Determination of Ferrous and Total Iron in Silicate Rocks by Automated Colorimetry David Whitehead and S. A. Malik Sedimentology Research Laboratory, Reading Universrty, Reading, England
Schafer’s ( I ) review of methods for the determination of ferrous iron in rocks and subsequent publications (2-5) indicate that manual methods are always used. The methods described below are for automated determinations of FeO and total iron (as Fez03) on aliquots of a single sample solution using a common set of standards. Provided that sufficient AutoAnalyzer modules are available, simultaneous determinations may be made. The manifolds described below are for separate determinations but require only slight modification for simultaneous analysis. The procedure for ferrous irop determination is based on Wilson’s (6) colorimetric method. In this method, ferrous ions are first oxidized by vanadate in acid solution during the decompositional stage to prevent aerial oxidation. The reaction may be represented by the equation:
The reaction is reversed by buffering the solution a t pH 5 and the violet-colored ferrous complex formed by addition of 2,2’-dipyridyl is measured a t 520 nm. Total iron is also determined colorimetrically with 2,2’-dipyridyl after reduction of ferric ions by hydroxylamine hydrochloride.
EXPERIMENTAL Reagents. All reagents are of Analytical Reagent grade and deionized water is used throughout. Bu~jer-Z,Z’-dipyridy1Reagent is made by separately dissolving 0.1 g 2,2’-dipyridyl in 30 ml of water containing 4 ml of concentrated HC1 (sp gr 1.18)and 34 g of hydrated sodium acetate in 5G0 ml of water, mixing the two solutions and diluting to 1 liter. The pH of this buffer solution is 4.9. Vanadate Solution is made by dissolving 10 g of ammonium metavanadate in 110 ml of 1:l H2.304 and diluting the solution to a liter. Hydroxylamine Hydrochloride is 3% w/v. Hydrofluoric Acid is 40% v/v. Boric Acid is 4%wlv. Hydrofluoric-Boric Acid Mixture is made by dissolving 40 g “BO? in water, adding 50 ml of 40% vlv H F and diluting to a liter. Standard Ferrous Solution is made by dissolving 0.2 g of highpurity iron in 10 ml of 2.5N HCl on a hot-plate and diluting the solution to a liter in a graduated flask. A set of ferrous standards to cover the range 5-25 pg Fe2+/ml is prepared containing additions of vanadate and HF/HzB03 mixture to give a matrix similar to the test solutions. Ferric Spike Solution is made by dissolving 0.2 g of high-purity iron in 10 ml of 2.5N HC) on a hot-plate, evaporating the solution to a small volume, and oxidizing the ferrous ions with a few drops of concentrated “ 0 3 and diluting to a liter with water. A 10 pg Fe”/ml solution is prepared by diluting the stock solution on the day of use. Procedure. Grind a sample to a grain size of 150 microns, or, when resistant minerals are present, to about 60 microns by grinding under acetone. Weigh out a sample to contain between 10-20 mg Fe2+. Transfer the dried sample to a polystyrene container (85 ml) containing 5 ml of vanadate solution. Add 10 ml of HF, cover the container, and leave overnight on a gently oscillating surface at room temperature keeping the reactants away from direct light. Transfer the decomposition to a liter graduated flask, using 200 ml of boric acid solution as wash, and make up to the mark with water. Rocks containing resistant minerals are decomposed in a 554
sealed Teflon bomb (7)a t 65 “C for 18 hours using the same reagents as for the cold method. FeO and total iron determinations are made using the manifolds shown in Figures 1 and 2. A sampling rate of 40 cups per hour is used with alternate cups acting as a water wash.
RESULTS AND DISCUSSION Initial tests were carried out using a manual spectrophotometer (Unicam SP500) to find the optimum pH range for both determinations. The stability of the ferrous-2,2’-dipyridyl complex a t various pH values was determined by measuring the absorbance of a 4 pg Fe2+/ml standard, containing only 2,2’-dipyridyl, buffered to a known pH. The change in absorbance with p H is shown in Figure 3; maximum sensitivity is achieved between p H 4-5. The optimum p H range for the FeO determination was found by measuring the absorbance of a set of ferrous standards, containing vanadate 2,2’-dipyridyl, and a mixture of HF/H?BO3, at a known pH. The calibration graphs determined a t different pH values are shown in Figure 4. There is a sharp decrease in the sensitivity below pH 4.6. The decrease in sensitivity is probably caused by the slowness of the reverse reaction a t the lower p H and not to incomplete reaction between ferrous ions and 2,2‘-dipyridyl since Figure 3 shows that quite high sensitivity is to be expected even a t pH 3.3 when vanadate is absent. This is also confirmed by the fact that no absorbance was observed a t p H 3.3 except when a reducing agent is present (hydroxylamine fiydrochloride) to convert all the Fe3+ to Fez+ (see Figure 4). This means that to use the same buffer-2,2’-dipyridyl reagent for the FeO and total iron determinations, the buffer solution must lie in the range pH 4.6-5.0. In the determination of FeO, Wilson (6)found that a ferric complex (possibily ferric acetate) causes enhancement. A correction was applied after measuring the absorbance of the ferric complex a t a wavelength (420 nm) a t which the color of the ferrous-2,2’-dipyridyl is much reduced. However, measurements made during the development of the present automated method have shown that, under the analytical conditions used, enhancement by ferric ions is not caused by simple overlap of spectral bands: for example, a 10 pg Fe2+/ml test solution containing 20 pg Fe3+/ml was enhanced by 0.075 absorbance unit when compared with a similar solution containing no Fe3+, whereas a 20 pg Fe3+/ ml solution containing no Fez+ produced an absorbance of only 0.004 unit. Tests were carried out to see whether or not a plateau of enhancement occurred after the addition of excess ferric ions to the system. A ferric solution was taken into the system along the Fe3+ reagent line shown in Figure 1;this ferric solution will subsequently be referred to as the spike solution. A series of spike solutions from 3-39 pg Fe3+/ml, in increments of 3 pg Fe3+/ml, were prepared. The absorbance of a set of ferrous standards was determined for each spike solution taken into the system. A plateau of enhancement occurred for all the standards for concentrations of spike
ANALYTICAL CHEMISTPY, VOL. 47, NO. 3, MARCH 1975
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for analysis
0 5 O
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Figure 1. FeO
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manifold Figure 3.
i
Absorbance graph of
Tube size (inches) Sample
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4 pg
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Figure 2.
Total iron manifold ,PH3
solution in the range 5-25 pg Fe3+/ml. The optimum concentration of the spike solution for intake into the system is about 10 pg Fe3+/ml; this ensures that the plateau is easily reached, yet the concentration is low enough to make the additional Fe3+ from rocks less than the critical 25pg Fe3+/ ml. The calibration graph is linear when a 10 pg Fe3+/ml solution is used (Figure 5, curve A ) . However, if de-ionized water is taken into the system instead of the spike solution, not only is there no enhancement, but the curve is non-linear (Figure 5 , curve B ). Enhancement in the absorbance of the standards is produced, however, by addition of extra time-delay coils to the system with a resulting increase in the linearity of the graph. This indicates that in the absence of extra time-delay coils, the back reaction is incomplete. The addition of excess ferric ions to the system as a spike increases the rate of the back reaction. In the manual method, and in the absence of ferric ions, a linear calibration graph i s obtained with a slope that is constant for a t least 24 hours showing the reaction to have gone to completion. Additional tests using the manual instrument showed that ferric ions also produce an enhancement by, complex formation. The slope of the calibration graph is increased when equal amounts of ferric ions are added to each of the ferrous standards. Also, there was continous enhancement when successively larger amounts of ferric ions (in the range 0-30 pg Fe3 /ml) were added to each standard. In contrast, in the automated method, increasing addition of ferric ions produce a plateau of enhancement which is taken to indicate that the formation of a ferric complex is slow and insignificant. However, excessive concentration of dipyridyl (0.4 g/l.) in the presence of the 10 gg Fe3+/ml spike solution caused a gradual base-line drift. Previous experience with other automated methods has shown that the
1
2
3
3
4
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Figure 4.
Dependence of the FeO calibration graph on pH
( 0 )Effect of pH on FeO standards. (+) Hydroxylaminehydrochloride added to the standards
/
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Figure 5.
Effect of 20 pg Fe3+/mlon t h e FeO standards
( A ) Standards with 20 Fg Fe3+/ml added. (B)Standards containing no Fe3+
formation of a precipitate adhering to the flow-cell walls can cause this. A concentration of 0.1 g dipyridyl/l. gave a perfectly stable base line. Table I shows that the FeO determinations on French standard rocks and Mica-Fe (8) by the proposed method compare favorably with those done by a modified Wilson's
ANALYTICAL CHEMISTRY, VOL. 47, NO. 3, M A R C H 1975
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Table I. comparative FeO Values for Silicate Reference Standards
Table 11.Comparative Total Iron Values for Silicate Reference Standards
Proposed automated method Standard reference sample
FeO, %
Sa
Modified Wilson’s tihimetric C.V.* m e t h c d F e O , %
Granite GA Mica-Fe
1.39 0.01 0.72 1.39 18.85 0.10 0.53 18.75 Basalt BR 6.72 0.04 0.60 6.66 Diorite DR-N 5.49 0.01 0.18 5.45 Standard deviation of 5 repeat determinations. C.V. = Coefficient of variation.
titrimetric method (9) (the latter method has been used for several years in these labordories). Total iron is determined on aliquots of sample solution by first reducing the FeJ+ (and V5+), followed by addition of the buffer-2,2’-dipyridyl reagent. The reduction was incomplete in the presence of vanadate when the reducing agent was added to the buffer as in Riley’s (10) method. In theory, it is not necessary for a reaction to go to completion when an AutoAnalyzer is used. However, reduction at pH 4.9 gave inaccurate results and poor precision. Table I1 compares the determination of total iron by the proposed method with those done by Riley’s method (10). Except for the total iron value obtained for granite GA (2.56% Fez03), the relative error between the two methods is less than 1%.To test whether this low value obtained for granite GA was caused by incomplete decomposition or by the method of determination, two powder aliquots were decomposed in a mixture of hydrofluoric and perchloric acids as described by Riley ( I O ) , followed by additions of vanadate and a mixture of HF/H3B03 to give the same matrix as a normal test solution, and the determinations done using the Fez03 manifold shown in Figure 2. A total iron
Standard reference
Proposed automated method Fe203, %a
sample
Sb
C.V.C
Riley’s method F e 2 0 3 , 160
Granite GA Mica-Fe Basalt BR Diorite DR-N
2.56 0.02 0.62 2.71 26.02 0.04 0.14 25.94 13.07 0.05 0.41 12.84 9.53 0.05 0.50 9.64 a Total iron as Fez03. Standard deviation of 5 repeat determi. nations. Coefficient of variation.
value of 2.73% was obtained (as compared to 2.71% obtained by Riley’s method) showing incomplete decomposition to be the cause. This was confirmed when a total iron value of 2.72% was obtained after decomposition in a Teflon bomb. ACKNOWLEDGMENT The authors thank J. E. Thomas for helpful criticism of the manuscript. LITERATURE CITED H. N. S. Schafer. Ana/yst (London),91, 755 (1966). E. M. Donaidson, Anal. Chem., 41, 501 (1969). B. R. Sant and T. P. Prasad, Talanta, 15, 1483 (1968). J. L. Girardin and R. Thiel, Bull. Rech. Pau., 4, 513 (1970). W. J. French and S.J. Adams, Analyst (London).97, 828 (1972). A. D. Wilson, Analyst (London),85, 823 (1960). B. Bernas, Anal. Chern., 40, 1682 (1968). H. de la Roche and K. Govindaraju, Method. Phys. Anal., 7, 414 (1971). A. W. Hounslow and J. M. Moore, Geological Paper 66-1, Carleton University, Ottawa, Canada, 1966. (IO) J. P. Riley, Anal. Chirn. Acta, 19, 413 (1958). (1) (2) (3) (4) (5) (6) (7) (8) (9)
RECEIVEDfor review June 10, 1974. Accepted October 10, 1974.
New Automated Colorimetric Method for the Determination of Chloride Using Chromotropic Acid Badar K. Afghan, Ricky Leung, Achut V. Kulkarni,‘ and James F. Ryan Analytical Methods Research Section, Water Quality Research Division, Canada Centre for Inland Waters, Burlington, Ontario, Canada
Chloride occurs in almost all natural waters and enters bodies of water through natural sources such as minerals, or may be derived from other sources such as human and animal sewage, agricultural wastes, effluents from paper works, petroleum refineries, etc. The presence of chloride is generally not harmful to human beings until very high concentrations are reached ( I , 2 ) . However, the presence of chloride and other substances in waters used in some industries may have a pronounced effect ( 3 ) .Water containing concentrations of chloride as low as 0.5 mg/liter accelerates stress erosion in power reactors ( 4 ). Chloride also has an immediate effect on corrosion of steel at concentrations as low as 3 m g h t e r (5, 6). In the majority of water quality laboratories in Canada and the USA, chloride is determined by the method based on displacement of thiocyanate ion from mercuric thiocyPermanent address, Analytical Chemistry Division, Bhabha Atomic Research Centre, Trombay, Bombay, 400085 India. 556
anate by chloride ion, and the subsequent reaction of the liberated thiocyanate ion with ferric ion to form a colored complex (7, 8). This method utilizes toxic reagents such as mercuric thiocyanate and, because of the increasing concern to reduce pollution from heavy metals such as mercury, it is desirable to have a method which does not require the use of toxic reagents, particularly when analyzing a non-toxic parameter such as chloride. During our studies for the development of a colorimetric rnethod for formaldehyde, chloride in the the presence of nitrate gave a positive interference during the determination. However, when these ions were present individually, they did not produce any significant error (9). Under the experimental condition used during formaldehyde analysis, chloride ion catalyzed the conversion of nitrate to nitrite and the resultant nitrite reacted with,chromotropic acid to produce strongly colored species. Therefore, it was decided to optimize this reaction to determine chloride and possibly replace the existing method.
ANALYTICAL CHEMISTRY, VOL. 47, NO. 3, MARCH 1975
