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July 2 2011 7 02 /07 /July /2011 13:34

Warning: This is a rather advanced experiment that requires the use of toxic gases and highly reactive and corrosive liquids. This procedure must be done in a fume hood or outside. Research thoroughly before completing. This experiment has not been verified by the author.

 

You will need:

 

Sodium hypochlorite (bleach)

Hydrochloric acid

Concentrated 98% sulfuric acid

Sulfur

Glass tubing

Stoppers

Erlenmeyer flasks
Sodium hydroxide

Source of heat

Dry ice

 

First, construct your apparatus. Flask 1 will have a tube running from it into Flask 2. Flask 2 will have a tube running from it to Flask 3. Flask 3 will have a tube running from it to Tub 1. Tubes are made of glass, but the connecting tubes can be made of rubber.

 

Add sodium hypochlorite to Flask 1. Add sulfuric acid to Flask 2 so that one tube is under the acid while the other is above. Add sulfur into Flask 3 with a small amount of dry ice. Add sodium hydroxide solution to Tub 1 and place the tube under the solution. Then, begin gently heating Flask 3. Because of the carbon dioxide, the sulfur should not ignite. It should instead melt into a light yellow thin liquid. Make sure one of the glass tubes is in the liquid sulfur. Keep the sulfur temperature low enough so that the glass tube does not get stuck in highly viscous hot molten sulfur. Then, add a small amount of hydrochloric acid to the bleach solution in Flask 1. You should get some bubbling in Flask 2, where the sulfuric acid dries the chlorine gas. The chlorine gas should then bubble into the sulfur, producing a chemical reaction to form disulfur dichloride if there is a little chlorine and sulfur dichloride if there is much chlorine. Continue adding hydrochloric acid and bleach, if necessary, to Flask 1 until the molten sulfur has become transparent but not clear. Then, wait for the sulfur chloride to cool. You should have formed a mixture of sulfur dichloride and disulfur dichloride.

 

Experiment 1

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July 1 2011 6 01 /07 /July /2011 21:43

Warning: Sulfur fumes are toxic. Do not inhale them.

 

You will need:
Sulfur

Sodium hydroxide solution

Glass container

Aluminium foil

 

First, burn the sulfur and create sodium sulfite as is shown here (http://lanthanumkchemistry.over-blog.com/article-production-of-sulfites-from-sulfur-78120865.html). Make sure the end pH is 9. Then boil sulfur powder with the sulfite solution and recrystallize. You will have formed sodium thiosulfate.

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June 30 2011 5 30 /06 /June /2011 16:48

Warning: Sulfuric acid at these concentrations is slightly corrosive. Be careful when handling. Sulfur dioxide fumes are poisonous. Do not inhale them. Complete this experiment in a well-ventilated area.

 

You will need:

 

Sulfur

Hydrogen peroxide, 3% (a greater percentage gives more concentrated sulfuric acid)

Glass container

Aluminium foil

 

Shape the aluminium foil into a tray. Fill it with sulfur. Ignite the sulfur by heating the tray over a flame. Place it in the container and close the lid. The jar will fill with sulfur dioxide and sulfur fumes. Open the jar, remove the tray, and add some hydrogen peroxide. Shake the jar. Reignite the sulfur and repeat the process until all of the sulfur has burnt. Now you have a dilute sulfuric acid solution. If you want it more concentrated, boil it down.

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June 29 2011 4 29 /06 /June /2011 14:27

Warning: Sulfur compounds can be irritating. Do not inhale them. It is best to do this reaction in a well-ventilated place. Heat is released during this reaction. Keep flammables away.

 

You will need:

 

Sulfur (flowers of sulfur)

Pliers

Aluminium foil
Glass jar

Container

Ammonia solution

 

I know that sulfites can be obtained easily, but they can also be made from sulfur, which can also be obtained easily. First, shape the aluminium foil into a shallow pan. Fill it with sulfur powder. Holding it with pliers, heat it in a flame. It will melt, then ignite. As soon as you see flames, place it in the glass jar and secure the lid. It will burn until the oxygen is depleted and then go out, leaving yellow-white sulfur fumes floating around. Open the jar and carefully pull out the foil pan. Add some ammonia. Close the jar and slosh around. Pour the ammonia solution into another container. Reignite the sulfur and repeat the process until all of the sulfur has burned. You will be left with a more or less basic ammonium sulfite solution. Evaporate it. Excess ammonia will leave, and crystals of ammonium sulfite will form. If you use sodium hydroxide instead of ammonia, you will need to measure it or you will get sodium hydroxide mixed with your sodium sulfite. To test ammonium sulfate, place it in a basic solution. Ammonia gas should be given off. Place it in an acidic solution. Sulfur dioxide gas should be given off. If you did not use enough ammonia, you may form ammonium bisulfite.

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June 28 2011 3 28 /06 /June /2011 14:17

Warning: Bleach releases smelly fumes during this reaction.

 

You will need:

 

Sodium hypochlorite bleach

Manganese dioxide from charged battery

Test tube

 

Add the bleach to the test tube. Add the manganese dioxide. After a few minutes, the solution should become light pink as the tinest trace of permanganate leaches out. After 24 hours, the solution should be brilliant magenta. If the battery was not charged, the manganese(III) is oxidized to manganese(IV), using up the bleach. If it is slightly discharged, the manganese(III) and the manganese(VII) will comproportionate, producing manganese(IV) again and the solution will become colorless after about 2 days. Here is the solution. It is very very dilute, but still brilliantly colored. If you are looking for a measurable yield from this reaction, do not complete it.

 

Sodium-permanganate-solution.JPG

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June 28 2011 3 28 /06 /June /2011 12:56

Warning: Chromates are toxic and carcinogenic. Destroy them with ascorbic acid before washing down the drain. Hydrochloric acid is corrosive. Do not apply to skin. Nickel compounds are also carcinogenic.

 

You will need:

18/10 stainless steel object

Hydrochloric acid

Sodium hypochlorite

Sodium chloride

Drinking glass

12VDC power supply

2 Alligator clips

Ordinary iron nail or screw

 

Stainless steel is well known for its corrosion resistance. Despite this, we will split up 18/10 stainless steel into its three constituent elements. The 18/10 stands for percent of chromium/percent of nickel. 18/0 stainless steel is more boring as there are only two elements to separate.

 

Let's analyze this problem to see whether you can design the experiment. Stainless steel is resistant to oxidation by oxygen but not by chlorine. Therefore, sodium bicarbonate electrolyte is out of the question as only oxygen gas will be produced. What compounds can produce chlorine? Sodium chloride is one of the most common ones. The acid produced should dissolve the protective layer of the stainless steel, exposing it to further corrosion. As you should know, the anode oxidizes, and the cathode reduces. So the stainless steel should go to the anode. Since the electrode is not affected during reduction, any metal can be used. Because of iron's ready availability, it is used as a cathode. The oxides or hydroxides are produced during this corrosion.

 

Here are the properties of the elements:

 

Iron: 2 common oxidation states, none higher than 3, Fe2O3 insoluble in ammonia

Nickel: 1 common oxidation state, Ni(OH)2 soluble in excess ammonia

Chromium: 2 common oxidation states, water soluble +6 oxidation state

 

So, to separate chromium, we need to oxidize it so we can leach it out of the remaining oxides. Bleach is an oxidant, so that can be used. To separate nickel, we need to leach it out as the ammine complex. This involves much washing with ammonia. Then, only iron is left.

 

Electrolyze the stainless steel. You will see a dark green precipitate that will gradually turn brownish-green. Electrolyze as long as you like, replenishing water as needed. Filter and dry the precipitate. After placing it in a container, add excess sodium hypochlorite (bleach). Stir and let it sit for a day. Stir it again. A yellow solution  of sodium chromatewill form. Filter and keep both the filtrate and the precipitate. Evaporate the filtrate until only a small amount is left. While doing this, wash the precipitate with water. Place the precipitate back in a container, then add an excess of ammonia. After a while, a blue solution of nickel(II) ammine complex should form. Keep this covered for several days as the ammonia leaches out the nickel. Then filter and keep both the filtrate and the precipitate. Evaporate the filtrate until only a small amount is left. Then rewash the precipitate, following the steps above. Continue rewashing until the solutions become colorless. Then you have pure iron oxide. Now you have three solutions and one solid: chromate, nickel ammine, and iron oxide. Add ascorbic acid to the chromate. It will turn light green and a chromium(III) solution will form. Concentrate the solution by boiling. Electrolysis it using a smooth cathode and any anode may form chromium metal. Electrolyze the nickel ammine solution using a smooth cathode and any anode. Nickel powder should form at the cathode. You may use a strong magnet to pick up the nickel powder if it gets mixed with the solution. Dissolve the iron oxide in hydrochloric acid and add ascorbic acid. A green solution should form of an iron(II) complex. Electrolyze the solution using a smooth cathode and any anode. Pure iron powder should form at the cathode. Use a magnet to take it out of solution.

 

If this big experiment works, you have separated stainless steel into its constituents using nothing but household chemicals and basic supplies.

 

There are better and more reliable ways of obtaining the metals, but they do not use solely common household products.

 

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June 27 2011 2 27 /06 /June /2011 14:32

Needed:

Copper(II) chloride or what is in italics below

(Copper wire, at least 6 inches

Sodium bicarbonate, 10g

Any source of power that can produce 6-12V at >100 mA )

Zinc strip (approx. 8 mm x 50 mm)

Zinc metal (powder, oxide, strip)

Copper or brass strip (approx. 8 mm x 50 mm), or carbon rod

Hydrochloric acid (concentrated, not less than 10% or 3M, 3 mL)

Sodium chloride, ½g

Water

Gloves

Twist ties

Tissues

Sand paper

 

CAUTION: Do not ingest any of these chemicals. Zinc, copper, sodium bicarbonate, and sodium chloride are not any more than mildly toxic, but they shouldn’t be swallowed. Hydrochloric acid is toxic, and causes burns. Keep off skin, clothes, and metals. Copper hydroxide is an irritant and it is toxic. Copper chloride is a mild oxidizing agent. It should not be allowed to contact metals; it will corrode them. It can cause burns.  Zinc chloride is poisonous. Get medical attention if swallowed. Keep off skin.

 

There are two ways to start. Either dissolve copper(II) chloride in water until the solution is bright green or do the following process in italics. (Dissolve the sodium bicarbonate in approximately 200 mL of water. Cut the copper wire into two equal pieces. Connect the positive and negative ends of the power source to the copper wires. Put them in the solution. Blue copper hydroxide will be produced. Periodically scrape it off. Some of it will be converted to black copper oxide. Allow it to run until the copper wire producing the blue substance gets thin. Then reverse the polarity. Allow it to run again, scraping periodically, until that wire gets thin. Filter the solution through a tissue. Collect the blue-black substance. You may discard the wire and the solution. Dissolve a pea-sized amount in about 8 drops of hydrochloric acid. A green solution (copper chloride) will form. Keep adding copper hydroxide (the blue substance) until no more will dissolve. Put this aside for now.)

 

Add 8 drops of hydrochloric acid to another container and dissolve as much zinc in it as you can. There should be some residue of zinc. Put this (zinc chloride) aside for now.

 

Dissolve the sodium chloride in water, and put it aside.

 

Fold tissues four times, and cut rectangles approximately ¾ in. x 1 ¾ in. out. Make two of these. Cut another rectangle approximately 1 in. x 2 in. out.

On one of the smaller ones, put enough green copper chloride solution to dampen, not wet, the paper. Apply it to both sides sparingly.

 

On the bigger one, dampen it with the sodium chloride solution.

 

On the second smaller one, apply the zinc chloride solution sparingly. Put them on a tissue to protect the working surface. The chemicals should not leak out or you have made the tissues.  

 

Sand the metal electrodes so that they are shiny and conductive. Roll the copper chloride paper around the copper strip (or carbon rod). Roll the sodium chloride paper around the copper chloride paper. Roll the zinc chloride paper around the sodium chloride paper. Place the zinc on the zinc chloride paper, and twist-tie the whole assembly together. You may wrap the cell in plastic wrap to prevent evaporation of chemicals. Otherwise, they will evaporate quickly. This cell can produce from 0.7V to 0.9V for the copper strip (1.1V to 1.5V for the carbon rod) at 50 to 400 mA. It may or may not run a small motor (try to get one that can run on 3V or lower). Light bulbs typically run dimly. Wire three or four cells if using the copper electrode (or two cells if using the carbon anode) in series to run an LED. The reaction going on is: Copper chloride + zinc à zinc chloride + copper. You will notice copper crystals growing in the paper if you run the cell. The zinc will start corroding. The green copper chloride will start becoming colorless. Copper may be deposited on the copper strip (or carbon rod). When the green chemical disappears from under the brass or copper strip and is replaced by copper crystals, the battery is dead. Add a couple of drops of green copper chloride to the area that contacts the copper strip to recharge the cell. If the green color leaks to the zinc strip, the cell’s capacity will be limited. A diagram of the cell is below:

 

 

 There are several ways to increase the capacity.

  1. 1.       Decrease the leakage of the chemicals to each other by making the NaCl sheet thicker.
  2. 2.       Increase the concentration of the zinc chloride, copper chloride, and/or sodium chloride by using more concentrated HCl or dissolving more in the water, respectively.
  3. 3.       Use a tying method so that the zinc and copper strip contacts the papers better.
  4. 4.       Using a less active metal (Ag instead of Cu) and/or a more powerful oxidizing agent at the cathode(AgNO3, NaClO, H2O2, KMnO4, MnO2 instead of CuCl2) may increase the capacity.
  5. 5.       Using a more active metal Mg or Al instead of Zn) and/or a more powerful reducing agent at the anode may increase the capacity.
  6. 6.       Increase the size of the electrodes so that they contact the electrolyte more.
  7. 7.       Increase the size of the electrolyte papers.

If you use methods 7, 6, or 3, you might increase the current. If you use methods 2, 4, and 5, you might increase the voltage. If you use method 1, you might increase the voltage and the current.

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June 24 2011 6 24 /06 /June /2011 12:57

Warning: Tin is slightly toxic. Copper compounds are toxic. Do not eat or drink them.

 

You will need:

Copper sulfate

Copper(II) chloride

Tin or high-tin alloy

 

Tin sits right between germanium, which forms a stable +4 oxidation state, and lead, which forms a stable +2 oxidation state. Because of this, its +2 oxidation state is stable, but a reducing agent. Tin can be oxidized twice.

 

Dissolve some copper sulfate in water. It is a blue solution. Add tin. The solution becomes more colorless. If it has a little excess sulfuric acid, the tin(II) sulfate will stay in solution. If the solution is slightly basic, tin(II) oxysulfate will precipitate as a white powder. Here, the tin(II) does not behave as a reductant. This is a picture of the reaction Sn + CuSO4 => SnSO4 + Cu:

Sn-and-copper-sulfate.JPG

Here is a picture of the final precipitate, filtered and dried:

Tin(II) sulfate

Tin reacts differently with copper(II) chloride. The difference is that tin(II) sulfate cannot be oxidized, while tin(II) chloride can be.

 

Dissolve some copper(II) chloride in water. The solution should be green. Add tin. At first, the reaction Sn + CuCl2 => SnCl2 + Cu happens, and red copper precipitates. But then, the solution starts becoming cloudy as SnCl2 + 2 CuCl2 => SnCl4 + 2 CuCl happens. White CuCl precipitates. The SnCl4 hydrolyzes, creating white SnO2 precipitate. Then, the copper(I) chloride oxidizes: 4 CuCl + O2 => 2 CuCl2 + 2 CuO, creating black copper(II) oxide and soluble green copper(II) chloride, which starts the process over again. In my case, the CuCl2 ran out, leaving behind a solution of tin(II) chloride and a precipitate of tin(IV) oxide and copper(II) oxide. The latter's color predominates.

 

In this first picture, the tin has just begun reacting, and copper is visible on it.

Tin-and-copper-II--chloride-1.JPG

In the second picture, the tin(II) chloride has been reducing the copper(II) chloride.

Tin and copper(II) chloride 2In the third picture, the reaction is complete.

Tin-and-copper-II--chloride-3.JPGThe yellow color is the result of impurities in the tin. The white square is a stain cut out of the background. 

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June 23 2011 5 23 /06 /June /2011 12:03

Warning: Nitric acid is very corrosive and toxic. Do not contact it or let it contact any organic materials. The fumes produced are highly toxic. Do not inhale them.

 

You will need:

 

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

Dilute nitric acid

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.

 

In an earlier post (http://lanthanumkchemistry.over-blog.com/article-reactivity-series-demonstration-part-2-77468092.html), I showed the reactivity series by immersing metals in hydrochloric acid (well, lithium floated) and observing their reaction. The corrosion is not observable for all metals, so nitric acid, as an oxidizing acid, will increase the number of metals able to dissolve. Aqua regia, a mixture of nitric and hydrochloric acids, is even stronger, but it does not dissolve silver because a passivating layer is formed.

 

Place nitric acid in all of the test tubes. Add the metals. Here the fun starts. The more reactive metals such as magnesium, calcium, zinc, aluminium, and lithium will release plain old hydrogen gas. The less reactive metals, though, such as copper, iron, lead, tin and nickel will begin releasing nitric oxide gas, which oxidizes to the brown choking nitrogen dioxide. Even silver dissolves, although gold is untouched. If you are using your silver jewelry, take it out before it completely dissolves in the nitric acid.

 

The End

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June 22 2011 4 22 /06 /June /2011 13:14

Warning: Hydrochloric acid is corrosive. Do not contact. Some of the metal-acid reactions are quite violent. Wear sufficient protection and conduct experiment in a safe place.

 

You will need:

 

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

Hydrochloric acid, concentrated

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.

 

In an earlier post (http://lanthanumkchemistry.over-blog.com/article-reactivity-series-demonstration-part-1-77387937.html), I showed the reactivity series by immersing metals in water and observing their reaction. The corrosion cannot be observed in a reasonable time span, so using hydrochloric acid instead of water should help. Hydrochloric acid should dissolve the passivating oxide layer, allowing the true reactivity of the metals to be seen.

 

Add hydrochloric acid to all the test tubes. Drop all of the metals in, starting with gold. Gold, silver, and copper have no reaction. Lead has an extremely slow to nonexistent reaction. Tin slowly reacts, releasing hydrogen gas bubbles occasionally. Nickel reacts a bit more noticeably, producing a measurable amount of hydrogen. Iron reacts quite noticeably. Zinc reacts vigorously, producing much heat. Aluminium reacts even more vigorously. Magnesium reacts even more vigorously. Calcium reacts even more vigorously. Lithium reacts even more vigorously. This trend is only observed when all acids are equal concentration and all elements are of equal shape and size.

 

Keep in touch for Part 3!

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