The reaction (actually an equilibrium) is:
(8)HOCl <==> (4)H2O + (2)ClO2 + (3)Cl2
|
Actually, from a practical point of view AndersHoveland may be correct, as per cited source below, page 554, "A parallel pathway involving HOCl
instead of ClO- is a 1000 times slower" referring to the multi-stage disproportionation reaction forming chlorate.
Also, same source, on page 553, "Photolysis of aqueous HOCl is also initiated by formation HO and Cl radicals that undergo a series of further
reactions producing hydrochloric and chloric acids and oxygen (87)." Note, this appears to follow the reaction I previously posted attributed to
Popper:
8HCl0 = 2HCl03 + 6HCl + 02
The reference (87) is: "A. J. Allmand, P. W. Cunliffe, and R. E. W. Maddison, J. Chem. Soc., 822 (1925); 655(1927); A. J. Allmand and W. W. Webb, Z.,
Phys. Chem. 131, 189 (1928); L. Bonnet, Rev. Gen. Mater. Color. 39, 29 (1935); K. W. Young and A. J. Allmand, Can. J. Res. 27B, 381 (1949); M. W.
Lister, Can. J. Chem. 30, 879 (1952). "
However, it turns out perhaps most interestingly that in the case of chloride free HOCl, the disproportionation reaction of HOCl proceeds to even
perchloric electrochemically. To quote "concentrated Cl-free HOCl can be oxidized electrochemically to chloric and perchloric acids (97)." Page 554.
The reference (97) is a patented process by World Pat. 9,114,614 (Oct. 17, 1991), D. W. Crawford and co-workers (to Olin Corp.).
Another interesting comment, same source, by the author is "Hypochlorite ion is oxidized to chlorate by ozone (142)." Page 559.
REFERENCE: "DICHLORINE MONOXIDE, HOCl, HYPOCHLORITES", Volume 8.
LINK
http://www.questscan.com/?tmp=redir_bho_bing&prt=Qstscan...
[Edited on 24-12-2011 by AJKOER]
|
|
|
AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
Wrong Link:
REFERENCE: "DICHLORINE MONOXIDE, HOCl, HYPOCHLORITES", Volume 8
http://www.scribd.com/doc/30121142/Dichlorine-Monoxide-Hypoc...
or, do a google search on: DICHLORINE MONOXIDE scribd
New users may have to register with scribd which I would highly recommend given its free and the high quality of the research provided.
|
|
|
| Pages:
1
2 |
|
![[*]](images/xpblue/default_icon.gif){kind=icon}
