xfusion44
Hazard to Others
Posts: 223
Registered: 6-8-2014
Location: Europe
Member Is Offline
Mood: Nostalgic
|
|
Problems with sodium hypochlorite production
Hi!
I'm trying to make some sodium hypochlorite solution, with electrolysis cell, but I always fail to make it. First I've tried with undivided cell
(35g/100ml NaCl was used) and two graphite electrodes. Voltage was 5V and current was 5A.
After 4.5h there was no yellow/green colouring to solution, but it smelled just like ordinary bleach, pH was also 14.
Then I've tried the same with 10g/100ml NaCl (same electrodes, same power), for 3.5h. Still no difference.
Today I decided to try bubbling Cl2 gas through solution of 20% bw NaOH, to see if that'd work, but after 3h, the solution was still colorless. For
Cl2 production I've used undivided cell (graphite electrodes, 20g/100ml NaCl, 300ml brine solution), voltage was 5V and current was 5 and up to 10A.
I just don't get it... What am I doing wrong?
I've seed other people on yt, using MnO2 and titanium electrodes, but I think it should work with graphite ones too? Am I wrong?
Thanks!
|
|
|
MeshPL
Hazard to Others
Posts: 329
Registered: 20-4-2015
Location: Universe
Member Is Offline
Mood: No Mood
|
|
Well, try using divided cell as your source of Cl2. Or acidify solution in your undivided cell. That should produce Cl2. You can bubble that through
NaOH and you'll have your bleach.
Hypochlorites are not efficiently produced by simple electrolysis. Nor is Cl2. To get hypochlorites you should preferably bubble Cl2 through basic
solution. To get Cl2 via electrolysis you need either divided cell or acidic enviroment (and chloride ions). That's because while on anode Cl2 is
formed, NaOH is "formed" on cathode. In divided cell that is not a problem (that won't go back and react with chlorine) and is not a problem when acid
is present, because than NaOH is neutralised by said acid. In undivided and nearly neutral cell however much Cl2 reacts with NaOH to form
hypochlorites and more chlorides. Although hypochlorites are your desired product, they can be further oxidised to chlorates, thus you will never
obtain as pure product as you may want.
Graphite is ok for electrodes. But remember it MAY not last as long as better anodes. Even when working as cathode it is slowly eroded (although not
as much as anode)
|
|
|
xfusion44
Hazard to Others
Posts: 223
Registered: 6-8-2014
Location: Europe
Member Is Offline
Mood: Nostalgic
|
|
Quote: Originally posted by MeshPL  | Well, try using divided cell as your source of Cl2. Or acidify solution in your undivided cell. That should produce Cl2. You can bubble that through
NaOH and you'll have your bleach.
Hypochlorites are not efficiently produced by simple electrolysis. Nor is Cl2. To get hypochlorites you should preferably bubble Cl2 through basic
solution. To get Cl2 via electrolysis you need either divided cell or acidic enviroment (and chloride ions). That's because while on anode Cl2 is
formed, NaOH is "formed" on cathode. In divided cell that is not a problem (that won't go back and react with chlorine) and is not a problem when acid
is present, because than NaOH is neutralised by said acid. In undivided and nearly neutral cell however much Cl2 reacts with NaOH to form
hypochlorites and more chlorides. Although hypochlorites are your desired product, they can be further oxidised to chlorates, thus you will never
obtain as pure product as you may want.
Graphite is ok for electrodes. But remember it MAY not last as long as better anodes. Even when working as cathode it is slowly eroded (although not
as much as anode) |
Thanks!
So, could I just add some HCl to undivided cell, to stop chlorine from reacting with NaOH?
Also, why didn't I get colored solution in undivided cell, as Cl2 reacts with NaOH - BTW, temperature was kept below 30°C?
[Edited on 7-2-2016 by xfusion44]
|
|
|
BromicAcid
International Hazard
Posts: 3237
Registered: 13-7-2003
Location: Wisconsin
Member Is Offline
Mood: Rock n' Roll
|
|
So you're just going by color to see if you're formed sodium hypochlorie? Have you tried any sort of qualitative testing? My guess would be that you
are forming sodium hypochlorite. What are you trying to accomplish, you say you performed electrolysis for 4.5 hours in the first post, that it was
pH 14 and that it smelled like bleach, sounds like you made sodium hypochlorite.
Are you trying to quantify the amount of sodium hypochlorite by color? Why not react with a known quantity of acetone and measure the chloform
formed? Not good for the environment but try it on a test-tube scale so you can measure your progress.
Did you have good stirring in your undivided cell to allow your chlorine formed to react with your hydroxide before it exited the solution? Chlorine
itself is not very soluble in water afterall. You keep mentioning your current at 5A and 10A, are you measuring it as such or are you using a power
supply and selecting that setting? What I am asking is, do you have an ammeter inline to measure the actual amps across the electrodes? Making
sodium hypochlorite using electrolysis is usually a pretty trivial matter, I am just trying to nail down where the issue is since there is no real
good reason that it should not have worked.
|
|
|
xfusion44
Hazard to Others
Posts: 223
Registered: 6-8-2014
Location: Europe
Member Is Offline
Mood: Nostalgic
|
|
| Quote: Originally posted by BromicAcid | So you're just going by color to see if you're formed sodium hypochlorie? Have you tried any sort of qualitative testing? My guess would be that you
are forming sodium hypochlorite. What are you trying to accomplish, you say you performed electrolysis for 4.5 hours in the first post, that it was
pH 14 and that it smelled like bleach, sounds like you made sodium hypochlorite.
Are you trying to quantify the amount of sodium hypochlorite by color? Why not react with a known quantity of acetone and measure the chloform
formed? Not good for the environment but try it on a test-tube scale so you can measure your progress.
Did you have good stirring in your undivided cell to allow your chlorine formed to react with your hydroxide before it exited the solution? Chlorine
itself is not very soluble in water afterall. You keep mentioning your current at 5A and 10A, are you measuring it as such or are you using a power
supply and selecting that setting? What I am asking is, do you have an ammeter inline to measure the actual amps across the electrodes? Making
sodium hypochlorite using electrolysis is usually a pretty trivial matter, I am just trying to nail down where the issue is since there is no real
good reason that it should not have worked.
|
Yes, I do have ammeter and both, voltage and amperage was measured. I agree, that I made NaClO, but since the solution wasn't yellow/green, I thought
it wasn't worth it, because I assumed that more concentrated solution has more yellow color to it. Yeah, that was exactly my plan - to use it, to make
chloroform. I thought it'd be more cheap to make NaClO at home, but now, I'm not really sure. No, I wasn't stirring the solution.
Thanks
|
|
|
AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
Try replacing NaCl with a concentrated CaCl2 solution.
The target is now, with occasional shaking (allowing the interaction of the formed chlorine water and a fine white suspension of Ca(OH)2 ), a visible
slightly soluble white precipitate, Ca(ClO)2.Ca(OH)2.
The latter is a more concentrated (and also less stable especially on exposure to air containing CO2 or heat) solid form of a hypochlorite that may
better fit your intended uses.
Related references, see http://www.google.com/patents/EP1966413A2?cl=en and http://www.google.com/patents/US3134641 . A more general reference is
http://www.google.com/patents/EP1966413A2?cl=en
[Edit] Note, replacing CaCl2 with MgCl2 would likely form a similar Magnesium compound (like dibasic magnesium hypochlorite, Mg(ClO)2.2Mg(OH)2.
[Edited on 11-2-2016 by AJKOER]
|
|
|
xfusion44
Hazard to Others
Posts: 223
Registered: 6-8-2014
Location: Europe
Member Is Offline
Mood: Nostalgic
|
|
| Quote: Originally posted by AJKOER | Try replacing NaCl with a concentrated CaCl2 solution.
The target is now, with occasional shaking (allowing the interaction of the formed chlorine water and a fine white suspension of Ca(OH)2 ), a visible
slightly soluble white precipitate, Ca(ClO)2.Ca(OH)2.
The latter is a more concentrated (and also less stable especially on exposure to air containing CO2 or heat) solid form of a hypochlorite that may
better fit your intended uses.
Related references, see http://www.google.com/patents/EP1966413A2?cl=en and http://www.google.com/patents/US3134641 . A more general reference is
http://www.google.com/patents/EP1966413A2?cl=en
[Edit] Note, replacing CaCl2 with MgCl2 would likely form a similar Magnesium compound (like dibasic magnesium hypochlorite, Mg(ClO)2.2Mg(OH)2.
[Edited on 11-2-2016 by AJKOER] |
Thanks for your suggestion!
|
|
|
|
![[*]](images/xpblue/default_icon.gif){kind=icon}
