INDUSTRIAL A N D ENGINEERING CHEMISTRY
86
Vol. 17, No. 1
Determination of Ferrous Iron i n Materials Containing Metallic Iron’ By C. E. Sims and B. M. Larsen NORTHWEST
EXPERIMENT STATION,
u.$. BUREAUOR MINES,I N
COBPERATION WITH COLLEGE OF
MINES,UNIVERSITY
OF WASHINGTON,
SEATTLE, WASH.
I
TABLEI
RON may be present in materials in three states of oxida-
Excess ferrous iron indicated
Metallic iron by titration above true tion-as ferric iron (Fez03), ferrous iron (FeO), and Ferrous weighed into Total titra- amount of ferrous iron + metallic iron. I n silicates, oxide ores of iron, slags, etc., iron hematite tion metallic iron present , Per cent Per cent no metallic iron is ordinarily present, all the iron being oxi- Sample Per cent 1 3.75 dized, either partly as ferrous iron, or completely as ferric iron. 2 3.75 3.80 In such materials the ferrous iron present can be easily de23.6 39.0 11.6 20 8 31.3 0.7 termined directly, ferric iron being then calculated by sub32.7 7.9 21.0 tracting the per cent of ferrous iron from the per cent of total iron in the sample. Evidently, the amount of ferric iron reduced to the The determination of ferA method has been developed for the accurate determiferrous condition in the acid rous iron in such cases is nation of ferrous iron in substances containing both solution is roughly proporsimple. The iron of the metallic and ferric iron. The method is important in all tional to the amount of sample is ordinarily disstudies of iron ore reduction. metallic iron present-that solved in hydrochloric acid No previous methods tested by the authors were accurate is, to the amount of hydrounder conditions such that in the presence of metallic iron, essentially because of the gen evolved during the soluair ‘is practically excluded reducing action of the latter during solution of the tion of the sample. from contact with the hot sample. solution and no oxidation The obvious course was In the method given, the metallic iron is removed by takes place. The solution to eliminate the metallic copper sulfate solution, the precipitated copper is then iron before treatment with is then cooled rapidly and removed by an alkali cyanide solution, and the residue, hydrochloric acid. I n an the ferrous iron titrated diconsisting of ferrous and ferric oxides and gangue, is attempt to do this, severaI rectly in the cold solution by dissolved in hydrochloric acid out of contact with air and samples of partially repotassium permanganate or the solution titrated directly for ferrous iron. The metalduced hematite were first dichromate solutions. lic iron may be determined in the filtrate from the treatboiled with a neutral copBlaira gives a somewhat inment with copper sulfate solution, total iron run on a per sulfate solution, as in volved procedure for this separate sample, and the ferric iron then found by differthe first stage of the metalmethod. For ferrous silience, The method is not accurate in the presence of lic iron determination by cates insoluble in hydrolarge amounts of ferrous sulfide. Williams and Anderson.4 chloric acid, the decomposiAfter boiling 20 minutes, tion may be carried out in platinum with the addition of hydrofluoric acid, as in the pro- the samples were filtered hot on a t e r paper and the residue of gangue, iron oxides, copper, and filter paper was placed cedure given by C r o o k e ~ . ~ in a covered beaker with distilled water and hydrochloric Experimental acid, boiled to complete solution of the oxides of iron, cooled, I n the course of extended work on the reduction of iron filtered, and titrated with potassium permanganate. The reores a t the Northwest Experiment Station, it has repeatedly sults on two analyses are shown in Table I1 (metallic iron been found desirable to determine the amount of ferrous was determined in the first filtrates). iron in samples which also contain metallic iron and ferric iron. A simple method of procedure would seem to be that TABLEI1 Ferrous iron Ferrous plus of first determining the metallic iron, and then the metallic Total iron Metallic iron indicated metallic iron iron plus ferrous iron by solution of another sample in hySamule Per cent Per cent Per cent Per cent A73.2 9.2 64.4 73.6 drochloric acid and direct titration. This gives high results 9.4 62.5 71.9 for ferrous iron owing to action indicated in the discussion B 76.4 30.3 56.3 86.6 30.5 53.0 85.5 of the hydrogen evolution method for metallic iron given in the paper by Williams and A n d e r ~ o n . ~Apparently, the “nascent” hydrogen given off by the dissolving metallic I n three out of four analyses the sum of ferrous iron indicated iron reduces some of the ferric iron in the solution. To by the titration plus the metallic iron was larger than the check these results a sample of pyrite cinder containing 54.1 total iron present. Apparently, the hydrochloric acid had per cent total iron, 1.7 per cent sulfur, and no metallic iron, acted upon the filter paper, taking organic matter into soluwas analyzed for ferrous iron by direct solution in hydrochloric tion, which reacted with the permanganate. Similar samples were then treated in the same way, but acid. Half-gram portions of this same hematite and varying amounts of pure iron wire were then treated in exactly the filtered in a Gooch crucible on asbestos after the copper sulfate treatment. In this case no organic matter was present in the same way. Results are given in Table I. solution to react with the permanganate, but still the ferrous 1 Received July 17, 1924. Published by permission of the Director, iron seemed too high. The results check fairly well, but U.S.Bureau of Mines. 3 “The Chemical Analysis of Iron,” 4th ed., 1901, p. 221. J. B. Lippinindicate almost no ferric iron present in these samples, as the cott c o . ferrous iron plus metallic iron almost checks with the total * “Select Methods in Chemical Analysis,” 4th ed., 1906, p. 611. Long- iron present. Further investigation disclosed the fact that mans, Green & Co. the metallic copper during solution in dilute hydrochloric 4 THISJOURNAL, 14, 1057 (1922).
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I N D U S T R I A L A N D ENGINEERING CHEMISTRY
January, 1925
acid acted as an even more efficient reducing agent for the ferric iron than the metallic iron it replaced, the copper being dissolved by ferric chloride according to the reaction 2 Feels
+ Cu = 2 FeClz + CuCla
It was therefore necessary to remove this copper before the acid treatment. Experiments showed that solutions of alkali cyanides would dissolve this finely divided copper fairly rapidly and quite completely, probably according t o the reactions 2 CU 4 KCN 2 HzO = 2 KCu(CN)a 2KOH Hz The procedure given below was then worked out on this basis. I n essentials, it consists simply in removing the metallic iron from the sample of copper sulfate solution and treating the residue with an excess of alkali cyanide solution, removing the precipitated copper, and leaving a residue containing only the oxides and gangue of the sample plus some asbestos from the filtrations. This residue can be treated by solution in hydrochloric acid and titrated directly for ferrous iron. Terres and Pongraczs analyzed partly oxidized iron samples for metallic, ferrous, and ferric iron. They removed the metallic iron with mercuric chloride solution by the method of Wilner and Merck,’ according to the reaction
+
+
Fe
+
+
+ HgClz = Hg + FeClz
The residue was dissolved in acid out of contact with the air and the solution analyzed for ferrous and ferric iron. Besides being rather $low and difficult, it seems doubtful that this method could give accurate results, for experiments have shown that ferric chloride dissolves the precipitated mercury, being itself reduced to ferrous iron, just as in the reaction with copper given above. A method could probably be worked out, removing the precipitated mercury before acid treatment. Because of the slowness and more difficult manipulation involved in the mercuric chloride method14 however, the method outlined below was considered the best solution of the problem. Procedure Weigh 0.5-gram samples of minus SO-mesh material into dry 150-cc. beakers. Add 15 cc. of a neutral 10 per cent copper sulfate solution, dilute to about 60 cc. with water, cover with watch glasses, and boil gently on hot plate for 15 to 25 minutes, maintaining the original volume by frequent additions of water. Remove and filter through Gooch crucibles on purified asbestos, washing four or five times with hot water. Place filtrates in 250-cc. beakers and run for metallic iron,4 if desired. Transfer residues to original 150-cc. beakers, washing out crucibles with water, add 10 to 15 cc. of a saturated, slightly ammoniacal solution of sodium cyanide, and dilute t o about 50 cc. Place covered beakers on a hot plate kept just w a r n (not boiling), and let stand overnight. Filter again on Gooch crucibles, washing thoroughly (8 to 10 times) to remove cyanides, and wash residues into 150-cc. Erlenmeyer flasks with about 20 to 25 cc. of distilled water. Cover flasks with watch glasses and heat to gentle boiling t o expel most of the air and displace it by water vapor. Add 25 cc. of hydrochloric acid and keep the covered flasks boiling gently until solution of the iron oxides is complete. Cool the flasks in running water. For titration, either the usual dichromate method should be used, employing outside indicator, or preferably the modification given by Knop,8 as indicated below, using diphenylamine as an inside indicator. The cold solution may be either filtered (removes asbestos and gives a clear solution for titration) or washed out di6
Prescott and Johnson, “Qualitative Chemical Analysis,” 7th ed., 1918, D.Van Nostrand Co. 2. Elektrochem., 26 386 (1919). Merck, 2.anal. Chem., 41, 710 (1902). J . A m . Chem. SOG.,46,263(1924).
p. 272. 0
’
.
87
rectly into 400-cc. beakers containing about 75 cc. water and 15 cc. titrating solution. Dilute to total volume of about 200 to 250 cc. Add 3 to 6 drops of diphenylamine indicator and titrate with a standard potassium dichromate solution (1 cc. equals 0.005 gram iron) until the green chrome solution just changes to a deep blue color that is unchanged by the addition of a few more drops of dichromate. Subtract 0.05 cc. from titration for the oxidation of each three drops of indicator used. Cc. of K2Cr20, = Per cent ferrous iron
Ferric iron present may be calculated by subtracting the sum of the metallic and ferrous iron from the total iron of the sample. Solutions Copper sulfate: 100 grams anhydrous copper sulfate per liter of water. Saturated solution of sodium cyanide. Standard potassium dichromate solution: 4.385 grams potassium dichromate per liter. Standardize against pure iron wire. 1 cc. equals 0.005 gram of iron. Titrating solution: 150 cc. of phosphoric acid plus 150 cc. of sulfuric acid, diluted with water to 1000 cc. Indicator solution: 1 gram of diphenylamine dissolved in 100 cc. of concentrated sulfuric acid. The solution becomes brown after some time, without losing its usefulness, however.
Accuracy of Method
It is nearly impossible to get an absolute check upon this method of analysis for ferrous iron, because neither pure ferrous oxide, pure reduced metallic iron, nor known mixtures of ferric oxide, ferrous oxide, and metallic iron can be obtained. The results in Table 111, however, give some confirmation of the accuracy of the method. A sample of high-grade hematite was analyzed for total iron. Four more portions of this sample were treated with copper sulfate and cyanide solution as in the procedure given above for ferrous iron. Two of these were titrated directly for ferrous iron and two reduced with stannous chloride and titrated for total iron. TABLE I11
Hematite sample
1 2 3
Direct total iron Per cent
Direct ferrous iron Per cent
Ferrous iron after preliminary treatment Per cent
Total iron given preliminary treatment as for ferrous iron but reduced with stannous chloride bebefore titration Per cent
68.4 63.4 0.25 0.3
4 5
0.4 0.3
6 7 8
63.4 63.2
The iron oxides were not appreciably affected by the preliminary treatment. The small amount of ferrous iron was probably present as sulfide of iron (sulfur content, 0.15 per cent). A sample of a nearly pure magnetite ore was then treated in a similar way, as indicated in Table IV. TABLE IV Ferrous Total iron iron by by direct direct solution solution Per cent Sample Per cent 1 66.4 2 66.4 3 21.5
4 5 6 7 8 9 10 11
21.5 21.4
Metallic iron Per cent
Ferrous iron by new method Per cent
Total iron remaining after treatment for removal of metallic iron and pptd. copper Per cent
0.25
0.2
21.2 21.1 21.1
65.9 65.8 66.0
The small amount of metallic iron present was introduced during the grinding and preparation of the magnetite sample.
I N D UXTRIAL A N D ENGINEERING CHElMIXTRY
88
Results in this case give a composition very close to the theoretical for magnetite (Fe304, or FeO Fez03). I n general, the results also indicate that the preliminary treatment with copper sulfate and cyanide solutions to remove metallic iron and the precipitated copper has no appreciable effect upon the iron present in the sample as either ferric or ferrous oxide. Consistent and reasonable results have been obtained with the method in nearly all analyses, some of the results of analysis of samples of hematite ore in varying stages of reduction by carbon monoxide gas being given in Table v. I n amounts above about 1 per cent, the presence of ferrous sulfide causes high results for both metallic and ferrous iron. Copper sulfate solution tends to dissolve the ferrous sulfide slowly, so that part of the sulfide iron is titrated as metallic iron. This dissolving action of the copper sulfate is not complete, and part of the ferrous sulfide remains to be dissolved by the acid. In this solution hydrogen sulfide is
Vol. 17, No. 1
evolved, which reduces part of the ferric iron in solution. I n samples containing around 1.0 per cent sulfur or less, however, nearly all the sulfide iron will analyze as metallic iron and the ferrous iron will not be appreciably in error. . TABLEV Sample 1
Total iron Per cent 70.4
Metallic iron Per cent 1.4
Ferrous iron Per cent 28.0 27.9 35.3 3.5.2 37.3 37.5 9.9
2
72.1
9.7
3
72.6
10.0
4
75.1
33.7
5
66.2
5.9
12.7 12.9
6
so.0
39.8
8.0
9.8
The Adsorption of Vapors b y Alumina' By I,. A. Munrb and F. M. G. Johnson MCGILLUNIVERSITY, MONTREAL, CANADA
Air was aspirated a t a The adsorption of vapors by alumina has been investiconstant rate through a purson2that alumina, pregated by a dynamic method. ifying train, then through pared by heating the The adsorptive power of the alumina depends on its state the liquid with which it was hydroxide in an open vessel of hydration. The curves show that the adsorption may to be saturated, and finally over a smoky flame, was a be divided into three classes, one of which denotes catathrough a tube containing very efficient drying agent. lytic decomposition of the vapor. the adsorbent, all a t room Dover and Marden,3as well Alumina is a very good adsorbent for all the vapors intemperature (20' C.). The as Marden and E l l i ~ t t , ~ vestigated. It can be used for the recovery of gasoline, course of the adsorpti'on was have investigated the dryether, benzene, the alcohols, sulfur dioxide, and ammonia followed by weighing the ing power of alumina pregas. It is not, however, suitable for the recovery of such a l u m i n a and saturating pared in this way. They vapors as acetone, ethyl acetate, amyl acetate, various liquid a t intervals. The found that the aluminawas a alkyl halides, and acetyl chloride and bromide. gain in weight of the alumore efficient drying agent mina and the loss in weight than calcium chloride, quicklime, or sulfuric acid. Phosphorus pentoxide was the of the liquid are expressed as cubic centimeters of vapor only desiccant of greater efficiency. Fisher, Faust, and a t standard temperature and pressure. Curves are obWaldenb used alumina in organic combustions, and ob- tained by plotting the volume of vapor supplied per gram tained the theoretical amount of hydrogen within 0.33 per of adsorbent against the volume of vapor adsorbed per cent. McIntosh, working in this laboratory, found on at- gram of adsorbent. These curves give a comparison of tempting to dry ammonia gas by alumina that the gas was the relative number of molecules of the different vapors adsorbed by it. I t therefore seemed probable that alumina adsorbed from air-vapor mixtures saturated at room temperawould be an adsorbent for other gases. Silica gel has already ture (20"C.). They also show the efficiency of the adsorbent been found to be a good adsorbent.6 The similarity of for each vapor under these conditions. These curves also alumina and silica gels suggested that the former would also give information as to the nature of the adsorption. be efficient in this respect. Preparation of Alumina This paper describes briefly the results obtained in a preI n the first experiments the alumina was prepared by heatliminary investigation on the adsorption of vapors by alumina, ing the pure hydroxide in an open vessel over a smoky flame. about thirty vapors having been studied. This method was soon discarded for a more definite one. General Procedure A U tube was filled with a known weight of the hydroxide in In an industrial application of alumina as an adsorbent, it the form of a fine powder. Small plugs of asbestos were used is probable that the process would involve adsorption from a to retain the alumina. The tube was placed in a bath of current of vapor-air mixture. For this reason a dynamic potassium and sodium nitrates a t 400" C. for 2 hours, while a t the same time dry air was drawn through the alumina. The method was chosen. resulting loss in weight gives the amount of water driven 1 Received May 28, 1924. off-i.e., the extent of dehydration. This is expressed as 2 J . A%. Chem. doc., 34,911 (1912). per cent of the initial weight. (Column 3, Table I) a I b i d . , 39, 1609 (1917). 4 THISJOURNAL, 7, 320 (1915). The total water content of the hydroxide was determined 8 I b i d . , 14, 1138 (1922). by blasting for several hours in a platinum crucible a t 8 McGavic and Patrick, J . A m . Chem. SOL, 42, 946 (1920); Miller, the highest temperature of the blast lamp. According to Chem. Mel. Eng., 28, 1155 (1920).
I
T WAS found by John-
