PART I
THE HYDROLYSIS OF PROPIONITRILE BY STRONG ACIDS
PART II
SECRET (War work)
A THESIS
by
3. D. McLean, B.Sc.
Submitted to the Faculty of Graduate Studies and Research in partial fulfillment of the . requirements for the degree of Master of Science McGill University,
Montreal, Canada.
September
1941 Physical Chemistry
M.Sc.
James Douglas McLean
The
Hydrolysi~
of Propionitrile by Strong Acids
A study of the strong acid catalysed hydrolysis
of nitriles has been undertaken for propionitrile, with
hydrobromic, nitric, and sulphuric acids.
The kinetics
of the reaction are similar to those found by previous
investigators for hydrochloric acid.
It was
found
that the rate of hydrolysis is not indeuendent of the
decomposition
of
the intermediate amide at acid
concentration above five normal.
A decrease in the
activation energy with increase in acid concentration
was found for all acids.
McGill University
Montreal, Canada
September 1941
PMTI
THE HYDROLYSIS OF PROPIONITRTLE BY STRONG ACIDS
TABLE OF CONTENTS
PAGE
ABSTRACT
1.
HISTORICAL
2 ..
NATURE OF PROBLEM
13.
EXPERIMENTAL
17.
Materials
Apparatus
Procedure
Method of analysis
RESULTS
24.
DISCUSSION OF RESULTS
41.
CONCLUSIONS
53.
BIBLIOGRAPHY
55.
1.
ABSTRACT
A study of the strong acid catalysed
hydrolysis of nitriles has been undertaken for
propionitrile, with hydrobromic, nitric, and
sulphuric acids.
The kinetics of the reaction are
similar to those found with hydrochloric acid (33).
It was found that the rBte of hydrolysis is not
independent of the decomposition of the intermediate
amide at acid concentration above five normal.
A decrease in the activation energy with increase
in acid concentration
W8.S
found for all acids.
2. HISTOHICAL Many theories have been proposed regarding
the part that catalysts play in chemical reactions.
The fact that catalysis covers a very wide field in
chemistry has attracted much interest as to what its
mechanism is in its different applications.
It is
known in the case of some reactions that the presence
or absence of the most minute quantities of foreign
substances will govern the velocity of the reaction,
or may even determine whether or not it will go in
finite time.
As a result of the work of van't Hoff,
Arrhenius, and Ostwald, in the nineteenth century, it
was believed that the catalytic effect of acids and
bases was independent of individua.l chemical character­
istics of the catalyst but was proportional to the
degree of dissociation.
Their catalytic activity was
attributed to free hydrogen and hydroxyl ions in
solution.
Ostwald stated that only free ions could
act as catalysts (30).
Certain discrepancies were
observed, but in general, the dependence of rate on
hydrogen ion concentration was established, especially
at low acid and low total concentrations.
This
dependence bas been so well established that ester
hydrolysis and sucrose inversion have been recommended
for the determination of hydrogen ion concentration.
From experimental evidence obtained by
Blanksma (4) and la.ter by Acree and Johnson
(2), who
studied the transformation of acetyl chloroamino
benzene into p-chloracetanilide, the dual theory of
catalysis was formulated to replace the defective
simple dissociation theory.
The above reaction is
catalysed by hydrochloric acid with a velocity proport­
ional to the square of the acid concentration, not to
the first power as demanded by simple hydrogen ion
catalysis.
It was concluded that the hydrogen ien is
not the only catalyst, but that a greater or smaller
effect was also due to the undissociated molecules;
hence the name "dual theory".
The dual theory of catalysis was developed
mainly by Taylor (39,40,41,42), and Dawson (8,9) and
his coworkers, from studies on the hydrolysis of
ethyl acetate.
Acree (1) has applied the ideas of the
dual theory of acid catalysis to basic catalysis.
He
concludes that the catalytic effect is due not only
to the hydroxyl ion, or negative ion, but also to the
undissociated molecules present.
dual theory of catalysis
Application of the
has been largely restricted
to comparative studies of catalysis in definite
solvents, and many inadequacies have been recognized.
4. Bronsted (5,6) has put forth an extended
theory of acid and base catalysis in which he
attributes the property of acid catalysis to all
molecules capable of giving off a bydrogen nucleus
or proton.
In other words, be defines an acid according
to the scheme:
Acid ~ Base + H+.
The effect of the
various acids would then depend upon the ease with
which the hydrogen nucleus was split off.
This in
turn would be dependent on the strength of the acid.
It follows that this is correct as long as the
catalysed reaction is not subject to any specific effect.
However there is the possibility that a reaction may
be catalysed by a certain acid and not by another; for
example, the transformation of acetochloroamino-benzene
into p-chloroacetanilide is catalysed by hydrochloric
acid and not by nitric acid.
Analogous conclusions can be drawn for basic
catalysis.
In general, basic catalysis is related
to the ease with which a hydrogen nucleus is taken up,
and all substances capable of so doing will exhibit
characteristics associated with basic catalysis.
In the classical theory of catalysis hydrogen
ion and hydroxyl ion hold a unique position.
Similarly
the dual theory considers the hydrogen ion and the
undissociated acid molecule as quite different, which
in turn, makes a difference between hydrogen
ion
5. catalysis and catalysis by acid molecules.
The
extended theory of Bronsted, however makes hydrogen
and hydroxyl ion catalysis only particular instances
of general acid and base catalysis.
From thermal
considerations, Bronsted points out that the hydrogen
ion must always be presumed hydrated; hence the
catalytic agent in hydrogen ion catalysis is not the
proton as such, but the oxonium ion, H30+.
The mechanism of catalysis has been described
as follows:
The primary condition for catalysis is
the formation of new complexes of the molecules of the
catalyst and of the substrate.
called "critical complexes".
The complexes are
The first step in the
reaction may be assumed to be the transfer of a proton
of the catalyst into the" substrate molecule.
The
complex may then decompose instantly, without fUrther
energy input, into product and catalyst.
The mechanism suggested predicts one
important thing in regard to the nature of the
substrate and that is that acid and basic catalysts
require substrates which have basic and acidic natures
respectively.
These properties do not need to be
developed to any great extent, in fact they do not
need to be developed to the extent where they can be
detected by any general means of measuring acidity
and basicity.
6. The dependence of the catalytic effect
upon the strength of the catalysing acids and bases
is an important question in the Bronsted theory.
It
would be expected that the catalytic effect would
increase with increasing strength of the acid and of
the base, since the ease with which a hydrogen nucleus
is given off by an acid and taken up by a base Can be
assumed to be greater, the greater the strength of
the acid and of the base.
Bronsted points out that the influence of
salt upon the reaction velocity (5) is not a kinetic
abnormality even in solutions of low concentration.
It is to be expected that there will be a kinetic
effect since the addition of a salt will change the
properties of the solvent.
One of the earliest
objections to the Arrhenius theory of hydrogen ion
catalysis was the fact that considerable variations
were observed in the reaction constant when a neutral
salt was added.
It is not necessary that the neutral
salt have an ion in common with the catalysing acid.
The kinetic salt effect may be divided
into two groups in dilute solution.
In some reactions
the change of velocity constant is proportional to
the concentration of salt; it is then said to be linear.
In other cases there is a relatively greater effect the
more dilute the salt solution.
J
Here the variation in
7. velocity of the reaction seems to be proportional to
some power of the salt concentration and this state
is termed exponential in contrast to linear (7).
The linear salt effect is peculiar to reactions
between molecules or between a neutral molecule and
an ion.
The exponential salt effect always occurs in
reactions between two charged molecules.
The above
observations have formed an important part in the
development of the general kinetic theory, in which
not only the concentrations of the reacting substances
but also their activity coefficients are supposed to
govern the course of the reaction.
Bence the primary
salt effect is a direct result of the changes in the
activity coefficient of the molecules involved, that
is, in changes in the nature of the solvent.
The
secondary salt effect is caused by a shift in the
equilibrium of the ions taking part in the reaction,
following the addition of the salt.
This change of
equilibrium may either cause an increase or a decrease
in the concentration of the ions which take part in
the reaction.
ACID HYDROLYSIS OF NITRILES:
While considerable success has been had
in dealing with the theoretical aspects of acid catalysis
in dilute solution, the same can scarcely be said of
strong acid catalysis.
Many cases of strong acid
catalysis have been studied, one of the most
8. interesting of which is the hydrolysis of nitriles.
In dilute acid solution, the rate of the
reaction is proportional to the hydrogen ion
concentration (12,16) but, in concentrated acid the
rate becomes very sensitive to acid concentraticn.
Indeed, the rate of nitrile hydrolysis probably
increases more rapidly with acid concentration than
the rate of any other acid catalysed reaction on
record.
For example, for the hydrochloric acid
catalysed hydrolysis of propionitrile it was found
(33) that a sixteen-fold increase in acid concentration,
from 0.5 N to 8.4 N, increased the rate by a factor
of 10 3 .
With cyanoacetic aCid, the rate increased
by a fEctor of 104 for a similar change in acid
concentration.
The hydrolysis of nitriles is known to be
made up of two successive reactions.
of amides from the cyanide.
(2)
(1)
The formation
The formation of
the corresponding acid and ammonia from the amide.
The equation in acid solution is as follows:
(1)
RCN + H20 + H+ ~ RCONH 2
RCONH 2 + H20 + H+ ~ RC0 2H + NH4 (2)
The second reaction is reported to be very
much faster than the first (16~20,43) and therefore
the first reaction, the hydrolysis of nitrile to amide.
9. is generally presumed to be the rate governing step.
The rate of hydrolysis is therefore usually followed
by determining the amount of ammonia produced in the
reaction.
Actually, as discussed later, work in these
laboratories shows that special analytical methods
are necessary in certain circumstances.
The equations for the alkaline hydrolysis
of nitriles are similar to the acid
cat~lysed
hydrolysis, with the difference that free ammonia is
liberated (34).
Krieble and McNally (19) have considered
the ion activity as a factor in the mechanism of the
hydrolysis.
They have investigated the hydrolysis
of hydrocyanic acid with hydrochloric acid and find
fairly good constancy of the ratio: k/(mean ion activity)2
at higher concentrations of acid; at lower acid
concentration the constancy does not hold.
Krieble and Noll (20) studied the hydrolysis
of acetonitrile, propionitrile, and hydrocyanic acid
with hydrochloric, hydrobromic, and sulphuric acids over
a wide range of acid concentration and at a temperature
of 65 0 C.
They found , for the hydrolysis of these
nitriles with hydrochloric acid, that k varied as the
square of the mean ion activity, although the agreement
was rather poorer than in the case of hydrocyanic acid.
10. According to Krieble (20,21,22) the rate at higher
acid concentration is a function of the square of the
mean ion activity; he postulates catalysis by
undissociated molecules.
In general he found for
hydrochloric and sulphuric acids that the velocity
constants were not directly proportional to either
the concentration of acid or to the hydrogen ion activity.
Krieble's conclusions are based essentially upon data
obtained at only one temperature.
The majority of the work on the acid catalysed
hydrolysis of nitriles has been concerned with finding
the variation of the reaction velocity constant with
change in acid concentration, and from these results
to determine the mechanism of the catalysis.
The only
study yet reported of the temperature coefficient of
nitrile hydrolysis has been made by Kilpi (16) who
worked with dilute acid solutions in a water-alcohol
medium.
Kilpi found that the .Arrhenius equation for
the temperature coefficient of reaction holds for
nitrile hydrolysis and he calculated the temperature
coefficient for various nitriles.
In general it
was found that reactivity decreased with increase in
the molecular weight of the nitrile except in the case
of propionitrile which hydrolyses faster than
acetonitrile.
11. The hydrolysis of aliphatic and aromatic
nitriles by concentrated sulphuric acid has been the
subject of a recent paper by Karve and Gharpure (14).
However the choice of glacial acetic acid as solvent
by these workers complicates the interpretation of
their results.
An extensive investigation of the hydrochloric
acid catalysed hydrolysis of nitriles at various
temperatures has been made in this laboratory (33).
Contrary to information reported in the literature it
has been found that there is an induction period with
hydrochloric acid above five normal, because the rate
of hydrolysis in concentrated acid solutions is not
independent of the hydrolysis of the intermediate amide.
Moreover a large decrease in activation energy for all
nitriles with increase in acid concentration has been
observed.
The values of E obtained are slightly
higher than those obtained by Kilpi (16) with hydrochloric
acid of concentration of approximately 0.5 normal.
However Kilpi employed water-alcohol mixtures as
solvent.
It is interesting to note that the change in
E is in the opposite direction to that found for amides
(32).
The general variation in activation energy with
concentration, for concentrated acid catalysis,
suggested by Leininger and Kilpatrick (26), is born
out by data presented (33), an example of which is
given in the following table.
12. PROPIONITRILE Normality of Hydrochloric acid 1.00
4.00
5.02
E
(ca1s)
25,600
24,800
23,900
2l,6UO
6.46
8.41
19,800
10.03
19,700
13. NATURE OF PROBLEM The present investigation constitutes a
study of the hydrolysis of propionitrile catalysed
by hydrobromic, nitric, and sulphuric acids.
investigation has been carried out over a
of acid concentration and over
a
temperatures, to determine if
the
The
range
wide range
change
activation energy found previously is a
of
of
general
phenomenon of acid catalysed nitrile hydrolysis or
is specific for hydrochloric acid; to determine how
the rate constant and the energy of activation vary
with various strong acids; and to correlate the
effects of the different. acids and determine the
nature of the variations of nitrile hydrolysis with
different catalysts.
In reactions between two molecules and a
catalyst, Lowry (27) supports the view of two
bimolecular reactions in which an activated complex
is formed which is the rate governing factor.
The
general type where A and B react with a catalyst to
form D follows.
A + B ... AB
AB + Catalyst
AB' .. D
~
AB' + Catalyst (rate governing)
(fast reaction)
14.
For similar reactions Boeseken supports
triple collisions.
A + B + Catalyst ~ ABCt
ABC'
~
(rate determining)
D + Catalyst
He considers that acceleration causeu by a
catalyst is not due to a diminution in the energy of
activation, but rather to an increase in the entropy
change with activation, in the sense of Schiffer and
Kohastamm (35), who assume a change in the steric
factor, the catalyst bringing about an increase in
the sensitive range of the molecule.
La Mer and
Kamner point out that quite generally in solution
the reaction constants are able to compensate
increases in E by reason of the entropy change with
activation, entirely apart from any steric factor.
As a result acceleration with enhanced activation
energies are not only possible but are often found.
Since uncatalysed reactions are usually bimolecular
and seldom trinlolecular, the view supported by Lovrry
is correct, although kinetic equations are the same
in both cases,
e.g.
~
dt
=k
[AJ [BJ[cl
It has been proposed that the action of the
catalyst in intermediate compounds is to loosen
atom bond necessary for reaction.
the
The catalyst must
attack the substrate in the right place
and
the
15. intermediate compound should not be too stable. Hence
mere formation of an addition compound does not
necessarily mean it is effective.
In gas reactions
intermediate products are limited and its life does
not affect the kinetic equations.
In solution, for
the above reason, the p0ssibilities are great and
hence the life period is of more interest.
The role played by intermGdiate products
is brought out in a study of consecutive re&ctions.
Suppose that we have substance A which forms an
intermediate compound C and that this in turn forms B.
If the rate of change of A to C is much faster than
C to B then the observed order of the whole reaction
A ...B will solely be fixed by the order of the slower
reaction
C~
B.
Whatever the rate of the first
reaction, C ... B alone is accessible to measurement.
The observed order in consecutive reactions is
determined by the slower.
At the initial stages of consecutive reactions
the rate A ... C will be a maximum.
During the first
period of the reaction the amount of C is gradually
increasing, and the rate of formation of B will increase
in a corresponding way.
Tho rate of formation of B
will be a maximum whon the system contains a maximum
of C.
This gives rise to an induction period.
16. The theory of intermodiate compounds yields
a type of velocity eg,uation observed in catalysis in
solution and hence is probable.
However definite
deductions regarding the exact form of the mechanism
and the relation of its steps cannot usually be made
from the observed kinetics, especially when the
observations represent an ambiguous limiting case of
general tendencies as we find with the most effective
catalysts (acid-base).
Only in rare cases where at
least the intermediate steps can be observed separately,
or in cases in 'which other physical observations permit
conclusions about a postulated step in the total
process is complete clarification of the mechanism
possible.
,
17.
EXPERIMENTAL
APPARATUS
A constant temperature bath operated at
l25 0 C.
consisted of a closed cylindrical copper boiler
fitted with a condenser at the top.
The boiler was
partially filled with isoamyl alcohol and heated.
Wells set in the top of the boiler and filled with
oil provided the receptacle for the tubes containing
the propionitrile-acid solution.
These wells assumed
the temperature at which isoamyl alcohol boils.
stirring was provided.
No
The condenser prevented any
isoamyl alcohol escaping from the boiler.
Pressure was
regulated and kept constant by means of a head of
compressed air and a mercury bubbler attached to the
top of the condenser.
A temperature of IOOoC.
was
obtained by using water instGnd of isoamyl alcohol in
a similar boiler.
All other temperatures were obtained by using
open oil baths, heated electrically and regulated by
means of a mercury thermoregulator and a pair of relays.
Stirring was provided in these baths.
In the open batu::.
constant to ±O.lo.
t\;;;wi>~rature w~s ke~t
In the boilers tem~erature was kept
18.
constant to within ± 0.2°.
Pipettes and volumetric flasks used in
measuring the various liquids were all standardised
for the temperature at which they were to be used.
Certain solutions such hs the propionitrile-acid
mixtures were handled at the temperature of an ice
water mixture, and the pipettes for this purpose 'l"lere
so calibrated.
Burettes used in the ammonia analYSis
were calibrated.
MATE:RIALS
Propionitrilc was obtained commercially
and was redistilled three times with a Vlhitmore
fractionating colurrill.
The fraction used had a range
of boiling point of O.2 oC.
Sodium hydroxide was C.P. grade pellets.
Analytical grades of hydrobromic, nitric, and
sulphuric acids 1."ere used without further purification.
These acids were standardised by specific gravity as
found with a specific gravity bottle; the percentAge of
acid present was found from values given in the
International Critical Tables.
This standardisation
was checked by titration of a diluted sample of the acid
19. dark bottles.
The sulphuric acid used in the ammonia
determination was standardised by running ammonia
analyses on a stc1Ddard amruonic.m sulphate solution,
in a manner similar to the determinations made on the
propionitrile solution.
The sodium hydroxide was
standardised by titrCl.tion ageinst the standard sulphuric
acid.
The standard acid and alkali were both
approximately 0.4 normal.
PROCEDURE
For a given run, a given volume of the
propionitrile was int.I.'oduced into a given volume of
the acid in a flask cooled in an ice bath.
To obtain
final acid concentrations of 1,2 , e.nd 4 normal,
solutions y.:ere made which, on dilution with an e;;ual
volume of the propionitrile solution would give the
required normality.
Account was taken of the volume
change of the acid on dilution.
The more concentrated
solutions used in the: investigation Vlere made by
adding a solution of propionitrile of a given strength
to the ice cold stock acid.
The normalities were then
celculated, account being tftken of the volume change
involved.
2u.
The solutions of acid and propionitrile were
thoroughly mixed and approximately eight pyrex glass
tubes holding from one to six cubic centimeters were
filled and sealed for each run.
The tubes were filled
so as to allow room for expansion of the liquia, but
dead space was kept a minimum to reduce the nitrile
in the gas phase to a negligible amount.
The tubes
were pIa ced simultaneously in a constcmt temperature
bath and the time noted.
Tubes were removed from time
to time, ceoled a.nd analysed.
If it was impossible to
analyse the contents of a tube immediately it Was
stored in an acetone-dry ice mixture.
The time required. for the tubes to assume
the temperature of the bath was determined experimentally,
by placing tubes with thermometers in the baths and
following the temperature rise.
It was then possible
to calculate the time that should be allowed for the
solutions to come to the temperature of the bath.
Since
the tubes on removal from the bath were immediately
cooled by immersion in an acetone-dry ice mixture, it
was assumed that tho reaction stopped instantaneously
and no correction for cooling was applied.
Since even
the shortest runs were approximately of two hours
duration the errors in calculating the correct time of
the reaction were negligible.
21.
METHOD OF ANALYSIS
Nitriles are hydrolysed to yield ammonia and
the corresponding acid according to the equation:
RCN + H20
~
RCONH 2 + H20
~
RC0 2H + NH3
The reaction may be followed by the determination
of ammonia.
Folin's aspiration method (11) wes used
in the present study.
The air was purified by passing
it through concentrated sulphuric acid and then
concentrated sodium hydroxide solution.
The soluticn
to be analysed was placed in a l&rge test tube which
could be easily attached to the aspiration train.
The air inlet was used as a pipette for introducing
a saturated solution of sodium hydroxide to the test
tube and at the same time preventing escape of ammonia
from the solution.
The air passed out of the test tube
and through a bubbler in a standard solution of
sulphuric acid in a long necked flask.
The
bubbler
was made by closing the end of a length of glass
tubing and then making a number of very small holes
in it by means of a hot platinum wire.
Two Kjeldahl
traps were used to protect the standard acid from
alkaline spray from the test tube.
Solutions of the more dilute acids, one to
four normal, were analysed, a beaker of ice water was
placed around the test tube containing the propicnitrile
22. so that hydrolysis of propionitrile during the
analysis would be negligible.
was fifteen minutes.
The time of aspiration
When higher concentrations of
acids were used an induction period was observed.
This
was eliminated by aspiration for twenty one minutes,
the test tube being surrounded by ice water for six
minutes and then by water at 70 0 C. for the remaining
fifteen minutes.
With this method of a.nalysis the
Kjeldahl traps are also heated to prevent excessive
condensation.
During the first six minutes all
unreacted propionitrile is removed from the s(;lution
and no error is introduced due to its hydrolysis
during the analysis.
During the remaining fifteen
minutes at 70°C. all propionamide remaining in the
solution is completely hydrolysed.
By analysin6 in
this manner it is possible to elirrinate the induction
period completely.
In both methods air was passed through the
system slowly for the first three minutes during which
time the greater part of the ammonia was removed.
The
rate of aspiration was increased for the remaining
time.
The times mentioned were determined by running
blanks on propionitrile and propionamide.
Such blanks
showed that no appreciable hydrolysis of propionitrile
occurred during the analysis and that the propionamide
would have reacted completely in the time allowed and
at the alkaline concentration used.
Such blanks
23 ..
showed also that the ammonia was completely retained
by the standard acid.
This acid is finHlly titrated
with the standard sodium hydroxide.
The propionitrile solutions used in the
investigation were approximately 0.2 normal.
24. RESULTS An induction period was observed for all
acids with an acid concentration above four normal.
This agrees with previous work (32,33) on the
hydrolysis of nitriles with hydrochloric acid where
it was found that the rate of amide hydrolysis is of
the same order as nitrile hydrolysis in acid
concentrations above five normal.
Here the rate of
hydrolysis as found by analysis of ammonia is dependent
upon the rate of amide hydrolysis and not independent
of it as is the case at lower acid concentrations
(16,20,43).
reactions.
This is an example of consecutive
Since amide hydrolysis is slower than the
nitrile hydrolysis, amide concentration builds up in
the solution and if the analysis is carried out by
determination of ammonia, it follows that a value will
be determined which is dependent on the rate of
formation of ammonia and not on the rate of hydrolysis
of nitrile to amide.
The induction period can be
eliminated by the high temperature analytical procedure,
as previously outlined.
In table I are shown some runs under
identical conditions which illustrate the induction
period obtained by analysis at 0 0 and its elimination
on analysis at 70 0 •
Figure A (page 27) shows pictorially
25. the relation tabled below and also the variation in
concentration of amide with time.
TABLE I
7.68 N
Hydrobromic acid at 68.1 o C.
Analysis carried out at:
70°C.
x 10 (moles/mole/litre/hour) k x 10
k
2.64
2.72
2.57
2.73
2.56
2.58
2.64
Ave.
13.0 N
1.27
1.63
1.81
2.29
2.39
2.69
Nitric acid at 60.5 0 C.
Analysis carried out at:
k
Ave.
x 10 (moles/mole/litre/hour) k x 10
1.54
1.51
1.50
1.53
1.49
1.47
1.51
0.057
0.176
0.201
0.245
0.248
0.276
0.270
26.
13.0 N
Nitric acid at 80.5 0 C.
Analysis carried out at:
k x 10 (moles/mole/litre/hour) k x 10
Ave.
8.62
9.39
9.10
8.74
8.72
9.46
9.00
10.6 N
1.20
1.82
1.78
2.06
2.30
2.25
2.37
Sulphuric acid at 8U.5 0 C.
Analysis carried out at:
o
O.C.
k x 10 (mo1es/mole/1itre/hour) k x 10
Ave.
2.05
2.07
2.06
2.02
2.12
2.09
2.11
1.98
2.05
1.71
1.77
1.88
1.93
1.98
1.95
2.02
27.
FIGURE A
Induction periods
••
4.
"
--
-~
_.
7
13.0
HNO~
'-'
at 60 . 5°
Sho~ing
v~en
10.6 N
variation of reaction lith time
analysis is carried out at 70° and at aOC.
28.
It is readily apparent from tho above tables
that the induction period is more marked with the
higher concentrations of acids.
With increasing
concentration of acid the rate of hydrolysis of nitrile
to amide increases, while the rate of amide hydrolysis
reaches a maximum and decreases again after an acid
concentration of about four normal has been passed.
Thus after a certain concentration of acid has been
reached the velocity of the first step relative to the
second will increase as acid concentration is increased
and consequently the induction period will probably
increase.
The unimolecular velocity constants recorded
in tables II--XI, which follow, are the mean of at
least five individual determinations, corresponding
usually to approximately 10% -- 80% reaction.
Some typical runs are given in table II.
It is seen that in general there is good agreement
among the individual k values for a given run; the
average deviation from the mean was usually less than
3%.
Those starred (A) values of k which deviate more
than 5% from the mean are not included in the average.
In this investigation no correction has been
ap~lied
for variation in collision number or for
changes in concentration of the solution with temperature.
29. The error will however be negligible.
TABLE II (Typical runs)
1.00 N
Hydrobromic acid at l13.S o C.
%reaction
k (uni) x 10
moles/mole/litre/hour
29.4
37.1
37.3
1.20
51.7
59.2
1.14
1.16
0.97 1l
1.16
1.19 65.3
Ave.
13.0 N
1.17 Nitric acid at 48.7 oC.
%reaction
k (uni) x 10
mo1es/mo1e/1itre/hour
20.9
0.496
0.487
30.1
36.9
0.495
0.490
0.485
43.3
65.7
65.7
Ave.
0.479 0.489 30. 18.9 N
Sulphuric acid at 60.5 0 C.
%reaction
k
(uni) x
10
mo1es/mo1e/1itre/hour
3.91
12.5
25.6
51.4
56.6
60.6
63.6
70.3
79.8
4.11
4.37
4.18
4.20
3.90
3.81
4.22 4.09 Ave.
TABLE III
1.00 N Hydrobromic acid
Temp&rature
C.
86.7
100.4
113.6
124.8
k
(uni) x
10
moles/mole/litre/hour
0.093
0.363
1.17
2.90
31.
TABLE IV
4.00 N
Hydrobromic acid
Temperature
k (urd) x 10
moles/mole/litre/hour
OCt
78.7
89.6
100.7
107.4
109.6
0.458
1.29
3.83
7.40
8.45
TABLE V
7.68 N
Hydrobromic acid
Temperature
°e.
k (uni) x 10
moles/mole/litre/hour
48.7
0.369
60.5
68.1
1.28
79.7
8.01
2.68
32.
TABLE VI
1.00 N
Nitric acid
Temperature
°e.
k
(uni) x
10
mOles/mo1e/litre/hour
100.4
113.6
0.0899
0.344
1.19
124.8
3.13
86.5
TABLE VII
4.00 N
Nitric acid
Temperature
°e.
78.7
89.6
100.4
107.4
k
(uni) x
10
moles/mo1e/1itre/hour
0.466
1.26
3.94
7.38
33.
TABLE VIII
13.0 N
Nitric acid
Temperature
°C.
k (uni) x 10
mo1es/mo1e/1itre/hour
48.7
0.489
1.51
2.80
60.5
68.1
80.5
9.00
TABLE IX
2.00
N
Sulphuric acid
Temperature
°C.
86.5
100.4
113.6
125.3
k (uni) x 10
mo1es/mo1e/1itre/hour
0.099
0.412
1.39
3.52
34. TABLE X
10.6 N
Sulphuric acid
Temperature
°e.
5B.5
k (uni) x 10
moles/mole/litre/hour
0.205
0.709
2.05
6.42
69.9
BO.6
92.4
TABLE XI
18.9 N
Sulphuric acid
Temperature
°e.
4B.7
60.5
64.6
6B.1
k (uni) x 10
mo1es/mo1e/litre/hour
1.25
4.09
6.68
9.43
35.
In table XII below, are given the values of
tne rate constant k at a temperature of 85°C.
as
determined from figures B, C, and D, pages 36, 37, and
38, for the various acids.
This table converts all
k values to standard conditions for comparison of the
different acids at various acid concentrations.
TABLE XII
k (uni) x 10 (at 85 0 C.)
moles/mole/litre/hour
Acid 1.00 N Hydrobromic II
4.00 N II
7.68 N 0.0813
0.832
.0
1.00 N Nitric fI
4.00 N
fI
13.0 N 0.0746
0.851 12.0
2.00 N Sulphuric n
10.6 N. It
18.9 N 0.0861
3.16
47.5
From the above table it is seen that for
hydrobromic and nitric acids an increase in
concentration of acid from one to four normal causes
the rate constant to increase by a factor of ten,
while at greater concentration of acid the change of
rate is still more marked.
A similar rapid change
of rate constant with acid concentration is observable
for sulphuric acid.
These changes of rate with
Z6 . FIGURE B 1 . 500
...
'f.
~
~
..
""
1.000
~
2 .000
0. 002 7
0 .00 29
0 . 0031
37. FIGURE 1 . 50
2. 50
2" . 00
L.....- _ _ _ _ _
O .OO~,
~-'-
0.0027
0_ ­
0.OC 29
1
T
0.003 1
38.
FIGURE
\
"l!.OO
'"
O.~O~02~'--------~
O.~OO~2=7--~----~O.~O~
O.='---------O~.~OO~'~1
/
t
/
39. concentration are clearly shown in figure E page 43.
From the data given in tables III to XI,
figures B, C, and D have been plotted.
In each
of the three figures it is seen that there are three
lines each one corresponding to one concentration of
that particular acid.
For example in figure B page 36,
there are three lines corresponding to the
concentrations 1.0, 4.0, and 7.68 normal hydrobromic
acid.
From the slope of the lines the values of E
have been calculated in the usual way from the
Arrhenius equation.
These values and also values of
log A of the Arrhenius equation have been tabulated
in table XIII.
It is noted that there
is
a
decrease in the activation energywlth increase in the
concentration for all acids.
For purposes of compar­
ison the table includes data for the hydrolysis with
hydrochloric acid (33).
40. TABLE XIII
Acid
1.00 N
4.00 N
7.68 N
E
Hydrobromic
Il
"
1.00 N
4.00 N
13.0 N
Nitric
2.00 N
10.6 N
18.9 N
Sulphuric
1.00
4.00
8.4
10.0
Hydrochloric
N
N
N
N
If If tI f! f!
"
"
(cal)
log A
25,800
25,400
22,200
13.7
14.8
16.7
26,400
25,600
20,500
14.U
14.5
12.6
26,300
24,300
23,4.00
14.3
10.0
25,600
24,800
19,800
19,700
13.6
14.1
12.5
12.8
14.0
41. DISCUSSION OF RESULTS
Consecutive reactions are those reactions
'Nhich take ple.ce in a series of intermedi8te stages.
The hydrolysis of nitriles in strong acid solution
affords a good example of consecutive reactions.
A
period of induction is characteristic of chemical
reactions which take place in a series of intermediate
stages (28).
An induction period has been found in
the hydrochloric acid catalysed hydrolysis of various
nitriles (33).
This induction period has dlso been
found in the present investigation in the case of
hydrobromic, nitric, and sulphuric acid at higher
concentrations.
It will be noted (table I) that in the case
of 7.68 normal hydrobromic acid and 10.6 normal
sulphuric acid that the rate constcmt, when analysis
is carl'ied out at
aOe, gradually increases until
when the reaction has gone approximately 80%
to
completion the rate constflnt is the same as the r8te
constant v:hen determined by anCtlysis at 7rPC.
In the
case of 18.9 normal sulphuric acid and 13.0 normal
nitric acid the rate constant when found by cnalysis
at OoC. does not approach that found by 70 0 analysis.
This is due to the fact that at these concentrations
of acid the rate of removal of amide is so slow
4 ",
~.
compared to its formation thEt the extrapol&ted values
of k for the nitrile do not approach its true value.
Figure A shows the variation of amide with time, fer
13 normal nitric acid and for 10.6 normal sulphuric
acid.
In the case of nitric acid the concentration
of amide does not fall to zero during the experimental
time, whereas with the relatively less concentrated
sulphuriC' acid the amide eventually reaches zero
concentration.
The catalytic efficiencies of the acids used
may te seen from table XII page 35 and figure
~
page 43.
They are approximately the same at the low concentrations.
Sulphuric acid usually functions as a monobasic acid
and its two normal solution may be compared with the
one normal solutions for the other acids.
Noll (20) have found e similar result.
Krieble and
The activation
energies at this concentration are also practically
the same.
At higher concentrations the individual
characteristics of each acid become more marked.
For
example it is noted that 7.68 normal hydrobromic acid
is similar in respect to catalytic effect to
normal nitric acid.
13.0
At large concentrations of acid
the difference in k values is very pronounced.
As
seen from table XII and figure E these acids are not as
efficient catalysts as is hydrochloric acid and the
43.
FIGURE E
2 .0
1.5
K
'.0
0.5
15
NORHAllTY
44. increase of rate with concentration is not as rapid.
The rate increases rapidly with increase in
acid concentration, especially at high acid concentration.
It is known that the rate of amide hydrolysis
in
cqncentrated acids is small and decreases with acid
concentration and if this is the case then, by choosing
the appropriate conditions we should be able to stop
the reaction of the hydrolysis of nitriles at the amide
stage.
This actually is the case
~hen
96% sulphuric
acid is used (13).
It has been proposed that in dilute solution
there is a direct proportionality betv'een the hydrogen
ion concentration and the velocity constant (16).
On
the other hand, in more concentrated solution, such a
relation no longer holds.
It has been reported (31)
that the reaction velocity at higher acid concentrations
increases considerably more rapidly than the hydrogen
ion concentration.
For example, in the present
investigation, an increase from 4.0 to 7.68 normal in
the case of hydrobromic acid causes an increase in the
reaction rate constant of more than fifteen times.
There must obviously be some other catalytic effect
than that due to the hydrogen ion.
It has been
postulated that the rate follows closely the square of
the mean ion activity (31).
only for hydrocyanic acid.
This relation was found
45. Krieble and Noll (20) have found thBt for
the hydrolysis of propionitrile, hydrochloric acid is
a better catalyst than sulphuric acid.
All of their
determinations have been made at only one temperature.
Krieble and IicNally (19) using hydrocyanic acid have
found that with four normal hydrochloric
and
hydrobromic acids the rate of hydrolysis with
hydrochloric acid is much greater than with hydrobromic
acid.
Peiker (31) has found with hydrocyanic acid
that hydrochloric acid is a better catalyst than
sulphuric acid, that v!ith five normal acids the ratio
is twenty-three to two.
The results of the present
investigation are in accord "lith these results.
A comparison of the results of the present
investigation and those of hydrochloric acid obtained
in this laboratory (33) show that there is a marked
difference between the two acids and that hydrochloric
acid is the more active.
compar~tive
A table of k values for
purposes of all acids follows.
Temperatures for sulphuric, nitriC, and
hydrobromic acids are all 85°C.
The velocity constants
are given in moles/mole/litre/hour x 10
46. ACID
NORMALITY
1.0
2.0
4.0
10.0
13.0
3.16
0.0861
H2 SO 4
T. °C.
10.6
18.9
47.5
85.0
HNO 3
0.0746
0.851
12.~2
85.0
HBr
0.0813
0.832
12.02
85.0
HCl
0.0761
83.4
79.7
0.571
"
tT
9.72
83.5
10.6
II
76.1
It is noted that after taking into account
the temperature coefficient, hydrochloric acid is a much
better catalyst than the other three acids.
The change of slope of the lines in figures
B, C, and D is in the direction of decrease in the value
of the activation energy with increase in the acid
concentration.
These changes are in agreement lliith
those found previously (33) for hydrochloric acid,
namely that an increase in the concentration of acid
catalyst lowers the activation energy in nitrile
hydrolysis reactions.
The nature of the change in the
activcJ.tion energy is not certain, but the magnitude
however is seen to be specific for each acid at
comparable acid concentraticns.
47.
From the literature have been collected a
number of examples of catalytic reactions which undergo
variations in E with concentration of catalyst.
They
include the hydrolysis of sucrose (25), the hydrolysis
of ethylal (26), the oxidation of organic acids by
chromic acid--sulphuric acid mixtures (36,38), the
decomposition of chromic acid in sulphuric acid solution
(37), and the hydrolysis of alpha propene chlorohydrin
(15).
In the hydrolysis of ethyl acetate E
is
independent of the addition of neutral salt (44).
In a
study of some ionic reactions, a variaticn of E with
neutral salt concentration has been reported (17,18).
It is noted that these reacticns are of the hydrolysis
or proton transfer type.
The above do not include a
number of recent studies in dilute soluticns on the
dependence of E on dielectric constant and ionic
strength.
It is almost universally accepted that the
energy of activation is constant with temperature, in
evaluating the Arrhenius equation.
In reactions
between ions it might happen that E would not be
independent of the temperature (23).
Hydrolysis
reactions are of a type which frequently exhibit
variation of E with temperature.
Leininger and Kilpatrick (25,26) have postulated
a general variation of activation energy with acid
concentration in strong acid catalysis.
obtained here support this view.
The results
They attribute the
change to a variation of the heat content of an
intermediate activated complex with acid concentration.
It should be kept in mind however that reactions in
concentrated acid solutions are in general more complex
and not as well understood as reactions in dilute
solutions.
The influence of such factors as degree of
hydration, dielectric constant, polarity of the medium,
et cetera, en the variation of E is difficult to assess.
It would appear that an unambiguous interpretation of
the phenomenon is difficult.
Interatomic attraction (24) might modify the
reactions by having an effect on the character and
number of collisions.
solution shows that
~
La Mer's evidence in dilute
of ionic reactions is dependent
on both concentration and temperature.
In the present
investigation E has been found to vary a great deal
with concentration, however no change has been noted with
temperature.
It may be that in the present investigation
the measurements were not accurate enough to detect a
temperature effect since it would in all probability
be small.
When E is plotted as ordinate and log k
as
abscissa as in figure F page 49, it is found that, in
general, there are a series of curves in which E decreases
49. FIGURE F .:16,000
E
24,000
22,000
.:1. 000
1.000
108
Leg nd
I<
01
1 normal acid
04
4 nOl'mal Cl cid
(-)8
8 normal acid
(31::::--12 normal C'cid
0 . 000
50. as k is increased.
Plotting the logarithm of k
against E would give a linear relation if E were the
sole determinant of k.
No such relation is found,
hence factors other than E must vary as the concentration
is va,ried.
These changes in E with change in acid
concentration are related to changes in the Arrhenius
constant A.
This is to be expected from the general
existence of a proximity effect (10) that usually
appears in reactions of aliphatic compounds.
With
hydrobromic acid the variation in E from 1.0 to 4.0
normal is small.
On the other hand with a change from
4.0 to 7.68 normal the drop in E is much more marked.
With nitric acid the changes in E are slightly more
pronounced.
18.9 normal sulphuric acid does not give
as low a value for E as would be expected from a comparison
of the values of the other acids; however if sulphuric
acid is considered as monobasic then the change of E
with normality more nearly approaches that of the other
acids,
In figure G page 52 are plotted the values
of E against normality as given in table XIII.
It is
seen that the rate of change of E with normality is
much more rapid in the case of hydrobromic acid than
in the case of nitric acid and similarly that the
change with nitric is more rapid than the change with
sulphuric acid.
In the case of hydrobromic and nitric
acids it is seen that as the normality is increased
51.
the slope of the line changes in a direction
indicating an increase in the rate of decrease of E.
Conversely in the case of sulphuric acid the rate of
decrease of E with increase in normality becomes less
with increase in normality.
52. FIGUEE G 26, 000
24.000
22 .000
20 .000 L-____________~~____________~--------------7.------------~
'1 0
NORMAL IT Y
I~
53. CONCLUSION
A period of induction has been observed for
the three acids in concentrations above four normal.
This induction period has been sho?m to be due to
consecutive reactions.
In general it is found that rise in temperature
results in an increase in the rate constant.
The
change is approximately two and one half to three times
for each ten degree change in temperature.
No change
of E with temperature change has been noted.
There is a marked variation of reaction rate
with acid concentration.
The rate constants for the
four acids: hydrochloric, hydrobromic, nitric, and
sulphuric are approximately the same at low concentrations,
but as the concentration of acid is increased the
specific catalytic effects of the different acids
beco~e
more pronounced and, at high concentration of
acid, the velocity constants for the different acids at
like concentrations are quite dissimilar.
Hydrochloric
acid is the most active, with hydrobromic next,
followed by nitric and then sulphuric acids.
The energy
of activation has been found to vary with the aCid.
There is a general decrease in the activation energy
with increase in concentration of acid.
54.
Considering that an investigation of
necessarily limited scope has been made, the number of
different aspects of solution kinetics which are
illustrated is quite remarkable.
Further study of
the nitrile hydrolysis should prove fruitful.
55.
BIBLIOGRAPHY
49, 345 (1913).
(1)
Acree, Arn.Chem.J.
(2)
Acree and Johnson, Am.Chem.J.
(3)
Acree and Johnson,
(4)
Blanksma, . Rec.P.B. 21, 366(1902), 22, 290(1903).
(5)
Bronsted,
Chern. Rev.
(6)
Bronsted,
Tr. Far. Soc. 24,
ibid.
37, 410 (1907).
38, 258 (1907).
5, 231 (1928).
(7). Bronsted and De1banco,
144, 248 (1925).
630
(1928) .
Z. anorg a11gem. Chern.
(8)
Dawson and Powis,
J. Chern. Soc. 103, 2135 (1913).
(9)
Dawson and Reimann, J.Chem.Soc.
107~
1426 (1915)
(10)
Dippy, Evans, Gordon, Lewis, Watson,
1421 (1937).
(11)
Folin and Farmer,
(12)
Grube and Motz,
(13)
Karrer, Organic Chemistry,
Company, (1938).
(14)
Karve and Gharpure,
Part 3, 139 (1939).
(15)
Kedrinski and Merson, Trans. Expt1. Research
Lab Khemgas, Materials on Cracking and Chemical
Treatment of Cracking Products, U.S.S.R.,
3, 311 (1936).
(16)
Kilpi, Z. physik Chern. 86, 641 (1914).
(17)
Kiss and Bossanyi, Rec. trav. Chim. 53, 903(1934).
(18)
Kiss, Magyr Chern. Folyoiat, 64, 13 (1938).
(19)
Krieble and McNally, J.Arn.Chem.Soc. 51, 3368(1929).
J. BioI. Chern.
J.Chem.Soc.
II, 493 (1912).
Z. physik Chern. 118, 145 (1925).
Nordemann Publishing
J. Univ. Bombay,
Vol. VIII,
56 •.
(20) Krieble and Noll, J.Am.Chern.Soc. 61, 560 (1939).
(21) Krieble and Peiker, J.Am.Chern.Soc. 55, 2326(1933).
(22) Krieble and Walker, J. Chern. Soc. 85, 1369 (1909).
(23) La Mer,
J. Chem. Phys.
(24) La Mer and Kamner,
1., 289 (1933).
J.Am.Chem.Soc. 57, 2662 (1935).
(25) Leininger and Kilpatrick, J.Am.Chem.Soc. 60,2891(1938).
(26) Leininger and Kilpatrick,
ibid.
61, 2510 (1939).
(27) Lowry, Class Book of Physical Chemistry, MacMillan
Company, (1929).
(28) Mellor,
J. Chern. Soc.
81, 1280 (1902).
(29) Mellor, Chemical Statics and Dynamics,
Green & Company (1914).
(30) Ostwald,
(31) Peiker,
J. prakt. Chern.
Longmans
(1) 27, 1 (1883).
Thesis McGill University, Montreal.
(32) Rabinovitch and Winkler, private communication on
the hydrolysis of amides in concentrated hydrochloric
acid solutions. (1941)
(33) Rabinovitch and Winkler, private communication on
the pydrolysis of nitriles in hydrochloric acid
solutions (1941).
(34) Reitz,
Z. physik Chern. (1938).
(35) Schiffer and Kohastamm, Proc. Amstend. Akud
13, 789 (1922).
(36) Sneth1age,
Rec. trav. Chim. 56, 873 (1937).
(37) Snethlage,
ibid.
57, 1341 (1938).
(38) Sneth1age,
ibid.
59, 111 (1940).
(39) Taylor,
(40) Taylor,
Medd. Nobelinst.
ibid.
2, Nr.
2, Nr. 35 (1913), Nr. 37 (1913).
(41) Taylor,
Z. Elektrochern.
(42) Taylor,
Medd. Nobelinst,
(43) Taylor,
J. Chern. Soc. 2741 (1930).
(LlLll Vlv~7Alowska.
34 (1913).
20, 201 (1914).
3, Nr.
Roczniki Chern.
1 (1915).
14, 1118 (1934).
PART I I SECRET (War work)
Physical Chemistry
1-1. Sc.
James Douglas McLean
Part II
W&r Res82rch
Some Reactions of SFs
The
deconpositi~n
attempt to prepare SzF1o.
of SFa has been studied in an
SFa and HzS ,.,ere passed through
SFa was pas sed through 8.n electric arc,
separately and concurrently with each of the following gases,
Hz S, Nz, E'tnd Oz.
SFa
WEcS
passed oyer molten potassium. , No
SzF10 was identified in the products of the above reactions.
McGill University, Montree.l, Canadft.
September, 1941. PART II
WAH RESEARCH REPORT
REPORT ON THE I}TVESTIGATION OF THE
C.E. 3
REACTIO~:
I~"'"TRODtJCTION
Since the great proportion of fluorine which is
passed over sulphur reacts to give sulphur hexafluoride,
rather than disulphur decafluoride, it was considered of
importance to investigate the decomposition of sulphur
hexafluoride.
The following investigation has therefore
been carried out with the idea that it might be possible
to break up the SF6 molecule in such a way that it .might
recombine to form S2FIO.
out previously.
SODle of the tests had been carried
However, they have.been repeated with special
emphasis being placed on the discovery of S2FIO, which if
formed in small quantities might have been overlooked by
previous investigators.
SF6 sublimes at -63.8°C. and decomposes at red heat.
S2FIO boils at 29°C. and melts at -92°C.
It
decomposes at 215°C.
Apparatus and Technique
A flow system was used in the determinations and
gaseous products were collected in a liquid air trap.
Advantage was taken of the difference in vapor pressures of
the possible components in order to separate them.
The gaseous
products and residues of the reaction which had been collected
in the liquid air trap were distilled from a trap surrounded
with a dry ice--acetone mixture to a trap surrounded by
-2­
liquid air, arter the system had been evacuated.
of
~vo
At the end
or three hours all SF6 would have eone into the liquid
air trap and any S2FIO, if present, would remain in the trap
surrounded by the dry ice--acetone mixture.
No attempt was made to find gaseous products or
residues other than SF6' H2S, and S2FlO. No attempt was
made to identify positively solid residues or products.
RESULTS
Reaction of SFS and H2S.
SF
and H S were passed through a furnace at 780 0 C.
Some decomposition took place as shown by a deposit of sulphur
in the line leading from the furnace to a liquid air trap,
and also in the trap.
Undecomposed SF6 and H2S were collected
in the liquid air trap.
No SZFIO was found.
The purpose of the HZS was to furnish extra sulphur
atoms for the fluorine that might be formed in the decomposition.
HZS decomposes at 300"C.
Since this temperature
is lower than that at which th test was performed it appears
that the sulphur deposit was chiefly caused by the decomposition
of HZS.
Due to the fact that at this time Mungen and Hugill
found that SzFIO decomposed at 215°C. no further tests of
this kind were made, as it is very unlikely that SZFlO could
be formed by this means.
-3­
Passage of SF6 through an electric arc.
An arc was established between platinum electrodes
in a glass cell, current being supplied through a transformer.
The glass cell was prevented from overheating by immersion
in a
beal~er
of vlater.
When SF was passed through the arc decomposition
took place.
This was noted by the formation of fumes in the
cell and the formation of a black deposit on the walls of
the cell.
After a short tirlle the arc would go out.
This
appeared to be caused by conduction through the black deposit
on the walls of the cell.
To overcome this various sizes of
cell and length of arc were used.
In general it was found
that the shorter the arc and the larger the cell the more
time was required for conduction on the walls to'take place.
When the cell was opened, parti cles of yellow sulphur were
noted on the black de:posit.
The black deposit is believed
to be black sulphur.
In order to slow dovm the rate of decomposition to
sulphur and hence the formation of the deposit on the walls
of the cell, the SF6 iNas dilutecl with other gases.
(a)
SF6 passed through an arc decomposed rapidly but
not completely.
deposit.
(b
t
)
The walls of the cell contained a black
The liquid air trap appeared to contain only SF6.
SF6 and H2S were passed through the arc in the
-4­
ratio of approximately 1:4.
rapid but not complete.
Decomposition was comparatively
Sulphur was found in the cell and
the line leading from the cell to the liq11id air trap as
well as in the trap itself.
the
li~uid
(c)
SF 6 and H2S were found in
air trap.
SF6 and N2 were passed through the arc in the ratio
of 1:4 and 1:6.
Decomposition, as shown by the formation of
the black deposit, was less ra:9id than in the above two cases.
The liquid air trap contained only SF 6 •
(d)
SF6 and 02 were passed through the arc in the ratio
of 1:4.
The rate of decomposition was of the sa'11e order as {cl.
The liquid air trap contained only SF6.
In all of the above cases no S,2F10 was found.
M. Berthelot (Ann Chim
Phys~,
(1900}) found that
SF6 is stable to the silent electrical discharge.
H. 11oiss8.n and p. Lebeau (Compt Rend 130, (1900))
found that SF6 is partially decomposed in the induction spark.
SF6
02 when strongly sparked yields a brown folcculent
solid.
They also found that SF6 and HBS react quite easily.
Passage of SF6 over molten potassi'llffi.
SF6 was passed over molten potassium. at 160°C. and
at atmospheric pressure.
indicating a reaction.
A film forr.1ed on the potassiura
It appeared to happen only when a
fIB-me aP:geared at the surface of the potassium for a few
seconds.
The liquid air trap contained on SF6.
nlis
-5­
was repeated at 170 0 0. and 20 cm. pressure.
The results were
the srune as above.
The S3lne type of experiment was repeated at 200 0 0.
and atmospheric :pressure.
was used.
A much larger surface of potassijm
A film for.Fled on the surface of the potassium
without visible flame and appeared to prevent further reaction.
The liquid air trap contained only SF6.
An apparatus was designed which provided for stirring
of the potassium at intervals.
SF6 was then passed through
the molten potassium at 80°0. and at atmospheric pressure at
a rate of approximately 30 cc. Imin.
Reaction
v.;1 th
potassiu.m
was evident as seen by the formation of a firm black mass and
occasional flarne.
The more active stages of the reaction as
evidenced by the flame were increased in frequency by rise in
temperature and rate of flow of SF6.
Stirring presented
unreacted surfaces of potassium to the SF 6 vapor.
At the end
of the run the residue consisted of a finn black mass containing
scattered :particles of potassium metal.
The liquid air trap
contained SF6 and a fraction of a drop of liquid which appeared
to be a solid at the temperature of a
d~J
and at a pressure of less than I cm.
The liquid turned milky
near room temperature.
ice--acetone mixture
At room temperature it was clear and
readily volatilized when the trap was warrJled with the hand.
There was not enough of this substance to make density or
molecular weight measurements.
A toxicity test was made.
-6­
If the substance were S2FlO the toxicity test would be
positive.
In this case the toxicity test was negative.
The last experiment was repeated using a larger
amount of potassium and passing SF6 through the potassium at
the S~le temperature and rate as before. The same firm black
mass v'ras noted in the reaction vessel at the end of the
reaction.
air trap.
SF 6 was the only product collected in the liquid