Chemical Engineering for Home Experimenters (Nov, 1939)

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Chemical Engineering for Home Experimenters

By RAYMOND B. WAILES

HOW would you like to transform your home laboratory into a miniature factory for obtaining useful products from minerals, with yourself as its amateur chemical engineer? You can roast ores to obtain important metals, convert plain-looking stones into sparkling white crystals, operate a rotating kiln, and magically separate wanted minerals from unwanted ones by a flotation process.

This may sound as though it would make rather a severe dent in your pocketbook. Nothing could be farther from the truth. Your rotating kilns or reaction kettles will be tin cans, supported horizontally and turned by your own hand. For a roasting furnace, frequently employed by metallurgical engineers, you will use a porcelain crucible or the bowl of a clay bubble pipe. Heating kettles may be squat tin cans, porcelain evaporating dishes, or beakers you can find on your laboratory shelf.

As for your raw materials, they will be the very same ones with which professional engineers have to deal. Perhaps you may live in a good locality for collecting some of the minerals that will be mentioned here. No matter if you don’t, for many dealers in chemicals can supply them. For variety, you will find an excellent source of supply in the “mineralogical sets” of assorted specimens for beginners, now on the market. For bulk consumption, the “pound material” for mineral collectors, sold by some firms at a few cents a pound, will prove ideal.

Would you like to prepare a widely used metal from a well-known ore ? You can easily extract lead from galena, a natural mineral which has been used in crystal radio receivers. Chemically, galena is a compound of lead and sulphur, called lead sulphide.

Professional metallurgists grind this ore and roast it at low heat, in a chamber resembling a blast furnace. During this process, the oxygen of the air converts part of the galena into lead oxide and lead sulphate. Then the temperature of the roasting chamber is raised. By reaction with the initial products, the unconverted galena is transformed into metallic lead, which is tapped off at the bottom of the furnace.

You can dispense with the preliminary step by mixing powdered galena (lead sulphide) with twice its weight of litharge (lead oxide). Place the mixture in a porcelain crucible or the bowl of a clay bubble pipe, and keep it at bright-red heat for about twenty minutes. The molten lead may then be poured off from the vessel. Hammering the metal will convince you that it is really lead.

In the purity of the material that you have to work with, you have an advantage over the professional chemist or metallurgist. His ore may be quartz rock in which the pure galena appears only as veins or streaks. However, in a clever process that he calls flotation, he has found a way to separate the valuable galena from the worthless quartz. You can try out his scheme for yourself.

Pulverize very finely, and mix thoroughly, equal portions of galena and common sand. Place a spoonful of this mixture in a pint bottle or a mayonnaise jar. Add a drop of pine oil, or cresol, and half fill the vessel with water. Shake vigorously, and then let the contents settle. The oil will rise to the surface of the water, carrying a large amount of the galena with it. Meanwhile the sand, which represents the worthless quartz in actual commercial practice, is left behind on the bottom.

An important ore of antimony metal is a mineral called stibnite, which usually is included in small mineralogy sets. It consists of a sulphide of antimony. In the commercial process, which you can easily duplicate, the antimony is extracted by heating stibnite with half its weight of iron filings.

Place this mixture in the bowl of a clay bubble pipe, or in a porcelain crucible. Also add a bit of sodium sulphate, and of powdered coal. These will act as a flux and assist the reaction between the iron and the stibnite. A bright-red heat is required for this reaction, which takes about forty-five minutes for its completion. The metallic antimony, which will be found at the bottom of the pipe bowl or crucible, is best removed by breaking open the container.

Lime, a material of great chemical and industrial importance, is produced rather simply. Limestone (calcium carbonate) is heated to redness in upright kilns, releasing carbon dioxide into the air and leaving the lime (calcium oxide) behind. To duplicate the process in your home laboratory, heat ordinary limestone or chips of marble (crystallized limestone) to redness for several hours. This may be done in a porcelain crucible, using a Bunsen burner; or, if you want to produce a larger amount of lime, place a few ounces of limestone or marble in a sand crucible and heat it for several hours in a coal fire.

Converting gypsum, a common mineral, into plaster of Paris will afford an enjoyable hour or two of chemical experimenting. Held in the hand, gypsum is just a common-looking, light-colored rock. Chemically it consists of calcium sulphate molecules, to each of which is attached a pair of water molecules. When gypsum is heated, some of this “water of crystallization” is driven off. The calcium sulphate that is left behind (with more or less water remaining, depending upon the dehydrating temperature) is the product known as plaster of Paris. When it is mixed with water, its molecules reunite with some of the water molecules and a hard mass is formed, closely approximating the composition of the original gypsum.

Rotating kilns, kettles, or muffle furnaces are used for making plaster of Paris commercially. One simple way you can prepare a small quantity is to heat powdered gypsum with a very low flame, from below, in an evaporating dish or a squat tin can. To control the temperature, stir the powder during the process with a chemical thermometer. When the material first becomes warm, you will observe “dew” or moisture being driven from the powder and condensing upon the thermometer. Keep the temperature at about 145 degrees C. (293 degrees F.) for half an hour, stirring every five minutes or so. When cold, the powder will set to a hard mass on being mixed with water. If the temperature used in preparing it is too high, the product will be “hard burned” and will take longer to set.

A MODEL rotating kiln, like those used in industry, may also be used to make plaster of Paris from gypsum. Pierce the lid and bottom of a baking-powder can and insert a stiff metal rod or wire, bent into a crank at one end, as illustrated. The rod is passed through the centers of the lid and bottom, and is attached to the metal with solder to keep it from slipping when the can is rotated with a heavy load.

Also punch a hole in the lid of the can, very near the rim, through which you can insert a glass tube. This tube should be about six inches long, almost half an inch in diameter, and sealed at one end. Half fill the tube with powdered aspirin tablets. With the sealed end inserted in the can, this apparently senseless accessory will serve as a temperature-indicating device!

Mount the “kiln” on improvised bearings, which may be drilled corks supported by laboratory stands, after filling it about a quarter full of powdered gypsum. Place a Bunsen burner beneath, with a flame-spreader attachment to distribute the heat. Start rotating the can, and adjust the burner flame until the aspirin in the glass tube remains molten, except for a small amount that solidifies in the exposed end of the tube. Since aspirin melts at approximately 135 degrees C.. the gypsum within the can will then be just about at the proper temperature of 145 degrees C. If the aspirin starts to solidify toward the inside of the can, increase the heat; if it tends to flow from the open end of the glass tube, lower the flame.

Rotate the kiln by the crank handle, while you continue heating, for half an hour or so. You need not revolve the can continuously; just a turn now and then is sufficient to expose fresh portions of the mineral to the hot surface of the can, until the desired plaster-of-Paris product is obtained.

From a rocky white mineral called witherite, you can prepare glistening white crystals that will serve you in many of your home-laboratory experiments. This mineral, not at all uncommon, consists largely of barium carbonate. By dissolving it in hydrochloric acid, you can obtain barium chloride, for which you will find frequent use.

Place the powdered mineral in a beaker and add its own volume of diluted hydrochloric acid (made by mixing one part of strong hydrochloric acid, by volume, with two or three parts of water). Carbon dioxide gas escapes and the mineral partially dissolves. It will not dissolve entirely, since it undoubtedly will contain considerable quartz rock.

Warm the solution gently and add several grams (or a teaspoonful) more of the powdered mineral. If the powder effervesces, not all of the acid has yet been used up in converting the carbonate to the chloride. Make fresh additions of the powdered mineral until they are no longer acted upon, showing that the solution may be considered free of acid.

If you were to stop at this point, filtering the liquid and letting it crystallize, the crystals of barium chloride would be colored a rusty brown by iron impurities. Therefore further treatment of the beakerful of liquid is first required.

Passing chlorine gas through the liquid, before obtaining the crystals from it, will dispose of the iron. This should be done outdoors. Fit a small flask with a one-hole cork, carrying a piece of glass tubing that dips into the liquid in the beaker. Place in the flask a small amount of potassium permanganate crystals, or a spoonful of bleaching powder or of liquid laundry bleach (sodium hypochlorite). Potassium permanganate will work best. Now add to the flask about half a fluid ounce of strong hydrochloric acid. Chlorine gas will be generated, and should be allowed to bubble for about half a minute through the beakerful of barium chloride solution and quartz residue. The result is that any iron impurities are precipitated at this point.

Now place the entire contents of the beaker, together with several grams more of powdered witherite, in a bottle and keep it stoppered for twenty-four hours or so. At the end of this time, filter out the worthless rock and the precipitated impurities, catching the filtrate in a porcelain evaporating dish. Wash the filter paper with fresh water, adding this filtrate to the first. When the liquid is now evaporated, gleaming white crystals of barium chloride will be obtained.

This final operation may be performed by heating the contents of the evaporating dish with a small flame, or with a homemade electric lamp-bulb heater previously described in this series (P.S.M., Aug., ’39, p. 182). If you use the lamp-bulb heater, you can place the dish upon it and let the liquid evaporate completely, leaving the barium chloride crystals behind. Otherwise, separate a quantity of the crystals first obtained; then discontinue heating, and let the remainder of the liquid evaporate at room temperature. The crystals may be bottled, and preserved for use in your future experiments. For using the lamp-bulb heater with large evaporating dishes, you can make a special attachment as shown on page 208.

A small collection of minerals, such as the mineral “sets” offer, will provide many other interesting chemical experiments. Barium, strontium, and calcium minerals exhibit typical “flame colors” when moistened with a drop of hydrochloric acid and held in a blue flame. Barium compounds tinge the flame green; strontium compounds, red; and calcium compounds, a brick red of different hue. Powdered cryolite mineral, shaken in a test tube of water and allowed to settle, becomes almost invisible because its refractive index, or light-bending power, so nearly equals that of water. Shake a bit of powdered orthoclase mineral (common feldspar), with a drop of phenolphthalein solution, in a little water, and no visible color change will occur. Although the mineral is alkaline and turns the phenolphthalein red, the color is removed by adsorption, adhering to the mineral. To prove that the clear liquid is alkaline, none the less, let the contents settle and add another drop of phenolphthalein. This time the result will be a visible red color.

3 comments
  1. DocScience says: January 22, 201311:00 am

    Today, the schools would be expelling and the FBI arresting these dangerous young terrorists in training.

  2. Stephen says: January 23, 20135:45 am

    How did the no-battery light on page 5 work? It was produced by General Electric, which suggests it wasn’t a scam or fantasy.

  3. GeorgeT says: January 23, 20138:16 am

    Stephen, it looks like a generator light — work the switch in and out with your, soon to be sore, thumb, and it spins a small generator inside. Some work that way, and some have a crank.

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