Download Expt 2-2 Bomb Calorimetry

Last
Modified:
9/23/08
Physical
Chemistry
Laboratory
(CHEM
336)
EXPT
2­2:
Bomb
Calorimetry
Introduction
Enthalpies
of
combustion
can
be
determined
reliably
using
bomb
calorimetry.
This
technique
was
once
widely
used
in
physical
chemistry,
since
it
provides
thermodynamic
data
in
a
very
direct
way,
and
is
still
used
in
the
fuel
and
food
industries.
Although
in
the
latter
case
its
use
is
commonplace,
enthalpies
of
combustion
cannot
be
used
uncritically
as
a
measure
of
the
energy
that
will
be
realized
by
the
body
when
the
food
is
consumed.
Not
all
foods,
for
example,
will
be
converted
by
the
body
into
exactly
those
products
formed
by
combustion
in
a
calorimeter.
To
measure
enthalpies
of
combustion,
we
burn
a
known
amount
of
material
in
a
bomb
calorimeter
and
determine
the
temperature
change.
The
bomb
is
pressurized
with
oxygen
to
ensure
complete
combustion,
and
sealed
to
prevent
escape
of
the
combustion
products.
The
compound
is
ignited
by
passing
a
current
through
a
fuse
wire
within
the
bomb.
Allowance
for
the
heat
capacity
of
the
bomb
is
made
by
igniting
a
known
quantity
of
a
substance
of
known
enthalpy
of
combustion
as
a
standard.
Heat
loss
to
the
surroundings
can
be
calculated
by
use
of
a
cooling
correction
curve,
or,
as
in
this
experiment,
prevented
by
use
of
a
jacket
around
the
calorimeter,
maintained
at
the
same
temperature
as
the
calorimeter
itself;
the
reaction
is
then
adiabatic.
The
procedure
is
described
in
great
detail
in
Expt.
6
of
Garland
et
al.
Also
read
pp.
145‐152
of
Garland
et
al.
to
refresh
principles
of
calorimetry.
A
neat
description
of
the
experiment
with
animation
can
also
be
found
at
http://www.chm.davidson.edu/ronutt/che115/Bomb/Bomb.htm
Theory
The
generic
chemical
equation
for
combustion
of
organic
compounds
at
low
temperatures:
CkHlOm
+
(k­l/4­m/2)O2(g)
‐‐>
k
CO2(g)
+
l/2
H2O
(l)
Depending
on
the
experimental
conditions
for
combustion
experiment,
the
heat
released,
q,
is
equal
to
the
enthalpy
change,
ΔHco
(at
constant
pressure),
or
the
internal
energy
change,
ΔU
(at
constant
volume).
In
a
bomb
calorimeter
we
measure
qv
(i.e.
ΔU),
which
can
be
used
to
calculate
ΔHco:
ΔHco
=
ΔU
+ Δ(pV)
The
last
term
can
be
evaluated
using
the
perfect‐gas
law,
making
the
overall
equation:
1
Last
Modified:
9/23/08
ΔHco
=
ΔU
+
RTΔngas
where
Δngas
is
the
increase
of
the
number
of
moles
of
gas
in
the
system.
Since
the
bomb
calorimeter
is
isolated
from
the
rest
of
the
universe,
we
can
define
the
reactants
(sample
and
oxygen)
to
be
the
system
and
the
rest
of
the
calorimeter
(bomb
and
water)
to
be
the
surroundings.
Then
the
change
in
the
internal
energy
of
the
reactants
upon
combustion,
ΔE,
can
be
calculated
from
ΔU
=
‐ΔUsurr
=
‐CcalΔT
where
Ccal
is
the
heat
capacity
of
surroundings
(i.e.
the
water
and
the
bomb)
and
ΔT
is
its
temperature
change.
We
assume
that
Ccal
is
nearly
independent
of
T
over
the
small
temperature
range.
Thus,
the
energy
of
combustion
is
relatively
easy
to
measure
by
monitoring
temperature
increase
in
a
course
of
reaction
in
a
calorimeter
with
known
heat
capacity,
Ccal.
The
latter
can
be
calibrated
using
another
reaction
with
known
heat
of
reaction
conducted
under
the
same
conditions.
In
a
bomb
calorimeter,
benzoic
acid
combustion
with
ΔUBA
=
‐
26.41
kJ/g
is
often
used
for
such
purpose.
Instructions
for
operating
Parr
Bomb
Calorimeter
1.
Cut
a
piece
of
iron
wire
approximately
3
inches
long
and
weigh
it.
Prepare
the
sample
by
weighing
approximately
0.5
‐
1.0
g
of
an
organic
compound
and
compressing
it
into
a
pellet;
weigh
the
pellet
accurately
after
it
is
pressed.
Do
not
exceed
3
metric
tons
of
pressure
when
pressing
the
pellet.
2.
Install
the
sample
into
the
ignition
pan;
make
sure
that
the
wire
does
not
touch
the
pan,
or
else
it
will
short
circuit
and
burn
through
rather
than
ignite
your
sample.
Approximately
10
cm
of
wire
is
required.
Weigh
the
wire
prior
to
assembly.
3.
Carefully
assemble
the
bomb
and
screw
down
the
cap
hand‐tight.
4.
Connect
the
bomb
to
the
oxygen
tank
and
carefully
fill
the
bomb
to
approximately
25
atm
(380
psi);
release
the
pressure
to
flush
and
refill
it
again;
check
the
bomb
for
leaks
by
immersing
it
in
water.
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Last
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5.
Dry
the
bomb
and
set
it
in
the
dry
pail;
then
set
the
pail
into
the
calorimeter.
6.
Fill
a
2‐L
volumetric
flask
with
water
at
20o
C
and
pour
it
carefully
into
the
pail;
avoid
splashing.
7.
Put
the
calorimeter
lid
in
place
and
clamp
the
thermometer
in
place
as
low
as
it
will
go.
8.
Plug
the
stirrer
and
ignition
cords
into
the
control
box;
turn
on
the
stirrer
and
make
sure
it
runs
smoothly
During
the
experiment
Begin
time‐temperature
readings,
reading
the
precision
thermometer
once
every
30
s
and
recording
both
the
time
and
the
temperature;
temperature
scale
should
be
read
to
the
nearest
0.01
°F.
After
recording
the
temperature
variation
for
approximately
5
min,
the
bomb
may
be
ignited
continue
T
vs
time
reading
for
about
15‐20
min.
After
completion,
release
the
bomb
pressure,
and
open
the
bomb.
Remove
and
weigh
any
unburned
iron
wire;
ignore
"globules"
unless
they
are
fused
metal
(oxides
will
crush)
if
the
inside
of
the
bomb
is
coated
with
soot,
the
combustion
was
incomplete
and
the
experiment
has
to
be
discarded.
Calculate
the
net
weight
of
iron
burned
wipe
dry
all
bomb
parts.
Plan
to
run
benzoic
acid
(two
times)
and
anthracene
(two
times).
If
you
have
time,
determine
the
enthalpy
of
combustion
of
Cheerios.
Calculations
The
heat
capacity
of
the
calorimeter,
Ccal,
should
be
found
by
determining
the
temperature
rise
ΔT
in
calibration
runs
with
benzoic
acid;
the
specific
energies
of
combustion
of
benzoic
acid
(BA)
and
iron
(Fe)
are
ΔUBA
=
‐
26.41
kJ/g
and
ΔUFe
=
‐
6.68
kJ/g,
respectively.
Write
the
balanced
equations
corresponding
to
ΔHco
for
all
substances
you
studied.
Using
the
calibration
value
of
Ccal,
calculate
the
energies
of
combustion
for
other
compounds
and
estimate
their
accuracy
from
your
experimental
conditions.
Calculate
ΔHco
for
the
compounds
you
studied
including
their
accuracies.
How
much
different
are
your
ΔHco
data
from
the
literarature
values?
Also
find
literature
values
of
ΔHco
for
naphthalene
and
phenanthrene.
Comment
on
the
differences
between
ΔHco
in
a
series
naphthalene,
anthracene
and
phenanthrene.
References
Garland
et
al.
Experiments
in
Physical
Chemistry
7th
Ed,
pp
145‐158.
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