Chapter 12
Chemical Kinetics
AP CHEMISTRY
MS. GROBSKY
Learning Objectives
 Describe the collision theory of reaction rates
 Use collision theory to describe the effect of reactant
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


concentration on reaction rate
Describe the effect of temperature on reaction rate using
the collision theory of reaction rates
Explain the concept of reaction rate
Derive the average and instantaneous rates of a reaction
from experimental information
Describe factors that affect reaction rate
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Concentration
Temperature
Catalyst
State of reactants
What Do We Mean When We Say
Chemical Kinetics?
 Principally interested in:
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

The rate of a chemical reaction
The factors that influence the rate
The mechanism by which a reaction takes place
 What do we mean by rate?

Change in concentration of a reactant or product per unit time


Can measure concentration spectroscopically (color change) and
production of a gas (pressure change)
Can be positive, negative or zero
Positive – concentration is increasing
 Negative – concentration is decreasing
 Zero – concentration is constant
𝑐ℎ𝑎𝑛𝑔𝑒 𝑖𝑛 𝑐𝑜𝑛𝑐𝑒𝑛𝑡𝑟𝑎𝑡𝑖𝑜𝑛 [𝐴] [𝐴]𝑓 −[𝐴]𝑖
𝑅𝑎𝑡𝑒 =
=
=
𝑐ℎ𝑎𝑛𝑔𝑒 𝑖𝑛 𝑡𝑖𝑚𝑒
∆𝑡
𝑡𝑓 − 𝑡𝑖

Collision Theory of Reactions
 Collision theory is a model that accounts for the observed characteristics of
reaction rates
 It states that for a reaction to occur:


Particles must collide
 Only two particles may collide at one time
Particles must have the correct geometry


Proper orientation of colliding molecules so that atoms can come in contact with each other to
become products
Collision must involve enough energy to produce the reaction
 Must overcome the electron/electron repulsion of the valence shell electrons of
reacting species
 Transform translational energy into vibrational energy in order to penetrate into
each other so that the electrons can rearrange and form new bonds
 The collision must equal or exceed the activation energy, Ea
 New collision product is at the peak of the activation energy hump and is called
the activated complex (transition state). At this point, the activated complex
can still either fall to reactants or to products
 Bottom line is that all collisions do NOT result in reactions!
Collision Theory of Reaction Rates
Factors that Affect Reaction Rates
 Nature of the reactants
 Some reactant molecules react in a hurry, others slowly

Physical state
 Gasoline (l) vs. gasoline (g)
 K2SO4 (s) + Ba(NO3)2 (s) → No Reaction!
• Will react in aqueous state

Chemical identity – what is reacting?
 Generally, ions of opposite charge react very rapidly
 Generally, the more bonds between reacting atoms in a
molecule, the slower the reaction rate
• Strong bonds react much more slowly
Factors that Affect Reaction Rates
 Concentration of reactants

More reactants, more collisions
 The possibility of a successful
collision increases!
 Temperature



“Heat ‘em up, speed ‘em up!”
 The faster they move, the more
likely they are to collide
Does not affect activation energy
 However, more particles now
have sufficient energy to
overcome the activation energy.
 Therefore, there is a general
increase in reaction rate with
increasing temperature
General rule of thumb: a 10°C
increase in temperature will
DOUBLE the reaction rate
Factors that Affect Reaction Rate
 Catalysts




Increase rate but are not used up
 Not part of the chemical reaction
 Regenerated
Change the rate by providing an
alternative reaction mechanism with
a different activation energy
Positive catalysts
 Increase reaction rate, lower Ea
 Ex. H2O2 decomposes
relatively slowly into H2O and
O2; however, exposure to light
accelerates this process AND
with the help of MnO2, it goes
extremely FAST!
Negative catalysts
 Decrease reaction rate, increase
Ea
 Ex. Food preservatives!
Factors that Affect Reaction Rate
 Surface area of reactants
 Exposed surfaces affect speed


Except for substances in gaseous state or solutions
 Reactions then occur at the boundary, or interface, between two
phases
The greater the surface area exposed, the greater change of
collisions between particles
Hence, the reaction should proceed at a much faster rate
 Ex. Coal dust is very explosive as opposed to a piece of charcoal

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Solutions have the largest exposure!
Inert gases do not affect reaction rate

Do not appear in rate law (more on this later!)
Chemical Reaction Rates
 The speed of a reaction is expressed in terms of its
“rate”

Some measurable quantity is changing with time
 Can be written in terms of reactant(s) disappearance
or product(s) appearance
𝑐ℎ𝑎𝑛𝑔𝑒 𝑖𝑛 𝑐𝑜𝑛𝑐𝑒𝑛𝑡𝑟𝑎𝑡𝑖𝑜𝑛 [𝐴] [𝐴]𝑓 −[𝐴]𝑖
𝑅𝑎𝑡𝑒 =
=
=
𝑐ℎ𝑎𝑛𝑔𝑒 𝑖𝑛 𝑡𝑖𝑚𝑒
∆𝑡
𝑡𝑓 − 𝑡𝑖
Chemical Reaction Rates
 Rate is not constant; it
changes with time
 Can be positive or negative


Product ALWAYS positive
Reactant ALWAYS negative
 Graphing the data of an
experiment will show an
average rate of reaction
 You can find the
instantaneous rate by
computing the slope of a
straight line tangent to the
curve at that time
Relative Reaction Rates
 We can consider the appearance of products along
with the disappearance of reactants
 Reactant’s concentration is declining; the products is
increasing
 Read the balanced equation

Divide the rate of change in concentration of each reactant by
its stoichiometric coefficient
aA + bB → cC + dD
1 𝐴
1 𝐵
1 𝐶
1 [𝐷]
𝑅𝑎𝑡𝑒 = −
=−
= +
=+
𝑎 ∆𝑡
𝑏 ∆𝑡
𝑐 ∆𝑡
𝑑 ∆𝑡
Chemical Reaction Rates
 Ex.
2 NO (g) + O2 (g) → 2 NO2 (g)
 Oxygen
can disappear
only half as rapidly as
the nitrogen monoxide
disappears while NO2
appears twice as fast as
oxygen disappears
Practice
 Because it has a nonpolluting product (water vapor),
hydrogen gas is used for fuel aboard the space
shuttle and may be used by Earth-bound engines in
the near future.
2H2 (g) + O2 (g) → 2H2O (g)


Express the rate in terms of changes in [H2], [O2], and [H2O]
with time
When [O2] is decreasing at 0.23 mol/L·s, at what rate is [H2O]
increasing?