There are a few things you can use lead for, well not really, I mainly want it in two forms: PbO, litharge, for a flux for ceramics and glass, and PbO2, lead dioxide, for making perchlorates by electrolysis. Impractical uses include making various salts for the heck of it, which I am also all for doing! :)
First off, you need to get lead metal. It's pretty cheap, and if not bought, can be found as wheel weights (with about 2.5% Sb), plumber's lead (if you know of a house being remodelled that has cast iron pipes, the joints are packed with 99.9% Pb) and a number of other things. Once you've got the metal, you need to oxidize it. Lead is pretty close to hydrogen (at -0.126V standard reduction potential), so you can't really dissolve it in acid, besides which most acids are inert to it! A good solution is electrolysis, but you need an electrolyte which ionizes well and forms a soluble lead salt (sodium sulfate and cold chloride need not apply). I haven't tried nitrate yet, because I don't have any. Chlorate and perchlorate work (I have tried chlorate), and acetate as well.
The above picture is with sodium acetate solution, made by neutralizing vinegar with baking soda. It isn't very conductive because acetate doesn't ionize well and the solution is weak concentration. The brown/white/gray tab floating in the upper-right is a piece of calcium sulfate as a buffer (any Pb++ ions in solution will react with SO42- ions and precipitate lead sulfate) to attempt to prevent metal plating through. Nonetheless some metal ions do make their way to the cathode and form a foamy lead sponge deposit which must be removed periodically. Lead hydroxide (among other things, such as the PbSO4 I mentioned) fall to the bottom to be washed later.
Lead sulfate is pretty insoluble, but lead carbonate is even less soluble, with the plus that it is more reactive. To remove impurities, I have to do a bit of conversion here, carbonating it to start. After boiling it for half an hour (in a 300ml flask, one of my few pieces of real labware!) with sodium bicarbonate solution, I decant and wash the precipitate.
Since lead chloride is almost entirely insoluble, while impurities such as sodium and iron are soluble, I can add muriatic (hydrochloric) acid until it stops bubbling, then wash the lead chloride.
What sets PbCl2 apart from say, PbSO4 is it is marginally soluble in boiling water. The center jar is crude PbCl2 while crystals are forming in the other two jars as they cool. After several treatments, all the lead chloride has been recrystallized.
...And here it is, drained and dried. The crystals are colorless flat needles, more like elongated, irregularly shaped plates. It is soft and slippery, almost like handling a fine fur. This was...155 grams.
After weighing a stoichiometric amount of caustic lye (NaOH) and precipitating (with a small excess of NaHCO3), I am left with a slightly greenish (possible iron impurity?) white lead hydroxide sludge to dry and calcine for PbO.
The old, and better, method involves burning lead in air, producing the oxide directly (purity depending on the metal, which is quite easy to get in more purity than I need).
So far my attempts have amounted to melting a pound of lead in a steel can (lead has no solubility for iron so it will not dissolve through, as zinc and aluminum will!), heating to redness (about 1100°F) and stirring for half an hour. At this temperature, the lead forms an oxide skin, which starts transparent, going through yellow, blue, purple, then metallic shades in a psychadelic wash as exposure goes on. In half an hour I collect maybe 1/2 pound of powdered, green material (which is yellow to red when hot). My problem is it doesn't seem to be doing anything! I once weighed 600 grams lead, poured the metal out from the oxides, weighed it at 435 grams, and assuming PbO stoichiometry, I should've had around 15 grams weight increase. So I put 165g on the balance, assuming it would fall over when the oxide was added ... and it leveled. Puuuuurfectly. WTF! I have also attempted to react it with hydrochloric acid, resulting in a little initial heat, after which it bubbles to (so far) no end, indicating lead metal. Apparently... I have lead droplets here, coated with a thin layer of lead oxide and/or suboxide (the suboxide is green, explaining the color). No amount of further calcining appears to be oxidizing it deeper. If anyone has experience in this subject, please e-mail me.
A past yield of lead oxide, obtained by reacting lead metal with boiling HCl and also electrolytic preparation with KClO3 solution (which is as slow as the sodium acetate above due to potassium chlorate's low solubility). This appears to have some amount of iron impurity, as some I used for preparing lead acetate left a red to brown syrup on crystallizing the solution. A later attempt produced a green solution, which on concentrating turned blue (that is, removing the lead acetate, which was itself stained green even in thick crystals).
Lead dioxide is in the +4 state. Not usually too easy to get to; you can reach a half-and-half compromize (red lead) pyrolytically, well somehow anyway. I've roasted pure lead hydroxide to an orange to pinkish red color before, but never a pure Pb3O4 state, and never starting from just lead metal. PbO2 is impossible through roasting with air as far as I know, although it might be possible in pure oxygen. It has to be oxidized in solution, for example by taking a lead salt like the acetate and adding an oxidizer like sodium or calcium hypochlorite, giving the reaction Pb(CH3COO)2 + 2NaOCl + H2O = PbO2 + 2NaCl + 2CH3COOH. (The typical basicity of bleach solutions provides something of a buffering action, combining with some of the freed acid; using a weak acid such as acetic I'm sure helps too.)
But this says nothing of this picture. Here I believe was a cool sodium chloride solution; a boiling solution produces reasonable amounts of Pb(OH)2, but without sufficient current flow to warm it up, it cools down and solubility for the anode's PbCl2 intermediate drops, slowing it down and ultimately, layers of PbCl2 and Pb(OH)2 passivate the anode. Some current still gets through (I'm guessing the Pb(OH)2 is slightly conductive), oxidizing it to PbO2, which forms a loose skin which is somewhat resistant to chloride but nonetheless flakes off. This layer is very resistive so is of no use to me. Sulfate solutions (including battery acid; in fact lead-acid batteries store energy as PbO2 and PbSO4 on the respective electrodes) also anodize a layer of PbO2, but not efficiently (much oxygen is evolved at the anode) and it too flakes off when used in a chloride solution.
Supposedly, electrolysis of a lead acetate solution is able to plate PbO2 on the anode, but the only result I've ever had was exactly what you'd expect: lead crystals growing off the cathode. However, an acidic lead nitrate solution (with some copper nitrate) *is* able to plate large, solid, continuous, stable layers of lead dioxide suitable for long-term use in electrolytic production of chlorate and perchlorate. I have *no* idea what keeps it from plating metal instead, all I know is, I have absolutely no nitrates around to do it!
The oxide section covers one route of PbCl2 synthesis. If you want some on hand, you can always take some lead oxide and react it with hydrochloric acid, or boil lead metal with acid. It will form a pasty suspension of PbCl2 which settles quickly and is insoluble. You can purify by washing (to remove solubles) and recrystallizing (to remove less soluble material such as excess Pb or PbO). It melts at a relatively low temperature and if some excess remains on your lead metal, when remelting, it forms a convienient flux which cleans the melt. Although it has a relatively high vapor pressure as chlorides go, it is lower than lead metal. Melt lead chloride outdoors or under a fume hood!
My lead(II) acetate, contaminated possibly with red to brown iron acetate (as mentioned above), which so far has refused to crystallize despite placing it in my MgSO4 dessicator, instead preferring to be an annoying stickiness between the hairy, sweet(!) lead acetate crystal lumps.
This can be made in two ways: one, by reacting lead oxide with acetic acid directly, or by cementating copper acetate with lead metal (displacing copper metal, which flakes off). This alone is kind of interesting because, if the solution is left undisturbed for a while, the lead reacts with only what copper ions can diffuse to it. The lead-heavy solution then remains on the bottom, with the metal. What happens is you get a blue copper layer floating on a clear lead layer!
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. The above material was also added, since by this time (hmm last update 8/8/05, today is 5/30/06, this was some weeks ago) it had given off a lot of acetic acid and was probably basic lead acetate anyway. Not to mention the sticky brown crud that had by then solidified the clumps into a hard cake. This product has been recrystallized, so much of the brownness has been removed. I may recrystallize it a third time to get a clean white product with bigger crystals. At the moment, this granular powder is free-flowing, but it cakes up after a few days and needs a shake. It still smells of acetic acid, so I should probably seal it up rather than let it continue to dry.
Here are the largest, clearest, most beautiful lead acetate crystals I have yet made. Maybe 25 grams worth. They are still slightly green, but white for the most part.