Sciencemadness Discussion Board

MgSO4 --> H2SO4

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Saerynide - 24-3-2004 at 02:14

AHHH!!!! STUPID NO SUBJECT THING!!! I have to retype it all now!!! I spent 1/2 hour typing that!!!! :mad:


There seems to be a "sulfuric acid production race" going on here, with axehandle and his catalyst + burner, darkflame with his copper sulfate and me with electrolysis of MgSO4 (yes, im still working on that. Im a stubborn girl:)). I think I should start a new thread on this, because the old Copper Sulfate one was getting messy and deviating away from CuSO4 :D

I attempted again to separate MgSO4 into H2SO4 and Mg(OH)2. I finally got some *real* carbon rods from D cells (not using those crappy pencil leads anymore :D) and I replaced the papertowel salt bridge with a pvc tube to reduce all possibilities of contamination. A 30% (by wieght) MgSO4 solution was divided evenly into 2 containers and connected by the pvc tube and electrolyzed over night.

I woke up this morning and found that the anolyte was orange/amber AGAIN!! The catholyte seemed to be coming out ok like Mg(OH)2 was forming. Last time, this happened, we assumed it was iron. But this time, I am absolutely sure there was no ferrous materials in the cell. I had only glass containers, a pvc tube, carbon rods and solder (tin+lead) wires connected to the rods. I even wrapped the wired parts of the rods in tape to keep them dry. The wires also did not look corroded at all.

The only possibility of contamination would be from the solder wires I guess. I suppose I could try using copper wires hooked up the same way and see if the solution goes blue/green.

I am extremely confused now. Im not as into getting my H2SO4 now than I am in to finding out what this mysterious contaminant is. Its bugging the crap outta me!! Unfortunately, I do not have any NaOH, so I can't mix it with the anolyte to test for iron :(

Any ideas?

Edit: Filling a salt bridge is a REAL PAIN!! :mad:

Edit2: Stupid Tripod.....

[Edited on 25-3-2004 by Saerynide]

[Edited on 25-3-2004 by Saerynide]

axehandle - 24-3-2004 at 06:37

If it is the solder, you could do as I did when trying to electrolyse FeSO4:

Have the connection wires be normal copper wires, stripped about 10mm, going to the electrodes through pvc tubes that barely slip over the end of the carbon rod, holding the copper wire firmly against the carbon rod as well as providing an airtight seal.

I'm sorry I don't have any pictures, I dismantled the electrodes after my spectacular failure (using a KNO3 based saltbridge made of paper towel pumping full bridge rectified mains 220V at ~10A through the towel, setting it on fire...)

PM me if you don't get what I mean and I'll draw you a picture, it's kind of hard to explain.

Edit: Pssssst...: This is a cleaner way of connecting the ATX PSU, it allows you to use multiple cells and higher currents (forces you to cut off the ATX mobo connector though):

Edit: Found the anode! Lousy pic quality, but:

[Edited on 2004-3-24 by axehandle]

[Edited on 2004-3-24 by axehandle]

[Edited on 2004-3-24 by axehandle]

[Edited on 2004-3-24 by axehandle]

Polverone - 24-3-2004 at 12:53

Here's a shot in the dark: that brown solution reminds me of what I get when I dissolve MnO2 in HCl. Since your electrodes came from a battery, is it possible that they are contaminated with manganese dioxide?

Does anybody do arc welding in Singapore? A place that sells welding supplies should sell carbon "gouging rods" that are cheap, clean, and designed to be used as electrodes (you may have to peel off the copper sheath though).

Edit: you should try to see if you can precipitate that brown junk. Try adding sodium carbonate and simmering it down. If it's coming from a metal salt, it should precipitate at some point. Then you might be able to figure out what's causing the discoloration.

[Edited on 3-24-2004 by Polverone]

hodges - 24-3-2004 at 16:15

Good idea, Saerynide - using the PVC pipe for a bridge vs using a paper towel. I assume you filled it with the same solution as in the two jars?

I really have no idea why the anode side turned brown. Are you using epsom salts for MgSO4? That should be pretty darn pure. You got exactly the same results using a paper towel and pencil leads? That would rule out contamination from the PVC and carbon rods.

Can you test your H2SO4 by adding a small amount of it to some baking soda and seeing how vigorous of a reaction you get?

Saerynide - 24-3-2004 at 20:02

Yeah I filled the pvc tube with the same solution in the jars.

Im out of baking soda now. I guess I'll try to run out at get some later today. To test for iron, I think I'll add vinegar to a bit of the solution and see if I get red insoluble ferric acetate. <-- Merk index is so useful :D

darkflame89 - 25-3-2004 at 02:01

This idea is soooo good. Anyway, i was with my separation of two soluble salts(CuSO4 and magnesium ethanoate)

I am going to try this out too as and when i have free time. The problem right now is that i cannot find graphite rods. Chemicals and equipment are pretty limited in Singapore. All we have is a homefix store and a pharmacy that caters to our health. I have got two choices: get carbon rods from batteries or buy those kind of charcoal drawing pencils, and try to get the leads out(i dun'y really know how:( )

As for the steel cathode, i am hoping to use a stainless steel bowl and connect this to wire. Is it ok to do so?

I also intend to do the flower pot method during the June holidays. Instead of using lead cathodes, can i use steel or carbon?

oh ya, Saerynide, could you post the pics again? The pics dun seem to be able to load correctly

[Edited on 25-3-2004 by darkflame89]

Saerynide - 25-3-2004 at 02:36

I tried charcoal pencils. They disintegrated really fast under high current or high voltage :( So I went for batteries. They seem to hold quite well (so far).

Yeah you can use carbon instead of lead, unless there's some strange mystical power of lead that I dont know of :o Stainless steel works too.

As for your stainless steel bowl, that sounds like a very good idea. I think Ill try it on my next run :D Thanks.

Saerynide - 25-3-2004 at 19:51

Ok, I tried adding vinegar with no luck :( I did not get any precipitate at all. Maybe it's because the H2SO4 in there keeps the ferric acetate dissolved, preventing it from precipitating?

Im heading out now on bike to get some baking soda to test the strength of the acid. :D

[Edited on 26-3-2004 by Saerynide]

darkflame89 - 26-3-2004 at 04:25

Ok. i bought my box of Epsom salts. I am so happy.:D Also bought 2 D batteries so that i can get the carbon electrodes.

Anyway, your brown solution looked like Fe(OH)3 to me...

Edit: Ok, not iron hydroxide as it is insoluble. Maybe some other iron (III) compunds..

[Edited on 26-3-2004 by darkflame89]

Saerynide - 26-3-2004 at 04:33

Fe(OH)3 is insoluble in water. I think it looks like FeCl3.

I highly doubt it FeCl3 though, the only anion in there should be SO4-, unless... the tap water was *that* impure?? (I used tap water on this try because I didnt feel like wasting dH2O on the first test run :P)

darkflame89 - 26-3-2004 at 04:39

How and what did you wrap your your cathode?

Saerynide - 26-3-2004 at 06:42

I wrapped them with clear mover's tape, but don't do that - it doesnt work too well. I suggest several layers of rubber electricians tape instead. I used that under water a few times and it kept stuff dry :D

I still havent gotten around to building my plastic-tube protected anode yet. I'll try doing it tonight and get something posted by tomorrow :D

t_Pyro - 26-3-2004 at 09:18

To protect the electrical connections between the electrode and connecting copper wire, I generally coil a length of the (stripped) copper wire around the electrode, wrap it with a few layers of tape, and then melt the tape so that it can fill in the spaces and make the connections air-tight. Unfortunately, this makes it kind of sticky, but an additional layer of tape makes handling easier.

Cl2 precense in the tap water.

axehandle - 26-3-2004 at 10:06

Many water purification plants (of older design) add Cl2 to the water to kill microorganisms. That may be where the Cl is coming from.

Just remembered it.

The MgSO4 could have a very slight contamination of FeSO4, as well. That would account for the iron.

As well, the water pipes in your house are possibly made of iron. That's another source. Can't think of anything else now...

But it sure <b>looks</b> like FeCl3.

[Edited on 2004-3-26 by axehandle]

[Edited on 2004-3-26 by axehandle]

[Edited on 2004-3-26 by axehandle]

Saerynide - 26-3-2004 at 11:41

Wow, Im not half bad at recognizing common chemicals :D Hehehehehe...

Thanks for those suggestions. They really helped put all the pieces together :)

Anyways, tomorow morning, I'll do the second run,with distilled water instead and with my new protected electrodes (if theyre dry by then).

If there's still any contamination... then, I'll be damned and its coming from the MgSO4 :o

Must head to bed now, its 3:40 am :D

hodges - 26-3-2004 at 16:43

Originally posted by Saerynide
Ok, I tried adding vinegar with no luck :( I did not get any precipitate at all. Maybe it's because the H2SO4 in there keeps the ferric acetate dissolved, preventing it from precipitating?

Try adding NaOH, or Na2CO3 to see if iron hydroxide or iron carbonate precipitate.

Organikum - 26-3-2004 at 17:54

btw. Saerynide - dont you have a "back" key on your keyboard? Because of retyping everything.....
Works everytime for me.

Also the browser might have such fancy buttons with meaningless arrows - at least mine has.

Saerynide - 26-3-2004 at 19:51

Those dont give you back what you typed :( They just reload the page.

[Edited on 27-3-2004 by Saerynide]

Organikum - 26-3-2004 at 21:18

Oh! Mine gives me back what I typed. Mabe its because I use OPERA? :o

Browser Cache

Turel - 26-3-2004 at 21:28

You can set the size of your browser cache in your tools options. When you hit back and it reloads the page, this is because your browser is not caching the previous page. Try increasing your browser cache by 10MB to fix the issue.

Saerynide - 27-3-2004 at 00:12

Blah! I have silicone all over my hands from doing the final sealing on my protected electrodes, so im typing with 3 fingers and its hard :mad:

I added the baking soda to the contaminated solution til slightly over saturation, it fizzed alot (there is H2SO4 :D) but, I got no iron carbonate precipitate. The excess soda sat at the bottom, and the solution stayed brown, but it got lighter in colour + murkier.

I just checked the Merck index, and it says that ferrous carbonate is soluble in CO2 saturated water..... So maybe there was FeCO3,but it stayed dissolved? I didnt know this before, so i dumped the whole thing down the fscking drain instead of leaving it to settle out!!!! :mad:

Now we'll never know, and it's all my fault :(

darkflame89 - 27-3-2004 at 21:49

OK, I need to sort this out; I'm new to electrochemistry, and i feel that electrolysis has a lot of potential to produce all kinds of acids, bases, and salts that i need.

In this case, at the cathode side, H+ ions and Mg2+ ions are present in the water. Mg(OH)2 is formed because of the discharge of hydrogen gas. Which leaves the hydroxide ions behind. This then reacts with the magnesium ions to form Mg(OH)2. At the anode side, the copper anode(if i use one) will dissolve to form copper sulphate. Correct??

Therefore, in such a setup, i can produce a base or alkali at the cathose side, and produce salt(if i use a active electrode) or acid(if i use a inert electrode) at the anode. Therefore i can produce all sorts of chemicals by this woderful method!:D

Esplosivo - 27-3-2004 at 23:09

Saerynide, on adding NaHCO3 you said that no precipitate of any kind was formed?! We can therefore exclude the fact that the salt was an iron (III) since that gives a rust brown ppt of iron (III) hydroxide (since Iron (III) carbonate is unstable and decomposes on formation). We can also exclude any Iron (II) Carbonate forming, which is a dirty green precipitate.

I doubt that the fact that carbon dioxide might increase the solution of iron (II) carbonate. On the contrary. The CO2 should reduce the solubility because of the 'common ion effect' on a slightly soluble compound. Even if true, the amount of iron (II) carbonate produced would still be insoluble in water, there would be too much of it. Next time boil the water after adding NaHCO3 to remove any CO2 gas dissolved.

Better still, next time ask for some better and more accurate qualitative tests. It would help you to know what is going wrong in that cell.

Saerynide - 28-3-2004 at 00:17

Darkflame, it should work, and does for most people. But the last time I tried to get CuSO4 like that, it didnt work. I dont know why but I just got a crapload of light blue copper hydroxide precipitate :( Maybe you'll have some luck or maybe I just messed up :D Tell me how it goes if you try it :)

Esplosivo, next time I get such a contamination, I'll be sure to:
1) not chuck the solution after the test
2) take a picture after the test
3) boil the solution
4) ask for more tests

If only I had NaOH :mad:

Oh well, Im setting up the second run now. Hopefully, we wont need to resort to the above 4 resolutions :D

Edit: Here's my new pvc protected electrodes. They're gonna be used for the HCl/NaOH cell Im planning on making this summer too :D

[Edited on 28-3-2004 by Saerynide]

[Edited on 28-3-2004 by Saerynide]

darkflame89 - 28-3-2004 at 04:47

You got copper hydroxide... Well, i tried the electrolysis of salt water( not really concentrated) with copper electrodes. I did not separate the solution into 2 bowls. What i got was a green insoluble solid which i think was copper hydroxide. I am quite sure it is copper hydroxide, cause the last time i reacted NaOH and CuSO4 i got a green solid too.

That is why i wanted to make the electrochemical reactions at the 2 half cells..

t_Pyro - 28-3-2004 at 10:07

Copper hydroxide is blue. What you have might be copper chloride.

[Edited on 28-3-2004 by t_Pyro]

Esplosivo - 28-3-2004 at 11:03

It cannot be copper chloride too, since it is soluble and cannot precipitate, though it does have a bright green colour. This salt is also hydgroscopic. Most chlorides are soluble (some exceptions like AgCl and PbCl2, the latter of which is also soluble in warm water). Copper Carbonate is light blue, and it is insoluble forming precipitates (colour differences might be due to impurities).

t_Pyro - 28-3-2004 at 11:12

Yeah, you're right about the solubility of the copper chloride- drying a sample of it takes ages for me! I can't think of any green inslouble compounds... Dark/dirty green could be ferric, dark green could also be a chromium salt. However, I've never heard of a green ppt on adding NaOH to CuSO<sub>4</sub>.

Organikum - 28-3-2004 at 12:24

Dont judge a copper salt by its color.....

I just added hot concentrated NaOH solution to a concentrated CuSO4 solution and got some dark-green precipitate which dissolved under the addition of more water to give a almost black solution.

Then I added some CuSO4 to an cold NaOH solution and got a deep-blue precipitate.

Temperature and concentration matters :D

[Edited on 28-3-2004 by Organikum]

Organikum - 28-3-2004 at 12:49

And after my quick and dirty experiments the hard truth from the MERCK-index:

- copper sulfate dibasic: Cu3H4O8S, blue-green, rhombic, bipyramidal crystals. Practically insoluble in water.

- copper sulfate tribasic: Cu4H6O10S, very finely divided, light-blue, gelatinous particles. Practically insoluble in water.

- Copper hydrate: Cu(OH)2, blue to blue-green gel or light blue crystalline powder. Stability is dependent on the method of preparation, may decompose to black CuO on standing a few days or upon heating. Practically insoluble in water. Sol. in concentrated alkali when freshly precipitated.


So I got CuO mostly in my first experiment and a mixture of everything in my second experiment I guess (no CuO up to now though)

hodges - 28-3-2004 at 14:47

Originally posted by Saerynide

Here's my new pvc protected electrodes. They're gonna be used for the HCl/NaOH cell Im planning on making this summer too :D

Your electrodes look good. You can get NaOH by electrolysis of NaCl solution. But at the anode you will get Cl2 instead of HCl. The Cl2 will escape into the air.

Saerynide - 28-3-2004 at 15:22

Yeah I know. Im making bubblers for the Cl2 :D

t_Pyro - 28-3-2004 at 21:11

Bubbling the chlorine through water will result in HOCl and some dissolved chlorine. Check with some blue litmus- it'll turn red, and immediately white. You might not even be able to detect the red.

Maybe you could make a jet of the H<sub>2</sub> gas to burn in an atmosphere of Cl<sub>2</sub> to get HCl, and then dissolve it in water (use an inverted funnel for this). The beauty of this method would be that the H<sub>2</sub> gas could be taken from the cathode, Cl<sub>2</sub> from the anode, so in effect, you'd actually be splitting up a salt of a strong acid/ strong base into the constituents, NaOH and HCl!

Saerynide - 29-3-2004 at 00:26

I wouldnt dare mix H2 with Cl2. It becomes explosive :o

Also, the HClO formed would break down into HCl and O2. Ive tried shaking Cl2 into water before, and it does yeild HCl. I saw the O2 bubbles coming out of the water :D

darkflame89 - 29-3-2004 at 01:52

Haha, that's what i intend to do anyway. I want to make the 3 most needed acids mostly by this method. HCl acid, H2SO4 acid, (dunno its possible to produce HNO3 this way).

Saerynide - 29-3-2004 at 02:10

There was some thread awhile ago saying that NO3- can be reduced at the cathode. Which makes making HNO3 this way would be very hard and tedious, but if there's a will, theres a way :D I forgot which thread it was though...

Esplosivo - 29-3-2004 at 02:12

Saerynide, as long as I remeber the decomposition of HOCl to HCl takes a long time to occur. Also Cl2 dissolves in water to give the so-called 'chlorine water'. The Cl2 dissolved in water will decompose to HOCl in the presence of light.

Organikum, sorry for my ignorance, but how can copper sulphate be di/tri basic?! I've never heard of these salts before. How are they made?

darkflame89 - 29-3-2004 at 02:12

Never mind i can always use sulphuric acid on nitrate salts to produce NO2 and dissolve them in water to get HNO3.

Or i can always try to biuld a Birkeland-Eyde reactor..

Saerynide - 29-3-2004 at 02:16

Im gonna put the HCl/NaOH maker outside anyways, so there'll be lotsa light for it. Im so not getting Cl2 poisoning again, so it wont be anywhere near indoors :D

darkflame89 - 29-3-2004 at 02:21

Liked to clarify somthing.. In this setup of the MgSO4, won't the OH ions move over to the anode side instead of producing Mg(OH)2?

Saerynide - 29-3-2004 at 02:53

Mg(OH)2 is insoluble and it precipitates out of the solution :D You get a slightly milky solution if its dilute, but you get a layer of white amorphous stuff if its concentrated.

Edit: Nooooooo!!!! Contamination again!!!! :mad: :mad: :mad:

How far do you hafta go to have no fscking contamination?!?! I already washed ALL the equipment with dH2O TWICE already :( Maybe the silicone doesnt make as much of a water proof seal as I hoped :o

The contamination's not as bad this time though. I tested the pH of the two solutions. It was 3 and 11 only :(

[Edited on 29-3-2004 by Saerynide]

Esplosivo - 29-3-2004 at 03:22

Come on Saerynide. Look at it on the good side, now you can carry out an analysis on the stuff :P. Its fascinating though. I cannot immagine the source of impurity. What water did you use? What was the source of the chemicals electrolyesed?

Saerynide - 29-3-2004 at 03:52

Lol, I guess I could have some fun with it now :D But no H2SO4 for making sugar turn to carbon or making oil of wintergreen :(

I used bottled distilled water from the supermarket and MgSO4.7H2O from the drug store, so it *should* be pretty pure. I also covered the cell with Glad wrap to keep dust from getting in.

I just thought of something. There would be an equilibrium in the cell because the electrodes can only draw so many anions and cations to them at any moment, so some ions would be crossing the salt bridge and reacting back with the acid and base to make the salt again, right? That would mean, the pH can only go so low, and would level off. So, if after each run, I filter the catholyte out to remove the Mg(OH)2, then run it again, and keep repeating this, I would eventually remove all the Mg2+ from the cell? :o That would leave me with a salt free acid?

Considering the fact that I am extremely tired/sleepy right now, it might be all gibberish :D

Esplosivo - 29-3-2004 at 04:02

I think you're right. Well, Mg2+ ions would move to the cathode, therefore there is not need to filter after each run. One must calculate the point at which the salt will be all used up (and I have no idea how this can be done, probably when Mg(OH)2 stops forming). Btw, I do not think that you will be able to get an acid totally free of the salt. Some would still remain, though in very small conc.

Just call in when carrying out the qualitative analysis. (I love analysis lol)

Saerynide - 29-3-2004 at 05:16

Originally posted by Organikum
And after my quick and dirty experiments the hard truth from the MERCK-index:

- copper sulfate dibasic: Cu3H4O8S, blue-green, rhombic, bipyramidal crystals. Practically insoluble in water.

- copper sulfate tribasic: Cu4H6O10S, very finely divided, light-blue, gelatinous particles. Practically insoluble in water.

- Copper hydrate: Cu(OH)2, blue to blue-green gel or light blue crystalline powder. Stability is dependent on the method of preparation, may decompose to black CuO on standing a few days or upon heating. Practically insoluble in water. Sol. in concentrated alkali when freshly precipitated.

Ohh!!!! The time when I tried to make CuSO4 using a paper towel salt bridge, MgSO4 electrolyte ,and a copper penny anode, I *did* get CuSO4 but I thought I got Cu(OH)2. So that light blue precipitate was the tribasic stuff?? :o

Btw, I looked up the definition of di/tribasic, and I dont understand how CuSO4 could fit the definition. Also, I dont understand why there are all those extra O's, H's and Cu's :o

Esplosivo - 29-3-2004 at 06:08

Same problem here. I've searched again and again and I cannot find any basic copper sulphate, neither di- not tri- basic. Are these complex salts or basic salts. Copper ions may form complexes having 'strange' colours. Well if they exist I would like a simple experiment for making one plz.

Organikum, not to doubt your knowledge, but are the chemical formulas stated previously correct. They do not seem correct to me, though I might be wrong.

[Edited on 29-3-2004 by Esplosivo]

t_Pyro - 29-3-2004 at 08:01

When you're talking of a dibasic/tribasic <i>acid</i>, it refers to the number of replaceable hydrogen atoms. For sulfuric acid, it's 2. I've never heard of any dibasic or tribasic salt, though what he might have meant might have been the oxidation state of the element in the compound. The "Cu3H4O8S" might be Cu<sub>3</sub>(OH)<sub>4</sub>SO<sub>4</sub>, where the copper is in the +2 state. This compound is new to me, though. I would have thought the hydroxyl ions would have been easily neytralised in the solution... The "Cu4H6O10S" might be Cu<sub>4</sub>(OH)<sub>6</sub>SO<sub>4</sub>, but I have the same misgivings against it...

Copper ions do form complexes, but I wouldn't call them "strange" coloured! A simple way to prepare tetramine copper sulfate is to add excess of ammonium hydroxide to a solution of copper sulfate. The dark blue solution is tetramine copper sulfate, a simple coordination compound. Nickel and cobalt salts respond to ammonium hydroxide in a similar manner, forming coordination compounds.

Coming back to chlorine and hydrogen: A <i>mixture</i> of H<sub>2</sub> and Cl<sub>2</sub> is explosive, as is a mixture of H<sub>2</sub> and O<sub>2</sub>. However, if pure H<sub>2</sub> is burnt as a jet in an atmosphere of O<sub>2</sub> or Cl<sub>2</sub>, the reaction is quite smooth.

[Edited on 29-3-2004 by t_Pyro]

Esplosivo - 29-3-2004 at 08:23

Ok lol. Sorry for my terminology. By 'strange coloured' I wanted to mean that it is not common to see copper (II) sulphate which is green, at least for me. I still don't know how those basic compounds named may be formed.

Saerynide - 30-3-2004 at 08:25

What should I try tomorrow? I could 1) filter out Mg(OH)2 and electrolyze again or 2) Try to plate out the metal ion contaminant.

Id like to try plating out the metal, because it will either remove the contaminant (best case scenario), or end up concentrating the acid to the point that the bit of metal is irrelevant (still good) :D

But, how am I gonna keep the freshly made metal from reacting with the H2SO4? More over, what if the H2SO4 reacts with my cathode?? :o

Analysis would have to wait til Friday. Ive got a ton of tests to study for :(

Btw, how do you guys fill your salt bridges? I had to shove a smaller tube into my bridge, dunk that under the solution, and use a syringe to suck up the liquid. This way is *very* painstaking cause I have to do it many times. Bubbles always get in and liquid goes into the syringe! :mad: I was so desperate once I even tried syphoning up the solution. That left me choking on a mouthful of nasty MgSO4 :(

Does anyone pity poor me?? *sniffle*

darkflame89 - 31-3-2004 at 01:37

You could try and plate out the metal. You will have to keep an eye on it. When the solution turns clear( i hope this does no take too long), you can just remove the cathode and anode. I dunno about this, hope it works.

About the filling of salt bridge, i am worrying over it too.

Another thing, i need to know how long i can run the setup if i have a 9 V battery or a 18 V series of batteries.

darkflame89 - 31-3-2004 at 02:54

got somthing to help in the quest for purer sulphuric acid:D

See attachment:

acidfilter.jpg - 22kB

Saerynide - 31-3-2004 at 03:06

Dont run on batteries. It'll cost you a fortune :o If you have an old computer that you're not using, take out its power supply, short the green to a black on the mobo connector and voila! Instant electrolysis powersupply :D To run a cell, hook up a yellow or red for the positive and a black for the negative. Check the key stuck on the power supply to see the voltages and amperages of each lead.

Nice idea for the acid filter :D But can we get activated charcoal here in drug stores? Also, were could we buy fiberglass cloth :o

darkflame89 - 31-3-2004 at 03:10

I dunno, but i will check up real soon in the stores. As for your old computer idea, its good but will it damage the computer permenently? :o

Saerynide - 31-3-2004 at 03:21

Well, I heard that power supplys can get killed from overloading and stuff, but mine has been abused for a long time and its still alive :D

Just take the power supply OUT of the computer to keep your hardware safe.

Ask axehandle, hes good with this stuff :)

Esplosivo - 31-3-2004 at 04:20

Activated charcoal can be bought at aquarium suppliers. It comes quite economically and is capable of absorbing many ions, especially ones which cause water to be coloured. What would its use be in filtering the H2SO4? To remove the colour?

And by the way, being prills, the charcoal can be place in a container with small holes at the bottom. No need for glass wool.

Saerynide - 31-3-2004 at 05:04

Sweet :D I must check my local aquarium shop :D

Btw, how does it "absorb" ions? Does it react with the ions? :o

Edit: About how many times can the carbon be reused?

[Edited on 31-3-2004 by Saerynide]

Esplosivo - 31-3-2004 at 05:31

The absorption of activated charcoal works because 'acticvated charcoal' is highly porous and has a very large surface area to which it can absorb chemicals. Activated charcoal absorbs mainly chloride ions and organic molecules. I don't think it reacts with the ions though, since it is no form of ion exchange substance used in ion-exchange columns.

Activated charcoal is also sold in pet shops. I cannot find a specific quantity to be used. Just filter through the solution until all the colour has been removed. When no change seems to occur replace the charcoal.

I am sure that certain ions like sodium and nitrates are not influenced by charcoal filtration, but ions like iron and chloride are. Activated charcoal is used in gas mask to filter our certain chemical warfare agents (mostly organic ones) - Just to show you it absorbing capability.

Organikum - 31-3-2004 at 09:17

Don´t wan to take the fun away Saerinyde but if you go to the forums ftp and download the "Kings Chemistry Survival Guide" you will find under "electrochemical methods" exactly what you are after:
Electrolysis of MgSO4 to yield magnesium hydrate and dil. H2SO4.

The book is the hit anyways!
Highly recommended!

t_Pyro - 31-3-2004 at 19:43

Silica gel is yet another good adsorbent. You can pass a solution through a column containing silica gel to adsorb the more easily adsorbed ions. To "recharge" the silica gel, it has to be treated with dilute HCl to get rid of the adsorbed ions, then dried in a microwave oven.

Activated charcoal can be "recharged" by heating it gently, if I remember correctly.

Activated charcoal and silica gel "adsorb" substances by forming weak bonds (due to Van der Waal's forces) between the surface of the adsorbent and the adsorbate. Finely divided metals adsorb certain gases by forming weak bonds with them at the surface.

Absorbents, on the other hand, "absorb", or incorporate into the bulk of the substance, the substance to be absorbed.

[Edited on 1-4-2004 by t_Pyro]

Esplosivo - 1-4-2004 at 04:12


Activated charcoal can be "recharged" by heating it gently, if I remember correctly.

No not gently, you need very high high temperatures. It must start glowing red hot for around 5 minutes.

Saerynide - 1-4-2004 at 06:22

My god, thats like burning it. I might as well buy some new stuff cause it'll probably cost less than the gas needed to heat it :o

Esplosivo - 1-4-2004 at 07:13

Ye you're right. Where I live aquarium fanatics used to 'reactivate' the charcoal by heating it in a baker's oven (the real big ones) and said it worked, though they lose quite a lot of it. It isn't worth the trouble, better buy some when it is needed.

Saerynide - 1-4-2004 at 08:23

I just read the part about electrolytic processes, and it said you have to "charge" the flower pot with MgSO4 even if you use electrolysing NaCl or anything else. Whats the point of this? Wont it just lead to contamination?

Organikum - 1-4-2004 at 08:50

I don´t think this will lead to contamination. Maybe the charging would also work with NaCl or something else, but I guess that the ions formed from MgSO4 are most prone for permeating the pot and thats what the procedure is good for: Charging the pot with ions - once charged also other ions will pass through with ease whats not sure - or may take a very long time otherwise.

Contamination is no issue regarding the amount of ions which will stay in the wall.....


Saerynide - 3-4-2004 at 12:01

I electrolyzed the contaminated H2SO4 for a few hours. The stainless steel knife was coated in a tan coloured stuff that looked like MgO which could be easily scraped off the knife. Then I filtered the solution 4 times, using more and more filters each time. The acid is now almost clear, and the pH has decreased slightly, from 4 to 3.5ish, judging from the indicator strip.

Tomorow, I will electrolyze it some more, then try to concentrate it to about half its volume :D

I just had to give this a shot

Quantum - 3-4-2004 at 20:19

I got a flower pot with the hole plugged and a plastic box where I placed 5000ml of cold tap water in the box and 1500ml in the pot. I took a sheet of pglass and drilled a hole in the middle to support the electrode. About 330g of MgSO<sub>4</sub> was disolved into the water in the box and 100g in the pot. both electrodes were made of lead from fishing weights.

I am running this at 5v/20amps. At first nothing seemed to happen but after a while small bubbles streamed off so the current densitry is low but this seems to be ok. The pH of the outer solution is red at about 4 using my paper. The solution is beautifly clear no hint of contamination from the electrodes. In the inner pot solution the pH is green at about 9 due to the pressance of Mg(OH)<sub>2</sub>. I don't see a persipitate of Mg(OH)<sub>2</sub> yet but I have only run the cell for a few hours so it may be yet to come.

My cell is huge; 6500ml total! If I can get good cheap clear H<sub>2</sub>SO<sub>4</sub> from it then I may never buy the stuff again! A good cheap source of Mg(OH)<sub>2</sub> will be handy as well. If this trial run is good then the "S Cell" may be a permanent addition to the garage/lab area. I could just push it into the corner. :D

The salt bridge Saerynide is using seems dubious to me. Filling it would be hard and it may contaminate the 2 solutions resulting in the formation of more MgSO<sub>4</sub>!:o

Next day I will pipit off some liquid and do my "Acid Test"(pun intended):D by heating my 200w soldering iron up and sticking the tip in the liquid. If all goes well I will see SO<sub>2</sub> fumes(hopefully not so much as to injure me <b>again</b>!

Wait till the 'morro for the latest results!

Saerynide - 3-4-2004 at 21:09

Yeah, I mentioned earlier that the salt bridge might lead to an equilibrium being formed between the two cells.

I just thought of something. There would be an equilibrium in the cell because the electrodes can only draw so many anions and cations to them at any moment, so some ions would be crossing the salt bridge and reacting back with the acid and base to make the salt again, right? That would mean, the pH can only go so low, and would level off.

And my suggestion for solving it would be:

So, if after each run, I filter the catholyte out to remove the Mg(OH)2, then run it again, and keep repeating this, I would eventually remove all the Mg2+ from the cell? That would leave me with a salt free acid?

The only I reason Im using this salt bridge is because I cant find an unglazed flower pot :mad:

Quantum - 3-4-2004 at 21:17

Living in Singapore can't be so good for free market capitalism; its the size of one of our states! It seems to me that rockwool inside pvc tube would be a better bridge and that a wooden bowl could act as a better bridge but not for strong acids/bases! It might contaminate the solutions even then. You might want to try a 'glazed' pot even then. I got one I thought was glazed(it felt smooth) but it works. Now if its shiny and a difrent color then its a safe bet it realy is glazed:(

darkflame89 - 4-4-2004 at 00:23

I have lots of flower pots at home. There's an orchid farm near here and my mum likes gardening. Ok the pots are quite big with the ability to hold at least 2L. I think i will get a small bucket to hold this. As for lead electrodes, its virtually impossible to get hold of them. So, copper electrodes are ok? I can always recrystallise the CuSO4 produced and electrolyse them later.

Saerynide, you might want to try the orchid farms near the northern part of the island. They have a wonderful supply of pots and $10 each.

darkflame89 - 4-4-2004 at 02:08

Incidentally, how much MgSO4 do i need to dissolve into the solution? Or will any amt do?

Saerynide - 4-4-2004 at 02:21

I'll give one of those orchid farms or those warehouse-sized flower shops a try. Probably wont have time to get one til around May cause of all the exams coming up :(

I heard that 30% salt (by wieght, I would assume) was best for electrolysis. I have no idea why though, maybe it has something to do with diffusion and equilibrium? Maybe someone here can explain it to me or correct me :)


Quantum - 4-4-2004 at 06:03

I ran the cell over night and the pH is 10 for the Mg(OH)<sub>2</sub> and 2 for the H<sub>2</sub>SO<sub>4</sub>!:o The pot has lots of milky white Mg(OH)<sub>2</sub> coating the inside of it and on the bottom. After doing the Acid Test:D by sticking a hot soldering iron in it and looking for SO<sub>2</sub> I saw some! Now I am going to boil the 5000ml down outside!

Nearly free Sulfuric Acid at last!

Saerynide - 4-4-2004 at 06:46

woooohoooo :D

hodges - 4-4-2004 at 13:02

Originally posted by Quantum
I ran the cell over night and the pH is 10 for the Mg(OH)<sub>2</sub> and 2 for the H<sub>2</sub>SO<sub>4</sub>!:o The pot has lots of milky white Mg(OH)<sub>2</sub> coating the inside of it and on the bottom. After doing the Acid Test:D by sticking a hot soldering iron in it and looking for SO<sub>2</sub> I saw some! Now I am going to boil the 5000ml down outside!

Nearly free Sulfuric Acid at last!

Good work! But you know how far you will have to boil this solution down? A pH of 2 implies 10^-2 moles of hydrogen ions per liter. At best, this solution is 0.01 molar (and it is actually worse than that, because H2SO4 has two hydrogens to ionize). To get a 6 molar solution, which is the concentration of battery acid, you will have to boil it down by a factor of 600! If you boil that down by a factor of 3 or 4 you will then have concentrated H2SO4!


Quantum - 4-4-2004 at 18:20

I have not yet been able to boil it down but somehow I don't think my yeilds will be very good:( I may be better off with buying it but I called the only hydroponics store nearby and they don't know whats in the acid solutions! They do however have 35%H<sub>2</sub>O<sub>2</sub>!

Do you guys think its worth buying some "pH Down" and then testing it by heating a drop till I hopefully see SO<sub>2</sub> fumes is a fairly safe bet? Im willing to try most anything.

Great Progress!

darkflame89 - 5-4-2004 at 02:04

Tried this crude version over the sunday. I set up 2 bowls with MgSO4 dissolved in them.The anode side ihas slightly more solution than the cathode side. For the salt bridge, i used a paper towel. Then i connected the whole thing to a 9V battery. I used copper electrodes as i do not have any carbon electrodes yet. Over 1 night, the anolyte turned blue and at the cathode side, there is a white coating of Mg(OH)2. Additionally, there are pieces of copper in the anolyte as the copper electrode corroded away.

The cell is still running now, and i am so excited!!:D:D There are pictures attached:

DSCF0452.JPG - 879kB

darkflame89 - 5-4-2004 at 02:05

And another one:

DSCF0453.JPG - 880kB

darkflame89 - 5-4-2004 at 02:07

And another one:

DSCF0454.JPG - 875kB

Esplosivo - 5-4-2004 at 02:21

Good work. Seems to be working pretty fine. After you've changed the electrodes to carbon rods just post a picture (1 not 3 :P) of the apparatus after a run. This is just to check for any brownish coloured solution, which might have been 'masked' by the copper (II) sulphate formed in your case. Thanks.

Saerynide - 6-4-2004 at 01:57

God dammit. I ran the cell again, and this time there was contamination AGAIN. Same tan/brown shit. I swear, the MgSO4 must be contaminated.... Shit like this doesnt hapen 3 times in a row :mad:

Im getting myself a NEW box of MgSO4 tomorrow :mad:

darkflame89 - 6-4-2004 at 02:18

Ok, liked to add this pH test for my catholyte is 10. And also i can have a nice blue colour for anolyte, with some pieces of copper inside.

I don't have any carbon rods at the moment, i tried to pry the metal cap off the battery but i can't!!!:mad:

Saerynide - 7-4-2004 at 02:32

You guys wont believe what happened last night. I had my cell running on my desk, and my cat was kinda frisky for some reason. He trampled all over my desk like he usually does, but he usually stays away from where I run my cell. Well last night, he wandered over there to take a look, and as he turned around to leave, he swipes the anolyte RIGHT OFF the desk before my very eyes :o I caught the container as it fell, so thankfully i didnt get a mess of broken glass. But I did get a huge splash (complete with long trickle marks) on the wall and a huge splash + puddle on the floor. I had to clean the whole mess before mom came up stairs and was rushing like freak to soak up the mess. Paper towels dont absorb salty + acid water very well :mad:

Esplosivo - 7-4-2004 at 21:52

I have set up and run the electrolysis for approximately 13 hrs right now. At the anode the pH is already acidic and there is no sign of contamination. There is also a white precipitate of magnesium hydroxide at the cathode. The only problem is that the carbon electrodes I've used are disintegrating into small black particles. The carbon electrodes are from battries. Does anyone have a suggestion to prevent this?! These particles can be easily filtered off though. If I will have the time I will carry out a titration and a simple analysis after another 12hrs.

Quantum - 7-4-2004 at 22:01

Like I said before I have already run my cell and it works fine but sadly its realy not good for making sulfuric unless you can get 10s of liters of the very dilute acid.

Now I have 5 liters sitting around along with some Mg(OH)<sub>2</sub>

I am going to pick up some battery acid soon along with a cheap hotplate.

hodges - 8-4-2004 at 06:12

Originally posted by Esplosivo
The only problem is that the carbon electrodes I've used are disintegrating into small black particles. The carbon electrodes are from battries. Does anyone have a suggestion to prevent this?!

Keep the temperature of the solution down and keep the voltage down. You might try putting a lamp of the same voltage as your power supply in series with your circuit to drop some of the voltage. You really only need a few volts across the electrodes; a higher voltage speeds up the proocess but produces fast erosion of the carbon rods and may also result in production of other products (such as reducing the sulfate ion to sulfite or even sulfide).

Organikum - 8-4-2004 at 08:07

I guess the main problem which arises here whilst electrolysing MgSO4 is connected to the ph.
A ph of 10 in the anode chamber just doesnt sound favoirable to me.

MgSO4 is just not very good for this, I would suggest to use FeSO4 so any available, which would for sure work much better and doesnt even require a diaphragm.

If you want to stick with the MgSO4 I would suggest to change the liquid in the anode chamber often, removing the basic MgOH and refilling with saturated MgSO4 solution. The cathode liquid should stay the same, H2SO4 in the cathode chamber favors the process.

god luck!

darkflame89 - 9-4-2004 at 00:19

Gaaargggghhhh!:mad:When i tried to prise off the metal cap of the battery today, some black stuff presumbaly carbon, spurted out of the can!!!:mad:It reached 1.5 metres from the table and hit the ceiling. Everything within 1 m radius was affected. I had to clean up!!:mad: Worse still, i was hit by it in the right eye!!!!

I will never ever try to get carbon rods from batteries.

Regarding the porous pot for use as diaphrahm in electrolysis, can i use multiple filter papers as a diaphramn?? Other than using ceramic pots.

Esplosivo - 9-4-2004 at 01:17

After running the cell for so long I have also got the same brownish tinge at the anolyte. Analysis ahead.

Thanks for the info. hodges. Will reduce the voltage next time, probably using some small resistors in parallel to the electrolysis set-up.

Could copper sulphate be used instead of magnesium sulphate. Copper sulphate is much more available and much more cheap (in my case at least). It contains only some impurities such as sodium sulphate and chloride in small quantities.

darkflame89, the black substance in the can is mostly MnO2. It is very difficult to clean. Next time try pulling the carbon rod out slowly, rotating the rod slowly as you pull it out. It only requires some patience :P As half the rod is out, the rest of the rod can be easily pulled out. The metal cap can be filed out.

[Edited on 9-4-2004 by Esplosivo]

Saerynide - 9-4-2004 at 05:30

OMFG... Im so sorry that you got sprayed right in the eye! Are you ok?? :o I feel so bad cause I suggested you to get them from batteries :(

Pressure mustve built up in that battery, like how faulty batteries explode for no reason.

Wow, youve gotten the brown tinge too? :o So Im not alone... I thought MgSO4 just didnt like me.

Btw, Organikum, what do you mean by H2SO4 in the cathode side?? Isnt it on the anode? :o

Edit: Someone said something about how sulfates can be oxidized by electrolysis. Maybe electrolysing it for so long changes it, so its not contamination?? :o

[Edited on 9-4-2004 by Saerynide]

Marvin - 9-4-2004 at 05:38


You learned the wrong lession. Wear eye protection when doing chemistry.

darkflame89 - 10-4-2004 at 00:19

Yah, learning that fire is dangerous by burning the hand is a hard way to learn.:( I going to have glasses on next time anyway. There must have been a pressure built up. But it is a new battery!!! Its a Energizer battery, the battery cap is not directly connected to the carbon rods. Therefore I still don't have the rods.

So now i have got to wait until after the mid year exams to buy some graphite drawing pencil and somehow get the leads(I have no idea how:( )

About the brown coloration, maybe it is because you all left it out for too long?

Saerynide - 10-4-2004 at 00:52

Dont use energizer! Those are alkaline. Get those el cheapo ones like Eveready (S$ 2.50 for 2 - the ones in my guide pics) or the S$1.20 for 2 ones (the red ones with silver rings/stripes).

Edit: I also need to learn to use eye protection. I once had a container of acetone I was distilling shatter and splash me right in the face :mad:

[Edited on 10-4-2004 by Saerynide]

Saerynide - 10-4-2004 at 03:36

Look at this:

All four sources say that H2SO4 can be brown, but depending on its purity. However, not a single source (Ive searched for a LONG time) *defines* "depending on purity". So, do they mean concentration?? Contaminants?? And if so, what *kind*??

So, maybe we do have H2SO4, but just the brown kind :o

Btw Esplosivo, since youre the analysis expert, tell me how the analysis goes. Im very eager to know your results :D

Esplosivo - 10-4-2004 at 04:04

This dark brown colouration of sulphuric acid is new to me. I have 98.9% GPR sulphuric acid at home and it is a colourless viscous liquid. Could any one provide more information about what this colouration of sulphuric acid is due to please.

Analysis is going to be carried, but not soon, since I'm am quite busy lately :P Currently I am also running the same set-up but using copper (II) sulphate. After 24 hrs the cathode was completely plated by copper and the anolyte gained a sort of greenish blue colour. I will let it run for an other couple of hours, then we'll see.

Saerynide - 10-4-2004 at 04:47

Finally! Something about the contaminants!


Clear, colourless, odourless, dense, oily liquid when pure; yellow to dark
brown when impure.

Sulfuric acid is sold or used commercially in different concentrations,
including technical (78 to 93%) and other grades (96, 98-99, and 100%).
Impurities include metals such as iron, copper, zinc, arsenic, lead,
mercury and selenium, sulfurous acid (as SO2), nitrates and


Organikum - 10-4-2004 at 04:57

The brown color comes from oxidized organic matter - simply burned stuff. It is not related to concentration but to purity - high purity H2SO4 is clear. But I know of no use of H2SO4 (except for analytical purposes) where the decolorization matters in any way. Industrial/technical grade H2SO4 is utmost always brownish to almost black, I guess this depends mostly on the containers in which it was transported.
Battery acid is even more contaminated also it looks clear - on concentrating this stuff I alsways had a yellow precipitate which had to be filtered out and the concentrated H2SO4 had a yellow tint whci I disliked. (looked like piss).

Saerynide - 10-4-2004 at 05:01

YESSSSS!!!!!! That means we've had H2SO4 all this time!!!! :D :o

Why didnt you tell us this earlier?? All this time Ive been trashing it thinking it was crap :mad: :(

All the money I wasted on MgSO4 just I was kept on starting over!! :(

Edit: Ooooooohhh!!! Maybe the organic crap is from the paper towel salt bridge!! I do realise half of my salt bridge goes brownish while the oher half stays white :o

[Edited on 10-4-2004 by Saerynide]

darkflame89 - 11-4-2004 at 00:33

So the brown coloration is nothing??!!yay!!:D

OK, got a new idea here, using the same setup, but instead of carbon electrodes ,we could use aluminium ones!!
Thats because, aluminium sulphate will be formed in the anolyte. And this hydrolyses in water to give you insoluble gelatinous white solid as Al(OH)3 and SULPHURIC ACID!!!:D The products are still the same at the cathode side.

This means as long as i can get rid of the aluminium hydroxide, i can get sulphuric acid!!!:D

And, that means i don't really need to use carbon electrodes!! And no, spraying of MnO2 paste into my eye!!:D

In fact, FeCl3, hydrolyses too to HCl and Fe(OH)3.

I got this from a book called Chemistry potpourri. Singapore hardly sells books like this. I got from Science Centre here.:)

Organikum - 11-4-2004 at 01:53

You wont get happy with Al electrodes, at least not in the H2SO4 compartment.

Esplosivo - 11-4-2004 at 04:07

The process is taking too long to produce a small quantity of acid. The carbon electrodes, which are correded at high voltages must be replaced, but with what?! If the carbon rods could be replaced by some other inert material (which is not gold and neither platinum :P) the process could be scaled up by increasing the voltage right? Does anyone have any suggestions.
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