Feedback See Periodic Table O See Hint The rate constant for the reaction below was determined to be 3.241x105 s-1 at 800 K. The activation energy of the reaction is 245 kJ/mol. What would be the value of the rate constant at 9.60×10² K? N,0(g) → N,(g) +o(g) x 10 -5 s-1 3.241 1st attempt

Feedback See Periodic Table O See Hint The rate constant for the reaction below was determined to be 3.241x105 s-1 at 800 K. The activation energy of the reaction is 245 kJ/mol. What would be the value of the rate constant at 9.60×10² K? N,0(g) → N,(g) +o(g) x 10 -5 s-1 3.241 1st attempt

Chemistry by OpenStax (2015-05-04)

1st Edition

ISBN:9781938168390

Author:Klaus Theopold, Richard H Langley, Paul Flowers, William R. Robinson, Mark Blaser

Publisher:Klaus Theopold, Richard H Langley, Paul Flowers, William R. Robinson, Mark Blaser

Chapter12: Kinetics

Section: Chapter Questions

Problem 63E: Hydrogen iodide, HI, decomposes in the gas phase to produce hydrogen, H2, and iodine, I2. The value...

See similar textbooks

Similar questions

Hydrogen iodide, HI, decomposes in the gas phase to produce hydrogen, H2, and iodine, I2. The value of the rate constant, k, fur the reaction was measured at several different temperatures and the data are shown here: Temperature (K) k (M -1 5-1) 555 6.23107 575 2.42106 645 1.44104 700 2.01103 What is the value of the activation energy (in kJ/mol) for this reaction?
The following statements relate to the reaction for the formation of HI: H2(g) + I2(g) -* 2 HI(g) Rate = it[HJ [I2J Determine which of the following statements are true. If a statement is false, indicate why it is incorrect. The reaction must occur in a single step. This is a second-order reaction overall. Raising the temperature will cause the value of k to decrease. Raising the temperature lowers the activation energy' for this reaction. If the concentrations of both reactants are doubled, the rate will double. Adding a catalyst in the reaction will cause the initial rate to increase.
The reaction 2 NO(g) + 2 H2(g) N2(g) + 2 H2O(g) was studied at 904 C, and the data in the table were collected. (a) Determine the order of the reaction for each reactant. (b) Write the rate equation for the reaction. (c) Calculate the rate constant for the reaction. (d) Find the rate of appearance of N2 at the instant when [NO] = 0.350 mol/L and [H] = 0.205 mol/L.
The initial concentration of the reactant in a tirst-order reaction A —» products is 0.64 rnol/L and the half-life is 30.0 s. Calculate the concentration of the reactant exactly 60 s after initiation of the reaction. How long would it take for the concentration of the reactant to drop to one-eighth its initial value? How long would it take for the concentration of the reactant to drop to 0.040 mol/L?
The reaction H2SeO3(aq) + 6I-(aq) + 4H+(aq) Se(s) + 2I-3(aq) + 3H2O(l) was studied at 0C, and the following data were obtained: [H2SeO3]0 (mol/L) [H+]0 (mol/L) [I]0(mol/L) Initial Rate (mol/L s) 1.0 104 2.0 102 2.0 102 1.66 107 2.0 104 2.0 102 2.0 10-2 3.33 107 3.0 104 2.0 102 2.0 102 4.99 107 1.0 104 4.0 102 2.0 102 6.66 107 1.0 104 1.0 102 2.0 102 0.42 107 1.0 104 2.0 102 4.0 102 13.2 107 1.0 104 1.0 102 4.0 102 3.36 107 These relationships hold only if there is a very small amount of I3 present. What is the rate law and the value of the rate constant? (Assumethatrate=[H2SeO3]t)
11.32 The following experimental data were obtained for the reaction 2A + 3 B—C + 2D [A](mol L 1) [B](mol L ’) Rate = A(C]/Af (mol L-1 s-1) 0.127 0.15 0.033 0.127 0.30 0.132 0.255 0.15 0.066 Determine the reaction order for each reactant and the value of the rate constant.
A study of the rate of dimerization of C4H6 gave the data shown in the table: 2C4H6C8H12 (a) Determine the average rate of dimerization between 0 s and 1600 s, and between 1600 s and 3200 s. (b) Estimate the instantaneous rate of dimerization at 3200 s from a graph of time versus [C4H6]. What are the units of this rate? (c) Determine the average rate of formation of C8H12 at 1600 s and the instantaneous rate of formation at 3200 s from the rates found in parts (a) and (b).
11.64 HBr is oxidized in the following reaction: 4 HBr(g) + O2(g) —• 2 H2O(g) + 2 Br,(g) A proposed mechanism is HBr + O2 -* HOOBr (slow) HOOBr + HBr — 2 HOBr (fast) HOBr + HBr — H2O + Bn (fast) Show that this mechanism can account for the correct stoichiometry. Identify all intermediates in this mechanism. What is the molecularity of each elementary’ step? Write the rate expression for each elementary' step. Identify the rate-determining step.
(Section 11-5) A rule of thumb is that for a typical reaction, if concentrations are unchanged, a 10-K rise in temperature increases the reaction rate by two to four times. Use an average increase of three times to answer the questions below. (a) What is the approximate activation energy of a typical chemical reaction at 298 K? (b) If a catalyst increases a chemical reactions rate by providing a mechanism that has a lower activation energy, then what change do you expect a 10-K increase in temperature to make in the rate of a reaction whose uncatalyzed activation energy of 75 kJ/mol has been lowered to one half this value (at 298 K) by addition of a catalyst?
Regular ?ights of supersonic aircraft in the stratosphere ale of concern because such aircraft produce nitric oxide, NO, as a byproduct in the exhaust of their engines. Nitric oxide reacts with ozone, and it has been suggested that this could contribute to depletion of the ozone layer. The reaction NO+O3NO2+O2 is first order with respect to both NO and O3 with a rate constant of 2.20107 L/mol/s. What is the instantaneous rate of disappearance of NO when [NO]=3.3106 M and [O3]=5.9107M?
The reaction NO(g) + O,(g) — NO,(g) + 0(g) plays a role in the formation of nitrogen dioxide in automobile engines. Suppose that a series of experiments measured the rate of this reaction at 500 K and produced the following data; [NO] (mol L ’) [OJ (mol L 1) Rate = -A[NO]/Af (mol L_1 s-1) 0.002 0.005 8.0 X 10"'7 0.002 0.010 1.6 X 10-'6 0.006 0.005 2.4 X IO-'6 Derive a rate law for the reaction and determine the value of the rate constant.
The following rate constants were obtained in an experiment in which the decomposition of gaseous N2O; was studied as a function of temperature. The products were NO, and NO,. Temperature (K) 3.5 x 10_i 298 2.2 x 10"4 308 6.8 X IO-4 318 3.1 x 10 1 328 Determine Etfor this reaction in kj/mol.