Oxidation of sulphide minerals—I

T&nro.

Vol. 24, pp. 251-254.

Pergamon

Press,

1977. Printed

OXIDATION

in Great Britain

OF SULPHIDE

MINERALS-I

DETERMINATION OF FERROUS AND FERRIC IRON IN SAMPLES PYRRHOTITE, PYRITE AND CHALCOPYRITE*

OF

H. F. STEGER Mineral Sciences Laboratories, Canada Centre for Mineral and Energy Technology, Department of Energy, Mines and Resources, Ottawa, Canada (Received

24 August

1976. Accepted

24 October

1976)

Summary-A method has been developed for determining small amounts of both ferrous and ferric iron in oxidized samples of pyrrhotite, pyrite and chalcopyrite. The oxidized iron is selectively dissolved in 10M phosphoric acid under reflux and can be determined with the accuracy generally accepted in chemical phase analysis.

As part of a general study to establish the “shelf-life” or stability of certified reference ores and concentrates,’ it is desirable to understand more fully the air-oxidation of sulphide minerals. In order to do so for pyrrhotite, Fe12S13, pyrite, Fe&, and chalcopyrite, FeCu$, it is necessary to be able to determine the ferrous and ferric iron in the oxidation productst in these minerals. For brevity, the expressions “ferrous and ferric iron from pyrrhotite, etc.” and “oxidized iron” will be used hereafter to describe the iron in the oxidation products and do not refer to the pure sulphide minerals. In general, the determination of ferrous and ferric iron in ores and minerals is subject to appreciable error if even a small amount of sulphide is present, because the H2S evolved on dissolution of these sulphides readily reduces ferric to ferrous iron.‘e3 It is evident, therefore, that the determination of small amounts of oxidized iron in pyrrhotite, pyrite or chalcopyrite samples could be very troublesome. Moreover. for samples of these minerals, a determination of the total oxidized iron can give high results because the H2S consumed in the reduction of ferric to ferrous iron is lost, so that a true correction for the amount of iron derived from dissolved sulphide mineral cannot be applied to the total oxidized iron. Because pyrite and chalcopyrite are relatively inert to non-oxidizing acids of moderate strength,495 the selective dissolution and subsequent determination of oxidized iron presents only minor difficulty. Pyrrhotite, however, is readily attacked by such acids and the selective dissolution of oxidized iron without some dissolution of the pyrrhotite is difficult. Of a series of non-oxidizing acids, both organic and inorganic, that were tested, 10M phosphoric acid was * Crown Copyrights reserved. t The oxidation products are iron oxide. hydroxide, carbonate, sulphate, thiosulphate, etc. and both ferrous and ferric iron can be present. 251

found to permit the determination of oxidized iron associated with pyrrhotite and, of course, also with pyrite and chalcopyrite. This paper describes the conditions for such a determination. EXPERIMENTAL Pretreatment

of mineral samples

Lumps of high-purity monoclinic pyrrhotite, pyrite and chalcopyrite (Wards Scientific) were crushed mechanicallv and the -20&325 mesh fraction (Tyler), i.e., 674 ~4, was collected for subsequent study. The non-sulphide impurities were removed from this fraction by density separation with di-iodomethane (s.g. 3.3). In order to remove as much as possible of the oxidation products formed during crushing, approximately 20 g of each sulphide mineral were shaken for 2 hr with 100 ml of ammonium sulphide solution6 freshly prepared by saturation of 400 ml of water and 14 ml of concentrated ammonia solution with HZS gas. The minerals were then filtered off, washed with deaerated water, anhydrous methanol and ethyl acetate, and dried under suction. The minerals were stored in vials (previously flushed with nitrogen) in a desiccator. Controlled

oxidation

of mineral samples

For each mineral, 5 I .5-g samples were weighed into each of five 30-ml crystallization dishes, which were then placed in a controlled temperature and humidity chamber set at 52” and 70% relative humiditv (RH). The samoles were removed from this chamber ai ;ari&s preselected times, dried in a desiccator for 18-24 hr and weighed to determine Aw, the change in weight due to oxidation, before analysis for oxidized iron. Two additional samples of pyrrhotite (- 2.5 g) were treated in this way for 1 and 2 weeks. An X-ray diffraction analysis of these latter samples of pyrrhotite indicated the presence of goethite, FeO(OH), and elemental sulphur. Determination

of oxidized

iron

All chemicals used were analytical grade. The reagents were 10M phosphoric acid (680 ml of concentrated acid diluted to 1 litre), ammoniacal cadmium chloride solution (23.7 g of CdClz.2:H20 dissolved in water, followed by addition of 200 ml of concentrated ammonia and dilution to 2 litres), sulphuric acid (1 + 3), iodate-iodide solution (2.44 g of KIOJ, 100 g of KI and 2 g of NaOH dissolved

252

H. F. STEGER

in water and diluted to 2 litres), and thiosulphate solution (22.7 g of Na,Sz03.5H20 and 0.2 g of Na2C03 dissolved and diluted to 2 litres with deaerated water. Recommended procedure

A portion (0.34.5 g) of each sample is transferred to a lOO-ml two-neck round-bottom flask equipped with a reflux condenser and a nitrogen inlet tube that extends to near the bottom of the flask. Then 25 ml of 10M phosphoric acid are added and the contents refluxed for 3.5 min. After cooling in ice-water for 5 min, the contents are transferred to a 200-ml volumetric flask, diluted to the mark and filtered through a Whatman No. 42 paper. The nitrogen used to purge the flask during both heating and cooling is passed through 60 ml of ammoniacal cadmium chloride solution in a 100-ml Erlenmeyer flask to collect any HIS evolved due to dissolution of the sulphide mineral. The HIS evolved is determined iodometrically. After the addition of 10.00 ml of iodate-iodide solution to the ammoniacal cadmium chloride solution, 25 ml of sulphuric acid (1 + 3) are added and the excess of iodine is backtitrated with thiosulphate. The titration was performed in this laboratory with a Mettler Automatic Titrator equipped with a Radiometer PP1311 twin platinum wire dead-stop electrode. The quantity of iron arising from the dissolution of the sulphide mineral is readily calculated from the amount of H2S evolved. This iron must be subtracted from both the ferrous and total oxidized iron derived from oxidation products of the sulphide mineral or other iron-bearing materials in the sample. The ferrous iron in the filtrate is determined spectrophotometrically as its l,lO-phenanthroline complex.’ After the determination of total oxidized iron, the ferric iron can be calculated by difference. In this study, the total oxidized iron was determined with a Techtron AA6 spectrophotometer, with standard solutions having a phosphoric acid content very similar to that of the oxidized iron samples. RESULTS

AND DISCUSSION

Some products such as iron sulphate, thiosulphate, etc. formed by oxidation of the iron sulphide minerals are, in general, soluble in water. Others, e.g., goethite, the major product from pyrrhotite, and probably formed to some extent from pyrite and chalcopyrite,

require acidic dissolution media. In fact, the goethite formed from pyrrhotite in this investigation seemed to be soluble only in acid that is more concentrated than OSM and heated under reflux. This material was, however, shown to be readily attacked by phosphoric acid. EfSect of phosphoric

acid concentration

Figure 1 shows the effect of phosphoric acid concentration in the determination of oxidized iron in a sample of pyrrhotite that was maintained at 52 and 70% RH for 1 week. The numbers below the points for total iron give the reflux time (min) required for complete dissolution of the goethite (visual observation) plus 1 min extra. Also depicted is the quantity of ferrous iron that can be attributed to the dissolution of pyrrhotite; this has been subtracted from the results for ferrous and total iron. Figure 1 illustrates that for a phosphoric acid concentration less than 9M the amount of pyrrhotite dissolved which can be accounted for from the H2S evolved is accompanied by a greater amount from which the H2S is not evolved but reduces ferric to ferrous iron. Moreover this additional dissolution of pyrrhotite causes high values for total iron. At a phosphoric acid concentration greater than 9M, no dissolution of pyrrhotite (as detected by the evolution of H2S) is observed and both the ferrous and total iron values appear to be constant within experimental error. The reason for this is probably the shorter time available for dissolution of any goethite. A phosphoric acid concentration of 10M was chosen for the procedure because it lies in the region of acidity where a constant value of ferrous and total iron is obtained but still permits the determination of ferrous iron.’ The use of a higher concentration would require partial neutralization before ferrous iron could be determined, and addition of base to the phosphoric acid filtrate containing the ferrous iron causes rapid oxidation of the ferrous iron to ferric. Effect of 10M phosphoric

-

‘IO

-

100

-90

Fig. 1. The effect of phosphoric acid concentration on the determination of ferrous and ferric iron in pyrrhotite. @-Ferrous iron calculated from HIS evolved. +-Ferrous iron. O-Ferric iron. A-Total iron.

acid on ferrous

iron

Because ferric iron has a strong affinity for phosphate, the oxidation of ferrous iron proceeds readily in phosphate medium. The possibility of oxidation of ferrous iron during the refluxing with 10M phosphoric acid, dilution to 200 ml or filtration was investigated and oxidation shown to occur. The calibration curves for the ferrous 1,lO-phenanthroline complex formed (a) from solutions of ferrous ammonium sulphate in water and (b) from equivalent solutions added to 25 ml of 10M phosphoric acid refluxed under nitrogen for 3.5 min, diluted and filtered, are parallel but whereas that for (a) is linear and passes through the origin, (b) gives an intercept on the ironconcentration axis. This possibility of oxidation is also present in the determination of ferrous iron in iron sulphide samples and it is therefore imperative that a calibration curve of type (b) be used for mineral analysis. It is also important to note that 0.25 mg of

Oxidation

of sulphide

Table 1. Determination of known quantities of Fe(B) and Fe(III) in pyrrhotite WI), Added 0 1.1 2.0 3.6 5.6

mg

253

minerals-I

Table 2. Effect of sample weight on Fe(II) results for pyrrhotite Fe(I1)

WIII), mg

Found

Found, minus pyrrhotite

Added

Found

Found, minus pyrrhotite

0.1 1.7 2.6 4.1 6.2

1.0 1.9 3.4 5.5

0 5.6 3.1 2.2 1.2

16.3 22.0 20.1 18.7 17.6

5.1 3.8 2.4 1.3

ferrous iron is the minimum that can be detected. Consequently a zero ferrous iron content obtained for a particular sample weight of mineral should be verified by another determination on an appreciably larger sample weight. Accuracy and reproducibility of recommended method The accuracy or even the applicability of the recommended method in determination of oxidized iron in iron sulphide minerals cannot be verified by comparison with results obtained by another existing analytical method. This must be done by the determination of known amounts of ferrous and ferric iron added to iron sulphide samples. Table 1 gives the results for a series of known additions of Fe(NH& (S0&.7H20 and Fe203 to 0.1800 +0.0005-g samples of the pyrrhotite treated at 52” and 70% RH for 1 week. It is evident that both the added ferrous and ferric iron can be determined with acceptable accuracy. Figure 2 shows the ferrous and ferric iron in the samples of pyrrhotite, pyrite and chalcopyrite as a function of the change of weight observed in one oxidation run. The good linear relationship between ferric iron and Aw again suggests that the recommended method is well suited to the determination of oxidized iron in iron sulphide minerals. The constancy of the ferrous iron content with increasing oxidation was unexpected and it was considered necessary to determine whether the 10M phosphoric acid exerted a levelling effect on the ferrous iron. The ferrous iron in the pyrrhotite treated

Sample

weight,

0.1007 0.3028 0.6013 l.Oc09

g

found,

mg

0.5 1.4 3.0 5.0

Fe(W, mglg 4.9 4.6 5.1 5.0

for 2 weeks was determined on 4 different sample weights. The results are summarized in Table 2. The good agreement for results from such a wide range of sample weights clearly shows the absence of any levelling effect. It can therefore be concluded that the quantity of ferrous iron is not increased under the conditions used in this study. The nature of this ferrous iron will be discussed in a future study of the oxidation process of iron sulphide minerals. The results in Table 2 suggest that the reproducibility of the recommended method is kO.03 mg for 4.9 mg of Fe(I1) per g of pyrrhotite. If, however, the ferrous iron values are compared for samples of sulphide minerals within any one oxidation run, the reproducibility is kO.5 mg for 3-5 mg/g for 27 samples of pyrrhotite in 5 different runs, f0.3 mg for 35 mg/g for 12 samples of pyrite in 2 runs, and f0.2 mg for l-l.5 mg/g for 12 samples of chalcopyrite in 2 runs. That is to say, there is an uncertainty of l&20% in the ferrous iron content. This uncertainty is high but still remains acceptable in chemical phase analysis, particularly for small amounts of the phase concerned.2*3 The reproducibility for ferric iron can, of course, be no better than that for ferrous iron because the former is given by the difference between total iron and ferrous iron. On the assumption that the determination of the total iron can be done with greater reproducibility than the phase separation, it seems reasonable that the overall reproducibility for the total iron might be similar to that for ferrous iron. Accordingly, the reproducibility for ferric iron should be approximately twice that for ferrous iron. Because the amount of ferric iron varies with the extent of oxidation, the uncertainty will decrease with increasing oxidation. For example, a reproducibility for Fe(II1) of -0.6 mg/g leads to an uncertainty of 20% at the 3 mg/g level but only 6% at 10 mg/g. Effect of dissolution of sulphide minerals

Fig. 2. Ferrous and ferric iron in pyrrhotite,' pyrite and chalcopyrite. Circles-ferric iron. Squares-ferrous iron.

The 10M phosphoric acid used is not a completely selective solvent for oxidized iron species in sulphide minerals. Occasionally some attack on the sulphide mineral is observed (as evolution of H,S), particularly for pyrrhotite. Such attack was not noted in 12 tests with chalcopyrite and only twice in 12 tests on pyrite, performed for this and subseq.uent oxidation studies. In both instances for pyrite, the ferrous iron derived from dissolution of the pyrite was 0.25 mg/g, which is only 5% of the total ferrous iron determined (5

254

H. F.

mg/g). It should be noted that this dissolution of pyrite is not accompanied by the undeterminable dissolution in which the H2S formed reduces ferric iron to ferrous (see Fig. 1) (the ferrous iron content for these two pyrite samples was the same as that of the other samples in the same oxidation run. Attack on pyrrhotite occurred in 8 of 27 experiments performed in this or subsequent studies. It was concluded that the extent of attack on pyrrhotite decreases with increasing extent of oxidation and, in fact, only one sample of pyrrhotite, treated for longer than 4 days at 52” and 70% RH, has thus far been found susceptible to attack by 10M phosphoric acid. The ferrous iron content of the pyrrhotite samples which were attacked by the acid was the same within the estimated reproducibility as that of samples of the same oxidation run not undergoing attack. This again suggests that ferric iron in 10M phosphoric acid is stable towards reduction by the small amounts of H2S arising from dissolution of the iron sulphide. F’yrrhotite that has been freshly treated with ammonium sulphide is very readily attacked by 10M phosphoric acid. Consequently, it was necessary to reduce the reflux time to 0.5 min, but even with this change the ferrous iron from the dissolution of pyrrhotite was several times greater than that due to oxidation. The latter was, however, the same within the estimated reproducibility of the recommended method as that for the pyrrhotite samples where no dissolution was observed. There is no danger of reducing ferric to ferrous because, in fact, ferric iron has, as yet, not been detected in 5 samples of pyrrhotite freshly treated with ammonium sulphide. The

STEGER

attack by 10M phosphoric acid on a freshly prepared sample of pyrrhotite and samples oxidized for short periods can probably be attributed to the relatively clean unoxidized surfaces of these samples. CONCLUSION

A method has been developed for determining ferrous and ferric iron in the oxidation products of pyrrhotite, pyrite and chalcopyrite. Although this method was applied only to an oxidation study of samples of these minerals which had undergone a special “cleaning up”, it can also be used for naturally-occurring samples of these minerals to determine their extent of oxidation, i.e., weathering. Moreover, there is some doubt about the safety of marine transport of chalcopyrite concentrate, and use of the described method for analysing this copper concentrate before and after transport could give a measure of the extent of oxidation undergone during transport.

REFERENCES

I. H. F. Steger. T&m.

1976. 23. 643.

2. R. S. Young, Chemicul Phase Analysis, Halstead-Wiley, New York, 1974. 3. H. F. Steger, Talanta, 1976, 22, 81. 4. Yu. V. Morachevskii and N. Kh. Pinchuk, I/em. Leningrad, Univ., 1956, 170. 5. A. M. Gaudin and N. P. Finkelstein, Nature, 1965, 201, 309. 6. S. Yu. Fainberg, Analysis of the Ores of Non-Ferrous Metals, Metallurgizdat, Moscow, 1953. I. H. Tamura, K. Goto, T. Yotsuyanagi and M. Nagayama, Tafanta, 1974, 21, 314.