Julie 20, 1952
TITANOUS I O N REDUCTION O F SODIUM
A N T H R A Q U I N O N E P-SULFONATE
3105
tion (Table I) for its half wave potential-pH rela- graphic wave in alkaline solution was investigated tionship is unlike the others and also in that a as a very sensitive check on this fact. Precipitasecond small wave is found (Fig. 1) after the main tion of an aluminum chelate as well as a decrease wave, It is possible that this additional wave in wave height was obtained with 3- and 4-methylcorresponds to reduction of the dimer formed as 8-quinolinols a t a pH of 12. No significant decrease the result of the first one electron reduction. An in wave height and no precipitation was observed unidentified dimer from the reduction of quin- with the 2-methyl and 2,3-dimethyl derivatives. aldine on mercury has been reported," and apparAcknowledgments.-This work was supported in ently a similar dimer was not found in the reduction part by a grant from the Research Corporation. of quinoline. This might account for the different The authors are also grateful to Joseph E. Seagram behavior of 8-hydroxyquinaldine compared to the & Sons, Inc., Louisville, Kentucky, in whose labother substituted 8-quinolinols. Since aluminum has been found not to form in- oratories most of this work was done, and also to J. Kimmel of the Seagram research departsoluble chelates with 2-substituted 8 - q ~ i n o l i n o I s , ~Edward ~~ the d e c t of adding aluminum ion on the polaro- ment for technical assistance with the polarograph. ( 1 1 ) V. V. Levchenko, Zhur. Obshchci K h i m . , 18, 1245 (1948).
[CONTRIBUTION FROM THE CHEMISTRY
LOUISVILLE 8, KENTUCKY RECEIVED DECEMBER 18, 1951
DEPARTMENT OF THE UNIVERSITY OF
CALIFORNIA, LOS ANGELES]
Oxidation-Reduction. 111. The Kinetics of the Reduction of Sodium Anthraquinone 0-Sulfonate by Titanous Ion and the Oxidation by Iodine BY CARLE. JOHNSON, J R , , AND ~ S. WINSTEIN RECEIVED AUGUST2, 1951 The kinetics of the reduction of sodium anthraquinone &sulfonate by titanous ion have been investigated spectroscopically.
In general, the rate-law for the production of the hydroquinone was found to be of the form: d(R)/dt = (Ti*II)[h(T)V:(R)l/2
+
+
k?(T)(R) K~(T)*/~(R)'/:] where T and R represent the quinone and hydroquinone, respectively. A mechanism has been proposed involving three reaction paths: (1) the reduction of S, the semiquinone, by Ti"'; (2) the reduction of S.S, a semiquinone dimer, by T W ; and (3) the reduction of TeS, a molecular complex composed of one molecule of the quinone and one molecule of the semiquinone, by TiIII, or such alternatives as the reaction between S and a complex ion containing one Ti111 ion and one T molecule. The rate of reduction of the quinone molecule itself by titanous ion is insignificant with respect to these three rate-determining reactions. The oxidation of sodium anthrahydroquinone @-sulfonateby iodine proved to be too rapid for accurate kinetic analysis. Under all of our experimental conditions it appears to be much faster than the reduction of the quinone by titanous ion.
The oxidation of hydroquinones by molecular rapid to lend itself to kinetic analysis by similar oxygen has been extensively investigated, princi- methods. pally by Weissberger and co-workers2 and by Lu0 Valle and Weissbergera who correlated their kinetic data on the basis of semiquinone theory. On the other hand, with the exception of the work of Dimroth4 on the relation between over-all free energy change and the rate of reduction of a series of quinones, little attention appears to have been paid to OH 0 the kinetics and mechanism of the reduction of I1 I quinones. For this reason, and because of the connection with the already reported5catalysis of the The Reduction by Titanous Ion.--Since the titanous chloride-iodine reaction by sodium anth- anthrahydroquinone I1 exhibits a sharp maximum raquinone p-sulfonate (I),we have studied separately in its light absorption a t 382 mp, while the anthrathe kinetics of reduction of sodium anthraquinone quinone I absorbs very much less and the titanous fl-sulfonate (I) by titanous ion and the oxidation chloride and other reagents are transparent in of the hydroquinone I1 by iodine. The first por- this region, the course of the reduction could be tion of this paper presents an analysis of the kinet- followed spectrophotometrically a t 382 mp. Figure ics of the reduction which was followed spectro- 1 shows optical density plotted vs. wave length for scopically. The last portion presents results ob- various combinations of sodium anthrahydroquintained for the oxidation which proved to be too one @-sulfonate I1 and titanous chloride in acid solution. It is seen that titanous ion, even at con(1) Department of Chemistry, University of Illinois, Urbana, Illinois. centrations higher than those employed in the rate (2) (a) T. H. James and A. Weissberger, THISJOURNAL, I S , 2040 determinations, and titanic ion have no specific (1937); (b) T. H. James, J. M. Snell and A. Weissberger, ibid., 80, effect on the absorption of the hydroquinone 11, for 2084 (1938); (c) A. Weissberger and G. Kornfeld, ibid., 61, 361 (1939); an identical absorption is obtained for the hydro(d) A. Weirsberger, J. E. LuValle and D. S. Thoma$, Jr., ibid., 86, 1934 (1943); (e) A. Wcissberger and J. E. LuVaIle, ibid., 89, 1676 (1947). quinone I1 whether produced by reduction of 1with ( 0 ) J. E. LuValle and A. Weissberger, ibid., 89. 1567 (1947). excess titanous ion or with zinc dust. Absorption (4) 0. Dimroth. 2. angew. Chem., 46 671 (1933). due to sodium anthraquinone 8-sulfonate (I) a t the ( 6 ) (a) C. E. Johnson, Jr. and S. Winstein, TEISJOUILNAL, 78, 2601 beginning of typical reductions amounted to ca. 3(1961); (b) ibid., 74, 756 (1952).
In reactions 1-12, excess titanous chloride was used so that the hydroquinone 11 produced a t infinite time was equal in concentration to the amount of \ quinone starting material. In reactions 13-18, 0.3 \ ' i excess sodium anthraquinone p-sulfonate (I) was used so that the concentration of the hydroquinone a t infinite time was limited by the amount of reducing agent added. The proportionality between AD and the hydroquinone concentration, (R), is shown by the constant value for AD/(R) which is recorded in the last column of the table. The possihility that the absorption measured at 382 mp is due in part to semiquinone or a similar intermediate 0.0 species seems to be ruled out by the fact that AD/ 300 400 500 600 7110 (R), as is shown in Table I, is not affected by the Wave length in millimicrons. presence or absence of excess quinone. SemiquinFig. 1.-Optical density 1's. wave length for socliuir~XIone concentration, (S), would be related to quinone thrahydroquinone 9-sulfonaLe: , sodiuni Linthra- concentration, (T), and hydroquinone concentrahydroquinone @sulfonate 5.00 X 10-s I I plus TiCla tion, (K), by equation 1, (S) having a maximum 0.0217 111; - - -, sodium anthrahydroquinone 6-sulfon- value when (T) equals (R). ate 5.00 X 10-5 X plus Tic13 0.010s;If; - - - - , sodium ( S ) = KB1/2(T)'/2(R)'/S (1) I
I
I
5
I -
anthrahydroquinone @-sulfonate5.00 X 10-6 i\f by reduc, TiCls 0.0217 A f ; tion of the quinone with zinc dust; , TiC13 0.0108 M. HCI was 0.325 M in all exses. Light path was 1.00 cni.
-----
------
5% of the final absorption a t 352 mp of the hydroquinone I1 after reduction was complete. If Beer's law is obeyed by both the quinone starting material I and the hydroquinone product 11, then AD, the change in optical density upon reduction, should be proportional to the hydroquinone concentration. Table I summarizes the results of several experiments which show that this relationship is obeyed. In each case, sodium anthraquinone p-sulfonate (I) was reduced by titanous chloride and the reaction allowed to go to completion. AD was obtained by subtracting the initial optical density from the final optical density a t 382 mp.
In measuring the progress of the reduction (R)t, the concentration of anthrahydroquinone I1 a t any time, t, was obtained from the relationship in equntion 2 (R)t = ( ADt / A D m )(T)o
TABLE I RELATIONBETWEEN INCREASE IN OPTICAL DENSITYUPON REDUCTION AND THE CONCENTRATION OF SODIUM ANTHRAHYDROQUINONE
(2)
where (T)o represents the initial concentration of sodium anthraquinone p-sulfonate (I). Since titanous chloride was always present in excess, the amount of R produced was always limited by the initial concentration of T. AD, was measured separately for each individual rate determination. The Rate Law.-When a reduction reaction is initiated, the optical density remains relatively constant for a short time and then begins to increase. The rate of change reaches a maximum a little short of the half-way point in the reaction and then the optical density gradually approaches a maxi-
0
0-SULFONATE PRODUCEDO
0
0
0.150 0
1 2 3
4 6 G 7 8 9 10 I1 19 13' 14' 15' 16' 17' 18'
20.00 5.00 2.50 2.50 2.50 2.50 2.50 2.50
ExcessC 0.608 ,809 Excess ,809 Excess Excess Excess Excess
.SO9 ,809 ,809
Excess
,405 ,201
Excess 2.50 Excess 2.50 Excess 2.50 Excess 2.50 Excess 40.00 2.58 20.00 2.64 20.00 2.64 2.58 20.00 20.00 2.42 2.74 20.00
20.00 5.00 2.50 2.50 2.50 2.50 2.50 2.50 2.50 2.50
,201 ,101 ,101 2.50 ,101 2.50 ,608 1 2 . 9 ,304 13.2 ,808 13.2 ,101 12.9 .lo1 12.1 ,201 13.7
0
0.00 1.387 .OO 0.345 .OO .165 .OO ,173 .OO .181 .OO .178 .OO ,172 .OO ,174 .OO .185 .OO .OO .OO
7.1 7.9
,171 ,175 .174 ,895 ,915 .Q10 ,895 .840
6.3
.950
27.1 6.8 6.8
6.94 6.90 6.60 6.92 7.24 7.12 6.88 6.96 7.40 6.84
7.00 6.96 6.94 6.93 6.89
0
2
4
"
0
0.100 0
+4
m
0
Q 0 0
0.050 0
6.94 6.94 6.93
e Ionic strength of 1.00 rt 0.01 M maintained by addition of KCl in all experiments. * D measured at 382 mp over a light path of 1.00 cm. A D is the change in optical density when the reduction is allowed to go t o completion. "EXcess" means more than enough t o reduce all the quinone present. d ( K I ) is 0.900 M . O(K1) is 0.164 M . ' ( K I )
i s 0.491 M.
0
0 C
0
oOooooo
0.000 10 15 Minutes. Fig. 2.--Optical density us. time curve for Run 7, Table 11.
5
June 20, 1952
TITANOUS IONREDUCTION OF SODIUM ANTHRAQUINONE &SULFONATE
mum value representing complete reduction to the hydroquinone. Optical density plotted against time has the general form shown in Fig. 2 which is a reproduction of the data for Run.7 of Table I1 which contains the details for all of the rate determinations carried out in this investigation. The speed of reaction was found to increase with increasing Ti"I concentration and with increasing initial anthraquinone concentration. TABLEI1 RATECONSTANTS AND CORRESPONDING CONCENTRATIONS FOR THE REDUCTION OF SODIUMANTHRAQUINONE 8SULFONATE BY TITANOUS CHLORIDE Run No.
(HM+)
( T i I I I ) (T) 104 M 106 M
1 0.101 2 . 8 3 2 . l o 1 2.77 3" .lo1 2.70 4 .201 2 . 8 3 5 .201 2.77 6 .405 5 . 5 4 7 .809 20.10 8 .809 2.83 9 .809 10.19 10 .SO9 20.38 11 .809 10.19 (KI) is 0.900 M.
2.50 2.50 2.50 2.50 2.50 2.50 2.50 2.50 5.00 2.50 2.50
ki
IO-Skz mole-' min.-1
I.* mole-2
min.-l
64 1 494 522 214.4 205.3 68.1 20.7 54.6 13.2 21.9 10.8
470 572 545 406 292 181 85.6 28.9 20.1 56.0 64.5
396 528 409 149 251 82.5 11.2 25.7 14.1 28.7 31.8
I. mole-1
1.2
10-6ki
min.-1
The D vs. time curves were analyzed by drawing a smooth curve through the experimental points with the aid of a flexible spline and measuring d(D)/ dt a t 5, 50 and 95% reaction with a tangent-meter. Values of d(R)dt were obtained with the aid of equation 2, and then values of d(R)/dt/(TilIT) were calculated for each of the points. The rate law shown in equation 3 was found to fit the data well.
+ K2(T)(R)+
d(R)/dt = (Ti"I)[K1(T)l/z(R)l/z
~ ~ ( T ) ' / s ( R ) ~(3) /P]
Since (T) and (R) are known a t any point on the D us. time curve, values for d(R)/dt, (Ti1I1), (T), and
(R) may be substituted in equation 3 a t 5 , 50 and 95% reaction yielding three equations which may be solved for kl,k2 and kB. The values of the constants are summarized in Table 11. By calculating d(R)/dt a t other points with the aid of these constants and comparing with observed values of d(R)/dt, it is seen that the constants reproduce the data well. Table 111 shows this comparison for a typical case (Run 4, Table 11). The calculated value of d(R)dt is tabulated beside the measured value for a number of points on the curve, the average deviation from the calculated values being 4.1%. A similar treatment of the data from Run 7 , Table 11, results in an average deviation of 5.55&. The constants k1, kz and k 3 were not tested for fit in every case. Presumably slightly better results could be obtained by a more laborious least squares treatment, but this was not carried out. Reference to Run 3, Table 11,indicates that substitution of potassium iodide for most of the potassium chloride does not change the rate constants significantly. With regard to the effect of hydrogen ion concentration on the rate con'stants, k l , kz and k3, a11 appear t o be roughly proportional t u I / (H+), but the relationship is not precise. This
3107
dependerlce on 1/(H+) parallels the analogous dependence on l/(H+) of the rates of reaction between titanous ion and iodine! TABLEI11 CALCULATED RATEOF PRODUCTION OF R COMPARED WITH THE EXPERIMENTAL VALUEAT 19 DIFFERENT STAGES OF REACTION.MEASUREMENTS TAKENFROM RUN 4 OF TABLEI1 No.
(TIIII) l o 4
M
(T) lo5
M
(R)
105 M
2.82 2.47 0.0270 2.82 2.45 .0541 3 2.81 2.42 .OM1 4 2.81 2.39 .lo82 5 2.80 2.34 ,162 6 2.78 2.23 ,270 7 2.73 1.96 ,541 8 2.67 1.69 .811 9 2.62 1.42 1.082 10 2.56 1.15 1.35 11 2.55 1.082 1.42 12 2.49 0.811 1.69 13 2.43 .541 1.96 14 2.38 ,270 2 . 2 3 15 2.36 .162 2 . 3 4 16 2.35 .lo82 2.39 17 2.35 .0811 2.42 18 2.34 .0541 2.45 19 2.34 .0270 2.47 Units are moles liter-' minutes-'. 1 2
IO'd(R)/df calcd.
106d(R)/dfa exptl.
0.50 .75 .94 1.11 1.40 1.86 2.60 2.95 3.02 2.85 2.78 2.36 1.78 1.05 0.71 .52 .42 .32 .20
0.51 .79 .99 1.15 1.48 2.01 2.54 2.89 2.99 2.86 2.77 2.32 1.66 0.96 .68 .51 .41 .30 .18
Mechanism.-The work of Michaelisa and others' has shown that the following equilibria are commonly encountered in quinone-hydroquinone mixtures. T + R z 2 S
s + s )r s.s Kd
T
+S
Kt.8 T.S
Here K , is the semiquinone formation constant, Kd is the semiquinone dimer formation constant, and K t , is the equilibrium constant for the formation of a T-S molecular complex from one quinone and one semiquinone molecule. Equations 1, 4 and 5 follow from these equilibria.
(s6)= kd(S)* = &K.(T)(R) (4) (T.S) = Lt.,(T)(S) = ~ ~ . ~ K ~ ' / ~ ( T ) * / z ( R( 5) ')/ z Referring now to the rate law, equation 3 and equations l , 4 and 5 above, it will be noted that: kl(T)1/z(R)1/2 is proportional to (S) and kl includes K,'/z; k2(T)(R) is proportional to ( S 6 ) and k2 includes &&; ka(T)a/z(R)l/ais proportional to (T,S)and KBincludes Kt and K s l / l . On the above basis the three contributing paths to the reduction reaction indicated by the rate law in equation 3 may be identified. The rate determining step corresponding to kl may be regarded (6) (a) L. Michaelis, Trans. Ekctrochem. Soc., 71, 107 (1937); (b) A n n . Rev. Biochcm., 7 , 1 (1938); (c) Cold Springs Harbor Symposium
on'QuantifofiucBiology, VII,33 (1939). (7) (a) W. M. Clark, B. Cohen a n d H . D. Gibbs, U.S.Publ. Health Kepts., Suppl. No. 64 (1926); (b) E. Frirdheim and L. Michaelis. J . Bsol. Chcm., Si, 356 (1031).
:1 1 os
CARL15. JOHNSON,
JR., AND
as a reaction hetween titanous ion and semiquinone. 'Jill1 1. s -....+ Ti"' + K The reaction corresponding to kz may be regarded ;is a reaction between titanous ion and semiquinone
dimer
- S.S -+ TiIV + S + K The reaction corresponding to k3 involves titanous Til11
ion, a quinone molecule, and a semiquinone species. Kinetics do not distinguish between such alternatives as reaction between titanous ion and a quinone-semiquinone complex
'rim -4- 7r.s
----f
l y v 4-'1.
+
I
