3309
CHARGE DENSITY ON THE PHOSPHORYL OXYGEN pends on The use of a smaller 6, if reasonable a t all, would result in that the diffusion limit is smaller than the rate observed. In this way, very important information about the mechanism of the reaction, namely the distance of the reactants in the activated complex, can be tested.
Acknowledgment. One of the authors (V. J. IF.) is very grateful to the Hungarian Academy of Sciences for the Graduate Fellowship awarded to him. The authors are very indebted to Dr. B. Rozsondai, who called their attention to the photographic method of the measurement of the diffusion coefficient.
The Charge Density on the Phosphoryl Oxygen in a Series of Phosphate Esters;
Tributyl Phosphate, a Monocyclic Phosphate, and a
Bicyclic Phosphate Ester1 by A. L. Mixon and W. R. Gilkerson* Department of Chemistry, University of South Carolina, Columbia, South Carolina 202'08
(Received April 1 I 1071)
Publication costs borne completely by The Journal of Physical Chemistry
The effects of adding a series of phosphate esters on the conductances of dilute solutions of piperidinium, N-methylpiperidinium, and N-ethylpiperidinium picrates in chlorobenzene at 25" have been measured. Cation-ligand association constants, KL, have been calculated for the N-methyl- and N-ethylpiperidinium cations with tributyl phosphate, n-octyl trimethylene phosphate, and l-oxo-4-ethyl-2,6,7-trioxa-l-phosphabicyclo [2.2.2]octane. The ratio of the cation-ligand association constant for the N-ethyl cation, &(Et) , to that for the N-methyl derivative, KL(Me),has been used as a probe of the electron density on the phosphoryl oxygen atoms in the esters, These results are compared with phosphoryl oxygen stretching frequencies, YPO, and with calculated (Huckel-MO) values of charge densities.
Introduction Bicyclic phosphate esters have been found2 to be poor extractants, compared to acyclic esters, for lanthanide metal ions. The same report contained the observation that these bicyclic esters had abnormally high phosphoryl oxygen stretching frequencies. Burger3 reported a linear relation between extractant ability and the phosphoryl oxygen stretching frequency, V P O : the lower the frequency, the better the extractant. Wagner4 carried out LCAO-MO calculations for a number of acyclic phosphoryl compounds and found an excellent correlation between the calculated phosphoryl PO a-bond orders and experimental values of phosphoryl oxygen stretching frequencies. The calculated values of the net charges on the phosphoryl oxygen also show excellent correlation with V P O . Presumably the extractant ability of a phosphoryl derivative is greater the greater the negative charge on the phosphoryl oxygen, and the larger the n-bond order of the phosphoryl bond, the smaller the net negative charge on oxygen and the larger the phosphoryl oxygen stretching frequency.
Recent Hucke 1-MO calculations by Collin6 indicate that the net negative charge on the phosphoryl oxygen decreases in the order ethylene phosphate anion > trimethylene phosphate anion > diethyl phosphate anion. Calculated values of the positive charge on phosphorus decrease more markedly in the same order. Boyde reported extended Huckel-XI0 calculations of both net charges on atoms and bond orders. Boyd's calculations indicate that net negative charge on the phosphoryl oxygen decreases in the order trimethyl phosphate > methyl ethylene phosphate. This is the reverse of the order given by Collin for the diester anions. The theoretical studies of both Collin and Boyd were concerned principally with the influence of the geo(1) This work has been subported in part by Grant GP 6949 from the National Science Foundation. (2) 8. G. Goodman and J. G. Verkade, Inorg. Chem., 5, 491 (1966). (3) (a) L. L. Burger, J . Phys. Chem., 62, 590 (1958); (b) J. L. Burdett and L. L. Burger, Can. J. Chem., 44, 111 (1966). (4) E. L. Wagner, J. Amer. Chem. Soc., 85, 161 (1963). ( 5 ) R. L. Collin, ibid., 88, 3281 (1966). (6) D. B. Boyd, ibid., 91, 1200 (1969).
The Journal of Physical Chemistry, Vol. 76,N o . 81, 1071
A. L. MIXONAND W. R. GILHERSON
3310 metrical constraints imposed by the presence of -(CH2).bridges between two of the alkoxy1 oxygens on the stability of the esters. The two studies are in agreement regarding the effect of the presence of rings on the positive charge on phosphorus but not on the effect on the negative charge on the phosphoryl oxygen. We believe that we have an experimental tool wherewith we can determine which in a series of phosphate esters has the greater negative charge on the phosphoryl oxygen atom and which has the least. We have found7J that the cation-ligand association constants, K L (eq l), for a series of ligands with piperidinium BH+
+L
BH+,L
KL = [BH+,Ll/[BH+l[Ll (1) (PipH+), 1-methylpiperidinium (hlePipH+),and l-ethylpiperidinium (EtPipH+) cations fell into two distinct classes: the first group, consisting of the free piperidines themselves, 2,B-dimethylpyridine (lut), and triphenylphosphine all had K L values which decreased in the expected order PipH+ > MepipH+ > EtPipH+, while the second group, consisting of triphenylphosphine oxide (Ph3PO) and tetrahydrofuran, had KL values which decreased in the order PipH+ > EtPipH+ > RlePipH+. We attributed the reversal of the order for ethyl- and methylpiperidinium cations with the second group of ligands to the presence of more than one lone pair of electrons on the oxygen atom. With one pair of electrons involved in the interaction with the =NH+ group, the other pair are available for interaction with a terminal CH proton on the end of the ethyl substituent on the cation, (1).
1
We propose that the ratio KL(Et)/K~(l\Ie), where KL(R) represents the K L value for the N-alkylpiperidinium cation, may be used as a probe for relative electron density on a coordinating atom in a series of ligands of not too different structure. We report here application of these ideas to the problem of the electron density on the phosphoryl oxygen in a series of trialkyl phosphate esters, tributyl phosphate (2)) n-octyl trimethylene phosphate (3), and 1-oxo-4-ethyl2,6,7-trioxa-l-phosphabicyclo [2.2.2]octane (4). These
2
3
The Journal of Physical Chemistry, Vol. 76, N o . $1, 1971
4
were chosen because of their stability, ease of preparation, and availability of other physical properties such as dipole moments and infrared spectra. The dipole moment of compound 3 has been determined in the course of this work.
Experimental Section Chlorobenzene (PhC1) was purified as before.9 Benzene (Baker and Adamson, reagent grade) was passed through alumina (Alcoa, grade F-20), kept for 1 day over sodium ribbon and then distilled from the sodium ribbon on a 38-cm column packed with glass helices. A middle cut was taken. Piperidinium picrate (PipHPi), N-methylpiperidinium picrate (MePipHPi), and N-ethylpiperidinium picrate (EtPipHPi) were prepared as described p r e v i ~ u s l y . ~Tri-n-butyl phosphate, 2 (Eastman Organic Chemicals, White Label), was distilled under vacuum using a Hickman molecular still; the bath temperature was maintained at 75'. n-Octyl trimethylene phosphate, 3 (RICP), was supplied by Dr. T. H. Siddal, 111, of the Savannah River Laboratory, E. I. duPont de Nemours and Co. This phosphate was purified by distillation at 148" ( 2 mm). A middle fraction was taken. The density of 3 was found to be 1.072 g/ml at 25', and the refractive index at 25" (Na D line) is 1.4480. l-Oxo-4-ethyl-2,6,7-trioxa-1-phosphabicyclo [2.2.2]octane, 4, the bicyclophosphate (BCP), was also supplied by Dr. T. H. Siddal, 111. Before each use it was recrystallized once from ethanol, mp 209-210' (lit.lo mp 207'). Precision conductance measurements were carried out a t 25.00' using bridge, constant temperature oil bath, and Kraus erlenmeyer conductance cells previously described. A Balsbaugh Laboratories Type 100T3 cell with nickel electrodes was used for measurements of dielectric capacitance. The air capacity of this cell is 104.9 pF. The capacitance measurements were made using a General Radio Type 716-C capacitance bridge and 716-P-4 guard circuit. The oscillator, operated a t 100 kHz, was a General Radio Type 1330-A, and the det,ector was a General Radio Type 1231-B amplifier and null detector. Density measurements were made using a Lipkin pycnometer. Refractive indices were measured on a Bausch and Lomb Abbe-SL refractometer. The dielectric constant of chlorobenzene is 5.621.l1 The viscosity of chlorobenzene is 0.752 cP12 and the density is 1,011 ~ / C C . ~All the foregoing physical constants are for 25.0'. All concentrations are expressed in units of moles per liter. (7) A. L. Mixon and W. R. Gilkerson, J . Amer. Chem. SOC.,89, 6410 (1967). (8) W. R. Gilkerson and A . L. Mixon, ibid., 89, 6415 (1967). (9) E. R. Ralph, 111, and W. R. Gilkerson, ibid., 86, 4783 (1964). (10) 0 . Neunhoeffer and W. Maiwold, Chem. Be?., 95, 108 (1962). (11) A. A. M a r y o t t and E. A . Smith, Nat. Bur. Stand. ( V . Sa),Circ., 514, 1 (1951). (12) R. L. McIntosh, D. J. Mead, and R. M. Fuoss, J . Amer. Chem. SOC., 62, 506 (1940).
3311
CHARGE DENSITY ON THE PHOSPHORYL OXYGEN
9
io3 [&I, M Figure 1. Conductance ratios, R, for PipHPi in chlorobenzene a t 25" with PhBPO, TBP, MCP, and BCP as addends.
Results We define a quantity, R = (g/g0)2, where g is the conductance of the salt solution a t a ligand concentration [L], and go is that in the absence of ligand. Figure 1 shows values of R for PipHPi plotted us. [L] for the phosphate ligands TBP (0.229 mM in salt), MCP (0.231 miW in salt), and BCP (0.240 mM in salt) as well as values already obtained' for Ph3P0 (0.244 mM in salt). The curves for the three phosphates are distinctly S-shaped, and close inspection of the curve for the phosphine oxide shows that it too is sigmoid. If a 1 : l cation-ligand complex is the only complex forming in a salt-ligand system, eq 1, then it has been s h ~ w n ~that , ' ~R is related to the ligand concentration by
R = 1
+ KL[L]
(2)
Obviously more complex interactions are occurring in these systems. We will not attempt to analyze these results to extract equilibrium constants but will discuss them qualitatively in the following section. Figure 2 shows the values of R for the salts PipHPi (0.229 mM), MePipHPi (0.309 mM), and EtPipHPi (0.306 mM) with tributyl phosphate (TBP) as ligand. The curvature upward in these graphs for the N-methyland N-ethylpiperidinium picrates indicates13 that a second ligand molecule is adding on to the cation-ligand complex, eq 3 BH+,L
Figure 2. Conductance ratios, R, for piperidinium picrate (H), AT-methylpiperidinium picrate (GI&), and AT-ethylpiperidinium picrate (C2Hi) with tributyl phosphate as addend in chlorobenzene a t 25".
I n such a case R is related to the ligand concentration by
R = 1
+ KL[L]+ KLKzL[LI~
(4)
Wemay then plot (R - 1)/ [L]us. [L]. The intercept a t [L] = 0 is K L and the slope is KLKu,. Values of KLSO obtained appear in Table I. The slopes were so small and/or uncertain that we do not report here any values of K ~ L . Values of KL for the methyl- and ethylpiperidinium
Table I : Ligand Association with Methylpiperidinium and Ethylpiperidinium Cations in PhCl a t 25' ---Ligands------
10-3 KL(Me), M-1 KL(Et)IKL(Me) YPO, cm-1 i4
D
PhsPO
TBP
MCP
150" 1.860 1200b 4.370
44 1.27 126OC 3.07'
125 1.14 1300d 5.50
BCP
110 1.12 1325c
7.1Oh
a Reference 7. b Reference 4. 0 References 2 and 3. Obtained in this laboratory. The ir spectrum was obtained using a neat sample in a KBr cell on a Perkin-Elmer 337 spectrophotometer. e G. M. Phillips, J. S.Hunter, and L. E . Sutton, J . Chem. Soc., 146 (1945). f G. K. Estok and W. W. Wendlandt, J. Amer. Chem. Soc., 77, 4767 (1955). 0 This work. T. L. Brown, J. S.Verkade, and T. S. Piper, J. Phys. Chem., 65, 2051 (1961). The dipole moment given is for the 4-methyl derivative rather than B C P which is the 4-ethyl derivative. The difference is expected to be negligible.
+L (13) J. B.Ere11 and W. R.Gilkerson, J. Phys. Chem., 72, 144 (1968).
The Journal of Physical Chemistry, "01. 76,No. 21, 1971
A. L. MIXONAND W. R. GILKERSON
3312 cations for both the monocyclic phosphate, 3 (MCP), as ligand and the bicyclic phosphate, 4 (BCP), as
ligand were obtained as described above and are listed in Table I. Also listed in Table I are values of KL for methyl- and ethylpiperidinium cations with triphenylphosphine oxide , Ph3PO previously repor ted.7 The dielectric constants of solutions of the monocyclic phosphate in benzene at 25" are related t o the mole fraction of solute, x, by the empirical equation, E = EO 4 2 . 1 ~ . The dielectric constant of pure benzene eo was taken as'l 2.274. The dipole moment of RlCP was calculated using the method of I3ede~trand.l~In this method, the density is taken to be a linear function bxz. If the of the mole fraction of solute, p = p1 volumes of solute and solvent are additive the value of b can be shown t o be given by b = ( M 2 - p1V2'>/V1', where M Pis the molecular weight of solute, and Vlo is the molar volume of pure component i. It was assumed that the molar distortion polarization of the solute at infinite dilution is equal to that of the pure solute, i.e., the molar refraction of MCP. This last should not be a critical assumption since this contribution is only 10% of the total molar polarization. We calculated the dipole moment of MCP to be 5.52 D.
+
+
Discussion The effects of added Ph3PO and the phosphate esters on the conductance of PipHPi are more complex than we had previously b e l i e ~ e d . ~We believe the sigmoid shapes of the R us. [L] curves to be due to the combined effects of the following reaction, in addition t o eq 1 and 3 PipHPi
+ L 1_ PipH+,L,Pi-
K2 =
(5)
[PipH+,L,Pi-] [PipHPi][L]
where PipH+,L,Pi- represents a complex between ligand and the ion pair 5. The ligands Ph3PO and the L
5
phosphate esters are interacting through the phosphoryl oxygen, which is sterically most favorably situated to add to the cation in the ion pair as shown. These are not the only equilibria involved, however. We are able to fit these data using one set of in the low [L] region but parameters KL, K z , and need to adopt another set in the high [L] region. One set will not allow a fit over the entire concentration region. It is not necessary for the main purpose of the present study to have values of K L for the piperidinium ion with these ligands, so we leave that topic. The Journal of Physical Chemistry, Vol. 76, N o . $1, 1971
We shall be concerned in the main with the values of KL(Me) for the N-methylpiperidinium cation and KL(Et) for the N-ethylpiperidinium cation with the ligands PhaPO, TBP, MCP, and BCP. We had previously found16 that with Bu3NH+ as the cation, KL generally increased in magnitude as the dipole moment of the ligand increased. Here with the phosphate esters, KL(RIe) values increase and then decrease as the dipole moments increase (Table I). The ratios KL(Et)/KL(Me) decrease as the moments increase. We believe the manner in which KL(Me) and KL(Et) change from ester to ester to be the result of two opposing effects. The first is a decreasing electron density on the phosphoryl oxygen as L changes from TBP to MCP to BCP (this will be discussed more fully below), and the second is the increase in dipole moments of the ligands in the order TBP < 1ICP < BCP as first two and then all three alkoxy groups are constrained to a conformation in which the ROP group dipoles have their negative ends roughly pointed in the same direction as the phosphoryl oxygen. The ROP group moments contribute proportionally much more to the total molecular dipole moments of the phosphate esters 1ICP and BCP than t o the electrostatic potential energy of interaction between the positive charge on the cation and the various group dipoles in the ester ligands. Our view of the cation-ligand complexes places the =N+H group from the cation along the phosphoryl oxygen-to-phosphorus axis. The center of the phosphoryl group moment then is closest to the site of positive charge on the cation, with the alkoxy dipoles further removed from the cation charge. Since iondipole interaction energy falls off as the inverse square of the separation distance, the alkoxy dipoles make a proportionally smaller contribution to the potential energy of the cation-ligand complexes in the case of MCP and BCP than to the total molecular dipole moments of these two ligands. The decrease in KL(Et)/KL(;CIe) as L is varied from Ph3P0 to TBP to MCP to BCP we interpret as indicating a decrease in net negative charge on the phosphoryl oxygen in the order Ph3P0 > TBP > MCP > BCP. Note that the values of VPO, the phosphoryl oxygen stretching frequency, included in Table I, increase in the same order. Such increases in VPO have been correlated elsewhere4 with decreasing negative charge on the phosphoryl oxygen. The relative negative charges on the phosphoryl oxygens in this series of phosphate esters, TBP, 1ICP, and BCP, are in an order similar to that found in Boyd's calculations (trimethyl phosphate greater than methyl ethylene phosphate) and are inverted when compared to the order calculated by Collin for phos(14) G. Hedestrand, 2.Phys. Chem. (Leipzig),B2, 428 (1929). (16) W. R . Gilkerson and J. B. Ezell, J . Amer. Chem. Soc., 8 9 , 808
(1967).
INTERMOLECULAR HYDROGEN BONDING phate diester anions. The comparisons between our experimental results and the results of both sets of calculations are not as direct as we should like. The theoretical calculations of Collin and of Boyd were carried out with model compounds which are not the same as any in our series but several in each set bear a strong resemblance to ones in our present set. We conclude from our experimental results that Boyd’s
3313 set of calculations6 is more valid than the other with the regard t o the net charge on the phosphoryl oxygen. We are more satisfied that our previous explanation* of the inversion of the values of K L for the methyl- and ethylpiperidinium cations in the case of the oxygen-containing ligands is a correct one in view of the excellent correlation of the values of the ratios KL(Et)/KL(Me) and VPO.
Intermolecular Hydrogen Bonding. I. Effects on the Physical Properties of Tetramethylurea-Water Mixtures1 by K. R. Lindfors,* S . H. Opperman, M. E. Glover, and J. D. Seese Department of Chemistry, Central Michigan University, Mt. Pleasant, Michigan 48868
(Received March $1, 1971)
Publication costs assisted by Central Michigan University
Vapor pressures, viscosities, densities, surface tensions, heat of mixing, and molar refractivities of watertetramethylurea mixtures were measured at temperatures from 25 to 85’. The data show that the tri- to pentahydrates are the most stable water-tetramethylurea species formed.
Introduction Recently there has been considerable interest in the interactions in hydrogen bonded binary mixtures. 2 , 3 The powerful interactions which occur between a protic and a dipolar-aprotic liquid produce marked changes in many chemical3band physical4properties. Tetramethylurea (TMU) and water form such a binary pair which has not been extensively studied. We report here studies on the TMU-water system using classical physical-chemical techniques. The effects of the TMU-water hydrogen bonds on density, viscosity, surface tension, vapor pressure, refractive index, and heat of mixing of these solutions were measured. The main purpose of the investigation was to determine the most prominent TMU-water complex species present in solution at various temperatures. When TMU and water are mixed, some of the hydrogen bonds between the water molecules are broken. New hydrogen bonds form between the water molecules and the TAW molecules. Thus the deviations in properties of the TMU-water mixtures from a linear interpolation between the properties of the pure components result not only from the formation of TMUwater complexes but also from the disruption of the water structure. However, large deviations in physical properties are strong evidence for the existence of interactions. The presence of a hydrogen bond donor,
the water, and of possible acceptors, the oxygen and nitrogens of the TMU, suggest that at least a major part of these interactions will be hydrogen-bond in nature. The chemical and physical properties of TMU have been reviewed in considerable detail previo~sly.~It is a clear, polar, aprotic liquid which is miscible in all proportions with water and all common organic solvents. TMU has been suggested as a useful reaction medium. I n its solvent properties, it resembles pyridine and dimethylformamide except that it has a higher boiling point (176.5’ (760 Torr)).6 It is an important medium for aryl deaminations’ and is the solvent of choice for higher temperature Ullmann reactions.$ TMU also increases the rate of alkylation reaction^.^ (1) Work supported in part by The Ott Chemical Co., Muskegon, Mich. (2) A. K . Covington and P. Jones, Ed., “Hydrogen-Bonded Solvent Systems,” Proceedings of a Symposium on Equilibrium and Reaction Kinetics in Hydrogen-Bonded Solvent Systems, University of Newcastle upon Tyne, Jan 10-12, 1968, Taylor and Francis Ltd., London, 1968. (3) (a) J. F. Coetzee and C. D. Ritchie, “Solute-Solvent Interactions,” Marcel Dekker, New York, N. Y., 1969; (b) A. J. Parker, Chem. Rev., 69, 1 (1969). (4) G. C. Pimentel and A. L. McClellan, “The Hydrogen Bond,” W. H. Freeman and Co., San Francisco, Calif., 1960. (6) A. Lllttringhaua and A. W. Dirksen, Angew. Chem. Int. Ed. Engl., 3, 260 (1964). (6) W. Mischler and C. Escherich, Chem. Ber., 12, 1162 (1879). (7) K. G. Rutherford and W. Redmond, Org. Syn., 43, 12 (1963). The Journal of Physical Chemistry, Vol. 76, No. 21, 1971
