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Chapter 16: Problem 50

What is the central idea of collision theory? How does this model explain the effect of concentration on reaction rate?

Short Answer

Expert verified
Collision theory states that reactions occur when reactant particles collide with sufficient energy and proper orientation. Increasing reactant concentration increases collision frequency, leading to a higher reaction rate.

Step by step solution

01

Define Collision Theory

Collision theory states that for a reaction to occur, reactant particles must collide with sufficient energy and proper orientation. This energy is known as the activation energy.
02

Understand Activation Energy

Activation energy is the minimum amount of energy needed for a collision to result in a reaction. Only collisions with energy equal to or greater than this threshold lead to successful reactions.
03

Explain Effective Collisions

Not all collisions between reactant particles lead to a reaction. Only those collisions that meet both the energy and orientation criteria are effective and result in the formation of products.
04

Effect of Concentration on Collision Frequency

Increasing the concentration of reactants increases the number of particles in a given volume. This leads to more frequent collisions.
05

Effect of Concentration on Reaction Rate

Since the reaction rate depends on the frequency of effective collisions, a higher concentration of reactants results in more frequent effective collisions, thereby increasing the reaction rate.

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Key Concepts

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

Activation Energy
Activation energy is a critical concept in chemistry. It represents the minimum energy required for a chemical reaction to occur. Imagine activation energy as the initial push needed to start rolling a heavy object up a hill. In a chemical reaction, this energy is necessary to break bonds in the reactants so new bonds can form the products. Only particles with energy equal to or greater than this activation energy can successfully collide and react. Therefore, even if many particles collide, only those with sufficient energy lead to a reaction. This concept helps explain why some reactions are slow or need a catalyst to proceed.
Effective Collisions
In collision theory, not all particle collisions result in a chemical reaction. Effective collisions are those that fulfill two main criteria:
  • Energy: The colliding particles must have enough energy to overcome the activation energy barrier.
  • Orientation: The particles must collide with the correct alignment to form products.
Imagine trying to fit a key into a lock—if the key is not aligned properly, it won't turn, even if you apply force. Similarly, effective collisions require the right orientation along with sufficient energy. This is why even at high concentrations, not all collisions are effective, and thus not all lead to a reaction.
Reaction Rate and Concentration
The reaction rate refers to how fast a reaction proceeds, and it is closely linked to the concentration of reactants.
Let's break this down:
  • Increased Concentration: When you increase the concentration of reactants, there's a higher number of particles in a given volume. This leads to more collisions between particles.
  • Frequency of Collisions: More particles mean more frequent collisions. However, keep in mind that these must still be effective collisions to accelerate the reaction.
  • Higher Reaction Rate: Because effective collisions occur more frequently with higher concentrations, the reaction rate increases. Think of it as increasing the number of people in a room; the more people there are, the higher the chances they'll bump into each other.
Overall, understanding these concepts helps explain how and why reactions happen. The interaction between concentration and reaction rate is a fundamental aspect of collision theory in chemistry.

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Most popular questions from this chapter

16.103 Even when a mechanism is consistent with the rate law, later work may show it to be incorrect. For example, the reaction between hydrogen and iodine has this rate law: rate \(=k\left[\mathrm{H}_{2}\right]\left[\mathrm{I}_{2}\right]\). The long-accepted mechanism had a single bimolecular step; that is, the overall reaction was thought to be elementary: \(\mathrm{H}_{2}(g)+\mathrm{I}_{2}(g) \longrightarrow 2 \mathrm{HI}(g)\) In the 1960 s, however, spectroscopic evidence showed the presence of free I atoms during the reaction. Kineticists have since proposed a three-step mechanism: (1) \(\mathrm{I}_{2}(g) \rightleftharpoons 2 \mathrm{I}(g)\) [fast] (2) \(\mathrm{H}_{2}(g)+\mathrm{I}(g) \rightleftharpoons \mathrm{H}_{2} \mathrm{I}(g)\) [fast] (3) \(\mathrm{H}_{2} \mathrm{I}(g)+\mathrm{I}(g) \longrightarrow 2 \mathrm{HI}(g) \quad[\) slow \(]\) Show that this mechanism is consistent with the rate law.

Express the rate of this reaction in terms of the change in concentration of each of the reactants and products: $$ 2 \mathrm{D}(g)+3 \mathrm{E}(g)+\mathrm{F}(g) \longrightarrow 2 \mathrm{G}(g)+\mathrm{H}(g) $$ When [D] is decreasing at \(0.1 \mathrm{~mol} / \mathrm{L} \cdot \mathrm{s}\), how fast is [H] increasing?

The overall equation and rate law for the gas-phase decomposition of dinitrogen pentoxide are \(2 \mathrm{~N}_{2} \mathrm{O}_{5}(g) \longrightarrow 4 \mathrm{NO}_{2}(g)+\mathrm{O}_{2}(g) \quad\) rate \(=k\left[\mathrm{~N}_{2} \mathrm{O}_{5}\right]\) Which of the following can be considered valid mechanisms for the reaction? I One-step collision II \(2 \mathrm{~N}_{2} \mathrm{O}_{5}(g) \longrightarrow 2 \mathrm{NO}_{3}(g)+2 \mathrm{NO}_{2}(g) \quad[\) slow \(]\) \(2 \mathrm{NO}_{3}(g) \longrightarrow 2 \mathrm{NO}_{2}(g)+2 \mathrm{O}(g)\) [fast] \(2 \mathrm{O}(g) \longrightarrow \mathrm{O}_{2}(g)\) [fast] III \(\mathrm{N}_{2} \mathrm{O}_{5}(g) \rightleftharpoons \mathrm{NO}_{3}(g)+\mathrm{NO}_{2}(g)\) [fast] \(\mathrm{NO}_{2}(g)+\mathrm{N}_{2} \mathrm{O}_{5}(g) \longrightarrow 3 \mathrm{NO}_{2}(g)+\mathrm{O}(g) \quad\) [slow] \(\mathrm{NO}_{3}(g)+\mathrm{O}(g) \longrightarrow \mathrm{NO}_{2}(g)+\mathrm{O}_{2}(g) \quad[\) fast \(]\) \(\mathrm{IV} 2 \mathrm{~N}_{2} \mathrm{O}_{5}(g) \rightleftharpoons 2 \mathrm{NO}_{2}(g)+\mathrm{N}_{2} \mathrm{O}_{3}(g)+3 \mathrm{O}(g) \quad[\) fast \(]\) \(\mathrm{N}_{2} \mathrm{O}_{3}(g)+\mathrm{O}(g) \longrightarrow 2 \mathrm{NO}_{2}(g)\) [slow] \(2 \mathrm{O}(g) \longrightarrow \mathrm{O}_{2}(g) \quad[\) fast \(]\) \(\mathrm{V} \quad 2 \mathrm{~N}_{2} \mathrm{O}_{5}(g) \longrightarrow \mathrm{N}_{4} \mathrm{O}_{10}(g)\) [slow] \(\mathrm{N}_{4} \mathrm{O}_{10}(g) \longrightarrow 4 \mathrm{NO}_{2}(g)+\mathrm{O}_{2}(g)\) [fast]

Acetone is one of the most important solvents in organic chemistry, used to dissolve everything from fats and waxes to airplane glue and nail polish. At high temperatures, it decomposes in a first-order process to methane and ketene \(\left(\mathrm{CH}_{2}=\mathrm{C}=\mathrm{O}\right) .\) At \(600^{\circ} \mathrm{C},\) the rate constant is \(8.7 \times 10^{-3} \mathrm{~s}^{-1}\). (a) What is the half-life of the reaction? (b) How long does it take for \(40 . \%\) of a sample of acetone to decompose? (c) How long does it take for \(90 . \%\) of a sample of acetone to decompose?

Enzymes are remarkably efficient catalysts that can increase reaction rates by as many as 20 orders of magnitude. (a) How does an enzyme affect the transition state of a reaction, and how does this effect increase the reaction rate? (b) What characteristics of enzymes give them this effectiveness as catalysts?
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