Amperometric Titration and Voltammetric Determination of Iodide With Rotated Platinum Wire Indicator Electrode I. M. KOLTHOFF AND JOSEPH JORDAN1 School of C h e m i s t r y , University of Minnesota, .Minneapolis 14, M i n n . The determination of to 10-3AM iodide, by measurement of its anodic diffusion current a t the rotated platinurn wire electrode, was carried out a t a potential of +0.65 volt US. SCE in a supporting electrolyte of 1iM sulfuric acid plus 0.01M hydrocyanic acid. .ipplication of the voltammetric charactcristics of the iodine-iodide system and various oxidizing and reducing agents led to the development of several amperometric titration methods for the determination of iodide in very dilute solutions. Iodide can be titrated with standard potassium iodate or permanganate in sulfuric acid solution to a n iodide or an iodine cyanide end point a t a poten-
I
T HAS been shown ( 4 )that the anodic diffusion current of iodide a t the rotated platinum wire electrode under certain conditions is proportional to concentration. From the knosledge of the voltammetric characteristics of the iodide-iodine system, combined with a study of the behavior of commonly used oxidizing and reducing agents, it is possible to predict suitable conditions for an accurate, rapid, and simple determination of iodide a t high dilutions by arnperometric titration, using the rotated platinum n ire electrode as indicator electrode. Several oxidimetric and reductimetric reagents have been investigated R ith this purpose in view. Methods have been developed for the direct titration of iodide with iodate, cerium(IV), or permanganate a t concentrations greater than 10-6.W. The sensitivity range of the amperometric titration can be extended to trace concentrations of iodide of the order of 0.3 to 0.4pM by conversion of the iodide to iodate and titration of the latter. At the rotated platinum wire electrode the reaction 1 2 2e$ 2 I- proceeds reversibly in acid solution ( 4 , 12). A solution containing both iodine and iodide yields a cathodic and an anodic ~ a v with e well defined diffusion currents as illustrated in Figurt 1 It is seen (Figure 1, curve I) that a solution of 5 X 10-5M iodine and lO-4M iodide in 1Jf sulfuric acid gives a composite wave with a half-wave potential of f0.5 volt. (All potential values reported are referred to the saturated calomel electrode.) The cathodic diffusion current of iodine extends over a potential range between 0 and +0.35 volt us. SCE. Because of the marked volatility of iodine in aqueous solution in the absence of iodide (or much bromide or chloride), it is not practicable to determincx iodine voltammetrically unless a complex former is present to decrease the volatility. With proper precautions against losses, the diffusion current of iodine in water (in the absence of complex formers) is proportional to concentration. In the presence of an excess of iodide practically all the iodine is converted into the triiodide ion, the diffusion current of which is equal to that of an equimolar concentration of free iodine ( 4 ) . The anodic diffusion current of iodide (oxidation to iodine) extends over the potential range between +0.60 and +0.75 volt 2,s. SCE. At still more positive potentials, an ill-defined second anodic wave is obtained corresponding to the irreversible electrooxidation of iodide (or iodine) to I+. Iodide can be determined
tial of +0.65 volt US. SCE. Titrations a t a potential of f0.2 volt US. SCE with ceric cerium in the presence of hydrocyanic acid and with permanganate i n the presence of either acetone or hydrocyanic acid yield good results, but are limited in applicability because other halides interfere. T w o precise methods, effective for extremely small concentrations of iodide, are based on the oxidation of iodide to iodate and the titration of iodate. Using the amperometric titration methods, a precision and accuracy of 0.2% were attained in the to 10-3,%fconcentration range and of 1 to 5 % a t concentrations as low as 10-7 to 10-6jM.
a t 25 C. by measurement of its anodic diffusion current in 0.1M perchloric acid a t a potential of +0.70 volt in a range of concentrations b e k e e n 10-5 and l O - 3 M , the precision being of the order of 2% in a 10-4M solution ( 4 ) .
If
1
+
1
On leave from the Hebrew University, Jerusalem, Israel.
I
I
1 +l.O
*O.S
POTENTIAL OF INDICATOR
J
0.0
ELECTRODE, VOLT 'e.SCE
Figure 1. Current-Voltage Curves Obtained a t Rotated Platinum Wire Electrode I. 5 X 10-6M iodine 11.
111.
+
10-4M potassium iodide in 1M sulfuric acid 2 X lO-4M potassium iodide in 0.1M hydrochloric 1M potassium chloride acid 2 X lO-4.M potassium iodide in 0.1M hydrochloric
+
,,-:a
1', II', 111'.
Residual current curves
In the present work it was found that the voltammetric determination is more precise (1% in a 10-4.V solution) when carried out in 1M sulfuric acid (instead of 0.1M perchloric acid) as supporting electrolyte, the diffusion current being measured a t a potential of +0.65 volt. Furthermore, the range of the method was extended to iodide concentrations as low as IO-eM by using an indicator electrode of an area about 10 times larger than that of the platinum wire electrode previously described. Bromide does not interfere, even if present in a concentration of 2M. I n the presence of chloride, high and erratic results are obtained. This is due to the fact that in the presence of much
1833
1a34
ANALYTICAL CHEMISTRY
chloride (1M and up), iodide yields an anodic wave with a poorly defined diffusion current ( 4 ) corresponding to the electrode reaction I2C1--+ICI-2 2 e. rln example of a wave of this type (in 1M potassium chloride plus 0.1M hydrochloric acid) is shown in Figure 1(curve 11) together with a current-voltage curve of IO-'M iodide obtained in the presence of 0.1M hydrochloric acid only (curve 111). The first of the two diffusion currents obtained in 0 . M hydrochloric acid with no additional chloride present (2 I-+ 1 2 2 e) is ill defined; the second (12 2C1- 4 2IC1 2 e ) is well defined, but its magnitude varies considerabIy with changes in chloride concentration.
+
+
+
+
+
0
.eo q
*.
!W
z
a a
3 0
+0.45 to +0.50 volt and are not affected by the presence of even 251 bromide or chloride. Iodoacetone and iodine cyanide yield no cathodic waves a t the rotated platinum electrode within the \?-orking range of potentials. Iodate, permanganate, and cerium(1V) were considered as poesible reagents for the direct oxidative titration of iodide to an amperometric end point. Iodate in 1JI sulfuric acid (also in the presence of 0.0lM hydrocyanic acid or of 20% acetone) yields no reduction current a t the rotated platinum wire electrode a t potentials more positive than +0.3 volt. Current-voltage curves of 10-~rM cerium(1V) (one-electron reduction) and of l O - ' M permanganate (five-electron reduction) in 151 sulfuric acid (not reported previously in the literature) are shown in Figure 3. Both solutions yield cathodic waves with diffusion currents extending over a potential range between +0.5 and 0.0 volt. The cathodic diffusion currents, measured a t a potential of +0.2 volt, are proportional to the concentration of cerium(1V) and permanganate, respectively, between 2 X 1O-sM and 10-3.11. The half-wave potential is +0.85 volt for the cerium wave and +0.82 volt for permanganate. The presence of 0.01M hydrocyanic acid does not affect the waves.
.40
50 40
3 I
4 1.0
I
1
00
05
POTENTIAL OF INDICATOR
/
I
ELECTRODE, VOLT vs.SCE
Figure 2. Current-Voltage Curves of 2 X Potassium Iodide Solutions at Rotated Platinum Wire Electrode
lO-4M
++
I. In M sulfuric acid 0.01M potassium cyanide 11. In 1M sulfuric acid 20% v./v. acetone 1', 11'. Residual current curves
Iodide can be determined voltammetrically in the presence of chloride if hydrocyanic acid is added to the supporting electrolyte. -4currenbvoltage curve of 10-4M iodide in 1M sulfuric acid plus 0.01M hydrocyanic acid is shown in Figure 2 (curve I). A single anodic wave is obtained which corresponds to the two-electron transfer IH C N & I C N -b H+ 2 e (1)
+
+
and which is not affected by the presence of chloride. Current-voltage curves of iodide in the presence of acetone have not been reported previously. It was found in this investigation that iodide yields a well defined anodic wave in 111.1 sulfuric acid plus 20% v./v. acetone. An example is plotted in Figure 2 (curve 11). It is seen that the diffusion current in the presence of acetone extends over the same range of potentials as in the presence of hydrocyanic acid (Figure 2, curve I)-Le., from +0.65 to +0.95 volt. The diffusion current in 20% acetone is proportional to iodide concentration in the range between 10-5 and l O - 3 X . The oxidation of iodide in the presence of 20% acetone and 131 sulfuric acid occurs quantitatively according to the over-all equation
I-
+ CHaCOCHs e CHJCOCHJ + H + + 2 e
(2)
The diffusion current in the presence of 20% acetone is about one sixth larger than that obtained in 0.01M hydrocyanic acid a t the same iodide concentration. Acetone has a similar effect on the wave heights of permanganate and cerium(1V). This effect cannot be accounted for by changes in the viscosity of the solution. While an increase in diffusion current would normally correspond to a decrease in viscosity, it was found the viscosity a t 25' C. of 1M sulfuric acid in water and in 20% acetone is equal to 1.08 and 1.38 centipoises, respectively. The waves in acetone and in hydrocyanic acid have a half-wave potential of about
POTENTIAL OF INDICATOR
Figure 3.
ELECTRODE, VOLT VS.SCE
Current-Voltage Curves in 1M Sulfuric dcid
I. 10-dM cerium(1V) 11. 10-dM potassium permanganate Residual current
For the indirect determination of iodide by transformation to iodate and titration of the iodine liberated by an excess of iodide ( I I ) , thiosulfate proved to be a suitable amperometric reagent in agreement with recent findings in the literature ( 1 , 5 ) . At POtentials more negative than +0.4 volt, thiosulfate did not give an anodic current under the experimental conditions. OUTLINE OF METHODS
The iodide-iodine cyanide diffusion current is ideally suited for the voltammetric determination of iodide a t the rotated platinum wire electrode. It is recommended that the current be measured a t a potential of +0.65 volt a t 25" C. and that an aqueous solution of 1M sulfuric acid plus 0.01M potassium cyanide be used as supporting electrolyte. To select suitable conditions for the amperometric titration of iodide, application was made of the classical stoichiometric oxidation reactions with iodate, cerium(IV), and permanganate. The direct oxidation of iodide must be carried out in acid solution and involves one or two equivalents. In the presence of acetone or cyanide, iodide is oxidized to the unipositive state (7-10). In the absence of acetone or cyanide, cerium(IV), permanganate, or iodate in acid solution oxidize iodide to elementary iodine. HOWever, the amperometric titration of iodide in dilute solutions to an iodine end point could not be carried out satisfactorily under the
1835
V O L U M E 2 5 , NO. 12, D E C E M B E R 1 9 5 3
I n the absence of bromide, a t chloride concentrations below 0
0.131 and in the presence of 0.01M hydrocyanic acid, iodide can
be titrated with either permanganate or with cerium(1V) a t a potential of +0.2 volt, where the diffusion current of excess reagent is measured and where iodide does not yield an oxidation current. Illustrations of the corresponding titration curves are given in Figure 5 (curves I and 11). However, since bromide and much chloride interfere, the titration a t +0.65 volt with permanganate in the presence of bromide is much to be preferred because of its general applicability. Iodide in 131 sulfuric acid can also be titrated with permanganate in the presence of 20% v./v. acetone a t a potential of +0.2 volt (Figure 5 , curve 111); this procedure, too, fails in the presence of bromide or more than 0.1M chloride. The use of cerium( IV) as reagent instead of permanganate is not practicable on account of a relatively slow reaction rate.
$ 0
-20
W -I
U
0 W
a W W
3
.*o
4
3 0
4 0
0.8
10
I5
ml. REAGENT ADDED
Figure 4. Amperometric Titration Curves of 50 311. of 10-4.M Potassium Iodide Solution with 2 X 10-3iM Potassium Iodate at +0.65 Volt vs. SCE 1. In 1M sulfuric acid In 1M sulfuric acid
11.
+-
40
9
5W
O.01M potassium cyanide
LL 0 W
experimental conditions with cerium( IT)or permanganate because the reaction rates were too slow. Good results were obtained using standard potassium iodate as titrant and 1-11sulfuric acid as supporting electrolyte. The reaction proceeds according to the equation 51-
+
103-
+ 6H'
-+
312
+ 3H20
+
103-
+ 3HCN + 3 H f 2 3 I C S + 3HAO
20
W +
0
2 W
0
(3)
The titration was carried out a t a potential of +0.65 volt corresponding to the iodide-iodine diffusion current. A titration curve of this type is shown in Figure 4 (curve I). If bromide is also present, it reacts with the excess iodate to form elementary bromine which yields a cathodic current a t the working potential. I n the presence of chloride, the titration line of iodide in 1.V sulfuric acid is distorted, because of the effect of chloride on the current-voltage curve of iodide (Figure l, curves I1 and 111). These interferences can be eliminated by making the supporting electrolyte 0.01,M in hydrocyanic acid. I n a medium of 1M sulfuric acid plus 0.0131 potassium cyanide, the ovidation of iodide proceeds to the unipositive state 21-
9
(4)
A titration curve of iodide with iodate in the presence of hydrocyanic acid is plotted in Figure 4 (curve 11). In this case, bromide (in concentrations as high as 2M) does not interfere because it is oxidized by excess iodate to bromine cyanide which does not yield a cathodic current under these experimental conditions. Chloride has no effect on the titration line, because iodine cyanide is considerably more stable than iodine chloride. Under the same conditions (in 1.11 sulfuric acid plus 0.01M potassium cyanide and a t a potential of 0 65 volt) permanganate and cerium (IT)can also be used as titrants, even though their diffusion currents are not attained a t the working potential of the indicator electrode. During the titration the anodic diffusion current of iodide decreases as in Figure 4, curve 11. At the end point the current is 0. After the end point a cathodic reagent current is obtained xhich, however, is not proportional t o the concentration of the excess titrate. The titration curve with permanganate becomes entirely similar to curve I1 in Figure 4 when the medium is 0.0lM in hydrocyanic acid, 0.531 in potassium bromide, and 1M in sulfuric acid. Excess permanganate is rapidly reduced by the bromide, with bromine cyanide being formed. Cerium(1V) is not a satisfactory reagent in this titration because it reacts too slowly with bromide.
L 00
I
I
I
os
10
I 5
I
mi. REAGENT ADDED Figure 5. Amperometric Titration Curves at a Potential of +0.2 Volt cs. SCE
+ + +
50 ml. of lO-4M potassium iodide in 1M sulfuric acid 0.01M potassium cyanide titrated with 2 X 10-841 potassium permanganate 0.01M 11. 50 ml. of 10-4M potassium iodide in 1M sulfuric acid potassium cyanide titrated with 10 - 2 M cerium(1V) 20% 111. 50 ml. of lO-4M potassium iodide in 1M sulfuric acid v./v. acetone titrated with 2 X 10-aM potassium permanganate IV. 50 ml. of 2 X 10-6M potassium iodate in 0.02.M sulfuric acid ewe98 potassium iodide titrated with 5 X IO-3.M sodium thiosulfate
I.
+
Reaction 3 was used in an indirect determination of iodide. The iodide was converted into iodate using an excess of chlorine water as oxidant. Subsequently, the iodate was titrated amperometrically by one of the following methods: Titration with standard potassium iodide solution in 1'11 sulfuric acid a t a potential of +0.65 volt. The corresponding titration curve is the converse of the one obtained in the titration of iodide by standard iodate in the same medium (Figure 4, curve I). The end point is determined by the excess reagent line of iodide-Le., by the iodide-iodine anodic diffusion current. Reduction of iodate to iodine by an excess of potassium iodide. The liberated iodine is titrated with standard thiosulfate solution a t a potential of +0.2 volt corresponding to the iodine-iodide cathodic diffusion current. An evample of a titration curve of this kind is plotted in Figure 5 (curve IV). EXPERIMENTAL
Materials. Reagent grade chemicals were used throughout. Solvents (conductivity water and acetone) were redistilled immediately before use. Cerium(1V) solutions were prepared by standard methods from ceric ammonium sulfate obtained from the G, Frederick Smith Chemical Co. Apparatus and Technique. A simple electrolysis cell, made of a 100-ml. borosilicate glass beaker ( 3 ) and a Hume-Harris saturated calomel reference electrode ( 2 ) were used in all the experiments. Current-voltage curves were recorded with a Sargent Model XXI polarograph in a thermostated water bath a t 25'
ANALYTICAL CHEMISTRY
1836 Table I.
Mediuq 1M HISO&,0.01M H C S 1M HzSOa
Concentration Range, Precision, and Accuracy of Iodide Deterniinations Approximate Sensitivity Limit Iodide, y per 10 MI.of C r - , Moles per Liter Sample Ab Be Ab B 10-5d io-ee 13 1-2 2 X 10-6 2 X 10-8 25 2-3
Method Neasurement of id at +0.65 volt Amperometrictitrationwithiodate a t +0.65 volt 1M HISOL 0 . 0 1 M H C S AinDerometric titration with iodate a i i-0.65 volt 1M HzSOd, 0.01M H C N Amperometric titration with Ce(1V) a t 4-0.2 volt 1M HzSO4, 20% acetone Amperometric titration with permanganate a t + 0 . 2 volt 1M HzSO,, 0 . 0 1 M H C S Aniperometric titration with perinanganate a t f0 . 2 volt 1M HzSO,, 0 01M H C S , Arnperometric titration with per0 . 5 M KBr manganate a t -0.65 volt 1M H~SOI Oxidation t o iodate and amper- 4 X ometric titration of 10s- with iodide a t +0.65 volt Oxidation t o iodate, addition of 3 X excess iodide, and titration of liberated iodine with thiosulfate a t + 0 . 2 volt 4 I n 5 X 10-SM solution for electrode A and in 5 X 10-6 solution b Small electrode 0 Large electrode.
Precision',
%
Ab 1
0.4
Bo 5 2
Sccuracy",
Maximum Tolerance To T O chloride bromide
%
Ab
1
0.5
BC
5 3
m
5 X 10-4M
m
Sone
10-5
io-
13
1-2
0.2
I .
0.~ .2
1
m
m
10-5
10-1
13
1-2
0.3
3
0.3
3
0.1X
10 X [ I -
10-5
10-8
13
1-2
0 2
1
0.1
1
0.1M
Xone
10-6
10-8
13
1-2
0.3
1
0.2
1
0.l.M
Kone
10-5
10-8
13
1-2
0.1
0.2
m
m
10-6
4 X 10-7
i)
0.5
0.4
1
10-1
3 X 10-7
4
0.4
0.2
0.2
for electrode B.
=t 0.02' C. Diffusion currents were measured a t the same temperature, wing a manual setup described previously (6). Amperometric titrations were carried out a t room temperature with a simple circuit. Potentials were applied using an uncalibrated slide-wire potentiometer which was adjusted to the desired value by compensating against a Leeds and Xorthrup type K1 potentiometer a t the electrodes. The potentiometer can be eliminated by short-circuiting the electrolysis cell with a half cell of suitable potential. The current was measured with a General Electric portabletype mirror and scale galvanometer which had a maximum sensitivity of about 10-*pa per mm. The working sensitivity of the current measuring instrument was adjusted with an Ayrton shunt and calibrated by measuring the potential drop across a standard resistance connected in series. Whenever current readings were taken, the cell circuit was closed for a brief period of time with a tap key. Between readings the circuit was left open to avoid depletion of the solution bv electrolysis. It has been shown ( 3 ) that significant errors may be caused by depletion upon continued electrolysis, especially when large electrodes and small volumes of solution are used. Two indicator electrodes made of platinum wire about 0.5 mm. in diameter but of different area were used. Thev were approximately 4 and 35 mm. in length and yielded diffueion currents of 251 i.1 and 2232 =t 20ga. per millimole per liter, respectively, for the thallous-thallic oxidation in 0.l.V sodium hydroxide. The electrodes were sealed into glass, and electrical contact was made with the aid of mercury. The platinumglass joints were carefully annealed, because irregularities at the metal-glass interface give rise to abnormally large residual currents. The electrodes were stored in 10.11 nitric acid when not in use. They were rinsed with water shortly before the start of an experiment and then immersed for 10 minutes in the experimental solution before a potential was applied. On applying an initial potential 2 t o 3 minutes elapsed before a steady current could be measured; thereafter the electrode responded instantaneously to changes in applied potential (\%hen measuring curreqt-voltage curves) and concentrations (in amperometric titrations). A synchronous motor was used to rotate electrodes a t 600 r.p.m. The electrode wire was bent eccentrically t o provide sufficiently constant stirring for rapid mixing during titrations. Residual currents appeared to be negligible in the working range of potentials. Standard solutions of titration reagents in the 10-3 and 10-4X concentration range were prepared shortly before use by dilution of 10-2.V solutions. The 0.01A' solutions were standardized daily by classical procedures The cerium(1V) solutions were made up in 1.11 sulfuric acid. Standard potassium iodide solutions were made oxygen-free by deaeration with purified Linde nitrogen and were stored under nitrogen. Procedure. In each titration, 25 or 50 ml. of sample was used Normalities of reagents used for titration were about 50 times those of the samples and were dispensed from 5-ml. precision microburets. For direct titrations with permanganate and cerium(1V) an aliquot of the sample was made 1.W in sulfuric
~~
~~
d
~
5
x
10-4.61 m
&-one m
Upper limit of concentration range is 10-aM. Upper limit of concentration range is 10-4M.
acid and 0.01X in potassium cyanide (or 20% v./v. in acetone) and titrated immediately. For the determinations involving the oxidation of iodide t o iodate, the sample solution was first neutralized with sulfuric acid to a methyl orange end point and then made 0.02M in sulfuric acid. Chlorine water was added until a permanent yellow color was obtained. The solution was boiled for 3 minutes, cooled, and deaerated for 30 minutes with nitrogen to remove excess chlorine. Subsequently, an aliquot \vas either acidified to 131 in sulfuric acid and titrated with standard potassium iodide, or excess potassium iodide was added (to make the solution about 0.1X in iodide) and the liberated iodine titrated with standard thiosulfate When using the smaller electrode, the reagent blanks proved to be negligibly small. In titrations with the large electrode, blanks were equal to about 593 of the reagent required for the sample of the smallest concentrations determined These blanks were applied as a correction to the evperimental results. R E S U L T S AND DISCUSSION
The methods were tested over a wide range of iodide concentrations from 10-3-lf dovn to extreme dilutions where the sensitivity of the current measuring instruments became the limiting factor. The lower limit to the applicability of the various amperometric titration procedures 'rvas set a t concentrations a t which the maximum current reading became less than a full scale deflection of the galvanometer a t maximum sensitivity (about 2 ~ ) . The large electrode carried the range of all methods to iodide concentrations smaller by one order of magnitude as compared to the small electrode. All results presented in the figures were obtained using the small indicator electrode. Concentration ranges, precision, and accuracy as determined with both electrodes are summarized in Table I. Interferences tested for were bromide and chloride. The limit of tolerance towards these halides is also given in the table. The voltammetric determination of iodide by the measurement of its diffusion current in the presence of cyanide can be carried out Lsithout interference from bromide and chloride. The method is limited to iodide concentrations between and 10-3Jf and its precision and accuracy are relatively poor. The titration methods give considerably better accuracy and precision. From the viewpoint of applicability in the presence of bromide and chloride, the titration a t +0.65 volt in a supporting electrolyte of 0,OlJf hydrocyanic acid plus 1Jf sulfuric acid with iodate as reagent is recommended. Permanganate can be used as well when the titration medium is made also 0.5X in bromide. The determination of iodide after oxidation to iodate and titration of the
V O L U M E 25, NO. 1 2 , D E C E M B E R 1 9 5 3
1832
liberated iodine with thiosulfate has the additional advantage that for a given amount of iodide the required amount of reagent equivalents is three times as large as in previous methods. Although the titration of iodide in concentrations greater than 10-3M has not been discussed herein, the methods described can be used with necessary changes for larger concentrations. ACKNOWLEDGMENT
Thanks are due Richard W. Ramette for measuring the current-voltage curves of permanganate and cerium. LITERATURE CITED
(1) Harris, E. D., and Lindsey, A . J., A n a l y s t , 76, 647 (1951).
(2) Hume, D. N., and Harris, W. E., IND.ENG.CHEM.,ANAL.ED.. 15, 465 (1943). (3) Kolthoff, I. M., and Jordan, J., J. Am. Chem. SOC.,74, 382 (1952). (4) Ihid., 75, 1571 (1953). ( 5 ) Kolthoff, I. M., and Lingane, J. J., “Polarography.” 2nd ed.,
Vol. 11, p. 946, New York, Interscience Publishers, 1952. (6) Laitinen, H. A., and Kolthoff, I. A f . , J. P h y s . Chem., 45, 1061 (1941). (7) Lang, R., Z.anorg. u. allgem. Chem., 122, 332 (1922). ( 8 ) Ibid., 142, 229, 279 (1925). (9) Ihid., 144, 75 (1925). . (10) Lewis, D., IND.ENQ.CHEM.,ANAL.ED., 8, 199 (1936). (11) Sadusk, J. F., Jr., and Ball, E. G., Ibid., 5, 386 (1933). (12) Vetter, K. J., Z . p h y s i k . Chem., 199, 22 (1952). RECEIIVED for review July 2, 1953.
Accepted September 17, 1953.
Titration of Weak Bases in Acetic Anhydride Solvent Mixtures JAMES S. FRITZ
AND MYRON
0. FULDA,
Zowa S t a t e CoZZege,
Ames, Zowa
The purpose of this investigation was to study the effect of titrating bases in organic solvents from which the last traces of water have been removed by the addition of excess acetic anhydride. For most tertiary amines and alkali metal salts, this increases considerably the sharpness of the break and rise in potential at the end point. A very large excess of acetic anhydride does not further increase the sharpness of the end point. The method proposed gives excellent results for tertiary amines including nitrogen heterocyclics of the purine, pyridine, pyridone, and thiazole type. Primary, secondary, and a few heterocyclic amines cannot be titrated. This method offers increased accuracy and broadens the scope of titrations in nonaqueous solvents to include weaker bases than formerly could be titrated.
I
T H.4P been recognized (3-5) that the presence of water de-
creases the sharpness with which bases can be titrated in acetic acid and other nonaqueous solvents. Because of this, the water in perchloric acid from which the titrant is prepared is often removed by addition of a calculated amount of acetic anhydride Some n ater still remains, however, since reagent grade glacial acetic acid commonly contains 0.3 to 0.5% water. The purpose of the present investigation was to study the titration of weak bases in nonaqueous solvents completely free from water. These conditions are practically obtained by titrating in solvents such as acetic acid or nitromethane containing 5 to 20% acetic anhydride. Titrations of tertiary amines in acetic acid containing acetic anhydride have been carried out (1, 7 ) , but no data are available concerning the effect of the acetic anhydride on the titration curve. REAGENTS AND SOLUTIOYS
Acetic acid, ACS grade glacial. Acetic anhydride, .4CS grade. Nitrobenzene, Matheson. Nitromethane, Eastman or illatheson. 1,s-Diphenyl-3-pentadieneone (dibenzalacetone). -4 0.1% solution in acetic acid. 1-Naphtholbenzein. -40.141, solution in acetic acid. Neutral red, 0.17, solution in acetic acid. Perchloric acid. A 0.1s solution is prepared by miung 8.5 ml. of i o to 72% perchloric acid n i t h enough acetic anhydride to remove all of the water. Let this mixture stand overnight, then dilute to 1 liter with acetic acid. Standardize against potassium acid phthalate. Triphenylmethanol (triphenylcarbinol). 0.1% solution in nitromethane. Samples were analyzed as received. Most were 98 to 100% pure. PROCEDURE
Dissolve a 0.3- to 0.8-meq. sample in 20 ml. of 4 to 1 nitromethane-acetic anhydride. Substances not readily soluble in
M O
,
I ML
1
I
