Table IV.
Zn+’ Concn., Mole/ Liter 0 00085 0 0100 0.0660 0 0660 0 0100 0.0235 0 0235 0 0118 0 0094 0 0235
*
Quinhydrone Taken, llIg./lOO 111. 5 1 5 0
6 0
Solubility Products from Zero-Current Potential Measurements
Addl. Supp. Electrolyte, Mole KXO3/ Liter 0.195 0.168 0
0,170 0.0294 0.0648 5 8 0.0718 5 2 0.0295 pH of buffer 6.80 iii all experiments. Prepared by author. Sample furnished by N. J. Zinc Co. 5 0 6 2 4 2
bility products for aniorphous zinc hydroxidr, where the solid phase is rapidly formed, and for zinc oxide. However, in the case of the slowly formed +hydroxide, of intermediate stability, unsatisfactory results were obtained; indeed literature values for the solubility product of this compound also must be regarded as doubtful. ACKNOWLEDGMENT
The author acknowledges with thanks the gift of a sample of orthorhombic zinc hydroxide by the ?Ten- Jersey
Solid Present
Temp., C. 24 24 24 21 25 O
e-Zn (OH), *
e-Zn(OH)2C e-Zn(OH)2c ZnO ZnO
25
24 24
E d . e , Zero Current, Volt us. S.C.E. Unknown In buffei~ $0.0084 I 0.0399 0,0741 + O 0°j5 0.0418 0.0613 0.0549 0.0624 0,0625 0.0498 0 0650 0,0735 0,0665 0.0815
Zinc Co. and many helpful discussions with T. R. Rubin. LITERATURE CITED
R. G., “Electrometric pH Determinations, Theory and Practice,” p. 74, R’iley, New York, 1954. (2) Ibid., p. 201. (3) Dietrich, H. G., Johnston, J., J. A m . Chem. SOC.49, 1419 (1927). (4) Feitknecht, W., Helv. Chim. Acta 13, 314 11930). (5) Ibid., 32, 2294 (1949). (6) Feitknecht, W.,Haberli, E., Ibid., 33,922 (1950). ( 7 ) Fricke, R., Meyring, K., 2. anorg. allgem. Chem. 230, 357 (1937). (1) Bates,
1
pH Unknown
k‘. p
7.83 7.26 6.72 7.19 6.98 7.08 6.71 6.60
3 . 9 x 10-16 3 . 3 x 10-16 1 . 8 x 10-16 2.4 X 2 2 x 10-16 1 . 7 x 10-16 2 . 5 X lo-” 3 8 X 10-l:
(8) Kline, W.D., J . Am. Chem. SOC.51,
2093 (1929). (9) Kolthoff, I. 11.. Lingane, J. J., “Polarography,” 2nd ed., p. 191, Interscience, New York, 1952. (10) I b i d . , pp. 246 et sep. (11) Kolthoff, I. M.,Orlemann, E. F., J . Am. Chem. Soc. 63, 664 (1941). (12) Muller, 0. H., Ibid., 62, 2434 (1940). (13) Muller, 0. H., Raumberger, J. P., Trans. Electrochem. Soc. 71, 169 (1937). (14) Tatl. Bur. Standards, “Selected Values of Chemical Thermodynamic Propert’ies,” 1952. RECEIVEDfor review 11ay 2 2 , 1958. Accepted July 14, 1958.
Pola rogra phy of M eta I-T hiosuIfate Cornplexes DONALD G. DAVIS Georgia Institute o f Technology, Atlanta ,The polarography of various metal ions was studied in thiosulfate media io establish analytical methods for such solutions and to determine the coordination numbers and dissociation constants of the metal ion complexes. Cadmium, lead, and copper(1) were reversible and determinable. Arsenic(lll), cobalt(ll), manganese, and zirconium gave no usable polarographic waves. Both zinc and nickel gave irreversible waves, but some useful information was obtained in each case.
P
studies in thiosulfate nicdin have been made for only a few inetal ions (10,12). However, thiosulfate iiiedia have been recommended for the electroplating of a number of metals ( 6 ) . Many studies have been made on metal-thiosulfate complexes (1, 4, 6, 14, 16), but a number of the results are mutually exclusive or have been obtained under rather extreme conditions. I n this work, conditions were used which are similar to those employed for OLSROGRAPHIC
73, Ga. electroplating and for the homogeneous generation of sulfide ion from thiosulfate ( 3 ) t o elucidate these processes and to establish analytical methods that could be used in conjunction with them. The thiosuIfate ion has the reputation of being a strong complexing agent mainly because of its extensive use in photography. The complexes of bivalent metal ions were rather weak, although the complex with cuprous is strong as might be expected. Little work was done on the complexes of iron, chromium, uranium, bismuth, and mercury because they decomposed readily to metal sulfides under the conditions used in this study. EXPERIMENTAL
All solutions were prepared from reagent grade chemicals. Stock solutions of t h e various metal ions studied were diluted with freshly prepared thiosulfate solutions of t h e appropriate concentration. T h e stock solution of cuprous ion was prepared immediately before use b y dissolving a weighed amount of previously asReagents.
sayed cuprous chloride in air-free 1.11 sodium thiosulfate solution. I n the studies of the eff’ect of complexing agent on the half-wave potential, the ionic strength was kept constant by the addition of potassium nitrate. Triton X-100 (Rohm & Haas Co.), 0.002%, was used as a maximum surpressor as recommended by Neites (11). Apparatus and Procedure. A Sargent Model 111 Polarograph was employed t o obtain t h e necessary polarograms, and potentials were measured with a Leeds & Xorthrup potentiometer. iiverage currents, corrected for residual current, were used for calculating the values of log i/id - i for the plots to test the reversibility of the electrode reaction. Values of the half-wave potential TT-err taken from these plots. An H-cell with a temporary plug of 370 agar saturated with potassium chloride was used. The agar plug had to be replaced weekly because of contamination and eventual darkening caused by the slow diffusion of thiosulfate ion toward the saturated calomel reference electrode (S.C.E.) and its eventual reaction with mercurous ion. The dropping mercury electrode had a value VOL. 30, NO. 1 1 , NOVEMBER 1958
1729
Table I.
Polarographic Characteristics of Metal-Thiosulfate Complexes in 1 M Sodium Thiosulfate
-Ei2, Volts Cd(I1) Pb(I1) Zn(I1) s'(I1) (IO) Tl(I1) ( I d )
0 0 0 1 0
i58
703 563
Slope 0 088 0 160 0 062
12
K
I
x 10-6 x 10-16 1 3 x 10-6
2.42 1 22 2 64 2 48 -3 16
Ligands
3
-3
7
50 5 0
63
of 1 7 ~ ~ ' 3 f of~ , ~1.SO. Oxygen 7vas removed with purified nitrogen, and all measurements were made a t 25" =k 0.1OC. The coiitrolled potential nork !vas done with manual apparatus as suggested by Lingaiie (9). A n agar plug vias used to back up the sintered-glass disk separating the compartments of a large H-cell. The cathode potential was measured against a saturated calomel electrode, t'he tip of which was just in contact IJ-ith the pool. Current integration was performed by the method of counting squares under the current time curve. The solution to be electrolyzed was deaerated before mercury n'as added to the cell. Potentials were measured with a Leeds & Sorthrup potentiometer, and currenbs n-ere nieasured n ith a calibrated milliamnieter. RESULTS A N D DISCUSSION
1
12
x
10-2
Figure 1. Polarogram of 0.804mM nickel in 1M sodium thiosulfate
was not found under the conditions used. The polarography of zinc is irrerersible, but the diffusion current is directly proportional to the concentrations. The formula Zn(Sz03)3-4 has been reported (4). The polarographic behavior of zinc in 1-11 thiosulfate is very similar t o that of zinc in 1.11 sulfate, the halfn-ave potential of the latter being about -1.06 X'olts US. S.C.E. Cobalt(II), nianganese(II), and arsenic(II1) gave no useful diffusion currents in 1 M thiosulfate. The waves of cobalt and manganese are masked by the reduction of thiosulfate ion which commences a t about -1.5 volts under
Polaropranis of 0.00111 cuprous: cadmium, and lead complexes were recorded in solutions of varying thiosulfate concatrations (1.0 to 0.2514. I n additioii, the diffusion current' constants of these metal ions of x-arious concentrations (0.002 to 0.0004X) in 1-11 thiosulfate were det'erniined. The results are recorded in Table I. I n each case, El was constant and independent of the concentration of metal ions (for a CobaltiIIII. in the form of the hexamgiven thiosulfate concentration), and the plot of Ed,,.US. log $(id - i) was a straight line whose slope corresponded to the appropriate n value. Therefore, these reductions were considered reversible. The values of the dissociation constants and the number of ligands were calculated in the usual \vay ( 7 ) . The values of the half-\yare potentials because of the sloivness of such disof the simple metal ions \Tere taken from placement reactions. Meites ( 2 1 ) with the exception of the The polarography of nickel in thiocuprous-cupric potential, the value of which was taken to be +0.143 yolt, as sulfate has unusual features. A typical g k e n by Onstott and Laitinen ( I S ) . polarogram is shown in Figure 1. As The number of ligands found for the concentration of thiosulfate ion is cadmium agrees with the dialysis nieasdecreased, the height of the first wave urements of Brintzinger and Eckardt (E1 = - 0.6 volt) decreases; and at low (4) made under similar condkions, but concentrations a second wave appears t,hey proposed the formula C U ~ ( S ~ O ~ ) ~ -a -t about -1.1 volts, which is the halffor the cuprous complex. Spacu and wave potential for the simple (aquo) LIurgulescu, however, found indication nickel ion. I n addition, the half-wave of Cu(S203)3-5 by potmtiometric nieaspotential of the first wave decreases a t urenients ( I 6 ) , nhich is in accord 11-ith first (becomes less reducing) and then the present study. KO values of the increases as the thiosulfate concentradissociat'ion constants were reported. tion decreases, attaining a minimum Uske and Levin (19) reported a value of value a t a thiosulfate concentration of 5 x for the dissociation of the about 0.511. This phenomenon is illustrated in Figure 2 . complex C U ( S ~ O ~ )but ~ - ~this : complex 1730
ANALYTICAL CHEMISTRY
This behavior and the depression of the diffusion current beyond about -0.8 volt is very similar to the polarography of nickel in thiocyanate media (18). The shifting of the half-wave potential with thiosulfate concentration can be interpreted using the method of Tanianiuslii and Tanaka (17). Although it is not possible to elucidate theye data in a strictly rigorous
iYi(S203)u(2--2w) +
where u > 2: > zc. The polarographic n-ave which appears a t about -0.6 volt is assumed to result from the reduction of S i ( S z 0 3 ) , ( 2 - 2 u\\-hich ~+ proceeds as follon-s:
-+
+
S l ( S ? O a ) ~ ' ? - - ? ~ )2 e Hg = Ni(Hg) zS203--
+
(1)
Si(Sz03),(*-2")~ is reduced by Equation 1 after dissociation according to Equation 2 :
+
Si(S20d)u(2-2u)+ = Si(S,03),(2-2u)+ (u -
L!!)S203--
(2)
The rate of this dissociation is assumed to be very large compared t o the diffusion rates of the various ions involved.
Likewise, Si(S203),12-2")Tis reduced according to Equation 1 after rapid association as folloii s: Xi(S2O3),(?-?") +
+
(L
- w)S2O3--
=
si(s,o,),(2-?') -
13)
The reduction of Si(S203)> f2--2L) is an activation-controlled reaction as defined by Tanaka (18). The halfwave potential should be independent of the concentration of the nietal complex, and the plot of log i'(id - i) 2's. E should be a straight line whose slope is greater than 0.0296 volt. These conditions were confirmed euperimentally. I n the media used, sodium thiosulfate-potassium nitrate mixtures of constant ionic strength, there apparently was a small kinetic contribution t o the diffusion current a < indicated b y the variation of the quantity id/\%, nhere h is the height of the mercury column. This effect appeared to be minor and was neglected in the interpretation of results, af wac the possible formation of nickel-nitrate complexes. T h e n the electrode process is activation controlled, Tanaka has phon n that the propertips of the imve are independent of the reverse reaction and that a (necessary because actiwtion control is an irreT ersible process) can be evaluated from the plot of log i/(id - i j 2's.E. Equations can then be derix rd to fit tlie particular situations, K h e n the thiosulfate ion concentration is high (1.0 t o 0.6J1), Equation 2 precedes Equation 1. The variation of half-wave potential may be represented b y the following equation: (4)
This equation predicts that the halfwave potential should be shifted to more negative values with increasing concentration of sodium thiosulfate. The slope of the plot of Ell2 V S . log CS,~.-- was found to be 0.038,which agrees \vel1 with the theoretical value of 0.036 when (u- u ) = 1 and when a is the average T-alue found for all the log i/(id- i) us. E plots in this concentration region. The value of CY used was 0.80. The situation changes when the thiosulfate concentration is less than about 0.536. I n this case. Reaction 3 precedes the activation-rontrolled electrode process (Reaction a), and the variation of the half-mve potential ivith concentration of the complexin.; agent is expressed by: 0.0591 w) __ Alog cs2c18-201 The slope of the line corresponding to this concentration region (Figure 2 ) is 0.015 when the average value of CY, 0.76, was used. The value of (v - w) is thus established at 0.39. This together
with the value of 1 for (u- v) indicates that the three nickel complexes in question are Si(Sz03)2--, NiS203,and Ni3(S203)2+L. The exact formula of the last complex is in some doubt, but the data obtained definitely show that the complex in question contains more nickel than thiosulfate ions. Complexes containing sulfate bridges ha\-e been proposed (6). Analogous thiosulfate compounds should be capable a t least of existence, although the possibility of bridge complexes requires that some of the nickel-thiosulfate bonds be made through oxygen. Nickel is not especially known for the ability t o form complexes in which oxygen acts as a donor, but the normal nickel polarographic nai-e n-as suppressed when eodiuni sulfate or sodium sulfite was used as a supporting electrolyte, presumablv because of complex formation. If the formulas of the nickel complexes proposed are compared v-ith those derived bv Tanaka for nickel-thiocyanate complexes ( 1 8 ) . the species that is directly rrduced is. in both cases, a neutral one. The polarogram of nickrl in 131 sodium thiosulfate (Figure 1) is very similar t o the polarogranis recorded by Tanaka for nickel in thiocyanate. This similarity extends even to the dppression of the limiting current fonnd at about -1.0 volt z's. S.C.E. T o establish exactly what electrode reactions were proceeding. a number of controlled potential electrolyses vere performed a t a large mercury pool. under conditions similar to those used for the polarography. When the potential was controlled at the top of the w v e (-0.7 volt), the reaction proceeded as expected-that is, an n value of 1.95 was found, which is in good agreement with the theoretical value of 2.00. When the potential was controlled a t the bottom of the dip (- 1.0 volt), a black precipitate of nickel sulfide was soon observed, and n values on the order of 3 to 6 were found. Precise measurement of n was not possible, because a large background current correction was necessary and because varying amount. of nickel remained insoluble as the sulfide and could not further participate in the reaction. Several electrode reactions are possible. For example :
+ 2 c - + Hg Xi(Hg) + S-- + 2H+ + SOa--
2H20 $- ?;iS20a
constant or increased slightly during the electrolysis. I n addition, no positive test for sulfate was obtained after electrolysis, but a detectable quantity of sulfite ion was found, The follon-ing reaction would account for the steady pH and the production of sulfite ion: Xi(S?O,) -L 4e-
+ Hg F' Si(Hg) + s-- + So$--
(7)
Each four-electron change produced two ions is--. SOa--), each of M hich can prevent one nickel ion from reaching the electrode. Thus an electron change of only four thirds per nickcl ion could be observed. The prevention of the reduction of nickel b y sulfite ions has been observed. Xickel gave no polarographic TI a1 e in 1-11 sodium sulfite solution. There are other possible explanations for the dip. Unfavorable kinetics may he entirely to blame, but this seems lccs likely than the nickel ions being prevented from reacting a t the electrode hv wlfide and sulfite ions. The &pression of the current could a1-o be euplained by the fact that a film of nickcl wlfide might form on the electrode curface and tpnd to prevent furthpr reaction. Film formation has been noticed in come c a w (8).but usually the drop characteristics are affected and the film if visible under a microscope. S o evidence for film formation n-as fold. rnfortunately a complete analysis of the limiting current depression was not pocsible, because major rariations of the factors nliich nould possibly affect the deprescion caused various complications in this region of the polarogram. For instance, L ariation of the nickel concentration changed the extent of the depression a t least partly, because the foot of the irreversible reduction of simple nickel ion extends into this region and is apparently not negligible even in 1M sodium thiosulfate. Polarography of nickrl in thiosulfate media I\ as unsuitable for analytical xork. Only in the case of very low ( < 0.13f) thiosulfate concentrations n-as any constancy of I I C observed. Even in this case, idhad t o be the total diffusion current of both the waves of the coniplex and the simple ion (mcasured at about -1.4 volts z's. S.C.E.).
=
LITERATURE CITED
(6)
Although this reaction would account for the decrease in limiting current, because the sulfide ion produced a t the electrode surface would precipitate nickel ions diffusing toward the electrode and sulfate would complex such ions, it can be discounted because the pH of the solution being electrolyzed remained
(1) Andersen, E. B., Z . physak. Chem. (Lezpzzg) B32, 237 (1936). (2) Bailar, J. C., "The Chemistrv of Coordination Compounds,'' p. 463, Reinhold, S e w York, 1956. ( 3 ) Banegas, R. F., Cnm. Chzle, fac. qiiim. y farm. Tesis Quim.farm.5 , 367
(1953).
(4) Brintxinger, H., Eckardt, \I7.> 2. anorg. u allgem. Chem. 227, 107 (1936).
( 5 ) Denny, T. O., Monk, C. B., Trans. Faraday SOC.47,992 (1951). VOL. 30, NO. 1 1 , NOVEMBER 1958
173 1
(6) Garnes, D. C., Lorenz, G. A,, Montillon, C . H., Trans. Electrochem. SOC. 77, 177 (1940). ( 7 ) Kolthoff, I. M., Lingane, J. J., "Polarography,,, 2nd ed., pp, 211-17, ~ ~ science, S e w Pork, 1952. (8) Ko'thoff, 1. hI.2 JIiller, C. S.2 J . Am. Chem. SOC.6 3 , 2732 (1941). (9) Lingane, J. J., IND.EXG. CHEM., AXAL.ED. 16, 150 (1944).
(10) Lingane, J. J., Meites, L., J . Am. Chem. SOC.73, 2165 (1951). (11). Meites, L., "Polarographic Techniques," Interscience, New York, 1955. bf. s., Shmaeva, T. (12) t M.,KovakovskiI, ~ Ukrain, ~ -Khim. Zhur. 20, 615 (1954). (13) Onstott, E. I., Laitinen, H. A,, J . Am. Chem. SOC.72, 4724 (1950). (14) Page, F, M., Trans. FaradazJ SOC. 50, 120 (1934).
(15) Saxena, R. S., Jain, N . L., J . Indian Chem. SOC.31, 319 (1954).
(16) Spacu, Gs, Murgulescu, J. G.1 Kolloid-2. 91, 294 (1940). (17) TamamuskiJ Tanaka, x'JBzilz' Chem. Japan 23, 110 (18) Tanaka, K.,Ibid., 23, 253 (1950). (19) Uske, E A , Levin, A. T., Zhur. Fzz. R'j
Khzm, 24, 1396 (1963,,
RECEIVEDfor review May 10, 1958. Accepted July 2, 1958.
Electrodes Consisting of Membranes of Precipitates ROBERT B. FISCHER and ROBERT F. BABCOCK' Deparfment o f Chemistry, lndianu University, Bloomingfon, Ind. ,This investigation was undertaken to devise a method of preparing membranes of inorganic precipitates and to evaluate them as potentiometric indicating electrodes. Suitable membranes have been prepared using paraffin as an inert binder. The electrical charge upon the surfaces of the inorganic particles is determined jointly by the precipitate and the electrolyte solution in contact with it. A potential i s developed across the membrane b y transport of oppositely charged, counter ions through the membrane. This type of membrane can be useful as an indicating electrode and also in measuring the surface electrical properties of the material of which each membrane i s composed.
E
consisting of membranes of precipitates are of possible interest in analytical chemistry for tn-o reasons: They may be useful as direct measurement of ion concentration or activity and for titrations involving precipitation reactions, and they may provide information concerning the surface electrical characteristics of the precipitates, thereby contributing fundamental knowledge concerning precipitation processes. Several other types of membranes have been studied and used as indicating electrodes. Foremost is the glass electrode for pH measurement. Some attempts have been reported to use glass of special composition to determine activities of other ions. Considerable research has been conducted on the preparation, theory, and application of electrodes consisting of membranes of collodion and similar materials as described, for example, by Teorell and others (12). An electrical potential apparently develops LECTRODEB
Present address, Technical Services Division, Standard Oil Co. of Indiana, Whiting, Ind.
1732
0
ANALYTICAL CHEMISTRY
across the membrane by means of preferential transfer of cations or anions through the membrane. Collodion membranes can be made selectively permeable to cations or t o anions, as desired, by introduction of negative or positive groups into the collodion structure (6). Clay minerals and synthetic ion exchange resins have also been formed into membranes and used as electrodes, with interesting results in single electrolyte systems but relatively little selectivity in mixed systems. Presumably the resin membranes contain small pores with numerous exchangeable groups along each pore. The pores are in intimate contact, so ions are effectively passed through the membrane, Excellent reviews of ion exchange membrane electrodes have been presented by Gregor ( 4 ) and Spiegler (11). Of particular interest in analytical B-ork is the recent study by Parsons of the use of cationic and anionic exchange resin membrane electrodes (9). H e used the electrodes both for the direct determination of certain ion activities and as indicator electrodes for the titration of sulfate by barium acetate; the presence of other cations and anions caused interference in both types of application. Sinha has also reported on the use of a n ion exchange membrane as a n indicating electrode for the titration of sulfate by barium ion as well as for other titrations (10). Inorganic materials have been investigated by several workers for possible use in membrane electrodes. Buchanan and Heymann used flakes of natural barite, but minute cracks developed and gave rise to simple liquid junction potentials (2). Kolthoff and Sanders prepared membranes of barium sulfate and of silver halide by a fusion technique (8). The silver halide membranes were used as electrodes for the titration of silver nitrate by potassium bromide. Results with
harium sulfate membranes were discouraging, perhaps because of minute cracks such as Buchanan and Heymann observed with their flakes of natural barite. Gregor and Schonhorn prepared membranes of multilayers of barium stearate, which proved to be reversible for passage of barium i'ons even in the presence of a high concentration of sodium ions ( 5 ) . It was suggested that ion transport through such a membrane may occur by exchange of barium ions between adjacent stearates through the membrane. Hirsch-Ayalon prepared a barium sulfate electrode by placing a sheet of cellophane between solutions of barium ion and sulfate ion, thus forming a membrane of precipitated barium sulfate embedded in cellophane ( 7 ) . H e then measured POtentials across the membrane as various electrolytes were added to the solutions on both sides. The purpose of the present research was to investigate methods of preparing membranes consisting of inorganic precipitates and to evaluate them as POtentiometric indicating electrodes. The chemical structure of a collodion or ion exchange membrane determines whether that membrane is preferentially permeable to cations or to anions. Such may not be the case, however, with membranes consisting of precipitates. The sign of the electrical charge fixed upon the surfaces of a precipitate is determined not only b y the precipitate itself but also by the kinds of ions in the electrolyte solutions in contact with it. Several inorganic materials have been used in the present study. Barium sulfate was selected for the most detailed study, primarily because considerable is already known concerning its zeta potential, so that interpretation of data should be less difficult than for materials whose surface electrical properties are not established.
