April, 1959
HEATSOF FORMATION OF FERROUS, FERRIC AND MANGANOUS CHLORIDES
Elemental analyses were obtained for the residues from the Ca, Sr, Ba, La and Bi compounds.s Unlike the results found in air,4v6where the ultimate residues are usually the metal oxides, the residues obtained here contained considerable amounts of carbon and nitrogen, and some hydrogen, as well as metal. An attempt was made to determine the infrared spectra of the Ba and Ca residues in KBr pellets. No recognizable absorption peaks were observed in the wave length region 3 to 15 fi, indicating that complete destruction of the starting chelates had occurred. It is striking that decomposition of the Ba and Sr chelates proceeds with only a small loss of weight. Most of the decomposition products must therefore be non-volatile. Because of the observed volatilities of the unchanged chelates it is difficult to compare the heat stabilities of the 8-hydroxyquinolates among themselves under the conditions used here. It is also difficult to compare the stabilities in vacuo with those observed in air.4J It is interesting, however, that some of the 8-hydroxyquinolates (notably the Mg and Mn compounds) show no decomposition at temperatures in excess of 400". Insofar as a comparison is possible, the 8-hydroxyquinolate chelates appear less stable in air than in v m o . This is reasonable in view of the possibilities of reaction with molecular oxygen in air. From Fig. 1 it will be noted that, for the divalent metal chelates which sublime, the temperature range over which sublimation takes place is a function of the metal present. In contrast to this behavior the trivalent metal 8-hydroxyquinolates all sublime in the same temperature range, 250 to
350". According to the Langmuir equationg 1 the rate of evaporation of a subliming substance in vacuo (8) Elemental analyses of the residue8 left by the 8-hydroxyquinolates of the indicated metals (% C, H and N, respectively): Ca, 03.1. 3.1, 7.0; Sr, 53.5, 1.6. 5.5: Ba. 43.1, 2.1, 6.1; La, 50.7, 1.7, 2.0; and Bi, -, -, 3.1. Only N waa determined for the Bi residue. (9) A. Weiasberger, ed., "Technique of Organic Chemistry." Vol. IV. Interscience Pub., Inc., New York, N. Y., 1951, p. 501.
605
at a given temperature should be a function of the vapor pressure, molecular weight and surface area of the substance. Since the molecular weights of many of the divalent metal 8-hydroxyquinolates are similar, the observed differences in temperature range for these compounds in Fig. 1 must be due primarily to differences in vapor pressure and/or surface area. Q =P/.\/2m
(1)
where Q is the evaporation rate per unit area, P is the vapor pressure of the compound, M is the molecular weight, R is the gas constant and Tis the absolute temperature. It is interesting that a relationship appears to exist between the temperature range of sublimation for the divalent metal 8-hydroxyquinolates and the electronegativitiesIO (XM)of the bonded metal ions. The order of decreasing XMfor these metals is Cu > Ni > Co >P b > Zn, Cd > Mn> Mg > Ca. With the exceptions of P b and Zn this is also the order of increasing temperature of sublimation. Since no relation between XM and surface area would be expected, the observed relation between XM and sublimation temperature probably is the consequence of a relationship between X M and vapor pressure for these compounds. Since the vapor pressure is determined, at least in part, by intermolecular forces in the solid compounds, these forces must decrease with increasing values of XM. The lack of dependence upon X M of the sublimation temperature range for the trivalent metal 8-hydroxyquinolates may be due to the additional shielding of the metal atom by the third 8-hydroxyquinolate residue present in these molecules. Acknowledgments.-The writers are grateful to Mrs. M. A. Pawlikowski for the preparation of some of the 8-hydroxyquinolates1 to Dr. H. Lady for the infrared determinations, and to Miss M. A. Knuth for help with some of the sublimation studies. (101 W . Cordy and W. J. 0. Thonim, J . Chem. Phua., 34, 439 (195D).
HEATS OF FORMATION OF FERROUS CHLORIDE, FERRIC CHLORIDE AND MANGANOUS CHLORIDE BY MARYF. KOEHLER AND J. P. COUGHLIN~ Contribution from the Minerals Thermodynamics Experiment Station, Region I I , Bureau of Mines, United States Department of the Interior, Berkeley, Gal. Receivad September 87, 1068
Heats of formation from the elements of anhydrous ferrous chloride, ferric chloride and manganous chloride were determined by measuring appropriate heats of solution and reaction in hydrochloric acid. The results obtained are as follows: ferrous chloride, -81.86 f 0.12; ferric chloride, -95.7 ==! 0.2; and manganous chloride, -115.10 f 0.12 (kcal./mole at 298.15'K.).
Previously existing2-svalues of the heats of formation of the chlorides of iron and manganese are ( 1 ) Formerly physical ohemiat, Minerals Thermodynamics Experiment Station, Bureau of Mines. Region 11, Berkeley, Calif. (2) F. D. Rossini, D. D. Wagman, W. H. Evans, S. Levine and I. Jaffe. Natl. Bur. Standards Circ. 500, 1952. (3) 0. Kubaachewski and F. L. Evans, "Metallurgioal Thermoohemistry," John Wiley and Sone, Ino., New York, N. Y., 1956.
ba&d upon rather old and somewhat uncertain Ohermochemical measurements. It is difficult to appraise the accuracy of these values. (Selected data in the two references cited differ by as much as 3.3 kcal./mole.) As a consequence, new determinatiQns appeared desirable. The present paper reports new values for anhydrous ferrous chloride,
MARYF. KOEHLER AND J. P. COUGHLIN
606
Vol. 63
ferric chloride and manganous chloride, obtained by solution calorimetry.
maintained 30.00 f O.0lo, and all results were corrected to exactly 30.00' (303.15"K.). Ferrous Chloride.-Table I gives the equations Materials.-Iron powder was prepared from ferric hydroxide, obtained by dissolving pure iron wire in sulfuric acid for the reactions used t o determine the heat of and precipitating with ammonia. The ferric hydroxide was formation of ferrous chloride and the mean values washed, dried a t 140" and crushed to - 100 mesh. It then of the heats and the precision uncertainties. Rewas reduced with hydrogen for 5.5 hr. a t 500°, after which the metallic clumps were broken up and again treated with actions 1 and 2 were measured consecutively in hydrogen for 5.5 hr. a t 500". The oxygen content of the the same batch of acid; reactions 3 and 4 were product as it was used in the measurements ranged from 0.10 measured consecutively in another batch of acid. to 0.22%. Thus the final solutions after conducting reactions The ferrous chloride was prepared by Kelley and Moore.4 1 and 2, and aft,er conducting reactions 3 and 4, are I t was retreated before use in the present measurements bg identical and AH1 AH2 - AH3 - AH4 yields the drying in a stream of hydrogen chloride for 5 hr. at 450 and 6 hr. a t 460-475". Tests for ferric iron (thiocyanate heat, AH6, of the resultant over-all calorimetric remethod) indicated 0.005% (e uivalent to 0.015% ferric action.
+
chloride). Analysis gave 44.06% iron and 56.10% chlorine, as compared with the theoretical 44.06 and 55.94%. Ferric chloride was prepared by the reaction of highpurit iron wire and dry hydrogen chloride gas near 350". The Lrric chloride sublimed in the hot zone and condensed downstream on the cooler parts of the tube wall. The condensate was removed and resublimed in dry hydrogen chloride. Tests for ferrous iron indicated 0.006% (equiva!ent to 0.014y0 ferrous chloride). Analysis gave 34.41% iron and 62.42% chlorine, as compared with the theoretical 34.43 and 65.57%. Reagent grade hydrogen peroxide was used at concentrations of 12.76 to 13.73%. Analysis of the hydrogen peroxide will be discussed later. T h e manganese was an electrolytic product furnished by the Boulder City, Nev., Station of the Bureau of Mines. The purity was stated to be 99.92%. It was degassed b j heating in a stream of pure helium for 1.5 hr. a t 990-1020 . Manganous chloride was re ared from reagent grade manganous carbonate by K. Eonway of this Laboratory. The carbonate was dissolved in hydrochloric acid and the solution was filtered and then evaporated until a good crop of crystals of the hydrated chloride was obtained. The crystals were transferred to a Pyrex tube and dehydrated by rolonged treatment with a stream of hydrogen chloride anznitrogen at 350". Analysis of the product gave 43.71% manganese and 56.37% chlorine (theoretical values, 43.66 and 56.34%).
8.
TABLE I HEATOF FORMATION OF FERROUS CHLORIDE
+
AHan.ir,
Reaction
ca1.
(1) F e ( 4 2H+(soln) = Fe++(soln) -20,820 f40 Hdg) (2) 2(HC1*12.731H20)(1)= 2H+(soln) 0 f 10 2Cl-(soh) 25.462H20 (soh) (3) FeClz(c) = Fe++(soln) -15,000 f 20 2Cl-(soln) - 2,030 i20 (4) 25.462H20(1) = 25.462H20(soln) - 3,790 =k 50 2(HC1~12.731H20)(1) = (5) Fe(a) FeClz(c) Hz(g) 25.462H20(1) At 298.15'K.; AH6 = -3,790 - 270 = -4,060 f 50
+
+
+
+
+
+
+
Six determinations of reaction 1 were made, yielding -20,860, -20,780, -20,780, -20,850, -20,810 and -20,870 cal. These results contain a correction of -338 cal. for heat absorbed in vaporizing water and hydrogen chloride by the escaping hydrogen and a correction of 4 cal. for the heat evolved in reducing a small amount of chloroplatinic Methods and Results acid used as a catalyst for the reaction. The correcThe heats of formation of ferrous and manganous tion for the oxygen content of the metal ranged from chlorides were obtained by calorimetric reaction - 3 to -8 cal., depending upon the time of storage schemes involving solution of the chlorides and the of the sample. Reaction 2 is necessary to maintain stoichiomemetals in hydrochloric acid. In the case of ferric chloride, an additional reaction step (the oxidation try. Only a single heat measurement was made to of ferrous ion by hydrogen peroxide) was necessary. confirm previous workg in which substantially this All measurements were made with previously de- same reaction was measured and found to have scribed apparzltus.6.6 The results are expressed virtually zero heat effect. Six measurements of the heat of reaction 3 were in defined calories (1 cal. = 4.1840 abs. joules). All sample weights were corrected to vacuum and conducted, the values being - 14,980, - 14,980, all formula weights accord with the 1954-1955 Re- - 15,000, - 15,000, - 15,010 and - 15,000. The calculated correction for the ferric chloride content port on Atomic Weights.? All reactions were conducted in 1936.2 g. of of the ferrous chloride is about 1 cal., a negligible 4.360 m hydrochloric acid; the preparation and amount. Substantially the same reaction as reaction 4 has standardization were discussed previously.8 The amounts of the other substances were 0.03 mole for been investigated previously. 9 Two additional the reactions used to obtain the heats of formation heat determinations were made to confirm the preof the iron compounds and 0.02 mole for the reac- vious results. In the first, the water was added to tions used to obtain the heat of formation of man- the calorimeter in three equal portions. The measured heats were -676, -672 and -666 cal., ganous chloride. Each individual heat of reaction measurement making a total of -2,014 cal. for 25.462 moles of was accompanied by a calibration of the cabri- water. In the second, the total required amount of metric system. The calorimeter surroundings were water was added in one step, and -2,060 cal. was t.he measured heat effect. These two measure( 4 ) K. E. Kelley and G. E. Moore, J . An. Chem. Soc., 61, 1264 ments and the three previous ones9 lead to -2,030 (1943). f 20 cal. as the heat of reaction 4. (5) J. C. Southard, Ind. Eng. Chem., 82,442 (1940). (6) J. P. Coughlin, J . Am. Chem. Soc., 77, 868 (1955). Reaction 5 represents the formation of ferrous (7) E. Wichers, ibid., 78, 3235 (1956). (8) J. P. Ceughlin, ibid., 78, 6479 (1956).
(9) J. P. Coughlin, THISJOURNAL, 62, 419 (1058).
I
HEATSOF FORMATION OF FERROUS, FERRIC AND MANGANOUS CHLORIDES
April, 1959
607
chloride, water and hydrogen from iron and hydro- acid makeup after reaction 9, to maintain rigid chloric acid. The heat was corrected to 298.15"K. stoichiometry. A single measurement was made to by means of the heat capacity data of KelleylO and confirm that the heat of reaction 10 was virtually Rossini. zero, as for the similar reaction 8. The heat of formation of ferrous chloride from the Six measurements of reaction 11 gave -24,440, elements was obtained from the heat of reaction 5 -24,470, -24,450, -24,480, -24,460 and and the literature value2 of the heat of formation of -24,440 cal. The calculated correction for the ferrous chloride content of the ferric chloride is neglithe hydrochloric acid (-38,900 f 50 cal./mole). gible (cu. 1 cal.). Fe(a) + ClZ(g) = FeClz(c) Three measurements of reaction 12 gave -3,630, AH288.16 = -81,860 f 120 cal. (6) - 3,620 and -3,630 for 45.483 moles &f water. Ferric Chloride.-Table I1 summarizes the reReaction 13 represents the formation of ferric actions used t o obtain the heat of formation of chloride, hydrogen, and water from iron, hydroanhydrous ferric chloride. Reactions 7-10 were chloric acid and hydrogen peroxide solution. Corconducted consecutively in the same batch of hy- rection of the heat to 298.15"K.WRS made by means drochloric acid and reactions 11 and 12 likewise of the heat capacity data of Kelley,l0 Rossini," were conducted in another batch of acid. The and Giguere and co-workers.12 summation, AH7 AH, AH9 4-AHlo- AHl1 The heat of formation of ferric chloride from the AH12, gives AHl$ for the over-all calorimetric reac- elements was obtained from the heat of reaction 13 and literature values2 of the heats of formation of hydrochloric acid (-38,900 f 50 cal./mole), waTABLEI1 ter (-68,320 f 10 cal./mole), and hydrogen perHEATOF FORMATION OF FERRIC CHLORIDE oxide solution (-45,6GO f 200 cal./mole).
+
+
+
AHsoi.18,
Reaction
aal.
Fe(a) 2H+(soln) = Fe++(soln) -20,820 f 40 Hzk) 2(HCl. 12.731Hz0)(1) = 2H +(s o h ) 0 f 10 2Cl-(soln) 25.462Hz0 (soln) -31,350 f 40 Fe++(soln) ~/z(Hz02~12.580H20)(l) H+(soln) = Fe+++(soln) 7.290HzO(soln) ( HC1~12.731Hz0)(1)= H +(soln) 0 i 10 Cl-(soln) 12.731H20(soln) -24,460 f 20 FeCldc) = Fe+++(soln) 3C1(soln) - 3,630 f 10 45.483H~0(1)= 45.483HzO(soln) 3(HC1.12.731H~O)(l) -24,080 f 70 Fe(a) '/z(HzOz'12.580Hz0)(1) = FeCldc) Hz(g) 45.483Hz0(1) At 298.15"K., A H I S = -24,080 - 410 = -24,490 f 70
+
+ +
+
+ +
+
+
+
+
+
+
+
Fe(a)
+ 3/2C&(g) = FeCla(c) AHz~s.Ic, = -95,700 f 200 cal. (14)
Manganous Chloride.-Table I11 summarizes the measurements for obtaining the heat of formation of anhydrous manganous chloride. Reactions 15 and 16 yield the same end products and concentrations as reactions 17 and 18. Consequently, AH19 = AH15
+
AH16
- AH17 - AH18.
TABLEI11 OF MANGANOUS CHLORIDE HEATOF FORMATION
+
AHaaa.16.
Reaction
cal.
(15) Mn(a) 2H+(soln) = M n + + -51,610 f 40 (sol4 HAg) (16) 2(HC1.12.731Hz0)(1) = 2H+(soln) 0 f 10 2C1-(soln) 25.462H20( s o h ) (17) hfnClz(c) = Mn++(soln) 2CI- 12,470 f 20 (soln) - 2,010 f 10 (18) 25.462H20(1) = 25.462HzO(soln) (19) Mn(ct\ 2(HC1.12.731Hz0)(1)= -37,130 f 50 MnClz(c) Hi(g) 25.462&0(1) At 298.15"K., AH19 = -37,130 - 260 = -37,390 f 50
+
+
+
+
Reactions 7 and 8 are identical with reactions 1 and 2 of Table I, and t,he same heats apply. In studying reaction 9 (the oxidation of ferrous ion by hydrogen peroxide) weighed amounts of the peroxide solution were sealed in glass bulbs, which were broken in the calorimeter a t the proper time. It was arranged that the amount .of peroxide was slightly less than that required to oxidize all of the ferrous ion present. Determinations of the ferrous ion content of the calorimeter solution before the peroxide was added, and immediately after reaction 9 was complete, served to determine the actual amount of peroxide used in reaction 9. This procedure was necessary because of the inherent instability of the peroxide. There always was some decomposition during filling, sealing and storing of the bulbs. This made it impossible to know their exuct peroxide contents, even though accurate weighings were conducted. Six measurements of the heat of reaction 9 gave -31,340, -31,370, -31,410, -31,370, -31,280 and -31,350 cal. Reaction 10 involves the necessary additional
Six measurements of the heat of reaction 15 gave -51,650, -51,540, - 51,650, - 51,590, - 51,620 and -51,590 cal. These results include a -329 cal. correction for vaporization of water and hydrogen chloride by the escaping hydrogen. The heat of reaction 16 would be expected to he virtually zero, in view of results for similar reaction 2. A single measurement was made to confirm this. Six heat determinations were made of reaction 17. The results were - 12,470, - 12,450, - 12,460, - 12,460, - 12,480 and - 12,490 cal. Three determinations of the heat of reaction 18 were made, -2,010, -2,010 and -2,000 cal. It is to be expected that the heat of this reaction would be virtually identical with that of similar reaction 4. Correction of the heat of reaction 19 to 298.15'K.
(IO) K. K. Kelley. U. 9. Bur. Mines Bull. 477, 1950. (11) F. D. Rossini, J . Research Natl. Bur. Standards, 4, 313 (1830).
(12) P. A. Giguere, B. G . RIorissette, A. W. Olrnos and 0. KnoiJ. Can. J . Chem., 33, 804 (1955).
+
+
+
S. V. R. MASTRANGELO
608
was made, using heat capacity data from the literature.lO*ll Combination of heat of reaction 19 with the heat of formation of hydrochloric acid2 yields the heat of formation of manganous chloride from the elements. Mn(a)
+ Clz(g) = MnC12(c) AH208.16
= -115,190 i 120 cal. (20)
Discussion of Results It is believed that these new heat of formation values are an improvement over previously existing data, especially from the viewpoint of assigning uncertainties. The value for ferrous chloride is -360 cal./mole more negative than the result listed in previously mentioned corn pi la ti on^.^^^ However,
Vol. 63
it is virtually identical with the value (-81,900 cal./mole) selected previously in connection with thermochemical measurements of fayalite, l 3 ferrous oxide14 and ilmenite.16 The value for ferric chloride is identical with the selection of Kubaschewski and Evans13but 1,100 cal./mole more positive than that of Rossini and co-workers.2 On the other hand, the value for manganous chloride checks the selection of the latter workers and is 3,190 cal./mole more negative than that of the former. (13) E. G. King, J . A m . Cham. SOC.,74, 4446 (1952). (14) G. L. Humphrey, E. G . King and K. K. Kelley, U. 8. Bur. Mines Rept. of Investigations 4870, 1952. (15) K. K. Kelley, 8. S. Todd and E. G. King, U. 8. Bur. Miaea Rept. of Investigationa 6059, 1954.
EQUATION OF SOLUBILITY OF NONELECTEOLYTES WHICH POSSESS A SPECIFIC INTERACTION BY S. V. R. MASTRANGELO Research Division Contribution No. 276, Jackson Laboratorg, Organic Chemicals Department, E . I . du Pont de Nemours and Company, Wilmangton, Delaware Received Septembsr IO,1068
An equation of solubility has been derived for two component systems which ossess one or more similar interactions per molecule. The equation redicts the number,f, of interactions per molecule anathe equilibrium constant, K. Good agreement with ex erimental fata on the solubility of chlorofluoromethanesand ethanes in tetraethylene glycol dimethyl ether and of chloroform in acetone is shown.
The solubilities of a number of simple chlorofluorocarbons have been measured in tetraethylene glycol dimethyl ether. Interaction presumably occurs by formation of hydrogen bonds. All of the systems studied gave a linear plot when the activity coefficient, +yr,of the solute was plotted against its activity, a, (defined as the ratio (P/Po) of its vapor pressure above the solution to the vapor pressure of the pure liquid solute at the same temperature). This linear relationship held in the region: 0 < a, < 0.5. When a, > 0.5, a downward curvature was noted in most cases. Assuming two experimental parameters, this suggests an equation of the form shown by equation 1. Yr
= Ki
+ Kz ar + f(KiJKz)ar2
(1)
f N a solvent sites and (Nr - 2) are free to occupy positions near the bound molecules in a manner that is energetically and statistically equivalent to that of the pure solute liquid. It is to be noted that in this rather simplified model, the elements of the solvent molecule which are not part of the active sites, e.g., the methylene groups in tetraethylene glycol dimethyl ether, are ignored. The fNa available ether oxygens are considered as separate entities. Use of a coordination’ number is thus avoided. Derivation of the Solubility Equation-Derivation of the solubility equation follows simply by application of standard methods of statistical thermodynamics. For this purpose, the “quasichemical” treatment of Fowler and Guggenheim (p. 358)2 was used with the “harmonic oscillator model” (p. 325)”for the liquid. The details of this treatment, using the proposed model, are straightforward after stating the partition functions for the bound and free solute molecules. These may be written according to equations 2 and 3, remectivel’v.
where K1 and Kz are constants. Since an equation of this type allows experimental solubility data to be plotted directly and correlated very simply, it was felt that some theoretical justification might not only yield some needed insight on solubility theory in general but also provide a useful interpretation of the constants of equation 1. Statistical Model.-The model which fits the systems studied regards the solvent as B number, fN,, of energetically identical sites, where N is the total number of solvent molecules in solution and f is the number of active sites per molecule. For example, the solvent used was tetraethylene glycol where dimethyl ether. This molecule has a theoretical , I = J’ (4) maximum functionality, f, of 5, one for each unhindered ether link, when a simple solute such as ) E. A. Guggsnheim. “Mixtures,” Chapter XI, Oxford Univordichlorofluoromethane (CHC12F) is used. The sity( 1 Press, Amen House, London, 1952. total number of solute molecules in solution is (2) R. H. Fowler and E. A. Guggenheim, “Statiatioal Thermotaken as N,, of which Z are bound to some of the dynamics,” Cambridge Univereity Press, Cambridge, lg39.
[Ela
