2098
JOHNA. MAGUIRE AND JOHNJ. BANEWICZ
A Spectroscopic Study of the Propionitrile-Iodine Molecular Complex by John A. Maguire* and John J. Banewicz Department of Chemistry, Southern Methodist University, Dallas, Texas 76222 (Received J u l y 19, 1971) Publication costs assisted by T h e Robert A . Welch Foundation
The formation equilibrium constant, K , for the molecular complex formed between propionitrilc and iodine in hcptane and carbon tetrachloride was determined from spectral data in the visible region at 25". Analysis of the data using a Scott type of equation yielded values of K that were wavelength dependent. This variation in the determined value of K was attributed to a variation in the optical properties of both the complexed and uncomplexed iodine molecules with increasing propionitrile concentration. An equation that corrects for these variations was developed and the analysis gave values of K that were constant to within +5%. The average value of K at 25" for the propionitrile-iodine complex in heptane was found to be 0.96 N - 1 .
Introduction Wtl havc been investigating the stabilities and thermodynamics of formation of the molecular complexes formcd bctwccw iodine and a series of aliphatic nitrilrs, using spectrophotometric methods. In thc course of this irivestigatiori w have noted discrepancies between the values of the formation equilibrium constant, K , for the propionitrilc-12 complex in carbon tetrachloride obtained in this laboratory and thosc rcportcd by Klaboe.' An examination of the results indicated that these discrcpancics could not be blamed entirely on experimental indetermination but werc due in part to the fact that thc equilibrium constants wcre determined at differcnt wavelengths. A detailed study of the propionitrile-12 system was undertaken to examine the apparent wavelength dcpendence of K .
Experimental Section Materials. Reagent grade iodine x a s purified by sublimation from potassium iodide before use. Eastman propionitrile, frec of isonitrile, E K 528, was purified by stirring with calcium hydride, followd by treatment with PtOj, then fractionation. The distilled nitrile was stored over a molecular sieve (Lindc 3A) until used. Thc propionitrile was checked by glpc and found to contain less than 0.2 ppt water. KOother impurities wcrc detectable by glpc. Baltcr GCSpectrophotometric grade carbon tetraehloridc and B & A Instrument grade hcptanc werc used without furthrr purification except that dry nitrogen was bubbled through the solvents before use. Preparatzon of Solutions and Procedure. All solutions wcrc prepared by \wight as described Jlixing and handling of tho solutions w r c carried out in a drybox to insure against moisture contamination. The spectra w r c determined using either a Becltman IIodcl DK2A recording spectrophotometer or a Beckman SIodel DU spectrophotometer. The temperature of thr solutions in the spectrophotometers was conT h e Journal of Physical Chemistry, Vol. 7G, .Yo. 15, 1972
trolled to within *0.1" using thermostated ccll compart ments. The general procedure used in obtaining the spectra has already been described.2 In this procedure the difference betmcen the absorbance ( A ) of the iodincnitrile solution and tho absorbance [Ao) of a rcferencc solution containing the same total iodine concentration was read directly. Because of the strong absorption of the nitrile in the ultraviolet, the charge-transfer band could not be used. Continuous plots of A - A0 were made in the visible region using thc DK2A spectrophotometer. These plots showcd that the magnitudc of A - A . reached maxima at about 535 and 460 nm. For the DU runs, selectod wavelcngt'hs were used in the rangc between 440 and 560 nm. All readings were taken at 25".
Results and Discussion From the value of A - A0 a t a particular nitrile concentration (CD) and a total iodine concentration (CA), the quantity [CDCA/(A - A o ) ] ( Xwas ) calculated at each wavelength. If only a 1 : 1 complex exists and Beer's law is obeyed by all specks, then a plot of X us. CD CA should yield a straight line according to the
+
CD
+ CA =
(€C
- EA)X - K-'
(1)
in which K is the formation equilibrium constant of the complex, EC is its molar absorptivity, and E A is the molar absorptivity of free 12. This equation, which is a modification of that first proposed by S ~ o t tis, ~the one usually used in analyzing spectroscopic data to obtain values of formation equilibrium constants. I n dcriving this equation, the absorptivity of the free nitrile is (1) P. Klaboe, J . Amer. Chem. Soc., 8 5 , 8 7 1 (1963). (2) J. A . ,Magtiire, A. Bramley, and J. J. Banewicz, Inorg. Chem., 6, 1752 (1967). (3) K. J. Rose and It. S. Drago, J . Amer. Chem. Soc.. 81, 6138 (1959). (4) R. L. Scott, Reel. T r m . Chim. Pays'Bas, 7 5 , 7 7 8 (1956).
PROPIONITRILE-IODINE MOLECULAR COMPLEX 4.2,
I
I
I
I
I
I
2099 11 -4.2
I
c p" I
P
Y
\
e
s 10
$2
CD
+ e,, M.
Figure 1. A plot of CDCA/(A - Ao) US. CD 4- CA at 25' for 460 (e), 475 (O), and 540 (0)nm. The left-hand ordinate is for 460 and 475 nm; the right-hand ordinate is for 540 nm.
assumcd to be zero and thc conccntration of complex small compared to the nitrile concentration. I n cases whcre the complcx concentration becomes important an additional term of the form, -(A - Ao)/(Ec - E A ) , is added to the right-hand side of cq 1. A value of K can be calculated from the intercept and ec - E A from the slope. Typical plots arc shown in Figure 1. I n the concentration rangc of CDfrom 0.2 to 1.1M , such plots appear to be lincar. The bcst values of K and EC - EA were calculated from eq 1 by the method of least squares and are listed in Table I. cc - CA for the Propionitrile-Iodine System Obtained from Low-Concentration Data a t 25" Using Eq 1
Table I: Values of K and
Wavelength, ---Solvent nm ,ya ,b
440 445 450 455 460 465 470 475 520 330 535 540 550 560 a
0.39 0.42 0.45 0.48 0.51 0.52 0.56 0.60 0.47 0.51 0.51 0.51 0.53 0.53
rC
f 0.02
f 0.02 f 0.02 f 0.02 f 0.02 f 0.02 f 0.02 f 0.01 f 0.03 f 0.01 f 0.01 f 0.01 f 0.01 i 0.02
K in M-I.
---Solvent
CCIp----
-
664 689 696 688 657 615 534 450 -557 -606 -612 -599 -515 -415
Ka,b
rAb
C?Hla--IC
f 25 0 . 7 1 i: 0.02 f 21 0.77 f 0.02 f 22 0 . 8 1 f 0.01
f 19 i: 17 f 17 f 12 f6 f 24 f9 f8 f8 f9 f 13
0.84 0.86 0.99 1.02 1.09 0.86 0.97
- 6Ab
579 610 642 f 0.01 660 663 f 0.01 615 f 0.02 566 f 0.02 ,500 i 0.02 f 0.03 -480 f 0 . 0 2 -565
0 . 9 8 f 0.02 1.01 i: 0.02 1 . 0 4 f 0.03
f 10 f9 f6 f6 f6 f7 f7 f5 f9 f7
-588 f 8 -535 i 6 -442 f 8
* Uncertainties are standard deviations.
For the propionitrile-12 complex in carbon tetrachloride, Klaboe' reports values of K at 20" of 0.41, 0.44, and 0.44 at 430, 435, and 440 nm, respectively. Becausc of the diffcrencc in temperature, a direct comparison of Klaboe's results and those in Table I cannot bc made. Howcver, if one uses the value of
AHo of -2.7 kcal/mol found for the ~ y s t e mK, ~can be corrected to 20". This corrected value at 440 nm is 0.43, in excellent agrcemcnt with Klaboe's value at this wavelength. The small but systematic variation of K with wavelength secn in Table I indicates that evcn though the plots shown in Figure 1seem to be linear, the assumptions on which cq 1 is based are not complctcly valid. Deranleau6 has examined in dctail thc theory of measurement of equilibrium constants for weak complexes and has emphasized the importancc of thc concentration range in analyzing data for such complexes. According to Deranleau, data should be obtained ovcr a widc rangc of the saturation fraction, s (in this casc the fraction of Ipcomplexed). I n the optimum rangc of s, between 0.2 and 0.8, relative errors in K and AE are minimized. A large portion of the saturation fraction region should bc covcred bcfore any modcl can be considercd proven and for low saturation fractions apparent linearities of plots similar to Figure 1 arc of little use in establishing a modcl. The results shown in Tablc I arc in thc s rangr from 0.2 to 0.3 in the case of heptanc and from 0.1 to 0.33 for carbon tctrachloride, assuming that the calculatcd values of K arc approximatcly correct. Table I1 lists the determined values of K and EC - E A for thc propionitrileiodine system in hcptane when thc s range is extended to about 0.7. For these high-concentration studies heptane was uscd as a solvent. I n this solvent the values of K arc higher than in carbon tctrachloride. I n addition, complications due to solvent competition, such as those dcscribed by Carter,' should bc minimized. The effect of thc range of s on the dctermincd valucs of K and EC - EA obtained using eq 1 can be seen by a comparison of Tables I and 11. The determined values of K are both conccntration and wavelength dcpendent. This is especially evident in the 400-nm region. Figure 2 shows some typical plots of X us. C D C A for these higher Concentration studies. Using the values of K and EC - EA listed in Tablc 11, values of A - A. m r e calculated for diffcrcnt donor conccntrations and the differences betwcen the calculated and observed values of A - A,, determincd. Figure 3 s h o w these differences in A - A. as a function of donor concentration at wavelcngths of 445 and 475 nm. Similar behavior was noted at the other wavclengths and indicates that, although Figure 2 appears to be linear, a definite curvature does exist. I n addition, the isosbestic point n a s found to blue shift by about 2 nm per unit molar increasc in CD. This shift in the isosbestic point is an indication that, in addition
+
(5) J. A . Maguire and J. J. Banewice, to be published. (6) D. A. Deranleau, J . A m e r . Cham. Soc., 91, 4044, 4050 (1969). (7) S. Carter, J. N . Murrell, and IC. J. Itosch, J . Chem. Soc., 2048 (1965).
T h e Journal of Physical Chemistry, Vol. 7 6 , N o . 16,1872
JOHN A. MAGUIRE AND JOHN J. BANEWICZ
2100
Table I1 : Analysis of Propionitrile-12 in Heptane a t 25". High-Concentration Data
a
A , nm
Ka,b
445 460 465 475 520 540 560
0.69 5 0.01 0.88 f 0.01 0.95 f 0.01 1 . 1 3 f 0.01 0.76 0.01 0.95 f 0.01 1.02 5 0.01
K in
M-1.
I'
0.5
aC
---Eq 2----(using KAV = 0.962
- eAb
658 f 7 653 5 4 614 5 2 483 f 3 -516 zt 4 -588 z!z 2 -442 f 1
Uncertainties are standard deviations.
I
I
I
I
B
0.90 0.99 0.92 0.96 0.92 0.99 1.05
527 593 631 545 - 441 - 573 -431 KA" = 0.96 zt 0.04c
B
a
34 17 -5 - 18 -20 -4 -3
502 607 611 543 - 427 - 583 -460
a
40 13
1
- 18 -24 -1 4
Uncertainty is average deviation.
i,
B
I 2.0
1.0 1.5 CD d-cn, M.
-
Ka
Figure 3. Differences in experimental and calculated values of A - AOobtained from eq 1: 445 ( 0 ) and 475 (0)nm.
+
Figure 2. A plot of CDCA/(A Ao) us. CD CA at 25' for 445 (a), 465 (a), and 540 (0)nm. The left-hand ordinate is for 445 and 465 nm; the right-hand ordinate is for 540 nm.
to the 1 ;1 complex, some new absorbing species is being formed by the interaction of the nitrile and iodine or that the free and complexed iodine are not obeying Beer's law. Although the influence of higher complexes on Scott plots has been extensively examined,s very little work has been done on the consequences of deviations from Beer's law.g The sensitivity of the visible I2 band to its environment has been noted by Voigt,'O who measured the spectra of 1 2 in a number of different solvents. I n most solvents the visible I2 band was found to blue shift from its gas-phase position. Voigt has proposed that in addition to charge-transfer and contact chargetransfer interactions, exchange repulsion of the excited state of Iz with nonbonding electron pairs may be a further cause of the blue shifts. This was proposed to account for the small blue shifts noted for Izin the alkyl chlorides and perfluorocarbons in which no contactcharge-transfer bands have been observed. The exact mechanism of this effect is not completely understood, but it seems to involve a weak, nonbonding, yet specific interaction between electron pairs on the solvent and the iodine molecule. The Journal of Physical Chemistry, Vol. 76, N o . 16,1078
I n order to cover a meaningful portion of the saturation fraction range, the mole fraction of nitrile must be increased to about 0.26. These high propionitrile concentrations will materially change the nature of the solvent. It would be unusual if the spectral properties of the free and complexed iodine were not altered. The effect that these changes will have on plots of X vs. CD CA will be determined by the functional dependence of ec and E A on nitrile concentration. A reasonable assumption is that the molar absorptivities of the free and complexed iodine vary in a linear fashion with CD. This would be expected if the interactions between excess nitrile and iodine were specific ones, similar to those proposed by Voigt. Thus if EC = €2 a C D and EA = EAO bCD, the equation relating x to CA becomes CD
+
+
+
+
CD
+
CA =
BX
+ aXCD - K-I
(2)
+
where B = ECO - EA" bK-'. I n obtaining this equation the concentration of complex is assumed to be small compared to the donor concentration. (8) G. D. Johnson and R. E. Bowen, J . Amer. Chem. Soc., 87, 1655 (1965). (9) R. Foster, "Organic Charge-Transfer Complexes," Academic Press, New York, N. Y . ,1969. (10) E. M. Voigt, J . Phys. Chem., 7 2 , 3300 (1968).
PROPIONITRILE-IODINE MOLECULAR COMPLEX The high-concentration data were analyzed using eq 2, and the results are listed in Table 11. If eq 2 is valid, then identical values of K should be obtained at the different wavelengths. Although experimental scatter does exist, the systematic variation in K with wavelength seems largely to have been eliminated and the values of K are the same to within about +5%. I n view of the small deviations from linearity for which eq 2 attempts to correct, the agreement among the values of K is considered to be satisfactory. An analysis of the data was also carried out using the method proposed by Scatchard." It has been shown6 that under certain conditions this method is more sensitive to complications thgn is the Scott treatment; in addition, the two methods weight experimental points differently. Equations 1 and 2 can be converted to the Scatchard form by multiplying both sides by K / X . The quantitative results obtained from the Scatchard treatment were essentially the same as those listed in Tables I and 11. This indicates that in this case the difference in the weighting of the experimental points does not significantly affect the results. The values of A - A . calculated by eq 2, using the parameters listed in Table 11, were in excellent agreement with the experimental values. For 24 of the 39 experimental points there was no difference between the experimental and calculated values of A - Ao,in most of the other cases the difference was 0.001 absorbance unit, and in no circumstance was the difference greater than 0.003. Furthermore, the discrepancies that were found showed no trend with donor concentration. Although the agreement between the calculated and experimental values is excellent, it should be pointed out that the additional term in eq 2 automatically ensures a better fit than could be obtained from eq 1. The sensitivity of A - A. to the value of K can be assessed using the sharpness of fit criterion of Conrow, Johnson, and Bowen.12 The sharpnesses were calculated at the individual wavelengths and found to range from a minimum of 10 at 450 nm up to a value of 410 at 475 nm. The average value of K was calculated to be 0.962. The best values of K and a calculated at each wavelength using this average value of K are also listed in Table 11. The magnitude and sign of a determine the curvature in plots of X vs. CD CA,which in turn gives rise to the variations in K found in Tables I and 11. The opposing shapes of the curves shown in Figure 3 illustrate the different curvatures produced by positive and negative values of a. It is of interest to note that the variation in a with wavelength is consistent with a blue shift of the visible complex peak with increasing donor concentration. Only in spectral regions where a is quite small can eq 1be used.
+
2101 Johnson and Bowen,8 in summarizing the criteria for the establishment of a 1 : l molecular complex, state that if a variation of K with wavelength is found it is a strong indication that termolecular complexes are also present. The general equations relating the absorbance of the solution to concentrations of donor and acceptor for the case of 1:l and 1:2'complexes have been reported by these investigators. Under the conditions where the concentration of donor is large compared to both of the complexes, the equation relating CD CAto X differs in functional form from eq 2 only in the addi~ The inclusion of this term permits tion of a C D term. an even more precise fit of the experimental data. However, this improvement was found to be beyond that permitted by the inherent experimental uncertainty. That is, the value of A - A0 cannot be read to any better than *0.001 OD unit, and any fit better than this is superfluous. We feel therefore that eq 2 can adequately account for our results, and the assumption of the existence of higher order complexes is not necessary. Interpretation of optical data can be complicated not only by the formation of termolecular complexes but also by deviations from Beer's law when only 1: 1 complexes are formed. I n systems involving the formation of weak complexeswhere high donor concentrations are needed to cover a sufficient range of s values, a great deal of caution must be exercised in interpreting optical data. I n particular, if equations similar to eq 1 are to be used, data must be obtained over a wide range of wavelengths and evidence of any curvature noted. Even when apparently linear plots are found, slight curvatures can produce rather drastic variations in the determined values of K with wavelength. When such evidence is found, more complex equations are required which correct for complications caused by higher order complexes or changes in the optical properties of the absorbing species. When the curvature is slight, it may not be possible t o distinguish unambiguously between these two cases. For the propionitrile-I2 system, the assumption that the optical properties of the absorbing species are altered as nitrile is added accounts for the experimental results.
+
Acknowledgment. The authors wish to thank the Robert A. Welch Foundation of Houston, Texas, for support of this project through Grants No. N-142 and N-056.
(11) G. Schatchard, Ann. N . Y . Acad. Sci., 51, 660 (1949). (12) K. Conrow, G. Johnson, and R. Bowen, J. Amer. Chem. Soc., 81, 6138 (1959).
The Journal of Physical Chemistry, Vol. 76, No. 16, 1973
