Determination of Hexaphenylethane GEORGE S. HAMMOND, ABRAHAM RAVVE', AND FRANK J. >IODIC2 Iowa State College, Ames, Iowa
U T H E course of a study of the reactions of the triphenyl'-methyl free radical it became necessary to have available some rapid, convenient method for the determination of the total concentration of the radical and its dimer in solution. Procedures ( 1 4 ) have been developed to accomplish this objective but the method presented here has been found particularly convenient and accurate. The oxygen absorption procedure (a)has been used in this laboratory with good results but its use in routine work requires the maintenance of a bulky gas absorption apparatus. The method of Preckel and Selwood ( 4 )is somewhat tedious. The authors have found that iodometric analysis, even with the Bachniann modification ( I ) , is somewhat less precise than the piesent method. This arises from the inconvenience, due to loss by volatilization, of handling iodine solutions in evacuated systems (necessary to prevent air oxidation). Incidental to this study of the iodometric method it was found that the presence of pyridine in the titration mixture renders the starch-iodine end point very uncertain unless the iodide concentration is maintained a t .I high level.
,-, -?I .
Figure 1.
Apparatus Used for Preparation of Free Radical Solutions
preparation, 6 grams of tIiphenylchloromethane, 10 grams of mercury, and 50 ml. of redistilled, sodium-dried, thiophenefree benzene were placed in compartment A . The apparatus was then evacuated to water pump pressure throu h the outlet B, filled with purified nitrogen (pyrogallol train? through C. The process was repeated twice and B and C mere sealed off. The mixture was then agitated for 22 hours with an all-glass stirrer, D. (Complete conversion of the halide to the free radical was observed to occur sooner in other runs; however, no attempt was made to find an optimum stirring time.) The ball joint, E, which was lubricated with mineral oil, was found to provide an effective seal for periods of several days. When the stirrer was stopped, the mercury and mercurous chloride were allowed to settle and the apparatus was tilted, permitting the solution to filter by gravity through the medium porosity, sintered-glass disk, F. Samples were removed u ith nitrogen-filled, calibrated syringes which were inserted through a rubber serum bottle stopper placed in a nipple a t G. A volume slightly larger than the desired aliquot was withdrawn and the volume was adjusted in the usual manner. The aliquot was then injected, through a serum stopper, into a 125-1113. Erlenmeyer flask which contained a known amount of peroxide. The reaction flask had previously been repeatedly evacuated and flushed with nitrogen. The peroxide was conveniently added as an aliquot from a standardized benzene solution. With the concentrations of peroxide and radical comparable to those reported below, the reaction was complete within 15 minutes. With more dilute solutions, it was necessary to allow longer periods of time for complete reaction. The excess benzoyl peroxide was determined iodometrically. The flask was opened and a few small pieces of dry ice, 25 ml. of Baker and Adamson C.P. glacial acetic acid, and 2 to 4 ml. of saturated aqueous potassium iodide were added in that order. The flask was swirled and allowed to stand for 5 minutes. The iodine liberated was titrated with a standard thiosulfate solution, with 100 ml. of distilled water and 2 ml. of 1yostarch solution being added as the end point was approached. Reaction of Triphenylmethyl with Potassium Iodide in Acetic Acid. When benzene solutions of triphenylmethyl were added to potassium iodide in acetic acid, varying amounts of iodine were liberated. The details of this work are not yet complete. However, it is important to note the necessity of ensuring the use of an excess of benzoyl peroxide in using the above analytical procedure. Otherwise the oxidation of iodide by excess triphenylmethyl will give spurious low values for the free radical. Results. The reproducibility of the results is indicated in the data summarized in Table I. The figures were obtained by analysis of 1-ml. aliquots of a single radical preparation, with varying amounts of peroxide and solvent. T h c rcwlts Kcre further checked by analyzing 2-ml. aliquots
It developed that one of the reactions which were being studied was, in itself, suitable for use as an analytical'method. Benzene solutions of triphenylmethyl react quantitatively a t room temperature with benzoyl peroxide according to the stoichiometric relationship,
Table I. Radical Prep 1
CiaHioOa = S(CQ"),C The products of the reaction, triphenylmethylbenzoate, benzoic acid, and tetraphenylmethane ( 3 , e), are formed in amounts which vary with the concentration of reactants ( 3 ) . However, the quantitative nature of the peroxide consumptiori is unaffected by this variation.
MI. in
0 421 .li Peroxide
Moles Radlcal/RIl x 104
Soh,
,\I1 in no
fi
Average
fin
7 00 6.70 6.72
+ 0.18
T a b l e 11. Comparison w-ith Analysis by- Oxygen Absorption
Preparation of Free Radical. The free radical solutions were prepared in the apparatus shown in Figure 1. In a typical
Radical Prep.
_2
Vol. of Assag Mixture,
2 00 2 00
EXPERVM E 1 l A L
I
Analysis of Solutions of Triphenylmethyl
2
2
Preqent address, Illinois Institute of Technology, Chicago, Ill Present addrew General Electric C'o Waterford, X Y
1373
Equivalents Radical/2-M1. Aliquot X 103 Oxygen Peroxide uptake assay 1.31 1 31 1.28 1.31
ANALYTICAL CHEMISTRY
1374 by both oxygen (2) absorption and this method. Table, I1 summarizes the data. Accuracy. Duplicate analyses in which larger aliquots of radical solutions were taken usually gave results which agreed more closely than would be anticipated from the data presented above. For this reason, and because of the difficulty involved in inserting the syringe through the serum stopper without occasionally losing a drop of solution, the authors believe that the accuracy of the method is limited by the inaccuracy in estimating the volume transferred by the syringes which are relatively crude volumetric apparatus. They believe that further improvement can be attained through the use of very fine needleand narrow-bore, precision syringes.
ACKNOW LEDGRIENT
The authors wish to acknowledge the financial support of this work by the Offic,e of Naval Research. LITERATURE CITED
(1) Bachmann. lt-. E., and Osborn, G.. J . O T ~Chem., . 5, 29 (1940). (2) Gomberg, hl., and Schoepfle, C . S., J . Am. Chem. SOC.,39, 1661 (1917). (3) Hammond, G. S.. Rudesill, J. T . , and Modie, F. J., Zbid., 73, 3929 (1951). (4) Preckel, R.,and Selwood, P. IY.,Ibid., 63, 3397 (1941). ( 5 ) Wieland, H., Ploetz, T., and Indest, H., Ann., 532, 175 (1937). (6) Ziegler, K., Orth, P., and Weber, IC., Zbid., 504, 131 (1933). RECEIVED for review February 27. 1951.
.ircepted May 12, 1952
Polarographic Determination of Iron in Nonferrous Alloys LOUIS MEITES, Yale University, New Haeen, Conn. H E polarographic literature contains descriptions of two Tmethods for the determination of iron. One ( 2 ) is based on the measurement of the height of the iron(II1) wave in an acid solution containing no complexing agent. This method is subject to interference by other ions which are also capable of oxidizing mercury (14). The other ( 6 ) requires separation of copper and reduction of iron to the +2 state, the concentration of ferrous iron being determined by measuring the height of its anodic wave in 1 F potassium oxalate a t a pH near 5. This method is specific for iron and is capable of giving very accurate results. It has not, however, been very widely adopted, presumably because of the care necessary to prevent air-osidation of the ferrous complex. The present communication describes a method which is free from both of these drawbacks. The final solution contains the iron in the form of a mixture of citrate complexes which are inert to both air oxidation and reduction by mercury, so that a rrelldefined wave is secured whose height remains constant for long periods of time. EXPERIMENTAL
All polarographic measurements rvere made with a calibrated Sargent-Heyrovskj. Model XI1 recording polarograph, using the visual scale of the instrument. A conventional H-cell modified by the insertion of a sintered-glass gas dispersion cylinder in the influent gas stream to permit rapid deaeration was used (IS). The hydrogen used for deaeration was freed from oxy en by passage through two vanadium(I1) perchlorate wash tottles (11, 12). A nTater thermostat was used to maintain a temperature of 25.00' f 0.03' C. A Beckman Model G pH meter was used for pH measurements, and a saturated solution of potassium hydrogen tartrate was used as the pH standard (1,s). All volumetric apparatus was calibrated by conventional methods. A stock iron(II1) solution was prepared by dissolving electrolytic iron in perchloric acid, eva orating to fumes of the acid, and diluting to known volume. T i e concentration of this solution was checked by reduction in a Jones reductor and titration with standard permanganate. All other chemicals were ordinary reagent grade and were not further purified except as described below. DATA AND DISCUSSION
The investigations of Lingane ( 5 ) and Meites (7-10) of the polarographic characteristics of iron( 111) and copper( 11) in tartrate, citrate, and oxalate media indicated that, in these media a t least, it was impossible to secure a satisfactory separation of the waves of these elements. This is also true in solutions containing certain other organic anions which form iron(II1) complexes sufficiently stable to have measurable half-wave potentials. ..le-
cordingly, it appeared that a satisfactory general method for the determination of iron must involve the separation of copper before the polarographic measurement. The most rapid and convenient method found for effecting this separation consists of the reduction of copper(I1) to the metal with amalgamated zinc. In a weakly acidic solution this reduction proceeds to completion within 30 seconds or less, which is a considerable advantage when many samples are to be analyzed. The reduced solution and the precipitated copper can easily be removed from the residual zinc by decantation, so that the zinc can be used for the reduction of another sample. Bismuth is quantitatively removed viith copper, which is of importance because bismuth also gives n-aves a t relatively positive potentials in the media under consideration, and these viould also interfere vith the iron wave. Attempts were made to apply this method of separation to solutions containing nitrate, such as would result from the solution of an alloy in a nitric-hydrochloric acid mixture. However, it R-as found that nitrate seriously impaired the efficiency of the reduction: Large amounts of nitrogen dioxide r e r e evolved unless the pH was so high that hydrous metal oxides precipitated. It is therefore necessary to remove nitrate from the solution priorto the treatment with zinc by a double evaporation to near dryness with excess hydrochloric acid. This both prevents the precipitation of tin(1V) oxide, which would retain appreciable amounts of iron, and causes the reduction with zinc to proceed smoothly. The reduced solution, after filtration through a coarse quantitative filter paper to remove metallic copper, bismuth, and silver, contains iron and tin in the +2 states. These are reoxidized by the addition of a small excess of potassium permanganate solution, and the excess permanganate is immediately destroyed by the addition of oxalic acid. A number of experiments were made in which a large excess of oxalic acid was added to serve as the supporting electrolyte. However, the difficulties encountered in this method outweighed the expected advantages. Much of the zinc precipitates, carrying down iron with it, and the re-solution of the zinc oxalate Fas troublesome and inconvenient. Tartaric and citric acids do not give insoluble zinc salts under these conditions, and are therefore to be preferred to oxalic acid. As there seemed to be no other grounds for choice, citric acid was selected because of its slightly lower cost. The reduced solution is accordingly treated with sufficient citric acid to make the final citrate concentration approximately 0.25 F , and the p H is adjusted to a value near 5 by the addition of concentrated ammonia. The diffusion current constant of iron(II1) in citrate media, measured a t -0.55 volt us. S.C.E., is independent of pH between about 3 and 6.7, so that the exact p H
