Overblog Follow this blog
Administration Create my blog
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

Repost 0
Published by LanthanumK - in Experiments
write a comment
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

Repost 0
Published by LanthanumK - in Experiments
write a comment
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.

Repost 0
Published by LanthanumK - in Experiments
write a comment
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.

Repost 0
Published by LanthanumK - in Experiments
write a comment
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

 

Repost 0
Published by LanthanumK - in Experiments
write a comment
June 9 2011 5 09 /06 /June /2011 15:10

Warning: Copper(I) chloride is toxic. Do not eat the precipitate or drink the solutions. Otherwise, everything is safe.

 

This is the beginning of the copper(II) chloride series of experiments. A wide range of experiments with copper(II) chloride can be done by the home chemist. Information on how to make copper(II) chloride can be found here: http://lanthanumkchemistry.over-blog.com/article-how-to-make-copper-ii-chloride-76079848.html

 

You will need:

 

Ascorbic acid (found as Vitamin C crystals)

Previously made copper(II) chloride

Filter paper, tissues, or coffee filter paper

A container

Copper wire

3% hydrogen peroxide, optional

 

Dissolve a small amount of copper(II) chloride in water. Add ascorbic acid. A color change from blue to green to a white precipitate is observed. This white precipitate is copper(I) chloride. Filter it. Copper(I) chloride is easily oxidized by air. Within a few minutes of being in air, the filter paper should start turning green again. Before it dries, it will be completely green. When I conducted this experiment, the copper(I) chloride took about 10 minutes to begin turning green. Make another small batch. Filter and let it dry until it just starts turning green. Add hydrogen peroxide to the filter paper. It should immediately turn green. The hydrogen peroxide oxidizes the copper(I) chloride to copper(II) chloride and copper(II) hydroxide. Both of these chemicals are green or blue-green.

 

Here are pictures of the reaction and the resulting white precipitate. Asc stands for ascorbate, the part of the ascorbic acid that reduces the copper(II) to copper(I). H+ stands for the acid part of ascorbic acid, which makes the solution turn greenish.

 

CuCl2 reduction stages

Method 2: Dissolve copper(II) chloride crystals in concentrated hydrochloric acid. Add copper metal and heat. This comproportionation reaction will occur: CuCl2 + Cu --> 2 CuCl. Add sodium bicarbonate to the resulting dark solution until a pure white precipitate falls out. The leftover solution will probably be greenish as a result of excess copper(II) chloride. This white precipitate is copper(I) chloride.

 

 

 

Repost 0
Published by LanthanumK - in Experiments
write a comment
June 8 2011 4 08 /06 /June /2011 13:40

Warning: Copper(II) chloride is toxic. Do not eat the crystals or drink the solution. Avoid getting the acidic solutions on your hands or face. To be safe, wear gloves and goggles. Keep crystals out of reach of children, unless they have extreme copper deficiency and need the copper (do not take this last sentence seriously).

 

This article will describe several processes that can be used for creating this chemical.

 

  • Mix hydrogen peroxide with hydrochloric acid. 30% peroxide is better than 3%. Add copper. A green solution of copper(II) chloride is produced. Evaporate and obtain the crystals. If you use 3% hydrogen peroxide, you may get next to nothing. As this experiment is very cheap, you can reproduce it with different levels of hydrogen peroxide and hydrochloric acid to find the ideal mixture. After trying this method, it works wonderfully. However, if not enough hydrogen peroxide is used, the copper reacts with the acidic copper(II) chloride solution to generate a dark green mixed copper(I)/copper(II) complex, which can cause problems as the salt dries, including the formation of an insoluble oxychloride or a monovalent chloride. Add a little more hydrogen peroxide if a dark green coloration is observed after dissolution is completed. If it changes color (gets more bluish), add more peroxide until it stops changing color. If the resulting blue solution is not clear, it means that there is not enough hydrochloric acid present. Add more HCl until the solution clears.
  • Dissolve sodium bicarbonate in water. Strip two copper wires. Place them in the solution. Connect a 12VDC power supply to them. Hydrogen and copper(II) hydroxide and copper(II) carbonate are produced. The hydrogen is seen as bubbles at the cathode. Oxygen may be released as bubbles at the anode. A blue precipitate of copper compounds begins to form at the anode. Scrape the wires periodically to remove loose precipitate. Eventually, the wire will become so encrusted with black copper(II) oxide that the electrolysis will stop. Filter and dry the precipitate. You may reuse the solution for more copper electrolysis experiments. This precipitate is mainly basic copper carbonate. Add hydrochloric acid to it until most of it is dissolved and put it out to evaporate. The rest of the copper carbonate should dissolve and the solution should be very dark green. After evaporation, beautiful copper(II) chloride crystals should be in the container. This method produces more CuCl2 than the previous method but is not as simple. The picture shows about 5 grams of green CuCl2 dihydrate produced by that process. I prefer this method because of its high yield and uncomplicated procedure (joke). Update: In reality, it is very low yielding and a highly complicated procedure.
  • Heat copper shavings and drop them in chlorine gas. This produces the brown anhydrous form of CuCl2.
  • For a solution containing both copper and chloride ions without any pure copper(II) chloride, just mix blue copper sulfate solution, which tends to be more readily available, with sodium chloride to form the more greenish copper chloro complex.

Copper-II--chloride.JPG

This copper(II) chloride will be used in later experiments.

Repost 0
Published by LanthanumK - in Experiments
write a comment
June 8 2011 4 08 /06 /June /2011 13:00

You will need:

3% hydrogen peroxide solution

Sodium bicarbonate (baking soda)

Mix ½ cup (115 mL) hydrogen peroxide and 10 g of baking soda in a cup. Let it sit until the hydrogen peroxide starts fizzling. Start your washer, pour the solution in, and let it run through a cycle (rinse only cycle is recommended). Add some scent if desired.

The chemical explanation of this process is: The hydrogen peroxide is adducted to the sodium bicarbonate, forming sodium percarbonate in situ. Sodium percarbonate is used in commercial washer fresheners.

Repost 0
Published by LanthanumK - in Experiments
write a comment
June 7 2011 3 07 /06 /June /2011 22:03

Warning: This experiment produces moderately toxic copper sulfate. Do not dump more than a small amount down the sink. Do not drink the solution or eat the crystals.

 

This article will detail the process of creating a dirt-cheap impure form of copper sulfate using these simple household substances: copper wire, magnesium sulfate (Epsom salts), a container, and a 12VDC power supply.

 

Dissolve the magnesium sulfate in warm water to make a concentrated solution. Strip two lengths of copper wire (stranded is better than solid) and immerse them in the magnesium sulfate electrolyte. Attach them to the 12VDC power supply. They should not be near each other. Turn the power supply on.

 

At once, large quantities of hydrogen bubbles should be seen coming from the negative electrode (cathode), while small quantities of oxygen come from the positive electrode (anode). A white flaky solid should form at the cathode, and a light blue solution should start forming at the anode. Remove as much of the white flaky solid, a magnesium hydroxide precipitate, as you can. This will prevent it from redissolving and stopping the copper sulfate from forming. When the white precipitate starts becoming bright blue, stop the electrolysis and remove the copper wires. What you have is a solution of copper sulfate mixed with a large quantity of magnesium sulfate. The copper wire functioning as an anode should be somewhat corroded, and the other wire should be unchanged. Evaporate the blue solution. It should become bluer and eventually slushy blue crystals should start precipitating. After a while of being exposed to air, they will turn a dirty green because the copper sulfate has lost part of its water of hydration.

 

If you want pure copper sulfate, just buy it. Making it will require purchasing sulfuric acid and electrolytically oxidizing the copper in a sulfuric acid bath (some copper may plate out again at the cathode). Making it is not very economical.

Repost 0
Published by LanthanumK - in Experiments
write a comment