AP Chemistry
5.2 Introduction to Rate Law
The gas-phase decomposition of nitrogen dioxide at \(500\text{ K}\) is represented by the equation below.
\[ 2\,\text{NO}_2\text{(g)} \rightarrow 2\,\text{NO(g)} + \text{O}_2\text{(g)} \]
The reaction is second order with respect to \(\text{NO}_2\text{(g)}\), with the rate law:
\[ \text{Rate} = k[\text{NO}_2]^2 \]
In an experiment at \(500\text{ K}\), when \([\text{NO}_2] = 0.050\text{ M}\), the initial rate of the reaction is \(2.5 \times 10^{-4}\text{ M}\cdot\text{s}^{-1}\). What is the value of the rate constant, \(k\), at this temperature?
\[ 2\,\text{NO}_2\text{(g)} \rightarrow 2\,\text{NO(g)} + \text{O}_2\text{(g)} \]
The reaction is second order with respect to \(\text{NO}_2\text{(g)}\), with the rate law:
\[ \text{Rate} = k[\text{NO}_2]^2 \]
In an experiment at \(500\text{ K}\), when \([\text{NO}_2] = 0.050\text{ M}\), the initial rate of the reaction is \(2.5 \times 10^{-4}\text{ M}\cdot\text{s}^{-1}\). What is the value of the rate constant, \(k\), at this temperature?
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