when a mask was inserted into the Analytrol to accommodate 1-mm. strips, the instrument was insufficiently sensitive. Wider strips or a more sensitive instrument should make this procedure practical. ACKNOWLEDGMENT
The authors thank Joseph Dorsa and Robert Palacios of the Research and
Development Laboratory for helpful suggestions in the design of the electrophoresis apparatus. LITERATURE CITED
(1) Block, R. J., Durrum, E. L., Zweig, G., “Manual of Paper Chromatography and Paper Electrophoresis,”
Academic Press, New York, 1958. (2) Edstrom, J. E., Nature 172, 908 (1953).
( 3 ) Karler, A,, Brown, C. L Kirk, P. L., Mikrochzm. Acta 1956, 1585. (4) Strain, H. H., ~ ~ N A LCHEM. . 30, 620 (1958). (5) Turner, B. M., Mikrochim. Acta 1958, 305. (6) Brown, C . L., Kirk, P. L., Ibid., 1956, 1729.
RECEIVEDfor review April 29, 1959. Accepted November 9, 1959. Work supported in part by U. s. Public Health Service Grant C43341.
Precipitation of Crystalline Iron(lll) Oxide from Homogeneous Solution SIR: The preparation of granular or macrocrystalline precipitates from aqueous solution is generally limited to species with large lattice energies and relatively small hydration energies. Because multivalent metal ions possess large hydration energies and usually exist in solution as hydrolytic species [Fe(OH) +2, Fe(OH)i, etc.], the direct separation of most metal oxides from aqueous solution is often assumed to be precluded by the hydrophilic character of such oxides. However, there is in principle no limitation to the formation of macrocrystalline oxides in aqueous solution. Energetically, such oxides are stable and may be expected to form. Kinetically, the mechanism for their formation is usually sufficiently complex to prevent the separation of a granular product. We report the preparation, by direct precipitation from aqueous solution, of crystalline iron(II1) oxide as p-FeO.OH. The product is granular, easily filtered, and not peptized upon washing. The procedure appears to be the first preparation of the granular rather than the hydrous (gelatinous) oxide which has been effected rapidly a t ordinary temperatures and pressures. The primary significance of the procedure is the demonstrated ability to influence profoundly the relative rates of crystal nucleation and aggregation for the hydrous oxides. The separation of microcrystalline iron(II1) oxide by slow hydrolysis of concentrated iron(II1) solutions has been reported under a variety of conditions, but the precipitates are invariably hydrous and are finely divided and easily peptized upon washing. Recently Gayer and Wootner ( 3 ) reported the preparation, by slow hydrolysis of ferric nitrate solutions over a period of weeks, of a microcrystalline iron oxide which is not peptized upon washing. 566
ANALYTICAL CHEMISTRY
Precipitation from homogeneous solution is particularly advantageous for the hydrous oxides, but prior application of this method to the preparation of metal oxides has not alleviated the hydrous character of the precipitates. I n the present procedure, the precipitations were performed by neutralizing in a suitable manner acidic iron(II1) solutions containing a suitable complexing agent. Precipitation from homogeneous solution invariably provided the most satisfactory method for the separation, and the following procedure is recommended. Five to 10 grams of urea are added to 100 ml. of a 0.001M iron(II1) solution containing 0.01M hydrochloric acid and 0.002 to 0.02121 N,N-dihydroxyethylglycine (DHEG). The solution is rapidly brought to boiling temperature, after which the iron precipitates in 1 to 15 minutes, depending upon the concentrations of the reagents. The solution is removed from the heat upon incipient precipitation of the oxide and filtered through porous glass or porcelain crucibles. The precipitate may be washed repeatedly without peptization and can be dried with alcohol or acetone. The crystalline product has been identified by x-ray analysis as the beta monohydrate of iron(II1) oxide, 8FeO.OH. Although a chloride medium appears to be essential for the formation of the 8-FeO.OH, the beta modification appears to be a distinct (though perhaps metastable) crystalline species (4). and the interplanar spacings and relative line intensities observed in this study agree well with those reported by Weiser and Milligan (8) and Kratky and Xowotny ( 5 ) . Visually the precipitates appear as clusters of hexagonal amber platelets whose diameters are 2 to 4 microns. The clusters are considerably larger, on the order of 20 to 60 microns in diameter. Because of occluded water,
the freshly prepared oxide usually weighs 25 to @yo more than pure FeO.OH. The pyrolysis curve is similar to that for the iron hydroxides (sic) (2) with the transition from pFeO.OH to a-FepOB occurring between 150’ and 185’ C. (8). Some of the conditions favorable to the preparation of the granular iron oxide have been investigated in order to provide information concerning the mechanism of the precipitation. The formation of a granular product is critically determined by the rate of formation of the primary nuclei, which is influenced primarily by the rate of change of p H during the neutralization. ilt room temperature, the hydrolysis of urea is slow, and a quantitative precipitation requires several days. The precipitation may be effected rapidly a t boiling temperature, but the product is more sensitive to hydration than a t lower temperatures, even though precipitates formed from homogeneous solutions are much less susceptible t o peptization than those formed by other procedures. By heating the solution rapidly from room to boiling temperature, a compromise between an initially slow rate of nucleation and a rapid aggregation (growth) of the nuclei is obtained which may permit an immediate quantitative recovery of the granular oxide. Often, however, the precipitation of the granular product is only 90 to 98y0 complete, and a quantitative separation may require aging at boiling temperature, which if prolonged may tend to peptize the precipitate. Apparently, it is the nucleation of the uncomplexed iron(II1) Thich determines the granular character of the precipitate. The complexing agent D H E G is a weak acid with pK2 = 8.10 (r), and studies on the stability of the iron(II1)-DHEG complex (1, 7‘) have indicated that in the p H range 2
t o 4, the iron(II1) is incompletely CC,,Jplexed. The influence of the initial pH of the solution on the granular charactzr of the precipitate suggests that the crystal nucleation for dilute iron solutions is completed prior to hydrolysis of the iron complex. If the initial concentration of iron(II1) is increased apprcciably, a hydrous oxide is precipitated quantitatively, but under these conditions the granular oxide cannot be prepared hpcause DHEG is too weak an acid to complex an appreciable fraction of the iron(II1) initially. For concentrated solutions. the formation of the granular oxide will undoubtedly be favorcd by employing an agent which more strongly complexes iron(II1) than does DHEG. The tetradentate DHEG ligand also appcars to be important in facilitating the formation of the nongelatinous oxide. While the role of the ligand has not been investigated extensively, preliminary experiments indicate that only tetradentate ligands are suitable for the preparation of a granular oxide.
Co u I o met ric
The empirical formula of the iron(II1) complex is Fe(OH)2G, where G- represents the ligand anion. If the primary nuclei of (FeO.OH), (6) condense b y a reaction of the type
+
(FeO.OH), HO-FeG-OH HO-FeG-O-Fe-(FeO.OH),l
=
+ H20
it may be recognized that tetradentate ligands are peculiarly adapted to the formation of the crystalline oxide by precluding excess water from the coordination sphere of the iron atoms. Other conditions favorable to the formation of a granular oxide have been described ( 1 ) . Attempts to alter the composition and boiling temperature of the solvent b y the addition of ethanol, ethylene glycol, and glycerol were not successful because the precipitation is much less complete in the presence of such hydroxylic solvents. ACKNOWLEDGMENT
The assistance of the Geigy Chemical
Corp. in providing the N,N-dihydroxyethylglycine used in this study is gratefully acknowledged. LITERATURE CITED
(1) Benck, R. F., M.S. thesis, Univemity of Nebraska, 1958. (2) Duval. C., “Inorganic Thermogrsvi-
metric Analysis,” Elsevier, New York,
1953. (3) Gayer, K. H., Wootner, L., J. Phys. Chem. 60, 1564 (1956). (4) Hofer, L. J. E., Peebles, W. C., Dieter, W. E.. J. Am. Chem. SOC.68, 1953 (1946). (5) Kratky, D., Nowotny, H., 2. Kn’st. A100, 356 (1938). (6) Krause, A., Kolloid 2 . 7 2 , 18 (1935). ( 7 ) Toren, P. E., Kolthoff, I. M., J. Am. Chem. SOC.77, 2061 (1955). (8) Weiser, H. B., Milligan, W. O., Ibid., 57, 238 (1935).
E. R. NIGHTINGALE, JR. R. F. BENCK Department of Chemistry University of Nebraska Lincoln 8, Neb. RECEIVED for review November 25,1959. Accepted February 1, 1960.
Oxidation of Boron
SIR: The method of oxidizing or reducing films on metallic surface has been extensively employed for the study of surface coatings and corrosion films (2, 3). The film is oxidized or reduced at constant current, and the potential of the metal is measured during electrolysis. The completion of the film dissolution is indicated by a rapid change in the potential. The total number of coulombs calculated from the current and time of electrolysis corresponds to the quantity of film. During the course of high-temperature studies on boron carbide, it was necessary to analyze microgram quantities of elemental boron deposited on platinum targets. Conventional wet chemical methods failed to remove the boron completely from the platinum base. Coulometric oxidation was then attempted as a method for dissolution of the boron film. Good correlation was obtained between the number of coulombs used and the amount of boron removed from the target. The results reported here are not extensive because of the secondary nature of the investigation. But in light of the increased interest recently in the chemistry of boron, the results may be of value t o the study of boron deposits and boron in alloys. This is also believed to be the first report of the oxidation of elemental boron by electrolysis. EXPERIMENTAL
Constant current was provided by
~~~~~~
Table I.
Q
b
Exposure
Current, Ma.
Coulombs
Stripping
3-1 3-2 3-7 3-8 3-10 3-1 1 5-16
1.52 1.52 1.082
0.213 0.505 0.203 0.529* 1.595 0.734 0.416
8 0
1 306 0.998 1.065
~
Results of Anodic Boron Stripping
18.9 7.8 19.7 59.6 27.4 15.6
Boron Found, Titrations
y
Residual.
7.3 16.5 7.9 14.9 51.2 25,l 13.2
2.8 2.2 0.5 3.5 29.3 18.7 10.2
Accuracy 1.0%. Summation of five different current levels.
four 45-volt B batteries connected in series and dropped across a selector m i t c h bank of large standard resistances for different current levels. The current levels were determined by mezsuring the voltage drop across a standard IO-ohm precision resistance with a Leeds & Northrup portable precision potentiometer. The samples analyzed were targets coated with boron. The boron was deposited in vacuum from a Knudsen cell containing boron carbide at about 2300O K. The target itself was a platinum disk ll/s inches in diameter by 0.005 inch thick. The boronplated disk was placed, face up, in a 30-ml. quartz beaker containing 5.0 ml. of 0.10A’ sodium hydroxide. A threeholed rubber stopper covered the beaker. It also supported two saturated calomel electrodes (S.C.E.) and a platinum wire probe. The probe was used for making electrical contact Rith the disk. The saturated electrodes were of the type used by Adams et al. ( 1 ) ; one served as a working cathode and the other a reference electrode.
The potential of the platinum disk S.C.E. was monitored by a Leeds & Xorthrup Speedomax recorder adjusted for a full scale deflection corresponding to 2.0 volts. The recorder operated a t a chart speed of 2 inches per minute. After the coulometric oxidation end point was reached, the electrodes and target were removed and the resulting solution was analyzed for boric acid by coulometric titration of the mannitol complex (4). Residual boron left on the targets after the coulometric stripping was determined by burning the targets in air, dissolving the resulting boron oxide (B203)in 0.10A‘ sodium hydroxide, and titrating as above. Spectrographic analysis proved that the burning process completelv removrd all boron from the targets (4). us.
RESULTS
A well defined S-shaped anodic potential-time curve was obtained for the boron dissolution. The initial potential of 0.30 volt us. S.C.E. corresponds to
+
V O L . 32, NO. 4, APRIL 1960
567
