CH353 - Physical Chemistry I
Spring 2013, Unique 52575

Lecture Summary, 15 April 2013

Kinetic Rate Laws: The rate of a reaction, v(t), is defined as the change in concentration of any of the species in the reaction as a function of time:
   v(t) = (-1/nu(Reactant))(d[Reactant]/dt) = (1/nu(Product))(d[Product]/dt).

Rate laws are constructed as v(t) = k[A]^m[B]^n...(etc) where m and m are unitless powers that take into account the order of each species in the reaction.  Rate laws must be experimentally determined, they cannot be deduced from the stoichiometrically balanced reaction.  The value of the rate law is always positive, and the rate constant, k, is an experimentally determined number with units necessary to make the units of v(t) mol/Ls.

Rate laws must be experimentally determined.  Today we discussed to methods for doing this, the method of isolation, and the method of initial rates.  From a practical perspective, both methods have their benefits and drawbacks, but both are experimentally quite useful.  However, just because we can put together a rate law from an appropriate set of experiments does not mean we understand how chemical reactions actually occur.  By exploring a simple reaction, the hydrolysis of an acid chloride, we see an example of how a seemingly simple reaction does not just happen in one step.  Instead, it proceeds in a series of discrete steps that each generate molecules that are not observed in the final products.  These molecules are intermediates, unstable, short-lived species that are produced during the reaction but then immediately consumed.  The sequence by which these intermediates are formed and consumed is called the reaction mechanism, which describes the exact physical steps process from moving from reactants to products.  Instead of describing our reactions in a single step, we are going to figure out how to separate reactions into these individual, so called elementary steps, that will help us figure out not just how fast reactions occur, but the physical changes that happen along the reaction coordinate.