Kinetics and Mechanism of Extraction of Zinc, Nickel, Cobalt, and Cadmium with Diphenylthiocarbazone, Di-o-To1 yl t hioca r bazo ne, a nd Di-a- Naphthylthiocarbazone

B. E. McCLELLAN and HENRY FREISER

Department of Chemistry, The University of Arizona, Tucson, Ariz.

,The kinetics of extraction of zinc, nickel, cobalt, and cadmium with di- phenylthiocarbazone, di-o-tolylthio- carbazone, and di-a-naphthylthiocar- bazone were studied. Each of these reactions were first order in metal ion and, with one exception, first order in reagent anion. The reaction of zinc with di-a-naphthylthiocarbazone was second order in reagent anion. The rate constants for each metal ion paralleled the rate at which water dissociates from the hydrated metal ion. In every case sterically hindering substituents on the phenyl rings of the reagents increased the rate of the reaction. A mechanism is proposed in which the rate-controlling step is the reaction involving the change in con- figuration of the monodithizonate complex from six- to-four-coordinate metal ion, accompanied by the disso- ciation of water. Values of the dis- tribution coefficients, KD,, of the rea- gents between CHC13 and water are reported.

N ADDITION to the use of solvent I extraction as a means of studying chelate formation equilibria, in favorable cases it provides a simple technique for the determination of rate constants of chelation reactions which are quite rapid. This is made possible by maintaining the concentration of ligand low by virtue of high reagent distribution coefficient (Knr) and p H control. Thus, when a dithizone solu- tion in chloroform of 10-4M is equili- brated with an aqueous phase, the aqueous dithizone concentration is lower by a factor of about lo6, the value of the KD,. The ligand, dithizonate anion, concentration would be still lower a t p H values less than the pK. for dithizone. The significance of the ability to control ligand concentration a t such low levels is, of course, that reactions have rate con- stants as large as lO7JI-'sec.-' and can he observed at reasonable time intervals (2-60 minutea) without elaborate ap- paratus.

In an earlier paper (4, the course of the reaction of zinc(I1) with dithizone was studied. The results indicated the probable importance of the role of rupture of zinc-water bonds in the rate- determining step. This suggested further study along two lines: first, to extend the zinc study to other dithi- zones, in which substituents are so situated as to make steric hindrance a factor in the overall chelate structure (the effect of steric hindrance on chelate formation thermodynamics has been previously investigated but rarely its effect on kinetics); second, to extend the study to other transition metal dithizonate extractions to learn whether the kinetics of these reactions could be correlated with the rate of water loss from the respective metal cations ( 2 , 3 ) .

EXPERIMENTAL

Apparatus. All extractions were performed in 50-ml. vials fitted with polyethylene liners and plastic caps. The samples were agitated on a n Eberbach box type shaker a t the low speed setting. The sample agitation was somewhat more vigorous than the maximum attainable using the Rurrell shaker (4 ) . The insulated shaker box was kept a t constant tem- perature (i 1' C.) by means of a copper coil through which water from a con- stant temperature bath was circulated. The coil was welded to a copper plate which served a$ the base of the shaker box. Measurements of radioactivity were .made with a Nuclear-Chicago Model DS5-5 well-type scintillation detector whose output was registered on a Model 192h scaler. A Beckman Model G pH meter was used for all pH measurements. This meter was stand- ardized against Beckman pH 7 buffer before each set of measurements. Spectrophotometric measurements were made with Beckman Model DU and DR spectrophotometers.

Materials. Specially processed, high activity zinc-65 in the form of ZnC12 was obtained from the Oak Ridge Yational Laboratory. The ac- tive zinc solution was diluted so that the final solution for extraction con-

tained a total zinc concentration of 2.5 X 10-6AlI and a chloride concentra- tion of 1 X 10-7.11. Reagent grade nickel perchlorate, cobalt perchlorate, and cadniiuni perchlorate obtained from G. F. Smith Chemical Co. were used to prepare 1 x 10-3.1f solutions of the metal ions. The isotopes C060

and Cd1I5 were obtained from the Oak Ridge Sational Laboratory. Reagent grade dithizone wa5 further purified by the method cited by Welcher (9). The 2,2'-dimethyl and di-a-naphthyl derivatives of dithizone were prepared and purified by standard methods in this laboratory by K. S. ;\lath. The reagent solutions were prepared by direct weighing of the carefully purified and dried reagents. The stability of the solutions was checked a t regular intervals by spectrophotometric methods using the following data:

Dithizone h max. = 606 mp, t = 4.06 x 104

2,2'-Dimethyl derivative h max. = 610 mp, t = 4.95 X lo4

Di-a-naphthyl derivative h max. = 685 mp, E = 2.15 X IO4

Dithizone solutions were stable for only about a week, while the 2,2'-di- methyl derivative solutions showed no evidence of decomposition even after a month. The di-a-naphthyl derivative solutions were very unst,able, and had to be prepared each day. All other chemicals were reagent grade. Water was purified by passing it through a mixed bed ion exchange column supplied by Crystal Research Laboratories, Inc., Hartford, Conn.

Partition of Reagent. Equal vol- umes of a solution of the reagent in chloroform and an aqueous phase consisting of phosphate buffer and sufficient sodium perchlorate to give an ionic strength of 0.25 were shaken for 15 minutes. The solutions were allowed to s tand for 10 minutes to ensure complete phase separation. The p H of the aqueous phase was measured. A portion of the organic phase was then transferred by pipet to a spectrophotometric cell and the absorbance, A , was measured.

2262 ANALYTICAL CHEMISTRY

From a plot of

values for K,/KD, were obtained (3). Determination of Kinetics of Reac-

tion. The extraction procedure was altered somewhat from that of Honaker and Freiser (4) . Their method in- volved taking micro diquots, at definite time intervals, on a single sample for a kinetic run. The standing time after agitation affected the results of the kinetic study. Since a t least a 5-minute standing period is necessary for phase separation, their method leads to large variations in the sta,nding time follow- ing agitation. Therefore the procedure was altered so that a separate sample was used for each time interval. This allowed each sample to stand the same length of time following agitation. The detailed procedure for zinc follows.

A11 extractions were made from aqueous solutions having an ionic strength of 0.25 produced by the ad- dition of sufficient tiodium perchlorate to the buffered aolutions. Acetate buffers were used for all the kinetic studies with zinc, but phosphate buffers were sometimes usNed for the other metal ions. All extractions were per- formed a t 25" =t :lo C. The initial aqueous zinc concentration in all cases was 2.5 X 10-6X. This zinc solution contained enough Zn"6 to give a count of about 5000 c.p.m. per ml. Similar concentrations and count rates were used with Co and Cd extractions. With N, 2.0 x ~ O - ~ J Z scilutions were used, since the Xi in the aqueous phase after extraction was spectrophotometrically determined using dimethylglyoxime ( 8 ) . The p H of the aqueous phase was ad- justed to the desired value by addition of the appropriate buffer solution, and its value was checked with the p H

Table I. Kinetic Data for Extraction of Zinc with Dithizone

[ H n ~ l o [H+] X l o4 x 104 Slope" kIh 0.97 2 82 0 0018 0.012 1 94 2 82 0 0040 0 013 3 00 2 sa 0 0078 0 017 5 00 2 82 0 0160 0 021 1 00 1 26 0 0082 0 024 3 00 1 26 0 0280 0 027 3 00 0 49 0 0630 0 024 5 00 0 49 0 1220 0 028 5 00 0 71 0 0750 0 024 2 00 1 32 0 0196 0 030 2 00 1 32 0 0160 0 024 2 00 4 90 0 0022 0 012 2 00 2 37 0 0048 0 014 2 00 I 26 0 0134 0 019 2 00 0 74 0 0258 0 022 2 00 0 45 0 0366 0 019 0 10 0 16 0 0090 0 033 0 10 0 06 0 0104 0 015

k'av = 0 021 u = 6 3 X l W 3

Q Slope of a plot of log ~ [ ~ ~ ~ ~ o us time.

. [H+l k' = 2.303 X slope X __ [Hnz)o

met,er. A 10-ml. aliauot was shaken for a definite time intkrval (from 2 to 60 minutes) with 10 ml. of a chloroform solution of the reagent. The time of reaction, measured with an electric timer, was taken from the time agita- tion was begun until it was stopped. Each sample was then allowed to stand 5 minutes to ensure complete phase separation. X 5-ml. aliquot was trans- ferred to a plastic tube for counting. The sample was then counted for a length of time necessary to obtain a total count of a t least 10,000. The ratio of the count rate of the sample to the count rate of the same volume of the original metal solution gave the metal concentration in the aqueous phase. A typical set of extraction data is shown in Table I.

Table II. Distribution Characteristics of Dithizones between CHC13 and

Water at 25" C.

slog Ka/ log Compound K D , pKab K D ,

Dithizone 10.18 4 . 3 5 . 9 2,2 '-Dimethyl-

dithizone 12.00 4 .7 7 . 3 Di-a-naphthyl-

dithiocarbazone 12 30 3 . 7 8 . 6

a Ionic strength = 0.25. Values of ( 6 ) lowered by 1.5 units

to correct the observed value obtained in dioxane-water mixture to aqueous medium. This correction is approxiniate but per- mits a valid comparison of these three related reagents.

RESULTS AND DISCUSSION

Distribution of Reagent. T h e dis- tribution of dithizone and its deriva- tives between a n aqueous and a n organic phase involves two equilibria according to the Nernst distribution law

K D , HDz(,) e HDz(,); K D , =

[HDzI, [HDa l w

[H+I 1Dz-I [HDzI,

-__

K. HDz(,, & H + + Dz-; K. =

The expression for the distribution of the reagent, D,, or distribution ratio, is

[HDzI, Dr = [HDz], + [Dz-lm

from which it can be shown that

Although very high values of D,, requiring very sensitive spectrophoto- metric assay of the low reagent con- centration in the aqueous phase, would be needed to determine KD,, a series of intermediate values of D, as a function of pH is more readily accessible and permits the determination of K a / K D , (the slope of a plot of 1/D, us. 1/ [H+]).

The values of Ku/KD, obtained in this manner for dithizone, the 2,2'- dimethyl, and the di-&-naphthyl deriva- tives of dithizone are shown in Table 11. Values of K D , that have been calculated from this data with the aid of inde- pendently measured K , values (6) are also included. The distribution coef- ficient of dithizone is quite high. The effect of methyl substitution is to in- crease the value of log K D , by 0.7 unit per carbon atom which is similar to that (0.6) found for a number of reagents. As expected, when the substituent is part of an aromatic ring, the rise in log KO, per additional carbon atom is smaller (0.34) in the naphthyl de- rivative.

Kinetics of Extraction. The reac- tion orders with respect to the metal ion were determined from plots of log

[Me+2], = 0 [Me+2It

us. time. This yielded straight lines, indicative of a first-order reaction with respect to metal ion in all cases. The change in the values of the slopes of these lines obtained a t various pH and [HRl0 values were used to obtain re- action orders with respect to [HR], and [H+] (Table 111).

Although the reaction orders of these extractions with respect to metal ion concentrations were exactly unity, the orders in reagent and hydrogen ion con- centrations deviate from this value. The reason for this deviation is con- sidered to arise from the greater un- certainty of the determination of these orders in all but the reaction of zinc with the di-a-naphthyl derivative where the orders are f 2 in reagent and - 2 in [H+I.

The steps involved in the reaction for the extraction of zinc with dithizone are as follows.

Table 111. Reaction Orders for Metal Dithizonates

R ~ ~ - Reac,tion order in Metal gent" [HR], IH +I Zinc A 1 15 -1 17

B 1 30 -1 13 C 1 91 - 1 70

Nickel A 1 10 -0 82 B 0 80 -1 04 C 1 3.5 -0 78

Cobalt A 0 9G -0 7 2 C 1 53 -0 7<5

A, B, and C refer to dithizone, its di-2-rnethy1, and di-a-naphthyl deriva- tives, respectively.

VOL. 36, NO. 12, NOVEMBER 1964 e 2263

HDz(,, e HDz(,, (1)

(2)

(3)

HDz(,, e H+ + Dz-

Zn+Z + Dz- e ZnDzT

ZnDz+ + Dz- e ZnDzz(%) (4)

ZnDzzcw] $ ZnDzz(.) (5 )

Honaker and Freiser ( 4 ) reported the forward reaction of step 3 to be the rate-controlling reaction. Substituting the equilibrium expression for steps 1 and 2 in

d [ Z I I + ~ ] dt

~~ - - k[Znt2].[Dz-] or

-d[Znf21 k'[Znf21 [HDzl, yields -~ - - - dt H +

(7 )

-d[Zn+2] K , d t Kn ,

= k - X

(8) [Zn+2] [HDz],

[H+I

Hence, dividing the value of k' by K , ' Ko, yields the rate constant, k , for the reaction. The values for k are 3.7 x 10*J31-'min~te-~ for dithizone, 4.5 x 10*M-lniinute-l for 2,2'-dimethyl diphenylthiocarbazone, and would be .till larger for di-cu-naphthj-lthiocarba- zone if the kinetic order were the same as for the other two. ,ilthough the zinc dithizone is considerably more stable (log K,, = 4.7) (these K , values are 1.5 log units lower than observed in dioxane-water mixture to make them more applicable to aqueous equilibria) than the methyl-substituted zinc chelate (log K J L = 3.0) ( f i ) , the zinc dithizonate does not form more rapidly than does the 2,2'-dimethyI diphenylthiocarba- zone chelate. In this case, the expected relationship between the stability of the chelate and the rate of formation is reversed. Thilky groups substituted on the phenyl rings in dithizone increased the rate of formation of the zinc che- lates.

The failure of the thermodynamic disadvantage of the presence of the sterically hindering substituents to be reflected in the rate of chelate formation

indicates that step 3 might be more complex than first envisioned. Several steps are involved in the formation of the chelate: loss of water molecules from the octahedral zinc hexahydrate, addition of the first and second ligands, and change of structure from the octahe- dral hexahydrate to the tetrahedral zinc dithizonate. In aqueous solution, zinc exists as the octahedral hexahydrate ( 7 ) . Coordination reactions of metal ions would involve either loss of co- ordinated water or expansion of the coordination sphere (1). The former alternative seems more likely in the case of zinc. Eigen (3) reports the rate of removal of water molecules from the zinc ion to be greater than loe minute-'. The rate a t which water molecules can be removed from the zinc ion would be expected to be the limiting rate of formation of the zinc chelates.

The following mechanism for the formation of the zinc dithizonates is presented as a modification of that given by Honaker and Freiser (4) :

Zn(HzO)s+2 + Dz- e Kf 1

Zn(H20)4Dz+ + 2Hz0 (9)

ki Zn(H20)4Dz+ -

Zn(HZ0)2Dz+ + 2Hz0 (10)

ki Zn(H20)zDz+ + De- -

Zn(Dz)z + 2H20 (11)

This alternative mechanism better explains the experimental results. First, the substituted dithizone derivatives show an increased reaction rate with zinc. The explanation for this might be that the bulky groups facilitate the re- moval of water molecules after the addition of the first ligand. Also, steric strain is relieved in changing from the octahedral to the tetrahedral structure. As k l increased, kz might be expected to become more significant in the overall rate expression. This would explain what has happened in the case of the di-a-naphthyl derivative. The rate equations for this mechanism are as follows:

-d [Zn+21 = kl[Zn(HzO),Dz+] dt

Table IV. Comparison of Rate Constants for Formation of Metal Chelates with Dithizone and Related Compounds with Rate of HzO Substitution in Metal Ions

Rate of H?O Rate constant (kJd-'minute-') >letal substitution, ~ ______ - .- ion k( minute -') HDz HMeDz cu-NHDz

Cd '-2 2 . 4 X 10'O 109 > 10'0 zn + % 3 . 0 x 109 3 . 7 x 108 4 . 5 ' X 108 5 . 8 X loga c o L 2 3 . 0 x 107 5 . 2 X lo6 4 . 8 X lo8 Ti+% 1 . 2 x 106 8 . 0 x 104 3 . 1 ' X lo6 7 . 6 X lo6

Q Arbitrarily cnlciilated in same manner as others for purposes of comparison.

or, upon substitution from Equation 9 :

- d [Zn + 2 ] - klK,K ,' [Zn f 2 ] [HDz 1, - dt Ko,[H+]

and, when k l is sufficiently high to make Equation 11 rather than Equation 10, rate determining:

- d [Zn +2]

dt _ _ ~ - - k~ [Zn(HzO)2Dz + ] [Dz-]

which upon substitution gives:

- - -d[Zn+2] dt

where

The values for k l , calculated for the reaction of zinc with dithizone and its 2,2'-dimethyl derivative based upon this mechanism, are lOt3 . \R minute-' for dithizone and minute-' for 2,2'-dimethyl derivative. The value of k z K T is 2.6 X 1016W-1minute-1 for di-a-naphthyl derivative. The value for K T is not known at the present, therefore k2 cannot be calculated. The value for KT would be expected to be quite large. This explains the abnor- mally high value that would be ob- tained for kz if the mechanism given by Honaker and Freiser ( 4 ) were used, since there the KT does not enter into the rate equation.

I t is rather unusual to find that the first formation step in metal chelation is slower than the successive steps in view of the natural advantage of greater charge neutralization. Our results would seem to indicate that the change in coordination number of zinc ion is reFponsible. In this connection, an attempt was made to observe the rate of reaction of zinc with diphenylcarbazone and with 8-hydroxyquinoline. In both cases, where zinc remains six coordinate, the reactions occurred too rapidly to study the kinetics by the extraction method. Further work with other sul- fur-containing ligands is under way to test the generality of this conclusion.

further implication of the zinc study is that water removal from the metal ion appears to play a key role in the rate- determining step. In this case, the reaction kinetics of other metal dithizonates migh, be expected to follow a course related to the rate of water loss from the respective metal cations (2 , 3). .-is seen in Table 111, the reaction orders observed with Ni and Co were the same as those for zinc. The rate constants, k , based on equations analogous to Equation 8 are listed in Table IV.

.In attempt was made to study the

2264 ANALYTICAL CHEMISTRY

kinetics of extraction. of cadmium with dithizone and its di-a-naphthyl deriva- tive. In both case?,, the reaction oc- curred too rapidly to be studied by the solvent extraction method. Equilibra- tion was attained in less than two ininutw in every r~iii. The data indi- cated the rate cons1 ant to be greater than 1 O9J1 -‘minute-’ in the case of dithi- zone and greater than 10’051-1minute-1 in the ($as(’ of the cy-maphthylthiocarba- zone.

From Tahle IV, the importance of the lability of the metal ion-water bond in the overall reaction kinetics may be seen. The rate of metal dithizonate formation increases in an approximately linear fashion with the rate of dissocia- tion of water from the various metal ions. Thus, it would seem reasonable to postulatr that loss of coordinated water is involved in the rate-determining step in each of the reaction systems studied. This view is supported by the effect, of

steric hindering bubstituents on the reagent molecules on the rate3 observed. These wbstituents can be expected t o facilitate the removal of coordinated water molecules.

Math et a / . (6) have shoivn that the chelate formation constants of nickel(I1) dithizonates do not decrease, as do those of zinc and cadmium with the presence of steric hindering substituents, probably because the nickel dithizonate involves an S-N (rather than N-S) structure. The rates observed here are of the same order of magnitude and dis- play the same reaction order as those obtained in the polyamine nickel coni- plexes ( 5 ) .

ACKNOWLEDGMENT

Appreciation is expressed to Quintus Fernando for helpful discussion during the course of this investigation.

LITERATURE CITED

(1) Basolo, F., Pearson, R. G., “hIech- anism of Inorganic Reactions,” p. 113, Wiley, Sew York, 1958.

(2) Eigen, M., Purerlppl. (‘hem. 6 , 9;-115 (1963).

(3) Eigen, AI., Z. Elektrochem. 6 6 , 107 (1962).

(4 ) Honaker, C. B., Freiser, H., J . Phys. Chem. 6 6 , 127 (1962).

(5) Margerum, D. \V., Rorabacher, 11. B., Clarke, J. F. G., Jr., Inorg. (‘hem. 2, 667 (1963).

(6) Math, K. S., Fernando, Q., Freiser,

( 7 ) Orgel, L. E., “An Introduction to Transition-Metal Chemistry Ligand- Field Theory,” p. 83, IViley, SeLv York, 1960.

(8) Welcher, F. J., “Organic ilnalytical Reagents,” Vol. 111, p. 189, \‘an Sostrand, Sew York, 1953.

(9) Ib id . , p . 466.

RECEIVED for review Alay 18, 1964. Accepted August 31, 1964. Financial support of this project by the Atomic Knergy Commission is grat,efully ac- knowledged.

H., AXAL. CHEhl. 3 6 , 1762 (1964).

Study of Kinetics of Iridium(III)-Chlorate Reaction by Stead y- Sta te Control led - Po te n tia I Co u I o me t ry G. A. RECHNITZ and J. E. McCLURE Departrn ent o f Chemistry, University of Pennsylvania, Philadelphia, Pa. J 9 I 04

b The use of steady-state controlled- potential electrolysis in the investiga- tion of kinetics of redox reactions has been examined with regard to types of systems amenable to study. The kinetics of the iridium(ll1)-chlorate redox system have been investigated by this method. The rate law has the form

= -k(llrCt6-3) (c103-) d( lrC16-3) dt

at a constant hydrogen ion concentra- tion of 0.2M with k .= (3.2 f 0.1) X 10-4MM-1second-1 at 45” C. and p = 1 .O. The effects of various reactants and products on the rate of the reaction were investigated over the concentration ranges (H+) = 0.05

(CI-) = 0.0 to 0.2M; (lrC16-3) = 8 X lop4 to 0.01M; and initial (lrC1,j-2) = 0.0 to 2.(8 X 1 O-3M. The rate constant is in good agreement with that determined polarographically (3.0 X 10-4M-1second-1). The Ar- rhenius activation en’ergy, determined over a temperature range of 3 5 ” to 50” C., is 3 0 kcal. The overall re- action i s thought to proceed through the rate-determining1 step lrC16-3 -k C103- -+ lrC16-2 $- C103-*. Aqua- tion of lrCIG-3 during the rate rneasure- ment was considered, and it was shown that l r C l ~ - ~ and IrC15(H20)-2 react with C103- at the same rate.

to 0.8M; (C103-) =: 0.02 to 0.15M;

NFORhfATIOPi ABOUT KIXETICS O f a I reaction can be useful in choosing optimum experimental conditions for carrying out that reaction. Applica- tion of controlled-potential electrolysis in the investigation of kinetics nf homo- geneous solution reactions has been approached through the rrlationship between kinetic parameters and the all- parent extent of oxidation or reduction during electrolysis (2 , 3, 6 , ?), and through the relationship between kinetic parameters and the steady-state current obtained during an electrolysis op- posed by a cheniical regeneration of the electroactive material in the bulk of the solution (9, 11). T h k paper discusses some of the limitations of the latter ap- proach and illustrates it by application of the method to an investigation of the kinet.ics of the iridiuin(II1)-chlorate redox system.

THEORY

In the following discussion, B, C, and D are not electroactive at the potential a t which .4 undergoes an electrode re- action.

If a bulk electrolysis is carried out on a substjarwe, A , to yield a product, E, at an electrode whose potential is main- tained constant and under limiting current conditions, the current will de- cay exponentially with time according to i = where io is the current a t time

zero and A is a constant depending upon prevailing mass transfer conditions, cell geometry, solution voluni~ , electrode area, and stirring (1).

If, however, the electroactive species, A , is regenerated by a chemical reaction bet,ween the electrolysis product and some species, C, in the bulk of the solu- tion ( B + C -+ d + D ) , the current, after an initial decay, will attain some steady-state value. This steady-state current, which results when the rates of the electrochemical process and chenii- cal reaction in the bulk of the solution become equal, can be used to obtain information about the kinetic paiam- eters of the chemical reaction. The relationship between the steady-state current, is,, and the rate of the cheniical reaction is shown by the folloiving.

The rate of the electrode reaction is given by

ki

where d is the concentration of the electroactive specics, i the electrolysiq current, n the electron change. I.’ the faraday, and T’ the volume of qolution in the electrolysis cell. The rate of the chemical reaction. for sini1)le ki- netics, is given by

VOL. 36, NO. 12, NOVEMBER 1964 2265