Kinetics+of+Cyclohexanone

Kinetics+of+Cyclohexanone - Chem
1310
HONORS


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Unformatted text preview: Chem
1310
HONORS
 A
Kinetics
Study
of
the
Iodination
of
Cyclohexanone
 Introduction

 Thermodynamics
 outlines
 the
 fundamentals
 of
 what
 drives
 chemical
 reactions
 to
 equilibrium
through
the
Gibb's
free
energy,
ΔG.
 ΔG
=
ΔH
‐T
ΔS
 (1)
 The
 contribution
 and
 interplay
 of
 enthalpy,
 ΔH,
 and
 entropy,
 ΔS,
 
 to
 the
 free
 energy
 are
 given
 in
 equation
 1.
 Some
 reactions
 readily
 occur
 with
 very
 negative
 Gibb's
 free
 energy
 while
 others
 are
 only
 spontaneous
 at
 higher
 or
 even
 lower
 temperatures.
 Thermodynamics
 dictates
 which
 factors
 are
driving
reactions,
but
enthalpy
and
entropy
are
state
functions
and
do
not
give
us
information
 about
the
chemical
process
while
it
is
occurring.

The
field
of
chemical
kinetics
 focuses
on
 the
time
 evolution
 of
 the
 process.
 
 The
 field
 of
 kinetics
 studies
 the
 rates
 of
 reactions
 and
 the
 pathways
 observed
for
given
reactions.
Many
chemical
reactions
are
complex
reactions
and
occur
through
a
 series
 of
 steps
 known
 as
 elementary
 reactions.
 These
 series
 of
 reactions
 compose
 an
 overall
 reaction
 mechanism.
 The
 ultimate
 goal
 of
 kinetics
 is
 to
 elucidate
 these
 mechanisms
 in
 order
 to
 obtain
 a
 better
 understanding
 of
 thermodynamic
 barriers
 that
 may
 exist.
 Such
 an
 understanding
 allows
for
control
of
chemical
reactions
and
has
led
to
the
production
of
catalysts
to
make
reactions
 occur
more
rapidly
or
spontaneous
under
feasible
conditions.

 Chemical
 kinetics
 requires
 experimentation
 to
 elucidate
 mechanisms
 and
 to
 completely
 understand
 the
 nature
 of
 a
 reaction.
 The
 rate
 of
 a
 chemical
 reaction
 is
 defined
 as
 the
 change
 in
 concentration
 of
 a
 reactant
 or
 product
 as
 a
 function
 of
 time.
 The
 rate
 of
 disappearance
 of
 any
 reactant
can
be
defined
as
well
as
the
rate
of
 appearance
of
any
product.
Rates
have
units
of
mol.L‐ 1.s‐1
 or
 M/s
 and
 are
 time
 dependent.
 When
 reactions
 first
 begin,
 a
 larger
 rate
 is
 observed
 and
 the
 rate
slows
as
equilibrium
is
reached.
This
is
because
the
reactant
concentrations
are
changing
with
 time,
i.e.
decreasing.
 A
rate
can
be
specifically
defined
by
the
following
expression:
 where
A
is
the
reactant
and
t
=
time

 (2)
 
 A
negative
sign
is
used
by
convention
because
the
reactant
is
disappearing.

The
rate
of
a
chemical
 reaction
is
proportional
to
the
reactants
and
a
rate
constant.
Given
the
reaction:
A
+
B
 C,
the
 following
 rate
 law
 can
 be
 written,
 where
 m
 and
 n
 describe
 the
 rate
 dependence
 on
 the
 concentrations
of
the
reactants
A
and
B:
 
 (3)
 In
 the
 expression,
 k
 is
 the
 rate
 constant
 (which
 is
 temperature
 dependent)
 and
 m
 and
 n
 are
 the
 reaction
orders
for
A
and
B,
respectively
which
denote
how
the
rate
changes
with
any
change
in
the
 concentrations
of
these
reagents.
In
most
chemical
reactions
(except
for
elementary
processes),
the
 values
of
m
and
n
are
 not
related
to
the
stoichiometric
coefficients
in
the
balanced
equation‐‐these
 order
numbers
must
be
determined
experimentally.
Furthermore,
m
and
n
are
not
always
integer
 values
(i.e.,
1/2,
3/4,
1/6,
etc.
are
possible
reaction
order
values).
 
 
 In
this
experiment,
the
following
reaction
will
be
studied:
 
 


(4)
 Cyclohexanone
 





Triiodide
 
 
 
 
 
 2‐iodocyclohexanone
 
 The
goal
of
the
experiment
will
be
to
determine
the
rate
law
which
requires
elucidation
of
 the
 concentration
 dependence
 of
 cyclohexanone,
 triiodide,
 and
 acid.
 This
 will
 be
 important
 in
 determining
the
mechanism
by
which
this
reaction
occurs.
The
overall
rate
law
will
have
the
form:
 


(5)
 Determining
 the
 rate
 law
 will
 also
 allow
 us
 to
 identify
 the
 rate
 determining
 step
 in
 the
 proposed
 mechanism.
Recall
that
the
rate
determining
step
is
the
slow
step
in
the
chemical
reaction,
and
the
 kinetics
of
the
reaction
depends
largely
on
how
quickly
this
step
occurs.

 
 Aqueous
triiodide
is
a
brown
liquid,
and
we
can
use
the
disappearance
of
color
to
assess
the
 kinetics
 of
the
reaction.
We
will
start
timing
the
reaction
as
soon
as
the
components
are
mixed,
and
 when
the
brown
color
has
completely
diminished,
we
will
stop
timing
and
make
a
note
of
the
time
 lapsed
 between
 the
 start
 of
 the
 discoloration
 and
 the
 completion.
 The
 disappearance
 of
 triiodide
 occurs
 early
 enough
 in
 the
 reaction
 to
 be
 used
 as
 an
 indicator
 of
 the
 initial
 rate
 of
 the
 overall
 reaction.
 
 The
 second
 part
 of
 the
 experiment
 will
 involve
 temperature
 variations
 to
 determine
 the
 activation
 energy
 associated
 with
 the
 rate
 determining
 step.
 The
 rate
 constant,
 k,
 is
 temperature
 dependant
 and
 encompasses
 the
 activation
 energy,
 Ea,
 for
 the
 system.
 The
 activation
 energy
 represents
the
minimum
amount
of
energy
that
must
be
accrued
by
the
system
for
the
reaction
to
 occur.
 The
 Arrhenius
 equation
 defines
 the
 relationship
 between
 the
 rate
 constant,
 activation
 energy,
and
temperature.
Recall
that
temperature
will
be
in
units
of
Kelvin
because
the
activation
 energy
has
units
of
joules.
 

(5)
 In
 the
 Arrhenius
 expression,
 A
 is
 called
 a
 pre‐exponential
 factor
 and
 includes
 factors
 such
 as
 the
 number
 of
 collisions,
 R
 is
 the
 energy
 gas
 constant
 (8.314
 J/molK),
 and
 Ea
 is
 the
 activation
 energy
 per
mole
(J/mol).

 
 
 Procedure


 Chemicals
and
Supplies:
 Stopwatchs
 1.0
M
HCl
 0.5
M
Cyclohexanone
solution
 0.02
M
I3‐
 pH
paper
 
 16x125mm
test
tubes
(to
hold
reagents)
 13x100mm
test
tubes
(to
run
trials
in)
 5.00
mL
serological
pipets
 #2
corks
 water
baths
 
 Part
A:
Determination
of
the
Rate
Law.

 When
completing
the
experiment,
note
the
concentrations
used
during
each
step.
 Using
pH
paper,
measure
the
pH
before
and
after
the
reaction.
 Record
the
temperature
in
the
room.
 A­1
 Obtain
~20‐mL
of
1.0
M
HCl,
~20‐mL
of
0.5
M
cyclohexanone,
~20‐mL
of
0.02
M
triiodide
and
 ~20‐mL
of
D.I.
water
into
four
medium,
16x125mm
test
tubes.

Obtain
4‐5.00mL
serological
pipets,
 one
for
each
tube
of
reagent.
All
reagents
can
be
put
into
the
13x100
mm
test
tube
except
the
I3‐.
 Add
the
I3‐
while
starting
the
stopwatch.

Cork
and
invert3‐4
to
mix.
 Make
a
"blank"
test
tube
with
6
mL
of
D.I.
water
and
peer
down
both
test
tubes
at
a
sheet
of
white
 paper.
When
the
reaction
tube
matches
the
blank,
record
the
time.

The
"time‐to‐colorless”
can
be
 used
as
the
inverse
of
the
initial
rate
of
the
reaction.

 Run
trial
1
 twice
to
test
the
reaction
conditions
and
you’re
reproducibility
to
determine
when
the
 triiodide
has
disappeared.
 A­2:
 Run
the
kinetics
trials
from
the
table
below.

 1.0
M
HCl
 (mL)
 1
 2
 0.5M
 Cyclohexanone
 (mL)
 2
 1
 1
 6
 0.02M
I3­
 (mL)
 D.I.
water
 (mL)
 Total
 Volume
(mL)
 Trial
 2
 3
 4
 
 1
 2
 2
 2
 1
 2
 1
 1
 0.5
 2
 2
 1.5
 6
 6
 6
 Dispose
of
all
trials
in
the
waste
container
for
halogenated
organics
 
 Part
B:
Determination
of
the
Activation
Energy
 B­1:
 Three
water
baths
will
be
in
the
lab,
two
warm
and
one
cold.
Select
a
kinetics
trial
and
heat
 up
all
reagents
except
I3‐
in
one
 test
and
 the
triiodide
 in
 another
test
 tube
in
 the
 water
 bath
 until
 equilibrated.
 Record
 the
 temperature
 of
 the
 water
 bath.
 Cyclohexanone
 is
 volatile,
 do
 not
 heat
 above
40
‐
45oC.
 B­2:
 After
 heating,
 with
 care
 simultaneously
 mix
 the
 solutions
 and
 record
 the
 time
 for
 the
 discoloration
of
triiodide.

Repeat
for
the
various
temperatures.
 B­5:
 Using
 the
 room
 temperature
 data
 from
 part
 A
 and
 data
 from
 part
 B.
 make
 a
 plot
 of
 ln(k)
 versus
1/T
(K‐1)
and
determine
the
magnitude
of
the
activation
energy.
 
 Dispose
 of
 all
 trials
 in
 the
 waste
 container
 for
 halogenated
 organics
 as
 well
 as
 any
 leftover
 cyclohexanone
and
triiodide.
 
 
 Post
­
Lab
Report
 Your
post
lab
report
consists
of
a
Data/Results
section
and
a
Discussion/Conclusion
section.
A
word
 processing
program
such
as
WORD
should
be
used.
Data
and
results
should
be
put
into
tables
and

 plots
should
be
made
using
Excel.
 Part
I:

Calculations
 1.
 Compute
 the
 rate
 law
 for
 the
 iodination
 of
 cyclohexanone.
 Show
 all
 mathematical
 calculations
with
the
final
answer
being
clearly
displayed.
 2.
 Make
 your
 plot
 in
 Excel
 for
 the
 reaction
 run
 at
 various
 temperatures.
 Make
 the
 plot
 show
 the
data
used
for
plotting
and
print
out.
Compute
the
activation
energy
for
the
chemical
reaction.
 Part
II:

Discussion
and
Questions
 1.
 2.
 3.
 What
 reactants
 (cyclohexanone,
 triiodide,
 and
 acid)
 contributed
 to
 the
 rate
 determining
 step?
Explain.
 Was
 acid
 a
 catalyst,
 intermediate,
 or
 neither?
 Explain
 your
 reasoning
 using
 the
 data
 collected.
 Given
 the
 following
 proposed
 mechanism
 for
 the
 iodination,
 answer
 the
 associated
 questions.
 

















 
 
 
 
 
 a.

Identify
all
intermediates,
if
any.
 b.

Identify
the
catalyst,
if
any.
 c.

Draw
an
energy
diagram
to
scale
(based
upon
the
calculation
of
the
activation


 




energy)
representing
this
mechanism.
 Data
Tables
 
 Part
A:
Rate
Law
Determination
 Temperature
 ___________________
 Trial
 
 [Cyclohexanone]
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 Observations/Comments:
 
 
 Part
B:
Activation
Energy
Determination
 
 Temperature
 
 
 
 
 
 
 
 
 [Cyclohexanone]
 [HCl]
 
 
 
 
 [Triiodide]
 
 
 
 
 Time
to
Colorless
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 [HCl]
 
 [Triiodide]
 
 Time
to
 Colorless
 
 pH
Observation
 
 
 Observations/Comments:
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 Pre
‐
Lab
 1.
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 2.
 
 
 
 
 
 3.
 
 
 
 
 
 (b.)
 (a.)
 
 
 
 
 
 
 Name________________________
 Given
the
reaction
of
cyanide
with
methyl
iodide,
the
following
was
observed:
 
 
 
 
 time
 42
s
 84
s
 84
s
 
 
 
 
 [methyl
iodide]
 0.02
M
 
 0.01
M
 
 0.02
M
 
 
 
 
 
 0.02
M
 0.02
M
 0.01
M
 [CN‐]
 What
is
the
reaction
order
with
respect
to
methyl
iodide?
(2
points)
 What
is
the
reaction
order
with
respect
to
cyanide?
(2
points)
 What
is
the
definition
of
rate
determining
step?
(2
point)
 What
is
the
difference
between
an
intermediate
and
a
catalyst?
(4
points)
 ...
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This note was uploaded on 01/26/2011 for the course CHEM 1310 taught by Professor Cox during the Spring '08 term at Georgia Tech.

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