e
THE IODOMETRIC TITRATION OF ARSENITE IN ALKALINE SOLUTIONS R. K. McALPINE University of Michigan, Ann Arbor, Michigan
THE
normal oxidation potentials of iodine and NazHPOa+NaH2POa to give a final ratio of 2:1, and arsenic acid are so close toget,her that in an acid solution (c) saturated H3BOa-NaB407. Weight burets were the reaction between arsenious acid and iodine is far used to measure the 0.1 N iodine and arsenious acid short of completion a t the equivalence point. How- solutions, the final end points and blanks being deterewer, the half-cell reaction for the oxidation of arsenious mined by using 0.0025 N arsenious acid and 0.005 N acid involves the liberation of H+. iodine solutions. Du~licatetitrations of annroximatelv 10 milliequivalents of arsenious acid gave an average deviation of 0.0025y0 and a maximum deviation of therefore, as the concentration of H + is decreased the 0.01%. Washburn's comment on these results is, strength of arsenious acid as a reducing agent rises. "It will be noticed that the conclusions of the foregoing There is little change in the strength of iodine as an calculations are completely justified." oxidizing agent with change in pH, thus in neutral or In paying tribute to the accuracy of the analytical alkaline solution the reaction between iodine and arseni- procedure developed by Washburn it should not he ous acid will take place with satisfactory completeness. overlooked that the upper and lower limits of acidity If the solution is too alkaline, however, there may be a proposed by him were derived wholly from calculations significant reaction between the hydroxyl ion and based on several equilibrium constants. No attempt iodine and the equivalence point may be considerably mas made to show that the titrations were equally acoverstepped by the time a starch end point is reached. curate when these limits were approached. The only The article which is universally referred to as estah- condition actually investigated was a pH of approxilishing the limits between which a satisfactory .titration mately 7. As a matter of fact, the insistence that in may be made i~ one by E. W. Washburn' on "The the hicarbonate-carbonic acid system the solution theory and practice of the iodometric determination should he kept saturated with C02 and the explanation of arsenious acid." For the acid range he calculated given for an experiment which he described in this conthat O.OOIYo of the arsenic would remain unoxidized at a nection would indicate that he placed very little relipH of 4 a t equilibrium under the following experimental ance on the alkaline limit, pH-9. conditions as derived from his own work: 13-=2X It is intexesting to note that the constant used for 10-I M, I-=0.086, and H,AsO4+H,AsO4-= 0.02 M. the hydrolysis of iodine to form hypoiodous acid and In the alkaline range he considered two possible reac- iodide ion was derived from an article by Sammet2 in tions of iodine with hydroxyl ion: which the statement was made that the value for this Washburn used 31~-+ 30H- = IOs- + 81- + 3H+ (1) constant is about lo-%to 10-'O. for his calculation. In 1910 this constant was experi1,- + OH- = H I 0 Lt 21(2) mentally determined by Bray and his value corrected Equilibrium constants were calcul~tedfor these two the next year by Bray and C ~ n n o l l y the , ~ final figure reactions by combining hydrolysis constants for iodine being 3X10-'3. When t,his value is combined with the with the water constant and with the ionization con- others used by Washburn and his calculations then restant for k-, then it was shown that, under equilibrium peated, it is found that the concentration of H + would conditions and the same experimental data as before, need t o drop to 4X10213-before the amount of H I 0 0.001% excess of iodine would be used to form 10%- present would equal 0.001% of the total amount of IBa t a pH of 10, and a similar excess would be converted used. to HI0 at a pH of 9. Washburn thus concluded that However, this calculation fails to take into account in his experimental work if the pH could be kept be- the fact that HIO, although a very weak acid, would distween the limits 4 and 9 the errors inherent in the reac- sociate to a very large extent under these conditions. tions would not exceed 0.001%. I t was assumed, Latimer ("Oxidat,ion Potentials," Prentice Hall, 1938, therefore, that the condition of maximum accuracy p. 57) gives 1x10-" as an approximate value for the would be at the geometric mean of these two points or a ionization constant of HIO, based on the experimental pH of 6.5. studies of Fiirth.4 Fiirth's own report gave values To set up a buffered solution that would approximate SAMMET, Z . phy8. Chem., 53, 640 (1905). a pH of 7 a t the end point, three different systems were 'BRAY,U '. C., AND E. L. CONNOLLY, J . Am. Chem. Soc., 33, tried. (a) 0.12 M NaHCO? saturated with CO.. (b)
..
' WASHBURN,E. W., J. Am. Chem. Sac., 30, 31 (1908).
1485 (1911). F ~ ~ R TA,, R , Z . Eleklroehem., 28, 57 (1922).
3152
JULY, 1949
363
ranging from 1.8 to 4X10-" and stated that 2 to 3X lo-" could be accepted with reasonable confidence. Experiments in this laboratory on the sensitivity of the iodme starch reaction in alkalme solutions favor the figure 3x10-". When this is used t o obtain the distribution of hypoiodite radical between the two forms H I 0 and IO-, it is found that with Washburn's data the formation of HIO+IO- would introduce an error of 0.001'% when the H + = 1.4X10-12. This figure is much below the olie calculated by Washburn for the interference due to formation of TO3(under equilibrium conditions at pH=lO the error would be 0.001%). Thus it seems probable that the titration of arsenious acid with iodine may be carried out satisfactorily in a more alkaline solution than is commonly believed and the first error t o be encountered will be due to the formation of TO3- rather than HIO+IO-. Of the modern texts only Kolthoff and Furman ("Volumetric Analysis," John Wiley & Sons, 1929) seem to recognize this situation. Based on experimental studies by Kolthoff in 1919 the statement is made, "The pH may be between 5 and 11 in the titration of arsenious acid with iodine if the solution contains some iodide. The case is different in the reverse titration of iodine with arsenious acid. UD to ~ H = 9 . 2a t the end of the titration the theoreticai quantity of arsenious acid is transposed, a t a higher pH, however, some iodine is oxidized to iodate and hence too little standard solution is used." In spite of this statement they give the procedure for the bicarbonate-carbonic, acid buffering system. Since the textbooks in analytical chemistry have generally overlooked the work of Kolthoff and have gone back to Washburn for experimental details and theoretical discussion. it seemed worth while to carrv out a careful series of titrations of arsenious acid in more alkaline range than usual. A standard solution of sodium arseuite was prepared by dissolving 4.1485 g. of A&O3in 120 ml. of 3 N NaOP and diluting to one liter. After thorough mixing, 50.00-ml. samples were placed in 500-ml. Erlenmeyer flasks, Created with 150 ml. of distilled water and 5 ml. of a freshly prepared starch solution, then various buffering agents were added and the solutions carefully titrated with an approximately 0.1 N solution of K13. After the titrations were completed, the pH's of the final solutions
a
were determined with glass electrodes by a pH meter. The conditions and results are summarized in Table 1. It will be observed that only in the first case was the common practice followed of using a bicarbonate-carbonic acid buffer. In the others the sodium hydroxide in.the original sample was converted t o sodium carbonate by the use of sodium bicarbonate, or to ammonium hydroxide by the action of ammonium chloride, and these conditions then modified by adding more sodium bicarbonate or sodium carbonate in the one series, or by using more ammonium chloride or ammonium hydroxide in the other. The results check as closely as one can work with a 50-ml. buret, the maximum variation being only 0.05 ml. Following these titrations it seemed worth while t o locate the point a t which alkaline errors would put in an appearance. A fresh solution of sodium arsenite was prepared and aliquot portions of this mere diluted and treated with NaHC03+N&C03.H20 or with N&HP0,.7H10 in amounts that would approximate final pH's of 11,11.5, and 12. The data and results are given in Table 2. TABLE 2 Titration of Sodium Arsenite i n More Alkaline Solutions
Buffwing agents added 5 g. NaHCO. 1 . 5 g. NaHC03 +.3.5 g. NsnCOI.HzO 1 . 5 g. NsHC03 13.6 g. NazCOs.HzO 1 . 5 g. NaHC03 45.6 g. NsCOs .H1O 29.3 g. NadP01.7HzO 11.1 g. N&HPO,. 7H10 5 . 3 g. NsnHPOr7Hn0
+ +
9.6
Ml. K13 solution used 40.94
11 .O
40.96
11.5
41.70
12.0 11.0 11.5 12.0
42.53 40.95 41.07 41.23
Applazide pH
It will be observed from experiments 1, 2, and 5 that the titration as carried out appears t o be accurate up to a pH of approximately 11. However, as the solution becomes more alkaline, significant amounts of the iodine are used up by a secondary reaction. The results with the sodium phosphate buffer are not as far off as in the sodium carbonate cases, possibly due in part to an actual difference in the pH's of the solutions, but also due in part to the fact that the solutions were swirled more frequently in the last two cases. To show how much difference this latter condition might make, two more solutions were prepared as duplicates of No. 3, and TABLE 1 the titrations repeated. In the first, the solution was Titrations of Arsenious Acid in Alkaline Solutions swirled almost constantly while the iodine was added in M1. KI, rapid drops. I n the second the iodine was run in a bit Buffering agents added Final pH solution used more rapidly but the solution was swirled only occasion5 ml. 5 N HCI + 5 g. N ~ H C O I 6.9 42.63 ally until close t o the end point. The volumes of 8.7 42.67 11.6 g. NaHCO. iodine used were 41.25 ml. and 43.25 ml., respectively. 1 . 4 g. NH4CI 9.3 42.66 9.6 42.65 2 . 6 g. NaHCO1 With rapid swirling to mix the iodine solution in very 0.96 g. NH,Cl + 7.5 ml. 5 N rapidly the secondary reaction took place to a conNHIOH 10.0 42.65 -siderably less extent than when the solution was 10.2 42.67 1 . 6 g. NaHCOa 4 g. 1 . 6 g. NaHCOs swirled less frequently. Na2COa.H%0 10.6 42.68 In attempting to account for these experimental facts
+
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it becomes necessary to consider in some detail the quantitative experiment. For this purpose a 50.00reactions which might occur when an iodine solution is ml. sample of the sodium arsenite solution used in added to an alkaline solution containmg an arsenite. Table 2 was treated with 1 g. KIOs and 5 g. NaHCOa, The following series of reactions, and the accompanying then diluted to 250 ml., and titrated with the iodme statements, are consistent with the facts. solution. The volume required, 40.95 ml., was identical with that used in titrations (I), (Z), and (5). From these considerations it. appears that when the iodme reagent is added to an alkaline arsenite solution there are two reactions competing with each other for the iodine. One involves the oxidation of the arsenite to acid arsenate ion, the other the formation of 103-and I- through 10- as an intermediate. For this purpose it is unimportant whether the oxidation of the arsenic is (a) Reactions (I), (3), and (4) are readily reversible done by Ia-, Is, HIO, or 10-. All are capable of doand:~each equilibrium very rapidly. ing it, although in a dilute aqueous solution between a (b) Reaction (2) in reversible and in alkalme solution pH of 7 and 10 it may be noted that the oxidizing iodine reaches equilibrium rapidly and goes practically to will be there chiefly in the forms 1%and IS-. The main completion. fact is that up to a pH of 10 the rate of formation of (c) Reaction (5) is not reversible in alkaline solution, 103- is so low that essentially all the iodine is used to it takes place rather slowly and is negligible a t a pH oxidize the arsenic and only negligible amounts are less than 11 where the concentration of H I 0 and its dis- changed to 103-. Even a t a pH of 11 when the titrasociation are so low that only extremely small amounts tion is carried out a t a moderate rate with reasonable of;IO- are present. swirling or stirring the amount of 103- formed is ( d ) In alkaline solutions 103-will not oxidize arseni- insignificant. However, a t a pH of 11.5 a considerable ous acid. fraction of the i o d i e is promptly changed to 10- and An experimental justification for statement (a) the formation of some 103-can then be recognized by may be demonstrated as follows: To 250-300 ml. of the excessive consumption of the reagent. distilled water add 1 g. KI, 5 ml. of starch solution, and Since the conversion of 10- to 103-and I- can now 1 drop of 0.1 N iodine solution. A deep blqe color take place a t a moderate rate even though it is much develops. Add 5 ml. of 3 N NaOH and swirl to mix. slower than thezoxidation of the arsenite, it is evident The color promptly disappears due to conversion of Is that the amount of 108-formed will vary considerably to 10- by reactions (3) and (4). Now add 6 to 7 ml. of with the conditions of the titration. If the solution is 3 N NHICl and swirl. Tlie blue color reappears with swirled or stirred vigorously during the titration so that equal promptness, showing the formation of Izagain by the i o d i e is distributed very..quickly throughout the reversal of reactions (3) and (4). solution, theq the amountswed in each of the reactions Statement (b) is justified by the fact that the reaction will be a fair measure of their relative rates. However, lends itself effectively to the use of As103 as a primary if local excesses of iodine are permitted to accumulate, standard in determining the concentration of iodine the arsenite in direct contact with these areas may be solutions. In slightly acid solutions equilibrium is completely oxidized, but the free 10- still present may reached rather slowly, but in alFaline solutions there is continue to change to 1O3- and.1;. Under these condino need for delay in approachmg the end point. tions, a t the end of the titration'considerably more of The statement concerning the llbnreversibility of the iodine will have been converted to 103-. This is reaction (5) in alkaline solution is justified as follows: the explanation of the experiment in which only 0.3-ml. a mixture of 1 g. KI, 1 g. KI03, and 5 g. NaHC08 is dis- excess of the iodine was used under the conditions of solved in 300 ml. of distilled water and starch solution rapid swirling, while with occasional swirling the excess added. The solution remains colorless. Even the was 3.3 ml. With moderate swirling, as in the ordinary addition of several grams of NH&1 fails to produce a titration, 1.75-ml. excess of t h e iodine solution was reblue color. The further fact that this reaction takes quired to reach the end point. place slowly is shown by reference to the previous exBy using the proper ionization constants it is possible periment in which the blue color was destroyed by addition of NaOH and then returned when NH4C1was TABLE 3 added. If the 10- reacted rapidly to form IOa- and IDistribution of Iodine Added between 10- and 1,- + Iz the solution would have remained colorless after treata t Higher pH's and Different Iment with NH&1. As a matter of fact, if the treatment with NaOH and then NH&l is repeated several times in succession it will be noted that the depth of the blue color falls off appreciably, indicating that a recognizable part of the iodine is being converted to 1 0 2 - while the. solution is colorless. The fourth statement requires justification by a
JULY, 1949
to calculate the distribution of the iodine reagent added among the forms Is-, 10-, and 12. However, since iodide ion is a product in the dissociation of Isand in the hydrolysis of 12,i t is evident that the distribution of the iodine will depend also on the concentration of I- in the solution. For these calculations two values were tried, 0.03 M I- and 0.1 M I-. The results are given in Table 3. At a pH of 11, with I-=0.03 M , the reagent added will be largely in the form of Is- and 4, with only a small fraction present as 10-. As the pH rises there is a marked shift to more 10- and less 1,-+I,, but as the concentration of I- rises the reverse effect is noted. From this it would appear that by increasing the concentration of I- the error due to the formation of 103should be reduced. To check this point experimentally two final samples of sodium arsenite lvere titrated under the same conditions as in Table 2, cases 3 and 4, except that 3.5 g. of K I were added to make the solution approximately 0.1 M with I- at the end of the titration. The results, 41.03 ml. and 41.72 ml., when compared with the amounts of iodine used without the excess K I (41.70 and 42.53, respectively) offersignificant confirmation of the point of view here expressed. There still remains one element of mystery to be solved. That is the problem of explaining the experiment described by Washburn in which the titration was carried out in a beaker, then when the end point was reached the solution was stirred for thirty seconds and the color disappeared. On adding a few drops of hydrochloric acid carbon dioxide was evolved and the color immediately returned. Washburn's own explanation of this, which seems to have been universally accepted, was that as the solution was stirred carbon dioxide was lost, the solution became more alkaline, and the iodine reacted with the OH- to form H I 0 and I-. On adding the HC1 this reaction was reversed. Others have reported seeing similar effects so the experimental facts may be accepted. Before attempting to offer an rtlternative explanation of the facts it may be observed that it is only when the amount of free iodme is extremely sdall that the color will disappear as quickly as was indicated. In the titrations described in this paper 500-ml. Erlenmeyer flasks were used for convenience in mixing and the end point was reached by adding fractional drops of approximately 0.1 N reagent. Under these conditions the concentration of free iodme a t the end point would be 3 to 5X10-6 M. In many cases the flasks were covered with small beakers and permitted to stand for a number of hours to test the stability of the color. In other cases the solutions were poured into beakers and stirred to see how readily the color might disappear under these conditions. In the open beaker the color did disappear more rapidly than in the covered flask although even in that case, with frequent stirring and standing, the color would nonnally persist for an hour or more except when dealing with the more alkaline solutions in the pH range of 11 or higher. This difference in times as compared with Washburn's figure may be due to the differencein
365
concentrations of free iodine at the end point. He claimed to have obtained end points with as little as 2x10-' M free iodine present. Obviously it would take less time to use up that amount of iodine by a slow reaction than would be needed for the larger amounts of iodine involved in the end points here reported. Washburn's titrations were carried out at a pH of approximately 7, i. e., NaHCOs was 0.12 M and the solution was saturated with Cop. Under these conditions it is extremely doubtful if the stirring would have reduced the concentration of Cot sufficientlyto exceed a pH of 7.5, while to reach a pH of 9 it would be necessary to remove all the free COz as well as part of that formed by hydrolysis of the NaHC03. I t is obvious, then, that the loss in color cannot be explained in terms of simple conversion of the iodine to H I 0 and I-. Three other possible ways in which the iodine might disappear may be considered. One involves the simple mechanical process of volatilization of iodine from the solution. A second deals with the reaction by which 10- changes to 1 0 2 - and I- by auto-oxidation and reduction. The third considers the possibility of absorption of oxygen from the air with oxidation of the iodine to 1 0 3 - in the alkaline solution. Of these, the second can be ruled out immediately, because it would involve reagents already present in the solution and should take place as readily in the covered flask as in the open beaker. It should even take place in a stoppered bottle, yet Washburn reported that under those conditions the color persisted for as much as two weeks. So far as the other two effects are concerned, the difference in time required for the solutions to become colorless in the covered flask and in the open beaker could be accounted for equally well by either of them. To distinguish between the.two'it is necessary to apply tests for 103-on the colorless solutions. This is easily done by merely acidifying the solution. If there is any 1 0 8present it will react with the I- to form free iodine and the blue color will again appear. When this was tried positive tests were obtained, showing that in the disappearance of color some oxidation to 103-had taken place. This, of course, does not completely exclude any loss of iodine by direct volatilization, but it does offer a basis for explaining the restoration of color when the solution is treated with hydrochloric acid. When hydrochloric acid is added to a slightly alkaline solution containing I-, I03; and starch solution, local areas are set up in which the solution is actually acid and iodine is liberated even though the amount of acid added is less than sufficient to acidify the whole solution. A simple demonstration of this may be carried out by adding about 1 g. of KI, 0.1 g. of KIO3,and 2 to 3 g. of Na*COs.Hz0 to 200 to 250 ml. of distilled water containing a little starch solution. Aft,er stirring to dissolve and mix thoroughly let 2 or 3 drops of 5N HC1 fall into the solution. Locally an effervescencecan be seen and on stirring a blue color is obtained, yet if the solution is tested it will be found to be quite alkaline. Actually the HC1 added is less than sufficient to change all the NarCOs to HC03-.
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JOURNAL OF CHEMICAL EDUCATION
Prohably most chemists have encountered these "local action" effects in the process of mixing solutions even though the theoretical discussions frequently seem to assume that the mixing is instantaneous. Another local action effect which is particularly pertinent to this paper and which is not well known, may be seen in the following experiment. Place 5 g. of NazC03.H20, 2 g. of KI, and 250 ml. of distilled water in each of two flasks. To the first add 5 ml. of starch solution and a drop of 0.1 N Iz solution. Swirl and note the formation of a blue color. To the second add the iodine solution first and then the starch solution. In this second case the reaction of the iodine with the OH- formed by hydrolysis of the Na2COs reduces the concentration of free iodine to the point where it will not give the blue color when the starch is added. In the first case, however, with the starch already present when the iodine is added, there are local areas in 15-hich the concentration of free iodine is sufficientlyhigh to cause the formation of the blue color. I n continuation of this experiment, it is instructive to let both of the flasks stand and observe them from time to time. Even after an hour the first solution will be colored and the second colorless. Evidently when the color has once been produced it will persist even when the concentration of free iodine has been lowered to a point where it will not cause the color to form. Eventually the blue solution will become colorless, due finally to oxidation of the iodine to 103-. SUMMARY AND CONCLUSIONS
a t p H = l l were identical with those obtained in less alkaline solutions, hut a t a pH of 11.5 a significant excess of iodine was used, and at a pH of 12 still more was required to obtain an end point. (4) In solutions sufficiently alkaline, and with low enough concentrations of I- to permit the formation of moderate amounts of IO-, a secondary reaction takes place at only a moderate rate involving an auto-oxidation reduction reaction of IO- t o give 1 0 3 I-. (5) The slow decolorization of a titrated solution on standing in an open beaker involves oxidation of the iodine to 1 0 3 - , probably combined with some volatilization of the iodine. (6) On the basis of these studies the following procedures were proposed and tried for the standardization of 0.1 N iodine solution against As203 as a primary standard. Weigh out several samples of As203, keeping the weight close to 0.23 g. to avoid using more than 50 ml. of the iodine solution. (a) Use 3 g. of NazCO3.Hz0 and 20 ml. of distilled l ~ a t e rto dissolve the As203. Heat to 60-70°C. and swirl gently until clear, rinsing down the walls of the flask to he sure all As203 is dissolved. Dilute the solution to 100 ml., cool, add 1 g. of NaHC03, and titrate with the iodine solution. By this procedure the pH will be approximately 10.4 a t the start and 9.5 at the end. (b) Use 10 ml. of 3 N NaOH (in place of the NazC03)to dissolve the A~203,keeping the solution cold. Swirl eentlv until dissolved and rinse d o m the walls of .. the flask to-he sure all the AszO3 is dissolved. This solution was then diluted to 100 ml. and two different methods of buffering were tried. In some cases 3 g. of NaHC03 were added, in others 2 g. of ?JH,Cl. Then starch solutiou was added and t%e titrat,ion carried out. The NaHCOaSmethod gave a pH of approximately 11 at the start and 10.5 a t the end, while t,he NH4C1gave approximately 9.9 at the start and 9.4 at the end. Duplicate titrations on each of two different batches of As203by all three of these methods gave identical results within t:he limits of weighing and reading the buret.
+
~
(1) It has been shown by use of Washburn's data and the proper equilibrium constants that the error due t o reaction of Ia or 13-with OH- t o form I- and IO-+HI0 would he only 0.0019/, a t a pH of 11.85. (2) Direct titrations of sodium arsenite solutions in the presence of various buffering agents gave identical results, within the accuracy of using a 50-ml. buret, when the final pH as measured by a pH meter ranged from 6.9 to 10.6. (3) On titrating sodium arsenite in a more alkaline range (pH calculated= 11.0, 11.5; and 12.0) the results
Oxidation
Max Epstein of the New Utrecht High School, Brooklyn, New York, some time ago sent us a suggestion for a demonstration. Melt some potassium chlorate in a large test tube held in a clamp over a burner. Drop in a small roll of steel wool-but keep out of the way! Complete oxidation. It may be added that a long wooden splint is a good substitute for the steel wool. It is always surprising that cold wood can actually be ignited in this way.
