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I am a fish
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| Quote: | Originally posted by Organikum
FeSO4 straight from the box... |
FeSO4 is pale green. It looks like yours has been oxidised into the Fe(III) state.
[Edited to fix typo]
[Edited on 18-4-2004 by I am a fish]
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Organikum
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Partially decomposed to ferric sulfate and this hydrolyzed - thats possible and probable, but it worked well so it will have been still mostly FeSO4.
I have some other boxes where the FeSO4 is - as you told - of this pale greenish color. As this was just a "proof of principle and setup" it
didnt matter, but you are right. I think I will post another picture after the next where the color is "right" - not to stirr confusion.
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chloric1
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FeII
If you have Battery electrolyte(H2SO4) you can moisten the FeSO4 with it and add roofing nails to regenerate the green color if you so wish. Also, a
little free H2SO4 will limit Cl2 solubility in water and may help to expel this. Just a thought. may not be practical if your time is as precious as
mine.
Fellow molecular manipulator
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Proteios
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Cl2 w NaBr was a good way of making Br2 (just collecting in vial in iced water). However Cl hardened all the tubing i ever got hold of (either
chlorination of the tubing, or the plasticiser), Br really eats tubing, presumably due to it propensity to condense as a liquid on almost everything.
just another happy memory
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ApprenticeCook
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Chlorine gas production
Addition funnel on RBF via a rubber tubing T-joint. Add HCl (~30%) to KMnO4 in the RBF while heating and stirring at high speed, the tube which is
connected to the T-joint has a CaCl2 inline dryer then is vented to the system that requires the now pretty much anhydrous chlorine gas.
Check Rhodiums site:
https://www.rhodium.ws/chemistry/eleusis/chlorine.html
https://www.rhodium.ws/chemistry/equipment/inline.gas.dryer....
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Organikum
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Please read the first post in this thread.
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BromicAcid
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This is just a spew of information, some new, some old.
From Preparative Inorganic Chemistry
MnO2*xH2O [~100g] + 4HCl [145.88 ] ---> MnCl2 + (x+2)H2O + Cl2 [70.91g]
Concentrated, air-free hydrochloric acid (d 1.16) is added dropwise to precipitated hydrated manganese dioxide (e.g., the 86% pure commercially
available material) [Note: PbO2, BaO2, KMnO4, are also listed] in a flask equipped with a dropping funnel and a gas outlet tube. The gas formation
may be regulated by moderate heating.
The chlorine thus formed is passed through water then H2SO4 then CaO then P2O5 and liquefied in a receiver cooled with Dry Ice-Acetone bath.
[Yeah I know you didn't want to pass though H2SO4 and you wanted an easy prep and this was already covered]
Electrolysis of an NaCl solution with HCl in an electrolytic cell described by Bodenstein and Pohl. The oxygen content of the Cl2 produced in this
manner is .01%.
From Complete Treatise on Inorganic and Theoretical Chemistry
Extremely pure Cl2 can be prepared via thermal decomposition of AuCl3.
To remove the last traces of HCl from Cl2 it is recommended to bubble though a CuSO4 solution.
You can bubble Cl2 into cold water to form the hydrate which keeps well below 9 degrees and in the dark. Heating slightly furnishes chlorine with
minimal impurities.
The action of potassium or sodium chlorate on hot concentrated HCl leads to chlorine formation but if the temperature is too low then chlorine oxides
are the favored product which may explode.
Heating a slurry of a chloride and a nitrate in concentrated H2SO4 furnishes nitrogen oxides and chlorine gas. The resultant mixture of gasses is
passed though H2SO4 where the nitrogen oxides are retained.
Action of HCl(aq) on Ca(OCl)2 can occur with or without a binding agent for the calcium hypochlorite. Heating the reaction increases the O2
concentration of the output gasses.
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Organikum
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Thanks Bromic!
"You can bubble Cl2 into cold water to form the hydrate which keeps well below 9 degrees and in the dark. Heating slightly furnishes chlorine
with minimal impurities."
This sounds like a very feasible idea for storage of Cl2 without compressing it.
A remark on the HCl and hypochlorite procedure: The use of dil. H2SO4 (battery acid) instead of HCl is advised as the sulfuric acid retains the
water.
I meanwhile overcame my aversion of using conc. H2SO4 for the drying of the Cl2 - it is not so bad as it looked like and gives superior results
compared to CaCl2 for example.
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I am a fish
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| Quote: | Originally posted by Organikum
Thanks Bromic!
"You can bubble Cl2 into cold water to form the hydrate which keeps well below 9 degrees and in the dark. Heating slightly furnishes chlorine
with minimal impurities."
This sounds like a very feasible idea for storage of Cl2 without compressing it.
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Though perhaps not the most dependable way of storing it:
1. Damn! There's a power cut.
2. Why has everything near the fridge gone green?
3. What's that odour? It smells like...
4. Arrrrgh! No! Cough! Help Me! Cough! Aaaaaccckkk!!!!
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Theoretic
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When iron is fused with KNO3 (I think NaNO3 could also work) iron ferrate is formed like so:
2KNO3 + Fe => K2FeO4 + 2NO
which could be reacted with HCl like so:
2K2FeO4 + 10HCl => 3Cl2 + 4KCl + 2Fe(OH)3 + 2H2O.
NO could be recycled and turned back into nitric acid, although the process could present technical difficulties.
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DDTea
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Just skimming through this thread again, I have not seen this mentioned--this was discussed in my Chemistry textbook, and I thought it was very neat.
Have you noticed that if you add even a little bit of HCl to Hypochlorite solutions, you get a flood of Cl2 gas?
This is because Cl2 is much more soluble in basic solutions than in acidic solutions (as you know, the NaOCl would be formed)... So, lowering the pH
causes the Cl2 to suddenly fall out of solution.
So, here's what I propose: dissolve as much CaOCl2 as possible in Warm water. Then the addition of a little bit of a strong acid, e.g.: HCl
(since we don't want to use H2SO4, and HCl is cheaper anyway). The resulting Chlorine would then be led into a separate container, immersed in
dry ice or alternately, sealed and placed in a freezer over night to freeze out the water.
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Prince_Lucifer
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Cl2 gas production.
That is a very clever idea Samosa, certainly worth investigating further 
I wonder if passing the generated Cl2 gas through Conc. H2SO4 would help remove H2O molecules?
Org, with the FeSO4/10%(aq)Ca(OCl)2 procedure, have you had a chance to incorporate an inline drying tube yet?
I would be interested to hear your results, and also your opinion on whether either method listed above would be a suitable Cl2 donor for toluene
chlorination reactions?!
I realise what im about to say is off topic, but it is related. How adversely would Cl2 with an EXCESS of H2O, affect toluene chlorination attempts?
Good work guys, theres no shortage of smart cookies lurking around 
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Organikum
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I am not sure how water will affect the chlorination of toluene - it doesnt hinder acetone chlorination, thus I know.
But oxygen which is often a byproduct of chlorine production WILL block the chlorination of toluene.
And thats the pont of my search:
A STEADY stream of anhydrous and oxygen-free chlorine.
Dropping HCl or H2SO4 onto bleach/bleaching powder/ hypochlorite will produce chlorine. This chlorine wont be evolved in a steady stream and it is wet
and probably not free from oxygen.
- I tried H2SO4 and hypochlorite and bleach and FeSO4. I encountered big problems with suckback and the flow was anything but steady.
- I tried the elctrolysis of zincchloride with a graphite and an Al electrode (to deposit the Zn there). This sucks.
- I tried the electrolysis of NaCl/HCl with graphite electrodes. This work best from all tried up to now but the graphite gets eaten fairly fast from
the aqueous HCl.
- I will try now the electrolysis of zincchloride with graphite electrodes and hope the Zn will separate in flakes and not deposit on the rod and
shorten the cell as it happened with Al (which got also eaten up...)
I am very sure now that chlorine production for organic chlorinations is best to be done by an electrolytic method. Which electrolytic method being
best I dont know by now.
Perhaps I am just an idiot, but every chemical method tried led to a severe chlorine intoxication (I have no fumehood) and ruined the reaction
intended by suckback/clogging of the inlet tube and/or sudden outbursts of Cl2 blowing the shit through the condensor - all chlorinations are highly
exothermic.
This is no fun at all and dangerous too. BzCl is a strong lachrymator and blowing the stuff all over you is all but funny or healthy.
The electrolytic method I am working on now is - as told - most promising. It produced fair results and gave a controllable reaction. And my lungs
still worked afterwards.
Of course I will post my final results here and actually I am sureI will be able to present an "save and easy to build" electrolytic
chlorine generator soon.
Up to then - patience please.
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Proteios
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solutions to suck back.
1)suck back arrestors can be purchased from suppliers.... they generally are not that good, and i would have doubts about their CL resistance
2) run with a carrier gas. The simplest is just attach an air pump to the CL generator. However if you are really keen to lose the O you can just
use a N or Ar cylinder. If gas consumption is too high then you can run on close circuit inert gas. However you will need good plumbing and a good
pump (chem. resistance wise). Carrier gas will also remove many of the problems of the exothermic reaction do to the lower conc. of Cl in the gas.
3) Cl liquifies low (ca -35). Some commerical freezers can get this low. Stage 1 liquify Cl. Stage 2 connect Cl ampoule to kit, and regulate Cl
flow with a hair dryer.(caution... with vapour pressure! There is a pressure explosion hazard here (pressures in excess of 6 bar may be expected),
but as long as the hazard is recognised, and accounted for, there should be no problem.)
4) convert Cl to bromine that is much easier to store. I dunno what you are trying to do with these clorides, but for most organic reactions bromine
will work just as well. again use hair dryer to regulate flow.
[Edited on 15-5-2004 by Proteios]
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Chemtastic
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Ahh, Chlorine...So Deadly, So Green
In this thread:
http://www.sciencemadness.org/talk/viewthread.php?tid=606
Haggis posted that:
| Quote: |
There are many methods for generating chlorine gas, but common bleach and sodium bisulfate does the trick.
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I've seen both in decently pure and concentrated solutions (12.5% NaClO to beat 5% bleach) at my local NAMCO pool store.
I'm still not sure about the reaction though...
NaClO + NaHSO4 --> ???
Anyway, isn't "wetness" of Cl2 just a physical property? If there's no chemical interaction, why not just stick your
chlorine in any freezer for drying...I would think any H2O vapor would form a layer of ice on the container bottom, especially with really cold
temperatures...
Liquefying the Cl2 would work too, but if you haven't any industrial refrigerators in your neighborhood, a dry ice enclosure would probably
work (what's that, -80C?). Finally, just quickly transfer the liquid to a pressure-safe container, like they do with the butane in BIC lighters.
I'd still like to know exactly how sodium hypochlorite and sodium bisulfate react, as, unless someone has a warning against it, I plan to
try it.
EDIT: If this works, I'll be somewhat less upset that NAMCO switched from HCl to NaHSO4 for their ph-Decrease.
[Edited on 27-6-2004 by Chemtastic]
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Chemtastic
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Today, I ended up trying a mixture of acetic acid (as 5% vinegar) and sodium hypochlorite (as 6% bleach) for the production of chlorine. I guess the
reaction would go something like this:
4H+ + 2OCl- --> Cl2 + 2H2O??
I started off with about 50mL of each solution, the vinegar being clear and the bleach a pale yellow (is NaOCl solution naturally yellow, because i
always see it in stores as such?). Upon mixing, there was no bubbling to indicate that anything was happening. However, the solution quickly changed
from pale yellow to a moderately deep amber. I also smelled the chlorine gas odor from about 3-5 feet away.
After the solution had sat on the deck (doing this outside) for about 5 minutes, the entire mixture had turned a shade of bright pink? What caused
this transition, and was it seperate from the initial change? Why was one so fast and the other so slow?
My hypothesis is that the initial reaction rapidly produced most of the possible chlorine gas, but most of this became dissolved in the solution,
giving it it's amber hue. Then, over the next five minutes, the HOCl in solution decomposed in the light (though it was cloudy...) to purely
dilute HCl. However, I don't know what would account for the colors.
Baking soda was added to the final pink mixture, and it did react fairly vigorously, but this could have been due to an excess of vinegar in the
original solution...i only used equal volumes...nothing fancy...
Does anyone have any ideas what could have been in the solution. particularly what caused the colors and why none of the gas supposedly produced was
released as bubbles?
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Saerynide
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Ahhh.... vinegar and bleach.
That was the lie I used to explain to the doctor how I got chlorine poisoned
I said I was cleaning the bathroom
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Esplosivo
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The Cl2 produced could have reacted directly with some of your vinegar forming chloroacids. The pinkish colour you stated seems strange to me - most
probably impurities. Oh btw, Cl2 is soluble in water, I think heating the solution will reduce its solubility and therefore form Cl2 gas.
Theory guides, experiment decides.
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hodges
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Bleach also contains some sodium hydroxide, and vinegar is a weak concentration of acetic acid, so its possible it did nothing more than neutralize
the NaOH. Try adding much more vinegar than bleach and see if you get Cl2 bubbles then.
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trilobite
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Electrolysis with a lead electrode
I believe chlorine will attack a lead electrode, but what will happen after that? Will the PbCl2 formed stick to the electrode or will it drop to the
bottom of the electrolytic cell, and if the former happens, will the PbCl2 layer grow until the cell resistance is so large that the electrolysis
stops?
Generating chlorine gas at a steady rate iinterests me too but so does electrolytic halogenation (bromine included) and platinum eletrodes aren't
my cup of tea either.
Yes, this thread is about new methods, but I'd also like to ask what is wrong with the MnO2 method? H2SO4 dropped to MnO2 and NaCl maybe?
That's the old way of making bromine, using NaBr/KBr instead. Being outdoors is just great for avoiding the worst case scenarios but often there
is no electricity available.
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Saerynide
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| Quote: | | Being outdoors is just great for avoiding the worst case scenarios but often there is no electricity available |
Extension cord!!!!
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ballzofsteel
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Why not just drip conc Hcl onto your TCCA?
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Theoretic
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Reacting hypochlorites with acids would get you HClO and not Cl2. Mix in an equimolar amount of NaCl to your hypochlorite and use twice as much acid,
that will work.
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Organikum
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For electrolysis where Cl or HCl is formed or electrolysed you have to use graphite electrodes, for electrolysis with sulfuric acid lead the material
of choice.
Hypochlorites and mineral acids produces very well chlorine. It is well referenced in the literature and was tried by myself with HCl and H2SO4, to
use the acids diluted is preferred for to avoid heating up.
TCCA and HCl will produce chlorine also. The use of 20% HCl avoids that to much HCl is expelled together with the chlorine.
There is nothing wrong with the MnO2/HCl method if one has lots of MnO2 and HCl available. It is volumetric not very effective for you need 4 mole HCl
to get one mole Cl2. With bleach or TCCA you get 1 mole Cl2 for 1 mole HCl, this makes your setup more compact - important so you want bigger amounts
of Cl2.
[Edited on 29-6-2004 by Organikum]
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ballzofsteel
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sulfuric
Alternatively,you could mix salt with your TCCA and drip sulfuric onto it.
Wouldnt this eliminate the need for drying.
The H2SO4 would release chlorine and absorb water,whilst forming HCl at the same time,which in turn would release more Cl2 and so on.
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