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Mass spectrometric determination of the heats of formation of the silane fluorides
CHEMICAL PHYSICS LETTERS
VOlUmc 51, number 2
MASS SPECTROMETRIC DETERMINATION OF THE SILANE
15 October 1977
OF THE HEATS OF FORMATION
FLUORIDES
Milton FARBER and R.D. SRIVASTAVA Spruce Sciences. Inc.. Monrou&za.C&fo&a QI016. USA Received
8 July 1977
The reaction of SiFa(g) with Ha(g) in the temperature range 1273-1693 K has been studied by means of the mass spectromctric method. Second and third law reaction enthaipies were obtained for SiFa(g) f Hz(g) = SiFsH(g) + HF(g), SiF$i(g) i- Hz(g) = SiF,H,O f HFW, and SiFzH2Eg) + Hz(g) = SiFHs(g) f HF
1. Introduction The high temperature reactions of SiF4(g) and Hz(g) were investigated mass spectrometrically to determine the heats of formation of the silane fluorides. The results obtained appear to be the first experimentally determined formation enthalpies for SiF3H(g), SiF2H2(g) and SiFH,(g). The investigation included simultaneous reaction studies of the various equiiibria involving the silane fluorides by means of effusionmass spectrometry in the temperature range 12731603 K, Second and third law bond energies and heats of formation were derived from the reaction enthalpies.
The dual vacuum chamber-quadrupole mass spectrometer system employed in these experiments has been described previously [ 11. The SiF,(g) and Ha(g) were contained in stainless steel bulbs and metered into the cell through variable leaks connected to 1 mm diameter alumina tubing. The alumina cell was 25 mm long, with an inside diameter of 6.8 mm, and employed an elongated orifice 0.75 mm diameter by 5-S mm long for beam collimation. Operating in the molecular flow regime of 0.1 mm Hg with these surface and orifice area dimensions of the cell ensured equilibrium
within the cell [2,33 _ The alumina cell was heated by the main carbon resistance furnace. Temperature measurements to i- 5% were made by optical pyromet~_ Details of the temperature measurements have also been discussed previously [ 1] _The reaction enthalpies for the silane fluorides were calculated for isomolecular gas phase reactions. The ion intensities were identified by their masses, isotopic distribution and appearance potentials. Details of the mass spectrometer resolution, as well as isotopic abundance ratios, have been presented in an earlier publication [4]_ Only the chopped, or shutterable, portion of the intensities was recorded, since the mass spectrometer was equipped with a beam modulator and a phase sensitive amplifier. In order to obtain the bond energies and the heats of formation for the sub~uo~des it was necessary to ascertain with a high degree of confidence that the measured ion intensities were those from the parent species and not from the fragments of the larger molecules. Therefore, operation of the mass spectrometer at an ionizing voltage 1 to 2 eV above the appearance potential will in nearly all cases allow only the formation of the ion from the parent species since a fragmentation process usually occurs at higher ionization voltages [2,5,6] _ A number of checks were made to ensure that SiF3H was produced from the reaction of SiF4 with H, and that the measured ion intensities were not the result of formation of the higher isotopic compounds 307
Volume 5 1, number 2
29SiF3 and 30SiF3, or the fragmentation of SiF4 to SiFl_ Since silicon has three naturally occurring isotopes (28Si = 92.2% , 2gSi = 4.7% and 3oSi = 3.1%), it would be possible that the 2gSiF3 peak at amu 86 would contribute to the 28SiF3H peak which a?so appears at amu 86. Employing published data [7 ] a calculation was made at a temperature of 1500 K to determine the contribution of SiF3 produced from the competing reaction of SiF4 and H2. The results showed that the ratio of SiF3/SiF3H was 0.62. However, at atomic mass number 86 the maximum contribution of the species 2gSiF3 would be 3% the intensity of_‘8SiF3H, producing an error of only 0.1 kcal in the reaction heat. If 2gSiF; were due to fragmentation of SiF4, its pressure would be independent of the reaction with H2. However, when the H2 flow was discontinued the amu 86 peak disappeared, indicating that no contribution was present because of the fragment 2ySiF;. The appearance potentials obtained at 1500 K were HF = 16 eV, H2 = 15.5 eV, SiF4 = 15.4 eV, SiF3H = 11.0 eV, SiF2H2 = 11.0 eV and SiFH3 = 13.0 eV (all + 1 ev). Values reported in the literature are SiF4 = 15.4 -C0.4 eV, H, = 15.4 f 0.02 eV and HF = 15.8 + 0.02 eV [8] _ Corrections were made to the ion intensities for their differences in relative cross sections and their molecular weight. The relative maximum ionization cross sections for single ionization were taken from Mann [9] _Those for the molecular species were calculated by multiplying the sum of atomic cross sections by an empirical factor of 0.7 [lo--121 _The relative cross-section values employed were SiF4 = 2.99, SiF3H ~2.29, SiF2H2 = 2.25, SiFH3 = 1.79, HF = 0.878 and H2 = 0.327. These values are prqbably accurate to * 30% and for isomolecular reactions would possibly yield errors of -+ 1 kcal/mole in the third law values. Due to the uncertainties in these cross-section estimates it is advisable to have second law corroboration for definitive thermodynamic values. Corrections were made for the electron multiplier gain, which is based on the square root of the molecular weight_ Cross-section and multiplier corrections were applied to the individual species intensities (example: pi = [~i((Or)si/lsi(or)i]Tsi)Details for calculations of second and third law thermodynamic data have been presented previously [2,10] .
308
15 October 1977
CHEMICAL PHYSICS LETTERS 3. Results and discussion 3. I. Heat of formation of SiF3H(g)
Equilibrium constants molecular reaction SiF,(g)
were calculated
for the iso-
+ H2(g) = SiF3H(g) + HF(g)
(1)
from the corrected relative ion intensities. The thermodynamic data as a function of temperature are listed in table l_ Thermal functions for the species were taken from the JANAF Tables [7]. An average third law AH298 value of 27.8 + 2 kcal/mole was obtained for reaction (1). Employing AH,,, values of -65.14 f 0.2 k&/mole for HF(g) and -385.98 f 0.2 kcal/ mole for SiF4(g) [7], a third ktW LUtEg8 vdue of -293.0 t- 2 kcal/moIe was calculated for SiF,H(g). The ion intensities for reaction (1) were employed for a least squares calculation of log& versus l/T (fig. 1), resulting in a second law AH value of 25.4 +3 kcal/mole at an average temperature of 1536 K. This value, when reduced to standard conditions and employing the thermal functions given above, yielded -293.6 f 3 kcaI/mole for the ANf29g of SiF3H(g). There are no other experimental data reported for the AH, of SiF,H(g); however, the JANAF Tables [7] contain a calculated value of -287 kcal/mole obtained from a linear interpolation of the bond strengths in SiF,(g) and SiH4(g). It has been found in several Table 1
Thermodynamic
data for the
reaction SiF4(g) + Ha(g) =
SiFaH(g) + HF(g) T WI
10sKP
--aW++&,a)~~1 (cal mo!-r K-r)
A&as (kcal mol-r 1
1273 1383
-3.07 -2.77 -2.53
7.62 7.61
7.46
27.6 28.0
-2.45 -2.28 -2.26 -2.26 -2.20 -2.12
7.42 7.38 7.37 7.37 7.35 7.33
27.9 27.8 27.8 27.8 27.8 27.8
1453 1500
1563 1570 1573 1600 1635 1653 1693
-2.10 -2.00
27.7
7.31
27.9
7.28
27.8 av.
27.8
CHEMICAL
Volume 5 1, number 2
PHYSICS LETTERS Table 2 Thermodynamic
-l.i
15 October 1977
data for the reaction SiFsH(g)
+ Hz(g) =
SiFt Hz (g) f HFW T(K)
-2.0 GB
'oMp
-~[lGO,-f&,W1 (cat. mol” K-l)
(kcal mol-’ 35.0 35.5 35.5 35.6 35.7 35.7 36.3 35.8 35.9 35.9 av. 35.7
1383 1453
-3.90 -3.72
1500 1563 1570 1573 1600 163.5 1653
-3.56 -3.37 -3.36 -3.36 -3.35 -3.19 -3.15
7.52 7.46 7.42 7.38 7.37 7.37 7.35 7.33 7.31
1693
-3.05
7.29
Aff& )
-25
-3.0
6.0
7.0
6.0
8.0
104 it/l Fig. 1. Plots of the logKI versus l/T for the gaseous isomolecalar equilibria resulting from the reaction of SiFe(g) with
The ion intensities for reaction (2) were used for a least squares calculation of log Q versus l/r (fig. I), resulting in a second law AH value of 29.7 + 5 kcalf mole at an average temperature of 1562 K. This vaIue, when reduced to standard conditions, yielded a AHfzg8 value of -195.6
4 5 kcal/mole
for SiFxH2(g).
A linear interpolation of bond strengths of SiFd(g) and SiH4(g) yielded an estimated value of -189 + 5 kcal/mole for the AHf2g8 of SiF2H2(g) [7] _
H2@.
3.3. Heat of formation of SiFH3(g) studies involving high temperature
stable species that when fluorine is present the molecule is slightly more stable than would be expected from linear bond calculations [13--161. Theoretical calculations may clarify this type of discrepancy. other
3.2. Heat of formotion of SiF2H2(g) Equilibrium constants were calcuIated from the ion intensities for the isomolecular reaction SiF, H(g) + Hz(g) = SiFZ Hz(g) + HF(g).
(2)
The thermodynamic data as a function of temperature are listed in table 2. An average third law AH298 value of 35.7 4 2 kc&l/mole was obtained for reaction (2). Employing -293.0 k&/mole for the A&z98 obtained in this work for SiF3H(g) and -65.14 f 0.2 kcal/ m&e for the Ai&98 of HF(g) [7], a third law A&29, v&e of -1922 -C2 kcaljmole was calculated for SiF2H2Cg).
Ion intensities were employed to calculate equilibrium constants for the isomolecular reaction SiF2H2{g) + Hi(g) = SiFH,(g)
+ HF(g).
(3)
The thermodyn~ic data for this reaction are presented in table 3. The thermal functions for the species were taken from the JANAF Tables [7] _ The average third Iaw value for reaction (3) was 27.6 t- 2 kcaljmole. This yielded a third law AHf2g, value of -99.4 + 2 k&/mole for SiFH,(g) employing -192.2 for the A%98 of SiF,H,(g). The ion intensities for reaction (3) were employed in a least squares calculation of logKI versus l/T (fig_ 11, resulting in a second law AH value of 25.1 & 4 k&/mole at an average temperature of 1562 K. This value, when reduced to standard conditions, together with thermal functions [7], yielded -100.4 4 4 kcaljmcle for the AH,,98 of SiFH,(g)_ A linear inte~olation of bond strengths of SiFd(g) and SiHd(g) 309
CHEMICAL
Volume 51, number 2
Table 3 Thermodynamic data for the reaction SiFHs(g) + HF(g)
T(K)
1383 1453 1500 1563
-3.54 -2.49 -3.31 -3.16
3.43 3.38 3.36 3.32
1570 1573
-3.15
3.31
-3.12 -3.07 -2.99 -2.95 -2.87
3.31 3.29 3.27 3.26 3.24
1600 163.5 1653 1693
[ 1 ] M. Farber,
A&,, (kcal mol-r )
K-i)
27.1 28.1 27.7 27.7 27.8 27.2 27.7 27.7 27.7 27.7
+ 5 kcal/moIe
for
Acknowledgement This research was sponsored by the Air Force Of-
Research (AFSC), United States Air
Force, under Contract
310
1977
M-A. Frisch and H-C. Ko, Trans. faraday Sot. 3202. (21 M. Farber, R-D. Srivastava and O.hi. Uy, J. Chem. Sot. Faraday 168 (1972) 249. [31 M. Farber and A.J. Darnell, J. Chem. Phys. 25 (1956) 526. r41 M. Farber and R.D. Srivastava, Combust. Flame 20 (1973) 33. ISI M. Farber and R-D. Srivastava, J. Chem. Sot. Faraday I 70 (1974) 1581. 161 M. Farber and R.D. Srivastava, J. Chem. Sot. Faraday I, to be published. Ta171 D.R. Stull and H. Prophet, JANAF Thermochemical bles, NSRDS-NBS37 (US Govt. Printing Office, Washington, 1971), revised 1974 and 1975. Potentials, Appearance Potentials, and Heats ]81 Ionization of Formation of Gaseous Positive Ions, US Department of Commerce, National Bureau of Standards Publication NSRDSNBS26, June 1969. PI J.B. Mann, J. Chem. Phys. 46 (1967) 1646. PO1 M. Farber and R.D. Srivastava, J. Chem. Sot. Faraday I 69 (1973) 390. 1111 R.F. Pottie, J. Chem. Phys. 44 (1966) 1646. (121 J.W. Otvos and D.P. Stevenson, J. Am. Chcm. Sot. 78 (1956) 546. 1131 O.M. Uy, R.D. Srivastava and M. Farber, High Temp. Sci. 4 (1972) 227. [I41 M. Farber and S.P. Harris, High Temp. Sci. 3 (1971) 231. [I51 R.D. Srivastava and M. Farber, J. Phys. Chem. 75 (1971) 1760. [I61 M. Farber and H.L. Petersen, Trans. Faraday Sot. 59 (1963) 836. 65 (1969)
yielded an estimated value of -90 the A&,,, of SiFH3(g) [7] _
15 October
f Hz(g) =
av_ 27.6
fice of Scientific
LETTERS
References SiF2H2(g)
-A[(+-&,)/TI (cal mol-r
PHYSICS
F44620-74-C-0075.
VOlUmc 51, number 2
MASS SPECTROMETRIC DETERMINATION OF THE SILANE
15 October 1977
OF THE HEATS OF FORMATION
FLUORIDES
Milton FARBER and R.D. SRIVASTAVA Spruce Sciences. Inc.. Monrou&za.C&fo&a QI016. USA Received
8 July 1977
The reaction of SiFa(g) with Ha(g) in the temperature range 1273-1693 K has been studied by means of the mass spectromctric method. Second and third law reaction enthaipies were obtained for SiFa(g) f Hz(g) = SiFsH(g) + HF(g), SiF$i(g) i- Hz(g) = SiF,H,O f HFW, and SiFzH2Eg) + Hz(g) = SiFHs(g) f HF
1. Introduction The high temperature reactions of SiF4(g) and Hz(g) were investigated mass spectrometrically to determine the heats of formation of the silane fluorides. The results obtained appear to be the first experimentally determined formation enthalpies for SiF3H(g), SiF2H2(g) and SiFH,(g). The investigation included simultaneous reaction studies of the various equiiibria involving the silane fluorides by means of effusionmass spectrometry in the temperature range 12731603 K, Second and third law bond energies and heats of formation were derived from the reaction enthalpies.
The dual vacuum chamber-quadrupole mass spectrometer system employed in these experiments has been described previously [ 11. The SiF,(g) and Ha(g) were contained in stainless steel bulbs and metered into the cell through variable leaks connected to 1 mm diameter alumina tubing. The alumina cell was 25 mm long, with an inside diameter of 6.8 mm, and employed an elongated orifice 0.75 mm diameter by 5-S mm long for beam collimation. Operating in the molecular flow regime of 0.1 mm Hg with these surface and orifice area dimensions of the cell ensured equilibrium
within the cell [2,33 _ The alumina cell was heated by the main carbon resistance furnace. Temperature measurements to i- 5% were made by optical pyromet~_ Details of the temperature measurements have also been discussed previously [ 1] _The reaction enthalpies for the silane fluorides were calculated for isomolecular gas phase reactions. The ion intensities were identified by their masses, isotopic distribution and appearance potentials. Details of the mass spectrometer resolution, as well as isotopic abundance ratios, have been presented in an earlier publication [4]_ Only the chopped, or shutterable, portion of the intensities was recorded, since the mass spectrometer was equipped with a beam modulator and a phase sensitive amplifier. In order to obtain the bond energies and the heats of formation for the sub~uo~des it was necessary to ascertain with a high degree of confidence that the measured ion intensities were those from the parent species and not from the fragments of the larger molecules. Therefore, operation of the mass spectrometer at an ionizing voltage 1 to 2 eV above the appearance potential will in nearly all cases allow only the formation of the ion from the parent species since a fragmentation process usually occurs at higher ionization voltages [2,5,6] _ A number of checks were made to ensure that SiF3H was produced from the reaction of SiF4 with H, and that the measured ion intensities were not the result of formation of the higher isotopic compounds 307
Volume 5 1, number 2
29SiF3 and 30SiF3, or the fragmentation of SiF4 to SiFl_ Since silicon has three naturally occurring isotopes (28Si = 92.2% , 2gSi = 4.7% and 3oSi = 3.1%), it would be possible that the 2gSiF3 peak at amu 86 would contribute to the 28SiF3H peak which a?so appears at amu 86. Employing published data [7 ] a calculation was made at a temperature of 1500 K to determine the contribution of SiF3 produced from the competing reaction of SiF4 and H2. The results showed that the ratio of SiF3/SiF3H was 0.62. However, at atomic mass number 86 the maximum contribution of the species 2gSiF3 would be 3% the intensity of_‘8SiF3H, producing an error of only 0.1 kcal in the reaction heat. If 2gSiF; were due to fragmentation of SiF4, its pressure would be independent of the reaction with H2. However, when the H2 flow was discontinued the amu 86 peak disappeared, indicating that no contribution was present because of the fragment 2ySiF;. The appearance potentials obtained at 1500 K were HF = 16 eV, H2 = 15.5 eV, SiF4 = 15.4 eV, SiF3H = 11.0 eV, SiF2H2 = 11.0 eV and SiFH3 = 13.0 eV (all + 1 ev). Values reported in the literature are SiF4 = 15.4 -C0.4 eV, H, = 15.4 f 0.02 eV and HF = 15.8 + 0.02 eV [8] _ Corrections were made to the ion intensities for their differences in relative cross sections and their molecular weight. The relative maximum ionization cross sections for single ionization were taken from Mann [9] _Those for the molecular species were calculated by multiplying the sum of atomic cross sections by an empirical factor of 0.7 [lo--121 _The relative cross-section values employed were SiF4 = 2.99, SiF3H ~2.29, SiF2H2 = 2.25, SiFH3 = 1.79, HF = 0.878 and H2 = 0.327. These values are prqbably accurate to * 30% and for isomolecular reactions would possibly yield errors of -+ 1 kcal/mole in the third law values. Due to the uncertainties in these cross-section estimates it is advisable to have second law corroboration for definitive thermodynamic values. Corrections were made for the electron multiplier gain, which is based on the square root of the molecular weight_ Cross-section and multiplier corrections were applied to the individual species intensities (example: pi = [~i((Or)si/lsi(or)i]Tsi)Details for calculations of second and third law thermodynamic data have been presented previously [2,10] .
308
15 October 1977
CHEMICAL PHYSICS LETTERS 3. Results and discussion 3. I. Heat of formation of SiF3H(g)
Equilibrium constants molecular reaction SiF,(g)
were calculated
for the iso-
+ H2(g) = SiF3H(g) + HF(g)
(1)
from the corrected relative ion intensities. The thermodynamic data as a function of temperature are listed in table l_ Thermal functions for the species were taken from the JANAF Tables [7]. An average third law AH298 value of 27.8 + 2 kcal/mole was obtained for reaction (1). Employing AH,,, values of -65.14 f 0.2 k&/mole for HF(g) and -385.98 f 0.2 kcal/ mole for SiF4(g) [7], a third ktW LUtEg8 vdue of -293.0 t- 2 kcal/moIe was calculated for SiF,H(g). The ion intensities for reaction (1) were employed for a least squares calculation of log& versus l/T (fig. 1), resulting in a second law AH value of 25.4 +3 kcal/mole at an average temperature of 1536 K. This value, when reduced to standard conditions and employing the thermal functions given above, yielded -293.6 f 3 kcaI/mole for the ANf29g of SiF3H(g). There are no other experimental data reported for the AH, of SiF,H(g); however, the JANAF Tables [7] contain a calculated value of -287 kcal/mole obtained from a linear interpolation of the bond strengths in SiF,(g) and SiH4(g). It has been found in several Table 1
Thermodynamic
data for the
reaction SiF4(g) + Ha(g) =
SiFaH(g) + HF(g) T WI
10sKP
--aW++&,a)~~1 (cal mo!-r K-r)
A&as (kcal mol-r 1
1273 1383
-3.07 -2.77 -2.53
7.62 7.61
7.46
27.6 28.0
-2.45 -2.28 -2.26 -2.26 -2.20 -2.12
7.42 7.38 7.37 7.37 7.35 7.33
27.9 27.8 27.8 27.8 27.8 27.8
1453 1500
1563 1570 1573 1600 1635 1653 1693
-2.10 -2.00
27.7
7.31
27.9
7.28
27.8 av.
27.8
CHEMICAL
Volume 5 1, number 2
PHYSICS LETTERS Table 2 Thermodynamic
-l.i
15 October 1977
data for the reaction SiFsH(g)
+ Hz(g) =
SiFt Hz (g) f HFW T(K)
-2.0 GB
'oMp
-~[lGO,-f&,W1 (cat. mol” K-l)
(kcal mol-’ 35.0 35.5 35.5 35.6 35.7 35.7 36.3 35.8 35.9 35.9 av. 35.7
1383 1453
-3.90 -3.72
1500 1563 1570 1573 1600 163.5 1653
-3.56 -3.37 -3.36 -3.36 -3.35 -3.19 -3.15
7.52 7.46 7.42 7.38 7.37 7.37 7.35 7.33 7.31
1693
-3.05
7.29
Aff& )
-25
-3.0
6.0
7.0
6.0
8.0
104 it/l Fig. 1. Plots of the logKI versus l/T for the gaseous isomolecalar equilibria resulting from the reaction of SiFe(g) with
The ion intensities for reaction (2) were used for a least squares calculation of log Q versus l/r (fig. I), resulting in a second law AH value of 29.7 + 5 kcalf mole at an average temperature of 1562 K. This vaIue, when reduced to standard conditions, yielded a AHfzg8 value of -195.6
4 5 kcal/mole
for SiFxH2(g).
A linear interpolation of bond strengths of SiFd(g) and SiH4(g) yielded an estimated value of -189 + 5 kcal/mole for the AHf2g8 of SiF2H2(g) [7] _
H2@.
3.3. Heat of formation of SiFH3(g) studies involving high temperature
stable species that when fluorine is present the molecule is slightly more stable than would be expected from linear bond calculations [13--161. Theoretical calculations may clarify this type of discrepancy. other
3.2. Heat of formotion of SiF2H2(g) Equilibrium constants were calcuIated from the ion intensities for the isomolecular reaction SiF, H(g) + Hz(g) = SiFZ Hz(g) + HF(g).
(2)
The thermodynamic data as a function of temperature are listed in table 2. An average third law AH298 value of 35.7 4 2 kc&l/mole was obtained for reaction (2). Employing -293.0 k&/mole for the A&z98 obtained in this work for SiF3H(g) and -65.14 f 0.2 kcal/ m&e for the Ai&98 of HF(g) [7], a third law A&29, v&e of -1922 -C2 kcaljmole was calculated for SiF2H2Cg).
Ion intensities were employed to calculate equilibrium constants for the isomolecular reaction SiF2H2{g) + Hi(g) = SiFH,(g)
+ HF(g).
(3)
The thermodyn~ic data for this reaction are presented in table 3. The thermal functions for the species were taken from the JANAF Tables [7] _ The average third Iaw value for reaction (3) was 27.6 t- 2 kcaljmole. This yielded a third law AHf2g, value of -99.4 + 2 k&/mole for SiFH,(g) employing -192.2 for the A%98 of SiF,H,(g). The ion intensities for reaction (3) were employed in a least squares calculation of logKI versus l/T (fig_ 11, resulting in a second law AH value of 25.1 & 4 k&/mole at an average temperature of 1562 K. This value, when reduced to standard conditions, together with thermal functions [7], yielded -100.4 4 4 kcaljmcle for the AH,,98 of SiFH,(g)_ A linear inte~olation of bond strengths of SiFd(g) and SiHd(g) 309
CHEMICAL
Volume 51, number 2
Table 3 Thermodynamic data for the reaction SiFHs(g) + HF(g)
T(K)
1383 1453 1500 1563
-3.54 -2.49 -3.31 -3.16
3.43 3.38 3.36 3.32
1570 1573
-3.15
3.31
-3.12 -3.07 -2.99 -2.95 -2.87
3.31 3.29 3.27 3.26 3.24
1600 163.5 1653 1693
[ 1 ] M. Farber,
A&,, (kcal mol-r )
K-i)
27.1 28.1 27.7 27.7 27.8 27.2 27.7 27.7 27.7 27.7
+ 5 kcal/moIe
for
Acknowledgement This research was sponsored by the Air Force Of-
Research (AFSC), United States Air
Force, under Contract
310
1977
M-A. Frisch and H-C. Ko, Trans. faraday Sot. 3202. (21 M. Farber, R-D. Srivastava and O.hi. Uy, J. Chem. Sot. Faraday 168 (1972) 249. [31 M. Farber and A.J. Darnell, J. Chem. Phys. 25 (1956) 526. r41 M. Farber and R.D. Srivastava, Combust. Flame 20 (1973) 33. ISI M. Farber and R-D. Srivastava, J. Chem. Sot. Faraday I 70 (1974) 1581. 161 M. Farber and R.D. Srivastava, J. Chem. Sot. Faraday I, to be published. Ta171 D.R. Stull and H. Prophet, JANAF Thermochemical bles, NSRDS-NBS37 (US Govt. Printing Office, Washington, 1971), revised 1974 and 1975. Potentials, Appearance Potentials, and Heats ]81 Ionization of Formation of Gaseous Positive Ions, US Department of Commerce, National Bureau of Standards Publication NSRDSNBS26, June 1969. PI J.B. Mann, J. Chem. Phys. 46 (1967) 1646. PO1 M. Farber and R.D. Srivastava, J. Chem. Sot. Faraday I 69 (1973) 390. 1111 R.F. Pottie, J. Chem. Phys. 44 (1966) 1646. (121 J.W. Otvos and D.P. Stevenson, J. Am. Chcm. Sot. 78 (1956) 546. 1131 O.M. Uy, R.D. Srivastava and M. Farber, High Temp. Sci. 4 (1972) 227. [I41 M. Farber and S.P. Harris, High Temp. Sci. 3 (1971) 231. [I51 R.D. Srivastava and M. Farber, J. Phys. Chem. 75 (1971) 1760. [I61 M. Farber and H.L. Petersen, Trans. Faraday Sot. 59 (1963) 836. 65 (1969)
yielded an estimated value of -90 the A&,,, of SiFH3(g) [7] _
15 October
f Hz(g) =
av_ 27.6
fice of Scientific
LETTERS
References SiF2H2(g)
-A[(+-&,)/TI (cal mol-r
PHYSICS
F44620-74-C-0075.