Note: Descriptions are shown in the official language in which they were submitted.
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TITLE: METHOD FOR PRODUCING HYDROGEN
FIELD OF THE INVENTION
This invention relates to the production of hydrogen gas from the reaction
of aluminum with water in the presence of sodium hydroxide as catalyst.
BACKGROUND OF THE INVENTION
Hydrogen energy is environment-friendly. Because of the actual human
ecology concerns, the exploitation of hydrogen as an universal fuel would
be greatly acclaimed. During the last two decades or so, the elaboration of
a hydrogen-based economy has made important progress on account of
numerous research projects such as the hydrogen fuel cell and the hydrogen
car. Although these important discoveries constitute milestones toward a
pollution-free society, more research is needed to obtain the hydrogen
easily and economically.
A convenient source of hydrogen is a reaction of aluminum with water to
split the water molecules into hydrogen and oxygen. The hydrogen is
released as a gas and the oxygen combines with the aluminum to form
aluminum oxide compounds. Aluminum is the third most abundant
element after oxygen and silicon in the earth's crust, and constitutes
approximately 8% by weight of the earth's crust. Aluminum is a safe
material and is commonly used in the food, cosmetics and medical fields.
Water is also abundant. Therefore, the reaction of these two elements to
produce hydrogen represents an interesting proposal to replace fossil fuels.
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Generally speaking, it is known that under certain conditions, aluminum
reacts with water to generate hydrogen and heat. It is also known, however,
that this type of reaction is not sustainable at ambient temperature. It is
believed that a protective oxide layer forms on a metal surface in contact
with water at ambient temperature and hinders the reaction. Therefore, it
has been accepted by those skilled in the art that the use of aluminum in a
reaction with water to generate hydrogen gas requires that the protective
oxide layer is efficiently and continuously removed, and that the reaction
is kept at an elevated temperature.
A number of hydrogen generators have been developed in the past. The
following patent documents constitute a good inventory of the devices and
methods of the prior art in the field of hydrogen gas generation using the
reaction of aluminum or alloys of aluminum with water.
US 909,536 issued on Jan. 12, 1909, and US 934,036 issued on Sept. 14,
1909, both issued to G. F. Brindley et al. These documents disclose several
compositions for generating hydrogen. The compositions comprise any
metal which can form an hydroxide when it is brought into contact with a
solution of a suitable hydroxide. For example, aluminum is reacted with
sodium hydroxide to release hydrogen and to produce sodium aluminate.
US 2,721,789,issued on Oct. 25, 1955 to Q.C. Gill. This document
discloses the structure of an hydrogen generator for reacting water with a
measured dry charge of aluminum particles and flakes of sodium
hydroxide. The reaction releases hydrogen gas and produces sodium
aluminate.
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US 3,554,707 issued on Jan.12,1971 to W.A. Holmes et al. This document
discloses a gas generator having bellows to raise or lower the level of water
in response to the pressure inside the generator. As the level of water
drops, the contact surface between the fuel cartridge and the water is lost
and the reaction is terminated.
US 3,957,483 issued on May 18, 1976 to M. Suzuki. 'This patent discloses
a magnesium composition which produces hydrogen upon contact with
water. The preferred magnesium composition comprises magnesium, and
one or more metals selected from the group consisting of iron, zinc,
chromium, aluminum and manganese.
US 3,975,913 issued on Aug. 24, 1976 to D.C. Erickson. This document
discloses a hydrogen generator wherein molten aluminum is reacted with
water. The generator is kept at a very high temperature to keep the metal in
a molten condition.
US 4,643,166 issued on Feb. 17, 1987, and
US 4,730,601 issued on Mar. 15, 1988 both to H.D. Hubele et al. These
documents disclose the structure of a fuel cell for producing heat energy
and hydrogen gas. The device has a reaction chamber containing a fuel
composition that is reactive with water. The fuel composition includes a
main fuel part of magnesium and aluminum in a molar ratio of 1:2, and the
second part is composed of lithium hydride, magnesium and aluminum in
equal molar ratio.
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US 4,670,018 issued on June 2, 1987, and
US 4,769,044 issued on Sept. 6, 1988, both to J.H. Cornwell. These
documents describe a log made of compressed wood waste and paper. The
log is coated with aluminum particles. Upon burning, the aluminum
particles react with moisture in the log to emit heat due to the generation of
hydrogen gas.
US 4,752,463 issued on June 21, 1988 to K. Nagira et al. This document
discloses an alloy which reacts with water for producing hydrogen gas. The
alloy material comprises essentially aluminum and 5 to 50% tin.
US 5,143,047 issued on Sept. 1, 1992 to W.W. Lee. This document
discloses an apparatus and a method for generating steam and hydrogen
gas. In this apparatus, an aluminum or aluminum alloy powder is reacted
with water to generate hydrogen gas. An electric power source is used to
start the reaction. The electric power source is used to explode an
aluminum conductor and to disperse pieces of molten aluminum into a
mixture of water and aluminum powder. A heat exchanger is provided to
extract useful heat.
US 5,867,978 issued on Feb. 9, 1999 to M. Klanchar et al. This document
discloses another hydrogen gas generator using a charge of fuel selected
from the group consisting of lithium, alloys of lithium and aluminum. The
charge of fuel is molten and mixed with water to generate hydrogen gas.
JP 401,208,301issued to Mito on Aug. 22, 1989. This document discloses
a process for producing hydrogen. Aluminum is reacted with water under
an inactive gas or a vacuum to produce hydrogen gas.
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CA 2,225,978 published on June 29, 1999 by J. H. Checketts. This patent
application discloses a hydrogen generation system wherein a coating on
reactive pellets is selectively removed to expose the reactive material to
water for producing hydrogen gas on demand. In one embodiment,
aluminum and sodium hydroxide are reacted with water to release hydrogen
gas and produce sodium aluminate.
DE 3,401,194 published in July 18, 1985 by Werner Schweikert. This
document discloses a device for utilizing energy from a chemical reaction
between various aluminum alloys and sodium hydroxide. The chemical
reaction occurring in this device generates heat, hydrogen gas, a direct
current and sodium aluminate as a residue.
FR 2,465,683 published in Mar 27, 1981 by Guy Ecolasse. This document
also discloses a process for producing hydrogen by the reaction of
aluminum on sodium hydroxide solution in water. A by-product of this
reaction is sodium aluminate.
Belitskus, David. 1970. Technical Note: "Reaction of Aluminum With
Sodium Hydroxide Solution as a Source ofHydrogen" J. Electrochem Soc.
(1970), (August), Vol. 117. No. 8, pp.1097-9, XP-002180270. This
technical paper describes several experiments wherein aluminum samples
including a cylindrical block, uncompacted powders and pellets of various
densities have been reacted with aqueous solutions of sodium hydroxide at
various concentrations to generate hydrogen gas. In t:hese experiments, the
formation of sodium aluminate was observed, as well as the regeneration
of sodium hydroxide through the precipitation of aluminum hydroxide.
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Stockburger, D. et al. 1991. "On-Line Hydrogen Generation from
Aluminum in an Alkaline Solution". Proc.-Electrochem. Soc. (1992), Vol.
92-5 (Proc. Symp. Hydrogen Storage Mater., Batteries, Electrochem), pp.
431-44, 1992, XP-001032928. This technical paper describes three sizes
of hydrogen generators in which aluminum is reacted with an aqueous
solution of 5.75 M sodium hydroxide. This technical paper also notes the
formation of sodium aluminate and the precipitation of aluminum
hydroxide that regenerates sodium hydroxide.
Although the chemical reactions of aluminum with water in the presence
of sodium hydroxide have been demonstrated in various projects in the
past, these reactions were not considered as being safe for use by the
general public. Sodium hydroxide is extremely corrosive and must be
handled according to particular safety procedures. Therefore, any chemical
reaction wherein sodium hydroxide is a consumable would not represent an
attractive source of hydrogen for use in vehicles or in household power
systems, for examples. As such, it is believed that a need still exists for a
method to produce hydrogen gas by the reaction of aluminum and water,
wherein the consumables are limited to aluminum and water.
SUMMARY OF THE INVENTION
Broadly stated, the process for producing hydrogen gas according to the
present invention consists of reacting aluminum with water in the presence
of sodium hydroxide acting as a catalyst.
In accordance with one aspect of the present invention, there is provided
a process for producing hydrogen gas, comprising the initial step of
providing an aqueous solution in a vessel. The aqueous solution contains
sodium hydroxide in a concentration between 0.26 M and 19 M NaOH.
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The next step consists of reacting aluminum with water at the surface of the
solution thereby generating a region of effervescence at the surface of the
solution and a precipitate sinking to the bottom region of the vessel. The
process also includes the step of maintaining the region of effervescence
separated from the precipitate at the bottom the vessel, to prevent the
precipitate from swirling and mixing with the aluminum in the reaction
zone at the surface of the solution. This process is advantageous because
it proceeds catalytically with the sodium hydroxide acting as the catalyst.
The process mentioned above is best carried out with an aqueous solution
containing between about 5M and 10 M NaOH. The process is also more
efficient when makeup water is added only after an initial amount of
aluminum has been consumed, and when the temperature of the aqueous
solution has reached a peak or 75 C.
In accordance with another aspect of the present invention, there is
provided a process for initiating and maintaining a catalytic reaction of
aluminum with water for producing hydrogen gas. The process comprises
the initial step of providing an aqueous solution in a vessel. This aqueous
solution contains a portion of NaOH and a portion of water. The next steps
consist of introducing a portion of aluminum in the aqueous solution, and
reacting that portion of aluminum with the portion of water. The process
also includes the steps of maintaining constant the portion of NaOH in the
vessel and adding additional portions of water and additional portions of
aluminum in the vessel according to the rates of consumption of the
aluminum and the water in the reaction.
Again, this process is best carried out with an aqueous solution of between
1.2 M and 19 M NaOH and at a temperature between 4 C and 170 C.
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In yet another aspect of the present invention, there is provided a process
for simultaneously producing hydrogen gas and alumina (A1203). This
process firstly comprises the step of providing an aqueous solution in a
vessel. The aqueous solution contains sodium hydroxide in a concentration
between 0.26 M and 19 M NaOH. The next step consists of reacting
aluminum with water at a surface of the aqueous solution and generating
hydrogen gas and alumina. The process also includes the step of
recovering hydrogen gas from the surface of the aqueous solution and
alumina from a bottom region of the vessel.
Other advantages and novel features of the present invention will become
apparent from the following detailed description.
BRIEF DESCRIPTION OF THE DRAWINGS
A preferred embodiment of the process according to the present invention
selected by way of examples will now be described with reference to the
accompanying drawings, in which:
FIG. 1 is graph illustrating a first reaction of aluminum with water to
produce hydrogen gas, in a 5.0 M sodium hydroxide solution,
carried out over a period of about 130 minutes;
FIG. 2 is a graph illustrating a second reaction of aluminum with water to
produce hydrogen gas, in a 4.95 M sodium hydroxide solution,
carried out over a period of about 100 minutes;
FIG . 3 is a graph illustrating a third reaction of aluminum with water in a
4.5 M sodium hydroxide solution, while keeping the temperature
relatively low;
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FIG. 4 is a graph illustrating a fourth reaction of aluminum with water in
a 4.5 M sodium hydroxide solution, while keeping the temperature
relatively low;
FIG. 5 is a graph illustrating a reaction of aluminum with water in a 1.2 M
NaOH solution;
FIG. 6 is a graph illustrating a reaction of aluminum with water in a 2.5 M
NaOH solution;
FIG. 7 is a graph illustrating a reaction of aluminum with water in a 3.9 M
NaOH solution;
FIG. 8 is a graph illustrating a reaction of aluminum with water in a 4.8 M
NaOH solution;
FIG. 9 is a graph illustrating a reaction of aluminum with water in a 5.5 M
NaOH solution;
FIG. 10 is a graph illustrating a reaction of aluminum with water in a 6 M.
NaOH solution;
FIG. 11 is a graph illustrating a reaction of aluminum with water in a 6.03
M NaOH solution;
FIG. 12 is a graph illustrating a reaction of aluminum with water in a 6.1
M NaOH solution;
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FIG. 13 is a graph illustrating a reaction of aluminum with water in a 6 M
NaOH solution, wherein the water was added continuously;
FIG. 14 is a graph illustrating a reaction of aluminum with water in a 6 M
NaOH solution, wherein the aluminum was added quickly;
FIG. 15 is a graph illustrating a reaction of aluminum with water in a 6.7
M NaOH solution;
FIG. 16 is a graph illustrating a reaction of aluminum with water in a 11.3
M NaOH solution;
FIG. 17 is a graph illustrating a reaction of aluminum with water in a
saturated 19 M NaOH solution;
FIG. 18 is graph illustrating maximum reaction temperatures obtained with
aqueous solutions ofvarious concentrations, and the responsiveness
of the reaction for solutions of various concentrations;
FIG. 19 is a graph showing the effects of adding water to the reaction as
opposed to adding aixed-molar NaOH solution to the reaction,
FIG. 20 is a partial cross-section view of an apparatus to produce hydrogen
gas, embodying some of the preferred conditions to carry out the
process according to the present invention;
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DETAILED DESCRIPTION OF THE PREFERRED EMBODIMENT
In the present invention it is believed that aluminum reacts with water
under certain conditions in the presence of sodium hydroxide as a catalyst.
It is believed that the reaction is carried out according to the equation (1),
as follows;
2A1 + 3H20 ~ A1203 +3Hz (1)
catalyst = NaOH
In one of the most pertinent prior art documents, the US Patent 934,036, it
is taught that aluminum reacts with water and sodium hydroxide according
to one of the following formulas;
2Al + 2NaOH + xH2O ~ Na2A12O4+ xI-I20 + 3H2 (2)
2A1 + 6NaOH + xHzO ~ Na6A12O6+ xH2O + 3H2 (3)
In other relevant prior art documents, Stockburger and Belitskus teach that
aluminum reacts with an alkaline solution of NaOH and that NaOH is
subsequently regenerated by the precipitation of Al(OH)3 as described by
the formulas (4) and (5) respectively.
2Al + 2NaOH + 6H20 =* 2NaAl(OH)4+ 3H2 (4)
2NaA1(OH)4 =~' 2NaOH + 2A1(OH)3 (5)
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It is also taught in Stockburger, that an optimum concentration of the
alkaline solution should be maintained at 5.75 M NaOH for an acceptable
reaction rate, and in Belitskus that the rate of precipitation of Al(OH)3 and
the regeneration of NaOH is insufficient to support rapid reaction rates.
The following experiments were carried out to demonstrate that under
certain conditions, the sodium hydroxide is not consumed in the reaction
but acts as a catalyst to the reaction as described in equation (1).
Experiments 1-1 to 1-8
A first series of eight experiments was carried out to measure the volume
of hydrogen gas produced in a typical reaction. In these experiments,
aluminum foil from Reynolds Aluminum Company of Canada was loosely
crumpled and placed in a one litre plastic bottle containing 500 ml of
catalytic solution of about 4.5M NaOH. The bottle was quickly capped
with a cover fitted with a tube which led to an inverted volumetric cylinder
filled with water. The bottle was immersed in a water bath to prevent
overheating.
The volume of water displaced by the gas produced was measured and
corrected to a gas volume at standard temperature and pressure (STP).
Atmospheric pressure on that day was obtained from a local weather office.
The corrected volume of gas produced was compared to the theoretical
quantity of hydrogen gas, which would be obtained according to the
equation,
2Al + 3H20 A1203 +3H2 (1)
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These experiments were carried out at a room temperature of 21 C and an
atmospheric pressure of 758 mm of Hg. In all cases the reaction started in
few seconds and continued for few minutes, until depletion of the
aluminum foil. It was noticed that a typical reaction with less than 5 grams
of loosely crumpled aluminum foil, is complete in less than 5 minutes. The
results of these experiments are shown in Table 1 below.
Table 1: Hydrogen Gas Production from Aluminum Foil
Exp. Al H2 H2 (1) H2 (1) Yield Deviation
(#) (g.) (1) (STP) Theoretical (%) (+/- %)
1-1 2.08 2.94 2.71 2.59 104 2.6
1-2 2.03 2.85 2.62 2.53 104 2.6
1-3 2.21 3.05 2.81 2.75 102 2.5
1-4 2.16 2.9 2.67 2.69 99 2.6
1-5 2.2 3.04 2.8 2.74 102 2.5
1-6 2.21 3.04 2.8 2.76 102 2.5
1-7 0.73 1.03 0.94 0.91 103 2.4
1-8 0.83 1.15 1.05 1.03 102 2.2
Ave. 102 2.47
The results from Table 1 show that the reaction is reproducible and
produces stoichiometric quantities of hydrogen gas. The 102% average
yield of hydrogen gas is considered to be within the measurement
uncertainty; however, there are at least two factors which might have
contributed to a slightly higher hydrogen yield. Firstly, the volume of gas
produced was corrected to STP. It is possible that the exhausted fume hood
in which the experiments were carried out could have lowered the reaction
pressure below the atmospheric pressure of 758 mm of Hg. This would
have increased the observed value for the volume of gas produced. An
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exhaust bench typically runs at 1 inch or 2 inches of water pressure. At a
maximum, this could have increased the measured volume by about 0.5%.
Secondly, the water used was tap water in all cases, in which dissolved air
may have been present. If any of this air had been released in the presence
of the warm hydrogen gas, this would have increased the volume of gas
measured. This would have affected the results by less than 1%. Since the
results are within the measurement error, and quantification of these two
sources of error would not significantly affect the results, no further
experiment was carried out in this area.
Experiments 1-9 and 1-10
The procedure used in the above experiments was repeated, with the
exception that the tube leading from the top of the reaction bottle was
connected to a gas sampling bag. Two samples of gas were obtained and
analysed. The results are presented in Table 2.
Table 2: Gas Analysis
Sample Hydrogen Oxygen &
Concentration Nitrogen
1 (1-9) 92 % balance
2 (1-10) 98 % balance
Table 2 shows that the purity of the hydrogen collected in the second
sample was 98%. This is close to what was theoretically expected. The
lower 92% concentration observed in the first sample was probably due to
the fact the system was not completely purged with hydrogen before the
sample was taken. By the time the second sample was taken, most of the
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air had been purged from the tube and the reaction bottle.
Experiment 1-11
The procedure used in the first mentioned experiments was repeated except
that the reaction bottle was placed in a water bath before the aluminum was
added to the water, and the hydrogen produced was bubbled through the
bath water. The temperature of the bath and the catalytic solution were
measured before and after the reaction, and at about four minutes after the
reaction was completed.
The water equivalent of the plastic containers for absorbing heat and their
specific heat were determined experimentally by adding a known quantity
of hot water to the reaction system at room temperature and then
calculating the heat transfer based on the final temperature.
The quantity of heat produced by the reaction was determined and
compared with the theoretical values. The results are shown in Table 3.
Table 3: Heat of the Reaction
Readings Temp. C Temp. C Temp. C Temp. C Time
Reactor Bath Reactor Bath
Start Start Finish Finish
1(1-11) 21.1 20.2 45.5 24.4 5.29
2(1-11) 21.1 20.2 38.3 25.3 5.33
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Table 3: Continued
Readings Al Heats of Heat of Heat Heat Efficiency
(g) Formation Formation Output Output (%)
A1203 H20 Actual Theoretical
kcal/mole kcal/mole kcal kcal
1(1-11) 9.52 -400.5 -68.3 33.3 34.5 96
2(1-11) 9.52 -400.5 -68.3 32.5 34.5 94
The results in Table 3 show that the observed heat released in the
production of hydrogen was 96% of the theoretical value. The 94% value
from the second reading can be attributed to the heat lost to the
surroundings during the time that lapsed between the readings.
The reaction has a net maximum heat production during hydrogen
generation of 195.6 kCal/mole. A further 204.9 kCal/mole will be released
if the hydrogen is burned with oxygen. Stated another way, 51 % of the
reaction energy is used to form hydrogen gas and 49% goes into the
production of heat.
Experiment 2-1: With 5.00 M NaOH Alkaline Solution
Sodium hydroxide (NaOH) pellets (40.63 g) from Wiler Fine
Chemicals were placed in a two litre Erlenmeyer flask. Tap water (200 ml)
was added to the flask. The mixture was swirled and allowed to stand on
the lab bench. The lab temperature was 25 C. After about an hour,
aluminum (Al) foil (30.72 g) was added in two portions. The first addition
of aluminum is referred to as time zero, the start of the reaction. The
temperature of the vapour coming from the top of the flask was measured
using a thermometer and was found to be 93 C four minutes after the first
half of the Al had been added. The flask was open to the atmosphere. The
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reaction was carried out for a period of 130 minutes. Additional quantities
of Al and water were added at regular intervals, and the temperature was
observed and recorded. The flask was swirled periodically to ensure the
solution was in contact with the Al. No further NaOH was added.
During this first experiment, a total amount of 98.7 g of aluminum was
added, and 650 ml of water was added to the initial volume. A graphic
illustration of this Experiment 2-1 is shown in FIG. 1. In this illustration,
the heavy curve labelled as `T' indicates the temperature of the vapour
coming out of the flask; the medium density curve labelled `A' indicates
the amount of aluminum added; and the lighter curve `W' indicates the
water added. The same labelling is used for all the experiments illustrated
herein.
The addition of Al to the NaOH solution resulted in the production of
vapour which issued from the neck of the flask at temperatures above 90 C.
Furthermore, this production of hot vapour started within a few minutes
(less than 4 minutes) of the Al being added. The reaction proceeded
vigorously with the addition of each charge of Al. Furthermore, even when
there was a delay between charges, such as at the 36 minute and 44 minute
additions, the reaction proceeded. Indeed, even when the addition of Al
was delayed for 41 minutes and the reaction mixture had been allowed to
cool, the reaction still proceeded vigorously (at about 128 minutes) when
Al was added and the mixture was swirled.
It is to be noted that the amount of aluminum consumed in this reaction is
about 3.6 times the amount predicted by the formulas (2) and (4), and about
10.8 times the amount predicted by the equation (3). These findings
confirm the catalytic nature of the reaction according to the present
invention.
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Experiment 2-2: With 4.95 M NaOH Alkaline Solution.
To tap water (100 ml) in a one liter suction flask was added NaOH pellets
(20.12 g) from Willer Fine Chemicals. The mixture was swirled to aid
solvation. The lab temperature was 23 C. Two thermocouples were
inserted through the suction inlet on the flask. The flask was open to the
atmosphere. Thermocouple 1(TC1) was placed in the NaOH solution
about one centimetre from the bottom of the flask. The junction of
thermocouple 2 (TC2) was placed in the flask neck at the same level as the
suction inlet. The thermocouples were read by a Scimetric System 200 data
recorder which stored the temperature readings at five second intervals.
After about half an hour TC1 and TC2 read 31 C and 21 C respectively.
After 53 minutes, TC 1 and TC2 read 26 C and 22 C, respectively, and
heavy duty Al foil (4.90 g) from Alcan Aluminum Limited was added to
the solution. There was vigorous reaction. This is referred to as time zero,
the start of the reaction.
After four minutes, water (35 ml) was added to the solution. Al foil and
water were added at five minutes intervals. The flask was swirled
periodically. No further NaOH was added. FIG. 2 illustrates the response
of this reaction in Experiment 2-2.
The reaction was monitored for about 90 minutes. During that period a
total of 53.32 grams of aluminum was consumed and 350 ml of water was
added to the initial quantity. The quantity of aluminum added corresponds
to about 3.9 times the amount predicted by formulas (2) and (4), and about
11.8 times the amount predicted by equation (3). Again this confirms that
the reaction according to the present invention proceeds according to
equation (1).
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In this experiment, there was no occasion in which Al was added that the
temperature of the vapour being emitted did not increase two to three
minutes of the Al being added. A sharp drop in the temperature was
observed about one minute after the addition of water. This is to be
expected since the water was at room temperature (- 23 C) and it was
poured in through the top of the flask. Thus, it would cool the system
momentarily.
In both Experiments 2-1 and 2-2, there was no indication that the reaction
would not have proceeded indefinitely if more Al had been added. The
regular addition of Al and the fact that the temperature of the vapour
remained above 80 C, except when water was added, indicate that the
reaction proceeded directly and no time was there a pause in the reaction
to permit the regeneration of any reagent species as predicted by the
formula (5).
The following Experiments 2-3 and 2-4 were carried out at temperature of
45 C or less to determine whether the reaction would be sustainable at
these temperatures. The composition of the precipitate forming in the
reaction at these temperatures, as well as the composition of the gases
emitted were also analysed.
Experiment 2-3:
Collection of Precipitate at an Early Stage of the Reaction
Sodium hydroxide (NaOH) pellets (39.92 g) from Wiler Fine Chemicals,
Lot # 14449, were placed in a two litre Erlenmeyer flask. Tap water
(182m1) was added to the flask. The mixture was swirled and allowed to
stand on the lab bench over night. Then it was swirled again to dissolve the
remaining NaOH and mix the solution. The solution was then transferred
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to a 400 ml beaker.
Commercial aluminum foil (24.23 g), namely Reynolds WrapTM, a
Registered Trade Mark of Canadian Reynolds Metal Company, Ltd., was
weighed, folded, and cut into portions that ranged in weight from 0.5 g to
1.5 g. The beaker containing the NaOH solution was placed in a water bath
which was cooled with ice cubes. Thermometers were placed in both the
water bath and the beaker. Care was taken to ensure the temperature of the
solution in the beaker was kept at or below about 45 C. In each case, the
temperatures of the solutions were read and recorded just before the
portions of Al were added. At 29, 44, and 50 minute the reaction beaker
was removed from the water bath to try to keep the reaction temperature as
close to 45 C as possible.
The Al foil was added in portions over a 59 minute period. When the Al
was added to the initial reaction mixture, gas bubbles were observed to
form after about 45 seconds. It was noted that when gas bubbles formed
on the surface of the Al, the piece of Al floated at or near the top of the.
reaction mixture. A small amount of fine black material was observed to
float in the reaction mixture after all the Al has been dissolved. By the 44
th
minute, the reaction mixture was observed to be very viscous because of
the formation of a solid material. At 50 minute, tap water (30ml) was
added to the reaction mixture. This Experiment 2-3 is explained
graphically in FIG. 3.
The solution was allowed to cool to room temperature, then it was filtered
through a porcelain suction funnel without any filter paper to ensure there
was no un-reacted Al present in the reaction mixture which could distort
the analysis of the precipitate. The cloudy, grey, viscous material which
passed through the funnel was filtered using a paper towel. It was washed
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with tap water. The final precipitate, a light grey solid, was allowed to
stand in the fume hood overnight, then a portion of it, labelled P3-1, was
dried in an oven at 102 C. After 45 minutes the sample was sealed in a
plastic bag and taken to an Electron Microscopy Unit for analysis. The
results of this analysis are presented in Table 4.
Experiment 2-4:
Collection of Precipitate at an Advanced Stage of the Reaction
A 400 ml beaker containing tap water (175 ml) in which NaOH pellets
(38.11 g) had been dissolved was placed in a water bath which was cooled
by ice cubes. Al foil (39.26 g) was weighed, folded and cut into portions
ranging up to 2 g. Both the NaOH and the Al came from the same source
as described in Experiment 2-3. The Al foil was added to the NaOH
solution over a 148 minute period following the same procedure as in
Experiment 2-3. Additional water, totalling 70 ml, was added in four
portions during the Experiment 2-4 at the times shown in FIG. 4.
The reaction mixture was allowed to stand in the fume hood for about two
hours after the addition of the last portion of Al, by which time the mixture
had stopped bubbling. Part of the mixture was then filtered through a fine
plastic mesh to ensure no un-reacted Al could contaminate the sample to be
analysed. The mixture which passed through the mesh was then filtered by
suction using qualitative filter paper. A sample of this grey precipitate was
taken without washing and labelled P4-1. The remainder of the precipitate
was removed from the filter paper and swirled with tap water in a flask,
then it was re-filtered and washed with tap water.
A sample of the washed precipitate was taken and labelled P4-2. Both
samples were dried in an oven at 102 C for about an hour then they were
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sealed in a plastic bag and taken to the Electron Microscopy Unit for
analysis. The results of the analysis are given in Tables 5 and 6.
The samples taken from Experiments 2-3 and 2-4 were analysed using a
JEOL-6400 Scanning Electron Microscope (SEM) equipped with a Link
eXL x-ray microanalyser. An accelerating voltage of 15 kV and a probe
current of 1.5 nA were employed, and spectral collection times were 200s
for sample P3-1 and 120s for samples P4-1 and P4-2. The results are
reported as oxide weight percent values, although oxygen was not analysed.
Oxide values were calculated from elemental analyses using specified
oxide stoichiometries. The minimum detection limits for NaOH under
these conditions are approximately 0.38 wt.% for sample P3-1 and 0.5
wt.% for samples P4-1 and P4-2.
Table 4: Sample P3-1.
SiO2 n.d. n.d. 0.23 0.32 0.18
TiOZ n.d. n.d. n.d. n.d. n.d.
A1203 59.71 67.63 80.08 57.50 70.11
FeO 0.27 0.28 0.33 0.30 0.44
MnO n.d n.d. n.d. n.d. n.d.
MgO n.d. n.d. n.d. n.d. n.d.
CaO 0.26 0.28 0.35 0.41 0.18
Na20 0.39 n.d. n.d. n.d. n.d.
K2O n.d. n.d. n.d. n.d. n.d.
CuO 0.38 0.46 0.42 0.70 0.41
Total 61.01 68.65 81.41 59.23 71.32
n.d. = not detected
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Table 5: Sample P4-1.
Si02 0.21 0.27 n.d. n.d. n.d.
Ti02 n.d. n.d. n.d. n.d. n.d.
A1203 63.05 54.87 62.40 63.02 74.57
FeO 0.26 0.26 n.d 0.29 0.32
MnO n.d n.d. n.d. n.d. n.d.
MgO n.d. n.d. n.d. n.d. n.d.
CaO n.d. n.d. n.d. n.d. n.d.
Na2O 8.03 11.21 9.46 4.02 4.25
K20 n.d. n.d. n.d. n.d. n.d.
CuO n.d. n.d. n.d. 0.37 n.d.
Total 71.55 66.61 71.86 67.70 79.14
n.d.= not detected
Table 6: Sample P4-2.
Si02 n.d. 0.29 n.d. n.d. 0.22
Ti02 n.d. n.d. n.d. n.d. n.d.
A1203 70.96 72.01 63.77 69.72 65.80
FeO 0.30 0.35 0.32 0.23 0.28
MnO n.d n.d. n.d. n.d. n.d.
MgO n.d. n.d. n.d. n.d. n.d.
CaO 0.18 0.10 0.16 n.d. 0.14
Na2O n.d. 0.69 n.d. n.d. n.d.
K20 n.d. n.d. n.d. n.d. n.d.
CuO 0.42 0.37 0.30 0.42 0.51
Total 71.86 73.81 64.55 70.37 66.95
n.d.= not detected
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The results presented in Tables 4-6 show the precipitate formed does not
contain sodium beyond what could reasonably be expected to be present in
an impure material precipitated from a concentrated NaOH solution. In no
case was the quantity of sodium in the precipitate present in amounts
exceeding 1.1 % of that required by the reaction products specified in
equation (2), (3) or (4). Therefore it may be concluded that the precipitate
formed is not Na/Al moiety, but is rather primarily an Al/Oxygen material,
which may contain some hydrogen in the form of hydroxyl groups or water
molecules.
The two samples collected and analysed in Experiment 2-4 show two
things, namely, the washing of the precipitate with water removes
significant amounts of Na; and that none of the five measurements on the
unwashed precipitate showed levels of Na which exceeded more than 34%
of that necessary to form the compounds given in equation (2), (3) or (4).
Indeed, the average sodium content of the five measurements was less than
one-fifth of that necessary to form the compounds given in equation (2),
(3), or (4). This removes any possibility that the Na/Al substances as
shown in equation (2), (3) or (4) was at one time present in the reaction
precipitate and was subsequently changed to an aluminum/oxygen species
by washing. If such were the case the Na:Al ratio from sample P4-1 would
have had to be at least 1:1. This was not observed. Therefore, it may be
concluded that even in the unwashed state the precipitate is primarily an
aluminum based compound.
The fact that washing with water readily removes most of the sodium
confirms that the sodium species present is water soluble as would be
expected for an ionic species containing sodium.
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Experiment 2-5: Activeness of the Filtrate
To a small amount (-50 ml) of the filtrate from the first filtration in
Experiment 2-4, was added Al foil (0.5 g). Within, about 60 seconds,
bubbling started and the Al completely dissolved, and a grey precipitate
formed in this previously clear solution.
Experiment 2-6: Collection of Gases.
To tap water (182 ml) in a four litre plastic bottle was added NaOH pellets
(40.15 g). The bottle was covered, shaken and the solution allowed to
come to room temperature after the NaOH had dissoaved. The bottle was
then placed in a water bath at 18 C. Al foil (14.8 g) was added in three
portions of about 5 g each. Both the Al and NaOH came from the same
source as described in Experiment 2-3. After the first portion of Al (4.63
g) was added, the bottle was capped with a lid fitted with a hose. Bubbles
started to form on the surface of the Al after about 10 seconds. Bubbles
came out of the hose, which was submerged in the water bath, after about
40 seconds. The Al had completely reacted within about three minutes.
The lid was removed from the bottle and a second portion of Al foil (4.98
g) was added and the bottle recapped. Bubbling from the hose started after
about 30 seconds, the hose was connected to a gas sampling bag and
sample P6-1 was collected. The addition of Al foil (5.19 g) was repeated
and gas sample P6-2 was collected. Both gas samples were analysed. The
analytic data and the normalized results are summarized in Table 7.
CA 02414135 2006-11-01
Table 7: Gas Analysis.
Observed Concentrations Normalized
Concentrations
Sample # P6-1 P6-2 P6-i P6-2
Oxygen 2% 1% 2% 1%
Nitrogen 7% 2% 7% 2%
Hydrogen 86% 92% 91% 97%
Total 95% 95% 100% 100%
Experiments 3-1 to 3-15
A series of fifteen experiments was carried out using NaOH concentrations
which ranged from about 0.25M to a saturated solution of NaOH in water
at room temperature. The saturated solution was about 19M. Thirteen of
these experiments were recorded on graphs, and are shown in the
accompanying FIGS. 5-17. On these graphs, the labels `T', `A', and `W'
designate the temperature of the reaction, and the aluminum and water
added respectively as in the previous graphs. The label `S' has been added,
however. The line `S' across each graph designates the amount of
aluminum that would react with the initial amount of NaOH if the reaction
would proceed according to the equation (2), (3) or (4). This amount is also
referred to herein as the stoichiometric amount of aluminum.
Solutions ofNaOH were typically cooled before starting the reactions. The
starting temperature for each reaction was often in the range of 4-10 C. The
reactions were carried out in glass vessels ranging in size from 25 ml to 500
ml. Solutions of NaOH were prepared by dissolving NaOH pellets from
BDH Inc., Toronto, Ontario, Canada, M8Z 1K5, Lot #128142-125228, in
tap water at room temperature. The heat of solvation was allowed to
26
CA 02414135 2006-11-01
dissipate and the portion of the solution to be used in the experiment was
cooled in an ice bath in the reaction vessel.
A thermocouple junction was placed in the solution about one centimetre
below the surface. The thermocouple reading was monitored continually
and recorded on a computer file every 15 seconds.
Aluminum foil (Reynolds Wrap from Canadian Reynolds Metals Company
Ltd., Montreal, Toronto, Calgary, Canada) was crumpled or folded and
added in portions ranging from 0.2 g to 1.1 g. Each portion of Al foil was
initially submerged in the solution using a glass stirring rod. Then it was
allowed to float to the top of the solution. The start time for every
experiment was the time when the first aluminum was added. Aluminum
was added in amounts to keep the temperature above 60 C.
Water was added in amounts up to 20 ml. Water was only added when the
reaction mixture became viscous and foamed more than one centimetre. In
most of the experiments, the addition of water started after about 75% of
the stoichiometric amount of A 1 was added. Water was added in sufficient
quantities to ensure that the level of the solution was at least one
centimetre
above the level of the precipitate. In most cases, water was added only
after the temperature had reached a peak or a value of at least 75 C. The
portions of water were also controlled so that the temperature of the top of
the solution did not drop more than 60 C when the water was added.
Aluminum and water were added until at least two times the stoichiometric
amount, based on equation (2), had been reached.
After the reaction had ceased the solution was cooled and the precipitate
was suction-filtered, and rinsed while still in the suction funnel with about
250 ml of tap water. Samples from the Experiments 3-1 to 3-15 were sent
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for elemental analysis of the precipitate and the hydrogen gas. The results
of these analyses are shown in Table 8.
Table 8. Catalytic Ratios and Product Analysis.
Test (#) [NaOHI Catalytic [A1203] [NaZO] H2
(M) Ratio (%) M (%)
3-1 0.26 3.0 98.3 <0.71
3-2 0.60 3.1 98.9 <0.71
3-3 1.2 4.2 96.3 1.14
3-4 2.5 3.3 98.7 0.7
3-5 3.9 3.9 96.8 <0.71
3-6 4.8 3.4
3-7 5.5 4.5 98.6 <0.71
3-8 6.0 2.6
3-9 6.0 3.3 98.6 <0.71 97
3-10 6.1 4.2
3-11 6.1 3.2 99.3 <0.71
3-12 6.1 3.8 98.3 <0.71
3-13 6.7 3.3 99.1 <0.71
3-14 11.3 2.7 97.3 <0.71 98
3-15 19 2.7 99.1 0.79 97
The expression "catalytic ratio" in the above table is calculated by dividing
the amount of Al that actually reacted by the amount that would have
reacted if the reaction were stoichiometric with respect to NaOH as in
equation (2), (3) or (4).
Table 8 also shows the results of the analyses of the precipitates filtered
from twelve of the experiments. In every case the concentration of the A 1
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species is larger than 96%. Sodium was detectable in only three of the
samples, and then at a maximum concentration of only 1.14% or less.
Thus, aluminum is present in the precipitate at levels that are two orders of
magnitude above sodium.
It may be concluded that the reaction according to the present invention is
catalytic in aqueous solutions from 0.26 M NaOH to 19 M NaOH. It
should be noted that although the 0.26 M and 0.60 M solutions showed a
catalytic reaction, the reaction temperature did not rise above 30 C during
those experiments. However, FIG. 5 shows that the temperature of the 1.2
M solution rose above 45 C even though the A 1 was added very slowly and
only after the previous portion had dissolved.
The results in FIGS. 5-17 show that the reaction can and does occur over
a temperature range from 4 C to 165 C. In one experiment with the
saturated solution a temperature of 170 C was observed. The molal boiling
point elevation constant will result in a higher boiling point for the more
concentrated solutions, ensuring that water does not boil offuntil the higher
boiling point is reached. In the case of the saturated solution from
Experiment 3-15, the boiling point elevation would have contributed to the
high boiling point of the solution. It was also noted that NaOH did not
precipitate from the solution even at the higher concentration, probably
because of the known higher solubility of NaOH in hot aqueous solutions.
It was found that at about 75% of the stoichiometric amount the solution
would become viscous and foaming with large longer-lasting bubbles.
Water was added at this point and often the addition of A 1 had to be slowed
down or an excess of un-reacted aluminum could be observed.
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The formation of a greyish-white precipitate would start between 75% and
100% of the stoichiometric amount. Once the precipitate started to form
it was necessary to keep the reaction zone above the precipitate a distance
of about 1 cm, or the precipitate would mix with the bubbling aluminum
and form a more viscous foam which on occasion overflowed the reaction
vessel.
Based on all the experiments described herein, it will be appreciated that
the present process to produce hydrogen is reproducible with aqueous
solutions from 1.2 M NaOH to 19 M NaOH and over a temperature range
from 4 C to greater than 170 C. Furthermore, the reaction is catalytic over
the same temperature range and over a NaOH conceritration range of 0.26
M to above 19 M. The reaction's by-product comprises high-purity
alumina (A1203).
Referring now to the graph in FIG. 18, there is shown therein a first curve.
30 showing the maximum temperatures obtained with different NaOH
concentrations. This best-fit curve was plotted from the data shown in
FIGS. 5-17, and is presented herein for illustrating the effect of NaOH
concentration on the maximum temperature of the reaction. FIG. 18
shows another curve 32 which represents the responsiveness ofthe reaction
to aluminum and water additions. This curve has been prepared by plotting
the time required to reach the initial maximum temperature of the reactions,
against the different NaOH concentrations studied. The resulting best-fit
curve is a complex inverted hyperbolic curve centred on a concentration of
about 8 M NaOH. This curve indicates that the reaction is highly
responsive to fuel additions, when the NaOH concentration is between
about 5 M and 10 M, and that the responsiveness decreases rapidly when
the NaOH concentration is adjusted away from this median region.
CA 02414135 2006-11-01
Referring now the FIG. 19, there is illustrated therein two curves. The first
curve 34 represents the effect of adding plain water to the catalytic reaction
of equation (1). As the reaction proceeds, the water is consumed, and
therefore, the concentration of NaOH increases, as shown by the segment
36, from its initial concentration 38. When water is added, as indicated by
segment 40, the concentration drops back to or below the initial
concentration 38. If water is added in portions to maintain a certain level
in a reaction vessel for example, the solution concentration fluctuates up
and down from the initial concentration 38, as generally represented by the
curve 34.
If someone is led to believe that the reaction proceeds as in equation (2),
(3) or (4), that person would logically add NaOH into the reaction vessel
with the makeup water. If NaOH is added to a reaction that actually
proceeds according to equation (1), however, the resulting NaOH
concentration of the aqueous solution in the reaction vessel would increase
as represented by curve 44. Whether the NaOH is added alone or in a
fixed-molar NaOH solution, as represented by segment 46, the NaOH
concentration of the solution in the reaction vessel would move quickly
toward saturation.
Reference is made again to the curve 32 in FIG. 18. It will be appreciated
that a regular addition of a fixed-molar NaOH solution to a reaction that
proceeds according to equation (1) would cause the responsiveness of the
reaction to move along the curve 32 as indicated by the series of arrows 48,
and quickly reach a region of very low responsiveness. Such migration of
the NaOH concentration toward a region of low responsiveness would
cause the reaction to cease or to appear to have ceased. The addition of
plain water, however, as taught herein, causes the responsiveness of the
reaction (1) to oscillate back and forth along the curve 32 toward and away
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from a more reactive state, as shown by arrows 50 and 52. These
oscillations 50, 52 are believe to stimulate the reaction, and to contribute
to some degrees to the catalytic feature of the reaction according to the
present invention.
The arrows 50, 52 and the corresponding theory explain the facts that in
some experiments, a water addition has caused the reaction to slow down,
according to the arrow 50, and in other experiments, the addition of water
caused an immediate response, as in 52. The same theory explains why
both events can occur in a same experiment, such as when the NaOH
concentration is maintained substantially in the median region, between 5
and 10 M NaOH. The curve 44 and arrows 48 on the other hand, explain
why prior inventors may have failed to observe a catalytic reaction with the
same elements.
It has been found that the reaction proceeds better when water is added
after an initial amount of aluminum has been consumed. This phenomenon
can also be explained using the curve 32 in FIG. 18. In a low
concentration solution, any delay in adding water causes the NaOH
concentration to move toward a highly responsive state, such as around 8M
for example. An addition of water at that time and a subsequent addition
of makeup water causes the NaOH concentration to oscillate within this
highly responsive region. On the other hand, if the initial concentration is
above 8 M for example, an addition of water brings the concentration back
to a highly responsive state, and therefore immediate results can be
observed.
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Additional Experiments
Additional experiments were carried out using aluminum wire of different
gauge sizes and aluminum flakes from the helical casing of armoured
electrical wire. Although these additional experiments were not recorded
in details, the catalytic effect was observed. Therefore, it is believed that
the reaction (1) is reproducible with aluminum flakes from beverage cans
and food packages, aluminum chips, shavings and sawdust found in
machine shop waste, and aluminum powder available commercially for
different purposes including fireworks, or other small aluminum particles
of the like. It is to be expected that the intensity of the reaction depends
upon the surface of contact between the aluminum and water. Aluminum
foil for example reacts faster than a heavy gauge aluminum wire, and
aluminum powder would react almost instantly to produce hydrogen gas.
Preferred Apparatus
A preferred hydrogen generator 60 is illustrated in FIG. 20. The hydrogen
generator 60, comprises a reaction vessel 62 made of non-corrosive
material, in which the reaction is carried out. A minimum amount of an
alkaline solution 64 is maintained in this vessel. During the operation of
the generator, it has been found that aluminum particles reacts with water
at the surface 66 of the alkaline solution and defines at and near the surface
66, a region of substantial effervescence. This region is defined as the
reaction zone `F'. The height of the reaction zone `F' vary with the
intensity of the reaction, and extends above and below the surface 66 of the
alkaline solution 64. During the operation of the generator, a precipitate 68
accumulates at the bottom of the reaction vessel 62. It is recommended to
maintain the reaction zone `F' at a height `H' of at least about 1 cm above
the precipitate 68, to prevent the precipitate from swirling into the reaction
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CA 02414135 2002-12-12
zone and mixing with the aluminum particles. Although this dimension can
be reduced in some installations, a dimension of one centimetre is
suggested herein to enable those skilled in the art to readily use the process
according to the present invention successfully.
A water bottle 70 is affixed to the side of the reaction vessel 62 and has a
piping system 72 connected to an array of nozzles 74 in the bottom of the
reaction vessel 62. Only one nozzle is shown for clarity. The introduction
of water through the bottom of the vessel 62 has the effect of capturing
some of the heat in the precipitate 68 to preheat the water entering the
reaction vessel. A second purpose for the feeding of water through the
bottom of the reaction vessel 62 is to entrain to the reaction zone `F', any
sodium hydroxide which may be present in the precipitate 68.
While the distance `H' of the reaction zone `F' above the precipitate 68
defines a low limit to the water content in the reaction vessel, the upper
limit should be defined as to maintain the concentration of the alkaline
solution over about 1 M NaOH, and more preferably, a concentration of 5M
NaOH. A sight glass 76 on the side of the reaction vessel 62 is provided
to monitor the minimum distance `H' of the reaction zone `F' above the
precipitate 68.
Aluminum particles 78 are delivered into the reaction vessel 62 from a
hopper 80 mounted on the top of the vessel 62, though an airlockTM rotary
feeder 84 and through a drop pipe 86 at the center of the reaction vessel 62.
A deflector 88 is mounted at the end of the drop pipe 86 to disperse the
aluminum particles 78 over the entire surface of the alkaline solution 64.
The hydrogen generated in the reaction vessel exits through the drop pipe
86 and the spout 90.
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The drop pipe 86 is preferably mounted through a large openable cap 92 on
the top of the reaction vessel. This cap 92 preferably covers a substantial
portion of the upper end of the reaction vessel 62 and provides access to the
reaction vessel for periodically cleaning the vessel. A bung 94 is provided
in the bottom surface of the reaction vessel 62 to recover the precipitate 68.
The aluminum particles 78 are preferably flakes, sawdust, milling shavings
and chips, powder or other similar small particles having a large surface
over volume ratio. It has been noticed that aluminum foil fragments for
example, have a tendency to float at the surface 66 of the alkaline solution
64. This is preferable and is explained by the buoyancy created by the foam
96 and the bubbling action generate in the reaction zone `F'. It is believed
that the bubbling action and the high temperature in this reaction zone is
ideal to prevent or reduce the formation of a protective oxide layer on the
surface of the aluminum particles. It is believed that the retention of the
aluminum particles in this reaction zone contributes largely to maintaining
the catalytic effect.
When relatively dense aluminum particles are used, it is recommended to
install a floating screen 98 at the surface of the alkaline solution 64, to
retain the aluminum particles in the reaction zone `F'.
As to other manner of usage and operation of the process according to the
present invention, the same should be apparent from the above description
and accompanying drawings, and accordingly further discussion relative to
these aspects is deemed unnecessary.
