The rate of the reaction: CO (g) + NO2 (g) à CO2 (g) + NO (g) was measured at several temperatures, and the following data were collected:
|
Temp (oC) |
K (M-1s-1) |
|
35 |
0.184 |
|
45 |
0.322 |
Using this data determine the value of Ea (energy of activation)
You have the following reaction: NO2 (g) + CO (g) ⟶ NO (g) + CO2 (g) The rate constant (k) at 701 K is 2.57 M-1s-1. If the activation energy is 150 kJ/mol, what is k at 895 K? R = 8.314 J/(mol*K) A) 680 M-1s-1 B) 443 M-1s-1 C) 2.58 M-1s-1 D) 0.950 M-1s-1 E) 6.52 M-1s-1
(a) 2 NO2(g) à 2 NO(g) + O2(g) The rate law for this reaction is 2nd order. The rate constant is k = 0.775 M¯¹s¯¹. How much time would it require for [NO2] to go from 0.06M to 0.05M? (b) A 1st order reaction has a rate constant of k = 1.0 X 10¯³ s¯¹ at 25ᴼC. If the reaction rate doubles (2X faster) at 35ᴼC, what is the activation energy (Ea) for this reaction?
The activation energy for the reaction NO2(g)+CO(g)⟶NO(g)+CO2(g) is Ea = 175 kJ/mol and the change in enthalpy for the reaction is ΔH = -375 kJ/mol . What is the activation energy for the reverse reaction?
The activation energy for the reaction NO2(g)+CO(g)⟶NO(g)+CO2(g) is Ea = 200 kJ/mol and the change in enthalpy for the reaction is ΔH = -200 kJ/mol . What is the activation energy for the reverse reaction?
The activation energy for the reaction NO2(g)+CO(g)⟶NO(g)+CO2(g) is Ea = 150 kJ/mol and the change in enthalpy for the reaction is ΔH = -375 kJ/mol . What is the activation energy for the reverse reaction?
If the mechanism behind the reaction NO2(g) + CO(g) Ó NO(g) + CO2(g) is : 1- 2NO2(g) à 2NO(g) + O2(g) (slow) 2- NO(g) + CO(g) + O2(g) à NO2(g) + CO2(g) (fast) Then its rate law is: A) Rate = k [NO2] . [CO] B) Rate = k [NO2] . [CO2] C) Rate = k [NO 212 D) Rate = k [co]2
The reaction NO2(g) + CO(g) CO2(g) + NO(g) has a rate constant of 2.57 M−1∙s−1 at 701 K and 567 M−1∙s−1 at 895 K. Find the activation energy in kJ/mol
The reaction rate of CO and NO2 in the reaction CO(g) + NO2(g) → CO2(g) + NO(g) is measured using the initial rates method. The results are tabulated below. [CO] (mol/L) NO2 (mol/L) -([CO]/Δt (mol/L·s) 8.00 10-4 5.50 10-4 8.40 10-8 8.00 10-4 7.78 10-4 1.68 10-7 1.60 10-3 5.50 10-4 1.68 10-7 Determine the rate expression and calculate the rate constant for the reaction.
The reaction 2 NO2(g) → 2 NO (g) + O2(g) has rate constants of 2.70 x 10-2 M-1s-1 at 227 oC and 0.240 M-1s-1 at 277oC. What is the activation energy of this reaction? (Given: Arrhenius equation, k = Ae-Ea/RT ) A) 99.6 kJ/mol B) 22.8 kJ/mol C) 49.8 kJ/mol D) -22.8 kJ/mol E) 65.3 kJ/mol I'm unsure on how to do it since you're not given the frequency factor
• For the reaction CO (g) + NO2 (g) → CO2(g) + NO (g) • Reaction rate = k[CO][NO2] • k = 1.9 L/mol•h • Determine the initial rate of reaction when [CO] = 3.8 x 10-4 mol/L and [NO2] = 0.650 x 10-4 mol/L. Answer: 4.7 x 10-8 mol/L•h know it is something simple I am not plugging in correctly, but I can't get this answer.