Sulfarnic Acid as a Standard of Reference in Acidimetrv J
SISTER 31. JOSETTA BUTLER, G. FREDERICK SMITH, University of Illinois, Urbana, Ill.
Sulfamic acid is a crystalline nonhygroscopic solid, available on the market in any desired quantity at a moderate price. It is a strong acid in aqueous solution and can be titrated with bases, using indicators with transition ranges varying from a pH of 4 to 9. Sulfamic acid is an excellent acidimetric standard of reference and should find widespread use in analytical chemistry. In precision and accuracy it compares well with other acidimetric reference materials. It can be purified and dehydrated easily and thus be obtained in uniform and exact composition.
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has been in progress in this laboratory for several years. I t was considered advisable to reexamine the usefulness of this acid as a primary standard, especially since it is now available in quantity both in the commercial and c. P. grades (technical grade from the Grasselli Chemicals Department of E. I. du Pont de Kemours B: Company). It is the purpose of the present investigation to describe such a study.
S A RECEST publication dealing with sulfamic acid,
NH,SO,H, Cupery ( 2 ) calls attention to its unusual physical and chemical properties, pointing out that “it is an iniportant addition to the group of comniercial acids represented by lactic, acetic, formic, tartaric, oxalic, and similar acids, and should be especially useful for applications in TI hich a highly ionized, non\-olatile acid is desired, or where precipitation of insoluble salts must be avoided.” I n discussing potential applications of sulfamic acid he states that “it should find extensive m e as an analytical reagent” and that “it has previously been recommended as a standard for titrimetric work because it is a nonhygroscopic crystalline acid which gives sharp end points with ordinary titration indicators.” Hoffmann (6) v a s the first to suggest that sulfamic acid might serve as an acidimetric standard. H e found that solutions of sulfamic acid could be titrated with potassium hydroxide and that either phenolphthalein or methyl orange could be employed as indicator. Herboth ( 5 ) subsequently recommended sulfamic acid for use in standardizing pharmaceutical solutions, but probably did not have a t his di-posal pure sulfamic acid. Several obvious misstatements occur in his publication, chief among which is the observation that an insoluble barium salt is precipitated when barium chloride is added to a solution of the acid. This is not in accordance with observations made in this laboratory; moreover, Cupery ( 2 ) gives the solubility of barium sulfamate as 34.2 grams per 100 grams of water a t 25” C. Since aqueous solutions of sulfamic acid are slowly hydrolyzed in accordance with the equation KH2S03H H?O + NHIIISOI, it is highly probable that the precipitate obtained by Herboth n-as barium sulfate and not barium sulfamate. Herboth also concluded that sulfamic acid as an acidinietric standard is not susceptible to any high degree of precision. I n the light of the experimental work reported below, it is again obvious that the discrepancies reported by him are due to the impurity of his materials. Misuch (7) also investigated sulfamic acid for use as a titrimetric standard. He was able to obtain comparable titers for solutions of sodium hydroxide and sodium carbonate when standardized with hydrochloric and sulfamic acids. A comprehensive study of the chemistry of sulfamic acid
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Physical Properties of Sulfamic Acid Sulfamic acid is a crystalline nonhygroscopic solid melting with decomposition at 205’ C. It is highly ionized in aqueous solution, as shorsn both by conductometric (8) and pH measurements ( 2 ) . It is indefinitely stable in the solid state at ordinary temperatures, but undergoes slow hydrolysis in solution in accordance with the equation given above. T h e solubility of the acid in 100 grams of water ranges from 13.68 grams at 0” t o 47.08 grams at 80” C. Sulfuric acidgreatly decrease? its solubility in \rater, a fact which may be used in its recovery from solution and in its purification ( 2 ) . The free acid is appreciably soluble in methanol and ethanol, slightly soluble in acetone, and practically in3oluble in ether (2). Such nitrogenous solvents as formamide ( 2 ) and liquid ammonia (observation in this laboratory) eshibit marked dissolving pomr for the acid. Many of the inorganic salts .of sulfamic acid have been prepared and practically all are soluble in water. Other physical data are presented in detail by Cupery ( 2 ) . Chemical Characteristics of Acidimetric Standards Sulfamic acid may properly be compared with such existing standards as benzoic and succinic acids, potassium acid phthalate, and potassium biiodate-nonhydrated crystalline compounds which are soluble in water and can be purified by simple recrystallization from water. Only potassium biiodate and sulfamic acid are strongly acidic and capable of use in connection n i t h a group of indicators having transition points over a wide range of p H values. A comparison of the properties of sulfamic acid and potassium biiodate is, therefore, pertinent. Potassium biiodate has the advantage of a high equivalent weight and, when properly prepared, possesses a definite, known hydrogen-ion content. However, a t least three recrystallizations from water are required to attain satisfactory purity. Since it contains a high percentage of iodine, potassium biiodate is not inexpensive. I t s solubility in water a t ordinary temperatures is low; i t is, therefore, difficult to prepare solutions more concentrated than 0.1 molar. Indicators having transition ranges between p H 5 and 9 may be employed. A solution of standard acid can be prepared using potassium biiodate, although this practice is not often followed. Sulfamic acid has a low equivalent weight (97.17). The properly purified product has a definite known hydrogen-ion content. It is easily purified by a single recrystallization from aqueous solution. The sulfamic acid of commerce is available in any amount a t a very modest price. It is readily soluble in water. Indicators with transition ranges between pH 4 and 9 may be employed. Sulfamic acid has thus far not been employed in t,he preparation of working solutions of known acid strength, although, as indicated by studies now in progress in this laboratory, by suitable procedure the preparation of such solutions might be made practicable.
DECEMBER 15, 1938
-4NALYTICAL EDITIOS
Preparation and Purification of Sulfamic Acid The sulfamic acid used in this investigation was obtained from a number of sources. Commercial samples produced by the sulfolysis of urea TI ere made available through the courtesy of AI. E. Cupery of the Experimental Station of E. I. du Pont de Xemours B Company. Samples mere prepared in this laboratory for the preliminary work by the sulfolysis of urea ( I ) , and by the action of sulfur dioxide upon hydroxylamine salts (3) and upon acetoxime ( 9 ) . seTera1 methods for the purification of the commercial material were tested before a satisfactory and reproducible piocedure TT as evolved. Simple recrystallization from hot water gave products whose hydrogen-ion values were consistently low to the extent of 0 1 to 0.2 per cent. It Ti-as apparent that some nonacid impurity was present in very small amounts-water or possibly urea, since the commercial material is made from urea. Recrystallization from a concentrated sulfuric acid solution was next attempted, but in no case was it possible to eliminate all traces of sulfate ion from the product. If urea or some other basic material nere present in the commercial product, it is conceivable that such a compound would be rather firmly fixed by sulfamic acid in the form of a salt and might concentrate in the first crystalline fractions obtained from solution Working on this theorv a 125-gram sample of the crude acid n as diqsolved in 300 grams of water preheated to 70' C. The solution nas filtered three times nith con-equent lonering in temperature, and each time the material crystallizing from solution (altogether about 25 grams) %\asdiscarded. The final filtrate was cooled rapidly to the temperature of an ice-salt mixture and alloned to stand for 20 minutes. The crystal> thus formed were removed by suction filtration, and washed nith a small quantity of ice Rater, then twice with cold ethanol, and finally nith ether.
The product was air-dried in an open dish for 1 hour, after which it was ground in an agate mortar and stored in a desiccator over Anhydrone. When samples of this material were analyzed, much more accurate results n-ere obtained. This method was therefore adopted for the purification of the four samples used to establish the value of sulfamic acid as an acidimetric standard of reference.
Standardization of Reference Solution of Barium Hydroxide A stock solution (18 liters) of barium hydroxide was prepared and standardized using constant-boiling hydrochloric acid prepared according to the method of Foulk and Hollingsworth (4). The barium hydroxide solution was stored in a large delivery bottle and protected from the carbon dioxide of the air by a tube containing Ascarite. Weight burets were used for all titrations and for n-eighing the hydrochloric acid.
Table I gives the results of this standardization. These show close agreement between individual determinations and the average hydrogen chloride value of the barium hydroxide solution a t the beginning and again a t the conclusion of this study. Several hundred subsequent titrations of various samples of purified sulfaniic acid n-ere carried out, and are reported consecutively with no omissions. Tables I1 and I11 were obtained using the same solution of barium hydroxide, which has a sulfamic acid equivalent of 0.010875 gram per gram of barium hydroxide solution required. Vacuum corrections were applied to all weighings of solutions and samples involved in this study.
Potentiometric Titration of Sulfamic Acid A series of p H titrations by sodium hydroxide using the glass electrode was carried out in order to demonstrate the strength of sulfamic acid and to determine suitable indicators
FIGURE 1. TYPICAL TITRATION CURVES I-IV. V, VI.
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Potentiometric titration of sulfamic acid b sodium hydroxide, using glass electrode. Comparative titrations of hydrochloric acidlwith barium hydroxide. Sharper breaks continuing over wider p H ranges are obtained with sulfamic acid.
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ordinary distilled water was sufficiently low (5.2 to 6) to shift the end point.
TABLEI. STANDARDIZATION OF APPROXIMATELY 0.1 N BARIUM HYDROXIDE NO.
1 2 3 4 5 6 7 8 AY.
1 2 3 4 5
Av.
(Using constant-boiling hydrochloric acid as reierence) Constant Boiling Gram HCl/Gram Deviation HC1 Ba(0H)z Ba(OH)z from Average Gram Grams % Standardization at Beginning of Investigation 0.020174 0.77235 38.284 -0.035 0.020184 0.73147 36,206 f0.015 0.020198 0.75848 37.552 f0.084 0.020183 0.75408 37.363 f O . 009 0.020181 0.77588 38.494 *o, 00 0.020164 0.77628 39.498 -0.084 0.020190 0.78810 39.039 f0.044 0.020178 0.78010 38.666 -0.015 .... 0.020181 0,036
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Standardization a t Conclusion of Investigation 0.020187 0.69460 34.409 0.020178 0.51275 25.362 0.020187 0.70935 35.187 0.020168 0.71285 35.346 0.020167 0.71475 35 443 ..... .... 0.020177
VOL. 10, NO. 12
f O , 049 f O . 005 f O . 049
-0.044 -8,049 0.039
OF PURITY OF SULFAMIC ACID TABLE11. DETERMINATION
(By titration with standard barium hydroxide using bromothymol blue as indicator. Individual titrational data) Gram Deviation NHISOIH Ba(0,H)r Acid/Gram from No. Taken Found Error Required Base Average Oram Gram Gram % Grams % Gram 1 0,39340 0.39346 f 0 . 0 0 0 0 6 0 . 0 2 3 6 . 1 8 0 0 010873 -0.000005 0 . 0 6 36.582 0.010881 t 0 . 0 0 0 0 0 3 2 0.39805 0.39783 -0.00022 3 0.39220 0.39194 -0.00026 0.06 3 6 . 0 4 0 0.010882 1 0 . 0 0 0 0 0 4 4 0.40245 0.40266 f 0 . 0 0 0 2 1 0 . 0 5 3 7 . 0 2 6 0.010869 -0.000009 0 . 0 7 36.540 0.010883 f0.000005 5 0.39765 0.39737 -0.00028 6 0.39945 0.39928 -0.00017 0 . 0 4 36.715 0.010880 f0.000002 0 . 0 4 3 5 . 9 0 4 0.010884 f0.000006 7 0.39065 0.39046 -0.00019 S 0.39940 0.39944 f 0 . 0 0 0 0 4 0 . 0 1 3 6 . 7 3 0 0 010874 -0.000004 Av. 0.00018 0 . 0 4 4 . . . . 0 010878 0.000005
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Table I1 gives a typical series of consecutive results. A summary of the average results of four such series is given in Table 111. A study of Tables I1 and I11 indicates that the average purity of the various preparations is approximately 99.945 per cent. The chief impurity is thought to be water of occlusion. The experimental results show a very satisfactory concordance from a series of individual preparations.
Discussion The experimental data demonstrate conclusively that sulfamic acid can serve as a n excellent primary standard. The fact that it can be purified merely by recrystallization from water and then need only be dried in air is a strong argument in favor of its use. The slightly low results, as compared with constant-boiling hydrochloric acid, are undoubtedly due to a very small trace of occluded moisture. Drying a t 105" C., either a t atmospheric pressure or in uacuo, did not change the values obtained. Higher temperatures are inadvisable, as samples heated to 135" C. always showed that some of the occluded moisture had been "fixed"-that is, a test for the sulfate ion was always obtained, presumably because of hydrolysis. Aqueous solutions of sulfamic acid are not stable, as hydrolysis occurs very slowly a t room temperatures with formation of ammonium acid sulfate. Since the hydrolytic product contains only one readily releasable proton, it is likely that the titers of such solutions would remain fairly constant. A study of the stability of aqueous solutions of sulfamic acid is in progress.
for its use. A number of typical curves (Figure TABLE111. DETERMISATION OF PVRITY OF QULFAMIC ACID depict the change in pH On addition Of inBy titration with barium hydroxide, using bromothymol blue as indicator. Composite creasing quantities of sodium hydroxide. For titrational data) Average Average purposes of comparison two curves obtained by Average Gram Deviation B a ( 0 H h .icid/Gram from titrating hydrochloric acid with barium hydrox&::? "%::if Ivesai?.e Error Required Base Average ide are also included. It is readily apparent Gram Gram Gram % Grams Gram from the form of these curves that sulfamic acid A 5 0.40358 o 40326 - 0 . 0 0 0 3 ~ o 0 8 3 7 . 1 0 ; o 0108:b: -O.OOOOOI B 9 0.40053 0 40037 -0.00016 0 04 36.81s 0.010880 +0.000002 can properly be with the strong C 8 0.40110 0 40099 -0.00011 0 03 36.123 0.010878 -0.000001 and that a whole series of indicators with transiD 8 0.39665 0 396.56 -0 00010 0.04 36 466 0.010878 -0 00000l ..... . . . 0.00018 0 047 .... 0 010879 . ., A'., tion points between p H 4 to 9 can be employed. Several indicators were used in conjunction with the pH titrations. I n the case of broAcknowledgnien t mothvmol blue, rosolic acid, Dhenolphthalein, methyl orange, and ðyl red-methylene- blue, the characteristic coior The authors desire to acknowledge their inclebtediieas to changes occurred within the indicated p H range. M. E. Cupery of the Experimental Station of E. I. du Pont de Bromothymol blue was selected for the subsequent inxemours k Company for a generous supply of technical vestigation of sulfamic acid because the color change occurs famic acid; and t o Harry Sisler for his help in preparing sulexactly at the experimentally determined equivalence falnic acid for use in the preliminary study of this problem point, ~h~ color change is also very sharp, yellow at a p~ of 6.4 and blue a t a p H of 7 . This indicatorhas the disadxintage of giving little warning as the end point is approached. Literature Cited This disadvantage is counterbalanced, however, by the added Baumgarten, P., Ber., 69, 1929 (1938); U. 3. Patent (to E. I precision attainable through its use. For rapid but less predu Pont de Nemours & Company) 2,102, 350 (Dec. 14, 1937), cise results the mixed methyl red-methylene blue indicator Cupery, h l . E., IXD. ESG. CHEU.,30, 627 (1938) is very useful. Divers and Haga, J . Chem. SOC.,69, 1834 (1896).
sziz.
Standardization of Barium Hydroxide Solution Approximately 0.4-gram samples of sulfamic acid were weighed in small weighing dishes. Each sample was dissolved in 100 cc. of distilled n-ater contained in a titration flask, and 6 drops of the bromothvmol blue indicator were added. Each solution was titrated directly Ivith barium hydroxide from a weight buret. Carbon dioxide-free air was bubbled through the solution during the titration. It was found necessary t o use boiled distilled water for washing down t h e sides of t h e flask, as the pH of the
Foulk and Hollingsworth, J . Am. Chem. SOC.,45, 1220 (1823). Herboth, L.. Arch. Pharm., 262,517 (1924). Hoffmann and Bielsalski, B e r . , 45, 1394 (1912). Misuch, K.,Farm. Zhur., 1928,310-13; C h e m . Z e n f r . . 99 ,111. 1129 (1928).
Sakurai, J., J. C'heitb. SOC.,69, 1654 (1596j. Schmidt, >I.. J . p r n k t . Chein., 44, 513-38 (18811 .iUgult 3 , 1838. ~ i e a e n t e dbefore the Dlvlslon Physical and Inorganic Chemistry a t the 96th lfeeting of the lrnerican Chemicai Society. Lfilxaukee. \\-is,, September s t o 9 , 1988.
RECEWED
