Polarographic Reduction of Diphenylthallium(lll) Cation in Aqueous Solution Joseph S. DiGregorio and Michael D. Morris Department of Chemistry, The Pennsylvania State University, University Park, Pa. 16802 The electrochemical reduction of diphenylthallium(lll) cation in aqueous solution has been investigated by polarography, controlled-potential electrolysis and cyclic voltammetry. A mechanism is proposed to account for the observed concentration dependence of the reduction. Investigations have been made of the decay schemes for the transient organothallium species formed during electrochemical reduction. The strong adsorption of the parent cation and of the electrogenerated intermediates formed at the electrode surface has been studied.
THALLIUM FORMS ionic organocompounds of the type RzTIX where R is an alkyl or aryl group and X is a halide, cyanide, o r other anion. The chemistry of dialkyl- and diarylthallium compounds has been reviewed recently by Yasuda and Okawara (1). Many species of these types are water soluble and very atable and unreactive (2). I n oxygen donor solvents, the dimethylthallium(II1) ion has a linear C-TI-C structure ( 3 4 and is iso-electronic with dimethylmercury. No studies o n the structure of the diphenylthallium(II1) cation have been made, but presumably this ion has a linear structure also. Dialkyl- and diarylthallium(II1) cations are electroreducible at the dropping mercury electrode. Costa ( 5 ) studied the polarographic reduction of dialkylthallium bromides, RzTIBr, where R is an ethyl, n-propyl, or n-butyl group, in aqueous propanol solution. He reported that reduction occurs to the bivalent and monovalent organothallium species and ultimately t o thallium amalgam. Coulometry was used to deduce the valencies of the unstable species formed, but n o attempt was made to identify the reaction products. We have examined in detail the mechanism of the polarographic reduction of the diphenylthallium(II1) cation, (CeH5)zT1C,and have proposed decay schemes for the transient organothallium species formed at the dropping mercury electrode. EXPERIMENTAL
Reagents. Diphenylthallium bromide was prepared by the Grignard synthesis (6). Solutions of diphenylthallium fluoride were prepared by reaction of diphenylthallium iodide (Metallomer Laboratories, Fitchburg, Mass.) and aqueous silver fluoride (7). The electrochemical behavior of diphenylthallium bromide and diphenylthallium fluoride is identical. Phenylmercury acetate, mp 149-151 “C (J. T. Baker), diphenylmercury, mp 124-125 “C (City Chemical (1) K. Yasuda and R. Okawara, Organometal. Chem. Rev., 2, 255 (1967). (2) G. E. Coates, “Organometallic Compounds,” John Wiley and Sons, Inc., New York, N. Y.,1960, p 158. (3) G. D. Shier and R. S. Drago, J. Organometal. Chem., 5, 330 (1966). (4) P. L. Goggin and L. A. Woodward, Trans. Faraday SOC.,56, 1591 (1960). (5) G. Costa, Ann. Chim. (Rome), 40, 559 (1950). (6) D. Goddard and A. E. Goddard, J. Chem. SOC.,1922,256. (7) E. Krause and A. V. Grosse, Ber., 58B, 272 (1925).
1286
ANALYTICAL CHEMISTRY
Corp.), and Triton X-100 (Rohm and Haas) were used as received. All other reagents were of ACS reagent grade. Supporting electrolyte solutions covering the pH range from p H 2.5 to p H 9.0 were prepared using appropriate amounts of phosphoric acid, potassium dihydrogen phosphate, and sodium hydroxide. All solutions were prepared using distilled water and contained 0.001 Triton X-100 as a maximum suppressor. Nitrogen was bubbled into all solutions in order to remove oxygen. Apparatus and Procedure. A Wenking potentiostat equipped with a conventional operational amplifier integrator to scan the potential was employed for polarographic measurements. Polarograms were recorded on a Honeywell XY Recorder. The dropping mercury electrode had a capillary flow rate of 1.64 mg/sec a t a mercury column height of 61.0 cm. Coulometric integration was performed with an operational amplifier difference integrator (8). Cyclic voltammograms were obtained at a slow dropping mercury electrode (400 sec drop time) or with a hanging mercury drop electrode (Brinkmann Instruments, Westbury, N. Y . ) with a Heath polarograph. A Hewlett-Packard 202A function generator was used to generate triangular wave potential scans for the Heath polarograph. A Tektronix oscilloscope or the XY recorder, depending on the sweep speed, was used to record the cyclic voltammograms. Three-electrode geometry was used in all electrochemical measurements. Temperature control for polarographic and cyclic voltammetric electrolyses was maintained by employing a waterjacketed cell, thermostated at 25.0 i 0.1 “C. Large scale coulometric electrolyses were run at room temperature. The saturated calomel reference electrode was placed in a separate compartment from the sample cell with electrical contact established through a cracked glass frit. A platinum wire counter electrode was employed. A mercury pool cathode was used for coulometry. A spiral platinum wire, contained in a separate compartment with electrical contact made through a fritted glass disk, served as the counter electrode for these measurements. All potentials are reported GS. the saturated calomel electrode. Well-defined polarograms are not obtained below the concentration limit 0.1 x lO-3M (C6H5)2Tl+. The upper concentration limit, 1.3 x lO-3M, is determined by the solubility of (CsH&T1+ in phosphate buffers. RESULTS
In the accessible concentration range, 0.1 X M to 1.3 X lO-SM, three polarographic waves are obtained. Figure 1 shows a n example of a polarogram obtained at 0.1 X 10-3M, and Figure 2 shows a typical polarogram obtained at 0.9 x lO-3M (CsH&Tl+. The limiting currents of waves 1 and 2 are diffusion-controlled, because plots of diffusion current us. the square root of the mercury column height are linear. At 0.1 x lO-3M the diffusion current on the third wave is independent of the mercury column height,
(8) George A. Philbrick Researches, Inc., “Philbrick Applications Manual-Computing Amplifiers,” Nimrod Press, Inc., Boston, Mass., 1966, p 46.
1.2
1.0 0.8
0.6
0.4
0.2
0.0
0.0
-0.2
Figure 1. Polarogram of 0.1 Capillary flow rate
= 0.95
-0.4
x
-0.6
-0.6
-1.2
-1.0
-1.4
-1.6
+ HP042-)a t p H 6.2
1O-V (CsH&TI+in 0.05F (H2P04-
mg/sec. 0.001%Triton X-100 present as maximum suppressor
10.0
L
-1
,I
8.0
6.0
4.0
2.0
0.0
1
1 0.0
1
1
1
-02
1
1
-0.4
1 -0.6
1
1 -0.8
1
1
1
-1.0
I -1.2
Figure 2. Polarogram of 0.9 X 10-3F (CeHs),TI+in 0.05F (H2P0,Capillary flow rate
=
0.95 mg/sec. 0.001
indicating a kinetic-controlled process for the third wave at this concentration in which the current is controlled by a chemical reaction occurring a t the electrode. Plots of log (idlid - i) us. E for the first and second waves in the concentration range studied are linear with slopes of 83 and 72 mV, respectively, which indicate that these reactions are irreversible. A similar plot for wave three yields a slope of 52 mV indicating electrochemical reversibility for the reaction corresponding to wave three.
I
I -1.4
I
1
1
1.6
+ HPO,*-) at p H 6.2
Triton X-100present as maximum suppressor
Table I shows the dependence of diffusion current on concentration of (C@H&Tl+for all three waves. The diffusion currents of the first and second waves are proportional to the depolarizer concentration over the range 0.1 X 10d3M to 1.3 x 10-3M. The diffusion current of wave three increases slowly at first but reaches a constant value as the concentration of (CsH&Tl+ is increased. The diffusion current on the second wave is slightly higher than the diffusion current on the first wave. This observaVOL. 40, NO. 8, JULY 1968
1287
tion is explained by the fact that on the first wave the current, as a function of time, does not increase monotonically, but reaches a maximum value during the lifetime of the drop and then decreases as the drop continues to grow. Thus the current-time curves on each drop are not t’” parabolas as predicted by classical theory. Instead the current o n the first wave increases less rapidly than t’I6. Therefore the diffusion current on the first wave is slightly lower than the diffusion current on the second wave. The shape of the current-time curves is determined by the potential dependent adsorption of both the parent cation and the intermediate species at the electrode surface (see discussion of electrocapillary curves below). Because the current-time curves are
Table I.
Variation of Diffusion Current with Concentration of Diphenylthallium(II1) Cationa
Conc. (CeH&Tl+ (mmoles/l)
Wave 1
Wave 2
Wave 3
0.13 0.25 0.51 0.76 1.02 1.27
0.42 0.93 1.80 2.75 3.80 5.20
0.52 1.20 2.55 3.90 4.90 6.20
0.38 0.45 0.45
id
(PA)
0.60 0.60 0.60
pH = 6.2; capillary flow rate = 1.64 mg/sec. 0.001% Triton X-100 present as maximum suppressor. 0
Table 11. Variation of Half-Wave Potential with Concentration of Diphenylthallium(II1) Cation”
Conc. (CsHs)zTl+ (mmolesil)
Wave 1
Wave 2
Wave 3
0.13 0.25 0.51 0.76 1.02 1.27
0.52 0.54 0.59 0.61 0.62 0.64
1.03 1.00 0.97 0.96 0.94 0.92
1.12 1.22 1.30 1.33 1.36 1.36
- E1/z(V
US.
SCE)
0 pH = 6.2; capillary flow rate = 1.64 mg/sec. 0.001 X-100 present as maximum suppressor.
Triton
Table 111. Variation of Half-Wave Potential with pHo
Conc. (Ce.H5)zTl+ (mmoles/l) A) 0.13
-
PH 3.1 4.3 5.2 6.2 7.0
B) 1.01
8.0 9.0 2.5 3.3 4.2 5.1 6.2 7.0 8.0
-E‘/z (V
Wave 1 0.52 0.52 0.52 0.52 0.52 0.52 0.52 0.62 0.62 0.62 0.62 0.62 0.62 0.60
SCE) Wave 2
US.
b
1.00
1.08 1.11 1.12 1.15 1.14 1.14 1.24 1.26 1.30 1.33 1.36 1.42 1.43
1.02 1.03 1.04 1.03 1.03 0.85 0.86 0.90
0.92 0.94 0.94 0.94
a Capillary flow rate = 1.64 mg/sec. 0.001 present as maximum suppressor. b Waves 2 and 3 not clearly separated.
1288
ANALYTICAL CHEMISTRY
Wave 3
b
Triton X-100
distorted by adsorption, undamped polarograms were recorded and maximum rather than average currents were measured. Table I1 shows the variation of half-wave potential with concentration of (C6H&T1+. As the concentration is increased above 0.1 x 10-3M, the half-wave potential of the second wave shifts to more anodic values, and the half-wave potentials of the first and third waves shift to more cathodic values. The dependence of the half-wave potential on pH is shown in Table I11 for 0.1 X 10e3M and for 1.0 X 10-3M (C6H5)zTI+. In both cases the half-wave potential of the first wave is independent of pH. The half-wave potentials of the second and third waves are p H dependent, indicating that hydrogen ions are involved in the reactions occurring at potentials on the second and third waves. There is also a slight dependence of the half-wave potential of the third wave on the drop time in the concentration range studied. An electrocapillary curve of a solution of (C6&)zTl+, 0.9 X 10-3M, is shown in Figure 3. Depressions corresponding to adsorption of the parent cation and of intermediate species at the electrode surface are observed. Strong adsorption occurs at potentials more negative than 0 V us. SCE. Inorganic and most organic cations are desorbed at potentials positive to the potential of zero charge (PZC). They are capillary-inactive in this region--i.e., they d o not lower the surface tension below the values obtained with a solution of the supporting electrolyte alone (9). Many organic cations are adsorbed at potentials negative to the PZC, but only aromatic cations are adsorbed at potentials positive to the PZC (10). As a result of the adsorption of an aromatic cation on the mercury surface, the electrocapillary curve lies below the curve obtained with a solution of the supporting electrolyte alone. As shown in Figure 3, a lowering of the drop time (surface tension) on the anodic side of the PZC is indicative of the adsorption of the (C&15)2T1+ at the mercury surface. Interaction between the pi-orbital electrons of the phenyl rings and the positively charged electrode surface causes strong adsorption on the anodic side of the PZC. Adsorption of linear CoH5-T1+-CsH5 occurs with the phenyl rings in a flat orientation against the electrode surface because of this pi-orbital interaction with the electrode. More pronounced lowering of the drop time in the regions of stability of the intermediate species is indicative of the stronger adsorption of these species than of the parent cation. Depressions in the potential region between - 0.60 and - 1.30 V indicate extensive adsorption of the intermediates of the reduction processes. Desorption occurs at -1.35 V, the half-wave potential of the third wave. The adsorption phenomena at the electrode surface lead to distortion of the current-time curves for individual mercury drops as mentioned above. Table IV shows the results of coulometric determinations of the apparent number of electrons transferred in the electrode processes. A value of n = 1.20 f 0.02 for the first wave is obtained over the entire range of ( C I H B )Tl+ ~ concentrations investigated. That n is significantly greater than 1.00 implies the presence of side reactions, discussed below, which yield electroactive products. The second wave yields a n apparent n value of 2.00 i- 0.02
(9) I. M. Kolthoff and J. J. Lingane, “Polarography,” Vol. 1, Interscience Publishers, Inc., New York, N. Y.,1952, p 141. (10) B. E. Conway and R. G . Barradas, Electrochim. Acta, 5, 319 (1961).
3.6
t
-
for all concentrations in the range 0.4 X lO-3M to 1.3 X 10-3M. Determination of apparent n values for the second wave below a concentration of 0.3 X lO-3M were not attempted because of the ill-defined nature of this wave below this concentration. As the concentration is increased from 0.1 X 10-3M to 1.0 x 10-3M, the apparent n value for the third wave decreases rapidly and approaches a lower limit of n = 2.10, as shown in Table IV. The expected value of n = 3.00 for the third wave at infinite dilution is obtained by extrapolation. The results of the analyses of the products formed during controlled potential electrolysis at a mercury pool cathode are shown in Table IV. Coulometry at potentials corresponding to limiting-current plateaus of all three polarographic waves at all concentrations of (C6H&T1+ yields one mole of thallium
\I \-
for each mole of (C6Ha)2TI+reduced. For the first wave slightly less than one mole of diphenylmercury, ( c ~ H s ) ~ H gis, formed for each mole of (C6Hs)2Tl+taken. Identification of the (C6H5)2Hg was made by comparison of its infrared spectrum with that of an authentic sample of (c~,Hs)nHg. Electrolyses at a potential corresponding to the second diffusioncurrent plateau were not attempted below 0.3 x lO+M because of the ill-defined nature of this wave below this concentration. Above 0.3 x lO-sM, in addition to the thallium produced, 0.5 moles of (CeH&Hg are formed for each mole of (C6H&TI+ reduced. Benzene, (C6H6),is also formed in this reduction. For the third wave at concentrations less g is too than 0.1 X 10-3M the amount of ( c 6 H ~ ) ~ Hformed small to determine accurately and principal products are benzene and thallium. Benzene was identified by its ultra-
Table IV. (C&)zTl+ taken. (mmoles X loz) 2.76 f 0.01 2.76 f 0.01 0.60 f 0.01 0.92 f 0.01 1.22 f 0.01 1.83 f 0.01 2.76 f 0.01 3.00 f 0.01 4.60 f 0.01 5.25 f 0.01
Coulometry of Diphenylthallium(II1) Cation in pH 6.2 Buffer (CsHs)*Hg found* -E (V US. SCE) napparen (mmoles X lo2) 0.8W 1.20 2.29 f 0.01 1.10f 2.00 1.33 f 0.06 d 1.@ 2.62 1,400 2.49 0.31 f 0.02 d 1.40Q 2.39 d 1.409 2.29 1.400 2.20 0.83 f 0.06 d 1.400 2.19 1.400 2.18 1.63 f 0.03 d 1.40g 2.16
TI" found. (mmoles x 10%) 2.80 f 0.06 2.85 f 0.06 d
0.93 f 0.02 d
d
2.85 f 0.06 d
4.80 f 0.10 d
50-ml sample. Determined by atomic absorption spectrometry. c Determined by coulometry. d Not determined. Wave 1. Wave 2. 0 Wave 3. (I
6
' 4
VOL. 40, NO. 8, JULY 1968
1289
6 -
1
1
1
1
1
1
1
1
1
1
1
1
1
1
1
1
Sweep reversal shows no evidence for electrochemical reversibility of this step. On successive sweeps, a peak current on the cathodic scan at -0.49 V (Peak 1) and a corresponding anodic peak at -0.43 V on the reverse scan are observed. This set of reduction-oxidation peaks indicates a reversible electrode process and is caused by the oxidation of thallium formed during the electrochemical reduction of (C&&Tl+ and the subsequent reduction of the thallous ion so formed. Peak 2 corresponds to polarographic wave 1 in Figures 1 and 2. If the scan rate is increased to 20 V/sec, a third cathodic peak appears at -1.34 V. A corresponding anodic peak separated by 60 mV from Peak 3 is also observed, indicating electrochemical reversibility for this step. If the scan rate is further increased to 40 V/sec, a voltammogram such as that shown in Figure 5 is obtained. The voltammogram is somewhat distorted because of the presence of uncompensated resistance in the fine capillary of the HDME. The cathodic peak current increases approximately linearly with scan rate. The ratio of anodic peak current to cathodic peak current increases slowly with increasing scan rate, but always remains below unity. The high background correction makes the data imprecise to analyze by standard methods (ZZ, 12) and it is impossible to tell whether Peak 3 is caused by reduction of an adsorbed film of material or is largely a manifestation of the rather large change in capacity observed in this potential region (see electrocapillary curve, Figure 3). As one would not expect an anodic wave corresponding to the third polarographic wave (see Equation 5 , below), it is probable that Peak 3 contains a large capacitative component.
-
c
c Q
L L
-2
a
0
-4
-6 -8
-1.0
-0.5
0.0
'4.5
Figure 4. Cyclic voltammogram of 0.6 X 10-SF (Cd-I&Tl+ at pH 6.2 Scan rate = 0.31 V/sec. 0.001 % Triton X-100present as maximum suppressor violet spectrum. Even though the electrolyses were carried out in a sealed cell, the recovery of C6Hewas not quantitative. As the concentration is increased above 0.1 x 10-3M the amount of (C&)zHg formed increases, but less than 0.4 moles of (C6H&Hg are formed for each mole of (C&,)zTl+ taken. These results indicate that two competing reactions occur at potentials on the third wave, and that the rates of these competing reactions are concentration dependent. The results of cyclic voltammetry at a slow dropping mercury electrode are shown in Figures 4 and 5 . Figure 4 shows a voltammogram obtained on a solution of (C6H&T1+, 0.6 x lO-3M, at a scan rate of 0.31 V/sec. During the first cycle only a cathodic peak at -0.63 V (Peak 2) is seen.
CONCLUSIONS
The following reaction scheme is proposed for the polarographic reduction of (CsHs)zTl+. At 0.1 X 10-3M the first electron transfer yields an adsorbed diphenylthallium(I1) (11) R.H.Wopschall and I. Shain, ANAL. CHEM., 39, 1514 (1967). (12) R. S. Nicholson and I. Shain, ibid., 36, 706 (1964).
96
64
2 32 3 c. +-
c0
o
L L
2 -32 -6 4 -9 6 l
1
0.0
1
1
1
1
1
1
1
-I .o
-0.5
Figure 5. Cyclic voltammogram of 0.6
x
I
1~ 1 ~ 1
~ 1
1
-I.5
10-SF (CgH5)*T1+at pG 6.2
Scan rate = 40 V/sec. 0.001% Triton X-100present as maximum suppressor
1290
ANALYTICAL CHEMISTRY
1
1
radical as shown in Equation 1. The cation is shown initially adsorbed, as required by electrocapillary evidence.
Hg,"
'Hg'
Transmetallation with the mercury electrode yields (C6Hs)zHg according to Equation 2. T1 , ,
\\
dsHs \\
C6H5
-.t
(C6Hs)aHg f TI"
(2)
,
I
Hg At least two side reaction sequences which would give apparent n values greater than 1.00 for the first wave are possible. Disproportionation of diphenylthallium(I1) radicals would yield diphenylmercury, thallium amalgam, and triphenylthallium. Triphenylthallium is readily hydrolyzed to benzene and diphenylthallium cation, which would be further reduced. However, such a disproportionation scheme is expected to give rise to a concentration dependent value of n, because disproportionation is second order in the thallium species. Hydrolysis- of the diphenylthallium radical t o give benzene and a transient inorganic thallium(I1) species is another possible side reaction. Any thallium(I1) formed would either be directly reduced to the metal or would decay to thallium(1) and thallium(III), which are also reducible at the control potential employed. Other possible sequences of side reactions can be formulated, but the available evidence is insufficient to distinguish among them. The second electron transfer yields an adsorbed phenylthallium(1) as shown in Equation 3.
T1
Tl'k\ I
\
d6Hs
C6H5
H + ,I
f e-
+
CsHs f C6Hs
(3)
must be applied before the rate of electrochemical reduction (Equation 5 ) approaches the rate of disproportionation. Benesch and Benesch (13) observed a similar change in half-wave potential with concentration in the polarographic reduction of organomercury compounds. Morris, McKinney, and Woodbury (14)reported similar behavior in the electrochemical reduction of tetraphenylantimony(V) cation. The second and third electron transfers, Equations 3 and 5 , show the involvement of hydrogen ions in the reaction scheme and explain the effect of p H on the half-wave potentials of these waves. The proposed reaction scheme explains the variation of the apparent n value for the third wave with concentration. At infinite dilution the adsorbed C&sTl(I) intermediates are far apart. The rate of disproportionation is very slow, whereas the rate of electrochemical reduction (Equation 5) is very fast, and an apparent n value of 3.00 would be observed. As the concentration is increased, the rate of disproportionation increases with the square of the concentration and the rate of electrochemical reduction decreases. As the concentration is increased further, competition between the two reaction paths, disproportionation and reduction, results in an apparent n value much lower than in the case where electrochemical reduction alone occurs--i.e., n = 3.00. The bridged transition state, TI /
/
solvent
solvent
Hg
\
, /
Hg
Rapid disproportionation yields (C6H5)2 Hg according to Equation 4.
//
\\
CsH5
C6H5 \
T1
\
/
/
203Hg (4)
Hg The third electron transfer corresponds to Equation 5 . T1 H+ d6H5
\
is similar to that proposed by Reutov and coworkers (15) for the isotopic exchange reactions of organomercury salts with metallic mercury. Pollard and Westwood (16) postulated a similar transition state for the isotopic exchange of (CsH&Hg with metallic mercury, \
Hg,'
\
f e-
+
C6H6 f TI" f Hg
(5)
Hg At 1.0 x lO-3M, the first and second electron transfers proceed as in Equations 1 t o 4 above. For the third electron transfer, the rate of disproportionation of adsorbed CeH5T1(I)t o (C6H&Hg (Equation 4) is fast compared to the rate of reducton to C&s and TI" (Equation 5 ) . This reaction scheme explains why the half-wave potential of the third wave shifts to more negative values as the concentration is increased. Because the rate of disproportionation (Equation 4) is proportional to the square of the concentration of the adsorbed CsH5Tl(I), a n increase in the concentration will increase the rate of disproportionation. Therefore, as the concentration is increased, a more negative potential
in which the labelled mercury metal is completely substituted by a n SEt process during which any one phenyl ring is never detached. There have been several reports in the literature of systems involving transmetallation of organometallic compounds with mercury during electrochemical reductions. Morris, McKinney, and Woodbury (14)reported arylation of the mercury electrode with aqueous solutions of reduced tetraphenylantimony(V) cation to yield ( c ~ & ) ~ H gand (CsHs)3Sb. Dessy and coworkers (17) studied the electrochemical behavior of triphenyllead acetate and diphenyllead diacetate in dimethoxyethane and reported arylation of the mercury elec(13) R. Benesch and R. E. Benesch, J . Amer. Chem. SOC.,73, 3391 (1951). (14) M. D. Morris, P. S. McKinney, and E. C. Woodbury, J . Electroanal. Chem., 10, 85 (1965). (15) 0. A. Reutov, P. Knoll, and U Ian-Tsei, Dokl. Akad. Nauk. SSSR,120, 1052 (1958). (16) D. R. Pollard and J. V. Westwood, J . Amer. Chem. Soc., 88, 1404 (1966). (17) R. E. Dessy, W. Kitching, and T. Chivers, ibid., p 453. VOL. 40, NO. 8, JULY 1968
1291
trode by products of the electroreductions to yield (C&)zHg in both cases. The transmetallation of (C6H5)2T1Brwith mercury has been observed in a nonelectrochemical system by Gilman and Jones (18). They reported an excellent yield (90%) of (C6Hs)2Hgafter refluxing in pyridine for eight hours. The polarographic results indicate that the intermediate formed during the second electron transfer (Equation 3) is an adsorbed phenylthallium species, C6HsT1, and not a phenylmercury radical, C6HsHg* . In the polarographic reduction of phenylmercury ions, both the formation of C6H5Hg. and its subsequent reduction to C6H6and Hg are irreversible (13, 19). Because the polarographic reduction to C6Hs and T1" (Equation 5) is reversible, it is an adsorbed C6HsT1 species, not C6HsHg., which undergoes rapid disproportionation to
(C6H&Hg (Equation 4) and reduction to C6H6and TI (Equation 5). Gilman and Jones (20) presented evidence for the formation of phenylthallium by pyrolysis of (C6Hs)3TI in xylene. An attempted electron paramagnetic resonance study to determine the nature of the intermediates formed during electrochemical reduction was unsuccessful. The formation of an insulating layer of (C6H5)2Hgprecipitate on the mercury electrode surface contained in an EPR aqueous flat cell inhibited the electrochemical reduction after a few seconds and no signals were observed. Stirring the mercury layer to expose fresh electrode surface to the depolarizer is not possible in such flat cells. It is possible that EPR studies of the reaction intermediates might be more successful in solvents in which diphenylmercury is soluble.
(18) H. Gilman and R. G. Jones, J. h e r . Chem. SOC.,61, 1513 (1939). (19) R. F. Broman and R. W. Murray, ANAL.CHEM.,37, 1408 (1965). (20) H. Gilman and R. G. Jones, J . Amer. Chem. SOC.,62, 2357
RECEIVED for review February 29, 1968. Accepted May 2, 1968. Presented in part, Division of Fuel Chemistry, 153rd ACS Meeting, Miami Beach, Fla., April, 1967. Work supported in part by National Science Foundation Grant GP6164.
(1940).
Potentiometric Determination of Boron as Tetrafluoroborate R. M. Carlson Department of Pomology, Unioersity of California, Dacis, Calif. 95616
J. L. Paul Department of Environmental Horticulture, Unicersity of California, Davis, Calq. 95616 The potentiometric determination of tetrafluoroborate with a liquid ion exchange membrane electrode i s described. The electrode can be used for tetraInterfluoroborate concentrations down to lO--6M. ference from several anions was estimated. Nitrate and iodide cause the greatest interference of those anions studied. Rate of formation of tetrafluoroborate in solution at 24 and 60 OC was studied. Complete formation of tetrafluoroborate was obtained in 0.28M hydrofluoric acid at 60 OC in 5 minutes. At lower temperatures and lower hydrofluoric acid concentrations more than 6 hours were required for complete formation. Using columns of the boron specific resin Amberlite XE-243 boron can be separated from interfering anions and the tetrafluoroborate ion can be formed in less than 15 minutes with a small volume of 10% hydrofluoric acid. The tetrafluoroborate is eluted with sodium hydroxide and measured potentiometrically. Application of the method to determination of boron in water samples is illustrated.
THE RECENT DEVELOPMENT of liquid ion exchange membrane electrodes that are selective for nitrate and perchlorate ions (1) suggests the possibility of a similar electrode that would respond to the tetrafluoroborate ion. The behavior of the three ions is similar in many ways. All three form sparingly soluble salts with nitron. Perchlorate and nitrate extract as tetrabutylammonium complexes into methyl isobutyl ketone along with tetrafluoroborate (2). It seemed plausible that the type of membrane used for the perchlorate and nitrate electrodes would be suitable for tetrafluoroborate. (1) G. A. Rechnitz, Chem. Eng. News,45 (25), 146 (1967). (2) W. J. Maeck, M. E. Kussy, B. E. Ginther, G. V. Wheeler, and J. E. Rein, ANAL.CHEM., 35, 62 (1963).
1292
ANALYTICAL CHEMISTRY
Current methods for determining boron are time-consuming and subject to interferences. A rapid potentiometric method sensitive to microgram quantities of boron would have application for boron determinations in many materials. This report deals with a liquid ion exchange membrane electrode sensitive to tetrafluoroborate ion. A procedure for isolating boron and rapidly forming the tetrafluoroborate ion on columns of the boron specific resin Amberlite XE-243 is described. EXPERIMENTAL. Apparatus. E M F measurements were made with an Orion Model 801 digital pH meter. All solutions were stirred during measurement of potentials. As the work was performed in a laboratory with a temperature of 24 f 0.5 "C, no further effort was made to control temperature. The tetrafluoroborate-sensitive electrode was prepared by modifying a nitrate-sensitive electrode manufactured by Orion Research, Inc. (Model 92-07). The liquid ion exchanger was converted to the tetrafluoroborate form by twice shaking 2 ml of exchanger with 50 ml of a molar solution of boric acid in 4 molar hydrofluoric acid and decanting the aqueous phase. The composition of the internal filling solution was 10-*M H3B03, 0.28M HF, and 10-*M KCI. The electrode was assembled according to manufacturer's directions. Electrode Sleeve. Because all solutions contained HF, a plastic sleeve was used to protect the saturated calomel reference electrode from attack. The sleeve was prepared by gently heating a '/*-inch i.d. polyethylene tube and drawing out a constriction. A single strand of size 50 mercerized cotton thread was passed through the constriction and sealed in place by reheating and twisting the tubing. The twisted
