THERMODYNAMIC PROPERTIES OF IODINE MONOCHLORIDE
2443
The Entropy of Iodine Monochloride. Heat Capacity from 17 to 322°K. Vapor Pressure. Heats of Fusion and Vaporization'
by G. V. Calder and W. F. Giauque Low Temperature Laboratory, Departments of Chemistry and Chemical Engineering, University of California, Berkeley, California
(Received March 1, 1965)
The heat capacity of IC1 has been measured from 17 to 322°K. The melting point and heat of fusion were found to be 300.53" and 2773 f 2 cal./mole, respectively. The entropy of ICl(s) = 23.41 gibbs/mole, and its free energy of formation from the elements is A F " = -3363 f 25 cal./mole, each at 298.15"K. Values of C,", So, (F" - H o 0 ) / T and , (H" H o o ) / Thave been tabulated. The total vapor pressure and the partial pressures of IC1 and C12 have been measured over stoichiometric liquid IC1, and the data have been used to calculate the partial pressure of IC1 over the solid and liquid from 250 to 330°K. AH", = 13,273 cal./mole for the vaporization process. A simple and accurate experimental method has been described for determining the equilibrium composition of the vapor over solutions with more than one volatile component. Equilibrium is preserved by extremely slow sampling through a capillary, with condensation in a liquid nitrogen cooled, evacuated collection vessel. The data of AlcRlorris and Yost for the reaction 90 IC1 = SOCl '/212(g) have been reconsidered and give AHOO = -5375 f 15 cal./mole and AF0z98 150K = -1216 f 15 cal./mole. Also the data of Beeson and Yost for the reaction NOCl = KO l/zClzhave been re-examined and give AH", = 8630 f 10 cal./mole and AF0z98.150R = 4897 =t10 cal./mole. Corrosion tests run on Hastelloy C alloy in liquid IC1 showed essentially complete resistance a t ordinary temperatures.
+
+
+
The thermodynamic properties of the iodine-chlorine system are not very accurately known. The situation is complicated by the presence of Clz and 1 2 gases in equilibrium with the several solid phases including IC1 and IC1, as was first shown by the work of Stortenb e k e ~ - . ~ ,Stortenbeker's ~ substantial investigations covered the freezing point curves and eutectic data, with some approximate values of vapor composition. A summary of the more important references to and properties of IC1 is given by Greenwood4 in a review of the properties of interhalogen compounds. The present work contributes the entropies of solid and liquid iodine monochloride, the partial pressure of IC1 over stoichiometric 1C1, and additional information which will help in the ultimate evaluation of the therniodynainics of this system. Iodine monochloride has two crystalline forms designated Q and p with melting points a t about 27.3 and 13.g0,respectively. The unstable is not very difficult
to obtain temporarily, in small amount, and it was observed during the course of preparatory work. We had hoped to obtain the heat capacity of both forms to supply one of the few cases that illustrate the approach to zero entropy a t OOK. by each of the crystalline forms. However, in this we were unsuccessful, and it appears that either the unstable p form cannot be supercooled very far, or a t least the chance of avoiding one or more centers of spontaneous recrystallization is very small in an amount of the order of 100 cm. ,. The X-ray structural determination^^^^ show that (1) This work was supported in part by t h e National Science Foundation. (2) W. Stortenbeker, Rec. trav. chim., 7, 152 (1888). (3) W.Stortenbeker, Z.physik. Chem., 3, 11 (1889); 10, 183 (1892). (4) N. N. Greenwood, Rev. Pure A p p l . Chem.. 1, 84 (1951). ( 5 ) H. K. Boswijk, J. van der Heide, A. Vos, and E. H . Wiebenga, Acta Cryst.. 9,274 (1956), for t h e form. (I
(6) G . B. Carpenter and p form.
S.M. Richards, ibid., 15, 360
(1962), for t h e
Volume 69, Number 7
J u l y 1966
2444
G. V. CALDERAND W. F. GIAUQUE
both forms consist of similarly ordered chain arrangements that should approach zero entropy at O'K. The staggered chains have the atomic order C1-I-IC1-1-1 with C1 branches on alternate iodine atoms in each strurture, so that the conversion of p to Q would appear to he a rather simple process, unfavorable to supercooling. A p p a r a t z , ~ . Most of the calorimetric apparatus a a s that uqed by Giauque and Egan,' including the standard copper-constantan thermocouple W-25, except that the gold calorimeter could not be used with corro$ive iodine monochloride. Corrosion measurements were made 011 a sheet of Hastelloy Type C alloy,* which is largely nickel, containing 110 (15.80), Cr (15.58), Fe (6.15), JT (3.75), Co (1.61), Mn (0.53), Si (0.52), Y (0.23), C (0.05), P (0.01), and S (0.004), where values in parentheses are percentages. The was sheet, which had a surface area of 23.1 cleaned and ininiersed in IC1 a t 28 to 30' for various times. As f.xpected, the first immersion showed a very small loss in weight followed by essentially complete resistance to IC1. The sheet was cleaned with pure acetone and dried in an oven a t 110' for 10 min. before each wighing. The results are given in Table I.
Table I : Corrosion Teat of Hastelloy C with IC1 Time, h r .
at.,g.
0 2 8 24 242
7.5596 7.5589 7.5590 7.5589 7.5587
After the calorimeter was constructed it was treated with liquid IC1 for 34 days, followed by washings with aqueous K I solution and distilled water before use. d 1 : 3 mixture of Kel-F 200 wax with Kel-F 210 grease9 was used in stopcocks and ground-glass joints. A 5.4123-g. sample of Kel-F 200 increased to 5.4189 g. after inimersion in IC1 at 30' for 3 hr. The small increase in weight was taken to indicate a slight solubility of IC1 TI the Kel-F which would have little effect considering the limited exposed surface in groundjoint use. The cylindrical calorimeter had a 4.0-cm. 0.d. and an 11.0-cni. wall 0.1 em. thick. It was 11.6 cm. in overall length owing to the rounded ends. There were no v a i i e ~in thti calorimeter for heat conduction since Hastelloy ha? a poor thermal conductivity. The effectiveness of the calorimeter would have been conqiderably ini1)roved by the addition of a soldered The Journal of Physical Chemistry
copper tube encasing the cylindrical walls to provide better heat distribution. The poor heat conductivity made thermal equilibrium periods two or three times as long as those in a gold or copper calorimeter. The . Hastelloy volume of the calorimeter was 144 ~ 1 1 1 . ~ A C tube of 0.254-cm. o.d., 0.17-cm. i.d., and 165-cm. length was attached to the calorimeter. All fabrication was by heliarc welding. A Hastelloy fitting at the top of the tube matched standard taper groundglass fittings to a vacuum system and to a glass weight buret for introducing the sample. The weight buret also had a standard taper fitting to a long Hastelloy capillary tube, 0.159-cm. 0.d. and 0.129-cm. i d . , which could be inserted into the upper end of the calorimeter. After the sample had been transferred into the calorimeter, the inlet capillary was removed and washed with CCl, to enable quantitative analysis of the small amount of IC1 which had adhered. The high-resistance thermometer-heater was wound from S o . 40 Au wire, which contained 0.1% Ag, in a manner similar to that described by Murch and Giauque,lo except that type BT 3477 Formvar" varnish was used. At turn 280 from the bottom a single turn of No. 30 double silk covered copper wire was laid down and soldered into the continuous 503 turns of the thermometer, The total thermometer resistance was 287.3 ohms a t 0'. The previously standardized thermocouple, W-25, was compared to the triple (13.95'1L) and boiling (20.37'K.) points of hydrogen and the triple (63.14'K.) and boiling (77.33'K.) points of nitrogen. I t was found to have values 0.10 and 0.12' high a t the triple and boiling points, respectively, for hydrogen, and agreed exactly a t the triple and boiling points of nitrogen. Appropriate corrections were applied. The temperature O'C. was taken as 273.15'K. and 1 defined calorie mas taken as 4.1840 absolute joules. Preparation of Iodine Monochloride. Iodine monochloride was made by direct combination of Jlallinckrodt iodine with chlorine. The iodine was stated to contain only 0.005% of C1 Br and 0.01% of nonvolatile material. The commercial grade 99.5% chlorine was purified by distillation in a low-temperature, vacuum-jacketed, silvered fractionating column, 180 cm. long, with a reflux ratio of 130: 1 maintained by means of a capillary take-off tube and condensation in liquid nitrogen. The column was packed with glass
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(7) W. F. Giauque and C. J. Egan, J . Chem. Phys., 5 , 45 (1937). (8) Union Carbide Stellite Co., Kokomo, I n d . (9) Minnesota Mining and Manufacturing Co., St. Paul, Minn. (10) L. E. Murch and W. F. Giauque, J . Phys Chem., 6 6 , 2052 (1962). (11) Westinghouse Inc., Benolite Dept., Manor, Pa.
2445
THERMODYNAMIC PROPERTIES OF IODINE MONOCHLORIDE
helices. The column condenser was cooled to about 195'K., with solid carbon dioxide in methylene chloride, which maintained a chlorine pressure of about 65 torr. The direct combination was accomplished in a bath a t 195'K. as recommended in Inorganic Syntheses.l 2 A weighed amount of iodine was cooled to 195'K. and evacuated to torr. A small excess of Clz was added, and the excess was evaluated by a second weighing and matched by the final addition of the stoichiometric amount of iodine. An IC1 sample weighing 1352 g. was made. The sample was purified by five successive fractional recrystallizations in a closed system. The amount was reduced to 600 g. as the temperature rose from 24.66' to near the melting point. Melting point measurements made later in the calorimeter indicated a liquid-soluble, solid-insoluble impurity of 0.010 mole yo. Analysis for the C1:I mole ratio was performed by weighing approximately 2-g. samples of solid IC1 in weighing bottles and then converting completely to the silver halide for reweighing. Conversion to the sodium halides was a(-coniplished by immersing the weighing bottles in a filtered dilute solution of n'aBH4, kept at 0' to control the reaction rate with the solid IC1. Additional T\'aBH4was added slowly as needed to complete the reaction. A blank test showed that no halide was present in the sodium borohydride. Four samples gave the C1:I mole ratio as 1.0014, 0.9985, 0.9997, and 0.9997, average 0.9998 f 0.0012. It would have been more accurate to convert the IC1 to Is- by means of aqueous K I solution followed by weight titration with aqueous NazSz03solution as was done in later nieasurenients. However, we believe the recrystallization ensured a 1: 1 ratio. The sample placed in the calorimeter weighed 442.6630 g. (in vacuo) after correction for the amount adhering to the capillary used to introduce it. The molecular weight was taken as 162.3574; thus, the amount of sample was 2.72647 moles. The Heat ('apacity of Iodine Monochloride. The heat capacity measurements involved no unusual problems, except that the previously mentioned poor thermal conductivity of the Hastelloy calorimeter increased the length of the equilibrium periods after energy additions and, thus, increased the corrections for heat leak. The data are given in Table 11. The measurements were continuous in the sense that intervals were not left between runs. The temperature increments can be inferred roughly from the separation of the mean temperatures. The correction for vaporization into the limited gas space was trivial because of the low vapor pressure. Melting Point of Iodine Monochloride. The melting
Table 11: Heat Capacity of Iodine Monochloride"mb Taw OK.
Cmesad
Tsv, OK.
Cmeasd
17.71 19.24 21.37 23.42 25.18 27.57 30.41 33.50 37.26 41.52 46.06 50.27 55.18 62,OO 69.06 75.95 82.27 88.49 95.62 102.46 109.66 117.60 125.82
1.91 2.18 2.57 2.96 3.31 3.67 4.16 4.65 5.20 5.79 6.31 6.74 7.19 7.71 8.26 8.64 8.93 9.19 9.45 9.72 9.96 10.20 10.53
134.23 142.55 151.94 161.94 171.19 180.48 190.23 200.06 209.53 218.32 226.99 236.28 245.53 255.88 266.91 277.50 287.22 295.04
10.66 10.55" 11.08 11.14 11.43 11.58 11.71 11.84 12.01 12.13 12.28 12.43 12.55 12.5OC 12.60d 12.95d 13.13d 13.74d
Liquid
303.68 310.45 317.76
24.63 24.69 24.61
a 0°C. = 273.15'K., molecular weight = 162.3574, C in gibbs/mole, 1 gibbs = 1 defined cal./defined deg. 442.6630 g. (in vacuo) = 2.72647 moles in calorimeter. Measurements not AT values of last four used in drawing the smoothed curve. rune on solid were 10.303, 9.620, 8.477,and 6.362' in order of increasing 5".
'
point of &-IC1 was measured as a fraction, f, of the sample melted. A plot of T vs. llfgives a good straight line with the intercept of l/f = 0 a t 300.53'K. The slope of the line indicates a liquid-soluble, solidinsoluble impurity of 0.010 mole %. This impurity is undoubtedly something other than nonstoichiometric chlorine or iodine since dissociation products of IC1
Table I11 : Melting Point of Iodine Monochloride" Reaietsnce thermometer T , OK.
Fraction melted, f
0.089 0.195 0.351 0.507 (
w)
Thermocouple T , OK.
300.446 300.44 300.491 300.51 300.506 300.51 300.514 300.50 (300.527) ... Accepted value = 300.53 f 0.05"K.
"0°C.= 273.15"K.
(12) Inorg.
Syl.,1, 166 (1939).
Volume 69,Number 7
Julg 1966
2446
G. V. CALDERAND W. F. GIAUQUE]
should not give a straight line for T us. l/f. Xo impurity eutectic melting was observed during the heat capacity iueasurenients. The melting point details are given in Table 111, where it may be noted that the relative temperature values as given by the gold resistance therniometer are precise to a much higher accuracy than can be claimed with respect to absolute temperature.
Table IV : Thermodynamic Properties of Iodine Rlonol:hloride",b -(Fo H'o) / T
T
CPO
S"
15 20 25 30 35 40 45 50 55 60 65 70 80 90 100 110 120 130 140 150 160 170 180 190 200 210 220 230 "0 250 260 270 273 15 280 290 298 15 300 300 53
1 368 2 333
0,592 1.125 1.744 2.415 3.106 3.806 4 502 5.182 5,845 6.488 7,112 7.716 8.863 9.926 10.921 11.854 12.737 13,570 14.362 15.113 15,829 16.514 17,172 17.802 18,406 18.987 19.548 20 092 20 620 21.133 21,630 22.113 22.263 22.583 23,041 23.405 23.487 23,510
2 2 2 2 3 3 4 5 5 6 6 7 7 8 8 9 9 10 10 10 11 11 11 12 12 12 13 13 13 13 13
170 337 554 807 085 383 690 007 320 646 966 283 910 520 111 682 234 766 281 778 259 725 176 613 038 451 852 241 621 992 353 705 815 050 386 655 716 733
Liquid 32.738 33 503 34.285 34.439
13 14 14 15
733 325 937 058
300 53 310 320 322 Solid was
3 4 4 5 6 6 7 7 7 8 8 9 9 9 io 10 10 10 11 11 11 11 11 11 12 12
284 101 885 599 216 723 180 579 970 312 830 243 619 980 287 556 785 990 193 417 584 716 810 976 154 336
1% 4%
12 12 12 12 L2 13 13 i3 13
614 740 856 89%5 9x0 102 200
223 229
24 588
24 695 24 561 24 501 n
crystalline form.
The Journal of Physical Chemistry
0 0 0 0 1 1 1
(Ha
-
Hoe)/
T
0 0 1 1 2 2 2 3 3 3 4 4 4 5 5 6 6 6 7 7 7 7 7 8 8 8 8 8 8 9 9 9 9 9 9 9 9 9
422 788 190 608 021 423 812 175 525 842 146 433 953 406 810 172 503 804 081 335 570 789 996 189 368 536 696 851 999 141 277 408 448 533 664 750 771 777
19 005
19 178 19 348 19 381
Units are gibbs/mole.
Heat of Fusion. As usual, the heat of fusion was measured by starting heat input somewhat below the melting point and ending somewhat above. Correction was made for any premelting which had occurred before each run started. During the measurement of a heat of fusion the last portion of the solid will be at the bottom of the calorimeter. To avoid excessive superheating of the liquid near the top of the calorimeter, heat was introduced by means of the lower section of the heater. However, a very small current was permitted to flow through the upper section so that a continuous record of its resistance and, thus, small change in the temperature were obtained. This was needed in order to evaluate the heat leak from this portion of the calorimeter surface during energy input. This procedure was especially indicated in the Hastelloy calorimeter because of the poor thermal conductivity. The SC,dT corrections were made using the C, values as listed in Table IV. The details are given in Table V.
Table V : Heat of Fusion of Iodine Monochloride" ____-__Run I
Ti, "K. Tz, "K. Total heat Premelting SCp(s)dT SCdl)dT AHm
a
_______ I1
I11
298 00 297 63 297 304 89 304 305 88 2916 0 2917 2933 9 3 3 5 2 9 -38 4 -38 -33 5 -107 4 -107 -131 8 2773 2775 2772 Average value = 2773 f 2 cal./mole
66 89 3 1 0 4
Units are cal./mole, m.p. = 300.53"K., 0°C. = 273.15"K.
The above data have been used to calculate the therniodynamic properties of solid and liquid IC1. The low temperature extrapolation of heat capacity was made on a C,/T us. T 2 plot. The results have been given in Table IV. The Vapor Pressure of Iodine Monochloride. The vapor pressure of IC1 was measured accurately over a short region above the melting point. lleasurements were made with a mercury manometer using a series of capillaries and mixing bulbs with nitrogen gas as a buffer to protect the mercury. The series consisted of six mixing bulbs each with a volume of 12 cc. Before each measurement a stopcock in the line to the calorimeter was closed, and all the bulbs were evacuated. The three bulbs next to the manometer were then filled with the balancing pressure of nitrogen, and the three next to the calorimeter were opened to the
THERMODYNAMIC PROPERTIES OF IODINE MONOCHLORIDE
IC1. This not only prevented any remote possibility of diffusion of nitrogen into the IC1 but also was designed to sweep any trace of inert gas out of the calorimeter. Since liquid IC1 cannot evaporate a t constant composition, the preceding procedure also ensured that only minor amounts were evaporated, and the composition should have remained a t the stoichiometric ratio. To prevent condensation in the connecting lines, it was necessary to cover the apparatus with a plastic tent and raise the temperature sufficiently above that of tJhe room. Thus, it was not practicable to extend the measurements to higher temperatures with this arrangement. A SociBtB GenBvoise cathetometer with a precision of 0.002 cm. was used as a comparison instrument with a standard meter bar. The usual corrections for capillary depression, temperature of the mercury and the meter bar, and the standard acceleration Of gravity, 980,665 from the local 979.973 cm./sec.z, were made. The observations are given in Table VI, where they are compared with the results of Cornog and Bauerl3!l4by means of an equation based on their results and derived by G r e e n w ~ o d . ~ Cornog and Bauer passed electrolytically generated hydrogen and oxygen over IC1. The amount corresponding to the partial pressure was obtained by absorption in soda lime and weighing. They made the incorrect assumption that the gas consisted only of undissociated IC1. While their data are probably the best previously available, their results over the liquid should have become increasingly low as they continued to remove a gas containing an appreciable excess of Clz, thus increasing the proportion of iodine in the remaining solution. Their lower values in Table VI are consistent with this.
Table VI: Vapor Pressure of Liquid Iodine Monochloride
T , OK., 303.16 P , torr, 39.31 P , torr, 38.90
304.05 41.83 40.83
306.39 48,90 (this research) 46.29 (Cornog and Bauer13)
Assuming that 1, and IC1 crystallize from their solutions as pure phases, St~rtenbeker's~ phase diagram and data indicate the following situations. Below the 12-IC1eutectic temperature, which Stortenbeker has given as 7.9', a dynamic method should produce Iz Clz in such small amounts that the 1 2 could not condense as a solid, and the small pressures of Iz and C1, should be equal. This statement is based on a calculation using the free energy equation for the dissociation of IC1 gas to be given later. Thus, within
+
2447
the limits of accuracy of the measurements of Cornog and Bauer, their total pressure below 7.9' should correspond to the partial pressure of ICl. However, their experimental sequence is not given, and it is possible that measurements at higher temperatures caused the accumulation of a considerable amount of iodine before the measurements below the eutectic were made. The results have been compared on both bases in Table VII. The partial pressure of IC1 is taken from data to be given in a later section. The sublimation pressure of solid iodine is taken from Shirley and Giauque,15who combined their low temperature calorimetric data with the highly accurate vapor pressure measurements of Baxter, Hickey, and Holmes16 and tables of ( F O - Hoo)T for Iz(g) based on band spectra. Table vII : Vapor pressure of Solid Iodine 3Tonochloride. comparison with Data of cornog and ~ ~ ~ ~ ~ 1 3 ' 1
t , OC.
- 15 - 10 -7 0 5
'€uta1 P,mm. (Cornog and Partial P,IC1 Bauer) (this research)
1.2 2.0 3.1 4.6 6.8
0.939 1.509 2.380 3.689 5.623
P,IC1 P,Iz(s)'
4
+ Iz(s)
(this research)
0.006 0.011 0.018 0.031 0.050
0.945 1.520 2.398 3.720 5,673
a The correct PI, has been given; however, Cornog and Bauer, in effect, would have obtained 1.563P1, since they divided all weights by the molecular weight of IC1.
The reason for the higher pressures obtained by Cornog and Bauer is not apparent since the technique is somewhat similar to that used so accurately by Baxter, Hickey, and Holmes for 1 2 . If solid solutions are formed instead of pure phases, the pressures should be lower rat>her than higher than those calculated. Bauer13 states that the data are averages with a maximum spread of 0.4 mm., and it seems probable that some undetected error must have entered. We see no reason why careful dynamic experiments on pure solid IC1 in the range near 0' should not yield accurate and AH",,. values of PICI Between the eutectic temperature and the melting point of IC1, the vapor pressure measured by Cornog and Bauer should be determined by solid IC1 in equilibrium with a liquid in which 1 2 is saturated with IC1. (13) J. Cornog and E. E. Bauer, J . Am. Chem. Soc., 64, 2620 (1942). (14) E. E. Bauer. Thesis, State University of Iowa, May 1942. (15) D. A. Shirley and W. F. Giauque, J . Am. Chem. Soc., 81, 4778 (1959). (16) G. P. Baxter, C. H . Hickey, and W. C. Holmes, ibid., 29, 127 (1907).
Volume 69, A'umber 7
J u l y 1966
2448
At the eutectic the pressure should be essentially that of IC1 alone, and, as will be shown, the pressure of Clz should gradually increase until it contributes about 5% of the total pressure a t the melting point. There are insufficient data to calculate this sequence, but it is evident that the data of Cornog and Bauer are too high near 1 he eutectic but do indicate the contribution of the Clz to the total pressure, approximately, near the melting point. Above the melting point, about the only statement which can be made is that a dynamic method will give pressures which will be increasingly lowered as evaporation proceeds; the iodine concentration of the liquid increases and the Clz pressure is suppressed, an effect which has been noted in Table VI and which is accurately shown in Table VIII, to be given later. Use of d P / d T as a measure of the heat of vaporization, as has been done by Cornog and Bauer and others, is riot valid since the gas and liquid change in coniposition during vaporization. Thus, the heats of vaporization and sublimation given by Cornog and BauerI3 are in error. This is shown by their value of the heat of fusion of ICl, obtained by difference as 1850 cal./niole, whereas the calorimetric result previously given is 2773 f 2 cal./mole. I t should also be noted that dynamic measurements such as those of Cornog and Bauer assume a knowledge of the molecular constitution of the vapor. They absorbed weighed amounts of vapor and divided the weight by the molecular weight of IC1. As we shall show, the vapor is rich in chlorine, and the true vapor pressure carinot be computed from the dynamic results without an evaluation of the amounts of such species as Clz and ICls, if present. The Hear of Vaporization of Iodine Monochloride. StortefibekerZhas shown that the vapor in equilibrium with a liquid with the stoichiometric coniposition IC1 contains a small amount of Clz. He niade a rather inaccurate determination of the proportions in a series of experiments at 30" that varied from 0.8 to 7.67,, average 4.3% Clz. The IC1-Cl2 mixture in equilibrium with the liquid must contain an extremely small pressure of 1 2 , and essentially all of the Iz,corresponding to the Clz formed, should remain in the residual liquid during the vaporization process. Thus, as vaporization continues, the proportions of IC1 and Clz in the vapor and the heat of vaporization should change. However, a series of such nieasurenients, extrapolated back to the start of the first increment removed, should correspond to the heat of vaporization of equilibrium vapor from stoichionietric IC1 liquid plus the heat of solution of the retained Iz. The Journal of Physical Chemistry
G. V. CALDER AND W. F. GIAUQUE
The vapor pressure of IC1 is almost the same as that of CHZNOZ,and Jones and GiauqueI7 found that the heat of vaporization of this substance could be readily measured by condensing it in a bulb cooled by liquid air, with the flow rate controlled by means of a capillary. A similar expedient was tried in the present case, but the exit tube always plugged with solid Iz despite attempts to have most of the pressure drop in an outside capillary system which could be heated locally. I t was then decided to utilize the (F" - H",)/T data based on the third law of thermodynamics along with the corresponding data, based on the band spectrum of IC1 gas, to calculate the heat of the process of evaporating stoichiometric IC1 liquid to give pure IC1 gas. For this purpose it was necessary to determine the partial pressure of IC1 over its liquid. A 250-cc. glass bulb containing 183.0423 g. of IC1 was immersed in a water thermostat held at 30.45". The liquid was stirred gently by means of a magnetic stirrer as samples of vapor were withdrawn very slowly for analysis. The flow rate was controlled by means of a very fine capillary which started at the surface of the thermostat and was electrically heated throughout its entire length to prevent condensation of iodine. There was no appreciable pressure drop in the apparatus at thermostat temperature. When the preceding arrangement was first tried, the observations were somewhat erratic, and it appeared that the heater used to prevent the deposition of IZ in the capillary had heated some of the gas near the capillary entrance. The results indicated that some Clz was diffusing back into the IC1. The addition of a train of two 0.07-mni. capillaries, each 1 cni. long and separated by a 1-cc. bulb just below the water surface of the thermostat, increased the gas flow rate sufficiently in the temperature gradient to prevent any back diffusion, while introducing a trivial pressure drop, and there was no further difficulty. The samples were condensed in bulbs cooled with liquid nitrogen or sometimes ice and weighed. Aqueous KI solution was then sucked into the air-free bulbs to convert all of the halogen to I3-. Titration with sodium thiosulfate in a weight buret evaluated the total halogen collected, and this with the weight enabled the C1:I ratio in the gas to be calculated accurately. The results fall very accurately on a smooth curve through the experiniental increments, arid the interpolation on a weight basis due to the loss of the analysis in run 111 should cause no error. (17) pi. M .Jones and W. F. Giauque, J. A m . Chem. Soc., 69, 983 (1947).
THERMODYNAMIC PROPERTIES OF IODINE MONOCHLORIDE
2449
Table VIII: Composition of Vapor Evaporated from Liquid Iodine Chloride a t 303.60%
wt. Run
removed, g.
G.-atom of halogen removed
0 1 2 3 4 5
0 4.8495 5.4626 5.0481 5.6222 5.8026
0.061418 0.069045 0.063633 0.070623 0.072525
G.-atom of C1 removed
G.-atom of I removed
0.032200 0.036079 0.033102 0.036524 0.037191
0,029218 0.032966 0.030531 0.034099 0.035334
Atom ratio, = Cl/I, in sample, av.
Halogen in remrtining liq., moles
Excess I in remaining liq., g.-atom
1,1048" 1.1021 1.0944 1.0840' 1.0711 1,0526
1 .12741b 1.09670 1,06218 1.03036 0.99505 0.95879
0 0,002982 0.006095 0.008666 0.011091 0.012948
r
Starting weight = 183.0423 g. a Limiting value of T = Cl/I, over pure IC1, obtained by extrapolation. Analysis lost. Data accurately interpolated on weight basis for continuity. 250-cc. flask.
+
Let ~ C I nI = n = total gram-atoms of halogen in a collected sample.
+
3 5 . 4 5 3 ~ ~ 126.904nI = weight of sample
r =
PIC1
nc1=
(1)
n1
weight - 35.453n 91.451
(2)
- 126.904n - weight nI weight - 35.453n 5
PIC1 =
126.904~~ - weight 91.451 ~
1
=
ratio r
(3)
To obtain the part'ial pressures we have the relations P
= PIC1
+ PCl, +
and
r =
PIC1 PIC1
PI,
+ 2PCh + 2P1,
(4)
(5)
Combining eq. 4 and 5 PIC1
2P r + l
= - - 2P1,
It will be shown later that PI, is trivial under the conditions of experiment; thus, the measured values of P and r a t 303.60"K. determine the partial pressure of IC1. The above equations accept the result of Stortenbeker2that ICls is unstable in the gas phase. However, the evidence that IC13 molecules do not constitute a portion of the saturated vapor is not conclusive, and it is desirable to show that our result is accurate even if IC13 is present. Assume p = Then
PIC1
+
PCla
+
PICIS
+ PI,
run, from smooth curve ...
1,0991 1,0899
1.0784 1.0618 1.0437
1.12741 moles of IC1 in a
+ 2PCh + 3PICI3 + + 2P1, PICl,
Eliminating P I C I Sfrom eq. 7 and 8
from which
=
PIC1
=
r a t end of
(7)
~
(3 - r ) P I (r - 1) PCl, - (' + 3 ) ~ 1 , (9) 2 2 2 ~
If we make PCI,= 0 as the most severe deviation from our assumption and introduce the experimental r = 1.1048, PICI= 0.948P - 2.05P1,, if all excess chlorine in the vapor is in IC13, instead of PIC^ = 0.95OP 2.00 PI, if all the excess chlorine is present as Clz. The difference is trivial; thus, if a small fraction or all of the excess chlorine is present in the vapor as IC18, it will not appreciably affect the accuracy of our measured value of the IC1 partial pressure. The observations and results are given in Table VIII. Each removal of soine 0.03 niole of halogen required about 24 hr. so that any lack of equilibrium between the liquid and gas phases is highly improbable. In the last column of Table VI11 values of r corresponding to the end of each run have been read from a smooth curve through the increments. These values have not been used in the present work but could be used in obtaining rather good estimates of the changing pressure of Clz as a function of iodine concentration. From the vapor pressure data in Table VI, the total pressure over stoichiometric liquid IC1 at 303.60"K. is 40.54 torr. Combining this with the limiting value = 38.52 torr and Pel, = 2.02 torr. of r = 1.1048, PICI For the process ICl(s)(l)
AF"/T
=
=
ICl(g)
(10)
R In (760/P) = A(F" - H",)/T
+ AHO~!T
(11)
Combining the ( P o - H o o ) / T data in Table IV with the (F" - H o , ) / T data for IC1 gas given by Evans, Munson, and Wagman,'* we find for process 10 Volume 69, Wumber 7
Julu 1966
2450
G. V. CALDER AND W. F. GIAUQUE
4.5758 log (760/38.32) = -37.794
+ AH00/303.60
from which A H o o = 13,273 cal./mole and = 1995 ca,l./mole. Also A H 0 2 9 8 . p , 0 ~ .=
298.15[A(H0 - H"o)/298.15]
AF"Z98.150K.
+ 13,273 (12)
and using the same thermodynamic tables fiH0298.150K. = -2.090 X 298.15 13,273 = 12,650 cal./mole for the solid. For the heat of vaporization of the liquid, which can easily be supercooled to 298.15"K.
+
ICl(1) = ICl(g) AH0sg8.150~.=
(10)
+
- AHIII,~OO.~VK.
A~sub,1g8.~50~.
ACp,fusion(300 53 - 298.15°K.)
=
2'38
11'360 AF0298.150K.
=
12,650 - 2773 =
"04
cal*'mole
(13)
1974 cal./mole
The partial pressure Of IC1(g) Over and liquid IC1 is obtained by substituting the several values of A ( F o - H",,)/T in eq. 11. A range of values is given in Table IX. Table IX : The Partial Pressure of IC1 over Stoichiometric IC1
a
T , OK.
P , torr
T ,OK.
250 260 270 273.15 280 290 298.15
0.42 1.12 2.74 3.69 6.55 11.44 26.27
300 300,53"
P , torr
29.96 31.13
Liquid 303.60 310 320 330
38.52 59.33 112.4 204.3
1Zelting point.
McMorris and Yost placed BaPtCls with weighed amounts of solid I2 in a reaction vessel of known volunie and studied the total pressure a t various temperatures on the assumption that BaPtCls would give fixed and known partial pressures of Cl2 corresponding to each temperature. This type of technique can fail badly if the decomposition produces microscopic particles of the solid product-in this case Pt. If a system has the mechanism for transporting material from microscopic particles, which have a high and variable free energy, to macroscopic phases, stable equilibrium may be obtained. I n the preceding case the results scattered widely and can be given no weight. In using the combination of reactions 13 and 16, McMorris and Yost made measurements on reaction 16 and a t first combined these results with those given by Dixon21 for reaction 15. Later, Beeson and Yosts2 remeasured the dissociation of NOCl in a very careful series of experiments, which included corrections for gas imperfection. Their interpretation of their data was handicapped by lack of accurate spectroscopic information and interpretation on SOC1, and the then necessary use of the slope of a log K vs. 1/T plot, a method which is correct in principle but lacking in discrimination and sensitivity and which often has led to incorrect results when small systematic errors have temperature trends. They gave the heat for reaction 15 as AH2980K. = 9030 f 100 cal./mole. We find in a re-examination of their excellent experimental data, by the much more powerful point by point method, AHZ~~.K = 9211 f 10 cal./mole. The ( F " - H o o ) / Tdata for Cls were obtained from Evans, Munson, and Wagman's and that for SOCl and X O from the "JAKAF Thermochemical Tables."23 A summary of the present analysis of the data for the three runs given by Beeson and Yost is given in Table
X. The Equiiibriuna Decomposition of ICI. We have also re-examined the several data which have been used to obtain the free energy of the reaction
ICl(g)
=
l/*clz
+ '/zIz(g)
(14)
,1fcL'Iorris and Yostlgcombined the reactions SOCl
=
NO
+ '/ZClZ
The seven observations of JIcMorris and YostlS for reaction 16 cover the range 410-452"K., and the equilibrium constant, K, a t 452°K. has decreased to one-fifth of its value a t 410'K. The reaction was described as "rapid and readily reversible." Values of (H" - H o o ) / Tand (F" - H o o ) / Tfor IC1 were taken
(15)
(16)
(18) W. H. Evans, T. R. Munson, and D. D. Wagman, J . Res. Satl. Bur. Std., 55, 147 (1955). (19) J. MoMorris and D. M. Yost, J . Am. Chem. SOC.,54, 2247
to give the net reaction (14). Also, AIcAlorris and Yost attempted to obtain the equilibrium in reaction 14 directly by assuming that they could use results of Gire,20 who had studied the decomposition pressure of BaPtCls.
(20) G. Gire, Ann. chim., 4 , 186 (1925). (21) J. K. Dixon, 2. physik. Chem. Bodenstein Festband, 679 (1931). (22) C. .\I. Beeson and D. M .Yost, J . Chem. Phys., 7, 44 (1939). (23) "JANAF Thermochemical Tables," Thermal Research Laboratory, Dow Chemical Co., 12lidland, Mich., KOCI, June 30, 1961; NO, June 30, 1963.
and
NO
+ ICl(g) = NOCl + '/*12(g)
The Journal of Physical Chemistry
(1932).
THERMODYNAMIC PROPERTIES OF IODINE L
M
Table X: Analysis of Data of Beeson and Yost
xoc1 = NO + 1/,Cl,n No of
Run
Temp iange, T , OK
I11 IV V
372-162 373-172 373-491
measurementa
AH'o, mean ab
Dev from mean
10 8628 11 8635 13 8628 Accepted AH', = 8630 f 10 cal./mole A H 0 2 ~ 81
6 0 ~=
AHo,
+
f20 fll
f8
- H",) = 8630 + 581 =
A(H02g8.i60K
9211 f 10 cal./mole A F 0 2 g 81 5 = ~ 4897 ~ f 10 cal./mole a
1/*12
C1
=
1/2Cl~
IC1
=
AH"0 = -28,540 cal./g. atom
I 4- C1
AH",
=
AH',
-17,770 cal./g.-atom =
49,640 cal./mole
from which ICl(g)
=
1/212(g) 4- '/ZClz
A H o @= 3330 cal./mole
lichlorris and Yostl9 used a similar method to derive an earlier value A H " , = 3280 cal./mole; however, while the extent of agreement is interesting, we believe that the value AHOo = 3255 cal./mole, derived from the combined equilibrium measurements of JlcJlorris and Yost'Qand Beeson and Yost,22should be accepted at full weight. In order to check on our assumption that the partial pressure of Izover stoichiometric IC1 is trivial, A F 3 0 3 . 6 0 0 K . is found to be 3689 cal./mole.
4'886
~
245 1 ~
torr. The correction to the partial pressure of IC1 derived from eq. 6 is less than 0.01 torr, and no readjustment of the calculations is necessary. At this point it is of interest to make some very approximate calculations as to the extent of the dissociation in liquid IC1 at 303.60"K. Using the heat of fusion of 1224as 3740 cal./mole at its melting point, 386.8"K., and estimating the ACp of fusion as 1.5 cal./mole on the basis of those for C12 (2.75) and Brz (2.0), we may write
AHm = 3160
from Evans, Nunson, and Wagman.18 The values of a H o o show a small trend at the three higher temperatures, and, on account of the considerable decrease in the value of K , we have accepted the four lower values which give AH',, = - 5375 f 15 (actual average f 13) cal./mole, AHC2981 6 0 g . = -5141 f 15 cal./mole, and AF0298 16 = - 1216 f 15 cal./mole. Combining the data for reactions 15 and 16, aHoo= 3255 f 25 cal./ mole, AH0298150K = 3277 i 25 cal./mole, AF0298160K = 3682 f 25cal./niole. Evans, Munson, and WagmanI8 summarize the data and references for the spectroscopic dissociation limits of Clz, 12,and IC'l =
~
Ids)
Cal./mole of XOCI.
I
~
=
14)
(17)
+ 1.55" cal./mole
(18)
d In P(I)/P(,) -AH, -~ dT RT2
(19)
from which PI,(l) = 2.570P1,(,) for supercooled liquid iodine at 303.60"K. From Shirley and Giauque,I5 PI,(,, = 0.4843 torr at 3O3.6O0K.,from which PI,(^) = 1.245torr at 303.60"K. Making the approximate assumption that Raoult's law holds for the whole range of composition in the torr and halogen mixture NI, X 1.245 = 3.59 X N I , = 0.00288 in liquid IC1 a t 303.60"K. If we now assume an equal molal concentration of Clz, and combine this with the measured partial pressure of Clz = 2.02 torr, a Raoult's law calculation gives 2.02/0.00288 = 701 torr as the vapor pressure of liquid chlorine at 303.60"K. This is at least some eightfold too low and suggests that the dissolved chlorine is held largely as IC13, which is not improbable since IC], exists as a solid when the concentration of Clz is increased sufficiently. We do not care to carry this kind of speculation further since the ideal of accurately determinable concentrations of specific molecular aggregates ill liquids is usually unattainable. In the present case one must ask the question as to whether some of the chain structure in solid IC1 does not persist into the liquid to complicate the ideal use of mole fractions. Sforeover, the long-known electrical conductivity of IC1 indicates some ionization. The evidence is summarized by Greenwood,* who quotes evidence by Fialkov and KaganskayaZ5for the reaction 2IC1 = I + IClZ-. Chlorine formed a t the anode and iodine a t the cathode would form IC1, and Is, respectively. The high heat capacity of liquid IC1 is undoubtedly due to the increasing dissociation with increasing tem-
+
(24) "Selected Values of Thermodynamic Properties," U. S. National Bureau of Standards, Circular 500. U. S. Government Printing
Volume 69, Number 7
Julu 1966
~
G. V. CALDER AND W. F. GIAUQUE
2452
perature. The heat capacities of solid Clz and IC1 at their melting points are 13.27 and 13.23 gibbs/ mole, respectively, whereas the corresponding values for the liquids are 16.03 and 24.59 gibbs/mole. Several additional useful relationships are given. Combining cq. 10 and 14 ICl(S) A H O O
=
=
+ '/zIz(g)
'/ZClZ
16,528
f
AF0298.150~. = 5677
(20)
25 cal./mole
f
=
AHoo = - 15,658 cal./mole AF0298.160K. = -4628 cal./mole
The Journal
oj
Physical Chemistry
ICl(s)
=
'/ZClZ
AHoo = 8699
f
AF0298.150K.= 3363
+ '/zIz(s)
(22)
25 cal./mole f
25 cal./mole
and ICl(1) = '/2C12
25 cal./mole
AF0Z98.150K. = 3342
and combining eq. 20 and 21 Iz(g)
we find
t '/zIz(s) f
(23)
25 cal./mole
(21)
Acknowledgment. We thank J. P. Chan, J. A. Duisman, and P. R. Siemens for assistance with the calorimetric measurements.
