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March 29 2012 5 29 /03 /March /2012 14:59

Silver is best dissolved by nitric acid, forming a colorless solution of silver nitrate. However, nitric acid is difficult to obtain, so acetic acid is used instead to dissolve it.

 

I dissolved some anodized (silver is made vulnerable, not protected, by anodizing) silver in acetic acid and hydrogen peroxide. It formed a milky solution of silver acetate (because of chloride impurities present in the vial). I then placed a piece of copper wire in the solution. Some silver formed as a black coating, but no silver crystals formed because of the insistent decomposition of hydrogen peroxide that was occurring from the copper wire. A purplish precipitate of what I assumed was micro-particulate silver accumulated on the bottom of the vial. When disturbed, it went into suspension. This same precipitate formed when ascorbic acid was reacted with the silver acetate solution.

 

Copper-wire-and-silver-acetate-2.JPGColloidal-silver--2-.JPG ae

 

I also precipitated some silver carbonate. It formed as a fluffy white precipitate. Once filtered and dried, it turned brown as silver oxide or even elemental silver (from light exposure) was formed.

 

Silver acetate decomposes on exposure to light. After a day in daylight, the silver acetate was colored slightly bluish because of silver metal that had begun forming from the decomposition.

 

Because of working with this silver solution, I discovered that just about every piece of equipment I have for chemical experiments is contaminated with chloride ions.

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March 28 2012 4 28 /03 /March /2012 15:42

Bismuth subcarbonate is formed by dissolving bismuth in a mixture of hydrogen peroxide and hydrochloric acid to form a colorless solution of bismuth(III) chloride. This is then reacted with sodium bicarbonate to precipitate white bismuth subcarbonate. The precipitate is filtered and dried. Some of this precipitate is placed on a piece of aluminium foil and heated. It turns black on the edges, then melts. The flame coming off the molten lump is slightly green. Once the lump cools, it is a yellowish color on the outside.

 

I then took another piece of dried bismuth subcarbonate and placed it on an iron loop. Again, it turned black and melted. I heated it until it was bright red hot and splashed it on the cold stove surface. The bead cooled to a yellow sphere (flat on one side) with bismuth metal visible on one end. It appears that the molten bismuth trioxide reacted with the iron wire to form iron(III) oxide and bismuth metal, which has a low melting point.

 

The bismuth subcarbonate decomposed to bismuth trioxide and carbon dioxide when heated. Bismuth trioxide melts at 817 C.

 

Here is the bead with the bit of bismuth metal visible.

 

  Bismuth-oxide-bead--1--copy-1.JPG

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March 23 2012 6 23 /03 /March /2012 12:41

I was looking for a solvent for bismuth metal, which I purchased from a gift shop for the Franklin Mineral Museum. I knew that acetic acid would only form the insoluble subacetate, so it cannot be used. I do not have nitric acid or sulfuric acid, so hydrochloric acid remains. Addition of hydrogen peroxide to the hydrochloric acid allowed the bismuth to be oxidized and the oxide to be dissolved, leaving a solution of bismuth(III) chloride. This solution is dense and colorless, visible as swirls when the dissolving solution is disturbed. The remaining bismuth is pitted and generally corroded.

 

Bismuth-and-chloride.JPG

A large amount of ascorbic acid crystals are added and some tincture of iodine is added as well. The tincture of iodine is decolorized, but no bismuth iodide is produced. I cannot seem to obtain bismuth iodide.

 

I then added water to the solution, to the top of the vial. The bismuth instantly hydrolyzed in the diluted solution, producing a white precipitate of BiOCl.

 

Bismuth-chloride-hydrolysis.JPG

When this precipitate is further neutralized with sodium bicarbonate, it turns light yellow as bismuth hydroxide is produced.

 

After disposal of the relatively nontoxic bismuth precipitate down the sink (bismuth solutions are too acidic to stay as ions in aquifers, so they precipitate and remove themselves naturally), I repeated the experiment. I took some of this bismuth solution and reacted it with an iodide. A yellow iodo complex is obtained.

 

Bismuth-iodo-complex.JPG

When this is diluted with water, it turns white again as the iodide hydrolyzes.

 

I then took some more of the bismuth solution and placed it on a piece of zinc. The zinc turned black as the bismuth chloride was reduced to elemental bismuth. The resulting layer was quite thick.

 

Bismuth-chloride-zinc-reaction--4-.JPG

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March 22 2012 5 22 /03 /March /2012 13:19

Copper forms both divalent (cupric) and monovalent (cuprous) compounds. Since most of the cuprous compounds are insoluble, they are relatively easy to prepare.

 

Copper(I) oxide: See here for the red form: Production To make the yellow form, react sodium hydroxide with copper(I) chloride. It will turn bright yellow. Another potential way to make the yellow form is to electrolyze concentrated salt water with a pure copper object as the anode. The anode initially turns white as copper(I) chloride is produced, then turns yellow as the NaOH from the cathode reacts with it. This may not make pure copper(I) oxide, which is susceptible to oxidation by any oxygen produced at the anode as well as air.

 

Copper(I) chloride: React copper(II) chloride (Production) with either ascorbic acid or sodium metabisulfite to form a white precipitate of copper(I) chloride.

 

 

Copper(I) bromide: React copper(II) sulfate with an equal amount of sodium bromide. Add ascorbic acid or sodium metabisulfite. White copper(I) bromide will precipitate.

 

Copper(I) iodide: React copper(II) sulfate and sodium iodide. A dark brown mixture of iodine and copper(I) iodide will form. Add ascorbic acid to reduce the iodine to water-soluble iodide and only a white precipitate should remain. You can also react copper sulfate and a mixture of tincture of iodine and excess ascorbic acid. A white precipitate of copper(I) iodide will form.

 

I reacted copper(II) chloride with a mixture of tincture of iodine and ascorbic acid. Several reactions occurred at once: reduction of copper(II) chloride to copper(I) chloride by ascorbic acid, reaction of copper(I) chloride with copper(II) chloride to form the dark brown complex seen here, reaction of copper(II) chloride with sodium iodide to produce brown iodine and colorless copper(II) iodide, and reaction of iodine with ascorbic acid to form colorless iodide. The end result, after all products were mixed together, was a light brown precipitate.

 

 

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March 21 2012 4 21 /03 /March /2012 13:08

I wanted to produce beryllium hydroxide as my example of a beryllium compound. To produce this solid, I placed beryllium in copper(II) chloride crystals, then added a few drops of water. An extremely vigorous reaction began, with steam being profusely evolved and the beryllium jumping around. When more water was added, the reaction slowed, but large amounts of copper and hydrogen were still being produced.

 

Beryllium-copper-II--chloride-reaction.JPG

 

I accidentally placed the lid on the container and it popped off 6 feet into the air after a few seconds. When the reaction was over, I let the copper particles settle, then decanted the cloudy (beryllium hydroxide was dissolved in its chloride solution) solution from the copper. Addition of ammonia seemed to solidify the solution. Beryllium hydroxide precipitated as an extremely firm and gelatinous precipitate, colored slightly blue by copper impurities locked inside the hydroxide structure.

 

Beryllium-hydroxide-flashbulb.JPG

 

Tilting the container up side down made the excess ammonia and liquid run out while leaving the structure of the hydroxide untouched. I added some water and vigorously shook it to dislodge some of the beryllium hydroxide. After stirring with a twist tie and using various other ways to break the hydroxide into small pieces, I finally got it onto filter paper, where it dried to a powder, just as I expected. (Do not breathe fumes, can cause berylliosis!) This powder can be used to generate other boring beryllium compounds.

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March 19 2012 2 19 /03 /March /2012 17:56
Silver, being a relatively inert metal, does not dissolve in hydrochloric or sulfuric acid. Even if it did, it would form the insoluble silver chloride or the slightly soluble silver sulfate. Here are some of my experiences with dissolving silver.
 
Aqua regia: Silver is insoluble in aqua regia because of the presence of chloride.
 
Nitric acid: This is the best solvent for silver metal. Silver dissolves easily in nitric acid, releasing nitrogen dioxide fumes and forming a colorless solution of silver nitrate. Here is a video by NurdRage on YouTube demonstrating the process.
 
Acetic acid and hydrogen peroxide: Although silver acetate is soluble, it barely passes as a soluble salt (where soluble is defined as excess of 1 g/100 mL dissolving, silver acetate is 1.02 g). Ordinary silver dissolves with difficulty in a 1:1 mixture of 5% acetic acid and 3% hydrogen peroxide. However, "anodizing" the silver in a salt water bath corrodes the surface, allowing it to dissolve in the aqueous solution. The solution is cloudy as the result of residual chloride ions in the metal from the "anodizing". The hydrogen peroxide decomposes as the silver dissolves, allowing bubbles to be used as an indicator of dissolution. This solution can be used to synthesize basic silver compounds. It might be non-ideal for the growing of a silver tree, as excess reagents in the solution can cause redissolution of the silver crystals from the copper wire. More concentrated reagents should be better.
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March 17 2012 7 17 /03 /March /2012 14:58
Lead iodide is a yellow solid insoluble in cold water but slightly soluble in hot water. It is easily produced from a few basic substances.
 
Caution: Lead iodide is toxic, as are all lead compounds. Do not breathe dust. Dispose of lead iodide at hazardous waste facility or just use small amounts of lead.
 
Needed: Lead metal, white vinegar, hydrogen peroxide, tincture of iodine, ascorbic acid (or sodium metabisulfite)
 
First, prepare solution 1. Add hydrogen peroxide to white vinegar in a 1:1 ratio. Add lead. Fizzing is observed and some of the lead dissolves. If any antimony, tin, or arsenic is present, it is left behind. Copper may dissolve but it will not effect the reaction. After 24 hours, filter the solution to remove the lead and any precipitate. This is a dilute solution of lead acetate.
 
Second, prepare solution 2. Add ascorbic acid (or sodium metabisulfite) crystals to tincture of iodine until it is colorless. Add about 1/5 more of the crystals to provide an excess of reducing agent. This is a dilute solution of sodium iodide, along with other compounds.
 
Third, mix the solutions. Begin by adding a few drops of solution 2 to solution one. A startlingly yellow precipitate will form from the two colorless solutions. Add more of the iodide solution until no more precipitate is formed upon addition. The excess hydrogen peroxide in the lead solution should be neutralized by the excess ascorbic acid (or sodium metabisulfite). If it is not, the resulting solution will be brownish and the lead iodide will begin turning brown. If this is the case, add more reducing agent to turn the lead iodide yellow again. Filter the precipitate and discard of the solution down the drain. (Since all of the lead was precipitated out, the solution contains little remaining lead and can be safely disposed of in this manner.)
 
Allow the precipitate to dry and place it in a small amount of water. Heat the slurry in a boiling water bath until the lead iodide dissolves. If it does not completely dissolve, add a little more water. Once it is almost completely dissolved (the residue is probably impurities), turn off the heat and place the lead iodide solution (now turned colorless) in an ice bath (salt is unnecessary). The solution will cool down and lead iodide will precipitate again. This lead iodide can be filtered and dried again. This lead iodide is now pure.
 
Here is a video of the reaction between the lead acetate and the sodium iodide solution.
 

 
 Here is a video of the purification in an ice bath.

 

Have fun with lead chemistry, but be safe with small amounts! All of my lead experiments (I also formed the sulfate, bromide, chloride, dioxide, monoxide, etc.) used about 200 mg of lead, which is quite a small amount of metal.

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March 15 2012 5 15 /03 /March /2012 15:01
Titanium, as a left side transition metal, is more prone to vigorous reactions at high temperatures than at low (e.g. room) temperatures. Actually, at low temperatures, titanium is an inert metal, hardly dissolving in or reacting with any acid, base, or chemical.
 
When copper(II) chloride crystals are placed on titanium and then heated, the blue-green copper(II) chloride dihydrate is first dehydrated to the brown anhydrous form. Then the excess HCl present in my copper(II) chloride is released and the CuCl2 begins heating. A blue flame color is observed, as well as some sparks. When the titanium gets hot enough, however, a red-orange flame shoots out of the titanium metal, and a cloud of white smoke is released. No copper(II) chloride remains after this reaction, but copper metal does, and underneath the black copper(II) oxide surface lies reddish-brown copper metal. A significant portion of the titanium reacted with the copper chloride, making the resulting titanium piece thin and brittle. What reactions occurred here? It seems that 2 CuCl2 + Ti + 2 H2O (from atmosphere) --> TiO2 (smoke) + 2 Cu + 4 HCl is a potential reaction. This is a net reaction, taking into account the instant hydrolysis of titanium tetrachloride in moist air. Unfortunately, the camera overexposed during the video, so much of the flame looks white instead of the wide range of colors that actually existed.
 
If titanium tetrachloride is really formed, then doing this reaction in a closed crucible may produce some of the volatile substance without hydrolysis.
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March 14 2012 4 14 /03 /March /2012 17:24
I recently obtained a magnesium firestarter. The magnesium is scraped into shavings by the accompanying saw and ignited with sparks from an attached mischmetal rod. The magnesium burns with a dull white flame. When this magnesium contacts any water, it both flash boils the water and reduces it to hydrogen, which ignites. The end result is a flash of light and an instant scattering of the magnesium shavings.
 
I thought about mixing some calcium metal in with the magnesium shavings. Although the calcium metal chunks were larger than the shavings, I had hope that at least some of them would ignite. They did. The calcium burnt quickly with a brilliant red-orange light, easily drowning out the magnesium fire. I took a video of this reaction; here it is.
 
I then, instead of scraping the magnesium block, scraped the mischmetal rod used to ignite the magnesium. Scraping it slowly prevents it from igniting. However, the pile of mischmetal flakes is very easily ignitable and burns with a relative dim white light, throwing off sparks and crackling in the process. It reacts with water in the same way that magnesium does, as magnesium holds many similarities with the rare earth metals.
Here is another video. This time the mischmetal shavings are mixed with a piece of lithium, which does not ignite.

 

  I will need to try a lithium - magnesium mixture to see whether I can get the lithium to ignite.
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March 13 2012 3 13 /03 /March /2012 14:36
After opening a dried-up lithium battery from which much of the lithium had already been depleted, I scraped some of the lithium off for a reaction with water. Then I heated the remaining lithium. A highly exothermic reaction occurred that did not, surprisingly, ignite the lithium to any degree as no flame was observed. The lithium metal only began glowing white hot as it oxidized completely to lithium oxide. No hydrogen was produced when the lithium ash was reacted with water. Here is the video:
 

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