Experiments

Friday 6 april 5 06 /04 /Apr 14:28

Many chemicals form hydrates, where water molecules are bound to the chemical molecules. These hydrates often have different colors and different properties. For instance, anhydrous stannic chloride is a corrosive, fuming liquid, while the pentahydrate is an acidic white solid.

 

Copper(II) chloride has two forms: a brown anhydrous form and a blue-green dihydrate form. Because of excess HCl, my copper(II) chloride hydrate is more green than blue. When it is heated gently, it turns brown and releases both water vapor and white fumes, showing that the excess HCl and the water have been released from the crystal structure.

 

Here is a picture of the resulting product, with the original dihydrate for comparison.

 

Copper-II--chloride-anhydrous-and-dihydrate.JPG

 

Since the dehydration required heat, the reaction is likely endothermic. Therefore, rehydration should be exothermic. It is. However, because the anhydrous copper(II) chloride had been heated too vigorously, there was some decomposition to HCl and CuO and the solution formed when it was placed in water was cloudy. Here is the video of the reaction when two drops of water are added to the bulk of the anhydrous substance.

 

I then tried burning a crystal of copper(II) chloride with a magnifying glass. The copper(II) chloride refracted the green coloration for a moment before they turned brown as they were dehydrated. Then they turned black and melted. The resulting liquid was quite mobile. When cooled, it forms a black amorphous solid.

 

Fused-copper-II--chloride.JPG

 

This dissolves in water, forming a mixture of copper(II) chloride solution (green), copper(II) oxychloride precipitate (green), and copper(I) chloride precipitate (white). The copper(I) turns to the oxychloride upon exposure to air. The black lump completely dissolves.

 

If I do the same with copper sulfate, I will post the results in this article.

By LanthanumK - Posted in: Experiments
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Wednesday 4 april 3 04 /04 /Apr 15:39
This experiment involves reduction of copper(II) oxide, a black water-insoluble powder, with magnesium metal. The reaction is CuO + Mg -> MgO + Cu. The reaction is quite exothermic. Do this away from all flammable materials. Do not get too close to the reaction as burning metal can spray all over with larger quantities. This reactions is commonly known as a thermite.
 
Grind a magnesium rod or bar into shavings or, better yet, powder using a knife or a file.
 
Precipitate copper(II) carbonate from a solution of a copper(II) salt by adding sodium bicarbonate. Filter, wash, and dry by heating. It will turn black as it decomposes to copper(II) oxide. If you just have copper metal, electrolytically oxidize the copper in a sodium bicarbonate solution to form copper(II) carbonate. Heat this paste after filtering and drying.
 
Using a mortar and pestle, grind equimolar amounts (approximately 3 g CuO to 1 g Mg) of these two chemicals together until the mixture is uniform.
 
Place on a metal surface in a small pile. Ignition using a ferrocerium flint can be dangerous as the hands will be very close to the mixture. A better method of ignition would be electric ignition using a very thin piece of titanium foil and a nine volt battery or a magnesium ribbon.
 
Despite the dangers, I used a ferrocerium flint to ignite the mixture. Since there was some wind, I was afraid that the thermite mixture would blow away after placing it out in the open piece of metal. Therefore, I ignited it right away.
If you will try this at home, please be more careful than I was.
The "Shoe for Comparison" picture malfunctions in the video.
 
I then tried a copper(II) sulfate reduction. Aqueous copper sulfate reacts quite vigorously with magnesium metal, but the crystals react yet more vigorously. I shredded some magnesium and ground the copper(II) sulfate - magnesium shaving mixture together with the end of a pen. I then placed it on the same metal dish in a small pile and added some more magnesium shavings on top. Since grinding magnesium is so difficult, I hardly produced enough magnesium on top to ignite the mixture, so I had to strike the ferrocerium for a couple minutes before I actually got it ignited. The top layer burnt with a dazzling white flame, but the bulk was still unburnt. It would not ignite. Here is the video of the ignition.
I retried this with more finely powdered magnesium and dehydrated copper(II) sulfate. The result was much better.


 
By LanthanumK - Posted in: Experiments
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Monday 2 april 1 02 /04 /Apr 17:17

Zinc is often used as a safe, stable, and dirt-cheap reducing agent for inorganic and preparative chemistry. For example, the easiest way to produce chromium(II) and vanadium(II) solutions is reduction of higher oxidation states with a mixture of zinc and hydrochloric acid. However, zinc can also function as an oxidizing agent. Stronger reducing agents such as magnesium and possibly aluminium are capable of reducing zinc(II) ion solution to solid zinc.

 

A couple of years ago, I dissolved chopped up pennies in vinegar to form zinc acetate. After several weeks, strong-smelling and damp crystals had separated out of the solution. I placed them in a container and never used them until recently. Then, I took some of these crystals, dissolved them in water, and placed the entire magnesium firestarter rod in them. Some gentle fizzing was observed, which is normal with metal salt-magnesium reactions. Then, a blue-gray film was beginning to form on the magnesium. I checked 12 hours later and noticed the piece of magnesium covered with crumbly growths of zinc. They were not spongy or metallic, and dried to a powder. After another day of drying, they were covered with a white hydroxide film because of their finely divided state. I photographed them in this state and then disposed of them.

 

    Magnesium-zinc-acetate--1-.JPG

 

 

By LanthanumK - Posted in: Experiments
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Saturday 31 march 6 31 /03 /Mar 14:04

I earlier tried to add hydrochloric acid to Pepto-Bismol. It decolorizes, but the solution remains too thick to allow for any extraction of the bismuth chloride solution. To fix this, I add Pepto-Bismol drops to a dilute hydrochloric acid solution and stir. The drops break up and decolorize, and the bismuth dissolves. The solution is filtered and reacted with zinc metal. The zinc metal turns black as a thin smear of bismuth forms on the surface.

 

However, although Pepto-Bismol is widely available and relatively cheap, this is not a good way to produce bismuth metal. The amount produced is so small and so finely divided that it just stains filter paper. Here is how the zinc looked.

 

DSCF0156 

By LanthanumK - Posted in: Experiments
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Thursday 29 march 4 29 /03 /Mar 15:04

I placed a piece of nickel metal from a spark plug in a moderately concentrated copper(II) sulfate solution and let it sit overnight. No reaction was observed in the morning.

 

I then placed the nickel in some water and added a copper(II) chloride crystal. It dissolved, leaving a highly concentrated layer of copper(II) chloride solution on the bottom of the vial. The nickel metal was right in this solution. It seemed that a tiny amount of copper was produced, but even if a reaction was occurring, it was so slow that it would take weeks to get any significant amount of copper from the nickel reaction. Even placement in a boiling water bath did not produce any reaction between the copper(II) chloride and the nickel metal.

 

Nickel is quite an inert metal, with an oxide coating protecting against oxidation to a significant extent. It seems that this oxide coating is quite resistant to copper compounds.

By LanthanumK - Posted in: Experiments
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