1756
F R E D CART.4N A N D CHARLES
cannot be interpreted in a quantitative way. Acknowledgment.-The authors thank Mr. J. N.
N.
CAUGHLAN
Vol. 64
Doshi for his assistance with the mass spectrometric analyses.
ELECTRIC MOMENTS OF THE SIMPLE ALKYL ORTHOVANADATES BYFREDCARTAN AND CHARLES N. CAUGHLAN Department of Chemistry, Montana State CoUege, Bozemn, Montana Received June 87, 1060
The electric moments of a series of alkyl orthovanadates have been determined in benzene solution a t 25". The following momenb were found: EtaV04, 1.18 D.;PrpVO4, 1.15 D.; PraiV04,1.23 D.; BupV04, 1.12 0 : ; BusiVOc 1.10 D.; BuNO4, 1.01 D.; BuNO4, 1.16 D.;h a t V 0 4 , 1.11 D. The densities, dielectric constants and refractive indices of the pure esters were also measured. The small electric moments observed show that rotation of the alkoxy groups is considerably restricted.
Introduction
BuatVO4 melts near 45", EtaVO4 near 0" and PrsVO4 and BuaaVOc 5-10' below zero. Vapor pressures were not The alkyl orthovanadates (R3V0,) are an measured but the esters boil from 30 to 140°, depending on interesting group of vanadium compounds with the ester, under approximately 1 mm. pressure. These compounds undergo thermal decomposition, "covalent" properties. I n 1887, Hall' prepared when impure. They are also decomposed by what was apparently ethyl vanadate. In 1913, especially exposure to ultraviolet light. Pure samples will keep for Prandle and Hew2 prepared and described several several months at 0" in the dark, but slightly impure of the alkyl esters. samples show noticeable decomposition. The branched Since the study by Prandl and Hess, only a few chain esters are slightly more stable than the straight esters. They are rapidly hydrolyzed in the presence publications have appeared investigating these chain of small amounts of moisture and gelatinous VZOSis procompounds. 3-.6 duced if water is present in excess. The esters are soluble This study was conducted to measure the elec- in most organic solvents. The purified esters were analyzed for vanadium content tric moments of these vanadium compounds and titrating samples as VO++ with standard KMn04. to help characterize them. It was hoped that in- byThe density, refractive index and dielectric constant were formation on the nature of the bonding and struc- measured on samples of the esters freshly purified by ?t ture of these esters could also be obtained. least a double vacuum distillation. The probable impunties, if present, could be traces of hydrolysis products, Experimental silicone hi-vacuum grease or parent alcohol or related subPreparation.-The ethyl ester was prepared by the re- stances not removed by vacuum distillation. action between VOCla and the sodium ethylate Density.-The densities of the pure esters were obtained using a single-necked capillary pycnometer of approximately VOCla 3NaOEt +Et3VO4 3NaCl (I) 30-ml. capacity. The densities obtained should be accurate The other esters were made using either of the following to *0.0001. reactions Refractive Index.-A Zeiss Abbe' refractometer waa used for measurement of 1 2 2 5 ~ . There was no indication V206 4-6ROH 2RavO1 3Hz0 (IIa) of hydrolysis of the esters during transfer and measurement. The refractive indices thus obtained should be accurate NH4VOs 3ROH 1_ RsVOi NHa 2Hz0 (IIb) to 10.0004. Prandl and Hess* used method IIa. With fused V205 Dielectric Constant.-The apparatus used has been the reaction is slow and using simple distillation to separate previously described.6 The cell used had a capacitance in the water the yields were low. It was improved in both air of approximately 25 ppf. and a cell volume of apprordrespects by using V Z O which ~ had not been fused and an mately 10 ml. Benzene and cyclohexane were used to efficient column to remove the azeotrope formed. The calibrate the cell. The apparatus operated a t 4.6 mc. substitution of NH4VOs for VZOs (IIb) further increased The results should be accurate to k0.008. The observed both speed and yield. dielectric constants were sensitive to traces of moisture and The usual procedure for reaction IIab was as described: this characteristic provided a convenient criteria of their A good grade of alcohol was purified by fractionation. purity. The results of these measurements are shown in The alcohol (in excess) and NH4VOs mixture was refluxed Table I. rapidly under a 90 cm. packed column. The azeotrope was removed aa formed. After several hours, the majority TABLE I of the excess alcohol was distilled. The contents of the PROPERTIES OF THE ALKYL ORTHOVANILDATES stillpot were centrifuged to eliminate unreacted solids and nysD l't decomposition products. The remaining alcohol was reEster Color and state d'jr moved with a vacuum distillation. The esters were puri- EtaV04 Yellow-orange liquid 1.1550 1.5059 3.333 fied by repeated vacuum distillation. Thermal decompo1.0752 1.4953 2.961 sition resulted if distillation was carried out at pressures PranV04 Yellow liquid Pra'VO4 Colorless liquid 1.0324 1.478 3.299 greater than 1mm. With either of the methods, yields of BuatVOc were poor. BupV04 Yellow liquid 1.0335 1.4898 2.780 Other esters were produced in yields of SO-SO%. Better Bu$V04 Pale yellow liquid 1.0189 1.4857 2.761 yields were produced by alcohols with large alkyl groups. Bua'VOr Colorless liquid 1.0083 1.4804 2.969 (1) J. A. Hall, J'. Chem. Soc., 51, 751 (1887). Bu3tV04 Colorless solid .... .... ... (2) W. Prandl and L. Heas, 2. onorg. Chem., 82, 103 (1913). AmstVO4 Colorless liquid 0.9863 1.4805 2.764
+
+
+
+
+
+
(3) M.G.Voronkov and Yu. I. Shorik, Imest. Akad. Nauk S.S.S.R. Otdel. Khim. Nauk, 503 (1958). (4) H. Funk, W. Weiss and M. Zeising, 2.enorg. ollgem. Chem., 296, 36 (lQ58). (6) N . F. Orlov and M. G. Voronkov, I m d . Akad. Nauk S.S.S.R. Oldrl. Khim. Nauk,933 (1959).
Electric Moments.-The total polarizations were obtained using the Hedestrand extrapolation formula.' (6) R. N. Crowe and C. N. Caughlan, J . A m . Chem. Soc., 72, 1694 (1950).
Nov. , 1960
1757
ELECTKIC MOMENTS OF THE S i w L E ALKYL ORTHOVANADATES TABLE I1 RESULTSOF ELECTRIC MOMENT CALCULATIONS
Ester
di
Ad/AN
c1
Ac/AN
I'T!
RD,
CC.
CC.
EtaVO4 0.8735 0.5019 2.2731 2.466 89.86 51.98 Pr.,nVO4 .8736 .4517 2.2748 2,226 102.24 G6.27 PraiVO, 8736 .4074 2.2709 2.427 107.06 66.96 BuanVOi .8736 .4412 2.2729 2.064 114.50 80.06 Bua'VO4 .8736 .4027 2.2703 1.957 114.01 80.55 Bu.PVO~ .8736 .3990 2.2716 1.698 110.36 80.71 BurtVO4 .8735 * 3574 2.2705 2.11 117.63 81.06" .4m,tV04 .8735 .3905 2.2731 1.924 128.35 94.66 This value calculated from tables of atomic refraction and values of EDfor the other esters.
.
0
The electronic polarizntions were considered to be an arbitrary fraction (96%) of the molar refractions. The atomic polarizations were estimated as an additive function of bond atomic polarizations. Values for bond atomic polarizations were obtained from calculated atomic polarizations.* Estimated atomic polarizations of 11-13 cc./ mole were used. Benzene was employed as the solvent. Solution Densities.-Since the Hedestrand equations requires Ad/AN, the magnetically controlled float densit method was used.9 This apparatus permitted rapid: precise measurements of density change under nearly anhydrous conditions. Ad/AN was obtained by least squares methods. Solution Dielectric Constant Measurements.-Ae/AN is also needed for the Hedestrand extrapolation. I n order to obtain this, a jacketed reservoir of approximately 100 ml. was added to the cell of the apparatus used for measuring the dielectric constants of the esters. Small weighed increments of the esters were added to a weighed quantity of benzene in the cell reservoir. The dielectric constant of the solution was measured after each addition. This provided data needed to calculate Aq'AN by least squares methods. Densities and dielectric constants of the solvent used in calculations of total polarization were obtained by extrapolation of €26 and dz54 to zero mole fract,ion. In addition to calculating the solution electric moments at 25", these moments were also obtained from e25, d25, and n 2 5 ~ using the modified Onsager equation.10 The results of these measurements are shown in Table 11. Cryoscopic molecular weight measurements were made in benzene. At the concentrations used in the electric moment measurements, the esters were monomeric. Using the RD values obtained for the esters and known atomic refractions, a VO group refraction of 15.86 =k 0 . 3 cc. was calculated.
Discussion An examination of models of these esters, const ructed with tetrahedral vanadium bond angles, indicates the possibility of restricted rotation in these molecules. A likely configuration (A) and an opposite, less favored, configuration (B) is indicated in Fig. 1. (Only two of the three alkoxy groups are shown.) Bond moments are indicated by labeled arrows, the point indicating the more negative partner. It should be noted that the less probable configuration (B) results in a combination of bond moments m2 and m l in a direction aiding bond moment, m3. The "A" position, however, shows a resultant moment opposing m3. Mutual restriction of rotation among the alkoxy groups should result in measured electric moments lower than the moments these molecules would possess if free (7) G. Hedeatrsnd, 2. phyrik. Chem., 2 8 , 428 (1929). (8) C. P. Smyth, "Dielectric Behavior and Structure," McGrawHill Book Co., New York, N. Y.,1955,p . 420. (9) C. N. Caughlsn and F. Cartsn, J . Am. Chem. Soc., 81, 3840 (1959). (10) See ref. 8.p. 226.
klonsaier.
D. 1.16 1.11 1.40 1.09 1.10 1.30
D. 1.18 1.15 1.23 1.12 1.10 1.10 1.16 1.11
soi in..
..
1.22
0
0
R
R 1-33
1
R
I
R
Fig. 1.-Organic vanadate vector model.
rotation of the alkoxy groups are possible. Estok" has suggested this effect in the orthophosphate esters. The esters can be considered as derivatives of vOc13 produced by replacing C1 atoms with alkoxy groups. VOC13 bond angles closely approximate tetrahedral angles.I2 Since the esters are found to be approximately monomeric a t the concentrations used for electric moment measurements, it seems probably that the vanadium bond angles of the esters are approximately tetrahedral. The oxygen bond angle in alcohols and ethers is near l l O o . l 3 The alkoxy oxygen bond angle in the esters is likely to be close to this value. Since 110' is effectively the tetrahedral angle (109'28'), tetrahedral bond angles may be used to simplify calculation of average free rotation electric moments. A general expression for the calculation of electric moments assuming free rotation has been developed by Eyring.14 Using tetrahedral angles and the notatlion for bond moments shown -in Fig. 1, the simplified equation becomes pz = 3mI2
+ 3mg2 -j-
+ ( m 2f mPm3+ + (2mlm? + mlrn3) - 6 cos? + mIz
ma2 - 6 cos
mlmz) -t 6 cosz
4 is the supplement of the tetrahedral angle (70'32'). -~ \ . -
Lone' -electron pairs in hybrid orbitals (e.g..
NFa, SR) can make sizeable contributions to t,he electric moment of a molecule. Were this the case for these esters such contributions should be included. Since the vanadium electrons appear to be engaged in bonding, we will assume the above equation to be valid. Using ml equal to 1.14 D,16 and since cos 4 is (11) G. K. Eetok and W. W. Wendlandt, J . Am. Chem. Soc.. 77,4767 (1955). (12) J. K. Palmer, ibid., 60, 2360 (1938). (13) See ref. 8,p. 297. (14) H. Eyring, Phys. Rev., 39, 746 (1932). (15) See ref. 8,p. 301.
NOTEB
VOl. 64
results of the solution measurements (Table 11) indicate moments near 1.2 D. Considerable reFrom this equation, the minimum average elec- striction to rotation of the alkoxy groups is, tric moment, possible, assuming free rotation and therefore, indicated. Acknowledgment.-We are indebted to the Office irrespective of the values of m2 and m3 is 1.86 Debyes. This is evident since a minimum calcu- of Ordnance Research, U. S. Army, for supporting lated p will occur n-hen mz - m3 = +0.38. The this work under Contract D-4-04-200-ORD-037. I/$,
this equation simplifies to ji' = (WL,
- ma)' - 0 . 7 6 ( ~ ~-2 m ~ + ) 3.611
NOTES THE DECL4RBOXYLATIOK OF THE TRICHLOROACETATE ION I N n-BUTYL ALCOHOL, %-HEXYL ALCOHOL AND n-CAPROIC ACID BY LOUISWATTSCLARK* Department of Chemistry, Saint M a r y of the Plains College, Dodge City, Kanaas Receined April 8, 1960
A large variety of organic reactions takes place as the result of an initial interaction between an electrophilic, polarized, carbonyl carbon atom (in aldehydes, ketones, esters, acids, amides and related compounds) and a nucleophilic agent. The extremely general Claisen condensation may be cited as an outstanding example of this princip1e.l It appears highly probable, also, that a great many reactions of biochemical interest, e.g., enzymic and immunization reactions, are of this same general type.2 Two closely related examples of this type of reaction which have been t'he subject of a great many kinet'ic studies over a relatively long period of time are the decarboxylations of malonic acid and the trichloroacetate ion in polar solvents. It has been demonstrated convincingly that the rate-determining st,ep for both reactions is the formation, prior to cleavage, of a transition complex involving coordinat>ion between the elect.rophilic carbonyl carbon atom of un-ionized malonic acid on the one hand, and that of the trichloroacetate ion on Dhe other, with an unshared pair of electrons on the nucleophilic atom of the solvent m ~ l e c u l e . ~The decomposition of malonic acid has been studied, to date, in approximately 53 non-aqueous s o l v e n t ~ ,that ~ of the trichloroacetat'e ion in more than a d ~ i i e n . ~ In those solvents containing mobile electrons in addition to the fixed electron pairs on the nucleophilic atom alii interest'ing difference in electron influence has been observed between these two species. It has been observed that, in the att'empted coordination between malonic acid and a polar molecule, the effective positive charge on the * Western Carolina College, Cullowhee. N . C. (1) R. Q . Brew-ster, "Organic Chemistry," 2nd Ed., Prentice-Hall. Inc., New York, N. I-., 1953,pp. 400 et seq. (2) J. H. Turnhull, Esperimtia, XV/8, 304 (1959). (3) (a) 0.Fraenkol, R. I,. Belford and P. E. Yankwich, J . A m . Chern. Soc., 76, 15 (1854); (h) L. W. Clark, Txrs J O U R K ~ L 68, , 99 (1959). (4) L. F.Clark, ibid., 64, 692 (1960). (5) L. W. Clark, ibid., 64, 917 (1960).
carbonyl carbon atom of the malonic acid tends to attract mobile electrons present in the solvent molecule, resulting in a decrease in AH*. On t h e other hand, the resonating negative ionic charge on the trichloroacetate ion tends to repel mobile electrons present in the solvent molecule, thus decreasing the electron density on the nucleophilic center a t the moment of reaction and resulting in an increase in AH*. The solvent molecule responds to the advances of the malonic acid, but resists the advances of the trichloroacetate ion. This phenomenon is not observed in polar solvents containing no mobile electrons.3b Although this principle has been tested already on a considerable number of compounds of various classes it was deemed worthwhile to carry out additional tests to see whether or not any exception could be found to it. For this purpose kinetic studies have been carried out in this Laboratory on the decarboxylation of the trichloroacetate ion in three additional polar solvents, namely, nbutyl alcohol, n-hexyl alcohol and n-caproic acid. The results of this investigation are reported herein. Experimental Reagents.-( 1) The potassium trichloroacetate used in this research was reagent grade, 100.0% assay. (2) The solvents were reagent grade chemicals which were freshly distilled a t atmospheric pressure directly into the reaction flask immediately before the beginning of each decarboxylation experiment. Apparatus and Technique.-The details of the apparatus and technique have been described previously.6 Temperatures were controlled to within fO.OlO, and were determined by means of a thermometer calibrated by the U. S. Bureau of Standards. The reaction flask was 100ml. capacity. Approximately 60 g . of solvent was used in each experiment, along with 330 mg. of potassium trichloroacetate (the amount required to furnish 40.0 ml. of COz a t STP on complete reaction).
Results In the decarboxylation of the trichloroacetate ion in n-butyl alcohol and in n-hexyl alcohol the log ( V , - V,) was a linear function of time over only approximately the first 75% of the reaction. In any one experiment no more than 85% of the theoretical volume of COz was collected. This probably is attributable to a partial absorption of the evolved COz by the alkoxide ions concomitantly formed. I n the decarboxylation of the trichloroacetate ion in n-caproic acid, on the other hand, (6)
L. W. Clark, zbzd., 60, 1150 (1956).
