Properties of ammonium nitrate


Shah, K.D.; Roberts, A.G.

Nitric Acid and Fertilizer Nitrates: 171-196

1985


Physical properties, thermodynamic data, and chemical properties of ammonium nitrate are given.

11
Properties
of
Ammonium
Nitrate
K.
D.
SHAH
and
A.
G.
ROBERTS
Imperial
Chemical
Industries
PLC,
Billingham,
Cleveland,
England
I.
Physical
Properties
172
A.
Crystalline
forms
172
B.
Solubility
and
hygroscopicity
173
C.
Density
175
D.
Boiling
point
175
E.
pH
of
aqueous
solutions
175
F.
Vapor
pressure
175
G.
Viscosity
and
surface
tension
179
II.
Thermodynamic
Data
180
A.
Reaction
heats
180
B.
Specific
heat
capacities
182
C.
Heats
of
phase
changes
182
D.
Heats
of
solution
and
dilution
182
III.
Chemical
Properties
184
A.
Decomposition
of
pure
ammonium
nitrate
184
B.
Steady-state
reaction
temperature
184
C.
Decomposition
rates
of
pure
ammonium
nitrate
187
D.
Catalytic
effects
on
decomposition
187
E.
Inhibiting
effect
of
urea
191
F.
Miscellaneous
chemical
reactions
192
G.
Self-heating
in
fertilizers
192
H.
Oxidizing
properties
192
References
193
171
172
/
Shah
and
Roberts
I.
PHYSICAL
PROPERTIES
Pure
ammonium
nitrate
is
a
white,
odorless
salt
with
a
melting
point
of
169.6°C,
of
molecular
formula
NH
4
NO
3
,
and
with
molecular
weight
80.
A.
Crystalline
Forms
Solid
ammonium
nitrate
occurs
in
five
different
stable
crystalline
forms
or
phases
(1).
A
sixth
phase
is
reported
to
exist
at
elevated
tempera-
tures
and
pressures
(2).
The
main
transitions,
summarized
in
Table
1,
are
accompanied
by
volume
changes
as
shown
in
Figure
1
(3).
Transi-
tions
V
IV
and
III
II
are
accompanied
by
3.0
and
1.3%
shrinkage,
respectively,
whereas
IV
÷
III
and
II
I
give
3.6
and
2.1%
expansion,
respectively.
A
number
of
unusual
features
have
been
discovered
in
these
transi-
tions.
Bridgman
(2)
found
that
phase
III
ceases
to
exist
at
pressures
above
860
kg/cm
2
,
and
as
the
pressure
is
increased
further
phase
II
disappears
at
9000
kg/cm
2
and
is
replaced
by
a
new
phase
VI.
Behn
(4)
discovered
that
in
exhaustively
dried
samples
of
ammonium
nitrate
the
32°C
transition
disappears
and
is
replaced
by
a
transition
at
50°C
eliminating
the
existence
of
phase
III.
This
transition
is
generally
be-
lieved
to
be
a
metastable
one.
Phases
III
and
IV
are
most
important
commercially
because,
firstly,
the
transition
between
them
occurs
at
a
temperature
close
to
ambient
temperatures
and,
secondly,
this
transi-
tion
is
accompanied
by
a
large
volume
change.
Repeated
phase
changes
can
easily
cause
physical
breakdown
of
product.
Therefore
to
obtain
a
stable
product
ammonium
nitrate
has
to
be
stabilized
against
this
phase
change.
A
number
of
methods
are
available
to
achieve
this
(see
Chap.
13).
Table
1
Crystalline
Forms
of
Ammonium
Nitrate
Temperature
Form
Crystal
system
range
(°C)
Liquid
Above
169.6
Phase
I,
a
Cubic
125.2-169.6
Phase
II,
d
Tetragonal
84.2-125.2
Phase
III,
y
Rhombic
32.3-84.2
Phase
IV,
13
Rhombic
-18—
32.3
Phase
V,
a
Tetragonal
Below
-18
Properties
of
Ammonium
Nitrate
/
173
Phase
Rhombic
1
*..
is
IV
Position
I
change
when
1
Phase
III
••••
of
phase
is
absent
Phase
Phase
Tetragonal
V
Tetragonal
II
Phase
Rhombic
III
Phase
cubic
I
liquid
4%14
\
1%14
-
50
0
50
100
150
200
Temperature,
°
C
Fig.
1
Density
changes
in
phase
transitions
of
ammonium
nitrate.
(Data
from
Ref.
3.)
B.
Solubility
and
Hygroscopicity
Ammonium
nitrate
is
extremely
soluble
in
water
and
the
solubility
in-
creases
rapidly
with
temperature,
as
shown
in
Fig
2
(5-7).
In
strong
solutions,
particularly
in
the
concentration
range
90-100%
ammonium
nitrate,
even
a
small
change
in
the
moisture
content
produces
an
ap-
preciable
change
in
the
crystallizing
temperature.
In
industry,
use
is
commonly
made
of
this
property
to
determine
the
moisture
content
in
concentrated
ammonium
nitrate
solutions.
Ammonium
nitrate
shows
considerable
solubility
in
many
organic
and
other
nonaqueous
solvents.
Of
particular
interest
to
the
fertilizer
in-
dustry
is
the
solubility
in
anhydrous
and
aqueous
ammonia
solutions
and
aqueous
nitric
acid
solutions.
These
systems
have
been
studied
(7-9).
1750
1700
1650
1550
1500
1450
1400
174
/
Shah
and
Roberts
,
a
a
a
WATER
VAPOR
ct
0
N
1\1.
ti
....
BOILING
41.
2
i:
SOLUTION
4
7-
P
v
C.)
Solution
,c-'
-'
0
+
solid
0\'
J
0Ammonium
'
Nitrate
Solution
+
ice
5
Ice
+
solid
Ammonium
Nitrate
20
40
60
80
100
Concentration
of
Ammonium
Nitrate,
wt
%
Fig.
2
Crystallization
and
boiling
point
curves
of
aqueous
ammonium
nitrate
solutions.
(Data
from
Refs.
5
-
7.)
Ammonium
nitrate
is
very
hygroscopic
but
does
not,
however,
form
any
hydrates.
The
relative
humidity
of
air
at
various
temperatures
when
in
equilibrium
with
a
saturated
solution
of
ammonium
nitrate
is
given
in
Table
2
(10).
Table
2
Hygroscopicity
of
Ammonium
Nitrate
Temperature
(°C)
5
10
15
20
25
30
35
40
45
50
Relative
humidity
(%)
82
75
70
67
63
59
56
53
51
48
Source:
Ref.
10.
240
220
200
180
160
140
°
120
1
1
1
1,1
100
a
I-
60
40
20
0
20
-
40
0
Properties
of
Ammonium
Nitrate
/
175
C.
Density
Solid
ammonium
nitrate
has
a
density
of
1.725
g/cm
3
at
room
tempera-
ture
and
it
is
affected
by
temperature
and
phase
transitions
as
shown
in
Figure
1.
Densities
of
the
aqueous
solutions
at
saturation
and
boiling
points
are
given
in
Table
3
(7)
and
more
extensive
data
are
given
in
Figure
3.
Densities
of
ammonium
nitrate
solutions
in
liquid
ammonia
have
also
been
measured
(11).
D.
Boiling
Point
The
boiling
point
of
pure
ammonium
nitrate
cannot
be
determined
di-
rectly
because
it
undergoes
decomposition
before
reaching
the
boiling
condition.
A
figure
of
230°C
has
been
estimated
by
extrapolation
of
the
boiling
points
of
aqueous
solutions
(see
Table
3).
The
atmospheric
boiling
point
curve
of
the
aqueous
solution
is
shown
in
Figure
2.
E.
pH
of
Aqueous
Solutions
The
pH
of
ammonium
nitrate
solution
is
an
important
control
parameter
in
the
manufacturing
process.
It
is
usually
expressed
on
the
basis
of
a
10%
(by
weight)
solution
of
ammonium
nitrate.
Neutral
solution
has
a
pH
in
the
range
4-5
and
about
the
neutral
point
a
small
addi-
tion
of
ammonia
or
nitric
acid
changes
the
pH
appreciably.
Figure
4
illustrates
this
effect
for
10%
solution
at
25°C
and
87%
solution
at
134°C.
F.
Vapor
Pressure
1.
Solid
Ammonium
Nitrate
Pure
ammonium
nitrate
exerts
vapor
pressure
due
to
reversible
ther-
mal
dissociation
reaction.
Only
at
higher
temperatures
do
the
decom-
position
reactions
become
important
(see
Sec.
III.A).
The
dissociation
pressure
can
be
estimated
from
the
following
equa-
tions
(12)
based
on
experimentally
determined
values
(12,13):
Solid:
log
p
=
10.708
4670
Liquid:
log
p
=
9.981
4360
T
where
p
is
the
total
pressure
of
NH
3
and
HNO
3
vapors
(mmHg)
and
T
is
absolute
temperature
(K).
Table
4
gives
the
dissociation
pressure
at
different
temperatures.
The
above
data
are
for
pure
ammonium
nitrate.
The
pressure
is
affected
by
any
excess
NH
3
or
HNO
3
present.
01
81.
1
aCI
OH
p
U
D
11
Dt
S
Table
3
Densities
of
Aqueous
Ammonium
Nitrate
Solutions
Saturated
solution
At
boiling
point
Temperature
(°C)
Concentration
(wt
NH
4
NO
3
)
Density
Temperature
(g
/m1)
(°C)
Concentration
(wt
NH
4
NO
3
)
Density
(g
/m1)
20
66.1
1.3115
110
51.9
1.177
40
73.3
1.3415
120
71.3
1.276
60
80.2
1.3519
130
81.4
1.328
80
85.9
1.3940
140
87.0
1.358
100
91.0
1.4145
150
90.75
1.375
120
94.7
1.4260
160
93.4
1.390
140
97.4
1.4320
180
95.9
1.398
160a
99.4a
1.4360a
200
97.55
1.401
170
a
100.0a
1.4Ma
220a
98.9a
1.403a
230
a
100
a
1.405a
a
Extrapolated
data.
250
200
0
0
ai
z
150
io
a)
0
100
ci)
Properties
of
Ammonium
Nitrate
/
177
100%
99%
97%
94%
90%
EkNi
e
IN
BOILING
1
3
,
to
ct
-
-
0
L
,0
1
%
0
0
2
%
30%
40%
50%
60%
70%
80%
.S
.
SY
49
\ \
S`
1400
Fig.
3
Densities
of
aqueous
solutions
of
ammonium
nitrate.
(Data
from
Ref.
7.)
Loiz
Linear
Log
0
1000
1100
1200
1300
Density,
kg/m
3
8
7
6
5
pH
8
7
S
pH
4
4
-3
3-
2
4
62
$
c
2
0
5
4
3
2
1
7
-5
.3
-2
1.0705
-0302
01
+
0
-01
.0203-0507
1
.2
3
-5
7
Acidity
moles
excess
HNO
3
/100
moles
NI
-
14NO3
ill
87%
134
t
2
3
4
5
0
Fig.
4
The
pH
values
of
aqueous
ammonium
nitrate
solutions.
(Cour-
tesy
of
Imperial
Chemical
Industries
PLC.)
1000
co
_0
ED:
100
CO
1
178
/
Shah
and
Roberts
Table
4
Ammonium
Nitrate
Dissociation
Pressure
Temperature
(°C)
120
140
160
180
200
220 240
Pressure
(mmHg)
0.07
0.25
0.84
2.27
5.8
13.7
30.3
(mbar)
0.093
0.333
1.120
3.026
7.731
18.262
40.390
2.
Aqueous
Solutions
Values
of
water
vapor
pressure
over
aqueous
solutions
of
ammonium
nitrate
are
summarized
in
Figure
5
(7,14,15).
The
vapor
pressure
over
saturated
solution
given
by
the
curve
ABC
goes
through
a
maxi-
mum
at
approximately
260
mbar
and
125°C.
Concentration
wt
%
1 f i l
%
/ /
/
///
I
/
A
,
---
Extrapolated
/
ii
i
7
—1
m
bar
=
100
N/m
2
=
0.0145
1
bike
.
4
/
. i
-
/ i
I
/
/
1
1
/
I
ii
°9-
1".
S .
k
--•
NOTE:Above
about
210°C
,...P
formed
as
a
product
water
is
99.9
0
1
of
ammonium
decomposition.
nitrate
___
___
___
A
0
20
Temperature,°C
Fig.
5
Vapor
pressures
over
aqueous
ammonium
nitrate
solutions.
(Data
from
Refs.
7,
14,
and
15.)
104.1
101.1
dou
3
x10
-3
E
Z
2
x10
a
0
p
5
Properties
of
Ammonium
Nitrate
/
179
G.
Viscosity
and
Surface
Tension
Viscosities
of
aqueous
ammonium
nitrate
solutions
at
various
tempera-
tures
are
given
in
Figure
6
(7,11,16).
Surface
tension
of
the
fused
ammonium
nitrate
at
different
tempera-
tures
and
of
the
aqueous
solution
at
100°C
and
different
concentrations
are
summarized
in
Tables
5
and
6,
respectively
(7).
1
Ns/m
2
=
1
kg/ms
=
10
3
cP
160
°C
So
ubility
limit
A
180
°C
200
°C
10
°C
0
°C
140
°C
10
°C
20
°C
120
°C
30
°C
40
°C
Atmospheric
pressure
boiling
limit.
60
70
80
100
°C
0
10
20
30
40
50
60
70
80
90
100
Concentration
wt
%
NH
4
NO
3
Fig.
6
Viscosities
of
aqueous
ammonium
nitrate
solutions.
(Data
from
Refs.
7,
11,
and
16.)
180
/
Shah
and
Roberts
Table
5
Surface
Tension
of
Fused
NH
4
NO
3
Temperature
(
°
C)
Surface
tension
(dynes/cm)
187.0
98.4
179.0
98.8
168.5
99.5
100.0
103.8
a
a
Extrapolated
data.
Table
6
Surface
Tension
of
Aqueous
NH4NO
3
Solutions
at
100°C
Surface
Surface
Concentration
tension
Concentration
tension
(wt
$
NH
4
NO
3
)
(dynes
/cm)
(wt
$
NH
4
NO
3
)
(dynes
/cm)
0
58.7
40
65.3
5
59.2
50
67.5
10
60.1
54
68.5
20
61.6
88
85.5
30
63.3
100
103.8
a
a
Extrapolated
from
surface
tension
of
fused
salt.
II.
THERMODYNAMIC
DATA
The
standard
free
energy
of
formation
of
ammonium
nitrate
at
25°C
is
-43.82
kcal/mol
and
its
entropy
at
25°C
is
36.0
kcal/mol
(13).
A.
Reaction
Heats
1.
Heat
of
Formation
(17)
2H
2
(g)
+
2
0
2
(g)
+
N
2
(g)
+
NH
4
NO
3
(s)
(Ali
=
-87.4
kcal/mol
at
25°C)
2.
Heat
of
Dissociation
NH
4
NO
3
(s/l)
o
NH
3
(g)
+
HNO
3
(g)
(1)
The
reaction
is
reversible
and
endothermic.
For
solid
NH4NO3,
OH
=
44.5
kcal/mol
at
25°C
(calculated
from
heats
of
formation)
(17).
For
liquid
NH
4
NO
3
see
Table
7.
Properties
of
Ammonium
Nitrate
/
181
Table
7
Heats
of
Dissociation
and
Decomposition
of
Liquid
NH4NO3
Temperature
(°C)
Heat
of
dissociation
reaction
(1)
tH
(kcal/mol)
a
Heat
of
decomposition
reaction
(2)
LH
(kcal/mol)
b
169.6
39.33
-13.21
200
38.92
-13.53
250
38.25
-14.03
300
37.65c
-14.58
350
37.12
c
a
From
Ref.
13.
bFrom
Ref.
18.
cBased
on
extrapolated
values
of
enthalpy
of
HNO
3
(gas).
Table
8
Specific
Heat
Capacities
Form
Specific
heat
capacity
(kcal/kg
°C)
Temperature
(°C)
NH
4
NO
3
Phase
V
0.31
-20
Phase
IV
0.40
7
Phase
III
0.37
57
Phase
II
0.34
105
Phase
I
0.34
147
Liquid
0.40
170
Aqueous
solution
(wt
%)
90%
NH
4
NO
3
0.45
100
70%
NH
4
NO
3
0.57
100
50%
NH4NO
3
0.69
100
30%
NH
4
NO
3
0.82
100
10%
NH
4
NO
3
0.94
100
Temperature
(°C)
Heat
change
(kcal/kg)
Transition
Reference
182
/
Shah
and
Roberts
3.
Heats
of
Decomposition
Reactions
Two
of
the
principal
reactions
are
given
below;
they
are
irreversible
and
exothermic:
NH
4
NO
3
(s/1)
+
N
2
0(g)
+
2H
2
0(g)
(2)
For
solid
NH
4
NO3,
A
H
=
-8.8
kcal/mol
at
25°C
(calculated
from
heats
of
formation)
(17).
For
liquid
NH
4
NO3,
see
Table
7
(18).
NH
4
NO
3
(s)
+
N
2
+
0.50
2
+
2H
2
0
(g)
Here
A
H
=
-28.2
kcal/mol
at
25°C
(calculated
from
heats
of
formation)
(17).
B.
Specific
Heat
Capacities
Data
on
specific
heat
capacities
from
different
sources
(13.15)
do
not
show
good
agreement.
Values
derived
from
enthalpy
data
(15)
are
summarized
in
Table
8.
C.
Heats
of
Phase
Changes
Heats
of
main
phase
transitions
are
summarized
in
Table
9
(15,19).
D.
Heats
of
Solution
and
Dilution
Heat
is
absorbed
when
ammonium
nitrate
dissolves
in
water.
The
heat
of
solution
when
dissolving
it
in
a
practically
infinitely
large
quantity
of
water
has
been
found
to
be
79
kcal/kg
at
18°C
(7).
Table
9
Phase
Transitions
Heat
Data
Phase
V—IV
-18
1.5,
absorbed
15
Phase
IV—III
32
5.0,
absorbed
Phase
III—II
84
3.9,
absorbed
Phase
II—I
125
12.8,
absorbed
Phase
I—liquid
170
18.3,
latent
heat
15
(fusion
or
melt-
of
fusion,
ab-
in
g)
sorbed
Solid—vapor
25
522,
latent
heat
19
(sublimation)
of
sublimation
absorbed
1
17
1
140
0
1
z
1
En
t
ha
lpy
o
f
so
lu
t
ion
kca
l
110
100
90
80
70
60
50
40
30
20
10
Properties
of
Ammonium
Nitrate
/
183
'
4
4
8
7
s.,
7,9
0
.
Isotherms
6
•••••
C
...
.
-..
SO
.
4tiN
ii‘
"...i
73
0
120
I S 4
\ \
7
60
cro146
=
At
L.
1
%
Ti
..--
-..-
0
600
..
—_
..
i
...-
#
1
110
\
00
Ailt
.,R\
'.-
A
.--
,
.....-
A.0
RN
- 70
0
0.1
02
0.3
04
05
06
07
08
09
1.0
Wt
fraction
NH
4
NQ
3
Fig.
7
Enthalpy—concentration
diagram
for
ammonium
nitrate—water
system.
170°C
125°C
83°C
32°C
184
/
Shah
and
Roberts
Enthalpy
concentration
data
for
the
ammonium
nitrate—water
system
are
summarized
in
Figure
7
(15).
III.
CHEMICAL
PROPERTIES
Ammonium
nitrate
is
thermally
quite
stable
at
ambient
temperatures.
It
begins
to
decompose
appreciably
at
approximately
200°C.
The
decomposition
reactions
are
complex
and
a
number
of
chemical
substances
affect
the
rates
of
decomposition.
Ammonium
nitrate
is
a
powerful
oxidizing
agent
and
therefore
a
strong
supporter
of
combustion.
Ammonium
nitrate
itself,
however,
does
not
burn.
These
properties
are
discussed
in
detail
in
the
fol-
lowing
pages.
The
explosive
properties
are
discussed
in
Chapter
15.
A.
Decomposition
of
Pure
Ammonium
Nitrate
The
thermal
decomposition
of
ammonium
nitrate
is
a
complex
phenomenon
with
many
intractable
factors.
Although
it
has
been
widely
investigated
by
a
number
of
research
workers,
a
complete
understanding
of
the
de-
composition
has
not
been
achieved.
When
ammonium
nitrate
is
heated
beyond
its
melting
point
(170°C),
it
begins
to
decompose.
At
high
temperatures
a
complex
series
of
re-
actions
occur.
Berthelot
has
described
seven
different
reactions
(19),
which
are
given
in
Table
10.
Later
investigations
have
shown
that
other
reactions
are
possible.
Reaction
(1)
is
reversible
and
endothermic,
whereas
all
the
other
reactions
listed
are
irreversible
and
exothermic.
It
is
generally
believed
that
in
any
decomposition
all
reactions
take
place
to
various
degrees
depending
on
the
conditions.
The
mode
of
decomposition
is
particularly
affected
by
the
temperature
and
its
rate
of
increase.
Up
to
about
250°C,
reaction
(2)
characterizes
the
decomposition
with
about
98%
of
the
nitrogenous
products
of
the
irreversible
reactions
be-
ing
nitrous
oxide
(20).
Reaction
(1)
becomes
significant
above
this
temperature.
Reactions
(3)—(6)
mainly
occur
above
300°C
and
are
usually
associated
with
explosive
and
detonative
behavior
of
ammonium
nitrate.
Reaction
(3),
which
generates
about
three
times
the
heat
of
reaction
(2),
is
believed
to
be
the
reaction
of
greatest
importance
in
a
detonation.
Reaction
(7)
was
found
to
occur
in
the
presence
of
spongy
platinum
and
gaseous
HNO
3
.
B.
Steady-State
Reaction
Temperature
As
described
in
the
previous
section,
when
ammonium
nitrate
is
heated,
it
decomposes
primarily
into
nitrous
oxide
and
water
exothermically.
Table
10
Decomposition
Reactions
of
Ammonidm
Nitrate
Number
Reaction
H
,
a
constant
P,
25°C
(kcal/mol)
Heat
release
,
b
constant
V,
27°C
(kcal/mol)
1
NH
4
NO
3
NH
3
+
HNO
3
+44.6
-41.7
2
NH
4
NO
3
+
N
2
0
+
2H
2
0
-8.8
13.2
3
2NH
4
NO
3
+
2N
2
+
0
2
+
4H
2
0
-28.2
30.5
4
2NH
4
NO
3
+
N
2
+
2NO
+
4H
2
0
-6.6
9.0
5
3NH
4
NO
3
+
2N
2
+
N
2
0
3
+
6H
2
0
20.8
6
4NH
4
NO
3
+
3N
2
+
2NO
2
+
8H
2
0
-22.2
29.8
7
5NH
4
NO
3
+
4N
2
+
2HNO
3
+
9H
2
0
-29.4
35.1
a
Approximate
values
at
constant
pressure
and
25°C
calculated
from
heat
of
formation
data
(17),
with
all
products
gaseous.
bData
from
Ref.
19.
Heat
release
at
constant
volume
and
27°C
with
all
products
gaseous.
P
ro
perti
es of
A
mmoni
um
N
i
t
rat
e
00
01
186
/
Shah
and
Roberts
Simultaneously
it
dissociates
endothermically
into
ammonia
and
nitric
acid.
The
combination
of
these
endothermic
and
exothermic
effects
results
in
a
steady
state
or
a
self-limiting
temperature
being
achieved
in
the
decomposing
mass,
provided
that
the
reaction
products
are
allowed
to
escape
freely.
This
phenomenon
has
been
investigated
by
Feick
and
Hainer
(18),
who
derived
the
following
relationship
for
reactions
(1)
and
(2):
3
A
H
v
P=
[1
+
where
P
is
the
total
pressure
on
the
reacting
mass,
p
is
the
dissocia-
tion
pressure
of
ammonium
nitrate,
A
H
v
and
H
r
represent
the
en-
thalpy
changes
per
mol
in
reactions
(1)
and
(2)
respectively
as
de-
scribed
in
Table
10,
and
Q
is
the
quantity
of
heat
added
to
the
sys-
tem
per
mol
of
ammonium
nitrate
decomposed.
The
effect
of
side
reactions
was
estimated
to
be
negligible.
The
relationship
was
tested
experimentally
for
the
adiabatic
case
(Q
=
0)
and
the
results
gave
good
agreement
between
the
calculated
and
meas-
ured
temperatures,
as
shown
in
Table
11.
Hainer
(21)
gives
the
following
simplified
equation
for
relating
the
total
pressure
P
(in
mmHg)
and
the
self-limiting
temperature
T(K):
-4.71
x
10
3
log
10
P
=
+
11.20
Free
escape
of
the
reaction
gases
is
a
very
important
requirement
in
this
mechanism
of
steady-state
temperature.
If
the
thermal
decom-
position
of
ammonium
nitrate
is
carried
out
in
situations
of
confinement
such
that
the
ammonia
and
nitric
acid
vapors
produced
from
reaction
(1)
cannot
freely
escape,
reaction
(1)
is
suppressed,
thereby
reducing
the
endothermic
effect.
Consequently
the
temperature
rises
and
can
reach
a
level
at
which
reactions
(2)—(6)
proceed
rapidly.
The
overall
Table
11
Variation
of
Steady-State
Temperature
with
Pressure
Under
Adiabatic
Conditions
Total
pressure
P
Temperature
Calculated
Measured
(cmHg)
(
°
C)
(°C)
38
266
271
76
289
292
114
304
306
152
314
315
2
Q
AH
r
p
Properties
of
Ammonium
Nitrate
/
187
effect
is
then
exothermic
and,
if
sustained,
can
develop
into
a
run-
away
situation.
C.
Decomposition
Rates
of
Pure
Ammonium
Nitrate
A
number
of
workers
have
studied
the
decomposition
kinetics
of
am-
monium
nitrate
(21-30).
Some
of
these
are
shown
in
Figure
8,
and
Bennett's
comparison
(30)
of
the
kinetic
data
of
different
workers
is
given
in
Table
12.
D.
Catalytic
Effects
on
Decomposition
A
number
of
substances
are
known
to
have
a
catalytic
effect
on
the
decomposition
of
ammonium
nitrate,
the
notable
ones
being
acid,
chloride,
and
chromates.
Water
is
also
shown
to
have
an
effect
on
the
decomposition.
1.
Effect
of
Water
Ammonium
nitrate
is
very
difficult
to
dry
completely
and
water
is
a
product
of
any
decomposition;
so
in
all
ordinary
circumstances
one
studies
the
decomposition
of
the
moist
substance.
A
small
amount
of
moisture
is
believed
to
have
a
profound
effect
on
the
decomposition
behavior
of
ammonium
nitrate.
Friedman
and
Bigeleisen
(31)
and
Keenan
(32)
found
that
thorough-
ly
dried
samples
of
pure
ammonium
nitrate
showed
remarkable
thermal
stability
and
could
be
heated
to
300°C
without
any
perceptible
decom-
position.
Introduction
of
a
small
amount
of
water
vapor
caused
decom-
position
at
180°C.
The
presence
of
a
small
amount
of
moisture
is
be-
lieved
to
be
essential
for
a
smooth
decomposition
of
ammonium
nitrate
at
moderate
temperatures.
The
roles
played
by
water
and
acid
have
been
investigated
by
Smith
(33).
More
recent
work,
however,
by
Rosser
et
al.
(23)
and
Bennett
(30)
suggests
that
moisture
and
water
vapor
can
have
an
inhibiting
effect
on
the
decomposition
of
ammonium
nitrate.
2.
Effect
of
Nitric
Acid
and
Ammonia
The
catalytic
effect
of
nitric
acid
on
the
thermal
decomposition
of
ammonium
nitrate
has
been
well
established
(22).
In
the
absence
of
added
acid
the
dissociation
of
ammonium
nitrate
to
ammonia
and
nitric
acid
produces
the
acid
required
to
initiate
the
decomposition
reaction.
Ammonia,
because
of
its
higher
volatility,
,escapes
more
readily
than
the
acid.
If
the
conditions
are
such
that
the
acid
accumulates
in
the
melt,
the
decomposition
is
autocatalytic.
If
the
vapor
can
escape,
the
concentration
of
nitric
acid
soon
attains
a
constant
value.
This
steady
concentration
of
acid
has
been
shown
to
be
a
function
of
temperature
(33).
10
-
10
-
ur
(0
0
:-E
0
c00
10
a
0
10
EC
la
10
188
/
Shah
and
Roberts
Reciprocal
absolute
temperature
(°K)
x
10
3
2
3
2.2
2.1
2.0
1.9
1.8
0
.
a
,9
io
,
a
Q-_
,
a
a
,
K\
Q.!
ti'
0
,-.•
._/
e
..
?
)
0
c
0
180
200
220
240
260
260
300
Temperature,
°C
Figure
8
Ammonium
nitrate
decomposition
rates
versus
temperature.
(Data
from
Refs.
21,
24-27,
and
30.)
P
ro
perti
e
s
of
A
mm
oni
um
N
i
t
r
at
e
I...I.
CO
CO
Table
12
Comparison
of
Rate
Constants
from
Various
Sources
Reference
Temperature
range
(°C)
Rate
constant
(sec
-1
)
Rate
constant
(sec
-1
)
At
200°C
At
240°C
22
200-300
10
14.2
c
exp
-31,0
40
a
0.49c
6.75c
RT
23
225-275
10
13.2
exp
,00
-41
0
1.15
x
10
6
5.62
x
10
5
RT
24
170-280
10
11.5
exp
-36,500
±
1800
4.37
x
10
6
9.12
x
10
5
RT
-39,000
±
3
000
7.76
x
10
6
19.5
x
10
5
25
240-300
10
12.9
exp
RT
26
240--360
10
13.8
exp
-40,500
±
2
500
12.6
x
10
6
35.5
x
10
5
RT
27
220-270
10
12.28
exp
-38,300
±
300
3.80
x
10
-6
9.12
x
10
-5
RT
28
240
±
1
3.30
x
10
5
29
230-265
10
16.478
exp
4
-49
R
,
T
50
0.425
x
10
6
2.70
x
10
5
30
170-200
,0
11.0
exp
-35,500
RT
±
1500
3.80
x
10
-6
7.24
x
10
-5
a
c
=
concentration
of
free
nitric
acid.
190
/
Shah
and
Roberts
Other
acids
also
generally
tend
to
enhance
the
decomposition
of
ammonium
nitrate.
Ammonia
and
similar
alkali
substances
have
an
inhibiting
effect
on
the
decomposition.
Wood
and
Wise
(22)
in
fact
found
the
reaction
to
be
almost
completely
inhibited
at
significant
ammonia
concentrations.
This
effect
is
claimed
to
be
due
to
the
neutralization
of
any
acid
form-
ed.
This
stabilizing
effect
of
ammonia
is
utilized
in
the
manufacture
of
ammonium
nitrate
by
maintaining
the
melts
and
the
final
products
on
the
alkaline
side
of
the
neutral
point.
The
effect
of
the
partial
pressure
of
ammonia
in
the
gaseous
phase
on
the
decomposition
of
ammonium
nitrate
has
been
recently
investi-
gated
(34).
3.
Effect
of
Chloride
The
presence
of
a
chloride
has
two
distinctly
different
effects,
de-
pending
on
the
level
of
inclusion.
When
present
as
an
impurity
in
small
amounts,
it
can
accelerate
the
thermal
decomposition
of
ammonium
nitrate.
When
mixed
at
much
higher
levels
(in
excess
of
about
10%),
it
can
produce
mixtures
capable
of
burning
in
a
self-sustaining
manner.
a.
Chloride
as
an
Impurity:
The
chloride
ion
has
a
strong
catalytic
effect
on
the
decomposition
of
ammonium
nitrate,
producing
about
a
10-
to
100-fold
increase
in
the
rate,
and
as
little
as
0.1%
can
produce
a
marked
effect.
This
field
has
been
extensively
studied
by
a
number
of
research
workers
(20,21,32,35-37).
For
chloride-catalyzed
decom-
position
to
proceed,
presence
of
acid
is
essential
(this
can
be
readily
produced
as
a
result
of
dissociation
if
not
added);
addition
of
ammonia
or
other
basic
substances
prevents
the
decomposition.
In
the
presence
of
chloride
the
gas
composition
changes
with
a
marked
increase
in
the
nitrogen
content
at
the
expense
of
nitrous
oxide.
Chlorine
gas
is
evolved
in
the
decomposition.
Before
the
vigorous
reactions
set
in,
there
is
usually
an
induction
period.
The
rate
of
decomposition
is
influenced
by
the
chloride
concentration.
The
induction
period
has
been
examined
in
detail
by
Guiochon
(24)
and
Colvin
et
al.
(37).
There
is
no
generally
agreed
precise
definition
of
the
induction
period.
It
can
vary
from
a
few
minutes
to
a
few
hours,
depending
on
the
conditions
such
as
temperature
and
presence
of
free
acid.
Guiochon
found
a
linear
relationship
between
the
logarithm
of
the
induction
period
and
the
reciprocal
absolute
temperature.
The
pres-
ence
of
free
acid
shortened
the
period
considerably.
The
chloride
con-
centration
did
not
affect
the
induction
period.
Colvin
et
al.
used
a
flushing
process
to
make
correct
measurements
of
induction
periods.
They
concluded
that
the
mechanism
consists
of
an
acid-autocatalyzed
adjustment
of
the
melt
acidity
from
one
steady-state
level
to
another.
When
catalyzed,
the
rates
of
decomposition
increase
substantially.
A
comparison
of
the
rates
measured
by
different
workers
is
not
easy,
Properties
of
Ammonium
Nitrate
/
191
as
the
data
have
been
obtained
under
different
conditions.
Heiner
(21)
found
the
rate
to
increase
to
approximately
1000
times
that
of
pure
ammonium
nitrate
at
175°C
when
catalyzed
by
1%
NaC1
after
an
induc-
tion
period.
Kerdivarenko
and
Davidovich
(38)
found
that
additions
of
0.06
and
0.03%
NH4C1
to
pure
ammonium
nitrate
increased
the
rates
by
75-
and
30-fold,
respectively,
at
195.6°C.
b.
Chloride
as
a
Component
in
Fertilizers:
Ammonium
nitrate
and
potassium
chloride
are
extensively
used
as
components
for
producing
a
wide
range
of
NPK
and
NK
compound
fertilizers.
A
third
component
is
usually
used
as
a
source
of
phosphate.
Certain
compositions
of
these
mixtures
are
capable
of
undergoing
self-sustaining
decomposi-
tion.
A
hot
source
is
required
to
initiate
the
reaction,
but,
once
initiated,
the
mixture
can
continue
to
decompose
even
after
the
hot
source
is
removed.
This
phenomenon
is
commonly
termed
as
"cigar
burning."
Toxic
gases,
for
example,
oxides
of
nitrogen,
hydrogen
chloride,
and
chlorine,
are
given
off.
The
cigar-burning
behavior
of
a
number
of
three-component
systems
based
on
ammonium
nitrate,
potassium
chloride,
and
a
third
component
have
been
studied
(39,40).
4.
Effects
of
Other
Materials
A
number
of
other
chemical
substances
have
shown
catalytic
effects
on
the
decomposition
of
ammonium
nitrate
(24,41).
Bromides
and
iodides
produce
catalytic
effects
similar
to
chlorides.
Fluorides
do
not
appear
to
have
any
significant
catalytic
effect
(24).
Compounds
of
chromium,
for
example,
chromates,
Cr
2
O
3
,
K2Cr
2
O
7
,
and
Cr(NO
3
)
3
,
appreciably
increase
the
decomposition
rates.
Zinc
and
compounds
of
manganese,
copper,
nickle,
and
cobalt
have
also
been
found
to
accelerate
the
rates.
Organic
and
carbonaceous
materials
appreciably
enhance
the
decom-
position
reactions
in
which
the
materials
are
oxidized.
E.
Inhibiting
Effect
of
Urea
Urea
has
an
inhibiting
effect
on
the
thermal
decomposition
of
pure
am-
monium
nitrate,
as
well
as
of
ammonium
nitrate
containing
chloride
or
organic
material
(42,
43).
Studies
of
the
stability
of
urea-ammonium
nitrate
mixtures
reveal
a
mutually
stabilizing
effect,
each
stabilizing
the
other
in
terms
of
thermal
decomposition
(44,45).
Urea
is
much
less
stable
than
ammonium
nitrate;
it
begins
to
decompose
even
at
130°C.
In
the
presence
of
a
large
excess
of
ammonium
nitrate
it
shows
con-
siderable
thermal
stability.
The
maximum
inhibiting
effect
of
urea
appears
to
be
produced
be-
tween
0.1
and
1%.
Increase
in
the
urea
content
beyond
about
1%
leads
to
a
reduction
in
the
thermal
stability
of
the
mixture
because
of
the
intensified
decomposition
of
the
urea
itself.
The
mechanism
of
the
in-
hibiting
effect
of
urea
is
complex
and
not
fully
understood.
192
/
Shah
and
Roberts
F.
Miscellaneous
Chemical
Reactions
Fused
ammonium
nitrate
rapidly
oxidizes
and
dissolves
many
metals,
and
its
attack
on
powdered
metals
is
often
violent.
Several
metal
oxides
dissolve
in
the
molten
salt
to
form
the
metal
nitrate
and
liberate
ammonia
and
water.
With
ammonium
sulfate
two
complex
salts
are
formed:
(NH
4
)
2
SO4
2NII4NO3
and
(NH
4
)2SO
4
3NH
4
NO
3
.
Ammonium
nitrate
forms
solid
solutions
with
a
number
of
ammonium
salts.
It
forms
complex
salts
with
a
number
of
nitrates.
G.
Self-Heating
in
Fertilizers
Pure
ammonium
nitrate
shows
no
tendency
to
self-heat
either
alone
or
when
packed
in
paper
bags
at
temperatures
in
the
0-75°C
range.
In
the
tests
carried
out
by
the
Association
of
American
Railroads
Bureau
of
Explosives
(46)
ammonium
nitrate
fertilizers
maintained
at
tempera-
tures
up
to
80°C
for
periods
over
30
days
failed
to
show
self-heating,
even
when
the
fertilizers
were
wax
coated.
A
specially
instrumented
8500-ton
load
of
wax-coated
ammonium
nitrate
fertilizer
in
a
ship's
hold
showed
no
abnormal
temperature
rise
over
a
period
of
14
days
in
a
trial
carried
out
by
the
U.S.
Bureau
of
Mines
(47).
The
above
re-
sults
are
consistent
with
the
decomposition
behavior
of
ammonium
ni-
trate.
The
fertilizers
are
produced
in
the
alkaline
state,
which
sup-
presses
the
decomposition
reactions.
Any
loss
of
ammonia
in
the
sub-
sequent
storage
under
normal
conditions
is
negligible
and
therefore
the
beneficial
effect
of
ammonia
continues
to
protect
the
fertilizer.
The
possibility
of
self-heating
taking
place
in
ammonium
nitrate
either
acidic
in
nature
or
containing
catalytic
impurities
cannot
be
ruled
out.
Quantitative
study
and
prediction
of
the
self-heating
effects
by
mathematical
modeling
techniques
are
very
difficult
in
the
case
of
am-
monium
nitrate
owing
to
the
fact
that
the
overall
reactions
are
complex
and
depend
on
a
number
of
factors.
Simple
models
have
been
developed
(21)
which
do
not
appear
to
have
taken
into
account
the
endothermic
effects.
H.
Oxidizing
Properties
Ammonium
nitrate
is
a
powerful
oxidizing
agent;
it
can
attack
a
large
number
of
organic
substances
and
reducing
salts
when
heated.
It
facilitates
the
combustion
of
combustible
substances
when
in
intimate
contact
with
them.
The
combustion
can
proceed
even
when
air
is
ex-
cluded.
Thus
ammonium
nitrate
can
present
a
fire
hazard
when
stored
with
combustible
materials,
even
though
it
itself
does
not
burn.
Reactions
of
ammonium
nitrate
with
organic
substances
accelerate
rapidly
owing
to
two
separate
effects:
Firstly,
organic
substances
appear
to
have
a
catalytic
effect
(21)
and,
secondly,
a
large
amount
Properties
of
Ammonium
Nitrate
/
193
of
heat
is
liberated
in
the
oxidation
reactions.
For
example,
2NH
4
NO
3
+
C
+
2N
2
+
CO
2
+
4H
2
0
OH
=
-75.2
kcal/mol
at
25°C)
whereas
the
main
decomposition
reaction
giving
N
2
0
and
H
2
O
only
produces
8.8
kcal/mol
of
ammonium
nitrate
decomposed.
Thus
mixtures
of
ammonium
nitrate
with
fuel
oil
and
similar
sub-
stances
have
the
ability
to
release
large
amounts
of
chemical
energy.
Such
mixtures
therefore
make
good
blasting
agents.
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