J. Phys. Chem. 1989, 93, 3836-3838
3836
These are the only evidences in our data of the existence of this transition. Phases I I I and III’. Thermally activated molecular reorientation occurs in these phases. In phase 111 above -1 15 OC motional averaging increases T2. Thus, the reorientation time 7r is about I O gs at -1 15 O C . The temperature dependence of T I indicates that, molecular rotations (at a rate much smaller than the resonance frequency) control T I . Conclusions NMR evidence is presented that cyclooctane rotor phase 1/11 supercools to -1 40 O C , transforming reversibly to nonrotor phase 111’. 111’ transforms on heating through -140 OC back to 1/11 which rapidly and irreversibly yields the nonrotor phase 111. 111 is shown to be the thermodynamically stable state below -104 OC; 111’ is always metastable with respect to 111. This picture is thoroughly confirmed by thermal data. In particular, the exo-
-
thermic irreversible transition 1/11 111 is evident. The DTA study also clearly shows the weak 11-1 transition near -89 “C. None of the ‘H NMR measurements can distinguish phases I and 11. The diffusion jump rate T ~ - ]is determined over a wide temperature range from both T2in the line narrowing region and T I D in the slow-motion regime. The activation energy is 4350 K (8.6 kcal/mol), &lo%. The observation of a Pake doublet in the 2H NMR is conclusive evidence that some molecules in phase I rotate less completely than spheres. Acknowledgment. We appreciate helpful discussions with J. R. Brookeman, F. A. L. Anet, and H. Strauss. The work was supported in part by NSF DMR 87-02847. The work was also supported by the generosity of the donors to the Petroleum Research Fund, administered by the American Chemical Society. Registry No. Cyclooctane, 292-64-8.
On the Question of Hypoiodlte Ion Formation in the Aqueous Solution of Iodine: Theoretical and Experimental Study of H2012Complex Jane6 Fonslick, Arshad Khan,* Chemistry Department, The Pennsylvania State University, Dubois, Pennsylvania 15801
and Brian Weiner Physics Department, The Pennsylvania State University, Dubois. Pennsylvania 15801 (Received: April 1I , 1988; In Final Form: September 19, 1988)
The possibility of hypoiodite ion formation by the hydrolysis of iodine has been examined by experimental as well as theoretical studies. The aqueous solution of iodine shows absorbance maxima near 205 nm (most intense band), 285 nm (weaker band), and 355 nm (weakest). The band at 355 nm had previously been attributed to 01- ions. INDO CI calculations on 01- do not give an absorbance band near any of the above wavelengths. Also, our experimental observations suggest that 01- ions do not form in the aqueous solution. The theoretical calculations predict a H2012complex in which two iodine atoms are arranged along the C, axis of H20and which has absorbance bands at 181 nm (most intense), 277 nm (weaker), and 365 nm (weakest). The shifts in wavelengths of maximum absorbance are explained in terms of solvation of the highly polar H2012 complex.
Introduction Starch-iodine solution in water shows a weak absorption band at around 355 nm, which has been believed to be due to the f ~ r m a t i o n l -of~ hypoiodite (01-) ions as a result of hydrolysis reaction of I2 (eq I ) . Mokhnach and Rusakova,j in their experiments, found that the addition of starch to an aqueous solution of iodine resulted in an increase in the absorbance at around 355 nm. They suggested that the increase in absorbance was due to the fact that 01- ions were formed more readily when starch was present and thus assigned the absorbance peak to 01- ions. I t has been found that the starch-iodine blue complex (absorbance band at 615 nm) does not obey the Beer-Lambert law and exhibits a positive i n t e r ~ e p t ~or- ~“threshold” on the concentration axis of the calibration plot (plot of absorbance due to starch-iodine complex vs I2 concentration). Drey] explained the starch-iodine threshold by considering formation of a minimum concentration of hypoiodite ions before
Even though this explanation of the threshold has been questioned by Hatch,6 the possibility of the hydrolysis reaction still existed as no other experimental or theoretical evidence exists to refute it. Ziegast2 et al., in explaining the starch-iodine blue complex formation in the aqueous solution (in absence of KI), suggested a slow conversion of Iz to I- ions by hydrolysis (eq 1) that then form I,- ions (eq 2) before being incorporated into the amylose helix of starch (eq 3). In order to examine whether 01- and Iions are formed by the hydrolysis of iodine (eq l ) , the present experimental and theoretical studies have been made.
Drey, R.E. A. Anal. Chem. 1964, 36, 2200. Ziegast, G.; Pfannemuller, B. Ini. J . Biol. Mucromol. 1982, 4, 419. (a) Mokhnach, V . 0.; Rusakova, N. M. Dokl. Akad. Nauk SSSR 1962, 143, 122. (b) Ibid. 1962, 145, 1290. (4) Lambert, J. L. Anal. Chem. 1951, 23, 1251. (5) Crouch, W. H . Anal. Chem. 1962, 34, 1698. (6) Hatch, G.L. Anal. Chem. 1982, 54, 2002.
Experimental Section In our experiments with iodine solution in water, N/10 iodine solution containing N/5 potassium iodide (Fisher Scientific) was diluted to an iodine concentration of 7.93 mg/L. To prevent evaporation of iodine and, hence, to obtain a stable absorbance reading, the iodine solution was kept at around 6 OC throughout
0022-3654/89/2093-3836$01.50/0
the reaction of iodine with starch can take place in the following manner:
I2 + H 2 0 = 01I2
I,-
+ 2H+ + I-
+ I- = 13-
+ starch = starch-iodine
0 1989 American Chemical Society
(1)
(2) complex
(3)
Theoretical and Experimental Study of H2012Complex the experiment. A Spectronic 601 spectrophotometer was used for the absorbance measurements in the wavelength range of 195-740 nm. The spectrum due to iodine in water was obtained by subtracting the spectrum due to KI solution from the spectrum due to iodine with KI solution. This shows three absorbance maxima. The most intense one is at around 205 nm, the second and relatively less intense peak at around 285 nm, and the weakest one at around 355 nm. The intensity ratio of these peaks was 8:2:1. In addition to the above experiments, spectra of starch-iodine and iodine in several solvents such as ethyl alcohol, 1-propanol, acetone, benzene, toluene, and xylenes were studied. Iodine solutions in nonaqueous solvents were made by dissolving resublimed iodine crystals (Fisher) in those solvents. All chemicals used were in their purest available forms. Different alcohols and acetone (oxygen-containing solvents) used in these tests gave an absorbance band at around 355 nm. Even though spectra due to iodine in most of the above-mentioned solvents could not be obtained in the UV range as those solvents absorb very strongly in the UV range, spectra due to iodine in ethyl alcohol and 1-propanol were obtained in the UV as well as in the visible range of Wavelengths. These spectra show remarkable similarities with that of the water solution. Three absorbance maxima were obtained at around 205, 285, and 355 nm. Even though the intensity of the second and the third peaks maintained a ratio of 2:1, the relative intensity of the first peak decreased significantly compared to that of water. Any trace amount of water as impurity in the nonaqueous solvents or in iodine crystal is not expected to give such well-defined peaks at these wavelengths. In the absence of KI, solubility of iodine in water is fairly small and hence the absorbance due to a trace amount of water (impurity) can be expected to be negligibly small. It should be mentioned that iodine in non-oxygen-containing solvents such as benzene, toluene, and xylenes gave spectra quite different from those of the oxygen-containing solvents. Theoretical studies were done to further substantiate these experimental findings about the role of oxygen in these spectra.
Theoretical Calculations and Discussion In order to examine the possibility of whether 01- ions are formed, INDO C I calculations were carried out on 01- with 43 single excited configurations. The atomic basis sets used in these calculations were obtained by the linear combination of Slater-type orbitals with different exponents. Two such functions were used for each of the valence s and p orbitals on oxygen and d orbitals on iodine. On the other hand, three such Slater-type functions were used to represent each of the valence s and p orbitals on iodine. Each of the core orbitals, however, was represented by a linear combination of two Slater-type orbitals. Detailed discussions on these basis functions can be found in several publications of Zerner and co-worker~.’~These calculations provided two absorbance maxima, the most intense band at 134 nm and a fairly weak band at 158 nm. It is, however, expected that 01ions in the aqueous solution will be solvated, and hence, spectral lines will be shifted to some extent. Because of the large difference between our experimental spectra and the theoretical spectra for 01- ion, it is unlikely that the 355-nm band is due to 01- ions. Several experimental results provide additional evidence against the existence of 01- ions in a neutral aqueous solution of iodine. For example, iodine in a highly alkaline solution also gives a weak absorbance at around 365 nm. This band had been attributed to 01- ions resulting from the following reaction:I2 I2
+ 2 0 H - = 01- + I- + H20
On the basis of a very low acid dissociation constant of HOI (2 (7) Bacon, A. D.; Zerner, M. C. Theor. Chim. Acta 1979, 53, 21. (8) Zerner, M. C.; Loew, G. H.; Kirchner, R. F.; Muller-Westerhoff, U. T. J . Am. Chem. SOC.1980, 102, 589. (9) Culberson, J.; Knappe, P.; Rosch, N.; Zerner, M. C. Theor. Chim. Acta 1987, 71, 21. (10). Zerner, M. C., Quantum Theory Project, University of Florida, Gainsville, FL 3261 1, private communication. (11) Blandamer, M. J.; Fox, M. F . Chem. Rev. 1970, 70, 59-93. (12) Haimovich, 0.;Trenin, A. Narure 1965, 207, 185.
The Journal of Physical Chemistry, Vol. 93, No. 9, 1989 3837
2.19 A‘
Figure 1. Perpendicular complex X lo-”) and a fairly weak absorbance band (at around 365 nm) in a strong alkaline solution, one would expect only an insignificant amount of 01- ion formation in a neutral solution to give any measurable absorbance value. In addition, 01- ions, if any exist at all as an impurity in a neutral aqueous solution of iodine, will be mostly in HOI form with only 0.02% in 01- form. Such a small amount of 01- ions is not expected to give the well-defined absorbance peak at 355 nm that we observed in our experiment with neutral iodine solution. Besides, iodine in several alcohols and acetone also gave an absorbance maximum at around 355 nm. 01- ion formation in these solutions seems unlikely, as for example, in order to obtain 01- from the acetone solution of iodine, it is necessary to break a carbon-oxygen double bond and then to form an 0-1 bond together with a doubly positive carbocation. Later in this discussion, an explanation is given for the existence of a 355-nm absorbance band for various oxygen-containing solvents used in this experiment. Theoretical and experimental evidence thus indicates that the absorbance band observed at 355 nm in the aqueous solution of iodine is not due to 01- ions. In order to identify what causes absorbance at 205, 285, and 355 nm in the aqueous solution of iodine, INDO CI calculations were carried out on possible complexes of iodine with water. Different configurations (143) for these calculations were obtained by permitting single excitations from the occupied Hartree-Fock ground state. The basis functions for oxygen and iodine were the same as those already mentioned, and hydrogen s functions were also included. The optimization of H2012geometry with 1-1 at right angles to the C2, axis of HzO (perpendicular geometry) provided a complex (Figure 1) with a small binding energy (31 kJ/mol). Here, the binding energy is defined to be the energy obtained by subtracting INDO S C F energies of H20 and 1, from that of H2012geometry. In this complex, the center of the 1-1 bond is at a distance of 2.19 8,from the oxygen atom and the 0-H and the 1-1 bond distances are 1.02 and 2.84 A, respectively. Both the 0-H and the 1-1 bonds are longer than those in H 2 0and 12. The bond angle H-0-H is 101.8’, which is smaller than that in water. Also, INDO CI calculations indicate that the absorbance bands due to this complex will be at 186, 500, and 610 nm. The 186-nm band is the strongest band followed by the 610- and 500-nm bands. On the other hand, optimization of H2012geometry with 1-1 along the C2, axis of H 2 0 (axial geometry) provided a stable complex (Figure 2) with a fairly large binding energy (228 kJ/mol). In this geometry, 0-1 (nearest to oxygen), 1-1, and 0 - H bond distances are 2.10, 2.87, and 1.02 A, respectively. Like the perpendicular geometry, 0-H and 1-1 bond distances are longer than those in H 2 0 and 12. Contrary to the perpendicular geometry, the H-0-H bond angle in the axial geometry is larger ( 108.5O) than that in the water molecule. The iodine atom closest to the oxygen has a small positive charge (+0.1) and the other iodine has a negative charge (-0.4), indicating polarization of the I2 unit in this complex. The I2 unit in the perpendicular geometry is, however, unpolarized. As expected, oxygen has a negative charge in each complex (-0.3 in perpendicular geometry and -0.4 in the axial) and each hydrogen has a positive charge ( f 0 . 2 in perpendicular and +0.3 in axial geometry). A larger positive charge on hydrogen may account for
3838 The Journal of Physical Chemistry, Vol. 93, No. 9, I989
I 2.87 A"
2.10A"
H Figure 2. Axial complex.
a larger H-0-H bond angle in the axial complex whereas a smaller positive charge may account for a smaller bond angle relative to that in the water molecule. INDO CI calculation on the axial complex gives absorbance maxima at 181 (most intense), 277, and 365 nm. The last two wavelengths match fairly well with the experimental values (285 and 355 nm) with an intensity ratio of 2:l. It should be pointed out that the above calculations were done on the isolated gas molecule. Hence, shifts in wavelengths (of maximum absorbance) in the experiment may be due to solvation of the highly polar complex in water. Since the far-UV range of the spectrum was not possible to obtain with our equipment, verification of the absorbance band at 181 nm was not possible. However, the most intense peak that we observed at around 205 nm may be the 181-nm peak of the gas phase, shifted due to solvation and hydrogen bonding. The reason for a larger shift of one band and relatively smaller shifts of the other two in the solution is not quite clear. Besides, theoretical calculations suggest a much larger band intensity at 181 nm than was obtained experimentally at 205 nm. One possible explanation may be that the 181-nm band arises from a transition to a state having energy close to that of the first ionization threshold. Accurate theoretical calculations, in such circumstances, would necessitate inclusion of Rydberg basis functions that are known to shift such spectral lines to a larger wavelength with a decrease in band intensity.*O It is interesting to note that the axial form of the complex is more stable than the perpendicular form by 197 kJ/mol. Since no absorbance bands at 500 and 610 nm were observed experimentally, one can conclude that the perpendicular complex is either present in a negligibly small amount or unstable in the solution. Several other geometries were tried with 1-1, making angles in between Oo and 90° with the C2"axis of HzO. These, however, did not provide any stable complex or spectra close to our experimental values. However, it should be noted that the INDO Hamiltonian can lead to predictions of spurious stable geometries; hence, considering the small binding energy and shallow potential
Fonslick et al. well for the perpendicular form, some doubt can be cast on the existence of such a complex. INDO calculations, with use of similar basis functions as ours, are, however, known to give spectra with excellent agreement with the experimental re~u1ts.l~ The analysis of the molecular orbitals (MO's) of the axial complex involved in electronic transitions reveal interesting features. The band at 18 1 nm was obtained by transition from a M O (5a,) containing predominantly iodine orbitals with a small percent of orbitals of oxygen and hydrogen. On the other hand, the absorbance band at 277 nm involved 2bl and 2b2 degenerate MO's and that at 365 nm involved 3bl and 3bz degenerate MOs, containing almost no hydrogen and hence orbitals of oxygen and iodine only with a major contribution from iodine. It is interesting that the 2bl and 2b2 orbitals are mostly concentrated on the iodine (90%) closest to the oxygen and 3bl and 3b2 are concentrated mostly on the other iodine (90%) farthest from the oxygen. The LUMO (lowest unoccupied MO) for each of these bands is the same antibonding orbital (6a1) with an appreciable contribution from each iodine atom (49% and 36%, respectively) and a small contribution from oxygen and hydrogen. Hence, polarization of iodine by oxygen and charge transfer during electronic transition from one iodine center to the other can explain the existence of the 355-nm band. Since most of these electronic transitions are internal to iodine in the H2012axial complex, a comparison of I2 spectra with those of the H2012complex would be useful. Close to our observed lines, the spectral lines for I2 exist at 215, 246, and 270 nm resulting from transitions from the ground state to different excited states of iodine.I4 Polarization of iodine during the complex formation together with the mixing of oxygen and hydrogen orbitals with those of iodine presumably shifts these absorbance bands to those corresponding to the complex. It is expected that the other oxygen-containing substances may give a similar complex by polarizing iodine and absorb at the 355-nm wavelength. This provides an explanation for Mokhnach and Rusakova's observation. An increase in absorbance at around 355 nm due to the addition of starch to an aqueous solution of iodine is presumably caused by the formation of iodine complexes with oxygen of the starch molecule. The absorbance maximum at around 355 nm for different alcohols and acetone, as observed in the present experimental study, can also be explained by an axial type of the iodine-oxygen complex formation. Further theoretical work is in progress on iodine complexes with different solvent molecules. In conclusion, it can be said that the absorbance band at around 355 nm in the aqueous solution of iodine is not due to 01- ions, and hence, the threshold, observed in the starch-iodine calibration line, is not due to the hydrolysis reaction of I2 in water. Further, the starch-iodine blue complex formation in the absence of KI cannot be explained by the hydrolysis of iodine where hypoiodite and iodide ion formation was assumed.z Acknowledgment. We thank M. Zerner for allowing us to use his INDO program for theoretical computations and L. Edel for his experimental assistance. We also thank D. Shaffer for generating plots for us. For partial support of this research, we acknowledge a grant from the DuBois Educational Foundation. Registry No. 01-, 15065-65-3; H,O.I,, 71034-54-3; I,, 7553-56-2; starch, 9005-25-8; ethyl alcohol, 64- 17-5; 1-propanol, 71-23-8; acetone, 67-64-1; benzene, 71-43-2; toluene, 108-88-3; o-xylene, 95-47-6; m-xylene, 108-38-3; p-xylene, 106-42-3. (13) Edwards, W. D.; Weiner, B.; Zerner, M. C. J . Am. Chem. SOC.1986, 108, 2196, and references therein.
(14) Huber, K. P.; Herzberg, G.Constants ofDiatomic Molecules; Van Nostrand Reinhold: New York, 1979.
