Thomas E. Kiovsky
and Richard E. Pincock University of British Columbia Vancouver, Canada
Demonstration of a Reaction in Frozen Aqueous Solutions
T h e revenible reaction of iodide ion with arsenic acid, which historically provided the first complete experimental demonstration of the relationship between reaction rates and equilibria, has recently been presented as a convenient student experiment in kinetics.'
+ 31- + 2HC
HaAs04
hl
HaAsOs
ha
+ Is- + Ha0
The rate of reaction in the forward direct,ion is given by the relation
and the rate of back reaction by
Measurement of both the rate constants, kr and ko, together with the position of equilibrium, allows an experimental verification of the relationship of an equilibrium constant to the rate constants for forward and reverse reaction (i.e., K = k,/kn). This reaction also provides a simple demonstration of the occurrence of reactions in frozen solutions. Shifts in equilibrium positions brought about by frcczing arc also illustrated. The Experiment
A solution of arsenic acid, perchloric acid, and potassium iodide, each a t 0.01 M concentration is boiled a few seconds to remove dissolved air, then divided into several samples (ca. 5 ml each) held in glass ampoules, and the ampoules are evacuated and ~ e a l e d . ~ These colorless samples, when placed in a dry iceacetone bath a t -SO0, change withm two minutes to yellow, then orange, and finally, to a yellow brown color. When thawed, a decidedly yellow solution results. I n about twenty minutes a t room temperature, or in a few minutes if warmed in hot tap water, the yellow color is discharged and the samples are again colorless. The cycle may then be repeated. A frozen sample held a t ca. -10' develops a bright orange color within twenty minutes, but turns grey in about two to three hours. Close inspection of the sample then shows the presence of iodine crystals imbedded in the ice. A pale yellow solution, which soon becomes colorless, results when the sample is thawed.
' BRITTON.D.. AND
HUGUS.Z. Z.. J. CAEM.EDUC.,40, 607 (1963 j. 'A permanent yellow color develops in samples not sufficiently degsssed.
A few small iodine crystals may he seen in the thawed solution; these slowly dissolve and the solution remains colorless. Samples stored at cca. -5 to -lo0, if not jarred, may frequently remain unfrozen. These samples remain colorless, while a yellow color develops in frozen samples held a t the same temperature. Frozen samples containing 0.1 M sodium chloride or ethanol (in addition to the 0.01 M reactants) turn a light yellow color a t -10". The samples containing sodium chloride produce a deep red brown color a t -SO0, but at this temperature only a light yellow color appears in the samples containing0.1 Methanol. Discussion of Results3
The freezing of a dilute aqueous solution of arsenic acid, perchloric acid, and potassium iodide results in crystallization of water, and the solutes are concentrated in the remaining liquid phase. As the rate of reaction in the forward direction is third order, the production of triiodide ion receives a great boost by the increase in concentration brought about by freezing. Although the temperature is lower and the rate constant (k,) decreased, the actual rate of production of 13- is increased by the high concentrations of H+, I-, and H&OI On the other hand, since the concentration of H + and I- appear in the denominator of the reverse rate equation, the rate of the back reaction is decreased by these greater concentrations in the liquid part of a frozen solution. The over-all result is that the reaction shifts to the right and the yellow color of 1sbecomes apparent. The normal equilibrium state, where a t 0' only about 0.02% of the iodide ion is oxidized to iodine, is reestablished on thawing. Another way of looking a t this reversible reaction is that the position of equilibrium is shifted by the decrease in volume brought about by freezing. The following relationship, where m denotes moles of reagent and Vhis the volume of liquid phase, points this out.
According to this equation, if the volume of the reacting system is decreased (i.e., by freezing out the solvent), the number of moles of H&03 and of 1 8 - must increase and the moles of H&Os, H + and I- decrease in order to attain equilibrium. I n addition, the equilihrium may be shifted farther to the right by the precipitation of solid iodine, which is especially insoluble a
Cf,PINCOCK, R. E., AND KIOVSKY, T. E., J. CHEM. EDUC., 43, 358 (1966). Volume 43, Number 7, July 1966
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a t the low temperature of these experiments. Inert solutes like sodium chloride act to increase the volume Vh and prevent the development of sufficient triiodide ion to give a great color change in the frozen solutions. Related to this effect is the fact that use of KI, rather than NaI, and HC104, rather than HCl, is essential for a most dramatic development of color on freezing. As KCIOa is insoluble relative to NaC104, NaCl, or KCI, it is precipitated out at low temperatures. Thus the potassium and the perchlorate ions do not con-
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tribute to the solutes present in the volume Vh, the volume is correspondingly smaller, and the shift in equilibrium farther to the right. Acknowledgment
The authors wish to express appreciation to the
U. S. Air Force Office of Scientific Research (Grant No. 1102-66) and to the National Research Council of Canada for support of this research.
