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Chapter 13: Problem 65

If there were no orientation requirement for collisions, would reactions be faster or slower than they are? Explain your answer.

Short Answer

Expert verified
Reactions would be faster if there were no orientation requirements for collisions, as it would increase the likelihood of successful collisions and result in a higher reaction rate.

Step by step solution

01

Understand Molecular Collisions

Molecular reactions occur when molecules collide with the right orientation and enough energy to overcome the activation energy barrier. The orientation requirement means that the molecules need to collide in a specific way or angle for the reaction to proceed.
02

No Orientation Requirement

If there were no orientation requirements for collisions, it would mean that molecules can react with each other without any specific alignment or collision geometry. It would increase the chances of successful collisions resulting in a reaction, as any collision with sufficient energy would lead to a reaction.
03

Faster or Slower Reactions

If there were no orientation requirements for collisions, the frequency of successful collisions would increase. Since the reaction rate depends on the number of effective collisions, having more successful collisions would lead to a higher reaction rate. Therefore, reactions would be faster in the absence of orientation requirements.
04

Conclusion

Reactions would be faster if there were no orientation requirements for collisions, as it would increase the likelihood of successful collisions and result in a higher reaction rate.

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Most popular questions from this chapter

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?

What do we mean by activation energy?

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?

A student says to you, "Catalysts are not used up in chemical reactions because they are not involved in the reactions." Is this statement true or false? Why?

In a chemical reaction, compound A is converted to compound \(\mathrm{B}\). In the process, energy is absorbed from the surroundings. Which compound is at a higher energy level? Explain your answer.
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