Michael G . Marenchic and Julian M. Sturtevant
544
0. kilocal/mole)
0.1
0.2
0.3 0 . 4 0.6
0.6 0.
r
a8 0.13
u
dl
012
013
0.'4
015
tentiornetric techniques, in particular experiments from which reliable information on polyions hydration and on the temperature dependence of polyions dimensions and charge densities for different CY values might be derived, diould prove valuable in this context.
0.6
00
0:s
Oi9
ll0
Acknowledgments. This work has been carried out with financial support of the Italian Consiglio Nazionale delle Ricerche. The authors wish to express their gratitude to Professor H. Morawetz for much stimulating advice and many helpful discussions.
stigatisn of the Association of ses in Aqueous Media'
. Marenchic2 and dulian M. Sturtevant" Depdrrment of Chemistry. Yaie University, New Haven. Connecticut 06520 (Received September 25. 19721
PubiicaCion costs assisfed b y the National lnsfifutes of Health and the Nationai Science foundation
The thermodynamics of association of 6-dimethylaminopurine in aqueous media have been investigated by flow microcalorimetry. The presence of charges on the solute molecules a t high and low pH decreases the association, but the changes in enthalpy and entropy are both of sign opposite to that expected on the basis of simple electrostatic considerations. Results obtained with added organic solvents suggest contributions to the binding free energy resulting both from hydrophobic forces (W. Kauzmann, Adunn. Prolein Chern., 14, 1 (1959)) and from surface forces ( 0 . Sinanoglu and S. Abdulnur, Fed Proc., Fed. Rnier Soc Erp. Biol., 24, 5 (1965)). It is not, however, possible to give a fully satisfactory analysis of the data in terms of presently accepted views of solute-solvent interactions. Preliminary data have been obtakrred for five additional purine derivatives.
~
~ t ~ ~ ~ u ~ t i ~ ~ The associai;ion in aqueous solution of pyrimidine and purine bases. ;Ind nucleosides and nucleotides derived from them, has received much study because of the bearing of this association on the interactions between the The Journal of Phy.jical Chemistry, Voi. 77, No. 4, 7973
stacked wrimidine and Durine rings in helical nucleic acids. Current information on this association has been ~ and his have well summarized by T s ' o . Gill -
1
Y
(1) From the Ph D thesis of M G Marenchic
Purine Bases in Aqueoins Media
545
shown that the thermodynamics of the association can be conveniently investi.gated by flow microcalorimetry. We have employed this method in further work on the association of certain purine derivatives, with emphasis on a study of the effects of pH, solvent composition, and nature and concentration of anionic species on the association of 6-dimethylarn.inopurine. Experimental Seetion The bases and nucleosides used were obtained from various sources, and were in most cases used without further purification since pap’er chromatography with n-butyl alcohol-water (86:143 gave only one spot. In some cases it was ascertained thab recrystallization led to no observable variation of dilution properties. Heats of dissociation were measured in a flow modification7 of the Beckman Model 190B microcalorimeter. Analysis of the resulting data is based on a treatment developed by Stoesser and Gill.4 This treatment assumes, first, that all solute species behave ideally; second, that polymerization of the monomeric solute proceeds to a n indefinite degree; and, third, that all polymerization steps are characterized by the same dissociation constant K and dissociation enthalpy AH. Stoesser and Gill employed molality as concentration unit, but if the first assumption is valid, there will be no volume changes on dilution and molarity can equally well be employed. It was shown that a t the concentration levels used in our work any volume changes on dilution were indeed negligible. It is found on this basis that the heat of dilution AQi from a .total monomer molarity of Mi to zero molarity is
AQi
AET
- (~~AH)”z~AQj/Mi)”z
(1)
From this expression it follows that for a finite dilution experiment from initial molarity Mi to final molarity Mf
AQ AQi .- A H .f (KAH>”’[(AQi - AQ)/M,]”’ ( 2 ) where AQi is the heat of dilution from Mi to zero concentration. As pointed out by Gill, et u L . , ~ if osmotic coefficient data are available, an estimate of AN is obtained by combination of osmotic arid calorimetric data, in effect using the former type of dahs for the evaluation of K , which may differ numerica.llj7 from the purely calorimetric estimate. For example, for B-methylpurine in water a t 25”, Stoesser and Gi1P reported K = 0.116 $1 0.030 m and AH = 5.6 A: 0.2 kcal rno1-I from calorimetric data alone, and K = 0.149 m and A H = 6.0 $1 0.4 kcal mol-1 on the basis of combined osmotic and calorimetric data. We have employed only calorimetric data in the evaluation of K and AN since 0smoti.c data are not available for most of the systems studied. In each set of experiments a stock solution was diluted to a number of :bw@r‘concentrations. Trial values of AQi were assumed until SI value was found which gave the best least-squares fit of the data to a straight line plot of A Q us. zG [(AQi -. Agj/Mfgliz, Values’of K and AH were then obtained from the intercept and slope of the straight line. Since more than one initial solution was usually einployed, each corresponding set of data points was treated as outlined above, anci then combined with the curve of highest initial concentration by addition to the experimental dilution heatis of a constant heat quantity equal to the weighted mean difference between the smoothed dilution curve for the lower initial con.centration and that for
x
A
0
I
I
0
12
IO
I
I
(y-)”2, 14
16
KCAL1’2
LiTERl’‘
18
20
MOLE-’
Figure 1. Data for two series of experiments on the heat of dilution of 6-dimethylaminopurine in water at 25” plotted to show adherence to eq 2. AQ is the observed heat of dilution per mole of monomer from a fixed initial molarity, 11/11, to a series of final molarities, M f , and AQi is a value for t h e heat of dilution from M i to zero molarity selected as outlined in the text to give optimum fit of the data to a straight line plot: 8 ,series I , M i = 0.0457 M, A Q = AQobsd; 0 , series 2, Mi = 0.0227 M , A Q = AQ,b,d i1.090 kcal mol-’; AQi = 5.086 kcal r~iol-’.The line is drawn for K = 0.0162 M, AH = 9.15 kcal mal-’.
the highest initial concentration. The combined data were then again treated by least squares to obtain refined values for K and AN. A typical group of experimental data for the dilution of 6-dimethylaminopurine in water is presented in Figure 1, and is seen to fit well the assumed polymerization scheme. It should be pointed out, however, that a simple dimerization would also lead to a linear dependence of A Q on (AQ/Mo)l/Z“ In the present case, indefinite polymerization seems more likely than dimerization.s In many cases the observable heats were too small, because of low solubility and/or small heat of dissociation, to enable a satisfactory selection of AQi to be made as outlined above. In these cases the observed heats of dilution were least squared to the equation
AQ
=
AQi
- AN(1 + ( K / 2 M , ) { 1 - [I -i- (4Mf/K)]”’]) (3)
which is based on the same model for the polymerization, for various assumed values of K. It was found that AH could then be expressed to better than a570accuracy as a function of K , over a full range of probable values for K , by the empirical equation
AN = A
+ ( B / K )+ GK
t 4)
where A , B, and C are constants. Thus, if the value for K for any of these cases is determined by other means such as osmometry, the value of AH, consistent with our caloriPredoctoral Trainee, USPHS GM No. 00748. from the National lnstitutes of Health. ~~
P. 0. P. Ts’o in “Fine Structure of Proteins and Nucieic Acids,” G.
D. Fasman and S. N. Tirnasheff, Ed., Marcel Dekker, New York, N. Y., 1970, p49. R. Stoesser and S. J. Gill, J . Phys. Chem.. 71, 564 (1966). S . J. Gill, M. Downing, and G. F. Sheats, Biochemistry. 6, 272 11967). E. L. Farquhar, M. Downing, and S. J. Gill, Biochemistry. 1, 1224 (1968). J. M. Sturtevant, Fractions. No. 1 (1969). P. 0. P. Ts’o, I . S. Melvin, and A. C . Olson, J. Amer. Chem. Soc., 85, 1289 (1963). The Journal of Physical Chemistry, Vol. 77, No. 4, 1973
Michael G. Marenchic and Julian M. Sturtevant
54
TABLE I: ~ ~ e r m o Parameters ~ y ~ ~ for ~ a~Single c Step in the Dissociation of Polymers of 6-Methylpurine in Water and of 6-Bimethylaminopurinein Various Solvents at 25" AS0, 10' K, M
Solvenr
AH,
kcal mol-'
AGO, kcal mal-'
cal deg-' mol-'
6-Methylpurine 12.9 i 1.6
5.56 f 0.05
1.21 f 0.07
14.6 f 0.1
f 0.01 f 0.01 f 0.01 f 0.01
22.5 f 0.3 21.9 rtr 0.2 22.5 f 0.3 22.8 f 0.6
& 0.02 f 0.02
23.3 f 0.6
6-Dimethylaminopurine HzO 0.01 M NaCl 0.1 ,w NaCI I .o rW NaCE 0.1 iW I'.la2SOn 1 .O iW Na2S04 0.1 M Na'rcxa 1.0 M Nal'CA 0.1 A4 NaQti 0.64 M dioxane 1.2 ,I4 c:I-l!&I\I 4.8 M c;H3GN
1.62 f 0.03 1.47 f 0.02 1.72 $1 0.04 1.71 0.04 1.72 & 0.05 1.07 f 0.04
*
2.36 f 0.06 6.92 f 0.02 27.8 f 1.2
3.80 f 0.05 3.81 f 0.07 21.6 f 0.1
9.7 9.02 9.1 9.2 9.3 9.5 9.1 7.3 5.2 9.4 9.8 8.0
fO.1 f 0.08 f 0.1 f 0.2 f 0.2 f 0.2 f 0.1 f 0.1 f 0.1 0.1
2.44 2.50 2.40 2.41 2.40 2.68 2.22
*
0.76 f 0.03 1.93 a 0.01
f 0.1 f 0.1
1.93 f 0.Ol 0.90 f 0.01
f 0.02 1.58 f 0.01
22.9 f 0 . 8 22.9 f 0.3 19.2 f 0.3
15.1 I 0.2 25.1 f0.2 26.2 h 0.3 23.7 f 0.3
a Sodium trichloroacetate
metric measurements, can be calculated with an accuracy largely determined by the accuracy of the measured equilibrium constant.
ssion The values of K and A H , and derived values for AGO and AS", are given in Table I for all systems for which a convergent calculation was possible. The values for 6methylpurine agree well with those reported by Stoesser and Gill* on the basis of calorimetric data alone, namely, M = 0.116 f 0.030 m and AH = 5.6 0.2 kcal mol-1. In all cases the standard entropy of dissociation is posi. tive. The cnatic contribution,g arising from the unit increase in mole number a t each dissociation step, lies between i-8.0 cal dpg-l mol-' in water and $7.7 cal deg-l mol-I in 4.8 is1 acetonitrile. Gill, et al ,s reported for the dissociation of polymers of purine itself, the quantities AH = 4.2 I 0.2 kcal mol-l, AGa = 0.44 kcal mol 1, and AS' = 12.6 cal deg-I mol-l. We thus see a steady trend in all three thermodynamic parameters toward more positive values in the series purine, 6-methylpurine, and 6-dimethylaminopurine. It is usually cons-dered3 that interactions of hydrophobic origin are of importance (n the association of purine derivatives in €320, and some part of the increased association in this series is presumably to be attributed to such interactions. However, that this cannot be the full story is indicated by and ' ASo change in the direction opthe fact that both ~!h posite to thst expected on the basis of increased orientation of water molecules on dissociation of the bases.1° After deduction of the cratic contributions from the entropies, and -5 to - 10 cal deg-I mol-I from the value for 6-di~nethylaniirropurineto allow for additional structurlng of water arc~undthe dimethylamino group, the entropy of dissociation for the substituted purine is seen to be fourfive times as large as that for the parent substance. The origin of this significant difference cannot be specified at the present time. A part of the iccrcased association of B-dimethylaminopurine as clampared with purine can probably be attributed to the increased polarizability of the a-electron system ofthe fermer substance."~l The Journai of Physical Chemistry, Vol. 77, No. 4, 7973
The data in Table I show that the dissociation process in the case of 6-dimethylaminopurine is not significantly affected by sodium chloride up to a concentration of 1 M , as expected for the dissociation of a substance into two uncharged products. The effects produced by sodium sulfate and sodium trichloroacetate can be rationalized if it is assumed that the stacked bases are unaffected by the added salts while the monomeric base is salted out by sodium sulfate and salted in by sodium trichloroacetate.12 On this basis
where f, and f i o are respectively the moiar activity coefficients in salt solution and in water, K, is the salting out constant, and C, is the salt concentration in moles per liter. The data for molar salt solutions give K, = 0.41 for sodium sulfate and K , = -1.5 for sodium trichloroacetate. From measurements of the solubility of adenine, Robinson and Grant12 found values of 0.24 and -0.31 for K , for these two salts. Robinson and Grant suggest that salting in results from direct i1-T 3raction of the salt anion with the base monomer. The decreased enthalpy and entropy of dissociation are consistent with this mechanism. In 0.1 M NaQH (pH -13) 6-climethylaminopurine carries a single negative charge, so that the tendency toward association is decreased by electrostatic repulsions. However, the changes observed for .IH and AS" are both of sign opposite to that expected on the basis of the simplest view of charged spheres interacting in a dielectric continuum, according to which
where BAS" = AS"(uncharged) - AS"(charged) and D is the dielectric constant of the continuum. The value of d In (9) R. W. Gurney, "Ionic Processes in Solution." McGvaw-Hill, New York, N. Y . , 1953, p 90. (IO) W. Kauzmann,Advan. Protein Chem., 14, 1 (1959). (11) H. DeVoeand I. Tinoco, Jr., J. Mol. Bioi., 4, 500 (1962). (12) D.R . Robinson and M. E. Grant, J. Biol. Chem.. 241, 4030 (1966)
Purine Bases in Aqueous Media
547
TABLE III: Empricielly Dellermined Constants for Eq 4 at 25"
--
l _ l _ l
System _ l s
4
A, cal mol - I
cal mol-2 I.
C, cal I.-'
Lower limit of K.a M
" I _
Adenosine-H$3 6-Chloropurine-H2(3 6-Cyanopurine-1-120
1050 1440
0.55 2.0
670
6-Rimethylanninl?pbirine-HCl (pH 2)
41 0 1295
1.6 1.5 0.80 0.65
2'-Deoxyadenosine-H20 6-~ethoxypurino-~l~Q a Minimum value of
1065
38,500 20,980 3,375 2,800 36,750 37,950
-
0.002 0.002 0.003 0.004
0.002 0.002
K lor 3~5% accuracy in AH
TABLE 1111: Thermodynamic Parameters Calculated Using the Constants of Table II with Indicateda Assumed Values
102K. M
Base
-
-
_ I -
Adenosine 6'-Chloropuriw 6-Chloropurine 6-Cyanopurine 6-Cyanopurine 2'-Deoxyadennsirie 2'-Deoxyadenasirie 6-Rimethyiarnjnoi)urine(pH 2) 6-Methoxypuririe 6-Methoxypurine
AH, kcal mol-' 9.6 5.4 6.8 4.4 5.7
22b 19
26 110 150 13c 21c 140
13 16
AGO. kcal mol -l
AS", kcai deg-'
mol-'
0.89
29
1.02
6.1
1.19
15 20 15 20 16
9.2 4.3 5.8 7.1
0.92
28
0.81
-0.06 -0.24
-0.2 1.23 1.09
15 15 20
*
a Assumed values are given in boldface type. Reference 17; K actually has the units moles per kg of water. Reference 17; K actually has the units moles per kg of water: the! lower value of K applies to concentrations below about 0.02 m, the higher value to higher concentrations.
0
AG*
I VOL % ORQANIC COMPONENT
I
I
I
70 60 50 40 SURFACE TENSION D Y N E S CM-'
Figure 2. Influences of the addition of organic components on the therrvlodynamic parameters of dissociation reactions at 25". Data of Crothers and Ratner14 (0) on the dissociation of the complex of actinomycin and deoxyguanosine in methanol-water mixtures; present data on the dissociation of 6-dimethylaminopurine polymers in ace!tonitrile-water mixtures (0) and in 0.64 M dioxane ( A ) . AH, TAS", and AG" are plotted as functions of the volume fraction of organic component in A, and of the solvent surface tension (interpolated to 25' (0); at 20' (0 and A ) ) in B.
Did In T for water at 25" is -1.36. Evidently, and not surprisingly, the effective dielectric constant for this system is very different from that of bulk water. In this connection it is intercsting that a communication of Andrews and Haydon quoted by Howarth, et al.,13 gives TASIAG = 4-0.45 for films of glycerol monoleate plus n-decane. If it is assumed that the entropy of dissociation a t low pH, where 6-dimethylaminopurine carries a p0sitiu.e charge, is the same a s that when it is negatively charged,
the data of Table I1 give K = 1.4 &I, AH = 4.3 kcal mol-l, AG" = -0.2 kcal mol-I, and A S o = +15 cal deg-l mol-1 (Table 111). However, even if this assumption is correct, this calculation represents a long extrapolation from the very small heats of dilution observed with solutions approximately 0.05 M in base. 6-Dimethylaminopurine is less associated in solutions containing dioxane or acetonitrile than in pure water. The initial increase in the enthalpy and entropy of dissociation followed by a decrease in both quantities as the concentration of acetonitrile is increased (Table I and Figure 2A) is similar to that observed by Crothers and Ratner14 for the dissociation of the complex of actinomycin with deoxyguanosine in methanol water mixtures. These authors interpreted the initial rise in the enthalpy and entropy of dissociation as indicative of a significant hydrophobiclo contribution to the association reaction. Sinanogld5.16 has proposed as an important source of binding energy for large molecules the surface energy required to form cavities around the molecules in the system. In an association reaction, the cavities around the two reactant species have a larger total surface area than the single cavity around the product, so that surface forces literally squeeze the reactants together. This picture leads to enthalpy and entropy changes of opposite sign to those expected to result from hydrophobic interactions, and the decrease of enthalpy and entropy observed in 4.8 M acetonitrile may thus be an indication that this mecha(13) J. V. Howarth, R. 0.Keynes, and J. M. Ritchie, J . Physiol.. 194, 745 (1968). (14) D. M. Crothersand D. I. Ratner, Bipchemistry, 7, 1823 (1968). (15) 0. Sinanoglu and,S. Abduinur, Federation Proc.. 24, 5 (1965). (16) 0.Sinanoglu in Molecular Associations in Biology," 8. Pullman, Ed., Academic Press, New York, N. Y., 1968, p 427.
The Journal of Physical Chemistry, \lo/. 77, No. 4, 1973
P. Bordewijk, M. Kunst, and A. Rip
548
nism is also operative. Additional support for this view may be drawn from the fact that our data and those of Crothers aind Ratner14 exhibit greater qualitative similarity when plotted as functions of surface tension (Figure 2B) than as functions of volume fraction (Figure 2Aj. Broom, et d , I 7 have reported equilibrium data for the association of adenosine ( K = 0.22 M ) and 2'-deoxyadenosine ( K = 0.13--0.21 M ) in water at 25". Using these values and &he constants in Table 11, we obtain the thermodynamic parameters given in Table 111. Preliminary values for 6-chloropurine, 6-cyanopurine, and 6methoxypurine me also given in Table 111, based on assignment of the values 15 and 20 cal deg-1 mol-1 for the entropy of dissociation. The equilibrium constant calculated in this way for cyanopurine is surprisingly high; as-
sumption of a lower value for the entropy lowers the equilibrium constant, for example, A 9 = 0 gives K = 0.37 M . It is to be emphasized that we are still very far from being able to give a satisfactory interpretation of the thermodynamics of association of molecules such as the purine bases in qualitative terms, to say nothing of quantitative explanations.
Acknowledgments. This research was aided by grants from the National Science Foundation (GB 23545) and the National Institutes of Health of the United States Public Health Service (GM 04725). (17) A. D. Broom, M. P. Schweizer, and P. 0. P. Ts'o, J. Amer. Chem. SOC..89, 3612 (1967).
e Association of Heptanol-1 in Carbon Tetrachloride from Static
ordewijk," M. Kunst, and A. Rip Gorlaeus Laboratories. Department of Physical Chemistry I / . University of Leiden. Leiden. The Netheriands (Received July 12. 1972)
Static dielectric constants have been determined of solutions of heptanol-1 in carbon tetrachloride in the concentration range from 0.1 to 3 mol 70 at temperatures from 25 to 45". Possible interpretations are a monomer-dimer-trimer equilibrium (heats of association for dimer and trimer 9 and 17 kcal/mol, respectively) and a monomer-dimer-tetramer equilibrium (heats of association for dimer and tetramer 9 and 26 kcal/mol, respectively). The measurements confirm Ibbitson and Moore's conclusion that the dimer is linear.
Introduction In spite of nnan:y investigations with different techniques, still no general agreement has been obtained concerning the structure of the multimers formed due to the self-association of the mono-alcohols.1-9 An important type of technique is the measurement of a physical quantity that is built up from contributions of the various types of multimers as a function of the formal concentration in a nonpolar solvent. The physical quantity may be the infrared absorption at some deliberately chosen frequency,l0-13 the chemical shift in nuclear magnetic resonance,14-17 the vapor p r e ~ s u r e ,or ~ the apparent value of the square of the molecular dipole moment.18-z0 Measurements of the lowering of the freezing point, of the partition coefficient with respect to water, and of the osmotic pressure fall under the same category. Also investigations combining different methods21-23 have been reported. The physical quantity under investigation need not depend on all. multimers present in all cases. For instance, the ir absorpltion may only depend on the concentration of open rnultimers, and both the vapor pressure in a nonvolatile solvent and the partition coefficient for organic solvents with respect to waterz2 depend only on the concenThe Journal ot Physical Chemistry, Vol. 77, No. 4, 797.3
tration of monomers. This simplifies the calculations, but not enough to get unambiguous results. Apart from incidental failures in the application of the measuring technique, and from unjustified a priori as-
(12) (13) (14) (15) (16)
S. N. Vinogradov and R. H. Linell, "Hydrogen Bonding," Van Nostrand-Reinhold, New York, N. y.,1971. P. Bordewijk, Thesis, University of Leiden, 1968. P. Bordewijk, F. Gransch, and C. ,I. F. Bbttcher, J. Phys. Chem., 73,3255 (1969). E. Tucker, S. B. Farnham, and S. D. Christian, J. Phys. Chem., 73, 3820 (1969). J. Crossley, Advan. Mol. Reiaxation Processes. 2,69 (1970). W. Dannhauser and A. F. Flueckinger, Phys. Chem. Liquids, 2, 37 (1970). J. Dos Santos, J. Biais, and P. Pineau, J. Chim. Phys.. 67, 814 (1970), J. Crossiey, L. Glasser, and C. P. Smyth, J. Chem. Phys.. 55, 2197 (1971). A. N. Fletcher, J. Phys. Chem., 75, 1808 (1971). A. Ens and F. C. Murray, Can. J. Chem.. 35,170 11957). W. C. Coburn. Jr.. and E. Grunwald. J. Amer. Chem. SOC.. 80. _. 1318 (1958) G Geiseler and E. Stockel, Spectrochrm Acta. 17, 1185 (1961). H Dunken and H Fritzsche, Spectrochrm Acta 20, 785 (1964). L K. Patterson and R. M Hammaker, Speciiach/m Acta 23A, 2333 11967). W. Storek and H. Kriegsmann, Ber. Bunsenges. Phys. Chem.. 72, 706 (1968). W. 0 . Dixon, J. Phys. Chem.. 74, 1396 (1970).
