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Reliability in standardization of sodium thiosulfate with potassium dichromate
Microchemical Journal 123 (2015) 9–14
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Reliability in standardization of sodium thiosulfate with potassium dichromate Toshiaki Asakai ⁎, Akiharu Hioki National Metrology Institute of Japan, National Institute of Advanced Industrial Science Technology, 1-1-1 Umezono, Tsukuba, Ibaraki 305-8563, Japan
a r t i c l e
i n f o
Article history: Received 23 March 2015 Received in revised form 12 May 2015 Accepted 12 May 2015 Available online 16 May 2015 Keywords: Coulometric titration Oxidation-reduction titration Certified reference material Potassium dichromate Sodium thiosulfate Potassium iodate
a b s t r a c t Sodium thiosulfate is often standardized with potassium dichromate. In the standardization, iodine (triiodide) liberated by potassium dichromate in an acidic potassium iodide solution is titrated with a sodium thiosulfate solution. The iodine liberation process significantly affects the titration results. In the present study, the accuracy of the liberation process was examined by assaying potassium dichromate through two different paths: assaying directly by coulometric titration with electrogenerated Fe(II), and assaying by gravimetric titration through the iodine liberation reaction with a sodium thiosulfate solution of which concentration was standardized by coulometric titration with electrogenerated iodine. The accuracy of the standardization of a sodium thiosulfate solution by potassium dichromate was discussed from the apparent assays of potassium dichromate under different measurement conditions. © 2015 Elsevier B.V. All rights reserved.
1. Introduction Sodium thiosulfate is a useful reducing agent for a titrant solution in volumetric analyses such as iodometry and the Winkler method to measure dissolved oxygen [1,2]. Oceanography community requires the accuracy and the reproducibility of the measurements of dissolved oxygen; the accurate standardization procedure of sodium thiosulfate is needed. The agent is often standardized with potassium dichromate and, on the other hand, is usually standardized with potassium iodate in Japan [3–6]. Iodine (triiodide ions) liberated by potassium dichromate or potassium iodate in an acidic potassium iodide solution is titrated with a sodium thiosulfate solution to standardize the thiosulfate [reactions (1), (2) and (3)]. K2 Cr2 O7 þ 6KI þ 7H2 SO4 →Cr2 ðSO4 Þ3 þ 4K2 SO4 þ 7H2 O þ 3I2
ð1Þ
KIO3 þ 5KI þ 3H2 SO4 →3K2 SO4 þ 3H2 O þ 3I2
ð2Þ
2Na2 S2 O3 þ I2 →2NaI þ Na2 S4 O6
ð3Þ
⁎ Corresponding author. Tel.: +81 29 861 6881; fax: +81 29 861 6890. E-mail address: [email protected] (T. Asakai).
http://dx.doi.org/10.1016/j.microc.2015.05.012 0026-265X/© 2015 Elsevier B.V. All rights reserved.
The iodine liberation process is significantly affected by the amounts of acid and potassium iodide added, the waiting time for the liberation, and light; therefore, the process plays a key role for the accuracy of standardization of sodium thiosulfate. The authors reported an accurate standardization procedure for a sodium thiosulfate solution with potassium iodate [7]. It has been pointed out that there are some biases in the standardization of sodium thiosulfate with potassium dichromate [3,8–12]. Some researchers reported that excess consumption of thiosulfate occurred in an acidic medium due to air oxidation of iodide ions [8,9]. Hahn described that the main reason of the excess consumption of thiosulfate in a weak acidic medium was not air oxidation of iodide, but a complex between Cr(III) and thiosulfate, which reacted with iodine slowly [3,12]. Kolthoff and Belcher reviewed the studies on the titration procedure [3] and mentioned the appropriate conditions to ensure complete and stoichiometric reaction with dichromate: the hydrogen ions concentration 0.2 mol L−1 to 0.4 mol L−1, 2% of potassium iodide, and standing for 10 min to 15 min before starting titration with thiosulfate. The impact of the biases has remained obscure because the accuracy of the standardization of thiosulfate has been discussed by using relative methods for conventional titration. Thiosulfate is usually standardized with sublimed iodine whose purity is hypothetically regarded as 100%. In the present study, the accuracy of the standardization was absolutely evaluated by coulometric titration (Fig. 1). Coulometric titration absolutely yields an accurate oxidimetric–reductometric assay based on Faraday's Law. The accuracy of the iodine liberation process was examined by comparing two assays of potassium dichromate using different
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1 μg, and potassium iodide and solutions were weighed with resolution of 10 μg. 3. Results and discussion 3.1. Determination of sodium thiosulfate by coulometric titration and uncertainty estimation
Fig. 1. Experimental design for the assay of potassium dichromate to investigate the iodine liberation reaction.
titrations based on coulometric titration. Firstly, potassium dichromate was directly assayed by coulometric titration with electrogenerated Fe(II) [13]. Secondly, potassium dichromate was assayed by gravimetric titration through the iodine liberation reaction with a sodium thiosulfate solution, of which concentration had been standardized by coulometric titration with electrogenerated iodine. The accuracy of the standardization of a sodium thiosulfate solution by potassium dichromate was discussed from the apparent assays of potassium dichromate under different measurement conditions.
2. Material and methods Constant-voltage biamperometry (a dead stop method) was utilized to measure the amount of liberated iodine under several conditions as a preliminary experiment. A DC source and a digital multimeter with a dual platinum-chip electrode (each platinum chip: 0.5 mm ϕ × 0.5 mm long) were used for constant-voltage biamperometry. Constant voltage was applied between the platinum pair, and the current was monitored. The current had good proportionality to the amount of liberated iodine [7]. A dual platinum-plate electrode (each 5 mm × 5 mm) was used in gravimetric titration to increase sensitivity as described previously [7, 13–15]. The experimental set-ups and the procedures of coulometric titrations for potassium dichromate and a sodium thiosulfate solution were described previously [7,13–15]. As described in the previous report [13], high purity potassium dichromate as a certified reference material (CRM) was established by coulometric titration with electrogenerated Fe(II). The certified value (mass fraction) of the CRM, NMIJ CRM 3002-a, is 99.974% ± 0.011%, where the value following ± indicates the expanded uncertainty with the coverage factor 2. Gravimetric titration was carried out to assay potassium dichromate with a sodium thiosulfate solution through the iodine liberation reaction in the following procedure: approximately 0.2 g of potassium dichromate were placed in a 200 mL tall beaker, it was dissolved in 100 mL of water, potassium iodide and 9 mol L−1 sulfuric acid were added, and the solution was titrated with a sodium thiosulfate solution. The concentration of the sodium thiosulfate solution was standardized in advance by coulometric titration with electrogenerated iodine [7]. A plastic syringe with a PEEK needle was used to weigh the sodium thiosulfate solution. Potassium dichromate was weighed with resolution of
The concentration of a sodium thiosulfate solution (3 g to 8 g) standardized by coulometric titration was, for example, 208.653 mmol kg−1 (RSD 0.007%, n = 5) (‘RSD’ means ‘relative standard deviation’, and ‘n’ is the number of measurements under a repeating condition). The uncertainty sources for the concentration of the solution (ca. 3 g of ca. 200 mmol kg−1) were as follows (the number in each parenthesis indicates the relative standard uncertainty): repeatability (0.0032%, RSD of the mean, n = 5), Faraday constant (0.000 002 2%), a standard resistor (0.000 65%, calibration and the influence of temperature), a frequency counter (0.000 05%), a voltmeter (0.000 15%) and weighing (0.0019%, sensitivity and linearity of the balance). Finally, the combined standard uncertainty was 0.0038% relative and the expanded uncertainty with a coverage factor 2 was 0.0076% relative. The coulometric titration allowed the authors to discuss the accuracy of the iodine liberation reaction by potassium dichromate in the order of magnitude of 0.01%. A sodium thiosulfate solution is stable enough to be required in the present study [16,17]. 3.2. Influence of the amount of potassium iodide added Approximately 0.2 g of potassium dichromate were dissolved in 100 mL of non-deaerated water followed by adding 1 g to 6 g of potassium iodide and 1 mL of 9 mol L−1 sulfuric acid in turn; after the acid addition, iodine liberation started. The iodine liberation process monitored by constant-voltage biamperometry using a dual platinum-chip electrode (applied constant-voltage 500 mV) is shown in Fig. 2; the indicator current corresponds to the amount of iodine liberation. Under the condition using 1 g to 2 g of potassium iodide and 1 mL of 9 mol L−1 sulfuric acid, lacking of iodide ions in the solution led to the deposition of solid iodine, and the indicator current drastically decreased in unforeseeable timing after the initial increase. Under the conditions using 3 g to 6 g of potassium iodide and the same amount of sulfuric acid, the
Fig. 2. The iodine liberation process monitored by constant-voltage biamperometry for 0.2 g of potassium dichromate using a dual platinum-chip electrode (applied constantvoltage 500 mV), stirring at 900 revolutions per minute (rpm). The sample was dissolved in 100 mL of non-deaerated water, and 1 g to 6 g of potassium iodide and 1 mL of 9 mol L−1 sulfuric acid were added.
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decrease of the indicator current did not occur; these conditions seemed to be enough to completely generate dissolved iodine. In the previous study [7], potassium iodate (0.15 g), whose redox equivalent is almost the same as that of potassium dichromate (0.2 g), liberated iodine completely in a couple of seconds after acid addition under the condition using 3 g to 5 g of potassium iodide and 1 mL of 9 mol L−1 sulfuric acid; consequently, no waiting time for the iodine liberation was necessary. Potassium dichromate conversely required more than 5 min for the complete iodine liberation according to Fig. 2. Gravimetric titration was performed to assay potassium dichromate in a darkroom using a sodium thiosulfate standardized by coulometric titration. The dependency of the apparent potassium dichromate assay on the amount of potassium iodide added is shown in Fig. 3. Approximately 0.2 g of potassium dichromate in a beaker was dissolved in 100 mL of deaerated or non-deaerated water followed by adding 1 g to 5 g of potassium iodide and 1 mL of 9 mol L−1 sulfuric acid in turn; liberated iodine was titrated at 10 min after the acid addition with a sodium thiosulfate solution, of which concentration was standardized by coulometric titration. The beaker was stoppered with Parafilm during the liberation to avoid the volatilization of iodine. Fig. 3 includes the certified value of potassium dichromate determined by coulometric titration with electrogenerated Fe(II) [13] and also the dependency of the apparent potassium iodate assay on the amount of potassium iodide added [7]. Liberated iodine by potassium iodate was titrated immediately after iodine liberation starting. The similarity of the assay values of potassium dichromate and potassium iodate was a coincidence because of high purity materials. The apparent assay of potassium dichromate with 1 g of potassium iodide and 1 mL of 9 mol L−1 sulfuric acid using deaerated or non-deaerated water was more than 101%. The apparent assays of potassium dichromate had a minimum for both the results using deaerated water and those using non-deaerated water; the minimum value with 2 g of potassium iodide was close to the certified value. The titration with less than 2 g of potassium iodide was not able to be accurately completed since the indicator current drifted upward during the titration. In this iodometry, the indicator current decreased with progress of titration, that is, the liberated iodine was reduced with a sodium thiosulfate solution. The end point of this type of reaction was at 0 A of the indicator current where the iodine just disappeared. The apparent assay of potassium dichromate corresponded to the mass of the sodium thiosulfate solution consumed until the end point. The upward drift of the indicator current indicated
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the production of iodine or oxidants with progress of time. The drift made it difficult to draw the titration curve and to determine the end point accurately. The apparent assays were also significantly affected by the speed of titration. Less amount of potassium iodide and slow titration led to more evaporation of iodine and more influence of light; accurate titration results were not expected. The iodine liberations were obviously not completed under the condition using 1 g of potassium iodide and 1 mL of 9 mol L−1 sulfuric acid for either 0.2 g potassium dichromate or 0.15 g of potassium iodate [7] as shown in the preliminary examination by constant-voltage biamperometry (Fig. 2). Consequently, unreacted potassium dichromate and potassium iodate remained in the solutions before starting a titration with thiosulfate. The authors reported that unexpected good results for potassium iodate under incomplete iodine liberation conditions were caused by resupplied iodide ions with progress of the reaction [7]. Such good results did not happen for potassium dichromate as shown in Fig. 3. Judging from Fig. 3, it was inferred that excess sodium thiosulfate was consumed in the presence of the unreacted potassium dichromate. There could be the following three possibilities: (1) sodium thiosulfate was insufficiently oxidized, (2) it was decomposed under the incomplete iodine liberation in an acidic medium, and (3) it interacted with potassium dichromate or generated Cr(III) and lost the reducing capability. If the remaining dichromate caused insufficient oxidation of thiosulfate, it would result in excessive consumption of thiosulfate. Rao and Sarma reported the direct standardization of thiosulfate by potassium dichromate with copper(II) sulfate as a catalyst [18] and mentioned that dichromate directly oxidized thiosulfate to tetrathionate state as usual. Therefore, excessive consumption of thiosulfate was not caused by its insufficient reduction. Though the possibility of its decomposition before reacting with remaining dichromate or iodine cannot be excluded, as mentioned above Hahn reported a phenomenon relevant to the third possibility [12]. It was the generation of the complex between thiosulfate and Cr(III) in the presence of dichromate; consequently, excess thiosulfate was apparently consumed. According to his report, Cr(III) once formed does not react with thiosulfate, but the complex between Cr(III) and thiosulfate is formed when the reduction of Cr(VI) takes place in the presence of thiosulfate. The reason of slight increase in the apparent assays of potassium dichromate under the condition using more than 3 g of potassium iodide could be the significant oxidation of potassium iodide during iodine liberation (10 min) in an acidic medium depending on the amount of potassium iodide added. Significant air-oxidation of iodide ions in an acidic medium without oxidants such as potassium iodate was not observed [7,15]; however, Cr(III) could accelerate the oxidation reaction of them (details are described in section Influence of background oxidation of iodide ions and vaporization of liberated iodine). 3.3. Influence of the amount of acid added
Fig. 3. Dependencies of the apparent potassium dichromate and potassium iodate assays by gravimetric titration with thiosulfate on the amount of potassium iodide added. Iodine liberation times after the acid addition were 10 min for potassium dichromate and 0 min for potassium iodate. ○: 0.2 g of potassium dichromate dissolved in 100 mL of nondeaerated water; 1 mL of 9 mol L−1 sulfuric acid. □: 0.2 g of potassium dichromate dissolved in 100 mL of deaerated water; 1 mL of 9 mol L−1 sulfuric acid. ×: 0.15 g of potassium iodate (NMIJ CRM 3006-a) dissolved in 100 mL of non-deaerated water; 1 mL of 9 mol L−1 sulfuric acid [7]. −: Certified value (99.974%) of potassium dichromate by coulometric titration with electrogenerated Fe(II) (not related to the abscissa axis) [13].
Approximately 0.2 g of potassium dichromate was dissolved in 100 mL of non-deaerated water followed by adding 3 g of potassium iodide and 0.5 mL to 5 mL of 9 mol L−1 sulfuric acid in turn; after acid addition iodine liberation started. The iodine liberation process was monitored in the same manner as done for the examination of influence of the amount of potassium iodide added (vide supra); the results are shown in Fig. 4. Under the condition using 3 g of potassium iodide and 0.5 mL of 9 mol L−1 sulfuric acid, the iodine liberation was not completed. Under the condition using 3 g of potassium iodide and 1 mL to 5 mL of 9 mol L−1 sulfuric acid, the iodine liberation was apparently completed. It seems that more acids accelerate the iodine liberation reaction. Gravimetric titration with thiosulfate was performed to evaluate the dependency of apparent assay of potassium dichromate on the amount of sulfuric acid used. A dual platinum-plate electrode was used for the end-point determination. Fig. 5 shows the results of the gravimetric titration together with the apparent assays of potassium iodate [7]. Approximately 0.2 g of potassium dichromate were dissolved in 100 mL of
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acid is required in order to accurately carry out the standardization of sodium thiosulfate with potassium dichromate. In addition, deaeration is desirable to avoid the excess oxidation of iodide by dissolved oxygen. 3.4. Influence of the iodine liberation time
Fig. 4. The iodine liberation process monitored by constant-voltage biamperometry for 0.2 g of potassium dichromate using a dual platinum-chip electrode (applied constantvoltage 500 mV), stirring at 900 rpm. The sample was dissolved in 100 mL of nondeaerated water, and 3 g of potassium iodide and 0.5 mL to 5 mL of 9 mol L−1 sulfuric acid were added.
deaerated or non-deaerated water followed by adding 3 g of potassium iodide and 0.5 mL to 4 mL of 9 mol L−1 sulfuric acid in turn; liberated iodine was titrated with a thiosulfate solution at 10 min after acid addition. The results by gravimetric titration using 100 mL of deaerated water and 1 mL or more of 9 mol L−1 sulfuric acid were consistent with those by coulometric titration with electrogenerated Fe(II); on the other hand, those by gravimetric titration using non-deaerated water increased with the increase in the amount of acid used. This fact suggested that dissolved oxygen oxidized iodide ions in an acidic media described in detail below [7]. The assay of potassium dichromate dissolved in 100 mL of deaerated or non-deaerated water, under the condition using 3 g of potassium iodide and 0.5 mL of 9 mol L−1 sulfuric acid, was more than 106%. The indicator current drifted upward during the titration under the conditions of being short on acids as under those of being short on potassium iodide (Fig. 3); accurate titration results were not ensured. To conclude, an appropriate minimum amount of
Fig. 5. Dependencies of the apparent potassium dichromate and potassium iodate assays by gravimetric titration with thiosulfate on the amount of sulfuric acid used. Iodine liberation times after the acid addition were 10 min for potassium dichromate and 0 min for potassium iodate. ○: 0.2 g of potassium dichromate dissolved in 100 mL of non-deaerated water; 3 g of potassium iodide. □: 0.2 g of potassium dichromate dissolved in 100 mL of deaerated water; 3 g of potassium iodide. ×: 0.15 g of potassium iodate (NMIJ CRM 3006-a) dissolved in 100 mL of non-deaerated water; 3 g of potassium iodide [7]. −: Certified value (99.974%) of potassium dichromate by coulometric titration with electrogenerated Fe(II) (not related to the abscissa axis) [13].
Gravimetric titration with thiosulfate was carried out to evaluate the dependency of apparent assay of potassium dichromate on the iodine liberation time. Fig. 6 shows the results of the gravimetric titration together with the apparent assay of potassium iodate (NMIJ CRM 3006a) [7]. Approximately 0.2 g of potassium dichromate was dissolved in 100 mL of deaerated or non-deaerated water followed by adding 3 g of potassium iodide and 1 mL of 9 mol L−1 sulfuric acid in turn; liberated iodine was titrated with a thiosulfate solution at 0 min to 20 min after the acid addition. The apparent assays by gravimetric titration at 10 min or more after the acid addition using deaerated or nondeaerated water were consistent with the certified value by coulometric titration with electrogenerated Fe(II); however, those by gravimetric titration at less than 10 min became higher. The apparent assay of potassium dichromate dissolved in deaerated or non-deaerated water by gravimetric titration at 0 min after the acid addition was more than 100.6%. The indicator current during titration also drifted upward under the conditions of being short on liberation time. It was also suggested that the positive bias at shorter iodine liberation time resulted from complex formation between Cr(III) and sodium thiosulfate induced by unreacted potassium dichromate [12] as mentioned above. On the other hand, potassium iodate oxidized iodide ions in a couple of seconds; therefore, even shorter iodine liberation time caused no bias. 3.5. Influence of background oxidation of iodide ions and vaporization of liberated iodine The background oxidation of iodide ions gives positive bias to the potassium dichromate assay and the vaporization of liberated iodine gives negative bias to it. Constant-voltage biamperometry was employed to estimate the background oxidation of iodide ions. Gravimetric titration with thiosulfate was performed to evaluate the vaporization of liberated iodine under a stoppered beaker and an unstoppered one during the libetaion. The authors previously reported the impacts of the oxidation of iodide ions in an acidic media [7,15]. In the previous report [7], indicator
Fig. 6. Dependency of the apparent potassium dichromate and potassium iodate assays by gravimetric titration with thiosulfate on the iodine liberation time in a stoppered beaker. The condition of 3 g of potassium iodide and 1 mL of 9 mol L−1 sulfuric acid was used. ○: 0.2 g of potassium dichromate dissolved in 100 mL of non-deaerated water. □: 0.2 g of potassium dichromate dissolved in 100 mL of deaerated water. ×: 0.15 g of potassium iodate (NMIJ CRM 3006-a) dissolved in 100 mL of non-deaerated water [7]. −: Certified value (99.974%) of potassium dichromate by coulometric titration with electrogenerated Fe(II) (not related to the abscissa axis) [13].
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current of constant-voltage biamperometry was monitored for 100 mL of non-deaerated water containing 1 g to 5 g of potassium iodide and 1 mL to 5 mL of 9 mol L−1 sulfuric acid without potassium dichromate nor potassium iodate in both the presence and the absence of room light. The current always increased acceleratively with time in a bright room and proportionally with time in a darkroom. The larger amount of sulfuric acid (within 1 mL to 5 mL of 9 mol L−1 in 100 mL of water) or surprisingly the smaller amount of potassium iodide (within 2 g to 5 g in 100 mL of water), the larger increasing rate of the current under each condition of a bright room and a dark one. Less than 2 g of potassium iodide caused smaller increasing rate of the current compared with 2 g of potassium iodide. The cause of the strange dependency of the current on iodide ion concentrations remains unclear. The impact of the air-oxidation to the assay of potassium dichromate under the used condition of acid concentrations was larger than that under the used condition of iodide ion concentrations; however, the impacts were always less than 0.01% at 10 min with below 5 g of potassium iodide and below 2 mL of 9 mol L− 1 sulfuric acid in a darkroom, and less than 0.03% at 10 min with below 5 g of potassium iodide and 3 mL to 4 mL of 9 mol L−1 sulfuric acid in a darkroom. In the examination of the dependency of the apparent potassium dichromate assay shown in Fig. 3, the assays slightly increased under the conditions using 3 g or more of potassium iodide. The impact of the background oxidation was less than 0.001% of the potassium dichromate assay under the condition using 5 g of potassium iodide and 1 mL of 9 mol L−1 sulfuric acid in non-deaerated water at 10 min after acid addition in a darkroom; therefore, the positive bias for higher concentrations of potassium iodide in Fig. 3 was not caused by the simple background oxidation of iodide ions without light. It is known that Cr(III) accelerates airoxidation of iodide ions [19]. Fig. 7 illustrates some examples of changes in background currents resulting from the oxidation of iodide ions. The changes in the indicator currents were monitored under the conditions of different amounts of 9 mol L−1 sulfuric acid and of potassium iodide in a darkroom. The indicator currents could be converted to the impact on the assays of potassium dichromate based on the titration curve. The effects of the oxidation of iodide ions under the conditions with 1 mL to 5 mL of 9 mol L−1 sulfuric acid and 1 g to 5 g of potassium iodide were
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negligible for the assays of potassium iodate or potassium bromate as the authors previously mentioned [7,15]. However, the effects under the presence of Cr(III) were significant in fact. The indicator currents significantly increased for the solutions containing both chromium(III) potassium sulfate dodecahydrate, equivalent number of Cr(VI) in 0.2 g potassium dichromate, and iron(II) sulfate, less than one-tenth number of the Cr(VI), which was used to reduce Cr(VI) of a possible impurity. Added Fe(II) might have consumed a part of dissolved oxygen, and even if the Cr(III) compounds contained any impurity oxidants such as Cr(VI) which could oxidize iodide ions, Fe(II) could have reducing effect on the impurities. Therefore, it was obviously confirmed that airoxidation of iodide ions were accelerated by the Cr(III) compound itself. If any solid Cr(III) compound is used, no dissolved oxygen is additionally introduced with Cr(III) and on the other hand there is a possibility of introducing any impurity oxidants. As a matter of fact, the similar results were obtained by directly using solid chromium(III) potassium sulfate dodecahydrate. In the case of 3 g of potassium iodide and 1 mL of 9 mol L−1 sulfuric acid with Cr(III) in Fig. 7, the impact of generated iodine on the assay of 0.2 g of potassium dichromate was over 0.3% at 10 min of the iodine liberation time. Chromium(III) had a role to accelerate air-oxidization of iodide ions; consequently, it led inaccurate standardization of sodium thiosulfate to some degree. Though these positive impacts were too large, it was concluded that the air-oxidation of iodide ions accelerated by Cr(III) was the main cause of the bias for 3 g to 5 g of potassium iodide shown in Fig. 3, because the actual experimental condition under which Cr(III) was generated with the progress of the titration was different from the condition of Fig. 3 under which Cr(III) always existed since the initial point. Gravimetric titration of potassium dichromate with thiosulfate in an unstoppered beaker was performed to evaluate the degree of iodine vaporization during 5 min to 20 min after acid addition under the condition using 3 g of potassium iodide and 1 mL of 9 mol L−1 sulfuric acid. The influence of the iodine vaporization to the potassium dichromate assay was ca. 0.02% per minute in an unstoppered beaker after the acid addition. Most of the liberated iodine was titrated within 30 s after removing the stopper for the usual titrations; therefore, the bias from iodine vaporization loss fell much smaller than half of 0.02%. According to our previous paper [7], the degree of iodine vaporization under the condition of 3 g of potassium iodide and 1 mL of 9 mol L−1 sulfuric acid in 100 mL of non-deaerated water containing 0.15 g of potassium iodate were 0.017% min−1 and 0.0006% min−1 in an unstoppered beaker and a stoppered one, respectively. Therefore, degree of iodine vaporization in a stoppered beaker is negligible; roughly similar degree is also seen from the results between 10 min and 20 min in Fig. 6. More acids and more potassium iodide could accelerate the iodine liberation; consequently, it could allow us to titrate potassium dichromate with a thiosulfate solution at 10 min or earlier after acid addition. Besides shorter liberation duration could be of advantage to avoid oxidation of potassium iodide. However, more acids and more potassium iodide increase the background and give the positive error to the assay of potassium dichromate. 3.6. Estimation of accuracy of standardization with potassium dichromate
Fig. 7. The iodine liberation process monitored by constant-voltage biamperometry without of potassium dichromate using a dual platinum-plate electrode (applied constantvoltage 150 mV), stirring at approximately 400 rpm. Increase of background current due to air-oxidation of iodide ions was monitored every 1 min after adding 1 mL of 9 mol L−1 sulfuric acid following the addition of 3 g or 5 g of potassium iodide to 100 mL of water with/without Cr(III) and Fe(II). Cr(III): chromium(III) potassium sulfate dodecahydrate, equivalent number of Cr(VI) in 0.2 g potassium dichromate; Fe(II): less than one-tenth number of the Cr(VI).
The authors developed NMIJ CRM 3006-a potassium iodate according to the article reported [7]; the certified value was 99.973% ± 0.018% (the value following ± indicates the expanded uncertainty with the coverage factor 2). The uncertainty value includes all components from the homogeneity, the stability, etc. The titration used for potassium iodate was very robust with respect to changes in the amounts of acid and potassium iodide and also the liraration time. The accuracy of the standardization of sodium thiosulfate with potassium dichromate was roughly estimated. An appropriate condition for the standardization with potassium dichromate was assumed as follows: 3 g of potassium iodide, 1 mL to 2 mL of 9 mol L−1 sulfuric acid in 100 mL of deaerated water, and 10 min liberation time. The uncertainty
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was calculated when conditions were changed within 2 g to 4 g of potassium iodide, 1 mL to 2 mL of 9 mol L−1 sulfuric acid and 5 min to 15 min of the liberation time in 100 mL of deaerated water. The possible biases of the assays at most were 0.024% for potassium iodide, 0.005% for acid and 0.026% for time. The bias of the standardization could be often up to 0.1% judging from this estimation. In the case of being short on acid and inadequate deaeration, the bias might exceed even 1%. These were all positive biases from the assays under the appropriate condition. Compared with potassium iodate, the visual endpoint detection (e.g., with starch) for potassium dichromate has a small disadvantage of hard color judgment. 4. Conclusions Potassium dichromate assays through two different paths were compared: direct coulometric titration with electrogenerated Fe(II) and gravimetric titration through the iodine liberation reaction with a sodium thiosulfate solution of which concentration was determined by coulometric titration with electrogenerated iodine. The accuracy of the latter gravimetric titration was discussed. When analysts utilize the reaction to standardize a sodium thiosulfate solution with potassium dichromate, the bias from the accurate concentration of the sodium thiosulfate solution might exceed even 1% unless an appropriate condition is chosen. The results of the gravimetric titration under an appropriate condition of iodine liberation were in good agreement with those of the former direct coulometric titration, though allowance of the appropriate condition for iodine liberation was very narrow. Impurities in reagents, sulfuric acid and potassium iodide, involved in oxidation-reduction reaction could affect the results of the standardization in principle. Analysts could need blank testing to estimate the effect of reagents used on the standardization. The testing could also give important information of the influence of light; the effect on the standardization was negligible because there was no dependency of the assays on the amount of sulfuric acid and potassium iodide with potassium iodate in the present study. Though potassium dichromate as a reference to standardize sodium thiosulfate is not superior to potassium iodate from the viewpoint of robustness against condition changes, it was confirmed in the present study that potassium dichromate was also reliably available.
References [1] L.W. Winkler, Die bestimmung des im wasser gelösten sauerstoffes, Ber. Dtsch. Chem. Ges. 21 (1888) 2843–2855. [2] ISO 5813, Water Quality — Determination of Dissolved Oxygen — Iodometric Method, International Organization for Standardization, Geneva, 1983. [3] I.M. Kolthoff, R. Belcher, Volumetric Analysis, Volume III Titration Methods: Oxidation-Reduction Reactions, Interscience publications, Inc., New York, 1957. [4] American Chemical Society, American Chemical Society Specifications — Reagent Chemicals, 9th ed. Oxford University Press, New York, 2000. [5] JIS K 8001, General Rule for Test Methods of Reagents, Japanese Industrial Standards Committee, Tokyo, 2009. [6] T. Asakai, M. Murayama, Scheme and studies of reference materials for volumetric analysis in Japan, Accred. Qual. Assur. 13 (2008) 351–360. [7] T. Asakai, A. Hioki, Investigation of iodine liberation process in redox titration of potassium iodate with sodium thiosulfate, Anal. Chim. Acta 689 (2011) 34–38. [8] C.R. McCrosky, The oxidizing action of potassium dichromate as compared with that of pure iodine, J. Am. Chem. Soc. 40 (1918) 1662–1674. [9] W.C. Vosburgh, Potassium dichromate as a standard in iodimetry and the determination of chromates by the iodide method, J. Am. Chem. Soc. 44 (1922) 2120–2130. [10] W.C. Bray, H.E. Miller, The standardization of thiosulfate solution by the permanganate–iodide methods, J. Am. Chem. Soc. 46 (1924) 2204–2211. [11] S. Popoff, J.L. Whitman, Standardization of solutions used in iodimetry II, J. Am. Chem. Soc. 47 (1925) 2259–2275. [12] F.L. Hahn, The reactions taking place in the iodimetric determination of chromates, J. Am. Chem. Soc. 57 (1935) 614–616. [13] T. Asakai, A. Hioki, Comparison of three electrochemical end-point detection methods to assay potassium dichromate by coulometric titration, Accred. Qual. Assur. 17 (2012) 45–52. [14] T. Asakai, A. Hioki, Precise coulometric titration of cerium(IV) as an oxidising agent with electrogenerated iron(II) and reliability in cerium(IV) standardisation with sodium thiosulfate, Anal. Methods 4 (2012) 3478–3483. [15] T. Asakai, A. Hioki, Potassium bromate assay by primary methods through iodine liberation reaction, Anal. Methods 5 (2013) 6240–6245. [16] T. Asakai, M. Murayama, T. Tanaka, Precise coulometric titration of sodium thiosulfate and development of potassium iodate as a redox standard, Talanta 73 (2007) 346–351. [17] T. Tanaka, H. Hayashi, Y. Komiya, H. Nabekawa, H. Hayashi, Stability of sodium thiosulfate solution and constant-current coulometric iodometric titration of potassium iodate, Bunseki Kagaku 56 (2007) 327–332. [18] V.P.R. Rao, B.V.S. Sarma, Standardization of thiosulfate using potassium dichromate by direct titration, Anal. Chem. 37 (1965) 1373–1375. [19] A.I. Vogel, Vogel's Textbook of Quantitative Inorganic Analysis, 4th ed. Longman, London, 1978. 376.
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Microchemical Journal journal homepage: www.elsevier.com/locate/microc
Reliability in standardization of sodium thiosulfate with potassium dichromate Toshiaki Asakai ⁎, Akiharu Hioki National Metrology Institute of Japan, National Institute of Advanced Industrial Science Technology, 1-1-1 Umezono, Tsukuba, Ibaraki 305-8563, Japan
a r t i c l e
i n f o
Article history: Received 23 March 2015 Received in revised form 12 May 2015 Accepted 12 May 2015 Available online 16 May 2015 Keywords: Coulometric titration Oxidation-reduction titration Certified reference material Potassium dichromate Sodium thiosulfate Potassium iodate
a b s t r a c t Sodium thiosulfate is often standardized with potassium dichromate. In the standardization, iodine (triiodide) liberated by potassium dichromate in an acidic potassium iodide solution is titrated with a sodium thiosulfate solution. The iodine liberation process significantly affects the titration results. In the present study, the accuracy of the liberation process was examined by assaying potassium dichromate through two different paths: assaying directly by coulometric titration with electrogenerated Fe(II), and assaying by gravimetric titration through the iodine liberation reaction with a sodium thiosulfate solution of which concentration was standardized by coulometric titration with electrogenerated iodine. The accuracy of the standardization of a sodium thiosulfate solution by potassium dichromate was discussed from the apparent assays of potassium dichromate under different measurement conditions. © 2015 Elsevier B.V. All rights reserved.
1. Introduction Sodium thiosulfate is a useful reducing agent for a titrant solution in volumetric analyses such as iodometry and the Winkler method to measure dissolved oxygen [1,2]. Oceanography community requires the accuracy and the reproducibility of the measurements of dissolved oxygen; the accurate standardization procedure of sodium thiosulfate is needed. The agent is often standardized with potassium dichromate and, on the other hand, is usually standardized with potassium iodate in Japan [3–6]. Iodine (triiodide ions) liberated by potassium dichromate or potassium iodate in an acidic potassium iodide solution is titrated with a sodium thiosulfate solution to standardize the thiosulfate [reactions (1), (2) and (3)]. K2 Cr2 O7 þ 6KI þ 7H2 SO4 →Cr2 ðSO4 Þ3 þ 4K2 SO4 þ 7H2 O þ 3I2
ð1Þ
KIO3 þ 5KI þ 3H2 SO4 →3K2 SO4 þ 3H2 O þ 3I2
ð2Þ
2Na2 S2 O3 þ I2 →2NaI þ Na2 S4 O6
ð3Þ
⁎ Corresponding author. Tel.: +81 29 861 6881; fax: +81 29 861 6890. E-mail address: [email protected] (T. Asakai).
http://dx.doi.org/10.1016/j.microc.2015.05.012 0026-265X/© 2015 Elsevier B.V. All rights reserved.
The iodine liberation process is significantly affected by the amounts of acid and potassium iodide added, the waiting time for the liberation, and light; therefore, the process plays a key role for the accuracy of standardization of sodium thiosulfate. The authors reported an accurate standardization procedure for a sodium thiosulfate solution with potassium iodate [7]. It has been pointed out that there are some biases in the standardization of sodium thiosulfate with potassium dichromate [3,8–12]. Some researchers reported that excess consumption of thiosulfate occurred in an acidic medium due to air oxidation of iodide ions [8,9]. Hahn described that the main reason of the excess consumption of thiosulfate in a weak acidic medium was not air oxidation of iodide, but a complex between Cr(III) and thiosulfate, which reacted with iodine slowly [3,12]. Kolthoff and Belcher reviewed the studies on the titration procedure [3] and mentioned the appropriate conditions to ensure complete and stoichiometric reaction with dichromate: the hydrogen ions concentration 0.2 mol L−1 to 0.4 mol L−1, 2% of potassium iodide, and standing for 10 min to 15 min before starting titration with thiosulfate. The impact of the biases has remained obscure because the accuracy of the standardization of thiosulfate has been discussed by using relative methods for conventional titration. Thiosulfate is usually standardized with sublimed iodine whose purity is hypothetically regarded as 100%. In the present study, the accuracy of the standardization was absolutely evaluated by coulometric titration (Fig. 1). Coulometric titration absolutely yields an accurate oxidimetric–reductometric assay based on Faraday's Law. The accuracy of the iodine liberation process was examined by comparing two assays of potassium dichromate using different
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1 μg, and potassium iodide and solutions were weighed with resolution of 10 μg. 3. Results and discussion 3.1. Determination of sodium thiosulfate by coulometric titration and uncertainty estimation
Fig. 1. Experimental design for the assay of potassium dichromate to investigate the iodine liberation reaction.
titrations based on coulometric titration. Firstly, potassium dichromate was directly assayed by coulometric titration with electrogenerated Fe(II) [13]. Secondly, potassium dichromate was assayed by gravimetric titration through the iodine liberation reaction with a sodium thiosulfate solution, of which concentration had been standardized by coulometric titration with electrogenerated iodine. The accuracy of the standardization of a sodium thiosulfate solution by potassium dichromate was discussed from the apparent assays of potassium dichromate under different measurement conditions.
2. Material and methods Constant-voltage biamperometry (a dead stop method) was utilized to measure the amount of liberated iodine under several conditions as a preliminary experiment. A DC source and a digital multimeter with a dual platinum-chip electrode (each platinum chip: 0.5 mm ϕ × 0.5 mm long) were used for constant-voltage biamperometry. Constant voltage was applied between the platinum pair, and the current was monitored. The current had good proportionality to the amount of liberated iodine [7]. A dual platinum-plate electrode (each 5 mm × 5 mm) was used in gravimetric titration to increase sensitivity as described previously [7, 13–15]. The experimental set-ups and the procedures of coulometric titrations for potassium dichromate and a sodium thiosulfate solution were described previously [7,13–15]. As described in the previous report [13], high purity potassium dichromate as a certified reference material (CRM) was established by coulometric titration with electrogenerated Fe(II). The certified value (mass fraction) of the CRM, NMIJ CRM 3002-a, is 99.974% ± 0.011%, where the value following ± indicates the expanded uncertainty with the coverage factor 2. Gravimetric titration was carried out to assay potassium dichromate with a sodium thiosulfate solution through the iodine liberation reaction in the following procedure: approximately 0.2 g of potassium dichromate were placed in a 200 mL tall beaker, it was dissolved in 100 mL of water, potassium iodide and 9 mol L−1 sulfuric acid were added, and the solution was titrated with a sodium thiosulfate solution. The concentration of the sodium thiosulfate solution was standardized in advance by coulometric titration with electrogenerated iodine [7]. A plastic syringe with a PEEK needle was used to weigh the sodium thiosulfate solution. Potassium dichromate was weighed with resolution of
The concentration of a sodium thiosulfate solution (3 g to 8 g) standardized by coulometric titration was, for example, 208.653 mmol kg−1 (RSD 0.007%, n = 5) (‘RSD’ means ‘relative standard deviation’, and ‘n’ is the number of measurements under a repeating condition). The uncertainty sources for the concentration of the solution (ca. 3 g of ca. 200 mmol kg−1) were as follows (the number in each parenthesis indicates the relative standard uncertainty): repeatability (0.0032%, RSD of the mean, n = 5), Faraday constant (0.000 002 2%), a standard resistor (0.000 65%, calibration and the influence of temperature), a frequency counter (0.000 05%), a voltmeter (0.000 15%) and weighing (0.0019%, sensitivity and linearity of the balance). Finally, the combined standard uncertainty was 0.0038% relative and the expanded uncertainty with a coverage factor 2 was 0.0076% relative. The coulometric titration allowed the authors to discuss the accuracy of the iodine liberation reaction by potassium dichromate in the order of magnitude of 0.01%. A sodium thiosulfate solution is stable enough to be required in the present study [16,17]. 3.2. Influence of the amount of potassium iodide added Approximately 0.2 g of potassium dichromate were dissolved in 100 mL of non-deaerated water followed by adding 1 g to 6 g of potassium iodide and 1 mL of 9 mol L−1 sulfuric acid in turn; after the acid addition, iodine liberation started. The iodine liberation process monitored by constant-voltage biamperometry using a dual platinum-chip electrode (applied constant-voltage 500 mV) is shown in Fig. 2; the indicator current corresponds to the amount of iodine liberation. Under the condition using 1 g to 2 g of potassium iodide and 1 mL of 9 mol L−1 sulfuric acid, lacking of iodide ions in the solution led to the deposition of solid iodine, and the indicator current drastically decreased in unforeseeable timing after the initial increase. Under the conditions using 3 g to 6 g of potassium iodide and the same amount of sulfuric acid, the
Fig. 2. The iodine liberation process monitored by constant-voltage biamperometry for 0.2 g of potassium dichromate using a dual platinum-chip electrode (applied constantvoltage 500 mV), stirring at 900 revolutions per minute (rpm). The sample was dissolved in 100 mL of non-deaerated water, and 1 g to 6 g of potassium iodide and 1 mL of 9 mol L−1 sulfuric acid were added.
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decrease of the indicator current did not occur; these conditions seemed to be enough to completely generate dissolved iodine. In the previous study [7], potassium iodate (0.15 g), whose redox equivalent is almost the same as that of potassium dichromate (0.2 g), liberated iodine completely in a couple of seconds after acid addition under the condition using 3 g to 5 g of potassium iodide and 1 mL of 9 mol L−1 sulfuric acid; consequently, no waiting time for the iodine liberation was necessary. Potassium dichromate conversely required more than 5 min for the complete iodine liberation according to Fig. 2. Gravimetric titration was performed to assay potassium dichromate in a darkroom using a sodium thiosulfate standardized by coulometric titration. The dependency of the apparent potassium dichromate assay on the amount of potassium iodide added is shown in Fig. 3. Approximately 0.2 g of potassium dichromate in a beaker was dissolved in 100 mL of deaerated or non-deaerated water followed by adding 1 g to 5 g of potassium iodide and 1 mL of 9 mol L−1 sulfuric acid in turn; liberated iodine was titrated at 10 min after the acid addition with a sodium thiosulfate solution, of which concentration was standardized by coulometric titration. The beaker was stoppered with Parafilm during the liberation to avoid the volatilization of iodine. Fig. 3 includes the certified value of potassium dichromate determined by coulometric titration with electrogenerated Fe(II) [13] and also the dependency of the apparent potassium iodate assay on the amount of potassium iodide added [7]. Liberated iodine by potassium iodate was titrated immediately after iodine liberation starting. The similarity of the assay values of potassium dichromate and potassium iodate was a coincidence because of high purity materials. The apparent assay of potassium dichromate with 1 g of potassium iodide and 1 mL of 9 mol L−1 sulfuric acid using deaerated or non-deaerated water was more than 101%. The apparent assays of potassium dichromate had a minimum for both the results using deaerated water and those using non-deaerated water; the minimum value with 2 g of potassium iodide was close to the certified value. The titration with less than 2 g of potassium iodide was not able to be accurately completed since the indicator current drifted upward during the titration. In this iodometry, the indicator current decreased with progress of titration, that is, the liberated iodine was reduced with a sodium thiosulfate solution. The end point of this type of reaction was at 0 A of the indicator current where the iodine just disappeared. The apparent assay of potassium dichromate corresponded to the mass of the sodium thiosulfate solution consumed until the end point. The upward drift of the indicator current indicated
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the production of iodine or oxidants with progress of time. The drift made it difficult to draw the titration curve and to determine the end point accurately. The apparent assays were also significantly affected by the speed of titration. Less amount of potassium iodide and slow titration led to more evaporation of iodine and more influence of light; accurate titration results were not expected. The iodine liberations were obviously not completed under the condition using 1 g of potassium iodide and 1 mL of 9 mol L−1 sulfuric acid for either 0.2 g potassium dichromate or 0.15 g of potassium iodate [7] as shown in the preliminary examination by constant-voltage biamperometry (Fig. 2). Consequently, unreacted potassium dichromate and potassium iodate remained in the solutions before starting a titration with thiosulfate. The authors reported that unexpected good results for potassium iodate under incomplete iodine liberation conditions were caused by resupplied iodide ions with progress of the reaction [7]. Such good results did not happen for potassium dichromate as shown in Fig. 3. Judging from Fig. 3, it was inferred that excess sodium thiosulfate was consumed in the presence of the unreacted potassium dichromate. There could be the following three possibilities: (1) sodium thiosulfate was insufficiently oxidized, (2) it was decomposed under the incomplete iodine liberation in an acidic medium, and (3) it interacted with potassium dichromate or generated Cr(III) and lost the reducing capability. If the remaining dichromate caused insufficient oxidation of thiosulfate, it would result in excessive consumption of thiosulfate. Rao and Sarma reported the direct standardization of thiosulfate by potassium dichromate with copper(II) sulfate as a catalyst [18] and mentioned that dichromate directly oxidized thiosulfate to tetrathionate state as usual. Therefore, excessive consumption of thiosulfate was not caused by its insufficient reduction. Though the possibility of its decomposition before reacting with remaining dichromate or iodine cannot be excluded, as mentioned above Hahn reported a phenomenon relevant to the third possibility [12]. It was the generation of the complex between thiosulfate and Cr(III) in the presence of dichromate; consequently, excess thiosulfate was apparently consumed. According to his report, Cr(III) once formed does not react with thiosulfate, but the complex between Cr(III) and thiosulfate is formed when the reduction of Cr(VI) takes place in the presence of thiosulfate. The reason of slight increase in the apparent assays of potassium dichromate under the condition using more than 3 g of potassium iodide could be the significant oxidation of potassium iodide during iodine liberation (10 min) in an acidic medium depending on the amount of potassium iodide added. Significant air-oxidation of iodide ions in an acidic medium without oxidants such as potassium iodate was not observed [7,15]; however, Cr(III) could accelerate the oxidation reaction of them (details are described in section Influence of background oxidation of iodide ions and vaporization of liberated iodine). 3.3. Influence of the amount of acid added
Fig. 3. Dependencies of the apparent potassium dichromate and potassium iodate assays by gravimetric titration with thiosulfate on the amount of potassium iodide added. Iodine liberation times after the acid addition were 10 min for potassium dichromate and 0 min for potassium iodate. ○: 0.2 g of potassium dichromate dissolved in 100 mL of nondeaerated water; 1 mL of 9 mol L−1 sulfuric acid. □: 0.2 g of potassium dichromate dissolved in 100 mL of deaerated water; 1 mL of 9 mol L−1 sulfuric acid. ×: 0.15 g of potassium iodate (NMIJ CRM 3006-a) dissolved in 100 mL of non-deaerated water; 1 mL of 9 mol L−1 sulfuric acid [7]. −: Certified value (99.974%) of potassium dichromate by coulometric titration with electrogenerated Fe(II) (not related to the abscissa axis) [13].
Approximately 0.2 g of potassium dichromate was dissolved in 100 mL of non-deaerated water followed by adding 3 g of potassium iodide and 0.5 mL to 5 mL of 9 mol L−1 sulfuric acid in turn; after acid addition iodine liberation started. The iodine liberation process was monitored in the same manner as done for the examination of influence of the amount of potassium iodide added (vide supra); the results are shown in Fig. 4. Under the condition using 3 g of potassium iodide and 0.5 mL of 9 mol L−1 sulfuric acid, the iodine liberation was not completed. Under the condition using 3 g of potassium iodide and 1 mL to 5 mL of 9 mol L−1 sulfuric acid, the iodine liberation was apparently completed. It seems that more acids accelerate the iodine liberation reaction. Gravimetric titration with thiosulfate was performed to evaluate the dependency of apparent assay of potassium dichromate on the amount of sulfuric acid used. A dual platinum-plate electrode was used for the end-point determination. Fig. 5 shows the results of the gravimetric titration together with the apparent assays of potassium iodate [7]. Approximately 0.2 g of potassium dichromate were dissolved in 100 mL of
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acid is required in order to accurately carry out the standardization of sodium thiosulfate with potassium dichromate. In addition, deaeration is desirable to avoid the excess oxidation of iodide by dissolved oxygen. 3.4. Influence of the iodine liberation time
Fig. 4. The iodine liberation process monitored by constant-voltage biamperometry for 0.2 g of potassium dichromate using a dual platinum-chip electrode (applied constantvoltage 500 mV), stirring at 900 rpm. The sample was dissolved in 100 mL of nondeaerated water, and 3 g of potassium iodide and 0.5 mL to 5 mL of 9 mol L−1 sulfuric acid were added.
deaerated or non-deaerated water followed by adding 3 g of potassium iodide and 0.5 mL to 4 mL of 9 mol L−1 sulfuric acid in turn; liberated iodine was titrated with a thiosulfate solution at 10 min after acid addition. The results by gravimetric titration using 100 mL of deaerated water and 1 mL or more of 9 mol L−1 sulfuric acid were consistent with those by coulometric titration with electrogenerated Fe(II); on the other hand, those by gravimetric titration using non-deaerated water increased with the increase in the amount of acid used. This fact suggested that dissolved oxygen oxidized iodide ions in an acidic media described in detail below [7]. The assay of potassium dichromate dissolved in 100 mL of deaerated or non-deaerated water, under the condition using 3 g of potassium iodide and 0.5 mL of 9 mol L−1 sulfuric acid, was more than 106%. The indicator current drifted upward during the titration under the conditions of being short on acids as under those of being short on potassium iodide (Fig. 3); accurate titration results were not ensured. To conclude, an appropriate minimum amount of
Fig. 5. Dependencies of the apparent potassium dichromate and potassium iodate assays by gravimetric titration with thiosulfate on the amount of sulfuric acid used. Iodine liberation times after the acid addition were 10 min for potassium dichromate and 0 min for potassium iodate. ○: 0.2 g of potassium dichromate dissolved in 100 mL of non-deaerated water; 3 g of potassium iodide. □: 0.2 g of potassium dichromate dissolved in 100 mL of deaerated water; 3 g of potassium iodide. ×: 0.15 g of potassium iodate (NMIJ CRM 3006-a) dissolved in 100 mL of non-deaerated water; 3 g of potassium iodide [7]. −: Certified value (99.974%) of potassium dichromate by coulometric titration with electrogenerated Fe(II) (not related to the abscissa axis) [13].
Gravimetric titration with thiosulfate was carried out to evaluate the dependency of apparent assay of potassium dichromate on the iodine liberation time. Fig. 6 shows the results of the gravimetric titration together with the apparent assay of potassium iodate (NMIJ CRM 3006a) [7]. Approximately 0.2 g of potassium dichromate was dissolved in 100 mL of deaerated or non-deaerated water followed by adding 3 g of potassium iodide and 1 mL of 9 mol L−1 sulfuric acid in turn; liberated iodine was titrated with a thiosulfate solution at 0 min to 20 min after the acid addition. The apparent assays by gravimetric titration at 10 min or more after the acid addition using deaerated or nondeaerated water were consistent with the certified value by coulometric titration with electrogenerated Fe(II); however, those by gravimetric titration at less than 10 min became higher. The apparent assay of potassium dichromate dissolved in deaerated or non-deaerated water by gravimetric titration at 0 min after the acid addition was more than 100.6%. The indicator current during titration also drifted upward under the conditions of being short on liberation time. It was also suggested that the positive bias at shorter iodine liberation time resulted from complex formation between Cr(III) and sodium thiosulfate induced by unreacted potassium dichromate [12] as mentioned above. On the other hand, potassium iodate oxidized iodide ions in a couple of seconds; therefore, even shorter iodine liberation time caused no bias. 3.5. Influence of background oxidation of iodide ions and vaporization of liberated iodine The background oxidation of iodide ions gives positive bias to the potassium dichromate assay and the vaporization of liberated iodine gives negative bias to it. Constant-voltage biamperometry was employed to estimate the background oxidation of iodide ions. Gravimetric titration with thiosulfate was performed to evaluate the vaporization of liberated iodine under a stoppered beaker and an unstoppered one during the libetaion. The authors previously reported the impacts of the oxidation of iodide ions in an acidic media [7,15]. In the previous report [7], indicator
Fig. 6. Dependency of the apparent potassium dichromate and potassium iodate assays by gravimetric titration with thiosulfate on the iodine liberation time in a stoppered beaker. The condition of 3 g of potassium iodide and 1 mL of 9 mol L−1 sulfuric acid was used. ○: 0.2 g of potassium dichromate dissolved in 100 mL of non-deaerated water. □: 0.2 g of potassium dichromate dissolved in 100 mL of deaerated water. ×: 0.15 g of potassium iodate (NMIJ CRM 3006-a) dissolved in 100 mL of non-deaerated water [7]. −: Certified value (99.974%) of potassium dichromate by coulometric titration with electrogenerated Fe(II) (not related to the abscissa axis) [13].
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current of constant-voltage biamperometry was monitored for 100 mL of non-deaerated water containing 1 g to 5 g of potassium iodide and 1 mL to 5 mL of 9 mol L−1 sulfuric acid without potassium dichromate nor potassium iodate in both the presence and the absence of room light. The current always increased acceleratively with time in a bright room and proportionally with time in a darkroom. The larger amount of sulfuric acid (within 1 mL to 5 mL of 9 mol L−1 in 100 mL of water) or surprisingly the smaller amount of potassium iodide (within 2 g to 5 g in 100 mL of water), the larger increasing rate of the current under each condition of a bright room and a dark one. Less than 2 g of potassium iodide caused smaller increasing rate of the current compared with 2 g of potassium iodide. The cause of the strange dependency of the current on iodide ion concentrations remains unclear. The impact of the air-oxidation to the assay of potassium dichromate under the used condition of acid concentrations was larger than that under the used condition of iodide ion concentrations; however, the impacts were always less than 0.01% at 10 min with below 5 g of potassium iodide and below 2 mL of 9 mol L− 1 sulfuric acid in a darkroom, and less than 0.03% at 10 min with below 5 g of potassium iodide and 3 mL to 4 mL of 9 mol L−1 sulfuric acid in a darkroom. In the examination of the dependency of the apparent potassium dichromate assay shown in Fig. 3, the assays slightly increased under the conditions using 3 g or more of potassium iodide. The impact of the background oxidation was less than 0.001% of the potassium dichromate assay under the condition using 5 g of potassium iodide and 1 mL of 9 mol L−1 sulfuric acid in non-deaerated water at 10 min after acid addition in a darkroom; therefore, the positive bias for higher concentrations of potassium iodide in Fig. 3 was not caused by the simple background oxidation of iodide ions without light. It is known that Cr(III) accelerates airoxidation of iodide ions [19]. Fig. 7 illustrates some examples of changes in background currents resulting from the oxidation of iodide ions. The changes in the indicator currents were monitored under the conditions of different amounts of 9 mol L−1 sulfuric acid and of potassium iodide in a darkroom. The indicator currents could be converted to the impact on the assays of potassium dichromate based on the titration curve. The effects of the oxidation of iodide ions under the conditions with 1 mL to 5 mL of 9 mol L−1 sulfuric acid and 1 g to 5 g of potassium iodide were
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negligible for the assays of potassium iodate or potassium bromate as the authors previously mentioned [7,15]. However, the effects under the presence of Cr(III) were significant in fact. The indicator currents significantly increased for the solutions containing both chromium(III) potassium sulfate dodecahydrate, equivalent number of Cr(VI) in 0.2 g potassium dichromate, and iron(II) sulfate, less than one-tenth number of the Cr(VI), which was used to reduce Cr(VI) of a possible impurity. Added Fe(II) might have consumed a part of dissolved oxygen, and even if the Cr(III) compounds contained any impurity oxidants such as Cr(VI) which could oxidize iodide ions, Fe(II) could have reducing effect on the impurities. Therefore, it was obviously confirmed that airoxidation of iodide ions were accelerated by the Cr(III) compound itself. If any solid Cr(III) compound is used, no dissolved oxygen is additionally introduced with Cr(III) and on the other hand there is a possibility of introducing any impurity oxidants. As a matter of fact, the similar results were obtained by directly using solid chromium(III) potassium sulfate dodecahydrate. In the case of 3 g of potassium iodide and 1 mL of 9 mol L−1 sulfuric acid with Cr(III) in Fig. 7, the impact of generated iodine on the assay of 0.2 g of potassium dichromate was over 0.3% at 10 min of the iodine liberation time. Chromium(III) had a role to accelerate air-oxidization of iodide ions; consequently, it led inaccurate standardization of sodium thiosulfate to some degree. Though these positive impacts were too large, it was concluded that the air-oxidation of iodide ions accelerated by Cr(III) was the main cause of the bias for 3 g to 5 g of potassium iodide shown in Fig. 3, because the actual experimental condition under which Cr(III) was generated with the progress of the titration was different from the condition of Fig. 3 under which Cr(III) always existed since the initial point. Gravimetric titration of potassium dichromate with thiosulfate in an unstoppered beaker was performed to evaluate the degree of iodine vaporization during 5 min to 20 min after acid addition under the condition using 3 g of potassium iodide and 1 mL of 9 mol L−1 sulfuric acid. The influence of the iodine vaporization to the potassium dichromate assay was ca. 0.02% per minute in an unstoppered beaker after the acid addition. Most of the liberated iodine was titrated within 30 s after removing the stopper for the usual titrations; therefore, the bias from iodine vaporization loss fell much smaller than half of 0.02%. According to our previous paper [7], the degree of iodine vaporization under the condition of 3 g of potassium iodide and 1 mL of 9 mol L−1 sulfuric acid in 100 mL of non-deaerated water containing 0.15 g of potassium iodate were 0.017% min−1 and 0.0006% min−1 in an unstoppered beaker and a stoppered one, respectively. Therefore, degree of iodine vaporization in a stoppered beaker is negligible; roughly similar degree is also seen from the results between 10 min and 20 min in Fig. 6. More acids and more potassium iodide could accelerate the iodine liberation; consequently, it could allow us to titrate potassium dichromate with a thiosulfate solution at 10 min or earlier after acid addition. Besides shorter liberation duration could be of advantage to avoid oxidation of potassium iodide. However, more acids and more potassium iodide increase the background and give the positive error to the assay of potassium dichromate. 3.6. Estimation of accuracy of standardization with potassium dichromate
Fig. 7. The iodine liberation process monitored by constant-voltage biamperometry without of potassium dichromate using a dual platinum-plate electrode (applied constantvoltage 150 mV), stirring at approximately 400 rpm. Increase of background current due to air-oxidation of iodide ions was monitored every 1 min after adding 1 mL of 9 mol L−1 sulfuric acid following the addition of 3 g or 5 g of potassium iodide to 100 mL of water with/without Cr(III) and Fe(II). Cr(III): chromium(III) potassium sulfate dodecahydrate, equivalent number of Cr(VI) in 0.2 g potassium dichromate; Fe(II): less than one-tenth number of the Cr(VI).
The authors developed NMIJ CRM 3006-a potassium iodate according to the article reported [7]; the certified value was 99.973% ± 0.018% (the value following ± indicates the expanded uncertainty with the coverage factor 2). The uncertainty value includes all components from the homogeneity, the stability, etc. The titration used for potassium iodate was very robust with respect to changes in the amounts of acid and potassium iodide and also the liraration time. The accuracy of the standardization of sodium thiosulfate with potassium dichromate was roughly estimated. An appropriate condition for the standardization with potassium dichromate was assumed as follows: 3 g of potassium iodide, 1 mL to 2 mL of 9 mol L−1 sulfuric acid in 100 mL of deaerated water, and 10 min liberation time. The uncertainty
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was calculated when conditions were changed within 2 g to 4 g of potassium iodide, 1 mL to 2 mL of 9 mol L−1 sulfuric acid and 5 min to 15 min of the liberation time in 100 mL of deaerated water. The possible biases of the assays at most were 0.024% for potassium iodide, 0.005% for acid and 0.026% for time. The bias of the standardization could be often up to 0.1% judging from this estimation. In the case of being short on acid and inadequate deaeration, the bias might exceed even 1%. These were all positive biases from the assays under the appropriate condition. Compared with potassium iodate, the visual endpoint detection (e.g., with starch) for potassium dichromate has a small disadvantage of hard color judgment. 4. Conclusions Potassium dichromate assays through two different paths were compared: direct coulometric titration with electrogenerated Fe(II) and gravimetric titration through the iodine liberation reaction with a sodium thiosulfate solution of which concentration was determined by coulometric titration with electrogenerated iodine. The accuracy of the latter gravimetric titration was discussed. When analysts utilize the reaction to standardize a sodium thiosulfate solution with potassium dichromate, the bias from the accurate concentration of the sodium thiosulfate solution might exceed even 1% unless an appropriate condition is chosen. The results of the gravimetric titration under an appropriate condition of iodine liberation were in good agreement with those of the former direct coulometric titration, though allowance of the appropriate condition for iodine liberation was very narrow. Impurities in reagents, sulfuric acid and potassium iodide, involved in oxidation-reduction reaction could affect the results of the standardization in principle. Analysts could need blank testing to estimate the effect of reagents used on the standardization. The testing could also give important information of the influence of light; the effect on the standardization was negligible because there was no dependency of the assays on the amount of sulfuric acid and potassium iodide with potassium iodate in the present study. Though potassium dichromate as a reference to standardize sodium thiosulfate is not superior to potassium iodate from the viewpoint of robustness against condition changes, it was confirmed in the present study that potassium dichromate was also reliably available.
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