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ANALYTICAL CHEMISTRY, VOL. 51, NO. 9, AUGUST 1979
t h e figure are compared t o a straight line with slope 59 mV/decade whereas the bottom line is compared to a theoretical line calculated with Equation 17. T h e lower data shows a n inflection at in agreement with theory for 3Dotlrbo greater than about 6.0 whereas the upper data approximated a straight line in agreement with theory for 3Dot/rb0 less than 6.0. Any comparison between theory and experiment for pseudopolarograms is more approximate than for classical polarography because currents can be estimated with much less precision than polarography. Nevertheless, agreement between experiment and theory is satisfactory and illustrates the deviation from linearity expected for electrodes and pre-electrolysis times that produce elevated amalgam concentrations.
LITERATURE CITED (1) W. D a v b n and M. Whimeld, J. Electronal. Chem. Interfaci3/€ktrochem., 75, 763 (1977).
W. L. Bradford, Limnol. Oceanogr., 18, 757 (1973). R. Ernst, H. E. Allen, and K. H. Mancy, Water Res., 9, 969 (1975). H.Bilinski, R. Huston and W. Stumm, Anal. Chim. Acta, 184, 157 (1976). T. A. O'Shea and K. H.Mancy, Anal. Chem., 48, 1603 (1976). J. Gardiner, in "Proceedings Int. Conf. Heavy Metals in the Environment", Toronto, Canada, 1975, Vol. I, p 227. (7) H. W. Nurnberg, P. Valenta, L. Mart B. Raspor, and L. Sipos, Fresenius' 2. Anal. Chem., 282, 357 (1976). (8) M. Branica, L. Sipos, S. Bubic, and S. Kozar, Nafl. Bur. Stand. ( U . S . ) , Spec. Pub/. 422, 917 (1974). (9) M. Branica, D. M. Novak, and S. Bubic, Croat. Chem. Acta, 49, 539 (1977). (10) A. Zirino and S. P. Kounaves, Ami. Chem., 49, 56 (1977). (11) T . S. Lee, J . Am. Chem. SOC.,74, 5001 (1952). ( 12) P. Debhay, "New Instrumental Methods in elect roc hem is^", Interscience, New York, 1954, pp 217-249. (13) J. W. Ross, R. D. DeMars and I . Shain, Anal. Chem.. 38, 1968 (1956).
(2) (3) (4) (5) (6)
RECEIVED for review January 29, 1979. Accepted April 2 5 , 1979. Work supported by funds from the Oceanographic Section, National Science Foundation, NSF Grant OCE 73-21045
Chemically Treated Tin Oxide Electrodes Responsive to pH and Sulfide H. A.
Laitinen" and T. M. Hseu'
Department of Chemistry, University of Florida, Gainesville, Florida 326 11
Films of antimony-doped tin oxide on glass, after treatment with 10 M NaOH, show Nernstian pH response from pH 2 to 10. The treated electrodes, after exposure to sulfur vapor, show Nernstian response to pS over the range of 6.0 to 15.2, depending upon the total sulfide concentration and pH. At a constant pS, the pH dependence is 59 mV below pH 6.5 and 30 mV from pH 6.5 to 10.65 in accordance with theory.
Electrically conductive, antimony-doped, polycrystalline tin oxide films have been of interest to analytical chemists as optically transparent and electrochemically inert electrodes (1-5). In connection with a study of specific adsorption of halide ions (6, 7) it became clear that the surface reactivity of tin oxide electrodes could be enhanced by pretreatment with caustic. The present study was undertaken to use pH response as a diagnostic measurement for the caustic pretreatment. Previous investigations (8 and references therein) have indicated a response relatively small compared to the theoretical 2.3 RT/F mV/pH, Green (9) considered theoretically two extreme cases of semiconductor-solution interfaces, one in which no surface states exist and the interfacial potential difference lies entirely in the space charge region of the semiconductor, and another in which a high concentration of surface states (donor-acceptor sites) leads to a metal-like behavior. Here the interfacial potential difference lies entirely in the solution phase, independent of carrier charge density. T h e first case would correspond t o nonresponse to solution composition, while the second would correspond to a Nernstian response. Kirkov (8)used a model involving both a space charge region and a Helmholtz solution layer to account for a non-Nernstian selectivity of tin oxide to several Present address, Department of Chemistry, National Taiwan University, Taipei, Taiwan. 0003-2700/79/0351-1550$01.00/0
ions. The present article describes the p H response of tin oxide electrodes before and after alkali treatment. It was also found that a sulfide ion responsive tin oxide could be prepared by exposure of the alkali-treated surface to sulfur vapor. The preparation and response of such electrodes at various p H values and total sulfide concentrations are reported here.
EXPERIMENTAL Preparation of Electrodes. The preparation of tin oxide films has been previously described ( I O , 11). A spray mixture 3.0 M in SnCl,, 0.03 M in SbCl,, and 1.5 M in HC1 was blown onto a Pyrex substrate heated on a hot plate, until a film thickness of ca. 0.5 pm was shown by interference fringes. The film was polished successively with 0.3- and 0.05-pm polishing alumina, washed successively with water, 1 M HNO,, isopropanol, and water. Alkali pretreatment consisted of soaking the freshly prepared electrodes with 10 M NaOH solution for 45 h at 50 O C , washing with several portions of distilled water and soaking in distilled water before use. The surface conductivity, as measured by the 4-point probe method ( I I ) was found to increase from 815 to 1530 Q-' cm-' upon pretreatment with NaOH. For conditioning to achieve sulfide response, a simple soaking of tin oxide electrodes overnight in 0.1 M Na2S led to a partial response, as evidenced by poorly reproducible potentials of less than Nernstian slope when plotted against sulfide concentration. Sulfur pretreatment, in contrast, led to theoretical response, as detailed below. For sulfur pretreatment, a 150-mL beaker containing solid Na2S4(96% purity, mp = 275 "C, Alfa products) was covered with a watch glass and heated gently to 275 O C in a glass mantle heater. At the bottom of the beaker, a dark colored, high viscosity melt layer formed, covered with a solid layer of Na2S4. The melt layer and solid phase were somewhat separated by a vapor phase of sulfur being continuously evolved. The tin oxide-coated glass, pretreated with concentrated NaOH solution, was placed on the solid Na2S4phase and exposed for 10 min to the sulfur vapor above it. Long exposure, or more severe conditions, led to the complete removal of the tin oxide film. The treated electrodes were washed in a stream of water, rubbing gently with paper towels, and usually were soaked in a dilute (ca. M) Na2S solution for an initial 'C 1979 American Chemical Society
ANALYTICAL CHEMISTRY, VOL. 51, NO. 9, AUGUST 1979
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O t
.,@@t I -2001
J. ----- .L 0
Time in min 2
6
4
8
1
0
1
2
1
4
PH
Figure 1. Cell emf as a function of pH. (A) Treated tin oxide electrode
(6)Untreated tin oxide electrode
’,
conditioning. The surface conductivity was 1687 V’ cm about the same as for the hydroxide-treated surface and about twice that of the untreated surface. A series of buffer solutions was prepared according to Clark and Lubs (12),and measured against a saturated calomel reference electrode using a Beckman Model 4500 digital pH meter. The tin oxide electrode measurements and the pH values of the buffer solutions were checked immediately before use with a Coleman Model 38 A expanded scale pH meter. A rubber insulated copper lead was attached to the tin oxide electrode with a mixture of copper paint (GC Walsco 37-02) and E-pox E 5 glue. All measurements were carried out in an air thermostat at 25 + 0.5 “C. Electromotive force vs. time curves for the dynamic response measurements were obtained with an Omni Scribe Series B-5000 recorder (Houston Instrument Co.). To examine sulfide ion response, similar measurements were carried out in a series of Clark and Lubs buffer solutions to which NazS had been added. To avoid absorption of atmospheric oxygen and carbon dioxide, argon gas was bubbled through the test solution. Electrode potentials stabilized rapidly and, in every case, equilibrium potentials were measured within 3 min after immersion of the electrodes in the test solutions. Stock solutions of NazS were prepared by dissolving reagent grade NazS.9 HzO in water, and standardized by using potassium iodate as an oxidant in alkaline solution (13). A series of stock solutions of NazS was prepared keeping the ionic strength constant at 0.1 M by addition of an appropriate volume of 1 M NaN03. pH readings were converted to hydrogen ion concentration using the activity coefficient of 0.76 for univalent ions (14) at an ionic strength of 0.1 M.
RESULTS AND DISCUSSION Potential-pH Relationship. The response of a tin oxide electrode treated with 10 M NaOH (curve A) and an untreated electrode (curve B) are shown in Figure 1. The experimental values of the slopes are 59.0 and 37 mV/pH respectively. The empirical equation for the response of the treated electrode for the pH range 2 to 10 is given by E = 0.495 - 0.0590 pH a t 25 “ C (1) when referred to the saturated calomel electrode. T h e “intercept potential” term in Equation 1 was found to depend upon the conditions of storing the electrode. When the electrode was stored in distilled water, the reproducibility of the measured potentials was within &3 mV in the pH range of 2 to 10. When the electrode was stored in 0.1 M NaOH solution, the slope remained unchanged but the potential increased by 20 mV over a 30-min period. Decreasing the pH from 13 to 1 using the buffer solutions yielded a line parallel t o the original calibration line but deviating by 20 mV over
Dynamic response characteristics of treated tin oxide electrode at extreme pH values
Figure 2.
the pH interval 2 to 10. The original response could be restored by washing the electrode and soaking in distilled water for about 20 min. The dynamic response characteristics were peculiar a t extreme p H values. In the pH region of 2 to 10, the pH-time curves were all smooth, reaching a stable reading in about 3 min. The response curves a t extreme p H values showed an “overshoot” phenomenon (Figure 2 ) , followed by a slow drift of emf. This behavior is presumably to be attributed to the amphoteric properties of tin oxide, which lead to slow surface acid-base reactions. Application of the treated tin oxide electrode to the potentiometric titration of 0.1 M HCl revealed that although there were minor differences between the curves obtained with tin oxide and glass electrodes a t extreme p H values, the same end point was observed in each case. Sulfide Ion Response. Sodium sulfide concentrations varying between 2.48 X lo-’ and 4.96 X lo4 M, a t an ionic strength of 0.1 M were used in test solutions, measured against the SCE with a 0.1 M N a N 0 3 salt bridge. pH readings were made with a glass electrode to permit calculation of the sulfide ion concentration using the equation
where [S2-Itis the total concentration of sulfide present ([S2-] [HS-] + [H,S]). The thermodynamic ionization constants K 2 = 1.20 X (15) for H2S, K , = 9.23 X can be converted to the ionization constants, K1 and K z , in terms of concentrations a t an ionic strength of 0.1 M by using the extended Debye-Huckel equation (14) Le.,
+
p K , = pK1’-
4 ~
2 f i and pKz = pK2’ ___
l+&
1+fi
the calculated value of K1 is 1.61 X lo-’ (pK1 = 6.80) and K 2 is 3.64 x (pK2 = 14.44) a t ionic strength of 0.1 M, respectively. The plotted data, over range of pS (-log [S2-]) from 7.0 ([$-I, = 8.27 X M, p H = 10.65) to 15.3 ([S2-],= 8.27 X lo4 M, pH = 5.65), follows a straight line. The response curve is Nernstian with a slope of 30.0 mV per concentration decade a t 25 “C and is within experimental error of the theoretical value given by Equation 3.
E = constant
- 0.296 log [S2-]
(3)
where the “constant” is the sum of the potentials of the saturated calomel electrode, the standard potential of the
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ANALYTICAL CHEMISTRY, VOL. 51, NO. 9, AUGUST 1979
-
8 1
0
3
1
5
7
9
1
1
1
PH
Figure 3. Cell emf as a function of pH and sodlum sulflde concentration. (A) [S2-], = 8.27 X M (B)[S2-], = 8.27 X M. (C) [S2-], = 8.27 x 10-5 M
treated tin oxide electrode, the liquid-junction potential between the test solution and the reference, and the activitity coefficient term of the sulfide ion (at 1.1 = 0.1). Since the measurements were carried out a t a constant ionic strength of 0.1 M, the junction potential was kept approximately constant. The intercept obtained from the extrapolation of this line on the potential axis is -0.525 V. In freshly prepared sodium sulfide solutions, Equation 3 is valid between a pS of 7.0 ( [ P I , = 8.27 X M, pH = 10.65) and 15.3 ([S2-], = 8.27 X lo4 M, p H = 5.66). The low p H of the latter can be attributed to C 0 2 absorption. The deviations of measured emf values from the straight line a t pS < 7 can be attributed to a change of the electrode surface composition under conditions of high OH- and S2-concentration. The more important practical situation is one in which the p H is determined by a buffer solution, and the sulfide ion concentration is related to total sulfide by Equation 2. At constant total sulfide concentration and constant ionic strength, the p H dependence of emf would have two limiting slopes, as shown by substitution of [S2-]from Equation 2 into Equation 3, namely 0.0296 V per pH unit in alkaline solutions where [H+]/k, >> 1 + ([H+]*/K1K2)and 0.0592 V per pH unit in acidic solution where [H+l2/K1K2>> 1 + ([H+]/K,). This behavior is experimentally verified in Figure 3 for three concentrations of sodium sulfide in various buffer solutions a t a constant ionic strength of 0.1 M. The plots of emf vs. pS shown in Figure 4 indicate a Nernstian response from pS between 7.0 to 16.2 in the pH range 5.2-10.6. The time response curves showed a smooth approach to a stable potential within 2' min in stirred solution or 3 min in unstirred solutions. M) sulfide The electrode, w h m stored in dilute (ca. solution and washed several times before each use, showed the same sensitivity after three months of use. In potentiometric titrations of AgNO, with NazS solution and the reverse, a large potential charge of about 600 mV occurred, with an end point a t 99.77% of the theoretical volume. The curves were quite similar to those observed with the silver sulfide membrane electrode (16-19). The sulfide electrode, when exposed to aqueous AgN03, behaved reversibly to silver ion concentrations from lo-' to M a t an ionic strength of 0.1 M, with a slope of 59.0 mV per log [Ag+] unit and an extrapolated intercept of 0.555 V vs. SCE. The extrapolated formal potential may be compared with a value calculated from EoAgtlAg = 0.7991 V and EsCE= 0.2415 V, using an activity coefficient of 0.76 for the silver ion, to be 0.5506 V. Evidently a layer of Ag,S is formed when the sulfide-sensitized electrode is treated with aqueous AgNO,. When Na2S was titrated with Pb(NO3I2solution, a large potential jump (from ca. 425 to ca. 575 mV) occurred after the first addition of lead ion, presumably due to the formation
-
4
6
8
IO
12
14
16
I8
PS
Figure 4. Cell emf as a function of pS at constant sodium sulfide concentrations and different (controlled)pH values. (A) [S2-], = 8.27 X M. (6)[S2-]! = 8.27 X M. (C) [S2-], = 8.27 X M. On each curve the linear region extends from point X (pH = 10.65) to point Y (pH = 5.16)
of PbS on the electrode surface. A large potential in the opposite direction occurred in the immediate vicinity of the equivalence point. T h e end point, taken as the point of maximal rate of change of potential, corresponded to 99.5% of the theoretical equivalence point.
CONCLUSIONS The pH response in the intermediate p H range may be useful as a diagnostic for the full development of surface sites for proton exchange. Nernstian response is to be expected only for tin oxide samples of sufficiently high charge carrier concentrations that the space charge region makes a negligible contribution to the total potential difference a t the interface. Sulfide-treated electrodes may prove useful as substitutes for silver sulfide membrane electrodes. LITERATURE CITED Cooper, W. Nafure (London) 1962, 194, 569. Kuwana, T.; Darlington, R. K.; Leedy, D.W. Anal. Chem. 1964, 36,2023. Propst. R. C. Anal. Chem. 1969, 4 1 , 910. Laitinen, H. A,; Watkins, N. H. Anal. Chem. 1975, 4 7 , 1352. Laitinen, H. A,; Conky, J. M. Anal. Chem. 1976, 4 8 , 1224. Yoneyama, H.; Laitinen, H. A. J . Electroanal. Chem. 1977, 75,647; 1977, 7 9 , 129. (7) Uchida, I.; Niki, K.; Laitinen, H. A. J . Ekctrochem. Soc.1976, 725,1759. (8) Kirkov, P. "Electrochemical Contributions to Environmental Protection"; Beck, T. R., Enke, C. G., Cecil, 0. B., McCallum, J., Wlodek, S. T., Eds.; The Electrochemical Society, Inc.: Princeton, N.J., 1972; p 57. (9) Green, M. J . Chem. Phys. 1959, 31,200. (IO) Mochel, J. M. U.S. Patent 2564707, August 21, 1951. (11) Kim, H.; Laitinen, H. A. J . Am. Ceram. Soc. 1975, 58,23. (12) Bower, V. E.; Bates, R. G. J . Res. Nati. Bur. Stand. 1955, 55, 197. (13) Bethge, P. 0. Anal. Chlm. Acta 1954, 10, 310. (14) Laitinen. H. A,; Harris, W. E. "Chemical Analysis", 2nd ed.;McGraw Hill: New York, 1975. (15) Koithoff, I . M. J . Phys. Chem. 1931, 35,2711. (16) Hseu, T. M.; Rechnitz, G. A. Anal. Chem. 1968, 4 0 , 1054. (17) Natarajan, K. A.; Iwasaki, I. J . Electrochem. Soc. Zndia 1974, 23,201. (18)Crombie, D. J.; Moody, G. F.; Thomas, J. D. R. Anal. Chim. Acta 1975, 80,1. (19) Kim. G.0.; Li, A. K.; Kim, S. A.; 'fan, S. D. Punsok Hwahak 1977, 75, 41. (1) (2) (3) (4) (5) (6)
RECEDTD for review January 15,1979. Accepted May 17,1979. The financial support of NSF grant 2CHE-75-08654 A02 is gratefully acknowledged. One of the authors (T.M.H.) also acknowledges a National Science Council, Republic of China, fellowship grant.
