Molecular Reaction Dynamics

Collision Theory of Kinetics With few exceptions, the reaction rate increases with increasing temperature temperature If we assume that a chemical reaction takes place due to collisions between reactant molecules i.e. rate  number of collisions / unit time A 2 + B 2  product rate = k 2 [A 2 ][B 2 ] “collision of A 2 and B 2 causes a reaction”

Rate constant vs. T

The Reaction Profile

The Activation Energy The minimum amount of energy need for initiation of a chemical reaction is the activation energy (E a ). Colliding reactant molecules possess kinetic energy > the activation energy or E a.

The Activated Complex Activated Complex the species temporarily formed by the reactant molecules as a result of collisions A small fraction of molecules usually have the required kinetic energy to get to the transition state the concentration of the activated complex is extremely small.

The Arrhenius Equation Arrhenius showed how the rate constant depended on temperature. The frequency factor The activation energy

The Arrhenius Equation

Estimating Rate Constants We can use collision theory to provide a basis estimating the rate constant. What if all collisions gave products?

The Energy Requirement Since we know only a fraction of collisions will give products

Converting to Molar Amounts

The Calculated Rate Law

The Geometrical Requirement There are many collisions of sufficient energy that do not yield products Define the steric factor P to account for local properties of the molecule Orientations during collisions Substituent effects

Effective Collisions

Ineffective Collisions

The Steric Factor and k 2 Note values of P can be quite small Usually in the range < 0.001

Limitations of Collision Theory Simple Collision Theory Best suited for studying reactions between simple species (atoms, diatomic molecules). P factor indicates how reactants collide becomes very important when the species get bigger.