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1. For the equilibrium reaction: 2NO2 - N2O4 a) calculate the literature values of ΔG°, ΔH°...
For reactions carried out under standard-state conditions, the equation ΔG = ΔH − TΔS becomes ΔG° = H° − TΔS°. Assuming ΔH° and ΔS° are independent of temperature, one can derive the equation: ln( K2 K1 ) = ΔH° R ( T2 − T1 T1T2 ) where K1 and K2 are the equilibrium constants at T1 and T2, respectively. Given that at 25.0°C, Kc is 4.63×10−3 for the reaction N2O4(g) longrightleftarrow 2NO2(g) ΔH° = 58.0 kJ/mol calculate the equilibrium constant...
Question text Calculate the equilibrium constants, KpKp and KcKc for the equilibrium reaction N2O4(g)⇄2NO2(g)N2O4(g)⇄2NO2(g) at 298 K. N2O4(g)N2O4(g) NO2(g)NO2(g) S0S0 (J/K/mol) 304.29 240.06 ΔfH0ΔfH0 (kJ/mol) 9.16 33.18 Select one or more: A. Kp=9.23Kp=9.23 , Kc=12.3Kc=12.3 B. Kp=0.563Kp=0.563 , Kc=0.33Kc=0.33 C. Kp=0.144Kp=0.144 , Kc=0.0058Kc=0.0058 D. Kp=0.355Kp=0.355 , Kc=1.23
For the reaction at 298 K, 2NO2 g → N2O4 g the values of ΔHo and ΔSo are -58.0 kJ and -177 J/K, respectively. What is the value of ΔG nonstandard at 298 K when the concentration of N2O4 is 1.69 M and the concentration of NO2 is 2.75 M? Remember to answer in correct units. ΔG=
Decomposition of nitrogen dioxide dimer N2O4 is described by the reaction: N2O4(g) = 2NO2(g) Concentration of N2O4 became 2 times less after 2,5⋅103 s. You have to calculate: a) the value of rate constant k of the reaction; b) the value of equilibrium constant Kp. You are given the value of standard Gibb’s energy of formation Goform: substance Goform, kJ/mol NO2(g) 51.6 N2O4(g) 98.4
Calculate the Entropy of the reaction equilibrium between 2NO2 - N2O4 given the equation: Delta G = Delta H - T Delta S where: Enthalpy = -47.5 Kj/mol Gibbs free energy = -3.59 Kj/mol Temperature = 298K
Consider the Haber synthesis of gaseous NH3 (ΔH∘f = -46.1 kJ/mol; ΔG∘f = -16.5 kJ/mol): N2(g)+3H2(g)→2NH3(g) What are the equilibrium constants Kp and Kc for the reaction at 350 K ? You may assume that ΔH∘ and ΔS∘ are independent of temperature.
The reaction N2O4(g) ⇌ 2NO2(g) has ∆G° = 13.3 kJ/mol at 25 °C. What is the value of ∆G, in kJ/mol, for this reaction at this temperature when [N2O4] = 3.0 M and [NO2] = 0.0010 M? Enter your answer to the tenths place. Do not include units. Include the sign if appropriate! Thank you for your help!
Consider the following reaction: 2NO(g)+O2(g)→2NO2(g) Estimate ΔG∘ for this reaction at each of the following temperatures and predict whether or not the reaction will be spontaneous. (Assume that ΔH∘ and ΔS∘ do not change too much within the give temperature range.) ANSWER MUST BE IN kJ!! a) 722 K b) 860 K
For the reaction N2O4(g)⇌2NO2(g), the value of K at 25∘C is 7.19×10−3. Calculate [NO2] at equilibrium when [N2O4]=6.90×10−2mol/L.
A reaction has an equilibrium constant of Kp=0.061 at 27 ∘C. Part A Find ΔG∘rxn for the reaction at this temperature. Find for the reaction at this temperature. 6.98 kJ 0.839 kJ -6.98 kJ 0.628 kJ Part B Above what temperature does the following reaction become nonspontaneous? 2 H2S(g) + 3 O2(g) → 2 SO2(g) + 2 H2O(g) ΔH = -1036 kJ; ΔS = -153.2 J/K 298 K 158.7 K 6.762 × 103 K This reaction is nonspontaneous at all...