Sciencemadness Discussion Board - The lead salts preparation thread!

Acetic acid is becoming more easily obtainable as an organic (green) weed killer. It is available as a 50% solution.
A mixture of Sodium Nitrate and HCl acid will dissolve Lead. Careful of the brown fumes.

Would Ammonium Nitrate do instead of the Sodium?

Dann2

Yes, but what would you obtain? Lead nitrate is far less soluble than lead acetate and it’s the latter I’m trying to make, in various ways.

My lead, which really is from the connector part that connects the cells, doesn’t even easily dissolve in hot 70 % HNO3: reaction starts on heating but more or less dies after about 30 min.

The green weed killers are most very weak solutions of pelargonic acid soaps (or fatty acids of similar length), a few percent. Acetic is used too but at 50 % that’s a ‘human repellent’! Various acetic acid strength grades are cheap on eBay.

Conversion of PbSO4 to lead basic carbonate (‘white lead’ - vernacular)

I’m interested in converting the lead (II) sulphate from my spent lead battery to lead (II) acetate.

Yesterday and today I explored the following displacement reaction:

3 PbSO4 (s, aq) + 4 Na2CO3 (aq) + 2 H2O (l) === > (PbCO3)2Pb(OH)2 (s)+ 3 Na2SO4 (aq) + 2 NaHCO3 (aq)

… taking advantage of the slight solubility of PbSO4 and the complete insolubility of ‘white lead’.

To that effect 3.25 g of PbSO4, 3.4 g anh. Na2CO3 and 17 ml of water were mixed in a 1” test tube. The amount of soda is almost 3 times the stoichiometric amount. The test tube was heated on steam bath for about an hour, and then allowed to cool and stand overnight.

The resulting precipitate was the filtered off, setting aside the first portion of filtrate and carefully washing the filter cake with multiple aliquots of hot DIW until it ran almost neutral and no bubbles formed with strong HCl (evidence that soluble carbonate had been eliminated).

The filtrate set aside was first neutralised slowly with strong HCl until no more bubbles formed and the solution reacted strongly acidic. It was then tested with Cu(NO3)2 for residual carbonates and tested negative, then tested again with conc. CaCl2 which caused a shed load of white CaSO4 to precipitate. It appeared the anticipated reaction had proceeded.

The filter cake had in the mean time been dried on filter, on a low setting hot plate and two medium pinches were added to 2 test tubes and a little water added: this stuff appears to be almost water repellent! The lead sulphate I started from did not show this.

Adding 38 % HNO3 to the first, fizzing started, then subsided and a perfectly clear solution resulted. To the second tube, some glacial acetic acid was added. That was noticeably slower and needed a little heat but it too resulted in a perfectly clear solution. All 3 test results are strong evidence that the displacement reaction had proceeded to completion.

The next step will obviously be to test this on battery lead (II) sulphate.

[Edited on 21-9-2011 by blogfast25]

sxl168 - 21-9-2011 at 21:08

I've done this before on a larger scale and it works fine. I removed 2 car batteries worth of positive plates and tossed them into a 5 gallon bucket. I then filled with water so the bucket was about half full. Tossed in an 5 LB bottle of pool pH plus from Wal-Mart, stir well for an hour or 2. Note that I didn't need to add heat as dissolving that whole bottle increased the temp of the solution to about 50-60C. I then let the bucket stand for a day, stirring once in a while. Decanting off the liquid, I recrystallized it twice to get a few pounds of crystal clear Sodium Sulfate. Some leftover Sodium Carbonate was also left after crystallization, but amounted to less than 50g or so. Treating the solids with vinegar/acetic acid produced large amounts of lead acetate.

I process the negative plates differently because they can be easily electrolyzed (in distilled water/dilute sulfuric acid) to remove the sulfate, and the process produces sulfuric acid at the same time. I use one of the battery posts as an anode when doing that. I did it slowly when doing mine, one plate at a time each taking 24 hours. The process is completed once hydrogen is evolved on the negative plate, or if you have a DMM, the voltage of the cell rises from the nominal 2.0-2.2 volts to 2.4-2.6 volts. I was using about 50 mA/cm^2 anode current density for those numbers. The voltages will be much higher if the ion concentration is low, such as when you first start the run. I ended up with about a 40% sulfuric acid solution once the last plate was finished.

blogfast25 - 22-9-2011 at 04:01

Quote: Originally posted by sxl168

I've done this before on a larger scale and it works fine. I removed 2 car batteries worth of positive plates and tossed them into a 5 gallon bucket. I then filled with water so the bucket was about half full. Tossed in an 5 LB bottle of pool pH plus from Wal-Mart, stir well for an hour or 2. Note that I didn't need to add heat as dissolving that whole bottle increased the temp of the solution to about 50-60C. I then let the bucket stand for a day, stirring once in a while. Decanting off the liquid, I recrystallized it twice to get a few pounds of crystal clear Sodium Sulfate. Some leftover Sodium Carbonate was also left after crystallization, but amounted to less than 50g or so. Treating the solids with vinegar/acetic acid produced large amounts of lead acetate.

Thanks sxl, that’s very much how I plan to do it today. I won’t be separating the lead from the lead sulphate and even heat might not be necessary although it would speed up things a bit. I was thinking of using an excess (as in the test) of soda, then simply reusing the solution several times.

Your pH pool plus was NaOH I gather from the temperature increase (solvation energy), right?

sxl168 - 22-9-2011 at 10:16

No, the pH Plus was ordinary Sodium Carbonate. It does generate a lot of heat when it dissolves in water, not as much as NaOH, but still significant when using the large quantities involved. I forget where I saw it now, either the label or the MSDS for the product clearly indicated that it will substantially heat water if the volume of water dissolving it is low. I'm sure some of the heat generated was the rapid chemical reaction between Sodium Carbonate and the fine particles of Lead Sulfate and leftover Sulfuric Acid too. I just remember it taking only about a minute for the water temp. to jump from 20C to 50-60C.

sxl168 - 22-9-2011 at 10:39

I just wanted to comment also that I came across some patents and research papers that use a similar process for wet process recycling of batteries and they warn that high pH levels over 11 along with high solution temperatures and long contact times would produce water soluble plumbates. You may want to be careful how much Sodium Carbonate excess you add because of that. From what I read, keeping solution temps. under 70C, contact time under 2 hours, or a final pH of under 11 keeps soluble plumbates from forming.

dann2 - 22-9-2011 at 11:19

Yes, but what would you obtain? Lead nitrate is far less soluble than lead acetate and it’s the latter I’m trying to make, in various ways.

....

The green weed killers are most very weak solutions of pelargonic acid soaps (or fatty acids of similar length), a few percent. Acetic is used too but at 50 % that’s a ‘human repellent’! Various acetic acid strength grades are cheap on eBay.

.....
[Edited on 21-9-2011 by blogfast25]

The Acetic acid weed killer is available at 50% solution strenght. You have to water it down yourself when it comes to actually using it as a weed killer. It a ready OTC source of Acetic acid if used in your neck of the woods.

Would the mixture of HCl acid + Ammonium Nitrate be like poor mans Aqua Reiga.
You would end up with Lead Chloride, then go to Lead Carbonate (via Sodium Carbonate or bicarbonate) and then into Lead Acetate. The Nitrogen and the Ammonia going away? It's just that Ammonium Nitrate may be easier to come by than Sodium Nitrate.

The origan recipe using the Sodium Nitrate + HCl comes from here:
http://www.oxidizing.110mb.com/chlorate/leaddiox/pbnitrat.ht...
Down the page a bit (not my work BTW)

You eventually get to the Lead Acetate via the Carbonate

Dann2

[Edited on 22-9-2011 by dann2]

blogfast25 - 22-9-2011 at 11:21

@ sxl:

Your 'pH up' must have been anhydrous Na2CO3 then: that does generate heat when rehydrated. I use Na2CO3 decahydrate, not much solvation enthalpy there. Some heat would come from the reaction, yes.

One of the reasons I used Na2CO3 was that it's less likely to generate plumbates, which acc. my lit. research really only form with 50 % NaOH (or KOH) or so. The CO32- is too weak a base to reach to high concentrations of OH-, needed to form Pb(OH)6(2-) anions.

[Edited on 22-9-2011 by blogfast25]

sxl168 - 22-9-2011 at 12:58

I found the pdf that I was trying to cite in the previous post. It can be found here and has decent info:
Recycling ULAB's

blogfast25 - 22-9-2011 at 13:20

Thanks both dann2 and sxl for these links, very useful indeed...

sxl: my excess of Na2CO3 would indeed appear to have been unecessary but also unlikely to lead to plumbites being formed. These require extremely high levels of OH- that even a strong solution of carbonate cannot deliver because carbonates are weak bases.

[Edited on 23-9-2011 by blogfast25]

blogfast25 - 23-9-2011 at 12:20

Well, well. It looks like my defunct lead battery hadn’t been discharged because there no lead (II) sulphate in the first cell I broke into. I found this battery years ago, drained off the acid and thought no more about it until recently. So that rather buggers my plan to convert the sulphate to white lead, unless I want to attempt to fully discharge the two remaining cells. But she may be faulty in other ways so I’m not chancing that.

That leaves only to try and improve dissolving process for the actual lead, something I’m having trouble with. I’ll have to revisit using home-brewed aqua regia, based on 1:3 nitric/hydrochloric but with 70 % nitric instead of fuming. Nitric alone dies quickly against this particular alloy. 50 % acetic with gradually added 35 % H2O2 might also be worth reconsidering.

Any decent ideas welcome…

sxl168 - 24-9-2011 at 12:45

I wouldn't judge the lead sulfate by visual inspection. I've seen nice silver-grey battery plates that were completely discharged as I found out when I charged the plates. The lead sulfate is a surface layer only on the Pb/PbO2 particles and is mostly transparent unless it is thick (Sulfated battery). If your battery has been sitting around for 2 years as you have said, it's pretty much guaranteed that it to be at over 50% discharged.

If you are just trying to get Pb into solution, household bleach worked for me in the past. The key is to stir while adding and only add 1% vol./vol. of bleach at a time and let it react. This will keep most of the Chlorine which will be generated in solution to eventually react with Pb/PbO2. It's a slow process but worked well for me. Works well on other metals too like Ni, Co, Fe, and Cu. Warm the solution up to 50C or so also to increase solubility of PbCl2, decant, cool to 1C and collect PbCl2 crystals. Reuse the liquid to dissolve more. Don't get too hot as solubility of Chlorine in water drops drastically above 50C. Your method of using H2O2 should work fine also, but I find H2O2 a bit expensive to use. You can use Epsom salt for a poor man's way to precipitate PbSO4 if you wish from PbCl2 solution, the reaction may take 15 min to kick in, but settles fast once it does.

sxl168 - 24-9-2011 at 13:00

I forgot one added bonus of the PbCl2 route is that so long as the crystals are cleansed with cold distilled water, you will be rejecting the Ca and Sb alloys present in the batteries.

blogfast25 - 25-9-2011 at 04:52

@sxl:

I know what you mean about the sulphate but I actually treated some of the (supposed) sulphate with carbonate and obtained absolutely nothing. On closer inspection of the material and a few tests with nitric acid, one can only conclude it is almost 100 % lead. On the plus side: that means I have also PbO2 which is easy to convert to sulphate plus oxygen with hot sulphuric acid.

Regards the dissolution of lead metal in acids, I’m probably a little unrealistic in terms of expected speed and my memory of past successful attempts must be playing up a little.

Still, a proper quantitative test with 8 g of battery lead and twice the stoichiometric amount of aqua regia (1:3 - 70 % nitric to 36 % HCl) showed again that this alloy is probably quite resistant to chemical dissolution. Reaction started at RT, became very swift for about ½ hour during simmering but then more or less died and little metal had actually dissolved. After hot filtering, diluting and icing the solution, about a flat teaspoon of PbCl2 was obtained. Disappointing yield! This is a slow boat to China…

I’m now pretty convinced, including from reading those links that the ‘fastest’ route to dissolving lead metal is probably as follows. Use a large excess of an acid (at least twice stoichiometry), HCl or HAc, and add the metal to it at 50 - 80 C (avoid boiling). Then start adding small aliquots of a suitable oxidiser (HNO3, H2O2, perhaps as you suggest hypochlorite), allowing each aliquot to react away before adding the next one. Use total reflux to maintain acid strength as much as possible. This should dissolve the metal while maximising the efficacy of the oxidiser. But this is nothing like dissolving for instance copper or nickel in nitric acid: it’s slower and more finicky.

I’ve just ordered some pure lead (99.5 % lead) from eBay to see if the alloying elements really have much influence on the rate of dissolution.

blogfast25 - 26-9-2011 at 08:38

I tested the ‘PbO2’ plates and it was quite disappointing. 2 x 5 g of material was subjected to strong H2SO4 and strong HNO3 (ample quantities) in very comparable conditions (back to back). Both react as known with evolution of oxygen, the nitric reacting much more vigorously. But there’s not much PbO2 there, it’s quite superficial and reaction stops after short time, leaving behind a bit of white PbSO4 and dissolved Pb(NO3)2 and much unreacted lead.

blogfast25 - 28-9-2011 at 08:43

A test with about 9 g of lead powder (99.5 %, 300 mesh) with 25 ml of 70 % nitric acid was quite revealing. The powder form makes it an unfair comparison of course but it’s interesting nonetheless. The powder was so reactive the acid had to be added gradually, at least at first. Reaction was vigorous and a whitish precipitate formed almost immediately. The reaction was strongly exothermic (aided also by the formation of a crystalline solid), then dies a little but was resuscitated on heating. I then gradually added small amounts of water, keeping the heat on. The precipitate gradually dissolved until at about 70 ml total volume only some turbidity remained. It hot filtered to clear (#1 filter) easily, then the filtrate started ‘snowing’ white Pb(NO3)2 crystals, even faster when half of the tube was iced. After about 1 h; about 10 g of snow white Pb(NO3)2:

Despite the bias introduced by the powder physical form, I think this conclusively proves that purish lead dissolves much more easily in nitric acid than the lead alloy used in my battery. It’d be worth confirming this with strong HAc + strong H2O2.

sxl168 - 1-10-2011 at 14:01

I was digging for more info and came across this info from Bulletin, Issues 157-161 By United States. Bureau of Mines
PbCl2 Solubility page 20 and onward. Some pretty detailed info for leaching Lead.

I also came across an old text via Google that mentions that the Antimony added to Lead does indeed poison the Nitric Acid when trying to leech Lead. A dictionary of chemistry and the allied branches of other sciences, Volume 7 By Henry Watts
Lead page 728.

Hope this helps out some people. I'm playing with the nearly saturated salt leaching solutions at the moment.

blogfast25 - 2-10-2011 at 05:58

Quote: Originally posted by sxl168

I was digging for more info and came across this info from Bulletin, Issues 157-161 By United States. Bureau of Mines
PbCl2 Solubility page 20 and onward. Some pretty detailed info for leaching Lead.

I also came across an old text via Google that mentions that the Antimony added to Lead does indeed poison the Nitric Acid when trying to leech Lead. A dictionary of chemistry and the allied branches of other sciences, Volume 7 By Henry Watts
Lead page 728.

Hope this helps out some people. I'm playing with the nearly saturated salt leaching solutions at the moment.

The solubility of Pb (II) in function of chloride concentration due to formation of chloride complexes PbCl3- and PbCl42- is shown clearly here:

http://en.wikipedia.org/wiki/Lead#Chloride_complexes

It would suggest higher solubility of lead in high concentration chloride media. PbCl2 has also been reported to be fairly soluble in conc. HCl.

Antimony as a ‘poison’ for nitric in alloyed lead? That would fit some facts here but elemental antimony reacts quite swiftly with nitric acid, forming the insoluble oxide. In an alloy that may be very different, though…

watson.fawkes - 2-10-2011 at 08:41

Antimony as a ‘poison’ for nitric in alloyed lead? That would fit some facts here but elemental antimony reacts quite swiftly with nitric acid, forming the insoluble oxide. In an alloy that may be very different, though…

Sounds like that oxide is forming a passivation layer. Given the different crystal structures of an alloy and of a pure metal, I find it plausible that one oxide form passivates and another does not.

blogfast25 - 2-10-2011 at 08:48

Quote: Originally posted by watson.fawkes

Sounds like that oxide is forming a passivation layer. Given the different crystal structures of an alloy and of a pure metal, I find it plausible that one oxide form passivates and another does not.

But that doesn't explain why aqua regia kind of suffers the same fate: antimony dissolves in AR with great gusto, to full solubility. Also pewter (about 95 Sn/5 Sb) does, without any residue. Both from experience...

The most reasonable assumption thus must be that alloying components like Sb do somehow profoundly alter the dissolution behaviour of lead alloys...

watson.fawkes - 2-10-2011 at 11:51

The most reasonable assumption thus must be that alloying components like Sb do somehow profoundly alter the dissolution behaviour of lead alloys...

I'm thinking by analogy with Cor-Ten steels, in which the surface oxides form a passivation layer, unlike ordinary steels. Admittedly that layer isn't as fully passivating as alumina is, but such a kind of layer would have the effect of slowing dissolution rates, even when the substrate alloy is otherwise reactive. My hypothesis is that it's crystal structures in the body of the alloy that are affecting the way that oxides behave on the surface. This seems very much like a material compatibility issue in semiconductor manufacturing, where lattice spacing is one of the reasons that one material might not adhere to another. As such, drawing the analogy to chemically-related species with different crystal structures doesn't work. In the present case, it' material compatibility between a lead alloy and its oxides. The difference in crystal structure between Pb and a Pb-Sb system is what's at stake. The crystals of Sb or Sn-Sb aren't particularly relevant for examining a change of behavior of the crystals of different lead alloys.

sxl168 - 2-11-2011 at 13:08

I got around to processing some of the positive battery plates today and found out they will convert to PbCl2 quite readily by adding HCl solution to them. Make sure to do this outside as a lot of Chlorine gas is given off and do it in small batches. Given the amount of Chlorine that I saw come off, I'd say this is not a reasonable route to use if processing large batches unless you have a method to catch the Chlorine. After reacting, I had a lot of fine PbCl2, some unreacted PbO2, and small chunks of the metallic grid that held the PbO2 in place at the bottom of my flask. The nice thing that I noticed is that the reaction seemed to keep going even though the solution clearly became PbCl2 saturated. I just made several passes at heating to boiling to decant off the solution, adding water to the solids, and repeating the process until the all of the PbCl2 dissolved out.

You could also probably avoid the whole Chlorine gas mess by heating the battery plates until they converted to red/yellow Lead Oxides. I'll try that when I get another chance.

Question: You could just bubble the Chlorine in cool water and use that Chlorinated water mixed with a dilute acid to dissolve other metals such as Copper, right?

blogfast25 - 3-11-2011 at 07:15

Bit of a useless bugger though, PbCl2, you'd have to convert it to litharge (easy). That chlorine came off means that it's largely PbO2, of course. So it's very wasteful in HCl: half gets blown off as chlorine, the other half you've gotta dispense with. Sulfuric is the better option here (for PbO2), IMHO.

Cl2 in water? A quite weak solution, not realy useful: even if it works on copper (no idea) you'd need tons of it...

S.C. Wack - 3-11-2011 at 15:23

I'd imagine that chlorine hydrate can be made to react with anything.

blogfast25 - 5-11-2011 at 03:56

Is 'anything' not a bit too broad? Certain metals will not be attacked appreciably, others so slowly it isn't a practical method.

Endimion17 - 5-11-2011 at 04:00

Has anyone tried to make very pure lead from the salts obtained? For example, a nice lead plate using electrolysis, for element collection.

blogfast25 - 5-11-2011 at 06:49

Quote: Originally posted by Endimion17

Has anyone tried to make very pure lead from the salts obtained? For example, a nice lead plate using electrolysis, for element collection.

I made lead by electrolysing molten PbCl2 with copper electrodes when I was about 16 but I've no idea how pure it was.

S.C. Wack - 5-11-2011 at 11:49

Is 'anything' not a bit too broad? Certain metals will not be attacked appreciably, others so slowly it isn't a practical method.

Obviously "made to react" is a small, off-topic comment, (I could have added that I also found PbSO4 in quite small crystals covering the electrodes of a very much dying battery, but this was already pointed out) which does not exactly mean adding water to chlorine hydrate, nor was I speaking of some practical amazing all-inclusive method for dissolving every substance known despite the use of anything. It was an attempt to point out that chlorine and water can be concentrated.

It is already known that no one appreciates brevity and hidden layers of meaning.

blogfast25 - 5-11-2011 at 13:52

You're just to clever by half...

sxl168 - 6-11-2011 at 07:26

Quote: Originally posted by Endimion17

Has anyone tried to make very pure lead from the salts obtained? For example, a nice lead plate using electrolysis, for element collection.

For most of my purposes, I don't really need purity, however if I do, I will get Lead dissolved in solution and add a sulfate to the solution to precip. PbSO4. Hot solutions speed this up a lot and produce larger particles. There are very few insoluble sulfate salts, so you should have rather pure PbSO4 after thorough washing. I do believe that BaSO4 is used in small quantities in batteries for seed crystal sites and so that would be the main contaminate. Do note that if sodium is present in moderate or high concentration, this will not work as PbSO4 is soluble in said solutions. From PbSO4 use Sodium Carbonate or Bicarbonate to convert it to PbCO3. From there you can turn it into the product of your choice.

As for plating out Lead, I've done that but I don't know of its purity and it also forms as a loose sponge mass/dendrites that is probably >90% electrolyte. Takes a while to flush out all the electrolyte and dry it before melting. I'm not aware of any method that produces compact deposits outside of a Fluroborate electrolyte. I'd definitely like to know if anyone has gotten a compact deposit besides the use of a Fluroborate electrolyte.

dann2 - 6-11-2011 at 13:53

Quote: Originally posted by sxl168

........... Do note that if sodium is present in moderate or high concentration, this will not work as PbSO4 is soluble in said solutions. From PbSO4 use Sodium Carbonate or Bicarbonate to convert it to PbCO3. ....................

Are you saying that Lead Sulphate is soluble in a solution of Sodium Sulphate?

That does not make sense to me.

Dann2

sxl168 - 7-11-2011 at 08:58

My bad for not proofreading before posting. Meant Sodium Chloride there.

dann2 - 31-5-2012 at 06:32

Interesting method using sugar to dissolve Lead Acid battery parts here:

http://www.sciencemadness.org/talk/viewthread.php?tid=20175

Lead Nitrate can be prepared with Ammonium Nitrate + Litharge (PbO). The PbO is added to a solution of AN and boiled whereby the Ammonia goes away and Lead Nitrate is left.
This was suggested by Plante1999.
PbO can be obtained from battery PbO2 by calcining (heating strongly) (Plante1999)

Dann2

[Edited on 31-5-2012 by dann2]

blogfast25 - 31-5-2012 at 06:44

Interesting method using sugar to dissolve Lead Acid battery parts here:

http://www.sciencemadness.org/talk/viewthread.php?tid=20175

Lead Nitrate can be prepared with Ammonium Nitrate + Litharge (PbO). The PbO is added to a solution of AN and boiled whereby the Ammonia goes away and Lead Nitrate is left.

Dann2

Thanks. I've been meaning to try that!

plante1999 - 31-5-2012 at 08:45

Lead Nitrate can be prepared with Ammonium Nitrate + Litharge (PbO). The PbO is added to a solution of AN and boiled whereby the Ammonia goes away and Lead Nitrate is left.

Dann2

I just make this process yesterday, and posted it in PbO2 anode thread this morning, few hour later the process is already in lead salt thread... Please give your reference next time (the reference is me).

MR AZIDE - 31-5-2012 at 11:32

Quote: Originally posted by garage chemist

PbO2, lead dioxide, is very easily made by electrolyzing 20% H2SO4 with two lead electrodes. At the cathode hydrogen is evolved, at the anode PbO2 is formed as a dark brown powder.
The electrolysis can be continued until the entire anode is converted into PbO2. The PbO2 is isolated by filtering the solution, washing and drying the PbO2.

PbO2 is a powerful oxidiser, capable of forming very energetic pyrotechnic mixes, especially with powdered metals.

Excellent method ...Ill need to try this one........I did have some nice chocolate brown Lead (IV) Oxide from and old accumulator, but this sounds a nice reaction.

Ammonium Chloro Plumbate (NH4)2.PbCl6
Adding PbO2 to ICE COOLED concentrated about (30% ish) HCl, PBO2 dissolves without oxidising the HCl to , to form a lead tetra chloride ligand. PbCL4 -2.

If I remember, the solution turns a yellow colour, but if the temp rises the PbCl2 will crystallize, and the unstable-ish ligand releases chlorine....

I think I remember reading that addition of ammonia/ammonium chloride to the ice cooled, Pb Cl4 complex should precipitate out ammonium chloro Plumbate.

PbO2 + 4HCL ----> PbCl4 + H20

PbCL4 + 2Cl- ------> PbCl6 -2

then

2NH4+ ----> PbCl6 -2 -------> (NH4)2.PbCl6

Says if this is added to ice cool H2SO4 then PbCL4 as a heavy yellow oil, hydrolysed with H2O......

blogfast25 - 31-5-2012 at 12:04

Your method is correct, MR A. The PbO2 must be freshly prepared though. (NH4)2PbCl6 is yellow. (NH4)2SnCl6 also exists.

MR AZIDE - 31-5-2012 at 12:12

The PBO2 I used was from the mesh on an old accumulator, which may have been contaminated with sulphates.

Ive found a big lump of lead sheet I found years ago, ( about 3mm thick and about 40 by 30 cm.
I going to try electrolysing the H2SO4 with Pb electrodes, to see how much PbO2 I can get........

Dont know what fun can be had with the ammonium chloro plumbate yet.

blogfast25 - 31-5-2012 at 13:15

Dont know what fun can be had with the ammonium chloro plumbate yet.

Not much. As you wrote, with cold (IIRW), dry, concentrated sulphuric a dark oil of PbCl4 separates out.

MR AZIDE - 31-5-2012 at 15:21

I HAVE just electrolyzed a strong NaCl solution with 2 lead electrodes, in a 100 ml beaker.

Using a MOT Tx in reverse, so the supply is 24 V.

Now, PbOH is supposed to be produced which is white, but at first, chlorine gas was being emmited, and the solution quickly has turned orange, then a dark brown precipitate has formed, which is settling at bottom of cell.

IS this precipitate likely to be Lead(IV) Oxide, or is it possible that it is a Lead Chlorate, is the Cl gas was mixing with the solution.? What colour are the LEad chlorate's,? Im assuming they are insoluble.

The elecrolysis, has now after a few mins, stopped producing any Cl gas, the Cathode still fizzing away.., and Now white PbOH is being formed at anode, with no gas evolution.
The current is falling due to the inhibition of the insolubility of the coating of Pb(OH)2,
IF the anode is wiped clean, the current is aver 2 Amp but drops to aony about 40mA.

Question is, why is the brown precip, now not happening, and its Pb(OH)2 .??

Its now disconnected, I think Ill repeat this tommorrow in a U-tube to keep the elecrtodes separate, and see what happens.

blogfast25 - 1-6-2012 at 05:54

I HAVE just electrolyzed a strong NaCl solution with 2 lead electrodes, in a 100 ml beaker.

Using a MOT Tx in reverse, so the supply is 24 V.

Now, PbOH is supposed to be produced which is white, but at first, chlorine gas was being emmited, and the solution quickly has turned orange, then a dark brown precipitate has formed, which is settling at bottom of cell.

IS this precipitate likely to be Lead(IV) Oxide, or is it possible that it is a Lead Chlorate, is the Cl gas was mixing with the solution.? What colour are the LEad chlorate's,? Im assuming they are insoluble.

The elecrolysis, has now after a few mins, stopped producing any Cl gas, the Cathode still fizzing away.., and Now white PbOH is being formed at anode, with no gas evolution.
The current is falling due to the inhibition of the insolubility of the coating of Pb(OH)2,
IF the anode is wiped clean, the current is aver 2 Amp but drops to aony about 40mA.

Question is, why is the brown precip, now not happening, and its Pb(OH)2 .??

Its now disconnected, I think Ill repeat this tommorrow in a U-tube to keep the elecrtodes separate, and see what happens.

Did you mean Pb(OH)2? The precipitate is PbO2. Lead chlorate would be white, I'm quite sure of it.

MR AZIDE - 1-6-2012 at 10:10

had a bit of difficulty in looking up lead chlorates/ perchlorates, but they are freely soluble in water.

I think what happened is the NACL was concentrated enough, that Pb(OH)2 was formed, and immediately reacted with the Cl gas, maybe changing the Pb(OH)2 to PBO2 .
I ve thrown it away now anyway.

I had much more interesting results electrolysing some HCL with Pb electrodes today, which I may post in the technochemistry bit.

here www.sciencemadness.org/talk/viewthread.php?tid=20284

[Edited on 1-6-2012 by MR AZIDE]

blogfast25 - 1-6-2012 at 12:58

I think what happened is the NACL was concentrated enough, that Pb(OH)2 was formed, and immediately reacted with the Cl gas, maybe changing the Pb(OH)2 to PBO2 .
I ve thrown it away now anyway.

[Edited on 1-6-2012 by MR AZIDE]

Check the half-potentials to see if:

Pb2+ + Cl2 === > Pb4+ + 2 Cl-

... is possible. Dunno off the top of my head, TBH...

Edit: I checked. This oxidation appears not to be thermodynamically favourable, chlorine isn't strong enough. But hypochlorite is and some of that may have formed with your electrolysis.

[Edited on 1-6-2012 by blogfast25]

MR AZIDE - 2-6-2012 at 10:21

as long as the Cl2 gas was liberated , the probable Pbo2 was produced, as definately hypochlorite will be being produced.......this oxidised the lead to from +2 to +4, .......

PbO2, Ive just read can be produced by shaking a solution of a hypochlorite with a suspension of Pb(OH)2 in water.

Lambda-Eyde - 2-6-2012 at 10:29

I had a paper on producing Pb(AcO)4 from Pb3O4, acetic acid and acetic anhydride. The yield could be maximized by oxidizing the remaining Pb2+ with chlorine gas IIRC.

blogfast25 - 2-6-2012 at 10:46

PbO2, Ive just read can be produced by shaking a solution of a hypochlorite with a suspension of Pb(OH)2 in water.

No need for shaking

, see the other thread on battery lead and my results. Hypochlorite easily oxidises Pb(OH)2 to PbO2.

[Edited on 2-6-2012 by blogfast25]

MR AZIDE - 2-6-2012 at 11:24

I see your source of NAOCl, was bleach.......is this like the cheap stuff you can buy in supermarkets.?

Ive seen like the cheap own brand stuff in the supermarket, and I read the labels just says '' contains sodium hypochlorite'', but it never says what concentration the sodium hypochlorite is.
Ive never got round to buying any to play with.

are most of these thick /thin bleaches( as they call it,) basically NaOCl in plain water?,,,, are there many additives........is the solution colorless of slightly yellow?

If so Id like to try at some point making KClO3 via nurdrage's method........

blogfast25 - 2-6-2012 at 12:10

I see your source of NAOCl, was bleach.......is this like the cheap stuff you can buy in supermarkets.?

Ive seen like the cheap own brand stuff in the supermarket, and I read the labels just says '' contains sodium hypochlorite'', but it never says what concentration the sodium hypochlorite is.
Ive never got round to buying any to play with.

are most of these thick /thin bleaches( as they call it,) basically NaOCl in plain water?,,,, are there many additives........is the solution colorless of slightly yellow?

If so Id like to try at some point making KClO3 via nurdrage's method........

Get 'no frills' thin bleach: 4 - 5 % NaClO (sodium hypochlorite). Weak concentration, strong oxidiser. Ca(ClO)2 ('pool shock') is 100 % (solid) but harder to get.

Thick = thin + glycerine. Avoid!

Fantasma4500 - 2-3-2013 at 13:12

Pb(ClO4)2 easily made

i happen to have access to NH4ClO4
aswell as i can make PbCO3 very easily (lead acetate + sodium carbonate, filter)
i cant really spot anything that should make this reaction impossible (NH4ClO4 + PbCO3 > (NH4)2CO3 + Pb(ClO4)2)
what i like about this is that NH4ClO4 is by what i understand alot more safe to store and alot easier to buy, PbCO3 is pretty easy to make, and most of all, the (NH4)2CO3 should decompose at ~60*C IN SOLUTION
i havent gotten around to do this.. yet...
i assume that it could possibly be explosive based on how Pb(ClO2)2 acts under heating, and also the fact that you shouldnt ever let a solution of Pb(ClO4)2 dry out (small amounts, more control, more knownledge)
and if it wouldnt be explosive by itself, it should act as a brilliant oxidizer..

Xenoid - 2-3-2013 at 14:23

Quote: Originally posted by Antiswat

.... i cant really spot anything that should make this reaction impossible (NH4ClO4 + PbCO3 > (NH4)2CO3 + Pb(ClO4)2)

Well... other than the fact that lead carbonate is virtually insoluble (.00011g/100ml) that is...

Eddygp - 2-3-2013 at 14:30

Does anyone have an interesting way to start with lead(III) ions?

12AX7 - 2-3-2013 at 14:40

Do you mean Pb(II) or Pb(IV)? And what are you starting exactly?

Lead doesn't have a +3 (unless there's some strange complex which exhibits it, but that's practically cheating).

Tim

Eddygp - 2-3-2013 at 14:48

yes, I meant lead +3. Pb(III), and I know it is not common, but I thought there would be some sort of temporary stable compounds like hypomanganates, which aren't always stable.

blogfast25 - 3-3-2013 at 05:42

AFAIK, no one here has attempted anything re. Pb (III).

Lead Dioxide

KernelPicnic - 23-3-2013 at 19:08

This afternoon I prepared some lead dioxide on a 100 mmol scale.

33g lead nitrate was dissolved in about 80mL warm distilled water. A mixture of 150mL 6% sodium hypochlorite and 25mL 10M sodium hydroxide was added at once. Instantly, a bright orange precipitate was formed, which slowly began to darken in color. The reaction mixture was heated just below boiling for about 15 minutes, until the precipitate turned a uniform dark brown. This was then gravity filtered and washed with ~400mL of tap water. The filter cake was dried with a desk lamp for about four hours. Yield was 21g, or 88%.

blogfast25 - 24-3-2013 at 06:07

Nice! With freshly prepared PbO2 you can make the interesting salt (NH4)2PbCl6 (ammonium hexachloroplumbate) which no one here seems to have attempted yet. Adding chilled conc. H2SO4 to it liberates PbCl4 as a dark, heavy, oily liquid.

Heuteufel - 24-3-2013 at 06:55

Well, I once made pyridiniumhexachloroplumbate (which is very similar to the ammonium salt) by saturating a suspension of lead(II)-chloride in HCl with chlorine and precipitating the product with pyridine. Then I decomposed the pyridiniumhexachloroplumbate to PbCl₄ with sulfuric acid: Lead Tetrachloride
I don´t remember a procedure, that uses lead dioxide.

Something interesting might be to attempt a synthesis of lead tetraacetate with the PbO₂ (usually red lead is used) and to use the product afterwards to do a cleavage of an 1,2-diol, like for instance: Synthesis of n-butyl glyoxylate

blogfast25 - 24-3-2013 at 07:06

Quote: Originally posted by Heuteufel

I don´t remember a procedure, that uses lead dioxide.

The following is from a *.pdf I picked up (I don't recall the url):

(b) Synthesis of (NH4)2PbCl6 from PbO2. It is imperative that the following reactions be performed in icecold
solutions. In a 50 mL beaker add a volume of saturated NH4Cl solution (10 mL), and in a 10 mL
beaker add concentrated (care!) HCl (5 mL). Cool both solutions thoroughly in ice. With the beaker
containing the HCl still in the ice bath, slowly add PbO2 (prepared above), with constant stirring (glass
rod); a yellow solution should be obtained. In some cases, a small amount of white precipitate will appear
at this stage; this does not influence the outcome of the reaction, however note its appearance and keep it
in mind when conducting the tests that follow. In the event that some white precipitate does appear after
adding all of the PbO2, leave the yellow solution in the ice bath for a few minutes so as to allow it to settle.
Next, quickly but carefully decant the yellow solution into the ice-cold NH4Cl solution (still in the ice
bath) with stirring, making sure to leave behind any white solid that may have previously appeared. Allow
the mixture to sit in the ice bath for several minutes in order to complete the precipitation, and then filter
the (NH4)2PbCl6 product (Buchner) from the mixture. Do not wash the product with water as it is very
soluble in this solvent. Allow the product to air dry (they may be stored in the locker until the next
laboratory session) and record the yield.

Fresh PbO2 is apparently quite soluble in concentrated HCl.

But your salt is interesting too.

[Edited on 24-3-2013 by blogfast25]

Heuteufel - 24-3-2013 at 12:44

Hmm, interesting! I once made lead dioxide from lead(II)-nitrate using sodium persulfate as an oxidiser (I prefer this method to the hypochlorite method) and I remember, I had serious problems cleaning the glassware particularly the glass frit I used for filtration. Cold HCl (37 %) had some effect, but not much, heating the acid helped to some extent...
Therefore I am somewhat surprised, but I don´t say, that the procedure doesn´t work. I must also admit, that I did the synthesis some time ago and thus my memories are somewhat vague...

DraconicAcid - 25-3-2013 at 08:45

Quote: Originally posted by Eddygp

yes, I meant lead +3. Pb(III), and I know it is not common, but I thought there would be some sort of temporary stable compounds like hypomanganates, which aren't always stable.

I'd be *very* surprised if you could make a stable lead(III) compound, as that would be a 6s¹ electron configuration. The s electrons (unlike the d electrons of manganese) usually come off as pairs.

blogfast25 - 25-3-2013 at 10:46

I'd be *very* surprised if you could make a stable lead(III) compound, as that would be a 6s¹ electron configuration. The s electrons (unlike the d electrons of manganese) usually come off as pairs.

Sure, but there are plenty 'exceptions' to these simple 'rules'. There are electron deficient compounds too, see boranes e.g. But Ive never heard of Pb(III).

DraconicAcid - 25-3-2013 at 12:00

I'd be *very* surprised if you could make a stable lead(III) compound, as that would be a 6s¹ electron configuration. The s electrons (unlike the d electrons of manganese) usually come off as pairs.

Sure, but there are plenty 'exceptions' to these simple 'rules'. There are electron deficient compounds too, see boranes e.g. But Ive never heard of Pb(III).

I suppose a diplumbane such as hexamethyldiplumbane would technically be lead(III), it's not an ionic compound. Neither Cotton & Wilkinson nor Greenwood's Chemistry of the Elements mention lead(III) as a possible oxidation state (although they do mention the Zintl-type anions such as Pb₅^2-.

blogfast25 - 25-3-2013 at 13:00

In hexamethyldiplumbane each Pb atom would share four electrons with other atoms: that's oxidation state IV by definition. The inter-Pb bond would still be an orbital with shared elctrons.

[Edited on 25-3-2013 by blogfast25]

DraconicAcid - 25-3-2013 at 13:45

I think that would be a valence of four by definition, but not an oxidation state.

Mailinmypocket - 26-3-2013 at 12:32

Quote: Originally posted by Heuteufel

I don´t remember a procedure, that uses lead dioxide.

It would seem that fresh PbO2 is not strictly required... I used an old commercial source of it and followed the instructions you quoted above, with nice results, just for fun

Dissolved 400mg PbO2 in 5ml conc. HCl, nice yellow solution was produced.

After dumping into ice cold saturated NH4Cl

Filter cake...

Product drying in a dish, final weight to be determined...

blogfast25 - 26-3-2013 at 13:20

Well, well, well, what a lovely confirmation there, MIMP, well done!

The yellow solution (second photo) must essentially be H2PbCl6 with excess HCl.

Now if you put some chilled conc. H2SO4 on dry, cold (NH4)2PbCl6 you should observe a dark oily liquid: PbCl4. Careful with that, I'm not sure how stable it is!

The homologous (NH4)2SnCl6 also exists and is easy to synth too. Named 'pink salt' it was once an important dye mordant for bright pinks, like Pink Madder or Cochenille.

[Edited on 26-3-2013 by blogfast25]

Mailinmypocket - 26-3-2013 at 13:30

Well, well, well, what a lovely confirmation there, MIMP, well done!

The yellow solution (first photo) must essentially be H2PbCl6 with excess HCl.

Now if you put some chilled conc. H2SO4 on dry, cold (NH4)2PbCl6 you should observe a dark oily liquid: PbCl4. Careful with that, I'm not sure how stable it is!

Well, I figured it would give me a reason to use my prehistoric bottle of PbO2! I'm thinking the same as you regarding the yellow solution, there was also quite a bit of chlorine gas produced during the addition of PbO2 to the HCl but not intolerable in the small amounts used. The preparation doesn't specify how much PbO2 was produced in "the previous experiment" and I got a little too spatula-happy and added too much. This resulted in a mucky dark suspension, so I re started the experiment. In retrospect I could have just added more HCl but had a dumb moment and didnt think to do so... Fail.

Otherwise it's pretty interesting to watch, when the PbO2 is added there is a vigorously fizzing dark liquid, and all of a sudden it clears to bright yellow and you can continue adding in small amounts. Fun!

Once the product is completely dry I am definitely going to try the H2SO4 carefully, and take some pictures of the results.

*edit* just saw your comment about the tin pink salt, I will definitely give that a try too. I love making inorganic salts for some reason

[Edited on 26-3-2013 by Mailinmypocket]

blogfast25 - 26-3-2013 at 13:42

Yes, these salts are quite satisfying to make. On the synth of 'pink salt' there's a thread of mine here somewhere. Search for 'hexachlorostannate', if you're interested. It starts from pewter but you can start from tin or SnCl2 too.

KernelPicnic - 26-3-2013 at 15:31

I attempted to make K2PbCl6 using the same procedure for the ammonium salt, except obviously with saturated KCl. Nothing precipitated when I added the PbO2/HCl, and only after sitting in the ice bath for about 30 minutes did a small amount of what looked like fine yellowish crystals drop out of solution. It was kind of hard to tell since the color of the liquid was still bright yellow. I'm thinking the precipitate might just be PbCl2 from partial decomposition. I'll try the ammonium salt whenever I get some NH4Cl.

[Edited on 27-3-2013 by KernelPicnic]

Mailinmypocket - 26-3-2013 at 16:39

Quote: Originally posted by KernelPicnic

I attempted to make K2PbCl6 using the same procedure for the ammonium salt, except obviously with saturated KCl. Nothing precipitated when I added the PbO2/HCl, and only after sitting in the ice bath for about 30 minutes did a small amount of what looked like fine yellowish crystals drop out of solution. It was kind of hard to tell since the color of the liquid was still bright yellow. I'm thinking the precipitate might just be PbCl2 from partial decomposition. I'll try the ammonium salt whenever I get some NH4Cl.

[Edited on 27-3-2013 by KernelPicnic]

If you can get some ammonia solution somewhere you could react that with the HCl you have and dry it out to get NH4Cl crystals

KernelPicnic - 26-3-2013 at 16:54

Quote: Originally posted by Mailinmypocket

If you can get some ammonia solution somewhere you could react that with the HCl you have and dry it out to get NH4Cl crystals

I've done that before, but with generic household ammonia the product turned out slightly yellow. I tried subliming it, but that never worked out very well, and the result just looked worse, so I ended up throwing it away. Whatever impurities may have been in there probably wouldn't make a difference anyway now that I think about it.

PbCl4

Mailinmypocket - 28-3-2013 at 12:23

Just for a small test I added some conc. H2SO4 to a few small granules of (NH4)2PbCl6. There was an immediate reaction, and a cloudy yellow mix was produced. The quantities used were so small that it was very difficult to see any oily substance, when I took a small sample out of the tube there was obvious fuming in the air. I will be scaling it up now that I was able to see that the reaction is not terribly violent. Just need to wait to have my fume hood back up and running before doing that.

Apologies for the small scale- it's a bit hard to see anything interesting and I wanted to have an obvious sample to show, but you get the idea! Ill attach a video of the fuming nature of the liquid later on once YouTube "approves" it.

blogfast25 - 28-3-2013 at 13:36

Water and sulphates may be the enemy of this reaction, though. Water because PbCl4 hydrolyses easily, sulphates because of insoluble lead sulphates. Better to 'keep your powder dry', as they say!

Poppy - 29-3-2013 at 17:39

Nice! With freshly prepared PbO2 you can make the interesting salt (NH4)2PbCl6 (ammonium hexachloroplumbate) which no one here seems to have attempted yet. Adding chilled conc. H2SO4 to it liberates PbCl4 as a dark, heavy, oily liquid.

Do you have a note on how long PbCl4 survives hermetically without overwhelming the jar?

blogfast25 - 30-3-2013 at 05:17

Quote: Originally posted by Poppy

Do you have a note on how long PbCl4 survives hermetically without overwhelming the jar?

AF Holleman claims that at room temperature it slowly decomposes to PbCl2 and Cl2.

I carried out my own synthesis of (NH4)2PbCl6 yesterday, this one starting from Pb(NO3)2 (oxidising it to PbO2 with thin bleach in very alkaline solution). Although I obtained the salt, yield was lower than expected because when I dissolved the PbO2 in conc. HCl (iced) quite a bit of white precipitate formed. This is due to partial oxidation of the HCl to chlorine (which I could smell) by the PbO2, I believe :

PbO2 + 4 HCl == > PbCl2 + Cl2 +2 H2O

If I find time this afternoon I might adapt the addition procedure to minimise this loss to PbCl2.

[Edited on 30-3-2013 by blogfast25]

blogfast25 - 1-4-2013 at 06:18

An attempt at adding ice cold conc. HCl drop by drop from a 10 ml burette onto an iced, concentrated slurry (in water) of PbO2 didn't improve yield of dissolved PbCl4 (as PbCl6(2-)) by much. Right from the start chlorine gas evolution could be smelled and at the end of the addition some white precipitate (PbCl2) was obtained. Subjectively I'd say yield of dissolved PbCl4 was about the sme as during the first experiment.

Also very subjectively I'd say about 30 % of the initial PbO2 is lost to PbCl2.

[Edited on 1-4-2013 by blogfast25]

KernelPicnic - 2-4-2013 at 11:54

Off topic, but yesterday I made about 30g of lead bromide. Apparently, the stuff is pretty light sensitive. When I finished drying the product, I noticed the powder adhering to the filter paper had turned dark grey. I also left the bottle near a window overnight and a thin layer had turned black. Kind of interesting since the chloride is not photosensitive at all.

blogfast25 - 3-4-2013 at 03:16

Lead bromide isn't known for being light sensitive, as far as I know. How did you prepare it? Is it possible that it's contaminated with silver bromide?

[Edited on 3-4-2013 by blogfast25]

KernelPicnic - 3-4-2013 at 03:30