Coulometric Generation of Hydrogen Ions in Water and Methanol Using Hydrogen-Permeable Membrane Electrodes THOMAS R. BLACKBURN and ROSALIND B. GREENBERG Department of Chemisfry, Wellesley College, Wellesley, Mass. 02 7 8 7
b The usefulness of hydrogen-permeable membrane electrodes for coulometric generation of hydrogen ions is explored through current-voltage curves and coulometric titrations of a primary standard base in water and methanol. Satisfactory currents, present only when the electrode is supplied with hydrogen, are observed in both solvents. The overall titration efficiency of constant-current coulometric titrations is 99.770 in wafer and 99.6% in methanol. No previous coul metric titration of bases in strictly nonaqueous solvents has been reported. An easily fabricated membrane electrode is described.
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Figure 1 . Construction of hydrogenpermeable membrane electrode A: Exploded view. C: Bakelite cap; G: rubber
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1935, Szebelledy and Somogyi ( 8 ) introduced coulometric titration of bases in water solutions via the halfreaction H 2 0+ 02 2H+ 2 e Streuli ( 7 ) generated hydrogen ions in acetonitrile by oxidation of a small amount of water added to the solvent. IF hile this method may, in principle, be extended to many solvents compatible with uater, i t is not suitable for rigorously nonaqueous systems or for solvents easier to oxidize than water. By exploiting the permeability of palladium and palladium-silver menibranes to hydrogen, we have constructed a protori-generating electrode whose functioning does not depend on the solvent. The electrode’s efficiency as a transmitter and oxidizer of hydrogen is greatly improved by surface pretreatment as described below. However, this improvement is fairly short-lived, probably because of loss of catalytic activity for the diqsociation of H2 and for adsorption and desorption of H atoms. Chloride ions and other complexants for Pd(I1) limit the useful range of operating potentials by promoting anodic dissolution of the palladium and loss of current efficiency for proton generation. As an illustration of the analytical potentialities of the electrode, a primary standard base, 2-amino, 2- (hydroxymethyl) - 1,3 - propanediol (6) has been titrated in water and in methanol. N
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gaskets;
P:
palladium or Pd-Ag foil, thickness
0.004 cm., exoosed area 1 cm.*; W(P): platinum wire; W(C): copper wire 6: Assembled electrode. S: Black Glyptal seals
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EXPERIMENTAL
Primary
standard
2-amino,Z-(hy-
dro~ymethyl)-1,3-propanediol(-4HPD) was obtained from Fisher Chemical Co., under the trivial name tris-(hydroxymethyl) aminomethane, or TH.411. R a t e r waq boiled or dpioniwd before use to minimize interference due to CO,. -4 basic ini urity was prewnt in the methanol (%&errs analyzed reagent grade), necessitating pretitration of the cell qolution to a potentiometric end point before the sample mas introduced. Pure palladium and 70% Pd-30% Ag alloy foil was supplied by Baker Platinum Division of Englehard Industries. =1 Beckman Duoqtat power supply was used as a constant-current source. The slow 0.1 to 0.2% drift in the output of this instrument was compensated manually, by monitoring the iR drop across a preciqion resistor with a potentiometer. Potential and pH readings were made with a Corning Model 7 pH meter. Time was registered by a Standard Electric Time Co. Model S-10 stopclock, which was operated simultaneously with the electrolysis current by a DPST relay. Other circuitry was of ordinary design ( 5 ) . Oxygen was removed from the cell by sweeping with prepurified nitrogen to avoid its reaction with the palladiumhydrogen electrode. Electrode Construction. Figure 1 shows t h e construction of t h e elec-
trode used in this study. T h e glass, rubber, and bakelite components of t h e electrode are parts from a n ordinary borosilicate glass condenser. Because of the frequent neccssi-y I r reju\ ellation uf the electrode surface, it is desirable to be a!de to dismantle the electrode and remove the palladium foil easily. The tube was flushed thoroughly with H2 before use, and the hydrogen pressure was maintained at 1-2 p.s.i.g. during use. Xccording to studies by Jewett and Makrides ( 4 ) , the rate of transfer of hydrogen through a palladium membrane is not a strong function of pressure a t pressures a t or above I atm. Choice of Membrane Composition. Early work in this study, and some of the titrations reported in this paper, were carried out using a pure palladium electrode. However, the palladium-silver alloy offers several advantages : The system Pd-H exhibits tlvo crystalline modifications ( a and p ) , with different thermodynamic properties, depending on the H content of the palladium (9). T h i l e the electrode is being used in the dynamic process of hydrogen ion generation, it is not convenient to control the foil’s hydrogen content which, in any case, may have a gradient’ from back to front. Thus the ciirrent-potential behavior of the elect’rode i. poorlv reproducilhle. The 307, Ag-70% Pd alloy’s crystalline form is independent of hydrogen content, and the electrode shows much more reproducible current-voltage curves for hydrogen ion generation than pure pilladium. Constant cycling of the palladiumhydrogen electrode through the a-0 phase transition during use results in warping, enibrittlement and, finally, cracking of the foils. The alloy foils are sturdier than pure palladium a t the outset and do not grow brittle with use, though they do undergo some distortion when they absorb hydrogen. The hydrogen permeability of 30% -4g-7070 I’d is less by a factor of about 5 than that of pure palladium (4). Since our hydrogen ion generation currents correspond to effective permeabilities much lower than those observed by Jewett and Makrides, other factors (probably surface poisoning) limit the overall rate of hydrogen transport. We observed somewhat larger current densities with alloy foils than with pure palladium. VOL. 38, NO. 7, JUNE 1966
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Figure 3. Decay of proton generation at an electrode after initial activation (see text)
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Figure 2. Current-voltage curves for Pd-Ag membrane electrode in aqueous phosphate buffer, pH 7
after activating pre-
treatment 0 After 48 hours exposure to air After 20 minutes exposure to pH 7 phosphate buffer After 30 minutes
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Electrode supplied with hydrogen
0 Hydrogen absent
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E , Volts us. S.C.E. RESULTS
Current-Voltage Curves. Figure 2 shows current-voltage curves for a palladium-silver alloy electrode in an aqueous phosphate buffer with, and without, hydrogen present. T h e electrode had been heated in a muffle furnace t o about 700-900' C. for 2 hours to coat the surface with a n oxide film, and this film reduced electrolytically to produce a finely divided surface (1, 2, 4). Just prior to the currentvoltage curves of Figure 2, the electrode had been further pretreated on both sides by cathodic deposition of a layer of palladium black from a Pd(I1) nitrate-nitric acid solution. A recent publication of Cleary and Greene (2) discusses in a thorough and s:;stematic way the treatment of pure palladium foils to obtain larger and longer-lived
Table I.
AHPD taken, mmoles 0,3084 0,3084 0,2815 0.2815 0,2470 0,2470 Table II.
currents for hydrogen oxidation than those reported here. When freshly treated as described above, the electrode is a relatively efficient transmitter of hydrogen and exhibits reversible behavior with respect to the solution; however, exposure to laboratory air or, mare severely, to solutions of ordinary cleanliness results in a shift of the curve to smaller currents a t a given potential. Note that in Figure 2, the current-voltage curve for the freshly treated electrode passes through the potential axis at the reversible potential without inflection. Figure 3 s h o w the decay of the currentvoltage curve to less reversible behavior with use and the passage of time. I n view of this decay, the oxidation-reduction treatment was given to each new foil, and the palladium black coat-
Coulometric Titrations of AHPD in Water
Supporting electrolyte: 0.5M XaN08 Generating Base current, ma. found, meq. 20 0.3106 20 0.3096 0,2807 25 25 0.2824 20 0.2479 20 0.2481
Rel. error, 70 +0.7
+0.4 -0.3 +0.3 +0.4 +0.4
Coulometric Titrations of AHPD in Methanol
Supporting electrolyte: 0.25M NaNOs AHPD Generating Base taken, mmoles current, ma. found, meq. Rel. error, 70 11 0 1585 0 1576 $0 6 -0 2 0 1440 12 0 1437 12 0 1445 0 1440 +o 4 12 0 1451 0 1440 +o 7 For the titrations in water, the average error is +0.3'7,-, standard deviation = 0.3%; in methanol, the average error is +0.47,, standard deviation = 0.47,. ~
878
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ANALYTICAL CHEMISTRY
ing renewed before each use of the electrode. Figure 4 shows current-voltage curves for a pure palladium electrode with, and without, hydrogen present, in methanol buffered with 0.lM each acetic acid and sodium acetate. Although the proton-generation currents obtained in methanol were generally less than those shown in Figure 2 for water, they are still usefully large. The anodic currents flowing a t the extreme anodic end of the curve in the absence of hydrogen probably result from the oxidation of the solvent to formaldehyde, formic acid, or a mixture. During long electrolysis at this potential, the p H of the electrolyte shows a slow decrease, and the characteristic odor of formic acid is obvious. This reaction is not suitable for titrating bases; in a single experiment, the titration curve for the titration of AHPD was poorly developed with the maximum pH drop occurring at about 50% beyond the theoretical equivalence point, calculated assuming 100% current efficiency for the oxidation of methanol to formic acid. Coulometric Titrations. Using currents chosen on the basis of the observed current-voltage curve to avoid oxidation of the solvent, the current efficiency for hydrogen ion generation in water and methanol was determined by coulometric titration of A H P D . For titrations in water, the auxiliary electrolyte compartment was filled with 0.5M Ag?\'Os SO that the cathode reaction would not interfere with the titration. This precaution was unnecessary for titrations in methanol. A glass electrode-aqueous S. C. E. pair was used for potentiometric end point detection in both water and methanol. We are aware that the apparent pH readings in methanol have no absolute significance, but assumed that the equivalence point coincided
method reported here should not have application for the titration of bases in any electrolytic solvent in which the potentials for the oxidations of hydrogen, palladium, and the solvent bear the same relationship to each other that they do in water and methanol.
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ACKNOWLEDGMENT
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The authors thank A. C. Makrides and D. N. Jewett and their colleagues for helpful discussions. -06
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Figure 4. Current-voltage curves for pure Pd membrane in methanol, buffered with 0.1M each acetic acid-sodium acetate
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Electrode supplied with hydrogen Hydrogen absent
with the maximum rate of change of pH with generation time ( t ) . This latter was located by the method of Gran (3) as the intersection of linear plots of At/ApH for short increments of 2 around the equivalence point. A summary of results is presented in Tables I and 11. The small but significantly positive titration errors shown in Tables I and I1 may be restated as negative deviations from 100% titration efficiency; failure to achieve 1 0 0 ~ ocurrent efficiency for hydrogen ion generation is the most
LITERATURE CITED
+I2
E , Volts us. S.C.E.
likely cause. The small anodic residual currents which flow in the absence of hydrogen (Figures 2 and 4) are of the same order of magnitude as the apparently inefficient portion of the generating current used in titrations. However, since the current-voltage performance of a given electrode changes with time in contact with solutions, there seems no exact way to correct for the residual current. Consequently, we have reported errors based on the total current. There is no a priori reason that the
(1) Blackburn,
T. R., Ph.D. thesis, Harvard University, Cambridge, &lass., 1962. (2) Cleary, H. J., Greene, N. D., Electrochim. Acta 10, 1107 (1965). (3) Gran, G., Acta Chem. Scand. 4 , 559 (1950). (4)Jewett, D. N., llakrides, A. C., Trans. Faraday soc. 6 1 , 932 (1965). (5) Lingane, J. J., “Electroanalytical Chemistry,” 2nd ed., Interscience, New York, 1958. (6) Riddick, J. A,, Ann. X. Y . Acad. Scz. 9 2 , 357 (1961); C . A . 56, 10884d (1962).
(7) Streuli, C. A., ANAL. CHEM.2 8 , 130 (1956). (81 Szebelledy, L., Somogyi, Z., 2. Anal. Chem. 112, 395 (1935). (9) Vasile, AI. J., Enke, C. G., J . Electrochem. SOC.112, 865 (1965). RECEIVEDfor review February 2, 1966.
Accepted April 11,1966. Work supported by National Science Foundation grant GP-3510 and by Wellesley College work scholarship to R. B. G.
Reversible Charge Transfer at the Tubular Platinum E ectrode W. J. BLAEDEL and L. N. KLATT Department of Chemistry, University o f Wisconsin, Madison, Wis.
b
Current-potential equations for the reversible charge transfer at the tubular platinum electrode (TPE) have been theoretically derived and experimentally verified with the ferricyanide-ferrocyanide system. At steady state, the dependences of the concentration profile upon distance along the tube and upon axial linear velocity are shown graphically. For flow rates of a few milliliters per minute in a 1 -cm. X 0.05-cm. diameter tube, the diffusion layer at steady state is comparable to that obtained in a potentiostatic experiment for times less than 1 second. For a given system, the half wave potential at the TPE differs slightly but definitely from that at the dropping mercury electrode.
H
electrochemical systems, where transport of the
YDRODYNAMIC
electroactive material occurs principally by convection, are of two general types: moving electrode systems, such as the rotating platinum wire electrode (14, I 7 ) , the rotating disk electrode (6,8), and the streaming mercury electrode ( 7 ) ; and moving solution systems, containing such electrodes as stationary wire electrodes (9, d 7 ) , a micro-bypass electrode ( d 6 ) , plane electrodes (19), spherical electrodes ( I S , 28), a microconical electrode ( I d ) , and tubular electrodes (4, 5, IO). Much of the work on these systems has been empirical or has been based upon the Nernst diffusion layer concept, which assumes a linear concentration gradient. Only a few cases have been treated rigorously, usually in the diffusion-limited region. Hydrodynamic electrochemical systems have great analytical and practical capabilities; however, a more
extensive theory of the convective transport process and of the associated charge transfer is required to understand the fundamental processes occurring at hydrodynamic electrodes. The following work is concerned with a theoretical description and experimental confirmation of the reversible charge transfer at the tubular platinum electrode (TPE). THEORY
Derivation of Equations. T h e general equation for mass transfer of a chemical species, Ci, in a tube of circular cross section, with radius p and length X,and with a laminar velocity regime is given by Levich (20).
VOL. 38, NO. 7, JUNE 1966
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