The reaction and rate law for the gas-phase decomposition of dinitrogen pentaoxide are 2N2O5(g) 4NO2(g) + O2(g) rate= = k[N2O5] Which of the following can be considered valid mechanisms for the reaction? I One-step collision -> II 2 N2O5(g) → 2 NO 3(g) + 2 NO2(g) 2 NO3(g) →2 NO2(g) +2 O(g) 2 O(g) → O2(g) [slow] [fast] [fast] = III N2O5(g) NO3(g) + NO2(g) [fast] NO2(g) + N2O5(g) → 3 NO2(g) + O(g) [slow] NO3(g) + O(g) → NO2(g) O2(g) [fast] IV 2 N2O5(g)=2 NO2(g) + N2O3(g) +3 O(g) [fast] N2O3(g) + O(g) 2 NO2(g) [slow] 2 O(g) → O2(g) [fast] V -> 2 N2O5(g) N4010(g) →> N4O10(g) 4 NO2(g) + O2(g) [slow] [fast] III IV none of the above

For the decomposition of nitrous oxide at 565 °C 2 N20(g) → 2 N2(g) + 02(g) the average rate of disappearance of N,0 over the time period from t = 0 s to t = 2.02×103 s is found to be 2.82×10 4 M/s. What is the average rate of appearance of 0, over the same time period? M/s.

For the reaction 3O2(g) + 4NO(g) → 2N2O5(g) the rate of change in concentration of O2, Δ[O2]/Δ t = 0.48 mol/L*sec. What is the rate of change of NO and what is the rate of change of N2O5? Show your calculations.

Rate constants for the decomposition of gaseous dinitrogen pentoxide are 3.7×10−53.7×10−5 s−1s−1 at 25 ∘C∘C and 1.7×10−31.7×10−3 s−1s−1 at 55 ∘C∘C.2N2O5(g)→4NO2(g)+O2(g What is the activation energy for this reaction in kJ/molkJ/mol? What is the rate constant at 35 ∘C∘C?

A proposed mechanism for the decomposition of hydrogen peroxide consists of three elementary steps: (1) H₂O2(g) →>> 2OH(g) (2) H₂O2(g) + OH(g) → H₂O(g) + HO₂(g) (3) HO₂(g) + OH(g) → H₂O(g) + O₂(g) If the rate law for the reaction is first order in H₂O2, which step in the mechanism is the rate-determining step? O Step 1 O None of the steps O Step 3 O Impossible to determine Steps 1 and 2 O All of the steps O Step 2 Hint: The slowest step in a mechanism determines the rate law. From each elementary reaction we can write a rate-law expression. Submit Question

Consider the following mechanism for the formation of nitrpgen dioxide: 2 NO(g) → N,O2(9) (1) N2O2(9) + O2(g) –→ 2 NO2(g) (2) Write the chemical equation of the overall reaction: O yes Are there any intermediates in this mechanism? O no If there are intermediates, write down their chemical formulas. Put a comma between each chemical formula, if there's more than one.

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Consider the hypothetical reaction A(g)  +  B(g)                  C(g) The four containers below represent this reaction being run with different initial amounts of A and B. Assume that the volume of each container is 1.0L. The reaction is second order with respect to A and first order in respect to B. Which of the containers, W,X,Y or Z, would have the greatest reaction rate? Justify your answer.

Suppose the formation of nitrogen dioxide proceeds by the following mechanism: elementary reaction 2NO (g) → N₂O₂ (g) step 1 2 N₂O₂(g) + O₂ (g) Write the balanced chemical equation for the overall chemical reaction. Write the experimentally- observable rate law for the overall chemical reaction. 2NO₂ (g) Suppose also k₁ « k₂. That is, the first step is much slower than the second. << Note: your answer should not contain the concentrations of any intermediates. rate constant 0 k₁ k₂ rate = k 1-0 010 X S