The standard enthalpy of formation of the aqueous iodide ion

A-004 J. Chem.7lhermodynamics 1977,9, 835-841

The standard enthalpy of formation aqueous iodide ion s,b

of the

GERALD K. JOHNSON Chemical Engineering Division, Argonne National Laboratory, Argonne, Illinois 60439, U.S.A. (Received 7 February 1977) The enthalpy of reaction of K&O,(c) with HI(aq), and the enthalpiesof solution of KBr(c) and Ia(c) in HI(aq) were measuredin an LKB Solution Calorimeter. Combination of the measuredvalues with data from the literature yielded the standard enthalpy of formation of the aqueous iodide ion: AHY(I-, aq, 298.15K) = -(I3572 & 0.017) kcalth mol-‘. This result was combined with a literature value for the enthalpy of solution of HI(g) to derive m:(Hl,g) = (6.333f 0.026) kcalti mol-1 and was also used to rederive the enthalpy of formation of aqueous XeO,, AZQXeOs.96.15Ha0, 298.15K) = (98.73 f 0.22) kcal,, mol-I.

1. Introduction The standard enthalpy of formation of the aqueous iodide ion, AH,“@-, aq, 298.15 K), which is numerically equal to the standard enthalpy of formation of aqueous hydroiodic acid, AH,“(HI, aq, 298.15 K), is an important basic thermodynamic quantity which enters into the enthalpies of formation of many metallic and nonmetallic iodides. Therefore, it is surprising that in spite of its importance, AH,“@-, aq) is not known with high precision. For example, the CODATA selection(‘) gives an uncertainty of +0.20 kcal,mol-’ for AH,“(I-, aq), whereas the uncertainties for AH,“(Cl-, as) and AH,“(Br-, as), are only rtO.021 and kO.035 kcal,,mol-I, respectively.? The CODATA selection”’ for AH,“@‘, as), -(13.60&0.20) kcal,, mol- r, is stated to be based on experimental work from Wu et a1.,(2) Vorob’ev et al.,(‘) and Howard and Skinner.(4’ Wu, Birky, and Heplert2) measured the enthalpies of reaction of Br,(l) and I,(c) with HI(aq) from which AH,“@-, aq) = -(13.69f0.19) kcal,,, mole1 can be derived. Vorob’ev, Broier, and Skuratov(3) measured the enthalpies of reaction of HI(aq) and 12(c) with H202(aq) from which AH,“@-, as) = -(13.45+_0.3) kcal, mol-’ can be derived. Howard and Skimrert4’ measured the enthalpies of reaction of I,(c) and HI(aq) with aqueous hydrazine from which AH,“@-, aq) = -(13.80+0.03) kcal,, mol-’ can be derived. Thus, the CODATA selection is essentially an average of these results with sufficient uncera This work was performed under the auspicesof the U. S. Energy Researchand Development Administration. b Presentedin part at the Fourth International Conference on Chemical Thermodynamics, Montpellier, France, 26 to 30 August 1975. t Throughout this paper caltb = 4.184J.

836

G. K. JOHNSON

tainty to overlap the extremes. Additional support for a value near - 13.60 kcaltb mol-’ for A&(1-, as) comesfrom the recent measurementsof the standard enthalpy of solution of HI(g) by Vanderzee and Gier.(‘) Their vaIue for AH,“,,,(HI, g), -(19.905+0.020) kcal, mol-‘, when combined with AH,“(HI, g) = (6.30+0.19) kcal, mol”(‘) gives AH,“(I-, as) = -(13.60+0.19) kcal,, mol-I. This result is in excellent agreement with the CODATA selection; however, the uncertainty attached to the value is still large. We are concerned about A&(1-, aq) becauseof the dependenceof A&(XeO,, aq) on the value chosen for AH;@-, aq). We had previously measured’@the enthalpy of reaction of XeO,(aq) with HI(aq) and based on Howard and Skinner’s(4r value for AH,“(HI, as) had derived AiY,(Xe03~96.15Hz0) = (99.94+0.24) kc&,, mol-‘. XeO,(aq) has been used in subsequent calorimetric studies(‘p*) because it is an excellent oxidizing agent and because its enthalpy of formation was considered to be well established. However, if we were to adopt the CODATA value for A&(1-, as) then the result for AH&eO,, aq) becomes (98.8fl.2) kcal,, mol-‘. Clearly, the enthalpy of formation of XeO,(aq) is not well established and hence its usefulnessas an oxidizing agent in calorimetric studies is diminished. These considerations caused us to undertake a determination of AH,“@-, aq) based on measurementsof the enthalpy of reaction of KBrO, with HI(aq) according to KBrO,(c) + 6HI(aq) = KBr *31, *3H,O(aq). (1) Measurements of the enthalpies of solution of KBr(c) and I,(c) into HI(aq) were also performed. Combination of the experimental measurementswith AH,“(H,O, I)(‘) and the enthalpy of decomposition of KBr03(c) into KBr(c) and O,(g), a quantity recently measured in this laboratory,(g) yields AH,“(I-, as>.

2. Experimental

MATERIALS The KBrO, sample was Merck Reagent-Grade material, which had been twice recrystallized from distilled water and dried under vacuum for five days at 400 K. The material was stored in a helium atmosphere glovebox and all subsequenthandling was done in this glovebox. The oxidizing power of the KBrO, was assayed by allowing it to react directly with iodide in dilute acid in the presence of a molybdate catalyst.(lo) The iodine formed was titrated with Na&O, solution that had been standardized against Mallinckrodt Primary Standard grade KI03 [manufacturer’s assay (100.00+0.05) moles per cent]. The oxidizing power of the KBrO, sample was equivalent to an assay of (lOO.Of0.05) moles per cent KBrO,. Spectrographic analysis indicated a mass fraction less than 30 x 10m6of metallic impurities, The KBr sample was purchased from Alfa-Ventron (Beverly, Mass.). This material was dried for 3 days at 400 K and transferred to the helium-atmosphere glovebox. Metallic impurities in this sample totalled less than 15 x low6 mass fraction. The iodine sample was part of a batch of high-purity crystalline material previously used at this laboratory for determining the enthalpies of formation of IFS and IF,(ll) as well as the enthalpy of solution of IZ(c) in HI(aq).@) Mass fractions of impurities in

ENTHALPY OF FORMATION OF AQUEOUS IODIDE ION

837

this sample, as previously determined,(“) were (x 10m6):Al, 1; B, 1; Cr, 2; Fe, 5; Ti, 10; Cl, 6; P, 2; S, 10; C, 26; N, 10; and 0, 2. This material was stored in the helium-atmosphere glovebox and ground in an agate mortar before use. Stock solutions of concentrated hydroiodic acid were obtained from two different sources. One stock solution was prepared by distillation, under argon, of commercial hydroiodic acid which contained approximately 1 mass per cent of H,PO, as a preservative. The other stock solution was prepared by distillation, under argon and in the presence of red phosphorous, of a solution prepared by dissolving HI(g) into distilled water. The stock solutions were then diluted with distilled water to concentrations of 0.20 mol dm-’ and 0.15 mol dm- 3 for the calorimetric experiments. All solutions were stored under argon in brown bottles to inhibit the decomposition into I1. Whenever a calorimetric solution began to turn brown, indicating the presence of I,, it was discarded and a fresh solution was prepared. CALORIMETRIC

SYSTEM AND PROCEDURES

The calorimetric measurements were carried out in a LKB-8700 Precision Calorimetric System. The 0.1 dm3 glass reaction vessel was motied so that temperature measurements could be performed with a quartz-crystal thermometer (HewlettPackard Model 2801-A). The gold stirrer-ampoule holder assembly was replaced with one constructed of Kel-F plastic. For the exothermic reactions of KBrO, with HI(aq), electrical calibration experiments were conducted before and after each KBrO, experiment, and the average value was used for s(calor). For the endothermic dissolutions of KBr and I,, electrical energy was added to the calorimeter in an amount approximately equal to the endothermic effect of dissolution. Thus, for these experiments the corrected temperature change of the system was very small. The mean temperature of the calorimetric system for all experiments conducted in this study was (298.15+0.01) K. The samples of KBrO,, KBr, and I, were loaded in the glovebox into weighed 1 cm3 glass ampoules equipped with glass fJling stems. The filled ampoules were reweighed, stoppered, removed from the glovebox, and promptly sealed with a flame. The rate of reaction of the KBr03 was initially quite fast, but then slowed down as the reaction proceeded. Approximately 20 min was required to reach the final drift rate. A molybdate catalyst(r’) was used in several trial experiments; however, it did not have a significant effect on the overall reaction time. The I,-in-HI experiments had reaction times which varied between 30 and 40 min. 3. Results Calorimetric results for the reaction of KBr03(c) with HI(aq) and for the dissolution of KBr(c) and I,(c) in HI(aq) are given in tables 1, 2, and 3, respectively. In the tables: m is the mass of sample; is the mean energy equivalent of the calorimetric system based on the electrical calibrations; A8, is the corrected temperature change of the calorimeter; AH,,, is the correction for saturating the helium in the free volume of the ampoule with water vapor; AHelectis the electrical energy

G. K. JOHNSON

838

added to the calorimeter to approximately offset the endothermic effects of the dissolution of KBr and Iz; and AH/M is the specific enthalpy of reaction or solution of sample. me AHva, corrections are based on the enthalpy of vaporization of water, 10.520 kcal,, mol- r ,(r) densities of 3.27, 2.75, and 4.93 g cmm3 for KBrO,, KBr and Iz, respectively, and the assumption that each ampoule had an internal volume of I cm3. The thermal effects of impurities in the samples were judged to be insignificant. TABLE 1. Results of the reaction of KBrOs(c) with HI(aq) a (calth = 4.184 J)

mUCBrO3) g 0.14017 0.13981 0.13967 0.13994 0.14124 0.14026 0.14009 0.14008



L\eo

calt,, K-l

K

104.979 105.066 105.012 104.952 105.053 105.003 105.022 104.942

( - A&)
0.98896

0.98633 0.98642 0.98813 0.99604 0.98940 0.98698 0.98853

AH,,,

calth -0.013 -0.013 -0.013 -0.013 -0.013 -0.013 -0.013 -0.013

--LWIM Cal,, g-l -740.76 -741.31 -741.74 -741.17 - 740.94 -740.79 - 740.01 -741.08

(AH/M) = -(740.98 f 0.18) Cal,, g-l b AH = -(123.744 f 0.067) kcalth mol-1 =ed a In each experiment, 99.41 cm3 of 0.20 mol dme3 HI solution was used. b Uncertainty is the standard deviation of the mean. c Uncertainty is the uncertainty interval. d The molar mass of KBrO3 was taken to be 167.0002 g mol-I. TABLE 2. Results of the dissolution of KBr(c) in HI(aq) D (cab = 4.184 J) m(KBr) g

Aec K

(-A&)
0.09963 0.09854 0.09961 0.09960 0.09970

-0.00038 0.00006 -0.00042 - 0.00020 -0.00022

0.040 -0.006 0.044 0.021 0.023

b

AHeleCt _I C&h

AH,,, calm

WM calth g- i

4.069 4.069 4.069 4.069 4.069

-0.013 -0.013 -0.013 -0.013 -0.013

41.11 41.10 41.16 40.93 40.91

= (41.04 rt: 0.05) cal, g-l 0 AH = (4884 f 15) calth mol-1 d.e



a Iu each experiment, 99.41 cm3 of 0.15 mol dmd3 HI solution was used. b A value of 105.0 calth K-l was used for in these experiments. c Uncertainty is the standard deviation of the mean. d Uncertainty is the uncertainty interval. 8 The molar mass of KBr was taken to be 119.002 g mol-l.

ENTHALPY

OF FORMATION

OF AQUEOUS IODIDE

839

ION

TABLE 3. Results of the dissolution of 4(c) in HI(aq) ’ (Cal,, = 4.184 J)

A@,

mU2)

ii-

g 0.63474 0.63355 0.63317 0.64013 0.63589 0.63794 0.63673 0.63697

(-A8,)(s(caior)) c&

-0.00038 -0.ooo50 o.ooo39 0.00009 -0.00021 -0.00036 0.00005 -0.ooo29

0.040 0.052 -0.041 -0.009 0.022 0.038 -0.005 0.030

b

AH,l,,t C&b 3.351 3.351 3.351 3.351 3.351 3.351 3.351 3.351

A&w

calth -0.012 -0.012 -0.012 -0.012 -0.012 -0.012 -0.012 -0.012

AH/M

caltng- 1 5.323 5.352 5.209 5.202 5.286 5.294 5.236 5.289

= (5.274 r!z 0.019) calCbg-’ c AH = (1339 * 10) caltn moI-ld.L, a In each experiment, 99.41 cm3 of a solution, which was (0.15 mol dmm3 HI + 0.0084 mol dm-s KBr) was used. b A value of 105.0 calth K-l was used for
TABLE 4. Reaction scheme for the derivation of tiF(HI, (c&h = 4.184 J)

aq) at 298.15 K a

= KBrOB(c) + 23.692(HI*275.05H,O); AH = (123.744 f 0.067) kc& molB. KBr(c) + 17.692(HI*368.50H20) = KBr .17.692HI .6519.485H~O; AH = (4.884 i 0.015) kcalth molC. 3&(c) + KBr *17.692HIs6519.485Hz0 = 31z*KBr+17.692HI*6519.485H,O; AH = (4.017 3~ 0.030) kcalth mol-1 D. 17.692(HI*275.05HZO) + 1653.317H20(c) = 17.692(HI.368.5OHaO; AH = -(0.212 f 0.002) kcalth molE. KBrOa(c) = KBr(c) + SO,(g); AH = -(8.10 ZIZ0.06) kcal, mol-’ F. 3Ha(g) + gOa = 3HaO(1); AH = -(204.945 ZIZ0.030) kcab mol-’ G. 6(HI .275.05Ha0) + 03Hz0(1) = 6(HI. coHzO); AH = -(0.819 f 0.008) kcal, mol-l AH = -(81.431 f 0.101) k&,, mol-1 H. X-b(g) + 3W) + d%OO) = 6(HI-a%O);

A.

31z.KBr.17.692HI.6519.485Ha0

o All species are aqueous unless designated otherwise.

The thermochemical cycle used to derive AH,“(HI, aq) = AH,“@-, aq) is given in table 4. Combination of the three experimental measurements,reactions A, B, and C, results in the removal of all of the product mixtures from the cycle. In table 4, the enthalpies of reactions A, B, and C were taken from tables 1, 2, and 3, respectively. The enthalpies of dilution of HI(aq), required in reactions D and G, were taken from the data of Vanderzee and Gier. w The enthalpy of decomposition of KBrO,, reaction E, was previously measured at this laboratory,(g’ and the enthalpy of for-

840

G. K. JOHNSON

mation of H,O reaction F, was taken from the CODATA selection.@)The standard enthalpy of formation of HI(aq) is one-sixth of reaction H or -(13.572&0.017) kcal,, mol - r ,[-(56.785kO.071) kJ mol-‘1.

4. Discussion The value obtained for AH,“(I-, aq) in this study is in good agreement with, but more precise than the CODATA selected value, -(13.60+0.20) kcal,, mol-r.(l) Our value is also, within the combined uncertainties, in agreement with the results of Wu, Birky, and Hepler, - (13.69+0.19) kcal,,mol-1,(2) and Vorob’ev, Broier, and Skuratov, -(13.45&0.3) kcal,, mol-1,(3) and in disagreement only with the result of Howard and Skinner, - (13.80+ 0.03) kcal, mol-1.(4) The weight of evidence suggests that Howard and Skinner’s value is too negative although it is difficult to explain how this could have occurred. In the samestudy, using the sametechnique, they derived a value for AE&(HBr, aq) which is in excellent agreement with the accepted value.(‘) In the introduction the CODATA selected value for AH,“(HI, g) was combined with Vanderzee and Gier’@ value for AH,“,,,,(HI, g) to derive a value for AH,“(I-, aq). The result, however, had an excessivelylarge uncertainty due to the large uncertainty assigned to AH,“(HI, g). It is possible to reverse this calculation and derive a value for AH,“(HI, g) based on our AH,“(I-, aq) value and AE&,,,(HI, g).(5) The result obtained, (6.333L-0.026) kcal,, mol- r, is in good agreement with the CODATA selection, (6.30f0.19) kcal, mol -I, but has a much lower uncertainty. The derived result is also in agreementwith A&(HI, g) values which can be derived from the equilibrium studies of Bodenstein,(r2),(6.43f 0.13) kcal,, mol- ‘, Rittenberg and Urey,(r3) (6.38kO.07) kcal,,, mol-‘, Taylor and Crist,(14) (6.27f0.01) kcal,, mol-‘, and Bright and Hagarty, (15) (6.35kO.07) kcal,,, mol-‘, as well as the spectroscopic result of Murphy,(‘@ (6.34kO.02) kcal,, mol- ‘. Thus the internally consistent data set, AH,“(HI, g, 298.15 K) = (6.333kO.026) kcal,, mol-r, A&,,,(HI, g, 298.15 K) = -(19.905~0.020)kca1,,mo1-‘,andAH,”(H1,aq,298.15K) = -(13.572+0.017) kcal,, mol-’ is recommended for HI. The AEJF(HI, as) value derived herein differs significantly from the value, -(13.79 AO.03) kcal,,, mol-‘,@I used previously by us to derive A.&(XeO, *96.15H20).(@ Thus, we take this opportunity to recalculate and update our result forAZ&(XeO,,aq). In addition to the change in AH,“(HI, aq), we have used the enthalpy of dilution data for HI(aq) recently reported by Vanderzee and Gier,(‘) and we have corrected the enthalpy of solution of I, in HI(aq) previously published, for a small bias in computing AtI,. The recalculated value is (1.213+0.027) kcal,, mol-‘, which differs somewhat from the value measured in the present study at twice the iodine concentration, (1.339&O.OlO) kcal, mol-‘. Several experiments were conducted during the present investigation at the iodine concentration used previously to recheck that result. The results were in good agreement, which establishes that the enthalpy of solution of I, in HI(aq) is strongly concentration dependent. Incorporation of these revisions into the Hess cycle(6) used to derive AHr(XeO,, aq) yields Aff,(Xe0,*96.15H20, 298.15 K) = (98.73f0.22) kcal,, mol-‘.

ENTHALPY

OF FORMATION

OF AQUEOUS IODIDE

ION

841

REFERENCES

1. CODATA RecommendedKey Values for Thermodynamics,1975.CODATA Bulletin 17 1976. 2. Wu, C.; Birky, M. M.; Hepler, L. G. J. Phys. C/rem.1963,67,1202. Vorob’ev, A. F.; Broier, A. F.; Skuratov, S. M. Dokl. Akad. Nauk SSSR 1966,173, 385. 2 Howard, P. B.; Skinner, H. A. J. Chem. Sot. (A) 1966,1536. Vanderzee,C. E.; Gier, L. J. J. Chem. Thermodynamics 1974,6,441. 2: O’Hare, P. A. G.; Johnson,G. K.; Appehnan, E. H. Znorg.Chem. 1970,9,332. 7. O’Hare, P. A. 0.; Hoekstra, H. R. J. Chem. Thermodynamics 1974,6, 965. 8. O’Hare, P. A. G.; Shinn, W. A. ; Mrazek, F. C.; Martin, A. E. J. Chem. Thermodynamics 1972, 4,401. 9. Johnson,G. K.; Smith, P. N.; Appehnan, E. H.; Hubbard, W. N. Irwrg. Chem. 1970,9,119. 10. Kolthoff, I. M.; Belcher, R.; Stenger,V. A.; Matsuyama,G. Volumefric Analysis. Interscience: New York. 1962, p. 269. 11. Settle, J. L.; Jeffes,J. H. E.; O’Hare, P. A. G.; Hubbard, W. N. J. Znorg.Nucl. Chem. 1976, Supplement,135. 12. Bodenstein,M. Z. Phys. Chem. 1899,29,295. 13. Rittenberg, D.; Urey, H. C. J. Am. Chem. Sot. 1935,56,1885. 14. Taylor, A. H.; Crist, K. H. J. Am. Chem. Sot. 1941, 63, 1377. 15. Bright, N. F. H.; Hagerty, R. P. 7bn.s. Faraday Sot. 1947,43, 697. 16. Murphy, G. M. J. Chem. Phys. 1936,4,344.