Sciencemadness Discussion Board - The lead salts preparation thread!

The lead salts preparation thread!

Pages: 1 2 3 4

dann2 - 20-7-2008 at 11:58

Hello,

How long would it take to make (say) 50 grams Monoxide?
I tried this once and it appeared to be very slow BUT my molten Lead was less than 600C.

Sitting and scooping for hours seems very boring!
Wonder if you were to set up a cheap aquarium pump to bubble air through the molten Lead using a glass tube. You could leave the whole lot sitting for a few days.
Care of one's eyes (and other bits of one's anatomy) of the splatter etc would be essential.

Dann2

Picric-A - 20-7-2008 at 13:36

Was your lead glowing red hot? maybe there wasnt enough oxygen.
I normally make less than 50g at a time but if you kept it hot enough it wouldnt take long, it takes me about 5mins to make 30g batches.
lol i have tried that, every time i tried it lead would always clog up the tube as the cold oxygen flowed through it.

Maybe if you blow a fast stream of air at it from a pump or something that wold speed it up for you?
tommorow i think i will make a low lying but wide iron cruicible and see how well that works.

12AX7 - 20-7-2008 at 20:45

Last time I tried that, I got what turned out to be a green granular substance (it was yellow to orange when hot) which, when treated with HCl, gave H2 and not much white stuff. It turns out the lead skin adheres quite nicely to the metal, so you get little beads coated in, I suppose, suboxide.

Tim

Picric-A - 21-7-2008 at 01:30

yea, you didnt heat it strong enough so it didnt seperate fully, hence the granular form
cant explain the green, like u say probs some suboxide

JohnWW - 21-7-2008 at 04:55

Preparations of PbO ("litharge"), (once used as a pigment in paints, now disfavored because of the neuro-toxicity of Pb compounds, and because of the variability of its color), are often light yellow or lime-green, even without any impurities except possibly a trace of Pb3O4 ("red lead", also used as a paint pigment). Higher concentrations of Pb3O4 in PbO would be indicated by an orange color. PbO is also used as a component of yellow or orange pottery glazes, and (being colorless when pure in bulk in siliceous glasses) in concentrations of up to 30 or 40% in lead glass, used for paperweights and in French lead crystal because of its weight.

A quicker preparation, if HNO3 is available, may be to dissolve Pb in HNO3 to obtain Pb(NO3)2, then heating that to drive off nitrogen oxides (caution - use fume hood) to leave PbO.

If carbonate is added to such a solution of Pb(NO3)2, insoluble PbCO3 (cerussite, an important Pb ore), or the hydrous form, 2PbCO3.Pb(OH)2 (hydrocerussite) can be precipitated out, leaving nitrate in solution. These can be used as an alternative to PbO in many reactions, and converted to PbO by heating sufficiently, to about 315 to 400ºC, to drive off CO2. It was the precipitated hydrous form, hydrocerussite, 2PbCO3.Pb(OH)2, that was called "white lead", and formerly used as the white pigment in white and pastel paints, until superseded by TiO2 because of the neuro-toxicity of Pb compounds and the superior opacity of TiO2.

[Edited on 22-7-08 by JohnWW]

Picric-A - 21-7-2008 at 12:14

yes then lead the NO2 fumes through water to regain the nitric used in the production of the Pb(NO3)2. been there done that lol

the one problem is if your desired product is Pb3O4 then you will need to reheat it to get it, so ultimatly its easier starting off with heating lead.
does anyone know of any other method of converting PbO to its higher oxide? eg, via a simple cost effective chemical reaction? thanks,
Picric-A

12AX7 - 21-7-2008 at 12:30

Quote:

Originally posted by Picric-A
yea, you didnt heat it strong enough so it didnt seperate fully, hence the granular form
cant explain the green, like u say probs some suboxide

I don't think it was that. I got similar results ranging from the temperature where rapid oxidation starts (the metal surface appears iridescent) to red heat, all the way up to the melting point of PbO, which tends to dissolve (lead can dissolve a few percent of the oxide) and attack my crucible (steel can) where it isn't in solution.

Quote:

Originally posted by JohnWW
Some preparations of PbO ("litharge"), which is usually white (and once used as a white pigment in paints until TiO2 became standard)

No, lead carbonate (or subcarbonate) was used as pigment, because it was cheaper than TiO2, which is better. PbO is always colored.

Quote:

Originally posted by Picric-A
does anyone know of any other method of converting PbO to its higher oxide? eg, via a simple cost effective chemical reaction? thanks,
Picric-A

Oxidation only works for some modifications, so I'm told. Heat it too far and it recrystallizes into the other form, that won't oxidize.

Even from freshly prepared Pb(OH)2, I've never gotten anything darker than light orange.

Tim

S.C. Wack - 21-7-2008 at 13:51

The Pb3O4 that I made IIRC by heating white PbCO3 to yellow PbO and heating further in air was only red when hot and is quite orange, but not in the same way as say dichromate. A duller, more ochre sort of dark orange. It did not take long convert at all. The color seems to be dependent on preparation as is told in books and can be seen in images from sellers of it. Methods of preparation chemical and otherwise can be found in some books that are available.

kilowatt - 1-8-2008 at 15:42

I tried heating lead to make PbO last night. I used my casting furnace with a crucible holding about 10-15lbs of lead. At first the PbO was coming off granular and mixed with lead, even at moderate red heat, maybe 600-700°C. It was orange-red while even barely hot but turned yellow/black as it cooled down. Later on in the process the lead started smoking, so I turned the burner down to just a low idle, barely visible incandescence on the surface, and liquid oxide was now floating on top. When I scooped it out it hardened to a glassy dark grey appearance. According to every MSDS for PbO the melting point is over 800°C, and I was nowhere near that at that time. Presumably Pb3O4 which has a much lower melting point was forming, or some mixture which has a low melting point. I was getting very little of the oxide on the surface so I put a tube down into the crucible connected to my air compressor with a needle valve and left it bubbling for a few hours, still only getting a very small amount of the liquid oxide.

Should I turn the furnace up to a high heat over 800°C to get better PbO yield? Does Pb3O4 decompose to PbO at higher temperature? I don't like how it smokes and I lose so much oxide this way (not to mention the fumes!!). I may be able to rig some sort of a trap for the smoke if I use the forced air from the tube to oxidize the lead, since it doesn't depend on air getting to the top of the crucible.

Edit: Well I'm running it at higher temperature right now, and it is making stuff, just not quite sure what. Probably PbO and/or Pb3O4. It is going slowly, but a little faster at least. It doesn't smoke too much, but I am still not going near it much because of the fumes. The stuff wets my crucible which is annoying. I am afraid it might slightly dissolve the crucible which is made of silicon carbide, but there doesn't seem to be much damage so far. Here's a pic of the progress.

[Edited on 1-8-2008 by kilowatt]

HPIM1439.JPG - 127kB

S.C. Wack - 2-8-2008 at 00:25

Silicon and PbO do not get along; molten PbO and lead oxidizes metals and reacts with Si cpds. including the carbides, and dissolves the product further. You do not have Pb3O4, 470-480C has been recommeded for its formation; higher is bad and it is completely decomposed below 650C.

In order for the lead to come into good contact with oxygen some sort of agitation is needed.

12AX7 - 2-8-2008 at 01:51

My experience is the SiC crucible itself just slowly burns, even without lead. The glazing is about gone on mine. Without protection, the graphite bonding burns nicely in air at these temperatures.

Tim

not_important - 2-8-2008 at 03:01

Pb3O4 is formed when PbO is heated in air at about 300-350 C, at red heat it decomposes back to PbO.

PbO formed at temperatures below bright red -too low to melt it - will be yellow when cool, if fused it will be more of a red colour; the yellow form is more reactive.

PbO is very easy to reduce. carbon monoxide will do so at a few hundred C. If you are using a fuel-fired furnace, as opposed to an electric one, do the best you can at keeping the flue gases away from the lead.

An iron container is likely the best to use, fused PbO wicks into pores in many ceramics that resist it's solvent action, and most silicate based ceramics are attacked to some extent by PbO.

Litharge usually does have some lead mixed in with it, unless care is taken when making it. You can collect the cold litharge and crush it, then heat in air with stirring at barely visible red heat to oxidise the unreacted lead.

kilowatt - 2-8-2008 at 05:15

Quote:

In order for the lead to come into good contact with oxygen some sort of agitation is needed.

Well I was continually bubbling air through it, I dunno how much more contact/agitation you could get. I will try a lower temp later today with an iron crucible with a lid on it.

Thanks for the advice.

[Edited on 2-8-2008 by kilowatt]

Lead Nitrate

kilowatt - 3-8-2008 at 19:52

A few minutes ago I set up a non-divided DC electrolysis cell with nearly saturated ammonium nitrate and lead electrodes to see if I could get lead nitrate (for nitric acid production) and ammonia. Immediately dark black material, which I would presume to be lead dioxide, started falling away from the anode, with the cathode quickly darkening as well. The cathode is definitely producing a lot of ammonia. Within just a couple minutes the anode was already noticeably worn down; this is a surprisingly fast process and I expect the electrode to be pretty much gone in just a few hours. I'm not sure about all the lead dioxide forming, is this going to want to react to lead nitrate on its own or am I going to have to decompose it and put it back in the electrolyte (which should be enriched with nitric acid)? Also does lead form any ammonia complexes?

Edit: The cathode is now surrounded by a spongy or gelatinous dark gray material. What is this? Some sort of ammonia complex? Or is spongy lead being electro-deposited around the cathode? When heated it converts quite readily to yellow PbO.

2nd Edit: After letting the cell rest for a few minutes the spongy material appears to have dissolved or settled out.

3rd Edit: I have unquestionably made lead nitrate with this setup; when a sample of the electrolyte was heated in a test tube beyond the decomposition of ammonium nitrate, copious amounts of N2O were liberated and yellow PbO was left behind. There is a lot of solid gray material at the bottom of the cell, presumably spongy lead. I will see if any dissolves after sitting overnight. It may need removed and recycled depending on how much or little the ammonium nitrate has been converted to nitric acid through loss of ammonia.

[Edited on 3-8-2008 by kilowatt]

S.C. Wack - 3-8-2008 at 20:08

300-350C sounds crazy low for Pb3O4 from PbO. Nothing that I've read gives such a low temperature. Have a reference?
The temperatures that I gave above were from Mellor.
In this article, where

"not only the effect of temperature upon the rate of oxidation of different starting materials but also the physical and chemical properties of the red lead formed from the different substances under different conditions of temperature have been studied. We have studied red lead made from litharge, lead hydroxide, white lead, lead sponge, metallic lead and lead tartrate.",

after 15 hours heated to 400C in a rotating jar in a furnace swept with air, they say that the litharge was less than half converted.

"The temperature best suited for the formation of red lead varies with the starting material used. About 425-430° is best for white lead, 450-470° for litharge and lead sponge, and about 450°. for converting lead hydrate and metallic lead to red lead. In fact, 450° may be taken as the temperature at which red lead can be economically formed from any suitable starting material."

Attachment: iec_4_867_1912.pdf (1.3MB)
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Taoiseach - 3-8-2008 at 22:01

@kilowatt

Why dont you simply fuse your lead with ammonium nitrate? Just heat the nitrate until it melts, then add the lead in small pieces. The mix will start bubbling vigourously and lots of ammonia is produced. When reaction is over, add hot water to dissolve the lead nitrate, filter and evaporate until crystals start to appear. Cool in the fridge to obtain your lead nitrate. Couldn't be easier.

Note that if your lead contains Sn, the ammonium nitrate will oxidize it to SnO2 appearing as a white precipate when water is added to the fused mass. Decant & filter the liquid to remove these impurities.

kilowatt - 3-8-2008 at 22:17

That would be quite convenient, thanks for the idea. I'm surprised it reacts so quickly and in that manner and will have to try this out on a small scale tomorrow. However molten ammonium nitrate can be fairly dangerous/unstable and I'm not sure I'm so keen on doing 5kg batches of this like I intend to with the electrolysis. After all, there is just a fine line between ammonium nitrate's melting and decomposition temperatures. I'm also not sure that it would be any more energy efficient than the electrolysis, but perhaps. Definitely something to look into.

I already intended to use this process to separate the scrap lead alloy, and just started a new thread with more specifics.

[Edited on 3-8-2008 by kilowatt]

not_important - 4-8-2008 at 05:13

S.C. Wack - from an older industrial chemistry tome, don't remember which. I'll try to find it again tomorrow. I do remember seeming to get lower yields of red lead at 450-500 then at 350-400 but this may have something to do with other factors.

kilowatt - 15-8-2008 at 17:23

Quote:

Why dont you simply fuse your lead with ammonium nitrate? Just heat the nitrate until it melts, then add the lead in small pieces. The mix will start bubbling vigourously and lots of ammonia is produced.

Back regarding this method, as I stated before fusing a large quantity of ammonium nitrate and randomly dropping lead into it initiating a vigorous exothermic reaction seems extremely dangerous. It wouldn't take much overheating from the reaction to send the AN melt into thermal runaway. Another problem is that pure lead nitrate has no melting point and is solid at the temperature of fused ammonium nitrate. Thus when the lead nitrate concentration gets sufficiently high the reaction stops, and much AN is not consumed. In my test runs I had to decompose it off which always results in a considerable report (probably a serious explosion if scaled up) as it nears the end just before nitric fumes come out.

It would be preferable to carry out this reaction in a minimal amount of high boiling solvent for the nitrates. This should be done by adding the fused ammonium nitrate from a heated addition funnel into a RBF containing the solvent and enough lead pellets to complete the reaction with a stirrer, as opposed to adding lead to the full amount of ammonium nitrate. DMSO looked tempting but it is not compatible with strong oxidizers. Perhaps there is a eutectic mixture of sodium/potassium or lithium nitrates that would work (pure lithium nitrate melts at 255°C, still too hot) but I have not been able to find any information or phase diagrams for alkali nitrate mixtures. Perhaps fused alkali nitrates would be reactive with lead anyways, I am not sure. Can anyone think of another solvent that would work here (stable liquid from 160°C to 200°C, compatible and miscible with lead nitrate and ammonium nitrate, unreactive with lead)?

ScienceSquirrel - 15-8-2008 at 17:52

Fused potassium and sodium nitrate react with metallic lead to form the alkali metal nitrite and lead oxides.

kilowatt - 15-8-2008 at 18:28

What about a liquid lead amalgam? As long as excess lead is used, the mercury should not nitrate as lead would reduce mercury nitrate. With thorough mixing and/or ammonium nitrate added to the bottom of the bath with a dip tube solid lead nitrate should float to the top. I am not sure how well it could be assured that all the AN has been reacted.

Taoiseach - 16-8-2008 at 03:23

Quote:

Yes the reaction IS exothermic. One just needs to melt the AN then remove the burner and put the lead in. If it is cut into sufficiently small pieces the reaction is selfsustaining i.e. the generated heat will keep the AN melted.

I never had problems with thermal runaway. I didnt try it on a large scale tough. From a theoretical standpoint it might happen tough - at high temperatures the NO3- might as well oxidize the NH4+ instead of the Pb.

There should be not nitric fumes tough! If you use stochiometric amounts then at least on a small scale the reaction will run smoothly, with the nitrate being reduced to nitrite which decomposes into N2 and H2O. The NH4+ reacts with the nitrate ion to ammonia and water. If you observe nitric fumes then they are from decomposing lead nitrate, not AN. Also the fact that you report a small explosion is quite worrying as there is nothing in this reaction which could explode! Lead metal does not form explosive ammines. You might have impurities of other metals present. AN also does not explode from heat, even when dropped onto a red hot heating plate all you get is a flame and a whoosh! sound.

Quote:

Can anyone think of another solvent that would work here (stable liquid from 160°C to 200°C, compatible and miscible with lead nitrate and ammonium nitrate, unreactive with lead)?

Yepp, its called water

I never even thought about this but this reaction should work perfectly well in aequous solution too. Water surely does not interfere with the reaction, as it is constantly produced by the reaction itself. Also it is known that even a cold AN solution does oxidize metals like iron and copper altough very slowly. You can try this easily: Just put a nail into a beaker with AN solution. After a few days you will observe a yellow precipate of iron hydroxide. When you use copper the solution slowly turns blue.

Thus you should try this: Make a conc. AN solution, heat to boiling and put in small pieces of scrap lead metal. Equilibrium should be driven to the right by decomposing ammonium nitrite into N2 and H2O and boiling off of ammonia. Advantages are: No solidification of fused mass at the end of reaction and thus no risk of decomposing Pb(NO3)2 (nitric fumes). Also the reaction temperature will be way below decomp. point of AN. The only problem I see is that it might react very very slowly. So make sure you cut your lead into very small pieces. Lead powder certainly would be best.

Here's the reaction:

Pb ---> Pb+2 + 2e-
NH4N+5O3 + 2e- ---> NH4N+3O2 + O-2
2N-3H4 + O-2 ---> 2N-3H3 + H2O
NH4NO2 ---> N2 + 2H2O
NH4+ ---> N + 3e- + 4H+
NO2- + 3e- ---> N + 2O2-

Thus:

Pb + 3NH4NO3 ---> Pb(NO3)2 + 2NH3 + 3H2O + N2

You need 1.2g AN for every gram of lead metal.

Good luck

kilowatt - 16-8-2008 at 14:50

Quote:

If you observe nitric fumes then they are from decomposing lead nitrate, not AN. Also the fact that you report a small explosion is quite worrying as there is nothing in this reaction which could explode!

Yes, that's what I was referring to. After the ammonium nitrate had been pretty much decomposed off I continued heating until the decomposition of lead nitrate, as a qualitative test for the presence of lead nitrate. In the test tube it always started with a considerable flash/pop followed by the controllable release of the nitric fumes. The decomposition of ammonium nitrate is quite capable of causing an explosion in such a situation where the temperature is quite high. It is quite possible the last amount of AN decomposes explosively especially given the rather rapid heating I used. I'm not sure about the nature of antimony nitrate or ammine complexes or if it would even form in the reaction, but there is certainly considerable antimony content in my mixed scrap lead alloy. Other metals capable of ammine complexing, such as copper, may be present in smaller amounts; the alloy comes from a mixture of random scrap lead articles.

I didn't realize the lead/AN reaction gave off nitrogen. The reaction can thus be considered a 1.5x waste of ammonium nitrate when compared to one that does not reduce nitrate, such as electrolysis. The electrolysis probably breaks down ammonia too, but it is obvious that a lot is still released from my test rigs; not sure what the yields are.

At this point I might be better off using magnesium oxide methods to break down the ammonium nitrate, later decomposing magnesium nitrate to release nitric vapors into a solution with lead filings. Perhaps there is no method for directly reacting the scrap lead and ammonium nitrate which has the characteristics I am seeking - a cyclic process or set of processes that produces nitric acid from ammonium nitrate without using up any additional acids/bases, and separates scrap lead into its constituent metals. The former goal should be achieved with magnesium nitrate. With that accomplished, the second goal should be achieved with the resultant nitric acid reacted directly with the scrap with no theoretical nitrate loss.

Thanks for all the info, and more ideas always welcome.

[Edited on 16-8-2008 by kilowatt]

Ephoton - 23-9-2008 at 18:31

my experiance with lead acetate.

I have tried a few ways to make this salt.
the best I found which is already stated here is.

acetic acid + peroxide treatement on metallic lead.

some ideas here are wrong though.

if you have ample acetic acid covering the lead acetate then temperatures
slightly over 100 C are not a problem.

you can distill off the water/acid till half of the liquid is removed.

if let to cool in another flask. you get very very big shards of lead acetate.

then pour liquid back into distillation flask and continue.

I can process half a kilo of lead salt in a one liter flask per go this way.
all salt is very very clear and water soluble.

add 300 grams of lead too 400ml of 85% acetic acid in a 1 liter flask setup for
reflux with equalized addition.

start water in condensor and drip 50% peroxide into solution through condensor
at a rate of around 3 to 5 drops a second.
continue this till solution reflux's (around 25 min) then slow down addition rate
to keep slight reflux.

when all of lead has been disolved filter and cool.

you will instantly get 300g of lead acetate percipitate when flask reaches room temp

filter then distill 150 ml of liquid.

place mother liquid into a flask too cool.

when crystals form (around 100g) pour liquid into distillation flask and
distill half of its volume again.

continue like this till you have all of the lead acetate crystalls.

never let the crystalls dry completely in the air or they will turn into lead carbonate
or at least basic lead acetate.

never boil solution to complete dryness.

keep adding peroxide too solution untill lead has disolved.

any excess is just given off as oxygen and does not hurt reaction.

do not add peroxide after lead has disolved or else you get a percipitate

a great mixture for crashing lead acetate from solution
is 30% ethyl acetate 70% ethanol.

[Edited on 24-9-2008 by Ephoton]

Pb(NO3)2 from Pb and N4NO3

UncleJoe1985 - 1-11-2008 at 11:03

kilowatt,

I'm interested in using your electrolysis method of making Pb(NO3)2 from NH4NO3 and Pb. I have difficulty finding PbO and could only obtain lead fishing sinkers. I have a steady supply of HNO3 from my arc furnace reactor. However, directly reacting lead with HNO3 is inefficient because 1/2 the acid decomposes into NO2.

My questions are:

1. What sized lead electrode did you use and how fast was the erosion?

2. How much current did you use?

3. How pure is the Pb(NO3)2 produced and what are the impurities?

Alternatively, if someone can tell me of a way to make Pb(NO3)2 with the chemicals at my disposal, or where to buy PbO, I'd like to hear it.

[Edited on 1-11-2008 by UncleJoe1985]

12AX7 - 1-11-2008 at 14:04

Why not lead the NO2 back into your acid reactor?

UncleJoe1985 - 1-11-2008 at 14:20

Nah, too much trouble unless I can absorb all of it. Somebody suggested using H2O2, but I don't know about its effectiveness yet.

12AX7 - 1-11-2008 at 14:30

If you simply use PbO, or PbCO3 or whatever, you'll have no problem. I suppose H2O2 might help, the solution will turn basic and the lead will dissolve, both in the presence of H2O2 which will decompose and oxidize fine in those conditions.

Tim

UncleJoe1985 - 5-11-2008 at 10:46

OK, I think I've narrowed down my choices in making Pb(NO3)2 using lead metal:

1. Pb = > PbCl2 => PbCO3 => Pb(NO3)2 as discussed here

2. Electrolysis of ammonium nitrate using lead as anode

For 1, can I use only HCl and no nitrate, or will that reaction be too slow? Could I speed it up by electrolysis? Also, I don't see how the chloride gets converted to carbonate. Assuming PbCl2 + Na2CO3 => PbCO3 + 2 NaCl is the path and it's reversible, it seems the products aren't favored because PbCl2 has very low solubility, except at high temperature, and PbCO3 is much more insoluble?

For 2, are there any impurities formed? How do I separate Pb(NO3)2 from the NH4NO3?

Please share any practical experiences

12AX7 - 5-11-2008 at 18:28

The products ARE favored because PbCl2 is very slightly soluble, PbSO4 is much less soluble, and PbCO3 a bit less soluble still. So, it is possible to make the sulfate or carbonate from the chloride in excellent yield, and the carbonate from the sulfate in reasonable yield.

Tim

UncleJoe1985 - 5-11-2008 at 20:08

12AX7,

I see you have a lot of experience with making lead salts on your web page. Can I make lead carbonate directly by electrolyzing lead with a sodium bicarbonate solution, which would be a lot easier than first making PbCl2? Then I can react the PbCO3 along with any lead that plated on the cathode with nitric acid to get what I want.

12AX7 - 5-11-2008 at 21:55

The anode may passivate. Inevitably a goodly amount of lead will dissolve (carbonate complex? high cathodic pH?) and deposit on the cathode as a spongy mess.

If you can prevent the deposition process (or don't mind screening it out periodically), it works.

But there's one more caveat. Bicarbonate is just fine, but as PbCO3 leaves solution, HCO3- turns to CO3(2-) and then CO3(2-) is depleted leaving more OH-. And lead dissolves in a hydroxide solution, enough that even more will be plated across. What's more, what does precipitate is Pb(OH)2, not the desired PbCO3 -- although that doesn't matter for your purpose at least.

The best approach, which I recall reading in a book in the SMDB library about industrial electrolytic production, involves adding CO2 to compensate. Ideally you'd have a bubbler, possibly under pressure, basically to maintain your electrolyte as salty club soda. That keeps the carbonate levels high, minimising pH (a soda solution concentrated with CO2 will hover around what, pH = 6?), keeping lead out of solution.

Tim

dann2 - 6-11-2008 at 03:21

Hello,

This may be of some use.

http://www.geocities.com/CapeCanaveral/Campus/5361/chlorate/...

Dann2

UncleJoe1985 - 6-11-2008 at 10:30

Quote:

That keeps the carbonate levels high, minimising pH... keeping lead out of solution

If I add HCl, will that keep the Pb ions from being reduced at the cathode?

12AX7 - 6-11-2008 at 20:51

HCl will increase chloride ion concentration resulting in PbCl2 fouling the anode. The only acid you want is carbonic, as that's what's being removed from solution.

Tim

dann2 - 10-11-2008 at 06:57

Hello,

Can accross the following here:
http://books.google.ie/books?id=0eGe5wDUEa0C&pg=PA13&...

Anothe way to Lead Chloride.

Dann2

Untitled-1 copy.jpg - 60kB

12AX7 - 10-11-2008 at 13:13

Sure, PbCl2 is nice and volatile. Just don't do it in a steel pipe (unless you want crusty holes and FeCl2 instead!).

Tim

dann2 - 11-1-2009 at 16:51

Link below to a book on Lead.
Worth a look for those big into this heavy stuff.

http://www.ila-lead.org/factbook/

Dann2

Swede - 24-3-2009 at 12:17

Bumpity - I am one of the lead dioxide nuts who can't stop attempting to plate this catalytic material for electrochemical purposes. Recently, I bought a home lead detection kit to see how far (and if) any Pb traveled from the reaction area, to settle on innocent surfaces.

The kit consists of tiny absorbent pads. You wet the pad with H20, then rub. When it encounters lead, the light orange color turns a brilliant scarlet. I verified this by wiping down my plating rig, which I knew would have Pb on it. The color was intense and unmistakable. A bit further away, I could still detect traces, but in general, the Pb did not travel far.

The question - anyone know the chemistry behind this kit? Two pads the size of a postage stamp was $10... I suspect the reagents used to make the kit costs maybe 8 cents. If something like this could be made at home, it would be both interesting and money-saving.

not_important - 24-3-2009 at 14:29

Quote: Originally posted by Swede

Recently, I bought a home lead detection kit to see how far (and if) any Pb traveled from the reaction area, to settle on innocent surfaces.

...

The question - anyone know the chemistry behind this kit? Two pads the size of a postage stamp was $10... I suspect the reagents used to make the kit costs maybe 8 cents. If something like this could be made at home, it would be both interesting and money-saving.

http://www.cpsc.gov/cpscpub/prerel/prhtml08/lead.pdf

http://www.ojp.usdoj.gov/nij/training/firearms-training/modu...

http://afte.org/Journal/aftejourW01shem.htm

rhodizonates
http://www3.interscience.wiley.com/journal/117952298/abstrac...

http://nvl.nist.gov/pub/nistpubs/jres/067/2/V67.N02.A06.pdf

Swede - 25-3-2009 at 05:19

Thank you, that is some excellent information. Sodium Rhodizonate... too bad it is not something simpler, but it is what it is.

Swede - 26-3-2009 at 07:32

not_important, thanks again. I found a source of Sodium Rhodizonate! Not too expensive for a somewhat exotic indicator. I already have tartaric acid and the correct buffering salts, plus HCl. It should be easy to prepare the test according to those interesting forensics websites.

I like the style of the commercial kits. I understand they are limited for certain applications, but for mine, which is simply wiping surfaces to find lead nitrate creep and mist, they're perfect. Heavy filter paper can be loaded with the reagents and activated with water.

no1uno - 3-4-2009 at 16:17

Lead carbonate is being used as a way of removing lead from used lead-acid battery electrolyte. They precipitate the lead carbonate from the solution by adding sodium carbonate to the lead sulfate/sulfuric acid mixture, forming sodium sulfate and lead carbonate.

Heating the lead carbonate to 400C or so should turn it to lead oxide/dioxide.

As used batteries aren't exactly difficult to get hold of, nor expensive, I cannot see why this route is not being taken up by home chemists.

Sedit - 3-4-2009 at 21:29

How does one distinguish between Lead(II)Acetate and Lead(IV)Acetate? They are both clear crystals IIRC so if one where to attempt to synthesis one from the other how would the know they achieved the results they desired?

[Edited on 4-4-2009 by Sedit]

not_important - 3-4-2009 at 22:07

Quote: Originally posted by no1uno

Lead carbonate is being used as a way of removing lead from used lead-acid battery electrolyte. ...
As used batteries aren't exactly difficult to get hold of, nor expensive, I cannot see why this route is not being taken up by home chemists.

Possibly because a liter of electrolyte has only a fraction of a gram of lead sulfate in solution.

As for distinguishing between plumbous and plumbic acetates, add a small bit to a ml or 2 of warm water, the Pb(II) salt will simply dissolve while the Pb(4) salt will decompose into insoluble brown PbO2.

Sedit - 3-4-2009 at 22:52

Easy enough. Thanks!

Swede - 4-4-2009 at 10:13

I have an odd request for lead chemists... there is a small but dedicated group of shooters who enjoy using legally-purchased silencers, especially for .22LR pistols and rifles. In the U.S., it requires a signoff by a local law-enforcement officer, and a $200 tax, but once those are done, Joe Civilian can buy his silencer.

The problem is buildup of metallic lead, especially in .22LR. These are dirty little cartridges that continuously vaporize lead and deposit it on the walls and in the chambers of the silencer. It is not unheard of for a 12 ounce silencer to weigh 64 ounces after a few thousand rounds have been fired; and it's all metallic lead, maybe a bit of carbon or nitrate residues.

There is NO traditional firearm cleaner that will get rid of such a horrendous lead buildup.

A few years back, someone suggested vinegar + peroxide to dissolve the lead, and turn it into lead acetate solution. It does work, and is what most guys are doing, but there was a lot of pooh-poohing when I mentioned the danger and toxicity of the remnants of this process.

My recommendation was to continue the recipe "as published" but to then add NaCl or some other salt to precipitate an insoluble form of lead, and then sequester those lead salts, rather than dump lead acetate down the drain, which is unfortunately what is probably happening, although no one will admit to it.

The silencers are made of aluminum, stainless steel, and occasionally titanium, and are usually sealed and impossible to take apart.

The challenge, then, is to improve upon the following: "Create a mixture of 50:50 household vinegar and OTC H2O2. Soak the suppressor for a week, agitating occasionally. Clean thoroughly with water when done." Something a bit more effective, safe, and responsible, and something that will not damage their toy. Any thoughts?

no1uno - 4-4-2009 at 11:54

not_important, there is NO shortage of dead batteries at the local dump and I suspect that this would not be all that unusual. They are just sitting there... Apart from that, how hard is it to dissolve additional lead in sulfuric acid? It kicks shit out of electrolysis and using nitrates... Couple of kilo's of sinkers, wait for them to go shiny & filter... Put in more, keep going until you have 20kg of clean, shiny sinkers

It's lead for god's sake, it will oxidise easily enough

dann2 - 4-4-2009 at 13:23

@Swede.
You could use Ammonium Sulphate (cheap readily available fertilizer from any gargen store) instead of NaCl so you will end up with insoluble (much less soluble than Lead Chloride) Lead Sulphate to dispose of.

Dann2

not_important - 4-4-2009 at 21:21

Quote: Originally posted by no1uno

not_important, there is NO shortage of dead batteries at the local dump and I suspect that this would not be all that unusual. They are just sitting there... Apart from that, how hard is it to dissolve additional lead in sulfuric acid?

Quite hard, once you reach the solubility limit of PbSO4, after which the sulfate will remain as a solid - a saturated solution is a saturated solution. Consider that PbSO4 is sitting there on the battery plates, a small amount dissolves in the electrolyte while the majority just sits there.

This is part of the reason lead containers were used to concentrate H2SO4 up to 78%, and why sulfuric acid plants used lead linings. Lead is near hydrogen in the electromotive series, the low solubility of PbSO4 gives the extra push that results in good resistance to less concentrated H2SO4.

4-Aminophenol - 5-4-2009 at 03:00

This is my Leadphosphate:

First I refluxed 169 g elemental lead with 113 ml of 65% Nitric Acid.
Than I dissolved the whole precipitate of Pb(NO3)2.
Now I dissolved 133 g of sodiumphosphate in Water, and added the solution to the Pb(NO3)2. I immediately received a white precipitate of Pb3(PO4)2. At least I filtered, washed with H2O and dryed the leadphosphate.

chief - 5-4-2009 at 05:09

How about using lead directly as anode in a chlorate cell ? Lead chloride as well as chlorate are insoluble, as well as lead hydroxide, oxide etc. . Any residual Lead-salts could be ppt. by adding as much as necessary Na2SO4-solution.

From the Pb-containing slurry of the bottom of the cell then Lead-salts could be made with weaker acids .... . The dissolution would be powered by the anodic oxidation and would be quicker ... ?

chief - 5-4-2009 at 05:59

A question would be if Pb-chlorate/perchlorate (if formed at all) might be a dangerous primary ... ?

Maybe even by regaining the Pb from the slurry (near 100 %), by heating with C, there could a Pb-Anode-cyclus be established, that would be quite much easier than the tedious plating of PbO2 onto graphite ... ?
As mentioned above any Pb can be ppt. out using Na2SO4, since PbSO4 is one of the most insoluble salts (from the standard acids) known to man. Only time is needed, to let the ppt. settle, and the rest of the solution will have maybe 1 mg/l or less of Pb.

With hot H2SO4 there is said to be some solubility of PbSO4, but thats no real wonder ...

Boiling the PbSO4 with conc. Soda-solution would most probably lead to PbCO3 ... ; maybe the others would react the same way: Hydroxide, dioxide ? Then this would be the best way to operate on the slurry, since the PbCO3 would be most useful for preparation of other salts ...

[Edited on 5-4-2009 by chief]

no1uno - 5-4-2009 at 19:53

Hmmmmmmmmm, bugger.... Ok, how to increase the solubility of the easily made lead salts (ie. those which can be made with the least effort and most accessible chemicals)?

Perhaps this might help:

Quote:

Description of the Leaching Process

The leaching process was developed by the Doe Run Company, independently of this research. This work has not been published in the open literature and is therefore summarized here. The process consists of contacting finely ground mattes and slags with a hot, saturated acidic brine solution containing approximately 25g/L ferric chloride. Ferric chloride serves to oxidise and solubilize metallic sulfides and is reduced to ferrous chloride in the process. Hot, saturated sodium chloride brine is employed to increase the solubilization of lead. Lead chloride is quite soluble in sodium chloride at elevated temperatures (~ 70g/L at 90C) whereas it has very low solubility (less than 1g/L) under conditions of low temperature and low chloride concentration

Taken from the Encyclopedia of Chemical Processing: http://tinyurl.com/d7yp9q

Does anyone have any lead chloride so they can check this? I have fuck all...

But it looks interesting, HCl is not exactly hard to get, NaCl, well....

Lead sinkers are everywhere (the dull coating on them is the carbonate usually)...

Use sodium hydroxide/carbonate to precipitate the lead hydroxide and then we could reuse the brine.

not_important - 6-4-2009 at 00:22

Strong solutions of alkali chlorides do somewhat increase the solubility of PbCl2, but not a great deal. Check the available books of solubilities.

Strong hydrochloric acid does better at increasing the solubility.

Lead is above hydrogen in the electromotive series, and will dissolve fairly quickly in non-oxidising acids if the resulting salt is soluble. Thus powdered lead dissolves in HNO3 and hot acetic acid, and in hot concentrated hydrochloric acid. A clever girl can use a lath to prepare lead turnings, or rid some wood rasps with a motor to ream off lead granules. Slowly add these to strong hot hydrochloric acid until it seems that the lead has stopped reacting, then cool and chill the solution, followed decanting and filtering off the PbCl2. The acid can be brought back up to strength and reused with further additions of lead.

Passing moist CO2 over lead turnings will slowly convert them to a basic lead carbonate. The Dutch process for white lead is based on the use of fermentation to produce CO2 and acetic acid, which react with metallic lead to form a basic carbonate containing a small amount of basic acetate.

chief - 6-4-2009 at 09:38

It's much easier to increase the surface of the lead: Just melting it and letting it drop into water will give grains with more or less flattened fringes. No mechanical dealing necessary.

not_important - 6-4-2009 at 17:02

I've had poor luck with mossy lead, there were always appreciable amounts of thicker sections that were slow to react, whereas the turnings reached rapidly. But you're correct in saying that it's easier to make mossy metals than to machine them.

einstein(not) - 27-4-2009 at 20:31

After a quick scan of the thread I haven't been able to find any info on what I might have made. If I missed it I apoligize in adavance.

100 grams Lead Dioxide and 100ml of glacial acetic acid were refluxed for 2 hours. Strong smell of acetic acid is gone from the clear liquid covering a brownish residue that appears to be equal in volume to the lead dioxide. At first I didn't think anything had changed but the lack of strong acetic acid smell must mean something did. Right?

Taoiseach - 27-4-2009 at 22:13

Anyone here who knows how to make lead nitrite, Pb(NO2)2?

I think I have accidentally made this already from NH4NO3 and Pb metal. I obtained a deep yellow solution. Adding a few drops of HNO3 and heating would make it colourless again.

Lead Carbonate from Lead Sulphate + Ammonium Carbonate

dann2 - 4-5-2009 at 17:00

Hello,
Anyone any comments to make on this patent. Lead Sulphate is easy to come by in old batteries. Does the Sulphate from old batteries contain any Antimony Sulphate?

EDIT:
From another thread(after U'ingTFSE):
have done this reaction to get lead acetate from old batteries. It does work but you have to grind the sulphate and boil it with lots of carbonate. Even then you don't get complete reaction but, if you settle or filter off the carbonate/sulphate mixture and leach it with acetic acid the carbonate disolves and you can repeat the process with the leftover sulphate.

Dann2

Attachment: US4220628.pdf (247kB)
This file has been downloaded 753 times

[Edited on 5-5-2009 by dann2]

dann2 - 21-2-2010 at 10:37

Hello,

PbO can be had by heating PbO2. I never actually seen this discussed but may have missed it.
Pb02 can be had from old Lead Acid batteries and should be fairly pure.
Will this PbO be difficult to get to dissolve in Nitric acid. I have read that it is difficult compared to PbO made by other chemical methods.
There is an article in References and translations needed(6) dated 21/2/10 (today) if anyone is interested.

Dann2

[Edited on 21-2-2010 by dann2]

Aqua_Fortis_100% - 21-2-2010 at 15:29

Today I made several long ingots of lead (from car batteries, lead seals, etc, all from scrapyard.. That should have some antimony , bismuth, etc but Im not very worried with it)

Tried to follow the info from patent US626330 "Process of Producing Peroxid of Lead". Used two Pb ingots, one as anode and the other as cathode. Used ~800mL of tap water, 2 teaspoons of homemade recristalized NaClO3 (containing some chloride) and 10 teaspoons of technical grade ammonium sulfate. +5V from PC PSU is already working on it. I dont know how many amps/current density because my multimeter is 'dead', but the electrodes are close, so I guess a decent current should be passing through it).

The anode is rusting and forming some white spots on the PbO2 layer but the electrolyte is clean, no PbO2 flaked in about 1 hour of electrolysis.

The only thing the patent did that I dont was bubbling air through the solution.. Maybe , I guess, that is just for stirring/flaking purposes, but I may be wrong.

Just for comparision, Im already doing a similair experiment but using plain NaCl as electrolyte (~100/L). When turning on power supply, winthin seconds appeared a rusty cloud on electrolyte.. And the PbO2 continues to flake off...

12AX7 - 21-2-2010 at 18:50

Quote: Originally posted by Aqua_Fortis_100%

<snip>
Just for comparision, Im already doing a similair experiment but using plain NaCl as electrolyte (~100/L). When turning on power supply, winthin seconds appeared a rusty cloud on electrolyte.. And the PbO2 continues to flake off...

Interesting. That was my experience also, although with fairly pure lead. Is it the alloy, or is lead (metal) suitable in chlorate solutions?

Tim

Aqua_Fortis_100% - 21-2-2010 at 19:15

Now the container with NaCl electrolyte have a decent layer of PbO2 on the bottom and the electrolysis keeps going.

Quote:

Is it the alloy, or is lead (metal) suitable in chlorate solutions?

Do you refer to chlorate -> perchlorate conversion? I dont tried ..But I have a 'feel' that this process wont work and if works, then will have poor eficience and will be messy. Maybe a little chloride in the home NaClO3 will be beneficial to make again PbO2 layer as it flake off.. Or maybe just chlorate will do, but havent tried this also.

In my ammonium sulfate/NaClO3 electrolyte is ppting a white substance,probably some sort of lead hydroxide, but almost none PbO2. So I think patent info dont works as described (unless the bubbling air step is really essential, but they didnt mention that)

[Edited on 22-2-2010 by Aqua_Fortis_100%]

12AX7 - 21-2-2010 at 21:32

Quote: Originally posted by Aqua_Fortis_100%

Ahh, then I misinterpreted this statement:

Quote: Originally posted by Aqua_Fortis_100%

T
The anode is rusting and forming some white spots on the PbO2 layer but the electrolyte is clean, no PbO2 flaked in about 1 hour of electrolysis.

If current were high, it would be interesting, but without a multimeter, you can't measure it accurately. Since it is forming a white skin (I would guess PbSO4), I take that to mean it's not working very well at all now?

Tim

Aqua_Fortis_100% - 21-2-2010 at 22:49

Oh sorry.. You is right.. I forgot that sulfate (lead-loving) ion was in solution, so cant be lead hydroxide.. The current keeps going, and the electrolyte have a temperature of about ~50°C.

Quote:

It dont forms a skin but some weird 'blobs'/clusters in some points of the anode, probably this is the reason that anode is still able to delivery current in solution.

Im wondering if the electrolytic method for producing carbonate/basic carbonate involving electrolysis in nitrate solution with CO2 bubbled through (already mentioned in some patent cited earlier) could really work well at home.

Im wanting somewhat pure that I can readily react with any acid (acetic, nitric, etc) without letting anything undissolved.
I have about a kilo of a mix of lead oxides, made few years ago from molten Pb + molten nitrate method.

http://i242.photobucket.com/albums/ff176/tnitrato/P9060068.j...

At time I was specting just nitrite and PbO, but reacting the well-washed/dried "PbO" with HCl gave chlorine that proved to have some PbO2/Pb3O4, etc in it..

So Im wondering if electrolytical methods will really give good products.. I dont think attempting PbCl2 route again since this salt has poor solubility, and therefore to purify it needs lots of hot water (I hate manipulating large amounts of water ahaha)

12AX7 - 22-2-2010 at 07:33

I've done that before, using chlorate. The transient species, lead chlorate, is soluble, and doesn't crystallize, so it isn't an explosion hazard. White Pb(OH)2 is the primary product. Without using a seperated cell and CO2 bubbler, you get some reduction of chlorate and lead, which should form an impurity of PbCl2 and lead sponge when it plates on the cathode, which is a nuisance.

Calcining the Pb(OH)2 produces a yellow to drab orange product, which should be PbO with some Pb3O4. I have never been able to produce bright orange Pb3O4 thermally.

Recrystallizing PbCl2 from hot water isn't so bad. You can reuse the water, it just takes a lot of passes.

Tim

UnintentionalChaos - 22-2-2010 at 08:35

Quote: Originally posted by 12AX7

Recrystallizing PbCl2 from hot water isn't so bad. You can reuse the water, it just takes a lot of passes.

Tim

Indeed. I'm under the impression, however, that it readily co-crystallizes with other lead halides. I had a mother liquor with traces of iodide (aka: my lead waste bucket) and I dropped out a crop of fairly white PbCl2. Double recrystallization from boiling distilled water did not seem to reduce the yellow tinge.

turd - 22-2-2010 at 13:10

Co-crystallization is a strange term in this context. It probably grows crystals of PbClI, a known compound (not a solid solution). Would be interesting to check under the microscope if there are different kinds of crystals or if this is really something like a solid solution with partial replacement of Cl by I.

dann2 - 22-2-2010 at 15:50

Hello AF100%,

Good to see you posting again!
There is a patent attached for making Lead Monoxide from not very pure Lead Sulphate. May be useful.

lead.JPG - 22kB

[Edited on 22-2-2010 by dann2]

Attachment: US4220628.pdf (247kB)
This file has been downloaded 698 times

Aqua_Fortis_100% - 22-2-2010 at 17:35

My chlorate/sulfate cell continues to, very slowly, produce some contaminated PbSO4. Some lead found their way in the cathode and formed a very thin lead sponge.

And in the chloride cell I scraped the anode and reverted the polarity on electrodes just trying to give a equal wear on both electrodes and produce more PbO2... But tomorrow I should stop the electrolysis since I will need the power supply.

Quote: Originally posted by 12AX7

Calcining the Pb(OH)2 produces a yellow to drab orange product, which should be PbO with some Pb3O4. I have never been able to produce bright orange Pb3O4 thermally.

Recrystallizing PbCl2 from hot water isn't so bad. You can reuse the water, it just takes a lot of passes.

Tim

Not sure if this was discussed before (is quite a time when I have read most of this thread, actually all of it in that time), but in the book Synthetic Inorganic Chemistry (fifth ed, pag 282) they describe the 'wet process' for producing Pb3O4, reacting PbO and PbO2 in basic solution, but even with this method in the end you have to heat the product (~350-400°C):

Quote:

WET METHOD

Materials: lead monoxide (litharge) PbO, 33 grams = 0.15
F.W.
lead dioxide PbO2, 24 grams = 0.1 F.W.
6N NaOH.
Apparatus: 600-cc. beaker.
5-inch funnel.
150-cc. casserole,
iron ring and ring stand.
Bunsen burner.

Procedure: It is rather difficult to adjust the temperature successfully for the dry method. Place 33 grams litharge, 24 grams of lead dioxide and 50 cc. of 6 N NaOH in a 600-cc. beaker. Stir thoroughly and leave in a warm place (80°), striring when convenient and adding water whenever the mass becomes dry, until the contents have become bright red. Finally wash the red lead thoroughly by decantation, and rinse on to a gravity filter in a 5-inch funnel. Let it drain over night. Lift the filter intact from the funnel, open it out on paper towels, and leave it on the hot plate until it is entirely dry. Detach the red lead from the paper by bending the paper; transfer the dry material to a 150-cc. casserole and heat it in a flame about 2 inches high, holding the casserole in the hand and rotating it in the flame. At the correct temperature (350-400°) the material becomes a dark reddish brown; after cooling it is a much more brilliant red than before heating. Great care must be taken to keep the material stirred during the heating so that the under layers do not become superheated and changed to PbO. Preserve the preparation in a 2-ounce corkstoppered bottle.

I dont have any experience in playing with Pb3O4, but this substance seems interesting to have lying around (IIRC some crackling star compositions are Pb3O4 based).

About PbCl2, I may try that, although Im now out of gloves

Tim, In your lead compounds page you seem to have a relevant work to manage liters of solution split in several pickle jars to purify PbCl2..(yeah DIY glassware was and still will be also my only (real fun!) way of producing compounds..)

Thanks Dann2, Im also glad to see you again and even more glad to see the amazing patent you posted.
The most interesting part IMHO, is that the reaction between ammonium carbonate and lead sulfate is usually quite rapid (less than 1 hour and most between 5 to 15 mins) and no heating is needed (actually I think heating probably cause part of ammonium carbonate being decomposed).

I only dont know if ammonium bicarbonate (more OTC here) is also suitable as direct reactant in place of carbonate. I was thinking in cautiously heating the solution to remove excess CO2 and make normal carbonate, but Im worried if that wouldnt destroy a significant amount of the bicarbonate.
If ammonium bicarbonate cant be used, maybe just using slight excess of bicarbonate in dilute ammonia water should result in converting most in ammonium carbonate.

BTW today I have visited a electronic components store to buy some diodes and as they have a trash specific for batteries, I asked to seller for some of it. He was glad to gave me a plethora of NiCd, Ni-MHD, MnO2, etc baterries and among them was a small 6V lead-acid battery. This battery reminded me instantly this thread.

[Edited on 23-2-2010 by Aqua_Fortis_100%]

[Edited on 23-2-2010 by Aqua_Fortis_100%]

12AX7 - 22-2-2010 at 17:58

Quote: Originally posted by Aqua_Fortis_100%

in the book Synthetic Inorganic Chemistry (fifth ed, pag 282) they describe the 'wet process' for producing Pb3O4, reacting PbO and PbO2 in basic solution, but even with this method in the end you have to heat the product (~350-400°C):

<snip>

Interesting. I am aware that PbO2(2-) and PbO3(2-) react, in fact I think it was in Brauer where, after a week or so, very large crystals are claimed to be obtained (>2mm) from such a solution. It seems like that process is being used as a catalyst here, which is a fine way to go. I would suppose you could reflux it as well, in which case you wouldn't have to add water, and it might go faster at ~100C.

Too bad about the final heating, though. 450-500C is pretty tight for "couple inch flame"...

Quote:

Yup! I had recrystallized a satiny white product, 100g or so. Nice crystals, not very distinct in shape but the way they pile up is cool.

Quote:

BTW today I have visited a electronic components store to buy some diodes and as they have a trash specific for batteries, I asked to seller for some of it. He was glad to gave me a plethora of NiCd, Ni-MHD, MnO2, etc baterries and among them was a small 6V lead-acid battery. This battery reminded me instantly this thread.

Yay for cadmium!

Now you can destroy your kidneys even easier :rolleyes:

Tim

Aqua_Fortis_100% - 23-2-2010 at 14:21

After decanting/washing several times Im now filtering the PbO2. Will treat it with some vinegar hoping to dissolve any oxide/hydroxide contaminating PbO2, then wash again, then dry.. The yield seems good for a "quick'n'dirty" electrolysis.

And, yeah, I was wanting some lithium and nickel based batteries, but cadmium comes along with it.. But I have no use in mind for cadmium, at least now

JohnWW - 23-2-2010 at 14:35

Cd is used in CdSe semiconductors (having a relatively low band gap voltage). I addition, a mixture of CdSe and CdS is used as a red paint pigment; and CdS is used on its own as a yellow paint pigment. The metal itself is used in some low-melting solder alloys, with Sn and Pb.

[Edited on 23-2-10 by JohnWW]

Aqua_Fortis_100% - 24-2-2010 at 13:25

Ive dried the PbO2 in the oven at 100°C but the brown powder contain some weird white granules. What can be? Pb oxide not leached by vinegar? Yield is about ~20-25g.

The filter containing PbO2 in the pores was treated with dilute HCl and the chlorine smell was perceptible. After, the solution was treated with excess baking soda and a small amount of a tan precipitate was filtered (but being a very small amount, I will just discard it in proper manner to a battery recycling center as "lead carbonate").

JohnWW thank for the info. However, Im wondering what the home chemist can make with cadmium rather than pigments and solder alloys (and of course, dont being poisoned). Some kind of catalyst for a crazy/interesting reaction, maybe? Or then another fun use, if any.

[Edited on 24-2-2010 by Aqua_Fortis_100%]

densest - 24-2-2010 at 16:36

Cd is used to color glass - some "crayon" yellows,reds, and oranges.

Panache - 25-2-2010 at 03:33

Quote: Originally posted by 12AX7

Calcining the Pb(OH)2 produces a yellow to drab orange product, which should be PbO with some Pb3O4. I have never been able to produce bright orange Pb3O4 thermally.

Tim

Why do you think this is? Are you certain of your Pb(OH)2?
When making some Pb(IV)Acetate late last year i had begun to go down this road as it seemed the logical way to Pb3O4, well I got as far as collating some lead scraps to begin when i found Pb3O4 for sale not shy of 1km from where i live. That made life much easier, but it appears from your experience i may have struggled making the Pb3O4 thermally.
For those interested in Pb(IV)Acetate the method in Brauer works as described, i used the method with acetic anhydride, am unsure regarding the method that is mentioned afterwards where the temperature is kept lower and no AA is used.

12AX7 - 25-2-2010 at 12:28

This is what I got from roasting:

Maybe it wasn't hot enough, maybe it was too hot, maybe it was too uneven. Dunno.

Tim

Panache - 25-2-2010 at 19:41

Quote: Originally posted by 12AX7

This is what I got from roasting:

[/img]

Maybe it wasn't hot enough, maybe it was too hot, maybe it was too uneven. Dunno.

Tim

I guess you weren't shocked when the commonly referred to procedure failed, how accurate is your temperature control in your kiln, can it be calibrated. I mean one really has to squint quite hard and be quite away from the srceen to see much orange in that sample. Whereas the commercial sample, which i dried at 200 overnight as per Braeur, is well, as you can see.

Image010.jpg - 18kB

[Edited on 26-2-2010 by Panache]

JohnWW - 25-2-2010 at 20:05

That looks like good pigment-quality "red lead", Pb3O4, Panache. Although out-of-favor in final-coat paints because of the toxicity of Pb, it is still used as a pigment in timber paint primers, and in pottery glazes.

[Edited on 26-2-10 by JohnWW]

Panache - 26-2-2010 at 02:10

Quote: Originally posted by JohnWW

That looks like good pigment-quality "red lead", Pb3O4, Panache. Although out-of-favor in final-coat paints because of the toxicity of Pb, it is still used as a pigment in timber paint primers, and in pottery glazes.

[Edited on 26-2-10 by JohnWW]

Spot on John
Came from an artisan pigment supplier, for those fine art painters wanting to grind their own oil paints. I thought it was cheap $30.00 for 500g. I mean some poor dude out the back of the store has to repackage it.
They began to give me the 'moron's' treatment before i cut them off deftly saying 'yes yes i know, but i'm not using it to paint with' as the cautions were in regards to grinding it.
'Oh what are you using it for?', i had the attention of three shopkeeps and a customer now.
'I'm making tetra ethyl lead' i replied
'Whats that used for?'
Oh how quickly we forget. One of them was older as well. tut tut.

Lead(II) acetate query

rajip - 1-3-2010 at 06:23

The photo 12AX7 posted of Lead acetate looks exactly like the photo at this site: http://webpages.charter.net/dawill/tmoranwms/Chem_Pb.html

This is how the person at the site posted above said he produced the Lead acetate:

Quote:

'Here's some more lead acetate. This was produced by dissolving stock lead oxide (the calcined stuff) in vinegar, however much a gallon of 4% acetic acid would take.'

SO...Can PbO (or is it PbO2) + vinegar/weak acetic acid be used to prepare lead (II) acetate or not?

watson.fawkes - 1-3-2010 at 08:14

Quote: Originally posted by rajip

The photo 12AX7 posted of Lead acetate looks exactly like the photo at this site: http://webpages.charter.net/dawill/tmoranwms/Chem_Pb.html

Holy non-secret identity!

rajip - 5-3-2010 at 04:57

Quote: Originally posted by watson.fawkes

Quote: Originally posted by rajip

The photo 12AX7 posted of Lead acetate looks exactly like the photo at this site: http://webpages.charter.net/dawill/tmoranwms/Chem_Pb.html

Holy non-secret identity!

I get the sarcasm but could someone clear up the air (pun intended) for me?

CFan PbO or PbO2 be produced in a furnace or kiln from lead metal, and if so how?

Wikipedia states that litharge (PBO) + acetic acid (does this include weak solutions like vinegar?) can be used to form lead (II) acetate, is that true?

12AX7 - 5-3-2010 at 06:39

PbO dissolves in vinegar just fine, that's how the ancients did it, and that's what I did on my website.

PbO can be produced directly (although also as indicated on my website, the yield can be poor, resulting mostly in suboxide and lots of granular metal), but probably not PbO2.

Tim

barbs09 - 18-7-2010 at 02:17

Hi, I have opened up a small lead-acid battery to see if I can find any easily extractable PbO2 and Pb metal. The battery was old and unable to crank much but prior to the dissection I gave this battery a good charge to concentrate the H2SO4 and to convert as much of the PbSO4 to PbO2.

The anode consisted of pressed brown PbO2 in a Pb grid which was housed between the cathodes in a plastic sheath which must be a porous membrane. Due to the damaged (old) state of the battery the Pb grid had collapsed and the plastic sheath held a mass of brown muck with pieces of Pb grid-no wonder it couldn’t hold a charge!!

My question is: how pure is the PbO2 and has anyone attempted to isolate it and obtain the usable PbO2. I would like to know if such recycling is worthwhile before I attempt this messy and potentially hazardous (many precautions taken) separation.

Also apparently the cathode consists of Pb sponge pressed into a Pb grid. These pressed pieces certainly rush to a grey powder in a pestle and mortar. Such finely divided metal would probably be more reactive which could perhaps facilitate the Pb + acetic acid preparation of lead acetate? One would have to consider the likelihood of antimony and other additives participating in this reaction.

Any thoughts welcome.

Barbs

thehappyhymenoptera - 24-7-2010 at 15:08

Has anyone had any success making lead tetroxide? I'm having troubles.

blogfast25 - 11-9-2011 at 09:42

Sorry to revive this old thread.

I want to make some lead (II) acetate (trihydrate) from scrap (battery) lead and acetic acid (I have glacial acetic so can prepare any concentration). I’ve trawled through the entire thread but little definitive is said on a workable recipe for the acetate.

Garage Chemist claimed success with dissolving lead in hot 25 % HAc, combined with 30 % peroxide (first page) but he wasn't very explicit. And someone else warned about over-oxidising the lead to lead (IV) acetate.

12AX7 (Tim) claims lead dissolves even in commercial vinegar but isn’t very specific either. That sound pretty slow to me, lead’s not hugely reactive…

Lead acetate does have a strong differential in hot/cold water solubilities, which should serve well for recrystallisation w/o boiling to dryness (not recommended with acetates anyroads).

Does this jog anyone’s memory on a prêt-a-porter recipe to fairly clean Pb(Ac)2.3H2O?

Magpie - 11-9-2011 at 13:43

If you want to use the route I did, ie, via PbO, you'll have to make the PbO first using KNO3. This can be done as I show here:

http://www.sciencemadness.org/talk/viewthread.php?tid=52&...

That red mud is surely PbO.

blogfast25 - 12-9-2011 at 06:46

Thanks Magpie, but I'm looking for something a little more direct.

For instance I've got lead nitrate, so I could e.g. go nitrate > carbonate > acetate but that seems such a waste of nitrate, if direct (and swift) dissolution of lead scrap in HAc + H2O2 is indeed feasible (as garage chemist suggests).

Guess I'll have to experiment, huh?

not_important - 12-9-2011 at 06:58

Don't even need H2O2, just bubble air through the acetic acid with the lead in that. The more surface air the lead has the better.

blogfast25 - 12-9-2011 at 07:04

Quote: Originally posted by not_important

Don't even need H2O2, just bubble air through the acetic acid with the lead in that. The more surface air the lead has the better.

Yes, that was mentioned higher up but my (limited) experience with air oxygen as an oxidiser is that it's pretty slow. I used it once to oxidise Fe (II) in hot (80 C) acid solution and although it worked it took a long time to get to completion. Not highly practical from that perspective.

Worth a try though...

blogfast25 - 15-9-2011 at 10:44

Lead Acetate from battery lead, glacial acetic acid and hydrogen peroxide

I took some pictures of this experiment but due to a Blogger glitch I can’t upload. Maybe tomorrow…

A 58.3 g lead button (about 0.26 mol) was obtained as part of a defunct car battery electrode and cleaned up (a bit) with some sanding paper.

I then mixed 47 ml of water, 36 ml of glacial acetic acid and 26 ml of 35 % H2O2. This is the theoretical stoichiometric amount for:

Pb === > Pb2+ + 2 e
H2O2 + 2 H+ + 2 e === > 2 H2O
2 x [ HAc === > H+ + Ac- ]

Pb + 2 HAc + H2O2 === > PbAc2 + 2 H2O

And would yield (at 100 % yield and no losses) a saturated solution of PbAc2 at 40 C (sol. 116 g/100 g water, acc. Wiki). Of course this is ‘theory’ and nothing more than a starting point…

The lead button was then added (everything at RT) and reaction started immediately, with strong evolution of gas (oxygen? hydrogen?). The reactor (an oversized tall 1 l pyrex beaker with a cooled lid for reflux) was then moderately heated on an electrical hot plate.

After simmering away for about 35 min the reaction more or less died, and the solution/suspension was dark grey with plenty lead left. In total I revived the reaction about three times, each time by adding 10 ml HAc and 10 ml H2O2 and stopped after about a total time of 105 min. I then hot-filtered the grey solution/suspension on a #1 filter: it filtered easily to a clear solution. But there’s plenty metal left: visually I’d say only about half of the lead had actually dissolved. The left-over sludge, left mainly in the 1 l beaker, was subjected to the same treatment as described above and also filtered. But even after that second treatment grey-black sludge with pieces of lead are left!

One ml of the first filtrate was treated with KI solution and tested very strongly for lead: lots of yellow PbI2 precipitated. And after about 30 minutes when the PbI2 had precipitated to the bottom completely, some I2 could also be seen: oxidation of iodide by residual peroxide? Oxidation of iodide by Pb (IV)?

Both solutions are now being iced overnight but I somehow doubt I will obtain PbAc2.3H2O crystals because I think the solubility limit at 0 C (19.8 g/100 g water, acc. Wiki) will not be exceeded.

So far this has not been a ‘quick and dirty route’ to lead (II) acetate.

Does any one know the main contaminants in electrode lead? It seems harder than sheet lead e.g.

[Edited on 15-9-2011 by blogfast25]

watson.fawkes - 15-9-2011 at 12:42

Quote: Originally posted by blogfast25

Does any one know the main contaminants in electrode lead? It seems harder than sheet lead e.g.

Not sure about your sample in particular, but antimony is the most common lead hardener for alloys. There's likely some others in there, although likely not at higher percentages than Sb. I'd suspect As and Sn as well.

Arsenical alloys for batteries: http://www.keytometals.com/Article88.htm

blogfast25 - 15-9-2011 at 12:58

Quote: Originally posted by watson.fawkes

Antimony: I thought as much. That would not dissolve in this mixture I think (but it may turn to oxide). Later on this week I’ll be dissolving the filter residues in aqua regia, that should roughly separate any Pb from Sn/Sb. Trouble is that the latter two aren’t easy to separate or test for with wet methods…

I was kind of surprised at the hardness of this lead: fairly though to hacksaw too…

Thanks for the link.

[Edited on 15-9-2011 by blogfast25]

dann2 - 15-9-2011 at 14:52

Be aware that the gas liberation will cause a mist of liquid to come out of the container (like a fizzy drink throws up droplets of itself out of a glass). This will be toxic. Put a cover on the beaker.

[Edited on 15-9-2011 by dann2]

blogfast25 - 16-9-2011 at 02:56

Quote: Originally posted by dann2

Be aware that the gas liberation will cause a mist of liquid to come out of the container (like a fizzy drink throws up droplets of itself out of a glass). This will be toxic. Put a cover on the beaker.

[Edited on 15-9-2011 by dann2]

There was. As reported.

not_important - 16-9-2011 at 06:08

Quote: Originally posted by watson.fawkes

...Not sure about your sample in particular, but antimony is the most common lead hardener for alloys. There's likely some others in there, although likely not at higher percentages than Sb. I'd suspect As and Sn as well.

...

Calcium is not uncommon, especially in the smaller sealed batteries but apparently also in extended life automotive types as well.

blogfast25 - 16-9-2011 at 06:36

Quote: Originally posted by not_important

Calcium is not uncommon, especially in the smaller sealed batteries but apparently also in extended life automotive types as well.

The alloys used for outer electrodes and pressed 'panels' may also be different. Sn/Sb/As affect charge retention, acc. Watson's link. Hardness is definitely more important for the exterior electrodes.

It'd be interesting to test acetic acid/H2O2 on a pure lead sample but I haven't got any. Lead shot, wheel balancing weights, fishing weights etc are all likely to be adulterated.

As expected no lead acetate crystals formed in the iced solutions, concentration being too low by a factor 2 to 4 I believe. But glistening smallish crystals can be seen on the dried filters...

[Edited on 16-9-2011 by blogfast25]

blogfast25 - 17-9-2011 at 09:30

Lead acetate update:

Below: the 54 g button of lead I started off with:

1. Crystallisation of the lead acetate:

The solutions of both tests above were combined, giving 272 g of solution which was first boiled down to about half the volume and then cooled. No crystals were obtained. Today I boiled it down further to 89 g of solution at which point it kind of had stopped boiling altogether. No turbidity developed. On cooling the mass congealed into a perfectly clear, stiff but jelly like substance:

This is likely to be a supersaturated solution or a supercooled liquid. I added 50 ml of DIW and dissolved the mass on the hot plate, then added a small amount of seed crystals (some lead acetate that had crystallised on a spatula). Within minutes crystals started to appear:

On cooling on ice, within minutes even more crystals were forming but I didn’t have time to take a shot.

2. Treatment of the filter residue with acids:

One part was treated with HNO3 38 %, another part with HNO3 38 % + HCl 37 % (aqua regia). In both cases the residue turned to a much lighter colour and some NO>NO2 developed.

With the nitric only, some white crystals of Pb(NO3)2 formed on cooling. The lighter colour of the residual nitric insoluble matter could point to SnO2 and/or Sb2O5, as both are highly insoluble.

With the aqua regia, plenty of PbCl2 precipitated on cooling. But the residue didn’t completely dissolve and a white mass was left: hard to say whether it was PbCl2 or the said Sn/Sb oxides or both and I didn’t attempt separation.

3. Aqua regia and the original lead alloy:

Determined to try and show presence of Sn or Sb, I tried to dissolve 25 g of the alloy in aqua regia (110 g HCl 37 % + 20 ml HNO3 70 % - this is a formula that works excellently for pewter). It worked surprisingly badly here. At RT there is slight attack, but at BP reaction is vigorous with loads of NO>NO2 evolved. But the lead alloy only dissolves slowly and after about 1 h the reaction died and most of the metal hadn’t dissolved. On cooling typical PbCl2 crystals formed but the quantity was well below what to expect from 25 g of lead (0.116 mol). It’s pointless to try and demonstrate Sn or Sb when so little alloy had dissolved. Tomorrow I’ll have another jab at it with strong HNO3…

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