J. H. STERN AND A. A. PASSCHIER
752
Vol. 66
THE HEATS OF FORMATION OF TRIIODIDE AND IODATE IONS BY J. H. STERN AND A. A. PASSCHIER Department of Chemistry, Long Beach State College, Long Beach 4, California Received November BO, 1961
The heat of reaction of solid iodine with excess aqueous potassium iodide (893 f 12 cal./mole) was measured microcalorimetrically at 25’. This value combined with other data yields -12.72 f 0.15 kcal./mole for the heat of formation and 56.0 e x . for the standard entropy of 18- (aq). Measurements of the heats of reduction a t 25” of separate solutions of potassium acid iodate and iodic acid with aqueous hydriodic acid to form triiodide ion were combined with the newly determined AHP of 13- and other pertinent heats of formation and dilution to yield AHfO of IO3- (as) = -54.8 ==! 0.5 kcal:/ mole. It is shown that the values of the heats of formation of triiodide and iodate ions are internally self-consistent within the limits of experimental error.
I. Introduction Different measurements of the temperature coefficient of the equilibrium constant for Idaq) -t I13yield values of the heat of complexing of aqueous iodine with iodide ion from -3.87 to -5.10 kcal./ mo1e.l These heats combined with the heats of formation of aqueous iodine‘ and iodide ion2 result in values ranging from - 11.7 to - 12.9 kcal./ mole for the heat of formation of triiodide ion. The Nationad Bureau of Standards2 lists -12.4 kcal./mole and a standard entropy of 41.5 e.u., which is not compatible with any of the above values when combined with pertinent thermodynamic data. Hence there is a considerable uncertainty associated with the heat of formation of triiodide ion and the heats of formation of other substances which are related to it. The National Bureau of Standards2 heat of formation of IO6- (as) appears to bo calculated from data that are between nearly fifty and more than eighty years old. Heats of formation of other halogen compounds and ions based on results taken from the older literature or from indirect measurements sometimes have been found t o be in conflict with more recent va1ues.l Thus direct and related calorimetric determinations a t 25’ of the heats of formation of these two important oxidation states of iodine were undertaken. That of triiodide ion was determined from the heat of reaction Iz(s)
+ I- (KI in excess) = Ia-
The heat of formation depends only on the measured heat and the well established heats of formation and dilution of potassium iodide solutions. The entropy was calculated from the results of these measurements and appropriate thermodynamic data. The heat of formation of iodate ion was determined from the heat of reduction with excess aqueous hydriodic acid combined with the newly determined 103-
+ 8HI = 318- + 3HzO + 2HC
heat of formation of I&- and other data. The stoichiometry of this reaction is well established. (1) Summarized by H. C. hlel, W. L. Jolly, and W. fix. Latimer, J . Am. (Ihsm. Sac., 76, 3827 (1953). (2) National Bureau of Standards, Circular 600, “Selected Values of Chemical Thermodynamio Properties,” Washington, D. C., 1952, Reprint of July, 1961. (3) I. PI. Kolthoff and E. B. Sandell, “Textbook of Quantitative Inorganio Analysis.” 3d Edition, The Macmillen Co., New York, N. Y., p. 694.
Solutions of iodic acid and potassium acid iodate in differing concentrations were used as separate sources of iodate ion in two different sets of measurements. Appropriate heat of dilution corrections also were determined. 11. Experimental A. Determination of AHfO of Is- (aq).-Materials.Iodine and potassium iodide were both reagent grade. All solutions of potassium iodide were freshly prepared with triply distilled water. Microcalorimeter and Experimental Procedure.-The measured heats were less than one calorie per run and were determined with a twin heat-leak microcalorimeter . 4 4 In this apparatus the heat of the reaction taking place within one of the calorimeters is exchanged with the bath. The heat exchange follows Newton’s law of conduction and
AH = K
1;
ATdt
where A H is the total heat exchanged between times ti and t z , K is the heat transfer constant, and AT’ is the temperature differencebetween thermostat and reaction calorimeter. AT is equal to zero before and after the reaction period. The integration is performed graphically by means of a recorder trace of A T vs. t , so that AH = K A A where AA is the area under the traced curve determined by planimeter and K by electrical calibration. The assumptions have been discussed in detail by Borchardt and Daniels . 4 The microcalorimeter consists of two identical glass vessels with flat bottoms, 35 mm. in diameter with a capacity of 80 ml. One vessel contained both reactants, the other potassium iodide solution only. Both vessels are capped by tight-fitting standard taper 40/5P Teflon covers 3 cm. thick. Passing through each cover is a thin thermistor probe, an off-center propeller-stirrer shaft, and a ramrod for breaking ampoules, all made of glass. The reaction was initiated by crushing a submerged ampoule containing iodine against the bottom of the reaction vessel. No heat effect was observed as a result of breaking an empty ampoule. A thyratron speed-controlled motor drives separate identical pulleys so that the contents of both calorimeters are stirred a t exactly equal rates. Temperature chan es as a function of time are plotted on a 10 mv. recorder &peedomax G), utilizing the amplified (Leeds & Xorthrup microvolt amplifier 9835-A) off-balance potential of a Wheatstone bridge network. The bridge consists of two variable five dial resistance decades set at 2 kilohms, and a matched thermistor (Fenwal G170, 2 kilohms each) in each of the two calorimeter vessels. It is powered by one mercury battery (Mallory RM-42R, 1.34 v.). The circuit is arranged so that the bridge remains balanced when temperature changes occur that are due to causes other than the reactlon itself. The recorder trace thus represents the temperature difference between the reactant system in one calorimeter vessel and the inert solution in the other as a function of time. (4) 15. J. Borchardt and F. Daniels, J. Am. Chem. floc., 79, 41 (1957). ( 5 ) F. D. Rossini, ed., ”Experimental Thermochemistry,” Ch. 12, E. Calvet, “Microcalorimetry of Slow Phenomena,” Interscience Publishers, Inc., New York, N. Y., 1956.
HEATSOF
April, 1962
FORMAnOIU OF
The calorimeter vessels are immersed to the level of the Teflon caps in a highly insulated enclosed 20-1. water-bath. The temperature is re ulated by a proportional controller (Electron-0-therm 148fT) and low-lag knife heater balanced m t h a submerged refrigeration coil. All runs were made at 25 000 f 0.0015’ as determined by a platinum resistance thermometer &ndMueller bridge. A heater coil (Manganin, 21.93 ohms) immersed in Octoil and contained in ti thin glass tube is used to determine the heat transfer constant K . The energy for the heater is furnished by two mercury batteries in series. The current is determined by measuring the potential drop across a standard 10-ohm resistance in series with the heater by means of a millivolt potentiometer (Leeds & Northrup 8691). The time of heating is given by an electric stopclock synchronized with the heater switch. Reactions were carried out in an excess of 2.0 M potassium iodide to ensure complete and reasonably rapid conversion to triiodide ion. In a run 50 ml. of potassium iodide was placed in each calorimeter vessel. Thermal equilibrium was attained in less than 5 min. after breaking the submerged ampoule containing a weighed quantity of iodine. The base line was unchanged before and after all runs or calibrations, with AT estimated to be less than 0.02’. B. Determination of AHP of IQa-(aq).-Materials.Primary standard KH(IO& (Hach Ver) and reagent .grade HIOa (B. and A . ) were heated for several hours prior to weighing and were dissolved in distilled water. Aqueous solutions of H I were freshly prepared and analyzed by methods discussed Calorimeter and Experimental Procedure .-The calorimeter has been described elsewhere.6 The reductions were carried out in an excess of aqueous hydriodic acid (200250 ml.). The iodate solution was weighed into a thin glass ampoule which then was submerged in the reaction solution. After crushing the ampoule, the reaction proceeded very rapidly and thermal equilibrium was attained in less than 60 see. Heat of dilution corrections were measured similarly. Observed temperature changes %*ere corrected for heat exchange7and electrical calibrations were made after each run in the range 26 f 1’.
111. Results All heats of reaction and dilution are given in the Tables I-IV. Precision indices associated with all average heats are standard deviations. Heat of Reaction of Solid Iodine with Aqueous Iodide Ion.-The details are given in Table I. TABLE I THEHEATOF REACTION AH1 Moles I~ x 103
0.9949 1.128 1.047 Av.
MICI
OF
12(s)WITH I-
Total energy change, cal.
AHi, cal./mole Is-
2.00 0.904 2 00 ,996 2.00 .934 A H I = 893 f 12 cal./mole It-
907 883 888
Heat of Reduction of Iodate Ion (KH(I03)2)with Hydriodic Acid in Aqueous Solution.-The details are given in Table 11. Heat of Reduction of Iodate Ion (HI03) with Hydriodic Acid in Aqueous Solution.-The details are given in Table 111. Heat of Dilution Corrections. -The appropriate heat of dilution corrections of iodic acid were determined in order t o correct, the calorimetric reaction results of Table I11 to standard states. The heat of dilution of potassium acid iodate was (6) J. H. Stern and A. A. Passchier, J . Chern. Eng. Data, 7 , 73 (1962). (7) .4. Weissberger, “Physical hlethods of Organic Chemistry,” Val. I, Third Edition, Interscience Publishers, Ino., New York, N. Y., 1959, p. 538.
TEIIODIDE AND
753
IODATE IONS
TABLR I1 THE HEAT OF ‘REDUCTION AHz OF KH(1Os)z Moles KH(I0a)n x 104
Moles HI
MH1
*ATcor. 0.0010
Total energy change, cal.
WITH
HI
-AH$, koa1 mole
Id*-
77.78 84.20 1.35 0.361 83.60 83.55 1.35 .385 80.14 83.45 1.35 ,375 AH2 = -83.73 It 0.44 kcal./mole 1 0 3 -
4.620 5.002 4.801 Av.
0.270 .270 .270
TABLE I11 THEHEATOB REDUCTION AH8 OF HIOa WITH H I hfoles HIOa
T cor.
Moles HI
Total energy change, oal.
-AHa, kcal zole
da
1 x 103 MHI ztO.O0lo 252.7 84.78 1.26 1.175 2.980 0.252 250.4 84.44 1.26 1.163 2.965 ,252 248.0 84.21 1.167 ,308 1.54 2.945 253.5 85.22 1.54 1.166 2.975 .308 Av. AHs = -84.66 f 0.45 kcal./mole IOa-
TABLE IV THEHEATOF DILUTIONAHc OF HIOd M~iosinitial
0.2536
MHIOS final X 106
No. of runs
hv. AH1 koal./mole HI08
5.00
2
-0.45 f 0.05
found to be negligible. Table IV lists summarized values. IV. Interpretation of Data Standard Heat of Formation and Entropy of 1%-(aq) .-Unless otherwise specified, all calculations were made using thermodynamic properties from the National Bureau of Standards Tab1ea.l The average AH1 (893 It 12 cal./mole) from Table I combined with appropriate heats of formation and dilution corrections yields AH* of 1 3 (0.02 &I) = -12.72 kcal./mole. Since the concentration of triiodide ion was very low, we assume this result also to be equal to AHtO of 13- (as) = - 12.72 kcal./mole, with an estimated over-all error of f0.15 kcal./mole. Combining this result with pertinent thermodynamic data we calculate Soof 13- (as) = 56.0 e.u. Standard Heat of Formation of Iodate 1on.Combining the average AH2 from Table I1 (KH(103)2)with the newly determined heat of formation of 1 3 - and other heats, we calculate AH? of IO3- (as) = -55.07 kcal./mole. Similarly the average AH2 from Table I11 combined with AH4 from Table IV (HIOJ, yields AHfo of IO3- (as) = -54.59 kcal./mole. We shall take our final value as the average of the above two, AHfo = -54.8 rt 0.5 kcal./mole. This value is in good agreement with those given by the National Bureau of Standards of -55.0 and -54.9 kcal./mole for 1 0 3 (as) and HIOs(aq) , respectively. This agreement shows that our value of AH+‘ of IS- (as) is internally self-consistent with experimental heats reported here and auxiliary heats used in these calculations. Acknowledgment.-Financial support of this work by the Research Corporation is gratefully acknowledged.
