Consider the following reaction: 2 NO(g) + 2 H 2 (g) → N 2 (g) + 2 H 2 O(g) (b) If the rate constant for this reaction at 1000 K is 6.0 × 10 4 M -2 s -1 , what is the reaction rate when [NO] = 0.035 M and [H 2 ] = 0.015 M? (c) What is the reaction rate at 1000 K when the concentration of NO is increased to 0.10 M, while the concentration of H 2 is 0.010 M?

Hello. In this problem, we are told the given reaction is second order with respect to the concentration of X. And at two R 98 Kelvin, the rate constant for the reaction is 0.683 per minority per second. Were asked to calculate the initial rate. When given the concentration of X is 1.135 Mueller will begin by writing our reaction rate law. So the rate is equal to our reaction rate constant times the concentration of X squared. Since its second order with regard to concentration of X, We have to find the initial rate. So we have our reaction rate constant which is 0.683 per minority per second times the concentration of X which is 1.135 Mueller. We're gonna square that. This works out to 0.880 hilarity squared. For morality second will simplify our units. One of the clarity units cancels were left within 0.880 similarity per second as our initial rate. Thanks for watching. Hope this help.

Ch.14 - Chemical Kinetics

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Brown 15th Edition

Ch.14 - Chemical Kinetics

Problem 30b,c

Chapter 14, Problem 30b,c

Consider the following reaction:
2 NO(g) + 2 H₂(g) → N₂(g) + 2 H₂O(g)
(b) If the rate constant for this reaction at 1000 K is 6.0 × 10⁴ M^-2 s^-1, what is the reaction rate when [NO] = 0.035 M and [H₂] = 0.015 M?
(c) What is the reaction rate at 1000 K when the concentration of NO is increased to 0.10 M, while the concentration of H₂ is 0.010 M?

Verified step by step guidance

Identify the rate law for the given reaction. Since the reaction is 2 NO(g) + 2 H2(g) → N2(g) + 2 H2O(g), the rate law can be expressed as: rate = k[NO]^2[H2]^2, where k is the rate constant.

For part (b), substitute the given values into the rate law. Use k = 6.0 × 10^4 M^-2 s^-1, [NO] = 0.035 M, and [H2] = 0.015 M. The expression becomes: rate = (6.0 × 10^4 M^-2 s^-1) × (0.035 M)^2 × (0.015 M)^2.

Calculate the rate for part (b) by evaluating the expression from the previous step. This involves squaring the concentrations of NO and H2, multiplying them together, and then multiplying by the rate constant.

For part (c), substitute the new concentration values into the rate law. Use k = 6.0 × 10^4 M^-2 s^-1, [NO] = 0.10 M, and [H2] = 0.010 M. The expression becomes: rate = (6.0 × 10^4 M^-2 s^-1) × (0.10 M)^2 × (0.010 M)^2.

Calculate the rate for part (c) by evaluating the expression from the previous step. This involves squaring the new concentrations of NO and H2, multiplying them together, and then multiplying by the rate constant.

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Key Concepts

Here are the essential concepts you must grasp in order to answer the question correctly.

Rate Law

The rate law expresses the relationship between the rate of a chemical reaction and the concentration of its reactants. For a reaction of the form aA + bB → products, the rate can be described by the equation Rate = k[A]^m[B]^n, where k is the rate constant, and m and n are the orders of the reaction with respect to reactants A and B, respectively. Understanding the rate law is essential for calculating reaction rates based on varying concentrations.

Order of Reaction

The order of a reaction refers to the power to which the concentration of a reactant is raised in the rate law. It indicates how the rate of reaction is affected by the concentration of that reactant. For example, if the reaction is first order with respect to NO, doubling the concentration of NO will double the reaction rate. Identifying the order of each reactant is crucial for accurately determining the overall reaction rate.

Units of Rate Constant

The units of the rate constant (k) depend on the overall order of the reaction. For a second-order reaction, like the one given (2 NO + 2 H2 → N2 + 2 H2O), the units of k are M^-2 s^-1. This means that the rate of reaction is proportional to the product of the concentrations of the reactants raised to their respective orders. Understanding these units is vital for correctly applying the rate law to calculate reaction rates.

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Consider a hypothetical reaction between A, B, and C that is first order in A, zero order in B, and second order in C. (e) By what factor does the rate change when the concentrations of all three reactants are tripled?

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The decomposition reaction of N₂O₅ in carbon tetrachloride is 2 N₂O₅ → 4 NO₂ + O₂. The rate law is first order in N₂O₅. At 64°C the rate constant is 4.82 × 10^-3 s^-1. (a) Write the rate law for the reaction.

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Textbook Question

The decomposition reaction of N₂O₅ in carbon tetrachloride is 2 N₂O₅ → 4 NO₂ + O₂. The rate law is first order in N₂O₅. At 64°C the rate constant is 4.82 × 10^-3 s^-1. (c) What happens to the rate when the concentration of N₂O₅ is doubled to 0.0480 M? (d) What happens to the rate when the concentration of N₂O₅ is halved to 0.0120 M?

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Consider the following reaction: 2 NO1g2 + 2 H21g2¡N21g2 + 2 H2O1g2 (d) What is the reaction rate at 1000 K if [NO] is decreased to 0.010 M and 3H24 is increased to 0.030 M?

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The reaction between ethyl bromide (C₂H₅Br) and hydroxide ion in ethyl alcohol at 330 K, C₂H₅Br(alc) + OH^-(alc) → C₂H₅OH(l) + Br^-(alc), is first order each in ethyl bromide and hydroxide ion. When [C₂H₅Br] is 0.0477 M and [OH^-] is 0.100 M, the rate of disappearance of ethyl bromide is 1.7×10^-7 M/s. (a) What is the value of the rate constant?