Kinetics and salt effects on the reaction between octacyanotungstate(IV) and periodate ions

Notes of the blue color, upon the addition of an excess of potassium azide to Co(I1) solutions in the alkali nitrate melt, showed that tetrahedral, azide rich Co(II) complexes form in the melt as well. However, these complexes decomposed rapidly at the temperatures of the melt. The above mentioned band shift from 548 to 513 nm, observed when potassium azide was added to Co(lI) solutions in molten alkali nitrates, showed that the octahedral symmetry was preserved in the species formed when CoOl) is in excess which indicates that the monoazido complex formed. The formation of Co(II) chloro complexes in alkali nitrate melt solution has been studied previously[2, 6-8]. Due to the relative positions of the chloride and nitrate ions in the spectrochemical series, band shifts to longer wavelengths took place regardless whether the chloride coordination resulted in the formation of a complex with tetrahedral symmetry or a complex in which the octahedral symmetry was preserved. Therefore, it was not possible to determine unequivocally whether the octahedral structure was preserved or not when the Co(ll) monochloro complex formed, Due to the higher ligand field strength of the azide ion. this diflSculty does not arise in the case of azide coordination and the observed shift of the absorption peak to shorter wavelength indicated clearly that the octahedral symmetry ~us preserved when the Co(II) monoazido complex formed. Some experiments were also conducted with molten KCNS as solvent. Among the first row transition metal ions Co(IlL Ni(II), Cr(Ill), and VO2(IV) dissolve in molten KCNS. Ni(ll) and Cr(lll) are octahedrally coordinated, Co(II) is tetrahedrally coordinated and VO2(IV) has a distorted octahedral structure [291. Upon potassium azide addition, azide coordination to the transition metal ions was indicated by slight band shifts and changes in the absorbances. In the case of VO2(IV), decomposition occurred when potassium azide was added. The azido complexes of the other ions mentioned above, formed when the transilion metal ion concentrations were higher than the azide ion concentrations, were thermally stable at the temperatures of the spectrophotometric work. In agreement with the positions of the azide group and the - NCS group in the spectrochemical series, band shifts to longer wavelengths took place upon azide coordination. In view of the relatively high [igand field strength of the NCS group, the formation of monoazido complexes can be expected to be highly favored at the azide ion concentrations considered in this study. The investigations discussed in this paper showed that transition metal azido complexes form in molten salt solutions and that monoazido complexes possess an unexpectedly high stability against thermal decomposition. The higher thermal stability of monoazido complexes, as compared to the corresponding corn+ plexes comprising more than one azide ligand, can be attributed to the fact that the preferred mechanism of the thermal decom+ postion of inorganic azides is the formation of azide radicals which interact in a bimolecular reaction to form three molecules of nitrogen[30, 31]. This mechanism cannot take place in the case of monoazido complexes.

l'S 4rmy MohilHv Equipmenl Research and Development Command ('ountermine Laboratory Fort Belroir v4 22060. U,S.A.

H.C. EGGHART

1391

REFERENCES I. D. M. Gruen, Nature 178, 1181 (1956). 2. D. M. Gruen, J. lnorg. Nucl. Chem. 4, 74 (1957). 3. D. M. Gruen, Spectroscopy of transition metal ions in fused salts. Fused Salts, (Edited by B. R. Sundheim), Chap. 5. McGraw-Hill, New York (1964). 4. G. P. Smith, Electronic adsorption spectra of molten salts. Molten Salt Chemistry, (Edited by M. Blander). lnterscience, New York (1964). 5. D. M. Gruen, Quart. Rev. XIX(4), 349 (1965). 6. 1. V. Yananayev and B. B. Dzhurinskiy, Dokl. Acad. Nauk SSSR 134, 1374 (1960). 7. D. M Gruen and R. L. McBeth, J. Phys. Chem. 63, 393 (1959). 8. D. M. Gruen, P. Graf, S. Fried and R. L. McBeth, Proc. 2nd Conf. Peaceful Uses of Atomic Energy, Geneva, 28, 112 (1958). 9. H. C. Egghart, J. lnorg. Nucl. Chem. 31, 1538 (1969). 10. V. Gulmann and O. Leitmann, Monatsheflefiir Chemic 97(3), 926 <1967). II. J. Badoz-l,ambling, Bull. Soc. Chim, Fr. 1195 (195(I). 12. R. M. Wallace and E. K. Dukes, J. Phys. Chem, 65, 2094 (1961). 13. D. Bunn, F. S. Dainton and S. Duckworth, Trans. Faraday Soc. 57, 1131 (1961). 14. D. Seewald and N. Sutin, lnorg Chem. 2,642 (1963). 15. P. Senise. J. Am. Chem. Soc. 81. 4196 (1959). 16. G, Saini and G. Ostacoli, J. lnorg. Nucl, Chem. 8, 346 (1958), 17. H. K EI-Shamy and M. F. Nassar, J. lnorg. Nucl, Chem. 16, 124 (1960). 18. F. (;, Sherif and W. M. Oraby, J. lnorg. Nucl. Chem. 17, 152 (1961). 19. F. G. Sherif :and W. M. Oraby, J. lnorg. Nucl. Chem. 24. 1373 (1962). 20. E Rhodes and A. R. Ubbelohde, Prec. R. Soc. London A251, 156 (1959). 21. See for example: T. M. Dunn, Modern Coordination Chemistry (Edited by J. Lewis and R. G. Wilkins), p. 266. Interscience, New York (1964). 22. F. A. Cotton, D. M. L. Goodgame and M. Goodgame, .L Am. ('hem. Soc. 83(23), 4690 (1961). 23. V. Gutmann and H. Lausegger, Monatshefte fiir Chemic 98~2), 439 (1967). 24. W. Beck. W. P. Fehlhammer, P. Pollmann, E. Schuierer and K. Feldl, Chem. Ber. 100, 2335 (1967). 25. S. Yamada and R. Tsuchida, Bull. Chem. Sot., Japan 26, 15 (1953). 26. J. Bjerrum, C. J. Ballhausen and C. K. Jorgnesen, Acta Chem. Scand. 8, 1275 (1954). 27. H. C. Egghart, Inorg. Chem. 6, 2121 (1967). 28. A. Abragam and M. H. L. Pryce, Proc. R. N,c. London A206, 173 (1951). 29. H. C7. Egghart. J. Phys. Chem. "/3,4014 (1969) and refs. therein. 30. B. [. Ewms, A. D. Yoffe and P. Gray, Chem. Rer. 59(4), 515 (19591. 31. A D+ Yoffe, Inorganic azides. In Development in Inorganic Nitr<~,en (*emistu (Edited by C. B. ColburnL Vol. I. Elsevier. Amsterdam (1966), and refs. therein.

f~)22 1902/8I/(~61391~3502.1D/0

L inore nml ('hem \ o ] 41 pp 1391-1~(1:1981 Prinled in Great Britain

PergamonPressLid.

Kinetics and salt effects on the reaction between octacyanotungstate(IV) and periodate ions (Received 2 May 1980: received for publication 19 Septemher 1980) The kinetics of the oxidation of octacyanotungstateIIV) by periodate ions have been studied in neutral and weakly alkaline solutions. The reaction is of the first order with respect to both reagents and it is catalyzed by the hydrogen ions, Inorganic

cations have an accelerating effect in the order: I.i" < Na +< K* < Rb÷
1392

Notes

anotungstate(IV) ions have been investigated to compare the redox properties of the cyanocomplexes of Fe(ll), W(IV) and Mo(IV). The oxidation of hexacyanoferrate(ll) by periodate has been previously studied [I].

40c

EXPERIMENTAL

70

K4W(CN)s'2H20 was prepared as recently described[2]. Solutions of this compound were prepared by weight and used immediately or else stored in the dark for no more than one day. Analytical grade KIP4 was used without further recrystallization. Buffer solutions were prepared'from KH2PO4 and KOH. Potassium nitrate was used to keep constant the ionic strength and the potassium ion concentration. All other reagents were of the Merck pro analysi quality. The reaction was followed by measuring the decrease in W(CN)s4- at 430 nm using a Beckman DBGT spectrophotometer. At this wavelength the extinction coefficient of the W(CN)84- is 110 M-' cm '. The temperature was controlled to within -+0.1 K. The pH's of the reaction solutions were measured on a pH-meter Radiometer Copenhagen.

~so % /

x

/

"

,

30c

50

5

RESULTS

~0

[H+].IOSM.

The determination of the stoichiometry of the reaction shows that the W(CN)84- and periodate react in the proportion of 2: 1, so that the overall reaction can be represented as follows:

Fig. 1. Dependence of k2 on [H+]. [W(CN)s4-]=4 • 10-3M, [104-1 = 5.10 -3 M and [K+I = 0.080 M.

IO4 + 2W(CN)84 + H20----~ 103 + 2W(CN)83- + 2OH-. results collected in Table 2 show that the reaction between periodate and octacyanotungstate(IV) ions is strongly cation catalyzed. The plots of k2 against the cation concentration are linear over the concentration range studied. The slopes of these plots, kcat, are indicated in Table 2.

(1) In neutral and alkaline media, where the kinetics of this reaction was investigated, the subsequent oxidation of W(CN)84by iodate is not thermodynamically feasible. The second order plots of I o g ( a - xl2)/(b-x) against time {a = [104 ]o, b = [W(CN)84 ]o and x = [W(CN)83-]} were linear to beyond 70% of reaction. Second order rate constants k2, were determined from the slopes of these plots. The results are collected in Table 1. The second order rate constant was found to vary linearly with the hydrogen ion concentration at various temperatures (Fig. 1), over the pH range 7.0-8.0. This dependence is given by eqn (2): k2 = k21+ kn+[H+].

DISCUSSION The results obtained for the stoichiometry and the kinetics of the oxidation of W(CN)s4- by periodate closely resembles the results obtained for the oxidation of Fe(CN)64 by periodate[1]. Frequently, these similarities in redox reactions of Fe(CN)64-, W(CN)64- and Mo(CN)64 ions with a common oxidant are observed and relationships between kinetic parameters and redox potentials have been reported[3,4]. These relationships suggest common reaction mechanisms. The following mechanism, similar to that proposed for the Fe(CN)64- oxidation by periodat~ may be assumed:

(2)

Below the pH 7 the linear dependence of k2 on [H +] does not hold. The values of k21 and k,, at 30°C are 1.58. M-~ sec -t and 1.66 (_+0.25). 10~ M-: sec -I respectively (ionic strength 0.1 M and K += 0.080 M), The heat of activation of the acid independent path is 17.5 _+1.2 kcal mole -1. The effect of the concentration of the various cations Li +, Na +, K +, Rb +, Cs + and Mg2+ on the reaction rate is indicated in Table 2. In all kinetic runs performed to measure the salt effects, no buffers were used in order to avoid indesirable salt effects from buffer salts composition. In these conditions the second order plots were satisfactory linear to beyond 10% of reaction. The

KI

H++ H4106- ~

H5106

(31

"H4IOE

(4)

'H4IO6-

(5)

KII

2H20 + I04-. r,, H++ H3IO62- .

Table 1. Second order rate constants for the reaction between periodate and octacyanotungstate ions

[w
pa = 7 , 5 8

5 ~ 1 5 5 5 i0 5

[ K ~ ] = O,O8O M

k2.1o2 < ~3..oi-I.n-i) 2,22 2,31 2,50 2,24 2,03 2,21 2,20 2,10 k 2 , av ~,18 ~ O,lO

• - 30,0 A 0,1

C

Notes

I

~,93

Table 2. Salt effects on the octacyanotungstate (IV)-periodate reaction

Salt added

LiNO 3 NaNO 3 ~03 RbNO 3 CsNO 3 Mg(N03) 2

Cone. salt added 0,025M O,050M 0,075M O,IOOM k + ~.I024£8] ~ 2 s -I) em 3 k2.102 ( dm3.mo~-l.s -I) (dm~mol~

1,86 1,92 1,98 2,12 2,26 3,35 a

2,00 2,16 2,28 2,53 2,8O 4,27 b

2,12 2,4~ 2,58 5,O5 5,4.0 5,13 e

[W(Cf{)~J;¢-~ lO.lO-3N , ~-"[I04J= 5.10-3M a) O,02M b) 0,O4 M e) 0,O6 M

W{CN)8 4

k~ +[04-----~W(CN)8 3 +IO42

(6)

k4 WICN U + HalO~ -----+ W(CN)83- + HJ|O6 2-

(7)

W{CN)/~ + HalO62 --~--~ W(CN)83 + H3IO63-.

(8)

Equilibriums (3), (4) and (5L have been reported[5] from the species present in aqueous solutions of periodate. From these equilibriums it is possible to consider that, over the {H+] used, the periodate species present are IO4 , H4IO6- and H3IO62 [6]. Reactions (6), (7) and (8) are possible rate determining steps, The I(VI) intermediates formed react with a further W(CN)84- ion to form l(V). According to this mechanism the reaction rate is:

2,32 2,66 2,87 3,50 3,97 -

0,060 0,i00 0,i19 0,186 0,229

0,03 0,19 0,89 1,50 2,60

_

T = 30,0 C

which is in agreement with the experimental eqn {2i where k21=k5 and kw =(k3KmK,~+k4Km). The assumption that (KmK{p+Km)[H+}~I seem to be correct only at pH :.7. Below the pH 7 the linear dependence of k2 on [H ~] does not hold. The results showing the variation of the reaction rate with the metal ion concentration and the nature of the cation indicate a specific cation effect. These results may be interpreted by' ion pair formation between the cation and any of the reacting species. In reactions between ions of the same sign, such as (6~-(8), the ion pairs are usually much more reactive than the free ions. The greater reactivity of the ion pairs is related with the lowering of the electrostatic repulsion between the reactants, However, the cation may be more deeply involved in the electron tran~ffer mechanism by forming a cation bridge which facilitates the process[7]. The existence of a roughly linear relationship between the polarizability of the cation[8] and 4,, (correlation coefficient 0.98) supports the above mechanism.

rate = {k~llO4 I ~ k4[H410~, ] + ks[H3IO62-l}[W{CN)84-], (9) When the concentrations of the periodate species are expressed in terms of the H4106- concentration and relating the last with the total concentration of periodate species, eqn (91 can be expressed as: rate

{k3KmKn I + k4Knl)[H+] + ks (KluKill + ~ [IO4-]r[W(CN)84-]. (10)

Equation (10) could be approximated to the experimentally determined dependence of k: upon [H+]. Thus when (KmKnl+ Km)[H+]< I, eqn (10) takes form (11)

k: : (k~KuiKil I + k4Km)[H +] + ks

(11)

-Author to whom correspondence should be addressed.

Departamento de Quimica Fisica Facultad de Quimicas Universidad de Sevilla Spain

P. G U ARDADO A. MAESTRE M. BAI,ON*

REFERENCES I. Y. Sulfab, J. lnorg. NucL Chem. 38, 2271 (19%). 2. J. G. Leipoldt, L. D. C. Bok and P. J. Cilliers, Z 4n,rg. Allg. Chem. 407,350 (1974). 3. J. G. Leipoldt, L. D. C. Bok. J. S. van Vollenhmen md J P Maree, Z. Anorg. Allg. Chem. 434, 293 (1977). 4. M. Balon, F. Ferranti, A. Maestre and A, lndelli, (ia:,:. (him. 1tal. 110, I29 (1980). 5. I. Pecht and Z. Luz, J. Am. Chem. Soc. 87, 4068 11%51 6. K. Kustin and E. C. Liebermam J. Phys. Chem. 68. ~,86q (1964). 7. J. Halpern and L. E. Orgel, Disc. Faraday Soc. 29,732 (1960) 8. C. N. R. Rao, in A Hand Book of Chemistry and Phvsi~ s. p 2~,4 Affiliated East-West Press Pvt. Ltd. (1967i.

041221902,8/;0613931}~$320t};{I Per~ar>,nPre':s~td

J mor# nm/ ChemVol 43. pp. 13931~95.1981 Prinled mGrea~Britain

Far IR spectra of Th(IV) halide complexes of some heterocyclic bases (Receired 27 May 1980: in revised form 8 August 1980: reveived for publication 16 October 1980) In the course of previous investigations[l, 2], we have reported the %'nthesis and IR spectra of Th(IV) perchlorato, nitrato and

thiocyanato complexes of some heterocyclic baseN. Halogens arc common ligands in coordination chemistry forming coordinate