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Chapter 13: Problem 128

Write the general rate law for each reaction, using \(x\) and \(y\) exponents as orders: (a) \(2 \mathrm{NO}+\mathrm{O}_{2} \rightarrow 2 \mathrm{NO}_{2}\) (b) \(2 \mathrm{H}_{2} \mathrm{O}_{2} \rightarrow 2 \mathrm{H}_{2} \mathrm{O}+\mathrm{O}_{2}\)

Short Answer

Expert verified
(a) \(Rate = k [NO]^x [O_2]^y\) (b) \(Rate = k [H_2O_2]^x\)

Step by step solution

01

a) Writing the rate law for 2 NO + O2 → 2 NO2

For the given reaction, we will be using the concentration of reactants: NO and O2. The rate law for a reaction involves the rate constant k multiplied by the concentrations of reactants, raised to the power of their respective orders in the reaction. Using x and y as the exponents, we have: \(Rate = k [NO]^x [O_2]^y\)
02

b) Writing the rate law for 2 H2O2 → 2 H2O + O2

For this reaction, we will be using the concentration of the reactant H2O2. Following the same process as before and using x as the exponent for the concentration of H2O2, we have the rate law as: \(Rate = k [H_2O_2]^x\)

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Most popular questions from this chapter

True or false? An energy-downhill reaction can always be expected to be faster than an energyuphill reaction. Explain your answer.

The value of \(\Delta E_{\mathrm{ryn}}\) for an exothermic reaction is always negative. (a) Why is this so in terms of \(E_{\text {reactants }}\) versus \(E_{\text {products }} ?\) (b) Why is this so in terms of bonds broken in the reactants versus bonds formed in the products?

A reaction \(\mathrm{A}+\mathrm{B} \rightarrow\) Product is run in a balloon. (Both A and B are gases.) The balloon has a volume of 1 L and is initially loaded with 1 mole of \(\mathrm{A}\) and \(1 \mathrm{~mole}\) of \(\mathrm{B}\). The reaction has the rate law Rate \(=k[\mathrm{~A}]\) The reaction is run again using the same amount of reactants, but this time in a balloon that has a volume of \(0.5 \mathrm{~L}\). How much faster will the reaction proceed in the smaller balloon? Explain your answer.

The rate law for a reaction involving \(\mathrm{A}(g)\) as the only reactant is: Rate \(=k[\mathrm{~A}]^{2}\) What happens to the rate when: (a) The volume of the reaction container is halved? (b) The concentration of \(\mathrm{A}\) is tripled?

Consider the transition state for a chemical reaction. (a) What is it (define it). (b) Can there be only imminent bond breaking in a transition state? Explain.
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