1111
V O L U M E 2 5 , NO. 7, J U L Y 1 9 5 3 gluconic and lactobionic, which can form lactones (see Table I). Attempts to measure carbonate salts were unsuccessful. I n addition t o the analysis of the organic salts reported, t h e method has been used for the determination of sodium chloride in dextran solutions. ACKNOWLEDGMENT
The authors are giateful to F. R. Stodola who supplied sample5 of many of the organic -alts reported.
1942. . 24, 300 (1952). ( 5 ) Pifer, C. IT., and Kollish, E. G., A N . ~ LCHEM., (6) Ibid., p. 519. ( 7 ) Ychrnall, JI., Pifer, C. IT., and Wollish, E. G., Ibid., 24, 144fi (1952).
(8)
Wiesenberger, E., Mikrochemie cer. Mibrochinz. d c t a , 30, 176-h (1942).
for rrvieiv February 11, 1953. Accepted .Ipril 2 7 , 1953. lien. tion of f i r m names or commercial products under zi proprietary nnnie or names of their manufacturer does not constitute an endorsement of sucli firms o i products by the U. 9 . Department of ;\griculture. RECEIVED
LITER-ITURE CITED
(1) Fritz, J . S., ASAL.(-’HEY.. 24,306 (1952). (2) Kleinzeller, A , . and Trim, A . R., -4rznZ!4.3t, 69,241 [I944
(3) Kunin, R., and Myers, R. J., “Ion Exchange Resins,” p. 127, Kew York, John Wiley 8- Sons, 1950. (4 1 Siederl, J. B., and Siederl, J’., “Organic Quantitative JIicroanalysis,” 2nd ed., pp. 62-5, Sew York, John Wiley & Soils,
1,
Reduction of Nitroguanidine and Its Derivatives by Titanous Chloride PAUL D. STERNGLANZ, RCTII C. THOMPSON’, ~ N DF . i L T E R L. SIVELL Chemical Research Dicision, Laboratory of .-idr.anced Research, Remington Rand, Inc., South Sorwalk, Conn. 7m c m and Hiilbert’s titanous chloride reduction method
h ( 2 ) for determining the nitro group fails when applied to
nitroguaiiidine and its derivatives, which belong to a class of organic compounds called nitroammonocarbonic acids ( I O ) . For example, Kouha and coauthors ( 6 ) obtained erratic results when they tried t o reduce nitroguanidine by this method. They were able t o analyze the compound satisfactorily by prolonging the boiling period and carrying out the reduction in a more concentrated acid solution. I n their modification, however. four instead of the theoretical six equivalents of titanous ions react with the nitro group. Zimmerman and Lieber ( I O ) then applied Kouba’s modification ( 6 ) to several nitroammonocarbonic acids. They found that nitroguanidine could not lie analyzed with any precision. and that nitroaniinoguaiiidine reacted with only 3.4 equivalents of titanous ions instead of the esprcted 4. However, xhen they modified Kouba’s method ( 5 ) 1))- adding bivalent iron as a reducing aid. they succtwicd in reducing the nitroammonocar1)onic acids completely, with a consumption of six equivalents of trivalent titanium. Since the amount of bivalent iron needed for each type of compound had to be empirically determined. they conceded that the method is of limited value. The existing titanous chloride reduction methods for nitroammonocarbonic acids (5, 6, I O ) are unsatisfactory; either they cannot be duplicated easilj- by other investigators or they are definitely empirical in nature (6, I O ) . .I study was made. therefow, to find a modification which does not have such limitations. In the proposed method the reduction proceeds rapidly a t room temperature with the consumption of six equivalents of trivalent titanium. The essential conditions are a 20070 miriim u m escess of titanous chloride and a weakly arid. buffered mrtliuni. REAGENTS
An approximately 0.3 S titanous chloride solution was prepired and stored as described by Siggia (8). Its strength was checked against that of the ferric alum solution before each series of experiments was run. Periodically the titanous solution was standardized againEt pure p-nitroaniline by the described method. An approximately 0.1 S ferric alum solution was used for back titration. The Reinhardt-Zimmermann procedure and the method described by Siggia (8) mere used interchangeably for standardizing the ferric solution. Buffer: potassium citrate, 800 grams per liter of aqueous rolution. Indicator: anlmoniuni thiocj-anate, 250 grams per liter of aqueous solution. . Tank carbon dioxide was used for conducting the titrations in :in inert, atmosphere. Contaminating osygen was removed by 1
Present address Wesleyan University, Sliddletown Conn.
passing the carbon dioxide through a 250-ml. gas washing hottlc. filled with vanadous sulfate solution ( 7 ) and through two 125-ml. gas nashing bottles filled with titanous chloride solution. IPP4R4TUS
A 250-ml. flat-bottom flask with outer 24/40 ground-glass joint was used. The inner joint member was sealed a t the top except for a hole, and a ring seal for a gas introduction tube. The hole served as a gas vent and was made large enough so that a funnel or buret tip could be inserted. The gas introduction tube served as a carbon dioxide inlet and reached to within 1 inch of the bottom of the flask. A4magnetic stirrer was used PROCEDURE
Take a sample 1.2 to 2.4 meq. in size [ I meq. = gram niolecular weight/( 1000 X 6 X number of nitro groups)]. Introduce the sample into the flask and dissolve in 10 ml. of 10qo acetic acid, heating on a steam bath, if necessary. For espediency, dissolve the two alkyl nitroguanidines in glacial acetic acid before diluting with water. Displace the air in the flask bv means of powdered dry ice (about 10 ml.) and purified tank carbon dioxide. The purpose of adding dry ice is to ensure faster displacement of air. When the dry ice has disappeared, bring the solution to room temperature; add 10 ml. of the potassium citrate buffer, followed by a measured excess of titanous chloride (200T0 minimum). Stir the solution approximately 3 minutes; add 20 ml. of hydrochloric acid (1 to 1) and 5 ml. ot ammonium thiocyanate solution, in the order given. Back titrate the excess titanou. chloride with the standard ferric :iluni solution. DISCUSSION
.In attempt W R S made first to reduce nitroanimonocarlionic acids by Kouba’s method (5) in a boiling titanous chloride-hydrochloric acid solution. LOK and erratic results were obtained, in agreement with those reported by Zimmerman and Lieher ( I O ) . It was reasonalile to assume that the reduction might go to completion if the reduction potential of tiivalent titanium \vere increased. One may to increase the reducing power of titanium is l)), raising the p H ( 3 ) . This effect has been clearly demonstrated on certain nitro compounds which Kolthoff and Ro1)inson ( 4 ) and Butts and coworkers ( I ) were investigating. -4 rapid, sisequivalent reduction a t room temperature occurred \Then the compounds were dissolved in a Imffered, weakly acid medium: in a mineral acid medium, the compounds would have required a t least a IO-minute boiling treatment for complete reduction. The rapid method was accordingly applied in this laboratory t o the nitroammonocarbonic acids. hgain, low7 and erratic results were obtained. Attrmpts to make the reductions complete
1112
ANALYTICAL CHEMISTRY
by boiling were abandoned when it was found that the blanks were high and not reproducible. The test results led to the opinion that the moderate excess of titanous chloride required in Kolthoff and Robinson's method ( 4 ) is insufficient for the refractory nitroammonocarbonic acids. Esperiments were made, therefore, to study whether complete reductions could be effected hy increasing the excess of titanous chloride. From the results (Table I ) , it was concluded that an excess of more than 100% titanous chloride is required for the complete reduction of the nitroammonocarbonic acids; 200'34 waa found to be a safe lower limit. Increasing the excess of titanms chloride to 700yo did not affect the reductions. This wide tolerance is an advantage over Zimmerman and Lieber's method ( I O ) , where either too high or too low results are obtained unless the amount of reducing aid is kept within the narrow range of its upper and lower limit.
Table 11. Reduction of Nitroammonocarbonic Acids by Proposed Method Approximate Excess Compound TiClr, % .V-n-Buty1-)V'-nitroguanidinea 200 200 200 500 500 500 S-Ethyl-SI-nitroguanidinea
Sitroguanidine b
200
Kitroaminoguanidine c
200 200 500 500 500 200 SO0
Table I. Effect of Varying the Per Cent Excess of Titanium Chloride Approximate Excess TiCIa, %
Compound
Nitro Compound Found, %" (Average of Duplicate -4nalyses) 89.4 95.0 97.OC
200 500 500 500
S i t r o Compound Sought, Found gram Gram % 0.0600 0.0577 96.2 0,0600 0.0586 97.7 0,0600 0.0582 97.0 0.0300 0.0291 97.0 0.0300 0.0292 97.3 0.0300 0,0289 96.2 0.0500 0.0482 96.4 0.0500 0.0485 97.0 0,0500 0,0482 96.4 0.0250 0.0240 96.0 0.0250 0.0242 96.8 0.0250 0.0242 96.8 0 0400 0.0391 97 8 0 0400 0.0391 97 8 0 0400 0.0394 98 4 0 0200 0.0197 98 3 0 0200 0.0196 97 8 0 0200 0.0196 97 8 0.0400 0.0385 96.2 0.0341 0.0400 97.8 0.0388 0.0400 97.0 0.0200 0.0194 97.2 0.0200 0.0193 96.5 0.0200 0.0192 96.0
Supplied through the courtesy of E . Lieher, Illinois Institute of Technology, Chicago, Ill. b Supplied by Matheson Co., Inc.; purified b r recrystallizing from water. c Supplied by Eastman Kodak Co.: purified b y recrystallizing from water. Q
96.8C 96.2
Sitroaminoguanidine 6
30 100
Kitrophenylacetic acidb
200 500 700 20 200 700
88.6
94.6 9i.OC 96.6C
97.2 99.1 99.5 99.5
a A reduction equivalent of exactly 6 titanous ions prr nitro group corresponds t o a result of 100%. b Described in Table 11. c Average of triplicate analyses.
Sitrophenylacetic acid, a nitro compound which does not belong to the class of nitroammonocarbonic acids, was included in Table I as an additional check on the proposed method. It can be reduced with 20% excess titanous chloride, using a method similar to Kolthoff and Robinson's ( 4 ) , or with 200% excess, using the proposed method. The proposed modification was then applied to other nitroammonocarbonic acids (Table 11). Six equivalent reduction took place in agreement with the equation
RNOz
+ 6Ti+++ + 6 H r --+ RNHz + 6 T i + + + ++ 2H20
It was found that the same reduction conditions are applicable to all the compounds. This is another advantage which the proposed procedure has over that of Zimmerman and Lieher ( I O ) , in which the amount of reducing aid necessary to complete the reaction is specific and changes in an unpredictable manner for different types of nitroammonocarbonic acids. Whereas the precision of the method is satisfactory (0.5% average standard deviation), all the results for nitroammonocarhonic acids have the tendency to be low (average 2.9%). (For ordinaiy nitro compounds the accuracy is of the same order as that which is obtained by the standard method. Note nitrophenylacetic acid, Table I.) This may be due to a slight hydrolysis of the compounds. As an example the hldrolyiis of nitroguaiiidine is given
YHN02
C=KH
Potassium citrate was added to raise the pH, to buffei the solution, and to complex the titanous and titanic ions u hich otherwise would be insoluble in the pH range used. This particular citrate salt was chosen because of its great water solubility. The amount used produced a negligible blank. The volume of titanous chloride which provides the required 200% excess (about 20 ml. for 0.3 S solution) lowers the pH of the sample solution from its initial value of about 6-i.e., pH of the mixture of 10% acetic acid and potassium citrate buffer-to about pH 4.5, If, for example, the normalitv of the titanous solution were only about 0.2 -V, about 30 ml. instead of 20 ml. would be required to provide the minimum 200% excess. The additional 10 ml. of titanous chloride solution nith the proportionally greater amount of hydrochloric acid would cause the pH to he decreased to about 3.5. This vaiiation in final pH (3.5 to 4.5) did not affect the results. Experiments proved that the 3-minute stirring time is not critical: equally good results were obtained when the stiiring continued for 15 minutes. Elimination of the boiling period used in pievious modifications (6,10) is an advantage since it has been found that variations in the boiling time affect the analysis ( 4 , 9 ) . The purity of the compounds investigated was established by melting point determination. bv microanalyses for carbon. hydrogen, and Dumas nitrogen, and by titration (Table 111). ACKNOW LEDGRIENT
The authors wish to express their appreciation to Frederic R. Benson, Chemical Research Division, Remington Rand, Inc.,
Table 111. Analyses of Compounds Compound S-n-Butyl-SI-nitroguanidine
l l . P . >' C . 8fi
Calculated. % C H
S
34.98 37 49 7,jj
X
Found, % C H
34.96 37.32 7.28
+ H2O +CO? + 2SH3 + KzO
'\ 2"
It is also possible that a side reaction occurs to a limited degree leading to a hydrouylamine-type reduction product which requires only four equivalents of trivalent titanium (10).
a
152 7.73 53.04 3.89 7.73 53.26 3.58 99.9% purity found by titration with standard sodium hydroxide solu-
tion.
V O L U M E 25, NO. 7, J U L Y 1 9 5 3
1113
for his helpful criticism. hppreciation is also extended to (4) Kolthoff. I. AI., and Robinson, C., Ibid., 45, 169 (1926). R. c., E~~~~~~ i ~ lllinois b ~ ~~ , ~of ~ ~~ ~ h,--hicago. ~ ~ ~ iI ~] I . ,~ ~~(5) Kouba, ~ ,. D. , L., ~ Kicklighter, ~ ~ and Becker, w. W., -&SAL. CHEM., 20, 948 (1948). for his kindness in submitting several compounds required for (6) J I ~ K -4, ~ F,, ~ ,them, fievs,,51, 304-5 (1962). this investigation. (7) Meites, Louis, and RIeites, Thelma, ASAL. CHEM.,20, 984 (1948). LITERATURE CITED
(1) Butts, P. G., Neikle, TV. J., Shovers, John, Kouba, D. L., and
Becker, W. W., AXAL.CHEX.,20, 947 (1948). (2) Knecht, E., and Hibbert, E., “New Reduction Methods in Volu-
metric Analysis,” New York, Longmans, Green and Co., 1925. (3) Kolthoff, I. hI.. Rec. trap. chim., 43, 7 6 8 (1924).
(8) Siggia, Sydney, “Quantitative Organic Analysis via Functional Groups,” p. 8 2 , New York, John Wiley 8: Sons, 1949. (9) Van Duin, C. F., Chem. Weekblad, 16, 1111 (1919). (10) Zimmerman, R. P., and Lieber, E., - 4 s ~CHEM., ~. 22, 1151 (1950).
R E C E I V Efor O review NorenibPr 6, 1952.
Accepted JIarch 12, 1953.
Effect of Cuprous Iodide on the lodometric Determination of Iron in Presence of Sulfate EDWARD W. K41MMOCK’ AND E R N E S T H. S W I F T Gates and Crellin Laboratories of Chemistry, California Institute of Technology, Pasadena, Calif. determination of ferric iron can be made in Tthe iodometric presence of high concentrations of sulfate if a t least 200 HE
mg. of cuprous iodide are added dissolved in the potassium iodide. In the iodometric determination of copper the necessity for the use of a solu1)le thiocyanate can be eliminated by the addition of an equimolal or greater quantity of ferrous iron. The presenee of moderate concentrations of sulfate cause the iodometric determination of iron to be less satisfactory in that higher concentration> of acid and of iodide are required, and the end points are less permanent ( 5 , 6). I n a study of the iodometric determination of ropper in sulfate-hydrogen sulfate buffers ( 4 ) , the authors observtld that an additional quantity of iodine -toichiometrically equivalent to any ferric iron present was liberated and that stable end points were obtained. Hahn and Windish (3) have reported the catalytic effect of small amounts of cuprous iodide on the ferric iron-iodide reaction, and during the progress of the work reported below) Hahn ( 2 ) proposed the use of cuprous iodide in the iodometric determination of iron, but in neutral solutions and in iodide concentrations sufficiently high, at least a t the beginning of the titration, to dissolve the cuprous iodide precipitate. Recently a procedure was described by Brasted ( 1 ) for the iodometric codetermination of copper and iron in which use is made of nitric arid ~olutionsand of sulfamic acid to eliminate nitrous acid. The present work was undertaken in order to investigate the effect of copper on the iodometric determination of iron in sulfate-hydrogen sulfate solutions similar to those used for the iodometric determination of copper. I n the course of this work additional observations were made on the effect of iron on the iodometric determination of copper. R E A G E h T S AND S O L U T I O N S
A copper solution, approximately 0.1 F (volume formal), was prepared from reagent grade copper sulfate pentahydrate, and was standardized against a standard thiosulfate solution by the procedure of Hammock and Swift (4). -4ferric chloride solution, approximately 0.1 F in ferric chloride and 0.1 F in hydrochloric acid, v,-as prepared from ferric chloride hexahydrate and hydrochloric acid. This solution was tested for ferrous iron and standardized against the same standard thiosulfate by the procedure outlined by Swift (6). The thiosulfate solution was prepared and standardized against Bureau of Standards primary standard potassium dichromate. The normality of the thiosulfate was checked at different times throughout the qeries of experiments, and no noticeable change was detected. -4hydrogen sulfate-sulfate buffer solution 1.80 F in ammonium sulfate and 0.060 F in sulfuric acid was prepared; the p H of this solution waa found by means of a glass electrode pH meter to be approximately 2 . 1
Present addresr. John Rluir College, Pasadena, Cahf.
Cuprous iodide was prepared by m i h g cupric sulfate with potassium iodide in a hydrogen sulfate-sulfate buffered solution, removing the excess iodine with thiosulfate, Jvashing the precipitate, dissolving the cuprous iodide in a saturated solution of potasqium iodide. removing all traces of iodine, and then reprecipitating the cuprous iodide by adding Lvater. The process was repeated if a portion of the cuprous iodide gave an iodine color with starch when dissolved in a solution of potassium iodide. The cuprous iodide thus prepared was washed with water, ethyl alcohol, and ethyl ether, then dried under a vacuum desiccator, and stored in a dark glass-stoppered bottle. At no time during this investigation did the cuprous iodide show an iodine color when dissolved in potassium iodide. (Subsequently the cuprous iodide did give an iodine color after it had inadvertently been heated to approximately 80” C.) PROCEDURE
Appropriate volumes of the copper and ferric solutions were pipetted into a 250-ml. “iodine flask,” 50 ml. of the acid sulfatesulfate buffered solution were added, the volume was made up to 100 ml., and 5 grams of potassium iodide dissolved in 5 to 10 ml. of water were added. The resulting mixture was titrated with the standard thiosulfate solution until the iodine color was indistinct, starch indicator solution vias added, and the titration was continued just to the disappearance of the starch-iodine color; this is designated the iodide end point. hpprouimately 5 grams of potassium thiocyanate were added, and if the starch-iodine color reappeared, the titration was continued to the disappearance of this color; this is designated the thiocyanate end point. Calibrated burets and pipets were used. The titration values shown in the tables are the result of a t least duplicate measurements; where the deviations are not shown, these did not exceed 0.1%. T I T R A T I O N O F C U P R I C 4ND F E R R I C SOLUTIONS
A series of titrations was made to determine the effect of varying the ratio of ferric iron and cupric copper in the iodometric determination of these substances. Hammock and Swift ( 4 ) reported that the presence of 60 mg. of iron in the iodometric determination of approximately 300 mg. of copper did not affect the
Table I. T i t r a t i o n s of Solutions w i t h Various C o n c e n t r a t i o n s of Iron(II1) a n d Copper(I1) Thiosulfate Used, MI.
Thiosulfate Calculated,
Cu(I1) ’ Fe(II1) end point end pointa MI. 4 185 0 000 43.71 3 765 0 413; 43.65 3 347 0 827 43.61 2 510 1 6.54 43.51 1 673 2 482 43 40 43.31 0 8365 3 309 0 418 3 732 43.26 0 000 4 136 43.23 a Average of 2 or more titrations. Maximum deviation was of t h e order of 2 parts per thousand. b A barely perceptible pink color, presumably the ferric thiocyanate complex, was visible after removing t h e starch-iodine color. This color increased in intensity a s t h e quantity of iron was increased.
