Chapter 12 - Chapter
12
Study
Guide


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Unformatted text preview: Chapter
12
Study
Guide
 Reaction
Rate:
the
rate
of
the
reaction,
how
quickly
a
reaction
will
go
to
 completion
(Example:
Think
of
it
as
MPH
for
chemicals)
 Collision
Theory:
theory
that
states
in
order
for
a
reaction
to
occur,
there
must
 be
a
collision
of
molecules,
molecules
must
have
enough
energy
to
overcome
 the
“energy
barrier”,
and
molecules
must
be
in
proper
orientation
 Activation
Energy:
Minimum
amount
of
energy
needed
to
overcome
the
 energy
barrier
 (Ex:
The
min.
 Amount
of
work
 You
need
to
do
to
 Pass
a
class.)
 Things
that
affect
reaction
 rates:
 Concentration:

increasing
concentration
increases
reaction
rates,
because
a
 larger
number
of
collisions
occur.
 Surface
area:

increasing
surface
area
of
a
solid
increases
reaction
rates
 because
as
surface
area
increases,
the
number
of
atoms
exposed
increases
as
 well,
allowing
more
collisions.
 Temperature:

a
temperature
increase
increases
the
average
kinetic
energy.

A
 increase
in
kinetic
energy
causes
more
collisions
to
occur,
and
also
increases
 the
fraction
of
effective
collisions.


 
 Reason
fraction
of
effective
collisions
increases:

the
increase
in
kinetic
 energy
allows
more
molecules
to
break
the
energy
barrier.

In
other
words,
 more
molecules
meet
the
activation
energy
requirement.
 Fraction
of
effective
collisions:

(the
number
of
effective
collisions)/(total
 collisions)
 Catalyst:
something
that
alters
the
pathway
in
which
a
reaction
occurs
without
 itself
being
consumed
in
the
reaction.
(Example:
Think
about
this.
You
need
to
 travel
a
10
mile
road
which
curves.
A
catalyst
would
be
like
a
shortcut
which
 cuts
6
miles
off
the
original
distance
but
gets
you
to
the
same
place.)
 
 Intermediate:

something
that
forms
temporarily
during
a
reaction,
 accompanying
a
catalyst.
 
 Ex:

Step
1)
O3(g)
+
Cl(g)
­>
ClO(g)
+
02
 
 Step
2)

ClO(g)
+
O3(g)
­>
Cl(g)
+
2O2(g)
 Overall:
2O3(g)
+
Cl(g)
+
ClO(g)
­>
ClO(g)
+
Cl(g)
+
3O2(g)
 Net:

2O3(g)

­­>

3O2(g)
 ClO
is
the
intermediate
in
this
reaction,
and
Cl
is
the
catalyst.

Cl
is
in
both
 reactants
and
products,
so
it
is
not
consumed
in
the
reaction.

ClO
is
the
 intermediate
because
it
is
temporarily
created
during
the
reaction
and
then
 consumed
in
Step
2.
 Chemical
Equilibrium
 Chemical
equilibrium:
state
in
a
chemical
reaction
where
concentrations
of
 the
reactants
and
products
remain
constant.

Products
are
converted
to
 reactants
at
an
equal
rate
as
reactants
are
converted
to
products.
(Example:
 The
mall
with
two
floors.
You
have
the
same
#
of
people
on
both
floors.
 However,
two
people
decide
to
leave
the
2nd
floor
while
two
people
come
up
 there.
As
long
as
the
number
is
the
same
on
both
floors,
equilibrium
is
 maintained.
It
doesn’t
have
to
be
the
same
people.)
 Equilibrium
Constant
Expression:
aA
+bB
 
cC
+
dD
 
 [C]c
[D]d
 Keq
=
 ______________
 
 [A]a
[B]b
 Keq
>
1

Products
are
favored
(lies
to
the
right)
 Keq
<
1
Reactants
are
favored
(lies
to
the
left)
 Keq
=
1
Similar
amounts
of
reactants
and
products
(lies
in
middle)
 Homogeneous
equilibrium:

equilibrium
in
which
reactants
and
products
are
 in
the
same
physical
state
(solid,
gas,
liquid)
 Heterogeneous
equilibrium:

equilibrium
in
which
reactants
and
products
are
 in
different
physical
states.
 In
heterogeneous
equilibriums,
solids
and
liquids
are
not
included
in
the
 equilibrium
constant
equation.


 CO(g)
+
H20(l)

 
CO2(g)
+
H2(g)
 Keq
=
[H2][CO2]

/
[CO]
 H20
is
not
included
because
it
is
not
a
gas
or
aqueous
and
because
it
is
a
pure
 substance.
 Le
Chatelier’s
Principle
 Le
Chatelier’s
Principle:
if
a
system
at
equilibrium
is
disrupted,
it
shifts
to
 establish
a
new
equilibrium.
(Example:
Let’s
say
you
have
a
scale
where
one
of
 the
sides
is
heavier
than
the
other.
You
want
to
equal
out
the
sides
so
you
take
 some
weight
from
the
heavier
side
and
move
it
to
the
lighter
side.
That’s
what
 happens
chemically
in
equilibrium
except
you
don’t
have
to
do
anything.)
 
 
 A(g)
+
B(g)

C(g)
+
D(g)
 
 Equilibrium
Shifts
due
to
Concentration
Changes
 Add
reactant,
equilibrium
shifts
right
 Add
Product,
equilibrium
shifts
left
 Remove
Reactant,
equilibrium
shifts
left
 Remove
product,
equilibrium
shifts
right
 
 Equilibrium
Shifts
Due
to
Volume
Change
 Relative
Number
of
 Gaseous
Molecules
 in
Balanced
 Equation
 Reactants
<
 products
 Reactants
>
 products
 Reactants
=
 products
 Example
 Increase
Volume
 Decrease
Volume
 N204(g)
 
2NO2(g)
 N2(g)
+
3H2
 
 2NH3(g)
 2NO(g)
 
N2(g)
+
 O2(g)
 Shift
right
 Shift
left
 No
Shift
 Shift
Left
 Shift
right
 No
Shift
 
 Effects
of
Temperature
Changes
on
the
Position
of
Equilibrium
 Type
Of
Reaction
 Endothermic
 Reaction
 Exothermic
 Reaction
 Equation
 Heat
+
A
+
B

C
+
 D
 A
+
B

C
+
D
+
 Heat
 Increase
 Temperature
 Shift
right.

Keq
 increases.
 Shift
left.

Keq
 decreases.
 Decrease
 Temperature
 Shift
left.

Keq
 decreases.
 Shift
right.

Keq
 increases.
 Treat
heat
like
another
reactant
or
product.


Increasing
heat
on
one
side
will
shift
the
 equilibrium
to
the
other,
and
removing
it
on
one
side
will
shift
equilibrium
toward
that
 side.
 
 ...
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