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Chapter 14: Problem 43

Explain why termolecular reactions are rare.

Short Answer

Expert verified
Termolecular reactions are rare due to the low probability of three molecules colliding simultaneously with the correct energy and orientation.

Step by step solution

01

Understanding Termolecular Reactions

Termolecular reactions involve three reactant molecules colliding simultaneously to form products. Unlike unimolecular and bimolecular reactions, these require three molecules to meet at the same place and time.
02

Analyzing Collision Probability

The probability of three molecules colliding simultaneously is very low compared to one or two molecules. The chance decreases exponentially as the number of molecules increases.
03

Effect of Concentration and Energy

For a successful termolecular reaction, the reactants need to have enough energy and be oriented correctly during the collision. This adds another layer of complexity and reduces the likelihood of such reactions occurring.
04

Conclusion

Due to the low probability of simultaneous collisions and the specific energy requirements, termolecular reactions are rare compared to unimolecular and bimolecular reactions.

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Key Concepts

These are the key concepts you need to understand to accurately answer the question.

Collision Theory
Collision Theory describes how chemical reactions occur between molecules. For a reaction to happen, reactant molecules must collide with each other.
Not every collision leads to a reaction though; only those with enough energy and correct orientation do.
This necessary energy is known as the activation energy.

In the context of termolecular reactions, three molecules need to collide at the same time for the reaction to proceed.
This is extremely rare.
  • Every collision must result in the correct orientation and energy required for the reaction.

The probability of all these factors aligning decreases significantly as more molecules are involved.
That's why termolecular reactions are scarce compared to unimolecular or bimolecular reactions.
Reaction Mechanisms
Reaction Mechanisms refer to the step-by-step sequence of elementary reactions by which a chemical change occurs.
Nature of a reaction mechanism can be complex in the case of termolecular reactions.

For such a reaction, there must be an event where three molecules meet simultaneously.
  • This is not just due to random chance; it's dictated by specific paths and intermediate states.
  • These mechanisms help chemists understand how certain reactions proceed and predict the behavior of reactions under different conditions.

While many reactions happen in steps, individual steps in a mechanism might include termolecular processes, but overall, these are rare.
This is due to the aforementioned low probability of three molecules colliding in the precise configuration needed.
Chemical Kinetics
Chemical Kinetics is the study of reaction rates and how different factors affect these rates.
In chemical kinetics, the concentration and energy of the reactants are essential for determining the speed of a reaction.

In a termolecular reaction, a high concentration of reactants is necessary to increase the likelihood of simultaneous collisions.
  • However, increasing concentration doesn't guarantee a reaction since the energy threshold and orientation conditions must also be met.
    This makes the reaction pathway even more sensitive to changes in conditions.
  • Even under ideal conditions, the reaction rate of termolecular reactions tends to be slower compared to simpler reactions.
This is because fewer effective collisions occur per unit time, making these reactions not only rare but also kinetically slow.
Understanding these dynamics helps chemists manipulate conditions to either favor or discourage certain reaction pathways.

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Most popular questions from this chapter

Reactions can be classified as unimolecular, bimolecular, and so on. Why are there no zero-molecular reactions?

As we know, methane burns readily in oxygen in a highly exothermic reaction. Yet a mixture of methane and oxygen gas can be kept indefinitely without any apparent change. Explain.

Define activation energy. What role does activation energy play in chemical kinetics?

For gas-phase reactions, we can replace the concentration terms in Equation (14.3) with the pressures of the gaseous reactant. (a) Derive the equation $$ \ln \frac{P_{t}}{P_{0}}=-k t $$ where \(P_{t}\) and \(P_{0}\) are the pressures at \(t=t\) and \(t=0\) respectively. (b) Consider the decomposition of azomethane $$ \mathrm{CH}_{3}-\mathrm{N}=\mathrm{N}-\mathrm{CH}_{3}(g) \longrightarrow \mathrm{N}_{2}(g)+\mathrm{C}_{2} \mathrm{H}_{6}(g) $$ The data obtained at \(300^{\circ} \mathrm{C}\) are shown in the following table: $$ \begin{array}{cc} \text { Time (s) } & \begin{array}{c} \text { Partial Pressure of } \\ \text { Azomethane (mmHg) } \end{array} \\ \hline 0 & 284 \\ 100 & 220 \\ 150 & 193 \\ 200 & 170 \\ 250 & 150 \\ 300 & 132 \end{array} $$ Are these values consistent with first-order kinetics? If so, determine the rate constant by plotting the data as shown in Figure \(14.7(\mathrm{~b})\). (c) Determine the rate constant by the half-life method.

How does a catalyst increase the rate of a reaction?
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