Solubility of Hydroxyl Cucurbit [6] uril in Different Binary Solvents

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Solubility of Hydroxyl Cucurbit[6]uril in Different Binary Solvents Lu Zhou, Changjun Zou,* Meng Wang, and Lu Li School of Chemistry and Chemical Engineering, Southwest Petroleum University, Chengdu, China ABSTRACT: The solubility of hydroxyl cucurbit[6]uril in pure water and another five binary solvents, including formic acid aqueous solution, acetic acid aqueous solution, hydrochloric acid aqueous solution, ammonia, and potassium hydroxide aqueous solution, was respectively measured in the temperature range from (273.15 to 353.15) K by the equilibrium method. The results indicate that the solubility of hydroxyl cucurbit[6]uril in all six solvents increases with increasing temperature. The ideal solution equation, Apelblat equation, and Yaws equation were used to correlate the experimental solubility data of hydroxyl cucurbit[6]uril and all exhibit good agreement. The dissolution enthalpy, entropy, and the change of Gibbs free energy of hydroxyl cucurbit[6]ril in these solvents were obtained from the experimental solubility data by the van’t Hoff equation and Gibbs−Helmholtz equation. cucurbit[n]uril modifications.26−28 For improving CB[6] water solubility, it has been reported that hydroxyl cucurbit[6]uril has successfully been synthesized by direct functionalization of CB[6] in the condition of strong oxidant.29 Nevertheless, there is no report about the solubility data of hydroxyl CB[6], which is very important in the application and industrialization of hydroxyl CB[6]. In this work, the solubility of hydroxyl CB[6] in pure water and another five binary solvents is measured by equilibrium methods in the temperature range from (273.15 to 353.15) K and correlated by three models, the ideal solution equation, Apelblat equation, and Yaws equation.

1. INTRODUCTION Cucurbit[6]uril (CB[6]) is a novel host macrocyclic compound followed after crown ether, calixarene, and cyclodextrin in the fields of supramolecule, synthesized by a condensation reaction of glycoluril and formaldehyde using acid as catalyst.1−3 Although its synthesis was first reported in 1905,4 the chemical structure of CB[6] was not known until 1981 when it was clearly characterized by Mock.5 The structural formula is given in Figure 1.6,7 On account of their hydrophilic outer wall,

2. EXPERIMENTAL SECTION 2.1. Materials. Hydroxyl cucurbituril (mass fraction > 99.5 %) used in the experiment was prepared and purified according to the literature.2,7,29,30 The other materials including hydrochloric acid, formic acid, acetic acid, ammonia, and potassium hydroxide were all provided by Kelong Chemical Reagent Factory (Chengdu, China) with analytical purity grades, which were utilized to prepare binary solvents with pure water. Their physical properties are given in Table 1, and the analytical balance used to prepare the solvents had an accuracy of 0.0001 g. The mass fractions for hydrochloric acid, formic acid, acetic acid, ammonia, and potassium hydroxide in the binary solvents were 100 w = (3.646, 4.603, 6.005, 3.504, and 5.611), respectively. 2.2. Apparatus and Procedure. The solubility of hydroxyl CB[6] in six kinds of solvents at different temperature was measured by the equilibrium method.31

Figure 1. Structural formula of cucurbit[6]uril.

oleophilic inner cavity, and strong rigid structure, all cucurbit[n]uril exhibit board application prospects in the fields of molecular recognition,8−10 self-assembly,11−13 catalysis,14−16 as well as biomedicine.17−19 In recent years, various mechanically interlocked molecules were synthesized such as rotaxanes,20 polyrotaxanes,21 molecular necklaces,22 rotaxane dendrimers,23 and rotaxane-based molecular switches24 using CB[6] as a molecular bead. However, CB[6] is clearly undissolved in common solvents except for strongly acid aqueous solutions, which severely limits the development of cucurbit[n]uril chemistry.25 Therefore, the supramolecular chemistry field has recently seen a wave of © 2014 American Chemical Society

Received: June 4, 2014 Accepted: August 21, 2014 Published: August 28, 2014 2879

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Table 1. Physical Properties of the Solvents Used in Experiments solvent

molecular formula

density, kg·m−3

purity, wt %

formic acid acetic acid hydrochloric acid ammonia potassium hydroxide

HCOOH CH3COOH HCl NH3 NaOH

1230 1050 1190 910 2044

88.0 99.5 36.0 25.0 82.0

A series of hydroxyl CB[6] aqueous solutions with known concentration were prepared, and their absorbance was determined at room temperature by 3802 UV−vis spectrophotometer (accuracy, 0.3% for absorbance) at the maximum absorption wavelength 300 nm, using pure water as a reference. Thus, a standard curve correlating absorbance and concentration was obtained.32 A glass bottle (100 mL) with a Teflon-coated magnetic stirrer and a condenser was utilized to prepare the saturated solution of hydroxyl CB[6].33 The solvents (20 mL) and excess solid solute were added into the glass bottle, and agitated at the speed of 300 rpm. At the interval of 0.5 h stirring and 0.5 h standing, the liquid was taken out from the bottle to be analyzed via UV−vis spectrophotometer at 300 nm until the absorbance of the liquid was invariant, thus indicating the dissolution equilibrium. It took 3 h to reach the equilibrium. The saturated solution of hydroxyl CB[6] in six solvents at the temperature from 273.15 K to 353.15 K was obtained according to the methods above with 3 h stirring. The temperature was controlled by thermostatic water bath (with uncertainty of ± 0.01 K). After 0.5 h standing, 0.1 mL (accuracy, ± 0.003 mL) of supernatant was sampled by Eppendorf pipet and distilled to 25 mL (accuracy, ± 0.01 mL) by volumetric flask with pure water against recrystallization, and then the absorbance of the diluted solutions was measured on a UV−vis spectrophotometer at room temperature at 300 nm. All experiments were carried out three times.

Figure 2. Standard curve of hydroxyl CB[6] aqueous solution.

curve also fits for the diluted solutions with other five solvents to calculate concentrations according to their absorbance, and then the concentrations of their saturated solutions, namely solubility, are obtained. The solubility of hydroxyl CB[6] in six solvents at different temperatures is calculated and plotted in Figure 3.

3. RESULTS AND DISCUSSION 3.1. Solubility of Hydroxyl CB[6] in Solvents. According to Bouguer−Lambert−Beer Law, a linear relationship exists between absorbance and concentration of solution; that is to say, in theory, the concentration of a solution can be obtained via a linear relationship when the absorbance is known. Therefore, the standard curve, correlating absorbance at 300 nm with the concentration for hydroxyl CB[6] aqueous solution, is obtained to calculate the solubility of hydroxyl CB[6] which is plotted in Figure 2. Its correlation equation is (1) A = 0.7135 + 1.1929C where C is the concentration of hydroxyl CB[6] solution in g/ g, A stands for the corresponding absorbance. The correlation coefficient (R2) of the equation is 0.9991, indicating good linearity. The measurement wavelength (300 nm) for the standard curve meets two conditions: the maximum absorption wavelength of hydroxyl CB[6] aqueous solution is at 300 nm; there is no absorption at 300 nm for the six solvents. Furthermore, the solutions of hydroxyl CB[6] in six solvents, used for UV− vis analysis, were diluted with pure water, which weakens and even eliminates the interactions between hydroxyl CB[6] and potassium hydroxide, acetic acid, ammonium, formic acid, or hydroxide acid. On the basis of these reasons, the standard

Figure 3. Solubility of hydroxyl CB[6] at different temperatures in the solvents of hydroxide acid aqueous solution (◆), pure water (■), formic acid aqueous solution (●), ammonium aqueous solution (▼), acetic acid aqueous solution (▲), and potassium hydroxide aqueous solution (★).

It can be seen from Figure 3 that hydroxyl CB[6] exhibits good solubility and the solubility of hydroxyl CB[6] in solvents increases with an increase in temperature, illustrating that its dissolution process is endothermic, which agrees well with the experiment phenomenon.34 Meanwhile, Figure 3 has also revealed that the solubility of hydroxyl CB[6] in six solvents is in the order of potassium hydroxide aqueous solution < acetic acid aqueous solution < ammonium < formic acid aqueous solution < hydroxide acid aqueous solution. The solubility of hydroxyl CB[6] in pure water is higher than that in acetic acid aqueous solution (6.005 wt %) and lower than that in ammonium (3.504 wt %) when the temperature is below 320 K, while the solubility in pure water is between that of hydroxide acid aqueous solution and that of formic acid aqueous solution if the temperature is higher than 320 K. 2880

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Table 2. Parameters (A, B, C), Root-Mean-Square Relative Deviation (RMSD) and Correlation Coefficient (R2) of the Ideal Solution Model, Apelblat Model, and Yaws Model Correlated According to Experimental Mole Fraction Solubility of Hydroxyl CB[6] in Different Solvents solvent pure water formic acid solution (100 w = 4.603) acetic acid solution (100 w = 6.005) hydroxide acid solution (100 w = 3.646) ammonium (100 w = 3.504) potassium hydroxide solution (100 w = 5.611) pure water formic acid solution (100 w = 4.603) acetic acid solution (100 w = 6.005) hydroxide acid solution (100 w = 3.646) ammonium (100 w = 3.504) potassium hydroxide solution (100 w = 5.611) pure water formic acid solution (100 w = 4.603) acetic acid solution (100 w = 6.005) hydroxide acid solution (100 w = 3.646) ammonium (100 w = 3.504) potassium hydroxide solution (100 w = 5.611)

A

A1 + B1 T /K

B2 ln x = A 2 + + C2 ln(T /K) T /K

lg x = A3 +

B3 C3 + T /K (T /K)2

RD% =

m1/M1 m1/M1 + m2 /M 2 + m3 /M3

3.1084 14.5330 7.8061 5.5920 6.3545 1.8550

0.9990 0.9676 0.9676 0.9563 0.9884 0.9884

14.77 14.41 14.31 14.35 14.32 14.27

9.9767 5.4580 1.6832 5.2994 3.8275 1.9444

0.9940 0.9935 0.9982 0.9955 0.9953 0.0040

44645.68 13498.45 381750.87 355491.55 218640.42 150775.08

4.0602 5.4580 1.1691 4.9160 2.0248 0.7412

0.9989 0.9949 0.9992 0.9962 0.9987 0.9994

xexp, i − xcal, i xexp, i

100

2 ⎤1/2 ⎡ N (x exp, i − xcal, i) ⎥ RMSD = ⎢∑ ⎢⎣ i = 1 ⎥⎦ N−1

(2)

(6)

(7)

where N represents the number of data points, xexp,i and xcal,i are the experimental and calculated values of solubility. From Table 2 and Table 3, it can be seen that RMSD and RD is small, indicating all equations including the ideal solution equation, Apelblat equation, and Yaws equation are appropriate to build the solubility mathematic model in our research. Furthermore, compared with the ideal solution equation and Apelblat equation, Yaws equation exhibited better agreement due to its R2 maximal value. 3.3. Thermodynamic Functions.37,38 The dissolution enthalpy (ΔHd) and dissolution entropy (ΔSd) can be calculated via the van’t Hoff equation:

(3)

(4)

where A1, A2, A3, B1, B2, B3, C2, and C3 are the regression parameters of the equations, T stands for temperature in K, and x is the mole fraction of solute. The x could be calculated via eq 5: x=

R2

C

Ideal Solution Model −1876.92 0.2994 −932.67 −3.2027 −894.23 −3.2027 −734.97 −2.8348 −888.83 −2.8805 −1096.80 −3.1954 Apelblat Model −100.0000 2948.46 −100.3634 3736.30 −100.2225 3738.29 −100.1136 3904.34 −100.0000 3748.14 −100.0000 3532.19 Yaws Model 0.5482 −1088.87 2.8167 −2964.36 2.2526 −2751.58 2.1710 −2522.88 0.8354 −1739.36 0.0456 −1407.83

3.2. Solubility Modeling. The effect of temperature on the solubility of hydroxyl CB[6] in different solvents has been correlated by three models, the ideal solution model (eq 2), the Apelblat model (eq 3), and the Yaws model (eq 4).35,36 ln x =

105 RMSD

B

ln xexp = −

(5)

ΔHd ΔSd + RT R

(8)

where xexp is the mole fraction of solute, R stands for the ideal gas constant, 8.314510 J·mol·K, T represents temperature in K. It can be seen from eq 8 that the natural logarithm of solubility in the mole fraction remains in a linear relationship with the reciprocal of temperature. According to the experimental data, the van’t Hoff plots correlate ln xexp and 1/T. These are drawn and shown in Figure 4. The change of Gibbs free energy (ΔGd) can be obtained by the Gibbs−Helmholtz equation:

where m1 and m2 stand for the mass of hydroxyl CB[6] and pure water respectively; m3 is the mass of other solvent component such as potassium hydroxide, acetic acid, ammonium, formic acid, and hydroxide acid aqueous solution; M1 and M2 represent the molecular weight of hydroxyl CB[6] and pure water respectively; M3 is the molecular weight of other solvent component. The parameters A1, A2, A3, B1, B2, B3, C2, and C3 were calculated by least-squares fit from the experimental data. The experimental and calculated values of solubility are listed in Table 2 with relative deviation (RD), and the parameters of solubility model are listed in Table 3 with root-mean-square relative deviation (RMSD). RD and RMSD were respectively defined as

ΔGd = ΔHd − T ΔSd

(9)

where T is set as 293.15 K. The calculated value of dissolution enthalpy, dissolution entropy, and the change of Gibbs energy are listed in Table 4. It 2881

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Table 3. Experimental and Calculated Mole Fraction Solubility (xexp and xcal) of Hydroxyl CB[6] in Six Solvents at Different Temperatures (T) from (293.15 to 353.15) K under Pressure of 0.1 MPa (P)a T K

a

Ideal Solution model 100xexp

293.15 303.15 313.15 323.15 333.15 343.15 353.15

0.2307 0.2704 0.3358 0.4032 0.4812 0.5726 0.6622

293.15 303.15 313.15 323.15 333.15 343.15 353.15

0.3260 0.3444 0.3779 0.3987 0.4375 0.4852 0.5514

293.15 303.15 313.15 323.15 333.15 343.15 353.15

0.2036 0.2115 0.2298 0.2475 0.2701 0.2991 0.3332

293.15 303.15 313.15 323.15 333.15 343.15 353.15

0.5044 0.5214 0.5477 0.5807 0.6423 0.6848 0.7538

293.15 303.15 313.15 323.15 333.15 343.15 353.15

0.2812 0.2973 0.3250 0.3521 0.3835 0.4233 0.4586

293.15 303.15 313.15 323.15 333.15 343.15 353.15

0.0009920 0.001110 0.001222 0.001345 0.001511 0.001681 0.001851

100xcal

Apelblat model

RD %

100xcal

Yaws model RD %

Pure Water 0.2236 3.08 0.2384 −3.34 0.2761 −2.11 0.2808 −3.85 0.3365 −0.21 0.3324 1.01 0.4051 −0.47 0.3952 1.98 0.4823 −0.23 0.4713 6.00 0.5683 0.75 0.5637 1.55 0.6635 −0.20 0.6755 −2.01 Formic Acid Aqueous Solution (100 w = 4.603) 0.3089 5.25 0.3240 0.61 0.3431 0.38 0.3452 −0.23 0.3785 −0.16 0.3717 1.64 0.4151 −4.11 0.4043 −1.40 0.4526 −3.45 0.4433 −1.33 0.4911 −1.22 0.4897 −0.93 0.5304 3.81 0.5443 1.29 Acetic Acid Aqueous Solution (100 w = 6.005) 0.1924 5.50 0.2016 0.98 0.2128 −0.61 0.2139 −1.13 0.2339 −1.78 0.2296 0.09 0.2554 −3.19 0.2487 −0.48 0.2776 −2.78 0.2718 −0.63 0.3001 −0.33 0.2992 −0.03 0.3231 3.03 0.3315 0.51 Hydroxide Acid Aqueous Solution (100 w = 3.646) 0.4790 5.04 0.5003 0.81 0.5199 0.29 0.5217 −0.06 0.5618 −2.57 0.5509 −0.58 0.6041 −4.03 0.5879 −1.25 0.6468 −0.70 0.6335 1.37 0.6897 −0.72 0.6881 −0.48 0.7329 2.77 0.7529 0.12 Ammonium (100 w = 3.504) 0.2705 3.81 0.2833 −0.75 0.2990 −0.57 0.3004 −1.04 0.3284 −1.05 0.3222 0.86 0.3585 −1.82 0.3489 0.91 0.3893 −1.51 0.3811 0.63 0.4208 0.59 0.4194 0.92 0.4529 1.24 0.4645 −1.29 Potassium Hydroxide Aqueous Solution (100 w = 5.611) 0.0009710 2.15 0.001020 −2.78 0.001099 0.99 0.001106 0.35 0.001234 −0.98 0.001212 0.86 0.001375 −2.23 0.001338 0.50 0.001522 −0.73 0.001489 1.45 0.001675 0.36 0.001667 0.81 0.001834 0.92 0.001877 −1.41

100xcal

RD %

0.2256 0.2768 0.3360 0.4038 0.4809 0.5678 0.6651

2.21 −2.37 −0.06 −0.15 0.06 0.84 −0.44

0.3283 0.3448 0.3691 0.4010 0.4409 0.4894 0.5473

−0.71 −0.12 2.33 −0.58 −0.78 −0.87 0.74

0.2035 0.2138 0.2284 0.2474 0.2709 0.2993 0.3327

0.05 −1.09 0.61 0.04 −0.30 −0.07 0.15

0.5023 0.5212 0.5491 0.5863 0.6326 0.6884 0.7542

0.42 0.04 −0.26 −0.97 1.51 −0.53 −0.05

0.2794 0.2998 0.3240 0.3521 0.3840 0.4201 0.4605

0.64 −0.84 0.31 0.00 −0.13 0.76 −0.41

0.0009950 0.001102 0.001223 0.001358 0.001507 0.001673 0.001854

−0.27 0.72 −0.08 −0.97 0.26 0.48 −0.16

Standard uncertainties u are u(T) = 0.01 K, u(P) = 10 kPa, ur(x) = 0.02, respectively.

can be inferred from ΔHd > 0 that the dissolution is an endothermal process, which is consistent with the experimental phenomenon. The result illustrates that the absorbed heat due to the solid solute overcoming the attraction between the molecule and ions is far greater than the released heat due to the solvation between solute molecule and solvent molecule. In addition, the change tendency of dissolution enthalpy in five kinds of solvents is almost opposite to the change of solubility

except for pure water. Similarly, the change trend of Gibbs energy is contrary to the variation tendency of solubility. Generally speaking, the dissolution process of most solutes is an entropy-driven process resulting from the increase of system disorder degree. The dissolution entropy of hydroxyl CB[6] in pure water is a positive value, while the dissolution entropy in the other five solvents is a negative value. Compared with the positive value of the dissolution entropy in the hydroxyl 2882

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AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected]. Tel.: +86 02883037327. Fax: +86 02883037305. Funding

This work was financially supported by National Natural Science Foundation of China, China National Petroleum Corporation Petrochemical Unite Funded Project (U1262111). Notes

The authors declare no competing financial interest.

■

Figure 4. Van’t Hoff plots of ln x versus 1/T for hydroxyl CB[6] in the solvents of hydroxide acid aqueous solution (◆), formic acid aqueous solution (●), ammonium aqueous solution (▼), pure water (■), acetic acid aqueous solution (▲), and potassium hydroxide aqueous solution (★).

Table 4. Thermodynamic Function Values of Hydroxyl CB[6] (293.15 K) ΔHd solvent pure water formic acid solution (1 mol/L) acetic acid solution (1 mol/L) hydroxide acid solution (1 mol/L) ammonium (1 mol/L) potassium hydroxide solution (1 mol/L)

−1

ΔSd −1

ΔGd −1

(J·mol )

(J·mol ·K )

(J·mol−1)

15771.0 7254.2 7434.6 6110.5 7389.7 9118.8

2.49 −21.60 −26.63 −23.57 −23.95 −26.57

15041.1 13586.2 15240.6 13019.6 14410.6 16913.6

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CB[6]/pure water system, the negative value can be explained from two aspects. On the one hand, hydroxyl CB[6] forms inclusion complexes with potassium hydroxide, acetic acid, ammonium, formic acid, or hydroxide acid aqueous solution via one of weak noncovalent interactions, namely, the strong ironcoordinated interaction, which leads to a reduction of disorder degree to some extent. On the other hand, the activity of water molecules is limited by the salvation of potassium hydroxide, acetic acid, ammonium, formic acid, and hydroxide acid, ultimately making the disorder degree decrease further.39

4. CONCLUSION The solubility of hydroxyl CB[6] in pure water, potassium hydroxide aqueous solution, acetic acid aqueous solution, ammonium, formic acid aqueous solution, and hydroxide acid aqueous solution was determined at different temperatures by the equilibrium methods. Hydroxyl CB[6] exhibited good solubility performance in all solvents used in the experiment, and its solubility increased with an increasing temperature. Three models, the ideal solution model, Apelblat model, and Yaws model are utilized to correlate and calculate experimental data in different solvents. The calculated values exhibited good agreement with experimental values. The measured solubility and mathematic model possess potential application value in the industrialization of cucurbituril and its derivatives. 2883

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dx.doi.org/10.1021/je5005033 | J. Chem. Eng. Data 2014, 59, 2879−2884