In view of the lower background level in the Beckman oxyhydrogen flame (see Figure 1) and the recently reported claim (34) that this mixture provides the best prospects for analytical purposes, it is of interest to compare several detection limits observed in this flame with those reported in Table I. Four typical elements (Co, Cu, Fe, Ni), which do not form stable monoxides and for which an oxy-hydrogen flame may be advantageous, were selected. All defining flame parameters were optimized for the oxy-hydrogen as well as for the oxy-acetylene flame. The data summarized in Table VI show that the interconal zone, oxy-acetylene spectra provided slightly superior detection limits and thus do not support Britske’s conclusions (34). Our measured values for the oxyhydrogen flame are typically an order of magnitude poorer than those reported by Gilbert (35). However, Gilbert’s criteria for detection limits differed markedly from ours. His values were recalculated from many sources. Also Gilbert based his detection limit on the concentration which “yields a net response above background of 1 mV across the load resistor when the background signal is at most 100 mV and amplification and flame adjustments are optimal.” This definition neglects the statistical fluctuations in the background signal. The erroneous assumption is occasionally made (32) that the emission spectrometric detection of low concentrations of the elements is limited strictly by the line~~
~
~~~
(34) M. E. Britske, Zuvodsk. Lab., 30, 1465 (1964). (35) P. T. Gilbert, Jr., in “Symposium on Spectroscopy,” ASTM Spec. Tech. Pub[. No. 269, Philadelphia, Pa., 1959, p. 73. (36) H. Kaiser, 2.Anal. Chem.,209, l(1965). (37) H. Kaiser, Ibid., 216, 80 (1966).
Table VI. Comparison of Several Flame Emission Detection Limits, Ngiml Turbulent
Element co
cu
Fe
Ni
Premixed oxy-hydrogen flame oxyIowa State Wavel@gth, acetylene University Literature A flame value (35) 3453.50 3247.54 3273.96 3719.94 3859.91 3524.54
1 0.1 0.2 0.7 0.6
2 0.4 0.6
2 1
0.2 0.02 0.02 0.1 0.2 0.12
to-background ratio. There is adequate documentation in the literature (4,19, 36,37) that the limiting factor is the ratio of line signal to statistical fluctuations in the background signal (spectral background fluctuations and detector noise). Because the detection limits reported in Tables I and VI take account of these statistical fluctuations, they should provide more realistic values than those recalculated on a less rigorous basis. ACKNOWLEDGMENT
The authors gratefully acknowledge the experimental assistance of Constance C. Butler and the valuable suggestions of Richard N. Kniseley. RECEIVED for review September 21, 1966. Accepted January 25, 1967. Contribution No.1969. Work was performed in the Ames Laboratory of the US. Atomic Energy Commission.
Polarography of Aquodiethyltin(1V) Cation in Aqueous Solution Michael D. Morris Department of Chemistry, The Pennsylvania State University, University Park, Pa. 16802 The polarographic reduction of aqueous solutions of the aquodiethyltin(1V) cation, (C,H6)2Sn(H,0)4+z, and its hydrolysis products proceeds by a reversible twoelectron transfer. The final product is a polymeric, probably cyclic, diet hylti n( II). Polarographic waves are distorted by the formation of the insoluble, polymeric reduction products on the electrode surface. Good polarograms can be obtained only at concentration levels of 1 x lO-4F aquodiethyltin(1V) cation and below. The pH dependence of the half-wave potential of the reduction yields values of pK, = 3.0 =k 0.1 and pK, = 5.4 st 0.2 for the first and second acid dissociation constants of the aquodiethyltin(1V) cation. From polarographic measurements of the formation constant for the chloroacetate complex, log Kf = 1.9 f 0.1. For the chloro complex, polarographic measurements yield log K, = 0.4 f 0.2. The pH dependence of at constant concentration (1 x lO-‘F) and the concentration dependence of E,,, at constant pH are consistent with a two-electron reversible reduction to an insoluble product. Plots of log ( i d - i) VI. Edrne have slopes of 31-39 mV. The abnormally high slopes are due to a decrease in the effective area of the electrode during the course of the polarogram by formation of an insoluble polymer on the electrode surface.
of a continuing investigation of the aqueous solution electrochemistry (1) and electroanalytical chemistry (2) of organometallic compounds, we have studied the polarography
As
PART
476
ANALYTICAL CHEMISTRY
of the aquodiethyltin(1V) cation, (CzH6)L3n(Ho0)4+2, in aqueous solution. Although this system appears to have no analytical utility, it is of interest here because of its relationship to the polarography of organometallic ions of analytical importance, such as the tetraphenylantimony and tetraphenylarsonium cations. Riccoboni and Popoff (3) investigated the polarography of diethyltin dichloride in unbuffered aqueous solutions of potassium perchlorate and potassium chloride. They concluded that the primary electrode process is reversible and is given by Equation 1.
+
Et2Snf2 2e- = EtzSn:
(1)
The diethyltin radical formed was presumed to undergo slow polymerization after the electron transfer step. Riccoboni and Popoff (3)found that the limiting current for the reduction process increases linearly with concentration up to about 5 X 1O-T EtzSnCl2and then remains fairly constant. They
(1) M. D. Moms, P. S. McKinney, and E. C . Woodbury, J . Electroanal. Chem., 10, 85 (1965). (2) M. D. Morris, ANAL.CHEM., 37,977 (1965). (3) L. Riccoboni and P. Popoff, Atti. Zst. Veneto Sci., Venezio, 107,Pt. IZ, 123 (1949); C.A.,44,6752a (1950).
I
I
I
1
I.6
I
I
1.4 12 I.o
q
0.8
*$ 0.6 0.4
02 0.0
-02
-0J
-0.2 4.3 -0.4 0 . 6 -0.1 0 . 7 -0,s -0.9 -1.0 -1.1 Edme,V US. SCE
-1.2 -13
Figure 1. Polarogram of 1.01 X 10-4Faquodiethyltin(1V) cation in a supporting electrolyte of 0.1FKCI, 0.1F HCl. 0.001 Triton X-100added as maximum suppressor
suggested that this behavior is the result of the formation of a layer of insoluble, polymeric diethyltin, through which electron transfer cannot occur. Saikina and coworkers also noted the polarographic reducibility of diethyltin in strongly acidic aqueous ethanol (4) and aqueous (5) media. Their observations on this compound are included in a. general survey of the polarography of organotin compounds. Recently Dessy and c:oworkers (6) have investigated the behavior of organometallic compounds in dimethoxyethane. They find that in this scllvent diethyltin dichloride gives a twostep wave and they propose the mechanism of Equations 2 to 4.
+ e-
Et2SnClz
= Et2SnCl'
(2)
This reduction is followed by a dimerization. 2EtzSnC1 = EtzCISnSnCIEtz
(3)
The second reduction wave is given by Equation 5 . EtSnCl'
+ e-
= EtzSn:
(4)
The final product is a polymeric diethyltin(I1). The structure and hydrolysis of organometallic cations in aqueous solution has been reviewed by Tobias (7). Dialkyltin(1V) salts dissociate in aqueous solution to give cations of the form RzSn(Hz0)4+z. The aquo ions may lose up to four protons, presumably one from each of the four coordinated water molecules. For the aquodiethyltin(1V) cation, only the (4) V. S. Toropova and hd. K. Saikina, Sb. Statei Obschei Khim. Acad. Nauk SSSR.,1, 210 (1953); C.A., 48, 12579g (1954). (5) M. K. Saikina, Uch. Zap. Karansk. Cos. Uniu. 116, 129 (1956); C.A., 51, 7191e (1957). (6) R. E. Dessy, W. Kitching, and T. Chivers, J . Am. Chem. Soc., 88, 453 (1966). (7) R. S. Tobias, OrganometaN. Chem. Rev., 1,93 (1966).
first acid dissociation constant has been reported (8),pK1 = 3.40 in 3FNaC104. I n addition to undergoing simple acid-base behavior, the partly deprotonated species can polymerize. The most important of these reactions is the dimerization of RzSn(H2O)OH+, given by Equation 5 . 2R2Sn(HZO)OHf = [R2Sn(H20)0Hf2I2
(5)
For Et2Sn(HrO)OH+ dimerization, Kd = 259 in 3F NaC104 (8). A large number of polymeric substances with the general formula [(GHa)2Sn]nhave been prepared. Recent studies on polymeric dimethyltin (9),diethyltin (IO)and diphenyltin (11) have shown that these polymeric tin(I1) compounds are either puckered rings of 6 to 9 RzSn units or long chain (10 to 20 units) compounds of the type R$n(SnRz),SnR3 or HR2Sn(SnRzIn SnR2H. The product or products obtained depend upon the methods of preparation. As Olson and Rundle point out (II), a monomeric RzSnwould have a vacant valence orbital and would be expected to be unstable. EXPERlMENTAL
Apparatus. A Wenking potentiostat was employed for coulometry and as a three-electrode polarograph. For this purpose a conventional operational amplifier integrator was used to apply the scan potential. Polarograms were recorded with a Honeywell strip chart recorder. Coulometric integration was performed with a Lingane titration coulometer (12) in series with the coulometric cell. (8) R. S. Tobias, H. N. Farrer, M. B. Hughes, and B. A. Nevett, Znorg. Chem., 5, 2052 (1966). (9) T. L. Brown and G. L. Morgan, Znorg. Chem., 2,736 (1963). (10) W. P. Neuman and J. Pedain, Ann. Chem., 672, 34 (1964). (11) D. H. Olson and R. E. Rundle, Inorg. Chem., 2, 1310 (1963). (12) J. J. Lingane and L. A. Small, ANAL.CHEM., 21,1119 (1949). VOL. 39, NO. 4, APRIL 1967
477
A water-jacketed polarographic cell, thermostated at 25.0 0.1’ C was employed. The reference saturated calomel electrode was placed in a separate compartment, behind a cracked glass seal. The counter electrode, a platinum wire sealed into glass tubing, was in the same compartment as the dropping mercury electrode. Controlled potential electrolyses were run at room temperature (25-30” C) using a large mercury pool cathode. The counter electrode was placed in a separate compartment, with electrical contact made through a fritted glass disk in the conventional manner. Reagents. Diethylin dichloride, m.p. 84-85 O C (City Chemical Corp.) was used as received. Luijten and van der Kerk (13) report m.p. 84.5-86’ C. Triton X-100 (Rohm and Haas) was used as received. All other reagents were of ACS reagent grade. Distilled water was used to prepare all solutions. The pH range from pH 0.5 to pH 10 was spanned with hydrochloric or perchloric acids, acetic acid, pyridine, tris(hydroxymethyl) aminomethane, boric acid, and bicarbonate buffers. Except in the case of the mineral acids, the total concentration of buffer components was 0.05F. Potassium chloride or sodium perchlorate was added to maintain an ionic strength of 0.1 to 0.2. All solutions contained 0.001 % Triton X-100 as a maximum supressor. &
RESULTS AND DISCUSSION
Polarography. A typical polarogram is shown in Figure 1. Similar polarograms are obtained in other supporting electrolytes, Well-defined polarograms are obtained for solutions of 1 X 10e4Faquodiethyltin(1V) cation. At 2 X 1 O - T sufficient distortion is present to prevent reliable half-wave potential measurements. At 5 x 10-4F and at higher concentrations, distortion is so severe that the wave appears to be split into several steps and the limiting current plateau is illdefined. At all concentrations, current-time curves for individual drops do not increase monotonically, but reach a maximum during the drop life and then decay as the drops continue to grow. Such behavior is characteristic of an electrode process in which an insoluble, insulating film is formed on the electrode surface (14). There is significant adsorption of both aquodiethyltin(1V) cation and its reduction products, as demonstrated by the depression of the electrocapillary curve of 0.1F HCl, 0.1F KCl in the presence of 1 x lO-4F aquodiethyltin(1V) cation (Figure 2). While adsorption may contribute to the distortion of the polarographic waves, the most important cause of this distortion appears to be the formation of insulating reduction products. Figure 3 shows the variation of limiting current (maximum value attained during drop life) with concentration of aquodiethyltin(1V) cation for reduction in 0.1F HC1, 0.1F KC1. The same results are obtained in other supporting electrolytes. The limiting current rises linearly with depolarizer concentration up to about 2 X 10-4F. The rise is less than linear up to about 4 x 1 O - Y and very slow beyond about 6 X lOW4F. This is the behavior expected when a film of insulating material is formed as the reduction product. Beyond a certain critical concentration, at which point the electrode is completely covered with the film, the limiting current increases slowly, if at all, with increasing depolarizer concentration.
Riccoboni and Popoff (3) report no increase in limiting current with concentration above 5 X 10-4Fdiethyltin dichloride. However, they measured average, not maximum limiting currents. If the formation of a definite quantity of insulating film blocks further reduction at the electrode, then the total quantity of electricity passes and thus, at constant drop time, the average current will become independent of depolarizer concentration, once this is high enough to produce the limiting quantity of film. The instantaneous current, on the other hand, is determined by the combined effects of diffusion-controlled, potentiostatic reduction of the depolarizer, the rate of growth of the mercury drop, and the continuing coverage of the drop by the insulating film. The maximum current attained depends on the relative rates of drop growth and drop coverage. The rate of coverage depends on the rate of diffusion to the electrode of the depolarizer and, therefore, on its bulk concentration. Thus, the maximum instantaneous current may continue to increase after the average current, corresponding to complete coverage of the electrode surface, remains constant upon further increase in depolarizer concentration. Over the entire pH range investigated, the diffusion current of l X lO-4Faquodiethyltin(IV) cation and its hydrolysis products remains constant (*3 %) under constant capillary conditions and varies linearly with the square root of mercury column height. Thus, the reduction is diffusion controlled. At high concentrations, where film formation is current-limiting the square root relationship does not hold. A plot of half-wave potential DS. logarithm of aquodiethyltin(1V) concentration is linear over the range 0.2 X lO-4F to 1.0 X lO-4F. The slope is -35 mV per decade of log concentration. The theoretical value is -29.6 mV per decade for a two-electron reduction. Above 1 X 10-4, the wave is too distorted to allow reliable half-wave potential measurements. Below 0.2 X 10-4F,the residual current is a sufficiently large component of the overall current, that the wave is not very well defined. The available concentration range represents a total shift in half-wave potential of less than 30 mV, so that the agreement with theory for a two-electron change is about as good as can be expected. Plots of Edme VS. log (id - i) are linear for 1 X 10d4Fsolutions in supporting electrolytes of different pH values. The slopes of these plots range from 31 to 39 mV with no apparent correlation with the buffer pH or composition. For a reversible, two-electron transfer the expected value is 29.6 mV. Our data is reconcilable with a two-electron reversible process if it is remembered that these logarithmic plots are very sensitive to small distortions in the rising part of the polarogram. At the foot of the wave, where very little reduction occurs, the polarogram is relatively undistorted. As the potential is made more negative, and more reduction occurs, the electrode becomes coated with an insulating film. Consequently, the maximum current during a drop life does not increase as rapidly as expected during the rising part of the polarogram. Thus, the slope of the log plot should be greater than 29.6 mV. The half-wave potential for the reduction wave can be described by Equation 6 at any pH and concentration of complexing agent, X.
EIi2= Eli2(acid solution) (13) J. G . A. Luijten and G. J. M. van der Kerk, “Investigations in the Field of Organotin Chemistry,” Tin Research Institute, Greenford, Middlesex, England, 1955, p. 103. (14) R. W. Schmid and C . N. Reilley, J . Am. Chem. SOC.,80, 2087 (1958).
478
ANALYTICAL CHEMISTRY
- 0.0296 log
Kl is the first acid dissociation constant of the aquodiethyltin (IV) cation and Kz is its second acid dissociation constant.
4.6
4.0
4.5
4
1
3.0
-3
4.4
d 2.0
m
2
.N
4.3
I .o
8 8
4.2
’+
4.1
a
1.0
a
ae
4.0
4.0 5.0 6.0 7.0 (Et2 SnCls), F X lo4
3.0
8.0
9.0 10.0
Figure 3. Variation of limiting current with concentration of aquodiethyltin(1V) cation. Supporting electrolyte, 0.1F KCI, 0.W HCI, 0.001 Triton X-100
3.9
3.8
2.0
1 -0.2 -0.4 -0.6 -0.8 -1.0
-1.2
-1.4
E, I’ us. SCE Figure 2. Electrocapillary curves a. 0.1F HCl, 0.1F KCl b. 0.1F HCI, 0.1F KCl, 0.001 Triton X-100 c. 1.01 X 10-4F aquodiethyltin(1V) cation in 0.1F KCI, 0.1F HCI, with and without addition of 0.001 Triton X100. Capillary flow rate, 1.79 mg/sec
z
z
KIis the constant for the formation of a 1 :1 complex between the aquodiethyltin(1V) cation and ligand X. The pH dependence is illustrated in Figure 4 for 1.02 X 10-4F diethyltin cation in various buffers at ionic strength 0.1 to 0.20. From these data we estimate pK1 = 3.0 & 0.1, pK2 = 5.4 + 0.2. It was assumed, in ekaluating the data, that pH = -1og(H+), and so the derived constant are mixed acid dissociation constants (15). That is, while concentrations of the acid and its conjugate base are t:mployed in the constant, hydrogen ion activities, values of pH, are used in the same expression. Tobias, Farrer, Hughes, and Nevett (8) report pK1 = 3.40 for aquodiethyltin(1V) cation in 3M sodium perchlorate. For the aquodimethyltin (IV) cation, Tobias and Yasuda (16) find pK1 = 3.54 in 3 M sodium perchlorate, 3.26 in 0.1M sodium chloride, and 3.1 3 in O.1Msodium nitrate. As we will show below, chloro complexation does not influence significantly the apparent value of Kl obtained for aquodiethyltin(IV) cation in media of 0.1 to 0.2 ionic strength, If we assume for the aquodiethyltin(1V) cation the shift of 0.3 to 0.4 unit of pK1 found for the aquodiethyltin(1V) system over the range ionic strength 3 to ionic strength 0.1, we find from the data of Tobias, Farrer, Hughes, and Nevett an extrapolated pK1 of 3.1 to 3.0 for aquodiethyltin(1V) cation in a medium of ionic strength 0.1. This is in good agreement with our value of 3.0 =t0.1 for ionic strength 0.1 to 0.2. No independent estimates of pK2 for aquodiethyltin(1V) are available. However. the similarity between this system and the extensively studied aquodimethyltin(1V) system, for which pK2 = 5.28 in O . l M sodium chloride suggests that our value of pK2 = 5.4 =t0.2 in media of 0.1 to 0.20 ionic strength is reasonable. The use of Equation 7 for the evaluation of shifts in half(15 ) A. Ringbom, “Complexation in Analytical Chemistry,” Wiley, New York, 1963, 13. 25. (16) R. S. Tobias and M. Yasuda, Can. J . Chem., 42, 781 (1964).
wave potentials is valid only if aqueodiethyltin(1V) cation undergoes no other chemical reactions than simple acid dissociation and complexation. The extensive studies of Tobias and his coworkers (7) have shown that the hydrolysis products of R & I ( H ~ O ) ~are + ~ extensively polymerized. The constant for dimerization of EtzSn(HzO)OHf,the most important of these reactions is Kd = 259 in 3 M NaC104(8). In the case of aquodimethyltin(1V) log Kd = 2.4 in 3M NaC104, but only 1.4 in O.lMNaCl(16). The tendency to form highly charged polymeric aggregates is significantly greater in the more concentrated electrolyte solution. Thus the literature value of Kd = 259 (3M NAC104) may be taken as a very conservative upper limit for the dimerization constant of aquodiethyltin (IV) in the media of 0.1 to 0.2 ionic strength employed in this study. The true value may be an order of magnitude lower. Using this upper limit we calculate that in 1 X 10-4F aquodiethyltin(1V) solutions less than 1 of the organotin species is dimerized if one assumes that the remainder is in the form of the precursor, Et%Sn(HzO)OH+. Therefore, it is reasonable to neglect the effects of polymerization on pH shifts in the present study. In the earlier stages of this work we attempted to use chloroacetate buffers to span the pH range from pH 2 to pH 4. However, chloroacetate forms a complex with aquodiethyltin(1V) cation and anamolously large values of Kl are obtained if one overlooks this complexation. From the shift in halfwave potential at pH 2.8 with increasing chloroacetate concentration we estimate that log Kr = 1.9 & 0.1 in a medium 0.1M KCl. Similary, chloride forms a weak 1:1 complex with aquodiethyltin(1V) cation. We estimate that log K r = 0.4 0.2 for the formation of this complex in a medium of ionic strength 0.5. Shifts in half-wave potentials were measured in perchloric acid-hydrochloric acid mixtures of total concentration 0.5F. The difference in half-wave potentials between supporting electrolyte of 0.5FHC104and 0.5F HCl is only 12 mV. Thus, the derived value of log Kr is highly imprecise. B y careful potentiometric measurements, Farrer, McGrady and Tobias (17) have found log K I = 0.38 for the formation of MezSnCl+in 3 M NaC104-HC104 medium at 25” C. This is in excellent agreement with our estimate of log K , for the formation of EtzSnC1+. A difference of about 3 mV is observed between the halfwave potentials for reduction of aquodiethyltin(1V) in 0.5F
*
(17) H. N. Farrer, M. M. McGrady, and R. S. Tobias, J . Am. Chem. Soc., 87, 5019 (1965). VOL. 39, NO. 4, APRIL 1967
479
I
I
I
,
I
I
1.0
2.0
I
I
I
I
I
I
I
I
, -1
-1.200 -1.100
w -1.000
5:
I
g -0.900
*
3 -0.800
q
-0.700
-0.6001
3.0
I
I 4.0
5,O
6.0
7.0
, 8.0
9.0 10.0
PH Figure 4. Variation of El/, with pH 1.02 X 10-4F aquodiethyltin(1V) cation. Supporting electrolyte, 0.1 F KCl, plus buffers (see text), 0,001 Triton X-100as maximum suppressor
HCIOl and in 0.25F HC1-0.25F HC104. Consequently, the effect of chloro-complexation is probably no greater than the error in measuring half-wave potentials, and is not a significant factor in the apparent value of pK1 or pK2 obtained in this work. We were unable to detect half-wave potential shifts in acetate buffers of pH 4.6 and varying acetate concentrations. If the formation constant for acetate complexation is the same as that for chloroacetate complexation, then the predicted shift in half-wave potential due to acetate complexation is less than 1 mV for the acetate buffers used in this study (0.05F acetic acid species). Coulometry. In order to verify that the reduction of the aquodiethyltin(1V) cation is a two-electron process, exhaustive controlled potential reduction of 6 X 10-3Fsolutions of aquodiethyltin(1V) cation in 0.1F KC1, 0.1F HCl were carried out. The control potential was -0.80 V us. SCE, well on the limiting current plateau of the polarographic wave of aquodiethyltin(IV), cation in this medium. Coulometry at this potential gives an n-value of 1.99 i 0.03 for the reduction. The product of the reduction is a greenish-yellow, oily substance which tends to form a thin film over the electrode surface, Because of this tendency of the electrolysis product to coat the electrode, complete electrolysis requires 5 to 6 hours. With the same apparatus and at the same concentration levels, other two-electron reductions of organometallics, where flocculent precipitates are formed, required 30 to 60 minutes. The product of the reduction is, at least initially, a diethyltin radical, which polymerizes, as suggested by Riccoboni and Popoff (3). We have attempted unsuccessfully to characterize the polymer formed. It is readily soluble in benzene,
480
ANALYTICAL CHEMISTRY
carbon tetrachloride, or cyclohexane to give a greenish solution. This is oxidized by atmospheric oxygen to diethyltin (IV) oxide in a few minutes. Consequently, attempts to measure its ultraviolet and infrared spectra have given very ambiguous results. The oxidized polymer is soluble in perchloric or hydrochloric acids, although dissolution is slow. The most likely initial polymer is a cyclic structure of 6 to 9 diethyltin units, such as has been observed for polymeric diethyltin and dimethyltin prepared chemically. Such cyclomers are readily susceptible to rearrangement and loss of tin (10). The material recovered after several hours at a mercury-aqueous solution interface may not be ihe same as that formed in a few seconds on a dropping mercury cathode. Thus, while it would probably be possible to isolate and characterize the polymeric diethyltin formed by coulometry, the results could not be related unambiguously to the polarographic data. Preliminary cyclic voltammetric measurements have shown that an anodic wave, presumably due to reoxidation of unpolymerized or partly polymerized diethyltin(I1) can be observed at sweep rates as low as 2 volts/sec (1 cps). Further cyclic voltammetric studies of the lifetime of diethyltin radicals are now in progress.
RECEIVED for review November 16, 1966. Accepted January 27,1967. Presented in part, Division of Analytical Chemistry, 152nd Annual Meeting, ACS, New York, N. Y.,September 1966. Work supported in part by National Science Foundation Grant GP-6164.
