摘要:

2808 CLOUGH OPTICAL ROTATORY POWERS AND RELATIVE CCCLXXXVI1.-The Relationship between the Opticat Rotatory Powers and the Relative Configurations of Optically Active Compounds. Part I I . The Relative Configurations of the Optically Active Mandelic Acids and P-Phenyl-lactic Acids. By GEORGE WJLLIAM CLOUGH. IN Part I * (J. 1918 113 526) it was shown that t,he formation of similar derivatives from the configuratively related compounds, Z-lactic acid Z-glyceric acid d-malic acid and d-tartaric acid was usually accompanied by changes of the same character in their optical rotatory powers. It was therefore assumed that as a rule, the introduction of the same substituent into similarly constituted, optically active compounds possessing the same relative con-figurations produced alterations of the same character in their optical rotatory powers.The somewhat vague term “ similarly constituted compounds ” was used in this connexion since it was not (and still is not) possible to state precisely the extent of the applicability of the rule. It applies to the members of such homo-logous series of optically active compounds as the aliphatic normal secondary alcohols and the simple a-hydroxy-acids. Indeed the higher members of a homologous series being derivatives from the lower members the rule would indicate that the optical rotatory powers of those optically active members of a series which are configuratively related would lie approximately on a curve. But not only do regularities occur in the optical rotatory powers of members of homologous series they are evident (with few excep-tions) in those of correspondmg derivatives from the optically active mono- and di-hydroxy-propionic and -succinic acids.The application of the principle to d- a-hydroxybutyric acid left no doubt that this compound possessed the same configuration it8 E-lactic acid but the same confidence could not be felt in the con-clusions which were drawn from the optical rotatory powers of compounds in which a phenyl or benzyl group was attached to the asymmetric carbon atom. In an endeavour to fix definitely the Errata in Part I Page 532 line 5 from bottom f o r “ X=Bz Y=Me ” read ‘‘ X=Bz Y=Et.” Page 532 last line for “ -123.6 ” read “ -247.” Page 640 line 3 for “ +50-6” ” read “ - 5 0 - 6 O . ” Since this paper was submitted for publication the author’s attention haa been directed to a paper by Freudenberg and Markert entitled “ Die Konfigur-ation der Mandekaure ” (Ber.1925,58 1753). These authors have employed the method indicated in Part I and have confirmed the suggestion made therein concerning the configurations of the optically active mandelic acide CONFIGURATIONS OF OPTICALLY ACTIVE COMPOIMDS. 2809 confipllrations of the optically active forms of mandelic and p-phenyl-lactic acids the optical rotatory powers of these compounds and some of their derivatives are considered in the present communi-cation. The measurement by Freudenberg Brauns and Siege1 (Ber. 1923 56 199) of the speciiic rotation of the amide of I-hexa-hydromandelic acid (prepared from I-maadelic acid) enabled those workera to confirm the suggestion that I-mandelic acid is a “ d ,’-a-hydroxy-acid (Part I p.534). (For compounds containing one asymmetric carbon atom it appears to the present author desirable to retain the prefixes d- and I- with their conventional significance, but to denote their relative configurations by the prefixes “ d ”- and “ I ”-; Ioc. cit. p. 534.) The optical rotatory powers of other derivatives from I-hexahydromandelic acid and from I-mandelic acid definitely indicate that these acids belong to the “ d ”- aeries of a-hydroxy-acids. This result is of importance in that it enables us to determine the configurations of I-benzoin the related optically active glycols and of amygdalin with reference to that of d-tartark acid.The experimental data quoted in tables I-V are also in the author’s opinion sufficient to justify the allocation of d- p-phenyl-lactic acid to the “ d ” - series of a-hydroxy-acids. If this con-clusion is accepted it is possible to assign configurations to the glycols derived from this acid and also to the four optically active phenylglyceric acids (provided the assumption is made that cis-addition of hydroxyl occurs on oxidation of the cimamic acids; compare Berner and Riiber Ber. 1921 54 1945). The molecular rotations of the standard “ d ”- acids in aqueous solution are lower * than those of the corresponding sodium (pow-ium or ammonium) salts (Table I). The same regularity is observed in the molecular rotations of I-hexahydromandelic acid I-mandelic acid d- P-phenyl-lactic acid and their salts.The molecular rotations of the lower esters of the same acids increase as the molecular weights of the esters increase (Table PI). The measurements by Wood and his collaborators of the optical rotatory powers and the optical dispersive powers of the optically active alkyl lactates and hexahydromandelates at various tem-peratures are especially valuable in comexion with the subject of this investigation. The molecular rotations of the amides of the same acids compared with those of the methyl esters further illustrate the relation of Throughout this paper an ’‘ incresse ” of rotation denotes an jncreaae inthe numerical value of a destrorotation or a decrease in the numerical value of a laevorotation. For the sake of clearness the tabulated valnes are given for the “d77-forms although in some ~ ~ J B S the measuremente wem actually made on the anentiomorphous forms 2810 CLOUGH OPTICAL ROTATORY POWERS AND RELATIVE Z-mandelic acid and of d-p-phenyl-lactic acid to the standard " d "-a-hydroxy-acids (Table 111).TABLE I. The Molecular Rotcations ([MID) of Some O p t i d l y Active Or-Hydroxp acids in Aqueous Xolution. Z-Lactic acid. 2-Glyceric acid. d-Mdic aid. d-Tartaric acid. Z-Heuahydro -Z-Mandelic acid. mandelic acid. d-p-Phenyl-lactic acid. Acid. - 2" (c = 5) - 2.3 (c = 20) ( c = 5 ) (c = 5 ) + 3.0 +21-3 - 21.3 - 240 (C = 1.6) (C = 2.7) (in alcohol) + 38.0 Sodium salt. + 13.2" + 20-6 + 16-2 + 59.9 + 13.6 + - 179 + 79.5 Potassium salt.References. +13.7" Purdie and Walker, J. 1895 67 630. + 23-7 Frankland and Apple-yard J. 1893 63, 311. 2268. 1073. +15-5 Stubbs J. 1911 99, +64-4 Landolt Ber. 1873,6, + Wood and Comley J., (ammonium) 1924 125 2639. -+ 75-7 Clough. TABLE 11. The dldecuhr Rotations of the Lower Esters of Some a-Hydroxy-acid-s. Z-Lactates.* [M1:,5" Z-Glycsrates. -f [MI::" d-Ma1ates.t [ M l y d-Tartrates. t [MI:" Z-Hexahydro- [MI?" Z-Mmdelates. [MI:' mandelates. $ d-8-Phenyl-lactates.§ [MF" Methyl. Ethyl. n-Propyl. n-Butyl. + S.6" + 13.4" +17*4" +19*6" $- 5.8 + 12.3 +19.1 +21*4 + 11-1 f 19.3 +25*3 +36-4 + 3.7 + 15.9 +29*7 f27.0 - 36.8 - 24.4 -15.5 -15.4 -276 -233 - -2209 (iso-) + 8.5 + 14.7 - -(at 17") * Wood Such and Scarf J.1923 125 601. t For references see Frank-§ Clough ; land and Gebhard J. 1905,87,865. McKenzie and Barrow J. 1911 99 1021. $ Wood and Comley loc. cit. TABLE 111. The Molecular Rotations of the Amides Methyl ester. Z-lactic acid. + 8.6" Z-Glyceric acid + 5-8 d-Malic acid. + 11-1 d-Tartaric acid. + 3.7 Z-Hexahyhmandelic acid. d-B-Phenyl-lactic acid. + 8.5 - 36.8 (at 30") Z-Mandelic acid. - 276 of Some u-Hydrory-acids. Amide.* + 19.6" (in water) + 66.2 (in methyl alcohol) + 52.8 (in water) + 164 (in water) + 65-4 (in aqueous alcohol) -144 (in water) + 104-2 (in ethyl alcohol) * For referencss see Freudenberg Brauns and Siegel Ber. 1923 56 195. -f McKenzie Martin and Rule J. 1914 105 1599 C O N ~ G ~ ~ ~ O N S OF O P T I C ~ Y ACTIVE COMPOUNDS 2811 From Table IV it is evident that acetylation of the methyl esters of the standard " d "-acids (except that of &tartaric acid) causes the molecular rotations to increase.The introduction of one acetyl group into methyl &-tartrate r a h the molecular rotation, but two acetyl groups in the molecule produce a depression in the molecular rotation. Acetylation of methyl d- p-phenyl-lactate is accompanied by an increase in the molecular rotation but although methyl Z-acetylmandelate possesm a higher spec& rotation (- 146') than methyl 1-mandelate (- 166") the molecular rotation of the former is lower than that of the latter compound. The molecular rotations of the benzoyl derivatives exhibit more irregu-larities; thus the complete benzoylation of methyl Z-lactate of methyl Z-glycerate or of methyl d-tartrate produces decreases in %he molecular rotations but the introduction of one benzoyl group only into ethyl d-tartrate and the benzoylation of methyl d-malab are accompanied by increases in the molecular rotations.The values in Table IV show that an increase in the molecular (or the specific) rotation of methyl d- p-phenyl-lactate is produced on benzoylation and that an increase in specific rotation (but a decreItse in molecular rotation) accompanies benzoylation of methyl Z-mandel-ate. Whilst the data for the benzoyl derivatives do not confirm or refute the conclusions drawn respecting the relative configurations of Z-mandelic and d- p-phenyl-lactic acids the optical rotatory power of methyl Z-phenylmethoxyacetate is in accordance with the view that Z-mandelic acid is a '' d "-a-hydroxy-acid.TABLE IV. TAe Molecular Rotations of the AcetyZ Benzoyl and Methyl Derivatives of Some a-Hydroxy-cacids. Acetyl. Methyl Z-lactate (+ 8.6'). + 76.4" Methyl Z-glycerate ( + 5.8). + 24-6 Methyl d-malate (+ 11-1). + 46.8 Methyl d-tartrate (+3.7). + 16.6 (mono- in water)f - 39.6 (di- in alcohol) Methyl Z-mandelate (-276). - 304 Methyl d-8-phenyl-lactate (+ 8.5). + 16-3 Benzoyl. Methyl. - 35.6" * +112.7O - 87.5 +103.8 + 15-0 + 92-4 - 280 + 180-0 (di- at 100') -382 -173 + 92-2 -(in acetone) 5 * Freudenberg and Rhino Ber. 1924 57 1556. t hudenberg and f McKenzie and Wren J. 1910 Bram Be?. 1922 55 1349. 97 484. It is worthy of note that the effect of a rise of temperature on the optical rotatory powers of the esters of Z-hexahydromandelic acid (Wood and Comley Eoc.cit.) is similar to that on the esters of Z-lactic acid E-glyceric acid d-malic acid d-tartaric mid and 3 Clough. VOL. CXXVII. 5 2812 OPTICAL ROTATORY POWERS AND RELATIVE CONFIGURATIONS. Z-mandelic acid. The influence of organic solvenfs on optical rotatory power will be discussed by the author in a later paper. The regularities observed in the influence of sodium barium and other halides on the optical rotatory powers of a-hydroxy-acids or their esters in solution have also been employed for determining the relative configurations of these optically active compounds. The data in Table V fully confkn the results already obtained. It should be pointed out that other inorganic compounds added to aqueous solutions of the acids in question do not always produce similar alterations in the rotatory powers.For example boric acid produces a diminution in the optical rotatory power of l-(" d "-)lactic acid (Henderson and Prentice J. 1902 81 658), but causes an elevation of that of &-tartaric acid (Biot). TABLE V. The InfEuence of Sodium Bromide on the Optical Rotatory Powers of Some Esters of a-Hydroxy-acids in MetAyE-cdcoholic Solution. In methyl-alcoholic sodium bromide ( N ) . In methyl alcohol. Methyl d-malate. + 8.7' (C = 6 ) - 8.0' (C = 5) Ethyl d - d t e . + 11.6 (C = 6) - 2.4 (C = 5) Methyl d-tartrate. + 4.6 (C = 5) - 8.4 (C = 6) n-Propyl d-tartrate. + 16.0 (C = 6) + 2-6 (C = 6) Methyl Z-mandelate .Ethyl 2-mandelate. -117 (C = 10) -150 (C = 10) Methyl d-p-phenyl-lactate. - 2.4 (C = 2.7) - 16.3 (C = 2.7) -142 (C = 3) -172 (C = 3) Ethyl d-8-phenyl-lactate. + 0.8 (C = 6) - 12.0 (c = 5) That the principle employed in this investigation also leads to correct deductions in other classes of compounds is shown by Karrer's confirmation of the present author's view that Z-asparagine, Laspartic acid and Z-leucine are configuratively related to d-(" Z "-) alanine (Helv. Chim. Acta 1923 6 957; see Part I p. 539). More-over a study of the optical rotatory dispersive powers of a number of corresponding derivatives from the optically active a-amino- and a- hydroxy-propionic acids has revealed regularities from which Freudenberg and Rhino have drawn the conclusion that I-('' d "-) alanine possesses the same configuration as I - ( '' d "-)lactic acid (Ber.1924 57 1551; see Part I p. 548). E X P E R I M E B T A L. l-Mcsndelic Acid-In water (c = 1.59) a s (I = 2) - 5-03', a:; - 5-28" agl - 5-98" a% - 13-55"; [a$& - 158" [a]$,- - 166", [a]& - 188" [a]S8 - 426". In aqueous sodium chloride (4N) (c = 1.59) a& ( I = 2) - 6.45"; [a]i4' - 203". Ethyl l-Mandektte.-+@' 1.128 ; txg3 ( I = 0.5) - 73*04O O L ~ - 76-55", c& - 88*10°; [a]& - 129.4" [a)& - 1357" [a]% - 156.4" THE ACTION OF SIUCA ON ELE-LYTES. PABT n. 2813 In methyl alcohol (c = 10) a& (I = 2) - 23-48" & - 24-38', a?& - 28-20' azs - 65.0"; [a% - 117*4" [ a x - 121*9', [a% - 141*0' [a% - 325". Inmethyl-alcoholic sodium bromide (37) (c = 10.12) a& ( I = 2) - 30.30" a& - 36.55"; [arc - 149*6' [a]& - 180.5".Methyl 1-phenylbenxoyloxymetde prepared by the action of benzoyl chloride on methyl I-mandelate in presence of pyridine boiled at 224-225"/18 111111. e' 1.217; Methyl I-P-phenyl-Ztwiate was prepared by McKenzie and Martin's a:. ( I = 0-5) - 86.05"; [a]$' - 141.4". method (J. 1913 103 117). 1.129; a:* ( I = 1) - 5.35"; [aE* - 4-74'. Methyl d-a-cccetoxy-g-pAenylprope'onccte m. p. 30-31" b. p. 185"/20 mm. was prepared by the action of acetyl chloride on methyl d- P-phenyl-lactate in presence of pyridine. 1.125; a:' (I = 1) + 8.23"; [a];. + 7-33". MethyI l-a-benzoyloxy-p-p~nyl~o~'onde b. p. 224-225'116 mm., was prepared by the action of benzoyl chloride on methyl I-8-phenyl-lactate in presence of pyridine. c* 1.161 ; ar ( I = 1) - 37.65"; The above esters required the correct amounts of sodium hydroxide for complete sapodication which was unaccompanied by change of sign of rotation. Some racemisation may have occurred in the preparation of methyl I-phenylbemoyloxyacetate ; Freudenberg and Markert (loc. cit.) give [a%; - 159.9" for this compound. [a%' - 32.45". The author desires to express his thanks to the Government Grant Committee of the Royal Society for a grant towards the expense of this investigation. ROYAL VETERINARY COLLEGE, LONDON N.W. 1. [Received Ocfober 5th 1925.

ISSN:0368-1645    

DOI:10.1039/CT9252702808

出版商:RSC    

年代:1925

数据来源: RSC

摘要:

THE ACTION OF SIUCA ON ELE-LYTES. PABT n. 2813 CCCLXXXVII1.-The Action of silica on Electrolytes. Part 11. By ALFRED FRANCIS JOSEPH and HENRY BOWEN OAKLEY. THE action of silica on two classes of electrolytes-acids * and salts-has already been studied (Joseph and Hancock J. 1923, 123,2022). Its effect on bases is now described. It must again be The experiments with acids have been repeated with purified ‘‘silica gel” with the same result. There has been some correspondence in Natecn (Jan. 31& March 28th and April 4th) on this wbject but the preaent authors have not yet obfained experimentd evidence req- 8 change of opinion. 5 c 3814 JOSEPH AND OAKLEY : emphasised that a high degree of purity of the silica is essential for work of this kind as the presence of a small quantity of basic impurity has an important effect on its properties.For example, a sample of silica gel which retained 6% of wafer aiter drying a t 180” retained only 302% if it had been washed with hydrochloric acid (followed by water) before being dried. Silicates of the AlMi and the AlMine-earth M&.-Mtky studies on these substances a t high temperatures have been carried out in recent years but data as to reactions taking place in presence of much water at the ordinary temperature are scanty. The following facts bear on what follows Potassium seems to form silicates of greater stability in water and lower solubility than sodium the properties of potash glass bear this out and also the fact (Morey J . Amer. Chem. Soc. 1914 36 230) that the compound KHSGO may be left for hours with water a t 100” without undergoing appreciable decomposition.Calcium silicate is less soluble in water than the barium salt. The sodium silicates, even when containing a high proportion of silica are all acted on by water to a noticeable degree. The equilibrium betwem silicic acid and sodium hydroxide was studied by constructing a titration curve for a “water glass ” solution containing 2 mols. of silica to 1 mol. of sodium oxide. The first points on the curve were obtained by the addition of hydrochloric acid and the later ones by the addition of sodium hydroxide. In each case the hydroxyl-ion concentration was determined by p measurement and the amount of sodium silicate present was found by subtracting the quantities of hydrochloric acid and hydroxyl from the total sodium.The results are in Table I. If the sodium oxide neutralised is plotted against p H , the curve shows a definite inflexion when the solution contains 2 mob. of neutralised sodium oxide to 1 mol. of silica. The pH a t which the ratio is half this is very nearly 10 and this corresponds to the value for the fist dissociation constant of silicic acid. The combination between the second molecule of sodium oxide takes place between ~ I I values of 11-6 and 12-6 indicating that the second dissociation constant of the acid is about 10-l2. The curve is typical of the combination of a strong brtse with a weak dibasic acid. TABLE I. Solution contains 0.0186 mol. of silica per litre. pa ..............................3.2 9.40 9-90 10.0 * 10.72 Mols. Na,O neutrdised.. . 0.0 0.0017 0-0043 0-0046* 0.0086 .............................. 11-77’ 12.28 12.31 12.53 12.63 Mols. Na,O neutralised ... 0.0112 0-0180 0-0278 0.0745 0.1570 The point marked * was obtained by interpolation THE ACTION OF SILICA ON ELECTROLYTES. PART II. 2815 The Action of AZMine Hydroxkk8 on SiZica.-When pure silica is acted on by an alkaline hydroxide neutralisation of the base commencw h e d i a t e l y and the extent to which the reaction proceeds is dependent on the nature and the concentration of the base the quantity of silica (at the surface of which the reaction takes place) as well as on the usual factors such as time of contact and temperature. The hen- of the silica will of course deter-mine the amount of reactive surface so that different specimens will not necessarily give the same quantitative results although Merent series will be comparable amongst themselves.The immediate effect is the production of a silicate at the surface of the silica. This is followed by its partial dissolution c a d by the solvent and hydrolytic action of the water. After a time, equilibrium is approached between the silicate in the solid and that in the liquid phase. Increase of concentration of either of the reagents causes an increase in the amount of silicate in each phase. For equal concentrations of different bases the amount of solid silicate formed depends on the solubility relationships of the silicate of the base used; being great where the solubility is small (as in the case of calcium) and small where the solubility is large (as in the case of sodium).These statements are illustrated by the results tabulated below. The amount of silica taken is given as g. per 100 C.C. in all other cases concentrations are expressed as g.-equiv. per litre. The silica was purified and the pE measurements were made M previously described. The reactions were carried out in wax-coated glass flasks. The amount of hydroxyl-ion remaining in solution was calculated from the pH measurements by means of the values of Michaelis for l.ilog K at the temperatures at which the experiments were made these varied between 33" and 38". From the hydroxyl-ion concentration the concentration of free base was calculated from its known degree of dissociation at the working concentration.Titration with standard hydrochloric acid gave the concentration in solution of the free base plus soluble silicate and this subtracted from the amount of base taken gave the amount of the base retained in the solid phase. Neither thermostats nor shakers were used in this work. TABLE 11. Effect of time on the progress of the reaction. Silica 0.5%. Base 0-0365N. Time ........................... 20 hours 48 hours 70 hours 12 days 96 NaOH neutralised . . . . . . 53 63 86 94 yo Ca(OH) neutralised ... 68 79 98 9 2816 JOSEPH AND OAKILEY : TABLE III. Comparison of the amounts of solid silicate formed from different bases after 1 day. Initial concentration of base 0.405N. Baae Fraction of Solid silicate Silicate Solid as Base.neut. base neutr. per litre. in soh. yo of total. NaOH 0.027 0-66 0.0014 0.026 5.2 KOH 0.029 0.71 0.0019 0.027 6.5 Ba(OH) 0.021 0.54 0.0063 0.014 30 Ca( OH) 0.034 0-84 0.034 0.000 100 The results follow the solubility relationships of the silicates of The barium experiment was made at a different time and the these four bases. initial concentration of the base was 0.0392. TABLE IV. Effect of the concentration of the base on the amounts of silicate formed in the solid and liquid phases after 2 days. Initial Final Total Solid Silicate in NaOH. NaOH. silicate. silicate. soh. 0.004 0-00016 04038 Small 0-0038 0.014 0-00058 0.0134 0.00024 0.0132 0-030 0.00159 0.0285 0.00032 0-0282 0.060 0.00828 0.0515 0-00112 0.0504 TABLE V. Effect of the amount of silica taken on that of the silicates formed in the solid and liquid phases after 8 days.Base. 0440N. KOH 0.040N. Ba(OH), 0.039N. CdfgFk NaOH, yo silics ~&CWL 0.5 1.0 0.5 1.0 2.0 0.5 1.0 0.5 1.0 Solid SiliCfbt0 silicate 04014 0-0027 0.0011 0-0021 0.0047 0.0063 0.0138 0-0349 0.0351 in SO^. 0.0366 0.0369 0,0379 0.0383 0-0358 0.0054 0.0092 0.0028 0.0027 neutr. 94 97 97 100 100 30 58 99 99 %B-(Note.-The experiments with sodium hydroxide were carried out with a different specimen of silica this does not affect the relationships which the above table is designed to show.) From the above two tables it appears that there is (a) a direct equilibrium between solid silica solid silicate and liquid phase (hydroxide and silicate in solution) which is the main factor in the case of sodium and potassium; and ( b ) an ordinary solubility relationship between solid and soluble silicates which is the main factor where as in the case of calcium and barium no easily soluble silicates exist.It follows from (a) that any reduction in th THII ACTION ox SILICA ON ELECTROLYTES. PABT n. 2817 amount of free silica should be accompanied by a reduction in the amount of solid silicate and this is shown by the following Sgures for the time effed of a stronger solution of sodium hydroxide on silica the gradual pa8sctge inta solution of the silica is accom-panied by a fall in the amount of solid eilicate. TABLE VI. Simultmeous hppearance into solution of silica and solid silicate. 1% Silica = 0.167 mol. per litre. NaOH 0*0101N.Time. PK. TOM silicate. Solid silicate per litre. 5 hours 11.94 0-0768 0.0057 1 day 11-41 0.0937 0-0046 4 $9 11.00 0.098 1 0.0039 8 9 ) 10-93 0.0986 0-0037 2 days 11-15 0-0972 04040 E8e.d of Neutrtd Salts on the Reaction.-The reactions between neutral salts and silica have been dealt with in the previous paper : acidity is always developed and some of the base goes into the solid phase. The addition of a salt to a mixture of silica and alkali should favour the formation of insoluble silicate by depressing the solvent and hydrolytic action of the water and this is found to be the case. Table VII gives the results of parallel experiments cmied out with and without the addition of the chloride of the base concerned. TABLE VII. 1% Silica and O.OQN-base.Time of standing 1 day. Conc. of salt soh. = N. Fraction of added base found in solid phase. Base. Without salt. With salt. %OH 0.030 0.078 KOH 0.035 0.149 0-70 0.85 0-161 0-210 Ca( OH) * BWH), Experiments carried out with sodium chloride and different amounts of silica showed that as before the amount of silica determined that of the solid silicate formed (which is of come, very small in the absence of free base). Titration methods could not be applied but the hydrogen ions liberated were equivalent to the solid silicate and this is seen to be roughly proportional to the silica taken. TABLE VIII Silica taken. p~ of mixture. H-ion conc. 4-01 o-Ooo10 3-74 0*00018 1% 2 2818 CRAPUN ISOMERIC CHANGE IN ILBOUTIC COMPOUNDS. summary. (a) The action of an alkaline hydroxide on silica results in the formation of a solid silicate part of which passes into solution. The amount of solid silicate formed is very small with sodium and potassium about 50% in the case of barium and nearly 100% with calcium. These results are in accordance with the solubilities of the silicates of these bases. ( b ) In the case of sodium and potassium the small amount of solid silicate produced is roughly proportional to the weight of silica taken although the amount of silicate in solution is only shghtly affected. These relationships show that the solid phases (silica and silicate) are directly concerned in the equilibrium. (c) The addition of a neutral salt increases considerably the amount of solid silicate produced. WELLCOME TBOPICAL RESEARCH LABORATORIES, KHARTOUM. [Received J d y let 1926.

ISSN:0368-1645    

DOI:10.1039/CT9252702813

出版商:RSC    

年代:1925

数据来源: RSC

摘要:

2818 CHAPMAN ISOMERIC CHANGE IN AROMATIC COMPOUNDS. CCCLXXX1X.-Isomeric Change in Aromatic Corn-pounds. Part I . The Conversion of Diacylanilides into Acyhrnino-ketones. By ARTHUR WILLIAM CHAPMAN. IN the come of an investigation on imino-aryl ethers it was found that when a current of dry hydrogen chloride was passed through fused dibenzanilide at 250-270" benzoyl chloride and benzanilide were produced : Diacetanilide decomposed in a similar manner yielding at 150-170" acetyl chloride and acetanilide. Diacylanilides when heated in presence of either hydrogen chloride or zinc chloride are converted into the corresponding mylamino-ketones by what has been regarded as an intramolecular change (Chattaway and Lewis J. 1904 85 386 589 1663; Angel, J. 1912 101 515; Derick and Bornmann J .Amer. Chem. Soc., 1913 35 1269). Migration does not occur in the absence of a catalyst of this kind and the necessary conditions are exactly those under which hydrogen chloride decomposes the diacylanilides. It seems then more reasonable to suppose that with hydrogen chloride as catalyst the change takes place by decomposition of the diacylanilide into acyl chloride and anilide and recondensation to yield the ketone and hydrogen chloride : C,H,*NAc + HCl+C,H,-NHAc + AcCl+C,H,Ac.NHAc + HCl. PhN(COPh) + HC1= PhNH-COPh + PhCOC1 PART I. THE CONVERSION OF DIACYWILIDES ETC. 2819 Zinc chloride especially if it has been allowed to become damp, usually contains a little free hydrogen chloride and this would account for ik catalytic activity. Diacetanilide remained almost unchanged when heated at 1-160" in presence of dry zinc chloride alone but was readily converted into p-acetylaminoacetophenone when hydrogen chloride wits passed into the mixture.The hydrogen chloride therefore appeared to be the essential factor although the zinc chloride no doubt assisted by acting as a condensing agent in the second stage of the conversion. The conversion of diacylanilides into acylamino-ketones furnishes yet another example of the numerous isomeric changes in the aromatic series which have been shown to proceed by fission of the mobile group followed by recondensation such as the Hofmann-Martius change (Beckmann and Correns Ber. 1922 55 852) the Fischer-Hepp change (0. Fischer Annalen 1895 286 145) the conversion of phenolic esters into hydroxy-ketones (Skraup and Poller Ber.1924 57 2033) and the isomeric change of imino-aql ether hydrochlorides (J. 1922 121 1676; 1923 123 1150). It is proposed to extend the present study fo other similar isomeric changes which have not yet been investigated from this point of view. E X P E R I M E N T A L . Decomposition of Dibenzunilide and of Diacetanilide by Hydrogen CiLEoride.-Dry hydrogen chloride was passed through the fused diacylanilide the temperature of which was gradually raised until distillation began and was then maintained constant until no more liquid came over. Diacetanilide (20 9.) yielded a colourless distillate (6-7 g.) b. p. 51-53" which reacted with p-toluidine to give acefo-p-toluidide (m. p. 147").The residue (11.5 g.) b. p. 286-290" crystrtllised from benzene or water in shining leaves (m. p. 114" alone or mixed with acetanilide). p-Aminoacetophenonepacetylaminoacetophenone and acetadide did not yield any acetyl chloride on treatment with hydrogen chloride even at 200-250". Dibemanilide (10 g.) yielded benzoyl chloride (2-1 g . ; b. p. 193-200"; converted by phenol and sodium hydroxide into phenyl benzoate m. p. 68-69') and benzanilide (4.8 g. ; m. p. after recrystahtion 159-160" not depressed by admixture with an authentic specimen). Conversion of Dimtanilide into p-Acetylamima&oph.-A mixture of diacetanilide (20 g.) and powdered dry zinc chloride (3 g.) ww heated at 140-160" while hydrogen chloride was led into 5 c 2820 THOBUS AND BARKER THE PARTIAL PRESSURES OF it for 44 hours.Since p-acetylaminoacetphenone is f i c u l t to isolate as such the product WM hydrolysed by boiling with con-centrated hydrochloric acid (150 c.c.) and the solution ww made alkaline with solid sodium hydroxide steam-distilled to remove aniline cooled and filtered. Boiling light petroleum extracted from the precipitate about 2 g. of paminoacefophenone (m. p. 106-107" not depressed by admixture with an authentic sample). In a similar experiment without hydrogen chloride no p-amino-acetophenone was isolated from the light petroleum extract but the alkaline aqueous liquor gave a deep red dye on diazotisation and coupling with 8-naphthol. In two control experiments in which diacetadide was heated alone no paminoacetophenone could be detected even by the diazo-reaction. p-Aminoacetophenone was also obtained by pwing acetyl chloride vapour through fused acetanilide and zinc chloride but was accompanied by much diphenylethenylamidine (m. p. 132-133"), formed by the condensation of 2 mob. of acetanilide with elimination of acetic acid. THE UNIVERSITY SHEFFIELD. [Received October 14th 1925.

ISSN:0368-1645    

DOI:10.1039/CT9252702818

出版商:RSC    

年代:1925

数据来源: RSC

摘要:

2820 THOBUS AND BARKER THE PARTIAL PRESSURES OF CCCXC.-The Partial Pressures of Water Vapour and of Sdphuric Acid Vapour over Concentrated Solutions of Sulphuric Acid at High Temperatures. By JOHN SMEATH THOMAS and W m F ~ C I S BARKER. THIS investigation was undertaken with the object of extending the measurements of Thomas and Ramsay (J.? 1923,123 3256). These authors determined the partial pressures of the sulphuric acid only. It seemed desirable also to obtain data regmding the partial pressures of the water vapour in equilibrium with concentrated solutions of the acid over the same range of temperature. At an early stage in the work it became evident that the results obtained in the previous investigation were uniformly low. Further examination revealed two sources of error (a) adsorption of sulphuric acid vapour on the glass wool employed to remove " &t," and (b) the possibility that the liquid was not in thermal equilibrium with the heating bath.Of these the former is the more serious ; it had been foreseen but the precautions taken against it now appear to have been inadequate. As regards the second possibility in the previous work the actual temperature of the acid was not measured ; it was assumed to have attained the temperature of the bath considerable time having been allowed for this purpose WATER VAPOTJR AND OF s m m c ACID vmm ETC. 2821 In this invdigation the apparafns was modified to eliminate these errors as far a~ possible and measurementa of the partial pressures of the water vapour and of the sulphuric acid were made a t a series of tempemtmx ranging from 180" to 300".The dynamical method (h. cit.) was again employed the partial pressures being calculated from the experimental data by the method of Foote and Scholea (J. AM. Chm. Soc. 1911 33 I=) but w h e w in the previous work the volume of the water vapour removed by the current of air was either neglected as in the case of the more con-centrated solutions examined or calculated * from Burt's measure-ments of the total vapour pressure (J. 1904,90,1339) the sulphuric FIG. 1. acid partial pressures of these more dilute solutions being assumed to be negligible in the present work it was obtained from the actual weight of water removed by the air. The partial pressure is then given by the formula ih which v represents the volume of the water vapour or the sulphuric acid vapour as the case may be wl the volume of air (reduced to N.T.P.) and w2 the combined volumes of the water vapour and the sulphuric acid vapour removed during the experiment.In calculat-ing these volumes the value 22-3 litres was taken for the gram-mole-cular volume. E x P E R I M E N TAL. P = 760 W l - v2), The rnod5cation.s in the apparatus can best be understood by reference to Fig. 1. volume of the water vapour was inadverbntly omitted. * In the previous communication reference to this correction for the 5 c * 2822 THOMAS AND BARKER THE PART= PRESSURES OF A current of air from the aspirator A was dried and freed from carbon dioxide by passing it through the vessels B and C containing concentrated sulphuric acid and soda-lime respectively.The aspirator was fitted with the manometer D in order that the air supply might be kept at constant pressurea pressure equivalent to 900 mm. Hg was found to be convenient. This greatly facilitated the experimental manipulation of the apparatus. From C the dried and purified air passed through the saturators E and F which con-tained the acid solution under investigation. These saturators each having a working capacity of approxi-mately 250 c.c. were constructed and arranged in the thermostat in the manner indicated in the diagram. They were half filled with short pieces of glass tube of narrow bore and worked very efficiently ; the tiny air bubbles produced by the constricted vertical air inlet tube followed a circuitous path and remained in contact with the acid for a considerable time.From the bottoms of the saturators vertical tubes passed through the bottom of the thermo-stat and to these taps were sealed by means of which acid was introduced and specimens could be withdrawn for analysis. To the upper portion of the final saturator a wide glass tube was sealed; a thermometer fitted into this by means of the ground joint G., dipped below the surface of the acid the actual temperature of which could thus be observed. Although no regulator was used this temperature remained very constant; variations in the voltage of the electrical supply occasionally gave rise to slight changes. During an experiment the temperature was observed at %minute intervals.These readings were plotted and from the graph the mean temper-ature was obtained. The thermostat consisted of a double-walled box the outer wall being of galvanised iron and the inner wall which was separated from the outer by an air space of about & inch of asbestos slate. Through each side of the box two insulated terminals passed leading to the heating units-spirals of nichrome wire-any of which could be used or cut out at will according to the temperature required. The final temperature regulation was effected by means of an adjustable elxternal resistance. The air saturated with acid vapour passed from the k a l saturator through the absorbing vessels J into the measuring device K. Two of these vessels were employed but only one is shown in the cliagram.Each had a capacity of about 300 c.c.; the volume between the fixed marks had been carefully determined and the lower limb was graduated. During the experiment water was run out at such a rate as to keep the air under atmospheric pressure. (a) the total Each experiment involved two determinations WATER VAPOTJR AND OF S U L P ~ C ACID VAPOUR ETC. 2823 weight of vapour (sulphuric acid and water) removed by a given volume of air and ( b ) the weight of the removed sulphuric acid vapour alone. For the determination of the total weight of vapour a series of absorption tubes was used. These contained solid sodium hydroxide followed by granular calcium chloride and were carefully weighed before and after each experiment with all the usual pre-cautions.For the estimation of the sulphuric acid vapour the conductivity method described (Zoc. cit.) was employed except in the case of the 99.23% solution ; no water vapour could be detected in the vapour from this and the increase in weight of the absorption tubes was assumed to be entirely due to sulphuric acid. Connexion between the absorbing vessels and the h a 1 saturator was made by means of the carefully ground joint H which was situated inside the thermostat thus obviating the premature con-densation of vapour. The inner portion of this joint projected into a small bulb not shown in the diagram the object of which was to prevent the creeping of acid along the surface of the glass into the absorbing vessels. Preliminary experiments having disclosed that the use of glass wool for the prevention of the mechanical transference of acid in the form of " mist " leads to serious errors this end was attained by using an extremely slow current of air ; in the final experimenta the rate of flow never exceeded 150 C.C.per hour. On examining the brightly illuminated space above the surface of the acid no trace of mist could be seen. The strength of the various acid solutions examined was deter-mined gravimetrically. Analysis of four samples from each saturator at the close of the series of determinations revealed no change in concentration. The Variatim of the Par'tial Pressures with the Temperature. Five series of measurements were made on sulphuric acid solutions of concentrations between 89-25 and 99.23% H,SO, in each case a t a number of different temperatures ranging from 180" to 295"; on account of the lower boiling points of the more dilute solutions measurements could not be made at the higher temperature.The results are in Table I in which the pressures are expressed in mm. of mercury. In calculating the values in columns 2 3 and 4 the assumption has been made that the sulphuric acid vapour is not dissociated, whilst the values in the next three columns are based on the assumption that complete dissociation of the acid vapour occm. The values for the dissociated acid cannot be obtained from the corresponding figures for undissociated acid by means of the expres 2824 moms rn BARKER rn PAR^ PRESSURES OF t . 183.0" 197.5 216-5 230.0 241.5 191.0 205-0 222.0 242.5 252.5 258.0 262.5 180.0 200.0 215.5 232.0 244.5 252.0 261-0 270.0 280.5 282.0 204.0 218.5 234.5 249.0 261.0 273.0 285-0 295.0 211.0 225.0 227-0 244.0 261.0 270-0 281.0 290-0 ?k&34* 0.5 1.3 2.1 3.8 5.3 0.6 1.9 4.5 6.4 11.3 13.6 16.3 2.1 4.8 8.5 13.4 19.9 20.0 27.9 39.9 52.0 52.6 5.9 9.8 14-7 28.6 38-8 61.9 91.6 132-3 33.2 49.9 55-4 84.1 163.8 229-8 272.3 381.5 TABLE I.Series A. 89.25% H,S04. Assuming complete dissociation. r ,-. > PESO. P. P'SOa. P'HIO- P'. 78.8 79.3 0-5 79.2 79.7 116-9 118.2 1.3 118.0 119.3 233.1 235.2 2-1 234.6 236.7 306.3 309.9 3.6 308.4 312.0 414.8 420.1 5.3 417.2 422.5 Series 33.50.7 84.7 158.5 271.6 385.3 448.7 41 1.1 91.26% H,SO, 51.3 0.6 51.3 51.9 86.6 1.9 86-1 88.0 163.0 4.5 162.1 166.6 278-0 6.4 275.7 282.1 396-6 11.1 390.7 401-8 462.3 13-4 454-2 467.6 427.4 15.9 419-0 434.9 Series C. 95.06% H,SO,. 10.1 21.2 46.5 91.9 120.1 156.5 180.7 254-9 310.0 350-2 12-2 26-0 55.0 105.3 140.0 176.5 208-6 294.8 362.0 402.8 2-1 4.8 8-4 13.2 19.4 19.5 26-9 37.8 48-7 49.2 Series D. 98.06% H,SO,. 0.0 1.5 3-2 2.6 5.0 5.3 11.8 14.7 5.9 11.3 17.9 31.1 43-8 67-2 103.4 147.0 5.9 9.7 14.4 27-5 36.9 57-2 81.8 112-7 Series E. 99.23% H,SO,. - 33.2 33.2 - 49.9 49.1) - 55.4 55.4 - 84.1 84- 1 I 163.8 163.8 - 229.8 229.8 - 272-3 272.3 - 381.5 351.5 12.2 25-8 54.4 103.5 136.4 172-0 201-1 279-6 351.9 376.5 5-9 11-2 17.6 30.0 41.6 62-5 92-2 126.2 33.2 49.9 55.4 84.1 163.8 229-8 272.3 351-5 14-3 30-6 62.8 116.7 155-8 191.5 228-0 317-4 400-6 425.7 11.8 20-9 32.0 57.5 78-5 119-7 174-0 237.9 66.4 99-8 110.8 168.2 327.6 459.6 544.6 763.0 sions fifiO4 = p'80 and Psi,.= 2pHrB04 + pEt0 where p' refers to acid in which dissociation is assumed. The reason for this lies in the method which is applicable whichever assumption is made, but which does not allow of results being translated from one basis to the other because of the volume change which must be taken into account. The following example taken from an actual experi WATER v ~ o m AXD OF s m m c ACID V I I P O ~ m.2825 ment on 91.26% H,SO, makes this point clear. In this example the question is simpMed because the volume of air used in the two runs chanced to be the same. 292-3 C.C. of air at N.T.P. were passed through the acid (262.5"). The increase in weight of the absorption tubes was 0-3554 g. and the weight of acid absorbed in conductivity water was 0.0630 g. Assuming no d i s s m i a t h : Vol. of 0.0630 g. %SO and 0-2924 g. -0 at N.T.P. = 14.3 and 362.2 c.c. respectively. Total vol. = 668.8 C.C. %Lo = 14.3 x 760/668-8 = 16.3 mm. and % = 362-2 x 760j668.8 = 411.1 mm. Therefore P=427.4 mm. Asmming the acid to be completely dissociated Weight of SO in the acid vapour removed = 0.0630 x 80/98 = 0.0514 g.and weight of water removed = 0.3554 -0.0514 = 0.3040 g. The corresponding volumes at N.T.P. are 14.3 and 376.6 c.c. respectively. Total volume = 683.2 C.C. plmI = 14.3 x 760/683-2 = 15-91 mm. and plHt0 = 376.6 x 760/683-2 = 419.0 221111. Therefore Pfdiar = 434.91 mm. But 2p~ie~ +-so = 443.7 mm. The difference between the results obtained by the two methods of calculation is not serious for the more dilute acids a t moderate temperatures. It becomes important at higher tem-peratures especially when- the concentration exceeds 95%. Although for the sake of uniformity all the experimental results are carried to the first decimal place this degree of accuracy is not necessarily claimed for each single determination. At low temperatures the relation between temperature and partial pressure of the water or of the acid vapour was approximately linear in both cases.At higher temperatures however the partial pressures increase with increasing temperature more rapidly than this simple relationship demands. The variation of the pa,rtial pressure with change of temperature can be represented with con-siderable llccuracy by expressions such as that used by Perman for solutiom of ammonia in water or better by Rankine's equation loglo p = a - PIT - y log T. The constants a p and y in the latter expression were obtained in the usual way from values taken from the smoothed curve. 99.23% H,SO p a = - 20.0946; 1)~t80 a = 8.2878; P H ~ a = 20.7109; a = 10.8170; {r a = -106.5692; a = 8.4658; Z ) H J ~ a = 48.2388; 91.26y0 H,SO mfi a = 65.2394; a = 66-8040; -,go a = 75.5303; 89.25% H,SO Z)ET.O a =- 32.5120; i P a =- 33.5235; 95.06% H,SO mo a = 22.4406; 8 = B = B = 8 = 8 = 8 = 8 = B = B = B = B = B = I s = 1695.86 ; 4319.7 ; 5415.31 ; 4518-21 ; 4541.0 ; 3516.97 ; 7026-67 ; 7354.24 ; 7483.28 ; 8529-90 ; 216.85 ; 166-663 ; -4828.25 ; y =- 9.3427.=- 0.5166. y = 3.6353. = 0.2571. 7 = -36.2945. = 4.2726. =- 0.1707. y = 12.4310. y = 17.8856. y = 18.3659. y = 21.4699. y =-13.114. y = - 134453 2826 THOMAS AND BARKER THE PARTIAL PRESSURES OF t. 180' 185 190 195 200 205 210 215 220 225 230 235 240 245 250 255 260 265 270 275 280 285 290 295 300 305 3 10 315 320 %¶804.msor' obs. calc. P. 10.0 9.5 19.5 12.0 11.7 24.0 14-2 14.0 28.0 17.2 16.95 34.2 21.0 20.5 41.0 24.5 24.6 48.0 29.0 29-5 58.2 34-2 35-3 68-5 41.0 42-1 81.0 49-0 50.0 94.5 58.0 60-1 110.5 67-5 70.6 128.0 80.0 83-1 149-7 94-0 97.7 173.0 111.0 115-5 198-2 131.0 134.5 227.0 153.0 158.0 257-5 179.0 184.3 291.0 209.5 214-9 327-5 243.0 250.5 368.0 282.2 286.6 412.0 328.0 326-6 459-0 380.0 389.1 600.2 440.0 448.4 565.5 505.2 517.2 626.2 577.0 596-2 692.0 659.5 684.7 -752.0 - - - 785.1 -95.06y0 H,SO,. -0,-Oh. 1.5 2.0 2.6 3-2 4.0 4-7 6.2 7.5 8.7 10.5 13-0 15.7 18-7 22.5 27-0 32-2 38.5 46.2 55.5 66-0 78.5 92.7 109.5 130.7 150.1 174-5 203.5 235-0 270-0 -04' calc .1.3 1-7 2.2 2.7 3.46 4.3 5.4 6.7 8.3 10.2 12.5 35.3 18-6 22.2 27-1 32.6 39.2 46-7 55.6 66- 1 78.1 92.3 109.1 132.7 151.7 174.2 202-8 235.5 272.7 P3HaO. P m O . obs. calc. P. 0.14 0-12 3.2 0-18 0.16 4.2 0-25 0.21 5.1 0-30 0-27 5.5 0.35 0.35 8-0 0.45 0.44 9-7 0.57 0.55 12.0 0-72 0.69 15.0 0.90 0.87 18.2 1-10 1.07 22.0 1.32 1.30 27.0 1.67 1.64 32-7 2.00 2.01 39.2 2.40 2.45 46.7 2.95 2.98 56.0 3-52 3.61 66.; 4-22 4.35 78-7 5.10 5.20 92-5 5-15 6-25 109.0 7.42 7.46 127.5 8.85 8-94 150.2 10.52 10-50 175.5 12.47 12-40 205.0 14.65 14.60 240.5 17.10 17.12 284-0 19.95 20.02 336.5 23-0 23.35 393.5 26.45 27.14 456.0 30.25 31-50 515.0 91*260, H,SO,. t . 180" 185 190 195 200 205 210 215 220 225 230 235 240 245 250 256 260 265 270 275 280 285 290 295 300 3.2 3.1 3.5 3.5 4.0 4-0 4.5 4.6 5-1 5.2 5.9 6-0 6.7 6.8 7.7 1-9 9.1 9.1 10.3 10-4 11.9 12.0 13.7 13-8 15.6 15.9 18.1 18.4 20.9 21.2 24.3 24.4 28-1 28.3 32-5 32.6 37.9 38.0 43.6 43.9 50-2 50.7 58-0 58-4 67.5 67.9 78.5 78.4 11.0 12.5 16.5 20.5 25.5 31.0 37.0 44-0 53-0 62-2 74.0 87.0 102.2 119-5 139-0 160.2 185.2 213.5 245-0 281.7 323.0 369-0 420.0 478.0 11-7 16.4 14.3 18.2 17.5 22.0 21-3 27-2 25.8 31-7 31.1 38-5 37.3 46-2 44.5 55.0 53.0 66.0 62-5 79.1 74-1 93-5 87.2 110-5 102-1 130.0 119.4 152-3 139.0 176-0 161.3 202-5 186.5 231.0 215-0 263.2 247-1 300.0 282.3 343-0 323.8 393-0 369-4 452-2 420-2 522.0 476-3 600-0 91.6 90.7 539.5 638.6 690-0 305 106.6 105.7 608.0 607.7 -0.47 0.62 0.82 1.05 1.35 1.71 2.14 2.65 3-27 3-98 4-85 5-87 7-07 8.55 10.25 12.25 14-60 17-30 20.30 24.9 0.51 33-5 0.66 41.0 0.84 48.5 1.07 59.0 1.35 70-0 1-70 85.4 2.13 102.6 2-63 122.0 3-24 146.2 3-98 172.0 4.84 200-9 5.88 233.0 7.12 271.0 8-55 313.0 10-22 360.5 12-16 413.0 14-45 469.2 17-01 532.5 20.02 600.0 22.42 675-2 31.6 39.0 49.0 58.4 70-7 85.4 102.6 120.4 144.3 170.4 199.2 233.0 270.6 312.1 360.1 413-0 470-8 535.5 604-5 684.9 35.0 42.0 50.0 61.2 78-1 88.0 106.2 127.2 152.0 179.5 209-5 2444 283.0 327.0 377.2 433.5 499-5 561-2 636-5 721.WATER VDUR AND OF s m m c ACID VAPOUE ETC. 2827 %b%4. t. obs. 180" 0.45 185 0.60 190 0-75 195 0-95 200 1-18 205 1.43 210 1-75 215 2-12 220 2.56 1)EfiOr-calc. 0.47 0-60 0-75 0.94 1-17 1-44 1.76 2- 15 2.59 Pmo-obs. 69-9 82.5 96.0 111.7 128-1 149.5 172.0 198-5 229.5 TABLE II (continued). 89.25% H,SO,. A , PH10. PE¶fk34*PHdo4* *. Pa@' calc. P. t. obs. calc. obs. calc. P. 69.4 70-7 225" 3-07 3.12 263.7 265.6 268-0 78.8 83.5 230 3-67 3.71 303.0 306.0 307.9 94.7 97.3 235 4.34 4-42 347.0 348.9 353.2 110-3 113.2 240 5-10 5.22 400.0 403.2 408.0 128.4 131.0 245 6-00 6-13 463-0 462.8 471-2 148-7 160-9 250 7.13 7.16 534.1 530.0 645.0 172.2 174.8 255 8-25 8.35 614.0 605.0 628.5 199.4 201.7 260 9.56 9-66 702.0 691.5 723.5 230.0 235.5 The values of the partial pressures and also of the total pressures taken from the smoothed curve which may be looked upon as the true experimentally determined values agree closely with the values calculated by means of these equations over the whole range.As was to be expected the divergences are greatest at low temperatures. The observed and the calculated values of the partial pressures are set out in Table 11. The values of P are obtained from the observed values on the assumption that the acid vapour undergoes complete dissociation. On extrapolating the total-pressure curves to 760 mm. the boiling points of the Merent solutions are obtained.Data concerning the boding points of concentrated solutions of sulphuric acid have been given by Marignac (see " Sulphuric Acid and Sulphur Dioxide," Wyld p. 194) and Beckmann (2. physikal. Chem. 1905 53 129). The values obtained in this investigation are in good agreement with those obtained by graphical interpolation from the above-mentioned sources. They are shown in Table III. TABLE 111. Conc. B. p. from B. p. calc. of V. P. from acid. curve. equation. 99.23% 315.2" 313-8O 98.06 349 352.8 95.06 307-5 307-6 91.26 278.0 278.1 89.25 261.5 263-2 B* p. from. 1)HsSO4 other at b. p. data. (mm.). 3 10" -331-338* 678.7 300 115 268.7 26.5 257 12-9 B. p. calc. at b. p. vapour (mm.). dissociated. - 309.75O 81.3 332.5 646 302.5 733-5 277 747- 1 26 1 PHrO assuming&SO, * The value 331.7' given by Landolt and Bornstein (" Tabllen," 5th ed., p.1433) on the authority of Beckmann for the boiling point of pure dphuric acid clearly should refer to the constant-boiling mixture containing 98.3% H,SO,. In columns 5 and 6 of the above table are given the paztial pressures of the water vapour and of the sulphuric acid vapour respectively at the bo- point 2828 THOMAS AXD BAEKFJ& THE PARTIAL PRESSURES OF Throughout this work it htts been assumed that the sulphuric acid vapour remains undissociated. Whilst this may be approxi-mately true a t the lower temperatures and especially for the more dilute solutions the vapour in equilibrium with which always contains a large excws of water molecules a t higher temperatures, particularly in the case of the more concentrated solutions this assumption cannot be justified.According to Dittmar (Chem. News 1870 20 258) pure sulphuric acid in the state of vapour is practically completely dissociated. If thh is so the values obtained for the sulphuric acid partial pressure will be lower than the real partial pressure of the dissociated acid and the boiling points obtained from these measurements will therefore be too high. The figures given in Table 111 show this to be the case although it should be pointed out that this influence even supposing it to exert its maximum possible effect is comparatively slight for solutions the concentrations of which lie below 95% &SO,. For example, whereas complete dissociation of the sulphuric acid vapour would increase the total vapour pressure of 98.06% acid a t 275" from 73.4 mm.to 127.5 mm. an increase of 73.774 in the case of 91.26% H,SO the increase a t the same temperature amounts to 3% only. Unfortunately the extent to which sulphuric acid vapour is dissociated a t different temperatures and in the presence of widely varying concentrations of water vapour is not known and c6me-quently no systematic correction for this factor can be applied to the values obtained in this research. In Table I1 values are given for the total vapour pressures based on the assumption that the sulphuric acid vapour is always completely dissociated. For solutions containing less than 95% H,SO, the increase in the total vapour pressure due to this factor is very slight.In the last column of Table 111 are given the values obtained for the boiling points when it is assumed that the sulphuric acid vapour is completely dissociated. These figures approach very closely indeed to the generally accepted values. The Variation of the Total Pressure and of the Partial Pressures with the Concentration. The mode in which the total vapour pressures and the component partial pressures over concentrated sulphuric acid solutions vary when the compositions of these solutions is changed is illustrated in Fig. 2. In this diagram the total pressure the pEoso4- and the pGo-isothermals have been drawn for temperatures of 260" 230°, and 200". In the case of both the pHtsO4- and the pW-isothermals a sudden and very marked change These isothermals are typical.The form of these curves is striking WATER VAPOUR AND OF SULPHURIC ACID VAPOUR ETC. 2829 of direction occurs when the solution contains 9802% of sulphuric acid. The total-pressure isothermals exhibit a pronounced minimum at the same composition. This concentration agrees very closely FIG. 2. C m p i t i o n of Solt&-imthermals. 100 99 98 97 96 95 94 93 92 90 91 89 650 600 550 8 -600 P 3 L 450 0' 440 350 22 2 g300 2 9 h 250 % 2 200 0) I-a 9 $ t 2 150 100 50 0 89 90 91 92 93 94 95 96 97 98 99 1ooo/, Cornpition of Solution-P-and ~~or-ieothemna18 only. B on ~p~to-isdhermale. @ Poi& m ~ - i e d h e W . x Poi& m P-ieothnnds. indeed with that of the constant-boiling mixture which is usually considered to contain 98.3% &SO4.At this concentration the composition of the liquid and gaseou 2830 TEOMAS AND BARKER THE PARTIAL PRESSURES OF phases must be the same and assuming the vapours in equilibrium with the liquid to behave as perfect gases the ratio of the two partial pressures should be identical with that of the molecular fmctions in the liquid. In Table IV these ratios are compared for a series of temperatures. TABLE IV. t . ?)Hr80r-1 SO" 1.5 200 3-7 220 10-5 240 19.6 2 60 39.7 280 82.2 300 159.0 PHto. 0.14 0-36 0.85 1-85 3.8 8.0 15.5 ?)Hasor IPHsO-10.7 10.5 10.23 10.05 10.45 10.27 10-30 M/(1 - 31). 10.01 10.0 1 10.01 10.01 10.01 10.01 10.01 When the nature of the partial-pressure curves and the very large variations in pressure which accompany small changes in composition in the neighbourhood of this point are taken into account the agreement is to be considered quite satisfactory.Since the partial pressures of both water and sulphuric acid vapour have now been determined it is possible t o test the applicability to concentrated sulphuric acid solutions under the conditions of these experiments of the Duhem-Margules equation in which pl and p2 represent the partial pressures of sulphuric acid and water vapours and M and 1 - 111 the molecular fractions of these substances in the liquid mixture. d log pl/d log M = d log p,/d log (1 - M ) TABLE T7 v . Conc. of H,SO,. t . yo. 180" 89 91 93 95 200 89 91 93 95 220 89 91 93 95 240 89 91 93 95 260 89 91 93 95 M .0.5978 0.6257 0-6864 0-7421 0.5978 0.6257 0-6864 0-7421 0.5978 0-6257 0.6864 0-7421 0.5978 0.6257 0.6864 0-7421 0.5978 0-6257 0-6864 P1. Pe-0.30 65 0.62 43 1.02 26 1-55 13 0-86 128 1.50 80 2-37 44-5 3-50 20.5 2-50 240 3.65 150.7 4.90 89-5 6.85 46.0 4.86 406 6-9 291 9.8 186.5 13.3 102.0 8.9 727 14-0 502 20.1 329 0.7421 27-1 188 dP1- -dP,-0.09 20.0 0-18 11.0 0.22 9.0 0.27 6.0 0.25 37-0 0.40 23.5 0.48 14.6 0-60 10.0 0.45 57.0 0.62 36.0 0.85 25.5 1-10 18.5 0-80 90.0 1.05 63-0 1.65 40.0 2-05 37-0 2.0 142.0 2-61 100.0 3.39 79.0 3.79 68.5 - _ dP1 dP%' 0.0045 0.0164 0.0246 0.0458 0.006 7 0.0170 0.033 1 0-06 10 0.0079 0.0172 0.0383 0-0594 0.0089 0.0167 0-0412 0.0553 0.0141 0.0261 0.0429 0.0553 fi.k M - . Pr M 0.003 1 0-0086 0.0179 0-04 14 0-005 1 0-0112 0.0243 0-0593 0.0069 0.01 11 0.0250 0.0517 0.008 1 0.0142 0-0240 0.0453 0.0082 0-0167 0.0279 0-050 WATER VAPOUR AND OF SULPHURIC ACID VAPOUR ETC. 2831 The integration of this expression involves certain assumptions regarding the relationship between the partial pressure of each component and the vapour pressure of that component in the pure state. If however the equation is written in the form - dPl =a .1-- M dP2 P2 iK it can be tested by substitution of the values obtained from the isothermals. The estimation of dpl and dp was carried out graphically by drawing the approximate tangent to the isothermal at the required point. The results obtained are summarised in Table V. The values obtained for pl/p2.(l - N)/M are invariably lower than those for - dpl/dp2 but in view of the fact that considerable dissociation of the sulphuric acid vapour probably occurs close agreement was scarcely to be expected. Summary. 1. The partial pressures of the water vapour in confact with acid-water mixtures containing from 89% to 99.3% of sulphuric acid have been determined a t temperatures ranging from 180" to 300" by a method previously described. 2. The previously determined values of the sulphuric acid partial pressures were vitiated by adsorption of acid vapour on the glaas wool employed to prevent mechanical transference of liquid and have been redetermined. 3. Both the water and the sulphuric acid partial pressures and also the total pressure can be represented by Rankine's equation for which the various constants have been calculated. 4. The b o h g points obtained from the total pressure by extra-polation are considerably too high. If however it is assumed that complete dissociation of the sulphuric acid vapour occurs values are obtained which are in good agreement with the accepted values. 5. The pHIBO,- pHto- and P-isothermals have been constructed for temperatures of 260" 230" and 200". The P-isotherm& exhibit a sharp minimum at a composition 98.2% €&SO4. The form of the partial pressure curves is in qualitative agreement with the Duhem-Margules equation. The quantitative agreement however is not close probably on account of dissociation of the acid vapour. logp = a - @/T - ylog T, THE UNIVEBSITY CAPE TOWN. [Received March 16th 1925.

ISSN:0368-1645    

DOI:10.1039/CT9252702820

出版商:RSC    

年代:1925

数据来源: RSC

摘要:

2832 LEDBURY AND BLAIR THE PARTIAL FORMALDEHYDE CCCXC1.-The Partial Formaldehyde VUPOUT Pressures By WILFRID LEDBURY and ETHELBERT WILLIAM BLAIR. THE investigation of the partial formaldehyde vapour pressures of aqueous formaldehyde solutions described in Part I (this vol., p. 26) showed that following a prolonged initial exposure a t 15", equilibrium between the liquid and rapour phases was only gradually attained when the temperature of a solution was subsequently changed to 20" or 0". An explanation of this was based upon the assumption that a new equilibrium between complex (polymerised, hydrated etc.) and simple molecules in solution depending on the temperature was gradually approached as time elapsed. The presence in aqueous formaldehyde solution of both the polymerised and monomolecular forms is made evident by the work of Auerbach and Barschall (Arb.Kais. Ges.-A. 1905 22 584; Chm. Centr., 1905 II 1081) who carried out molecular-weight determinations by the cryoscopic method. These determinations indicated that in solutions containing 37-38% (by vol.) of formaldehyde there probably exist polymerides of complexity greater than trimolecular . Auerbach and Barschall concluded that the equilibrium between the different molecular forms of formaldehyde is reversible so that the condition of aqueous solutions of formaldehyde a short time after preparation depends onIy on the concentration and temperature. In order to study further the influence of temperature and solution concentration on the partial formaldehyde vapour pressures of aqueous formaldehyde solutions determinations have been carried out a t 35" and 45" and these new data correlated with those pre-viously obtained at 0" and 20".It has been previously pointed out by the authors that methyl alcohol when present in small amounts (compare the proportion of methyl alcohol to formaldehyde in commercial formalin) con-siderably enhances the partial formaldehyde vapour pressures of formaldehyde solutions a t 20". Formaldehyde in methyl-alcoholic solution and in a mixture of water and methyl alcohol containing a large proportion of the latter has since been investigated in this connexion and also the effect of the presence of small amounts of the alcohol in formaldehyde solutions a t 35" and 45". of Apww Solutions of Formaldehyde.Part 11. E x P E R I M E N T A L. Method and Appratus.-Aqueous solutions of formaldehyde free from methyl alcohol were prepared either by the method of con-tinuous refluxing previously described or by dissolving solid para VAPOUB PBESSUBES OF AQUEOUS SOLUTIONS ETW. 2833 foddehyde uncontaminsted with methyl alcohol in boiling water. The solutions 80 prepared were kept for 2 or 3 weeks at room temperature ; it was then assumed that the equilibrium between the difFerent molecular species of formaldehyde in solution had attained a value dependent only on the concentration and temperature of exposure irrespective of the method of preparation (Auerbach and BarscI.mIl loc. cit. ; Auerbach and Pliiddemann ibid. 1914,47,116). To ascertain the effect of the presence of small amounts of methyl alcohol on the partial formaldehyde vapour pressures of the solutions, the alcohol having an acetone content of less than O.lO% was added to a series of the solutions in such amounts as to provide a constant ratio CH40/CH,0.For the preparation of formaldehyde in methyl-alcoholic Solution a current of dry nitrogen was passed over para-formaldehyde in a silica tube at 170". Formaldehyde vapour, cam& forward by the gas-stream was absorbed in methyl alcohol, as free aa possible from acetone. The solution was filtered from any insoluble polyoxymethylenes and kept Beveral weeks before use. This solution has an exceptionally irritant action on the membranes of the eyes and nose. Unlike aqueous formaldehyde, it produces this effect after a minute or so.The formaldehyde content (p. 15%) of this solution was estimated both before and after refluxing with dilute sulphuric acid and the Uerence in the values showed that the equivalent of about 5 g. of formddehyde per 100 C.C. was combined as methylal. When a methyl-alcoholic solution of formaldehyde was prepared by heating an excess of pure paraformaldehyde with methyl alcohol in a seded tube for 6 hours at 120" the major portion of the aldehyde in solution was combined as methylal. The free formaldehyde content was 6.63 g. per 100 c.c. whilst the methylal present waa equivalent to 12-0 g. per 100 C.C. Refluxing paraformaldehyde with methyl alcohol at the ordinary pressure for 2 days yielded a solution con-taining. 5.16 g. of free formaldehyde and 144 g.of combined formaldehyde per 100 C.C. As in the corresponding determinations at 20" and 0" the " dynamic " or " flow " method was employed for the series of experiments at 35" and 45". The apparatus used was similar to that described in Part I except that a heating unit of 500 watts was substituted for the small carbon-filament lamp. Since accurate estimations of small amounts of formaldehyde are possible by the use of the iodometric method of Romijn (Anatyst 1897 22 221; see also Ckm.-Ztg. 1901 25 74; Ber. 1898 31 1979; 1901 34, 2817) it was necessary to pass volumes of air of only 2 or 3 litres a t the most for each vapour pressure determination. During the course of a run involving several successive determinations on th 2834 LEDBURY AND BLAIR THE PAR- FORMALDEHYDE same solution the alteration of the strength of the formaldehyde solution in the " carburettors " was thus made negligible even in the case of concentrated solutions at 45" (compare Part I).De Waal ( P k m . WeekbZad 1907,44,1207; J . Phcbrm. Chim. 1907,26,498) has shown that in presence of air increase of temperature promotes oxidation of formaldehyde in solution but that this is slight even at 50" after an exposure for 400 hours. In the present series of determinations air was passed at the rate of about 1 litre in 3 hours. All estimations of formaldehyde and methyl alcohol were carried out 88 previously described and a similzr procedure of experi-mentation was followed. F d b y & Vapour Pressures of Aqueous Fumxddehyde Solutions For solutions of formaldehyde not containing methyl alcohol, the vapour pressure values at 35" obtained after the passage of several litres of air (for successive determinations) were slightly hqgher than those obtained at the outset.This increase however, wit8 not nearly so marked as at 20° and the eventual steady maxi-mum was more rapidly reached. Solutions containing methyl alcohol each gave a series of practically identical values from the beginning and this was also the case with both sets of solutions at 45". It is evident that at the temperatures under consideration, the new equilibrium between the various molecular species in solution is attained much more rapidly than at 0" or 20". The equilibrium values of the formaldehyde vapour pressures for form-aldehyde solutions free from methyl alcohol at 35" are in Table I, and the corresponding values for solutions containing methyl alcohol (CH,O/CH,O = 0.13) in Table II.(In Tables I-V g = grams of formaldehyde in 100 C.C. of formalin solution mt = mg. of formaldehyde vapour in 1 litre of issuing air at to and p = partial pressure of formaldehyde vapour in mm. of Hg.). Curves A and B in Fig. I represent the relationship between concentration of solution and partial pressure at 35' for pure solutions and for solutions containing methyl alcohol respectively. at 35" and 45". TABLE I. g ............ 1.09 5.15 11.8 18.6 20.8 31.0 39.5 T I Z ~ ~ ~ ......... 0.27 1.13 2-06 2-87 3.17 4.27 4.58 p ............ 0.166 0.695 1-29 1-80 1-94 2.48 2.81 TABLE 11.g (CH,O/CH,O = 0.13) ......... 5-24 13.05 23.7 38-3 ms6o ................................. 1.200 2.52 3.97 5.39 p ....................................... 0-735 1-56 2.45 3-3 vmom PmssmEs OF AQUEOUS SOLUTIONS ETC. 2835 The data obtained from solutions at 45" are in Tables I11 a;d Iv, and the curves showing variation of formaldehyde vapour pressure wifh solution concentration are plotted in Fig. 2. TABLE 111. g ....................................... 10.8 20.4 28.75 m4; ................................. 3-77 6:17 7.70 p ....................................... 2.30 3.79 4-72 TABLE IV. g (CH,O/C&O = 0.13) ......... 9.76 19.15 27.0 m450 ................................. 3-88 6.62 8-62 p ....................................... 2.39 4.07 5.3 FIG.2. 4.0 :::E 1.0 0 1 1 j v i I 50 20-0 30.0 40.0 10.0 20.0 30.0 40.0 G m . of formaldehyde per 1OOc.c. of eoh4timt. A methyl &hot ob8mt. B CH401CH,0=0.13. Variation with Ternperdure of the Partial Formaldehyde Vapou~ Pressures of Aqueous F d d e h y d e Solutions. From data derived from the curves drawn to represent the variation of partial formaldehyde vapour pressure with solution concentration at particular temperatures vix. 0" 20" 35" and 45", the graphs of Fig. 3 have been plotted to illustrate the manner in which the partial pressures change with temperature. The latter curves are plotted for solutions containing respectively 5 10 20, 30 and 40 g. of formaldehyde in 100 C.C. It is apparent that the differences between the partial pressures of the more concentrate 2836 LEDBUBY BND B= THE P ~ T I A L FORMALDEHYDE solutions at a given temperature are less marked than in the case of the weaker solutions.The data from molecular-weight deter-mimtiom the low values of the partial formaldehyde vapour pressures at the temperatures under consideration and the gradually dimini8hing dplclg of the vapour pressure-concentration curves demonstrate the increasing preponderance of polymerides in formalin solutions with increasing concentration. The relatively high FIG. 3. Temperature. concentrations of polymerised formaldehyde in the stronger solu-tions explains the particular characteristic of the curves referred to above. When the logarithms of the partial formaldehyde vapour pressures (log p ) are plotted against the reciprocals of the absolute tem-perature ( l / T ) for solutions of given strengths a series of almost parallel straight lines is obtained.In Fig. 4 the curves indicate the straight-line relationships between log, p and 1/T for solutions having formaldehyde contents of 10 20 30 and 40 g. per 100 c.c. VAPOUB PBESSUBES OF AQUEOUS SOLUTIONS ETC. 2837 respectively. Thus the formaldehyde partid vapour pressure of my aque~us formaldehyde solution up to a 40% concentration can be expressed for the temperatures under consideration by 8 simpMed form of Rankine's equation (Edinburgh New Phil. J., July 1849) connecting vapour pressure md temperature : log, p = a - B/T, where ale% = 9-47 ~ ~ 2 0 % = 9-70 &30% = 9.81 a@% = 9.87 (approx.) and p (for all solutions up to a%) = 2905 (approx.).A efraight-line relationship is somewhat deviated from when loglop for any given solution strength is plotted against T. The formula log p = a + pT + yT2 + . . . . which Perman (J. 1903, FIG. 4. 0*00370 0.00360 0*00350 F; 0*00340 040330 040320 0.003 10 88 1168) successfully applied to aqueous solutions of ammonia, was found to be inapplicable t o aqueous solutions of formaldehyde. blethyl-alcoholic Solution of FormuZd&y&.-Prior to the determin-ation of the partial formaldehyde vapour pressure at 20" of a methyl-alcoholic solution (15.9 g. CE,O per 100 c.c.) blank determin-ations were carried out on a sample of the methyl alcohol employed, since although this alcohol was the purest obtainable it waa not quite free from acetone (less than 0.10%).At 20° the blank values were equivalent fo about 0-1 mm. of formaldehyde estima-tions being carried out iodometrically. Further vapour-pressure determinations have been carried out at 20" and 0" on formaldehyde in methyl alcohol-water solution the alcohol being present in preponderating amount. This latter solution was obtained b 2838 LEDBTJRY AND BLAIR THE PARTIAL FORMALDEHYDE diluting with water the original 15% methyl-alcoholic solution in order to obtain a liquor containing 10 g . of free formaldehyde per 100 C.C. Blanks were performed at 20" and 0" on a mixture of methyl alcohol and water in which the proportions of the two components were the same as in the formaldehyde solution examined. The data obtained are in Table V solution A contained 15.97 g .of free formaldehyde (equivalent of 5.0% formaldehyde as methylal) in 100 C.C. of methyl-alcoholic solution; solution B contained 10 g. of free formaldehyde 59 g. of methyl alcohol and 26.7 g. of water in 100 C.C. TABLE V. Solution. t . mp. P. A ................................. 20" 3-68 2-27 B ................................. 20 1-93 1.16 B ................................. 0 0.390 0.234 It will be seen that at 20" the partial formaldehyde vapour pressure of an approximately 15% solution of formaldehyde in methyl alcohol is almost five times the value of that for an aqueous solution of corresponding strength. For a solution containing about 60% of methyl alcohol there is an approximately threefold increase at both 20" and 0" over an aqueous solution containing a corresponding amount of formaldehyde (10 g.per 100 c.c.). In all cases practically constant vapour pressure values were obtained from the outset. Since methylal vapour carried forward from the methyl-alcoholic solution to the absorption worms does not affect the iodometric estimation of the formaldehyde fixed by the water (Bergstrom, J. Amer. Chem. Soc. 1923,45 2150) the figure obtained represents the partial vapour pressure of uncombined formaldehyde in methyl-alcoholic solution. If it be assumed that the methylal (which is present to the extent of the equivalent of 5 g. of formaldehyde per 100 C.C. of 15% methyl-alcoholic solution) does not exert a very pronounced influence in its solvent capacity on the partial vapour pressure of the free formaldehyde in solution then it is found that the application of the simple mixture rule to a 15% aqueous solution in which CH,O/CH,O = 0.13 by no means accounts for the enhanced value a t 20" brought about by the addition of the alcohol to the original aqueous formalin i.e.from 0.49 to 0.59 mm. Hg. It has been pointed out that the removal of methyl alcohol from commercial formalin (40% by vol.) increases the tendency of para-formaldehyde to separate from solution. Experiments have shown, however that paraformaldehyde is more readily soluble in water than in methyl alcohol ; thus the precipitation of paraformaldehyd VAPOUR PRESSURES OF AQUEOUS SOLUTIONS ETC. 2839 from formalin which has been freed from methyl alcohol cannot be attributed to the removal of a constituent having a greater solvent capacity for formaldehyde.It therefore appears that the presence of methyl alcohol in aqueous formaldehyde solutions must bring about an alteration of the equilibrium between simple and complex molecules in such a way as to enhance considerably the concentra-tion of mono-molecular formaldehyde. The consequent decrease in the concentration of the polymerides naturally decreases their liability to precipitation as paraformaldehyde. Such an effect on the above-mentioned equilibrium serves to explain why methyl alcohol present in small amount in an aqueous formaldehyde solution increases the partial formaldehyde vapour pressure to an extent beyond that anticipated from the application of the simple mixture rule.It is here assumed as hitherto that the formaldehyde vapour pressure is almost entirely dependent on the concentration of the simple molecular form of the aldehyde. Xummar y . The partial formaldehyde vapour pressures of aqueous solutions of formaldehyde free from methyl alcohol have been determined by the dynamic method at 35" and 45". Corresponding determinations, carried out on solutions of formaldehyde containing small amounts of methyl alcohol (CH,O/CH,O = 0.13) have shown that the influence of the alcohol in enhancing the vapour pressures is similar to that noted at 20". The effect of increased temperature in bringing about a more rapid adjustment of equilibrium conditions has also been shown. From the vapour pressure values a relationship connecting functions of the vapour pressure and the temperature has been derived. The partial formaldehyde vapour pressure of a methyl-alcoholic solution of formaldehyde has been determined and the fiuence of methyl alcohol on the formaldehyde vapour pressures of aqueous formaldehyde solutions considered. The authors desire to express their thanks to the Department of Scientific and Industrial Research and to the Admiralty for permission to publish these results. ROYAL NAVAL COBDITE FACTOBY, HOLTON HEATH. [Received S e w e r lab 1926.

ISSN:0368-1645    

DOI:10.1039/CT9252702832

出版商:RSC    

年代:1925

数据来源: RSC

摘要:

2840 OLDHAM TRANSFORMATIONS OF THE SUGAR NImTES. CCCXCII.-Transforrnation~ of the Sugar Nitrates. By JOHN WALTER HYDE OLDHAM. ALTHOUGH the nitrates of the sugar group are as's rule definitely crystalline and eaaily adaptable to synthetical operations little use seems to have been made of them since they were studied by Will and Lenze (Ber. 1898 31 68) and Koenigs and Knorr (Ber. 1901, When 2 3 5-trimethyl glucosan dissolved in chloroform is acted on by fuming nitric acid under the conditions described in the experimental part tt.imethy1 glucose dinitrate is obtained. The result is striking particularly M the action of the same reagent on 1 6-diacetyl2 3 5-trimethyl glucose yielded a syrup which failed to crystallise. Trimethyl glucose dinitrate has no action on Fehling’s solution either before or after acid hydrolysis.The nitrate group in position 1 can be replaced by methoxyl by boiling with methyl alcohol in presence of barium carbonate the reaction yielding trimethyl methyZglucoside 6-mononitrate. This mononitrate is converted on hydrolysis into the same trimethyl methylglucoside as is obtained from trimethyl glucosan and the nitrate group waa definitely allocated to the 6 position by the fact that on treatment with sodium iodide the same trimethyl methylglucoside iodohydrin wtts readily produced as that already obtained from Fischer’s acetodibromoglucose (Irvine and Oldham this vol. p. 2729). This sequence of changes indicatea that the nitrate group possesses considerable reactivity yet the compound was recovered unchanged (as wm also methylglucoside 6-mononitrate) when heated with alcoholic ammonia as described by Wallach (Ber.1881 14 422). In the parallel case of triacetyl glucosan the action of fuming nitric acid was more complex and a yellow syrup showing a high nitrogen content was invariably produced together with trkwtyl glucose 1 6-dinitrateY which wm isolated in fair yield. The com-pound had likewise no action on Fehling’s solution before or after acid hydrolysis but was readily converted into derivatives of different types. On boiling in glacial acetic acid solution with sodium acetate it was converted into tetra-acetyl glucose 6-mono-nitrate ; and this on hydrolysis yielded a tetra-acetyl glucose isomeric with that already obtained by Fischer in that the sixth hydroxyl group in place of the ikst is unsubstituted.Further although triacetyl glucose dinitrate reacted with methyl alcohol only with great difficulty yields of 5cr-sO~0 of the corresponding triacetyl rnethylglucoside 6-mononitrate were obtained. This compound was converted into triacetyl methylglucoside on hydrolysis and into 34 957) OLDHAM TRANSFORMATIONS OF THE SUGAE "RATES. 2841 triaCetyl methylglucoside 6-icniohydrin on trestment with d u r n iodide in acetone in the manner described by Irvine and Oldham iKethyZgZuca&le 6-mononitrate obtained by the action of ammonia on the corresponding acetyl derivative could not be isolated as it could not be separated from acebmide. By the action of dimethyl-amine and removal of the acetodimethylamide however the mono-nitrate of the parent glucoside w a isolated &s a syrup which failed to crystallise.By means of the silver oxide reaction this compound was converted into the trimethyl methylglucoside 6-mononitrate already described thus estabhhing the constitution of the complete series of new derivatives. Triacetyl methylglucoside 6-iodohydrin, on treatment with dimethylamine in the manner already described, gave a somewhat poor yield of m&7&lwide 6-iodohydrin. Finally, incidental evidence was obtained in the come of large-scale pre-parations of triacetyl methylglucoside 6-mononitrate that degrad-ation to the pentose seriea had taken place to a limited extent. Synthetical and constitutional studiea based on the method of nitrate formation now communicated are in progress.(loc. cit.). E x P E R I M E N T A L. 2 3 5-Trimethyl Glucose 1 6-Dinitrate.-A solution of trimethyl glucosan (5 g.) in a mixture of 60 C.C. of fuming nitric acid and 40 C.C. of chloroform containing a few grams of phosphorus pentoxide was kept for 19 hours at the ordinary temperature. The specific rotation increased from -64" to + 140.7" calculated on the change in concentration due to the addition of two nitrate groups. The mixture was poured slowly with constant stirring into ice the lower layer separated washed once with ice-cold water and taken to dryness below 60" in a vacuum. The dinitrate crystallised from absolute alcohol in colourless needles m. p. 86" (Found OMe, 30.3; N 9.1. C,H,,OI& reqces OMe 29-8; N 9.0).The compound is sparingly soluble in alcohol insoluble in water or light petroleum and soluble in other solvenfs. [.ID + 149.3" in chloroform (c 2.3193) + 151.7" in acetone (c 1.463) + 144-8" in methyl alcohol (c 1.508) and + 147.2" in a mixture of 60% of fuming nitric acid and 40% of chloroform by volume (c 2.915). Trimethyl Methylglucoside 6-Mommitrak-The dinitrate WM boiled in 5% solution in methyl alcohol for 4 hour in presence of barium carbonate. The product isolated in the usual manner (yield W%) was drained on a tile and recrystallised from light petroleum; it then melted at 53-54" % of the superfused sub-stance = 1.4565 (Found C 42.7; H 6.7; OMe 44.1 ; N 4.95. C,&,,O$ requires C 42.7 ; H 6-8 ; OMe 44.1 ; N 5.0%). Th 2842 OLDHAM TRANSFORMATIONS OF THE SUGAR NITRATES.mononitrate is insoluble in water but soluble in all other solvents. [=ID - 5.2" in chloroform (c 2=0653) - 4.4" in acetone (c 3.5845), and - 1.3" in methyl alcohol (c 3-3413). Conversion of Trimethyl Methylglucoside 6-Mo?wnitrate into priwm%yl MethyZqlucoside.-The compound was hydrolysed by boiling with iron dust in glacial acetic acid in not more than 5% concentration until a few drops of the liquid gave no trace of blue colour on standing for some time with a strong solution of diphenyl-amhe in concentrated sulphuric acid. (This reaction is given by all these compounds and constitutes an extremely delicate test for the presence of a nitrate group.) The solution was then filtered the reaidue washed with glacial acetic acid and the iiltrate poured info water containing a fair quantity of dissolved sodium acetate the object of which was to suppress the hydrogen chloride set free by the chloroform with which the solution was repeatedly extracted.The trimethyl methylglucoside thus removed on recrystallisation from light petroleum showed the correct melting point and mixed melting point. The yield waq good. Conversion into the Cmaprmding ITodohydrin.-On treating the mononitrate with sodium iodide in the manner already described (%e and Oldham this vol. p. 2729) a good yield of crystals was obtained m. p. 31-33" which showed no depression on mixing with a sample of the authentic iodohydrin. Preparation of Triacetyl Glucose 1 g-Dinitrate.-The best con-ditions as yet found are the following Triacetyl glucosan (5 g.) was dissolved in a mixture of 30% of chloroform with 70% of redistilled nitric acid * by volume solid nitrogen pentoxide (10 g.) added and the whole made up to 100 C.C.with chloroform and kept at the ordinary temperature for about 110 hours. The specific rotation of the solution was then about + 138". Sometimes the liquid separates into two layers; should these not disappear towards the end of the reaction the system can be made homogeneous by the addition of a small measured amount of fuming nitric acid. The product was isolated in the manner described for the corresponding methylated derivative and washed with cold absolute alcohol (yield about 65%). On recrystallisation from hot absolute alcohol this yield was diminished by nearly half and the product melted at 132-133".The substance is insoluble in water light petroleum or cold ethyl alcohol sparingly soluble in ether or cold methyl alcohol, * In this paper the expression " fuming nitric acid " refers to the product obtained by distilling ordinary concentrated nitric acid over its own volume of concentrated sulphuric acid. " Redistilled nitric acid " refers to a product, d J.54 obtained by distilling "fuming nitric acid" over twice its volume of concentrated sulphuric acid OLDHAM TBANSFORXATIONS OF THE SWGB XITBATES. 2843 and soluble in other solvents. The rotation of the pure compound in a mixture of fuming nitric acid and chloroform similax to that used in the preparation wag [ah + 144.2" for c = 5.0. The substance has not been mlpd as it is diilicult to obtain a perfectly pure specimen, but in view of the derivatives into which it can be transformed there can be little doubt that it consists of triacetyl glucose 1 6-dinitrate.Tetra-cmtyl CZumse 6-blonondtPate.-The dinitrate was boiled for 1-14 hours with a mixture of acetic acid acetic anhydride and sodium acetate. If the dinitrate waa pure scarcely any colour ww developed; but if not the solution coloured deeply. In the former case the reaction was practically quantitative as waa shown by the specific rotation falling from + 144" to + 23.3" the corresponding value of the pure substance in acetic acid being + 23.2". The product was poured into water when about 80% of it separated almost immediately in a cryatalline condition; the remainder, however sometimes took several days to separate.The product after recrystallisation from absolute alcohol melted at 142-143", reduced Fehling's solution strongly on boiling ww insoluble in water or light petroleum sparingly soluble in ether or alcohol and soluble in other solvents (Found C 42-55; H 5.0; N 3-65. glacial acetic acid (c 2.494) and + 27.2" in chloroform (c 2-038). Tetra-acetyZ GZucose.-The 6-mononitrate waa hydrolysed with iron dust in glacial acetic acid exactly as described in the case of the corresponding methylated compound. Tetra-acetyl gZme crystal-lised from ether in colourless needles m. p. 1266-1276" [.ID + 9.8" in chloroform (c 0-816) (Found C,H40, after allowance for the alkali used by the reducing sugar 68.7. Calc. C2Ha02, 68.9 %).The substance is insoluble in cold water or light petroleum, sparingly soluble in ether soluble in hot water and in organic solvents generally. "he above reaction is rather uncertain as sometimes the product failed to crystallise. TriaCetyZ Methylglucoside 6- Mononitrate.-Triacetyl glucose di-nitrate unlike both acetonitroglucose and the corresponding methylated derivative reacted only with great diEculty with methyl alcohol in presence of barium carbonate even when a small quantity of pyridine was added as recommended by Koenigs and h o r n (h. cit.). On boiling the mixture for 24 hours the specific rotation of the solution became constant at + 5.0". The solution WM atered and the solid well washed with chloroform. The washings were used to dissolve the residue left on evaporation of the methyl-alcoholic filtrate in order to purify it from trams of barium nitrate.After removal of the chloroform the red residue crystallised twice C&,O1& requires C 42.7; H 4.8; N 3.6%). [.ID + 23.2" in VOL. CXXVII. 5 2844 OLDELAX TRANSFORMATIONS OF THE SUGAR NITRATES. from absolute alcohol melted at 1336--1346" (yield 50-60%), and was insoluble in water or light petroleum sparingly soluble in ether or cold ethyl or methyl alcohol and soluble in other solvents (Found C 42.8; H 5.35; N 3.8; OMe 8.6; CH3-C02H 66.0. C,H190,,N requires C 42.7 ; H 5.2 ; N 3.8 ; OMe 8.5 ; CH3*C02H, 65.7%). [a], - 14.3" in chloroform (c 5.964) and - 14.1" in acetone (c 206913). If the substance is boiled with methyl-alcoholic sodium methoxide the specific rotation falls to - 115" allowance being made for the change in concentration due to loss of acyl groups, which corresponds to the formation of about 80% of Fischer's anhydro-methylglucoside.Trkeetyl .M&hylglutmide.-On hydrolysing the nitrate with iron dust in glacial acetic acid in t'he manner already described t r k t y Z methylglutmide was obtained in good yield in colourless crystals, melting after recrystallisation from ether a t 136134.5'; [a]= in chloroform = - 19.1" for c = 1.514 (Found OMe 9-5; CH3*CO&I 54.6. c13H&g requires OMe 9.6 ; CH3*C02H 56.2%). The compound is insoluble in cold water or light petroleum but soluble in other solvenh including hot water. Met h ylgluwside 6 - Mononit rate. -Tr iacet y 1 met hylgluc oside 6 -mononitrate was treated with a 5% methyl-alcoholic solution of dimethylamine and the acetodimethylarnide produced was distilled off in a vacuum at 100"; the residual syrup failed to crystalbe (Found N 5.9; HNO, 25.7.C,H130& requires N 5.85; €€NO3, 26.3%). A specimen of the substance was twice methylated by the silver oxide reaction ; the product crystallised on nucleation with tri-methyl methylglucoside 6-mononitrate obtained as described above. &r draining on a tile and recrystallisation from light petroleum, the compound showed the correct melting point and mixed melting point thus proving the constitution of this series of substances. Triacetyl Methylglucmide 6-Iodohydrin.-For the sake of compari-son triacetyl methylglucoside 6-mononitrate was treated with sodium iodide in acetone in the manner already described.The product after recrystallisation. from aqueous alcohol consisted of colourless needles m. p. 111-112~5" which were insoluble in water or light petroleum but soluble in other solvents. The yield was poor [Found OMe 7.4; I 28.0; CH,*CO,H (I being calculated as CH3*C0,H) 564. C,Hl9O8I requires OMe 7.2; I 29.5; CH3*C0,H 55.8y0). The specific rotation in chloroform was + 0-9" for c = 3.027. -Methylglumside 6-Iodohydrin.-On treatment with dimethyl-amine in the manner already described the precedmg triacetyl derivative was converted into the parent glucoside in bad yield. The product which consisted when pure of colourless crystals bu WEEKS LEAD DIEYDRIDE AND LEAD TE-RZDE.2846 was usually slightly pink melted at 157-158" was insoluble in ether or light petroleum very sparingly soluble in chloroform, slightly more soluble in acetone and soluble in other solvents. The best recrystallking medium was a mixture of chloroform and ethyl acetate (Found OMe 10.3; I 41.7. C,H130,1 requires OMe, 10.1; I 41.7%). Acetonitropentose.-On taking to dryness the mother-liquors from the recrystallisation of triacetyl methylglucoside 6-mononitrate, boiling for 4 hour with acetic anhydride and sodium acetate and pouring into water a precipitate wits obtained which after very many recrystallisations from absolute alcohol melted at 168-169". The compound did not reduce Fehlmg's solution and was a t first thought to be the a-form of triacetyl methylglucoside 6-mononitrate. It contained no methoxyl however and the analytical figures are in close agreement with those required for a triacetyl pentose mononitrate (Found C 41.2; H 4.7; N 4.2. C,lH,,Ol,,N requires C 41.1 ; H 4-7 ; N 4.4%). The substance had the same solubilities as triacetyl glucose dinitrate and showed spec& rotation in chloroform = + 92-0" for c = 1.214. [.ID in chloroform = - 16.1" for c = 1.677. The author wishes to express his indebtedness to the Carnegie Trust and also to Principal Sir James C. Irvine for much valuable advice and for the kindly interest he has always shown in the work. UNITED COLLEGE OF ST. SALVATOR AND ST. LEONARD, UNIVERSITY OF ST. ~ D R E W S . [Received October 26th 1925.

ISSN:0368-1645    

DOI:10.1039/CT9252702840

出版商:RSC    

年代:1925

数据来源: RSC

摘要:

WEEKS LEAD DIEYDRIDE AND LEAD TE-RZDE. 2846 CCCXCII1.-Lead Dihydride and Lead Tetrahydride. By EDWARD JOSEPH WEEKS. SOLID hydrides of arsenic (Ckm. News 1924,129 31) antimony (this vol. p. 1069) and bismuth (ibid. p. 1799) having been pre-pared the existence of a solid lead hydride was investigated. To a solution of alkali plumbit'e made from lead acetate and caustic soda or potash pure aluminium foil was added. The grey deposit obtained was washed many times with caustic potash solution and h l l y with water until the washings were neutral. It waa filtered off and dried in a vacuum desiccator over sulphuric acid for 3 4 5 days both operations being performed in an atmosphere of hydrogen. The action follows the equation 2KHPb0 + 2Al= 2KAl0 + Pb& The deposit contained lead and hydrogen only and on heating in a vacuum gave off hydrogen and left metallic lead.Details of the method of analysis have already been given for bismuth (h. cit.). 0.3253 G. gave 20 C.C. of H measured at It.7'2. and 0-3243 g. of 5 ~ 2846 GLASSTONE AND RIGGS CO-X FOEBfA!L'ION Pb. H 0.5; Pb 9901%. 0.1504 G. gave 10 C.C. of H at N.T.P. and 0.1422 g. of Pb. H 0.6 ; Pb 99.5%. On combustion 1.9550 g. gave 0.0662 g. of H,O. H 0.5% (Pb,H requires Pb 99.5; H, 0.5%). Lead dihydride heated in a tube in the absence of air gave lead and hydrogen only. No trace of PbH appeared to be formed as no lead deposit could be obtained on heating the issuing gas. The dihydride oxidises rapidly in the air and therefore must be kept in an inert gas. Fused potassium nitrate reacted vigorously with it (as with As2H2; Sb,H,; Bi2H2) but only slowly oxidised h e l y divided lead. To obtain the tetrahydride first prepared by Paneth the dihydride was heated in a silica tube in an atmosphere of pure hydrogen and the issuing gas was passed through a heated tube ; a deposit of lead was then obtained. With hydrogen alone no deposit was formed and hence it is concluded that the reduction of Pb,H follows the equation Pb2H + 3H2 = 2PbH,. BA~ERSEA GEAMMAB ScHooL, LONDON S.W. 11. [Received November 2nd 1925.

ISSN:0368-1645    

DOI:10.1039/CT9252702845

出版商:RSC    

年代:1925

数据来源: RSC

摘要:

2846 GLASSTONE AND RIGGS CO-X FOEBfA!L'ION CCCXCIV.-Complex Formation in Lead Nitrate Solu-tions. Part I I . The Quaternary System Potassium Nitrate-Lead Nitrate-Barium Nitrate- Water. By SAMUEL GLASSTONE and ERKEST J. RIGGS. THE primary object of the present work was to investigate the quaternary system lead nitratebarium nitrate-potassium nitrate-water a t 25" and at 50" and hence it was required t o have a knowledge of the three ternary systems KN0,-Pb(NO,),-H,O KN0,-Ba(NO,),-H,O and Pb(N0,)2-Ba(N0,)2-H,0 at the same temper-atures. The h t of these three systems has been investigated by Glasstone and Saunders (J. 1923 123 2134) a t 25" and 50"; the second has been examined at 9.1" 21.1" and 35" by Findlay Morgan, and Morris (ibid. 1914,105 779) and to a limited extent a t 25" by Foote (Amer.Chem. J . 1904 32 251); the 25"-isothermal for the third system was investigated by Fock (2. Kryst. 1897 28 337). Before the quaternary system was examined therefore Foote's measurements on the ternary system KNO,-Ba(NO,),-&O at 25" were exfended and the 50"-isotherm was determined completely ; Fock's work a t 25" was repeated because this author did not use the Schreinemakers residue method for the determination of the com-position of the solid phase and the 50O-isotherm for the syste TX LEAD "RATE SOLUTIONS. PABT II. 2847 Pb(N0,),-Ba(N03)2-H,0 was also determined. Finally the quaternary system Pb(N0,)2-Ba(N03)2-KN03-€&0 was investi-gated as completely as possible at 25" and 50". It was hoped to obtain through this investigation further evidence for the existence of a double or complex salt of lead nitrate and potassium nitrate, the presence of which has already been suspected in solutions con-taining these two salts (see Glasstone and Saunders loc.cit.). Lead and barium nitrates have an almost identical crystal structure (Vegard 2. Physik 1922 9 395) and separate from solutions con-taining both salts as a continuous series of mixed crystals; further, potassium and barium nitrates form a double salt 2KN03,Ba(N03), (Foote h. cit.) and consequently it was thought possible that when this double salt separated from a solution containing lead barium, and potassium nitrates there might be a tendency for it potassium nitratelead nitrate double salt with a similar structure which has no stable existence at the ordinary temperature to separate with it as a mixed crystal.This idea was based on the well-hown existence of ~ 0 7 H 2 0 in the form of a mixed crystal with FeS04,7H20, and %he recent preparation by Richards and Meldnun ( J . Amr. C h . Soc. 1921 43 1543) of Na,S0,,4€&0 which does not exist by itself as a mixed crystal with Na2Cr04,4H,0. The resulb of the present work indicate that lead nitrate can exist in some form which is isomorphous with either potassium nitrate or with the potassium-barium double nitrate; the most probable form is as the double salt 2KN03,Pb(N03), which hm no stable existence at the ordinary temperature but might separate as a mixed crystal with E X P E R I M E N T A L . General Prdure.-In the case of the second of the ternary systems mentioned above the procedure was the same as that followed in the first part of this work (Glasstone and Saunders h.cit.); for the third ternary system and the quaternary system, however owing to the deposition of mixed crystals it was essential that very little of the solid phase should separate and consequently the mixed solids were made fo dissolve in water at a temperature just above that at which the isotherm was being determined and the solution was stirred in a thermostat until sufficient solid for analyais had separated as the result of evaporation. Method of AnaZysis.-The amounts of water in the saturated solutions and wet solids were determined by drying known weights first on the water-bath then in an air-oven at l l O o and finally at 130".The dry solid was then dissolved in water made up to a known volume and an aliquot portion taken for analysis. If the ~ ~ O ~ Y ~ ( N O ~ ) 2848 BLASSTONE AND RIGBS COMPLEX FORMATION solution contained only potassium and barium nitrates the barium was precipitated directly M sulphate with sulphuric acid and the barium sulphate was washed dried and weighed. When the solu-tion for analysis contained lead and barium nitrates only hydrogen sulphide was passed into the warmed solution until all the lead was precipitated; the precipitate of lead sulphide was filtered off, washed with a solution of hydrogen sulphide and the barium in the filtrate estimated as sulphate. The amount of lead was usually found by difference but in several cases the sulphide was dissolved in hot dilute nitric acid and the lead estimated as sulphate by the method described in Treadwell’s ‘‘ Quantitative Analysis ” (1919, p.174). When the solution to be analysed contained all three nitrates both the lead and the barium were determined by the methods described above and the potassium was obtained by difference. The analytical method was tested beforehand on mixtures of known weights of the various nitrates and found to be quite satisfactory IN LEAD NFI!BATE SOLUTIONS. PdaT II. R&.-These are all given at3 percentages. /--KNO,. 27.39 27.67 17-14 9-13 3.80 0.00 *14-8 k0,. 45.45 44.88 44.66 40.3 1 35-75 30-08 28-86 24-14 18.76 10-53 3.93 0.00 2849 KN0,-Ba(NO,),-H,O at 25".Solution. Rest. Ba(NO,),. H,O. KNO,. Bs(NO,),. H,O. Solid phase. - / -.-- -. -0 0.00 72.61 - - -2.44 69-89 78-65 4-84 16.51 KNO + D.S. 4-88 77.98 39-31 49.67 11.02 D.S. 6.6 78-6 - - - D.S. + Ba(NO,), 6-57 84.30 1-18 92.26 6.56 Ba(NOS)* 7.72 88.48 1-05 91-18 7.77 I9 9.28 90-72 - - -D.S. refers to the double salt 2KNO,,Ba(NO,),. * Taken from Foote Eoc. cd. 99 Solution. c BWO,),. 2.34 4-22 5.11 5.64 6.99 10-16 10-81 10.43 10.14 10-92 12.73 14.63 KN03-7 H,O. 52-21 50.90 50.23 54-05 57.26 59-76 60.33 65-43 71-10 78-55 83-34 85-37 -Ba(NO,),-H,O at Rest. . - .. ___ KNO,. Ba(NO,),. 87.05 0.65 84.20 1.25 74-36 3-87 42.78 50.55 42-38 46.05 41-56 50.93 2-18 92-55 0.02 95.39 2-28 93.48 1-33 95.39 1.07 95.85 - -50".H,O. 12.30 14.55 21-77 6.67 11.57 7-51 5-27 4.59 4.24 3.28 3.08 -Solid phase. I&O + D.S. D.S. D S . + Ba(NO,), =o, S Y Ba(NOd2 The results obtained for this system are represented in Fig. 1. Pb( NO,),-Ba( NO,),-H,O at 25". Solution. Pb(N0a)p Ba(N0A. 34.60 1.39 31-13 2.65 22.73 4.23 16.21 5.29 8.96 6.66 6-26 7-31 1.63 8.64 -Solution. - Pb(N0a)v Rs(h'O,)z 40.82 2.57 37-78 3.83 29.31 6-38 17-05 9.12 10.34 11-05 5.72 12-85 -7 HtO. 64-01 66.22 73-04 78-50 84-38 8643 89.73 Rest. Solid phese. Pb(NOa)a* 8 1-83 66.90 38.08 15.06 10.15 3.44 -Ba(NO,)p 11.26 18.90 56.67 80-02 84.28 89-54 -Y HSO. 6.91 14.20 5-25 4-92 5-57 7.02 -Pb(NOJ2.87.54 76-56 39.27 15.44 10.42 3-59 -Ba(NO;h. 12-46 23-44 60.73 84-66 89-58 96-41 -Pb (NO,) 2-Ba (NO,) 2-H,0 at 50". Rest. Solid phase. _- - - -_ H,O. Pb(NOa)z. Ba(N0,):. HSO. P/b(NO,h. Ba(XO;)*. 56-61 77-06 9.54 13.40 88.31 11.69 58-39 63.40 26-10 10.50 69-02 30.98 64-37 40.31 55.14 4.55 41.25 58-75 73.83 78-61 8-92 83-87 7-21 8.78 91.22 81-60 6-60 87.70 5.70 6.67 93.33 - I - - 2850 GLASSTONE AND ~ Q B S COMPLEX FORMATION The results obtained for this system we represented in the usual mmner in Fig. 2. In the columns headed " solid phase " me given ma. 2. H20 FIG. 3. FIG. 4. 25' yo Pb(NO,) in Pb(NO,) + Ba(NO,) in solutions. Circles rqreent point8 from the ternary system; square.8 points from the q u a t e m r y aystem.the compositions of the mixed crystals which separate from the various solutions in the dry state ; these results were obtained fro IN LEAD KITRATE SOLUTIONS. PART II. 2851 the amlpb of the wet solid by allowing for the amounts of salt present in the adherent mother-liquor. Fig. 3 gives the gram-percentage of lead nitrate in the mixed crystals against the gmm-percentage of lead nitrate in the total salt dissolved in the solution from which the crystals separate at 25" whilst Fig. 4 gives the corresponding curve at 50". KN03-Pb(N03)2-Ba(N03)2-H20 at 25". Solution. Dry solid. 7- - KNO,. Pb(N0,)- Ba(NO,)*. H,O? &Ow Pb(N0,)I. Ba(NO,).. phases. 15.26 2.12 6-05 76-57 43.23 0.10 56.67 D.S.+ M.C. 17.39 11-44 3-40 67.77 41-98 0.43 57.59 ,, 15-99 13.82 4.76 65-43 43.45 0.80 55.75 ,, 16.01 21.25 4-33 58-41 43-51 1.70 54.79 ,, 15.65 26.59 4.06 53-70 0.88 31-36 67.76 ,, 17.87 35.90 2.46 43.77 23-10 14.95 61.95 ,, 23-41 37.68 2.11 36.80 23.78 38-10 38.12 25.75 34-65 1.16 38.44 69.45 4-42 26-13 D.S.? M.C.+ I(N0, 26.78 23.26 1.76 48.20 92-31 0.61 7.08 99 27-23 15-28 2-32 55-17 96.06 0-23 3.71 Y Y 27.69 9.85 1.87 60.59 92-13 0.10 7-77 1 3 27.74 4-95 1.92 65-39 63.28 0.30 36.42 25.22 40.46 0-72 33.60 45.44 51-86 2-70 mo:'+ M.C.25-00 39.84 1-18 33-98 2-11 28-02 69.87 Y Y 25-00 38.00 1-59 35.41 10.61 76.19 13.20 9 9 24-91 37-48 1.80 35.81 94.90 1.31 3.79 9 , 25.74 32.63 1.92 39.71 91.23 1-36 7.41 D.S.+ -0, M.C. refers to Pb(NO,),-Ba(NO,) mixed crystal. KN03-Pb(N03)2-Ba(N0,),-H,0 at 50".Solution. Dry solid. ENO,. 27.40 28.28 32-61 35.48 36.91 38-06 39-43 40.90 41-39 41-86 33.25 33-41 32-64 Pb(NOa)** 15-89 27.00 38-93 33.62 29-22 24-98 19-96 14.64 13.01 10.35 40.88 40.19 39-82 Ba(NOdI. 8.11 6-68 2.15 2-01 2.24 2.59 2-85 3-32 3.44 3-69 1-02 1-51 2-40 - HZO. 48-60 39-04 26-31 28-89 ' 3 1.63 34.37 37-76 41-14 42.16 44.10 24.85 24-89 25.14 7 mow 1-04 9.4 1 15-02 91-53 92.65 65.31 57-61 64-03 71.12 97.69 53.25 34.54 46.76 Pb(NOt)r 6.15 11-94 47-42 1-02 0-00 1-62 0.94 0.65 0-44 0.19 43-15 35.58 46-12 B ~ ( N o ~ . Phases. 92.81 D.S.+ M.C. 78-65 ,, 37-56 7.45 D.S.? KNO, 7-35 9 , 33.07 Y Y 41.45 Y Y 45.32 9 , 28-44 9 , 2-12 3-60 KNO:'+ M.C.29.88 9 9 7.12 9 , The two isotherms in this quaternary system have been drawn from the above results by the method of Schreinemakers as this was found to give the least confusing diagram and are represented in Fig. 5. The points A By and C represent the composition of ternary liquids which are in equilibrium with two solid phases-potassium nitrate and double salt double salt and lead nitrate and potassium and lead nitrates reapectively-whilst the points on the 5 D 2852 GLASSTONE AND RIGGS COMPLEX FORMATION curves AD BD and CD give the composition of quaternary solutions in equilibrium with double salt and potassium nitrate double salt and mixed crystal and potassium nitrate and mixed crystal re-spectively.The point D represents the composition of the quater-nary isothermal invariant solution. The composition of the dry solid given in the above tables was obtained by calculating from the analysis of the solutions the amounh of the three nitrates dissolved in the water which was present in the wet solid and deducting FIG. 5. Ba"O,)* KNO, WNO,), these from the total amounts found by the analysis of the dried residue. An examination of the results given above for t'he quaternary system shows that when the two solid phases in equilibrium with saturated solution should consist only of potassium nitrate and double salt there is also present a small amount of lead salt; the results are so consistent as to rule out the possibility of experimental error.The presence of lead has been confirmed by washing the wet solids with a little water and grinding the residue at 30-35" with a solution containing roughly 25% of potassium nitrate and 2 IN LEAD NITRATE SOLUTfONS. PART 11. 2853 of barium nitrate which dissolves lead nitrate very readily. The solid was then filtered off and ground up with a fresh portion of the solution; this process was repeated several times but even after ten treatments the solid residue still contained an appreciable quantity of lead salt. These results indicate that lead nitrate in some form is separating as a mixed crystal either with potassium nitrate or else with the double salt. The most probable explana-tion is that the double salt 2KN03,Pb(N03)2 which has no stable existence alone under ordinary conditions is separating out from solution as a mixed crystal with the barium double salt of s i m h formula.Other explanations are of course possible but they seem to be less probable. In order to obtain if possible further evidence for the existence of the double salt of lead and potassium nitrates the results for the systems from which the potassium-barium double nitrate and mixed crystals of lead and barium nitrates separate were examined &B follows. It was msumed as a first approximation that all the potassium nitrate in the solid phase was in the form of the double salt with barium nitrate and so the aaount of barium nitrate present as double salt was calculated; the residual barium nitrate was assumed to be present as mixed crystal with lead nitrate.In this way the ratio of lead nitrate to barium nitrate in the mixed crystal was determined and compared with the ratio of these two salts in the quaternary solution. The results are given below. Pb(NO*) Pb (NO,), Pb(NO,) + Ba(NO,) %* '0° Pb(NO,) + Ba(NO,) %* 25' Solution. Mixed crystal. Solution. Blixed crysfal. 94-69 82.49 94.75 72.35 86.76 32.00 82-62 15-09 77-07 12-78 66.20 6.32 25.95 10-67 These results have been plotted in Figs. 3 and 4 for 25" and 50" respectively for the purpose of comparison with the results for the ternary system. It is seen that for a given ratio of lead nitrate to barium nitrate in the solution the mixed crystal separating from the ternary solution contains relatively more lead than does that separating from the quaternary system ; if a lead-potassium nitrate double salt had been present then the reverse would have been expected. It is possible however that the presence of pohsium nitrate in the solution alters the ratio of lead to barium in the mixed crystal and so no definite conclusiom can be drawn frgm the results. Summary. (1) The ternary systems KN0,-Ba(N0,)2-H20 and Pb(N03),-Ba(NO3),-H2O have been investigated a t 25" and 50". 5 D * 2854 KNECHT AND EIIBBERT THE BE€IAVIOUR OF GLUCOSE AND (2) The quaternary system KN03-Ba(N03)z-Pb(N03)z-H20 has been investigated at 25" and 50". (3) There is shown to be some evidence for the existence of a double salt 2KN03,Pb(N0,), in the form of a mixed crystal with the double salt 2KNO3,Ba(hTO3),. UNIVERSITY COLLEGE EXETER. [Received November 3 4 1925.

ISSN:0368-1645    

DOI:10.1039/CT9252702846

出版商:RSC    

年代:1925

数据来源: RSC

摘要:

2854 KNECHT AND EIIBBERT THE BE€IAVIOUR OF GLUCOSE AND CCCXCV.-The Behaviour of Glucose and Certain Other Carbohydrates towards Dyestufjs and towards Potassium Ferricyanide in an Alkaline Medium. By EDMUND KNECHT and EVA HIBBERT. ALTHOUGH we are well informed as to what changes take place in the dyestuff when indigo or some other dyestuff is reduced by glucose in presence of alkali no attempts to elucidate the character of the oxidation of the glucose have hitherto been recorded. Many attempts have however been made to ascertain the character of the oxidation products formed by the action of Fehling’s solution and of cupric hydroxide in presence of caustic alkali as well as of other inorganic oxidising agents in an alkaline medium the results of which indicate that the reactions are of a complicated character.The most complete researches on this subject are those of Gaud (Cornpt. rend. 1894 119 604) who employed Fehling’s solution as the oxidising agent and obtained formic oxalic tartronic lactic and glyceric acids along with some catechol and the elaborate work of Nef (AnnaZen 1907 357 214) who using copper sulphate and a slight excess of caustic soda obtained hexonic trihydroxybutyric, glyceric glycollic formic and carbonic acids. By means of red mercuric oxide and barium hydroxide Herzfeld (Annalen 1888, 26 32) obtained gluconic acid as the main product. By using sodium hypoiodite as the oxidising agent Romijn (2. a d . Chm., 1897 36 19) obtained this acid alone and Willstatter and Schudel (Ber. 1918 51 780) showed that the reaction may serve for the estimation of glucose.The authors have shown ( J . SOC. Dyers Col. 1925 41 94) that when glucose is boiled with excess of methylene-blue in presence of caustic soda oxidation takes place very rapidly and to a definite degree (three atoms of oxygen to one mol. of glucose). Further work in this direction has now been carried out. Methylene-blue was not entirely suitable for the purpose on account of its great liability to undergo decomposition by the action of the alkal CERTNN OTHER CARBOHYDRATES TOWARDS DYESTUFFS ETC. 2855 (Bernthsen Annden 1885 230 73) so potassium indigotintetra-sulphonate which reacts in the same way as methylene-blue is much more stable towards alkalis and can be estimated volu-metrically with rapidity and accuracy was used for most of the oxidations given below.The amount of indigotin reduced was invariably equivalent to three atoms of oxygen per mol. of glucose, four atoms of oxygen per mol. of laevulose and as might have been expected 3& atoms of oxygen in the case of invert-sugar. Galactose and glucosamine hydrochloride behaved like glucose each taking up three atoms of oxygen. Maltose was apparently hydrolped, since the molecule (C,2H22011 + H20) reduced an amount of indigo-tin equivalent to six atoms of oxygen. Thioindigodisulphonic acid (Ber. 1907,40,3821) gave with glucose and laevulose results identical with those obtained with potassium indigotintetrasulphonate. Potassium ferricyanide can replace indigotin in the titration of glucose laevulose and glucosamine giving identical results ; the reaction is very rapid being complete in 15 seconds.The extent of the oxidation was in the &st instance determined by titrating the ferrocyanide in the cooled solution midiiied with sulphuric acid with permanganafe to the appearance of a pink colour. The result was corroborated by estimating the excess of ferricyanide by means of standard sodium hydrosulphite with methylene-blue as indicator. Ultimately it waq however preferred to estimate the excw ferricyanide by titration with titanous chloride in presence of a trace of ferric chloride the disappearance of the blue colour indicating the end-point. There was a close agreement of the results in all three cases. In view of the constancy of the results obtained with methylene-blue potassium indigotintetrasulphonate thioindigodisulphonic acid and potassium ferricyanide and the exact stoicheiometric character of the reactions it was surmised that in the case of glucose and galactose the oxidation had resulted in the formation of hexane tetrahydroxy-dicarboxylic acids corresponding to sac-chasic and mucic acids whilst in the case of hvulose the oxidittion would just suffice to produce a mono- and a di-carboxylic acid in equimolecular proportions.In spite of numerous attempts it has, however not been possible to identify any of the oxidation products. The difEculty lies mainly in the removal of the large quantity of colouring matter required to effect the oxidation; for instance, 1 g. of glucose requires more than 16 g.of potassium indigotin-tetrasulphonate which will be present after oxidation is complete, partly as such but mainly in the reduced condition. Also, through the action of the alkali considerable isomerisation o 2856 KNECHT AND EIBBERT THE BEHAVIOUR OF GLUCOSE AND the oxidation products probably takes place and these do not possess the uniformity of the product of oxidation in an acid medium. If glucose which is neutral to phenolphthalein is converted into a dicarboqlic acid (or acids) during the reaction the extent of such conversion should be measurable by titration for the resulting acid would require two equivalents of caustic alkali for its neutralis-ation. On similar grounds the acids resulting from the alkaline oxidation of lamulose should require three equivalents of alkali.It was obviously impossible to use dyestuffs as oxidants to test the validity of this reasoning. Pure potassium ferro- and ferri-cyanides are neutral to phenolphthalein and the latter was employed as the oxidant in the experiments. In carrying out the estimations it had to be taken into consideration that potassium ferricyanide in becoming reduced to ferrocyanide takes up two molecules of potassium hydroxide for each atom of oxygen supplied. The results confirmed exactly the above reasoning. It is known that in the titration of glucose by means of Fehling’s solution the extent of the oxidation as measured by the amount of copper reduced is represented by rather less than 2$ atoms of oxygen per molecule of glucose. On the other hand it has been shown that by employing in place of Fehling’s solution the solution of copper carbonate in potassium carbonate and bicarb-onate advocated by Soldaini (Gazzetta 1876 6 322) the degree of oxidation is almost doubled (see also Ost Ber.1890 23 1035, 3003).* It was therefore considered possible that the use of an alkaline carbonate in the titration of glucose with indigotin might also give rise to a higher degree of oxidation. The results obtained, however were in no way modified by substituting for the potass-ium hydroxide an amount of potassium carbonate and bicarbonate equivalent to that employed in the alkaline copper carbonate solution. In carrying out the titration it is important that the substances be added in the sequence given. If the caustic potash is added to the boiling glucose solution before the indigotin 30 seconds’ boiling is sufficient to lower the result by 50% or more.This seems to be due to the conversion of some of the glucose into lactic acid (Nencki J . p. Chem. 1881 24 498). If the boiling is continued for 2 minutes the amount of alkali neutralised corre-sponds exactly to the formation of two molecules of lactic acid. Laevulose behaves in this respect like glucose. Neither lactic saccharic mucic nor gluconic acid reduced * No satisfactory explanation of this great difference in the behaviour of the two solutions has hitherto been advanced CERTAIN OTHER CARBOBYDRATES TOWARDS DYESTUFFS ETC. 2857 methylene-blue when boiled with this dyestuff in presence of was substituted for methylene-blue or potassium indigotintetra-sulphonate as oxidant both glucose and laevulose took up exactly two atoms of oxygen per molecule.This would correspond to the formation of glycuronic acid from the first and of hydroxygluconic acid from the second of these carbohydrates. Superimposing an inQotin titration on the kitone-blue titration of glucose resulted in a further atom of oxygen being taken up. From its behaviour towards methylene- blue and towards indigotintetradphonic acid, it might have been expected that lamdose after being oxidised with kitone-blue would take up two more atoms of oxygen on superimposing a methylene-blue titration. Only one further atom of oxygen was taken up. b i n d d i n e 2B when employed as the osidising agent supplied only one atomic proportion of oxygen to glucose.By super-imposing a methylene-blue titration no further oxidation took place; and although the oxidation product was not identified, this circumstance strengthens the opinion that the oxidation had resulted in the quantitative formation of gluconic acid. E x P E R I M E N T A L. In principle all the titIratiom are carried out in the same way. A known weight of the carbohydrate is heated with a known volume of a standard solution of the oxidising dyestuff which must be present in considerable excess in a current of nitrogen. When the mixture boils caustic potash is added and the boiling continued for the time specified. The mixture is now a t once acidified and the excess of dyestuff titrated wit'h titanous chloride in a current of carbon dioxide.The quantity of carbohydrate most suitable for a titration is 0-01-0.05 g. and the quantity of caustic potash required is about 0-5 g. or 5 C.C. of a 10% solution. The following special points should be noted in carrying out the titrations. Methylene-blue. A suitable strength of this dyest& (working with a titanous chloride solution of which 1 C.C. = 0.002 g. Fe or thereabouts) is 10 g. of the medicinal product per litre. The time of boiling in presence of alkali should in no case exceed 5 seconds; beyond this limit the methylene-blue loses strength through decomposition and the results become untrustworthy. A suitable strength is 40 g. Potassium ir,digotintetrasuIp~nate 2858 KNECHT AND -BERT rn BEHAVIOUR OF GLUCOSE AND of the crystallised dyestuff per litre.The procedure comkts in adding to an aqueous solution of a definite weight of glucose excess of the standard indigotin solution. The mixture is boiled in a current of nitrogen caustic soda is added and the boiling continued for 8 minute. The mixture is then acidified with dilute sulphuric acid excess of tartrate of soda added and the hot solution titrated with titanous chloride until the blue colour disappears. Ferricyanide. A suitable strength for a standard solution is 50 g. per litre. The quantity of caustic potash added should be more than sufficient to allow for the formation of ferrocyanide and for the neutralisation of the organic acids formed. The time of boiling is 15 seconds. The mixture is rapidly cooled in running water before being acidified with dilute sulphuric acid.The excess ferricyanide is then titrated with standard titanous chloride as described. It is not necessary to employ either nitrogen or carbon dioxide. Boiling for 30 seconds suffices to complete the oxidation after which the solution is acidified with tartaric acid in slight excess before the excess of dyestuff is titrated with titanous chloride. Strength and procedure as for kitone-blue. Rosinduline G also may be used for the purpose. Crystal-scarlet was not found suitable as an alkaline oxidising agent the end-point not being sharp. Like the rosindulines this dyestuff supplies one atom of oxygen to glucose. Titration of Glycuronic Acid.-A solution of pure euxanthic acid (0-201 9.) in sulphuric acid (d 1.735) was diluted boiled and filtered, and the separated euxanthone was washed so that the filtrate contained the whole of the glycuronic acid.The solution thus obtained reduced indigotin equivalent to 26.4 C.C. of Ticla (1 C.C. TiC1 = 0402001 g. Fe) indicating that almost exactly one atom of oxygen had been used. Expressed as percentage of glycuronic acid in euxanthic acid the result gives 45.5% as against 45.9% calculated for Cl,H,,Olo + H,O. Almost the same result was obtained in the osazone process 0.035 g. of the same specimen of euxanthic acid was hydrolysed with sulphuric acid as before and the resulting glycuronic acid converted into the osazone; this required 14.4 C.C. of TiCl (1 C.C. = 0.001924 g. Fe) (Found 4507% calculated on the euxanthic acid employed).Action of Sodium Hydroxide on Glucose and Lceevulose.-An aqueous solution of 0.5045 g. of glucose was boiled with 25 C.C. of N-sodium hydroxide for 2 minutes and the excess of alkali then titrated with N-sulphuric acid and phenolphthalein. The amount of caustic soda neutralised was 5.6 c.c. which is equal Kitone-blue. A suitable strength is 10 g. per litre. RosinduZine 223 CERTAIN OTHER CARBOHYDBATES TOWLELDS DYESTIEFIB ETC. 2859 to 44.4% on the weight of the glucose (calc. for N-NaOH, 0.5115 Gram of h d o s e similarly treated required 5.7 C.C. of Y-sodium hydroxide which represents 4404% on the weight of the laevulose. In another experiment the boiling with caustic soda waa continued for 5 minutes; this did not affect the result.Acidi'metric Titration of the Oxidation Products of Glucose and Lmlose.4lucose (0-5 g.) was oxidised with 7-5 g. of potassium ferricyanide and 50 C.C. of N-potassium hydroxide. Titration with N-sulphuric acid and phenolphthalein then showed that 1.24 g. of potassium hydroxide had been neutralised in the reaction. The amount required according to the equation is 1-245 g. On similar treatment of 0.5 g. of lmulose 1.45 g. of potassium hydroxide were neutralised (calc. 1.40 g.). Indig&i?a Titration supem'mped on Kitone-blue Titration.-0-02084 Gram of glucose required kitone-blue equivalent to 13.6 C.C. of T i 3 (1 C.C. = 0.001916 g. Fe) which (for 2 atoms of oxygen supplied) corresponds to 0.02094 g. of glucose or 100.4%. In the superimposed indigotin titration an amount of indigotin equivalent to 6.9 C.C.of TiC1 was required corresponding (for one additional atom of oxygen) to 0.02123 g. of glucose or 101.3%. 4445%). Titre of TiCl ( 1 C.C. 3 =- x10-r TiCI, Sugar. Oxidant. (g.). g . Fe). (c.c.). Glucose ......... ......... ......... ......... ......... ......... Laehose ...... 9 , 9 9 9 9 9 , ...... ...... ...... ...... Sucrose (inv.). .. Galactose ...... Arabinose ...... Glucosamine hydrochloride Glucoesmine hydrochloride Glycuronic acid in euxanthic acid ............ 9 , 9 9 ... ... 9 ...... M I Th F R K-b M I Th F K-b M I F M I I I F I 0.03956 0.0400 0.03733 0.02029 0.0400 0~0200 0.04006 0.02023 04400 0.0390 0-0200 0.0291 0.0227 0.0227 0.0388 0-03304 0*0400 0.05232 0.02505 0-2010 27.42 18.38 19-23 19-23 19.98 19-98 20.81 19-15 19.23 19-23 19-98 15-72 20-01 20.01 17-24 18.45 20-66 19.23 17-49 20.01 26.7 40.5 36-2 19.7 12-4 12.5 48- 1 26-2 51-5 50.5 26.2 42.5 25.9 26.0 42-0 33-4 42.4 42.3 22.2 26.4 sugar %of found sugar (g.).taken. 0.03922 99.15 0.03988 99.7 0.037286 99.8 0.02032 99.8 0.03984 99.6 0.020076 100.4 0.04023 100.4 0.02016 99-7 0-03976 99.4 0.03898 99.9 0-03033 100.3 0-02914 100-1 0,02261 99-6 0.02269 99.9 0-03879 99.8 0.03300 99.8 0-03911 97.8 0.05219 99.7 0.02493 99.5 049150 45.5 M methylene-blue ; I potassium indigotintetrasulphonate ; Th thio-indigodisulphonic acid ; F potassium ferricyanide ; R rosinduline ; K-b kitone-blue 2860 KNECHT AND HIBBERT THE BEHAVIOUR OF GLUCOSE ETC.Various amounts of glucose with potassium indigotintetrasulph-onate as oxidising agent. Glucose taken C.C. of TiCl, k-). d. 0.052675 50.1 0.04214 40.2 0.02 107 20.0 0.010535 10.1 04052675 5.0 1 C.C. TiCl = Glucose found yo of glucose k.1- taken. 0.052505 99.7 0.042 1 29 99.9 0.02096 99.5 0.0 10585 100.4 0.00534 99.6 0.001956 g . Fe. Various amounts of glucose with potassium ferricyanide as oxidising agent. 0*0900 65.3 0-0798 99.8 0.0500 41.0 0.0500 100-05 0*0400 32.6 0.0399 99.7 0.0198 16.2 0.0198 100.0 0.0099 5.1 0.0099 100.0 1 C.C. TiCl = 0.002272 g.Fe. Sir mniaq. 1. It is shown that in an alkaline medium glucose lmdose, and certain other carbohydrates are oxidised to a definite degree by methylene-blue and by potassium indigotintetrasulphonate. Glucose glucosamine and galactose take up under the conditions, exactly three and lzvulose four atomic proportions of oxygen. Potassium ferricyanide can replace the dyestuffs in these estim-ations. 3. When boiled for two minutes with excess of caustic soda in an atmosphere of nitrogen both glucose and laevulose neutralise an amount of alkali corresponding exactly to the formation of two molecules of lactic acid. 3. When glucose is boiled with excess of potassium ferricyanide and caustic potash the amount of alkali neutralised (over and above that required to form potassium ferrocyanide) represents exactly three equivalents of caustic potash. In the case of Itemlose, the amount of caustic potash neutralised is approximately four equivalents. 4. Glycuronic acid can be estimated quantitatively by titration with potassium indigotintetrasulphonate. The volumetric osazone titration method may also be employed for the purpose. 5. In presence of caustic alkali glucose is also oxidised quanti-tatively by kitone-blue and by rosinduline. In the former case, two and in the latter one atomic proportion of oxygen is supplied by these dyestuffs. MUNIC~AL COLLEGE OF TECHNOLOGY, UNIVERSITY OF M~NCHESTER. [Received J d y 23rd 1925.

ISSN:0368-1645    

DOI:10.1039/CT9252702854

出版商:RSC    

年代:1925

数据来源: RSC

摘要:

CARTER THE SALTING-OUT EFFECT. 2861 CCCXCVI.-The Salting-out E8ect. The Influence of Electrolytes on the 8olubility of Iodine in Water. By JOHN STANLEY CARTER. THE fact that the solubility of non-electrolytes in water is lowered by addition of electrolytes has attracted considerable attention, mid various empirical or semi-empirical equations have been put forward fo express the relation between the change in the solubility and the concentration of the electrolyte. In most of the earlier work the latter rarely exceeded one gram-equivalent per litre and over this range the Solubility lowering is approximately proportional to the concentration of the electrolyte. Experiments in which the electrolyte concentration was varied over a wider range show, however that this proportionality by no means holds.Measurements of the solubility of phenol in sodium sulphate solutions (Dawson J . Soc. Chem. Ind. 1920 39 151~) of the solubility of ethyl ether in sodium chloride solutions (!€home J., 1921 119 262) and of quinol and quinone in various salt solutions (Linderstrom-Lang Cmpt. rend. Trav. Lab. Curlsberg 1924 15 4) show that over a considerable concentration range the influence of electrolytes on the solubility of non-electrolytes can be expressed by the exponential formula 8 = $&kc first suggested by Seischenov (2. physiM. Ghem. 1889,4 117) where so and s denote respectively the solubilities of the non-electrolyte in m-ater and in a salt solution, the Concentration of which expressed in terms of the volume of the solution is represented by c and k is a constant characteristic of the dissolved electrolyte.In the present investigation the validity of the exponential formula has been tested by measurements of the solubility of iodine in aqueous solutions the concentration of which ranged up to saturation at the temperature of observation. In previous measurements (Jakowkin 2. physikd. Cherra. 1896,20,19 ; Dawson, ibid. 1906 56 606) in which the salting-out effect was examined by determining the partition of iodine between carbon disulphide and aqueous salt solutions the concentration of the electrolytes wa8 varied only between comparatively narrow limits. The num-bers thus obtained are consequently of little use for the present purpose. The experimental data now recorded have reference to the effects produced by sodium nitrate sulphate and dihydrogen phosphate and by nitric and sulphuric acids.These are all readily soluble in water and do not react with iodine. The salts of weak acida ar? generally inadmissible on account of hydrolysis and the haloge 2862 CARTER THE SA.L!KNG-OUT EFFECT. !FHE INFLUENCE OF salts are unsuitable because of the readiness with which they form perhalogen compounds. Various other salts had to be rejected on account of their oxidising properties which introduce difliculties in the estimation of the iodine by means of thiosulphate. Measure-ments were also made with sodium chloride allowance for the effect of perhalide formation being made in a manner which will be described later. E X P E R I M E N T AL.Iodine and the salts employed were p d e d by the customary methods. Since Bray ( J . Amer. Chem. Soc. 1910 32 932) has shown that the hydrolysis of iodine is repressed in faintly acid solution the values of p H for the solutions of sodium nitrate, sulphate and chloride were adjusted to 5-5-5 by the addition of small quantities of the corresponding acids. In the determinations of solubility every precaution was taken to ensure complete saturation efficient atration and to prevent loss of iodine by volatilisation. In the experiments with nitric and sulphuric acids the dissolved iodine was extracted from a known weight of the saturated solution by repeated shaking with carbon tetrachloride; the latter was then freed from mineral acid by shaking with dilute aqueous sodium acetate.Freshly prepared O.OlN-sodium thiosulphate was used for the titrations. In order to obtain solutiom of sufliciently high salt concentration, the first experiments (with sodium sulphate) were made a t 35", but subsequent measurements were made a t 25" and 10". Determinations of the solubility of iodine in water with a slight acid reaction (px 5 ) gave the following results (millimoles per litre) : At 10" .................. 0-835 0-837 At 25" .................. 1.319 1.321 1.322 At 35' .................. 1.808 1.809 The values at 25" agree with the value 1-320 recorded by Bray (Zoc. cit.) but are lower than that 1.334 given by Hartley and Campbell (J. 1908 93 741) which is considered by Bray to be a little too high. In Table I the solubilities experimentally determined are com-pared with those calculated from the exponential equation which wa8 put into logarithmic form, the value of the constant k' being given by the slope of the line in the graph of log s-concentration.In the application of this formula to the results obtained for the more concentrated solutions the mode of expressing the concen ELECTROLYTES ON THE SOLWIUTY OF IODINE IN WATER. 2863 tration of the dissolved electrolyh is a matter of great hp%nm. The range of validity of the logarithmic formula is very much greater when the concentration of the electrolyte is referred fo a hed weight of water than it is when the concentration is expressed in terms of a ked volume of solution. In the latter case deviations appear when the concentration of the electrolyte reaches 2-3 gram-equivalents per litre and the magnitude of these deviations increases continuously with the concentration of the dissolved mlt.For this reason the concentrations of the electrolytes c represent the number of moles (equivalents in the case of sodium sulphate) per lo00 moles of water. The solubility of iodine is expressed in the same way 8' being the measured value and 8 the value cal-culated from equation (1). The values of so at 25" and 35" respect-ively are 0.0238 and 0.0327. The ordinary weight percenttbge composition of the various solutions is given under '' w%." TABLE I. Sodium Nitrate at 25". k' = 040296. W yo. 8'X 10s. 8 X loa. W yo. 8'X 10s. 8 x lo3. W %. 8'X 10'. 8 X l(r 5-89 11-10 15-31 1-98 4-48 6.85 5.66 8.64 1-05 2-31 4.93 7.53 14.40 20-64 22.2 21.8 25.-67 14-5 14-5 39-83 9-11 20.2 19.9 29.75 12-5 12.9 44-16 7-53 18-5 18.4 35.43 10.8 10.8 Sodium Nitrate at 35".k' = 040243. 32.7 32.3 15.89 26.8 26.2 50.1 10.0 32.0 31.4 29.06 20.1 20.1 30-2 30.0 40.17 14-3 14-7 Sodium Sulphate at 25". k' = 0.0090. 17.3 17.4 16-24 8.55 8.60 19.66 6.56 14.4 14-5 Sodium Sulphate at 35". k' = 04087. 31.8 31.1 9.13 20-2 19.8 25-31 5-78 29.3 29.0 14.79 13.3 13.5 31.80 3.05 25.2 25.2 21-10 8-39 8-42 33-1 2-58 Sodium Dihydrogen Phosphate at 25". k' = 0.0062. 19.4 20.0 26-40 11-1 11-1 45.26 11-81 16.1 16.6 31-66 8.93 8-85 13.5 13.7 41-06 [4-3] 5.35 9.16 7.60 9.94 6-59 5-85 3.06 2.66 3.98 Comparison of the observed and the calculated solubilities of iodine shows that there is close agreement over the entire range of possible salt solutions in the case of sodium nitrate and sulphate; and also in the case of sodium dihydrogen phosphate excepting the two most concentrated solutions for which the measured solubilities are smaller than the calculated values.This divergenc 2864 CARTER THE SALTING-OUT EFFECT. THE INFLUENCE OF suggests the operation of some factor other than that responsible for the normal salting-out effect. An explanation may possibly be found in the acid character of the electrolyte in this case for, as will be seen later the behaviour of nitric and sulphuric acids is quite different from that of the corresponding salts. Iodine dissolved in a solution of sodium chloride exists in part as free iodine and partly in the form of a complex polyhalide.The influence of polyhalide formation is such that the solubility of iodine passes through a maximum when the concentration of the sodium chloride increases (Table 11 column 2). FIG. 1. 4.0. 3.5 d Z 3.0. u - 2i 2 3 2.0 *g 2.5 -.r 8 8 1-5 Iu 1.0 0-5 0 10 20 30 40 50 60 70 Percentage of nitric acid. Solubility 6f iodine in aqueous solutions of nit?& acid. In order to obtain the value of k’ in equation (1) for sodium chloride it was assumed that the polyiodide equilibrium NaCl+ I NaClI conforms to the requirements of the mass-law equation and that the variations in the value of the expression [NaCl][IJ maClI,] = a are due to the salting-out effect of the sodium chloride on the iodine.By plotting these values against the sodium chloride concentration and extrapolating to zero salt content a limiting value of a = 12.0 was obtained and this was assumed to be the true value of the mass-law constant. In calculating the concen-trations of free and combined iodine from this number advantage was taken of the fact that the total salt concentration may b ELECTROLYTES ON THE SOLUBILITY OF IODINE IN WATER. 2865 substituted for [NaCl] in the mass-law equation without introducing any serious error. TABLE 11. Sodium Chloride at 25". E' = 040575. Conc. of total yo NaCl. dissolved iodine. 4.52 0.04403 7-43 0.05295 14-14 0-06289 20.31 0.0631 1 23-15 0-06105 25-95 0.05 7 90 Conc. of combined iodine = [NaClI,].0-02405 0-03560 0-05085 0.05475 0-05405 0.05210 Conc. of free iodine =d. 8. 0.0200 0.0197 0.01 74 0.0172 0.0121 0.0122 040835 0.00843 0.00700 0.00698 0.00580 0.00671 The close agreement between the numbers in the last two columns of Table I1 affords clear evidence that the salting-out effect of sodium chloride is exactly similar to that of the other salts examined when due allowance is made for the disturbing influence of plyiodide formation. Tables III and IV contain the data for solutions of nitric and sulphuric acid. As there is no simple relation between the solubility of the iodine and the concentration of the dissolved acid the con-centration of the iodine is expressed in millimola per litre (b). Solubility loo.- % mo,. b. 11-80 0.969 23.88 1.035 32.35 1.020 40.82 0.985 49.57 0-952 55.82 0.962 64-30 1.08 70.00 1.78 TABLE 111. of Iodine in Nitric Acid Solutions. 26". /- yh HNO,. 10.91 23-16 30.05 33-62 42.28 52.77 60-36 63-81 67-41 69.65 72-60 Y b. 1.60 1-79 1-85 1-86 1-85 1-89 2.10 2-36 2.78 3.12 3.94 35". 7-. ?& IINO,. b. 3.31 1.87 6-50 2.04 15.50 2.33 28-72 2.60 42.63 2-75 43-13 2-73 58-34 3.68 TABLE IV. Solubilit,y of Iodine in Sulphuric Acid Solutions at 25". Colour of sat. sol. % Colour of % H,SO,. b. sat. sol. H,S04. b. 11.87 1.02 Brown. 40.42 0.545 Brown with pink tinge. 22-73 0.820 Brown. 59.53 0.380 Pink with brown tinge. 33-25 0.615 Brown. 72-02 0.270 Pink. S7-37 0,200 Pink.The somewhat complicated relations exhibited by the data for solutions of nitric acid are more clearly represented in the diagram 2866 DOBSON : The curve for 35" shows that the solubility increases continuously with the concentration of the acid but the rate of increase diminishes markedly in the middle portion of the curve. The factors which give rise to this diminution have obviously a much greater influence a t lower temperatures in that the curve for 10" shows a maximum and a minimnm. At 25" the relations are intermediate in character ; the form of the curve corresponds with the behaviour expected from the results a t 10" and 35O. In connexion with these results, it may be noted that solutions containing more than 60% m03, although brown when viewed in bulk have a distinctly pink shade.In sulphuric acid solutions the solubility of iodine decreases con-tinuously as the concentration of the acid increases from 0 to 87%, but the data do not conform a t all to the requirements of the logarithmic formula which expresses the salt effects. The colour of the solutions changes gradually from brown to pink. There is apparently no connexion between the effects observed with nitric and sulphuric acid solutions and those obtained for neutral salt solutions. If the mineral acids exert any action which iS comparable with that of the salts it is more or less completely masked by other effects which play no part in the cme of the neutral salts. It seems probable that the solvent power of nitric acid for iodine is in part responsible for the complicated behaviour exhibited by nitric acid solutions. Recent measurements (Manchot Jahrstorper and Zepter 2. u w g . Chem. 1924 141 45) of the solubility of nitrous oxide in aqueous solutions also have shown that the influence of the free mineral acids is totally different from that of the corresponding salts. My thanks are due to Professor H. M. Dawson for suggesting UNIVERSITY OF LEEDS. the lines of this research and for his helpful criticism and advice. [Received October 3&A 1925.

ISSN:0368-1645    

DOI:10.1039/CT9252702861

出版商:RSC    

年代:1925

数据来源: RSC