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THE REACTIONS OF METAL OXIDES AND CARBONATES
WITH FUSED ONIUM SALTS
by
THOMAS MARTIN SOUTHERN, B . S .
A THESIS
IN
CHEMISTRY
Submitted to the Graduate Faculty of Texas Technological College
In Partial Fulfillment of the Requirements for
The Degree of
MASTER OF SCIENCE
At ~
T3
No, /4^
4E6-W035
ACKNOWLEDOIENTS
I am deeply indebted to Dr. W. W. Wendlandt for
his help in directing this thesis, and to the R, A. Weloh
Foundation and the United States Air Force, Office of
Scientific Research, for their financial support.
ii
TABLE OP CONTENTS
Page
ACKNOWLEDGMENTS 11
LIST OF TABLES v
LIST OP FIGURES vl
INTRODUCTION 1
Chapter
I. INSTRUMENTATION AND SPECIAL TECHNIQUES. . . l|
Automatic Recording Thermobalance i. Differential Thermal Analysis Apparatus. . ij. Mass-SpectrometrIc-Gas Evolution Analysis. 5 Isothermal Kinetics Apparatus 6 Reaction Stolchiometry Studies 10 Magnetic Susceptibility Studies 11 Reflectance Studies « 11
II. REACTIONS OP FUSED ONIUM SALTS WITH METAL CARBONATES 13
Methylammonlum Chloride and Barium Carbonate 13
Dlmethylammonium Chloride and Barium Carbonate 16
Trimethylammonlum Chloride and Barium Carbonate 32
Methylammonlum Chloride and Strontium Carbonate , 32
Dlmethylammonium Chloride and Strontium Carbonate 5l
Trimethylammonlum Chloride and Strontium Carbonate 55
III. REACTIONS OP METAL OXIDES IN FUSED ONIUM SALTS 68
Methylammonlum Chloride and Zinc Oxide . . 68 Methylammonlum Chloride and Cadmium Oxide. 72 Methylammonlum Chloride and Magnesium
Oxide 75
111
Chapter p^g^
Methylammonlum Chloride and Cobalt(III) Oxide 80
IV. CONcliUSIONS , 88
LIST OP TABLES
Table Page
1. A. Reaction Stolchiometry Data for methylammonlum Chloride with Barium Carbonate 1?
B» Kinetic Data 1?
2. A. Reaction Stolchiometry Data for Dlmethylammonium Chloride with Barium Carbonate, . • , • . 26
B. Kinetic Data 26
3. A. Reaction Stolchiometry Data for Methylammonlum Chloride with Strontium Carbonate 38
B. Kinetic Data 38
I4.. A, Reaction Stolchiometry Data for Dlmethylammonium Chloride with Strontium Carbonate. 514-
B. Kinetic Data • . » ^ Sk
5« A. Reaction Stolchiometry Data for Trimethylammonlum Chloride with Strontium Carbonate 62
B. Kinetic Data 62
6. Reaction Stolchiometry Data for Methylammonlum Chloride with Zinc Oxide 71
7. Reaction Stolchiometry Data for Methylammonlum Chloride with Cadmium Oxide 75
8. Reaction Stolchiometry Data for Methylammonlum Chloride with Magnesium Oxide 80
9. A. Magnetic Data for Methylammonlum Chloride with Cobalt (III) Oxld 8l|.
B. Reaction Stolchiometry Data for Methylammonlum Chloride with Cobalt (III) Oxide . 81;
LIST OP FIGURES
Figure Page
1. Diagram of Isothermal Kinetics Apparatus. . . . 9
2. Instrumental Curves for the Reaction of Methylammonlum Chloride with Barium Carbonate . , . 15
3» Rate Curves for the Reaction of Methylammonlum Chloride with Barium Carbonate. . . . . . . . 19
l.. Arrhenius Curves for the Reaction of Methylammonlum Chloride with Barium Carbonate . . . 22
5. Instrumental Curves for the Reaction of Dlmethylammonium Chloride with Barium Carbonate • • • 2I4.
6. Rate Curves for the Reaction of Dlmethylammonium Chloride with Barium Carbonate . . . 28
7. Arrhenius Curves for the Reaction of Dlmethylammonium Chloride with Barium Carbonate . . . 31
8. Instrumental Curves for the Reaction of Trimethylammonlum Chloride with Barium Carbonate 3I4.
9. Instrumental Curves for the Reaction of Methyl
ammonlum Chloride with Strontium Carbonate. , 37
10. Typical pH vs. Time Plot Ij.1
11. Rate Curves for the Reaction of Methylammonlum Chloride with Strontium Carbonate l\}^
12. Rate Curves for the Reaction of Methylammonlum Chloride with Strontium Carbonate 1.6
13. Arrhenius Curves for Methylammonlum Chloride with Strontium Carbonate,Reaction 1 kQ
11 .. Arrhenius Curves for Methylammonlum Chloride with Strontium Carbonate, Reaction 2 50
vl
vll
Figure Page
15« Instrumental Curves for the Reaction of Dlmethylammonium Chloride with Strontium Carbonate 53
16. Rate Curves for the Reaction of Dlmethylammonium Chloride with Strontium Carbonate, . 5?
17. Arrhenius Curves for the Reaction of Dlmethylammonium Chloride with Strontium Carbonate, . 59
18. Instrumental Curves for the Reaction of Trimethylammonlum Chloride with Strontium Carbonate'. 61
19. Rate Curves for the Reaction of Trimethylammonlum Chloride with Strontium Carbonate. . 65
20. Arrhenius Curve for the Reaction of Methylammonlum Chloride with Strontium Carbonate, . 67
21. Instrumental Curves for the Reaction of Methylammonlum Chloride with Zinc Oxide 70
22. Instrumental Curves for the Reaction of Methylammonlum Chloride with Cadmium Oxide 7k-
23. Instrumental Curves for the Reaction of Methylammonlum Chloride with Magnesium Chloride , . 78
2l|.. Instrumental Curves for the Reaction of Methylammonlum Chloride with Cobalt(III) Oxide. , . 82
25» Reflectance Studies of the Reaction of Methylammonlum Chloride with Cobalt(III) Oxide . . 87
INTRODUCTION
It has been known for some time that onluro salts
(ammonium and substituted ammonium salts) act as strong
acids in the fused state. The solution of several metal
oxides in fused pyrldlnlum chloride was first observed by
Long and Audrleth(l), They found that copper(II) oxide,
barium oxide, magnesium oxide, lead oxide, etc, dissolved
In the fused salt yielding the corresponding metal chloride,
pyridine, and water. Since no reaction stolchiometry data
were given, it was not known whether the reactions were quan
titative or whether coordination accompanied the reaction.
Audrieth, Long, and Edwards(3) observed that metals
such as aluminum, zinc, copper, etc, also dissolved in
fused pyridiniura chloride to give the corresponding metal
chloride and hydrogen. The metal oxides, as well as the metal
chloride salts, were found to dissolve readily In the fused
salt. Again, no reaction stolchiometry data were given for
the reaction.
Audrieth and Schmidt(2) observed the solution of
some metal oxides, metal carbonates, and pure metals in
fused ammonium nitrate. They found that metal oxides such
as copper(II) oxide, magnesium oxide, lead oxide, calcium
oxide, nickel oxide, etc., were soluble whereas aluminum oxide
chroralum(III) oxide, Iron(III) oxide, tin(II) oxide, and
the more acidic oxides did not dissolve. It was found, how
ever, that the nitrate salts were soluble. It was also
found that the metal carbonates were soluble as well as the
metals above hydrogen in the electromotive series. Again
no reaction stolchiometry data were presented.
In another paper Schmidt and Audriethdi.) reported
some high temperature acid-base reactions in fused ammonium
chloride. They found that iron(III) oxide, magnesium oxide,
lead(II) oxide, lead(IV) oxide, cadmium oxide, etc, reacted
and dissolved in an ammonium chloride melt to give the metal
chloride, ammonia, and water. They also found that calcium
oxide and barium oxide reacted slowly, while nickel(II) oxide,
copper(II) oxide, and vanadium(V) oxide gave products of
Indeterminate composition. Zinc oxide reacted forming a
product that could be distilled at red heat and was presumed
to be Zn(NH^)Clp. The carbonates of these metals also re
acted in the ammonium chloride melt; however, they published
no analytical data to support the results they gave.
Scott and Coe(5) observed similar results using a melt
of pyrldlnlum chloride. They studied the displacement of
metals from their chloride salts by other metals. For ex
ample, gold was displaced by mercury, bismuth, and antimony;
arsenic by mercury and bismuth; mercury by antimony; and bis
muth by antimony and silver. These displacement reactions
occurred in a yield of greater than 95^ of the desired metal.
They also found that mercury was displaced by bismuth and
3
silver; antimony by silver with yields of the desired metal
of 50^ to 95^. It was mentioned that this would be a good
way to purify these metals.
Since most of the above studies were only of a pre
liminary nature, a more extensive investigation was under
taken in this laboratory, Methylammonlum, dlmethylammonium,
and trimethylammonlum chlorides were chosen as the reacting
onium salts since the free amine evolved in the acid-base
reaction in each case would be a gas and would make analy
sis of the systems simpler. These reactions were studied
using the techniques of differential thermal analysis (DTA),
thermogravimetrlc analysis (TGA), gas evolution analysis
(GEA), and mass spectrometric analysis (MSA), as well as
by conventional chemical methods.
CHAPTER I
INSTRUMENTATION AND SPECIAL TECHNIQUES
Automatic Recording: Thermobalance
The thermobalance employed in these studies has pre
viously been described by Wendlandt(6). The instrument
consisted of a beam type analytical balance equipped with
a Plsher Recording Balance Accessory (Pisher Scientific
Co., 711 Forbes Ave,, Pittsburgh, Pa.), It had a sample
mass capacity up to 70 mg and was programmed to increase
the temperature of the furnace at a linear rate of about
8°C per minute.
Differential Thermal Analysis Arparatus
The differential thermal analysis (DTA) Instrument
employed in this work was built in this laboratory. It
consisted of a cylindrical metal block into which two 3
mm holes were drilled into the top and a 6 mm hole through
the center of the block, A fifty watt electric heater car
tridge was placed in the center hole to heat the block and
was programmed to increase the temperature of the block at
a linear rate by a temperature programmer similar to the
one described by Wendlandt(7). The sample holders were 3
mm outside diameter (o.d.) Pyrex glass tubing, 5 cm long
and sealed at one end. About 30 mg of sample was then
placed in the sample holder, and a Chromel-Alumel thermo
couple was inserted in the open end of the tubing so that
the thermocouple Junction was in intimate contact with the
sample. The entire sample holder was then placed In one
of the 3 tnm holes drilled into the heating block. In the
other 3 "i hole was placed another sample holder-thermocouple
combination with aluminum oxide replacing the sample. The
difference in the outputs of the two thermocouples was fed
into a microvolt D.C, amplifier and the amplified signal
was fed into the "Y" axis of an X-Y recorder. The tempera
ture of the block was taken as the temperature measured by
the aluminum oxide reference thermocouple and was fed into
the "X" axis of the same X-Y recorder,
A glass bell Jar covered the entire heating block
to reduce air convection currents around it. The heating
block was mounted on an aluminum base by means of two ce
ramic supports. All connections were made through the
aluminum base.
Mass Spectrometric-Gas Evolution Analysis
The apparatus employed in this work for simultaneous
mass spectrometric analysis (MSA) and gas evolution analy
sis (GEA) has been previously described(8), Sample sizes
ranged in mass from 100 mg to 150 mg and were heated at a
linearly increasing rate of approximately 8° to lO^C per
minute.
Isothermal Kinetics Apparatus
The apparatus used in determining the kinetics of the
various reactions is shown in Figure 1. Nitrogen was used
as the inert carrier gas to flush the reaction chamber.
The gas entered the system through flow meter, PL, which
carefully measured its flow through the system. The re
action chamber consisted of a Pyrex glass tube, I4. cm in diam
eter and 38 cm long, with ground glass Joints at both ends,
A 10 cm long Nlchrome resistance wire wound furnace, PU, was
built around one end of the tube. An aluminum block, A13,
was placed inside the Pyrex tube and inside the furnace to
act aa an infinite heat sink to retard drastic temperature
changes in the furnace when the sample and cup were intro
duced into the furnace. The aluminum block had a 13 mm
square groove cut in its bottom to allow the sample cup, S,
to slide in, A 3 mm hole was drilled in one end of the block
to permit insertion of a Chromel-Alumel thermocouple. The
temperature of the furnace was then continuously detected
and recorded against time on recorder R2. The temperature
of the furnace was controlled using a temperature controller,
CI, similar to the one described by Wendlandt(7).
A 500 to 700 mg mass of sample was placed into a
Coors 13-G porcelain boat, S, and was pushed into and re
moved from the furnace by a magnet attached to the outside
of the boat. Another magnet on the outside of the Pyrex
tube could then be used to move the boat to the desired
position in the furnace. The flowing nitrogen gas stream
carried the gaseous products from the furnace through a
glass fritted bubbler, BBL, into a beaker containing a
stirred 0.5N hydrochloric acid solution. The pH of the
solution was continuously measured by a glass electrode and
standard saturated calomel reference electrode, EL, and re
corded using a Heath recording pH meter, Rl. Hence, the prog<
ress of the reaction was followed by monitoring the change
in pH of the acid solution as a function of time. The
kinetic data could then be calculated from this plot. All
rate constants and activation energies obtained were cal
culated using a least squares program with an IBM 1620
Model 2 computer,
A possible error in this determination of the free
amine is the hydrolysis of the alkylammonium chloride. The
hydrogen ion concentration from the acid was 1 x 10"^ mmoles
per ml and the hydrogen ion concentration from the hydroly
sis of the methylammonlum chloride was calculated to be 6.76
X lO" mmoles per ml. It is seen that the hydrogen ion con
centration from the hydrolysis of the methylammonlum chloride
is very small compared to the hydrogen ion concentration
from the acid remaining in solution. Therefore, for all
practical purposes, the hydrogen ion concentration contri
buted by the hydrolysis reaction could be ignored and would
not affect the pH of the solution appreciably.
The same problem arises with dlmethylammonium chloride
8
Figure 1
Diagram of Isothermal Kinetics Apparatus
CM
10
and trimethylammonlum chloride and their hydrolysis reac
tions. The hydrolysis constant for dlmethylammonium chloride,
I.9I4. X 10" , is somewhat smaller than that of methylammon
lum chloride; therefore, it will contribute almost the same
amount to the final hydrogen ion concentration. In the case
of trimethylammonlum chloride, the hydrolysis constant should
be in the same order of magnitude as methylammonlum chloride
and dlmethylammonium chloride. It is seen then that hydroly-
sis during the kinetic studies does not play an appreciable
part.
The concentration of amine liberated may then be cal
culated by the difference in hydrogen ion concentration be
tween the initial pH and the pH at time t.
Reaction Stolchiometry Studies
The same apparatus as described in the proceeding
section was used in the reaction stolchiometry studies ex
cept that the furnace temperature was programmed to Increase
at a linear rate of 8°C per minute. The evolved amine was
collected in a l\.% boric acid solution and titrated with a
standard solution of hydrochloric acid using bromocresol
green as the end-point indicator. After completion of the
reaction, the residue in the boat was dissolved in water
and the chloride content determined by the Mohr method.
Carbon dioxide was determined by absorption in a "U"-tube
filled with Ascarit^ after the amine and water in the effluent
gas had been removed by scrubbing with concentrated sulfuric
11
acid.
Magnetic Susceptibility Studies
The apparatus used in the magnetic susceptibility
studies has previously been de3cribed(9). The instrument
records simultaneously the mass-change curve and the mag
netic susceptibility of the sample, employing the Faraday
method, from room temperature to 500®C.
The mass magnetic susceptibility was calculated
using the formula:
A^u ~ A'^c ^s A/y -Ar
^ A"s - A^c ^u
where X ^ " ^ mass magnetic susceptibility of the unknown,
*^^ is the magnetic susceptibility of the standard (32.3 x
10"^) in c.g.s, units for (NHK)2Pe(S0K)2 6H2O, ^m^ is the
mass change for the unknown, A'^s ^ ^ .e mass change for
the standard, and A^^c ^ ^ ® mass change for the empty
cup. The magnetic moment may also be calculated using
the formula:
where yCC is the magnetic moment of the unknown sample,
MWy is the molecular weight of the unknown sample, f^ is the
molar magnetic susceptibility of the unknown sample, and T
is the absolute temperature at which the reading was made.
Reflectance Studies
The reflectance spectra of the compounds were obtained
12
using a Beckman Model DK-2A spectroreflectometer. Samples
were placed in a 2.524- cm diameter circular groove contained
in a 5*1 cm square aluminum plate and covered with a 5*0
cm square piece of Pyrex glass. Magnesium oxide was used
aa the reference substance.
CHAPTER II
REACTIONS OP FUSED ONIUM SALTS WITH METAL CARBONATES
Methylammonlum Chloride and Barium Carbonate
The mass-loss (TGA), differential thermal analysis
(DTA), gas evolution analysis (GEA), and mass spectrometric
analysis (MSA) curves are shown in Figure 2, curves A through
D, respectively.
The mass-loss curve showed that the reaction between
methylammonlum chloride and barium carbonate occurred at
200°C to 230°C with a theoretical mass-loss of 21.77/^ of
the original sample mass as calculated from the acid-base
reaction:
2MeNH3Cl -»- BaCO^ »-2MeNH2 + BaClg +H2O + CO2 (1)
where Me = CH- , The actual mass-loss obtained from the
mass-loss curve was 21.7%. The DTA curve showed but one
endothermic peak with a AT ^ ^ ^ of 2l5°C indicating that any
other reactions which may be taking place are occurring sim
ultaneously with the acid-base reaction.
The GEA curve also showed but one peak with a T^^^ max
of 2lj.5 C. The difference in temperatures of the MSA-GEA and
DTA peaks and the mass-loss curves was due to the different
13
11
Figure 2
Instrumental Curves for the Reaction of Methylammonlum ;oride with Barium Carbonate
(A), TGA (B). DTA (C). GEA (D). MSA
if)
o
H 5
10 mg
15
®
I -<3
3 -J O >
CO <
I mv
©
CH3
CO2
NH.
HgO
C/)
LU
UJ
®
100 200 TEMPERATURE ( C)
3 0 0
16
heating rates of the two instruments.
Three products were found by MSA-GEA analysis to be
carbon dioxide, water, and methylamlne. At no time during
the reaction was chlorine or hydrogen chloride observed in
the products.
The results of the stolchiometry studies are shown in
Table lA. They are seen to be in excellent agreement with
the acid-base reaction as shown in Equation (1).
The results of the kinetics investigation of the re
action of methylammonlum chloride and barium carbonate are
shown in the rate curves. Figures 3A and 3B. The reaction
obeyed zero order kinetics which indicated that the rate
determining step in the reaction was the solution process
of the barium carbonate in the methylammonlum chloride melt.
The kinetic data are shown in Table IB, while the Arrhenius
curve is shown in Figure L;.. The activation energy was found
to be 25 i 14- kcals per mole.
Dlmethylammonium Chloride and Barium Carbonate
The mass-loss, DTA, GEA, and MSA curves are shown in
Figure Sf curves A through D, respectively.
The mass-loss curve showed a one step mass-loss from
190° to 200®C, The percentage mass-loss of the total initial
sample was calculated to be 31.7^f which compares favorably
with ^2,2l\.%, the theoretical mass-loss calculated from the
acid-base reaction:
17
TABLE 1
The Reaction of Methylammonlum Chloride with Barium Carbonate
A. Reaction Stolchiometry Data
P r o d u c t T h e o r e t i c a l F o u n d N o r m a l i z e d Analyzed (mmoles ) (mmoles ) to Equation
MeNH2 1.77 1.72 1.91; BaCl2 0,89 0,88 0.99 CO2 1.00 1.00 1.00
B, Kinetic Data
1/T X 10^ -in K E' (Kcal/mole)
2,10 9.73 25 t I; 2.01 8.88 2.02 8.71 1.99 8.51 1.97 8.28 1.95 7.98
18
Figure 3
Rate Curves for the Reaction of Methylammonlum Chloride with Barium Carbonate
19
T=5I2»K
as--
2JO'
X
CSJ
X
o 1.5 •
O
<
10" UJ o z o o
0.5-•
^^0
T=496*K
T=477»K
10 15 20 TIME (Minutes)
25 30
20
25 I
2 0
10
g X
z fO
O
u. o
1-5
< IT I -Z l UJ
o z O O
LO-
05
Ts502»K
T=489»K
00- r To fe 20 TIME (Minutes)
25 30
21
Figure I4.
Arrhenius Curves for the Reaction of Methylammonlum Chloride with Barium Carbonate
22
6J0+
7.0-•
1.90 2JOO
l/T XIO 2J0
23
Figure 5
Instrumental Curves for the Reaction of Dimethylammonlu Chloride with Barium Carbonat
(A). TGA (B), DTA (C). GEA (D). MSA
m e
2k
100 - 200 TEMPERATURE (°Q
300
25
2Me2NH2Cl + BaCO^ >-BaCl2 + 2Me2NH + COg + H2O (2),
The DTA curve showed three endothermic peaks with
AT^in ®^ k-^ * 1^0 » ®^^ 185 C, respectively. The endother
mic peak at I4.O C was a solid-solid phase transition of the
dlmethylammonium chloride, as seen from the DTA curve of
the pure substance, and the endothermic process at I60 C
was the melting of the dlmethylammonium chloride since
there was no gas evolution at this point. Gas evolution
peaks are shown with T^g^j^ of 175^ and l85^C,
The MSA results are shown in Figure 5^. These curves
showed the gaseous products to be dimethylamine, water, and
carbon dioxide. At no time were chlorine or hydrogen chlo
ride observed in the mass spectrum.
The results of the stolchiometry studies are shown
in Table 2A, As seen from the data, the reaction was found to
be quantitative.
The reaction kinetics were found to be zero order with
respect to dimethylamine concentration indicating a solution
process as the rate determining step. The rate curves are
shown in Figure 6A and 6B, These curves gave the rate of
solution of the barium carbonate in the dlmethylammonium
chloride melt. The Arrhenius activation energy for the
solution process was determined to be 3 - 0,7 kcal/mole,
as calculated from the curve shown in Figure 7.
26
TABLE 2
The Reaction of Dlmethylammonium Chloride with Barium Carbonate
A. Reaction Stolchiometry Data
Product Analyzed
Theoretical (mmoles)
i 'ound (mmoles)
Normalized to Equation
Me2NH BaCl2 CO2
1.82 0.91 0.92
1.80 0 .92 0 .91
1.98 1.01 0 .99
B, Kinetic Data
1/T X 10^ -In k E* (Kcal/mole)
2.22 2.12 2.03 2.00 1.97 1.95
7.01 6.92 6.80 6,66 6,81 6.70
3 - 0.7
27 •.fr->*.H,ft, •mi,t1^f^lt0ll0i^0ltll^ '•*,»*.,
Figure 6
Rate Curves for the Reaction of Dlmethylammonium Chloride with Barium Carbonate
28
3 4 TIME (Minutes)
29
TIME (Minutes)
'i'
'y^' i'
30
Figure 7
Arrhenius Curves for the Reaction of Dlmethylammonium Chloride with Barium Carbonate
31
32
TrImethylammonlum Chloride and Barium Carbonate
The mass-loss, DTA, and GEA curves are shown in Figure
8, curves A through C, respectively.
The mass-loss curve showed a single step mass-loss from
100° to 250°C. The theoretical mass-loss, as calculated from
Equation (3), was found to be 38.14-3/ of the total initial
sample mass.
2Me NHCl + BaCO *-Zi^e^li + BaClg + H2O + CO2 (3),
The actual mass-loss calculated from the mass-loss
curve was 2S.9%» The DTA curve contained two endothermic
peaks with AT ^ ^ at 50 and 225 C, respectively, while the
GEA curve showed but one peak. The endothermic peak at 50°C
was caused by a solid-solid phase transition of trimethyl
ammonlum chloride. It was found from the attempt to obtain
the mass spectra of the products that the trimethylammonlum
chloride sublimed at the temperatures employed. The rate
of sublimation was comparable to that of the reaction. No
mass spectra of the reaction products were obtained since
the sublimed trimethylammonlum chloride plugged the inlet
system to the mass spectrometer each time an analysis was
attempted.
Since the reaction was not quantitative, reaction
kinetics studies were not attempted.
Methylammonlum Chloride and Strontium Carbonate
The mass-loss, DTA, GEA, and MSA curves are shown in
^j!
33
Figure 8
Instrumental Curves for the Reaction of Trimethylammonlum Chloride with Barium Carbonate
(A). (B), (C).
TGA DTA GEA
3k
100 200 300 400 _, 500 600 TEMPERATURE m
35
Figure 9, curves A through D, respectively.
The mass-loss curve revealed a single mass-loss start
ing at 180°C and changing slope in the 270°C region. Reaction
continued at a slower rate to 335°C, where the reaction
ceased and thermally stable products were formed. Calcula
tions from Equation (U) give a theoretical percentage mass-
loss of 26.33^» as compared to the actual mass-loss of 25.2^.
2MeNH^Cl + SrCO^ »• 2MeNH2 ^ SrCl2 - H2O + CO2 (U).
The DTA curve of the reaction revealed that there were
three endothermic processes involved with peaks at 230°, 310®,
and 325^0, respectively. The 310°C peak was seen as a shoulder
on the 325°C peak.
The GEA curve showed that gaseous products were evolved
during all three of the reactions. The products found for
the two main reactions, using the technique of MSA, were
carbon dioxide, water, and methylamlne. At no time was
chlorine or hydrogen chloride observed.
The results of the reaction stolchiometry studies are
shown in Table 3A. The data show that the overall reaction
is the reaction given by Equation (I4), However, attempts
to obtain the Intermediate composition by conventional analy
tical methods and mass-loss studies resulted in failure be
cause of the overlap of the two reactions.
Determination of the reaction kinetics of these two
reactions was fairly simple since both reactions were found
36
Figure 9
Instrumental Curves for the Reaction of Methylammonlum Chloride with Strontium Carbonate
(A). (B). (C). (D).
TGA DTA GEA MSA
37
200 300 TEMPERATURE ( C)
38
TABLE 3
The Reaction of Methylammonlum Chloride with Strontium Carbonate
A, Reaction Stolchiometry Data
Product Analyzed
Theoretical (mmoles)
Pound (mmoles)
Normalized to Equation
MeNH2 SrCl2 CO2
2,21 1.11
1.23
2.014. l.Oli.
1.114-
1.85 0.9U 0.93
B. Kinetic Data
1/T X 103 -In ki -In k2 E^ (Kcal/mole) E2 (Kcal/mole
2.09 2.05 2.00 1.98 1.89
10.13 9.51 8.97 8.29 8.1;2
10.17 9.14-7 9.1^6 9.39 8.57
5 ^ 1 2 1 - 3
39
to be zero order with respect to methylamlne concentration.
A typical pH vs. time plot is shown in Figure 10. The curve,
obtained under isothermal conditions, showed a very fast
evolution of methylamlne. The curve then changed slope in
dicating a slower reaction in which methylamlne was evolved.
The rate of the second reaction was easily determined from
the slope of the pH y^. time curve. This represents the
rate of the second reaction only, since the first reaction
was at this time completed. The rate obtained from the first
portion of the pH XS« time curve can be shown to be the rate
of the first reaction plus the rate of the second reaction.
The rate equation for the rate determining reaction is:
dC ^
"dt"
dG2 ^
ki or J"dCi = J'kidt (a)
^1 * ^1* "*" 1
dt = K2 or fdG2 = f^2^^ ^ ^
Cp— Kpt +• Zp
where C^ is the concentration of amine liberated by the first
reaction, C2 the concentration of the amine liberated by the
second reaction, t the time, k ^ the rate constants, and Z ^
and Z2 the Integration constants. Now, if equation (a) is
added to equation (b), and X is defined as the total concen
tration of methylamlne given off in the first reaction:
ko
Figure 10
Typical pH vs. Time Plot
k1
k2
dX 60-^ dC^ • — ^ = K, + K^ (c)
dt dt dt 1 2
JdC^ + jdC^ = /K-j dt + f^2^^ ^^^
C^ + C2 = (k t + Z;,) + ( k2t + Z2) (e)
°1 - 2 = (k^ + k2)t + (Z^ -H Z2) (f)
^lotting C^ 4- C2 vs . t, the slope of the line was k^ +
k2> and the intercept of the line was Z-, -h Zp or the sum of
the two integration constants. Since k2 was easily obtained
from the second portion of the pH vs_. time curve, k^ was ob
tained by taking the value of k ^ + k2 from the first rate
curve and subtracting the value of k2 obtained from the second
rate curve. The rate curves of k^ + k2 and the rate curves
of k2 are shown in Figures 11 and 12, respectively. The data
for the Arrhenius curve is shown in Table 3B, and the Arrhen
ius curve is shown in Figure I3. The Arrhenius curve for the
activation energy of the second reaction is shown in Figure
ill. The activation energy obtained for the first reaction
was 5 - 1 kcal per mole, and the activation energy for the
second reaction was found to be 21 1 3 kcal per mole.
It is believed that the first reaction (5) Involves
the formation of a metal chlorohydrogencarbonste and the
second reaction (6) the decomposition of it into carbon diox
ide, water, and metal chloride.
13
Figure 11
Rate Curves for the Reaction of Methylammonlum Chloride with Strontium Carbonate
Reaction 1
kk
4.0--
35
ro O X 3 0 f
CVJ
X
-I" o U.2^" O
O
<
a:
2JO-
UJ o z 8 i- t
bO
0.5
ao|
T=530'l
TIME (Minutes)
l5
Figure 12
Rate Curves for the Reaction of Methylammonlum Chloride with Strontium Carbonate
Reaction 2
1;6
13 14 15 16 17 18 19 20 21 T\':/iE (Minutes)
22 23 24 25 26
kl
Figure 13
Arrhenius Curves for Methylammonlum Chloride and Strontium Carbonate
Reaction 1
U8
1/T X 10
•m
k9
Figure ll .
Arrhenius Curves for Methylammonlum Chloride with Strontium Carbonate
Reaction 2
50
7.a
8J0-»-
C
I
9.0'
JOCr r_v
"•'fe 1.95 5i 2.05 1/T XIO^
51
MeNH^Cl + SrCO^ ^MeNH2 + Sr(HC03)Cl (5)
MeNH^Cl + SR(HC03)C1 • MeNHg + SrCl2 + H2O + COg (6)
No data, however, were obtained that substantiated this
series of reactions. It does seem logical that the carbonate
would decompose through the hydrogen carbonate, and the dif
ference in the rates of Equation (5) and Equation (6) would
determine whether the hydrogencarbonate would be seen or not.
Dlmethylammonium Chloride and Strontium Carbonate
Dlmethylammonium chloride and strontium carbonate were
found to react at a temperature of 170°C, The mass-loss
curve, shown in Figure 15A, indicates a mass-loss of 3S»3%
of the original sample mass, as compared to the theoretical
mass-loss of 35.87/^ as calculated from the acid-base reaction:
2Me2NH2Cl + SrCO^ ^ 2Me2NH + SrCl2 + H2O + CO2 (7)-
The results of the reaction stolchiometry studies,
as shown in Table i A, substantiate Equation (7).
The DTA studies of the above reaction revealed three
endothermic peaks with a AT^^j^^ at 60°, 165°, and 195^0, re
spectively, as shown in Figure l5B. The endotherm at 60°C
is a solid-solid phase transition of dlmethylammonium chlo
ride since no gas 'evolution peak was observed at that tem
perature. Mass spectrometric analysis found that the gases
evolved during these reactions were dimethylamine, water, and
carbon dioxide, as shown in Figure 15D.
52
Figure 15
Instrumental Curves for the Reaction of Dlmethylammonium Chloride with Strontium Carbonate
(A). TGA (B). DTA (C). GEA (D). MSA
53
C/) CO o
10 mg ®
<3
o
O > UJ
CO <
1.8*
Imv
> -
co
UJ I -
(CH3)2NH
CO2
HgO
UJ Q:
100 200 TEMPERATURE (* C)
300
5lf
TABLE I4.
The Reaction of Dlmethylammonium Chloride with Strontium Carbonate
^. Reaction Stolchiometry Data
Product i'heoretlcal Pound Normalized Analyzed (mmoles ) (mmoles) to Equation
Me2NH 2.26 SrCl2 1.13 CO2 l.ll;
B. Kinetic Data
1/T X 1 0 ^ -In k E ^ (Kcal/mole)
2.3I4- 8,19 8 i 2 2.22 7.78 2.11 7.39 2,01+ 7.39
2.21 1.12 1.10
1.96 0.99 0,97
55
The results of the Isothermal rate studies are shown
in Figure 16 and in Table 1;B. The kinetics were found to
be zero order with respect to dimethylamine concentration,
and the Arrhenius curve, as shown in Figure 17, gave an
activation energy of 8 - 2 kcal per mole for the reaction.
Trimethylammonlum Chloride and Strontium Carbonate
Trimethylammonlum chloride was found to react quanti
tatively with strontium carbonate. The mass-loss, DTA, and
GEA curves are shown in Figure l8, curves A through C, re
spectively.
The mass-loss curve revealed that the reaction between
trimethylammonlum chloride and strontium cerbon«?te took place
from 150° and 250°C, with a resulting mass-loss of 31;.0 ,
as compared to a theoretical mass-loss of 3U.70^ calculated
from the acid-base reaction:
2Me3NHCl + SrC03 >• 2Me3N + SrCl2 + CO2 + H2O (8).
The DTA curve showed four endothermic peaks with AT^j^-
at [4.0°, 150°, 195°, and 3l5°C, respectively. The gas evo
lution curve showed no gas evolution at I4.O C indicating this
to be a solid-solid phase transition of trimethylammonlum
chloride. Gas evolution peaks with T ^ ^ at 150 , 195 » and
3l5°C, respectively, were observed.
The reaction stolchiometry studies showed that the
products of the acid-base reaction were produced in quanti
tative amounts, as shown in Table 5A. It was felt at this
point that in view of the preceding reactions and the analysis
56
Figure 16
Rate Curves for the Reaction of Dlmethylammonium Chloride with Strontium Carbonate
11.0-
57
2 4 8 10' 12 ^ 16 TIME (Minutes)
18 20 22 24 26
58
Figure 17
Arrhenius Curves f o r the Reaction of Dlmethylammonium Chloride wi th Strontium Carbonate
59
- I n - k ' -
60
Figure 18
Instrumental Curves for the Reaction of Trimethylammonlum Chloride with Strontium Carbonate
(A). TGA (B). DTA (C), GEA
61
100 200 TEMPERATURE ( C)
300
62
TABLE 5
The Reaction of Trimethylammonlum Chloride with Strontium Carbonate
A. Reaction Stolchiometry Data
Product Analyzed
Theoretical (mmoles)
Found (mmoles)
Normalized to Equation
Me3N SrCl2 CO5
2.07 i.ou 1.85
2,02 1.02 1.92
1.95 0.98 l.Oli.
B. Kinetic Data
1/T X io3 -In k E ' (Kcal/mole)
2.12 2.07 2.05 2.02
8.22 7.63 7.I1.6 7.06
2 7 - 1 ^
63
thus far, the mass spectrometric data were not needed to
characterize this reaction fully and were not obtained.
Reaction kinetics were found to be zero order with
respect to trimethylammonlum chloride. The results of the
isothermal rate studies are shown in Figure 19 and In Table
5B. With these Isothermal rate data, an Arrhenius curve
was plotted and the activation energy determined to be
27 " 14- kcal per mole, as shown in Figure 20.
61+
Figure 19
Rate Curves for the Reaction of Trimethylammonlum Chloride with Strontium Carbonate
65
0.0, 4 t—t—> A .k TIME (Minutes) 13
66
Figure 20
Arrhenius Curve for the Reaction of Methylammonlum Chloride with Strontium Carbonate
67
-€J'-
l/T X 10"
CHAPTER III
REACTIONS OF METAL OXIDES IN FUSED ONIUM SALTS
Now that the reactions of metal carbonates in fused
onium salts have been characterized, it was interesting to
see if the metal oxides also react under the same conditions,
since they are much more stable than the carbonates. It
was also interesting to speculate on and to predict the pro
ducts of the reactions between onium salts and transition
metal oxides, since the metals used in the precedinp- discus
sion were non-coordinating metals.
Methylammonlum Chloride and Zinc Oxide
It was found that reaction took place between methyl
ammonlum chloride and zinc oxide at 130°C. The mass-loss
curve, shov;n in Figure 21A, revealed a three step mass-loss.
The first mass-loss resulted in the loss of 3.9^ of the
total original sample mass. The theoretical percentage
of water calculated from Equation (9) was 3*k%*
2MeNH3Cl + ZnO ^ZVie'^E^ + ZnCl -»• H O (9),
The second mass-loss was found from 200° to 500°C and re
sulted in the loss of 11.3^ of the total original sample
mass. The theoretical mass-loss resulting from the evolu
tion of methylamlne, according to Equation (9), was calculated
68
69
Figure 21
Instrumental Curves for the Reaction of Methylammonlum Chloride with Zinc Oxide
(A). TGA (B), DTA (C), GEA (D). MSA
70
100 200 300 4 0 0 500 TEMPERATURE ( C)
600 7oa
71
to be 11.99^. The third mass-loss observed from 500° to
700 C was due to the conversion of zinc chloride to zinc
oxide by air. The slight discrepancies in the calculated
and actual mass-loss were explained by the following:
The DTA study of the reaction showed four endothermic
peaks with AT^^^ at 85°, 100°, 130°, and 300°C, respect
ively, as shown in Figure 21B. The gas evolution curve, as
shown in Figure 21C, revealed that the DTA endotherms at
^"^min °' ®^^ ®^^ 100°C, respectively, were not accompanied
by gas evolution peaks. The mass spectra of the reaction
products, shown in Figure 21D, indicated that the first re
action evolves water and a small amount of methylamlne.
This accounts for the fact that the mass-loss curve cal
culations were a little high. The second reaction then
involved the evolution of only methylamlne.
The results of the reaction stolchiometry studies are
shown in Table 6. The data lead to the following reaction
sequence:
2MeNH3Cl + ZnO P- Zn(MeNH2)2Cl2 + H2O (10)
Zn(MeNH2)2Cl2 >-ZnCl2 -»• 2MeNH2 (11).
TABLE 6
Reaction Stolchiometry Data
Product 'i'heoretical -t'ound Normalized Analyzed (mmoles ) (mmoles ) to Equation
MeNH2 ^-^2 ^'^ 2.00 ZnCl2 2.1+1 2.I4. 1.00
72
Methylammonlum Chloride and Cadmium Oxide
A similar reaction took place when a mixture of methyl
ammonlum chloride and cadmium oxide was heated at a linear
rate. The mass-loss curve, as shown in Figure 22A, revealed
that the reaction involved two mass-losses; the first from
200° to 225°C, and the second from 225° to 360°C. The first
mass-loss resulted in the loss of 7.9^ of the total original
sample mass. The theoretical mass-loss due to water, accord
ing to reaction equation (12), was found to be k»0^% of the
total original sample mass.
2MeNH3Cl -•- CdO >-CdCl2 + H2O + 2MeNH2 (12).
The second mass-loss resulted in the loss of 12,l\%
of the total original sample mass. Calculation of the theo
retical percentage mass-loss due to methylamlne resulted in
a percentage mass-loss of 13.97^ of the total original sample
mass. The reason for these discrepancies will be obvious
later.
The DTA curve of the reaction, as shown in Figure 22B,
revealed four endothermic peaks with AT^^^^ of 170°, 200°,
225°, and 270°C, respectively. The endotherm at 170°C is
believed to be due to the fusion of the methylammonlum chlo
ride since the gas evolution curve, as given in Figure 22C,
showed no gas evolution at this temperature. The remainder
of the endothermic peaks were accompanied by gas evolution.
The mass spectrometric analysis of the gases, as given in
Figure 22D, showed that methylamlne and water are liberated
during both reactions.
73
Figure 22
Instrumental Curves for the Reaction of Methylammonlum Chloride with Cadmium Oxide
(A). TGA (B). DTA (C). GEA (D). MSA
71+
100 2 0 0 300 400 500 TEMPERATURE ( C)
600
75
The reaction stolchiometry results are shown in Table
7. With this data It was easily seen why the calculations
from the mass-loss curve could not be correlated.
TABLE 7
Reaction Stolchiometry Data
Product Theoretical Found Normalized Analyzed (mmoles ) (mmoles) to Equation
MeNH2 2.83 2.8 1.98
CdCl2 I.I4I 1.2 0.85
If it is pointed out that a one to one mixture of
methylammonlum chloride to cadmium oxide was used, and it
is noted that this is a 100^ excess of cadmium oxide, then
the series of reactions that follow will explain the data
very well. The initial reaction can proceed by one of two
paths:
2MeNH3Cl . 2C60-^'''^''''^2^2'h ^ "("^^S (13)
CdCl2 + H2O -K 2MeNH2 - CdO
Then the following reactions occur:
Cd(MeNH ) CI »-CdCl2 + 2MeNH2 (II4-)
Cd(0H)2 »-CdO + H2O (15).
Methylamm.onium Chloride and Ma.c nesium Oxide
Magnesium oxide and methylammonlum chloride were mixed
in a one to one mole ratio and heated. The mass-loss curve,
as given in Figure 23A, showed that the reaction took
76
place with three d i s t i nc t mass-losses: the f i r s t from 200°
to 275°C; the second from 275° to 1|00°C; and the third from 1 o o
I;00 to 725 C, The first mass-loss resulted in the loss
of 28.5^ of the total initial sample mass. The theoretical
percentage mass-loss due to the evolution of methylamlne
from the acid-base reaction was 28.72% based on the total ini
tial sample mass. The second mass-loss resulted In the loss
of 10.0^ of the total initial sample mass. The theoretical
percentage mass-loss due to evolution of water from the
acid-base reaction was 9.7U^ based on the total initial
sample mass. The third mass-loss then, by necessity, was
the oxidation of magnesium chloride by the air to magnesium
oxide and chlorine gas and resulted in the loss of 25.6^
of the total original sample mass. The theoretical mass-
loss due to the oxidation of the magnesium chloride, was
25.55^ of the total original sample mass.
The DTA curve, as given in Figure 23B, revealed four
endothermic peaks with AT^^^ ®^ 190°, 210°, 250°, and 350°C,
respectively. The first endotherm at 190°C is the melting
of the methylammonlum chloride followed immediately by the
reaction endotherms at 210° and 250°C, respectively. The
fourth endotherm at 350°C was found to be the decomposition
of magnesium hydroxide to magnesium oxide and water.
The gas evolution curve, as shown in Figure 23C, re
vealed that gaseous products were evolved at all the endothermic
77
Figure 23
Instrumental Curves for the Reaction of Methylammonlum Chloride with Magnesium Oxide
(A). TGA (B). DTA (C). GEA (D), MSA
78
100 200 3_0p 400 TEMPERATURE
500 600
79
DTA peaks indicating that reaction is immediate after the
melting of the methylammonlum chloride.
The MSA curve, as given in Figure 23D, showed that
methylamlne and water were evolved from the first three re
actions, but only water was evolved from the last reaction.
This shows very clearly the decomposition of magnesium hy
droxide.
There is, however, an apparent contradiction between
the mass-loss results and the mass spectrometric data, since
the mass-loss suggests that all the water went to form mag
nesium hydroxide while the mass spectrometric results sug
gest that some water is given off during the initial reac
tion. It must be remembered at this point that the two
experiments were carried out under different conditions.
The mass loss studies were run in a static air atmos
phere while the mass spectrometric studies were carried
out under a very high flow rate of helium, greater than
100 ml/minute. This high flow rate would tend to remove
any evolved water vapor while the water vapor formed in
the sample mass-loss would have to diffuse through the mag
nesium oxide and would remain in the sample longer allow
ing time for the reaction to form magnesium hydroxide not
formed in the initial reaction.
With the reaction stolchiometry data, as shown in
Table 8, and the above data, it is possible to formulate
the following reaction to explain the data:
80
2MeNH3Cl + 2MgO *-2MeNH2 + Mg(0H)2 + MgCl^ (16)
Mg(0H)2—^MgO + H^O (17)
2MgCl2 + O2 :»-2MgO + 2CI2 (in air) (I8).
TABLE 8
Reaction Stolchiometry Data
Product 'i'heoretical Found Normalized Analyzed (mmoles ) (mmoles) to Equation
MeNH2 1.37 1.1 1.61
MgCl2 1.37 1.1 0.80
Methylammonlum Chloride and Cobalt(III) Oxide
Cobalt(III) oxide was found to react with methylam
monlum chloride at 200°C by means of an oxidation-reduction
reaction. The mass-loss, DTA, GEA, and MSA curves are shown
in Figure 2l+, A through D, respectively.
The mass-loss curve showed a two step mass-loss, the
first resulting in the loss of 25.3^ of the total original
sample mass and the second being continuous to 500 C. The
second mass-loss was due to the sublimation of excess
methylammonlum chloride used in order to get complete re
action for the magnetic susceptibility studies. Since this
sublimation was a continuous process above 200°C, no mean
ingful calculations could be made from the mass-loss curve.
However, the change in magnetic moment serves as a good
analysis for the cobalt compounds formed.
81 - V )tf*ili«w .»*yi**
Figure 2I4.
Instrumental Curves for the Reaction of Methylammonlum Chloride with Coba l t ( I I I ) Oxide
(A), (B) , (C), (D).
TGA DTA GEA MSA
