Chemical Education Today
Letters Standardizing Iodine The use of ascorbic acid as a standard (1) for volumetric iodine requires some warnings. It has been written “starch solution retards the reaction of ascorbic acid [with iodine] so that the titration is ‘extended’” (2). This is my own recollection of this titration when I did it routinely half a century ago. The difficulty may be circumvented by not using an indicator but relying on the color of the iodine: this is deep enough to be observed with a one-drop excess of 0.1 N I2 in an otherwise colorless solution. Alternatively, the same authors suggest (quoting ref 3) “a sharp end-point is obtained if Variamine Blue is used instead of starch solution as indicator”. That “the end point was determined when a pale blue color remained with the first excess of iodine” (1) was therefore a source of error. A number of substances have been sold commercially in primary-standard quality for standardizing iodine but they have not become popular. GFS Chemicals marketed them until recently but has discontinued all except arsenic(III) oxide, which is too toxic to be entrusted to students in introductory chemistry and has an inconveniently low equivalent weight. It is probably possible to prepare primary-standard-quality ascorbic acid: the work reported in ref 1 may be a motivation to do so. The material would have to be carefully stored. The problem can be circumvented altogether. Elemental iodine is economically available from Aesar or GFS Chemicals with 99.8% purity. Standard iodine solutions can therefore be made directly and conveniently with better accuracy than novices can standardize a solution. Still higher purities are available at unrealistic prices. An even better method would be to make the solution from Primary-Standard-grade potassium bi-iodate (99.95%) and reduce it to iodine with KI and HCl before filling the volumetric flask to the mark. This would avoid the difficulty of weighing a volatile substance. There is a more important and more general problem. The statement “its water content was very low (about 0.5%) even before drying, and so the standardization of the iodine solution is not affected when this procedure is carried out
with undried ascorbic acid” implies that a determinate error of 0.5% is acceptable. This is added to another determinate error, since no commercial ascorbic acid of which I am aware is guaranteed to be more than 99% pure. This raises the question of what accuracy should be demanded of introductory students. In my first job as an analytical technician (newly graduated from high school) I was instructed that volumetric analyses should be accurate to within 0.1%. With the passage of the years, I realized that this was unrealistic and that 0.2% repeatability should be expected from a professional and 0.3% from a novice; 1% would have led to reprimand if not dismissal. Both the authors and the reviewers must have been convinced that lower standards are appropriate in an introductory course. Are they right? Are we justified in encouraging practice that is less than professional? Harris found in a comparative study that students performed better in analytical chemistry labs if they had not done titrations in introductory chemistry labs, where they presumably learned sloppy technique. When we teach introductory chemistry we must necessarily teach some parts that are outside our own field of expertise, committing errors that we learned from our teachers and passing them on in good faith to our students and hence to their students. No solution has been found for this problem. Literature Cited 1. Silva, C. R.; Simoni, J. A.; Collins, C. H.; Volpe, P. L. O. J. Chem. Educ. 1999, 76, 1421–1422. 2. Strohecker, R.; Henning, H. M. Vitamin Assay—Tested Methods; translated by D. D. Libman; Verlag Chemie: Weinheim, 1965; p 228. 3. Erdey, L.; Kaplar, L. Z. Anal. Chem. 1958, 180. Stephen J. Hawkes Department of Chemistry Oregon State University Corvallis, OR 97331-4003
JChemEd.chem.wisc.edu • Vol. 77 No. 12 December 2000 • Journal of Chemical Education
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