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March 7 2012 4 07 /03 /March /2012 15:59

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.

 

Lead-dissolving-in-HAc---H2O2.JPG

 

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.

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March 7 2012 4 07 /03 /March /2012 15:34

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.

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March 6 2012 3 06 /03 /March /2012 18:17

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.

 

Silver-halides.JPG Silver-halide-light-exposure.JPG

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.

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March 5 2012 2 05 /03 /March /2012 21:14

 

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.

 

Beryllium copper and chloride Beryllium copper sulfate

 

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.

 

Aluminium copper sulfate Aluminium copper and chloride

 

Youtube video

 

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).

 

Magnesium copper sulfate Magnesium copper and chloride

 

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.

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March 5 2012 2 05 /03 /March /2012 17:24

I recently purchased some titanium foil from GalliumSource. They sent magnesium. My first inkling of a problem was an extremely vigorous reaction between the foil and hydrochloric acid. It also formed small bubbles when placed in water and reacted violently with copper chloride. It was not shiny, showing a film of oxidation typical of magnesium. It was remarkably light but not remarkably strong, unlike titanium. After a short email discussion, I received word that I was receiving a replacement piece of titanium foil.

 

http://www.youtube.com/watch?v=3a0cSP5pSkc&feature=youtu.be shows several properties of this magnesium foil, demonstrating obvious differences from titanium.

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March 4 2012 1 04 /03 /March /2012 02:52

Sometime in the past I passed 50 volts of electricity through a thin piece of zinc. The hot arc melted and boiled the zinc, sending molten zinc everywhere. Once the zinc even ignited and burned with its characteristic blue-green flame. One piece of molten zinc landed on my brother's hand and was captured by this still image from a video. Soon after he says "ouch". The piece of zinc is between the alligator clips, barely visible. Zinc has a low melting and a low boiling point, making this process especially splatter-prone.

 

Snapshot---8.jpg

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March 3 2012 7 03 /03 /March /2012 15:22

I had two old mercury thermometers. Now there is one. Outdoors, I broke open a mercury thermometer and placed the majority of the mercury in a vial. To seal the vial, I placed petroleum jelly around the lid and tape over the petroleum jelly. Unfortunately, the subsequent playing with the mercury bead dislodged quantities of the petroleum jelly and spread it around the vial. The mercury was difficult to see.

 

Mercury-in-vial.JPG

I opened the vial (the petroleum jelly had countless tiny beads of mercury stuck in it) and placed it in another vial previously shown to have a good lid seal. Now the mercury is much easier to see. Here is a closeup of the bead. I plan to do nothing with the mercury in the near future except keep it with the element collection.

 

DSCF9885.JPG

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March 2 2012 6 02 /03 /March /2012 14:51

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.

 

Copper-I--and--II--chlorides.JPGCopper-dissolving.JPG

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.

Ferric-and-iron-comprop-1.JPGFerric-and-iron-comprop-2.JPGFerric-and-iron-comprop-3.JPGIron-II--carbonate-product.JPG

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.

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March 1 2012 5 01 /03 /March /2012 18:35

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.

 

DSCF9857.JPG

 

 

Here is the indium immediately after immersion in the copper(II) chloride solution.

 

DSCF9860.JPG

Here is the indium after a couple of minutes in the solution. The thickness of copper sponge is seen.

 

DSCF9864.JPG

 

Here is the completed reaction. This indium solution is removed and prepared for the next reaction.

 

DSCF9865.JPG

Here is the indium solution with a piece of zinc in it.

 

DSCF9867.JPG

 

Here is the indium solution with the zinc after 8 hours.

DSCF9868.JPG

DSCF9869.JPG

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March 1 2012 5 01 /03 /March /2012 18:21

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!

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