1861
STABILITY OF OXYGENIODIDE CHARGE-TRANSFFIR COMPLEX their respective zwitterionic forms. Consequently, we must consider the relative free energies of the possible intermediates involving proton bound to oxygen as opposed to those in which it is bound to nitrogen. Finally, a noncyclic but synchronous hydration mechanism provides an additional alternative. It should be noted that all the deuterium isotope effects reported in this paper are less for 2-pyridinecarboxaldehyde than for 4-pyridinecarboxaldehyde. Although one might be prompted to associate the difference in isotope effects with the proximity of nitrogen and the aldehydic group in 2-pyridinecarboxaldehyde, it should be recalled that the spontaneous rate coef-
ficient for the hydration of 4-pyridinecarboxaldehyde is actually somewhat higher than the corresponding value for the 2-pyridinecarboxaldehyde hydration (Table 11). This latter observation could imply lack of importance of cyclic transition states involving the ring nitrogen in the hydration of 2-pyridinecarboxaldehyde, transition states which are otherwise conceivable due to the unique environment a t the site of hydration of this latter s u b ~ t r a t e . ~It is hoped that studies which are presently being carried out in these laboratories pertaining to the activation parameters associated with these two reactions will further rationalize these differences.
The Spectrum and Stability of Oxygen Iodide Charge-Transfer Complex by H. Levanon and G . Navon Department of Physical Chemistry, The Hebrew University, Jerusalem, Israel
(Received August 1 5 , 1 9 6 8 )
Extra absorption spectra due to oxygen iodide charge-transfer interaction in different solvents were studied. Attempts to determine tl 3 stability constant of the 021- complex led to an upper limit of K I. 0.3 M-l. Heats of formation of 1.2,0.0,and -0.7 kcal in water, ethanol, and ethyleneglycol,respectively, were obtained. From the solvent dependence of the optical absorption intensity and the heats of formation, an outer-sphere type oxygen iodide complex was proposed. From the energies of the optical transitions, a vertical electron affinity of Eozv < 5 kcal for oxygen was estimated.
Introduction When oxygen is dissolved in solutions containing halide and thiocyanate ions, a new optical absorption in the uv region takes place.' This absorption has been attributed to a charge-transfer interaction between the oxygen acting as an electron acceptor and the halide ions as electron donors. Charge-transfer interaction of oxygen with various organic solvents has been observed before,2--B and the vertical electron affinity of oxygen has been estimated from its spectra.? The aim of this work is to investigate in some detail the charge-transfer interaction between oxygen and iodide ion, t o measure the stability of this complex, and to estimate the vertical electron affinity of oxygen from its spectrum.
Experimental Section Materials. Matheson Extra Dry grade oxygen and Prepurified grade nitrogen were used. Water was triply distilled. Acetonitrile was from Matheson Coleman and Bell, spectroscopically pure. All other chemicals were AR grade without further purification.
Procedure. In all cases, blank solutions were prepared by using nitrogen in place of oxygen. All measurements at atmospheric pressure were carried out by bubbling the gas through two vessels cont,aining the solutions in series. The first solution was used to saturate the gas with the vapors of the solution and was discarded. The second solution was poured after the bubbling into a 10-cm silica spectrophotometric cell. The whole procedure was carried out in a thermostated water bath. For the high-pressure measurements, a high-pressure cell with an optical path length of 5 cm was used.* The gas was dissolved in the solutions by shaking the cell under the appropriate pressure. The spectrophotometric measurements were carried (1) G . Navon, J . Phys. Chem., 6 8 , 969 (1964). (2) D. F. Evans, J . Chem. Soc., 345 (1953);J . Chem. Phys., 2 3 , 1424 (1955). (3) A. U. Munck and J. F. Scott, Nature, 177, 587 (1956). (4) L. J. Heidt and L. E. Ekstrom, J. Amer. Chem. Soc., 79, 1260 (1957). (5) H. Tsumbomura and R. 8. Mulliken, i b i d . , 85, 5966 (1960). (6) J. Jortner and U. Sokolov, J . Phys. Chem., 6 5 , 1633 (1961). (7) J. Jortner and U. Sokolov, Nature. 190, 1003 (1961). (8) Purchased from Research and Industrial Instruments Co., England.
Volume 75, Number 6 June 1.960
1862
H. LEVANON AND G. NAVON
0.5)-
Figure 1. Extra absorption spectra of 0 2 1 - in aqueous solutions. Oxygen a t atmospheric pressure; KI concentrations: (1) 0.2 M , (2) 0.4 M , (3) 0.5 M , (4) 0.6 M , ( 5 ) 0.8 M , (6) 1.0 M .
Figuiue 3. Benesi-Hildebrand plots for the results given in Figure 1.
out with a Beckman DU spectrophotometer, and some of them were repeated with a Cary Model 14 spectrophotometer. Every spectrophotometric result is the absorption difference of the solution of iodide and oxygen and the solution of iodide and nitrogen and the absorption difference between the solvent with oxygen and the solvent with nitrogen. To avoid any error due to stray light, the total optical density was kept under 0.9. Any solution in which oxidation was found, as
O 0.12 0.10
t
-
0.06 -
where 1 is the optical path length, D is the optical density, e is the molar extinction coefficient, and K is the stability constant given by
0.08
0.04
\a.02
inferred from the absorption of 11- at 352 mp, was discarded. It is convenient to evaluate the stability constant and the molar extinction coefficient of a complex by a linear plot when one component is in excess of the other^.^ In our case, two sets of experiments were carried out. The first was with atmospheric pressure of oxygen and varied excess concentrations of iodide, and the second was with a small concentration of iodide and large concentrations of oxygen obtained by dissolving the gas under high pressure. Benesi-HildebrandO plots for the above two sets of experiments, assuming 1:1 complex, are respectively
[IO,-] K = ([I-] - [IO,-]) ([O,]
-
- w2-1)
(3)
Except for the temperature dependence measurements, all solutions were kept a t a constant temperature of 24.0 f 0.5".
I
I
300
1
1
I
I
380
340 XJ rnP
Figure 2. Extra absorption spectra of 0 2 1 - in aqueous solutions. KI concentration 2 X 10-8 M ; oxygen pressure: (1) 20 atm, (2) 60 atm, (3) 80 atm, (4) 100 atm. The JournaZ of Physical Chemistry
I
Results Spectra i n Aqueous Xolutions. The extra absorption spectra due to oxygen iodide interaction of solutions (9) H. A. Benesi and J. H. Hildebrand, J . Amer. Chem. SOC.,71, 2703 (1949).
1863
STABILITY OF OXYGEKIODIDE CHARGE-TRANSFER COMPLEX I
I
I
-
I
yq
0.7
-
1
I
I
I
I
I
I
I
I
I
a9
0.5
0.3
‘\\
0.1
Figure 4. Benesi-Hildebrand plots for the results given in Figure 2.
saturated with oxygen under atmospheric pressure and varied concentration of K P J l are shown in Figure 1. The spectra of the complex with various oxygen pressures are given in Figure 2 . To make sure that the maxima in Figure 2 are not due to instrumental error, we repeated the measurements of trace 4 (Figure 2 ) , using the Cary 14 spectrophotometer and obtained the same results within experimental error. BenesiHildebrand plots of these two sets of experiments are given in Figures 3 and 4. According to eq 1 and 2, these plots should give straight lines with slopes of 1/Ke and an intercept 1/e. The same results are plotted in Figure 5 using Scott’s12modification
Z[I-][021
.D
- 1 I KE
[I-I
(4)
E
We did not find any significant change in the spectrum of iodide nor in triiodide under pressures up to 100 atm of nitrogen, in agreement with Fyfe’s results for iodide.13 Thus in our considerations, we ignore any change of the 10, spectrum under such pressures.
I
nc I
1
0.2
0.4
1.0
(I-) Figure 5. Scott plots for the results given in Figure 1. Wavelengths: 0 , 3 2 0 mp; X, 310 mp; A, 305 mp; 0,300 mp; 0, 290 mp.
Figure 6. Extra absorption spectra of 021- in ethanolic solutions. Oxygen a t atmospheric pressure; NaI concentrations: (1) 0.050 M , (2) 0.075 M , (3) 0.10 M , (4) 0.15M , (5) 0.20 M .
The stability constant K is the intercept-slope ratio in the Benesi-Hildebrand plot, or the slopeintercept ratio in the Scott plot. From Figures 3 to 5, it is evident that the stability constant of the oxygen iodide complex is too small to be estimated from our results; however, an upper limit of K 2 0.3 M-’ can be given.14J5 Spectra in Organic Solvents. Since solutions of oxygen in ethanol have an intense absorption due to charge-transfer interaction with the solvent, we have made spectral measurements only under atmospheric pressure of oxygen taking varied concentrations of sodium iodide. These are shown in Figure 6. The corresponding Benesi-Hildebrand and Scott plots are given in Figures 7 and 8. In chloroform and acetonitrile, the absorption of oxygen is relatively small and high-pressure experiments are feasible. Because of the low solubilities of iodide salts in chloroform, tetraethylammonium iodide salt was used. The extra absorption spectra of oxygen iodide complex in chloroform are shown in Figure 9. Owing to the high solubility of oxygen in chloroform, the sensitivity of the Benesi-Hildebrand plot for the determination of the stability constant in high pressures (10) The oxygen concentrations in various salt solutions were calculated from the relation log Ca/C, = k . [ S ] , where COand C . are oxygen concentrations in pure solvent and salt solutions, respectively ; [SI is the molar salt concentration, and k , is the salting-out constant which is 0.133 M-1 for aqueous KI solutions and 0.10 M-1 for ethanolic NaI solutions a t 24O.11 (11) A. Ben Naim, H. Levanon, and G. Navon, unpublished. (12) R. L. Scott, Rec. Trav. Chim., 75, 787 (1956). (13) W. S. Fyfe, J. Chem. P h y s . , 37, 1894 (1962). (14) Note should be taken that the Benesi-Hildebrand and the Scott
relations were derived without taking into account participation of solvent molecules in the stoichiometric equation. A treatment which includes solvent molecules was given by Carter, et al.16 However, the uncertainty of K in our result is large enough to be insensitive to such a correction. (15) 8. Carter, J. N. Murrell, and E,J. Rosch, J. Chem. Soc., 2048 (1965). Volume 79, Number 6 June 1060
1864
H. LEVANON AND G. NAVOM 0
0 03 0.01
'41-1 Figure 7. Benesi-Hildebrand plots for the results given in Figure 6.
Figure 9. Extra absorption spectra of 0 2 1 - in chloroform. Oxygen pressure of 60 atm; (Et)dNI concentrations: (1) 2 X M , (2) 4 x 1 0 - 4 ~ .
of oxygen is higher than in water solution. From Figure 10, one can see that the best straight lines that could be fitted to the experimental points give an apparent negative intercept. Such an intercept might not be necessarily due to experimental error. Carter, et ai.,15 pointed out that a negative intercept could be obtained in principle for weak complexes when solvent molecules are displaced during the complex formation. However, the difficulties in obtaining much higher pressures in this system do not allow us to draw conclusions from this behavior. In acetonitrile the results were irregular, probably because of oxidation processes; nevertheless, we could observe a reproducible absorption peak at about 286 mp. Determination of A H . I n all the experiments for the determination of the stability constant, the optical density was proportional to the oxygen and to the iodide concentrations. The expression of the optical density for very small stability constants has the form
D
= la[IOi-]
= KEL[OJ[I-]
(5)
Figure 10. Benesi-Hildebrand plots for 0 2 1 - in chloroform. (Et)kNI concentration, 2 X 103 M .
In this expression for the optical density only K and
LI
6
k I
s
Q 3
B can vary with temperature. However, usually E varies very little with temperature. In such cases, a plot of log D against the inverse temperature gives the heat of formation of the oxygen iodide complex, AH. From Figures 11 to 13 we estimate AH values of 1.2 f 0.2, 0.0 f 0.1, and -0.7 f 0.1 kcal in the solvents water, ethanol, and ethylene glycol, respectively. For the comparison given in the discussion we determined AH of the formation of the triiodide ion. Using the Benesi-Hildebrand equation we found values for K at different temperatures. Figure 14 describes log K vs. 1/T and a value of AH = -6.8 f 0.1 kcal was found. I n the above calculations the complexing of iodine with the solvent16was neglected.
(16) P. A. D. De Maine, J. Chem. Phys., 26, 1192 (1957). The Journal of Physical Chemietry
1865
STABILITY OF OXYGENIODIDE CHARGE-TRANSFER COMPLEX
: :I
-0.3
i1
-1.0
-'i
-1.41
0
0
I 3.3
I
32
I
I
3.5
3.6 10")
Figure 11. Logarithm of the extra absorption spectra of 021- in aqueous solution vs. the reciprocal temperature. Oxygen a t atmospheric pressure; K I concentration 0.5 M ;wavelengths: 300 mp; 295 mp; 0 , 3 2 0 mp; 0 , 3 1 0 mp; A, 305 mp; 0 , 2 9 0 mp; X, 285 mp.
+,
log 0
-0.1
I
.*0;I
. I
I
I
3.4
I/t
I
I
v,
-0.f
-
- 0.:
-
Discussion The Stability of the 021- Complex. Although we were not able to measure the stability constant of the oxygen iodide complex, we could, using eq 5 , evaluate the product K for every solvent. If E is independent of the solvent, which is the usual case, this implies a strikingly small variation in Ke (cf. Table I). For comparison, the stability constants of the triiodide complex in water," ethanol,'* and ethylene chloride,l9 are 7.7 X
I
2
3.3
I 3.4
3.5
Figure 13. Logarithm of the extra absorption spectra of 0 2 1 - in ethylene glycol, Oxygen a t atmospheric pressure; wavelengths: 0,330 mp; 0,320 mp; A, 310 mp; 0 , 3 0 0 mp; A,290 mpq
lo2,3.3 X lo4, and l o 7 M-l, respectively, i.e., a change of lo4 in going from water to ethylene chloride compared to a change of 5 for 021-in going from water to chloroform. These facts indicate that the oxygen does not replace any solvent molecule from the first coordination sphere of the iodide. Thus, we conclude from this discussion that the oxygen and iodide interact as an outer-sphere complex. This should be compared to the contact charge-transfer complexes suggested by Orgel and MullikenZ0 and discussed by others,l5,21 Such an outer-sphere model is in agreement also with the small heat of formation measured for the 0 2 1 complex (cf. Table 11). These heats of formation show a definite solvent dependence, but owing to the limited
I
3.2
I
I
3.3
3.4
I 3.5
I l/r.
] 10-3
Figure 12. Logarithm of the extra absorption spectra of 021- in ethanol. Oxygen a t atmospheric pressure; NaI concentration 0.15 M; wavelengths: A,320 mp; X, 310 mp; 0, 300 mp; A, 290 mp; 0 , 2 8 0 mp.
(17) M. Davies and E. Gwynne. J . Amer. Chem. Soc., 74, 2748 (1952). (18) P. Job, Compt. Rend., 182, 1621 (1926). (19) R. E. Buckles, J. P. Yuk, and A. I . Popov, J . Amer. Chem. Soc., 74, 4319 (1952). (20) L. E. Orgel and R. S. Mulliken, U M . , 79, 4839 (1957). (21) J. E. Prue, J . Chem. SOC.,7534 (1965).
Volume 78, Number 6 June 1969
1866
H. LEVANON AND G. NAVON
Table I
Solvent
Salt
Water
KI
Water Ethanol Chloroform Acetonitrile K I Ethylene glycol
KI
Salt concentrations, mM
Oxygen pressure, atm
Oxygen solubility,‘ mM/atm
200-1000 2 50-200 0.2-0.4 0.2-0.8 0.5
0.91 20-100 0.91 60 90 0.91
1.2-0.93 1.27 9.7-9.4 9.1 1oc 0.48
NaI (Et)aNI KI KI
Ke, X,.
mp
M-2
...
cm-1
