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Chapter 11: Problem 19

The central idea of the collision model is that molecules must collide in order to react. Give two reasons why not all collisions of reactant molecules result in product formation.

Short Answer

Expert verified
One reason why not all collisions between reactant molecules result in product formation is insufficient energy. The colliding molecules may not have enough energy to overcome the activation energy barrier, which is the minimum energy required for a reaction to occur. Another reason is the incorrect orientation of the colliding molecules. A productive collision requires the reactant molecules to collide in a specific orientation that allows their atoms to rearrange and form new bonds, leading to the formation of products. If the orientation is not appropriate, the reaction will not proceed.

Step by step solution

01

Reason 1: Insufficient energy

One reason why not all collisions between reactant molecules result in the formation of products is that the colliding molecules may not have enough energy to overcome the activation energy barrier. Activation energy is the minimum energy required for a reaction to occur; if the colliding molecules' combined kinetic energy is less than the activation energy, the reaction will not proceed and the molecules will simply bounce off each other.
02

Reason 2: Incorrect orientation

Another reason why not all collisions between reactant molecules result in the formation of products is that the molecules may not collide in the proper orientation. A productive collision requires the reactant molecules to collide in such a way that their atoms can rearrange and form new bonds (leading to the formation of products). If the orientation of the colliding molecules is not appropriate for this rearrangement to occur, the reaction will not proceed and the molecules will separate without reacting.

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Most popular questions from this chapter

The activation energy for a reaction is changed from \(184 \space\mathrm{kJ} /\) mol to \(59.0 \space\mathrm{kJ} / \mathrm{mol}\) at \(600 .\) K by the introduction of a catalyst. If the uncatalyzed reaction takes about 2400 years to occur, about how long will the catalyzed reaction take? Assume the frequency factor \(A\) is constant, and assume the initial concentrations are the same.

The decomposition of iodoethane in the gas phase proceeds according to the following equation: $$\mathrm{C}_{2} \mathrm{H}_{5} \mathrm{I}(g) \longrightarrow \mathrm{C}_{2} \mathrm{H}_{4}(g)+\mathrm{HI}(g)$$ At \(660 . \mathrm{K}, k=7.2 \times 10^{-4} \mathrm{s}^{-1} ;\) at \(720 . \mathrm{K}, k=1.7 \times 10^{-2} \mathrm{s}^{-1}\) What is the value of the rate constant for this first-order decomposition at \(325^{\circ} \mathrm{C} ?\) If the initial pressure of iodoethane is 894 torr at \(245^{\circ} \mathrm{C},\) what is the pressure of iodoethane after three half-lives?

Hydrogen peroxide and the iodide ion react in acidic solution as follows: $$\mathrm{H}_{2} \mathrm{O}_{2}(a q)+3 \mathrm{I}^{-}(a q)+2 \mathrm{H}^{+}(a q) \longrightarrow \mathrm{I}_{3}^{-}(a q)+2 \mathrm{H}_{2} \mathrm{O}(l)$$ The kinetics of this reaction were studied by following the decay of the concentration of \(\mathrm{H}_{2} \mathrm{O}_{2}\) and constructing plots of \(\ln \left[\mathrm{H}_{2} \mathrm{O}_{2}\right]\) versus time. All the plots were linear and all solutions had \(\left[\mathrm{H}_{2} \mathrm{O}_{2}\right]_{0}=8.0 \times 10^{-4} \mathrm{mol} / \mathrm{L} .\) The slopes of these straight lines depended on the initial concentrations of \(\mathrm{I}^{-}\) and \(\mathrm{H}^{+} .\) The results follow: The rate law for this reaction has the form $$\text { Rate }=\frac{-\Delta\left[\mathrm{H}_{2} \mathrm{O}_{2}\right]}{\Delta t}=\left(k_{1}+k_{2}\left[\mathrm{H}^{+}\right]\right)\left[\mathrm{I}^{-}\right]^{m}\left[\mathrm{H}_{2} \mathrm{O}_{2}\right]^{n}$$ a. Specify the order of this reaction with respect to \(\left[\mathrm{H}_{2} \mathrm{O}_{2}\right]\) and \(\left[\mathrm{I}^{-}\right]\) b. Calculate the values of the rate constants, \(k_{1}\) and \(k_{2}\) c. What reason could there be for the two-term dependence of the rate on \(\left[\mathrm{H}^{+}\right] ?\)

Describe at least two experiments you could perform to determine a rate law.

Define what is meant by unimolecular and bimolecular steps. Why are termolecular steps infrequently seen in chemical reactions?
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