Transcript Kinetics and Thermodynamics of Simple Chemical Processes

2-1
Kinetics and Thermodynamics of Simple
Chemical Processes
Chemical thermodynamics: Is concerned with the extent that a
reaction goes to completion.
Chemical kinetics: Is concerned with the speed that a reaction
goes to completion.
Thermodynamic Control: The most stable products are formed.
Kinetic Control: The fastest product is formed.
Equilibria are governed by the thermodynamics
of chemical change.
When the concentrations of reactants and products no longer
change, the system is said to be at equilibrium.
A system at equilibrium is described mathematically:
A large value of K indicates that the reaction goes to
completion.
The equilibrium constant can be related to the
thermodynamic function Go:
Go  RT ln K
When Go is negative, the equilibrium constant is greater
than 1 and the products of the reaction are favored over the
reactants.
When Go is positive, the equilibrium constant is less that 1
and the reactants of the reactions are favored over the
products.
The free energy change is related to changes in
bond strengths and the degree of order in the
system.
The Gibbs free energy change for a reaction is related to the
enthalpy change and the entropy change for the reaction:
G  H  T S
o
o
o
The enthalpy change, Ho, can be estimated:
H o 
  Bond Energies     Bond Energies 
Reactants
Products
The entropy change, So, is related to the amount of disorder in
the system. The entropy of a substance phase is much larger in
the gas than in the liquid phase.
In a chemical reaction where all substances are in the same
phase, the entropy will increase if there are more product
molecules than reactant molecules.
The rate of a chemical reaction depends on the
activation energy.
The potential energy of the system as a chemical change occurs
can be plotted vs. time:
The higher the activation energy, Ea, the slower the reaction.
Collisions supply the energy to get past the
activation-energy barrier.
The average kinetic energy of molecules at room temperature is
about 0.6 kcal mol-1. The kinetic energies of individual molecules
can be plotted as a Boltzmann distribution curve:
As can be seen from the curves, there are more molecules having large
kinetic energies at high temperature than at low temperature.
Since the energy required to reach the transition state in a chemical
reaction comes from molecular collisions, the rate of chemical reactions
always increases with rising temperatures.
The concentration of reactants can affect reaction
rates.
The rate of a chemical reaction can be expressed as a rate law.
The rate law must be experimentally determined; it cannot be
derived directly from the balanced chemical equation.
First Order Reaction:
A  B, Rate = k[A]
Second Order Reaction:
A + B  C + D, Rate = k[A][B]
The Arrhenius equation describes how temperature
affects reaction rates.
The rate constant, k, depends upon temperature according to the
Arrhenius equation:
k  Ae
 Ea / RT
In general, raising the reaction temperature by 10 oC will increase
the rate constant by a factor of 2 or 3.