I have been away from chemistry of all kinds for a long time, hence my absence from this forum. I hope you will welcome me back!
Well, recently I decided to start doing chemistry again and since I have tried producing perchlorates before, unsuccessfully, I chose to try this out
again.
My setup is a 1.5L glass jar with a plastic lid, in which a titanium cathode and platinum anode is mounted, only spaced 5 mm. apart. The lid of course
also has a vent hole for the gasses produced to escape. My power supply is a PSU with all the black wires (negative) clamped together and all the red
wires (positive) clamped together for optimum current supply. The label on the PSU says that the red wires together with the black yields 5V and
30Amps. I assume this is max. output. Nevertheless, the platinum anode is connected to all of the red wires and the titanium cathode to all of the
black wires.
My electrolyte consists of 450g of pure sodium chloride in 1250ml deionised water. No dichromate/persulfate additives was added.
The cell was located outside on a table, the outside temperature being between -5-0 degrees Celsius. No insulation was added around the cell, which
gave me an average cell liquor temperature of about 0-5 degrees Celcius (I know, it is a far too low but from what I've read, it only results in loss
of voltage(?))
My calculations for the cell run time was as follows:
NaCl --> NaClO3: 190 hours (assuming 40% efficiency), which eventually would give me 820g NaClO3 (without taking loss into account)
NaClO3 --> NaClO4: 35 hours (assuming 40% efficiency), which eventually would give me 943g NaClO4 (again, without taking loss into account).
That should add up to give a total run time of about 225 hours (40% efficiency), however I decided to run the cell for 334 hours, since I did not have
time for the purification steps at the moment the run was supposed to be finished, and figured that the extra run time would not hurt.
Today, I decided to process my cell liquor. The solution was completely clear and had a pale yellow tint. I filtered it just to be sure that no
particles was present, and the solution was transferred to a 1.5L Erlenmeyer flask and a sample was taken for test for chlorates (Indigo Carmine +
HCl; resulted in a slow colour change from dark blue to an almost not present faint green; imprecise result due to hypochlorite) and the solution was
boiled quite vigorously for an hour. At the end, the solution was completely transparent and without any tint, and the volume of water lost was only
about 50 ml. due to condensation in the neck of the flask. However, a large amount of precipitate had appeared (about 1 cm. on the bottom of the
flask). I presume this is sodium chloride from the decomposition of sodium hypochlorite into sodium chloride and sodium chlorate (3NaClO --> NaClO3
+ 2NaCl). A sample was taken from the solution, as before, and added to the exact same testing mixture (Indigo Carmine + HCl). However, this time the
test reacted violently changing to yellow and producing a large amount of chlorine dioxide (from the reaction between NaClO3 + HCl).
This all tells me that the cell liquor consisted mainly of sodium hypochlorite before boiling (please tell me if I am wrong), and the boiling solution
(including the undissolved NaCl) from the Erlenmeyer flask was transferred back into the cell (this time wrapped in cloth to serve as insulator).
Now the cell is running (and producing a lot more gas than before) and I am waiting to get my hands on some methylene blue to determine whether
perchlorate production has started yet or not. Do you have any ideas at all about what I should do here? I cannot find out where I screwed up (maybe
with the calculations, cell temperature,...?). And what can be the cause of the sudden precipitation of that much NaCl when the volume lost is only 50
ml? It must be NaCl since all other sodium salts possibly produced are more soluble in water than NaCl. I will keep this thread updated with my
findings.
Many thanks in advance!!
- laekkerBoy
[Edited on 23-1-2013 by laekkerBoy]laekkerBoy - 23-1-2013 at 14:49
When thinking of it, the evaporation that happens during the run of the cell could possibly have maintained a saturated solution of the lower-soluble
compounds. Then, when the hypochlorite is broken down, the less soluble sodium chloride would precipitate out. Obvious. I do not know why I did not
think of that before.
However, I am still concerned about my very ineffective cell. I have run it several times before without any problems, though under different
temperatures. Could the ineffectiveness be due to the low temperature now allowing the decomposition of the hypochlorite, hence preventing the cell
from going further on into the process of creating more chlorate and eventually perchlorate?
If you spend time in study of the materials in the site listed above likely all your questions will be answered. Or one might say this site is
required reading for anyone building these cells.
By the way there is a long thread on this subject on this site if you will search for it. Possibly you may find your answer by reading it.
[Edited on 1-24-2013 by IrC]Ral123 - 14-2-2013 at 00:33
I've been trying to make NaClO3 from quite a while. I use graphite electrodes, 5v power supply and a small flask operating quite hot. I've never been
able to produce any product with oxidising properties. I've produced few grams of nice NaClO3 from NaOH, HCl and KMnO4. In my last attempts I ran the
cell till 1/3 of the graphites are in the solution, a lot of distilled water has been added. Then I filter the graphite and take part of the mixture.
I bubble chlorine trough it, as much as it would take it. In the end, it still didn't oxidise anything. My bad luck doesn't stop there. When I made a
H2/O2 gas generator, I used NaCl, left it working for a while and then assumed the gas is almost pure H2/O2. I was wrong, my chemically made H2/O2
gave much more powerful bang. I bought KCl to make a control sample with KClO3 and it worked like a piece of cake. It quickly gave me dozens of grams
of product, much more violent as oxidizer then my commercial KClO4. I now have very tough graphites from the contacts of these electric buses. How is
the standard way of making NaClO3 from salt with graphites? Everybody gives guide lines but looks like nobody has did it.phlogiston - 14-2-2013 at 03:02
laekkerBoy, at low temperatures, the hypochlorite can be oxidised to chlorate and eventually perchlorate by anodic oxidation (rather than
autooxidation), but this is inefficient (in terms of electricity). The higher temperature and pH optimisation are aimed at maximising the rate of
hypochlorate auto-oxidation. Perhaps this was a factor in that run?
Ral123, After electrolysis, how do you isolate the NaClO3? You need an extremely concentrated solution to get it to crystallise. Try adding some KCl
solution to your NaCl electrolyte, to see if any KClO3 precipitates. If so, the chlorate is there, you just have problems getting crystalising it from
the solution.Ral123 - 15-2-2013 at 13:33
I let a lot of crystals NaCl form, and then take the solution. If I've added HCl to the electrolysis, I assume I must have solution of NaClO3. If I
didn't add anything, but distilled water while the electrolysis, I enrich in chlorine and then assume I must have NaClO3 solution. Is there something
special with the graphites or the voltage or temperature with NaClO3? jock88 - 15-2-2013 at 15:00
