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June 14 2011 3 14 /06 /June /2011 13:23

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

 

Nitrogen is a colorless, odorless, tasteless, relatively inert gas. It reacts directly and easily with very few elements, lithium being the only metal that reacts with nitrogen at room temperature. Some other metals form nitrides when they burn in air, such as magnesium. Nitrogen forms a wide array of compounds. Ammonia is used in cleaning. All proteins contain amino acids, which contain nitrogen. Hydrazine is used as a rocket fuel. Sodium azide is used in airbags. Nitrous oxide is used in canned whipped cream. Nitrites are used to preserve food. Nitrates are used as oxidizers in pyrotechnics. Most explosives contain nitrogen compounds.

 

In element form: Heat ammonium dichromate. The corrosion of a metal in a test tube inverted in water consumes the oxygen, leaving mostly nitrogen behind.

 

In compound form: Nitrates are common fertilizers. Ammonia is one of the most common nitrogen compounds.

 

Here is my sample of nitrogen gas. It was created by the corrosion of steel wool in a test tube, absorbing oxygen and leaving 95% nitrogen behind.

 

Nitrogen.JPG

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June 13 2011 2 13 /06 /June /2011 14:32

Warning: Copper compounds are toxic. Do not eat or drink. Hydrochloric acid is corrosive. Wear gloves.

 

You will need:

 

Copper(II) chloride

Hydrochloric acid

Test tube

 

Copper(II) chloride forms chloro complexes with hydrochloric acid. They range in color from green to yellow. Add hydrochloric acid to a copper(II) chloride solution. It will turn yellow. This is the CuCl4(2-) ion. When diluted with enough water, it will form the normal Cu(H2O)(2+) ion again, which is green-blue.

 

CuCl2-equilibrium.JPG

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June 13 2011 2 13 /06 /June /2011 14:21

Warning: Copper compounds are toxic. Do not eat or drink them. Alkalis are corrosive. Wear gloves.

 

You will need:

 

Copper(II) chloride

Sodium bicarbonate

Sodium carbonate

Sodium hydroxide

 

Dissolve copper(II) chloride (how to make it is found at http://lanthanumkchemistry.over-blog.com/article-how-to-make-copper-ii-chloride-76079848.html) in water. Divide it into three parts. To one, add sodium bicarbonate. You will form copper(II) bicarbonate, which decomposes instantly, releasing carbon dioxide and converting to copper(II) carbonate.

 

To the next solution, add sodium carbonate. A blue-green precipitate of copper(II) carbonate will be formed. Basic copper(II) carbonate, where some of the carbonate is replaced with hydroxide, may be formed as well.

 

To the next solution, add sodium hydroxide. A gelatinaceous blue precipitate of copper(II) hydroxide forms, which decomposes when wet for an extended period of time to black copper(II) oxide. It reacts with carbon dioxide in the air to make basic copper carbonate.

 

Filter and dry all of these precipitates. Heat them in a test tube or an aluminium foil boat (if you are cheap). They will turn black and create copper(II) oxide, releasing water vapor, carbon dioxide, or both.

 

The picture on the left is copper(II) carbonate. The picture on the right is copper(II) oxide.

 

Copper-II--oxide--2-.JPGCopper(II) carbonate

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June 13 2011 2 13 /06 /June /2011 14:14

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

 

Carbon has several physical forms, or allotropes. Graphite is black, soft, electrically conductive solid. Diamond is a clear, hard, electrically insulating solid. There are other forms like carbon nanotubes and fullerenes that are interesting but too expensive for such a common element. Graphite is used in pencils and carbon-zinc batteries. Diamonds are used as abrasives and jewelry.

 

In element form: Obtain graphite from an art supply store. Buy an artificial diamond. Open a carbon-zinc cell and remove the carbon rod. Use the impure carbon in a pencil lead.

 

In compound form: All organic compounds contain carbon.

 

Here is my sample of carbon. It is composed of graphite rods from a carbon-zinc cell.

 

Carbon rods

 

 

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June 11 2011 7 11 /06 /June /2011 16:48

Warning: Chloralkali process can release toxic chlorine vapor and sodium hydroxide, which is corrosive. Research conditions before attempting electrolysis on a substance. Only release halogens outside, in a fume hood, or in a test tube.

 

The chloralkali process can produce sodium hydroxide, hydrogen, chlorine, sodium hypochlorite, and sodium chlorate from sodium chloride and water. If potassium chloride is used, then the potassium compounds are formed. Here is a sample design to get you started in building chloralkali apparatus.

 

You will need:

 

Two glasses

Salt

Water

U-shaped hose

Carbon rod

Screw, nail, etc.

Two wires

24VDC power supply

 

Take two sturdy glasses that do not tip over easily. Fill one with salt water and the other with fresh water. Take a piece of hose shaped like a U and fill it with salt water. Plug the ends with tissues. This is called a salt bridge. Insert it into the two solutions so it connects them like a liquid wire. Obtain a carbon rod and any iron object. Insert the iron object in the fresh water and the carbon rod in the salt water. Attach wires to the electrodes and run them to a 24VDC power supply. Higher voltages may be used, but there is a risk of severe electrical shock. Run the fresh water electrode to the negative and the salt water electrode to the positive. Turn on the power supply. Nothing appears to happen at first. Very small amounts of bubbles start forming at the cathode (negative electrode). Eventually, sodium ions start trickling in from the anode (positive electrode) container. The reaction speeds up and keeps at a moderate speed. Sodium hydroxide and hydrogen is produced at the cathode. Chlorine gas is produced at the anode. If you do not want the chlorine gas, find a large silver-colored screw. Use that instead of the carbon rod. The chlorine gas will oxidize the iron and form large quantities of iron(III) chloride, which hydrolyzes to iron(III) hydroxide. The screw will be eaten away. Even stainless steel cannot withstand this corrosive action. The reaction produces heat. Heat evaporates the water. Continue to replenish the water.

 

Modifications: To make chlorates, do the electrolysis in one glass of salty water with two carbon rods. The electrolysis is very fast and large amounts of chlorine gas are released. The glass will get hot. Make sure the glass is steaming hot or hypochlorites will form instead. To make hypochlorites, do the first modification, but place the glass in an ice bath.

 

There are many alterations that can be made to this process to make it more efficient, quicker, etc. Bromides and iodides can be used as well as other cations.

 

Here is a picture of my chloralkali process. This is the no-chlorine modification. The iron(II) and iron(III) hydroxides are visible in the left container, which was stained permanently. The salt bridge is visible in the middle.

 

chloralkali.JPG

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June 11 2011 7 11 /06 /June /2011 15:33

Warning: Copper compounds are toxic. Do not eat or drink them. Do not get ammonia on your hands.

 

You will need:

Sodium carbonate

Sodium citrate

Copper sulfate

Copper(II) chloride

Sodium sulfite

Aluminium

Ammonia

 

In this page (http://lanthanumkchemistry.over-blog.com/article-experiments-with-copper-i-chloride-76297411.html), the fourth experiment showed that microcrystalline copper(I) oxide was formed when copper(I) chloride was reacted wtih a strong alkali. There is another form of copper(I) oxide though, that is red.

 

Method 1: Produce a solution of Benedict's reagent by adding 100 mg of sodium carbonate, 173 mg of sodium citrate, and 17.3 mg of copper sulfate pentahydrate to 1 mL of water. Add some glucose, fructose, or sucrose. The sucrose must have previously been heated with hydrochloric acid, or it will not work. The blue solution should turn dark red. Add water to the solution to dissolve any excess sugar and filter the precipitate.

 

Method 2: React a solution of copper(II) chloride with sulfur dioxide (can be made from sodium metabisulfite crystals + concentrated hydrochloric acid). This will make copper(I) oxide. Source: Wikipedia. Actually, try sodium sulfite instead of sulfur dioxide.

 

Method 3: React copper(I) chloride with any strong alkali. Yellow copper(I) oxide will form.

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June 11 2011 7 11 /06 /June /2011 15:20

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

 

Boron is an element that exists in two forms. One is an extremely hard black crystalline substane. The other is a red amorphous powder. Both are slightly toxic. Boron is a semimetal that forms covalent compounds with the halogens. Boric acid is the only acid safe to put in one's eyes. Boranes can be used as fuels. Boron is difficult to extract and does not have many uses as an element.

 

In element form: Neodymium magnets have about 6% boron in them. This could be considered an alloy or a chemical compound, but it behaves more like a alloy than a chemical compound.

 

In compound form: Borax is made of sodium borate. Heat resistant glass is made of borosilicates. Boric acid is used in some insecticides and eye washes. Boron carbide is used in some bulletproof vests.

 

Here is my sample of boron. This brown boron powder, insoluble in hydrochloric acid, was made by dissolving a neodymium magnet in hydrochloric acid.

 

    Boron-powder.JPG

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June 10 2011 6 10 /06 /June /2011 14:10

Warning: Strong bases can turn your fat into soap. Even if you want to lose weight, do not apply strong bases to your skin. Ammonia vapors are somewhat toxic. They rise in air, so always keep an ammonia bottle on a high shelf.

 

There are two alkalis used in the home lab: a strong alkali and ammonia. That strong alkali can be sodium, lithium, or potassium hydroxide. There are several ways to obtain these alkalis.

 

For ammonia: Just buy it. It is very cheap. Always buy unscented ammonia. If you feel like making ammonia, heat an ammonium salt with a base and absorb the ammonia gas in water. You may also want to distil the household ammonia to produce concentrated ammonia.

 

For lithium hydroxide: Take apart a AA or coin cell lithium battery and add water on the lithium. A white water-soluble solid will be formed, along with some choking fumes. Either keep the white solid in a sealed container and use it when needed, or dissolve it in water and keep that solution in a sealed container for use when needed. If it comes out of solution after being in solution for a few days, it is absorbing carbon dioxide from the air around your unsealed containre. Now it is lithium carbonate, which is not an alkali. Keep it sealed! Lithium hydroxide is used to absorb carbon dioxide in space ships.

 

For sodium hydroxide: Obtain drain cleaner or oven cleaner that contains sodium hydroxide. Some drain cleaners also contain aluminium particles, so be careful with your choice. The chloralkali process (seen here) can be used to make sodium hydroxide solution. It can also be bought from soap-making suppliers. Sodium hydroxide can also be made by adding calcium hydroxide to a sodium carbonate solution until the solution (not the precipitate) no longer fizzes when vinegar is added. Filter and keep the solution. Keep it out of air to prevent conversion to sodium carbonate.

 

For potassium hydroxide: Open all of your dead alkaline batteries and pour the contents into water. The potassium hydroxide electrolyte will leach into the water. Since potassium hydroxide has a strong attraction for water, you can keep adding battery powder to your solution, filtering, adding more battery powder, filtering, etc. Do not keep the potassium hydroxide solution in air for any length of time, or it will convert to potassium carbonate. Potassium hydroxide is deliquescent, so it is best to leave it in solution.

 

Although calcium hydroxide is not an alkali, it is used to make alkalis. For small amounts of calcium hydroxide, purchase pickling lime from the grocery store. For large amounts, purchage quicklime used to sweeten soil. This is actually calcium oxide. If it gets hot when it is mixed with water, it is calcium oxide. Adding water to calcium oxide produces calcium hydroxide, which is known as slaked lime.

Lithium-hydroxide--2-.JPG

This is lithium hydroxide produced by the Li-water reaction. It is stored in a closed container.

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

Lithium-after-reacting-with-water.JPG

Lithium from a CR2450 battery after reacting with water.

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

I used LiOH for concentrated alkali procedures as my KOH had turned to K2CO3 before I knew to keep it in a closed container. Dilute NaOH solution from the chloralkali process was also available.

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June 10 2011 6 10 /06 /June /2011 13:51

Warning: Copper(I) chloride is toxic. Do not eat the precipitate or drink the solution. Ammonia vapors are strong. Duck when opening an ammonia bottle as ammonia vapors are lighter than air. Hydrochloric acid can burn skin if splashed on it. Use gloves. Alkalis can burn skin if splashed on it. Use gloves.

 

In an earlier post (http://lanthanumkchemistry.over-blog.com/article-how-to-make-copper-i-chloride-76194624.html), copper(I) chloride was made. I know you probably used it all up, so go ahead and make it again. Keep it in solution this time.

 

Materials necessary:

Containers or test tubes

Sodium hypochlorite (household bleach) for experiment 1

2 copper wires and

12VDC power supply for experiment 2

Household unscented ammonia for experiment 3

Lithium or sodium hydroxide for experiment 4 (how to make alkalis: http://lanthanumkchemistry.over-blog.com/article-how-to-obtain-alkalis-76299009.html)

Hydrochloric acid for experiment 5

 

Copper(I) chloride is a white insoluble compound formed by the reduction of copper(II) chloride with ascorbic acid. In this set of experiments, I will show that:

 

Bleach is an oxidizer.

Copper(I) has a different color than copper(II).

Copper(I) can be reduced further to copper metal.

Copper forms three oxidation states.

Copper(I) forms an ammine complex.

Copper(I) forms an oxide with bases.

Copper(I) forms a chloride complex.

 

First: Take some copper(I) chloride solution and add sodium hypochlorite. It will oxidize the white copper(I) chloride to green copper(II) hydroxide.

Second: Place the two copper wires in a copper(I) chloride solution. Attach them to your power supply. Turn it on. There should be red copper metal forming at the negative electrode, along with hydrogen. There is still the white copper(I) in the middle. You may see copper(II) forming at the anode as the copper wire is oxidized. Here are all the three oxidation states of copper.

Third: Add ammonia to copper(I) chloride solution. It clears up. After a while, it starts turning blue on the surface. This is the oxidation of the colorless copper(I) ammine complex to the bright blue copper(II) ammine complex by the oxygen in tha air.

Fourth: Add your alkali solution to the copper(I) chloride solution. A bright yellow to orange precipitate is formed. This is microparticulate copper(I) oxide. It will oxidize in air just like copper(I) chloride. In a later experiment, the red copper(I) oxide will be formed.

Fifth: Add hydrochloric acid to the copper(I) chloride solution. It will dissolve. This is a copper(I) complex with chloride ions.

 

In this experiment, you have seen oxidation by oxidizing agent, oxidization and reduction by electrolysis, formation of an ammine complex and oxidation of that ammine complex, formation of an oxide by reaction with a base, and a water-soluble chloride complex.

 

This picture is the microcrystalline copper(I) oxide formed by reaction of copper(I) chloride with alkali:

 

Cuprous-oxide-microcrystalline-2.JPG

 

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June 10 2011 6 10 /06 /June /2011 13:26

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

 

Beryllium is an insanely light, strong, gray metal. It is the first alkaline earth metal. It has properties similar to aluminium, but it is much more expensive. Beryllium seems unreactive because it forms a protective oxide layer, similar to aluminium. It is very toxic when powdered and inhaled, and somewhat toxic when ingested. Its compounds are sweet, but do not eat them! I initially did not have any beryllium, but gave in and purchased a small piece of this expensive metal.

 

In element form: X-ray windows are made of pure beryllium foil. Some pieces of aircraft and other military equipment are also made of beryllium. 0.5% to 3% beryllium is used to harden copper in an alloy called beryllium copper. This alloy is used to make nonsparking tools. Some high-end speakers use beryllium or beryllium alloys.

 

In compound form: The safest form of beryllium is the mineral beryl, beryllium aluminium silicate.

 

This is my sample of beryllium. It is a 1.1 gram beryllium lump purchased from GalliumSource for $15.00 plus shipping.

 

    Beryllium (4)

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