GIL NAVON AND GABRIEL STEIN
3630
the intensity relation are indicative of the large asymmetry parameter. The observed signals of VI and vI1 were narrow, the maximum-slope width being about 0.3 kc. According to a theoretical treatment by Leppelmeier and Hahn, 31 the magnetic dipole-dipole interaction is quenched for 14Xnuclei having an integral nuclear spin when they are placed in a field gradient having a large asymmetry. When signals were observed by Zeeman modulation, the intensity of V I was several times as high as that of P . This intensity
pattern is characteristic of large values of the asymmetry parameter and is due to the effect of the Zeeman field.32 This is explained by a recent theoretical calculation carried out by Negita33on urea, for which q = 32.3%. ~~
~
~
(31) G. W. Leppelmeier and E. L. Hahn, Phys. Rev., 141, 724 (1966). (32) M. Minematsu, J . Phys. SOC.Japan, 14, 1030 (1959). (33) H. Negita, J . Chem. Phys., 44, 1734 (1966).
The Reactivity of Some High- and Low-Spin Iron(II1) Complexes with Atomic Hydrogen in Aqueous Solution
by Gil Navonl* and Gabriel SteinIb Department of Physical Chemistry, Hebrew University, Jerusalem, Israel
(Received June 23* 1966)
The rate constants of the reactions of H atoms (produced by the electrodeless high-frequency dissociation of hydrogen gas) with some high- and low-spin complexes of FeI" were determined. The results are compared with those obtained for related H atom reactions with Fe"' complexes, using radiation chemical techniques, with reactions of such complexes with cas-, and with the rates of reactions of other metal ion complexes. The results are consistent with mechanisms involving H atom addition and, in the case of high-spin complexes, specific effect of the ligands on the resulta.nt electron transfer.
The mechanism of reduction processes involving electron transfer in aqueous solution is a subject of much experimental and theoretical interest. Atomic hydrogen in aqueous solution is a reducing agent having a simple structure. Since H atoms are neutral, extensive ionic interaction with the dipolar solvent medium is not involved. Following our previous investigation2 of the reduction of cobalt(II1) complexes by H atoms, we investigated the reactions of H atoms with various complexes of iron(II1). The technique is identical with that previously developed, and in a preceding publication3 we critically examined the assumptions involved in deriving the rate constant by means of this technique. There3 The Journal of Physical Chemistry
we also compared the results obtained using different techniques, e.g. , using ionizing radiations for obtaining H atoms. In the present work we investigated the reactions of H atoms with two groups of iron(II1) complexes. The first group includes low-spin com~ + ,Feplexes of iron(II1) : Fe(CN)63-, F e ( d i p ~ ) ~ and (phen)3*+which we compare with the results previously reported4 for the low-spin complex iron(II1)-cyto-
(1) (a) Bell Telephone Laboratories, Murray Hill, N. J.; (b) at present NSF Visiting Professor at Chemistry Department, Boston University, Boston, Mass. (2) G. Navon and G. Stein, J . Phys. Chem., 69, 1390 (1965). (3) G. Navon and G. Stein, ibid., 69, 1384 (1965).
REACTIVITY OF SOME HIGH-AND LOW-SPINIRON(III)COMPLEXES
chrome-C. The second group includes the high-spin complexes of iron(II1) : Fe(H20)2+, and the pentaaquo complexes FeF2+, FeOH2+, and FeC12+. These results we correlate with the rate constant estimated from the results previously obtained5 for FeS04+. Some of the experimental techniques and the formal treatment of the results are quite different for the two groups. These differences result from the fact that for low-spin complexes H atoms do not reoxidize the ferrous complex formed. I n the case of the high-spin complexes reoxidation occurs and must be taken into account.
Experimental Section H atoms were produced by a 30-Rlcps high-frequency electrodeless discharge in pure Hz gas a t 27 mm as described p r e v i ~ u s l y . ~The ~ , ~ dose rate of H atoms reacting with the solutions was determined using the ferricyanide do~irneter.~"Dose rates were of the M min-I in the present experiments. order of Exact dose rates were usually determined several times a day during experimental runs. Dose rates and results were usually reproducible within f10% in the course of the present investigation. For each experiment, 25 ml of solution was used with the reaction vessels and experimental arrangements described p r e v i o ~ s l y ~abt ~4". ~ Rate constants reported are calculated for 25", according to the details given in ref 2.
Results A . Low-Spin Complexes. I n these cases, the ferro complexes found are not reoxidized by H atoms. Therefore the reactions could be followed by determining the yield of ferrous ion formed or ferric ion that disappeared. Rate constants were then obtained by following the formation of the ferro form, undisturbed by reoxidation, and then calculating the rate constants using the procedure3 previously given. In this procedure one considers competition between the recombination of H atoms (for which a rate constant of 3 X 1O'O i1I-l sec-' is accepteds) and the reaction between H atoms and the scavenger employed. The kinetics are then inhomogeneous, diffusion-controlled competition between a second-order and a first-order reaction. The reliability of the absolute rate constants thus derived was critically examined in the previous work,* and it was shown that within a factor of 2 these values agreed with those obtained for the same H atom reactions using the techniques of, e.g., radiation chemistry which produces essentially homogeneous scavenging conditions. The rate constants obtained by the present method are themselves reproducible to within 10%.
3631
Fe(CN)63-. Some years ago the reduction of the hexacyanoferrate(II1) ion (ferricyanide) by H atoms was investigated and the order of magnitude of the rate constant e~timated.~"In the absence of the techniques now available3 for derivation of rate constants from the present method, earlier work paid attention mainly to the fact that, a t high enough concentrations of ferricyanide, it is capable of scavenging practically all H atoms reaching the solution and may therefore be used as a dosimeter. In the present work we reinvestigated the reactions of H atoms with ferricyanide in detail. Since ferricyanide may indeed serve as a dosimeter in investigations of the reactions of atomic H, we tried to establish the best conditions for the use of this dosimeter. Solutions of potassium ferricyanide were used. In Table I the results are shown. T o establish the real dose rate of introduction of H atoms into the solution, the concentration of ferricyanide was varied between 2.0 X and 1.88 X M , keeping the dose rate constant. Under these conditions the yield of ferrocyanide changed from 0.102
Table I: Reduction of Fe(CN)68- Ions" Initial concn, 10-4 M
Reduction yield, 10-4 M min -1
19.5 15.3 10.36 6.27 3.29 2.24 1.26 0.925 0.505 0.275
0.964 0.940 0.888 0.834 0.680 0.500 0.400 0.262 0.153 0.102
----lO-Qfi, b
M-1 sec-l---, C
d
4.6 1.9 2.1 1.4 1.5 1.9 1.8 1.3 1.5
1.7 1.5 1.4 1.7 1.3 1.4 1.8 1.7 1.2 1.5
0.72 0.79 0.92 1.28 1,07 1.6 1.7 1.5 1.1 1.4
10-alc av 2.0 h v dev, 70 30
1.5 10
1.2 24
'Time of reaction 1.5 min. b - d Rate constants calculated using the diffusion model, assuming the following dose rates: A = 0.95, 1.0, 1.1 X M min-1, respectively. N o t included in the average.
(4) G. Czapski, N. Frohwirth, and G. Stein, Nature, 207, 1191 (1965).
( 5 ) (a) W. G . Rothschild and A. 0. Allen, Radiation Res., 8 , 101 (1958); (b) G. Csapski and G. Stein, J . Phys. Chem., 63, 850 (1959). (6) G. Csapski, J. Jortner, and G. Stein, ibid., 65, 956 (1961).
(7) (a) G. Czapski and G. Stein, ibid., 64,219 (1960); (b) J. Rabani and G. Stein, J . C h a . Phys., 37, 1865 (1962); (c) A. Appleby, G. Scholes, and M. Simic, J . Phys. Chem., 67, 1610 (1963); (d) E. Hayon and M. hloreau, J. Chern. Phys., 391 (1965). (8) H. A. Schwara, J. Phys. Chem., 67, 2827 (1963).
Volume 70. Number 11 November 1966
3632
to 0.964 X M min-'. Columns b, e, and d in Table I show the rate constants in units of 109 4d-l sec-I-, assuming that the true dose rate is 0.95, 1.0, 11 min-'. The rate constant should or 1.1 X remain constant over the entire concentration range of the scavenger ferricyanide when the correct dose rate is used. This is so over the entire 100-fold range of concentration with the choice of c, in Table I, L e . , A = 1.0 X 114 min-I-, with the mean deviation being only hlOyc. Thus we do not observe any possible effect of scavenger depletion. The reaction was followed by measuring the change in optical density at 420 mp, takingi& E 1000 M-' cm-'. The value of k = 1.5 X IO9 A1-l sec-' thus obtained is in good agreement with rate constants obtained in radiation chemistry which range from 1.0 to 4.0 x 109 AI-1 sec-'.ib-'d From these results the best choice for the ferricyanide dosimeter is to employ a M solution. This will have an OD of 1.0 in a 1-em spectrophotometer cell. The yield of the dosimeter varies with the dose rate. In Figure 1 the dependence of the yield of the actinometer on the dose rate is shown, calculated according to the diffusion model, as discussed in our previous work. JI7e checked this theoretical curve against experiments at different dose rates and found that the calculated rate constants did not vary with the dose rates. From the correlation shown in Figure 1, one can now determine the dose rate, using 1 mM ferricyanide solution as the dosimeter. All the above reactions were carried out in the absence of buffers at near neutral pH, without the addition of salts other than potassium ferricyanide. Addition of up to 1 J I S a C l at neutral pH had no effect on the reaction rate. In the previous worki&it was shown that ferrocyanide is not reoxidized to ferricyanide by atomic hydrogen, even in acid solution in the presence of 0.3 N HISOJ. We have n o x reinvestigated and confirmed this. We a130 confirmed that the rate constant for the reduction is independent of pH between pH 2 and 9. It was of some interest to see whether in the reduction process replacement of one of the cyanide ligands by water occurs. This could be the case if in the reduction process atomic hydrogen removes one CK radical, forming H C S , resulting in the formation of a ferroaquopentacyanide ion, in a mechanism of reduction involving group transfer. Ferroaquopentacyanide can be determined by the "lor reaction with nitros~benzene.~We standardized this reaction by synthesizing Ka3Fe(CK)6.Hz0 according to the procedure of Hoffnian'O and confirmed the value of Ej28 5700 11-1 c m - ~giveIl in the literaturell for the reaction product with nitrosobenzene. The Journal of Physical Chemistry
GIL NAVONAND GABRIELSTEIN
i.t/
A (10% rnin:')
i
4
/
t
2tt.
/
/
/
t / 0
1 0.1
0.2
0.3
0.4
'''DA
0.5
per min.
Figure 1. Calibration graph for dosimetry using 10-3 M ferricyanide solution.
In experiments where up to 80% of an initial concentration of M ferricyanide solution was reduced to ferrocyanide, a t neutral or at alkaline pH, no pentacyanoaquo complex was found. In acid solutions, Ad) pentacyanocontaining 1 N H2S04,some aquo complex was found. However the same amount was formed when H atoms acted on the equivalent concentration of ferrocyanide, a t the same acidity. Therefore it is likely that the replacement reaction occurs not as a result of the reduction process, but as an independent reaction between atomic hydrogen and ferrocyanide. These results are of interest for the interpretation of experiments in the photochemistry and radiation chemistry of ferro- and ferricyanide solutions, where the formation of the pentacyanoaquo complex has been Recent on the photochemistry of ferrocyanide ion under conditions where H atoms are formed from eaq- supports the view that reduction of hexacyanoferrate(II1) by H atoms does not lead directly to C S displacement. F e ( d i p ~ ) ~and ~ + F e ( ~ h e n ) ~ ~Since + . the ferri complexes of a,a-dipyridine and of o-phenanthroline are not stable in neutral solution, unlike the ferro com(9) I. Murati and S. Asperger, Bnal. Chem., 33, 809 (1961). (10) G. Hoffman, Ann., 312, 1 (1900). (11) E. RIasri and If.Haissinsky, J . Chim. Phys., 397 (1963). (12) s. Asperger, I. L\luratj, and J . Chem. sot,, 730 (1960). (13) AT. Shirom and G. Stein, Nature, 204, 778 (1964). (14) G. Stein, Advances in Chemistry Series, No. 50, American Chemical Society, Washington, D. C., 1965, p 230; M. Shirom and G. Stein, t o be published.
.,
REACTIVITY OF SOMEHIGH-AND LOW-SPIN IRON(III) COMPLEXES
plexes, it was necessary to prepare them freshly for each set of experiments. This was done by using a stock solution of ferrodipyridyl or phenanthroline prepared from ferrosulfate and the respective compound. Before experiments an aliquot of this was made 1 N in sulfuric acid and treated a t 0" with excess PbO2 powder. The remaining Pb02 and the PbS04 formed were separated by centrifugation. The resulting oxidized solutions were kept below 5" until and during the experiments. Between preparation of the solutions and completion of the experiments not more than 1 hr elapsed. Under these conditions decomposition of the solutions was less than *5%. All reduction experiments for these two complexes were in the presence of 1 N HzS04. To determine the concentrations of the various complexes, when only one form was present, we made use of the absorption maxima and extinction coefficients for the oxidized and reduced complexes in both cases, as determined by S ~ t i n , whose '~ values we rechecked and confirmed. I n the case of the dipyridyl complex, the reduction mas followed by determining both the tervalent complex remaining and the divalent formed. For the divalent the absorption a t 522 mp was used, applying the relation CII = D ~ $ 3 6 5 0 . At 522 mp the tervalent absorbs negligibly. For the tervalent, a t 620 mp, me allowed for the weak but not negligible absorption by the divalent complex, using the expression CIII = (D620 - 205C11)/320. CII and CIII are the concentrations of the divalent and tervalent forms, respectively. In Table I1 the results are given as calculated from determinations of the divalent complex formed, since this procedure was the more accurate and reproducible one.
3633
I n the case of the phenanthroline complex, we found it convenient to use for the determination of the divalent complex the absorption a t 490 mp, rather than a t the peak of 510 mp. Go10,200, not much lower than a t the peak of 510 mp, while the absorption by the tervalent complex decreases from 300 to 4;: 100, so that it may be neglected a t the latter wavelength. The concentration of the tervalent complex was also determined, not a t the position of the maximum, but a t 620 mp, where ;:;E 814 is not much less than a t the maximum of 602 mp, while the absorption of the divalent decreases from 300 a t 602 mp to 83 a t 620 mp. We then calculated C I I ~= ( 0 6 2 0 - 83C11)/814. As in the previous case, good agreement was obtained between determinations based on divalent formed and tervalent that disappeared. Once again determinations based on the amount of divalent formed proved to be more accurate and reproducible, and these are shown in Table 111. Table 111: Reduction of Fe(phen),a+ ' Initial concn, 10-4 M
Reduction yield, 10-4 M min-1
Y/A
W
M -1 sec-1
8.46 8.04 4.18 2.65 1.70 1.34 0.875 0.80 0.507 0.40
2.34 2.38 1.98 1.33 0.895 1.05 0.72 0.59 0.436 0.25
0.840 0.854 0.714 0.564 0.322 0.378 0.259 0.212 0.157 0,098
0,072 0.065 0.17 0.39 1.6 1.75 2.6 3.8 6.8 15 5
2.7 3.2 3.5 3.2 2.0 2.9 3.6 2.9 3.3 2.0
lO-gk,
Av Time of reaction 1.0 min. min-1.
Table 11: Reduction of Fe(dipy)33+ a Initial concn, 10-4 M
14 7.4 5.9 3.0 1.80 1.71 0.99 0.57 0.52
Reduction yield, 10-4
2.9st0.5
Dose rate A = 2.8 X 10-4 M
lo-%,
M
M
-1
min -4
Y/A
w
sec -1
3.2 2.52 2.32 1.54 1.15 1.05 0.75 0.67 0.37
0.864 0.688 0.629 0.416 0.312 0.284 0,203 0.181 0,100
0.057 0.206 0.280 0.932 1.70 1.33 4.22 5.33 15.0
2.3 2.0 2.0 2.0 2.4 2.0 2.6 2.5 2.0
Av 2 . 2 k 0 . 2 Time of reaction 1.0 min. min-1.
In Table IV the values of the rate constants obtained for the three low-spin complexes now investigated are given together with that obtained3 in the case of the low-spin complex ferricytochrome-C a t neutral pH. They are compared with the rates of electron exchange with the ferrohexaaquo complex, when the rates are known. In the case of ferricytochrome-C the rate constant with H atoms is slightly pH dependent due most probably to pH-dependent changes in the protein moiety.
Dose rate A = 3.7 X lo-' M (15) 51. H. Ford-Smith and N. Sutin, J . Am. Chem. Soe., 8 3 , 1830 (1961).
Volume 70, Number 1 1
November 1966
3634
GIL NAVON AND GABRIEL STEIN
Table IV: Absolute Rate Constants for Reduction of Iron(II1) Complexes by Atomic Hydrogen" and Fez@
Fe3+ FeF2' FeSOa FeOH2 FeC12+
