440
ANALYTICAL CHEMISTRY
Eomewhat poorer. Because of the limited sample size, primary amine contents of less than 0.01% in the presence of secondary and tertiary amines cannot be determined by this method. The method was successfully applied to the determination of ethanolamine hydrochloride without prior neutralization of the Eample. The absorbance of the hydrochloride is proportionally the same as the free amine. LITERATURE CITED
(1) Duke, F. R., IND. ENG. CHEM., ANAL.ED. 17, 19G (1945). (2) Haddock, L. A,, Evers, N., Analyst 57, 495 (1932).
(3) Hawkins, W.,Smith, D. bl., Nitchell, J., Jr., J . Ani. Chern. SOC. 66, 1GG2 (1944). (4) Johnson, J. B., Funk, G. L.. ANAL.CHEM.,submitted for publi-
cation. (5) Martell, A. E., Calvin, II.,“Chemistry of the Metal Chelate Compounds,” p. 273, Prentice-Hall, New York, 1953. (6) Siggia, S., Hanna, J. G., Kervenski, I. R., ANIL. CHEY.22, 1295 (1950). (7) Van Slyke, D. D., J . B i d . Chern. 9 , 185 (1911). (8) Wagner, C. D., Brown, R. H., Peters, E. D., J . Am. Chem. SOC. 69, 2609 (1947). (9) Woelfel, 15’. C., BNAL. CHEW20, 722 (1948).
RECEIVED for review October 21,
1955.
Accepted February 7, 1956.
coulometric Titration of Ammonia with Hypobromite G. MYRON ARCAND and ERNEST H. SWIFT, California Institute o f Technology, Pasadena, Calif.
A coulometric niethod has been developed for the titration with electrolytically generated bromine of niicrogram quantities of ammonia in a solution having a pH of 8.5. The effects of pH and of certain metal ions on the titration have been investigated. Quantities of ammonia ranging from 14 to 230 y have been determined with an average error of less than 0.2 y.
W
ILLARD and Cake (14) and Ilolthoff and Stenger (7) have developed volumetric methods for the determination of ammonia by oxidation with hypobromite according to the following equation: 2NH3
+ 3Br0-
= 3Br-
+ N* + 3Hz0
(1)
More recently, aniperonietric end points have been described for titrations which are based on the same reaction (8, 9). Bromine disproportionates in basic solutions (6) according to the equation Br2
+ 20H-
= BrO-
+ Br- + HzO
(2)
Thus, it seemed that a coulometric determination of ammonia might be developed which utilized the electrolytic generation of bromine and was based on the reactions represented by Equations l and 2. The possibility that such a method could be used for the direct determination of the ammonia in a Kjeldahl digestion solution justified the investigation desciibed. EXPERIMESTAL
Reagents. All solutions were prepared from reagent grade chemicals. Kater was distilled from alkaline potassium permanganate in order t o remove most of the reducing materials normally found in the tap distilled water. This special water was prepared in lots of about 2 liters each. Buffer solutions were prepared by dissolving 4.8 grams of sodium tetraborate decahydrate in 500 ml. of water to give solutions which were 0.025VF (volume formal) in sodium tetraborate. The pH was adjusted to the desired value by adding sodium hydroxide pellets or concentrated perchloric acid as required. Standard solutions of ammonium chloride from 0.033 to 0.074VF were prepared by dissolving weighed quantities of dried Ealt (Baker’s analyzed reagent) in the specially distilled water in a volumetric flask. Aliquots of these standard solutions were diluted to give concentrations ranging from 8.34 X 10-5T’F to 1.34 X ! O - V F . Bromine solutions were prepared by adding 5 grams of sodium bromide and 3 ml. of saturated bromine mater t o 1 liter of distilled water. These solutions were standardized by coulometric titration with cuprous copper as described below. Apparatus. All volumetric equipment was calibrated. The titration apparatus mas similar t o that described by Aleier, Myers, and Swift (10) as modified by Ramsey, Farrington, and Sn-ift (12). New equipment was constructed which had the following changes.
The current regulating vacuum tubes, Type IC;, were replaced by Type GAU6, because the latter have more stable opcrating characteristics and give more dependable service. The 11/2-volt dry cell which supplied the filament voltage for the vacuum tubes was replaced b y a &volt storage battery (Willard Signal Corps Battery BB-207/U); and the 11/2-volt dry cell which supplied the indicator potential was replaced by a 2-volt storage cell (Willard Signal Corps Battery BA-54.4). The rubber stopper which served as a cap for the titration cell during operation mas equipped with a short piece of glass tubing flared a t the top to allow addition of solutions to the cell without exposing the electrodes to the air. The indicator electrodes were made of pieces of sheet platinum 2 cm. by2.5 cm. A Leeds & Northrup box-type reflecting galvanometer, shunted with a rheostat to provide variable sensitivity, was used for indicator current readings near the end point. The generating current mas determined by measuring the voltage drop across a Leeds & Northrup standard mercury resistor (90.90s absolute ohms) through which the current was passed. This current corresponded to 4.0782 X equivalent equivalent per second a t the medium rate and to 1.0046 X per second a t the high rate. The voltage applied across the two indicator electrodes (the “indicator potential”) was measured with a Gray Instrument Co. Queen potentiometer. The solution in the isolated electrode compartment was 0.3VF perchloric acid. End Point Methods. Experinients showed that the indicator current resulting from a given quantity of bromine in an alkaline solution was not sufficiently reproducible for use as a means of determining the end point. Therefore the current in the alkaline solution was used only to indicate xhen a slight excess of bromine had been generated. Two methods for determining this excess were then considered. I n the first and recommended method the sohtion is acidified and bromine again generated until a predetermined current value is obtained. In the second method the solution is acidified, cupric copper added, and the excess bromine back-titrated with cuprous copper as described by Buck and Swift (a). Selection of Indicator Potential. An indicator potential of 150 mv. was used in both methods of determining the end point. In the first method described above this value was found to be sufficiently high to give an indication of the presence of excess bromine in the alkaline solution without causing an excessive current when the solution was acidified. Farrington, Meier, and Swift ( 5 ) found that the optimum applied indicator potential for the determination of the minimum indicator current in their cupric copper-bromide system was 60 mv. Fluctuations of indicator potential should then have a minimal effect on the minimum indicator current. On the other hand, Buck and Swift ( 2 ) used an indicator potential of 200 mv. for a simple system in which an arbitrary current
V O L U M E 28, N O , 4, A P R I L 1 9 5 6 caused by an excess of bromine was used. However, the bromide concentrations used in the experiments reported here are considerably higher than those used by Buck and Swift and caused higher minimum currents. In order to avoid high niinimum currents, a low indicator potential is desirable. Expeiiinents showed that the indicator current \vas no longer a linear function of concentration after -1 or 5 seconds of generation when the indicator potential was as low as 90 mv., but that linearity was maintained over the desired range when the indicator potential was greater than 120 mv. Most experiments weie performed with an indicator potential of I50 mv. All p H measurements were made with a Beckman Model F pH meter. Generator Current Efficiency in Alkaline Solutions. That hypobromite could be produced with 100% current efficiency in alkaline solutions was demonstrated by generating bromine for a known time under the conditions of a titration, then acidifying the solution, adding cupric copper, reversing the polarity, and titrating the bromine by generation of cuprous copper. The second end point method described above v a s used, and an arbitrary current value was taken as the end point of thc latter titration. An end point correction was made as follows. To 25.0 ml. of mater, 15.0 ml. of buffer (pH 0.3), and 10.0 nil. of 5VF sodium bromide n-ere added 2.0 ml. of 9VF perchloiic acid and 1.0 ml. of 1 V F copper sulfate. The resulting acid concentration was 0.3VF. Bromine was gcnerated in 1.0 second intervals and the stable indicator current recorded after each generation. The indicator current was considered to be stable when it did not vary by more than 0.2 to 0.3 pa. in 30 seconds. 11 straight line was passed thiough the plotted points and a predetermined arbitrary current value was selected as the end point current. The time Corresponding to this current was designated a s the end point correction time. Then, to check the generator current efficiency, a solution was prepared which containcd 25.0 ml. of water, 15.0 ml. of buffer (pH 9.3), and 10.0 ml. of 5 V F sodium bromide; this solution had a final p H of 8.8. In all cases, the p H of the prepared titration solution was 0.3 to 0.5 unit less than that of the buffer. The ionic strength of the buffer was about O.O5Vilf, and that of the titration solution was about 1VdI. Bromine was generated for 30.0 seconds, whereupon 2.0 ml. of 9 V F perchloric acid and 1.0 ml. of copper sulfate were added, bringing the acid concentration to 0.3VF. The polarity of the generator electrodes was reversed and cuprous copper was generated until the indicator current decreased to about 40 pa. The indicator current was recorded when it became stable. Generation of cuprous copper was continued in 1.0-second intervals until the indicator current had decreased below the arbitrary end point value. A striliglit line was drawn through the plotted points and the time corresponding to the end point current was taken as the titration time. The end point correction time was subtracted from the titration time to yield a value from which to calculate the equivalents of cuprous copper generated. I n general, the number of equivalents was not calculated, but comparison was made by considering the anodic and cathodic generation times. As the result of such checks, it was found in every case that less cuprous copper than bromine mas generated. The apparent loss of bromine was equivalent to about 7 to 8 seconds of generation, and was reproducible to within 0.2 second. Therefore, two sets of evperiments weie perfornied to attempt to determine if the error was due to low? generation efficiency or if bromine as being lost in some manner. End point correction determinations were made as previously described. Titration solutions rsere prepared by taking 15.0 ml. of vater, 15.0 ml. of buffer solution, 10.0 nil. of 5 V F sodium bromide, and 10.00 ml. of standard bromine solution. The pH of the solutions was 8.8. The solutions were allowed to stand for about 30 seconds, then 2.0 ml. of 9VF perchloric acid and 1.0 ml. of I V F copper sulfate were added and the bromine was titrated as described previously. Again, the losses of bromine
441 were equivalent t o about 7 seconds of gcneration. Similar espeiiments were performed in which bromine was added to the cell v*hen the platinum electrodes were present t o see if the electrodes had a catalytic effect; also the quantity of bromine added was varied to see if the effect was dependent on bromine concentration. In all cases the loss was the same, suggesting that impurities in the solution ivere being oxidized by hypobromite and that bromine could be generated in alkaline solution with 100% current efficiency. Because thi- losses were reproducible, the decision was made to use a combined reagent and end point correction procedure rather than to attempt to purify all the reagents. PROCEDURE
Reagent and End Point Correction. The, blank solution contained 25.0 ml. of vater, 15.0 ml. of buffer solution (pH 8,9), and 10.0 ml. of 5VF sodium bromide. The pH of the solution \ms 8.5. Bromine was generated for 8 to 10 seconds and the indicator current allowed t o stabilize. I t found that if the current vias.less than 10 pa., the reducing impmities were not completely removed. When the current stabilized at abo$ 10 pa., 2.0 nil. of 9 V F perchloric acid were added, causing a rise in the indicator current t o about 25 pa. The current was recorded and bromine was generated in 1.0-second interv&, the stable current being recorded a t each interval. A straight line was drawn through the plotted points and the time corresponding t o 30 p a . was taken as the reagent and end point correction time. Correction determinations were repeated until three consecutive determinations agreed within 0.2 second. Titration. The titration solution contnincd 15.0 nil. of water, 15.0 ml. of buffer solution (pH 8.9), 10.0 ml. of 5VF sodium bromide, and 10.00 ml. of standard ammonium chloride solution. The ammonium chloride solution was added just before the cell was mounted, so that losses through volatilization might be reduced. As bromine was generated, the indicator current usually rose to about 20 pa. and leveled off a t about that value for most of the titration. Near the end of the titration, the current hegan to rise and the titration was stopped when the rate of rise was about 1 pa. in 5 seconds (usually 5 to 15 seconds before the end point). The current mas alloiwd to decrease until stable and, if it xas lower than 10 pa., generation was continued in short intervals until a stable current of about 10 pa. W D S obtained. Then 2.0 ml. of 9 V F perchloric acid were added and the generation of bromine was continued as described for the reagent and end point correction. The reagent and end point correction time was subtracted from the titration time to obtain a corrected time; this time was then used to calculate the quantity of ammonia titrated. DISCUSSION A N D RESULTS
Effect of pH. Previous workers (8) have found that the accuracy of the volumetric titration of ammonia with hypobromite is critically depeiident upon the pH of the solution. Among the factors vhich are involved is the stoichiometry of the titration. Clusius and Rechnitz ( 4 ) and Riley and others (IS)have shown that under certain conditions significant quantities of nitrous oxide may be formed by the following reaction: 2XH3
+ 4Br0-
=
N20
+ 4Br- + 3H20
(3 )
The stoichiometry of the titration is also affected by the extent to which the reaction
3 BrO- = BrO3-
+ 2Br-
(-5)
takes place, since bromate ion docs not oxidize ammonia. Tinally, the possibility of volatilization of ammonia from the solution becomes greater with increase in pH. Because of these considerations a series of experiments was made in order to establish the permissible pH limits for the eoulometric titration. The end point correction and titration procedurcs used were as described above, with the exception 1hat the p H of the system was varied. The results of a series of experiments in which the p H was varied between 8.2 and 10.0 are shown in Table I. It can be seen that 8.5 and 8.2 are the best p H values of those
442
ANALYTICAL CHEMISTRY
Table I. PV-0.
Effect of pII on Determination of Ammonia of
Detn.
pH
3 7 5 5 6 1
10.0
Standard Dev.. Y
55.72 226 2 56.62 227.3 56.98 227.7 57.17
10.14 f0.3 10.12 10.2 10.12
8.5
516.81 227.4 56.81 227.5
8.2
227.4
8.8
56.81 56.81
3
Table 11.
A m m o n i a , ~ Taken Found Error -1.1 -1.2 -0.19 -0.2 f0.17 +0.3 +0.36
10.20
Effect of Excess IIypobronlite on Detcrniirlatiorl of Ammonia
(Ammonia sample WRS added t o t h e solution after a quantity of brominc had been generated. Equivalents of bromine initially generated are shown in second column; total generated for titration are shown i n third column. pH of each solution was 8.8.)
x0,of
Detn. 1
4 3 7
Bromine Generated, E q u i v . X lo7 Initial Total 230 200 100 0
401 401 401 401
Ammonia, y Taken Found 227.4 227.4 227.4 227.4
228.8 228.0 227.5 226.2
Error f1.4
+0.6 +0.1
-1.2
Std. Dev., Y
f0.2 10.2 10.3
tried. Hovever, a system a t a pH of 8.2 is not convenient, because attainment of constant indicator currents a t this pH value was excessively slow. Kolthoff and others (7) found that a pH of 8.5 was most satisfactory for their procedure. It is seen that the observed errors shift from negative to positive as the pH of the system decreases. Although the reason for this is not definitely known, it is possible that a t the higher pH values loss of ammonia may be the predominant source of error, while formation of nitrous oxides may predominate a t the lower values. It is also possible that compensation of these errors may give the best results a t pH S.5. Other buffer systems were considered for the control of the pH during the titration. Tile carbonate-bicarbonate system is impractical for the method used here because the formation of carbon dioxide on acidification of the solution would tend to sweep bromine from the solution. Phosphate buffers were not used because the available reagents apparently contained animonia or other reducing impurities, which would liave caused the blank corrections to be unreasonably large. Effect of Excess Hypobromite. The fact that fairly large negative errors \\-ere found when 227 y of ammonia were titrated a t p H 8.8 suggested the possibility of adding the sample after some quantity of bromine had been generated, thus allowing less time for the ammonia to escape from the solution. A series of experiments v a s performed in which bromine was generated for certain fixed times, after which the sample was added and the titration continued. The results are shown in Table 11. It can be seen that the positive error increased as the initial excess of bromine (or hypobromite) was increased. This may be due to an increased formation of nitrous oxide when the hypobromite is in excess. End Point Method. The optional end point methods have been described. An end point in the alkaline solution would have eliminated the necessity for adding acid to the solution with the resulting high indicator current. However, it was found that the in'dicator current in alkaline solution was not reproducible and did not give a reliable measure of the amount of bromine which was present. Another reason for determining the end point in acid solution is that it allows the possibility of back-titration with cuprous copper if the end point is badly overrun. However, if this backtitration method is used, care must be exercised to use exactly the same technique for the blank as for the titration. A few runs were made in which the blank correction was determined as described above and the titration was accomplished by a back-
titration method. In these cases, the errors were as large as +0.6 y , compared with an error of 0.2 y when the end point was determined as described in the procedure. The slope of the indicator current-time curve was greater when the bromine concentration was decreased by generating cuprous copper than when the bromine concentration was increased by generating bromine. The reason for this difference is not known. However, it was sufficient to cause observed differences in the accuracy of the titration. The concentration of sodium bromide used (approximately 1 V F ) was that which experiments had shown to be satisfactory for the determination of bromate by addition of bromide and cupric copper to an acid solution and titration by generation of cuprous copper. These experiments were made t o provide for the possibility that significant quantities of bromate would be formed during the titration of ammonia with bromine in an alkaline solution. Buck and Swift ( 2 ) used a bromide concentration of 0.1VF for the determination of aniline; therefore, a comparable concentration may be practical for the determination of ammonia. Experiments in connection with the bromate determination had shown that the lowest practical acid concentration was 0.3VF. The procedure used by Laitinen (9),in w-hich a single indicator electrode is employed, has the advantage that the end point can be determined in allraline solution. However, if small quantities of bromate are formed, positive errors result, because it was found that the bromate does not oxidize ammonia under the conditions of the titration. Effect of Other Metal Ions. If the procedure described above could be utilized to eliminate the distillation process in the microKjeldahl determination of nitrogen, it would greatly increase the ease with which that procedure is carried out. Therefore, experiments were made to determine if the materials commonly employed as catalysts in the Kjeldahl digestion would interfere Kith the coulometric determination of ammonia. The procedure just described was used, except that amounts of copper sulfate, mercuric nitrate, and selenium were added which were typical of the amounts which n-ould be found in a micro-Kjeldahl determination (5, 11).
Table 111. Effect of Metals on Determination of Ammonia [Concentration of CuSOd was 6 X lO-rVF, and of Hg(KOa)z, 2 X 1 0 - 4 V F ~ I .4dded Metal Cu(I1)
Taken 227.5
56.81
HsW)
227.5 56.81
Ammonia, y Found 227.8 227.2 Av. 2 2 7 . 5 57.13 57.06 56.82 Av. 5 7 . 0 0 227.1 56.64 56.93 5 6 . 66 Av. 5 6 . 7 4
Error +0.3 -0.3 fO.0 f0.3
t0.2
0.0 t0.2 -0.4 -0.2 fO.1
-0.2 -0. I
Selenium metal was dissolved in hot concentrated sulfuric acid, then slowly diluted with water, and the solution n-as neutralized to a pH of 4. On dilution, some selenium reprecipitated. A portion of the mixture was filtered and experiments were performed with small quantities of the mixture and of the filtrate. The results wit.h copper sulfate and mercuric nitrate are shown in Table 111. It can be seen that neither copper sulfate nor mercuric nitrate affects the accuracy of the ammonia determination.
V O L U M E 28, N O . 4, A P R I L 1 9 5 6
443
The presence of selenium affects the titration, however. No actual titration was made with selenium compounds present, because it was found that the selenium in solution reduced hypobromite as rapidly as i t was formed. Although the species of selenium involved is not known, experiments made u-ith selenious acid showed that selenite reduced hypobromite. It is evident that if selenium is to be used as a catalyst, some way must be found t o eliminate interference before the ammonia determination is begun. Experiments 11-ere performed in which metallic lead was added to the sulfuric acid solution containing selenium in the hope that the selenium would be reduced to thc metal. Although most of the selenium was removed from the solution in this way, 4 days were required to complete the reduction. Reduction with sulfur dioxide ( 1 ) mas not considered because of the time required for removal of the excess.
Table IV. No. of Detn.
Taken
Confirmatory Titrations .4mmonia, y Found
Error
Std. Dev., Y
Confirmatory Titrations. Table IV shows the results of a series of confirmatory titrations made by the procedure described above. I n all cases, the average error was less than 0.2 y, mhile the standard deviation was about the same. Thus, one might expect a maximum error of 0.4 y in a determination involving quanti-
tities of ammonia between 14 y and 230 y. Both the relative and absolute accuracies are somewhat better than those reported by Laitinen (9). ACKNOWLEDGiMENT
Preliminary experiments made by Fred C. Anson have been helpful in this work. G. Myron Arcand is indebted to the General Electric Co. for a predoctoral fellowship during the academic year 1954-55. LITERATURE CITED
(1) Adams, C. I., Spaulding, G. H., ANAL.CHEM.27, 1003 (1955). (2) Buck, R. P., Swift, E. H., Ibid., 24, 490 (1952). (3) Clark, P. P., “Semimicro Quantitative Organic Analysis,” p. 40, Academic Press, New York, 1943. (4) Clusius, K., Rechnitz, G., H e h . Chim. A c t a 36, 59 (1953). (5) Farrington, P. S., Meier, D. J., Swift, E. H., ANAL. CHEY. 25,591 (1953). (6) Jones, G., BaeckstroGm, S., J . Am. Chenz. SOC.56, 1517 (1934). (7) Kolthoff, I. AI., Stenger, lr. A., I X D . I h G . CHEW ANAL. E D . 7 , 7 9 (1935). (8) Kolthoff, I. M., Stricks, W., hIorren, L., Analyst 78, 405 (1953). (9) Laitinen, H. A., IVoerner, D. E., AXAL.CHEY.27, 215 (1955). (10) hIeier, D. J., Myers, R. J., Swift, E. H., J . A m . Chem. SOC.71, 2340 (1949). (11) Milton, R. F., Waters, W.A., “Nethods of Quantitative bIicro Analysis,” p. 79, Edward Arnold and Co., London, 1949. (12) Ramsey, IT. J., Farrington, P. S., Swift, E. H., AXAL.CHEM. 22,332 (1950). (13) Riley, R. F., Richter, E., Rotheman, A I . , Todd, N., Myers, L. S., Jr., Nusbaum, R., J . Am. Chem. Soc. 76, 3301 (1954). (14) Willard, H. H., Cake, W.E., Ibid., 42, 264G (1920).
RECEIVED for review July 18,
1955. Accepted ,January 10, 1956. Contribution No. 2016 from the Gates and Crellin Laboratories of Chemistry.
Coulometric Titrations with Electrically Released Ethylenediaminetetraacetic Acid Titrations of Calcium, Copper, Zinc, and lead CHARLES N. REILLEY and WILLIAM W. PORTERFIELD Department o f Chemistry, University o r North Carolina, Chapel Hill, N. C.
Successful coulometric titration of calcium, copper, zinc, and lead ions was accomplished by the indirect electrical generation of ethylenediaminetetraacetic acid released upon reduction of mercuric-ethylenediamine tetraacetate chelate at a mercury pool. The pertinent equilibrium conditions are summarized in a potentialpH diagram. This diagram, in conjunction with polarograms for kinetic effects, furnished valuable information concerning the solution conditions desired. Extension of this method to the titration of many other metal ions appears feasible.
7 OULOMETRIC titrations with constant current have b een employed in various t,ypes of oxidation-reduction, precipitation, and neutralization reactions. The present investigation was undertaken to test the feasibility of extending the application of coulometric titrations into compleximetry. If a general method for the coulometric generation of ethylenedianiinetetraacetic acid [EDTA, HiY, (ethy1enedinitrilo)teti-aacetic acid] could be found, it would be highly probable that the coulometric method could be used for the successful titration of
(A
the alkaline earths, the rare earths, and such metal ions as manganese(I1 j, iron(II), copper(II), mercury(I1j, cobalt, nickel, zinc, cadmium, lead, aluminum, gallium, indium, thorium, palladium(II), and bismuth. The solution of i,his problem required the discovery of a mode of generating ethylenediaminetetraacetic acid with lOO7, current efficiency, anti the development of a compatible and sensitive end-point detection device. The properties of the end-point method selected are discussed in greater detail elsewhere (8). Mode of Generation. Several modes of coulometric titration of metal ions through the formation of stable chelates with ethylenediaminetetraacetic acid suggest themselves. The formation of the reagent, ethylenediaminetetraacetic acid, through an electrode process such as the osidation of ethylenediaminetetraacetaldehyde, was ruled out on the basis of requiring reagents which were not commonly available in pure form and because doubt existed about the tjuccessful generation a t 100% current efficiency of ethylenediaminetetraacetic acid in such a manner. It seemed that either of the following two indirect titration procedures might be employed. The first would require the coulometric generation of base in order to titrate the acid liberated when an excess of ethylene-
