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Host-Guest Inclusion Complexation of #-Cyclodextrin and Triiodide Examined Using UV-Vis Spectrophotometry Janet L Pursell, and Christopher J Pursell J. Phys. Chem. A, Just Accepted Manuscript • DOI: 10.1021/acs.jpca.6b00982 • Publication Date (Web): 21 Mar 2016 Downloaded from http://pubs.acs.org on March 24, 2016

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The Journal of Physical Chemistry

Host-Guest Inclusion Complexation of α-Cyclodextrin and Triiodide Examined Using UV-Vis Spectrophotometry

Janet L. Pursell1 and Christopher J. Pursell1,2* 1

Department of Chemistry, Trinity University, One Trinity Place, San Antonio, Texas 78212,

United States. 2

Okinawa Institute of Science and Technology Graduate University, 1919-1 Tancha, Onna-son,

Kunigami-gun, Okinawa, Japan 904-0495. *

Corresponding Author: [email protected]; (210)999-7381

Abstract The historically relevant host-guest complexation of α-cyclodextrin (α-CD) and triiodide (I3-) in aqueous solution was examined using a systematic UV-Vis spectrophotometric approach.

This particular system is experimentally challenging because of the coupled

equilibria, namely, I2 + I- ⇌ I3- and α-CD + I3- ⇌ α-CD·I3-.

We therefore developed a unique

experimental approach that allowed us to determine the concentration of all iodine species. This enabled us to unequivocally demonstrate that the large increase in the UV absorbance with added α-cyclodextrin is due to an increase in the overall triiodide concentration as α-CD essentially converts iodine to triiodide according to the coupled equilibria.

Herein we report (a)

the complexation stoichiometry is 1:1 (i.e. the host-guest complex is α-CD·I3-), (b) the binding constant is KH-G = (1.35 ± 0.05) x 105 M-1 at room temperature, and (c) the binding constant is temperature dependent with ∆H = -31.0 ± 0.9 kJ/mol.

1

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Introduction Host-guest complexation by cyclodextrins (CDs) has a long and rich history as recently reviewed by Crini.1

These interesting shallow, truncated, cone-shaped molecules are cyclic

oligosaccharides with hydrophobic cavities and hydrophilic exteriors.

The most common and

extensively studied are α, β, and γ-CD with six, seven, and eight glucopyranose units, respectively.

Inclusion of guest species by cyclodextrin hosts can result in remarkable

modifications in the physical, chemical, and biological properties of the guest.

Practical

applications are now found in many industries including pharmacy, food, chemistry, chromatography, catalysis, biotechnology, agriculture, cosmetics, hygiene, medicine, textiles, and the environment.1-3 While Schardinger had originally used the reaction of iodine to distinguished α-CD from β-CD,1 pioneering work on cyclodextrin host-guest binding was conducted by French and his co-workers and concerned the complexation of α-CD and triiodide.4

Using the method of

continuous variation or Job’s method, they spectrophotometrically determined that for the complex aqueous system of α-CD with iodine, iodide and triiodide, the α-CD preferentially binds the triiodide as α-CD·I3-.

While they demonstrated the 1:1 stoichiometry for the

complex, their experimental method did not allow them to determine the binding constant (KH-G) or its temperature dependence (i.e. ∆H).4 2


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Many years later, Diard et al.5 used potentiometric methods to study the binding of α, β and γ-cyclodextrins to iodine, iodide and triiodide.

While the quality of the fits to their data

(using multiple equilibrium constants) is somewhat poor, they did discover the preferential binding of α-CD and triiodide with 1:1 stoichiometry and KH-G = (3.3 ± 0.3) x 105 M-1 at room temperature.5

Next, Kitamura et al.6 conducted very careful isothermal titration calorimetry

(ITC) experiments examining the complexation of cyclodextrins with aqueous iodine solutions. They also observed preferential binding of α-CD and triiodide with 1:1 stoichiometry and KH-G = (1.45 ± 0.02) x 105 M-1 at 20 ºC, along with the enthalpy of binding ∆H = -28.2 ± 0.4 kJ/mol.6 Shortly thereafter, Minns and Khan7 reported in this Journal a spectrophotometric and theoretical study of the α-CD and triiodide aqueous phase host-guest complexation system. Surprisingly, they determined a 2:1 stoichiometry (i.e. α-CD2·I3-) with KH-G = 7.0 x 108 M-2 for both 15 and 25 ºC (i.e. no temperature dependence).7

They also reported the observation of a

very large increase in the UV absorbance of triiodide upon inclusion in the α-CD (cf. 2-3 fold increase).

Using this increase in triiodide absorbance with added α-CD, saturation plots (for

both 15 and 25 ºC) appeared to confirm the 2:1 stoichiometry and the binding constant was calculated based upon this complexation stoichiometry and the experimental concentrations. Their semi-empirical quantum mechanical calculations appeared to further confirm the 2:1 stoichiometry.

However, they did not include the contribution of the aqueous-phase solution, 3



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which is critically important for host-guest systems that involve water and hydrophobic-hydrophilic interactions.8 These results reported by Minns and Khan are rather surprising.

In particular, the

following three findings are rather unexpected: (1) the large increase in UV absorbance of triiodide upon inclusion by α-CD, (2) the 2:1 complexation stoichiometry, and (3) the temperature independence of the complex binding constant between 15 and 25 ºC.7 Furthermore, these authors did not provide an explanation for the supposedly large increase in the molar absorptivity of triiodide when bound by α-CD, nor did they explain the apparent temperature independence of the binding constant.

Thus, the report by Minns and Khan7 in this

Journal was the motivation for the present study. Herein we describe our results from a systematic spectrophotometric study of this interesting and historically important host-guest complexation of α-CD and triiodide in the aqueous phase.

After examining the experimental approach by Minns and Khan7, along with

the methodology presented by French4, we developed a unique spectrophotometric approach that allowed us systematically examine this host-guest systems that involves coupled equilibria. This enabled us (a) to understand the apparently large increase in UV absorbance of triiodide when complexed with α-CD, (b) to establish the 1:1 complexation stoichiometry, and (c) to

4


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measure the KH-G binding constant at room temperature, the temperature dependence of KH-G, and ∆H for complexation.

Experimental Solid iodine, potassium iodide, and α-cyclodextrin were purchased from Sigma Aldrich. All aqueous solutions were prepared using standard analytical techniques with nanopure water and stirred for at least 30 minutes.

Because iodine is only slightly soluble, those solutions were

gently warmed and stirred for 5 hours to one day.

All spectra were collected with a Hitachi

scanning UV-Vis spectrophotometer with 2 nm resolution and nanopure water as the reference. A temperature controlled cuvette holder maintained the solution temperatures to ± 0.2 ºC as measured with a thermocouple suspended in the 1 cm UV quartz cuvette.

Results and Discussion In order to study the host-guest inclusion complexation of α-CD and triiodide in aqueous solution, we first examined the equilibrium associated with the iodine species, namely: + ⇌

(1)

Molecular iodine is only slightly soluble in water and is converted to triiodide by the addition of iodide.

While both I2 and I3- display UV-Vis absorption spectra as shown in Figure 1, their 5



peak molar absorptivities differ by a factor of 34 (note the different concentrations reported in Figure 1).

Calibration spectra, Beer’s Law plots, and a Table with molar absorptivities for

iodine and triiodide are given in the Supporting Information. The equilibrium constant for Reaction 1: = [

[ _ ]

][

(2)

]

was determined by using the initial experimental condition [I2]0 = [I-]0, such that under equilibrium conditions [I2] always equals [I-], thereby allowing us to determine the concentration of all iodine species.

Keq can then be expressed in the linear form: [ ] = [ ]

(3)

where the equilibrium concentrations [I3-] and [I2] were spectrophotometrically measured.

In

this way, we determined Keq = 783 ± 16 M-1 1

at room temperature (see Supporting

0.8

Absorbance

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0.6

-

I3

I2

Information), which agrees with the literature

0.4

(cf. 714-760 M-1).5,9-11 Using this same

0.2 0 250

350

450 550 Wavelength (nm)

approach, the equilibrium constant was

650

Figure 1. Reference UV-Vis spectra for triiodide and iodine in aqueous solutions at room temperature. Solid line: [I3-] = 20 µM (made from 20 µM I2 and 20 mM I-); dashed line: [I2] = 800 µM (with 50 mM HIO3 to force any I- to I2).

determined over the temperature range 15-40 ºC and a linear van’t Hoff plot yielded ∆H = -19.1 ± 0.1 kJ/mol (see Supporting 6


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Information), which agrees with the literature (cf. – ∆H = 22-17 kJ/mol).9,10,12

It is very

important to understand that because of this rather modest equilibrium constant value, one may or may not convert all iodine to triiodide depending on the initial concentrations and the excess iodide. Concerning our experimental methodology and any potential systematic error due to the naturally occurring hydrolysis of iodine in aqueous solution, we note that the hydrolysis reaction is very fast13 but the equilibrium constant is exceedingly small (cf. K = 5.4 x 10-13 M2).14

As

evidenced in the calibration spectra for iodine in Figure S1, there is a very small I3- “impurity” peak at 352 nm due to this interfering reaction.

However, according to concentrations

calculated using the measured absorbance and molar absorptivities, this represents only a 0.6% error in the concentration of iodine.

The hydrolysis of iodine under our experimental

conditions is therefore unimportant.

Furthermore, our experimental methodology is validated

by the linearity of our plots, along with the agreement of our K and ∆H values with the literature values. The host-guest inclusion complexation of α-CD and triiodide, assuming 1:1 stoichiometry (vide infra, to be confirmed through the analysis), is given by the following reaction and equilibrium expression: α + ⇌ α ⋅ 7


(4)


[α⋅ _ ]

= [α][ ]

(5)

This complexation equilibrium is coupled to Reaction 1 through the I3- guest species.

The

binding of I3- by α-CD (according to Reaction 4) causes iodine to be converted into triiodide (according to Reaction 1) and increases the overall concentration of I3- in solution (as “free” I3and “complexed” αCD·I3-).15,16 100x[I2]0. and 2).

This is demonstrated by the spectra in Figure 2 where [I-]0 =

At first glance, the spectrum appears to represent just triiodide (compare Figures 1

However, for these experimental concentrations, only 38% of the iodine has actually

been converted to triiodide by the large excess of iodide (as calculated with our Keq value and in agreement with the absorption peak values).

0.02

0.3

Absorbance

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The addition of α-CD pulls more I3- into

0.01

0.2 0 425

solution as it forms the αCD·I3- complex,

525

0.1

giving rise to a large increase in the absorption 0 250

350


650

Figure 2. Spectra showing the increase in absorbance with the addition α-CD to a solution with triiodide. Solid line: [I3-] = 3 µM (made from 8 µM I2 and 800 µM I-; 38% I2 converted to I3-); dashed line: with the addition of 10 µM α-CD. Inset shows decrease in absorbance due to iodine conversion to triiodide with added α-CD.

spectrum (see dashed spectrum in Figure 2). The inset clearly demonstrates a decrease in absorbance between 450-550 nm due to the loss of iodine with the addition of α-CD.

In an effort to convert all (or nearly all) iodine to triiodide we used [I-]0 = 1000x[I2]0 as shown in Figure 3.

This represents almost 90% conversion of I2 to I3- (according to Reaction 1 8


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and our Keq value). Now the addition of α-CD 0.02

0.3

Absorbance

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causes only a very slight increase in absorbance 0.01

0.2

and a very small red-shift.

0 425

For this

525

0.1

concentration of α-CD, 50% of the I3- is 0 250

350


650

Figure 3. Spectra demonstrating the similarity of the molar absorptivities of “free” I3- and the complex αCD·I3-. Solid line: [I3-] = 7 µM (made from 8 µM I2 and 8 mM I-; 86% I2 converted to I3-); dashed line: with the addition of 10 µM α-CD. Inset shows negligible change in absorbance in the iodine region.

essentially the same.

complexed to α-CD (as calculated according to our KH-G value, vide infra).

These spectra also

demonstrate that the molar absorptivities for “free” I3- and “complexed” αCD·I3- are

It should also be noted that while there is a very large concentration of

iodide in solution, there is an insignificant amount of I- binding by α-CD since that binding constant is so small (ca. KH-G = 19 M-1 for αCD·I- complexation at room temperature).3 To further examine and quantify the host-guest inclusion complexation of α-cyclodextrin and triiodide in aqueous solution, our experimental approach involved the coupling of the two equilibrium reactions (Reactions 1 and 4 above) in a systematic way that allows us to spectrophotometrically evaluate the concentrations of all iodine species (i.e. I2, Iand I3-) and thereby determine the complexation stoichiometry and the KH-G binding constant. We utilized the unique experimental condition with [I2]0 = [I-]0, such that at equilibrium [I2] always equals [I-], which greatly simplifies the analysis.

Additionally, because the

9



concentrations of all iodine species are of

0.4 0.1

isosbestic pt

Increasing Absorbance

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0.3

[αCD]

↑ 0.2

↓

0.05

similar magnitude, only the complexation of α-CD with triiodide needs to be considered

0 385

485

0.1

(since the binding constants for α-CD with 0 250

350


650

Figure 4. Representative spectra at room temperature demonstrating the spectrophotometric approach for determining the complex stoichiometry and binding constant for triiodide and α-CD. [I2]0 = [I-]0 = 50 µM; [α-CD] = 5-50 µM. Inset: the isosbestic point indicates the well-behaved conversion of I2 to I3- with the addition of α-CD.

iodine and with iodide are so much smaller than the binding constant for α-CD with triiodide).

The following relationships (again

assuming 1:1 stoichiometry) then apply for

determining the equilibrium concentrations and the KH-G binding constant from the spectrophotometric determination of the equilibrium concentrations of [I2] and [I3-]: [ ] = [ ] + [α ⋅ _ ]

(6)

Combining Expressions 3 and 6, gives: [α ⋅ _ ] = [ ] − [ ]

(7)

[α] = [α] − [α ⋅ _ ]

(8)

Last,

These relationships allow us to confirm the 1:1 stoichiometry and determine the KH-G binding constant, where [I3-]total and [I2] were spectrophotometrically determined.

Example spectra

demonstrating this experimental approach are shown in Figure 4. Note the isosbestic point (in 10


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the inset) that indicates this is a well-behaved system involving the simple conversion of iodine to triiodide with the addition of α-CD.

The KH-G binding constant at room temperature was

determined as KH-G = (1.35 ± 0.05) x 105 M-1 using a linear form of Equation 5, as demonstrated in Figure 5.

This value agrees with the previous ITC value by Kitamura et al. (cf. KH-G = (1.45

± 0.02) x 105 M-1 at 20 ºC).6

The 1:1 stoichiometry is strongly supported by the linearity of

Figure 5 for the 21 different experimental concentrations utilized.

A similar data treatment and

plot using 2:1 stoichiometry displays large scatter and non-linearity with a correlation coefficient of 0.23 (see Figure S9).

Additionally, the binding constant for 2:1 stoichiometry was

algebraically determined as KH-G = (4 ± 7) x 1010 M-2 (using the 21 experimental conditions), with the standard deviation being almost twice R² = 0.9793

[α αCDI3-] (x 10-5 M)

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3

as large as the average value.

Thus, the 1:1

2

stoichiometry is confirmed.

For

1

completeness, we also report here that in another spectrophotometric study we examined

0 0

0.5

1

1.5

2

[α αCD] [I3-] (x 10-10 M2)

Figure 5. Plot of concentrations for determining the complex binding constant at room temperature according to Equation 5. The linearity for the 21 different concentration combinations confirms the 1:1 stoichiometry. The slope is the binding constant (cf. KH-G = (1.35 ± 0.05) x 105 M-1 at room temperature). The error for each individual measurement is 5%.

the host-guest complexation of molecular iodine and α-cyclodextrin (using the blue-shift that occurs in the I2 spectrum upon binding to α-CD) and determined the binding constant 11



KH-G = (5.93 ± 1.00) x 103 M-1 at room temperature, which agrees with literature values (ca. KH-G = 2-20 x 103 M-1).5,17

Thus, under our experimental conditions, there is no interference due to

the complexation of α-CD with iodine or iodide. Next, the host-guest inclusion complexation of α-cyclodextrin and triiodide in aqueous solution was examined as a function of temperature from 15-40 ºC using the same experimental approach as presented above (and taking into consideration the temperature dependence of Reaction 1, see Supporting Information).

Figure 6 displays one example of the resulting

spectra, while Figure 7 is the van’t Hoff plot of the data for four different experimental conditions.

These results clearly demonstrate that the host-guest complexation of α-CD and

triiodide in aqueous solution is temperature dependent with ∆H = -31.0 ± 0.9 kJ/mol, which agrees with the previous ITC literature value 0.6

Absorbance

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Increasing Temp

0.4

(ca. ∆H = -28.2 ± 0.4 kJ/mol).6

Concerning a

↓

determination of the entropy, since there is

0.2

typically a large error associated with an

0 250

350

450

550

650

Wavelenth (nm)

extrapolation of the intercept from van’t Hoff Figure 6. Spectra demonstrating the temperature dependence of the complexation of triiodide and α-CD. T = 15-40 ºC; [I2]0 = [I-]0 = 50 µM; [α-CD] = 20 µM.

plots, we prefer to reference the literature value of ∆S = -64 J/Kmol.3

12


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We now turn to the results reported by

12.5 R² = 0.9967

Minns and Khan7, namely: (1) the large increase 12.0

ln KH-G

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in UV absorbance of triiodide upon inclusion by

11.5

α-CD, (2) the 2:1 complexation stoichiometry, 11.0 0.00315

0.00325

0.00335 1/T (K)

0.00345

Figure 7. van’t Hoff plot for the temperature dependence of the complexation of triiodide and α-CD. Each point is the average of four different experimental conditions (cf. [I2]0 = [I-]0 = 50 and 100 µM; [α-CD] = 5 and 20 µM). The slope yields the enthalpy of complexation ∆H = -31.0 ± 0.9 kJ/mol. The error for each individual measurement is 5%.

and (3) the temperature independence of the complex binding constant between 15 and 25 ºC.

Initial inspection of Figure 2 above does

appear to suggest only the presence of I3- in solution and the absorbance does appear to

greatly increase with the addition of α-CD, as Minns and Khan reported.7

It is important to

note that the experimental conditions used for the spectra in Figure 2 (cf. [I3-]0 = 8 µm, [I-]0 = 800 µm, yielding 38% conversion of I2 to I3- before the addition of α-CD) are similar to the conditions used by Minns and Khan (cf. [I3-]0 = 60 µm, [I-]0 = 600 µm, yielding 33% conversion of I2 to I3- before the addition of α-CD), such that Reaction 1 has not been forced to completion by a large enough excess of iodide.

Unfortunately, Minns and Khan had assumed that the

excess iodide (i.e. [I-]0 = 10x[I2]0) had effectively converted all iodine to triiodide.

And, as we

have demonstrated above, the large increase in UV absorbance is not due to a change in the molar absorptivity of I3- when complexed with α-CD, but is the result of an increase in the total 13



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Page 14 of 22

amount of I3- in solution as α-CD essentially converts iodine into triiodide (according to Reaction 1) as it forms αCD·I3- (according to Reaction 4).

Thus, these authors did not properly

account for the complications associated with the coupled equilibria.

The misinterpretation of

the spectrophotometric behavior unfortunately affected the determination of the complex stoichiometry, the binding constant KH-G, and the temperature dependence of KH-G as reported by Minns and Khan.7

Furthermore, their semi-empirical quantum mechanical calculations appear

to be incomplete since they did not include the important contributions of the aqueous solution.8

Conclusions We have carefully and systematic examined the host-guest inclusion complexation of α–cyclodextrin and triiodide in aqueous solution using a unique spectrophotometric experimental approach.

Taking into account and utilizing the coupled equilibria, namely, I2 + I-

⇌ I3- and α-CD + I3- ⇌ α-CD·I3-, we have conclusively demonstrated that α–cyclodextrin and triiodide form a 1:1 host-guest complex in aqueous solution with a large binding constant of KH-G = (1.35 ± 0.05) x 105 M-1 at room temperature. with ∆H = -31.0 ± 0.9 kJ/mol.

This binding constant is temperature dependent

These results agree with the pioneering work by French’s

group4, in which they utilized a spectrophotometric approach called the method of continuous variation or Job’s method to determine the 1:1 stoichiometry.

Additionally, our results agree

14


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very well with the isothermal titration calorimetric (ITC) studies by Kitamura et al. (ca. KH-G = (1.45 ± 0.02) x 105 M-1 at 20 ºC and ∆H = -28.2 ± 0.4 kJ/mol).6

The linearity of our plots and

the agreement of our KH-G and ∆H values with the ITC calorimetric values validates our unique spectrophotometric methodology that utilizes the coupled equilibria of this host-guest system. Our study of this interesting and historical host-guest complexation system has regretfully revealed the spectrophotometric misinterpretations by Minns and Khan7.

In

particular, we have conclusively demonstrated that (1) the large increase in the UV absorbance is not due to a change in the molar absorptivity associated with the αCD·I3- species, but is the result of an increase in the overall concentration of triiodide in solution (i.e. as “free” I3- and complexed αCD·I3-) due to the increased conversion of iodine to triiodide with added α-CD; (2) the host-guest complex stoichiometry is 1:1 in aqueous solution; and (3) the equilibrium binding constant is temperature dependent, demonstrating simple van’t Hoff behavior. Host-guest binding of anions involving ionic-hydrophobic interactions in aqueous solutions has been a challenge for researchers and many questions remain unanswered.18

For

example, the reason for the often observed negative entropy for host-guest binding is still not fully understood.19 While exciting advancements have been reported recently,19-24 including efforts to understand hydrophobicity in terms of the chaotropic effect,24 continued work in this important area is warranted. 15



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Finally, while we are unable to perform computational studies of the host-guest inclusion complexation of α–cyclodextrin and triiodide in aqueous solution in order (i) to explain the 1:1 stoichiometry, (ii) to determine the host-guest complex structure, and (iii) to understand the reason for the strong binding, we are hopeful that another research group with those skills will tackle this interesting and historically important problem.

16


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Acknowledgments The authors thank Trinity University, the Dreyfus Foundation, and the Welch Foundation (Department Grant W-0031) for their financial support of this research project.

We

also acknowledge the important contribution of Caitlin McEathron in bringing the research of Minns and Khan to our attention. JLP thanks Dr. Jennifer Holt for helpful assistance and partial supervision.

This research was conducted by JLP under the supervision of CJP at Trinity

University, and the manuscript was written by CJP while on academic sabbatical at the Okinawa Institute of Science and Technology Graduate University (OIST).

CJP therefore thanks

Professor Mukhles Sowwan and OIST for financial support during his academic sabbatical.

Supporting Information Calibration spectra for iodine and triiodide; Beer’s law plots; spectra for the iodine-triiodide equilibrium; equilibrium constant plot; van’t Hoff plot; saturation plot; 2:1 stoichiometry plot; table of molar absorptivities.

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References (1) Crini, G. Review: A History of Cyclodextrins. Chem. Rev. 2014, 114, 10940-10975. (2) Szejtli, J. Introduction and General Overview of Cyclodextrin Chemistry. Chem. Rev. 1998, 98, 1743-1753. (3) Rekharsky, M. V.; Inoue, Y. Complexation Thermodynamics of Cyclodextrins. Chem. Rev. 1998, 98, 1875-1917. (4) Thoma, J. A.; French, D. Studies on the Schardinger Dextrins. X. The Interaction of Cyclohexaamylose, Iodine and Iodide. Part I. Spectrophotometric Studies. J. Am. Chem. Soc. 1958, 80, 6142-6146. (5) Diard, J. P.; Saint-Aman, E.; Serve, D. Potentiometric Association Constant Measurements of α, β or γ-Cyclodextrin Complexes Involving Iodide, Tri-iodide or Iodine Species. J. Electroanal. Chem. 1985, 189, 113-120. (6) Kitamura, S.; Nakatani, K.; Takaha, T.; Okada, S. Complex Formation of Large-ring Cyclodextrins with Iodine in Aqueous Solution as Revealed by Isothermal Titration Calorimetry. Macromol. Rapid Commun. 1999, 20, 612-615. (7) Minns, J. W.; Khan, A. α-Cyclodextrin-I3- Host-Guest Complex in Aqueous Solution: Theoretical and Experimental Studies. J. Phys. Chem. A 2002, 106, 6421-6425.

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(8) Liu, L.; Guo, Q-X. Use of Quantum Chemical Methods to Study Cyclodextrin Chemistry. J. Incl. Phenom. Macrocycl. Chem. 2004, 50, 95-103. (9) Awtrey, A. D.; Connick, R. E. The Absorption Spectra of I2, I3-, I-, IO3-, S4O6= and S2O3=. Heat of the Reaction I3- = I2 + I-. J. Am. Chem. Soc. 1951, 73, 1842-1843. (10) Ramette, R. W.; Sandford, Jr., R. W. Thermodynamics of Iodine Solubility and Triiodide Ion Formation in Water and in Deuterium Oxide. J. Am. Chem. Soc. 1965, 87, 5001-5005. (11) Ramette, R. W. Equilibrium Constants from Spectrophotometric Data. J. Chem. Educ. 1967, 44, 647-654. (12) Jones, G.; Kaplan, B.B. The Iodide, Iodine, Tri-iodide Equilibrium and the Free Energy of Formation of Silver Iodide. J. Am. Chem. Soc. 1928, 50, 1845-1864. (13)

Lengyel, I.; Epstein, I. R.; Kustin, K. Kinetics of Iodine Hydrolysis. Inorg. Chem. 1993,

32, 5880-5882. (14) Allen, T. L.; Keefer, R. M. The Formation of Hypoiodous Acid and Hydrated Iodine Cation by the Hydrolysis of Iodine. J. Am. Chem. Soc. 1955, 77, 2957-2960. (15) Kitamaki, Y.; Jin, J-Y.; Takeuchi, T. Determination of Inorganic Anions via Postcolumn Reaction with Iodide in Ion Chromatography. J. Pharm. Biomed. Anal. 2003, 30, 1751-1757.

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(16) Hiroi, T.; Jin, J.; Takeuchi, T. Photometric Detection of Cyclodextrins in Liquid Chromatography by using Iodine Generated Electrochemically In-Situ. Anal. Bioanal. Chem. 2005, 381, 1089-1094. (17) Thoma, J. A.; French, D. The Dissociation Constant for the Cyclohexaamylose – Iodine Complex. J. Phys. Chem. 1958, 62, 1603. (18) Sessler, J. L.; Gale, P. A.; Cho, W.-S. Anion Receptor Chemistry; Royal Society of Chemistry: Cambridge, 2006. (19) Yawer, M. A.; Havel, V.; Sindelar, V. A Bambusuril Macrocycle that Binds Anions in Water with High Affinity and Selectivity. Angew. Chem. Int. Ed. 2015, 54, 276-279. (20) Fox, J. M.; Kang, K.; Sherman, W.; Heroux, A.; Sastry, G. M.; Baghbanzadeh, M.; Lockett, M. R.; Whitesides, G. M. Interactions between Hofmeister Anions and the Binding Pocket of a Protein. J. Am. Chem. Soc. 2015, 137, 3859-3866. (21) Wu, Y.; Shi, R.; Wu, Y.-L.; Holcroft, J. M.; Liu, Z.; Frasconi, M.; Wasielewski, M. R.; Li, H.; Soddart, J. F. Complexation of Polyoxometalates with Cyclodextrins. J. Am. Chem. Soc. 2015, 137, 4111-4118. (22) Sokkalingam, P.; Shraberg, J.; Rick, S. W.; Gibb, B. C. Binding Hydrated Anions with Hydrophobic Pockets. J. Am. Chem. Soc. 2016, 138, 48-51.

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(23)

Lisbjerg, M.; Nielsen, B. E.; Milhoj, B. O.; Sauer, S. P. A.; Pittelkow, M. Anion Binding

by Biotin[6]uril in Water. Org. Biomol. Chem. 2015, 13, 369-373. (24) Assaf, K. I.; Ural, M. S.; Pan, F.; Georgiev, T.; Simova, S.; Rissanen, K.; Gabel, D.; Nau, W. M. Water Structure Recovery in Chaotropic Anion Recognition: High-Affinity Binding of Dodecaborate Clusters to γ-Cyclodextrin. Angew. Chem. Int. Ed. 2015, 54, 6852-6856.

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TOC Graphic

0.4 0.1

isosbestic pt

Increasing Absorbance

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

0.3

[αCD]

↑ 0.2

↓

0.05

0 385

485

0.1

0 250

350


650

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Page 22 of 22

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