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 (k) to activation energy (Ea) and temperature (T): k=Ae−Ea/RT
A is or frequency factor, represents frequency of collisions with proper orientation
R is universal gas constant (8.314 J mol−1 K−1)
To calculate rate constant:
Determine activation energy (Ea) and pre-exponential factor (A) experimentally or from literature
Substitute values of Ea, A, R, and T (in Kelvin) into Arrhenius equation
Solve equation for k, the rate constant
Activation energy and reaction rates
Activation energy (Ea) 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 Ea could decrease rate by factor of 10)