1785
Reaction of Ammonia and Sulfur Dioxide
Thermodynamics of the Reaction of Ammonia and Sulfur Dioxide in the Presence of Water Vapor' Ronald Landreth, Rosa G. de Pena, and Jullan Helcklen' Departments of Chemistry and Meteorology and Center for Air Environment Studies, The Pennsylvania State University, University Park, Pennsylvania 16802 (Received August 22, 1974: Revised Manuscript Received May 27, 1975) Publication costs assisted by the Environmental Protection Agency
The reaction of "3, S02, and HzO vapor to produce particulate matter was examined at 14, 24, 35, and 44'. At the three highest temperatures a white solid is produced consisting of two NH3 molecules, one SO2 molecule, and one H2O molecule. Infrared studies have shown the solid to be (NH4)zSOs. Thus the reaction is 2NHs(g) SO2(g) HZO(g) s (NH4)2S03(s).The enthalpy and entropy of the reaction are -53 kcal/ mol and -113 cal/(mol deg), respectively. There is about f200h uncertainty in both values, so that the data are consistent with the enthalpy of -60.9 kcal/mol for the normal solid form of (NH4)2S03. At 14', there is less HzO in the solid, and presumably some (NH4)2S205 is produced.
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Introduction
Due to the presence of ammonium sulfate in the atmosphere and the strong desire to determine how it is formed,2 the reaction of ammonia, sulfur dioxide, and water vapor is of importance. The reaction of SO2 and "3 under anhydrous conditions has been reported previously.2a,3,4The history of the water reaction is closely related to that of the anhydrous reaction and was reviewed by Scott et a1.2aFriend et aL5 could not produce particles in a similar experiment but their gas pressures were much smaller than used in the other investigations. Neither the chemical structures nor the thermodynamics of the adducts are precisely known. McLaren et aL4 have reported a AH of reaction of -34 kcal/mol for the reaction of SO2 and NH3 in wet nitrogen and -45 kcal/mol in 0 2 . No temperature ranges for the experiments were stated. ScargilP identified two products from the NH3-SOz-HzO reaction as (NH&S03 and (NH4)2S205 using chemical analysis and infrared spectroscopy. The dissociation constants of both compounds were measured by a transpiration method from 273 to 298'K. The values found for A S and AH for the two sublimations were 152.2 cal/(mol deg) and 64.8 kcal/mol, respectively, for (NH4)2S03 and 191.7 cal/(mol deg) and 81.0 kcal/mol, respectively, for (NH4)zSz05. In this paper, we have examined the reactions from 14 to 44O between Son, "3, and H2O vapor in Nz. The stoichiometry and thermodynamic functions are reported. Experimental Section The reaction was carried out in a 2.9-1. cylindrical cell 149 cm long and 5 cm in diameter. The cell was capped a t each end by a quaptz window. Along the length of the interior of the cell was a narrow tube with holes every 10 cm through which the gases entered to allow thorough mixing. Temperature regulation was made with a constant-temperature water bath surrounding the cell. During the experiments SO2 absorption at 3340 b; was monitored photometrically. The 3340-b; radiation was from
* T o whom correspondence should be addressed a t the Department of Chemistry, The Pennsylvania State University.
a Hanovia Utility 30620 Hg lamp which had been passed through a Jarrell-Ash 82-410 monochromator set at 3340 A. The light beam was split and half of it passed through the reaction vessel to a RCA 935 phototube; the other half impinged directly on a second RCA 935 phototube. When the reaction vessel was empty, the signals from the two beams were balanced by a Wheatstone bridge. Thus with SO2 in the cell the difference in beam signals was a direct measure of the SO2 pressure. The SO2 and NH3 gases were from the Matheson Co. and the N2 gas was from Phillip and Sons Inc. Before use, the N2 was passed through a trap packed with glass wool at -196' to remove water and impurities. Both SO2 and NH3 were twice purified by distillation from -98 to -196'. Sometimes further distillation was performed from -98 to -130'. Distilled water was used after being degassed a t various temperatures from -10 to -196'. The gases were handled in a mercury and grease-free vacuum line containing Teflon stopcocks with Viton "0" rings. The vacuum line could be evacuated to less than 1 X Torr. The gases were introduced into the cell and the experiments performed as follows. Water was the first gas introduced and was allowed to stand for 5 min to obtain a constant-pressure reading. A known prepared mixture of SO*, "3, and Nz (the pressures of SO2 and NH3 were below the point at which they react in the absence of Hz03) was added in increments to the water in the cell. All gas pressures were read on a calibrated National Research Corp. 820 alphatron vacuum guage. After each incremental addition of the S02-NH3-Ng mixture, the SO2 absorption increased and then remained constant if particles were not produced. Ultimately the reactant pressures became sufficiently high so that particles were produced and the SO2 absorption decreased slowly over a period of a few minutes after the sharp increase accompanying the incremental addition. This peak was taken as the nucleation pressure. Then the contents of the cell were expanded by 4.5% into a small evacuated volume connected by a stopcock to the reaction vessel. If the particles sublimed to restore the SO2 pressure, then the expansion was repeated until the SO2 pressure could not be restored. At this point all of the solid had just sublimed, and the NH3 and HzO pressures could be computed from their The Journal of Physical Chemistry, Vol. 79, No. 77. 1975
1706
Landreth, de Pena, and Heicklen
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initial pressures and the expansion factor. These pressures were taken as the equilibrium vapor pressures. When Nz was used as a diluent, the ratio of its pressure to that of SO:! varied from 77 to 292 and to that of NH3 from 81 to 146. Reactions without nitrogen were done in a similar manner as above except that water was not necessarily added first. The order of addition of the reactants did not affect the results.
Results When SOz, "3, and Hz0 gases were mixed, a white solid was produced instantly if the gas pressures were sufficiently high. The reaction was reversible and when the gas pressures were reduced, the solid sublimed back to its gaseous reactants. When Nz was used as a diluent there was a slight delay and a noticeable slowing of the reaction. Gas pressures could be measured to within 7-10% accuracy. T o be sure that the reaction was not affected by the monitoring light, runs were done with the light off using it only to make a measurement. The results were the same as with continuous monitoring. The pressures of SO:! and H20 in equilibrium with the solid a t 24O are shown graphically in log-log plots in Figure 1 for three mixture ratios of [SO:!]/[NH3]. The plots are linear, and as the ratio is lowered, the lines lie successively lower. The results in N2 are the same as without it. All the plots of SO2 pressure vs. H20 pressures are linear and can be fitted with a line of slope -Ih. The equilibrium pressures of SO:! and HzO at 14, 24, 35, and 44O are shown graphically in log-log plots in Figure 2 for a 1:1 ratio of [SO:!]/[NH3]. As the temperature is lowered, so are the points on the plot. At the three higher temperatures, the points can be fitted to a line of slope -3s. However, at 14O the plot cannot be fitted to a line of slope -lh but is best fitted by a line of slope -yS. The supersaturation ratio needed for particle nucleation (i.e., the absorbance of SO2 needed to form particles divided by the absorbance of SO:! a t equilibrium) was obtained in each experiment. A plot of this ratio vs. H20 pressure in any series showed considerable scatter. For some of the series there was no trend in the data. For other series, there may have been a slight downward trend as the Hz0 vapor pressure was raised. However the scatter was sufficiently large so that this trend may not have been meaningful. The average values for the six series of runs are tabulated in Table I. There is considerable scatter, but no obvious trend in the data. The average value for the six series is 1.50 f 0.10.
+ yHnO(g) z=
(NHd,SOz(HzO)y
A
262
PRESENT
0
IO0 0.333 1.00
PRESENT PRESENT ABSENT
0
c
1.0
0.01.
1
J
I
Flgure 1. log-log plots of the vapor pressures of SO1 vs. the final
H 2 0 pressure in equilibrium with the solid at 2.62:1, 1:1, and 1 3 mixtures of SO2 and NH3 at 24'.
SYMBOL
0 0
A
t
TEMP,'C 14 24 24 35
N2 PRESENT PRESENT ABSENT PRESENT PRESENT
44
1.0:
a
.
e -
- . N
m 0
u
-
-
0.1
r
0.01 0.01
Discussion The heterogeneous equilibrium under study is wNHa(g) + SO&)
N2 --SYMBOL [SO21 / [NH,]
I
I
0.1
10
I 10.0
[HeO], TORR
(1)
Flgure 2. log-log plots of the vapor pressures of SO2 vs. the final H 2 0 pressure in equilibrium with the solid at 14, 24, 35,and 44'. A 1:l ratio of [SO2] to ["a] was used in all reactions.
The equilibrium expression is
K = l/[NH3]" [SOz][H20]Y
(2)
where K is the equilibrium constant a t a given temperature. Equation 2 can be rearranged to give log [SO,] = - X log [NHs] - y log [H:!O]
- 1% K
(3)
If [HzO] is held constant, a log-log plot of [SO:!] vs. iNH3] will give a plot of slope - x . Such a plot is shown in Figure 3 for [HzO] = 1 Torr. The [SO:!] values were taken from Figure 1, and [NHs] was computed from the known value of The Journal of Physical Chemistry, Vol. 79, No. 17, 1975
the [SOZ]/[NH~] ratio. The plot is linear with a slope of -2; thus x = 2. Now if we consider those series of runs with [SOz]/[NH3] = 1.00, then eq 2 can be rearranged to log [SOZ] =
-log K
- JJ log [H20] l+x
(4)
The log-log plots of [SO21 vs. [HzO] shown in Figure 2 have slopes -)s at 24,35, and 44O and -)bj a t 14O. Thus y = 1 a t
Reaction of Ammonia and Sulfur Dioxide
1787
TABLE 11: Equilibrium Constants for the Reaction 2NH3(g) + SOz(g) + HzO(g) s ( N H ~ ) Z S O Z - H Z O ( S ) ~
TABLE I: Average Values of the Supersaturation Ratios Needed for Nucleation Temp, "C
ISO;!I/[",I
Supersatn ratio ~~
1.00 2.62 1.00 0.33 1.00 1.00
14 24 24 24 35 44
1.51 1.39 1.62 1.63 1.33 1.54 Av 1.50 i 0.10
24 35 44
~
~
+1
460 25.5 1.54
*1 *1
a l n K(atm-4) = A S / R A H = -53 kcal/mol.
-
154 8.51 0.51
AH/R7', AS = -113 cal/(mol deg);
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t
o'osd.03
0.b5 '
I
'
' 0.I
1
0.5
[NH3],TORR
Figure 3. log-log plot of [SO,] vs. [NH3] in equilibrium wlth 1 Torr of H 2 0 and the solid at 24'.
the three highest temperatures and 'h a t the lowest temperature. The formula of the solid a t 24' and above is (NH3)2S02. H20. From the intercepts of the plots in Figure 2, the equilibrium constants can be evaluated, and they are listed in Table 11. The equilibrium constants in Table I1 are plotted on the van't Hoff plot, Figure 4, since In K(atmW4)= AS/R - m / R T (5) A straight line can be drawn reasonably through the points. The slope gives AH = -53 kcal/mol and the intercept gives A S = -113 cal/(mol deg), both with an uncertainty of is considerably lower than about f2W. Our value of the value of -34 kcal/mol reported bjr McLaren e t a1.: and our thermodynamic functions are somewhat higher than those of Scargill.6 Also shown in Figure 4 is the plot based on Scargill's thermodynamic functions. I t can be seen that, in spite of the differences in thermodynamic functions, which are quite inaccurate because of the small temperature ranges involved, the equilibrium constants from the two studies are in excellent agreement. Infrared measurements in our laboratory have shown that the solid ( N H ~ ) Z S O ~ - is H ~(NH4)2S03.' O The enthalpy of formation of this compound is known to be -211.6 kcal/ For the gases the enthalpies of formation, AHo~g8 (kcal/mol), are:9 "3, -10.970; H20, -57.798; SOz, -70.947. Thus the enthalpy for the reaction 2"3(g)
+ SOn(g) + HzO(g) s ("4)2SOd~)
is -60.9 kcal/mol, a value intermediate to those found here
10' IT,
OK-'
Figure 4. Semilog plot of the equilibrium constant K vs. the reciprocal temperature for the reaction PNHs(g) SOdg) H2O(g) s
+
+
("3)2S02"20(~).
0'
0.01 0.01
1.0
0.10
5.0
[HzO], TORR
Figure 5. Comparison of 14' data for a 1:l mixture of SO2 and NH3 with the equilibrium vapor pressure curves expected for (NH4)2S03 and (NH4)&05 from Scargill' and for ("&SO3 extrapolated from the higher temperature data of this study.
and those found by Scargill; but within the uncertainty limits of both measurements. The formula for the solid a t 14' is not established. The data are replotted in Figure 5 along with the expected lines The Journal of Physical Chemistry, Vol. 79,No. 17. 1975
1788
Nakashima, Kume, and Mataga
for (NH4)&03 and (NH4)2S206 from Scargill’s results and the extrapolated plot from our higher temperature data for (NH&S03. The solid formed should correspond to that with the lowest equilibrium vapor pressure. The data, though badly scattered, can conform to either our extrapolated plot for (NH4)2S03 or Scargill’s plot for (NH4)2S205. Probably both compounds are present and this may account for the scatter. Scargill’s plot for (NH4)2S03appears to lie too low to be consistent with the data.
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Acknowledgment. The authors wish to thank Drs. K. Olszyna and L. Stockburger for their assistance. This work
was supported by the Environmental Protection Agency through Grant No. 800874, for which we are grateful. References and Notes (1) CAES Report No. 369-74. (2) (a) W. D. Scott, D. Lamb, and D. Duffy. J. At#& Sci., 26, 727 (1969): (b) C. S.Kaing, D. Stauffer, and V. A. Mohnen, Nature (London), Phys. Sci., 244, 53 (1973). (3) R. Landreth, R. G. de Pena, and J. Heicklen, J. fhys. Chem., 78, 1378 (1974). (4) E. McLaren, A. J. Yencha, J. M. Kushnir, and V. K. Mohnen, Tellus. 26, 1 (1974). (5) J. P. Friend, R. Leifer, and M.Trichon, J. Atmos. Sci., 30, 465 (1973). (6) D. Scargill, J. Chem. SOC.A, 2461 (1971). (7) I. C. Hisatsune and J. Heickien, unpublished results, 1974. (8) Nat. Bur. Stand. (US.), Tech. Note, No. 270-3 (1968). (9) “JANAF Thermochemical Tables”, Dow Chemical Co., Miland, Mich., 1966.
Electronic Excitation Transfer between the Same Kind of Excited Molecules in Rigid Solvents under High-Density Excitation with Lasers Nobuaki Nakashima, Yujl Kume, and Noboru Mataga’ Department of Chemistry, Facutfy of Engineering Science, Osaka University, Toyonaka, Osaka, Japan (Received November 1, 1974; Revised Manuscript Received June 3, 1975)
--
By means of high-density excitation with a Q-switched ruby laser, the electronic excitation transfer between excited pyrenes in the fluorescent state (SI SI So S,) and that between excited perdeuteriophenanthrenes in the phosphorescent state (TI TI SO + T,) in rigid solvents at 77OK have been observed by directly measuring the luminescence decay curves. Comparing the theoretical decay functions derived in the present work with the experimental ones, it has been proved that the excitation transfers in these systems are due to Forster’s dipole-dipole coupling mechanism.
+
Introduction
Bimolecular quenching of fluorescence due to the interaction between excited aromatic molecules was first found by Tolstoi and Abramov’ as well as by Bergman et aL2 I t has been confirmed by many other investigations3& following the above ones that this phenomenon is occurring in general under high-density excitation. This bimolecular quenching seems to consist of the following two processes: (1) mutual approach of excited molecules to a distance where they can interact with each other; (2) an actual quenching process due to the interaction between the excited molecules. However, in many cases, including the excited molecules in solution and excitons in crystals, the actual quenching process is masked by the rapid “bimolecular collision”. The bimolecular collision process has been confirmed for many cases, and its rate constant has been determined. Such a study has proved to be a powerful mean especially for the investigation on the diffusion of excitons in crystal^.^ As for the mechanisms of the actual quenching process, Abramovl has insisted that the overlapping between the S, SIabsorption and SI SOfluorescence bands is essential for the quenching process to occur. Babenco et al.6 as well as Kobayashi and Nagakura6 have also insisted that
-
-
The Journalof Physical Chemistry. Vol. 79, No. 17, 1975
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the quenching occurs owing to the electronic excitation transfer between the excited states due to the Forster mechanism.’ However, there has been no direct experimental proof for that so far as we know. In order to elucidate the mechanism of this bimolecular interaction, we have derived decay functions of excited states for the systems fixed in rigid solvents assuming the excitation transfer between excited states due to Forster’s very weak interaction mechanism of the dipole-dipole coupling7 and have compared them with observed decay curves of excited molecules fixed in rigid organic solvents a t 77OK under high-density excitation with a Q-switched ruby laser. We have proved that the mechanism of the bimolecular quenching in these excited molecules is Forster’s dipoledipole coupling. Experimental Procedure The second harmonic of a Q-switched ruby laser was used for exciting the rigid solution. The ruby laser system was the same one as used b e f ~ r e .Its ~ , maximum ~ output a t 3472 %, was 2 x 1017 photons/pulse and the pulse width was ca. 15 nsec. Fluorescence decay curves were observed by using a Nalumi R21 monochromator, an RCA 1P28 photomultiplier, and a Tektronix 585A oscilloscope. The re-
