TRANSPORT PROCESSES IN HYDROGEN BONDING SOLVENTS
3985
Transport Processes in Hydrogen Bonding Solvents. 111. The Conductance of Large Bolaform Ions in Water at 10 and 25" by T. L. Broadwaterl and D. Fennel1 Evans2 Department of Chemistry, Case Western Reserve University, Cleveland, Ohio
44106
(Received Mau 1, 1969)
Precise conductance measurements were carried out on Bu3N(CH2)8NBu3Br2,Pr&(CH2)gNPr3Br2, Et3N(CH.J4NEtaBrz, MeN(CH2CHz)SNMeBr2,and MeN(CH2CH2)aNMeIz in water a t 10 and 25' and on B U ~ N ( C H ~ ) ~ NBu3Br2in methanol in 10 and 25'. From a comparison of the limiting ionic mobilities and concentration dependence of the tetraalkylammonium ions and these large bolaform ions, it is concluded that cation-cation pairing does not make a significant contribution to the concentration dependence of RIN+ ions in dilute aqueous solutions. I n addition, when the concentration dependence for these salts in aqueous solution is compared to that observed in other hydrogen bonding solvents, many similarities are observed. It is likely that part of their peculiar behavior which has been attributed to the structural features of aqueous solution is a general property of hydrogen bonding solvents.
Introduction Water structure effects have been shown to be an important factor in determining the magnitude of limiting ionic proper tie^.^ However, the extent to which they are responsible for the anomalous concentration dependence in aqueous electrolyte solutions is not understood. I n the case of the tetraalkylammonium salts, several explanations have been advanced based on ionic aggregation induced or enforced by water struct ~ r e . ~ It- ~has been suggested that cation-cation pairing could be responsible for these effects and several recent papers ha,ve lent support to this argument. Wen and Nara studied the volume change associated with mixing of solutions of R4N halides and alkali metal halides.1° Analysis of the data using Friedman's ionic solution theory showed that when the R4S ions were large, cation-cation interaction was found to predominate. Wen, Nara, and Wood calculated the volume change upon dimerization of Bu4N+ and found it to be - 11 cc/mol.ll Due to its labile nature, the isolation and direct study of a cation-cation pair is impossible. This difficulty can be circumvented, at least in part, by the study of appropriate model compounds. In a previous paper the preparation and partial molar volume measurements on B U ~ N - ( C H ~ ) ~ S B U ~a (large B ~ ) bolaform ~, electrolyte, were reported.'2 This compound can be likened to two tetraalkylammonium ions held in close proximity. The change of dz with Concentration for DiBuBrz was similar to that observed for Bu4NBr,with 42 decreasing to a minimum d2(min) = DiBu.55Hz0 and increasing with concentration. The ratio 72'(DiBu2+)/Vz'(BuzN+) was found to be 1.91 and the volume change upon dimerization was estimated to be -11.5 cm3/mol. If the value of the partial molar volume for DiBuBrz is corrected by a factor of 2, then
the curve of ~$2 for the bolaform salt may be almost superimposed on that for BudNBr. I n an effort to explore further the implications of cation-cation pairing in aqueous solutions of the tetraalliylammonium ions, we have studied the conductance of a number of bolaform ions of the general These form (C,Hz,+&N +(CdLJN+(C,H2, +1)3. bolaform cations will be abbreviated as DiBu2+,DiPr2+, and DiEtz+for the butyl, propyl, and ethyl derivatives, respectively. As suggested by Fuoss,13 DMD2+will be used for the dimethyl derivative; it differs from the others in that two charged nitrogens are paired through three side chains. Experimental Section The dibromide salts of DiBu2+ (mp 123-125°),12 DiPr2+(mp 227-229'), and DiEth2+ (mp 268-269.5') were made by refluxing an excess of trialkylamine with the appropriate alkyldibromide in methanol or ethanol. Recrystallizations were from methanol-ether mixtures. DMDBrz and DhSDIz were made by reaction of 1,4-diazocyclo(2,2,2)octane in methanol with the (1) For further details, cf. T. L. Broadwater, Ph.D. Thesis, Case Western Reserve University, Sept 1968. (2) To whom all correspondence should be addressed. (3) R. L. Kay and D. F. Evans, J . P h y s . Chem., 70,2325 (1966). (4) S. Lindenbaum and G. E. Boyd, ibid., 68,911 (1964). ( 5 ) W. Y .Wen and S. Saito, ibid., 68,2639 (1964). (6) F. Franks and H. T. Smith, ibid., 68,3581 (1964). (7) F. Franks and H. T. Smith, Trans. Faraday SOC.,63, 2586 (1967). (8) B. E. Conway, R. E. Verrall, and J. E. Desnoyers, ibid.,62, 2738 (1966). (9) D. F. Evans and R. L. Kay, J . P h y s . Chem., 70,366 (1966). (10) W. Y . Wen and K. Nara, 71,3907 (1967); 72,1137 (1968). (11) W. Y . Wen, K. Nara, and R. H. Wood, ibid., 72, 3048 (1968). (12) T. L. Broadwater and D. F. Evans, ibid., 73, 164 (1969). (13) J, E. Lind, Jr., and R. M.Fuoss, ibid.,66, 1749 (1962).
V o l u m e 78, N u m b e r 11 November 1969
3986
T. L. BROADWATER AND D. FENNELL EVANS
Table I : Equivalent Conductances of the Diammonium Salts 104~
104~
A
HnO, 25' DiMDBrz l O l O 7 ~= 1.49 A = 0.197 2.436 4.418 7.585 10.831 15.650 22,708 31.469
l
149.59 147.82 145.67 143.95 141.81 139.29 136.76
DiEthBrz O 7 ~= 1.84 A = 0.094 0.507 119.71 1.426 118.26 2.485 117.22 4.320 115.84 6.707 114.45 9.751 113.01 13.985 111.36 20.029 109.45
l O 7 ~=
DiPrBrz 1.16 A = 0.063
0,668 1.298 2.218 3.667 5.745 9.050 13.115 18,247 24.421
107.78 106.90 105.95 104.83 103.65 102.13 100.63 99.06 97.47
A
DMDIz 7 ~= 1.95 A = 0.221 (Run I) 5.457 145.43 10.339 142.51 15.755 140.04 20,797 138.13 25.264 136.65 28.998 135.52 33.171 134.35 37.505 133.24
1.82 A = 0.221 (Run 11) 1,499 149.08 2.740 147.66 4.378 146.20 6.047 144.97 8.168 143.63 11.943 141.64 22,623 137.42
1 0 7 ~=
l O 7 ~=
DiBuBrz 2.00 A = 0.065 (Run I )
1.177 3.687 6.729 11.067 16.036 21.186 27,014 31.964
102.28 100.22 98.60 96.89 95.35 94.02 92.70 91.72
DiBuBrz 2.29 A = 0.065 (Run 11) 0.721 102.79 1.784 101.45 3.348 100.22 5.042 99.27 7.407 98.13 9.790 97.17 13.280 95.97 17.392 94.77 23.161 93.33 28,739 92.12
l O 7 ~=
appropriate methyl halide. Both salts were purified by several recrystallizations from hot methanol. All of the salts were dried overnight in a vacuum oven at 60" and, T h e Journal of Physical Chemistry
104C
A
104~
A
HzO, 10' DMDBrz DiEthBrz 1 0 7= ~ 0.67 A = 0,197 l O 7 ~= 1.16 A = 0.094 0.853 108.67 0.675 83.95 2.074 106.98 1.357 83.25 3.528 105.92 2.277 82.53 5.337 104.87 4.368 81.41 7.643 103.77 6.485 80.51 10.918 102.52 9.240 79.56 15.564 101.09 14.565 78.07 25.880 98.56 20.80" 76.68 DiPrBrz 1 0 7= ~ 1.06 A 0.574 1.407 2.340 3.997 6.197 8.919 12,599 17.565 23.385
= 0.063
75.69 74.85 74.25 73.45 72.59 71.70 70.71 69.60 68.50
DiBuBra A = 0.065 (Run I) 0.506 72.37 1.296 71.57 2.507 70.83 4.874 69.79 7.650 68.86 10.866 67.98 14.213 67.20 18.941 66.25 24.141 65.35 29.563 64.52 35.226 63.75 41.066 63.01
DiBuBrz A = 0.065 1 0 7= ~ 0.73 (Run 11) 1.108 71.81 1.810 71.32 3.231 70.54 5.337 69.68 8.175 68.75 11.858 67.77 16.763 66.70 22.085 65.72
MeOH, 25" DiBuBrz l O 7 ~= 0.79 A = 0.224 0.181 102.98 0.372 100.87 0.570 99.29 0.788 97.87 0.999 96.69 1.279 95.33 1.675 93.64 2.076 92.17 2.938 89.53 3.815 87.31 4.994 84.85 6.327 82.54 7.895 80.26 10.018 77.70 13.347 74.53
MeOH, 10" DiBuBrz 1 0 8 ~= 0.98 A = 0.224 0,366 81.50 0.751 79.35 1.161 77.68 1.586 76.23 2.084 74.77 2.684 73.28 3.408 71.76 3.992 70.67 4.833 69.28
l O 7 ~= 1.36
except for DMDIZ, were analyzed by Galbraith Laboratories, Knoxville, Tenn. The preparation of the solvents, the electrical
3987
TRANSPORT PROCESSES IN HYDROGEN BONDING SOLVENTS Table 11: Conductance Parameters for the Diammonium Salts Solvent
Salt
DMDBrz
HzO
DMDIz
HzO
DiEthBrz
Hz0
DiPrBrz
HzO
DiBuBrz
HzO
MeOH
T,'C
10 25 25 25 10 25 10 25 10 10 25 25 10 25
Ao
equipment, conductance cells, cup dropping device, and general techniques were as previously described. ' l4
Results The measured equivalent conductances and corresponding molar concentrations are given in Table I. The specific conductance of the solvent and density constants, A , are also included in the table for each run. Density measurements were carried out on the most concentrated solution from each conductance run and the density was assumed to follow the relationship d = do Am, where do is the density of the solvent and m the concentration in moles of salt per kilogram of solution. The A value obtained at 25" was also used to calculate density changes at lo", a valid assumption for TAA salt s o l u t i ~ n s . ~ The data were analyzed with the Fuoss-Onsager conductance theory15 for nonassociated electrolytes in the form
+
A = A. - SC'I2
+ EC In C + J C
1.23 f 0.24 2.00 i: 0.06 1.87 i 0.03 1.59 Z!C 0.08 1.52 f 0.11 1.93 =t0.11 1.84 i 0.07 2.22 i 0.08 2.09 i 0.03 1.97 f 0.04 2.50 f 0.04 2.38 =!= 0.07 1.15 f 0.17 0.43 i 0.14
109.93 i 0.16 154.05 f 0.06 152.35 f 0.05 152.54 f 0.05 85.26 =!= 0.04 121.31 3~ 0.05 76.83 f 0.02 109.58 f 0.04 73.38 f 0.02 73.46 f 0.01 104.70 f 0.02 104.55 f 0.04 85.20 f 0 . 1 0 106.40 f 0.20
J
c
Li
0.29 0.09 0.06 0.09 0.07 0.09 0.04 0.08 0.04 0.02 0.04 0.07 0.11 0.15
53.78 75.83 75.37 75.56 29.11 43.09 20.68 31.36 17.23 17.31 26.48 26,33 39.70 49.95
The bolaform ions show the same temperature dependence as R4N+ ions, the larger ones being structure makers (C1Oz5> 1) and the smallest, DMD,f2 being a structure breaker. DiEth2+ shows little if any temperature dependence, a behavior similar to that observed for Et4N+ ion. The corresponding values for the R4N+ ions are 1.049, 1.033, 1.002, and 0.979 for the butyl through methyl derivative^.^ Table 111: Mobility Ratios of 10 and 25" Ion
T,OC
DiBu
25 10 25 10 25 10 25 10
DiPr DiEth DMD
Xa
clop6
M
1.040
1.36
1.034
1.36
1.011
1.34
0.962
1.70
+
26.40 17.27 31.36 20.68 43.09 29.11 75.56 53.78
(1)
The parameters A. and d are recorded in Table I1 and were calculated with a computer program written by Kay.16 The program was modified to correct the terms a, p, K , and b for the case of unsymmetrical electrolytes; these functions are defined elsewhere."16 The required values of the dielectric constant and viscosity of the solvents and the limiting conductance of the anions were taken from the l i t e r a t ~ r e ,except ~ Xo(Br-) = 45.5 in MeOH at 10".17 Table I1 also contains u, the standard deviation, and J . Because DiBuBrz appeared to be associated in methanol, only M were analyzed. those points below 2 X
Discussion Comparison of Limiting Ionic Conductance. Values of Xo+r) at 10 and 25" are recorded in Table I11 for each salt, along with the value of
In order to compare the mobilities of the tetraalkylammonium ions and the corresponding dimers on a similar basis, the ratio, M , defined as 114 =
Xo(DiR2+)
Xo(RrN+)
was calculated at 25") and the values are given in Table 111. Except for DMD2+,the mobilities of the bolaform ions are larger than those for the RdN+ ions by about 35%. This discrepancy can be largely accounted for by the difference in geometry of the bolaform ions as (14) (a) D. F. Evans, C. Zawoyski, and R. L.Kay, J.Phys. Chem.,69, 3878(1965); (b) J. L. Hawes and R. L. Kay, ibid., 69, 2420 (1965); (c) C. G. Swain and D. F. Evans, J . Amer. Chem. Soc., 88,383 (1966). (15) R. M. Fuoss and F. Accasoina, "Electrolytic Conductance," Interscience Publishers, Inc., F e w Yorlr, N. Y . ,1959. (16) R. L. Kay, J . Amer. Chem. Soc., 82,2099 (1960). (17) R. L. Kay, private communication.
Volume 73, Number 12
November 1969
T. L. BROADWATER AND D. FENNELL EVANS
3988 compared to the corresponding symmetrical R4N+ ions. Using the formula of Perrinl8
1 6rNA
RJSX were mixed with (a/2)C parts of (R4N)2X2, assuming that they are both strong electrolytes. Denoting the conductance of RdNX, the monomer, by A m and (R4N)ZXZ, the dimer, by AD, then
p = -
A s o ~ n=
where p is the mobility and A is a geometrical coefficient, the mobility of the bolaform ions can be predicted. If the bolaform ions are considered to be rigid ellipsoids, then
(1 - a ) A m
+ AD
(VI)
Ai0 - Sal?'/*where, if C is in moles per liter, qa/2)c c = (2 These substitutions give
NOW,Ai
r
= Z
Aicoln =
=
C =~ (1 -
a)c+
(1 - a)[Amo - Sm(2
+
+ a)c,
+ O~)'/~C'/~] +
A AD' - S D ( 2 f (Y)1/2c1/2] (VII) AD" is larger than A m o , but the Onsager term S D is where a and b are the semiminor and semimajor axes, larger than Xm. If the formation of dimers is to explain respectively. In the calculation, b was set equal to 2a, the observed concentration dependence, then the dimer where a is the radius of the tetraalkylammonium ion term on the right of eq VI1 must decrease more rapidly monomer u n k 3 Using these relationships, the values with concentration than the monomer term. It follows of XO can be calculated from XO = (0.65 X 1O4)Ziep/3O0 if this does occur, that there is a concentration CE, at and are found to be Xoe11(DiBu2+) = 28.4, XOe"(DiPi2+) which the two terms on the right are equal to each = 31.0, Xoe11(DiEth2+) = 35.0 and XO~~~(DR!ID~+) = other. Above this Concentration, dimerization will 40.4 in HzO at 25". Comparison of the calculated result in a larger contribution to the conductance than values with those in Table I1 shows that for the larger that caused by the monomer alone, and below CE, ions the agreement is within 8%. There is serious dimerization will result in a smaller conductance. At disparity in the case of the smallest ion, DMD,2+ CE which does not have the same geometrical relationship A m 0 - Sn,(2 ~)'/*CE'/* = to its tetraalkylammonium analog. It should be noted that dimerization of the Bu4N+ ions in a, BurNBr [ADO - S D ( 2 ~)"zcE"z]z (VIII) solution at infinite dilution would result in an increase in l - a the conductance since the equivalent conductance and solving for CE'/~with Q = a / ( l - a ) and P = would change from 97.53 to 104.70. (2 gives ConcentTation Dependence. If cation-cation pairing is to explain the concentration dependence of these CE'" = (ADOQ - Amo)/(XDQ - X m ) p aqueous solutions then it must lead to a substantial For Bu4NBrg and DiBuBrz, A m o = 97.5, S m = 58.2, decrease in the conductance with concentration. The ADO = 104.6, and SD= 90.1. Assigning a the value 0.1 extent to which this explanation is compatible with gives C = 1.5 M . At a = 0.01, C = 1.41 and as a experimental observation can be seen from the following approaches zero, C approaches 1.40 M . c a l c ~ l a t i o n . ~In~ this discussion, RIN+ will be denoted This result shows that in the concentration range by the subscript 1, (RJi)22+by 2 and 2 - , the anion, by 1 1 4 in which the larger tetraalkylammonium to 3. At some finite concentration, the conductance of salts were studied, the formation of cation-cation pairs KQ, the solution can be expressed as Ksoln = K I f K Z will only lead to an increase in conductance instead of where K~ is the specific conductance X 1000. The the decrease observed. This indicates that previous equivalent conductance is obtained by dividing K ~ by~ I ~ explanations based on cation-cation pairing should be the concentration of electrolyte in equivalents per liter. reexamined. However, the magnitude of the deviaThe equivalent conductance for a solution of several tions from expected behavior, even in dilute solutions species is defined by and for many different types of measurements, imply that some type of ionic aggregation does take place. i Most of the discussion concerning the tetraalkylwhere Ziis the absolute magnitude of the valence and ammonium solutions have focused attention on the role C, the concentration of the species in moles per liter. of the cation, and most of the explanations have emphaIf C is the concentration of salt in moles per liter and 01 sized the unique structural properties of water. Howthe fraction of RJS+ ions dimerized to (R4hT)2,'+then ever, the deviation from ideality in conductance meaC1 = (1 - a)C, CZ= ( a / 2 ) C and CB= C, and surements were shown to be independent of temperaA s o ~ n= (1 - a)X1 ax2 Xa (VI
+
+
+
+
+ +
In eq V, Asoln is the conductance which would be measured at some concentration C, if (1 - a ) C parts of The Journal of Physical Chemistrv
(18) F. Perrin, J. Phys. Radium,7,l(1936). (19) The cross terms necessary to describe exactly a tertiary salt solution have been ignored in this approximate calculation.
TRANSPORT PROCESSES IN HYDROGEN BONDING SOLVENTS ture, whereas properties dependent upon structural effects have been found to change with temperature. Furthermore, the magnitude of the deviations from ideality as determined from a number of physical m e a ~ u r e m e n t s 4 ~is9 ~ extremely ~~ sensitive to the nature of the anion, and any satisfactory explanation must account for this. It is this aspect of the problem that we wish to discuss in some detail. Studies of other hydrogen bonding solvents, such as the alcohols21 and amides,24suggest that some of the features which heretofore have been ascribed to the unique properties of aqueous solutions may in fact be more general properties of hydrogen bonded solvents. In solvents of lower dielectric constant, such as ethanol ( E 24.33),22 propanol ( E 20.45),22and butanol (e 17.45),23 association constants for the tetraalkylammonium salts can be determined with some assurance, and it is found that for a given cation, KA increases with increasing anionic size as can be seen in Table IV. As the di-
3989
As the dielectric constant becomes still larger ( E >50), reliable values of KA can no longer be obtained. In this case, deviations from ideality can best be studied by using a rearranged form of eq I A' = A
- Ao + S d i - EClog c
=
( J - F&)c
(XII)
A plot of A' vs. c should give a straight, line of slope (J - FAo) for nonassociated electrolytes. J is a function of the ion-size parameter, d, which should increase with increasing ionic size; however, its magnitude cannot be calculated a priori. FA0 corrects for the effect of the electrolyte on the viscosity of the solution. The corresponding equation for associated electrolytes is A' = A
- A,
f
SdG - Ecy log cy
=
(Jr - FAoy - KAyf*A)C (XIII) However, when K A is small and
y
close to unity,
K A y f 2 A becomes essentially linear in c and cannot be Table IV : Association Constants for the Tetraalkylammonium Salts in the Alcohols Salt
MeOHa
EtOHb
PrOHb
BuOHC
..
39
640
3
75 123
149 266 415
1180
770
2200
BuaNCl Bu,NBr Bu4NI Bu~NCIOI a
See ref 21.
16
* See ref 22.
860
See ref 23.
electric constant increases, K A decreases as expected. However, the dependence of K A upon size remains opposite to that predicted by coulombic theory25and to that observed in where hydrogen bonding is not a predominant factor. For the larger anions, K A is larger than can be accounted for by simple coulombic theory. An explanation consistent with these facts is one based on a multiple-step association process involving desolvation of the anion in the second step.23
R4N+
+ X(ROH),-
Ki
Ka
I_ RJYROHX(ROH).-1 J_ R4NX(ROH).-1
+ ROH
(IX)
Since neither of the ion pairs contributes to the conductance, the total association constant, K A ,is equal to
K A = Ki(1 -k Kz)
(XI
The equilibrium constant for the first step, K1, can be calculated2efrom K1 = 2.524 X 10-3(aaeea/arkT)
(XI)
which gives k-1 values of 13, 25, 40, and 65 in methanol, ethanol, propanol, and butanol,23respectively. Values of log k - 2 appear to vary linearly with l / ~ . ~ 7
separated from J . Comparison of eq XI1 and XI11 shows that when an associated electrolyte is analyzed by eq XII, the slope of the line in a A' plot will be smaller the larger the value of K A . I n formamide, a hydrogen bonding solvent of high dielectric constant (e log), plots for the tetrabutylammonium and tetrapropylammonium halides give slopes which are smallest for the iodides and largest for the chlorides.24 This is the same pattern observed when conductance data for methanol solutions are analyzed by eq X I I ; however, the deviations are greater than those with formamide and allow association constants to be explicitly evaluated. The same type of A' plots were found for the tetraalkylammonium halides in aqueous solution, but the deviations from ideality did not increase at lower temperature in DzO solution, a result that strongly indicates that water structural factors are not involved. It is reasonable to assume that cation-anion pairing of the type given in eq XI can quantitatively account for the concentration dependence of conductance in dilute solution.27 As a consequence, at least part of the peculiar concentration dependence of the tetraalkylammonium salts in aqueous solution is not the result of water structure effects, but rather a general feature of hydrogen bonding solvents. (20) R. H. Wood, H. L. Anderson, J. D. Beck, J. R. France, W. E. de Urg, and L. J. Soltzberg, J . Phys. Chem., 71,2149 (1967). (21) R. L. Kay, C. Zawoyski, and D. F. Evans, ibid., 69, 4208 (1965). (22) D. F. Evans and P. Gardam, ibid., 72, 3281 (1968). (23) D. F. Evans and I?. Gardam, ibid., 73, 158 (1969). (24) D. F. Evans and J. Thomas, to be published. (25) R. M.Fnoss, J . Amer. Chem. SOC., 80,5059 (1958). (26) G. G. Hammes and J. I. Steinfield, ibid., 84, 4639 (1962). (27) R. L. Kay, D. F. Evans, and G. P. Cunningham, J . Phys. Chem., 73,3322 (1969). Volume 79, Number 11 November 1060
3990
CHARLES L. CRONAN AND FRIEDEMANN W. SCHNEIDER
However, a number of measurements in more concentrated solutions including partial molar vo1umes,6t1z heats of mixing,Z0 and volumes of mixing10 provide strong evidence for some kind of higher aggregation, Furthermore, the similarities in behavior in water and the alcohols do not extend to higher concentrations; in water P2for BueNBr decreases with concentration while in methanol” and ethanolz8it increases. We believe that an explanation that is applicable to the more concentrated solutions and which is compatible with the results for dilute solutions is one that involves a reaction of the ion pairs (either the solvent separated or contact ion pairs of eq IX) with another cation to give a triple ion.
R8X
+ RIN+
[RINXRIN]+
Unlike the triplicates found in media of low dielectric constant where stabilization is entirely coulombic in nature,z9 part of the stabilization of this entity in aqueous solution must be attributed to a hydrophobic effect.
Acknowledgment. This work was supported by Contract No. 14-01-001-1281 with the Office of Saline Water, U. S.Department of the Interior. (28) W.Y. Wen, in “1966 Saline Water Conversion Report,” U. 9. Department of the Interior, Washington, D. C., 1966, p 13. (29) H. S. Young and C. A. Kraus, J . A m e r . Chem. Soc., 73, 4752 (1951).
Cooperativity and Composition of the Linear Amylose-Iodine-Iodide Complex l*$b by Charles L. Cronan and Friedemann W. Schneiderlo Department of Chemistry, University of Southern California, Los Angeles, California 90007 (Received August 19,1968)
Adsorption isotherms of iodine-iodide in the helical cavity of amylose (D.P. = 800) have been determined by spectrophotometric measurements of titration curves at various iodide concentrations (10-5 to 2.5 x 10-2 M KI) and temperatures. A new description of the stoichiometry of the bound species is given in terms of two limiting models: A t very low iodide concentrations the predominantly bound species is considered to be Iz whereas at high iodide the bound species is 1 3 - (the theoretical Ising lattice treatment predicts Id2-). Both models apply in the intermediate transition region where competition for binding sites occurs between Iz and Is-. The over-all composition of the bound species can be expressed approximately as Ia.Ib- where b varies between zero and unity. Statistical analysis of the isotherms using the exact one-dimensional Ising model (n = 15) with first nearest neighbor interactions (between - 1.6 and - 3.0 kcal/mol) shows that the degree of cooperativity is moderate. The main source of the stability of the complex originates from the intrinsic binding of a species by one turn of the amylose helix. The enthalpy and entropy of binding have been determined at i W KI. Factors affecting the position of A, of the blue band are attributed to the aggregation of partially or completely filled amylose helices.
Introduction Although the blue amylose-iodine-iodide (AI) complex has been known for over 150 years, its detailed structure and composition are not yet completely understood. Basic to the understanding of its over-all stability is an insight i n t o the role played by cooperative interactions between bound species which for many years have been suspected to be importanta2 I n previous work a quantitative model treatment of cooperaM KI and 20.0” tivity in the linear AI complex a t has been given for the first time.3 Complex formation was treated as a one-dimensional adsorption process of iodine species into the helical cavity of amylose, and it was shown that cooperativity in the linear complex is surprisingly moderate, the over-all stability being due T h e Journal of Physical Chemistry
mainly to the intrinsic binding of an iodine species to a helical turn of ca. eight anhydroglucose units. This work complements the previous equilibrium dialysis measurements and presents spectrophotometric titration studies which cast new light on the stoichiometry and thermodynamics of the complex. The degree of (1) (a) Presented a t the Pacific Conference on Chemistry and Spectroscopy, Anaheim, Calif., Oct 1967. (b) Abstracted in part from the h1.S. Thesis of C. L. Cronan, University of Southern California, 1968. (c) T o whom correspondence should be addressed. (2) (a) For a general review of the complex, see H. Morawetz, “Macromolecules in Solution,” Interscience, New York, N. Y., 1965; (b) R. L. Whistler and E. F. Paschall, Ed., “Starch: Chemistry and Technology,’’ Vol. 1, Academic Press, New York, N. Y., 1965, Chapter 1; (c) J. Szejtli and S. Augustat, Staerke, 18, 38 (1966). (3) F. W. Schneider, C. L. Cronan, and S. K. Podder, b. P h y s . Chem., 72,4563 (1968).
