A. Consider the reaction mechanism shown belowStep 1.H₂O2 +JH₂O + 10.slowStep 2.IO + H2OzH₂O + O2 + IfastWhat is the overall balanced chemical equation?Group of answer choicesH₂O2 → H₂O + O22 H2O2 → 2 H₂O + O2none of the answers are correct

Answered: Image /qna-images/answer/e20a3fbd-18eb-4b79-aeb6-ef2fec030074.jpg

A. Consider the reaction mechanism shown below Step 1. H₂O₂ +J H₂O + 10- slow Step 2. IO + HzQz H₂O + O₂ + I fast What is the overall balanced chemical equation? Group of answer choices H2O2 → H₂O + O2 2 H₂O2 → 2 H₂O + 02 none of the answers are correct

A. Consider the reaction mechanism shown below Step 1. H₂O₂ +J H₂O + 10- slow Step 2. IO + HzQz H₂O + O₂ + I fast What is the overall balanced chemical equation? Group of answer choices H2O2 → H₂O + O2 2 H₂O2 → 2 H₂O + 02 none of the answers are correct

Chemistry: Matter and Change

1st Edition

ISBN:9780078746376

Author:Dinah Zike, Laurel Dingrando, Nicholas Hainen, Cheryl Wistrom

Publisher:Dinah Zike, Laurel Dingrando, Nicholas Hainen, Cheryl Wistrom

Chapter16: Reaction Rates

Section: Chapter Questions

Problem 6STP

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A reaction is believed to occur by the following mechanism: Stepl: 2AI (Fast equilibrium) Step 2: I + B C (Slow) Overall: 2 A + B C What experimentally determined rate law would lead to this mechanism? (a) Rate = k[A][B] (b) Rate = k[A]2[B] (c) Rate = k[A]2 (d) Rate = k[I][B]
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.
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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)
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Can a reaction mechanism ever be proven correct? Can it be proven incorrect?
The ozone in the Earths ozone layer decomposes according to the equation 2 O3(g) 3 O2(g) The mechanism of the reaction is thought to proceed through an initial fast equilibrium and a slow step: Show that the mechanism agrees with this experimental rate law: Rate = (1/2)[O3]/t = k[O3]2[O2].
Consider the following statements: In general, the rate of a chemical reaction increases a bit at first because it takes a while for the reaction to get warmed up. After that, however, the rate of the reaction decreases because its rate is dependent on the concentrations of the reactants, and these are decreasing. Indicate everything that is correct in these statements, and indicate everything that is incorrect. Correct the incorrect statements and explain.
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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.
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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?