CHEM 120 Lecture Notes - Lecture 7: Kinetic Theory Of Gases, Collision Theory, Activated Complex

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10 Feb 2016
CHEM 120 – Lecture 7 – Collision Theory & Reaction Mechanisms
14.8: Theoretical Models for Chemical Kinetics: Collision Theory
Collision theory = the rate of a chemical reaction is proportional to the number
of collisions between reactant molecules; the more often reactant molecules collide, the
more often they react with one another, and so the faster the reaction rate is
Kinetic- molecular theory can be used to calculate the collision frequency of a gas
oIn gases at STP, ~10^30 collisions occur per second in 1L! huge!
If each reaction proceeded at that rate/if each collision produced a
reaction, then the rate would be huge! ~10^30/Na!
oActual rates in gas reactions are generally much slower a typical value being
10^-4M/s thus only a small fraction of collisions produce a reaction
Activation Energy (Ea)
Not all collisions result in a chemical reaction (ie. Only collisions between molecules
whose combines KE is above the minimum will react)
For a reaction to happen, there must be a redistribution of energy sufficient to
“reorganize” certain bonds in the reaction molecule(s)
Activation Energy (Ea) = the minimum KE that molecules need to bring to their collisions
for a chemical reaction to
Kinetic energy distribution:
(only some particles react
= shaded area = have
met/overcome Ea)
Shaded area is
proportional to e-Ea/RT,
therefore dependent on
Collision theory related to
Ea & KE: if activation
barrier (Ea) is high, only a
few molecules have sufficient KE to react; orientation of molecules is also very important
for a reaction to occur
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