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Chapter 12: 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
Not all collisions of reactant molecules result in product formation due to two main reasons: firstly, the collisions might not have sufficient activation energy, which is the minimum energy required for a successful reaction; secondly, the reactant molecules might collide at an incorrect orientation, preventing efficient overlapping of the reacting orbitals and thus, hindering the bonding process.

Step by step solution

01

Understand the Collision Theory

The collision model proposes that for a chemical reaction to occur, the reactant molecules must collide with each other. However, not all collisions will lead to a reaction, and certain conditions must be met for the reaction to proceed.
02

Reason 1: Insufficient Activation Energy

The reactant molecules need to have a minimum amount of energy for a successful reaction, which is called the activation energy. Activation energy is the energy barrier that needs to be overcome for a reaction to occur. When molecules collide with less than the required activation energy, they may simply bounce off each other and not react.
03

Reason 2: Incorrect Orientation

The reactant molecules must also collide at the correct orientation (approach angle) for a successful reaction. The collision orientation determines the efficient overlap between the reacting orbitals of the molecules, allowing the bonding process to take place. If the reactant molecules collide at an incorrect orientation, they may bounce off each other and not form a reaction.

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Most popular questions from this chapter

The reaction $$ 0^{\circ} \mathrm{C}, $$ These relationships hold only if there is a very small amount of \(\mathrm{I}_{3}^{-}\) present. What is the rate law and the value of the rate constant? (Assume that rate $=-\frac{\Delta\left[\mathrm{H}_{2} \mathrm{SeO}_{3}\right]}{\Delta t} )$

The activation energy for some reaction $$ \mathrm{X}_{2}(g)+\mathrm{Y}_{2}(g) \longrightarrow 2 \mathrm{XY}(g) $$ is 167 \(\mathrm{kJ} / \mathrm{mol}\) , and \(\Delta E\) for the reaction is $+28 \mathrm{kJ} / \mathrm{mol}$ . What is the activation energy for the decomposition of XY?

Consider the hypothetical reaction $\mathrm{A}_{2}(g)+\mathrm{B}_{2}(g) \longrightarrow\( \)2 \mathrm{AB}(g),$ where the rate law is: $$ -\frac{\Delta\left[\mathrm{A}_{2}\right]}{\Delta t}=k\left[\mathrm{A}_{2}\right]\left[\mathrm{B}_{2}\right] $$ The value of the rate constant at \(302^{\circ} \mathrm{C}\) is $2.45 \times 10^{-4} \mathrm{L} / \mathrm{mol}\( \)\mathrm{s},\( and at \)508^{\circ} \mathrm{C}\( the rate constant is 0.891 \)\mathrm{L} / \mathrm{mol} \cdot \mathrm{s}$ . What is the activation energy for this reaction? What is the value of the rate constant for this reaction at \(375^{\circ} \mathrm{C} ?\)

Experiments during a recent summer on a number of fireflies (small beetles, Lampyridaes photinus) showed that the average interval between flashes of individual insects was 16.3 \(\mathrm{s}\) at \(21.0^{\circ} \mathrm{C}\) and 13.0 \(\mathrm{s}\) at \(27.8^{\circ} \mathrm{C}\) a. What is the apparent activation energy of the reaction that controls the flashing? b. What would be the average interval between flashes of an individual firefly at \(30.0^{\circ} \mathrm{C} ?\) c. Compare the observed intervals and the one you calculated in part b to the rule of thumb that the Celsius temperature is 54 minus twice the interval between flashes.

What are the units for each of the following if the concentrations are expressed in moles per liter and the time in seconds? a. rate of a chemical reaction b. rate constant for a zero-order rate law c. rate constant for a first-order rate law d. rate constant for a second-order rate law e. rate constant for a third-order rate law
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