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June 21 2011 3 21 /06 /June /2011 17:13

Warning: The lithium hydroxide produced when lithium reacts with water is caustic. Do not contact. Only the lithium and the calcium cannot be recovered unharmed form the water. 

You will need:

 

One or more of these: lithium, calcium, magnesium, aluminium, zinc, iron, nickel, tin, lead, copper, silver, gold

Water

Lots of test tubes

 

For consistency, all of the metals should be about the same shape and size. Zinc powder will react more vigorously than an aluminium chunk, even though aluminium is higher on the reactivity series.

 

There is a series called the reactivity series. It ranks the metals from most reactive to least reactive. The most reactive elements react with water, oxidizing acids, and nonoxidizing acids. They are isolated by electrolysis. The second group reacts with oxidizing and nonoxidizing acids. They can be isolated by a carbon reduction. The third group reacts only with oxidizing acids. They are very easy to isolate. The list of metals above is ranked in the order of the activity series.

 

Add warm water to all of the test tubes. Add the metals, starting with gold. Gold and silver have absolutely no reaction. Copper and lead very slowly corrode because of the dissolved oxygen in the water (not because of the water), both forming a protective layer that inhibits further corrosion. Tin and nickel do not react. Iron and zinc corrode a little more readily, although zinc forms a partially protective layer. This is still corrosion by the dissolved oxygen, not by the water itself. Aluminium has a tough protective layer, making it completely unreactive. Magnesium has a protective layer not as strong as aluminium. Since magnesium is so high on the reactivity series, it is corroded by water itself, not by the dissolved oxygen in water. Small hydrogen bubbles can be seen escaping from the magnesium. Calcium forms an incomplete protective layer, but still reacts quite vigorously with water. Lithium, the highest on the reactivity series, forms no protective layer and (if it is not covered in oil) reacts instantly and vigorously with water.

 

The nature of the oxide (crumbly or adhesive) affects the reactivity of the metal as well as its placement on the reactivity series. The oxide layer can be removed by treatment with hydrochloric acid and the true reactivity of the metal can be seen.

 

 

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June 20 2011 2 20 /06 /June /2011 12:51

Warning: Copper(II) salts are toxic. Do not eat or drink.

 

You will need:

 

Copper sulfate

Sodium chloride

Aluminium foil

 

Copper forms a chloro complex with sodium chloride as well. This chloro complex, as we saw before, is reactive toward aluminium.

 

Dissolve the copper sulfate and sodium chloride in water. Add a piece of aluminium foil to each solution. On the copper sulfate side, you may see a slow reaction. Mix the solutions by pouring one into the other. This will form the green copper chloro complex instead of the hydrated copper complex. The aluminium will start bubbling and dissolving faster.

 

You can try making bromo complexes as well and testing their reactivity. Use sodium bromide instead of sodium chloride.

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June 17 2011 6 17 /06 /June /2011 15:03

Warning: Many metal oxide fumes and other oxides are harmful for inhalation. Do not burn yourself on the hot materials. Arsenic, selenium, lead, and tellurium are quite toxic.

 

You will need:

One or more of these: lead powder, bismuth powder, tellurium powder, selenium powder, antimony powder, arsenic small piece, titanium powder, magnesium ribbon, lithium piece, zinc small piece, copper powder, molybdenum, tungsten, sulfur

Source of heat

 

Many elements burn or oxidize in air when heated. This produces a colored flame, which may be the same or different from the flame seen in a flame test.

 

Lead powder burns with a grayish flame, making highly toxic reddish smoke of lead(II,IV) oxide. The smoke is not visible in this picture.

 

Lead-powder-burning--5-.JPG

 

Bismuth burns in air with a blue flame, making yellowish smoke of bismuth(III) oxide.

 

Tellurium burns in air with a green-blue flame, making white smoke of tellurium(IV) oxide. If the smoke is heated enough, it melts to a red liquid.

 

Selenium burns in air to make variously colored (yellowish white, reddish white, or white) smoke of selenium(IV) oxide. It has a garlicky odor and is highly toxic.

 

Antimony burns in air to make white fumes of antimony(III) oxide. It dissolves slightly in water.

 

Arsenic burns in air with a pale lavender-colored flame to make white fumes of arsenic(III) oxide, which smells like garlic and is very toxic.

 

Titanium burns in air to make titanium(IV) oxide and titanium nitride. It burns nitrogen as well as oxygen.

 

Magnesium burns in air with a blindingly bright white flame to make white fumes of magnesium oxide. If oxygen is limited, magnesium nitride is formed.

 

Lithium burns in air with a reddish-white flame to make white lithium oxide.

 

Lithium-burning.JPG

 

Zinc burns in air with a blue-green flame to make white zinc oxide, which is yellowish when hot.

 

Zinc-burning.JPG

 

Copper powder oxidizes without a flame in air to make either red copper(I) oxide or black copper(II) oxide.

 

Molybdenum burns to make white or slightly bluish molybdenum(VI) oxide.

 

Tungsten burns when heated to make yellow-green tungsten(VI) oxide.

 

Sulfur burns when heated with a blue flame to make gaseous sulfur(IV) oxide.

 

Compare the metallic character of the element with the melting point of the oxide. The more nonmetallic an element is, the lower the oxide's melting point is. The alkali metals are an exception. The higher oxidation state an oxide is in, the lower the melting point for metals and semimetals and the higher the melting point for nonmetals. Try to find other examples of periodicity as well. Don't breathe the fumes!

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June 16 2011 5 16 /06 /June /2011 13:36

Sodium-tetrachlorocobalt-II-.JPGWarning: Cobalt(II) chloride is toxic and potentially carcinogenic. Do not eat or drink. Do not dump cobalt chloride down the drain. Place it in the garbage if you have a very small amount.

 

You will need:

 

Cobalt(II) chloride

Sodium chloride or potassium chloride

Heat

Test tube

 

Dissolve the cobalt(II) chloride in water. It will make a rose-red solution. Add your chloride. You probably will not see any change. Heat it. You will see it turn blue. This is evidence for excess chloride ions. When it cools, it will turn back to rose again.

 

This is caused by a principle known as Le Chatalier's principle. This reaction is driven by heat: Co(H2O)6(2+) + 4 Cl(-) <=> CoCl4(2-) + 6 H2O It moves to the right when heated and to the left when cooled.

 

Sodium-cobalt-chloride.JPG

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June 16 2011 5 16 /06 /June /2011 13:12

Warning: Flames can burn. Keep flammables out of flames. Many metals and their salts are toxic. Do not eat or drink them. As long as you do not make more than 1/2 gram of manganese(II) chloride, you may dispose of it down the drain.

 

You will need:

 

One or more of these: copper(II) chloride, lithium compound (LiCl is best, but LiOH can be used), sodium compound (NaCl is best), potassium compound (whichever one is most abundant, KCl is best), calcium compound (CaCl2 is best), strontium compound, barium compound, zinc compound, iron compound, cesium compound, rubidium compound, boron compound, manganese compound.

Nichrome wire

Flame

 

Flame tests are an interesting part of home chemistry. They are colors produced when certain substances get heated. These colors are the result of heat exciting the electrons in the atoms of the chemical. They soon fall back down, emitting electromagnetic radiation. Some just happens to be in the visible spectrum.

 

A little word of preparation on the above chemicals:

 

CuCl2: Already stated here (http://lanthanumkchemistry.over-blog.com/article-how-to-make-copper-ii-chloride-76079848.html)

 

LiCl: React LiOH or Li2CO3 with HCl until pH is about neutral. Dry with heat.

 

NaCl: Table salt.

 

KCl: React KOH or K2CO3 with HCl until pH is about neutral and dry. Or use salt substitute.

 

CaCl2: Some ice melts have this in them. Or CaCO3 or Ca(OH)2 with HCl until pH is about neutral and dry.

 

SrCl2: SrCO3 from pottery store in HCl until pH is about neutral. Dry.

 

BaSO4: Ask a doctor for a barium meal and dry some of it.

 

ZnCl2: Dissolve zinc in HCl and dry with heat.

 

FeCl2: Dissolve iron in HCl and dry without much oxidation.

 

CsCl: Difficult to obtain.

 

RbCl: Difficult to obtain.

B2O3: Purchase boric acid eye wash.

 

MnCl2: Dissolve MnO2 in HCl, filter and dry.

 

You may already have most of these salts. If you do, then dip a nichrome wire in HCl and heat it. Repeat this process several times to get the wire clean. Then moisten the wire and dip it in the chemical. Some crystals should stick on. Heat it in the flame. You should see the flame color. CuCl2 bright green, LiCl crimson, NaCl yellow, KCl light lilac, CaCl2 red-orange, SrCl2 red, BaSO4 yellow-green, ZnCl2 light green, FeCl2 golden yellow, CsCl blue, RbCl red, B2O3 light green, and MnCl2 light yellow-green.

 

Later, an experiment similar to this will be done; only the later experiment is combustion and flame color, not just heating and flame color.

 

A flame test can also be made by adding some of each salt to FancyHeat(R) methanol gel. The salt will dissolve in the gel and produce the respective flame color when the gel is ignited.

 

 

 

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June 15 2011 4 15 /06 /June /2011 15:44

Warning: In this experiment, googles and gloves were worn. Hydrochloric acid can burn skin and dissolve several substances. Make sure the bottle does not tip over.

 

My younger brother wanted to make a torch. He decided that hydrogen would be the easiest obtained fuel. He made a hole in the cap of a shampoo bottle and ran a hose through a valve. The hose ended in a glass pipe. He took some HCl and added it to a shampoo bottle. Then the zinc was added. The hydrogen cleared the bottle of oxygen. Then he took a heating element and ignited the hydrogen. The first try was not too good since he used too little zinc. The next try had a large, diffuse flame. In this experiment, safety procedures were observed.

 

He could improve it by allowing the oxygen to mix with the hydrogen in the tube before ignition. This would produce a more concentrated flame. The shampoo bottle also softened as a result of the reaction's heat. A glass bottle would be more durable.

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June 15 2011 4 15 /06 /June /2011 14:29

Warning: Copper salts are toxic, and zinc is slightly toxic. Do not eat or drink.

 

You will need:

 

Copper(II) chloride or copper sulfate

Zinc

 

Dissolve copper(II) chloride or copper sulfate in water. Add zinc. You will see a red precipitate of copper forming on the zinc. Depending on the acidity of your solution, there may be hydrogen produced. The solution will eventually turn colorless if you added enough zinc. The colorless solution is zinc chloride or zinc sulfate. The copper was reduced and the zinc was oxidized.

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June 14 2011 3 14 /06 /June /2011 13:52

Warning: Copper salts are toxic. Do not eat or drink.

 

We have already seen the copper(I) ammine complex, which is colorless. It was formed when copper(I) oxide dissolves in ammonia. When it was oxidized, a blue solution of the copper(II) ammine complex was formed. Here, though, you will create a concentrated copper(II) ammine complex and see its real appearance.

 

You will need:

 

Copper(II) chloride or copper sulfate

Household ammonia

 

Dissolve copper(II) chloride in water. The color from the hydrated copper(II) ions is blue for dilute solutions and green for concentrated solutions. Add ammonia. First, the hydroxide ions react to make a blue precipitate of copper(II) hydroxide. Then, the ammonia molecule itself removes the water molecules from the hydrated copper(II) ions, making the copper(II) ammine complex. This is indicated by a dramatic darkening of the solution to a navy blue.

 

Here is a relatively dilute solution of tetramminecopper(II). It is dark blue, almost black in the thick parts. The blue can be seen in the thin parts.

 

Tetramminecopper-sulfate.JPG

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June 14 2011 3 14 /06 /June /2011 13:42

Warning: Many of these salts are toxic. Do not eat or drink them. Nickel compounds are known carcinogens.

 

You will need:

 

One or more of these: copper(II) chloride, copper sulfate, nickel sulfate, magnesium sulfate, cobalt(II) chloride, iron(II) sulfate,

Test tube

Source of heat

 

Many metal salts form hydrates, where water molecules are incorporated in the crystal structure of the salt. The water molecules can be disconnected by heating, in most cases. The appearance of the salts can change based on their water content.

 

Copper(II) chloride dihydrate is green or green-blue. Heat it. It turns brown. This is the anhydrous form.

Copper sulfate pentahydrate is blue. Heat it. It turns white. This is the anhydrous form.

Nickel sulfate heptahydrate or hexahydrate is blue-green. Heat it. It turns light green. This is the anhydrous form.

Magnesium sulfate heptahydrate is colorless. Heat it. It turns white. This is the anhydrous form.

Cobalt(II) chloride hexahydrate is burgundy. Heat it. It turns blue. This is the anhydrous form.

Iron(II) sulfate heptahydrate is green. Heat it. It turns white. This is the monohydrate.

Extremely optional: Chromium(III) chloride hexahydrate is green. React it with the dangerous thionyl chloride. It will form the purple anhydrous form.

 

 

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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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