Kinetics

Reaction Rates

Reaction Rates

Collision theory

Factors affecting reaction ratePotential energy diagrams

temperature

concentration

Surface area

catalystsActivated complex

Activation energy

Reaction Rates

A measure of how quickly a reaction occurs An experimental, measurable quantity Rate = change in property/change in time

Example: Speed = miles/hour Speed = distance/time

Chemical kinetics: the study of reaction rates and the factors that affect them

What could we measure for a reaction?

Easily measured properties include: Change in mass of a solid Change in concentration Temperature changes pH changes Gas volume changes Color changes

We must also measure changes in TIME!

Writing Rate Expressions For a general reaction

aA + bB cC + dD General form:

We need to modify the rate expression to compensate for stoichiometry The reaction has only one rate for a given set of

conditions Convention: all reaction rates are positive

time

Dmoles

dtime

Cmoles

ctime

Bmoles

btime

Amoles

aRate

1111

Measuring Rates

Average Rate

Initial Rate Calculate average rate

for early part of data when plot is nearly linear

time

molesXrateaverage

What happens to the rate over time?

Compare average rate at beginning vs. average rate at end

Reaction rates typically slow down over time Why?

There are fewer moles of reactants left, and therefore fewer collisions.

Collision Theory

Molecules must collide in order to react. They must collide with

the correct orientation. “Effective collision”

Has appropriate orientation; molecules may react.

“Ineffective collision” Doesn’t have needed

orientation; particles will separate.

Collision Theory, cont.

Molecules must collide in order to react. They must have enough

energy to react. Activation Energy, Ea

The minimum energy that reactants must have for the reaction to occur

Potential Energy Diagrams

Activation Energy: from reactants to top of “hill”

Transition State Aka Activated Complex High energy state, where

bonds are broken and new bonds are formed

Hrxn = energy of products – energy of reactants

Potential Energy Diagrams

Which reaction would you expect to be fastest? Slowest? Why?

Reactions with a smaller activation energy will occur

more quickly than reactions with a larger Ea.

Collision Theory

Basic premise: More collisions = faster reaction rate More collisions = greater likelihood for effective

collisions

How can we speed up the rate of a reaction?

Increase temperature Particles move more quickly, so more possible

collisions More particles are likely to have enough energy to

overcome activation energy barrier

Increase concentration More particles, so more possible collisions

Increase surface area More particles are exposed, so more collisions are

possible

Catalysts Speed up reaction rates,

without being consumed Homogeneous vs.

heterogeneous catalysts Enzymes Catalytic RNA Catalytic antibodies Catalytic converter in car

engine Effectively lower the

activation energy of the reaction May even change the

mechanism of the reaction