ANALYTICAL CHEMISTRY
1618 The optical density of neutral solutions of these proteins was measured, and the specific extinction coefficients, k , were calculated. These are found in column 2 of Table I. -4rsenic analyses on the coupled protein (column 3) led to the calculation of the moles of azo group per gram of arsaniloazo protein. These values are found in column 4 of Table I. If the specific extinction coefficients in column 2, Table I , are divided by 10.0 X 103,the molecular extinction coefficient for arsaniloazotyrosine, there are obtained the number of moles of azo group per gram of coupled protein based on optical measure ments. These values are found in column 5 of Table I. The number of azo groups calulated from the extinction coefficient for arsaniloazotyrosine are reasonably close to the number calculated from the direct arsenic analyses. Even where the agreement is least satisfactory, the two methods yield results which differ only by a few groups per mole of coupled protein. Such agreement would be adequate for almost all the n-ork done R-ith azoproteins. SU3I3IARY AND CONCLUSION
an extinction coefficient in very good agreement with that obtained Ji-ith the purified product. The molecular extinction coefficient of arsaniloazotyrosine has been measured. I t may be used to calculate the number of azo groups in any arsaniloazo protein for which the specific extinction coefficient has been measured. It is therefore possible to measure the number of coupled azo groups in any such protein by comparing the light absorption of the azo protein with that of the same diazotized amine coupled with an excess of tyrosine. LITERATURE CITED (1) Boyd, W.C., and Hooker, 5. J., J . Biol. Chem., 104, 329 (1934). (2) Eagle, H a r r y , a n d Vickers, Percy, Ibid., 114, 193 (1936). (3) Magnuson, H . J., a n d Watson, E. B . , IND. ENG.CHEM.,d s . 4 ~ . ED.,16, 339 (1944). (4) P a u i y , Herm, Z.physiol. Chem., 42, 508 (1904); 44, 159 (1905); 94,284 (1915). (5) Willard, H . H . , Merritt. L. L., a n d Deane, J. A., “ I n s t r u m e n t a l Methods of Analysis,” Sew r o r k , D. Van Nostrand Co., 1948. 24, 1952. Funds for this work were prol-ided by a grant from The Research Council of Rurgers University, the State Unirersity of S e w Jersey.
RECEIVED for review .4pril 2 , 1952. Accepted June
The quick method of preparing arsaniloazo p-cresol yields
lodometric Determination of Copper LOUIS MEITES Sterling Chemistry Laboratory, Yale Unicersity, i\-eu; Hacen, Conn.
HE conventional iodometric method for the determination of Tcopper(II) requires the addition of just enough potassium iodide to form solid copper(1) iodide and to maintain the liberated iodine in solution. The direct titration of tbis mixture with standard thiosulfate leads to low results, because an appreciable fraction of the free iodine is tenaciously adsorbed by the precipitate, which consequently becomes violet or brownish in color, thus tending to obwure the end point. Foote and Vance (2)
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intensity of color of the precipitate is considerably decreased, EO that the end point can be more precisely located. The aveiage error in the titrations of pure solutions reported by Foote and Vance was i0.06%. This error, the authors said, was due principally to errors in judging the end point, which may be much larger than this unless the analyst is familiar with the titration and takes great care to defer the addition of starch and thiocyanate until the latest possible moment. Furthermore, triiodide ion is known to oxidize thiocyanate slowly, and some elror may therefore result from the addition of thiocyanate even near the end of the titration. Scott ( 4 )devised a method for the performance of this titration in which enough iodide was added to retain the copper(1) in solution throughout the titration, presumably as CuI2-. This would eliminate all the errors generally asEociated with the separation of a solid phase during a titration, and because a clear colorless solution should be obtained a t the end point, it seemed that this method should be both more accurate and more convenient than that of Foote and Vance. The results of the experiments leported here confirm these expectations and indicate that the littleknown method of Scott is considerably superior to that of Foote and Vance. Its single disadvantage, that of increased cost of the reagents, is offset by the greater speed and accuracy Kith n hich the titration can be made. EXPERIaI ENTA L
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Figure 1. \‘ariation of Titration Error with Iodide Concentration a t End Point
recommended the addition of a small excess of ammonium (better potassium) thiocyanate just before the end point. This reacts a t the surface of the copper(1) iodide to form the less soluble copper(1) thiocyanate; as the latter has much less affinity for triiodide ion, the previously adsorbed iodine is released into the solution and may be titrated with the thiosulfate. -4s areeult, the
An 0.23686 F stock solution of copper(I1) sulfate was prepared by weight from very pure electrolytic copper. Weighed aliquots of this solution (usually about 20 grams) were taken for the titrations. The 0.1 AT sodium thiosulfate solution was prepared by dissolying sodium thiosulfate in boiling distilled water, in the bottle in which the solution was to be stored. hfter boiling for a few additional minutes, a trace of sodium carbonate was added and the solution was allowed to cool protected from dust. It was standardized against weighed portions of a nearly saturated solution of five times recrystallized potassium iodate, which had been prepared by weight. Replicate titrations usually agreed to = t O . O l % or better. Most of the titrations were made by adding 0.1 A’ sodium thiosulfate solution from a weight buret until the iodine color was nearly completely discharged, then finishing the titration rn ith
V O L U M E 24, NO. 10, O C T O B E R 1 9 5 2
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Figure 2.
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Variation of Titration Error with pH at End Point
0.004 iV thiosulfate, using a calibrated volume buret The dilute thiosulfate was prepared by weighing a portion of the stock solution into a calibrated volumetric flask and diluting with freshly boiled water: such solutions were a h ays discarded after 6 hours. I n nearly every case the end point could be located within about 2 drops of the dilute thiosulfate. As the total weight of 0.1 A' thiosulfate used was usually about 50 grams, this corresponds to an uncertainty of about =kO.Ol% .i volume buret alone was used for a small number of titrations in the presence of interferin elements. For these titrations the volume normality of the tfiosulfate a-as determined from its density and weight normality. Magnetic stirring was used throughout. -411 titrations were niade in Erlenmeyer flasks open to the air. The results were corrected for a trace of oxidizing impurity (0.0002%, calculated 3s iodate) in the potassium iodide used. No oxidizing or reducing impurities were found in the other chemicals used. RESULTS
Effect of Iodide Concentration. The results of titrations in ca. 0.1 F acetate buffers of pH 4.4 =t0.1 are shown in Figure 1, wi1ei.e the titration error is shon-n as a function of the iodide conceijtration a t the end point. The copper concentration a t the etiil point was about 0.03 F . At iodide concentrations below ahout 1.7 F , the copper(1) iodide either fails to redissolve completely when the potassium iodide is added or reprecipitates as thiosulfate solution is added. The solid copper(1) iodide thus pre-ent a t the end point adsorbs triiodide ion and leads to low results exactly as in the classical method. With concentrations of potassium iodide above about 3.5 F , a vihite microcrystalline piecipitate forms as thiosulfate is added. This could not be isol a k d for identification] as it is soluble in fresh 4 F potassium iodide anti washing with water is naturally out of the question. I t does not form if sodium iodide is used instead of potassium iodide. Figure 1 shows that the mean error under these conditions is at lout j ~ 0 . 0 2 ' 3 a~t all potassium iodide concentrations greater than that required to maintain the copper(1) in solution. All subsequent titrations were carried out in 2.2 F potassium iodide, corresponding to 35 grams per 100 ml. of solution. Effect of pH. Aliquots of the stock copper(I1) solution were treated with varying amounts of sulfuric acid, acetic acid, and ammonia to give solut,ions whose pH values a t the end points of
1619
the titrations varied from 0.4 to 8.3. The calculated volume of water to give a total volume of 175 ml. at the end point, and 65 grams of solid potassium iodide, were added, and the solution was titrated as above. The results, shown in Figure 2, demonstrate that an accuracy of &O.O2a/, can be secured a t any pH between 0.4 and 6.3. If the pH is higher than 6.3, the color of the solution begins to disappear temporarily after each small addition of thiosulfate when the titration is only 90 to 95% complete. The result shown in Figure 2 a t pH 7.7 was based on an end point which lasted for 5 minutes and which occurred nearly 7% too soon. If the pH is as high as 7.9, a dark green precipitate forms during the t'itration and iodine can no longer be detected in the supernatant liquid [which, if the pH is as high as 8.3, may even be colored blue by unreduced copper(II)] after only SOY0 of the ctllculated amount of thiosulfate has been added. Should such :I precipitate appear during a titration, it can be redissolved by the addition of a few drops of glacial acetic acid and the titration can be continued without ill effects. I n view of the frequently encountered statexuent that copper catalyzes the air-oxidation of iodide, it is somewhat surprising that no such effect could be detected in these esperinients. Even with 4 F potassium iodide and 0.25 F sulfuric acid, the mean error of four titrations was +0.01 i 0.03%, and the solutions remained colorless a t the end point for at least 10 minutes. These data led to the adoption of the follon-ing procedure, which was used for the determination of copper(I1: in the presence of ions of a number of ot'her metals. RECOMMENDED PROCEDURE
Evaporate a nitric acid solution containing not more than 150 mg. of copper nearly to dryness, add 3 ml. of concentrated sulfuric acid, and evaporate to incipient dryness. Repeat the second evaporation if the residue is large and nitrate may still be present. [Evaporation with sulfuric acid would give the s l o ~ ~ - lsoluble y anhydrous sulfates of iron(II1) and chromium(VI), n-hich might retain some copper and would have to be redissolved in hydrobromic acid. To avoid this, perchloric acid ma\ be used instead of sulfuric acid if chromium is knonn to be absent. Perchloric arid may be used even in the presence of chron~iumif the chromate produced is removed with barium chloride or if the result is corrected for the amount of chromium present.] Dissolve the residue in 25 ml. of water and add 1 gram of sodium acetate, 4 grams of sodium fluoride (unless tin and iron are known to be absent), and 25 grams of potassium iodide. Swirl for 10 to 15 seconds, then titrate with 0.05 to 0.1 -V thiosulfate. Starch indicator may be added near the end of the titration if desired: i t is unnecessary with pure solutions of copper salts. There is no need to standardize the thiosulfate against copper, for the titration is stoichiometrically exact. If the neight of copper present is greater than 150 mg., either the amount of potassium iodide or the concentration of the thiosulfate must be increased to prevent precipitation of copper(1) iodide. EFFECTS OF OTHER ELEMEhTS
Iron. I n the presence of the recommended concentration of fluoride, iron(II1) is nearly completely precipitated as ( S a , K ) 3 FeFa, and does not interfere with the titration. Four titrations of portions of the stock copper(I1) solution containing 0.15 to 0.3 gram of copper were made in 2.2 F potassium iodide a t pH 5.5 in the presence of 1 F sodium fluoride and 0.05 to 0 9 gram of iron(111). The mean error was 2 ~ 0 . 0 4 %and ~ the end points were stable for a t least 2 hours. I n this medium iron(I1) is a v e ~ y powerful reducing agent and must be completely okidized before the addition of iodide. Arsenic. Fiom the standard potentials of the arsenic (111-V) and iodide-triiodide couples selected by Latimer (S), and a t concentrations of iodide and triiodide a t the end point of 2 and 1 X 10-6 F , respectively [the latter figure corresponds to an error of 0.020, in the titration of 0.1 F copper(I1) with 0.1 S thiosulfate]. the relative concentrations of the two oxidation states of arsenic are given by the equation [As(V)I 4 x 10-7
ANALYTICAL CHEMISTRY
1620 Accordingly, if equal numbers of moles of arsenic (1’) and copper(11)were present in a mixture, the pH a t the end point must be a t least 4.9 if the titration error is to be less than 0.1%. The minimum pH value thus calculated actually corresponds to an arsenic to copper ratio seldom encountered in practical analyses. I n titrations of solutions containing about 400 mg. of copper and 100 mg. of arsenic, the titration error was less than 0.05% a t all pH values between 4.7 and 6.3. At lower pH values the titration error increased somewhat more rapidly than the equation given above predicts, reaching +1.5% a t pH 3.9. Bismuth. One frequently encounters the statement that bismuth(II1) does not interfere in the classical titration except by consuming iodide. This is not entirely true unless the bismuth concentration is very small (when its iodide consumption would be insignificant), for the basic bismuth salts formed during the adjustment of the pH are converted into Bi14-, whose deep orange color makes the location of the end point extremely difficult. The same interference is, of course, encountered in the present titration. Even 0.02 m M bismuth noticeably affected the precision with which the titration could be performed, and with as little as 2 mM bismuth (corresponding to 40 mg. of the element in the recommended procedure) the end point could only difficultly be located with 2 drops of 0.1 X thiosulfate. For precise results with either this or the classical method, the sample should contain less than about 10 mg. of bismuth per 100 ml. of solution a t the end point. Lead. Large amounts of lead may also interfere in the classical procedure owing to the formation of yellow lead( 11) iodide. This, however, is readily soluble in 2 F potassium iodide, giving a perfectly colorless solution which presumably contains the lead as the Pb14-- ion. By the recommended procedure, 300 mg. of copper(I1) in the presence of 1 gram of lead(I1) can be determined within &0.05%. Antimony. At or above p H 4.8, antimony(V) does not liberate any detectable amount of iodine. However, as in the method of Foote and Vance ( I ) , the solubility of antimonic acid in the titration medium is relatively low (ca. 25 mg. per 100 ml.), and a precipitate of this material tends to occlude appreciable amounts of copper, giving low results. There is no evident remedy for this difficulty. Tin. I n the steps involved in the preparation of the solution, tin(1V) is nearly completely precipitated as the hydrous oxide, which is known to retain small amounts of copper. A solution containing 350 mg. of copper and 175 mg. of tin was twice evapo-
rated to near dryness s i t h concentrated nitric acid and the mixture was then analyzed by the recommended procedure; the error was -0 05y0. The relatively small error may be due to partial re-solution of the hydrous stannic oxide by the fluoride added later. Silver and Mercury. Slightly more than 1 gram of silver(1) dissolves in 100 ml. of 2 F potassium iodide to give a colorless solution. Using a microburet, 5 mg. of copper(I1) were determined in the presence of 900 mg. of silver with an accuracy of f0.3C7,. The solubility of mercury(I1) in this medium is over 20 grams per 100 ml.; such a solution is only faintly yellow and the end point can easily be perceived. Other Elements. Barium, calcium, cadmium, magnesium, manganese, strontium, titanium, tungsten, and zinc do not interfere. [If manganese is known to be absent, the pH may be adjusted to 4.7 to 6.0 with ammonia and acptic acid in the usual way ( I ) . ] Chromium (111),cobalt, and nickel do not interfere unless their concentrations are PO high as to mask the end point. [The anhydrous chromium(II1) sulfate formed on evaporation with sulfuric acid does not retain detectable amounts of copper. ] Molybdenum and uranium interfere, but they Nould rarely be present in significant amounts. Vanadium will be largely converted to the +5 state: this oxidizes iodide and the +4 vanadium then formed is reoxidized by the air, leading to a large oveiconsumption of thiosulfate. I n a typical analysis of a solution containing 300 mg. of copper(I1) and 5 mg. of vanadium(V), the titration error was nearly +8%, only one fourth of which corresponds to the direct reduction of the vanadium. Fortunately the determination of copper in the presence of vanadium is rarely necessary. Perchlorate, phosphate, pyrophosphate, and sulfate do not interfere, and most other anions would be destroyed in the preparation of the solution. LITERATURE CITED (1)
Foote, H. TI’., and Vance, J. E., ISD ENG.CHEM.,ANAL.ED.,8, 119 (1936).
(2) Foote, H. TT., and Vance, J. E., J . Am. Chem. Soc., 5 7 , 8 4 5 (1935). (3) Latimer, TT. hI., “Oxidation States of the Elements and Their Potentials in Aqueous Solution,” S e w York, Prentice-Hall,
Inc., 1938. (4) Scott, W.TT., “Scott’s Standard Methods of Chemical Analysis,” edited by N. H. Furman, 5th ed., p. 368, New York, D van Kostrand Co., 1939. RECEIVED for review March 6, 1952. Accepted June 12, 1952. Contribution 1129 from the Department of Chemistry of Yale University
Dielectric Method of Determining Solvent Content of Dewaxing Mixes H. D. LEROSEN, L. V. WIKE’,
ARD
s. w. DENTON, T h e Texas Co., P o r t
IELECTRIC constant measurement has had many applicaD tions in analysis, best known of which is the determination of moisture content. Moisture has been so determined in oilseeds ( 7 ) ,in powders ( 2 , 5 ) ,in grain (IO), and in pulp (I). Using dielectric constant measurements, Freymann ( 3 ) identified and differentiated vegetable and mineral oils, Perumova (6) determined bivinyl and isobutylene in gases, and Schaafsma (8) determined both solvent and oil in the phases obtained when extracting a mineral oil with polar solvents. As a further application of dielectric constant measurement, the determination of dewaxing solvent in dewaxing solvent-waxy oil mixtures used in plant scale dewaxing operations 1% as undertaken, since a rapid method was desirable. On investigation, it was found that the dielectric reading of a dewaxing solvent composed of methyl ethyl ketone and toluene was lowered by the addition of a waxy oil; the lowering was proportional to the amount of 1%-axy oil added. 1
Present address, Hughes Aircraft Co., Culver City, Calif.
Arthur, Tex.
APPARATUS AND REAGENTS
Dielectric-type moisture meter made by Tagliabue Instruments Division, TVeston Electric Instrument Corp., or an equivalent instrument. Water bath, suitable for melting waxy oil-deIvaxing solvent blends. Pipet, 200-ml. Thermometer, approximate range 80” to 180” F. Dewaxing solvent, of like composition to that used in plant operations. Waxy oils, of stocks used in plant operations. CALIBRATION OF INSTRUMENTS
Several synthetic blends of dewaxing solvent and waxy oil stocks were prepared by volume; these covered the range of proportions used in plant operations. Both the dewaxing solvent and the waxy oil stocks were heated in preparing the blends. This necessitated the measurement of volumes a t temperatures considerably above that a t which volumetric ware is calibrated;
