摘要:

EQUILIBRIA ACROSS A COPPER FERROCYANIDL ETC. 1313 CXXII. -Eyuilibria Acrom a Copper Ferrocyanide and an Amy1 Alcohol Membrane. By FREDERICK GEORGE DONNAN and WILLIAM EDWARD GARNER. DONNAN and Allmand (T. 1914 105 1941) investigatd the dis-tribution equilibrium of potassium and chlorine ions acro8s a copper ferrocyanide membrane. I n the method adopted a solution of potassium ferrocyanide was placed on one side of the membrane and a solution of potassium chloride on the other. The results, whilst affording undoubted confirmation of Donnan’s theory (Zeitsch. Elektrochem. 1911 17 572) were complicated by the uncertainty as to the manner of ionisation of potassium fern-cyanide. It was considered that this difficulty would be overcome if solutions of two ferrocyanides were employed on thO two sides of the mEmbrane especially if these salts were ionised to the same extent.I n the present investigation the sodium potassium and calcium salts were found to be suitable and mixtures of these were used to test the validity of the theory. With a mixture of sodium and potassium ferrocyanides equil-ibrium will be set up by an interchange of sodium and potassium ions since t-he membrane is not permeable to ferrocyanogen ions. Assuming that the ions obey the laws of ideal solutions the equa-tion for the equilibrium is given by where the symbols indicate molar ionic concentrations. This equation was derived from thermodynamical considerations by Donnan (Zoc. cit.) and its derivation is also possible from the kinetic theory as follows The number of ions of any one kind penetrating the membrane is proportional to its ionic concentration, C,; the number of ions exchanging across the membrane is pro-portional not only to the concentration C, but also to the concen-tration of the ion C2 which interchanges with the first across the membrane that is, N = K .C . C,. I n the case of sodium and potassium ions a t distribution equil-ibrium four kinds of interchanges across the membrane are possible. Two of these (a) and (6) (exchange of like ions) do not affect the final concentrations in any way. The four interchanges are shown below 1314 DONNAN AND QARNER EQUILIBRIA ACROSS A COPPER ( b ) [KO] +I- [K'] I A t equilibrium the number of exchanges due to ( c ) that is, N,=K[Na,']~,'] must equal that due to (d) that is, Nd = K[Na,'] [Klo] therefore [Na,'] [K,'] = [Naz'] [K,'].When the calcium salt is substituted for the potassium salt the equation of equilibrium becomes : ( a ) [Ca,'.] [Na,*l2= [Ca '][Na1'I2, (71) [Cal] [Na,I2 = [Ca,] [Na& and if the degree of ionisation of the two salts is the same. I n the case of the sodium-potassium cells which were investi-gahd the experimental results showed good agreement with the requirements of the theory. The degrees of ionisation of the two salts are very similar so that the ratio of thel molar concentrations of the salts is the same as that of the ionic concentrations. The sodium-calcium cells however gave unexpected results. Whereas the equation ( b ) which refers to the concentrations of the two salts holds within thO limit of the experimental error of the analysis it was found &hat equation ( a ) above does not accurately represent the relationship between the ionic concentrations of the calcium and sodium salts on the two sides of the membrane.The activities of the ions in this case appear to be more closely related to the molar than to the ionic concentrations. This result also may indicate that adsorption plays an important part in the trans-ference of the ions across the membrane. Further experiments were carried out in order to find a liquid membrane which would be permeable to one electrolyte and impermeable to a second which contains an ion common to the first. With this purpose in view the solubilities of several salts in moist organic solvents were determined.Amy1 alcohol was found to be the most satisfactory of these solvents and potassium and lithium chlorides the most suitable electrolytes. Since lithium chloride is readily soluble in amyl alcohol (a saturated solution is 1 -83N) and potassium chloride is practically insoluble (a saturate 1FICBBOCYANIDE AND AN AMYL ALCOHOL MEMBRANE. 1316 solution is 0*0048N),* it was hoped to set up a cell of the following I. KCI LiCl in water I Amy1 alcohol I LiCl in water 11. which is of the same type as that investigated by Donnan and Allmand with the copper f errocyanide membrane. It was not however practicable to use en amyl alcohol membrane owing to the slow rate of diffusion of lithium chloride through the amyl alcohol. The problem was therefore approached in an indirect manner.Determinations were made of the distri-bution concentrations of lithium chloride between amyl alcohol and water a t 2 5 O . Aqueous solutions of lithium chloride and lithium and potassium chlorides were shaken with amyl alcohol, and the two layers separated and analysed. The concentration of the lithium chloride in 11 which is in equilibrium with a mixture of the two chlorides in I was calculated from these results. The calculation of the ionic concentrations is complicated by the high values of the viscosity of the solutions and by changes in the state of hydration of the lithium ion with concentration. Green (T.? 1908 93 2023) has deduced the degree of ionisation of concentrated lithium chloride solutions from measuremenla of the conductivity of solutions of lithium chloride of which the viscosity has been increased by means of sucrose.The chief objw tion to the values which are obtained in this way lies in the hydr-ation of the lithium ion which will be affected by the addition of sucrose to the solutions. The concentrations of the ions and the undissociated part of the electrolyte have however been calculated using Green's values for the degree of ionisation of lithium chloride and the results are in fairly satisfactory agreement with theory. The agreement is better in those cases where the total concentration of the electro-lytes is below I>Tv'. [Li,'] . [Cl,/] = [Li,'] . [Cl,/]. The distribution-coefficient of lithium chloride between amyl alcohol and water has been calculated and it appears that lithium chloride occurs in amyl alcohol solution as double molecules.The coefficient is however only a constant over a small range of con-centration and above 5N the coefficient increases. The increase is probably associated with errors in the degree of ionisation due to the dehydration of the lithium ion. The experimental work in this paper leads to the conclusion that the same equilibrium relationships are established whether the * Moist &my1 alcohol. type : Thus [LiCI],=~iCl] and 3 D 1316 DONNAN AND GARNER EQUILIBRIA ACROSS A UOPPER equilibrium is brought about by the transference of ions as is the case with the copper ferrocyanide membrane or by the transference of the undissociated part of the electrolyte as is the cage with the amyl alcohol membrane.EXPERIMENTAL. The osmometer vessels used in the determination of the ratios of Che ions were those described by Donnan and Allmand (T. 1914, 106 1944). The copper ferrocyanide membrane in parchment paper was damped in position between two shallow cylindrical vessels and was separated from the supporting rim by rubber bands. The vessels which were of Jena glass were fitted with side-tubea t o facilitate the introduction of the solutions and the volume of each vessel was about 100 C.C. The membranes were prepared by the meth.od described by Donnan and Allmand (Zoc. cit.) and the parchment paper was usually left in contact with tbe solutions for two days. The membranes were tested for leaks by placing a ferrocyanide solution on one side and an isotonic solution of sucrose on the other and no leakage occurred over a period of six weeks.The ferrocyanides which were used in the investigation were purified by crystallisation from water. The calcium and ammonia ferrocyanides were prepared from hydrof errocyanic acid (No yes and Johnston J . Amer. Chem. Soc. 1909 31 991). Since potassium ferrocyanide forms insoluble double salts with magnesium and calcium ferrocyanides of the type it was not possible to use potassium ferrocyanide against these salts in the cells. The precipitation of the insoluble salts takes place slowly a t the ordinary temperature but quicker on heating, as if a chemical change were taking place. The double salts with sodium ferrocyanide are soluble in water.R,K2FeCy,,3H,O, Po tassium-Sodium Ferro cyanide Cells. The usual pro-cedure was t o place a solution of potassium ferrocyanide in one side of the cell and a solution of sodium ferrocyanide in the other. The time required for the attainment of equilibrium was deter-mined by conductivity measurements and no change in the con-ductivity could be observed after an interval of one week. The cells were however allowed to remain with occasional shaking, over a period of three t o five weeks in which time equilibrium was certain to have been reached. Solutions were used of a strength 0.025 molar FERBOCYANIDE AND AN AMYL ALCOHOL MEMBRANE. 1317 I n order to prevent changes in the ferrocyanide soluhioner with time several precautions were necessary-(1) When pieces of well-washed copper ferrocyanide membrane were placed in solutions of ferrocyanides it was observed that the strength of the solutions diminished several per cent.in three or four days and a t the same time the presence of sulphates in the solutions was detected. The change was almost entirely due t o that side of the membrane which was last exposed to the copper sulphate solution. The adsorbed copper sulphate on this side of the membrane reacted with the ferro-cyanide solutions with the forma-tion of sulphates and a slight in-crease in the thickness of the mem-brane. The copper f errocyanide membrane after being clamped in position was on this account washed for three t o four days with f errocyanide solutions of the same concentrations as those to be used in the experiment.(2) Another source of trouble was the oxidation of the ferro-cyanides by the small quantity of air enclosed in the osmometer vessels. The oxidation was also considerable if the conductivity of the solutions was measured from time to time in the ordin-ary conductivity vessels. To make the change due to this cause small as possible a special con-ductivity vessel was constructed (Fig. 1) which could be filled with nitrogen. The amount of air in the osmometer vessel was also reduced to a fraction of a c.c. and in the majority of the experiments of which the results are given in the tables the cells were not opened until immediately before analysis. (3) When solutions of different concentrations were employed on the two sides of the membrane osmosis of water was prevented by the addition to the solution of the requisite amount of sucrose.The amounts which were added were calculated from conductivity 3 Ds 1318 DONNAN AND GARNER EQUILIBRIA AOROSS A COPPER data. and Allmand have shown (Zoc. &.). two or three days. The effect due to this cause is however small as Donnan The solutions were kept in the dark and the celIs shaken every Method of Andysis. The solutions of the ferrocyanides were decomposed with con-centrated sulphuric acid and the sulphates of the alkali metals converted into chlorides by the precipitation of the iron (twice) with ammonia and of the sulphate with a slight excess of barium chloride. After the removal of the barium as carbonate the mixed chlorides in the solution were obtained by evaporisation and weighed.The potassium was determined as perchlorate and the sodium calculated by difference. The f errocyanide concentrations were determined before filling into the cells and after the equilibrium had been reached by titration against potassium permanganate solution. These analyses serve as a check on the result8 obtained by the gravi-metric analysis The variation in the ferrocyanide concentrations as determined by the three methods outlined above rarely exceeded 1 per cent. A method of analysis based on a conductivity method was not found to give the requisite degree of accuracy. The results for the potassium-sodium cells are given in table I. The weights of potassium chlorate and the mixed chlorides are given (in order to indicate the possible errors of the analyw) and in columns 6 and 7 are included the molar concentrations of the sodium and potassium on both sides of the membrane*A and B .In column 8 is found the total concentration of the metals and in 9 and 10 four times the total concentration of ferrocyanogen before and after the experiment. The results show that no large amount of oxidation or absorption of the salts has taken place during the period of the experiment. The ratios of the sodium and the potassium in the solutions on the two sides of the cell are compared in the last column and it will be observed that the ratio is the same for (a) and ( b ) within experimental error. The ratio of the ionic concentrations will be but little different from those given in the table since the degrees of ionisation of the sodium and potassium f errocyanides are very similar.The conductivity of 0.025 molar solutions of potassium sodium and ammonium ferrocyanides was determined a t 26O and the degree of ionisation calculated. The results are given in table 11 Cell No. 1 a .................. b .................. 2 a .................. b .................. 3 a .................. b .................. Pa .................. b .................. 5 a .................. b .................. TABLE I. Potassium and Sodium Ferrocyanide Weight8 of I ’F KCIO,. 0.3123 0.3146 0.2064 0.2069 0-1984 0.2181 0.1839 0.2163 0.1731 0.1785 pa. 0.2669 0.2666 0.1992 0.1982 0-1973 0.2181 0- 1970 0.2332 0.1936 0.1993 NaCl.0.0979 0-0965 0.0881 0.0869 0.0905 0-1007 0-0980 0.1168 0.1004 0-1033 KC1. 0.1680 0.1691 0.1111 0.1113 0.1068 0.1174 0.0990 0.1164 0-0931 0.0960 Na. 0.0336 0.0410 0.0603 0.0496 0-0516 0.0576 0.0569 00666 0.0673 0.0884 K. 0.0461 0-0667 0.0497 0.0498 0.0477 0.0624 0.0443 0.06206 0.0416 0.064 1320 DONNAN AND GARNER EQUILIBBIA AOROSS A COPPER TABLE 11. Degree of Zonisation of Ferrocyanides. salt. +40. rctusch. 1000. (NH.J,Fe(CN) ..................... 383.1 742.0 61.8 (Na),Fe(CN) ..................... 337.1 647.6 6% L (K),Fe(CN) ........................ 393*0* 742.0 62.9 A0 Kohl-* Noyes. Thus the relationship given below has been proved to be correct, that is, Since the activities of the potassium and sodium ions are prob-ably very similar the equation deduced by Donnan has been shown to hold.S o d ~ ~ m - ~ 4 m r n o n ~ t ~ m C e h . Ammonium ferrocyanide solutions slowly attack the copper ferrocyanide membrane. The membrane thickens and changes in colour from a dark brown to a reddish-brown It does not how-ever appear to break down as on on0 occasion a cell was made up of a solution of ammonium ferrocyanide on the one side and an isotonic sucrose solution on the other. After five weeks the sucrose solution was tested and it was found that no ferrocyanide had diffused through the membrane. The' concentration of ferro-cyanide had however diminished and the colour of the me,mbrane on the one side had changed to a reddish-brown.I n consequence of these irregularities only one sodium-ammonium cell was examined. The ratio ____ on the two sides were found to be 0.8480 and 0.8595 respectively. "H41 Sodium-Calcium Cells. The cells were made up as described previously. The amounts of sucrose used to prevent osmosis were the same as those used with the potassium and sodium solutions of the same concentra-tion. The concentrations of sucrose are probably too great in the case of the calcium solutions sinoe the calcium salt is less ionised in solution than the potassium salt but no appreciable osmosis occurred. On the other hand when the concentrations of sucrose were calculated from Sherrill's equations ( J . Amer.Chern. SOC., 1910 32 742) a considerable amount of osmosis occurred. The rate at which equilibrium is reached is about ita rapid a FBRROCYANIDE AND AN AMYL ALCOHOL MEMBRANE. 1321 in the case of the sodium-potassium cells. In some cases the changes in concentration were followed by conductivity measure ments. In Fig. 2 are given the changes of conductivity of a sodium-calcium and a sodium-magnesium cell respectively. Time indays. 0. 1. 3. 4. 5. 6. 7. 8. 10. 12. Bridge-reading. -Na-Mg . 2974 3358 3800 - 3994 - 4040 - - 4060 / -. Na-Ce . 3470 4082 - 4600 - 4635 - 4643 4635 -Time in days. From the curves and the above table it will be readily seen that the rate of exchange of sodium and calcium ions is approxim 1322 DONNAN AND UARNER EQUILIBRIA ACROSS A OOPPER ably equal to that of sodium and magnesium ions and that a constant reading is obtained in about ten days.The slight fall in the bridge-reading a t the end of the experi ment (Na-Ca) is probably due to the oxidation of the ferrocyanide, Method of Analysis. The analysis of the solutions of calcium and sodium ferro-cyanides gave rise to considerable trouble owing to the small volume (100 c.c.) which was available for analysis. The ferro-cyanide solution was evaporated to dryness and decomposed with concentrated sulphuric acid. The mixed sulphates were dissolved in dilute hydrochloric acid and the iron was removed as hydroxide with ammonia. The calcium was precipitated as oxalate and con-verted into oxide. Traces of iron were sometimes present in the oxide which was on this account dissolved in hydrochloric acid and the iron precipitated.The calcium was then weighed as sulphate. The filtrate containing sodium sulphate was evaporated to dryness traces of iron were removed and the sodium was weighed as sulphate. The results of the analyses are shown in table 111 and it is found that where the concentrations represent the total concentrations of the calcium and sodium atoms in the solution. The ratio of the equilibrium concentrations of the calcium ferro-cyanide on the two sides of the membrane is slightly higher than the ratio of the squares of the equilibrium concentrations of the sodium ferrocyanide but the variations are of the same order as those due to errors of analysis. The change in the degree of ionisation with concentration is known for potassium and calcium ferrocyanides but not for sodium f errocy anide.lO0u K,Fe(CN),. Ca,Fe(CN),. (1) 0.06 Mol- ............... 48.6 22.1 (2) 0.026 Molm ............ 63.1 23.5 (Noyes and Johnston J . Amer. Chem. Soc. 1909 31 1010). Similar figures have been obtaihed for calcium ferrocyanide in the course of this work. Assuming that the sodium salt resembles potassium f errocyanide, and that the ionisation of the mixed sodium and calcium salts is the game as in solutions of the pure salts with concentratioa corr TABLE 111. Calcium-Sodium Cells. Normality of 1 - r-4Fe(CN),. 2Ca. 00573 0.0453 00615 0.0349 0.0643 0.0397 0652 0.0362 0.1287 0832 Na. 0.0640 0.0573 0.0867 0.0654 0.0726 0.0629 0.0893 0.0670 0.0686 0.0661 I.01213 0.1026 0.1482 0.1003 0.1269 0.1026 0.1646 01043 0.1973 0.1393 Ir. 0.1196 0.1034 0-1495 0.1030 0.1286 0.1049 0.1634 0.1042 0.2140 0- 1303 111. 0.1185 0.1029 0.1473 0.1026 0.1023 0.1236 0.1645 0.1032 0.1995 0.1 420 CC%l rn 1.266 1.762 1.368 1.802 1.547 “%I’ “%I’ 1.248 1-767 1.332 1.776 1495 “bll C%I 1.117 1.326 1.1154 1.335 1.22 1324 DONNAN AND QARNER EQUILIBRIA ACROSS A COPPER sponding with the ferrocyanogen-ion concentration (Arrhenius) it is possible to calculate the ratios of the .calcium-ion concentration and the ratios of the squares of the sodium-ion concentrations on the two sides of the membrane.These are given in the table. The results (table IV) show that the relationship does not hold so strictly as equation (1). TABLE IV. Calcium-Sodium Cells. [ X I "all' [Ca;'I* "%'I: No. [Ca,I' "4 [Ca; 7 [Na2']*' 1 ..................... 1.27 1.25 1-25 1.19 2 ..................... 1.76 1.76 1.68 1.58 3 ..................... 1-37 1-33 1.33 1-25 4 ..................... 1-80 1.78 1.73 1.61 5 ..................... 1.65 1.50 1-49 1.36 The activities of the ions of these two salts thus appear to be more nearly proportional to the molar concentrations than to the ionic concentrations. Cell 3 which was opened once during the experiment shows that 2-3 per cent. of oxidation has taken place and in cell 5 a change due t o osmosis, occurred. Neither of these changes appears to affect the ratio to any great extent.I n cell 5 it should be noted that the concentration of the calcium ferrocyanide is greater than that of the sodium ferro-cyanide; in the other cells the reverse is the case. Irregularities occurred in two of the cells. Amy1 Alcohol Nembrane. The results of some preliminary experiments on this membrane are given below. Materials.-Amy1 alcohol (b. p. 131.5O) was obtained by repeated fractionation of fuse1 oil through a six-bulb fraction-ating column. One sample of the alcohol was used throughout the work. The lithium chloride was free from calcium and was completely soluble in amyl alcohol. Its solution in water was neither acid nor alkaline. It was also analysed by conversion into lithium sulphate followed by the estimation of the sulphate in this sub-stance as barium sulphate BERBOCYANIDE m D AN AMYL ALCOHOL MEMBRANE.1326 Method.-Aqueous solutions of lithium chloride were shaken with amyl alcohol in stoppered bottles in a thermostat kept a t 2 5 O , and when equilibrium had been reached the two layers were analysed. A dysis.-The aqueous solutions of lithium chloride were estimated volumetrically with N / 10-silver nitrate and the results were checked by analysis of the lithium as sulphate. A known volume of the amyl alcohol layer was placed in a Jena-glass distilling flask and the amyl alcohol distilled off. Water was added to the residue and the solution titrated with silver nitrate. The solutions containing the potassium and lithium chlorides were analysed according to the method employed by Gooch (Proc.Amer. Acad. 22 N.S. 14 177). A known volume of the aqueous solution was evaporated in the presence of 10 C.C. of amyl alcohol, and a little hydrochloric acid added to convert any lithium hydr-oxide into chloride. The lithium chloride dissolves in the amyl alcohol and potassium chloride is left behind. The residue is collected and washed with hot amyl alcohol. The lithium is then estimated as sulphate and the potassium chloride dissolved in water and estimated with silver nitrate. To check the results the solution of the mixed chlorides was titrated directly with silver nitrate. Results.-The concentrations of lithium chloride in the two layers are given in table V. The concentration of the lithium chloride in amyl alcohol diminishes rapidly with decrease in the concentration in the aqueous solution.The degree of ionisation of lithium chloride solutions cannot be given with any accuracy. The viscosity of the solutions is so great that allowance must be made in the derivation of the degree of ionisation from the conductivity results. Green (loc. cit.) has determined the TABLE V. Distribution of Lithium Chloride between Amy1 Alcohol and Water.* LiCIA [LialArn. [Licl]Aq. [LiclJAm. Total. Total. 100~. undissociated. [E 12.54N 8.49 7.77 6.68 5-00 3.14 3.00 2.7 1 1-735N 0.903 0.683 0.387 0.1266 0.0342 0-0314 0.0251 36.9 47.8 50.6 54-2 60.1 66.3 66.7 67.7 7,9137 4-43 3.85 3.06 2.00 1-06 1.00 0.88 0-0273 0.0428 0.041 8 0.0366 0.0277 0.0256 0.0262 0.0268 * In the calculation of the equilibrium constant a correction has been made or the d q p e of ionisation of the LiCl in the amyl dcohol solution 1326 DOHITAX ANXI QABNER EQUILIBRIA ACROSS A OOPPER viawity and conductivity of solutiona of lithium chloride ov0r a wide range of concentration and haa corrected for the effect of viscoeity by the addition of sucrose to the solution.The con-ductivity at infinite dilution was calculated over a wide range of viscosities ; the degree of ionisation is obtained directly from the A equation a=- where h is the conductivity of a solution of lithium nr chloride and & the conductivity a t infinite dilution of a solution of lithium chloride containing sucrose and with the same viscosity aa the first solution.Washburn ( J . Amer. Chem. SOC. 1911 33, 1461) finds that the relation between the degree of ionisation the conductivity and the viscosity is given by the relation a=. fo 4 l f ' 0 where f repreeents the fluidities and m=0'94 but this holds inly from 0-1'ON. For more concentrated solutions m varim with the concentration. In table VI m is given for the concentrations TABLE VI. Degree of Ionisation of Concentrated Lithium Chlam'de Solutions. Normality. 0.0 2.0 3-0 4.0 5.0 6.0 7.0 8.0 9.0 10.0 11.0 12.0 f* 11 1-67 84-87 73-84 63.66 64.59 46.27 38.60 31-55 25.28 19.76 16.07 11.42 m. -0.900 0.858 0.832 0.812 0.798 0.786 0.771 0-769 0.749 0-738 0.729 A.115-3 61.63 52.58 44.76 ;I:; 26.04 21.05 16.725 13.225 10.395 8.149 loo [$)" 100a(Green). -68.4 66.2 61.9 68.7 55.6 62.0 48.4 44.8 42.0 39.6 37.2 -69.9 86.7 63.6 60.1 66.8 63.1 4945 46.0 43.1 40.6 38.2 2-12N and the value of rn decreases from 0.90 to 0.73. The degree of ionisation calculated from the above equation and the data given by Green will be found in columns 5 and 6. The values are the same as those obtained by Green t o within 1-2 per cent. The method of calculation is open to the objection that the lithium ion is probably hydrated in solution and that the conductivity at infinite dilution Aj is given by lithium ions which are prob-ably hydrated to a different extent than is the caw in a pure solution of lithium chloride with the same viscosity.The calcu-lated values of the degree of ionisation will therefore be the more accurate for the more dilute solutions. The values of a in table VI column 6 are used in tables V and VII. In table V the amount of undissociated lithiu FERBOOYANIDE AND AN SMYL BLCOHOL MICMBBANE. 1327 chloride (column 4) is calculated from these values of a. From the figurea in the last column it will be seen that lithium chloride is associated i n amyl alcohol solution to double molecules. A constant for [LicllAm- ~ [LiUlX,. is only obtained between 2N and 5N. Above this concentration the coefficient rises from 0.0277 to 0.0428, and then falls. The most probable cause of this deviation is dia-cussed above and it appears that the calculated concentrations of the undissociated molecules are too low.The molar concentrations of the solutions of mixed chlorides in I. (p. 1315) and the corresponding equilibrium concentrations of lithium chloride in the amyl alcohol are given in table VII. TABLE VII. Results with Amy1 Alcohol Membrane. Undissociated Total Total LiCl in Total LiCl [Li’] x [Cl’]. KC1. LiCl. amyl LiC1. - No. I. I. alcohol. 11. I. 11. I. 11. 1 0.944 3.504 0.554 3-78 1-44 1.36 5.94 646 2 1.200 2.613 0.303 2-95 0.945 0.976 4.046 3.900 3 0.962 5.45 0.2236 5-80 2.39 2-46 11.02 11-17 From the amount of salt dissolved by the amyl alcohol and the data in table V the corresponding values of lithium chloride in 11. are calculated and given in column 5.These figures represent the concentrations of lithium chloride in equilibrium with the solu-tion of mixed chlorides in columns 1 and 2 across the amyl alcohol membrane. The ionic concentrations are obtained from Green’s values for lithium chloride ‘and from Kohlrausch and Grotman’s values for potassium chloride. The degree of ionisation for the higher concentrations of potassium chloride are obtained by extra-polation from the latter values. The degrees of ionisation of lithium and potassium chlorides are apparently very similar. The ionic concentrations of the solutions of the mixed chlorides are calculated on the assumption made in the case of the calcium-sodium ferrocyanide cells. Cells 1 2 and 3 show good agreement with the equations [LiCl] = [LiCl] and [LiJ [Cl,’] = [Lii] [Cl,’]. The agreement, which is better than would be eqected supports the values for the degree of ionisation of lithium chloride which were obtained by Green. Summary. Determinations have been made of the equilibrium concentra-tione of solutions of sodium and potassium ferrocyanides an 1328 NIEEENSTEIN THE COLOURINQ MATTER OF (b) I. Sodium ferrocyanide ( c ) I. Sodium ferrocyanide Calcium f errocyanide 11. Ammonium ferrocyanide 11. The solutions in (a) ( b ) and ( c ) were in the neighbourhood of 0-025N. A liquid membrane has been investigated; amyl alcohol was chosen as the most suitable solvent and the electrolytes employed were potassium and lithium chlorides. (’) I* Lithium 1 Amy1 alcohol 1 Lithium chloride 11. Potassium chloride Lithium chloride gives rise to double molecules in amyl alcohol solution and a constant is obtained for the partition-coefficient up t o 5 N . So far as the preliminary experiments go the equilibrium con-centrations of the lithium and chlorine ions and the undissociated part of the electrolyte agree with Donnan’s theory. ~VERSITY COLLEQE, GOWER STREET, W.C. 1. [Received September 22nd 1919.

ISSN:0368-1645    

DOI:10.1039/CT9191501313

出版商:RSC    

年代:1919

数据来源: RSC

摘要:

1328 NIEEENSTEIN THE COLOURINQ MATTER OF CXXII1.-The Colouring Matter of the Red Pea GaZl. By MAXIMILIAN NIERENSTEIN. THE colours of oak galls are very varied and rich. They range from white and cream through all tints of yellow to deep orange, from pale green t o a rich dark hue and through almost every shade of red some being very beautiful and attractive. These red colours are generally ascribed to the presence of anthocyanins, which are supposed t o be derived from the tannins present in galls (compare Gertz “ Studien ofver Anthocyans,” 1906 ; Connold, “ British Oak Galls,” 1908 ; Eiister “ Die Gallen der Pflanzen,” 1911 ; Magnus “ Die Entstehung der Pflanzengallen,” 1914). Our knowledge of the anthocyanins has been fundamentally increase THE RED PEA GALL. 1329 by the recent investigations of Wheldale Willstiitter Everest and others (compare Petkin and Everest ‘‘ The Natural Organic Colouring Matters,” 1918).Their researches have conclusively proved that the anthocyanins are derived from the different flavones present in plants. This suggested an inquiry into the colouring matter of the so-called anthocyanin of the “red pea gall,” fre-quently found on the leaves of different British oak trees especially Quercus pedunculata when galled by Dryophanta diuisa Adler. It seemed reasonable to expect that the anthocyanin of this gall would in all probability be derived from cyanidin the anthocyanin of quercitin and if so it might furnish some evidence regarding the much discussed question as to the relationship between the pathological products produced by the gall and those normally present in the plant (compare Dekker ‘ I Die Gerbstoffe,” 1913).It was incidentally also thought possible that an anthocyanin derived from a gall might prove t o be closely allied to quercetone or isoquercetone both anthocyanin-like oxidation products of quercetin described by Nierenstein and Wheldale (Ber. 1911 44, 3487) and Nierenstein (T. 1915 107 869; 1917 111 4) as it was probable that the accelerated oxidative processes common to larva and imagines (compare Krogh “ The Respiratory Exchanges of Animals and Man,” 1916) which in addition to numerou~ inquilines are present in large numbers in galls (compare Connold, Zoc. cit.; Kuster Zoc. cit.) would favour the production of an oxidation product such as quercetone and not that of a reduction product such as cyanidin (compare Everest Proc.Roy. SOC. 1914, [B] 87 444). The investigation of the red colouring matter derived from the “red pea gall ” has however t o some extent proved disappointing. It was found that dryophantin the name suggested for this pig-ment was in no way allied either t o the flavones or to the anthocyanins but that it consisted of purpurogallin and two mole-cules of dextrose. On the other hand it must be mentioned that purpurogallin has not previously been found in nature. Dryo-phantin is derived from pyrogallol like gallotannin and is there-fore of pathological origin like the latter. Dryophantin how-ever cannot be regarded as an anthocyanin and probably the same can be said of the other so-called anthocyanins derived from plant galls.It is therefore proposed t o classify these red pigments in a new group of natural organic colouring matters to which the name gallorubrones is assigned 1330 NIERENSTEIN THE COLOURJNCI MATTER OF Preparation of Dryophantin. The galls used in this investigation were collected in the vicinity of Bristol and East London during the months of August and September 1913 1915 1917 and 1918 and care was taken to avoid admixtures with the different galls of the Neuroterus species frequently met with on the same leaves as the galls of Dryophanta diuksa. I n all 94 grams of the galls were collected and the dried material was powdered and extracted in a Soxhlet apparatus a t first with ether and subsequently with chloroform so as to remove W ~ X chlorophyll and the so-called gall-fats.The carefully dried powder was then again extracted in a Soxhlet apparatus with alcohol which dissolved both the colouring matter and the tannins. The alcohol was distilled off in a vacuum and the viscid residue redissolved in water. The cold aqueous solution made up to 150 c.c. was shaken with 5 grams of fat-free caseinogen to remove the tannins (compare Korner and Nierenstein Chem. Z e i t . 1911, 36 31) filtered and extracted with ether. The ether left on evaporation only traces of a tarry substance apparently a by-product. The aqueous solution was evaporated under diminished pressure a t about 5 5 O (water-bath temperature) and the residue dissolved in boiling alcohol and filtered.The red alcoholic extract, after being evaporated to a small bulk was poured into water the mixture extracted several times with ether and the small quanti-ties of alcohol and ether present were removed from the aqueous liquid by prolonged heating on a boiling-water bath. The solu-tion on cooling became semi-solid owing to the separation of crystals ; these were collected and washed repeatedly with ether and dilute alcohol. The deep red product obtained in this way was purified by several crystallisations from dilute and finally absolute alcohol. The air-dry substance was dried a t 130° for analysis without apparent loss of weight. The total amount of dryophantin thus obtained corresponded with about 4 grams and there was no apparent difference if fresh or old material (about six months old) was used which showed that there was apparently no deterioration on keeping.Found C=50*2 50.4; H=5*5 5.4. Dryophcamtin was obtained in deep red glistening needles with a bronzy lustre. It was almost insoluble in cold alcohol sparingly soluble in hot water but fairly readily so in boiling methyl and ethyl alcohol and in larger quantities of boiling acetone. It sinbred a t 216O and melted a t 219-220° to a viscous liquid. The C2sH,0, requires C = 50.6 ; H = 5.2 per cent THE RED PEA GALL. 1331 addition of ferric chloride to its alcoholic solution produced a brick-red precipitate and a similar precipitate but slightly darker in colour was also obtained by the addition of lead acetate. A trace of ammonia turned the alcoholic solution deep blue which became red on acidification.These colour changes could be pro-duced in an unlimited number of times in the same solution with-out affecting its sensitiveness to these reagents. Similar blue solu-tions were also obtained by the addition of sodium potassium or barium hydroxides to alcoholic solutions of dryophantin. In this connexion it must be mentioned that similar colour changes are also given by purpurogallin itself (compare Wichelhaus Ber., 1872 5 848; Struve Annalen 1872 163 164; Hooker Ber., 1887 20 3259). On repeating these observations it was found, however that the colour changes are not so permanent in the case of purpurogallin as in the case of dryophantin. Hydrolysis of Dryophan t i i z . Experiments having shown that dryophantin was a glucoside, its decomposition with acid was studied in the following manner.0.5246 Gram dissolved in 550 C.C. of boiling water was digested with 5 C.C. of sulphuric acid for two hours. A deep red crystal-line product commenced to separate and more of it was deposited on cooling. This was collected in a Gooch crucible washed with cold water so as to remove all traces of sulphuric acid and dried a t 160O. I n this way 0.1928 gram of purpurogallin was obtained. Found Purpurogallin = 36.7. C,H,,O, requires purpurogallin = 40.4 per cent. The low value obtained for purpurogallin is due to its sparing solubility in water and i t was found that the filtrate recovered on hydrolysis of dryophantin to which had been added the washings of purpurogallin contained 2.2 per cent.of purpurogallin when determined colorimetrically by Willstiitter and Stoll’s method for the estimation of small amounts of purpurogallin (Annulem 1918, 416 46). The total amount of purpurogallin from dryophantin corresponded therefore with 38.9 per cent. which is 1.5 per cent. below the theoretical if the hydrolysis of dryophantin is expressed as : C,,H,801 + 2H,O = C,,H,O + 2C6Hlz06. A second experiment gave 37.1 per cent. of purpurogallin grnvimetrically and 1 - 9 per cent. colorimetrically corresponding with 39.0 per cent. of purpurogallin which is 1.4 per cent. below the theoretical 1332 MCBAIN AND KAM: The purpurogallin recovered from dryophantin was recrystallised from glacial acetic acid and had the correct melting point of 274-275O generally given for purpurogallin (Found C = 59.8 ; H=3*7.Calc. C=60.0; R=3*6 per cent.). The acetyl deriv-ative which had been prepared by digesting with acetic anhydride, cryatalhed from alcohol in orange-yellow ’ needles melting at 179-180° and the melting point was not depressed after mixing with the tetra-acetyl derivative of purpurogallin (Found : C=59*2; H=4.7. The filtrate from the first hydrolysis was quantitatively tested for dextrose by Fischer and Freudenberg’s method (Ber. 1912, 46 915) and the dextrose estimated volumetrically in several portions of the hydrolysate by Bertrand’s method (Bull. SOC. chim., 1906 [iii] 35 1286) as used by Geake and Nierenstein (Ber., 1914 47 893) for the estimation of dextrose in gallotannin. Calc. C=58.7; H=4*1 per cent.). Found Dextrose=62.8 63.1 63.0. C,,H,80, requires dextrose = 63.3 per cent. The filtrate of the second hydrolysis was prepared as in the experiment for the quantitative estimation of dextrose and then concentrahd to a small bulk. It was subsequently converted into dextrosazone which crystallised from dilute alcohol in glistening, yellow needles melting a t 203-204O (Found N = 15.8. Calc. : N=15*6 per cent.). The author begs to acknowledge his indebtedness to the Govern-ment Grant Committee of the Royal Society for a grant from which much of the cost of the investigation was defrayed. BI o- CHEMICAL LABORATORY, CHEMIOAL DEPARTMENT, UNIVERSITY OF BRISTOL. [Received October 6th 1919.

ISSN:0368-1645    

DOI:10.1039/CT9191501328

出版商:RSC    

年代:1919

数据来源: RSC

摘要:

1332 MCBAIN AND KAM: CXXIV.-The Efect of Salts on the Vapour Presswe and Degree of Dissociation of Acetic Acid in Solution. An Experimental Refutation of the Hypothesis that Neuti-al Salts Increase the Dissocia-tion Constants of Weak Acids and Bases. By JAMES WILLIAM MCBAIN and JAMES EAM. IR 1899 Arrhenius proposed a modification of his classical dis-sociation theory to the effect that salts increase the disaociatio THE EFFECT OF SALTS ON THE VAPOUR PRESSURE ETC. 1833 constants of weak acids present with them in solution as if either the water had acquired greater dissociating power or the salt itself was acting as a dissociating medium. The exppimental evidence he adduced was the rate of inversion of sucrose by weak acids in the presence of salts. Arrhenius’s idea was very generally accepted and developed, particularly in the field of non-aqueous solutions; but in 1914 it was called in question by McBain and Coleman’s re-interpretation of the direct experimental evidence.On recalculation of the data given by Arrhenius in conformity with presenbday conceptions of this reaction they found that the supposed effect was entirely absent thus reversing the significance of the experiments. In other words the dissociation constant of weak acids is not affected by the presence of salts. They followed this up by a review of all the available experi-mental evidence bearing on this subject and they found that it supported only the simple form of the classical dissociation theory. One isolated group of experiments was left outstanding inas-much as in this particular case the measurements were conflicting in their evidence ; these were certain determinations of hydrogen ions by the method of electromotive force.The potential of the hydrogen electrode in solutions of acetic acid was greater than that predicted when sodium chloride was present although this was not the case when sodium acetate was the added salt. (1) Object of t h e Present Tnvestigatkon. The present communication adduces a hitherto unsuspected but general effect of such salts as sodium chloride on undissociated acetic acid which would account for the apparent results derived from the measurements of hydrogen electrode potential. The equilibrium under discussion is HAc H*+Acl, where HAc stands for a weak acid such as acetic acid.The potential of the hydrogen electrode in this solution is admittedly boo great when sodium chloride is present. This has hitherto been interpreted as a real increase in acidity; in other words a displacement of the equilibrium to the right-an enhancement of the dissociation constant itself. The electrical potential of the hydrogen electrode however, measures the product of the chemical potential and the concen-tration of the hydrogen ion. Instead of assuming that the increase in this product is due to increase in concentration we here submi 1334 MCBAXN AND KAM: experimental evidence for the alternative explanation that the other factor the chemical potential has been enhanced. If the concentration of the hydrogen ion has remained unaltered, but its chemical potential or reactivity has been increased it is necessary for the continuance of equilibrium that the chemical potential or reactivity of the substance on the left-hand side of the chemical equation should likewise have increased.Such increase in reactivity or potential of undissociated acetic acid in the solution must be accompanied by a parallel increase in the partial pressure of acetic acid in the vapour phase. This is open t o direct experimental test and we find that a remarkable increase is actually exhibited fully accounting for the electromotive force data observed. This removes the last evidence in favour of Arrhenius’s proposed modification of his classical dis-sociation theory. The effect here discovered has to be taken into account in most determinations of electromotive force.(2) The Experimental Method. The simple experimental method adopted was the distillation of aqueous solutions of acetic acid with and without addition of various salts. Pipettes burettes and measuring flasks were carefully Cali-brated. The distillations were carried out in a flask of fused silica of about 1500 C.C. capacity heated directly by contact with a large Bunsen flame. The neck and upper half of the flask were covered with a lagging of magnesia and asbestos so as to avoid as much as possible fractional distillation in the flask. The d i s tance between burner and flask was kept constant through all distillations as was also the flame itself. The quantity of distilling liquid mas each time 1000 c.c.and the distillate was collected in four to five fractions of about 75 C.C. each. After each fraction the distillation was quickly inter-rupted for the abstraction of a similar quantity of about 75 C.C. from the residue in the flask; the first residue was dbstracted as soon as the liquid began to boil and just previous to the collection of the first fraction of the distillate. The residues were left to cool in glass-stoppered bottles vaseline being used on the stoppers to prevent ingress of carbon dioxide from the air. Samples of 20 C.C. of each of the distillates and residues were titrated against standard solutions of sodium hydroxide of approximately equal strength with phenolphthalein as indicator ; the usual precautions were taken to avoid the vitiating effmij of carbonic acid.Thus for each distillate the ratio R2 could b THE EFFECT OF SALTS ON THE VAFOUR PRESSURE ETC. 1336 determined between the concentrations of acetic acid in that dis-tillate and the mean concentration of acid in the residue in the flask before and after. Thus the ratio R was determined for solutions of acetic acid of concentrations varying from 0.05 to 0.5N. It appeared that, a t least between these limits R increases only very slightly with the concentration as is apparent from the curve No. I of Fig. 1, which shows the value of R plotted against the acid concen-trations, The same operations were repeated with solutions of sodium, potassium and lithium chlorides potassium thiocyanate potassium nitrate sodium sulphate and sodium acetate in 0.05 to 0.4N-solu-tions of acetic acid.The concentration of the salts was deter-mined by careful evaporation of 20 C.C. of each residue in a porce-lain evaporating dish in a hot-air oven a t temperatures depending on the nature of the salt in question. The acetic acid was titrated as before in distillate and residue to obtain the ratio R of the concentration. The values of the ratios R are dependent on the concentration of the salts but they are independent of the concentration of acetic acid. Sodium acetate differs from the other salts investi-gated in that it has scarcely any effect on the distillation of acetic acid. In this manner values of R derived from a very large number of distillations of solutions of pure acetic acid became the standard of comparison for a number of distillations of acetic acid contain-ing added salts.(R1 - R') gives the percentage increase of The expression 100 the ratio R caused by the addition of salt. The experimental data here presented comprise well above one hundred such determinations. R2 (3) Method of Calculation of the Distillation Data. Distillation was selected for the measurement of the partial oapour pressure of acetic acid merely for the sake of convenience and accuracy. It must be borne in mind that the composition of the distillate shows only the relative proportion of acetic acid and water vapour in the vapour phase above the solution. What is required is the absolute magnitude of the partial vapour pressure of acetic acid a t a definite temperature say looo.The polymerisa-tion of the acetic acid vapour may be neglected for the present purpose since it amounts to only a few per cent. a t these lo 1336 UCBAIN AND KAM: partial presaureg. Even this slight effect ia largely eliminated in comparing R with R,. Ordinary variations in barometric pressure and the concomitant alterations in boiling point have no appreciable influence on the composition of the distillate. Hence all the distillations may without error be regarded as having been in effect carried out a t loo* even where much salt has been added. In order to obtain the absolute instead of the relative magnitude of the partial pressure of acetic acid in the vapour distilling over a t looo the actual partial pressure of the water has to be evaluated.Now the partial pressure of the water which in pure water wits 760 mm. has been diminished by two effects for which allowance has to be made. The first correction may be termed the (( osmotic correction.” The vapour pressure of the solvent has been reduced in familiar fashion through the osmotic activity of the substance in solution. Hence in all cases the observed concentration of the acetic acid in the distillate must be diminished accordingly before use. This is readily done with sufficient accuracy for the present purpoee by taking the lowering of vapour pressure of the solvent to be 1.80 per cent. per mol. of total solute (ions and undissociated acid and salt). It results from the effect of the appreciable partial pressure of the acetic acid in lowering the pressure a t which the water actually distils over instead of this occurring a t a partial pressure of water vapour equal to 760 mm.Here again then in order to base the relative magnitude of the partial pressure of acetic acid on the constant value of 760 mm. for that of water vapour throughout the vola-tility correction has to be applied so as to diminish the observed concentration of the acetic acid in the distillate accordingly. This consists in the reduction of the latter by 0.12 per cent. for a 0 . W -solution of acetic acid and taking this correction as proportional to the concentration of the acetic acid in the distillate. Since the acetic acid was at most N / 2 this correction in no case exceeded 0.6 per cent. A third and final correction had tc be made this time in the apparent composition of the solution undergoing distillation.This is the ((correction for dissociation” of the acetic acid in the solu-tion. The actual ratios measured were those between the con-centrations of distillates and the corresponding solutione in the distilling flask (the (I residues ”). What is required is the ratio based on the actual concentration of undissociated acetic acid in the flask. In the case of solutions containing only acetic acid. The second correction is the ‘( volatility correction. THE EFFECT OF SALTS ON TEB) VAPOUB PRESSURE ETC. 1337 this consisted simply in subtracting the known amount of dis-sociated acetic acid from the total concentration of acetic acid in the flask. The dissociation constank of acetic acid was taken to be 1.11 x 10-5 a t looo.The correction for dissociation involves much calculation where salt is present since in order to determine the actual concentration of undissociated acetic acid it is necessary to calculate the amounb of the various ions and undissociated Salk present including those formed by metathesis. For example with common salt the follow-ing molecular species were present H’ Act Na’ Cl’ HAc HCI, NaC1 and NaAc. This was done by Arrhenius’s method which is based on the principle of isohydrism (loc. cit.) and agrees with the method of Sherrill ( J . Amer. Chem. SOC. 1910 32 741). The calculation is laborious and involves successive approximations Fortunately, the exact degree of dissociation of the various salts has but little influence on the results since the really important values appear in both numerator and denominator of Arrhenius’s equations ; con- I ductivity data a t the ordinary temperature could theref ore be employed failing the existence of such a t higher temperatures and concentrations.Indeed the calculation for solutions of one salt in acetic acid might have been applied to the case of any other salt of the same concentrations except of course in the case of sodium acetate. The effect is chiefly dependent on the relative concentration of acid and added salts. ( 4 ) The Distillation of Solutions of Pure Acetic Acid. Following the method already described fifteen distillations were carried out with N / 20- to A7 / 2-solutions of acetic acid involving more than seventy determinations.The object in view was to determine the ratio R, that is the ratio concentration of acid in distillate t o concentration of acid in residue. C c r Thus R,= 2, Cd being the concentration of the distillate C being the mean value of the concentrations of the residue before and after the separation of the distillate. For any one concentration the results agreed to within 1 per cent. The values of R were corrected as described in Section (3) above and they increase by only about 1 per cent. over the whole range of concentration. The relative concentration of double molecules, which has not been allowed for changes from about 5 to 8 per cent. over this same range. The experimental evidence is summarised in table I 1338 MoBAIN AND ?CAM: TABLE I.(Curve I.) Ratio of Concentration i n Distillate and Residues in Aqueous Acetic Acid. Concentration of acid in Rz R2 flask. (uncorrected). (corrected). N/20 0.662 0.67 1 N/10 0.666 0.67 1 0.673 0.674 0.680 0.677 0-686 0.678 N15 N /3 N12 The corrected values of R increase very slowly indeed. Plotting them against the concentrations we obtain a straight line (Curve I in Fig. 1). FIG. 1. I. CH,*CO,H COW. (B=0*6S!; C=0.678). 11. NaCl uncorr. Percentage ancreme due to salt. 111. NaCl corn. 9 ) ?1 9 ) 9 ) IV. KCI ? 9 9 * 1 1 V. KCNS , 9 9 I 9 1 9 ?1 VIA.Czcrue logcJ = [{1-BR,).log3.2-3026 + R . K . Co 2.3026 01 A0 = v,,- 1. Ratios of acetic acid in distillate and reeidue with and without ao?ded salts. The corrections for volatility and osmotic effect can be taken from graphs; for example in the case of N/20-acetic acid the THE EFFECT OF SALTS ON THE VAPOUR PRESSURE ETC.1339 amount to -0.06 and -0.09 per cent. respectively. The third correction for degree of dissociation of the acetic acid is here 1.50 per cent. Thus the total correction in this case is an increase of 1'35 per cent. on the observed ratio R,=0.662. Hence we obtain This value is smaller than the one arrived a t by Lord Rayleigh (Phil. Mag. 1902 [vi] 4 535) R2=0.73. Corrected it reduces sligh€ly to 0.725 which is still considerably greater (the reduction in this case being caused by the method of calculation of R from the residues). The discrepancy is however probably due to the different manner of heating.Lord Rayleigh kept the neck and upper part of the distilling flask hotter than the boiling liquid in order to prevent condensation. We found however that this involved the quantitative evaporation of drops splashing up f roan the boiling liquid which produces the same error as if they had splashed directly into the distillate. It was for this reason that we relied on good heat insulation and fairly rapid distillation. R (corrected) = 0.662 x 1.0135 = 0.671. T o t e on the Calculation of the Composition given Stage of the Distillation of Following Lord Rayleigh (Zoc. cit.) but residues expressed in mols. per litre, where C0 and C are the concentrations and in the distilling flask before and after the For example for R,=0.671 and --O- c 1 c,- 9 of the Residues at any Acetic Acid.using concentration of: Vo and Vl the volumes distillation, 5 = 8 . 2 2 1 . Hence in v, order to double the concentration of the resiiual acid solution in 7.921 the flask -5.- =On878 or almost 88 per cent. of the volume of the 6.221 solution must be distilled over. For R = 0.73 (Rayleigh's value in Nll0-solution) this quantity would amount to 92 per cent. Again if 40 per cent. of the original volume of solution is dis- c tilled off that is if Vl=O*6 Vo,+ =1*183 which means that the concentration of the residue will have increased by 18.3 per cent. 6 0 (5) The Experimental Data for Added Salts. Having determined the ratios R for solutions of acetic acid the corresponding ratios R, were obtained in exactly the same manner, but with a salt added to the solution.VOL. cxv. 3 I! 340 MCBAIN AND KAM: The salts used were the purest obtainable in 1913 and are named in Section (2) (above). The distillates were in each case tested for traces of the stronger acid formed by metathesis but these were found to be negligible. The corrections were calculated and applied in the same manner as before except that the acetic acid destroyed by metathesis was calculated according to Arrhenius's principle of isohydrism (Zeitsch. physikal. Chem. 1899 30 208). The results are summarised in tables eII-VIII which require no further explanation except to note that in tables I1 and VTI space is saved by averaging the figures for all the values obtained over certain ranges of concentration.The number of experimental values so averaged is given in the last column. With the exception of sodium acetate all these salts cause a remarkable increase of the ratio cd and sodium chloride shows this increase more than any other of the salts investigated. Sodium acetate on the other hand appears to have no appreciable effect even in l.0N-solution. TABLES I1 .-VIII. Ratios R of Concentration of Distillates and Residues of Apue0u.e Acetic Acid with added Salts and Increase thereof over R, of Table I . c,' TABLE 11. (Curves 11. and 111.). Sodium Chloride. HAc. 0-1-0*2 0.1-0.3 0.1-0.3 0.1-0.3 0.13 0.076 0.078 0.077 Sdt. (0.12) (0.21) (0-29) (0-41) 0.661 1-04 1.38 2-30 RP 0.688 0.708 0.721 0.742 0.790 0.830 0,904 1-076 R2-0.669 0.67 1 0-672 0.67 1 0-668 0.663 0.664 0.664 Rl (corr.). 0.690 0.7 10 0.716 0.721 0.715 0-820 0.885 1.033 R2 %-3 (corr.). [ R ] Expts. 0.673 0.673 0.673 0.673 0.67 1 0.671 0.67 1 0.671 2.6 ~ 7 5.4 16 6.6 15 9.6 4 16.9 1 22.2 1 32.1 1 53.9 1 TABLE 111. (Curve IV. Fig. 1). Potassium Chloride. 0.22 0.2-0.3 0.7146 0.673 0.7111 0.674 5.55 4 0.2540 0.4673 0.7389 0.675 0.7309 0.676 8.20 1 0.2157 (0.70) 0.7741 0.673 0.759 0-674 13-55 2 0.2053 0.9228 0-8121 0-673 0.7930 0.674 17.70 1 0.22 1.17 0.8383 0-673 0.8133 0.674 21.55 2 0.2194 1.953 0.9363 0.673 0.8872 0.674 31.60 1 TABLE IV. (Curve V. Fig. 1). Potassium Thiocyanate. 0.2-0.23 0.2-0.34 0.6900 0.673 0.6868 0.674 1.90 6 0.2673 0.6590 0.7297 0.675 0.7181 0.675 6.40 1 0.2366 1-663 0.7693 0.674 0.7343 0-674 8.90 1 0.2596 2.451 0.8710 0.675 0.8112 0.676 20-10 THE EFFEUT OF SALTS ON THB BAkpOUR PRESSURE ETC.1341 HAc. 0.22 0.22 0.25 0.25 0.27 0.29 0-21 0.22 TABLE V. (Curve VI. Fig. 2). Sodium Sulphtc. S d t . (0.07) (0.13) (0.22) (0-35) (0-61) 1-10 1.40 1.64 RIB 0.669 0.674 0.674 0.683 0.7 10 0.762 0.807 0.839 R,* 0.673 0.673 0.675 0.675 0.677 0.678 0.673 0.673 Rl. (COAT.). 0.660 0.673 0.67 1 0.678 0.700 0.743 0.786 0.813 R,. (con.). 0-674 0.674 0.674 0,674 0.675 0.675 0.674 0.674 - 0.2 8 - 0.4 7 + 0.5 6 + 3.8 2 + 10.0 1 + 16.6 1 + 20.6 1 TABLE VI. (Curve VII. Fig. 2). Lithizcm Chloride.0-2-0.22 0.2-0.25 0.7143 0.673 0.7116 0,674 5.57 3 0.2269 0.2799 0.7188 0.673 0.7147 0-674 6.00 1 0,2469 0.3930 0.7450 0.674 0.7375 0.674 9.40 1 FIG. 2. .i % VI. Na,,SO, cow. VIII. D O s cow. VII. LiCl , IX. NaCH,'COB GOW. Percentage invreaae in ratios of acetic acid in diabillate and residue due to added s&. TABLE VTI. (Curve VIII. Fig. 2). Potassium Nitrate. 0.21 0.22 0.6944 0.673 0-6940 0.674 1.13 2 0.24 0.3-0.5 0.7198 0.674 0.7110 0.674 2-32 2 TABLE VIII. (Curve IX. Fig. 2). Sodium Acetate. 0.11 (0.14) 0.6822 0.667 0.6769 0.671 0.8 3 0*1-0*23 0.4-0*67 0.6948 0.671 0.6799 0.673 0.9 7 0.14.28 0*9-1*1 0.7001 0.673 0,6799 0.673 1.3 3 0.2900 1.810 0.7244 0.678 0.6834 0.675 1-3 1 The increase of R, the relative concentration of the distillate, for the same sollation of aoetic acid that is the expression 3 ~ 1342 M c B m AND KAM: 100 1- amounted to no less than 62 per cent.(observed) (” isR2) in the c a k of 2*3N-sodium chloride. The increase for most salt8 seems proportional to the concentra-tion of the salt and independent of the concentration of acetic acid. Curves 11-V of Fig. 1 and VI of Fig. 2 show this percentage increase plotted against the concentrations and they point con-vincingly towards a straight line function between these two values. Sodium aulphate differs from all the other salts in that the experi-mental evidence shows a slightly negative effect a t lower concen-trations up to about 0.3N. Note on the Calculation. of Residues in the Presence of Salts which exhibit a Straight Line Function of the Ratio Increase.The calculation is similar to the one for the pure acid solution, but we must introduce the functional relation of R with regard to the concentration of the salt. We have found experimentally for all salts except the sulphate and acetate 100(Bd2) = lOOKp or R =R2(Kp + I) 4 in which p is the concentration of the salt and K is a constant. Hence R,=R2 - + 1 where V is the volume in litres containing 1 mol. of salt. I f there are y mols. of acetic acid in V litres of the solution in the flask and a quantity dv of the solution containing dy of the acetic acid distils over we may set the concentration of the dis-(; ) tillate equal to dY - d V = RIP= Y R2(; + 1);. Since c = 3 and by v‘ Integrating, En 5 = R*X (F) + (R - qzrt 5, C1 v Vl Vl or where Co and C are the concentrations and Vo and V the volumes: in the distilling flask before and after distillation THE EFFEUT OF SALTS ON THE VAPOUR PRESSURE ETC.1343 An alternative formula deducible in the same way is Either of these expressions can be used to calculate the con-centration of the acetic acid in the distilling flask if the initial value of the salt concentration and either its increase or the relative volume of the residual solution are given. They contain two constants. lOOR is the one which is peculiar to the added salt and it is simply the percentage increase by 1-ON-salt. The other R, is the ratio of acetic acid in distillate and residue for the same solution in the absence of salt. It is convenient t o express VT1 as a fraction of V, the volume a t the beginning of the distillation containing 1 mol.of salt ; but the initial concentra€ion 1 of salt in mols. per litre= -. vo Inspection of the equations show that they are identical with that deduced above for solutions of pure acetic acid except for the correcting factor R,lp('"- "I) This factor of course dis-appears for large values of Vo that is for very low concentration of added salt. Conversely for very high concentrations of salt it is predominant as is evident from the consideration that for say, 2-6N-potassium chloride no separation takes place owing to the concentration of acetic acid in the distillate having been so raised as to equaI that of the residue. Above such concentrations the residue becomes weaker instead of stronger.The peneral behaviour of a distillation is shown graphically in curve VIA of Fig. 1 which assumes that 1.0N-sodium chloride was initiallv present ( Vo= 1). AS the distillation proceeds the salt accumulates and the concentration of acetic acid in the residue slowly rises to a maximum where 50 per cent. of the liquid has C distilled over (V,=0*5). At this point l o g 2 = 0.03 whence the increase of concentration is 7-2 per cent. At V,=0.27 when 73 per cent. has been distilled over the concentration of the acid is again the same as it was before the ,distillation. Beyond this point the value for log CJ assumes rapidly negative values, owing to the high concentration of salt in the residue. The general equation may be tested by one of our experimental resulte.Taking the last pair of values in table I11 for potassium chloride 1 *ON-pota8sium chloride increases the ratio of acetic acid by 18.5 per cent. hence K=0.185. Further R,=0.674, v0.v1 Q, C 1344 McBAIN AND KAM: 1.128 V,, Vo=-. 1 These values inserted in the equation 1.1 28 v1 = - 1.953 lead to the prediction that C’,=1.O78C0. Experimentally Co was 0.2118N (less 1 per cent. for metathesis) and C after distillation 0.2194N (plus 5 per cent. for the three corrections). Hence C = -Go = 1.0950,. In general the concentrations observed appear to agree with the predicted values within about 2 per cent. To sum up the process of concentration or separation of the constituents of a binary mixture by means of distillation may thus be considerably accelerated or retarded by the addition of a salt, and will largely depend on the values of the constants R and R, that is on the nature of the mixture and of the added salt.0.230 0*210 (6) Discussion of the Results. The remarkable effed of a salt on the partial vapour pressure of acetic acid must be evident from the preceding. Comparing the slopes of the various graphs showing the relationship between salt concentration and percentage ratio increase it; appears that the effect is greatest for the chlorides of lithium sodium and potassium and least for sodium acetate. The effect is in all cases independent of the concentration of the acetic acid. The increase of the concentration of the residue in the flask for acetic acid in aqueous solution during the distillation not very rapid in itself is still less if salt is also present.At a concentra-tion of 2-3N-sodium chloride the residue becomes weaker ; in other words the vapour phase in the flask contains more acetic acid than the liquid from which it originates. For sodium sulphate up to about 0*35N the experimental evidence for the ratio increase seems somewhat complicated. I f anything there is a negative effect as the course of curve VI in Fig. 2 indicates. Beyond 0*35N the effect is decidedly positive. Between O.35N and l-lN the graph is practically a srtraicht line, but beyond l * l N its slope appears to increase until at 1*55N the limit of solubility is approached. Sodium acetate up to 1.8N shows only a very slight effect;. Although the effect of the cation is undeniable (note for example, the greater slope of the ratio increase for sodium chloride as com-pared with the one for potassium chloride) the influence of the anion seems t o be the predominating factor.The series sodium acetate potassium nitrate potassium thiocyanate sodium eulphate, pofsssium chloride lithium chloride sodium chloride shows the inaresse in a progressive degree. It is evident that this order is uot that of the Hofrneister or lyotropic series THE EFFECT OF SALTS ON THE VAPOUR PRESSURE ETC. 1345 It is well known that there is a general qualitative similarity between the effect of neutral salts on such various phenomena as solubilities of gases and non-electrolytes surf ace tension compressi-bility maximum density of water viscosity dielectric constant, imbibition and gelatinisation of gels and increase or decrease of rate of catalysis.The explanation of this undoubted parallelism is wholly unknown and in each individual case there are pro-nounced exceptions. I n the present instance the exceptions are the acetate and the sulphate although they may be paralleled by certain cases of catalysis. Possibly in the case of the sulphate our correction for metathesis (formation of HSO,’) has not been suffici-ently great. The effects are reconcilable with a solvate form of the dissociation theory. With regard t o the main thesis of this paper it has now been proved that the reactivity of the undissociated acetic acid is increased by addition of such salts as sodium chloride. A 0-2N-solution is affected to the extent of 5.5 per cent.by 0.2N-sodium chloride. It was pointed out in the earlier paper (Zoc. cit.) that Walpole’s measurements of electromotive force for this particular case gave a result for the hydrogen ion which was 7 or 8 per cent. too high. These two effects 5.5 and 7 or 8 per cenk. are equal within the experimental error and thus the effect on the dissocia-tion constant deduced for acetic acid cancels out and leaves that constant unchanged by the presence of the salt. Thus the electromotive force data may be regarded as agreeing with all the other data bearinq on this subject i n d the experi-mental evidence all points to the conclusion that the dissociation constants of weak substances are not appreciably affected by the addition of salts.One point we have not investigated namely the effect of salts on the chemical potential of the acetate ion. Summary. (1) It is shown experimentally that many salts enhance the partial vapour pressure of acetic acid in aqueous solution by very appreciable amounts. I n the case of 2-3N-sodium chloride the increase amounts to no less than 62 per cent. (2) Since this partial pressure is a measure of the reactivity of the undissociated acid in the solution. the undissociated acid must be regarded as exhibiting enhanced chemical potential in the presence of such salts. ‘This is parallel with the available data for the effect of such salh on the measurement of hydrogen ion by electromotive force. The enhancement is thus discovered to be operative on both sides of the chemical equation and hence to leave the dissociation constant of acetic acid sensibly unaltered 1346 RIVETT AND O’CONNOR SOME TERNARY These experiments remove the only remaining evidence (apart from the ambiguous behaviour of certain insufficiently investigated colloids) for the view that salts might have been regarded as increasing the strength of weak acids. (3) Whereas a number of salts increase the partial pressure of acetic acid to an extent proportional to the concentration sodium sulphate exhibits a more complicated behaviour whilst sodium acetate has only a very slight effect. All electromotive force data on weak acids in the presence of salts other than sodium acetate require to be corrected for the effeck here described. THE CHEMIOAL DEPARTMENT, UNIVERSITY OF BRISTOL. [Received September 17th 1919.

ISSN:0368-1645    

DOI:10.1039/CT9191501332

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年代:1919

数据来源: RSC

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1346 RIVETT AND O’CONNOR SOME TERNARY CXXV.-Some Ternar,y Systems containing Alkali Oxalates and Water. By ALBERT CHERBURY DAVID RIVETT and EDMUND ARTHUR O’CONNOR. THERE has been considerable discussion a t various times regarding the alleged formation of certain double oxalates of the alkali metals. Wenzel is quoted by some early writers as maintaining the existence of t.he double salts K,C204,Na,C,04 and but Rammelsberg ( A I ~ . Pltys. Chem. 1850 [ii] 79 662) has thrown doubt on the one case and Souchay and Lenssen (,4nnalen, 1856 99 31) on the other. Foote and Andrew (Amer. Chent. J. 1905 34 164) drew the same conclusions as Rammelsberg and Souchay and Lenssen and claimed to have shown that these double oxalates do not exist a t 25O in the solid state. They state that the solid monohydrates of potassium and ammonium oxalates remain in equilibrium with a common saturated solution and that the same holds for anhydrous sodium and hydrated potassium oxalates.A few years later, Barbier (Bull. SOC. chim. 1908 [iv] 3 725) described a double oxalate of potassium and ammonium stating that it might readily be formed by adding a concentrated solution of potassium oxalate to a saturated solution of ammonium carbonate. The analysis he quotes points to a pure 1 :I anhydrous double salt. w 2 0 4 (NH4)2C204 SYSTEMS CONTAINIETQ ALKALI OXALATBS ASD WATER 1347 None of the authors mentioned made complete investigationa by the solubility method. Pooh and Andrew (Zoc. c&.) obtained a few solubility figures but to determine the solid phases present they relied on a general principle laid down in a previous paper (ibid.p. 153) that when two salts are mixed in varying propor-tions and treated with water a t a constant temperature the resi-due of undissolved solid remains constant in composition and the solution varies if a pure double salt is present whilst on the other hand the residue varies and the solution remains constant in composition when a mere mixture of the two single salts is present. This rule is liable to mislead if only because it does not take into account the possibility of the formation of mixed crystals (solid solutions). It has seemed worth while therefore to apply the solubility method more fully to some of these ternary systems of alkali oxalates in water. Knowledge of the solid phases present has been obtained by the customary graphic method of plotting in a triangular diagram the percentage compositions of pure solution and of the moist solid (or “residue ”) in equilibrium with it and extrapolating the straight line joining the two poinb to the com-position of the pure solid uncontaminated with adhering solution.The systems investigated are those containing potassium sodium, and ammonium oxalates in pairs with water. The solid phases of the individual salts stable a t these temperatures with their own aqueous solutions are respectively K,C,O,,H,O (NH,),C,O4,H20, and N%C20,. EXPERIMENTAL. I. System K,C,O,-(NH,),C,O,-H,O at 2 5 O and 50°. The system potassium oxalate ammonium oxalate and water has been examined at 2 5 O and 50° and the figures obtained are given in tables I and 11 respectively and plotted in Fig.1. $uitable mixtures of the salts (monohydrates) and water were heated in bottles which were sealed and placed in a thermostat, in which they were continuously rotated for about forty-eight hours. Undissolved solid was allowed to settle and clear solution drawn into a pipette through a small plug of cotton wool held in rubber tubing. A known weight was diluted to a suitable volume for subsequent analyses. Residues were obtained by pouring solu-tion and suspended solid on to a Buchner funnel drawing the solution through rapidly by means of a pump but disconnecting the pump before more than a very small amount of air had been 3 E 1348 RIVETT AND O'CONNOB SOME TIBRHARY drawn through the moist solid.A slight loss of waterr vapour a t 50° ia inevitable but with rapid working it can be made almost negligible. Total oxalate was determined by titration of a fraction of the stock aolution with standard potassium permanganate ammonium by distillation with alkali and absorption of ammonia in standard acid and potassium by difference. Concentrationa have been ex-pressed in percentages by weight but as densities have been deter-mined in all cases figures for concentrations in other terms are readily obtainable. TDLE I. 250. No. 1 2 R 2 3 R 3 4 R 4 6 R 5 6 7 8 9 R 9 10 R 10 11 R 11 12 R 12 13 14 16 16 Density. 1.021 1.040 -1.058 1.068 1.087 1.1 07 1.124 1.128 1.137 --1.166 1.185 1.204 1.217 ---1.216 1.216 1.216 Percentage composition of solution or residue (R).K2C2OI. -2.67 0.44 4.32 2-41 6.6 1 2-75 9-48 2.62 12.10 14.18 16.37 16.54 2.14 19-39 4-37 21.9 6.70 24.3 13.4 26.9 26.3 26.8 27-2 (NH*),CZO'~ 5-01 4-72 81.3 66.4 61.1 72-3 4.48 4-38 4-16 4.01 3-78 3.68 3-67 3.32 3-10 2.90 2.76 80.4 '76-4 76.2 61.9 1.83 0.85 c H20. 96.0 92.6 18.3 91.2 41.2 89-1 36.1 86.4 26.1 83.9 81.0 80.9 79-9 17.2 77.3 19.2 78.0 19.2 72.8 24.7 7 1.3 71.0 72-3 72.8 Solid phases. (NH,),C*O,H,O. Solid solution o SYSTENS UONTBZNINQ ALKALI OXALATES AND WATER. 1349 TABLE 11.50°. NO. 1 2 3 R 3 4 R 4 6 R 6 6 R 6 c 8 9 R 9 10 Percentage composition of solution or residue (R). h Density. K4C20,. (NH4),C,04. q0. 1.034 - 9.63 90.4 1.080 7.99 8-44 834 1.136 1-164 1.187 1403 1.254 ---16.20 4.20 17.99 6-72 22.4 13.3 24.4 16.2 30.4 7.10 6-79 6-10 6.76 4-78 70.0 66.7 55.9 66.3 76.7 26.8 76.2 27-6 7 1.6 30.8 69.8 28-5 64-8 1.261 31.0 3.34 65-7 1.262 31.6 2-64 65.9 7 76.6 0.47 24-0 1.252 33-1 - 66.9 FIU. 1. With the exception of some of the reaidues these rwults are plotted in Fig. 1 and show distinctly that only two solubility 3 E" 1350 RIYETT AND O’OONNOR SOM’E TERNARY curves are obtainable a t each of these temperatures the two meet-ing sharply a t a quadruple (or condensed triple) point.There is no evidence a t all of the existence of a double salt. On the other hand it is quite apparent from the relations between compositions of solutions and corresponding residues (shown only for 50°) that the solid phase present in those complexes in which excess of ammonium oxalate is taken (curve CB) is not this pure solid but contains in addition some potassium oxalate. The pro-portions of the two in the solid vary according to the composition of the solution in equilibrium and the evidence is definite that mixed crystals of the two salts are produced. The more potassium oxalate there is in the solution the more there is in the solid. The same must hold with regard to the solids in equilibrium along the curve A B .These solids will be mainly potassium oxalate with steadily increasing proportions of ammonium oxalate. The amounts of ammonium oxalate in solutions along this curve are, however so small that a slight error in the analysis of the residue may make the extrapolation method uncertain in showing the solid phase. Some of the mixed crystals were dried by draining on a porous tile in a closed $essel immersed in the thermostat. Analysis proved the two constituents of these mixed crystals to be the respective monohydrahs. The form of the isotherms with a sharp change of direction a t B shows that two distinct solid phases must be present at this a univariant point. It follows that the series of mixed crystals is not complete but that there is a limit to the solubility of each solid in the other.The compositions are also not given by such points as B and E in Fig. 1 those of the exact solid solutions in equilibrium with the corresponding liquid solutions a t the other ends of the tie-lines. Some of the crystals may have had a core of the pure maih con-stituent so that the mean composition represented by D and E may be low in potassium oxalate as compared with the solid solu-tion itself. All that is established is the existence of mixed er ystals . The work of Foote and Andrew (Zoc. cit.) is not quite extended enough to show this. They obtained only the points A B and C a t 25O. It has been found impossible t o repeat the work of Barbier (Zoc. cit.) which pointed to the formation of an anhydroua double salt IC,CzO4,(NH4),C20,.By following closely the method which he desoribed for isolating the compound crystals were precipitated which after drying on a porous tile contained 81.1 per cent. of These limits have not been determined SYSTEMS (JONTAININB ALKALI OXALATES AND WATER. 1351 ammonium oxalate 5'8 per cent. of potassium oxalate and (by difference) 13.1 per cent. of water. After washing with a solution of ammonia as recommended by Barbier the proportion of potassium oxalate decreased slightly. It appears certain that this precipitate consists of mixed crystals of the two hydrates. It is of interest to note that Souchay and Lenssen (Zoc. cit.), following Wenzel's instructions for preparing the alleged double salt obtained crystals which they stated to be ammonium oxalate, but containing 0.9 per cent.potash. 11. System Na,C20,-(NH,),C,0,-H,0 at 2 5 O and 50". Sodium oxalate differs from both the potassium and ammonium salts in crystallising anhydrous from aqueous solution; hence it is less likely that mixed crystals will be formed between it and either of the other two. The figures in tables I11 and IV show that neither double salt nor mixed crystals occur in the sodium-ammonium system a t 2 5 O or 50°. The method of analysis was similar to that adopted for the previous system. TABLE III. 26O. Percentage composition of solution or residue (R). No. 1 2 3 R 3 4 6 6 7 8 9 R 9 10 11 Density. 1.027 1.030 1.033 1.037 1 to39 1.043 1.047 1.043 1.035 1.028 1.021 --Na,C,04.3.73 3-69 3-65 3-50 3.51 3.46 3.41 2.86 1-82 0.40 0.89 63.9 -(m4)2cZ04. H 2 0 . - 96.3 0.74 95.6 1.49 94.6 0.53 35.6 2-48 94.0 2-89 93.6 3-77 92.8 4.74 91.8 4-76 92.4 4.81 93-4 66.9 32.7 4.88 94.2 6-01 96.0 Solid phases. Na,C,O, 1352 NO. 1 2 3 4 R 4 5 6 R 6 7 8 R 8 9 10 11 Density. 1.023 1*M1 1.036 1-044 1.049 1.066 1.063 1.059 1.051 1,042 1-034 --I TABLE IV. 50°. Percentage oomposition of solution or reeidue (R). r A \ Na,C,O,. (NH&C,Op H2O. 96-6 4-54 -4-46 1.59 94.0 4.37 3.14 92.5 4.28 4.64 91.1 70.0 1-47 28.5 4-29 6.12 89.6 4-13 7.86 $8.0 66.7 3-64 39-8 4.05 9.19 86.8 3-57 9-21 87-2 €047 63-4 36.1 2-44 9.32 88.2 1.25 9-46 89.3 9.63 90.4 -Fra.2. Solid phases. Na,Ca04. The results are plotted in Fig. 2 where on account of the sparing solubilities of the two components only a single angle of the triangle is shown. The residues are omitted. 111. System .- K,C20,-Na,C20,-H20 at 25O. Foote and Andrew (Zoc. c i t . ) concluded that anhydrous sodium oxalate and the monohydrate of potassium oxalate can exist sid SYSTEMS CONTAJBINQ ALKALI OXALATES AND WATER. 1363 by side in equilibrium with a common saturated solution and the figures in table V plotted in Fig. 3 confirm this showing that neither double salts nor mixed crystals are formed. Solutions and residues were analysed by determining total oxalate by titration and total anhydrous salts by weighing after evaporation and dry-ing a t 125O a t which itemperatwe potassium oxalate monohydrate is readily dehydrated.As the amount of sodium oxalate present is always relatively small this indirect method is less accurate than that employed in the previous two cases. FIG. 3. N O . 1 2 3 4 R 4 5 6 R 6 7 8 9 10 Density. 11215 1.218 1.223 1.226 1.228 1,178 1.136 1.084 1-067 1.026 -K,C;04. 27.2 26.8 26.3 26-2 81.9 26.1 19-6 14.4 7-00 8-10 3.99 -Na2C2O4. H2O. - 72.8 0.77 72.4 2.17 71.6 0.73 17.4 2.50 71.4 1-71 7,2.0 3.21 77-2 65.3 27.7 3.21 82.4 3-40 88.6 3.71 92-3 3.7 1 96.3 TAfiLi. -v. 25" Percentage composition of solution or residue (R). K,C204,7H20 and Na2C,01. Na,C,O4 1364 PRICE THE DEUOMPOSITION OF Summary. (1) It has been stated by some authors and denied by others that potassium and ammonium oxalates form a double salt. Iso-therms have been obtained a t 2 5 O and 50° and show that a t these temperatures mixed crystals of the monohydrates are formed. (2) Anhydrous sodium oxalate has been shown to exist in equil-ibrium with ammonium oxalate monohydrate and a common, saturated solution at 2 5 O and 50°. Neither double salts nor mixed crystals are formed. (3) The same has been shown to be the case at 2 5 O with anhydrous sodium oxalate and potassium oxalate monohydrate. W N I V E I C S ~ ~ ~ OF MELBOURNE. [ReeCived Nwember lat 1919.

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DOI:10.1039/CT9191501346

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年代:1919

数据来源: RSC

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1364 PRICE THE DEUOMPOSITION OF CXXV1.-The Decomposition of Carbamide in the Presence of Nitric Acid. By TUDOR WILLIAMS PRICE. WHILE investigating the use of dilute nitric acid as a nitrating agent a t elevated temperatures it was found that nitration was entirely inhibited by the addition of carbamide although all the acid was used up. This reaction was considered worthy of further investigation the results of which are given in the present com-munication. The action of concentrated nitric acid on carbamide nitrate in the solid state at the ordinary temperature has been studied by Franchimont (Rec. trav. chim. 1884 3 216) who found that a gas was given off slowly consisting of equal parts of carbon dioxide and nitrous oxide. The volume of gas was such that it contained all the carbon of the carbamide as carbon dioxide and half the nitrogen as nitrous oxide the other half forming ammonium nitrate.No account has been found in the literature of the action of dilute nitric acid on carbamide although the action of hydro-chloric acid and sulphuric acid has been studied by Fawsitt (Zeitsch. physikal. Chem. 1902 41 601) and of hydrochloric acid by Werner (T. 1918 113 84). Fawsitt (Zoc. cit.) found the reaction between carbamide and dilute hydrochloric and sulphuric acids to be unimolecular whic CARBAMIDE IN THE PBESENCE OW NITRIC ACID. 1356 he explained by the assumption that the former is first transformed to a small extent into ammonium cyanate at a meaeurable rate, and if the inverse change is neglected according to a unimolecular reaction The ammonium cyanate is thereupon decomposed into the ammonium salt of the acid used and carbon dioxide a t a much greater rate than that a t which ammonium cyanate is formed and thus the decomposition of carbamide in the presence of hot acids is almost identical with the rate a t which it is transformed into ammonium cyanate.Werner (Zoc. cit.) on the other hand states that the above assumption is quite unnecessary and that the velocity of the reaction is regulated by the rate of dissociation of carbamide (at looo) into ammonia and cyanic acid when both products of dis-sociation are removed practically as fast as they are generated. Whichever scheme of the mechanism of the reaction is correct, the final products are the same namely carbon dioxide and the ammonium salt of the acid used.If the action of hot dilute nitric acid on carbamide is the same as that of hydrochloric acid, the reaction should be unimolecular and the products should be carbon dioxide and ammonium nitrate. Veley (Proc. Roy. SOC. 1892 52 27) has shown that nitric acid decomposes slowly a t high temperatures forming nitrous acid, the weaker the acid the higher being the temperature required for decomposition; also it is well known that carbamide decom-poses nitrous acid very readily forming carbon dioxide and nitrogen. Hence it is possible that the true explanation of the disappearance of nitric acid when heated with carbamide in solu-tion is the alternate formation of nitrous acid from nitric acid, and decomposition of the nitrous acid by carbamide.In this case, nitrogen will be present in the evolved gas as well as carbon dioxide. There are thus three separate methods of decomposition of carbamide by nitric acid the gaseous products in each case being different. An analysis of the gas produced will then be of extreme importance. As will be seen in the experimental part the reaction between carbamide and nitric acid is unimolecular and the gas evolved consists entirely of carbon dioxide. E X P E R I M E N T A L . The carbamide used was ordinary commercially pure carbamide recrystallised four times from absolute alcohol. The nitric acid solutions were made from pure distilled nitric acid and containe 1356 PRICE THE DECOMPOSITION UF only a trace of nitrous acid.A normal solation of carbamide is taken to be one containing half its molecular weight in grams in a litre of water; all the other solutions of carbamide were made up on this barris. Ten C.C. of the mixed solutions of carbamide and nitric acid were placed in hard-glass test-tubes which were then sealed and placed in water a t the required temperature. After various periods of heating the tubes were withdrawn cooled opened and their contents titrated with standard sodium hydroxide solution, using methyl-orange as indicator. The majority of the experiments were made a t looo but the reaction between N / 2-csrbamide and N / 2-nitric acid was studied a t 70° 80° 90° and looo. The influence of certain salts on the velocity of the reaction was also examined. In order to obtain figures comparable with those of Fawsitt the velocity-constants for a unimolecular reaction were calculated according to the equation instead of the correct equation, a - Xl loge - k=- 1 d 8 - t a-x, where a is the initial concentration of carbamide x1 the amount of carbamide decomposed in time t, and x2 the amount decomposed in time t,.In every case t was fixed a t sixty minutes so as t o obviate the error due to the time taken for the tube to reach the desired temperature. u zl and x2 were obtained from the titre of sodium hydroxide immediately after mixing the two solutions and the titre after times t and t,. In one case the constants for a bimolecular and termolecular reaction were calculated. Results.-The results of two experiments are given in detail, and all are summarised in table I.It will be noted that the velocity-constant tends to increase towards the end of an experi-ment and this is much more marked with 1- and 2N-solution than with the more dilute ones CARBAMI.DE TI!l T3E PRlPsENUE OF NXTRIC ACID. 1357 Ezperimcat 3. N-Curbamide + N-Nitric Acid at 100°. a = 50-40. Time in minutes. 60 120 1 80 270 360 480 600 900 1600 0 - x. 46-60 40.61 36-06 29.53 24-2 1 17-84 13.36 5.18 1-21 kl x 106 k x 108 & x l(r (mimolecular). (bmolecular). (termoleoular). - - -84 46 84 84 60 101 89 66 127 91 64 163 - - - - ( 97) ( 98) (106) (109) Mertn value b x lo6 = 87. - - - -Egperiment 5. N/ 2-Carbamide + N/ 2-Nitric Acid at looo. a = 25.43.60 22-75 -120 19.81 101 180 17.25 100 270 13.96 101 360 11.39 100 480 8-29 106 600 6-36 102 900 3.00 104 1200 1-42 107 Mean value k x 10s = 102. Time in minutes. a - z. k x 106. 1500 0.89 (169) TABLE I. 2N-N-N/2-N/2-N / 4 -NIP-N/8-N/16-N/2-N/2-N/2-N/2-N/2-N/2-N/2-Showing Vetocity-constants obtained in the Decomposition of Carbamde. Maxi-m? Dma- vane-tion of No. of Mean tion of Concentration of reaction Tern- expt. in observa- vahe of k from lnirrfure. perature. minutes. tions. k x 109 mean. W-Carbamide + 2N-HNO, ... 100’ 420 7 73 2.0 , +N-HNOs ... 100 270 6 105 4.0 , + N-HN08 100 360 5 87 4-0 , + N-HNOs 100 1800 9 88 3.0 , + NIS-HNOS ... 100 1200 9 10.2 5-0 , + N/4-HNOa ... 100 420 7 111 3.0 , + N/8-HNOa ...100 420 6 116 6.0 , + NIlG-HNOS 100 3 60 7 133 8-0 122 5.0 91 + N/2-NH,N03} loo l5o0 10 900 9 118 6-0 420 5 118 3.0 600 8 116 3-0 , + NIS-HNOI ... 100 600 5 108 2.0 -l- N/S-rnO* ... $- N/2-HNO ... ’ 9 + N/2-KW08 .. . } loo + NP-HMO ..*} 100 9’ -I- NIB-KCI ...... , + N)2-HNOS ...- 89 1890 9 23.0 2-8 , + N/2-ECNO9 ... 80 2790 9 9.3 0.7 , + N/2-HNO ... 79 9360 10 2 4 0. 1358 PBIOE THE DECOMPOSI!MON OR Experiment 17 .-Twenty-five C.C. of N-carbamide and 25 C.C. of N-nitric acid were heated a t looo in a boiling tube fitted with a reflux condenser. A delivery tube connected the top of the con-denser to a nitrometer. The total air space from the level of the liquid in the boiling tube to the nitrometer was 34 C.C. One hundred QC. of gas were collected in the nitrometer and rejected.It was considered that by this time all the air had been swept out of the apparatus and collection of the gas for analysis was com-menced. Fifty C.C. of this gas were almost entirely absorbed by a piece of moist potassium hydroxide only a minute bubble being left. The gas was thus composed entirely of carbon dioxide and did not contain any nitrous oxide or nitrogen the small residue being either air which had not been completely swept out of the apparatus or nitrogen from the trace of nitrous acid present in the nitric acid. The heating of the carbamide and nitric acid solution was con-tinued for fifty hours a t the end of which the solution was evaporated to dryness. On analysis the residue was found t o consist of ammonium nitrate with a little udchanged carbamide.Discussion of Results. The results show that the reaction between carbamide and nitric acid in dilute solution is undoubtedly unimolecular. The products consist entirely of carbon dioxide and ammonium nitrate. Hence the reaction is analogous to the decomposition of carbamide by hydrochloric or sulphuric acid and is not due to the preliminary decomposition of nitric acid into nitrous acid with subsequent decomposition of the latter by carbamide. On comparing the results with those obtained by Fawsitt for hydrochloric acid in table 11 i t will be seen that a t all dilutions the velocity of reaction is greater with nitric acid thm with hydro-chloric acid. TABLE 11. Showing Comparison hetween Velocity of Reaction at looo of Carbamide with Nitric Acid and Hydrochloric Acid.k x 109 - Remtion mixture. HNOS. Ha." - 2N-Carbamide + 2N-acid ...... 73 N- , + N- , ...... 07 50 2712- , + N12- 9 ...... 107 77 N/4- , + N/4- , ...... 111 90 Xf0- , + N/8- , ...... 116 101 Nll6- , + N/16-, ...... 133 101 * Fawsitt's results CARBAMIDE IN THE PBESENCE OF NITBIO ACED. 1369 The velocity of reaction diminishes regularly with an increase in the concentration of the nitric acid and no maximum point was found such as Fawsitt found for hydroahloric acid (Eoc. cit. p. 612). Fawsitt and Werner both state that only free carbamide is active. According to Werner the equilibrium between carbamide and its nitrate can be represented as follows: and the reaction between carbamide and nitric acid on the analogy of the reaction between carbamide and hydrochloric acid can be represented thus : Phase 1.HN:C<lH3+ HN03=NH,N0,+ (HNCO LZ HO*CN). Phase 2. (HNCO HO*CN) + H,O + HN0,=NH,N03+ CO,. The diminution of the velocity with increase in concentration of the acid can thus be explained since it is only “free” carbamide, and not carbamide “fixed” as its nitrate which takes part in Phase 1 of the reaction. Fawsitt found that the decomposition of carbamide in the presence of hydrochloric acid was retarded slightly by the addition of ammonium chloride but was accelerated in the presence of sulphuric acid by the addition of ammonium sulphate. He also found that the decomposition of carbamide in the presence of water alone was accelerated by the addition of ammonium carbonate sodium chloride and potassium chloride whilst ammonium chloride and ammonium hydrochloride had a decided retarding effect.In the presence of nitric acid the present author has found that ammonium nitrate potassium nitrate ammonium chloride, and potassium chloride all have a distinct accelerating effect on the decomposition of carbamide a t N / 2-concentration. According to the dissociation theory of the decomposition of carbamide the first action of heat on it is the production of ammonia and cyanic acid; in the presence of acids this.is followed by combination of ammonia with acid forming an ammonium salt, and by hydrolysis of cyanic acid forming an ammonium salt and carbon dioxide. The addition of the ammonium salt or of any salt containing an ion common with the ammonium salt t o the reaction mixture should therefore have a retarding effect on the velocity of decomposition.As shown by the experiments this is not the case when carbamide is decomposed in the presence of nitric acid and hence it would seem that the dissociation theory o LEWIS STUDIES IN CATALYSIS. PART XEI. the decomposition of carbamide is not applicable in the preeence of nitric acid. The great influence of temperature on the velocity will be seen from table III. TABLE 111. Showing EfJect of Temperatwe on ‘Vetocity of Reaction between. N/2-Carbamide and N/2-Nitric Acid. looo 102.0 89 23.0 80 9.3 70 2.4 Temperature. k x los. Below 80° the velocity of decomposition of carbamide is small, and a t 30° to 40° would be negligible; hence the use of excess of carbamide for the removal of nitrous acid from a mixture of nitrous acid and nitric acid will not be accompanied by loss of nitric acid as such if the temperature is not allowed to rise above 40°. In conclusion the author’s best thanks are due to Meesrs. Nobel’s Explosives Company Limited and to Mr . William Rintoul Manager of the Research Section for the facilities afforded for carrying out this work and for permission to publish the results. THE RESEARCH LABORATORIES, ARDEER . [Received November 6th 1919.

ISSN:0368-1645    

DOI:10.1039/CT9191501354

出版商:RSC    

年代:1919

数据来源: RSC

摘要:

LEWIS STUDIES IN CATALYSIS. PART XEI. CSXV1I.-Studies in Catahpis. Part XII. Catalytic Cderia and the Radiation Hypothesis. By WILLIAM CUDMORE MCCULLAGH LEWIS. TEE criteria which have been suggested from time to time as apply-ing to the phenomenon of catalysis are as follows (compare Rideal a.nd Taylor “Catalysis in Theory and Practice,” Chap. 2): (I) The chemical composi$ion of the catalytic agents is un-changed on completion of the reaction process. (2) Minimal amounts of a catalytic agent are adequate for the transf armation of large quantities of the reacting substances. (3) A catalyst does not affect the final state of equilibrium. (4) A catalyst modifies the velocity of two inverse reactions to the same degree LEWIS STUDIES IN CATALYSIS. PART XII. 1361 (5) A catalytic agent is incapable of starting a reaction; it c&~l only modify the velocity of the reacbion.Criteria (I) and (2) are closely related (2) in fact being the corollary of (1). Both would be accepted a t once provided secondary effech are excluded. To this extent they state a fact of experience and indicate that catalysis is simply a special case of ordinary chemical reactivity. Criteria (3) to (5) which form a group by themselves are in a different category as representing generalisations which may or may not be true. Criterion (4) is the corollary of (3) so that the group contains two distinct criteria. Considerable difference of opinion exissts a t the present time regarding the validity of these conclusions according to the point of view adopted as the basis of criticism.It is of some interest therefore to examine criterion (3) or ( 4 ) and criterion (5) from the point of view of the radiation hypothesis of chemical reactivity. On the radia-tion hypothesis the possibility of a reaction occurring depends on the existence of radiation of a type or frequency absorbable by the reacting substance the quantum of which radiation is suffici-ently large t o communicate the necessary critical increment to the molecule. I n the case of thermal radiation which is the kind of radiation envisaged in the quantum theory theoretically all possible wavelengths or frequencies are represented a t any temperature. Consequently the type of radiation necessary for any reaction is present in the space occupied by the matter and therefore every reaction is correspondingly possible.This mu* include the so-called catalytic reactions as well as those to which this name is not applied. From this point of view therefore c a ~ we conclude that a catalyst does not initiate but simply accelerates a process which would occur although under certain conditions infinitely slowly ? This cannot be affirmed withont qualification. In the form in which criterion ( 5 ) is stated it is evidently assumed that the same process may occur whether the catalyst be'present or not but this assumption is not necessarily true. It seems necessary to ascribe in certain cases if not in all a definite stoickeiometric molecular mechanism to a catalyst just as one would to any other reactant and consequently by adding such 8 catalyst a new process commences (the origin of which is the field of radiation) that happens to give rise to certain end-produh, which we believe might be attained in the absence of the catalyst.The fact appears to be that criterion (5)) as ordinarily stated, involves a false antithesis. From the point of view of the radis-We shall consider criterion (5) in the first place 1362 LEWIS STUDIES IN CATALYSIS. PART XII. tion hypothesis a catalyst may be said either ta render a reaction possible by supplying the necessary matter or it may merely accelerate according to circumstances. In no case is it the funda-mental initiator of a process. The r81e played by the catalyst may be conveniently illustrated by the catalytic effect of an acid in the inversion of sucrose or the hydrolysis of an ester.Prior to the addition of the acid the reaction is possible involving reaction between a molecule of sucrose or ester and either a molecule of water or its ions probably the undissociated molecule. On addi-tion of the catalysing acid the hydrogen ion accelerates the process already begun by the hydrogen ions already present. The undis-sociated molecule of the acid may also accelerate the reaction but in doing so it is almost certain that it produces an intermediate substance which was not formed in its absence. In so far as the intermediate stage is concerned the molecule of the acid has rendered a new intermediate process possible although the final produds are independent of the nature of this intermediate stage.The real source or origin of initiation of any reaction on the radiation hypothesis is the radiation itself. The material catalyst, if it acts simply as a transformer hastens a reaction which radia-tion has already initiated. The catalyst may also act as a mole-cular reactant giving rise under the stimulus of radiation to new intermediate products. The validity of conclusion (5) depends, therefore on the particular view adopted regarding the mechanism of the process. It seems that two distinct modes of mechanism are possible and are apparently realised in the well-known acid catalysia. On one mode the catalyst simply accelerates; on the other it renders a new mechanism possible from the material point of view. Turning now to criterion (3) or (4) which possesses much greater practical significance the radiation hypothesis leads t o the con-clusion that as a general principle criterion (3) or (4) is certainly not true.Let us take the simplest possible case of reversible reaction, represented by A B. The substance A is characterised by being capable of absorbing radiation of frequency v, as a result of which it is transformed into B. The substance B is likewise capable of absorbing radiation of frequency vE as a result of which the process is reversed. The heat evolved & on passing from A to B is then given by the expression where N is the Avogadro number and & is referred to one gram-molecule of .4 transformed. Q = Nh(v - V A ) LEWIS STUDIES IN CATALYSIS. PART XII. 1363 Let us consider the special case in which vA=vB or approxim-ately so.A catalyst acting as a transformer will in this case be unable to distinguish between the two types of molecules A and B since each is capable of absorbing the same type or approximately the same tspe of radiation. It will therefore catalyse both the direct and the reverse reaction equally. That is the opposing velocity constants will be equally increased and the equilibrium constant will remain unaffected by the presence of the catalyst. This result is in harmony with the criterion. If on the other hand the heat of the process is considerable, that is v A differs considerably from vB then it no longer follows that a positive catalyst will equally accelerate both reactions. In general it would not be expected to do so and consequently in general the equilibrium point will be affected by the catalyst.It is a significant fact in view of the conclusion just drawn that those reactions such as esterification or hydrolysis in which the equilibrium point is not sensibly affected by the catalyst are precisely those in which the heat effect is small. In the above case we have been considering mainly homogeneous catalysis by means of ions. Let us now take the case of catalysis by the undissociated molecule such as the molecule of hydrochloric acid which is generally regarded as functioning through the form-ation of an intermediate ternary compound. Thus Falk and Nelson ( J . Amer. Chem. SOC. 1915 37 1732) represent the inter-mediate oxonium complex in the case of hydrolysis of esters aa (ester,HCI,H,O).I n the reverse process the corresponding com-pound is (carboxylic acid,HCl,alcohol). These two compounds are tautomeric and may be identical. I f they are identical as Falk and Nelson assume then the hydrochloric acid molecule will equally affect the direct and the reverse process and thus leave the equilibrium point unchanged. This explanation of the mechanism of the effect produced by the undissociated molecule of the cata-lysing acid has certainly the advantage of simplicity. It has this implication however. Such additive compounds are generally formed rapidly compared with the rate of any further decomposi-tion which they may undergo If this is so and if the same inter-mediate compound is formed in the hydrolysis as in the esterifica-tion it would follow that the velocity constants should be the same, and the equilibrium constant should therefore be unity.This is not in agreement with experiment although it is significant that the value of K is not greatly removed from unity. Thus experi-ment has shown that In this case Q is zero or approximately so. [methyl acetate] x [water1 [methyl alcoholl x [acetic acid] = 4.6 R 1364 LEWIS STUDIES IN CATALYSIS PART XII. (compare Part V of this series of papers T. 1916 189, 71). This ratio means that the velocity constant of es'teriflcation is 4-6 times the velocity constant of hydrolysis. On the radiation view this ratio is mainly determined by the relative value of the exponential terms that is by where Nhvl is the critical increment of esterification and Nhv is the critical increment of hydrolysis that is, ~ M ~ Y z - v1 )IB T = 4.6.A t 300° absolute we find therefore that vZ-vl=1 x 1013. Both v1 and v do not however lie very far from the value 2 x 1014 so that the difference in respect of position of absorption of infra-red radiation is extremely small being of the order of one-twentieth of the absolute value of either frequency. The same idea is con-veyed by saying that the heat effect does not exceed 1000 calories. The fact therefore that the equilibrium constant possesses a value not unity but not far removed therefrom means on the radiation basis that the intermediate compounds are not identical but tauto-meric and further that both kinds of molecule absorb almost the same frequency so that any change in the equilibrium constant introduced by altering the concentration of the catR1yst is insensible.( I t may be noted that the relatively large change in R observed by Lapworth has its origin as Lapworth has shown, in what is virtually a distinct reaction not directly connected with the actual esterification-hydrolysis process itself .) I n the case con-sidered we conclude therefore that as a practical guide the criterion (3) or (4) is true that is in those cases in which the equilibrium constant is not far removed from unity. The con-clusion is obviously comparable with that drawn in connexion with ion catalysis. In addition to honiogeneous catalysis by dissolved substances in a given solvent it is well known that different solvents themselves exert their own catalytic effect.From the point of view of radia-tion we conclude that in this type of catalysis criterion (3) or ( 4 ) cannot in general be even approximately true since each solvent is characterised by its own electromagnetic properties that is by its power of absorption at different wave-lengths which differs from solvent to solvent and consequently entails a different dis-tribution of radiation density. The particular type of radiation required by the reactant is therefore present to a different extent, according to the nature of the solvent and consequently th LEWIS STUDIES IN CATALYSIS. PdRT XII. 1365 velocity constant and equilibrium constant is a function of the solvent. Specific differences are further introduced by the mutual interaction of solvent and solute whereby the effective frequency itself is altered to a slight extent.Finally as regards heterogeneous catalysis evidence has been collected and presented by Bancroft ( J . Physical Chem. 1917 21, 573; 1918 22 433) to show that the catalyst affeda the equil-ibrium point of the process. Heterogeneous catalysis has been examined in a preliminary manner from the point of view of radiation.(T. 1919 115 182) and it is concluded that the equil-ibrium point must be a function of the nature and extent of the catalytic material owing t.0 the alteration in the values of the critical incremenh which is introduced by the presence of the catalyst. The catalytic layer here considered is only one molecule, or possibly two molecules in thickness that is it is of the order 10-8 cm.In this layer the final amounts of reactants and resultants will differ in general from the true equilibrium amounts characteristic of the homogeneous gas phase for in the adsorption layer the relative amounts are determined by the relative adsorp-tion capacities. If this is a complete statement of the phenomenon, i t is evident that criterion (3) or (4) is inapplicable. There is, however a further possibility to be considered. The adsorbed reactants and resultants are in an activated con-dition as long as they are actually in the adsorption layer. I n time they necessarily pass out into the homogeneous gas phase, owing to desorption. If these molecules a t the moment of leav-ing the adsorption layer lose the extra energy which they possess and become immediately transformed into molecules of normal energy content it follows that their relative concentration in the homogeneous gas phase becomes identical with that in the adsorp-tion layer for in general the reaction in the homogeneous phase is extremely slow.There is the possibility however that the activated reactants and resultants on leaving the true adsorption layer do not immediately revert to the inactive state but may retain their activity for a short space corresponding perhaps with a layer 10-6 cm. in thickness. If this is the case there will be a rapid chemical change in this extra-adsorption layer which will tend t;o bring the concentrations of the reactants and resultants into the ratio required by the law of mass action for the homo-geneous phase.Whether the true equilibrium point would be attained or not would depend on the average life of the activated molecules. A t moderately high temperatures the average mole-cular velocity may be taken to be 105 cm. per second; hence it would require 10-11 second for a molecule to traverse a distanc 1366 LUMSDEN CRITERIA OF THE DEGREE OF of 10-6 cm. It is possible that the activate’d state may be main-tained for a longer period of time than this and therefore the more likely is the process to attain the true equilibrium position. The more selective the nature of the adsorption material &he further in general will the “ equilibrium ” of the adsorption layer depart from the true mass action equilibrium. Hence even if such a compensating effect as that suggested above actually operates, criterion (3) or (4) cannot be regarded as valid. It is concluded therefore that criterion (3) or (4) is in general, not true; in homogeneous systems it approximates more closely to experiment the smaller the heat effect accompanying the reaction ; in heterogeneous systems it is not certain whether even this approximation to validity holds good. M u s ~ u r r LABORATORY OF PHYSICAL DD ELECWRO-UHEMISTRY, UNIVERSITY OF LIVERPOOL. [Receiued October 31at 1919.

ISSN:0368-1645    

DOI:10.1039/CT9191501360

出版商:RSC    

年代:1919

数据来源: RSC

摘要:

1366 LUMSDEN CRITERIA OF THE DEGREE OF CXXT'III .- Criteria of the Degree Commercial Toluene. By JOHN SCOTT LUMSDEN. THE following investigation was undertaken o j Purity of for the Explosives Department of the Ministry of Munitions for the purpose of find-ing some easily applied method for estimating the degree of purity of commercial toluene and is published with the permission of the Department. Commercial toluene contains varying amounts of a liquid which cannot be nitrated and has distilled along with the toluene during rectification. This liquid has not been isolated but in the follow-ing tests the assumption is made that it may be represented by a paraffi mixture which boils close to the boiling point of toluene. By fractionation of petrol a quantity of such a liquid was obtained boiling at 108-112°.Estimation of Toluene b y Specific Gravity.-Pure toluene a t 1 5 O wa,s found to have D 0.8712; the p a r a f i mixture had D 0.743. Since there is no change in volume on mixing toluene and paraffin, a graph was prepared to show the density of mixtures. From this graph it mas found that the presence of 1 per cent. of paraffin produced a lowering of the density of 0.0013. Whilst such a graph may not be depended on to give the accurat PURITY OF COMMERCIAL TOLUENE. 1367 percentage of impurity present because the toluene may be moist and the impurity may not be like that assumed it is certainly useful in indicating poor samples. For example of two samples of commercial toluene (a) had D 0.872 and ( b ) 0'8613. Both were fractionated with the following results Sample (a) showed a trace of moisture then boiled steadily a t l l O o .All fractions right to the end had D 0.8716. This was therefore a very good sample of toluene. Sample ( b ) showed a little moisture and then boiled half a degree below alowly rising to half a degree above l l O o . The fractions had D 0.8623 0-8621 0.8622 0.8619 0.8612 0.8586 and 0.8570. The original low density and the decreasing density of the fractions pointed to a considerable admixture of paraffins; from the graph the indication is 7.7 per cent. Estimation. of Toluene b y Nitration.-To effect nitration, potassium nitrate and sulphuric acid were employed as described later. It was soon found that the nitration of toluene cannot be made to stop when the mononitro-compound is formed but that any excess of acid is used up producing the dinitro-compound; it was however proved as the result of many experiments that nitration stops quantitatively when all the toluene has been con-verted into the dinitro-compound.Using nitration as a means of estimating toluene there are obviously two methods which may be adopted namely by nitrating with a weighed excess of potassium nitrate and determining the weight of nitric acid remaining after nitration the amount of acid used is obtained and from that the weight of toluene nitrated, or by isolating and weighing the dinitro-compound. Adopting the first method it was found that the residual acid could be satis-factorily determined by the nitrometer but that by the second mekhod the dinitro-compound could not be completely recovered, because a small but definite amount remains in solution in the acid liquid from which the substance is filtered.A 100 C.C. flask is weighed, and into this is weighed accurately by dropping from a fine pipette 2 grams of the sample to be nitrated. Five grams of pure dry powdered potassium nitrate are weighed in a watch-glass or aluminium scoop and transferred with great care to the flask. The flask is cooled in water and 30 C.C. of pure concentrated sulphuric acid are poured in rapidly the flask being shaken con-tinuously to prevent the formation of solid lumps which if formed, take aome time to dissolve. When the action is over a pale yellow, homogeneous liquid is obtained. To cause the dinitro-compound t o separate about 60 C.C.of water are added the first few C.C. being dropped in alowly lest The nitration process is as follows 1368 LUMSDEN CRITERIA OF THE DEGREE OF the heat evolved should volatilise any nitric acid and at the same time the flask is cooled in water. The solid is obtained in soft flakee which are nearly white and when the contents of the flask are quite cold ihe separation may be considered complete. Filtration is effected by the aid of the pump; using a small porcelain funnel with a small filter paper ; the preparation flask is washed out with several small amounts of wabr which are poured through the funnel and the solid is washed free of acid. The filtrate is transferred to a graduated cylinder, and together with the washings of the filter flask is made up to exaotly 150 C.C.The whole operation of nitration prweds smoothly and the only point where care is needed is when adding water to the nitratad liquid great precautions being then necessary to avoid rise of temperature. Calculations C,H,-CH + 2KN0 5 C,H,(NOS),-CH,. 92 202.2 182.0 2 4.3956 3.95 (a) Two grams of toluene require for nitration 4.39 grams of potassium nitrate. One per cent. of this-0.0439 gram-remain-ing unused represents 1 per cent. of impurity in the toluene. 0.004522 Gram of potassium nitrate gives 1 C.C. of gas in the nitro-meter ; 0.0439 gram therefore represents 9-72 C.C. Working with 2 grams of toluene 9-72 C.C. of gas in the nitrometer a t N.T.P. represent 1 per cent. of impurity.( 6 ) Five grams of potassium nitrate are used to nitrate 2 grams of toluene but only 4.3956 grams are required. This is an excess of 0.6044 gram which in the nitrometer would give 133.65 C.C. of gas. ( c ) Two grams of toluene give 3-95 grams of dinitrotoluene. One per cent. of this-0.0395 gram-as a deficit in the yield repre-sents 1 per cent. of impurity in the toluene. (a) Estimation by Nitrometer.Two grams of toluene were nitrahed with 5 grams of potassium nitrate. The liquid was made up to 150 c.c. and 15 C.C. were taken for the nitrometer estimation. (The 15 C.C. were measured from a small burette into the cup of the nitrometer and 10 C.C. of pure sulphuric acid were used to liberate the gas.) 19.8 C.C. of gas a t 16O were obtained or 18-71 a t N.T.P.For the whole 150 C.C. this was 187.1 c,c.; deducting the known excess used namely 133.65 c.c., leaves 53-45 C.C. As 9.72 C.C. represent 1 per cent. of impurity, 53.45 c,c. represent 5.5 per cent. of impurity. Examples of Nitration. The following are the results of other estimations PURITY OF COMME~IAL TOLUENE. 1369 Pure toluene. 150 C.C. liquid 16 C.C. taken. 13.27 C.C. g a ~ at N.T.P. 132.7~~. For 160 C.C. liquid, From excess taken 133.6 -. -0.9 C.C. - 0.9/9-72 = - 0.1 per Toluene 100-1 per cent. cent. Good toluene. 176 C.C. liquid 20 C.C. taken. 16.7 C.C. grts Elf N.T.P. 137.37 C.C. For 176 C.C. liquid, From exces8 133.65 3.72 C.C. 3.72/9.72 = 0.38 per Toluene 100 - 0-38 = cent. 99.62 per cent. Poor toluene. 160 C.C.liquid 15 C.C. taken. 18.64 C.O. g m at N.T.P 186-4 C.C. For 150 C.C. liquid, From excess 133.6 -62-8 0.0. 52.8/9.72 = 6.4 per Toluene 100 - 6.4 = cent. 94.6 per cent. The sample of pure toluene shows more than 100 per cent.; this is due to experimental loss. It will be realised that if the potassium nitrate is not pure and dry or if there is any loss before the filtrate is made up to a definite volume less nitric acid is found, and this is indicated as a small increase in the percentage of toluene. (b) Estimation by Weighing the Solid.-The solid dinitxo-com-pound is transferred from the filter funnel to a watch-glass and set in a desiccator over sulphuric acid or the watch-glass may be set over a beaker of boiling water when the solid melts to a clear liquid with water beneath it.On cooling the solid forms a firm mass from which water is readily removed by filter paper; then, after a much shorter time in the desiccator the weight becomes constant. It was found that when the volume after nitration was made up to 150 c.c. approximately 0.11 gram of solid remained in solu-tion and this amount is added to the weight found in the follow-ing examples. Two grams of “pure ” toluene gave 3.8337 grams of solid; add-ing 0.11 gram the total dinitro-compound was 3.9437 grams; 100 per cent. toluene should give 3.9564 grams therefore the deficit is 0*0127 gram. 0.0395 Gram deficit represents 1 per cent. of impurity and 0.0127 gram 0.32 per cent. The toluene was there-fore 100 - 0*32 = 99-68 per cent.Two grams of a poor sample of toluene gave a deficit of 0.2573 gram that is 0.2573/0*0395=6*5 per cent. The toluene was therefore 100 - 6.5 = 93.5 per cent. Estimation of Toluene by Miscibility with Acetic Acid. Whilst engaged on this investigation a report by Professor Orton was received describing the experimental work detailed in his recent paper (this vol. p. 1055). He there showed that whe 1370 LUMSDEN CRITERIA OF TEE DEQBEE OF 88 to 90 per cent. acetic acid is mixed with hluene m the propor-tions of 1 C.C. of acid to 0.8 C.C. of toluene tws layers are formed, and on raising the temperature a point of complete miscibility is reached which is definite for the same strength of acid and so sharply defined that it can be read to one-tenth of a degree. He showed also that the temperature of miscibility of toluene rises regularly with the dilution of the acetic acid and further that for a given strength of acid the temperature of miscibility of toluene containing paraffin rises in direct proportion to the amount of paraffin present.Working with two strengths of acetic acid 87.9 and 89.5 per cent. and using pipettes of certain volumes he prepared curves from which the percentage of a mixture of toluene and paraffin could be found when the temperature of miscibility with one of these acids had been determined. The method is susceptible of great accuracy but it is very troublesome to prepare acetic acid of a definite strength and almost impossible to keep a concen-trated acid without absorption of moisture from the air and more-over it requires so much care to carry out the process as set down by Professor Orton that it could scarcely be used as a commercial test.The folfowing modification consisting in the changing of an absolute t o a comparative method may however be carried out in any commercial laboratory. A quantity of concentrated acetic acid of unknown strength is taken and with equal volumes of this the temperatures of misci-bility with equal volumes of pure toluene toluene with 5 per cent. of paraffins and the sample of toluene to be tested are found. Then since the percentage of paraffins is proportional to the increase in the temperature of miscibility the comparison of the rise of temperature of miscibility of the sample with that of the 5 per cent.paraffin mixture gives the paraffin content of the sample. To carry out the test there is required: (1) A quantity of pure toluene which need not be synthetically pure but might be considered as 100 per cent. quality. (2) A mixture of this toluene with 5 per cent. of paraffins made by mixing 95 grams of toluene with 5 grams of paraffins of the same boiling point prepared by the distillation of petrol. (3) Acetic acid prepared by taking 100 grams of glacial acetic acid adding a few C.C. of water and by trial with toluene in the proportions described below finding if two layers are formed and if the temperature of miscibility is somewhere between 25O and 30°. (4) The sample to be tested. The proportions of acetic acid and toluene found by Professo PURITY OF COMMERCIAL TOLUENE.1371 Orton as suibable for the test may be adhered to and two pipettes prepared one delivering 1 C.C. and the other 0'8 c.c. or quantitiea in these proportions but they need not be exact. The pipettes are made from narrow glass tubing and must have fine capillary points to deliver very slowly. A test-tube 1.2 cm. in diameter serves as a vessel in which the miscibility point is determined and while hhe test is being made it may for convenience be fixed to the thermometer by a rubber band. The test is carried out as follows. With the larger pipette acetic acid is run into the fest-tube and pure toluene is added from the smaller pipette. The tube is then fixed to the thermometer which must be graduated in tenths of a degree. Water is heated in a beaker holding about 2 litres the temperature being raised speedily and by stirring with the thermometer and tube an approximation t o the miscibility point is found by noting the temperature when the two layers disappear and a homogeneous liquid is obtained.A little cold water is then added t o the beaker the flame is turned very low the water thoroughly stirred and as the tempera-ture slowly rises the thermometer and tube are moved continuously in the water. As the miscibility point is approached the two layers disappear the liquid becomes opalescent and then auddenly becomes transparent. This is the point a t which the temperature is noted. By cooling the water in the beaker lo the test may be repeated. The same procedure is gone through with the 95 per cent.toluene and with the sample being examined the tube being washed out thoroughly each time and dried in an air-oven. The following are results obtained : Pure 95 Per cent. toluene. toluene. Sample. (1) The miscibility temperature wm 32.0' 38.3" 414O (2) Same sample on another day, the acid being different ............ 33.4 40.3 44.0 From (1) 38.3" - 32" = 6.3" and 41.4" - 32" = 9.4". A rise of 6.3" represents 5 per cent. of impurity 9.4" represents 7.4 per cant. From (2) 40.3" - 33.4" = 6.9" and 44.0" - 33.4" = 10.6'. A rise of 6.9" represents 5 per cent. of impurity 10.6" represents 7.6 per cent. It is seen that the strength of the acid does not require to be known it has only €b remain constant during the time that three portions are withdrawn and the pipettes do not require to be of any definite volume although they must deliver precisely the same volume each time.The three tests might be carried out simultaneously by having three test-tubes attached to the thermometer one with each mixture, VOL. cxv. 3 1372 BAXTER AND FARGHER 1 3-BENZODIAZOLEARSINIC and the three points of miscibility determined as the temperature of the water is raised. When the same sample of toluene was tested by the different methods the following results were obtained : From specific gravity curve .............................. 92.3 per cent. From nitrometer estimation after nitration ............ 92.9 ,, By weighing dinitro-compound ........................... 92.6 ,, By temperature of miscibility with acetic acid ...... 92.5 ,, Conclusions.-The specific gravity of commercial toluene gives a fair indication of the amount of impurity present; the estimation of the temperature of miscibility with acetic acid is quickly done, and although a comparative method if skilfully carried out is trustworthy; the nitration and estimation of the excess of nitric acid by the nitrometer is believed to be the most accurate test and indicates the total nitration which has taken place and the weighing of the solid after nitration gives the actual practical yield of nitro-compound . By none of the methods under ordinary conditions of working, can an accuracy closer than one-half per cent. be expected. THE TECHNICAL COLLEGE, DUNDEE. [Received October 16th 1919.

ISSN:0368-1645    

DOI:10.1039/CT9191501366

出版商:RSC    

年代:1919

数据来源: RSC

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1372 BAXTER AND FARGHER 1 3-BENZODIAZOLEARSINIC CXXIX.-1 3-Benxodiazolearsinic Acids and their Reduction Products. By ROBERT REGINALD BAXTER aiid ROBERT GEORGE FARGHICR. SINCE the discovery of salvarsan (Ehrlich and Bertheim B e y . 1912, 45 756) many attempts have been made to prepare derivatives or closely allied products which would render unnecessary the some-what elaborate technique involved in the clinical use of the original compound which owing to its pronounced acidity has to' be con-verted into the disodium salt in turn alkaline in reaction before use. Neosalvarsan the sodium N-methylenesulphinate of salvarsan, introduced to avoid these difficulties possesses many advantages, such as ready solubility and neutral reaction but in spite of these, there is a consensus of opinion in favour of salvarsan as its spire chaeticidal action appears to be stronger and more certain.In this and in other cases the attempts have led to the formation of acid derivatives giving neutral salts with alkalis. The authors desired 021 the other hand to obtain hydrochlorides soluble in water and sufficiently less acid in reaction than salvarsan to admi ACIDS AND THEIR. REDUCTION PRODUCTS. 1373 of their direct use and it was with this object in view that the present investigation was commenced since it was considered prob-able that the hydrochlorides of the arsenobenzenes derived from 1 3-benzodiazole (benzoglyoxaline) would fulfil these requiremenh. Further considering the properties of the glyoxaline nucleus they might be expected to exert acidic as well as basic properties and thus closely resemble salvarsan itself.Additional interest would attach to such compounds in view of their relationship to 3 4 5 3’ 4’ 5’-hexa-aminoarsenobenzene and its N-methyl derivatives (D.R.-P. 286667 286668 286854 286g55, 294276 ; E.P. 7488 and 8041 of 1913 ; U.S.P. 1081079 ; Benda Ber., 1914 47 1316; Karrer Ber. 1916 49 1448) which are stated t o possess the unexpected property of dissolving in alkali hydrogen carbonates f ormiiig carbamates with the same degree of alkalinity as that of normal blood serum and a relatively low toxicity So far as the authors are aware the beiizodiazolearsisic acids have not been examined Bertheiin (Rer. 1911 44 3092) who described 3 4-diaminophenylarsinic acid and several of its deriv-atives making no mention of them.They are however readily obtained by the well-known reaction of ortho-diamines with formic or acetic acid whereby the acyl derivative first formed suffers loss of a molecule of water if a monoacyl or of a molecule of acid if a diacyl derivative with consequent closing of the ring (compare Ladenburg Ber. 1875 8 677; Wundt Ber. 1878 11 826). The comparative stability of 3 li-cliacet?/lan~i.no~her~~lal.sinic acid, which crystallises from boiling water unchanged and does not alter save for slight decomposition when heated a t 250°/20 mm. and, generally the stability of diacyl-o-diamines (compare Bistrzycki and Ulffers Bcr. 1890 23 1876) seems t o indicate that in the absence of anhydrides the reaction proceeds with intermediate :‘ormation of a moiioacyl derivative.5 6-l)iamino-m-tol?/lars~?~ic acid ( I ) which together with 3 4-diaminophenylarsinic acid (Bertheim lw. cit.) formed the starting material for the investigation was obtained by the reduction of 6-nitro-6-amino-~~z-tolylarsinic acid (this vol. p. 989) with sodium hyposulphite. It undergoes all the condensation3 typical of an (D.R.-P. 269660; E.P. 1667 of 1914). /\ CH,NH I I N \/ () /\/ I )AsOP, NH2 / \ q N H20,AJ ’ ‘ / /\/\/ CH,(\JNH, \/ \rLN AsOs K2 () CH, (1.1 01.1 (111.) 3 ~ 1374 BAXTER AND FARGHER 2 ~-BENZODIIIZOJiWLRSINIU o-diamine giving for example with sodium nitrite in acid solution, 7-methyl-1 2 3-benzotriazole-5-arsinic acid (11) and with phen-ant h raquinone 4-m e t h y l phenarzt h raph enaaine - 2-ar sinic acid ( 111).The latter derivative shows the colour reactions charachristic of phena~thraphenazine. With boiling glacial formic acid the acids readily yielded 7-methyl-l 3-benzodiazole-5-arsinic acid (IV) and 1 3-benzodiazole-5arsinic acid (V) respectively. CH NH / A /\/\CH I \1 "HCH H208As\/-N H20,Ad I - - - h (V. 1 \/ (IV.) The action of acetic acid did not proceed quite so smoothly. 3 4-Diaminophenylarsinic acid although stated (Bertheim loc. cit.) t o be readily soluble in acetic acid is far from being so and on boiling with acetic acid acetic anhydride or a mixture of the two, gives rise to very highly coloured products which are difficult to purify. If the reaction is carried out as described in the experi-mental portion of this paper the main product is 3:4-diacetyE-aminophenylarsinic acid which on heating with water in a sealed tube a t 130° yields 2-methyl-1 3-benzodiazole-5-arsinic acid (VI).I n the case of the higher homologue the desired reaction takes place predominantly with the formation of 2 7-dimethyl-1 3-be~zo~zazole-5-arsinic acid (VII). Incidentally it may be men-tioned that it has since been f ouiid that if 3 4-diaminophenylarsinic acid dissolved in the requisite amount of sodium hydroxide (40 per cent. solution) to form the monosodium salt is treated with sufficient acetic anhydride to neutralise the sodium hydroxide and form acetic acid with the water present and then boiled the main product of the reaction is the benzodiazole derivative.(VII.) NH NH (VIII.) The reduction of the above acids with hypophosphorous acid, although it proceeds normally leads to gelatinous products difficul ACID8 AND TREIR REDUOTION PBODUOTS. 1376 to handle and consequently the reduction was carried out by means of sodium hyposulphite. This gives pale yellow arsenobenzenes, insoluble in water which were converted into hydrochlorides by precipitating their solutions in acetic acid with excess of hydro-chloric acid or by treating the suspensions of the bases in water with just sufficient hydrochloric acid t o bring about solution and precipitating as before. The hydrochlorides form pale yellow powders soluble in water but very sparingly so in the usual organic solvents. Their aqueous solutions react strongly acid to litmus, but neutral to methyl-orange.It was expected that this acidity would prove too great for pur-poses of intravenous injection and that it would be difficult to distinguish between the effect due t o the acidity and that due to the arsenic compound. This was confirmed by experiments with 5 Eil-arseno-l 3 1' 3I-benzodiazo7e (VIII) by Miss Soref of the Wellcome Physiological Research Laboratories. EXPERIMENTAL. 3 4-Dinmino~~eny7arsinic A cid. The preparation of considerable quantities of this acid was carried out according t o the directions given by Bertheim (Zoc. cit.), save that i t w'as found advisable to omit the boiling of the solution after the reduction with sodium hyposulphite as this diminished the yield and gave a more highly coloured product.1 3-Ren,zodinzo?e-5-nl.siltic Acid (V). 'Ten grams of 3 4-diaminophenylarsinic acid were boiled under a reflux condenser for six hours with 100 C.C. of glacial formic acid. After removal of the excess of formic acid by distillation water was added when 7.5 grams of crystalline material rapidly separated. To remove adhering traces of colouring matter the product was treated with charcoal in dilute alkaline solution. On making neutral to methyl-orange crystallisation commenced a t once practically the whole being recovered in a pure state. 1 3-Benzodiazole-5arsinic acid crystallises from water in which it is sparingly soluble in clusters of minute flattened prisms which are anhydrous. On heating it gradually darkens above 250° and decomposes rapidly about 297O (corr .).Found N=11-4; As=30*8. C,H,O,N,As (242.1) requires N = 11.6 ; As = 30.96 per cent 1376 DAXTER AND FARUHER 1 3-BENZODIAZOLEARSMIO 5 ; 51-Arseno-1 3 ; 1' 3t-benzodiazole (VIII). A solution of 5 grams of 1 :3-bcnzodiazole5-arsinic acid in 100 C.C. of water containing 1.2 grams of sodium hydroxide was treated with 50 grams of sodium hyposulphite and 11.5 grams of magnesium chloride dissolved in 300 C.C. of water and the mixture heated with stirring at 60° for two hours in an atmosphere of carbon dibxide by which time the precipitation of the yellow arsenobenzene was complete. After cooling the product was collected well washed with water suspended in 90 C.C. of water, dissolved by the addition of sufficient hydrochloric acid to form the dihydrochloride filtered from traces of undissolved matter and precipitated by pouring into an equal volume of concentrated hydrochloric acid.The hydrochloride was collected well washed with alcohol and ether and dried in a vacuum over sulphuric acid. The yield amounted t o 3.9 grams. 5 5t-14 rseno-l 3 1' 3/-benzodinzoZe is obtained as a bright yellow powder practically insoluble in water or the usual organic solvents, sparingly if a t all soluble in methyl alcohol containing hydro-chloric acid but fairly readily so in 50 per cent. acetic acid. The c7ihydrochloride forms a pale yellow powder soluble in water very sparingly so in methyl alcohol or ethyl alcohol and quite insoluble in acetone or ether. The aqueous solution reacts acid to litmus, but neutral to methyl-orange.After drying in a vacuum over-night it retains solvent equivalent to two molecules of water. Found N = l l . 3 ; As=30*2; C1=13*8. C,IH,oN,As.,2RC1,2TT,0 (493.1) requires N = 11 - 4 ; As= 30.4 ; C1= 14.4 per cent. 3 4-nincc.tylnniil.~op7~en?/larsinic Acid and 2-Methyl-l 3-henzo-diazole-5-a~sinic A cid (VI). Experiments in which diamiiiophenylarsinic acid was heated with glacial acetic acid with acetic anhydride with acetyl chloride and pyridine or with acetic anhydride in the presence of a trace of sulphuric acid were unsuccessful owing to the considerable amount of decomposition which occurred. It was found however that i f the acid dissolved in methyl alcohol was treated with a mixture of acetic acid and acetic anhydride and the methyl alcohol then removed as completely as possible by distillation on the water-bath, the resulting solution after heating to active boiling for four hours, gave an excellent yield of a product which from its composition, proved to be the diacetyl derivative of diaminophenylarsinic acid.3 ; 4-Diacetylami~iol~hen.yJn~sin i c acid crystallises from water i ACIDS AND THEIB REDUUTION PRODUOTS. 1317 which it is fairly readily soluble as a felted mass of fine needles which retain from 2 to 2.5 per cent. of solvent. Found loss at l l O o (two specimens) 2.6 2.2. I n dried material N=8*8; As=23*8. C,oH1,O,N,As (316.1) requires N,= 8.9 ; As = 23.7 per cent. As the product did not lose acetic acid when heated to 250°/ 20 mm. it was heated with ten times its weight of water in a sealed tube for four hours a t 130° when the desired reaction took place almost completely.2-Met hyl-1 3-benzodiazole-5-arsinic acid crystallises from water, in which it is somewhat sparingly soluble in minute needles contain-ing two and a-half molecules of water of crystallisation the) last half molecule being removed with difficulty a t l l O o but fairly readily a t 120O. After drying a t l l O o it darkens on heating above 250° and decomposes a t about 270° (corr.). Found loss a t 110°=15*5. I n dried material C = 37.3 ; H = 3.8 ; N = 10.7 ; As = 29.3. C,Hg0,N2As,2~H20 requires H,O = 15.0 per cent. C,H,O,N,As (256.1) requires C = 37.5 ; H = 3-5 ; N = 10.9 ; As = 29-3 per cent. 2 21-Dimethyl-5 51-ameno-1 3 11 3~-benzodiazole.This was prepared in the same manner as 5 5'-arseno-l 3 1' 31-benzodiazole which it resembles very closely in its properties, The dikydrochloride forms a pale yellow powder soluble in water, but almost completely insoluble in methyl alcohol ethyl alcohol, ether or acetone. After drying in a vacuum for twenty-four hours it retains solvent approximating t o three molecules of water. In aqueous solution it reacts acid to litmus but neutral to methyl-orange. Found N=10*3; C1=12*7; As=28*0. C~6H14N4As,,2HCl,3H20 (539.1) requires N = 10.4 ; C1= 13.1,; As = 27.8 per cent. 5 6-Diamino-i~-toZylarsin.ic A cid (I) A solution of 8-6 grams of 5-nitro-6-amino-~n-tolylarsinic acid in 75 C.C. of water and 6.2 C.C. of ION-sodium hydroxide was cooled t o -lo and treated in one operation while stirring vigorously, with 20.4 grams of sodium hyposulphite.When the reaction was complete as shown by the change in colour the solution was filtered treated with 8.1 C.C. of hydrochloric acid (D 1-12) and the re~ulting acid which separated in the course of a few minutes 1378 BAXTER AND FARGHflR I ~-BENZODIAZOSfEArRSMIC collected washed with water and recrystallised from watef in which it is sparingly soluble in the cold but fairly readily so on warming. It separates in colourless needles which contain between one and a-half and two molecules of water of crystallisation. It dissolve3 fairly readily in methyl alcohol or acetic acid but is very sparingly soluble in ether benzene or petroleum. I t s solution in dilute hydrochloric acid gives a characteristic deep violet color-atiofi with a drop of a dilute solution of potassium dichromate.It appears to decompose slowly on keeping. Loss a t l10°=11*2. C7H,,0,N,As,l~H20 requires 9.9 per cent. C,H,,03N,As,2H,0 requires 12.8 per cent. C,R,,O3N,As (246.1) requires N = 11.4 ; As=30.5 per cent. In dried material hT=11.3; As=30.4. 7-Methyl-1 2 3-benzotriazoEe-5-arsinic acid (11) was obtained by treating a solution of 3 4-diamino-m-tolylarsinic acid in dilute hydrochloric acid with a molecular proportion of sodium nitrite. It forms a crystalline powder very sparingly soluble in water but m p e readily so in alcohol. It crystallises from 50 per cent. alcohol in minute colourless glistening needles which are anhydrous and gradually decompose on heating above 280O.Found N = 16.6. C7H,0,N,As (257.1) requires N = 16.4 per cent. 4-Met h ybphennn t Ii rn ph encl(2ine-2 -urshic acid (111) was obtained by the condensation of molecular proportions of phenanthraquinone and 3 4-diamino-v~-tolylarsinic acid in acetic acid solution. It forms an amorphous yellow powder sparingly soluble in water and the usual organic solvents. With concentrated sulphuric aeid it develops an eosin-red colour with nitric acid a cherry-red and with hydrochloric acid an insoluble red compound. The colour is dis-charged in all three cases by dilution with water. It dissolves in dilute sodium carbonate or sodium hydroxide but a flocculent pre-cipitate of the sodium salt is thrown down on the addition of excess of the reagent.Found N = 6 * 4 . C2,H,,CJ3N,As (418.2) requires N = 6.7 per cent. 7-Methyl-1 3-he.nzodiazole-5-a~sinic Acid (IV). Two grams of 3 4-diamino-m-tolylarsinic acid were boiled wit,h 20 C.C. of glacial formic acid for six hours. The isolation of the product was carried out as in the case of 1 3-benzodiazole-5-arsinic acid save thatq the treatment with charcoal was found to be mor AOIDS AND THlIR REDUCTION PRODUCTB. 1379 effeative in acid solution. The air-aried product contains only traces of moisture and resembles the lower homalogue very closely in its properties. It dissolves sparingly in water and crystallism in minute prismatic needles which darken when heated above 280° and melt and decompose a t about 300° (mrr.), Found loss a t 110°=1.3.I n dried material N = 10.8 ; As = 29.4. C,H,O,N,As (256.1) requires N = 10.9 ; As= 29.3 per cent, 7 71-Dimethyl-5 51-arseno-1 3 If 31-benzodiazole. This was obtained by the reduction of the above acid with sodium hyposulphite in the manner previously described (p. 1376) as a pale yellow powder insoluble in water and the usual organic solvents but moderately soluble in acetic acid. The corresponding dihydrochloride forms a pale yellow granular powder soluble in water but insoluble in methyl alcohol ethyl alcohol ether or acetone. After drying for twenty-four hours under greatly reduced pressure it retained solvent corresponding with two molecules of water. Found N=10.6; As=28*6; C1=13*1. C,,H14N4As~,2HC1,2H,0 (521.1) requires N = 10.75 ; As= 28.8; Cl= 13.6 per cent.2 7-Dirnethyl-1 3-benzodiuzole-5-arsinic ,4 cid (VII). This acid was obtained by the prolonged action (six to eight hours) of boiling acetic acid on 5 6-diamino-m-tolylarsinic acid, the isolation being carried out in the usual manner. It is some-what soluble in boiling water but. sparingly so in cold and separates from the former in minute prismatic needles which contain two molecules of water of crystallisation. Found loss a t 110°=12'2. I n dried material N = 10.6 ; As = 27.9 27.5. C9Hl1O3N,As,2H,O requires 11.8 per cent. C,H,,QsN,As (270.1) requires N= 10.4 ; As= 27.75 per cent. 2 7 21 71-Tetranaethyl-5 5'-a.rseno-l 3 1/ 31-benzodiazoZe. The above acid was readily reduced to the corresponding arseno-benzene by sodium hyposulphite under the conditions previously employed. The urseno b enzene closely resembles those already described in its physical and chemical properties. The dihydro-chloride forms a pale yellow powder soluble in water but insoluble 3 $ 1380 COWABD AND WILSON THE EQUILIBRITTM in the usual organic solventa. After drying for twenty-four hours in a vacuum it retains solvent corresponding with approximately two molecules of water. Found N = 10.0 ; As = 27.3 ; Cl = 12.3. C,sH,sN4As,,2HC1,2H,0 (549.1) requires N = 10.2 ; As = 27.3 ; C1= 12.9 per cent. WELLUOME CREMICAL RESEARCH LABORATORIES, LONDON E.C. 1. [Received October 27th 1919.

ISSN:0368-1645    

DOI:10.1039/CT9191501372

出版商:RSC    

年代:1919

数据来源: RSC

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1380 COWABD AND WILSON THE EQUILIBRITTM CXXX .-The Equilibriuin between Carbon Hydrogen, and Methane. By HUBERT FRANK COWARD and STANLEY PIERCE WILSON. EQUILIBRIUM in the system C + 2H = CH is more suited to experi-mental study in one respect than the more commonly investigated high-temperature dissociations such as those of water carbon dioxide and hydrogen chloride. The proportion of methane in equilibrium with carbon and hydrogen diminishes with increasing temperature and hence when equilibrium is once established in the zone of reaction the cooling of the gaseous mixture on its removal from the carbon is not accompanied by an alteration in composition. It is a simple matter to ensure the absence of elementary carbon from any but the zone of measured temperature in which the equilibrium is attained.The synthesis of methane from its elements was achieved by Bone and Jerdan (T. 1897 71 41; 1901 79 1042) who found that when a stream of hydrogen was passed over carbon a t about 12000 the resulting gaseous mixture was hydrogen containing, roughly 1 per cent. of methane and no other hydrocarbon. This synthesis was confirmed by Bone and Coward (T. 1910 97 1219), who obtained an almost theoretical yield of methane from 0.0824 gram of carbon. The decomposition of methane into carbon and hydrogen with-out the permanent formation of any detectable amount of any other hydrocarbon was demonstrated by Bone and Coward (T.? 1908 03 1197). The above-mentioned experiments the forerunners of those now to be described indicated that the equilibrium mixture at 10O0-12OO0 contained roughly 1 per cent.of methane. The determination of equilibrium values in the system has bee BETWEEN CARBON HYDROGEN AND METHANE. 1381 attacked in two other quarters. Mayer and Altmayer (Ber., 1907 40 2134) carried out experiments with the aid of a catalyst a t temperatures ranging from 475O to 625O and declared that Bone and Jerdan's methane formed a t 1200° could not have arisen by the direct union of carbon and hydrogen. Pring (T. 1910, 97 498) found that a carbon rod lightly coated with platinum gave the value 0.55 per cent. of methane in equilibrium a t 1200O and 0.30 per cent. a t 1500O. Finally Pring and Fairlie (T. 1912, 101 91) conducted experiments under high pressures and assum-ing the law of mass action to hold down to atmospheric pressure, calculated the following equilibrium values : Amorphous carbon.Graphite. 1200" 0.36 per cent. of methane 0.24 per cent. of methane 1500" 0.21 , 9 0.07 , 7 9 The series of experiments the results of which are recorded below were conducted in essentially the same way as those described in detail by Bone and Coward (Zoc. c i t . ) except that the gaseous products of reaction were passed backwards and forwards over the carbon used until analysis showed the attainment of con-stancy of composition. The porcelain tube used to contain the reacting substances was as before set up coaxially with a wider one through which dry hydrogen was passed. The wider tube was heated electrically by means of a platinum wire resistance.The temperature of the inner tube was a t l l O O o constant within f 6 O , for a length of 10 to 15 cm. (in different furnaces) which was ample to contain the carbon used. The latter was held in a quartz boat or in an open quartz tube just narrow enough to pass freely into the porcelain tube. The temperature of the reaction zone was measured by means of a platinum and platinum-rhodium thermo-junction lying in the inner porcelain tube but encased in a thin quartz tube; one wire was insulated from the other by running i t through a fine quartz tube The thermo-couple was standardised by means of lithium chloride (m. p. 605O) antimony (m. p. 628O) an alloy of 80 per cent. of copper and 20 per cent. of tin (m. p. 738O), potassium carbonate (m. p. 8 3 5 O ) and electrolytic copper (in.p. 1 0 8 4 O ) . The gaseous products were collected direct from the furnace over a mixture of glycerol and water and before being re-passed into the furnace were dried by passage through a U-tube containing fragments of dry stick potassium hydroxide. The electrolytic hydrogen and the carbon were prepared as described by Bone and Coward. The methane was obtained by interaction between aluminium carbide and water washed with 3 Iff 1382 COWARD AND WILSON THE EQUILIBBmM ammoniacal cuprous chloride and sulphuric acid and collected after air had been displaced from the apparatus used for the prepara-tion. Hydrogen was then removed by liquefaction of the gas by means of liquid air and subsequent distillation. This procedure was proved to afford a satisfactory means for removing hydrogen from large volumes of methane.It cannot however be recom-mended for removing air from methane; methane must be pre-pared free from air in the firsb place unless large quantities of the gas are sacrificed during the fractional distillation after liquefaction. Experiments at 1 1 0 0 O . Synthesis of Methane.-Pure dry hydrogen was passed over 1 gram of carbon (ash 0.06 per cent.) a t the rate of about 1 litre per hour until from 5 to 6 litres of gas had been collected. A litre of this was measured and treated with palladium foil; the resulting concentrate was measured and analysed. The residual 4 to 5 litres of the gas once passed over carbon were re-passed a t the same rate as before and a sample was taken for analysis after concentration.The residual gas was then passed twice again over the heated carbon and again sampled for analysis. The resulh of the analysis calculated back to the unconcentrated mix-tures are (in percentages) : 1st passage. 2nd passage. 3rd passage. Carbon monoxide ......... 0.07 0.16 0.51 Methane ..................... 0.46 0-70 0.65 Nitrogen .................. 0.21 0.51 0.67 The experiment was repeated with the same sample of carbon: 1st passage. 2nd passage. 3rd passage. 4th passage. Carbon monoxide ...... 0.13 0.23 0.32 0.47 Methane .................. 0.43 0.68 0-64 0.67 Nitrogen ............... 0.26 0.49 0.56 0.69 Each experiment shows ail accumulation of methane during the first and second passage of the gas over the carbon but thence-forth the methane content remained practically constant.The nitrogen content is a measure of the leaking in of air during the experiment and analysis and in itself is unimportant. The carbon monoxide due iii part t o air and in part to a slight oxidising action by the glaze of the heated tube raises the question as to whether carbon monoxide may play some important part in the synthesis. It is just conceivable although contrary to the results of other investigations,* that carbon monoxide would be * Bone and Coward (T. 1908 93 1987) showed that when a mixture of 98 per cent. of hydrogen and 2 per cent. of carbon monoxide was passed through a porcelain tube at l l O O o in the absence of carbon no more than 0.03 per cent. of methane was present in the issuing gas BETWEIR OARBON HYDROBIN ARD METHANE.1383 redudd by the large excesa of hydrogen present to methane and water; the latter would regenerate carbon monoxide by reaction with carbon. If carbon monoxide acted as a catalyst in the synthesis then the two reactions-whatever they might be precisely -indicafred as A and B in the scheme would each be fast in comparison with the rate of the direct synthesis. Hence the accumulation of carbon monoxide exhibited in the experiments would only be possible provided that the reversion of methane to carbon and hydrogen were a very rapid reaction indeed when the methane was increased above about 0.7 per cent. That this is not the case is shown by experiments (below) on the decomposition of methane a t 1100O.The carbon monoxide is therefore produced as a side reaction between carbon and water vapour and plays no part in the formation of a t least the great proportion of the methane found. Decomposition of Methane.-A mixture containing 2.0 per cent. of methane and 98.0 per cent. of hydrogen was passed over the same sample of carbon eight times in all. The gas was sampled from time to time for analysis as before. The results were (in percentage) : 2nd passage. 5th passage. 8th passage. Carbon monoxide ......... 0.45 0.65 0.92 Methane 1.18 0.64 0.61 Nitrogen .................. 0.49 0-52 1.04 ..................... Here again the steady formation of carbon monoxide is evident, whilst the methane-content becomes constant a t or before the fifth passage and shows the attainment of equilibrium as in the synthetic experiments.Synthesis and Decomposition of Methane in the Presence of Nickel.-One gram of the pure carbon was intimately mixed with 1 gram of nickel oxide obtained by igniting the nitrate and the mixture heated to 300° in hydrogen until no more water vapour was formed and then raised gradually to 1100O in a current of hydrogen. The gas issuing from the tube was collected and re-passed three times; analysis showed i t to contain carbon mon-oxide = 1.29 methane = 0.61 nitrogen = 0.49 per cent. The decomposition of methane was conducted with a mixture containing 5.7 per cent. of methane and 94.3 per cent. of hydrogen. After five passages over the same mixture of carbon and nickel the gas contained carbon monoxide = 1.27 methane = 0.57 nitrogen 1384 COWARD AND WILSON THE EQUILIBRIUM 1-33 per cent.The increase in carbon monoxide is ascribed to the well-known difficulty of reducing finely divided nickel oxide com-pletely in hydrogen. C!oZZection of Results at llOOO.-The equilibrium values indicated above are: Carbon-Nickel Pure carbon. mixture. Synthesis .................. (i) 0.66 ; (ii) 0.67 0.61 Decomposition ............ 0.61 0.57 The meail value for the equilibrium amount of methane a t l l O O o is therefore 0.62 per cent. Experiments at 1000°. Synthesis of Methane.-The synthesis of methane from the highly purified carbon used in the experiments a t l l O O o proved to be too slow for the attainment of equilibrium a t 1000°. A less pure sample containing 0.5 per cent.of incombustible matter (silica) gave the following percentage results in an experiment con-ducted in a similar manner t o the corresponding experiment a t 1 1 0 0 O . 1st passage. 2nd passage. 3rd passage. 4th passage. Carbon monoxide ...... 0.05 0.10 0.17 0.28 Methane .................. 1.06 0.97 0.96 1.06 Nitrogen ............... 0.15 0.44 0-32 0.31 Equilibrium is evidently established from the first. The carbon monoxide-content is much lower than in the experiments a t llOOo, and although increasing throughout has no appreciable influence on the methane-content of the gases. The high value for nitrogen in the second experiment is most probably due to a small in-leakage of air during the concentration of tKe sample of gas analysed.An experiment with 1 gram of carbon containing 4 per cent. of platinum finely disseminated gave similar values : 1st passage. 2nd passage. Carbon monoxide ...... 0.07 0.16 Methane .................. 1.09 1.08 Nitrogen ............... 0.09 0.19 Decomposition of Methane.-A mixture containing 2.0 per cent. of methane and 98.0 per cent. of hydrogen gave in two separat'e experiments : First experiment. Second - experiment. 3rd passage. 5th passage. 2nd passage. Carbon monoxide ......... 0.19 0-41 0.55 Methane .................... 1-16 1.13 1.06 Nitrogen ............... 0.83 0.84 0.6 BETWEEN CARBON HYDROQEN AND MITHANE. 1386 Synthesis and Decom,position of Methane in Presence of Nickel. -The Bame sample of a carbon-nickel mixture as was used for the experiments a t l l O O o gave the following resulfs a t 10000.Starting with pure hydrogen samples of the unconcentrated gases showed on analysis an approximately constant content of methane from the second to the sixth passages. The methane then present was determined accurately after concentration of the mixture : carbon monoxide = 2.2 methane = 1.23 nitrogen =0*7 per cent. The decomposition of a 3.7 per cent. methane mixture gave in similar circumstances a constant content of methane from the third to the seventh passage. The final composition was carbon monoxide = 1 - 1 methane = 1 * 18 nitrogen = 1 *2 per cent. Collection of Reszdts at 1000°.-The equilibrium values indicated above are: Carbon (0.5 per Carbon-4 per cent. of ash). cent. of platinum.Carbon-Niokel. Synthesie . . . . . . . . 1.01 1.09 1.23 Decomposition ...... 1.13 ; 1.06 - 1.18 The mean value for the equilibrium amount of methane at 1000° is therefore 1.12 per cent. Experiments at 850°. The reaction between carbon and hydrogen proved to be too slow a t 850° for equilibrium t o be reached a t that temperature within a reasonable period of time in the apparatus employed. Experiments with the nickel-carbon mixture were more successful. Synthetic experiments with a slow stream of gas showed that from the sixth to the eighteenth passage the methane content remained constant according to the approximate results of analyses of un-concentrated samples of the mixture. The methane present a t the end of the eighteenth passage was determined after concentration : carbon monoxide =0.5 methane = 2.60 nitrogen= 1.3 per cent.Corresponding experiments on the decomposition of methane, carried out with a very slow stream of gas showed constant com-position between the second and fifth passages of the gas carbon monoxide = 0.2 methane = 2.4 nitrogen = 0.6 per cent. The equilibrium amount of methane a t 850° is therefore 2.5 per cent. Ex peyimen t .c n t 65 Oo . Hydrogen and methane were separately heated for prolonged periods in contact with carbon a t 650O. I n no case did the analyses of the resulting mixtures indicate that equilibrium could be attained within any reasonable period of time although attempt ,1386 TITB EQUJLIBRlVM BETWEEN CAftBON HYDROdEN ETC. were made to activate the carbon by means of finely divided nickel, (a) by grinding together nickel oxide and pure carbon and reducing in hydrogen at 300O; ( b ) by igniting a mixture obtained by evaporating nickel nitrate in the presence of excess of finely divided carbon and heating finally in hydrogen a t 300O; ( c ) by absorbing nickel nitrate in small pieces of porous porcelain igniting,' and depositing carbon by the prolonged passage of methane over the heated $did.It is not concluded that it is impossible to prepare carbon in a sufficiently adive state to attain equilibrium a t this temperature, but the claim of Mayer and Altmayer (Zoc. c i t . ) to have succeeded a t still lower temperatures is much weakened by the fact that they record no attempt to discover whether their gaseous products con-tained carbon monoxide.If as i t seems fair t o assume therefore, they overlooked carbon monoxide much of their so-called methane may have had no actual existence That there must have been carbon monoxide present in their products can scarcely be doubted in view of first the presence of large amounts of nitrogen and, secondly their use of somewhat large amounts of nickel which as is known is difficult to obtain free from nickel oxide by reduction of the latter a t a low temperature. Summary and Conclusions. The percentage composition of the equilibrium mixtures of methane and hydrogen in contact with amorphous carbon a t atmo-spheric pressure is 1100". 1000". 850". Methane ............ 0.6 1.1 3.5 Hydrogen ......... 99.4 98.9 97.5 These values will doubtless be capable of correlation by means of a thermodynamical equation when the specific heats of the three reacting substances are known together with the heat of formation of methane.Information as t o the specific heat of methane at high temperatures is however wanting a t present and the heat of combustion of methane has apparently not been determined with a . gas of sufficient purity. Calculation of the mean value of the latter between 850° and l l O O o from the experimental values for the equilibrium by means of the integrated form of the equation dlogK - Q _ _ _ - - _ _ _ d T R T2' givea the value 18,000 calories with a possible error of perhaps a couple of thousand calories THE PROPEETIES OF AMMONIUi3.l WITRATE. PABT I. 1887 In conclusion the authors wish to acknowledge their indebted-ness to assistance from the Government Grant Committee of the Royal Society to Measrs. Johnson Matthey and Co. for the loan of palladium and to Mr. T. R. Bradbury M.Sc. for valuable co-operation during one period of the work. FACULTY OF TECHNOLOGY, MANCHESTER UNIVERSITY. [Received November let 1919-

ISSN:0368-1645    

DOI:10.1039/CT9191501380

出版商:RSC    

年代:1919

数据来源: RSC

摘要:

THE PROPEETIES OF AMMONIUi3.l XQmggTE. PABT I. 1887 CXXXL-The Properties of Ammonium Nitrate. Part I. The Freezing Point and Transitim-temperatures. By REGINALD GEORGE EARLY and THOMAS MARTIN LOWRY. AMMONIUM nitrate is remarkable amongst salts both for its ewy fusibility and for the fact that the salt exists in not fewer than five crystalline forms the ranges of stability of which are marked by a series of four well-defined transition-temperatures. These different forms are conveniently named in the same way as the various forms of iron by Greek letters starting with the form which is stable a t the lowest temperatures. The crystalline forms of these modifications and the published data in reference t o the transition-temperatures are set out in the following table.TABLE r. Changes of State in Ammonium Nitrate. Form of Crystalline Change 1 t toliq. Liquid 127" 124-125" 124" 123*5-125.8', 1 6 to E { 125.6" 122.6" 126.2" 1244-L-125*6O 125.0' 125.2". 82*5-86" 81" 82*5-86.5" ) yt,o 6 {8722*S0 83.0" 86-" 85.S0 84" 84O, salt. system. of state. Transition temperature. 165-106° 165" 168O 1652O. -t Cubic 6 "Rhombohedra1 82-26". } B to (36" 31-35' 35O 31-35' 32.4" y TRhombic 1~ Rhombic I 32*2" 35*0° 35.0" 32.5". a Tetragonal a to - 16" - 18". * Or tetragonal. j- Or monosymmetric. I n view of the fundamental character of these constants and of the fact that the temperatures are all easily accessible a redeter-mination was made which has had the effect of fixing the freeain 1388 EARLY AND LOWRY THE PROPERTIES OF point with an estimated error of a few tenths of a degree at 169*6O, whilst the three upper transition points have been determined probably within a few hundredths of a degree at 125'2O 8 4 .2 O and 32*1°. I. Freezing Point at 169.6O. The difficulty of determining the correct melting point or freez-ing point of the salt and the fact that values much below the correct figure have so often been recorded arise from two charac-teristic properties of the salt namely the very great sensitiveness of the freezing point to the influence of small quantities of moisture (details in reference to which will be given in a later paper) and the tendency of the salt to retain its moisture even when attempts are made to remove it by somewhat drastic methods. On account of these properties of the salt we obtained in our earlier experiments a long series of concordant values for the freezing point which were afterwards proved to be more than 20 too lbw.These freezing points were determined by observing the arrest of temperature which occurred immediately after crystal-lisation had begun in about 40 grams of the molten salt the salt being stirred vigorously in a glass tube surrounded by an air-jacket in a bath of boiling water in order to prevent over-rapid cooling. The temperatures were measured by a thermometer graduated in fifths of a degree from Oo to 200° and recently re standardised t o 0 . 0 2 O a t the National Physical Laboratory ; this therrnometar was of the compensated type and the zero had remained constant within Om020 over a period of about ten years.Two auxiliary thermometers were used to record the temperatures of the " exposed column " of mercury inside and toutside the glass tube containing the salt. A series of freezing points determined in this way gave very concordant figures the corrected values being as follows: 166*95O 166*95O 166.90° 167*15O 167*0°. Mean 167.0°. This was the freezing point of a sample of the salt which had been recrystallised twice from water and dried first in a steam-oven during several hours and then during a day and a night over sulphuric acid in a desiccator exhausted with the help of a Gaede mercury pump. Doubt was first thrown on the correctness of these readings when it was discovered that a commercial sample of ammonium nitrate which had been dried and ground and then re-dried in the usual way in the laboratory gave a freezing point of 168*8O that is, nearly two degrees higher than our best laboratory specimen of the salt.As this higher freezing point could obviausly not be du AMMONIUM NITRATE. PART I. 1389 to impurities which would tend to lower the freezing point of the salt it aould only be attributed to the fact that the drying had been accompanied by a grinding which was absent in the labora-tory process of purification already described. In order to secure equally favourable conditions in the labora-tory the purified salt was heated in a steam-oven cooled and crushed in a mortar and left overnight in a vacuum desiccator, this cycle of operations being repeated three or four times.A series of freezing points of material prepared in this way gave the following figures : 169*16O 169*14* 169*04O 169-05O 169.06O. Mean 169.09O. The first two values were for ammonium nitrate from Hopkin and Williams twice recrystallised; the next two are for the same sample after storing for some weeks in a desiccator over calcium chloride; the last value is for a commercial sample of Norwegian nitrate recrystallised from water and dried by the method just described. These experiments showed that the freezing point of the salt was a t least 169.1° and might be as high as 169.2O in view of the first two values recorded in the list. A further improvement in the freezing point was effected by sealing up in an exhausted flask connected with a bulb of phos-phoric oxide some of the salt prepared as described above.The nitrate was shaken up repeatedly to expose a fresh surface and a t the end of a week the freezing point was determined. I n plotting the cooling curves for these determinations of the freezing point, two modifications were made in the practice previously adopted, namely (i) the outer water-jacket was kept a t 80° instead of looo in order to reduce the amount of water vapour round the apparatus; (ii) the stirring of the molten salt was stopped as soon as the temperature showed signs of becoming steady in order to reduce the loss of heat from the1 molten mass and so to prolong as much as p i b l e the horizontal arrest in the cooling curve. The freezing points determined in this way were 169.58O 169.55O.After three months' further drying by phosphoric oxide no further rise of the freezing point was produced but on the contrary it appeared to have receded slightly. The freezing point of the pure dry salt may therefore be placed a t 169.6". If the freezing point of the salt is expressed in whole numbers the figure 170° is more exact than any hitherto given since the experimental numbers recorded above are unlikely to be too high but may still be slightly below the true maximum figure. The drying of ammonium nitrate which is not really a difficul operation when dealing with the aolid salt cannot be effected merely by stronger heating since the salt which begins to decom-pose at 200° retains even a t this temperature the moisture pro-duced by decomposition.Thus two samples which had been heated to vigorous effervescence a t 225O,'froze a t 166.8O and 167*0° when cooled again to the freezing point of the molten salt. Again, a sample of the nitrate which melted originally a t 169O froze a t 1 6 5 ~ 5 ~ after being heated to 230° to 240° for a few minutes and at 163*4O after a second heating. The molten salt indeed appears to cling to its water quite as tenaciously as concentrakd sulphuric aaid and it is only by grinding the solid to a fine powder that complete drying is readily effected. 11. Transition-t emperature at 125.2O. When the molten salt is allowed to cool from the melting point in a tube provided with a thermometer dipping into the salt the highest of the transition-temperatures referred to in table I gives rise only to a retardation of cooling a t about 1 2 5 O .I n order to produce slower cooling and to secure if possible a horizontal arrest in the cooling curve in spite of the poor thermal conductivity of the salt the following method was adopted. About 90 grams of the purified salt were melted and poured into a small beaker which was lowered into the centre of a vacuum-jacketed cylinder A standard thermometer was inserted and the salt was allowed to set round the bulb. The stem of the thermometer passed through a large rubber stopper which served to close the top of the cylinder and 80 prevent loss of heat by convection to the outer air. The cylinder in its turn was immersed as far as possible in a bath of hot water stirred mechanically and maintained a t a fairly constant temperature by a flame.A Bunsen valve was provided for the escape of air by expansion from the cylinder whilst two auxiliary thermometers gave the temperatures of the " exposed column " of the standard thermometer. An attempt was made to retard the cooling still further by exhausting the air from the interior of the cylinder but this procedure did not appear to serve any useful purpose and was abandoned because of the additional trouble caused by the frotxng up of the salt to a porous mass during exhaustion. When observing the transition point a t 125O the temperature of the outer water-bath was maintained a t 95O but it was then generally allowed t o cool to 7 5 O in order to observe the further transition a t 84O to which reference is made in Section I11 below.The readings obtained in this way with air in the vacuum cylinder were 125*Z0 125*6O 125*8* 125.1" 125'6" 125*2O whils AMMONIUM NITRATE. PART I. 1301 with the air exhausted from the cylinder the readings were 125.2O and 125.40. ThO mean of the whole series is 125’4O. Zawidzki (Zeitsch. physiknl. Chern. 1903 47 721) obtained a mean value of 125*0° f o r arrests in the cooling of ammonium nitrate alone and mixed with silver nitrate. I n the case of the transition point a t 84O the poor conductivity of the salt may be compensated by stirring the powder in a small revolving drum heated externally by means of a vapour-jacket but this method cannot be employed a t the higher temperatures as the powder begins to cake into tough lumps.A much more effective method of securing good conduction during heating and cooling consists in stirring the powdered salt in a bath of liquid not differ-ing too widely from i t in density. The liquid selected for this purpose was a mixture of tribromoethane (D 2-62; b. p. 1 8 8 O ) and xylene (D 0.86; b. p. 140°) the density of ammonium nitrate a t this temperature being about 1-6. Good results were also obtained by using nitrobenzene (D 1.22; b. p. 211O). This method of improving the thermal conductivity has the advantage that the thermometric measurements can be made within a maximum of accuracy. The procedure was as follows. The liquid was made into a (‘mush” with powdered ammonium nitrate in a glass tube 20 cm. long and 3.7 cm. in diameter provided with a stirrer and a split cork to carry a standard thermometer; as before two auxiliary thermometers were used to record the temperature of the The glass tube was surrounded by a metal cylinder 20 cm.long and 5 cm. in diameter closed at the top with cotton wool and forming a narrow air-jacket round the tube. This cylinder was supported by a bung in the axis of a larger metal cylinder 32 cm. long and 10 cm. in diameter con-taining paraffin of high boiling point to serve as a heating or cool-ing bath and provided also with a thermometer. I n plotting a heating or cooling curve the paraffin bath was adjusted by hand t o a temperature a few degrees above or below the transition point. The ((mush” in the inner tube was stirred a t intervals of thirty seconds to two minutes according to the velocity of heating or cooling and the thermometer was read ‘immediately after.The arrest-temperatures observed in this way during heating and cooling together with the conditions under which the heating and cooling were carried out are shown in table 11. exposed column ” of mercury 1392 EARLY AND LOWRY THE PROPERTIES OF TABLE 11. ArTest-teirLperutzcres of Anz?nonium iVitl.de Suspe~Zed in a bath of Liquid. (corrected). outer bath. Composition of liquid. Arrest-points Temperature of Heating 126.21 130" Nitrobenzene. 128.16 140 Xylene and ethylene bromide. 125.38 140 Xylene and ethylene bromide. 125.22 130 Xylene and tribromoethane. 125.25 130 Nitrobenzene. Mean 126.24" Cooling 126-27" 120" Xylene and ethylene bromide.126.21 120 Xylene and ethylone bromide. 125.09 120 Xylene and tribromoethane. Mean 125.19' These values may be compared with those givea by Vogt (Physi-I d . Zeitsch. 1911 12 ll29) who obtained with a dilatometer the upper and lower limits 125.25O and 125*13O mean 125-2O. The eight values now recorded range from 125'09O to 125.38O'. The mean of the five readings obtained by heating the salt is 125*24O whilst the mean of the three readings obtained by cooling is 125-19O. The general mean of all the readings is 125.22O and this is probably the best value for the transition-temperature. I n view however of the difficulty of reading the temperatures to O*0lo we prefer to give the transition-temperature to a tenth of a degree at 125.2". 111. Transition-temperature at 8 4 .2 O . The transition a t 84O is accomplished by an abrupt expansion which frequently made itself manifest by fracturing the glass vessel in which a cast sample of nitrate was being cooled. There is how-ever a very strong tendency for over-cooling to occur and in many cases the salt was cooled to 32O without any indication that this change of crystalline form had taken place. . I n order to ensure the conversion of the 6- into the y-form i t was necessary to inocu-late the surface of the block with particles of the salt which had been heated to 60° and to assist the conversion by scratching the surface with a sharply pointed glass rod. Even under these con-ditions the over-cooling was usually very pronounced ; the change of crystalline state (even in the vacuum-jacketed apparatus used successfully to record the change of state a t 125O) only began when the temperature had fallen to 82O or below and the latent heat was then not sufficient t o restore the temperature to the transition point or to produce a horizontal arrest in the cooling curve.I AMMONIUM NITRATE. PART I. 1393 every case therefore the transition merely produced a sinuous curve the highest temperature recorded for the maximum on this curve being 8 2 - 7 O . The first well-marked arrest of temperature a t the transition point was observed when heating a sample of the powdered salt in a machine in which the nitrate was thrown over and over in such a way that the bulb of a thermometer was constantly bathed in the falling nitrate.A small machine in which this principle was embodied gave well-defined arrests on heating a t 84*3O 84*Z0 84'3O, and 8 3 ' 3 O the first three values being concordant within one-tenth of a degree. It is remarkable that this transition-temperature, which was found to be so exceptionally difficult to locate by the methods first employed proved to be by far the easiest of the transition points to determine exactly when once the proper con-ditions were established. The most f avourable conditions for determining this transition point were those already described in connexion with the change of state a t 125*2O namely to compensate fpr the lack of thermal conductivity by stirring the powdered salt in a bath of liquid of almost equal density. The liquid used for this series of experi-ments was a mixture of ethylene bromide (D 2.18; b.p. 131O) and xylene (D 0.86; b. p. 140O). Special precautions were again needed to prevent " lag " in the change of state; the '' mush" was therefore inoculated during heating with crystals heated to looo, and during cooling with crystals heated to 60°. The inoculation was carried out just before reaching the transition point and was followed by gentle stirring. The arrest points recorded in this way which are lower than the mean of Zawidzki's arrests a t 85'4O, but higher than the dilatometric readings 82.16O to 82.36O of Vogt, are set out in the following table: TABLE 111. Arrest-temperatures of Ammonium Nitrate suspended in a Bath of Liquid. (corrected). outer bath. Composition of liquid.84.16 90 Y 9 9 9 84.25 90 9 9 Y Y 84-20 90 9 9 9 9 Arrest-points Temperature of Heating 84.20" 90" Xylene and ethylene bromide. Mean 84.20' Cooling 84.28" Falling slowly. Xylene and ethylene bromide. 84.13 9 9 7 9 9 9 9 9 84-18 9 7 9 9 9 9 9 9 Mean 84.20 1804 EA.RLY AND LOWRY THE PROPERTIES OF In this case the means of the four values obtained during heat-ing and of the three values obtained during cooling are identical. The transition point may therefore be fixed probably within a few hundredths of a degree a t 84*20° or may be given to a tenth of a degree as 84.2". I V . TransitiolL-temperature at 32-10. (a) Heating and Cooling Curves.-The transition-temperature in the neighbourhood of 32O is more easily observed than either of those occurring a t a higher temperature since a prolonged arrest in the neighbourhood of 3 2 O can always be detected when the nitrate is cooled through this temperature; indeed both this arrest and that a t 8 4 O are constantly encountered when the nitrate is handled commercially.I n spite of this fact exceptional difficulty is experienced in securing an exact determination of this transition-temperature. This difficulty was ultimately traced to the fact that over a range of about a quarter of a degree on either side of the true transition point the velocity of change of state is so slight as to be practically imperceptible with the result that the con-version usually takes place a t a temperature definitely below the transition point on cooling and a t a temperature definitely above it on heating.Large numbers of cooling curves were plotted in order to deter-mine this transition-temperature accurately. I n some experiments only a sinuous cooling curve was obtained but in others well-marked horizontal arrests were recorded a t the following tempera-tures : Cast blocks ............... 31*9O 32'0° 32*0°. Pressed blocks ............ 31*7O 31*5O 31*4O 3 1 ~ 8 ~ . Loose crystals ............ 31-6O. Horizontal arrests were also1 recorded sometimes a t lower temperatures for example 31.0° 29-5O 29-6O. Arrests during heating were always a t a higher temperature and the curves wen3 generally of a sinuous form rising to a maximum value before falling again to a minimum approximating to the transition-temperature of the salt.These minima in the heating curves were observed a t the following temperatures : Cast blocks ........................ 33.2O. Pressed blocks ..................... 33*8O 33*1° 33'2O. In one experiment in which the sample had become very much over-heated before the change of state set in a still higher read-ing was obtained a t 35*3O (compare Zawidzki Zoc. cit. who obtaine AMMONIUM NITRATE. PART I. 1398 tan average of 35-0° for seven arrests in heating ammonium nitrate alm0 and mixed with silver nitrate). I n view of the fact that the horizontal arrests or maxima in the cooling curves were always very much more fully developed than the minima in the heating curves it was believed that the former could be assumed to give a correct value for the transition-tempera-ture which was located provisionally a t 32O; actually however the data now quoted can only be used to prove that the transition-temperature lies within certain limits for example between 32-0° and 33*1° and i t was not possible to secure an absolute deter-mination by this method in view of the fact that in no case were the arrests on heating and cooling within one degree of each other.Attempts to secure more accurate readings of these transition-temperatures by using larger quantities of nitrate up to a kilo-gram resulted in failure the conductivity of the salt being so low as to prevent the effective flow of heat from one part of the mms to another; better results were indeed always obtained by heat,-ing or cooling much smaller quantities of ammonium nitrate, insulated as carefully as possible for example in a vacuum vessel, in order to reduce the rate of heating or cooling t o a minimum.A distinct improvement in the heating curves was however, obtained as in the case of the transition a t 84O by stirring about 700 grams of powdered nitrate in a small drum surrounded by a steam-jacket in such a way as to produce a constant flow of nitrate past the thermometer. By using this method the following arrests in the heating curves were recorded: 32.6O 32.75O 32.7O 32*9O 33*0°. These readings are definitely lower than those recorded previously when cast or pressed blocks of the nitrate were heated but no improvement could be effected in the cooling curves so that there still remained a gap of about 0*6O between the highest arrest on cooling and the lowest arrest on heating, The method of stirring the powdered nitrate in a bath of liquid of equal density which had proved so successful a t the higher temperatures was a complete failure when applied to the deter-mination of the transition-temperature at 32O.The temperature recorded on the thermometer immersed in the liquid frequently failed to show any arrest a t all. Even after inoculation the cool-ing curves were extremely erratic and only on two occasions were satisfactory arrests observed a t 31-8O and a t 31.6O; on no occasion was an arrest oherved in the heating curve. It may be noted, however that Muller (Zeitsch. physikd. Chem. 1899 31 354) obtained a satisfactory arrest a t 32*20 by cooling from 80° a mix 1396 EARLY AND LOWRY THE PROPERTIES OF ture of 100 grams of ammonium nitrate with 15 to 20 grams of water and that this temperature lies within 0*lo of our final value for this transition point.(b) Experiments with the BiZatometer.-The unexpected difficul-ties which were encountered in trying to determine the exact posi-tion of the transition-temperature a t 3 2 - 1 O can be traced to the relative slowness with which the change of state takes place immedi-ately above or below the transition point. even when assisted by inoculation as contrasted with the much greater velocity of the changes a t 84O and 1 2 5 O . This is in accordance with the general rule that changes of this character become more and more sluggish as the temperature falls, by reason of the decreasing mobility of the molecules and the increasing resistance which the rigidity of the material opposes to molecular rearrangement.Under such conditions the thermal method becomes difficult or impossible and i t is usually necessary to fall back on some method of determining the transition-tempera-ture in which ample time can be allowed for the change of state to reveal itself. I n the case of ammonium nitrate the most promising method was to follow by means of a dilatometer,* the expansion or contraction which accompanies the change of state, instead of relying on the absorption or liberation of latent heat to produce an arrest in the heating or cooling of the salt. Experiments which were made 'on these lines gave us our first trustworthy value for this transition-temperature and also pro-vided valuable information as to the velocity of the change in the irrmediate neighboarhood of this point'.The solid used in the dilatometer was made by fusing pure dry ammonium nitrate pouring into a mortar breaking the cast lump into pieces about 0.3 cm. in diameter and sieving to free it from dust. This form of the salt was used in order to secure good thermal contact between the solid and liquid and a t the same time t o avoid the risk of fracturing the bulb by the sudden expansion of a closely-packed powder. The dilatometer held about 60 c.c. and the bulb was sealed off after filling about three-quarters full with fragments of nitrate. The liquid was a paraffin of high boiling point which had been treated with concentrated sulphuric acid to free i t from olefines and then dried over metallic sodium; it was introduced in an air-free condition by making use of the apparatus shown in Fig.1 of a paper by Wade and Merriman (T. 1912 101 2430). I n a dilatometer filled in this way there are a t the transition-* Compare vctn't Hoff Zeitsch. ph?jsikat. Chem. 1895 17 130 and '' Vorlesungen,' 1898 vol. i. p. 18 AMMONIUM NITRATE PART I. 1397 temperature two alternative positions fox the meniscus ; one which may be called TTB is the position when all the nitrate is still in the stable low temperature or &form; the second which may be called Vy is the position to which the menisous rises when all the nitrate has been converted (without change of temperature) into the lighter y-form.The position 17s can be determined by heat+ ing the dilatometer from the atmospheric temperature to a point just below the transition point; Try can be determined by heating the dilatometer say to 50° until the whole of the nitrate has passed into the y-form and then cooling it to a point immediately above the transition point at 32O; very little extrapolation is then required to give the exact position of these two points. I n using this method to determine the transition-temperature of ammonium nitrate it is esseiitial that the dilatometer should contain both forms of the salt in order that change of state may take place quite readily in either direction. This condition was secured by heating the dilatometer t o 50° when the meniscus rose to a point well above Vy; on immersing the dilatometer in cold water part of the salt reverted from the y- to the &form a change which was revealed a t once by the appearance of white patches on the lumps of nitrate; this change took place long before the contents of the bulb as a whole had time to cool to the transition-temperature.Having made sure in this way that the nitrate in the bulb contained a substantial proportion both of the 8- and of the y-forms of the salt it was possible to find a range of tempera-tures (just above the transition point) a t which the meniscus tended to settle down in the neighbourhood of Vy by reason of the complete conversion of the contents to the y-form and a range of temperatures (just below the transition point) a t which the meniscus would settle down in the neighbourhood of V S owing to the complete conversion of the nitrate to the &form.The actual behaviour of the dilatometer when heated a t different temperatures after the preliminary treatment described in the preceding paragraph was as follows : (i) When the dilatometer was immersed in a thermostat a t 32*2O the meniscus settled down very quickly to a definite posi-tion between VB and Vy and during the course of two hours showed no tendency to rise to Vy or to fall towards VB. The change of state appeared in fact to be arrested a t the point to which it had been brought by the more drastic preliminary treat-ment of the salt. A precisely similar behaviour was observed a t 32*1° 32.0° 31*9O 31*8O and 31-7O.(ii) When the temperature of the thermostat was reduced to 31.6O the meniscus for the first time began to show a tendency t 1308 EARLY AND LOWRY THE PROPERTIES OF fall towards VB although the change a t this temperature was so slow as to be almost imperceptible; on repeating the experiment with the thermostat set to 31*5* however a definite movement towards VCrp could be seen. These two temperatures are therefore definitely below the transition point of the nitrate. (iii) When on the other hand the temperature of the thermostat was raised to 32.3O a definite but exceedingly slow upward movement of the meniscus towards Vy was observed and a more pronounced movement when the temperature was raised to 3 2 ~ 7 ~ . These experiments on account of the extreme sluggishness of the change of state in the immediate neighbourhood of the transi-tion point failed to fix the exact position of this temperature, although they served to locate it between 31-6* and 32.3O.These two limits agree quite closely with those arrived a t from a study of the cooling and heating curves which had shown arrests below and above tw40 corresponding limits a t 32*0° and 32.6O; rather clmer limits were recorded by Vogt. who observed equal and o p p site slow changes of volume a t 32*40° and 32.62O. (G) Quantitatiue Experiments with the Di1atometer.-In order to determine the exact position of the transition point a series of quantitative experiments was made on the velocity of the change of state as shown by observations with the dilatometer a t tempera-tures above and below the transition-temperature.The preliminary treatment of the nitrate was much the same as in the previous experiments and was carried out as follows. The two extreme positions of the meniscus a t the transition-temperature were first located on the scale of the dilatometer as follows : V/3=20 cm.; Vvy=48 cm. Starting with cold nitrate in the 8-form the dilatometer was next immersed in a bath of water a t 34O in order to initiate the change from the r& to the y-form. A t this temperature (on account of thermal expansion) the two extreme positions of the meniscus would be about 22 and 50 cni.; when therefore the meniscus had risen to about 36 cm. on the scale of the dilatometer it was clear that the two forms were present in roughly equal quantities whilst the thermal conditions of the dilatometer and its contents had not been aeriously disturbed.After this preliminary treatment the dilatometer was immersed in a thermostat set to a definite temperature just above or below the transition point. Thermal equilibrium was quickly estab-lished and exact measurements could be made of the rate of movement of the meniscus consequent on the change of state fro AMMONIUM NITRATE PART I. 1380 to y. or vice versa; the uniformity of this movement O V ~ T con-siderable periods of time was such as to justify the view thafi the rates recorded in heavy type in table IV are definite physical con-stants of the change of state over the range of temperatures from 3 1 . 2 O to 33-0°. The readings represent the movement in mm.of the meniscus during successive intervals of ten minutes. The mean steady velocities (obtained by averaging the numbers which are shown in heavy type in table IV) are recorded in similar unite. 65 50 45 40 35 30 25 20 15 10 5 31 The final column shows the percentage of the salt which would undergo change of state in one hour as deduced from the volume changes recorded in the previous columns. The thermometer readings are given to tenths of a degree but are all subject to a correction of +0*03O. It is remarkable that a t temperatures a whole degree above or below the transition point nearly an hour is required to change half of the material from the B- to the y-form, or vice versa. When all these velocities are plotted out in a diagram as i 1400 EARLY AND LOWRY THE PROPERTIES OF the figure a curve is obtained which is almost symmetrical 011 either side.According to the well-known principle which is used in determining the ‘( critical volume ” of liquids and vapours the transition-temperature may be determined accurately by ruling a series of horizontal lines across the velocity diagram reading off the temperatures given by the intercepts of these lines with the curve and working out an average value for each pair of tempera-tures. These averages as set out in table V below are as follows: 32.09O 32*11” 32.13O 32*13O 32*12O 32-11° 32.11O. The mean value 32*11° is much more accurate than any of the values deduced either from arrests in the heating and cooling curves of the salt or from merely qualitative observations with the dilatometer and is probably correct to within a few hundredths of a degree.For practical use the second decimal may be omitted and the transition-temperature given to onetenth of a degree as 32.1”. (d) Form of Velocity Cu.rue.*-Whilst the use of the dilatometer in determining transition-temperatures has been a well-known standard method for more than twenty years and has been applied repeatedly when the change of state is too slow to be followed by the thermal method (see especially E. Cohen’sr experiments on the allotropy of metals); the complete velocity curve shown in the figure does not appear to have been plotted in any of these cases. The most striking feature of this curve is its complete symmetry as proved by the constancy of the temperatures shown in the last column of table V which only vary over a range of +0*02O.This complete symmetry which could not be predicted makes it possible t~ locate the transition temperature very accurately and i t is doubtful if any other method is capable of giving equally exact results; for comparison it may be noted that the use ol the “recti-linear diameter ” t o determine critical volumes depends on the existence ol a skew-symmetry oiily in tlie curve of specific volumes for the liquid and saturated vapour. In view of the regularity of the curve it is of interest t o inquire into its mathematical form. The data now recorded can be expresed by the equation f ( t - to) = log,*( .\/ u+ l), where t =temperature, to = transition temperatur,e, v =velocity of change (percentage changed per hour), Added 6/12/19 AMMONIUM NITRATE.PART r. 1401 The agreement of this formula with the experimental results is shown in the following table : 21. 1 2.5 5 10 20 30 40 -J(t - to) obs. 0.33" 0.4 1 0.50 0.62 0.72 0-79 0.86 r t ( t - to). calc . 0-30" 0.41 0.51 0.62 0.74 0.81 0.86 (e) 9no77zctlies i n the Heating mid Coolitig C'urves.-At a very early period in the investigation it was noticed with some surprise that very steady arrests of heating and cooling might occur a t temperatures which were obviously not exact transition points, since they were scattered over a considerable range on either side of a mean value which could be regarded provisionally as the correct transition-temperature.This was an anomaly for which no explanation could be suggested a t the time. General experience in such matters has shown that when deal-ing with rapid transitions such as are observed in iron a t 895O and 766O the arrest in the cooling curve becomes blurred when-ever liberation of latent heat fails to compensate for loss of heat by cooling; a lowered arrest point is therefore nearly always revealed by its sinuous character. I n dealing with ammonium nitrate many sinuous arrests have been observed a t temperatures either below or above the real transition point but in several cases the arrests recorded a t these lower or higher temperatures were perfectly sharp. These observations can all be interpreted in the light of the data now given for the velocity of change of state.Thus in the first place i t is obvious that since this velocity of change is imper-ceptible from 31*B0 to 32*3O there can be no marked liberation or absorption of latent heat and no arrest in the cooling or heating between these limits. Outside these limits of temperature the flow of latent heat may arrest the cooling or heating but the con-ditions are such that the arrest point cannot be regarded as a fixed temperature but must be considered as a variable tempera-ture depending directly on the rate of heating or cooling. Thus, taking the latent heat of the transition as 5.02 calories per gram, and the velocities of change of state as recorded in table IV the temperatures of arrest for the different rates of coding and heating shown in column 2 of table V are given in columns 3 and 4 of that table TABLE IV.Velocity of Change of State above and Motion of meniscus in mm. per 10 minutes. Temperature.* M A 31.2" 31-4 31.5 31.6 31.7 31.8 32.0 32-2 32.3 32.4 32.5 32.6 32.7 32.8 33.0 33.8 8.4 11.0 1.4 0.8 10.5 -6.5 7-0 -13.7 4-8 9.0 t8-4 21.7 25.0 20.0 9.0 8-3 2.1 2-2 1.2 1.9 0.8 0-4 1.1 0.3 0.1 0.1 0-5 0.0 1.2 0.3 1-7 1.6 4.7 3.0 7.6 3.5 7.9 7.0 ts-0 77.7 22.9 23.8 21.0 -8-2 8.1 3-0 3.7 1.9 2.0 0-4 0.7 0.2 0.0 0.0 0.0 0.0 0.0 0.7 0.3 0.1 0.4 2.3 2-2 2.3 2.8 5.7 5-0 t7.2 t7.2 M.3 24.4 -7.5 4.2 2.2 0.7 0.0 0-0 0.0 0-2 0.2 1-7 2-5 4.8 t7-2 24.1 - - -7.5 - -4-5 4.5 4.3 2.5 2-4 2.4 0-6 0.7 0.6 0.0 0.0 0.0 0.0 0.0 0.0 0.0 0-1 0.0 0.3 0.2 0.2 0.2 0.3 -1.5 1.5 1.4 2.0 1-7 1.9 4-0 4.0 4-0 @-8 f6-8 fl-4 - - ---4.7 2.4 0.6 0.0 0.0 0.0 0.2 -1.4 1-7 4-0 fl.3 -* Correction at each temperature = + 0.03" AMMONIUM NITRATE.PABT 1. 1403 TABLE V. Velocity of Change of State of Ammonium Nitrate at Diferent Temperatures. Per cent. changed per hour. 1 2.5 5 10 20 30 40 Rate of cooling or heating. Calories per gram per hour. 0.050 0.126 0.251 0-602 1.00 1.51 2.01 Temperature (arrest-points). - Cooling. Heating. 31.76' 32.42' 31.69 32.62 31-63 32.63 31.61 32.75 3140 32.84 31.32 32-90 31.25 32.96 Average temperature.32.09' 32.1 1 32.13 32-13 32.12 32-1 1 32-11 Mean 32.11" The arrest points shown in the table are the temperatures a t which the latent heat liberated by a change of state proceeding with known velocity would exactly balance the heat gained or lost by radiation conduction etc. Imperfect conduction of heat in the mass may displace these temperatures still further from the true transition point and a t the same time destroy the sharpness of the arrests. The latter phenomenon is observed also in the transitions a t 84O and a t 125O where sinuous arrests a t abnormally low. temperatures are frequently observed but the velocity of change of state a t these two points is so much greater that a pro-longed or " horizontal '' arrest has never been observed except a t a temperature agreeing very closely with one or other of these transition points.Summary. (1) The freezing point of ammonium nitrate is very sensitive to the influence of traces of moisture; by careful purification and drying it has been raised to 169-6"-(2) The highest of the transition points has been determined from the arrests of temperature on heating and cooling the aalt when suspended in a liquid of similar density. The temperatures recorded were on heating 125'24O; on cooling 125-19O; mean value 125.2". (3) The second transition point was determined by the same method the observed temperatures being on heating 84'20O; an cooling 84.20° ; mean value 84.2". (4) The lower transition point cannot be determined in this way as the change of state is too slow to be detected over the range fzom 31'8O to 32'2O. It was determined dilatolmetrically by measuring the rate of change over a range of temperatures; the VOL. axv. 3 1404 KING THE PRODUCTION OF curve of velocities proved to be symmetrical and the transition-temperature was therefore found by taking the average of pairs of temperatures a t which the change proceeded with equal velocities in opposite directions. I n this way the transition-temperature was fixed a t 32.1". (5) The form of the velocity curve for the change of state can be expressed by the simple empirical equation f ( t - to) = k log, ( 4; + 1). (6) The arrest points due to the change of state a t 3 2 ' 1 O never coincide with this temperature the arrest point being determined b y t h e rate of loss or gain of heat. A table is given showing the temperatures of arrest on oooling or heating a t fixed rates expressed in calories per gram per hour. GUY'S HOSPITAL, LONDON S.E. 1. [Recci.ved October 13th. 1919.

ISSN:0368-1645    

DOI:10.1039/CT9191501387

出版商:RSC    

年代:1919

数据来源: RSC