Determination of Iodide by Oxidation with Nitrous Acid JOHN K. JOHANNESSON Wellington City Council laboratory, Wellington, New Zealand
Iodide i s oxidized quantitatively to positive iodine-i.e., I+CI--by nitrous acid at high concentrations of hydrochloric acid. Addition of sulfamic acid, following the nitrite oxidation, effects a rapid and complete destruction of the excess nitrous acid, nitrogen being evolved. Following a fourfold dilution with water to avoid any subsequent aerial oxidation in the presence of concentrated acid, iodide i s then added, and the resultant triiodide ion is titrated with sodium thiosulfate using the dead-stop end point. Under these conditions each gram atom of iodide originally present yields 2 gram equivalents of free iodine for titration.
I
x AS earlier paper (j),it x a s shown that sulfamic acid did not react with iodine in acid solution, and could be used to eliminate interference due to nitrite when free iodine was being estimated colorimetrically by the o-tolidine reagent. It was believed that iodine could be estimated by oxidation to free iodine in acid solution by using sodium nitrite (Y), follou-ed by removal of the excess nitrite by reaction with sulfamic acid. The iodine so formed could then be titrated in thr usual manner with sodium thiosulfate. The reaction of sulfamic acid n i t h nitrous acid has been used for the rstimation of both sulfamic acid and nitrite (3)by gasometric (6)and amperometric (4) methods. This reaction is represented by Equation 1, and is quantitative, in that no trace of residual nitrous acid is found in the presence of an excess of sulfamic acid. NHzS020H
+ HNOz
-+
+ HzSOa + H20
(1)
In dilute acid solution oxidation of iodide proceeds according to Equation 2. "--2C+Iz
(2)
The molecular iodine produced is in equilibrium with hypoiodous acid and iodine according to Equation 3, I?
+ H,O
+
HOI
+ HI
(3)
and in acid solution there is a further reaction as in Equation 4 HOI
+ H+
-+
H2+01
(4)
-4review of the literature (6) shows that in strongly acid solution oxidation with IO3- yields positive iodine, commonly regarded as I+Cl- but which should probably be regarded as H201+.C1-in acidaqueoussolutions (1). Experiments were made to determine the effect of acid upon the quantitative nature of the oxidation of iodine with nitrite. EXPERIMENTAL
Iodide solutions were adjusted to varying degrees of concentration of acid strength with hydrochloric acid. An excess of sodium nitrite was then added, and, after a minute, excess of solid sulfamic acid was added in small portions, allowing the effervescence to cease before each addition. This was continued until a fresh addition gave no further effervescence. The sides of the apparatus were washed carefully with dilute sulfamic acid, potassium iodide was added, and the triiodide so formed was then titrated with standard thiosulfate solution. At low acid concentrations starch gave a good end point, but a t high acid concentrations it was not satisfactory; shaking with chloroform gave a satisfactory end point. The dead-stop method using two smooth platinum electrodes mas superior to other methods and was used. The addition of potassium iodide to the solution of positive iodine was essential for obtaining an accurate result with the dead-stop method. If iodide is not added, the galvanometer reading steadily rises as the titration is performed-Le., as I +is reduced to I-, which then reacts with the I + in forming triiodide ion. Finally a maximum is reached and the reading begins to decline. The titration result, however, is low. Possibly the strong oxidizing ability of the iodinium ion may oxidize thiosulfate beyond tetrathionate, thus increasing its apparent titer. It is clear from Table I that a t low acid concentrations the reaction is incomplete, but as the acid concentrations are increased, the formation of I+C1becomes quantitative. Dilution of the solution following destruction of excess nitrite avoids interference due t o atmospheric oxidation of the added I- in strongly acid solutions. Oxidation of the iodide in the pres-
ence of perchloric acid instead of hydrochloric acid did not result in the formation of positive iodine. The presence of hydrochloric acid appears essential. INTERFERENCES
Bromide is not oxidized by nitrite under these experimental conditions. Reducing substances are oxidized by the nitrous acid a t the same time the iodide is oxidized. Oxidizing substances are readily removed by treatment with excess sodium sulfite, The sulfite remaining is oxidized by the nitrite.
Table I. Effect of Acid Concentration on Degree of Oxidation of Iodide to Positive Iodine (Ten milliliters of 0.01N potassium iodide
Acid Concn.,
% V.,N. 0.i5
2.75
taken) Titration Ml.
0.01 Thio
7.65 10.1
4.0 10.0
11.3 17.9
20.0
20.0 20.0
30.0
-% Recovery 12
76.5 101 113
179 200 200
I+
38.2 50.5 56.5 89.5 100 100
SUGGESTED METHOD
The electrode system consisted of two pieces of smooth platinum foil connected through a microammeter to a 1.5-volt dry cell battery shunted n-ith radio resistances so that 0.15 volt is applied to the electrodes. The solutions are stirred magnetically. Procedure. To a suitable volume of cold solution a quantity of concentrated hydrochloric acid is added, so t h a t the solution contains 25y0 by volume of the concentrated acid. Sodium nitrite solution ( l O ~ o is ) now added slowly until an excess is present. This may be gaged by observation of the color, which will become less brown as the free iodine first liberated is converted to the pale colored iodinium ion. Addition of as much nitrite again will ensure an excess. The above reaction is rapidly complete, and the nitrite is then destroyed by small cautious additions of solid sulfamic acid. The solution effervesces and additions are continued until no further effervescence takes place. The sides of the beaker or flask are now carefully rinsed down with dilute sulfamic acid solution, and distilled water is added to increase to approximately four VOL. 30, NO. 9, SEPTEMBER 1958
1535
times the original volume of the test solution, and 1 gram of potassium iodide added. The solution is then titrated with suitable strength of sodium thiosulfate, The dead-stop method is used to detect the end point. However. in the absence of suitable electrical equipment, the titration may be conducted in the presence of a small amount of chloroform, the end Doint being ivhen the last violet color disappears from the Organic layer. Alternatively, the end point may
be observed by using starch if the final dilution is great enough to avoid oxidation by air. I n the presence of an oxidizing agent, sufficient sodium sulfite is added to ensure complete reduction, warming if necessary. After cooling, iodide is estimated.
(4)Hirozawa, S. T., Brasted, R. C., ANAL.
CHEM.25, 221 (1953). (5) Johannesson, J. K., Ibicl., 28, 1475 f19B6).
(6j-0esperJ R. E., ‘(Newer Methods of Volumetric ChemiLal Analysis,” p. 71, Van Kostrand, Yew York, 1938. ( 7 ) Reiss, R., Arzneimittel-Forsch. 6 (2), 77 (1956).
LITERATURE CITED
(1) Bell,
(2) Carson, W.S., Jr., AXAL. CHEM.23, 1016 (1951). (3) Cumming, K. AI,, Alexander, W,-4., Analyst 68, 273 (1943).
R. P., Gellis, E., J. C h e m SOC.
( L o n d o n ) 1951,2735.
RECEIVEDfor review ilugust 25> 1957. Accepted May 6, 1958.
Colorimetric Determination of Tetra methy1phosphoniu m Ion JAMES KOLMERTEN and JOSEPH EPSTEIN Sanitary Chemistry Branch,
U.S. Army Chemical Warfare Laboratories, Army Chemical Center, Md.
b Micro quantities of tetramethylphosphonium chloride (TMPC) can b e estimated colorimetrically as the orthophosphate ion b y oxidation in dilute aqueous solution with ammonium persulfate. Chloride ion interferes with the oxidation in acid or neutral solution, but not in alkaline solution.
T
conversion of various orgenophosphorus compounds to the orthophosphate ion in dilute aqueous solution with ammoniuni persulfate is under investigation in these laboratories in connection with the analytical program. The usefulness of ammonium persulfate for conversion of the nerve gases Sarin (isopropylmethylphosphonofluoridate) and Tabun (ethyl dimethylphosphoramido-cyanidate) to the orthophosphate ion was discovered approximately 10 years ago (S), and has, since then, been extended because of its ability to oxidize easily many phosphite, phosphonate, and phosphate esters. The orthophosphate ion resulting from the oxidation can be estimated colorimetrically by any of several methods. In the present work, it has been found that tetramethylphosphonium chloride (TMPC), a compound reputedly very resistant to oxidation ( I ) , can be converted to the orthophosphate ion in dilute aqueous solution, quantitatively and under normal laboratory conditions. The phosphomolybdate method of Dickman and Bray ( 2 ) has been used by the present authors to estimate the orthophosphate concentration. Although tetramethylphosphonium chloride in distilled water can be quantitatively oxidized to the orthophosphate ion, the conversion is incomplete when
Table
ANALYTICAL CHEMISTRY
Recoveries of TMPC Orthophosphate ton
as
(After treatment J\ith ammonium persulfate solutions) XaC1, TMPC, Mg. Mg. Added Recovered
HE
1536
1.
0
0 1030 0 1045 A 0 0056”
0 1 40 2 82 4 68
0 0 0 0
2760 2760 2760 2760
0 0 0 0
2590 1905 1716 0469
Average of seven individual determinations. 0
the solution contains a n excess of chloride ion (Table I). The less than theoretical recovery n’as shown not to be due to the interference of chloride ion in the formation of the molybdenum blue color from orthophosphate ion. From 98 to 99% recovery of orthophosphate ion (0.0738 mg.) was achieved even when as much a s 700 nig. of sodium chloride was added. It is believed that chloride ion interferes in the initial oxidation of tetramethylphosphonium chloride, and the interference may be overcome by performing the oxidation in strongly alkaline solution (Table 11). The initial concentration of sodium hydroxide needed for complete oxidation is dependent upon the quantity of chloride ion in the sample. For example, theoretical conversion to orthophosphate by ammonium persulfate in this procedure is obtained when the initial sample contains 0.1 mg. of tetramethylphosphonium chloride and 6 mg. of sodium chloride, and is 0.471M with respect to sodium hydroxide. On the other hand, theoretical recoveries (within the limits of experimental
error) are obtained even with solutions containing as much as 90 mg. of sodium chloride in a sample of 0.1 mg. of tetramethylphosphonium chloride, if the initial concentration of sodium hydroxide is 1.0~14. Other phosphorus compouiids in which the phosphorus atom is contained in the anion nioiety of the compound, or neutral phosphorus compounds, can be separated from tetramethylphosphonium chloride by using cationic exchange resins ( I ) . The tetramethylphosphonium ion, which is exchanged and affixed to the resin, is then eluted with hydrochloric acid. Results of experiments using a n ion exchanger (Dowex 50, hydrogen state) are shown in Table 111. REAGENTS AND APPARATUS
Ammonium molybdate-hydrochloric acid reagent and stannous chloride reagent were prepared according to the method of Dickman and Bray ( 2 ) . Ammonium persulfate reagent, 0.25M aqueous solution of ammonium persulfate (ACS grade). Sodium sulfite reagent, 0.25M aqueous solution of anhydrous sodium sulfite (A4CSgrade). Potassium phosphate (dibasic, anhydrous), Fisher, certified reagent, dried to constant weight at 11.5’ C. and stored in a desiccator. A stock solution of this 0.01 material (0.05622 mg. per ml. mg. of phosphorus per ml.) was used in preparation of a calibration curve. Sodium hydroxide reagent, i . O N aqueous solution, stored in a paraffinlined bottle or in borosilicate glass. Klett-Summerson photoelectric colorimeter, S o . 60 filter. PROCEDURE
Calibration Curve. Aliquots of the
