Potentiometric Study of Copper(1) Complexes with Univalent Anions in Acetonitrile John K. Senne and Byron Kratochvil Department of Chemistry, University of Alberta, Edmonton, Alberta, Canada
The standard potential of the Cu(l)-(0) couple in acetonitrile was found to be -0.5842 & 0.0005 V vs. the standard silver electrode. The stepwise formation constants of several complexes with Cu(l) in acetonitrile were determined by potentiometric titration of Cu(l) solutions with salts of these anions using a saturated copper amalgam indicating electrode. Log p1 and p2 values were: CI-, 4.3 and 10.2; Br-, 3.5 and 7.3; I-, 3.1 and 5.8; SCN-, 3.6 and 7.2. Nitrate association with Cu(l) is negligibly small. Evaluation of the potentiometric data required knowledge of the ion association constants of the titrant salts in acetonitrile. For NaSCN and Et4NN03these data were determined by high precision conductance measurements. Analysis of the conductance data by the Fuoss-Onsager equation gave for NaSCN, K, = 88* 4, A. = 189.8 i 0.2; and for Et4NN03, K, = 7 =t2, A. = 191.37 f 0.03.
THE BEHAVIOR of copper salts in acetonitrile is of interest since copper(I1) perchlorate has been shown to be an analytical oxidant of some utility in this solvent (1-3). The oxidizing strength of copper(I1) is a function of the anions present, their concentrations, and their interactions with copper(1) as well as copper(I1). To date the only formation constants which have been measured are those of chloride with copper(1) and copper(I1) by polarographic and spectrophotometric methods (4), and those of perchlorate, hexafluorophosphate, and tetrafluoroborate with copper(1) by conductance measurements (5). The main purpose of this work was to measure the formation constants of selected monovalent anions with copper(1) in acetonitrile. In conjunction with this study, the standard potential of the copper(1)-(0) couple was accurately determined, and the ion association constants of sodium thiocyanate and tetraethylammonium nitrate were measured. EXPERIMENTAL
Apparatus. A Sargent Model SR Recorder with a 10-mV full scale deflection was used for measurements of the copper (1)-(0) standard potential. It was provided with a variable and accurately known bucking potential and allowed continuous monitoring of cell potentials with an accuracy of 1 0 . 2 mV. Potential measurements in the complexation studies were made to the nearest 0.1 mV with a Leeds & Northrup 7554 Type K-4 Potentiometer. Three cells were used in the determination of the copper (1)-(0) standard potential. All were of the standard H.cell design, with three compartments separated by ultrafine sintered-glass frits. Electrical contact to the amalgam electrodes was made with a platinum wire sealed into the end of soft glass tubing. (1) B. Kratochvil, D. A. Zatko, and R. Markuszewski, ANAL. CHEM., 38, 770 (1966). (2) B. Kratochvil and P. F. Quirk, ibid., 42, 535 (1970). (3) H. C . Mruthyunjaya and A. R. V . Murthy, Indian J. Chem., 7 , 403 (1969). (4) S . E. Manahan and R. T. Iwamoto, lnorg. Chem., 4, 1409, (1965). ( 5 ) B. Kratochvil and H. Yeager, J . Phys. Chem., 73, 1963 (1969).
The indicating electrode for all potentiometric work was a saturated copper amalgam, prepared by electrolyzing an aqueous 0.1M CuS04 solution, slightly acidified with HzS04,at a current density of about 0.05 A/cm2 until the mercury cathode contained about 1% copper by weight. An anode of pure copper wire was used in the electrolysis. When first prepared the amalgam appeared homogeneous but within an hour or two a solid phase separated [the solubility of copper in mercury is about 0.00025% (@I. The amalgam was stored under aqueous 0.1M H2S04in a weighing bottle and exposed to air only when the bottle was opened for removal of amalgam. Silver electrodes were prepared by sealing platinum foil into soft glass and electroplating silver on the platinum from 50 ml of an aqueous solution containing 0.85 gram of AgN03, 1 gram of KCN, and 2 grams of KzC03. A current density of 30 mA/cmZ was applied for 10 min, after which the electrodes were rinsed with dilute HC104, water, and finally acetonitrile. When not in use the electrodes were stored in 0.01M AgN03 in acetonitrile. Of four electrodes prepared, one gave significantly different results relative to the other three and was rejected. The average potential difference between any two of the remaining three electrodes when immersed in 0.01MAgN03in acetonitrile was 0.2 mV. Unless otherwise specified, the reference electrode was a saturated silver-mercury amalgam in contact with silver perchlorate in acetonitrile. This electrode attained equilibrium quickly. Its reproducibility equals or exceeds that of the silver, 0.01M silver nitrate electrode, and it is not affected by jarring or shaking. The amalgam was prepared by placing 25 to 35 cm2 of silver foil into a glass tube with about 30 ml of mercury. The solubility of silver in mercury is about 0.035% (7), so the relative amounts are not critical. The tube was sealed with a flame and heated for 8 hours at 110 “C with occasional shaking; it was allowed to cool to room temperature overnight before being opened. The amalgam was stored in contact with silver foil in a dropping bottle, and portions were removed as needed with an eyedropper. Protection from the air was not necessary. Solid copper electrodes were prepared from pure copper wire by polishing with fine emery cloth immediately before insertion into previously deaerated test solution. The apparatus, procedure, and method of data analysis for the determination of association constants from conductance measurements have been described previously (5). Solution densities, d, were calculated from the equation d = do ATii, where do is the solvent density, f i is the molality, and A is a constant for each salt. The value of A was determined by measuring the density of a 0.01M solution of the salt in acetonitrile. For sodium thiocyanate and tetraethylammonium nitrate, the values of A were 0.054 and 0.062, respectively. The density of solid tetraethylammonium nitrate was determined as 1.14 g/cm3. Chemicals. Acetonitrile was purified by the method previously described (5). Copper (I) perchlorate was prepared by reaction of copper
+
(6) M. Hansen and K. Andergo, “Constitution of Binary Alloys,” McGraw-Hill, New York, N. Y . , 1958, p 588. (7) Ibid., p 24.
ANALYTICAL CHEMISTRY, VOL. 43, NO. 1, JANUARY 1971
79
metal with copper(I1) perchlorate in acetonitrile (8). It was recrystallized once from purified acetonitrile and dried under vacuum. Analysis for copper by EDTA titration gave 99.8% CuC104.4CHsCN (9). Silver perchlorate was prepared as previously described (5).
Tetraethylammonium perchlorate was prepared by metathesis from sodium perchlorate and tetraethylammonium bromide. It was recrystallized once from water and dried under vacuum. Silver nitrate tests for bromide were negative. Tetraethylammonium nitrate was prepared by neutralization of a 10% aqueous solution of tetraethylammonium hydroxide (Eastman Organic Chemicals, Rochester, N. Y.) through dropwise addition of concentrated nitric acid to pH 7.0. The resulting solution was evaporated to near dryness on a hot plate and dried under vacuum at 110 "C overnight. This material was analyzed in two ways, by passage through a strong cation exchange resin in the hydrogen form with subsequent titration of the nitric acid liberated, and by a redox method based on the reduction of nitrate ion by excess Fe(I1) in strong acid (10). Both methods gave purities of 100 i 1%. The above material after analysis was recrystallized once from purified acetonitrile, pulverized, and dried at 110 "C under vacuum before use. Sodium thiocyanate was purified by recrystallization of reagent-grade material first from purified acetonitrile, and then from water. The resulting crystals were dried under vacuum at room temperature, pulverized, and redried under vacuum at 110 "C. A purity of 99.9% was obtained by coulometric titration at constant current with electrochemically generated silver(1). Tetraethylammonium iodide was prepared by mixing equal volumes of 0.22M sodium iodide and tetraethylammonium chloride in technical grade acetonitrile. The sodium chloride precipitate was filtered off and the filtrate reduced to about 50 ml by evaporation on a hot plate. The crystals of tetraethylammonium iodide obtained on cooling were recrystallized once from purified acetonitrile and dried under vacuum at 110 "C. Reagent grade sodium perchlorate was recrystallized once from purified acetonitrile and dried under vacuum at 110 "C. A silver nitrate test for halides was negative. Copper(1) nitrate was prepared as described previously (8). Because of the instability of CuN03.4CH3CN,the pure solid could not be isolated. However, a 0.1M solution in purified acetonitrile was colorless and clear for several hours after filtration through Whatman No. 2 filter paper; after that time, a blue solid precipitated from solution (11). Dilute solutions were prepared by dilution of aliquots of the freshly filtered solution. The exact concentration of copper was determined by EDTA titration after evaporation of the acetonitrile and oxidation of the metal with dilute nitric acid. Procedure. For the determination of the standard potential of the copper(1)-(0) couple, the following cell was used:
The central compartment contained a solution identical to that in the copper half-cell and served to prevent diffusion of silver(1) into the copper half-cell. The solution above the copper amalgam was deaerated by bubbling argon presatu_
_
_
~
(8) B. J. Hathaway, D. G. Holah, and J. D. Postlethwaite, J. Cliem. SOC.,1961,3215. (9) J. S. Fritz, J. E. Abbink, and M. A. Payne, ANAL.CHEM., 33, 1381 (1961). (10) W. Leithe, ibid.,20, 1082 (1948). (11) H. H. Morgan, J . Clzem. SOC.,1923, 2901. 80
rated with acetonitrile through the solution for 15 minutes and then passing the gas over the surface for the duration of the experiment. The silver amalgam half-cell did not require deaeration. The cell potential was continuously monitored for 8-12 hours with a recorder and the best linear portion was extrapolated to zero time. Potential drift during this time was typically on the order of 2 to 3 mV. Data for the calculation of formation constants were obtained by titration of copper(1) with the appropriate ligand in the following cell:
The titrant contained, in addition to the ligand, the same formal concentration of copper(1) perchlorate, generally 5 X lo-", and supporting electrolyte as the copper halfcell, as suggested by Manahan and Iwamoto (4). By maintaining a constant concentration of copper throughout the titration, treatment of the data is simplified. The ligand concentration typically was varied over the range zero to 4 X 10-2M. Titrant was added in milliliter increments to the copper half-cell. The solution in the cell was deaerated continuously by passing argon gas, presaturated with acetonitrile, through it. As before, the central compartment contained solution identical to that in the copper half-cell. RESULTS AND DISCUSSION
Standard Potential of the Cu(1)-(0) Couple. In the determination of standard potentials, it is important to minimize liquid junction potentials. This can be accomplished by constructing a cell without transference using an electrode of the second kind, or by following one of several extrapolation procedures (12). Because suitable electrodes of the second kind are not presently available in acetonitrile, the second method was used. Three series of measurements were made on the cell
For simplicity the central compartment is not indicated. In the first series Mi was varied while the sum of Mi and MZ (the ionic strength) was held constant. As M i approaches zero, the compositions of the solutions approach that of pure supporting electrolyte and the junction potential approaches zero. Extrapolation of a plot of cell potential us. MI to M I = zero (Figure 1) yields a potential at a particular ionic strength from which the junction potential has been eliminated. Other series of measurements were taken in the same way as the first but at different values of ionic strength. Slopes and y-axis intercepts were calculated for each set of data by leastsquares, along with standard deviations for the intercepts. The values obtained at varying ionic strengths were: p = 0.03, E"' = 0.6811 f 0.0005 V; p = 0.01, E"' = 0.6815 f 0,0002 V ; and p = 0.006, E"' = 0.6811 f 0.0005 V. Each point on Figure 1 is the average of three measurements, each with a different cell. It should be possible to obtain the standard potential from an additional plot of potential us. a function of the ionic strength. However, such a plot for this system is not pos-
(12) R. G. Bates, "Determination of pH," John Wiley and Sons, New York, N. Y.,1954, p 47.
ANALYTICAL CHEMISTRY, VOL. 43, NO, 1, JANUARY 1971
682 I
I
I
I
I
I
I
I
I
I
I
1
0
681
CL
0
= 0.03
680 679 678 677 6 76;
I
I
I
I
I
I
I
I
I
I
1
2
3
4
5
6
7
8
9
10
C u ( 1 ) , A g ( I ) MOLARITY x lo3 Figure 1. Cell potential as a function of copper(1) and silver(1) concentrations at various ionic strengths
sible because the three intercepts in Figure 1 are the same within experimental error. This indicates that the activity coefficients of the copper(1) and silver(1) species are of the same magnitude since the potential for the cell reaction Cu(Hg) is given by
+ Ag+
Ag(Hg)
Copper(1) Complexes in Acetonitrile. Overall formation constants, PI and P2, for copper(1) with the halides, thiocyanate, and nitrate were calculated by combining the formation constant expressions with the mass balance equations for total ligand, [L] O, and total copper, [CUI'. Ion association of the cation of the supporting electrolyte, C+, with the ligand was also taken into account by including the association constant for C+L-, Ka,to give
+ Cu+
+ 0.0591 log ([Ag+]/[Cu+])+
Eoeii= Eoeeii
0.0591 log
(l) By maintaining identical concentrations of silver(1) and copper(I), the second term on the right hand side of the equation becomes zero. At the ionic strengths used neither Y A ~ + nor ycU+ will be unity. However, upon dilution the ratio of Y A ~ +to ycU+ would be expected to approach unity more quickly than the individual activity coefficients, and the cell potential should not be strongly dependent upon the ionic strength. The standard potential of the cell was therefore taken to be the average of the three intercepts, 0.6812 i 0.0002 V. To obtain the standard potential of the copper(1)-(0) couple, a correction for the free energy of amalgamation of copper is necessary. The value reported in water by Oku (13) is 236 f 2 calories. This quantity should be independent of the solvent. To determine the free energy of amalgamation of silver, measurements were made on the cell 0.01M AgN03 in A g ! acetonitrile A value of 0.0870 f 0.0005 V was obtained, which corresponds to a Gibbs free energy of -2006 10 calories. Taking into account the above values, and arbitrarily referring all potentials to the silver, silver(1) couple at unit activity, gives:
*
Half-Reaction AgC e = Ag(Hg) Ag+ e = Ag Ag+ (0.01MAgNOa) e = Ag Cu' e = Cu(Hg) Cu+ e = Cu
++ + +
+
[L]" - [CULI - 2[CuLZ-l [L-1 = ___ 1 K&+l
YAs+ (YC"+)
E", volts 0.0870 k 0.0005 0.0000 -0.133 f 0.002 (14) -0.5942 dz 0.0005 -0.5842 =t0.0005
+
where [L-] is the free ligand concentration. Where C+ was sodium ion, the calculations had to include association with perchlorate [ K , for sodium perchlorate in acetonitrile = 10 (15)], but no correction was necessary for tetraethylammonium perchlorate, which is completely dissociated in acetonitrile (16). The free copper concentration was obtained from the Nernst equation and the cell potential. Using these equations, and assuming the concentration of CuL to be negligible at large ratios of ligand to copper(I), PZ may be calculated either point by point or by a plot based on the equation
AE
=
-0.0591 log [CUI"
+ 0.0591 log [CuLz-]/Pz
-
0.0591 Q log [L-I (3) which is readily derived from the Nernst expression and the equations at hand. AE is the potential shift from the value observed before any ligand has been added. A graph of A E us. log [L-] will have a slope of 0.0591 Q where Q is the subscript of the highest order complex formed. PZ may then be calculated from the intercept. The data were analyzed by the latter method and in all cases except nitrate, Q was found to be two. Values of PI were calculated on a point by point basis from data taken at ratios of ligand to copper of less than two. Combining expressions for PI, P2, Ka, and total ligand concentration gives PI
=
[L]" - [L-I - 2Pz[Cu+l[L-12 - Ka[C+IL-l [Cu+I[L-l ~~
(4)
The uncertainties in the above values are standard deviations. (13) M. Oku, Sci. Rep. Tohoku Imp. Utziu.Ser. 1 , 2 2 , 2 8 8 (1933). (14) B. Kratochvil, E. Lorah, and C. Garber, ANAL.CHEM.,41, 1793 (1969).
(15) R. L. Kay, B. J. Hales, and G. P. Cunningham, J . Phys. Chem., 71, 3925 (1967). (16) I. Y. Ahmed and C. D. Schmulbach, ibid., p 2358.
ANALYTICAL CHEMISTRY, VOL. 43, NO. 1, JANUARY 1971
81
Table I. Complex Formation Constants of Some Anions with Copper(1) in Acetonitrile K , for Supporting titrant electrolyte, Titrant salt 0.1M Log PI Log P a Et4NCl 35 (Ref. Et4NC104 4.23,4.35O 10.20, 10.24 20)
EtrNBr
10 (Ref.
Et4NC104 3.47, 3.52
7.36, 7.29
Et4NI
17) 5 (Ref.
EtrNC104 3.10, 3 . 1 3
5.86, 5.96
NaCIO,
5.69, 5.67
17, 18)
NaI
0 (Ref. 19) NaSCN 88c Et4NNOa 7d
3.07
NaC104 3.57, 3.57 7.33, 7.18 Et4NClOd Unassociated a Previously reported value, 5.9 (4). * Previously reported value, 10.8 (4). c This work. d This work. Previous estimated value was 25 (20).
Table 11. Conductance Data for NaSCN and Et4NN03 in Acetonitrile at 25 "C 104c
NaSCN 7.580 11.182 14.005 16.776 19.668 22,529 25.159 29.029
NaSCN 6.991 11.248 16.098 19.718 24.994 28.948 33.581 36.687
A 1 0 8 ~= 170.34 165.00 161.44 158.33 155.35 152.62 150.36 147.25
AA 3.98 0.02 -0.04 -0.02 0.03 0.03 -0.03 0.02 -0.01
1 0 8 ~= 3.32
171.30 165.03 159.25 155.40 150.71 147.52 144.10 142.08
0.00
-0.01 0.05 -0.08 0.04 0.02 -0.06 0.02
104c
EtdNNOa 3.524 5.674 7.704 9.687 11.842 14.152 16.349 19.134
Et4NNOs 3.338 5.573 7.302 9.679 11.133 12.570 13.648
A AA 108~= 7.54
183.78 181.61 179.91 178.58 177.24 175.88 174.73 173.40
0.02 -0.01 -0.06 0.02 0.04 0.00
-0.01 -0.01
1 0 8 ~= 7.50 183.99 0.00 181.77 0.00 180.34 0.00 0.00 178.65 177.74 0.01 176.86 -0.01 176.27 0.00
Table 111. Conductance Parameters and Averaged Conductance Values Salt Trial AO ii K, NaSCN 1 189.63 f 0 . 1 6 2 . 5 i 0 . 5 84 i. 3 NaSCN 2 1 8 9 . 9 4 f 0.13 3 . 4 i 0 . 6 91 i 3 EtaNNOa 1 1 9 1 . 4 0 f 0.09 4f 1 9 f4 EtrNNOa 2 1 9 1 . 3 4 4 ~0.04 3 . 5 f 0 . 7 5 f2
uA 0.05 0.03 0.04 0.01
(17) A. C. Harkness and H. M. Daggett, Jr., Can. J . Chem., 43, 1215 (1965). (18) G. Kortum, S . G . Gokhale, and H. Wilski, Z.Phys. Chem. (Frankfurt am Main) 4, 86 (1955). (19) R. P. T. Tomkins, E. Andalaft, and G. J. Janz, Trans. Faraday Soc., 65, 1901 (1969). (20) C. W. Davies, "Ion Association," Butterworths, London, 1962, p 96.
82
Similarly, expressions for PI, Pz, and total copper(1) concentration give [CUI" - [Cu+] - P*[Cu+][L-]2 Pl = (5) [Cu+l[L-I Combining Equations 4 and 5 such that PZ is eliminated results in a quadratic expression from which the free ligand concentration may be calculated.
+ + Ka[C+I)[L-l - Ll" -
Pdc~+l[L-l~ (1
+
[CU+] [CUI"= 0 (6) The free copper concentration may be calculated from the change in cell potential, thus all quantities necessary for the calculation of PZ are available. Values of p1 and pz are given in Table I for the anions investigated. Conductance measurements of copper(1) salts of the symmetrical anions perchlorate, tetrafluoroborate, and hexafluorophosphate give association constants of 0, 9, and 15 while the corresponding silver(1) salts are unassociated (5). However, silver nitrate has an association constant in acetonitrile of 71 despite appreciable silver(1)-acetonitrile interaction. Therefore the association constant of copper(1) nitrate is of interest. This salt cannot be isolated in sufficiently pure form for high precision conductance work, and so cannot be studied by the technique used for the other copper(1) salts. The potentiometric method described here gave a K, value of 7 for copper(1) nitrate, but random errors in the measurement technique coupled with possible coordinating trace impurities indicate that the uncertainty in this value is on the order of + l o . Thus, 1 ppt chloride in the tetraethylammonium nitrate would be sufficient to give an association constant of 50. As an alternate approach, solutions of copper(1) nitrate in acetonitrile were prepared as described in the experimental section and used for potential measurements on the cell CuC104,0.0100M CuNOs, 0.01088M Cu(Hg) Et4NC104, 0 . lOOM EtdNC104, 0. lOOM A potential difference of 2.5 mV was observed, which could be accounted for completely by differences in concentration and by a liquid junction potential of 0.1 mV, estimated by the Henderson equation. We conclude that the best value for the association constant of copper(1) with nitrate is zero, with an uncertainty of 5 . This is on the same order as the association of copper(1) with the symmetrical anions listed above, and indicates that the relatively large K, value for silver nitrate is not due to unusual behavior of the nitrate ion, but to specific ir.teraction by the silver. Conductance Measurements. The conductance data for NaSCN and Et4NN03are given in Table I1 and the results as analyzed by the Fuoss-Onsager equation in Table 111. ~
~
RECEIVED for review July 17, 1970. Accepted October 27, 1970. Financial support by the National Research Council of Canada and the University of Alberta is gratefully acknowledged.
ANALYTICAL CHEMISTRY, VOL. 43, NO. 1, JANUARY 1971
