CHAPTER 13 HONORS HOMEWORK
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The reaction below was studied by performing two experiments. NO(g) + O3(g) → NO2(g) + O2(g) In the first experiment (results shown in left table) the rate of disappearance of NO was followed in a large excess of O3. (The [O3] remains effectively constant at 1.0×1014 molecules/cm3.) In the second experiment [NO] was held constant at 2.0×1014 molecules/cm3. The data for the disappearance of O3 are in the right table. The rate of reaction is given by Rate = k[NO]x[O3]y Determine x and y. (Both need to be correct for credit.)
Alrighty kiddos I’m gonna be real all this stuff is gonna be hard, but its okay because I went through the struggle for ya. So we need to figure out how the molarity of NO and O3 increases throughout the experiment.
What is the value of the rate constant obtained from each set of experiments? For the first, Rate = k'[NO]x. What is k’ (in s-1)?
For the second, Rate = k”[O3]y; what is k” (in s-1)?
What is the value of the rate constant for the overall rate law? Rate = k[NO]x[O3]y. What is k (in cm3 molecules-1 s-1)?
In the gas phase the production of phosgene from chlorine and carbon monoxide proceeds by the mechanism given below. The rate of reaction is Rate = k[CO]x[Cl2]y Determine x and y.
The decomposition of NO2(g) occurs by the following bimolecular elementary reaction. 2NO2(g) → 2NO(g) + O2(g) The rate constant at 273 K is 2.3 x 10-12 L mol-1 s-1, and the activation energy is 111 kJ/mol. How long will it take (in s) for the concentration of NO2(g) to decrease from an initial partial pressure of 2.70 atm to 1.50 atm at 457 K? Assume ideal gas behavior.