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ANALYTICAL CHEMISTRY
difficulty will be encountered in handling this problem should it arise. The following procedure was carried out to test for the actual 'presence of sodium bisulfite in the solution as it arrived from 'Oak Ridge. The sample tested was stated to contain in a 0.045 N basic solution of sodium bisulfite. The 0.1-ml. aliquots of the solution which were examined would therefore have contained 468 micrograms of sodium bisulfite. The aliquot was diluted to 5.0 ml. and, after the addition of 1 drop of sulfuric acid, waa shaken immediately with an equal volume of carbon tetrachloride containing in one instance 5.0 micrograms and in another 20 of iodine. The distribution of the radioactivity between the water and carbon tetrachloride was determined, and the radioactivity found to be very largely in the carbon tetrachloride in both cwm. However, when
50 micrograms of sodium bisulfite from a freshly prepared solution were added and the same water and carbon tetrachloride were reshaken, the radioactivity was found to be almost entirely in the aqueous phase. A 0.1-ml. aliquot thus failed to contain sufficient sodium bisulfite to reduce as little as 5.0 micrograms of iodine. LITERATIJRE CITED
(1) Burgus, W. H., and Davies, T. H., U. S. Atomic Energy Commission, Declassified Documents 2549-C (1949). (2) Grdak, B., Biochem. Z., 270, 291 (1934). (3) Leblond, C. P., and Sue, P., Am. J. Physiol., 134, 549 (1941). (4) Sandell, E. B., and Kolthoff,I. M., Mikrochim. Acta, 1,Q(1937). RECEIVED July 14, 1D4R.
Reund- Table Discussiens One of the special features of the program of the Division of Analytical Chemistry at the national meeting of the A.C.S. in Atlantic City last Septemberwas the inauguration of round-table discussions. Realizing the potential interest and value in this type of meeting the division's officers arranged for three discussion groups. The subjects covered were Polarographic Behavior of Organic Compounds, The Karl Fischer Reagent for Use in Determination of Water, and Determination of Carbon in Ferrous Alloys. All three sessions were well attended.
The splendid work of the moderators and their panels and the active participation of the audience in each case resulted in stimulating sessions. Analytical Chemistry obtained complete stenographic reports of the discussions and turned them over to the moderators for condensation and rewriting. We are happy to publish the resulting reports on the following several pages. We also take this occasion to express the hope that this type of round-table sessions with lively discussion from the floor will be continued and expanded at divisional meetings.-The Editors
Polarographic Behavior of Organic Compounds Moderator: PHILIP J. ELVING, The Pennsylvania State College, State College, P a . Panel: OTTO H. M ~ L L E R Syracuse , University, Syracuse, N . Y . STANLEY WAWZONEK, Iowa State University, Iowa City, Iowa MELVIN J. ASTLE, Case Institute of Technology, Cleveland, Ohio
LOUIS MEITES, YaG University, Xeuj Haven, Conn.
T
HE round-table discussion centered around the methodology for studying the behavior of organic compounds by polarographic measurements. The role of various factors and techniques was considered in the light of recent developments in polarography. The principal topics considered included: (1) the effect of environmental factors such as ionic strength, buffer capacity, specific buffer components, maxima suppressors, and pH; (2) criteria for reversibility ; (3) the deduction of electrode reaction mechanisms from the variation of half-wave potentials with pH, and the value of n, the number of faradays of electricity involved in the reaction a t the electrode of 1 mole of the active species; (4)factors determining reducibility or reactivity a t the dropping mercury electrode; and ( 5 ) current-controlling processes. ROLE OF THE ENVlRONMENT
In his discussion of the role of the environment, Muller pointed out that the usual excess of indifferent electrolyte ensures a high ionic strength in the solution and that the so-called excitation phenomena are not too well described and have not been too important in practical analysis. In discussing buffer systems, difficulties were described which may arise when the rate of equilibrium attainment in the buffering system is slow compared to the limited time spans involved in using the dropping mercury electrode; specifically, the failure was pointed out of bicarbonate
buffers to react fast enough to maintain pH constant at the mercury-solution interface. In unbuffered solutions, quinhydrone gives two separate waves, because the pH a t the electrode varies from 2 to 10 as a result of the electrolytic processes which consume or liberate hydrogen ions. In the presence of a buffer of low capacity the curve obtained is something in between those obtained in buffered and unbuffered solutions; actually, the curve represents an acid-base titration at the electrode face. The need for more buffers in the alkaline region was mentioned. Discussion brought out the general nature of the phenomenon occurring in solutions of poor buffering capacity-Le., the interchange of hydrogen ions between the constituents of the buffer and the solution s e e m to be generally rather slow. Even with perfectly adequate buffers, the slowness of dissociation and association of the reactants and products may be the cause of the appearance of several waves if the reactant that is being reduced is involved in the slow process. It was emphasized that buffers that are suitable for use with the dropping mercury electrode may be ineffective because of inability to reach equilibrium in time when used with electrodes in which the reactions have to go on much faster. Such electrodes include the rotating platinum electrode, the bypass electrode in which the solution streams past a microelectrode, and the systems used in cathode-ray oscillographic polarography. The inability of the
V O L U M E 2 2 , NO. 3, M A R C H 1 9 5 0
483
buffer to maintain the p H constant a t such electrodes may be due to rate phenomena rather than to buffering capacity. The effect of the environment, Astle emphasized, may account for the lack of agreement between different investigators on polarographic data, particularly concerning the half-wave potentials of compounds reduced at the dropping mercury electrode. The latter potentials will vary with the nature of the buffer constituents, even though the pH is kept constant. Addition of small amounts of an organic solvent such as alcohol may result in profound change in the reduction potential of a compound. The substituted benzaldehydes give two waves in a buffer containing glycine but only one wave in buffers of the same pH containing phthalate or citrate. Thus possibility of interaction of the buffer ions with the reactive system should be considered. In the reduction of methylene blue in the redox system, methylene ieuco-methylene blue, in a phosphate buffer, the lower part of the curve obtained in an acetate buffer is missing; this is apparently due to the formation of an insoluble product of the phosphate with the leuco form of the methylene blue. It was pointed out that the type of minimum which had been reported in the polarographic wave of phenolphthalein could also be observed in substituted aromatic nitro compounds in glycine buffers (but not in citrate buffer), in certain inorganic compounds, and in almost any system by the addition of sufficient surface-active material. Discussion emphasized that a buffer would provide adequate capacity only if used a t its pK value where pK represents the negative logarithm of the ionization constant of the active component of the buffer system-that is, the effective buffering range of a buffer system is the pH range of (pK - 1) to (pK 1 i.
+
pound of similar size; (2) the chemical properties of the compounds concerned; (3) the variation of half-wave potentials with pH; and (4) the slope of the wave. The determination of n is carried out most exactly by coulometric measurements during a reduction a t controlled potentials and is the most aceurate method when the number of electrons involved is large, as with picric acid. In simple cases involving only a few electrons per unit, n can be determined by substituting estimated values of the diffusion coefficient in the IlkoviE equation. The value of n will give an indication of the final product formed a t the electrode. Thus, if n is 4 in the reduction of nitrobenzene, the product formed must be phenylhydroxylamine. The chemical properties of the compounds concerned are important. The products postulated as being formed a t the dropping mercury electrode should not be reducible further a t the electrode or should not react with the electrolyte used. An example in which this point may be important is the reduction of picric acid in 0.1 A' hydrochloric acid. Coulometric measurements indicate a 17-electron change. Reduction to triaminophenol requires 18 electrons. To explain these facts, it has been suggested that the reduction proceeds to the substituted hydrazobenzene. 4 question can be raised about the stability of this product toward 0.1 -1' hydrochloric acid, since similar conditions bring about the benzidine rearrangement with hydrazobenzene itself. The benzidine if produced would be stable to further reduction a t the drop, Pihile the hydrazobenzene represents a compound which can be reduced further, a t least chemically, to an amine. The shift of half-wave potentials to more negative values with an increase in pH can be explained by either a direct reaction of the compound with acid to form an intermediate ion which is more easilv absorbed or reduced,
+ + H' eRZC-OH ~ ~ 6 -+ 0e + ~ ~ ~ 2 - 0 ~
POLAROGRAPHIC REVERSIBILITY
In brginning the discussion of the criteria for reversibility in polarographic phenomena, Muller discussed the meaning of reversibility. The electrochemist studying nonpolarixable or indicator electrodes in oxidation-reduction systems considers it system reversible if the same potentials are obtained, whether the oxidized form is reduced to known amounts of the two forms or the reduced form is oxidized to the same ratio of the two forms; if a constant and reproducible value is obtained for the redox potential in a 50:50 mixture of oxidized and reduced forms, the system is considered reversible. The situation in polarography is confused, inasmuch as the end products of polarographic reductions are rarely polarographically oxidizable. The use of oscillographic techniques in presenting simultaneously both oxidation and reduction processes a t the mercury electrode seems advantageous in evaluating reversibilit.y, although oscillographically observed reversibility and irreversibility may not he identical to the ordinarily observed polarographic phenomena or to potentiometric measurement. It must be kept in mind that irreversible and reversible steps may accompany or follow each other. Thus, the reduction of oxygen to hydrogen peroxide is reversible, while the subsequent conversion of hydrogen peroxide t o hydroxyl ion or water is irreversihle. REACTION MECHANISMS
The discussion of the question of what actually occurs a t the dropping mercury electrode was opened by Wawzonek, who summarized the present status of the experimental approach t o the subject and the usual theoretical interpretation. Reaction mechanisms a t the dropping mercury electrode are best determined by isolating the products formed. In cases where such a procedure is not carried out, the factors important in deducing a reaction mechanism a t the dropping mercury electrode in their order of importance are: (1) the determination of n by coulometric measurements, from the IlkoviE equation or from a comparison with the polarographic behavior of a com-
RZC=O
or hy a reaction of the molecule directly followed by a reaction of t h e anion with water in the following manner: R2C=0
Rzc-0-
+ e +R2q-O-
+ H20 e R,;-OH
+ OH-
In both cases the intermediate radical may dimerize or be reduced further to the carbinol. The speed of the reaction of anion with water would determine whether the curve would be dependent upon pH or not. The slope of the wave gives an indication of the potentialdetermining step. This value turns out to be usually around 0.06 or higher. The higher values are usually considered as evidence of the irreversibility of the reaction. A possible mechanism of reduction, suggested by Pearson for nitro compounds which may be more general and which would give such values for the slope of an irreversible system involving hydrogen ions, is the following:
A
H + + e e H(Hg) (overvoltage) + nH +AH, (in alkaline or neutral solution) A + nH + H + +AH;+;,, (in acid solutions)
The potentials for such systems would be dependent upon the diffusion coefficientof the reducible material and upon its reaction rate with hydrogen atoms. The difficulty of answering questions concerning reversibility and reaction mechanisms in polarographic studies was stressed by Meites, who used examples of the behavior of inorganic substances to develop his discussion. Thus, surface-active materials of all types will, if present in sufficiently high concentration, shift the half-wave potentials of polarographic systems, the magnitude of the shift varying with the concentration of the surface-
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ANALYTICAL CHEMISTRY
active material-eg., the addition of sufficient gelatin to an alkuline tartrate solution (pH 14) of copper shifts the half-wave potential of the copper by 1.3 volts while the wave still retains its reversible character. The effect of surface-active materials in producing two-step waves was also considered. A theory was outlined which attempts to explain certain observed changes in half-wave potential and diffusion current on the basis of the formation of highly hydrated micelles whose concentration can be correlated with the addition of surface-active agent and the vapor pressure of water above the solution. The rate of micelle formation and difference in rate of diffusion of micelles of varying degree of hydration may account for some types of two-step waves. I t was suggested that rather than to try to suppress maxima by the addition of a substance such as gelatin, m-hich might result in unpredictable changes in polarographic behavior of the substance under study, it would be preferable to remove the maxima by increasing the dilution so that the concentration of the active species was not more than a few hundredths millimolar. It was brought out in subsequent discussion that readily interpietable waves are very difficult to obtain for organic (*ompourids
when present in concentrations below 0.1 to 0.5 millimolar. I t was also indicated that pH will alter the effect of varying the concentration of gelatin in a solution. Arguments were advanced for and against the explanation of the effect of gelatin as being due to adsorption of the gelatin on the surface of the drop. In a discussion of the coulometric measurement of n values, microcoulometric cells using dropping mercury were considered as a possible way of evading the assumption usually made that the behavior at a large stirred surface is comparable to that which occurs at the mercury drop. I t was pointed out that one could not always depend on stirring by the falling mercury in the microcoulometric cell to avoid local depletion of the solution. Among those taking extensive part in the discussions were J. K. Taylor, Sational Bureau of Standards; G. A. Crowe, Hercules Powder Company; Louis Lykken, Shell Development Company; W. W. Davis, Eli Lilly & Company; K. L. Metcalf, E. I. du Pont de Nemours & Company; and P. A. Geary, Smith, Kline & French Laboratories. RECWVED December 19, 1949.
Karl Fischer Reagent for Use in the Determination of Water 3Ioderator: JOHN MITCHELL, JR., E. I . d u Pont de Nemours 6% Company, Inc., Polychemicals Department-
Chemicals, Wilmington, Del. Panel: L. R. KhNGAS, Hercules Powder Company, Wilmington, Del.
WILLIAV S E i l I i N , Calco Chemical Dicision, American Cyanamid Company, Bound Brook, N . J .
T
1IE Karl Fischer reagent is a solution of iodine, sulfur dioyide, and pyridine in methanol. Each of these components enters into the basic two-step reaction for water. Normally, sulfur dioxide, pyridine, and methanol are present in excess. Therefore, the strength of any preparation is dependent on the iodine concentration. For general laboratory use the reagent may be prepared to contain the components in the molar ratios, 112:3S02:10CbHJ ill methanol, a t a concentration equivalent to 3 to 4 mg. of water per nil. of reagent. This composition gives a reagent suitable for the titration of samples of extremely low to high water concentrations, minimizes some types of interference, is a good general solvent, does not degrade to darkly colored end products, and assures an excess of components other than iodine. However, for specialized purposes other ratios may be better suited-for example, a reagent containing a higher concentration of pyridine (20) is better suited for the determination of water in acetone. Of course, the same effect can be obtained by employing pyridine as an inert diluent for the sample. A preparation containing less pyridine, such as 4 to 5 moles per mole of iodine, is satisfactory for analyses for water in alcohol. A reagent of about one third the water ryuivalence is useful for titrations of samples containing trace quantities of water (8). Contrary to some reports, the Fischer reagent may be employed for the titration of samples containing amounts varying from minute traces of moisture to pure water. Obviously, as the concentration of water becomes high more care must be exercised i n weighing. Normally, titrations are made by delivering the complete reagent from a desiccant-protected buret into the sample to be analyzed. Because the reagent is subject to parasitic side reactions, which effect a gradual reduction in strength, daily standardizations are necessary. These can be made by titrations of a metha-
nol or ethanol solution containiiig a known quantity of water, water-saturated alcohols, a weighed amount of water, or a stable hydrate. Incidentally, the reagents prepared to be equivalent to only 1 to 2 mg. of water per ml. are more stable than preparations of higher iodine concentrations (8). Thus, Wiberley reported that a reagent equivalent to 1 mg. of water per milliliter was stable. By employing desiccant tubes packed with phosphorus pentoxide on asbestos, he found that usually only weekly standardizations were necessary (21). Reagent left in the buret appears to degrade more rapidly than that in the reservoir, possibly because of the increased exposure to light (18). Fischer reagent is an extremely aative desiccant. Adequate protection against exposure to atmospheric moisture must be maintained for both the stored reagent and the reagent being delivered from a buret into the sample to be analyzed. In areas of high humidity titration in a closed system might be beneficial. A rubber sheet or polythene film can be used to cover the titration flask. Bdequate protection is maintained by inserting the buret tip through a hole in the sheet (6, 16). -4modification of the titrimetric technique was suggested by Johannson (S), who proposed a divided reagent. A solution of pyridine and sulfur dioxide in methanol was added to the sample flask, and a second solution of iodine in methanol was delivered from a buret. On contact with the contents of the sample flask active Fischer reagent was formed which reacted immediately with the water in the sample. In this way a stable reagent was obtained. Seaman and his co-workers (16) made a more thorough study of this method, and found that the stable iodine in methanol solution could be standardized against sodium thiosulfate. No evidence of significant degradative side reactions wm observed, provided no great excess of iodine wm added over that required t o react with all the water in the flask. Several investigators have mentioned some difficultyin finding
