Kinetics and Mechanism.





The Collision Theory of Chemical Reactions.



I. A chemical reaction involves the reorganization of atoms. This reorganization requires:



1. A collision between reacting species.



2. The correct orientation for the collision.



3. Energy for the reacting sites to get close, for bonds to be stretched (or broken), for electrons to move into antibonding orbitals.





II. The whole process is entirely reversible.






1. The Collision: Molecularity.


Unimolecular: one species only involved.



Bimolecular: two species collide.



Termolecular: three species collide simultaneously.



Note the statistics means that most reactions will occur by unimolecular or bimolecular processes, the chance of three species colliding simultaneously with the right orientation and enough energy is small. For more that three it is impossibly small.





The ramification of this observation is that complex reactions cannot take place at once: they must involve several successive collisions and atom reorganizations.



Each such collision and reorganization is considered to be a step in the overall process, and writing out the process in the steps gives the step-wise mechanism.



Step-wise mechanisms are theoretical.



More than one such mechanism may be written for a chemical process. Investigations of mechanisms try to discover which set of steps best describes the reaction.








The overall rate of a step-wise process.



Note that in a step-wise process, the steps are successive. Step 1 must occur before step 2 before step 3, etc.



The consequence of this is that the overall rate is no faster than the slowest step in the process, which in effect constitutes a "bottle neck".





In a step-wise process the overall rate of reaction is equal to the rate of the slowest step.





A consequence of this is that only the species reacting in the slowest step dictate the rate of the reaction.





The slowest step of a mechanism is called the rate-determining step.










Use of Kinetics to determine the step-wise mechanism.




Kinetics deals with how fast a chemical reaction is occuring.



In practical terms the rate of a chemical reaction is measured by measuring:



the rate of disappearance of reactant or

the rate of appearance of product:





For product: d[P]/pdt (where p is the coefficient of P in the balanced chemical equation)



For reactant: -d[R]/rdt (where r is the coefficient of R in the balanced chemical equation)


In terms of the collision theory, the rate of a reaction will depend on:


the number of collisions occuring that can lead to a reaction.





Note that to lead to a reaction a collision must:

have the right amount of energy in the right place.

occur in the right orientation.








The Variation of Number of Collisions.


The number of collisions taking place depends on:



1. The concentration of the colliding species



2. The molecularity of the process.









1. Concentration.



As the concentration of reactants drops, the number of collisions will drop, so the rate of the reaction will drop. That is, as a reaction proceeds and the reactants are used up, the rate of the reaction will drop.






2. Molecularity.


When the concentration changes, that change affects each species in the collision:



for a unimolecular process the effect is just for the one species.



for a bimolecular process the effect is for each species, so the effect is squared.



for a termolecular process the effect is cubed.







Because the overall rate of a step-wise reaction depends on the rate-determining step (the slowest step), the way that the rate of reaction changes with concentration changes will depend on the molecularity of that rate-determining step.





By using practical kinetics to give an idea of the molecularity of the rate determining step, some idea of the step-wise mechanism for a reaction can be formulated.










Practical Kinetics.



Experimentally it is found that



rate of reaction (d[P]/pdt) varies as [concentrations]Z.



or



d[P]/pdt = k[concentrations]Z. Equation 1.





[concentrations] refers to a number of reactant (product) concentrations.



Z refers to a power the concentrations are raised to. Z cannot be obtained from a balanced chemical equation, that is the equation for the overall reaction, since as we have deduced the rate of the reaction depends only on the reactants in the slowest step.



Equation 1 is called a Rate Equation.



The constant, k, in equation 1 is called the rate constant.



Z is called the order with respect to a particular reactant.



The sum of all Z values is called the overall order of the reaction.










Typical Rate Equations.








Zero Order: Rate = d[P]/pdt = k
First Order: Rate = d[P]/pdt = k[A]
Second Order: Rate = d[P]/pdt = k[A]2.

Rate = d[P]/pdt = k[A][B]