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Author: Subject: PbO2
axehandle
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[*] posted on 3-6-2004 at 22:39
Ceramic substrate PbO2 electrode construction


This really belongs in Technochemistry, but I didn't want to create a new topic.

I found an (once you sift through the dribble) easy method to construct a ceramic substrate PbO2 anode in this patent (US 4,008,144):
4,008,144

[Edited on 18-8-2004 by vulture]




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[*] posted on 27-7-2004 at 12:26


Finally I have obtained 1kg of copper carbonate, which I've converted to copper acetate with acetic acid. The MSDS for cupric acetate states that skin irritation can be caused by skin contact. I just thought I should mention to anyone trying this out that skin irritation is just the first name. It, in fact, itches like nobody's business! I'm going crazy, it itches like hell! Arrgghhh!!! (I know I shouldn't have gotten it on my skin, but being the slobbering idiot I am.....)

[Edited on 2004-7-28 by axehandle]




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[*] posted on 27-7-2004 at 13:01


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 :P

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.




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[*] posted on 28-7-2004 at 11:30


isn't PbO2 that dark brown powder inside all these lead acid batteries I have been taking apart ?

I pretty sure it is.
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[*] posted on 28-7-2004 at 11:33


Yes it is. It might contain some lead sulphate though. Anyway that is the stuff.



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[*] posted on 2-8-2004 at 18:24


Quote:

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.


I sure will. Right now I have a vat containing 10 litres of water, 500g of copper carbonate converted to cupric acetate, and a huge chunk of lead. The displacement is slow going, but it seems to work. I expect it to take a few weeks though.

Edit: (next day) The lead lump was half eaten through! Just dumped in a freshly cast ingot. This is faster than I thought.


[Edited on 2004-8-3 by axehandle]




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[*] posted on 2-8-2004 at 21:58
Re: Pb compounds


I have several kilograms of metallic Pb lying around my place, in the form of old Pb pipes and roof flashing that was removed when I did some renovations and alterations, along with fishing sinkers. Now I know what to do with it!

BTW: PbO2 and soluble plumbates e.g. Na2PbO3 are very powerful oxidants, capable of oxidizing Mn(II) to MnO4- in alkaline solution, and probably also Fe(II) to FeO4--. MnO2 is much weaker.

Plumbate, and also bismuthate(V), are in fact used in a colorimetric method for determining Mn in water, by oxidation to MnO4-, the absorbance of which in about the middle of the visible spectrum is then measured. However, it is interfered with by the presence of Fe, which is oxidized at the same time to FeO4--, which has substantial absorbance in the same part of the visible spectrum.

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[*] posted on 13-8-2004 at 12:38


There, the Pb(Ac)2 via copper carbonate --> copper acatate --> lead acetate route was successful. I'm right now boiling in the 10 litres of water it was done in --- very boring.

What I noted was that the lead chunks in the vat were eaten up at a very satisfactory rate: over 1kg of solid lead was eaten up in 3 days! The displaced copper doesn't stick, BTW, it falls to the bottom in small pieces; these were very easy to remove using a simple coffee filter.

The copper carbonate as obtained from my pottery supplier wasn't completely pure. Some insoluable junk, white to white-blue in color, it left at the bottom of the vat. Hasn't interfered with the process though AFAIK.




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[*] posted on 13-8-2004 at 15:26


I think that the best way of obtaining Pb(IV) for use as a reagent in solution is by firstly dissolving PbO (litharge, obtainable from plumbous salts or heating finely divided metallic Pb in air), or else metallic Pb itself (with the evolution of hydrogen) in concentrated excess NaOH or KOH to form plumbite, PbO2--; then electrolysis of an alkaline solution of this at an appropriate voltage to obtain a solution of plumbate, PbO3--. Na or K bismuthate(V) is is obtained in a similar manner. Acidification of the solution at a low temperature with a strong mineral acid would precipitate out PbO2.

It, and bismuthate, is used in colorimetric chemical analysis for Mn by oxidizing the latter to MnO4-, although any Fe present is also oxidized to FeO4-- and interferes.

A commercial use of Na or K plumbate solution is to obtain Ca or Zn plumbates, which are used as corrosion-preventing metal paint primers especially for galvanized iron, as precipitates by adding Ca or Zn salts.

I am rather surprised at the posts describing Pb dissolving in non-oxidizing acids like HCl, which should not happen in the absence of a strong oxidant because Pb is below H in the electrochemical series. Perhaps it was not pure Pb, but solder? - the Sn content in it would dissolve in HCl.

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[*] posted on 13-8-2004 at 16:20


No... Pb is ABOVE H in the electrochemical series, at least in my (handwritten) variant:
Li, K, Ba, Ca, Na, Mg, Al, Mn, Zn, Cr, Fe, Co,
Ni, Sn, Pb, H, Cu, Hg, Ag, Au, Pt

And Pb(Ac)2 is a lead(II) salt, not a lead(IV) salt....

And Pb doesn't dissolve in NaOH, I tried it 5 weeks ago.

Your post is making me confused, but it could be you (or me) being drunk. I know I am!


[Edited on 2004-8-14 by axehandle]




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[*] posted on 13-8-2004 at 17:18


Indeed, Pb does have a negative standard potential. Yet it doesnt dissolve in dilute acids (except HNO3), not even HF < 60%. It dissolves in hot conc. H2SO4, forming a complex. But it does dissolve i.e. in HAc, in the presence of oxygen (slowly of course). Or HCl - although even slowlier because PbCl2 is rel. insoluble.

Regarding your plumbate route - sure, that may well work - dissolution of any PbII salt in NaOH to form the [Pb(OH)4]2- ion, then electrolysis to [Pb(OH)6]2-. This in turn has then to be converted to PbO2, presumably with an acid. Surely not as simple as simply adding alkaline NaOCl to PbII salt.

Burning fine Pb - how realistic is that? HNO3- always the hassle with NOx and wasting HNO3. The copper acetate route sounds still best to me!

In case you misunderstood this JohnWW, the issue was about PbO2, not soluble PbIV compounds!




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[*] posted on 14-8-2004 at 01:21


Axehandle - where did you get your electrochemical series from, viz.:
Li, K, Ba, Ca, Na, Mg, Al, Mn, Zn, Cr, Fe, Co, Ni, Sn, Pb, H, Cu, Hg, Ag, Au, Pt ?

Some versions on the internet which I found by searching on Google for "electrochemical series" (using also -amazon -buy -purchase -payment -price as negative search parameters to filter out book advertisements) put Pb just above H (by -0.13 volt), while others put Pb just below H.

In any case, even if the very small negative potential for Pb -> Pb++ over H of -0.13v is correct, this small potential, plus the poor solubility of PbCl2, would result in the reaction of Pb with HCl being initially slow, and then arrested after a short time due to the layer of PbCl2, except at high temperatures with an excess of HCl. Similarly, the solubility of PbSO4 in aqueous solution is even less, being infinitesimal even in boiling water, which explains why the H2SO4 in lead-acid batteries does not simply dissolve away all the Pb.

(BTW They agree that Cr (-0.74v) and Ni (-0.24v) are above H, which means that the corrosion resistance of stainless steels, and of "hastelloys" containing principally Cr and Ni, is due entirely to the formation of impervious oxide layers.)

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[*] posted on 17-8-2004 at 14:49


Quote:

Axehandle - where did you get your electrochemical series from, viz.:
Li, K, Ba, Ca, Na, Mg, Al, Mn, Zn, Cr, Fe, Co, Ni, Sn, Pb, H, Cu, Hg, Ag, Au, Pt ?

A Swedish college chemistry testbook printed in 1960. I'm fairly sure my notes are correct, I usually double check when copying. Can't vouch for the accuracy of the book though...

I'm ready to make PbO2 from the acetate now using NaOH + NaOCl (or KMnO4?), just wanted to ask those who have done this: The PbO2 precipitate formed, does it sink to the bottom or float? (This is crucial for me to know...)




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[*] posted on 18-8-2004 at 03:12


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 :P



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[*] posted on 18-8-2004 at 06:36


That was what I was afraid of. I suppose adding only a stochiometrically balanced amount of oxidizer, and then separating the precipitated PbO2 by sedimentation would be best then...

I'll see whether KMnO4 can oxidize Pb(OH)2 to PbO2, it would have the advantage of being its own colour indicator...


[Edited on 2004-8-18 by axehandle]




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[*] posted on 18-8-2004 at 08:58


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 :P

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]




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[*] posted on 18-8-2004 at 09:40


I oxidized my Pb acetate in acetic acid with Ca(OCl)2. This formed a fairly easy-to-filter precipitate of PbO2, perhaps because I was using a concentrated, solid oxidizer instead of a dilute aqueous solution. The calcium waste products dissolved in the acid. After the first filtration, I washed the precipitate with water several times before letting it dry.



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[*] posted on 18-8-2004 at 09:45


Thank you both for the suggestion of washing the mix repeatedly. Stupidly enough that was something I didn't even consider -- just goes to show that I lack lab skill (or perhaps that I'm an idiot, I sure feel like one...)....



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[*] posted on 18-8-2004 at 10:09


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.

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[*] posted on 18-8-2004 at 12:16


I applied this oxidizer to a solution of lead acetate in aqueous acetic acid, so the system was never alkaline. I obtained a heavy, dark brown powder.



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[*] posted on 18-8-2004 at 12:39


Check your books, John. McAlpine and Soule:
"PbO2...is formed...(b) by treating Pb(II) in alkaline soln with oxidation agents, as halogens, H2O2, KMnO4, NaOCl, etc."

EDIT: Hawley's: "PbO2...Derivation: By adding bleaching powder to an alkaline solution of lead hydroxide."

Merck: "PbO2...Lab prepn from Pb(II) acetate and Ca(OCl)2: Newell, Maxson, Inorg Syn 1, 45 (1939)"

[Edited on 18-8-2004 by S.C. Wack]
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[*] posted on 19-8-2004 at 15:54


What about mixing with molten KNO3? Don't you end up with PbO and KNO2 (also useful)?
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[*] posted on 19-8-2004 at 16:42


Quote:
What about mixing


Mixing what with nitrate?

I'll guess: PbO + NaNO3 -> PbO2 + NaNO2
However: 3 PbO + NaNO3 -> NaNO2 + Pb3O4

and:
3 Pb + 4 NaNO3 = Pb3O4 + 4 NaNO2
Pb3O4 + 4 HNO3 = 2 Pb(NO3)2 + PbO2 + 2 H2O

AFAIK, that 1st equation is better for PbO2 than nitrite.
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[*] posted on 19-8-2004 at 16:48


No, he meant Pb + KNO3 --heat--> PbO + KNO2.
The KNO2 is soluble in water, while PbO isn't.
It works, but it isn't a terribly economic method. Plus I wouldnt want to play around tooo much with molten KNO3 - unless of cause I was to prepare KNO2 - for which there aren't many methods that are easier than this.

Edit: I am not entirely sure about whehter Pb3O4 is produced, if Pb is added to molten KNO3. At least I haven't heard of it. I guess, as usual it's a question of conditions.

[Edited on 20-8-2004 by chemoleo]




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[*] posted on 19-8-2004 at 18:01


Quote:
Originally posted by S.C. Wack
Check your books, John. McAlpine and Soule:
"PbO2...is formed...(b) by treating Pb(II) in alkaline soln with oxidation agents, as halogens, H2O2, KMnO4, NaOCl, etc."

EDIT: Hawley's: "PbO2...Derivation: By adding bleaching powder to an alkaline solution of lead hydroxide."

Merck: "PbO2...Lab prepn from Pb(II) acetate and Ca(OCl)2: Newell, Maxson, Inorg Syn 1, 45 (1939)"

[Edited on 18-8-2004 by S.C. Wack]


McAlpine and Soule are definitely WRONG about KMnO4 oxidizing Pb(II) to PbO2 in alkaline solution. As for H2O2, I suspect they are wrong again, because even small amounts of heavy metals merely catalyse the autodecomposition of H2O2 with the evolution of oxygen. At best, it would take an enormous excess of H2O2. However, NaOCl and dissolved Cl2 can do it, though.

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