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INDUSTRY GIVES A LABORATORY TO AMERICA’S YOUNG SCIENTISTS

YOUTHFUL, IMAGINATION, an inexhaustible national resource, is being developed along scientific lines by the American Institute of the City of New-York. This organization, chartered in 1828 and devoted throughout its existence to the promulgation of science and the encouragement of American industry, established its junior branch in 1928 and recently has intensified its efforts in this direction through the American Institute Laboratory at 310 Fifth Avenue, New York.

Its aim is to direct and utilize the imaginative faculties of youth which, since the founding of the institute, have been turning more and more toward science and mechanics. Under its wing are more than 730 juvenile science clubs, scattered throughout the United States, its possessions, and foreign countries. Some meet in high schools, some in settlement houses, and some are spontaneous youthful organizations with cellar or attic laboratories and club rooms. In the aggregate there are more than 30,000 youthful club members.

They experiment with model airplanes, bacteria, telescopes, radio, tropical fish, light, sound, animal-breeding, and in numerous other fields. Their ambition is limited only by their own knowledge and the cost of equipment, and it was to obviate the latter difficulty to some degree that the American Institute Laboratory has been established with the cooperation of the International Business Machines Corporation, which gave the use of two floors of a New York City office building, and of the Westinghouse Electric & Manufacturing Company, which supplied the equipment.

There is room for thirty to work at a time, and the laboratory is used by three shifts daily. One uses it from 9 a.m. to noon; one from 2 p.m. to 6 p.m., and one from 6 p.m. to 9 p.m. It is open six days a week. Ordinarily, a student has the use of it for two periods a week.

Members of the junior activity clubs of the American Institute are eligible to use the laboratory. They are boys and girls from twelve to eighteen years old. Membership in their club, which pays dues of $2 a year to the institute, is the only requirement necessary except the ability of the student and the suitability of his project.

The student desirous of getting working space in the laboratory writes to the institute describing his project, its purpose, the equipment which will be necessary, and the time it will take. Allotments of space are made as it becomes available. The laboratory has projection microscopes, aquaria, a darkroom, drafting and drawing boards and equipment for their use, a wood-working shop with power sanders, lathe, drill presses, and other machinery, and departments fitted for special projects in radio, aviation, and the physics of sound.

There are tables fitted for glass-blowing, and other equipment with which students may manufacture some of the devices which may be necessary for the work they plan to do. A tool kit is issued to each student when he enters the laboratory, and at the end of his work period he replaces it, in condition to be used again immediately should a student in the next shift be engaged in the same kind of work.

There is a reference library, and students have access also to the library of the American Institute at its headquarters at 60 East Forty-second Street. The laboratory also has an advisory board of scientists in various fields who will answer students’ questions and give technical information.

Students are contributing constantly to the equipment of the laboratory. One is engaged in making a blueprinting machine, and another is working on a mimeographing outfit. Another is custodian of one of the stockrooms, working on his own project in his spare time.

Some of the budding scientists have domestic difficulties which interfere to some extent with their careers. One, whose mother is dead, has to leave a little early every day to get home in time to cook supper. So far as is known, supper never has been late, but an experiment he is conducting in hydroponics, to determine how onions thrive under varying conditions, suffered once for lack of sufficient attention.

Another fled to the laboratory as a sanctuary with his white mice. He had been breeding the animals to study the Mendelian characteristics of succeeding generations and about Christmas time last year, when he had reached the twelfth generation in his tests, his mother rebelled. Enough was enough, she said, and twelve generations of mice were altogether too many mice. She was exceedingly firm about it, too, and the young scientist had to lead an immediate exodus of his highly bred subjects. He found a temporary home for them with a neighbor until he gained admission to the laboratory. There the mice are housed in a cage built for just such experiments by one of the junior activity clubs of the institute in Maiden, Mass.

The clubs all over the country are engaged in just such work as is going on at the laboratory, though generally without the equipment that is available there. The institute plans to establish other communal laboratories in centers where they may be used by several clubs. As far as possible, projects are undertaken at the New York laboratory with a view to helping clubs at a distance. Cultures, for instance, are being grown there in large quantities so that they may be sent to outlying clubs.

The American Institute has a Science Fair every year in the Education Hall at the American Museum of Natural History in New York, at which members of the affiliated clubs exhibit their handiwork. The institute sends its own technicians to aid in setting up the more elaborate exhibits. Leading scientists and educators are among the judges at the fair. Last year on the opening day the attendance was 7,222.

Airplane models naturally are among the more popular projects of club members, and Richard Walton, a youthful aeronautical engineer who won a prize at the Pittsburgh Science Fair, designed and manufactured a wind tunnel with which to test the model craft. Alan Goodman designed a seaplane bomber which carried torpedoes in its pontoons, reducing air resistance. It carried machine guns in the wings and a cannon on each side of the propeller.

Wallace Cloud, fourteen years old, a student at the Grover Cleveland High School, is working at the institute’s Fifth Avenue laboratory on the distillation of household refuse. His experiments might put a high value on the garbage pail, as they indicate the possibility of extracting chemicals valuable both in medicine and for explosives.

Judges at the Science Fair have at their disposal $3,000 in prizes to be awarded for conspicuously good work. The Veterans’ Wireless Operators’ Association offers the Marconi Memorial Award Scholarship to institute members.

With the establishment of its laboratory in New York for boys and girls with a scientific bent, the American Institute feels that it has taken a long step forward in the program it undertook at the completion of its first hundred years for the training and development of the imagination of America’s youth. When other such laboratories have been established, the youth organization of the institute will become a close-knit national training school.

Amateur Chemist’s Robot
Hyman Cordon, chemical student, of Boston, with a “man” he built out of rubber, glass, and other scraps. It eats food and digests it in human fashion, having heart, intestines, lungs, bladder, etc. It was exhibited at a recent “science fair.” (Int. News)

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England Now Has Gasoline Made from Coal

British motorists may now enjoy the novelty of buying gasoline made from coal, which has just been placed on public sale. The event marks the beginning of a great chemical industry by which England hopes to put 65,000 men to work and to end her dependence upon imported petroleum. A monster plant now rising at Billingham-on-Tees will transform 1,000 tons of coal daily into the synthetic fuel, using a process already in successful operation in a smaller experimental plant at the same site. In this process, known as hydrogenation, powdered coal is mixed with heavy oil and the resulting paste is fed, with hydrogen gas, to a converter. The mixture undergoes a chemical transformation under tremendous heat and pressure, yielding a mixture of hydrocarbons from which pure gasoline is recovered by distillation. Another of the products is Diesel oil, which may also be changed into gasoline by an additional conversion treatment with hydrogen. Both the hydrogen and heavy oil used in the process are obtained in the course of producing the gasoline, leaving coal as the chief raw material required. Results of production indicate that approximately a gallon of gasoline may be obtained from twenty-four pounds of coal, and the large-scale plant under construction should show an output of 80,000 gallons of gasoline a day.

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Spectacular Fireworks

By STANLEY STEWART

IN making fireworks, if the experimenter will always remember that he is dealing with explosives that may pop off at any moment, and therefore exercises constant caution, the various spectacular night displays outlined in the accompanying article are not any more dangerous than playing with matches. At all times, care must be exercised in grinding the ingredients. Always use a clean mortar; always powder each chemical separately; when mixing, dump the required portions on a sheet of dry paper and use a wooden spatula, or gently rock the contents of the paper back and forth. Although the author is only fifteen years old he has been making fireworks for years and has not yet had one of them go off accidentally. The formulas contained in this article have all been tried and tested, and will be found to work perfectly.

Aerial Maroon To make a mortar, fill a 5″ by 1″ cardboard tube at least 1/8″ thick and to the depth of 1/2″ with plaster of Paris. When dry, punch a hole in the tube large enough to accommodate a salute fuse, just above the plaster.

Two kinds of propellant may be used—either flashlight powder or rifle powder. (See note at end.) To make a shell, use a tube 2″ high with a diameter slightly less than that of the mortar. Seal to the depth of 1/2″ with plaster of Paris, leaving through the plaster a hole large enough to accommodate the type of fuse used in roman candles. When dry, place 1/2″ of flashlight powder in the shell.

To make the flashlight powder: 2 parts potassium perchlorate (NOT potassium chlorate) and 1 part red phosphorus. Fill rest of the shell with plaster of Paris; when this is dry, place it, fuse-end down, on a spoonful of flashlight powder in the mortar. Pack a wadding of paper on the top of the shell.

If you do not wish to prop up the mortar with bricks, paste a cardboard disc on the bottom of the mortar.

American Beauty Bomb Use an 8″ by 1-1/2″ mortar and 5 spoonfuls of flashlight powder propellant.

To make the shell, use a cardboard tube 3-1/2″ high, the diameter slightly less than that of the mortar. Paste a cardboard disc on one end of the shell. Fill the shell with this compound: one part sulphur, 2 parts powdered charcoal, 3 parts strontium chlorate, mixed with shellac to form a paste.

Place the shell in the mortar, sealed end up. Add paper wadding.

Aurora Rocket Fasten with wire an 8″ x 1″ cardboard tube to a wooden stick 48″ long, and fill the tube with plaster of Paris to the depth of 1″; leaving in the plaster of Paris a hole 3/4″ in diameter. Run a fuse through this hole; and pack paper around it to secure the fuse. On top of the plaster of Paris, place the following compound to the depth of 4″: 2 parts of potassium chlorate, 1 part sulphur, 3 parts powdered charcoal, 2 parts powdered emery. Then add plaster of Paris to the depth of 1″ leaving through the middle a fuse hole in which to run a black-powder fuse. Add 1/2″ flashlight powder, then 6 “Star-balls.”

To make the star-balls, mix with shellac, to form small balls, 2 parts potassium chlorate, 1 part sulphur, 1-1/2 parts powdered moth-balls, 1 part powdered iron.

Add 1″ plaster of Paris.

Battle In the Clouds Put flashlight powder propellant to the depth of 1″ in any size mortar. Place on top two rolled-up strings of Chinese firecrackers, and add paper wadding.

Cannonade Shell Into a mortar 18″ x 2″, put 2″ of rifle powder. Use a cardboard tube shell, 8″ high and slightly less in diameter than the mortar. Run a 9″ black powder fuse through the shell, leaving 1″ outside. Put a 1″ plug of plaster of Paris in the bottom of the shell. Add 1″ of flashlight powder; then a 1″ plug of plaster of Paris; another inch of flashlight powder, and so on to the end of the shell. Place the shell, fuse-end down, in the mortar. Fill to the top with paper wadding.

Combination Chainlight Shell Make three cardboard tubes, 2″ long and Y2″ wide. Put a cork in the end of each. Tie them on a string so that they will be twelve inches apart. Fill the first tube with 2 parts of strontium chlorate, 1 part sulphur, 2 parts powdered charcoal. Fill the second tube with powdered charcoal (2 parts), 2 parts barium chlorate, 1 part powdered sulphur. Fill the third tube with 4 parts potassium chlorate, 2 parts sulphur, 2 parts powdered copper, 1-1/2″ parts copper sulphide, 3 parts black copper oxide. Mix each filler with shellac and press into its tube. When they harden, group the tubes in your hand, with the string end up. Place them in a mortar just large enough to accommodate them; pack parachute and string on top of the tubes, and add paper wadding. Use black-powder propellant.

Dragon Rocket Use a mortar with diameter slightly larger than that of a large spool, and a flashlight powder propellant. Paste cardboard over one end of a large spool. Mix 2 parts potassium chlorate, 1 part sulphur, 1 part powdered emery, 2 parts powdered iron with shellac, and press into the spool hole. When dry, place in mortar over 2 spoonfuls of flashlight powder, with the cardboard end up. Add 2″ of paper wadding.

Emerald Bomb Make this like the American Beauty Bomb, but substitute barium chlorate for strontium chlorate.

Flitter Bomb Make like the American Beauty Bomb, except for powdered iron instead of powdered charcoal, and potassium chlorate for strontium chlorate.

Fiery Tail Salutes Make like the American Beauty Bomb, but substitute potassium chlorate for strontium chlorate. Before putting the compound in the shell, place three firecrackers in the bottom of the shell.

Golden-Star Mine Candle Effect As golden-star roman candles are not sold, you will have to make such candles. Use two old roman candle tubes 12″ long, and be sure the interiors of the tubes are clean; fill each tube to the depth of 1/2″ with plaster of Paris. When dry, punch a hole in the side of each tube, just above the plaster of Paris, and insert a salute fuse in each hole. Put 1/2″ of rifle powder in the bottom of each tube, and one golden star on top; pack 1/2″ of filler powder on top of the star. On top of this, pack 1/2″ of rifle powder, add one golden star; then put in another 1/2″ of filler powder, and continue in this manner until you have reached the top of the tubes.

Make the stars in cylinder shape, 1/2″ long, the diameters should be slightly less than the inside of the tubes. Mix the following compound with shellac to form a very thick paste: 2 parts sodium chloride, 1 part napthalene, 4 parts powdered charcoal, 4 parts potassium chlorate.

Filler Powder Mix lightly the following on a sheet of paper: 3 parts potassium chlorate, 2 parts iron (reduced by hydrogen), 1 part sulphur.

(Before filling the tubes burn small samples of the filler powder on an iron plate, to make sure that no residue is left after burning. If there is, vary the quantities of the chemicals in the filler powder until no residue is left.) For mortar, use a cardboard tube 12″ high, 2″ wide. Cut a disc of wood 1/2″ thick, to fit snugly in the bottom of the mortar. Drive 8 tacks through the cardboard tube into the wooden disc, and punch 2 holes opposite each other in the mortar just above the disc. Use strong glue to fasten the roman candles on the outside of the mortar, so that the two fuses will protrude through the two holes in the mortar. Put 2″ of black powder in the mortar, next, 15 golden stars; pack 6″ of paper wadding on this. Cut a piece of film 4-1/2″ long and 1/2″ wide. Glue, with shellac, each end of the film strip over the top of each roman candle.

Bury this mortar half-way in the ground. Light the strip of film exactly in the middle. After the candles have quit shooting, do not approach until the mortar has fired. Do not shoot under a tree or overhead obstruction.

Jewel Mine Make two roman candles as described for the Golden Star Mine, only using the following 12 stars in each candle. To make all the stars, mix the compound with gum arabic in water to form a paste.

First Star: 3 parts potassium chlorate, 1 part sulphur, 3-1/2 parts powdered charcoal.

Second Star: 3 parts potassium chlorate, 1 part sulphur, 3 parts powdered iron.

Third Star: 3 parts potassium chlorate, 1 part sulphur, 2 parts powdered antimony, 2 parts powdered arsenic.

Four Star: 3 parts potassium chlorate, 1 part sulphur, 2 parts sodium chloride, 1 part sodium nitrate.

Fifth Star: 3 parts potassium chlorate, 1 part sulphur, 2 parts powdered indigo.

Sixth Star: 3 parts strontium chlorate, 1 part Sulphur, 2 parts powdered charcoal.

Seventh Star: 3 parts potassium chlorate, 1 part sulphur, 1 part barium nitrate, 1 part barium hydroxide.

Eighth Star: 3 parts barium chlorate, 1 part sulphur, 2 parts powdered charcoal.

Ninth Star: 3 parts potassium chlorate, 1 part sulphur, 2 parts black copper oxide.

Tenth Star: 3 parts potassium chlorate, 1 part sulphur, 2 parts powdered copper.

Eleventh Star: 3 parts potassium chlorate, 1 part sulphur, 1-1/2 parts powdered copper, 1-1/2 parts copper sulphide.

Twelfth Star: 3 parts potassium chlorate, 1 part sulphur, 2 parts lime.

Substitute one of each of these stars in the mortar, in place of the Golden Stars.

Iris Bomb Obtain a mortar 12″ high and 1-1/2″ wide. Put 1-1/2″ of rifle powder in the bottom, and on this 12 stars, 6 of Composition A and 6 of Composition B. Composition A: Mix with shellac to form a paste, 3 parts potassium chlorate, 1 part sulphur, 2 parts powdered copper. Composition B: Mix with shellac to form a paste, 3 parts potassium chlorate, 1 part sulphur, 2 parts powdered aluminum. On top of the stars, pack 5″ of paper wadding.

Magnesium Bomb Use a mortar 12″ high and 2″ wide. Fill to the depth of 11/2″ with flashlight powder. On this place the following shell: Fill a cardboard tube size 1″ x 3″ with the following: 1 part sulphur, 3 parts potassium chlorate, 2 parts magnesium (size of confetti), 2 parts gum arabic. Moisten with water and press into the tube. Paint, with ordinary house paint, over one end of the tube; and place the painted end up in the mortar. Pack in 3″ of paper wadding.

National Color Bomb Drill three holes, 1″ in depth and 1″ in diameter, in a “two by four” 12″ long. One hole should be 2″ from one end; another 2″ from the other end, and the third hole in the middle. Fill to the depth of 1″, with plaster of Paris, three cardboard tubes 12″ high and 1″ wide. When dry place tubes, plaster of Paris-end down, in the three holes. Punch two holes opposite each other, just above the plaster of Paris in each tube. Run a continuous black powder fuse through all holes except the last, and plug this with a match stem. Put 1/2″ of rifle powder in the bottom of each tube. In tube 1, place a red star, in tube 2 a white star, and in the other tube place a blue star. On top of each star place 1″ of paper wadding. (Directions for making stars under Jewel Mine.) Parachute Bomb (Apple Green) Make a mortar 12″ high with a diameter slightly larger than a 1-ounce pill can, remove lid and tie a small green parachute to the bottom of the can. Fill the can with this compound: 4 parts potassium chlorate, 1 part sulphur, 3 parts barium hydroxide, 2 parts barium nitrate, 2 parts gum arabic. Moisten with water and press into can. When dry, place can, open end down in mortar. Pack parachute on top of can. Pack 3″ of paper wadding on top of this. Be sure to fire this in open country.

Sapphire Bomb Like Magnesium Bomb, but substituting lycopodium for chips of magnesium.

Serpent Mine Use 24 empty .22 caliber cartridges, long-rifle. Moisten black powder and gum Arabic and press into the cartridges. Substitute the cartridges for the stars in the Golden Star Mine.

Serpent Shell Like Aerial Maroon, but substitute the cartridges, as in Serpent Mine, for half the powder in the Aerial Maroon shell.

Shooting Star Rocket Like Aurora Rocket, but substitute varied colored stars as described for Jewel Mine in place of the fire balls of Aurora.

Geysers Make a sharp-pointed wooden cone, with a base 1-1/2″ in diameter. Obtain a very thick-walled cardboard tube, 2″ in diameter and 1 foot long, and cut 4 cardboard discs that will fit over the end of the tube. Glue 2 of these discs together, and glue this on one end of the tube; then glue the wooden cone on this. Stick the pointed wooden cone in the ground. Put a firecracker in the bottom of the tube, and cover with 1″ of rifle powder. On top of this put 1-1/2″ of the following compound: 2 parts potassium chlorate,1-1/2 parts lithium nitrate, 1 part sulphur, 1-1/2 parts powdered charcoal. Then fill the rest of the geyser with the following compound: 3 parts potassium chlorate, 2 parts iron (reduced by hydrogen), 1 part sulphur. Glue together the other two cardboard discs; punch a hole 1/2″ in diameter in the center and glue this on the end of the geyser. Be sure the powder comes right up to the hole. Cut a piece of friction tape, 1″ square, and stick this over the hole. To fire, pull off the tape, light the powder at long range, and step back quickly.

Hanging Chain Rocket Make a rocket as described in the Aurora Rocket; but leave out the “Star-balls” and substitute parachute and lights as described for the Combination Chain Light Shell. In order to fire all the rockets, bury a narrow bottle up to its neck in the ground and place the end of the rocket stick in this. Never shoot the rockets described in this article at an angle.

Parachute Bomb (Silver Flare) Make like above, but with this compound: 3 parts potassium chlorate, 2 parts powdered magnesium, 1 part sulphur.

Parachute Rocket (Red Star) Same as Aurora Rocket but substitute for the golden star-balls the following: 3 parts strontium chlorate, 1 part sulphur, 2 parts powdered charcoal.

Parachute Rocket Special Effect Make exactly like Aurora Rocket except for star-balls. In place of them, use a cardboard tube 3″ long and 3/4″ wide. Fill to the depth of 1/2″ with plaster of Paris. Punch a hole 1/4″ in diameter just above the plaster of Paris, and fill the hole with a mixture of black powder and shellac. When this is dry, fill the tube to the depth of 2″ with the following compound: 3 parts powdered iron, 2 parts sulphur, 5 parts potassium chlorate. Fill the rest of the tube with plaster of Paris. Tie a string around the middle of the tube. Attach string to a small parachute. Place tube, hole-end down, in rocket, pack parachute on top of tube, and put a cork in the end of the rocket.

Ruby Bomb Like as Magnesium Bomb but with strontium chlorate instead of potassium chlorate.

Reporting Star Mine Use a heavy cardboard tube, 1″ wide and 48″ long. Fill to the depth of 1-1/2″ with plaster of Paris. When dry, put 1″ of rifle powder in the tube. Place one reporting star (see below), with open end down on this. Add 1-1/2″ of filler powder. Place 1″ of rifle powder on this and one reporting star and 1-1/2″” of filler powder. Continue in this manner to the top of the tube.

How To Make Reporting Stars Use a cardboard tube 1/2″ in diameter and 2″ long. Cover a baby Chinese firecracker with glue, and place fuse-end down in tube. Fill the space around the firecracker with plaster of Paris. Fill the space above the firecracker with a mixture of black powder and gum arabic, slightly moistened with water.

(NOTE—Since the above was put in type, Mr. Stewart has added other information, as follows: “The ‘rifle’ powder I use is taken from 16-gauge shotgun shells; it is smokeless and burns rapidly. The pasteboard tubes are about 0.2″ thick, but exactness is not important. They come from packages, etc.; the best are those upon which oilcloth comes rolled. Reduce the quantity of potassium chlorate in the filler powder if it leaves a gummy residue; too much iron may do so also. I do not advise shooting up more than a 1-1/2″ snuff can full of powder; a larger can may be still burning when it lands. I find that gum arabic and water is better than shellac for stars, etc. A good quick burning fuse is made by dissolving 2 oz. of potassium chlorate in 150 cc. of boiling water, and soaking 1/2″ strips of blotting paper in this; then dry. For the apple-green parachute bomb, use ‘rifle’ powder or this; 3 parts potassium chlorate, 2 parts charcoal, 1 part sulphur”)

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FUN with the HALOGENS

HOME EXPERIMENTS WITH A FAMOUS CHEMICAL FAMILY

By RAYMOND B. WAILES

WHENEVER the members of the halogen family put on an act, you can be sure there will be something doing in the way of entertainment. The nimblest of the family quartet undoubtedly is chlorine. You have seen this actor in several roles before—bleaching dyes, and attacking metals with accompanying pyrotechnics, for example—if you have followed this series of articles. Iodine has made a personal appearance before you as a chemical detective, revealing latent fingerprints on paper. Another member of the family, fluorine, has shown you its remarkable power of etching glass when teamed with hydrogen. The remaining member of the quartet, bromine, is an irritating, rascally sort of character if encountered alone. However, when handcuffed to hydrogen, its behavior is so satisfactory that you should let it enter your home chemical laboratory and allow it to perform for you.

Uniting bromine with hydrogen yields hydrogen bromide, or hydrobromic acid—just as chlorine and hydrogen form hydrogen chloride, or hydrochloric acid. Hydrobromic acid reacts with substances to form bromides, as does its better-known relative, hydrochloric acid, to form chlorides.

Unlike most acids or acid-forming gases, however, hydrogen bromide cannot be prepared for practical purposes by heating corresponding salts with strong sulfuric acid. It is formed in the reaction, to be sure, but it is rapidly decomposed by the oxidizing action of the sulphuric acid. This difficulty is overcome by heating a bromide with strong phosphoric acid, which does not decompose the product.

To make hydrogen bromide, place a half ounce of potassium bromide or sodium bromide in an Erlenmeyer flask or a retort, with a capacity of sixty to 200 cubic centimeters (two to seven fluid ounces). Cover the chemical to a depth of an eighth of an inch or so with strong phosphoric acid, which you can buy at any drug store under the name of eighty-five-percent, or sirupy, phosphoric acid.

Arrange tubing to lead from the flask or retort to the bottom of an empty side-necked test tube, which will serve as a catch bottle to condense water vapor distilled from the phosphoric acid. At the test tube’s side neck, attach more tubing that will conduct the hydrogen bromide vapor to the bottom of a test tube for collecting the gas. This test tube may be left open, and the gas, being heavier than air, will displace it and fill the tube.

Apply heat to the flask or retort, with a Bunsen burner, and hydrogen bromide gas will be generated. It will pass through the system into the last test tube. Meanwhile a teaspoonful or so of useless condensate will collect in the side-necked test tube.

Collect a tubeful of hydrogen bromide gas and then pour it, as if it were water, into the air. A white cloud forms as the gas combines with moisture in the air.

Hold an object moistened with ammonium hydroxide (household ammonia may be used) in the stream of hydrogen bromide gas from your apparatus. Dense white clouds of ammonium bromide will be formed. This resembles the formation of similar clouds of ammonium chloride, when hydrogen chloride (hydrochloric acid gas) comes in contact with ammonia.

If hydrogen bromide is heated, it decomposes into its constituents, hydrogen and bromine. To show this, soak a sheet of filter paper in an alcoholic solution of fluorescein, and dry it. Then dampen the yellow-stained paper with water and wad it into the mouth of a test tube filled with hydrogen bromide gas. Hold the test tube in a Bunsen flame. The heat will release free bromine, which will change the hue of the paper to a reddish coloration. The bromine reacts with the fluorescein to form eosin, a red dye. This test for free bromine can also be adapted to tell whether a salt is a bromide. Usually it is sufficient to heat the salt with strong sulphuric acid, in a test tube plugged with the fluorescein test paper as before. If the salt is a bromide, hydrogen bromide will be formed and the sulphuric acid will decompose it, releasing free bromine and turning the paper pink or red.

Close a test tube of hydrogen bromide gas with your thumb, invert it, and open it under water. As the gas dissolves, water will rise in the tube. The solution of hydrogen bromide in water is known as hydrobromic acid—just as hydrogen chloride, dissolving in water, forms hydrochloric acid. So soluble is ¦ hydrogen bromide gas that 612 volumes of it can be dissolved in one volume of water.

To make hydrobromic acid solution for your tests, you could simply let the gas from your apparatus bubble through water in a test tube. A better way, however, is to fit a distilling flask to the side-necked test tube by means of a cork attached to the side neck. The distilling flask should contain about ten cubic centimeters (three teaspoonfuls) of water, and its arm, pointing downward, should dip into the same quantity of water in a test tube, as shown in the illustration. Hydrogen bromide gas from your apparatus first passes into the bulb of the distilling flask, where the water greedily absorbs it. Any gas not recovered here will dissolve in the water-filled test tube below. After letting the gas bubble through the system for several minutes, disconnect the distilling flask, and combine the solutions that the distilling flask and test tube contain.

This solution of hydrobromic acid, you will find, has distinctly acid properties. In fact, it is strong enough to dissolve metals such as zinc and magnesium. Drop small quantities of these metals into portions of the liquid, and hydrogen gas will be evolved. The metal itself is converted into a bromide salt.

Hydroxides of the various metals are easily dissolved by hydrobromic acid. You can make some copper hydroxide for this test by adding a small amount of ammonium hydroxide to a clear solution of copper sulphate, and filtering to recover the resulting precipitate. This solid residue of copper hydroxide will readily dissolve when you pour some of your hydrobromic acid solution upon it.

With iodine, another member of the halogen family, you can carry out a mysterious and spectacular experiment. This test calls for solid iodine crystals (not the liquid “tincture of iodine” used as an antiseptic, which is a solution of the crystals in alcohol) and should be performed outdoors.

Mix a quantity of the iodine crystals with about twice their volume of metallic aluminum powder, such as is used in aluminum paint, by thorough stirring in a bone-dry porcelain crucible or tin-can lid. So far, no reaction has taken place. Now add a drop of water to the mixture. In several seconds, when the water wets the aluminum metal, things commence to happen.

The iodine-aluminum mixture becomes warm. Suddenly it glows with a soft, red hue. Purple fumes of iodine vapor issue from the mass. (It is to dissipate these fumes that the experiment is performed outdoors.) While the vapor is being emitted, the mixture will continue to glow. Then, as it starts to cool, the dying glow of the aluminum oxide—formed when the aluminum burns in the air—is spontaneously rekindled to brightness. This phenomenon is known as “recalescence.” Finally a cold, twisted residue is left.

Curiously, the drop of water took no chemical part in the reaction; it simply acted as a catalyst or “trigger” to set things going. No one seems to know just how or why a catalyst works. But in some mysterious way it makes certain chemicals interact, simply by its presence.

Here is an experiment with another compound of the halogen family—a hypochlorite —which shows that two catalysts are sometimes better than one.

For the hypochlorite used in this test, you can make a solution of calcium hypochlorite by dissolving (Continued on page 230) about ten grams (two teaspoonfuls) of bleaching powder in 100 cubic centimeters (three and a half fluid ounces) of water, and filtering to obtain a clear solution. Or you can use the straight, undiluted solution of sodium hypochlorite sold under various trade names at drug and grocery stores, for bleaching and for whitening clothes.

Half fill three test tubes with either of these hypochlorite solutions. To one tube, add about five cubic centimeters (a teaspoonful and a half) of a dilute solution of copper sulphate. To the second tube, add some ferrous (iron) sulphate. To the third tube, add five cubic centimeters of copper sulphate solution and an equal quantity of ferrous sulphate solution. Shake each tube and let them stand.

Soon you will see gas bubbles forming in the third tube, containing both copper sulphate and iron sulphate, although there will be practically no evolution of gas in the two other tubes. The gas is formed by the decomposition of the hypochlorite solution. Insert a glowing straw in the third test tube, and it will flare up and burn with a vivid light, showing that the gas is oxygen. Here is an instance in which two substances, neither of which could be considered a catalyst alone, do a nice bit of teamwork as catalysts when placed together.

Boy Chemist “Eats Up” Course in Foodstuffs

Relationship between the fields of chemistry and cookery is the research project that interests seventeen-year-old Edgar Friedenberg, the youngest man ever to appear on a program of the American Chemical Society. Friedenberg is pictured below taking time off from his studies in synthetic foodstuffs to try a little practical work with the frying pan.

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Experiments With Oxygen FOR THE AMATEUR CHEMIST

A few common chemicals supplied by the druggist and simple apparatus is all that is required to produce these interesting experiments with oxygen.

by VERNON TRACEY

OXYGEN experiments form a very interesting field of adventure for the amateur chemist due to the fact that oxygen is one of the most active of the chemical elements. It readily combines with most any other element to form many different compounds. These compounds of oxygen and other elements are known as “oxides” and the process of combination is called “oxidation,” or more commonly known as burning. We see examples of oxidation every day in the burning of fuel, but this is not very active when one considers the fact that the air is only one-fifth oxygen, the rest being mainly nitrogen and a small percentage of other gases.

The amateur chemist can produce pure oxygen by heating a mixture of potassium chlorate—a white powder which can be bought cheaply from any druggist or dealer in chemical supplies, and manganese dioxide—a black powder. The mixture is composed of two parts of potassium chlorate to one part of manganese dioxide and is heated over a bunsen burner or alcohol lamp. The test tube is supported over the flame by a ringstand and the open end of the tube is fitted with a rubber cork. An elbow made from glass tubing is fitted into the cork and a length of rubber tubing leads to a bottle in the pneumatic trough where the gas is collected.

The pneumatic trough consists of a pan fitted with a wooden shelf about an inch from the bottom. The trough is filled with water high enough to come about an inch above the shelf. The shelf has a small hole drilled through it and a medicine dropper or piece of glass tubing drawn to a point and bent to a right angle is fitted into the end of the delivery tube and leads up through the hole in the shelf.

To collect the gas, fill a bottle with water, cover its mouth with a piece of glass, invert it in the pneumatic trough and remove the glass again. The mixture is heated in the test tube and after allowing the gas to bubble up for a few seconds, invert the bottle of water on the shelf over the end of the delivery tube. The gas will bubble up into the bottle and displace the water. After the bottle is full, cover its mouth with a piece of glass and remove it from the shelf. It can be set out on the table if the glass is left over its mouth to prevent the gas from escaping.

Eight grams of the chlorate mixture will make several large bottles of oxygen. After the gas ceases to flow, remove the test tube and allow it to cool, otherwise water will be drawn back through the delivery tube and break the glass. The contents of the test tube —fused potassium chloride and manganese dioxide can be removed by dissolving in water.

In removing the oxygen the potassium chlorate is reduced to potassium chloride but the manganese dioxide which is used as a “catalytic agent” to speed up the reaction undergoes no change itself. If so desired it can be filtered out and used again. A sheet of blotting paper fitted into a funnel will serve as a filter.

Steel wool will flare up and burn with a brilliant light if held in a bottle of oxygen as can be seen in the photo. A wad of it is held with forceps over a bunsen flame until it begins to glow, the cover is removed from the oxygen and the wool is thrust into the bottle where it immediately bursts into flame. Bits of molten steel will fall and spatter on the bottom of the bottle creating the effect of fireworks and care should be taken not to get the face too near the mouth of the bottle. Iron oxide will be deposited on the bottom of the bottle at the end of this experiment.

Oxygen will combine with steel wool however at ordinary temperatures, although much more slowly. To demonstrate this, force a wad of steel wool into a test tube, fill it with water and collect it full of oxygen in the manner just described. Support it over a glass of water on a ringstand so the mouth of the test tube is submerged to a depth of about an inch. Allow it to set all night and upon examination the next day the water will be found to have risen up to a considerable height in the test tube. The steel wool has rusted and is covered with a coat of iron oxide. This shows that the oxygen has combined with the steel wool and left a partial vacuum in the test tube which in turn draws the water up from the glass.

We see examples of iron oxidizing at ordinary temperatures every day but this takes place more slowly than the steel wool did in the pure oxygen. There is just as much heat produced at the end of this slow oxidation as when the wool was burned in the bottle of oxygen; but so slowly that it is dissipated as fast as it is produced.

If, however, this heat given off by slow oxidation cannot escape as fast as it is produced as in the case of a mow of damp hay or a pile of oily rags the heat keeps accumulating until the kindling temperature is reached and the material bursts into flame. This is known as “spontaneous combustion” and is a common cause of fire in farm buildings.

You have no doubt heard the joke about making the match burn twice—well, here’s how to do it: Light a match and allow it to burn for a few seconds. Blow out the flame and if a glowing ember remains on the end, thrust the match into a bottle of oxygen and it will immediately burst into flame again with a slight harmless explosion. This action can be repeated several times until the oxygen becomes used up.

A glowing charcoal dropped into a bottle of oxygen will burn vigorously. Cover the bottle with a glass while the charcoal is burning to collect the carbon dioxide formed. After the burning has ceased, pour about an ounce of lime-water into the bottle and shake. The lime-water will turn milky; this indicates the presence of carbon dioxide.

Similarly if sulphur is burned in oxygen, weak sulphurous acid will be formed. Test with blue litmus paper which turns red on contact with acid.

Chemists who have no gas supply or who wish to heat large flasks will find a one burner electric stove very convenient. Flasks should be supported an inch or so above the heating element by means of a ringstand as shown in the photo. There is little danger of breaking the glass as long as it does not come in direct contact with the red-hot wires.

Ordinary drinking glasses will be found very suitable for collecting small quantities of oxygen. They are easier to handle than test tubes and hold enough oxygen for ordinary tests. In collecting gases in this manner, slide a piece of glass over the mouth of the vessel before withdrawing it from the water to prevent the gas from escaping, as shown in the photo.

Whether he has much knowledge of chemistry or not, the amateur chemist should be interested in the practical use of this gas.

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Dangerous ACIDS MADE SAFELY BY Home Chemist

By Raymond B. Wailes

BECAUSE they enter into a wide variety of reactions, acids form an interesting and important group of chemicals. By preparing them in small quantities, the home experimenter can learn a great deal about chemistry and its many mysterious reactions and valuable processes.

The fact that many acids are considered dangerous should in no way dampen the amateur chemist’s ardor. Handled cautiously, they are as safe and harmless as a sharp knife in the hands of one who is careful and dexterous. They should, of course, be stored in glass bottles and kept away from clothing and hands. If some acid is spilled accidentally, it should be neutralized immediately by applying a base such as ordinary baking soda.

When diluting a strong acid, always pour the acid into the water, adding it slowly and stirring the mixture with a glass tube or rod. Never pour the acid in quickly. If you do, enough heat may be generated when the two liquids mix to form steam bubbles that will blow the acid and water out of the container.

Although the amateur chemist with his meager supply of equipment cannot prepare concentrated sulphuric acid in his home laboratory, he can manufacture it in a weak form that will illustrate the method and serve to introduce an important chemical phenomenon called catalysis.

To prepare sulphuric acid, you will need some sulphur, water, calcium chlo- ride, and iron (ferric) oxide. The experiment is a simple one and requires only homemade apparatus consisting of a bottle, a flask, glass tubing, a few corks, a glass funnel, a gas burner, and rubber tubing. The parts should be arranged as shown in the illustrations. Flowers of sulphur placed in the shallow lid from a tin can is burned under the funnel at the extreme right. The sulphur dioxide formed together with some air is collected by the funnel and then passes through a drying bottle, containing the calcium chloride, to the horizontal tube of hot iron oxide. The presence of the hot iron oxide causes the sulphur dioxide to steal oxygen from the air and become sulphur trioxide. Because in this reaction, it induces a chemical change in another substance and is unchanged itself, the iron oxide is said to be a catalyst.

Finally, the sulphur trioxide formed is bubbled through water in the absorbing flask at the left. Being soluble, it combines with the water and a weak solution of sulphuric acid results.

Unaided, the original sulphur dioxide formed by the burning sulphur would not follow the desired course through the various tubes and bottles. To pull it through the system, suction must be applied to the mouth of the absorbing flask. This can be done by allowing water to siphon from a gallon jug and applying the suction formed in the jug to the absorbing flask by means of a length of rubber tubing as shown in the drawing.

To prepare the iron oxide catalyst for this experiment, soak some asbestos fiber or pumice stone in iron chloride or some other iron chemical solution until the mass is well saturated. Then add ammonium hydroxide (ordinary household ammonia will serve). This will precipitate iron hydroxide in the pores of the asbestos or pumice. The liquid then can be poured off, fresh water added and shaken and also poured off.

Next heat the impregnated pumice or asbestos in a crucible or tin-can lid over a gas burner. This final operation will convert the iron hydroxide into the desired iron oxide. The finished catalyst then is placed in the horizontal tube and heated gently with a gas burner as the sulphur dioxide is pulled through.

After burning about a teaspoonful of the sulphur, remove the absorber from the system and test the liquid with a piece of blue litmus paper. If an acid is present, the paper will turn pink. To prove that it is sulphuric acid, place a small quantity of the liquid in a test tube and add two drops of hydrochloric acid followed by several drops of barium chloride solution. If sulphuric acid is present, a white precipitate will be formed.

Although sulphuric acid made by this simple process will be weak, it should dissolve bits of magnesium and attack pieces of zinc to produce tiny bubbles of hydrogen gas. Of course, the concentration of the liquid can be increased by boiling but even then the home chemist will find that the acid will be too weak -to be of any great value for experimental purposes. ‘ It is interesting to note, however, that this same type of contact process is used commercially to manufacture sulphuric acid. Of course, a more expensive substance, usually a form of platinum, is used as the catalyst.

While the home chemist will be interested particularly in the chemical uses of sulphuric acid, he can perform a novel experiment to illustrate one of its important physical properties. In a concentrated form, sulphuric acid is capable of absorbing large quantities of moisture from the air. For this reason, it is often referred to as being hygroscopic. To understand this action more clearly, place some strong sulphur- ic acid in a small vessel and expose it to the air. The acid will absorb so much water from the surrounding air that it soon will overflow the container.

Besides many of its other valuable uses, concentrated sulphuric acid can be used to produce another useful chemical —nitric acid. This is done by placing some sodium nitrate or potassium nitrate in a glass retort containing a quantity of sulphuric acid made by mixing equal parts -of the acid and water. When the chemicals are heated, nitric acid vapors will be given off and can be condensed to a liquid by cooling.

To condense these vapors, the best procedure is shown in the photograph. Insert the end of the retort outlet tube in the mouth of a flask and rest the flask in a glass funnel. A stream of water directed on the upper face of the flask then will serve to cool it and condense the vapors leaving the retort. The funnel will serve to catch the cooling water which can be led through a rubber tube to a drain or a large pan or bottle placed on the floor.

Nitric acid manufactured by this method will be found to be quite energetic in its action with metals, carbonates, and other chemicals. Because of its activity, it should be stored in glass-stoppered bottles. It will attack both cork and rubber.

By using sulphuric acid and a small amount of iron sulphate solution, the home experimenter can test for the presence of nitric acid or nitrates. Simply place about a quarter of an inch of the sulphuric acid in a test tube, add an equal amount of iron sulphate solution, being careful not to shake the tube, and then slowly add the liquid to be tested by allowing it to run down the walls of the tube. If a brown ring is formed when the solution reaches the area between the acid and the iron sulphate and gentle heating causes the ring to disappear, it is proof that either nitric acid or a nitrate is present.

Hydrochloric acid, a third member of the important acid family, can be produced by adding ordinary table salt to sulphuric acid and heating the mixture.

Like nitric acid, hydrochloric acid also should be made in an all-glass retort. The end of the exit tube dipped into a water-cooled flask of water then will lead the gas through the water where it will be dissolved to form liquid hydrochloric acid. Although the home chemist can manufacture hydrochloric acid by this method, it will be less expensive and troublesome to use commercial muriatic acid (slightly impure hydrochloric acid).

It is a simple matter to test the distillate formed for hydrochloric acid. If a drop of silver nitrate solution is added to any solution of a chloride, a white curdy precipitate will be formed. Exposed to the sunlight, this precipitate of silver will change to a dark brown owing to decomposition.

An interesting experiment showing how heating may decompose a substance can be performed with some sal ammoniac (ammonium chloride). Being produced when hydrochloric acid gas comes in contact with ammonia gas, it can be made to break apart again by applying heat.

To separate the two gases when they are set free, the home chemist must employ a niterlike wad of asbestos fibers or other nonflammable substance rammed into a glass tube to form a plug. Ammonium chloride then is inserted into the tube at one side of the plug and the tube is mounted horizontally above the small flame of a gas burner.

IN A few seconds, the ammonium chloride will begin to decompose to form hydrochloric acid gas and ammonia gas. Being lighter than the hydrochloric acid gas, the ammonia will diffuse, spread, or travel faster and will issue from the open end of the tube nearest the porous plug. The presence of the gas can be shown by holding a moist strip of red litmus paper near the mouth of the tube until it turns blue. Similarly, the hydrochloric acid gas will issue from the other end of the tube and will give evidence of its presence by coloring damp blue litmus red. In these experiments with acids, and in fact in any experiment where a chemical in a long tube must be heated evenly, the flame-spreading attachment shown in the photograph will form a valuable addition to your gas burner. If you made the burner previously described (P.S.M., May ‘33, p. 53) you will recall that the stack was made from a six-inch piece of three-eighths-inch iron pipe. To make a flame spreader, simply select a three-eighths-inch pipe cap, saw three slots across the top of the cap sixty degrees apart, drill holes at the ends of each slot, and finally screw the cap into place on threads cut in the upper end of the burner.

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How Scientists Are Taking the Pinch Out of America’s Billion-Dollar Shoe Bill

New Tanning Discoveries Will Bring You Cheaper Footwear By John Walker Harrington

WELL-SHOD feet are among the essentials of health and long life,” declared Dr. John B. Huber in a recent article in POPULAR SCIENCE Monthly.

The magnitude of our national shoe bill is revealed in this story of new discoveries in tanning, which hold forth hope of a coming fall in every family’s expenses for footwear.

THE expense of keeping the feet of the American people off the ground has been so high since leather went to war that every shoe pinches a pocket-book. The enormous cost of keeping us shod may be comprehended when we consider that according to the last available census of manufactures, $1,149,560,000 worth of boots and shoes are produced annually in the United States.

In fact, this sum nearly equals the cost of providing our daily bread; for the value of wheat flour made annually in the United States, according to the same census, was $1,436,388,000, or only about $300,000 more than the sum at which the boots and shoes were appraised.

New Discoveries in Tanning Science, however, is now at work to slice America’s shoe bill. In the chemical laboratory and in the machine-shop speedy processes of tanning leather are being substituted for the slow and expensive methods handed down through generations. New materials that can be made available for footwear are being discovered and developed. As a result, there is substantial hope that new, improved processes will help to cut at least 25 per cent from the cost of shoes, saving $287,-390,000 annually.

If all the boots and shoes worn by the men were placed end to end, allowing 12 inches as the length of each, they would reach four times around the earth and still there would be enough left to keep a few regiments on a war footing.

Suppose that some giant cobbler should work all the leather used in all the footgear of the men, women, and children of this country into one big shoe! He would need the entire 800 square miles of the island of Cuba on which to put it. This would not, of course, include the slippers and the babies’ shoes, so he would have to save the island of Manhattan in the center of New York for the infant size.

Some of the younger generation wear cloth and fiber shoes, of which 8,000,000 pairs are made a year, but in the main the hides of the ox and the bull, the skin of calves and of kids is used in the manufacture of footgear. The day is not far distant when the shark and the porpoise, or the white whale, will yield a far larger percentage of their hides to the tannery than they do at present.

As science takes a still deeper interest in our shoes, many more kinds of skins and hides will be made available for our wear. Indeed, the time may come when we shall be satisfied with mashed paper heels and when rubber will be adapted to all kinds of weather.

Looking into the future, there is much hope for less costly footgear in the many new processes for tanning leather quickly that are now being perfected by up-to-date chemistry. Of all the trades, that of tanning leather seems to have been about the last to come under the control of science.

Every commercial variety of skin or hide has to be prepared in some way before it can be used for human wear. The pelts of animals, when removed, will dry up or become hard, or even rot, unless they are prepared to withstand decay. Tanning is the chemical process that preserves them, makes them comfortable to the human foot, and keeps out moisture. When a hide has been tanned, it has under- gone certain changes in its fibers that make it a new product. Chemistry has waved its magic wand over it, and has made it new.

Despite this fact, however, chemistry has had little to say about tanning until the past few years. In the past, the formulae for making the tanning liquors have been handed down from father to son. If hides have been spoiled, the moon has been to blame and the public has paid for the damage.

Old Methods Are Costly What a hard working citizen is the old time tanner, in his rubber boots and his stained apron! As stubborn as the blind mule that grinds the bark for the vats, he wades around in tons of wet bark and hauls away at the dripping hides for months before he is ready to say that he has made leather. Even in some of the most up to date tanneries, from three to six months is required to prepare a hide for the shoemaker. All this is tremendously expensive.

With the new methods, it is possible to tan some hides in 24 hours, and the quality of the goods is as high, and even higher, than it was before science lent a hand in leather manufacture. The college professor has been leading the leather industry into ways of economy by employing substances that will work the chemical changes of tanning quickly.

It took the college professors in the great leather school at the University of Leeds, England, to show the tanner how he could save time and money. Within the past few years, the American scientists have been coming to the fore, as was shown by the arrival of many delegates from England to attend recent meetings of the Section of Leather Chemistry of the American Chemical Society held at Columbia University, New York City.

Certain manufacturers have created a substantial fund for the investigation of the facts of chemistry that apply to tanning, and have placed the reports of those researches at the disposal of their competitors. Part of the fund is used at Columbia by Professor Arthur W. Thomas in running the smallest tannery on earth, which occupies only a square yard. Small bits of hide are in this tiny plant put through all the processes of the big establishment, and their structure is studied under the microscope. At Pratt Institute, in Brooklyn, is another center of the new tanning art, where up-to-the-minute tanners are pupils.

Exact Knowledge for the Tanner One of the new wrinkles in the tanning of skins is the gaging of tanning materials by the electrical discharges that come from them. This process gives to the tanner an exact knowledge of the work that his materials will accomplish. By studying the nature of the structure of the hides more closely in the light of the new knowledge of the colloids, that is of substances similar to gelatine, the tanner of the new school has his material tested in advance, so that he can tell just how it will act, instead of depending upon mere luck.

The modern tanner obtains extracts of his vegetable tanning ingredients, instead of making use of the substances themselves. He has learned, too, that the astringent action is due to tannin, its active principle, and that the making of the astringency is connected with the electrical charge of those tannin particles. The higher the charge, the greater the astringency. As a result of such discoveries, blind use of certain vegetable extracts is no longer necessary, since the properties peculiar to one kind can be obtained from an entirely different extract by simple treatment, according to the principles of colloid chemistry.” One of the most difficult problems in tanning is known as the regulation of bating. When hides are soaked in lime and water for the purpose of removing the hair, they swell, and must be reduced to normal, otherwise they will have a soft texture which will make them unfit for good leather. They are reduced by the bating process, which consists of soaking them in a “pickling” solution. Formerly this process was so uncontrolled, that many valuable hides were spoiled in the bating. By means of the microscope and the hydro- gen electrode, the nature of the action of the bating liquors has been established and a method devised for managing and directing the process that has changed it from the most mysterious to the most scientific process in the tannery.

Hides Tanned in Eight Hours The leather chemist of the present day advocates the chrome process of tanning, where speed is desired. This consists in submitting the structure of the hide to certain chemical changes, by adding bichromate or potash or such substances to the vat liquors. By this method, bides have tanned within eight hours. For many purposes, the chrome leathers are preferred. Certainly they are made in a manner that does not keep the money of the tanner tied up in his vats. In the opinion of the leather chemists, the whole art of tanning is likely to undergo still more radical changes in the next few years.

Through such described shortening of schedules, this revolutionary move will make its influence felt at the cobbler’s bench and in the retail store so that the American public may have better and at the same time cheaper shoes.

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Amuse Friends with CHEMICAL Stunts

DO YOU like to dabble with chemicals? It was a hobby with Thomas A. Edison during his youth and formed the basis of an education that later brought thousands of new inventions into the world. Far from being a “dry” science, chemistry can be very amusing and entertaining. How many people would believe that you could pour a little drinking water into a china bowl and cause it to burst forth with flames several feet high—without the use of matches?

Flower Experiment Is Harmless The experiment is harmless with ordinary care. Use a heavy bowl, of china or metal, and in it place a teaspoonful or more of an inflammable substance, such as benzine. In the center place a bit of potassium metal the size of a pea, no larger. All that you have to do is pour in a little water to make the benzine ignite, as in Fig. 4. Store your stock of metallic potassium in a bottle of kerosene and do not handle it with wet fingers.

A paring knife cuts it as easily as if it were cheese. Perhaps you’d like to repeat this experiment in the same way using a substitute for potassium, a compound that bursts into flame and gives a loud report at the same time. Make only a small quantity at a time, and handle it carefully with a dry slip of paper. First powder a tea-spoonful of antimony-potassium tartrate (the common “tartar emetic”) and mix with it some lampblack equal to the bulk of a pea. Place the powder in the bottom of a test tube, sprinkle a layer of charcoal on top and stopper the tube well. Heat the mixture over a bunsen burner adjusted to give as high a heat as possible without melting the glass. After an hour or so the compound may be allowed to cool; transfer the chunk of material to a well stoppered bottle and after it has crumbled to a powder it is ready for use. Simply pour water on a bit of it in a bowl.

Sulphuric acid can be used to light a candle, as in Fig. 2. The acid is on the end of the partly burned match, while the candle wick is prepared by first dipping in a thick paste of two parts potassium chlorate and one part sugar, which is allowed to dry before the experiment. Powder the sugar and potassium chlorate separately and then mix with a small amount of water.

Here is another amusing chemical experiment. Dissolve a chunk of paraffin the size of a hickory nut in two ounces of benzole and using it with a pencil-size brush draw a picture of a caricature with a very large nose on paper or cardboard.

Fill in the nose portion with a water solution of cobalt nitrate. When dry the drawing will be practically invisible but it can be restored by passing a soft brush, dipped in lampblack, carefully over the lines. The lampblack will adhere to the paraffined lines, as in Fig. 3. Slightly heat the “picture” and the nose will turn red. If, at the time of preparing the drawing, you fill in the nose with a blue solution of cobalt chloride it will be blue until heated; then it will turn red.

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PRACTICAL AND MYSTIFYING HOME TESTS YOU CAN MAKE WITH IRON

By Raymond B. Wailes

MYSTIFYING and spectacular ‘effects give a keen interest to home experiments with iron and its compounds. The amateur chemist can make paint, produce molten iron from a simple mixture, and perform many other stunts that show why iron is man’s most useful metal.

Iron betrays its presence everywhere. Our blood gets its red color from the iron it contains. Soils, clays, bricks, and stones are colored by the iron in the earth’s crust.

A handful of ordinary nails or tacks will serve as the starting point for the home chemist’s experiments. From them he can produce several interesting iron compounds.

Iron, as we know, is attacked or dissolved by acids. This can be demonstrated by placing several iron nails in a test tube of sulphuric acid. To speed up the reaction, warm the tube gently. Soon gas bubbles will be given off and close inspection will discover tiny black specks in the liquid. The gas is hydrogen and the black specks, when filtered out, will prove to be small bits of carbon. Carbon is present in most forms of iron as an impurity. Cast iron, which contains more carbon, will produce more black specks than the wire nails.

As the reaction continues, you may notice a peculiar odor. This is caused by the escaping gas which not only contains odorless hydrogen but acetylene or other hydrocarbon gases as well. The carbon, chemically combined with the iron, reacts to form a carbide which, in the presence of water or a dilute acid, forms acetylene gas easily recognized by its characteristic odor.

When the tiny specks of carbon have been filtered from the solution, the remaining liquid will be a beautiful green. This filtrate is iron (ferrous) sulphate, sometimes called “copperas.” This is an unfortunate name, however, for contrary to its spelling and sound, copperas contains no copper.

If this solution is allowed to evaporate, an interesting mass of green crystals will be formed. However, being particularly efflorescent (drying out rapidly) they will soon crumble to a white powder due to their water of crystallization. To preserve them for future use, place them in a tightly stoppered bottle. Study the crystals closely; they play an important part in the manufacture of paints, inks, and dyes.

PLACE a few of the green crystals of iron sulphate in a porcelain crucible or evaporating dish and heat them. Sulphur dioxide and sulphur trioxide will be driven off as they decompose, leaving a soft reddish-brown powder of iron oxide that is particularly valuable as a polishing material, known as jeweler’s rouge, and as a pigment in the manufacture of red paint. To demonstrate its usefulness as a paint pigment, mix the powder with some linseed oil, thin it with turpentine, and add a drop or two of drier. In preparing the iron oxide be sure to continue the heating of the crystals until all the fumes have been driven off. During this process, stir it continuously to expose all the material.

If, instead of dissolving the iron nails or tacks in sulphuric acid, hydroahloric acid is used, iron chloride will be formed. Crystals of iron chloride contain so much water of crystallization that they often conglomerate into one solid mass. Placed in a beaker and heated without water, they will melt in their own water of crystallization.

When in this fluid state, the substance has a remarkable property. It will dissolve a long strip of thin paper, making it disappear as it is fed into the liquid. The amateur chemist can amuse his friends by preparing a small quantity of the liquid iron chloride and feeding two or three feet of tissue paper into the beaker. Bit by bit, the long strip will be consumed. Then if the solution is poured into a large quantity of water, the paper will reappear as small hairlike particles.

Small iron objects can be coated with copper in the home laboratory merely by immersing them in a solution of copper sulphate. The iron is said to be plated by immersion to distinguish it from the coating obtained by electroplating. Unfortunately, the coating is extremely thin and soon disappears.

NOT every iron surface will exhibit this property, however. Armed with a few simple chemicals, the amateur can treat the surface of a piece of iron to prevent the formation of the copper coating.

First, clean the strip of iron thoroughly with sandpaper and immerse it in nitric acid until the evolution of reddish-brown gases indicate that the iron is being attacked. Then remove the iron, wash it in water, dip it in a solution of potassium dichromate, and wash it again. The iron then is in a passive state, the treatment tending to alter the surface iron. If the strip is immersed in the solution of copper sulphate, no coating of copper will be deposited on its surface. As only the surface of the iron is changed, however, it can be brought back to normal by striking it a blow or scratching it with a sharp pointed instrument. The blow or the scratching breaks the passive skin effect and again causes the copper to be deposited where the skin has been broken.

This peculiar quality makes it possible for the amateur chemist to perform a surprising experiment. Treat the surface of a strip of iron as described so it is in the passive state. Then, using the sharp point of a nail, write your name on the surface of the strip. The passive skin effect will be broken where the point of the nail touches the surface and, by immersing the strip in copper sulphate, the words will be made visible by the line of copper deposited. The use of chromates and dichromates to protect the surface of iron from rust and corrosion is put to practical use. Another method of protecting iron, somewhat similar to the method of obtaining surface passivity, is called “parkerizing.” It consists in causing the surface of the iron to become coated with a thin layer of phosphate of iron. One simple way of doing this is to heat, in a solution of phosphoric acid, the iron that is to be corrosion-proofed. The home chemist can park-erize in a small way by dipping the iron into the ordinary phosphate solution used at soda fountains.

SINCE iron combines readily with oxygen to form iron oxide, it would seem logical to suppose that iron oxide could be reduced to form iron. This is true. In fact, this is what is done when iron is obtained originally from the iron ore. Iron oxide (iron ore) is heated in a large blast furnace with carbon in the form of coke. The carbon combines with the oxygen in the iron oxide, leaving pure iron. A process of this type is called reduction; the iron oxide being reduced to iron.

ALTHOUGH this particular process would be too bulky to duplicate in the home laboratory, the amateur can simulate the reduction of iron oxide by using easily obtained hydrogen gas in place of the carbon. Place the iron oxide, which you obtained by heating the iron sulphate, in a non-metallic tube. Heat the tube and pass hydrogen gas through it. The hydrogen will combine with the oxygen and reduce the iron oxide to iron. If your attempt has been successful, a magnet will attract the small particles of iron formed.

The fact that aluminum also exhibits the property of reducing iron oxide to iron forms the basis of a very interesting and useful industrial process. In factories and repair shops, this simple chemical reaction is known a* the “thermit” process of iron welding. Iron oxide is mixed with aluminum and ignited. This causes the two chemicals to react and give off a great amount of heat. In fact, so great is the temperature that the iron, freed by the reaction, flows from the base of the crucible in a molten, white-hot state.

How this is used to form a weld is shown in the drawing. A mold first is made around the portions to be welded. Then the thermit mixture of aluminum and iron oxide is placed in a conical crucible fitted into the entrance gate of the mold. When the mixture is ignited by means of a strip of magnesium, the reduced iron, heated to an enormously high temperature, flows into the mold.

With some aluminum paint powder, a small amount of powdered magnetic iron oxide, and an old flowerpot, the amateur chemist can reproduce the thermit process on a miniature scale. Mix the iron oxide and the aluminum powder in the ratio of one to two and place the mixture in a three-inch flowerpot. It will be necessary, of course, to plug the hole in the base of the flowerpot.

To ignite the thermit charge, a priming starter will be needed. For this starter, the experimenter can use five parts of barium nitrate, two and one-half parts of aluminum powder, and one part of sulphur by weight. Place this in a small depression in the top of the thermit. To light the charge, use two or more large matches held in a pair of pliers.

Once started, the reaction continues with intense brilliancy and heat. Few experiments are as spectacular as the reduction of iron oxide by the aluminum in the thermit process. The mass will glow and sparks will jump as the aluminum burns by removing the oxygen from the iron oxide.

When experimenting with the thermit mixture, always work with one or two ounces of the mixture and confine the reaction to some heat-resistant and insulating receptacle like the flowerpot. Good results will not be obtained if the mixture is fired in piles as, in that case, the heat is dissipated too rapidly.

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Hints for Beginners in Amateur Chemistry

Join in the Fun of Experimenting at Home! This Article Tells How Easy It Is to Start

By RAYMOND B. WAILES

IF YOU have been following this series of articles for some time, you probably have already set up a more or less complete chemical workshop in which to carry on your experiments. However, there is always a new crop of beginners coming along—newcomers who would like to join the fun and who need some simple advice on equipment and working methods. Old-timers surely won’t begrudge this space to help others get started in the fascinating pastime of amateur chemistry—and perhaps their own memories will be refreshed with a pointer or two.

Let’s suppose, then, that you are starting from scratch—and would like to have a place where you can perform “magic” tricks with chemicals, test household preparations, and carry out a great variety of beautiful and spectacular chemical experiments. Where do you start, and how?

First, you will want a corner of your home where you can permanently arrange your paraphernalia, out of the way of others. You shouldn’t have to put up with the inconvenience of using the bathroom or laundry as an improvised laboratory—even if the rest of the family are broad-minded enough to let you! The attic, the cellar, or a spare room will give you a place where you can work undisturbed, and leave equipment for an interesting experiment set up as long as you wish. You can also use part of the garage, but solutions will have to be protected against possible freezing in cold weather.

In many ways, the basement makes the best choice. Gas, the perfect laboratory fuel, can be tapped from the pipes here and led to your chemical bench. Usually electricity will also be available. The ideal home laboratory would be one supplied with gas, electricity, and running water. If necessary, however, you can get along without them. Instead of a gas-burning Bunsen burner, for example, you can use an alcohol lamp for heating test tubes and flasks and for bending glass tubing. Electric heat will also serve. Where high temperatures are called for, a gasoline or alcohol blowtorch may be used.

Even if it is nothing more than a packing case with a board nailed on for a shelf, you will need some sort of a chemical workbench. Once you have caught the “bug” of experimenting with chemicals, you will aspire to a more commodious piece of laboratory furniture, and you can easily make it for yourself. An old kitchen table, or a new one that you can buy cheaply unfinished, makes a first-class foundation for a chemical bench. On this base you can attach substantial shelves, bookcase-fashion, to hold your stock of chemicals and a few pieces of laboratory glassware that you use most frequently. It is a good plan to construct three shelves, five inches wide and of seven-eighths-inch stock, running along the back of the table. Supported by uprights of five to six-inch width, from the same stock, the shelves may be four feet long without sagging when filled with chemicals. Tallest and heaviest bottles go beneath the first shelf, which should be eight inches above the table top, while six-inch spaces suffice between the upper shelves. Two coats of battleship-gray paint will enhance the job. Pegs that may be added for holding and drying flasks, as shown at the right of the workbench illustrated, should be left un-painted.

If you lack a supply of running water, a substitute may be improvised by fitting a gallon jug as a siphon. Insert a two-hole stopper carrying a short length of glass tubing as an air vent, and a longer section of tubing that reaches to the bottom of the jug. To the latter, attach rubber tubing and a pinch clamp. Set the jug on a shelf above your bench, and you can draw off water as needed by squeezing the pinch clamp. Smaller quantities of water may be dispensed from a “wash bottle,” a useful aid described in a later paragraph.

A gallon crock, or a metal pail with several interior coats of paint, will serve as a receptacle for waste. Spent solutions, used filter papers, burnt matches, and cork borings may be thrown directly into it. An added convenience is a drain made by mounting a funnel, or the inverted top cut from a large bottle, at the side of the bench. With rubber tubing leading to the waste bucket, this makes a handy sink for pouring off liquids. What makes a suitable assortment of chemicals and apparatus to start your experiments with ? Often, a novice at stamp collecting makes a good beginning by purchasing one or more inexpensive “packets” containing a large, mixed collection of varieties. Likewise, there are excellent chemical “sets” or “kits” on the market that provide a beginner with a representative variety of materials at very low cost. You can purchase either an assortment of chemicals alone, or a kit that also includes such permanently useful laboratory accessories as test tubes, test-tube holders, an alcohol lamp, and a beaker or flask, together with a tripod for heating its contents.

If you are assembling your apparatus separately, your principal needs in your first experiments will be a number of test tubes, a test-tube holder for handling them above a flame, and some kind of a heater—a Bunsen burner, alcohol lamp, or electric stove. Test tubes four to six inches long, and half to three quarters of an inch in diameter, are good standard sizes. Useful items of equipment also include beakers, flasks, glass funnels, a graduated cylinder or two, a test-tube rack, a porcelain crucible, an evaporating dish, and an assortment of cork stoppers, cork borers, and glass tubing. A small photographic balance, preferably with gram weights, will also come in handy for weighing out chemicals.

Of course you don’t have to buy all this at once, but can add as you go along. Many interesting and practical experiments require only a test tube or two and a few inexpensive chemicals.

Suppose, let’s say, you want to know whether the hydrogen peroxide in your medicine cabinet has lost its strength. Just add a drop or two of hydrochloric acid to a sample of it in a test tube, then several drops of potassium di-chromate solution, and heat the contents of the tube. If a blue or green color appears, the peroxide is still good.

By equally simple experiments, the presence or absence of many substances in an unknown compound may be confirmed. Moisten baking powder with water, wait until the bubbling stops, and add a drop or two of a solution made by mixing ten drops of tincture of iodine with six teaspoonfuls of water. If a blue coloration is formed, the baking powder contains starch. Carbonates, like marble or washing soda, effervesce when you add an acid. Ammonium compounds, such as sal ammoniac, can be recognized by the odor of ammonia when you heat them with an alkali.

Apparatus for more complicated experiments need not necessarily be purchased ready-made. You will be surprised to find the variety of equipment that you can put together from odds and ends.

An empty paste jar makes a fine alcohol lamp, when you solder a metal tube to the screw top and pass a round cotton wick through it. Transformed as shown in one of the illustrations, a half-dollar electric stove becomes a highly serviceable laboratory heater. A bent metal clamp, drilled with two holes, attaches an iron rod or laboratory support to the base. The same illustration shows a tricky way of using this heater in evaporating a large quantity of a solution, with an inverted flask arranged to give an automatic feed.

A piece of wire with loops twisted in its ends will hold a test tube, for heating its contents over a flame. Always apply the heat near the surface of the liquid in the tube, which should be held nearly horizontal and with the mouth pointing away from you. This will prevent cracked test tubes, and keep any liquid that spatters from striking you or your clothing.

Bore holes in one of a pair of thin boards, mount them one above the other in a wooden stand with the bored one uppermost, and you will have a serviceable rack for your test tubes. Small bottle brushes, from the five-and-ten-cent store, will help clean them after use.

If you have never tried cutting and bending glass tubing, you will be surprised to find how easy it is to make pieces to order for connecting your apparatus. Nick one side crosswise with the edge of a three-cornered file, press with your thumbs on the opposite side of the tubing while you hold it in your hands, and it will break cleanly in two. Twirl the cut end in a flame and it will be smoothed or “fire-polished.” To bend glass tubing, roll it in the flame until it is softened, and it will then take any desired form. A wide flame will avoid a flattened, constricted bend; use a flame spreader with your Bunsen burner, or make three miniature alcohol lamps from medicine vials, by fitting the corks with metal or glass tubes and wicks, and mount them side by side in a wooden block. To form a small nozzle from tubing, heat it glowing hot and draw it out like taffy as you remove it from the flame, cutting off the tip at the point desired. Connections between pieces of glass tubing may be made with short lengths of rubber tubing.

For your wash bottle, mentioned earlier, fit a good-size jar or chemical flask with a two-hole stopper carrying two bent pieces of glass tubing. The longer, reaching to the bottom of the water-filled flask, should have a small nozzle for a tip. When you blow into the other, which reaches only to the bottom of the cork, water will squirt from the nozzle into a test tube or other vessel.

To store your laboratory apparatus, you can press into service a discarded kitchen cabinet, a chest of drawers, or an old wardrobe fitted with wooden shelves. An ideal storage cabinet would be a double-size steel locker, or kitchen cabinet, of a type sold widely in department stores.

Chemicals and apparatus may be purchased according to your requirements, as you progress with your hobby, from a number of chemical-supply houses that handle mail orders. Some drug stores in large towns also specialize in stocking a wide variety of laboratory supplies, and even the nearest corner pharmacy will be able to provide a number of the chemicals that you may need.

A list of some of the chemicals most frequently used in home experiments might read as follows (the strong acids listed are to be handled with particular care) :

Ammonium chloride; ammonium hydroxide (household ammonia can be used); barium chloride; calcium carbonate (marble); calcium oxide (lime); cobalt chloride; cupric chloride; cupric oxide (black copper oxide); cupric sulphate; ferric chloride; ferrous sulphate; ferrous sulphide; hydrochloric acid; lead acetate; magnesium metal (in ribbon form); magnesium sulphate (Epsom salts); manganese dioxide; manganese sulphate; nickel ammonium sulphate; nitric acid; phenolphthalein (one-percent alcoholic solution); potassium chlorate; potassium di-chromate; potassium iodide; potassium nitrate; potassium permanganate; potassium thiocyanate; silver nitrate; sodium bicarbonate (baking soda); sodium bisulphate; sodium carbonate (washing soda); sodium ferro-cyanide; sodium hydroxide (lye); sodium silicate (water glass solution); sodium thio-sulphate; sulphur; sulphuric acid; zinc metal.

In coming articles, some of the fascinating experiments that you can perform with these and other items of your chemical equipment will be described.

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How TO CONVERT OLD ELECTRIC LIGHT BULBS INTO CHEMICAL GLASSWARE

By Earl D Hay

EXPERIMENTS in an amateur chemical laboratory are much more interesting when they are made with the same kind of apparatus as that used in professional laboratories. As a rule, however, the home chemist experiences a great – shortage of flasks and endeavors to use various kinds of bottles as makeshifts, little realizing that he may make from burned-out electric light bulbs a great variety of useful flasks like those sold by chemical supply houses at from 20 to 75 cents each. The lamps used in the average home vary in size from 25 to 200 watts and are suitable for small Florence or boiling flasks. Larger flasks are made from 300-, 500-, and 1,000-watt lamps, which can be obtained from the janitors of stores and linemen of the city lighting companies.

The methods of working all sizes are the same, and only a few minutes are required to complete a flask.

The method of making a Florence flask washing bottle from a 300- or 500-watt lamp will be described. The flat bottom is made first. Cut off the connection of the center wire on the cap with a knife and break off the end of the slender tube which was used in evacuating the bulb when it was made. This is necessary in order to equalize the air pressure on both sides of the glass wall. Next screw the light into a drop-cord socket to provide a handle for holding the bulb while the large end is being heated over a large laboratory gas burner or a gasoline blowtorch.

The bottom of the bulb is carefully warmed up and then heated evenly to a light red color. Now quickly place it in a vertical position on a level wooden block or an asbestos pad and bear down gently. The spherical bulb will flatten on the bottom. If heated too hot, the bulb will wrinkle and become distorted; if not heated enough, too much pressure will be required and the bulb will be broken.

After the bulb has cooled sufficiently to be handled, remove the brass cap from the neck by filing through the threads on a diagonal line as shown in one of the photographs in order not to scratch the glass with the file. Pull the split cap off with a pair of pliers, and scrape off the sealing wax that lies between the brass cap and the glass, taking care not to destroy the two copper wires leading into the center of the bulb.

As the bulb becomes quite hot while the neck is being shaped, it is necessary to provide some adequate means of holding it. If a pair of heavy asbestos mittens are not at hand, a satisfactory holder can be made from a piece of strong cloth by cutting a round hole in it large enough to admit the neck of the bulb. The neck is inserted through the hole, and the cloth folded back over the bulb.

The end of the bulb neck is now carefully heated until the glass becomes red and plastic. With a pair of pliers, seize the two copper wires and carefully remove the glass core by pulling straight out on the wires as the bulb is rotated in the flame to keep the entire circumference at the same temperature.

Next take a round, soft pine stick with a conical point and begin to open up the mouth of the neck and roll a bead on it by rotating the neck in the gas flame and rubbing the plastic edge out and down with the wooden stick. This enlarging process is continued until the neck will take the desired size of cork or rubber stopper.

The flask is now complete and ready for use. If it is to be used for a washing bottle, a heavy rubber band or stout cord wound around the neck will make it much stronger in resisting the stopper pressure and more convenient to handle.

If the bulb is to be made into a boiling or a receiving flask, the bottom will not need to be flattened, and the brass top may be removed and the throat enlarged to the proper size at once. If a heavy smooth lip is desired, it can be made by making a mold of some heat-resisting material as shown in the drawings at the end of this article, and the lip turned down against it. This mold or form must be made in halves and clamped around the neck of the bulb. It must be warmed carefully before use, otherwise the glass will crack.

If a lipless flask is desired, the small end may be removed by placing a string saturated with kerosene around the neck and allowing it to burn away, then quickly plunging the neck into water up to the heated ring. This will cause the glass to contract and pop off at the line where it was heated. The broken edge is then smoothed by carefully grinding it down on a smooth grinding wheel and finishing it with a fine sharpening stone.

If a glass-tube cutter is available, the end of the neck can be removed without difficulty. This method is more reliable than the use of the kerosene string.

Long-necked flasks may be made by welding test tubes or necks from broken flasks to the necks of light bulbs. After a little practice this can be accomplished without difficulty. First be sure the ends to be joined are of the same diameter and fit all way around. This can be accomplished by grinding the ends on a smooth oilstone. Heat the ends carefully and evenly in the gas flame until plastic; then bring them into contact and exert a slight pressure.

FOR A DISTILLING flask, it will be necessary to weld a tube to the neck of one of the larger flasks at a downward, angle of approximately 75 deg. to the neck. A hole is first made about halfway down the neck of the flask by heating the side of the neck to a red heat over the gas burner and then punching the hole from the inside by using a redhot wire with a right-angled hook on the end. A piece of tube of the desired bore and length should be selected, and the free end heated sufficiently to smooth off the sharp edges. The end to be welded to the flask is next heated and flanged. This flange is turned out at a right angle to the tube and should extend about 1/8 in. all the way around it. The neck of the flask and the tube are next brought to a welding heat in the same flame; then the tube is carefully centered over the hole in the flask and the two gently pressed together. Very little pressure can be used or the flask will become distorted. The joint is now heated quite hot and the flange gently smoothed down to make the joint stronger and neater in appearance. In doing this, be careful to support both tube and flask or they will tend to sag out of shape.

AFTER the joint has been completed, the hot flask should be placed in a heated oven and allowed to cool very slowly. This will temper the glass and remove the strains set up in the welding operation. If all the flasks are given the hot-oven cure, they will be less liable to crack in use, especially when heated over a gas flame.

Sometimes the necks are cracked because of careless heating. If carefully cut off, the lower halves of such bulbs make transparent covers and shallow dishes.

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Dry Ice-Capades

Dry ice is very interesting stuff! Get yourself a chunk (handling it with gloves) and perform the simple experiments illustrated here.

DRY ice is solid carbon dioxide. It’s very interesting stuff. For one thing, it sublimes at room temperature; that is, although a solid, it evaporates to form a gas without passing through the liquid state. The mist you see formed by dry ice is water “squeezed” out of the air because it has been chilled below the dewpoint.

Dry ice will readily freeze water and other liquids, and is sometimes used to “quick-freeze” food. The water in plant or animal tissues, under proper conditions, freezes very rapidly and the hard, frozen tissue then is brittle and shatters when struck. The “burns” caused by dry ice are really areas where the body fluids have been frozen. Since ice formation is often accompanied by the growth of needle-like crystals, one can see that these frost-bites can be both painful and dangerous. For the sake of safety, dry ice should be handled with gloves or tongs, not with bare hands.

On the following pages are photographs illustrating a few experiments that may be performed with dry ice. Try them!

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Glass Making Easy for Home Chemist

By Raymond B. Wailes

BECAUSE of its importance in glass making and other industries, silicon opens a particularly interesting experimental field to the home chemist. In nature, silicon is almost as plentiful as oxygen. Yet, it hides itself well in its compounds. It never is found free and uncom-bined and can be separated from its associates only through clever chemical thievery in the laboratory.

Industrially, silicon is obtained by heating sand—a compound of silicon and oxygen—and coke to a high temperature in an electric furnace. The white-hot coke steals the oxygen from the sand to form carbon monoxide and frees the silicon. Although the amateur chemist will have no electric furnace in which to duplicate this process, he can obtain a similar result by heating sand and powdered magnesium over his ordinary laboratory gas burner.

First grind some pure white sand in a mortar and mix the powder that results with about half its weight of powdered magnesium. Place the mixture in a small crucible and cover it with a sheet of tin to exclude the air. The cover should not be clamped on but merely rested in place. Finally, heat the crucible with the blue flame of your burner. Soon after the heat is applied, the magnesium in the mixture will burst into flame. This will increase the temperature until finally the entire mass will glow brightly. The reaction that takes place will be vigorous and spectacular but entirely harmless.

When the flame has subsided, allow the crucible to cool and dump its contents into a small beaker of water. Hydrochloric acid then should be added and the resulting liquid filtered. The silicon freed by the reaction will be left on the filter paper in the form of a grayish powder. Like the coke in the commercial process, the magnesium powder acts as a reducing agent, which robs the sand (silicon dioxide) of its oxygen. The hydrochloric acid is added to dissolve the magnesium oxide formed by the reaction, leaving the silicon free.

Through the use of alkalies, the home chemist can perform other interesting experiments in which sand is torn apart to form new silicon compounds. For instance, by fusing sand with sodium hydroxide (lye) or sodium carbonate (soda ash) and leaching out the products, sodium silicate can be formed. The home chemist may recognize this chemical more readily by the more familiar name, “water glass.” When sand is fused with limestone and sodium carbonate, glass results. Again, because of the comparatively low temperature developed by the laboratory gas burner, the amateur may find it difficult to prepare glass from these three chemicals. However, by using sand, sodium carbonate (or bicarbonate), and lead monoxide (litharge), he should be successful in producing enough glass beads to demonstrate the process. The crucible should be porcelain. A mixture of equal parts of sand, sodium carbonate, and lead monoxide should be placed in a crucible and heated over a gas burner until a clear white or yellowish liquid results. This liquid will be molten glass and can be formed into tiny beads by pouring it onto a tin bottle top and allowing it to cool.

If plain white sand is used in preparing the mixture, colorless or slightly yellow glass will result. By adding metallic oxides to the mixture, however, the glass can be colored to correspond with the metal used. Cobalt, for instance, will impart its characteristic blue color while copper or selenium will color the glass red.

Incidentally, a handy tool for lifting crucibles and other hot containers can be made from an inexpensive serving fork of the type having a finger that can be moved to push meat or vegetables from the tines. Simply cut off a portion of the prongs and the movable finger and bend them as shown in the illustration.

By duplicating the first experiment, in which pure silicon was prepared, in a slightly modified form, another useful chemical—magnesium silicide—can be produced. The two processes differ only in the amount of magnesium powder used. In this experiment, the magnesium and sand should be mixed in the proportions of two to one by weight (instead of one to two) to provide an excess of magnesium.

Grind the sand as before and add the magnesium powder. Then place the mixture in the crucible and again cover it with a sheet of tin. Although, as before, the cover should not be clamped in place, it can be a tighter fit. Heat the mixture to start the reaction. When it is completed, examine the crucible carefully. The magnesium silicide it contains can be scraped out with a knife and bottled for future use.

A combination of silicon and hydrogen, called silicon hydride, can be prepared by adding hydrochloric (muriatic) acid to a small quantity of the magnesium silicide you have made. The product, a mysterious gas, is particularly interesting because it ignites or explodes spontaneously as soon as it is released in the air.

To demonstrate this harmless spontaneous reaction, the home chemist should arrange a simple generator consisting of a glass flask, a two-hole stopper to fit, some glass and rubber tubing, a funnel or reservoir made by cutting the bottom from a bottle, and a pinchcock. The apparatus should be arranged as shown so that water poured into the upper reservoir flows into the flask to displace the air. The glass tube leading from the reservoir should extend almost to the bottom of the flask.

Place the magnesium silicide in the flask, replace the stopper, tighten the screw clamp over the rubber section of the outlet tube, and finally fill the system with water. Ten or fifteen cubic centimeters (about a half fluid ounce) of hydrochloric acid then should be added by pouring it into the reservoir. Because of its weight, it will sink to the bottom of the flask where it will soon come in contact with the magnesium silicide.

In the reaction that follows, bubbles of silicon hydride gas will be given off. However, being prevented from escaping by the pinchcock, the gas will collect in the flask, gradually pushing the water and acid back into the reservoir.

When a quantity of the gas has collected, loosen the screw clamp. As the gas reaches the outer atmosphere it will burst into flame and burn with a bright yellow light. Take particular notice of the smoke that is formed. It contains small particles of silicon dioxide or sand formed by the reaction.

The property of silicon hydride to ignite spontaneously will be illustrated further when the apparatus is taken apart. As each small gas bubble trapped in the tubing and bottle comes in contact with the air it will explode with a harmless crackling and popping.

Like many elements in the chemical family, silicon has a brother. It is called boron. Ordinary household borax and boracic (boric) acid both contain boron and it is with these two inexpensive and easily obtained substances that the home chemist can perform many interesting experiments.

Boric acid is a weak acid and because of its mild antiseptic properties is widely used as an eye wash. When heated, solid boric acid froths as its water of crystallization is driven off, finally melting into a crystal glasslike substance. This glass is boron trioxide and is chemically akin to ordinary sand. Like sand, it will be reduced when ignited with powered magnesium, giving free boron. In demonstrating this reaction, however, it will be best to use commercial boron trioxide, since the homemade product may contain impurities.

Boric-acid crystals can be made in the home laboratory from ordinary household borax. First dissolve a small quantity of borax in water. When tested with litmus, this solution will display an alkaline characteristic by turning red litmus blue. Then add just enough hydrochloric acid to redden a strip of blue litmus dipped in the liquid. This will convert the borax into boric acid.

To obtain the boric acid in crystal form, heat the solution to concentrate it. The boric acid crystals which separate out from the liquid when it cools can be identified by the fact that they will feel greasy to the touch. Finally, dry the crystal by placing them on a sheet of blotting paper.

By making up simple test papers, the amateur can test any solution for the presence of boric compounds. These test papers are made by immersing strips of blotting or other absorbent paper in a solution of turmeric in water and allowing them to dry. You can obtain powdered turmeric from grocery stores.

When using the test paper, first add a few drops of hydrochloric acid to the solution to be tested. Then place two or three drops of the liquid on the paper and allow the strip to dry. This can be done by placing it on the sides of a hot flask of boiling water. If the strip, which originally is yellow, turns pink where the liquid was applied, boron is indicated. As a double check, place the pink portion of the paper in ammonia water. If the first test is correct, the spot will turn black. Since boron is an ingredient of many eye washes, hair lotions, and throat gargles, the experimenter can use the turmeric papers he has prepared to test for its presence.

Another common compound of boron also is sold by druggists as a mouth wash or oral antiseptic. It is called sodium perborate and its cleansing properties are due to the fact that it gives off quantities of oxygen. This can be demonstrated by dissolving some of the powder in water and heating it, testing the gas by placing a small smoldering string close to the liquid. The glowing of the string will be evidence of the oxygen present.

Because of its high oxygen content, sodium perborate also will bleach the color from cloth. The colored portion of wood match boxes can be completely bleached out in a few seconds with a solution of the chemical.

By adding salts to a solution of borax, the home experimenter can produce various precipitates. Copper sulphate solution, for instance, added to borax solution, will form green copper borate. In a similar way, borates of nickel, iron, chromium, cobalt, magnesium, calcium, zinc, manganese, and aluminum can be made.

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Scientific Experiments with Toys

By Raymond B. Wailes

Many Novelty, Toy and “Jokers” Supply Stores sell small glass “meters” or “thermometers.” as they are called, attached to a card supposed to represent the quantity of intoxicating liquor the individual can consume, a state of health, denote a fortune, etc. The items are designed to provoke mirth and hilarity, but they operate on a scientific principle and can be used admirably for demonstrating some physical laws. What to do and how to conduct the experiments are details covered in the accompanying text.

THE glass “meter,” used in these experiments, consists of a glass bulb to which is welded a glass stem; the latter projects into but does not quite touch the interior bottom of the bulb.

The whole device is partially filled with colored alcohol and evacuated. When the hand grasps the bulb, the liquid within the bull) expands and boils up into the stem of the tube, pulsating as if alive. Such devices are sold usually attached with a funny card purporting the instrument to be a “drunk meter” with which the degree of inebriety can be judged. They are also sold in the form of “Storkmeters,” “Fortune Telling” thermometers,” “Love Thermometers,” etc.

A very simple experiment showing that water absorbs infra-red or heat rays, can be shown by holding a small flat, water-filled pill – bottle between a candle flame and one of the meters. The dampening effect on the pulsating liquid contained in the meter is at once noted. The water acts as an absorber, or screen, for the infra – red rays. Alum solution works better than water (Fig. 1).

It is very simple to show that heat rays can be reflected similar to light rays. This can be done by forming a sheet of metal foil (such as is wrapped about tobacco, candy or photo film) into the shape of a concave reflector, and placing a candle flame at the focal point of the metallic screen, as shown in Fig. 2. When the screen throws the maximum amount of light upon the bulb of the meter, the meter will pulse very rapidly. The absorption of heat can be improved further by cautiously and momentarily thrusting the cool glass bulb into a smoking candle flame or the soot from a burning lump of camphor. The coating of soot absorbs the rays of heat more readily. This experiment also shows why dark clothes are “hotter” in the summer.

Alcohol and water become warm when they are mixed. The evolved heat is simply due to the heat of solution. If the tip of the stem of a novelty meter is snipped off, air enters and the vacuum is spoiled. In this form, novelty meter now is really an air thermometer, for if the air within the bulb is heated ever so slightly, the colored liquid in the bulb will rise in the stem; hence it can be used to detect heat.

For the experiment under discussion, place a metal nut over the stem of the air thermometer, made in the manner described, and place the device in a small glass or beaker filled with half an inch of water. The nut serves as a sinker. Now pour alcohol into the water, as illustrated in Fig. 3. It will mix, dissolve, and the heat produced will be shown by a rise of the colored liquid in the stem. Rubbing alcohol, denatured alcohol, grain alcohol, or automobile radiator alcohol, can be used with success in this experiment.

Using a meter with a weight attached to the bottom to float it upright in water, the principle of the hydrometer can be shown. On adding salt, sugar, or an acid, to the water, and stirring, the specific gravity, or density, of the water is raised by converting it into a solution, and the floating meter will rise or protrude farther from the liquid than before. By calibrating the stem, using thread and water-proof cement, a practical hydrometer can be made. Fig. 4 shows the hydrometer.

If the stem of a novelty meter be immersed upside down in a mixture of ice and salt, the volatile vapor of the liquid in the bulb above and outside the cold bath will condense into a colorless liquid in the stem. The heat of the air about the bulb causes more liquid to evaporate and soon the entire liquid which was previously in the bulb is now in the stem, quite colorless. The absence of color is due to the fact that the solid dyestuff does not vaporize. The experiment (Fig. 5) illustrates distillation at room temperature.

A humidity meter, or hygrometer, can be made by wrapping a cloth about the stem of a meter and immersing the tail or loose end of the cloth in a glass of water (Fig. 6). The water wets the cloth on the stem, and evaporates, thus cooling it, and the cooling effect is perceived by a pulsation of the liquid in the stem. It is the difference in temperature between the bulb and the stem which causes pulsation of the liquid within, so by timing the number of throbs per minute, you can arrange a scale of temperatures which will be fairly accurate, on dry days.

The last experiment can be made the basis of a funny little character which constantly appears to suck the colored liquid up into the stem throughout the day and night. The glass of water in the last experiment can be substituted by a narrow vial of water concealed inside a suitable figure which can be obtained at a toy store for a few cents. The cloth wrapped tube enters a hole made in the mouth of the figure. His hands can be bent to grasp the stem also. As long as there is water in the vial and the cloth wrapped about the stem of the meter is kept wet, the colored liquid pulses up and down in the tube. The motion continues for hours, and is dependent on the amount of moisture in the air. On a sultry, foggy, humid day, the movement of the liquid will be somewhat slow because the water of the cloth does not evaporate fast enough to produce a moderately cool temperature of the tube or stem.

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Weird Stunts with Aluminum in the Home Laboratory

Electrical Experiments You Can Perform with This Most Useful Metal—An Easy Way to Purify Water Containing Sediment

By Raymond B. Wailes

OUTWARDLY aluminum is one of the least spectacular elements of the earth. Yet in the home laboratory, weird stunts reveal the strange properties that make it one of the world’s most useful metals.

Although at one time worth its weight in silver, chemistry has made aluminum one of our commonest metals. According to leading scientists, its uses will continue to grow. Even now railroads, steamships, and airplanes make use of its physical qualities for lightness combined with strength.

Most important of its chemical properties is its unquenchable thirst for oxygen. Pure aluminum left in the air soon becomes coated with an oxide. It is this characteristic that makes its impossible to obtain the metal in its free state and also forms the basis of thermit welding (P.S.M., Aug. ‘33, p. 50) and many other modern processes in industry.

To the home chemist, this fast-forming oxide of aluminum offers the means of performing two novel electrical experiments. For the first, immerse two sheets of aluminum foil in a small jar or beaker containing a solution of baking soda (sodium bicarbonate). Connect one sheet directly to one side of the house lighting circuit and the other sheet through a series-connected lamp to the other side.

Then turn on the current. The series lamp will light and a brilliant display of sparks will appear on the surface of the two aluminum sheets. When viewed in a dark room, these sparks will dart and flicker like a swarm of bluish-white lightning bugs.

As the experiment continues, the sparking will grow less and less until finally both aluminum sheets will become incased in a ghostly, soft-white glow. Turning off the current, will stop the glow but it will reappear when the current is again turned on.

Soon, the series-connected lamp, that once was brilliantly lighted, will get dimmer and dimmer. Finally it will go out. A formation of oxide on the aluminum sheets becomes thicker and thicker until it forms a non-conducting wall that cuts down the current.

By substituting a strip of carbon or lead for one of the aluminum sheets, you can transform your novel glow cell into a simple liquid rectifier. Connected to an alternating current source, the cell will act as a one-way street, allowing only direct current to pass.

To test the current flowing through the rectifier circuit, you need only cut one of the wires and place the bared ends on a piece of white paper wetted with a solution of salt water to which a few drops of phenolphthalein solution have been added. If direct current is flowing, the paper around the negative wire will turn red. On the other hand, if the current is alternating, the paper around both wires will turn red. In preparing this experiment be sure the current is shut off when wire is cut. While you are at it, you may as well make up a batch of this prepared paper for future use in your electrical work. Simply place the paper in the salt-water-phenolphthalein solution, allow it to dry, and place it in a tightly stoppered bottle. When you want to make a polarity test, tear off a piece of the paper, wet it, and bring it in contact with the two terminals of the circuit.

BEFORE breaking up your electrolytic rectifier, lift the two electrodes out of the solution and study their surfaces. The aluminum will be covered with a dull white film of oxide. It is this oxide that allows the current to pass only in one direction.

Around the home we find aluminum and its compounds in many of its varied forms. Most common, of course, is as a metal in the large assortment of kitchen utensils. However, when aluminum is combined with potassium, sulphur, and oxygen, it becomes potassium aluminum sulphate or alum—the main ingredient of the styptic pencil you carry in your shaving kit. Liquid deodorants for excessive perspiration also contain aluminum in the form of aluminum chloride. Incidentally, a good product of this type can be made by dissolving about a tablespoonful of the aluminum chloride in half a tumbler of water.

All solutions containing aluminum can be identified by the jellylike precipitate formed when ammonium hydroxide (ordinary household ammonia will do) is added. As a test, make up an aluminum solution by adding a piece of styptic pencil or a crystal of alum to a tumbler of water. When the ammonia water is added, the liquid will cloud up as the thick aluminum hydroxide precipitate is formed.

Many aluminum compounds will react with ordinary water without the addition of the ammonium hydroxide to form the hydroxide of aluminum. It is this curious fact that makes it possible for us to purify turbid water simply by adding some compound of aluminum such as aluminum sulphate or alum.

This action can be shown in a striking way. Select two similar jars or beakers and fill one with water. Drop a pinch of dirt and some household cleaner into the water and pour the resulting liquid back and forth from one jar into the other until the foreign matter becomes well suspended. Then place an equal amount of the liquid in each jar, stir one with a styptic pencil, and set them aside.

In about eight or ten hours compare the two jars. The one treated with the alum will be clear while the other still will be a cloudy, turbid solution. In settling, the jellylike precipitate formed by the addition of the alum will have carried all the dirt to the bottom of the container.

In the dye industry, this amorphous hydroxide of aluminum performs another important task. Many dyes will not enter the texture of some cloths directly. For this reason, the material is first soaked in baths of aluminum sulphate and ammonia water. This causes the aluminum hydroxide to be precipitated onto the fibers where it forms an adhesive for the dye. Chemically speaking, the aluminum hydroxide adsorbs the dye and holds it “fast.” In the industry, substances used in this way are called “mordants,” and the combination of the color and the aluminum hydroxide are referred to as “lakes.” Besides its many other uses in the home laboratory, ordinary alum serves as a particularly good substance for use in the study of crystals. Make a strong solution of alum in hot water and filter it. Then suspend a short length of string into a beaker of the hot liquid. As the solution cools, beautiful jewel-like crystals will form on the string. After several days it will resemble a necklace of clustered stones.

By using an ordinary styptic pencil, the amateur chemist can make use of the crystals of alum to perform a mystifying experiment in magic writing. Words or sentences can be made to appear on a perfectly clean sheet of glass merely by pouring a solution (cold) of alum in water over its surface. The sheet of glass is first prepared by writing some simple word on its surface with the tip of a styptic pencil. The writing can be so light that it will be invisible to the casual observer. However, when the microscopic particles of alum left by the pencil come in contact with the alum solution, they serve as a starting point for a rapid crystal growth. Picking up alum from the solution, these tiny crystals grow until the writing appears as a broad white line.

ALUMINUM powder such as is used in “aluminum” paints, fireworks, and flashlight powders is often put to another practical use that will prove a timesaver for the home experimenter. Combined with an adhesive mixture of the type obtained when celluloid is dissolved in acetone or amyl acetate, a so-called plastic solder is formed.

You can make another type of aluminum cement by heating the aluminum powder with sulphur. Mix one part by volume of the aluminum powder with three parts of flowers of sulphur or rolled sulphur (brimstone) and heat the mixture in an iron container. For small quantities around the home workshop, you can place the mixture in the top of a sleeve-top can and heat it over the laboratory gas burner. Be careful not to overheat it, however. If it should burst into flame, extinguish it quickly by smothering it with a sheet of tin.

Stir the mixture thoroughly during the heating. When it has become molten, pour it into a simple rectangular mold made by bending a narrow strip of sheet metal. To use the -solder” you have made, heat the stick with a match and allow the molten drop to fall into the hole or crack to be puttied. Bear in mind, however, that a metallic putty of this type cannot be used in all cases where soft solder is recommended.

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Crime-Detection Tests FOR THE Home Chemist

How Hidden Fingerprints May Be Found by Using Iodine Vapor — Forgeries Also Are Revealed by This Remarkable Element

By Raymond B. Wailes

NEW thrills await the home chemist who experiments with iodine. Besides its queer properties and varied uses, it serves as the gateway to a new branch of chemistry—the mysterious and interesting art of scientific crime detection.

With iodine, the amateur experimenter can transform his home laboratory into a miniature crime bureau. In a few hours, he can master some of the chemical tricks that aid the modern sleuth in his search for hidden fingerprints, clever check alterations, and forgeries.

First, however, the amateur must learn how to obtain this active element in its free state. For years, it was recovered commercially from a giant type of seaweed called kelp. Now it is obtained from the solutions left behind when Chile saltpeter is crystallized in large quantities.

Its solution in alcohol, commonly known as “tincture of iodine,” furnishes the home chemist with an easily obtained source. To free the solid iodine, it is necessary only to heat the tincture slowly and carefully to drive off the alcohol.

This is best accomplished in a large test tube. After the alcohol has been entirely evaporated, pungent-smelling gases will rise from the tube. These will be the characteristic violet-colored vapors of iodine.

Although this heavy vapor should not be breathed in any quantity, its peculiar odor reminds us of our experiments with chlorine. In fact, although it is a solid at ordinary temperatures, the characteristics and activity of iodine closely parallel those of gaseous chlorine.

No scrubbing bottle or absorber will be needed in our experiments with iodine since the dangerous vapors condense readily on the cool walls of the test tube. Chemically speaking, iodine sublimes, that is changes from a solid to a gas and back to its solid form again.

It is the violet, pungent vapor of iodine that forms an important weapon in the scientific detective’s bag of tricks. To demonstrate one of its most important uses, press your thumb against a sheet of white paper. No image of the lines and pores of your skin can be seen. However, if the paper is brought in contact with vapors of iodine, the thumb print will appear, well defined and clear cut. Even after several months, this system can be used to uncover hidden clues.

These same vapors of iodine also are used by the scientific sleuth to discover forgeries. Untouched paper can be identified from paper rubbed with bread crumbs or an eraser. Dry paper can be distinguished from paper that has been wet and redried, by the color developed by the iodine vapors. Iodine also brings to light any marks or depressions made in paper with a blunt object. The invisible indentations will stand out clearly in a strong violet color when placed in the vapor.

When experimenting with iodine, it is best to obtain the solid chemical from one of its compounds. While tincture of iodine can be used, the solution contains such a small amount of the element that results are not always satisfactory.

Potassium iodine, lor instance, is an excellent source of free iodine since it can be obtained at any drug store. Simply mix it with manganese dioxide and add strong sulphuric acid. Immediately, the violet-colored vapors will appear. By heating the tube or flask containing the mixture, the quantity of iodine produced can be increased.

A simple piece of apparatus for making and collecting the iodine is shown in the illustrations. The mixture is placed in a test tube and an inverted glass funnel is so mounted above the tube that the iodine vapors released will travel up the funnel stem and condense. It is a simple matter then to scrape out the crystals that are formed and store them in a glass-stoppered bottle. If you find that the vapors condense on the upper portions of the test tube instead of in the funnel, heat that portion of the tube also.

Like chlorine, iodine combines readily with many other elements. Zinc powder (zinc dust) and iodine crystals when mixed react slowly to form iodide of zinc. If, however, a drop of water is allowed to fall on the freshly prepared mixture, the combination is instantaneous. A hissing noise is heard and the violet-colored vapors of free iodine can be seen.

Aluminum powder (aluminum bronze used in making paints) and iodine crystals heated in a test tube unite to form aluminum iodide. As the reaction takes place, a vivid glow will be produced. When cool, a few drops of water added to the test tube will decompose the substance with the evolution of heat.

Mysterious iodine explosions can be set off in the home laboratory by mixing nitrogen and iodine (nitrogen iodide.) To make this compound, crush or grind some iodine crystals in ammonium hydroxide (household ammonia will serve). Be sure to keep the crystals under the liquid and stir frequently for an hour or so. Then pour off the liquid and collect the black nitrogen-iodide crystals, placing them on pieces of paper. As these crystals dry, you will note a peculiar property. Each crystal will explode at the slightest touch and a violet cloud of iodine will puff up from the chemical. In fact, the substance often will explode spontaneously. The explosions, which are harmless, will sound like those of small potash caps.

This strange chemical phenomenon is caused by the violent decomposition of the nitrogen iodide. Moist, it is a stable compound, but dry, it decomposes so rapidly that it explodes. When mercury and iodine unite to form mercuric iodide, they open the way to many interesting experiments in both chemistry and physics. At room temperature, mercuric iodide is a red powder. However, when heated to about 150 degrees Centigrade, it changes mysteriously to yellow. Because it displays this property, mercuric iodide is often referred to as being enantiotropic. This color change is due to a change in the crystalline form of this mysterious substance.

In time, the yellow mercuric iodide will return to its original red color. Unaided, this change may require two or three days but it can be brought back in a few seconds by “painting” the substance with a dry brush or by stroking it lightly with your finger.

Although mercuric iodide can be made by heating a mixture of mercury and iodine in a test tube, such a process would be costly; especially, since it can be made in large quantities simply by adding a solution of mercuric chloride (bichloride of mercury) to a solution of potassium iodide.

When the mercuric-chloride solution is added, a red precipitate will be formed.

As this dissolves and disappears add mercuric chloride solution to form more of the precipitate. To get a complete reaction, test the top liquid from time to time, as the precipitate settles, by adding more mercuric chloride. If a precipitate forms, more of the solution must be added. Continue adding and testing until no precipitate is formed.

When the reaction is finally completed, allow the red precipitate to settle, pour off the clear liquid, and add fresh water. Repeat this process several times to wash the precipitate and remove any chemicals still in solution. Finally, the mercuric iodide can be filtered off in a paper funnel, dried, and placed in a bottle.

The precipitate can be scraped from the filter paper with a spatula made by rounding the edges of a strip of thin celluloid. An iron or metal knife should not be used since it is likely to combine with the chemical.

A simple yet mystifying experiment with heat can be performed with this red powder. To prepare the apparatus, fasten a cross or other figure cut from thin copper to a thin square of wood, using a cement of the type formed by dissolving scraps of celluloid in acetone. Then cement a sheet of white paper over the metal figure and wood base and apply a thin coat of red paint made by rubbing some of the red mercuric iodide you have made with cement or weak shellac.

When the paint has dried, hold the red square near the flame of an alcohol lamp or gas burner. Gradually, the portions of the paper not in contact with the metal will turn a vivid yellow, while the paper covering the cross will remain unchanged. This color-changing property of mercuric iodide also can be used by the amateur chemist to show the relative heat conductivity of metals. First, obtain wires of several different metals and coat them with shellac or some variety of paint. Then after they have dried, coat them with a paint made by mixing iodide with some base such as shellac.

When this final coat has dried, arrange the wires so they project into the flame of an alcohol or gas burner. The ends of the wires will soon become hot and the rapidity with which the heat is conducted along their lengths will be shown by the change in color from red to yellow.

A peculiar, heavy liquid can be made by dissolving mercuric iodide in potassium-iodide solution. The resulting fluid will have such a high specific gravity that stones, glass stoppers, and other heavy objects placed on its surface will float.

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Magic in Chemistry, Chemistry in Magic

Prove you’re a man to be reckoned with—and the only man who can make the gal in the photo (Fig. 1) blush. Prepare her for the test by painting her cheeks with phenolphthalein solution (from the drug store), and be sure the cheeks are slightly moist when you perform the trick. Ordinarily this solution is colorless, but when a finger (yours) moistened with household ammonia is brought near it, the reaction of the fumes with the solution causes it to turn pink. When the ammonia evaporates, the cheeks lose their color.

Whoosh goes this miniature volcano (Fig. 2) you’ve made from a small, pointed mound of ammonium dichromate (from chemical supply houses or large photographic stores). Just center the chemical on an asbestos pad (for heat protection) and light the tip with a match (Fig. 2A). When burning starts, turn out the room lights. In darkness the burning chemical looks like a miniature volcano, complete with sparks and lava-like material tumbling down its sides. When the lights flash on again your audience will see that a mountain of green powder (chromium sesquioxide) has replaced the little heap of ammonium dichromate.

Show your guests you can cool a bottle sans refrigerator and sans ice (Fig. 3). All you need are a large fruit juice can, a turkish bath towel, 1 lb. of common photographic hypo, water, and a couple of rubber bands (Fig. 3A.) Wrap and rubber band the folded towel around and under the can; then pour in 1 qt. of the coldest water you can get. Dissolve the hypo in it by rapid stirring, and insert the bottle. Bottle temp, should drop about 25°F. Later, after the moment of glory, bottle hypo solution for later photographic use.

Make your friend’s name turn into a caricature of himself (Fig. 4). Fortify yourself ahead of time by drawing the caricature in invisible ink made by mixing a pinch each of potassium iodide and cornstarch in a tablespoon of water, then heating for several minutes to dissolve the starch. Next, write the name with any ink that can be removed with ordinary ink remover, but dilute 1 part ink into 15 parts of water (Fig. 4A).

It’s a whiff of chlorine gas in the jar that does the trick. It bleaches the ordinary ink and releases the brown-colored free iodine in the invisible ink. To make the chlorine, cover the bottom of the jar with sodium hypochlorite bleach, such as Chlorox or Linco, and then add a little hydrochloric acid. Lay a square of stiff cardboard over the jar to confine the gas while it is being generated, and don’t inhale the gas.

Want to make an ink with which you can write a message that is invisible in humid weather but which turns blue when the weather is fair —and the fairer the bluer (Fig. 5)? Just dissolve a few crystals of cobalt chloride in a little water. Write the message with a blunt instrument, and apply plenty of solution (Fig. 5A). Invisible when even slightly moist, the writing appears whenever the weather is dry enough, or the paper is dried artificially under an electric light.

Show the folks how pure oxygen facilitates combustion (Fig. 6). Pour about an inch of hydrogen peroxide into a test tube, and add a few grains of powdered manganese dioxide which acts as a catalyst to liberate the oxygen. After a few seconds, insert into the upper part of the tube a wood splinter or a piece of cord with a spark at its tip. Instantly, the spark will burst into flame. A small piece of steel wool, heated red hot, will burn brightly if lowered into the tube.

You can bring a dozing audience (someone else’s, naturally, not yours) to quick attention by changing wood alcohol to that evil-smelling gas, formaldehyde (Fig. 7). Customarily used in water solution to preserve biological specimens and to harden photographic film, the gas is made commercially by oxidizing methyl, or wood alcohol by means of heat and a surface catalyst. You can do this, however, by immersing a test tube containing a teaspoonful of the alcohol in hot water. When alcohol warms, heat a 1/4-in. coil of bare copper wire in a gas flame and plunge it into the alcohol vapor. The copper causes the vapor to unite with oxygen from the air and from the film of oxide on itself, and the smell of wood alcohol changes into the pungent odor of formaldehyde.

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Generating SMOKE and STEAM for Amateur Theatricals

By Kenneth Malcolm

CURLING wisps of smoke rising in a fireplace, great smoke-gusts bursting in from an offstage forest fire, steam issuing from grotesque modernistic machinery or even from the spout of a humble teakettle—all the realistic steam and smoke effects which so often add to the interest of professional dramatic productions can be easily duplicated, at least on a moderate scale, by the amateur.

The apparatus to be described is a simplified version of that used in the professional theater, and costs not more than a dollar or two. The smoke—produced chemically by uniting ammonia gas with chlorine—is harmless and may be generated instantly wherever desired.

Obtain three 1-pt. fruit jars with screw caps, about 3 ft. of 3/8 in. outside diameter glass tubing, some sealing wax, and 6 or 7 ft. of 3/8 in. inside diameter rubber tubing. Except for the chemicals and perhaps a box or rack, these are all the materials necessary.

From the glass tubing cut three pieces 6 in. long, and three 3 in. long. This may be done with a tube cutter or simply by notching the tubing with a small triangular file and—with the tubing held in your two hands so that the notch is away from you—breaking at the notch.

Beneath the cap of each jar will be found an inset of white glass. As this cannot be drilled with ordinary drills, carefully break it out. Then, through the top of the caps, drill two 3/8-in. holes, as indicated in the drawing.

Inside of each cap now melt a thin layer of sealing wax by heating the cap over a spirit lamp or a low gas flame.

This is to prevent the chemicals from eating the metal. Next, seal one long and one short tube into each cap by applying a generous mound of wax on the underside of the cap. Allow the tubes to project about 2 in. above the top.

AT A druggist’s or a chemical supply A house, buy about 4 oz. each of concentrated ammonia, commercial hydrochloric acid, and glycerine. Pour the acid in one jar, ammonia in another, and glycerine in the third. To each add water until the solution reaches the middle of the jar. Do not put the chemicals into the jars until the covers are ready to be put into place, because the ammonia and acid give off very penetrating and disagreeable odors (perhaps even dangerous) and soon lose their strength.

When the caps are in place, the jars must be air-tight. If rubber rings or gaskets are lacking, a heavy coat of vaseline applied to the thread of the caps will make an adequate seal.

The three jars should be connected with two short lengths of rubber tubing as shown—the longer glass tube of the center jar being connected to the shorter tube of the first, and the shorter tube of the center jar being connected to the longer tube of the third. It is very important that they be arranged in correct order —first ammonia, second acid, and third glycerine. The glycerine solution acts as a sort of filter.

About 2 ft. of rubber tubing should be connected with the long tube of the first jar. A short length of glass tubing should be inserted in the other end for a mouthpiece. One end of the remaining length of rubber tubing should be pushed over the short tube of the third jar.

Blowing into the mouthpiece will cause white smoke to pour from the tube at the other end of the apparatus. It may be led by the tube wherever desired. Instead of coming from a single point, the smoke may be distributed. An attachment for this purpose may be made by taking a 2-ft. length of rubber tubing, corking one end, inserting a short glass tube in the other end for aid in connecting, and cutting a line of holes at intervals of 1-1/2 in.

To assist in transporting and storing the apparatus and prevent the jars from being overturned, it is advisable to construct a rack or a complete case.

Commercial smoke apparatus is generally operated by air that has been compressed in a tank by a hand pump. This arrangement may be imitated by amateur builders, if so desired, but lung power is much cheaper and less complicated.

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Safe Stunts with Fire FOR THE HOME CHEMIST

By Raymond B. Wailes

OF ALL home chemistry experiments, tests with combustibles offer the most in spectacular fun and harmless excitement. For even after some 60,000 years of use, fire still holds a mysterious fascination.

Although we are accustomed to kindling a fire with a match or some other small flame, a spark or a flame are by no means necessary to start some substances burning. Many materials ignite spontaneously when subjected to nothing more than a slight rise in temperature. Carbon disul-phide, a liquid often used as an ant exterminator, is one of these substances and for this reason presents a serious fire hazard if not handled carefully.

To demonstrate the mysterious action of spontaneous combustion, heat the bulb of a chemical thermometer (Centigrade) until the mercury climbs to the 250 mark. Then hold the bulb near the surface of a small amount of carbon disulphide contained in a shallow dish. Almost instantly, the liquid will burst into flame. No actual source of fire will have been used, yet the liquid will be ignited as readily as if it had been touched off with a match. Finally, extinguish the flame with a tin cover placed over the dish and repeat the experiment using an iron rod which is hot but not visibly hot. Again the liquid will burst into flame.

A similar and more common action takes place when oily rags, especially cloths used by painters, are stored in a warm place. Before long, slow oxidation of the oil takes place and the heat generated by the reaction raises the temperature to the kindling point of the cloth.

By using a simple collection of materials including two olive bottles or large test tubes, the home experimenter can illustrate spontaneous combustion graphically. Mold a small quantity of cotton waste or scrap thread into two loose balls, making each about the size of an egg, and stuff one into each of the olive bottles. Pour some ordinary machine oil into one container and a mixture of linseed oil and regular paint drier into the other. Then embed the bottles in a large container insulated with wads of cloth to prevent the conduction or radiation of heat. After placing a thermometer in each bottle, watch the two columns of mercury.

In a short time, the bottle containing the linseed-oil-soaked cotton will show a steady increase in temperature. This will be caused by the slow oxidation of the oil, the chemical reaction being aided by the addition of the drier. At this point, the mixture will be going through the early stages of spontaneous combustion and if allowed to continue until the temperature reached the kindling point, the oil-soaked cloths would burst into flame.

On the other hand, the thermometer in the second bottle will show little change in temperature. This can be explained by the fact that machine oil, a form of mineral oil, does not oxidize as readily as a vegetable oil and therefore does not produce large quantities of heat in the process.

As this experiment will show, spontaneous combustion is a collective reaction. At the start, the vegetable oil begins the cycle by combining with the oxygen in the air. This produces heat which in turn promotes a more rapid combination or oxidation and produces more heat. Naturally, it is only a matter of time before the temperature builds up to the point where the oil-soaked cloth takes fire.

Just as carbon disulphide is a dangerous liquid that bursts into flame with the slightest increase of temperature, another liquid carbon compound, carbon tetrachloride, is equally active as a fire extinguisher. When thrown or squirted on a fire, it cools the flame and blankets the base of the fire with a heavy cloud of gas that soon cuts off the necessary supply of air and oxygen.

Strange as it may seem, however, even carbon tetrachloride can be made to burn under the right conditions. When placed in contact with powdered zinc and sand, for instance, and ignited with a magnesium fuse, the combination will burn to give off large quantities of heat and larger quantities of smoke. It is this mixture which forms the basis of the smoke pots used by armies in war time and because of the large quantities of smoke given off, the experiment demonstrating it should be performed out-of-doors.

First select a small tin can and fill it three quarters full with a mixture of equal amounts of extremely fine sand, road dust, or fuller’s earth, and finely powdered zinc. Carbon tetrachloride then should be poured into the can until the mixture is thoroughly soaked. Any excess not absorbed by the powder can be poured off. Finally, make a small conical depression in the top of the damp mixture, fill it with powdered magnesium, and top it off with a strip of magnesium to act as a fuse.

Once the magnesium is ignited, it will prime or ignite the mixture and start the chemical action which produces the billows of dark gray smoke. What actually happens is this: The zinc in burning combines with the carbon tetrachloride to form zinc chloride and particles of black carbon. The zinc chloride then reacts with the moisture in the air to form white zinc oxide which, together with the particles of black carbon, make up the dark gray smoke. Incidentally, miniature smoke pots of this type are a valuable property for amateur theatricals.

By combining the two carbon liquids used in the experiments so far, the home chemist can produce an almost heatless flame. Mix three parts by volume of carbon disulphide with eight parts of carbon tetrachloride and light the resulting solution. The temperature of the flame will be so low that a piece of newspaper, generally considered as being particularly inflammable, held in it will not burn. It may char, depending on conditions, but it will not burst into flame.

In experiments to determine the combustibility of inflammable materials, fire department officials have found that heavy vapors often flow along surfaces for many feet to be ignited by some distant flame. A simple home-laboratory experiment that shows the heaviness and flowing qualities of gasoline vapor can be performed by pouring a half teaspoonful of liquid gasoline into a beaker. In a short time, the beaker will be filled with a heavy vapor of gasoline. If it is then carefully tipped, the heavy vapors can be poured into a second beaker. Finally, being careful to keep away from the first beaker containing the liquid gasoline, invert the second beaker containing the vapor over a lighted candle or match. The gasoline vapor will literally pour out of the beaker and take fire with a lazy, floating flame.

From your experiments with spontaneous combustion, it must not be assumed that mineral oils do not burn. Nothing could be further from the truth. It is the mineral type of oil that is used as a fuel for heating. In fact, two specifications on which mineral oils are graded are their flash point, the temperature at which their vapors when mixed with air will explode, and their fire point, the temperature at which they will take fire and burn.

As a practical experiment in combustibles, the home chemist should obtain several lubricating oils and test them for flash and fire points. The equipment required is particularly simple and the results obtained are comparatively accurate. First place a shallow evaporating dish or the friction top of a tin baking soda can on the support of your laboratory stand as shown in the photographs, arranging it so that it is held over the tip of your regular gas burner. Then rig a thermometer vertically and allow its bulb to dip into a small sample of the oil placed in the dish. Also arrange a small pilot light by fitting the spout of a small oil can to a piece of rubber tubing leading to your gas supply. When you have lighted both the gas burner and the pilot light you are ready to proceed. The burner should heat the oil slowly and the pilot flame should be no larger than the head of a large match.

Stir the oil slowly and whisk the tiny pilot flame across the surface of the oil at ten-second intervals. As the temperature increases, watch the oil carefully. Sooner or later, whisking the pilot across the oil will cause the oil vapor to ignite with a momentary flash. The reading of the thermometer then will be the flash point of the oil.

Continue the heating and resume the whisking process with the pilot light. The moment that the oil catches fire and continues to burn, read the thermometer again. This second reading will be the firepoint.

The differentiation between the two points is easily recognized. At the flash point only the vapor given off by the oil takes fire and it will stop burning the instant the pilot flame is removed from the vicinity of the oil. At the fire point, on the other hand, the oil will continue to burn of its own accord until it is extinguished by smothering it with a sheet of tin.

To prove that oils differ in characteristics, the home experimenter should test various types and grades of oils for their flash and fire points. If a test is repeated, a fresh sample of the oil should, of course, be used. When accurate results are desired, these flash and fire tests should be performed in the absence of drafts. A direct draft on the test apparatus will tend to cool the surface of the oil and may raise the flash and fire points. Also, it is well to use the same size pilot flame in each of a series of comparative tests.

Just as chemicals can be used to promote combustion, they also can be used to prevent it. Both wood and cloth can be fire-proofed through the use of simple chemicals. To demonstrate this, dissolve about five teaspoonfuls of ammonium phosphate in two or three tea-spoonfuls of water. Immerse a small square of cloth in the liquid, allow it to soak for a minute or so, and then hang it up to dry. Finally, try to ignite the cloth with a match. Although it may scorch or char, it will not burst into flame as readily as cloth of the untreated variety. If ammonium phosphate is not available, ordinary alum dissolved in water can be used.

Wood can be fire-proofed in a similar way by using sodium silicate (water glass). As an experiment, paint a narrow band of the chemical around a large kitchen match stick about an eighth inch or so in back of the head. When the water glass has dried, strike the match. It will burn brightly until the flame reaches the treated wood. Trick matches that will go out as soon as they are lighted can be made in this way by the home chemist whose friends are continually borrowing a light.

Dynamite Made from Corn
Production of a highly explosive dynamite from corn is one of the latest developments of the chemical laboratory. It is the result of the recent discovery at the University of Iowa of an inexpensive method of extracting inositol, a sugarlike substance, from corn. Inositol is a non-explosive form of alcohol but when nitrated becomes a powerful solid explosive. It can be produced from the waste by-products of the manufacture of cornstarch.

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HOW CHEMISTRY CREATES A PHOTOGRAPH

What goes on in the emulsion that coats film is shown by simple test-tube experiments.

By TRACY DIERS

THE film in your camera is thinly coated with one of the most unstable chemicals known to man. Silver bromide is its name, and from the moment of its birth it is kept in a cradle of darkness until in your camera a swift shaft of light seeks it out. The intricate and far-reaching changes brought to silver bromide by that flash of light are in part still secrets of nature. Much of what happens in your camera and in the darkroom is known, however, and can be shown at home with a few chemicals in a test tube.

Dissolve a crystal of silver nitrate in a test tube half full of water, and dissolve a crystal of potassium bromide in another. If you place both test tubes in direct sunlight, nothing will happen. Mix the two together, and immediately a yellow precipitate of silver bromide falls to the bottom of the tube. Expose this to strong sunlight or to a photoflood bulb, and it will immediately turn purple. This purplish-colored material is chemically pure silver, freed from the bromine it was formerly attached to. You might say that the reaction was this: Silver bromide plus light equals silver plus bromine. The latter goes off as a gas.

Since it isn’t practicable to coat films or paper with a lumpy precipitate, we had to invent some method of spreading the chemical evenly and thinly over the surface we wished to sensitize. The answer was to suspend the silver bromide in an emulsion. Dissolve, with gentle heating, 20 grains of household gelatin in a test tube half full of water. Add a pinch of potassium bromide and then pour in a dilute silver nitrate solution. Silver bromide is again formed, but this time it remains evenly suspended.

In photography the action of the silver bromide is speeded and intensified by the use of a developer. Stock developing solutions consist of a developing agent such as pyro or hydroquinone, an alkali such as sodium carbonate or borax, a little potassium bromide as a restrainer, and some sodium sulphite as a preservative. To show how a developer works, put several drops of any stock developer solution in a flask of silver bromide emulsion that has been exposed only briefly, and not sufficiently to turn dark. The emulsion now turns black in a few seconds—much faster than when light does all the work of separating the silver.

On exposed film, you would now see an image consisting of dark areas of metallic silver and light areas of unexposed silver bromide that the developer has not affected. A fixing bath is next needed to keep these light areas from darkening upon subsequent exposure to light, for if they did the entire negative would turn black and the image would be lost. The fixing bath is sodium thipsulphate, the hypo familiar to every photographer.

You can see how it works by dissolving a few small crystals of sodium thiosulphate in water in a test tube and pouring the solution into half a test tube of unexposed silver bromide emulsion. Immediately the yellow suspension disappears, and the tube con-tains a clear, transparent liquid.

What happens is this. Sodium thiosul-phate reacts with silver bromide to form silver thiosulphate, a clear solution; then, if there is plenty of fresh hypo present, the reaction continues, forming complex salts that are dissolved by the remaining hypo.

A simple taste test illustrates the point. Fresh hypo has a salty, bitter taste, and silver sodium thiosulphate is almost as sweet as sugar. Dip a sheet of ordinary writing paper in a solution of 1 oz. of hypo and 4 oz. of water, rinse it in water for a moment, and taste the salty hypo remaining. Then dip a small piece of enlarging paper in the hypo for a few seconds, rinse it quickly, and taste the sweet emulsion side. Return the enlarging paper to the hypo for a few minutes longer, remove and rinse it, and you taste the salty hypo again. Expert photographers avoid stale hypo so that the reaction, or fixing, will be complete and no silver sodium thiosulphate will remain.

Every photographer has taken negatives that were too light or too dense to afford good prints. Too light an image may be due to underexposure in the camera, or underdevelopment in the darkroom. A negative of this kind can sometimes be salvaged by intensification. Too dark a negative can be improved by reduction.

To show the action of a common reducer, first prepare half a test tube of silver bromide emulsion, expose it, develop with a few drops of stock developer, and add four or five drops of glacial acetic acid (the photographer’s short stop) to halt development. Then add 6 cc. of hypo solution. To the dark emulsion you now have, slowly add a reducing solution consisting of 25 grains of potassium ferricyanide in 1 oz. of water. Watch the dark silver particles lighten as they are changed into silver ferrocyanide and in that form are dissolved by the hypo. This is the well-known Farmer’s reducer.

A common method of intensification is to combine the silver deposits of the negative with some denser metal such as chromium or mercury. To show how a typical intensi-fier works, first make up the usual emulsion, expose, and develop it as before. Add a few drops of glacial acetic acid short stop. Oddly enough, the first step in intensification is bleaching. Add 20 grains of mercuric chloride to 1 oz. of water and pour in enough of this solution to bleach out completely the black metallic silver, which in the process is converted into white silver chloride.

On a negative, the image seems lost at this step. It is brought back by redevelopment. To your bleached test-tube emulsion, add a little ordinary ammonia water. The emulsion becomes even blacker than it was before; it has been intensified. The new blacker chemical is a very complex salt, silver mercury ammonium chloride.

The process of toning photographs, in which the black areas are converted to sepia or some other color, also may involve bleaching and redevelopment. Prepare the usual small quantity of emulsion and develop it after exposure. Add the acetic acid short stop. Then prepare a small flask of bleaching solution consisting of 110 grains of potassium ferricyanide, 110 grains of potassium bromide, and 2 oz. of water. Also dissolve separately in a test tube 50 grains of sodium sulphide in 2 oz. of water. This is the redeveloper.

Pour some bleaching solution into the developed emulsion. The black silver particles are quickly bleached out into a yellowish substance. The silver first becomes silver ferrocyanide and then silver bromide.

Now pour in some of the redeveloper and at once the emulsion turns brown—in a photo, sepia toning would have resulted. The sodium sulphide in the redeveloper has changed the silver bromide to silver sulphide, the compound that appears as tarnish on silverware, and this is the sepia tone.

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Fun with Black Light for Home Chemists

By RAYMOND B. WAILES

CHEMICALS that glow with magic colors in the dark, under invisible illumination with “black light,” have been applied to theatrical costumes and decorations with spectacular effect. Your own home laboratory can be the stage for equally striking experiments with these substances, which possess the curious property known as fluorescence. Also, you can prepare other substances that shine in the dark through the phenomenon called phosphorescence—which is distinguished from fluorescence by the fact that phosphorescent chemicals continue to glow for some time after removal from the light that excites them.

To make a phosphorescent preparation, simply break off the point of an indelible lead pencil and pulverize it thoroughly with a pinch of powdered boric acid (boracic acid), which you may find awaiting use as an eye lotion in the family medicine chest. Heat the mixture in a tin-can lid with the aid of a Bunsen burner, an alcohol lamp, or the kitchen range. The pulverized mass will twist and squirm, finally becoming shiny and glasslike. When the lid and its contents have cooled, hold the product close to an electric light. Then turn off the light, and if the room is dark you will see the preparation glow vividly. Its light will gradually dim and disappear.

You can obtain phosphorescent light of different colors by heating boric acid powder with household dyes, or with a bit of fluorescein, and varying the heating so that the boric acid is only partly fused.

For black-light experiments with fluorescent chemicals, you will need a source of ultra-violet rays. The handiest is an argon lamp bulb, which you can obtain from any large electrical store for about half a dollar. Screwed into any lamp socket providing the ordinary house current of 110 volts, it emits a quantity of invisible ultra-violet light, besides some visible violet light. The lamp bulb may be fitted with a cup-shaped shade, to concentrate its rays upon your chemical bench, and to screen the visible light from your eyes so that it will not interfere with your observations.

Many other sources of ultra-violet light are available for your experiments. Even direct sunlight will enable you to observe the fluorescence of a few chemicals, though in most cases the brightness of the visible part of the sunshine masks their fluorescent glow. A photo-flood lamp will provide you with black light if you inclose it in a well-ventilated box, lined with asbestos paper, and let light emerge only through a window of special ultraviolet filter glass. This purple-black glass, opaque to practically all but ultraviolet rays, may be obtained from dealers in scientific apparatus and chemicals. If your home laboratory does not have 110-volt current on tap, you can produce ultra-violet light of sufficient intensity for many experiments with an iron spark gap, connected to a battery-operated spark coil.

To start your black-light experiments, you need look no farther than your own household for fluorescent materials. Ordinary machine oil, exposed in a darkened room to your source of ultra-violet rays, will glow with blue light. So will petroleum jelly from the medicine cabinet. Calomel, or mercurous chloride, in solid form, will fluoresce under like conditions with a peculiar orange-gold color. Pyrogallic acid, familiar as a photographic developer, also exhibits fluorescence upon exposure to black light. However, more spectacular tricks of chemical magic can be performed with anthracene, a coal-tar product that will glow vividly with a yellowish-green color in ultra-violet light. Dissolved in benzine (benzol), gasoline, or xylene, it makes an invisible ink with which you can write secret messages on paper. When the writing dries, it cannot be seen in ordinary light. Hold it before ultra-violet rays in a dark room, however, and it will glow in letters like fire.

HERE is an entertaining variation of the stunt, to perform before your friends. Darken a room in which you have set up a white sheet of cardboard as a “blackboard.” Dip a stick of chalk momentarily in a xylene or benzene solution of anthracene, and let the chalk dry. Then write with the prepared chalk upon the white “blackboard,” which must be illuminated meanwhile by black light. The message will be traced in fiery, glowing letters, with mystifying effect that you can enhance by clever showmanship.

Small dolls, whose clothing has been dipped in anthracene solution and dried, will also fit well into any program of chemical magic. They glow weirdly when you make them dance under black light. For a two-color effect, impregnate only part of the clothing with the anthracene solution, and the rest with a solution of sodium salicylate in water. When dry, the latter fluoresces under black light with a pale-blue color.

Even in daylight, a solution of fluorescein in water will visibly fluoresce. A mere speck of the chemical, no larger than a pinhead, will suffice to make a quart of solution. To help dissolve the fluorescein, add a gram or quarter teaspoonful of sodium hydroxide (lye) to the water. The fluorescence of the solution becomes far more vivid under black light. Use the liquid to fill a round-bottomed flask, or a burned-out electric light bulb with the base removed, and you will have a “crystal-gazing globe” that will magically glow in the dark when illuminated by ultra-violet rays.

To make another preparation whose fluorescence can be seen in sunlight, add a solution of ammonium fluoride to a solution of uranium nitrate until no more precipitate forms. Filter off the precipitate, which is a uranium compound called uranyl ammonium fluoride. Wash it with alcohol while it is still on the filter paper, let it dry, and then promptly stopper it in a bottle. This solid substance, like the fluorescein solution, fluoresces much more brilliantly in the dark, under an argon lamp.

Increasing the power of your black-light source enhances the luminous response of fluorescent substances, and you can do this by mounting several argon lamps close together on a board. Wire them in parallel, or so that each receives its supply of current independently of the rest. As the argon bulbs consume only two watts apiece, a dozen or two of them would use no more current than a standard twenty-five or fifty-watt bulb.

To observe fluorescent effects at their best, another expedient is to place the substance under observation in a homemade viewing box, which is used in conjunction with your source of black light. A sliding cover on the box will facilitate inserting and removing your specimens. Fit one end of the viewing chamber with a window of ultra-violet filter glass, of the special type previously mentioned, which is available in several sizes. While the window admits

black light to the specimens in the box, you can view them through a hole cut in the top. Any leakage of visible light into the box may be prevented by surrounding the peep hole with a masklike shield that fits your nose and forehead.

Such a viewing box will enable you to observe the fluorescence of many more substances. Common buttons glow with a blue color, which cannot be seen by the unscreened argon lamp because its own visible light masks the faint blue fluorescence. Uranium-tinted glassware, yellow-green in color and found in many households, will also fluoresce in the viewing box.

DYES in particular are apt to exhibit fluorescent properties under the rays. An alcoholic solution of rhodamine will glow red. A solution of uranin in water will shine with yellow light. Many minerals, too, show fluorescence. Some substances—including the preparation of boric acid and a dye, described in an earlier paragraph—are both fluorescent and phosphorescent.

Camera fans know that light rich in ultraviolet rays will act much more strongly than visible light alone upon a photographic film. You can easily make a light-sensitive solution that will respond to different kinds of light in similar fashion. Dissolve about five grams, or a teaspoonful, of ammonium oxalate in 100 cubic centimeters (about three and a half fluid ounces) of water. Make another solution of five grams, or a half teaspoonful, of mercuric chloride in the same amount of water. Mix the two solutions and filter. Place equal amounts of the filtered mixture of liquids in each of three test tubes.

Expose one tube overnight to light from an argon bulb, and another tube, for the same length of time, to the light of an ordinary electric-lamp bulb. A screen may be placed between the bulbs so that each tube is illuminated by only one type of light. Meanwhile keep the third tube in a dark place to serve as a “blank” or reference standard.

After twelve or more hours, you will find a considerable deposit of white crystals in the tube exposed to the argon lamp. The test tube lighted by an ordinary bulb will contain only a third or fourth of this amount, and none at all will be found in the tube kept in the dark. Ultra-violet rays, most abundant in the first case, promoted a reaction between the ammonium oxalate and mercuric chloride to produce the white precipitate, which consists of mercurous chloride.

Sunburn is another sort of chemical action produced by black light. “Sun-tan ointments,” popular at this time of year, prevent sunburn by shielding the skin from an overdose of the ultra-violet rays in sunlight. To show how they filter out the rays, heat about five cubic centimeters (a teaspoonful and a half) of liquid paraffin oil or petroleum jelly. Dissolve two or three salol (phenol salicylate) tablets in the hot oil or jelly. Mix the product with about ten grams of lanolin to form an ointment.

Now smear a bit of this preparation on a plate of glass, just as you would apply such a cream to the skin. Hold the coated plate between an argon lamp and some fluorescing substance. The fluorescence will stop, showing that the ultra-violet rays have been absorbed.

A much better “sun-tan” cream for practical use can also be made easily in your home laboratory. Melt twelve and a half grams (two and a half teaspoonfuls) of lanolin and thirty-seven and a half grams (seven and a half teaspoonfuls) of petroleum jelly in an evaporating dish. Add 100 cubic centimeters or about three and a half fluid ounces of rose water, a little at a time, stirring meanwhile. Then stir in fifteen grams (about three teaspoonfuls) of calamine powder, which is a zinc compound. A drug store can supply you with these ingredients.

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How Evaporation Steals Heat

EVERY time a liquid evaporates into a gas, it snatches a definite amount of heat from its container and surrounding air, cooling both below their original temperatures. This law of physical chemistry has long been useful to the human race as a means of cooling foods or drinks. Primitive man found that water placed in unglazed earthenware vessels would seep through the pores, evaporate, and cool the water remaining inside. Campers and country dwellers still cool water in this way.

Today, all our mechanical refrigerators, electric and gas alike, harness the cooling effect of evaporation. Alternately compressed into a liquid and allowed to expand into a gas, the refrigerant absorbs heat during each evaporation cycle.

The following experiments demonstrate how evaporation lowers temperatures.

Ether Freezes Water. Since some liquids vaporize at a lower temperature than water, they produce a correspondingly greater degree of cold. Ether is such a liquid.

This may be shown by filling a test tube about a third full of ether. After equipping the tube with a two-hole stopper, insert a length of bent glass tubing until one leg almost reaches the bottom. Into the other hole thrust a short delivery tube fitted with rubber tubing long enough to carry the ether vapor out an open window.

Now moisten the outside of the test tube with water and blow steadily into the inlet tube. The ether will quickly evaporate and ice will form on the test tube.

Insecticide Produces Frost.

Some liquids evaporate at temperatures even lower than ether. An example is di-chlorodifluoromethane, a liquified gas used for refrigeration and in insecticide bombs.

With such a bomb, you can easily see what spontaneous evaporation of this chemical will do. Hold the end of a glass tube, bent into a coil or zigzag shape, to the nozzle while the bomb is in action. (Be sure you hold the tubing with a pair of tongs or a spring clothespin, for any liquid accidentally sprayed on your fingers might freeze them). Evaporating and expanding as it escapes through the tube, the chemical quickly causes frost to form on the outside.

Liquefying a Gas.

In an electric refrigerator, a motor-driven compressor liquefies a refrigerant gas, such as sulphur dioxide, by squeezing it powerfully. For purposes of experimentation, you can liquefy this gas by subjecting it to intense cold.

To generate the gas, put 15 grams of sodium bisulphite into a flask arranged, as at the right, with dilute hydrochloric acid (1 part acid to 2 parts water) in a dropping funnel above it. To dry the gas, fill the horizontal tube with lumps of calcium chloride, packing cotton loosely in each end before inserting stoppers with entry and exit tubes. Lead the latter into a large test tube supported in a freezing mix- ‘ ture made by filling the beaker half full of acetone and adding dry ice slowly. (Caution: keep your fingers away from the mixture.) Now open the clamp, allowing acid to fall drop by drop, and apply gentle heat. When a quantity of liquid sulphur dioxide has collected, lift the test tube from the freezing mixture in a well-ventilated room. The liquid will quickly evaporate, forming frost as at the right above.

Gas Refrigerator Uses Ammonia. Altered as at the right, the apparatus will generate and liquefy ammonia, the refrigerant found in gas refrigerators. Put in the flask a dry mixture of 15 grams ammonium chloride and 30 grams of calcium hydroxide, insert a one-hole stopper, and in the drying tube place a loose layer of calcium oxide. Omit the hydrochloric acid; gentle heat will start the reaction. When some liquid ammonia has collected, remove the test tube from the freezing mixture. Again, frost will form.

Chemicals Produce Cold. A few ounces of ammonium nitrate crystals will show that cold also may be produced chemically—merely by dissolving certain substances in water. A commercial device for cooling bottled beverages on picnics makes use of this principle.

To demonstrate how it works, put a few drops of water on a block of wood or a cork coaster and set a thin-bottomed glass on the water. Then put equal amounts of cold water and ammonium nitrate into the glass and stir rapidly. The coaster will freeze to the bottom of the glass. In dissolving, the chemical absorbs heat from its surroundings.

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Home Science Stunts with Soap

by Prof. Victor Lewitus

Make a strong soap solution by mixing shaving soap and water. After taking a puff on a cigarette, blow the smoke through a bubble pipe to make a soap bubble. The inside of the bubble then will contain the white smoke, and when it breaks, it does so with a puff, furnishing a very striking experiment. A clay or corncob pipe will be more suitable than the briar variety, inasmuch as the soap mixture probably will make the pipe unsuited for further smoking.

Into a glass of water place two tea-spoonsful of bicarbonate of soda (baking soda). After it has dissolved, add one or two teaspoonsful of soap chips, which should be dissolved before proceeding with the experiment. Finally add some ordinary vinegar and notice the thick, voluminous foam that forms. This type of mixture may be used to extinguish a fire, and is very successful because of its tenacious character. If a sufficient quantity is available it will act as a blanket covering the burning area.

Into a narrow bottle place a few crystals of copper sulphate or “bluestone.” Cover these with a few inches of water. Then add several pinches of soap chips and shake the mixture for an instant. It will be found that a beautiful emerald-green solidified mass develops. This is caused by the soap solution reacting with the copper sulphate, which has an acid reaction.

Below—Fill three drinking glasses of the same size half full of water. Use ordinary tap water in the first; the same water for the second into which has been placed some salt; and distilled or rain water for the third. Place small equal quantities of soap into each glass of water and shake each for the same length of time. Notice that a thick lather will form quickly in the distilled or rain water while practically no suds will appear in the salt water. The tap water probably will have suds but not as many as in the rain water. This experiment will show the softness of tap water as compared to the other extremes.

Below—Into a small glass receptacle place two ounces of water. Then pour into the water a half an ounce of mineral oil. Shake the combination and note that within a short time after the shaking, the oil and water will separate. Add one ounce of soap chips to this solution, and this time after shaking, the oil and water will not separate. This shows that oil and water will mix under the proper conditions.

Place two quarts of water in each of two pans and add some salt to one of them. Ask someone to wash his hands in the water to which salt has been added. He will find that try as he, will, no lather will form, while if you.

wash in the plain water, an abundance of lather will be obtained.

This burn-aid soap will be found to be a very useful preparation to have on hand in the home at all times. Children particularly are subject to many minor burns, and because of the pink color of the material, they probably will have no objections to having it applied when necessary. Into a glass vessel pour three ounces of lime water (obtainable at the drug store). Add to this, with constant stirring, linseed oil until you have incorporated three ounces. Mix thoroughly by shaking, in a six- or eight-ounce bottle. During the shaking, the material will turn to a milky-appearing substance. To this add about ten drops of mercurochrome solution (obtainable at the drug store). This last addition

will turn the solution a pink color. These simple experiments will be found interesting and simple for the amateur chemist, and give to the experimenter a broader knowledge of soap—one of the most common and useful things in life. While in continual use by every person, few understand its workings other than for washing purposes.

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Three Magic Metals

Producing Cold With Electricity and A “Quicksilver Heart” That Beats Are Only Two of the Amazing Tests You Can Perform Easily With Simple Substances

By Raymond B. Wailes

YOU are accustomed to seeing an electric element in a toaster or radiant heater grow red-hot when current passes through it—but did you know that when electricity flows through joints of certain metals, it produces a cooling effect? Have you ever made a drop of murcury behave as if it were alive or prepared a pair of magical alloys that are solids when separate, and a liquid when mixed?

These are a few of the fascinating experiments that you can perform with metals, using three in particular that you may not have employed before in your home laboratory—mercury, antimony, and bismuth.

You may already have discovered that many of the metals needed in your experiments can be found in your home. They should be prized, bottled, and labeled like any other chemicals. The shell of an old dry cell will furnish you with zinc, and worn-out pots and pans with aluminum while the mesh scrapers used to clean them supply copper in handy form. Likewise, mercury, which is expensive to buy, may be available right at hand. It can be salvaged from a broken or discarded thermometer, provided, of course, the instrument is not one of the kind that uses alcohol colored with a dye. The quantity of mercury will be small, but it will be ample for a number of tests, and it is easy to clean and use over and over again.

Only a drop of the liquid metal is required for a striking demonstration known as the “mercury-heart” experiment. Place the mercury in a small, shallow vessel— a glass caster well from the ten-cent store

will do nicely—and cover the drop with a dilute solution of sulphuric acid. One part of strong acid to six parts of water makes a suitable solution, which should be colored faintly purple by the addition of a drop of potassium permanganate solution.

Now thrust a sewing needle into the solution from the side, jabbing the point into the drop of mercury, and you will receive a surprise. The drop will hump itself up, as if alive, and retreat from the needle. No sooner has it done so, however, than it flattens out again, repeating the pulsation each time it comes in contact with the needle point.

The same materials will serve for a “tidal-wave” experiment. Only enough of the colored acid should be used, this time, to encircle the drop of mercury, leaving its upper sur-face uncovered. Hold the needle vertically and touch it to the surface of the mercury drop. Then draw the needle sideways until it just meets the solution. Immediately the mercury gathers itself up about the needle, while the solution backs away. The mercury then relaxes and flattens out as before. The pulsation will continue for a considerable time. Changes in the surface tension of the mercury, caused by the electrical action of the metals and the acid, account for the remarkable behavior of the drop of metal in these two experiments. Alloys of mercury with other metals are called amalgams and one of the most curious of these is a double amalgam known as Mackenzie’s alloy. To make it, grind together in a mortar one part of mercury and two parts of bismuth metal, by weight, until a homogenous product is obtained. This is a bismuth amalgam. Make a lead amalgam in the same way, using three parts of mercury and four parts of lead, again measured by weight. The bismuth amalgam and the lead amalgam are both solids at ordinary temperatures. Place some of each in your palm and rub them together. Presto! They are transformed into a liquid alloy that you can pour freely from hand to hand.

An amalgam of magnesium metal and mercury may be made by rubbing the two together in a mortar with a pestle. Considerable heat is liberated as the metals unite, if the magnesium is in powdered form. The magnesium alloy that results is notable for its ability to decompose water, releasing hydrogen gas. Heating the water will make the effect more marked. The magnesium interacts with the water to form magnesium hydroxide per oxide, while the mercury can usually be reclaimed at the end of the experiment.

A few hints on working with mercury in the laboratory will not be amiss. Because of its propensity for forming amalgams with other metals, a wise precaution is to remove any valuable rings from the fingers before handling it. Gold and silver, by contact with mercury, quickly acquire a silvery coating of amalgam. If a piece of jewelry is made entirely of gold, however, and contains no stone or part that might be damaged by heat, the mercury may be volatilized and driven off by heating the article gently.

Spilled mercury is elusive, but may be picked up with a thin scoop of stiff paper, if the drop is first wetted. Mercury can be cleaned by filtering it through chamois skin, applying pressure with the fingers if the quantity is small. Shaking mercury with weak nitric acid (about an eight-percent solution) is another way of purifying it. This tends to dissolve any foreign metals that may be present as impurities. Two of the most interesting metals for home experiments are antimony and bismuth—a pair so alike in their properties that they might be called chemical brothers. You might search your house high and low without finding either of these elements for they appear in everyday life only in the form of a few compounds. Ask for bismuth at a drug store, for example, and you are likely to get the subnitrate or subcarbonate, which are used medicinally for certain stomach disorders. Antimony is contained in potassium antimony tartrate, more familiarly known as tartar emetic. The most striking experiments require the metals themselves, however, and these may be obtained from dealers in chemicals, usually in the form of lumps and powder mixed.

By passing an electric current through a couple or junction of antimony and bismuth, you can produce either a heating or a cooling effect at will. This strange phenomenon is known as the Peltier effect, after the French scientist who discovered it in 1834. To demonstrate it, you will need a fair-size lump of each metal. The pieces should be attached with bare copper wire to a pair of long, metal knitting needles passing through the cork of a flask or bottle and serving as supports. Adjust them so that the lumps of antimony and bismuth are in contact with each other. If the lumps available are not large enough, you can form pieces of the desired size by melting the powdered metals and casting them in a paper mold. Pass a glass tube through a third hole in the cork and place the flask tightly on the cork. The apparatus should be arranged so that the glass tube dips into a beaker or small wine glass filled with colored water. The whole arrangement will act as a thermometer.

Now connect two or three dry cells, as shown in the diagram, and attach wires from them to the knitting needles, thus closing the circuit and setting up a flow of electric current through the antimony-bismuth couple you have made. When the direction of the current is from the antimony to the bismuth, the couple will be heated, the air in the flask will expand, and you will see bubbles of air emitted from the tube that dips into the beaker. But if the current is made to flow in the opposite direction, the junction of the metals is chilled, and the air in the flask contracts, as evidenced by the water rising in the tube. You are observing a remarkable phenomenon, direct cooling by an electric current which, if it could be applied practically to a device like an electric refrigerator, would eliminate all moving parts and produce a silent, efficient apparatus that would never wear out!

Once brought to its ignition point, metallic antimony burns in the air almost as readily as paper. You can show this by heating a pellet of antimony upon a charcoal block until it begins to burn, and then tossing it upon the inverted lid of a paper box. Rolling and bouncing from side to side, it continues to burn in the air as the coating of oxide that would smother it is continually knocked off. Its temperature becomes high enough to scorch the paper of the box lid, leaving tracks that record the movement of the large globule and the smaller ones that break off from it. If the antimony is simply melted and dropped upon the box lid, it will smolder for about a minute, emitting white smoke and leaving a trail of the oxide behind it. White fumes of the oxide are also produced if antimony is heated upon a charcoal block with a pointed flame, such as that of a blow-pipe. Finely powdered antimony or bismuth burns rapidly if it is thrown into a Bunsen flame, and a pinch of antimony tossed upon molten potassium nitrate that you have heated in a tin-can lid or a crucible takes fire with a shower of sparks.

Antimony and bismuth, like iron, decompose water when heated red-hot. They react chemically when heated with sulphur or when dropped into a vessel of chlorine gas, producing heat and sometimes light. Sulphides and chlorides of the metals are formed as a result.

When you try to dissolve antimony chloride or bismuth chloride in water, you will notice that a white precipitate is always produced unless a drop or two of hydrochloric acid is also added. The acid dissolves the white precipitate, which is a basic salt of the metal being used and is known as an oxychloride. Nearly all antimony and bismuth chemicals produce precipitates of this kind unless enough acid is present to keep the oxychloride in solution.

YOU can apply this fact in an effective little chemical trick. Make a solution of the chloride or nitrate of either antimony or bismuth, using just enough acid to give a clear, waterlike liquid. Hold half a glassful of this solution under a faucet, and add enough water to fill the glass. The contents change from “water” to “milk” as the white precipitate of oxychloride appears, due to the reduced concentration of acid. To one who is not in on the secret, it looks as if you drew a glass of milk from the water faucet.

A peculiarity of bismuth oxychloride is that it is somewhat photosensitive, turning gray in sunlight. A chemical difference between the oxychlorides of antimony and bismuth can also be shown. Adding a few crystals of tartaric acid will cause a precipitate of antimony oxychloride to redissolve, while a precipitate of bismuth oxychloride is unaffected. This may be used as a test to distinguish between antimony and bismuth.

Place a strip of iron or zinc in a solution of an antimony or bismuth salt, to which a little acid has been added to prevent formation of the oxychloride, and metallic antimony or bismuth will be deposited upon the foreign metal. A strip of copper placed in a solution of antimony chloride, which has been acidified with strong hydrochloric acid, becomes covered with a curious violet-colored mass. This is known as Reinsch’s test for antimony.

Antimony may also be detected by a method closely resembling Marsh’s test for arsenic, described in an earlier issue (P.S.M., Dec. ‘34, p. 56). Hydrogen gas, generated in a flask from zinc and sulphuric acid, is led through a drying tube containing anhydrous calcium chloride and then through a horizontal piece of glass tubing about eight inches long, which is gently heated with a small flame during the experiment. After hydrogen has been generated for five or ten minutes to clear the apparatus of air, admit the solution to be tested for antimony to the generating flask, through a thistle tube or a separatory funnel.

ANY antimony that is present in the chemical will combine with a part of the hydrogen, forming a gaseous compound of antimony and hydrogen called stibine. When the stibine gas reaches the heated outlet tube, it will be decomposed, and a metallic mirror of antimony will be deposited upon the inner surface of the glass. If the gas issuing from the outlet tube is ignited, a cold porcelain dish held in its path will also receive a deposit of antimony. The zinc used in this experiment, if not of high purity, may contain some arsenic, and this will produce a deposit of similar appearance. The two are readily distinguished, however, by a simple test. The antimony stain will not dissolve in a solution of bleaching powder, while the arsenic deposit will. Try the test first upon a compound that you know contains antimony, for practice. If the gases are not being burned at the end of the outlet tube, the room should be kept well ventilated while you are performing this experiment.

Stibine, formed in the test just described by the combination of antimony and hydrogen, is known as a hydride. The gas has been detected issuing from storage batteries while they are being charged, as a result of the combination of hydrogen liberated during the charging process with antimony present in the lead plates of the batteries.

Melt some bismuth in a crucible, and you can obtain a sample of the crystals produced by this metal. Let the molten bismuth cool until a hard crust forms on the surface. Then break the crust with an iron rod and pour out the remaining liquid. The crystals will be found adhering to and covering the inside walls of the crucible.

Antimony and bismuth form a variety of peculiar and interesting alloys with other metals. One, containing antimony and copper, has a beautiful purple color! Another, consisting of bismuth, tin, and lead, will enable you to “silver” the inside of a flask or a bulb blown from glass tubing.

MELT four parts of lead, by weight, in a crucible; then add six parts of tin. When all is molten, stir in ten parts of bismuth. A piece of glassware in which the resulting alloy is placed, gently heated, and swirled about, will receive a silvery inner coating and will make an interesting exhibit to add to your chemical museum.

Though their names may be unfamiliar to the layman, antimony and bismuth have a number of commercial uses. Type metal is made of about seventy-five parts of lead, twenty of antimony, and five of tin; unlike most substances, it expands upon solidifying, thus giving a clear impression of the type mold. Babbitt metal and Britannia metal are other alloys containing antimony. Bismuth alloys, which melt at remarkably low temperatures, are employed in automatic sprinkler systems for fire-fighting. Salts of bismuth are employed medically, as already noted, and are also used in the manufacture of hand lotions, cosmetics, artificial pearls, porcelain enamels, and certain paints.

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Girl Chemist

Jackie Bates works harder, has lonelier life than most of her ex-classmates, but makes more money, likes her profession

Chemistry, once strictly a man’s profession, has become increasingly hospitable to women. The expansion of industrial chemistry has helped. Women are particularly in demand for delicate laboratory work that requires small hands, finger dexterity and painstaking attention to detail. With job opportunities opening in the field, more college girls than ever before have been preparing for careers in chemistry.

Today, seven percent of all chemists in industry are women. The ratio is a good deal higher at the Merrimac Division of the Monsanto Chemical Company at Everett, Mass. Here, nearly 20 per cent of the research staff of 109 is female. One of them, Jacqueline Bates, is seen on these pages in a typical day’s work. She is one of four women who make up the analytical laboratory staff. Their job is (1) determining the identity and purity of organic compounds, (2) establishing methods for control of production and application of chemicals, and (3) evaluating new methods of analysis lor organic and inorganic compounds.

At 22. Jackie Bates has made chemistry her career. Although it is a lonely, tense, exacting, sometimes frustrating profession, she enjoys it. She finds her work satisfying, her day full, her advancement altogether satisfactory. After 18 months on the job she regards herself as a veteran: “The sulphur dioxide smell doesn’t bother me any more.” (It’s common at a sulfuric acid plant, but visitors unaccustomed to it notice it.) Jackie’s working day begins at 7:30, when Robert Voigt, a chemical engineer at the laboratory, picks her up at her home in Arlington, Mass., for the 15-mile drive to the huge, sprawling Monsanto plant on the outskirts of Boston. She makes a hurried change from street dress to white uniform and is at her post in the lab shortly alter eight. She works steadily through the morning with time out only for coffee, which she has in the lab. Lunch at noon comes on a tray from the company cafeteria, if experiments in progress cannot be left un-watched. Back at work before one, Jackie normally finishes her day by 4:45. Attractive, full of good humor, Jackie enjoys dancing, swimming, bowling and no shortage of male admirers. During her first year at Simmons College in Boston. Jackie planned to be a nurse after graduation, but by sophomore year shifted her interest to biology. In her third year she settled on chemistry. This choice withstood the test of a summer job at Monsanto, which convinced Jackie that she wanted to make chemistry her career. Personnel at the laboratory were impressed with her ability, and Monsanto again became her employer in July. 1947, after she had received her BS. The success of girl chemists like Jackie has opened another career door for college girls.

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NEW FEATS OF Chemical Wizards REMAKE THE WORLD WE LIVE IN

By ALDEN P. ARMAGNAC

IMAGINE a ball of fiber, weighing only one pound, of so fine a texture that if unrolled it would reach from the Atlantic to the Pacific! This marvel of chemistry, exhibited when American chemists recently assembled at Kansas City, Mo., to compare their achievements, is the latest kind of rayon, or artificial silk. A garment made from it can be hidden in the palm of the hand. To produce it, laboratory workers have gone the silkworm one better—for it measures one third thinner than natural silk. Improvements in methods of purifying the wood pulp that serves as its raw material, and in the chemical solutions and machinery used in its manufacture, have combined to make its production possible.

As much like Arabian Nights tales read the stories of other feats that chemists here and abroad are accomplishing today. Your home, your clothing, your car, and the whole world about you are benefiting from the wizardry of their touch.

From water in which corn has been steeped during the manufacture of corn starch, they have found, comes a sugar with a strange dual personality. In its pure form it is sweet to the taste, and is being tested clinically as a substitute for ordinary sugar in diets for diabetics. Treat it with nitric acid, however, and it becomes an explosive more powerful than nitroglycerin! It has the advantage that no inert material need be added to prepare it for use; nitroglycerin, a liquid, must be molded with earth to be usable in the form of dynamite, but the “corn dynamite” is naturally a solid. Applications are foreseen for it in quarrying, excavating, and tunneling. As much as 1,-000,000pounds of the sugar, it is estimated, can be recovered annually as a by-product of starch manufacture through a new chemical process devised by Prof. Edward Bartow, president of the American Chemical Society.

BY SEEKING a substitute for air, Dr. J. Willard Hershey of McPherson College, McPherson, Kans., demonstrates that the audacity of chemists knows no bounds. Recently he reported that he had found something better for human lungs than the natural atmosphere!

Air that we breath contains nitrogen, oxygen, and small amounts of helium and other rare gases. Would any other gaseous mixture support life as well? Shutting mice, guinea pigs, cats, and monkeys in hermetically sealed glass jars, Dr. Hershey experimented to see how long he could keep them alive in gases and gas mixtures of every possible variety, constantly supplied through tubes. Pure air, he found, came out second best in his tests! Animals thrived in a mixture of helium and oxygen, which received the highest rating. Sufferers from diseases which cause difficulty in breathing may be the first to benefit from his discovery.

Your motoring habits may be changed in two important particulars—tires and fuel—by new discoveries. Current reports from Germany announce that chemists there have mastered the production of artificial rubber—but American experimenters have already shown the way. A factory at Deep Water Point, N. J., is now producing synthetic rubber at the rate of a million pounds a year. Special properties give it important advantages over natural rubber in many applications, and in time of war, chemists foresee, it might replace the natural product altogether. For the discovery of its basic raw material—a compound of ordinary acetylene, named vinylacetylene— Father J. A. Nieuwland of the University of Notre Dame has received the Nichols Medal, one of the highest American honors in chemistry. “This field of research requires greater courage than that of the battlefield,” declared the chemist who made the award. Experts well know the extreme danger of an explosion during experiments with little-known derivatives of acetylene, and Father Nieuwland, fully aware of his peril, risked his life countless times in the tests that led to his success.

MAKING gasoline from coal is an achievement of foreign chemists, and huge plants for the purpose have recently been set in operation in England and Germany. Meanwhile, in this country, the idea of blending gasoline with alcohol for motor fuel has aroused chemists to lively controversy.

Advocates of the plan maintain that it would enable farmers to dispose profitably of huge quantities of surplus farm products such as corn, which would be chemically converted into alcohol. Blending the product with a considerably larger proportion of gasoline, they assert, would yield a motor fuel at least as good as pure gasoline, if not actually superior in economy and power output. While chemists are not in agreement as to the value of the proposed blend, U. S. Bureau of Standards experts hold that it would be a satisfactory motor fuel, with one important proviso —that, for best results, engines should be especially designed for its use. At any rate, motorists will soon have an opportunity to judge for themselves, for as this is written a plant is being opened at Atchison, Kans., to produce 10,000 gallons of alcohol from corn daily. The product is to be used to produce blended motor fuel, which is expected to sell at the same price as ordinary gasoline.

Photographers may benefit by an accident that recently befell four young research chemists, engaged in preparing a batch of photographic emulsion. To their surprise, printing paper coated with the emulsion produced black where the white parts of the picture should have been, and white where they expected black. Investigating, they found that they had stumbled upon a formula for a new kind of material for photographers, which permits direct photographs to be made, without requiring the production of a negative as an intermediate step. The new emulsion is declared suitable for films, plates, or paper, and is developed by standard methods and solutions.

New metals are emerging from the laboratories of modern alchemists. Superior blades for safety razors are promised by a steel alloy developed especially for the purpose by chemists of the Mellon Institute of Industrial Research at Pittsburgh, Pa. Other combinations of metals have recently yielded a tungsten alloy that replaces lead as a shield against the powerful rays of radium; an alloy of iron, aluminum, nickel, and cobalt from which the most powerful permanent magnets in the world are now being made; and stainless-steel alloys, combining beauty with strength, for building railway cars. So bewildering is the variety of new alloys constantly being developed that chemists themselves are hard put to it to keep up with advances made by their own colleagues. To index what is known today about iron and steel alloys alone, in handy form for reference, the Engineering Foundation of New York City has put 150 men to work on a monumental search of the whole world’s technical literature—an undertaking believed to be unprecedented in scientific history.

STRANGE as it may seem, one of the metals about which chemists know the least is iron itself! Pure iron is almost a myth. That the iron we know bears little resemblance to it, however, was demonstrated not long ago when experimenters produced the purest specimens on record, by heating them in hydrogen flames. The iron they obtained does not rust in pure oxygen and water, even after months of exposure.

Unfamiliar colors will greet the eye in the strange new world being created by scientific investigators. British chemists recently announced the discovery of a new blue coloring pigment, for use in paints and printing inks. Until now, ultramarine, discovered in 1704, and Prussian blue, discovered in 1826, have enjoyed a virtual monopoly for the production of this shade. Neither, however, has possessed all the qualities prized in a pigment—brightness, strength of coloring, and fastness to light and heat, as well as to acids, alkalies, and other solvents. The new pigment, christened “monas-tral fast blue,” is hailed as satisfying every one of these tests. In addition, it is declared the nearest approach yet made to an ideal shade of blue for color printing. Random examples like these show how the triumphs of chemists are affecting every branch of life. Some of the most remarkable transformations wrought by their magic, however, may occur right in your own home.

Wooden furniture, for instance, may become out of date before long. Things that have always been made of wood or metal —radio cabinets, bottle caps, bowling pins —are now being fashioned from synthetic materials known as plastics, created in the chemist’s test tube. That, experts say, is only a beginning. Imagine tables, chairs, and beds made of these glistening plastics, easy to keep spotless and difficult to scratch or mar! At this very moment, only the slightly higher cost of the synthetic materials stands in the way of their universal use. If it can be pared down so that plastics can compete with wood and metal on a price basis—and this is quite within the realm of possibility, according to William Haynes, New York chemical expert—their possible applications become startling. “Just one industry I can find,” another prominent chemist, Dr. John E. Teeple of New York, remarks jestingly, “where the disappearance of wood might be a horrible calamity. I cannot see how the manufacture of antique furniture could continue without wood!”

EVEN the supremacy of glass as a material for windows is threatened by new transparent materials of the plastic type. While their value for use in the home remains open to speculation, they have a number of desirable qualities. One is their flexibility, permitting them to be bent into curved shapes—a feature that has already led to their adoption for airplane windows.

Will steam heat go into the discard one of these days? Chemists are developing a preparation to take the place of steam in a heating system.

Its base is a white, flaky compound known as diphenyl—a chemical relative of synthetic geranium perfume—which turns to vapor at about 500 degrees F. Since it holds more heat than steam, and can be raised to a greater temperature without developing dangerous pressure, the new heat-carrying material has already found industrial applications.

HOME refrigeration, too, has come in for attention from the chemical engineer. Ice boxes employing “dry ice,” or solidified carbon dioxide, as a refrigerant have recently been introduced, particularly for use in hot regions of the country where ice factories are remote and where electricity is not available. The dry ice is placed in an insulated inner compartment so that it will not withdraw heat too rapidly, as its temperature of 109 degrees F. below zero would otherwise freeze solid the whole contents of the refrigerator. Its chilling effect, transmitted through metal fins on the compartment, can be regulated to keep the ice-box temperature within the desired limits. A novel advantage resulting from evaporation of the refrigerant is the atmosphere of carbon dioxide formed within the ice box, which is said to retard bacterial growth and also to check the spread of food odors.

Frying pans of glass with superior heat-resisting qualities, for cooking on top of the stove, are the result of a recent chemical improvement upon the glass used in standard oven ware. Behind this development lies the story of chemists who turned cooks in a Corning, N. Y., laboratory to test glassware made from as many as 1,500 promising new formulas. Tons of potatoes and countless hamburger steaks sizzled in their dishes. Hungry dogs, more pleased than the scientists themselves with some of the first results, got many of the meals. Some of the food was burned black—purposely—to see what the glassware would stand. Eventually the experimenters arrived at the formula they were seeking, which is embodied in the glassware that has just reached the market.

No article used about the home is too inconsequential to attract the interest of skilled chemists. One has just produced a “nonskid” floor wax by impregnating ordinary wax with rubber, preventing falls on a freshly polished floor. Another has improved cedar chests by perfecting a transparent exterior coating which retains both the natural oil of the wood and its moth-repelling aroma. Thus, even to the smallest details, chemists are helping to make the world a better place to live in.