Solid solutions of RuO2 and IrO2

1778

Notes Table 4. IR absorption frequencies (cm-~) of Th(IV) thiocyanate complexes

Complex

]C_-C, C-N and ring stretchings of heteroeyclic ring

"9(c~)

"~(cs)

~(scs)

.etaz-s

stretching

J

Th(NCS)4.4(Py)

1625s,1603s, 1530s, Dr, 1470m

2035vs

8t~Om

480m

300m

Th(NCS) 4.4(Pie)

162Oa, 161Om, 15~Om, 1480s

2o40vs

840m

kTOm

310=

Th(NCS)4.4(NH2PY)

166Ova,1620s, 1550s,br,1480m

203~vs

8h-Ore

470w

318=

Th(NCS) 4.4(2,4LN) 1620vs, br, 16OOsh, 1530m

2045vs, br

820m

475=

312m

ThCNCS)4.4( 2, 6LN) 1642s, 1618s, 1575vs, 1472S

204Ovs

820m

465w

31Om

Th(NCS)~..k(Q)

1632vs, 1592s, 1551s, 1490m

20kOvs

830m

465w

37 lm

Th(NCS) 4.4(Isoq)

1620s, 1580m, 1550m, 1~-80m

20t~3~s

840m

470m

359m

Th(NCS) 4" 2(Bipy)

1620m, 1605s, 1585s, 1568m, ll+)+Os

2045vs

840,.

470w

362"

Th(NCS) 4" 2(Phen)

1590s ,1495m, 1425m

201+Svs

840m

465m

330m

In all the thiocyanato complexes the three fundamental frequencies C-N stretch (v0, C-S stretch (v3) and N-C-S bending (v2) lie in the range 2080-2040, 860-760 and 480-465era-' clearly indicate that in all the complexes the thiocyanate group is N-bonded[3, 12, 13]. The metal-ligand vibration in these complexes has been tentatively assigned in 360-300cm-I region[1, 14].

2. C. E. F. Rickard and D. C. Woollard, lnorg. Nucl. Chem. Lett. 14, 207 (1978). 3. R. K. Agarwal, Mahesh Srivastava and A. K. Srivastava, J. lnorg. Nucl. Chem. in press. 4. B. C. Smith and M. A. Wassef, J. Chem, Soc. (A), 1817 (1%8). 5. V. V. Savant and C. C. Patel, J. Less Common Metals, 24, 459 (1971). Acknowledgement--The authors are thankful to U.G.C. New 6. S. S. Singh, Ind. J. Chem. 7, 812 (1969); Z. Natur[orchung. Delhi, for financialsupport. 24A, 2015 (1969); Israel J. Chem. 7, 471 (1969); Z. Anorg. AIlg. Chem. 384, 81 (1971); Labdev J. ScL Tech. 9A, 198 Chemistry Department R.K. AGARWAL* (1971); 10A, 14 0972). Meerut College A.K. SRIVASTAVA 7. B. J. Hathaway and A. E. UnderhiU, J. Chem. Soc. 309 Meerut MAHESH SRIVASTAVA (1961). India NEELU BHAKRUt 8. S. S. Krishnamurthy and S. Sounderajan, Can. J. Chem. 47, T. N. SRIVASTAVA? 995 (1969). 9. J. R. Ferraro, J. MoL Spectra 4, 99 (1960). 10. C. C. Addison and N. Logan, Adv. lnorg. Chem. and RadioREFERENCES chem. 6, 95 (1964). l. T. N. Srivastava, S. K. Tandon and Neelu Bhakm, J. Inorg. 11. T. Ueki, A. Zalkin and D. Templeton, Acta Cryst. 20, 836 Nacl. Chem. 40, 1180(1978). (1966). 12. J. L. Burmeister, Coord. Chem. Rev. 1, 205 (1966); 3, 225 (1968). *On leave from: L. R. (P. G.) College, Sahibabad (Ghaziabad) 13. Z. M. S. AI-Kazzaz, K. W. Bagnall and D. Brown, J. lnorg. and for correspondence. Nucl. Chem. 35, 1501(1973). fChemistry Department, Lucknow University, Lucknow, 14. R. K. Agarwal, P. C. Jain, Veena Kaput, Sunita Sharma and India. A. K. Srivastava, Trans. Met. Chem. 5, 237 (1980).

I. inorf.,ucl. Chem.Vol.42,pp. 1778-1781 © PetlpmmnPressLtd.,1900. PrintedinGreatBritain

0o22-190218011201-177815ol0ol0

Solid solutions of RuO2 and lrO2 (Received 30 November 1979; received for publication l0 March 1980)

The lower oxides of the platinum group metals have been widely tronic conductivity[l], chemical stability[2], and catalytic proused as catalysts. The dioxides of ruthenium and iridium in perties toward a great number of commercially important particular are of considerable utility because of their high elec- processes,particularly in the area of electrode reactions.

Notes Both RuO2 and lrO2 adopt the futile crystal structure at atmospheric pressure[3]. This similarity in the structure of the oxides, the similarity in crystallographic radii of the +4 ions[4], and the similarity of unit cell dimensions[3] suggest that solid solutions of the two oxides could be stable. Such solid solutions could provide useful combinations of catalytic properties at a cost less than that of IrO2. No information on the mutual solubility of these oxides are available in the open literature. The similarity in oxide structure does not assure mutual solubility in all proportions. Work in this laboratory and a published study [5] on the RuO2-TiO2 system, again a pair of rutile oxides, found miscibility restricted to a narrow range. Firing of RuO2 has been shown to markedly improve the corrosion resistance of RuO2 prepared from aqueous solution[6]. No information was available on the structure changes induced by thermal stabilization of RuO2 or IrO2 prepared by other methods. Of particular interest to us was the influence of thermal history on the morphology. For these reasons a study was undertaken to investigate the nature of this mixed oxide system in terms of the limits, if any, to mutual solubility and characterization of the morphological changes which occur during thermal stabilization. EXPERIMENTAL

Oxide preparation and characterization The oxide materials studied were prepared by the Adams fusion method. Appropriate quantities of RuCl3.2H20 and IrCI3. nH:O (Engelhard Industries) were ground together with 500% (w/w) of reagent NaNO3 (Fisher). This mixture was fused in an open Vycor dish and held at 500°C for 180 min. The fusion was permitted to cool in air and, after solidification, the reaction mixture was leached with distilled water to separate the oxide. The oxide was further purified by three distilled water

,

I

~

i

I

I

?

0>

1779

washings at room temperature, each followed by a 24 hr period to collect the oxide by settling. Thermal stabilization was carried out in open Vycor or quartz dishes in a muffle furnace. All of the materials used for the composition study were fired for 60 min at 550°C. Surface area and porosity measurements were made by nRrogen adsorption using a Micrometics Digisorb 2500 instrument. X-Ray diffraction measurements were performed using a Debye-Scherr camera with Ni-filtered Cu radiation. Ruthenium and irridium analyses were made by atomic absorption following a peroxide fusion to bring the metals into solution. RESULTS ANDDISCUSSION X-Ray diffraction measurements gave evidence of a single rutile phase to be present over the entire range of composition (Rut.x)Irx)O2. In particular, there was no evidence of a discrete IrO2 or RuO2 phase. Thus it appears that the two oxides are mutually soluble in all proportions. The RuO2 unit cell undergoes a smooth expansion to accommodate the slightly larger IP ÷ ions into the lattice; presumably the iridium is statistically distributed. Figure 1 presents some results of the X-ray diffraction studies. The unit cell volume is plotted as a function of the iridium content in the bulk phase. This parameter increases smoothly from pure RuO2 to that of frO2. Table 1 lists the lattice parameters for the compositions studied. Also listed are densities calculated from the unit cell, assuming a statistical distribution of Ir. There appears to be a slight high bias to the density data, as the literature value of RuO2 density is 6.97 g cm -3 as compared to 7.1 gcm -3 calculated. Measurements of the surface area and porosity of the (Ru,_x) lh) 02 materials are listed in Table 2. There is a consistent trend toward a reduced pore volume as the iridium content of the mixed oxide increases. This is probably not an artifact of the individual samples, as the pore volume shows a steady decrease despite considerable variation in surface area within the series. The (Ruo.~ Iro.t3) O2 system was selected for detailed examination of thermal stabilization effects. Table 3 summarizes the results of this part of the study. The prolonged high tem-

Table 2. Surface area and porosity of the mixed oxides Surface M3/g

Cumulative Pore Volume c c / ~

0.0

79.8

0.31

4.0

79.05

0.28

At % Ix IMetal Basis

62

I

I

I

I

I

20

4O

60

8O

100

62.5

0.20

73.1

0.16

49.6

50.7

0.18

68.0

44.8

0.11

100. O

88.7

O. l l

ATOMIC % IR

Fig. I. Unit cell volume as a function of IR content.

15.5 47.5

Table 1. Crystal parameters of IrOrRuO: system lattice parameters At ~ Ir

a

c

v

,,,~

0

4,48

3.104

62.30

3 7.1 g/cm

5.4

4.50

3.118

63.14

7,26

19.0

4,50

3,120

63.09

7.91

50.8

4,51

3o132

63.70

9.32

65.4

4,51

3,134

63.7

10.04

4,50

3,154

63.86

11.66

100

Notes

1780

Table 3. Influence of thermal treatment on solid morphology Stabillzatlon Temperature °C

Stabillzatlozz Time Min

Surface A[ea

Pore Volume

mL/g

cm3/g

0.235

None

None

550

60

62

0.198

550

480

30

0.171

700

60

18

0.152

800

60

5.6

0.027

900

60

3.3

0.004

900

480

2.4

0.002

100

l

128

I

l

I

I

(d) 900°C - 80 Min.

J

50

0

I

100

I

J ~

I

I

I

I

(c) 700°C • 60 Min. w

50

>o

£

01

I /

l

100


IE (hi

-

I

•.

X

50

0

I

I

I

I

I

100

I

I

I

I

I

(a} Post Fusion - NO Thermal Stabilization

I

I

I

50

100

200

I 300

I 50O

PORE DIAMETER (A~

Fig. 2. Influence of thermal history

perature treatment results in a very large loss of surface area and pore volume. Figure 2 presents the pore size distributions of several of the samples in Table 2. Increasing temperature and increasing time at temperature cause a loss fine porosity and eventually leads to the loss of all defined pore structure below

500 A. There was no evidence in the X-ray diffraction measurements for oxide phase separation or reduction to metal in the series listed in Table 2.

Acknowledgement--The authors wish to thank General Electric Company for permission to publish this work.

GeneralElectric Company Direct Energy Conversion Programs 50 Fordham Road Wilmington, MA 01887 U.S.A.

E. N. BALKO C. R. DAVIDSON

Notes REFERENCES I. W. D. Ryden and A. W. Lawson, Phys. Rev. BI, 1494(1970. 2. A. B. Nikol'skii and A. N. Ryabov, Russ. J. Inorg. Chem. 20, I (1%5). 3. D. B. Rogers, R. D. Shannon, A. W. Sleight and J. L. Gillson, Inorg. Chem. 8, 841 (1%9).

1781

4. G. V. Samsonov, Ed., Handbook of the Physico-Chemical Properties of the Elements. IFI/Plenum, New York 0%8). 5. Yu. E. Roginskaya, V. I. Bystaov and D. M. Shub, Zh. Neorg. Khim. 22, 201 0977). 6. C. Iwakura, K. Hiroo and H. Tamura, Electrochim. Acta 22,335 (1977).

o022=~9o2/80112oi-1781/$o2.o0/0

I. inoeg, nud Chem. Vol. 42, pp. 1781-1782 © PergamOn Press Ltd., 1980, Printed in Great Britain

Reactions of iron fluorides with hydrazine (Received 26 October 1979; recievedfor publication 10 March 1980) Fe(II) halides tend to coordinate two hydrazine molecules yielding chain complexes of the type [Fe(N2H4hX2],, where X = CI, Br, I[1]. Some authors reported the existence of [Fe(NzH4hCI2]'2H20, [Fe(NzH+hBr2].H20 and [Fe(NzH4h I2]'HzO[2]. The corresponding fluoride complexes have not been isolated so far, although similar compounds with ammonia exist; FeFz.SNHH3"H20, FeFz.NH3"HzOand FeFz.1/2 NH3'H20 were described long ago[3]. Hydrazines of Fe(III) halides have not been described so far. Yet the existence of the N2HsFeFs complex indicates, that such compounds might exist in spite of the reducing nature of hydrazine[4]. In this investigation we undertook a systematic study of the FeFz-N2H4 and FeF3-N2H+systems. The preparation of anhydrous hydrazine, the experimental technique and the methods of analysis were described elsewhere[5]. Iron was determined by EDTA titration with sulfosalicilic acid as indicatur[6]. Potentiometric determinations of hydrazine in the presence of divalent iron yielded too high results. We modifiedthe standard method[7] by adding 1,2-diaminucyclohexane-tetra,acetic acid into the starting solution at pH = 6.5; the DCTA complex was stabilized by using a 20% excess of hydrochloric acid for the reaction of N2H+ with KIO3. This modificationenables accurate determinations of Fe(II) with the same precision as the original method.

frequency thus indicating stronger bonds between the metal ion and the hydrazine ligand. That could he interpreted in terms of the stronger field effect of the fluoride ions as compared with the chloride ions. The (N-N)-stretching frequency also proves that hydrazine is bidentate and acts as a bridging ligand[9]. Weak hydrogen bonds are stronger in the fluoride complex, causing stretching vibrations to move to lower frequencies and deformationvibrations to higher ones as compared with Fe(NzH+hCI2.

(a) FeF2-N2H+system FeF2 was prepared by repeated hydrofluorination of commercial FeF2 (Koch-Light) or metallic iron; gaseous HF or molten N2HtF2 were used for that purpose. FeF2 reacts readily with anhydrous hydrazine; in few days the white solid turns first to light-brown and then to greyish-green solid. We noticed no gas or heat evolution during the reaction. Several experiments were carried out performing the reaction at 25 and 500C,removing the excess of liquid anhydrous hydrazine in vacuum between -30 and +25*(2, changing reaction times between one week and six months. The same greyish-greensolid was obtained with all the qualities of FeF2. The product was found to contain 40.1% N2H+, 35.5% Fe and 24.8% F thus suggesting the formulation Fe(NzH+hF2 (Calc: 40.58% N2H4, 35.36% Fe and 24.06% F). Fe0q2I~hF2 is stable in vacuum and in dry air. In open air iron hydrolyses giving off free hydrazine which protonates in acids. On heating Fe(NzH~F2 in argon flow, the decomposition starts slowly at 130 and hastens at 170"12.At 210"(2a strong exothermic peak occurs in the ffrA curve; all hydrazine is lost at that temperature and white FeF2 results. The decomposition in vacuum follows the same course; the first decomposition gases were observed between 160 and 170°C. The Mtsshauer spectrum of Fe(N2H+hF2 indicates divalent iron. The IR spectrum is similar to the spectrum of the chloride complex Fe(N2H+hCIz[g]. The (N-N)-stretching vibration, however, occurs at lower frequency and the (FEN) one at higher

(c) FeF3-N2H4system FeF~ was prepared by hydrofluorination of FeCI3'6H20, by fluorination of FeF2 with elemental fluorine at 500°C or by the reaction of iron po~vderwith XeF2[10]. When FeF3 reacts with an excess of anhydrous hydrazine, the colour of FeF3 changes from green over light-brown,dark-brown to light-green.The reaction is very vigorous at room temperature and the evolved gases cause severe pressure build ups, often leading to breakage of the glass reaction vessel. Therefore, the reaction is preferably carded out at lower temperatures (from -5 to +12')(2)and the decompositiongases are pumped off. The solid has 40.4% N2H(, 31.3% Fe and 20.1% F. Its Mtsshauer spectrum indicates that iron is in a high spin divalent state. Its IR spectrum is identical with the spectrum of Fe(NzH+hF2, but exhibits some + + additional hands and shoulders in the (NH 3)-deformation,(NH 3)bending, (N-N)-stretching and (N2H~)-torsionregion, indicating the presence of N2H~ ions. The product is homogenous and its IR spectrum differs from the spectrum of N2HsF. Thermal decomposition of the solid is entirely different from the thermal decomposition of Fe(N2H+hF2 with the exception that white FeF2 is the end product in both cases. It decomposes earlier then FeOq2H4hFr-at 60*(2,slower and in three steps: the first (DTA peaks at 192 and 212"(2) and the second (at 246°C) ones ate exothermic; the third one (at 2690C)is endothermic. The decomposition is finished at very high temperature (415°C), indicating different and stronger bonds as compared with Fe(N2H,hF2.

(b) FeFr-N2H+-H20 system Solid FeF2 was reacted with 80% solution of hydrazine hydrate. The reaction is exothermic; the dark-brown solid, which is formed at the beginning, changes its colour to green after prolonged reaction. The product is Fe(N2H4)2F2again with the typical composition 39.8% N2H4, 35.6% Fe and 22.7% F. IR spectrum of the product is identical with the one of the product prepared in the anhydrous system. The adduct reacts with open air. The presence of N2H~ and OH- species is observed in the IR spectrum of the resulting ochre solid; trivalent iron is indicated by the M6sshauer spectrum. Simultaneous oxidation and hydrolysis attains its equlibrium after 7 months: the product then contains 6.1% NzH,, 39.4% Fe and 20.0% F, corresponding to the formation of (N2Hs)o.3FeF,.s(OH)I.8"H20.