Preparation of Reagent Grade A. N. Mulokozi University of Dor es Soloom Tanzania
I~oN(W) Oxide from Iron Ore An experiment for freshman chemistry
T h e task of making laboratory experiments more meaningful and interesting t o studen& i s a t present accepted a s a serious challenge t o chemistry educators all over t h e world. T h e preparation of reagent grade Fez03 is a n experiment hased on the research approach. T h e advantages of this experiment are many. 1) The materials and equipment are easily obtained. Suitable mineral samples may be obtained through mining and gealogical institutions in the country, or through commercial firms dealing with analytical standards such as Bureau of Analysed Samples Ltd. Newham Hall, Newhy, Middleshrough, Teesside, England. 2) For the last two years we have been offering at the University of Dar es Salaam a beginners' lecture course bearing the title: "Minerals as a Source of Inorganic Industrial Materials and Chemicals." The etbusiasm shown by our students in this lecture was far beyond our expectations. This encouraged us to incorporate the preparation of pure compound from minerals in the chemistry laboratory course. 3) Finally an experiment involving the preparation of pure chemical compounds from minerals will find appreciation in developing countries such as Tanzania where orders of chemicals from overseas suooliers . take sometimes 6 man to arrive. Much delay of important work may be avoided when the desired chemicals can be prepared in the laboratory from minerals. ~
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Types of Suitable Ores and Their Description
Oxide Ores: Those include hematite FezO3, magnetite Fe301, and ilmenite FeTiOs. T h e first two mentioned are the most suitable starting materials for preparation of iron compounds. Titaniferrous iron ore is another convenient . is mainlv comnosed of ilmensource of iron c o m ~ o u n d s It ite a n d magnetite. Sulfide Minerals: T h e important iron mineral in this groud is iron pyrites FeS2. copper pyrites with the ideal composition corresponding to the formula CuFeSz is less suitable for preparation of iron compounds. I n all the minerals mentioned here, there are often varying amounts of sands and other siliceous matter known a s gangue.
tion, 20 ml of concentrated hydrochloric acid is added. The mixture is heated to boiling. Stannous chloride solution is added carefully to the solution until it is completely deeolorized. Stannous chloride should only be in slight excess. The solution is prepared by dissolving 3 g SnC12.2Hz0 in a boiling mixture of 25 ml concentrated hydrochloric acid and 75 ml distilled water. The solution for titratian is cooled by placing the flask under a stream of cold water from the tap. A saturated solution of mercuric chloride (10 ml) is added. A slight white precipitate should form if enough stannous chloride solution was used. If the precipitate formed is grey or hlaek too much stannous chloride was added, and the sample must be discarded. The solution is now diluted with 100 ml distilled water followed by an addition of concentrated sulfuric acid and phosphoric acid (5 ml of each). The indicator (0.3 ml of aqueous diphenylamine sodium sulfonate) is added, and the solution is titrated rapidly with 0.1 N K2Cr20r until the color changes to permanent violet blue (1 ml 0.1 N K2Cr207 = 5.585 mg Fe). The Recovery of lron(lll)Oxide from lron Ore Metals show marked variation in their complex-forming behavior. Based on this property are some bf the most convenient methods of separation of metal ions in solution. Those methods are listed below, 1) The interfering ions may be kept in solution through complexing while the desired metal ion is precipitated with a suitable reagent. Iron may be precipitated from the solution of copper pyrites with ammonia solution in which copper remains soluble as tetrammine complex cation Cu(NH3)2+. 2) The desired metal ion may he removed from solution by camplexing with appropriate reagents followed by extraction with oxanrt solvents ~sulventextracllon). In hydrochloric acid solurmns won forms a chloro complex kC1.- which r a n be easily exrrncted with dierhyl ether or hetter still wrth methyl rsubutyl ke. tone. 3) The desired ion may he absorbed on a suitable inn-exchange resin in column chromatographic separation. In the present experiment, the separation of iron is effected by absorption of the chloro complex anion FeClr- on a strongly basic anion-exchange resin. L
Analysis of the Ore
The Ion-Exchange Separalion
Before taking the trouble to prepare chemical compounds from minerals, it is important first to establish their suitability as the source of the compounds desired. For this purpose chemical analysis of the minerals has to he carried out. A rapid procedure far the determination of iron in iron ore is now described. Alternative methods are described elsewhere (1-3). A finely ground sample (0.7-1.0 g) of the ore is decomposed with a boiling mixture of equal volumes (30 ml each) of concentrated perchloric acid and phosphoric acid for 45 min with constant shaking of the flask. If hy this time there are still some hlaek particles in the residue, heating must be continued. The decomposition is considered completed when the sample dissolves entirely, or when only greyish white residue is left hehind. Hot perchlarie acid reacts violently with carbon and organic eompounds. Care must be taken when working with this acid. The procedure described so far is suitable for decomposition of oxide ores. Sulfide minerals are best dissolved by heating the sample in a mixture of hydrochloric acid and nitric acid. The solution on cooling is transfered into a 250-ml volumetric flask. Distilled water is added to fill the volume up to the mark. From this solution aliquots of 50 ml are taken for titration. To 50 ml of the solu-
T h e ion-exchange procedure employed here is based on the interaction between the hexaaquoiron(II1) cation complex and chloride ions i n acidic medium. A t high proton concentration t h e protonation of the coordinated water molecules occurs. This has two effects. First the water molecules lose their coordinating ability since in each protonated water molecule a n electron pair is donated t o the proton. Second on protonation the water molecule aquires a positive charge so t h a t it is repelled hy the positively charged metal ion. T h e protonation of water a t high acid concentration renders it ineffective a s a coordinating ligand permitting thereby easy access of the chloride ion t o the metal. It is now easily understood why chloro com-
634 / Journal of Chemical Education
The experiment is hased in part on a paper presented at the conference on teaching of University chemistry held at Nairobi, Kenya, December 14-18.1971.
plexes of iron(m) are only formed in strong hydrochloric acid solutions. The formation of the chloro complexes follow the pattern indicated Fe(HzO)ss+ + 4 HCI
-
Fe(HzO)&lr-
+ 4 HzOf
(1)
Both chloro complexes of iron shown in eqns. (1) and (2) are known to exist in hydrochloric acid solutions of iron(111) compounds (4). The reactions above make a suitable subject for discussion among students who already have had lectures on coordination chemistry. Some of the factors affecting the stability of metal complexes may be demonstrated here. Whereas with fluoride, iron forms a hexafluoro complex, the substitutions with chloride lead to formation of a tetrachloro complex. This is because chloride is more easilv oolarized. leadine to an accumulation of a negative charge on the metal.- his unfavorable situation is remedied bv a change of coordination number from six to four. The absorption of various metal ions on strongly basic ion-exchange resins in hydrochloric acid solutions has been extensively studied by Kraus and coworkers (5). Based on the data of these investigators a procedure for separation of iron from other metal ions was proposed (6). The same separation method is adopted after making a few alterations to suit the present purpose. The Separation Procedure Perchlorate and phosphate ions interfere in the ion.exchanyr separation of iron hg means of strongly hasic anion-exchange resins. For this rcaaon the sollrtion obtained hy thc drcomposirion of the ore with a mixture of perchlorie and phosphoric acid cannot he used directly in ion-exchange separation of i n n The sepnration of iron from the intertrring ions ma" be achieved by P I P C I ~ I tation of iron hydroxide with excess sodium hydroxide. The procedure is however time consuming. For this reason an alternative procedure involving decomposition of the ore by fusion with sodium peroxide is described here. A mixture of the ore (4 g) and sodium peroxide (12 g) is carefully fused in a nickel crucible, using a Bunsen burner far the purpose. The mixture is heated gently at first, to avoid loss of sample through frothing, and then strongly for 30 min in order to obtain a clear homogeneous liquid phase. The entire process takes about 45 min. After cooling, the sample cake is decomposed with 200 ml of 6 N HC1 in a 500-ml beaker. On heating the mixture for 30 min on a water bath, nearly all the iron is dissolved. The solution is filtered hot to remove preeipitated sodium chloride. The solution after addition of 400 ml eancentrated hydrochloric acid is passed through a 40 mm X 700 mm ion-exchange column packed with 300 ml Amberlite IRA 4001 previouslv washed with concentrated hvdrochloric acid. The flow rate is: maintained at 10 ml mi". T h e w h n n is then washed four times wrth lilu-ml porlrons of 8.5 h' HCI. TI, remove iron r'rum the resin 400 ml diatillrd water is passea through the column. Thr break-through for iron is shown by the appearance of a yellow effluent solution from the column. Only the fraction containing the yellow solution is collected. From the solution iron hydroxide is eentrifuobtained by precipitation with ammonia solution. ~ f t & gation, iron hydroxide is isolated, and transferred into a porcelain crucible and ignited at 700'C in an electric furnace. Ignition of iron hydroxide at higher temperatures should be avoided, hecause of the tendency of iron(IE) oxide to decompose to iron (11)oxide.
Amberlite IRA 400 is a product of Rohm and Haas,Philadelphia, Pennsylvania, and may be obtained through BDH, Poale, England. 2Titaniferous iron ore from Liganga in southwestern Tanzania was used. The author carried out three determinations and ahtained 50.8%, 50.5% and 50.9% iron. This is in agreement with results of other analysts (9).
Although Erdey (7) recommends the ignition of imn(II1) hydroxide at 8CQ°C, slight decomposition of iran(II1) oxide to iron(I1) oxide may possibly occur at temperatures above 700°C. (8)Quantitative conversion of ferric hydroxide to ferric oxide is effected by heating the hydroxide precipitate at 700'C for 1 hr in an electric furnace. After cooling in a desiccator, the product is weighed and the yield is calculated. A weighed sample is then dissolved in warm hydrochloric acid, and the solution obtained is used for the determination of iron with diehromate according to the procedure described previously. The purity of the product is assessed using the results of the analysis. Students Results A total of 35 students a t Dar es Salaam University carried out this experiment during the academic year 197273. Locally obtained titaniferous iron ore containing 51% iron and 13% rutile was provided.2 The results obtained by students were on the whole in good agreement with those expected. The results are summarized below. Number of students
90 Fe found
3
48-49 49-50 50-50.5 50.5-52 aver 52
9 14
5 4
The recovery of iron oxide from the ore by students was satisfactory. All 35 students obtained yields over 50%. To minimize the time spent on the experiment, repeated fusion which is often recommended for complete decomposition of the ore was not encouraged here. Therefore, 100% recovery of iron oxide was not anticipated. The results for iron recovery are shown below. Number o f students
Weight o/ Fez03 ~ng
% FezOa 9 1.52-1.75 52-60 15 1.75-2.05 60-70 8 2.05-2.16 70-75 3 2.16-2.35 75-81 Ferric oxide obtained by students was in general of satisfactory quality. The results of iron determination by 26 students gave values between 69.4% and 70.3%. The calculated percentage of iron in Fe.0~ is 69.94.
Conclusion Many students after completion of the experiment were of the opinion that it was the first time they did something serious in the laboratory. In a way a long experiment made the students feel they were doing something similar to research, where each part of the work had a significant connection with the whole experiment. The experiment was carried out in an average of 12 laboratory hours. Literature Cited 11) Furmsn, N. H., "Standard Methods of Chemieal Analysis," Vol. 1. D. Van Ncs-
trandCo,,Inc., Tomnto, 1962. 121 Hiliebrand, W. P., and Landell, G. E. P.. "Applied Inoqanic Analysis: John Wileyand Sona, h e . , New York. 1953. 13) Cumming and Kay, "Quantitative Chemical Analysis.: Oliver Boyd. London, 1956.
14) Cotton. F. A , and Wilkinaon. G.. "Advanced inoqmie Chemistry: Intencience. NewYork, 1 9 6 7 , ~859. . IS) Kraus. K. A,, and Nelson, F.. Pmr. 1st United Notions Conf on Peocdul Uses o/ AiornieEnergy. I, 11311953). Comparewithref. (6) 16) Waiton, F. H., "Principles and Methods of Chemical Analysi.: Rentice-Hall of Indialtd., NewDeihi. 1966~. 154. 17) Erdey. L.. "Grsvimetrie Analysis. Part n!' International Series of Monographs on Analytiesl Chemistry. Vol 7, PergamonPress. Oxford. 1965. (8) Mailer. G. 0.. "Praktikum der Qusnfifativen Chemischen Analwe: S. Hinel. Iripzig, 1962 221. (9) Harris. J. F..'P~ummaryof the Geology of Tanganyib, Part IV: Government Press. Oar es Salaam. 1951. p. 89.
Volume 50, Number 9, September 1973 / 635
