335
Thermodynamics of Caffeine Aqueous Solutions data in Figure 8B, it may be seen that (at least) two relaxation regions are discernible, one approximately comparable to the one for the solvent mixture ethanol-THF and an additional one at lower relaxation frequency. A quantitative analysis of these data is not possible because of the unknown contribution of the solvent T H F to the investigated relaxation processes. T o this purpose it would be necessary to extend the frequency range to about 100 GHz. It is possible, however, to draw some qualitative conclusions from the above data. It seems that the relaxation a t the higher frequency (Figure 8B) is comparable to the relaxation of the ethanol-THF mixture and the lower relaxation (centered around the relaxation frequency f~ 23 0.5 GHz) is the result of the addition of NaC104 to the solvent mixture ethanol-THF. This relaxation frequency f~ 0.5 GHz is lower by a factor of 4 than the average relaxation frequency shown by NaC104 0.1 M in pure T H F ( f =~ 2 GHz). Since the viscosity increases only by 5% (with respect to pure THF) by the addition of ethanol, as in the ultrasonic studies, it seems likely that ethanol intervenes in
-
solvating the ion pairs, consequently lowering their rate of rotation. References a n d Notes H. Wang and P. Hemmes, J. Am. Chem. Soc., 95,5115 (1973). P. Jagbdzinski and S.Petrucci, J. Phys. Chem., 78, 917 (1974). R. M. Fuoss and C. A. Kraus, J. Am. Chem. Soc.. 55,2387 (1933). D. N. Bhattacharyya, C. L. Lee, J. Smid, and M. Szwarc, J. Phys. Chem., 69, 608 (1965). (5)C. Micheletti. Thesis, University of Paris, 1972. (6) H. Wang and P. Hemmes. J. Am. Chem. Soc., 95, 51 19 (1973). (7) S. Petrucci in "Ionic Interactions", Vol. li, Academic Press, New York, N.Y., 1971, Chapter II. (8) J. Lamb in "Physical Acoustics", Vol. 11, W. P. Mason, Ed., Academic Press, New York, N.Y.. 1965, Part A, Chapter 4. (9) H. Farber and S . Petrucci, J. Phys. Chem., 79, 1221 (1975). IO) R. M. Fuoss, J. Am. Chem. SOC.,80, 5059 (1958): N. Bjerrum, K. Dan. Vidensk. Selsk., MatJys. Medd., 7 , 9 (1926). 11) M. Eigen and L. DeMaeyer in "Techniques of Organic Chemistry", Vol. VIII, Part I),S . L. Frisss, E. S . Lewis, and A. Weissberger, Ed., 2nd ed, Interscience, New York, N.Y., 1963. 12) J. P. Badiali. H. Cachet, A. Cyrot, and J. C. Lestrade, J. Chem. Soc., Faraday Trans. 2, 69, 1339 (1973). 13) Justifications for the approximation inherent in assuming a onedimension system were discussed in the original theory: S . H. Glarum, J. Chem. Phys., 33, 639 (1960). (1) (2) (3) (4)
Thermodynamics of Caffeine Aqueous Solutions A. Cesaro, E. Russo, and V. Crescenzl' institute of Chemistry, University of Trieste, 34 127 Trieste, ita& (Received August 2, 1974; Revised Manuscript Received November 3, 1975) Publication costs assisted by the Consigiio Nazionale delle Ricerche
A rather complete description of the thermodynamic properties of caffeine in aqueous solution for a wide concentration range is afforded on the basis of calorimetry, vapor pressure osmometry, solubility, and density experiments. The whole set of data is compatible with an interpretation of the concentration dependence of properties considered in terms of multiple association equilibria of the solute. Data on complex formation between caffeine and the dye ethidium bromide in aqueous solution are also reported. The thermodynamics of caffeine self-dimerization as well as of the caffeine-ethidium bromide adduct formation is discussed and compared with that for a few different species known to dimerize and/or aggregate by "vertical stacking'' in dilute aqueous solution.
Introduction Pyrimidine and purine bases, as well as a number of their derivatives, are known to self-associate in aqueous solution with the formation of dimers or higher aggregates by "vertical-stacking". A thermodynamic characterization of this behavior has been achieved in many instances with considerable accuracy by means of vapor pressure osmometry,1 ultracentrifugation,2and ~ a l o r i m e t r y . ~ ~ ~ A common feature, apparently independent of the degree of stacking of the compound under consideration, is that both the enthalpy and entropy df each association step are quite large and negative. In particular, the negative (unitary) entropy change of such association, for which hydrophobic forces would be relevant, has been a matter of debate.4 In the context of this type of studies, caffeine (1,3,7-trimethyl-2,6-dioxopurine) seems a particularly suitable com-
pound. The study of the concentration dependence of its properties should, in fact, be free from complications that arise in other aggregating systems due to electrostatic interactions and due to the possible existence of protolytic products in solution. Moreover, caffeine is reported to form 1:1 mixed molecular complexes with different organic compounds containing aromatic ring systems, a fact which has possible pharmacological implication^.^^^ There is, however, still some lack of quantitative information on the behavior of caffeine in water (for instance, according to Guttman and Higuchi16 the self-association of caffeine would proceed to the dimeric and tetrameric species only, a conclusion questioned by Gill et al.3), and nothing is known on the thermodynamics of its interactions with aromatic compounds in aqueous solution. Pursuant to our interest on the physical chemistry of nonionic compounds in aqueous s 0 1 u t i o n ~and ~ ~ in the light The Journal of Physical Chemistry, Vol. 80, No. 3, 1976
336
of what was mentioned above, we have carried out a series of solubility, vapor pressure osmometry, density, and microcalorimetry experiments aiming a t a better characterization of the caffeine-water system. The interaction of caffeine with the dye ethidium bromide and with urea has also been studied, principally by means of calorimetry. Experimental Section
(a) Materials. Urea and caffeine (C. Erba, Italy) were recrystallized from water-ethanol (1:l) and from water, respectively. Hydrated caffeine crystals were obtained in this way at around room temperature. Thermogravimetric and differential scanning calorimetric analysis revealed that such crystals lose 1 mol of water per mole of caffeine a t around 82 "C. Sublimation of caffeine leads to anhydrous crystals exhibiting x-ray powder diffraction spectra quite different from those of the monohydrate crystals. Crystalline forms were used in the calorimetric heats of solution measurements of both hydrated and anhydrous materials. Sublimed caffeine (anhydrous) was employed in most of the other experiments. (b) Methods. (1) Osmotic coefficients, @, were obtained a t 29.8 and 35.0 "C by vapor pressure osmometry (PerkinElmer, molecular weight apparatus Model 115) using sucrose (A-grade) solutions (0-0.05 m) as standards, writing @ = ( AR/m)caffei,,e/( AR/m)sucrose where AR (ohm) is the difference in thermistor resistance readings for the solution of molality m and for water. The (ARlm),,,,,,, value was independent of m , and equal to the value reached by ( AR/ m),,ff,i,, for m = 0. In our experience the reproducibility of the measurements rapidly deteriorates with decreasing working temperature using the above apparatus because of its poor temperature control near 25 "C (water as solvent). For this reason we shall not report @ values for 25 "C as we consider these values of insufficient reliability. (2) Saturated caffeine solutions were obtained by prolonged equilibration under continuous shaking with pure solid caffeine in unsaturated or supersaturated solutions in a thermostated bath, The equilibrium molality of caffeine was derived either by direct uv analysis (after appropriate dilution of a weighed amount of the equilibrated solutions) using a molar extinction coefficient of 6 9900 at 272 nm,9 or by evaporation to constant weight at 90 'C. A faster method, following a modified Davies et al.1° procedure, has yielded the same solubility data, well within experimental errors (f0.0003 m from 16 to 30 "C). (3) Density measurements were carried out using an A. Parr Model DMA-02-D microdensimeter connected with a thermostated bath at 25 f 0.0001 'C. Water and air were employed for routine instrumental calibration using literature datal1 for their densities at the temperature and pressure of the experiments. (4) Calorimetric measurements were made a t 25 'C with a LKB 10700 batch-type microcalorimeter with glass cells. In the heat of dilution experiments weighed amounts of caffeine solutions of known concentration (ca. 4 g) were a i x e d with weighed amounts of water (0.5-2 g) in the calorimetric cells. Experimental values of the heat of dilution ranged from 0.2 (at the highest dilution) to 55 mcal, with an estimated uncertainty of 0.1 mcal. In the heat of solution experiments solid anhydrous caffeine (ca. 3 mg) weighed with a semimicro analytical balance (Gibertini G-15; precision f 1 0 pg) was mixed with weighed amounts of water (or urea solution). The complete dissolution of crystalline caffeine takes place in less than 4 min under the The Journal of Physical Chemistry, Vol. 80,No. 3, 1976
A. Cesaro, E. Russo, and V. Crescenzi
experimental conditions employed. The final concentration was checked by uv absorption analysis. Caffeine hydrate crystals were kept in a saturated solution and wiped with a filter paper before use. The amount of caffeine hydrate in the heat of solution experiments was determined by uv analysis of the final solution. The interaction of caffeine (2 X to 5.3 X M) with ethidium bromide (4.84 X M) was studied with a LKB flow-type calorimeter.12 The net heat of interaction was obtained during the same experiment by subtracting the experimental heats of dilution of each substance. The enthalpy of complex formation ( A H )was obtained from the relation
AH = Q/MV where Q is the net heat of interaction (in cal/s), M is the complex concentration in the flux V (in our experiments LIS). The concentration of the 1:l complex was 9.58 x evaluated using the equilibrium constant obtained from spectral measurements. ( 5 ) Uv measurements were performed at room temperature with a Hitachi Coleman EPS-3t spectrophotometer using quartz cells of appropriate pathlengths. Spectra of ethidium bromide (1-2 X lo-* M) in the presence of different caffeine concentrations ( to M) were recorded using 1-cm thermostated quartz cells. Results and Discussion
(A) Caffeine in Water. The results of the measurements carried out on aqueous solutions of caffeine are listed in Table I. Assuming that deviations from ideality of the water-caffeine system are essentially due to solute-solute interactions leading to caffeine association, the experimental osmotic coefficient, @, values (Table I) were used to calculate the association constant with the equation1
K = (1 - @.)/mOZ (1) assuming a multiple equilibrium with K = K1 = Kz = . . . K,-1. The results are K = 9.0 f 0.1 m-l (29.8 "C), and K = 7.8 f 0.1 m-l (35.0 "C). From the density data of Table I one deduces that the apparent molar volume of caffeine goes from GVo = 144.2 ( m 0) to aV = 142.15 cm3/mol for m = 0.1, a t 25 OC. We interpret this decrease in terms of association and write
-
@" =
c
xivi,
i= 1
where X i is the mole fraction of solute species of aggregation number i having a molar volume Vi. To solve the problem one can assume that the molar increment in volume passing from i to (i 1)is a constant, u , independent of i, i.e. that Vi = avo ( i - 1 ) u . On the basis of our density data and using the association equilibrium constant, K , given above, one then easily calculates u = 138.3 cm3/mol. This implies, for instance, that the molar volume of caffeine dimers in water would be 282.5 cm3/mol of dimer, which is less by 5.9 cm3/mol of dimer than would be expected from simple additivity of the apparent volume of the monomer. Our results are in qualitative agreement with dilatometric data reported by Kasarda.13 With the same assumptions made in the analysis of the data, the enthalpy of association, AH",can be evaluated using the relative apparent molal enthalpy ( k )data of Table I and
+
+
*
AH" = cp,/(l
- cp)
(2)
337
Thermodynamics of Caffeine Aqueous Solutions TABLE I : Physicochemical Properties of Caffeine Aqueous Solutionsa Osmotic coefficientc
Solubilitvb msatdr
t
mol/kgH,O
16.60 19.30 20.50 23.80 26.08 26.44 28.88
0.0714 0.0827 0.0905 0.1026 0.1126 0.1166 0.1309
rn
29.8 "C
0.1116 0.0765 0.0590 0.0528 0.0386 0.0258
Apparent re1 molal enthaluv - - 25 OCd
35.0 "C
-
Density 25 "Ce -_I__
._
cp
m
cp
m
cal/mol
m
d , g/cm3
a,,, cm3/mol
0.622 0.683 0.723 0.745 0.781 0.835
0.1027 0.0895 0.0702 0.0509 0.0296 0.0203
0.656 0.680 0.720 0.755 0.837 0.877
0.1051 0.0632 0.0357 0.0298 0.0247 0.0169 0.0125 0.0092 0.0084 0.0044 0.0028
1150 970 71 0 640 550 420 340 240 230 110 60
0.00626 0.01054 0.02018 0.02864 0.05419 0.05477 0.07629 0.10472 0.10754
0.997360 0.997577 0.998069 0.998500 0.999840 0.999859 1.000965 1.002456 1.002558
143.1 143.8 143.5 142.7 142.5 142.6 142.5 142.0 3.42.2
-@L,
Heats of solution in water at 25 "C Crystalline anhydrous
Crystalline hydrate
__
I _ _ _ _ -
mol of caffeine
x
lo6
g H2O
38.12 46.52 47.82 46.44 38.73 0.1 kcal/mol
kcal/mol
AH,
3.31 3.54 3.49 3.50 3.33
29.0 9.12 14.5 8.72 8.74
g H,O
Q, mcal
AH,
kcal/mol
4.4290 148.6 5.08 4.0647 43.49 4.78 4.5195 76.02 5.22 3.9289 43.36 4.9'7 4.1301 44.80 5.12 i. Mean AHIS = 5.0 i 0 . 1 kcal/mol a In the case of solubility, osmotic coefficient, and density data, each value is an average of at least four determinations. b Maximum deviations from the mean value are t 3 X m. C Maximum deviations from the mean value are i 5 x d Maximum deviations from the mean value are * 10 cal/mol. e Maximum deviations from the mean value are ?: 8 x l o T hg/ cm3. Concerning the @L data, i. 1 0 cal/mol represents the maximum deviation of the average @L values from a graphical best fit interpolation of @ p ~ against m plot. 11.5 13.1 14.1 13.3 11.6
4.1162 4.5024 4.0713 4.0593 4.0778 Mean AHIS = 3.4
Q, mcal
mol of caffeine x lo6
using extrapolatedz3values of CP. We obtain AHo= -3.17 f 0.10 kcal/mol, a t 25 O C . Elaboration of the same experimental data according to the procedure of Stoesser and Gill14 yields a nonlinear plot for the highest caffeine concentrations; nevertheless the values -3.4 kcal/mol and 9.4 m-l for the enthalpy and the constant of association (at 25 "C),respectively, could be extrapolated from the low concentration region. Concerning the solubility vs. temperature data it is important to point out that equilibrium was attained in each case between saturated solutions and solid hydrated caffeine. From these data (Table I) we derive d log msatd/d(l/ T) = 1.86 X lo3 K. Using for the differential heat of solution, AHDs,the relationship
and introducing the CP and dWdm values extrapolatedz3 from the data of Table I (msatd = 0.1092; asatd = 0.602, and dWdm = -1.32 f 0.05), one obtains AHDS = 3.90 f 0.05 kcal/mol. The same figure, within experimental error, can be independently derived from the calorimetric data of Table I. In fact a t 25 O C AHDS = AHIS + @L(m,,td) = 3.9 f 0.1 kcal/ mol where AHIS is the heat of solution at infinite dilution.24 Given the calorimetric heat of solution data for both the hydrate and anhydrous form of caffeine we finally evaluate (see Table I) AH = 1.6 f 0.1 kcal/mol for the enthalpy of the reaction a t 25 O C .
-
(caffeine-HzO)~ (caffeinels
+ (H20)~
(4
This heat should reflect the breaking of the H bond linking one caffeine molecule (through its Ng) with one water rnolecule in the hydrate,15 assuming that the crystal packing of caffeine is little affected by reaction 4. Unfortunately, however, no crystallographic data seem to be available for anhydrous caffeine in order to substantiate our qualitative conclusion. In conclusion, the whole set of results reported here may be interpreted in terms of association of caffeine in aqueous solution. Our osmotic coefficient data cannot, however, support the lack of trimer formation suggested by Guttman and Higuchi16 and are instead more compatible with the hypothesis of multiple association with identical equilibrium constants. Our heat of dilution values a t 25 OC are in good agreement with those reported by Gill et ale: but they do imply a nonlinear plot of @L against (CPL/m)1/2which, in turn, would suggest that A H z o # A H 3 O # . . . . As a matter of fact, using the limiting ( m 0) value of AHo = -3.4 kcal/ mol, one calculates K = 9.4 m-l in agreement with the osmotic coefficient result, while higher K values are obtained with increasing m. From our data it also appears that the entropy of caffeine dimerization (ca. -7 eu) may result solely from the contribution arising from the unit decrease in mole number upon association (ca. -8 eu, in water). Therefore the dimerization of caffeine is entirely enthalpy driven, and one can safely assume that the source of dimer stability resides in the optimization of van der Waals contacts between stacked molecules whose planes would be at a distance of 3.4 X cm apart with CZ symmetry.17
-
The Journal of Physical Chemistry, Vol. 80, No. 3, 1976
338
Comparison of our data with those from the literature indicates that the extent of dimer formation in aqueous caffeine solutions a t 25 "C is similar to that of purine3 and 6methylpurine4 but smaller than in the case of 6-dimethylaminopurine or of 6(methylamino)-9-methylpurinea t the same temperature. The latter compounds with the exclusion of purine are characterized, however, by entropies of dimerization which range from nearly twice to almost three times as large (in absolute value) as that found for caffeine. In particular, the entropy of dimerization amounts to -13 eu for 7-methylaminopurine, and eu for p ~ r i n e ,-14.6 ~ -22.4 eu for both 6-dimethylaminopurine and G(methy1amine)-9-methylpurine2in water at 25 "C. For these purine derivatives van der Waals interactions between stacked molecules would thus be only one of the factors governing a rather complicated mechanism of association in which water molecules are also heavily involved (but in a manner poorly understood as yet). In this context caffeine would represent a relatively simple case. (B) Interaction of Caffeine w i t h E t h i d i u m Bromide and with ilrea. The cationic dye ethidium bromide (EB) has received considerable attention because of the strong, selective interaction it can establish with DNA. Thermodynamic and spectroscopic aspects of EB dimerization in waterls and of the binding of EB by synthetic polyelectrolytesI8 and by DNA19 in dilute aqueous solution have recently been studied in this laboratory. Concerning the behavior of EB in water, it is useful to recall that (1) EB is rat,her reluctant to dimerize ( K = 41 M-l at 25 "C) in comparison to most common dyes, (2) the EB absorption spectrum is red-shifted upon dimerization, an infrequent phenomenon for dye aggregation, and (3) the E B dimerization ent,halpy (AH = -7.6 kcal/mol) differs by no more than about 15% from the average value of -6.5 kcal/mol calculated on the basis of enthalpy of dimerization data for other different dyes.18 Pursuant to our interest in a quantitative description of the capability of caffeine to form 1:l complexes with aromatic compounds in water, it was considered worthwhile to begin with EB, a well-characterized and interesting case itself. Investigating the spectral properties of EB in aqueous caffeine solutions we find that the dye spectrum is hypochromic and red-shifted with respect to that of the free mo480 nm; €480 5270). For example, nomeric dye in water (A, in 0.1 M caffeine the spectrum of E B (1.8 X lo-* M) exhibits a maximum absorption a t 510 nm with an apparent molar extinction coefficient of €510 4380. The hypochromic effect and red shift are qualitatively similar to those found for E B dimerization as well as for E B binding by polycarboxylates and by DNA.18 Assuming that a 1:l complex is formed between caffeine and EB, an equilibrium constant KEB= 126 M-l has been calculated at 25 "C on the basis of series of spectral measurements for different EB/caffeine molar concentration ratios,20 using the extrapolated value of €480 4050 for the EB-caffeine 1:l complex. In these experiments self-dimerization of caffeine has been neglected, M. since its concentration was always below Calorimetric measurements, carried out as explained in the Experimental Section, have yielded for the enthalpy of EB-caffeine complexation the value AHEB = -5.4 f 0.1 kcal/moI of E B bound at 25 "C (Table 11). This figure is just the numerical average of the enthalpies of self-dimerization of caffeine and of EB. The total entropy of complexation then results, ASEB = -8.6 eu, corresponding to the The Journal of Physical Chemistry, Vol. 80, No. 3, 1976
A. Cesaro, E. Russo, and V. Crescenzi
TABLE 11: Heats of Mixing Aqueous Caffeine (caf) with Aqueous Ethidium Bromide (EB)a at 25 "C Mrnfx 103
Q, call&
A H , kcal/mol -
~
52.96 38.62 23.19 12.54 1.97
-28.8
X
x 10-6 -18.3 X -15.2 X -4.6 X ~ H " E=B-5.4 i 0.1 kcal/mol ~ M E =B 4.84 x M. b Flow rate 9.58 x -21.0
-5.41 -5.28 -5.42 -5.49 -5.34 l./s.
cratic contribution well within estimated uncertainties in the K E B and AHEBvalues. Energetically favorable r-electron cloud overlap between E B and caffeine would then be the major source of stability for their 1:l (stacked) complex in aqueous solution. A similar explanation was afforded for the strong complexation of the dye acriding orange with sarcosine anhydride in dilute aqueous solution once the possible relevance of hydrophobic interactions was ruled out.8 Finally, it may be worth reporting that we have also found by means of measurements of the heat of solution of caffeine (2-6 X lov3 M final molarity) in aqueous urea (0.2-5 M) that such calorimetric data are consistent with the formation of a caffeine-urea 1:ladduct in solution. The heat of solution data do, in fact, obey the simple equation AH,,,/M = KAHuree- KAH,,,, which follows immediately from the mass law, and where AH,,,is the "excess" heat of solution of caffeine in a urea solution of molarity M . The best fit parameters are K = 0.50 f 0.05 M-l and AHurea= -3.60 f 0.10 kcal/mol. Lacking conclusive spectroscopic evidence, interpretation of these data is still open to question, as caffeine and urea might associate either by hydrogen bridges or by a sort of stacking mechanism.21 More work is being carried out with urea and urea derivatives to elucidate this point.
Acknowledgment. This work has been sponsored by the Italian Consiglio Nazionale delle Richerche, Rome. References and Notes (1)(a) P. 0.P. Ts'o and S. I. Chan, J. Am. Chem. Soc.. 85,1289 (1963):(b) P. 0. P. Ts'o, I. S. Malvin. and A. C. Olson, ibid., 86,4176 (1964). (2)R. Bretz, A. Lusting, and G. Schwarz, Biophys. Chem., 1, 237 (1974). (3)S.J. Gill, M. Downing, and G. F. Sheats, 8iochemistry, 6,272 (1967). (4)M. G. Marenchic and J. M. Sturtevant, J. Phys. Chem., 77, 544 (1973). (5) T. Higuchi and D. A. Zuck, J. Am. Pharm. Assoc., 42, 132,138 (1953). (6)E. Shefter, J. Pharm. Sci., 57, 1I63 (1968). (7)F. Quadrifoglio, V. Crescenzl, A. Cesaro, and F. Delben, J. Phys. Chem., 75, 3633 (1971). (8)V. Crescenzi, A. Cesaro, and E. Russo, lnt. J. Peptide Protein Res., 5, 427 (1973). (9)N. H. Ishler, T. P. Finucane, and E. Borber, Anal. Chem., 20, 1162 ( 1940). (10)F. J. C. Rossotti and H. Rossotti, "The Determination of Stability Constant", McGraw-Hill, New York, N.Y., 1961,p 190. (1 1) "Handbook of Chemistry and Physics", 47th ed, The Chemical Rubber Co.. Cleveland, Ohio, 1966. (12)F. Quadrifoglio and V. Crescenzi, Biophys. Chem., 2, 64 (1974). (13)D. D. Kasarda, Biochim. Biophys. Acta, 217, 535 (1970). (14)P. R. Stoesser and S. J. Gill, J. Phys. Chem., 71, 564 (1967). (15) D. J. Sutor, Acta Crysfdbgr., 11, 453 (1958). (16)D. Guttman and T. Higuchi, J. Am. Pharm. Assoc.. 46,4 (1957). (17)J. N. Klkkert, G. R. Kelly, and T. Kurucsev, 8iopo/ymers, 12, 1459 ( 1973). (18)V. Crescenzi and F. Quadrifoglio, Eur. Polyrn. J., 10, 329 (1974). (19)F. Quadrlfoglio, V. Crescenzi, and V.,Giancotti, Biophys. Chem., 1, 319 (1974). (20)R. W.Ramette, E. A. Dratz, and P. W.Kelly, J. Phys. Chem., 66,527 (1962). (21)W.P. Jenks, "Catalysis in Chemistry and Enzymology", McGrawHill, New York, N.Y., 1969,p 461.
Communications to the Editor
339
(22) J. H. Stern, J. A. Devore, S. L. Hansen, and 0. Yaruz, J. Phys. Chern., 78, 1923 (1974). (23) The values of Q and dQldm used in eq 2 and 3 have been extrapolated from the Q data at 29.8 and 35 ' C . The estimated errors for this procedure are reported.
(24) Recently a value of 3.60 & 0.06 kcal/mol has been reported for the heat of solution of caffeine In water." This value was reported to be independent of concentration up to 0.01 m.'* These findings are not in agreement with evidence reported in the present paper as well as with previous data given by Gill et al?
COMMUNICATIONS TO THE EDITOR
Electrometric Study on the Chloride Ion Inclusion into the Poly(a,L-glutamic Acid)-Acridine Orange Complex
Sir: Neutral salts such as potassium chloride are known to affect the absorption spectra, optical rotatory properties, and kinetic aspects of the poly(@-glutamic acid) (PLGA)acridine orange (AO) c ~ m p l e x . ~Sat0 - - ~ et al. interpreted these phenomena by the assumption that (i) the chloride ion (Cl-) associates with the A 0 dimer as follow^:^^^^^ 2AO+
+ C1-
f
(AO)ZCl+
(1)
and then (ii) this (A0)2C1+ is bound to the ionized carboxyl groups of PLGA. On the other hand, Schwarz assumed that the A 0 dimer is present as a divalent ~ a t i o n . ~ , ~ , ' If AO+ associates with C1- according to (l),the activity of C1- would decrease. Thus the validity of their mechanisms could be clearly elucidated from measuring the C1activity in the present ~ y s t e m . This ~ , ~ communication reports the results of emf measurements of the C1- activity in the PLGA-A0 system. Acridine orange was used as the hydrochloride salt (AO. HC1). PLGA was used as the sodium salt. The following electrochemical cell was used:
-5
-4
-3
-2
-1
L o g a r i t h m of total c h l o r i d e i o n concentrat i o n Figure 1. Reversible emf of the cell as a function of the logarithm of
the total chloride concentration added as AO.HCI or NaCI: ( 0 )NaCl only; (0)AO-HCI.
HgClz(s)/Hd satd KClIjAO-HCl(aq) f PLGA(aq)lAg/AgCl T h e emf was measured with a Radiometer p H meter 4d. All solutions were kept a t 298 f 0.5 K. The p H value was adjusted a t p H 4.53 f 0.03 by 5 mMIO acetic acid-sodium acetate buffer. In Figure 1 the emf values are plotted against the logarithm of the total amounts of C1- added as AOnHCl. It should be noted that the emf values linearly decrease with log[AO.HCl] a t all concentrations between 1 X lom5and 5 X M.l0 In this concentration region the dimer forma~ , ~ these data, it is tion of A 0 takes place m a r k e d l ~ . From concluded that both AO+ and (A0)z2+ negligibly associate with C1-. This conclusion agrees with what was obtained by Padday on other dye^.^,^ Figure 2 exhibits the emf values of the following systems: (a) NaCl in water (pH 4.531, (b) PLGA-NaCl in water (pH 4.53), (c) PLGA-AO-HCl in water (pH 4.53), where in system c the glutamyl residue to dye ratio (R/D) was kept constant at R/D = 10. The measurements were performed in
-5
-4
-3
Logarithm of total chloride i o n c o n c e n t r a t i o n Figure 2. Reversible emf of the cell as a function of the logarithm of the total chloride concentration added as AO-HCI or NaCI: ( 0 )(a) NaCl only; (A)NaCl immediately after measuring PLGA-NaCI; (0)(b)
PLGA-NaCI at pH 4.53; (A)(c) PLGA-AO-HCI at pH 4.53; R/D = 10.
the order of a, b, and c. In the presence of the PLGA, the emf values shift down appreciably below [AO-HCl] = M. Washing the Ag-AgC1 electrode with water a whole day after the measurements for the PLGA-NaC1 system caused The Journal of Physical Chemistry, Vo/. 80, No. 3, 1976
