V O L U M E 27, NO. 3, M A R C H 1 9 5 5 Table 111.
371
Effect of Certain Compounds on Determination of Elemental Sulfur
Compound Added n-Butyl sulfide Ethyl disulfide n-Butyl mercaptan
wt. of Compound Ilg. 200 40 20 20 8000
Petroleum ether
Table I\’.
Wt. Present
s, M g . Found
17.80 15.60 23.04 33.05 15.60
17 87 15.60 21.05 2G. 27 15.60
Analysis of Acetone Solutions of Elemental Sulfur Relative Error,
Sulfur, Ala. Taken
Found
7.80 10.10 15.60 15.60 15.60 15.60 23,40 31.20
7.79
70 0.1 0.1 0.5 0.0 0.3 0.3 0.4 0.0
10.11
15.53 1.5 . IiO 15.64 l5.5,j 23.49 31.20
cyanide to precipitate and become unavailable for reaction with the sulfur. Concentrations of water higher than 20% appeared to cause the rate of the reaction to decrease appreciably. Interferences. The effect of a number of substances on the proposed procedure was investigated. Table I11 shows that an excess of either a typical aliphatic sulfide or disulfide does not cause any alteration of the results. On the other hand, me!captans do interfere. This interference undoubtedly results from the reaction between the mercaptans and elemental sulfur which is catalyzed by the basic reagent. Metallic ions such as silver, mercury, or cadmium interfere with the titration by forming stable complexes with the cyanide reagent. Apparently the complesed cyanide ions do not readily react with the sulfur. Results. Table I V shows result< obtained by the proposed method of aliquots of acetone solutions containing known quantities of elemental sulfur. The average relative error of these determinations mas found t o be 0 . 2 7 . The standard deviation was 0 3%.
Average 0 , 2 LITERATURE CITED
cyanide in a solvent consisting of 80% by volume of isopropyl alcohol and 20y0by volume of water were found to be remarkably stable This is illustrated in Tahle 11. These data indicate that the alcoholic cyanide solutions are considerably more stable than simple aqueous solutions of sodium cyanide which are reported to decrease in formality by 0.3% per day (5). Effect of Water Concentration on Titrations. It was found desirable t o have between 15 and 20T by volume of water in the acetone solvent at the beginning of the titration. With smaller amounts of water, there was a tendency for sodium
(1) Am. SOC.Testing Materials, Philadelphia, P a . , ”Book of ASTLI Standards,” P a r t 6, p. 52, 1952. (2) Bartlett, J. K . , and Skoog. D. .I AKAI.. .,CHEM..26, 1008 (1954). (3) Castiglioni, h..Z.a n d . Chem., 91, 32 (1932). (4) Hardrnan, A. F., a n d Barhehenn. H. E., IKD. ENG. CHEW., i l x a ~ED.,7, 103 (1935). (5) Kolthoff, I. lI., a n d Stenger V. A.. ”Volumetric Analysia,” Vol. 11. pp. 282-3, Interscience. S e w York, 1947. (6) M a r k . G. L., and H a m i l t o n , J. 31.. IVD. E m . CHEW.,A K ~ L E D . , 14, 604 (1942). (7) Tuttle, J. B., “Analysii of R u b b e r , ” p. 58, Chemical C a t a l o g Co., Xew York, 1922.
RECEIVED for review September 7, 1954. Accepted November 15, 1954.
I r o n W Perchlorate as a Reductant in Glacial Acetic Acid 0.N. HINSVARK
and K. G. STONE
Kedzie Chemical Laboratory, Michigan State College, East Lansing, M i c h .
A reducing agent was required which could be used for nonaqueous titrations of oxidants. Iron(I1) perchlorate in glacial acetic acid was satisfactory for the determination of chromium trioxide and sodium permanganate in glacial acetic acid without the addition of water. .4n amperometric end point with two active electrodes was most suitable.
S
TUDIES still in progress in this laboratory led t o the investigation of iron(I1) perchlorate as an analytical reagent for the determination of oxidants in glacial acetic acid. Aqueous iron(I1) perchlorate has received limited attention as a reductant in the establishment of oxidation potentials of various cerium( IV) complexes ( 1 ) . Iron(I1) chloride has been used in glacial acetic acid (3))but its low solubility in this solvent limits its applicability. Iron(I1) pel chlorate is much more soluble; solutions in ewess of 0.1S n i t h respect t o iron(I1) are easily prepared. Using acetic acid solutions of iron( 11) perchlorate standardized against aqueous potassium dichromate, solutions of sodium . permanganate and chromium trioxide in glacial acetic acid can be analyzed M ithout the introduction of aqueous reagents. REAGENTS AXD APPARhTUS
Baker’s analyzed acetic acid was further purified by distillation away from chromium trioxide followed by a second distillation from potassium permanganate. The iron(I1) perchlorate hesa-
hydrate and perchloric arid ( T O % ) were obtained from the G. Frederick Smith Chemical Co. Baker’s analyzed chromium trioxide and Fisher Scientific Co. C.P. grade sodium permanganate were used as oxidants. A Fisher Elecdropode, sensitivity 0.025 pa. per scale division, equipped with 2-em. 18-gage platinum wire electrodes, was used for detection of the equivalence point ( 2 ) . The solution being titrated was stirred under a nitrogen atmosphere with a magnetic stirrer. The iron(I1) perchlorate solutions were prepared in the following manner: Acetic anhydride in slight excess over that necessary to react with the water present in the iron(I1) perchlorate was added to the acetic acid. After the acetic acid had been flushed with nitrogen, the approximate weight of iron(I1) perchlorate hexahydrate was added. This solution was left under a nitrogen atmosphere for a minimum of 2 hours but frequentlj. much longer. A4measured volume of this solution was added to a solution of 5 ml. of 85% phosphoric acid in 20 ml. of water The resultant solution was analyzed by titrating to the diphenylamine color change Ll-ith a solution of primary standard potassium dichromate ( 4 ) . .4 potentiometric titration showed the equivalence point to be coincident with the color change when the iron(I1) solution was titrated under these conditions. I n the experimental work the iron(I1) perchlorate solution9 were used for t h e analysis of acetic acid solutions of sodium permanganate and chromium trioxide. I n the permanganate studies an approsimate weight of sodium permanganate was added to the desired amount of acetic acid. Sodium permanganate was used in preference to the corresponding potassium salt because of its much greater solubility in this medium. The actual concentration of the solution prepared in this manner was found by titrating a weighed quantity of
ANALYTICAL CHEMISTRY
372 primary standard sodium oxalate in the usual manner. This titration was done in acetic acid as well as in aqueous media; in an aqueous medium acidified with 2 ml. of sulfuric acid in 25 ml. of water, it gave a better end point. Regardless of the medium, the fading of the permanganate color a t the equivalence point makes the end point somewhat indeterminate. By observing the persistence of the permanganate color for 45 seconds, the values for the sodium permanganate concentration were obtained. The solutions of chromium trioxide were prepared by adding the approximate weight of chromium trioxide to the desired quantity of acetic acid. The chromium trioxide was standardized by adding a measured volume of the acetic acid solution to an excess of 10% aqueous potassium iodide in a glass-stoppered flask. The reaction mixture was left in the dark for 30 minutes. The liberated iodine was then titrated to a starch end point with aqueous standard sodium thiosulfate. STABILITY OF IRON(I1) PERCHLORATE
When no precautions are taken, acetic acid solutions of iron( 11) perchlorate are slowly oxidized by air. This is to lie espected, because oxygen is much more soluble in glacial acetic acid than in water. Evidence for the relative stability of the iron(I1) solutions stored under nitrogen arid under air is presented in Table 1. I n order to minimize decomposition, it is desirable to store the solution urider an atmosphere of nitrogen. When passed through the solution being titrated, nitrogen minimizes air oxidation of the iron(I1) which would lead to high values for the concentration of oxidant. Utilizing nitrogen, the equivalence point is sharp and definite, yielding more accurate and reproducible results.
Table I. Days
Stability of Iron(1I) Perchlorate
Under Air Normality 0 0265 0 0258 0 0240
0 1 3
Table 11.
50 75 ..
100 125 150 175 200 a 50
0
0 026.5 0 0264 0 02G5
1
3
Sensitivity at Equivalence Point"
Potential, MV.
L-nder Kitrogen Days Normality
_Ai_
PA. I l I I . ' 1\11. 1 2 3 ij .i . .i 7 9
10 2 10.4 10 3
ml. of 0.0102N Fe(C104)z titrated with 25.20 nil. of 0.0194S Na1\1nO4.
DETECTION OF EQUIVALENCE POINT
Potentiometric measurements with platinum-calomel electrodes were too unstable for practical use. The solution being titrated was too highly colored under the conditions employed to permit the utilization of color indicators because of the deep red-brown color of the iron(II1) formed during the titration. However, an amperometric method m-ith two active electrodes was possible for detecting the equivalence point. To increase the sensitivity, 0.5 ml. of 70% perchloric acid was added to the solution and a potential of 150 mv. was applied across the electrodes. This value was obtained experimentally from the data given in Table 11, 150 mv. being the minimum potential giving the greatest sensitivity under the conditions employed. Sulfuric acid and 85% phosphoric acid were tried in lieu of perchloric acid, but the results were unsatisfactory. When sulfuric acid was used, a precipitate was formed, causing erratic results. The solution remained clear when phosphoric acid was used, but the titrations gave values too low compared with the standardization and the results were not reproducible. DETERMINATIONS WITH IRON(I1) PERCHLORATE
Sodium permanganate was determined by titrating a measured volume of the iron(I1) perchlorate solution Tit11 the permanganate
Table 111.
Determination of Sodium Permanganate
N of Fe(CIO4z
Days
0.1103 0.1103 0.1103 0.1103
N of NaMnOl from Fe(Clo4)~ Titration 0.0901 0.0900 0.0637 0.0636
0 0 1.5 1.5
Table IV. N of Fe(C10dz by KzCrzO:
A. of
NaMnO4 from NapCnO4 0.0897 0.0903 0.0635 0.0637
Determination of Chromium Trioxide N of CrOa from Fe(C10dz
Titration
N of CrOz from Iodonietric Determination
Titration of Standard Fe(C104)z Solution 0.0679 0.0679 0.0679 0.0710 0.0710
0.1087 0.1088 0.1087 0.0683 0.0684
0.1089 0.1089 0.1089 0.0684 0,0684
Titration of CrOa with Standard Fe(C104)z 0.0710 0.0710 0.0710 0.0710 0.0510 0.0510
0.0685 0.0684 0.0683 0.1087 0.1159 0.1158
0.0684 0.0684 0.0684 0,1089 0.1161 0,1161
solution or by adding excess iron(I1) to the permanganate solution and back-titrating the excess iron( 11) solution. This procedure was adopted in preference to the direct titration of permanganate with the iron( 11) solution, since the deconiposition of the sodium permanganate is accelerated by the addition of perchloric acid. A desired volume of the standardized iron( 11) perchlorate solution was pipetted into t,he titration beaker containing enough acetic acid to cover the electrodes xhile nitrogen was being p:issed through the solvent,. To this solution 0.5 ml. of 70% perchloric acid was added, and the titration was carried out with thr sodium permanganate solution of unknown concentration. The titration is conducted by adding the reagent rapidly a t first and then dropwise as the equivalence point is approached. The approach to the end point is indicated by the magnitude of the decrease in the galvanometer reading. In the course of the titration, the current flow reaches a maximum, then decreases, and finallj- falls off to zero a t the end point. With this value a calculation of the permanganate concentration is made. The results are presented in Table 111. The values dei,ived from the nonaqueous titration of permanganate with iron(I1) perchlorate solutions are in good agreement with those found from oxalate titrations within t,he limits of error of the methods. The two values obtained are indicative of the stahilit'y of the permanganate solutions. I n spite of the precautions taken for purifying the acetic acid, a significant amount of decomposition occurs in a relatively short time. Restandardization immediately prior to use is therefore necessary. Table IV presents the results found in the analysis of chromium trioxide solutions, These data show the feasibility of a direct titration of the chromium trioxide solution with one of iron(I1) perchlorate or by the reverse procedure-Le. , titrating a known concentration of iron(I1) with the chromium trioxide solution. With both titrations the results are in good agreement' n-ith those obtained iodometrically. The titrations were conducted in essentially the same nianner as in the permanganate analyses. The solution being titrated is acidified with 0.5 ml. of perchloric acid and the titrant is added with nitrogen flowing throughout the titration. When the iron( 11) solution is being titrated, the galvanometer behaves in the same manner as in the permanganate titration, the current flow passing through a maximum and proceeding to zero a t the end point. I n titrating the chromium trioxide solution, the galvanometer shows the current passing through a small maximum
.
V O L U M E 27, NO. 3, M A R C H 1 9 5 5
373
and falling ofl to zero as the end paint is reaohed. The next addition of the iron(II) solution oauses a large increase in the current flow.
’
(2) Stono. K. G.. and Scholten. H. G., (1952).
ANAL. CKE-M.,24, 671-4
(3) Tomioek, 0.. and Heyrovskl, A,. Collection Csechoslou. Chem. Comnmns., 15, 997-1020 (1950). (4) Willard, 11. IT.. and Furman, N. H.. “Elementary Quantitative Analysis,” 3rd ed.. p. 241, Van Nostrand, New York,
ACKNOWLEDGiMENT
It is a pleasure to thank the G. Frederick Smith Chemical Go. for the gift of iron(II) perchlorate hexahydrate.
1940.
RECEIVED for review September 11. 1954. Aoeepted November 24,
1954.
Abstracted from B portion 01 the thesis presented by 0. N. Hinavark in partial fulfillment of the requirements for the degree of doctor of philomphy and supported in w r t by a grant. NSF-281, by the National Science Foundrtion.
LITERATURE CITED
(1) Smith, G . F., “Cerate Oaidimetry,”’p 24, ti. Frederick Smith Chemical Go., Columbus. Ohio, 1942.
Measurement of Refractometric Dry Substance of Sucrose Solutions D. F. CHARLES
and
P. F. MEADS
waiian Sugar I;
wkott, Calif.
Accuracy or remaemmetric readings 1s impurrau~LU controlling purity and concentration of liquid sugar products. An evaluation has been made of the aouraey and precision of readings for two commercial refractometers. The standard deviation of the error of observation is 0.03% solids for the Bauseh and Lomh precision sugar refractometer and 0.07% solids for the Zeiss sugar refractometer. The study indicates a small error in the international s a l e of refraotive indices of sucmae solutions in the range from 50 to 75% sucrose. about 0.07yo solids at 66% solids. This is in qualitat i v e agreement with findings of others and points to the need for additional work to provide a sound basis for revision of the scale.
are employed widely in the sugnr industry to measure the dissolved solids content of both pure and impure sucrose solutions. Generally, a sugar industry lahoboratory may he expected to have several of these instruments, SomctimeS including a variety of modelr. Refractometers may be located a t plant operating stations or in routine control laboratories t o provide information for prompt process or product control, and they may find application in research laboratories for process study or other investigabions. The result of a refi’act,ometriomeasurement may he used directly as an expression of concentration in per cent solids by weight or it may he used as part of other laboratory procedures-for example, in purity determinations or an adjustment of solutions for color readings. Although the principal emphasis of the present discussion is in reference to the sugar industry, t.he data can be recalculated for use in other processing industries. Almost without exception, the industrial instruments in use in these industries in this country fall into the classification of critical angle refractometers. Refractometers which are now commercially availahlo give very satisfactory results, particularly for routine analytical work. On occasion, hon.ever, when more precise results are desired, certain questions of procedure and instrument accuracy arise. Some of these problems have been investigated recently in this hhoratory and this paper discusses the results. In particular, two factors have been of primary concern: the determinahion and correction of errom in sero point adjustment; and the relative accuracy of the Zeiss sugar refractometer (Figure 1) and the Bausch & Lomb precision refractometer (Figure 2). I n investigating the latter problem, discrepancies were encountered in the international scale for converting refractive indcx to per cent sucrose. Inasmuch as this experience parallels that of other observers, this paint has ako been investigated in some detail.
GENERAL BACKGROUND
The refractive index of 8r sugar solution is dependent on the concentration of the solution. Consequently, tabulations have been developed relating indes of refraction to per cent sucro~ein pure sucrose solutions (7) and the merose content of such solutions may he determined by reading the refractive index. In fact, mme instruments have n per cent sucrose scale mounted directly on the instruments. While this scale is correct for pure suorose only, the effect exerted upon refractive index by other sugars and soluble impurities is near enough to that for sucrose to make the reading on this sucrose scale extremely useful even for impure solutions. Refractometers are accordingly u t i h e d to measure the solids content of process liquors and of final liquid products in cane sugar refincries and beet, sugar factories. They are also used as
Figure 1. Zeiss sugar kefraetometei
