Complexation constants diiodine

Table 1.13 Logarithm (to base 10) of complexation constants (I mol ) of diiodine with bases at 25° C in the indicated solvents.

The fact that the determination of diiodine complexation constants is dependent on a second unknown makes values more uncertain than the hydrogen-bond formation constants. A revision of the statistical evaluation of diiodine complexation constants obtained by the popular Benesi-Hildebrand method [53] shows [57] that the confidence interval is always much larger than previously reported. Examples of revised 95% confidence limits are (in 1 mol ) 0.32-0.40, 0.53-1.86 and 1.10-1.32 for the complexation constants of diiodine with 1-bromobutane, benzene and dioxane, respectively. However, better 95% confidence intervals can be obtained. A careful application of the Rose-Drago method to the complexation of diiodine with carbonyl bases gives [62], for example 0.53 0.04 (benzaldehyde), 1.12 0.06 (acetone), 8.1 0.7 (AA -dimethylbenzamide) and 15 0.4 (A,A-dimethylacetamide) (in 1 mol , in heptane at 25 °C). 

The measurement of diiodine complexation constants has generally been performed in alkanes or chlorinated solvents. It has been observed [63] that the value of the equilibrium... 

When the diiodine complexation constants cannot be measured either in alkanes or in CCI4, solvents such as CH2CI2 or CHCI3 are generally used. To establish relationships between the Ka values determined in alkanes and those determined in CH2CI2 or CHCI3, there are only enough literature data for pyridines and thiocarbonyl compounds. The equations found are as follows ... 

Table 5.4 Diiodine complexation constants (I mol j measured in heptane and in various alkanes.

Figure 5.7 Family-dependent relationships between diiodine complexation constants measured in alkanes, pKg 2, and in CCI4, logKc (in CCI4) nitriles (n), phosphoryls (o), it bases (x), thioethers (<)), pyridines fBj and selenoethers ( ). The dashed line of slope unity is drawn for comparison.

Table 5.5 Family-dependent relationships (Equation 5.11) between diiodine complexation constants measured in alkanes and in CCI4.

Table 5.5 Equilibrium constants (I mol 0 and diiodine basicity scales pKgo 3/7 J AG° (kj mol 0 for the complexes of diiodine with six-membered N-heteroarenes.

Table 5.9 Equilibrium constants (I ) and diiodine basicity scales pKgi and AC° (kj moh ) for the complexes of diiodine with five-membered N-heterocycles (mainly heteroarenes).

Table 5.12 Equilibrium constants Kc (I mol j and diiodine basicity scales pKsi2 3-nd AC° (kj mol j for the complexes of diiodine with carbonyl bases.

The V (1—1) frequency shifts upon complexation of diiodine with about 100 Lewis bases. The complex is formed in cyclohexane at 20 °C. Although the diiodine stretch cannot be observed by IR spectroscopy, since it does not produce a dipole moment variation, the polarization of the I—I bond by complex formation renders the vibration IR active. The frequency shifts are calculated from the free diiodine value of 210 cm measured by Raman spectroscopy. The highest frequency shift is found for piperidine (39.5 cm ). In the interpretation of these results, it must be noted that the normal coordinate describing v(I—I) is expected to mix with v(B- 1) because the two bands are generally close. Hence, the frequency shift v(I—I) is not simply related to the change in force constant of the I—I bond upon complexation. 

Determination of the Complexation Constant of Diiodine with lodocyclohexane by Visible Spectrometry... 

Haloalkanes are weak Lewis bases towards all Lewis acids. For this reason, they have attracted little attention in the construction of Lewis basicity scales. In this experiment, their basicity towards diiodine is studied by measuring the complexation constant of diiodine with lodocyclohexane in cyclohexane and comparing the result with those for other haloalkanes. 

Visible spectrometry is the method of choice for measuring diiodine basicity because the visible band of diiodine and the blue-shifted band of its complexes have rather high absorption coefficients ( 10001 moC cm ). Consequently, there are significant variations in absorbance on complexation, even when small quantities of complex are formed, and accurate complexation constants can be obtained for weak Lewis bases. A disadvantage of electronic spectrometry is the width of the absorption bands. Hence, even significant blue shifts on complexation cannot prevent the overlap of the visible bands of free and complexed diiodine. Therefore, there are two unknowns, the complexation constant and the absorption coefficient of the complex, in the equation relating the equilibrium constant to the absorbance and the initial concentrations of diiodine and base. 

Table 7.10 Calculation of the complexation constant of diiodine with iodocyclohexane in cyclohexane at 25 °C. ...

The best conditions for obtaining accurate diiodine affinities by visible spectrometry from the temperature variation of the complexation constant are studied in this experiment using the example of the complexation of diiodine with dimethyl sulfoxide (DMSO) in CCI4. The result will be compared with those of two previous studies by Drago et al. [16] and Klaeboe [17], The importance of the solvation term in the diiodine affinity measured in CCI4 will be evaluated by comparison with a value measured in cyclohexane, a less solvating medium. 

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