Thenoyltrifluoroacetone: Polarographic and Spectrophotometric Behavior, and Dissociation Equilibria Mechanisms of Electrolytic Reduction PHILIP J. ELVING and PHILOMENA G. GRODZKA' Department o f Chemistry, University o f Michigan, Ann Arbor, Mich.
b Over the usual pH range, thenoyltrifluoroacetone (TTA) gives four polarographic waves, whose presence and properties depend upon the pH, the specific buffer used, and the buffer component concentration. The report of only two TTA waves in a previous study is due to the large pH intervals between measurements. Spectrophotometric examination of TTA solutions, before and after controlled potential electrolysis at a massive mercury cathode, has made possible an explanation of the complex polarographic behavior on the basis of the various tautomeric and acid-base equilibria involving TTA species, including assignment of the polarographic waves to the various TTA species present. The observed variations in the polarographic behavior of TTA in different alkaline buffers are probably due to specific catalytic effects of buffer components upon the specific rate constants for the conversion of ionized, nonreducible (at the potentials involved) TTA species to un-ionized, reducible species and not to borate complexation as previously suggested.
tulated to explain the polarographic behavior of TTA in the alkaline region. I n a previous study (3) of TTA, the hydrated form (11) and borate were reported to form a complex at p H 9. However, the ultraviolet spectrum of TTA in pH 8 borate buffer indicates it to be essentially all in the enolate form (IV) (8). Obviously, these conclusions are in apparent contradiction. The differing natures of the borate species between p H 8 and 9 would not be a
RI- C -C H g - C - R 2
X
2
O0-
S
of various heavy metal ions involving thenoyltrifluoroacetone (TTA, compound I of Figure 1) as chelating agent have been extensively investigated in recent years. Selectivity of action, as in extraction, is frequently possible by p H control because of the multiplicity of possible TTA forms due to hydration, enolization, and subsequent dissociation. The interrelationships of these various TTA species were investigated in the present study largely in terms of the electrochemical behavior of TTA, with considerable aid from its absorption spectra. The equilibria involved are shown in Figure 1; part of this scheme is based on the literature (2, 3, 7 , 9, 16, 19, 80) and part on the present studye.g., forms V and VI have been posEPARATIOKS
Present address, Department of Chemistry, Smith College, Northampton, 1
Mass.
2
0
ANALYTICAL CHEMISTRY
R,-C
- CH2-CIxa
dH
9-
+ti+
Ri-c-cH2-k-R~
significant factor (cf. subsequent discussion). Resolution of this contradiction and of the resulting implications required re-examination of both the spectrophotometry and the polarography of TTA under comparable conditions. It is apparent from the literature on the latter topics that the effect of experimental variables has not been sufficiently considered. The results of the present study and the literature data have revealed the basis for the
R l - C - C H 2 - CI - R 2 -H+
= -HzO tHgO
R2
RI-C-CHg-C-Rg 1Xb
OH
ILl
R~-C=CH-C-R~
I I
m
P
R
9-
$
0I RI-C-CH=C-Rn-
i
+ RgCOOH
F1 Ri-C-CH3
I
OH
X I I '
I
=
I
+
R2COOH
Figure 1. Summary of postulated polarographic reduction mechanisms, and of dissociation, tautometric, and hydrolytic equilibria for thenoyltrifluoroacetone (TTA) and related and derived species R1
= the 2-thenoyl group, HC-CH
I1
HC
/I
C-
's' R j = the trifluoromethyl group,-CFs Rectangles represent electrolysis reactions producing the various waves and are designated by wave numbers; I' indicates wave I in alkaline soh.
contradiction and have allowed the assignment of TTA species giving rise to the various polarographic waves. To save space in subsequent discussion, the carbonyl group site adjacent to the thiophene ring in TTA will be referred to as the thenoyl carbonyl: the carbonyl group site adjacent to the trifluoromethyl group will be cited as the gamma carbonyl. SPECTROPHOTOMETRIC BEHAVIOR
OF TTA
Interpretat'ion of the polarographic data for TTA necessitates accurate knowledge of the chemical species present. The present study revealed that previous reports (2, 9, 20) of the TTA species in solut'ion were based on spectrophot'omet'ric data in which solvent effects were not sufficiently evaluated. Coneequently, these effects n-ere investigated. Assignments of absorption maxima, reported for various TTA species and related compounds in different media and found in the present study, are summarized in Table I. In aqueous acid solution the spectrum of TTA (actually TT.I.H20 with the gamma carbonyl hydrated) resembles that of 2-acetylthiophene with absorption maxima in t'lie regions of 260 to 270 and 285 t o 295 m p ; these have been assigned to the undissociated hydrate form (11) (20). In benzene solution TT.4 exhibits a maximum a t 325 mp with a side shouldpr a t 355 mp, which has been ahsigned to the enol form (111) (20); the benzene cutoff int'erferes with observation o. absorption in the 260- to 295-mp region. TTA. HzO dissolved in benzene does not absorb a t 325 and 355 mp. In aqueous alkaline solution TT-1exhibits a new band a t 340 mp, which increases exponentially n i t h pH and which has been assigned to the clnolat,e form (IV) (2, 20). The infrared cqecbra of solid TTA and TTA.HZ0, measured in KBr disks, provide clear-cut evidence for structures in which strong hydrogen bonding is present-i.e.,
h
v h
13
? m m 2 N
h
co
v
13
cc
m 13
ce N
' I
,si-C=CH--&-CF. ', i o-H...o TTA 0 . ..H-0
U-(!-CHZ-(!-CF' I . /
Hcj TTA.HzO
Solut'ions of TTA in solvents more polar than benzene, with the exception of water, would therefore be expect'ed to show spectra of the diketo (I) form ( 5 ) ; TTA.H20 spectra Ivould not be expected to undergo much change with increased solvent polarity, since tautoVOL. 33, NO. 1, JANUARY 1961
0
3
meric changes are blocked in this compound. To evaluate the spectral assignments, especially that of the enolate and a possible wavelength shift in going from nonaqueous to aqueous media, TTA was examined in nonaqueous media of varying polarity. I n 1 to 1 benzene-methanol, both TTA and TTA . H 2 0 unexpectedly showed a maximum a t ca. 338 mp (presumed enolate absorption). In absolute ethyl alcohol, TTX showed a broad absorption peak (maximum a t 330 mp; shoulder a t 265 mp) ; its spectrum in 95y0 ethyl alcohol was similar to that in alkaline aqueous solution; increasing amounts of water decreased the 340-mp absorption. I n ethylene glycol dimethyl ether, the TTA spectrum was similar to that in absolute ethyl alcohol; TTA. HzO, on the other hand, showed one similar to that of TTA in acidic aqueous solution with maxima a t 262 and 290 mp. After a day’s standing, the spectrum of the former TTA solution showed that the anhydrous TTA was being converted to the hydrate, probably because of traces of water present in the solvent. The foregoing data indicate that either of two situations exists. If the original spectral assignments (20) are correct, alcohols and the glycol ether can act as bases in converting TTA to the enolate form. On the other hand, if the 338-mp absorption band present in polar solvents and aqueous alkaline solution does not represent a TTA species (enolate) different from the TTA species (enol) present in benzene solution alone, the difference in spectra obtainedin polar and nonpolar solvents must be due to physical solvent effects. For instance, mesityl oxide, (CHa)2C=CH--CO--CHa, exhibits K and R bands which tend to coalesce with inceasing solvent polarity (6). The TTA enol bands a t 325 and 355 mp in benzene could possibly shift towards each other in the more polar solvent, water, so as to appear as one band a t 338 mp. If the first situation exists, addition of compounds which act as bases in benzene should produce an enolate spectrum. Aniline addition did not alter the spectrum. However, piperidine addition did produce a large symmetrical absorption band (maximum a t 340 mp; small shoulder a t ca. 280 to 290 mp) in accord with the expectation that piperidine would be a stronger base than aniline in benzene. Evidently, solvent effects do not shift maxima except for the conversion of one TTA form to another due to solvent interaction.
presence of only one wave in borate buffer was attributed to complexation with buffer component. The first TTA wave was assigned to reduction of the gamma carbonyl; its kinetic nature was attributed to conversion of hydrated TTA (nonreducible a t the potentials involved) to the reducible enolate form. The second wave was ascribed t o a l e to 2e diffusion-controlled reduction of the thenoyl carbonyl to pinacol, carbinol, or both as the pH varies; in alkaline solution, reduction of thenoyl carbonyl may have contributed to the first wave. Figure 2 summarizes the polarographic behavior of TT.4, observed in the present investigation, as a function of pH; the validity of using solutions 0.045M a in buffer system and 0.0005M in TTA, upon which Figure 2 is based, is subsequently justified. Detailed discussion of Figure 2 will be deferred until the influence of buffer component concentration is considered, except to note the presence over the pH range of four different TTA waves, which, to facilitate subsequent discussion, will be designated as I, 11, IIIa, and IIIb; the composite behavior of waves IIIa and IIIb, which fuse into one a t higher ionic strength, will be referred to as that of wave 111. The behavior previously assigned (3) to the second TTA wave is actually a composite of that shown by waves I1 and
111. Although wave I11 definitely appears a t pH 6.9, it may be present, though
2 POLAROGRAPHIC BEHAVIOR
OF
TTA
TTA has been reported (3) to give two polarographic waves in chloride, acetate, and ammonia buffers; the
4
0
ANALYTICAL CHEMISTRY
4
small and masked by hydrogen discharge, a t lower pH-e.g., a definite inflection is noted in the polarographic curve a t pH 6.7. Variations of current with drop time (corrected mercury height) and with temperature indicate that wave I is, in the main, kinetically controlled in the acid and alkaline regions, tending toward diffusion control in the neutral region. Wave I1 shows diffusion control; wave I11 would also seem to be diffusion controlled. Use of Low Concentration Buffers. I n view of the subsequently discussed TTA-buffer component interaction, which affects the wave I height, i t seemed advisable t o keep the buffer concentration as lorn as possible, so t h a t the p H dependence of the polarographic behavior could emerge more clearly. Consequently, data were obtained a t the lowest level advisable for polarography-i.e., concentration ratio of TTX to buffer of 1 to 100. Since buffer concentration affects wave I least in acetate and borate buffers, any variation in TTA behavior in low concentrations of these systems should be due only to pH. The agreement of data obtained a t 0.045M buffer concentration in phthalate, chloride, and phosphate systems with those for acetate and borate buffers would indicate that buffer-TTA interaction is small and that the observed behavior is mainly tb function of pH. The adequacy of such loiv concentra-
6
8
IC
PH Figure 2. Variation of apparent diffusion current constant, I, and of €112 of ITA with pH Subscript A on the wave number Indicates that data were obtained in ammonia buffer
tions for buffering a t the electrode interface in the acid region is indicated by the fact that a t high acetate concentration the wave I1 height is less than a t low concentration. If the buffering capacity of the 0.045M system were not adequate, the wave I1 height would be expected to be lower in these solutions than in the more concentrated solutions, since the p H a t the electrode surface would then be higher than in the bulk solution; there is no doubt that increasing pH decreases the wave I1 height even in buffers of high ionic strength. In the alkaline region the buffering capacity is judged to be adequate because wave I shows only minor variations between 0.045 and 0.1M borate buffers. The effect of ionic strength variation n i t h pH for the 0.045M phosphate and phthalate buffers is negligible. The behavior of TTA in these systems agrees with that observed in 0.045M acetate buffer over the same p H range; in alkaline phosphate buffer the behavior of TTA is about the same as in borate buffer, xhich is 0.045M in KC1. However, the data for 0.045M ammonia buffer indicate that buffer-TTA interaction, resulting in a larger specific rate constant for the conversion t o reducible species than in the other 0.045M buffers, is still appreciable, as judged by the differing El and i values (cf. subsequent discussion). Specific Buffer Component Effect. Table I1 summarizes the effects of buffer component nature and concentration. Such behavior is readily explained on the basis of the subsequently discussed reduction niechanisms. The decrease in t5ave I1 ITith increased acetate concentration is, however. difficult to explain. If the decrease results from increased specific conversion rates or a varying ratio of undissociated to dissociated forms, w v e I would be expected to vary more than is actually observed. On the same basis, the decrease in wave I11 with increased borate concentration does not arise from increased conversion rates or a varying ratio of undissociated to dissociated forms. The increases in waves I and I1 with increased phosphate concentration are probably due to the increased specific conversion rates of nonreducible ionized TTA species to reducible, un-ionized species; the increased wave I and decreased anionic wave I11 with increased ammonia concentration are similarly explainable. El(a values for waves I and I1 are only slightly affected by increasing buffer component concentration; the effect on El,, of wave I11 is difficult t o decipher (Figure 3). The shape and division into distinguishable waves of wave I11 depend upon the specific buffer used; in the 0.045M buffer
systems, the split into two waves becomes discernible a t about pH 7.5. At this p H the shape of wave I11 approximates that of curve D in Figure 3. Because of the difficulty of resolving I I I a and IIIb, quantitative statements about them individually are apt to be misleading; consequently, only their qualitative behavior will be described. Their heights and .Elrz values in a particular buffer system appear t o remain constant over the p H range where their separation is discerniblei.e., 7.8 to 9.2-with the wave IIIQ height being about twice that of IIIb. El (uncertainty of 10.03 to 1 0 . 0 5 volt) is -1.53 volts for I I I a and - 1.67 for IIIb in borate and phosphate buffers, and - 1.46 and - 1.59 volts in ammonia buffer. Salt Effects. for kinetically controlled wave I is essentially identical in ammonia, borate, and phosphate buffers a t the same p H for a considerable ionic strength range. However, the wave height increases appreciably with increased buffer component concentration in ammonia and phosphate buffers, b u t not in acetate and borate; the possibility that this is due to salt effects depending primarily upon the charge and concentration of an ionic buffer component species seems doubtful. Increasing ionic strength by KCl addition to borate buffer does not appreciably affect the wave I height, and the two TTA waves in (NH4)zS04 buffer (pH 9.3) are essentially identical in all respects to the waves in NH&l buffer. Plots of the apparent wave I diffusion current constant LIS. the concentration ratio of A-/HA in the different alkaline buffers indicate differing dependenceLe., the wave I processes do not show general catalysis. If TTA-buffer interaction in ammonia and phosphate buffers affects the TT-4 species equilibria so as to favor formation of the species whose reduction produces r a v e I, the implication is that either a larger equilibrium concentration of a nonreducible TTA species is converting to a reducible species or the specific rate of conversion to the reducible species is greater in these buffers than in borate buffers of the same pH. The latter was shown to be the case by spectrophotometric ex-
Table 11.
8
19
17
r---7-
-1
1
_
._
I1
15
1.3
II
~ 13
_
~- ~
_ ~~
19
15
I? POTENTIAL VOLT’S
Figure 3. Polarograms of TTA in alkaline solutions; variation of wave 111 appearance with borate buffer concentration (TTA concn. is 0.48 to 0.56mM) Polarogrom Buffer A 0 . 0 4 5 ~H ~ B O0~. ,4 5 ~ KCI-NaOH 8 0.225M Borax-HCI C 0.025M Borox, 0.45M KCI-NaOH D 0.05M NaHzPO4-NoOH E 0.05M NHz-NHiCI
PH 8-44 8.43 8.36 7.94 8.40
amination of TTA in pH 8.8 ammonia and borate buffers. The absorbances a t 265 mp (due t o TTA.HzO and to acetylthiophene produced by hydrolysis) and a t 340 mp (due to TTA enolate), measured right after mixing, were about the same in both solutions. Furthermore, the absorbances of TTA
Summary of Effects of Buffer Component Concentration on TTA Wave Heights
Change in Wave Height with Increasing Buffer Concn: Buffer System Wave I Wave I1 Wave 111 Decreases NW Acetate NW Decreases Borate Increases Decreases Phosphate Increases NW Decreases Ammonia Increases A f sign indicates only a minor variation; NW indicates nonappearance of wave in buffer concerned.
**
VOL. 33, NO. 1, JANUARY 1961
5
solutions of widely varying ammonia concentration were identical. It will be recalled that increased ammonia buffer component concentration increases wave I height, whereas a similar borate increase has a negligible effect. These data rule out the possibility of a larger concentration in ammonia than in borate buffer of a TTA species, which could convert to a reducible species. Possible Borate Complexation. The previously suggested ( 3 ) complexation between a boron species and either the tn-o hydroxyl oxygens of the hydrated gamma carbonyl of TTA HaO or one hydroxyl oxygen and one fluorine, implies that the acetylthiophene portion of the molecule remains intact and that the complex should hare an absorption spectrum similar t o that of TTA in aqueous acid solution with the maxima someJvhat displaced. Since such a spectrum is not obtained, a TTA. H20borate complex seems unlikely, although an equilibrium of the type 0-
+ borate species
\ n1-!-CH=d-CF3 S’ -+ +-
1\&I
e- ,
0-borate
CH,-C-CF,
1
OH
1
(I)
would show an enolate absorption a t 340 nip. The latter has been observed ( 2 ) ; however, the equilibrium of Equation l is doubtful, since a fivefold change in borate concentration a t pH 8.4 and 8.8 in the present study did not appreciably affect the absorption; a twentyfold change in borate concentration a t p H 8.4 did not appreciably affect the height and El,, of wave I. A similar equilibrium involving T T h HzO seems doubtful for the same reason. The order of activity of borate ions for complexing with nonionizable polyoxy compounds is HIBOO -CO-CHa
0
0
0
If the TTA-hydrate ion (V or VI) is not reducible a t potentials a t which ketohydrate (11) and diketo (I) are reduced (the usual case with acid and conjugate base), the cause of the lack of correspondence between the decrease in wave I height with time and extent of hydrolysis in different alkaline buffers is evident. Spectral identification of the postulated TTA-hydrate ion (V or VI) is not feasible. Removal of a proton from ketohydrate (11) would be expected only to shift somewhat the absorption due to the acetylthiophene moiety, since no new resonance bands would result. Unfortunately, the large enolate ion (IV) and acetylthiophene (produced by the hydrolysis) absorption a t high p H would mask any such shift. However, some support for existence of the TTA-hydrate ion is offered by measurements of the TTA hydrolysis rates a t different OH- concentrations via the decay of the 338-mp enolate absorption peak, which indicate more than one hydrolyzable TTA species to be present (20). In a paper which appeared after the present work had been completed, slow conversion of a 0I
OH- fast
At pH 8.6: +-CH-CII~-CO-CH3 I
+-CH-CH,-CO-CH, and +-CH2-CHA--CO-CHI The corresponding electrolysis products of XVI, unfortunately, were not identified. If the polarographic behavior of the TTA enols (111) is analogous to that of XVI and XVII, with due allowance being made for the TTA species equilibria, TTA wave I1 a t pH 4 to 5 is due in part to reduction of the small amount of free radical (IXb) produced by the reduction of enol from IIIb (wave I). The reduction products of wave I1 are mainly the free radical (X) produced from the ketohydrate form (11),which can then dimerize, and some carbinol produced from reduction of the free radical (1x6). The subsequent decrease of wave I1 as wave I grows is probably due to the nonreducibility a t the potentials involved of the stable free radical (IXa) produced from the reduction of enol from IIIa in wave I. On this basis, the predominating enol species formed from ketohydrate (11) in acidic aqueous solutions must be IIIa. Further support for IIIa being the enol species predominating in the wave I reduction is obtained by comparing (Figure 4) the percentages of undissociated (HA) and dissociated acid (A-) with per cent I for the various waves (% I = I / I m a x )(10) and noting the close agreement of the percentages of HA and I of wave
dFor TTA, therefore, the analogous path from form (V) to (VI) to acetylthiophene and trifhoroacetic acid can be postulated.
Figure 4. Variation with pH of undissociated and dissociated acid forms of
TTA
- - - -Calculated -Calculated
from pKa of 6.38 (2). from ratios of diffusion current constant, I, for wave indicated to maximum value attained for I for wave
dihydrate ion to the stable enolate form is postulated to account for the slow rate of ionization of hexafluoroacetylacetone (17). Waves IIIQand IIIb appear a t pH ranges where TTA is in the enolate (IV) and TTA-hydrate ion (V and VI) species. Since wave I11 is largely, if not entirely, diffusion controlled, the reduction processes must involve these three forms rather than enols (111). The strong inductive effect of the trifluoromethyl group would be expected to make the gamma carbonyl relatively more negative than the thenoyl carbonyl and therefore more difficultly reducible. The decrease in wave I11 when wave I height is increased by increased ammonia buffer concentration also supports the conclusion that the thenoyl carbonyl is involved in the wave I11 reduction. The choice of the enolate form (IV) rather than the hydrated ion form (V) as being the form responsible for wave IIIa is based on the resonance possibilities in the free radicals produced upon reduction; the one (VII) produced from the enolate form (IV) would present more opportunity for resonance stabilization than the corresponding radical (XII) produced from the VOL 33,
NO. 1, JANUARY
I961
9
hydrated ion (V); consequently, the reduction of IV t o VI1 would be more facile than that of V to XII. The reduction mould continue to form VI11 (or XIII). Wave 11% may be due to reduction of enolate (IV) to the doubly charged carbinolate ion (XI) or of the hydrated ion form (V) to XI1 and XIII. Relation to Previously Postulated TTA Reduction Mechanisms. The present study generally agrees with the previous study (3) as to the species involved in the production of polarographic maves I and I1 in the acidic region. However, the p H intervals between measurements in the previous study were so large that the dependence of the wave heights on p H in the neutral and alkaline regions \vas misinterpreted; wave 111, misinterpreted as wave 11, was attributed to reduction of TTA HzO (11), while the present study indicates that it results from the ionized TT.4 forms V and VI. In addition, the effect of buffer component concentration on wave I height was not taken into consideration with the result that borate complexation was postulated to account for the lack of wave I in alkaline borate media. EXPERIMENTAL
Apparatus. Spectrophotometric measurements were made a t 25' i 2' C. on a Cary Model 11 spectrophotometer equipped with 1-em. glassstoppered quartz cells and a Beckman Model DU equipped with a hydrogen discharge lamp and matched 1-em. silica cells. A Leeds & Northrup Model E ElectroChemograph was used in the polarography with a jacketed electrolysis H-cell (25' 5 0.1" C.); the saturated calomel reference half-cell was connected to the test solution leg by a saturated potassium chloride-agar bridge. The cell resistance was about 900 ohms for 0.045M background electrolyte solutions, equivalent to a n iR drop of 0.004 volt for the largest currents obtained (4 pa.); iR corrections were not considered justifiable. I n O.1M KC1 (open circuit; 2 5 O ) , the four capillaries used had drop times of 3.8 to 4.9 seconds, and me/~t'/~values of 1.5 and 1.8. The coulometric apparatus had a
NO.
Table 111. PH 1.6-1.9 2.8-3.8 3.9-5,i 4.5-5.2 5.9-8,0 7.7-9.8
water-jacketed electrolysis cell (25' =t 0.1" C.), and a silver coulometer; the external anode (a coil of about 40 inches of No. 10 gage silver wire immersed in saturated KC1 solution) was connected to the cell via a saturated KC1-agar bridge. The potential of the mercury cathode was automatically controlled us. a S.C.E. by a Fisher controlled potential Electro-Analyzer, The operation of the apparatus was checked by electrolyzing known solutions of trichloroacetic acid. p H measurements, accurate to +0.05 p H unit, !yere made with a Beckman Model G meter. Chemicals and Solutions. TTA (99.57, pure, Midcontinent Chemicals Corp.), was used without further purification; good polarographic reproducibility was obtained on samples supplied over a period of several years. Sitrogen used for deoxygenating was purified and equilibrated by successive passage through alkaline pyrogallol solution, sulfuric acid, water, and a portion of the teqt solution. All other chemical. were of C.P. or reagent grade. Alcoholic stock solutions were prepared by dissolving TTA in ethyl alcohol (the final solution was usually 10% ethyl alcohol) in a volumetric flask and then adding water. Aqueous stock solutions were prepared by dissolving the TT.4 in warm water, cooling, and then diluting to the required volume. Once, on preparing an alcoholic stock TTA solution, a red oil separated out, which dissolved on shaking; formation of the oil could not be again obtained under similar conditions. A 31mM TTA solution (60% alcohol), which had been standing for about 6 months, had a red-orange color; the spectrum of this solution, diluted to appropriate concentration with miter, was the same as that of a freshly prepared solution. The red form wa5 also readily produced in benzene solutions containing a high TTA concentration; appropriate dilution with benzene for spectrophotometry resulted in disappearance of the red color and spectra were exactly similar to those obtained with colorless TTA solutions. Such behavior would indicate a n equilibrium, yTTA F? red compound ( A X ) (11)
where q is probably large, since the red form appears a t high concentration and disappears on dilution.
Buffer Systems Employed
Buffer Composition
0 05M KCl-HC1 0 05M KHCsHhOrHCl 0 05M NaOA4c-HOAc 0 0531 KHCsH4O4-?JaOH 5 0 05M SaH2POrljaOH 6 0 05M HaBO,, 0 05M KC1-HC1 or KaOH i 8.2-9.4 0 05-M NHz-NHjCl 1,72 1.11 KC1-HC1 8" 5.02 1X SaOAc-HOAc 90 10" 6.52 1M KH2POA-SaOH 110 8.01 0 05M Borax, 0 9M KC1-HC1 Buffers 8 to 11 Tere used in coulometric runs and subsequent spectrophotometry. 1 2 3 4
10
ANALYTICAL CHEMISTRY
Stock buffer solutions were generally prepared by adding either concentrated acid or base (unmeasured component) to a salt (measured component) solution until the desired p H was obtained and then diluting to volume. Buffer compositions are given in Table 111. Stability and Polarographic Reproducibility of TTA Stock Solutions. The 10% alcohol-TTA stock solutions were quite stable; identical polarograms were obtained in acetate and ammonia buffers of p H less than 9 for freshly prepared stock solutions and for those which had stood for a number of days. ;ibove p H 9 the current reproducibility n as very poor because of extensive hydrolysis of TTA in the test solution, although the general features of the hydrolytic and polarographic behavior were discernible. The 1% alcohol concentration of the test solutions has no apparent effect on the polarographic behavior of TTA in 0.045M buffered solutions-e.g., results for 1% alcoholic and aqueous test solutions were similar. Polarographic Procedure. The test solution was usually prepared immediately before the run by diluting 5.00 nil. of stock solution to 50.0 ml. in a volumetric flask with the desired buffer; p H was measured on a portion of the test solution both before and after the run; no significant variation was ever found. I n cases where the sequence of adding one solution to another was thought important, the point was tested; the order was always found to be insignificant. Nitrogen was bubbled through the test solution for 10 minutes, the mercury electrode inserted, and nitrogen passed for 5 more minutes. A nitrogen atmosphere was kept above the solution during electrolysis. I n the case of multiple waves, currents were corrected for electrocapillarity. Temperature coefficients were calculated by a binomial expansion. Coulometric Procedure. Forty-five milliliters of buffer solution was added to the cell and then deoxygenated for 10 minutes; mercury was introduced and the solution electrolyzed a t a potential about 0.3 volt more negative than that at which the electrolysis was to be run, until the current fell to less than 1 ma. The mercury was removed; 5.00 ml. of TTA stock solution added; nitrogen bubbled through the cell solution for 5 minutes; the mercury reintroduced; the coulometer connected; and electrolysis a t the desired controlled potential started. A nitrogen atmosphwe \vas maintained over the solution during electrolysis. Solutions appropriate for spectrophotometry n-ere made by diluting 5.0 ml. of the electrolysis solution to 50 ml. with the same buffer as used in the run or with high concentration buffer of another p H of interest. ACKNOWLEDGMENT
The authors thank the U. S. Atomic Energy Commission and the Horace H. Rackham Graduate School of The
University of Michigan for support of the work described. They also thank James L. Johnson of The Upjohn Co. for the infrared spectra of TTA.
(6) Hartough, H. “Thiophene and Its Derivatives,” Interscience, S e w York, 1952. (7) Irving, H. AI., Quart. Revs. 5, 224 (1951). (8) Kemp, P. H., “Chemistry of Borates,”
Borax Consolidated, Ltd., London,
LITERATURE CITED
(1) Campaigne,
E., Diedrich, J. L., J . A m . Chem. SOC.73,5240 (1951). ( 2 ) Cook, E. H., Taft, R. IT., Jr., Ibid., 74,6103 (1952). ( 3 ) Elving, P. J., Callahan, C. AI., Zbid., 77,2077 (1955). (4) Elving. P . J.. Leone. J. T.. I b i d . . 80,1021 (1958). ’ ( 3 ) Gillam, -4., Stern, E. S., “Electronic
Absorption Spectroscopy,” .irnold, Ltd., London, 1954.
Edward
1956. (9) King, E. L., Reas, IT. H., J . A m . Chern. SOC.73, 1806 (1951). (10) Kolthoff, I. AI., Liberti, ii., Ibid., 70, 1888 (1948). (11) Larsen, E. >I., Terry, G., Leddy, J., Ibid., 75,5107 (1953). (12) Lowry, T. M.,Lishniund, R. E., J . Chem. SOC.1935, 1313. (13) Pasternak, R., Helv. Chim. Acta. 31, 753-76 (1948). (14) Pearson, R. G., Nayerle, E. A., J . -1rn. Chem. SOC.73, 926 (1951).
(15) Rasmussen, R. S., Tunnicliff, D. D., Brattain, R. R., Ibzd., 71, 1068 (1949). (16) Reid, J. C., Calvin, M., Ibid., 7 2 , 2948 (1950). (17) Stewart, R., Van der Linden, R., Can. J. Chem. 38,399 (1960). (18) Tzuzki, Y., Bull. Chem. SOC.Japan 16,23 (1941). (19) Kish, L., Bolomey, R. A., J . A m . Cheni. SOC.72, 4486 (1950). (20).Zebroski, E., Atomic Energy Commission Rept. BC-63 (1947); Ph.D.
thesis, University of California, Berkeley, Calif. (21) Zuman, P., Suture 165,485 (1950).
RECEIVEDfor review July 25, 1960. -4ccepted October 12, 1960. Division of Analytical Chemistry, 137th Meeting, ACS, Cleveland, Ohio, April 1960.
Effect of Electrode Configuration and Transition Time in Solid Electrode Chronopotentiometry ALLEN J. BARD Deportment of Chemistry, University of Texas, Austin, Tex.
b To determine the optimum electrode and transition time range for chronopotentiometric analysis, the transition ”C0, was measured time constant, io+ for the reduction of silver(1) and lead(II), and the oxidation of iodide and hydroquinone, over a transition time range of 0.001 to 300 seconds. The transition time constant increased at long transition times, due to spherical contributions to diffusion and natural convection. Increase at short transition times was ascribed to charging of the double layer, electrode oxidation, and roughness of the electrode. By employing a horizontal electrode, with a glass mantle, oriented so that density gradients were not produced, i,+iz/Co was maintained constant to =t0.2% with transition times of 7 to 145 seconds.
A
the theoretical basis of chronopotentiometry is N ell established. and several preliminary expcrimental studies indicate its general analytical applicability, relatively feir applications to actual analytical problems have been made. Unlike polarography. in which current is a dcpendent variable, chronopotentiometry requires the choice of a current density, nhich may vary over a wide range, depending upon the transition time desired. Several papers recommend choosing the current density to obtain short transition times (0.3 to 1 second) ( I O , 1 9 ) , while other authors prefer transition times in the range of 10 to 60 seconds ( 3 , l/t, 18). Chronopotcntiometry with solid electrodes also allows the choice of an electrode LTHOUGH
from a number of different types that have been used, nhich include: disk electrodes (platinum disk embedded in glass) (3. 9, I O , 20), foil electrodes (9, I O , 14, 18, 20). cylindrical wire electrodes ( 9 , 10, 1 9 , BO), and nire loop electrodes (80). This paper discusses factors which affect choice of electrode and transition time range. The fundamental equation of chronopotentiomctry is
where io is the current density (microamperes per square centimeter), T is the transition time (seconds), C” is the concentration of the electroactive species (millimoles per liter), n is the number of faradays per mole of reaction, D is the diffusion coefficient of the electroactive species (square centimeters per second), and F is the faraday. This equation is derived under the assumption that semi-infinite linear diffusion is the only means of mass transfer of the electroactive species to the electrode surface. This equation predicts that i , ~ l ’ ~ / C ’[the transition time constant ( I S ) ] should be constant over any range of current densities or transition times for a given reaction a t the same electrode. Although no previous evaluation of the transition time constant over a very wide range of transition times has been made, practically every chronopotentiometric study has noted a decrease in i o ~ l ’ * j C viith o decreasing transition time (increasing current density). Delahay ( 4 ) ascribed
this dccrease to heating of resistors in the current source during a trial, leatling to a decrease in current density, but other authors found this variation, even when the current was held constant. Lingane (14) suggested t h a t failure to take into account the blank transition time for the supporting electrolyte alone may contribute to this variation. Howercr, even a t a constant concentration of the electroactive species, C”, i , ~ ~is’ *found to vary with current density. Davis and Ganchoff (3) recently studied this variation in the reduction of h1nO4-, Cr207-2, Cc(IV), Fe(III), and T’(V) a t a platinum electrodr. The variation was apparent in all of these reductions (transition times of about 5 to 80 seconds) and could not be eliminated by any choice of graphical method for evaluating T. The authors resorted to the expedient of making all mrasurements a t a giwn transition time (30 i 2 seconds) ; yielding an average deviation of 0.6% over a fivefold concentration range. To ascertain the cause of this variation and to define the conditions under which the transition time constant is truly constant, i O T l i 2’c”was measured with plane electrodcs of different construction and orientation in the solution over a very nide range of current densities and transition times. EXPERIMENTAL
Cell. The cell used in these studies comprised a 400-ml. electrolysis beaker with a platinum foil auxiliary electrode (3 x 3 cm.) enclosed in a chamber separated from the main cell VOL. 33, NO. 1, JANUARY 1961
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