In actual practice, cyanide is often destroyed by chlorination under alkaline conditions. To demonstrate that intermediate products, such as CNO-, do not interfere, the silver indicator procedure was used to follow the titration of a KCN solution by hypochlorite (Figure 2). When the end point is passed, all of the cyanide has been destroyed, including the indicator itself. The fact that a single, smooth curve is obtained, with an extremely large end-point break, suggests that none of the reaction products interfere with the electrode response. The method should perform well in the analysis of both chlorinated waste waters and untreated samples. The successful application of this method to the laboratory
analysis of actual cyanide waste samples has been reported (10). The time required per analysis is about 5-10 min for those samples which do not contain inert transition metal cyanide complexes, such as Ni and Cr. Inert complexes increase the time of analysis by approximately 5-10 min. per analysis. No prior distillation should be necessary for samples of practical interest. RECEIVED for review May 8, 1972. Accepted July 10, 1972. ~_____ ~~~~
(10) L. E. Lancy, Lancy Laboratories,Zelienople, Pa., personal com-
munication, 1972.
Some Characteristics of Several Commercially Available Cation-Responsive Glass Electrodes Sonny Phangl and B. J. Steel Department of Physical and Inorganic Chemistry, University of Adelaide, Adelaide, S . Australia ALTHOUGH THE LITERATURE abounds with reports of various uses and properties of cation-responsive glass electrodes, it appears that there is a need for more information regarding the properties of commercially available cation-responsive glass electrodes ( I , 2). At present most of the reports concern electrodes which have been made from glasses which were either supplied by the manufacturers or which have been produced by the workers themselves in their laboratories. This report contains the characterization and the conditions necessary for such electrodes to be useful in the determination of activity coefficients. EXPERIMENTAL Apparatus. The low impedance side of the circuit consisted of a micro-step potentiometer (Type 44248, Cambridge Instrument Co. Ltd., U.K.), one terminal of which was connected to the Ag, AgCl electrode while the other was grounded. The high impedance side of the circuit consisted of a Cary Model 31CV vibrating reed electrometer. The latter was used as a null-detector. A grounded water-bath was used throughout this investigation and an opaque brass lid was used to cover the bath. It acted as both a light and an electrical shield. Table I shows the type of glass electrodes used in this work. The GEA33, GEA 33/C, and the Beckman 39278 electrodes are sodium-responsive and are designed for use in routine measurement of sodium ion concentrations. The Beckman 391 37 Cationic Electrode is meant generally for monovalent cations. Reagents. Recrystallized AR sodium chloride was used. Nitric acid and sodium hydroxide were of AR grade. The hydrochloric acid was obtained from a constant boiling mixture. All the solutions were prepared by using doubly distilled water. 1 Present address, School of Chemical Sciences, Science University of Malaysia, Penang, Malaysia.
(1) “Glass Electrodes for Hydrogen and Other Cations,” G. Eisenman, Ed., Marcel Dekker, Inc., New York, N . Y . , 1967.
(2) G. Mattock, Chimia, 21,209 (1967). 2230
Table I. Glass Electrodes Used in This Work Manufacturer Label on electrode Type EIL GEA 33 1 Beckman 39278 3 Beckman 39278 4 Beckman 39137 5 Beckman 39278 6 39137 Beckman 7 EIL GEA 33 8 EIL GEA 33/C 10
Procedure. The electrodes were tested at 25 “C. Earlier tests had shown no electrical leakage by using a water-bath and thus the inconvenience of an oil-bath was avoided. Each cell was magnetically stirred and only aged thermal electrolytic type Ag, AgCl electrodes were used. RESULTS AND DISCUSSION Selectivity Constant. By considering a cell of the type,
Ag, AgCl~HClor NaCl soln~GlassElectrode and the glass electrode potential written as
E
=
constant
+ RT - ln(aH + K H N ~ C J N ~ ) F
it can be shown ( 3 , # )that El - E2 = RT -In F
(1)
KHN~
El is the emf of the cell in O.lm NaCl solution in the limit as H+ ions + 0. E? is the emf of the same cell in O.lm HC1 solution. , how effective an elecThe selectivity constant, K H N ~shows trode “sees” the Na+ ions relative to the H+ions for Na-re(3) G. Eisenman, D. 0. Rudin. and J. U. Casby, Science, 126, 831 (1957). (4) G. Eisenman et a/., “The Glass Electrode,” Interscience Publishers, John Wiley and Sons, New York, N.Y., 1962, p 232.
ANALYTICAL CHEMISTRY, VOL. 44, NO. 13, NOVEMBER 1972
Table 11. Selectivity Constants of the Glass Electrodes Emf of the cells (mV) 0. l m NaCl KEN& Electrode 0. l m HCl pH 9 . 0 K H N ~ (Literature) 2.69 1 31.68 3 1 3
Av emf, 1 Avemf, 3 4 5
4 5
311.91 35.38 310.55 33.52 311.23 235.00 174.97 235.95 175.66 235.48 175.32
139.08 2.16 136.19 2.43 137.64 83.98 101.64 108.01= 102.64
0.546 0.034
Av emf, 4 102.14 Av emf, 5 Non-equilibrium value after 4 days. NAS 11-18 glass (5).
0.06*
0.241
0
I
80
40
mo
120 TIUE,mkutas
+
200
260
Figure 1. Photoelectric effects on cation-responsive glass electrodes
.
P = Electrode 7, Q trode 8
=
Electrode 6, R
=
Electrode 10, S
=
Elec-
A = Light on B = Light off
sponsive glass electrodes. For example, when KHS* = 10, the Na+ ions are ten times as effective as the H+ ions on a molefor-mole basis in determining the electrode potential. Thus using Equation 2, the selectivity constants of electrodes 1, 3, 4, and 5 were determined. Two cells, one containing exactly O.lm NaCl solution (adjusted to p H 9 using NaOH) and the other, O.lm HCI solution, were set up in the water bath. The glass electrodes were transferred from one cell t o the other and each time they were rinsed with the solution in which they were to be immersed. The final emf values together with the values are shown in Table 11. The latter also indicates the steps taken in measuring the emf’s. For example, electrodes l and 3 were first measured in O.lm HCI solution and then transferred to the O.lm NaCl solution. The average emf arrived a t in each solution was used in the calculation of K H N ~ All . the electrodes responded to within a few millivolts of the equilibrium value in less than 1 hour. However, each took several hours to reach the equilibrium value. An emf change of less than 0.02 mV in 15 minutes was considered as having reached the equilibrium state. Electrode 1 (GEA 33) appeared to equilibrate much faster than the others. Such a n experiment also appeared to impose a considerable strain on the glass electrodes. Electrodes 3 and 4 (both Beckman 39137), showed the symptom of “peeling”i.e., flakes of glass appeared to peel off at the bulb towards the end of the experiment. Electrode 4 became erratic and drifted with time without reaching the equilibrium state even after 4 days in the NaCl solution. Presumably, as a result of this experiment, Electrode 4 never completely recovered and soon had t o be discarded. The values of K H S were ~ compared with Eisenman’s NAS 11-18 sodium responsive glass (NanO,11 mole per cent; A1203, 18 mole per cent; S O ? , 71 mole per cent). All the KHS& values indicate that the electrodes responded more towards H+
~
_
_
(5) Reference 1. p 426.
_
~
Table 111. Response of Electrode to pH Changes in 0.5m NaCl Solution pH adjusted using nitric pH adjusted using hydrochloric acid acid __Emf of electrodes, Emf of electrodes, mV _- mV 1 3 PH 1 3 PH 74.18 74.09 74.03 74.06
196.50 196.61 197.19 219.71
1
5
73.64 73.44 73.56
173.75 173.68 173.54
a
9.00 7.12 6.27 4.22
72.76 72.60 72.53 72.60 1
5
8.72 6.22 4.52
72.23 72.27 72.32 12.70
169.30 169.79 170.28 170.99
189.48 188.25 188.63 205.80
7.10 8 . lja 7.25 5.05 9.63 7.06 6.30 5.40
NaOH was used to increase the pH.
ions than Na- ions. Among the electrodes tested, Electrode 1 was apparently the best Na-responsive electrode. It is of interest to note that Savage and Isard (6, 7) have shown that glass of composition NAS 3.3-0.4, NAS 2.75-0.75, and LAS 2.5-0.5 have K H X &values greater than 1. K13xa values for commercial glass electrodes are not readily available for a useful comparison to be made with those from this work. Response to pH Variations. Electrodes 1, 3, and 5 were chosen for the experiment. The aim was to determine how p H variations may contribute to errors when measuring the sodium ion activity. For example, in a 3-component system like H,O-NaCl-X, the introduction of the third component, X, may alter the p H value of the existing 2-component solution. The electrodes were placed in 0.5m NaCl solution and kept at 25 “C in the water-bath. Specially prepared solutions of H N 0 3 , HCl, and NaOH were made up in order that the p H could be changed by using a minimum number of drops of acid or base. The p H was measured using a conventional p H meter and the results are shown in Table 111. Of the three electrodes tested, Electrode 1 was the least affected by pH variations, followed by Electrode 5 , and then Electrode 3. Such a behavior was consistent with the K H N ~
(6) J. A. Savage and U. 0. Isard, Pliys. Cliem. Classes, 3, 147 (1962). (7) Zbid.,7, 60 (1966).
ANALYTICAL CHEMISTRY, VOL. 44, NO. 13, NOVEMBER 1972
0
2231
~ for Elecvalues found earlier. From Table 11, K H =~0.546 trode 1, 0.241 for Electrode 5, and 0.034 for Electrode 3. Mattock ( 2 ) has shown that above 0.01 M Na+ ion concentration and for pH >5, the effects of pH changes would become unimportant for BH 68, BH 104, and NAS 11-18 glasses. His measurements were carried out using a conventional pH meter and the background medium of his solutions was enthanolamine hydrochloride. Since Electrode 1 (GEA 33) is known to be made from BH 104 glass (8),the findings here are in accord with Mattock's. Response to Light Variations. Electrodes 6 , 7, 8, and 10 were tested using a Mazda incandescent tungsten-filament lamp (60 watt, 240 V). The electrodes were placed in a reaction cell containing O.lm NaCl solution and magnetically stirred in the water-bath at 25 "C. We found that Ag, AgCl electrodes of the thermal and thermal-electrolytic type were not affected by light emitted from the Mazda incandescent lamp. Recently work by Moody et al. (9), and Milward (10) have confirmed our findings. The reference electrode was therefore left unshielded from the light. A glass window in the water-bath allowed the exposure of the electrodes to the lamp which was held about 20 cm away.
The results are shown in Figure 1. A represents the time when the light was switched on, and B when the light was switched off. The emf at zero time was taken as the reference value. The emf was showing a steady value before the light was switched on. The subsequent increase in emf was more pronounced for the Beckman than for the GEA electrodes, In fact, for Electrode 7, the emf was still increasing after the light was switched off. Electrode 8 (GEA 33) was the least affected. Milward (10) has reported the effects of light on glass electrodes, He observed a decrease in the pH when the electrodes were illuminated with an Osram mercury vapor lamp, a Desaga Heidelberg T.L. illuminator, a Mazda infrared lamp, and sunlight. It should be noted that if one had plotted pNa us. time, a graph of similar shape as that shown by Milward would be obtained. In our case, an increase of about 6 mV would correspond to a decrease of 0.1 pNa unit. The effects of light on the glass electrodes as shown in Figure 1 are serious when accurate activity coefficientsare to be measured.
(8) A. K. Covington and T. H. Lilley, Phys. Chem. Glasses, 8, 88 ( 1967). (9) G. J. Moody e f a/., Analyst, 94,803 (1969). (10) A. F. Milward, ibid., p 154.
RECEIVED for review May 16, 1972. Accepted July 18, 1972. One of us, S. P., is grateful to the University of Adelaide for a University Research Grant.
ACKNOWLEDGMENT
The authors thank R. A. Robinson for reading the manuscript and making several helpful suggestions.
Electrochemical Reduction of Mercury(I) and Mercury(II) on Platinum in Fused Sodium-Potassium-Nitrate Eutectic at 250 "C H. S. Swofford, Jr., and James Dietz School of Chemistry, Unicersity of Minnesota, Minneapolis, Minn. 55455
THEELECTROCHEMISTRY OF Hg(1) and Hg(I1) at stationary platinum microelectrodes in fused chloride media has been studied recently. We report below some of the results of our studies in non-complexing fused media (Na-KNO,, 250 "C). Laitinen and Liu ( I ) and Laitinen, Liu, and Ferguson (2) have observed a one-step, two-electron reduction of Hg(I1) to Hg(0) at a platinum microelectrode in Li-KCI eutectic at 450 OC. Hg(1) was reported to be unstable and to disproportionate to Hg(I1) and Hg(0). In the low-melting (70 "C) ternary eutectic A1C13-NaC1-KC1, Hames and Plambeck (3) have observed the reduction of Hg(I1) to occur in two, two-electron steps at a stationary tungsten microelectrode. Using half-wave potentials obtained in their work, they calculated the equilibrium constant for the disproportionation reaction (1) H. A. Laitinen and C. H. Liu, J. Amer. Chem. SOC.,80, 1015 (1958). (2) H. A. Laitinen, C. H. Liu, and W. S. Ferguson, ANAL.CHEM., 30, 1266 (1958). (3) D. A. Harnes and J. A. Plarnbeck, Can. J . Chem., 46, 1727 (1968).
2232
Hgz+
+ Hgo -+
Hgz*+
(1)
to be 3.6 X lo3. In a similar system Torsi and coworkers ( 4 , 5 ) cite both electrochemical and spectrochemical evidence for the existence of the Hga2+ion. The voltammetric behavior of Hg(1) and Hg(II), at platinum electrodes, in fused nitrate solvents has not been reported. Mazzocchin et al. (6) potentiometrically determined the standard electrode potentials for the Hg/Hg,*+, Hg2*+/Hg2+ couples in Li, NaKNO, eutectic at 150 "C, and reported values of $0.060 V and $0.270 V, respectively, cs. a 1.0-molal Ag/Ag+ reference electrode. Mercury(1) was indicated as stable in the melt for at least a few hours, and a value of 2.24 X l o 2 was reported for the equilibrium constant for Reaction 1 above. The electrochemical oxidation of mercury electrodes in fused Na-KNO, eutectic (7) and in fused Na-KNOJ eutectic con(4) G. Torsi and G. Marnantov, Inorg. Nucl. Chem. Lett., 6, 843 ( 1970). (5) . , G. Torsi, K. W. Fung, G. M. Begun, and G . Marnantov, Inorg. Chem. 10,'2285 (1971); ( 6 ) G. A. Mazzocchin. G. G. Bornbi. and M. Fiorani, J . Electroanal. Chem., 17, 95 (1968). (7) H. S. Swofford, Jr., and C. L. Holifield, ANAL.CHEM.,37, 1513 (1965). .
I
ANALYTICAL CHEMISTRY, VOL. 44, NO. 13, NOVEMBER 1972
