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Precise determination of ferrous iron in silicate rocks
Geochimica et Cosmochimica Acta, Vol. 66, No. 6, pp. 1085–1093, 2002 Copyright © 2002 Elsevier Science Ltd Printed in the USA. All rights reserved 0016-7037/02 $22.00 ⫹ .00
Pergamon
PII S0016-7037(01)00809-2
Precise determination of ferrous iron in silicate rocks TETSUYA YOKOYAMA and EIZO NAKAMURA The Pheasant Memorial Laboratory for Geochemistry and Cosmochemistry, Institute for Study of the Earth’s Interior, Okayama University at Misasa, Misasa, Tottori-ken 682-0193, Japan (Received October 20, 2000; accepted in revised form August 22, 2001)
Abstract—We have developed a highly precise method for the determination of ferrous iron (Fe2⫹) in silicate rocks. Our new method is based on Wilson’s procedure (1955) in which surplus V5⫹ is used to oxidize Fe2⫹ into Fe3⫹ while equivalently reducing V5⫹ into V4⫹. Because V4⫹ is more resistant to atmospheric oxidation than Fe2⫹, Fe2⫹ in the sample can be determined by measuring unreacted V5⫹ by adding excess Fe2⫹ after sample decomposition and then titrating the unreacted Fe2⫹ with Cr6⫹. With our method, which involves conditioning the sample solution with 5 M H2SO4 in a relatively small beaker (7 mL), the oxidation of Fe2⫹ or V4⫹ that leads to erroneous results can be completely avoided, even in 100-h sample decompositions at 100°C. We have measured the concentration of FeO in 15 standard silicate rock powders provided by the Geological Survey of Japan (GSJ). Analytical reproducibility was better than 0.5% (1) for all but those samples that had small amounts of Fe2⫹ (⬍1.5 wt.% of FeO). Fourteen of these samples gave FeO contents significantly higher than the GSJ reference values. This likely indicates that the GSJ reference values, obtained by compiling previously published data, contain a large number of poor-quality data obtained by methods with lower recovery of Fe2⫹ caused by oxidation or insufficient sample decomposition during analyses. To achieve accurate determinations of Fe2⫹ in our method, several factors besides the oxidation must be considered, including: (1) long-term variations in the concentration of Fe2⫹ solution must be corrected; (2) excess use of the indicator must be avoided; and (3) the formation of inert FeF⫹ complex must be avoided during titration when using boric acid as a masking agent. Copyright © 2002 Elsevier Science Ltd oxidation of Fe2⫹. In our preliminary analyses of basaltic, andesitic, and rhyolitic standard rock powders using Pratt’s method, we obtained relative values 2 ⬃ 5% lower than the reference values, although the accuracy of the reference values themselves is difficult to assess. One of the most effective ways to achieve accurate determinations of Fe2⫹ is Wilson’s method (Wilson, 1955) or Peters’ method (Peters, 1968), in which excess V5⫹ is added to the sample before decomposition. This method is based on the following reversible reaction:
1. INTRODUCTION
Determination of ferrous iron (Fe2⫹) and Fe2⫹/Fe3⫹ ratio in silicate rocks is important in geological, geochemical, and petrological studies for evaluation of weathering (Koch et al., 1995), determination of oxygen fugacity in magma (Sack et al., 1980; Thornber et al., 1980; Kilinc et al., 1983; Carmichael and Ghiorso, 1986; Christie et al., 1986; Kress and Carmichael, 1988; Kress and Carmichael, 1991), and understanding chemical equilibrium between minerals and silicate melt (Roeder and Emslie, 1970; Beattie et al., 1991; Beattie, 1993). Recent analytical progress using x-ray fluorescence (XRF) methods enables highly precise determination of total iron in silicate rock samples with analytical error less than 0.5% (1) (relative standard deviation; e.g., Takei and Nakamura, in preparation). Determination of Fe2⫹ concentrations with the same precision as that for XRF analyses of total iron would make it possible to more precisely address the above petrological issues. Several analytical techniques for the determination of Fe2⫹ in silicate rocks are reviewed elsewhere (Bock, 1979; Potts, 1987; Amonette et al., 1999). Although high-precision analyses for Fe2⫹ are relatively easily obtained, the accuracy of these analyses is more questionable. Fe2⫹ can easily be oxidized during analytical procedures, resulting in erroneously low values. In the classical method of Pratt (Pratt, 1894), which is commonly used in many laboratories, the sample is decomposed by mixtures of HF and H2SO4 in a covered platinum crucible that excludes atmospheric oxygen. This relatively difficult technique often yields low concentrations owing to the
* Author to whom correspondence ([email protected]).
should
be
Fe2⫹ ⫹ V5⫹ ⫽ Fe3⫹ ⫹ V4⫹
(1)
which was first found to be effective in determination of Fe2⫹ by Ishibashi and Kusaka (1950). This reaction proceeds to the right side in the condition of strong acidity, and the amount of the V4⫹ reduced from V5⫹ is then the same as the amount of Fe2⫹ in the sample. Because V4⫹ is extremely resistant to oxidation and does not back-react to V5⫹ during the analytical procedure, Fe2⫹ concentration in the sample can be determined by measuring the unreacted V5⫹ by adding excess Fe2⫹ after sample decomposition and then titrating the unreacted Fe2⫹ with Cr6⫹. Whipple (1974) improved Wilson’s method and evaluated its accuracy and precision by determining the recovery of Fe2⫹ using ammonium iron(II) sulfate hexahydrate as a starting material. He achieved determinations of 7 mg of Fe2⫹ with precision of ⬃0.3%. In his method, however, a 1% correction was required due to small amounts of oxidation of Fe2⫹ or V4⫹ during decomposition, even though he deoxidized the reagents by N2 purging for 4 h before analysis. Whipple also attempted to decompose the sample in a closed N2 box, but this resulted in up to 0.5% loss of Fe2⫹.
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Amonette and Scott (1991) developed a method that improved upon Peters’ method, and achieved better determinations of Fe2⫹ (0.2 to 0.9%, 1) in 5 geochemical reference samples and 3 mica samples. However, they recommended that sample decomposition must be carried out at 60°C within 24 h, because they found that higher temperature and longer decompositions cause loss of Fe2⫹ due to oxidation. These conditions are, however, insufficient to completely decompose some minerals containing Fe2⫹ such as pyroxene, garnet and spinel, these being more acid-resistant than mica. In this study, we have reevaluated the methods of Wilson (1955) and Whipple (1974) and developed more routine, more accurate methods (better than 0.5% precision, 1) while overcoming some problems that occurred in previous studies. We provide new determinations of the FeO concentration (wt.%) in 15 standard silicate rock powders. 2. EXPERIMENTAL SECTION
All of the experiments in this study were carried out at the Pheasant Memorial Laboratory (PML), Institute for Study of the Earth’s Interior in Misasa.
2.1.4. Indicator 0.2 g of sodium diphenylamine-4-sulfonate (first grade, Kanto Chemical) was dissolved in water and diluted to 100 mL resulting in 0.2% (w/v). This solution was stored in a 125-mL polyethylene dropping bottle. The titration is complete when the color changes from gray to violet. 2.1.5. H2SO4–H3PO4 mix solution 400 mL of 18 M H2SO4 was slowly added to 400 mL of water in a draft chamber while cooling in a water bath. After cooling to room temperature, 200 mL of 15 M H3PO4 was carefully mixed. 2.1.6. HF-H2SO4 mix solution 33.3 mL of 18 M H2SO4 was slowly added to 86.7 mL of 27 M HF in a draft chamber while cooling in a water bath. After cooling to room temperature, it was stored into a 125-mL polyethylene dropping bottle. This produces a 5 M concentration of H2SO4 in this solution. 2.1.7. 0.3 mmol/g Fe2⫹ in 1 M H2SO4 solution
2.1. Reagents and Apparatus 2.1.1. Water Water was deionized and purified using a mixed-bed resin and filters (WL21P, Yamato, Japan). Typical specific resistance of the water was better than 16.7 M ⍀䡠cm. 2.1.2. Hydrofluoric, sulfuric, and phosphoric acids 27 M HF (special grade, Wako, Japan), 18 M H2SO4 (trace analysis grade, Kanto Chemical Co., Japan), and 15 M H3PO4 (special grade, Wako) were used without any further purification. These acids are highly toxic and corrosive. They irritate the skin and mucous membranes of the body. Thus, to avoid inhalation, ingestion, and contact with the skin or eyes, one should wear protective gloves, glasses, and clothing while using a chemical fume. 2.1.3. Titrant 1.96 g of potassium dichromate reference material for volumetric analysis (Kanto Chemical, 99.98 ⫾ 0.01%) was dissolved in water and diluted to 1000 mL using measuring flask, resulting in 0.04 N K2Cr2O7 solution. We used this solution as a primary standard throughout this study. The solution was stored in a tightly sealed 1000-mL glass bottle. Although the concentration of the solution did not change after 1 month, it increased by 0.3% after 7-month storage (probably due to the evaporation of water), based on a calibration by ammonium iron(II) sulfate hexahydrate. Therefore, it is recommended to use solid ammonium iron(II) sulfate hexahydrate as a secondary standard by determining its absolute Fe2⫹ concentration with newly conditioned K2Cr2O7 solution, and calibrate the titrant when it is stored over a few months.
13.8 g of ammonium iron(II) sulfate hexahydrate (analytical grade, Merck, Germany) was dissolved into 100 mL of 1 M H2SO4 (⬃103.5 g) and stored in a 125-mL polyethylene dropping bottle. The molar concentration of this solution was determined as follows: Approximately 0.7 g of the solution was weighed in a Pyrex glass beaker, followed by addition of 5 mL of H2SO4–H3PO4 solution. The resulting mixture was then diluted to 40 mL by the addition of water. One drop of the indicator was added to the beaker and the solution was titrated with 0.04 N K2Cr2O7. The mean value of five successive titrations was used as the molar concentration. 2.1.8. 0.3 mmol/g V5⫹ in 5 M H2SO4 solution 4.69 g of sodium metavanadate (special grade, Kanto Chemical) was dissolved into 100 mL of 5 M H2SO4 (⬃123.5 g) and stored in a 125-mL polyethylene dropping bottle. The molar concentration of this solution was determined as follows: approximately 0.7 g of the solution and 1.4 g of Fe2⫹ solution were weighed in a Pyrex glass beaker. The subsequent procedure was the same as that described in section 2.1.7. 2.1.9. Saturated boric acid 100 g of special grade boric acid (Kanto Chemical) was dissolved in 500 mL of hot water (80°C). After cooling to room temperature, the supernatant was stored and used in the analytical procedures. 2.1.10. Burette system A piston burette (Titronic 96, Schott, Germany) was used for titration in the course of this study. This system has a 20-mL glass cylinder and enables spewing of the solution in 0.01-mL aliquots. The calibration of the dosing unit was carried out by repeatedly spewing and weighing 5.00, 10.00, and 15.00 mL of water, and the accuracy of the volume was within 0.15%
Precise determination of Fe2⫹ in silicate rocks
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relative to the digitally displayed volume. The precision of the system was less than 0.05% when 5.00 mL of water was spewed, and it decreased as the volume increased. 2.1.11. Weighing system A microbalance with 5 decimal places (R200D, Sartorius, Germany) was used for weighing samples. Typical uncertainty of this balance was ⫾0.00010 g (1). It is important to eliminate static electricity of the sample before weighing, and a static eliminator (Endstat II, Otsubo Electronics, Japan) was used for this purpose. 2.2. Procedure 2.2.1. Determination of Fe(II) in silicate rocks Fifteen GSJ (Geological Survey of Japan) standard silicate rock powders were analyzed in this study, including JP-1 (peridotite); JB-1b, JB-2, and JB-3 (basalts); JA-2 and JA-3 (andesites); JR-1, JR-2, and JR-3 (rhyolites); JH-1 (hornblendite); JGb-1 and JGb-2 (gabbros); JG-1a and JG-3 (granodiorites); and JG-2 (granite). Elemental concentrations of these rock powders are given in Terashima et al. (1998) for JB-1b, Imai et al. (1999) for JR-3, JGb-2, and JH-1, and Imai et al. (1995) for the rest of the samples. All rock powders were analyzed in the stock in which they were provided by GSJ, without any further pulverization. Starting material containing less than 0.15 mmol of Fe2⫹ was weighed into a 7-mL clean Teflon beaker. The amount of sample used for the determinations did not exceed 0.4 g. Approximately 0.7 g of V5⫹ solution was then added to the beaker and carefully weighed. The HF–H2SO4 solution was subsequently added to the sample at a proportion of 1.5 mL/0.2 g of silicate sample. The beaker was tightly sealed and agitated in an ultrasonic bath for 20 min. The beaker was then heated overnight on a hot plate at 100°C. All the rock powder was decomposed under such a condition, but a certain proportion of spinel and zircon remained undissolved, as discussed later. After sample decomposition, ⬃1.0 g of Fe2⫹ solution was carefully weighed into a clean Pyrex beaker and 5 mL of the H2SO4–H3PO4 solution was added. Afterwards, saturated boric acid was added to the Pyrex beaker at a proportion of 10 mL/1.5 mL of the HF–H2SO4 solution that was used for sample decomposition. All of the decomposed material was moved into the Pyrex beaker by carefully washing the inside of the Teflon beaker with H2O. The recovered solution was diluted to 40 mL with H2O, then one drop of indicator was added before titration with 0.04 N K2Cr2O7. The concentration of Fe2⫹ was calculated by the following equation: FeO (wt.%) ⫽ (VCr 䡠 CCr ⫹ WV 䡠 CV ⫺ WFe 䡠 CFe) 䡠 10⫺3 Wsamp ⫻ 71.85 ⫻ 100
(2)
where VCr was the amount of titrant (mL); WV, WFe, and Wsamp were weights of the V5⫹ solution, Fe2⫹ solution, and sample, respectively (g); CCr was the normality of the titrant
Fig. 1. Relationship between recovery yield of Fe2⫹ and heating time of the sample with 2 M (}), 3 M (f), and 5 M (F) H2SO4 conditions.
(N); and CV and CFe were the molar concentrations of V5⫹ and Fe2⫹ solutions, respectively (in mmol/g). 2.2.2. Recovery yield measurement of Fe2⫹ Approximately 0.06 g of solid ammonium sulfate hexahydrate (⬃0.15 mmol of Fe2⫹) was weighed into a 7-mL clean Teflon beaker, after which ⬃0.7 g of V5⫹ solution was added. 1.0 mL of HF, together with variable amounts of H2SO4, were then added to examine the effect of varying acidity. The sample was treated the same as described above except the heating times varied from 3 to 100 h to examine the effect of the side reaction during the procedure. The amount of Fe2⫹ in the Teflon beaker was used to calculate recovery yield, normalizing to the initial concentration of Fe2⫹ in the ammonium iron(II) sulfate hexahydrate (2.550 mmol/g). 3. RESULTS AND DISCUSSION
3.1. Performance of the Method The most plausible source of analytical error in the determination of Fe2⫹ by our method is from side reactions such as the oxidation of Fe2⫹ during sample decomposition. To evaluate the accuracy of our method, we repeatedly measured the recovery yield of Fe2⫹ throughout the chemical procedure by using ammonium iron(II) sulfate hexahydrate as a starting material while varying both the acidity and the heating time. As shown in Figure 1, the recovery yield is strongly affected by both of these variables. 100% of the Fe2⫹ was recovered within 100 h heating when the molar concentration of H2SO4 was 5 M. The 3 M condition was effective for 100% recovery unless the beaker was heated for more than 42 h. However, it was difficult to obtain correct values when the acidity was 2 M. These results can be explained by the fact that stronger acidity drives reaction (1) to the right side. In the condition of 5 M, reaction (1) must be kept completely to the right side, and all Fe2⫹ is oxidized
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20 mL V5⫹ solution in 4.5 M H2SO4 ⫹ 5 mL HF (corresponding to 3.6 M H2SO4 overall) at 60°C for 48 h. Conversely, we found no Fe2⫹ loss in 3 M H2SO4 for 42 h. This is probably due to the difference in the size of the decomposing system (i.e., 125-mL beaker with 25 mL of solution for Amonette, 7-mL beaker with 3⬃5 mL of solution for ours), reflecting a difference in the extent of atmospheric oxidation. The reduction of the decomposing size is achieved by reducing the volume of V5⫹ solution while increasing its concentration compared to the previous methods, and determining its amount by weighing. Table 1 shows the relationship between the standing time and the concentration of FeO in JB-3 determined by our method with a constant acidity in 5 M H2SO4. The nearly constant value of FeO suggests that the side reaction with oxygen at 5 M condition does not occur, even in analyses of silicate samples. Another possible problem that yields an incorrect FeO value in analyses of silicate samples is the coprecipitation of Fe2⫹ with insoluble fluorides associated with Mg, Ca, or Al, which form during sample decomposition with HF. Langmyhr and Kringstad (1966) investigated the formation of insoluble fluorides during HF digestion by artificially synthesizing them from various combination of elements, and noted the occurrence of Fe2⫹(Al, Fe3⫹)F5䡠3.7 H2O in the Al-Fe-HF system. However, this system does not contain Mg or Ca, which are commonly abundant in silicate rocks. No insoluble fluoride containing Fe was detected by X-ray diffraction analysis in the HF digestion of mafic rocks (Yokoyama et al., 1999) and felsic rocks (Takei et al., 2001). Instead, CaAlF5, CaMg2Al2F12, Na0.88Mg0.88Al1.12(F,OH)6䡠H2O, MgF2 and AlF3 were the dominant fluoride minerals. Thus, we conclude that the formation of insoluble fluorides does not affect our FeO determinations.
Table 1. FeO (wt.%) in JB-3 with changing heating time. Hours
FeO
3 3 9 15 15 18 22 22 26 26 26 42 96 96 Average 1
7.98 8.04 8.02 8.00 8.02 8.02 7.97 7.94 7.98 8.02 8.01 8.02 8.00 8.04 8.00 0.03
into Fe3⫹ by reducing an equivalent amount of V5⫹ into V4⫹. From our results, it can be concluded that V4⫹ is very resistant to oxidation in the 5 M condition even when dissolved oxygen exists in the system. In our method, therefore, predeoxidization of reagents such as V5⫹ solution and HF is not required. However, at weaker acidity, a small proportion of the Fe2⫹ remains unconsumed by reaction, resulting in lower recovery of Fe2⫹ owing to its oxidation at longer heating times. Our results explain the results of Whipple (1974) that demonstrated 1% loss of Fe2⫹ after 20 h of standing time. In his experiment, the whole-sample acidity was lower than that of the 2 M H2SO4 condition, because he used 4 mL of V5⫹ solution in 5 M H2SO4 with 6⬃8 mL of HF added. In our method, the loss of Fe2⫹ in the 2 M condition with 24-h standing was 3% (mean value), slightly larger than that of Whipple’s experiments. This is probably due to heating of our samples, which was not done in Whipple’s study. Even without heating, the oxidation of Fe2⫹ could not be avoided when the sample acidity was less than 2 M H2SO4. On the other hand, Amonette and Scott (1991) showed 2% loss of Fe2⫹ in a sample that was decomposed in a mixture of
3.2. Determination of Fe2ⴙ in the GSJ Standard Rock Powders Table 2 shows the concentration of FeO for 15 GSJ standard rock powders determined in this study. Also shown are the GSJ reference values obtained from Imai et al. (1995) and Imai et al.
Table 2. FeO (wt.%) of 15 GSJ standard rocks and their recommended or preferable (asterisked) values. Sample JB1-b JB-2 JB-3 JA-2 JA-3 JR-1 JR-2 JR-3 JP-1 JGb-1 JGb-2 JG-1a JG-2 JG-3 JH-1 a b
FeO (1)
1 (%)
n
Recomm. (⫾1)
n
Diff. %
5.15 ⫾ 0.01 10.04 ⫾ 0.03 8.00 ⫾ 0.03 3.75 ⫾ 0.01 4.99 ⫾ 0.02 0.510 ⫾ 0.004 0.448 ⫾ 0.004 1.99 ⫾ 0.01 6.81b ⫾ 0.06 10.89 ⫾ 0.03 5.54 ⫾ 0.02 1.40 ⫾ 0.01 0.655 ⫾ 0.005 2.016 ⫾ 0.002 8.20 ⫾ 0.01
0.23 0.34 0.35 0.22 0.49 0.80 0.90 0.38 0.90 0.30 0.43 0.92 0.79 0.09 0.09
5 5 14 5 5 4 5 4 4 4 4 4 6 4 4
5.16a 9.98 ⫾ 0.30 7.85 ⫾ 0.17 3.69 ⫾ 0.12 4.83 ⫾ 0.19 0.49 ⫾ 0.10 0.44 ⫾ 0.09 1.86 ⫾ 0.10 5.99 ⫾ 0.42 9.43 ⫾ 0.47 5.41 ⫾ 0.24 1.36 ⫾ 0.22 0.57 ⫾ 0.07 1.83 ⫾ 0.21 8.09* ⫾ 0.26
40 20 15 13 21 17 6 13 24 6 22 14 11 4
⫺0.2 0.6 2.0 1.7 3.4 4.0 1.7 7.0 13.6 15.5 2.4 2.9 14.9 10.2 1.4
After Terashima et al. (1998). Not recommended nor preferable value. The contribution of residual spinel is not included. See the text.
Precise determination of Fe2⫹ in silicate rocks
(1999), and the differences between our data and reference concentrations. Standard deviations of reference values are obtained from the database available at the GSJ internet website, http://www.aist.go.jp/GSJ/HomePage.html. JB-1b, a recently prepared GSJ standard material, does not have GSJ reference value for FeO, and we thus compared our data to the value obtained from Terashima et al. (1998). Analytical precision, evaluated based on the analytical reproducibility for each sample, was generally better than 0.5% (1) error. In some cases, however, the precision was poorer, approaching 1.0%. This poor precision is due to limitations related to sample size (up to 0.4 g) for JR- and JG- series samples with low FeO contents (as low as 1.5 wt.%). Larger amounts of sample would decrease the error, although the beaker size and the amounts of reagents would need to be increased. All of our data, except that for JB-1b, gave results different from the reference values, with these differences exceeding our analytical reproducibility. The differences are positive in all cases and range to ⫹15.5% of the reference values. In the case of seven samples (JB-3, JA-3, JR-3, JP-1, JGb-1, JG-2, and JG-3), the statistical test shows our results to be distinct from the reference data at the 99.5% confidence interval. Because the reference values of GSJ standards are obtained by averaging various published (and unpublished, in some cases) data analyzed by different methods, the difference between our data and the reference values, especially for the above seven samples, indicates that poor-quality data, which should be outliered, are included in the data source. The positive differences might indicate that such poor-quality data are obtained by a method with insufficient recovery of Fe2⫹ owing to its oxidation during measurement. One serious problem is the existence of spinel in JP-1, which is not only very resistant to HF, but also contains Fe2⫹ that can not be neglected. Because GSJ standard material is delivered as a powder, it is difficult to estimate the modal abundance of the spinel or the amount of Fe2⫹ in the spinel. We attempted to continue heating at 100°C on the hot plate for 4 days in the case of JP-1 to ensure dissolution of spinel, and even in this experiment we noted the existence of a small amount of undissolved spinel. We thus separated the spinel that remained in the beaker by hand-picking under a microscope and weighed it, and observed that ⬃0.5 mg remained when 110 mg of JP-1 was decomposed. Takazawa et al. (1996) reported that spinel in Horoman lherzolite, which is collected from the same location as JP-1, contains 15 wt.% FeO. Using this value and the amount of undissolved spinel in our experiment, the wt.% of FeO in JP-1 was corrected from 6.81% into 6.88%. The relatively higher uncertainty observed in our JP-1 analyses without correction might indicate the difference of the amount of undissolved spinel. We conclude that a correction using the amount of remaining spinel is necessary to obtain accurate values of FeO content when samples contain significant amounts of this mineral. Zircons in granodiorite and granite are also HF-resistant. However, the trace Fe contained in zircon is usually trivalent (Deer et al., 1992). Therefore, we consider our FeO value of granodiorite and granite to be reliable. We also analyzed the FeO concentration of 15 GSJ samples based on the method described in Wilson (1955). Details of this experiment are given in the Appendix. Table 3 shows that our
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Table 3. FeO (wt.%) of 15 GSJ standard rocks determined by Wilson’s method (average of duplicate analyses), and their relative different percentage from our value. Sample JB1-b JB-2 JB-3 JA-2 JA-3 JR-1 JR-2 JR-3 JP-1 JGb-1 JGb-2 JG-1a JG-2 JG-3 JH-1
Wilson’s
This study
Diff. %
5.04 9.91 7.94 3.74 4.81 0.42 0.35 1.85 6.34 9.48 5.12 1.35 0.56 1.93 7.70
5.15 10.04 8.00 3.75 4.99 0.510 0.448 1.99 6.81 10.89 5.54 1.40 0.655 2.016 8.20
⫺2.2 ⫺1.3 ⫺0.8 ⫺0.2 ⫺3.7 ⫺19 ⫺22 ⫺7.1 ⫺6.9 ⫺13 ⫺7.6 ⫺3.8 ⫺14 ⫺4.2 ⫺6.0
results using “Wilson’s” method for the standards show FeO abundances that are 0.2 to 22% lower than those that are obtained by our “original” procedure. Two major reasons are considered to explain such discrepancies between our method and that of Wilson (1955). The first one is the atmospheric oxidation during sample decomposition. Because no H2SO4 is included in the sample decomposing solution in Wilson’s method, it would cause an imperfect reaction (to the right side in reaction (1)), leaving Fe2⫹ unreacted with V5⫹ in the solution. Extraordinary large differences observed in JR-1 and JR-2 are, therefore, attributed to long-term oxidation during overnight sample decomposition. Because these two samples are rhyolitic obsidians, they were instantaneously decomposed by HF attack, and were subsequently affected by atmospheric oxidation for a relatively longer time compared to the other samples. In JG-2, the main source of Fe2⫹ is mica, which also decomposed rapidly causing a relatively longer oxidation period, resulting in –14% difference. The second possibility is the insufficient decomposition of rock samples, which also causes lower recovery of Fe2⫹, because the sample solution is not heated during decomposition in Wilson’s method. Actually, even in zircon-free samples, undissolved minerals were found in JP-1, JGb-1, JGb-2, and JH-1 after overnight decomposition. Such undissolved minerals were identified as spinel for JP-1 and orthopyroxene for JGb-1, JGb-2, and JH-1, respectively, using a SEM-EDX at PML (Horiba EMAX-7000 energy-dispersive X-ray spectrometer assembled into a Hitachi S-3100H scanning electron microscope). For these four samples, we also investigated rock powders before their decomposition using SEM-EDX, and found the existence of several Fe2⫹-containing minerals besides spinel and orthopyroxene (olivine for JP-1; clinopyroxene, hornblende, magnetite, and ilmenite for JGb-1; clinopyroxene and ilmenite for JGb-1; olivine, clinopyroxene, and hornblende for JH-1). We conclude that spinel and orthopyroxene are significantly acid-resistant compared to other rock-forming minerals, and preferentially remain when the sample solution is not heated. These observations clearly explain the large negative differences of these samples analyzed by Wilson’s method relative to results obtained by our method. Also, the lower FeO values of GSJ reference data might contain
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⫺d[Fe2⫹]/dt ⫽ k⬘[Fe2⫹]a
(3⬘)
Table 3 shows the rate constants for Eqn. 3⬘ supposing zero-, first-, and second-order reactions, which were calculated from the observed long-term variation of [Fe2⫹]. Only the zero-order reaction yielded similar rate constants irrespective of the initial concentration of Fe2⫹, indicating that [Fe2⫹] itself does not kinetically control the oxidation of Fe2⫹ in our experiments. Using the average of the k0 values in Table 4, the variation of [Fe2⫹] is expressed as [Fe2⫹]t ⫽ [Fe2⫹]0 ⫺ 0.00005 䡠 t (mmol/g)
(4)
where [Fe2⫹]0 represents the initial concentration of Fe2⫹ and t is time (in hours), respectively. In our analyses of silicate samples, [Fe2⫹] is determined once a day before the sample analyses. In some cases, analytical time within a single day exceeds a few hours, resulting in significant decrease in [Fe2⫹] and necessitating corrections using Eqn. 4. Fig. 2. Long-term variations of the concentrations of Fe2⫹ (}) and V5⫹ (f) solutions. Each error is determined from a standard deviation out of more than 5 times’ analyses.
analyses with insufficient sample decomposition and, therefore, sample digestion by HF with heating is indispensable for precise analyses of FeO in silicate rocks. 3.3. Fluctuation of the Concentration of Reagents Long-term stability of the concentrations of V5⫹ in 5 M H2SO4 and Fe2⫹ in 1 M H2SO4 solutions is important because these concentrations directly affect the concentration of FeO determined by our method. Although the concentration of V5⫹ solution was constant within analytical error over 10 days, the Fe2⫹ concentration showed a long-term decrease exceeding analytical error, and this decrease was clearly due to oxidation of Fe2⫹ by oxygen dissolved in the solution (Fig. 2). To determine the oxidation rate of the solution, we prepared several solutions of Fe2⫹ in 1 M H2SO4 with different concentrations of Fe2⫹, and examined their long-term variations. The Fe2⫹ oxidation reaction kinetics can be explained by the following equation: ⫺d[Fe2⫹]/dt ⫽ k[Fe2⫹]a 䡠 [O2]b 2⫹
3.4. Effect of the Amount of Indicator In the titration of K2Cr2O7 using sodium diphenylamine-4sulfonate as an indicator, surplus Cr6⫹ exceeding the equivalence point is required for the oxidation of the indicator itself. Thus, the amount of Cr6⫹ consumed up to the end point, which is different from the equivalence point, results in a positive error in the determination of FeO concentration. To evaluate the effect of the indicator, we have titrated 0.1 g of solid ammonium iron(II) sulfate hexahydrate while changing the amount of the indicator. As shown in Figure 3, the surplus Cr6⫹ can be ignored only when one drop of the indicator solution, which contains ⬃0.2 mol of sodium diphenylamine-4-sulfonate, is added. Otherwise progressively larger errors are produced as the amount of indicator increases. Furthermore, too much addition of the indicator would result in the side reaction, namely, the formation of green insoluble material by a reaction between the indicator in reduced and oxidized forms. This side reaction also consumes excess Cr6⫹, potentially leading to erroneous results. Because the above experiment was carried out with a total of 40 mL of titrated solution, two drops of the indicator would be permissible without correction if the whole solution is doubled.
(3)
where [Fe ] and [O2] are their concentrations in the solution, k is the rate constant, and a and b are the order of reaction, respectively. Provided that oxygen dissolved in the solution is always at equilibrium with air, [O2] can be considered constant. Equation 3 then becomes more simple:
3.4.1. Effect of the amount of HF We also found that coexisting HF in the solution interferes with the titration of Fe2⫹. As was described in section 3.4, we have examined the titration of 0.1 g of solid ammonium iron(II) sulfate hexahydrate while adding varying amounts of HF. As
Table 4. Rate constants (k) based on the assumptions of zero-, first-, and second-order reactions with various concentration of initial Fe2⫹. [Fe2⫹]0 mmol/g
k0 mmol 䡠 g⫺1 䡠 day⫺1
k1 day⫺1
k2 mmol⫺1 䡠 g 䡠 day⫺1
0.5664 0.4144 0.3346 0.3171
0.00127 0.00121 0.00123 0.00122
0.00227 0.00293 0.00374 0.00391
0.00407 0.00713 0.01133 50.01249
Precise determination of Fe2⫹ in silicate rocks
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or 101.427 (Arthur et al., 1999). HF dissociates weakly in the solution as: HF ⫽ H⫹ ⫹ F⫺
K2
(6)
where K2 is the acid dissociation constant, which is 10⫺3.17 (Ho¨ gfeldt, 1979). Combining these two equations yields: K1 䡠 K2 ⫽
[FeF⫹] [H⫹][F⫺] [FeF⫹][H⫹] 䡠 ⫽ [Fe2⫹][F⫺] [HF] [Fe2⫹][HF]
(7)
where [ ] represents molar concentrations. In our experimental condition, [H⫹] is controlled by the concentration of H2SO4, and that was 1.8 mol/L. Thus, [Fe2⫹] ⫺ [FeF⫹] ⫽ 1 ⫺ K 1 䡠 K 2 䡠 [HF]/1.8 [Fe2⫹]
Fig. 3. Relationship between recovery yield of Fe2⫹ and the amount of indicator varying from 0.2⬃1.1 mol, which corresponded to 1⬃5 drops of the indicator solution. The total amount of titrated solution is fixed as 40 mL.
shown in Figure 4, the recovery yield of Fe2⫹ shows a negative correlation with the amount of HF. This indicates that the existence of HF interferes with the redox reaction between Fe2⫹ and Cr6⫹ by forming an inert complex with Fe2⫹. The cation complex FeF⫹ is one cause of the problem. The formation of FeF⫹ is given as: Fe2⫹ ⫹ F⫺ ⫽ FeF⫹
K1 0.83
where K1, the equilibrium constant, is 10
(Ho¨ gfeldt, 1979)
(8)
The left side of this equation can be approximated by the recovery of Fe2⫹ titrated if [FeF⫹] is inert. Dashed and bold lines in Figure 4 are the theoretical lines of recovery yield calculated from Eqn. 8, when K1 equals 100.83 and 101.427, respectively. Our data broadly match the behavior of the theoretical lines, supporting the existence of inert [FeF⫹] in the solution. Although some data fall off these lines, those data can be explained by complicated factors involving coexisting species such as Cr6⫹, Cr3⫹, and Fe3⫹ in the solution. The most effective way to solve the problem of the formation of FeF⫹ is to use boric acid as a masking agent. We found that the use of 10 mL of saturated boric acid/1 mL of coexisting HF was effective in preventing the formation of inert complexes (Fig. 5). Overuse of boric acid, exceeding the above rate, did not interfere with the titration. However, boric acid did not work as a masking agent when HF attacked Fe2⫹ before the addition of the boric acid. Therefore, in the actual determination procedure for natural samples, 0.3 mmol/g Fe2⫹ solution is first weighed into the Pyrex beaker, and then saturated boric acid is added before the addition of the decomposed sample containing HF. 3.5. Detection Limit and Blank
Fig. 4. Relationship between recovery yield of Fe2⫹ and the amount of coexisting HF. Two lines indicate theoretical recovery of Fe2⫹ when K1 is 100.83 (dashed) or 101.427 (bold).
Detection limit of the method is largely attributed to the uncertainty of the amount of oxidation-reduction reagents actually used, namely, Fe2⫹, V5⫹, and K2Cr2O7 solutions. This can be estimated from the propagation of errors accompanied with each reagent, and it generally corresponds to ⬃0.00045 mmol of Fe2⫹ in this study. This means that Fe2⫹ contained in the starting material can not be distinguished from the error of reagents when it is less than 0.00045 mmol, which we define as the detection limit of our method. We carried out blank runs by determining the amount of Fe2⫹ in the procedure described in section 2.2.1 without adding any starting material. 1.5 mL of HF–H2SO4 mix solution was used, as in the case for decomposition of 0.2 g of silicate rock powder. We totally tested 3 sets of blank runs, in which each run comprised of a set of 5 blank measurements. All the blanks were less than ⫾0.00030 mmol of Fe2⫹, indicating that the blank does not exceed the detection limit. Therefore, any element that interferes with the series of redox reactions in our chemistry is absent in the reagents used in this study.
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Fe2⫹. To test the validity of such a system, we measured JG-1a that showed the worst reproducibility in Table 2 together with procedural blank. Four analyses gave 1.406 ⫾ 0.009 (1) wt.% of FeO, which is not only the same value as Table 2, but also resulted in better reproducibility, which improved from 0.92% to 0.63%. Because the average of triplicate blank runs was less than the detection limit (⫺0.00017 mmol of Fe2⫹), it is concluded that the dilution of the concentrations of Fe2⫹, V5⫹, and K2Cr2O7 solutions is effective not only to decrease the minimum range of the analysis, but also to improve the precision for samples with extremely low FeO contents. 4. CONCLUDING REMARKS
Fig. 5. The effect of saturated boric acid for masking HF. Open circles indicate 100% recovery of Fe2⫹ in the titration with coexisting HF, and crosses show the failure of the titration. 10 mL of saturated boric acid/1 mL of coexisting HF is found to be enough.
3.6. Modification of the Method In this study, we used less than 0.15 mmol of Fe2⫹ in starting materials while decomposing less than 0.4 g of rock samples. Providing 0.1 g of the sample for decomposition avoids sample heterogeneity, the measurable maximum FeO concentration in a sample with this method is calculated to be 10.8 wt.%. Because some natural samples (e.g., ferrobasalt) contain ⬃20 wt.% FeO, we evaluated the measurement of 0.22 g of JB-2 (containing ⬃0.30 mmol of Fe2⫹) when increasing the V5⫹ solution as 1.17 g (⬃0.35 mmol). Although the final solution became more blue during titration (due to larger amount of V4⫹) than the normal case, we obtained 10.07 ⫾ 0.01 (1) wt.% of FeO as a result of five analyses, which is identical to that of routine analyses in Table 2 when errors are considered. Therefore, our method can be applied to samples containing ⬎20 wt.% FeO by simply increasing the amount of V5⫹ solution. On the other hand, the minimum range of our method is restricted by the detection limit or blank of the method. Assuming that the blank of the whole procedure has a composition of the detection limit (0.00045 mmol), at least 0.09 mmol of Fe2⫹ is required as the starting material if one would like to suppress the effect of the blank to less than 0.5% of Fe2⫹ contained in the sample used. This corresponds to 1.6 wt.% of FeO when 0.4 g of sample is used, and the effect of the blank increases as the FeO content in the sample decreases. One simple way to further improve the minimum range is to enhance the resolution of the titration by diluting the concentrations of Fe2⫹, V5⫹, and K2Cr2O7 solutions. When all the concentrations of these solutions were diluted to half of the usual case (0.15 mmol/g for Fe2⫹ and V5⫹ solutions, and 0.02 N for K2Cr2O7 solution), the detection limit estimated by the same way as described in section 3.6 is ⬃0.00021 mmol of
We have established an accurate and precise method for the determination of Fe2⫹ in silicate rocks with analytical error better than 0.5% (1). In this method, the oxidation of Fe2⫹ can be avoided by keeping the sample solution at higher acidity (5 M H2SO4). Sample decomposition in a relatively small beaker (7 mL) with a small amount of sample solution is very effective for preventing atmospheric oxidation. Fe2⫹ concentrations in 15 GSJ standard silicate rock powders were determined in this study, yielding values higher than those previously reported (except for one sample), indicating that the GSJ reference values were determined with lower recovery of Fe2⫹ owing to its oxidation or insufficient sample decomposition during measurement. In addition to the necessity of avoiding oxidation of Fe2⫹, we have demonstrated several important factors that determine the accuracy of our method: (1) correction of the long-term variation of the concentration of Fe2⫹ solution; (2) the amount of indicator; (3) the use of boric acid as a masking agent to protect Fe2⫹ from forming inert FeF⫹ complexes during titration. Acknowledgments—We would like to thank N. Imai (Geological Survey of Japan) for offering GSJ standard rock powders. We are grateful to G. E. Bebout and M. J. Walter for improving the English used in this paper. We also thank A. Makishima, H. Takei, N. Takeuchi, and all the other member of PML for their analytical support and useful discussion. We would like to express our appreciation to the Associate Editor, B. R. Frost and anonymous reviewers for their efforts to improve the quality of this paper. This research was supported by Grants-in-Aid for Scientific Research from the Ministry of Education, Science, and Culture, Japan (Monbusho) and the Japanese Society for the Promotion of Science (JSPS) to E.N. Associate editor: B. R. Frost APPENDIX Wilson’s Procedure 0.5 g of powdered sample and 0.1 g of solid ammonium metavanadate (special grade, Kanto Chemical) were weighed into a 100-mL polypropylene beaker, and 10 mL of 27 M HF was added. The mixture was kept at room temperature overnight for complete sample decomposition. After adding 30 mL of 5 M H2SO4, the contents were transferred into a 500-mL glass beaker with 250 mL of saturated boric acid. After adding 7 drops of 0.2% sodium diphenylamine-4-sulfonate, the solution was titrated by 0.07 mol/L Fe2⫹ solution, which was conditioned by dissolving 27.5 g of ammonium iron(II) sulfate hexahydrate into 1 L of 0.5 M H2SO4. The solution containing only ammonium metavanadate was also treated the same as the sample and titrated to determine the actual concentration of the Fe2⫹ solution.
Precise determination of Fe2⫹ in silicate rocks REFERENCES Amonette J. E. and Scott A. D. (1991) Determination of ferrous iron in non-refractory silicate minerals. 1. An improved semi-micro oxidimetric method. Chem. Geol. 92, 329 –338. Amonette J. E., Khan F. A., Gan H., Stucki J. W., and Scott A. D. (1999) Comparison of oxidimetric, spectrophotometric and Mo¨ ssbauer-spectroscopic methods for determination of Fe(II) in nonrefractory minerals. Proceedings Eleventh International Clay Conference, Ottawa, Ontario, Canada, 1997, pp. 277–286. Arthur R. C., Sasamoto H., Shibata M., Yui M., and Neyama A. (1999) Development of Thermodynamic Databases for Geochemical Calculation. Japan Nuclear Cycle Development Institute. Beattie P., Ford C., and Russell D. (1991) Partition coefficients for olivine-melt and orthopyroxene-melt system. Contrib. Mineral. Petrol. 109, 212–224. Beattie P. (1993) Olivine-melt and orthopyroxene-melt equilibria. Contrib. Mineral. Petrol. 115, 103–111. Bock R. (1979) A Handbook of Decomposition Methods in Analytical Chemistry. International Textbook Co. Carmichael I. S. E. and Ghiorso M. S. (1986) Oxidation-reduction relations in basic magma: A case for homogeneous equilibria. Earth Planet. Sci. Lett. 78, 200 –210. Christie D. M., Carmichael I. S. E., and Langmuir C. H. (1986) Oxidation states of mid-ocean ridge basalt glasses. Earth Planet. Sci. Lett. 79, 397– 411. Deer W. A., Howie R. A., and Zussman J. (1992) An Introduction to the Rock-forming Minerals. Longman House. Ho¨ gfeldt E. (1979) Stability Constants of Metal–Ion Complexes: Part A: Inorganic Ligands. Pergamon Press. Imai N., Terashima S., Itoh S., and Ando A. (1995) 1994 compilation values for GSJ reference samples, “Igneous rock series.” Geochem. J. 29, 91–95. Imai N., Terashima S., Itoh S., and Ando A. (1999) 1998 compilation of analytical data for five GSJ geochemical reference samples: The “Instrumental Analysis Series.” Geostand. Newslett. 23, 223–250. Ishibashi M. and Kusaka Y. (1950) A new method for the determination of ferrous iron in silicates (in Japanese). Nihon Kagaku Zasshi 71, 160 –162. Kilinc A., Carmichael I. S. E., Rivers M. L., and Sack R. O. (1983) The ferric–ferrous ratio of natural silicate liquids equilibrated in air. Contrib. Mineral. Petrol. 83, 136 –140. Koch C. B., Mørup S., Madsen M. B., and Vistisen L. (1995) Ironcontaining weathering products of basalt in a cold, dry climate. Chem. Geol. 122, 109 –119.
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Kress V. C. and Carmichael I. S. E. (1988) Stoichiometry of the oxidation reaction in silicate melts. American Mineral. 73, 1267– 1274. Kress V. C. and Carmichael I. S. E. (1991) The compressibility of silicate liquids containing Fe2O3 and the effect of composition, temperature, oxygen fugacity and pressure on their redox states. Contrib. Mineral. Petrol. 108, 82–92. Langmyhr F. J. and Kringstad K. (1966) An investigation of the composition of the precipitates formed by the decomposition of silicate rocks in 38 – 40% hydrofluoric acid. Anal. Chim. Acta 35, 131–135. Peters A. (1968) Ein neues Verfahren zur Bestimmung von Eisen(II) oxid in Mineralen und Gesteinen. Neues Jahrb. Mineral. Monatsh. 3/4, 119 –125. Potts P. J. (1987) A Handbook of Silicate Rock Analysis. Blackie. Pratt J. P. (1894). On the determination of ferrous iron in silicates. Am. J. Sci. 48, 149 –151. Roeder P. L. and Emslie R. F. (1970) Olivine–liquid equilibrium. Contrib. Mineral. Petrol. 29, 275–289. Sack R. O., Carmichael I. S. E., Rivers M., and Ghiorso M. S. (1980) Ferric–ferrous equilibria in natural silicate liquids at 1 bar. Contrib. Mineral. Petrol. 75, 369 –376. Takazawa E., Frey F., Shimizu N., and Obata M. (1996) Evolution of the Horoman Peridotite (Hokkaido, Japan): Implication from pyroxene compositions. Chem. Geol. 134, 3–26. Takei H., Yokoyama T., Makishima A. and Nakamura E. (2001). Formation and suppression of AlF3 during HF digestion of rock samples in Teflon bomb for precise trace element analyses by ICP-MS and ID-TIMS. Proc. Japan Acad. 77, 13–17. Terashima S., Taniguchi M., Mikoshiba M., and Imai N. (1998) Preparation of two new GSJ geochemical reference materials: JB-1b and coal fly ash JCFA-1. Geostand. Newslett. 22, 113–117. Thornber C. R., Roeder P. L., and Foster J. R. (1980) The effect of composition on the ferric–ferrous ratio in basaltic liquids at atmospheric pressure. Geochim. Cosmochim. Acta 44, 525–532. Whipple E. R. (1974) A study of Wilson’s determination of ferrous iron in silicates. Chem. Geol. 14, 223–238. Wilson A. D. (1955) Determination of ferrous iron in rocks and minerals. Bull. Geol. Surv. G. B. 9, 56 –58. Yokoyama T., Makishima A., and Nakamura E. (1999) Evaluation of the coprecipitation of incompatible trace elements with fluoride during silicate rock dissolution by acid digestion. Chem. Geol. 157, 175–187.
Pergamon
PII S0016-7037(01)00809-2
Precise determination of ferrous iron in silicate rocks TETSUYA YOKOYAMA and EIZO NAKAMURA The Pheasant Memorial Laboratory for Geochemistry and Cosmochemistry, Institute for Study of the Earth’s Interior, Okayama University at Misasa, Misasa, Tottori-ken 682-0193, Japan (Received October 20, 2000; accepted in revised form August 22, 2001)
Abstract—We have developed a highly precise method for the determination of ferrous iron (Fe2⫹) in silicate rocks. Our new method is based on Wilson’s procedure (1955) in which surplus V5⫹ is used to oxidize Fe2⫹ into Fe3⫹ while equivalently reducing V5⫹ into V4⫹. Because V4⫹ is more resistant to atmospheric oxidation than Fe2⫹, Fe2⫹ in the sample can be determined by measuring unreacted V5⫹ by adding excess Fe2⫹ after sample decomposition and then titrating the unreacted Fe2⫹ with Cr6⫹. With our method, which involves conditioning the sample solution with 5 M H2SO4 in a relatively small beaker (7 mL), the oxidation of Fe2⫹ or V4⫹ that leads to erroneous results can be completely avoided, even in 100-h sample decompositions at 100°C. We have measured the concentration of FeO in 15 standard silicate rock powders provided by the Geological Survey of Japan (GSJ). Analytical reproducibility was better than 0.5% (1) for all but those samples that had small amounts of Fe2⫹ (⬍1.5 wt.% of FeO). Fourteen of these samples gave FeO contents significantly higher than the GSJ reference values. This likely indicates that the GSJ reference values, obtained by compiling previously published data, contain a large number of poor-quality data obtained by methods with lower recovery of Fe2⫹ caused by oxidation or insufficient sample decomposition during analyses. To achieve accurate determinations of Fe2⫹ in our method, several factors besides the oxidation must be considered, including: (1) long-term variations in the concentration of Fe2⫹ solution must be corrected; (2) excess use of the indicator must be avoided; and (3) the formation of inert FeF⫹ complex must be avoided during titration when using boric acid as a masking agent. Copyright © 2002 Elsevier Science Ltd oxidation of Fe2⫹. In our preliminary analyses of basaltic, andesitic, and rhyolitic standard rock powders using Pratt’s method, we obtained relative values 2 ⬃ 5% lower than the reference values, although the accuracy of the reference values themselves is difficult to assess. One of the most effective ways to achieve accurate determinations of Fe2⫹ is Wilson’s method (Wilson, 1955) or Peters’ method (Peters, 1968), in which excess V5⫹ is added to the sample before decomposition. This method is based on the following reversible reaction:
1. INTRODUCTION
Determination of ferrous iron (Fe2⫹) and Fe2⫹/Fe3⫹ ratio in silicate rocks is important in geological, geochemical, and petrological studies for evaluation of weathering (Koch et al., 1995), determination of oxygen fugacity in magma (Sack et al., 1980; Thornber et al., 1980; Kilinc et al., 1983; Carmichael and Ghiorso, 1986; Christie et al., 1986; Kress and Carmichael, 1988; Kress and Carmichael, 1991), and understanding chemical equilibrium between minerals and silicate melt (Roeder and Emslie, 1970; Beattie et al., 1991; Beattie, 1993). Recent analytical progress using x-ray fluorescence (XRF) methods enables highly precise determination of total iron in silicate rock samples with analytical error less than 0.5% (1) (relative standard deviation; e.g., Takei and Nakamura, in preparation). Determination of Fe2⫹ concentrations with the same precision as that for XRF analyses of total iron would make it possible to more precisely address the above petrological issues. Several analytical techniques for the determination of Fe2⫹ in silicate rocks are reviewed elsewhere (Bock, 1979; Potts, 1987; Amonette et al., 1999). Although high-precision analyses for Fe2⫹ are relatively easily obtained, the accuracy of these analyses is more questionable. Fe2⫹ can easily be oxidized during analytical procedures, resulting in erroneously low values. In the classical method of Pratt (Pratt, 1894), which is commonly used in many laboratories, the sample is decomposed by mixtures of HF and H2SO4 in a covered platinum crucible that excludes atmospheric oxygen. This relatively difficult technique often yields low concentrations owing to the
* Author to whom correspondence ([email protected]).
should
be
Fe2⫹ ⫹ V5⫹ ⫽ Fe3⫹ ⫹ V4⫹
(1)
which was first found to be effective in determination of Fe2⫹ by Ishibashi and Kusaka (1950). This reaction proceeds to the right side in the condition of strong acidity, and the amount of the V4⫹ reduced from V5⫹ is then the same as the amount of Fe2⫹ in the sample. Because V4⫹ is extremely resistant to oxidation and does not back-react to V5⫹ during the analytical procedure, Fe2⫹ concentration in the sample can be determined by measuring the unreacted V5⫹ by adding excess Fe2⫹ after sample decomposition and then titrating the unreacted Fe2⫹ with Cr6⫹. Whipple (1974) improved Wilson’s method and evaluated its accuracy and precision by determining the recovery of Fe2⫹ using ammonium iron(II) sulfate hexahydrate as a starting material. He achieved determinations of 7 mg of Fe2⫹ with precision of ⬃0.3%. In his method, however, a 1% correction was required due to small amounts of oxidation of Fe2⫹ or V4⫹ during decomposition, even though he deoxidized the reagents by N2 purging for 4 h before analysis. Whipple also attempted to decompose the sample in a closed N2 box, but this resulted in up to 0.5% loss of Fe2⫹.
addressed 1085
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Amonette and Scott (1991) developed a method that improved upon Peters’ method, and achieved better determinations of Fe2⫹ (0.2 to 0.9%, 1) in 5 geochemical reference samples and 3 mica samples. However, they recommended that sample decomposition must be carried out at 60°C within 24 h, because they found that higher temperature and longer decompositions cause loss of Fe2⫹ due to oxidation. These conditions are, however, insufficient to completely decompose some minerals containing Fe2⫹ such as pyroxene, garnet and spinel, these being more acid-resistant than mica. In this study, we have reevaluated the methods of Wilson (1955) and Whipple (1974) and developed more routine, more accurate methods (better than 0.5% precision, 1) while overcoming some problems that occurred in previous studies. We provide new determinations of the FeO concentration (wt.%) in 15 standard silicate rock powders. 2. EXPERIMENTAL SECTION
All of the experiments in this study were carried out at the Pheasant Memorial Laboratory (PML), Institute for Study of the Earth’s Interior in Misasa.
2.1.4. Indicator 0.2 g of sodium diphenylamine-4-sulfonate (first grade, Kanto Chemical) was dissolved in water and diluted to 100 mL resulting in 0.2% (w/v). This solution was stored in a 125-mL polyethylene dropping bottle. The titration is complete when the color changes from gray to violet. 2.1.5. H2SO4–H3PO4 mix solution 400 mL of 18 M H2SO4 was slowly added to 400 mL of water in a draft chamber while cooling in a water bath. After cooling to room temperature, 200 mL of 15 M H3PO4 was carefully mixed. 2.1.6. HF-H2SO4 mix solution 33.3 mL of 18 M H2SO4 was slowly added to 86.7 mL of 27 M HF in a draft chamber while cooling in a water bath. After cooling to room temperature, it was stored into a 125-mL polyethylene dropping bottle. This produces a 5 M concentration of H2SO4 in this solution. 2.1.7. 0.3 mmol/g Fe2⫹ in 1 M H2SO4 solution
2.1. Reagents and Apparatus 2.1.1. Water Water was deionized and purified using a mixed-bed resin and filters (WL21P, Yamato, Japan). Typical specific resistance of the water was better than 16.7 M ⍀䡠cm. 2.1.2. Hydrofluoric, sulfuric, and phosphoric acids 27 M HF (special grade, Wako, Japan), 18 M H2SO4 (trace analysis grade, Kanto Chemical Co., Japan), and 15 M H3PO4 (special grade, Wako) were used without any further purification. These acids are highly toxic and corrosive. They irritate the skin and mucous membranes of the body. Thus, to avoid inhalation, ingestion, and contact with the skin or eyes, one should wear protective gloves, glasses, and clothing while using a chemical fume. 2.1.3. Titrant 1.96 g of potassium dichromate reference material for volumetric analysis (Kanto Chemical, 99.98 ⫾ 0.01%) was dissolved in water and diluted to 1000 mL using measuring flask, resulting in 0.04 N K2Cr2O7 solution. We used this solution as a primary standard throughout this study. The solution was stored in a tightly sealed 1000-mL glass bottle. Although the concentration of the solution did not change after 1 month, it increased by 0.3% after 7-month storage (probably due to the evaporation of water), based on a calibration by ammonium iron(II) sulfate hexahydrate. Therefore, it is recommended to use solid ammonium iron(II) sulfate hexahydrate as a secondary standard by determining its absolute Fe2⫹ concentration with newly conditioned K2Cr2O7 solution, and calibrate the titrant when it is stored over a few months.
13.8 g of ammonium iron(II) sulfate hexahydrate (analytical grade, Merck, Germany) was dissolved into 100 mL of 1 M H2SO4 (⬃103.5 g) and stored in a 125-mL polyethylene dropping bottle. The molar concentration of this solution was determined as follows: Approximately 0.7 g of the solution was weighed in a Pyrex glass beaker, followed by addition of 5 mL of H2SO4–H3PO4 solution. The resulting mixture was then diluted to 40 mL by the addition of water. One drop of the indicator was added to the beaker and the solution was titrated with 0.04 N K2Cr2O7. The mean value of five successive titrations was used as the molar concentration. 2.1.8. 0.3 mmol/g V5⫹ in 5 M H2SO4 solution 4.69 g of sodium metavanadate (special grade, Kanto Chemical) was dissolved into 100 mL of 5 M H2SO4 (⬃123.5 g) and stored in a 125-mL polyethylene dropping bottle. The molar concentration of this solution was determined as follows: approximately 0.7 g of the solution and 1.4 g of Fe2⫹ solution were weighed in a Pyrex glass beaker. The subsequent procedure was the same as that described in section 2.1.7. 2.1.9. Saturated boric acid 100 g of special grade boric acid (Kanto Chemical) was dissolved in 500 mL of hot water (80°C). After cooling to room temperature, the supernatant was stored and used in the analytical procedures. 2.1.10. Burette system A piston burette (Titronic 96, Schott, Germany) was used for titration in the course of this study. This system has a 20-mL glass cylinder and enables spewing of the solution in 0.01-mL aliquots. The calibration of the dosing unit was carried out by repeatedly spewing and weighing 5.00, 10.00, and 15.00 mL of water, and the accuracy of the volume was within 0.15%
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relative to the digitally displayed volume. The precision of the system was less than 0.05% when 5.00 mL of water was spewed, and it decreased as the volume increased. 2.1.11. Weighing system A microbalance with 5 decimal places (R200D, Sartorius, Germany) was used for weighing samples. Typical uncertainty of this balance was ⫾0.00010 g (1). It is important to eliminate static electricity of the sample before weighing, and a static eliminator (Endstat II, Otsubo Electronics, Japan) was used for this purpose. 2.2. Procedure 2.2.1. Determination of Fe(II) in silicate rocks Fifteen GSJ (Geological Survey of Japan) standard silicate rock powders were analyzed in this study, including JP-1 (peridotite); JB-1b, JB-2, and JB-3 (basalts); JA-2 and JA-3 (andesites); JR-1, JR-2, and JR-3 (rhyolites); JH-1 (hornblendite); JGb-1 and JGb-2 (gabbros); JG-1a and JG-3 (granodiorites); and JG-2 (granite). Elemental concentrations of these rock powders are given in Terashima et al. (1998) for JB-1b, Imai et al. (1999) for JR-3, JGb-2, and JH-1, and Imai et al. (1995) for the rest of the samples. All rock powders were analyzed in the stock in which they were provided by GSJ, without any further pulverization. Starting material containing less than 0.15 mmol of Fe2⫹ was weighed into a 7-mL clean Teflon beaker. The amount of sample used for the determinations did not exceed 0.4 g. Approximately 0.7 g of V5⫹ solution was then added to the beaker and carefully weighed. The HF–H2SO4 solution was subsequently added to the sample at a proportion of 1.5 mL/0.2 g of silicate sample. The beaker was tightly sealed and agitated in an ultrasonic bath for 20 min. The beaker was then heated overnight on a hot plate at 100°C. All the rock powder was decomposed under such a condition, but a certain proportion of spinel and zircon remained undissolved, as discussed later. After sample decomposition, ⬃1.0 g of Fe2⫹ solution was carefully weighed into a clean Pyrex beaker and 5 mL of the H2SO4–H3PO4 solution was added. Afterwards, saturated boric acid was added to the Pyrex beaker at a proportion of 10 mL/1.5 mL of the HF–H2SO4 solution that was used for sample decomposition. All of the decomposed material was moved into the Pyrex beaker by carefully washing the inside of the Teflon beaker with H2O. The recovered solution was diluted to 40 mL with H2O, then one drop of indicator was added before titration with 0.04 N K2Cr2O7. The concentration of Fe2⫹ was calculated by the following equation: FeO (wt.%) ⫽ (VCr 䡠 CCr ⫹ WV 䡠 CV ⫺ WFe 䡠 CFe) 䡠 10⫺3 Wsamp ⫻ 71.85 ⫻ 100
(2)
where VCr was the amount of titrant (mL); WV, WFe, and Wsamp were weights of the V5⫹ solution, Fe2⫹ solution, and sample, respectively (g); CCr was the normality of the titrant
Fig. 1. Relationship between recovery yield of Fe2⫹ and heating time of the sample with 2 M (}), 3 M (f), and 5 M (F) H2SO4 conditions.
(N); and CV and CFe were the molar concentrations of V5⫹ and Fe2⫹ solutions, respectively (in mmol/g). 2.2.2. Recovery yield measurement of Fe2⫹ Approximately 0.06 g of solid ammonium sulfate hexahydrate (⬃0.15 mmol of Fe2⫹) was weighed into a 7-mL clean Teflon beaker, after which ⬃0.7 g of V5⫹ solution was added. 1.0 mL of HF, together with variable amounts of H2SO4, were then added to examine the effect of varying acidity. The sample was treated the same as described above except the heating times varied from 3 to 100 h to examine the effect of the side reaction during the procedure. The amount of Fe2⫹ in the Teflon beaker was used to calculate recovery yield, normalizing to the initial concentration of Fe2⫹ in the ammonium iron(II) sulfate hexahydrate (2.550 mmol/g). 3. RESULTS AND DISCUSSION
3.1. Performance of the Method The most plausible source of analytical error in the determination of Fe2⫹ by our method is from side reactions such as the oxidation of Fe2⫹ during sample decomposition. To evaluate the accuracy of our method, we repeatedly measured the recovery yield of Fe2⫹ throughout the chemical procedure by using ammonium iron(II) sulfate hexahydrate as a starting material while varying both the acidity and the heating time. As shown in Figure 1, the recovery yield is strongly affected by both of these variables. 100% of the Fe2⫹ was recovered within 100 h heating when the molar concentration of H2SO4 was 5 M. The 3 M condition was effective for 100% recovery unless the beaker was heated for more than 42 h. However, it was difficult to obtain correct values when the acidity was 2 M. These results can be explained by the fact that stronger acidity drives reaction (1) to the right side. In the condition of 5 M, reaction (1) must be kept completely to the right side, and all Fe2⫹ is oxidized
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20 mL V5⫹ solution in 4.5 M H2SO4 ⫹ 5 mL HF (corresponding to 3.6 M H2SO4 overall) at 60°C for 48 h. Conversely, we found no Fe2⫹ loss in 3 M H2SO4 for 42 h. This is probably due to the difference in the size of the decomposing system (i.e., 125-mL beaker with 25 mL of solution for Amonette, 7-mL beaker with 3⬃5 mL of solution for ours), reflecting a difference in the extent of atmospheric oxidation. The reduction of the decomposing size is achieved by reducing the volume of V5⫹ solution while increasing its concentration compared to the previous methods, and determining its amount by weighing. Table 1 shows the relationship between the standing time and the concentration of FeO in JB-3 determined by our method with a constant acidity in 5 M H2SO4. The nearly constant value of FeO suggests that the side reaction with oxygen at 5 M condition does not occur, even in analyses of silicate samples. Another possible problem that yields an incorrect FeO value in analyses of silicate samples is the coprecipitation of Fe2⫹ with insoluble fluorides associated with Mg, Ca, or Al, which form during sample decomposition with HF. Langmyhr and Kringstad (1966) investigated the formation of insoluble fluorides during HF digestion by artificially synthesizing them from various combination of elements, and noted the occurrence of Fe2⫹(Al, Fe3⫹)F5䡠3.7 H2O in the Al-Fe-HF system. However, this system does not contain Mg or Ca, which are commonly abundant in silicate rocks. No insoluble fluoride containing Fe was detected by X-ray diffraction analysis in the HF digestion of mafic rocks (Yokoyama et al., 1999) and felsic rocks (Takei et al., 2001). Instead, CaAlF5, CaMg2Al2F12, Na0.88Mg0.88Al1.12(F,OH)6䡠H2O, MgF2 and AlF3 were the dominant fluoride minerals. Thus, we conclude that the formation of insoluble fluorides does not affect our FeO determinations.
Table 1. FeO (wt.%) in JB-3 with changing heating time. Hours
FeO
3 3 9 15 15 18 22 22 26 26 26 42 96 96 Average 1
7.98 8.04 8.02 8.00 8.02 8.02 7.97 7.94 7.98 8.02 8.01 8.02 8.00 8.04 8.00 0.03
into Fe3⫹ by reducing an equivalent amount of V5⫹ into V4⫹. From our results, it can be concluded that V4⫹ is very resistant to oxidation in the 5 M condition even when dissolved oxygen exists in the system. In our method, therefore, predeoxidization of reagents such as V5⫹ solution and HF is not required. However, at weaker acidity, a small proportion of the Fe2⫹ remains unconsumed by reaction, resulting in lower recovery of Fe2⫹ owing to its oxidation at longer heating times. Our results explain the results of Whipple (1974) that demonstrated 1% loss of Fe2⫹ after 20 h of standing time. In his experiment, the whole-sample acidity was lower than that of the 2 M H2SO4 condition, because he used 4 mL of V5⫹ solution in 5 M H2SO4 with 6⬃8 mL of HF added. In our method, the loss of Fe2⫹ in the 2 M condition with 24-h standing was 3% (mean value), slightly larger than that of Whipple’s experiments. This is probably due to heating of our samples, which was not done in Whipple’s study. Even without heating, the oxidation of Fe2⫹ could not be avoided when the sample acidity was less than 2 M H2SO4. On the other hand, Amonette and Scott (1991) showed 2% loss of Fe2⫹ in a sample that was decomposed in a mixture of
3.2. Determination of Fe2ⴙ in the GSJ Standard Rock Powders Table 2 shows the concentration of FeO for 15 GSJ standard rock powders determined in this study. Also shown are the GSJ reference values obtained from Imai et al. (1995) and Imai et al.
Table 2. FeO (wt.%) of 15 GSJ standard rocks and their recommended or preferable (asterisked) values. Sample JB1-b JB-2 JB-3 JA-2 JA-3 JR-1 JR-2 JR-3 JP-1 JGb-1 JGb-2 JG-1a JG-2 JG-3 JH-1 a b
FeO (1)
1 (%)
n
Recomm. (⫾1)
n
Diff. %
5.15 ⫾ 0.01 10.04 ⫾ 0.03 8.00 ⫾ 0.03 3.75 ⫾ 0.01 4.99 ⫾ 0.02 0.510 ⫾ 0.004 0.448 ⫾ 0.004 1.99 ⫾ 0.01 6.81b ⫾ 0.06 10.89 ⫾ 0.03 5.54 ⫾ 0.02 1.40 ⫾ 0.01 0.655 ⫾ 0.005 2.016 ⫾ 0.002 8.20 ⫾ 0.01
0.23 0.34 0.35 0.22 0.49 0.80 0.90 0.38 0.90 0.30 0.43 0.92 0.79 0.09 0.09
5 5 14 5 5 4 5 4 4 4 4 4 6 4 4
5.16a 9.98 ⫾ 0.30 7.85 ⫾ 0.17 3.69 ⫾ 0.12 4.83 ⫾ 0.19 0.49 ⫾ 0.10 0.44 ⫾ 0.09 1.86 ⫾ 0.10 5.99 ⫾ 0.42 9.43 ⫾ 0.47 5.41 ⫾ 0.24 1.36 ⫾ 0.22 0.57 ⫾ 0.07 1.83 ⫾ 0.21 8.09* ⫾ 0.26
40 20 15 13 21 17 6 13 24 6 22 14 11 4
⫺0.2 0.6 2.0 1.7 3.4 4.0 1.7 7.0 13.6 15.5 2.4 2.9 14.9 10.2 1.4
After Terashima et al. (1998). Not recommended nor preferable value. The contribution of residual spinel is not included. See the text.
Precise determination of Fe2⫹ in silicate rocks
(1999), and the differences between our data and reference concentrations. Standard deviations of reference values are obtained from the database available at the GSJ internet website, http://www.aist.go.jp/GSJ/HomePage.html. JB-1b, a recently prepared GSJ standard material, does not have GSJ reference value for FeO, and we thus compared our data to the value obtained from Terashima et al. (1998). Analytical precision, evaluated based on the analytical reproducibility for each sample, was generally better than 0.5% (1) error. In some cases, however, the precision was poorer, approaching 1.0%. This poor precision is due to limitations related to sample size (up to 0.4 g) for JR- and JG- series samples with low FeO contents (as low as 1.5 wt.%). Larger amounts of sample would decrease the error, although the beaker size and the amounts of reagents would need to be increased. All of our data, except that for JB-1b, gave results different from the reference values, with these differences exceeding our analytical reproducibility. The differences are positive in all cases and range to ⫹15.5% of the reference values. In the case of seven samples (JB-3, JA-3, JR-3, JP-1, JGb-1, JG-2, and JG-3), the statistical test shows our results to be distinct from the reference data at the 99.5% confidence interval. Because the reference values of GSJ standards are obtained by averaging various published (and unpublished, in some cases) data analyzed by different methods, the difference between our data and the reference values, especially for the above seven samples, indicates that poor-quality data, which should be outliered, are included in the data source. The positive differences might indicate that such poor-quality data are obtained by a method with insufficient recovery of Fe2⫹ owing to its oxidation during measurement. One serious problem is the existence of spinel in JP-1, which is not only very resistant to HF, but also contains Fe2⫹ that can not be neglected. Because GSJ standard material is delivered as a powder, it is difficult to estimate the modal abundance of the spinel or the amount of Fe2⫹ in the spinel. We attempted to continue heating at 100°C on the hot plate for 4 days in the case of JP-1 to ensure dissolution of spinel, and even in this experiment we noted the existence of a small amount of undissolved spinel. We thus separated the spinel that remained in the beaker by hand-picking under a microscope and weighed it, and observed that ⬃0.5 mg remained when 110 mg of JP-1 was decomposed. Takazawa et al. (1996) reported that spinel in Horoman lherzolite, which is collected from the same location as JP-1, contains 15 wt.% FeO. Using this value and the amount of undissolved spinel in our experiment, the wt.% of FeO in JP-1 was corrected from 6.81% into 6.88%. The relatively higher uncertainty observed in our JP-1 analyses without correction might indicate the difference of the amount of undissolved spinel. We conclude that a correction using the amount of remaining spinel is necessary to obtain accurate values of FeO content when samples contain significant amounts of this mineral. Zircons in granodiorite and granite are also HF-resistant. However, the trace Fe contained in zircon is usually trivalent (Deer et al., 1992). Therefore, we consider our FeO value of granodiorite and granite to be reliable. We also analyzed the FeO concentration of 15 GSJ samples based on the method described in Wilson (1955). Details of this experiment are given in the Appendix. Table 3 shows that our
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Table 3. FeO (wt.%) of 15 GSJ standard rocks determined by Wilson’s method (average of duplicate analyses), and their relative different percentage from our value. Sample JB1-b JB-2 JB-3 JA-2 JA-3 JR-1 JR-2 JR-3 JP-1 JGb-1 JGb-2 JG-1a JG-2 JG-3 JH-1
Wilson’s
This study
Diff. %
5.04 9.91 7.94 3.74 4.81 0.42 0.35 1.85 6.34 9.48 5.12 1.35 0.56 1.93 7.70
5.15 10.04 8.00 3.75 4.99 0.510 0.448 1.99 6.81 10.89 5.54 1.40 0.655 2.016 8.20
⫺2.2 ⫺1.3 ⫺0.8 ⫺0.2 ⫺3.7 ⫺19 ⫺22 ⫺7.1 ⫺6.9 ⫺13 ⫺7.6 ⫺3.8 ⫺14 ⫺4.2 ⫺6.0
results using “Wilson’s” method for the standards show FeO abundances that are 0.2 to 22% lower than those that are obtained by our “original” procedure. Two major reasons are considered to explain such discrepancies between our method and that of Wilson (1955). The first one is the atmospheric oxidation during sample decomposition. Because no H2SO4 is included in the sample decomposing solution in Wilson’s method, it would cause an imperfect reaction (to the right side in reaction (1)), leaving Fe2⫹ unreacted with V5⫹ in the solution. Extraordinary large differences observed in JR-1 and JR-2 are, therefore, attributed to long-term oxidation during overnight sample decomposition. Because these two samples are rhyolitic obsidians, they were instantaneously decomposed by HF attack, and were subsequently affected by atmospheric oxidation for a relatively longer time compared to the other samples. In JG-2, the main source of Fe2⫹ is mica, which also decomposed rapidly causing a relatively longer oxidation period, resulting in –14% difference. The second possibility is the insufficient decomposition of rock samples, which also causes lower recovery of Fe2⫹, because the sample solution is not heated during decomposition in Wilson’s method. Actually, even in zircon-free samples, undissolved minerals were found in JP-1, JGb-1, JGb-2, and JH-1 after overnight decomposition. Such undissolved minerals were identified as spinel for JP-1 and orthopyroxene for JGb-1, JGb-2, and JH-1, respectively, using a SEM-EDX at PML (Horiba EMAX-7000 energy-dispersive X-ray spectrometer assembled into a Hitachi S-3100H scanning electron microscope). For these four samples, we also investigated rock powders before their decomposition using SEM-EDX, and found the existence of several Fe2⫹-containing minerals besides spinel and orthopyroxene (olivine for JP-1; clinopyroxene, hornblende, magnetite, and ilmenite for JGb-1; clinopyroxene and ilmenite for JGb-1; olivine, clinopyroxene, and hornblende for JH-1). We conclude that spinel and orthopyroxene are significantly acid-resistant compared to other rock-forming minerals, and preferentially remain when the sample solution is not heated. These observations clearly explain the large negative differences of these samples analyzed by Wilson’s method relative to results obtained by our method. Also, the lower FeO values of GSJ reference data might contain
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⫺d[Fe2⫹]/dt ⫽ k⬘[Fe2⫹]a
(3⬘)
Table 3 shows the rate constants for Eqn. 3⬘ supposing zero-, first-, and second-order reactions, which were calculated from the observed long-term variation of [Fe2⫹]. Only the zero-order reaction yielded similar rate constants irrespective of the initial concentration of Fe2⫹, indicating that [Fe2⫹] itself does not kinetically control the oxidation of Fe2⫹ in our experiments. Using the average of the k0 values in Table 4, the variation of [Fe2⫹] is expressed as [Fe2⫹]t ⫽ [Fe2⫹]0 ⫺ 0.00005 䡠 t (mmol/g)
(4)
where [Fe2⫹]0 represents the initial concentration of Fe2⫹ and t is time (in hours), respectively. In our analyses of silicate samples, [Fe2⫹] is determined once a day before the sample analyses. In some cases, analytical time within a single day exceeds a few hours, resulting in significant decrease in [Fe2⫹] and necessitating corrections using Eqn. 4. Fig. 2. Long-term variations of the concentrations of Fe2⫹ (}) and V5⫹ (f) solutions. Each error is determined from a standard deviation out of more than 5 times’ analyses.
analyses with insufficient sample decomposition and, therefore, sample digestion by HF with heating is indispensable for precise analyses of FeO in silicate rocks. 3.3. Fluctuation of the Concentration of Reagents Long-term stability of the concentrations of V5⫹ in 5 M H2SO4 and Fe2⫹ in 1 M H2SO4 solutions is important because these concentrations directly affect the concentration of FeO determined by our method. Although the concentration of V5⫹ solution was constant within analytical error over 10 days, the Fe2⫹ concentration showed a long-term decrease exceeding analytical error, and this decrease was clearly due to oxidation of Fe2⫹ by oxygen dissolved in the solution (Fig. 2). To determine the oxidation rate of the solution, we prepared several solutions of Fe2⫹ in 1 M H2SO4 with different concentrations of Fe2⫹, and examined their long-term variations. The Fe2⫹ oxidation reaction kinetics can be explained by the following equation: ⫺d[Fe2⫹]/dt ⫽ k[Fe2⫹]a 䡠 [O2]b 2⫹
3.4. Effect of the Amount of Indicator In the titration of K2Cr2O7 using sodium diphenylamine-4sulfonate as an indicator, surplus Cr6⫹ exceeding the equivalence point is required for the oxidation of the indicator itself. Thus, the amount of Cr6⫹ consumed up to the end point, which is different from the equivalence point, results in a positive error in the determination of FeO concentration. To evaluate the effect of the indicator, we have titrated 0.1 g of solid ammonium iron(II) sulfate hexahydrate while changing the amount of the indicator. As shown in Figure 3, the surplus Cr6⫹ can be ignored only when one drop of the indicator solution, which contains ⬃0.2 mol of sodium diphenylamine-4-sulfonate, is added. Otherwise progressively larger errors are produced as the amount of indicator increases. Furthermore, too much addition of the indicator would result in the side reaction, namely, the formation of green insoluble material by a reaction between the indicator in reduced and oxidized forms. This side reaction also consumes excess Cr6⫹, potentially leading to erroneous results. Because the above experiment was carried out with a total of 40 mL of titrated solution, two drops of the indicator would be permissible without correction if the whole solution is doubled.
(3)
where [Fe ] and [O2] are their concentrations in the solution, k is the rate constant, and a and b are the order of reaction, respectively. Provided that oxygen dissolved in the solution is always at equilibrium with air, [O2] can be considered constant. Equation 3 then becomes more simple:
3.4.1. Effect of the amount of HF We also found that coexisting HF in the solution interferes with the titration of Fe2⫹. As was described in section 3.4, we have examined the titration of 0.1 g of solid ammonium iron(II) sulfate hexahydrate while adding varying amounts of HF. As
Table 4. Rate constants (k) based on the assumptions of zero-, first-, and second-order reactions with various concentration of initial Fe2⫹. [Fe2⫹]0 mmol/g
k0 mmol 䡠 g⫺1 䡠 day⫺1
k1 day⫺1
k2 mmol⫺1 䡠 g 䡠 day⫺1
0.5664 0.4144 0.3346 0.3171
0.00127 0.00121 0.00123 0.00122
0.00227 0.00293 0.00374 0.00391
0.00407 0.00713 0.01133 50.01249
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or 101.427 (Arthur et al., 1999). HF dissociates weakly in the solution as: HF ⫽ H⫹ ⫹ F⫺
K2
(6)
where K2 is the acid dissociation constant, which is 10⫺3.17 (Ho¨ gfeldt, 1979). Combining these two equations yields: K1 䡠 K2 ⫽
[FeF⫹] [H⫹][F⫺] [FeF⫹][H⫹] 䡠 ⫽ [Fe2⫹][F⫺] [HF] [Fe2⫹][HF]
(7)
where [ ] represents molar concentrations. In our experimental condition, [H⫹] is controlled by the concentration of H2SO4, and that was 1.8 mol/L. Thus, [Fe2⫹] ⫺ [FeF⫹] ⫽ 1 ⫺ K 1 䡠 K 2 䡠 [HF]/1.8 [Fe2⫹]
Fig. 3. Relationship between recovery yield of Fe2⫹ and the amount of indicator varying from 0.2⬃1.1 mol, which corresponded to 1⬃5 drops of the indicator solution. The total amount of titrated solution is fixed as 40 mL.
shown in Figure 4, the recovery yield of Fe2⫹ shows a negative correlation with the amount of HF. This indicates that the existence of HF interferes with the redox reaction between Fe2⫹ and Cr6⫹ by forming an inert complex with Fe2⫹. The cation complex FeF⫹ is one cause of the problem. The formation of FeF⫹ is given as: Fe2⫹ ⫹ F⫺ ⫽ FeF⫹
K1 0.83
where K1, the equilibrium constant, is 10
(Ho¨ gfeldt, 1979)
(8)
The left side of this equation can be approximated by the recovery of Fe2⫹ titrated if [FeF⫹] is inert. Dashed and bold lines in Figure 4 are the theoretical lines of recovery yield calculated from Eqn. 8, when K1 equals 100.83 and 101.427, respectively. Our data broadly match the behavior of the theoretical lines, supporting the existence of inert [FeF⫹] in the solution. Although some data fall off these lines, those data can be explained by complicated factors involving coexisting species such as Cr6⫹, Cr3⫹, and Fe3⫹ in the solution. The most effective way to solve the problem of the formation of FeF⫹ is to use boric acid as a masking agent. We found that the use of 10 mL of saturated boric acid/1 mL of coexisting HF was effective in preventing the formation of inert complexes (Fig. 5). Overuse of boric acid, exceeding the above rate, did not interfere with the titration. However, boric acid did not work as a masking agent when HF attacked Fe2⫹ before the addition of the boric acid. Therefore, in the actual determination procedure for natural samples, 0.3 mmol/g Fe2⫹ solution is first weighed into the Pyrex beaker, and then saturated boric acid is added before the addition of the decomposed sample containing HF. 3.5. Detection Limit and Blank
Fig. 4. Relationship between recovery yield of Fe2⫹ and the amount of coexisting HF. Two lines indicate theoretical recovery of Fe2⫹ when K1 is 100.83 (dashed) or 101.427 (bold).
Detection limit of the method is largely attributed to the uncertainty of the amount of oxidation-reduction reagents actually used, namely, Fe2⫹, V5⫹, and K2Cr2O7 solutions. This can be estimated from the propagation of errors accompanied with each reagent, and it generally corresponds to ⬃0.00045 mmol of Fe2⫹ in this study. This means that Fe2⫹ contained in the starting material can not be distinguished from the error of reagents when it is less than 0.00045 mmol, which we define as the detection limit of our method. We carried out blank runs by determining the amount of Fe2⫹ in the procedure described in section 2.2.1 without adding any starting material. 1.5 mL of HF–H2SO4 mix solution was used, as in the case for decomposition of 0.2 g of silicate rock powder. We totally tested 3 sets of blank runs, in which each run comprised of a set of 5 blank measurements. All the blanks were less than ⫾0.00030 mmol of Fe2⫹, indicating that the blank does not exceed the detection limit. Therefore, any element that interferes with the series of redox reactions in our chemistry is absent in the reagents used in this study.
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Fe2⫹. To test the validity of such a system, we measured JG-1a that showed the worst reproducibility in Table 2 together with procedural blank. Four analyses gave 1.406 ⫾ 0.009 (1) wt.% of FeO, which is not only the same value as Table 2, but also resulted in better reproducibility, which improved from 0.92% to 0.63%. Because the average of triplicate blank runs was less than the detection limit (⫺0.00017 mmol of Fe2⫹), it is concluded that the dilution of the concentrations of Fe2⫹, V5⫹, and K2Cr2O7 solutions is effective not only to decrease the minimum range of the analysis, but also to improve the precision for samples with extremely low FeO contents. 4. CONCLUDING REMARKS
Fig. 5. The effect of saturated boric acid for masking HF. Open circles indicate 100% recovery of Fe2⫹ in the titration with coexisting HF, and crosses show the failure of the titration. 10 mL of saturated boric acid/1 mL of coexisting HF is found to be enough.
3.6. Modification of the Method In this study, we used less than 0.15 mmol of Fe2⫹ in starting materials while decomposing less than 0.4 g of rock samples. Providing 0.1 g of the sample for decomposition avoids sample heterogeneity, the measurable maximum FeO concentration in a sample with this method is calculated to be 10.8 wt.%. Because some natural samples (e.g., ferrobasalt) contain ⬃20 wt.% FeO, we evaluated the measurement of 0.22 g of JB-2 (containing ⬃0.30 mmol of Fe2⫹) when increasing the V5⫹ solution as 1.17 g (⬃0.35 mmol). Although the final solution became more blue during titration (due to larger amount of V4⫹) than the normal case, we obtained 10.07 ⫾ 0.01 (1) wt.% of FeO as a result of five analyses, which is identical to that of routine analyses in Table 2 when errors are considered. Therefore, our method can be applied to samples containing ⬎20 wt.% FeO by simply increasing the amount of V5⫹ solution. On the other hand, the minimum range of our method is restricted by the detection limit or blank of the method. Assuming that the blank of the whole procedure has a composition of the detection limit (0.00045 mmol), at least 0.09 mmol of Fe2⫹ is required as the starting material if one would like to suppress the effect of the blank to less than 0.5% of Fe2⫹ contained in the sample used. This corresponds to 1.6 wt.% of FeO when 0.4 g of sample is used, and the effect of the blank increases as the FeO content in the sample decreases. One simple way to further improve the minimum range is to enhance the resolution of the titration by diluting the concentrations of Fe2⫹, V5⫹, and K2Cr2O7 solutions. When all the concentrations of these solutions were diluted to half of the usual case (0.15 mmol/g for Fe2⫹ and V5⫹ solutions, and 0.02 N for K2Cr2O7 solution), the detection limit estimated by the same way as described in section 3.6 is ⬃0.00021 mmol of
We have established an accurate and precise method for the determination of Fe2⫹ in silicate rocks with analytical error better than 0.5% (1). In this method, the oxidation of Fe2⫹ can be avoided by keeping the sample solution at higher acidity (5 M H2SO4). Sample decomposition in a relatively small beaker (7 mL) with a small amount of sample solution is very effective for preventing atmospheric oxidation. Fe2⫹ concentrations in 15 GSJ standard silicate rock powders were determined in this study, yielding values higher than those previously reported (except for one sample), indicating that the GSJ reference values were determined with lower recovery of Fe2⫹ owing to its oxidation or insufficient sample decomposition during measurement. In addition to the necessity of avoiding oxidation of Fe2⫹, we have demonstrated several important factors that determine the accuracy of our method: (1) correction of the long-term variation of the concentration of Fe2⫹ solution; (2) the amount of indicator; (3) the use of boric acid as a masking agent to protect Fe2⫹ from forming inert FeF⫹ complexes during titration. Acknowledgments—We would like to thank N. Imai (Geological Survey of Japan) for offering GSJ standard rock powders. We are grateful to G. E. Bebout and M. J. Walter for improving the English used in this paper. We also thank A. Makishima, H. Takei, N. Takeuchi, and all the other member of PML for their analytical support and useful discussion. We would like to express our appreciation to the Associate Editor, B. R. Frost and anonymous reviewers for their efforts to improve the quality of this paper. This research was supported by Grants-in-Aid for Scientific Research from the Ministry of Education, Science, and Culture, Japan (Monbusho) and the Japanese Society for the Promotion of Science (JSPS) to E.N. Associate editor: B. R. Frost APPENDIX Wilson’s Procedure 0.5 g of powdered sample and 0.1 g of solid ammonium metavanadate (special grade, Kanto Chemical) were weighed into a 100-mL polypropylene beaker, and 10 mL of 27 M HF was added. The mixture was kept at room temperature overnight for complete sample decomposition. After adding 30 mL of 5 M H2SO4, the contents were transferred into a 500-mL glass beaker with 250 mL of saturated boric acid. After adding 7 drops of 0.2% sodium diphenylamine-4-sulfonate, the solution was titrated by 0.07 mol/L Fe2⫹ solution, which was conditioned by dissolving 27.5 g of ammonium iron(II) sulfate hexahydrate into 1 L of 0.5 M H2SO4. The solution containing only ammonium metavanadate was also treated the same as the sample and titrated to determine the actual concentration of the Fe2⫹ solution.
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