KINETICS OF THE OXIDATION OF IRON(II)BY BROMINE
Oct., 1952
877
THE KINETICS OF THE OXIDATION OF IKON(I1) BY BROMINE B Y PAUL
R.
CARTER A N D NORM.4N
DAVIDSON
Gates and Crellin Laboratories of Chemistry of the California Institute of Technology, Pasadena, California Received March IS, 1068
The kinetics of the oxidation of iron( 11) by aqueous bromine has been investigated spectrophotometrically at an ionic strength of 1.00 maintained with sodium ions and perchlorate ions. In solutions with (H+) = 0.5-0.8 M , (Br-) = 0.05-0.5 = 10-6 MI (Fe+++)o= 0-10-3 111, the rate law is -d(Z Brz)/dt = k2(Fe++)(Br3-)[l M , (Fe++)o and (E (k7(Br-)(Fe+++)/ka(Fe++))]-l;ke = 34.0 (A0.7)(liter/mole sec.) a t 29.8', 21.8(A0.3) at 20.4', kz = 107.69(+0.a0ex ( -8400( *500)/RT); (k7/k3) = 0.11( AO.01) a t 29.8". The probable mechanism is given by equations (2) and (3). T f e inhibition by ferric ion is strong evidence for this reaction path and eliminates an alternate path involving iron(1V) as the unstable intermediate. The available thermodynamic estimates give for reaction (2), AF'zss = 8.8kcal., AH'zss = 1.0. In solutions with ( H + ) = 0.14.9 MI (Br-) = 1-3 X l O - 8 M , there are addiiional terms in the rate law, k6(Brz)(Fe++) kl4(HOBr)(Fe++),for reactions of molecular bromine and hypobromous acid; kg = 0.76(A0.2); kl4 = 2.l(f0.4) X lo4 a t 29.8'. Phosphoric acid catalyzes the reaction; in solutions with ( H + ) = 0.1-0.7M , (Br-) = 0.2-0.5 M , (H3P04)= 0.1-0.5 M , the main additional rate term a t 29.8' is 2.4(&0.4) X lo3 (Br3-)(Fe++) (&POI-).
+
+
The reaction 2Fe++
+ Bra- (or Brz) +
2Fe+++
+ 3Br- (or 2Br-)
(1)
is of interest because it presumably proceeds via the two one-electron steps kz
+ Br- + Br2F e + + + Br2- +Fe+++ + 2Br-
Fe++
+ Bra-
Fe+++
(2)
kr
k3
(3)
involving the formation of the unstable species, Br2- (or a bromine atom). The kinetics of the oxidation of iron(11) by aqueous hydrogen peroxide1p2and by i ~ d i n e and , ~ of titanium(II1) by iodine4 indicate that these reactions proceed by analogous mechanisms involving OH radicals and Iz-. The induction of the isomerization of maleic to fumaric acid by the reaction of dipositive iron with aqueous bromine is evidence for the occurrence of bromine atoms or Brz- in reaction (1).6 Several reactions due t o chlorine atoms or C12- have been induced by the addition of iron(I1) to chlorine solutions.6 Reactions of OH radicals, induced by the addition of dipositive iron to hydrogen peroxide (Fenton's reaction), have been extensively studied. Several observers have made the qualitative observation that the reactions of aqueous bromine or chlorine with dipositive iron are not "instantaneou~."~-'~We have accordingly found that it is possible to measure the rate of reaction (1) by rapid manual operation of a Beckman spectrophotometer when the reagents are present a t concentrations in the range 10-3-10-5 hd. The kinetic results con(1) F. Haber and J. Weiss, Proc. Roy. SOC.(London), 8147, 332 (1934). (2) J. H. Baxendale, M. G. Evans and G. 9. Park, Trans. Faraday Soc., 42, 165 (1946);W. G. Barb, J. H. Baxendale, P. George and K. R. Harprave, ibid., 47, 462 (1951). (3) A. V. Hershey and W. C. Bray, J. A m . Chem. Soc., 68, 1760 (1936); this paper gives references to the many significant earlier investigations of this reaction. (4) C. E. Johnson and 6 . Winstein, ibid., 78, 2601 (1951): D. AI. Yost and 8. Zabaro, ibid., 48, 1181 (1926). (6) F. Wachholts, 2. Elektrochcrn.. 88, 545 (1927). (6) H.Taube, J . A m . Chem. Soc., 66, 1876 (1943); 68, 611 (1946). (7)'P. Farrington, Ph.D. thesia. California Institute Tech., 1950. (8) H.von Halban and H. Eisner, Helu. Chim. Acta, 18, 724 (1935). (9) R. N. J. Saal, Rec. trav. chirn., 4'7, 73 (1928). (10) A. W. Francis ( J . A m . Chem. Soc., 48, 655 (1926)) concluded, on the basis of competition experiments with organic reducing agents, that the rate constant for the reaction between aqueous FeSO, or FeClr and Bra is ca. 4 X 104 (liters/mole sec.). This number is a factor of 10' greater than the rate constant obtained by us.
firm the reaction path consisting of steps (2) and (3) above, provide some information about the rates of reaction of molecular bromine and hypobromous acid as well as of the tribromide ion, and disprove an alternate mechanism based on the formation of tetrapositive iron as an intermediate. Experimental Materials.-Weighed quantit,ies of ferrous ammonium sulfate were dissolved in 1 M perchloric acid. There was less than 2% air oxidation per month of these ca. 10-3 M Fe++ solutions (iron(II1) in the solutions was estimated after addition of phosphoric acid by the ultraviolet light absorption of the iron( 111)-phosphate complexes).** Weighed quantities of sodium bromate were dissolved in an excess of sodium bromide and perchloric acid to give the tribromide ion solutions. Solutions of molecular bromine, prepared by the addition of the C.P. liquid to mixed wdiumperchlorateperchloric acid solutions, were standardized iodometrically. C.P. sodium bromide was dissolved in distilled water and a small amount of bromine was added. Excess bromine was boiled away and the solution filtered. Iron(II1) perchlorate (Fe( C104)3.9H20)was prepared by dissolving ferric chloride in 60% perchloric acid and fuming. After two recrystallizations from water, the product wm free of chloride ion. Solutions in perchloric acid were standardized iodometrically. Sodium perchlorate was prepared by neutralization of C.P. sodium bicarbonate with perchloric acid. Since sodium perchlorate is frequently used as an inert salt to adjust ionic strength, it is important to note that the G. F. Smith Co. anhydrous sodium perchlorate was unsatisfactory for our studies in that 0.5 M solutions were about 10-4 N in oxidizing impurities that reacted with Fe++. Procedure.-About 1.5 ml. of an iron(I1) solution was rapidly discharged into a 1-cm. path length, glass stoppered, quartz cell containing an equal volume of a bromine solution, or vzce versa. The iron(I1) solution was added from a calibrated pipet driven by a hypodermic syringe. The pipetting gave almost complete mixing, but the cell was stoppered and shaken for a few seconds to ensure mixing, and then placed in a thermostated cell compartment of a Beckman model DU spectrophotometer. The first optical density reading was usually obtained about 24 seconds after mixing, and the half-time of a typical experiment was about 3 minutes. The response time of the spectrophotometer to small changes in optical density is lcss than a sccond. Extinction Coefficients.-The formal extinction cocflicients, E = D/( Z Br2), of bromine solutions, as measured in this research, are listed in Table I.
TABLE I FORMAL EXTINCTIONCOEFFICIENTS, e, OF BROMINEAT 29.8' IN 0.5 M HC1O4, 0.5 ikl NaBr (e = (l/cl) loglo l o / Z ) h (md t (liter/mole em.) X lo-*
x c
360
340
330
320
310
300
(1.97 1.48 2.70 4 21 6.62 10.2 290 280 270 265 260 255 250 245 15.9 24.3 31.8 32.2 29.3 23.7 16.7 10.8
(11) P. R. Carter and N. Davidaon, to be published.
P. R. CARTER AND N. DAVIDSON
878
The data of Griffith, MoKeown and Winn12 for the dissociation constant of the tribromide ion, when extrapolat,ed t o p = 1.0 and T = 29.8', give KD = (Br2)(Br-)/(Br3-) = 0.0615 mole/liter. The molecular extinction coefficients of the tribromide ion are therefore 1.123 times the e's of Table I. (The contribution of molecular bromine to the light absorption listed in Table I is negligible.) The formal extinction coefficients of bromine in bromide solutions at 20.4" agreed with those calculated from the above data usin KD = 0.0570. The formal extinction coefficient of iron(III7 in 0.5 M NaBr-0.5 M HClO, was determined as 275. This value was used in correcting the observed optical densities for the small contribution due to iron(II1) formed in the reaction. The molecular extinction coefficient of bromine in 0.5 M HClO4 and 0.5 M NaC104 at 452 mp was measured as 93.6. The absorption spectra of iron(II1) in phosphoric acid will be reported later.11
Results and Discussion All reaction rates were measured at the constant ionic strength of 1.0, maintained with sodium ions and perchlorate ions, after the desired acidity and bromide ion concentration were selected. The kinetic results for any particular run at a constant acidity and bromide ion concentration conformed t o the rate law -d(ZBrZ)/dt
= k(ZBrz)(Fe++)
(4)
where ( 2 Brz) is the total (formal) concentration of bromine. Figure 1 displays the results of a typical run.
solutions with nitrogen did not affect the results. The rate constants are independent of hydrogen ion concentration over the limited range studied. The values of kl in Table I1 are calculated from k by the relation kl = k (1 4- KD/(Br-)), where KD, the dissociation constant of the tribromide ion at p = 1.0, is taken as 0.0615 a t 29.8" and 0.0570 at 20.4'. The constancy of kl, when the bromide ion concentration is varied from 0.05 t o 0.5 M , shows that, in this concentration range, the tribromide ion is the reactive bromine species. TABLE I1 RESULTSAT (Br-) = 0.05, 0.2 AND 0.5 M , 29.8 and 20.4' All runs were at p = 1 .00,maintained with Na+ and ClOa-; all runs were at (H+) = 0.50, unless otherwise noted; spcctrophotometric observations a t h = 300 mp.
h
3
Br-) nioles/liter
(ZBrz) (Fe++) Iiioles/liter X 106
6.27 6.27 7.42 12.54 12.54 12.54"+ 12.54"' 6.27 6,27' 12. 54d 12.54 18.80 25.09
1.06
VOl. 56
a
T = 29.8" 2.15 0.479 4.30 ,483 5.62 .483 2.27 .479 4.56 ,483 4.58 .483 4.48 ,483 for (Ur-) = 0.5, kl (av.) 5.36 0.196 5.10 .196 3.92 ,209 for (Br-) = 0 . 2 , IC] (av.) 12.93 0.0498 12.54 ,0498 12.89 ,0498 for (Br-) = 0.05, Icl (av.)
F1
T
v
A 0.90
k
ki
litors/iiiolo see.
29.6 33.4 30.5 34.4 30.6 34.5 29.6 33.4 29.6 33.3 30.4 34.3 31.0 34.9 = 33.8 f.0 . 5 26.5 34.8 25.3 33.3 26.6 34.4 = 34.2 f 0 . 7 14.9 33.1 15.7 35.1 15.3 34.2 = 34.1 =t0 . 9
= 20.4'
0.500 19.0 21.2 .500 18.5 20.7 .500 19.4 21.7 ,500 19.7 22.0 ,500 19.6 21.9 ,500 19.7 22.0 k i (av.) = 2 1 . 8 f 0 . 3 M. 1.33 X A!. 0 (SOa') = 6.G X ( H +) = 0.80 M. * ( H +) = = e Excluded froill uverage. 6.27 6.27' 12.54 12.54 18.81 18.81
+0)
5 M I
0.86
3.18 5.14 3.21 5.10 3.18 5.10
&s$J.
0.76
0
25 50 75 100 200 300 Fig. 1.-Rate of reaction of iron(I1) with aqueous bromine: 20.4", 0.5 M NaBr, 0.5 M HCIO,, (Fe++)o = 1.88 X 10-4 M , ( Z Br2)o = 3.18 X 111 0,29.8', 0.5 M HClO,, 0.5 M NaC104, (Br-)o = 1.02 X 10-3 M , (Fe++)o= 4.15 X lO-SM, (E Br& = 3.39 X 1 0 - 3 M ; the vertical log (Fe++)/(Z Br2) for this run. coordinate is 1
-8-
-e-,
The values of ICIa t 20.4' (21.8 f 0.3) and 29.8" (34.0 f 0.7)have been used t o derive the equation kl =
107.'69(M.3@J exp
(-8400 ( f 5 0 0 ) / R Y )
The errors quoted are probable errors, not estimated safe limits of errors. Inhibition by Ferric Ions.-The results of the The Uninhibited Tribromide Ion Reaction.-The previous section are consistent with the reaction pertinent results are collected in Table 11. At a path of equations (2) and (3) and the assumption fixed bromide ion concentration, the constant k is that ka >> Icz. (A mechanism composed of reaction independent of the initial concentration of re- (2) followed by Br2Brz- ---f Br3Br- is actants. Experiments with small amounts of consistent lvith the rate data, but implausible.) added sulfate showed that the amount added as If this be true, the rate constant, ICl, of Table I1 ferrous ammonium sulfate did not affect the rate; is equal to k,. Assuming that, AF and AH for the check experiments, not listed here, showed that, reaction 1/2 Rrz 4 73r arc the same in solution as there was no photochemical reaction clue to the in the gas phase, one calculates for the reactionsI3 light of the spectrophotometer. Deaeration of the (13) "Selected Values of Che~riicrl Therniodynsinic Properties,"
+
+
(12) It. I). Criffitli, A . Rlolicown aiid A. Yoc., 88, 101 (1032).
G. Winii, Truns. Parudity
+
Natiriiial I31iroau of Standards. Waslririgtoti, I). C.; Scriob: 1, 'I'ablc 11-1, Juiic 3 0 , 1048; l'aldc 47-1, Junc 30, lY,i!?.
Oct., 1952
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