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April 11 2012 4 11 /04 /April /2012 15:59

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Antimony oxidation: Antimony powder is placed in hydrochloric acid. No reaction occurs since antimony is not reactive enough to dissolve in hydrochloric acid. When hydrogen peroxide is added, however, the antimony dissolves, forming either a colorless solution of antimony trichloride or a white precipitate of the trioxide. The reaction was recorded.


Calcium reactions: I recently obtained 10 grams of calcium for $13.00 from Gallium Source LLC (shipping included). The calcium is golden-colored and dry. Some of this calcium is placed in warm and cold water. The warm water reaction is faster than the cold. Addition of acetic acid rapidly speeds up reaction. A 1% solution of acetic acid reacts vigorously with calcium. Here are some small calcium chunks in the solution.


Indium properties: When indium is hammered, it easily picks up dirt. Granules of indium were hammered into a foil, and then folded to demonstrate flexibility. With one tap of a hammer, a granule completely flattens. By the way, the line in the indium granule is a gentle fingernail press.


Alkaline earth metals in acid: Beryllium does not react with 5% acetic acid, while calcium reacts vigorously. Beryllium, however, reacts vigorously with concentrated hydrochloric acid, just like calcium. The resulting piece of etched beryllium is blackish with pronounced crystal structure. It has a strange sweetish smell. I was sure to wash my hands after handling. Here is the piece of beryllium.

Galinstan alloying: This is a very haphazard experiment. Galinstan is first cleaned by a hydrochloric acid bath. Then the galinstan is washed with water and applied to a scratch in a wad of aluminium foil. Some bubbling is observed as the galinstan removes the passivation layer on the aluminium, allowing it to react with the residual water. Most of the water is pipetted away. Another bit of galinstan on unscratched aluminium foil did not absorb in the time interval of this experiment. The galinstan bead forms a thick gray coating containing a mix of galinstan and aluminium oxides. This floats to the top of the bead and can be scraped off, although it begins reforming instantly. Then the galinstan bead was mixed with the aluminium oxide paste. A golden-colored paste was formed, probably as the result of an oxide film. Then all the powder and paste was tossed in hydrochloric acid. A galinstan bead immediately reforms and the aluminium oxide, along with the galinstan content, dissolves. The galinstan bead is removed to another lid and zinc is tossed in the container. Violent fizzing erupts which quickly slows. For some reason, galinstan-impregnated zinc reacts about 10 times slower with hydrochloric acid than normal zinc, a good indication of the alloy’s formation. When the fizzing has sufficiently slowed, the zinc is placed on the lid where the galinstan bead is residing, hungrily awaiting the next bite of metal, be it zinc (I did not know that) or aluminium. When a large, unresisting, and recently cleaned (by the acid bath) piece of zinc already weakened by gallium reduction from the hydrochloric acid is added, the temptation to the galinstan is too great to resist. My first inkling that an invasion was occurring was the creeping of a silver-colored area on the opposite side of the zinc metal. This caught my attention and it was videotaped as it grew. It grew to about two-thirds of the zinc’s surface area before I lifted the zinc. Half of the galinstan was absorbed into the zinc, while the other half was flattened against the zinc in the process of being absorbed. However, the zinc was becoming saturated with galinstan. When the metal was tilted, a pocket of liquid galinstan formed on the lower side. By this time, the metal was being held by pliers in a location that was intentionally spared from the acid bath and was farthest from the site of the invasion by the next door neighbor. These various features of the metal were taped, using the last of my tiny memory card. I accidentally touched part of the zinc that did not appear to be amalgamated and the piece instantly fractured. I quickly placed the zinc in a Petri dish to prevent it from cracking out of the pliers and shattering on the floor. The end that was being held broke off. I can only imagine how brittle the silvery galinstan-saturated section is. Meanwhile, a thin oxide coating has begun forming on the zinc, turning it a milky yellow of Tyndale effect lore. Here are the pictures: Top is the zinc-galinstan alloy. In the flashbulb light, the shiny galinstan-soaked portion appears dark, probably because reflection is more specular (learn physics). Below, the difference between normal zinc and galinstan-impregnated zinc in reaction with hydrochloric acid is seen. Bottom left is the galinstan-aluminium oxide paste that was a golden color, along with the powder. Bottom right is the appearance of the galinstan in the cleaning bath.

Indium: I stepped on a stray indium bead. It prostrated itself so completely to the ground that it never rose again. No other stable metal is so submissive and shapeable.


Calcium burning: Small pieces of calcium are burnt in a wire loop and the results are videotaped.

Aluminium alloying: A bead of galinstan is placed on lightly scratched beverage can aluminium. The scratching was not enough to remove the thin transparent protective layer on the inside of the can. More vigorous scratching at a new location and moving the bead of galinstan to that location begins the alloy process. A shell of aluminium oxide forms around the liquid galinstan. The aluminium is manifesting evidence of corrosion on an edge about ¾ inch from the original introduction of the galinstan, showing that the aluminium is probably thoroughly impregnated.  The red dashed arrow shows the aluminium oxide shell. The blue circles show the galinstan bead in its shell, top view (bottom) and bottom view (top). The red solid arrow and red circle both show evidence of corrosion in areas of aluminium exposed to the air.


Indium dissolution: A small piece of indium foil is placed in hydrochloric acid. It is completely dissolved within 12 hours. The dissolution begins extremely slowly but evidently increases in speed as any oxide layers are removed from the metal. Two short videos are taken. Later, sodium bicarbonate is added to the solution. A white precipitate forms which easily passes through filter paper. This is most likely indium(III) hydroxide. Like aluminium(III) hydroxide it is very gelatinous.


Indium reactions: A piece of In is placed in CuCl2solution. The indium gets covered by a spongy Cu layer immediately. The Cu thickens quickly, and a dilute solution of CuCl2 is cleared of Cu in about 1 minute (when the indium is shaken, the reaction is much quicker). Here are some pictures of the reaction.    

The first picture shows the piece of indium foil. The second shows the indium foil floating on the copper(II) chloride solution. The solution has touched the edges and the underside, turning them both reddish-brown. The third shows the indium foil immediately after submersion. The copper layer is thin and dark. The fourth picture shows the indium foil about 10 seconds later. The copper layer has grown in thickness, although the solution has not experienced any significant de-coloration. (The difference in solution color is fluorescent vs. flashbulb light.) The fifth picture shows the thickness of the copper around the indium piece. The last picture shows the solution after the reaction has run to completion and been agitated. Pieces of spongy copper have broken from the indium piece, leaving the blackish-looking indium partially exposed. The solution has decolorized, turning into indium(III) chloride. The reaction that has occurred is 2 In + 3 CuCl2 à 2 InCl3 + 3 Cu. Indium is a quite reactive metal.

Indium reduction: The resulting indium chloride solution from the above experiment is reacted with zinc. Initially, no reaction is apparent, but a layer of indium later is seen to be forming on the zinc. It appears spongy, though not as spongy as the copper layer above. When the indium sponge is later compressed, it behaves like a metal, just like the lead sponge previously made, which also compresses to a solid metal. However, it was lost. These are pictures of the reaction. The first (left) is the zinc just after immersion in the indium chloride solution. The second shows the zinc about 8 hours afterwards. So with the third (the indium sponge is visible on the edges) and the fourth (compare to the first). The fifth shows the indium powder and the zinc outside of the solution. Some of this powder was pressed into the pellet which soon was lost.


Iron comproportionation: Iron(III) oxide is dissolved in an excess of hydrochloric acid to create a yellow solution. A machine bolt is then placed in the solution. A brief and swift dissolution of the electroplated zinc coat occurs, and the hydrogen production stops. It seems as if all of the iron is undergoing this reaction (Fe + 2 FeCl3 à 3 FeCl2) instead of the typical Fe + 2 HCl àH2 + FeCl2. The solution, after 30 minutes, has noticeably turned greener. The reaction will be left overnight. Here is the first set of pictures. The first picture shows the iron(III) oxide completing its dissolution in hydrochloric acid. The second picture shows the machine screw 30 seconds after starting. The third picture shows the solution after 30 minutes.

Beryllium copper(II) chloride reaction: Because of beryllium’s many similarities to aluminium (relatively high melting point, dissolution in alkalis, protective oxide coating), I wondered whether beryllium would undergo the same vigorous reaction that aluminium undergoes with copper(II) chloride. However, I did not want to exhaust, contaminate, or ruin my beryllium, so I only used a highly dilute copper(II) chloride solution (boring). The beryllium reacted more vigorously than the aluminium would under similar circumstances, forming hydrogen gas, beryllium chloride (somehow it stays in solution from the copper(II) chloride’s excess acidity), and copper metal. The beryllium was covered in a layer of dark brown copper, which smeared as it washed off, giving the impression that the beryllium was corroding. However, upon rubbing, the strange dark luster of corroded beryllium metal shone again. The brief (to prevent excess beryllium dissolution) reaction was videotaped.



Iron comproportionation: The reaction was deemed complete by morning, after about 10 hours. The solution was neutralized with sodium bicarbonate to precipitate the iron. A white precipitate formed. This white precipitate is iron(II) carbonate, which is white when pure and oxygen-free as a result of the vigorous carbon dioxide bubbling through the solution. Some of the solution not in contact with the iron was smeared near to the top of the vial, and it retains the brown color of iron(III) oxide (iron+++ doesn’t form a carbonate). This shows that the comproportionation reaction was completely successful and the resulting solution only contained ferrous, not ferric, ions. After a few minutes, the edges and top of the iron(II) carbonate had turned either brown or dark green, the products of aerial oxidation. Ordinary tap water, which contains dissolved oxygen, was added to the solution, and the iron(II) carbonate darkened to a greenish color. Addition of hydrogen peroxide turned it brownish amid fizzing. The left picture shows the resulting iron(II) chloride solution, containing excess hydrochloric acid. The center picture shows the white iron(II) carbonate precipitate. The brown spot at the top is the iron(III) oxide from the unreacted iron(III) chloride solution. The right picture shows the iron(II) carbonate after standing a few minutes.



Titanium reactions: Titanium foil piece from GalliumSource is placed in hydrochloric acid. A vigorous reaction occurs and the titanium dissolves as fast as a piece of aluminium would. Sodium bicarbonate is added. A precipitate occurs even in a strongly acidic solution (much fumes, pH < 1). Wasn’t titanium supposed to dissolve slowly in even boiling hydrochloric acid? Well, the heat produced by this reaction was enough to boil away the hydrochloric acid. The translucent gel on the side of the vial in the picture below is similar to aluminium hydroxide, which would not exist in such a strongly acid solution. The precipitate, however, is curdy and relatively heavy.





“Titanium” reactions: The foil reacts slowly with water, forming small bubbles of hydrogen gas and a white precipitate. This is definitely magnesium.





Metal and copper reactions: The reactions of magnesium, aluminium, and beryllium with copper sulfate and copper chloride solutions are compared. Aluminium does not react with copper sulfate, while it reacts vigorously with copper chloride. Magnesium and beryllium only exhibit slightly increased signs of reactivity with my excess-acid copper chloride; the excess acid is in all likelihood the only reason for the increased reaction rate. Therefore, only aluminium exhibits the strange phenomenon of being so much more reactive with the chloride than with the sulfate. The pictures are of Al-chloride (top left), Al-sulfate (top center), Be-chloride (top right), Be-sulfate (center left), Mg-chloride (center right), and Mg-sulfate (bottom).





Magnesium reactions: Magnesium reacts vigorously with concentrated copper sulfate solution, just like its reaction with copper chloride minus the acidity. Magnesium also reacts violently with a copper sulfate-ascorbic acid mixture, which turns green for some reason (complex formation?). Brown copper mud is produced in the reaction. Magnesium reacts moderately with plain ascorbic acid, showing a rare instance where ascorbic acid functions as an oxidizing agent. Top left is the magnesium-copper sulfate reaction. Top right is the magnesium-copper sulfate-ascorbic acid reaction. Bottom is the magnesium-ascorbic acid reaction, with small bubbles escaping along the sides.




Copper sulfate – ascorbic acid reaction: This try was not good. Too much ascorbic acid was present.





Copper sulfate – ascorbic acid reaction: Copper sulfate solution, which is sky blue, reacted with ascorbic acid to form a green solution. A tiny amount of fine copper metal dust has precipitated.



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April 11 2012 4 11 /04 /April /2012 15:47

I obtained iron-free manganese dioxide (probably mixed with carbon) from a tantalum capacitor. It is quite uniformly black, with some silver electrodes connecting it to the current. When placed in cold hydrochloric acid, it releases a small quantity of chlorine gas and produces a brown solution. I initially thought that this was an indication of iron, but the solution decolorized when heated, showing that the manganese dioxide was relatively pure. It seems that manganese(III) is relatively stable as a chloro complex in cold solutions, but decomposes to chlorine and manganese(II) in hot solutions.


If you have a hard time bringing a hot plate or any other heat source outside to perform the reaction, you may absorb the chlorine produced by a tissue soaked in either ascorbic acid or sodium metabisulfite mixed with sodium carbonate. No chlorine smell will be produced until the tissue is exhausted of its chemicals. After the dissolution, pure manganese(II) chloride solution will remain. If you desire to keep a source of manganese ions, precipitate the manganese as the carbonate, which is resistant to further aerial oxidation, unlike the hydroxide. Manganese(II) chloride can be evaporated but it needs to be stored in a tightly sealed container to avoid any deliquescence.


Manganese dioxide from an alkaline battery or a carbon-zinc battery is more problematic. In these, large quantities of iron impurities are generally present, which can be difficult to extract. Occasionally, you may find a battery that happens to be free of iron impurities. If so, a colorless solution will be formed upon heating the manganese dioxide with hydrochloric acid. Otherwise, a dirt brown solution will form. This may be remedied by dissolving the manganese dioxide in sodium metabisulfite solution. The sulfite will reduce the manganese(IV) to soluble manganese(II), while the iron(III) remains untouched.


You may get purer manganese dioxide from 3 volt lithium coin cells. I generally have had better experience with these than with alkaline batteries, but they are much more expensive than the aforementioned batteries.   Once I got a batch of pure manganese(II) chloride from an unknown alkaline battery.




Extracting pure manganese salts from batteries is difficult, but the fact that manganese(IV) is a stronger oxidizer than iron(III) is useful to know.

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April 10 2012 3 10 /04 /April /2012 15:42

Titanium is a rather unreactive metal, making it difficult to dissolve. There are several ways to oxidize titanium.


Hydrochloric acid: Boiling concentrated hydrochloric acid will slowly dissolve titanium, forming a purple solution of titanium trichloride. When I tried this, hardware store hydrochloric acid was unable to dissolve the titanium. It might work if it was boiled for an extended period of time, however.


Sulfuric, acetic, nitric acid: Titanium is quite resistant to all of these acids. It may dissolve extremely slowly in boiling sulfuric acid but that is a generally poor way of dissolution.


Bromide dissolution: I found somewhere on the internet that titanium has a weakness for bromine. Therefore, I decided to dissolve it by electrolyzing it in a sodium bromide solution with a nine volt battery. I was really amazed that absolutely no bromine was produced. All of the bromine reacted with the titanium, which appeared for form a soluble tetravalent complex with the bromine. When this contacted the sodium hydroxide produced at the cathode, a clumpy white precipitate began forming. The titanium gradually grew thinner until it was just a brittle fragment remaining that was no longer electrically connected. I noticed that the titanium hydroxide had a purplish tint to it, showing that some titanium was found in the trivalent state.




The titanium foil completely dissolved within a short period of time, showing the efficacy of this method. The product is also soluble in hydrochloric acid, forming a colorless solution which turns bright red upon addition of hydrogen peroxide. However, a smell of bromine is noticed, so is the red coloration the titanium peroxo complex or the bromine water? I will need to do more research into this.



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April 9 2012 2 09 /04 /April /2012 15:32

What is the best way to store lithium?


I tried extra virgin cold-pressed heart-healthy olive oil. The lithium appears to react initially, producing a few bubbles. However, the lithium was covered with its foamy electrolyte and so it was not clear whether it was the lithium or the electrolyte reacting. The reaction appears to stop after a while. Olive oil primarily consists of oleic and palmitic acid esters, which do not seem to react with the lithium metal. So olive oil is a good, though not ideal, way to store lithium metal because of its lightness. Lithium is liable to get stuck on the side of the container after any disturbance because of its density. After the oil drains from the surface, it is susceptible to oxidation and therefore is destroyed.


Plastic wrap is the first method that I used to store lithium metal. However, the lithium's reaction with almost all atmospheric gases guaranteed that some gas would find its way into the plastic wrap. The lithium did have somewhat of a restriction on reacting; I detected a distinct smell of ammonia when the resulting white powder was placed in water, signifying that nitride was produced due to the lack of oxygen.


I then tried petroleum jelly. Lithium keeps very well in this medium. Since it does not float in petroleum jelly, there is no worry about lithium being isolated on the side of a container. A large piece of lithium can be covered with petroleum jelly and pressed against the side of a container to make an excellent lithium display for an element collection. A disadvantage is that the petroleum jelly coats the surface of the lithium, making it less likely to react and contaminating any reaction.


LIthium can remain in air for about a day before it starts significantly oxidizing (turning white). This is for normal air with about 40% humidity. With extra-humid air, however, the time period is much shorter. If you are opening a lithium coin cell for only one quick experiment, it is best to skip preservation and just use the lithium as-is. For a collection or display, use petroleum jelly. For further experiments, use petroleum jelly but be sure to completely wipe off the lithium before use.

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April 7 2012 7 07 /04 /April /2012 15:58
I have become interested in mini-thermites, thermites that are relatively safe and interesting. Magnesium, being more reactive than aluminium and still quite stable, appears to be the reducing metal of choice. The heat wave created by a magnesium thermite is so intense that it blows the chemicals apart after ignition, prevent any coalescing of liquid metal. This is bad for most thermite applications but good when intense heat is quite undesirable. Grinding magnesium on a file is relatively easy; the provided scraper with the firestarter makes larger pieces (less desirable) with more difficulty than the file.
I had previously prepared iron(III) oxide hydrate by electrolytically oxidizing iron in aqeuous solution and filtering the precipitate formed. The powder is mixed and ground with the magnesium dust to form a brown mixture that looks to be highly reactive and a lot of fun. The following sentence is not true, but was written initially: Ignition will be with a magnesium strip, not the dangerous mischmetal flint, because of the decreased size of the magnesium.
This thermite failed with titanium thread ignition, steel wool ignition, and mischmetal - magnesium ignition. This was because first, there was too much excess iron(III) oxide. Second, the iron(III) oxide was hydrated, making it unlikely to burn.
So I made a new mixture. I heated the iron(III) oxide to dehydrate it. I then mixed it with an excess of magnesium shavings, since they will burn even when no iron(III) oxide is present. I then scraped up a large amount of magnesium shavings and powder. I placed the thermite mixture on a brick, then piled the shavings on top. On top of that, I placed the easily ignitable powder. I then ignited the powder. Since the shavings take a second to ignite, it gave me enough time to retreat before the reaction began. I ended up using a mischmetal flint just because it was the easiest way at the time. The reaction was quite silent and slow compared to copper, but it was much hotter because of this.
Here is the video of the second trial.

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April 6 2012 6 06 /04 /April /2012 14:28

Many chemicals form hydrates, where water molecules are bound to the chemical molecules. These hydrates often have different colors and different properties. For instance, anhydrous stannic chloride is a corrosive, fuming liquid, while the pentahydrate is an acidic white solid.


Copper(II) chloride has two forms: a brown anhydrous form and a blue-green dihydrate form. Because of excess HCl, my copper(II) chloride hydrate is more green than blue. When it is heated gently, it turns brown and releases both water vapor and white fumes, showing that the excess HCl and the water have been released from the crystal structure.


Here is a picture of the resulting product, with the original dihydrate for comparison.




Since the dehydration required heat, the reaction is likely endothermic. Therefore, rehydration should be exothermic. It is. However, because the anhydrous copper(II) chloride had been heated too vigorously, there was some decomposition to HCl and CuO and the solution formed when it was placed in water was cloudy. Here is the video of the reaction when two drops of water are added to the bulk of the anhydrous substance.


I then tried burning a crystal of copper(II) chloride with a magnifying glass. The copper(II) chloride refracted the green coloration for a moment before they turned brown as they were dehydrated. Then they turned black and melted. The resulting liquid was quite mobile. When cooled, it forms a black amorphous solid.




This dissolves in water, forming a mixture of copper(II) chloride solution (green), copper(II) oxychloride precipitate (green), and copper(I) chloride precipitate (white). The copper(I) turns to the oxychloride upon exposure to air. The black lump completely dissolves.


If I do the same with copper sulfate, I will post the results in this article.

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April 4 2012 4 04 /04 /April /2012 15:39
This experiment involves reduction of copper(II) oxide, a black water-insoluble powder, with magnesium metal. The reaction is CuO + Mg -> MgO + Cu. The reaction is quite exothermic. Do this away from all flammable materials. Do not get too close to the reaction as burning metal can spray all over with larger quantities. This reactions is commonly known as a thermite.
Grind a magnesium rod or bar into shavings or, better yet, powder using a knife or a file.
Precipitate copper(II) carbonate from a solution of a copper(II) salt by adding sodium bicarbonate. Filter, wash, and dry by heating. It will turn black as it decomposes to copper(II) oxide. If you just have copper metal, electrolytically oxidize the copper in a sodium bicarbonate solution to form copper(II) carbonate. Heat this paste after filtering and drying.
Using a mortar and pestle, grind equimolar amounts (approximately 3 g CuO to 1 g Mg) of these two chemicals together until the mixture is uniform.
Place on a metal surface in a small pile. Ignition using a ferrocerium flint can be dangerous as the hands will be very close to the mixture. A better method of ignition would be electric ignition using a very thin piece of titanium foil and a nine volt battery or a magnesium ribbon.
Despite the dangers, I used a ferrocerium flint to ignite the mixture. Since there was some wind, I was afraid that the thermite mixture would blow away after placing it out in the open piece of metal. Therefore, I ignited it right away.
If you will try this at home, please be more careful than I was.
The "Shoe for Comparison" picture malfunctions in the video.
I then tried a copper(II) sulfate reduction. Aqueous copper sulfate reacts quite vigorously with magnesium metal, but the crystals react yet more vigorously. I shredded some magnesium and ground the copper(II) sulfate - magnesium shaving mixture together with the end of a pen. I then placed it on the same metal dish in a small pile and added some more magnesium shavings on top. Since grinding magnesium is so difficult, I hardly produced enough magnesium on top to ignite the mixture, so I had to strike the ferrocerium for a couple minutes before I actually got it ignited. The top layer burnt with a dazzling white flame, but the bulk was still unburnt. It would not ignite. Here is the video of the ignition.
I retried this with more finely powdered magnesium and dehydrated copper(II) sulfate. The result was much better.

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April 2 2012 2 02 /04 /April /2012 17:17

Zinc is often used as a safe, stable, and dirt-cheap reducing agent for inorganic and preparative chemistry. For example, the easiest way to produce chromium(II) and vanadium(II) solutions is reduction of higher oxidation states with a mixture of zinc and hydrochloric acid. However, zinc can also function as an oxidizing agent. Stronger reducing agents such as magnesium and possibly aluminium are capable of reducing zinc(II) ion solution to solid zinc.


A couple of years ago, I dissolved chopped up pennies in vinegar to form zinc acetate. After several weeks, strong-smelling and damp crystals had separated out of the solution. I placed them in a container and never used them until recently. Then, I took some of these crystals, dissolved them in water, and placed the entire magnesium firestarter rod in them. Some gentle fizzing was observed, which is normal with metal salt-magnesium reactions. Then, a blue-gray film was beginning to form on the magnesium. I checked 12 hours later and noticed the piece of magnesium covered with crumbly growths of zinc. They were not spongy or metallic, and dried to a powder. After another day of drying, they were covered with a white hydroxide film because of their finely divided state. I photographed them in this state and then disposed of them.





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March 31 2012 7 31 /03 /March /2012 14:04

I earlier tried to add hydrochloric acid to Pepto-Bismol. It decolorizes, but the solution remains too thick to allow for any extraction of the bismuth chloride solution. To fix this, I add Pepto-Bismol drops to a dilute hydrochloric acid solution and stir. The drops break up and decolorize, and the bismuth dissolves. The solution is filtered and reacted with zinc metal. The zinc metal turns black as a thin smear of bismuth forms on the surface.


However, although Pepto-Bismol is widely available and relatively cheap, this is not a good way to produce bismuth metal. The amount produced is so small and so finely divided that it just stains filter paper. Here is how the zinc looked.



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March 29 2012 5 29 /03 /March /2012 15:04

I placed a piece of nickel metal from a spark plug in a moderately concentrated copper(II) sulfate solution and let it sit overnight. No reaction was observed in the morning.


I then placed the nickel in some water and added a copper(II) chloride crystal. It dissolved, leaving a highly concentrated layer of copper(II) chloride solution on the bottom of the vial. The nickel metal was right in this solution. It seemed that a tiny amount of copper was produced, but even if a reaction was occurring, it was so slow that it would take weeks to get any significant amount of copper from the nickel reaction. Even placement in a boiling water bath did not produce any reaction between the copper(II) chloride and the nickel metal.


Nickel is quite an inert metal, with an oxide coating protecting against oxidation to a significant extent. It seems that this oxide coating is quite resistant to copper compounds.

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