Collision Theory of Chemical Reactions

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Collision Theory of Chemical Reactions

This lesson aligns with NGSS PS1.B

Introduction

Chemical reactions are intricate processes that involve the transformation of reactants into products. The Collision Theory, a cornerstone in the realm of chemical kinetics, provides insights into the microscopic events that govern these transformations. This article delves into Collision Theory, exploring its principles, Collision and Activation energy as well as Arrhenius Equation.

Fundamentals of Collision Theory

The Collision Theory, first proposed by Max Trautz and William Lewis in the early 20th century, postulated that the molecules constituting the reactants are considered to be rigid spheres, and reactions are assumed to take place only upon the collision of these spheres (molecules) with each other.

Therefore, it became essential to quantify the occurrences of collisions leading to the product formation in order to gain a comprehensive understanding of the reaction. This necessity gave rise to the term "collision frequency."

Collision Frequency

Collision Frequency is defined as the count of collisions occurring per second within a unit volume of the reacting mixture and is typically symbolized as Z.

Role of Collision Frequency:

The frequency of collisions between reactant molecules is a crucial factor. While not all collisions lead to a reaction, the more frequent and energetic the collisions, the higher the likelihood of successful reactions. Collision frequency is influenced by factors such as temperature and concentration.

Orientation of Collisions:

Molecules must collide in a specific orientation for a reaction to occur. Not all collisions result in a successful reaction, as proper alignment is essential. To illustrate, let's examine the subsequent bimolecular elementary reaction:

P+Q→Product

According to the collision theory, the rate of the above reaction can be expressed as:

Here:

[math]Z_(PQ)[/math] denotes the collision frequency of reactants P and Q.
[math]E_a[/math] represents the Activation Energy.
R is the Universal Gas Constant.
T signifies the Temperature in absolute scale.
ρ is the steric factor.

Additionally, another significant parameter that plays a role in the context of collision theory can be identified.

Another critical factor influencing the rates of chemical reactions is activation energy ([math]E_a[/math]). This was given by Arrhenius, activation energy refers to the minimum energy requirement that reactants must attain for the successful formation of products during a chemical reaction.

In accordance with the Arrhenius Equation, molecules possessing energy equal to or surpassing the activation energy will engage in collisions leading to product formation. However, it became evident that this postulate did not universally hold true for all reactions, as substantial deviations were observed, particularly in reactions involving complex molecules.

In a nutshell, we can say that in a collision theory, both activation energy and effective collision govern the rate of a reaction.

Collisions and Activation Energy:

According to the Collision Theory, for a chemical reaction to take place, colliding molecules must possess energy equal to or greater than the activation energy.

Activation energy is defined as the minimum energy required by the particles involved in a specific reaction for the reaction to occur. In the absence of particles colliding with adequate energy to fulfill the activation energy, the reaction remains inactive. Thus, the activation energy must be provided before the initiation of a reaction.

Initiating a chemical reaction involves the breaking of chemical bonds within the reactants, a process that demands energy. The energy necessary to commence this reaction is referred to as activation energy. In certain scenarios, the activation energy is relatively low, allowing the reaction to commence at room temperature without the need for additional heating.

Arrhenius Equation

Temperature has a significant influence on the rate of chemical reactions, with a consistent trend of increasing rates due to rise in temperatures. Typically, the rate of chemical reactions tends to double with each 10 K increment in temperature. Although in some instances, this augmentation may surpass a twofold increase. The temperature coefficient, representing the ratio of rate constants at two distinct temperatures, commonly falls within the range of 2 to 3, calculated as:

Svante Arrhenius formulated a mathematical relationship between temperature and the rate constant. This dependency is expressed algebraically as follows:

This equation is Known as the Arrhenius equation. Here A is known as frequency factor. Ea is the activation energy , R is gas constant and T is the absolute temperature.

Summary

The Collision Theory postulated that the molecules constituting the reactants are considered to be rigid spheres, and reactions are assumed to take place only upon the collision of these spheres (molecules) with each other.
Collision Frequency is defined as the count of collisions occurring per second within a unit volume of the reacting mixture, and is typically symbolized as Z.
Collision Frequency between reactant molecules is a crucial factor. The more frequent and energetic the collisions, the higher the likelihood of successful reactions.
Activation energy is defined as the minimum energy required by the particles involved in a specific reaction for the reaction to occur.

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