FORMATION CONSTANTS OF SILVER(I)CYANIDE COMPLEXES
3763
Formation Constants of Silver (I) Cyanide Complexes in Equimolar Sodium-Po tassium Nitrate Me1ts1
by H. Ti Tien Department of Chemiatry, Northeastern University, Boston 16,Masaachusetts (Received April 19,1966)
Using a gallium-filled silica glass reference electrode in conjunction with a silver electrode, the behavior of silver(1) in fused equimolar sodium-potassium nitrate in the presence of cyanide ions was studied potentiometrically. The change of e.m.f. of the cell before precipitation of silver cyanide is ascribed to the formation of complex ions in the melt, The species formed were assumed to be Ag+-CN- and Ag(CN)2-. The formation constants thus obtained for Ag+-CN- and Ag(CN)2- are, respectively, 1.1 X lo4 and 1.1 X lo8 in molarity units at 521'K. The discrepancy of formation constants for Ag(CN)2- between those previously reported and the one given here is pointed out. The use of galliumfilled glass reference electrodes in fused salt systems at high temperatures is briefly mentioned.
Introduction Formation constants of Ag(1) and cyanide ions in fused equimolar sodium nitrate-potassium nitrate mixtures were reported recently by Manning and Blander.2 They evaluated the formation constants for AgCN, Ag(CN)2-, and Ag2CN+ from the analysis of e.m.f. changes in concentration cells. Manning and Blander were apparently the first to report the fonnation constant of AgCN in any medium. The formation constant for Ag(CN)2- given by Manning and Blander is 3.1 X 1Olo in mole fraction units at 246' which is in excellent agreement with the value given by Jordan and Pendergrast. Previously, Flengas and Rideal' using a similar type of concentration cell (and under essentially identical conditions) carried out electrometric titration of Ag(1) with potassium cyanide. The formation constant for Ag(CN)2- reported by Flengas and Rideal is 8.5 X lox2in mole fraction units at 250'. In view of the large discrepancy between the reported values for Ag(CN)z-, the Present study Was initiated. The purpose of this study, in addition to re-evaluating the formation constants of silver(1) cyanide complexes, Of the glass was to demonstrate the reference electrode to studies of this nature."' Experimental Section All reagent grade C h ~ c a l Were s used without fur ther purification other than oven drying at 150' when-
ever feasible. The electrolytic cell assembly, furnace, and general procedure were essentially the same as described previously6with minor changes noted below. A larger crucible furnace (Type M-506, Hevi-Duty Electric Co., Milwaukee, Wis.) was used. The temperature of the furnace was controlled by a Wheelco temperature regulator. Instead of using a Hg-filled Pyrex glass reference electrode, the reference electrode was constructed with Vycor glass and filled with gallium. The advantage of a Ga-filled Vycor glass electrode permits measurements at much higher temperatures. The construction of the gallium-filled glass reference electrode was essentially the same as previously reported.6 Instead of using platinum or tungsten wire to establish electrical contact, carbon rods had to be used since metal wires were rapidly attacked at elevated temperatures. Additions of AgN08 and NaCN to the melts were made using a slightly modi(1) Presented at the 149th National Meeting of the American Chemical Society, Detroit, Mich., April 4-9, 1965. (2) D*L*Manning and M. Blander, I W V . Chem., 1,694 (1962). (3) J. Jordan and J. Pendergrast, Proceedings of the 7th International Conference on Coordination Chemistry, Stockholm, 1962, p. 102. (4) 8.N. Flengas and E. Rideal, Proc. Roy. SOC.(London), A233,443 (1956). (5) G. W. Harrington and H. T. Tien, J. Phys. Chem., 66, 173 (1962). (6) H. T. Tien and G. W. Harrington, Inorg. Chem., 2 , 369 (1963). (7) H. T. Tien, Anal. C h m . , 36, 929 (1964).
Volume 69, Number 11 November 1966
H. TI TIEN
3764
140
120
100
d
3
80
60
+
AgN03 sNaCN in Ga-filled glass (1) Ag equimolar Na-K reference electrode nitrate melt
40
20
0
El = E'
0
RT
-In C A ~ F (3)
+
AgN03 xNaCN in AgN03 in Ag (4) Ag equimolar Na-K equimolar Na-K nitrate melt nitrate melt
RT --
CA~ In F [&+I
(5)
The results of the measurements for typical runs at 248" with initial silver(1) ion concentrations at 2.76 X (crossed circles), and 4.23 X (circles), 3.52 X in molarity units are shown in Figure 1. The use of molarity is a matter of convenience since the present data will be discussed in terms of Bjernun's ion association theory.8 Assuming that the change of e.m.f. of the cell (4) as a function of the added cyanide ions is due to complex ion formation, the over-all formation constants are calculated according to a modification of the mathematical treatment developed by Bjerrum and Leden.v The results of these calculations and the graphical The Journal of Physical Chemistry
10 15 NaCN added, M
x
20
25
30
106.
Figure 1. Variation of cell potential "ith concentration of NaCN.
in the absence of cyanide ions (Le,,x = 0) and
LE=
5
extrapolations are presented, respectively, in Table I and Figure 2. The extrapolated formation constants for AgCN (PI) and Ag(CN)2- (p2) are, respectively, 1.1 f 0.1 X lo4and 1.1 f 0.3 X lo8 in molarity units. The values obtained in this study are in good agreement with the formation constants reported by Manning and Blander and by Jordan and Pendergrast. The formation constant pz appears to be inconsistent with the result obtained earlier by Flengas and Rideal. A summary of these values is given in Table 11. It is interesting to note that Jordan and Pendergrast used a nonelectrochemical method (thermometric enthalpy titration) for obtaining their formation constants. Two significant points should be made in regard to this study. First, the cell as represented by (l), in which silver ion concentration was changed, was shown to be quite stable in the absence of added cyanide ions. For a typical run, the experimental slope is 105 k 1 m.v. at 248" as compared with the theoretical slope 104.5 calculated according to eq. 2. The data for the run, together with a schematic diagram of the glass reference electrode, are shown in Figure 3. The use of (8) H. T. Tien, Abstracts, 149th National Meeting of the American Chemical Society, Detroit, Mich., April 1965, p. 7M; details to be published. (9) H. T. Tien and G. W. Harrington, Ivwrg. Chem., 3, 215 (1964).
FORMATION CONSTANTS OF SILVER(I) CYANIDE COMPLEXES
3765
~~
Table I : Electromotive Force Data Used to Determine the Formation Constants for Species in Equimolar NaNO~-KNO~ Melt" AE,
-Log
-
mv.b
CCN-
n
0 0.8 1.6 2.3 3.2 4.0 5.7 7.5 9.5 11.2 14.8 18.5 22.2 30.0 38.2 46.7 60.0 68.5
0 5.70 5.40 5.22 5.10 5.00 4.82 4.70 4.60 4.52 4.40 4.30 4.22 4.10 4.00 3.92 3.82 3.77
IAg +I lor
KN-1 x 10s
3.50 3.42 3.38 3.33 3.26 3.20 3.08 2.95 2.83 2.72 2.51 2.31 2.13 1.79 1.49 1.23 0.914 0.757
0.153 0.306 0.420 0.589 1.34 1.12 1.46 1.80 2.14 2.79 3.45 4.07 5.34 6.53 7.77 9.54 10.76
0
s
0 0
([CN-I) x 10-
x
. .. 0.0134 0.0268 0.0514 0.0603 0.0760 0.109 0.155 0.201 0.245 0.345 0.442 0.550 0.760 0.990 1.21 1.56 1.78
3
i
1.53 1.16 1.21 1.25 0.70 1.22 1.27 1.31 1.35 1.41 1.49 1.58 1.79 2.07 2.90 2.97 3.37
28.1 1.96 2.62 2.55
.I
...
1.07 1.17 1.17 1.17 1.11 1.13 1.18 1.29 1.53 2.32 1.96 2.11
e Data taken from the smoothed curve (E= Figure 1). silver ion concentration: 3.50 X 10-6 (molarity units).
' Initial
0
2
4
[CN-I
x
+ CN61 =
Ag+
=
8
8
10
101 M.
Figure 2. Bxtrapolation of over-all formation constants.
-
DETAIL Reference Electrode
Table 11: Values of Formation Constants (in molarity units)" in Equilmolar Sodium Nitrate-Potassium Nitrate Melt Ag+
0
0
S i l i c a or A h m i n o e11icate G l ~ ~ r ,
Ag+-CN-
IAa +-CN-l jr + CN-1
+ 2CN-
= Ag(CN-)*
[Ag(CN-)zl
Gallium Netal
" = [Ag+][CN-]* 82
PI
Flengss and Rideal (ref. 4Y Jordan and Pendergrmt (ref. 3)b Manning and Blander (ref. 2)O This studyd
...
2.1
x
10'0 (2600)
...
8.0
x
107 (2600)
-110
-140
'
1.1 x 104 7.8 X 109 (248") (1.1 f 0.1) X lo4 (1.1 f 0.3) X 10s (248O) -170
*
, The density of the melt is estimated to be 1.9 g./cc. based on the data given by Smithells, which are used to convert all formation constants: C. J. Smithefls, ''Metal Reference Book," Butterworth and Co. Ltd., London, 1955, p. 620. bOriginal values of formation constants given in molality units. Original values of formation constants given in mole fraction units. Solutions were prepared by weight.
Figure 3. N e m t plot of potential between equilibrated glass reference electrode and silver electrode 08. silver ion concentration. Theoretical slope, 104.5; experimental, 105 f 1; temperature 521°K.
gallium in place of mercury greatslyextended the usefulness of such reference electrodes in fused-salt solutions. To date, gallium-filled glass reference electrodes
have been successfully tested well above 360". It is anticipated that, if suitable ion-sensitive material can be found in place of glass, the suitability of this
1 5.20
5.00
4.80 4.80 -Log [Ag 'I.
4.40
4.20
4.00
Volume 69,Number 11 Novsm6er 1966
H. TI TIEN
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type of reference electrode would be practicable up to the boiling point of gallium (2070”at 1atm.).1° A relevant comment should be made in regard to the mechanism of the operation of the glass reference electrodes. As discussed previously,6 the equilibration of reference electrodes in fused salts is of paramount importance. Equilibration seems to follow an exponential rate suggesting an exchange process is taking place between the glass and the melt. That the superficial washing with water and drying in the air have no apparent deleterious effect on stability strongly suggests that more than the surface layer is involved during the equilibration. Some of these facts are reminiscent of the glass electrode used in pH determinations. It is a well-known fact that the glass electrode has a socalled salt “error” at high pH values, which implies that the glass membrane is sensitive to other ions in the solution. This fact has been amply demonstrated by the work of Lengyel and Sammt,” who used a sodium glass electrode to measure the sodium ion concentration in molten PbC12. Prior to this work, Horovitz12 has suggested that the glass membrane behaved like a solid electrolyte, and the operation of the glass electrode depended on the exchange reactions of cations, which were limited to those ions able to migrate into the glass, These ideas were fully justified and supported by Sollnerla (ion-exchange membrane electrodes), Delimarskii14 (sodium glass electrode in fused salts), and Rudin, et ~ 1 6 . ~ 5(sodium glass electrode in aqueous solutions). The suitability of a glass reference electrode in various fused salt solutions containing either Li, Na, K, or Na-K ions has been tested. I n the cases studied, the cell reactions obeyed the Nernst equation. Thus, it is clear that the glass electrode must act like a cation electrode reversible to alkali ions of the fused salts used. That the reason for equilibration is important may be explained on the basis of ion-exchange reaction between the glass and the melt. Considering the glass as a solid electrolyte-as first suggested by Horovitz, in a way not unlike synthetic ion-exchange resins, the residual charge of the silicon-oxygen lattice is balanced by cations (usually Na or K). Upon immersion in the melt, it is highly likely an exchange process takes place between the cations present in the glass and those in the melt. However, it seems that not all of the cations can participate in this exchange reaction, presumably on account of charge and ionic radius of the cation involved. Once equilibrated, the glass electrode behaves like a specific cation electrode for Li, Na, or K. I n the melt containing Na-K, as in the present case, a mixed-electrode function must have developed. The operation of the glass electrode as a reference The Journal of Physical Chemistry
electrode in different fused-salt solutions is thus apparent since the potentiometric measurements were carried out in the melt in which the glass electrode had been equilibrated. The constancy of the reference electrode potential is, therefore, believed to be the result of the exchange reaction of cations between the glass and the melt, whose composition was practically invariant during the course of a run. l6 Secondly, in the case studied the mathematical manipulation of data for the evaluation of formation constants according to the procedure previously d e scribedg appears to be valid and sensitive as evidenced by the results, which are in good agreement with the published values obtained by two different methods2va (see Table 11). The internal consistency and the accuracy of the formation constants thus obtained can be checked readily by substituting the results in the fundamental equation relating the formation constants and the equilibrium concentrations of the ligand and the metal ion (eq. 7 of ref. 9). For the present case the equation may be written as
The results of calculations are given in Table I11 using the data presented in Table I and the extrapolated 0values. Table I11 Bi[CN-]
0.306 0.420 0.589 1.12 1.46 1.80 2.14 2.79 3.45 4.07
0.0355 0.0510 0.0736 0.136 0.186 0.237 0.287 0,395 0.515 0.644
+ BnlCN-1‘
0.0347 0.0481 0.0686 0.137 0.184 0.234 0.287 0.393 0.511 0.629
(10) For detailed discussions, see (a) R. W. Laity in “Reference Electrodes,” D. J. G . Ives and G . J. Jans, Ed., Academic Press Inc., New York, N. Y., 1961, Chapter 12; (b) J. D.Corbett and F. R. Duke in “Technique of Inorganic Chemistry,” Vol. 1, H. B. Jonassen and A. Weiasberger, Ed., John Wiley and Sons, Inc., New York, N. Y., 1963,Chapter 3. (11) B. Lengyel and A. Sammt, 2. physik. Chem. (Leipzig), 181, 65 (1937). (12) K.Horovits, Nature, 127,440 (1931). (13) K. Sollner, J . Phya. Chem.. 49, 171,266 (1946). (14) Yu. K.Delimarskii, Ukr. Khim. Zh., 21,449 (1966). (15) G . Eisenman, D.0. Rudin, and J. U.Casby, S c k c e , 126, 831 (1957). (16) H. T. Tien, Ph.D. Thesis, Temple University, 1962.
FORMATION CONSTANTS OF SILVER(I)CYANIDE COMPLEXES
The physical significance of Table I11 is that the data obtained in this study can be adequately described by a two-parameter equation (eq. 6), thereby substantiating the postulated species in the melt for the concentration range investigated. It can be seen however that, at low cyanide concentrations, eq. 6 is only approximately correct in describing the experimental data. This fact can be traced in part to the failure to take into consideration the possibility of Ag2CN+formation at very low cyanide concentrations. The existence of Ag,CN+ and its formation constant have been reported by Manning and BlanderS2 It is interesting to note that, if data were independent of Ag+ concentration or were extrapolated to zero Ag+ concentration, the formation constant 81 can be readily estimated from Figure 1, according to the methods developed by Braunstein, Blander, and Lindgren" and by Braunstein, et aZ.'* Thus, it can be readily shown that from the limiting slope of the curve in Figure 1, a value 1.0 X lo4 for 81 is obtained as compared with the extrapolated value 1.1 X lo4in Figure 2. Braunstein, el al., state that their methods of calculat-
3767
ing formation constants Iead to a more reliable extrapolation with equal weights to all experimental points. It should be mentioned however that the method proposed by Bjerrum, which was not discussed in the paper of Braunstein, et al., and on which the present calculation is based, is also fairly simple and does not require more than one successive approximation. For interested readers, an excellent analysis of Bjerrum's mathematical formulation, together with those suggested by Leden and by Fronaeus, may be found in a paper by Sullivan and Hindman.lB
Acknowledgments. edge the assistance taining some of the also grateful to Dr. helpful discussions.
The author wishes to acknowlgiven by Lawrence Gray in obpreliminary data. The author is George W. Harrington for many
(17) J. Braunstein, M. Blander, and R. M. Lindgren, J. Am. Chem. Soc., 84, 1529 (1962). (18) J. Braunstein, H. Braunstein, and A. 9.Minano, Inorg. Chem., 3,
1334 (1964). (19) J. C. Sullivan and J. C. Hindman, J. Am. Chem. SOC.,74, 8091 (1952).
Volumo 60, Numbst 11 November 1066
