HEATSOF COMBCSTIOS OF Resa ASD Re&
May, 1962
thiomalic, P-hydroxypropionic, glycolic, and thioglycolic acids, as shown in Table I. The stability (of chelates with analogs of acetic, propionic, and succinic acids decreases for any given acid as tho donor group changes fromNHs t o OH to SH. This decrease parallels the order of decreasing basicity in any given series. Considering acids which contain the same donor group (NH2 or OH or 8H) decreasing stability is observed in the analogs of succinic, propionic, and acetic acids, in that order. This order is found to be unusual in that the six-member ring chelates of the propionic acid analogs are more stable than the five-member ring chelates of the acetic acid analogs. Succinic acid analogs, although the weakest bases, form the strongest chelates in any group. These would be expected to be the least stable inasmuch as a decrease in ligand basicity results in a corresponding lower affinity for metal ions. The behavior of succinic acid analogs may be attributed to the formation of either a bidentate or tridentate complex. The bidentate complexes possibly may possess either a five- or six-membered ring, thus taking on the structure of acetic acid analogs with a mbstituted carboxylmethyl group, or of propionic acid analogs with a substituted carboxyl group. I n either case, the bidentate chelates would be expected to be only slightly stronger than tbose of the acetic acid or propionic acid analogs, whereas tridentate formation would be expected to give a large increase in the stability constant (-2 log units).22 Such an increase is observed only for malic scid. Lumb and Martel123 ( 2 2 ) M. Cefola, A. Tompa, A. InoTg. Chem., 1, 290 (1962).
V. Celnano, and P. S. Gentile,
791
have observed that aspartic acid forms bidentates with alkaline earth ions and for these chelates the stability constant was 0.2log unit higher than the glycine, the small increase being attributed to the inductive effect of the negative carboxylate group in aspartic mid, which can lend stability to the structure by increasing the basicity of the donor group toward the metal ion. The difference between the stabilities of the aspartate and p-alinate chelates of uranyl ion is approximately 0.2 log unit, thus indicating formation of a bidentate chelate with aspartic acid. I n all probability aspartic acid forms a six-membered ring since the p-alinate is more stable than the glycinate. Table I1 lists the relative position of U02t2 in a series of metal ions for the formation of glycinates and aspartates. TABLB I1
STABILITY CONSTANTB OF VARIOUSMETALSWITH GLYCINE AND
-Glycine---Metal ion
Cu(I1) UOZ Xi( 11) Zn(I1) Co(I1) Pb(I1) Mn( 11) Ag( 11)
log
ASPARTICACID
Ki
8.62(24) 7.53 6.18 (24) 5.52(24) 5.23(24) 5.47 (24) 3 44 (24) 3.51 (24) I
-Aspartic AMetal ion
Cu(I1)
uoz
Ni( 11) Co(I1) Zn(I1) Cd(I1) Hg(1I) Ca( 11)
acidlog Kt
8.57(25) 8.00 7.12 (25) 5.90(25) 5.64(25) 4.37 (25) 2.43 (23) 1.60 (23)
(23) R. F. Lumb and A. E. Martell, J . Phys. Chem., 87, 690 (1953). (24) G. Monk, TTans. Faraadag Soe., 47, 297 (1951). (25) S. Chaberek and A. E. Martell, J . Am. Chem. Soc., 74, 6021 (1952).
THE HEATS OF COMBUSTION OF Resz AND RezS7 AND THE THERMODYNAMIC FUNCTIOKS FOR TRASSITION METAL SULFIDES'~2 BYJ. E. MCDONALD~ AND J.
w.COBBLE
Department of Chemistry, Purdue University, Lufayette, Indiana Received Julv Zh. 1861
The heats of combustion of Re& and Red37 have been measured by bomb calorimetry. The heat and free energy of formation a t 25' of Re& are -42.7 f 1.2 kcal. mole-l and -41.5 =!= 1.2 kcal. rnole-l, res ectively. The similar quantities for Re& are - 107.9 d: 1.8 kcal. mole-' and - 101.0 f 1.8 keal. mole-', respectively. &ith data on these key sulfides, along with recalculated heate of formation from previous vapor pressure measurements, it has been possible to estimate the thermodynamic functions for a number of sulfides not previously available.
I. Introduction Very few thermodynamic data have been available on transition element sulfides in spite of their potential interest in solid state and high temperature chemistry. Certain vapor pressure meas(1) This research was supported by the United States Air Force through the Air Force Office of Scientific Research of the Air Research and Development Command under contract AB 18(600)-1525. Reproduction in whole or part is permitted for any purpose of the United States Government. (2) This constitutes communication number VI11 in our previous series on rhenium and technetium chemistry. For the previous paper in this series see J. A m . Chem. Soc., 82, 2111 (1960). (3) From the Ph.D. thesis of J. E. McDonald, Purdue University, 1961; Dow Chemical Co. Fellow, 1959-1960; Procter and Gamble Fellow, 1960-1961.
urements of Biltz, Juza, and their co-workers fix the heats of formation of Pt&, PtS,4 IrzS8, ReSz,s Rust,' and OSSZ~ if the assumption 18 made that the only vaporizing species is gaseous sulfur. Since Re& can be directly sublimed at -lOOOo, this assumption needs independent verification. The purpose of this communication is to report the calorimetric determination of the thermody(4) W. Bilta and R. Juza, 2.anorg. Chem., 190, 161 (1930). (5) W. Biltz, J. Laar, P. Ehrlich, and K. Meisel, zbid., 283, 257 (1937). (6) R. Juza and W. Biltz, 2. Elektrochem., 37, 498 (1931). (7) R. Juza and W.Meyer, 2. anorg. Chem., 213, 273 (1933). (8) R. Juaa, ibiad., ZlS, 129 (1934).
J. E. MCDOPL'ALD AND 5. W. COBBLE
792
namic functions for Resz and Reps?. With these data it now is possible to check the nature of the vaporization process in these and similar sulfides. Further, because of the central and key position of rheiiium sulfides among the other. transition elements, thermodynamic functions for many other sulfides can be estimated which have not been available previously.
.
*
J.
111. Experimental Results
The combustion of Rez%produces both SO2and
so3
-
+ 21 O d d
Re2S7(c) ReBdc)
+ 1402 (g)
--t
+
Re~Odc) 7S02(g) (1)
RezOdc)
+ 7SOs(g)
(2)
Experimentally, the fraction of sulfur which appears as SOz in the combustion products was found to be 0.883 0.019, independent of the amount of sulfide burned. The thermal data for the sum of reactions 1 and 2 are summarized in Table 1. I h e average AE for the processes in eq. 1 and 2 is - 1.1743 f 0.0019 cal. mg.-l of Rez&. Corrected to 100% SO3, and changing AE to AH,lo the heat of combustion for reaction 2 becomes AHzO = -849.9 f 1.7 kcal. mole-l Res&. Using necessary auxiliary thermodynamic data" and a heat of formation of Re207(c)of -296.7 i 0.8 kcal. rno1e-',l2 the heat of formation of Rez% becomes -107.9 1.8 kcal. mole-1 at 25'.
*
11. Experimental Chemicals.-Rhenium heptasulfide was prepared by precipitation from a strongly acid perrhenate solution with hydrogen sulfide gas. The black precipitate was washed repeatedly with water, carbon bisulfide, and finally ether, and dried under vacuum over P z O ~The . carbon bisulfideether extractions were repeated until the dried solid had attained a constant composition. The material was analyzed for both rhenium and sulfur content, the former by reduction to the metal with hydrogen a t 700". The sulfur was determined by oxidation with aqueous hydrogen peroxide and precipitation of the sulfate so formed as barium sulfate. The composition of the product was: Re, 62.24%; S,37.05.%, corresponding to Re& 4 9 . Rhenium disulfide was prepared by the thermal deco2position of Re2ST under a stream of dry nitrogen a t 450 The product was analyzed by reduction to the metd with hydrogen as before. The composition, Re, 74.39%, corresponded to a material ReSz.02. Equipment .-The calorimeter consisted of a Parr microbomb of 50-ml. capacity immersed in a 2-1. calorimeter vessel. The vessel was a brass container with a double wall which could be evacuated, and had a lucite cover lid. The calorimeter assembly was submerged in a 15-gal. temperature-controlled (EtO.OO5O) thermostat. Access tubes for the stirrer, heater, and connections were provided in the lucite cover. The bomb was held in a bearing-mounted collar which could be rocked or rotated by a pulley arrangement through the access tubes. This feature was not utilized in the present experiments because of subsequent chemical side reactions. The bomb calorimeter system (hereafter designated CMC-1) was calibrated using a benzoic acid sample from the National Bureau of Standards (obsd.: 6319 =t19.0ral.g.-', NBS: 6318.3 cal. g.-I). The thermistor temperature measuring and recording system w a ~essentially identical to that previously described.$ However, the thermal range was expanded somewhat to accommodate larger heats (up to 90 cal. per mv.). The firing system was of a conventional type and firing energy corrections for the fuse wire (2 in. of #32 gage platinum) and the fuse ( I in. of #SO cotton thread) under the combustion conditions were determined to be 4.4.5 0.16 C R I . Procedure.-The bomb containing the sample in a small platinum boat was assembled and flushed a few times with oxygen. The combustions all were made at a standard pressure of 30 atm. The calorimeter system first was electrically calibrated, then a combustion run was made, followed by a second electrical calibration. The average electrical equivalent from the two calibrations was used in the calculations. After the final calibration, the bomb was removed and the combustion gases slowly vented into a series of washing bottles containing known amounts of a standardized basic hypochlor te solution. The contents of the wash bottles then were analyzed t o determine the change in hydroxide and hypochlorite concentrations. On the assumption that the gases were either SO2 or SO3, these analyses along with the original weight of sulfide provided the information necessary to fix the combustion reaction. Examination of the contents of the bomb after each run demonstrated that no solid residue remained in the case of the Resz combustions, although quantities of solids varying from 1-2.5 mg., which were shown to be Re2&, sometimes were recovered from the heptasulfide combustions. A final test of the stoichiometry was made by spectrophotometric analysis of the perrhenate ion obtained by washing out the contents of the bomb. (9) J. E. McDonald, 6 4 , I345 (19601,
Vol. 66
P. King. and J. W.Cobble, J . Phys. Chem.,
*
TABLE I THEENERGY OF COMBUSTION OF Re& Sample wt. (md
Heatn
AT
25'
AE
(cal.mg. -1)
(Gal.)
1.1802 41.814 1.1702 36.814 59.201 1.1765 52.171 I . 1737 40,331 1.1707 Av. 1.1743 f O.0O1gb cal. mg.-l Standard deviation. a Corrected for firing energy. 35.43 31 -46 50.32 44.45 34.45
The combustion of rhenium disulfide takes place according to reactions 3 and 4
+ 2ReSz(c) +
2ReSdc)
Odd Odg)
-3 RezOdc)
+ 4Soz(g)
(3)
ReDdc) 3. 4SOa(g) (4)
I n this case the fraction of sulfur which forms SO, is 0.817 i0.019. The thermal data for Re% are summarized in Table TI. From this data, the average value for the energy of combustion is -254.71 + 0.73 kcal. mole-1 of Re&. Changing this quantity to the heat of combustion,1° and correcting to 100% SOs, AH40 = -294.58 f 1.13 kcal. mole-1 Re%. Using auxiliary thermodynamic data,111i2 the heat of formation of Resz becomes -42.7 =k 1.2 kcal. mole-l at 25'. IV. Discussion Volatilization Processes for ReSz(c).-By making certain e~tirnationsl3~14 of the ACp of vaporization and the entropy of the solid sulfides, it is possible to predict the vapor pressures of Res2 and Re&, a t high temperatures, Assuming for the process ReSdc) = Re(c)
+ Sdg)
(5)
(10) The Wnshburn corrections are estimated t o be small compared t o t h e experimental error. (11) U. S. Bureau of Standards, Circular 500, "Selected Values of Chemical Thermodynamic Properties," 1962. (12) G . E. Boyd, J. W. Cobble, and W. T. Smith, Jr., J . Am. Chem. Soc., '76, 5783 (1953). (13) Based upon data given in K. K. Kelley. U. 5. Bur. of Mines Bull. 584 (1960). (14) K. K. KeLey, U. S. Bur. of Mines Bull. 406 (1937)
May, 1962
HEATSOF COMBUSTION OF Resz AXD Re&
TABLE I1 THEENERW OF COMBUSTION OF ReSz AT 25" Sample wt.
(me.)
AE (cal. g. -1)
IIeat" (oal.)
65.984 57.089 6%. 463 62.336
1.0169 1.0244 1.0104 1.0162 Av. 1.0170 I 0.O02gb cal. g.-l b Standard deviation. a Corrected for fii4ng energy. 64.89 55.73 61.82 60.35
is - 1.5 between 0 and 298'K., AHoO for the vaporization process is calculated to be 73.2 kcal. mole-'. If AT, for the same process over the temperature range 0 and 1400-1500OK. is -3.0 and AS6 a t 298°K. is 42.5 gbs. mole-1, then it follows -log
Pbtrn) =
15'6 T 'Os
+ 2.16log T - 15.5
(6)
793
evident that in this region of the periodic table, the direct molecular volatilization may contribute significantly to the total pressure. Consequently, the heats of formation commonly calculated for these sulfides, particularly for OsSz and RuS2, from total pressure measurements alone may be subject to some errors. Vaporization Process for RezS7(c).-The thermal data obtained on Re2S7(c)are consistent with the observations on the rapid decomposition of RezS7(c) into ReSz(c) in the 700-800°K. temperature range. Assuming that the decomposition can be represented as RezST(c) --+-
2ReSp(c)
+ 3/2s~(g)
(9)
AH0 a t 298OK. is 67.4 kcal. mole-' from the thermal experiments. Estimating the entropy of Re& to be 48 gbs., and reasonable AC, v a l u e ~for ~~?~~ eq. 9 it follows
This equation should be valid u p to 15OOOK. -log P(atrn) = 14" lo3+ 2.01 log T - 23.0 (10) However, vapor pressures calculated from equation 6 are not in agreement with the data of Juza and This equation predicts AHoO = 67.9 kcal. mole-' Biltz,s by about a factor of -10 (see Table 111). and vapor pressures of 0.01 and 3.2 mm. at 700" One concludes that either the vaporization process and 800°K., respectively, and therefore is in for Res2 is not wholly described by eq. 5, or that agreement with the chemical behavior of Re2&. the direct vapor pressure measurements are in error. I n this respect, it was reporteds that the TABLE IV quartz container used in the va,por pressure meas- THERMOCHEMICAL DATAFOR SELECTED TRANSITION ELEurements showed evidence of chemical attack. MENT SOLID SULFIDES AT 25'" If volatile components (such as SO2) were pro-AH@ - AFfO SQb Compound (kcal. mole-') (koal. mole-') (gbs. mole-1) duced, . . then the observed vapor pressures would be too high. TcS~ 53.5 51.6 17 It also is possible to estimate the effect of molec66.0 62.8 20 TcS, ular vaporiLation of ReS2(g) 73.5 69.4 21 TcS8.s
13 1.1 0.1 1383 1462 55 4.2 0.6 1498 96 7.3 1.2 Experimental vapor pressures of Jusa and Biltz, ref. 6. Calculated from Calculated from thermal data, eq. 6. eq. 8.
39.2 38.5 12 47.0 44.1 12.4 (18) 12 37.2 36.2 40.8 39.9 16 42.5 41.1 18 35.8 34.4 __ 12 39.0 37.1 18 42.7 __ 41.5 20 49.8 47.2 23 53.95 50.5 24 oss 26.4 25.4 12 37.4 05s; 37.2 22.3 (20) ossa 42.6 40.3 23 OSSd 48.0 44.3 26 IrS 23.0 21.7 12 15.1 28.1 27.9 IrS1.t 3__ 10. IrSe' 31.5 __ 25.6(20) 12 0 20.2 18.6 PtS' PtSZC 26.6 24.3 17.3(20) a Underlined values are from experimental measurements. In general, entropy values have been estimated from additivity rules and comparison with corresponding oxides. When experimental values have been available, the estimated value is given parenthetically, and serves to indicate the magnitude of errors involved in both the experimental determinations and estimation of the entropy. From the compilation of K. K. Kelley, ref. 14; see, however, text for possible sources of error due to molecular volatilization.
The conclusion from the vapor pressures given in Table I11 is that the original vapor pressure data in reference 6 are not correct. However, it also is
Summary of Thermochemical Data.-Table IV contains a summary of the thermodynamic functions for rhenium sulfides determined in the
ReSz(c) = Res&)
(7)
I t was noted6 that ReSz only slowly sublimed a t 1 0 0 O O . This probably indicated that the partial pressure of ReS2(g) at this temperature was about 10-5 atm. AS7 = 45 is probably a reasonable estimate for the vaporization process and although there is very little data upon which to make reasonable heal, capacity estimates, A C p 7 = 0 will not cause SL serious error. Therefore, the chemistry of the situation enables one to predict that AHoo = 86.5 kcal. mole for the molecular vaporization given in eq. 7, and
This equation also predicts vapor pressures too low to account for the data of Juza and Biltz.6 The situation can be summarized in Table 111. TABLE I11
HIGHTEMPERATURE VAPORPRESSURES FOR Resz T,OK.
Poh.d.,a
mm.
P d o d . , b mm. SZ(9)
Pcalcd..c
mm.
ReSk)
RuS RuSzC RhS RhSi.6 RhSe PdS" PdSz Res2 Ress Res3.&
c
_
-
794
ROGERc. M I L L I K 4 N
present research. These data also make it possible to estimate similar functioiis for neighboring sulfides which have not been available previously. Some other values available in the literature also
Vol. 66
have been included for convenience and reference. Acknowledgment.-The authors are indebted to Professor Alan W. Searcy far his helpful comments on the manuscript.
NON-EQUILIBRIUM SOOT FORMATION I N PREbTIXED FLAMES BY ROGER C. MILLIKAN General Electric Research Laboratory, Schenectady, N . Y. Received August I?’, 1961
Soot deposition has been studied in a series of large, flat, premixed ethylene-air flames burning on porous metal burners. The mixture ratios bracketed those at which soot first appears. Using combined optical and probe sampling techniques, spatially resolved measurements were made of C2H2, CH,, OH, and soot concentrations as functions of burned gas temperature. It was found that acetylene and methane (in smaller amounts) are produced in the reaction zone and survive to form about 1% of the burned gas. The pyrolysis of these hydrocarbons competes with an oxidation reaction involving OH. The bdance between these reactions determines whether or not soot is set free. The dark space betrsTeen the reaction and soot zones is a region of high oxidizing power due to the OH concentration in excess of equilibrium. The critical mixture ratio a t which soot appears is temperature dependent. This temperature dependence leads to a value of -34 f 10 kcal./ mole for the difference in activation energy between the soot oxidation and formation reactions.
Introduction Soot formation occurs in many premixed flames when it should not-if chemical equilibrium were attained. The onset of soot deposition in certain ethylene-air flames can be observed for mixtures with a carbon to oxygen atom ratio as low as 0.58. If equilibrium prevailed, soot would not be set free until this ratio exceeded unity. That is, even though there is more than enough oxygen present to oxidize all the carbon to CO, this does not happen: Instead soot is liberated. Such non-equjlibrium behavior was established by Street and Thomas1 for premixed flames of a variety of fuels burning in air. This aspect of soot formation is but one of the puzzles2which adds scientific interest to a problem of obvious industrial importance. The purpose of this work was to find an explanation for the occurrence of non-equilibrium soot deposition, and to elucidate the mechanisms involved. Theories of soot formation are many. They have been well summarized,i.2 and will not be recounted here. The difficulty lies in the lack of sufficient experimental data on a well characterized system to limit the possibilities to one mechanism. We have chosen for study a set of large, flat ethylene-air flames vhose mixture ratios bracket those at which soot first appears. By using a combination of optical and probe sampling techniques, an experimental description of these flames has been obtained which is complete enough to severely limit speculation on the mechanisms of soot formation. A key point in the present work has been a determination of the temperature depeiidence of soot formation. This turned out to be an important factor which has not been recognized in earlier work. Experimental The flames studied were burned on porous metal burners of the type developed by Kaskan.8 Burners used ranged (1) J. C. Street and A. Thomas, Fuel, 81, 4 (1956). (2) A. G . Gaydon and H. G. Wolfhard, “Flames, Their Structure Radiation and Temperature,” Second ed., Chapman and Hall, Ltd., London. 1960, pp. 175-209.
from a 2.5 X 5 em. rectangular one to circular ones 7 cm. in diameter. All were of the compound design so that the flame could be shielded from its surroundings by a sheath of Nz ga;. The use of such porous metal burners aided this work in two ways They gave large flat flames which could be probed optically with good spatial resolution. More important is that on such burners the burned gas temperature becomes a variable independent of mixture ratio. Consequently we have been able to observe the temperature dependence of soot formation at constant over-all composition. Since the gas flow is normal to the burner a t a calculable velocity, the scale of distance from the burner can be converted to a time scale. For all the flames considered here this conversion is of the order of 1 msec. per mm. Measurements have been made of burned gas temperature, GH2, CH,, OH, and soot concentrations. In addition, the critical mixture ratio a t which soot luminosity appears has been determined. Most of the techniques of measurement have been described previously. The following is only a brief account of the procedures used. A. Apparatus.-The ethylene was of C.P. grade. Air was supplied from a 2000 p.s.i. reservoir and compressor. Both were metered with calibrated critical flow orifice meters to within 0.5%. Details of the narrow beam optical setup have been given.*r6 B. Flames.-The ethylene-air flames used all fell within the region defined by the following ranges of these variables: mixtureratio 0.5 < atomic C/O < 0.7, linear gas velocity 7 < uZ5 < 20 cm./sec. NTP, temperature 1600 < T < 2000 K. The characteristics of some of the indivldual flames used are given in Table I. Noted there for each flame is the atomic C/O ratio +, the linear gas velocity entering the flame, vTg, and the maximum burned gas temperature. C. Temperature Measurements.-Gas temperatures were measured in one case by a fine wire thermocouple. Such couples soon became brittle and broke when exposed to these reducing flames. Subsequent measurements were made by the sodium &line reversal m e t h ~ d . They ~ are believed accurate to f 5 0 ” K . Soot luminosity and scattering from even the richest of our flames were so feeble that no errors in the temperature measurements arose from that source. D. Concentration Measurements.-The OH radical concentration was determined by the ultraviolet line absorption method of Kaskan.6 Since the level of OH was low, it was necessary to pass the light beam twice through a 7 cm. (3) W. E. Kaskan, “Sixth Symposium (International) on Combustion,” Reinhold Publ. Corp., New York, N. Y., 1957, p. 134. (4) R. C . Millikan, J. Opt. Soc. Am., 61, 535 (1961). (6) R . C. Millikan, Combustion and Flame, 6, 349 (1961). (6) W. E. Kaskan, ibid., 2, 229 (1958); J . Chem. P h w , 29, 1420 (1968).
