equation (at 245‘ C), are 4.0 X [for Pb(II)] and 6.6 X 10-6 sq cm per second for [Cd(II)] in comparison with 2.3 X 10-6 sq cm per second for Pb(I1) and 5.1 X 10-6 sq cm per second for Cd(I1) (both at 264’ C) obtained by Inman and Bockris (9). Copper (11). Cu(I1) is a better oxide ion acceptor in molten NaN03-KN03 than either Cd(I1) or Pb(I1) (18). Apparently the precipitation of CuO, more important at higher temperatures (13,has in some cases (12, 13) prevented the voltammetric studies of Cu(I1). Voltammetry was made possible by either the addition of an acid (KHSO,) (32) which competes successfully with Cu(I1) for the oxide ion, or the use of CuC12 as a solute (12) indicating the formation of a chloro complex. Within the concentration range studied, we have obtained a linear dependence of the peak current (first wave) on concentration (Figure 3), using either C U ( N O ~ ) ~ or CuClz as a solute. A second wave at - - 0 . 2 volt, which was not concentration-dependent, was observed for both C U ( N O ~and ) ~ CuClz; a third wave at -0.5 volt was observed only for the case of CuClz or upon addition of KC1 to the melt containing CU(N03)z. Both the second and the third waves are much better defined at faster scan rates (Figure 6). This behavior is qualitatively similar to the case of multi(31) A. D. Graves, G. J. Hills, and D. Inman, “Advances in Electrochemistry and Electrochemical Engineering,” P. Delahay, ed., Vol. 4, Interscience, New York, 1966. (32) Yu. S. Lyalikov, Zh. Anal. Khim., 8, (1953).
step charge transfer with catalytic regeneration of the reactant (33),suggesting that Cu(I1) may be regenerated by the reaction of Cu(1) with NOs- ions. The observed decrease of i,/u1’2 with u (Figure 4) is also in qualitative agreement with the theory (33). The theory of the above mechanism is not directly applicable if partial precipitation of Cu(I1) and Cu(1) takes place. In summary, firm conclusions regarding the reduction of Cu(I1) in this medium cannot be made until additional chemical studies on both Cu(I1) and Cu(1) in molten nitrates have been performed. Indium(II1). The addition of I n c h to molten NaN03K N 0 3 results in a yellow precipitate, possibly InOC1. The voltammogram at a platinum wire electrode is drawn out. Much better definition is obtained at an indium pool electrode [which was also used for anodic formation of In(III)] (Figure 1). The limiting or peak current increased with the concentration of In(III), although in view of precipitation, the concentration dependence of peak current was not studied thoroughly. The above observations indicate that In(II1) is probably similar to Al(II1) and other trivalent cations (18) in being a very good Lux-Flood acid in molten NaN03-KN03. RECEIVED for review August 25, 1967. Accepted December 26, 1967. Work partially supported by the Atomic Energy Commission, Contract AT-(40-1)-3518. (33) D. S. Polcyn and I. Shain, ANAL.CHEM., 38, 376 (1966).
Polarographic Reduction of OxaIuric Acid Analytical Application Glenn Dryhurst’ and Philip J. Elving The University of Michigan, Ann Arbor, Mich. Oxaluric acid (oxalic acid monoureide, carbamyloxamic acid) is polarographically reduced in a two-electron process apparently to glyoxyllic acid monoureide. The reduction occurs by two distinct mechanisms; the first, operative in acid solution (pH 1 to 5), is pHdependent and involves reduction of the undissociated acid. Above pH 7, the oxalurate anion is reduced in a pH-independent process. Between pH 5 and 6, both processes occur, the heights of the respective waves depending on the extent of dissociation of the acid and the rate of recombination of anion and hydrogen ion. The factors involved in the polarographic measurement of oxaluric acid in analytical situations are discussed.
DURINGAN INVESTIGATION of the electrolytic oxidation of purines, it became necessary to identify, characterize, and determine oxaluric acid, which was suspected of being a hydrolysis product of one of the primary oxidation products, parabanic acid. Preliminary investigation indicated that oxaluric acid (I), was readily reduced at the dropping mercury
EXPERIMENTAL
Apparatus. Polarograms were recorded on a Sargent Model XXI polarograph, using a water-jacketed H-cell (2) maintained at 25” i 0 . 2 ” C, which contained a saturated calomel electrode (SCE) in one leg; all potentials cited are us. SCE. Potentials were checked with a Sargent S-30260 laboratory potentiometer. The apparatus for cyclic voltammetry has been described (3), as has the preparation of the 1 Present address, University of Oklahoma, Norman, Okla.
COOH
“2
I
I N
492
electrode (DME) and could be readily determined polarographically. The only previous reference to the polarography of oxaluric acid is Hladik’s report ( I ) that it was reduced in two steps at such negative potentials that wave heights and half-wave potentials could be only approximately estimated. The general lack of clarity and incompleteness of this earlier work led us to undertake a more thorough study of the polarographic behavior of oxaluric acid.
ANALYTICAL CHEMISTRY
(1) V. Hladik, Sb. Merzinarod Polyarog. Sjezdu Praze, 1st Congr., Part I, 686 (1951). (2) J. C . Komyathy, F. Malloy, and P. J. Elving, ANAL.CHEM., 24, 431 (1952). (3) G. Dryhurst and P. J. Elving, ANAL.CHEM., 39,606 (1967).
POTENTIAL, VOLT Figure 1. Polarograms of 0.5 mM oxaluric acid A . pH 1.15 chloride buffer B. pH 2.9 chloride buffer C. pH 4.7 acetate buffer D. pH 6.1 McIlvaine buffer E. pH 7.9 McIlvaine buffer F. pH 9.2 ammonia buffer G. pH 11.2 chloride bufkr
pyrolytic graphite electrode (4). The hanging mercury drop electrode was prepared by collecting two drops from the DME onto an amalgamated platinum disk sealed into glass. The capillary (marine barometer tubing) had rn = 1.857 mg per second in distilled water at open circuit at h = 58 cm; drop times, measured at the potentials of interest, were between 3 and 4 seconds. Chemicals. Oxaluric acid was prepared following Blitz and Schauder (5) [m.p., 206-8" (d.)]; stock solutions (1 or 2 mM) were prepared in water and stored for only 1 day because of the reported instability of the acid in both alkaline (6) and acidic (7) solution after extended periods of time. Buffer solutions were prepared from analytical reagent grade chemicals. Water-saturated argon was used for deoxygenating the test solutions. Polarographic and Voltammetric Procedures. Test solutions were prepared by diluting a known volume of stock solution with an appropriate buffer solution to give a final ionic strength of 0.5M;the pH was measured. Appropriate volumes were transferred to the H-cell, oxygen was removed by purging with argon, and the polarogram or cyclic voltammogram was recorded in the usual fashion. RESULTS AND DISCUSSION
Stability. In pH 11.2 chloride-hydroxide solution, the height of the oxaluric reduction wave decreased noticeably in the course of an hour and had almost disappeared after about 2 hours, presumably because of hydrolysis of oxaluric acid (I) to oxalic acid and urea (6): "2
COOH I+H2O-+LO+
I
1
COOH
(1)
2"
Below pH 11, oxaluric acid appeared to be relatively stable, at least for 30 to 60 minutes, the time necessary for three to four replicate polarographic determinations. Oxalic acid is not reduced at the D M E over the normal pH range. (4) L. Chuang, I. Fried, and P. J. Elving, ANAL. CHEM.,36, 2426 (1964). (5) H. Blitz and H. Schauder,J. Prakt. Chem., 113,77 (1926). (6) A. Strecker, Ann., 52, 113 (1860). (7) H. Blitz and G. Schauder, 1.Prakt. Chem., 106,147 (1923).
Polarography. Between pH 1 and 5, oxaluric acid exhibits a single, well formed polarographic wave, which,, although close to background discharge, presents no diffir culties in diffusion current and half-wave potential measurement (Figure 1). In pH 6.0 acetate and pH 6.1 McIlvaine solutions, the two waves produced are SO close that measurement of their individual wave heights and half-wave potentials (Eli2)was at best approximate; the more positive wave was about one fifth the size of the more negative wave. For convenience, the height and Eli2 of the composite wave were measured; the composite current values line up satisfactorily with those for the single wave, indicating that an additional faradaic contribution is not involved and that the two waves result from a split into two of the single wave seen at the lower and higher pH. Between pH 1 and 6 , E112 varies linearly with pH: ElI2= -0.895 -0.108 pH. Above 6.1, a single essentially pHindependent wave is again produced: Ell2 = -1.650 i 0.080 (mean f range); the variation probably reflects changes caused by the buffer composition. The diffusion current constant, Z = il/Cm2/3t1/6, is constant between pH 1 and 11 at 3.78 f 0.37 (mean rt standard deviation; the magnitude of the deviation is largely due to the values at pH 4.7 and 7.0); the value of Z indicates that the over-all faradaic process involves two electrons per molecule of electroactive species. The temperature dependence of the current between 25" and40"C (1.4,1.9,andl.6z per degree in pH 1.15 chloride, pH 6.0 acetate, and pH 8.0 McIlvaine buffers, respectively) indicates that the waves are diffusion-controlled, except, perhaps, for a slight kinetic effect at p H 6.0 acetate, where the kinetic nature of the first wave would be minimized by its being only a sixth of the composite wave at 25" and about a third at 40" C (the latter increase corresponds to 4.5z per degree). The dependence of the limiting current on drop time in these media also supports effective over-all diffusion control-i.e., linear dependence on h1/2. The slope of the wave, based on the calculation of (Elir - Ea/J = 0.056/n for a reversible process at 25" (cf. also Equation 3), approaches at pH 3 to 5 and 9 to 11 that expected for a reversible l e process (Table I). However, cyclic voltammetry at the hanging mercury drop electrode showed no reversible couple at any pH. Oxaluric acid was neither VOL 40, NO. 3, MARCH 1968
493
“I
Table I. Effect of pH and Background Composition on Rate-Determining Step in Polarographic Reduction of Oxaluric Acid (Ell4
Backgroundo Chloride
volt
1.1 2.1
0.108 0.108 0.065 0.065
2.9 Acetate McIlvaine Ammonia KOH-KCI
- Ed412
PH
3.7 4.7 6.0 6.1 7.0 7.9 9.2
11.2
0.055
0.105 0.095 0.098 0.085 0.040 0,065
B
P*
anao 0.52
0.59
1 .o 1 .o 1.6 1.6 1.8 1.o 1.1
0.58
d
0.52 0.87 0.87 1.oo 0.54
7l-
--
II:
0.66 1.4 0.87
All backgrounds 0.5M in ionic strength. Average value of three replicate determinations. c See text for method of calculation. d Above pH 6, wave was pH-independent; hence, p
5
0
I
I
0.2
0.4
1
I
0,6
0.8
I
I.o
8-
b
0
=
0.
a -
A
+a
3
-
V
oxidized nor reduced in the available potential range on cyclic voltammetry at the pyrolytic graphite electrode. Concentration Dependence and Analytical Utility. Measurement in pH 0.75 chloride and pH 4.7 acetate solutions over the range of concentration normally used in polarography showed a nearly linear dependence of concentration on wave height, especially in the more acidic solution (Figure 2); these two pH values were selected because the currentconcentration relation, as judged by the diffusion current constant values, seemed normal at pH 0.75 and low at pH 4.7. Actually, the experimental points at each pH form a separate smooth curve, whose projection passes through the origin with a more pronounced initial curvature at pH 0.75 than at pH 4.7. The curvature of the limiting current-concentration relationship is best seen in the variation of the ratio of il/C with C, which is plotted on a somewhat magnified scale in Figure 2 with a minimum 2% variation assumed for current reproducibility. The systematic decrease of the il/C ratio with increasing concentration is probably associated with the natures of the electroactive species and of the electrode processes involved at the different pH values, including acid-base equilibria. From the viewpoint of quantitative analysis, oxaluric acid can be satisfactorily determined by the preparation of a calibration curve for the concentration region of interest. The pH used for analysis can be selected in the first place on the basis of a pH where the oxaluric acid wave does not fall in the same potential region as waves due to other electroactive components in the sample. Where a choice of pH is available, that pH should be selected where the oxaluric acid wave has the most readily measured shape as well as a satisfactory il/Cratio-e.g., pH 0.75 would normally be preferable to pH 4.7 in the range of 0.2 to 1 m M oxaluric acid, whereas pH 4.7 might be preferable for concentrations below 0.2 m M because of the more nearly linear current-concentration relationship in that region. The pH region in which the wave is split-Le., around pH &-should generally be avoided for quantitative analysis. The spread of the currents at each concentration reported in Figure 2 was similar to that generally encountered in polarography-i.e., f 3% or less of the mean. 494
0
ANALYTICAL CHEMISTRY
CONCENTRATION, m M
Figure 2. Variation with oxaluric acid concentration A . Limiting current B. Ratio of limiting current to concentration I. pH 0.75 chloride solution II. pH 4.7 acetate solution Each point is the mean of three replicate determinations
MECHANISM OF OXALURIC ACID REDUCTION Calculation of the an, value both by a log plot method (8) and by utilizing the (E114 - E310 wave slope, using Equations 2 and 3 (based on average current during the drop life at 25”), respectively, Ed.e.
E
E112
Ell4-
0.05915 i log .- ___ an, id - i
E314
= O.O564/an,
(3)
gives essentially identical results ; data based on Equation 3 are reported (Table I). Since the slope was measured on a composite wave near the pK. region (cf. subsequent discussion), the values in that region (ca. pH 6 ) have little significance. The rate of change of Ellzwith pH along with ana data can be used to calculate p , the number of hydrogen ions involved in the rate-determining step (8): (4)
(8) L. Meites, “Polarographic Techniques,” 2nd ed., Chap. 4, Interscience, New York, 1965.
Values o f p so calculated are reported in Table I. On the basis of CY being ca. 0.5, the data suggest a one-electron, one-proton rate-controlling process at pH 1 to 2, a twoelectron, two-proton controlling process between p H 3 and 5 , and a return to a one-electron but still two-proton process at pH 6 ; the rate-controlling process is pH independent above pH 8 and seems to involve one electron at pH 8, but two electrons at higher pH. Between pH 1 and 5 , therefore, the rate-determining step can be represented by Equations 5 and 6 (I represents oxaluric acid):
coo-
”2
I
I
H
fast
I
I
(5)
I+H++e------t N
I11
+ 2e
slow
rate detg.
At pH 1 to 2, the process represented by Equation 5 would be followed by a second rapid one-electron, one-proton reduction to I1 (glyoxyllic acid monoureide). Since the dissociation constant for oxaluric acid could not be located, a so-called polarographic dissociation constant or pK, of ca. 5.3 was estimated from the relative heights of the acid and anion reduction waves at pH 5.98 and 6.10 and the lowering of the acid wave at pH 4.7, which is approximately 17% of the average value (the anion wave is masked by background discharge). Such calculated pK. values are usually somewhat higher than the actual value because of the tendency to overestimate the undissociated acid concentration, since the wave due to reduction of the latter contains an increment arising from the recombination in the layer adjacent to the electrode of anion with protons during the life of the mercury drop as the acid form concentration at the electrode surface is decreased because of its reduction; the anion concentration is, correspondingly, underestimated. The operational or conditional pK, of oxaluric acid is accordingly about 5 or slightly less. Kinetic factors in the reduction of acids have been abundantly discussed-e.g.,
l H
fast I
(7)
coo-
YHZ
H
COOH
i
+ e + 2H+
anion of I1
COOH
“2
yoo-
1
+ 2H+
(8)
anion of I1 The process represented by Equation 8 may be more complicated than indicated, since the reduction product of the oxalurate anion(II1) is shown as a dianion at the reduced carbonyl group. While such a dianion has been shown to exist in the reduction of a carbonyl group in pyridine (IO), its existence in a proton-available medium such as water would be at most very short and protonation would be practically simultaneous with the acquisition of the second electron. At pH 5 to 6, the acid will be only partially dissociated and both the acid reduction (Equation 5 or 6), producing the first wave, and the anion reduction (Equation 7 or 8), producing the second wave, will occur with resultant wave splitting. Consequently, for the composite wave is slightly high with respect to the pH-dependent region and slightly low with respect to the pH-independent region; the slope of the composite wave would be similarly distorted. The small acid wave at pH 6 (most positive wave, Figure 1) is indicative of the degree of dissociation of the acid and the relatively slow recombination of proton and anion under the exgerimental conditions. The low value of the diffusion current constant in pH 4.7 acetate solution is also indicative of the degree of dissociation and slow recombination; since the anion reduction occurs at potentials on or behind background discharge, only the wave due to reduction of the undissociated acid is observed.
(9).
At pH 6 and above, the acid would be almost completely dissociated and the reduction probably proceeds according to Equation 7 or 8: (9) R.BrdiEka, V. Hanu:, and J. Kouteckf, “Progress in Polarography,’’ P. Zuman, Ed., Vol. I, Chap. VII, Interscience, New York, 1962.
RECEIVEDfor review August 18, 1967. Accepted January 2, 1968. The authors thank the National Science Foundation and the Office of Research Administration of The University of Michigan, which helped support the work described. (10) R. F. Mitchielli and P. J. Elving, J. Am. Chem. Soc., in press. (1968).
VOL 40, NO. 3, MARCH 1968
495
