PbO2

Hmm maybe NaOCl wouldnt be the best. It normally contains NaCl (obtained via bubbling Cl2 into NaOH), plus it is normally basic.
Not only may your precipitate contain PbCl2, but especially Pb(OH)2 . Better to use Ca(OCl2) which is free of base. On the other hand, a great excess of NaOCl would probably do the job.

For dissolving lead to get some lead salt, without using nitric acid etc, you could go via the copper acetate route - copper is deposited, and lead acetate is formed. you cant use copper sulphate as the corresponding lead salt is insoluble. to make copper acetate, either dissolve CuCO3 in acetic acid, or make it with Calciumacetate/copper sulphate, filter off the calcium sulphate and you are left with copper acetate.
Neat method, made 1.5 kg worth of lead acetate that way, with minimal effort

[Edited on 28-4-2004 by chemoleo]

Please spoonfeed me.

I guess I can spoon feed, but only because you posted this in the reagents and apparatus aquisition area and not general chemistry

un-be-fucking-lievable

Sucking up.

Well axehandle, there is nothing stopping you! It's a nice and satisfying experiment. I really really hope you do not claim that you can't find lead sheets/airgun bullets/fishingweights/lead solder anywhere in Sweden!


PS H2O2 does not necessarily work, I think Bromic only thought it might work too!

[Edited on 28-4-2004 by chemoleo]

another method

I tried it - success :)

I happened to have some lead carbonate-acetate at hand - so rather than just looking at it, I used it!

1. I dissolved exactly 10 grams in an excess of acetic acid, and filtered off to get rid of a slight turbidness.

2. To the clear solution NaOH(aq) was added. Once neutralised, a fine pigment white precipitate (Pb(OH)2) was produced, which, upon adding more NaOH turned yellow. This was reversible however, I acidified with acetic acid once again, and everything went back into solution. Then I adjusted the pH to pH 10 with NaOH, and collected the white precipitate.

3. The precipitate was resuspended in H2O dest, and the suspension was halfed.

4. To one half an excess of 30% H2O2 was added. The precipitate turned a pale lemmon yellow, with a few bubbles appearing. Thinking, maybe the reaction is not complete, I started heating it. Now that was a bad idea. At still low temps, a runaway suddenly occurred, and HUGE amounts of gas/smoke was produced - which thankfully cleared rapidly, indicating it was mostly steam and oxygen (for fear of lead in the air I departed the scene for a few minutes, and let fresh air do its job)... the yellow precipitate did not change colour, after the runaway, either.

5. To the second solution I added a large excess of NaOCl (with NaCl). Immediately a rust brown precipitate formed. No gas evolved this time, no chlorine evolution, thankfully. No accidents this time.

Conclusions?

Well I checked some MSDS's:

PbO2: Appearance: Dark brown powder. Incompatibilities: Aluminum carbide, sulfides, hydrogen peroxide, hydroxylamine, combustible and organic materials.

Pb3O4: Appearance: odourless red or orange powder

PbO: Appearance: yellow-orange powder. Stable. Reacts violently with hydrogen peroxide, strong oxidizing agents, aluminium, zirconium, halogens, sulphur trioxide, boron, silicon, sodium, zinc.


Well well the conclusion is obvious lol...
NaOCl did indeed produce PbO2, seemingly clean (no PbCl2 precipitate, I would have noticed the crystal needles). I won't mix it with H2O2, thank you very much

Now the other reaction, Pb(OH)2 is not as conclusive. Obviously something happened to it, else the colour change wouldnt have occurred. But it's definitely not PbO2. I guess most likely it's PbO.. or some modification of Pb3O4. I know it doesnt agree with the MSDS description, but I seem to remember a commercial batch of Pb3O4, which had this lemmony pale colour, too.

Anyway - axehandle, there you have it. No more sucking up to get some H2O2 pls... you dont need it!

PS it's currently drying, hoping to weigh it later.

chemoleo,

If anyone was worried...

Looky, looky.

I had 5 litres of 30% HCl and decided to make some PbCl2 just because possession of lead salts is illegal here.

In this picture you can barely make out 1284g of pure Pb dissolving in 1500g of HCl(aq). It's really, REALLY evolving H2!

Yuck!

Polverone is right. Make this outside.. even with open windows HCl may destroy some sencitive electronics (some circuits may have very thin leads and locally no protection against nearby air..).. and not only electronics..
Oh, and small sprinkles from H2 evolution may carry away some lead salts in the air

PbO2 production is scary..
Btw, how about taking it from car battery? Although PbSO4 should be removed somehow..

There's much more joy in making it yourself!
Read above, easy methods are all there!
Still, axe... you are gonna get lots of PbCl2 which is not well soluble. What are u going to do once the lead is covered with it, and the solution is saturated?
Well I guess as long as you can prevent the PbCl2 coating, its fine... as you have to process your your PbCl2 with NaOH anyhow. And that should work regardless... providing you work with weighed stoichiometric amounts, and you stir it throroughly and for a good while ...

PS yes I would seriously watch it with HCl fumes. My dad a few years ago nearly killed me when he found all his tools were rusty

[Edited on 17-5-2004 by chemoleo]

Make sure (for a change ) you dont inhale the fumes coming off it. They really ARE nasty!
Plus, be careful not to explode the NH4NO3. That is even NASTIER!

Ah, yes, thanks for the concern. I wasn't intending to heat a mixture of NH4NO3 and Pb DRY though.... (mixing it with diesel oil or perhaps water will be much safer...

Ceramic substrate PbO2 electrode construction

Pay attention when using such salts. Lead acetate is worse since it is readily absorbed through the skin, and I suppose none of us want an increased lead concentration in the blood right

Well anyway, I just wanted to ask you to document the separation of the formed PbO2 when you have oxidised the Pb acetate. Chemoleo told me he used centrifugation, but I don't possess a large enough centrifuge to seperate resonable quantities of ppt. Filtration prooved useless, since most of the PbO2 simply passed through the filter (but the reson surely is that I used a cheap coffee filter). Suggestions are welcome.

Re: Pb compounds

It first forms as a sort of 'suspension' but after some time it settels on the bottom. I found it very difficult to extarct this ppt from the liquid, filtering only resulted in the collection of a very small quantity of PbO2 as I said previously in this thread. Wish you all the luck you will need

Well, in principle it is quite simple nonetheless.
PbO2 is bound to have a high density, so it will settle rather quickly. Definitely faster than CaSO4, which is equallly a pain to filter.
So, I would mix/heat your solution with NaOCl, like in a large batch.
Then, using a high glass/ beaker, let teh products settle over night.
Decant.
Repeat this 2-3 times. Then boil the remaining water off, and put on a radiator to get rid off the last bits of water. You could resuspend this in H2O (to get rid of traces of NaCl and such), and filter. I am sure the filtering wont be a problem then, the PbO2 clumps up and won't go into suspension like before.

Alternatively I suggest a centrifuge ... then the waiting/decanting wont be necessary several times, but prob. only once

Edit: As to the KMnO4 - I guess you ought to keep it basic, too. I am not sure whether you might get a precipiate of MnO2, so this might be mistaken (or mix) for PbO2.

[Edited on 18-8-2004 by chemoleo]

KMnO4 cannot oxidize Pb(II) to Pb(IV), at least not under neutral or alkaline conditions. Pb(IV) is usually the stronger oxidant. In fact, a colorimetric method of analysis for Mn in water involves oxidizing it to MnO4- with Na or K plumbate(IV) in alkaline solution. (It also oxidizes Fe to FeO4--, which unfortunately interferes with the analysis).

As for oxidizing Pb(II) to a precipitate of PbO2 with Ca(OCl)2 ("chloride of lime", which gives an alkaline solution: are you quite sure that you did not end up with calcium plumbate, CaPbO3, instead? If the precipitate was light gray in color (which is the color of CaPbO3 primer paint for galvanized iron), this would be indicated.

John W.

I don't know if anyone believed this or not, but I'll rant just in case. I suggested that he check the literature, and he calls the literature BS. I checked the literature and found journal refs to PbO2 from:
alkaline Pb acetate and H2O2
alkaline Pb nitrate and H2O2 (quite pure, btw)
alkaline PbO and KMnO4
But one of Rhodium's funnier lines applies here: "However, to obscure their discovery as much as possible for the rest of the world, they encrypted their findings by publishing it in french (PGP was not available at the time)." And if not French, German. Unfortunately the volume of Mellor that covers PbO2 (vol 7, pp. 681-687) is the only one that the library doesn't own, and his "Modern Inorganic Chemistry" is only so helpful here.

Authors such as McAlpine and Soule (the excellent "Prescott and Johnson's Qualitative Chemical Analysis", the writers for Mellor - and farther back, Thorpe, Roscoe and Schorlemmer, and Watts - go by and reference actual papers published in respected journals. When something is questionable they find a way to say so. They state facts that are known, not make things up.

Those who disagree should publish their own book to counter the BS artists such as Inorganic Syntheses: "Hypochlorite is the most commonly used oxidant, but permanganate, sodium peroxide, ferricyanide, and many other oxidizing agents can be used." However, in order to get published, one would have to do better than mumble something about oxidation of Mn(II), something which happens in acid, and in any case I don't see how it applies here. So, elaborate if you can, please.

OK then - first I dissolved some Pb in dilute HNO3. Used my high-tech sublimation apparatus for this - a pint beer glass sealed with a RBF, and a "warm"plate. But the last of the lead was slow to dissolve. I got impatient so I converted it all to sulfate and carbonate. I wanted the sulfate to test the easiness of conversion to PbO by fusion with Na2CO3, and will try something like US3644090.

The carbonate was converted by heat to a dull flowerpot-colored oxide, and some of this was heated more strongly to a bright dark orange oxide that was red-violet when hot. Heating the bright oxide in vinegar gave PbO2 as a dark brown precipitate and half the Pb as soluble acetate, which was precipitated as the sulfate with H2SO4.

The experiment with the lower oxide and H2O2 was first tried at perhaps 50C with a 25% excess of conc. NaOH. Dropwise addition of unheated H2O2 in concentrations between 10 and 30% gave an immediate formation of brown sludge at the site of reaction, but disappeared no matter how much more peroxide was added and at what rate. The H2O2 was made by evaporation of grocery store 3% and had a pH of 3.38. Should have added some base to it before addition.

Repeating this at nearly boiling temps was dramatically different, and dark brown PbO2 formed and remained easily. Not much H2O2 was required. HCl was added with the formation of white PbCl2 and Cl2.

The experiment with KMnO4 was done next, also with a 25% excess of oxidant and base. This was done at boiling, and an unheated satd soln of KMnO4 was added dropwise. There certainly was no "autodecomposition" of the KMnO4, and the excess was decanted off and the precipitate washed many times. It gave the sensitive color test for Mn, by fusion with NaOH and KNO3 in a porcelain crucible. It gave Cl2 and white precipitate with HCl. MnCl2 is quite soluble, and the filtrate was chlorine green-yellow.

1st batch of PbO2

This was made from:

1) mixing bleach with a lump of PbAc2

2) letting it react at a 75 degrees C oven for an hour

3) letting it sit overnight, then filtering through a coffee filter

4) letting the filter dry, then scraping it out

No purification has been done, but I'm now confident that household bleach (mix of NaOCl and NaOH) <i>can</i> indeed be used in this process.

EDIT: I know the camera makes it look black, but it's really dark brown -- camera does bad job with big light contrasts and bad lighting.

I managed to get hold of 5 litres of bleach today (some sort of outdoor furniture cleaner / algae killer) and will begin preparing a proper PbO<SUB>2</SUB> batch shortly.

This is fun!

[Edited on 2004-9-10 by axehandle]

Well, it's an amazing synth! No suspicious things needed, only pottery pigment, metallic lead, vinegar and bleach!

Possession of lead salts/compounds here basically carries a caning sentence so they are impossible to get, for any price.

However....:

1) Copper carbonate: cheap
2) Vinegar: cheap
3) Lead: almost free
4) Bleach: cheap
5) Look on the face of government chemical dept.: priceless

Guess I should send them some

[Edited on 2004-9-10 by axehandle]

EDIT: I'd really like to do this using a centrifuge -- at this rate the sedimentation will take days done properly sigh...


[Edited on 2004-9-11 by axehandle]

Illegal?

VERY weird

Can anyone explain this:

(I've done the "let sink down to half the height then pour off water then add water" cycle about 10 times without this happening!!!)


[Edited on 2004-9-17 by axehandle]

PbO2

Alternative route PbAc2 --&raquo; PbO2

I just conducted a small electrochemical experiment:

Approximately 10g of PbAc<SUB>2</SUB> was dissolved in 300 ml of hot tap water. Two titanium strips approximately 6x100x1 mm respectively were submersed in the solution, approximately 7cm apart and connected to the +5V and GND lines of a 300W ATX PSU. The cell was left running for 5 minutes.

Result: Metallic lead was formed at the cathode in needle-like protrusions which fell off when the cathode was disturbed; the anode was covered with a dark brown substance which I will assume being PbO<SUB>2</SUB>. It was easy to scrape off with a knife, proving that is doesn't adhere to titanium (which is obviously an advantage in this synthesis).

Conclusion: This is a workable alternate route to PbO<SUB>2</SUB> from PbAc<SUB>2</SUB>.





And here, observe the needle-like solid lead:


[Edited on 2004-9-19 by axehandle]

I shouldn't have evaporated the water on a hotplate -- damned PbO<SUB>2</SUB> decomposed partly to Pb<SUB>3</SUB>O<SUB>4</SUB> (?)...

So now I have to do it all over again. This time I'm getting <i>pure</i> NaOCl though, not that soap-ridden crap.

But I <i>did</i> manage to coax 152g of PbO<SUB>2</SUB> from a new (as in never acid-filled) motorcycle battery, so <b>anode, here I come!</b>

Axehandle, I've been tinkering a fair bit with this Pb(Ac)2 method.

For my test anode, I cut a strip of pvc from some pipe (got heaps of it), cleaned it with pvc primer and coated the convex side in everyday pvc glue. Then I just pressed/rolled it in the PbO2 powder. It sticks very well, although I didn't use enough glue and it's a bit patchy. But it does conduct.


As for the plating solution, I'm not sure how much luck you'll have that way. I certainly haven't been successful.
At http://www.geocities.com/CapeCanaveral/Campus/5361/chlorate/... he says,

Interesting that vinegar is about 0.5 M, hey So I've spent the last couple of weeks making some solid Pb(Ac)2. You can't evaporate the solution too quick as it decomposes at 100 deg C.

Hope this helps.
janger

[Edited on 20-9-2004 by janger]

My steaming vat of poison

Not steaming, actually, but it sounded more dramatic that way. Here's a picture of my 2nd run, making PbAc<SUB>2</SUB>.

The thing under the surface to the left is a sheet of Pb, the ring in the back (and at the front, although you can see only the top) was a disc of lead I cast in an old cooking pot lid. The circuference was thickest, and only a ring remains.

The black device in the middle is a submersible aquarium water pump to improve the circulation.

This run is complete, the liquid should now be almost pure aqueous solution of PbAc<SUB>2</SUB>.

Look at all copper fragment sludge at the bottom --- yuck!

EDIT: More info... total volume is close to 10 litres, I used 500g of copper carbonate and the stochiometrically necessary amount of HAc times ~1.5, don't remember the exact amount... Right now I'm filtering the liquid through carbon, then it's boiling time.


[Edited on 2004-9-22 by axehandle]

The picture I <i>intended</i> to attach, instead of that horrible closeup:

And here's the result of the very first purification by recrystallization step, 228g of lovely sweet PbAc<SUB>2</SUB>!

I'm boiling more, this is SO fun!!!

EDIT: 521g so far...

EDIT2: 890g, and still more than half the liquid left to process...


[Edited on 2004-9-22 by axehandle]

I don't think any lead salts except straight organics (with a carbon-lead bond) are absorbable though the skin to an appreciable extent. Additionally calcium uptake is preferred over lead for incorporation into the skeletal matter and muscles. As long as you are not deficient in this, you should at least have an edge, take a calcium supplement or drink lots of milk

Plastic

Try dissolving it in HCl instead. I'd bet a euro or two that this will work. Acetic acid may just not be aggressive enough, CuCO3 is fairly stable and is used in paints, hence i assume if somewhat crystallised, or hydroxy carbonates (double salt) it is more active. It won't help you directly with your problem, but at least you know that you got the Cu carbonate all right. Well you could always preciptate from the CuCl2 solution again, with carbonate or hydroxide, and dissolve the now *activated* hydroxide or carbonate in HAc. Bit of a waste of reagents, but they are all cheap

Did you try boiling the carbonate in HAc?
Really though you ought to use pure HAc (obtainable here as 25% concentrate), you are introducing all sorts of impurities here, including other organic acids!


On a separate note, I remember that CuO wouldnt dissolve that greatly in HAc, and NiO takes a long time in strong HNO3 to dissolve. I reckon it's just due to the relative chemical stability of the (oxy) carbonates.

Re. the white powder - it could be Cu(1+).
It can form from Cu(2+) + Cu --> 2 Cu(1+). It should disappear w. time. Otherwise it's impurities, in either the lead or your reagents, or, as you say, some hydroxy-acetate precipitate. My commercial batch is insoluble too - which is a big pain btw, pure lead acetate crystals, as nice and heave as they are, become first turbid, then white on the surface. I couldnt find a to keep them shiny.


PS For conventions sake, acetic acid is CH3COOH - carboxylic acids R-COOH in general... makes it easier to read


[Edited on 8-6-2005 by chemoleo]

Cyrus: When I performed this procedure, I used 24% HAc, and I had to stir the mix of HAc and CuCO<SUB>3</SUB> for quite some time for all of the carbonate to be converted to acetate / dissolve. Lots of CO<SUB>2</SUB> was produced. Perhaps it's your low HAc concentration?

Also, you don't have to basify the Pb(Ac)<SUB>2</SUB> solution with NaOH prior to mixing it with bleach -- the bleach contains more than enough NaOH to do the job. The same seemed to be the case with bleaching powder (Ca(OCl)<SUB>2</SUB>, although the base there obviously is Ca(OH)<SUB>2</SUB>.

Chemoleo, I probably should have used nice conc. acetic acid, but I don't have any. I'll have to get some, any suggestions?

I added a whole bunch of lead, and after 2 days the blue color is almost all gone. (I stirred a few times a day) But there is now a lot of very fine white powder floating around when I stir it. (The pH is between 3 and 4) I guess I'll just wait around and see if it goes away, or I could always filter it out.

Now here's a dumb question; when I've got my lead acetate solution is there any particularly safe way to turn that into lead acetate crystals. (I don't like the idea of boiling it, as when I boiled sodium acetate I could really smell the sodium acetate coming off)
Slow evaporation?

I boiled in about 10dm<SUP>3</SUP> of Pb(Ac)<SUB>2</SUB> solution, and I'm alive...

I boiled it down to a syrupy fluid which I poured into several paper plates (the kind you eat from). I then placed the paper plates on ice, whereas the Pb(Ac)<SUB>2</SUB> crystallized out. Then I dumped the crystals + remaining fluid (discolored by impurities) into a funnel equipped with a coffee filter; the discolored fluid went through the funnel, the Pb(Ac)<SUB>2</SUB> crystals (big lumps of them actually) stayed in it. The result looked pretty pure, and the procedure is repeatable...

Someone probably has something not so nice to say about my lab skills, but my method worked...

[Edited on 2005-6-9 by axehandle]

[Edited on 2005-6-9 by axehandle]

Heh

"I boiled in about 10dm3 of Pb(Ac)2 solution, and I'm alive... "

Yeah, but I want to be alive and SANE...

I was thinking, perhaps it's not even necessary to get lead acetate crystals. One could get PbO2 just by adding bleach. As for using the lead acetate "as is" for electroplating, I don't know if the impurities would hurt anything though.

edit-
~
Impurities in the bath, such as iron or cobalt compounds, can drastically reduce the strength, density, and surface smoothness of lead dioxide deposits
~

This was taken from the link right above this post.


[Edited on 9-6-2005 by Cyrus]

Quote:

Originally posted by chemoleo Well here it goes... the dissolution of lead metal with cheap OTC materials. I copied parts over of what I posted at E&W already: Quote: A while back I managed to solubilse a kilogram of lead with a neat little method. I didnt want to waste HNO3 for the dissolution, neither NH4NO3 etc. Neither did I have any existing lead salts, or lead oxide. But I did have about 20 kg of solid lead metal, plus 10 kg of copper sulphate (CuSO4), obtained from a gardeners store. I got it to work as follows: 1. Calcium carbonate is easily obtained, I neutralised a known amount with acetic acid (conc. vinegar), yielding a solution of calcium acetate. 2. I dissolved an equimolar amount of Copper sulphate CuSO4, and added to it the stoichiometric amount of calcium acetate. What you get is copper acetate and calcium sulphate 3. THe precipitating calcium sulphate is filtered off, then the filtrate is cooled to precipitate more CaSO4, and filtered off again. One is left with a fairly clean solution of copper acetate, which has a lovely dark green colour and tends to crystallise at the surface of the the liquid. 4. This is probably the longest step - to the solution of copper acetate, one adds a large excess of solid lead metal! the finer the lead pieces, the better (faster). Over the course of a week or two, the copper acetate colour (dark green) slowly disappears, and becomes completely colourless. What happened meanwhile is that solid copper deposited as thin sheaths on the solid lead. I found no stirring is needed during those two weeks, the copper deposited anyway. 5. The solid copper is scraped off the remaining lead, the lead pieces are taken out, and the solution is filtered once again to remove copper pieces and any more CaSO4 6. One is left with a clear solution of lead acetate. Made from 100% over the counter materials. Slow evaporation of the water in the solution leads to LARGE crystals of lead acetate, which are very heavy and slowly become milky/powdery on the surface at air! Quote: Follo-up: When the Ca acetate started reacting over the course of a week or two, I noticed that an insoluble yellowish substance accumulated at the side of the reaction vessel. Thinking, oh, this must be because the HAc evaporates (which is, as I found out later, because the double salt of the hydroxid/acetate forms, plus possibly PbO), I added 500 ml of 20% HAc to the 5 litres in the reaction vessel - and promptly, the yellowish stuff disappeared, and it was a clear green (later clear altogehter) solution again. From that point onwards, I always made sure I keep the solution acidic, with excess acetic acid! Well axehandle, there is nothing stopping you! It's a nice and satisfying experiment. I really really hope you do not claim that you can't find lead sheets/airgun bullets/fishingweights/lead solder anywhere in Sweden! PS H2O2 does not necessarily work, I think Bromic only thought it might work too! [Edited on 28-4-2004 by chemoleo]