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OF THE CHEMICAL SOCIETY, PAPERS READ BEFORE THE CHEMICAL SOCIETY. I.-First Report to the C‘iLemicul Society on “Researches on some Points in Chemical Dynamics.” By C. R. ALDER WRIGHT, DSc. (Lond.), Lecturer on Chemistry, and A. P. LUFF, Demonstrator of Chemistry, in St. Mary’s Hospital Medical School. 9 1. INTRODUCTORY. From certain theoretical speculations, it appeared probable to us that in cases of actions of single decomposition typified by the equa- tion- A + BC = AB + C, the temperature at which the action commences is a function (I) of the physical condition of the bodies experimented with ; (11) of the heat-disturbance (i.e., evolution or absorption-warmetonung) occur- ring during the reaction ; and (111) of other chemical habitudes of the substances concerned, possibly expressible as numerical values con- stant for each substance.(I.) That the temperature of initial action of a given body A, on a given compound BC, varies to an appreciable though not an enormons extent with the physical condition of the compound BC, is not only highly probable, h priori, but has been demonstrated to be the case in the instance of carbon oxide acting on ferric oxide, by I. Lo w t h i an Bell,” with whom one of us co-operated in the inquiry. According as the ferric oxide had been prepared by a wet process, or by ignition of ferrous sulphate, and according as it was pure or ictermixed with * “Chemical Phenomena of Iron Smelting,” by I. Lowthian B e l l , F.R.S., M.P. Routledge and Sons. Page 13, et sep. VOL. XXXIII. B2 WRIGHT AND LUFF OK RESEARCHES OX inert earthy matters (as in the case of various iron ores), it was found that the temperature at which pure carbon oxide first began to take away oxygen from the ferric oxide forming carbon dioxide varied between 141” and 211”.It is a matter of everyday experience that, hard compact substances are little if at all affected by various chemical agents, under circumstances where action is readily brought about if these substances be in different physical states as to aggregation of particles, &c., although determinations of the temperatures at which action commences in these different instances have for the most part not yet been made. (11.) As regards this speculation, little if any trustworthy information of an experimental kind is extant.Lowthian Bell has shown (7oc. cit., p. 91) tliat carbon oxide does not act on zinc oxide at 420°, whilst it reduces lead oxide completely at that temperature. Now the heat disturbance during the reaction- ZnO + CO = Zn + C02, must have a lower value than that during the reaction- PbO + CO = Pb + C02, because the “ heat of combustion ” of zinc is greater than that of lead, whence less heat must be evolved in the reduction of zinc oxide by a given reducing agent than in that of lead oxide. In order to see if it is a general rule that, ccete~is paribus, the temperature at which a reducing action commences is lower the greater the algebraic value of the heat disturbance taking place during its occurrence (i.e., the less the heat absorption, or the greater the heat evolution taking place), the experiments detailed in this report were made.(111.) That the temperature at which action commences is influenced by other chemical habhdes of the bodies involved besides their “ heats of combustion” (and their physical state), is rendered probable by the observation of Lowthian Bell (Zoc. cit., pp. 13 and 91), that stannic oxide is unaffected by carbon oxide at 420°, whilst ferric oxide is acted o u at much lower temperature (141” to 211°, according to the physical state) ; for the “ heat of combustion ” of tin is nearly the same as that of iron (reckoned per equal quantities of oxygen consumed): viz., according to Aiidrews :- Tin. Iron, 4230 4153 Heat evolved in uniting with 1 gram of oxygen . . from which it might be expected that the action of carbon oxide on tin oxide and on iron oxide would commence at temperatures not far apart. It would result from this that if numerical values, “constat~ts of chemical activity,” be really assignable to each substance, the constaut, for sitannic oxide must differ considerably from that for ferric oxide.SOMX POINTS IN CHEMICMJ DYNAMICS.3 In order to trace out the relationships (if any) which exist between the heat disturbances during reactions of the form- A + BC = AB + C, and the temperatures at which these actions are first noticeable, a series of observations have been made as to the temperatures at which the reducing actions commence of the three agents, carbon oxide, hydro- gen, and finely divided carbon, on the oxides of copper and iron in dif- ferent physical conditions, with the result of substantiating the general correctness of the rule, that the lower the algebraic value of the heat dis- turbance, the higher the temyeratwe of initial actiovt, so far as these bodies are concerned.Thus during the reactions- (1) ...... CUO + GO = cu + co, (3) ...... 2 c u o + c = 2cu + coz, (2) ...... CuO + H, = Cu + H20 the heat disturbances are (per 1 6 grams of oxygen transferred) + 30.05, + 19.52, and + 9.48 kilogram heat-units respectively ; with a given specimen of copper oxide (i.e., with a constant physical state), the reducing action of carbon oxide is uniformly manifested at a lower temperature than that of hydrogen, and that of hydrogen is noticeable at a lower temperature than that of carbon. With cuprous and ferric oxides, and these three reducing agents, the same order of sequence is observed; whilst on comparing the temperatures of initial action of these reagents on oxide of copper and oxide of iron in analogous physical states, it is found that each reducing agent begins t o act on copper oxide at a lower temperature than on iron oxide, the heat disturbances during the reduction of ferric oxide by these three agents being respectively + 1.!,0, -8.63, and -18.61, or considerably less in each instance than with copper oxide and the same reducing agent respectively, $ 2. CALCULATION OF HEAT DISTURBANCES DURING REDUCTION OF OXIDES OF COPPER AND IRON BY CARBON OXIDE, HYDROGEN, AND CARBON.The numerical values just cited as the calculated heat disturbances are deduced from the following data ; they represent the heat evolu- tions or absorptions that would ensue if the metallic oxidcs were reduced in such a way that all the materials and products were examined at the same temperature, via., about 15" C.; as shown further on, little difference in the relative values is caused by recalculating the numbers, on the supposition that the action takes place at some more elevated temperature, say 200" or 300". Let H, be the heat evolved by copper in uniting with 16 grams of oxygen to form cupric oxide. B 24 WRIGHT AND LUFF ON RESEARCNES ON H, that evolved by CO in uniting with 16 grams of oxygen to form H, that evolved by hydrogen in uniting with 16 grams of oxygen to And H4 that evolved by carbon in uniting with 16 grams of oxygen Then the heat disturbances taking place during the reactions- co,.form HaO. to form CO,. (1) ...... CUO + co = Cn + co, (3) ...... 2cuo + c = 2cu + COa, (2) ...... CUO + Hz = CU + HZ0 are respectively :-(1) Hz-HI ; (2) H, -HI ; and (3) H4--H1. For HZ the following values have been found:- Now H, was found by Andrews to be 38-30 kilogram heat-units. 28 x 2.431 (Andrews.) 2,s x 2.40'3 (Favre and Silbermann.) 28 x 2.490 (Dnlong.) Mean 28 x 2.441 = 68.35. Again, for H3 the value 2 X 34.275 = 68.55 is found as t,he mean ~7alue deduced from the experiments of Hess, Dulong, Grassi, Favre, and Silbermann, J u l i u s T homsen, Andrews, and Joule." This number expresses the " heat of combustion of hydrogen," assuming that water is formed; if, however, aqueous vapour be pro- duced, as is the case in the experiments that follow, this value is too high by 18 x 0.596 = 10.73, where 0.596 is the latent heat (in kilo- gram heat-units) of aqueous vapour at 15" (Regnault).Hence H3 = 68.55 -10.73 = 57.82. Similarly for H4 the following values have been found (for amor- phous carbon) :- 6 x 8.080 (Fnvre and Silbermann.) 6 x 7.912 (Despretz.) 6 x 7.900 (Andrews.) Mean 6 x '7.964 = 47.78 Q Vide Alder W r i g h t , Phil. Zag., Dec., 1874. This number is practically the same as t,hose found recently by von T h an, and by S c h u 11 e r and v. Wart, h a (Dew5 Chem. Bes. Bey., 187'7, 947 and 1298) ; the first of whom obtained the values 2 x 33.982 and 2 x 34.041, whilst the latter observers found 2 x 34126, when 2 grams of hydrogen and 15.96 grams of oxygen (S tas) at 0" are burnt to water a t 0".SOME POINTS IN CHEMICAL DY NAISIKCS.5 Hence the heat disturbances during the above three reactions are- (1) H, - HI = 68.35 - 38.30 = + 30.05 (2) H, - HI = 57.82 - 38.30 = + 19.52 ( 3 ) H4 - Hi = 47.78 - 38.30 = + 9.48 I n just the same way the heat disturbances during the reduction of iron oxide are calculated; the "heat of combustion" of iron, per 16 grams of oxygen added on, was found by Andrews to be 16 X 4.153 = 66-45 kilogram heat-units ; whence the heat disturbances during reduction by carbon oxide, hydrogen, and carbon are respec- tively- 68.35 - 66.45 = + 1-90 57.82 - 66.45 = - 8.63 47.78 - 66.45 = - 18.67 I n order to recalculate from these numbers the heat disturbances taking place at any other temperature, say To, the following formula must be employed :- rr, = HI, + 72.1 + h.2 - (h, + h*), where Hk is the recalculated heat disturbance at T".H15 is that found as above at 15". hl is the heat required to raise the temperature of the metallic oxide h2 is the heat required to raise the temperature of the reduciiig ha is the heat required to raise the temperature of the reduced metal And h4 is the heat required to raise the temperature of the comple- h,, 1x2, h g , and h* are obtained by means of the formula- from 15" to T. agent from 15" to T, from 15" to T. mentary product from l 5 O to T. h = W x s x (T-15) Where W is the weight of substance heated (in kilos.), and S its mean specific heat between 15" and To. Since W and S are tolerably small fractions, the value of h,, h,, h3, ha are necessarily not very k r p , so that the differential correction, hl + h2 -(h3 + h4) is but small Thus if '1' be taken as 215" (so that T - 15 = 200", this number being taken to simplify calculation) the values of H, for the three cases of reduction of copper oxide by carbon oxide, hydrogen, and carbon become respectively-6 WRIGHT AND LUFF ON RESEARCHES ON For Carbon Oxide- Value of W.Value of s. + 0.028 x 0,245 x ZOO}={ - + 1*37}=~,~+0.52 1.21 - 0.0635 x 0.095 x 200 - 0 . 0 4 x 0.216 x 200 - 1.90 H215 = HI5 + 0.0795 x 0.142 x 200 €115 + 2.26 For Hydrogen- H21j = H15 + 0.0795 x 0.142 x 200 H,, + 2.26 + 0.002 x 3.409 x ZOO)=( - + 1*36}=-~+0.68 1.21 - 0.0635 x 0.095 x 200 - 0.018 x 0.4805 x 2,OO - 1.73 For CaThon- Hz15 =HI, + 0.0795 x 0.142 x 200 H,, + 2.26 + 0.006 x 0.2415 x ZOO>=( - + O*?y)_ir..+o.3g 0.21 - 0,0635 x 0.095 x 200 - 0.022 x 0.216 x 200 - 0.95 It is manifest that the three values H,, + 0.52, HI, + 0.68, and HI5 + 0.39 (or 30.57, 20.20, and 9-87) are in the same order of mag- nitude as the values of HI5 itself, viz., 30.05, 19.52, and 9-48.The 0.52 + 0.68 + 0.39 = 0.53, is awerage increment of the three cases, 3 so near to each one increment severally, that the difference is much less than may readily be attributed t o inexact values being taken for S, the values assumed above not representing exactly the mean spec/@ 7wafs between 15" and 215", but only approximations thereto, being mainly the mean specific heats found for somewhat different ranges of temperature. 9 3.PREPARATION OF METALLIC OXIDES I N DIFFERENT PHYSICAL STATES. (I.) Cupric Oxide. Three varieties of cupric oxide in as pure a form as possible were employed throughout the whole series of experiments ; these were obtained as follows :- Specimen A .-Electrolytic copper was dissolved in nitric acid and precipitated by a solution of caustic potash, well boiled, filtered, and thoroughly washed and dried at 130". In order to prevent the taking up of organic matter by the copper oxide, it was found absolutely indispensable to employ washing water freed from organic matter by careful re-distillation, and potash that had been fused at a red heat before use ; glass-wool, or asbestos filters only were employed. In several preliminary experiments much time was lost through the useSOXE POIXTS IN CHElIIChL DYNAMICS.7 of oxides of copper and iron containing traces of organic matters de- rived from one source or other ; these oxides were found t o evolve small quantities of carbon dioxide on heating to rednws in VCLCZLO, which necessarily vitiated the determinations of the temperature a t which carbon begins to reduce these metallic oxides (8 6). Specinze~ B.-Copper nitrate (from electrolytic copper), carefully ignited till completely converted into oxide and retaining no oxide of nitrogen. Spechen C.-Electrolytic copper, in thin slips, roasted for many days a t a bright red heat in a porcelain tube, through which air was aspirated. The lumps of oxide thus produced were crushed, metallic particles picked out, and the whole again roasted in a current of air, and finally in oxygen, a t a bright red heat, for several hours.Even after several repetitions of the roasting, it was not found practicable to obtain a material containing exactly the amount of oxygen to form cupric oxide ; the substance ultimately obt,ained was finely powdered, thoroughly intermixed, and analysed with the result of finding that the oxygen present was exactly 98.5 per cent. of that required to form CuO, i.e., that the composition was Cu2m0191, o r 194Cu0,3Cu20: the deficiency in oxygen was determined by dissolving known weights in a mixture of hydrochloric acid and ferric chloride, diluting with boiled water, and titrating by permanganate. As stated below (§ 5 ) , this method was found t o give very constant results, and was employed, in many cases, to decide whether deoxidation of copper oxide had commenced, and to what amount it had gone on.Besides these, several other specimens of copper oxide were used for comparison in some of the experiments (vide $ 4). (11.) Caprous Oxide. According t o Andrews, the " heat of combustion " (per constant quantity of oxygen) of copper to cuprous oxide is almost identical with that to cupric oxide, viz., per 16 grams of oxygen :- Copper to cupric oxide . . . . . . . . 38-30 kilogram heat-units 7 9 cuprous ,, .. .. .... 36.61 ,, 7 7 T t should therefore result, if the hypothesis be correct that the temperature of initial action is ( c u h r i s ycirlbus) lower the grc,ater the heat evolution (algebraically), that cuprous and cupric oxides should be acted on by each reducing agent at about tlie same temperature, their physical states being identical.As shown below, this appears to be the case. Caprous oxide mas prepared perfectly free from organic matter by dissolving cane-sugar in water, filtering, and inverting by boiling8 WRIGHT AND LLWF ON RESEARCHES ON with sulphuric acid ; electrolytic copper converted into nitrate was boiled with potash (recently fused a t a red heat) and this inverted sugar, and the red ouprous oxide precipitated, thoroughly washed with water free from organic matter. Several other samples of cuprous oxide were prepared in other ways, such a s reduction by a filtered solution of commercial glucose, $c. ; but these invariably re- tained more or less organic matter, and evolved CO, on heating to redness in vacuo; some absorbed oxygen whilst drying, and niucli labour was expended before a good sample was obtained.On analysis, the pure cuprous oxide was found to contain 99.7 per cent. of CE~O, the process adopted being titration with permanganate as above de- scribed ; it gave off no trace of CO, on heating to redness in a Sprengel vacuum. (111.) Ferric Oxide. Three varieties of ferric oxide were used in the experiments with this substance. Specimen A.-Pure ferrous sulphate was recrystallised several times, dehydrated, and strongly calcined for some hours till wholly con- verted into ferric oxide ; the compact granular substance produced was ground in a mortar to a fine powder. Specimsrz B.-Ferric oxide prepared by calcining the sulphate was dissolved in boiling hydrochloric acid, and the solution poured into excess of ammonia water : the precipitat,e was boiled and thoroughly washed till no trace of chlorine was present in the wash waters ; after drying in the water-bath, it was rendered anhydrous, or almost abso- lutely so, by long-continued heating a t 180-200".Specimen C.-Ferric oxide 13 was strongly ignited in a platinum crucible for a few minutes; it shrank considerably in bulk, and became darker and less readily soluble in hydrochloric acid. Several other specimens were previously prepared by precipitation of ferric salts, such as the filtered solutions of commercial chloride, &c. ; but these were found to contain traces of organic matter derived either from the materials or the wash water, as they gave off more or less carbon dioxide on heating to redness in a Sprengel vacuum.The above three specimens, however, were found to be free from organic matter by this test ; and after heating in a test-tube to 300-400", they were not i n any way reduced, no blue colour being given on solution in hydrochloric acid and addition of ferricyanide. 5 4. DETERXIEATION OF TEMPERATURE OF INITIAL ACTTON OF CARBON OXIDE ON COPPER AND IRON OXIDES, The metallic oxides were placed in a test-tube arranged vertically in a paraffin bath, about 0-5 gram being taken for each observation :SOME POINTS IS CHEMICAL DTNANICS. 9 the mouth of the test-tube was closed by a cork with three perfo- rations ; one for a thermometer, the bulb of which reached down to the bottom ; one for an entrance-tube for leading in pure carbon oxide, also reaching to the bottom; and the third for an exit-tube just passing though the cork, and bent externally to a right angle and coupled by an india-rubber joint to a Will and Varrentrapp’s ammo- nia-bulb apparatus filled with clear baryta-water.In order to make sure that the entering carbon oxide had the same temperatiire as the bath, the entrance-tube was bent over downwards and upwards so as to form a U reaching down to about the bottom of the test-tube, and lying close to the outer surface ; the upper bend of this tube between the cork and the U was wrapped with lamp-wick to avoid aGrial cooling. By keeping the baths a t varions fixed temperatures and passing carbon oxide through for a few minutes, the temperature of initial action was approximately found by noticing at what temper- ature the baryta-water first became turbid ; and then by repeating the observations several times with fresh portions of metallic oxide, still nearer temperature-readings were obtained : the differences observed between several careful determinations rarely exceeded a degree above or below the mean.The carbon oxide employed was generated from ferrocyanide and sulphuric acid being washed through caustic soda, silver nitrate, and mlphuric acid, and stored in a gasholder over water : the gas thus kept in stock invariably contained a little air, probably derived from that absorbed by the water. As some of the experiments showed (§ 4 infra) that a partially reduced metallic oxide sometimes causes a kind of catalytic oxidation of carbon oxide to dioxide when a mix- ture of carbon oxide and a little air is led over it, it is possible %hat the original unreduced metallic oxide might to some extent possess the same power ; wherefore, before use, the free oxygen in the carbon oxide was remove4 by passing the gas from the gasholder through a tube filled with fragments of well-burnt charcoal, and heated red-hotl by a small gas-furnace, and then through a series of Liebig’s bulbs containing caust,ic potash and potassium pyrogallate.In this way carbon oxide, absolutely free from all trace of admixed oxygen, was obtained, but containing a small percentage of free nitrogen. It is noteworthy that a red-hot tube filled with pumice-stone did n o t cause the removal of all free oxygen (by conversion into carbon dioxide) from the air containing carbon oxide, whilst the same tube filled with charcoal fragments compjetely removed all free oxygen unless the stream of gas were considerably rapid.The issuing gas, however, contained carbon dioxide in notable quantity, the reducing action of the charcoal on the carbon dioxide not being suSciently great to re-convert all of that gas into carbon oxide; hence, potash-bulbs toLO KRIGHT AND LUFF ON RESEARCIES ON purify the issuing gas from carbon dioxide were indispensable, although the pyrogallate tubes were only requisite as an additional precaution. I n this way, the following numbers .were obtained as the mean results from several trials in each instance :- C1L)IYZ'C O.?~id& Specimen A.Specimen B. Specimen C. Temperature at which turbidity 1 } 60' 125" 146" was first produced in baryta- water after from three to five minutes. . . . . . . . . . . . . . . . . . . . J Temperature a t which action was bidity produced by a few bub- bles of issuing gas). . . . . - . . . . Very similar numbers were obtained with other samples of cupric oxide prepared in similar ways, but not so pure ; although some of these contained traces of organic matter, yet none of them evolved any carbon dioxide on heating in a current of air t o a temperature higher than that a t which carbon dioxide was freely formed on heat- ing in carbon oxide. Thus, three other specimens gave these num- bers : D, prepared by precipitation from recry stallised commercial copper sulphate, washing, and drying a t 130" ; E, by ignition of commercial copper nitrate; and F, roasted copper wire, as sold for combustion analyses, powdered in a mortar.Temperature of first turbidity . . . . 68' 125" 137" Temperature of well marked action 85" 133" 142" well marked (considerable tur- } 68" 133" 150" Specimen D. Specimen E. Specimen F. Cq~rnus Oxide. Besides the pure cuprous oxide described above, two other speci- mens were also examined so far as the temperature of initial action of carbon oxide is concerned ; one prepared by precipitation of copper nitrate by glucose and potash, and one purchased from Messrs. Hop- k i n and Williams : each of these two contained small quantities of organic matter, but gave off no carbon dioxide on heating in a current of air to temperatures above those where the reducing action of carbon oxide is well marked.Pure cuprous I m p r e cuprous oxide. oxide. -1 -7 From H o 1) li i 11 and W i 11 i a 111 s. r 33s glucose. Temperature of first turbidity.. . . 110" 128" 140" Temperature of well marked action 120" 136" 150"SOME POINTS IN CHEMICAL DYNANICS. 11 Specimen A. Specimen B. Spec*iinni C . Temperature of first turbidity . . . . 202" 90" 220" Temperature of well marked action 210" 100" 2:30° Two other specimens of ferric oxide, prepared by precipitation, but retaining traces of organic matter, were first acted on a t 85" and a t 97", the action being well marked a t 100" and 105" respectively. The above numbers show how considerable an influence on the tern- perature of initial action is exerted by difference in physical state ; this is especially well instanced by specimens B and C of ferric oxide, where simple ignition of B for a few minutes raises the temperature cf initial action by 130'.I n one experiment some of the copper oxide B had been heated in a porcelain crucible over a Bunsen burner, and the surface had become reduced by the diflusion into the crucible of gases from the flame ; on allowing to cool with the lid off, the partially reduced oxide again became black and apparently reoxidised ; after cooling, however, the portion constituting the upper film was found to be first acted on by carbon oxide a t 97", the action being well marked at 100" ; so that a lowering of some 28" in temperature of initial action was brought about by tlie alteration in physical structure due to partial reduction and reoxidation a t a temperature not very elevated.The numbers found by Lowthian B e l l (Zoc. cit.) as the tempera- ture of initial action of carbon oxide on ferric oxide, prepared by cal- cining the sulphate, and by precipitation and gentZe ignition, were respectively 208" and 141" ; the above results agree well with the first of these numbers, whilst tlie numbers found for precipitated oxide dried a t 180", and the same strongly ignited, lie on opposite sides of the value found by B e l l in the latter instance. It deserves notice that the action of carbon oxide on oxide of copper prepared by precipitation, well marked a t temperatures conside~nbly below loo", becomes extremely energetic a t 100" : tubes of 50 to 100 C.C.capacity, containing about a gram of the copper oxide and filled with carbon oxide, were sealed and heated to 100" for a few hours ; a t the end of the time the carbon oxide was wholly converted into carbon dioxide ; in one case this result was observed after four hours' heat- ing, in another after two hours a t 100". I n order to gain some idea as to the rapidity with which the action goes 011 a t loo", a slow stream of purified carbon oxide was lcd over 0.5 gram of cnpric oxide (speci- men A) in a tube heated in boiling water, and the carbon dioxide pro- duced collected and weighed ; in this way the iollowing numbers were obtained :-12 WRIGHT AND LUFF ON RESEARCHES ON Oxygen removed, Time.that originally present being 100". 15 minutes 22.6 60 7 7 49.3 90 > 7 71.8 120 7 7 79.8 The loss of oxygen deduced from the loss of weight of the cupric oxide a t the end of the time was 82.4, the slight difference being due in all probability to the carrying away of a milligram or two of moisture from the potash bulbs for absorption of CO, by the issuing excess of CO. The reduced copper oxide was black and distinctly pyrophoric whilst warm. In the course of some other experiments of the same kind, it was noticed that when the carbon oxide was not purified from small quan- tities of admixed air by the method above described, the amount of carboil dioxide collected greatly ezceedstb that corresponding to the loss of weight ; thus in two experiments the following numbers were obtained :- Oxygen removed, calculated Oxygen removed, Time of exposure.from COO found, that calculated from originally present bzing looo. loss of weight. (I) 90 minutes 155.0 71.9 (2) 220 ,) 113.0 72.0 whence it is evident that either the partially reduced oxide acted catalytically in causing combination between the free oxygen and the carbon oxide in the gaseous mixture, or the nearly reduced oxide took up the free oxygen, and was then again reduced in virtue of the reactions- cu,o, + 0 = CaxOy+l cu,o, + 1 + co = cu,o, + co,, thus forming more carbon dioxide than that corresponding to the difference between the original oxygen in the cupric oxide employed, and that left associated with the copper after the action was com- pleted.§ 5. DETERMINATION OF TEMPERATURE OF INITIAL ACTION OF HYDROGEN ON COPPER AND IRON OXIDES, The same arrangement was adopted as that employed as above described in the case of carbon oxide, the baryta-water bulb being omitted ; in the case of some few of the observations where thc tem- perature was above about 260" (so that there might be some danger of the cork being charred and vapours given off which might perhapsSOME POIBTS IN CHENICXL DYNAMICS. 13 reduce the metallic oxide), a simpler arrangement was employed con- sisting of a narrow glass tube passing horizontally through a hot-air bath, the metallic oxide being placed in the central part of the tube, and the bulb of a thermometer just level with i t : to avoid currents of air i n the bath, the tube and thermometer-bulb mere enclosed in a small inner chamber with perforations just big enough to allow the ends of the tube and the thermometer-bulb to pass through.The hydrogen was prepared from zinc a'nd sulphuric acid, being passed successively through (J -tubes containing pumice-stone soaked in silver nitrate, caustic potash, and sulphuric acid. In order to de- termine as nearly as possible the temperature of initial action, qnanti- ties of about 0.3 gram of metallic oxide were used for each obseri ation ; the air-bath having been heated to some given temperature, the hydro- gen was turned on and allowed to pass through for 15 minutes ; the tube containing the oxide was then removed from the bath and allowed to cool, a current of hydrogen still passing through.The cooled con- tents were then boiled with hydrochloric acid (in the case of the copper oxides with hydrochloric acid and ferric chloride), the soliltion diluted with boiled water and tested by ferricyanide or by permanganate ; generally weighed quantities of metallic oxide and a standard per- manganate solution were employed. Careful tests showed that neither the cupric oxide prepared by precipitation (A), nor that by ignition of the nitrate (B), nor any one of the three kinds of ferric oxide, caused the faintest trace of blue with ferricyanide when tested in this way before exposure to hydrogen or after heating per se. As to the copper oxide from roasting of metal (C) and the cnprous oxide, weighed quantities were taken for each trial, and the amount of permanganate corresponding subtracted from that ultimately consumed ; it was found that absolutely concordant numbers were obtained in several consecu- tive blank experiments made to determine the amount of permanganate solution required for the copper oxides before exposure to hydrogen.In this way the following numbers were obtained, 15 minutes' exD0- sure to hydiogen a t the case :- Temperat me. All temperatures up I ternperat'ure indicated being allowed in each Coyper Oxide (A). Percentage of oxygen removed (that in original oxide being 100'). to 83" No trace. At ............ 87-90' 1 - 7 ,, .............. 100" 7.0 ,, .............. 150" Almost complete reduction ; much red spongy metal formed not readily dissolved by hydrochloric acid and ferric chloride.14 WRIGHT AN) LUFF ON RESEARCHES ON Cupuic Oside (B).Temperature. Percentage of oxygen removed. All temperatures up to 170" At .......... 178-180" 4.3 ), .............. 200" 8.8 No trace. Cupuic Oxide (C). All temperatures up to 170" Nil. At .............. 175" 16.7 ,, .............. 240" Much metal reduced, undissolved by a few minutes' boiling with hydrochloric acid and ferric chloride. Cupous Oxide. All temperatures up to 150" At .............. 160" 2 7-5 ,, .............. 180" Much metal left undissolved by hydrochloric acid and ferric chlnride. No trace. Ferric Oxide (A). All temperatures up to 255" At .......... 2(j.5-2;oo ,, .............. 290" ,, .............. 310" ,, 360" (in mercury va- pour). Femic Oxide (B). All temperatures up to 190" At ..............200" ................ 210" ................ 2,50" All temperatures up t o 240" At .............. 250" ,, 360" ( i r mercury va- pour). No trace. 1.8 5.6 8.0 16.5 No trace. 0.7 1.4 i s 7 No trace. 1-2 14-5 From these observations the following temperatures may be taken as close npprciximations to the temperatures of initial action of hydro- gen 011 these metallic oxides :-SOME POlNTS IN CHEMICAL DYNASIICS. 15 Oaides of Copper. Cupiic oxide (A). Cupric oxide (B). Cupric oxide (C). Cuprous oxide. 85" 175" 172" 155" Fewic Oxides. Specimen (A). Specimen (B) . Specimen (C). 260" 195" 24.5" It is noticeable that by plotting out some of the above results, taking the temperatures as ordinates and the percentages of oxygen removed as abscisse, approximations are obtained to curves which indicate the relative degrees of energy of reducing action of hydrogen at different temperatures.It is proposed to multiply and extend observations of this nature, carefully conducted under precisely the same conditions, save as regards temperature, so as to obtain the data for accurately- determined curves of this kind, not only with hydrogen and copper and iron oxides, but also with other substances. Three experiments mere made by Lowthian B e l l (Zoc. cit., p. 118) by exposing calcined Clevehd ironstone to the action of hydrogen a t 104-127" for 30 minutes, a t 199-227' for 30 minutes, and at a bright red-heat for four hours respectively. In the first case no reduction whatever was perceptible ; in the second a trace of reduction had ap- parently taken place ; whilst in the third much of the ferric oxide was reduced, though not all.The evidence of the trace of reduction in the second case was that the material, after exposure, slightly reduced permanganate when dissolved in acid : it is manifestly possible that this action was due to the presence, in the substance treated, of a minute grain of imperfectly peroxidised ferrous carbonate ; still, ad- mitting that the reducing action of hydrogen in this form of impure ferric oxide is just perceptible between 199" and 227", this result is fairly in accordance with the above observations. Stromeyer found (Pogg. AnnuZerL, vi, 471 [1826]) that ferric oxide (how prepared is not stated) is readily completely deoxidised by hydrogen at a red-heat, reduction commencing a t considerably lower temperatures, but not proceeding very rapidly.31 a g n u s found (il, id., vi, 509) that whilst ferric oxide is not appreciably acted on by pure dry hydrogen at the temperature of boiling water, and whilst no visible production of water ensues even on heating in boiling rape oil, conzpZete deoxidation ensues in two hours a t a temperature close to that of boiling mercury, and tit any rate below the temperature of meltiug zinc.16 WRIGHT AND LUFF ON RESEARCHES ON 6 6 . DETERMINBTION OF TEMPERATURE OF INITIAL ACTION OF CARBON ON COPPER AND IRON OXIDES. The first preliminary experiments in this direction were made by heating intimate mixhres of the carbon and metallic oxides (ground together in an agate mortar) to various temperatures, and then en- deavouring t o measure the reduction by solution in hydrochloric acid (or hydrochloric acid and ferric chloride) and titration by perman- ganate ; but so many difficulties occurred in the way of obtaining any- thing like trustworthy results, that this plan had t c be abandoned.In the first place, all kinds of carbon tried produced more or less reduc- tion of ferric chloride when boiled t'herewith without any metallic oxide a t all ; and the amount, of reduction produced was different after heating to, say, 300" from what it was originally, and was not con- stant after the heating, so that it was not possible to applya correction by using weighed quantities of carbon ; this was ultimat,ely traced to the power which carbon possesses of occluding gases, and in particu- lar carbon oxide, which in this occluded form seems to act readilyas a, reducing agent, much as hydrogen, occluded by palladium, possesses high reducing powers.For this same reason, indications of reduc- tion by carbon were obtained a t temperatures far lower than those subsequently determined by more accurate methods, the heat causing a certain amount of the occluded carbon oxide to be evolved, and so to act on the metallic oxide. (It is not unlikely, moreover, that occluded air on heating would act on the carbon, thus producing addi- tional carbon oxide.) Direct experiments proved, as described below, that the gases occluded by carbon are not entirely given off in a Sprengel vacuum a t the ordinary temperature, notable quantities of carbon oxide and dioxide being extracted by this means on heating after complete exhaustion a t the ordinary temperature.The mode of operation finally adopted was the following :-Batches of carbon were prepared in quantity sufficient to last for a whole series of experiments : the amounts of gas given off by weighed quantities of each kind on heating through certain ranges of temperature after pre- vious complete exhaustion in a Sprengel vacuum at the ordinary tern- perature were determined once for all (on samples drawn from the bulk after thorough intermixture). Weighed quantities of carbon and metallic oxide were then thoroughly incorporated by con h u e d grinding together in an agate mortar ; the mixture was transferred to a suitable tube connected with a Sprengel pump, and the whole ex- hausted till no more gas was extracted by 5 or 10 minutes' action of the pump.The tube was then heated to a given temperature mid maintained thereat for 15 to 30 minutes, and the gas evolved pumpedSOME POINTS IN CHEMICAL DYNAMICS. 17 out, measured, and analysed (when necessary). As long as the bulk of gas thus extracted did not amount to appreciably more thari that furnished by the same weight of carbon alone under the same con- ditions, it was inferred that there was no action between the carbon and the metallic oxide, and the temperature of initial action was taken to be that temperature a t which more gas began to be furiaished thaw, that due to the carbon alone, the evolution of gas going om continuously foi.some time. I n thus manipulating it was found absolutely necessary to confine the mixture of carbon and metallic oxide in the closed end of the tube by a plug of recently ignited asbestos firmly rammed down, otherwise the evolution of gas from the carbon on exhausting caused a dancing about of particles of carbon in the tube, thus effectinq a partial mechanical separation of carbon from metallic oxide ; and ivhat was of still more serious moment!, caused particles of carbon to be drawn over into the tubes of the Sprengel pump, thus utterly vitiating all measurements by reabsorbing, to a greater or lesser extent, the gas evolved. It was found that after strong ignition the asbestos used for the plug retained so little occluded air, &c., that when the tube 'was rendered vacuous while cold it practically remained so on heating if nothing but asbestos were present.Many of the earlier observa- tions were spoiled, however, by our omitting this precaution and using asbestos which parted with sensible quantities of gases on heating after rendering the tube vacuous when cold. Several kinds of carbon were experiniented with a t first, but many of these were rejected because they contained considerable percentages of hydrogen (even after strong ignition), or on account of inorganic impurities. T.wo varieties were ultimately selected, which, although not absolutely pure amorphous carbon, were, we believe, as pure as that substance can be practically obtained. One of these was a hard dense coke-like charcoal, prepared by heating recrystallised sugar in closed crucibles, powdering the light charcoal thus obtained, and strongly igniting it in small portions at a time in well-closcd pla- tinum crucibles, then igniting it in it current of chlorine for two hours, and finally again igniting it in closed platinum crucibles over a blowpipe for six hours until no trace of hydrochloric acid escaped.The other was an extremely light porous carbon, obtained by passing carbon oxide over ferric oxide at about 400450" for many hours, whereby (as Lowthian B e l l has shown, loc. cit., p. 44) rednction of the carbon oxide to carbon is brought about, to a large extent, by the agency of the lower oxides of iron first formed, thus :- Fe,O, + CO = C + FexOy+,, VOL. XXXIIL. C18 WRIGHT ASD LUFF ON RESEARCHES ON the higher oxide thus formed being again reduced by another portion of carbon oxide. Fe,O,+ + GO = CO, + Fe,O,.The black mass thus obtained was boiled with dilute hydrochloric acid, thoroughly washed, again boiled with concentrated hydro- chloric acid, washed on the pump-filter till the washings were neutral, and finally dried at loo", and gently ignited in a closed crucible to expel moisture. On analysis these two samples gave the following numbers :- From sugar; 0.2575 gram gave 0.99130 GO, and 0.0105 HzO, and left 0*0040 ash; 0,2375 gram gave 0.8285 CO, and 0.0160 H,O, and left 0.0040 ash. From carbon oxide ; C.2250 gram gave 0.7'360 CO, and 0.0145 H,O, and left 0.0020 ash. From sugar. From carbon oxide. Carbon .............. 96.17 95.1g 96-49 Hydrogen............ 0.84 0.75 0.72 Oxygen (by difference). 1.43 2.44 1.90 Ash ............... 1.56 1-68 0.89 -- 100*00 100*00 100*00 The hydrogen was probably present, a t least to some extent, in tb form of moisture taken up during cooling and weighing ; charcoa, especially of the light pnlverulent character of that from carbon oxide, is far more hygroscopic than oxide of copper, which, as is well known, is very difficult to get perfectly free from hygroscopic moisture. The oxygen was necessarily present to s'ome extent as moisture, the rest being probably in the form of occluded carbon oxide and dioxide ; any occluded air would also be reckoned as oxygen in the above analyses. I n order t,o obtain the corrections for the amounts of gas evolved from these two samples of carbon on heating after pumping out all gases expelled in a Sprengel vacuum a t the ordinary temperature, weighed quantities were placed in tubes with asbedos plugs, and treated exactly as the mixtures of carbon and metallic oxides sub- sequently examined. In about half an hour after first setting the pump a t work, all gas capable of coming off a t the ordinary tempera- ture bas expelled, no appreciable amount being collected during five minutes' action of the punip, the mercury clicking thoroughly.On then heating to 100" in boiling water a little gas was evolved; in a few minutes the vacuum became again perfect, no more gas being evolved in five minutes more. I n the same way a further quantitySOME POISTS IN CHEMICAL DYNAMICS. 19 was obtained on heating to 360" in mercury vapour, the vacuum again becoming perfect in a few minutes.Similarly a little more gas ma$ obtained on successively heating to about 420" and again to near 450", these temperatures being deduced from the effect of heating in a bath of melted solder in which mere inserted glass tubes of the same size and thickness as those employed for the carbon, containing fragments of pure zinc (melting point, close upon 420") and pure antimony (melting point close upon 450"). This mode of approximating to the tem- perature to which the tube was heated was found to be fairly reliable, aiid not subject to greater errors than any other method of determin- ing the temperature that could be conveniently adopted (S i em e n's electrical pyrometer was not tried, however).In this way the following quantities of gas mere obtained from each sample of carbons reckoned in c.cs. a t 0" and i60' per gram of carbon :- Gas collected. Range of temperature. Sugar carbon. Carbon from CO. .............. 1.00 0.20 100" to 360" 0.98 360" to 420" (lead just melted) 0.10 420" to about 460" (somewhat 0.60 3-00 moV) 2.38 4-20 The following analyses were made of the total gas thus collected :- .............. Oe70 1 15" to 100" above melting point of anti- - - * Another sample of carbon was prepared from carbon oxide by passing over pure ferric oxide (Specimen A), boiling the resulting mass with hydrochloric acitl, washing, drying, and heating in hydrogen to redness (any iron compound not washed out), boiling again with hydrochloric acid, thoroughly washing, drying, and gently igniting. This was found to give off a rather smaller quantity of gas, posibly because the additional heating in hydrogen had rendered it somewhat less porous.Gas collected per gram of carbon. 15" to 360". ....................... 0 *60 360" to 420'. ....................... 0.20 420' to about 480' .................. 2 -00 2 *SO Range of temperature. - This gas consisted of- Carbon dioxide ............ '71 -02 Carbon oxide.. ............ 21 -01 Other gases.. .............. '7 *97 100 *oo This sample was not employed in any of the experiments described below.20 WRIGHT ASD LUFF ON RESEARCHES ON Sugar carbon. Carbon from CO. Carbon dioxide .............. 48.2 28.6 Carbon oxide (absorbed by cu- prous oxide) ..............43.4 33.3 Other gases ................ 8.4 38.1 - - 100.0 100.0 3 decigrams of carbon and 6 of metallic oxide were employed in every observation, the temperature determinations being made precisely as above described. Just as with the carbon alone, a practically perfect vacuum was obtained after each further heating of the tube after a few minutes' working of the pump, until a temperature was attained when gas began to come off continuously ; this temperature mas then taken as the temperature of initial action ; on repea,ting the experi- ments the results were never further from those previously obtained than might readily be due to the comparative roughness of the means of estimating the temperature adopted, whilst in most cases almost absolute agreement was noticed.The following are the numerical results obtained (it being under- stood that in every instance the tube was pumped vacuous at the ordi- nary temperature before any heat was applied, and that the tempera- ture WRS never raised above the particular limit assigned till long after a vacuum had been again established, i.e., until the mercury clicked perfectly and no appreciable amount of gas was collected in five minutes ; the temperature was judged to be 420" when the zinc in the companion tube was partially, but not wholly melted, and similarly at 450°, when the antimony was partly, but not entirely, melted; it was estimated a t about 430" when the zinc was just thoroughly melted, and 440" when the temperature was rising and the antimony showed signs of incipient fusion ; a t 460" and upwards, when the anti- mony was thoroughly melted, the temperature still rising I n a few instances temperatures a little below 360" were obtained by a hot air- bath furnished with a thermometer as described in § 5.The volumes of the gases are all reduced to 0" and 760'.SOME POIKTS IN CHEMICAL DYNAMICS. 360-420" 21 420-450* COPPER OXIDES. Carhon. from Sugar. 30 .. .. 0 *50 0 *50 -- Nil. Specimen (A), Cupric Oxide. 15 15 .. lo -90 .. 0 '10 0'55 11.00 0.53 0 -61 ___------ -- Nil. 10.39 __ Specimen (B). .. .. 0 *4 0 50 Nil. 15-360" 30 2 -05 4 *30 0.05 0 '10 2 .lo 4 *40 0 *53 0 -61 ------ 1 -57 3 *'is 360-420" 360450" 15 1 30 15-360" 1360-430" 30 15 ---- .. .. .. .. ---__- 15-360" 430460' 15 7 -90 0 .lo Range of temperature .. Time of heating in mi- nutes .............. 350-360' (hot-air bath) 360" (in mercury vapour) 30 .. .. -- 0 -40 0 .80 0 '10 ----- 30 2 -20 0 .lo 2 *30 0 '30 I__________ 2 '30 Carbon dioxide colIected Other gases Y Y - * ----- Total gas collected. ..... Gas due to occlusion . . , . Gas produced from ac- tion of carbon on metallic oxide .... --.__-- I Total gas collected.. .... Gas due to occlusion.. .. Gas produced by action of carboii on metallic oxide ............ - ~ - - (Specimen (C) . 15-360- 860-420" 1 Range of temperature . . Time of heating in mi- nutes .............. I 30 I 30 - 12 *30 0.50 12-50 0-50 1 0.61 --I__- - 1 0'20 ---- Carbon dioxide collected Otlicr gases 7 7 * . Total gas collected ....... Gas due to occlusion.. .. Gas produced from ac- tion of carbon on metallic oxide...... Nil. I 7'39 Nil. Nil. 11 *89 Carbon f r o m CO. Specimen (B) . Specimen (A). 15-360" 30 360420" 15 15-300" (hot-air bath) 30 Range of temperature . . Tinip of heating in mi- nutes .............. Carbon dioxide collected - Other gases 7 . .. .. .. .. .. 5 '40 0 *lo 5 *so 0 -40 -- 0 -25 0 'YO 0 *20 0 *so Nil. 5 -10 Nil.22 WRIGHT AXD LUFF ON RESEARCHES ON Cnrbosz from CO. I 15-340O (hot-air bath) 30 Specimen (C). I 340-350' (hot-air bath) 30 Cuprous Oxide. 15 ---- .. .. -- 0.40 0 -36 ~ _ _ _ _ _ - I 15 5 -5 0'1 5 -6 about 1 -0 . 36O0 (in mercury Tapour) 30 ----__I__ Carbon dioxide collected. Otlier gases > 7 * * Total gas collected., .... Gas duc to occlusion. ... --- Range of temperature . .I 15-360' I 360-420" 1 420--440' .... ~- 0 *30 0 *30 .. .. -__I 0 *25 0 '30 0 50 0 '10 0 '60 0 *30 Range of temperature. ........... Time of heating in minutes ...... Carbon dioxide collected ........ Other gases ........ 7 ) ~~ Total gas collected .............. Gas due t.0 occlusion ............ 15-360" 15 -- .. .. 0.40 0 *50 About, 360' (in Hg vapour) 30 --~- .. .. __- 0 -50 0 -50 360-420" 15 .. .. 0 -55 0 *53 Gas produced from action of carbon on metallic oxide.. ............ Kil. Nil. Nil. Range of tempernture. ........... Time of heating in minutes Carbon dioxide collected ........ Other gases ........ ...... ____- 7 9 15-360' 15 -8. .. About 360' (in Hg vapour) 30 -- 360-420' 15 Total gas collected .............. Gas due to occlusion ............ C -45 0 5 0 Gas produced from action of carbon on metallic oxide.. ............Nil. Nil. Kil. 1 *70 0 '10 1 *SO 0 *30 -~ 1-50 FERRIC OXIDES. Ci-crbon f r o m Szcgar. Ferric Oxide (A). 420480" 15 2.90 0 -10 3 -00 0 -71 2 -29 I Ferric Oxide (B). 420-480' 15 2 *60 0 -10 .. .. * * I .. 0 -50 0 *55 0'50 1 0.53 2.70 0 9.1 1 '99SOME POINTS IN CHERlICiiL DPXAXICS. 23 Cnrbm f r o m sup^. ,. .. 0 -40 0 '50 Ferric Oxide (C). .. .. -- 0.50 0 *50 15-360' 15 About 360' (in Hg vapour) 30 .. .. 0 *30 0 *30 -- Nil. About 360' (in Hg vapour) 30 360-420' 420-440" 15 15 -- 2.50 0-45 1 2.70 0 -443 1 -00 .. Nil. I 1.70 360-420' 15 Time of heating in minutes ...... Carbon dioxide collected ........ -- Other gases ........ 7 9 _ _ _ _ ~ ~ _ _ _ Total gas collected .............. Gas due to occlusion ............Range of temperature .......... Time of heating in minutes ...... 15 .. .. 0 *25 0 *30 420-480" 15 2 *40 0 *10 -- Gas produced from action of carbon on metallic oxide.. ............ Carbon dioxide collected ........ Other gases ........ 7) Nil. .. .. -- 0 *60 0 -53 -- Nil. -~ Total gas collcctrd .............. Gas due to occlusion ............ Gas produced from action of carbon on metallic oxide.. ............ - e__ --I-- Nil. I Nil. I I Carbon fionx CO. Ferric Oxide (A). Ferric Oxide (B). (Not determined, the specimen being entirely used up.) Ferric Oxide (C). About 360' (in Hg vapour) 30 Range of temperature, ........... 420-440" 15 2.60 0 -20 2 *80 1-00 -- --* -- 1 -80 360420" 15 -- .. .. G -4.0 1 0.40 15-360' 15 Time of heating in minutes Carbon dioxide collected ........Other gases ........ ...... -~ > 7 .. .. -- 0 -20 0 *30 -- Nil. .. .. n 0 -30 0 -30 Total gas collected.. ............. Gas due to occlusion ............ Gas produced from action of carbon on metallic oxide.. ............ Nil. I Nil.24 WRIGHT AKD LUFF OK RESEARCHES ON From these observations the following values are deduced as the temperatures of initial actions :- Copper Oxides. Specimen (A). Specimen (B) . Specimen (C). Cuprous oxide. Sugar charcoal.. . . 390" 430" 440" 390" Carbon from CO . . 350 390 430 345 Iron Oxides. Sugar charcoal . . . . . . 450" 450" 450" Carbon from CO . . . . . . 430 - 430 Specimen (A). Specimen (B). Specimen (C). The above numbers indicate further how energetic is the action of carbon oxide on copper and iron oxides ; the gas collected was always practically wholly carbon dioxide, indicating that all carbon oxide evolved from occlusion as well as any due to the action of carbon on carbon dioxide, or on the metallic oxide, was wholly oxidizcd by the metallic oxide.An experiment was made by Lowthian B e l l (Zoc. cit., p. 53), which appears a t first sight to indicate that an extremely intimate mixture of partially reduced ferric oxide and carbon, prepared by leading carbon oxide over heated ferric oxide, is capable of evolving gas at a temperaiure of 250-265", and therefore that carbon begins to reduce this form of iron oxide a t that temperature. One gram of the mixture was placed in a tube, air displaced by a current of' nitro- gen, and the tube sealed and heated in an air-bath for 44 hours ; a t the end of that time, on opening the tube under rtiercury, it was found to contain about 7.0 c.cs.of carbon oxide and 0.16 of carbon dioxide. As the mixture cont]aiiied carbon 52.14, iron 35-13, and oxygen 12.73 per cent., this amount of gas represents about 13.4 c.cs. of carbon oxide and 0.3 of carbon dioxide per gram of carbon. Admitting that this gas was really formed by t h e action of the carbon on the partially reduced iron oxide, it would not in any way vitiate any coiiclusions drawn from the results abo-re described, noT is it necessarily opposed to them in any way, since it is highly probable that the oxide of iron, partially reduced at about 400" by carbon oxide, would be excessively spongy and in an entirely different physical state from the ferric oxide examined by us, and therefore that action would commence a,t a lower temperature ; but the following experiment ren- ders it somewhat doubtful as to how far tlhis gas really did arise from the action of the carbon on the iron oxide, and how far it WRS simply occluded gas displaced by the heating and not perfectly reabsorbed as the tube cooled. We prepared an analogous mixture containing-SOME POIXTS IN CHEJIICBL DYNARIICS.25 Carbon ....................................... 39.4 Metallic iron (soluble on digesting with iodine and boiled water) ................................ 23.3 Iron existing as oxide (insoluble in iodine water) .... 31% Oxygen (by difference) .......................... 5.7' 100.0 - and then placed a gram of it in a tube, displaced all air by nitrogen, and exhausted cold with the Sprengel pump.After subtracting from the total bulk of gas collected that due to the capacity of the tube, &c., a residue of 1.2 c.cs. appeared, and on analping the expelled gas just 1.2 c.ca. of carbon oxide were found, or 3.0 c.cs. per gram of carbon present, and not more than traces of carbon dioxide. On successively heating to loo", to 250", and to 360", gas was pumped out at each heating, the tube speedily becoming practically vacuous till the temperature was again raised. The following quantities of gas mere thus extracted, calculated per gram of carbon, so as to com- pare the figures with Lowthian Bell's results :- At ordinary temperatures. 15' tco 100". 100" to 250'. 250" to 360". Totd.Carbon dioxide. . Nil. 2.0 2 *o 3.3 7.3 Carbon oxide. ... 3.0 0.5 traces. 0.5 4-0 11.3 -- It is evident from the circumstance that carbon oxide was actually extracted at the ordinary temperature, that this gas, at any rate, must have been occluded ; whilst the fitful expulsion of the remaining gaq on raising the temperature by successive stages, renders it much more probable that the rest of the gas was also occluded than that it was due to the action of the carbon on the iron oxide. The gradually in- creasing amounts of carbon dioxide are doubtless due to the partial oxidation of the carbon oxide by the hot. iron oxide, The much smaller amount of occlusion observed in the carbon experimented with by us as above described, is probably due simply to alteration of texture during the process of freeing from reduced iron, &c.§ 7. DISCUSSION OF THE ABOVE RESULTS. It is abundantly manifest that the temperature of initial action of carbon oxide and of hydrogen on all the metallic oxides exaniined, varies considerably with the physical state of the metallic oxide ; that of carbon varies not only with the physical state of the metallic oxide, but also with that of the carbon, a light pulverulent carbon beginning to act at a much lower temperature than a hard, dense, coke-like26 WRIGHT AND LUFF ON RESEARCHES, ETC. charcoal, even though the latter be reduced to very fine powder. In the case of ferric oxide, although a difference in physical structure of the carbon produces a notable difference in temperature of initial action, yet no appreciable difference in temperature appears to be caused by any initial difference in the physical state of the ferric oxide, the reason for this presumably being that at a temperature of 400' and upwards the " precipitated " oxide alters its texture, and becomes in structure the same, o r nearly the same, as that produced in the first instance by ignition processes. This alteration in physical texture by heat alone does not seem t o take place, or at any rate not to so great an extent, with oxides of copper. Nextly, on comparing the temperatures of initial action on a, given kind of metallic oxide of different reducing agents, it is invariably found that that reducing a p z t b e g i m t o act at the lowest temperature which has the greatest heat of combustion); so that the heat disturbance during its reaction has (algebraically) the greatest value. Thus hydrogen always begins to act a t a lower temperature than carbon, and carbon oxide at a lower temperature than hydrogen :- Carbon oxide. Cupric oxide A . . .. 60" ,, C . . . . 146 Cuprous oxide .... 110 Ferricoxide A .... 202 ,, B . . .. 125 ,, B .... 90 ,, c .... 280 Hydrogen. 85" 175 172 155 260 195 245 Sugar carbon. Carbon from CO. 390" 350 O 430 350 440 430 390 345 450 430 450 - 450 430 Again, when cupyic and ferric oxides, prepared by analogous pro- cesses, and, therefore, presumably in much the same physical state, are compared, it is uniformly found that the temperature of initial action of a given reducing agent is lower o n oxide of copper than o n oxide of iron; i e . , that the action commences a t a lower temperature the greater (algebraically) the value of the heat disturbance :- Jetcnllic Oxides pwpared by precipitation. Carbon osidc. Hydrogen. Sugar carbon. Carbon from CO. Copper .......... 60" 85" 390" 350" Iron.. ............ 90 195 450 430 ?* Jltallic Oxides prepared by ignition of Xalts, PG. Copper .......... 125' 175" 430" 390" 202 260 4.50 430 { 220 245 450 430 Iron. ............. Q Deduced from the results with the other specimens.MUIR ON THE INFLUENCE EXERTED, ETC. 27 That this rule is not invariable for all metallic oxides has been already indicated (§ 1). It is noticeable that cuprous oxide, having much the same “heat of formation’7 as cupric oxide (0 3, ii)7 and, therefore, causing much the same heat disturbance during its rednction by a given reducing agent, is uniformly first reduced at a temperature no further removed from the temperature of initial action on one or other form of cupric oxide than may readily be due to difference in physical state.

ISSN:0368-1645    

DOI:10.1039/CT8783300001

出版商:RSC    

年代:1878

数据来源: RSC

摘要:

MUIR ON THE INFLUENCE EXERTED, ETC. 27 II.-On the In$u,ence ezerted by Time and Mass irt certain Reactions in, which insoluble Salts are produced. By M. M. PATTISON MUIR, F.R.S.E., Przlector in Chemistry, Caius College, Cambridge. 1. IF two salts can, by an interchange of acids and of bases, produce an insoluble compound, that insoluble compound is actually produced when solutions of the salts in question are mixed. The first action of the twQ salts upon one another results in the production of a small quantity of the insoluble compound, which is immediately removed, by precipitation, from the sphere of action ; the original conditions being thus partially restored, a further quantity of the insoluble compound is produced, and this alternate maintenance and disturbance of equi- librium is carried on until the whole of the insoluble compound which can be produced has actually been produced.2. In a paper published in 1857 (Chenz. Xoc. J., ix, 54), Gladstone has shown that when two salts are mixed under circumstances such that all the resulting bodies are free to act and react, the proportions in which the various bases and acids enter into combination are partly dependent upon the mass of each substance in the mixture, and that " an alteration in the mass of any of the compounds present alters the amount of every one of the other compounds, and that in a reguIarly progressive ratio further, that this " equilibrium of affinities '' is in certain cases not attained " for hours, or even days." Now, although when an insoluble compound is produced, the eyui- librium which was first established is destroyed by the removal of the new salt from the sphere of action ; and, although the process of pre- cipitation usually takes place with so much rapidity as to forbid any measurement of the time occupied, chemists, nevertheless, generally, I think, believe that the ma.ss of the so-called precipitant influences the amount of the precipitate, and also the time during which the process28 XCIR ON THE INFLUESCE EXERTED BY of precipitation proceeds.I n a note to the paper referred to, Glads tone remarks: “ I t is easily conceivable that when the affinity for each other of the two substances which produce the insoluble compound is very weak, the action may last some time, and become evident to our senses.Is not, this actually the case wheii . . . . carbonate of soda in solution is added to chloride of calcium ? ” &c. 3. In the following quantitative measurements of the influence of time in precipitation, aiid of mass of the precipitant on the amount of insoluble compound produced, I have availed myself of the suggestion made in the foregoing quotation from G1 a d s t o n e ’s paper. A solution of calcium chloride of known strength was prepared by dissolving a weighed quantity of pure calcspar in hydrochIoric acid, evaporating repeatedly to dryness after successive additions of water, care being taken to avoid loss by spirting, and making up to a deter- minate volume. Solutions of pure potassium and sodium carbonates, of known strength, were also prepared.A certain volume of the carbonate solution was diluted with a fixed volume of water, a mea- sured quantity of calcium chloride solution was run in, and the contents of the beaker were shaken up once ; after standing for a fixed number of minutes the precipitate was collected, washed, and dissolved in it measured volume of standard hydrochloric acid, the liquid was boiled to expel carbon dioxide, aid, after cooling, the excess of acid was determined by means of half-normal ammonia solution. The results can only be regarded as approximate ; nevertheless, for the object which I set before myself, they are, I thiuk, almost as in- structive as if their perfect accuracy were unimpeachable. The circumstances preventing great exactness in the results, are, principally, the fact that the process of filtration occupied an appreci- able time, usually five to six minutes, aiid the fact that calcium car- bonate is slightly soluble in water, The inaccuracy introduced by the first circumstance attaches equally to all the experiments, inasmuch as the volume of liquid t o be filtered was in each case the same, and the actual duration of the process of filtration did not vary more than one minute.The amount of wash water employed was also approximately the same in each case, hence the inaccuracy arising from the second circumstance mentioned above also attaches equally to all the experi- ments. The actual amount of calcium carbonate dissolved could not in any case cause a greater error than 1 per cent. on the final result. The results are, I think, strictly comparable among themselves, and the absolute errors are not considerable.4. I will first consider the influence of time :-TIME AND MASS IN CERTAIX REACTIOXS, ETC. 2 9 TaUe I. Total volume of liquid = 70 C.C. Proportion CaC12 : Na2C03 : : 1 :.l molecule. 10 C.C. CaC1, solution = 107.062 rngni. CaC1,. 14 C.C. INaJ3.3, solution. Percentage of CaCO, produced, calculated on total theoretical yield. Time. 5 minutes. 74.53 30 7 7 84.68 180 ? 7 93.1 7 46 hours. 94.2 7 60 3 7 85.91 Table 11. 10 C.C. CaCI, solution. 28 C.C. Na2C03 solution. Na2C0, : CaC12 : : 2 : 1 molecule. Volume = 70 C.C. Time. 5 minutes. 30 1 3 60 ? 7 180 , Percentage of CaCO, produced. 99-08 93.1 7 94.27 100.00 Table 111. 10 C.C. CaC1, solution. 42 C.C. Na2C03 solution.Na,CO, : CaCl, : : 3 : 1 molecule. Volume = 70 C.C. Time. Peraentage of CaCO, produced. 5 minutes. 94.27 60 37 96.87 30 > ? 95.3 7 180 > > 100~00 Table IT. 10 C.C. CaCl, solution. 56 C.C. Na2C03 solution. Na2C03 : CaC1, : : 4 : 1 molecule. Volume = 70 C.C. Time. 5 minutes. Percentage of CaCO, produced. 100.00 Table V. 10 C.C. CaCl, solution. 19.5 cc. Kz,CO, solution. K2C03 : CaC1, : : 1 : 1 molecule. Volume = 70 C.C. Time. 5 minutes. 30 ? 9 60 7 ) 180 7 ) 46 hours. Percentage of CaC03 produced. 81-11 85.00 85.91 94.2 7 94.2 730 MIUI1C ON TKE INFLUENCE EXERTED BY Tuble VT. 10 C.C. CnCI, solution. 39 C.C. K,CO, solution. KR,C03 : CaC12 : : 2 : 1 molecule. Volume = 70 C.C. Time. Percentage of CaCO, produced. 5 minutes. 94.27 30 ?, 98-65 60 7 9 100~00 T d e VII.10 C.C. CaCl, solution. 58.5 C.C. K,CO, solution. K,C03 : CaC12 : : 3 : 1 molecule. Volume = 70 C.C. Time. Percentage of CsC03 produced. 5 minutes. 96-45 30 ,> 98.65 60 7 9 100*00 Table VIII. K&O3 : CaCl, : : 4 : 1 molecule. 10 C.C. CaCI, solution = 57.045 mgm. CaC12. 40.9 C.C. K&O3 solution. Volume = 70 C.C. Time. 5 minutes. Percentage of CaC03 produced. 100*00 The results contained in Tables I and I1 are represented graphically by curves C and D respectively: the results in Tables V and VI by curves A and B respectively. It is evident that the greater portion of the chemical change is pro- duced during the first five minutes of the action ; after the expiry of that time the action very much decreases in rapidity ; the amount of change in the next 25 to 30 minutes is, however, always greater than the amoxnt accomplished in the second period of the same duration.The action proceeds, as it were, with a rush at first ; it then gradually becomes more and more slow. 5 . The numbers obtained also illustrate the influence exerted by in- creasing the mass of one of the salts. This influence will be rendered more apparent by arranging the numbers in a somewhat different form :- Table IX. Details as before. Time = 5 minutes. 1 : 1 mol. 74.53 2 : 1 do. 92-08 3 : 1 do. 94.2 7 4 : 1 do. 100~00 NazCO,; : CaCI,. Percentage of CaCO, produced.30 MIUI1C ON TKE INFLUENCE EXERTED BY Tuble VT. 10 C.C. CnCI, solution. 39 C.C. K,CO, solution. KR,C03 : CaC12 : : 2 : 1 molecule. Volume = 70 C.C. Time. Percentage of CaCO, produced.5 minutes. 94.27 30 ?, 98-65 60 7 9 100~00 T d e VII. 10 C.C. CaCl, solution. 58.5 C.C. K,CO, solution. K,C03 : CaC12 : : 3 : 1 molecule. Volume = 70 C.C. Time. Percentage of CsC03 produced. 5 minutes. 96-45 30 ,> 98.65 60 7 9 100*00 Table VIII. K&O3 : CaCl, : : 4 : 1 molecule. 10 C.C. CaCI, solution = 57.045 mgm. CaC12. 40.9 C.C. K&O3 solution. Volume = 70 C.C. Time. 5 minutes. Percentage of CaC03 produced. 100*00 The results contained in Tables I and I1 are represented graphically by curves C and D respectively: the results in Tables V and VI by curves A and B respectively. It is evident that the greater portion of the chemical change is pro- duced during the first five minutes of the action ; after the expiry of that time the action very much decreases in rapidity ; the amount of change in the next 25 to 30 minutes is, however, always greater than the amoxnt accomplished in the second period of the same duration.The action proceeds, as it were, with a rush at first ; it then gradually becomes more and more slow. 5 . The numbers obtained also illustrate the influence exerted by in- creasing the mass of one of the salts. This influence will be rendered more apparent by arranging the numbers in a somewhat different form :- Table IX. Details as before. Time = 5 minutes. 1 : 1 mol. 74.53 2 : 1 do. 92-08 3 : 1 do. 94.2 7 4 : 1 do. 100~00 NazCO,; : CaCI,. Percentage of CaCO, produced.TIME AXD 31ASS IN CERTAIN RELACTIONS, ETC. 31 Time = 30 minutes. 1 : 1 mol. 85.00 2 : 1 do. 93.1 7 3 : 1 do.95.37 4 : 1 do. 100.00 Time = 60 minutes. 1 : 1 do. 85.91 2 : 1 do. 94.2 7 3:l do. 96-87 4: I do. 100*00 Time = 180 minutes. 1 : 1 do. 93.17 2 : 1 do. 100~00 Time = 46 hours. 1 : 1 do. 94.2 7 2 : 1 do. 100.00 Table X. Details as before. Time = 5 minutes. 1 : 1 mol. 81.11 2 : 1 do. 94-2 7 3 : 1 do. 96.45 4:l do. 100~00 Time = 30 minutes. 1 : 1 do. 84-68 2 : 1 do. 98.65 3 : 1 do. 98.65 4 : 1 do. 100*00 Time = 60 minutes. 1: 1 do. 85.91 2 : 1 do. 100*00 Time = 180 minutes. 1 : 1 do. 94-2 7 2 : l do. 100*00 Time = 46 hours. 1 : 1 do. 94.27 2 : 1 do. lO0~00 K&03 : CaC12. Percentage of CaC03 produced. The first two sections of Table IX are represented graphically in the curves marked F ; the first two sections of Table X are represented in the curves marked E.33 MUIR OX THE INFLUENCE EXERTED.BY 6. These numbers show that the equation- CaC1, + iU,CO, = 2MC1 + CaCO,, does not furnish a full expression of the action of sodium or potassium carbonate upon calcium chloride. When the two salts are mixed in the proportion expressed by their respective formul;t3, the action repre- sented in the equation is n o t completed, even after the expiry of so long a time as 46 hours. In order fully to realise the equation, the maps of the alkaline carbonste must exceed. that which -would be actually required to transform the wholc of the calcium chloride into carbonate were the equation strictly true. When two molecules of potassium carbonate are presentl to one molecule of calcium chloride, the action is complete after the lapse of an hour ; were sodium carbonate employed, it would be necessary to use rather more than three molecules for one of the alkaline earth chlode, in order to complete the action in the same period of time.If the mass of the alkalinecarbonate be four times that represented in the equatior,, the action is complete in five minutes. 7. So soon as the solutions are mixed an atomic interchange begins, resulting, we must believe, in the formation of calcium carbonate, which, being insoluble, is at once removed from the sphere of action, and of sodium or potassium chloride, while at the same time portions of sodium or potassium carbonate and of calcium chloride remain un- decomposed. As the mass of sodium or potassium chloride increases, the action of the small residual portions of sodium or potassium car- bollate upon the remaining calcium chloride becomes slower and slower, until it filially ceases.It may, perhaps, be allowable to imagine the few calcium chloride and sodium or potassium carbonate molecules as so widely separated from one another by the great mass of sodium chloride molecules, that among their excursions t o and fro they rarely, if ever, come into colli- sion with one another. The addition of a further quantity of the sodium or patassium carbonate molecules of course increases the chances in favour of any one of these molecules meeting with a molecule of calcium chloride, and hence the chances in favour of the production of fresh quantities of calcium carbonate. Again, it is not improbable that a secondary reaction may take place to a limited extent, represented by such an equation as- (1.) CaCO, + 2MC1 = CaC1, + M2C03.If this reaction does take place TF: should expect that so soon as the M,CO, produced reaches a certain amount, the converse reaction- (2.) CaC1, + M,CO, = CaC03 + 2MC1 Tvould again begin, and that these two reactions would maintain the composition of the mixture constant. But if we increase the mass ofTIME AND MASS I N CERTAIN REACTIONS, ETC. 33 &co3, it is clear that the second reaction will preponderate over the first, and if we'ensure that when the second reaction is finished, a con- siderable number of M2C03 molecules remain unacted upon, the con- ditions under which reaction (1) can be realised will no longer prevail. 8. Increasing the mass of the individual molecules of the precipitant tends to increase the amount of chemical change brought about in a stated period of time.The molecular weights of sodium and potassium carbonates (sup- posing the ordinary formulae to be really molecular formuke) are to one another as 53 is to 69.1. The numbers obtained show that when potassium carbonate is used, a larger quantity of calcium carbonate is formed in a given time than when sodium carbonate is employed, except when the action is allowed to proceed during very extended periods. After 46 hours the amount of calcium carbonate formed by the action of sodium carbonate was the same as that produced by the action of potassium carbonate. In the following table I have arrangcd the numbers so as to illustrate the point under consideratiofi.Table XI. Details as before. Percentage of CaC03 produced. / A \ Time = 5 minutes. Time = 30 minutes. Proportion of -- -- M2C03 : CaC12. Na2C03. K,C03. Na,C03. K2C03. 1 : 1 mol. 74.53 81.11 84.68 85.00 2:l ,, 92.08 94.3 7 93.1 7 98-65 3:1 ,) 94.27 96.45 95.37 98-63 4 : l ,, 100~00 100*00 100*00 100.00 Time = 60minutes. Time = 180minutes. P -- 1:l ,, 85.91 85.91 93.17 94.27 2 : l ,, 94.27 100*00 100.00 100*00 3:l ,) 96.87 100.00 4 : l ,, 100~00 100*00 Time = 46 hours. 1:l ,, 94.27 94.27 9. If the hypothesis put forward in paragraph 7 be correct, it is evident that any circumstances which would increase the excursions of the residual molecules of sodium or potassium carbonate and of calcium chloride, would also increase the chances in favour of collision occur- ring between these molecules, and hence the chances in favour of the production of larger quantities of insoluble calcium carbonate.On the other hand, any circumstances tending to decrease the excursions of VOL. XXXIII. D34 MUIR ON THE INFLUENCE EXERTED BY these residual molecules, or tending to increase the number of molecules which take no part in the production of calcium carbonate, and hence to decrease the chances of the sodium or potassium carbonate and the calcium chloride molecules coming into collision, would also tend to diminish the amount of calcium carbonate produced. Raising the temperature a t which the two salts are allowed to react upon one another would tend to increase the number of excursions of the differ- ent molecules ; diluting the solutions would tend to decrease the chances of collision between the two sets of molecules.I have carried out a few experiments under each of these conditions. The experiments already detailed were performed at the ordinary tem- perature, vie., 15" to 20". TabZe XII. Details as before. Time = 5 miuutes. Temperature = 40". Percentage of calcium carbonate produced. M2C03 : CaC1,. Na,CO,. KZCOS. 1 : 1 mol. 88.50 85.50 2 : l ,, 93.17 95.36 3 : l ,, 96.45 96.45 4 : l ,, 100~00 100~00 Table XIII. Details as before. Volume = 140 C.C. Percentage of calcium carbonate produced. r f-\ Time = Time =: Time = 5 minutes. 30 minutes. 60 minutes. M ~ C O ~ : CaC12. N&CO,. K,C03. Na2CO3. Na2COy. 1 : 1 mol. 63.58 63.58 74-05 75.64 2 : 1 ,, 78.92 87.69 86.84 86.84 3 : 1 ,, 86.84 - 88.25 91.05 Increase of temperature causes in every case an increase in the amount of calcium carbonate produced.Dilution causes a very marked decrease. A small proportion of this decrease may be traced to the solvent action of the larger volume of water upon the calcium carbon- ate precipitated ; but there can be no doubt that dilution, apart alt,o- gether from the solvent action of the diluent upon the precipitate, very materially decreases the amount of chemical chmge taking place in a definite interval of time. 10. If sodium or potassium chloride were employed as diluent inTINE AXD JiBSS IN CERTAIY REACTIONS, ETC. 35 place of water, we should expect to find a marked decrease in the amount of calcium carbonate formed, inasmuch as the alkaline chlor*ide would not only act as a diluent, but it would also tend to increase the action represented in equation (1) in paragraph 7.The following table contains the results of some experiments performed with the view of testing the reasonableness of this expectation. TcLble XIV. 10 C.C. CaC1, solution, with 16, 32, and 48 C.C. K,CO, solution : volume made up to 70 C.C. with dilute NaCl solution (1 C.C. = 8.5 mgm. NaC1). Percentage of CaC03 produced. ~~ M2C03 : CaC1,. Time = 5 minutes. Time = 30 minutea. 1 : 1 mol. 64.43 6 4-43 2 : l ,, 84.05 92.44 3 : l 7 , 92-49 95.24 It is well known that alkaline chlorides exert a solvent action upon calcium carbonate ; but the decrease in the amount of calcium carbon- ate produced in these experiments as compared with that formed in experiments where water alone was employed as diluent, cannot be entirely traced to such an action.The alkaline chloride appears rather to prevent the formation of calcium carbonate. Even were we to regard the action of the chloride as that of a solvent, we must surely believe that the calcium carbonate is transformed info chloride, and passes into solution in this form. 11. T t occurred to me that discontinuous addition of one of the solutions would probably bri,ng about the production of larger quanti- ties of calcium carbonate than are found when the whole of one solu- tion is added to the whole of the other at one time. A few experi- ments were tried ; the results are summarised as follows :- Table XV. Details as before. CaClz solution added in small successive portions during 34 minutes ; liquid shaken up after each addition ; whole allowed to remain at rest during 18 minutes after final addition of calcium chloride.Percentage of CaC03 produced. I ~~ MM,C03 : CaC12. Na2C03. K2C03. 1 : 1 mol. 89-00 89.65 2 : l ,, 100.00 100-00 CaCl, solution added in small successive quantities during 20 mi- nutes, whole allowed to remain a t rest during additional 10 minutes. D ’ 236 MUIR OX THE ISFLUENCE EXERTED BY Na2C03. &GO.+ 1 : 1 rnol. 94.2 7 94.2 7 2 : 1 do. 100-00 100*00 CaC1, solution added in small successive quantities during 150 minutes, whole allowed to remain a t rest during additional 30 minutes. 1 : 1 mol. 94.2 7 94.2 7 2 : 1 do. 100*00 100~00 Na2CO3 KZC03 These results show that when one of the reacting sohtions is added discontinuously, the action reaches a maximum more quickly than when the solutions are mixed at one time, but that the maximum so reached is no greater than that which is finally attained under the latter conditions.12. I have also carried out a few experiments using ammonium, sodium, and potassium oxalates as precipitants, and determining the amount of calcium oxalate produced by dissolving in warm sulphuric acid and titrating with potassium permanganate. The results are not, however, of much interest, as the action cf time and of mass is not so marked as in those experiments which hare been already detailed. The chemical change is much sooner completed when oxalates than when carbonates are employed as precipitants.I give the results of my experiments so far as they have been pursued. Tuble XVI. 20 C.C. CaCl, solution = 114.09 mgm. Temperature = 15" to 20". Volume = 70 C.C. Time = 5 minutes. Percentage of CaC204 produced. 7 ~~ M,C,O~ : CaCl,. (NHJ~C~O,. ~ x%c,o,. K2C204. 1 : 1 mol. 93.85 93-85 93.85 2 : 1 do. 100*00 100*00 100.00 3 : 1 do. 100~00 100~00 100.00 13. I n a paper entitled " On the Application of Liquid Diffusion to produce Decompositions " ((Them. SOC. J., iii, 601, Graham describes certain experiments from which " it follows that cold solutions of sul- phate of lime and chloride of sodium may be mixed without decom- position, or without any sensible formation of sulphate of soda." Graham further remarks, that "it would be interesting to submit such a mixture t o a diffusion experiment after being kept for different periods.The effects of time and temperature are so often convertible that we might anticipate a gradual formation of sulphate of soda." I have carried out Graham's idea in a somewhat modified form.TIME AND MASS IN CERTAIN REACTIONS, ETC. 37 Although the results of the experiments to be now described cannot strictly be brought under the heading of the present paper, inasmuch as no insoluble salt is produced in the reaction, nevertheless, as the results illustrate the influence of time and mass, and are also interest- ing from another point of view, I venture to detail them. Solutions of calcium sulphate and sodium chloride were mixed and allowed to remain in covered beakers during four weeks. Measured portions were then withdrawn, and to each a volume of alcohol, fully equal to the volume of the liquid itself, was added.The precipitated calcium sulphate was collected, washed with alcohol, ignited, and weighed. An equal volume of the calcium sulphate solution was heated with alcohol in the same way, and the calcium sulphate con- tained therein was thus determined. The results axe contained in the following table:- Table XVII. (1.) 100 C.C. CaSOa solution = 191.0 mgm. CaS04 +- 10 C.C. NaCl solution = 233.3 mgm. NaC1. Amount of CaS04 in solution after 28 days = 167.0 mgm. Amount of CaS06 in solution after 28 days = 103.0 mgm. Amount of CaSOa in solution after 28 days = 64.0 mgm. (2.) 100 C.C. CaS04 solution + 50 C.C. NaCl solution. (3.) 100 C.C. CaS04 + 100 C.C.NaUl solution. A very considerable decomposition of the calcium sulphate by the sodinm chloride had evidently taken place. If the results be arranged in the form adopted in many of the pre- ceding tables, the influence of mass on the amount of chemical change becomes more apparent. Temperature = ordinary. Time = 28 days. CaSO, : NaC1. 1 : 1.42 mols. 1 : 7.1 mols. 1 : 14-2 mols. The amount of chemical change varies almost directly with the mass of the sodium chloride employed. I intend t o carry out a series of experiments upon the influence of time and mass in bringing about the change under consideration, and hope to communicate the results to the Society at no distant date. Graham’s experiments show that on heating a mixture of calcium sulphate and sodium chloride, sodium sulphate is formed, and con- tinues to exist in the cold solution. Amount of CaSO, decomposed = 3.80 per cent. of Amount of CaS04 decomposed = 19.15 per cent. Amount of CaSOI decomposed = 32-98 per cent. the total CaS0, present.38 RIXGZETT O S THE The experiments detailed above show that a similar formation of sodium sulphate occurs, without heating, if solutions of the two salts be allowed to react upon one another for several weeks. The secondary consideration which lends interest to these experi- ments may be stated in Graham’8 words : “ If such formation of sodium sulphate be the case, we have an agency in the soil by which the alkaline carbonates required by plants may be formed from the chlorides of potassium and sodium, as well as from the sulphates of potash and soda : for the sulphate of lime generally present will con- vert these chlorides into sulphates.” This observation derives its force from the facts demonstrated by Graham, that sulphates of potassium and sodium are deconiposed by lime-water, yielding diffusates which contain caustic potash and caustic soda respectively, but that clilorides of potassium and sodium are not so decomposed.

ISSN:0368-1645    

DOI:10.1039/CT8783300027

出版商:RSC    

年代:1878

数据来源: RSC

摘要:

38 RIXGZETT O S THE 111.-The Chemistry of Cocoa Butter. Part I. Two New Fatty Acids, By CHARLES T. KINGZETT. SOME few months ago I commenced, at the suggestion of my friend, Dr. B. H. P a u l , an investigation upon the chemistry of cocoa butter, and although my work has had many interrupt'ions, some interesting and important results have been attained ; so important indeed have been these results, that I am induced to publish this paper, which otherwise must be regarded 8s of a preliminary nature. The sample of cocoa butter upon which I worked was obtained from Messrs. Cad b u r y and Sons. It was a hard, imperfectly transparent, slightly yellowish solid, melting at about 30" C., and when once melted, retaining the liquid state for some time a t a temperature below that just mentioned.To alcohol of 85 per cent. and maintained at the boiling point, it gave up but traces of soluble matter, and from the fact that hot ammonia water extracted from it but traces of matter, I conclude that it is practically or entirely free from fatty acids, as such. It was proved by specia,l experiments to be also entirely free from volatile or soluble fatty acids. 1st Experimeizt.-40 grams were saponified by caustic soda solution, the mixture nearly neutralised with dilute sulphuric acid, and the precipitate of sodium salt washed after filtration. It was dissolved in much boiling water, and was then precipitated by barium chloride solu- tion.CHEMISTRY OF COCOA BUTTER. 39 An attempt was made to estimate the oleate of barium contained in this salt, by extraction with boiling alcohol ; but that employed (of 85 per cent.) dissolved some considerable quantity of a barium salt other than the oleate, whose free acid was solid at ordinary temperatures.1 also attempted to estimate the oleic acid by converting the barium salt into lead salt, and extracting this by ether, but in course of the extraction the salt became yellow, and something more than mere oleate of lead dissolved, The lead precipitate insoluble in ether was reconverted into barium salt, through the intervention of an ammurium soap made from the free acids and the barium salt analymd- (1.) 0.495 grm. dry at 100" C. gave 1048 grm. CO,, and 0,427 grrn. H,O. (2.) 0.4190 grm. gave 0.1490 BaSo4. Synopsis of Analysis :- c ........ 57.74 H.. ...... 9*5b Ba ...... 20.90 0 ........11-78 If we deduct the barium we get by calculation the following per- centages for the free acid :- C ........ 73.00 H ........ 12.11 0.. ...... 14-89 The barium salt was now decomposed by hydrochloric acid in pre- sence of ether, and the free acids obtained by distillation of the ether were recrystallised from alcohol. The first crop was lost ; the second, which was evidently impure (being slightly colonred), was analysed ; 0.212 grm. gave 0.545 grm. CO, and -237 grm. HzO. c ........ 70.11 H.. ...... 12.42 0 ........ 17.47 2 . d Eqeriment.-Another quantity of cocoa butter was now sapo- nified and the sodium salt made into barium salt. This was decom- posed by hydrochloric acid in presence of ether, and the fatty acids obtained by distillation of the ethereal solution crystallised from 85 per cent.alcohol containing animal charcoal in suspension ; the last mother-liquor was retained. The solid acid obtained, as described, was a snow-white powdery substance when dry. I t was recrystallised from alcohol, two successive crops being ob- tained ; the first by deposition ; the second after concentratiw.40 KINGZETT ON THE The first crop had a melting point of 65". (a.) 0.164 grm. dried at 100" C., and burnt with PbCr04, with the (t.) 0.125 grm. gave 0.147 grm. HzO and 0.353 grm. CO,. aid of KC103, gave 0.197 grm. H,O and 0.464 grm. GO,. Synopsis :- a. b. C.. ...... 77-16 77.01 H ...... 13.35 13.06 0 ...... 9-49 9.93 This same mixture of acids (for such it eventually proved to be) was afterwards fractionated by recrystalIisation from spirit! into three portions, 1, 2, and 3, which were analysed. Fraction 1.Melted at 68". (a.) 0.185 grm. gave 0.223 grm. HzO, and 0.521 grm. COz. (6.) 0.159 grm. gave 0.188 grm. HzO. C ........ 76.81 H.. ...... 13*35 and 13.14. 0.. ...... 9.84 Fraction 2. Melted at 61". 0.138 grm. gave 0.3960 GO, and 0.162 H20. C ........ 78.26 H... ..... 13.04 0 ........ 8.70 &Taction 3. Melted at 58". 0.157 grm. gave 0.410 grm. CO,, and 0.182 H,O. c ........ 71-22 H ........ 12.87 0 ........ 15.91 The second crop (following that which furnished these three fmc- tions) was fractionated into two portions, each melting at 58". The oleic acid contained in cocoa butter was isolated from the last mother-liquor alluded to above, and obtained after these successive crops of solid fatty acids.It was made into ammonium salt, and this into barium salt, which was dried and extracted by alcohol. The powder thus obtained on cooling of the alcoholic extract was snow- white. 0.322 gave 0.105 BaSOa = 19.25 per cent. Ba. Inasmuch as oleate of barium gives 19-59 per cent. Ba, and further seeing this acid had also the other properties of ordinary oleic acid, there is no reasonable doubt left regarding its true nature, 3rd Eyerimed.-Another quantity of cocoa butter was now sapo- nified; the sodium soap washed slightly, and decomposed by sul-CHEMISTRY OF COCOA BUTTER. 41 phuric acid ; the free acids were washed by repeated meltings with hot water, then dried and weighed. It was recrystallised from alcohol in the presence of charcoal, and the mother-liquors treated as described below.Four successiva crops of fatty acids were obtained : ( a ) shortly after cooling to some extent; ( b ) on standing ; (c) by concentration of mother-liquor ; ( d ) after last fraction on standing. The first crop was further fractionated by alcohol into three portions, and the third crop into two portions; the mother-liquors being for the time dis- regarded. Weight = 37 grms. Examination, of the three fractions from 3rd crop :- (1.) Melted at 70.5". (2.) Melted at 70.8". (3.) Melted at 61.0". The first portion melting at 70.5" C. was aKain fractionated with two crops. (a.) Melting at 71-71.5" C., and ( b ) melting at 71-7'1.5" C. Analysis of ( a ) :- 0.179 grm. gave 0.518 C 0 2 , and 0.219 H,O.C ........ 78-92 H ........ 13.59 0 ........ 7.49 This product was eventually further fractionated into two portions, It melted at 7 2 ~ 2 " ~ solidified again at 69". (a.) 0.170 grm. gave 0.515 GO2 and 0.213 H,O. (b.) 0-069 ,, 0.208 ,, ,, 0.088 ,, (a.) C. 82-62 ( b . ) 82.21. H. 13.92 14.16. 0. 3.46 3.63. the first of which was analysed. !#?he second crop melted at 66" C. The third crop melted at 56.5" C. This last crop was further fractionated into two portions and a Portion ( a ) melted at 59 to 59.5" C., and ( b ) at 57.5 to 58" C. Analysis of ( b ) :- 0204 grm. gave 0.542 CO, and 0.257 H,O. mo ther-liquor. C ........ 72-46 H ........ 13-99 0 ........ 13-55 To recapitulate, we may here show in tabular form the percentage composition and melting points of the foregoing products :-42 KlSGZETT ON THE 1.M. 11.. . 57-58' C .... 70.11 H .... 12.42 0 .... 17'47 2. 3. 1 4. 5. 6. 7. 8. ------- _------- --- 58.5' 68" 65' 61" 71.5' 57'50 i - 71 '22 '72.46 I $3.00 $76'81 77-16 78-26 78.92 12 '87 13.99 , 12.11 13-14 13.06 13'04 13-59 15.91 13'55 14.89 10*03 9.88 8.70 7'49 - 9. 72 '2 82 -62 13 9 2 3 *46 Neglecting the intermediate terms, and taking the extremes reacted by the method employed, these latter are tolerably well represented by the formulx?, C12H2,02 and C6,H,,02 :- c 12H?40?. Cfi4H12802* c ...... 72.00 82.75 H ...... 12.00 13.79 0 ...... 16-00 3.46 100.00 100.00 Now of the products whose analyses are described above, only one was a t all coloured, and that was the one containing 70.11 per cent. ; moreover, this was the only one whieh was burnt with PbCr04 with- out the supplemental aid of potassic perchlorate, therefore, neglecting this, the nearest analyses approaching those just given are- C12H2102.C6,&3O,. C ...... 71-22 82-62 H ...... 12.87 and 13.92 0 ...... 1591 3.46 100-00 100~00 -- C,,H2,02 is the formula of lauric acid, which, however, has a melt- ing point of 43". As the lowest melting point I have yet observed among the fatty acids from cocoa butter is 57.5", the fatty acid to which it belongs. if lauric acid, must contain some quantity of an acid of higher melting point than lauric acid. Therefore, the acid itself would be lower in the series CnH2n02 than lauric acid. Further, the highest acid previously known in the series C,H2,02, is melissic acid, C30H6002, and whereas I have described a definitely crystalline substance of the empiric formula C61H,2802, it is by far the highest acid term in the series, yet discovered.But I would not claim to have attained as yet the ultimate limits in either direction ; the higher acid may contain more than 64 atoms of carbon to the molecule, and the lower one may contain less than 12 atoms to the molecule. I have, however, exhausted the capabili- ties of the only process yet used by me ; in the future, I intend to apply whatever other methods are available among those known, or in my power to devise.CHEMISTRY O F COCOA BUTTER. 43 It onlyremains for me to add that I have prepared barium, calcium, and silver salts of the higher acid, but refrain from publishing their percentage composition until more ultimate possibilities may have been attained.I may mention, however, that the silver salt is colonred only slightly by light, and is deposited by precipitation from an ammoniacal solution by argentic nitrate as a gelatinous mass, which, when dry, is highly electric and quite insoluble in alcohol of 85 per cent, and in ether. The higher fatty acid itself also appears to be somewhat electric; it crystallises in granules from a concen- trated alcoholic solution, but in microscopic needles from a more dilute solution. At high temperatures it distils apparently unchanged. The crystalline character of the lower acid is more distinct, and some- times it is obtained in beautiful pearly plates, and at others in fine long needles. In the third experiment I determined, approximately, the amount of oleic acid contained in the 37 grams of total fatty acids from cocoa butter, by separating, as far as could be, all crystallisable solid ftitty acids from the ultimate mother-liquors after recrystallisation.It amounted to about 20 per cent. of the total. So far as I have ascertained, and merely for the sake of a calcu- lation, it may be stated that a compound glyceride of the composition C J L 0 2 C&L CtxHin02 { Cl8H3.302 requires exactly 20 per cent. oleio acid to be furnished by the total acids derivable from it. At the same time, it may be that cocoa butter is a mixture of the t'hree glycerides, or even of more, whose fatty acids in such case form part of the intermediate crops not yet analysed. Before concluding the paper, it should be stated that I have in vain searched chemical literature with the view of discovering any publi- cation bearing upon my subject.But the only matter I have been able to find is a statement made in nearly all works to the effect that cocoa butter yields almost exclusively stearic acid, and the evidence of its purity when prepared is a permanent melting point of 69 t o 70" C. It appears, therefore, that the melting point has hitherto been accepted as sufficient evidence of the presence of stearic acid ; but, as will be seen from my investigation, this is entirely incorrect. Compared with the formula of the highest product I have obtained from cocoa butter, the melting point of 72.2" appears to be low, but it is impossible that the acid can contain less than 64 atoms of carbon, and should its melt- ing point be regarded in any way as a difficulty, the acid could, of necessity only, be still higher in the series than so far ascertained. Without waiting until its molecular weight shall have been definitely44 WARINGTON OX NITRIFICATION. established, I propose for it the name of Theobromic acid, which recalls the source from which it is obtained, namely, the fat of the seeds of Theobronaa Cacao. The seeds are known to yield from 30 to 50 per cent. of this fat.

ISSN:0368-1645    

DOI:10.1039/CT8783300038

出版商:RSC    

年代:1878

数据来源: RSC

摘要:

44 WARINGTON OX NITRIFICATION. IT.- OTL Nitr@icatio?L. By Ro B E R T WAR I N G TON. I HAVE the pleasure of submitting to the Society a report of some experiments on nitrification recently carried out in the Rothamsted Laboratory. The formation of nitrates in soil, under favourable conditions as to moisture, temperature, and &ration, is a reaction well known and frequently studied. It is a process of great technical importance, being the only natural source of saltpetre. It is also of the highest interest to the scientific agriculturalist, nitrates being invaluable as plant-food, and their loss by drainage one of the most serious difficuI- ties with which a farmer has to contend. The bodies yielding nitric acid in the soil are, firstly, the various nitrogenous organic substances which arise from the decay of vege- table or animal matter; and secondly, ammonium salts, either pro- duced i n small quantity during the decay of organic matter, or carried t o the soil by rain, or, in some cases, applied intentionally as manure.A further source of the nitrates contained in soil is to be found, according t o some writers, in the free nitrogen of the atmosphere ; but of any supply from this source, other than the ready-formed nitrates contained in rain, there is atpresent no substantial proof. The production of nitric acid, both from ammonia, and from the organic matter of arable soil, is well illustrated by the field experi- ments at Rothamsted. The winter drainage-water from the various plots in the experimenkal wheatfield, is found to contain nitrates nearly in proportion to the quantity of ammonium-salts applied the preceding autumn (Jour.Roy. Agri. Soc., 1873, 334, 337).” While the drainage-water, rich in nitrates, which is obtained at the drain- gauges from soil which has been unmanured for the last eight years, clearly derives the greater part of its nitrogen from the organic matter of the soil. Many investigations have been made as to the mode in which nitri- fication takes place. A mixture of nitrogen and oxygen is known to yield nitrous acid when subjected to repeated electrical discharge, but * For further details see Jour. Chem. Soc., 1871, 286 ; also Sixth Report of the Riceys’ Pollution Commission, 58-68 ; and a lecture by Dr. Gilbert, at the Science Conference a t Sonth Kensington, ‘‘ On Some Points in connection with Vegetation,” page 39.WARINGTON ON NITRIFICATION. 45 ozone has been shown to be incapable of oxidising gaseous nitrogen. Ammonia, may he converted into nitrous and nitric acid either by electric discharges in the presence of oxygen, or by the action of ozone (L.Carius, Ann. Chenz. Pkarm., clxxiv, 31), but not by ordi- nary oxygen. Ammonia may also be oxidised into nitric acid, accord- ing to some observers (Pesci, Thenard, Knop), by prolonged con- tact with ferric oxide at ordinary teniperatures. With regard to the production of nitric acid from nitrogenous organic bodies, very little seems to be known. It is generally assumed that when such bodies decay in a porous medium, offering a sufficiently large surface for oxidation, nitrates must necessarily be formed, especially if a sali- fiable base be present ; but this view, as far as I am aware, has never been confirmed by exact experiments. The action of ordinary oxygen upon organic matter is supplemented, in the theory of some writers, by the action of ferric oxide and of ozone.In February of the present year S c h 108 s i n g and MU n t z laid before the French Academy (Conzpt. Tend., lxxxiv, 301) the results of an ex- periment, proving, in their opinion, that nitrification was due to the action of an organised ferment ; that, in fact, it was probably a function of some low form of vegetable life. This theory of nitrification they further state mas regarded as probable by Pasteur as far back as 1862. P a s t eur, in view of the active oxidation induced by mycoderms in various kinds of organic matter, expressed the opinion that nitrifi- cation required t o be studied over again from this point of view. As the experiment of Schloesing and Miintz formed the starting point of the trials at Rothamsted, it will be necessary to give a brief account of their research.The primary object of these experimenters was to ascertain if the presence of humic matter was essential to the purification of sewage by soil. To this end they took a glass tube one metre in length, and filled it with a mixture of ignited quartz sand and powdered lime- stone; 5 kilos. of sand, and 100 grams of lime-stone were employed. The column of purely mineral soil thus prepared was supplied with sewage at such a rate that the sewage took eight days to pass through the soil.During the first 20 days the sewage passed through unaltered ; but after that time nitric acid began to appear in the filtrate, and rapidly increased till the filtered sewage contained no longer any ammonia, but nitrates only. The authors ask-If it was a simple case of oxida- tion, why did 20 days elapse before the commencement of the reaction ? If, on the other hand, the seeding and growth of germs was necessary before the soil could become effective, the reason of the delay in nitri- fication is at once apparent. After passing sewage four months with complete oxidation of its46 WARINGTON ON NITRIFICATION. ammonia, the authors placed a small vessel of chloroform at the top of the column, and allowed the vapour of the cliloroform to pour con- tinuously upon the soil.&I3 n t z had shown in an earlier investigation (Cornpt. rend., lxxx, l250), that chloroform effectually suspends the action of orgsnised ferments, while it has little or no action upon soluble ferments. I n 10 days after the introduction of chloroform vapour all nitrates had disappeared in the exit water, and the sewage passed through the soil quite unchanged ; nitrification had thus plainly been effected by some agent, the activity of which was destroyed by the presence of chloroform. After supplying chloroform for 15 days, the vessel containing it was withdrawn. Nitrification did not resume with the removal of the chloroform. After about seven weeks, during which the sewage passed unchanged, the attempt was made to seed the colnmn of sand afresh with the organisms necessary for nitrification.For this purpose 10 grams of a vegetable soil, known to nitrify with ease, were washed with waster, and the turbid washings poured upon the column of sand. Eight days after this application nitrates again appeared in the exit water. As the sewage took eight days to traverse the column of sand, this was the shortest time in which nitrification could become appaPent. They con- clude that in the instance before them, and in other cases of rapid nitrification, oxidation is produced through the action of a living organism ; but they are careful not to assert that this is the only mode in which nitrification can take place. The importance of this new theory of nitrification is clearly very p e a t , As the amount of nitrates occurring in soils of known history is a t present being investigat'ed at Rothamsted, it was thought advisable to submit the new theory without delay t o the test of further experiment.The conclusion of Schlcesing and Miintz is clearly based on two distinct lines of proof. Firstly, the action of antiseptic vapours in preventing nitrification. Secondly, the possi- bility of inducing nitrification by seeding with R substance already nitrifying. We will describe, in the first place, the experiments made under the former head of inquiry, A soil was taken from the surfaceof a kitchen garden, the collection being made after heavy rain, when there was, consequently, but little nitrate present. The soil was air-dried, crushed, arid sifted, and the fine soil employed for the experiments.Two series of experiments were made with this soil, but as the general method was in each case the same, one description of the plan followed will suffice. Four glass tubes were taken, in each of which 100 grams of soil were placed. I n the first experiment these tubes were of unequal Such is the investigation of S chlcesin g a,nd JM u n t z.WARISGTOX ON NITRIFICATIJN. 47 length ; in the second trial the tubes were all 1 2 inches in height. In the first experiment the tubes were filled with the dry soil, aiid 10 C.C. of water gradually added afterwards ; the result was that only the upper half of the column of soil became visibly wet. In the second experiment the soil was uniformly mixed with 20 C.C.of water before its introduction to the tubes, the whole of the column was thus in a moist condition. The lower end of each of the four tubes was connected with a small water-pump, which served as an aspirator, while the upper end was connected with a U-tube containing pumice drenched with very dilute sulphuric acid. This U-tube ensured the supply of moist air to the soil, and cut off any ammoniacal vapours that might be present in the laboratory. In the case of three out of the four soil tubes a 10 oz. bottle, containing a few pieces of sponge, and provided with a cork and two tubes, was interposed between the soil tube and that containing pumice and sulphuric acid. The sponge in one of these bottles was moistened with a small quantity of liquid carbolic acid.The sponge of a second bottle received bisulphide of carbon. The third bottle received chloroform. The carbolic acid being little volatile, was not renewed during the course of an experi- ment. Of the bisulphide of carbon and chloroform 5 C.C. were used on the first day of the experiment, and on every seventh day a fresh quantity of 2 C.C. was placed on the sponge. As these liquids are very volatile, it is evident that during a great part of both experiments the amount of bisulphide of carbon and of chloroform carried to the soil must have been extremely small. The result of the whole arrange- ment was tthat through one soil tube moist air free from ammonia could be aspirated ; and that in the case of the remaining three tubes the moist air reached the soil more or less charged with antiseptic vapours.Aspiration was carried on for 1 to 1+ hour, through each tube, nearly every day. The first series of experiments commenced on June 29th, and con- cluded on August 7th, lasting in all 39 days. The second series commenced on September 24th, and lasted till November 9th, or 46 days. At the conclusion of both experiments the soil in each tube was extracted with hot water, the chlorine removed with sulphate of silver, and the nitric and nitrous acid present finally measured as nitric oxide gas after the method of Crum and Frankland.* A consider- able amount of froth was in most cases produced in the shaking tube, and in some instances the whole of the gas could not be removed from the shaking tube to the eudiometer ; the determinations which are too * The above is but an outline of the process used ; the exact mode of treatment found most suitable in determinations of nitrates in soil will probably be described on a future occasion.48 WARINGTON ON NITRIFICATION.low from this cause are marked in the following table with an asterisk. The determinations belollging to the second series of experiments wer'e, as the result of more experience, more accurately done than those in the first series. A summary of the results will be found in the following table :- Nitrogen as Nitrates and Nitrites per Million of air-dried Soil. First Second Original soil.. ................... 6.12" 8.91 Air passed ...................... 4G.87 50.86 History of Soil. Experiment. Experiment.,, with carbolic acid ...... 17.20 40.77 9, ,, carbon bisulphide . . 6.70* 9-75 7, ,, chloroform ........ 9.48 '7.86" I t is evident at once on a review of these figures, that a very large amount of nitrification was determined in each case by aspirating ordinary air through the garden soil, the amount being greatest in the second experiment, in which the soil contained more water, and was equally moistened throughout. The effect of carbolic acid vapour in preventing nitrification is in both cases imperfect, though much less so in the first experiment than in the second. Carbolic acid enters but slowly into vapour at ordinary temperatures, the vapour is also soluble in water, I n the second series of experiments the tem- perature of the laboratory was lower, the tube containing the soil was longer, and the soil uniformly wet.A smaller amount of carbolic vapour would thus enter the soil than in the first, experiment, and the quantity so entering would be retained in the upper portion of the moist soil, while nitrification went on unhindered below. In fact, when the tube was taken down at the end of the experiment, it was found that the cotton wool placed at the lower end of the tube was covered with mould, and no smell of carbolic acid could be perceived; while the wool at the upper end of the tube was quite clean, and the odour of carbolic acid manifest. In the first experiment, as already mentioned, only the upper half of the column of soil was really wet ; here nitrifi- cation was checked by the carbolic acid, while below nitrification could proceed but slowlyfrom lack of moisture.We learn from this experi- ment, that carbolic vapour cannot be succespfully employed as an antiseptic when any penetration of a moist porous body is demanded, whereas bisulphide of carbon fulfils such requirements excellently. In the case of the soils treated with chloroform and bisulphide of carbon, the cotton wool was found perfectly clean at both ends of' the column, and the odour of the vapour was perceivable at the bottom as well as at the top. The quantity of nitrate found in these soils is so little different from that present in the soil before aspiration, that we may The fact admits, I believe, of explanation.WARIKGTOX OK KITRIFICATIOS . 49 safely conclude that in these two cases no nitrate, or at least an insig- nificant amount, has been produced in the course of the experiment.The result of this branch of the investigation is, therefore, to prove that chloroform vapour will effectually prevent nitrification in a soil rich in vegetable matter, a fact directly confirmatory of Schloesing and Mun t z's result obtained with sewage. We also learn that bisul- phide of carbon is equally effective with chloroform, and that carbolic acid is probably effective to the extent in which it comes in contact with the soil. It thus appears that antiseptics, as a class, are inimical t o nitrification ; a result which is hard to explain if this process be one of simple oxidation only. We turn now to the experiments made at Rothamsted upon the second branch of the inquiry.The object aimed at here was to ascer- tain if nitrification could be induced in a weak solution of chloride of ammonium, by seeding with a small quantity of a nitrifying body. Four stoppered pint bottles of pale green glass were taken ; each was half filled with the dilute solution of chloride of ammonium used in Nesslerising (1 C.C. = *00005 gram ammonia), t o which a small quan- tity of acid phosphate of pot'assium was added. The object of the phosphate was to provide food for the organism to be presently intro- duced. Two of these bottles were placed in the window, and two in it dark cupboard. To one of the bottles in each series was added 1 C.C. of a mixture made by stirring together 1 gram of kitchen garden soil, and 5 C.C. of water.The soil used in this experiment, was rich in nitric acid ; it had, however, been kept for two or three years in an air- dry condition. On June 2nd all the bottles were tested, but no nitric acid found. Thinking that the slight acidity of the solution might be unfavourable to nitrification, a little freshly precipitated carbonate of calcium was at this date added to each bottle. On July 11th the solutions were again tested, but no nitric acid found. In the bottle seeded with garden soil, aiid exposed to daylight, a film of green growth had commenced on the bottom of the bottle. Thinking that the solutions might be too strong, or that the germs of the garden soil had lost their vitality by long keeping, the bottles were now in every case filled with water to the shoulder, and the two bottles already seeded received in each case 1 gram of surface soil, obtained three weeks previously from a " fairy ring." It need hardly be stated that the " fairy-rings" found in meadows contain in the circle of luxuriant grass a more or less vigorous growth of fungi.A sample of the soil beneath thering had in this case been taken to the depth of 12 inches ; this sample contained in its upper part a small quantity of mycelium. The whole 12 inches was air-dried, crushed, and sifted, and from the fine soil 1 gram taken for each of the seeded bottles. The experiment commenced on May 19th. VOL. XXXIII. E50 WARINGTON ON NITRIFICAT 10N. The bottles were next examined on October 11th. No nitric acid was found in either of the unseeded bottles.In the seeded bottle standing in daylight the green growth had increased. No nitric acid was found in this bottle ; the application of the Nessler test proved that ammonia was still abundant. With the seeded bottle preserved in darkness the case was wholly different ; here abundance of nitric acid was found,* amounting to 18.89 parts of nitrogen per million, while the Nessler test showed that ammonia had completely dis- appeared. The exact strength of the chloride of ammonium solution before nitrification was unfortunately not known, as a t the previous examination it had been diluted with about its own volume of water. If the dilution had been exactly to twice its volume, it would have contained 20.58 parts of nitrogen per million. We have here, then, a complete nitrification of an ammonium salt, brought about without any attempt to aikate the fluid, by the simple addition of a little soil from a fairy-ring, and of small quantities of phosphates, potash, and lime.The action also takes place exclusively in the dark, the perfectly similar mixture exposed to daylight exhibit- i n g no nitrification. It was evident that i f the solution just obtained was really full of nitrifying organisms, it ought to be capable of acting itself as seed, and of setting up nitrification in other solutions. On October 15th t h bottle contaicing the completely nitrified solution was shaken, an 1 C.C. of the contents added to each of the bottles-one in daylight, and the other in darkness-which had as yet remained unseeded. On November 6th the solutions were examined.That in the window showed 110 nitric acid ; that in the cupboard contained a trace, amount- ing t o about 0.72 parts of nitrogen per million. Thinking that the nitrifying organism might require organic carbon for its growth, which these solutions did not supply, 50 milligrams of acid tartrate of potassium were added to each bottle. That in daylight still showed no nitric acid. That in darkness now contained nitric acid equal to 2.04 of nitrogen per million. On December 3rd the nitric acid in the solution kept in darkness had risen to 11.89 of nitrogen per million ; but in the corresponding bottle in the window none was found. The experiment is still in pro- gress.? * The examination for nitric acid, and its quantitative determination, was made by means of indigo (Chem.N e w , xxxv, 46, 5’7). This method is admirably suited for such a purpose. 10 C.C. of the solution were sufficient for qualitative testing, and a few successive experiments with 10 C.C. sufficed for a quantitatke determination. -f On December 19th the nitric acid had incrzased to 20.06 parts of nitrogen per million ; a trace only of ammonia wa8 present. The bottle in daylight still contained no nitric acid. On November 19th the solutions were again examined.YctUNG ON THE GAS OF TJ3E GROTTA DEL CANE. 51 We have here, then, fresh proof that nitrification can be started in one solution from another, when without such seeding no nitrifica- tion would take place. The conclusions of S c hloesin g and Nun t z are thus again confirmed, with the addition of the important fact that darkness is apparently essential to the action of the nitrifying germs. That the nitrification observed in soils and waters is due to the growth of some mycoderm, must now be considered as highly probable. The process thus becomes perfectly analogous to the acetic fermenta- tion, in which alcohol is oxidised by the growth of Mycoderma aceti. The mycoderm producing nitrification has probably special characters, and might be sought for by a microscopical examination of the nitrified solutions already described. This part of the investigation must, how- ever, be left t o other observers. POSTSCRIPT.-January 4th, 1878.-In a paper read before the Royal Society, on December 6th, it is shown by Messrs. A. Downes and T. P. Blunt, that light is inimical t o the development of bacteria, a, few hours’ exposure to daylight being in many cases sufficient to destroy all the germs existing in an organic fluid, while similar solu- tions kept in darkness developed bacteria freely. The bearing of this obw-atinn on the fact already recorded as to the influence of light on nitrification is obvious. The influence of light on nitrification was not apparently quite unknown ; it is twice hinted at in Gmelin’s Chemistry (Cavertdish i3ociety’s Translation, iii, 68; vii, 92), but is not mentioned by most writers on the subject.

ISSN:0368-1645    

DOI:10.1039/CT8783300044

出版商:RSC    

年代:1878

数据来源: RSC

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YctUNG ON THE GAS OF TJ3E GROTTA DEL CANE. 51 \-.-The Gas of the Grotta del Cane. By T. GRAHAM YOUNG. SOME time ago my attention was drawn to the abstract of a, paper which appeared in the Chemical News (vol. xxxv, page Zl), giving the results of some analyses of the gas from the Grotta del Cane, near Naples, by E. A. Finot. The results of these analyses tend to show that the residual mixture of nitrogen and oxygen, after absorption of the carbonic acid, contains a larger proportion of oxygen than common air. Being desirous of accounting for the excess of oxygen, I took the opportunity of being in the neighbourhood of the grotto to procure samples of the gas in sealed tubes. The temperature of the air in the cave above the layer of gas is very different from that of the gas itself, as the following table shows :-52 YOUNG ON THE GAS OF THE GROTTA DEL ChXE.April 12, 1877. Temperature outside cave in sliade. .......... Temperature about 2 yards inside the cave, and 2 feet above layer of gas .............. About 2 inches above l a p of gas.. .......... About 2 iiiches below layer of gas ........... On floor of cave .......................... About 2 yards still further iuto cave on floor., April 27, 1877. 21" 25.5" 3 2" 36" 40" 21" - - 3 8" 40" April 28, 1877. 20" -- 25' 39" 40" - - The gaseous layer is completely saturated with aqueons vapour. The floor of the cave inclines downwards, and likewise the roof, so that, after proceeding about five yards, the gas is level with the roof, which prevents further ingress. On carrying a lighted torch into the cave, its smoke gradually falls, till it reaches the layer of gas, npon which it settles ; and on looking in, the surface of the gaseous layer is seen, resembling that of water, and appears covered with beautiful undulations.On holding the head below the level of the gas, holding the breath, and keeping the eyes open, an intolerable prickling sensation is pro- duced upon the eyes by the carbonic acid. A dog brought into the cave, as is the custom there, appears, as it were, to drink the gas, lapping with ih tongue. Then its eyes begin to dilate to an unnatural size, and its lapping becomes more spas- modic; beyond this it does not seem to suffer. While in the cave, also, the dog was able to stand, but when carried and set on its feet outside in the fresh air, it fell, and lay struggling zs if in paroxysms of suffocation, but recovered in two or three minutes. I was told, how- ever, that the animal gets into such a nervous state with the prospects of its frequent ordeals, that it has to be killed in three months. On nnalysing the gas, I get results contrary to those obtained by M. F i n o t , as regards the excess of oxygen. .I find no excess of oxygen, but an amount nearly corresponding with that contained in common air ; also a much larger proportion of carbonic acid than that found by 31. F i n o t . The following is a comparison of the results :- The carbonic acid varies from 61.5 to 71 per cent. Accoiding to Yo u n g. According to F i n o t . CO, ...... 61.50 71'00 Residual air 25.38 25.69 Residual air 0 ........ 7.80 5.80 2025 20.00 18-46 20.13 24-74 27.10 N ........ 30.70 23.80 79.75 80.00 56.16 5418 75-26 72.90 100~00 100~00 100*00 100~00 100~00 100~00 100-00 100~00

ISSN:0368-1645    

DOI:10.1039/CT8783300051

出版商:RSC    

年代:1878

数据来源: RSC

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53 VI.-On some Derivatives of Allylacetone. By J. I(. CROW, B.Sc., Dalton Scholar in the Chemical Laboratory of Owens College. p H , ( CB5) I I ALLYLACETONE, CH,-CO-CHp(C3H5), or 4 CO I t , was first pre- jCH, pHred by Zeidler, who obtained it by decomposing allylacetoacetic ether with potash, in the usual manner in which all acetoacetic ethers are decomposed. As this ketone contains a double linking, it ap- peared interesting to try the action of nascent hydrogen upon it, for tbe purpose of determining whether the addition of hydrogen takes place in the position of the doublelinking in the chain, orwhether this reaction is the same as in other cases, an alcohol being formed. The acetone wits prepared according to Zeidler's method, and diluted with its own volume of ether, and the whole transferred to a, flask surrounded by cold water, and containing also a quantity of water abdut twice as large in volume as that of the ethereal solution of the acetone.The flask was connected with a reversed condenser, and sodium cut in thin slices dropped gradually in, until about half as much more of the metal had dissolved as was theoretically needed for the reaction. The ethereal solution was then separated by a funnel from the watery part, and dried with potassium carbonate. After the ether was distilled off, the greater part of the remaining liquid came over from 135" to 140'. This portion, after repeated fractional dis- tillations, yielded the pure substance boiling at 138-139". It was dried over anhydrous copper sulphate, a :J analysed with the following results : - (a,) 0.1064 gram substance yielded on combustion 0.2795 gram C02, and 0.1148 gram H20.H20. formula C6H120. ( b . ) 0.1152 gram substance gave These numbers point to the Calculated for this formula. c ...... 72-00 H ...... 12-00 0 ...... 16-00 100~00 0.3025 gram C02 and 0.1255 gram conclusion that the liquid has the Pound. a. 6. '71.64 71-61 11.99 12.10 The liquid appears t o be an alcohol of the ally1 series, as it possesses VOL. XXXIII. F54 CROW ON SOME DERIVATIVES O F ALLYLACETONE. the power of combining very violently with bromine, and yields an acetate when heated with acetic anhydride. I t has a sweet, but a t the same time rather pungent odour, somewhat resembling that of allyl alcohol. It is soluble to a slight extent only in water, but freely in alcohol and ether.It appears to be a secondary alcohol, and a homo- logue of allyl alcohol, and has the graphic formula- The specific gravity of the a,lcohol was found to be 0.842 at 16*2", compared with water at 17.5". Its acetate was obtained by gently heating the alcohol in a flask connected wit,h a reversed condenser, with acetic anhydride in excess. The product was then washed with a dilute solution of potassium carbonate, dried, and distilled ; the acetate was obtained, boiling between 147-149", as a colourless liquid of a pleasant refreshing odour. (a.) 0.1991 gram of the substance gave 0.491 gram GO2 and 0,1766 gram HzO. ( b . ) 0.1570 gram of substance gave 0.3877 gram CO, and 0.1397 gram H,O. It was analysed, wit'h the following results :- Calculated for the Found.formula C,H,,0C2H30. a. 6 . c ...... 67.61 67-26 67.35 H ...... 9.86 9.86 9.88 0 ...... 22.53 - - 100-00 As before remarked, the alcohol combines very readily with bro- mine. In order to obtain the bromine compound, a mixture of the alcohol with chloroform as placed in a flask surrounded by cold water, and connected with a condenser. Bromine was then gradually added from a dropping funnel until it effected a permanent coloration of the liquid; the chloroform was then distilled off under reduced pressure, to prevent the bromo-compound from decomposing ; and the resulting product, which was thick and slightly brown, was dried over sulphuric acid and lime. The mass does not crystallise, and cannot be distilled ; it has a most pungent odour, its vapoinr atta,cking the nostrils.The following analyses show that it has the formula, CgHl2Br20, and is an addition-product of one molecule of bromine to one of alcohol. 0.1604 gram of substance gave 0.2196 gram AgBr and 0.0088 gram Ag.CARNELLEY AND O'SHEA ON TETRARROMIDE OF TIN. 55 0.1038 gram of substance gave 0.1487 gram AgBr and 0-0007 (a) 0.2305 gram of substance yielded on combustion 0.2314 gram ( b . ) 0.410 gram of substance e v e 0.4105 gram CO, and 0.1685 formula C6HI,Br20H. a. 6 . gram Ag. CO, and 0.0947 gram H,O. gram HzO. Calculated from Found. C ...... 27.69 27.38 27.31 IT ...... 4-61 4.56 4.57 Br ...... 61.54 61.97 61.45 0 ...... 6-16 - 100~00 This dibromide was hssted with silver acetate in sealed tubes, to obtain a diacetin, but without success, as the body decomposed ton rapidly on heating. An attempt to substitute hydroxyl for bromine by heating the substance in sealed tubes with water also failed. When this body is boiled with K2C03 solution, a clear oily liquid distils over with the steam; it has a smell very similar to that of the dibromide, but contains an atom of bromine less. On analysis 0.2804 gram of substance gave 0.1484 gram AgBr and 0.0064 gram Ag, which corresponds with 41.03 per cent. of bromine, the percentage required by the empirical formula C6H,,Br(OH)2 being 40.40.

ISSN:0368-1645    

DOI:10.1039/CT8783300053

出版商:RSC    

年代:1878

数据来源: RSC

摘要:

CARNELLEY AND O'SHEA ON TETRARROMIDE OF TIN. 55 VII.--Note ow, Tetrabromide of Tin. By THOB~AS CARNELLET, D.Sc., Assistant Lecturer on Chemistry in The Owens College, and L. T. O'SHEA. IT having been necessary to prepare a quantity of tetrabromide of tin for a research on the melting points of metallic salts, in which one of us is engaged, it is thought that a short account of this body will not be without interest, especially as, with two exceptions, there has not been published any previous account concerning it. One of these occurs in Balsbrd's well-known research on bromine (Ann. Chim. Phys., xxxii, 337). His account of it, however, is very incomplete, neither melting point, boiling point, nor analytical numbers, &c., being given, for he merely says that by heating tin in bromine a white crys- talline substance is obtained, which is easily fusible and volatile.The IF 256 CARNELLEY AND O'SHEA ON TETRABROMIDE OF TIN. second reference is that of Bcedeker, in his Beziehungen Zwischelz Dichte u n d Zusnmm ensetzung bei f e s t e n zcnd liguiden S t o f e n (Leipsig, 1860), where he simply gives the specific gravity and melting point. In order to prepare stannic bromide, a piece of combustion-tubing a was bent thus, \AA/, and a quantity of metallic tin placed in b c d the bend c, and heated to the melting point by a lamp ; bromine was then gradually dropped into the bend h, by means of a tap-funnel attached to n. On the bromine passing over the tin, the latter burnt with a very faint violet flame, while the tin bromide formed condensed in the bend d.The burning of the metal can be kept up continually by properly regulating the flow of bromine, and would serve as a very good lecture experiment. The distillate obtained was freed from excess of bromine by redistilling repeatedly, till the liquid which passed over was quite colourless. On slightly cooling, or even without, it solidified to a mass of beautiful colourless crystals of pure stannic bromide. Tetrabromide of tin melts at 30' C., and boils without decomposi- tion at 201" (uncorrected) ; a large quantity distilled entirely at this temperatme, without the thermometer showinq the slightest altera- tion. On exposure to the air a t the ordinatry temperature it does not fume, and is only very slowly decomposed, but on gently warming i t gives off white vapours of oxide of tin.It dissolves in cold water without decomposition, but after standing several hours it is gradually decomposed, hydrate of tin being precipitated ; boiling, however, hastens the decomposition, as does also the addition of nitric acid, which in the hot solution precipitates the whole of the tin almost im- mediately, with a slight evolution of bromine. The specific gravity, according to B cedeker (vide supru) of a specimen of tin tetrabromide melting a t 39" was 3.322. Adysis.-The tin was determined by evaporating to dryness with nitric acid, igniting, and weighing the oxide of tin produced. For the estimation of the bromine a weighed portion was boiled with dilute nitric acid in a, flask provided with a tube dipping under am- monia, so as to retain any hydrobromic acid that might be given off.On cooling, the hydrate of tin was filtered off, and the filtrate, to which the ammoniacal solution had been added, mixed with sulphur- ous acid to reduce any free bromine, and the bromine determined in the ordinary way with silver nitrate. (1.) -9244 gram of substance gave -3188 gram SnO, = 27-13 per cent. of Sn. (2.) -7572 gram of substance ga,ve -2610 gram SnO, = 27-11 per cent. Sn.MILLS ON POTABLE WATERS. 57 (3.) ,9523 gram of substance gave 1.6210 gram AgBr and 9044 gram Ag = 72.78 per cent. Br. Found. Calculated. (1) (2) (3) Mean. Sn = 26-94 27.13 27.11 - 27.12 Br = 73-06 - L 72.78 72-78 100*00 99-90 -- Determiiaation of Vuyour Density.-& the compound distilled with- out the least sign of decomposition, its vapour density was determined by Dumas' method, with the following results :- I. 11. Weight of bulb + air .... = 2.4697 4.985 Temperature of air.. ...... = 16" C. 13' c. Weight of bulb + vapour . = 2.6415 5.157 Capacity of bulb.. ........ = 17.1 C.C. 19.5 Barometer .............. = 741 731 3 , oil bath.. .. = 228" 260" The small bubble of air was measured and allowed for in the second experiment, but not in the first. Found. Calculated f o r Mean. SnBr4. (2) 229 219 Vapour density.. = 230.6 227 (1)

ISSN:0368-1645    

DOI:10.1039/CT8783300055

出版商:RSC    

年代:1878

数据来源: RSC

摘要:

MILLS ON POTABLE WATERS. 57 VIIL-On Potable Waters. By EDMUND J. MILLS, D.Sc., F.R.S., “Young” Professor of Technical Chemistry in Anderson’s College, Glasgow. A COMPLETELY satisfactory method of examining potable waters in- volves the solution of a considerable number of important chemical problems, and must long continue to require the thought and labour of mmy experimenters. These problems have occupied much of my attention during the last two years, in consequence of my undertak- ing to report every month on the quality of the water which is derived from Loch Katrine and supplied to Glasgow. With respect to the organic constituents of waters two principal pro- cesses have been proposed. In the first of these (Wanklyn, Chap- man, and Smith, Chem. SOC. J., xx, 445) the unconcentrated mixture58 MILLS ON POTABLE WATERS.of organic impurities (for a mixture we must assume it to be) is boiled with alkaline permanganate, and the ammonia thereby evolved is regarded as a measure of the impurity present. This process is simple and elegant, and will probably form a leading feature in our future methods, for it admits of graduation, and may serve to dis- tinguish one impurity from anc, ther. Unfortunately, however, it deals with bodies a t present unknown, in a manner of which we are igno- rant ; its results therefore admit of no precise interpretation, and (as has been justly remarked) (Suttmb’s VuluirLetrlc hzaZysi.s, 3rd edit,, p. 345) can aid only to distinguish a bad water from a good one. Accord- ing to the second process (Chem. SOC.J., xxi, 77) the water is concen- trated, usually after the addition of a reducing agent, and the residue submitted to combustion in a vacuum with cupric oxide in presence of copper. Now substantially satisfactory evidence has been adduced that there is no material loss during the evaporation. It is easy therefore, by a suitable mode of collecting and analysing the gases of combustion, to deteymine exactly the total amounts of carbon and nitrogen extractable from the organic mixture present in the water, This process admits of no graduation, and can never distinguish between one organic impurity and another associated therewith. Happily, however, it fur- nishes two familiar fundamental data in a manner well understood by chemists, and its results are precisely interpretable. Hence it forms the philosophical starting-point in an endeavour to solve the problems of water-analysis.It will be referred to exclusively in the following pages. I. EBRORS AND BL,~NKs. Probable 3:rror qf the Results.-In their first paper on this subject (Zoc. cit., pp. 93, 94), Frankland and Armstrong gave various determinations of organic carbon and nitrogen in small quantities of sugar, urea, and hydric hippurate, &c. Assuming 100 as the weight of organic carbon or nitrogen taken, I find their analyses lead to the following figures :- No. of deter- Prob. error per cent. of Substance. minations. one determination. Organic carbon .......... 10 5.96 nitrogen. ......... 6 7.35 I entertain little doubt, from my own experience, that this nncer- tainty must, have rapidly diminished, though it would naturally enough occur in the case of anyone inventing this process and then using it for the first time.In 1874 (Sixth Report, of the Rivers’ Corn- niissicn, p. 505), F r a n k l a n d obtained further results with some quinine solutions. Three of these deal with too minute quantities toMILLS ON POTABLE WATERS. 59 be fairly taken into account for probable error; the remaining five furnish the following results :- No. of deter- Prob. error per cent. of Substance. minations. one determination. Organic carbon .......... 5 2.84 ,, nitrogen.. ........ 5 2-56 These carbon determinations are all rather too high, the nitrogen determinations all rather too low : they are evidently affected by con- stant error, for which I cannot give any account.Hence the probable errors in the quinine experiments are greater than those of the process itself. Judging of the results simply by their own mean, the probable error per cent. per single determination reduces t o 1.36 for the carbon and 1.44 for the nitrogen, An intermediate number, say 2 per cent., would probably be a very fair estimate. In order to appreciate the significance of this result, it will be neces- sary t o compare i t with others, derived from common analytical figures. Gmelin gives three analyses of cane-sugar (Watts’s translation, xv, 247), four of urea (ibid., vii, 363), and five of hydric hippurate (ibid., xii, 72). Considering the theoretical amount of‘ carbon or nitrogen in each case as 100, we arrive at the following conclusion:- No.of deter- Prob. error per cent. of Substance. minations. one determination. Organic carbon ...... 12 ,, nitrogen. ..... 8 The carbon determinations are all too low ; the nitrogen determina- tions are apparently free from constant error. Gmelin also gives (ibid,, xvii, 269) five analyses of quinine, of which only the carbon determinatioiis are available for calculating probable error ; the necessary data for the sulphate are also lacking. The probable error per cent. of single detwmination is, in this case, 0.55. According, then, to Fr a n kl and’s quinine experiments, the com- bustion method of water analysis may be taken as now giving a probable error of about 2 per cent. per determination ; and OUY know- ledge of the amount of carbon and nitrogen extractable from cannose and urea with hydric hippurate is, as a crholle, no more accurate t.han this, according to the analyses cited by Gmelin.On the other hand, the analyses of quinine given by Gmelin are more than thirteen timea as valuable as those which Frankland adduces. This disparity, how- ever, is considerably diminished by the fact that quinine is easily brought into a condition for exact weighing, but that its sulphate is not. From several hundred analyses made in my laboratory, in con-60 MILLS ON POTABLE WATEBS. nection with which systematic computations of the above character have been made, ‘it appears that, taking grnvimetric and volumetric methods together, the probable error per cent. of a single determina- tion slightly exceeds, as a rule, 1 per cent.Large numbers of deter- minations, however, brought forward by Gmelin as evidence of the composition of various bodies, have a considerably greater probable error than even 2 per cent. It appears, then, that the accuracy of Frankland and Armstrong’s method compares favourably with that of the ordinary combustion process, and may be safely used as a basis of inference. Blank Experimerzts.-The water is, as a rule, evaporated with a con- stant quantity of hydrosodic sulphite dissolved in a large excess of aqueous hydric sulphite, a trace of ferric salt having been first added if necessary. It is convenient to have the two sulphurous solutions then mixed together. These lead to a constant excess a in the carbon and nitrogen, which excess can be determined by n special combustion of their residue.We may term a the chemical error in the blank. When I& cubic centimeters of purified water are evapo- rated with the sulphites, in order to imitate as closely as possible the proceeding of an analysis, a further excess of carbon and nitrogen is introduced. This is evidently proportional to the number of cubic centimeters of water present, and may be termed the “ volume error v.” Hence the general form of correction for a blank experiment is a + mu, v being the amount of error introduced per cubic centimeter of water. In order to obtain purified water, it is convenient to heat distilled water to near boiling for about twenty-four hours with alkaline per- manganate, and then, after excessive acidification with hydric sul- phate, to heat it again for an equal time; in this way the organic matter has a double solicitation t o decompose. The acid liquid is finally distilled.A litre of water thus purified from the Loch Katrine snpply adds nothing to the chemical error during an evaporation lasting twenty-four hours. It seems probable that the volume error is in most cases very small. I have found the following errors in blank experiments :- (1) .000048 carbon, -000111 nitrogen (a). (2) .000047 carbon, ,000114 nitrogen ( a + nv). (3) -000090 carbon, .000085 nitrogen ( a : new chemicals). InJlzcence of Size of Evaporating Dish.-In the original method, as still carried out by its inventors, the water was evaporated in a round glass dish, having a diameter at the orifice of 4 inches (10.2 centims.).It was proposed by Bischof (Proc. Phil. SOC. Glasgow, ix, 21) to in- crease the diameter of the dish to 6 inches (15 centims.), in order to accelerate the rate of evaporation. The dishes employed in my labo-MILLS ON POTABLE WATERS, 61 ratory have the minimum available diameter, namely, 3 inches (7.6 centims.). In the following table their performkcc* is contrasted with that of a 6-inch dish, operating at the same time, in the same apartment, and on the same waters :- Org. carbon. Org. nitrogen. inch dish . . . . *22 7 -01 5 -203 -027 -235 -020 9 , ,7 * * " -214 -018 c ..., {; > 9 7 ) * . - * *444 *039 9 9 7 9 ' + . * *438 -041 ,, 2, ' ' 0 ' *375 -050 Means { 9, >> -323 -024 9 ) , j - - - -307 -034 Differences -.. . . - so16 + -010 G 9 , 1 7 A . . . . Be,.. { 7 9 7 9 . * " D a m . , {g 9 ) 9 ) * * * . -387 -022 It appears, then, that there is a loss of carbon and a gain of nitrogen by using a &inch dish. Owing to the irregular manner in which the organic impurity of the residue is protected by the drying sulphite, the effect also is naturally irregular. A special experiment was made with two pairs of residues in order to ascertain the source of the increase of nitrogen. The residues were added to water, free from ammonia, and distilled with sodic carbonate, as in ordinary ammonia determinations.. The results were- Ammonia. The greater part of the ammonia here found was produced from the peaty matter present in the water by the action of the sulphites with which it was evaporated.The difference between the two results must represent ammonia absorbed from the atmosphere. Its mean amount is -011, corresponding to -009 " organic nitrogen," a number very nearly identical with the mean excess (-010) actually found above. It is evident, then, that the use of a, 6-inch dish leads to un- certain and inaccurate results. The areas of a 3, 4, and 6-inch dish are to one another as the numbers 1, 1.8, and 4 ; and the error pro- bably increases rapidly with the area. The problem in fact is, how to evaporate a water in the narrowest possible dish in the shortest possible time. * I am indebted to Mr. Macnab for the careful carrying out of these experi- ments.62 MILLS ON POTABLE WATERS. 11. APPARATUS. Evaporator.-The evaporator I employ is represented in the accom- panying figure at about one-fourth its natural size. It is constructed of copper, and consists of a cylindrical water-bath, a, surmounted bya hemispherical cover, b, which stands in a water-joint.In the platform of the bath is placed a 3-inch glass dish, c ; this is held down steam tight against an india-rubber band by the three clamps d. EVAPORATOR. (SCALE t.)MILLS ON POTABLE WATERS. 63 There is an aperature, e, in the cover, through which and a simple stuffing arrangement enters the constant feed of water to be analysed. The cover is kept cool by a continuous supply of cold water through a number of jets in the circular pipe,J: This water becomes warm by the time it has flowed down to the water-joint; which, in its turn, delivers it to a discharge pipe, g, and thence to the constant feed, h, for the water-bath. The remaining orifices of the bath are three.Of these, 7~ is the steam exit ; Z is the entry, m is the exit for a currelit of aspirated air. This air, having been first passed over purified pumice moistened with somewhat dilute hydric sulphate, traverses two coils at the base of the interior of the bath ; thus heated, it ascends a chimney whose summit is a little above the rim of the dish on the right hand side of the figure. It then strikes to the left across the dish; and descending a similar pipe whose orifice is level with the dish’s rim, passes through 71% to an aspirator, which is fixed to the tap. The air is thus unfavourably disposed for giving up anything absorbable by the evaporating liquid ; and, in fact, has no chemical in- fluence on the process.The cooled cover condenses 6 cub. centims. an hour. With a current of air amounting to 360 litres an hour, the total evaporation is rather more than 45 cub. centims. an hour ; with a current of 1062 litres, the total evaporation wonld be 100 cub. cenhims. an hour. Convectiordess Feed.-in Bischof’s apparatus (Zoc. cit., p. 22), the supply passes from the reservoir to the evaporating dish through a straight tube. There is, consequently, convection through this tube to the reservoir ; and that is necessarily attended with much loss of heat. It is, moreover, objectionable to heat the delicate organic matter of waters (and especially of peaty waters) any longer thiin is strictly necessary.I have theref ore contrived a convectionless feed. Its principle is very simple. The delivery tube, instead of being straight, is bent in two places ; the consequence of which is that, at the end of each regurgitation, two long bubbles of air, pzn, always re- main entrapped. Between these and the tube, the layer of water is so thin, that convection to their left becomes impossible. The heat that would otherwise escape is thus locked up within the evaporator, and the water to be analysed is not heated until the last moment. So effectually is convection arrested, that, with the first made (and rather imperfect) feed of this kind, and a rate of evaporation considerably exceeding 100 cub. centirns. an hour, it took four hours for some rather strong hydrochloric acid in the dish t o redden blue litmus in the reservoir.The evaporator also works as a whole so steadily that, if a scale be applied to a reservoir of suitable shape, the diurnal changes of the barometer can be accurately integrated by its means.64 MILLS ON POTABLE WATERS. 111. NATURAL CONSTANTS. Living plants and animals consist of a large number of compounds which undergo transformation after death; and it is an accepted be- lief, for which much evidence can be adduced, that in such transforma- tion those compounds are greatly reduced both in number and com- plexity. The natural oxidising action whereby this end is atfained is known to be very efficient when conducted in the soil, and mora par- ticularly through great depths of soil. Hence we may expect that in water gathered considerably below the eart’h’s burface, or in some equivalent situation, there will be very few organic compounds, or even one alone; in short, that we shall be dealing with a case of what Chapman designated “ limited oxidation.” The ratio of organic carbon to organic nitrogen must, under this condition, approximate to one or more of the ordinary chemical ratios, C, : C, : and ; if constant, it may be termed a ‘‘ natural constant in potable waters.’’ The Sixth Report of the Rivers’ Commission (London, 1874), con- tains an enormous number of analyses of waters taken from every variety of geological formation, and must ever remain one of the principal monuments of scientific industry in connection with natnral waters.The result is, that the great majority of the returns show no satisfactory evidence of chemical constants :-the polluted or turbid waters for an obvious reason; the clay effluents, because oxidation in clay is too languid; springs from certain coal measures, because they are in fact polluted by coal, &c.; in some cases, the figures are promising but too few. There are, however, three groups of data that do lead up to the object of which I was in qnest, as might indeed have been ex- pected of them; these refer to unpolluted (“ clear” or “ slightly turbid ”) waters from deep wells, and from land unmanured and un- cropped,-to which may be added the results of an investigation into the effect of filtration through spongy iron. (a) I. Deep Wells in the Chalk. (Report, pp. 99-lOl.)-The follow- ing table contains the organic analyses of fifty waters :- I have carefully examined the whole of that Report.MILLS ON POTABLE WATERS.65 ~ ~~ NO. I. 11. IV. V. VI. VII. VITI. IX. X. XI. XII. XIII. XIV. XV. XVI. XVII. XIX. xx. XXI. XXTI. XXIII. XXIV. xxv. XXVI. XXVII. XXVITI. XXIX. xxx. XXXI. XXXII. XXXIII. xx XIV. xxxv. XXXVI. XXXVII. xx XVIII. XXXIX. XL. XLT. XLII. XLIII. XLIV. XLV. XLVI. XLVII. XLVIII. XLIX. L. m. xvrrr. Locality. Brighton. Bury St. Edmunds. Canterbury. Deal. 3 ) 7 7 9 7 Dorchester. Dover. Dunbridge. Great Grimsby. Hull. Ipswich. Norm-ich. Ramsgate. Sudbury (Suff.). Taverham. Thames Basin. 7 7 l 3 7 7 >> 7) 9 ) 3 1 >) 3 ) IJ 77 >) Y ) 17 97 7) Y 3 7 7 7 7 9 9 >) >) 9 , 7 ) 3 7 Yl )l Y ) >> J J 77 Winchester.TABLE I. - hganic k b o n . -048 -055 '089 '012 '032 -050 056 -040 '034 -028 '044 *025 '032 -093 '064 '050 '045 '059 -076 '041 '029 '036 '041 '034 -027 -071 '020 '028 '049 '031 *040 -048 -044 '056 '030 -058 -064 '01 9 -026 '081 '051 '033 '131 '05 2 -036 '030 -048 '074 '052 '042 hganic itrogen -009 ,011 '020 '012 '013 -007 '024 '010 '008 -005 '005 '007 '012 -016 -017 '027 '013 '012 -007 '024 '009 '016 .007 '030 '026 '018 '006 '026 '006 '009 '006 '005 -007 -005 '007 '01 1 '009 '018 *0l7 '012 -008 '01 1 -017 '014 '010 '024 '010 -015 '011 '018 Remarks. Clear and palatable. 79 )) > 7 9 ) 7 ) 3lightly turbid ; palatable. Clear and palatable. 9 ) Y7 Y 7 3 ) >> 9 ) )> 3 3 9 ) Slight,ly turbid ; palatable. Clear and palatable. >7 3 ) ) l 77 Slightly turbid ; palatable.Clear and palatable. 77 97 Y ? 9 ) > Y >7 >7 97 9 ) 9 ) J? Y > )) > I l I Y ) >) Slightly t,urbid ; palatable. Clear and palatable. Y ) 37 11 >l 9 , 2 9 ~- Assuming thaL all the determinations have the same '' weight," the method of least squares (which will be used throughout) leads to a66 MILLS ON POTABLE WATERS. org. carbon ratio, = 3.089, with a probable error of *26 (or 8.3 per org. nitrogen cent.) for the mean. When it is remembered that, in nineteen of the nitrogen determinat,ions, the amount actually obtained is in the third place of decimals of a part per hundred thousand, such a probable error appears exceedingly moderate. If we were to omit determina- org. carbon tion XLV as exceptional or accidental, we should find org.nitrogen = 2.993, with a probable error of -2.3 (or 7.6 per cent.) : but, in the absence of direct evidence, this proceeding would be illegitimate. (Report, p. 220).- There are fourteen experiments, the original water having been the Chelsea supply. 11. Efect of Filtration through Spongy Iron. TABLE 11. N O . I. 11. 111. LV. v. VI. VII. VIII. IX. X. XI. XII. XIII. XIV. - Organic Carbon. *025 -046 -063 -070 -060 -120 -097 -060 -113 -077 -048 -073 '077 -089 Organic Nitrogen. -004 -015 *036 -047 -008 .031 -023 *015 '027 -021 -009 -013 -012 -020 Remarks. Clear. TU';bid. Clear. Sli'ihtly turbid. Clear. Contained suspended particles. Slightly turbid. 1 , 9 9 9 , Clear. 9 9 Hence we obtain the ratio 3.015, with a probable error of -40 (or 13.4 per cent.). This quotient is substantially the same as the pre- ceding." Weighting" the two in accordance with their probable errors, we arrive as a mean result at 3.067, with a probable error of -22 (or 7.2 per cent.). (p) I. Deep Wells in the Devonian, Millstone Grit, and Coal Men- sures. (Report, pp. 89--92.)-There are eleren available determina- tions. TABLE 111. Unpolluted Deep Wells (Devonian). Locality. 71 Mansfield. -021 111. Brbroath. Remarks. Slightly turbid ; palatable. Clear and palatable. 9 ,MILLS ON POTABLE WATERS. 67 Organic Nitrogen. ---- -045 1 '049 1 -033 -037 *075 -054 '042 :;;; i No. No. I Locality. Locality Organic Carbon. Organic Organic Carbon, Nitrogen. l l I Remarks. Uwpolluted Deep Wells (Millstone &it). TV. I Glossop. I *092 1 -020 Clear and palatabIe. 1.Unpolluted Deep Wells (Magnesian Limestone). Pon t efract . Mansfield Woodhouse. VII. Sunderland. 97 Uiipolluted Deep Wells (Coal Measures). VIII. Accrington. -045 *Ol7 Clear and palatable. Bedlington. Blnckburn. 1 '??; 1 ':z 1 7) f f XI. Castle ford. -133 -045 Slightly turbid ; palatable. The ratio is 2.534, with a probable error of -26 (or 10.1 per cent.). 11. Dminage f r o i n Land Unmanzcred nnd Uncropped. (Report, p. 62.) -The land referred to had bee! manured with guano and superphos- phate in 1868. The determinations, of which nine only can be selected, extend from 1872 to 1874. TABLE IV. I-----I- -- I. 11. 111. IT. v. VT. VII. VIII. IX. Rothametcsd. -108 '147 '127 *114 -096 *098 ,174 *I17 -098 Remarks. Clear. Sliiht.17 turbid. f ? )) 9 ) 9 9 Clear.$ 9 The ratio is 2-tj20, with a probable error of -08 (or 3.1 per cent.). Combining this with that found in PI, the result is 2.521, with a pro- bable error of -07 (or 2.9 per cent.). pi. Deep Wells in the New Red Sandstone. (Report, pp. 93-94,)- There are twenty-four available determinations.68 NO. -- I. 11. 111. IT. v. VI. VIT. VIII. IX. X. XI. XII. XIII. XIV. XV. XVI. XVII. XVIII. XIX. xx. XXI. XXII. XXIII. XXIV. MILLS ON POTABLE WATERS. Locality. --- Birkenhead. Birmingham. 9 7 > I )I 9) 7 9 Kidderminster. Liverpool. WaCasry. Wolverhampton. > Y >) 9 9 Worksop. >) TABLE V. Organic Carbon. '041 -047 '052 '034 -031 '009 '037 '01 5 -066 .020 '076 '018 '050 '027 '039 '017 '026 '030 '041 '068 '032 -064 -026 -048 Organic Nitrogen.'038 '015 '016 -006 *007 '004 -01 2 '004 '024 '020 '033 '013 '020 -007 '035 '013 '019 '008 -014 '010 '011 '023 *006 -007 Remarks. __---- Clear and palatable. > 9 Slightly t,urbid ; palatable. Clear and palatable. 9 9 2) Slightly turbid ; palatable. Clear and palatable. Slightly turbid ; palatable. Clear and palatable. 9 ) 7 ) 7 9 >, Slightly turbid ; palatabIe. Clear and palatable. 7 9 >> > ? > 9 Slightly turbid ; palatable. Clear and palatable. 9 ) 9 9 ~~ ~ ~~~~~~ The ratio is 2-056, with a probable error of -27 (or 13.2 per cent.). The following table contains a summary of the whole of the preced- ing evidence as to the natural constants of waters. No. of TABLE VI. Prob. error Calculat.ed. of constant* C I z + N, = 3.429 64 a = 3.067 -22 20 p = 2.521 -0 7 + Na = 2.571 3,000 observations.Constant. 24 n/ = 2.056 -27 C,, + Ng = 2.057 The value of these results may be appreciated as follows. It has already been shown (p. 59) that, in certain instances of the ordinary combustion process, the probable error of a single determination of organic carbon is 1.21, of organic nitrogen, 2.72, per cent. If we suppose it required to ascertain the ratio between the two quantities, the quotient will be 1.00, with a probable error of 2.98 per cent. The total evidence for the constant 2.571 is thus rather better than a single determination of the ratio of carbon to nitrogen by the com- bustion process, as above referred to.* When, however, we find this * The following instance of the state of our knowlecige in soine C ' X Y ~ S of quaiititatire According to the c+dcnw giren b2- Gmelin anal>sis inay be adduced at this point.MILLS ON POTABLE WATERS.69 constant to be evidently one of a series of three others, its probability and theirs become very considerably increased. Hence, while trust- ing that many additional analyses may be brought t o bear on dimi- nishing the probable error of these constants, 1 cannot but think that their establishment has been fairly secured. Can we find a porous oxidiser, such that it will reduce the whole of the nitro- genous organic matter to a body or bodies having one of the ratios a, 6, or 7 ? If so, there can be little doubt that the alkaline perman- ganate process would then give as definite and certain a result as it does now wihh an alkaloid.* The ratio of nitrogen to oxygen in the air is, with small variations, a constant quantity.Its action upon living beings has to a great extent been elucidated, and appears t o involve a closed cycle of ex- changes between animals and plants; but without this cycle is the region of decomposing beings, its action whereon has hitherto con- tinued unexplained. Now, upon this latter subject, the analyses of the organic matter in waters throw considerable light. The principal constituents of animals, the proteins, have an nC, formula ; the priiicipal constituents of plants, the celluloids, have also an 1aC6 formula. By oxi- dation in presence of water and the porous earth, all these are, to some extent, consumed ; but, in the long run, as we have seen, the ?%c6 formula is preserved, in a set of bodies whose number does not necessarily ex- ceed three. These bodies have a constant composition and, in all proba- bility, a practically constant weight ; their continued production, more- over, involves a definite rate of organic change. The mass of the oxygen concerned in their formation must, therefore, be definite. We are, as yet, ignorant of the manner in which nitrogen takes part in the natural che- mistry of the earth’s surface ; but the total amount in the air is finite, and if it be the complement or related to the complement of the amount existing organically in potable waters, its mass must also be definite. We can, therefore, conceive an origin for the constancy of the compo- sition of the air; and seeing that the carbon-nitrogen “ratios” in waters are three at least, we can understand how it has chanced that the oxygen-nitrogen ratio in air is not a definite chemical ratio. A material alteration in the mass of atmospheric oxygen, or of the porous soil, or of the soil’s temperatiwe, would involve a vital che- mistry based on some other modulus than nC6, and a total change in the aspect of terrestrial existence. (iii, 321) for the percentage of water in so common a substance as potash-alum, the mean result is 2.51 per cent. above theory; the probable error of this mean is 1-80 per cent., and of a single determination, 3.60 per cent. These errors are consider- ably enhanced if calculated from the theoretical, instead of the mean experimental number. These results lead to some not unimportant suggestions. * I have already performed aome experiments in this direction. VOL. XXXIII. G

ISSN:0368-1645    

DOI:10.1039/CT8783300057

出版商:RSC    

年代:1878

数据来源: RSC

摘要:

70 IX.-Ort a Fourtk New Method for Estimating Bismuth Vohmetrically . By M. M. PATTISOK MUTR, F.R.S.E., Praelector in Chemidry, Caius College, Cambridge. 1. WHEN a saturated solution of oxalic acid is added to a solution of bismuth in nitric acid, a white crystalline precipitate slowly forms this precipitate, according to Souchay and Lenssen (Ann. Chern. Pharm., cv, 245) is normal bismuth oxalate, Bi,3C20a + ag. This reaction appeared to afford a mean whereby bismuth might be accurately and readily determined. I f the oxalate were produced under conditions which might be realised without too much difficulty, and if this oxalate were of constant composition, every requisite of an accurate volumetric process would be fulfilled. 2. I carried out several experiments, using a solution of bismuth in a small excess of nitric acid ; precipitating with a saturated solution of oxalic acid ; washing by decantation ; dissolving in warm dilute sulphuric or hydrochloric acid ; titrating with standard permanganate ; and calculating on the assumption that three molecules of oxalic acid found corresponded with two atoms of bismuth thrown down as oxa- late.The results were: not, however, very satisfactory, the amounts of bismuth found being generally too low, and the numbers exhibiting considerable discrepancies . 3. Sonchay and Lenssen (loc. cit.), have shown that the action of hot water upon normal bismuth oxalate results in the production of a basic oxalate having the composition Bi2O3.2C2O3 + aq. By taking advantage of the production of this basic oxalate I thought it would be possible to estimate bismuth accurately.To a measured volume of a solution of a weighed quantity of bis- mnthous oxide in nitric acid, an excess of a saturated solution of oxalic acid was added; the supernatant liquid was poured off from the, precipitate, which soon settled to the bott)om of t<he vessel ; the precipitate was boiled with successive small quantities of water, until the decanted liquid ceased to exhibit an acid reaction ; the residue was dissolved in dilute hydrochloric acid ; and standardised permanganate of potassium was run in-the liquid being maintained at a tempera- ture of 60" or so-until a permanent pale pink colour was produced. The results obtained were accurate, and showed no discrepancies. I quote a few :- This oxalate is stable ; it is only slightly soluble in nitric acid.MUIR ON ESTIMATING BISMUTH VOLUMETRIGALLY.71 (1.) 50 C.C. bismuth-solution = 0.28404 gram Bi : used 51 C.C. stan- dard permanganate = 0.17008 gra,m H2C,0a.2H20 = 0.28348 gram. Bi. difference = - .00056 gram. (2.) 50 C.C. bismuth-solution = 0.28404 gram. Bi.: used 50.7 C.C. permanganate = 0.28182 gram Bi. difference = - -00222 gram. (3.) 10 C.C. bismuth-solution = 0.056364 gram B i : used 10.2 C.C. permanganate = 0.056697 gram Bi. difEerence + *000333 gram. The calculation is based on the assumption of the accuracy of the formula assigned by Sonchny and Lenssen to the basic oxalate of bismuth. I have, by independent analysis, proved the accuracy of their formula. One molecule of oxalic acid found corresponds with one atom of bismuth.4. The solution containing bismuth must be free from hydro- chloric acid as the basic oxalate is readily soluble in that, acid. A large excess of nit.ric acid must also be avoided. Oxalic acid must be added in considerable excess. If the precipitate be thoroughly shaken up with the liquid, and the vessel be then set aside, the precipitate quickly settles, and the superna,tant liquid may be poured off through a filter in a very short time. I find that if the precipitate be boiled for 5 or 10 minutes with successive quantities of about 50 C.C. of water, it is quickly transformed into the basic salt. So soon as the supernatant liquid ceases t o show an acid reaction, the transformation is complete. It is well to employ a solution of permanganate so dilute that at least 50 C.C.are required for the titration. The basic oxalat2 may be dissolved in dilute sulphuric acid in place of hydrochloric ; it is more soluble, however, in the latter acid ; if the solution contains but little hydrochloric acid there is no danger of chlorine being evolved during the process of titration. 5. In applying this process to the estimation of bismuth in solution containing other metals, it is necessary if the solution contain sub- stnncescapable of acting upon, or of being acted on by permanganate, to separate the bismuth from the other metals present. This is easily done by precipitating in a partially neutralised solution with much warm water and a little ammonium chloride. The precipitate must be dissolved in nitric acid, and the liquid boiled down once or twice with addition of the same acid in order to expel all hydrochloric acid, before precipitating as oxalate.The liquid should contain just su&- cient nitric acid to prevent precipitation of the basic nitrate before oxalic acid is added. 6. The failure of those experiments in which I sought to determine bismuth after precipitation as normai oxalate, is evidently to be traced to the partial decomposition of the normal salt by the wash-water with the production of varying quantities of basic oxalate. My own 0 272 THOSlAS ON CUPROUS CHLORIDE AND THE experiments have convinced me that the normal oxalate very readily undergoes partial decomposition by contact even with cold water. 7. Of the various methods which I have devised for the volumetric estimation of bismuth, I regard the oxalate method as the most generally applicable. No special standard solution is required, as permanganate solution is in constant use where volumetric analysis is practised ; the process is simple and accurate, and may be quickly carried out : bis- muth is readily separated from other metals and brought into a con- dition suitable for estimation by this method.

ISSN:0368-1645    

DOI:10.1039/CT8783300070

出版商:RSC    

年代:1878

数据来源: RSC

摘要:

72 THOSlAS ON CUPROUS CHLORIDE AND THE X.-Orrz Cuprozcs Chloride and the Absorption of Cnrboizic 0,:iJe and Hydrochloric Acid Gas. By J. W. THOMAS. SONE years ago I noticed considerable irregularity in the determination of carbonic oxide by absorption when using the reagent in the liquid form, simply passing the solution by the usual pipette into the labora- tory tube or absorption tube, and I came to the conclusion that the incorrectness of my determinations was due to the strength of the acid solution of cuprous chloride which I employed. On trying the cuprous chloride solution, supersaturated with ammonia, concordant results were obtained, and I have since used the re-agent in this form. Although the percentage of carbonic oxide (CO) in a gaseous mix- ture is very accurately determined by explosion with oxygen, the direct determination of the same by absorption is very important when it is required to know the constituents of a gas of unknown composition containing various carbon compounds.I n order there- fore to discover the reason of the irregularity in the absorption of CO by cupi-ous chloride, I undertook the experiments the results of which I now have the honour of submitting to the Society. The method recommended for making cuprous chloride ( Cu,CI2) for use in gas analysis (Sutton’s Volumetric Analysis, p. 28.5) is to strongly acidulate a saturated solution of cnpric chloride (CuCl,) with hydrochloric acid (HCl), and place the solution iu a bottle with copper filings or turnings, leaving it until the liquid is colourless. Thia end can be accomplished by adding almost any quantity of acid, provided the solution is allowed to stand for a sufficient time.In endeavouring t o find a more expeditious method by which a Cu,CI., solution could be made and the acid strength known comparatively, IABSORPTION OF CARBONIC OXIDE, ETC. 73 tried that of Wohler (Watt's Diet., 1st sup., p. 493), dissolving the cuprous chloride in hydrochloric acid of known strength. Although a good method, the process is tedious. By accident, when trying some experiments with hydrochloric acid of different but unknown strengths upon CuCI,, I succeeded in making a colourless solution of Cu2C12 jnstantly, hut although I have tried many times since I have not suc- ceeded in doing it again. The result of successive trials brought out the following method, which is very expeditious, although little more than an old method perfected.A long, narrow stoppered bottle, hold- ing about 4 ozs., is filled thvee parts full with closely-packed. copper tnrnings, previously boiled in caustic soda to remove grease. Intro- duce 6 grams of the ordinary hydrous cuprous chloride, and add 20 C.C. of hydrochloric acid (ordinary strong acid, mur. pur.), and shake for a minute until the cupric chloride is dissolved and forms the black cuproso-cupric chloride, then add 10 C.C. of water and let it run into the bottle slowly so as to float on the surface of the acid liquid; re- place the stopper of the bottle and give the latter a violent shake, when the solution becomes colourless instantly, and entirely converted into cuprous chloride, a large quantity of the white compound being deposited. This is a striking illustration of the action of water in break- ing up unstable chemical compounds.The colourless solution, which is very strongly acid, is saturated with cuprous chloride, and after the addition of a further 30 C.C. of water, is ready for use (sol. B). The fol- lowing is another good method for making cuprous chloride rapidly, the operation being conducted as before in a bottle three parts filled with clean and, in this instance, dry copper turuiigs closely packed. Six grams of anhydrous cupric chloride is placed in the bottle and 20 c.c, of the strongest hydrochloric acid added; on shaking for a minute or two the liquid will become colourless and the acid saturated with cuprous chloride (sol.A). For the purpose of finding the best strength of cuprons chloride solution for use in gas analysis, and the most accurate method of using it, 1 prepared some pure carbonic oxide from formic acid. Soh- tion A and R, and in fact any solution of' cuprous chloride in hydro- chloric acid, no matter how concentrated or reasonably dilute, absorbs carbonic oxide with facility, and the irregularity in the estimation of CO by absorption is due to the subsequent treatment of the solution with potash or soda to neutralise the free hydrochloric acid. When solution A or B, or one of intermediate or weaker strength was used, the absorption of carbonic oxide was rapid, especially if the absorption tube was agitated: when, however, a solution of potash was added, notably if passed from a quick-delivering pipette, a considerable evolu- tion of gas takes place.Much of the gas so liberated is carbonic anhy- dride, expelled from the potassium carbonate present in the potash-solu- .74 THOMAS ON CUPROUS CHLORIDE AND TEE tion by the free hydrochloric acid, but I invariably found that some carbonic oxide was set free, and that the quantity liberated depended upon the manner in which the KHo was added. If the potash (a saturated solution was employed) was passed up slowly and allowed t o form a stratum underneath the acid solution (separated by a film of cuprons hydrate) and, when it was added in excess, the absorption tube is briskly and suddenly agitated, as much as 63 per cent.of the carbonic oxide can be set free and will remain unabsorbed. The car- bonic acid liberated is of course absorbed when the canstic alkali is in excess. When the potash is slowly added and the mercury is allowed to drop through the liquid, or the absorption hbe is feebly agitated, the quantity of CO liberated is very much decreased, and it is possible to reduce it to 4 per cent.; the reaction is not, however, very con- trolable. I next tried whether the liberation of carbonic anhydride caused a similar evolution of carbonic oxide, and, to this end, removed as much carbonic anhydride as possible from the potash solution by quicklime, but did not obtain any better results. Potash-solntion of varying strengths was tried in order to see if the saturated solution was too concentrated, but the results obtained were very similar, car- bonic acid being liberated according to the manner in which the acid solution was neutralised, and the carbonic oxide libeyated in 6 experi- ments amounted to 42% per cent,, 37.4 per cent., 14.4 per cent., 5-2 per cent., 3.8 per cent., and 23-8 per cent.I endeavoured to liberate as much as possible in the two first experiments. Apart from the action of potash, the absorption of carbonic oxide by cuprous chloride, although very rapid at first until from 90 to 95 per cent. has been taken up, is by no means complete under half an hour if left at rest, or 15 minutes if well agitated, as the last traces are very slowly absorbed, more especially when mixed with a large volume of hydrogen, which gas appears to retard the absorption of carbonic oxide more than any which I have tried.Although the hydrochloric acid solution of cuprous chloride may be used to detect the presence of carbonic oxide, it cannot be relied upon for a quantitative determination when potash or soda is used subsequently to neutralise the free hydro- chloric acid. The above experiments relate only to the use of cnprons chloride in the liquid form in apparatus of the Frankland and Ward model, and have little bearing on the absorption of that gas by a papier machi6 ball saturated with the re-agent, and the acid vapour subsequently removed by a potash ball. I found, however, that by allowing the cuprous chloride solution to wet the sides of the absorption tube and the potash ball afterwards to come in contact with the liquid, that the accuracy of the absorption, even by this method of procedure, was perceptibly impaired.ABSORPTION OF CARBONIC OXIDE, ETC.75 The method which I long employed was to add ammonia to the cuprous chloride solution until it was strongly alkaline and the pre- cipitate at first formed waB redissolved, and after the absorption of carbonic oxide had taken place, to add very dilute sulphnric acid until the blue colour of the absorbent was nearly destroyed. There is, how- ever, a liability to introduce error by proceeding in this manner, especially if the hydrochloric acid solution of cuprous chloride is strongly supersaturated with ammonia, and if the latter contains much carbonate, on account of the carbonic anhydride which would be disengaged by the sulphuric acid and the risk of adding too much of the latter to set hydrochloric acid free.I have recently made some experiments with a view of using the ammonio-cuprous chloride solu- tion in a form which shall be su5ciently neutral to dispense with the necessity of employing any re-agent to remove the vapour of ammonia. The following method was found to give very concordant results when tried on the same sample of a gaseous mixture. A small pipette (straight) is required, having a long narrow point with an orifice amaller than a wash-bottle jet. Some ammonia (strength one NH3, -880, to one of HzO) is sucked into the pipette and a little of the cuprous chloride solution (B) is poured into a small porcelain crucible or a test-tube on foot, cut short, and stirred with the pipette containing ammonia.As soon as the slightest shade of blue remains permanent, on stirring, fill the usual bent pipette with the solution and introduce the latter into the absorption tube at once, the gas being previously t-eady for the absorption. The ammonia when added to the hydrochloric solu- tion of cuprous chloride forms ammonio-cuprous chloride (in addition to ammonium chloride) which soon separates in the form of very minute transparent crystals, and only a very little, of what appears to be a hydrate, is formed, and the slightest excess of ammonia over that required to neutralise the acid is indicated by the liquid assuming a blue colour. It is well known that when ammonia is added in excess to a colourless solution of cuprous chloride, in the absence of oxygen, the solution remains colourless, but if the cuprous chloride solution is exposed to the air for a few seconds, even some of the cuprous chloride is oxidised, and as quickly converted by the free acid into cupric or cuproso-cupric chloride, which gives a blue colour with ammonia as soon as the free acid is neutralised.By proceeding thus there will be no free ammonia to pervade the gas (the first blue colour being due to ammonio-cuproso-cupric chloride), and consequently no necessity to add acid subsequently, and although much of the copper compounds are precipitated, the re-agent absorbs carbonic oxide nearly as well as the hydrochloric acid solution of cuprous chloride. It is unnecessary to give details of the experiments made, but I may mention that they76 THOMAS ON CUPROUS CJXLORIDE AND THE were tried on mixtures of carbonic oxide and hydrogen, carbonic oxide and marsh-gas, carbonic oxide and ethane, and coal gas (free from carbonic anhydride and oxygen), and in all instances gave very good results, but the absorption is not complete under half an hour, unless the vessel is briskly agitated. A large volume of a mixture of hyditogen and marsh-gas was measured and then transferred over the ammonio-cuprous chloride solution prepared as above in order to see if the pressure of the gas varied, the difference in the two readings was *03 per cent.The hydrochloric acid solution of cuprous chloride (B) answers well, but it should be kept colourless, and to this end the copper turnings must always be above the liquid and the bottle well stoppered, else the solution soon becomes less acid, owing to the formation of fur- ther quantity of cuprous chloride, through oxidation and subsequent reduction.Since the above was written I have made further experiments, with a .view of using phosphoric acid in the place of sulphuric acid for neutralising an alkaline (NH,) solution of cuprous chloride, and I have also endeavoured to arrive a t some method of neutralising the acid solution of cuprous chloride by other reagents than potash and soda, Phosphoric acid does not appear to liberate hydrochloric acid from ammonium chloride in the cold, and only traces of hydrochloric acid are given off when a concentrated solution of ammonium chlo- ride is boiled with phosphoric acid.The free hydrochloric acid in an acid solution of cuprous chloride is neutralised by phosphate of ammonia, phosphate of soda, oxalate of ammonia, oxalate of soda, sulphate of soda, sulphate of ammonia, and alkali tartrates. It is stated (S utton’s Volumetric Analysis, p. 367) that sodic phos- phate and sodic snlphate owe their absorbent powers (for HC1 gas) to the water of crystallisation which they contain, and in Watts’s Diet. of Chem., i, 282, “the ball of sulphate of sodium is recom- mended to be large, else, if much HC1 gas be present, the sulphate of sodium is apt to become deliquescent.” I find that phosphate of soda and sulphate of soda do not absorb HCl gas by virtue of their water of crystallisation, but by giving up half their Na to form acid salts and sodium chloride, When HC1 is added to a saturated aqueous solution of sodic phosphate a t the boiling temperature, the whole of the acid combines with the sodium of the phosphate and forms sodium chloride, providing there be two equivalents of sodium to one of HCl in the solution.From experiments, it appears probable that in the presence of water at the ordinary temperature hydrochloric acid displaces all the sodium i n hydric-disodic phosphate. Either sodic sulphate or disodic phos-ABSORPTION OF CARBONIC OXIDE, ETC. 77 phate may be used to neutralise the €ree hydrochloric acid in the solu- tion of cuprous chloride after the absorption of carbonic oxide, but although a saturated solution of the former contains more sodium, a solution of the latter does not apparently contain free hydrochloric acid when the quantity of that acid present is nearly equivalent to the total sodium.A saturated solntion (aqueous) of ammonium sulphate is, however, the best adapted for neutralising the free acid in a cuprous chlo- ride solution, as this salt is very soluble in water, and also possesses the advantage of holding the cuprous cliloride in solution without destroy- ing its absorbent properties, so that the solution of ammonium sulphate might, if required, be added to the acid solution of cuprous chloride before introducing it into the laboratory tube. The hydrochloric acid solution of cuprous chloride (B) requires about 80 per cent. of its volume of a saturated solution of sodium sulphate to neutralise the hydrochloric acid, but it is best to add volume f o r volume.Per- fectly dry ammonium sulphate absorbs HC1 gas very readily. When R satmated solution of sodium sulphate or disodic phosphate is added to a HCI solution of cuprous chloride after the absorption of carbonic oxide, no carbonic oxide is liberated, but a white crystalline precipitate of cuprous chloride falls down and is afterwards partly dissolved in the sodium chloride formed. If, after a saturated solution of disodic phosphate is added, the tube containing the cnprous chlo- ride, &c., is allowed to stand over mercury for several days, the undis- solved cuprous chloride becomes converted into cupric phosphate by the reduction probably of orthophosphate to pyrophosphate, but no carbonic oxide is liberated. When a gaseous mixture containing hydrochloric acid is analysed in an apparatus on Dr. F r a n k l a n d or Prof. M'Cleod's model, a few drops of a saturated solution of sodium sulphate may be used to re- move hydrochloric acid gas, and potash and pyrogallic acid may be subsequently introduced to absorb carbonic anhydride and oxygen without washing out the laboratory tube, but owing to the moisture which invariably remains in the eudiometer the result will be much more accurate if the absorption of the hydrochloric acid gas is con- ducted in an ordinary absorption tube, and the remainder of the analysis done in the apparatus mentioned. The experiments relat- ing to the action of hydrochloric acid upon sodium sulphate, &c., will be described in another paper.

ISSN:0368-1645    

DOI:10.1039/CT8783300072

出版商:RSC    

年代:1878

数据来源: RSC