Twin Discovered for Carbon

CARBON is the latest chemical element to be shown to have a twin. Last winter two California physicists showed that oxygen, long supposed to be single, was not only double, but triple. Now Dr. Arthur S. King, of the Mt. Wilson Observatory, and Dr. Raymond T. Birge, of the University of California, have found a kind of carbon that is heavier than the ordinary form. Carbon is one of the most essential elements in living matter. These experimenters heated carbon in a vacuum in an electric furnace to a temperature around 5,000 degrees Fahrenheit. When the light that it emitted was analyzed with a spectroscope, the usual bright bands of the spectrum appeared.

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Unlocking Fortunes from Atoms

by Jay Earle Miller

Now that chemists have discovered the last element, it remains for the research worker to find practical uses for substances which are at present mere laboratory curiosities. Somebody will make a fortune one of these days by finding ways of using gallium, germanium, tellurium, and many other “unpopular” elements.

THE last missing thing that goes to make up our known world was detected recently. Almost simultaneously the next to the last of the ninety-two elements, which had been located last year, was isolated in the form of a metal, and isolation of the last may be expected shortly.

With those two discoveries chemistry has finished the job of determining what the earth and every animate and inanimate thing in it is made of. Ninety-two things—some gases, some metals, some chemicals of different sorts, ranging from hydrogen at one end of the table to uranium at the other—have now been found. . Last Two Elements Discovered And because uranium and several of its near neighbors are radio-active, so complex in their atomic structure that they break down spontaneously, there is reason to believe that nothing more complex than they are can exist on this particular planet, and therefore that the table is complete.

Four research specialists working at Alabama Polytechnic Institute, located No. 85, the last missing element, in sea water, in potassium bromide, and in several well-known minerals by a process of super chemical analysis.

Two of the four last year had detected No. 87, and on October 15th a professor at Cornell isolated samples of it in a substance known as samarskite, a lustrous, velvet black mineral from Norway. And within ten days after that event a prospector located deposits of samarskite in the West, so if a useful field is found for the new element the United States will have supplies of it.

Uses Must Be Found For Elements The interest of the scientific world centered in the fact that the atomic table of the ninety-two elements which, supposedly, include everything that exists in this world of ours, has finally been completed, in accordance with the prediction and rules laid down by John Dalton, an English chemist, 128 years ago.

But for the practical man the discoveries open two new roads to fortune—if a use can be found for the newly-found substances. The story of fortunes made from the identification of new elements is a long one.

Von Welsbach was just a research scientist busy exploring the unknown world of the so-called “rare earths” until he conceived the idea of utilizing zirconium, which had been discovered by Klaproth in 1789, to make the original Welsbach gas mantle. Later he found that a combination of thorium and cerium, the former discovered by Berzelius in 1828, and the latter by Klaproth in 1803, made a still better mantle.

Aluminum Once a Curiosity Aluminum, known to the ancient Romans, was just a laboratory curiosity until Charles Hall, an American, found a way to extract it from its ores on a commer cial basis and so founded the great aluminum industry.

Edison searched the world and spent thousands of dollars seeking a practical material from which to make incandescent lamp filaments, while all the time tantalum, which had been found by Berzelius in 1820, and tungsten, first isolated by K. W. Scheele in 1781, were waiting to be put to work.

For a century the world went along believing it knew everything about air, and then, in 1898, after four years’ work, Sir William Ramsay, partly with the assistance of Lord Rayleigh and partly with the aid of Travers, proved that the air was not composed entirely of oxygen and nitrogen, but contained four additional rare gases. Those discoveries gave birth to argon, neon, krypton and xenon.

Rare Gas Builds a Fortune Then came Dr. Claude, seeking a cheaper method of extracting oxygen and making acetylene gas, and found himself with so much of the rare gases on his hands that he invented the neon tube and sign industry to dispose of the waste product, and so made several fortunes.

A spectroscopic photograph taken during an eclipse of the sun in 1868 showed a new and unidentified line, and the new element was named helium. In 1903 Sir William Ramsay proved the element, which had not then been discovered on this earth, was a by-product of the disintegration of Pierre and Mme. Curie’s newly-discovered radium. Next the gas was found in American natural gas wells, and so the helium airship was born.

Those are just samples of what has been done. There is plenty of fame and fortune waiting for the man who discovers new uses for* any one of a long line of elements. An enormous amount of research work is being done trying to find new practical uses for ruthenium, one of the series of so-called “platinum metals” which all belong to the platinum family. Large stocks of ruthenium have accumulated as a by-product in the purification of other metals, and, aside from some limited use of ruthenium red as a dye in the silk industry, no commercial market has been found.

Columbium, first found in 1801 by C. Hatchett in New England, and so named tor the United States as the first element discovered in American ores, is another example of a metal waiting for a market, it is scarcely used in the arts. A steel-gray metal, nearly as hard as wrought iron, malleable and capable of being welded, it is strongly resistant to all acids save a warm and highly-concentrated solution of sulphuric acid. It belongs to an- other family of elements, the other members of which are vanadium, widely used in making special steels, and tantalum, which was used for a time for incandescent light filaments, until replaced by tungsten.

Gallium Wants a Job One of Lecoq de Boisbaudran’s discoveries—and he made several—may shortly be put to use in several fields. It is gallium, which he discovered in 1875. Because it boils at a temperature of 1,700 degrees Centigrade it has recently been suggested that this metal might make a« excellent liquid for high temperature thermometers. Other suggestions are that it be used with aluminum in alloys for optical mirrors and as cathodes in metal vapor tubes. New uses for gallium are important to industry, for it is produced in considerable quantities from the residues of the zinc smelters at Bartlesville, Oklahoma.

Germanium, which, as its name indicates was first found in Germany (by C. Winkler, in 1886), is another metal for which the Bartlesville zinc mines would like to find new uses, as it, too, is a byproduct of the zinc smelter. Germanium is a relative of tin and lead.

Rarest Elements Found by Americans The United States lagged in the discovery of missing elements, but out of seven found in the last six years the last three, and therefore the most difficult of all to find, because they were the rarest, have fallen to Americans. Until the discovery of illinium by Prof. B. S. Hopkins of the University of Illinois in 1926 no new element had first shown itself to an American investigator—columbium, while coming from a sample of New England rock, having been isolated and identified by a foreigner.

Last year Dr. Fred Allison and Edgar J. Murphy, at Alabama Poly, identified, but did not isolate, a sample of No. 87. This year they, with Prof. Edna H. Bishop and Anna L. Sommer, utilized a remarkable new method of super-chemical analysis to identify No. 85 in samples of sea water, potassium bromine and a number of well known minerals. At almost the same time Prof. Papish, at Cornell, succeeded in isolating samples of No. 87 in a substance known as samarskite, a lustrous, black mineral first found in Norway and Wyoming.

One Pint Weighs 16 Tons There is considerable evidence that far more complex elements than we know must exist elsewhere in the space, notably in the “white dwarfs” of the star system. Sirius, the dog star, for example, has a companion star which apparently consists of some substance or substances with an average weight of about 16 tons to the pint. It is impossible to conceive of anything with such a weight, yet the size of the body, and its computed attraction for Sirius shows that that must be its density.

The answer may lay in things we do not know about gravitation and magnetism. Einstein recently announced a new “fifth dimensional” system of mathematics to correlate gravity and magnetism, and it is possible that under certain conditions of temperature and pressure those factors may explain the seemingly impossible elements apparently existing in numerous stars.

Changing Lead to Gold Transmutation, the dream of the ancient alchemists who believed that all substances might be reduced to gold, seems nearer and nearer. The discovery by Ramsay that helium was a by-product of the breaking down of radium, was the first instance of transmutation in nature. Theoretically, if given sufficient power, it is possible to knock enough electrons out of the atoms of mercury, No. 80 in the periodic table, to get gold, No. 79, or to go a step farther and transmute gold to platinum, No. 78. In fact it has been claimed that the first step has been done on an experimental scale, the only drawback being that the power consumption made the resulting gold the most expensive metal ever produced.

Reaching out into the future science is speculating on the possibility that the relation of electricity to all things in nature may eventually be proven. There already is considerable evidence that disease, for example, is simply a change in the electrophoretic potential of the cells of the body. Dr. Falk, of the University of Chicago, several years ago found that the various types of pneumonia germs differed from each other according to their electrical potential, and he utilized this fact to develop a quick method of identifying the type of germ taken from any patient.

The solution of germs was placed in a tube leading across the field of a microscope, with electrical connections in either end. The current was then turned on and the observer timed the germs across the microscope field with a stop watch. The greater their negative potential the faster they were attracted to the positive electrode, and the faster they moved the more dangerous and deadly they were, because increased negative potential has something, as yet unexplained, to do with the deadliness of the disease.

Cells Made Artificially When Dr. George W. Crile, of Cleveland, announced the discovery of auto-synthetic cells last winter that same phenomenon of reaction to potential changes was observed. Dr. Crile does not claim that his manufactured cells are living bodies, the same as the cells in your body, but they do act the same, in that they grow and multiply and perform under the microscope just like any other living cell. The materials he used were the ashes of fats, proteins and body ash, mixed in distilled water.

But, getting back to the elements and their possible uses, if you want to go in for chemistry and seek a fortune where so many others have found it, the field is almost unlimited. There are two score or more of the 92 elements which remain laboratory curiosities because no one has found either a way to utilize them or a way to produce them in commercial quantities.

There is titanium, which is unique because it is the only new element ever discovered by a minister of the gospel, its discoverer having been the Rev. William Gregor, who found it in 1789. Titanium is extensively used in the paint and dye industry, and titanium white has been marketed as a rival for white lead and zinc white, but if additional uses could be found the available supply is sufficient to fill them.

Cobalt, a familiar metal widely used in steel alloys was known to the alchemists, who gave it that name because it resembles a metallic ore, yet yielded no metal when smelted. It was not until 1733 that G. Brandt first separated the pure metal, and today it is widely used in telephone and other magnets, as well as in highspeed cutting tools.

Strontium, discovered by Cruikshank in 1787, has found a wide use in separating sugar from molasses. Some of its neighbors in the table— krypton, one of the rare gases of the air, rubidium, yttrium, one of the rare earths, and zirconium, first used by Welsbach in gas mantles, all are awaiting commercial exploitation.

Molybdenum has found wide use in special steels, but the next three, masurium, ruthenium, which has already been mentioned, and rhodium are available for research work. Palladium, discovered by Wollaston in 1802, one of the platinum metals, has been used to coat mirrors, in dentistry, as a substitute for platinum in jewelry during the war, in airplane parts quite recently, and is also the element which, when added to yellow gold, turns it white, and so produces the well-known white gold of the jeweler.

Chromium Made Fortunes Chromium is an example of one of the most recent of the great fortune makers. First identified by L. N. Vauquelin in 1798, it was not isolated until 1859, and later found some use in special hard steels, for, next to the diamond, chromium is the hardest of all known things. Within the last few years the development of workable processes of chromium plating have made fortunes for several people, and threatened to practically displace the nickel plating industry—not that that has harmed nickel sales, for chromium usually is plated over a nickel plate base.

Cesium, lanthanum, praseodymium—one of the rare earths discovered by Von Welsbach before chance led him to make his fortune with the gas mantle—neodymium, another of his rare earth discoveries, illinium, samarium, europium, and gadolinium are others needing exploitation. Gadolinium is especially interesting at this time, for it is found chiefly in samarskite, the same mineral which has just yielded up element No. 87, and which has been found in Wyoming in quantities.

Selinium, discovered by J. J. Berzelius in 1817, went begging for a user for nearly a half century until a patient worker discovered that it was sensitive to light to such an extent that its electrical resistance changed when exposed to light rays. Another half century went by, and then television was born on the strength of that discovery.

Dysprosium is another element which has possibilities. Discovered in 1886 by de Boisbaudran, and first isolated in the pure state by G. Urbain in 1906, it is noted because its salts are the most magnetic of all mineral salts, and therein may lie the clue to someone’s fortune.

Hafnium, discovered at Copenhagen in 1923, rivals radium in the quick application found for it, for, within two years, several patents had been issued for its use in radio tubes, where its high melting point makes it valuable.

Smithson Tennant, an Englishman, got his name into the table of discoverers twice, once in 1803 when he discovered osmium, and again in 1804 when he isolated iridium. He didn’t live long enough to cash in, but his two discoveries have made fortunes for others. For osmium was used for some time to tip fountain pen points, and iridium is the metal used to alloy platinum and make it hard, for it is, in its pure state, too soft to be useful. The famous standard metre, preserved in the vaults of Paris as the standard of measure, consists of 90 per cent platinum and ten per cent iridium.

The University of Wisconsin has done considerable work trying to find new uses for tellurium, which was first discovered way back in 1798 by M. H. Klaproth. American miners of other metals accumulate as much as 125,000 pounds of tellurium a year as a by-product which must be isolated in purifying their ores. It was used as a crystal detector in early radio sets, and for a time found a place as an antiknock fluid for automobile engines, before tetraethyl lead displaced it there.

Cerium, one of the materials now used in gas mantles, also has a place in medicine—as have many other elements—its oxalate being used to prevent seasickness.

This list suggests just a part of the available material out of which new fortunes may be made. And it explains why chemistry is becoming an increasingly important and popular study in American universities, and why great industries are spending fortunes every year in research work.

Name Elements 99 and 100

Two great scientists who died within the last year, Albert Einstein and Enrico Fermi, have been honored by the naming of elements 99 einsteinium and 100 fermium. The symbol for einsteinium is plain E, that for fermium, Fm. Now all discovered elements are named, since 101 was previously named mendelevium after the Russian, D. Mendeleev, who announced the periodic system of the elements in 1869.

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A “Down the CELLAR” Chem Lab

by FREDERICK O. SCHUBERT

Here are some interesting experiments you can perform with simple chemicals, with notes on building the beginnings of your own basement chemistry lab. More next month!

NOW that we’ve succeeded in shoving Andy, the grease monkey, and the rest of the “hangar gang” over a bit for the lab boys, let’s get together and make real use of our “chem” pages. It’s not going to be hard if you spatula wielders and test tube wrestlers will get behind and push the crate along. After all, it’s to satisfy you birds who like to fiddle with specific gravity and distillation rather than angles of incidence or variable pitches that we’re goin’ into the mechanics of chemistry.

The major part of you fellows undoubtedly have some sort of a workshop either in the basement, up in the attic or out in the barn. But to get under way we’ve got to give all the gang a chance to start from scratch. So, at the risk of being considered elementary, we’re going to start by talking about the “down in the cellar chem lab”

that we have designed. Perhaps you seasoned experimenters will catch an idea or two.

As you probably have noted already, no dimensions are given in constructing this lab. Your own best judgment—and the size of space available—will have to govern its construction. Too, the location of the nearest gas, electric and water outlets must be considered. Your own ingenuity will enable you to adapt these plans.

An ordinary kitchen table is perhaps the handiest to start with because it can easily be obtained. It is already equipped with a “drawer and can quickly be set up. The details of shelving likewise are simple and require no great detailed discussion. Enough to say that as your experiments continue and you add more and more equipment and chemicals, the value of adequate shelf-room will be appreciated. They may look empty for a while, but Oh Boy, when you get going!

Much of the testing and mixing equipment must be purchased. If you’re on good terms with the local druggist you can get him to order much of your stuff wholesale. Then—you can make a lot of it yourself.

Take the centrifuge. In recent years this piece of apparatus has saved much time in separating liquids of different specific gravity from each other and solids from liquids when they are held in suspension in such a way that they cannot be filtered. What formerly required days is now done in a very few minutes.

This you can make from an old phonograph. By extending the pin that holds the playing disc and equipping it with a double cross-arm that has slots for test tubes at all four ends you’ve got the finest centrifuge in the world. Wind ‘er up, throw the switch and she is off while you’re doing something else. Although the number of revolutions per minute is not as great as electrically operated machines, it will serve every requirement of the cellar chemist.

A piece of marble can easily be obtained from any junk-shop. This serves well as an ideal mixing board for pasty materials. Bottles, too, can be salvaged from the heap in sizes to meet every need. The same holds true with corks, wire for test tube holders and scraps of pure metals for lab work and analysis. The illustrations show how you can make your own alcohol lamp.

Among the equipment that you will have to buy are graduates calibrated in cubic centimeters, spatulas, mortars and pestles, glass tubing for stirring rods and droppers, funnels, scales, test tubes, filter paper, litmus papers, thermometer, hydrometer and a gas or electric hot-plate. To this should be added a pair or two of rubber gloves for working with acids and a good, substantial rubber apron. A number of good books on chemical experimentation, too, are not amiss for in them will be found much of interest and value.

The question of what chemicals and preparations to purchase is a difficult one to answer in this limited space. When it is considered that the average pharmacopeia contains more than 25,000 chemical preparations to which additions are made almost daily, we can be pardoned our wish to refrain from specifically stating that you should have this or that acid, powder or liquid. As we continue in our monthly rambles we’ll have plenty of opportunity of adding to the chemical end of our lab.

One thing in chemistry that is always interesting is the large number of tricks you can perform. Take the case of glowing pictures. They can easily be made by drawing a picture on white paper with a solution of 40 parts saltpeter and 20 parts of gum arabic in 40 parts of warm water. An ordi- nary pen is used and all lines should be connected. Extend one line out to the edge of the sheet and mark the spot where it runs off with a light pencil. When a burning match is held to this spot the design— formerly invisible—begins to glow and appear in brown.

There is another little stunt that will make a permanent display for your living room. It’s a chemical garden and simple to make. Prepare a small jar full of cold, saturated solution of Glauber’s salts and into the liquid suspend a kidney bean and a non-porous marble, stone or piece of glass by means of silk threads. Cover the jar and in a few days small crystals of sulphate of sodium will be seen radiating from the bean increasing it in size and giving it the appearance of a sea urchin. The non-porous body will remain unchanged.

An explanation of this stunt states that the bean appears to have a special partiality for the crystals, which is due to the absorption of water by the bean, but not the salt. In this way a supersaturated solution is formed in the immediate neighborhood of the bean and the crystals, in forming, attach themselves to its surface.

A good “weather vane” can be made from white blotters saturated in a solution of one ounce of cobalt chloride, half ounce sodium chloride, 75 grains of calcium chloride, quarter ounce of acacia and three ounces of water and left to dry. The amount of moisture in the air is roughly indicated by the changing color of the blotters. Rain is denoted by a rose-red tint, lavender-blue bespeaks dry weather and a bluish-red announces a change in weather conditions.

Speaking about the weather, here is a handy little weather glass that is even more accurate than the papers. Dissolve 2-1/2 drachms of camphor in 11 drachms of alcohol and 38 grains each of saltpeter and sal ammoniac in ’9 drachms of water. After mixing the two solutions pour them into test tubes and cork them airtight—or better yet, draw out the tube until only a pin hole remains. If you use the corking method run a red-hot wire through the corks so that you will have a very small hole about the size of a pin.

When the camphor appears soft and powdery and almost fills the tube you can expect rain and south or southwest winds. When the substance is crystalline expect fine weather with winds from the north, northeast or northwest. When only a portion crystallizes on the side of the glass, wind can be expected from that direction.

During fine weather the substance remains clear and at the bottom of the tube. When the substance begins to rise and a small star can be seen swimming around in the liquid you can be sure that rain is on the way.

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“HOT DOGS” IN THE LAB

by Harry M. Schwalb

Condensed from The Laboratory In 1939 the hot dog hit the front pages of the international press when President Roosevelt’s wife served it to the king and queen of England. And as 1955 draws to an end, Americans, by devouring over 8y2 billon tangy “red hots,” have made the frankfurter or wiener (formulation’s the same, though the wiener is a bit shorter) a major phase of the meat industry, outranking everything but ice cream in popularity on the national menu.

Few people realize that the humble frankfurter (which got its official name when a German butcher popularized it in Frankfurt a century ago) was commonplace 1,000 years before Christ; fewer still are aware that it has finally arrived in the transforming halls of the modern laboratory.

For example, in the American Meat Institute Foundation laboratories on the University of Chicago campus, you can watch bacteriologist Eileen Felton and physicist William Powers create a genuinely hot hot dog. They stuff the frank into a metal “bun,” lower the sample into a cobalt-60 furnace, and study the effect of gamma radiation on the processing of the foodstuff.

When atomic sterilization is perfected to the point where the franks maintain the desired ruddy color and develop no side flavors or odors, the food industry will have an effective new method of preparing franks for long-range non-spoilable keeping for soldiers and civilians the world over.

In still another Foundation laboratory, scientists study not the new techniques but the traditional ones (smoking franks to an internal temperature of 155 degrees Fahrenheit for two hours). And in their laboratory-sized smokehouse, chemists control the gas velocity, flow direction, temperature, humidity and smoke density with amazing delicacy.

From even the miniature lab smokehouse it is a tremendous jump in scale to the bench of histologist Hsi Wang, who uses a “micromanipulator” with his microscope to segregate the components of a single bacterial cell. Bacteriology is an important part of hot-dog science since the organisms that attack frankfurter surfaces, though non-toxic, are unusually cantankerous: they’re not only unafraid of the heavy salting that frankfurters get, they actually grow in media containing 10 percent salt. These bacteria produce hydrogen peroxide; while this turns brunettes into pretty blondes, it turns the ruddy frankfurter surface into a non-marketable green. These “greening” epidemics have been probed by frankfurter biochemists who, through their directions as to plant cleanliness and new smoking-room temperatures, just about have the intruders licked.

Supermarkets have boomed the success of the hot dog; they have also introduced a new problem: the intense artificial light of the retail counter. The problem initiated the study of meat pigments and what happens to them during processing, during exposure to light, and in the presence and absence of oxygen. Chemicals like disodium phosphate and ascorbic acid may solve the problem once and for all.

Frankfurter casings are still another lab achievement. For years, sheep and hog casings from China, Russia, and Australia were used, despite the fact that no genuinely clean, hole-free and uniformly-calibrated sheaths could be obtained. A fellowship was set up at Pittsburgh’s Mellon Institute with the object of devising an edible synthetic substitute. For ten years, chemists investigated all edible film-forming materials—gelatins, casein plastics, agar agar, algin, starches, cellulose.

Selecting cellulose, they next had to find a way to eliminate that stiff “papery” quality. After experimenting with every known softening treatment, they created an untreated soft, pliable film merely by making the film ultra-thin. Today, “skinless franks” outsell the other kind.

Until recently, little data had been developed on the nutritional properties of the hot dog; the thing was such a pleasure to eat, it was- almost too much to expect that it could be good for anyone, too. Especially unknown was the proportional nutrient-loss due to processing.

So, biochemists descended on retail stores in Chicago, scooped up samples, subjected them to hundreds of analyses. Result: the hog dog has been shown to be as adequate a source of 18 essential amino acids, plus thiamin, riboflavin, niacin and iron—as fresh beef, pork and lamb.

*The Laboratory, Vol. 24, Number 25; published by Fisher Scientific Co., 717 Forbes St., Pittsburgh 19, Pa.

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ADVENTURES of the POISON SQUAD

by James Nevin Miller

IN THE city of White Plains, N. Y., not so long ago, more than 700 people suddenly were stricken with a mysterious ailment. City authorities thought the case was food poisoning. But just what kind, puzzled them. True enough, it was learned that all the victims had eaten chocolate eclairs, cream puffs or Boston cream pies. However, none of the custard-filled pastries appeared to be “spoiled” although it was suspected that contaminated custard filling might have been the source of the poisoning.

Therefore, the city officials turned over samples of the pastry to the New York office of the U. S. Food and Drug Administration. These Government experts in turn sent the pastry to the Washington laboratories of the Administration.

Very quickly, the Federal investigators, aided by local authorities cleared up the situation by tracing the source of the spoiled food to one manu- facturing bakery in Westchester County, seizing and destroying all shipments sent out on the same day as the food poisoning outbreak, and identifying the deadly bacteria that caused the large-scale illness.

Thrills are commonplace to these health sleuths, who are known unofficially as Uncle Sam’s “Poison Squad.” Their job is to guard the safety of your pantry by tracking down, armed with ingenious laboratory and field methods, outstanding cases where moldy fruits and vegetables, contaminated sea foods, spoiled canned goods, and doubtful-looking imported edibles offer a serious menace to the health of the nation.

Just a few of the Poison Squad’s recent achievements, besides the settlement of the poisoned pastry case, include: the rounding up, during a three months period, of nearly a quarter million shipping cases of canned salmon, an appreciable percentage of which was spoiled; the investigation of a food poisoning outbreak in Philadelphia somewhat similar to the one at White Plains; and the discovery of the cause of a strange “epidemic” of the parasitic disease known as trichinosis, at Williamsville, N. Y.

Seizure of poisonous, decomposed or filthy foods is like stopping a murderer’s bullet in flight, says Dr. A. C. Hunter, chief bacteriologist of the Poison Squad. Oftentimes such seizure has been criticized as only a worthless gesture that may be compared to the arrest of a murderer’s revolver after it has slain its victim. Dr. Hunter does not agree. He says it is better to protect the public health by confiscating dangerous foods before they can cause injury than it is to prosecute after the damage is done. Seizure is the prompt and effective weapon which is the first reliance of the Food and Drug Administration.

Were it not for the vigilance of the Poison Squad your family right now might be eating apple butter or apple jelly containing lead or arsenic. Recently a jury, after studying a thorough investigation by the Federal food sleuths, found a big fruit company in the state of Washington guilty of a violation of the Food and Drugs Act in shipping interstate stocks of apple scrap which contained residues of poisonous lead and arsenical sprays. About 46,000 pounds of the scrap had been seized by the undercover men assigned to the case.

Surprisingly enough, nearly one-third of the time, money and effort expended by Uncle Sam’s food policemen is being devoted to protecting the public from the danger of poisons used in sprays to combat insect pests and diseases that attack fruits and vegetables. Every year thousands of carloads of fruits and vegetables are given painstaking laboratory examination to detect traces of such health-destroying residues. However, many of the states are cooperating and the situation is improving rapidly from year to year, officials say.

But, to complete the story about the White Plains poisoned pastry case. Aided by local health officers, the Federal food sleuths, as mentioned earlier, traced the source of the spoiled food to a single manufacturing bakery in Westchester County and rounded up and destroyed all shipments sent out on the same day as the food poisoning outbreak.

Meantime New York agents of the Food and Drug Administration made a thorough inspection of the bakery that made the pastry and found not a single trace of any unsanitary conditions. To this day the case is somewhat enshrouded in mystery, but it is believed by the Government scientists that the cream-filled pastry became poisonous mainly because it became unduly exposed to warm temperatures without proper refrigeration.

This outbreak and many other somewhat similar ones emphasizes that cream-filled pastries, since they are ideal for the growth of bacteria, should be produced, handled and refrigerated with extraordinary care if they are to be held any length of time before consumption.

A good many months ago New York City members of the Poison Squad tracked down 15 cases of the parasitic disease known as trichinosis, at Williamsville, N. Y. The outbreak, like others occurring in the United States, was traced to the eating of raw, or improperly cooked pork that was infested with the parasite known as Trichinae. Although 8 of the 15 persons affected were confined to hospitals, no deaths were reported. This fact is attributed to the prompt action of the Federal health sleuths, aided by Dr. Myron Metz, local health officer, who obtained a list of the buyers of the infected pork and advised each person to call a physician in case he felt sick.

A case of food poisoning in North Dakota, in which 12 persons died from eating home-canned peas, has prompted the United States Department of Agriculture to call attention again to a method of canning non-acid vegetables in the home to guard against the deadly botulinus poison.

In the canning of non-acid vegetables—peas, asparagus, beans, corn, beets, and spinach—the only safe course is to destroy all bacteria that may be present by canning under steam pressure, according to the Bureau of Home Economics. In the case of acid vegetables and fruits, such as tomatoes, apples, peaches, and gooseberries, the bacteria are killed at boiling temperature (212° F.) but with non-acid vegetables there is no assurance that the botulinus organisms will be killed by processing in boiling water unless the material is heated for six hours or longer. Obviously, a 6-hour treatment of peas or similar vegetables would result in a very unattractive product. A much shorter heating time is required at a temperature of 240° or 250° F., such as may be obtained in a pressure cooker.

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Inventions Needed in Field of Electrochemistry

An interview with Professor Colin G. Fink
Head, Division of Electrochemistry Columbia University

by Richard H. Parke

“THE young inventor looking for new worlds to conquer would do well to investigate the vast but little-explored domains of electrochemistry. Hundreds of new products and inventions difficult or impossible to discover during the countless ages of the past with mechanical skill alone are today readily possible through the combined power of electricity and chemistry Thus Professor Colin G. Fink of Columbia University presents an invitation — and a challenge—to inventive minds everywhere.

He sees no reason why young men with courage and imagination should not play an important part in future scientific development and he believes his own field offers unlimited possibilities for advancement.

When I called on him in his laboratory at Columbia, I told him the readers of Modern Mechanix were particularly interested in needed inventions and I asked him to name a few of the problems which, in his opinion, demanded the greatest attention at this time.

“Our basic engineering metal is iron,” he began. “It is the most abundant and has the most valuable properties, such as high tensile strength, high fatigue value, etc. The world produces more of iron and steel than all the other metals put together. For every 100 tons of iron and steel, only 2 tons of copper are produced—and copper occupies second place in quantity and importance among metals!

“With ever rapidly increasing utilization of steel in railroads, automobiles, bridges, buildings, high transmission towers and other articles, the problem of eliminating rust and protecting steel products and costly structures against corrosion is undoubtedly the biggest problem confronting the chemist and engineer. Metals rust as a result of electrochemical action. If we eliminate this, we eliminate rust! The young men of the coming generation will solve this problem; marked progress towards a solution already has been made.

“We are also badly in need of an electric source of illumination suitable to our eyes which will operate at 90 per cent or better in efficiency instead of less than 10 per cent as at present. Electricity is modern man’s indispensable and versatile servant; eliminate it and civilization would drop back 100 years, But, although electricity has increased the efficiency of old hand-operated machines and innumerable devices a thousandfold or more, it has so far not accomplished anything near as much for artificial light.

“The tungsten lamp in the average man’s home furnishes but 5 cents’ worth of light for every dollar’s worth of electricity. The balance, 95 cents, is lost as useless heat. Modern homes and towns need a lamp that will furnish at least 90 per cent light and only 10 per cent heat. The little firefly, or glowworm, operates its light at figures even better than 90 per cent. The lamp of the future may very well be based on principles similar to those of the firefly’s glow.

“A simple means of electrically controlling rainfall—keeping it out of the cities—is another important problem to solve. The Cottrell electrical precipitation process is used at factories all over the world to precipitate fumes, dusts, smoke and vapors of all kinds. It requires very little development to apply this same process to the precipitation of rain in localities outside of towns and cities. Millions of dollars’ worth of shoes, clothing and other goods are annually destroyed in the cities on account of rain.

“We should systematically investigate the application of electric currents in the stimulation of the growth of living cells and the formation of many organic compounds. The production of fruits, vegetables and other food materials in a few localities and then shipping them half way around the globe appears to us an awful economic waste and extravagance. With the aid of electricity, truck farms located in the outskirts of cities, or on the roofs of skyscrapers, will produce any, or all, fruits and vegetables. Why not?

“One other important problem is a means to convert ores into finished metal products easily and efficiently. Metals occur in nature combined with other elements, notably sulphur. The old process that dates back at least 5,000 years consists in roasting the sulphur compounds, thereby producing oxides; and then mixing these oxides with coke or charcoal and heating to high temperatures to reduce the oxides to metal— usually very impure at that. Electrically, it will be commercially possible some day to produce pure metals free from sulphur, phosphorus and other deleterious or objectionable impurities.”

Dr. Fink is, himself, working on two important developments at the present time. He is experimenting on a way to convert the sun’s rays into electric power and he is trying to find a satisfactory method of extracting gold from the seas.

In explaining his “power from the sun” theory, he told me that mankind was rapidly approaching the need for additional power resources.

“There is not enough water power on earth to supply more than 10 per cent of our energy demands,” he said, “and since coal, oil and gas necessarily produce most of the remaining 90 per cent, a substitute must be found. I believe this can be accomplished through the use of photo-electric generator cells and I predict that the day will come when these cells will be placed, for example, on the roofs of large apartment buildings and will be used to operate electric refrigerators, flat irons, toasters, vacuum cleaners and other appliances. Still greater uses will of course eventuate.”

Dr. Fink’s new photo-electric generator cell produces about 25 per cent more current than devices formerly used and also is more sensitive to light. The current obtained is still far too small to be practical for commercial power generation but he expects a vast improvement will be made in the near future.

Pointing out that this cell is not the “electric eye” of the motion picture and television, he said it was composed of a sheet of metal in a salt solution or in contact with a salt which, like silver bromide, is sensitive to the sun’s rays. Opposite to this sheet of metal, but kept in the dark, is a second sheet of metal. While the sun shines on the first an electric current flows from one sheet of metal to the other.

“It has been known for a long time,” he explained, “that some chemical compounds are changed in composition upon exposure to the sun. Now certain of these compounds change in one direction only. Thus white silver bromide is converted to a black product when exposed to the sun, as all amateur photographers know. Another well known case, especially for the Sunday hiker, is the way green pop bottles will be found to have changed to a lavender shade after having been discarded and left to lie on the ground exposed to the sun’s rays.

“The chemical salts or compounds we are particularly interested in at the present are those that will change to a new compound in the light but will change back to the original when placed in the dark.”

The “gold from the sea” problem is one that has baffled scientists for many years. One of the main difficulties in this field has been the fact that the gold particles, or ions, that carry a positive charge of electricity, are unable to get close enough to be deposited on the negative electrode, or cathode, because the alkaline film surrounding the cathode is too thick. This alkaline film is brought about by the decomposition of sodium chloride (ordinary table salt) at the cathode, the sodium ions being converted to sodium hydroxide (lye).

Dr. Fink has succeeded in overcoming this film obstacle by rotating the cathode at high speed, an action that reduces the thickness of the alkaline film.

He uses a copper disk for his cathode, which is electrically propelled and spins around at a terrific speed.

It is interesting to know that his original apparatus was the familiar malted milk stirrer. He took off the little nut at the end of the rod and replaced it with a copper disk the size of a half dollar.

But the speed of the malted milk machine was not high enough and he substituted his present apparatus which, while considerably larger and heavier, looks not unlike its smaller counterpart in any drugstore!

In a typical experiment, Dr. Fink employed a disk of 5-centimeter diameter, spinning at the rate of 8,500 revolutions per minute. The disk was set in three liters of a 3 per cent salt solution containing three miligrams of gold. At the end of a half hour, more than 90 per cent of the gold was plated out on the disk.

Dr. Fink pointed out, however, that hope of recovering billions from the seas must be dispelled for the present because the cost of the electricity to operate the cathode is about five times the value of the gold recovered.

He added that at the same time this procedure was more profitable in the case of radioactive metals. The metal polonium, for example, can readily and profitably be recovered from waste solutions.

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unusual compounds find uses because

Electronics Tells The Chemist

By Shirley Motter Linde

THERE are about 750,000 known organic chemical compounds. Less than one percent of these have any known medical or industrial use!

The other 99 percent are a huge potential of untapped applications. They represent hundreds of thousands of chemicals sitting idle on laboratory shelves when they might possibly be useful in curing cancer, fighting viruses, killing insects, giving more gas mileage, making rocket fuels for space vehicles, producing new synthetics, etc.

With present methods, it’s a long, time-consuming, laborious process ferreting out practical uses for new compounds. Let’s say a company comes up with a new chemical as a by-product. They may spend millions of dollars in research trying the compound in one application,

then another trying to find a use. After years of hit or miss research, they might be successful.

Or the company may catalog the fundamental characteristics of the compound and pass the information on to other industries in hopes of receiving suggestions on how the substance might be developed.

Or, if the compound looks like it might work as, say, an insecticide, the company might ask an insecticide expert to compare its properties with known bug killers and decide whether it’s feasible.

All this set scientists at the Midwest Research Institute, Kansas City, Mo., to wondering whether a program might be developed to determine, with the aid of electronic computer techniques, the structural and physical properties which correlate the uses of chemical compounds.

Institute chemist Richard A. Carpenter said, “The mental processes of the expert led us to think of employing rapid electronic equipment to simulate his decisions. Since the expert bases his decisions on the comparison of the new compound with the properties and structures of materials previously found promising, this information could be coded and processed into a memory device. The results would be a sort of ‘super chemist’ armed with millions of facts, correlations and significant patterns from past experience and research. This ‘super chemist’ would be equipped with a never-failing memory able to make many comparisons simultaneously without bias and at a high rate of speed.,,

Like the expert, the computer could calculate, compare, correlate, and come up— in minutes—with possible uses for the new compound.

For example, says Carpenter, the computer can be used to pick out likely chemicals to treat cancer. So far, he points out,

cancer research has been a matter of blindly trying everything that comes along in hopes of a cure. Over 20,000 compounds have been tried already. The computer, on the other hand, can analyze the common characteristics of the more promising substances and pin-point similar but different substances offering the most promise for further experimentation.

Much preliminary work was necessary even before the computer could be called into play. The research team wrote to 1,000 companies for catalogs and data sheets on their compounds. The material poured in. Hundreds of thousands of facts on boiling points, densities, formulas, and uses had to be coded and punched onto computer input cards. A system had to be devised for transcribing that data onto a magnetic tape. This process is still under way.

Monthly trips are made to Washington, D. C, to make use of the huge Vanguard computation center. Only about five minutes a month are allowed the researchers, however, because the computer priority is alloted to satellite problems. But those five minutes are the equivalent of 25 years of desk computing time. It takes researchers three weeks to get enough information together to take up the five minutes on the computer.

After two years of painstaking work, the research team has assembled and coded facts about almost all the 3,000 compounds with known uses. These facts will serve as the computer’s memory. Within a few months the entire memory system will be complete. Then the electronic computer, with its high-speed classifying system, will start on practical problems.

What will be the impact of this electronic brain on chemistry? Says Carpenter: “It will certainly stimulate research, reveal better chemicals for many old and new uses, cut development costs by making research more efficient and productive, generate new knowledge by revealing normally obscured correlations, and most significantly, it will substitute a methodology for the current confusion.”

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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 preciéitate 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 havt 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.