• Abstract and Figures
• Public Full-text

THE CERIOMETRIC DETERMINATION OF HYDRAZINE AND ALUMINUM oy CLIFFORD JAMES DERNBACH A THESIS submitted to the OREGON STATE COLLEGE in partial fulfillment of the requirements for the degree of DOCTOR OF PHILOSOPHY A3?ROTED: Redacted for Privacy Profegsor of Chentetry Redacted for Privacy Eeail of De nt of Chenlstry Redacted for Privacy Chatrman of Sohool, 0r Corurlttee Redacted for Privacy Chalrnan of State College Oraduate Conaoll ACKNOLEDGEMENTS The writer wishes to express his utmost thanks to Dr . J. P. Mehl ig for his constructive suggestions throughout this work . Thanks are also given to Dr . E. c. Gilbert for informing the writer of the existence of a published article which had a bearing on part of this work. TABLE OF CONTENTS Page I. THE DETERMINATION OF HYDRAZINE 1 PREPARATION OF SOLUTIONS 3 PROCEDURE 3 EFFECT OF AiiKAI·INITY 6 EFFECT OF EXCESS POTASSIUM FERRICYANIDE 6 EFFECT OF T!ME 7 EFFECT OF FINAL ACIDITY 8 RECOMMENDED PROCEDURE 8 RESULTS 9 DISCUSSION 11 S~Y 11 A NOTE ON THE DETERMINATION OF PHENYLHYDRAZINE 12 BIBLIOGRAPHY 13 LIST OF TABLES TABLE I. 5 TABLE II. 6 TABLE III. 7 TABLE IV. 10 TABLE V. 10 ii Page II. THE DET:ERMINATION OF ALUMINUM 14 PREPARATION OF SOLUTIONS 15 PROCEDURE WITH CERIC SULFATE 16 PROCEDURE 19 DETERMINATION OF UNKNOWNS 23 DISCUSSION 26 SUMMARY 28 BIBLIOGRAPHY 30 LIST OF TABLES TABLE I. 18 TABLE II. 19 TABLE III. 22 TABLE IV. 23 TABLE V. 24 TABLE VI. 25 TABLE VII. 26 1 flm CERIOMETRIC DETERMINATION OF HYDRAZINE AND ALUMINUM I THE DETERMINATION OF HYDRAZINE The use of eerie sulfate as a standard oxidant in titrimetry has become more and more widespread in the last few years, due, in part, to the, availability of an increas­ ing number of good reversible redox indicators (8). The many advantages of this reagent, among which are: its extreme stability, its simplicity of chemical change, and its use in hydrochloric acid solution, make it far superior to potassium permanganate, and, in many respects, to potassium dichromate. There are many different methods in the literature for the determination of hydrazine. Ray and Sen (9) were the first to determine the quantitative relationship between hydrazine and alkaline ferricyanide solutions. The reaction proceeds as follows: N 2H4 + 4 K3Fe(CN )6 + 4 KOH ---+ N2 + 4 :K4Fe(CN )6 +- 4 H20 They made this reaction the basis for a gasometric method for the determination of hydrazine. The reaction was found to take place without the formation of ammonia as is the case when many oxidizing ag ents react with hydrazine. Cuy and Bray (5) in a further study showed that the oxygen error is negligible if the alkali is added after the ferri~yanide. They determined conclusively that in 2 alkaline solution it is the atmospheric oxidation of the hydrazine that produces a weakening of _the solution and not the decomposition of the hydrazine. They also sug­ gested that a titrimetrio method for hydrazine could be used, in which excess of standard potassium ferrioyanide solution is added to the alkaline hydrazine solution and the excess ferrioyanide determined iodometrioally. However , since several more rapid methods were available they did not recommend this procedure for the determination of hydrazine. Benrath and Ruland (2) state that hydrazine is oxidized by eerie sulfate to nitrogen and ammonia according to the follo ing equation: 2 Ce(S04)2 + 2 N2fl4 ---;a.. N2 + (N~)2S04 + Ce2(S04)3 However, Lang (7) determined hydrazine and hydrazoio acid together by adding an excess of standard eerie sulfate solution, treating the excess with an excess of standard arsenious acid, and completing the titration with stand­ ard eerie sulfate. The hydra~ine was then determined in a separate sample by the iodine cyanide procedure. The purpose of the present work was to develop a simple method for the determination of hydrazine by making use of potassium ferricyanide and eerie sulfate. The method consists in the oxidation of hydrazine with an alkaline ferricyanide solution, and the titration with a standard eerie sulfate solution of the ferrocyanide 3 which is produced. Potassium ferrocyanide is readily titrated with standard eerie sulfate, either potent­ iometrically (I, 6 , 11) or with indicators (3, 12). PREPARATION OF SOLUTIONS Hydrazine Sulfate . An approximately 0.1 N solution was prepared by dissolving about 3.3 grams per liter. This solution was standardized by the well known iodate method of Bray and Cuy (4). Potassium Ferricyanide. A 0.5 M solution. Ceria sulfate. An approximately 0.1 N solution was prepared by dissolving 80 grams of eerie sulfate in 500 ml . of ~.0 N sulfuric aoid and diluting to one liter. This solution was standardized against pure potassium ferrocyanide follo ving the directions of Furman and Evans (6) Sodium Hydroxide. A 6 M solution. Hydroo~lorio Acid. A 6 M solution. PROCEDURE The procedure followed consisted of adding to a measured volume of hydrazine sulfate solution an excess of potassium ferricyanide, making alkaline with sodium hydroxide, shaking for one-half minute, and letting stand I 4 for a definite time. The mixture was then acidified with hydrochloric acid, diluted to 100 ml., and titrated with standard eerie sulfate solution. The endpoint was determined by the visual method which Furman and Evans (6} showed coincided with the potentiometric endpoint. There is sufficient ferric iron present in the eerie ammonium sulfate, and most other sources of eerie ions, so that colloidal ferric ferrocyanide forms during the titration of the ferrocyanide. At the endpoint the green color, caused by the colloidal ferric ferrooyanide, sharply disappears and the solution has the brownish color of the ferricyanide ions. The writer prefers to add a few drops of 0.5 M ferric chloride near the end of the titration to give a sharper endpoint, espcially in the presence of excess ferricyanide ions. If the ferric chloride is added too soon, a blue precipitate of ferric ferrocyanide is formed and thus some ferrooyanide is lost. Diphenyl­ amine may also be used as indicator (3), but the solution must be diluted to about 0.01 N ferrocyanide concentration before titration with eerie sulfate. The writer found some difficulty in obtaining a sharp en~point with diphenylamine in the presence of excess ferricyanide. Hexanitrato ammonium aerate may also be used as standard oxidant as indicated by Smith, Sullivan, and Fr.ank (10). However, since there is not enough iron present in this 5 highly purified substance, it will be necessary to add a few drops of ferric chloride near the endpoint in order to form the green colloidal ferric ferrocyanide. EFFECT OF AI·KALINITY The effect of alkalinity on the completeness of the oxidation of the hydrazine was studied. As seen from the results of Table I, a very low alkalinity will suffice for the complete oxidation, and a high alkalinity up to 3 M sodium hydroxide, yields good results. Table I Effect of Alkalinity Volume of 0.5 14 K3Fe t(.CN)6 added- 10 ml. Time being 2 minutes. Total volume being 100 ml. No. HYdraz ine NaOH HCl Ce(SOt)2 Normality Sulfate 6 M 2M 0.105 N Hydrazine ml . ml. ml. ml. 1 19.84 eo 40 19.98 0.1063 2 19.84 15 35 19.96 0.1062 3 19.84 10 30 19.98 0.1063 4 19.84 7 27 19.98 0.1063 5 19.84 4 24 19.96 0.1062 6 19.84 2 22 19.98 0.1063 The normality of the bydrazine sulfate solution obtained by the Bray and Cuy iodate method (4) was found to be 0.1062 N. The procedure used is as follows: To 6 10 ml. of 6 N sulfuric acid in a ground glass stoppered flask add a measured amount of 0.1 N iodic acid or of potassium iodate. 30 to 50 percent in excess of that needed is used, then add the hydrazine. After five minutes add excess potassium iodide and titrate with standard thiosulfate. EFFECT OF EXCESS POTASSIUM FERRICYANIDE Theoretically 20 ml. of 0.1 N hydrazine sulfate solution should, require 4 ml. of 0.5 M potassium ferri­ cyanide for complete oxidation. Table II shows results obtained using varying amounts of ferrieyanide, other conditions remaining constant. Table II Effeot of Ferrioyanide Concentration 10 ml. of 6 M NaOH added. Time 2 minutes. 30 ml. of 6 N HCl added. Final volume 100 ml. No. HYdrazine K3Fe(CN)3 Ce(S04)i Normality sulfate 0.5 M 0.1054 of ml. ml. ml. Hydrazine 1 19.84 5 19.12 0.1016 2 19.84 6 19.80 0.1052 3 19.84 7 19.96 0.1062 4 19.84 8 19.98 0.1063 6 19.84 10 19.98 0.1063 6 19.84 15 ----- -----­ 7 In numbers 1 and 2, containing 1 and 2 ml. more than the theoretical amount of potassium ferricyanide, the oxidation was incomplete in the two-minute interval. Even after a five-minute interval, on a run containing 6 ml. - of ferricyanide solution, the oxidation was incomplete. In number 6 the color change was not sufficiently sharp to indicate the endpoint. Too large an excess of ferrioyanide imparts a dark red-brown coloration, and the transition from green to red-brown is not distinct enough to serve as the endpoint. However , in a potentiometric titration, this excess would not interfere with the endpoint. EFFECT OF T!ME Table III shows the effect of time on the reaction. Table III Effect of Time 10 ml. of 6 M NaOH, 10 ml. of K3Fe(CN)6(0.5 M), and 30 ml. of 6 M HCl added.lOO ml.

Load more

  35

Copy Link