+
erium(IV) Man anese(Ii) z (Hi) System 6,A. ~ ~ C ~G. ~N. IRAO, ~ and ~ , 6. l P. R A 0 2 Department of Chemisiry, Stale University of New York, Buffalo, New York
b A spectrophotometric investigation of the Ce(lV) -/- Mn(ll) Ce(lll)
+
Mn(lll) cross-reaction has revealed the rate expressons
The kinetics of oxidation of manganese(I1) by cerium(1V) have been briefly studied by hspray, Rosseinsky, and Shaw ( 2 ) , by following the appearance of manganese(II1) spectrophotometrically; but no detailed investigation of the important reversible h'ln(I1) $ Ce(II1) systeni Ce(1V) hln(1II) has been reported, although a few kinetic studies of reactions involving manganese(II1) as an oxidant may be found in the literature (6, 13, 16). In this paper we report the results of a detailed kinetic study of the cerium-manganese systeni in acidic sulfate media, from the viewpoint of both the forward and the reverse reaction; a combination of these results with the corresponding data yields a full description of the system and should be useful for the selection of optimum analytical conditions.
+
for the main forward path, and
for the main backward path, where the apparent second order rate constants for the forward and backward reactions are 0.43M-1 sec.-l and 5.9M-1 sec.-I, K@SpeCtiVely,at 25' C. in 3 M H&31. From the solvent dependence of the rate expressions, the principal oxidizing species appear to be Ce(S0&, in the forward reaction, and Mn(OH)+2in the backward reaction. The measured equilibrium constant. Cor the cross-reaction k in good agreement with the quotient of the apparent rate constants in the various media studied. HE IMPORTAKCE of kinetic considerations in determining the stoichionietry and general utility of analytical oxidation-reduction reactions has been ever more widely recognized (16) in recent years and has led to the realization that detailed kinetic studies must be performed on key solution reactions if their full analytical potential is to be attained. Among the most frequently used reactions in analytical chemistry are those involving cerium(1V) as an oxidant; for this reason, we have previously investigated the reaction of cerium(1V) with oxalate (8),and other workers have studied the reactions of this oxidant with iron(I1) (f , Y ) , chromium(II1) (IS),tin(I1) ( 5 ) , mercury(1) ( I $ ) , and uranium(1V) (3).
Alfred P. Sloan Fellow.
2 Present address, Indian Institute of Science, Bangalore, India.
e
ANALYTICAL CI-IEMISTRY
I42 7 4
+
EXPERIMENTAL
All chemicals used were of reagent grade. The concentration of ceric sulfate solution (G. F. Smith Chemical Co.) was determined by titration with standard ferrous ammonium sulfate, as was the concentration of potassium permanganate solution used. Manganese(II1) ( 2 , 16) solutions were prepared by oxidation of manganese(I1) with potassium permanganate in the presence of excess manganese(I1) in strong sulfuric acid medium. Under these conditions the disproportionation of manganese(II1) t o manganese(I1) and manganese(1V) is effectively suppressed, and such solutions were found to be stable for at least five days. The concentration of manganese(II1) solutions were determined by titration with standard ferrous ammonium sulfate solution. A Beckman NIodel DB spectrophotometer equipped with thermostat and a Photovolt Model 43 recorder was used to record absorbance us. time plots. The disappearance of cerium(1V) in the cerium(1V)-manganese(I1) reaction and the appearance of cerium(1V) in the reverse reaction were followed by monitoring the absorbance of cerium ( 1 1 7 ) a t 400 mp (molar absorptivity &Z 1000). Beer's law was followed satisfactorily at this wavelength by cerium (117) solutions. Because manganese(II1) also absorbs to a small extent at this 45), wavelength (molar absorptivity
the absorbance readings were corrected for the contribution of nianganese(II1) in all experiments. The experimental procedure consisted of placing 2 ml. of cerium(1V) solution containing requisite reagents in the spectrophotometer cell and injecting, with a syringe, 1 nil. of manganese(I1) solution. Cerium(II1) solutions were injected into manganese(II1) solutions in the study of the reverse reaction; 90% of the reaction occurred in the first one minute in most of the experiments. Kinetic data were analyzed by the usual methods and reduced to rate constants by means of graphical methods. All experiments were conducted a t 25' =t0.2' C., unless otherwise stated. Each of the k values given is a mean of three experimental runs. The average deviation in the rate constants is estimated t o be 5 5 % and that in the eauilibrium constants to be 510%;. The concentrations of hydrogen ion, bisulfate ion, and sulfate ion in the reaction media have beer calcL!lajed employing the 1.alue oi the di:*sociation constant of biiuliate ion computed from the Raman spectral ivork o i Smith , I ? ) . It is assumed that the dependence of this dissociation constant on ionic strength is nor altered by the presence of perchloric acid, cerium, nianganesp: and sodium ions in solution and that perchloric acid and sodium bisulfsre are completely ionized. RESULTS
The Cerium(1V)-Manganese(I1) Reaction. The rate law for the overall reaction is that of a first order forward reaction opposed by a second order reverse reaction when the concentration of manganese(I1) is kept constant and in excess ( 2 ) . The relevant empirical rate expression is given by
where [C], [ D ] , and [ E ] are the concentrations of cerium(1V) , cerium(II1) , and manganese(I1I) at time t , and kl and k z are the apparent first order and second order rate constants, respectively. Since [Do] = [E,] = 0, [ D ] = [ E ]= C, = C a t time t , where o subscript
refers to concentrations a t the beginning of the reaction. If C, is the equilibrium concentration of cerium(1V) , kl
. c, =
kz . (C,
- C,)2
(2)
a t equilibrium. Substituting for the value of k z in Equation 1, integrating, and rearranging yields
Hence, the plot of log CO2- c.C,/ C,(C - C,) us. t should be a straight line according t o this equation. Figure 1 shows the plot of log Co2 - C.C,/ C,(C - C,) for a typical experiment. The predicted rate lam is followed over 80% of the course of the reaction. Additional results obtained a t different initial concentrations of cerium(IV), keeping other conditions constant, are summarized in Table I. The agreement between the rate constants indicates that the reaction is first order with respect to cerium(1V). The results of experiments conducted with different initial concentrations of manganese(I1) keeping other conditions constant are given in Table 11. The direct proportionality between the rate constant and the concentration of manganese(I1) indicates that the reaction is first order with respect to manganese(I1) also. The equilibrium constants show good agreement with those obtained at different initial concentrations of cerium(1V) , indicating that equilibrium is attained in all the experiments. Effect of Temperature. The effect of the temperature on the rate constant was studied in the presence of both 3,OOM and 1.50X sulfuric acid. The reaction follows the Arrhenius equation in the temperature range studied. Figure 2 shows the plot of
c2 - C'C, vs. time for cerium(IV)-manganese (II) reaction
Plot of log 2
Figure 1 .
CdC
log k l , us. 1/T, where T is the absolute temperature (kl, in 1. mol.-1 sec.-l) in 3.0041 sulfuric acid. ( k l , is the apparent second order rate constant for the forward reaction.) Brrhenius activation energy values obtained in presence of 3.00M and 130M sulfuric acid are 13.7 and 12.6 Kcals., respectively, the corresponding frequency factors being 1010.44 and The values of entropy AS* and heat AH" for the activation process are calculated to be -17 e.u. and 13.0 Kcal. in 3.00X sulfuric acid at 25' C. employing the equations
kl, =
k.T h ~
. K*,
=
TAS" and AG*
- R T In K*
--1
0.60-
0 El,
'
=
-
Effect of Sulfuric Acid and Ionic Strength. The effect of variation of
sulfuric acid concentration from 1.OO to 3.00X on the rate constant kl is shown in Table 111. S o attempt is made t o keep the ionic strength constant in these experiments. The rate constant decreases with increasing concentration of sulfuric acid under these conditions. The rate constant is found to decrease with increasing
Table 1. Effect of Cerium(lV) Concentration on Rate Constant (&I) n'langanese(I1) = 66.7mM, sulfuric acid = 3.00M, T = 25" C., p = 5.2
Concn. of Ce(IV), mM 10 k l , sec.-l
I
0.80-
-
AH*
(where lc, h, T , K * , AG", AH", R, and AS" have their usual significance),
1.00,
3
- C,)
Cerium(lV) = 1.036mM, monganere(l1) = 66.7mM, sulfuric acid = 3.00M, T = 25' C.
0.40-
020-
33
34
35
J- x IO4 T
Figure 2. Plot of logarithm of rate constant kl, vs. reciprocal of temperature Cerium(lV) = 0.691 mM, manganese(l1) = 66.7mM, sulfuric acid = 3.00M
a
Equilibrium condtant, Kfa
0,415 0.289 0.553 0,293 0.691 0,289 1.036 0.288 1.175 0,298 1,382 0.306 [Ce(III)! [Mn(III)]
KI
=
0,055 0.0% 0.060 0.058 0.055 0.057
[Ce(IT')][Mn(II)] '
Table It. Effect of Manganese(l1) Concentration on Rate Constant (kl)
Cerium(1V) = 0.691mMJ sulfuric acid = 3.00M, T = 25" C., p = 5.2 Equilibrium Concn. of constant, Mn(II), mM 10 k l , see.-' Kf 20.0 26.7 33.3 40.0 53.4 66.7
VOL. 38,
NO. 13,
0.095 0.124 0.143 0,191 0.247 0.289
DECEMBER 1 9 6 6
0.061 0.058 0,056 0.061 0.057 0.060
e
1901
1.20
1.00
0.80 YN
-
Cll
0
Figure 3.
I
I
I
I
I
I
I
0.10
0.20
0.30
0.40
0.50
0.60
0.70
Plot of log
- X*XJ ( x , - x)aZ
x,[a2
0.80
0.60
0.40
0.20 YS.
time for manganese0
(III)-serium( 111) reaction Ceriurn(lll) = 1.70mM, rnangonese(1ll) = 1.70mM, sulfuric acid = 3.00M, manganese(1l) = 66.7rnM,
ionic strength a t any given concentration of sulfuric acid (see Table IV). Sodium perchlorate is used to vary the ionic strength in these experiv-ents. Effect of Hydrogen Ion Goncentration. The results of experiments conducted with different concentrations of sulfuric acid and sodum bisulfate while keeping the sum of bisuifate and sulfate concentrations
Table 111. Effect of Sulfuric Acid Concentration on Rate Constant (kl)
Cerium(1V) = 0.691mM, manganese(I1) = 66.7mL21, T = 25" C. Concn. Equilibrium of sulfuric constant, acid, A I 10 kl, sec.-l Kl 1.00 0.986 0.125 1.25 0.758 0.134 0.622 0.122 1.50 0.449 0,081 2.00 0.340 0,064 2.50 0.289 0.060 3.00
33
34 I 7 x
T = 25' C.
IO4
Figure 4. Plot of logarithm of rate constant temperature
constant a t 3 . 0 X are given in Table V. The concentrations of bisulfate ion and sulfate ion unavoidably vary to a small extent in these experiments, because of the dissociation of bisulfate ion. Keglecting these small variations, the rate constant kl is found to vary linearly with the concentration of hydrogen ion. Effect of Bisulfate Concentration. Mixtures of perchloric acid and sulfuric acid are employed for this purpose while keeping the ionic strength a t 5.2. The concentration of hydrogen ion varies to a small. extent unavoidably in these experiments. The rate constant is found to vary linearly with the inverse of bisulfate ion concentration. This relationship, of course, is indistinguishable from an inverse dependence of rate constant with [S04-2]. The results are given in Table VI.
Effect of Ionic Strength on Rate Constant
Cerium(1V) = 0.691mM, manganese(I1) = 66.7mM, T Concn. of sulfuric acid, M Ionic strength 10 kl, sec.-l 0.622 1.50 2.7 3.2
1.50
1.50 2.00 2.00
3.7 3.6 4.1
2.00
4.6
0.35s 0,283
0,449 0.3'23
0.273
=
(k~)
Cerium (111)-Manganese (111) Reaction, The rate law for this reaction is obtained a t different initial concentrations of cerium(II1) and manganese(II1) as follows. If u and b are the initial concentrations of manganese(II1) and cerium(II1) and 5 is the amount of cerium(1V) formed after time t , then dCe(1V)
--at
- k z ( ~- ~ ) ( -b 2 ) - k i ~ (4)
The reaction between cerium(1V) and manganese(I1) is effectively made first order by using an excess of manganese(I1). At equilbrium,
k z ( ~- x8)(b-
e
=
ANALYTICAL CHEMISTRY
2,) =
klze
(5)
where xe is the concentration of ceriuni(1V) a t equilibrium; substituting for kl in Equation 4, integrating and rearranging, one obtains
(x, -
Z)U
*
2,)
.b
=
kz
(U
. b - $2) Xe
25' C.
Equilibrium constant, Kf 0.122 0.054 0.033
0.oSi 0,063 0,034
Effect of Hydrogen Ion on Rate Constant (kl) 0.69lmM, manganese(I1) = 66.7mM, T = 25" C., JI, = 5.3 Concn. of Rate const,ant Equilibrium sodium Concn. of 10 k l , see.-' constant, K f bisulfate, Af sulfuric acid, AI 0.060 0.289 0.00 3.00 0.044 0.246 0.50 2.50 0.023 0.210 2.00 1.00 0.017 0.165 1.50 1.50 Table V.
Cerium(IV)
kz vs. reciprocal of
Cerium(ll1) = 1,70rnM, rnanganese(ll1) = 1.70mM, sulfuric acid = 3.00M, rnanganese(l1) = 66.7mM
In $,(a. b - 2 Table IV.
35
This equat'ion simplifies to
In 2 , ( d - 2xe) = ( 2 , - x)u2
kz
(a2 - xa2) Xe
t
(7)
when the initial concentrations of manganese(II1) and cerium(II1) are equal-i.e., when a equals b. A plot of log x,(u' - Z X , ) / ( X , - E ) U * us. t should be a straight line according to this equation. Such a plot for a typical experiment is shown in Figure 3. I n general, the rate law is obeyed over 80% of the course of the reaction. Apparent rate constants calculated from slopes of such plots for different initial concentrations of manganese(II1), cerium(III), and manganese(I1) are pre-
sented in Table VII. The agreement between the rate constants is satisfactory and shows that the reaction is first order with respect to both manganese(II1) and cerium(II1). The initial presence of manganese(I1) has no appreciable effect on the rate constant. The internal agreement between equilibrium constants is fair. Effect of Temperature. The systen1 obeys the Arrhenius equation in the temperature range (12' to 35' C.) studied. The plot of log k2 us. 1/T, where T is the absolute temperature, is shown in Figure 4; Arrhenius activation energy calculated from the plot is 16.5 Kcal., and the frequency factor is 10l2.8*in a medium of 3.0J4 sulfuric acid. The values of entropy and heat of the activation process are calculated to be -1.7 e.u. and 15.9 Kcal., respectively. Effect of Sulfuric Acid and Ionic Strength. I n presence of 3.00, 2.50, and 2.00.11 sulfuric acid, the values of rate constant obtained a t 25' C. are 5.89, 5.77, and 6.16 1. mol.-' sec.-l, respectively. No attempt was made to maintain the ionic strength constant in these experiments. The rate constant is found t o increase with increase in ionic strength at any given sulfuric acid concentration when the ionic strength is varied by the addition of sodium perchlorate. Effect of Hydrogen Ion Concentration. I n Table VI11 are shown the results of experiments carried out a t varying hydrogen ion concentration, keeping the total concentration of sulfate and bisulfate constant at 3.0M. The rate constant kl is found t o vary linearly with the inverse of hydrogen ion concentration in the concentration range studied. Small changes occurring in the concentrations of bisulfate ion and sulfate ion are considered to have no appreciable effect on this result. Effect of Bisulfate Ion Concentration. The results obtained with mixtures of perchloric and sulfuric acid, keeping the ionic strength a t 5.2 are given in Table IX. The rate constant is found to vary linearly with the inverse of bisulfate ion concentration. This result is experimentally indistinguishable from an inverse relationship of the rate constant with [S04-2].
Equilibrium Constant. It is gratifying to note that equilibrium constants for the reaction Mn(II1)
+ Ce(II1) + Rln(I1)
+ Ce(1V)
obtained experimentally starting with Mn(II1) and Ce(II1) (15.0 f 1.0) agree well with the value calculated from experiments starting from Ce(1V) and Mn(I1) (16 f 1.0). The value
computed from the ratio of the individual rate constants is slightly lower (g13.5). A value for the equilibrium constant can also be calculated from the standard potentials of the cerium (111)-cerium(1V) couple (- 1.44 volts) and the manganese(I1)-manganese(II1) couple (-1.51 volts) (11). From the equation
Table VI.
Effect of Bisulfate Ion on Rate Constant (kl)
Cerium(1V) = 0.691mM, manganese(I1) = 66.7mM, T = 25" C., p = 5.2, hydrogen ion = 4.2kf Concn. of Rate bisulfate constant, Equilibrium ion 111 10 k ~ sec.-l . constant. K-, I 2.05 0.289 0 060 1.63 0,341 0.057 1.34 0.382 0.062 1.02 0.436 0.063 I
where E o is the potential of the cell and K the equilibrium constant, the value for the equilibrium constant for the oxida-
Table
VII. Apparent Second Order Rate Constants
(k2) for Oxidation of Cerium(ll1) by Manganese(ll1) in Sulfuric Acid Medium
T = 25" C., sulfuric acid Cerium(III), mJf 1.67 1.67 1.67 1.67 1.67 1.67 1.67 1.70 2.21 2.55 2.89 3.40
Table VIII.
=
3.0011.1, p
hIanganese(III), Manganese(II), mM mM 1.28 12.8 1.70 17.0 3.40 34.0 0.85 66.7 1.28 66.7 1.70 66.7 2.55 66.7 1.70 66.7 1.70 66.7 1.70 66.7 1.70 66.7 1.70 66.7
=
5.2
le^, lit. mol.-1 sec.-l
Equilibrium constant, Kb'
5.61 5.59 5.59 5.76 5.73 5.84 5.62 5.89 5.78 5.57 5.90 5.56
15.4 15.7 16.5 15.2 15.6 14.2 16.1 14.4 14.1 14.0 14.3 14.7
Effect of Hydrogen Ion Concentration on Rate Constant (k2)
Cerium(II1) = 1.70mM, manganese(II1) = 1.70mM, T = 25' C., manganese(I1) = 66.7A1, p = 5.3 Concn. of Concn. of sodium bisulfate Equilibrium sulfuric acid, added, I1.l kp, 1. mol.-1 sec.-' constant, Kb 3.00 0.00 2.75 0.25 2.50 0.50 2.25 0.75 2.00 1.00 1.75 1.25 1.50 1.50
tion of Ce(II1) by iCln(II1) is calculated to be 15.3, which agrees well with the experimentally observed values. Though the AH value for the reaction Ce(II1) F), Nn(I1) Xn(II1) Ce(1V) is positive (2.9 Kcal.), the change in entropy of f 1 5 e.u. helps the reaction to move t o the right hand side.
+
+
DISCUSSION
Hardwick and Robertson (10) have shown that cerium(1V) exists mainly as Ce(S04)a-2 in sulfuric acid solutions of concentration greater than one molar, and described the major equilibrium involved as
Table IX.
EfFect of Bisulfate Ion on Rate Constant (kz)
Manganese(II1) = 1.70mM, cerium(II1) = 1.70mM, manganese(I1) = 33.3mM, T = 21' C., p = 5.2, hydrogen ion = 4.2M Rate Concn. of constant kz, bisulfate 1. mol. -1 Equilibrium ion, M set.-' constant, Kb
VOL. 38, NO. 13, DECEMBER 1966
1903
Kal
Ce(S04)s-2
a t 25" C. in 3.00M sulfuric acid, compares well with the value reported by Aspray, Rosseinsky, and Shaw (2), 0.215 1. mol.-1 sec.-l a t 20.2' in 4.50X sulfuric acid, considering the variation of experimental conditions. The equilibrium constants also agree with the values of Aspray et al. Dulz and Sutin (7) conipared the rate constants experimentally obtained for the cerium(1V)-iron(I1) and cerium(1V)-iron(I1) phenanthroline reactions, with those calculated from the Marcus theory, and found reasonable agreement. It would be interesting to attempt a similar comparison for the cerium(1V)-nianganese(I1) cross-reac-
+ H+ +
Ce(SO4)z 3. HSO4-
(8)
where K,, = 5 X We assume Equation 8 to be major equilibrium involved in the present study for cerium(1V) because the molar absorptivity values found in presence of 1.5 to 3.OM sulfuric acid for cerium (IV) agree well with those reported by Hardwick and Robertson for Ce(S04)s-2 species under similar conditions, Cerium(II1) has also been shown to be complexed with sulfate in sulfuric acid solutions, the predominant species at high concentrations of sulfuric acid being C ~ ( S O P ) ( ~6-). ilIanganese(II1) has been shown to exist mainly as hlnF+2 in solutions containing perchloric acid and fluoride by Fackler and Chawla (9). I n view of the ability of manganese(II1) to complex with ions like OH-, F-, and C1-, we may assume that manganese(II1) exists predominantly as hInS04+species in sulfuric acid solutions employed in the present study. In the concentration range investigated, the rate constant of the cerium(1V)-manganese(I1) reaction is found to be linearly related to the hydrogen ion concentration and to the inverse of bisulfate ion concentration. These solvent effects are consistent with a mechanism involving Ce(S04)~ as the major reacting species; the predominant reaction path may then be formulated as
[LInON+z] =
-+ hIn(I1) +
*
+
dt
kla[Ce(S03z1[NIn(II)1 (10)
and because the concentration of [Ce(S04)2]in turn is given by
[Ce(IV)l[H+l. Ke, (11) [HSOI-I through expression 8, the over-all rate law for the forward path may be given by -dCe(1V) dt
which corresponds to the observed rate law. The value of the apparent second order rate constant (ha) obtained in thc present work, 0.435 1. mol.-1 sec.-1
4
e
ANALYTICAL CHEMISTRY
+ HzO Kee
2
hPnS04+
tion studied here. According to the X b r c u theory, the relationship k a b = (kaka Kabf)"2 (13)
because
kakb where logf = (log K,b)'/4 log - (14)
and
* H + + s04-'
KHSOa-
HSO4-
22
-dLCIn(ll )
dt
+
MnOH+2 HS04- (15) the concentration of [MnOH+2]would be given by
=
[H
+
1[S04-2 I
KHsoa-
(18)
The reaction is formulated
+
kzb
nlnOHZ+ Ce(S04)z--+ activated complex -, Ce(1V) LIn(I1) (19)
+
Thus, the rate will be given by
- k 2 b [RlnOH+2]
=_ dCe(lV) ___-
dt
[HSO4-]
(17)
~
[ce(So,),-]
(9)
The rate of the forward reaction will thus be given by -dCe(IV)
[S04-21
I
where kaband Kab are the rate constant and equilibrium constant of the oxidation-reduction reaction, k , and kb the rate constants for the individual electron exchange reactions, and z is the collision frequency of two uncharged molecules in solution (taken as 101lM-l sec.-l), may be used t o predict the reaction rates of appropriate reactions. er
activated complex Ce(II1) NIn(II1)
[H+l
tration range investigated. This dependence may be accounted for if NInOH+2 is assumed to be the major reacting species in these solutions. Assuming the equilibrium in sulfuric acid solutions of manganese(II1) as
[MnS04+][HzO]* K,, - K,, [R/In(III)] [HSOh-] [HS04-']
kl.
Ce(S04)t
The rate constant of the reverse reaction between manganese(I.11) and cerium(II1) is found to be linearly related 1 1 to __ and also t o ___ in the concen-
A knowledge of the cerium(1V)ceriuni(II1) and manganese(II1)-manganese(I1) exchange rate constants in sulfuric acid medium a t 25' C. is needed for the calculation, as the corresponding equilibrium constants are available from this study. Diebler and Sutin (6) calculated the rate constant for manganese (11)-manganese(II1) exchange reaction in 3F perchloric acid a t 25' C. as 3 X lo-4F-1 sec.-l from indirect experiments. Assuming this value to be also applicable in 3 X sulfuric acid and employing 4.4X-1 sec.-l for cerium(IV)ceriuni(II1) exchange rate constant (7), the rate constant for the cerium(IV)manganese(I1) reaction is calculated to be 0.75 X 10-2 1. mol.-1 sec.-l and may be compared with the experimental value of 0.435 1. mol.-l sec.-l In view of the many assumptions which had to be made in finding the calculated value, the agreement is rather good and gives some support to the view that cerium (1V)-manganese(I1) reaction may be of the outer sphere type in sulfuric acid media.
which is in accordance with the observed rate law. The formulation of [MnOH+2] as the reactive species is reasonable, because manganese(II1) is known to be hydrolyzed even in presence of six molar perchloric acid (9, 19). Taube (14) has pointed out that in the reaction of CrClf2 with manganese(III), manganese (111) chooses a path involving OH- in preference to C1-, even when the acidity is high. Hence it is not surprising that R.InOH+z is the reactive species in sulfuric acid solutions of manganese (111). Once again, the specific rate constant may be compared with a value calculated from the Marcus theory. Using an exchange rate constant for the manganese(I1I)-manganese(II) couple determined for perchloric acid medium, the calculated rate constant for the manganese(II1)-cerium(II1) reaction is 0.3M-' sec.-l, which compares well with the experimental value 5.9M-l sec.-l
(7) D L ~G., , Sutin, N., Inorg. Chem. 2 ,
LITERATURE CITED
~F. s.,~ i 1 worth, P., Trans. Faraday SOC. 61,
(1) Adamson,&I.G., ~
(1965).&I. J., Rosseinsky, D. R., Bspray, (2)689 Shaw, G , B., them. ~ ~ ( d ~. ~ 1963, p. 911 (3) Baker, F. B., Newton, T. W., Kahn, 31.)J . Phys. Chem. 64, 109 (1960). (4) Blatz, L. A, Ibid., 66, 160 (1962). (5) Brubacker, H., jr,,court, 8.J , , J . Am. Chem. SOC.78, 5530 (1956). (6) Diebler, H., Sutin, N., J . Phys. Chem. 68, 174 (1964).
91'7 (1963). ~(8) ~ El-Tantawy, ~ -~ P. ~ A , , Rechnitz, ~ G. , A., ANAL.CHEM.36, 1774 (1964). (9) Fackler, J. P., Jr., Chawla, I. D., Inorg. Chem. 3, 1130 (1964). ~(10)dHardwick, ~ ~ T. ) J., Robertson, E., Can. J . Chem. 29, 828 (1951). "Oxldatbn Po(11) Latimer,
tentials, Prentice-Hall, 2nd ed., New Pork, 1952. (12) nIcCurdy, IT. H., Jr., Guilbault, G. G., J . Phys. Chem. 64, 1825 (1965). (13) Nicol, ?If. J., Itosseiiisky, D. R., Chem. I n d . (London) 1963, p. 1166.
c,
ent for Automated Synth
(14) Ogard, A. E., Taube, H., J . Phys. Chem. 62, 357 (1958). (15) ltechnitz, G. A,, ANAL. CHEX 36, 453R (1964): 38, 513R (1966). (16) Itosseinsky, D. R., J . Chem. SOC. 1963, p. 1181. (17) Smith, J. H., "T? Structure of
Electrolytic Solutions, W.J. Hamer, ed., p. 51, John Wiley and Sons, 1959. (18) Tong, J. Y. P., King, E. L., J . Am.
Chem. SOC.82, 3805 (1960). (19) TJ-eh, c. F., Davies, G., Nature 205, 692 (1965). RECEIVEDfor review May 20, 1966. Accepted September 19, 1966.
is of Peptides
R. B. MERRIFIELD, JOHN MORROW STEWART, and NILS JERNBERG The Rockefeller University, New York,
N. Y.
7 0027
b An instrument which can perform automatically all of the operations involved in the stepwise synthesis of peptides b y the solid phase method is described in detail. The synthesis of the peptide chain takes place on a solid polymer support and all of the reactions are conducted within a single vessel. The apparatus is composed of two main parts-the reaction vessel and the components required to store, select, and transfer reagents, and the programmer which controls and sequences the operation of the various components. The operation of the instrument and its application to the synthesis of several peptides are described. in methods O f isolation, purification, analysis, and structure determination of peptides and proteins have outdistanced our synthetic achieyements in this area. To cope with nianj- of the new problems which haye arisen, a greatly accelerated and simplified approach to peptide synthesis was required. Solid phase peptide synthesis was devised ( I S ) , and developed (9, 10) with these objectives as guides. The principles of the method and the special features which make it adaptable t o an automated process have been reviewed (11, IS), and an apparatus designed for automated peptide synthesis was constructed and briefly discussed (16). This article describks in detail the instrument which can perform automatically all of the operations involved in the stepwise synthesis of polypeptides. XTENSIVE BDVANCES
GENERAL PRINCIPLES
The method is based on the fact that a peptide chain can be synthesized in a stepwise manner while one end of the chain is covalently attached to an
Figure 1.
Apparatus for automated peptide synthesis
insoluble solid support. During the intermediate synthetic stages the peptide remains in the solid phase and can therefore be manipulated conveniently without significant losses. All of the reactions, including the intermediate purification procedures, are conducted within a single reaction vessel. It is this feature which permits convenient automation of the process. The problem in essence is simply to introduce the proper reagents and solvents into the vessel in the proper sequence a t the proper times. The solid support is a chloromethylated styrene-divinylbenzene copolynier bead. The C-terminal amino acid is coupled as a benzyl ester to the resin and the peptide chain grows one residue a t a time by condensation a t the amino end with N-acylated amino acids. The tert-butyloxycarbonyl group has been
the protecting group of choice and activation has usually been by the carbodiimide or actil-e ester routes. Since each of the reactions in the synthesis can be modified in a variety of ways it was important t o design the apparatus with sufficient flexibility to cope with a wide range of reactions and conditions. APPARATUS
The apparatus is composed of two main parts, the first being the reaction vessel with the components required to store and select reagents and t o transfer them into and out of the vessel, and the second being the programmer which automatically controls and sequences the operation of the various components. A photograph of the complete instrument is shown in Figure 1, a schematic drawing is given in Figure 2, and the wiring diagram of the proVOL. 38, NO. 13, DECEMBER 1966
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