something that`s always confused me is the Halate cell.
take NaCl for instance, you want to make NaClO3.
Chlorine gas is given off during this reaction, and yet other than the addition of 3 oxygens to the original NaCl there`s no difference.
and it seems a waste of good Chlorine that could be put back into the reaction by the addition of NaOH.
where does this extra chlorine come from anyway, the product of NaClO3 after destruction is NaCl again???
I`m sure I`m missing something quite simple here, I just can`t see it woelen - 22-11-2006 at 07:01
In an ideal cell there is no extra chlorine, nor any loss of chlorine. A perfectly constructed chlorate cell does not expell any chlorine. If you have
an ideal cell, with perfect mixing of the chemicals in the cell, then the net reaction is:
NaCl + 3H2O ---> NaClO3 + 3H2, written in ionic form this is
Cl(-) + 3H2O ---> ClO3(-) + 3H2
So, no loss of chlorine, nor the required addition of chlorine from an external source. All chlorine is coming from the table salt.
The reaction mechanism is as follows:
Anode absorbs electrons: 2Cl(-) - 2e --> Cl2
Cathode gives electrons: 2H2O + 2e --> H2 + 2OH(-)
Now, if there is perfect mixing, then you will have the following reaction:
Cl2 + 2OH(-) ---> Cl(-) + ClO(-) + H2O
On heating, the hypochlorite disproportionates to chloride and chlorate:
In a real cell, there is loss of chlorine, because it bubbles out of solution. With vigorous stirring, this effect can be made less severe. Also,
having the cathode below the anode helps, because it makes the OH(-) from the cathode move along the anode, making the absorption of Cl2 and formation
of ClO(-) and Cl(-) faster, hence less loss of chlorine.
Another cause of loss is the production of oxygen at the anode. This is an unwanted side reaction:
2H2O - 4e --> 4H(+) + O2
If extra NaOH is added, or excess chlorine is lost by bubbling away and excess NaOH builds up in solution, then this unwanted side reaction is
strongly promoted, because OH(-) ions are much more easily oxidized than water:
4OH(-) - 4e ---> 2H2O + O2
So, once you have too much NaOH in solution (either added on purpose, or because of excessive chlorine loss), then the cell becomes very inefficient,
due to oxygen production. But the good news is that this reduces OH(-) content again, and makes the cell more efficient again. So, fortunately there
is a certain feedback, which works in the right direction.
Yet another cause of loss is the back-reduction of hypochlorite and chlorate at the cathode, with the following net reactions:
ClO(-) + 2e + H2O ---> Cl(-) + 2OH(-)
ClO3(-) + 6e + 3H2O ---> Cl(-) + 6OH(-)
Addition of a pinch of a dichromate strongly suppresses this back-reduction. I also have the impression that the use of a large cathode with lower
current density per square cm also reduces the effect of back-reduction.
Making bromate is easier than making chlorate, I have found out practically. Bromate cells do not suffer from losses due to bubbling, because Br2 is a
liquid under the conditions of the cell, it simply slowly moves away from the anode. Bromate cells in general also suffer much less from anode erosion
and corrosion of the leads towards the cell. A pinch of dichromate again helps suppressing the effect of back-reduction at the cathode, but it can
lead to very hard to remove chromium (III) contamination of your product, it becomes pale green and not snow-white. With platinum anodes, the problem
of formation of green chromium (III) does not occur, I also found out that practically.
[Edited on 22-11-06 by woelen]YT2095 - 22-11-2006 at 07:08
actualy it`s a Bromate cell I`m running at the minute, voltage 5.48 current 1.14A at 27 celcius.
the liquid is quite Tea like brown now, I`m stirring roughly every 3 mins with the thermometer, there`s little color change though?
in the case of Real cells with the chlorine loss, is making the soln a little more Basic going to help lock up this Cl2 as it forms a little better?woelen - 22-11-2006 at 07:26
Quote:
in the case of Real cells with the chlorine loss, is making the soln a little more Basic going to help lock up this Cl2 as it forms a little better?
No, don't add NaOH. It may help lock up the Cl2, but its negative effect of promoting the formation of oxygen outweighs the positive effect of
enhanced absorption of chlorine.
The tea-like brown color of the bromate cell I recognize very well. It is a mix of orange/yellow from dichromate and green from chromium (III)
compounds. If you continue doing electrolysis (let it run for 12 hours for each 20 grams of KBr in your solution at a current of appr. 1 A), then
finally the liquid will become dark green, due to a mix of finely dispersed carbon and chromium (III). The carbon will settle at the bottom within a
few hours and the chromium (III) remains in solution. If you use KBr at sufficiently high concentration, then you will see the solid KBrO3 settle at
the bottom, looking very dirty drak green. It contains carbon, trapped in the crystals and chromium (III).
I took the green liquid, with the dirty crystals and put it in the fridge to separate out more of the KBrO3, and then I vigorously stirred it. The
crystalline mass quickly settles again, the carbon remains in the liquid much longer. I quickly decanted the liquid. Next, I rinsed the dirty crystals
three times with ice cold distilled water. After that, I dissolved the crystals in as little as possible of almost boiling hot distilled water. I kept
the liquid hot and allowed solid stuff (carbon) settle and decanted the clear liquid in a clean pre-heated beaker. Then I let it cool down slowly.
This yields light green crystals. I finally put this in the fridge again to obtain the last amount of KBrO3. I repeated this recrystallization another
time, and the final product is still light green, but it is totally free of carbon-crap and it also is free of bromide. If I add the crystals to
dilute sulphuric acid, then I do not get a yellow color. If they still contain bromide, then on addition of acid you get Br2: BrO3(-) + 5Br(-) + 6H(+)
--> 3Br2 + 3H2O.
EDIT: I did not stir my bromate cell every three minutes, I placed the cathode somewhat under the anode, such that bubbles of hydrogen (taking OH(-)
with them) moves along the cathode. The convection through the liquid does the rest and gives sufficient stirring. Every hour or so, I checked the
liquid and did a little shaking. I used a 2 A current, halving the time needed for the electrolysis. The liquid became quite hot, I estimate somewhere
between 50 and 60 C.
[Edited on 22-11-06 by woelen]YT2095 - 22-11-2006 at 07:35
I`m not using dichromate in the Bromate cell, that`s only something I use in Chlorate cells as I PPT out with a potassium salt later so the Cr3
contamination isn`t an issue.
with the Bromate cell I`m already using the Potassium salt, so I want to keep it as Chemicaly pure as possible, carbon`s not a real issue here nor is
my anode lifespan. I did the same thing before on Minuature scale to make a half gram of of the bromate for the RP tests woelen - 22-11-2006 at 08:03
Let me know how the cell fares in the long run. Without the dichromate I can imagine that the production of hydrogen gas at the cathode comes to a
halt after some time (this is something which garage chemist observed in his cell). If this is the case, then the cell still takes a lot of current,
but what you gain at the anode is destroyed again at the cathode and your cell then effictively is nothing more than an expensive and risky way of
heating your lab.
If you have had it running for several hours, then carefully watch how much hydrogen is produced at the cathode. Btw, what material is your cathode
made of and how thick is the cathode. I have the impression that a thick/large cathode with large surface area makes the back-reduction effect less
severe, but I'm not sure about that.YT2095 - 22-11-2006 at 08:12
it`s been running now for several hours, the Cathode is quite "happy" and fizzing merrily away.
I`m using 12mm Graphite rods, both anode And cathode.
the deep tea color builds up along the bottom and then a quick but vigorous stir later it`s all back to the normal brown color and no change overall
between stirings.
the current and voltage haven`t changed at all, beyond a slight fluctuation during the stiring.
the temp is not 29c though, so Im still well within the cell efficiency margin. if it does go much above 30c then it`s either active cooling or I`ll
shut it down for a while.
it will be filtered and heated after this though, and then returned back to cell for further Punishment YT2095 - 22-11-2006 at 11:49
interesting you should mention the cathode stopping!
I powered down the cell for 40 mins as its temp was 32c and I also had to leave the lab for a short while.
when I got back and tried to fire it up again, almost Nothing happened, same power drain, but little in the way of any visble reaction?
I had to "Jump-Start" it by direct power from a car batt for 30 secs and try again, even then it was at a craw, several mins later though all seems to
back up and running as normal.
I did notice that when I gave it a stir the cathode side seemed to die down a bit, so maybe he switched his cell off for a while also?
I know Now that will temporarily kill a Bromate cell, so I`m just leaving it switched on now, screw the temp
[Edited on 22-11-2006 by YT2095]woelen - 22-11-2006 at 23:36
What you observed is the nasty back-reduction effect. The cell did work, but instead of reduction of water to hydrogen and hydroxide ions, you had
reduction of bromate ions to bromide ions again and hence you see no bubbling of hydrogen. This is what I mentioned in my previous posts. Even if the
cell does seem to function properly (production of H2), there still may be considerable back-reduction. You cannot tell the difference between 100%
formation of H2 and 50% back-reduction and 50% formation of H2. This effect is very bad for the cell efficiency, and hence the addition of a pinch of
dichromate. The dichromate prevents back-reduction. I do not know why, however.YT2095 - 23-11-2006 at 09:13
well I let run all night, filtered and heated then crystalised slowly all day today, I have Beautifull clusters formed, it almost looks like bacterial
cultures in a petri dish, but in 3D like shrubs or bushes
I`ll return the liquer back to the cell after this crop and let it run some more, at the moment it looks like a Good harvest, with as much by volume
produced as the Bromide I put in. they`re also pure white/clear too.
the only thing I can gleen from experimentation thus far is to Never switch a Bromate cell off half way through when using graphite electrodes.woelen - 24-11-2006 at 00:37
Interesting. I'll try again with graphite rods next weekend and see what yield I obtain. It would be good if the dichromate is not needed. That makes
the synth even more accessible for a larger group of people.
On my webpage about KBrO3 synth I do advice the addition of some K2Cr2O7, but this is based on the experiences of garage chemist. Hence, I did not
even try without in my newer experiments.
Is your bromate really purely white, or does it still show some yellowish or gray color? I also thought that my first batch is white, but compared to
my second batch, it certainly is not white. Look at the difference in this picture:
The left seems white, but in contract with the right sample it definitely is not white.YT2095 - 24-11-2006 at 01:34
its very difficult to say at the moment, as I`m letting the soln cool VERY slowly to get larger crystals and so it`s just been put into the fridge now
(about 20 mins ago) to stay there for another 8 hours, but at face value they seem to be white enough when looked at from underneath the flask.
something Interesting Has occured though, when I shut the cell down, I washed the electrodes in plain water and left them out to dry, the cathode is
Covered in a crop of white crystals now!?
and nothing at all on the anode, normaly I boil them afterwards and then dry them in a bunsen flame until no further halogen smell exists, I wasn`t to
fussed this time as the remains of the Bromate soln will be retured back to the cell for round 2.
it`s Odd that these crystals formed so quickly and so many of them though, it might even be impossible to remove them entirely, and be best to use new
electrodes for each different cell, I think I`ll label these as Bromate Only now.YT2095 - 24-11-2006 at 10:01
my crystals didn`t turn out white at all, they have a yellow tinge, the same color as the soln did, and even after 2 washes it remains locked up in
the clusters, so I may as well have use Dichromate afterall.
not to worry though, I returned that soln back to the cell and the washing water and added another 2 5ml scoops to it and some dichromate this time.
interestingly the same "Stall" occurs if you turn the cell off for a minute, even with the chromate added, so it`s not That which seems to be the
cause IMO, it has to be something else. I tried something different to "jump-start" it this time, I reversed polarity on the electrodes, a whitish
cloudy mist comes off the cathode into soln. then when wired back to correct polarity, it still doesn`t fire up, I had to shove 12v car batt power
through it again for 30 secs, then resume normal power.
Also, whilst running this cell the temp started to creep up and got past 30c so I had an idea of Cooling it in an ice bath and still running it, the
Voltage remains constant but the current drawn drops by roughly 1 quarter of original, to draw an Amp I have to run at 7.1 volts.
then something unexpected occured the cell had been running at 1A at 9celcius for about 30 mins, when I went to stir it I noticed white clusters on
the cathode, like little round shrub/bushes roughly between 2 and 4mm around, and yet on the Anode, there where clear needle like crystals at 90
degrees perpendicular to the elctrode All pointing towards the cathode and Only on the side facing the cathode, and not in little clumps either but
totaly Covered.
it`s wholely remarkable to see! it`s things like this that make make Science/Chem worth it, you just couldn`t Pay for a day like today woelen - 24-11-2006 at 10:45
Why isn't the cell allowed to heat up beyond 30 C? I let my cell run, regardless of temperature. I ran 2 A through just 40 ml of liquid and this
liquid becomes so warm that I can hardly bear touching the glass container. I think it is going well beyond 50 C and maybe even 60 C, but who cares? I
think it even is advantageous. The disproportination reaction from hypobromite to bromate+bromide is faster at these high temperatures, and the
solubility of potassium bromate is much higher at those temperatures. During electrolysis you do not want any crystals, certainly not on the
electrodes. After electrolysis you let it cool down in order to get all crystals.YT2095 - 24-11-2006 at 11:07
well according to data that I`ve read, a perchlorate cells effeciency drops off quite drasticly at 40c so I`m applying this to all cells for now as
it`s Expermental, also I try to avoid this range where possible as I don`t have the current to waste, this PSU`s not able to sustain above 1.5 A.
so in my case it`s largely due to safety factors over effciency. nothing else
[Edited on 24-11-2006 by YT2095]woelen - 24-11-2006 at 14:02
Ah, yes, for a perchlorate cell it does matter. At too high a temperature, you just produce oxygen at the anode, instead of making perchlorate. But
for halate cells it is no issue. You produce the free halogen anyway, regardless of temperature. For a chlorate cell, I can imagine that a low
temperature is somewhat advantageous, due to the better solubility of Cl2 in cold water (less losses due to bubbling of Cl2 out of solution), but for
bromate that is not an issue. With good mixing, it also hardly is an issue for chlorate production.woelen - 25-11-2006 at 07:11
I have had a cell running without dichromate added. I dissolved 20 grams of KBr in appr. 40 ml of water. After 4 hours I still did not have any
crystals of KBrO3, while in my previous 2 runs I already had crystals of KBrO3 after 2 hours, while using the same concentration.
The production of hydrogen is slow. After these 4 hours I decided to add some K2Cr2O7, while keeping the cell switched on. I dissolved 50 mg of
K2Cr2O7 in a ml of water and poured this into the liquid. At once, the production of hydrogen becomes much faster. I now have it running and will keep
it running for 8 hours or so.YT2095 - 25-11-2006 at 09:42
20g in 40ml!
wow, I was only doing about 8g in 150ml woelen - 25-11-2006 at 12:13
Now I have it running for 5 hours or so, after adding the dichromate. Now I have my liquid totally clogged with the solid KBrO3. A thick crust is
formed between anode and cathode and I had to remove the electrodes a few time, because they became covered by the solid more and more. Putting the
electrodes back into the liquid hardly can be done, there is a really thick layer of KBrO3 at the bottom now, it is more voluminous than the initial
volume of KBr I added. But this of course is understandable, because more matter (oxygen is added), while the density only marginally increases. The
color of the raw KBrO3 is dirty green.woelen - 25-11-2006 at 14:21
I just stopped my cell, and rinsed the KBrO3 somewhat with icecold distilled water. The yield is very good, the material is almost free of bromide,
even without recrystallizing from water (it only gives a very faint yellow color to dilute H2SO4), but its color is definitely green. Even darker than
my first batch.
Now I have a new idea on why there is such a bad color. I think the color is due to carbon, which absorbs dichromate as well. The yellow/black mixture
makes the green color. I noticed that there is coarse carbon, which easily could be removed, but I think there also is VERY finely divided carbon,
almost colloidal, which is really hard to remove. It cannot be filtered and simply remains in solution, it does not settle. Tomorrow I will change my
website on the synthesis of KBrO3, I really think this is the cause of the green color.
I have my last batch drying. I will not recrystallize it, I'll convert it all to Ba(BrO)3, which is only very sparingly soluble (IIRC 0.4 grams per
100 ml of water in the cold) and very easy to isolate. I did some testing with a Ba(BrO3)2/S mix (ratio appr. 3 : 1 by weight). I can only say wow at
the color of the light from this mix, when it burns. A really deep green, saturated like a good green LED with a slight blue tinge. I seldomly saw
such a brilliant and saturated color from a pyro-mix.
[Edited on 25-11-06 by woelen]vulture - 25-11-2006 at 15:20
Quote:
It cannot be filtered and simply remains in solution, it does not settle.
Dissolve the KBrO3 in water and filter over atleast 5cm of fine silica. You could also add some ammonia to precipitate any Cr(III) which is then going
to stick to the silica as well. (I'm assuming you don't have basic alumina)YT2095 - 26-11-2006 at 07:53
Woelen what Ba salt are you using that displaces the K cation in in the bromate cell?
I can`t think of many that would drive that reaction unless you`re using NaBrO3, I make my BaClO3 form the chloride of barium.
are you making Bromic acid and then doing the reaction?woelen - 26-11-2006 at 12:36
I do not use any Ba-salt in the bromate cell. I first make KBrO3 and lateron I make Ba(BrO3)2.
Making Ba(BrO3)2 is one of the simplest syntheses. Take any soluble Ba-salt (I use Ba(NO3)2, but BaCl2 also will do) and dissolve in as little as
possible of hot, near 100 C water. Take some KBrO3 and also dissolve in as little as possible of hot, near 100 C water. Mix the two hot solutions and
let cool down. When cooled down to room temperature, put in the fridge and let cool down to close to 0 C. You'll get a white crystalline precipitate
of almost pure Ba(BrO3)2. Decant the water. Rinse two times with ice cold distilled water. Now you have almost pure Ba(BrO3)2. For pyro-purposes you
do not even need to rinse, just press out most water and let dry.
Yield is near 100%. The very good thing of Ba(BrO3)2 is that its solubility is quite good at 100 C. IIRC more than 10 grams per 100 ml of water, while
at 10 C its solubility only is 0.4 grams or something like that, and at 0 C its solubility is less than 0.3 grams.
[Edited on 26-11-06 by woelen]YT2095 - 27-11-2006 at 05:44
I concur my kBrO3 is also filthy, coloidal carbon seems quite likely the cause, 2 washes later and they`re still grey.
back to the BaBrO3, what actualy drives the reaction here?
I would have thought the addition of KCl or NO3 would have moved the Ba from the bromate, not the other way around.
it certainly doesn`t seem to with the Chlorate.woelen - 28-11-2006 at 13:26
Yes, it is difficult to get clean KBrO3 when carbon rods are used. But the impurity is very specific. Except for the carbon (possibly with adhering
chromate/chromium), the KBrO3 is very pure. Funny to read that you get grey KBrO3. That supports my theory of the carbon impurity. You did not use
dichromate, so you only have C-impurity, making it grey. With dichromate added, the additional chromate/dichromate, absorbed by the carbon, makes it
appear green.
The formation of Ba(BrO3)2 is driven by its low solubility. Of any combination KCl/KBrO3/BaCl2/Ba(BrO3)2, the Ba(BrO3)2 by far is the least soluble,
and that is what separates from solution. The same is true for the nitrate variation.YT2095 - 29-11-2006 at 05:08
well that`s Handy to know. I wanted to make some Calcium Bromate and Chlorate, it`s suposed to give a nice Pink flame, somewhat different to K and Sr
mixes.
thus far the only method I have is electrolysis with periodic current reversal every 30 secs to stop the Cathode from getting caked in crap, it would
nice to have a more elegant method woelen - 29-11-2006 at 14:38
Unfortunately, I only have found a satisfactory way to make KBrO3, AgBrO3, and Ba(BrO3)2. Other bromates are very soluble, and making calciumbromate
is really hard. It simply is not formed from solution, because any other calcium salt will settle first from solution.
One could make Ba(BrO3)2, dissolve this in boiling hot water, add dilute sulphuric acid to precipitate BaSO4 and have HBrO3 in solution, and then add
CaCO3 or Ca(OH)2 to this, one could also add other metal carbonates or hydroxide/oxides. But going this way is very cumbersome. Preparing a solution
of HBrO3 without traces of Ba(BrO3)2 or traces of excess H2SO4 will be really hard, unless you have very precise equipment and want to take the
painstakingly precise effort to prepare a solution of H2SO4, with very precisely determined concentration.
Making Ca(ClO3)2 or Ca(BrO3)2 by electrolysis of a solution of CaCl2 or CaBr2 seems a very cumbersome thing to me also. You get solid Ca(OH)2 all the
time around the cathode, which really clogs up the thing. This is not the way to go for making the chlorate or bromate to my opinion.
Btw, I changed my page on making KBrO3, it now has the explanation of the fine carbon particles in it. If I have some silica, then I will try the
method of Runaway on a test tube scale. But first I need to obtain suitable silica.woelen - 30-11-2006 at 14:49
