ANALYTICAL CHEMISTRY
1374 by both oxygen (2) absorption and this method. Table, I1 summarizes the data. Accuracy. Duplicate analyses in which larger aliquots of radical solutions were taken usually gave results which agreed more closely than would be anticipated from the data presented above. For this reason, and because of the difficulty involved in inserting the syringe through the serum stopper without occasionally losing a drop of solution, the authors believe that the accuracy of the method is limited by the inaccuracy in estimating the volume transferred by the syringes which are relatively crude volumetric apparatus. They believe that further improvement can be attained through the use of very fine needleand narrow-bore, precision syringes.
ACKNOW LEDGRIENT
The authors wish to acknowledge the financial support of this work by the Offic,e of Naval Research. LITERATURE CITED
(1) Bachmann. lt-. E., and Osborn, G.. J . O T ~Chem., . 5, 29 (1940). (2) Gomberg, hl., and Schoepfle, C . S., J . Am. Chem. SOC.,39, 1661 (1917). (3) Hammond, G. S.. Rudesill, J. T . , and Modie, F. J., Zbid., 73, 3929 (1951). (4) Preckel, R.,and Selwood, P. IY.,Ibid., 63, 3397 (1941). ( 5 ) Wieland, H., Ploetz, T., and Indest, H., Ann., 532, 175 (1937). (6) Ziegler, K., Orth, P., and Weber, IC., Zbid., 504, 131 (1933). RECEIVED for review February 27. 1951.
.ircepted May 12, 1952
Polarographic Determination of Iron in Nonferrous Alloys LOUIS MEITES, Yale University, New Haeen, Conn. H E polarographic literature contains descriptions of two Tmethods for the determination of iron. One ( 2 ) is based on the measurement of the height of the iron(II1) wave in an acid solution containing no complexing agent. This method is subject to interference by other ions which are also capable of oxidizing mercury (14). The other ( 6 ) requires separation of copper and reduction of iron to the +2 state, the concentration of ferrous iron being determined by measuring the height of its anodic wave in 1 F potassium oxalate a t a pH near 5. This method is specific for iron and is capable of giving very accurate results. It has not, however, been very widely adopted, presumably because of the care necessary to prevent air-osidation of the ferrous complex. The present communication describes a method which is free from both of these drawbacks. The final solution contains the iron in the form of a mixture of citrate complexes which are inert to both air oxidation and reduction by mercury, so that a rrelldefined wave is secured whose height remains constant for long periods of time. EXPERIMENTAL
All polarographic measurements rvere made with a calibrated Sargent-Heyrovskj. Model XI1 recording polarograph, using the visual scale of the instrument. A conventional H-cell modified by the insertion of a sintered-glass gas dispersion cylinder in the influent gas stream to permit rapid deaeration was used (IS). The hydrogen used for deaeration was freed from oxy en by passage through two vanadium(I1) perchlorate wash tottles (11, 12). A nTater thermostat was used to maintain a temperature of 25.00' f 0.03' C. A Beckman Model G pH meter was used for pH measurements, and a saturated solution of potassium hydrogen tartrate was used as the pH standard (1,s). All volumetric apparatus was calibrated by conventional methods. A stock iron(II1) solution was prepared by dissolving electrolytic iron in perchloric acid, eva orating to fumes of the acid, and diluting to known volume. T i e concentration of this solution was checked by reduction in a Jones reductor and titration with standard permanganate. All other chemicals were ordinary reagent grade and were not further purified except as described below. DATA AND DISCUSSION
The investigations of Lingane ( 5 ) and Meites (7-10) of the polarographic characteristics of iron( 111) and copper( 11) in tartrate, citrate, and oxalate media indicated that, in these media a t least, it was impossible to secure a satisfactory separation of the waves of these elements. This is also true in solutions containing certain other organic anions which form iron(II1) complexes sufficiently stable to have measurable half-wave potentials. ..le-
cordingly, it appeared that a satisfactory general method for the determination of iron must involve the separation of copper before the polarographic measurement. The most rapid and convenient method found for effecting this separation consists of the reduction of copper(I1) to the metal with amalgamated zinc. In a weakly acidic solution this reduction proceeds to completion within 30 seconds or less, which is a considerable advantage when many samples are to be analyzed. The reduced solution and the precipitated copper can easily be removed from the residual zinc by decantation, so that the zinc can be used for the reduction of another sample. Bismuth is quantitatively removed viith copper, which is of importance because bismuth also gives n-aves a t relatively positive potentials in the media under consideration, and these viould also interfere vith the iron wave. Attempts were made to apply this method of separation to solutions containing nitrate, such as would result from the solution of an alloy in a nitric-hydrochloric acid mixture. However, it R-as found that nitrate seriously impaired the efficiency of the reduction: Large amounts of nitrogen dioxide r e r e evolved unless the pH was so high that hydrous metal oxides precipitated. It is therefore necessary to remove nitrate from the solution priorto the treatment with zinc by a double evaporation to near dryness with excess hydrochloric acid. This both prevents the precipitation of tin(1V) oxide, which would retain appreciable amounts of iron, and causes the reduction with zinc to proceed smoothly. The reduced solution, after filtration through a coarse quantitative filter paper to remove metallic copper, bismuth, and silver, contains iron and tin in the +2 states. These are reoxidized by the addition of a small excess of potassium permanganate solution, and the excess permanganate is immediately destroyed by the addition of oxalic acid. A number of experiments were made in which a large excess of oxalic acid was added to serve as the supporting electrolyte. However, the difficulties encountered in this method outweighed the expected advantages. Much of the zinc precipitates, carrying down iron with it, and the re-solution of the zinc oxalate Fas troublesome and inconvenient. Tartaric and citric acids do not give insoluble zinc salts under these conditions, and are therefore to be preferred to oxalic acid. As there seemed to be no other grounds for choice, citric acid was selected because of its slightly lower cost. The reduced solution is accordingly treated with sufficient citric acid to make the final citrate concentration approximately 0.25 F , and the p H is adjusted to a value near 5 by the addition of concentrated ammonia. The diffusion current constant of iron(II1) in citrate media, measured a t -0.55 volt us. S.C.E., is independent of pH between about 3 and 6.7, so that the exact p H
1375
V O L U M E 24, NO. 8, A U G U S T 1 9 5 2 of the h a 1 solution is not critical. At p H values below 3, however, Z increased slightly, while as the p H is raised above 6.7 the wavc height at -0.55 volt decreased and a new wave appears a t more negative potentials (10). I t is convenient to perform this neutralization in the presence of methyl red, adding ammonia until the color of the indicator just begins to deviate from red. The methyl red here serves three important purposes. First, it provides an easy method for the neutralization of the solution t o a p H within the optimum range. Second, it eliminates a maximum formed a t high iron concentrations in the absence of a maximum suppressor: this leads to an increasing value of the ratio i d / C with increasing iron(111) concentration above about 0.8 mM. In the presence of 0.0005% methyl red the diffusion current a t -0.55 volt is accurately proportional to the concentration of ferric iron over the entire range from 0.0095 to 19.2 mM (Table I). Finally, it greatly reduces the rate of the photocatalytic reduction of ferric iron (6), presumably because the acid form of the methyl red removes the components of the incident light which are most effective in promoting this reaction. In the absence of methyl red the cathodic iron(II1) diffusion current may decrease a t the rate of several tenths of a per cent per hour when the solution is allowed to stand in diffuse daylight before the analysis, which is clearly disadvantageous in practical analyses. If the p H is so high that most of the methyl red is converted to the alkaline (yellox) form, this last advantage is lost: Either the pH should be lowered again by the addition of hydrochloric acid or the solution should be stored in the dark until it is analyzed.
mium, cobalt, magnesium, manganese, nickel, and zinc indicate that none of these elements will interfere d t h a diffusion current, measurement a t -0.55 volt us. S.C.E. in 0.25 F citrate a t a pH near 5. Lead gave a wave for which El,? = -0.765 volt: ita slope is sufficiently great so that even 0.8 grain of this element would cont,ribute less than 0.01 ramp. to t’he diffusion current measured in t,he recommended procedure. Tin( IY) appears to be polarographically inert in this medium. In some solutions prepared from reagent grade tin or t,in salts a wave was obsewed for which E,/, was -0.95 volt, but it was much too small to correspond t,o complete reduction to tin(I1) and hence might have been due to impurities in the chemicals used. Stannous tin, which probably would interfere ( 4 ) , is completely oxidized by the t.reatment with permanganate. Uisrnut’h gives a wave for which E l / ? = -0.5% volt, but this element, is quant,itativcly removed in the reduction stcp.
I Table I
I
I
I
1
Diffusion Current Constant of Iron(II1) in 0.25 F Ammonium Citrate
(0.000570 methyl red, pH 5.4. E d , e , = - 0 . 5 5 volt us. S.C.E., 1 = 2.760 seconds. vi = 3 . 146 mg./sec., r n Z / 3 t ’ / 5 = 2 . 5 4 3 nig.z/a sec.-l, 2. Diffusion currents have bern corrected for residual current) Ferric Diffusion Iron, Current, 1\Iillimola~ fiamp. 0.00953 3.1; 0.030 0.0190 0 060 3 . I5 0,0380 0.121 3.18 3.136 0.1365 0.428 3.18 0.2167 0.689 0.2832 3.15’ 0.892 0,381 1,198 7.14l 0.734 2.320 16: 3.127 1.062 3.32 6.14 1,920 8.34 2.628 3.81 12.03 . 3.16 23.20 3.13’ 10.62 33,30 60.13 3.13; 19.20
b.
Mean i d / C = 3.156 i 0 . 0 1 8 i d / C r n e / a t 1 / 6 = 1 . 2 4 1 3~ 0 . 0 0 7
As shown in Table I, the diffusion current constant of iron(II1) in the h a 1 solution was equal to 1.241 f 0.007 pamp./mmole/ liter/mg.*/ sec. - 1 1 2 . This is considerably higher than the value, 0.93 =k 0.01, found by Lingane ( 5 )in 0.5 F citrate, p H 5 to 12, containing 0.005% gelatin. The difference is not due to the differing citrate concentrations, for Z was found to vary only about &2% as the citrate concentration was changed from 0.05 to 0.9 F, either in the presence of 0.0005% methyl red or, if the iron concentration a a s less than 0.8 mM, in the absence of any maximum suppressor. However, the presence of gelatin exerted a considerable effect on the value of Z. This is shown in Figure 1, in which the effects of gelatin and methyl red are compared. The value of I in the presence of 0.005% gelatin interpolated from this Figure is in fair agreement with that reported by Lingane ( 5 ) . The disadvantages which would be associated with the use of gelatin in this procedure are evident. The possible effects of the other elements which may be present in nonferrous alloys deserve discussion. The known polarographic characteristics of aluminum, antimony, areenic, chro-
In connection with this work it was of int’erest to check the iron contents of samples of reagent grade citric acid purchased froiii several manufacturers. In every case the label reported the presence of not more than 0.0005% iron. The blank due to so small an amount of iron could easily be corrected for, and in most cases could even be safely ignored. Actually, however, polarograms of solutions of the various preparations showed that t,he percentage of iron was, with the exception of one brand, considerably in escess of that claimed. The significance of this observatioii for practical work is that the residual current n u s t be measured in a solution containing the same concentration of citric acid as that to be used in later analyses: The common graphical method of correcting for the residual current, may give seriously erroneous result’sfor alloys low in iron. The following procedure has becn found t o give nearly coiiiplete removal of iron from several samples of citric wid Containing appreciable amounts of this element. Suspend a pound of reagent grade citric acid in about 100 inl. of water and heat nearly to boiling. Cool, discard the liquid phase (which contains most of the iron), wash the residual solid twice with small port,ions of water, and finally the acid recrystallizes from ice-cold eth~iiol. Methyl red is also reducible a t -0.55 volt, so that iii the analysis of a low-iron alloy it is advantageous to use only enough methyl red to serve as an indicator in the adjustment of the p H instead of the 0.0005% recommended below. These considerations led to the development of the following procedure. RECO.\IMENDED I’KOCEDURE
Weigh 1 to 10 grams of the alloy into a 250-ml. Erlerimeyer flask, add 25 ml. of concentrated hydrochloric acid and 10 nil. of concentrated nitric acid, and warm gently until decomposition is complete. Evaporate nearly to dryness, add 10 ml. of concentrated hydrochloric acid, and repeat the evaporation. Add about 20 ml. of distilled mater, swirl until solution is complete
13P6
ANALYTICAL CHEMISTRY
Table 11. Analyses of National Bureau of Standards Samples Iron, %
a
Sample Description and Number Aluminum brass 164 Aluminum alloy (casting) 86c Manganese bronse 62b Phosphor bronze 63b Sheet brass 37 Ounce metal 124b Sheet brass 37d Tin-base bearing metal 54, Cast bronze 52b Zinc-base die-casting alloy 94a Magnesium-base alloy 171 Purified citric acid, no methyl red.
Certificate value 2.52 0.90 0.82 0.47 0.29 0.26 0.076
0.033 0.032
0.013 0 002
Polarographic value 2.50 0.92 0.825 0.48 0.28 0.26 0.078 0.031 0.036 0.017 0.0015a
and deaerate with oxygen-free nitrogen or hydrogen. Measure the total current a t -0.55 volt 8s. S.C.E., and subtract the current measured under the same conditions in a solution containin the same amounts of oxalic and citric acids, methyl red, and ammonia. The difference between these currents is the diffusion current of iron, which is related to the percentag: of iron in a sample weighing W grams by the equation yo iron = 0.450 id l W, ’3t’ 16. RESULTS
Table I1 shows the results of analyses of a number of standard samples of nonferrous alloys issued by the National Bureau of Standards. In every case the polarographic value was in good agreement with the certificate value. LITERATURE CITED
(neglect any residue of lead chloride), and add 10 to 50 grams of amalgamated zinc containing not less than 1% of mercury. Agitate for 1 minute, adding a few drops. of hydrochloric acid if the solution becomes turbid ( recipitation of hydrous metal oxides). Decant the liquid and t\e precipitated copper through a rapid filter paper (Whatman 41H or equivalent) into a 100-mi volumetric flask. FVash the Erlenmeyer flask containing the zinc with several small portions of water pouring them through the filter paper into the volumetric flask. Discard the precipitate, and reserve the zinc for re-use. To the combined filtrates add 0.2 M potassium permanganate dropwise until the permanganate color ersists for a t least 5 seconds, then immediately add 2 ml. o f saturated (ca. 1 M ) oxalic acid. Add 5 grams of citric acid and 0.5 ml. of 0.1% methyl red, then sufficient concentrated ammonia to give the first perceptible color change, and finally dilute to the mark. Transfer a portion of the solution to a polarographic cell
E., Miller, R. G.. a n d Smith, E. R., J . Research AVatl.BUT.Standards, 47, 433 (1951). (2) Hohn, H a n s , 2.Elektrochem., 43, 127 (1937). (3) Lingane, J. J., IND.ENQ.CHEM.,ANAL.ED.,19, 810 (19471 (4) Lingane, J. J.. J . Am. Chem. Soc., 65, 866 (1943). (6)Ibid., 68, 2448 (1946). (6) RIeites, Louis, ANAL.CREM.,20, 895 (1948). (7) XIeites, Louis, J . Am Chem. Soc , 71, 3269 (1949). (8) Ibid., 72, 180 (1950). (1) Bates, R. G., Bower, V.
(9) (10) (11) (12)
Ibid., p. 184. Ibid., 73, 3727 (1951). Ibid., p. 4479.
& k i t e s , Louis, a n d hfeites, Thelma, ANAL. CHEW..20, 984 (1448) (13) Ibid., 23, 1194 (1951). (14) Mnich, Edmund, Z. Elektrochem., 44, 132 (1938). \ _ _ _ _ ,
RECEIVED for review May
Accepted May 12, 1952
1, 1962.
Determination of Iron in Water in the Presence of Heavy Metals K. L. MORRIS State Hygienic Laboratory, University of Iowa, Iowa City, Iowa
VAILABLE methods for the determination of iron in water
A are unsatisfactory in the presence of many heavy metals
encountered in mining and industrial waste waters, and require tedious separation techniques for elimination of these ions. Moss and Mellon (1) reported the formation of a colored complex of ferrous tripyridyl with a procedure for determination of iron. This procedure like previous methods was subject to severe interference by heavy metal ions such as copper, cobalt, nickel, and zinc. rl reIatively specific procedure involving the ferrous irontripyridyl complex with the sequestering of normally diverse heavy metals by ethylenediamine has been developed
and u.ed full qtrength. pared (2).
Standard ferric iron wlution was preAFP 4RATUS
Spectrophotometric investigation was done on a Beckman Model DU spectrophotometer using 1-em. Corex cells. Colorimetric studies were performed on a Fisher electrophotometer and with matched 100-ml. Sessler tubes. A Beckman Model H2 pH meter was used for all pH measurements.
A
06.
PRINCIPLE OF METHOD
Tripyridyl f o r m a reddish purple colored complex with ferious iron detectable visually in concentrations as low as 0.01 p.p.ni., and proportional t o the concentration of iron present. Experimental results indicate that one molecule of ferrous iron reacts with two molecules of tripyridyl. The p H range for this color formation varies from 1.5 to 12.0 with no change in hue over the entire range, and the colored complex is stable for at least 30 days. Ethylenediamine is used to sequester the heavy metals other than ferrous iron at a pH of approximately 9.6. Figure 1 shon s the spectrophotometric characteristics of the color system.
05.
f
3 0 PPU/
_lL h
ZOPPY
FE
i
1\ \
REAGENTS
The tripyridyl solution was prepared bv dissolving 0.1 gram of 2,2’,2”-tripyridyl (G. Frederick Smith Chemical Co.) in 100 ml. of 0.1 N hydrochloric acid. Hydroxylamine hydrochloride solution, used 89 the reductant, was prepared by dissolving 10 grams of the solid in 90 ml. of distilled water. Ethylenediamine (60%, practical grade) --as obtained from Eastman Kodak Co.
I
480
470
490
510
530
550
570
590
MILLIMICRONS
Figure 1. Spectrophotometric Characteristics of Ferrous Iron-Tripyridyl Complex
I
