I heated a chunk of indium to red heat. It initially melted, then seemed to solidify. In fact, the surface developed a skin of indium oxide which prevented the metal from dripping all over. After this skin was formed, there was no more oxidation. I can try heating a piece of thin indium foil and seeing whether it oxidizes more thoroughly. Here is what the indium looked like after heating. The indium is still soft and malleable under the coating.
Note: Since my battery is simultaneously wet (squeezing it will make it drip electrolytes) and dry (no beakers full of electrolyte), I will just call it "my cell".
In the past, I had experimented with a variety of cells. My best zinc-copper(II) chloride cell obtained a total voltage of 1.6 volts at 80 milliamps. This was after quite a bit of work and previous trials. Magnesium, which I did not have at that time, is more electropositive than zinc, and should create a higher voltage. I soaked a tissue (which did not soak very well) with dilute copper(II) chloride solution and wrapped it around a fat carbon rod from a D size carbon-zinc battery. I then wrapped a NaCl-soaked tissue around the CuCl2 tissue. A piece of magnesium foil was placed on the outside and the whole assembly was twist-tied together. This battery, which hardly took 5 minutes to assemble and had many problems, produced 1.95 volts at 120 milliamps. Magnesium is a better cathode than zinc when cost is not a concern. I could improve this design by:
The battery was able to run a red LED, which drew 3.0 milliamps.
Pictures will be seen later. Look here for more information about electrochemistry and voltaic cells: http://chemwiki.ucdavis.edu/Analytical_Chemistry/Electrochemistry/Voltaic_Cells
Lead is a unique metal to dissolve as its chloride and sulfate are relatively insoluble. Therefore, hydrochloric and sulfuric acids are not the ideal choice for dissolving lead.
Sulfuric acid: Sulfuric acid dissolution of lead does not work without an electric potential, such as that occurring in a lead-acid battery.
Hydrochloric acid: Lead dissolves extremely slowly in hydrochloric acid. In warm hydrochloric acid, it dissolves very slowly (a minor improvement) and crystals of lead(II) chloride are precipitated when the solution is cooled. Even the addition of hydrogen peroxide does not help much.
Acetic acid: The ancient Romans used this method. Acetic acid placed in lead pots and exposed to the air for several months created lead acetate. This is a very slow method.
Acetic acid and hydrogen peroxide: Besides nitric acid, this seems to be the best and safest method for dissolving lead, if anything regarding lead can be safe. The lead dissolves quite quickly, especially if the acetic acid is boiled down. This dissolution is occurring in a 1:1 mixture of 5% acetic acid and 3% hydrogen peroxide. The dissolution is very rapid when the dilution of the solutions are accounted for.
Nitric acid: This dissolution is the best, but nitric acid is hard to obtain. Soluble lead nitrate will be obtained and nitrogen dioxide fumes will be given off. Since I do not have nitric acid, I cannot do this method.
Electrolytic oxidation: A lead anode connected to a power supply and placed in a sodium chloride solution should produce a large amount of lead chloride. A sodium bicarbonate solution is probably better as the basic carbonate will be formed, which can be dissolved in a variety of acids to form the corresponding lead salts. Be careful with halogen release; when I electrolytically oxidized lead in sodium bromide, bromine was formed as an intermediate, producing a stink during the electrolysis. The lead soon absorbed the bromine and formed pale white insoluble lead bromide.
Copper sulfate undergoes complex reactions with ascorbic acid. When a normal solution of copper sulfate is reacted with ascorbic acid, the solution turns a nice green. A tiny amount of copper metal is produced as well, showing that some copper(I) sulfate may have temporarily formed. Pictures and further experimentation will be seen later.
Silver halides are light sensitive, making them useful for old-time film photography. Silver chloride, a white solid, has the quickest reaction with light. It turns purplish after about one minute in strong sunlight. Upon further exposure, it turns black, as chlorine gas is released (in rates too small to be significant) and silver metal is left behind. Silver bromide, a pale-white solid, reacts after about 15 minutes in strong sunlight. Silver iodide, a yellow solid, hardly reacts with light at all.
The first picture shows silver chloride (top), silver bromide (center), and silver iodide (bottom). Unfortunately, the sodium iodide used to produce the silver iodide has ascorbic acid in it, so the silver iodide was reduced to gray silver metal in the second picture. Some yellow silver iodide flakes are visible around the edges of the black puddle. The center pool of AgBr has turned slightly gray. The flakes of silver chloride on the right side of the bottom picture are quite dark gray.
Other silver compounds also react with light. Silver nitrate is known to decompose when exposed to light in the presence of organic compounds. Silver carbonate seems to decompose to silver oxide in light, which may further decompose to silver metal. Silver acetate solution slowly decomposes to colloidal silver metal when exposed to sunlight. 8 hours is enough to impart a faint Tyndale effect coloration to a dilute solution of the salt in acetic acid.
Copper(II) chloride has a unique ability to partially remove the protective oxide layers from aluminium. Copper sulfate, a more common copper compound, does not behave this way with aluminium. Why does copper(II) chloride have this ability? Does copper(II) chloride have the same ability with other reactive metals protected by thin oxide coatings, like beryllium and magnesium (well magnesium is only partially protected)?
Beryllium reacts rapidly with copper(II) chloride solution, even in low concentrations. Hydrogen gas, copper metal, and the metal hydroxide or chloride is set free. Beryllium also reacts with copper sulfate solution. The difference in reactivity is actually questionable here. The first shows a copper and chloride solution; the second a copper and sulfate solution. The beryllium got a soft copper plating on it from the reaction.
Aluminium reacts rapidly with copper(II) chloride solutions, forming similar products to beryllium. If the copper(II) chloride is pure, aluminium hydroxide is produced as a gelatinous white precipitate. If some residual acid is present, the resulting solution of aluminium chloride is clear. Dilute copper sulfate has no effect on aluminium. The difference in reactivity is illustrated most dramatically with this metal.
Magnesium also reacts rapidly with concentrated copper(II) chloride solutions, forming hydrogen, copper, and magnesium chloride. It reacts a little more slowly with copper sulfate solutions (left) than with copper chloride solution (right).
Solutions of copper(II) and chloride ions have the same effect as copper(II) chloride solutions, showing the formation of a copper chloro complex which is responsible for the reactions. Only for aluminium, however, is the difference seen so strongly; the chloro complex has a doubtful effect on the reaction rates of the other two metals.
A comproportionation reaction sounds extremely complicated. In reality it is quite simple.Two differently charged atoms orr ions of the same element react to form two equally charged ions. Three common examples will be discussed.
First: Copper. When an acidic copper(II) chloride solution is reacted with copper metal, the solution darkens. This is the result of the formation of copper(I) ions. The reaction that is occurring is: Cu(0) + Cu(2+) --> 2 Cu(+) This dark solution is a complex of copper(I) ions with additional copper(II) ions. When the solution is made neutral, the copper(I) chloride precipitates out. These pictures shows the difference in color between the bluish green copper(II) chloro complex and the brownish green copper(I)-copper(II) chloro complex.
Second: Iron. When an acidic iron(III) chloride solution is placed in contact with iron, the iron reduces the iron(III) to iron(II). This reaction is occurring: Fe(0) + 2 Fe(3+) --> 3 Fe(2+) Yellow iron(III) chloride is easily obtained by dissolving rust in hydrochloric acid or by simple purchase. This reacts with iron metal to form a yellow-green iron(II) chloride solution. The test for iron(III) is a brown precipitate when sodium bicarbonate is added. Iron(II) forms from a dark green to a white precipitate. Since my solution was very acidified, all of the oxygen had been displaced from solution by the baking soda - acid reaction, and the precipitate was quite white. Top left: Original solution at experiment start. Top right: Solution after 30 minutes. Bottom left: solution after 9 hours. Bottom right: Iron(II) carbonate. The brown spot is a control, a dab of iron(III) chloride that escaped the iron metal. It forms the brown iron(III) precipitate when reacted with baking soda.
I did this experiment just for my readers.
Third: Tin. When a tin(II) chloride solution is desired, it of course must be pure. However, tin(II) chloride oxidizes to tin(IV) chloride and tin(IV) oxychloride in air. To prevent this oxidative contamination of the solution, tin metal is added. It reacts with any tin(IV) formed, converting it back into tin(II). This reaction occurs: Sn(0) + Sn(4+) -> 2 Sn(2+) All of these solutions are colorless, so pictures are pointless. The formation of a precipitate of tin(II) or (IV) oxychloride may occur, but this precipitate is out of solution and does not act as a contaminant of concern. It is only the tin(IV) in solution that is a problem.
Indium is placed in a dilute copper(II) chloride solution. The indium instantly becomes coated with a steadily thickening layer of copper metal. The copper solution gradually clears over the course of a couple minutes. What remains is a solution of indium chloride and a large amount of spongy copper precipitate. Indium, despite its electrode potential of -0.34 V, behaves partially like aluminium in dissolving quickly in copper(II) chloride solution. This could be further proven by dissolving indium in copper sulfate and comparing the reaction rates.
The resulting indium solution is pipetted off and reacted with zinc metal. The zinc slowly reduces the indium chloride to the metal, resulting in the formation of a spongy layer on the zinc. This bears a similarity to the deposition of tin by zinc-tin(II) chloride, although the indium precipitate is much less spongy.
Here is the initial piece of indium metal.
Here is the indium immediately after immersion in the copper(II) chloride solution.
Here is the indium after a couple of minutes in the solution. The thickness of copper sponge is seen.
Here is the completed reaction. This indium solution is removed and prepared for the next reaction.
Here is the indium solution with a piece of zinc in it.
Here is the indium solution with the zinc after 8 hours.
I saw an interesting experiment where cellulose (paper or wood) is dissolved in a cuprammonium hydroxide solution. This is then squirted into hydrochloric acid to form thin threads of the fabric rayon. The Cu(NH3)OH (simplification) solution is made by adding a small amount of ammonia to a copper salt to precipitate the hydroxide. Filter the precipitate and react it with a large excess of ammonia. The resulting deep blue solution is also known as Schweizer's reagent. The cellulose-containing substance is added to this solution. The solution is then acidified to destroy the ammonia complex and regenerate the cellulose fibers. This video (http://www.youtube.com/watch?v=5aRBn-9yV8Y) states that the cellulose takes an extremely long time to dissolve, so this is not a good demonstration, but it is interesting to make fabric out of some scrap paper!