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White Yeti
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[*] posted on 22-8-2011 at 02:15
Small scale chlorine generation.


Hello everyone,
I have a hallogen collection and my chlorine sample is getting very fade, probably because the chlorine escaped over time. I would like to replace my sample with a new vial with a small sample of chlorine.

Chlorine can be generated via electrolysis and the chloralkali process, but I would like to fill the vial with only a small amount of chlorine.

I thought about mixing three acids together to obtain chlorine, but I'm not sure if it will work.

First, mixing concentrated aceitic acid with sodium hypochlorite will yield hypochlorous acid and sodium acetate. Then I thought about adding miruatic acid so that there is a mix of HClO and HCl. These acids exist in equilibrium with water and chlorine gas.

But I ran into a problem, cooling the mix increases the solubility of chlorine gas and heating the mix will cause disproportionation.

How can I shift the equilibrium so as to obtain chlorine gas?

Thanks in advance and thanks for reading.
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[*] posted on 22-8-2011 at 02:19


Use a standard protocol, such as addition of hydrochloric acid to potassium permanganate, manganese dioxide, or trichloroisocyanuric acid (TCCA).
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[*] posted on 22-8-2011 at 03:20


You can make chlorine simply by adding excess HCl to bleach. No acetic acid needed as intermediate step. You will get enough Cl2 to make a nice green cloud, even from 5% household bleach. The equilibrium is far on the side of Cl2.

If you want to make your own sample of Cl2 and want to seal it, then assure that it is superdry. Some contamination with air or CO2 is no problem, you won't see the difference, but water is killing for your sample. It makes it look less attractive (condensation on the glass) and over time, the sample fades, due to an irreversible reaction in which HCl and O2 are formed).

I made my sample of Cl2 by generating Cl2 from calcium hypochlorite and HCl. I made this in a 300ml erlenmeyer. Then I took a syringe with PVC tube, connected to the tip. In the syringe I had some powdered P4O10 and in this I sucked 60 ml of Cl2 gas and swirled this around for a few minutes to have this in contact with the P4O10. Then I replaced the PVC tube on the tip by a fresh and dry one and I pressed the Cl2 in a 20 ml glass ampoule, which I then quickly sealed with a hot propane torch.

You can see my sample on the Wikipedia page of chlorine:

http://en.wikipedia.org/wiki/Chlorine

This sample is a few years old and it still has a nice green color. No fading at all. I think it is 80...90% chlorine, the rest being CO2 from the calciumhypochlorite and air.




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[*] posted on 22-8-2011 at 03:58


So the wiki photo of chlorine is a sample by woelen, huh? That's quite an honour!



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[*] posted on 22-8-2011 at 05:30


Keep it simple. Just use bleach and hydrochloric acid. Pass the resulting gas mixture through concentrated aqueous solution of potassium permanganate and then through a pipe full of coarse grade calcium chloride. The permanganate oxidizes any hydrogen chloride leftovers, and CaCl2 removes the water.

You can prepare a nice ampoule using a test tube, and just let the gass fill it up untill it starts coming out, which is evident by the smell (use caution!).
Then slowly remove the glass tubing you used for filling, while still passing the gas, and then quickly seal the neck in the flame. That should give you pretty pure chlorine sample.

[Edited on 22-8-2011 by Endimion17]




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[*] posted on 22-8-2011 at 05:51


My personal experience is that CaCl2 alone does not dry sufficiently. When a sample, prepared in this way, is put in a cold place, then the inside of the glass can become somewhat turbid, due to condensed water. The sample will not fade, but the condensed water on the inside of the glass is not nice at all. When the chlorine is dried with P4O10, then a sample, prepared at room temperature, remains clear even when the temperature goes down to the freezing point of water.

If you don't have any P4O10, then I expect that first passing it through some CaCl2 and then bubbling the gas through a long column of concentrated H2SO4 also makes it sufficiently dry. CaCl2 is a nice drying agent, but for some purposes you need 'stronger' stuff. I sometimes use a trick which saves a lot of P4O10. I then take 90% CaCl2 and 10% P4O10 and shake the whole mass and use that for drying purposes. This works nearly as good as pure P4O10 (pure P4O10 quickly becomes less effective, it becomes covered with a sticky/glassy layer, while the inside still is powdery dry).

[Edited on 22-8-11 by woelen]




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[*] posted on 22-8-2011 at 08:53


How about a cold trap? Easy enough to make, especially for just getting water vapor out.
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[*] posted on 22-8-2011 at 10:00


Quote: Originally posted by woelen  
My personal experience is that CaCl2 alone does not dry sufficiently. When a sample, prepared in this way, is put in a cold place, then the inside of the glass can become somewhat turbid, due to condensed water. The sample will not fade, but the condensed water on the inside of the glass is not nice at all. When the chlorine is dried with P4O10, then a sample, prepared at room temperature, remains clear even when the temperature goes down to the freezing point of water.

If you don't have any P4O10, then I expect that first passing it through some CaCl2 and then bubbling the gas through a long column of concentrated H2SO4 also makes it sufficiently dry. CaCl2 is a nice drying agent, but for some purposes you need 'stronger' stuff. I sometimes use a trick which saves a lot of P4O10. I then take 90% CaCl2 and 10% P4O10 and shake the whole mass and use that for drying purposes. This works nearly as good as pure P4O10 (pure P4O10 quickly becomes less effective, it becomes covered with a sticky/glassy layer, while the inside still is powdery dry).

[Edited on 22-8-11 by woelen]


My drying column was half meter long and filled with CaCl2, and the ampoule never gets below some 15 °C, so I never got to see any condensation...
I have the pentoxide, but I hate working with it.




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[*] posted on 22-8-2011 at 11:12


Thanks for the feedback!
I will definately try adding bleach to muriatic acid. Between calcium sulphate and chloride, which one is better for drying chlorine? I have lots of both chemicals on hand, and anydrates are not difficult to make. I don't have phosphorus pentoxide, so I have no choice but to use the calcium salts. Would it be overkill if I flushed the vial with helium gas before filling? I don't like working with hallogens, so I want to make a sample that will last a long time.

That chlorine sample sure looks nice woelen, I'd be happy if mine even gets close:)
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[*] posted on 23-8-2011 at 11:18


You can also use Bleach + NaHSO4 in place of HCl. Volumes of Cl2 generated.

Sodium bisulfate or Sodium hydrogensulfate is commonly sold as a pH balancer for a spa.

If you have no strong acid, but Ca(OCl)2 mixed with CaCl2 (Bleaching Powder), then use carbonated water per the reactions:

[1] Ca(OCl)2 + CO2 + H2O => CaCO3 + 2HOCl

[2] HOCl + Cl- => Cl2 + OH-

and in the presence of CaCl2, Chlorine is generated as the Ca(OH)2 falls out of solution:

2 HOCl + CaCl2 --> 2 Cl2 (g) + Ca(OH)2 (s)

per the source:
http://www.scienceforums.net/topic/34876-potassium-chlorate/

I believe the reaction [2] is more correctly written as reversible and Cl2 is only generated in the case of an insoluble hydroxide. This explains why CO2 does not generate Cl2 for a solution of NaClO + NaCl (e.g., Chlorox Bleach).

Another source gives:

Ca(OCl)2 + CaCl2 + 2 CO2 --> 2 CaCO3 + 2 Cl2

which is say to be accelerated in the presence of water and temperature. Also, for dry bleaching powder:

Ca(OCl)2 + CaCl2 --> 2 CaO + 2 Cl2

which the author describes as the reverse of the chlorination reaction. Hence, my reversible reaction comment previously.

The authors also note on heating the solid, the primary reaction is a decomposition:

Ca(OCl)2 --> CaCl2 + O2

which is accelerated in the presence of Fe or Mn, and over 150 C may occur explosively. Otherwise, in solution, the disproportionation is 70% of the reaction with the production of Calcium Chlorate and CaCl2.

Source: Handbook of Detergents: Production, Volume 142 by Uri Zoller and Paul Sosis, page 446.
http://books.google.com/books?id=dXn3aB1DKk4C&pg=PA247&a...





[Edited on 23-8-2011 by AJKOER]

[Edited on 23-8-2011 by AJKOER]

[Edited on 23-8-2011 by AJKOER]
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[*] posted on 21-9-2011 at 16:35


Mr Chlorine comes to the rescue.
There are many ways to make chlorine, don't mind if I repeat some reactions listed above.

2HCl + H2SO4 --> Cl2 + H2SO4 + 2(H+) // This reaction I don't generally like because you waste H2SO4 but it is a way. Where the H2SO4 ionizes the Hydrogen and leaves chlorine to react with itself.

NaClO + HCl ----> NaOH + Cl2 // A Good one I learned by Nurd Rage, but I don't generally like the Idea of 2 liquids combined to make a gas, because concentrations are extremely hard to predict.

Ca(ClO)2 + 2HCl ----> 2Cl2 + Ca(OH)2 //This one is by far the most efficient because they are easy to buy and very reliable. This is the one I most Recommend because of it's solid and liquid mix, it makes Chlorine generation very simple and safe.

Techniques:

Normal apparatus how a practical chemist would go about this is use a round bottom flask or other simple construct made out of glass.
I'm not a practical chemistry, in fact I'm probably the most unorthodox chemist on this forum. So How I go about this.....

Is use a Plastic bottle, that was used for holding water. You poke a hole at the top of the cap and hot glue in a tube. The tube was acquired from Ethernet tubing that holds the 8 wires. When you add the chemicals that produce chlorine, you should add in the Solid, and add in Diluted amounts of the other reactant at a time(See **Note). Adding the cap with the tube is self explanatory. When you need the amount of chlorine you need you just slowly squeeze the bottle. When you are all finished with the reactant, Either Cap the bottle with a non punctured hole, and throw away making sure that it isn't going to react any more, or just let it vent outside.

**NOTE: Please use dilute amounts of HCl when you add in. You don't want excessive amounts of it to be lost just because you added to much HCl molecules at a time.

[Edited on 22-9-2011 by MrTechGuy1995]
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[*] posted on 21-9-2011 at 17:25


Quote:
2HCl + H2SO4 --> Cl2 + H2SO4 + 2(H+)


This isn't a balanced equation (where did the positive charge come from, i.e. where did the electrons go?), nor do I believe it would be correct if you just substituted H2 for 2(H+). Further I doubt this proceeds easily (though H2SO4 *is* an oxidizer), since HCl may be produced from various salts in combination with H2SO4, and the product gas is not normally contaminated with Cl2. Also, Brauer suggests drying HCl gas with H2SO4. So, where did you get this idea? Reasoning that if HBr does this surely HCl must also?

Quote:
NaClO + HCl ----> NaOH + Cl2


Not really (if someone uses your stoichiometry they will not use enough HCl). Things do not become more basic as you dump in HCl. The full equation here is
NaOCl + 2HCl -> NaCl + H2O + Cl2 . Your equation for calcium hypochlorite contains the same error.

Quote:
Mr Chlorine comes to the rescue.


Self-promotion is hard when you make elementary mistakes.
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[*] posted on 21-9-2011 at 17:29


Quote: Originally posted by bbartlog  

Quote:
Mr Chlorine comes to the rescue.


Self-promotion is hard when you make elementary mistakes.


Verry funny!!!




I never asked for this.
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[*] posted on 21-9-2011 at 18:10


Quote: Originally posted by bbartlog  
Quote:
2HCl + H2SO4 --> Cl2 + H2SO4 + 2(H+)


This isn't a balanced equation (where did the positive charge come from, i.e. where did the electrons go?), nor do I believe it would be correct if you just substituted H2 for 2(H+). Further I doubt this proceeds easily (though H2SO4 *is* an oxidizer), since HCl may be produced from various salts in combination with H2SO4, and the product gas is not normally contaminated with Cl2. Also, Brauer suggests drying HCl gas with H2SO4. So, where did you get this idea? Reasoning that if HBr does this surely HCl must also?

Quote:
NaClO + HCl ----> NaOH + Cl2


Not really (if someone uses your stoichiometry they will not use enough HCl). Things do not become more basic as you dump in HCl. The full equation here is
NaOCl + 2HCl -> NaCl + H2O + Cl2 . Your equation for calcium hypochlorite contains the same error.

Quote:
Mr Chlorine comes to the rescue.


Self-promotion is hard when you make elementary mistakes.




I'm not thinking in terms of the side products, that's the reason I write it like that. It is that is that when you have two molecules of HCl mix with 1 molecule of NaClO, It does make NaOH and Cl2.
When you get into higher numbers with a high concentration of NaOH being produced over the NaOCl then the HCl that can't react with the NaClO as frequently, therefore it goes to NaCl and H2O. That is where the Acidity and Alkalinity turns to a neutralization.
Same concept with the Ca(ClO)2.
It's theoretical vs. Experimental

As for the H2SO4, It wants to disassociates the HCl---> (H+)+(Cl-).
The (H+) remains in solution and could be oxidized by the H2SO4. It gets complex from there, that I don't understand how it would react. H2SO4 is hydrophilic, and likes to attract H+ ions.

When you get to the Replacement of the hydrogens in the (SO4-2), with the group 1 or 2 chlorides, then you will get HCl gas.
It's where the H2SO4 wants to release the H2, because it is more reactive than hydrogen.

I probably made a mistake in there. It's hard to explain in text, I'd rather use a whiteboard.

[Edited on 22-9-2011 by MrTechGuy1995]

[Edited on 22-9-2011 by MrTechGuy1995]
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[*] posted on 21-9-2011 at 18:21


Quote: Originally posted by MrTechGuy1995  

2HCl + H2SO4 --> Cl2 + H2SO4 + 2(H+) // This reaction I don't generally like because you waste H2SO4 but it is a way. Where the H2SO4 ionizes the Hydrogen and leaves chlorine to react with itself.


Hydrochloric acid does not react with sulfuric acid, at least not at room temperature. Mixing hydrochloric acid with concentrated sulfuric acid will just cause all the HCl to come out as a gas. Indeed, this is an excellent way to prepare anhydrous HCl.

It would, however, likely be possible to oxidize HCl by passing the gas into boiling 40% sulfuric acid.

Hot concentrated sulfuric acid oxidizes HBr very readily, and 98% concentrated sulfuric acid oxidizes hydroiodic acid, HI, at room temperature, although dilute sulfuric acid does not.

Quote: Originally posted by MrTechGuy1995  

NaClO + HCl ----> NaOH + Cl2
but I don't generally like the Idea of 2 liquids combined to make a gas, because concentrations are extremely hard to predict.


The reaction is actually:
NaOCl + (2)HCl --> NaCl + H2O + Cl2

Because bleach actually contains both chloride ions and hypochlorite ions, you can actually just add in some NaSO4H, or dilute H2SO4, into the solution, and it will give off chlorine.
This is because the bisulfate is acidic, and in the presence of chloride ions (such as from NaCl), it transiently forms HCl, which can be oxidized by the hypochlorite.

Solutions of chlorine in water actually have an equilibrium:

Cl2 + H2O <==> HCl + HOCl

(a little trivia for those readers with more advanced chemical knowledge, there is actually also a very small equilibrium with dichlorine monoxide, Cl2O, in water solutions saturated with Cl2. This small equilibrium is normally insignificant, but can have significant effects in certain reactions, such as the oxidation of ferricyanide)

Also, if you want to prepare oxygen gas, just mix hydrogen peroxide with bleach.

NaOCl + H2O2 --> NaCl + O2

Hydrogen peroxide even reacts with chlorine, the net equation in water is:

H2O2 + Cl2 --> (2)HCl + O2

(more trivia, hydrochloric acid causes H2O2 to gradually decompose to oxygen. If 30% concentrated HCl is used in excess, there will be a slight, but distinct, odor of chlorine. Nearly all of the escaping gas, however, will still be oxygen. The reaction is fairly slow, even with concentrated acid there is only tiny bubbles, which last for several hours)

The reaction with of H2O2 with bromine is just the opposite, however. Hydrogen peroxide readily oxidizes hydrogen bromide.

(2)HBr + H2O2 --> (2)H2O + Br2

[Edited on 22-9-2011 by AndersHoveland]
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[*] posted on 21-9-2011 at 18:29


Quote: Originally posted by AndersHoveland  
Quote: Originally posted by MrTechGuy1995  

2HCl + H2SO4 --> Cl2 + H2SO4 + 2(H+) // This reaction I don't generally like because you waste H2SO4 but it is a way. Where the H2SO4 ionizes the Hydrogen and leaves chlorine to react with itself.


Hydrochloric acid does not react with sulfuric acid, at least not at room temperature. Mixing hydrochloric acid with highly concentrated sulfuric acid will just cause all the HCl to come out as a gas. Indeed, this is an excellent way to prepare anhydrous HCl.



Not to sound like a troll... But When I mixed Conc H2SO4 with 20% HCl (aq), a greenish gas appeared, it didn't smell like exactly like chlorine made in other methods, but I think thats because of the H2SO4 taking the H20 out of the air, making it less pungent.

If that isn't chlorine I haven't a clue what it was.
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[*] posted on 21-9-2011 at 18:36


This could be possible. I am not entirely sure.

If sulfuric acid is capable of oxidizing hydrochloric acid to any extent, the sulfuric acid would have to be extremely concentrated (probably >98%).

For more information about this, you might see:
http://www.chemguide.co.uk/inorganic/group7/halideions.html
The site, however, is not very forthcoming about the specifics of the reactions, and the exact concentration of the sulfuric acid is probably the determining factor. So one should be somewhat sceptical about the information in the link.

[Edited on 22-9-2011 by AndersHoveland]
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[*] posted on 22-9-2011 at 03:36


Alkaline batteries have a fair bit of manganese dioxide in the black paste, and there is a lot of that paste in a lantern battery.

It is also not hard to get to that paste, as the casing is zinc, not steel. A pipe cutter works brilliantly for opening batteries, I can get the lithium ones open in about 15 seconds.

Dump the paste into a coffee paper and pour water over it to wash the base out. If you then pour hydrochloric on, you'll be met by a nice filthy cloud of chlorine. Enough will come off a lantern battery's paste that you'll be able to see the green cloud in an open beaker - meaning there is a good deal of it there. In fact, there's enough from one lantern battery to do you some serious harm. Blogfast termed it 'the chlorine factory' when discussing making manganese salts from battery paste, which is an amusingly accurate description. There is no way you can do an entire lantern battery in a beaker in the house, it'll raise the concentration in the air far too much for it to be okay.

Note that even 100% atmospheres of chlorine do not have much colour! They are not bright green, they are pale straw yellow / green.

A manganese dioxide generator I was using for some sulphur chloride fun!


The generator was filled with this colour within minutes, and it has now been running for 12 hours. That is as green as it gets.


Quote:
If that isn't chlorine I haven't a clue what it was.


Put a pinch of table salt in a test tube and then drip a drop of concentrated sulphuric acid onto it. Give it a smell.

If the smell matches, you have made hydrogen chloride; as that is the result you'll get for concentrated sulphuric on sodium chloride.

Hydrogen chloride is as much a sensation as a smell; a sensation of a skewer going up ones nosey wosey. It does smell a little like chlorine, but the two smell quite different, and chlorine does not produce such an immediate and sharp sting.

The combination of concentrated sulphuric and hydrochloric does produce hydrogen chloride. As the hydrochloric already has as much hydrogen chloride dissolved in it as will go, the dehydrating effect of the sulphuric forces it to leave the solution. It will also leave, partially, due to the heat generated when the two are mixed.

However, hydrogen chloride is not green. It is colourless. It tries to redissolve it's self as soon as it leaves the solution, which it does by going into the moisture in the air. So you will often see it as a white steam or clouds. Fire fighters responding to an industrial release of it will walk towards the leak spraying the air with the hose, to knock it all back onto the floor as hydrochloric acid; blocking the drains and putting things down to soak it up as it goes.

It is stupidly corrosive, and will stain or rot all the metallic fittings around your beaker (including the 'stainless' fittings), so best done outside.

I have not seen chlorine coming off the two, or smelt it, and I am usually using 98% sulphuric and 36% hydrochloric, so I'm not sure about why yours is green. Do you mean the gas it's self?

If you can see chlorine, the smell of it will knock you sideways; it should absolutely stink of indoor municipal swimming pools.

Quote:
Would it be overkill if I flushed the vial with helium gas before filling?


Helium is overkill yeah, because it's very expensive, a rare element and it's lighter than air, so it is prone to floating out of containers.

Argon is far cheaper, just as unreactive and sinks, so it'll sit still; which is why the majority of chemistry using inert gases is done with nitrogen or argon. Nitrogen can form nitrides with some things, but is still used quite frequently since it's cheaper again.

A bigger issue is moisture.

If you want it to last a long time, you will need to connect the vessel to high vacuum and then warm it up with a hot air gun or blow torch; glass clings onto a microscopic layer of water on it's surface. There are also OH groups dangling off the surface of the glass, which higher halogens can react with. That needs cooking under vacuum to get it all out.

Ideally, you need to allow the chlorine back into the ampoule without disconnecting anything, or moisture will get back in there as the chlorine goes in.

My manganese woes;


Dioxide works fine, but producing known moles of chlorine from the eBay and battery paste stuff can be difficult.

Battery paste is full of activated carbon, so it is not appropriate to assume it is manganese dioxide you are measuring.

The eBay stuff can be recycled, repackaged, battery paste. NurdRage got a bag of it full of sand. :D

To get a well known molar volume from it, it is best to save the mess from previous generators. Dump it out into a coffee paper, collect the filtrate and boil it down to recover manganese chloride (it'll probably be yellow from iron chloride contamination). On heating, the manganese chloride will release yet more chlorine and go back to dioxide. But all the activated carbon and other cutting rubbish (e.g. eBay sand) will now be stuck in the filter paper and can be binned***, so the result will work much more nicely when reused. There are ways to get rid of the iron contamination as well.

This may seem tedious, but you can apply the same logic to pool chemicals, in that the tablets and lumps may not be pure material or the concentration they state. At least with the manganese, ensuring it is pure and measuring out a precise amount is not particularly difficult after it's been through the process once.

***I have some sitting around in the middle of that process actually, I'll see if I can take some photos in a minute.

{edit}Found 'em;

Hmmmmm, not sure about that


What's left in the filter paper. I tried dripping some more hydrochloric over it. Nil poi! That'll be activated carbon, as found in Evereadius C'sellius.


The filtrate, full of manganese chloride. The 'waste' here will actually be purer than the original, so reuse it.


[Edited on 22-9-2011 by peach]




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[*] posted on 22-9-2011 at 05:23


I've noticed two things and I have to speak up. :D

1. Lots of people here use their noses to smell stuff. Skewer up the nose? If that happens, your olfactory nerve is for a treat...
Never use nose inhaling. Use mouth and nose, and inhale just to let the vapors reach yout tongue and just to touch the olfactory nerve. You'll get much more information, with reduced danger. Letting the vapors/gases to flow over the nerve can be devastating more than you think.

2. Coffee filters? Come on. That's recycled paper. At least use apothecary filter paper sold in large sheets. It's not expensive.
Recycled paper can leech stuff, and white filter papers used in laboratories are designed to be more inert. Or just buy those round papers that are tested against conc. HCl/HF.




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[*] posted on 22-9-2011 at 05:54


Quote: Originally posted by Endimion17  
2. Coffee filters? Come on. That's recycled paper. At least use apothecary filter paper sold in large sheets. It's not expensive.
Recycled paper can leech stuff, and white filter papers used in laboratories are designed to be more inert. Or just buy those round papers that are tested against conc. HCl/HF.


I have to slightly disagree. Coffee filters, even paper towels will remain the choice of quite a few hobbyists. I started with them and still use them quite often. Coffee filters have some disadvantages but leeching is probably not that bad: remember that this is a food product.

As regards apothecary paper you certainly can't get it from pharmacies here in the UK, these have all become quite dumbded-down places where little or no preparation goes on and where it's all nothing more than buying and selling. If I asked the one around the corner for some AP she'd look at me as if I'd come straight from Mars...




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[*] posted on 22-9-2011 at 11:48


There is another easy reaction that can be used to produce chlorine. Chlorates and hydrochloric acid react to liberate chlorine.

KClO3 + (6)HCl --> KCl + (3)H2O + (3)Cl2

Potassium perchlorate, however, does not react with hydrochloric acid. This is a useful reaction to remove traces of chlorate from lab prepared perchlorate.
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barley81
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[*] posted on 22-9-2011 at 11:56


Actually, chlorate and hydrochloric acid react to produce chlorine dioxide along with the chlorine. Woelen's site has an experiment which involves this.
http://woelen.homescience.net/science/chem/exps/clo2/index.h...
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AndersHoveland
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[*] posted on 22-9-2011 at 12:11


If excess hydrochloric acid is used, and it is dilute, there will only be chlorine, since ClO2 actually oxidizes HCl.

Woelen writes that:
"The two idealized 'extremes' for the reaction equation are.
ClO3– + 5Cl– + 6H+ → 3Cl2(g) + 3H2O
5ClO3– + Cl– + 6H+ → 6ClO2(g) + 3H2O"

And I would tend to agree with his view of this. The reactions are quite complicated, and a discussion of the reactions can be found another thread in this forum:
"Chlorite and hypochlorite interactions"
http://www.sciencemadness.org/talk/viewthread.php?tid=16378



[Edited on 22-9-2011 by AndersHoveland]
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[*] posted on 22-9-2011 at 12:17


From Woelen's website:
Quote:

The reaction of chlorate with concentrated hydrochloric acid is very complex. Two possible idealized 'extremes' can be given for the reaction equation. The net real reaction will be a linear combination of these two extremes, resulting in a gaseous mixture of approximately 50% Cl2 and 50% ClO2 and a higher percentage of ClO2 in the liquid (this gas dissolves better in water than does Cl2). The exact ratio cannot easily be determined and depends on the concentration of the acid and on the temperature. Pure chlorine or pure chlorine dioxide can, however, never be obtained with the reaction between a chlorate and hydrochloric acid.


And chlorine dioxide (at least with sodium chlorite and hydrochloric acid in a test tube) apparently does not oxidize HCl. If it did, the reaction between sodium chlorite and hydrochloric acid would not produce pure chlorine dioxide, but a mixture of chlorine dioxide and chlorine gas.

[Edited on 22-9-2011 by barley81]
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[*] posted on 22-9-2011 at 12:21


Quote:

"Pure chlorine or pure chlorine dioxide can, however, never be obtained with the reaction between a chlorate and hydrochloric acid."


I would have to disagree with this.
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