J. Phys. Chem. 1903, 87, 3349-3354
The essentially naked character of the M(CO)5species in C7F14 is reflected in their high reactivity toward CO, Nz, cyclohexane, or M(CO)6. In each case the rate constant is within 1 order of magnitude of the rate constant expected for diffusion-controlled processes in C7F14. It is noteworthy that the rate constant of reaction of M(CO)5 with CO is about 3 orders of magnitude faster in C7F14 than those found in cyclohexane. This is expected from the observed rapid coordination of M(CO)5to cyclohexane and the fact that the spectrum indicates that equilibrium 11 M ( C ~ ) & C ~ H ~ M(C0)5 Z) -k C6Hiz (11) lies to the left at room temperature. The strength of the interaction between M(CO)5 and cyclohexane could be determined if values of the equilibrium constant for equilibrium 11could be obtained at various temperatures. Such a determination has not been possible in the present study. However, our preliminary results do show that Cr(CO)&C6HlZ) is very labile (Itdise1 106 8-9 and there also appears to be evidence for an associative reaction of CO with C ~ ( C O ) ~ ( C ~ H ~ Z ) . (21)In a recent paper on the picosecond flash of Cr(CO)8 in cyclohexane, benzene, and methanol it was shown that the Cr(CO)5 (solvent) complex was formed within 25 ps of excitation. (22)J. A. Welch, K. S.Peters, and V. Vaida, J. Phys. Chem., 86,1941 (1982).
3349
The observation that M(CO)5reacts with M(CO), to give M2(CO)11could also be relevant to studies in alkane solvents as it had been a concern in earlier studies that Mz(CO)ll might be formed and possibly be the reason for the observed “impurity complexes”. Our data show that Crz(CO)ll is labile (kdimL lo6 s-l) and therefore it is extremely unlikely that the formation of these species plays an important role in the reactions of Cr(C0)5 in dilute alkane solutions of Cr(CO)6. In this connection it is also worth noting that we have found no definite evidence for the isocarbonyl complex M(C0)50C. (From matrix studies it is expected that Cr(CO)50Cabsorbs at about 460 nm.9) In conclusion the results presented here demonstrate the high reactivity of 16-electron coordinatively unsaturated species and show how their reactivity is modified by the formation of solvent complexes even in hydrocarbon solution. Acknowledgment. We thank the National Board for Science and Technology (Ireland) and the C.N.R.S. (France) for travel and subsistence expenses. C.L. acknowledges postgraduate awards from the Department of Education and Trinity College, Dublin. Registry No. Cr(C0)6, 13007-92-6; MO(CO)~, 13939-06-5; W(CO)6,14040-11-0; Cr(CO)5,26319-33-5;Mo(CO),, 32312-17-7; w(co)s, 30395-19-8; co, 630-08-0; Nz, 7727-37-9; C&12,110-82-7.
Cycloamylose Complexation of Inorganic Anions Robert I. Geib, Lowell M. Schwartz,’ Mlchael Radeos, and Daniel A. Laufer Deperlment of Ctmmlsby, Unverslty of Massachusetts, Boston, Massachusetts 02125 (Received: June 1, 1982; I n Final Form: January 28, 1983)
Formation constants for inclusion complexes of cyclohexaamylose (a-cyclodextrin) and cycloheptaamylose (/3-cyclodextrin)with inorganic anions ClO;, SCN-, I-, Br-, NO3-, and 10,- in aqueous solution are determined at various temperatures by a novel pH potentiometric method. Complexes with chloride ion even at 1 M C1could not be detected. AH” and AS” values for cyclohexaamylose complexation of C104-, SCN-, and I- are found to conform with a previously reported correlation based on corresponding complexation with a variety of substrate species including aliphatic and aromatic carboxylic acids and carboxylate anions and substituted phenols and phenolate anions. This correlation serves as a basis for theorizing a common binding mechanism for the complexes, Le., polar interaction between cyclohexaamylose and substrate. Thus, the same binding mechanism is indicated for small inorganic ion substrates.
Complexes of cycloamyloses (also cyclodextrins, to be denoted by Cy) and their derivatives have been reported with substrates of widely varying structures. These complexes have been detected by several physical and chemical techniques but pH potentiometry has proven to be particularly successful in terms of accuracy and precision. We have reported pH potentiometric determinations of complexation constants of several aqueous organic acids, phenols, anions and small solvent molecules.14 Others5s6 (1)Gelb, R. I.; Schwartz, L. M.; Johnson, R. F.; Laufer, D. A. J. Am. Chem. SOC.1979,101, 1869-74. (2)Gelb, R. I.; Schwartz, L. M.; Laufer, D. A. Bioorg. Chem. 1980,9, 450-61. (3)Gelb, R. I.; Schwartz, L. M.; Cardelino, B.; Furhman, H. S.;Johnson, R. F.; Laufer, D. A. J. Am. Chem. SOC.1981,103,1750-7. (4)Gelb, R. I.; Schwartz, L. M.; Radeos, M.; Edmonds, R. B.; Laufer, D.A. J. Am. Chem. SOC.1982,104,6283-8. 0022-3654/83/2087-3349~0~ .50/0
have also employed pH potentiometry similarly. The method involves measuring the pH of an aqueous acidbase conjugate which will also be called the buffer components and denoted by HB and B-. If cycloamylose is added to such a solution and forms a complex with either or both buffer components, the buffer/H+ equilibrium will be shifted and a pH perturbation will be observed except in the unlikely case that complexes of exactly equal strength are formed with both buffer components. The extent of pH perturbation, indeed, depends on the difference in such complex strength. Experimentally one measures the pH of a series of solutions containing variable amounts of buffer and cycloamylose and then fits the data (5)Miyaji, T.;Kurono, Y.; Uekama, K. Chem. Pharm. Bull. 1976,24, 1155-9. (6)Connors, K. A.; Lipari, J. M. J . Pharm. Sei. 1976,65, 379-83.
0 1983 American Chemical Society
3350
Gelb et at.
The Journal of Physical Chemistfy, Vol. 87,No. 17, 1983
to a set of model equations' including activity coefficient correlations. This procedure usually yields the complex formation constants for cycloamylose with both buffer components. In this communication we seek to study complexes with inorganic anions of strong acids and, since these anions do not exist in equilibrium with appreciable concentrations of their conjugate acids in dilute aqueous solution, the procedure described above must be modified. The inorganic anions (to be denoted by A-) are introduced into a solution of cycloamylose and HB/B- buffer. If Acomplexes with Cy, Cy is withdrawn from previously established equilibria involving HB and/or B- and so the pH is perturbed. The extent of perturbation depends upon the amount of A- added and upon the strength of the CyAcomplex, so that this strength is calculable. Because the substrate anion A- perturbs the pH measurement only by means of the intermediate HB/B-/Cy equilibrium, we call this modification "pH perturbation of the second kind" by analogy with the classification of electrodes into those of the second kind which measure species in solution using some intermediate cell reaction. Also we will call the inorganic anions A- "second-kind" substrates to distinguish them from the buffer components HB/B- which we will call "first-kind" substrates.
Model Equations The complex formation constants of interest are adjustable parameters in a set of coupled nonlinear equations describing the equilibria and conservation relationships involved in each aqueous solution. These equations and the data treatment procedures for experiments involving fiist-kind substrates only have been described earlier.' We now extend those equations to include second-kind substrates as well. Each aqueous solution is prepared with known analytical concentrations of buffer acid Fm, buffer base FB-, cycloamylose Fey, a second-kind substrate FA-, and alkali-metal cation FM+ equivalent to FB- and FA-. Equilibrated solutions of these components contain uncomplexed species H+, OH-, HB, B-, M+, Cy, A- and the complexes CyHB, CyB-, and CyA-. Certain other complexes are conceivable but not included. These are, for example, (1)ternary complexes Cy2HBand Cy,B- which are avoided by judiciously selecting a buffer of benzoic acid/benzoate which is known to form no such complexes,' (2) associations between HB and A- whose absences are confirmed by independent experiments of adding A- to HB/B- buffer and detecting no significant change in pH other than those ascribable to activity coefficient effects, and (3) mixed CyHB-A- and ternary Cy,A- complexes which are eliminated by inference during the data treatment procedure as will be explained later. The hypothesized equilibria are for the buffer [H+I[B-I YH+YB[HBI YHB for complexes with first-kind substrates [CyHBl Y C ~ H B K C ~ H=B [CY][HBI Y C ~ Y H B K, =
(1)
(2)
for complexes with a second-kind anion (4)
and the water autoprotolysis equilibrium. Bracketed
species here denote molar concentrations and subscripted y quantities are molar activity coefficients. The relevant conservation equations are for the components FHB
+ FB- = [HB] + [B-] + [CyHB] + [CyB-]
FCy
= [Cy]
+ [CyHB] + [CyB-] + [CyA-]
FA-
= [A-]
+ [CyA-]
and an electroneutrality relation [H+] + F M + = [B-] + [CyB-] + [A-]
+ [CyA-] + [OH-]
The molar activity coefficients which appear above are provisionally taken as follows: those referring to uncharged species are assumed to be unity and those referring to ions are estimated from the Debye-Huckel correlation with temperature-dependent parameters from the tabulation of Robinson and Stokes' and with ion size parameter values in nanometers of 0.9 for H+, 0.35 for OH-, 0.6 for benzoate ion, 0.3 for I-, and 0.35 for SCN- and for C104-. The experimental data of pH vs. analytical concentration were fitted to the model equations by an interactive nonlinear regression computer program described elsewhere.s The adjustable parameter of interest here is the complex formation constant KcyA-but we also allow K, to adjust as well since any inaccuracy in the pH meter calibration propagates into the value of K,. If the known K, value for benzoic acid were fixed in these equations, then any such inaccuracy would be an uncompensated determinate error and would probably prevent a satisfactory fit of the experimental data to the model equations. However, we do fix previously measured values of KCyHB and KCyB-. In these experiments we have employed two forms of cycloamylose: a-cyclodextrin or cyclohexaamylose, which will be denoted 6-Cy, and P-cyclodextrin or cycloheptaamylose, which will be denoted 7-Cy. As explained above, preliminary experiments are required to determine the complex formation constants of cycloamylose, both 6-Cy and 7-Cy, with the buffer components benzoic acid/ benzoate over a range of temperatures. Values of these constants for 6-Cy were reported previously' and results of measurements with 7-Cy will be reported and explained later in this communication.
Confirmatory Experiments 6-Cy Complexes with I-, ClO,-, and SCN-. Using pH perturbation of the second kind we measured KGQA- values at 25 "C for iodide, perchlorate, and thiocyanate ions. Successive incremental portions of known 6-Cy and Asamples were added to a buffer of 1 mF each of benzoic acid and benzoate and the pH was recorded after equilibration after each addition. By this procedure the 6-Cy concentration varied from 0 to 0.03 F and the A- concentration from 0 to 0.05 F in 10-15 increments. The pH vs. analytical concentration data were fitted to the model equations detailed above and yielded residual standard deviations within the statistical scatter expected from pH meter and volumetric uncertainties. Furthermore, the pattern of residuals appeared random during the course of the titrations, which indicates that the set of model equations was an accurate representation of the physicochemical behavior of these solutions. This absence of uncompensated determinate error serves to confirm our a priori hypothesis that ternary complexes such as Cy2A~
(7) Robinson, R. A.; Stokes,R. H. "Electrolyte Solutions",2nd revised ed.; Butterworths: London, 1965. (8) Schwartz, L. M.; Gelb, R. I. Anal. Chem. 1978, 50, 1571-6.
Cycloamyiose Complexatlon of Inorganic Anions
TABLE I: Complex Formation Constants of 6-Cy with Perchlorate, Iodide and Thiocyanate at 25 "C anion C10,1-
SCN-
PH perturbation conductance 45.8 i 0.9' 19.0 * 0.3 34.6 i 0.4
*
51 S b 32 i 12 41 t 10
lit. values 28.9,c 2gd 12.4,c 13.5e 18.7c
a Standard error estimates based on standard deviations of 0.002 pH unit and 0.2-mg gravimetric uncertainties ropagated through the nonlinear regression calculations. Standard error estimates based on 0.2%conductance and 0.002-mL volumetric uncertainties. Reference 10. Reference 11. e Reference 12.
or CyHB-A- were not present. Any significant concentration of these species would have invalidated the conservation equations as written and so would have precluded satisfactory fits of the data to the model equations. The results of these experiments appear in Table I under the heading of "pH perturbation". In order to confirm the experimental and theoretical methodology described here, we measured the complexation constants for these same systems using a completely independent method, conductometry. The conductance of a solution of the A- salt is decreased upon addition of Cy because (1)the equivalent conductance of the larger CyA- ion is significantly less than uncomplexed A- and (2) because added Cy increases viscosity of these solutions thus decreasing the equivalent conductance of all ionic species. Correction for this latter effect is made by noting the decrease in conductance of sodium chloride solutions whose ions do not form complexes with CY.^ The experimental and calculational procedures employed in these measurements have been detailed previo~sly.~ Results obtained here for complexes of 6-Cy with perchlorate, iodide, and thiocyanate ion are shown in Table I. The conductance measurements yield complexation constants with far less precision than by pH perturbation but the two methods c o n f i i each other to within statistical uncertainty. Table I also shows corresponding measurements reported by other workers. These literature values are somewhat lower than ours but the reasons for the discrepancies are not apparent.
Conditional Formation Constants for Weak Complexes The determination of thermodynamic formation constants requires that ionic activity coefficients be estimable from correlations such as the Debye-Huckel equation. However, preliminary experiments with several inorganic ion substrates showed that very weak complexes are formed with 6-Cy or 7-Cy. In order to determine these small formation constants with satisfactory precision, it is necessary to use such large A- concentrations that the solution ionic strength exceeds the range of validity of the activity coefficient correlations. It is thus necessary to forego the determination of thermodynamic formation constants in these instances in favor of conditional formation constants measured at specified ionic strength levels. The methodology of pH perturbation must, therefore, be altered somewhat by defining and determining values for conditional formation constants K b y H ~ (9) Gelb, R. I.; Schwartz, L. M.; Murray, C. T., Laufer, D. A. J. Am. Chem. SOC.1978,100, 3553-9. (10) Wojcik, J. F.;Rohrbach, R. P. J. Phys. Chem. 1975, 79,2251-3. (11) Cramer, F.;Saenger, W.; Spatz, H.-Ch. J.Am. Chem. SOC.1967, 89, 14-20. (12) French, D.Adu. Carbohydr. Chem. 1957,12, 189-260.
The Journal of Physlcal Chemistry, Vol. 87, No. 17, 7983 3351
TABLE 11: Conditional Formation Constants of Complexes of Cycloamylose with Benzoic Acid and Benzoate Ion at 25 "C ionic strength.
M-
K I s - c Y ~ ~K f 6 - c y ~ -
K'7-CyHB
K'7-CyB-
808i 10 1 1 . 8 2 1 . 0 5 4 2 i 14 833 i 19 12.8 f 1 . 8 0.10" 819 i 19 10.2 i 2.1 0.50' 876k 15 1 5 . 7 i 1 . 2 5 9 2 i 15 1.0' 915 f 16 17.3 i 1 . 4 673 2 6 ' Ionic strength adjusted with KC1.
241 2
0.003
'
305 2 37 i 1
and K'cyB- of Cy with the buffer components. These conditional constants are
which equations show the relationships to the thermodynamic constants defined in eq 2 and 3. In order to measure K/CyHB and K b y B - it is necessary to establish the specified ionic strength with an electrolyte whose ions do not form complexes with Cy. We have shown previouslyg by 13C NMR spectrometry that chloride ions form no such complexes with 6-Cy up to 0.5 M levels. In order to confirm that no complexes form with 7-Cy also and up to 1M levels we measured K'cyHB and K'cyB- with both 6-Cy and 7-Cy in KC1 solutions using conventional "first-kind" pH perturbation methodology' but by fixing activity coefficients at unity. The results of these experiments are shown in Table 11. We notice that values of both K ' w Y H B and K',GyHB are about 15-20% greater at 1 M ionic strength than in dilute solution. It is important to demonstrate that this variation is due to the effect of ionic strength on activity coefficients rather than to complex formation of Cy with chloride ion. Activity coefficients affect K/CyHB through the quotient YC~YHB/YC HB. Both CyHB and Cy are structurally very similar, wodd be expected to interact with the solvent similarly, and thus are likely to have equal activity coefficients. Therefore, K'cyHB variation should be proportional to YHB variation and this latter effect can be determined independently by measuring the solubility of benzoic acid in aqueous KC1 solutions. We equilibrated solid benzoic acid with water and with 1.1F KC1 at 25 "C and then analyzed the supernatant liquid for dissolved acid by titration with standardized NaOH. After correcting for the concentration of dissociated acid, we found the concentration of molecular benzoic acid decreased from 0.026 M in pure water to 0.022 M in 1.1 F KC1 solution. This decreased solubility corresponds to an approximate 159% increase in yHB and so accounts almost completely for the observed increase in K'cyHB with increasing KC1 concentration. The fact that increasing chloride concentration leads to increasing KbyHBand K b y B - values indicates that negligible chloride complexation occurs. Such complexation would effectively decrease these conditional formation constant values. If chloride ion were complexing, the concentration of Cy available to complex with HB and Bwould be diminished and so the perturbation of the buffer pH by Cy additions would be less. This effect if not taken into account would appear as lesser values of K'cyHB and K', B-. Thus, we are confident that C1- does not complex with 6-Cy or 7-Cy up to 1M and the values tabulated in Table I1 represent the effect of ionic strength on the conditional formation constants KLym and K L y ~ - .Hence,
3352
The Journal of Physical Chemistry, Vol. 87, No. 17, 1983
we will utilize these conditional constants in the secondkind pH perturbation experiments seeking the conditional complex formation constants of inorganic ions. The only experimental modification is to introduce both the substrate A- and the buffer components into the initial solution and then add increments of solid Cy, taking pH readings after each equilibration. The buffer components each have concentrations approximately 1mF and the Asalt is at 0.5 F. Because the added solid Cy makes a negligible increase in the solution volume and because the reaction of A- with Cy leaves the total ionic concentrations unchanged, the ionic strength of the solutions are practically invariant at 0.5 M. Results of these conditional measurements are shown in Table 111. Second-kind substrates are bromide, iodide, iodate, thiocyanate, nitrate, and perchlorate ion and both 6-Cy and 7-Cy are used. The temperature is varied between 20 and 59.5 "C. Also shown in Table I11 are temperature-dependent thermodynamic constants of iodide, thiocyanate, and perchlorate complexing with 6-Cy and conditional constants at 0.1 M ionic strength reported by Rohrbach et al.I3 Uncertainty Estimations We notice that our conditional formation constant values differ substantially from those reported in ref 13, more than can reasonably result from the differences in ionic strength used. Because the standard error estimates quoted along with our values in Table I11 do not comprise the full extent of their uncertainties, we undertook a further propagation-of-determinate-error analysis. We chose one particular as typical, the complex between 7-Cy and C104-at 40 "C. Each source of determinate error and our estimate of its magnitude is listed in the first two columns of Table IV. The upper part of the table refers to the experiment and calculation done to determine the conditional formation constants of benzoic acid and benzoate complexes, i.e., K \ x y H B and K ' 7 q y B - . These uncertainties then contribute systematic error to the determination of the conditional formation constant for the C10, complex and this latter error analysis is shown in the lower part of the table. The estimated error from each source parameter is propagated into the formation constant by simply recalculating that constant with the parameter perturbed from its estimated best value. The observed perturbation in the formation constant is shown in the columns headed AK' in Table IV. In addition to experimental sources of error, we have also recognized that one of the assumptions in our methodology cannot be strictly correct and have attempted to estimate the magnitude of this error under the entry "specific ion effect errors". By this we are referring to the procedure whereby we determine K b , , H B and K b y B - in 0.50 M ionic strength solution established with KC1 and then employ these same values in 0.50 M ionic strength solutions established with A- to find KbyA-on the assumption that K ' C ~ Hand B K ' C ~ Bdepend only on ionic strength, not on specific ion effects. This practice is equivalent to assuming that the activity coefficient factors Y H B Y C ~ / Y C ~ H Band YB-YC,,/YC~B- are only ionic strength dependent. This is certainly not true at the 0.5 M level for individual ionic activity coefficients and probably not strictly true for these factors either. However, the magnitudes of the systematic errors introduced here are difficult to quantify and so we have taken a somewhat arbitrary but reasonable variability of 2% as a basis for the error analysis in Table IV. (13) Rohrbach, R. P.; Rodriguez, L. J.; Eyring,E. M.; Wojcik, F. J. J . Phys. Chem. 1977,81, 944-8.
Gelb et ai.
-4-
,/e
c
PRCP4YCI
/'- BUT4\2L i'----
~8
+4
2
4
-~~ 1 __ -E -12 -I6 23 24 28 2S' c 3 - O W K
Figure 1. Correlation of AH" and A S " of 6-Cy complexation with 28 substrates from ref 3 and 4 shown as circles. The solid line is the least-squares fit AH" = (410 f 15)AS" - (1.2 f 0.2) X lo3 cal mol-' The squares indicate the positions of A H " , A S " for the 6-Cy complexes with three inorganic ions reported here.
In an effort to provide some experimental verification of this error analysis, we made two separate series of experiments to determine K',-CS10,- and both series are shown in Tables I11 and IV. Both series are done in 0.50 M ionic strength but one uses 0.10 F KC104and 0.40 F KC1 while the other uses 0.50 F NaC104. K',.CyC1O,values at corresponding temperatures differ by 5.15%. Our error analysis in Table IV estimates approximately 13-18% and so we must conclude that our conditional formation constants have uncertainties at least in the range 10-20%. Actually, we have found by similar error calculations that the weakest complexes studied here such as those with bromide ion yield conditional constants with uncertainties up to 30%. Because of the relatively large uncertainties of the conditional formation constants, we felt that any further calculation of enthalpies and entropies from the temperature variations would be untenable. However, the thermodynamic formation constants of 6-Cy with I-, SCN-, and C104- seem to be sufficiently reliable to allow meaningful consideration of the thermodynamic parameters AH" and AS" of complexation and so we proceeded to make van't Hoff plots which yielded results shown in Table V. In this table we also show AH" and ASo for the buffer components benzoic acid and benzoate as reported in ref 1. Conditional A€€and A S values in 0.50 M ionic strength are given in parentheses for comparison. Discussion In an earlier communication3 we reported the existence of a correlation between AH" and AS" of complexation of 6-Cy with 20 substrates including such diverse species as aliphatic and' aromatic carboxylic acids and anions and substituted phenols and their anions. We can now augment the original 20 species with hydroxide ion3 and several small organic molecules: acetonitrile, 1,4-dioxane, ethanol, 2-propanol, 2-methyl-2-propano1, cyclohexanol, and dimethyl sulfoxide. Recalculation of the least-squares line with all 28 AH",ASodata points yields the correlation
AH" = (410 f 15)AS" - (1.2 f 0.2) X lo3 where AH" is given in cal mol-' and AS" in cal mol-' K-l. The uncertainties are standard error estimates based on scatter of the observed points from the line. This correlation line is drawn in Figure 1 along with the circles representing the 28 species used in its calculation. The
'he Journal of Physical Chemistry, Vol. 87, No. 17, 1983 3353
Cycloamylose Complexation of Inorganic Anions
?
d:??
+I N
+I
0
0 0 0 +I
9,
+I
Z$??
%$e
-$Z""
. I
(9".
(9"?
0 0
000
+I +I
+I
F?(9
9m0
+I
+I
U
am* m i 3
4,-
*m
222
CDn? 00
+I +I
+I
19
99"
+I +I
mmt-
*4,
wrl
d 2
i
E
,
2 +I
m
E
w
L"??
2
000
N
+I
+I
+I
F"?"
??
mm
*
wwm w i i
m+-
a
H +I
+I +I
Y I
+I +I
+I +I
+I
+I
+I
+I
+I
+I +I
2: 222 aJ
Y
s e93 3 .-u0 E
n
3 +I
+I
+I
'0. t-. t-em
m.
+I
3
i
i
+I
+I
0y
oot3
i
mmm
U
mim i
d
i
+I
+I +I
3
t-1m +I
+I
+I
mmw
m o m
30000
* N 3
Po)* i
omm
+I
i
+I
i
+I
4,
t-m
+I
+I +I
N o m m t-m bi4, m m m w w m v me4
o m
m m
'0. o o m * o m
3354
The Journal of Physical Chemistry, Vol. 87, No. 17, 1983
Gelb et at.
between 6-Cy and the substrate provide the driving force for the complexation reaction and these interactions result in torsional constraint of the 6-Cy structure reflected in propagated estd uncertainties, % the substantially negative AS" values for the complexation magnisource of error tude, % AK',-c~HB A K ' , . ~ ~ ~ - reactions. The mechanism of "hydrophobic" interaction whereby release of water molecules from the uncomplexed 0.4 0.4 0.6 s o h vol reactanta or from the 6-Cy cavity provides free energy for 1 0.2 0.6 FHB the reaction does not seem to play an important role in 1 nil nil FBthese reactions. The observed substantially negative ASo 0.5 0.6 0.7 purity of 7-Cy values in these complexations as well as the compensation max probable error -1 -2 temperature value of 410 f 15 K indicate that differential propagated uncertainties .~ solvation and hydrophobic effects are of minor importance for aK',-cyclO;, % in &Cy complexations of the inorganic anions studied here. estd magniin 0.10 F in 0.50 F This view seems to be supported by examination of source of error tude, % C10; ' c10,conditional AH and AS values for complexations of benzoic acid and benzoate, iodide, and thiocyanate anions calcu4.1 3.3 K 7-CyHB 1.2 lated from conditional constants in Table 111, and shown K PCYB' 1.9 1.6 0.5 1.1 0.5 s o h vol 0.4 in parentheses in Table V. It seems clear that the conFHB 1 0.2 0.1 ditional AH and A S values which refer to a highly ionic 1 0.2 0.1 FBmedium of 0.5 M ionic strength do not differ substantially 1 1 11 FC10, from the dilute-solution AH" and AS" values. Because purity of 7-Cy 0.5 1.5 0.8 hydrophobic interactions would presumably differ markspecific ion effect edly between pure water solvent on one hand and the ionic errors strength 0.50 M on the other, such differences would be i n K'7-CyHB 2 6.9 5.5 in K ,.cyB2 1.7 1.8 reflected in marked changes of AH and A S with varying ionic strength. We interpret the relative invariance of max probable 20 15 error these parameters as further evidence that the complexations are the result of specific interactions between 6-Cy a Observed K',-CyC1O; = 6.3 in 0.10 F KClO,, 0.40 F and the substrate species and are not substantially influKCl. Observed K',-cyclo; = 6.6 in 0.50 F NaClO,. enced by hydrophobic forces. TABLE IV:
K
7-cyc10;
ProDaeation of Determinate Errors into U n c e i t i n t y a t 40 "C
~
-
-
TABLE V: Enthalpy and Entropy Changes of Cyclohexaamylose Complex Formation with Perchlorate, Iodide, Thiocyanate, Benzoate Ions and Benzoic Acid substrate
A H " , kcal mol-'
AS",cal mol-' K-'
-13.5 f 0.8 -6.3 f 0.2' -5.9 f 0.2 -13.7 t 0.6 (-7.0 f O . l ) b (-18.2 f 0.4) thiocyanate -6.8 f 0.3 -15.9 f 1.1 (-6.3 i 0.3) (-14.5 f 0.9) benzoateC -3.9 i 0.3 -8.4 i 1.1 (-3.6 i 0.4) (-6.8 i 1.2) benzoic acidC -10.2 t 0.1 -20.9 t 0.5 (-9.9 i 0.1) (-19.9 t 0.5) Uncertainties are standard error estimates based on scatter of points from least-squares van't Hoff lines. Entries appearing in parentheses are conditional values at 0.50 M ionic strength. Other entries are based on thermodynamic formation constants. From ref 1. perchlorate iodide
positions of the complexes reported in ref 4 are labeled and the three squares represent the inorganic ion complexes from Table V. The apparent fit of these thermodynamic parameters for the 6-Cy complexations of perchlorate, iodide, and thiocyanate ions leads us to conclude that the binding mechanisms in these complexes are of the same nature as those studied previ~usly.~Polar interactions
Experimental Section Cyclohexaamylose and cycloheptaamylose reagents were obtained from the Aldrich Chemical Co. 6-Cy samples were equilibrated with the atmosphere for about 1week before use. Prolonged vacuum drying at 100 "C indicated that the aerated material consisted of the hexahydrate amylose. 7-Cy samples were twice recrystallized from water and air dried for several days to yield the octahydrate. Most inorganic salts were recrystallized from water and dried in vacuo at 100 "C overnight. All materials were reagent grade. pH measurements employed an Orion 801 pH meter equipped with conventional glass and reference electrodes. Particular care was taken to assume thermal equilibration. Electrodes were equilibrated at the measurement temperature until drift rates of less than 0.001 pH/15 min were observed. Measurements with perchlorate solutions employed a reference electrode especially made up in NaCl solution to avoid KC104 precipitation at the liquid junction. Acknowledgment. We gratefully acknowledge the support of the National Institute of General Medical Sciences, U.S. Public Health Service (Grant No. GM 26004). Registry NO.ClO,, 14797-73-0; SCN-, 302-04-5; I-, 20461-54-5; Br-, 24959-67-9; NO;, 14797-55-8; IO3-, 15454-31-6;a-cyclodextrin, 10016-20-3;0-cyclodextrin, 7585-39-9.
