Scientific American Supplement, No. 288 - July 9, 1881
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NEW YORK, JULY 9, 1881

Scientific American Supplement. Vol. XI, No. 288.

Scientific American established 1845

Scientific American Supplement, $5 a year.

Scientific American and Supplement, $7 a year.

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I. ENGINEERING AND MECHANICS—Dry Air Refrigerating Machine. 5 figures. Plan, elevation, and diagrams of a new English dry air refrigerator

Thomas' Improved Steam Wheel. 1 figure

The American Society of Civil Engineers. Address of President Francis, at the Thirteenth Annual Convention, at Montreal. The Water Power of the United States, and its Utilization

II. TECHNOLOGY AND CHEMISTRY.—Alcohol in Nature. Its presence in earth, atmosphere, and water. 6 figures. Distillatory apparatus and (magnified) iodoform crystals from snow water, from rain water, from vegetable mould, etc.

Detection of Alcohol in Transparent Soaps. By H. JAY

On the Calorific Power of Fuel, and on Thompson's Calorimeter. By J.W. THOMAS

Explosion as an Unknown Fire Hazard. A suggestive review of the conditions of explosions, with curious examples

Carbon. Symbol C. Combining weight. 12. By T. A. POOLEY Second article on elementary chemistry written for brewers

Manufacture of Soaps and their Production. By W. J. MENZIES

The Preparation of Perfume Pomades. 1 figure. "Ensoufflage" apparatus for perfumes

Organic Matter in Sea Water

Bacteria Life. Influence of heat and various gases and chemical compounds on bacteria life

On the Composition of Elephant's Milk. By Dr. CHAS. A. DOREMUS. Comparison of elephant's milk with that of ten other mammals

The Chemical Composition of Rice. Maize, and Barley. By J. STEINER

Petroleum Oils. Character and properties of the various distillates of crude petroleum. Fire risks attending the use of the lighter petroleum oils

Composition of the Petroleum of the Caucasus. By P. SCHULZENBERGER and N. TONINE

Notes on Cananga Oil. or Ilang-Ilang Oil. By F. A. FLUeCKIGER. 1 figure. Flower and leaf of Cananga odorata

Chian Turpentine, and the Tree which Produces It. By Dr. STIEPOWICH. of Chios, Turkey

On the Change of Volume which Accompanies the Galvanic Deposition of a Metal. By M. E. BOUTY

Analysis of the Rice Soils of Burmah. By R. ROMANIC, Chemical Examiner, British Burmah

III. PHYSICS AND PHYSICAL APPARATUS.—Seyfferth's Pyrometer. 7 figures.—Pyrometer with electric indicator.—Method of mounting by means of a cone on vacuum apparatus.—Mounting by means of a sleeve.—Mounting by means of a thread on a tube.— Mounting by means of a clasp in reservoirs.—The pyrometer mounted on a bone-black furnace.—Mounted on a brick furnace

Delicate Scientific Instruments. By EDGAR L. LARKIN. An interesting description of the more powerful and delicate instruments of research used by modern scientists and their marvelous results

The Future Development of Electrical Appliances. Lecture by Prof. J. W. PERRY before the London Society of Arts.—Methods and units of electrical measurements

Researches on the Radiant Matter of Crookes and the Mechanical Theory of Electricity. By Dr. W. F. GINTL

Economy of the Electric Light. W. H. PREECE'S Experiments Investigations

On the Space Protected by a Lightning Conductor. By WM. H. PREECE.—5 figures

Photo-Electricity of Fluor Spar Crystals

The Aurora Borealis and Telegraph Cables

The Photographic Image: What It Is. By T. H. MORTON. 1 figure.—Section of sensitive plate after exposure and during development

Gelatine Transparencies for the Lantern

An Integrating Machine. By C. V. BOYS.—1 figure

Upon a Modification of Wheatstone's Microphone and its Applicability to Radiophonic Researches. By ALEX. GRAHAM BELL,—2 figures

IV. ARCHITECTURE.—Suggestions in Architecture, 1 figure.—A pair of English cottages. By A. CAWSTON

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A Chemist of merit, Mr. A. Muentz, who has already made himself known by important labors and by analytical researches of great precision, has been led to a very curious and totally unexpected discovery, on the subject of which he has kindly given us information in detail, which we place before our readers.[1] Mr. Muentz has discovered that arable soil, waters of the ocean and streams, and the atmosphere contain traces of alcohol; and that this compound, formed by the fermentation of organic matters, is everywhere distributed throughout nature. We should add that only infinitesimal quantities are involved—reaching only the proportion of millionths—yet the fact, for all that, offers a no less powerful interest. The method of analysis which has permitted the facts to be shown is very elegant and scrupulously exact, and is worthy of being made known.

[Footnote 1: The accompanying engravings have been made from drawings of the apparatus in the laboratory of which Mr. Muentz is director, at the Agronomic Institute.]

Mr. Muentz's method of procedure is as follows: He submits to distillation three or four gallons of snow, rain, or sea water in an apparatus such as shown in Fig. 1. The part which serves as a boiler, and which holds the liquid to be distilled, is a milk-can, B. The vapors given off through the action of the heat circulate through a leaden tube some thirty-three feet in length, and then traverse a tube inclosed within a refrigerating cylinder, T, which is kept constantly cold by a current of water. They are finally condensed in a glass flask, R, which forms the receiver. When 100 or 150 cubic centimeters of condensed liquid (which contains all the alcohol) are collected in the receiver, the operations are suspended. The liquid thus obtained is distilled anew in a second apparatus, which is analogous to the preceding but much smaller (Fig. 2). The liquid is heated in the flask, B, and its vapor, after traversing a glass worm, is condensed in the tube, T. The operation is suspended as soon as five or six cubic centimeters of the condensed liquid have been collected in the test-tube, R. The latter is now removed, and to its liquid contents, there is added a small quantity of iodine and carbonate of soda. The mixture is slightly heated, and soon there are seen forming, through precipitation, small crystals of iodoform. Under such circumstances, iodoform could only have been formed through the presence of an alcohol in the liquid. These analytical operations are verified by Mr. Muentz as follows: He distills in the same apparatus three to four gallons of chemically pure distilled water, and ascertains positively that under these conditions iodine and carbonate of soda give absolutely no reaction. Finally, to complete the demonstration and to ascertain the approximate quantity of alcohol contained in natural waters, he undertakes the double fractional distillation of a certain quantity of pure water to which he has previously added a one-millionth part of alcohol. Under these circumstances the iodine and carbonate of soda give a precipitate of iodoform exactly similar to that obtained by treating natural waters.

In the case of arable soil, Mr. Muentz stirs up a weighed quantity of the material to be analyzed in a certain proportion of water, distills it in the smaller of the two apparatus, and detects the alcohol by means of the same operation as before.

The formation of iodoform by precipitation under the action of iodine and carbonate of soda is a very sensitive test for alcohol. Iodoform has sharply defined characters which allow of its being very easily distinguished. Its crystalline form, especially, is entirely typical, its color is pale yellowish, and, when it is examined under the microscope, it is seen to be in the form of six-pointed stars precisely like the crystalline form of snow. Mr. Muentz has not been contented to merely submit the iodoform precipitates obtained by him to microscopical examination, but has preserved the aspect of his preparations by means of micro-photography. The figures annexed show some of the most characteristic of the proofs. Fig. 1 shows crystals of iodoform obtained with pure water to which one-millionth part of alcohol had been added. Fig. 2 exhibits the form of the crystals obtained with rain water; and Fig. 3, those with water. Fig. 4 shows crystals obtained with arable soil or garden mould. The first of Mr. Muentz's experiments were made about four years ago; but since that time he has treated a great number of rain and snow waters collected both at Paris and in the country. At every distillation all the apparatus was cleansed by prolonged washing in a current of steam; and, in order to confirm each analysis, a corresponding experiment was made like the one before mentioned. More than eighty trials gave results which were exactly identical. The quantity of alcohol contained in rain, snow, and sea waters may be estimated at from one to several millionths. Cold water and melted snow seem to contain larger proportions of it than tepid waters. In the waters of the Seine it is found in appreciable quantities, and in sewage waters the proportions increase very perceptibly. Vegetable mould is quite rich in it; indeed it is quite likely that alcohol in its natural state has its origin in the soil through the fermentation of the organic matters contained therein. It is afterward disseminated throughout the atmosphere in the state of vapor and becomes combined with the aqueous vapors whenever they become condensed. The results which we have just recorded are, as far as known to us, absolutely new; they constitute a work which is entirely original, which very happily goes to complete the history of the composition of the soil and atmosphere, and which does great credit to its author.—La Nature.

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By H. JAY.

It appears that every article manufactured with the aid of alcohol is required on its introduction into France to pay duty on the supposed quantity of this reagent which has been used in its preparation. Certain transparent soaps of German origin are now met with, made, as is alleged, without alcohol, and the author proposes the following process for verifying this statement by ascertaining—the presence or absence of alcohol in the manufactured article: 50 grms. of soap are cut into very small pieces and placed in a phial of 200 c.c. capacity; 30 grms. sulphuric acid are then added, and the phial is stoppered and agitated till the soap is entirely dissolved. The phial is then filled up with water, and the fatty acids are allowed to collect and solidify. The subnatant liquid is drawn off, neutralized, and distilled. The first 25 c.c. are collected, filtered, and mixed, according to the process of MM. Riche and Bardy for the detection of alcohol in commercial methylenes, with 1/2 c.c. sulphuric acid at 18 deg. B., then with the same volume of permanganate (15 grms. per liter), and allowed to stand for one minute. He then adds 8 drops of sodium hyposulphite at 33 deg. B., and 1 c.c. of a solution of magenta, 1 decigrm. per liter. If any alcohol is present there appears within five minutes a distinct violet tinge. The presence of essential oils gives rise to a partial reduction of the permanganate without affecting the conversion of alcohol into aldehyd.

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By J.W. THOMAS, F.C.S., F.I.C.

A simple experiment, capable of yielding results which shall be at least comparative, has long been sought after by large consumers of coal and artificial fuel abroad in order to ascertain the relative calorific power possessed by each description, as it is well known that the proportion of mineral matter and the chemical composition of coal differ widely. The determination of the ash in coal is not a highly scientific operation; hence it is not surprising that foreign merchants should have become alive to the importance of estimating its quantity. While, however, the nature and quantity of the ash can be determined without much difficulty, the determination of the chemical composition of coal entails considerable labor and skill; hence a method giving the calorific power of any fuel in an exact and reliable manner by a simple experiment is a great desideratum. This will become more obvious when one takes into consideration the many qualities and variable characters of the coals yielded by the South Wales and North of England coal fields. Bituminous coals—giving some 65 per cent, of coke—are preferred for some manufacturing purposes and in some markets. Bituminous steam coals, yielding 75 per cent, of coke, are highly prized in others. Semi-bituminous steam coals, yielding 80 to 83 per cent, of coke, are most highly valued, and find the readiest sale abroad; and anthracite steam coal (dry coals), giving from 85 to 88 per cent, of coke (using the term "coke" as equivalent to the non-volatile portion of the coal) is also exported in considerable quantity. Now the estimation of the ash of any of these varieties of coal would afford no evidence as to the class to which that coal belongs, and there is no simple test that will give the calorific power of a coal, and at the same time indicate the degree of bituminous or anthracitic character which it possesses.

In order to obtain such information it is necessary that the percentage of coke be determined together with the sulphur, ash, and water, and these form data which at once show the nature of a fuel and give some indication of its value. To ascertain the quantity of the sulphur, ash, and water with accuracy involves more skill and aptitude than can be bestowed by the non-professional public; the consequence is that experiments entailing less time and precision, like those devised by Berthier and Thompson, have been tried more or less extensively. In France and Italy, Berthier's method—slightly modified in some instances—has been long used. It is as follows:

70 grammes of oxide of lead (litharge) and 10 grammes of oxychloride of lead are employed to afford oxygen for the combustion of 1 gramme of fuel in a crucible. From the weight of the button of lead, and taking 8,080 units as the equivalent of carbon, the total heat-units of the fuel is calculated. This experiment is very imperfect and erroneous upon scientific grounds, since the hydrogen of the fuel is scarcely taken into account at all. In the first place, hydrogen consumes only one quarter as much oxygen as carbon, and, furthermore, two-ninths only of the heating power of hydrogen is used as the multiplying number, viz., 8,080, while the value of hydrogen is 34,462. In other words, one-eighteenth only of the available hydrogen present in the fuel is shown in the result obtained. Apart from this my experience of the working of Berthier's method has been by no means satisfactory. There is considerable difficulty in obtaining pure litharge, and it is almost impossible to procure a crucible which does not exert a reducing action upon the lead oxide. Some twelve months ago I went out to Italy to test a large number of cargoes of coal with Thompson's calorimeter, and since then this apparatus has superseded Berthier's process, and is likely to come into more general use. Like Berthier's method, Thompson's apparatus is not without its disadvantages, and the purpose of this paper is to set these forth, as well as to suggest a uniform method of working by means of which the great and irreconcilable differences in the results obtained by some chemists might be overcome. It has already been observed that a coal rich in hydrogen shows a low heating power by Berthier's method, and it will become evident on further reflection that the higher the percentage of carbon the greater will be the indicated calorific power. In fact a good sample of anthracite will give higher results than any other class of coal by Berthier's process. With Thompson's calorimeter the reverse is the case, as the whole of the heating power of the hydrogen is taken into account. In short, with careful working, the more bituminous a coal is the more certain is it that its full heating power shall be exerted and recorded, so far as the apparatus is capable of indicating it; for when the result obtained is multiplied by the equivalent of the latent heat of steam the product is always below the theoretical heat units calculated from the chemical composition of the coal by the acid of Favre and Silbermann's figures for carbon and hydrogen. On the other hand, when the heating power of coal low in hydrogen is determined by Thompson's calorimeter, much difficulty is experienced in burning the carbon completely; hence a low result is obtained. From a large number of experiments I have found that when a coal does not yield more than 86 per cent, of coke, it gives its full comparative heating power, but it is very questionable if equal results will be worked out if the coke exceeds the above amount although I have met with coals giving 87 per cent. of coke which were perfectly manageable, though in other cases the coal did not burn completely. It will be noted that the non-volatile residue of anthracite is never as low as 86 per cent., and this, together with the very dry steam coals and bastard anthracite (found over a not inextensive tract of the South Wales Coal field), form a series of coals, alike difficult to burn in Thompson's calorimeter. Considerable experience has shown that in no single instance was the true comparative heating power of anthracite or bastard anthracite indicated. With a view to accelerate the perfect combustion of these coals, sugar, starch, bitumen, and bituminous coals—substances rich in hydrogen—were employed, mixed in varying proportions with the anthracitic coal, but without the anticipated effect. Coke was also treated in a like manner. Without enlarging further upon these futile trials—all carefully and repeatedly verified—the results of my experiments and experience show that for coals of an anthracitic character, yielding more than 87 per cent. of coke, or for coke itself, Thompson's calorimeter is not suited as an indicator of their comparative calorific power, for the simple reason that some of the carbon is so graphitic in its nature that it will not burn perfectly when mixed with nitrate and chlorate of potash. A sample of very pure anthracite used in the experiments referred to, gave 90.4 per cent. of non-volatile residue, and only 0.84 per cent. of ash. This coal was not difficult to experiment with, as combustion started with comparative ease and proceeded quite rapidly enough, but in every instance a portion of the carbon was unconsumed, and consequently instead of about 13 deg. of rise in temperature only 10 deg. were recorded.

Since the calorific power of a coal is determined by the number of degrees Fahrenheit which a given quantity of water is raised in temperature by a known weight of fuel, it follows that every care should be taken that the experiment be performed under similar atmospheric conditions. The oscillation of barometric pressure does not appear to affect the working, but the temperature of the room in which the work was done, and especially that of the water, are most important considerations. It has been observed by some who have used this apparatus—and I have frequently noticed it myself—that the lower the temperature of the water is under which the fuel is burnt the higher is the result found. This has been explained on the assumption that the colder the water used, the greater is the difference between the temperature of the room and that of the water; hence it would be expedient that in all cases when such experiments are made the same difference of temperature between the air in the room and the water employed should always exist. For example, if the temperature of the room were 70 deg., and the water at 60 deg., then the same coal would give a like result with the water at 40 deg. and the room at 50 deg.. This has been regarded as the more evident, because the gases passing through the water escape under favorable conditions of working at the same temperature as the water, and are perfectly deprived of any heat in excess of that possessed by the water. Under these circumstances it would seem only reasonable that this assumption should be correct. It was, however, found after a large number of experiments upon the same sample of coal that this was not the case. 30 grammes of coal which raises the temperature of the water 13.4 deg., when the water at starting was 60 deg. and the room at 70 deg., gives 13.7 deg. rise of temperature with the water at 40 deg. and the room at 50 deg.. Conversely, when the water is at 70 deg. and the room at 80 deg., a lower result is obtained. The explanation appears to be this: The gas which escapes from the water was not in existence in the gaseous form previous to the experiment, and the heat communicated to the gas being a definite quantity it follows that the more the gas is cooled the greater the proportion of chemical energy in the shape of heat will be utilized and recorded as calorific power.

In order, therefore, to make the experiment more simple and workable at all temperatures, a sample of coal was selected, which should be perfectly manageable and readily consumed. Appended is an analysis of the coal employed (from Ebbw Vale, Monmouthshire):

Composition per cent.

Carbon...............................88.33 Hydrogen............................. 5.08 Oxygen............................... 3.28 Nitrogen............................. 0.55 Sulphur.............................. 0.70 Ash.................................. 1.26 Water (moisture)..................... 0.80 ——- 100.00

In the following experiments the standard temperature of the water was taken as 60 deg. F., and as the coal gave 13.4 deg. of rise of temperature, 67 deg. F. was selected as the standard room temperature. The reason for this room temperature is obvious, for, whatever heating effect the higher temperature of the room may have upon the water in the cylinder during the time occupied by the first half of the experiment, would be compensated for by the loss sustained during the second half of the experiment, when the temperature of the water exceeded that of the room. The mean of numerous trials gave 13.4 deg. F. rise of temperature, equal to 14.74 lb. of water per lb. of coal. When the water was at 50 deg. and the room at 57 deg., the mean of several experiments gave 13.5 deg. rise of temperature. When the water was 40 deg. at starting and the room at 47 deg., 13.65 deg. was the average rise of temperature. Trials were made at intermediate temperatures, and the results always showed that higher figures were recorded when the water was coldest. With a view of getting uniformity in the results it was thought well to make experiments, in order to find out what temperature the room should be at, so that this coal might give the same result with the water at 50 deg., 40 deg., or at intermediate temperatures. Without going much into detail, it was found that when the temperature of the room was at 40 deg. and that of the water 40 deg., and the experiment was rapidly and carefully performed, 13.4 deg. rise of temperature was given; but this result could be obtained without special effort when the room was 42 deg. and the water 40 deg. at starting. It is evident that the cooling effect of the air in the room upon the water cylinder is very appreciable when the water has reached 13 deg. above that of the room. When the water was at 50 deg. and the room at 55 deg., the coal gave 13.4 deg. rise with ease and certainty, and it would not be out of place to remark here that with those coals which burn well in Thompson's calorimeter, the results of several trials are remarkably uniform when properly performed. With the water at 70 deg. and the room at 80 deg., a like result was worked out. Experiments at intermediate temperatures were also carried out (see table in sequel). It is true that the whole difference of temperature we are dealing with in making these corrections is only 0.25, but 0.2 in the result, when multiplied by 537 to bring it into calories, as is done by the authorities in Italy, makes more than 100 heat units—a serious difference when 5d. per ton fine is attached to every 100 calories lower than the number guaranteed.

Taking the latent heat of steam as 537 deg. C., and multiplying this number by 14.74, the evaporative power of the coal used in these experiments, its equivalent in calories is 7,915. From the analysis of this coal, disregarding the nitrogen and deducting an equivalent of hydrogen for the oxygen present, the total heat units given by Favre and Silbermann's figures for carbon (8,080) and hydrogen (34,462) will be 8,746. It will be seen, therefore, that the calorific power, as determined by Thompson's apparatus, gives a much lower result when multiplied by 537 than the heat units calculated from the chemical composition of the coal. When I used Thompson's apparatus in the chemical laboratory at Turin to determine the evaporative power of various cargoes of South Wales coal, it was agreed by mutual consent that the temperature of the water at starting should be 39 deg. F. (the temperature at which the heat unit was determined). The temperature of the room was about 60 deg., but this varied, as the weather was somewhat severe and changeable. Under these conditions, with the water at 39 deg. and room 60 deg., the coal which gives 14.74 lb. of water per lb. of coal, will give as high as 15.88 lb. of water per lb. of coal. This result multiplied by 537=8,496 calories, approaching much more nearly to the theoretic value. This method of working is still practiced abroad, but experience has shown that very widely differing results follow when working in this manner, especially if the temperature of the room is changeable, as it naturally is where ash determinations and other chemical work is proceeding simultaneously. The time the experiment lasts, taking the reading on a quickly rising thermometer and other considerations, render the experiments anything but trustworthy when 0.2 of a degree makes a difference of more than 100 calories. In the instructions supplied with Thompson's calorimeter nothing is said as to the temperature of the room in which the experiment is performed, but simply that the water shall be at 60 deg. F. If, with the water at 60 deg., a room were at 50 deg., as it often is in winter, a good coal would give 14 lb. of water per lb. of coal as the evaporative power; but if in summer, the room were at 75 deg. and the water at 60 deg., the same coal would give 15 lb. of water per lb. of coal. If further evidence were needed of the effect of temperature consideration of the experiments already referred to will show how necessary it is that some general rule shall be adopted. Considerable stress is laid (in the instructions) upon the quantity of oxygen mixture used being determined by rough experiments. This I have found leads to erroneous conclusions unless a number of experiments are tried in the calorimeter, as it often happens that the quantity which appears to be best adapted is not that which yields a trustworthy result. There are many samples of South Wales coal, 30 grains of which will require 10 parts of oxygen mixture in order to burn completely, but since a little oxygen is lost in drying and grinding, and few samples of chlorate are free from chloride, it is not safe to use less than 11 parts of oxygen mixture, but this amount is sufficient in all cases, and never need be exceeded. I have made numerous experiments with various coals (anthracite, steam, semi-bituminous, and bituminous, including a specimen of the ten yard coal of Derbyshire), and find that with 11 parts of chlorate and nitrate of potash, they are all perfectly manageable and yield the best results. It is quite clear that the excess of chlorate is decomposed in all instances, and the latent heat of the oxygen evolved, but those coals which are best to experiment with did not yield results that differed when the quantity of oxygen mixture was reduced to nearly the limit required for combustion of the coal. Under these circumstances, therefore, the constant use of 11 parts of oxygen mixture—a suitable quantity for all coals exported—would enable operators to obtain similar figures, and make the test uniform in different hands.

The following is a brief outline of the method of procedure recommended: Sample the coal until an average portion passes through a sieve having 64 meshes to the square inch. Take about 300 grains (20 grammes) of this and run through a brass wire gauze having 4,600 meshes to the square inch, taking care that the whole sample selected is thus treated. One part of nitrate of potash and 3 parts of chlorate of potash (dry) are separately ground in a mortar, and repeatedly sifted through another wire gauze sieve, having 1,000 meshes to the square inch, in order that the oxygen mixture shall not be ground to an impalpable powder, as this is very undesirable. It absorbs moisture rapidly, and interferes with the regularity of the combustion when very fine. 330 grains of the powder are weighed out (after drying), and intimately incorporated with 30 grains of coal—better with a spatula than by rubbing in a mortar—and then introduced into a copper cylinder (31/2 inches long by 3/4 inch wide, made from a copper tube), and pressed down in small portions by a test-tube with such firmness as is required by the nature of the coal, not tapped on the bottom, since the rougher portions of the oxygen mixture rise to the surface. As the temperature of a room is almost invariably much higher than the water supply, a little hot water is added to that placed in the glass cylinder, until the difference of temperature between the water and the room is about the mark indicated in the following table:

Room at The water should be

80 deg. F. 70 deg. F. 72 64 67 60 60 54 55 50 50 46 42 40

Say, for example, the room was at 57 deg. and the water placed in the cylinder was at 46 deg.: add a little hot water and stir with the thermometer until it assumes 52 deg.. By the time the excess of water has been removed with a pipette until it is exactly level with the mark, and all is ready, the temperature will rise nearly 0.5 deg.. Let the thermometer be immersed in the water at least three minutes before reading. The fuse should be placed in the mixture, and everything at hand before reading and removing the thermometer. After igniting the fuse and immersing the copper cylinder in the water, the apparatus should be kept in the best position for the gases to be evolved all around the cylinder, and the rate of combustion noted. Some coals are very unmanageable without practice, and samples of "patent fuel" are sometimes met with, containing unreasonable proportions of pitch, which require some caution in working and very close packing, inasmuch as small explosions occur during which a little of the fuel escapes combustion.

In order that the experiment shall succeed well, experience has shown that the nature of the fuse employed has much to do with it. Plaited or woven wick is not adapted, and will fail absolutely with dry coals, unless it is made very free burning. In this case not less than three-quarters of an inch in length is necessary, and the weight of such is very appreciable. I always use Oxford cotton, and thoroughly soak it in a moderately strong solution of nitrate of potash. When dry it should burn a little too fast. The cotton is rubbed between two pieces of cloth until it burns just freely enough; then four cotton strands are taken, twisted together, and cut into lengths of 3/4 inch and thoroughly dried. Open out the fuse at the lower end when placing it in the mixture so as to expose as much surface as possible in order to get a quick start, but carefully avoid pressing the material, and use a wire to fill up close to the fuse. A slow start often spoils the experiment, through the upper end of the cylinder becoming nearly filled up with potassic chloride, etc.

By paying attention to such details, and following the method recommended, the apparatus yields very satisfactory results with bituminous and semi-bituminous coals.—Chemical News.

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Words pass along with meanings which are simple conventionalities, marking current opinions, knowledge, fancies, and misjudgments. They attain to new accretions of import as knowledge advances or opinions change, and they are applied now to one set of ideas, now to another. Hence there is nothing truer than the saying, "definitions are never complete." The term explosion in its original introduction denoted the making of a noise; it grew to comprehend the idea of force accompanied with violent outburst; it is advancing to a stage in which it implies combustion as associated with destruction, yet somewhat distinct from the abstract idea of the resolution of any form of matter into its elementary constituents. The term, however, as yet takes in the idea of combustion as a decomposition in but a very limited degree, and it may be said to be wavering at the line between expansion and dissociation.

Strictly, in insurance, fire and explosion are different phenomena. A policy insuring against fire-loss does not insure against loss by explosion. It thereby enforces a distinction which exists, or did exist, in the popular mind; and fire, in an insurance sense, as distinct from explosion, was accurately defined by Justice McIlvaine, of the Supreme Court of Ohio (1872), in the case of the Union Insurance Company vs. Forte, i.e., an explosion was a remote cause of loss and not the proximate cause, when the fire was a burning of a gas jet which did not destroy, though the explosion caused by the burning gas-jet did destroy. Earlier than this decision, however (in 1852), Justice Cushing, of the Supreme Court of Massachusetts, in Scripture vs. Lowell Mutual Fire Insurance Company, somewhat anticipated later definition, and pronounced for the liability of the underwriter where all damage by the explosion involves the ignition and burning of the agent of explosion. That is, for example, the insurer is liable for damage caused by an explosion from gunpowder, but not for an explosion from steam. The Massachusetts Judge did not conceive any distinction as to fire-loss between the instantaneous burning of a barrel of gunpowder and the slower burning of a barrel of sulphur, and insurance fire-loss is not to be interpreted legally by thermo-dynamics nor thermo chemistry. While the legal principles are as yet unsettled, the tenor of current decisions may be summed up as follows: If explosion cause fire, and fire cause loss, it is a loss by fire as proximate cause; and if fire cause explosion, and explosion cause loss, it is a loss by fire as efficient cause. Smoke, an imperfect combustion, damages, in an insurance sense, as well as flame, which is perfect combustion; and where there is concurrence of expanding air with expanding combustion, the law settles on the basis of a common account. It's all "heat as a mode of motion."

Explosions are the resultants of elemental gases, vaporization, comminution, contact of different substances, as well as of the specifically named explosives. With new processes in manufacture, involving chemical and mechanical transformations, and other uses of new substances and new uses of old substances, explosions increase. The flour-dust of the miller, the starch-dust of the confectioner, increase in fineness and quantity, and they explode; so does the hop-dust of the brewer. In 1844, for the first time, Professors Faraday and Lyell, employed by the British government, discovered that explosion in bituminous coal mines was the quickening of the comparatively slow burning of the "fire-damp" by the almost instantaneous combustion of the fine coal-dust present in the mines. The flyings of the cotton mill do not explode, but flame passes through them with a rapidity almost instantaneous, yet not sufficient to exert the pressure which explodes; the dust of the wood planer and sawer only as yet makes sudden puffs without detonating force. Naphtha vapor and benzine vapor are getting into all places. One of the latest introductions is naphtha extracting oil from linseed, and then volatilized by steam superheated to 400 deg. F. This combination reminds us, as to effectiveness, of the combination at the recent Kansas City fire, when cans of gunpowder and barrels of coal oil both went up together.

But it is the unsuspected causes of explosion which make the great trouble, and prominent among these is conflagration as itself the cause of explosion, and such explosion may develop gases which are non-supporters of combustion as well as those which are inflammable. You throw table salt down a blazing chimney to set free the flame-suppressing hydrochloric acid, you discharge a loaded gun up a blazing chimney to put out the fire by another agency; still the salt, with certain combinations, may be explosive, a resinous vapor may be combustive in a hydrochloric atmosphere, and gunpowder isn't harmless when thrown upon a blaze—in fact, our common fire-extinguisher, water, has its explosive incidences as liquid as well as vapor.

Gases explosive in association may be set free by the temperature of a burning building and get together. In respect to the old conundrum, "Will saltpetre explode?" Mr. A. A. Hayes, Prof. Silliman, and Dr. Hare's views were, as to the explosions in the New York fire of 1845, that in a closed building having niter in one part and shellac or other resinous material in another, the gaseous oxygen generated from the niter and the carbureted hydrogen from the resins mingling by degrees would at length constitute an explosive mixture. A brief consideration of specific explosives uniting may serve to illustrate this phase of the subject.

Though the explosion of gunpowder is the result of a chemical change whereby carbonic acid gas at high tension is evolved (due to the saltpeter and the charcoal), the effect and rapidity of action are greatly promoted by the addition of sulphur. On the contrary, dynamite, now so important, and various similar explosives, are but mixtures of nitro-glycerine with earthy substances, in order to diminish and make more manageable the development of the rending force of the base. The explosive power of any substance is the pressure it exerts on all parts of the space containing it at the instant of explosion, and is measured by comparing the heat disengaged with the volume of gas emitted, and with the rapidity of chemical action. In the case of gunpowder, the proper manipulation and division of the grains is important, because favoring rapid deflagration; but in a purely chemical explosion, each separate molecule is an explosive, and the reaction passes from the interior of one to the interior of another, suddenly driving the atoms much further apart than their naturally infinitesimal vibrations.

Purely chemical explosives like nitro-glycerine, gun-cotton, the picrites, and the fulminates, present a terrible danger from the unknown mode of the new union of atoms, and reaction of the particles within themselves, in spontaneous explosions happening in irregular manner. Some curious circumstances attend the manufacture and use of gun-cotton,[1] nitro-glycerine, and dynamite. Baron von Link, in his system of the artillery use of gun-cotton, diminishes the danger of sudden explosion by twisting the prepared cotton into cords or weaving it into cloth, thereby securing a more uniform density. Mr. Abel's mode of making gun-cotton, which explosive is now used more than any other by the British government, includes drying the damp prepared cotton upon hot plates, freely open to the air. If ignited by a flame, however, in an unconfined place, gun-cotton only burns with a strong blaze, but if confined where the temperature reaches 340 deg. F., it explodes with terrific violence. Somewhat similar is the action of nitro-glycerine and dynamite, which simply burn if ignited in the open air, while the same substance will explode through a very slight concussion or by the application of the electric spark; a red-hot iron, also, if applied, will explode them when a flame will not. With care, nitro-glycerine can be kept many years without deterioration; and it has been heated in a sand-bath to 80 deg. C. for a whole day without explosion or alteration. One curious experiment is deserving of mention: If a broad-headed nail be partly driven into pine wood, and then some pieces of dynamite placed on the head of the nail, the latter may be struck hard blows with a wooden mallet without exploding the dynamite so long as the nail will continue to enter the wood.

[Footnote 1: The purest gun-cotton may be regarded as a cellulose, in which three atoms of hydrogen are replaced by three molecules of peroxide of nitrogen.]

Taking gunpowder as the unit, picrate of potash (picric acid and potassium) has five times more force, gun-cotton seven and a half times, and nitro-glycerine ten times more force. There are others still more powerful, but less known and used, and some explosives are quite uncontrollable and useless.

But the particular object of these remarks is to refer to articles of merchandise non-explosive under general conditions, but so in particular circumstances, as the two fire-extinguishers, water and salt, are explosive under given conditions. The memorable fire which, in July, 1850, destroyed three hundred buildings in Philadelphia, upon Delaware avenue, Water, Front, and Vine streets, was largely extended by explosions of possibly concealed or unknown materials, the presence of the generally recognized explosives being denied by the owners of the properties.

"The germ of the first knowledge of an explosive was probably the accidental discovery, ages ago, of the deflagrating property of the natural saltpeter when in contact with incandescent charcoal."[1] Although much manipulation is deemed necessary to form the close mechanical mixture of the materials of gunpowder, it has never been proved that such intimate previous union is necessary to precede the chemical reaction causing explosion; indeed, some explosions in powder works, before the mixture of the materials, or just at its commencement, seem to point to the contrary. It is also certain that in the manufacture of gunpowder the usual nitrate of potassium (saltpeter) can be replaced by the nitrates of soda, baryta, and ammonia, also by the chloride of potassium; charcoal by sawdust, tan, resin, and starch; and though a substitute for sulphur is not easily found, the latter, or a similar substance, is not an absolute necessity in the composition of gunpowder.[2]

[Footnote 1: Encyclopaedia Britannica, new edition, viii, p. 806.]

[Footnote 2: Vide Abel's Experiments in Gunpowder, as detailed in Phil. Trans. Eoy. Soc, 1874.—Vide also Bull. Soc. d'Encouragement, Nov., 1880, p. 633, Sur les Explosives.]

The generally received theory of the chemical action which makes gunpowder explosive is that it is due to the superior affinity of the oxygen of the niter (KNO3) for the carbon of the charcoal, and the production of carbonic acid gas (CO2) and carbonic oxide (CO) suddenly and in great volume. The latter extinguishes flame as well as the former, unless its own flammability is supported by the oxygen of the atmosphere until the degree of oxygenation CO2 is reached. Considering that water (H2O) is composed of two volumes of hydrogen and one of oxygen, and that under an enormously high temperature and the excessive affinity of oxygen gas for potassium or sodium (freed from nitrate union), dissociation of the water may be possible, aided by its being in the form of spray and steam, we would hesitate to deny that an explosive union of suitable crude salts could occur during the burning of a building containing them when water for extinguishment was put on. Any one who has seen the brilliance with which potassium and sodium burn upon water can easily imagine how such strong affinity of oxygen for these substances might aid in severing its union in water in their presence and under extraordinary heat. It might be safe so say that the presence of water under very high temperature may be as aidful to form an explosive among such salts as have been named, as sulphur is for the rapid combustion of gunpowder.

In the review for August, 1862 (Saltpeter Deflagrations in Burning Buildings and Vessels—Water as an Explosive Agency), it was shown that Mr. Boyden's experiments in 1861-62 proved that explosions would occur when water was put upon niter heated alone, and stronger explosion from niter, drywood, and sulphur; also explosion when melted niter was poured on water. The following points we reproduce for comparison: If common salt be heated separately to a bright heat, and water at 150 deg. F. poured on it, an explosion will occur. Niter mixed with common salt, placed upon burning charcoal, and water added, produce a stronger explosion than salt alone. Heating caustic potash to a white heat, and adding warm or hot water, produces explosion. At a Boston fire small explosions were observed upon water touching culinary salt highly heated. Anthracite coal and niter heated in a crucible exploded when sea water was poured on them.

The production of explosion by the putting of water on nitrate of potassium and chloride of sodium arises from the union, at high temperature, of the oxygen of the water with the potash and soda. Of the three liberated gases, hydrogen only is inflammable, and the other two suffocative of flame; but together the nitrogen and chlorine are not to be undervalued, for chloride of nitrogen is ranked as the most terrible and unmanageable of all explosives. Chlorine is a great water separator, but in the present case its affinity for hydrogen would result in hydrochloric acid, a fire extinguisher.

What happens in chemical experiment may be developed on a large scale in burning grocery, drug, or drysalters' stores, when great quantities of materials, such as just mentioned, including common salt, almost always present, are heated most intensely, and then subjected to the action of water in heavy dashes, or in form of spray or steam.

Picric acid, the nature of which we have several times previously mentioned, and which explodes at 600 deg. F. (only 28 deg. above gunpowder), may also be an element in such explosions during fires. Its salts form, in combinations, various powerful explosives, much exceeding gunpowder in force; and they have been used to a considerable extent in Europe. Picric acid, now much employed by manufacturers and dyers for obtaining a yellow color, is always kept in store largely by drysalters and druggists, and generally by dyers, but in smaller quantity.

In a very destructive fire which occurred in Liverpool, Eng., in October, 1874, involving the loss of several "fire-proof" stores, repeated explosions of the vapor of turpentine rent ponderous brick arched vaults, and exposed to the flames stocks of cotton, etc., in the stories above. This conflagration was started by the carelessness of an employee in snuffing a tallow candle with his fingers and throwing the burning snuff into the open bung-hole of a sample barrel of turpentine, of which liquid there were many hundreds of barrels on storage in the buildings. Turpentine vapor united with chlorine gas may not produce explosion, but by spreading flames almost instantly throughout the burning buildings, such burnings have practically equaled, if not excelled, explosions, which may sometimes be fire-extinguishers. In such cases detonation may be prevented by there being ample space to receive the suddenly ignited vapor, lessening the tension of it, but carrying the flames much more rapidly than otherwise to inflammable materials at great distance.

If disastrous results have arisen from the vapor of turpentine as a fire spreader in vaults without windows, it is possible that if a quantity of hot water were suddenly converted into steam in closely confined spaces, effects of pressure might be observed, less destructive perhaps, but resembling those which other explosives might produce. If the immense temperature attained in some conflagrations be considered—sufficient to melt iron and vitrify brick—it is possible to conceive of water as being instantly converted into steam. Even a very small quantity of water thus expanded could produce most disastrous results. While such formation of steam, if it happened, would certainly extinguish most flames in direct contact, the general phenomena shown would be explosive.

A curious circumstance occurred at the Broad street (N.Y.) fire in 1845, previously mentioned. The fire extended through to Broadway, and almost to Bowling Green. A shock like a dull explosion was heard, and by many this was attributed to the effects of gunpowder and saltpeter. Several firemen were, at the moment of the shock, on the roof of the burning building, when the whole roof was suddenly raised and then let down into the street, carrying the men with it uninjured. One of the firemen described the sensation "as if the roof had been first hoisted up and then squashed down." Query: Was this like the common lifting and falling back of the loose lid of a tea-kettle containing boiling water? Was it from steam—at a low pressure perhaps—seeking vent through the roof in like manner to the raising of the kettle-lid? Without dilating on this part of the subject, we mention it as a possible cause of minor explosions—doubtless to become better known in future. It may even be that explosions happening from steam acting in close spaces may have been attributed to gunpowder, or to niter and other salts, separate, but suddenly caused to combine in chemical reaction.—American Exchange and Review.

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By T.A. POOLEY, B.Sc., F.C.S.

This element, which next deserves our attention, is one of great importance and wide distribution; it occurs in nature in both the free and the combined states, and the number of compounds which it forms with other elements is very large. Unlike the previous elementary bodies we have studied, carbon is only known to us in the solid form when free, although many of its combinations are gaseous at the ordinary temperature and pressure. Carbon is known to exist in several different physical states, thus illustrating what chemists call allotropism, which means that substances of identical chemical composition sometimes possess altogether different outward and physical appearances. Thus the three states in which pure carbon exists, viz., diamond, graphite, or plumbago, and charcoal are as different as possible, and yet chemically they are all exactly the same substance. The diamond is the purest carbon, and occurs in the crystalline form known as a regular octahedron; the diamond is one of the hardest substances known, and is therefore, utilized for cutting glass; it has also a very high specific gravity, namely, 3.5, which means that it is three and a half times heavier than water, and it is far heavier than any of the other allotropic modifications of carbon. Graphite or plumbago, the second form in which carbon occurs, is widely distributed in nature, and the finer qualities are known as black lead, although no lead enters into their composition, as they are composed of carbon almost as pure as the diamond; the specific gravity of graphite is only 2.3. Charcoal, the third allotropic modification of carbon, is by far the most common, and is formed by the natural or artificial disintegration of organic matters by heat; we thus have formed wood charcoal, animal charcoal, lamp-black, and coke, all produced by artificial means, and we may also class with these coal, which is a natural product, and which contains from 85 to 95 per cent. of pure carbon.

Wood charcoal is made by heating wood in closed vessels or in large masses, when all the hydrogen, oxygen, and nitrogen are expelled in the gaseous state, and the carbon is left mixed with the mineral constituents of the wood; this form of carbon is very porous and light, and is used in a number of industrial processes.

Animal charcoal, as its name implies, is the carbonaceous residue left on heating any animal matters in a retort; and contains, in addition to the carbon, a large proportion of phosphates and other mineral salts, which, however, can be extracted by dilute acids. Animal charcoal possesses to a remarkable degree the property of removing color from solutions of animal and vegetable substances, and it is used for this purpose to a large extent by sugar refiners, who thus decolorize their dark brown sirups; in the manufacture of glucose and saccharums for brewers' use, the concentrated solutions have to be filtered through layers of animal charcoal in order that the resulting product may be freed from color. The decolorizing power of animal charcoal can be easily tested by any brewer, by causing a little dark colored wort to filter through a layer of this material; after passing through once or twice, the color will entirely disappear, or at all events be greatly reduced in intensity. Animal charcoal also absorbs gases with great avidity, and on this account it is utilized as a powerful disinfectant, for when once putrefactive gases are absorbed by it, they undergo a gradual oxidation, and are rendered innocuous, in the same way animal charcoal is a valuable agent for purifying water, for by filtering the most impure water through a bed of animal charcoal nearly the whole of the organic impurities will be completely removed.

Lamp-black is the name given to those varieties of carbon which are deposited when hydrocarbons are burned with an insufficient supply of oxygen; thus the smoke and soot emitted into our atmosphere from our furnaces and fireplaces are composed of comparatively pure carbon.

Coal is an impure form of carbon derived from the gradual oxidation and destruction of vegetable matters by natural causes; thus wood first changes into a peaty substance, and subsequently into a body called lignite, which again in its turn becomes converted into the different varieties of coal; these changes, which have resulted in the accumulation of vast beds of coal in the crust of the earth, have been going on for ages. There are very many different kinds of coal; some are rich in hydrogen, and are therefore well adapted for making illuminating gas, while others, such as anthracite, are very rich in carbon, and contain but little hydrogen; the last named variety of coal is smokeless, and is therefore largely used for drying malt.

Carbon occurs in nature also in a combined state; limestone, chalk, and marble contain 12 per cent. of this element. It is also present in the atmosphere in the form of carbonic acid, and the same compound of carbon is present in well and river waters, both in the free state and combined with lime and magnesia. All animal and vegetable organisms contain a large proportion of carbon as an essential constituent; albumen contains about 53 per cent., alcohol contains 52 per cent., starch 44 per cent., cane sugar 42 per cent., and so on. The presence of carbon in the large class of bodies known to chemists as carbohydrates, of which starch and sugar are prominent examples, can be easily demonstrated. If a little strong sulphuric acid be added to some powdered cane sugar in a glass, the mass will soon begin to darken in color and swell up, and in the course of a few minutes a mass of black porous carbon will separate, which can be purified from the acid by repeated washings; the sugar is composed of carbon, hydrogen, and oxygen, the two last-named elements being present in the exact proportion necessary to form water; the sulphuric acid having a strong affinity for water, removes the hydrogen and oxygen, and the carbon is then left in a free state.

Carbon forms two compounds with oxygen—carbon monoxide, commonly called carbonic oxide, and carbon dioxide, commonly called carbonic acid; and the last-named, being of most importance, will be studied first.

_Carbon Dioxide, or Carbonic Acid, Symbol CO_2_.—Carbonic acid occurs, as we have already stated, in large quantities in combination with lime and magnesia, forming immense rock formations of limestone, chalk, marble, dolomite, etc.; it also issues in a gaseous state from volcanoes, and it is always present in small quantities in the atmosphere; it is found dissolved in well and river waters, and it is a product of the respiration of animals. Brewers also are well aware of the existence of this body, for it is evolved in enormous quantities during the alcoholic fermentation of saccharine fluids. When carbonaceous substances are burnt the bulk of the carbon is converted into carbonic acid, and thus our furnaces and fireplaces are continually emitting enormous quantities of carbonic acid into the atmosphere. With these different sources of supply it might reasonably be thought that carbonic acid would be gradually accumulating in our atmosphere; the breathing of animals, the eruption of volcanoes, the combustion of fuel, and the fermentation of sugar, are ever going on, and to a fast-increasing extent with the progress of civilization, and yet the proportion of carbonic acid in our atmosphere is no greater now than it was at the earliest time when exact chemical research determined its presence and quantity. A counteracting influence is always at work; nature has beautifully provided for this by causing plants to absorb carbonic acid, holding some of the carbon, and allowing the oxygen to escape again into the atmosphere to restore the equilibrium of purity. This mutual evolution and absorption of carbonic acid is continually going on; occasionally there may be either an excess or a deficiency in a particular place, but fortunately any irregularity in this respect is soon overcome, and the air retains its original composition, otherwise animal life on the face of the globe would be doomed to gradual but sure extinction.

Carbonic acid can be prepared for experimental purposes by causing dilute hydrochloric acid to act upon fragments of marble placed in a bottle with two necks, into one neck of which a funnel passing through a cork is fixed, and into the other a bent tube for conveying the gas into any suitable receiver. The evolution of carbonic acid by this method is rapid, but easily regulated, and the gas may be purified by causing it to pass through some water contained in another two-necked bottle, similar to the generator. The chemical change involved in this decomposition is expressed by the following equation:

CaCO3 + 2HCl = CO2 + H2O + CaCl2 Calcium Hydrochloric Carbonic Water. Calcium Carbonate. Acid. Acid. Chloride.

By referring to the table of combining weights given in a previous paper, it will be seen that 100 parts of calcium carbonate will yield 44 parts of carbonic acid. Instead of hydrochloric acid any other acid may be used, and in the practical manufacture of carbonic acid for aerated waters sulphuric acid is the one usually employed. Carbonic acid is colorless and inodorous, but has a peculiar sharp taste; it is half as heavy again as air, its exact specific gravity being 1529; one hundred cubic inches weigh 47.26 grains. It is uninflammable, and does not support combustion or animal respiration. Under a pressure of about 38 atmospheres, at a temperature of 32 deg. F., carbonic acid condenses into a colorless liquid, which may also be frozen into a compact mass resembling ice, or into a white powder like snow. Carbonic acid is soluble in water, and at the ordinary pressure and temperature one volume of water will hold in solution one volume of the gas; under increased pressures, far larger quantities of the gas can be held in solution, but this is rapidly evolved as soon as the excess of pressure is removed. Upon this property the manufacture of aerated waters depends. The presence of free carbonic acid can be easily detected by causing the gas to pass over the surface of some clear lime-water. If any be present a white film of carbonate of lime will at once be formed. In testing carbonic acid in a state of combination, the gas must first be liberated by acting upon the substance with a stronger acid, and then applying the lime-water test. The presence of large quantities of carbonic acid in a gaseous mixture can be readily detected by plunging into the vessel a lighted taper, which will be immediately extinguished. This ought always to be adopted in a brewery, where many fatal accidents have happened through workmen going down into empty fermenting vats and wells without first taking this precaution.

The presence of carbon in this colorless gas can be demonstrated by causing some of it to pass over a piece of the metal potassium placed in a hard glass tube, and heated to dull redness; the potassium then eagerly combines with the oxygen, forming oxide of potassium, and the carbon is liberated and can be separated in the form of a black powder by washing the tube out with water.

Carbon Monoxide, or Carbonic Oxide. Symbol CO.—This is formed when carbon is burnt with an insufficient supply of oxygen, or when carbonic acid gas is passed over some carbon heated to redness. This gas is continually being formed in our furnaces and fire-places; at the lower part of the furnace, where the air enters, the carbon is converted into carbonic acid, which in its turn has to pass through some red-hot coals, so that before reaching the surface it is again converted into carbonic oxide; over the surface of the fire this carbonic oxide meets with a fresh supply of oxygen, and is then again converted into carbonic acid. The peculiar blue lambent flame often observed on the surface of our open fire-places is due to the combustion of carbonic oxide, which has been formed in the way we have just described. Carbonic oxide is a colorless, tasteless gas, which differs from carbonic acid by being combustible, and by not having any action on lime water.—Brewers' Guardian.

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The thermometers and pyrometers usually employed are almost all based on the expansion of some fluid or other, or upon that of different metals. The first can only be constructed with glass tubes, thus rendering them fragile. The second are often wanting in exactness, because of the change that the molecules of a solid body undergo through heat, thus preventing them from returning to exactly their first position on cooling.

The principle of the Seyfferth pyrometer is based on the fact that the pressure of saturated vapors, that is, vapors which remain in communication with the liquid which has produced them, preserves a constant ratio with the temperature of such liquid, while, on the other hand, the temperature of the latter when shut up in a vessel will correspond exactly with that of the medium into which it is introduced.

This instrument is composed of a metallic vessel or tube which contains the liquid to be exposed to heat, and of a spring manometric apparatus communicating with the tube, and by means of which the existing temperature is shown. The dial may be provided with index needles to show minimum and maximum temperatures, as well as be connected with electric bells (Fig. 1) giving one or more signals at maximum and minimum temperatures. The vessel to contain the liquid may be of any form whatever, but it is usually made in the shape of a straight or a bent tube. The nature of the metal of which the latter is made is subordinate, not only to the maximum temperature to which the apparatus are to be exposed, but also to the nature of the liquid employed. It is of either yellow metal or iron. To prevent oxidation of the tube, when iron is employed, it is inclosed within another iron tube and the space between the two is filled in with lead. When the apparatus is exposed to a high temperature the lead melts and prevents the air from reaching the inner tube, so that no oxidation can take place.

Pyrometers filled with Ether.-These are tubular, and constructed of yellow metal, and are graduated from 35 deg. C. to 120 deg.. They are used for obtaining temperatures in vacuum apparatus, cooking apparatus, diffusion apparatus, saturators, etc. Figs. 2, 3, 4, and 5, show the different modes of mounting the apparatus according to the purpose for which it is designed.

Pyrometers filled with distilled water are used for ascertaining temperatures ranging from 100 deg. to 265 deg. C., 80 deg. to 210 deg. R., or 212 deg. to 510 deg. F.

Pyrometers filled with mercury are constructed for ascertaining temperatures from 360 deg. to 750 deg. C.


The temperature necessary for the complete carbonization of the organic substances of animal charcoal is from 430 deg. to 500 deg. C. In order to transmit this temperature from the cylinder to the charcoal it is indispensable that the air surrounding the cylinder be heated to 480 deg. to 550 deg.. If the heating of the animal black exceeds 500 deg. the product hardens, diminishes in volume, and loses its porosity. There are two methods of ascertaining the temperature of the red-hot bone black by means of the pyrometer: First, by inserting the tube of the instrument into the black. (Fig. 6, a.) Second, by finding the temperature of the hot gases in the furnaces (Fig. 6, b.). In the first case, the plunge tube should be of sufficient length to allow its extremity to penetrate to the very bottom layer of the red-hot black. This mode of direct control of the temperature of the black is only employed for ascertaining the work accomplished by the furnace, that is to say, the ratio existing between the temperature of the hot air surrounding the cylinder and the black itself. This calculation being effected, it is useless to note the differences of temperature which arise in the spaces between the cylinders of which the furnace is composed.

The position that the pyrometer should occupy is subordinate to the construction of the furnace. Fig. 6 shows the type which is most employed.

In a furnace with lateral fire-place, cc are the heating cylinders, and dd the cooling cylinders. C D is the plate on which are mounted vertically the former, and from which are suspended the latter, b shows the pyrometer, the length of which must be such that the manometric apparatus shall stand out one or two inches from the external surface of the wall, while its tube, traversing the wall, shall reach the very last row of heating cylinders.

That the apparatus may form a permanent regulator for the stoker it is well to adapt to it an arrangement permitting of a graphic control of the work accomplished and signaling by means of an electric bell when the temperature of the gases in the furnace descends below 480 deg. C. or rises above 550 deg. C.


The operation of heating brick furnaces is generally performed according to empirical methods, the temperature having to vary much according to the products that it is desired to obtain. It is necessary, however, for a like product to maintain as uniform a temperature as possible. These observations are particularly applicable to continuous furnaces such as annular brick furnaces, etc., in which a uniformity of temperature in the different chambers is of vital importance to perfect the baking. In these furnaces the tube of the pyrometer is inserted through one of the apertures at the top, as shown in Fig. 7. The dial is graduated up to 750 deg., which is more than sufficient, since the temperature of the upper part of a compartment fully exposed to the heat rarely exceeds 670 deg. to 680 deg. C.

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Potash soaps are generally superior to soda soaps for most purposes, but more especially in washing wool and woolen goods. The difference between the use of a potash and a soda soap for these purposes is very marked. Potash lubricates the fiber of the wool, renders it soft and silky, and to a certain extent bleaches it; soda, on the other hand, has a tendency to turn wool a yellow color, and renders the fiber hard and brittle. It cannot be too strongly insisted upon, therefore, that nothing but a potash soap (or some form of potash in preference to soda if an alkali alone is employed) should be used in washing wool in any form—either manufactured or unmanufactured. This is fully borne out by nature, who invariably assimilates the most appropriate substances. Wool when growing in its natural state is lubricated and protected by a sticky substance called "grease" or "suinte;" this consists to the extent of nearly half its weight of carbonate of potash, hardly a trace of soda being present. It is very evident, therefore, that potash must be more suitable for washing wool than soda, as the teaching of nature is always correct.

There are certain prejudices against the use of potash soap, which have, to a great extent, prevented its more extensive use. Many consumers of soap fancy that because a potash soap is soft it necessarily must contain more water than a soda soap; this, however, is quite an erroneous notion. A potash soap is soft, because it is the nature of all potash soaps to be so, just in the same way that on the other hand all soda soaps are hard. As an actual fact a good potash soap contains less water than many quite hard soda soaps that are now in the market. Another reason is that soapmakers have had every interest in using soda in preference to potash—particularly when latterly soda has been so cheap.

Potash not only is a more expensive alkali, but its combining equivalent is greatly against it as compared with soda; that is to say, that thirty-one parts of actual or anhydrous soda will saponify as much tallow or oil as forty-seven parts of anhydrous potash. It will be evident, therefore, that the use of potash instead of soda is decidedly more advantageous to the soapboiler, and more particularly in the present age, when the demand is for cheap articles, often quite without regard to the quality or purpose for which they are to be used. As far as consumers are concerned, this has been a mistake. Potash soap, though it may cost more, is in most cases actually the most economical. Soap is never used in exact chemical equivalents, but an excess is always taken. Potash soap is much more soluble than a soda soap; it therefore penetrates the fiber, and consequently removes dirt and grease much more quickly. Notwithstanding, also, that its chemical combining equivalent is greater than that of soda, it is, nevertheless, the strongest base, and always combines with any substance in preference to soda. For these reasons—probably combined also with the fact that in the whole realm of the animal and vegetable kingdoms, to which all textile fabrics belong, potash is more naturally assimilated than soda—a smaller quantity of potash soap will do more practical work than a larger quantity of soda soap.

There are other reasons why potash soaps have not been used; originally soft soap was made either with fish oil or olive oil. Fish oil is objectionable, as the strong smell imparted to the soap renders it unfit for many finishing purposes. Nothing can be better than olive oil soap, but it is a costly article, and only can be used for finer purposes. There are now, however, many of the seed oils that are much cheaper. Linseed, rape seed, and cotton seed all produce a good soap. Cotton seed oil is particularly suitable for the purpose; the manufacture of this oil during the last few years has been brought to great perfection, and the cost is now much less than that of tallow or of any other seed oil. It is now difficult to distinguish a well refined cotton seed oil from olive oil; it is therefore in every way suitable for making soft soap. One of the chief causes, however, why potash soap has not been more generally made is that a convenient form of potash has been unobtainable. For many years the only source of potash was from the ashes of burnt trees. These ashes are collected, mixed with lime, lixiviated, and the resulting lye boiled down. The result is a very impure form of potash, also of a very variable composition, depending upon the trees used for the purpose. Canada has been the principal source of supply of this form of potash; hence the commercial name of Montreal potashes. The classification of "firsts," "seconds," and "thirds" is from the inspection at the warehouse there; this, however, is exceedingly superficial, the ashes being simply tested for their alkaline strength, with no discrimination between potash and soda, which is a difficult and delicate chemical test. Soda being now far cheaper than potash, and also the alkaline equivalent, as previously explained, being greatly in favor of soda, there has been every inducement to "enterprising" producers of ashes to adulterate them with soda, which, in many cases, has been largely done. Another source of potash has been beetroot ashes, very similar to wood ashes, and also German carbonate of potash, which latter about corresponds to a common soda ash, as compared with caustic soda; with these articles, a tedious boiling process, very similar to the old process for the production of hard soap, had to be adopted, the ashes, or carbonate of potash, previously being dissolved and causticized with lime by the soap maker. The production of a first-class soft soap was also a very difficult operation, as the impurities and soda contained varied considerably, often causing the "boil" to go wrong and give considerable trouble to the soapboiler.

During the last two years, however, caustic potash has been introduced, that manufactured by the Greenbank Alkali Co., of St. Helens, being very nearly pure. With this article there is no difficulty in producing a pure potash soap, either for wool scouring, fulling, or sizing, by a cold process very similar to that described for the production of hard soda soap with pure powdered caustic soda.

The following directions will produce an excellent soap for wool scouring: Fifty pounds of Greenbank pure caustic potash are put into eight gallons of soft water; the potash dissolves immediately, heating the water. This lye is allowed to cool, and then slowly added, with continual mixing, to 20 gallons of cotton seed oil, mixed with 20 pounds of melted tallow, the whole being brought to a temperature of about 90 deg. F. After stirring for some minutes, so as to completely combine the lye and oil, the mixture is left for two days in a warm place, when a slow and gradual saponification of the mass takes place. If when examined the oil and lye are then found not completely combined, the stiff soap is again stirred and left two days, when the saponification will be found complete, the result being the formation of about 330 pounds of very stiff potash soap, each pound being equal to about two pounds of the ordinary "fig" soap sold. The requisite quantity is thrown into the scouring vat with about five per cent of its weight of refined pearl ash to increase the alkali present, the weight depending somewhat upon the kind of wool washed on purpose for which the soap is required. If the wool is very dirty or greasy, rather a stronger soap is sometimes advisable. This can easily be attained by reducing the quantity of oil used to 18 gallons.

The advantages to be gained by the wool scourer or other consumer making his own potash soap are that a pure, uniform article can always be thus produced at a less cost than that at which the soap can be bought. Potash soap, like soda soap now sold, is much adulterated, in addition to all the impurities originally contained in the potash used, and which, unlike soda soap, cannot be separated by any salting process. Many other adulterations are added to increase the weight and cheapen the cost. Silicate of potash, resin, and potato flour are all more or less employed for this purpose, to the gain of the soap maker and at the expense of the consumer.

The production of potash soap for fulling and sizing, and the most suitable oils and tallow for the production of the various qualities required for these purposes, must be reserved for the next issue.—Textile Manufacturer.

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We have, on a previous occasion, described the process of "maceration" or "enfleurage," that is, the impregnation of purified fat with the aroma of certain scented flowers which do not yield any essential oil in paying quantities. At present we wish to describe an apparatus which is used in several large establishments in Europe for obtaining such products on the large scale and within as short a time as possible. The drawing gives the idea of the general arrangement of the parts rather than the actual appearance of a working apparatus, for the latter will have to vary according to the conveniences and interior arrangements of the factory.[1]

[Footnote 1: Our illustration has been taken from C. Hofmann, "Chemisch-technisches Universal-Receptbuch," 8vo, Berlin, 1879, p. 207.]

A series of frames with wire-sieve bottoms are charged with a layer of fat in form of fine curly threads, obtained by pressing or rubbing the fat through a finely-perforated sieve. The frames are then placed one on top of the other, and to make the connection between them air-tight, pressed together in a screw press. A reservoir, E, is charged with a suitable quantity of the flowers, etc., and tightly closed with the cover, after which the bellows are set into motion by any power most convenient. Scented air is thereby drawn from the reservoir, E, through the pipe, G B, toward the stack of frames containing the finely divided fat, which latter absorbs the aroma, while the nearly deodorized air is sent back to the reservoir by the pipe, D, to be freshly charged and again sent on its circuit. This apparatus is said to facilitate the turning out of nearly twenty times the amount of pomade for the same number of frames and the same time, as the old process of "enfleurage." It might be called the "ensoufflage" process.—New Remedies.

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At a recent meeting of the London Chemical Society, Mr. W. Jago read a paper "On the Organic Matter in Sea-water." On p. 133 of the "Sixth Report of the Rivers Commission," it is stated that the proportion of organic elements in sea-water varies between such wide limits in different samples as to suggest that much of the organic matter consists of living organisms, so minute and gelatinous as to pass readily through the best filters. At the suggestion of Dr. Frankland, the author has investigated this subject. The water was collected in mid-channel between Newhaven and Dieppe by the engineers of the London, Brighton, and South Coast Railway in stoppered glass carboys. The author has used the combustion method, the albuminoid ammonia, and in some cases the oxygen process of Prof. Tidy. To determine how the various methods of water-analysis were effected by a change of the organic matter from organic compounds in solution to organisms in suspension, some experiments were made with hay-infusion. The results confirm those of Kingzett (Chem. Soc. Journ., 1880, 15). the oxygen required first rising and then diminishing. The author concludes that the organic matter of sea-water is much more capable of resisting oxidizing agents than that present in ordinary fresh waters, and that the organic matter in sea-water is probably organized and alive.

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W. M. Hamlet, in a paper before the London Chemical Society, said: Flasks similar to those of Pasteur ("Etudes sur la Biere," p. 81), holding about 1/4 liter, were used. The liquids employed were Pasteur's fluid with sugar, beef-tea, hay infusion, urine, brewers' wort, and extract of meat. Each flask was about half filled, and boiled for ten minutes, whereby all previously existing life was destroyed. The flask was then allowed to cool, the entering air being filtered through a plug of glass wool or asbestos. The flask was then inoculated with a small quantity of previously cultivated hay solution or Pasteur's fluid. Hydrogen, oxygen, carbonic oxide, marsh-gas, nitrogen, and sulphureted hydrogen, were without effect on the bacteria. Chlorine and hydric peroxide (about 7 per cent, of a 5 vol. solution) were fatal to bacteria. The action of various salts and organic acids in 5 per cent, solution was tried. Many, including potash, soda, potassic bisulphite, sodic hyposulphite, potassic chlorate, potassic permanganate, oxalic acid, acetic acid, glycerin, laudanum, and alcohol, were without effect on the bacterial life. Others—the alums, ferrous sulphate, ferric chloride, magnesic and aluminic chlorides, bleaching powder, camphor, salicylic acid, chloroform, creosote, and carbolic acid—decidedly arrested the development of bacteria. The author has made a more extended examination of the action of chloroform, especially as regards the statement of Muentz, that bacteria cannot exist in the presence of 21/2 per cent, of chloroform, which substance is therefore useful in distinguishing physiological from chemical ferments. The author concludes that amounts of chloroform, phenol, and creosote, varying from 1/4 to 3 per cent., do not destroy bacteria, although their functional activity is decidedly arrested while in contact with these reagents. To use the author's words, bacteria may be pickled in creosote and carbolic acid without being deprived of their vitality. The author concludes that the substances which destroy bacteria are those which are capable of exerting an immediate and powerful oxidizing action, and that it is active oxygen, whether from the action of chlorine, ozone, or peroxide of hydrogen, which must be regarded as the greatest known enemy to bacteria.

Mr. Hamlet, in replying to some remarks of Messrs. Kingzett and Williams, said that in all cases the solution which he had used had been completely sterilized by exposure to a temperature of 105 deg. for ten minutes. The India-rubber tubing he had used was steamed. Carbolic acid solution must contain at least 5 per cent, of carbolic acid to be fatal to bacteria. He was quite aware of the importance of distinguishing between the action of the substances on various kinds of bacteria, and was quite prepared to admit that a treatment which would be fatal to one kind of bacterium might not injure another.

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