R. Landreth, R. G. de Pena, and J. Heicklen
1378
Thermodynamics of the Reactions ( N H 3 ) n * S 0 2 ( ~2 ) nNH3(g) iS02(g)' Ronald Landreth, Rosa G. de Pena, and Julian Heicklen" Deparfments of Chemistry and Metrology and Center for Air Environment Studies, The Pennsylvania State University. University Park, Peilnsylvania 76802 (Received October 29, 7973; Revised Manuscript Received April 17. 7974) Publicahon costs assisted by the Environmental Protection Agency
Anhydrous SO2 and NH3 were allowed to react to form their solid adducts. Equilibrium vapor pressures for each gas were measured at 5, 15, 24.5, 35, and 45". The reaction is NH3-S02(s) ~1 NH3(g) + SO2(g) with In K(atm2) = AS/r - A H / R T where K is the equilibrium constant, AS = 15.3 cal/mol OK, and AM = 9.5 kcal/mol to within *lo% uncertaiy. The equilibrium was unaffected by the presence of excess 0 2 or N2. With excess NH3 at the lowest two temperatures, there was thermodynamic evidence for the production of the ( N H ~ ) z . S Oadduct. ~ For the reaction (NH3)2.SOz(s) ~t 2NH3(g) + SO2(g) the thermo82 cal/mol "K and AH 23 kcal/mol with at least dynamic functions are crudely estimated to be A S k20% uncertainty.
-
Introduction The anhydrous reaction between NH3 and SO2 has been known for a very long time, and the history of this reaction has been reviewed by Scott, et aL2 Above lo", the reaction between NH3 and So2 produces a 1:l adduct which is a yellow solid. Below lo" an additional white solid is produced which is the adduct of two molecules of NH3 and one molecule of S02. Both reactions are reversible, and as the pressure is reduced, the solids sublime to form NH3 and SOz. Neither the chemical structures nor the thermodynamics of the adducts are precisely known. Scott, et a1.,2 measured the vapor pressure of the solids between -10 and -70". From their da.ta they estimated the enthalpy of sublimation to be -1 and -15 kcal/mol, respectively, for the 1:l and 2:l adducts. McLaren, et report a heat of reaction of -30 kcal/ mol for anhydrous NH3 with SO2 both in the presence and absence of 0 2 . They did not specify their temperature range and it is not known whether this value refers to the 1:l or 2:l adduct. Friend, et ~ l . worked , ~ at very low pressures and found no evidence for solid formation in the absence of ultraviolet radiation. In this paper we have examined the reaction in the temperature range 5-45' and at partial pressures very much larger than those employed by Friend, et aL4 We show, from thermodynamic arguments, that the solid is the 1:1 adduct in this temperature range, except with excess NW3 at the lowest two temperatures. Furthermore, AH and A S are determined to be 9.5 kcal/mol and 15.3 cal/mol "K, respectively, for the reaction NA,.SO,(s)
NH,(g)
+ SO,(g)
Experimental Section The reaction was carried out in a 2.7-1. cylindrical cell 137 cm long and 5 cm in diameter. The cell was capped at each end by a 2-mm thick quartz window to allow passage of ultraviolet light. Inside and along the length of the cell was a small tube with holes every 10 cm through which the gases entered to allow thorough mixing of the gases. Temperature variation was accomplished by changing the The Journal of Physical Chemistry, Vol. 78, No. 14, 1974
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temperature of a constant temperature water bath around the cell. Particle formation was determined by light scattering. The light from a Hanovia Utility 30620 ultraviolet quartz lamp was passed through a Jarrell-Ash 82-410 monochromator set at 3660 A and through the cell. The light was detected by a RCA 935 phototube and the signal sent through a General Radio Co. 1432-M resistance box to a Texas Instruments Inc. 1-mV recorder. S02, NH3, and 0 2 gases were from the Matheson Co., and N2 gas was from Phillip Wolf and Sons Pnc. Before use, the N2 was passed through a trap packed with glass wool a t -196" to remove water and impurities. The same was done with 0 2 at -98". Both SO2 and NHs were twice purified by distillation from -98 to -196". The gases were handled in a mercury- and grease-free vacuum line containing Teflon stopcocks with Viton "0" rings. The vacuum line was connected to the reaction cell and the presTorr. sure could be reduced to 1 x The gases were introduced into the reaction cell as follows. For experiments with SOz and NH3 only, one gas was added to the cell and the pressure read on a National Research Corp. 820 alphatron vacuum gauge. The alphatron gauge had been calibrated by an oil monometer for each gas. After having closed off the cell and frozen down the excess gas, the second gas was added in increments until light attenuation was noticed on the recorder. At 5" light attenuation was not seen on the recorder because the particles settled on the wall. Particle formation was determined visually. For runs with 0 2 or N2 present, one reactant was placed in the cell, and the other reactant diluted in N2 or 0 2 was then added.
Results Anhydrous SO2 and NH3 gases were mixed, and a yellow solid was produced instantly if the gas pressures were sufficiently large. The reaction was completely reversible: when the gas pressures were reduced sufficiently, the solid disappeared. When one of the reactants was diluted in 0 2 or Nz, there was about a 30-sec delay before particle production, if particles were produced at all. In these experiments, we waited 5-10 min after each addition to be sure that particles either were or were not produced.
Thermodynamics of t h e Reactions between NH3 and SOa
t I
0 24 5 i I * C e 24 s i. I T 0 15 i. I'C
I"
0
111
o2
8"
o*
0 0
'5
* IEC
5 f IPC 5 f IT
1379
\
02 hi,
U L 2 - U 31
30
I
, L L . . L L L u ,
10 IN.,],
1 100
Twr
33
32
35
34
36
+
Figure 1. Log-log plots of the vapor pressures of SO;!and NH3 in equilibrium with thelr solid product at 5, 15, 24 5, 35, and
SOz(g).
45"
TABLE I: Equilibrium Constants for the Reaction NHs.SOz(s) NHa(g) SO,(g).
Experiments were done in which the lower pressure reactant was placed in the reaction vessel and then the other reactant, either alone or diluted in N2 or 0 2 , was added. Several additions of the second reactant were made until the solid was produced. The point at which solid appeared was taken as the equilibrium condition. For some runs, the total pressure was then reduced by 10% and the solid disappeared. When 0 2 or N2 was used as a diluent, the ratio of its pressure to the reactant pressure was varied between 2.45 and 12.1. The pressures of the gases could be ascertained to within 5% accuracy for pressures IO0 Torr. To be sure that the reaction was not photocheniically induced by the monitoring lamp (SO2 absorbs weakly at 3660 A), runs were done with the lamp off while the mixture was equilibrating. The lamp was then turned on just to take the measurement. The results were the same as with continuous exposure. The equilibrium conditions, i. e., the pressure limits at which solid is just formed, are shown graphically in loglog plots in Figure 1 at five temperatures between 5 and 45". As the temperature is raised the data points lie successively higher. The results in the presence of diluent gases are identical with those with no diluent added in agreement with McLaren, et aL3 The plots of SO2 pressure us. NH3 pressure at the three highest temperatures are linear and can be fitted with a slope of -1.00. At the two lower temperatures, the data points are fitted with a line of slope -1.00 in excess SOz. However, in excess NH3, deviations OCCUF and the data points lie below the linear extension of the results in excess S 0 2 .
Discussion The equilibrium of interest is the heterogeneous one nNH3.S02(s)
q*
S02(g) 4- nNH3(g)
where n is either one or two. The equilibrium expression for the reaction is
37
io3/ T , OK-' Figure 2. Semilog plot of the equilibrium constant K vs. the reciprocal temperature for the reaction N H ~ . S O ~ ( Si=) NH3(g)
+
Temp, OC
45 I 1 35 I 1 24.5 I 1 15 1 1 5 1 1 a Ln K(atmz) = A S / R kcal/mol.
K , Torr2
IOOK, atmz
425
736 386 243 149 83
223 140 86 48 - AH/RT, A S
= 15.3 cal/mol
OK,
AH = 9.5
where K is the equilibrium constant at each temperature. Consequently, a log-log plot of the equilibrium pressure of SO2 us. that for NH3 should be linear with a slope of -n. The plots in Figure 1 clearly show that n = 1 for the three higher temperatures under all conditions and for the lower two temperatures in excess S02. The equilibrium constants are the products of the equilibrium reactant gas pressures, and they are listed in Table I. They range from 48 Torr2 at 5" to 425 Torr2 at 45". The equilibrium constant, K, is related to the entropy, AS, and enthalpy, AH, of reaction through the wellknown expression In K(atm2) =
AS/R - AN/RT
where T is the absolute temperature. Figure 2 i s a semilog plot of K us. 1/T. The plot is linear. From the slope AH is found to be 9.5 kcal/mol, and from the intercept A S is found to be 15.3 cal/mol "K. These values have a -+IO% uncertainty. The falloff from linearity in Figure 1 for runs with excess NH3 at the two lowest temperatures indicate production of the adduct of two NH3 molecules with one SO2 molecule. Thus the slope becomes steeper and approaches two. From the very limited data the equilibrium constants atms) a t 5" are estimated to be 225 Torr3 (0.51 X atm3) at 15" for the reaction and 1030 Torr3 (2.34 X
-
(N&)~.SOLS) a 2NH3(g)
Thus A S 82 cal/mol "K and AN at least a &20% uncertainty.
+ SO,(g)
-
23 kcal/mol, with
The Journal of Physical Chemistry, Vol. 78, No. 14, 1974
Murray L. Jansen and Howard L. Yeager
1380
It should be noted that our values for A H for both reactions are greater than those reported by Scott, et ~ l . but , ~ lower than the value of 30 kcal/mol reported by McLaren, et el.^ If the value of McLaren, et al., refers to the 2:l adduct, it and our value may agree to within the large experimental uncertainty. Ack&,wEedgment, ~l~~ authors wish to thank Drs. K . Olszyna and Ihil. h r i a for their help. This work was sup-
ported by the Environmental Protection Agency through Grant No. 800874, for which we are grateful. References and Notes (1) CAES report No. 320-73. ( 2 ) W. D. Scott, D. Lamb, and D. Duffy, J. A m . Sci.. 26, 727 (1969). ( 3 ) E. McLaren, A . J. Yencha, and J. Kushnir, 1.U.G G . Symposium on Trace Gases, Germany, April 2-5, 1973. (4) J. P. Friend, R Leifer. and M .Trichon. J . ~ t m Sci. 30,465 (1973).
Conductance Measurements of Alkali Metal Trifluoroacetates in Propylene Carbonate Mtrrray L. Jansen' and Howard L. Yeager" Deparfmenf of Chemistry, The University of Calgary, Calgary, Alberta, Canada
(Received February 6, 7974)
Publication costs assisted by the National Research Council of Canada
Precise conductance measurements are reported for the alkali metal trifluoroacetates, tetrapropylammonium chloride, and tetrabutylammonium nitrate in propylene carbonate. Lithium trifluoroacetate is extensively associated, with the formation of ion aggregates at higher concentrations. The degree of association for the trifluoroacetates decreases with increasing size of the alkali metal ion. Single ion mobilities are reported for chloride, nitrate, and trifluoroacetate ions in propylene carbonate.
Introduction Conductance measurements for a variety of alkali metal and tetraalkylammonium halides and perchlorates in propylene carbonate (PC) have been reported.2 All salts were found to be essentially unassociated in this solvent of moderately high dielectric constant (64.92 at 25").3 Single ion mobilities were derived and discussed in terms of the ion solvating ability of PC. We have extended this work to study the conductance behavior of the alkali metal trifluoroacetates. Wu and Friedman4 have measured the heats of solution of these salts in PC. Results indicated that the lithium salt is extensively associated, and the formation of ion pair dimers was suggested. Conductance measurements indicated that the lithium and sodium trifluoroacetates are associated, but no equilibrium constants were calculated .4 Precise conductance measurements are reported here for these salts, in order to accurately evaluate the influence this more strongly basic anion has on association processes in PC. Association trends are discussed in terms of the relative sizes of the alkali metal ions. In addition, conductance measurements are reported for tetrapropylammonium chloride and tetrabutylammonium nitrate in order to estimate single ion mobilities for the respective anions. Trends for anion mobilities in PC are discussed. Experimental Section The purification of PC has been described previously.2 Tetra-n-propylammonium chloride (Eastman Kodak) was precipitated three times from acetone with ether and dried under vacuum at 150" for 40 hr. Analysis by silver nitrate titration: 99.7%. Tetra-n-butylammonium nitrate The Journal of Physical Chemistry, Vol. 78. No. 14, 1974
was prepared by metathesis from tetra-n-butylammonium chloride (Eastman Kodak) and silver nitrate (Engelhard Industries Ltd.) in an ethanol-water mixture. The filtrate remaining after removal of silver chloride was evaporated to yield the crude product. The salt was recrystallized three times from ethyl acetate and dried under vacuum a t 85" for 72 hr. Anal. Calcd: C, 63.12; , 11.92. Found: C, 63.15; H, 11.70. Lithium trifluoroacetate was prepared from recrystallized and dried lithium bromide (Research Organic/Inorganic Chemical Corp.) and dried silver trifluoroacetate (Eastman Kodak) in ether. The filtrate remaining after removal of the silver bromide was reduced in volume, and the precipitated salt was dried under vacuum at 60" for 36 hr, and then at 100" for 72 hr. Analysis by cation exchange and acid titration: 99.2%. Sodium, potassium, rubidium, and cesium trifluoroacetates were prepared by adding equivalent amounts of trifluoroacetic acid (Eastman Kodak) to aqueous solutions of the corresponding carbonates, followed by careful reduction in the volume of the solutions to initiate crystallization. Sodium trifluoroacetate was recrystallized three times from a 6:l dioxanemethanol mixture and then dried under vacuum at 80" for 24 hr and a t 125" for 48 hr. Analysis by cation exchange and acid titration: 99.8%. Potassium trifluoroacetate was recrystallized from dioxane and ethanol and dried under vacuum a t 130" for 72 hr: analysis 100.0%. Rubidium trifluoroacetate was recrystallized twice from a 1:1 dioxaneethanol mixture and dried under vacuum at 70" for 1 week: analysis 100.5%. Cesium trifluoroacetate was recrystallized twice from ethyl acetate and dried under vacuum at 60" for 12 hr: analysis 100.3%. The purification of the alkali metal trifluoroacetates proved to be a formi-
