spectrometric surface measurements in the far infrared region to characterize the organomercuric compounds which are undoubtedly formed with oxidized graphite. The foregoing work shows rather conclusively that it is generally disadvantageous to do ASV in acidic solutions and particularly in acidic solutions containing oxidants with the electrodes described. Based upon several years of experience with ASV, if a t all possible, it is better to work a t pH values 2 3 and preferably on the p H interval between 4 and 7 . Work in the low p H range will necessitate the frequent repolishing of the electrode, thus necessitating the frequent decontamination of the cell and electrode if lead is the metal of interest and especially if it is present a t 50.1-ppb level. Workers who study chemical speciation will always be forced to work in acidic media, however, and will have to continually contend with the problems described in the introduction in addition to the problems of adsorption described in a previous paper ( 2 ) . Others who, as we are, are interested only in total metals bound to organic particulaie matter or in filtered water samples may be benefitted by the new ozone oxidation methodology presently under development a t this laboratory, because the natural pH of the sample is essentially unchanged after the oxidation. Also, naturally occurring organic compounds which cause electrode adsorption problems are destroyed.
ACKNOWLEDGMENT The author thanks W. M. Garrison and A. S. Newton, both of this laboratory, for their advice and comments during the course of this work and T. Novakov, also of this laboratory, for making the PES analyses. Thanks are also due the following Lawrence Livermore Laboratory personnel: R. H. Sanborn for making the infrared analyses, and D. C. Camp for supplying the neutron-damaged graphite speci-
mens. Thanks also go to R. C. Fox of Chevron Research, Richmond, Calif., for supplying us with some of the amines tested.
LITERATURE CITED (1) (2) (3) (4) (5) (6) (7) (8) (9) (10)
R. G. Clem, G. Litton, and L. D. Ornelas, Anal. Chem., 45, 1306 (1973). R. G. Clem and A. F. Sciamanna, Anal. Chem., 47, 276 (1975). W. R. Seitz, Ph.D. Thesis, M.I.T. (1970). R. G. Clem and W. W. Goldsworthy, Anal. Chem., 43, 918 (1971). W. W. Goldsworthy and R. G. Clem, Anal. Chem., 43, 1718 (1971). W. W. Goldsworthy and R. G. Clem, Anal. Chem., 44, 1360 (1972). T. J. DeBoer. Ph.D. Thesis, Groningen (1953). R. F. Nystromand W. G. Brown, J . Am. Chem. SOC.,69, 1197 (1947). G. R. Hennig, Prog. lnorg. Chem., 1, 125-201, (1959). V. L. Snoeyink and W. J. Weber, Prog. Surf. Membr. Sci., 5 63-113 (1972). (1 1) R. E. Nightingale, "Nuclear Graphite", Academic Press, New York, N.Y., 1962. (12) J. J. Lingane, "Electroanalytical Chemistry", 2nd ed. interscience, New York, 1958, p 209. (13) K. W. Pepper, J . Appl. Chem., 1, 124(1951). (14) H. P. Boehm, Adv. Catal. Relat. Subj., 16, 179 (1966). (15) G. R. Hennig, "Proc. Conf. Carbon 5th", Vol. 1. 1961, p 143. (16) Y. A. Zarif'yanz, V. F. Kiselev, N. N. Lezhnev, and D. V. Nikitina, Carbon, 5, 127 (1967). (17) L. G. Makarova and A. N. Nesmeyanov, "Methods of Elemento-Organic Chemistry", Vol. 4, North-Holland Publishing, Amsterdam, 1967. (18) T. M. Florence, J. Nectroanal. Chem., 27, 273 (1970). (19) J. A. R. Samson and R. B. Cairns, Appl. Opt., 4, 915 (1965). (20) J. G. Calved and J. N. Pitts, "Photochemistry", Wiley, New York, 1967. (21) W. R. Matson, Ph.D. Thesis, M.I.T. (1968). (22) M. L. Studebaker, E. W. D. Huffman, A. C. Wolfe, and L. G. Nabors. lnd. Eng. Chem., 48, 162 (1956). (23) F. C. Whitmore, "Organic Chemistry", Vol. 2, 2nd ed., Dover, New York, 1961, pp 675-676. (24) M. Stulikova, J. Nectroanal. Chem., 48, 33 (1973). (25) J. S. Mattson and H. B. Mark, "Activated Carbon", Marcel Dekker. New York, 1971, p 135. (26) F. J. Miller and H. E. Zittel, Anal. Chem., 35, 1866 (1963).
RECEIVEDfor review February 13, 1975. Accepted June 12, 1975. Work performed under the auspices of the U S . Atomic Energy Commission.
Polarographic Measurement of Diffusion Coefficients for Potassium Ion in A/,"-Dimethylformamide-Water Mixed Solvents Makoto Saito Department of Chemistry, College of Liberal Arts, Toyama University, 3 190 Gofuku, Toyama 930, Japan
The diffusion coefficients of K+ ion in the mixed solvents, DMF and water, were determined by a polarographic method. It was concluded that the diffusion coefficient was closely related to the density of the solvent rather than to its viscosity. The relationship of the difference of apparent and D'), vs. the percentmeasured diffusion coefficients, (D" age of water in the solvent showed the same trend as that of the difference between the measured and apparent density, ( d - d'), vs. the percentage of water. From this, it is concluded that the diffusion coefficient is closely related to the structure of the solvent.
-
The electrochemical behavior of electrolytes in homogeneous aqueous-organic solvent mixtures has been studied for a long time, and many investigations of the electrochemical behavior of metal ions in mixed solvents have 1784
been carried out in media containing alcohols (1-4) as the amphiprotic solvents. Recently studies have been extended to include ions in water-DMSO (5-7). There have been few investigations on the electrochemical behavior of metal ions in DMF-water mixed solvents. Gaur et al. (8-10) studied the polarography of Mn(ClOJ2, Ni(C104)2, and Co(C104)2 in 0, 20, 40, 60, and 80% water and DMF mixed solvents containing 0.1M NaC104 as supporting electrolytes. They discussed the number of reduction electrons and the diffusion coefficients of these metal ions from the polarograms obtained. Takaoka et al. (11) investigated the polarographic behavior of alkali metal ions in DMF-water mixed solvents. In their paper, they have recommended a common electrode potential scale to be applied to all solvents in which the potential of a dropping mercury electrode in a DMF-water mixed solvent is measured vs. an aqueous calomel electrode. The present paper summarizes the following results: the
ANALYTICAL CHEMISTRY, VOL. 47, NO. 11, SEPTEMBER 1975
I
Table I. Half-Wave Potentials of Potassium Ion in Various DMF-Water Mixed Solvents
o/
-2.120c
E 1 / 2 ( V ) vs. S C E VOl. no
H2@
0 10 25 50 75 100
~
1
X
10-3YK'
-2.044 -2.044 -2.045 -2.071 -2.096 -2.127
2
X
iO-3hfKc
-2.043 -2.044 -2.042 -2.068 -2.095 -2.130
4
X
iO-3!dK+
-2.042 -2.041 -2.039 -2.065 -2.103 -2.126
6
X
iO+YK*
-2.043 -2.042 -2.048 -2.075 -2.100 -2.130 -2.0401
hydration number for potassium ion; the relationship between the diffusion current and the concentration of potassium ion; the relationship between the diffusion coefficient of potassium ion and the density of the mixed solvent (water-DMF).
EXPERIMENTAL Solvents a n d Materials. DMF was purified by the same method described in a previous paper (12), involving treatment with calcium hydride (shaking and standing for 48-72 hr) and then twice distilled under reduced pressure. The boiling range was 37-39 "C a t 9-10 mm Hg pressure. The specific conductivity was in the range 7.0-7.5 X ohm-' cm-', which is in agreement with the lowest values reported elsewhere (13-15). Karl Fischer titrations indicated the water content in the distillate to be less than 0.03%. There were no impurities, which caused a detectable current a t negative potentials less than +0.5 -2.7 V vs. SCE in the polarogram. Tetraethylammonium perchlorate (TEP) was used as the supporting electrolyte, and was prepared by reaction of tetraethylammonium bromate with perchloric acid, followed by recrystallization from water and drying a t 60 O C in a vacuum oven for 24 hr. Other reagents were analytically pure grade reagents used without further purification. Apparatus. The electrolysis cell which is a conventional Hshaped cell has one diaphragm of sintered glass and consists of three electrodes. Polarograms were obtained with a dropping mercury electrode polarograph, Yanagimoto P8-DP. The reference electrode used was SCE with a salt bridge devised by Takaoka ( 1 1 ) . Three different capillaries were constructed for these experiments, and their characteristics, m and t established. Capillary I (open circuit): m = 1.46 mg/sec and t = 4.12 sec for h = 60 cm; Capillary I1 (open circuit): m = 1.50 mg/sec and t = 4.22 sec for h = 70 cm; Capillary I11 (open circuit): rn = 1.72 mg/sec and t = 4.30 for h = 50 cm. Procedure. A sample of 0.1M TEP-DMF-water containing K+ ion was placed in the cell, and dried argon gas was passed through it for about 20 min. DMF and the agar bridge connected with SCE were inserted into the cell. Polarograms were recorded 5-6 times for each test solution at 25.0 f 0.1 OC. Apparent half-wave potentials (PI/*) and diffusion currents (id) were determined from the polarogram obtained by application of Hohn's "Schnittpunkt Methode" ( 1 6 ) .The density of DMF-water mixed solvent was measured with a pycnometer. The pycnometer volume was determined from the value of direct measurement of pure water a t 25.0 f 0.1 "C and the literature value of the density of pure water a t 25 "C ( 1 4 ) .
-
RESULTS AND DISCUSSION Half-Wave Potentials of the Polarographic Wave of Potassium Ion. The half-wave potentials of potassium ion in various solvent mixtures listed in Table I are independent of the concentration of potassium ion, but are significantly changed by any variation in the water content of the solvent. The result for 1 X 10-3M potassium ion is shown in Figure 1, in which the half-wave potential does not change until the percentage of water in the solvent reaches 25%; and the direction of the shift in the half-wave potential is toward the negative potential with an increasing percentage of water. According to Takaoka ( I I ) , the half-wave potential did not change until water comprised about 50%.
-2.020 0
25
75
50
100
% of water
Figure 1. Relationship between half-wave potential and percentage of water in solvent
Table 11. Hydration Number of Potassium Ion in Mixed Solvent Percentage of water
&umbers of hydration
25-50 50-75 75-100
1.5 2.6 4.2 2.8
Av.
The latter result does not agree with the result obtained in this experiment. Hydration Number for Potassium Ion in Solvent. T h e hydration number for potassium ion in DMF-water mixed solvent was determined by Lingane's method ( I 7). Half-wave potential of a complex metal ion should shift with changing activity of a complex forming substance according to
a t 25 "C, where A E 1 / 2 is the difference of half-wave potential caused by formation of a complex ion; A log Cd,, the change in activity for the complex-forming species; p , the number of ligands per complex ion; and n is the number of Faradays of electricity required per mole of electrode reaction. From Equation 1,the following equation is finally obtained, 0.0591 n
0.0591 n
(E1/2)c- (E1/2)s = -log K c - p -log C x (2) where ell^)^ and ( E I / ~are, ) ~respectively, half-wave potential of complex ion and single ion, K , is the stability constant of complex ion, and C, is the molar concentration of the complex-forming substance. As seen from Equation 2, the linear relationship exists between ( E ~ / Z - )( ~E I / Z ) ~ and log C,; and K , is obtained from extrapolation to C, = 1. The relationship between ( E I / z ), ( E I / ~and ) ~log C, a t the region from 25-100% of water was found to be linear, and p = 2.32 and K , = 1.58 X was obtained. The hydration numbers for potassium ion in solvents containing 25-50, 50-75, and 75-100% water were calculated from Equation 1. The results are shown in Table 11; and the general trend is one in which the hydration number for potassium ion increases with increasing percentage of water in the solvent.
ANALYTICAL CHEMISTRY, VOL. 47, NO. 11, SEPTEMBER 1975
1785
Table 111.Relationships between Diffusion Current and/or
id X ~ ~ 1 a ’n d 2
Concn of K
i x 10-3 Vol. “I H 2 0
id
2 .oo 1.68 1.36 1.75 2.03 3.50
0.
10 25 50 75 100
ion ( M )
2 x 10-3
i d x n,.ll2
1.90 2.06 2.01 2.83 3.15 3.50
id x n r 1 / 2
id
3.79 3.24 2.89 3.27 4.46 6.80
3.60 3 -97 4.28 5.04 6.11 6.80
4 x
x
Vol. 94 H 2 0 measured
0 10 20 25 30 40 50 75 100 a
0.88 0.62 0.51 0.50 0.50 0.60 0.66 1.27 1.83
x nr112
X
lr3
id
id
6.80 7.71 8.66 10.79 12.95 15 .OO
20
3
measured
1.845”
id
id
d(9/cm )
(cm2/sec)
6
7.17 6 -30 5.85 6.65 9.45 15.00
Table IV. Measured Values of the Diffusion Coefficient ( D ’ )of Potassium Ion a n d Densities of Various Mixed Solvents D’
Percentage of Water in Mixed Solvent
10.96 9.36 8.90 10.20 13.30 21 .oo
x nrl12
nr1/2
10.40 11.46 13.17 16.22 18.22 21 .oo
0.944 1.224 1.470 1.622 1.370 1 .ooo
1 1.50
0.9446 0.9638 0.9779 0.9836 0.9880 0.9939 0.9968 0.9967 0.9970
5s 1.00
0.50
0.9970b
Ref. 19. Ref. 14. O l
10.00
Table V. Computed a n d Measured Values of Diffusion Coefficients of Potassium Ion in Various Mixed Solvents D x Vol. @i H 2 0
0 10 20 25 30 40 50 75 100
Measured
0.88 0.62 0.51 0.50 0.50 0.60 0.66 1.27 1.83
100
vr”2,
Figure 2. Relationship between id X qrl’z, of water in the solvent
and the percentage
(cm2/sec) Computed
...
0.8
-
“‘:j/-\
0.59 0.45 0.43 0.43 0.48 0.59 1.20
-
0.2 0.ok
...
T h e Ilkovic Equation Applied to Potassium Ion. The relationship between the concentration of potassium ion and the diffusion current of the polarographic wave a t various ratios of DMF and water in the solvent is shown in Table 111, and the results confirm the Ilkovic equation. The relations between the product of the diffusion current and the square root of relative viscosity, and the percentage of water in the solvent with various concentrations of potassium ion are shown in Figure 2. It is evident that a linear relationship holds between (id X vr1/2) and the concentration of potassium ion or the percentage of water in the solvent. The slopes of these straight lines increase with increasing concentration of potassium ion and in the percentage of water in the solvent. In the case of the relationship between (id X vr1I2) and the percentage of water in the solvent, the slopes were respectively 1.80, 3.54, 6.74, and 10.68 at the various concentrations of potassium ion. Current-Voltage Curves of Potassium Ion. The log plots of current-voltage curves of the polarographic waves of potassium ion in the solvents were linear with slopes of 1786
50 75 % o f water
25
0
0.03
L
.
,
,
0
Figure 3. Relationship between (G” centage of water in the solvent
,
‘
I
50 % o f water
I
1
100
- 03 or ( d - dq and the per-
70.35 mV at 10% water, 60.30 mV at 50% water, and 70.30 mV a t 100% water in the solvents. From the values of the slopes, the calculated value of n in each solvent were 0.90, 0.98, and 0.84, respectively, indicating that the electrode reaction of potassium ion was essentially a one electron reduction. Diffusion Coefficient-Density Relationship. The values of the diffusion coefficients of potassium ions were calculated from the Ilkovic equation (Table IV). The plots of the relationship of (D” - D’) or ( d - d ’ ) to the percentage of water in the solvent are shown in Figure 3. As is evident from Figure 3, the changes of (D” - D’) and ( d - d ’ ) with the percentage of water in the solvent indicate the
ANALYTICAL CHEMISTRY, VOL. 47, NO. 11, SEPTEMBER 1975
LITERATURE CITED
same tendency, suggesting that the substantial change in the diffusion coefficient of potassium ion (D” - D’) is controlled by the destruction of a bulk of the structure of the solvent ( d - d’). The value of (D” - D’)per unit density (1 ‘g/cm3)was examined, and it can be seen from Table V that the value of (D”- D’)/(d - d ’ ) is almost constant the mean being 2.6 X lod4 (cm5/g sec). In the case of a large percentage of water in the solvent, it seems that the structure of water is disordered by the molecule of DMF; and, on the other hand, in the case of a large percentage of DMF which is bonded with a weaker intermolecular force than that of water, the structure of DMF is destroyed by the water polar molecule (18). The fact that the degree of destruction of the structure of DMF by water molecules is larger than that of the destruction of the structure of water by DMF molecules, is estimated from the slope of ( d - d ’ ) curve on Figure 3. As is evident from Figure 3, the ratio of DMF to water in the solvent where the greatest destruction of the structures of DMF and water takes place, is in the range of 30-40% water.
(1) A. E. Brodsky. Fresenius’Z. Phys. Chem.. 121, l(1926). (2)J. Sancho, A. Aldaz, and A. PuJante, J. Electroanal. Chem., 25, 505 (1970). (3) F. N. Kozlenko and S. P. Miskidzh’yan, Russ. J. Phys. Chem., 39, 506 (1965). (4)G. Kugler and G. A. Rechnitz, Fresenius’ Z. Anal. Chem., 214, 405 (1965). (5) J. P. Morel, D S c . Thesis, Lniversity of Clermont-Ferrand, 1969. (6)K. H. Khoo, J. Chem. Soc., A, 1177 (1971). (7)N. A. Izmailov, Russ. J. Phys. Chem., 34, 1142 (1960). (8)J. N. Gaur, Electrochim. Acta, 11, 939 (1966). (9)J. N. Gaur, Electrochim. Acta, 12, 1489 (1967). (10)J. N. Gaur, Electrochim. Acta, 15, 519 (1970). (11)K. Takaoka, Rev. Polarogr. (Kyoto),14, 63 (1966). (12)M. Saito and M. Tanaka, J Coll. Liberal Arts, Toyama Univ., 7, 27 (1975). (13) C. M. French and K. H. Clover, Trans. faraday SOC.,57, 1975 (1961). (14)D. P. Ames and P. G. Sears, J. Phys. Chem., 59, 16 (1955). (15)L. R. Dawson, M. Colbent, G. M. R. Leader, and H. K. Zimmerman, Jr., J. Electrochem. Soc., 99, 28 (1952). (16)I. Tachi, “Polarography,” lwanami Shoten. 1954,p 196. (17)J. J. Lingane, Chem. Rev., 29, l(1941). (18)J. A. Riddick and W. B. Bunger, “Organic Solvents,” 3rd ed., Wiley-lnterscience, New York, 1970. (19)R. A. Robinson and R. H. Stokes, “Electrolyte Solutions,” Butterworths, London, 1959,pp 513-5.
ACKNOWLEDGMENT The author thanks K. Takaoka for the relative viscosity data and H. Kanda for helpful assistance.
RECEIVED for review, November 14, 1974. Resubmitted April 21, 1975. Accepted May 27, 1975. Presented a t the 23rd Annual Meeting of the Japan Society for Analytical Chemistry, Osaka, October 2-4, 1974.
Fabrication and Evaluation of Glass Microbulb Electrodes J. D. Czaban and G. A. Rechnitz Department of Chemistry, Sfate University of New York, Buffalo,N.Y. 74274
Routine preparation of glass microelectrodes with sensing bulb sizes in the 100- to 500-pm range is made possible by a newly designed, seml-automatic fabrication apparatus which provides reproducible control of critical fabrication parameters. Evaluation of the resulting miniature electrodes reveals excellent response characteristics and suggests novel analytical and biomedical applications.
We also describe an inexpensive preamplifier which permits these electrodes to be used in conjunction with any conventional p H meter. The system has been tested by evaluation of the electrodes and instrumentation under a variety of circumstances with solution standards and in acrylamide gels for the measurement of pH gradients in isoelectric focusing procedures.
EXPERIMENTAL Glass electrodes are generally available in two size ranges, the conventional bulb type electrode having a 5- to 10-mm diameter and the ultra-micro type with sensing tips in the micron size region for intracellular use (1-3). The latter type of electrode is extremely delicate and difficult to use. There is a need for miniaturized glass electrodes in an intermediate size range, e.g., 100- to 500-pm tip diameter, having the ruggedness and convenience of the common macroelectrode type. We now describe techniques for the fabrication of such electrodes with sensing bulb diameters as small as 175 pm. Evaluation of the resulting electrodes shows excellent pH response, durability, and dynamic behavior. Using the experimental apparatus detailed below, such electrodes can be routinely prepared in the laboratory. The dimensions and properties of the miniaturized electrodes are such that they may be used as sensors for biomedical monitoring and other analytical purposes. For example, the electrodes can be mounted inside a 26-gauge hypodermic needle for “in vivo” monitoring of muscle pH, a procedure which requires surgery ( 4 ) with conventional electrodes.
Apparatus. Electrode Fabricator. Figure 1 shows a block diagram of the electrode fabricator. Briefly, the system operated as follows: when the foot pedal was depressed, the injector moved the electrode into the microflame for a preset time interval. After an independent time delay, the air pressure generator formed the bulb by forcing air into the hot electrode glass via the two-way automatic valve and the bleeder valve. Once the bulb was shaped and as the air pressure was slowly released, the electrode was removed from the flame and allowed to cool quickly. The entire process was completed in a few hundred milliseconds, or less, depending on the size of the electrode. The timing functions were generated in the control circuit with the heating and delay intervals controlled by calibrated variable resistors mounted for convenient adjustment. These two parameters were simultaneously displayed on a dual trace storage oscilliscope and the pressure developed in the electrode was monitored with a miniature manometer (not shown). The microtorch was constructed by soldering telescopic sections of stainless steel tubing together until the desired orifice was obtained. The fuel gas (propane) and oxygen were proportioned by a pair of precision metering vales (Nupro, B-4MG) and mixed before entering the microtorch. The electrode injector which moved the electrode into the flame is shown in Figure 2. The stop lug on the injector arm was used to adjust the total travel of the arm. If this lug was set such that the arm travel was greater than the travel of the solenoid plunger, the momentum of the spring-loaded injector arm would cause it to ov-
ANALYTICAL CHEMISTRY, VOL. 47, NO. 11, SEPTEMBER 1975
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