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Simple Explanation:
Introduction to Elementary Steps
In chemical reactions, there are individual steps that take place. Unlike an equation for an overall reaction, elementary processes show what happens to individual particles in a reaction using representations called elementary steps.
Molecularity
Unimolecular Step: When a particle breaks down (decomposes).
AB → A + B, Rate = k[AB]
Bimolecular Step: When two particles collide and react.
AB + C → AC + B, Rate = k[AB][C]
A + A → AA, Rate = k[A]²
Termolecular Step: When three particles react simultaneously (rare due to low probability).
Concepts of Elementary Processes
Rate-Determining Step: In reactions with multiple elementary steps, the slowest step determines the rate of the reaction since the reaction can only proceed as fast as its slowest elementary process. This step is called the rate-determining step, and its position in the sequence of elementary processes for a reaction changes how it could be used to find the rate law.
The slowest elementary step requires the most activation energy.
Intermediate Species: The intermediate species of a reaction are substances that are produced in one step and then is consumed in an upcoming step, which means that it doesn't appear as a final reactant or product in the overall reaction.
Example with a Potential Energy Diagram
Eₐ (Activation Energy): The amount of energy needed to initiate a reaction.
ΔH (Enthalpy): The amount of heat energy gained/lost in a reaction (measured in KJ/mol).
The peaks on the diagram represent the transition states, where chemical bonds start to break and form, indicating elementary steps.
If the potential energy of products are lower than the reactants, the overall reaction would be exothermic. Otherwise, it would be endothermic.
The hump with the highest Eₐ is the slowest step in the reaction mechanism.
Examples
Mechanism Example #1:
F₂ → 2F (slow)
F + H₂ → HF + H (fast)
Solve for overall reaction equation: The intermediate species is F, so all of the F in the elementary steps cancel out. We then combine what is left of the processes to determine the overall reaction equation. The remaining reactants are F₂ and H₂.
Balanced Overall Reaction: F₂ + H₂ → 2HF
Solve for rate law: We must first find the rate-determining step, which is F₂ → 2F. Since the step is the first of the sequence, we then examine the reactants of that step to construct our rate law.
Rate Law: Rate = k[F₂]
Mechanism Example #2:
2NO → N₂O₂ (fast)
N₂O₂ + O₂ → 2NO₂ (slow)
Solve for overall reaction equation: The intermediate species is N₂O₂, so all of them in the elementary steps cancel out. We then combine what is left of the processes to determine the overall reaction equation. The remaining reactants are 2NO and O₂.
Balanced Overall Reaction: 2NO + O₂ → 2NO₂
Solve for rate law: We must first find the rate-determining step, which is N₂O₂ + O₂ → 2NO₂. Since the step is the last of the sequence, we need to substitute for the intermediate species of the reaction:
Before substitution: Rate = k[N₂O₂][O₂]
Identify intermediate species: N₂O₂
From 2NO → N₂O₂, we know that 2NO makes N₂O₂, so we substitute [N₂O₂]: Rate = k[NO][NO][O₂]
Simplify: Rate = k[NO]²[O₂]
Rate Law: Rate = k[NO]²[O₂]
Mechanism Example #3:
I₂ → 2I (fast)
I + H₂ → H₂I (fast)
H₂I + I → 2HI (slow)
Solve for overall reaction equation: The intermediate species are H₂I and I, so all of them in the elementary steps cancel out. We then combine what is left of the processes to determine the overall reaction equation. The remaining reactants are I₂ and H₂.
Balanced Overall Reaction: I₂ + H₂ → 2HI
Solve for rate law: We must first find the rate-determining step, which is H₂I + I → 2HI. Since the step is the last of the sequence, we need to substitute for the intermediate species of the reaction:
Before substitution: Rate = k[H₂I][I]
Identify intermediate species: H₂I and I
From I + H₂ → H₂I, we know that I and H₂ makes H₂I, so we substitute [H₂I]: Rate = k[H₂][I][I]
Simplify: Rate = k[H₂][I]²
From I₂ → 2I, we know that I₂ makes 2I, so we substitute [I]²: Rate = k[H₂][I₂]
Rate Law: Rate = k[H₂][I₂]
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