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woelen
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[*] posted on 4-11-2007 at 08:37


Yes, electrolysis is a slow process. You could have known if you had looked at the quantities I showed on my webpage ;).

I doubt that you can convert the 120 g of KBr quantitatively into KBrO3 in one run. I would first let it run for one day or so, and collect the KBrO3. Also have a frequent look at your cell. The anode tends to be clogged by crystals of KBrO3 somewhat later in the process and this may completely bring the process to a standstill. Every now and then, you may need to cleanup the anode, scraping off crystals of KBrO3.

If you have been running your cell for some time, then let the liquid cool down in the fridge, such that you have maximum yield of KBrO3. The liquid simply can be decanted (don't throw it away, it can be used with more KBr and water to make a new cell) and the solid KBrO3 must be rinsed with ICE-COLD water, and then recrystallized. You'll see that it can be made amazingly pure with only a single recrystallization. Dissolve the crystals in as little as possible of hot - near boiling - water and then let cool down agian. Finally put it in the fridge. This gives a yield of 90% or so.
The remaining 10% in the liquid can be retrieved as well, with impurity, but this is useful for the more raw pyrotechnic experiments, or for making bromine.




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[*] posted on 4-11-2007 at 08:38


Electrolysis in quantity easily runs days. Is that a problem?

I have some H2PtCl6. The solid powder is orange, while the solution ranges from dilute yellowish green (much like dilute alkaline chromate) to yellow.

Tim




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[*] posted on 4-11-2007 at 10:23


The solution is very orange red, dilute it is yellow.

Palladium solution in nitric acid or aqua regia actually a very very potent yellow colour that looks brown. I will show you all what I mean by this shortly.




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[*] posted on 4-11-2007 at 10:46


Thanks Tim. No, its not a problem.... but I reserve the right to consider it a pain in the bum :P I want it now!!! :D

I was concerned initially about the dissolution of the anode as I said before but it appears to be sitting pretty 3.5hrs into the process. There is brown scum on the surface but I think that is a titanium compound. The outside of some of the the Ti wire was inadvertently oxidised during the glassing process. It went all multicoloured like many metal when they are heated in air and some of it may have oxidised more than that. It didn't matter because I ran 8 wires into the solution with a greater surface area by a factor of at least 2 than the anode. 10ft of the wire, including postage from England, cost far less than a tenth of the platinum so I didn't care about wasting it. Anyhow, its clean now and there is black scum so Im guessing that the Ti wire is to blame.

@wolen- Firstly, I read your page but was more concerned with procedure than numbers. And I knew that it was bad, I had the number 10000 in my head for columbs per mole (hey, its the right order of magnitude). What I didn't do mentally was multiply it by 6 electrons and realise that dividing by 3600 was still a lot.

No need to worry, I'm pretty good when it comes to recrys. Ive had way too much practice. I'm also obsessive compulsive about purity, so it will have at least 2 recrystallisations, probably one each for each batch isolated and another for the whole lot all added together.

Also- and in one sense this speaks highly of my cell- it is stone cold. I added a bunch of ganged small resistors to push the current up to 3.6A but it is not going to make that much of a difference. I wish it was running hot because then I could isolate a fair bit of bromate by cooling. It is sitting on a hotplate so if necessary I will add a water bath when crystals start to appear.

Is there any disadvantage to running a cell like this hot? Will it significantly increase the wear on the anode? And IIRC, it will increase the current, so I will have to be careful of runaway.... Is this right?

[Edited on 5-11-2007 by Antwain]
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[*] posted on 4-11-2007 at 10:49


Quote:
Originally posted by Fleaker
The solution is very orange red, dilute it is yellow.

Palladium solution in nitric acid or aqua regia actually a very very potent yellow colour that looks brown. I will show you all what I mean by this shortly.


Yes, I have made it. Well its brown by definition because it looks brown. And I assure you it stains stuff brown. But yes, when diluted it becomes yellow and remains coloured to a very high dilution.
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[*] posted on 4-11-2007 at 11:11


Quote:
Originally posted by Antwain

I suppose that it kind of helps to demonstrate that we shouldn't be surprised that molar quantities of reducing and oxidising agents can explode with a large release of energy.


Absolutely right! When I burn my chlorate impregnated filter papers after thorough drying, the energy release is terrific. Twenty to 30 grams of oxidizer being reduced in less than a second. The heat is intense as I have seen NaCl vapor condense on room temperature surfaces!:o:o



The first virgin run of my 2 liter sodium chlorate cell assuming 50% was 14 days! I only have 6 amps to play with at this writting. A 12 amp setup is in testing. I boiled the solution to half volume and reran another 7 days to get 600 or 700 grams of sodium chlorate. With 10 or 12 amps I could complete the first run in a week and secondary run in 4 days. But still this takes a while.

[Edited on 11/4/2007 by chloric1]




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[*] posted on 4-11-2007 at 11:33


At the moment I am runing the shameful battery charger, rated for 6 amps. But I just grabbed a bare transformer my father had custom wound (why I do not know :o) in the 70s. It is supposedly good for 2* 20A @8v but probably not parallel bridgeable. I will need to get a case for it, a fuse and power cord, but this seems like it is up to it is up to any challenge I could reasonably set for it. And the diode bridges attached to it :o. If god needed diode bridges, these are the ones he would use ;), they are huge. Now I just need to build a variable resistor that can dissipate a gazillion watts.
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[*] posted on 4-11-2007 at 20:34


Bah, you haven't seen anything 'til you get your hands on hockey puck rectifiers. :D Four kiloamperes, anyone?...

Conductivity of the solution will probably about double by the 60C range. My cell runs about 30A when the solution is changed, 60A when it's reached equilibrium around 60C or whatever it sits at. Which is less these days, as the anode is down to about 3/4" dia..

Burning NaClO3 is indeed quite impressive. It contains much more oxygen than the potassium salt, and let's not forget the wildly luminous sodium d-lines giving that intense yellow firey look to it. I've got a cardboard drying tray that's encrusted with NaClO3. One of these days, I'm going to take it into the street and light it. With a piece of fuse, mind you. I have two KClO3 trays, but I know they won't be nearly as bright or fast.

Tim




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[*] posted on 5-11-2007 at 10:29


Occasionally a nice 200A, 12V adjustable Lambda power supply will turn up on Ebay. :D


Good for electroplating, electroforming, and electrowinning also. (answer for Woelen V)

[Edited on 11-5-2007 by Eclectic]
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[*] posted on 5-11-2007 at 10:34


Why all those high currents? If you don't have high currents, then use lower currents in cells, connected in series. You easily can take 24 volts and add 4 cells to it. In many cases that is more efficient.

I, however, just am patient. I only need small amounts, and I then can live perfectly with 2 A.




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[*] posted on 5-11-2007 at 15:27


After crystals came out I heated it a bit and then 7 hrs later more crystals appeared. So I took a first crop and its a bit ugly, definitely a candidate for 2 recrystallisations. Next time I think I will just let it run, because although the anode was covered in crystals the current had not dropped by more than 0.1A if that. I added some more dichmorate because I don't know if there was enough there. Interestingly bromine (the bad smell not visible amounts) was liberated but it shouldn't have been because the solution was alkaline... or is this a case of dichromate (bichromate) acting as an acid? It was about another 1/2 a gram added to only 20mL of solution so that may have changed the pH.

Something dodgy is going on too... for some reason the electrolysis seems to increase the solubility of (presumably) the bromate. When I stopped the current and heated the solution quite hot the net result was MORE precipitation, although it was actually a lot of precipitation followed by some dissolution on heating. Does anyone have an explanation for that? I have stopped mechanical stirring and only dredge it up sometimes so it is probably heterogeneous in the jar.

Also, based on what happened after my latest dichromate addition, I think the brown/black muck is a chromium compound, probably a hydrated oxide. There is too much for it to be dissolution of either electrode. How does dichromate work if it doesn't stay in solution?

Also, will the electrodes suffer damage if the bromide concentration becomes too low?(I believe so but just checking). And if so then how will I know when this happens. Bubbles at the anode will be hard to see. Will the current suddenly change?

Incidentally, the current hasn't changed by more than ~5% across a wide temperature range. In fact less, it is staying between 4.5 and 4.8A. Always.
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[*] posted on 5-11-2007 at 15:30


@tim

That sounds like a lot of fun. For goodness sakes give a good distance. I have not seen many oxidizers that burn with such vigor. The low end roaring sound sounds like detonation is imminent.




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[*] posted on 5-11-2007 at 15:48


Actually, I may have found a good solution to the crystals problem. The electrodes are now together at one side of the top of the jar and a very slow convection is taking place because of the hydrogen evolution. As such the bottom of the jar is cooler (feels about 10*C) so crystals are now growing on the bottom. Before, the bottom half was in heated water so the top was cooler.
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[*] posted on 5-11-2007 at 16:05


I know that this thread is 'bromate synthesis', but considering some of the off topic stuff so far this is more closely related.

How powerful an oxidiser is bromate? I have seen a reference somewhere comparing it to permanganate. The standard reduction potentials never tell the story, because concentrations and especially pH differences change the character of redox reaction quite a lot.

In acid solution, is bromate or chlorate the more powerful oxidiser?
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[*] posted on 5-11-2007 at 18:46


Quote:
Originally posted by Antwain
After crystals came out I heated it a bit and then 7 hrs later more crystals appeared. So I took a first crop and its a bit ugly, definitely a candidate for 2 recrystallisations.


What's ugly about it, particulates or just green? Particulates are hot filtered, green can be neutralized (H2O2, if it doesn't react with bromate) or lost to the recrystallizing wash.

Quote:
Next time I think I will just let it run, because although the anode was covered in crystals the current had not dropped by more than 0.1A if that. I added some more dichmorate because I don't know if there was enough there. Interestingly bromine (the bad smell not visible amounts) was liberated but it shouldn't have been because the solution was alkaline... or is this a case of dichromate (bichromate) acting as an acid?


Regardless of pH, my cell continuously gives off chlorine fumes. Even though the solution may become alkaline as a result, it only takes a few ppm (if even that) to smell. There is plenty of hypohalide present to form a small equilibrium concentration of dissolved Cl2 regardless of pH, and plenty of gas bubbling through (i.e., the hydrogen) to sparge it out.

After all, bleach smells noticably "chlorine-ey", despite its pH of 10 or so.

I wouldn't worry about pH, it will always force itself back up as a result of halogen release. After crystallizing, I add acid to the liquor until the chromate changes color (pH ~ 5), or until too much chlorine gas (pH 7-8) is being produced from the hypochlorite that hasn't decomposed.

Quote:
When I stopped the current and heated the solution quite hot the net result was MORE precipitation, although it was actually a lot of precipitation followed by some dissolution on heating. Does anyone have an explanation for that?


Hypobromite decomposing from heat?

Quote:
Also, based on what happened after my latest dichromate addition, I think the brown/black muck is a chromium compound, probably a hydrated oxide. There is too much for it to be dissolution of either electrode. How does dichromate work if it doesn't stay in solution?


Is the solution, after filtering, green? If not, then it's probably some chromium compound. I don't know what it would be to be black.

When electrolyzing a sodium chlorate solution with platinized titanium anode and copper cathode, I noticed a large amount of brown to black material forming, and often being attracted to the cathode, either as a result of deposition or electrostatic attraction. (Interesting thing that attraction, I've had particulates in certain solutions which were slowly attracted to metals placed in the solution. Electrostatic collection without all the kilovolts!) It turns out the copper cathode is actually attacked by reactive species in solution, regardless of its negative potential. Since the solution is so reactive, copper sponge doesn't form on the cathode, only more black crud, making it appear to agglomerate.

Quote:
Incidentally, the current hasn't changed by more than ~5% across a wide temperature range. In fact less, it is staying between 4.5 and 4.8A. Always.


What voltage?

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[*] posted on 5-11-2007 at 20:28


@Tim- 3.08v across the electrodes plus 4m of 15A rated wire, I can't measure the voltage across the cell directly.

Yes, it may just have been cold enough before heating for hypobromite to exist.

There is very negligible bromine coming off now, but it stank really well when I heated the dichromate solution.

I took a picture of the gunk but my camera wouldn't focus. It is blackish and solid. If I didn't know better I would suspect it was carbon.

[Edited on 6-11-2007 by Antwain]
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[*] posted on 6-11-2007 at 00:03


Antwain, do not add so much dichromate. You really only need to add a pinch to your solution. In my experiment, I only used 50 mg or so for 40 ml of liquid. Dichromate is reduced at the cathode, giving green chromium(III) species, and these give a green color to your product. If you have so much dichromate, then you get a lot of fine particles, which may be very hard to remove (too small for filtering, and they act as nucleation sites for crystal formation).

KBrO3 in aqueous solution is a potent oxidizer, not because of its really high redox potential, but because it acts fast. Just for fun, add a little pinch of solid KBrO3 (may be wet or impure, does not matter) to concentrated HCl. A vigorous reaction starts, in which a mix of Cl2 and Br2 is produced. Unfortunately, bromate is not a clean oxidizer. It can be reduced to bromine, or to bromide, depending on whether it is present in excess amount or not. It only works at low pH. At high pH, bromate is not capable of oxidizing much. I would say, that compared with chlorate in aqueous solution, bromate is more powerful. Not really stronger, but acting faster.

In organic reactions, bromate is not the oxidizer of choice, because of bromine formation. An important unwanted side reaction may be bromination of compounds, and not only oxidation. E.g. toluene, suspended in a H2SO4/KBrO3/water mix will be oxidized to benzoic acid, but one get a funny sweet smell as well, most likely this is some brominated compound of benzene or toluene. Not what you want. With KMnO4 you only get benzoic acid.

Bromate also is a very powerful (and dangerous!) oxidizer in dry mixes, compositions based on bromate are more sensitive than comparable compositions, based on chlorate. So, for real production work it is not suitable, simply too dangerous and too high a risk of unwanted ignition. On Usenet:rec.pyrotechnics, it was compared to chlorate as follows: "Playing with chlorates is like playing with a deadly poisonous snake. Playing with bromates is like playing with a deadly poisonous snake, which is slightly pissed off.".

I use the bromate mainly for making bromine. It is a fantastic source for making bromine. Simply mixing it with a 5-fold molar amount of KBr or NaBr and adding this to 20% sulphuric acid makes the bromine drop out and simply pipetting is suitable for obtaining already 75% yield. For the remaining 25% you'll need a distillation setup. I do not store bromine, I make it when I want to experiment with it.

If you want to protect your anode, then simply stop electrolysis, when 75% to 80% of the calculated time for full conversion has passed. In this way, you certainly will not perform electrolysis in a cell, which hardly has any bromide. The liquor then can be kept for future cells. No need to throw it away, it will be a nice starting point for a new cell. I indeed would be careful with electrolysis when hardly any bromide is left. There will be much more oxidative strain on the anode in that case, because oxygen must be formed, instead of bromine.

Finally, I would not add any acids or whatever stuff to your cell for purification. Just collect the solid matter and recrystallize. Acid also adds contamination.




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[*] posted on 6-11-2007 at 03:40


Yes, the dichromate was a bad idea. Well, you learn these things. My solution is a mess :(, well it can be cleaned.

So as I mentioned in another topic solution got up my anode and started to corrode it. I sealed it with wax but this is temporary at best. I have put the connecting wire through hell it is way TOO flexible for my liking. The wax has come off a bit and has clearly reacted, a bit, but I can deal with that; there isn't much and it can be gotten out.

Now I find out that my cathode is leaking like a sieve too. At least the solution in that isn't causing damage in there, just putting a few bubbles up the glass tubing. I cant see why.... well I can, it was because in each case I sealed multiple strands of wire into the glass. But for the cathode I melted it so long and hard that the copper oxidised a lot and the oxide dissolved in the glass. You can't get more forceful than that and it still wasn't enough to actually seal it.

I think that after this I am about to give up on sealing stuff in glass, it just isn't happening for me. At least the platinum seems to be doing ok.
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[*] posted on 6-11-2007 at 07:44


Oh well, may as well give the blow by blow. Even if noone wants to see it now it may help someone in the future if I am thorough.

After about 12-18hrs, maybe 6hrs room temp and 12hrs at ~30-40*C, the solution started depositing crystals again. They are a pain to get out so I decided to freeze it and then return it for more electrolysis. The first 6 hours when nothing deposited was because I did a dirty thing and added the bromide/bromate I made as described in the bromine thread a while back, maybe 30-45g in the 5:1 ratio, then got out much of the bromate because common ion stopped it from dissolving (if it has KBr in it I will get that on the recrys). The solution volume also increased a bit as stuff was washed into the jar.

Anyhow, when crystals started appearing at room temp I gave it the heat with waterbath treatment. And just now it started to crystallise warm so I poured it out and iced it. I don't know the density of my product as there is still liquid in it, but I would guess 15-20g came out. Pretty poor by current, but at least its working. The following picture shows the 2 crops I have obtained thus far. You can clearly see the new and old stuff because the new stuff is still very yellow, but last time the bromate settled out and left yellow solution when it warmed up so hopefully it will again. Also there is a large amount of brown/black gunk now thanks to the dichromate, but it does settle, so worst case I will be doing a decant and have a lot of water to boil down.

electrolysis - bromate-a.JPG - 22kB
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[*] posted on 6-11-2007 at 18:43


Quote:
Originally posted by woelen
I would say, that compared with chlorate in aqueous solution, bromate is more powerful. Not really stronger, but acting faster.


@woelin-the percieved activity of the bromate over the chlorate is probably due the decreased stability of bromine oxides. Most chlorate oxidations seem to go like:chloric acid,chlorine dioxide,chlorine/chloride.




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[*] posted on 10-11-2007 at 03:11


Ok several updates. Much of this will seem obvious to the experienced electrochemist, but I am pulling myself up by the bootlaces here and trying to detail everything.

Don't use a steel lid on your jar, it will corrode and run into the solution. It can be separated with only a loss of time, but I am sure it was hurting current efficiency. I am using a plastic lidded jar now and its going good.

Don't place the cathode under the anode. I don't know if Pt adsorbed hydrogen was screwing things up, but its going better now. There will still be lots of bubbles to react with Pt catalyst but its not being blasted with them.. I only did this, btw, to keep the electrodes close in my original bad-geometry setup.

Next to each other is good and off the bottom works well too. The bromine sinks and there is a distinct layer stopping a mm below the bottom of the cathode. This is also good because it keeps the electrodes slightly warmed by their own power and precipitation happens at the bottom.

I estimate previous current effiency to have been <10% and maybe <5%. Just from the precipitate now seen after several hours, it is clearly much higher. 5-10 fold higher.
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[*] posted on 10-11-2007 at 13:10


2 questions. Firstly, there are some bubbles forming on the anode. Much less than the cathode, by like 20 fold or more. There was still orange dense fluid pouring off it at a good rate. Does this mean that my current density is too high or that I am out of bromide, or could it be either?

This question is kind of in unison with the last. I have a very bad feeling about my cell. Does KBrO3 dissolve with a very noticeable cooling of the solution, like KBr? If not, then I am very concerned that I have been isolating KBr from my cell :(. Sometimes it seems like there are two different solids there sometimes just one. It should not be possible to be isolating KBr, based on the current volume of solution and the quantity of KBr used initially, regardless of whether the stuff I have pulled out so far is KBr or KBrO3. But for something which isn't soluble in cold water it is acting like something which is very soluble, not so much in terms of how much dissolves (the crystals are too small to know how much exactly I have) but 5g/100mL into 20mL shouldn't be able to produce this much cooling. Or should it?

Actually, one piece of additional information. Despite having gotten out a reasonable amount of solid. Should a MAXIMUM of 100g of KBr, dissolved in 180mL of solution be able to crash out enough solid to half fill the beaker with slurry when cooled in an ice bath?

Ok, so I panicked. I just killed 2 birds with one stone, doing a solubility test while recrystallising all the product I have so far. It must not have been packed as tight as I thought, because it was ~70-80mL of what seemed like near dry crystals, but at density>3 that is impossible. Say it was at least 100g. It would not dissolve in 200mL of water despite best efforts... As it dissolved from the bottom it was snowing down from the top which was being evaporatively cooled by stirring. It dissolved with difficulty in 300mL of boiling water. Definitely not KBr then.

[Edited on 11-11-2007 by Antwain]

[Edited on 11-11-2007 by Antwain]

[Edited on 11-11-2007 by Antwain]
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[*] posted on 10-11-2007 at 14:47


Antwain, actually, testing KBrO3 purity is very easy. Just take some of the crystals and add them to just 1 ml of 10% H2SO4 (not HCl, use H2SO4). If the solution remains colorless, then you have really pure KBrO3, if the solution turns ligt yellow, then you have some KBr in your KBrO3. If the solution turns orange and you see even some vapor of Br2, then you have very impure KBrO3. Also, try your solid, mixing it with some powdered S and/or C, and ignite. If you have KBrO3, then you'll definitely notice, even if it is impure ;).

If you see a steady stream of bubbles of oxygen at your anode, then it is best to stop. The concentration of bromide then has dropped considerably. Let the contents of the cell cool down in the fridge in order to get most of your KBrO3 out of it, and keep the liquid for further production of KBrO3 (just add new KBr and maybe some water).




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[*] posted on 10-11-2007 at 17:01


Quote:
Originally posted by woelen

I use the bromate mainly for making bromine. It is a fantastic source for making bromine. Simply mixing it with a 5-fold molar amount of KBr or NaBr and adding this to 20% sulphuric acid makes the bromine drop out and simply pipetting is suitable for obtaining already 75% yield. For the remaining 25% you'll need a distillation setup. I do not store bromine, I make it when I want to experiment with it.




Woelin- I did the math and it looks great. I assume your 20% sulfuric acid is cool as not to vaporize your bromine. Question is when you make up your bromate/bromide mix in water how much water do you need? Or do you use the crystaline salts into the acid? Is this reaction exotheric to any great extent?
I once made bromine by adding 30% H2O2 to concentrated HBr. The was a slight delay and then a fast reaction occured vaporizing most of the bromine sending me running for higher ground.




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[*] posted on 11-11-2007 at 04:00


I use the solid salts for making the bromine. I mix solid KBrO3 and KBr (or NaBr) in a molar ratio of 1 : 5. I use a slight excess of bromate, in order to avoid the problem of bromine, dissolving very well in solutions, which contain bromide.

As a start, try the following for making 10 grams of bromine, of which appr. 7.5 grams can simply be pipetted from below the aqueous layer.

Take 3.5 grams of KBrO3 and 12 grams of KBr (or 10 grams of NaBr), and finely grind the chemicals and mix them thoroughly. No need to fear any reaction, as long as no acid is added.

Take 40 ml of cool 20% H2SO4. Add the solid mix, 2 gram at a time and stir after each addition. Soon, you'll see bromine dropping out. When you have added all solid, then cap the container and carefully shake, in order to have all salt mix dissolved. Lots of bromine drop out. After 30 minutes or so, you'll have a clear liquid (bright red/orange) with some white crap (mostly KHSO4/K2SO4) and a big blob of bromine at the bottom. With a long pasteur pippette you simply suck up this bromine. Yield: 75% to 80%, appr. 2.5 ml. The remaining bromine is dissolved in the red/orange liquid. You also should keep this liquid, bromine water also is quite interesting for many experiments. If you want to go to the max, then you can distill the bromine water to get the remaining 20% of bromine. I, however, do not take that effort, I simply keep the bromine water.


You can scale up this procedure, the reaction is only slightly exothermic. Use 2 gram lots for each addition. So, on scaling up, you'll need more time for making the bromine, but it works equally well.

Be careful with this, LOTS of bromine vapor are created, and when swirling the solution, keep it capped, or covered with a piece of glass!




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