The influence of HgII, CdII and Cl− on the formation and dissociation rates of FeNCS2+

Z inorg, nucl, Chem., 1976, Vol, 38, pp. 267-272. Pergamon Press. Printed in Great Britain

THE I N F L U E N C E OF H g n, C d II AND CI- ON THE FORMATION AND DISSOCIATION RATES OF FeNCS 2+ GARY HIGDON and D. L. LEUSSING Chemistry Department, Ohio State University, Columbus, OH 43210, U.S.A.

(Received 26 May 1975) Abstract--The influence of Cd~, Hgn and CI on the rates of formation and dissociation of FeNCS 2÷ has been examined in 0.5 M HCIO~, 0.05 M M(CIO,):. 0.05 M Cd~÷ exhibits no direct kinetic effect but inhibits the rate of formation of FeNCS :* by sequestering NCS- in an unreactive form. Fe~÷ reacts readily with Hg(SCN)2 and Hg(SCN)~-, water loss from the primary coordination sphere of Fe~÷ comprising the rate determining step. The reaction FeNCS2* + HgSCN+ ~ Fe~*+ Hg(SCN)2 exhibits a high second order rate constant which appears to arise from a very favorable pre-equilibrium reaction forming a dinuclear complex, FeNCSHgSCN ~÷,followed by a slow rate of Fe-NCS splitting. CI acceleration appears to procede via the formation of a ternary complex, Fe(NCS)CI+. Rate constants are given and their implication discussed. INTRODUCTION MANY examples are known of reactions involving the metal ion assisted removal of ligands from the primary coordination spheres of inert metal ions, e.g. (NH3)~ComCI 2÷ + Hg2+~ (NH3)sCo(H:O) 3÷ + HgCl+ [1]. These reactions are characterized as electrophilic substitution reactions, SE, and a review concerning them has b e e n published within the past few years [2]. Very little has appeared in the literature regarding the analogous reactions that involve complexes of labile metal ions. In a stopped flow study, Gibson and Carlyle [3] investigated the influence of Hg" species on the rate of dissociation of FeC1 :÷. Although the speed of the reaction prohibited determining rate constants for the reaction of Hg 2÷ with FeCI 2÷, they did find that HgCF and HgCI: are both reactive, the former species being especially so. For the reaction, Fe 3*+HgCI2~HgCF+ FeCl 2÷, they report a forward rate constant of 25 M ~sec-' while the reverse rate constant is 1.7 × l07 M-~sec '~ in 0.45 M HCIO4, I = 0.5 M, 25 °. In another study performed at an ionic strength of 3.0 M, Campion, Conocchioli and Sutin[4] found that the presence of the divalent cations Mg 2÷, Mn 2~, Co 2÷, Ni 2. and Zn 2÷ did not influence the rate of dissociation of FeCl 2÷ in any significant way. It was found, however, that 2+ • • Fe d~d increase the rate by operating through an electron transfer pathway. The present study was undertaken to gain more information regarding the nature of Se reactions of labile metal ions. The influefice of Cd u and Hg ~ on the rates of formation and dissociation of FeNCS ~÷ were studied. Cd 2* has a higher affinity for NCS- than those divalent metal ions studied by Campion et al. [4] have for CI-, therefore, the opportunity of characterizing electrophilic activity with a metal ion other than Hg H is somewhat better in Cd2+-NCS - systems. Also, because Fe 3÷ is normally coordinated to the nitrogen end of NCS- [5], the sulfur end is free to coordinate to a second metal ion. An ion with a high affinity for sulfur, such as Hg H, might then be found to be especially effective in promoting the 2+ dissociation of FeNCS . Finally, it has been reported that the presence of CI- [6] qualitatively promotes the dissociation of FeNCS 2+, but the kinetics of the reaction had not been determined. The present investigation was undertaken to examine the electrophilic activity of Cd" and Hg H

in the Fe3+-NCS - system. The effect of CI was, also, investigated. EXPERIMENTAL

Stock solutions of Fe(CIO,)~, Mg(C104)2, Cd(CIO4)2 and HCIO, were prepared from G. F. Smith Chemical Co., Reagents and were standardized using accepted methods. A known stock solution of NaC10, was prepared by carefully neutralizing NazCO3 with HCIO,, driving off CO2 and diluting to a known volume. Standard solutions of KNCS were prepared by weight after carefully drying a sample of Baker's reagent grade material. Hg(CIO4)2 solutions were prepared by dissolving a weighed amount of reagent grade HgO in a slight excess of HCIO4. Kinetic runs were performed at 25.0° using a Durrum Gibson stopped flow apparatus. The photomultiplier output was digitized and the results were stored in a Nova minicomputer for further processing. One of the syringes of the stopped flow apparatus contained the Fe 3+ solution and the other contained NCS-. Both solutions contained the inert electrolyte at the same concentration to obviate the spurious effects which tend to be encountered when solutions of widely different densities are mixed in a stopped flow apparatus. In most experiments the inert electrolyte comprised 0.50M HC104 and 0.05 M Mg(CIO,)2. In the experiments to determine the influence of CdH on the rates, the Mg(CIO4),~was replaced entirely by 0.05 M Cd(CIO4)2. In studying the influence of H + concentration, HCIO4 was replaced by an equivalent amount of NaCIO,, and in the studies where the effect of CI- was being examined, NaC1 was added while the Mg(CIO4)7 level was decreased in such a manner as to maintain the ionic strength at 0.65 M. The Hg" experiments were performed under conditions where the excess NCS- was present in an accurately known concentration over the amount required to quantitatively form Hg(SCN),. Lower concentration levels than these resulted in rates that were too fast to be measured. The exact details of each set of experiments are presented in Tables 1--4. Reactions were in general monitored by measuring the absorbance changes of FeNCS2-" at 500 nm. In the CI studies that were performed in the absence of NCS-, a wavelength of 380 nm was used. The relaxation time, ~-, of a reaction was obtained by a least squares fit of the digitized experimental absorbance-time curve to an equation of the form A, =A~+A~e -~/~ For the Fe2+--CI--NCS- system, where two relaxations exist, an equation of the form A, = A®+ A~ e-"~ + A2 e-"~2 was employed.

267

268

GARY HIGDON a n d D . L . LEUSSING

RESULTS The results obtained for the Fe3+-NCS - reaction under our conditions are presented in Table 1A. Previously, the rate law for the formation and dissociation of FeNCS 2÷ was found[6] to conform to the reaction,

Table 1. Relaxationtimes in the Fem-NCS- system A• Medium:

0,5OM HCI04, O.05M Mg(CIO4)a, ~.0 x IO-4M ENCS Felll tot

i/%bs

i/~eale

M

sec -I

sec -I

a

.0125,

3.1

3.03

•oo62

1.9

2. Ol

•o o 3 1

1.5

i. 51

kI

Fe 3++ NCS-<

'FeNCS 2+, KFosCS.

(1)

k 1 .oo~5

1.4

1.4o

•OOI2~

1.~

1.19

.00063

1.2

1.32_

R e a c t i ~ Path:

fomxd:

rJ+~+ Rcs" ~, r~es++, ~ e ~

= ~I/~-~ (1)

k~ = 167 / ¢ ' l s e c - t , k f l = 0.9@ see -1, l~eNO~q = i0 e'as M -I

Medlt~:

HCI04 + NaCIO 4 = 0.50M, O.O5M Mg(Cl04)e; 2.0 x IO'4M ~ C S

III Fet ot M

HCIO4 M

i/"re aleb see -I

i/%b s sec "I

.oo5o

.3o

2.o

2.13

.oo31

.3o

1•8

1.77

•OO1~4

.30

1.35

1.42

•0 0 0 6 2

.30

1.3

1.31

.0025

.1¢0

1.45

1.51

''

.30

1.6

1.65

''

.20

1.9

1.96

"

.lO

2.95

2.88

In analyzing the data of Table 1A this pathway was assumed• The calculations were straightforward and uncomplicated because it was valid to neglect the presence of Fe(NCS)2 + and higher complexes owing to their relatively low stabilities[8]. Using CORNEK II a search for k~* and KroNcs was made (k', is determined by the quotient k~/KFoNcs)and the "best" fit yielded values of 167 M-'sec-' and 102.53M-L Calculated rates obtained using these results are seen in Table 1A to lie in excellent agreement with the observed rates (presented as l/r). Further good agreement is obtained between the value of KFeNCS found here and those reported in the literature for approximately similar conditions: 10TM M-' (I = 0.4 M, 250) 6, 102.20 M -~ (I = 0"3 M, 200) 9 and 10 TM M -1 (I = 0'5 M, 25°) [10]. The rate of FeNCS 2+ formation has been reported 6 to exhibit an inverse dependence on the H + concentration. The results of a brief study of this effect under our reaction conditions are presented in Table lB. In agreement with the earlier conclusions[6], it was found possible to resolve kl into two components according to the equation, k] = k, + k~"l[H+].

Rate Equaticm:

k~ = kx + k~[H+]; fo%m~:

kl = 12~ M'isec "I,

k~ = 21.5, see -J-, k.l = 0.73 see-l, k.~ = .1214 see "I

Table 2. Influenceof Cd" on the FeI"-NCS- relaxation times Medi~:

0.5M HCI04, O.050M Cd(ClO4)e, 2.0 x IO'4M KNC8 Fe III tot

i/%bs

M

sec "I

a

•O050

1.3

i/tale see -I

L59

.0031

l.e

i.~4

.0010

1.1

1.08

.00068

I.i

1.o 5

.0oo50

1.o

1.o4

k~ Reaeti~s:

Fe +++ + NCS I-

Cd ++ + NCS"

~

fast ~

FeNCS ~

CdNCS +

(1) (2)

rc~cs Assumed:

KCdNC~+ = i0 I'3s M "l, (Ref• 12) k i = 1-67 ~J-see -J', k_~' = 0.9@ nee, ( ' ~ X e :r.A)

Further numerical analysis to extract rate and equilibrium constants from the relaxation-concentrationdependencewas made usinga computerprogram CORNEKII. This programis similarto a version described in the literature[7], but which has been modifiedto permitvalues of equilibriumconstants to be obtainedas well as rate constants. Complexationreactions of Cdxtand Hg n are considerably faster than those of Fem, so it was valid to assume that the complexation reactions of these divalent metal were sufficientlyfast to comprisean equilibriumstep coupledto the Fe" reactions.

Values of kl equal to 124 M-~sec -' and k~" equal to 21.5 sec-' were found. Below, Connick and Coppel[6] report corresponding values of 127 M-'sec-' and 20.2 sec -1 for media at an ionic strength of 0.4. Wendt and Strehlow [11] have found values of 150M-'sec-' and 45 sec-' for I = 0.5, 25° from measurements on solutions at considerably lower [H +] levels than employed for the experiments described in Table lB. Despite large differences in pH between the two sets of reaction conditions even these last results are in fair agreement with the present values. The effect of Cd2+ on the reaction rates was investigated by replacing the 0.05 M Mg(ClO4h in the reaction medium with 0.05 M Cd(C104h. Under these conditions approx• 50% of the NCS- that is not bound to Fem is bound as CdNCS +. A comparison of the rates observed in the presence of Cd H (Table 2) with those obtained in the presence of MgH (Table 1A) shows that Cd2+ causes the relaxation times to be slower• It was found possible to account for this inhibition simply by assuming that Cd~+ and CdNCS + are kinetically inactive and that the sole effect of Cd H is to reduce the amount of NCS- available for reaction with Fe 3+. Using a value of KCdNCS of 10~'33M-~ as interpolated for I = 0 . 6 5 M from data published by Gerding[12] and the values of k] and k'-i found here, the calculated rates shown in the last column of Table 2 were obtained. The good agreement with the observed results demonstrates the negligible kinetic activity of Cd H under the reaction conditions employed. The upper limit for the forward rate constant of the reaction k~ Fe w +CdNCS-,

)FeNCS + Cd2+ kL2

(2)

The influence of Hg", C d " and CI- on the formationand dissociation rates of FeNCS>

269

Table 3. Influence of Hg(SCNh on the Fe'"-NCS- relaxation times Medi~:

0.5M HCIO~, O.050M Mg(Cl04)2, 1.00 x IO-SM Hg(SCN)2

III Feto t

KNCS

i/~ob s

1/ ~ealea

i/'realeb

M

M x i0~

s ee-I

see-i

sec-i

4.21

.oo~o5

h.o

4.25

4.~2

.oo31o

''

3.8

5.79

3.81

.000992

''

2.9

2.8o

2.95

.o00682

,'

2.8

2. ? 5

2,8}

.000310

,,

2.8

2.68

2.68

1.0

8.6

8,58

8,62

• ooo99"2

' '

6. ~

6.5o

6.39

. ooo682

''

6.1

6.19

6. o7

,000310

''

5.7

5.67

5.70

.00310

~. Model A

kl Fe+++ + NCS- ~ k.i

Re&ctiens:

Fe +++

FeNCS ++

kg ~ FeNCS ++

+ ag(s~)2

(i)

+ Hg(SCN)+

',3_.)

FeNCS ++ + ~(SCN)2

(4_)

+++ Fe

+ ~(SCN)~k_~ fast

ng(SC~)2 + SCNfound:

:_

1ag(SC~)~ , ~-~SCN~

k~ = 16.9 M-Zsec -~, k.~ = 1.0 x i0~ M'Isec "I, k~ = i00 M-isec -I k.,~ = 4.5 x i0~ M'~sec -~, KHgSCNs = i o~'~9 M'~

b. Model B Resctlens:

FeS+

k~ ~ k-{

FeNCSe+

Fe a+ + Hg(SCN)~

Fe s + + Hg(SCN)~

(1)

~

FeNCS 2+ + Hg(SCN) ~+

(3)

~4 f-

FeNCS2 ÷ + H g ( S ~ I ) a

(4)

k.g

FeNCS ~÷ + H ~ C N +

fo%~Id:

> fast

k~ = 57,7 M'Zsec "l, k ~

Fe(NCS)Hg(SCN) s+

= i. 3 X 107 M'lsec -l

k~ = 206 M-lsec -I, k.~ = 8.4 x i0e M-Zsec "l KD = i07 .19 M-I

assumed:

H ~ + + NCS"

~

Hg(S~) I+ + NCS"

HgSCNI+' ~ g S C N = lOS'°8 M-I ~

Hg(SCN)2~ K ~ S C N 2 = 107.78 M -I

~ ( s ~ ) ~ + Ncs- a ~(sc~)~', ~ s c ~ ~(sc~)~- + ~cs"

(Ref. 13)

a ~(scm~-, ~ s c ~ ,

(Ref. 13)

= i°2"84 ~-ic

(Re~. 13)

= z°~'8~ M'I

(Ref. 13)

kl, k-1 as given in Table IA c.

Ezeept for Model A where this constant was fit.

appears to be about 3M-~sec -' and for the reverse reaction about 0.4 M-lsec ~. The first two ligands to be acquired by Hg xxgenerally occupy a linear coordination geometry and are tightly bound. Higher complexation induces a tetrahedral coordination geometry in which the ligands are most loosely held. These trends in stability can be seen in the stepwise stability constants determined by Ciavatta and Grimaldi[13] for the HgrLNCS - system given in Table 3. Although an ionic strength of 1.0 M was employed in determining these values, they do not differ much from results reported by Tanaka, Ebata and Murayama[14] for a lower ionic strength of 0.2 M. It appears then that little error has been introduced in using the values obtained at the higher ionic strength to analyse the present data. Inspection of Table 3 shows that the presence of Hg" markedly increases the reaction rates in the Fe>-NCS JINCVol. 38,No. 2--F

system, the more so, the lower the concentration of NCS-. The latter observation shows that Hg" complexes having a small number of bound NCS- ions are highly active kinetically. Preliminary computations revealed that only the reaction paths, k~ Fe >

+ Hg(SCN)2,

kL3

~FeNCS > + Hg(SCN), +

(3)

and, k; Fe > + Hg(SCN)3-,

kL4

)FeNCS 2+ + Hg(SCNh

(4)

need be included in the reaction scheme in addition to path (I). Hg > reactivity was not found, not because it is inactive, but because it was present at very low

270

GARY HIODON and D. L. LEUSSIN(3

Table 4. Influenceof CI- on the Fe'U-NCS- relaxation times A.

Med/~a:

HClO~ = O.5OM , NaCI+3 Mg(CIO~) e = O.15M

III Feto t M

NaCI M

i/mcalc a see "I

O.OLOO

.o5o

17.55

]-7.~9

0.0050

''

17.4

17.3~

0.0025

''

17.1

17,24

O. O100

.i00

19.6

19.67

0.0075

''

19.4

19.60

0.0025

''

19.6

19.48

0.OO10

''

19.6

19.45

Reaction Path:

found:

B.

i/mob s see -1

Medium:

Fe +++ + C1 ~-

**

(~)

FeC1 ++

k~ = 45 M-~sec -z, k.~ = 14.6 sec "z, 6Feel = i0 °'~e M "~

HCI04 = 0.50M, NaCl+3 Mg(ClO4)e = O.15M, 2.0 x IO'4M KNCS

Fe I I I tot M

~1

.OLOO

.o5o

3.2

~.43

3.68

.oo75

,,

2.8

2.07

3.14

.0050

''

2,k

1.71

2.60

.0025

''

2.0

1.36

2.O6

.0010

"

1.7

1.15

1.74

,O100

.i00

4.6

2.26

4.60

.o075

"

4.1

1.94

3.96

.0050

"

3.5

1.63

3.31

.0C25

''

3.0

1.52

2.68

2.6

1.13

2.30

1/~c~eb

i/%b,

M

slow, sec -1

.0010

slow, see -I

1/%~c c see -I

slow,

K D

FeNCS 2++ HgSCN '+,

k~ Fe z+ ÷ NCS"

Reactions:

~

FeNCS 2+

(i)

M Ye s+ + CI-

~

FeCI a* + NCS"

FeCI 2+

~

(5)

FeNCS e+ + Cl"

(6)

cs&culated aas1~ing k~ = k!e = 0

fotmd;k~ = 6 x 10a M'Xsec "I, k ~

a s s u m e d for b. and c. :

= iO M'isec "I

k~, k_~ as givem in Table IA k4, k_~ as ~

above.

Table 5. Rate and equilibrium constants for the formation of FeCI2÷in 0.5 M H+ Fes+ + CI" ~ FeC1++ k_6

M

~_~

M-isec "~

sec "I

1o8 tree1

I

M -~

ace.

M

45

15

0.48

0.65

this work

~5"

~

o.6o~

0.5o

al

51 a

13

0.60

0.~o

Ii

56 a

14

0.61

1.0

22

84

21

(b)

0 - 5 ( 0 . 4 5 ~ +)

3

68 a

13

0.72

i.o

23

o.k6

1.o

24

........

m.

Calc~ted

for 0.SM H + using the ex~easicla

:~

+~/~ +

and the reported values of ks and k~

b.

The value ~iv~a in Ref. 21 w~s used.

concentration levels under the experimental conditions described in Table 3. The concentration of Hg(SCN)/was also not important and reaction paths involving this complex were likewise not observed. In the initial attempts to fit a reaction model to the data, the stability constant of Ciavatta and Grimaldi[13] were used and a search for the "best" values of k~ and k~ was made. It was found possible to achieve only a fair fit to the observed rates, and a concentration dependent trend in the difference between the observed and calculated rates was evident. Another series of trials was then made in which KH,~sc~>~,the stepwise stability constant for the formation of Hg(SCN)f, was also considered to be a variable and was fit in the least squares sense. The results of these calculations, are listed in Table 3, Model A. A very good fit of the calculated rates to those observed is shown. Unfortunately, the value of K,g(scN~ which was obtained, 10~39M-I, is considerably higher than the literature results. These are in agreement with a value less than 1029M-L While Model A is therefore indicated to be a poor model, the calculations indicate that a higher degree of complexation occurs in the mixed system than is predicted by considering only the stability constants which apply to the binary Fe~+-NCS- and Hg2+-NCS systems. It was then assumed that a stable dinuclear complex is rapidly formed via the reaction, ~FeNCSHgSCN 3+.

(5)

Precedent for the formation of a complex of this type has been given: Bifano and Link [16] have found evidence for the formation of cis-Co(en)2C12Hg3+, and Orhanovic and Sutin[17] have demonstrated the existence of CrNCSHg '+. Using the Ciavatta and Grimaldi[13] stability constants for the Hg2+-NCS complexes, a search for the "best" values of kL k~ and Ko was made. The results are given in Table 3 under Model B where it is seen that the calculated rates agree with the observed somewhat better than is the case with Model A. The high value of Ko which was found, 107q9M-1, shows that Hg(SCN) + has almost as great a coordinating affinity for the terminal sulfur atom of FeNCS 2+ as it has for uncomplexed NCS-. Indeed, the difference between Ko and K,,scN2 is almost entirely accounted for by the coulombic differences between FeNCS ÷ and NCStowards HgSCN ÷. The high value of Ko also suggests that the complex is singly bridged with the HgIx prossessing a linear coordination geometry. While a double NCSbridge would tend to enhance the stability of the polynuclear complex the pronounced bond weaking as the Hg(II) ion goes from a linear to a tetrahedral coordination geometry would likely result in a considerable net destabilization. Consistent with this line of thought, no evidence was found for the formation of appreciable concentrations of higher complexes, such as Fe(NCS)2Hg(SCN)2+ or Fe(NCS)Hg(SCN)/÷, even though reduced charge repulsion would be expected to stabilize these species. Fe 3+ exhibits almost identical rate constants in its reactions with NCS- and Hg(SCN)3-, the values being 1.7 and 2.1 × 105 M-'sec-', respectively. These rate constants lie in the range determined by rate limiting water loss from the primary coordination sphere of an outer sphere complex [18, 19],

The influenceof Hg", Cd" and CI- on the formation and dissociationrates of FeNCS2+ K

Fe(H20)63+ + L ° , %Fe(H20)~... L" fast

k 3+

Fe(H.,O)6 ...L

n-

ex

~Fe(H20)sL

3 n

+H20.

The observed second order rate constant for formation is equal to the product Ko~k,x. Because k~ is independent on L" and Ko~ is roughly the same for the two monovalent anions, nearly identical values of k~ and k~ result. The lower value found for k; is the natural consequence of a decrease in Ko~ arising from a reaction involving a neutral species. Indeed, the observed decrease in the second order rate constant is just about the same as that quantitatively predicted from the changes in Ko~ brought about by coulombic factors within the errors of this type of calculation [19]. From this the rates of reaction of both Hg(SCNh and Hg(SCN)3 with Fe ~* appear to be limited by the rate of FeH~-OH2 bond breaking despite a vast difference between the Hg"-SCN bond strengths in the two Hg lI complexes. A further coulombic induced decrease in rate is expected for the reaction of CdNCS + with Fe 3.. However, the upper limit determined for the rate constant for this reaction is about 30-50% smaller than that predicted on this basis alone. A possible explanation for the additional decrease in reactivity may be that Cd 2+ ion in N-bonded to NCS , leaving the sulfur end available for attachment to Fe 3~. The equilibrium position of the reaction. CdNCS ++ Fe3+~CdNCSFe 4+, is expected to lie far to the left because of the low affinity of Fe 3~ for sulfur, in addition to electrostatic considerations. Intermediate CdNCSFe "+, when it is formed would tend to break down into reactants rather than products. In contrast to the situation with CdNCS*, the high value of the second order rate constant, kL3 for the reaction path, FeNCS 2++ HgSCN + ~ Fe 3++ Hg(SCN)z, (k'-3 = 1.3 x 10VM-'sec ~) seems to arise almost entirely from a highly favorable pre-equilibrium constant. For the reaction sequence. KD

FeNCS 2++ HgSCN', fast ~FeNCSHgSCN 3+ ~'~', Fe 3++ Hg(SCN)2, a value of kd~s equal to 0.87 sec -~ is obtained from the quotient k'-~/KD. This result is identical, within the experimental uncertainties, to the first order rate constant k'~ found for the uncatalyzed dissociation of FeNCS 2+. Thus, the coordination of the Hg(II) atom to the free end of Fem complexed NCS- does not appear to appreciably influence the cleavage rate of the iron(III)--nitrogen atom bond. Gibson and Carlyle[3] have found the forward and reverse rate constants for the reaction, Fe 3++ HgCb~FeCI2++HgC1 +, to be 25M -~ sec -~ and 1.7x 107 M ~sec ', respectively in 0.45 M H +. These values are very close to those found here for the analogous NCS paths. It is expected that the forward rates will be similar because in both cases the slow step is water loss from the primary coordination sphere of Fe 3+. The nearly identical reverse rates are a consequence of the nearly identical equilibrium constants for the two reactions. While C1- is more weakly bound to Fe "~ than is NCS-, it is also more weakly bound to Hg ~J by roughly the same factor. Weaker bonding explains why Gibson and Carlyle[3]

271

found no evidence for the formation of polynuclear species in their C1- systems. It would be very difficult to detect a complex such as FeCIHgC1~+ if its stability constant were one or two order of magnitude lower than the value of KD found here for NCS . On the other hand, because the over-all rate constants for the reaction FeX 2++ HgX + ~ Fe3++ HgX2 are observed to be nearly the same in the two cases, a lower value of KD implies that an intermediate such as FeC1 HgCI3+ dissociates into Fe 3÷ and HgCI2 at a considerably faster rate than does the corresponding NCS complex. Thus, the coordination Hg ~ directly to the atom that is coordinated to Fe ~"greatly assists its removal, whereas Hg Hcoordination to a remote site has only a small effect. Gibson and Carlyle[3] report rate constants for the reaction, Fe 3++ HgCI3 ~ FeCf + + HgCI2, that are considerably higher than the values of k~ and k'-4 found here for the analogous NCS- reaction. These high values arise from secondary trends shown in their rate data[3] which we were unable to reproduce despite repeated efforts by two different workers in our laboratory [20]. Our results indicate no kinetic activity of HgCI3 under the reaction conditions employed by Gibson and Carlyle, however our rate constants found for the principle path, that involving HgC1+, are in very good agreement with theirs. We have no explanation for the origin of the slightly different trends observed in these two laboratories, except to suggest that they may be instrumental. No difficulty was encountered by us in reproducing results reported by Campion, Conocchioli and Sutin[4]. A brief study of the rate of formation of FeCf ~ was also made under the reaction conditions employed in our NCS- studies, and the results are shown in Table 4. At the C1 concentration levels employed, 0.50 and 0.10 M, the Fe3+-CI relaxation times are about an order of magnitude faster than those observed for the Fe3+-NCS system. The data of Table 4A yielded values of 45 M-'sec ~for k;, the forward rate constant for the formation of FeCf ~and 10°48M-~ for the stability constant. Generally, good agreement exists between these results and literature values, which are given in Table 5 for comparable conditions. The variations between the values reported by the various authors have little effect on the conclusions drawn here concerning the ternary Fe3~-NCS-CI system. The data for this last system are given in Table 4. The following reaction paths are assumed to be present in a reaction mixture containing Fe 3", CI- and NCS under the concentration conditions defined in Table 4B, k~ )FeNCS 2+

Fe3+ + NCS-(

(1)

k- I

k~ Fe3÷C] ~k i s FeCf +

(5)

k~ FeCf++NCS ~ ' F e N C S 2 - + C I .

(6)

k- 6

Because FeClf and Fe(NCSh + are not sufficiently stable to be formed in appreciable concentrations under these reaction conditions, it was assumed that the concentrations of the mixed complex, Fe(NCS)CI +, could, also, be neglected. It is then predicted that this reaction system will show two relaxation times which are the eigenvalues of the matrix, a~l-A at2 l a21 az~_- A /

272

GARY HIGDON a n d D . L . LEUSSING

where all= kl([Fe 3+]+ [NCS-]) + k' ,[FeC12+]+ k'-,[C1-] an = (k[ - k~)[NCS-] - k'-6[FeNCS2+] a2~ = (k~ - k~)[Cl-] - k' 6[FeC13+] a22 = k[([Fe3+1+ [CI-]) + k'5 + k~[NCS-] + k '-6[FeNCS 2+] and the concentrations are the equilibrium values. Because Fe3+-C1- reactions are about an order of magnitude faster, the slower relaxations essentially originate solely from the reactions of the FeNCS 2+ system. Preliminary calculations of the theoretical rates were made assuming that k~ and k'-6 in the above reaction scheme are zero. A comparison of the results of these calculations, given in the fourth column of Table 4B, with the observed values given in the third column, shows that the observed rates are faster, the more so the greater the C1- concentration. These results indicate that a CIpromoted path, such as that designated as (6) in the above reaction scheme, operates at a significant rate. Assuming values of k l and k~ and the appropriate stability constants as found in this study a search for the "best" value of k~ yielded 6x l0 s M-~sec -1. The theoretical rates obtained with this model are presented in the last column of Table 4, and can be seen to agree well with the observed rates. The rate constant for the formation of Fe(H20)5(NCS) 2+ from Fe(H20)sCI2+ has been found here to be about 2-3 times faster than its rate of formation from Fe(H:O)6 ~+. This enhancement results from the well known effect of a coordinated ligand to markedly increase the water exchange rate constant, kex, of a metal ion[25]. In contrast to the increase observed in kL the rate constant for the back reaction, k'-6, indicates that the rate of formation of FeC12+via the attack of CI- on FeNCS 2+is slower than the attack of CI- on Fe 3÷. This observation does not necessarily imply that coordinated NCS- inhibits the exchange rate of H20 molecules coordinated to Fem. The reaction path via the ternary species is complicated and likely involves a sequence such as, XFe(H20)5 2+ + Y - (

fast

>XFe(H20)sY +

KOS

XFe(H20)sY ÷ ~

k_

+

XFe(HEO)4Y + H20 k X

H:O + XFe(H20)4Y+v-*X(HEO)Fe(H20)4Y + k-H20

1/K6s

X(H20)Fe(H20)4Y+~

'X-+(HzO)Fe(H20)4y2L

The intermediate mixed complex, XFe(H20)~Y ÷, can dissociate to give either X- or Y-, and the ratio of the moles of X- produced to the amount of Y- is the ratio of the rate constants k-x/k-v. Since the rate of dissociation of ligand is a function of the metal-ligand bond strength, It

is more probable that a less tightly bound ligand, such as CI-, will dissociate from the mixed complex than a more tightly bound one, such as NCS-. Thus, the over-all rate of replacement of C1- by NCS- via the mixed complex will procede easily, but the reverse reaction will be much more difficult. FeNCS 2+ formation via the hydroxy complex procedes very rapidly even though OH- is bound more firmly to Fem than NCS-. Using a value of 2.80 as the pK, of Fe ~+[26] the second order rate constant for the reaction of NCS- with FeOH 2+ is calculated from k, H to be 1.4x 104 M-~sec -~. It is unlikely that this reaction path involves the direct replacement of O W by NCS-. Rather, a rapid protonation reaction in an intermediate mixed complex, Fe(NCS)(OH)(H20), + + H +~Fe(NCS)(H20)5, will give the final product without the necessity of breaking a metal ion-oxygen atom bond. This provides a very fast path for the conversion of a coordinated OHion to a coordinated water molecule. Acknowledgement--The National Science Foundation has kindly

supported this research. REFERENCES 1. J. N. Bronsted and C. E. Teeter, J. Phys. Chem. 28, 583 (1924). 2. M. M. Jones and H. R. Clark, J. Inorg. Nucl. Chem. 33, 413 (1971). 3. D. Gibson and D. W. Carlyle, lnorg. Chem. 10, 2078 (1971). 4. R. J. Campion, T. J, Conocchioliand N. Sutin, J. Am. Chem. Soc. 86, 4591 (1964). 5. D. P. Fay and N. Sutin, Inorg. Chem. 9, 1291 (1970). 6. J. F. Below,Jr., R. E. Connick and C. P. Coppel, J. Am. Chem. Soc. 80, 2961 (1958). 7. V. S. Sharma and D. L. Leussing, Talanta 18, 1137 (1971). 8. Stability Constants (Compiled by L. G. Sillen and A. E. Martell), The Chemical Society, London, Special publication No. 17 (1964), Special Publication No. 25 (1971). 9. V. N. Kumok and V. V. Serebrennikov,Zhur. Neorg. Khim. 9, 2148 (1964); quoted in Ref. [8]. 10. M. W. Lister and D. E. Rivington, Can. J. Chem. 33, 1572 (1955). 11. H. Wendt and H. Strehlow, Z. Elektrochem. 66, 228 (1962). 12. P. Gerding, Acta Chem. Scand. 20, 2771 (1966). 13. L. Ciavatta and M. Grimaldi,Inorg. Chim. Acta 4, 312 (1970). 14. N. Tanaka, K. Ebata and T. Murayama, Bull. Chim. Soc. Japan 35, 124 (1%2). 15. C. J. Nyman and G. S. Alberts, Anal. Chem. 32,207 (1960). 16. C. Bifano and R. G. Linck, Inorg. Chem. 7, 908 (1968). 17. M. Orhanovicand N. Sutin, J. Am. Chem. Soc. 90, 538 (1%8). 18. R. G. Wilkins and M. Eigen, Mechanisms of Inorganic Reactions (Edited by R. F. Gould), Advances in Chemistry, Vol. 49, American Chemical Society (1%5). 19. K. Kustin and J. Swinehart, Inorganic Reaction Mechanisms (Edited by J. O. Edwards). Wiley, New York (1970). 20. M. Schwartz and G. Higdon, these laboratories. 21. 1~. E. Connick and C. P. Coppel, J. Am. Chem. Soc. 81, 6389 (1959). 22. T. Yasunaga and S, Harada, Bull. Chem. Soc. Japan 42, 2165 (1969). 23. J. K. Rowley and N. Sutin, J. Phys. Chem. 74, 2043 (1970). 24. M. J. M. Woods, P. K. Gallagherand E. L. King, Inorg. Chem. l, 55 (1962). 25. J. P. Hunt, Coord. Chem. Rev. 7, 1 (1971). 26. R. M. Milburn, J. Am. Chem. Soc. 79, 537 (1957).