534
HAROLD A. SCHWARZ [CONTRIBUTION FROM
THE
Vol. 79
CHEMISTRY DEPARTMENT, BROOKHAVEN NATIONAL LAUORAI'ORY]
The Radiation Chemistry of Ferrous Chloride Solutions BY HAROLD A. SCHWARZ RECEIVED AUGUST17, 1956 The radiation induced oxidation of ferrous ion in air-saturated 0.4 iM HC1 differs from the oxidation iu 0.4&I IizS04 iii that the ferric iron production is not linear with dose. The initial yield is essentially the same, G ( F ~ I I I(ions ) oxidized per 100 e.v. absorbed) being 15.8, but the kinetics indicate a competition between ferric iron and oxygen for H atoms. The effect of the competition is to decrease the yield, and the effect increases with increasing chloride ion concentration. The ratios of the rate constant for H atoms reacting with ferric iron and oxygen a t the various chloride ion concentrations have been separated into the ratios for the individual species present, giving kFe+++/ko, = 0.004, k F & + + / k O Z = 0.23 and k F e C 1 , + / ko, = 0.48. Some of the work was performed with added ferric iron Rigg, Stein and Weis9 have suggested t h a t the kinetics observed in the oxidation of FeS04 solutions initially, in which case the irradiated solution was compared directly with the unirradiated solution in the spectrophoby X-rays in the p H range 1-3 are affected by tometer, giving the change in optical density of the solution the presence of ions in solution capable of complex- instead of absolute values of the optical density. Thus ing the ferric ion. The addition of fluoride and small changes in ferric concentration of the order of 10-2070 phosphate ions t o the solution apparently makes the could be determined with a precision of about 1%. In order to check the spontaneous air oxidation of ferrous ferric iron less liable t o reduction by H atoms. ill FeClz in 0.4 M HCl was satchloride, a solution of The difference observed in the kinetics of oxida- urated with 0 2 and allowed to stand overnight. The initial ferric iron concentration was 52 fiM increasing to 57 p M in tion of ferrous ion by y-rays when the substrate is changed from sulfuric acid solution to hydrochloric 18 hours. Since all solution preparations and analyses were performed in less than six hours, this source of error acid solution apparently is also due to the form of is negligible. complexing of the ferric ion. In air-saturated Results and Discussion 0.4 M &Sod, ferrous ions are oxidized with a yield JI of 15.5 ions per 100 e.v. absorbed in the ~ y s t e m . ~ In Fig. 1, ferric production in solutions Ferric production is linear with dose until all but a in ferrous ion is given as a function of dose and of few per cent. of the oxygen in solution is used up.4 substrate composition. In 0.4 M HC1, air satuIn air-saturated 0.4 M HCI, though the initial rated, the rate of ferric production decreases as the rate of oxidation is essentially the same, the ferric dose increases. In oxygen-saturated solution, ferproduction is not linear with dose, falling off as the ric iron production is almost linear with dose in the range studied and the initial rates (calculated by a reaction proceeds. The present paper is an elucidation of the ki- method to be given later) are given in Table I. netics of the ferrous chloride oxidation in terms of the TABLE I nature and extent of complexing of the ferric ion. FERROUS IONOXIDATION YIELDS Conditions Yield Experimental Reagent grade materials were used without further purification. The water was distilled four times as described previously.6 Solutions were equilibrated with air by shaking and n i t h oxygen by bubbling. They were irradiated in a cylindrical Com source a t a dose rate of 3.02 X 10l6 e.v./cc.-min. The dose rate was determined with the ferrous sulfate dosimeter taking the yield of ferrous oxidation as G F ~ ( I I I(ions ) per 100 e.v.) = 15.45 =k 0.11.3 All analyses were performed on a Beckman DU ultraviolet spectrophotometer. Ferric iron was determined as the appropriate complex. A t 23.8', the follouing extinction coefficients were determined: in 0.4 M H2SO4, €305 2170; in 0.246 M HCl, 0.20 M HClO4, €335 1026; in 0.442 M HC1, t335 1358; in 0.434 M HCl, 1.00 M NaC1, 2010; in 0.441 M HC1, 0.60 M NaC104, e335 1366; in 0 40 M HClO4, €240 4185. The solutions containing 1 M C1- were diluted 1:1 uitli NaCl solution such that the final C1- concentration was 1.436 M . The extinction coefficient of iron in chloride solutions has a temperature coefficient of about 273 per degree6; hence a thermostated Beckmann was used and the samples were allowed to equilibrate before reading. The extinction coefficient in HC104 agrees well with that reported by Bastian, Weberling and Palilla; €240 4160.7 (1) Research performed under the auspices of the U. S. Atomic Energy Commission. (2) T. Rigg, G. Stein and J . Weiss, Proc. R o y . Soc. (London),ball, 375 (1952). (3) (a) C. J . Hochanadel and J. A. Ghormley, J. Chenb. Phys., 21, 880 (1953); (b) R.H . Schuler a n d A. 0. Allen, i b i d . , 84, 56 (1956). (4) Jerome Weiss, H . A. Schwarz and A. 0. Allen, International Conference on t h e Peaceful Uses of Atomic Energy, Geneva, Switzer]:and, 19.55. I > cY(Fe'")o; ;.e., ( 0 2 ) = (040- a(Fe"I)
I
(9) J. A. Ghormley and C. J. Hochanadel, Radiation Research, 3, 227 (1955). (10) W. G. Barb, J . H. Baxendale, P. George and K . R. Hararave, TYQW F a r a d a y Soc., 47, 462 (1951).
Fig. 2.-Rate of oxidation of ferrous ion in ferric~ as the reciprocal oxygen mixtures. G F ~ ( I I I ) - ~ GisHplotted coordinate: m, 0.43 M HC1, 1.00 M NaC1, air saturated; 0, same, 0 2 saturated; a, 0.43 M HC1, 0.60 M NaC1, air saturated; 0,0.43 MHCl, air saturated; 0, same, 0 2 saturated; (3, 0.43 M HC1, 0.6 M NaClOa, air saturated; A, 0.245 M HCl, air saturated; A, same, 0 2 saturated; 0 , 0 . 4 1 MHClOd air saturated; P I same, 0 2 saturated.
Milling, Stein and Weiss" have observed that the ferrous oxidation yield is increased by the addition of HCIO4 in deaerated solution. This effect amounts to 2 or 3y0 in 0.4 M HC104, and is attributed t o direct action of the radiation on the perchlorate ion. This effect is not evident in this work, perhaps due to the presence of oxygen. The slopes of the curves in Fig. 2 (which are proportional to k a / k b ) are seen to increase with increasing chloride ion concentration. The extent to which the ferric iron is complexed also increases (11) B. Milling, G Stein and Joseph Weiss, .2'atirre, 170, 710 (1952).
HAKOLD A. SCIIWAKZ
536
Vol. 70
with increasing chloride ion concentration. RabinTABLE I1 owitch and Stockmeyer have determined the as- SOLUTION COMPOSITION AND FERRIC-~XYGEX COMPETITION sociation constants for the first three chloride com- CONSTANTS AS FUNCTIOSS OF CHLORIDEION CONCENTRAplexes to be (at 1-1 = 1 and 25") K1 = 3.7, Kz = 1.1 TION AND IONICSTRENGTH and K , 0.03.6 They also determined the varia( H + ) = 0.40 111 tion of K1 with ionic strength. The relative solu- (C'L) LA Fe+'+ FeClt+ FeCh+ Feela kdkb tion compositions a t the various chloride ion concentrations (calculated assuming Kz to vary with 0.000 0.42 1.000 ... ... ... 0.004 p in a similar manner t o K1) are given in Table I1 .245 .45 0.445 0.436 0.118 0.001 ,159 along with values of k,/kb determined from Fig. 2. .443 .45 ,274 ,486 .237 ,003 .22G t
1.5
~
--___--~
~
T---
,436 1.037 1.436
1.03 1.05 1.45
,292 .10.i
.Os4
,471 ,403 ,313
,234
,477 ,607
,003 ,013 ,020
,218 .33F ,355
In Fig. 3, the first member of this equation is foulid to be linear with (FeClZ+)/(FeCl++). k J k b and ks/kb are found to be 0.23 and 0.48, respectively, from the intercept and slope. There is no large effect of ionic strength, which is not surprising, as the activity coefficients are probably near their minimum value. No errors are given for these values since it is difficult to assess the errors in the solution composition. The relative magnitude of L 1 these values undoubtedly is dependent on the activity coefficients of the complexes and reasons for 0 0.5 1.0 1.5 2.0 k2 being twice k1 are subject to criticism on this FeClz+/FeCl++. basis. Fig. 3.-Ferric iron-oxygen competition constant as a There is insufficient quantitative information on function of extent of single and double complexing: 0, p = rate constants of H atoms reacting with iron com0.45; e, = 1.05; n, = 1.45. plexes t o make a significant correlation or devise a In order to account for the reaction of H atoms general mechanism. Several qualitative observawith the various ferric complexes, the fourth step tions can be found in the literature. Rigg, Stein and Weiss studied the effect of fluoride and phosin the mechanism should be replaced by the series phate on the FeS04 oxidation a t high pH,* but unH + Fe+++--+ H + + F e + + ko fortunately their system was complicated by the I€ + FeCl+++ H + + Fe++ + C1- k l presence of the FeS04+ complex which they could H + FeCI2++H + + Fe++ + 2C1- kt not include in their calculations, although they susetc. pected its presence. However, their results indicate that k F e O H + + / k O , is of the order of unity while k , is then given by k F e F + + / k o l and kFepO,/kol < < 1. Since ferric k,(FexI1) = ko(Fe+++)+ kl(FeCl++) kt(FeCL+) 4- . . . iron production is a linear function of dose in the (2) irradiation of FeS04 in 0.4 ill until the 0 2 All terms above kz have been dropped, since the concentration is reduced to a few micromolar, it concentration of the higher complexed species in may be concluded that kFeSOh+/kOz < 10-3.4 solution is negligibly small. Fe(CK)6" production in k4Fe(CN)6 solutions is In HC104 solutions, the ferric iron is probably just equal to the Hz yield (G 0.4)12 indicating not complexed. Hence Ra/kb = k o / k b = 4 X that any oxidation by OH radicals is counter10-3. Since this is small compared t o k a / k b ob- balanced by reduction by H atoms. Tentatively, served in chloride solutions, we neglect ko(Fe+++) this suggests that kFe(Cxi),' is large. in equation 2 and find UPTON, N. Y.
+
-
(12) H . Fricke and E. J . Hart, J . Chem. Phrs.. 3, 596 (1935).
