quickly at room temperature when the thiocyanate is brought in contact with RuOc as in this method, Nitric acid, which is a serious interference in most spectrophotometric methods for ruthenium, can be tolerated in this method. A number of analyses were made of a ruthenium nitrate solution containing 200 pg. of ruthenium per ml. Aliquots containing from 40 to 60 pg. of ruthenium were analyzed (Table 111). The standard deviation obtained is for the most favorable concentration range and may not apply to different ruthenium concentrations.
ACKNOWLEDGMENT
The authors gratefully acknowledge the suggestions of R. R. Rickard and P. F. Thomason during this work. Thanks are also due to Lonas Guinn, Grant Hickey, and R. L. Lewis for some of the data presented.
LITERATURE CITED
(1) Beamish, F. E., McBryde, W. A. E., Anal. Chim. Acta 9, 349 (1953); 18,
562 (1958).
(2) Furman, N. H., ed., “Scott’s Standard Methods of Chemical Analysis,” 5th ed., Vol. I, p. 740, Van Nostrand, New York, 1950. (3) Martin, F. 8.. J . Chem. SOC. 1954. 2564. (4) Sandell, E. B., “Colorimetric Determination of Traces of Metals,” 3rd ed., p. 779, Interscience, New York, 19.59. -__I.
(5) Surasiti, C., Ph.D. thesis, University of Minnesota, 1957. (6) Westland, H. D., Beamish, F. E., ANAL.CHEM.26,739 (1954). (7) Yaffe, R. P., Voigt, A. F., J . Am. Chern. SOC.74, 2500 (1952). RECEIVEDfor review August 4, 1960. Accepted March 27, 1961.
Titrimetric Determination of Ferrous and Ferric Iron in Silicate Rocks and Minerals CHARLES V. CLEMENCY and ARTHUR F. HAGNER Department o f Geology, University o f Illinois, Urbana, 111. The automatic derivative spectrophotometric titration method of Malmstadt and Roberts ( 5 ) for the determination of iron in titanium and its ores, using coulometric generation of titanous ion as the titrant, has been adapted successfully to the determination of total iron in silicate rocks, magnetite, pyrrhotite, pyrite, or mixtures of these, and to the determination of the ferrous-ferric ratio of silicate rocks, magnetite, or mixtures of these. The method measures the amount of ferric iron present. Using a rapid method of sample dissolution b y a hydrofluoric-sulfuric acid mixture, both total iron and the ferrous-ferric ratio ccn b e obtained within one hour. U. S. Geological Survey rock samples G-1 and W-1 were used as rocks of known iron content. Evidence is presented that reduction of ferric iron during dissolution of the sample may b e a major cause of error in the conventional wet method of analysis for the ferrous-ferric ratio. Results of higher precision and accuracy are obtained in a shorter time using smaller sample sizes than with conventional methods.
iron in small samples of silicate rock, The method described is an adaptation of the automatic spectrophotometric titration method of Malmstadt and Roberts (6) for the determination of iron in titanium sponge, alloys, and ores. Because of the importance of the ferrous-ferric ratio of silicate minerals and rocks as a parameter in the interpretation of likely environments of formation of these rocks, the method
+ 7
BERZELIUS
PLAT F O R M
I
N A STUDY of the Sterling Lake, N. Y., magnetite deposit by Hagner and Collins (3) and Hagner, Collins, and Aye (4),it was necessary to determine the total iron content of 100-mg. samples of metamorphic silicate minerals and rocks. Gravimetric methods similar to those described by Washington (8) proved unsatisfactory for routine analysis of such small samples. The purpose of this investigation was to develop a new method for the analysis of total
888
ANALYTICAL CHEMISTRY
Figure 1 . in beaker A.
Arrangement of electrodes
Platinum foil half cell electrode, Sargent Catalog No. 5-3051 5 6. Platinum gauze electrode, Sargent Catalog NO.S-29672 Gauze screening 1 ’/, Inches high and 1 inches In diameter
was extended to provide a new and better means of obtaining this ratio than was previously available. PROCEDURE FOR DETERMINATION O F TOTAL IRON I N SILICATE ROCKS
Apparatus. A Sargent-Malmstadt automatic titrator and titrator control unit; a Sargent Model I V coulometric current source; and a commercially available Sargent platinum gauze electrode were used (Figure 1). Reagents. Titanium Sponge Solution. Dissolve 40 grams of low-iron content titanium sponge (obtained from Cramet, Inc., Chattanooga, Tenn.) in a mixture of 1200 ml. of 5 M sulfuric acid and 100 ml. of 48% fluoboric acid and dilute to 2 liters with water. A 50-ml. aliquot contains 1 gram of titanium. Standard Iron Solution. 1.00 mg. per ml. Mallinckrodt 99.97% pure iron wire in hydrochloric acid. Leuco Methylene Blue Solution (and other reagents) as described in (6). Blanks. Blank values were obtained by running two standard iron solutions at the beginning, and one a t the end of the day. The standard solutions each contained 5.00 mg. of iron. The number of microequivalents needed t o titrate these solutions was calculated (e.g., 89.5 at 100 ma.) and the standards were run for total iron. The difference between the average number of microequivalents actually used to titrate the standards (e.g., 93.5) and the calculated number was taken as the blank. In this case the blank would be 93.5 - 89.5 = 4.0 peq. This procedure also checks on the efficiency of the electrodes. When the electrodes need regeneration, the number of microequivalents needed to titrate the same standard solution increases. When the number becomes
more than normal, the electrodes should be regenerated. Regeneration of Electrodes. With the solutions used by Malmstadt and Roberts (6) the electrodes remained trouble-free indefinitely, but under present conditions, it was necessary to regenerate them every 2 to 4 weeks depending upon amount of use and probably also upon type of sample. Some elements in the samples may poison the coating on the electrodes, causing a reduction in efficiency. To provide the necessary coating on the electrodes, remove the electrodes from the instrument, soak them in hot chromic acid cleaning solution for 5 minutes, rinse thoroughly in water, heat to redness in a flame, and replace in the instrument. Make the outer gauze electrode the anode by reversing the cell lead plug a t the back of the current source. Electrolyze a solution prepared from 50 ml. of saturated potassium permanganate, 1 mi. of concentrated hydrochloric acid, and 100 ml. of water for 15 minutes. Rinse electrodes thoroughly, make the outer electrode the cathode once again, and regenerate the electrodes, as described in (6), by electrolyzing a regular blank solution containing *a few milliliters of methylene blue for 2 hours. The main cause of poor instrument performance was almost always traceable to the electrodes. Periodic regeneration of electrodes every 2 to 4 weeks or whenever indicated by excessively high blank values should keep difficulties a t a minimum. Regeneration should be the first action taken to remedy any trouble whose origin is not known. Poor end point detection,, premature or late cutoff a t the end point, or continuous shutting off of the instrument before the end point was reached is caused by poor positioning of electrodes within the beaker. Once the electrodes have been properly positioned and perform well, their position should be carefully noted so that they may be returned to the same position after removal. See instruction manuals for further information. Preparation of Instruments for Titration. The instruments were warmed u p for a t least 15 minutes prior to titration. The wave length filter was set on 650 mp, and a yellow cutoff filter was taped over the lens leading to the CdS photocell. The multiplier switch was set to 0.1 (100 ma.). The isolated anode compartment was filled with 0.1N sulfuric acid and kept full during subsequent titrations. Sample Preparation. Only porcelain, mullite, or other iron-free grinding apparatus should be used to grind the samples. Ignore the precipitate of insoluble matter which usually forms while the sample is dissolving in the platinum crucible and which may be mistaken for undissolved sample. This precipitate usually dissolves in later steps. However, when rocks contain large amounts of calcium and magnesium, the titrating
solution may be slightly turbid. This has no effect on the end point or on the amount of iron found, as it would in a colorimetric analysis, so long as the amount of turbidity remains reasonably constant throughout the titration. The ability to use such turbid or dirty solutions is of great benefit in the analysis of silicates which may contain a great many constituents, some of which may be insoluble under the desired conditions. During dissolution most of the silica is volatilized as silicon tetrafluoride. If the sample contains magnetite or pyrrhotite, add a further 1 to 2 ml. of concentrated hydrochloric acid. If pyrite is present, take the hydrofluoricsulfuric mixture to fumes and heat until all the pyrite is dissolved. If large amounts of pyrite are present or if pure pyrite is to be analyzed, the pyrite should be dissolved in 1:1 nitric acid in a beaker, after which the nitrates must be removed by a double fuming with sulfuric acid before proceeding with the titration or additional treatment with hydrofluoric acid. Procedure for Total Iron. A 0.1to 0.2-gram sample of finely ground rock powder (less than 100 mesh) was weighed out directly into a platinum crucible of about 40-ml. capacity. Three milliliters of concentrated hydrofluoric acid and 1 ml. of concentrated sulfuric acid were added to the sample. If magnetite is present, 1 ml. of concentrated hydrochloric acid may be added in addition. Using platinum-tipped tongs, the covered crucible was placed on a hot plate protected by an asbestos gauze and swirled occasionally while heating for 10 minutes just below boiling. The crucible was cooled, the inside washed down with water, and the contents transferred quantitatively to a 200-ml. tall-form (Berzelius) beaker containing 10 ml. of a saturated solution of boric acid and 1 ml. of concentrated hydrochloric acid. The beaker was covered and contents boiled for 2 to 3 minutes to dissolve most of the precipitated material. Fifty milliliters of titanium solution were added and the ferrous ion and titanous ion were oxidized by adding dropwise and with stirring a saturated solution of potassium permanganate until a pink color of excess permanganate persists. The sides of the beaker were washed down and a few tenths of a gram of solid sodium azide were added to destroy excess permanganate. Carefully and with stirring, 25 ml. of concentrated sulfuric acid were added and the sample diluted with water to within 1 inch from the top of the beaker. A glass hook and watch glass were placed on top of the beaker and the solution was boiled for 5 minutes. The beaker was cooled in a pan of cold water to is not about 75" C.-temperature critical. The watch glass and hook were rinsed off, 15 drops (always use a constant amount) of leuco methylene blue indicator solution were added, and the beaker was placed in the titration cell. The beaker platform was raised into
position. With the current ~ource power switch off, the solution was stirred for a few seconds. After stirring, the timer was reset to zero and the autdmatic button was pressed, starting the stirrer, timer, and generation of titanous ion. The instrument shuts off automatically a t the end point when the blue color is reduced to colorless. If the machine shuts off while the solution is still blue, the titration can be continued using the manual control and visual end point detection. The number of microequivalents was read from the timer and the amount of iron present calculated by the equation: 55.85 X
yo Fe
=
X (no. micro-
equivalents - blank) X 100 Wt.of sample in grams
Procedure for Determination of Ferric Iron in Silicate Rocks. A titrating solution was prepared by placing 50 ml. of titanium solution, 10 ml. of a saturated solution of boric acid, and 1 mi. of concentrated hydrochloric acid in a 200-m1. tallform (Berzelius) beaker. The titanium ion was oxidized by adding dropwise and with stirring a saturated solution of potassium permanganate until a pink color persisted. The sides of the beaker were washed down with water and the excess permanganate destroyed by adding a few tenths of a gram of soiid sodium azide. Twentyfive milliliters of concentrated sulfuric acid were added while stirring, and then the solution was diluted to within 1*/2 inches from the top of the b q k e r . A glass hook and watch glass were placed on top of the beaker and the solution was boiled for 5 minutes. The solution was cooled to about 75" C. in a pan of cold water, after which the beaker was removed and wiped dry. At this point the solution was ready for the addition of the sample. While the titrating solution was boiling, a 0.1- to 0.2-gram sample of rock powder was weighed out into a platinum crucible. While the titrating solution was cooling, the leuco methylene blue solution was prepared and all instruments were made ready to titrate. To the sample in the platinum crucible were added 3 ml. of concentrated hydrofluoric acid immediately followed by 1 ml. of concentrated sulfuric acid. The covered crucible was placed on a hot plate covered with an asbestos mat and heated just below boiling for 31/2 minutes. The time was measured from the addition of the sulfuric acid. The crucible was removed from the hot plate, the sides of the crucible were washed down with boiled water and the contents transferred quantitatively into the cooled titrating solution. Fifteen drops of leuco methylene blue solution were added and the solution'was titrated immediately as described in the procedure for total iron. The number of milliequivalents were read and the ltmount of ferric iron present was Calculated. To obtain the ferrous iron content, the ferric content (as metal) was substracted from the total iron VOL. 33, NO. 7, JUNE 1961
889
I
3.co
I
n
z
3
0
0 ”
“ 0
IL“
N
8
LL
cu
2!
8
0.50
4 5
0
IO
15
20
25
30
0.50Y
5
IO
HEATING TIME (rnin.)
I5
20
25
30
HEATING TIME (rnin)
Figure 2. Curves showing variation of amount of iron found Fez03)with heating time
Figure 3. Effect of varying acid concentrations and oxidizing conditions on Fe203 found in 0.2000 grams of
A. 0.2000-gram samples of W-1 dissolved in inert atmosphere 8. 0.2000-gram samples of G-1 dissolved in uncovered crucible
See text for explanation of curves
(70
70
w-1
while exposed to air C. 0.2000-gram samples of W - I dissolved in same manner as those in B Horizontal lines, actual amount of Fez03 present in G-1 and W-1
content (as metal). These figures can then be converted to oxides if desired. If the sample contained magnetite, 1 to 2 ml. of either concentrated hydrochloric or hydrobromic acid was added in addition to the reagrnts mentioned in the procedure. If pure magnetite was to be analyzed, it n as best dissolved in 2 ml. of concentrated hydrobromic acid which is an excellent and rapid solvent for magnetite. This solution was heated for 3 t o 5 minutes and treated as any other sample. (Xote: Hydrobromic acid cannot be used to dissolve magnetite in the total iron procedure because bromine is formed when potassium permanganate is added t o bromide ion and this bromine is not destroyed by sodium azide. Other reducing agents that reduce bromine were tried, but these also reduced ferric ion and caused low resultjs.) Unfortunately, pyrite and pyrrhotite, and presumably other sulfides, interfered with the ferric determination. Evidently sulfur gases released during dissolution caused reduction of ferric ion. Therefore, sulfides must be removed or be absent from the samples. Interferences. Interferences are the same as described in (5), namely, vanadium, molybdenum, and copper, in addition to the presence of sulfides, graphite, and organic material such as is usually found adsorbed on clays. EXPERIMENTAL RESULTS
Preliminary work was done by preparing ten synthetic samples, each containing 5.00 ing. of iron, and analyzing them for total iron. These yielded an average value of 5.00 mg. of iron found with a range of 4.98 to 5.04 mg. The standard deviatiorr was 0.4YG and the relative error 0.4%. After initial work was completed, it mas desired to analyze rocks of known 890
ANALYTICAL CHEMISTRY
iron content as close in mineral composition as possible to the amphibolites and gneisses under investigation. Since no such standards are available, U. S. Geological Survey rock samples G-1 (granite) and W-1 (diabase) were used. These samples have been the subject of exhaustive chemical studies by Fairbairn et al. ( 1 ) and Stevens et al. ( 7 ) . Although they are not classed as standards, they do provide rocks whose composition is fairly well known. Fairbairn (2) lists “recommended values” which arc taken in this paper to be the true values of total iron, ferrous, and ferric oxide content of G-1 and W-1. For G-1 these values are: 0.93y0 ferric oxide, 0.99% ferrous oxide, and 2.03% total iron (calculated as ferric oxide). For W-1 these values are: 1.50% ferric oxide, 8.71% ferrous oxide, and 11.17% total iron (as ferric oxide). Stevens’ (7) average values for iron oxides are, in general, slightly lower than those of Fairbairn, and may be slightly more accurate because of additional analyses incorporated into the averages. In general. results of analyses of G-1 and W-1 made during this investigation yielded slightly lower figures than FairLairn’s “recominended values,” tending to agree more closely with Steyens’ values. Total Iron Analyses. Five 0.2000gram samples of G-1 were analyzed for total iron content by the procedure described above. The average total iron found was 1.94yo (as ferric oxide) with a range of 1.89 t o 2.00%. The standard deviation was 0.05% and the relative error was 4%. Seven 0.2000gram samples of W-1 were run for total iron and yielded an average value of I l . O S ~ O(as ferric oxide) n-ith a range
of 10.99 to 11.20’%, The standard deviation was 0.07% and the relative error was 1%. Five 0.1000-gram samples of W-1 were run for total iron to compare the effect of sample size on error. These yielded an average value of 10.93% total iron nith a range of 10.80 to l l . O i 7 c . The standard deviation was 0.12% and the relative error was 2%. Five 30-mg. samples of pure magnetite were run for total iron. These analyses yielded values of 68.7, 71.0, 70.6, 70.0, and 69.67, total iron. Two samples containing known weights of magnetite plus IT-1 were prepared and run for total iron. The samples contained 26.2 and 24.2 mg. of total iron, respectively. The amounts recovered mere 26.0 and 24.6 mg. This shoved that total iron can be determined satisfactorily on pure magnetite and on magnetitesilicate mixtures. Additional 15 ork showed that total iron can also be dctermined on pyrrhotite, pyrite, and mixtures of these with silicates by slight changes in the arjd treatment as described in the procedure. Ferric Iron Analyses. -4series of eight samples each of G-1 and IT-1 were run for ferric oxide content using different heating times betn-een 2 and 4 minutes. These results are summarized in Tables I and 11. Each sample within a series yielded essentially the same results. Additional data for other heating times are represented by the points forming curves B and C of Figure 2. All samples n-ere dissolved as described in the procedure above. The ferrous oxide content is obtained by difference. The amounts of ferric oxide found in G-1 and IT-1 are in good agreement with the recommended values of Fairbairn ( 2 ) and of Stevens ( 7 ) . Pure magnetite was analyzed for ferric oxide content, after which mixtures of magnetite and W - 1 were pre-
pa’red having a known ferric oxide content. Five 50-mg. samples of purc magnetite yielded values of 47.7, 47.2, 47.0, 47.5, and 48.3% ferric oxide. ‘ho samples of magnetite plus W-1 containing 10.2 and 10.9 mg. of ferric oxide yieldcd values of 9.5 and 10.9 mg., respectivelp. This shor\s that the ferric content of mineralic mixtures can be determined and that the amount found will be equal to the sum of that present in each mineral constituent. Care should be taken not to use as test material magnetite with any substantial amount of titanium present. Study of Oxidation and Reduction During Sample Dissolution. A major problem in the determination of ferrous and ferric iron is control of t h e oxidation state during sample dissolution. In t h e conventional method of analysis of silicates for ferrous oxide content as described, for example, by Washington ( 8 ) , precautions are usually taken to prevent oxidation by the atmosphere. As Riley and JTilliams (6) state: “It is generally agreed that the principal source of error in these methods is aerial oxidation of the ferrous iron during the decomposition process, or to a lesser extent, during the titration.” It was thought that aerial oxidation of ferrous iron would be the primary source of error in the analysis of the ferrous-ferric ratio; therefore, the rock samples were dissolved in an inert atmosphere. A small Plexiglas dry box was built with a self-contained heating unit. The heating unit was a porcelain crucible wound with 4 feet of No. 30 Chrome1 wire (6.44 ohms per foot) and imbedded in alumina cement; the temperature was controlled by Variac. The Variac settings were calibrated by placing a few milliliters of water and a thermometer in the platinum crucible. Samples were dissolved in a platinum crucible set in the well of the heating unit. The box &-aspurged with nitrogen gas follon ed by sulfur hexafluoride gas. Sulfur hexafluoride gas (obtainable from The Matheson Co., Joliet, Ill.) is inert, nontoxic, almost insoluble, unreactive with w-ater solutions, and about five times heavier than air. Temperature n-as maintained a t about 75” to 85” C. Samples 11ere dissolved by adding 3 ml. of concentrated hydrofluoric acid and 1 ml. of concentrated sulfuric acid through a hole in the cover and then heating. After dissolution, the samples mere poured into a previously prepared titanium solution and the ferric iron was titrated. Results n-ere erratic, but always low. It was found that the low results mere not due to incomplete solution of the sample as was first suspected, but to reduction of ferric iron during dissolution, perhaps by other elements in the rock. It was presumed
that reduction v a s the cause of thp low Table I. Results of Analyses of G-1 results rather than pi ecipit:ition, comfor Ferric Oxide Content plexation, or some other effect \\hhich could also decrease the ferric iron conAll samples tceighed 0.2000 gram tent. That the low results were not due Heatto incomplete sample dissolution was 1ng Error Sam- Time Fezor, % verified by running a total iron analysis ple (Min.) Present Found (%) after a certain dissolution time. If 1 2 0 93 0 90 -0 03 essentially all the iron was iecovered, it 0 85 -0 08 2 2 could be assumed that the sample was 0 84 -0 09 3 3 completely dissolved within that time. 4 3 0 92 -0 01 After 2 minutes 93y0 of the total iron +O 11 5 3 104 6 3 0 92 -0 01 was recovered and after 3 minutes 97% 7 4 0 80 -0 13 of the total iron was recovered. 0 93 000 8 4 Curve A of Figure 2 summarizes how AV. 0 90 0 06 reduction takes place mhen samples of Std. dev. 0 07 W-1 were dissolved in an inert atmosRel. error 6% phere and how the amount of reduction varied with time. Little oxidation occurred after standing 30 minutes in the Table II. Results of Analyses of w-1 box. for Ferric Oxide Content Because reduction was apparently A411 samples weighed 0.2000 gram taking place rather than oxidation, the inert atmosphere was considered unHeatnecessary and some samples of G-1 and ’3g Sam- Time Fe203r %Error W-1 wore dissolved while exposed to ple (Min.) Present Found (yo) air. Curves B and C of Figure 2 show 1 2 1.50 1.63 +0.13 that between 0 and 2 minutes the sample 2 2 1.40 -0.10 mas in various stages of dissolution, as 1.57 $0.07 3 2 indicated by increasing ferric content 1.39 -0.11 4 3 1.40 -0.10 5 3 with time. B e h e e n 2 and 4 minutes 1.53 +0.03 6 3 the amount of ferric iron found was 1.16 -0.34 7 4 approximately the same as the recom8 4 1.63 +0.13 mended value and analyses were fairly Av . 1.46 0.13 reproducible. After 4 to 5 minutes a Std. dev. 0.16 rather sharp decrease in ferric content Rel. error 9% indicated that reduction m s occurring. From this stage on, reproducibility became poor. K i t h further time, reducTable 111. Effect of Using Covered and tion reached a minimum a t which point Uncovered Crucibles on Amount of air oxidation began t o predominate. F e z 0 3Found in 0.2000 Gram of W-1 Kot all the ferric iron was reduced. The Samples were dissolved by the recomferric iron content then increased ~ i t h mended procedure time until the recommended value was Heating exceeded. Reproducibility was very Fez% % Time poor a t this stage but the trend was Present Found (Min.) Crucible clear. 1.50 1.60 3.5 Uncovered When another series of samples was 1.47 3.5 Covered run under the same conditions, with the 1.04 6.0 Uncovered 0.46 6.0 Covered crucible covered, slightly better reproducibility was obtained during the 3- to 4-minute heating time. The increase in ferric content with time because of air oxidation was sharply curferric ion \yere recovered after heating tailed. Table I11 shons different rethan in the original solution, showing sults obtained for t n o different heating that no reduction occurred during the times using covered and uneorered treatment of a solution. The fact that crucibles. the ferric iron is not completely reOccurrence of this unexpected reduced indicates the presence of a limited duction, and subsequent oxidation by amount of reducing agent, such as the air, may help to explain the nide another component in the rock. variations reported in ( 1 ) for the ferrousThe results of analyses of G-1 and ferric content of G-1 and K-1. AlIT-1 for percentage of Fe20, were in though the reducing agent is unknown, good agreement with the recommended the authors suspect some component of values when the exact procedure dethe rock to be responsible, such as pyscribed above w’as followed. Further rite which releases sulfur dioxide during investigations showed that the results of dissolution, although titanium (111) is the analyses were sensitire to changes also a possibility. Khen a pure iron in conditions. (Curves A through D solution of known ferrous-ferric content of Figure 3). These curves were prewas treated in the manner of a rock pared by analyzing 0.2000-gram samples sample, only slightly higher amounts of of R - 1 using essentially the same pro~~
VOL. 33, NO, 7 , JUNE 1961
* 891
cedure as described above, but changing one of the conditions. When Such a series of samples was run substituting 2 ml. of 2:1 sulfuric acid for the 1ml. of concentrated sulfuric and heating in an uncovered crucible, curve A , Figure 3 was obtained. The amount of Fez03 found was always high because of aerial oxidation. Using this 2:l acid but heating in a covered crucible, curve B, Figure 3 was obtained. The masking effect of aerial oxidation has been decreased and the reactions can be followed in more detail. From 0 to 6 minutes the sample d i e solves a t a much slower rate than with concentrated acid. After 6 minutes reduction begins, before the sample is completely dissolved. After 10 minutes reduction slows, and oxidation begins to occur fairly rapidly partly because a large part of the hydrofluoric acid has been evaporated, allowing air to replace the vapors in the crucible. With the same acid mixture used for curves A and B of Figure 3, and in addition adding 1 ml. of concentrated hydrochloric acid, curves C and D resulted. Curve C was obtained with an uncovered crucible and curve D with a covered one. When the crucible was uncovered, oxidation predominated and high and erratic results were obtained. The rate of oxidation appears to have been greater in the presence of hydrochloric acid than in its absence. When a covered crucible was used the sample dissolved more rapidly in the presence of hydrochloric acid and reduction began earlier, followed by oxidation after about 10 minutes. The presence of hydrochloric or hydrobromic acid did not seem to interfere when concentrated sulfuric acid was used. The foregoing emphasizes the effect of conditions on the amount of ferric oxide found and the need to maintain constant and uniform conditions if true and reproducible results are to be obtained. Apparently reproducible results can be obtained as long as conditions are held constant for each sample. However, these results may be far from the true content. The amount of ferric iron found depends on the competing rates of a t least three major reactions which are taking place simultaneously. These are: rate of dissolution of the sample, which controls the amount of ferric and ferrous
892
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ANALYTICAL CHEMISTRY
iron in solution and the amount of reductant available, which itself is presumed to be a component of the sample; rate of oxidation of ferrous iron by aerial oxidation and perhaps also by components of the samples; rate of reduction of ferric iron by the unknown reducing agent. Each of these reactions proceeds a t its own rate, which is dependent on many factors. The curves of Figures 2 and 3 illustrate the results of these three reactions for the specific set of conditions described. Any factor affecting any of these reaction rates will affect the shape of the curve, and an apparently small change in conditions may have a very large effcct on the resultant curve. CONCLUSION
Conventional methods of analysis for ferrous and ferric oxide content of silicate rocks may not yield true valucs under some conditions. The common belief that aerial oxidation is the major source of error may not be so in all cases. Evidence is presented that reduction of ferric iron by an unknown reducing agent occurs to some extent when silicate rocks such as G-1 and W-1 are dissolved. Such reduction has no doubt been observed before, but a low result with the conventional method invariably is attributed to incomplete dissolution of the sample. Further investigation into the nature of the unknown reducing agent, with a complete study of the oxidation-reduction behavior of ferrous and ferric ions in various solutions would be desirable. For example, indications are that a pure solution of ferrous sulfate a t room temperature is not very much oxidized by bubbling air through it for several hours, but if other ions are present, the oxidation rate increases. No effort was made to find out how small a sample size could be used or to improve the precision and accuracy from that obtained with the sample sizes and iron contents used. Undoubtedly with further experimentation these could be improved. For example, the Model IV current source permits selection of currents of 200, 100, 50, 10, and 5 ma., of which only the 100 setting was used. By analyzing for ferric rather than ferrous iron as conventional methods
do, this method provides a check on the conventional methods in the determination of ferrous-ferric ratios. The authors believe that the usual methods for analysis of silicate rocks are in need of improvement with respect to speed, accuracy, and simplicity. Although the difficulties inherent in a complete and unified scheme of analysis for a complicated system such as silicate rocks are fully appreciated, advantage should be taken of new methods, techniques, and instruments in attempts to improve rock analysis. Methods utilizing the instrument described in (5) have been published for the analysis of calcium, magnesium, potassium, and titanium in various materials, and these may possibly be adapted to silicate analysis. At present, the methods and instruments described are especially useful when one or two specific elements are sought. However, accumulation of a whole series of such new methods may eventually lead to a new unified scheme of analysis for silicate rocks. ACKNOWLEDGMENT
The authors thank H. V. Malmstadt of the Department of Chemistry, University of Illinois, for the use of instruments and laboratory facilities and for reading the manuscript. They also thank D. M. Henderson for his interest and many valuable suggestions during this study and in the writing of the manuscript. LITERATURE CITED
(1) Fairbairn, H. W., et al. U. S . Geol. Survey Bull. No. 980 (1951). (2) Fairbairn, H. W., Geochim. et Cosmochim. Acta 4,143 (1953). (3) Hamer. A. F.. C ~ l l h L. ~ .G.. Science 122,3182, 1230 i1955). ’ (4) Hagner, A. F., Collins, L. G., Aye, T., Mzning Ens. 10, No. 12, 1246 (1958) \ - I
(abstract).
( 5 ) Malmstadt, H. V., Roberta, C. B., ANAL.CHEM.28,1412 (1956). (6) Riley, J. P., Williams, H. P., Mikrochzm. Acta ( V i m )516 (1959). (7) Stevens, R. E., et al., U. S . Geol. Survey Bull. No. 1113, (1960). (8) Washington, H. “The Chemical Analvsis of Rocks, 4th ed., Wiley, New-York, 1930.
si,
RECEIVED for review December 12, 1960. Accepted March 29, 1961.
