AL

Chemistry Mechanisms and Reaction Rates

Molecular Intermediates in Chemical Reactions

  • Intermediate Molecule (c)

    • Formed when reactants (a and b) come together.

    • Reacts further with another b to produce final product (d).

    • Key concept: the molecule (c) can exist separately as it splits from reactants or interacts further in the reaction.

  • Molecularity

    • Unimolecular: Involves one molecule changing to form products.

    • Bimolecular: Involves two molecules (e.g., a and b) reacting together to form an intermediate (c).

    • Trimolecular: Involves three molecules synchronizing to form products.

    • Examples visualize how molecules interact in sequence, contributing to reaction outcomes.

  • Reaction Mechanism

    • Definition: The series of steps that describe the pathway from reactants to products.

    • Rate Order: Determines how the rate of a reaction depends on the concentration of reactants.

    • Breakdown into steps can reveal rate-limiting factors and guide expectations for reaction speed.

  • Mechanistic Pathway Example

    • Step 1: a + b → c (formation of intermediate)

    • Step 2: c + b → d (final product formation)

    • Overall Equation: a + 2b → d

    • Note: Intermediate (c) does not appear in the overall equation as it is consumed during the process.

  • Energy Level Diagram

    • Y-axis: Potential Energy

    • X-axis: Reaction Progress

    • Reactants start at some energy level (R).

    • First activation energy leading to the formation of intermediate (c).

    • Structure contains two 'humps':

    • First Hump: Activation energy (Ea) for the slow step, which is the rate-determining step.

    • Intermediate Energy Level: c is at a higher energy state than reactants but lower than final products.

    • Second Hump: Activation energy for fast step (Ea').

    • Lower than the first hump, indicating a faster reaction speed in step two.

    • Note: Ea' is differentiated to indicate a second activation process not as significant as the first.

  • Stability of Intermediate (c)

    • Intermediate has a higher energy level than reactants but less stability than the final products.

    • Its formation indicates a transient state in the pursuit of reaction completion.

  • Rate Determining Step

    • Significance: The slowest step influences the overall reaction rate.

    • Identifying the rate limiting step is crucial for understanding kinetic behavior and controlling speed of reaction.

  • Calculating Rate Orders

    • Context: Experimental data influences how rate orders are defined.

    • Mechanism’s elementary steps can inform rate orders when backed by experimental observations.

    • Requirement 1: The sum of elementary steps must match the overall balanced chemical equation.

    • Requirement 2: The rate determining step must produce the same rate law as determined from experimentation.

  • Example Reaction

    • Reaction: H₂ + 2 ICl → 2 HCl + I₂

    • Given rate law: rate = k [H₂]¹ [ICl]¹ (second order overall)

    • Upcoming discussions will involve mechanisms that align with this data to confirm validity.

  • Key Takeaways

    • Always compare derived mechanisms with experimental rate laws to validate reactions.

    • Use visual representations, such as energy diagrams, to grasp complexities of chemical kinetics clearly.