Chemical Shifts of 2-(2,2,2-Trifluoroethoxy)ethanol
The Journal of Physical Chemisfty, Vol. 83, No. 15, 1979
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Fluorine Chemical Shifts and Microheterogeneity in Aqueous Solutions of 2- (2,2,2-Trif luoroethoxy)et hanol Norbert Muller Deparfment of Chemistry, Purdue University, West Lafayetfe, Indiana 47907 (Received March 26, 1979)
The title compound was prepared in the expectation that its water solutions would show interesting structural effects similar to those suggested by a variety of measurements for aqueous tert-butyl alcohol and aqueous 2-n-butoxyethanol and that the nature of these effects might be elucidated by following the fluorine chemical shift as a function of concentration and temperature. The material is miscible with water above room temperature, but an upper critical solution temperature was found between 1 and 2 "C. Chemical shift-composition curves determined at five temperatures spanning the interval from 1 to 61 "C reveal a change in the nature of the solutions when the mole fraction of the nonaqueous component is about 0.18. It appears that in the dilute region the cosolvent molecules are preferentially solvated by water, especially at the lower temperatures, while the more concentrated solutions are microheterogeneous, with rapid exchange of the organic species between water-rich and cosolvent-rich domains. It is very likely that the properties of aqueous solutions of tert-butyl alcohol, 2-n-butoxyethanol, and 2-(2,2,2-trifluoroethoxy)ethanol indeed reflect essentially the same type of behavior.
Introduction From fluorine chemical shifts in mixtures of water and organic cosolvents containing 1 vol % of the probe, 6,6,6-trifluoro-l-hexanol, it was inferred that the organic liquids could be arranged in an order reflecting their increasing tendency to form microheterogeneous mixtures with water.l Dioxane, ethylene glycol, and 2-methoxyethanol stand a t one extreme, with little or no indication that their aqueous solutions are anything other than random molecular mixtures a t any concentration. For aqueous solutions of acetone, 1,2-dimethoxyethane, methanol, and dimethylformamide,appreciable deviations from random mixing appeared when the volume fraction of the organic component was in the range 0.4-0.8. The most striking deviations were found a t volume fractions between 0.2 and 0.8 in the systems water/tetrahydrofuran and waterltert-butyl alcohol (t-BuOH). At almost the same time a variety of other experimental evidence was reported which likewise points to the existence of solute clusters or other local ordering in several water/organic solvent systems. Most of these results were obtained with water/ t-BuOH mixture^,^-^ which were first recognized as microheterogeneous systems on the basis of low-angle X-ray scattering data several years earlier.6 Some were obtained with water/tetrahydrofuran solutions.' Quite recently, it was reported that for aqueous 2-n-butoxyethanol the concentration dependence of the apparent molar volume, expansibility, and heat capacity closely resembles that of aqueous t-BuOH, so that it should be possible to rationalize the behavior of both types of solution with a single model.8 T o obtain further insight into the nature of these mixtures it was desirable to find a fluorinated organic fluid which might be expected to give aqueous solutions resembling those of tetrahydrofuran, t-BuOH, or butoxyethanol, since such solutions could be studied by fluorine magnetic resonance without the need of introducing an extraneous component as a probe. The readily available analogue of t-BuOH, 2-trifluoromethy1-2-propano1, is not suitable because of its limited water solubility. Instead, 2-(2,2,2-trifluoroethoxy)ethanol(I) was chosen for study because it is relatively easy to prepareg and structurally related to butoxyethanol; moreover, noting that both 2-trifluoromethyl-2-propanol and l,l,l-trifluoro-2-propanol 0022-3654/79/2083-2041$01 .OO/O
are much less water soluble than their unfluorinated analogues, it is reasonable to expect that I and 2-n-butoxyethanol may not differ too greatly in hydrophobicity. Fluorine chemical shifts of aqueous solutions of I as a function of composition a t several temperatures, presented below, indeed reveal deviations from random mixing which appear to be closely similar to those which occur in aqueous t-BuOH or 2-n-butoxyethanol.
Experimental Section Compound I was prepared from trifluoroethanol and ethylene oxide as described by Brey and Tarrant.g The crude yield was about 90%, and a fraction boiling a t 69.5 "C (80 mmHg, uncorrected) was retained for subsequent work. (Anal: Calcd for C4F3H702: C, 33.34; H, 4.90. Found: C, 33.34, H, 5.01.) A somewhat higher boiling point, 84 "C (80 mmHg), was reported earlier,g but the density of our product, 1.289 f 0.002 at 25 "C, agrees with the reported valueg of 1.2902. The material is completely miscible with water at room temperature, and an aqueous solution containing 50% I by volume could be heated in a sealed tube to 180 "C without showing phase separation. This would suggest that I is less hydrophobic than 2-nbutoxyethanol, which in water has a lower critical solution temperature8of 49 "C. However, cooling aqueous solutions of I reveals the existence of an upper critical solution temperature lying between 1 and 2 "C. The volume fraction of I a t the consolute point is near 0.4; detailed information about the coexistence curve was not obtained. For the NMR determinations solutions of I and distilled deionized water were prepared in 5-mm sample tubes by means of volumetric syringes, and a sealed capillary of trifluoroacetic acid was placed in each tube to provide a reference signal. NMR spectra were obtained with a Perkin-Elmer R-32 spectrometer operating at 84.669 MHz and equipped with a variable-temperature accessory. Experience with this instrument indicates that temperatures may be expected to be within a t most *l "C of the nominal values. Small temperature deviations are probably the main source of error in the chemical shifts, which should be good to about fO.O1 ppm. The fluorine signal is a triplet as expected, with J H F N 9 Hz, the exact value being slightly solvent dependent. For neat I, the signal is downfield from the reference, and it is shifted @ 1979 American Chemical Society
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Norbert Muller
The Journal of Physical Chemistry, Vol. 83, No. 15, 1979
VOLUME FRACTION, Figure 2. Chemical shifl of I in water at 1 OC for volume fractions between 0.15 and 0.75 with results of model calculations (see text). 02
04
06
VOLUME FRACTION,
08
10
#)
Flgure 1. Chemical shift (downfield from external trifluoroacetic acid) as a function of concentration for I in water at several temperatures. To aid in distinguishing the curves, they have been displaced as follows: top curve, 1 OC, undisplaced; second curve, 16 O C , 0.2 ppm added to each chemical shift; third curve, 37 OC, 0.4 ppm added; fourth curve, 46 OC, 0.6 pprn added; lowest curve, 61 OC, 0.8 ppm added. The dashed region in the top curve corresponds approximately to the solubility gap.
farther downfield as water is added. The results obtained at five equally spaced temperatures between 1 and 61 “ C are presented in Figure 1.
Discussion The chemical shift of a trifluoromethyl group in an ideal binary mixture is expected to depend linearly on the volume fraction.lJO The curves in Figure 1 show large deviations from this behavior, and these are not of the type that could be attributedl to the negative AV of mixing generally found for aqueous organic systems. (Experimental values of AV of mixing for the solutions used in this work are not now available.) Unlike the hexane/ perfluorohexane system,l0where the deviations vary quite smoothly with composition, the curves in Figure 1 show fairly sudden changes in slope when the volume fraction of I, abbreviated 4, is near 0.2. In the most dilute region the shift varies linearly with 4, and the initial slope is strongly temperature dependent. When 4 > 0.2 the curves are very similar, and it is especially noteworthy that if the 1 “ C curve is continued smoothly through the solubility gap it closely resembles all the others. It thus appears that the change in slope near 4 = 0.2 is not directly connected with the existence of the critical point near 4 = 0.4. Then it should be permissible to discuss the overall shape of the curves without explicitly allowing for the small solubility gap a t 1 “C. Water structure effects are expected to be maximal a t the lowest temperatures, and this suggests looking first at the 1 “ C data. Here the behavior in the highly aqueous region provides strong support for a clathrate model of the sort proposed earlier for aqueous t-BuOHIS,lland aqueous 2-n-b~toxyethanol.~ If each alkoxyethanol molecule occupies a separate cavity in a fluid but clathratelike network, the net effect is a preferential solvation of the organic species by water, even though the organic molecules are predominantly hydrophobic. The chemical shift, which depends almost entirely on interactions between the fluorinated groups and their nearest neighbors, should then remain independent of 4, as observed. This can continue only as long as sufficient water molecules are available to form a nearly complete cage about each molecule of I. The data imply that this cutoff is reached when 4 0.185 or, in terms of mole fractions, X = 0.0354.
The behavior of the chemical shifts a t higher concentrations suggests that the solutions have become microheterogeneous, consisting of fairly small concentrated and dilute domains with rapid exchange of material between them. It is not possible to specify the size of the domains or their composition, but it seems plausible to regard the phenomenon as being analogous to a phase separation,lP2p8 taking the dilute domains as a microphase rich in water and in some sense “saturated” with I, while the concentrated microphase is rich in I and similarly “saturated” with water. Then the composition of the domains should coincide with the values of 4 which define the upper and lower boundaries of the pseudo-two-phase region. These data d a not reveal how much water must be added to pure I before the dilute microphase first appears; it seems most reasonable that small amounts of water would be present as monomers or very small oligomers hydrogen bonded to ether oxygens or hydroxyl groups of I. In this region, the rate of change of the chemical shift with 4 should be rather small, as observed, because the intercomponent hydrogen bonding would reduce the chances of contact between water molecules and trifluoromethyl groups. The following two model calculations were made with the arbitrary assumption that the upper limit of the pseudo-two-phase region is a t 4 = 0.75, corresponding approximately to the composition of a dihydrate of I, in order to show that such a microheterogeneous system would indeed give chemical shift curves resembling those found here. 1. By assumption, a unit total volume of solution with 0.185 < 4 < 0.75 contains a volume Vd of dilute material with c $ ~ = 0.185 and a volume Vc of concentrated material with C#J~= 0.75. If the chemical shift of I in a dilute domain is ad = 4.395, and the corresponding value in a concentrated domain is aC = 3.123, one can calculate the shift as a function of the overall volume fraction from the simultaneous equations vc Vd = 1 (1) V“C
+ + VQJd
6 = (6CV“C
=4
+ 6dVdrp)/4
(2) (3)
Substitution of the appropriate numerical values yields 6 = 2.707 + 0.31305/4 (4) Shifts calculated in this way are shown as curve A in Figure 2. The curve has approximately the right shape, but the predicted values lie too far upfield in the middle of the concentration range. 2. A need for modifying this model emerges when it is recalled that in a microheterogeneous system the high state of subdivision of the microphases demands that a significant amount of material must be found a t any moment
Chemical Shifts of 2-(2,2,2-Trifluoroethoxy)ethanol
The Journal of Physical Chemistry, Vol. 83, No. 15,
a t the boundary between domains. This suggests a three-site model, with Vc, Vd, and v' representing the volumes of concentrated, dilute, and interfacial material. No rigorous procedure exists for calculating v' or the corresponding composition and chemical shift, q5i and 6'. For an exploratory calculation it is natural to take qY = (@ 4d)/2 and 6' = (SC fid)/2 and to adopt some simple formula for v' that allows perhaps 15% of the material to be in the interfacial region when Vc = Vd. A second requirement is that v' must approach zero when either Vc or Vd does so. The relation V = VcVdhas these properties, and although it cannot be expected to give an accurate result it should provide a test of whether or not an effort to allow for the interfacial material represents an improvement in the model. Instead of (1) through (3) one now has
+
+
+ v d + VCVd = 1 V"C + Vdfp + VCVd(C$C + $d) / 2 = q5 6 = (6CVCq5C + 6dVdf#P+ 6iv'$i)/q5 vc
(5) (6)
(7)
With the numerical parameters used before, the results now define curve B of Figure 2, which lies much closer to the experimental points than curve A. It seems obvious that the data could be fit as nearly as desired by adjusting details of the model calculations, such as the values taken for Cbc and 6c or the recipe used to evaluate v'. Such an adjustment of parameters would not be rewarding because the resulting refinement would still represent a severe oversimplification. However, the quality of the fit obtained with curve B strongly supports the conclusion that the microphase separation model is useful at least as a first approximation to describe the nature of these solutions. Very similar conclusions were drawn from our earlier work with trifluorohexanol in aqueous t-BuOH and tetrahydrofuran' and supported by noting that a large body of other experimental workl1J2 with t-BuOH is consistent with the existence of alcohol-rich domains when the mole fraction exceeds 0.04 (volume fraction 0.18). Quite recently, laser light scattering experiments3showed an abrupt increase in the magnitude of the concentration fluctuations when the mole fraction of t-BuOH exceeds a threshold value near 0.05. These observations extended up to a mole fraction of 0.5 (volume fraction 0.84), and the anomalously large concentration fluctuations appeared to persist even a t the highest concentrations. Turning now to the curves for the higher temperatures, one finds that the major observed changes are a progressive increase in the initial slope and a broadening of the transition corresponding to the first appearance of the concentrated microphase, together with a gradual shift of this transition t o lower concentration. The change in the initial slope is consistent with the supposition that the clathratelike structures which are prevalent a t 1 "C gradually break down on heating. Consequently the probability of contact between pairs of organic molecules increases, and the chemical shift should come to depend on the concentration about as observed. At 61 "C the initial slope is high enough to indicate that the tendency of I t o be preferentially solvated by water in dilute solutions has largely disappeared. The changes in the nature of the transition between the dilute region and the microheterogeneous region suggest that clathratelike hydration of I and microphase separation are t o some degree competing processes. At 1 "C, the clathratelike structures are so stable that their presence dominates the behavior as long as sufficient water is
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available, i.e., up to q5 = 0.18. When the hydration cages begin to break down on heating, the tendency toward microheterogeneity remains and is able to express itself more freely, and its effect on the chemical shift is noted a t lower concentrations. Thus the ability of water to encage hydrophobic molecules may not in itself be responsible for the microphase separation phenomenon, but it does account for the fact that a t low temperatures microheterogeneity is not observed until a threshold concentration is exceeded. The abrupt onset of microheterogeneity above a "critical" concentration is reminiscent of micellization in surfactant solutions,1$8as is the fact that the concentration fluctuations do not progress to the point where bulk phase separation is observed. Indeed, as already noted by ROUX, Perron, and Desnoyers,8 the presence or absence of a consolute point seems to have no direct bearing on the occurrence of microheterogeneity. However, there is a fundamental difference between the two phenomena, in that for a typical long-chain surfactant the critical micelle concentration is reached long before the mole fraction of water falls to the point where it becomes difficult to build a separate clathration shell around each hydrophobic chain. In principle, butoxyethanol could be converted to a nonionic surfactant of the n-alkyl poly(oxyethy1ene)glycol monoether type13 by appropriately enlarging both the hydrophilic and hydrophobic ends of the molecules, and one might be able to observe a gradual transition between the sort of microheterogeneity found here and an ordinary micellar solution. If it is supposed that the solution at r$ = 0.185 and 1"C is more extensively structured than water itself and that this structure "melts out" rather rapidly on heating as suggested by the chemical shifts, it might be expected that the partial molar heat capacity of I would pass through a maximum a t this concentration which would gradually disappear on raising the temperature. Just such behavior was reported8 for the water/2-n-butoxyethanol system, where the apparent molar heat capacity a t 4 "C shows a sharp peak a t a mole fraction of 0.025 (volume fraction 0.16), which becomes less pronounced and shifts to lower concentration on heating. Somewhat similar findings were also p r e ~ e n t e dfor ~ , ~aqueous t-BuOH, except that here the heat capacity maximum occurs a t a higher mole fraction, near 0.07 (volume fraction 0.3), where other data1J1J2 indicate that both microphases are already present. In conclusion, this work vividly supports the emerging view that a number of binary aqueous mixtures with moderately hydrophobic organic liquids undergo a significant structural change when the volume fraction of the nonaqueous component is about 0.18. In the dilute region the cosolvent molecules are kept apart by the formation of hydration cages about each of them, but this effect is weakened when the temperature is raised. The more concentrated solutions consist of water-rich and cosolvent-rich domains. This behavior may play a key role in determining some of the dramatic changes in the spectroscopic, kinetic, or thermodynamic properties of additional components when they are dissolved in such binary mixed solvents.lJ1J2 Acknowledgment. The author thanks Mr. Joseph C. Deaton for making many of the chemical shift determinations. References and Notes (1) N. Muller, J . Magn. Reson., 28, 203 (1977). (2) C.de Visser, G. Perron, and J. E. Desnoyers, Can. J. Chem., 55, 856 (1977). (3) K. Iwasaki and T. Fujiyama, J. Phys. Chem., 81, 1908 (1977).
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The Journal of Physical Chemistry, Vol. 83, No. 15, 1979
(4) M. A. Anisimov, V. S. Esipov, V. M. Zaprudskii, N. S.Zaugol'nikova, G. I. Ovodov, T. M. Ovodova, and A. L. Seifer, Zh. Strukt. Khim. (Engl. Trans/.),18, 663 (1977). (5) K. Iwasaki and T. Fujlyama, J . Phys. Chem., 83, 463 (1979). (6) F. Franks, "Water. A Comprehensive Treatise", F. Franks, Ed., Vol. 2, Plenum Press, New York, 1973, p 23. (7) 0. Kiyohara and G. C. Benson, Can. J . Chem., 55, 1354 (1977). (8) G. Roux, G. Perron, and J. E. Desnoyers, J . Solution Chem., 7, 639 (1978).
(9) M. L. Brey and P. Tarrant, J . Am. Chem. Soc., 79, 6533 (1957). (10) N. Muller, J . Phys. Chem., 83, 1393 (1979). (11) D. N. Glew, H. D. Mak, and N. S.Rath, "Hydrogen-Bonded Solvent Systems", A. K. Covington and P. Jones, Ed., Taylor and Francis, London, 1968, p 195. (12) E. M. Arnett, "Physico-Chemical Processes in Mixed Aqueous Solvents", F. Franks, Ed., Elsevier, New York, 1967, p 105. (13) J. M. Corkill, J. F. Goodman, and R. H. Ottewill, Trans. Faraday Soc., 57, 1627 (1961).
Internal Rotation in Octafluorotoluene. A Variable Temperature Carbon-13 Spin-Lattice Relaxation Study Jozef Kowalewskl" and Anders Erlcsson Department of Physical Chemisfry, Arrhenius Laborafory, University of Stockholm, S- 106 9 I Stockholm, Sweden (Received December 18, 1978) Publication costs assisted by the Swedish Natural Science Research Council
Carbon-13 spin-lattice relaxation times and nuclear Overhauser enhancement factors are reported for neat octafluorotoluene at six temperatures in the range from 268 to 346 K. The dipole-dipole relaxation times are separated and the random jump and the stochastic diffusion models are applied to calculate the rate of internal rotation of the CF3 group. The temperature dependence of the internal rotation rate is found to follow Arrhenius-type behavior with an activation energy of about 5.5-6 kJ-mol-l.
Introduction Since the introduction of commercial FT NMR spectrometers, carbon-13 spin-lattice relaxation time (5"') measurements have become increasingly popular.' In particular, relaxation measurements for proton-bearing carbons, where the dipole-dipole mechanism is usually dominant, have been an invaluable source of the information about the dynamic behavior of chemical systems. On the other hand, only a few papers dealing with 13C T1 measurements for fluorine-bearing carbons have appeared in the literature. The first report of the measurements of both TI and nuclear Overhauser enhancement factors (NOEF) in some simple fluorocarbons was given by Hawkes and AbrahamS2 More recently, Lyerla and VanderHart3 reported a I3C T1 study of perfluoroalkanes and were able to conclude that internal chain motion in these compounds is considerably more restricted than in their hydrocarbon counterparts. In this paper we apply the carbon-13 T1measurements to octafluorotoluene in order to quantify the internal motion of the CF3 group. As is usual in this type of work, we assume that the relaxation rate of fluorine bearing 13C nuclei due to dipole-dipole interactions with fluorines (l/TIDD) may be entirely assigned to the directly bonded atoms, that the extreme narrowing conditions apply, and that the modulation of the dipolar interaction may be described by a single effective correlation time, rceff,for each carbon. These assumptions allow a simple calculation of the effective correlation time from measured 13C T I values by means of eq 1,where nF is the number of fluorine
atoms directly bonded to the carbon atom under consideration and rCF is the bonding distance. For the simplest case of isotropic reorientation of a rigid body the effective correlation time is simply related to the rotational 0022-365417912083-2044$0 1.OOlO
diffusion constant, D1 (eq 2). For the treatment of the rceff = 1 / ( 6 D )
(2)
internal motion of the CF, group we adopt the formalism developed originally by Woessner4 and modified for the case of 13C in the methyl group by Kuhlmann and Grantq5 Basically, the model assumes a random reorientation of a spin pair (in this case 13C19Fin the CF, group) about an axis which, in turn, undergoes random tumbling. We assume the overall tumbling to be isotropic and apply the random jump and stochastic diffusion modelsS4In both these models the expression for r,efffor the CF3carbon is given by eq 3. Here A , B, and C are geometrical factors.6 rceff = B + C A + (3) 6D 6 D + D[ 6 D + D[' If tetrahedral angles are assumed between the CF bonds in the CF3 group, the following values are obtained: A = 0.111, B = 0.296, and C = 0.593. D( and Dill are related to the internal rotation rate and depend on the model chosen. If random jumps are assumed to occur between the six equivalent equilibrium positions we obtain D{ = Dill = 3Ri, where Ri is the number of reorientational jumps per second. In the stochastic diffusion model the relation between D[ and D[l is ' l 4 D [ / = D[ = Di, where Di is the internal rotation diffusion constant (in rad2 s-'). Thus, TIDDvalues for the fluorine-bearing ring carbons may be applied to the calculation of the overall rotational diffusion constant, by using eq 1and 2. Following that, TIDD values for the CF3 carbon may be used to extract the internal rotation rate parameters by means of eq 1 and 3. In accordance with usual procedures we assume that all diffusion and rate constants follow Arrhenius-type behavior with respect to temperature. Thus, the jump rate is expected to follow eq 4a. The activation energy for this Ri = R t exp(-Er/RT)
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