Arthur Dent

Look for "Cameo" anti-tarnish cleaner. At first it looks like plain Comet or Ajax scrubbing powder, but it has this strange licorice aroma and after scrubbing it on steel with a wet sponge for 2 or 3 minutes, it does wonders even with the most deeply stained sinks. I had a sink that was utterly destroyed by corrosive reflux from the drain that did its work during a whole weekend... well that Cameo powder made it look nice and shiny! It's about $3 to $5 a bottle and is available at many drugstores and hardware stores. Robert

I'll refer you to this thread: http://www.scienceforums.net/forum/showthread.php?t=6679 Where you'll find resources of companies that specialize in selling elements of the periodic table. There are also members on this thread who collect such elements and might have "spares" for trade and sale. Robert

Thanks jdurg, I'll take this into account, although Lithium Nitride decomposes instantly when in contact with water and generates ammonia gas: Li3N + 3 H2O → 3 LiOH (aq) + NH3. Been busy for the past few days so I never got to the final step of turning the LiOH into a nitrate. I'll do that during the weekend. My thoughts were to dilute my HNO3 with water so that the reaction isn't too intense, and add a small surplus of acid so that all of my LiOH turns into LiNO3. I want to have relatively pure stuff to work with. Robert

With the advice above, I have just prepared the LiOH. Prepared a becher with 80 ml of distilled water and i went outside in my backyard and started dropping tiny chunks in the water (about 2mm x 2mm). It is indeed extremely exothermic, ant my first attempt resulted in the first chunk rapidly clinging to the side of the becher where it caught fire (nice red flame) and melted and cracked the Becher's side. Aww, dammit, I loved that 250 ml becher... So I emptied the content in a 500 ml becher and used a different technique of stirring the becher in a slow motion after dropping the pellets, It prevents the metal chunks from clinging to the side and prevents the hydrogen from catching on fire. One interesting note is that as I was dropping pellets, they were dissolving slower and slower, until the last one that took over 90 seconds of stirring before disappearing completely. The resulting solution looks very dense and slightly syrupy. Even though the metal was very blackened, it left no deposits at the bottom of the becher and it looks perfectly clear. I'll let the solution cool down (It got really really hot during the process), and i'll try to add the HNO3 later. Again, I imagine that the acid and the base together will create a highly exothermic reaction, so I think I'll put the becher in a tray filled with water and ice chunks. Should I use the pure HNO3 (42 BE) and pour it drop by drop or should I dilute it 50% with distilled water? Robert

Excellent, thanks! I'll just gently throw little bits of lithium in water until I have a fairly concentrated LiOH solution, i'll leave it to settle so that any traces of mineral oil come up to the surface, and then add nitric acid (I guess until the solutions becomes neutral to slightly acidic). Then i'll dessicate the solution. Robert

Hi, I recently acquired some metallic lithium stored in mineral oil. I was wondering about two things concerning this reactive metal... 1) Before using the metal to prepare a Lithium salt, I do need to remove the mineral oil on its surface (as not to contaminate my preparation). Is there a solvent I can use? Even just in air, it turns nearly black in a matter of seconds. 2) Just dropping the metal in water will result in creating Lithium Hydroxide, but how do I go about creating either Lithium Nitrate or chloride? I thought of just dropping very very small quantities at a time directly into Nitric acid or Hydrochloric acid, but i'm too scared about the intensity of the exothermic reaction with concentrated acids. 3 is there a more passive way of producing these chemicals? Can I use the LiOH to produce LiCl or LiNO3? Thanks all Robert

Correct me if i'm wrong, but that would be tectosilicates, which makes up a major part of the earth's surface, and includes quartz, feldspath and a good majority of the rocks we step on.

Thanks for the advice... I do have HNO3 but don't want to create a mixture of nitrate and chloride so I guess i'll let things simply sit. My goal would be to have relatively pure stuff, so I'd like to avoid using chemicals that would contaminate my final product. @rogerxd45, I did something similar to your watchglass protecting the mixture, I used a piece of duct tape (I'm canadian, heh!)... As for buying Nickel Chloride, I do have a nice chem provider in my neighborhood (Prolab Scientific) but the prices are a bit high and I just like the challenge of making my own stuff up! (especially for small quantities) I usually order stuff from ProLab that I can't make myself and is relatively inexpensive. Robert

Hm, okay then. I have to admit I'm always a bit weary of heating-up strong acids on a hotplate (I do have one for outside use), but I thought maybe if I added a bit of H202, it would accelerate the process. So i'll try just heat this weekend, hoping to keep the reaction under control. After two days, the becher exhibits a very pale greenish color, and there are quite a few bubbles on the nickel itself. Robert

In my never-ending quest to prepare inorganic salts for my various experiments in electroplating, I was attempting to prepare a small amount of Nickel Chloride by using shards of 99.9% nickel with concentrated HCl, but I guess that what's missing is some sort of oxidizer because the reaction is nil... nothing, nada! There's some very weak bubbling after an hour, but nothing to write home... no "green" color at all. A while ago, I had prepared succesfully some Zinc Chloride using the same basic process, but in that case, the reaction was vigorous and when it subsided, I had some relatively pure ZnCl2. So here I have a becher with a few grams of nickel and about 20 ml of HCl (20 BE). Is there a reagent that I can safely and carefully add to the content to kickstart the reaction? Robert

Yes. http://www.elementsales.com/pl_element.htm#br Robert

I remember reading that a Titanium/Zirconium alloy was one of the strongest alloys possible, in hardness and tensile strenght. After googling it, I see it's used extensively in medicine for implants. And I guess that some ceramics, like the ones used in hi-end bulletproof vests, are extremely tough. But of course, a material that would exhibit the properties of our friend Wolverine's claws doesn't exist, although if it did, it would save a lot of lives. Robert

I will use Zinc Chloride to test for the presence of gold in solutions. Thanks for your suggestion, however liquid nitrogen is not something I can easily access nor do I want to play with, even though I know of a distributor in my neighborhood. But maybe dry ice (CO2) would work?

I just tested the dessication method using CaCl2 in the form of a household "damp-trap" box sealed in with an amount of Zinc Chloride solution sealed in a ziploc bag... The ZnCl2 won... after a full week in the bag, there is still 30 ml of solution and it hasn't dropped even 1 ml, it doesn't work at all. Or am I too hasty? Anyway, I just took my becher with 50 ml of Copper Sulfate solution, who already has beautiful spiky crystals growing at the bottom and see if this damp-trap process will dessicate the solution faster. I'll use a hotplate to dry up the ZnCl2 instead. - Robert

@ John Cuthber: When you mentioned CaCl2, I remember a long time ago that I use some of these inexpensive "Damp Trap" boxes in my basement a few years ago. The only ingredient in there is anhydrous Calcium chloride sealed in a plastix box. Once you open the seal, it does a pretty good job at removing moisture. I went to the dollar store and grabbed a few of those boxes to experiment with. Below is a "damp-trap" sealed in a ziploc bag with a few ml of Zinc Chloride solution... i'll let 'em fight it out for a few days and we'll see who wins! And I guess that I can also recycle the CaCl2 when it's wet by heating it up a bit. I'll keep you posted if it works. I'll aslo try it on my solution of CuSO4 to see how fast it works. Robert Robert

I wonder if the dessication method would work with the OP's green goo? What if the syrupy CuCl2 in a becher was sealed-up in a ziploc bag with a becher of anhydrous Sodium Hydroxide, or even better, a couple of those small "dessicant" bags found in electronic gear boxes? Or is the CuCl2 much more hygroscopic and it would do the reverse process? - Robert

Oh yeah, NO2... sorry I confused with N2O. Good advice!, i'll chill the solution and retrieve the crystals as they form. I don't have any Copper Hydroxide or Carbonate but I guess I can sacrifice a little bit of the solution to have the purest CuSO4 possible. Crystals are already starting to grow on the lattice of copper wire (I left the solution outside overnight so it gasses off). Many thanks! Robert

Hi all, I just prepared a batch of Copper Sulfate... I have prepared a solution of 60ml distilled water, 6ml HN03 and 10ml H2S04 and put a big ball of fine telephone copper wire (unsheated). The reaction was energetic and produced a beautiful blue solution. Now it has settled and there is a bit of copper left, but the solution is deep blue. Even though the HNO3 is the oxidant and most of it escapes as it turns into nitrous oxide, can there be some Copper Nitrate impurities in my CuSO4 ? And now that the solution is settled, I know that there is a bit of sulfuric acid left in the solution, what's the easiest way to get rid of the H2SO4 to get clean crystals? Thanks Robert

Thanks for the info, that makes sense. I also read that pure zinc chloride solution is acidic (a 6M solution has a ph of 1!). I added a few thin strips of Zn to see it it would react and it doesn't, so I think i'm pretty close to a good solution of ZnCl2. And after evaporation and dessication, I figure the remaining traces of HCL will evaporate too. My next task will be to filter out the solution first to get the tiny bits of undissolved zinc, zinc oxide and gunk. Then on to the hotplate! - Robert

Hi, Quick question about the preparation of Zinc Chloride... I have been attempting to prepare a small batch of ZnCl2 using zinc cut up in small strips and some HCl (25ml of 32% HCl and 25ml distilled water) in a becher. Zn + 2 HCl → ZnCl2 + H2 This is my first attempt at this and at first, the reaction was quite energetic, so I figured that when it would subside, the acid would be neutralized and the result would (mostly) a pure zinc chloride solution. But 24 hours after the beginning of the process, a strip of ph paper indicates that the content is still strongly acidic, even though the remaining little pieces of zinc have stopped bubbling. Should I gently heat the solution to complete the process and get the ph down? Knowing that ZnCl2 is very hygroscopic, is there any way to keep it in solid form, even if a bit "wet"? Will a glass bottle with a teflon seal be enough to keep the moisture to a minimum? Thanks in advance for your input

Maybe there are some nickel impurities in your HCl... If I recall, nickel chloride is a yellowish-green salt. The RONA brand HCL I buy at the hardware store in my area is perfectly clear, 20BÉ , and comes in a tightly sealed plastic bottle in a polythene bag. EDIT: mmh. just saw you posting saying the HCl was clear afterall. Maybe it was a reflection from the label or something...

Hi folks, Glad I found this place! Looking forward to some interesting discussions related to inorganic chemistry, lasers and optics! I'm a graphic designer by trade, concert organizer in my free time, and general science enthusiast/hobbyist. I'm a big fan of science fiction, and classic/contemporary progressive rock. - Robert in Montreal