Chemical Kinetics and Catalysis: Mechanisms, Rate Laws, and Energy Profiles

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Chemical Kinetics and Reaction Mechanisms

Elementary Reactions and Molecularity

Chemical reactions can occur in a single step or through multiple steps. An elementary reaction is a reaction that happens in one step, and its molecularity refers to the number of reactant molecules involved in that step.

Unimolecular: Involves one reactant molecule. Example: Cl2(g) → 2 Cl(g)
Bimolecular: Involves two reactant molecules. Example: NOCl(g) + Cl(g) → NO(g) + Cl2(g)
Termolecular: Involves three reactant molecules. These are rare.

The rate law for an elementary reaction directly follows its stoichiometry.

Molecularity	Elementary Reaction	Rate Law
Unimolecular	A → products	Rate = k[A]
Bimolecular	A + A → products	Rate = k[A]2
Bimolecular	A + B → products	Rate = k[A][B]
Termolecular	A + A + A → products	Rate = k[A]3
Termolecular	A + A + B → products	Rate = k[A]2[B]
Termolecular	A + B + C → products	Rate = k[A][B][C]

Table of elementary reactions and their rate laws

Reaction Mechanisms

A reaction mechanism is the sequence of elementary steps that make up the overall reaction. Each step may produce or consume intermediates, which are species that appear in some steps but not in the overall reaction equation.

Intermediates: Produced in one step and consumed in another; do not appear in the overall reaction.
Overall Reaction: The sum of all elementary steps.

Reaction mechanism with intermediate species

Rate-Determining Step and Multistep Mechanisms

In multistep reactions, the rate-determining step is the slowest step, which controls the overall reaction rate. The observed rate law is often determined by this step.

If the first step is slow, the rate law matches the first step.
If a later step is slow, earlier steps may affect the concentrations of intermediates, leading to more complex rate laws.

Toll plaza analogy for rate-determining step

Energy Profiles and Reaction Progress

Potential Energy Diagrams

An energy profile shows the potential energy changes during a reaction. Multistep reactions have multiple maxima (activation energies) and minima (intermediates).

Activation Energy (Ea): The energy barrier for each step.
Intermediates: Correspond to valleys between peaks.
Rate-Determining Step: The step with the highest activation energy.

Energy profile for multistep reactions Potential energy profile for a multi-step reaction Potential energy diagram with intermediates and rate-determining step

Catalysis

Definition and Types of Catalysts

A catalyst is a substance that increases the speed of a chemical reaction without being consumed. Catalysts are present at the start and regenerated at the end of the mechanism.

Homogeneous Catalyst: Same phase as reactants (e.g., liquid, gas).
Heterogeneous Catalyst: Different phase from reactants (usually solid).

Definitions and types of catalysts

How Catalysts Work

Catalysts provide an alternative reaction pathway with a lower activation energy, increasing the reaction rate.

Catalysts do not affect the overall thermodynamics (ΔE) of the reaction.
They are consumed in one step and regenerated in another.

Catalyst lowers activation energy in reaction pathway

Examples of Catalysis

Homogeneous Catalysis: Decomposition of Hydrogen Peroxide

The decomposition of H2O2 can be catalyzed by Br- and H+ ions, which are regenerated at the end.

Overall Reaction: 2 H2O2(aq) → 2 H2O(l) + O2(g)
Intermediates: Br2(aq)
Catalysts: Br-(aq), H+(aq)

Catalyzed decomposition of hydrogen peroxide Potential energy diagram for catalyzed and uncatalyzed reactions

Heterogeneous Catalysis: Hydrogenation of Ethene

Heterogeneous catalysts, such as metals, facilitate reactions on their surfaces. For example, the hydrogenation of ethene to ethane occurs on a metal catalyst.

Reactants adsorb onto the catalyst surface.
Reaction occurs via surface intermediates.
Products desorb from the surface.

Heterogeneous catalysis mechanism

Catalytic Converters

Catalytic converters in automobiles use heterogeneous catalysts to convert harmful gases into less toxic products.

CO(g) or CxHy(g) + O2(g) → CO2(g) + H2O(g)
NO(g) or NO2(g) → N2(g) + O2(g)
Common catalysts: Pt, Pd, Rh

Catalytic converter mechanism

Enzymes: Biological Catalysts

Enzymes are biological catalysts that operate via a "lock and key" mechanism, binding substrates and lowering activation energy by straining bonds.

Highly specific for their substrates.
Examples: Amylase (converts starches to sugars).

Lock and key mechanism of enzyme catalysis

Applications and Environmental Chemistry

CO2 Capture and Metal-Organic Frameworks (MOFs)

CO2 capture is important for climate change mitigation and fuel generation. MOFs are porous materials designed to bind CO2 efficiently, even at high temperatures.

Challenges: Weak dispersion forces, low-temperature capture, energetic inefficiency.
Research focuses on creating reactive sites within MOFs to bind CO2 strongly.

Structure of a MOF for CO2 capture CO2 binding to MOF structure CO2 capacity of MOF at different temperatures

Catalysis in the Atmosphere: Ozone Depletion

Catalytic mechanisms are involved in atmospheric chemistry, such as the destruction of ozone by chlorofluorocarbons (CFCs).

Ozone protects against UV radiation.
Catalytic cycles involving Cl radicals lead to ozone depletion.

Summary

Elementary reactions and their rate laws are fundamental to understanding reaction kinetics.
Reaction mechanisms consist of sequences of elementary steps, with intermediates and rate-determining steps.
Energy profiles illustrate activation energies and intermediates in multistep reactions.
Catalysts increase reaction rates by lowering activation energy and are crucial in industrial, environmental, and biological processes.