Coulometric Titrations with Iodine Titration of Arsenic and Use of a n Amperometric End Point WILLIAM J. RAMSEY, PAUL S. FARRINGTON,
AND
ERNEST €1. SWIFT
Calqornia Institute of Technology, Pasadena, calij. The conditions under which electrolytically generated iodine can be used for secondary coulometric titrations and the end point determined by an amperometric method have been investigated. Tripositive arsenic in quantities from 64 to 1200 micrograms has been titrated in solutions having a pH value of 8 with an average error without regard to sign of 0.6 microgram.
T
Limited range pHydrion indicator paper was used instead of the pH meter for determining pH values greater than 10. In all the above solutions, reagent grade chemicals were used. The laboratory distilled water was often found to contain small amounts of an unknown oxidizing agent. This oxidizing agent n as removed by boiling the distilled water for about 20 minutes, and then bubbling nitrogen through it for about 30 minutes. Only water treated in this way was used. Apparatus. The apparatus wm essentially the same as that described by Aleier, Myers, and Swift (4),with the following changes: In place of the laboratory direct current supply, the voltage of which was found to vary slightly, use was made of a simple voltage-regulated rectifier which was connected to the alternating current line through a voltage-regulating Sola transformer. The generating cathode was enclosed in a shield rather than the anode. Finally, except where noted, nitrogen gas, after being passed successively through 6 VF sodium hydroxide and soda lime, was maintained above the solutions being titrated Preliminary Adjustment. The current of the generation circuit was determined by measuring the voltage drop across a standardized 199.87-ohm resistance through which the generation current was passing. At the low rate of generation, this current vas found to correspond to the oxidation of 1.0385 X low8equivalent per second, while the high rate current corresponded to 1.0393 X 10-7 equivalent per second. The indicator potential was measured with a Gray Instrument Companv Queen potentiometer. One yolume formal sulfuric acid was placed in the cathode shield. Before each set of titrations the indicator electrodes were ihorted to the generator anode, a potassium iodide solution n-as placed in a titration cell (a 40 X 80 mm. weighing bottle), and iodine was generated in the solution for 100 seconds a t the high rate. Titration Procedure. The initial step in carrying out a set of titrations was the determination of the blank time. In order to do this, a number of blanks were run. For each blank, 5 ml. of 1 T'F potassium iodide, 20 ml. of the buffer of pH 8, and 25 ml. of water were placed in a titration cell. The indicator potential was adjusted to an initial value of 0.15 volt. The generation current, passing through a dummy resistance, was then adjusted to the desired value. Finally, iodine was generated for short intervals of time depending upon the rate of generation used (0.5 second for the high rate, 2.5 for the low). Twenty seconds were allowed after each interval of generation for equilibrium conditions to be approached, and the indicator current and the generation time were recorded, A plot of indicator current us. time of generation was made, and the linear portion of this plot, from approximatelv 10 t o 40 pa., was extrapolated to zero indicator current. The average value of the generation time intercepts for a group of blanks was designated as the blank time. To maintain the sensitivity of the indicator electrodes, they were shorted to the generator anode, and iodine lyas generated for 50 seconds a t the high rate after each blank and every titration. Then the solution was removed from the apparatus, and the electrodes were rinsed with water. After the blank time had been determined, a group of titrations TI as made. Twenty-five milliliters of a standard arsenic solution nere pipetted into a titration cell. To this were added 20 ml. of the buffer of pH 8 and 5 ml. of 1 VF potassium iodide, and the cell was rapidly attached to the apparatus. The generation current was adjusted, and titration was begun. Occasionally during titrations, small adjustments had to be made on this current.
HE advantages of the secondary type of coulometric process, in which an intermediate half-cell reaction is caused to take place a t the electrode, have been discussed by Szebelledy and Somogyi ( 7 ) and by Meier, Myers, and Swift (4).A number of studies have been made of the use of bromine as an intermediate oxidizing agent for the titration of reducing agents such as hydrazine (Q),hydroxylamine ( f O ) , thiocyanate (8), thiodiglycol ( 6 ) , tripositive arsenic ( 5 ) ,tripositive antimony ( 1), and iodide ( 11 ) . The use of cuprous copper as an intermediate reducing agent for the titration of chromate and vanadate has also been investigated
(4). In order to extend the application of such processes, especially to selective titrations, it seemed desirable to investigate the application of other intermediate half-cell reactions having various standard potentials and capable of being used in solutions over an extended range of p H values. The criteria establishing the suitability of any proposed intermediate half-cell reaction are the following: The active state of this intermediate must be capable of being produced electrolytically with 100% current efficiency. The state thus produced must then be capable of rapidly and stoichiometrically reacting with the agent to be titrated. Lastly, there must be a method of determining an end point, and if the amperometric method is chosen, an excess of the active state of the intermediate must cause a current between the two indicator electrodes which for the required concentrations and conditions is directly proportional to the concentration of that state. An investigation of the iodide-iodine half-cell as such an intermediate has been made and the pH range has been established under which the titration of tripositive arsenic can be made with iodine and the excess iodine determined amperometrically. EXPERIMENTAL
Chemicals. Standard solutions of tripositive arsenic were prepared in the following way: A 1-gram sample of Bureau of Standards arsenious oxide which had been dried for 2 hours a t 120" C. was accurately weighed out and dissolved in 20 ml. of water containing 2 grams of sodium hydroxide. The resulting solution was acidified by adding 10 ml. of 3 volume formal ( V F )sulfuric acid and was diluted to approximately 100 ml. The weight of the resulting solution was determined. Solutions prepared in this way were used as stock solutions, none of which was kept for more than 4 weeks. Weighed samples of these stock solutions were diluted to appropriate volumes to provide the standard solutions, and none of these was kept for more than 2 days. Solutions of 1 V F potassium iodide were prepared by dissolving weighed samples of the solid in 0.005 VF sodium carbonate. A solution buffered to p H 8 was prepared by dissolving solid sodium hydroxide in a solution 0.25 VF in sodium dihydrogen phosphate until the correct p H was obtained as measured with a Beckman pH meter. Other buffered solutions were prepared in the following way: The weight of a salt corresponding to one of the ionic species used in the buffer solution sufficient to give a solution a t least 0.2 VF in that salt was dissolved in water. The corresponding acid or sodium hydroxide was added until the dcsired pH was obtained. 332
333
V O L U M E 22, NO. 2, F E B R U A R Y 1 9 5 0 When the indicator current began t o rise, the titration was stopped, and a plot of indicator current us. time of generation was constructed from data obtained in exactly the same manner as in the case of the blanks. The linear portion of this plot was also extrapolated to zero indicator current, and the generation time intercept was designated as the titration time. The blank time was subtracted from the titration time to give the corrected titration time. The corresponding weight of arsenic was calculated from the values of the corrected titration time and the rate of geneiation.
0.10 V F in potassium iodide. The desired pH value was obtained by taking 45 ml. of either 0.5 VF sulfuric acid, 0.10 VF sodium hydroxide, 1.0 VF sodium hydroxide, or one of the following buffer solutions: pH 2, sulfate-monohydrogen sulfate; p H 3, dihydrogen phosphate-phosphoric acid; pH 5, acetate-acetic acid; pH 7 , monohydrogen carbonate-carbonic acid; pH 8, monohydrogen phosphate-dihydrogen phosphate; pH 9 and pH 10, dihydrogen borate-boric acid; p H 11, carbonate-monohydrogen carbonate; pH 12, phosphate-monohydrogen phosphate.
DISCUSSION
Potential Difference between Indicator Electrodes. In order to minimize effects due to changes in the internal and external resistance of the indicator circuit, it is desirable to choose an indicator potential which is within the range wherein the curve of indicator current us. indicator potential is flattest, and which is as high as practicable within that range. This initial potential value was determ i n d in the following way: Iodine was generated in a. blank solution until an indicator current of 20 pa. was obtained when the potential difference between the indicator electrodes was 191 mv. Then this potential difference was varied, and both it and the indicator current were recorded. The data thus obtained were plotted, and the resulting graph is shown in Figure 1. This graph is essentially in agreement with similar ones obtained by Brunner ( 2 ) . Because the flattest portion of the curve lies between 100 and 160 mv., an indicator potential of 150 mv. was chosen and was used for all subsequent WOlk.
:: 2 g
2
' 10 GENERATION
Figure 2.
TIME
20 (SECONDS)
Indicator Current VS. Time of Generation at Various Hydrogen Ion Concentrations
Figure 2 shows typical data obtained a t p H 9 or less. The plots have been corrected for the presence of oxidizing or reducing agents in the reagents used in preparing the buffers, by extrapolating the linear portions of these curves t o zero indicator current, and by calling the intercept on the generation time axis zero generation time. That the curves are linear over this range of p H valucs is evident, and the current readings were stable to *0.2 microampere for a t least 3 minutes. No conclusions can be drawn from the slopes of these curves because the sensitivity of the indicator electrodes varied sufficiently over a period of a few d a y to account for the different slopes shon-n.
+ z K W
a K 0
Table I. Rate of Decrease of Indicator Current
0
L PH 10 11 12 I
0.1 02 POTENTIAL D I F F E R E N C E B E T W E E N INDICATOR E L E C T R O D E S ( V O L T S )
Figure 1. Variation of Indicator Current with Potential Difference between Indicator Electrodes Iodine concentration 2 X 10 -8 V F
Effect of pH on Indicator Current. In order to determine the effect of p H on the magnitude and constancy of the indicator current, blanks were run in solutions of various pH values in the manner described above, and plots of indicator current us. generation time were constructed.
In all these blanks, a potential difference of 0.15 volt was used, the total volume of the solution was 50 ml., and the solutions were
Indicator Current, pa. 40 25 48 37 45 40
Rate of Decreasc ra./Sec. 0.020 0.014 0.07 0.036 0.09 0.07
In solutions having pH values of from 10 to 12 the indicator current was not stable, and its rate of decrease could be determined. A t given values of pH and current these rate measurements agreed to an accuracy of * 15%, and a series of such measurements is shown in Table I. TThen indicator current readings were made with only 10 seconds being allowed after each period of generation for stabilization of the current, the curves of indicator current us. generation time were found to be linear. Indeed, a t all pH values of less than 12 only between 1 and 5 seconds were required after each period of generation for the indicator current to reach a value within 0.2 microampere of the value taken. In solutions in which the concentration of sodium hydroxide was greater than or equal to 0.10 V F , the rate of decrease of the indi-
334
ANALYTICAL CHEMISTRY
catoi current was so high that no valid determination of a 1lne:ri dependence could be made-for example, in 1.0 VF sodium hydroxide solution, the rate of decrease of this current was 0 25 microampere per second. I t is reasonable to assume that these decreases of indicator ruirent with time a t pH values of 10 or greater are due to decreases in the total iodine-Le., zero oxidation state iodine-concentration, in view of the fact that the indicator current is linearly propoitional to this total iodine concentration. Three mechanisms are possible by which this might occur: Iodine vapor may be volatilized from the solution; iodine may hydrolyze, forming iodide and hypoiodite or iodate ions; and iodine may be oxidized to iodatc by atmospheric oxygen, as postulated by hlchlpine ( 3 ) . The first of these mechanisms cannot be the rate-determining one, because no comparable rate of decrease of indicator current was ohserved a t higher hydrogen ion concentrations. The conclusion of McAlpine that air oxidation of the iodine is responsible for its disappearance in solutions of low hydrogen ion concentrations e m not be accepted because it was found that in a solution buffered at a pH of 12 the rate of decrease of the indicator current wa, the same within =t5% when an atmosphere of ovygen was maintained above the stirred solution as nhen an atmosphere of nitrogen nas so maintained. Therefore hydrolysis is apparently responsible for the removal of iodine from solutions having pH values of 10 or greater. The above observations shon. that the amperometric method for the determination of excess iodine may be used in solution., in which the hydrogen ion concentration is greater than or equal to mole per liter and with but little difficulty may be extended to include solutions in which it is equal to 10-10 mole per liter The reproducibility with which the rate of decrease of the indicator current could be observed suggests that this amperonietric procedure could be used to study the rate of hydrolysis of iodine a t very l o v concentrations.
solutions buffered to a pH of 8 and with an indicator potential of 0.150 volt, and then plotting indicator current against generation time. Though the slopes of the threc lines shown differ noticeably from one another, it was found that if more blanks were run, the slopes of successive blanks would tend to become equal. When a sufficient number of blanks, on the order of three or four, had been
Table 11. Confirmatory Titrations Sllinhel
Ib 1 2
'
Taken 69,43
3
AV.
1 ;
618,O
3 4
.Iv.
111 1
1233.1
9
3
-i .1v ,
IIIC 1 2 3
1233 1
4
.Iv.
Arsenic, Microgranis round Error 69 3 69,4l 69.3: 69.38
E ! 618.9 619..: 618.9
-0.12 0.05 -0.08 -0.05
% Error -0.17 0.07 -0.12 -0.07
0.6 0.9 0.6 1.3 0.9
0.10 0.15 0.10 0.24 0.15
203, g
0.0
0.00
12.53. 1254.: 1253.6 1263 4 1253, R 1253.7
0.8
0.9 0.5 0.3 0 5 0.6
0.06 0.07 0.04 0.02 0.04 0.05
1253. 1252.: 1253, 5 1252., 1263 .'I
0.4 -0.3 0.4 -0.4 0.0
0.03 -0.02 0.03 -0.03 0.00
~
Roman numerals indicate stock solutions used. Low rate of generation used for this group only; high rate used for all o t h e r groups here represented. C I n this g r o u p only, no atmosphere of nitrogen above solution undergoing titration. a b
0
,. Y
~
a
L B o C
u 3
N
c
z
W
C
: e C
t-
u 5 0
GENERATION
Figure 3.
TIME
(SECONDS1
Indicator Current T S . Generation Time for a Group of Blanks
Effect of Iodide Concentration on Indicator Current. I n order to determine whether any significant effect was caused by increasing the iodide concentration above the usual value, the following experiment was carried out: In a blank solution, enough iodine was generated to give an indicator current of 25 pa. Eight grams of solid potassium iodide were dissolved in the solution and the indicator current was again noted. The indicator current was found to increase 1 or 2 pa. Thus changes in iodide concentration above 0.1 V F have only a small effect upon the indicator current. Accuracy of Determination of Iodine Concentrations. Figure 3 shows the curves obtained by running three successive blanks in
run, total concentrations of iodine between 2 X 10-6 and 1 X equivalent per liter could be determined to an accuracy of *l%. However, the slopes of the curves shown have only a second-order effect on the blank time, so that for the purposes of this titration only three or a t most four blanks were run to determine each blank time. Effect of pH on Coulometric Titration. The upper limit of the hydrogen ion concentration range a t which arsenic may be coulometriexllj- titrated using iodine as the intermediate is determined by the rate of the reaction betx-een iodine and arsenious acid. -At a pH value of 7 , this rate was found to be sufficiently ~ of iodine generation. However, hjgh for titrations a t the 1 0 1 rate the high rate of iodine generation could not be used a t this pH value because eycess iodine accumulated in the solution before the equivalcrice point, causing a premature rise in the indicator current. At a hydrogen ion concentration of mole per liter neither rate of generation could be used. I n order to investigate the possibility of carrying out the selective titration of other reducing agents in the presence of tripositive arsenic, experiments were made in which it was found that in solutions 0.1 VF in perchloric acid as much as 1100 micrograms of tripositive arsenic had no effect on the blank time or on the indicator current-generation time curve. At a p H of 3 the iodine reacted with the tripositive arsenic slowly, but a t such a rate that a stable indicator current could not be obtained. The lower limit of hydrogen ion concentration is determined by the effect of the rate of hydrolysis of iodine on the determination of the end point; this effect has been discussed above.
335
V O L U M E 2 2 , NO. 2, F E B R U A R Y 1 9 5 0 Confirmatory Titrations. Table I1 shows the data obtained from confirmatory titrations which were carried out as described above. I t is seen that the titration is accurate to approximately *0.294 over the range of quantities of arsenic investigated. I t is believed that the factors limiting the accuracy of these titrations are: the oxidation of tripositive arsenic by atmospheric oxygen, the preparation and dilution of the standard solutions, the determination of the end point, and the measurement and control of the generation current. Preliminary titrations of 70-microgram quantities of tripositive arsenic, over which no atmosphere of nitrogen was maintained, gave results as much as -0.6% in error. In addition, two groups of titrations (111) were made with t,he same stock solutions ant1 carried out under as closely identical conditions as possible, except that the second group was not done under an atmosphere of nitrogen; the average value for the first group is 0.6 microgram greater than that for the second. These effects are attributed to air oxidation, and in all the other titrations shon-n in Table I1 :in atmosphere of nitrogen v a s provided. It is thought that a source of error in the above data was the accuracy with which standard solutions could be prepared. There is considerably better agreement between titrations within any one group than between the average titer of a group and the calculated value of the titer for that group.
The maximum spread in the values obtained for the blank time as 0.04 second for the high rate of generation and 0.2 second for the low rate. Two factors are involved in this spread: the accuracy with which the measurements and extrapolations could be made, and the random contamination of blank solutions. I t is probable that essentially these s:rme factors enter into the detcrmination of the end point. The arcuracy to which the generation cur1Pnt could be determined n as about * 0.0370, and it could easily be controlled 15 ithin this limit. LITERATURE CITED
Brown and Swift, J . Am. Chpm. Soc., 71, 2717' (1949 . Brunner, Z. p h y s i k . Chem., 56, 391 (1906). XIcAlpitie, J. Chem. Education, 26, 362 (1949). Meier, Myers, and Swift, J . Am. Chem. Soc., 71, 2340 [ 1949). ( 5 ) Myers and Swift, Ibid.. 70, 1047 (1948). (6) Sease, Xiemann, and Swift, -4s.~~. CHEJI.. 19, 197 (194;). (7) Szebelledy and Somogg-i, Z . arid. Chem., 112, 313 (193R1.
(1) (9) (3) (4)
( 8 ) Ibid., p. 385. (9) Ibid.. D . 391. ( l O j I b i d . ; 400.
i.
(11) Wooster,
Farrington, and h i f t , AXAI..CHEM.,21, 14.57 (1949).
RECEIVED August 13, 1949. Contribution 1317 from the Gates and Crellin Laboratories of Cheriiiatry, California Institute of Technology.
Effect of Ammonium Salts on Determination of Nicotine C. L. OGG, C. 0. m-ILLITS, AND CONSTANTINE RICCILTI Eastern Regional Research Laboratory, Philadelphia 18, P a .
Ammonium salts interfere in the determination of nicotine by the silicotungstic acid method by retarding the precipitation. Graphs are presented to show the effects of concentrations of nicotine, salt, and reagent and time of digestion on recovery of nicotine. Procedures are proposed to minimize the effect of the salt.
D
ETERMISATIOX of nicotine by the gravimetric silicotungstic acid procedure is t,heofficial method of the Association of Official Bgricultural Cheniists ( 1 ) . Hon-ever, if other all
