J. Phys. Chem. 1083, 87,325-328
325
Effect of Ethanol on the Conductivity of Supercooled Aqueous Potassium Chloride Solutions R. J. Speedy,’ J. A. Ballance, and 6. D. Cornlsh Chemistry Department, Victoria University of Wellington, Wellington, New Zealand (Received August 4, 1982; I n Final Form: August 16, 1982)
Adding ethanol (up to 3 M) to 0.1 M aqueous KC1 solutions increases both the resistivity (by up to 300%) and its anomalous temperature coefficient. The effects are more pronounced at lower temperatures. The temperature dependence of the resistivity can be described by r = r,(T/T, - l)?,where r is the resistivity, r2 and y are constants, and T, is the temperature at which r extrapolates to infinity. The notion that the “iceberg”structure induced in water by ethanol may be the same as the structure of the low-density species whose cooperative formation is thought to underlie the above temperature dependenceof many properties of pure water is discussed. We conclude, with some qualifications, that ethanol, unlike polar or ionic solutes, enhances the cooperative process.
Introduction The intention of this work was to test the idea that the hydrophobic cage or structures induced in water by nonpolar solutes are related to the structures which exist in pure water and which give rise to its lowtemperature anomalie~.~Guided by the principle of parsimony one expects common cause for the anomalies of water and hydrophobic effects. The unusual properties of supercooled water can be interpreted, in a qualitative way, by supposing that there are some low-density species which form cooperatively in the random network of hydrogen-bonded water molec u l e ~ . ~The - ~ cooperative element is required to generate clusters of the low-density species which grow in size as the temperature is reduced, giving rise to the long-range density-density correlations which are implied by the observed behavior of the thermodynamic response functions in supercooled water.5 Stillinger6p7has described a model in which the low-density species are identified as polyhedral cages of water molecules. The necessary element of geometrical cooperativity is present in the model because adjacent polyhedra can stabilize each other by sharing a face. Stillinger’s model accounts plausibly for the phenomenon of hydrophobic association by noting that hydrophobic solutes are likely to reside in the cages and, therefore, to be associated because the cages tend to be associated. Many measured properties, X,of supercooled water have a dominant temperature dependence of the f ~ r m ~ . ~ X = Xo(T/T, - 1)Y (1) where X o and y are constants and T, is the temperature at which X extrapolates to infinity. In the stability limit interpretation4p5T, is the temperature at which the decay length for density-density correlations diverges, corresponding in the terms of the above structural ideas to the
temperature at which a cluster grows to macroscopic size. If there is a direct relation between the “iceberg” structure about a solute and the cooperatively forming structures which produce the singularity at T,, then one would expect the addition of hydrophobic solutes to water to enhance the anomalous effects. Alternatively, if the “iceberg”structure is not compatible with the cooperative structures one expects suppression of the anomalous effects. A signal feature of a cooperative process is its sensitivity to perturbing influences. The dramatic effects of ionic and polar solutes, which suppress the anomalous behavior of water, have been demonstrated by Angell’s g r o ~ p . ~ ~ * ~ ~ In the percolation modello,” the anomalous transport properties of water are a consequence of the increasing size of patches of connected four-bonded water molecules in the random network. The singularity at T , corresponds to the percolation of the patches. The addition of a hydrogen-bonding solute like ethanol is expected to have a disconnecting effectll and therefore to suppress the anomalous effects. As a test of these ideas we have chosen to measure the effect of adding ethanol on the electrical conductivity of 0.1 M aqueous KC1 solutions. Conductivity measurements are simple, precise, and show a large temperature dependence, and resistivities conform to eq 1 in the absence of ethanol.12 Ethanol has the characteristics of solute which induces “iceberg”formation.13J4 Adding 6 mol % ethanol to water increases the viscosity by 50% at 25 O C 1 5 and shifts the density maximum to higher temperature^,'^ for instance. There is evidence that a 6 mol % solution of ethanol would form an aqueous clathrate crystal at -73 “C with the stoichiometry C2H50H-17H20.14 The Walden product for ions dissolved in an ethanol/water mixture passes through a maximum near this composition.16
(1)W. Kauzmann, Adu. Protein Chem., 14, 1 (1959). (2)F. Franks In “Water-A Comprehensive Treatise”, Vol. 4, F. Franks, Ed., Plenum Press, New York, 1975,Chapter 1. (3)C. A. Angell in “Water-A Comprehensive Treatise”, Vol. 7, F. Franks, Ed., Plenum Press, New York, 1981. (4)R. J. Speedy and C. A. Angell, J. Chem. Phys., 65,851 (1976). (5)R. J. Speedy, J. Phys. Chem., 86,982 (1982). (6)F. H.Stillinger, Science, 209,451 (1980). (7)F. H.Stillinger in ”Water and Polymers”, S. P. Rowlands, Ed., American Chemical Society, Washington, DC, 1980;ACS Symp. Ser. No. 127,pp 11-22.
(8)M.Oguni and C. A. Angell, J. Chem. Phys., 73,1948 (1980). (9)M.Oguni and C. A. Angell, J. Chem. Phys., in press. (10)H.E. Stanley, J.Phys. A, 12,L329 (1979). (11)H.E. Stanley and J. Teixeira, J. Chem. Phys., 73,3404 (1980). (12)R. J. Speedy, J.Phys. Chem., preceding article in the this issue. (13)F. Franks and D. J. Ives, Q.Reu. Chem. Soc., 20,1 (1966). (14)A. D. Pot& and D. W. Davidson, J. Phys. Chem., 69,996(1965). (15)R. L.Kay and T. L. Broadwater, J.Solution Chem., 5,57 (1976). (16)R. L. Kay in “Water-A Comprehensive Treatise”, Vol. 3, F. Franks, Ed., Plenum Press, New York, 1973,Chapter 4.
0022-3654/83/2087-0325$01.50/00 1983 American Chemical Society
Speedy et al.
326 The Journal of Physical ChemWy, Vol. 87, No. 2, 1983 TABLE I : Resistivity of KCl/Ethanol/Water Solutions Fitted b y r = r,(T/T, - 1 ) - y
solution 0.1 M 0.1 M 1M 0.1 M 3M
KC1 KCIethanol KC1ethanol
Trangel'C
r2/ ( a cm)
y
-27 t o 30 -20 to 3 5
16.14 16.16
1.56 1.64
-35 t o 3 5
14.10
2.02
TABLE 11: Resistivity of KCl/Ethanol/Water Solutions Fitted b y r = A e x p ( E / T ) ( T / T ,- 1)-3'2 a %
R
T,/K
std dev
220 222
1.0 0.17
221
1.3
solution
0.1 M 0.1 M 1M 0.1 M 3M
KCI KCIethanol KCIethanol
Trangei'C
A
TJK
E/K
std dev
-27 t o 30 -20 t o 35
15.06 11.99
220 224
25.5 118.7
1.0 0.19
-35 t o 3 5
3.73
225
568
1.5
a Data f o r 0.1 M KCI are the combined results for the two independent runs reported in ref 12.
a Data for 0.1 M KCI are the combined results o f the two independent runs reported in ref 12.
Conde, Teixeira, and Paponx7have studied the effect of ethanol on the sound velocity in aqueous solutions to -20 "C. The adiabatic compressibility, K,, decreases dramatically with increasing ethanol concentration. In an X-ray scattering study Bosio, Teixeira, and Stanleyxahave shown that the increase in the densitydensity correlation length, which is evident for HzO and D20 at low temperatures, is less apparent in ethanol/water solutions.
As noted previously,12there are no doubt several different factors with different temperature dependences which influence the ionic mobility. Only near the limit T T, will the factor which apparently gives rise to the above temperature dependence be totally dominant. Thus the variations in T, and y in Table I may reflect the fact that the equation is being forced to accommodate "background" terms. The variation in y can be eliminated if a background term of the Arrhenius form or the Vogel-Tamman-Fulcher is included. Table I1 shows the best-fit parameters in the equation
Experimental Section The apparatus and calibration procedure are described in more detail elsewhere.12 The conductivity cell was a pyrex capillary, about 17 cm long, sealed at each end into stainless steel pressure tubing which also acted as electrodes. The central 15 cm length of the capillary was thermostated. The resistance was measured by applying 950 V dc across the cell and a series dc voltmeter (input impedance lo7 Q f 0.1%). The current passing through the cell and through the input impedance of the voltmeter was measured as a potential on the voltmeter. For calibration purposes the conductivities of the solutions were also measured in a conventional conductivity cell, calibrated with KC1 solution^,^^ over the range 0-35 "C. Results Solutions were made up with 0.100 mol of AR KC1 plus 1.00 or 3.00 mol of ethanol and double distilled water to 1 dm3 at 20 "C. The ethanol concentrations correspond to about 2 and 7 mol %, respectively. The resistivities of these solutions are shown in Figure 1,which includes previously reported results12for 0.1 mol dm-3 KC1 solutions for comparison. The precision is about 0.2% for the 1 M ethanolic solution and 1.2% for the 3 M solution. The difference arises because the 1 M solution was studied in a 28-pm bore capillary and froze at -20 "C, whereas the 3 M solution was studied in a 7-pm bore capillary and froze at -39 "C. The cell resistance was above 10" D below -30 "C and some of the scatter is attributed to stray parallel conductive paths at the 1013-Qlevel. When the solution froze at -39 "C the resistance only increased by a factor of 5. Some measurements were made below -35 "C, but they were not reproducible and are not included. Analysis. The measurements are well fitted by the equation3p4J2 r = r,(T/T, - 1)Y with the best-fit parameters in Table I. (17) 0. Conde, J. Teixeira, and P. Papon, J. Chem. Phys., 76, 3747 (1982). (18) L. Bosio, J. Teixeira, and H. E. Stanley, Phys. Reo. Lett., 46, 597 (1981). (19) G. Jones and B. C. Bradshaw, J.Am. Chem. Soc., 55,1780 (1933).
-
r = A exp(E/T)(T/T, - 1)-3/2
Table I11 shows the best-fit parameters in the equation r = A exp(B/(T - T,))(T/T, - l)-] when Tois fixed at 160 K. (The exponent y = 1 has been derived in a mean field approximation for the shear viscosity.22 It can also be obtained simply by assuming a Taylor expansion of r-l about T, along an isobar.) The precision of the fits to these different equations is about the same, so we have no way of discriminating between alternative forms of the background term, and no particular significance can be attached to the parameter values. The trends in the values of E and B indicate that some of the increase in r with increasing ethanol concentration can be assigned to background effects; but despite this, the anomalous term (T/T, - 1)Y increases with increasing ethanol in each case. Viscosity measurements are included in Table I11 since the fit to Hardy and C ~ t t i n g t o n ' sresults ~ ~ is very good. However, the extrapolation to -35 "C with this equation is about 30% lower than the value reported by Zheleznyi et al.27 The approach used by Oguni and Angel13,8,gto separate anomalous and background contributions to the thermodynamic properties of water could also be applied to the transport properties.
Discussion Figure 1 and its inset show that the resistivity and its anomalous temperature dependence are both enhanced by the addition of alcohol. The magnitude and form of the effect suggest that alcohol couples positively with the cooperative process, even though we cannot separate anomalous and background contributions unambiguously. The enhancing effect of alcohol contrasts with the effects of ionic and polar solutes which suppress the anomalies of ~ater.~*~J~ (20) H. Vogel, Phys. Z.,22, 645 (1922). (21) E. W. Lang and H. D. Ludemann, Angew. Chem., Int. Ed. Engl., 21, 315 (1982). (22) P. Papon and P. H. E. Meijer, Physica, 101A,477 (1980). (23) R. C. Hardy and R. L. Cottington, J. Res. Natl. Bur. Stand. U.S.A.,42, 573 (1949).
Conductivity of Supercooled Ethanol-KCI Solutions
The Journal of Physical Chemistry, Vol. 87, No. 2, 1983 327
TABLE 111: Best-Fit Parameters in the Equation X = A e x p ( B / ( T- 160 K))/(T/T, solution
a
property
T range/'C
pure waterb VICP 0.1 M KC1 rl(n cm) rl(n cm) 0.1 M KC1-1 M ethanol rl(n cm) 0.1 M KC1-3 M ethanol Data for 0.1 M KCl are the combined results of
A
0.082606 0 to 125 11.43 -27 to 30 -20 t o 35 10.29 -35 to 35 7.35 the two runs reported in ref 12.
lp Ts/K
BIK
% std dev
226 222 224 226
171.08 116.7 152.4 227.1
0.07 0.9 0.16 1.2
Reference 23.
large to be accounted for simply by the presence of uncorrelated low-density species. Stanley and Teixeira's plausible estimate of this contribution'l is an order of magnitude smaller than the observed a n ~ m a l y Thus . ~ ~ ~it is necessary to invoke the idea of a cooperative clustering effect3-" which can generate long-range density correlations. In terms of the compressibility e q ~ a t i o n ~ ~ ! ~ ~
2 500
2000
&TKT = 1
+ p l ( g ( r ) - 1) dr
r/ncm
1500
I t
-20
0
20 TIo(
1000
500
0
- 40 - 20 0 20 Tl'C F@re 1. The resistivii of 0.1 M aqueous KCI solutions containing (0) 1 and (0)3 mol dm-3 ethanol. The line with no points shows the resistivity of 0.1 M aqueous KCI wlth no ethanol. The dashed extrapolations were calculated with the parameters in Table I. Inset. The Ratio R of the resistivities of the solutions containing ethanol to the resistivities of solution with no ethanol calculated with the parameters in Table I. Thus the results seem to support the idea that the anomalies of supercooled water are due to the clustering of cagelike structures which have a close geometrical relationship to the structures which are thought to form about the hydrophobic part of the alcohol molecule in aqueous solution. It is necessary, then, to explain the paradoxical evidence that (1) the addition of alcohol suppresses the long wavelength anomaly in the structure factor of supercooled water;18 (2) the adiabatic compressibility of supercooled water is dramatically reduced by the addition of ethanol;" and (3) the limiting partial molal compressibilities of solutes (including hydrophobic solutes) are decreasing and become negative as the temperature reduces below 25 OC." In other words, the isothermal compressibility of a supercooled solution is less than that of pure water. It may be possible to understand the paradox in terms of the cage model. First, it is essential to recognize that the compressibility anomaly of supercooled water is too (24) S. Cabani, G. Conti, and E. Matteoli, J. Solution Chem., 8, 11 (1979).
(where p = N / V is the density, kg is Boltzmann's constant, and g(r) is the radial distribution function) KT can become large if g(r) is greater than unity for large values of r. For KT to diverge, as suggested by the stability limit conject ~ r e , 4density ,~ correlations must persist over macroscopic distances,26since the amplitude of g ( r ) cannot diverge. We note that the percolation modellOJ1needs to be modified to account for the anomalous density correlations. In its present form it predicts enhanced density correlations at the nearest-neighbor separation of about 3 A, but not at greater distances (see Appendix A, part B, of ref ll), and so it does not account for the compressibility anomaly: or for the long wavelength anomaly in the structure factor.18 In the cage model g(r) becomes long range because the low-density cages are self-stabilizing and form correlated clusters which grow in size as the temperature is reduced.6~~ The anomalous increase in the viscosity is attributed to the longer structural relaxation times of the clusters, and we suppose that the same effect is reflected in the rapid decrease in ionic mobilities observed in supercooled water. Some evidence for long-range influences on ionic mobilities is given in ref 12. A plausible resolution of the paradox noted above can be achieved if we allow the added ethanol to fill the cages. Filling a cage with a solute molecule restores the local density to near its average value and thus decreases the amplitude of the local density fluctuation. This effect would reduce KT. However, filling the cages does not necessarily disrupt them. Since the resistivity anomaly is attributed to the structural (i.e., cage-cage) correlations rather than to the density-density correlations per se, it is not suppressed by filling the cages. Indeed, for those solutes which stabilize the cages we expect the transport anomalies to be enhanced, as observed.
Conclusion The picture of water structure which is beginning to emerge is that of a random tetrahedral network of hydrogen-bonded water molecules in which there is a cooperative process which causes clusters of a particular structure to differentiate from the bulk fluid and grow, reversibly, as the water is cooled. There is no direct evi(25) J. P. Hanson and I. R. McDonald, "Theory of Simple Liquids"; Academic Press, London, 1976, p 43. (26) H. E. Stanley, "Introduction to Phase Transitions and Critical Phenomena"; Oxford University Press, New York, 1971, p 98. (27) Yu A. Osipov, B. V. Zheleznyi, and N. F. Bondarenko, Russ. J. Phys. Chem., 51, 1264 (1977).
J. Phys. Chem. 1983, 87,328-331
328
dence to indicate the nature of these structures, except that they are of lower density than the bulk. The present measurements suggest that there is a positive coupling between these clusters and dissolved ethanol. The simplest intermetation is that the low-densitv structures in water and {he hydration shell of hydrophobic solutes are both in cagelike arrangements Of water structure to the water in crystalline clathrates.
Note Added in Proof. Halfpap and Sorensen28aand Oguni and Angel128bhave independently reached a very similar conclusion. Registry No. KC1,7447-40-7; EtOH, 64-17-5; H20, 7732-18-5. (28)(a) B. L. Halfpap and C. M. Sorensen, J. Chem. Phys., 77, 466 (1982); (b) M.Oguni and C. A. Angel], J. Phys. Chem., submitted for publication.
Interactions of Ruthenium( I I)Photosensitizers with Triton X-I 00 Krlsnagopal Mandal, B. L. Hauensteln, Jr., J. N. Demas,' The Department of Chemlstv, Unlversifj' of Vlrglnia, Charlonesvllle, Vlrglnia 2290 1
and B. A. DeGraff The Department of Chemistry, James Madlson Unlverslfy, Herrlsonburg, Vlrglnia 22807 (Received June 15, 1982; In Final Form: September 17, 1982)
Using excited-state lifetime measurements as a probe, we have studied the interactions between a series of ruthenium(I1)photosensitizers and Triton X-100 surfactant. The complexes are weakly bound to the surfactant assemblies,and the presence of the photosensitizers influences the assembly process. There is a good correlation between the strength of the binding and the nature of the ligands on the metal complex.
The use of surfactant assemblies in the study of photochemical processes has become widespread in recent years.' Most research has concentrated on using ionic micelles to bind photosensitizers and to produce a separation of photoproducts. We were interested in a surfactant system where electrostatic interactions alone would not be responsible for the binding so that we could examine more closely the interactions between the surfactants and the photosensitizers. Also, we wished to have a system which would exhibit more varied binding regions than the charged surfactants provided. This feature would provide more flexibility in the regions for binding the photosensitizers, the quenchers, and the reaction products. We report here the results of our study of the interactions between several tris(a-diimine)ruthenium(II)complexes and the neutral surfactant Triton X-100 (octylphenoxypolyethoxyethanol; C8H1,CGH4(0CH2CHz),0H, x = 9,lO). The neutral Triton should provide at least three distinct regions: the relatively polar potentially coordinating polyether region, the hydrocarbon region, and the potentially n-bonding aromatic region. We have found that the use of lifetimes of photosensitizers as a function of surfactant concentration is a powerful probe of photosensitizer surfactant interactions. In our earlier studies using cationic and anionic surfactants, we found the binding of the complexes to the micelles to be very tight with little change in the luminescence properties above the cmc. With Triton X-100, however, the binding between the micelle and the probe molecule is relatively weak, and the presence of the metal complex influences the nature of the surfactant assembly. Further, we can correlate the structure of the ligands with the strength and nature of the interactions between the com(1) For recent reviews see (a) Gratzel, M. Acc. Chem. Res. 1981, 14, 376. (b) Turro, N.J.; Gratzel, M.; Braun, A. M. Angew. Chem., Int. Ed. Engl. 1980,19,675. (c) Yekta,A.; Aikawa, M.; Turro, N. J. Chem. Phys. Lett. 1979, 63,543.(d) Kalyanasundaram, K. Chem. SOC.Rev. 1978,4, 453. ( e ) Thomas, J. K. Acc. Chem. Res. 1977, 10, 133.
TABLE I: Properties of Ruthenium( 11) Photosensitizers with Triton X-100"
RU(~PY)~I~+ Ru( phen),]l+ Ru( Clphen),lz Ru(Mephen),] ,+ Ru( 4,7-Me2phen),] R~(5,6-Me,phen),]~+ R~(Me,phen),]~+ R~(Ph,phen)(phen),]~+ Ru(Ph,phen),lz+ Ru( (SO,Ph,),phen)(phen),] R u ( ( S 0 , P h ) hen),]" Ru(Me2bpy)3217 , +
+
0.64 0.95 5.25 5.70 10.2 33.3 56 1 5.05 9.26 7.26
0.585 0.923 0.97 1.39 1.72 1.97 1.81 3.18 3.41 3.38 3.75 0.350
1.58 2.18 2.85 2.92 2.16 4.69 5.32 3.79 5.34 0.401
1.49 1.61 1.55 0.250 0.830 0.315 0.070 0.00 0.00 0.00
a Experimental fits were typically made over the Triton X-100 ranges of 0-7 or 0-25 mM.
plexes and the surfactant assemblies.
Experimental Section The ligands, and our abbreviations, are as follows: 2,2'-bipyridine (bpy) 4,4'-dimethyl-2,2'-bipyridine (Me2bpy), 1,lO-phenanthroline (phen), 5-chloro-1 , l O phenanthroline (Clphen), 5-methyl-1,lO-phenanthroline (Mephen), 4,7-dimethyl-l,lO-phenanthroline(4,7Me2phen), 5,6-dimethyl-l,lO-phenanthroline(5,6Me2phen), 3,4,7,8-tetramethyl-l,lO-phenanthroline (Me4phen),4,7-diphenyl-l,lO-phenanthroline (Ph2phen), and disulfonated 4,7-diphenyl-l,lO-phenanthroline ((S03Ph)zphen). AU ligands were from G. Frederick Smith Chemical Co. and were used without further purification. [Ru(Phzphen)],C12was prepared by the method of Watts and Crosby2and purified by chromatography on alumina. The remaining tris homochelated complexes were prepared (2)Watts, R. J.; Crosby, G. A. J. Am. Chem. SOC.1971, 93,3184.
0022-365418312087-0328$01.50/00 1983 American Chemical Society
