Solutions for Colorimetric Standards VII. Aqueous Solutions of Salts of Elements 23 to 29 C. T. K4SLINE WITH M. G. MELLON, Purdue University, Lafayette, Ind.
AN
INSPECTION of the Bohr form of the periodic table of the chemical elements reveals the interesting fact that
the elements forming colored ions in aqueous solution include those assigned to the so-called transition group, together with a few others situated adjacent to or near them. Thus we find no color until element 22 (Ti) is reached. It continues then through the successive elements, 23 to 29 (V, Cr, Mn, Fe, Co, Ni, Cu), disappearing sharply with 30 (Zn). Aqueous solutions of certain salts of part of these elements have been used for various colorimetric purposes in a number of different ways. Their use in colorimetric analysis has been reviewed before (1, 4, 6, 6). Most of the applications have been in the preparation of permanent standards for use in determining substances whose properties are such as to prohibit their own use for comparison solutions. Thus the 1933 edition of the American Public Health Association’s “Standard Methods of Water Analysis” specifies such inorganic standards for the determination of color, ammonia, chlorine, iron, and silica. I n determining the specifications for such a series of standards, the general practice has been to match a known amount of the constituent to be determined with a solution containing one or more of the substances forming colored ions. The match has usually been made on the basis of visual comparison of the two systems. It frequently happens that solutions appearing to be matched with the eye yield quite different
spectral transmission curves. Then, if the two systems are observed under a source of illumination having a spectral energy distribution distinctly different from that of the source used for the original matching, the systems may not appear matched. In view of the present usage of such systems for colorimetric standards and of the possibility of applying them to some of the many new colorimetric methods, i t seemed worth while to make a more exhaustive study of the characteristics of various systems showing promise of possible usefulness in this direction. The present work was limited to elements 23 to 29, together with cerium, and was confined to compounds that seemed most likely to be of value. Many previous studies have been made on the absorption spectrum of solutions of these compounds. I n general, the objective of the work was different, and usually the methods of measurement employed were not capable of giving the quality of results now obtainable with photoelectric instruments.
Experimental Work MATERIALS.Wherever it was feasible, the salts t o be used were recrystallized at least twice from conductivity water, and the latter was used to make all solutions. Concentrations were determined by standard analytical methods. The best analytical practice was followed in preparing solutions of potassium permanganate and ceric sulfate, the latter being made from a double
I 4
k
2
0
$0
160
ad0
SO0
Si0
160
580
wave T e n g ~ h
600
620
610
660
Wuve L enqtb
680
FIGURE2. SPECTRAL TRANSMISSION CURVESFOR YELLOW SOLUTIONSCONTAINING SALTS OF VANADIUM,CHROMIUM, CERIUM,OR IRON
FIGURE1. SPECTRAL TRANSMISSION CURVESFOR RED AND PURPLESOLUTIONS CONTAINING SALTSOF COBALT OR MANGANESE
(The formula f o r the vanadium compound is uncertain.)
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INDUSTRIAL AND ENGINEERING CHEMISTRY
FIGURE3. SPECTRAL TRANSMISSION CURVESFOR GREEN SOLUTIONS CONTAINING SALTSOF NICKEL,COPPER,OR CHROMIUM
VOL. 8, NO. 6
FIGURE5. SPECTRAL TRANSMISSIOX CURVES FOR DIFFERENT CONCENTRATIONS OF TETRAMMINO NICKELSULFATEAND CHLOROPENTAMMINO COBALT CHLORIDE
ammonium sulfate. A c. P. grade of vanadium pentoxide, free of iron, was dissolved in concentrated sulfuric acid and then diluted in making the yellow solutions. The blue solutions of
vanadyl sulfate were made by suspending vanadium pentoxide in hydrochloric acid and bubbling sulfur dioxide or hydrogen sulfide through the solution to accomplish the reduction. The solution was then filtered, analyzed, and diluted. Ferric chloride was made by passing chlorine over hot iron wire and subliming the product. APPARATUS. The hotoelectric spectrophotometer, built by the General Electric &mpany, has been described by Hardy ( 2 ) . The method of use, together with advantages of the instrument, mas outlined in a previous paper by the authors ( 3 ) .
DATA. Since spectral transmission curves reveal the colorimetric characteristics of a system more quickly than tabular data, the graphical method has been used here, plotting per cent transmittancy as ordinates and wave length as abscissas. All data were calculated to a basis of 10.0-mm. cell thickness. Figures 1 to 5 illustrate the results obtained. In Figures 1 to 4 the curves are grouped according to the following hues: reds and purples, yellows, greens, and blues. Each curve is one of a series for different concentrations a t a given acidity, or the same concentration at different acidities. Figure 5 shows a limited selection of curves from two such series. Discussion
FIGURE4. SPECTRAL TRANSNISSION CURVESFOR BLUE CONTAINING SALTSOF COPPER,COBALT,NICKEL, SOLUTIONS OR VANAD1tJ.M
The curves as such require little comment. The systems selected for measurement and the range of acidities covered in the solutions seemed most likely to include those of value. Thus, interesting solutions, such as the alkali ferrates and cupric thiocyanate, were not considered worth-while possibilities, but potassium permanganate was included because of its unique characteristics. Someone may discover how to stabilize it.
NOVEMBER 15, 1936
ANALYTICAL EDITION
The pronounced effect of variations in acidity is shown for solutions such as cupric chloride and bromide and cobalt chloride. Presumably this depends upon the extent to which complex ions are formed. While certain other complexes were included, such as tetrammino cupric sulfate and chloropentammino cobaltic chloride, no attempt was made to exhaust the possibilities among the large number of complex compounds of cobalt or chromium. While a solution containing trivalent chromium is a desirable green, the system becomes somewhat dichromatic and the hue changes toward the violet on standing, with the ultimate attainment of equilibrium. At the higher acidities, the change is retarded. As any data obtained for ferric chloride merely corroborated bhose presented in a n earlier report, they are not included here.
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Summary Spectral transmission curves are presented for the visual region for aqueous solutions of ceric sulfate and of various salts of elements from atomic number 23 to 29. These include a considerable range of concentrations and acidities,
Literature Cited (1) ilrny, Am. J. Pharm., 104, 2 7 2 (1932). (2) Hardy, J. Optical Soc. Am., 25, 305 (1935). ENG.CHEM.,Anal. Ed., 7, 187 (1935). (3) Kasline and Mellon, IND. (4) Mellon, Proc. Indiana Acad. Sci., 32, 164 (1922). (5) Snell and Snell, “Colorimetric Methods of Analysis,” New York, D. Van Nostrand Co., 1936. (6) Yoe, “Photometric Analysis,” Vol. 1, New York, John Wiley & Sons, 1928. RECEIVED July 23, 1936. Abstracted from 8 portion of a dissertation submitted by C. T. Kasline to the graduate school of Purdue University in partial fulfillment of the requirements for the degree of doctor of philosophy.
Determination of Iodine and Bromine in the Presence of Each Other LADISLAUS SPITZER, Peterdi ucca 11, Budapest, Hungary
I
T I S WELL known that bromine is a more powerful oxidizing agent than iodine. Formic acid, for example, is not oxidized by iodine, as is apparent from the nature of Romijn’s method for determining formaldehyde (1) ; bromine, on the other hand, oxidizes not only the acid but also the formates to form carbon dioxide ( 2 ) as is indicated in the equations HCOOH Brz = COz 2 HBr (1) HCOONa Brz = COz HBr NaBr (2)
+
+
+
+
+
On the basis of these reactions and the well-known ability of bromine to displace iodine from a n iodide, it appeared possible to establish methods (1) for the determination of iodine and bromine in the presence of each other, and (2) for the determination of iodide in the presence of bromide. 1. Sodium formate is added to one aliquot portion of the solution of bromine and iodine; after the bromine has been reduced to bromide, the iodine is titrated with standard thiosulfate eolution. Potassium iodide is added to a second aliquot portion of the solution and the total iodine is titrated with thiosulfate solution. The quantity of bromine may be calculated from the difference between the two titrations. 2. The iodide in a mixture of bromides and iodides is liberated by adding bromine; the excess of bromine is reduced by adding sodium formate solution and the iodine is then titrated with thiosulfate.
Experimental IODIKE SOLUTION. Iodine (1.542 grams) was dissolved in 25 per cent potassium bromide solution and diluted to 1 liter with the potassium bromide solution. Ten milliliters of the solution were equivalent to 12.1 ml. of 0.01 1V sodium thiosulfate. BROMIKE SOLUTION. Bromine (1.052 grams) was made up to 1 liter with 25 per cent potassium bromide solution. Ten milliliters of the solution were equivalent to 13.1 ml. of 0.01 N sodium thiosulfate. POTASSIUM IODIDE.Potassium iodide (1.82 grams) was made up to 1 liter with potassium bromide solution (25 per cent). SODIVM FORMATE. A 5 per cent solution was prepared from formic acid that had been purified by distillation and from sodium hydroxide of reagent grade. The sodium formate was purified by recrystallization.
Procedure for Bromine in the Presence of Iodine From 0.25 to 0.35 gram of the substance to be tested is weighed and made up to 250 ml. in a 25 per cent potassium bromide solution. Two 25-ml. portions of this solution are placed in flasks and each portion is diluted to 100 ml. To one of the aliquots is added 1 to 1.5 ml. of the 5 per cent sodium formate solution. The flask is shaken vigorously and then allowed to stand for 10 minutes. The solution is then titrated with sodium thiosulfate solution, using starch indicator. The second aliquot is treated with 5 ml. of 10 per cent potassium iodide solution and then the amount of thiosulfate that is equivalent to the sum of the bromine and the iodine, is determined by titration.
The procedure was tested by preparing mixtures containing from 1 to 20 ml. of each of the halogen solutions. The amount of the sodium formate solution ranged from 0.5 to 2 ml. and the excess of sodium formate from 24 to 82 mg. Each mixture of the bromine and iodine solutions was diluted to 100 ml. A period of 10 minutes was allowed for the interaction of the bromine and the sodium formate in each case. From 2 to 4 ml. of 10 per cent potassium iodide were used prior to titrations for the sum of bromine and iodine. The results are summarized in Table I. TABLEI. DETERMINATION OF BROMINE AND IODIXE IN PRESEKCE O F E A C H OTHER Iodine Present Mg 15.35 23.02 15.35 23.02 7.67 7.67 1.53
.
Bromine Present
0.01 N Thiosulfate for Iz
Mg.
MZ .
10.47 5.23 5.23 1.04 15.70 20.94 10.47
12.10 18.10 12.05 18.15 6.05 6.00 1.20
THE
0.01
N Thiosulfate for Br2
Iodine Found
MI. 25.15 24.65 18.60 19.45 25.65 32.25 14.25
15.35 22.97 15.29 23.03 7.67 7.61 1.52
Mg.
Bromine Found
MMg 10.43 5.23 5.23 1.04 15.70 20.98 10.43
The p H ranged from 3.0 to 6.9, depending upon the amount of bromine present. This variation does not appear to have affected the accuracy of the results. The presence of sodium formate, even if in tenfold excess over the theoretical amount, does not affect the accuracy of the determination of iodine. The sodium formate must be pure.
