ing the value by successive approximation using Equation 18 or 19, The values of Kl and K2 calculated in this manner are given in Table 11. The mean values of Kl and K2 are found to be log Kl = 2.90 f 0.03 and log K2 = 0.77 i 0.05 a t the ionic strength of 1.OM. The marked consistency in the values of K 1 and K2 over a wide range of acidities can be taken as providing strong support to the understanding that the fluoride-selective membrane electrode senses selectively the free fluoride ions even in acidic solutions. Calculations were made to test whether the electrode did show any response at all to the HF2- ion. In this connection, use was made of the equation
E
=
Eo
-E In { [F-1 + K[HFz-]), F
(20)
to calculate the selectivity constant, K , by successive approximations. The results did not show any improvement in the agreement among the values of K1 and K2. On the contrary, there was a greater divergence in the values of K2. I t is, therefore, concluded that the response of the fluoride electrode, if any, to the HF2- ion is negligible. RECEIVED for review October 19, 1967. Accepted December 6, 1967. The financial support of Grants NIH GM-14544 and NSF GP-6485 is gratefully acknowledged.
Catalytic Reactions at Tubular Electrodes L. N. Klattl and W. J. Blaedel Department of Chemistry, University of Wisconsin, Madison, Wis. 53706
Current-flow rate equations for the mass transferlimited catalytic regeneration of the reactant at a tubular electrode have been theoretically derived and experimentally verified. The dependence of the concentration profile upon axial distance, flow rate, and rate constant of the chemical reaction are shown graphically. This steady-state hydrodynamic electrochemical system may be used to study reactions with pseudofirst-order rate constants greater than 0.1 sec-’.
Two MAIN ADVANTAGES are associated with hydrodynamic electrochemical systems. First, the forced controlled convection greatly increases the rate of mass transfer, producing a proportionate increase in the current and in the sensitivity when applied to analytical measurements. Second, measurements in these systems may be made under time-independent steady-state conditions, as opposed to the transient, time-dependent measurements that are required of the more conventional techniques. The need for high speed electronics in the readout circuits is diminished. Also, the steady-state nature of the hydrodynamic systems eliminates the charging current, which often limits the sensitivity of the time-dependent techniques. Theoretical description of electrochemical systems applied to the study of chemical kinetics has employed time-dependent techniques almost exclusively. Kinetic studies in hydrodynamic systems have been rare, partly because of success of the time-dependent techniques, and perhaps difficulties associated with the mathematical treatment and experimental measurements in hydrodynamic systems. Galus and Adams (1) considered by means of the reaction layer concept (2, 3) an irreversible first-order chemical reaction succeeding a reversible charge transfer reaction occurring a t the rotating disk electrode. Levich ( 4 ) treated the preceding and catalytic 1 Present address, Chemistry Department, Southern Illinois University, Carbondale, Ill. 62901
(1) Z. Galus and R. N. Adams, J. Electroanal. Chem., 4,248 (1962). (2) K. Wiesner, 2.E/ekrrochem., 49, 164 (1943). (3) R. Brdicka and K. Wiesner, Collection Czech. Chem. Commun., 12, 138 (1947). (4) V. G. Levich, “Physicochemical Hydrodynamics,” PrenticeHall, Englewood Cliffs, N. J., 1962. 5 12
ANALYTICAL CHEMISTRY
chemical reactions a t the rotating disk, neglecting the convective terms in the mass transfer problem by assuming very large rates for the chemical reactions. Albery ( 5 ) and Albery and Bruckenstein (6-9) applied the rotating ring-disk electrode to an irreversible first-order succeeding chemical reaction. With the recently improved techniques for continuous measurements in flowing solutions, and with an improved understanding of the convective mass transfer-charge transfer interaction, the advantages of making electrochemical kinetic studies in hydrodynamic systems are now more realizable. The following work gives a theoretical description and experimental confirmation of the catalytic mechanism occurring a t a tubular electrode. THEORY
Derivation of Equations. In the discussion that follows the process is considered as a reduction; however, the extension to an oxidation is obvious. The catalytic mechanism involves the heterogeneous reduction of an oxidized species 0 to a reduced species R , which in turn reacts homogeneously with an electroinactive species Z , present in the bulk of the solution, regenerating substance 0. This mechanism may be depicted by Reactions 1 and 2.
In order to treat this mechanism occurring a t a tubular electrode of circular cross section, with radius p and length X,and with a laminar flow regime, the previously considered equations of convective mass transfer ( 4 , IO, 11) must be modified to accommodate the coupled chemical reaction. This modification involves the addition of kinetic terms expressing the
(5) W. J. Albery, Trans. Faraday SOC.,62, 1915 (1966). (6) W. J. Albery and S. Bruckenstein, Ibid., 62, 1920 (1966). (7) Ibid., p. 1932. ( 8 ) Ibid., p. 1938. (9) lbid., p. 1946. 38, 879 (1966). (10) W. J. Blaedel and L. N. Klatt, ANAL.CHEM., (11) L. N. Klatt and W. J. Blaedel, Ibid., 39, 1065 (1967).
rate of the chemical reaction as substances Oand R are transported axially along the tube, as given by Equation 3. (3) va is the axial linear velocity. The boundary conditions are:
y=o: :
y+
x = 0:
co=o co+co* co= c o *
(4) (5) (6)
The tube parameters and coordinate system (origin centered at the tube wall at the entrance; axial distance, x ; distance normal to the wall, y ) are more fully defined by Blaedel and Klatt (IO). In these equations C,, CR,and C z are the molar concentrations of substances 0, R, and Z, respectively. C,* and Do are the bulk concentration and diffusion coefficient of substance 0, respectively, and k’ is the bimolecular rate constant corresponding to Reaction 2. The ability to study the chemical reaction in the mass transfer-limited region of the current-potential curve of the electroactive species (Equation 4) has eliminated the problem posed by the nonuniform potential distribution at the electrode surface previously encountered ( I I ) . Equation 3 is a second-order differential equation with three dependent variables, C,,CR,and Cz. A complete mathematical description of Reactions 1 and 2 would require solving simultaneously Equation 3 and analogous ones for substances R and Z. Because of the enormous mathematical difficulties encountered in this approach without a substantial increase in information concerning Reactions 1 and 2, the following simplifying assumptions are made. If Cz >> CR,then Cz is independent of x and y , and Reaction 2 becomes pseudo-first-order in CR,with a rate constant k = k‘Cz
(7)
Also, providing that the simultaneous occurrence of the chemical reaction does not change the basic mode of mass transfer, the relationship between CRand Co may be given by Equation 8.
CR = ( D o / D ~ ) ? ’ ~ ( co *CO)
6, is the Laplace transform of C,, s is the Laplace parameter, and Ai(g) is the Airy function (13). To obtain the solution, it was assumed that Do = DR. The flux of substance 0 to the wall of the tube, obtained by differentiating Equation 14 with respect to y and evaluating at y = 0, is given by
KO(W )are modified Bessel functions of the second kind of fractional order (13). The general inverse Laplace transform of Equation 15 is not known; however, special cases may be obtained readily. When 1
(“)3’2>>
Do
30,s Equation 15 simplifies to
CO*fi
Do ($)v=o
=
S
Obtaining the inverse Laplace transform of Equation 17, and integrating the flux over the surface of the tube, the current, i, is given by i = nF 2 r p X C , d z k
( 1 8)
This result is applicable for k > 1 sec-I. A second special case of Equation 15 may be obtained when
For this case Equation 15 reduces to
(8)
Defining u =
c,*
- c,
(9)
and applying appropriate rules of differentiation to Equation 9, Equations 7 to 9 and Equations 3 to 6 may be combined to give
r(a) is the gamma function (13). Proceeding as in the previous case, the mass transfer-limited current, id, is obtained (14). id = 5.31
x
105~D~2’3X2’3V1~3Co*
(21)
V is the volume flow rate. For laminar flow in a tube, the volume flow rate is given by
v = rp20a/2 (22) A result applicable for k < 1 sec-lmay be obtained by expand-
with boundary conditions y = 0 : u = c,*
(11)
y + a :
u+o
( 12 )
x=o:
u = o
( 13 )
The boundary value problem (Equations 10 to 13) describing Reactions 1 and 2 has been reduced to one with a single dependent variable, and may be solved by means of the Laplace transformation operator (12) to give (12) R. V. Churchill, “Operational Mathematics,” 2nd ed., McGraw-Hill, New York, 1958.
ing the modified Bessel functions in terms of their appropriate infinite power series (13) and dividing these power series. The result in terms of the current is given by the two-term expansion
i = id
+ 2.50 X 10~nD,~‘3p2X4~~Co*k/V~’3 (23)
(13) M. Abramowitz and I. A. Stegun, Eds., “Handbook of Mathematical Functions,” National Bureau of Standards, Washington, D. C., 1964. (14) W. J. Blaedel, C. L. Olson, and L. R. Sharma, ANAL.CHEM., 35, 2100 (1963). VOL. 40, NO. 3, MARCH 1968
513
W
J
0
I
DISTANCE FROM ELECTRODE SURFACE, CM.X 1000
5
I
DISTANCE FROM ELECTRODE SURFACE, CM. X IO00
Figure 1. Dependence of concentration profile upon axial distance
Figure 2. Dependence of concentration profile upon Row rate
co = 10.0 cm/sec, Do = 5.0 X P = 0.05cm
x = 0.5 cm, Do = 5.0 X 10” sq cm/sec, k = l.O/sec, p = 0.05 cm
10-8 sq cm/sec, k
=
l.O/sec,
Equation 21 defines id. The error resulting from termination of this series after two terms increases with k, approaching 6 as k approaches unity. Concentration Profiles. The general inverse Laplace transform of Equation 14 is not known. The dependence of the concentration of substance 0 upon the independent variables x and y , the experimentally controllable axial flow rate, ,.v and the rate constant of Reaction 2 are shown in Figures 1, 2, and 3, respectively. These curves were obtained by numerically solving the boundary value problem (Equations 10 to 13) through use of a finite difference algorithm (15). Large Rate Constants. Equation 18 may also be obtained from the boundary value problem by setting the left side of Equation 10 equal to zero. This condition may be achieved by allowing either the flow rate or du/bx to equal zero. The former condition is readily achieved. The latter condition is achieved when the rate of the chemical reaction becomes large enough to maintain the concentration of substance 0 unchanged with respect to the axial distance, x . The independence between current and flow rate predicted by Equation 18 may be explained by considering the interaction between the mass transport and the chemical reaction. The laminar flow regime has velocity vectors in the axial direction only, with molecular diffusion being the only means of mass transport normal to these axial vectors. The region of changing concentration is much smaller when Reaction 2 is occurring than it would be when Reaction 2 is not occurring (Figure 3) and in this case, an increase in the flow rate does not significantly increase the flux of substance 0 to the electrode surface. When the chemical reaction is fast, the flux of substance 0 to the electrode surface is contributed to principally by the regenerating chemical reaction, and only negligibly by diffusion. This result may be achieved in principle for any rate constant by simply increasing the electrode length-Le., decreasing s in Equation 15. A result identical with Equation 18 is obtained in the time(15) E. L. Stiefel, “Introduction to Numerical Mathematics,” Academic Press, New York, 1963.
5 14
ANALYTICAL CHEMISTRY
dependent transient electrochemical systems when the rate of the chemical reaction maintains the concentration distribution independent of time, giving a steady-state current. Small Rate Constants. Equation 23 ptedicts that for small rate constants, increasing the flow rate decreases the kinetic portion of the current. This is intuitively correct: For small rate constants, increasing the flow rate gives the chemical reaction less opportunity to occur during passage through the electrode, and thereby less opportunity to contribute to the kinetic current. EXPERIMENTAL
Apparatus and Procedure. The basic apparatus and experimental procedure have been described (11). The tubular gold electrode, 0.930 cm long and with a radius of 0.0508 cm, was cut from seamless gold tubing, and assembled using the procedure for the tubular platinum electrode described previously (11). The mercury film tubular platinum electrode (MTPE) was prepared according to the procedure of Oesterling and Olson (16). All work was done at ambient room temperature, 25” f 2”
c.
All current data were corrected for background currents. Materials. The ferric and ferrous solutions were prepared from reagent grade hydrated salts, and were checked for the respective iron content by the conventional dichromate titration. A stock solution of 0.005F Fe2(S0& in 0.25F H2SO4was used to prepare all ferric solutions. A solution of 0.01F FeS04 in 0.25F H S 0 4 prepared before each experiment was used to prepare all ferrous solutions. Hydrogen peroxide solutions were prepared by appropriate dilution of Superoxol (hydrogen peroxide 30 %, Merck and Co., Rahway, N. J,), and were analyzed for H202 content by permanganate titration (17) immediately after completion of an experiment. (16) T. 0. Oesterling and C. L. Olson, ANAL. CHEM., 39, 1543 (1967). (17) C. E. Huckaba and F. G . Keyes, J. Am. Chem. Soc., 70, 1640 (1948).
The hydroxylamine sulfate was twice recrystallized from water and dried at room temperature under a vacuum. Assays based upon its redox and acidic properties indicated purities of 99.11 and 99.15%, respectively, as (NH20H)2HzS04. The potassium titanium oxalate was twice recrystallized from water and dried at room temperature under a vacuum. Assays based upon oxalate and titanium contents indicated purities of 99.97 and 99.70 %, respectively, as KzTiO(C204)2. 2 HzO. All other chemicals were of reagent grade quality and were used without further purification.
I .o
*O
\v 0.5 0”
RESULTS AND DISCUSSION
Choice of Model System. For Reaction 2 to proceed, substance Z must be a stronger oxidizing agent than substance 0. Thus, in order to observe the catalytic mechanism the heterogeneous electrolytic reduction of substance Z must proceed with a large overvoltage. A chemical reaction suitable for experimental verification of the catalytic mechanism must not only meet this requirement, but it must also be first-order with respect to substance R, and it must have a sufficiently large rate constant (k‘ > O.l/mole/sec). A number of systems following the catalytic mechanism of Reactions 1 and 2 have been reported in the electrochemical literature, and it was supposed early in the study that selection of a demonstration system would be relatively straightforward. However, in each case investigated, the kinetic system either failed to meet the requirements stated above or was more complicated than the proposed theoretical model. Ti(IV)-NH*OH System. The reduction at a mercury electrode of Ti(1V) to Ti(II1) followed by its chemical reoxidation in the presence of hydroxylamine has been studied by numerous workers and is well represented by Reactions 1 and 2 (18-25). However, at the platinum electrode Ti(1V) was not reducible while hydroxylamine was. Ti(1V) was not reducible at the tubular gold electrode either. At the MTPE hydroxylamine was not reducible, while Ti(1V) gave a limiting current which was only about 13% of its theoretical value. The inability to obtain mass transfer-limited currents for the reduction of Ti(1V) eliminated systems in which chlorate instead of hydroxylamine serves as the chemical reoxidant (26, 27). Fe(III)-NHZOH System. The alkaline reduction at a mercury electrode of the triethanolamine complex of Fe(II1) followed by its chemical reoxidation in the presence of hydroxylamine (28) occurs beyond the cathodic potential range of the platinum, gold, or MTPE electrodes. Furthermore, the alkaline reduction is complicated by a succeeding chemical reaction that is coupled with the charge transfer reaction (29). (18) A. Blazek and J. Koryta, Collection Czech. Chem. Commun., 18, 326 (1953). (19) P. Delahay, C. C. Mattax, and T. Berzins, J . Am. Chem. SOC., 76, 5319 (1954). (20) 0.Fischer, 0. Dracka, and E. Fischerova, Collection Czech. Chem. Commun., 26, 1505 (1961). (21) H. B. Herman and A. J. Bard, ANAL.CHEM., 36, 510 (1964). (22) J. H. Christie and G. Lauer, Zbid., 36, 2037 (1964). (23) J. M. Saveant and E. Vianello, Electrochim. Acta, 10, 905 (1965). (24) P. J. Lingane and H. H. Christie, J. Electroanal. Chem., 13, 227 (1967). (25) C.V. Evins and S. P. Perone, ANAL,CHEM., 39, 309 (1967). (26) J. Koryta and J. Tenygl, Collection Czech. Chem. Commun., 19,839 (1954). (27) D. E. Smith, ANAL.CHEM., 35, 610 (1963). (28) D. S. Polycyn and I. Shain, Zbid., 38, 376 (1966). (29) A. A. Vleck, “Progress in Inorganic:Chemistry,” F. A. Cotton, Ed., Vol. 5, Interscience, New York, 1963.
I
5 DISTANCE FROM ELECTRODE SURFACE, CM. X 1000 Figure 3. Dependence of concentration profile upon the rate constant x = P =
0.5 cm, ca
=
10.0 cm/sec, Do
=
5.0 X
sq cm/sec,
0.05 cm
Fe(III)-H202 System. The reduction at a mercury electrode of Fe(II1) to Fe(I1) followed by its chemical reoxidation by hydrogen peroxide in acid medium was thoroughly studied by Kolthoff and Parry (30),and evaluated in terms of catalytic currents at a dropping mercury electrode by several workers (31-33). A variation of this system using the ethylenediaminetetraacetic acid complex of Fe(II1) has also been reported (34, 35). The Fe(III)-H202 system could not be used to verify the catalytic mechanism equations with a tubular platinum electrode because of the spontaneous decomposition of H20zon platinum (36, 37). Exploratory studies in acid medium at the tubular gold electrode showed that H202 did not undergo spontaneous decomposition, but that it was reduced electrolytically at potentials corresponding to the mass transferlimited reduction of Fe(II1). A somewhat potential-dependent, flow rate-independent reduction current of 0.5 to 2 pA was observed for 0.001FHz02. Further studies on the reduction of Fe(II1) in 0.25F H2S04 gave rather nonreproducible mass transfer-limited currents. At the MTPE, the behavior of hydrogen peroxide was about the same at the tubular gold electrode. The reduction of Fe(II1) in 0.25F H2S04 gave mass transfer-limited currents that were reproducible and constant within 1 at a particular flow rate. In spite of the possible complications caused by the slight heterogeneous reduction of hydrogen peroxide at the MTPE, this system was chosen for experimental verification of the theoretical equations. (30) I. M. Kolthoff and E. P. Parry, J . Am. Chem. Soc., 73, 3718 (1951). (31) P. Delahay and G. L. Stiehl, Zbid., 74, 3500 (1952). (32) S. L. Miller, Zbid., p 4130. (33) Z . Pospisil, Collection Czech. Chem. Commun., 18, 337 (1953). (34) B. Matyska and D. Duskova, Zbid., 22, 1747 (1957). (35) B. Matyska, Zbid., 22, 1758 (1957). (36) G. Bianchi, F. Mazza, and T. Mussini, Electrochim. Acta, 7, 457 (1962). (37) J. E. Harrar, ANAL.CHEM., 35,893 (1963). VOL. 40, NO. 3, MARCH 1968
515
Table I. Flow Rate Study for k
Curve A B C
-I
< 1 Sec-'"
LL
c
Slope and probable error
Intercept and probable error
7 . 4 i0.1
124 i 1
5.1 i 0.4
127 i 2
0.45
45
+1
0.26
43 43
2.9
* 0.2
132 Average 127
k,
k I,
sec-1
0.66
M-lsec-1 42
P xo
u
Data from Figure 4.
0
1 Table 11. Flow Rate Study for k 3 1 Sec-lu V, ml/min
i, PA
11.3 7.58 6.58 4.95 4.29
47.8 46.0 44.5 44.5 44.6
a Conditions. As in Table I, but with 43.7 mF H202and 0.5mF Fe(II1).
Fe(III)-H202 System for k < 1. Equations 23 and 21 predict that for k < 1 sec-l a plot of i/V113C0* vs. l/Vz'3 should be linear, with the slope a function of the rate constant and an intercept related to the diffusion constant of Fe(II1). In Figure 4 such plots are shown for three solutions of different compositions. The good linearity of these plots supports the correctness of Equations 23 and 21. Data from the plots of Figure 4 are analyzed quantitatively in Table I. The intercept is independent of composition, as required by Equations 23 and 21. Further, the magnitude of the intercept permits calculation of Doas 4 X sq cm/sec in 0.25F H2SOa, which compares reasonably well with 4.2 X lop6 sq cmjsec in sq cmjsec in 0.25F 1F HSO, at 25" C (11) and 7.3 X HzS04at 30" C (31),lending additional support to the correctness of Equations 23 and 21. Calculation of the bimolecular rate constant from the slopes gives a constant value of 43 However, literature values of the bimolecular rate M-'sec-'. constant are 7OM-l sec-'(in 0.25F H2S04), (30), 61M-l sec-l (in 1F H2S04)(38), and 62M-I sec-' (in 0.5F HzSOI) (33) at 25" C, considerably different from the calculated value. This discrepancy is explained in the next section. Fe(III)-H202 System for k > 1. Equation 18 predicts that for k > 1 sec-l the current should be independent of the flow rate. In Table 11, currents are listed for various flow rates for a system with a considerably higher pseudo-first-order rate constant than the systems of Table I (achieved by using a higher H202concentration). Although there is a slight trend in the data, the current is essentially independent of flow rate, in agreement with Equation 18. However, the bimolecular rate constant calculated from the data on Table I1 averages around 21M-1 sec-1, in disagreement with the value of 43 M-' sec-l found from the data of Table I. Further studies with systems in which H202:Fe(III) decreased from 172 to 40 gave bimolecular rate constants that increased proportionately from 11 to 43 M-' sec-'. (38) 3. H. Baxendale, M. G. Evans, and G. S . Park, Trans. Faraday SOC.,42, 155 (1946).
516
ANALYTICAL CHEMISTRY
Figure 4. Plot of i/V'/aC,* us. 1/V2/3 Conditions: 0.25 F H2SO4, p = 0.0508 cm, X = 1.07 cm, E = f0.100 volt cs. SCE A . 7.88 mF H2O2 - 0.2 mF Fe(II1) 8. 5.00 mF H202 - 0.2 mFFe(II1) C. 2.97 mF HzO? - 0.1 mF Fe(I1I)
Although slight reduction of hydrogen peroxide occurs directly at the cathode, the magnitude of the current corresponds to only about 2 of a one-electron mass transportlimited current, and alone cannot explain the above trend. However, the mechanism first proposed by Haber and Weiss (39) for the oxidation of ferrous ions by hydrogen peroxide in neutral and acidic media does provide an adequate explanation :
+ OH- + O H . HzO + 0zH. OH. + H202 0 2 H . + HaOz O2 + H 2 0 + O H . Fe+2 + O H . Fe+3 + OH-
Fe+2
+ H202
+
Fe+3 +
4
+
(25) (26) (27)
Reaction 24 constitutes the rate-determining step. Reactions 25 and 26 represent a free-radical chain reaction leading to the decomposition of hydrogen peroxide, which is therefore catalytic. Reaction 27 is the chain-terminating step. At comparable concentrations of hydrogen peroxide and ferrous ions Reactions 24 and 27 proceed practically exclusively (38). Since the polarographic studies of Kolthoff and Parry (30) a r d Pospisil (33), and the homogeneous study of Baxendale et at. (38) clearly show that the reaction is first-order in both hydrogen peroxide and ferrous ion, the reaction between hydrogen peroxide and ferrous ions obeys all the postulates of the theoretical model corresponding to Reactions 1 and 2, and this reaction may be used to test the correctness of the theoretical model. However, in the presence of high proportions of HzOZ, Reactions 25 and 26 become important compared to Reaction 24, reducing the concentration of H202 especially in the region near the electrode surface where the free radicals are generated, and where Reaction 24 occurs. Since the bimolecular rate constant is calculated from the bulk H202concentration, these competing reactions are responsible for the decrease in bimolecular rate constants as the (39) F. Haber and J. Weiss, Naturwissenschaftetz, 20, 948 (1932).
H*Oi-Fe(III) ratio increases. The same effect has been observed by Kolthoff and Parry (30, 40). The numerical discrepancy between the calculated and literature values of the bimolecular rate constant may be attributed to the slight heterogeneous reduction of hydrogen peroxide, providing an additional source of free radicals which increases the rates of Reactions 25 and 26.
Although the demonstration system is complicated by these side reactions, the agreement between theory and experiment is excellent.
RECEIVED for review September 1, 1967. Accepted December 11, 1967. Work supported in part by the National Science Foundation, through Grant No. GP-3190, and in part by the Atomic Energy Commission, through Grant No. AT(11-1-)1082. A fellowship to Leon N. Klatt from the American Oil
Study of Impregnated Silicone Rubber Membranes for Potential Indicating Electrodes E. B. Buchanan, Jr., and James L. Seago’ Department of Chemistry, Uniuersity of Iowa, Iowa C i t y , Iowa
Electrodes fabricated from silicone rubber membranes and impregnated with various insoluble transition metal salts and chelates were investigated for their responses toward transition metal cations. Only electrodes fabricated from membranes impregnated with metal salts, and not chelates, responded. The crystalline form, the degree of hydration, and the associated anion of the imbedded salt had no effect on a membrane’s response or selectivity. These electrodes were nonspecific in their responses toward cations. The presence of additional electrolytes in the test solutions leveled the responses of the electrodes to the concentrations of the added salts. The primary purpose of the imbedded material is to adsorb water and ions during the conditioning process and to provide sites for rapid metal ion exchange at the membrane interface. The transition metal responsive electrodes were found to possess suitable mechanical properties but the desired selectivity was not attained.
SUCCESSFUL DEVELOPMENTof ion-selective electrodes for alkali and alkaline-earth metal cations has created a renewed interest in direct potentiometry as an analytical scheme. Few studies have been made of electrodes responsive to transition metal ions. Electrodes of this latter type would have direct application for the determination of transition metal ions found in biological specimens and industrial wastes. Rechnitz ( I ) , Hill ( 2 ) , Lakshminarayanaiah (3),and Pungor (4)have reviewed the various types of membranes employed in the attempted development of a specific ion electrode. Most reported works have dealt with the responses of these membranes in solutions of alkali or alkaline-earth metals. Chatterjee ( 5 ) prepared clay membranes that responded to 1 Present address, Procter & Gamble, Ivorydale Technical Center, Cincinnati, Ohio.
(1) G. A. Rechnitz, Chem. Eng. News, 45 (25), 146 (1967). (2) G. J. Hill, in “Reference Electrodes,” Ives and G. J. Janz, Eds., Academic Press, New York, 1961, chapter 9. (3) N. Lakshrninarayanaiah, Chem. Rel;., 65, 491 (1965). (4) E. Pungor, ANAL. CHEM., 39, 28A (1967). (5) B. Chatterjee and D. K. Mitra, J . Indian Chem. Soc., 32, 751 (1955).
52240
copper, molybdenum, and cobalt. Recently, MorazzaniPelletier and Baffier (6) published work on the use of certain insoluble transition metal salts embedded in collodion and paraffin membranes. They found membrane response was affected by the number and nature of the anions in the solution, the ionic strength of the solution, and hydrostatic effects on the membrane. These membranes possessed a high rate of water transference and the resultant potentials were subject to drift and were not selective. Shatkay (7) has reinvestigated the paraffin electrodes and has found that they are not completely permselective nor specific in their responses, At least two companies, Orion Research Laboratories and Corning Glass Works, are marketing specific liquid ion-exchange electrodes. Pungor (8-13) developed impregnated silicone rubber membrane electrodes which show selective responses to chloride, bromide, iodide, and sulfate ions. Rechnitz (14, 1 9 , utilizing Pungor’s silicone rubber membrane electrodes impregnated with the silver halides, characterized the responses of these electrodes toward chloride, bromide, and iodide ions. The present study is an attempt to evaluate those factors that contribute to the selectivity and sensitivity of impregnated membrane electrodes toward transition metal ions. The suitabilities of paraffin and silicone rubber as matrix materials are compared. Membranes containing transition metal compounds (chelates and ionic salts) in varying degrees of (6) S. Morazzani-Pelletier and M. A. Baffier, J . Chim. Phys., 62,429 (1965). 39, 1056 (1967). (7) A. Shatkay, ANAL.CHEM., (8) E. Pungor and K. Toth, Acta Chim. Acad. Sci. Hung., 41, 239 (1964). (9) E . Pungor, J. Havas, and G. Madarasz, Fr. Patent No. 1,402,343 (1965). (IO) E. Pungor, J. Havas, and K. Toth, Z . Chem., 5,9 (1965). (11) E. Pungor, J. Havas, and K. Toth, Inst. Control Systems, 38, 105 (1965). (12) E. Pungor, J. Havas, and K. Toth, Acta Chim. Acad. Sci. Hung., 48, 17 (1966). (13) E. Pungor, J. Havas, and K. Toth, Mikrochim. Acta, 1966, p 689. (14) G. A. Rechnitz, M. R. Kresz, and S. B. Zamochnick, ANAL. CHEM., 38, 973 (1966). (15) G. A. Rechnitz and M. R. Kresz, Ibid., p 1786. VOL. 40, NO. 3, MARCH 1968
a
517
