NIOBIUM SULFOSALICYLATE COMPLEXES
Jan., 1961
145
4 SPECTROPHOTOMETRIC STUDY OF THE NIOBIUM SULFOSALICYLATO COMPLEXES BY ORVALE. AYERSAND JAMESE. LAND (Contributedfrom the Ross Chemical Laboratory, Auburn University,Auburn, Alabama Receiaed JuZv Il, 1960
Niobium pentoxide, dissolved in NaOH solution, was found to react with sulfosalicylic acid in a stepwise fashion to form a water-soluble complex compound, which produced a greenish-yellow solution with maximum absorption at 310 mp. By using spectrophotometric data, the Nb: (SSal) ratio was postulated to be 1:2 in the predominant species and 1:1in the minor species present in a solution 10-6 to 10-3 M in sulfosalicylic acid. The stepwise formation constants, kl and kz, and the over-all formation constant, K, for the complexes in these solutions a t pH 2-4 were calculated to be 1.08 X lo4,4.18 X lo3and 4.52 X lo7,respectively. A solid compound was isolated from solution but its formula could not be determined.
Introduction The use of sulfosalicylic acid to form a coordination compound with niobium has been previously reported but mainly in connection with analytical schemes. Schwartz2 effected the separation of niobium and tantalum with it while Das Gupta and Dhars employed it for estimating microgram quantities of niobium in the presence of moderate excesses of tantalum and titanium. Sudarikov and Busarov4 reported that solutions of niobium(V) sulfate and niobium(V) chloride became greenish-yellow in color when sulfosalicylic acid was added and from these solutions isolated a compound to which they assigned the formula, (NbO)z(SSa1)3~4Hz0.5 It was our aiin in this study to determine the chemical constitution of the niobium sulfosalicylato complex species in aqueous solution. From absorption measurements in the ultraviolet and by using a modification of the methods developed by Newman and Hume,6 along with those of Mukherji and Dey,? and Brown and Land8r9 it was possible to postulate that two complex species were present in the concentration range studied and to calculate formation constants for the equilibria involved.
Experimental Measurements.--The equipment and conditions for obtaining the measurements were the same as previously described,s except that for the ultraviolet absorption measurements 2 em. matched silica cells were used. Materials.-The sulfosalicylic acid (Merck and Co., Inc., Rahway, N. J.), labeled 99.5-100.3% pure, was used without further purification. Solutions for ultraviolet absorption studies were prepared as follows: 0.066 g. of Nb206(Fairmount Chemical Co., h-ewark, N. J . ) and 1-2 E;. of KHSO, were fused together in a Vycor crucible until a clear melt was obtained. Thc hot crucible was removed from the heat and rotated in such a manner that the melt solidified on the sides of the crucible. The slightly cooled crucible with its contents was placed in (1) Based upon Orval E. Ayers' Iv1.S. thesis research. Presented at the Southeastern Regional Meeting of the American Chemical Society, November 5-7, 1959. (2) V. Schwartz, Z . ongea, chem., 47, 228 (1934). (3) A. K. Das Gupta and S. K. Dhar, J . Sea. Ind. Res., l2B, 396 (1953). (4) B. N. Sudarikov and Yu. P. Busarov, Neow. Khzm. (U.S.S.R.), 2 , 702 (1957). (5) The symbol (SSal) has been employed throughout this paper for the sulfsalioylato ligand. (6) L. Newman ana, D. N . Hume, J . A m . Chem. Scc., 79, 4571 (1957). (7) A. K. Mukherji and A. K. Dey, J. Inorg. NucI. Chem., 6 , 314 (1958). (8) S A. Brown arid J . E. Land, J . A m . Chem. Soc., 81, 3185 (1959). (9) S. A. Brown and J. E. Land, J . Less-Common MetcsZe, 1, 237 (1968).
about 200 ml. of hot 2 N HC1 solution which was maintained just below the boiling point. The precipitated oxide was allowed to digest for about 20 minutes and then filtered on a buchner funnel using No. 50 Whatman filter paper. The freshly precipitated oxide was washed with 200 ml. of hot water and rinsed thoroughly with acetone. The precipitate was then dried by drawing air through it until all the acetone was removed. The dried Nb& was added to 100 ml. of a 0.6% NaOH solution and refluxed for about one hour to effect solution of the oxide. This NaOH-niobate solution was allowed t o cool to room temperature, filtered to remove any particles of filter paper or unreacted oxide and the volume made up to 500 ml. with distilled water. An aliquot of this solution was analyzed for Nb. A 10-3 M sulfosalicylic acid solution was prepared by dissolving 0.1271 g. of sulfosalicylic acid dihydrate in a small quantity of water and then diluting to 500 ml. Solutions having a constant Nb concentration and varying sulfosalicylic acid concentrations were prepared for spectrophotometric scan as follows: using a buret, seven constant volume portions of the NaOH-niobate stock solution were measured into 25-ml. volumetric flasks. Varying amounts of standard sulfosalicylir acid solution and approximate1 2 N H C 1 4 were then added in order t o make, when eacz solution was diluted with water, final solutions between pH 2 4 and the snlfosalicylic acid concentrations of the re2.4 X lo-', 3.2 X spective solutions 1.0 X 6.0 X ;O+, 8.0 X 10-4 and 1.0 X M. 4.0 X Seven reference solutions were prepared in a similar manner except that they contained no Nb. Perchloric acid was used to make the final solution acidic because the perchlora.te ion is known to show little, if any, tendency to form complexes with metal ionic species. The concentration of the Nb in a given series of solutions was between 1 X and 2 X 10-6molar.
0 30
1c c
READ L E F T S C A L E /
'D
csso,
I
io4.
Fig. 1.-Curve 1, absorbance as a function of sulfosdicylato ligand concentration; curve 2, ligand concentration divided by absorbance as a function of ligand concentration.
146
ORVAL
E. AYERSAKD JAMES E.
Vol. 65
LAND
tions in each case were the same in composition and concentration except that the niobium was absent. Data from two typical runs are plotted in Fig. 2. Analytical .-Analysis of the NaOH-niobate solutions used in forming the niobium sulfosalicylic acid complex was accomplished gravimetrically. An aliquot of the solution to be analyzed was acidified with 2 N HC1 solution whereupon the Nb205 precipitated. After a short period of digestion, the precipitate was filtered with suction, using Whatman No. 42 paper, and thoroughly washed with hot water. The filter paper and precipitate were placed in a previously weighed crucible and heated slowly a t first, and finally to red heat to drive off all the carbon and volatile matter. The residue was weighed as Nb205 and then converted to Nb by the gravimetric factor 0.6990.
0.31_L
0.2
0.4
0.6
0.8
Ratio C N b / C T o t a l Fig. 2.-Continuous
06-
variations plots for niobate sulfosalicylato mixtures.
Results and Discussion Although it is known that the method of continuous variations possesses certain limitations it was tried for the determination of the maximum number of ligands per metal atom in the complex species. The results presented in Fig. 2 indicate a 1:2 ratio of the niobium to the sulfosalicylato group in the lop3to 7.5 X 10-4 M concentration range. Only one absorption peak a t 310 mp was noted in the absorbancy measurements made on the solutions used in these continuous variations studies and it was reasoned that it indicated only the concentration of the highest complex species in the over-all equilibrium constant equation K =
(1)
C[NbO(ES,l),]/(CNb)(c~S~l)2
For the calculation of this over-all constant K we modified a method previously employed by Rlukherji and D ~ Y .For ~ simplicity, let "x" equal the concentration of the complex species and ''a" and "b" the initial concentrations of the niobium species and the chelating agent, respectively, in a given solution. Using two concentrations, al, a2 and b l , bz, of the reactants, having the same value of x, ie., the same absorbancy ( A = 0.6 in Fig. 2), we have
a
IC
=
or 4(a2
z / ( a , - z)(bl - 2z)2 = z / ( a , - z)fb2 - 2 ~ () 2 )~
+ bz - a1 - b ~ ) ( z+) ~(4a1bl - 4 a A + b12 - bZ2)(2)
+ ( a h 2 - alb12) = 0
(3)
From the graph (Fig. 2) it was determined that a1 = b, = 4.2 x 10-4, a2 = 6.3 x and 3.3 X So equation 3 becomes bz = 3.7 X IO-Y(X)~ - 3.385 X lO-'(z) l o g CSSa,.
Fig. 3.--Deterininat,ion of the formation constant k, with two species present and one absorbing. The solution,s prepared by this procedure were placed in matched 2 cm. silica absorption cells and scanned from 400 t o 290 mp. The resulting curves showed the complex to absorb a t 310 mp and the absorption at this wave length increased as the concentration of the sulfosalicylic acid increased. Data taken from a typical set of curves are plotted in Fig. 1, curve 1. For the continuous variation studies, solutions of the NaOH-niobate and the sulfosalicylic acid solutions were prepared by dilution so that the molarity of each solution was the same. A series of 7 to 9 solutions for absorption measurements were then prepared by mixing varying amounts of these two solutions so that in each case the combined volume was constant. The pH was adjusted to approximately 4 by adding small amounts of HClA, but the quantity added did not cause any significant volume change in the niobium sulfosalicylic acid solution. Reference solu-
+ 2.803 X lo-"
=
0
(4)
and L was calculated to be 1.94 X or 1.444 X Using the latter value for substitution into equation 2 , since the other value gives an unreasonable answer, K was computed to be 4.52 X lo7. This over-all formation constant can be considered to represent the product of the two separate formation constants, lcl and kz, for the equilibrium involved in the formation of this complex species. For the evaluation of these individual step constants, absorption measurements were made on a series of solutions where the concentration of the niobium was constant and the conceiitration of the sulfosalicylic acid was progressively increased. Curve KO. 1 of Fig. 1 shows the increasing absorbance a t 310 mp with increasing concentration of the ligand. As no other absorption peak than the 310 mp was noted in these solutions and since the relative concentration of the ligand was so much
THERMODYNAMICS OF THERMOCELLS WITH FUSED OR SOLIDELECTROLYTES
Jan., 1961
greater than that of the niobium, it was considered that the main equilibrium would be the step, NbO m(SSa1) = [NbO(SSal)2], and k2 = (SSal)+, c[NbO(Ssal)l]/(CNbO (SSal)2-m) (CSSal) Using a method previously d e ~ c r i b e d it ~ , ~was possible to calculate Ao, the limiting absorbance, for this complex species from the straight line plot of A US. Cssa,/’A. With A0 known a modificationg of equation C13 of Kewman and Hume,6log(Ao ,4)/A = --(m)log Cssal - log IC,, could be used under these circumstances to evaluate k,, and here n = 2. In Fig. 3 a plot of log (A0 - A ) / A us. log CSSal is noted to give a straight line with a slope for m of unity (actually 1.025 using the method of least squares), which indicates that one (SSal) group is
+
147
being added in this equilibrium step. From this plot kz was calculated to be 4.18 X lo3, then kl is given by K / k z and equals 1.08 X lo4. When the NaOH-niobate solution was treated with the sulfosalicylic acid solution the greenishyellow color reported by Sudarikov and Busarov4 was observed. Treating this resultant solution with an equal volume of acetone caused a fmely divided greenish-yellow solid to precipitate. The results of the analysis of this precipitate, however, could not be interpreted to indicate any definite empirical formula. The infrared spectrum of this solid precipitate, using the KBr wafer technique, showed a-small absorption peak a t 10.9-11.0 which was believed t.0 represent t,he NbO bond.8
THERMOIlYNAR!tICS OF THERMOCELLS T171TH FUSED OR SOLID ELECTROLYTES BY KENNETH S. PITZER Department of Chemistry and Lawrence Radiation Laboratory, University oj California, Berkeley
4, California
Received J u l y ti?1060
The thermodynamic principles related to thermocells are reviewed. The equation for the potential of a cell with a single component electrolyte, such as a fused or solid salt, is similar to that for a cell with aqueous solution electrolyte after the Soret equilibrium has been established. In particular the total “transported entropy” of the ionic species reacting at the electrodes may be obtained nithout any complication from transference numbers, and values of this quantity are given for several cella. The meaning and measurability of the partial molal entropies of single ions in electrolytes are considered. It, is found that the transported entropy of metal ions in fused salts is approximately equal to estimated values of the partial molal entropies of these ions; hence the entropies of transfer are small. I n solid electrolytes the entropies of transfer are sometimes large and are discussed in terms of the probable conductance mechanisms.
Through the years several thermocells with fused salt or solid salt electrolytes have been investigated.’-5 Recent work in other fields gives us now a clearer picture of the conduction mechanism in the solid electrolytes as well as additional data for the liquids. Also some of the previous discussions of these thermocells are confused unnecessarily by apparent uncertainties in transference number. Hence it seems worthwhile to look again at the information on these systems. Thermodynamic Relationships.-The thermodynamics of a thermocell was derived by Eastman6 and in greater detail by Wagner.7 We shall follow the definitions and terminology of Agar and Brecks and confine our attention to cells with pure metal electrodes and simple MX, electrolytes where X is a halogen or nitrate. The e.m.f. of the cell & is defined as the electrical polential of a wire attached to the hot electrode less the potential of a similar wire attached to the cold electrode. The electrical work for v equivalents of electricity is determined (1) L. PoincarB, A n n . chzm. phys., [ 6 ] 21, 289 (1890). (2) H. Reinhold a n d A. Blachnr, Z. EEektrochem., 39, 290 (1933); H.Reinhold, %. anor&.allgem. Chen., 171, 181 (1928). (3) H. Holtan, Thesis, Utrecht, 1953; Tgds. Kjsmz Berouesen Met., 12, 5 (1952). 63, L ,419 (4) B. R. Sundheim and J . Rosenstreich, THISJ U U R X ~ (1959). (5) B. F. Markov, Doklady Akad. Nauk S.S.S.R., 108, 115 (1956). (6) E. D.Esstman, J . Am. Chem. Sac., 50, 292 (1928). (7) C. Wagner, Ann. Physzk, 3, 629 (1929); 6, 370 (1930). (8) ,J. N. Agar and W. G. Breck, Trans. Faraday Sac., 63, 167 (1957).
by the entropy absorbed from the heat reservoir surrounding the hot electrode when this positive electricity passes through the cell from the cold to the hot electrode. This entropy is equal to the sum of the entropy abso_rbed in the electrode reaction, in this case XM - S M +v - vS.+ ( M ) and the entropy transported away from the hot’ electrode region, in this case - 8 * M + Y - z 8 T e - ( ~ ) . Here EM is the molal entropy of the metal, SM and S*M are the partial molal entropy and the entropy of transand fer, respectively, of the M+. ion, and S,-(M) 2Pe- (M) are the corresponding properties of the electron in the metal M. These terms may be combined to yield +Y
+Y
where ??M tu, the total “transported entropy” of the ion M + Y , is the sum of the partial molal entropy of the ion and the entropy of transfer. The quantity ge-(M) is, similarly, the “transported entropy” of electrons in the metal M. It is desirable to note a t this point the relationship to aqueous solution thermocells. After the Soret equilibrium is established, equation 1 is applicable to the corresponding aqueous solution cell, but a t uniform electrolyte concentration in each half cell the expression contains the additional term t- X (S*M+Y vS*x-)where t- is the transference number of the negative ion. In the aqueous solution the sum of the two ionic entropies of transfer is the
+
