J . Phys. Chem. 1988, 92, 4172-4175
4172
Enthalpies of Formation of Dinitrogen Pentoxide and the Nitrate Free Radical A. H. McDaniel, J. A. Davidson,* C. A. Cantrell, R. E. Shetter, and J. G . Calvert National Center f o r Atmospheric Research, Boulder, Colorado 80307 (Received: November 30, 1987: In Final Form: February 1 , 1988)
The enthalpy of formation for solid N205 was determined via the reaction with H20(1) and was found to be -12.91 f 0.61 kcal/mol. The enthalpy of sublimation was determined from the vapor pressure of N205 measured as a function of temperature over a range of 21 1-273 K and found to be 14.104 f 0.075 kcal/mol. From these new data and the enthalpy of reaction for N20,(g) ==NO3(& + N02(g) determined from the temperature dependence of the equilibrium constant, the enthalpies of formation for N205(g)and N03(g) are 1.19 f 0.61 kcal/mol and 15.39 f 0.75 kcal/mol, respectively. To minimize experimental uncertainties, precise measuring devices were used to monitor the temperature and pressure. Calculated enthalpies rely upon the currently reported enthalpies of formation for N02(g) and HN03(aq). The enthalpy of formation of the NO, radical reported here is significantly less than the currently accepted value.
Introduction
The oxides of nitrogen, such as NzO5 and NO3, are recognized as important species in atmospheric chemistry. NzOSacts as a temporary reservoir for NO2and NO3, and may also act as a sink for these species through hydrolysis reactions on surfaces. NO, may contribute to nighttime acid generation and depletion of organic compounds in urban polluted tropospheres.' Because of the importance of these two species in atmospheric chemistry, accurate thermodynamic properties, such as enthalpies of formation, are required to better understand the contribution of each specie in proposed reaction mechanisms. Several studies regarding the thermodynamic properties of these oxides, N205 in particular, have previously been carried out. In 1957 Ray and Ogg2 measured the AHf(N205,g)by the reaction ) measured by Ogg3 in of NzO5 with N O . The A H f ( N 2 0 5 , swas 1947 via the reaction of N205 with H20(1). The most recent vapor pressure data on N2O5 was taken by Daniels and Bright4 in 1920. These early determinations resulted in the currently accepted values for the thermodynamic properties of N2O5. Unfortunately the accuracy with which the thermodynamic properties of N 2 0 5are measured limits the accuracy of subsequent calculations involving the nitrate radical. At present there is a substantial uncertainty associated with the M f ( N 0 3 , g )reported in the JANAF Thermochemical Tables5 This uncertainty may be due, in part, to the uncertainties associated with the early measurements of N205 thermodynamic properties. Other factors, such as the temperature dependence of the equilibrium constant for the reaction
also contribute to the uncertainty in determining the heat of formation of the nitrate free radical. This study evaluates the thermodynamic properties of N2Os ) using modern calorimetric techniques. The A H f ( N 2 0 5 , swas determined through the reaction N205(s)+ H20(1)
-
2HN03(aq)
(2)
which is known to proceed stoichiometrically and is quite exothermic. In addition, the vapor pressure of N2O5 was measured directly above the crystalline compound over a range of 21 1-273 K. This vapor pressure data provide a measure of the enthalpy (1) Wallington, T. J.; Atkinson, R.; Winer, A. M.; Pitts, J. N., Jr. J . Phys. Chem. 1986, 90, 4640 and the references therein. (2) Ray, J. D.;Ogg, R. A. J . Chem. Phys. 1957, 61, 1087. (3) Ogg, R. A. J. Chem. Phys. 1947, IS, 337. (4) Daniels, F.; Bright, A. C. J . A m . Chem. SOC.1920, 42, 1131. (5) Chase, M. W. Jr.; Davies, C. A,; Downey, J. R., Jr.; Frurip, D. J.; McDonald, R. A,; Syverud, A. N. J . Phys. Chem. Ref. Data 1985, 14, 1537, 1561 (JANAF Thermochemical Tubles, 3rd ed., 1985; Supplement).
0022-3654/88/2092-4172$01.50/0
change for the sublimation of N205 and thus permit the calculation of AHf(N205,g). Recent research done by Cantrell et a1.,6 regarding the aforementioned equilibrium constant, combined with the data from this study provide the information necessary to accurately determine the enthalpy of formation of the nitrate free radical in its standard state at 298 K. Throughout this study special attention was given to the reported uncertainties in laboratory measurements and their propagation in all subsequent calculations. Experimental Section
Heat of Formation of N205.The AHdN2O5,s) was determined by measuring the AHm for reaction 2. A Hewlett Packard 2804A digital quartz thermometer enabled temperature measurements to f 0.002 "C over the range of temperatures employed in this study. To minimize heat loss and thermal drift, all experiments were performed in a 5.0-L vacuum-jacketed Dewar flask. Three liters of deionized water served as both the reaction medium and heat-transfer fluid. A stainless steel cooling coil, stirring paddle, 9 . 9 6 4 50-W resistor, and quartz oscillator thermometer were immersed in the deionized water. The Dewar flask was then fitted with a rigid foam lid to insulate the system. Before and experiments or calibrations were performed, the temperature of the entire system was cooled to approximately 278 K. Va!ves positioned outside the Dewar on the inlet and outlet to the coil allowed the temperature regulating fluid to be shut off during the experiment. The rationale of working at reduced temperatures (278 K) was to minimize the rate of thermal decomposition of the NZO5. This also served to lower the pressure of N2Os above the solid and thus minimized the excess heat evolved due to the reaction of gaseous N205. The heat capacity of the system was determined before each experiment by placing a 20.0-V potential across the 9 . 9 6 4 resistor and monitoring the temperature change over a short period of time ( 5 min). Both the voltage and resistance were measured by means of a digital multimeter. An IBM-XT computer equipped with a data collection and manipulation routine collected digital output from the quartz oscillator thermometer at 1 Hz. A linear regression routine was used to fit the temperature as a function of time for both the background temperature data, collected before and after the applied voltage, and the temperature data resulting from the applied voltage. The averaged background temperature vs time slope, obtained from the fit routine, was then subtracted from the slope resulting from the applied voltage, yielding the dT/dt due to the heat capacity of the system. The power input divided by the difference of the temperature vs time slopes gave the CJSYS) = 3.328 f 0.033 kcal/'C, averaged over 12 determinations. The uncertainty reported for the heat capacity of ( 6 ) Cantrell, C. A.; Davidson, J. A.; McDaniel, A . H.; Calvert, J. G.; Shetter, R. E. J . Chem. Phys. 1988, 88, 4991.
0 1988 American Chemical Society
The Journal of Physical Chemistry, Vol. 92, No. 14, 1988 4173
Enthalpies of Formation of N2O5 and NO3 5.5
I
TABLE I: Summary of Enthalpy of Formation Data for N205Solid from This Study mass of N205,g AT, 7.1350 7.1813 8.7180 9.8001 10.5173
h
0
e
5.3 -
E c
F
5.2
+5
-
0
200
400 600 Time ( s e c o n d s )
800
Figure 1. Plot of temperature versus time illustrating the thermal drift (background temperature slope), and temperature change for the hydrolysis reaction. N 2 0 s ( s ) was added at 220 s.
the system represents the precision associated with averaging over the twelve determinations at the 95% confidence interval. N2O5 was prepared by using a modified synthesis of Davidson et al.’ described elsewhere by Cantrell et aI.* The estimated H N 0 3 impurity was less than 1% as determined by infrared spectroscopy. While not in use, the N 2 0 5was stored at 193 K. Dinitrogen pentoxide was transferred, under vacuum, to a tared glass container at liquid nitrogen temperature. The glass container was tared while evacuated. All masses were measured with an analytical balance to f0.001%. Before weighing, the container was allowed to reach a temperature whereby external condensation of water was not a problem. Any NO2 and O2 formed by the decomposition of N205during this time was pumped away and then the container and its contents were weighed. During the time required to weigh and introduce the N2O5 into the calorimeter, it is estimated that less than 1% of the N205 decayed. N o attempt was made to correct for this decomposition because it is partially self-compensating since NO2 reacts exothermically with H 2 0in a manner similar to N,05. The container was then placed in the calorimeter and allowed to reach thermal equilibrium. After 2-3 min, the vessel was smashed thereby introducing N2O5 into the system. Temperature changes were calculated by using a linear regression routine to fit the background temperature data, as a function of time, collected before and after the reaction took place. Then the resultant lines were subtracted point by point with their differences averaged to give the final change in temperature (see Figure 1). Depending on the mass of dinitrogen pentoxide reacted, AT ranged from 0.35 to 0.53 O C . As an added check, aliquots of the resulting HNO,(aq) solution were titrated with a 0.0980 M standard of NaOH. In all cases, the mass of dinitrogen pentoxide measured with the analytical balance agreed, within experimental uncertainty, with the mass calculated from the titration. The determinations by weight had a smaller experimental uncertainty and therefore only these were used in the calculation of enthalpy change. N2O5 Vapor Pressure. The vapor pressure of N2Os was measured directly above the crystalline compound over the temperature range of 21 1-273 K. Experimental apparatus consisted of a 10-L vacuum-jacketed Dewar flask equipped with a stirring paddle, quartz oscillator thermometer, 2.0-L glass container, and a copper coil through which temperature regulating fluid was continuously circulated. Ethanol acted as the heat-transfer fluid (7) Davidson, J. A.; Viggiano, A. A;, Howard, C. J.; Dotan, I.; Fehsenfeld, F. C.; Albritton, D. L.; Ferguson, E. E. J . Chem. Phys. 1978, 68, 2085. (8) Cantrell, C. A.; Davidson, J. A.; Shetter, R. E.; Anderson, B. A,; Calvert, J. G . J . Phys. Chem. 1987, 91, 6017.
O C
0.351 0.355 0.426 0.487 0.530
AH,,” AHdHNO?, . kcalrmol ‘kcal/mol -17.617 -17.704 -17.511 -17.815 -18.070
-49.4879 -49.4817 -49.4827 -49.4800 -49.4784
.
AHdN20J, $1,‘ kcal/moI -13.044 -12.956 -13.139 -12.830 -12.572 -12.91 f 0.54
“The calculation of the AH2involved two corrections, the difference between the heat capacity of deionized water and dilute HNOJ, and the additional contribution of a small amount of N2OS(g)to the exothermicity. The corrections totaled less than 1% in all cases (see text). JANAF thermochemical data on heats of formation for various dilute HNO, solutions were fit to a second-order polynomial, over a small mole ratio range, to yield the desired AHf(HN03, as). The enthalpy of formation (kJ) was fit versus the mole ratio (moles of H20/mole of HN03) to give the following quadratic equation, -206.918 - 1.29 X 10-4x 1.53 X 10-*x2. CTheenthalpy of formation for N20s(s)was calculated by using JANAF thermochemical data for the AHf(H20, I) at 298 K. The uncertainty reported here is only due to precision at the 95% confidence interval.
+
inside the Dewar flask. A rigid foam lid insulated the system from further thermal losses. The 2.0-L glass container was fitted with a stainless steel valve and plumbed into an external vacuum system. The container, valve included, was then immersed in the temperature-regulated ethanol. A feed-through in the lid allowed an MKS Baratron pressure transducer to be positioned a short distance from the valve. For pressures below 10.120 Torr, a 10-Torr capacitance manometer monitored the pressure of the container. Above 10.120 Torr, a 100-Torr capacitance manometer was used. Errors associated with the reported pressure measurements were determined one of two ways. For measured values of pressure at and below 0.454 Torr, the reported error was *0.001 Torr. For values above 0.454 Torr, the reported error was fO.15% of the reading. Two sets of error limits were used because at low pressures the measurement was limited by the five figure digital display while at higher pressures the accuracy of the Baratron determined the reported uncertainty. In either case, the largest value of uncertainty was carried through all subsequent calculations. The total volume between the valve and pressure transducer was estimated at 20 mL. Using a small volume between the pressure transducer and the 2.0-L glass container served two purposes. First, and most importantly, it minimized the perturbation to the internal flask temperature due to expanding the N 2 0 5 vapor into the pressure measuring device. Since a relatively small increase in volume (1%) was associated with the expansion, the change in flask temperature due to the AHsubbecame insignificant. Second, the small volume minimized thermal decomposition associated with exposing the N2O5 vapor to room temperature during measurements. As before, N,05 was initially transferred into the container under vacuum. A Neslab ULT-80 low-temperature bath circulator regulated the temperature in increments of 5 deg. After the temperature was increased, a strip chart recorder was used to determine that thermal equilibrium had been established before taking any data. Any NO2 formed through thermal decomposition was removed by vacuum prior to taking data. All temperature and pressure data were displayed on digital read-out and recorded. The data were reduced with a weighted least-squares routine described elsewhere by CvetanoviE et The routine included a weighting parameter which accounted for the effect on the uncertainty in the measured pressure due to a linear to natural log transformation of the dependent variable. The final result of the fit was a logarithmic form of the Arrhenius equation giving the vapor pressures as a function of temperature. The enthalpy of sublimation for the process (9) CvetanoviE, R. J.; Singleton, D. L.; Paraskevopoulos, G. J . Phys. Chem. 1979, 83, 52-54
4174
The Journal of Physical Chemistry, Vol. 92, No. 14, 1988 N205(s)
N2°5(g)
(3)
was then obtained directly from the slope of the fit line.
Results Heat of Formation of N 2 0 , . The results of this study on the AHf(N205,s)are summarized in Table I. Two corrections were necessary to arrive at the reported AH,,, for reaction 2. First of all, the heat capacity of the entire system was corrected to take into account that the heat-transfer fluid was no longer deionized water but a dilute HNO, solution. Using an equation for the apparent molal heat capacity of various dilute nitric acid solutions provided by the C R C Handbook,Io the C,(H,O) at 5 OC was subtracted from the initial C,(SYS). Then the apparent molal heat capacity for nitric acid was added to the difference thus giving the heat capacity for a system with a dilute nitric acid medium. The value reported earlier in the Experimental Section is the corrected heat capacity of the system. The other correction to the AH,,, deals with the presence of vapor phase N2O5 in the glass container used to introduce dinitrogen pentoxide into the system. Because of the vapor, the measured AT and hence the measured AH,,, for the hydrolysis reaction would be increased by an amount proportional to the AHs,,,(N20,). To correct for this increase, 4.347 X lo-, kcal was subtracted from the AHrxn(in kcal). This number was obtained by using the ideal gas law for a 0.070-L volume, AHsubdetermined in this experiment, and vapor pressure data also determined in this experiment. Both corrections were minor at the temperatures employed in this study; combined they amounted to less than a 1% change in the measured enthalpy of reaction. Using Hess' law and basic mathematical manipulations of the following equations yields the reported enthalpies l,f2
CJSYS) = Rl(
$ 2) -
AHf(N205,s)= 2AHf(HN03,aq) - AHf(H20,1) - AH2 AHf(NzO,,g) = AH3 + AHf(N,O,,s) AHr(NO,,g) = AH,
+ AHf(N@,,g)
(E) (F)
- AHf(N02,g) (GI
where V is the applied voltage, R, is the resistance, Tb is the background temperature, Cvapis the correction for the presence of vapor phase N 2 0 , equal to 4.347 X kcal, ml is the mass of N 2 0 5 , M I is the molecular weight of N,05, and R is the universal gas constant. From eq E, AH2 from Table I, and the current thermodynamic properties for H 2 0and HNO,, the AHf(N20S,s)is endothermic by 12.91 f 0.61 kcal/mol. The AHf(N,O,,g) is calculated from eq F using 14.104 0.075 kcal/mol for AH,, determined in the vapor pressure portion of this experiment, and found to be 1.19 0.61 kcal/mol. Finally, AHf(N03,g) is calculated from eq G by using 22.1 1 0.43 kcal/mol for A H l , determined by Cantrell et al. through the temperature dependence of the equilibrium constant for reaction 1. The enthalpy of formation for the nitrate free radical at 298 K is 15.39 f 0.75 kcal/mol. Both the magnitude and error, respectively, associated with this determination are significantly less than those of the currently accepted value. All reported uncertainties are combined random and estimated
*
*
(10) Parker, V. B. CRC Handbook of Chemistry and Physics; CRC Press: Boca Raton, FL, 1984; p D-123 (NSRDS-NBS 2).
McDaniel et al. systematic errors at the 95% confidence interval. In this study, although special care was taken to reduce random and systematic errors, some possible sources remain. The presence of nitric acid in the N 2 0 5would effectively decrease the calculated enthalpy of formation, as illustrated by eq C and F. A correction due to vapor-phase N 2 0 , has already been applied; however, inaccuracies may exist in this estimate. Certainly inaccuracies exist among the current estimates of the thermodynamic properties of HNO, and NO,. The enthalpy of formation of N 2 0 5relies heavily upon the accuracy with which these values are determined. All of the terms in eq E and F used to calculate the AHH,(N20,,g) are between 1 and 2 orders of magnitude larger than the enthalpy of formation for gas-phase N 2 0 S . As a result, seemingly insignificant fractional errors even as small as a few percent propagating through eq E and F may have significant effects on the calculated AHf(N,O,,g) and its associated uncertainty. The enthalpy of hydrolysis obtained in this experiment differs from those of previous efforts by approximately 2.5 kcal/mol. Ogg measured the enthalpy of hydrolysis and found it to be -20.2 kcal/mol which yielded a AHf(N,O,,s) of -9.6 kcal/mol. It appears that this value may be in error. A likely source of error is Ogg's failure to correct for vapor-phase N 2 0 , present in his system prior to the reaction with water. A 2.5 kcal/mol increase in the enthalpy of hydrolysis may be accounted for if 1 g of solid dinitrogen pentoxide were contained in a relatively small volume (40 mL) at room temperature. Unfortunately the experimental procedures are described in too little detail to properly assess sources of error. Ray and Ogg measured the enthalpy for N20, reacting with NO. From their measurements they determined the enthalpy of formation of gas-phase N 2 0 5to be 3.06 0.20 kcal/mol. This value differs by 1.87 kcal/mol from that obtained in this study. The value determined in this study is clearly outside the precision reported by Ray and Ogg. The reason for this discrepancy is unknown. Because of the importance of this value in determining the thermodynamic properties of the nitrate free radical, additional experiments should be performed. The most significant result of this study is the determination of the enthalpy of formation for the NO3 radical. Previous determinations for the enthalpy of formation range from 17.0 to 17.6 k c a l / m ~ l . ~ ~ IThe ' - ~se ~ values rely upon the older estimates of the thermochemical properties for N205.The value reported here is between 1.6 and 2.2 kcal/mol less than previous determinations. This discrepancy, in large part, is due to the AHf(N205,g)determined in this experiment. The previous estimates for the temperature dependence of the equilibrium constant for reaction 1 are essentially in agreement and would yield values for the enthalpy of formation of the nitrate free radical similar to that reported in this work. The 2.2 kcal/mol difference between the past and current determinations may have profound effects on the calculated enthalpies of certain reactions. Processes currently considered thermoneutral that involve the nitrate radical as reactants may become thermally unfavorable. For example, calculated equilibrium constants involving NO, in hydrogen abstraction reactions are approximately 1/40 of that obtained by using the currently accepted enthalpy of formation for NO,. N,O, Vapor Pressure. Table I1 summarizes data from this study as well as that obtained by previous experimenters. A comparison is made between actual vapor pressures and pressures calculated from best fits to data over a temperature range of 211-305 K. It appears that in the high-temperature regime (273-305 K), this study is in good agreement with both Daniels and Bright' and Russ and PokornyL4(see Figure 2). However, in the low-temperature regime (21 1-273 K) it is apparent that the previous experimenters' data and fit lines deviate from this study (see also Table 11). The reason for discrepancy stems from extrapolating beyond the valid range of the previous experiments' (11) Schott, G.; Davidson, N. J . Am. Chem. SOC.1958, 80,1841. (12) Burrows, J. P.;Tyndall, G. S.; Moortgat, G. K. Chem. Phys. Lett. 1985, 119, 196. (1 3) Graham, R. A,; Johnston, H. S . J . Phys. Chem. 1978, 82, 265. (14) Russ, F.; Pokorny, E. Monatsh. Chem. 1913, 34, 1027.
Enthalpies of Formation of N2O5 and NO3
The Journal of Physical Chemistry, Vol. 92, No. 14, 1988 4175 Temperature (K)
TABLE 11: Summary of N206Vapor Pressure Data (in Torr) and Corresponding Calculations and Extraoolations from Fit Functions
temp, K 21 1.66 213.21 216.61 216.78 221.60 221.66 221.70 226.23 226.55 230.67 231.55 236.33 237.53 242.79 242.95 243.00 246.50 247.41 247.71 252.00 252.59 252.60 252.67 257.61 257.62 258.00 262.30 262.57 263.00 267.54 268.00 272.67 273.00 278.00 28 1.70 283.00 283.50 288.00 293.00 298.00 303.00 305.50
this work
Daniels Russ and and Bright Pokorny
0.025 0.034 0.053 0.059 0.113 0.114 0.120 0.223 0.229 0.405 0.454 0.849 0.989 1.894 1.799 2.3 3.3 3.294 3.391 6.3 5.908 5.870 5.985 10.240 10.120 13 18.6 16.89 21 25.40 32 44.90 51 79
51.5 111.2
118 132 183 279 420 620 760
-calculated Daniels Russ this and and work" Brightb Pokorny' 0.026 0.033 0.055 0.056 0.1 15 0.116 0.117 0.221 0.231 0.405 0.455 0.846 0.985 1.882 1.918 1.930 2.922 3.248 3.363 5.478 5.850 5.857 5.902 10.115 10.126 10.546 16.556 17.023 17.792 28.128 29.438 46.336 47.817 76.327 106.74 119.84 125.26 185.23 282.08 423.54 627.48 760.08
0.179 0.208 0.289 0.294 0.466 0.469 0.471 0.724 0.747 1.101 1.196 1.870 2.090 3.397 3.448 3.463 4.771 5.183 5.321 7.857 8.286 8.294 8.346 12.997 13.008 13.457 19.710 20.186 20.967 31.241 32.523 48.812 50.225 77.223 105.87 118.22 123.33 180.19 273.48 413.31 622.02 761.90
0.034 0.043 0.072 0.074 0.150 0.151 0.152 0.285 0.298 0.515 0.578 1.057 1.226 2.299 2.342 2.356 3.520 3.899 4.033 6.455 6.877 6.884 6.936 11.636 1 1.648 12.1 10 18.643 19.145 19.970 30.891 32.258 49.600 51.100 79.465 108.93 121.43 126.57 182.50 270.00 393.50 565.38 674.24
"These data were calculated from a weighted linear fit of In P (atm) versus 1000/T (K) with slope of -7.0982 and a y intercept of 23.2348. *Thesedata were calculated from the equation log P = 1244/T + 34.1 log T - 85.929 given by Daniels and Bright. Since no data were taken below 258 K, the values reported between 211 and 258 K were extrapolated. CThesedata were calculated from the equation log P = 3161.2/T + 1.75 log T - 0.00606/T + 10.679 given by Russ and Pokorny. Values above 283 K and below 243 K are extrapolations beyond the region where measurements were actually taken. fit functions and does not indicate that the results obtained by the previous authors should be discarded. A logarithmic form of the Clausius-Clapeyron equation giving the natural log of the vapor pressure as a function of 1000/T (K) is obtained from the fit function described earlier. The value of the slope is -7.0982 with the intercept equal to 23.2348. The d(ln Pa,,)/d(lOOO/T (K)) of this fit function yields the AH,,, for dinitrogen pentoxide. As expected from the Clausius-Clapeyron equation, a temperature-independent value of 14.104 f 0.075 kcal/mol is obtained. The enthalpies of sublimation reported by JANAF for Russ and Pokorny, and Daniels and Bright are 13.37 f 0.06 and 13.25 f 0.12, kcal/mol, respectively. These earlier determinations differ by approximately 0.85 kcal/mol from the enthalpy of sublimation determined in this experiment. However, there seems to be a discrepancy between the values originally reported by the cited authors and the JANAF Thermochemical Tables. In one case, good agreement between this study and an
300
275
250
' % -1 h
VI
Y
-3
.K
n
-5 c
0
.
-
-
1 ;;-
d y - -7 -
- 11 3.20
..*.
,
This Work Daniels &
Bright
Puss & Pokorn;
3.60
,
,
,
\ ,
4.00 4.40 1000 / T (K)
,
~
4.80
Figure 2. Plot of logarithm of vapor pressure in atmospheres versus 1000/T (K) over a temperature range of 211-310 K. TABLE 111: Thermodynamic Parameters Recommended from This Studv" AH,, kcal/mol -17.74 f 0.61 AH3, kcal/mol 14.10 f 0.07 AHf(N,OS,s), kcal/mol -12.91 i 0.61 W ( N 2 0 5 ,g), kcal/mol 1.19 f 0.61 15.39 f 0.75 AHf(N03,g), kcal/mol All reported uncertainties are combined random and estimated systematic errors at the 95% confidence interval.
earlier determination is obtained for the value for the enthalpy of sublimation at 298 K. For instance, Daniels and Bright indicate a temperature dependence associated with their experimentally determined AHsub. At 298 K, however, they report an enthalpy of sublimation equal to 14.04 kcal/mol which is in good agreement with this study. A value of 14.5 kcal/mol may be obtained from the d(ln Patm)/d(1/ T ) of Russ and Pokorny's fit function, which closely parallels this study (see Table 11). It is not clear how JANAF arrived at their reported values but any inaccuracies may have a significant effect on subsequent calculations especially on the enthalpy of formation of gas-phase dinitrogen pentoxide.
Conclusion This study evaluates the thermodynamic properties of N 2 0 5 using modern analytical techniques and experimental equipment. Special attention is given to reducing the experimental uncertainties associated with laboratory measurements of pressure and temperature. The values obtained for the enthalpy of formation of N205,in both the solid and gas phases, are significantly different from the current determinations reported by JANAF. From these new thermochemical data, the enthalpy of formation of the nitrate free radical in the gas phase at 298 K is evaluated and found to be 15.39 f 0.75 kcal/mol. This determination is significantly less endothermic than previous determinations and as a result may have profound effects on thermochemical calculations involving the nitrate radical. The thermodynamic parameters recommended from this study are summarized in Table 111. Acknowledgment. This work was supported in large part through project W16042 of the Upper Atmosphere Research Program, Earth Science and Applications Division, The National Aeronautics and Space Administration. We are grateful to Dr. Pat Zimmerman and John Lind for many useful discussions related to this work. The National Center for Atmospheric Research is funded by the National Science Foundation. Registry No. N,O,, 10102-03-1; NO3, 12033-49-7
