Feb., 1963
S O L U B I L I T I E S OF
TYPICAL COMPOUNDS
28 1
IN N - h t E T I I Y L A C E T . 4 M I D E
SOLVENTS HAVING HIGH DIELECTRIC COISSTAKTS. XIV. SOLUBILITIES AND REACTIOM OF SEVERAL TYPICAL INORGAXIC AND ORGANIC COMPOUNDS IN K-METHYLACETARUDE' B Y L. R. DAWSON, J. E. BERGER, J. W. VAUGHN, ASD H. C.
ECKSTROM
Department of Chemistry, University of Kentucky, Lexington, Kentucky Received July 10, 1961 Studies of several oxidation-reduction reactions and mathematical reactions of inorganic compounds in N-methylacetamide (NMA) a t 40" have been made. In general, behavior paralleling that in water has been observed. I n some cases the reactions are slower in NMA. New compounds which have been prepared include AlC13.6NMA, Al13ra~10NMA,and NMA.H&O,. Solubilities of numerous inorganic and organic compounds in NM-4 have been determined. The data provide evidence that XMA is a more "universal" solvent than is water.
Xumerous previous papers from this Laboratory by Dawson, Sears, et al. , have described the electrochemical properties of solutions of electrolytes in N-methylacetamide (SillA). As a solvent K M h possesses unusual properties. Its very high dielectric constant (16.5 a t 40') strongly enhances its dissolving power for inorganic solutes and its organic constitution makes it a good solvent for organic compounds. Other desirable properties include a broad liquid range (from 29.8 to 206') and stability even a t the boiling point. The purpose of this study was to obtain solubility data for a series of types of compounds in NMA and to study several typical reactions in this solvent.
Experimental Conductance Measurements.-Adequate descriptions of the apparatus and procedures used for making conductance measurementa have been given previously.*-' Solvent.-The preparation and purification of NMA has been described in earlier papcrs.2-4 The specific conductance of the NMA used in this investigation ranged from 0.6 X to 2 X ohm-' cm.-1. Salts.-Reagent grade salts were used in the reaction studies. These were dried in a vacuum desiccator over anhydrous magnesium perchlorate and used without further treatment. Determination of Solubilities.-Approximate solubilities were determined by use of a large test tube fitted with a stopper and covered with aluminum foil into which was placed 10 ml. of solvent. From 0.5 to 1.0 g. of the solute being studied waa added in small incrementa with constant shaking. When the solution appeared t o be saturatcd, the tcst tubes were heated to 60" in a water bath with frequent shaking. If solute precipitated upon cooling to 40°, the solution was assumed to be saturated. Solubilities determined in this manner probably are within 5% of the correct value. To determine quantitative solubilities, solutions were prepared containing 25 ml. of solvent and excess solute in g1- stoppered flasks. The flasks were placed in an oven a t 40' for one week and shaken frequently during this time. They were transferred thvn to a thermostated bath and held for an additional week with frequent shaking to ensure saturation. To determine the concentration of solute a small amount of the solution was withdrawn and transferred to a tared conductance eel1 containing NMA. The cell with its contents was weighed again and plwed in a thermostated bath where it was held for from 12 to 24 hr. with frequent shaking. The conductance of this solution then was determined in the usual way. All samples were run in duplicate. Conductance data for all of the salts used in this study are available in this Laboratory. Most of these data have been published.4-5 From thcse data the conccntrations of the saturated solutions were calculated. (1) This work was supported in part by the U. S. Atomic Energy Commission under Contract No. AT-(40-1-)-2451. (2) L. R. Dawson, P. G. Sears, and R. 11. Graves, J. Am. Chem. Soc., 71, 1986 (1955).
(3) L. R. Dawson, E. D . Wilhoit, and P. G. Sears, ibid., 78, 1569 (1956). (4) L. R. Dawson, E. 13. Willioit, Iz. R. Holmes. and P. G. Sears, ibid., 79, 3004 (1957). ( 6 ) I.. It. Ihwson, el al., ibid., 80, 4233 (1958), and others.
Reactions in NMA. Ferric-Iodide System.-A 0.1-g. sample of ferric chloride dissolved exothermally in 11 ml. of NMA and 0.1 g. of anhydrous potassium iodide was dissolved in a second 11-ml. portion of the solvent. Upon mixing the solutions, a reartion proceeded readily. Aftcr shaking the mixture for sevcml minutcs, excess aqueous silver nitratc solution was added to stop the reaction. Free iodine was extractcd with chloroform. Its presence indicated the occurrence of the reaction 2FeClS 2KI 2KC1+ 2FeClz 1 2 . Cupric-Iodide Reaction.-A 0.15-g. sample of anhydrous cupric sulfate was placed in 10 ml. of NMA; not all of i t dissolved. Then 0.25 g. of anhydrous potassium iodide waa dissolved in a second 10-ml. portion of solvent. The two solutions were mixed and Inaintained a t 35". A slow reaction was followed visually by noting the appearance of the characteristic color of iodine. After 7 min. the temperature was raised to 80" and maintained for 20 min. The resulting dark brown solution then was quenchcd and free iodine was extracted with chloroform, 4KI + 2CuI 2x2giving evidence of the reaction 2CuS01 so4 11. Sulfide-Iodine Reaction.-Anhydrous hydrogen sulfide 'sas bubbled into 10 g. of NMA a t 40' for several minutes. Then 2.4 g. of iodine was dissolved in 11 g. of NMA. Twenty drops of the resulting solution wcighed 0.5 g. and contained 0.09 g. of iodine. The hydrogen sulfide solution was titrated by adding the iodine solution dropwise. The iodine color waa discharged rapidly until 20 drops of oxidant had been added. Sulfur prccipitated from the mixture. Twenty drops of iodine was added in excess and hydrogen sulfide bubbled through it until the "end-point" was reached again. This caused the dissolution of all of thc free sulfur. The process was repeated several t m e s by addition of successive amounts of oxidant and reductant. Apparently the S. rcaction was HzS IZ 2HI Iodine-Sulfite Reaction.-A 0.1-g. sample of anhydrous sodium sulfite was added to 10 g. of NMA; not all of i t dissolved. Upon adding an iodine solution in N M A the iodine color was discharged gradually, indicating a rather slow reaction. At 80' the reaction wm more rapid. No attempt waa made t o identify the products 1 2 + 2NaI but it is probable the reaction was 2NazSOa NazSzOs. Dichromate-Iodide Reaction.-A 0.3-g. sample of anhydrous potassium iodide was added to a solution of 0.05 g. of potassium dichromate in NMA. There was no visible change. The addition of 0.36 g. of NMA.HZSO4 immediately caused the appearance of free iodine. After several minutes there was considerable free iodine extracted with chloroform. Iodide-Iodate Reaction.-A solution was prepared consisting of 0.1 g. of anhydrous potussium iodate, 0.42 g. of potassium iodide, and 12 g. of S M A . Addition of 0.55 g. of (NMA)z.HC1 produced the gradual liberation of free iodine. Thiosulfate-Iodine System.-An iodine-NMA solution wa8 decolorized rapidly upon being addcd dropwise to a solution containing 0.1 g. of sodium thiosulfatc pentahydratc in 10 g. of NMA. Ferrous-Dichromate Reaction.-A saturated solution of ferrous sulfate hcptahydrate in NMA reacted rapidly with potassium dichromate with or without the prescnee of anhydrous sulfuric acid. Carbonate-Acid Reaction.-A 0.1-g. sample of calcium carbonate was addcd to 10 g. of NMA. Hydrogen chloride was supplied by adding 0.36 g. of (NMA)*.HCl. Gas was evolved only when the mixture waa heated. The speed of the reaction wu9 limited
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282
L. R. DAWSON, J. E. BERGER, J. W. VAUGHX,AND H. 0. ECKSTROM
apparently by the low solubility of the carbonate in NMA. A similar system was treated with YMA.HZS04. Carbon dioxide was evolved slowly a t room temperature and rapidly a t elevated temperatures. Bicarbonate-Acid Reaction.-When sodium bicarbonate in NMA was treated with SMA.HZSO4, carbon dioxide was evolved rapidly. NMA-Hydrogen Chloride Complex.-The preparation of a white crystalline material having the formula (NMA)2 HC1, reported earlier by D'Alelio and Reid,6 was confirmed. The crude product was recrystallized easily from benzene and dried over magnesium perchlorate under reduced pressure. The following analytical dFta were obtained: equivalent weight, 183.9, 183.7 (calcd. for (NMA)Z.HCl, 182.7); % chlorine 19.28, 19.30 (calcd. for (NMA)z.HCl, 19.41); yo nitrogen by Kjeldahl, 15.17, 15.27 (calcd. for (NMA)Z.HCl, 15.34). NMA-Sulfuric Acid Complex.-Concentrated sulfuric acid was added dropwise t o NMA a t room temperature. An exothermic reaction occurred with a concomitant marked increased in viscosity. When the system became quite sirupy, the mixture w~t8 cooled and the entire mass crystallized. Attempts to recrystallize the compound from benzene and ether were not entirely successful. In boiling benzene the crystals melted to a colorless liquid phase having a density greater than that Qf benzene. After digesting a t the boiling point, the mixture was cooled and the more dense phase crystallized. A slurry was prepared using a second portion of benzene and the solid was filtered from the system in a dry atmosphere. Analytical data: equivalent weight, 85.3, 85.3, 85.4 (calcd. for NMA.HzSO4, 85.6); yosulfate 56.32, 56.30, 56.24 (calcd. for NMA*H2504, 56.12); yo nitrogen by Kjeldahl 8.07, 8.27 (calcd. for NMA.H2S04, 8.18). The compound is very deliquescent. Aluminum Chloride-Hexa-N-methy1acetamide.-A 6.78-g. sample of anhydrous aluminum chloride was added slowly t o 96 g. of ?;MA. A vigorous exothermic reaction occurred. A solid became caked a t the bottom but after heating 2 hr. on a water bath, the solid phase consisted entirely of flocculent white material. This solid was filtered in a dry atmosphere, washed with acetone, and dried in a vacuum desiccator over magnesium perchlorate. Analytical data: equivalent weight, 193.3, 195 (calcd. for AlCla.GNMA, 190.6); % chlorine 18.34, 18.26 (calcd. for AlC&.GNMA, 18.60); % nitrogen by Kjeldahl, 14.48, 14.35 (calcd. for AlCla.GNMA, 14.70). Aluminum Bromide-Deca-N-methylacetamide.-Aluminum bromide is quite soluble in NMA; concentrations as high as 0.7 molar were obtained a t room temperature. The more concentrated solutions are very viscous. Despite the use of various added solvents, it was not possible t o induce crystallization of the solid phase. A quantity of the solution was shaken with a 15-fold excess of acetone. Upon cooling this system to -80' a small quantity of heavy oil separated. The acetone was decanted and the residue heated a t 50' under a current of clean dry air. The remaining liquid had an equivalent weight of 326 (calcd. for A1Br3.lONMA, 332.6) and contained 24.5y0 bromine (calcd. for AIBra. lONMA, 24.03). Boron Trifluoride-NMA System.-Boron trifluoride was prepared7 by reaction of sulfuric acid, boric oxide, and ammonium fluoroborate. The gas was passed through an appropriate purifying train and bubbled slowly into NMA. A mildly exothermic reaction occurred and the solution became viscous; however, all attempts to isolate a definite complex were unsuccessful. This behavior parallels that reported by Muetterties and Rochow* for boron trifluoride in other amides. In general, the courses of the various reactions in NMA seem to parallel those in water. In numerous cases the reactions are slower in S M A . This may be attributable in part to its greater viscosity.
Results Solubilities of Inorganic and Organic Compounds in NMA.-Table I contains approximate solubility data for several inorganic compounds in NMA at 40". These data are believed to be within 5% of the accurate (6) G. D'Alelio and E. Reid,
J. Am. Chem. Soc., 69, 109 (1937). (7) €1. F. Walton, "Inoiganic Preparations," Prentice-Hall, Ino., NeYork, N. Y . , 1948, p. 99. (8) E. L. Muetterties and E. G. Rochow, J. Am. Chem. Soc., 76, 490 (1953).
Vol. 67
valnes. Quantitative solubility data for other compounds are given iii Table 11. TABLE I APPROXIMATE SOLUBILITIES IS MOLESPER LITER OF SEVERAL INORGANIC COMPOUNDS I N N-METHYLACETAMIDE AT 40" Compound
Solubility
&( NO3)3.9HzO
1 12
BaBrz 2H20 Ba(NOd2 Ba12.2H20 CaC12 CaBrp Ca( N03)2.4H20 Cr( S03)3 8H20 MgCL 6H20 Mg(N0s)z 6H20 YiClt 6Ha0 ICBi-03
0.75 0 36 0 40 0 27 0 40 2 46 1 19 0 10 0 23 0 21 0 03
Compound
Solubility
K2S04 KPi NaBr XaI NaSCN XaC104 NaPi SaBr03 Sax01 SrBrz Sr(C104)~ Sr(N03)~
0 03 0 51 1 85 2 86 3 64 5 00 0 48 0 10 0 65 1 64 0 87 0 71
TABLE I1 QUANTITATIVE SOLUBILITIES I N hlOLES PER LITER O F SEVERAL IXORGANIC COMPOUNDS IN S-METHYLACETAMIDE AT 40' Compound
NHdC1 NH4Br NHJ SH4SCN NHaCIO4 NH4N 0 3 KCl
Solubility
0 1 2 4
896 745 470 626 1 349 3 281 0 120
Conipound
KBr KI KSCS KC104 KNOa LiCl NaCl
Solubility
0.429 1 356 2 170 0 413 0 247 0 413 0 339
These data show that for a common anion the order of increasing solubilities with respect to the cation is
Kf
< ?;a+ < Li+ < XH4+
For ammonium salts the order of increasing solubility is C1-
< C104- < Br- < I- < ?SOs- < SCY-
The corresponding solubilities in water are C104-
< C1- < Br- < I- < KO3- < S C S -
It is apparent that the only difference in the two sets of data is a reversal of the chloride-perchlorate solubilities in ?;MA. This same reversal is exhibited by potassium salts for which, in K M h , the order is C1-
< S O a - < C104- < Br- < I- < SCN-
The corresponding solubilities in water are C104-
< NOs- < C1- < Br- < I- < SCN
The following salts were found to be quite insoluble in
NMA: carbonates of sodium, calcium, barium, and magnesium; sulfates of sodium, beryllium, calcium, barium, and magnesium; and phosphates of sodium, calcium, barium, and magnesium. Calcium fluoride is insoluble and the solubility of calcium chloride is low. To be noted especially are the very low solubilities of the carbonates and the limited solubilities of the chlorides. The following organic compounds dissolved to the extent of 30% by weight or more in S M A a t 40': dioxane, cyclohexane, toluene, pyridine, nitrobenzene, phenol, acetone, chloroform, aniline, carbon tetrachloride, naphthalene, ethanol, acetonitrile, anisaldehyde, methyl ethyl ketone, pyrogallol, benzaldehyde, nbutylamine, acetophenone, ethyl acetate, and di-nbutyl phthalate. Normal pentane exhibited the lowest
Feb., 1963
OXIDATION OF GLYCEROL BY CERIUM(IV)IN SULFURIC AND PERCHLORIC ACIDS
283
From these data it appears that N-methylacetamide is a more “universal” solvent than is water.
solubility of the organic compounds tried; it dissolves only to the extent of about 5% by weight at 40’.
MECHANISM AND KINETICS OF THE OXIDATION OF GLYCEROL BY CERIUM(1V) IN SULFURIC AND PERCHLORIC ACIDS BY G. G. GUILBAULT AND W. H. MCCURDY, JR. Princeton University, Princeton, New Jersey, and University of Delaware, Newark, Delaware Received July 12, 1962 The rate of oxidation of glycerol by cerium(1V) sulfate in sulfuric acid was found to follow the rate expression: -d[Ce(IV)]/dt = 4.0 X 10-8 [Ce(IV)][glycerol] a t 20.0” in 2.0 F sulfuric acid, where concentrations are in moles per liter and time is in minutes. The rate also was found to be first order in perchloric acid a t low acid concentrations, and inversely proportional to HSOA-. As in previous oxidations discussed, a prior equilibrium step is believed to occur: Ce(S04)z H + e CeS04+2 HSOd-. I n perchloric acid the oxidation of glycerol by ceric perchlorate proceeds via an intermediate complex between one Ce(1V) ion and one glycerol molecule. The existence of the complex was proved independently by spectrophotometry, and the proposed mechanism fits the kinetic data.
+
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further oxidized to formic acid by excess Ce(IV) a t elevated temThe oxidation of organic substances by cerium(1V) peratures, and (2) by keeping the concentration of Ce(1V) lower reagents is found to follow different mechanisms, dethan glycerol, the possibility of further complexes being formed is pending upon the type of acid media used. In sulfuric eliminated. acid, oxidation of the polyhydric compounds takes All solutions were thermostated a t 20.0 =k 0.1’ before mixing, and during reaction. place via a free radical mechanism.2~3 However, in perchloric acid a preliminary complex between cerium Results and Discussion and the glycol is first formed, followed by dispropor(A) In Sulfuric Acid. Effect of Perchloric Acid and ~ tionation of the complex as the rate-determining ~ t e p . ~ ,Bisulfate on the Rate.-An effect similar to that obIn the present study, elucidation of the kinetics of served in the uncatalyzed mercury(1)-cerium(1V) reacthe glycerol oxidation was attempted, with the aim of tions’ is observed in the glycerol-cerium(1V) reaction determining the manner in which cerium(1V) oxidizes in sulfuric acid. At constant concentrations of ceriumthis substance. The rate of reaction was studied both (IV), glycerol, and bisulfate, an increase in the perin sulfuric and perchloric acids, and an independent chloric acid concentration a t low amounts (0.5-4.0 F ) , check of the mechanism in the latter acid was performed produces a nearly linear increase in the rate. At spectrophotometrically by proving the existence and higher concentrations, precipitation of cerium(1V) structure of the intermediate complex. salts occurs, forming a heterogeneous system, with a Experimental decreasing rate (Table I, B). Also, an increase in the bisulfate concentration from 0.2-2.0 F causes a deReagents and Solutions.-Cerium(1V) sulfate solution, 0.025 F, was prepared by heating 13.2 g. of cerium(1V) sulfate [Cecrease in the rate of oxidation. Again, the prior eyuilib(HS04)a (G. F. Smith ChemicalCo.)] with6.0ml. of 96% sulfuric rium reaction acid. The resulting paste was carefully heated with stirring for 5 min. and gradually diluted to 1 1. with distilled water containing 82.5 ml. of 72y0 perchloric acid. These solutions were standardized against primary standard arsenic(II1) oxide. Ceric Perchlorate Solution.-Soluble, 0.01 F , 100 mesh ceric hydroxide, Ce(OH)r (G. F. Smith Chemical Co.), was dissolved in 115 ml. of 72y0 perchloric acid, and the resulting solution diluted to 1 1. with distilled water. This reagent was stored in a dark reagent bottle and standardized as described above. A 1 M solution of glycerol was prepared by dissolving the proper amounts of C.P. glycerol in 1-1. portions of distilled water. The solution was standardized with periodic acid. Rate Measurements.-The rate of the reaction
Ce(1V)
+ glycerol -+
3CH20
+ 2Ce+3+ 2H+
was measured either by titration of the remaining Ce(1V) or by measuring the decrease in the absorbance of the solution with a Beckman Model DB a t 280 and 400 mp. The titrimetric method finally was abandoned and only the spectrophotometric method was used, since it allowed the use of lower cerium(1V) concentration. This was desirable because (1) it minimizes the error due to formaldehyde which is produced during the reaction and can be (1) W. H. McCurdy, Jr., and G. G. Guilbault, J . Phzls. Chem., 64, 1825 (1960). (2) G. I-Iargreaves and L. H. Sutcliffe, Trans. Faraday Soc., 61, 1105 (1955). (3) G. Mino, S. Kaizernian, and E. Rasmussen, J. A m . Ckem. Sac., 81, 1494 (1959). (4) M. Ardon, J . Chem. Soc., 1811 (1957). (5) F. R. Duke and R. F. Bremer, J . A n . Chem. Soc., 73, 5179 (1951).
Ce(SO&
+
+ H+
CeS04+2 HS04-
is believed to take place, and this CeS04+2 is the reactive species in the glycerol oxidation. The notation Ce(1V) will again be used to denote the total cerium(1V) concentration. Effect of Cerium(1V) and Glycerol.-At a constant glycerol concentration of 0.200 M , a bisulfate concentration of 0.20 M , and 0.20 M perchloric acid, an increase in the cerium(1V) concentration over the range tested, 0.00861-0.06888 M , produced a linear increase in the rate, indicating a first-order effect. Likewise, an increase in glycerol over the range 0.0266-1 .OO M at constant cerium(IV), bisulfate, and acid concentrations of 0.01722, 0.20, and 0.20 M , respectively, caused a linear increase in the rate. The mechanism proposed to account for this observed rate dependence is Ce(1V)
H i H + HC-;-C--CH2 I , / OH
OH
ki
+
1 OH
H-C-H
II
0
+ C--CH2 H + H + + CefS OH I OH 1
