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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 19 2012 2 19 /03 /March /2012 14:51

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

 

Krypton, like all other noble gases, is a colorless, odorless, and tasteless gas. It is quite heavy and quite inert, but does form a few compounds. Krypton is much rarer in the atmosphere, from which it is extracted by distillation. Therefore, it is much more expensive. Krypton produces a gray-green light when introduced into an electric arc.

 

In element form: Krypton is used in many flashlight bulbs.

 

In compound form: No sources found.

 

Here is my sample of krypton. It is a couple of flashlight bulbs.

 

Krypton.JPG

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March 17 2012 7 17 /03 /March /2012 15:21

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

 

Bromine is a dark red liquid, one of the two elements liquid at STP on the periodic table. Bromine is a highly reactive member of the halogen group. For example, aluminium reacts violently with bromine when placed in it. Bromine has a high vapor pressure and will quickly evaporate. Bromides are similar to chlorides, although more easily oxidized. Bromine dissolves in water to form a red solution. Bromine forms a series of oxyanions which are strong oxidizers and somewhat toxic. Bromates were used to treat flour before they were banned or discouraged in many places. Bromine is easily extracted from bromide by reaction with concentrated sulfuric acid.

 

In element form: Sodium bromide reacts with sulfuric acid to make bromine vapor and liquid. It needs to be distilled in all-glass apparatus. Bromine water, still corrosive but much safer, can be made in an impure form by reacting a mixture of sodium bromide, sodium hypochlorite, and hydrochloric acid.

 

In compound form: Sodium bromide is used as a bromide reserve in spas and pools and can be cheaply bought.

 

Here is my sample of bromine. It is a clear plastic vial of bromine water, which later gets cloudy and corroded.

 

Bromine-water.JPG

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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 16 2012 6 16 /03 /March /2012 15:43

 
Interesting properties of titanium include its flammability and relatively low conductivity. Therefore, electrical current can easily provide the heat necessary to ignite titanium. A thin strip (0.5 mm) of titanium foil easily ignites when applied to a nine volt battery, for example. Since titanium has the strength of steel, this strip is not fragile and can be bent to form interesting shapes, which are followed by the flame. For example, it can be shaped into letters and placed on a piece of wood. When a switch is thrown, the titanium ignites at one end (hopefully, if it is cut to the correct thickness) and someone's name, for instance, is traced out by the flame. It can also be used as a fuse as well, where the thinnest part near the top electrode is ignited and the titanium burns down the wire.
 
Here is a YouTube video of the burning.
 
 
I also tried cutting a thin strip of niobium and connecting it to a nine volt battery. Despite its similar thickness and resistivity, the niobium did not burn. Instead, the niobium became red hot when current was applied. This shows that titanium is more flammable and susceptible to oxidation than niobium.
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March 16 2012 6 16 /03 /March /2012 15:38

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

 

Selenium, like most nonmetals, is quite a unique and interesting element. Selenium is necessary in small quantities in the human body, but toxic in larger amounts. The element itself has three forms: a glassy, black, nonmetallic form; a red, powdery, nonmetallic form; and a gray, malleable, semimetallic form. This gray form is the most useful and the most common form of selenium. Selenium has the interesting property of changing its conductivity upon exposure to light. Therefore, it was used in the first photodetectors (cadmium sulfide is used now). Selenium forms a highly toxic and smelly dioxide which is the primary selenium compound. Selenium forms binary molecular compounds as well as anionic covalent complexes like selenite and selenate. Selenides replace a small amount of sulfide in many ores, making selenium a common byproduct in processes like copper production.

 

In element form: Selenium rectifiers contain a thin layer of selenium metal, as well as old photocopier machines and light meters.

 

In compound form: Selenium sulfide is used in Selsun Blue shampoo. Organic selenium complexes, as well as the inorganic selenites and selenates, are both used in selenium vitamin supplements. Selenium toners contain sodium selenite.

 

I do not have any verifiable elemental selenium. I do have a few specks that precipitated out when some selenium vitamin supplement slurry was electrolyzed, but of course they are not selenium. I plan to purchase pure elemental selenium in the near future, after which this page will be updated.

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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 15 2012 5 15 /03 /March /2012 14:48

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

 

Arsenic is a highly toxic metalloid that comes in several allotropes. The nonmetallic forms of arsenic are least stable, such as yellow arsenic, while the common metallic gray arsenic predominates. Gray arsenic is a dull solid. When heated in air, it burns to the trioxide. Arsenic dissolves in nitric acid to form the oxide as well. Arsenic is found in a wide range of minerals, and is almost always extracted as a byproduct. For example, arsenopyrite is iron arsenide sulfide. When this is roasted, both arsenic and sulfur oxides are produced. Arsenic was a common poison until an extremely sensitive test called the Marsh test was developed for it. This produces the gas arsine (AsH3) from arsenic-containing materials and later deposits the arsenic on a glass tube. Even a tiny trace of arsenic will produce a notable coloration on the glass surface.

 

In element form: Lead wheel weights for cars contain about 1/4% arsenic.

 

In compound form: Old treated wood contains Chromated Copper Arsenate as the treatment chemical. Infrared LEDs contain gallium arsenide as the semiconductor. High frequency microwave circuits may use gallium arsenide components. Chicken meat also can contain arsenic in the form of roxarsone.

 

Here is my sample of arsenic. It is one of innumerable lead wheel weights that are continually leaching into the environment from road abrasion after being loosened from vehicle tires.

 

Wheel-weight.png

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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 14 2012 4 14 /03 /March /2012 17:17

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

 

Germanium is a hard, inert, and brittle metalloid from Group 14 on the periodic table. Germanium has properties intermediate between tin and silicon. In the early days of electronics, germanium was the primary semiconducting material because purity is not essential for it to work. Silicon, because of its great abundance, has since superseded germanium in almost all electronics applications. Germanium forms quadrivalent covalent compounds with very low boiling points. Germanium can be used as a nutritional supplement but this use is questionable. Germanium dioxide is a white amphoteric solid that can dissolve in both acids and bases.

 

In element form: Germanium windows were used for motion sensors in the past. Germanium diodes are still found in old electronics equipment.

 

In compound form: Germanium compounds are used in the phosphors for mercury vapor lamps. Germanium dioxide is used in fiber optic cables.

 

Here is my sample of germanium, small germanium crystals from old diodes.

 

Germanium-diodes.png

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