G. Rarnarnurti, K. Renganathan, and L. R. Ganesan Madura College Madurai-625011, India
II
The Instability Constant of the 13II2 System
+
A student experiment
A reference to the commonly used textbooks on practical physical chemistry ( I ) reveals that the study of the triiadide-iodide-iodine system is invariably presented as a n exercise in the distribution of iodine between two liquid phases, water and CCl4. The experiment is in two parts, in the first of which the distribution ratio of iodine between water and CClr is determined. This information is used in the second part of the experiment for the evaluation of the instability constant of the triiodide-iodide-iodine system. In experimental courses where the aim is two fold, (1) to determine the equilibrium constant for the homogeneous equilibrium 13- F! I12 and (2) to measure the distribution ratio of iodine between two solvents, both parts of the exoeriment are necessarv. However if the aim is (1) . . stated above, an alternative shorter procedure is available which is simpler since it involves fewer titrations, avoids contamination of the aqueous sample withdrawn for titration by drons of the verv much more concentrated ~ h a s eof iodine in C C I ~(2), a n d dispenses with the somewhat difficult titration of t h e organic liquid layer which "presents some problems" being "slow" with an inherent difficulty in "avoidine overtitration" (3). Such an alternative procedure viz., equilibrating potassium iodide solutions of known concentration with solid iodine is so simple that one wonders why i t has not also been mentioned in the text books along with the usual longer procedure. A quick survey of the standard books on analysis shows that the following difficulties would be encountered.
+
1) Sdlid iodine dissolves surprisingly slowly in dilute potassium iodide media (4). 2) Triiodide solutions are inherently unstable hecause of the
volatility of iodine. 3) There is a possibility of onidation of idide by dissolved oxyKen. This is howe"er not really serious for essentially neutral
triiodide solutions. 4) Iodine solutions stored in an ordinary glass container for a long period of.time may contain a small amount of iodate. Glass may give off a little alkali which disproportionates iodine into iodide and iodate. 5) There is evidence for the existence of species such as 4-and 17- in very concentrated solutions (5). In the experiment described helow, an attempt was made to assess the usefulness of the alternative procedure as a student experiment. The error due to the volatility of iodine was minimized by using well-stoppered bottles. Alkali contamination by glass was avoided by using Pyrex bottles.
326 / Journal of Chemical Education
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Experimental Data Bot¶le No.
Concentrafion of KI = le)
moles liter-'
Concentration of total iodine
11,l + I!,-)
=
mole liter-'
6
e'/e
I /c liter mole-'
Temperature of experiment -26.4-C. Mean value of C'IC = 0.39844. Mean value of I/C = 5.07109 I mole-'.
Vigorous and continued shaking was done to reduce equilibration time.
Experimental Seven Pyrex bottles containing potassium iodide solutions of known concentrations and an eighth bottle containing pure water were shaken, after adding excess of solid iodine to each, while being kept in a thermostat. After allowing sufficient time for equilibration, the shaking was stopped and the solutions allowed to settle. Five milliliters of the solution was withdrawn from each bottle and titrated against standard sodium thiosulfate solution. Each, bottle was returned to the thermostat after withdrawal of the sample for titration andshaking was continued. This procedure was repeated till two successive titer values for the same bottle agreed. It was found that when shaking was stopped, the iodine crystals quickly settled at the bottom because of their high density. Thus withdrawing the clear liquid above the solid presented no problems. With good shaking, consistent titer values could he obtained in two or three hours. The table gives the results. Theory The equilibrium constant for the system IQ- I- 1% is
= +
If we use the approximate form of t h e - ~ e h y e - ~ i i c k e l Theory for the activity coefficients of the ionic species (6)
We h?ve YI- = YIZ-. Also YI, is nearly equal to unity because the low of (12) in the system. Thus
Kz
E
(-1( W) IzW) -
K-
Since the system is in equilibrium with solid iodine, a h =
--
Constant (12). If we take this as x and (13-1 as y, then, concentration of total iodine = (12) (Is-) = C' = (X + y). If the concentration of the potassium iodide taken for the experiment is (I-) (I3-) = c then Kc = x(c - y)ly. This gives by rearrangement,
+
+
"
=. X
(f) + (L). K,+x
variation of llc. This situation is quite satisfactory for a student experiment. The mean value of c'lc, when used in the calculation of K, with the mean value of l/c and with x = 1.00339 X M (experimental) yields a mean value of (xl(K,,+ x)) = 0.39335, whence K, = 1.5475 X M. This very nearly agrees with the value calculated for Kc from the values available a t different temperatures.
Discussion
Literature CHed
The very low magnitude of x implies that (c'lc) should he essentially constant. That this is true is shown by the values in column 4. The very slight upward drift in the magnitude of c'lc, especially noticeable for higher values of c is possibly due to the presence of very low concentrations of other species such as 1s- and 17- in the system. Whether this is the true cause of the variation in c'lc cannot be decided unless a much wider range of concentration is investiin ("Ic) is (' 4% of gated' Fortunately the the mean value even in the extreme cases) for a four-fold
(1) Shamsker, David P.. and Carisnd, Carl w., "Experimenrp in P h @ d Chemistry." 2nd Ed., McGraw-Hill. Kogekushe Ltd., 1967, p. 180; Farrington. Daniek, et al., "Experimental Physical Chemistry." 7th. Ed., McGraw-Hill. Kogakusha Company, 1970, p. 113; Wilson, J. M.. et al.. '"Experiments in Physical Chemistry." 2nd. Ed., Pergamon Presa, 1968, pp. 39and 40: Jamn, Arthur M.."Practical Phyaieal Chemist,try."2nd. Ed.. J end A Churchill Lfd.. London, L961, p. 276. I21 ShormaLu. David P..and Garland, Carl W.. "Erperimonrs in Physical Chemis~y," 2nd. Ed., McGraa-Hill, Kogakuahe Ltd.. 1967, p. 182. 13) Farrington. Daniels, cf d . , "Experimental Physical Chemistry." 7th. Ed., Mffirawu:,,~V-"-,..."L." ,". . " s ~ - - . . ~ u " " . v ~ , rL>.",p. , .*.. 14) Fischcr, R. B., and Peters. Dennis G..l'Quantitative chomicdmyala: W. B. sun. ders Co.. 1968, p. 574. I51 Shoemaker, David P.. and Garland, Carl W.."Experiments in Physical Chemistry." 2nd ~ d .~. e ~ ~ ~ ~~ - ~ ~i ~lt dl . 1,, ~~6 1P., 180.k ~ h ~
.....,
-
Volume 53,Number 5, May 1976 / 327
