Glossary

6.2 Transition state theory fundamentals

3 min read•july 22, 2024

Theory explains how chemical reactions occur through an . This high-energy intermediate forms when reactants collide with enough energy, determining the reaction rate. Understanding this process helps predict and control reaction speeds.

The theory postulates a quasi-equilibrium between reactants and the activated complex. It differs from by focusing on the complex's formation and decomposition. Calculating rate constants and understanding 's role are key aspects of this theory.

Transition State Theory Fundamentals

Activated complex in reactions

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  • High-energy, unstable intermediate formed during a chemical reaction when reactant molecules collide with sufficient energy and proper orientation
  • Located at the highest point on the diagram, representing the maximum potential energy
  • Critical point in a that determines the rate of the reaction
    • Point at which the reactants are partially converted into products
    • Rate of the reaction depends on the concentration of the activated complex and the rate at which it decomposes into products (products could be molecules like H2O or NH3)

Transition state theory postulates

  • Reactants and activated complex are in quasi-equilibrium
    • Concentration of the activated complex is proportional to the concentrations of the reactants
  • Activated complex can convert into products or revert back to reactants
    • Rate of product formation depends on the rate of decomposition of the activated complex
  • Activated complex passes through the transition state only once, does not oscillate back and forth around the transition state
  • Differs from collision theory:
    • Considers formation of an activated complex, while collision theory does not
    • Assumes quasi-equilibrium between reactants and activated complex, while collision theory does not consider equilibrium
    • Focuses on decomposition of activated complex into products, while collision theory emphasizes collision frequency and orientation of reactant molecules (like two billiard balls colliding)

Rate constant calculation methods

  • Arrhenius equation relates (kk) to activation energy (EaE_a) and temperature (TT): k=Ae−Ea/RTk = A e^{-E_a/RT}
    • AA is or frequency factor, represents frequency of collisions with proper orientation
    • RR is universal gas constant (8.314 J mol−1^{-1} K−1^{-1})
  • To calculate rate constant:
    1. Determine activation energy (EaE_a) and pre-exponential factor (AA) experimentally or from literature
    2. Substitute values of EaE_a, AA, RR, and TT (in Kelvin) into Arrhenius equation
    3. Solve equation for kk, the rate constant

Activation energy and reaction rates

  • Activation energy (EaE_a) is minimum energy required for reactants to form activated complex
    • Higher activation energy means fewer reactant molecules will have sufficient energy to form activated complex
  • Rate of chemical reaction is inversely proportional to activation energy
    • As activation energy increases, rate of reaction decreases because smaller fraction of reactant molecules will have enough energy to overcome activation energy barrier
  • Relationship between activation energy and reaction rate is exponential, as described by Arrhenius equation
    • Small change in activation energy can lead to significant change in rate constant and reaction rate (doubling EaE_a could decrease rate by factor of 10)

Key Terms to Review (14)

Activated Complex: An activated complex is a transient, unstable arrangement of atoms that forms during a chemical reaction, representing the peak energy state along the reaction pathway. This complex exists between reactants and products, indicating that the reactants have sufficient energy to overcome the activation barrier, allowing for the transformation into products. Understanding this concept is crucial for grasping how reactions proceed and why certain conditions influence reaction rates.
Activation Energy: Activation energy is the minimum amount of energy required for a chemical reaction to occur. It represents the energy barrier that reactants must overcome to be transformed into products, linking the concepts of kinetics and thermodynamics in the context of chemical reactions.
Collision Theory: Collision theory is a fundamental concept in chemical kinetics that explains how reactions occur and why reaction rates vary. It posits that for a reaction to take place, reactant particles must collide with sufficient energy and the correct orientation. This theory connects to various aspects of reaction dynamics, including the interplay between energy, molecular structure, and the speed of reactions.
Electronic effects: Electronic effects refer to the influence of electron distribution in a molecule on its reactivity and stability, particularly during chemical reactions. These effects are crucial in understanding how substituents on a molecule can either stabilize or destabilize transition states, impacting the rate and mechanism of a reaction. By analyzing electronic effects, chemists can predict how changes in molecular structure can affect the energy barrier to reaction, which is a key element in transition state theory.
Orientation Factor: The orientation factor is a numerical value that quantifies the likelihood of reactant molecules properly aligning during a collision to result in a successful reaction. This factor is crucial in understanding how the geometry of the colliding species impacts reaction rates and pathways, linking it to the broader concepts of molecular interactions and energy barriers.
Potential Energy Surface: A potential energy surface (PES) is a multi-dimensional graphical representation of the potential energy of a system as a function of its molecular configurations, which is crucial for understanding chemical reactions. The PES illustrates how the energy changes with different arrangements of atoms and helps identify stable states, transition states, and reaction pathways, linking directly to concepts like reaction coordinates, collision dynamics, and quantum chemical calculations.
Pre-exponential factor: The pre-exponential factor, often denoted as A in the Arrhenius equation, represents the frequency of collisions or interactions between reactant molecules that lead to a reaction. It reflects the likelihood of reactants overcoming the energy barrier for a reaction to occur, serving as a crucial component in determining the rate constant of a reaction. This factor is influenced by factors such as temperature, molecular orientation, and the nature of the reactants, playing a significant role in understanding reaction kinetics across various environments.
Rate Constant: The rate constant is a proportionality factor in the rate law that quantifies the speed of a chemical reaction at a given temperature. It connects the concentration of reactants to the reaction rate, showing how quickly the reaction proceeds. The value of the rate constant is influenced by factors such as temperature, activation energy, and the presence of catalysts, making it a key element in understanding reaction kinetics and dynamics.
Rate-determining step: The rate-determining step is the slowest step in a reaction mechanism that limits the overall rate of the reaction. This concept is crucial because it helps to identify which specific reaction step controls how fast a chemical reaction occurs, influencing the formulation of rate laws and determining the kinetics of reactions.
Reaction Coordinate: A reaction coordinate is a hypothetical construct that represents the progress of a chemical reaction, typically illustrating the energy changes that occur as reactants transform into products. It serves as a way to visualize the transition states and intermediates involved in a reaction, making it essential for understanding the kinetics and thermodynamics of chemical processes.
Reaction pathway: A reaction pathway is the series of steps that a chemical reaction follows from reactants to products, including all intermediate states and transitions. Understanding this pathway is crucial for predicting how reactants transform into products, which can involve changes in energy and molecular structure. The pathway is often represented visually through potential energy surfaces or reaction coordinate diagrams, helping to elucidate the mechanisms behind various reactions.
Steric hindrance: Steric hindrance refers to the repulsion that occurs when atoms or groups within a molecule are forced close together, leading to a decrease in the reactivity of the molecule. This phenomenon is crucial in determining how easily a reaction can occur, as larger substituents can block access to reactive sites and impede the transition state formation necessary for a chemical reaction.
Thermodynamic versus Kinetic Control: Thermodynamic versus kinetic control refers to the two different ways that chemical reactions can be governed, with thermodynamic control focusing on the stability of products and kinetic control concentrating on the pathway and rate of formation. In thermodynamic control, the most stable product is favored at equilibrium, while in kinetic control, the product that forms the fastest under the given conditions is prioritized. Understanding these concepts is crucial in analyzing reaction mechanisms and predicting product distributions.
Transition state: The transition state is a temporary, high-energy arrangement of atoms that occurs during a chemical reaction, representing the point of maximum energy along the reaction pathway. This state is crucial as it determines the activation energy required for the reaction to proceed and connects the reactants and products through an energy barrier.
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