T = 298 K
ΔGo = -16.668 KJ/mol
ΔGo = -16668 J/mol
use:
ΔGo = -R*T*ln Kc
-16668 = - 8.314*298.0* ln(Kc)
ln Kc = 6.7276
Kc = 8.351*10^2
Answer: b
QUESTION 12 Given that AG for NH=-16.667999999999999 kJ/mol, calculate the equilibrium constant for the following reaction...
Calculate the equilibrium constant for the following reaction at 25 °C and 150°C. 11. (2) NH CI() NH()HC1(g) NH,C1 (s) -202.87 NH( -16.45 HC1(8) -95.30 AG (kJ-mo1-1) AG [-16.45(-95.0 -202.87] 91.12 kJ At 298 K: K#e*(- Gran/RT) e*(91120/(8.314 x 298) 1.1 At 423 K: K= e"(-aGrin/RT)-e*(91 120/(8.314 x 423) = 5.6 × 10-16 10-12 Calculate the equilibrium constant for the following reaction at 25 °C and 150°C. 11. (2) NH CI() NH()HC1(g) NH,C1 (s) -202.87 NH( -16.45 HC1(8) -95.30 AG...
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The equilibrium constant for the equation Ag+ (aq) + 2NH, (aq) =[Ag(NH),]+(aq) is Kf = 2.5 x 103 M-2 at 25.0 °C. Calculate the value of AGixn at 25.0 °C. C AG = k Is the reaction spontaneous under standard conditions? kJ/mol J/mol no O yes Calculate the value of AGrxn at 25.0 °C when [Ag+] = 0.00138 M, (NH2) = 0.438 M, and [[Ag(NH3)2]+] = 0.00739 M. Is the reaction spontaneous under these conditions? AGrxn = kJ/mol O yes...
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Given the values of AGfº given below in kJ/mol, calculate the value of AG in kJ for the reaction at 298 K.: 3 NO2(g) + H20(1) > 2 HNO3(0) + NO(g) AGº (NO2) = 49. AG (H2000) = -237. AGRO (HNO3) --77. AGP (NO) = 85.
Calculate the equilibrium constant of the reaction CO(g) + H2(g) ⇌ H2CO(g) given that for the production of liquid formaldehyde ΔrG< = +28.95 kJ mol−1 at 298 K and that the vapour pressure of formaldehyde is 1500 Torr at that temperature.
1. a) Show that: , equals: AGpuC")_ACC"= Hem , assuming AHxn is constant over these ranges. b) Calculate AGxn of the reaction at 500K given AHPx = -92.22 kJ/mol and AG n=-32.9 kJ/mol at 298 K. N2 (g) + 3 H2 (9) → 2 NH3(g)
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