The following initial rate data are for the reaction of hypochlorite ion with iodide ion in 1 M aqueous hydroxide solution: _IO_IO_I+_100 Experiment [OCI], M Initial Rate, [I], M M. 8-1 1 0.00240 0.00313 5.36 × 10-4 2 0.00240 0.00626 1.07 x 10 -3 3 0.00479 0.00313 1.07 x 10-3 4 0.00479 0.00626 2.14 x 10-3 Complete the rate law for this reaction in the box below. Use the form k[A] [B]", where '1' is understood for m or n and concentrations taken to the zero power do not appear. Don't enter 1 for mor n. Rate= k = M-1 8 -1

The following initial rate data are for the reaction of hypochlorite ion with iodide ion in 1 M aqueous hydroxide solution: _IO_IO_I+_100 Experiment [OCI], M Initial Rate, [I], M M. 8-1 1 0.00240 0.00313 5.36 × 10-4 2 0.00240 0.00626 1.07 x 10 -3 3 0.00479 0.00313 1.07 x 10-3 4 0.00479 0.00626 2.14 x 10-3 Complete the rate law for this reaction in the box below. Use the form k[A] [B]", where '1' is understood for m or n and concentrations taken to the zero power do not appear. Don't enter 1 for mor n. Rate= k = M-1 8 -1

Chemistry for Engineering Students

4th Edition

ISBN:9781337398909

Author:Lawrence S. Brown, Tom Holme

Publisher:Lawrence S. Brown, Tom Holme

Chapter11: Chemical Kinetics

Section: Chapter Questions

Problem 11.33PAE: The following experimental data were obtained for the reaction of \'I14* and NOf in acidic solution....

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The acid-catalyzed iodination of acetone CH3COCH3(aq) + I2(aq) CH3COCH2I(aq) + HI(aq) is a common laboratory experiment used in general chemistry courses to teach the method of initial rates. The reaction is followed spectrophotometrically by the disappearance of the color of iodine in the solution. The following data (J. P. Birk and D. L Walters, Journal of Chemical Education, Vol. 69, p. 585, 1992) were collected at 23 C for this reaction. Determine the rate law for this reaction.
The oxidation of iodide ion by the hypochlorite ion in the presence of hydroxide ions I(aq) + ClO(aq) IO(aq) + Cl(aq) was studied at 25 C, and the following initial rates data (Y. Chia and R. E. Connick, Journal of Physical Chemistry, Vol. 63, p. 1518, 1959) were collected: (a) Determine the rate law for this reaction. (b) One mechanism that has been proposed for this reaction is the following: Show that the rate law predicted by this mechanism matches the experimentally determined rate law in part a. (Note that when writing the expression for K the equilibrium constant, [H2O] is not involved. See Chapter 15.)
For the reaction 2N2O5(g)4NO2(g)+O2(g) the following data were collected, where Rate=[N2O5]t Time (s) T = 338 K [N2O5] T=318K [N2O5] 0 1.00 101 M 1.00 101 M 100. 6.14 102 M 9.54 102 M 300 2.33 102 M 8.63 102 M 600. 5.41 1O3 M 7.43 102 M 900. 1.26 1O3 M 6.39 102 M Calculate Ea for tills reaction.
The following experimental data were obtained for the reaction of \'I14* and NOf in acidic solution. NH/(aq) + NO2-(aq) — N;(g) + 2 H,O(f) INH/I (mol L1) [NO21 (mol L-1, Rate = A[NJ/At (mol L-1 s’) 0.0092 0.098 3.33 X IO"7 0.0092 0.049 1.66 X 10‘7 0.0488 0.196 3.51 X 10"6 0.0249 0.196 1.80 X 10-6 Determine the rate law for this reaction and calculate the rate constant.
Instantaneous rates for the reaction of hydroxide ion with Cv+ can be determined from the slope of the curve in Figure 11.3 at various concentrations. They are (1) At 4.0 105 mol/L, rate = 12.3 107 mol L1 s1 (2) At 3.0 105 mol/L, rate = 9.25 107 mol L1 s1 (3) At 2.0 105 mol/L, rate = 6.16 107 mol L1 s1 (4) At 1.5 105 mol/L, rate = 4.60 107 mol L1 s1 (5) At 1.0 105 mol/L, rate = 3.09 107 mol L1 s1 (a) What is the relationship between the rates in (1) and (3)? Between (2) and (4)? Between (3) and (5)? (b) What is the relationship between the concentrations in each of these cases? (c) Is the rate of the reaction proportional to the concentration of Cv+? Explain your answer.
The label on a bottle of 3% (by volume) hydrogen peroxide, H2O2, purchased at a grocery store, states that the solution should be stored in a cool, dark place. H2O2decomposes slowly over time, and the rate of decomposition increases with an increase in temperature and in the presence of light. However, the rate of decomposition increases dramatically if a small amount of powdered MnO- is added to the solution. The decomposition products are H2O and O2. MnO2 is not consumed in the reaction. Write the equation for the decomposition of H2O2. What role does MnO2 play? In the chemistry lab, a student substituted a chunk of MnO2 for the powdered compound. The reaction rate was not appreciably increased. WTiat is one possible explanation for this observation? Is MnO2 part of the stoichiometry of the decomposition of H2O2?
The rate of photodecomposition of the herbicide piclo- ram in aqueous systems was determined by exposure to sunlight for a number of days. One such experiment produced the following results. (Data from R.T. Hedlun and C.R. Youngson, “The Rates of Photodecomposition of Picloram in Aqueous Systems," Fate of Organic Pesticides in tbe Aquatic Environment, Advances in Chemistry Series, #111, American Chemical Society (1972), 159—172.) Exposure Time, t (days) [Pidoram] (mol L_1) 0 4.14 X 10-6 7 3.70 X 10-6 14 3.31 X 10-6 21 2.94 X 10~6 28 2.61 X 10~6 35 2.30 X 10-6 42 2.05 X 10-6 49 1.82 X 10"6 56 1.65 X 10-6 Determine the order of reaction, the rate constant, and the half-life for the photodecomposition of picloram.
Label the elementary processes for the reaction between H2 and O2 see section 20.7 as initiation, propagation, branching, or termination reactions.
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Use the References to access important values if needed for this question. The following initial rate data are for the reaction of hypochlorite ion with iodide ion in 1 M aqueous hydroxide solution: OCT+1 → or + Cr Experiment [OCI , M Io. M Initial Rate, M 8-¹ 0.00243 0.00184 3.30 × 10-4 0.00243 0.00486 0.00486 1 2 3 4 Complete the rate law for this reaction in the box below. Use the form k[A] [B]", where '1' is understood for 0.00367 6.58 × 10-4 0.00184 6.60 × 10-4 0.00367 1.32 × 10-³ m or n and concentrations taken to the zero power do not appear. Don't enter 1 for m or n. Rate = k = M
Cc.43.
The Arrhenius equation shows the relationship between the rate constant k and the temperature T in kelvins and is typically written as k = Ae-E/RT where R is the gas constant (8.314 J/mol K), A is a constant called the frequency factor, and Da is the activation energy for the reaction. However, a more practical form of this equation IS In = (1-12) Ea R which is mathmatically equivalent to In Ea F (1/2 R T₂ k2 T1 where ki and k2 are the rate constants for a single reaction at two different absolute temperatures (T₁ and T₂). ▼ Part A The activation energy of a certain reaction is 42.1 kJ/mol. At 20 °C, the rate constant is 0.0150s¹. At what temperature in degrees Celsius would this reaction go twice as fast? Express your answer with the appropriate units. T₂ = Submit Part B Value k₂ = Request Answer Submit Given that the initial rate constant is 0.0150s¹ at an initial temperature of 20 °C, what would the rate constant be at a temperature of 200. °C for the same reaction described in Part…