In chemical kinetics, the concepts of reaction order and molecularity play a crucial role in understanding how reactions proceed. For elementary reactions, these two terms are often considered equal. But what does this mean, and why does it hold true? This content will explore the definitions of order and molecularity, their differences in complex reactions, and why they are equal in elementary reactions.
Understanding Molecularity
Definition of Molecularity
Molecularity refers to the number of reactant molecules involved in a single step of a reaction. It is always a whole number because it represents the actual number of molecules that collide to initiate a reaction.
Types of Molecularity
-
Unimolecular Reaction (Molecularity = 1)
- Involves a single reactant molecule undergoing decomposition or rearrangement.
- Example: The decomposition of ozone (O3 ? O2 + O).
-
Bimolecular Reaction (Molecularity = 2)
- Involves the collision of two reactant molecules.
- Example: The reaction between nitrogen dioxide and carbon monoxide (NO2 + CO ? NO + CO2).
-
Termolecular Reaction (Molecularity = 3)
- Involves three molecules colliding simultaneously.
- These are rare because the probability of three molecules colliding at the same time is low.
- Example: The formation of nitric oxide (2NO + O2 ? 2NO2).
Understanding Reaction Order
Definition of Reaction Order
The order of a reaction refers to the sum of the powers of the concentration terms in the rate law. It determines how the rate of reaction depends on the concentration of reactants.
General Form of a Rate Law
For a reaction:
The rate law is expressed as:
where:
- k is the rate constant
- [A] and [B] are the concentrations of reactants
- m and n are the reaction orders with respect to each reactant
- The overall order is m + n
Why Are Order and Molecularity Equal for Elementary Reactions?
Characteristics of Elementary Reactions
An elementary reaction is a single-step reaction where reactants directly form products without intermediate steps. Since the reaction occurs in one step, the molecularity (number of colliding molecules) directly determines the rate law.
For an elementary reaction:
- The rate law is derived directly from the molecularity.
- The reaction order is equal to the sum of the exponents in the rate law, which matches the molecularity.
Examples of Elementary Reactions
-
Unimolecular Reaction
A ? Products- Molecularity = 1 (only one molecule involved)
- Rate law: Rate = k[A]
- Order = 1 (matches molecularity)
-
Bimolecular Reaction
A + B ? Products- Molecularity = 2 (two molecules collide)
- Rate law: Rate = k[A][B]
- Order = 2 (matches molecularity)
-
Termolecular Reaction
A + B + C ? Products- Molecularity = 3 (three molecules collide)
- Rate law: Rate = k[A][B][C]
- Order = 3 (matches molecularity)
When Do Order and Molecularity Differ?
Complex Reactions
For non-elementary or complex reactions, the reaction does not occur in a single step. Instead, it proceeds through multiple intermediate steps. In such cases:
- The rate-determining step (slowest step) controls the overall reaction rate.
- The reaction order is determined experimentally, not from molecularity.
- Reaction order can be a fraction or zero, unlike molecularity, which is always a whole number.
Example of a Complex Reaction
Consider the decomposition of hydrogen peroxide:
The rate law is:
- The order is 1, but molecularity cannot be directly assigned because the reaction proceeds through intermediate steps.
Key Differences Between Order and Molecularity
| Feature | Molecularity | Order of Reaction |
|---|---|---|
| Definition | Number of molecules colliding in an elementary step | Sum of exponents in the rate law |
| Value Type | Always a whole number (1, 2, or 3) | Can be a fraction, zero, or whole number |
| Determination | Based on reaction mechanism | Determined experimentally |
| Applicability | Only for elementary reactions | Applies to both elementary and complex reactions |
For elementary reactions, the reaction order and molecularity are always equal because the reaction occurs in a single step. This direct relationship makes it possible to determine the rate law simply by analyzing the reaction mechanism. However, for complex reactions, the order must be determined experimentally, as it depends on the rate-determining step rather than the total number of molecules colliding. Understanding these principles is fundamental in chemical kinetics, helping scientists predict reaction rates and optimize industrial processes.