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February 8 2012 4 08 /02 /February /2012 15:58

As this blog is far more successful than my current one, I plan to switch back to this one.

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September 2 2011 6 02 /09 /September /2011 21:13

As I begin college, I plan to start a new blog here that is not just about chemistry and experimenting. It will be found at http://lanthanumk.over-blog.com.

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July 17 2011 1 17 /07 /July /2011 00:06

I am not able to keep up this blog anymore. I would like to thank my readers. Please continue in your chemistry endeavours! You can still contact me by email, which is viewable on my public profile. I will be glad to answer your chemistry questions.

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July 8 2011 6 08 /07 /July /2011 14:14

Warning: Thionyl chloride is one of the only easily obtainable chemicals that rates a 4--Extreme Risk on the NFPA diamond for health. Do not bring it in the house in anything but a strong, sealed container. Noxious fumes are created in this experiment.


You will need:


Lithium-thionyl chloride battery (I obtained mine from an EZpass device in the woods)




Very very sealed container


Remove as much from the lithium-thionyl chloride battery as you can. These things are built very strongly so a drill is needed. After you remove all the extra stuff from the battery, you should see a carbon rod sticking out on one side. Drill it out. As you drill, you will see smoke and smell sulfur. This is the thionyl chloride around the carbon rod reacting wih the water in the air; it is not fire. After a nice hole is drilled, remove your smoking drill bit. Then, place the battery hole down in the sealed container and let the thionyl chloride drip out. The container may cloud up. You may have to squeeze the battery to get significant amounts of thionyl chloride out. Now you have thionyl chloride. It can be used for producing the anhydrous forms of several metal compounds. It also reacts vigorously with water. Place a drop of it in water. Bubbles of sulfur dioxide will leave, along with some HCl fumes. Both of these gases are very annoying to the nasal passages. Do what you want with the rest of the thionyl chloride.

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July 8 2011 6 08 /07 /July /2011 14:01

I have compiled a list of sources for the elements that are available to the amateur chemist. Cobalt will be discussed here.


Cobalt is the first metal discovered since antiquity. It is a hard silvery ferromagnetic metal that is somewhat reactive chemically. Its +2 aqua ion is pink-red and its +2 chloro ion is blue. Cobalt compounds are quite toxic. Cobalt blue was used as a glass colorant and blue pigment. Cobalt chloride is an indicator in color-changing Drierite.


In element form: Cobalt(II) chloride can be acidified and reduced with magnesium to obtain pure cobalt metal. Cobalt plating is used in a few things. Some tools are made of cobalt steel, which contains 5% cobalt. Alnico magnets contain from 5 to 24% cobalt.


In compound form: Indicating Drierite contains cobalt(II) chloride, as well as blue-to-pink moisture testing papers. Cobalt blue glass contains trace amounts of cobalt aluminate. Most lithium ion batteries contain lithium cobaltite.


Here is my sample of cobalt. It is reduced from acidic cobalt(II) chloride solution by magnesium metal.



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July 7 2011 5 07 /07 /July /2011 15:00

Warning: Copper and zinc salts are slightly to moderately toxic. Do not ingest them. Hydrochloric acid is corrosive.


You will need:


Ascorbic acid or sodium metabisulfite

Sodium bicarbonate

Carbon rod


Electrolytically oxidize the brass in sodium bicarbonate solution with a carbon cathode. Green copper carbonate and white zinc carbonate are formed. Filter and dissolve the precipitate in HCl. A dark green solution is formed. Add sodium bicarbonate until the pH of the solution is around 2 or 3. Then add ascorbic acid or sodium metabisulfite. White copper(I) chloride will precipitate. Filter. Now you have a copper(I) chloride precipitate and a zinc chloride solution. You have separated brass.

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July 7 2011 5 07 /07 /July /2011 14:55

I have compiled a list of sources for the elements that are available to the amateur chemist. Iron will be discussed here.


Iron is the most common structural metal in the world. It is very strong, cheap, and sturdy. It is the second most common metal in the earth's crust. It is rarely found in its pure form, but in the form of oxides, which may be reduced to the metal by carbon. Alloys of iron with other metals are ubiquitous. One of iron's biggest problem is its vulnerability to corrosion. Iron rusts readily in moist air, and much research has been done to find out how to make iron corrosion resistant. Iron compounds are used to process water. They are also used as pigments and in vitamin supplements.


In element form: Find an electromagnet core or electrolyze an iron(II) solution with an iron anode to obtain pure iron. All steels are made of more than 50% iron.


In compound form: Rust is iron oxide. Iron(III) chloride is used to etch circuit boards. Iron(II) sulfate is used as an iron supplement.


Here is my sample of iron. It is iron filings, one of the many sources of iron. I have several other iron compounds and objects but they are not featured here. Fe filings

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July 6 2011 4 06 /07 /July /2011 13:52

Warning: Hydrochloric acid is corrosive. The resulting metal chlorides are toxic or carcinogenic.


You will need:


Alnico magnet

Hydrochloric acid

3% or more concentrated Hydrogen peroxide

Sodium bicarbonate

Sodium hypochlorite

Ammonia, concentrated

Sodium hydroxide


Alnico magnets are made of an iron-aluminium-nickel-cobalt alloy. The process for separating them is detailed here.


Add hydrogen peroxide to concentrated hydrochloric acid in a 1:1 ratio. Add the alnico magnet. Bubbling is observed and a colored solution is formed. This solution contains these components:


Iron(III) chloride, yellow

Nickel(II) chloride, green

Cobalt(II) chloride, blue (in this HCl concentration)

Aluminium(III) chloride, colorless


Here are some properties of the chlorides:


Iron(III) chloride: Forms alkali-insoluble hydroxide, is not oxidized by bleach, hydroxide does not dissolve in ammonia

Nickel(II) chloride: Forms alkali-insoluble (at normal temperatures) hydroxide, is not oxidized by bleach, hydroxide dissolves in ammonia

Cobalt(II) chloride: Forms alkali-soluble hydroxide, is oxidized by bleach, hydroxide dissolves in ammonia

Aluminium(III) chloride: Forms alkali-soluble hydroxide, is not oxidized by bleach, hydroxide does not dissolve in ammonia


Based on these properties, a separation mechanism is thought out:


Add sodium bicarbonate to reduce the pH to around 2.

Add excess sodium hydroxide. Cobalt and aluminium dissolve, while nickel and iron remain behind. Filter.

Acidify the cobalt and aluminium solution to neutral and then add excess ammonia. A cobalt ammine complex will form and the aluminium hydroxide will precipitate. Filter and separate.

Add excess ammonia to the nickel and iron precipitate. Nickel should dissolve as the nickel ammine complex. Filter and separate.




Nickel(II) ammine chloride, blue solution

Iron(III) oxide, brown precipitate

Cobalt(III) ammine complex, orange solution

Aluminium hydroxide, white precipitate


You have successfully separated an alnico magnet.


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July 6 2011 4 06 /07 /July /2011 13:20

I have compiled a list of sources for the elements that are available to the amateur chemist. Manganese will be discussed here.


Manganese is a silvery hard corrosion-vulnerable metal. It is not commonly used as an element. Ferromanganese, the form in which it is sold, is about 85% manganese and 15% iron. Manganese compounds are used in batteries, pyrotechnics and pigments.


In element form: Reduction of manganese(II) chloride produced by dissolution of manganese dioxide in hydrochloric acid should produce manganese. (If the solution is yellow not light pink then there is contamination and you will reduce iron not manganese.) Manganese steel has 10-15% manganese in it.


In compound form: Impure manganese dioxide is found in alkaline and carbon-zinc batteries. Manganese chloride can be made by dissolving purified manganese dioxide in hydrochloric acid.


Here is my sample of manganese. It is produced by magnesium reduction of a manganese chloride solution.





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July 5 2011 3 05 /07 /July /2011 13:19

Warning: Cobalt compounds are toxic and potentially carcinogenic. Do not ingest or inhale them. Bleach fumes may be a problem. The procedures that release chlorine should be done under a fume hood or outdoors.


You will need:

Test tube

Cobalt(II) chloride

Bleach (sodium hypochlorite solution)  

Sodium carbonate

Source of heat


Hydrogen peroxide



Dissolve the cobalt(II) chloride in some water. You will see the burgundy color of cobalt aqua ions. Add bleach. It instantly turns brown-black and cobalt(III) oxide precipitates.  Some chlorine is produced. Filter and dry the cobalt(III) oxide. This oxidation state is only stable in the presence of certain ligands and in a basic state. Acidify the precipitate with hydrochloric acid. Chlorine is released and a cobalt(II) ion is formed. Depending on the amount of excess hydrochloric acid, the color of the resulting solution can be red, blue, or green.




Dissolve more cobalt(II) chloride in water and precipitate it with sodium carbonate. Filter and dry the cobalt carbonate produced and heat it. It decomposes to black cobalt(II,III) oxide when heated. It also becomes calcined, making it insoluble in acid. I placed the resulting substance in HCl and it only produced a tiny amount of cobalt(II) chloride before the acid evaporated.




Dissolve more cobalt(II) chloride in water. Add ammonia. A blue precipitate of cobalt(II) hydroxide will form. This is only one form of cobalt(II) hydroxide, but that is beside the point. Then add some hydrogen peroxide. The solution will turn greenish, then brown as cobalt(III) hydroxide precipitates. The precipitate will be light, fluffy, and insignificant, almost like aluminium hydroxide.




Do this reaction outside. Dissolve more cobalt(II) chloride in water to produce a concentrated solution. Use a carbon anode and an iron cathode. A vigorous electrolysis will result in the formation of a brownish-green precipitate of probably cobalt(II) oxide.





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