J. CHEM. SOC. FARADAY TRANS., 1992, 88(4), 591-594 59 1

Kinetic Study of the Oxidation of Iodide by Hexachloroiridate(1v) in Concentrated Electrolyte Solutions

Amalia Rodriguez, Maria Bejarano, Elena Fernandez-Boy, Maria del Mar Graciani, Francisco Sanchez and Maria Luisa Moya* Departamento de Quimica Fisica , Universidad de Sevilla , C/ Profesor Garcia Gonzalez s/n, 4 10 12 Sevilla , Spain

Rate constants are reported for the oxidation of iodide by hexachloroiridate(iv) in various concentrated electro- lyte solutions containing up to 6 rnol dm-3 salt at 298.2 K. The conclusion is that the salt effects are mainly due to ion-solvent interactions. This is supported by considering the activity coefficients of the reaction participants, estimated by using the Stokes-Robinson hydration treatment. The study of kinetics at 333.2 K shows that the distinct behaviour of potassium nitrate seems related not only to the different structural character of this salt with respect to the rest of the electrolytes studied, but also to specific interactions which are operative favouring the reaction.

We have recently been concerned with medium effects on reactions between ions.'-'' Over the years different approaches have been used to examine medium effects on reactivity, one of the more widely used approaches being the dissection of medium effects into initial and transition states." Traditionally, this approach has been used in pure solvents or in binary solvent mixtures; however, it seems interesting to examine the influence of initial-state and transition-state solvation changes in determining overall reactivity in concentrated electrolyte solutions, given that the medium properties can also be changed by adding a salt. In other words, it is of interest to test if a unitary treatment of salt and co-solvent effects would be possible. To carry this out we have chosen the oxidation of iodide by hexa- chloroiridate(rv), a well characterised outer-sphere electron- transfer reaction,'2T'3 which offers the additional advantage of being pH independent.', This reaction has been studied by Blandamer et a/.', in various binary aqueous solvent mix- tures. The present work can give a deeper insight into medium effects on this reaction, including salt as a special co-solvent case. At first our idea was to carry out the kinetic study at 298.2 K. However, after the specific behaviour observed for potassium nitrate, we decided to study the process of 333.2 K. The reason for this will be given in the Discussion.

Experimental Materials

The products employed in this study were commercial P.A. grade and were used without further purification. All solu- tions were made up with deionized water (conductivity < S cm-').

Kinetics

The rate measurements were performed using a Hi-Tech stopped-flow instrument. The thermostatic system main- tained the temperature in a range of fO.l "C. The kinetics were followed at 478 nm, the wavelength of the maximum in the complex IrCli-.

mol dmP3 and [I-] = 2.5 x lo-, mol dm-,. All runs employed an excess of sodium iodide so that pseudo-first-order condi- tions were followed. This pseudo-first-order condition was also convenient in order to ensure that the process studied

Reagent concentrations were [IrCli-] = 2.5 x

was the oxidation of the iodide and not that of tri-iodide, 13, which is formed from iodine, a product of the reaction, and the iodide. An iodide excess makes sure that the oxidant is consumed mainly during the iodide oxidation process. This avoids the complication caused by two parallel paths of reac- tion for the disappearance of the oxidant since the rates of oxidation of iodide and tri-iodide are different.15 The pseudo- first-order rate constants, kobs, were obtained by plotting ln(A, - A, ) us. time, where A, is the measured absorbance at time t . The first-order kinetic plots were linear over at least three half-lives. The rate constants were found to be repro- ducible within a precision of ca. 5%. A second-order rate law, first order in respect to each reactant, was found for the dif- ferent electrolyte solutions in agreement with previous results.'2*'

To verify the reliability of the experimental results, we have reproduced data from the literature. So, for the reagent con- centrations CIrCli-1 = 2.5 x loe5 mol drn-,, [I-] = 4 x lo-, mol dm-, and p = mol dm-, (LiClO,), at 298 K, we obtained k = 328 dm3 mol-' s-l in close agree- ment with the value k = 320 dm3 mol-' s-l given by Bland- amer et a/. '

Concentrations of added salts covered the 1.0-6.0 mol dm-, range for LiNO, and NaNO,, 1.0-2.5 mol dmP3 for the less soluble KNO,, and 0.5-3.0 mol dm-, for

, CaWO,), and Sr(N03)2.

Density and Viscosity

Density and viscosity measurements were carried out by standard methods. The maximum error in the density mea- surements was f 1 x lop4 g ern-,. The viscometer was cali- brated with high-purity conductivity water. The accuracy of flow times was f0.25 s. The reliability of the viscosity mea- surements at 333.2 K was tested by reproducing data from the literature.16

Results Table 1 summarizes the second-order rate constants for the different salt solutions at 298.2 K. The rate constants at 333.2 K are listed in Table 2.

The electrolyte concentrations, C, and the values (q/qo - l)/C1'2 at 333.2 K, where q/qo is the relative viscosity of

the solution, are listed in Table 3 for each salt. The Jones- Dole B coefficients are also listed in this table. B values were

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592 J. CHEM. SOC. FARADAY TRANS., 1992, VOL. 88

Table 1 Rate constants for the oxidation of iodide by hexachloroiridate(1v) in various concentrated electrolyte solutions at 298.2 K

k/103 dm3 m o l - ' ~ - ~

[salt]/mol dm - ' NaNO, LiNO, KNO, Sr(N03)2

0.5 1.0 1.5 2.0 2.5 3 .O 4.0 5.0 6.0

-

3.05

5.40

8.68

-

-

11.8 15.5 20.3

- -

2.11 7.63 11.6

3.54 14.9 22.2

5.39 5.37 -

11.8 15.8 -

-

- -

-

~ ~ ~ ~ ~ _ _ _ _ _ _ _ _ _ _

2.59 2.18 4.38 4.15 6.32 6.18 9.94 8.47

16.7 11.5 23.5 14.1 - -

2.72 4.52 5.93 8.15

10.5 13.1 -

- -

calculated from viscosity data by using the equation:'

1

(q/q() - 1) $y = A + BC",

where A and B are characteristic parameters for the solvent and the electrolyte.

Discussion The detailed oxidation mechanism proposed by Blandamer et ~ 1 . ' ~ can be written in the form:

+ 1rc13 - + I*, 21 - I - (I)

The redox step is thermodynamically unfavourable. Using the literature values in water E*(II-C~;-'~-) x 0.96 V and E*(I-/I.) x 1.4 V,14 the ratio ka/kb may be computed as ca. 4 x mol-' dm3. After the initial redox reaction, the recombination (kb) within the solvent cages competes with the diffusion (k , ) of 1'. The separation of the individual rate constants is not possible, but since these reaction rates are close to the diffusion control limit, one can assume that k, x k , . Taking into account that the observed rate constant is of the form k , = kakc/(k, + kc), then k , % O S k , . Therefore, the observed rate constant is a measure of the rate of the initial redox process. This is in agreement with the fact that the rates of oxidation of iodide by different iron(u) complexes and by IrClz- can be rationalized by Marcus' e q ~ a t i o n . ' ~

The results given in Table 1 show the existence of a posi- tive salt effect for all the electrolytes studied. This effect is specific in the sense that it depends on the nature of the salt

present in the reaction medium. These results can be inter- preted by taking as a starting basis the expression:

in which yA and yB represent the activity coefficients of the reactants, y # being that for the activated complex. If the activity coefficient for a given ionic solute is related to salt concentration, C , using a Guggenheim-Guntleberg type of equation,I8 the final dependence of log k on electrolyte con- centration takes the form shown in eqn. (3).

AC'I2 (3) log k = log k" + - 1 + C'P + sc

As written here, parameters A and S are adjustable, obtained by fitting the data to eqn. (3). Term AC", represents the Coulombic effects of added salt upon the participants in the reaction and the term SC accounts for the non-Coulombic effects of the salt. For all the electrolytes studied log(k/dm3 mol- s - ') varies linearly with salt concentration, indicating that the third term of the right-hand side of eqn. (3) is much greater than the second term, in agreement with the idea advanced earlier that changes in Coulombic interactions are not important in concentrated solutions.' 8-21 Therefore, the cause of the observed salting process upon the reaction par- ticipants is considered to be mainly due to non-Coulombic interactions. For a given salt the term S seems to be constant.

In our opinion, slope S would be a measure of the effects of hydration of added salts on the reaction. Considering that all the electrolytes used have common anions, it seems reason- able to expect that the difference in the slope values is due to the different nature of the cations. Indeed, Fig. l(a) is a plot of S us. B, the Jones-Dole coefficient for salts, at 298.2 K. As can be seen, the kinetic salt effect of a given electrolyte, as measured by S, is proportional to its B coefficient, which is a

Table 2 Rate constants for the oxidation of iodide by hexachloroiridate(1v) in various concentrated electrolyte solutions at 333.2 K

k/103 dm3 mol-I s-'

[salt]/mol dm - NaNO, LiNO, KNO, Mg(No 312 Ca(N03)2 Sr(No 312

0.5 1 .o 1.5 2.0 2.5 3.0 4.0 5.0 6.0

9.77 - - -

10.6 7.94 18.6 14.6 - - 24.5 20.9 13.8 10.0 30.9 28.8 - - 36.3 36.3

49.0 20.8 15.3 25.1 17.4 39.8 24.0 49.0 30.2

- - - - -

- -

10.7 14.8 18.2 22.9 28.8 37.1 -

11.5 15.5 19.5 24.5 29.5 38.0 -

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J. CHEM. SOC. FARADAY TRANS., 1992, VOL. 88 593

Table 3 Molal electrolyte concentrations, C, (q/qo - l)/C'/* values and Jones-Dole coefficients, B, for various salts at 333.2 K

NaNO, 0.50 - 0.0545 0.10 - 0.0206 0.15 - 0.049 1 0.25 0.0148 0.50 0.0434 0.60 0.0509 0.75 0.0642 1 .00 0.08 15

B = 0.145 dm3 mol-'

Mg(N03)2

0.50 0.141 0.10 0.156 0.15 0.180 0.25 0.228 0.50 0.308 0.60 0.335 0.75 0.3 72 1 .OO 0.426

B = 0.370 dm3 mol-'

LiNO,

0.50 0.001 66 0.10 0.0329 0.15 0.0429 0.25 0.0556 0.50 0.0789 0.60 0.0865 0.75 0.0990 1 .00 0.1 18

B = 0.107 dm' mol-'

Ca(NO,), 0.50 0.120 0.10 0.148 0.15 0.156 0.25 0.176 0.50 0.22 1 0.60 0.237 0.75 0.264 1 .OO 0.301

B = 0.238 dm3 mol-'

0.50 0.10 0.15 0.25 0.50 0.60 0.75 1 .OO

0.50 0.10 0.15 0.25 0.50 0.60 0.75 1 .oo

KNO,

-0.0732 - 0.0376 - 0.0292 -0.0152

0.00333 0.00636 0.0121 0.0191

B = 0.0970 dm3 mol-'

Sr(N03)2

0.124 0.136 0.164 0.175 0.213 0.225 0.244 0.272

B = 0.194 dm3 mol-'

measure of the salt-solvent interactions.22 Fig. l(a) shows that the point corresponding to the potassium salt deviates from the general correlation. Taking into account that this ion belongs to a different type to the other cations in the Nightingale classification of ionic hydration,23 the conclusion could be made that not only the degree to which the cations are hydrated is important in relation to the salt effects

0.38

7 0.30 - E

m g 0.22 cr,

0.14

0.38

0.30 I - E 6 0.22 m

G-

0.14

0.00 0.1 0 0.20 0.30 B/dm3 mol-'

LiNO3

0.13 0.23 0.33 0.43 B/dm3 mol-'

Plot of the slopes S us. the Jones-Dole coefficients B for the Fig. 1 different electrolytes studied. T = (a) 298.2 and (b) 333.2 K

observed, but also the form in which they are hydrated. We will consider the potassium behaviour further.

The increase of the rate constants with increasing salt con- centration can be due to either a greater destabilization of the initial state with respect to the transition state or to the lower stabilisation of the initial state with respect to the transition state. In order to investigate this, activity coefficients of IrC1;- and I - ions were estimated by using the Stokes- Robinson hydration treatment.24-26 These authors obtained an equation for activity coefficients by superimposing on the Debye-Huckel treatment an allowance for the ion-solvent interaction in terms of a hydration model (water bound to ions is no longer a part of the bulk solvent), which can be applied in concentrated electrolyte solutions. The activity coefficient of the transition state was estimated from eqn. (3). The vaIues of the relative activity coefficients for the different electrolyte solutions are listed in Table 4. These results illus- trate that the increase of the rate constants with increasing salt concentration is due to a greater destabilization of the initial state than of the transition state.

Solvent structure is distorted by the presence of supporting electrolytes. This distortion can be thought of as due to com- peting forces from solvent structure in the bulk and from the ionic field and/or the oriented molecules associated with the ion. At 298.2 K potassium nitrate behaves as a structure breaker, whereas the other salts used in this study are struc- ture makers. However, with increasing temperature there will be less competition between the ionic field and the reduced solvent structure, therefore potassium nitrate could become a structure maker at a high enough temperature. Thus, the study of the reaction at high temperature seems interesting in order to investigate whether the different structural character of the potassium salt at 298.2 K is responsible for the devi- ation shown by this electrolyte in Fig. l ( ~ ) . Table 2 sum- marises the rate constant values for the oxidation of iodide by hexachloroiridate(1v) at 333.2 K. At this temperature log(k/dm3 mol- I s- I ) varies linearly with salt concentration for the electrolytes studied. This linear behaviour indicates that also at this temperature the salting effect on the reaction participants is due mainly to non-Coulombic interactions [see eqn. (3)]. Jones-Dole B coefficients summarised in Table 3 show that at 333.2 K all the electrolytes used in this study behave as structure makers. However, the plot of the slopes S us. B at 333.2 K [Fig. l(b)] shows that potassium nitrate also

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Table 4 Relative activity coefficients for several salt concentrations

[MNO,]/mol dm-3

M activity 1.0 1.5 2.0 2.5 3.0 4.0 5.0 6.0

Y1,Cl;- 1.0 - 1.1 - 1.3 1.6 2.1 -

Na+ { yI- 1.0 ~ 1.1 ~ 1.2 1.4 1.7 ~

Y # 1.0 - 0.7 - 0.6 0.6 0.7 -

Y1,Cl:- 1.0 - 1.2 - 1.6 2.2 3.6 6.6 1.1 - 1.3 1.7 2.3 3.4

1.0 - 0.8 - 0.8 1.0 1.5 3.0

Ylrc& 1.0 1.0 1.0 1.1 - - - ~

1.0 1.0 1.0 1.1 - - - -

1.0 0.7 0.6 0.4 - - - -

[M(NO,),]/mol dm -

M activity 0.5 1.0 1.5 2.0 2.5 3.0

71Kd - 1.0 1.1 1.5 2.0 3.0 5.2 MgZ+ { 7 , 1.0 1.1 1.4 1.6 2.1 3.0

'Jf 1.0 0.7 0.8 0.8 1.0 1.7

YtKli ~ 1.0 1.1 1.3 1.6 -~ -

caz+ { yl- 1.0 1.1 1.2 1.5 - Y# 1.0 0.6 0.6 0.6 - -

1.0 1.1 1.2 1.5 2.0 2.7 1.0 1.1 1.2 1.4 1.8 2.3 1.0 0.7 0.7 0.8 0.9 1.3

deviates from the general behaviour at this temperature, although the deviation is less pronounced than at 298.2 K. This result seems to indicate that the deviation of potassium nitrate at 298.2 K is not due only to its different structural character.

In conclusion, we can say that for the reaction studied the solvation of the supporting electrolyte is the dominant cause of the salt effects observed in concentrated solutions of elec- trolytes. Only in the case of potassium nitrate are specific interactions, whose nature is unknown to us at the present, apparently operative. We are investigating salt effects in con- centrated salt solutions in other electron-transfer processes (cation-cation and anion-cation) to elucidate the behaviour of potassium nitrate.

Junta de Andalucia and Patronato de la Fundacion Camara de la Universidad de Sevilla.

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The authors acknowledge the financial support of the GICYT (PS 8901 18), Consejeria de Educacion y Ciencia de la Paper 1/02457D; Received 24th May, 1991

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