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February 29 2012 4 29 /02 /February /2012 03:06

The disposal of some chemicals requires their destruction in many cases to make their disposal safer.

 

Cyanides: Cyanides are prone to oxidation. Reaction of cyanide-containing waste with sodium hypochlorite will form relatively nontoxic cyanate compounds. These can be flushed down a sink.

 

Chromates: Chromates are oxidizers. Therefore, reaction with ferrous sulfate (available in impure form as moss killer) will reduce the chromium to much less toxic chromium(III). Addition of sodium bicarbonate (baking soda) will further isolate the chromium as the completely harmless and insoluble chromium(III) hydroxide.

 

Copper, cobalt, nickel: You may want to consider precipitating copper metal from large amounts of pure copper waste as copper commands a high price. However, if the copper solution is contaminated, large amounts may be disposed of by reaction with sodium bicarbonate to produce the insoluble copper(II) carbonate, which can be disposed of in solid waste. Do a similar process for cobalt, nickel, titanium, and molybdenum.

 

Silver: Large amounts of silver are certainly undesirable to dispose of, but if they must be disposed, reduce the silver solution with sodium hydroxide and glucose or zinc before disposal. Forget about disposal and recover the silver metal.

 

Permanganates: Permanganates are quite unstable and will spontaneously decompose. But if brown stains are undesirable, add iron(II) sulfate solution, along with some hydrochloric acid. This may produce chlorine, but in the process the permanganate will be reduced to light pink manganese(II) solution. This can then be neutralized with sodium bicarbonate and the resulting precipitate disposed of in solid waste.

 

Mercury, cadmium, lead, thallium, arsenic: It is best to take these to a solid waste disposal area or use quantities of such small size (10 mg per experiment for mercury, thallium, arsenic) that no significant pollution could occur.

 

Antimony: Antimony naturally exists as stibnite, antimony trisulfide. Reaction of an antimony solution in dilute hydrochloric acid with iron(II) sulfide will precipitate antimony in this relatively harmless and reddish form.

 

Zinc: Zinc is not extremely toxic, therefore for reasonable amounts no special disposal means are necessary.

 

Magnesium, sodium, calcium, strontium, etc: These can be flushed down a sink in great quantities as they are completely harmless.

 

Barium: Add Epsom salts to precipitate barium as the insoluble barium sulfate. Wash away.

 

Beryllium: Beryllium naturally exists as the silicate. Precipitation as the silicate is likely the wisest method of disposal of small quantities of this valuable material.

 

The rare earth metals, as well as aluminium, gallium, indium, the alkali metals, silicates, phosphates, may be flushed in reasonable quantities down a sink without causing any guilt as to toxic effects on environmental habitats.

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February 28 2012 3 28 /02 /February /2012 23:35

Today in my chemistry class I performed a lab which involved reaction of iron(III) nitrate with potassium thiocyanate solution to form a colored complex which is measured spectroscopically. Earlier, the formation of this complex was used as a qualitative test for iron, but it did not work well. The addition of the thiocyanate in that lab to a solution shown later to only contain manganese ions produced a white precipitate. Do any other metals form thiocyanates that are not as well known? Cobalt(II) thiocyanate is soluble, however, and so other transition metal thiocyanates can be considered to be soluble as well. The composition of that precipitate is a mystery.

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February 28 2012 3 28 /02 /February /2012 00:54

The production of sodium periodate is very simple. You need tincture of iodine and sodium hypochlorite. Add the sodium hypochlorite to the tincture of iodine. The triiodide ion is oxidized to hypoiodite, whichly quickly oxidizes to iodate, then gradually to periodate: I3(-) + 9 ClO(-) -> 3 IO3(-) + 9 Cl(-)  IO3(-) + ClO(-) -> IO4(-) Sodium periodate is sparingly soluble and crystallizes out as colorless crystals after a few minutes. Here is the sodium periodate. The tincture of iodine coloration has completely disappeared.

 

Sodium-periodate-crystals-2.JPG

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February 28 2012 3 28 /02 /February /2012 00:48

I once dissolved sodium bromide in a concentrated HCl-acidified solution of copper(II) chloride. The solution turned brown as a copper(II) bromo complex was produced. When diluted with water, a dramatic color change occurred, ultimately forming blue copper(II) aqua ions (blue).

 

Copper bromo complex

 

Learn more here: http://woelen.homescience.net/science/chem/exps/copper_halogenide/index.html

 

Copper forms a variety of interesting complexes. Here is more information from woelen's website: http://woelen.homescience.net/science/chem/exps/cu-redox/index.html

 

Complexes of copper(I) often assume strange colors. I recognized the yellow color of the copper(I) oxide as the one produced when copper objects are electrolytically oxidized in aqueous sodium chloride solution.

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February 27 2012 2 27 /02 /February /2012 15:34

I am currently dissolving a bit indium in hydrochloric acid. The bubbling starts slow (like the speed of tin) and then speeds up to  the speed of the dissolution of iron. This is expected based on the electrode potential of indium metal. This reaction is occurring: 2 In + 6 HCl --> 2 InCl2 + 3 H2

 

Indium-dissolving.PNG

 

The indium chloride solution is then neutralized with sodium carbonate to form indium(III) hydroxide, which is difficult to filter.

 

Indium-III--hydroxide.JPG

 

Indium chemistry is very boring: only the element is interesting. Reduction of indium chloride by zinc might produce indium powder which can be pressed together to form a piece of indium metal; this could be an interesting demonstration of the softness of indium.

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February 26 2012 1 26 /02 /February /2012 02:54

I tried dissolving my beryllium lump. Placement in 5% acetic acid resulted in no reaction. In this regard beryllium (and its protective layer) are very similar to aluminium. Placement in concentrated hydrochloric acid, however, resulted in a vigorous reaction. The surface coating was etched and the unique crystal structure of the beryllium is visible. The beryllium is $18.00, so any significant period of dissolution for no purpose is a waste of money. I plan to make some beryllium hydroxide paste in the future, however. The peculiar sheen and darkness of the beryllium after dissolution is not captured well on camera, yet here are some pictures of it.

 

Beryllium-after-dissolution--1-.JPGBeryllium-after-dissolution--2-.JPG

 

 

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February 26 2012 1 26 /02 /February /2012 02:42

Recently I produced zinc- and aluminium-gallium alloys. There are two ways to do this with zinc. The first method is to react zinc with an acidic solution of a gallium salt. The gallium produced dissolves in the zinc metal to form the alloy containing a vast majority of zinc. It is highly brittle. Another way is to clean zinc by short immersion in hydrochloric acid and then place galinstan on the zinc. The galinstan will absorb into the zinc, forming an alloy containing a large minority of galinstan. The properties of this high-galinstan zinc alloy have not been determined yet.  Aluminium alloys are made by scratching aluminium and applying a bead of galinstan. The aluminium dissolves in the galinstan as the galinstan dissolves in the aluminium. The protective layer is removed and grayish aluminium oxide growths appear. A low-galinstan aluminium alloy slowly oxidizes and is very structurally unsound. A high-galinstan aluminium alloy is liquid or near liquid and reacts vigorously with water, meanwhile oxidizing rapidly in air. First is aluminium strip with galinstan on it. The aluminium oxide has completely covered the galinstan in a hemispherical shape. The next is aluminium foil with galinstan on it. The aluminium oxide coating has been partially mixed with the galinstan. The third is zinc with galinstan on it. The shiny areas have high galinstan concentration.

 

Aluminium-alloying-galinstan-bead-and-house.JPGAluminium-oxide-galinstan-mixture.JPGZinc-galinstan-alloy-natural-light.JPG

 

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February 24 2012 6 24 /02 /February /2012 19:47

Magnetite is a common iron ore. This ferromagnetic mineral has the chemical formula Fe3O4. Use a strong magnet covered in paper to sift through sand or dry clay powder. A significant number of magnetite particles should be sticking to the paper over your magnet. Place the particles in concentrated hydrochloric acid. They will dissolve to form a yellow solution. Add sodium bicarbonate (baking soda until a precipitate just begins forming, then add one drop of hydrochloric acid to dissolve the precipitate. This slightly acidic solution of iron(II) and (III) chlorides can then be reduced with zinc, magnesium, or electrolysis to produce a magnetic precipitate of iron metal.

 

Of course, this complete process has no industry uses because 1) zinc or magnesium is much more expensive to extract than iron 2) zinc is much rarer than iron 3) wet chemistry is not as easy to perform as dry chemistry 4) iron rapidly corrodes in wet environment. However, pure iron is produced by electrolysis of an aqueous solution of an iron salt.

 

Taking this a step further in the engineering direction will involve obtaining coke and limestone and producing a miniature blast furnace.

 

Taking this experiment a step in the interesting direction will involve obtaining fine aluminium powder, grinding it in a mortar and pestle with the magnetite particles, and igniting with magnesium ribbon in a crucible (thermite reaction, research before reacting)

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February 24 2012 6 24 /02 /February /2012 19:18

1-18 - Use Control + F to find what you want in here

Bismuth reactions: The bismuth precipitate has turned pure white upon drying. Reaction with bleach seems to produce no oxidation to the tetroxide. Addition of hydrogen peroxide, however, turns it yellow. Is this yellow solid a peroxo complex of bismuth?  Below is the “tetroxide” (left), the peroxo complex (center), and the precipitate (right).

Cobalt precipitation: The magnesium has completely precipitated the cobalt from the solution of cobalt(II) chloride.

Bismuth work: I dissolved some of the bismuth precipitate in hydrochloric acid. It appears to have retained no carbonate ions, and is almost wholly bismuth(III) hydroxide. I then added sodium iodide solution to the bismuth solution. A deep yellow solution was obtained. Upon neutralization with sodium bicarbonate, a bright yellow precipitate formed which got duller yellow as the pH rose. This solution is likely a solution of bismuth triiodide in hydrochloric acid. All of this bismuth chemistry was just done with a few tiny pieces of bismuth. The largest is barely thicker than a pencil lead. These bismuth chunks were placed in sodium hypochlorite solution for half an hour. They were then removed, washed, and immersed in hydrochloric acid. No chlorine gas was produced, showing that no higher bismuth oxide was produced by the hypochlorite.

Bismuth dissolution: The remaining bismuth pieces are placed in a 2:1 mixture of 3% hydrogen peroxide and 5% acetic acid. Dissolution appears to be very slow, with a small amount of cloudiness forming. Cloudiness is most likely the result of bismuth acetate hydrolysis. Neutralization with sodium bicarbonate yields no precipitate.

1-19

Bismuth-copper(II) chloride reaction: Bismuth is reacted with aqueous copper(II) chloride solution. A small amount of gas is evolved, and the solution becomes cloudy. The reaction then ceases. Addition of hydrochloric acid dissolved the precipitate, which is probably a mixture of copper(I) chloride and bismuth(III) oxychloride, but no further reaction with the bismuth occurs. When precipitated with sodium bicarbonate and excess ammonia, a cloudy copper(II) ammine complex is formed. This means that the bismuth(III) hydroxide is behaving like a gelatin. It might precipitate out later, but it is unlikely that it will do so without first occluding some copper compounds.

Chlorate production: Bleach is warmed to disproportionate it. Then production is stopped.

1-20

Bismuth-copper(II) chloride reaction: A small amount of bismuth(III) hydroxide, contaminated with copper, settled out on the bottom of the copper(II) ammine solution. It was washed with ammonia. As is visible, all of the copper washes away, leaving only the white insoluble bismuth hydroxide behind.

Capacitor dissolution: The dielectric from an unknown capacitor was immersed in hydrochloric acid. No reaction occurred. Therefore, it is just an ordinary ceramic capacitor.

Chromate production: The metal oxide mixture produced by oxidizing the 18/0 stainless steel spoon is placed in sodium hypochlorite solution. A yellow coloration begins coming off the oxides, showing that some chromate is produced.

Tantalum capacitors: I obtained three tantalum capacitors. I cracked one open and it consists of an outer manganese dioxide layer with an inner core of pure tantalum. The manganese dioxide can be dissolved, leaving the clean tantalum behind. However, in the cold near-freezing outdoor environment (because of  chlorine production), the manganese dioxide dissolves extremely slowly. The tantalum core is hit with a hammer and seems shatter-proof. The wire in the center is actually made of pure tantalum. After dissolution for a while, the solution is washed down the drain. Despite the tantalum’s high density, it is taken along with the water flow and is lost.

Samarium-cobalt magnet: A cheap speaker verified to contain a samarium-cobalt magnet is ripped open and the magnet, which closely resembles a neodymium magnet, is removed. It dissolves in acetic acid similarly to a neodymium magnet as well. Supposedly the most common magnet contains a small percentage of iron (ugh) and copper as well. The iron will probably begin to dissolve in the acetic acid along with the samarium, just like in the mischmetal.

1-21

Periodate reduction: Addition of hydrochloric acid to wet sodium periodate crystals produces a yellow solution which seems to be iodine monochloride. NaIO4 + 8 HCl -> NaCl + ICl + 3 Cl2 + 4 H2O is a possible reaction. A smell of chlorine was produced, but no fizzing was observed as the crystals were damp and the hydrochloric acid dissolves chlorine well.

Tantalum cleaning: The dirty tantalum was placed in hydrochloric acid. Some chlorine production is observed, and a dark brown manganese(III) chloride solution begins seeping off the tantalum (because it is cold). Eventually, however, this solution will decompose releasing its chlorine and turn colorless. It does so faster when warmed, even in a hand. So it is warmed by a hot water bath. When the dissolution has been going on for several hours, it is stopped. The greenish-blue iridescent color produced when tantalum is anodized to form the dielectric layer is visible through the thin film of remaining manganese dioxide. The tantalum wire protruding from the pellet is also iridescent, showing that it was anodized as well.

Lead chromate production: Lead acetate solution is produced by dissolving lead in 1:1 mixture of acetic acid/hydrogen peroxide. This solution is reacted with the sodium chromate solution. A bright yellow precipitate is formed. This is filtered and dried.

Chromium peroxide formation: Hydrochloric acid is added to the chromate solution. No visible chlorine is evolved, showing that the bleach must have been almost entirely consumed. Hydrogen peroxide was then added. A deep purple-blue solution of chromium(VI) peroxide formed, which quickly decomposed into a green chromium(III) chloride solution with the evolution of oxygen. Addition of ammonia formed no precipitate of chromium(III) hydroxide because the solution was too dilute.

Basic lead chromates: When lead chromate is added to ammonia, it turns orange, showing the formation of a basic lead chromate. Sodium hydroxide will probably make it red.

Acidic lead chromates: When hydrochloric acid is added to lead chromate, it turns white, and a yellow solution is leached off.

Dichromate formation: Acetic acid is added to sodium chromate solution in an attempt to produce a dichromate solution but a yellow precipitate of what appears to be lead chromate (eyedropper contamination?) is thrown down instead.

Chromate reduction: A non-acidified chromate solution is reacted with hydrogen peroxide. The typical chromium(VI) peroxide forms, which is decomposed into a yellow-brown solution. Addition of hydrochloric acid produces a gray-green solution of chromium(III) chloride. More chromate solution is reduced with ascorbic acid. A purple-gray solution of chromium(III) chloride is formed, which becomes less purple as it ages. The remainder of the chromate solution is reduced with ascorbic acid and flushed down a drain.

Samarium-cobalt magnet dissolution: The magnet is removed, and the solution of samarium and iron chlorides is light brown. Ammonia is added, which precipitates the all-too-familiar dark green iron(II) hydroxide. Hydrochloric acid is added, and a red-orange solution that does not appear like an iron(II) or iron(III) chloro complex is produced. Zinc is added to reduce the iron and any other non-rare earth metals. The solution lightens amid fizzing of hydrogen. Once it is light, it is neutralized with sodium bicarbonate. A very light brown precipitate is formed which does not oxidize. Because of the presence of the zinc, any trivalent iron would be reduced to divalent iron, and a color change would be visible upon exposure of the hydroxide to air. However, there was none, and so the residue must be iron-free samarium carbonate. However, the precipitate appears to have darkened, but not turned browner.

End of experiments: This is probably the end of experiments for my winter break.

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February 24 2012 6 24 /02 /February /2012 19:18

This is the diary of all of the experiments I have done over the winter break. It would be nice if the dozens of pictures found in the diary could be posted online; however, it is a huge effort. Furthermore, many of the pictures are only embedded in the document. If you find any interesting experiments that you would like to duplicate, by all means do so. If you have further questions, ask under "CHEMISTRY QUESTIONS" in the Links section on bottom of the right column on this blog. I will be glad to inform you to the extent of my knowledge.

 

Use Control + F to find certain chemicals, like chromate, if you do not feel like scrolling through so many pages.

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