Part A The second-order reaction 2 Mn(CO)5 → Mn₂(CO)10, has a rate constant equal to 3.0 x 109 M1 s¹ at 25°C. If the initial concentration of Mn(CO)s is 1.0 x 105 M, how long will it take for 90.% of the reactant to disappear? O 3.0 x 10-4 s O 3.3 x 10-16 s O 3.7 x 10-15 s 3.0 x 10³ s

For the reaction:A(g) + B(g) →AB(g) the rate is 0.30 mol/L.s, when [A]0 = [B]0 = 1.0 mol/L. If the reaction is first order in B and second order in A, what is the rate when [A]0 = 1.9 mol/L and [B]0 = 4.3 mol/L?

The decomposition of N,O5 to NO, and O, is first order, with a rate constant of 4.80 x 1o4/s at 45°C. If the initial concentration of N,O5 is 4.18×102 mol/L, what is the concentration of N205 after 548 s? How long would it take for the concentration ofN,O5 to decrease to 1.00 x 10-2 mol/L from an initial concentration of 4.18x102 mol/L?

Dinitrogen pentoxide, N205, decomposes by a first order reaction with an activation energy of 110 kJ/mol. At 65 °C, the rate constant is k = 5.2 x 10-3 s-1. What is the initial rate of reaction when 8.01 g of N,O5 is confined in a 0.75 liter container and heated to 65 °C? A. 3.98 x 10-4 M/s B. 5.14 x 10-4 M/s C. 0.56 x 10-5 M/s D. 2.31 x 10-2 M/s

A particular reactant decomposes with a half-life of 103 s when its initial concentration is 0.275 M. The same reactant decomposes with a half-life of 211 s when its initial concentration is 0.134 M. Determine the reaction order. 1 What is the value and units of the rate constant for this reaction? Units k = M-1.s-1 M·s-1 M-2.s-1

It was found that for the reaction 2 A + B → 3 C the rate law is Rate = k[A][B] · %3D At the initial concentrations [A]o = 0.100 M and [B]o = 0.200 M the initial rate is 0.0500 M/s. %3D What would be the initial rate if the initial concentrations were [A]o = 0.500 M and [B]o = 0.0750 M? 0.0352 M/s 0.0937 M/s O 0.0711 M/s 0.00533 M/s O 0.469 M/s

Initial rate: 1.6 •10^-7mol/L•s K=6.4 •10^-9L/mol•s Rate law: k[HI]^2 Can you please show me the steps.

13.102 Chlorine oxide (CIO), which plays an important role in the depletion of ozone (see Problem 13.101), decays rapidly at room temperature according to the equation 2C10(g) Cl₂(g) + O₂(g) From the following data, determine the reaction order and calculate the rate constant of the reaction. Time (s) 0.12 x 10-³ 0.96 x 10-³ 2.24 x 10-³ 3.20 x 10-³ 4.00 x 10-³ [CIO] (M) 8.49 x 10-6 7.10 x 10-6 5.79 x 10-6 5.20 x 10-6 4.77 x 10-6

The rate constant for this zero-order reaction is 0.0170 M ·s at 300 °C. A → products How long (in seconds) would it take for the concentration of A to decrease from 0.860 M to 0.320 M? t = S

In the presence of excess thiocyanate ion, SCN, the following reaction is first order in iron(III) ion, Fe+ The rate constant is 1.27/s. Fe+ (ag) + SCN(ag) → Fe(SCN)2+ (ag) If 35.7% of the reaction is required to obtain a noticeable color from the formation of the Fe(SCN)2+ ion, how many seconds are required? Time =