Molten potassium thiocyanate: The reactions of some compounds of iron, cobalt, nickel, copper and zinc

inorg, nucL Chem., 1976, Vol. 38, pp. 1795-1798. Pergamon Press. Printed in Great Britain

MOLTEN POTASSIUM THIOCYANATE: THE REACTIONS OF SOME COMPOUNDS OF IRON, COBALT, NICKEL, COPPER AND ZINC D. H. KERRIDGE and S. J. WALKER Department of Chemistry, The University, Southampton SO9 5NH, England (Received 27 November 1975)

Abstract--The electronic spectra of four transition metal cations stable in molten potassium thiocyanate [Fe(II), Co(II), Ni(II) and Cu(I)] are reported together with their reactions and those of three unstable cations [Fe(III), Co(IIl) and Cu(II)] with a variety of anions including cyanide, cyanate and sulphide. A novel reduction of zinc(II) sulphate is also reported.

INTRODUCTION

RESULTS AND DISCUSSION

A RECENT review of the chemistry of molten potassium thiocyanate[1] has revealed a number of inconsistencies and deficiencies in existing knowledge. In particular, with comparatively well known first row transition metals there have been conflicting reports on stabilities and spectroscopic measurements while other rather obvious reactions have not been studied at all. This paper reports studies on 15 compounds of five such transition metals in a total of eight oxidation states.

Iron(lI) chloride was readily soluble (-0.15M) in potassium thiocyanate at 200°, giving a pale yellow solution with two absorption maximum at 8900 c m ~ (molar absorptivity (E) 21.5-+0.5 1mole ' cm ~) and 6000 cm -~ (E = 23.5 -+0.5 1mole -L cm ~). Though these maxima are of somewhat lower wavenumber but of higher absorptivities than those of the only previous report by Jacmirski et al. [7] (10000 cm-~(16), 7900(13)), they are still considered indicative of octahedral coordination. The red shift of 1000cm ~' from the analogous absorption in aqueous solution is paralleled by the similar difference found with chromium(III)[8] which was attributed to the second sphere environment of potassium ions lowering the ligand field of the thiocyanate ions. The absorptions are thus assigned to the ~Eg ~ 5T2~ transition split by the JahnTeller effect. The absorption rose steeply beyond 25000 cm J, no doubt due to the charge transfer band at 34300 cm t (6340) [9]. There was no evidence of precipitation of iron(II) sulphide from solutions of more than 0.1 M as reported previously [10]. However such precipitation was obtained by two reactions not previously reported. Addition of sodium sulphide to an 0.08 M solution of iron(II) chloride caused the formation of a black precipitate (analysed as Feo.99sSo.995) which did not dissolve on addition of excess sulphide. Addition of potassium cyanate to a similar iron(II) chloride solution caused the slow formation of the same black precipitate (analysed as Fe0.996S0998). The rate of this reaction was much increased by the addition of glass wool, all the iron then being precipitated leaving a pale yellow melt. The reaction evidently involved an overall reduction by cyanate ions, i.e.

EXPERIMENTAL

Materials, Potassium thiocyanate was prepared as previously described[2]. Cobalt(II) and nickel(II) chloride were dehydrated with thionyl chloride [3] and copper(II) chloride with dry hydrogen chloride, the analyses being (Found: Co, 45.2; CI, 54.4. Calc. for COC12: Co, 45.4; CI, 54.6%). (Found: Ni, 45.0; C1, 55.1; Calc. for NiC12: Ni, 45.3; CI, 54.7%) and (Found: Cu 47.3, CI, 52.6. Calc. for CuCI2: Cu, 47.3; CI, 52.7%). Zinc(II) chloride was also dehydrated with hydrogen chloride but in the molten state[4] (Found: Zn, 48.0; CI, 51.9. Calc. for ZnC12: Zn, 48.0; CI, 52.0%). Iron(II) chloride was prepared by reaction of the metal with hydrogen chloride and iron(IIl) chloride sublimed in chlorine (Found: Fe, 44.2; C1, 55.6. Calc. for FeCI2:Fe, 44.0; CI, 56.0%. Found: Fe, 34.3; C1, 64.6. Calc. for FeC13:Fe, 34.4; C1, 65.6%). Copper(I) oxide was prepared by reducing copper(II) oxide with sulphur dioxide and extracting copper(II) sulphate with water (Found: Cu, 88.9; Calc. for Cu20: 88.8%). Copper(II) oxide and zinc(II) sulphate were dried at 120° for 24 hr and 300° for 12 hr, respectively (Found: Cu, 79.7. Calc. for CuO: Cu, 79.9%) and (Found: Zn, 40.3; SO4, 59.5%. Calc. for ZnSO4: Zn 40.5, SO4 59.5%). Six other compounds prepared by standard methods analysed as follows: (Found: Co, 22.0; CI, 39.9. Calc. for [Co(NH3)6]C13: Co, 22.1; CI, 39.8%). (Found: Co, 14.6. Calc. for Na3[Co(NO2)6]: Co, 14.6%). (Found: Ni, 24.3. Calc. for K2Ni(CN)4: Ni, 24.3%). (Found: Cu, 64.1, CI, 35.8, Calc. for CuCI: Cu, 64.2, CI, 35.8%). (Found: Cu, 52.3. Calc. for CuSCN: Cu, 52.3%). (Found: Cu, 35.1. Calc. for Cu(SCN)2: Cu, 35.3%). Procedure. Reactions were carried out in "Pyrex" tubes, with B.24 joints, which snugly fitted into an electric furnace connected to an Ether Transitol controller. The reaction tubes were closed with silica-gel drying tubes except when they were connected to the gas sampling apparatus. The temperature was maintained at 2000 unless otherwise stated. Anion analyses were performed using the tests developed earlier [5]. Physical measurements. Thermogravimetric analysis was carried out as described earlier[2], as was the X-ray powder photography[6] and gas analysisl5].

Fe2+ + SCN + 2 e ~ F e S + C N 2CNO ~ (CNO)2 + 2e.

(la) (lb)

Since part at least of the reaction was surface catalysed, it may have involved polymerisation or thermal decomposition of the oxycyanogen initially formed. Iron(IIl) chloride was immediately reduced to a pale yellow iron(II) solution and an orange-brown precipitate of parathiocyanogen (SCN)x, which analysed as S~oo2Co9~N0.999 and formed as previously reported[l 1]. The reaction was shown to be quantitative since the

1795

1796

D.H. KEV,RUX~Eand S. J. WALKER

absorptivities of iron(II) calculated from the iron(III) added were found to be 21.0-+0.5 and 23.2-+ 0.51 mole-' cm -1 at 8900 cm -~ and 6000 cm ', respectively. Cobalt(II) chloride dissolved readily to give clear blue solutions with three absorption bands (Table 1) which were slightly lower in energy but intermediate in intensity between the values previously reported/8, 12, 13]. The coordination is tetrahedral with nitrogen bonded thiocyanate and the 16800cm-' and 16500cm-' absorptions assigned to the transition 4T,(P)~4A2 split by spin-orbit coupling and the 7300cm-1 absorptions assigned to the 4T~(F) ~ 4A2 transition. The solutions were found to be stable even when of > 0.1 M concentration, in contrast to the claim [10] that cobalt sulphide was precipitated. Addition of potassium cyanide however produced a black precipitate which analysed as Co,.~So.gsr. The blue solution becoming colourless when the cyanide/cobalt ratio was ->2.5 from which point half the original cobalt was found to be still in solution as the hexacyanocobalt(III) complex. The stoichiometry was thus: 2Co 2÷+ SCN- + 5CN- ~ CoS + [Co(CN)~] 3-

(2)

spectroscopic cell. The reaction was also considerably speeded up on the addition of glass wool to the melt in a 2.5 cm diameter tube. A surface catalysed reduction process (similar to that postulated for iron(II) was evidently occurring with cyanate apparently acting as both the reductant and the source of oxygen to form cobalt(II) oxide Co 2++ OCN- + 2e-o CoO + CN-

(5)

which further reacted by an exchange process already observed [5]. O 2- + SCN- ~ S2- + OCN-.

(6)

Cobalt(III) complexes other than the hexacyano anion were found to react in potassium thiocyanate. Hexam-: mino cobalt(Ill) chloride being immediately and quantitatively reduced at 2000 with the evolution of ammonia. The blue cobalt(II) formed having [Co(NH3)6]CI3 + x KSCN ~ [Co(NCS)4] 2+ (SCN)x + NH3 + CI- (7)

rather than that proposed by Hennion et al. [9]. absorptivities calculated on the basis of the initial cobalt(III) of 16800 cm -~ (sh) (~ = 205 -+4), 15500 (280-+ 5) and 7300 (44-+ 1) in agreement with Table 1. The orangeThe black sulphide precipitate was not found to dissolve brown precipitate was confirmed as being on addition of excess potassium cyanide, up to a parathiocyanogen (SCN)x (analysed as So.999Co.99,No.9~). cyanide/cobalt ratio of 10, in contrast to the reaction A different reaction occurred with sodium hexanitpreviously claimed to occur in the presence of excess rocobaltate(III), reduction being immediate at 200° but cyanide [9]. producing a grey-black precipitate identified as cobalt(II) oxide. ["d" values by X-ray diffraction 2.13 (100), 2.42 2 C o S + S C N - + llCN-~2[Co(CN)6]3-+3S 2-. (4) (65) and 1.51 (50). A.S.T.M. Index for CoO 2.14 (100), 2.44 (75) 1.52 (50)] and evolving nitrogen dioxide, as well as Black cobalt(II) sulphide was also found to be precipi- forming a blue cobalt(II) solution and a deposit of tated when sodium sulphide was added to 0.2 M solutions, parathiocyanogen (So.9~Co.99~No,gD on the walls of the analysing as Co,.~So987 and showed no tendency to tube. Attempts to measure the absorption of the blue redissolve even with a sevenfold excess of sodium sul- solution after filtration were unsuccessful, as more greyblack precipitate was formed and it continued to be phide. Addition of potassium cyanate to cobalt(II) solutions produced for several hours until the solution was colourproduced no immediate reaction but over 24 hr the blue less. The difference in stability as compared with the colour faded and a black precipitate became visible, hexammino salt was evidently due to nitrite, a conclusion which initially had cobalt analyses of up to 76% decreas- supported by the observation that addition of sodium ing with time (an aged precipitate analysing as Col.0osSo.gs~, nitrite to cobalt(II) solutions produced a fairly rapid i.e. 65.3%Co). The colourless melt contained sulphide and precipitation of all the cobalt as the grey-black oxide. Nitrite has been found to act as a source of oxide in cyanide as well as cyanate but no sulphate, thiosulphate, nitrate or nitrite and showed no absorption other than an molten nitrites [6], e.g. edge at >31000cm -~. Thermogravimetric analysis of these melt solutions indicated no weight loss if the 3CoCh + 6NO2- ~ Co304 + 6C1- + 2NO2 + 4NO. (8) reactants were perfectly dry. It was not possible to measure the absorption spectrum of the blue solutions This is thought to occur via the equilibrium after the addition of cyanate since the black precipitate NO2- ~ NO + + O 2(9) started forming immediately after transfer to a 1 mm 3Co 2++ SCN- + 11CN--o CoS + 2[Co(CN)6] 3-.

(3)

Table 1. Cobalt(II) chloride spectra in molten potassium thiocyanate Present Work cm-I

c l.mole-lcm -I

cm-I

16800(sh)

210 ~

17hoo(sh)

16500

29o~ 5

160OO

7300

h5 ~ i

c 1 .mole-lcm-I

Eggh~rt 8

De Haas 13

Harrington and Sundheim 12 cm-I

¢ l.mole-lcm -I

cm-1

E l°mole-lcm -1

17800(sh) 99

16ooo

h8o

-

16oo0

51o

770o

6h

1797

Molten potassium thiocyanate displacement to the right in thiocyanate melts being assisted by the precipitation of the cobalt oxide and the reaction of the nitrosyl cations with thiocyanate or excess nitrite anions, xNO + + x S C N -*(SCN)~ + x N O

(10)

NO* +NO2 -*N203~NO2+NO

(11)

or

followed by oxidation in air NO + ~O-~ NO2.

(12)

Nickel(II) chloride dissolved to give a green solution with an absorption maximum in the visible wavelength regiorr closer to the value of Harrington and Sundheim[12] than those of Egghart[8] or De Haas et al.[13] though that in the IR was more similar to the quoted values. The molar absorptivities agreed most closely with those of Egghart. The values found were 13700cm -~ (E19.8-0.8) and 8300 (17.5-+0.7), as compared to the published values of 14000 (4.6) [12], 14300 (18) and 8550 (13)[8], 14300 (2.4) and 8300 (> 2.4)[13] respectively and may reasonably be assigned to the 3Tlg ~ 3A2, and 3T2g 4"-3A2g transitions of a nitrogen bonded octahedral complex. Potassium cyanide additions to a 0.2M solution of nickel(II) chloride or to a solution of potassium tetracyanonickelate(II), gave a yellow green solution with an absorption edge at > 23500 cm -~ with no indication of a precipitate. These observations are in better agreement with those of Hennion et al. [9] who measured an absorption maximum at 34400cm ' (e = 6160) rather than of Harrington and Sundheim[12] who reported a black precipitate. Sodium sulphide additions to nickel(II) chloride solutions produced a black precipitate (analysed as Nio 997So.997),which was insoluble even in a fivefold excess, together with a pale yellow melt which contained no nickel. Potassium cyanate in an 0.2 M nickel(II) solution reacted slowly to produce a similar black precipitate (Nio999So990 but much more rapidly in a spectroscopic cell or when glass wool was present. Thermogravimetric analysis indicated no weight change up to 400°C while cyanide and sulphide ions were formed in the melt. A similar reaction sequence is postulated as for iron(lI) and cobalt(II). Copper(I) chloride or thiocyanate dissolved readily at 200 ° to give a pale yellow solution with an absorption edge at >30000 cm -~, similar to a charge transfer transition at 34200 cm -~ (715) reported by Hennion et al. [9]. Addition of sodium sulphide gave a black precipitate of sulphide immediately (analysing as CuL999So.99s)which was insoluble in a fivefold excess. Copper(I) oxide was insoluble and unreactive to 350°. Copper(II) chloride and thiocyanate were both soluble and reactive producing an orange-brown precipitate of parathiocyanogen (51.002C0.996N0.999)and a pale yellow copper(I) solution. All the original copper was precipitated as copper(I) thiocyanate on dissolving the frozen melt in water (Cuo.999SCN). Copper(II) oxide showed no solubility or reactivity up to 300°, indicating that the early report [14] of the reaction 2CuO + 2KSCN--> Cu2S + K2S + 2CO + N2 must have occurred at higher temperatures.

(13)

Zinc(II) chloride was sparingly soluble (-0.005 M at 200 °) and did not react up to 370°, which extends the range of thermal stability stated[15] to reach 3000 and may be compared to the report [10] that a > 0.1 M solution reacted to form the sulphide above 195°. By contrast zinc(II) sulphate was found to be insoluble, though it commenced to react at 3400 producing an off-white precipitate with evolution of sulphur dioxide (IR absorption 1360cm ~) and a yellow sublimate of sulphur (Found: S, 99.8%). The remaining solution was pale yellow with no absorption bands from 4000cm 1 to an absorption edge at >25000 cm 1 which contained cyanide and-sulphide but no cyanate, sulphate, sulphite, thiosulphate, nitrate or nitrite. The precipitate which contained free sulphur, was analysed after extraction with carbon disulphide and shown to be zinc(II) sulphide (analysed Zr~1999Su~6).Thermogravimetric analysis showed a weight loss of 72.5-+ 0.5% beginning at 300° with a maximum rate at 390° and substantially completed by 420 °. Thus the overall stoichiometry of the reaction can be expressed as ZnSO4+2SCN ~ Z n S + 2 C N

+S+SO2+O~

(14)

since the weight loss found varied between the 59.5% calculated for the loss of sulphur dioxide and oxygen and the 79.3% if all the sulphur was volatilised. In this reduction of sulphate the presence of zinc is essential, since potassium sulphate (which was found to be moderately soluble -0.5 M at 200°) did not react up to 400 °, the decomposition temperature of the melt itself. It is noteworthy that cyanate and other sulphur or nitrogen oxidation products were not formed and thus the reduction process was not similar to that of the transition metal tetraoxyanions investigated earlier[2,5]. The detailed mechanism of this novel reduction was evidently more complex than indicated by eqn (14) since for example the melt also contained sulphide and its pale yellow colour and absorption edge are probably to be attributed to sulphide and polysulphide ions [sodium sulphide was shown to have an absorption maximum at 26200 cm (E 13.21 mole ~cm- t)] which are likely to play a role in the mechanism of the reduction. CONCLUSIONS The chemistry of compounds of five transition metal in eight oxidation states dissolved in molten potassium thiocyanate has been elucidated and conflicting reports [1] resolved. The d - d electronic absorption bands of three stable oxidation states, iron(II), cobalt(II) and nickel(II) have been measured and assigned to octahedral, tetrahedral and octahedral nitrogen bonded complexes respectivelyl These cations have been found stable at higher concentrations than had been reported [I0], but formed their insoluble sulphides on reaction with sulphide or cyanate anions. Potassium cyanide also produced some sulphide with cobalt(II) but oxidised half to the hexacyanocobaltate(IIi) complex. Copper(I) was stable, displaying a charge transfer absorption but could be precipitated with added sulphide ions. Copper(II) cations oxidised thiocyanate to parathiocyanogen as did some higher oxidation state compounds of iron(III) and cobalt(III), in the case of the latter when hexanitrocobaltate(III) was used cobalt(II) oxide was also precipitated. Zinc(II) was found to be stable up to 370° but, in an at present unparalled reaction, formed the sulphide in the presence of sulphate, the latter

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D. H. KERRIDGE and S. J. WALKER

oxidising thiocyanate to cyanide, sulphur and sulphur dioxide.

Acknowledgements----Grateful thanks are extended to the Science Research Council for granting and to Borax Consolidated Ltd., for sponsoring a C.A.P.S. award to S.J.W. REFERENCES 1. D. H. Kerridge, In Advances in Molten Salt Chemistry, Vol. 3, Plenum, New York (1975). 2. D.H.KerridgeandM.Mosley, J. Chem.Soc.(A),352(1967). 3. J. H. Freeman and M. L. Smith, J. Inorg. Nucl. Chem. 7, 224 (1958). 4. D. H. Kerridge, J. Chem. Soc. 1178 (1963). 5. D.H. Kerridge and M. Mosley, Jr. Chem. Soc. (A), 1874 (1967).

6. D. H. Kerridge and S. A. Tariq, Inorg. Chim. Acta 4, 499 (1970). 7. H. B. Jacmirski, S. W. Wokkow, I. I. Mostowska and N. J. Buriak, Roez. Chem. 46, 1999 (1972). 8. H. C. Egghart, Z Phys. Chem. 73, 4014 (1969). 9. J. Hennion, J. Nicole and G. Tridot, Analusis 1, 48 (1972). 10. A. Eluard and B. Tr~millon, J. Electroanal. Chem. 13, 208 (1967). 11. D. H. Kerridge and M. Mosley, J. Chem. Soc. (A), 429 (1969). 12. G. Harrington and B. R. Sundheim, Ann. N.Y. Acad. Sci. 79, 950 0960). 13. K. S. De Haas, P. A. Brink and P. Crowther, J. Inorg. Nucl. Chem. 33, 4301 (1971). 14. J. Milbauer, Zeit. Anorg. Chem. 42, 433 (1904). 15. T. Yanagi, K. Hattori and M. Shinagawa, Rev. Polarog. (Kyato) 14, II (1966).