oxidised, add water and filter. Evaporate the filtrate until fumes of sulphuric acid are again visible, when more water must be added, and the liquid again filtered if necessary. To the clear solution now add tartaric acid, then ammonia in excess and sulphide of ammonium. Evaporate the liquid, containing the precipitate suspended in it, in a dish, until all the smell of sulphide of ammonium has disappeared. Filter, evaporate to dryness, and ignite to get rid of the organic matter. Powder the black ash, boil it in moderately strong hydrochloric acid, filter, add a crystal of chlorate of potash, and boil for a minute. Now add chloride of ammonium and ammonia, and boil for five minutes. If, at the end of that time, any precipitate is observed it will be alumina. From the filtered solution, if oxalate of ammonia be added, the lime will be precipitated; and if to the filtrate from this, ammonia and phosphate of soda be added, the magnesia will come down." PRELIMINARY INVESTIGATION OF THE By HENRY MORTON, Ph.D., Double Sulphates. When solutions of equivalent weights of uranic sulphate and rubidium sulphate are brought together, and the proper degree of concentration has been reached, warty concretions of minute crystals form, possessing a hue of remarkable beauty and a brilliant fluorescence. So strong is this that the white porcelain dish in which the crystals form appears pink by contrast. Thallio-uranic sulphate forms in a precisely similar manner, but is of a golden yellow colour with a fine lustre, but very slight fluorescence. It is not readily soluble in water, but is very stable, being easily re-crystallised from hot solution. Analysis of these salts has not been completed in time for insertion in this preliminary notice. We will now pass to the optical study of these salts in their alphabetical order. Ammonio-Uranic Sulphate, U2O3SO3+NH2OSO3+2HO. -The spectrum of this salt has been already described, but for completeness we will here reproduce it with some additional data. The bands of this substance are distinguished by great abruptness on their more refrangible side, rising as it were suddenly to a narrow brilliant line, and then fading off gradually on the lower side in a Their manner suggesting a rounded convex surface. positions are shown in 1 of Fig. 20. I In the first experiments made with this salt it was heated to about 200° C. to drive off its water, but subsequent experience showed that the same result might be reached by a continued application of a temperature of 100° C. Two atoms of water in this salt would amount to 671 per cent; and it was found that a specimen Uranic sulphate has heretofore been known to form the bottle placed in a hot-water oven continued to lose weight following double salts: U203SO3 + NH4OSO3+2HO. The last-named was obtained by Berzelius, but Ebel. men and Péligot could not reproduce it. To this list we have added the following new compounds: 2(U2O3SO3)+NH40 SO3. U2O3SO3 + MgO U2O3SO3 + MgO SO3+7HO. (?) SO3+4HO. SO3+2HO. The history of the ammonio-diuranic sulphate has been already given in a previous part of this paper, and we need only add here that we have as yet been unable to prepare it otherwise than by the decomposition of the for about twelve hours. At the end of this time its total Ammonio-Diuranic Sulphate, 2(U2O3SO3)+NH4OSO3. Ammonio-uranic sulphate. Magnesio-uranic sulphate. Rubidio-uranic sulphate. Sodio-uranic sulphate. ammonio-sulphate as there described. Both the magnesium salts were formed under conditions seemingly identical, when a cold solution of the mixed sulphates was allowed to concentrate over sulphuric acid. * Communicated by President Morton. without further decomposition, yields a spectrum such as is represented in 5 of Fig. 20, and whose bands are as bright and sharply defined on the upper edge as those of the hydrated ammonio-sulphate; from which they are to be distinguished only by their position, and perhaps by their greater breadth descending more gradually into the NEWS dark spaces on the lower side. The bands near 70 and 80 of the cut are not correctly shown, the first being too low, and the second too high. They should show an even spacing, one having its upper edge at 70.5 and the other at 80. The absorption spectrum of this salt is shown at 2 of Fig. 19. Magnesio-Uranic Sulphate.-As we have already noticed, a mixture of uranic and magnetic sulphates will form two compounds, one of these, which seems easy to reproduce, having the formula U2O3SO3+ MgOSO3+4HO. The other we have only succeeded in obtaining once in small quantity. This and the fact that it was then mixed with magnesium sulphate, render an accurate analysis impossible, but a determination of the water and sulphuric acid seems to indicate that the body has the formula U203SO3+MgOSO3+7HO. We would, however, only present this as a suggestion. The two salts yield fluorescent spectra, which are entirely distinct in the positions of their bands, although both are alike, and of what we have called the normal form in their general character. The following table will show this 404 480 55'5 65'7 746 84.8 916 448 530 664 708 810 90'5 96'0 The measures here given are of the upper edges of bands' except the 8th, which is measured at its centre. The absorption spectrum of the normal salt, A, is given at 3 of Fig. 22, and is characterised by the absence of the lower bands found in other double sulphates. Potassio-Uranic Sulphate, U2O3SO3+KO,SO3+ +2HO. This salt readily crystallises out when a solution of the mixed sulphates in atomic proportions is allowed to evaporate in the air; it forms warty concretions of minute crystals of a yellow-green colour and very bright fluorescence. In this state it contains 2 atoms of water, and shows the spectrum represented in I of Fig. 24, which is of the normal character, with a sharp termination of each band on its upper edge, and the bands peculiarly broad and bright. This salt suffers no change by heating or drying at 100° C., but if dried at 150° C. it loses all its water, and its spectrum changes to that shown at 2 of Fig. 21, in which the bands are much rounded, and displaced a little upward in the spectrum. It would thus appear that this salt can exist, and display fluorescent action as a bihydrate and anhydrate, but we have obtained no evidence of its forming a monohydrate. The absorption spectrum of the hydrated salt will be seen at 1 of Fig. 19, being indistinguishable from that of the ammonio-salt. The spectrum of the anhydrate is more difficult to observe, but by using the substance in powder with a little oil between slips of glass we can make out bands whose centres are at 93'7, 101'O, III7, and 1244 respectively, and which are therefore quite unlike those of the normal salt. Rubidio-Uranic Sulphate, U2O3SO3+RbOSO3+ +2HO.-The preparation of this salt we have already described, and we will therefore pass at once to its fluorescent spectrum. In this the bands are much blended or rounded, and are decidedly lower than the corresponding ones of the potassium salt. Their brightest parts are located as follows: Fluorescent Spectra of Rubidio-Uranic Sulphate. 3. 4. 5. 6. 7. 8. 340 416 492 576 667 76.0 85.2 94'8 This substance loses 2 atoms of water if it is dried at 100° C., but its fluorescent spectrum is not changed as regards the position of its bands; the brightness of its fluorescence is, however, reduced. An exposure to a yet higher temperature causes (180° C.) slight loss of weight (0.6 of 1 per cent), and yet further reduces its fluorescence without other effect. The absorption spectrum of the normal salt will be found at 4 of Fig. 19. a Sodio-Uranic Sulphate, U203SO3+NaOSO3+5HO.This salt presented more difficulties at the outset than any other, and, as might be expected, has yielded some very curious results. Thus a certain specimen was observed to yield the peculiar spectrum shown at 2 of Fig. 22, while the rest of the crop of crystals in another bottle from which it had been taken showed nothing of the sort. A prolonged study of this body has evolved the following facts, and has shown that many more await further investigation:-This salt in its normal state, containing 5 atoms of combined water yields the spectrum shown at 1 of Fig. 19, which is | of the normal type. A portion of this water is, however, lost with great ease, occasioning the formation, under ordinary conditions, of mixtures of several hydrates. Some of these we have been unable to isolate and determine, but we have found that, by gradually drying at a temperature of 150° C., we obtain a monohydrate whose spectrum is that shown at 3 of Fig. 19. If the salt in a damp state is placed suddenly in the oven at 150° C., it will lose almost all its water, and give a spectrum sensibly continuous; in this condition it may be dried at 200° C. without suffering any change. When heated to 250° to 290° the salt loses all its water, and then acquires a spectrum such as is shown at 4 of Fig. 19, in which each band looks like a prismatic column. Under various conditions, which we have not yet been able to determine with certainty, intermediate amounts of water are lost, and mixtures of other hydrates are produced, in which, no doubt, salts with 2, 3, or 4 atoms of water are involved. These, once formed, will, like the corresponding hydrates of uranic sulphate, maintain themselves in the presence of desiccating treatment, which would deprive the normal salt of much more water. The absorption spectrum of the normal salt is shown at 5 of Fig. 19. It is, perhaps, unnecessary to state that the peculiar spectrum shown at 2 of Fig. 22 is believed to be produced by three or more overlapping spectra belonging to as many mingled hydrates. Thallio-Uranic Sulphate, U203SO3+TIO,SO3+3HO.This substance has a very faint fluorescence. Bands can be made out at 92'4 (?), 35·6, and 76'0; below the first of these three seems to be a faint continuous spectrum. This salt loses 3HO and all its fluorescence by drying at 100° C. Its absorption spectrum is shown at 6 of Fig. 19. NOTES OF WORK BY STUDENTS OF PRACTICAL CHEMISTRY IN THE LABORATORY OF THE (No. II.) Communicated by J. W. MALLET, Professor of General and Applied Chemistry in the University. (1). On the Best Mode of Converting Calcium Oxalate into Carbonate in the course of Analytical Work. By Mr. J. R. McD. IRBY, of New Orleans, Louisiana. CALCIUM precipitated as oxalate is sometimes weighed as such after drying at 100° C., is sometimes converted into carbonate by heating to a carefully regulated temperature just short of redness, sometimes converted into lime by heating to bright redness or beyond, and sometimes given the form of sulphate by treatment with sulphuric acid or ammonium sulphate. Of these four modes of procedure, although accurate results can be obtained by any one of them, the second is, on several grounds, to be preferred for general use. If the oxalate itself be weighed, an unburnt filter, previously tared when dried at 100° C., has to be weighed with it, and the errors which may be allowed to arise from aygroscopic moisture affecting this filter at either of its weighings, and the tube or watch-glasses used to contain it are more likely to influence the result to a serious extent than those from similar causes when but the ash of a filter and a small crucible are concerned. The conversion into caustic lime requires a very strong heat, generally obtained by means of a blast lamp, the blowing being kept up for some time. The last traces of carbon dioxide are driven off with difficulty, and the platinum crucible is liable to alter slightly in weight, while the risk of mechanical loss from the blast, and the great readiness with which moisture is taken up from the * See paper by Aug. Souchay in Fresenius's Zeitschr. f. Anal. Chem.' 10 jahrg., 3 heft, s. 323, and remarks upon same by Fresenius in same periodical, s. 326. atmosphere by the lime during cooling and weighing, are not to be altogether overlooked. The acid fumes given off during the conversion into sulphate and final evaporation are annoying, and the. heating is tedious and requires very careful watching to prevent loss. Pure calcium carbonate is undoubtedly the most stable and desirable form in which to obtain the final product for weighing, and if at once obtained by carefully managed heating of the oxalate, as may be accomplished in welltrained hands, leaves nothing to be desired; but if the temperature be allowed to rise a little too high, and a little caustic lime be formed, the necessary evaporation with solution of ammonium carbonate is tedious and troublesome, and cannot be hastened without almost certain loss from spirting. Moreover, while a very small quantity of lime can thus easily be restored to the condition of carbonate, if any considerable amount of material have to be dealt with the moistening and evaporation will often need to be repeated more than once before the weight becomes constant. The object in heating is, therefore, to so regulate the temperature as to ensure the complete destruction of all oxalate, and to avoid altogether the decomposition of the carbonate. The statement of the writers on analytical chemistry, that the proper temperature is represented by very low redness, or should be just short of redness, is wanting in precision. Working at night or in the daytime, by bright sunlight or on a dark, cloudy afternoon, one's estimates of barely visible redness will represent by no means a small range of temperature, and it needs a good deal of personal supervision to teach a new laboratory student exactly how to proceed, so as to obtain at once, without the delays above referred to, a result so often called for as an accurate determination of calcium. In order to simplify and give greater precision to the details of the process, it was suggested to Mr. Irby to get some better measure of the temperature really required, and to find, if possible, some empirical rule for its ready production and regulation. Not having at command one of Siemens's electric resistance pyrometers, it was attempted to estimate the temperature needed for the decomposition of the oxalate by comparison with the melting-points of some of the metals. Calcium oxalate (specially prepared and found to be quite pure, and fully dried at 100° C.) heated to the melting-point of lead for a short time, was found, on cooling, to have begun to decompose, but the extent of the change was quite small. At just the melting-point of zinc the weight could easily be reduced to within one or two per cent of the theoretical amount, but several hours' exposure to this temperature scarcely produced complete reduction to carbonate. Several metallic alloys were tried, in order to get from some one of them a barely-fused bath of the proper temperature, but they all gave too much trouble from surface oxidation and the tendency to separate into a more and a less fusible portion. Finally, it being ascertained that a heat but very little beyond the melting-point of pure zinc was required, the following arrangement was found practically successful. A solid cylinder of cast-iron, smoothly turned, 53 m.m. high and 66 m.m. in diameter, had a cylindrical hole of 40 m.m. deep and 40 m.m. diam., drilled into the upper end, thus producing a sort of crucible with walls and bottom 13 m.m. thick. A turned disc of cast-iron, 58 m.m. in diameter and 6.5 m.m. thick, with a little knob in the middle of the upper surface to serve as a handle, formed a cover; and in the upper surface of this cover, two hemispherical cavities, each 10 m.m. in diameter, were drilled at opposite sides, the centre of each 15 m.m. distant from the edge. Round the outside of the cylinder a little groove was turned, which enabled the whole to be supported over a lamp by a stout iron wire triangle. Fresenius, in his excellent "Anleitung zur Quant. Chem. Anal.,' 5 aufl., s. 201, has given minute directions as to the details. The figure shows this thick walled crucible and cover in vertical section. The total weight was about 1050 grms. In the little cavities (a a) bits of metallic zinc of about a gramme each were placed, a piece of porcelain, such as a small crucible cover, was placed in the bottom of the cast-iron vessel, and upon this a platinum or porcelain crucible (the latter being found to answer best, on account of its inferior conducting power) containing the calcium oxalate to be heated. Gas from a Bunsen burner with tube of 9 m.m. diameter was used as the source of heat, and the position of the burner was so regulated that, when the stopcock was fully opened, the fame played a over the bottom of the cast-iron block and about one fourth up the outside all round. With quantities of I to 15 grm. of oxalate, the full flame of the lamp was turned on at once, the mass of iron in the block ensuring sufficiently gradual heating, and it was then only necessary to notice when the bits of zinc in the cavities of the iron cover had fully melted; the decomposition was then complete. This took about thirty minutes, but during that time no attention on the part of the operator was needed. The calcium carbonate left in the crucible was quite free from caustic lime. The foliowing are two examples of the results : Calcium oxalate taken. Grm. 1'0388 1'5765 Calcium Calcium carbonate carbonate obtained. calculated. Grm. Grm. 0'7109 0'7115 I'0792 1'0798 Lime (derived Lime calcu gradually and to keep up the heat longer, about three quarters of an hour being required, while, for quantities of 3 grms., an hour was necessary. In the course of actual analysis, the oxalate being upon a filter, as much as possible of the substance should be detached from the paper and treated as above described, while the filter itself is burnt upon the lid of the platinum or porcelain crucible, the minute residue left treated with one or two drops only of strong solution of ammonium carbonate, which small quantity can be dried up quite quickly and easily on a water-bath or gently heated sandbath, and the cover then introduced into the cast-iron cylinder along with the crucible; the two to be taken out, cooled, and weighed together. (old atomic weights of copper and oxygen), or 3 molecules of cupric oxide, 1 of cupric chloride, and 4 of water, the author remarking that one or two analyses give one-half more water. Rammelsberg ("Handb. d. Mineralchem.," s. 191) makes three varieties of the mineral, each containing 3 molecules of cupric oxide and 1 of the chloride, but therewith 3, 4, and 6 molecules of water respectively. In view of this discrepancy of results as to the amount of water, and there seeming to be no recorded analysis of Australian atacamite, a portion of a very fine, wellcrystallised specimen, of rich dark green colour, from South Australia, was placed in the hands of Mr. Cabell for analysis. Sp. gr. of mineral =4'314. The following were the results, neglecting traces of insoluble siliceous residue, ferric oxide, and alumina :— : formulaThe figures in the last column are calculated from the 3CuH2O2+CuCl2, or Cu2 ¡ (HO)3, from oxalate) lated directly which probably represents the composition of normal as calculated from carbonate. Per cent. 38.32 38.33 from oxalate. 38.36 It proved to be important that the decomposition of the greater portion of the oxalate should take place without the carbon monoxide gas given off taking fire, since the additional heat produced by its combustion through and on the mass sufficed to burn a little of the carbonate into lime, and the tendency to separation of carbon (rendering the mass dark in colour) was much greater when the gas took fire, while carbon once separated could not be well burnt off again without increasing the heat too much. The effect of too rapid heating, attended with the burning off of carbon monoxide, is shown by the following results, obtained in experiments in which the maximum temperature derived from the lamp was no higher than in others of entirely satisfactory character : Percentage of Lime obtained from Oxalate. 38.18 38.19 Calculated. } 38.36 Hence, when the oxalate amounted to about 2 grms., it was found better to turn on the gas to the lamp somewhat the only noteworthy point being the replacement of a little of the antimony by arsenic. This does not seem to have been before observed for bournonite, although so common in other minerals containing the sulphides of these metals. The perfectly crystallised condition of the specimen examined negatives the idea of an admixture with tetrahedrite. (4). Analysis of an Iron Slag of Fine Blue Colour from Barrow Iron Works, Lancashire (England). By Mr. J. R. McD. IRBY. This slag was very compact, tenacious, and harddeclared by a lapidary to be fully equal in hardness to the most refractory jasper; fracture splintery and imperfectly conchoidal. Its colour was a remarkably fine blue, quite comparable with the darker varieties of lapis lazuli, so that a specimen cut en cabochon made a handsome stone for a finger ring. Sp. gr. = 2.883. When reduced to an impalpable powder, it was completely decomposed by strong hydrochloric acid, with loss of colour, evolution of hydrogen sulphide, and separation of flocculent silica. It was ascertained that none of the sulphur present existed in oxidised form. Analysis gave Quotients on dividing by Molecular Weights. The mineral is, therefore, opal silica-fiorite or siliceous sintra. The trace of chloride of iron comes in quite naturally from its volcanic origin. The amount of water retained after drying at 100° C. agrees very closely with the results of Gottlieb's experiments (Journ. Prakt. Chem. [2], vol. vi., 185-196) upon hydrated silicic acid artifically prepared by decomposition of silicon fluoride. (6). Analysis of "Novaculite," or "Ouachita Whetstone," from Hot Springs, Arkansas. By Mr. C. E. Wait. A very pure, snowy white specimen of this beautiful material, which has been fairly described by Dr. D. D. Owen in his "Second Geological Report on the State of Arkansas, as 46 equal in whiteness, closeness of texture, and subdued waxy lustre, to the most compact forms and white varieties of Carrara marble," of sp. gr. 2.649, proved to consist of— Silica Alumina Magnesia Sodium oxide Potassium oxide Iron 99.635 (by diff.) 0'778 5.769 0'056 Ferrous oxide The ratio of the oxygen in the bases to that in the silica is about 6 10, the slag approximating to a simple metasilicate of calcium. While the general composition is quite different from that of lapis lazuli, the resemblance in colour and reaction with hydrochloric acid strongly suggest the probable presence of the same or an analogous compound of sulphur with the one to which the fine blue tint of that mineral is generally ascribed. There is enough of the alkaline metals present to allow of the existence of their sulphides, according to the older view of the constitution of lapis lazuli and ultramarine, and more than enough aluminium for its sulphide assumed by Stein to be the cause of the colour.. (5). Analysis of a Siliceous Crust on the Surface of Decomposing Obsidian from the East Side of St. Castagna, Lipari. By Mr. J. A. CABELL. A light grey obsidian, exposed to the action of heated sulphur dioxide, air, and vapour of water, had been decom. posed to the depth of 10 to 12 m.m., producing a white, opaque, porous mass, on the outside of which a delicate, warty, and vesicular crust of 5 to 10 m.m. appeared, resembling hyalite, quite colourless, most of it transparent, some portions slightly milky. This outside transparent crust was the subject of examination. Sp. gr. (in powder)=2.062. Dissolved to the extent of one-third its weight by boiling for three The silica, or at any rate nearly all of it, appears to be in the crypto-crystalline, not in the amorphous or opaline form, as on boiling for three minutes with a 20 per cent solution of sodium hydrate but 163 per cent of the mineral was dissolved, and thirty minutes' boiling only led to 3.56 per cent being taken up. (7). Analysis of Electric Calamine from Wythe County, Virginia. By Mr. J. R. McD. IRBY. This mineral was from the land of the late Mr. David Graham, on New River, about 10 miles above the point where it is crossed by the Atlantic, Mississipi, and Ohio Railway. It occurred in irregular masses, for the most part made up of contorted sheets 6 or 8 m.m. thick, botryoidal, and slightly stained by ferric oxide on the outer surface; pure white and nearly opaque, with radiated structure within; hardness, a little over 5; sp. gr., 3'338 at 21° C. It slowly and imperfectly gave up its water at 100° C. |