LIFE SUPPORT MATERIALS IN A MANNED SUBMERSIBLE
T. C. Wang
Harbor Branch Foundation, Inc. RR 1, Box 196 Fort Pierce, Florida 33450 February, 1982
HARBOR BRANCH TECHNICAL REPORT NUMBER 45 LIFE SUPPORT MATERIALS IN A MANNED SUBMERSIBLE
T. C. Wang
Harbor Branch Foundation, Inc. RR 1, Box 196 Fort Pierce, Florida 33450 February, 1982
HARBOR BRANCH TECHNICAL REPORT NUMBER 45 Table of Contents
CARBON DIOXIDE REMOVAL
I. Alkali Metal Hydroxide
1. General review
2. Lithium hydroxide
3. Sodalime
4. Baralyme
5. Discussion - water vapor, temperature, humidity, porosity, residence time.
II. Alkali Metal Superoxides, Peroxides
1. Sodium or potassium superoxide
2. Alkali metal peroxides
III. Regenerative Type of CO2 Absorbent 1. Molecular sieves
2. Metallic oxides
B. Silver oxide
3. Solid amines
4. Liquid amines 5. Carbon dioxide reduction
A. Sabatier process
B. Bosch reaction 6. Miscellaneous
A. Membrane separation
B. Freeze-out OXYGEN SOURCES
I. High Pressure Gas Storage
II. Crogenic Storage
III. Chemical Generation
1. Chlorate candle
2. Alkali metal superoxides and peroxides
3. Decomposition of hydrogen peroxide
4. Miscellaneous
A. Oxygen generation
B. Photosynthesis
TRACE CONTAMINANTS REMOVAL
I. Contaminant Sorption by Solids
1. Activated charcoal
2. Molecular sieve
3. Silica gel, activated alumina
II. Chemical Conversion of Contaminants
1. Hopaclite
2. Purafil
GAS PROPERTIES
I. Carbon Dioxide
II. Nitrogen
III. Oxygen
IV. Air
V. Helium
VI. Carbon Monoxide
REFERENCES
ii CARBON DIOXIDE REMOVAL
I. Alkali Metal Hydroxide
1. General review
The oxides, hydroxides, peroxides and higher oxide states of the alkali and alkaline earth metals react with carbon dioxide to form carbonate, bicarbonate or hydrates of carbonates and bicarbonates. The theoretical capacities of the various possible CO 2 sorbents is shown in Table 1.
Lithium monoxide has a greater CO 2 capacity than any of the other oxides. The use of lithium oxide as a carbon dioxide absorbent has been investigated for short duration space flights. Its dynamic absorption characteristics are not as good as those of lithium hydroxide. Alkali metal hydroxides other than LiOH have a high capacity for CO 2, However, their high degree of hygroscopicity eliminates them from consideration as solid absorbents.
Alkaline earth metal oxides such as MgO and CaO absorb carbon dioxide very slowly, while the addition of water vapor and conversion of the oxide to hydroxide increases the rate of absorption. Dedman and Owen (1962), in studying the reaction of CaO with CO 2, showed an initial rapid uptake due to chemisorption, followed by a slower reaction rate controlled by the diffusion of CO 2 into the pores of the solid.
Comparison of alkaline earth metal hydroxides (see
Table 1) shows Mg (OH)2 to have a high capacity for CO 2 on a unit weight basis. Little work has apparently been done on
1 the absorption of CO 2 by magnesium oxide under conditions which are applicable to the removal of C02 from expired gases.
The absorption capacity of calcium hydroxide is somewhat less
than that of Mg(OH)2' but the literature suggests that the
CO 2 absorption rate of hydrated lime is very slow. This slow reaction rate has been circumvented through the use of
additives of other hydroxides to the basic lime structure.
The addition of 4.5 wt % sodium hydroxide to lime, containing
17.5% water, results in a product which absorbs CO 2 signifi cantly more rapidly than Ca{OH)2 alone. This material,
commercially known as soda lime, has been used for many years
in anesthesiology for the absorption of CO 2 . Another modification of lime is Baralyme, which is a mixture contain
ing 20 wt % Ba{OH)2 • 8H20 in lime. By far, lithium hydroxide, soda lime and Baralyme are the most effective C02 absorbents
in enclosed space.
2 • Lithiurn hydroxide
LiOR absorbs CO2 from a gas mixture and in the presence of water vapor. The reaction is given by the following equations:
2LiOH + 2H20 2LiOH •H20 ------(1) 2LiOH. H + CO Li + 3H ----(2) 20 2 2C03 20
The reaction of C02 with LiOH requires the presence of water
in an amount sufficient to produce LiOH • H20 prior to or
simultaneously with the CO2 reaction (Williams, 1969; Wang,
1979). The properties of LiOH as a CO 2 absorbent are listed below:
2 Properties of LiaR
Chemical formula LiOH
Molecular wt. 23.95
Melting point 862°F
Density of crystals 0.0918 Ib/in3
Bulk density of granular LiOH 0.0145-0.0162 Ib/in3
Theoretical CO 2 absorption capacity 0.919 Ib C02/1b LiaR
Heat of absorption (H20 gas) 875 Btu/lb C02
Water of reaction 0.409 Ib/lb CO 2 or 0.375 Ibjlb LiOR
The dynamic removal efficiency, which expresses the reduction in CO 2 concentration across the bed, varies with time in the operation of LiOH as the absorbent is used up.
The theoretical capacity of LiOH for CO2 absorption can not be achieved in an atmosphere control system because of the bed dynamics, characteristics, canister dimensions, air superficial velocity and the process air inlet temperature.
The LiOR particle size is generally kept between 4 to 8 mesh size for high utilization efficiencies on low bed pressure drops. Normally, no channeling occurs with this particle
size at gas velocities up to 1.0 ft/sec (Bentz, 1976).
The CO2 absorption capacity of LiOE varies with various temperatures and humidities (Wang, 1975). The chemical
reaction associated with water production and the rate of
dehydration affect the C02 absorption capacity (see
discussion) .
3 3. Sodasorb (Soda Lime)
Sodasorb consists essentially of hydrated lime
[Ca(OH)2] , sodium hydroxide (NaOH), potassium hydroxide (KOH) ,
14 to 19 percent moisture content, and some inert. It is in the form of granules and contains an indicator which changes from colorless or white to blue when the NaOH has reacted to form Na2C03' The chemical reaction of CO 2 with Sodasorb is as follows:
CO 2 + H20 + H2Co 3 ... (3) 2 H2C03 + 2 NaOH + 2 KOH + Na2Co3 + K2C0 3 + 4 H20 .•. (4)
Na2C03 + K2C03 + 2 Ca(OH)2 + 2 CaC03 + 2 NaOH + 2 KOH ... (5) The water is added as free moisture in the manufacturing process and is necessary for efficient removal of CO 2, Sig nificant water is lost through evaporation if properly sealed storage is not maintained and the C02 absorption capacity will be decreased. The theoretical C02 absorption capacity and physical properties of Sodasorb are shown in Tables 1 and 3.
4. Baralyme
Baralyme is a mixture of barium hydroxide octalhydrate
[Ba(OH)2 • 8H20] , calcium hydroxide [Ca(OH)2] and potassium hydroxide (KOH). The eight waters of crystallization present in the Ba(OH)2 serve to fuse the mixture into a homogeneous mass which will hold its shape and form under varying conditions of heat and moisture. The water is used as shown in the chemical equations:
4 Ba(OH)2° S H2O + CO 2 -+ BaC03 + 9 H20 ... (6)
9 H2 0 + 9 CO 2 -+ 9 H2C03 ...... (7)
9 H2 C03 + 9 Ca(OH)2 -+ 9 CaC03 + IS H20 ... (S)
2 KOH + H2C03 -+ K2C03 + 2 H20 ... (9)
Ca(OH)2 + K2C03 -+ CaC03 + 2 KOH ... (lO) When temperature is above 150°F, water has a tendency
to liberate from Baralyme granules and CO 2 absorption capacity will be decreased. Tables 1 and 3 show the properties, characteristics and CO 2 absorption capacity of Baralyme.
5. Discussion
Lithium hydroxide, Sodasorb and Baralyme are the most widely used absorbents used for CO2 removal. Experience has
shown that the CO 2 removal is related to the porosity, granule size, surface area, moisture content, scrubber design and other environmental conditions such as temperature, humidity,
pressure, etc. (Boryta, 1967; Cook, 1972; Mi Ll.er , 1979; Wang,
1975; Williams, 1969). Insufficient water vapor in the air
stream allows only partial CO 2 reactions as shown in the chemical equations, while an excessive amount of water vapor
forms a water film barrier around the absorbent granules also
resulting in an incomplete reaction between CO 2 and absorbents (Williams, 1969; v7ang, 1979, 19S0). Higher temperature
normally favors higher reaction rate and therefore has a
higher water production. Too high a temperature in the
chemical reaction produces excess amounts of water that
could decrease the available surface area for CO 2 reaction.
5 When this occurs, the CO 2 absorption capacity of absorbents decreases.
For a given CO 2 content in the alr stream at any given temperature, there is a corresponding humidity of the feed
stream which results in maximum C02 absorption efficiency.
When the relative humidity in the air stream is below 14%, anhydrous LioH has a higher CO 2 absorption capacity at higher temperatures. However, above 40% relative humidity, lower temperatures are favored as shown in Figure 1 (Wang, 1979).
The CO2 absorption capacity of Sodasorb and Baralyme at various temperatures and humidities are shown in Figures 3-5
(Wang, 1980). When temperature is less than 65°F, water content in the air stream does not significantly affect CO 2 absorption capacity. When temperature is greater than 65°F, the CO2 absorption capacity of both Sodasorb and Baralyme are increased with increased temperature (Wang, 1980).
Porosity with surface area in the granules also controls the CO 2-absorbent reaction rate (Wang, 1978, 1980). The more porous granules provide more available surface area for the
CO2 reaction with the absorbents. Water and C02 can more easily diffuse into the inner parts of the granules which R results in a higher CO 2 absorption capacity. Sodasorb Hp is about two times more porous than Sodasorb as shown in Table 3.
The CO2 absorption capacity of Sodasorb HP was found to be 2.5 times more efficient than Sodasorb granules (Wang, 1978,
1980) .
6 Residence time is another factor to affect CO 2 absorption capacity of various absorbents. Residence time
is defined as the time the gas is exposed to the absorbent material. A short residence time normally causes an incom
plete chemical reaction. Therefore, the CO 2 absorption capacity of the absorbent is decreased. On the other hand,
a long residence time requires a low scrubbing rate through
the scrubber. This will result in C02 accumulation in the enclosed space when the C02 scrubbing rate is lower than the
C02 production rate. The proper residence time and scrubbing
rate is therefore needed in order to maintain a desirable CO 2 concentration in the submersible. Generally, residence time of one second or more have been used in scrubbing design.
However, various tests have shown that residence time can
vary from 1 to 0.1 seconds with respect to different absorbents
and still operate (Bentz, 1976). Wang (1980) has shown (Fig. 6)
that lithium hydroxide has the highest CO 2 absorption capacity when the residence time is about 0.8 seconds. The C02
absorption capacity of both Sodasorb and Baralyme is increased
as residence time increases. However, when the residence time
is less than 1.0 seconds, the CO 2 absorption capacity of both absorbents also is greatly reduced.
II. Alkali Metal Superoxides and Peroxides
1. Sodium or potassium superoxide
Alkali metal superoxides such as potassium superoxide
or sodium superoxide act as both carbon dioxide absorbents
and oxygen sources. The theoretical CO 2 absorption capacity
7 of superoxides are shown in Table 2. This makes them very attractive for enclosed space such as a submersible or a space vehicle (Schneider, 1965; Keating, 1980). The use of such materials in atmospheric control is generally similar to that of alkali metal hydroxides. They are packaged as granular beds through which gas is circulated. Superoxides are highly reactive with strong oxidizing and alkaline properties, and they appear to provide for odor control and sterilization of process air. Although this factor makes such materials promising for CO 2 removal, this advantage is offset by handling problems incurred by their chemical reactivity. This material is very hygroscopic and reacts very readily with water vapor in the air or liquid water. This reaction releases oxygen and heat. With sufficient heat the combustible material could ignite and, ln the presence of the pure oxygen, a rapid fire would result.
When air is passed through a bed of K02, the moisture in the air reacts with K02 to produce oxygen and potassium hydroxide. The latter substance then absorbs carbon dioxide in the air.
The main reactions are:
2 K02 + H2 O ~ 2 KOH + 3/2 °2 ------(11)
2 KOH + CO2 ~ K2C0 3 + H2O ------(12)
KOB + CO 2 ~ KBC0 3 ------(13) The following hydrating reactions also compete with the above reactions:
8 KOH + 3/4 H2O -+ KOH' 3/4 H2O ..... (14 )
KOH + H2O -+ KOH' H2 O ...... (15)
KOH + 2 H2O -+ KOH' 2H20 ...... (16) 2 K2C03 + H2O -+ 2 (K2C03 • 1/2 H2O) (17)
2 K2C03 + 2 H2O -+ 2 (K2C0 3 • 3/2 H2O) (18) The chemical equilibria of the system are somewhat
complex. Major parameters appear to be the input rate of moisture and carbon dioxide and the temperature level sus
tained in the chemical bed. The principal reactions may be
combined into the overall system equation:
+ 3 02 + 131.82 K cal/mol... (19)
Man's respiratory quotient (RQ) is defined as the volume
ratio of CO 2 formed to oxygen consumed. The RQ for the metabolism described above is 0.82. A similar figure may be
calculated for potassium superoxide which generates three
volumes of oxygen while consuming two volumes of carbon
dioxide. It chemical RQ is 0.67. Potassium superoxide cannot
by itself counter balance man's RQ. Effective removal of CO 2 results in overproduction of oxygen, or, conversely, a
balancing oxygen production results in underabsorption of
carbon dioxide. Neither of these conditions can be tolerated
for an extended period of time in a closed system. A program to investigate the use of potassium and sodium
superoxide for oxygen control in manned space vehicles was
completed by Kunard and Rogers (1962). This study used the
canister approach, Le., passing the moist, C02-enriched air
9 through a canister of K02 granules (usually 4-6 mesh). The
results showed that there was no reaction of dry CO2 with
K02' This indicated that the CO 2 did not react directly with the K02 but only with the reaction product KOH.
With high concentrations of CO 2 and water vapor, they found that the initial oxygen production rate was high (Kunard
and Roger, 1962). As the reaction proceeded, the 02 produc
tion rate decreased while the water adsorption rate remained
essentially constant. This indicated that hydration of the
reaction products was competing with the reaction of K02. Kunard and Rogers (1962) concluded that by controlling
the inlet C02/H20 mole ratio at 2/1, the ratio of C02 uptake
to 02 release could be controlled at about 0.82, matching the
human RQ. They also found that the absolute water concentra
tion was the controlling factor in the rate of oxygen
production. A high initial water concentration in the inlet
stream resulted in a high initial rate of oxygen evolution.
This rate dropped with time, even though the water concentration
was maintained at the original high level, since the surface
was becoming coated with the reaction products. After the
K02 particles become coated, the evolution rate becomes dependent on the diffusion rate of oxygen out of, and water
into the granule.
Kunard and Rogers (1962) concluded that it is desirable
to operate a K02 bed with the lowest water concentration which will provide suitable 02 evolution. This will prevent mushing of the canister bed, due to overloading with water, and will
10 also prevent over-production of 02. Bolles et ale (1964) reported that a microcontactor with potassium superoxide, and probably other superoxides, for air revitalization is feasible. A microcontactor is a device for producing finely divided superoxide particles by grinding a block of high density material, and then providing intimate contact between the finely divided superoxide particles and a dynamic air stream. No over-production of oxygen was encountered, as is commonly the case when using canisters of the superoxide. The microcontactor matched respiratory quotients (RQ) between 0.6 and 1.1 required by man. Steady state conditions with respect to 02 and CO 2 concentrations were obtained after only a few minutes of operation.
Recently, Li (1979) developed a life support system that used canisters of potassium superoxide and a canister of lithium hydroxide to release oxygen and at the same time to remove exhaled carbon dioxide in a manned submersible.
This all-chemical life support system will enable the submersible Deep Quest to remain under water six times as long as previously possible, for a total of 72 hours.
2. Alkali metal peroxides
Lithium peroxide, potassium peroxide and sodium peroxide are also of interest to remove CO2 in closed space.
Lithium peroxide (Li202) is of interest as an air vitali zation material because in the presence of moisture it can be used to react directly with carbon dioxide to yield oxygen and lithium carbonate (Markowitz, 1964; Miller, 1966).
11 Thus, it is possible to remove 0.96 Ib of C02 with each pound of lithium peroxide from a closed breathing
system and at the same time, to release 0.35 Ib of oxygen to
the system. The RQ for a system employing only lithium
peroxide would be 2.0. As a result, the use of this chemical would require an additional source of oxygen. The theoreti
cal capacity of lithium peroxide for C02 removal is about 4%
greater than the capacity of lithium hydroxide for CO 2. In the presence of water vapor, carbon dioxide
absorption and oxygen generation by lithium peroxide does occur, but oxygen generation lags far behind the amount
anticipated on the basis of equation (20). However, the
absorption of CO2 and the evolution of oxygen proceed by two different reactions. Lithium peroxide and water vapor
first react to yield the active C02 absorbents, LiOH,
LiOH • H20 and H202.
~ Li202 + 2 H20 2 LiOH + H202 (21) LiOH + H20 ~ LiOH· H20 (22)
C02 is absorbed via:
2 LiOH + CO 2 and
2 LiOH· H20 + CO 2 ~ Li2C0 3 + 2 H20 (24) Oxygen is evolved as a result of the decomposition of the H202:
~ H20 + 1/2 02 (25)
12 It has been shown that in order to achieve theoretical yields of oxygen, it will be necessary to develop a catalyst to insure the decomposition of all H202 formed in equation 21.
The chemistry of lithium peroxide and preliminary respiratory exchange of this compound has been reviewed
(Ducros, 1966). The state of art of lithium peroxide, as an air revitalization material, is not nearly as advanced as it is for superoxides.
III. Regenerative Type Of C02 Adsorbents
1. Molecular sieves
Molecular sieve adsorbents were introduced to industry in 1954. They have been applied by the process industries to the drying and purification of a large variety of gas and liquid streams. Molecular sieve adsorbents are synthetic crystalline zeolites (Sherman, 1977). The structure formula of the type A zeolite is:
Me x/n (A102) x (Si02) v : m H20 Me represents exchangeable cations of charge. The absorbents of present commercial interest are the sodium and calcium ion-exchanged forms of this structure.
The zeolite is precipitated in a hydrothermal process as a fine, white powder with an average particle size of about 2~. The powder is bonded with 20% clay and extruded as 1/16 and 1/8 in. diameter pellets or formed into roughly spherical beads from 4 to 12 inches in size.
13 The zeolites have a relatively high affinity for CO 2 but a still higher affinity for water. Thus, water will be picked
up in preference to carbon dioxide, displacing carbon dioxide
previously absorbed. The capacity of the zeolites for water
adsorption is shown in Figure 7. Silica gel is commonly used
to dry the air stream before it is introduced into the zeolite
beds. The carbon dioxide adsorbents can be regenerated by
application of a vacuum with the bed at ordinary temperature.
A vacuum less than 50 microns is required for desorption at
reasonable rates. Desorption of carbon dioxide and water could also be carried out with heat to 450°F and a direct hot air purge countercurrent to the adsorption flow.
Figure 8 shows the comparative carbon dioxide adsorption capacity of molecular sieves at a temperature of 77°F. Other
adsorbing materials such as silica gel and alumina also are
shown. The temperature at which the adsorption process takes place has an effect on the molecular sieve loading capacity.
Figure 9 shows the isothermal adsorption equilibrium capacity of molecular sieve, Type SA, at various temperatures. Under dynamic conditions, the usable capacity is substantially less than the values given by the equilibrium curves. Superficial velocity, carbon dioxide partial pressure, temperature, bed length, particle size, and process gas composition could influence the adsorption capacity of the molecular Sleves.
2. Metallic oxides
The separation process for recovering low concentrations of carbon dioxide from atmospheric gas streams by regenerative
14 metallic oxides has been studied (Rutz, 1963). Carbon dioxide absorption capacity of metallic oxides are presented in Table 1. Magnesium and calcium are among those metallic elements possessing the highest capacity for carbon dioxide.
They also have relatively low regeneration energy require ments. Only lithium and beryllium oxides have higher theoretical carbon dioxide capacities. The high toxicity of beryllium compounds excludes them from use in manned systems, while lithium oxide is extremely unstable, especially in the presence of water. Magnesium oxide and silver oxide are potentially valuable regenerative CO2 sorbents for enclosed space application because they are nontoxic solids that react with atmospheric CO2 at room temperature to form nontoxic solid carbonates. The temperature at which metallic carbon ates decompose is lower and the decomposition energy is also lower than for other reversible-carbonates systems.
A. Magnesium oxide
The presence of water with magnesium oxide is shown to be necessary for the efficient absorptions of carbon dioxide. The water takes part in the reaction as follows:
MgO + H20 - Mg (OH) 2 ..•.. (27) CO + H -H (28) 2 20 2C03
Mg (OH) 2 + H2C03 - MgC0 3 • 2H20 (29) Thus, the magnesium oxide is converted to the hydroxide, which readily combines with the carbonic acid and forms a hydrated carbonate. The affinity of carbon dioxide is due to magnesium oxide pellet density and porosity. The pellets
15 can be formed with magnesium carbonate mixed with approximately three times its weight of water and allowed to saturate for
1 hr before the paste is extruded into pellets. The pellets are then converted to magnesium oxide by heating them in a muffle furnace. Regeneration of this magnesium oxide is effected at 500°C for 1 hr. Vacuum and purging could further reduce the required regeneration temperature (Calombo, 1966).
B. Silver oxide
The use of silver oxide as a reversible CO 2 sorber was first studied in the 1940's as part of a program sponsored by the u.S. Navy. Results of this work indicated that Ag20 could be used as a reversible CO 2 sorber, but only if the Ag20 was present in a form that displayed a large amount of surface. It appeared that the reaction of CO 2 with
Ag20 was limited by the rate of diffusion of CO 2 gas through the Ag2C03 surface layer on the Ag20 particles. Chandler et al., (1962) have described work on using silver oxide as a C02 sorbent. Silver oxide was produced by the co-precipitation of silver carbonate and aluminum oxide and subsequent conversion of the carbonate to silver oxide by heating. This product removed 1% concentration of CO 2 effectively from air when used in a sorption column. Foster
(1965) developed a method to: (1) form silver foil into individual sorption modulars, (2) oxidize a surface layer of the foil to silver oxide, and (3) convert the silver oxide layer to an active C02 sorbent. It is concluded that the
16 si1ver oxide - si1ver system removed substnatially a 11 the CO 2 from air containing 1% CO 2 until at least 50% of silver oxide was converted to silver carbonate. Substantially all of the
CO 2 could be recovered by prolonged heating at 210oC.
3. Solid amines
Another type of chemical CO 2 adsorption is based on the use of ion exchange resins. The resins are polymeric materials that have chemical reactivity built in via the addition of activated functional groups. The amine resins are classified into strong (quarternary amine salts) or weak base (primary, secondary and tertiary amines). These com- pounds are basic and react with CO 2 to form carbonates. With the solid amines bed cooler than approximately 140°F, the adsorption process takes place according to the following reactions (Sutton et al., 1972).
Primary Amines - RNH 3 • OH (30)
Secondary Amines R2NH + H2O --- R2NH 2 • OH ...... (32) R2NH2 • OH + CO 2 ---- R2NHi HC03 ...... (33) Tertiary Amines
011- R3N + H2O -- R3NH' R3NHOH + CO 2 .....-.- R3NH' HC03 Tepper et al. (1969) conducted a screening to determine
CO 2 adsorption capacity from a flowing air stream for each class of amine base ion exchange resins. The results are
17 shown in Table 4. The strong base resins (IRA-400 and lRA-910)
have the greatest capacity for carbon dioxide. Dynamic CO 2 adsorption and desorption processes, as well as equilibrium
CO 2 bed loading conditions are sensitive to the amount of water present. IR-45 resin increases in bed water content up
to as high as 40% weight resulting in correspondingly
increased adsorption efficiencies. The physical and chemical
properties of IR-45 resin are shown in Table 5.
Regenerations of resins can be performed with heat/
vacuum at temperatures less than 100°C for weak base resins.
Strong base quaternary ammonium resins, which have a signifi
cantly greater C02 capacity during adsorption than the weak
base materials, could not be thermally/vacuum regenerated
even after prolonged heating.
4. Liquid amines
Liquid amines have been used in industrial applications
and on submarines. These systems usually employ monoethanol
amine (~mA) (Raven, 1962; Smith, 1960), or alkazid M (Goan,
1960; Allen, 1967; Gadomski, 1967), an alkaline metal salts of amino acid. The mechanism of C02 adsorption by organic acid salts seems to follow a course.
H20 + CO 2 ~ H2C03------(34) ~ H2CO + 2 RM 2 RH + M2C03---(35) M + H ~ 2MHC0 2C03 2C03 3------(36) R is the organic acid radical and M is the metal ion. Figures
10 and 11 show the carbon dioxide absorption and desorption
18 rates of both MEA and Alkazid M solutions. Comparison of both CO 2 absorbents are summarized as follows: PROPERTY RESULTS OF EXAMINATION
Ease of handling Alkazid M slightly superior to f.1EA
Vapor pressure Alkazid M greatly superior to (atmospheric contamination) MEA
Oxidation resistance MEA very slightly superior
Viscosity MEA considerably less viscous
CO 2 absorption capacity MEA has greater capacity
Stripping rate MEA slightly superior
Toxicity Alkazid M comparatively non toxic
Corrosion MEA less corrosive
5. Carbon dioxide reduction
A. Sabatier process
Carbon dioxide is mixed with hydrogen and passed into a reactor where it is reduced with a catalizat at 400°C to 700°F. Sabatier reactions procedure is:
The precious metal ruthenium, rhodium and iriclium are the most effective catalyzing materials for promoting the reduc- tion reaction, with 0.5 percent rutherium-on-alumina the best of these (Thompson, 1967). The Sabatier stoichoimetric
H2:C02 ratio of 4:1 for complete conversion of C02 to methane is not optimum in an actual unit. Optimum reaction rates have been experimentally determined to be a molar ratio of
4.35:1 (Yakut S. Bowker, 1969). These catalysts for the
19 reaction are easily poisoned by halogen or sulfur-containing compounds, so these must be excluded from the feed gas.
B. Bosh reaction
In the presence of an iron catalyst at temperatures of 1100° to l800°F, carbon dioxide reduction reaction is represented by the equation:
CO 2 + 2 H2 ~ C + 2H20 Unfortunately, side reaction from the Bosh reaction results in the production of carbon monoxide and methane as follows:
CO2 + C ~ 2CO
2H2 + C ~ CH 4 The resulting gases are usually recycled to achieve higher degree of conversion.
6. Miscellaneous
A. Membrane separation
Membrane separations are receiving active attention
(Ketteringham et al., 1970; Major 1966). The power and area requirements still appear excessive. But in combination with other methods, this single phase operation looks promising.
B. Freeze-out
The most attractive freeze-out system is one in which both the water entrained by the process gas and the C02 are removed by freezing, subsequently, the solids are sub limated to vacuum and the heat of sublimation recovered in a regenerative manner to cool the incoming gas. The disadvantages
20 of this method are: (1) the relatively high power requirement at sink temperature above 420°C; (2) the sensitivity of the system to heat and leak, (3) the low overall system reliability associated with the use of Freon refrigeration loops.
21 OYXGEN SOURCES
I. High Pressure Gas Storage
The basic role of gaseous storage systems appears to be that of supplemental storage for pressurization of the
chamber. The need for high delivery rate in pressurization
also favors gaseous storage.
High pressure gas storage without incurring excessive pressure shell weight minimizes the container volume. It can be shown that if the fluid stored acts like an ideal gas, the weight of container designed to hold a given charge is essen
tially independent of pressure while container volume is
inversely proportional to pressure. Very-high-pressure
storage appears to be the ideal goal. However, gas compres sibility factors begin to limit the weight efficiency of storage. At pressure above several thousand p.s.i., gases become less compressible. The decrease in compressibility is less serious for helium than for oxygen and nitrogen.
Gaseous oxygen may be stored at any pressure up to
5000 p.s.i. The container must comply with the specification for the pressure at which it will be used and be suitable for oxygen storage.
II. Cryogenic Storage
The cryogenic storage of liquid oxygen offers some advantage when space and weight are prime considerations in the submersible. These advantages are a higher fluid storage density at low to moderate pressure, reduced container weight
22 per unit of stored mass, provision of potential refrigeration of cooling sources as heat sinks for liquid oxygen.
The major defects are the sensitivity to unexpected heat leaks and the complexity of delivery at different pressures and temperatures. These defects require special attention to insulation needed, fluid expulsion, phase separation for venting and quantity measurement. Cost, devel opment time, servicing equipment and limited expulsion capability are other disadvantages.
III. Chemical Generation
Table 6 shows some of the pertinent physio-chemical properties of oxygen-producing chemicals for close environ ment use. Potassium superoxides, chlorates and hydrogen peroxide are compounds of primary interests.
1. Chlorate candle
The composition of chlorate candles are 92% sodium chlorate, 4% barium dioxide, and 4% steel fiber (or 4% of iron powder) (Smith, 1960). The ingredients are mixed with
3% by weight of water and pressed into shape after the water is removed by drying.
The principle utilized in chlorate candles is the decomposition of sodium chlorate at high temperatures (700
800°C) into sodium chloride and oxygen. The heat for this decomposition is supplied by the oxidation of iron wool mixed with the chlorate. Since there is a tendency toward libera tion of a small amount of chlorine, barium peroxide is added
23 to provide an alkaline medium for recurring it; barium peroxide also decomposes hypochlorite intermediate to yield oxygen. Iron wool in chlorate candle acts as a fuel and holds the burning front together. It also promotes flame penetra tion into the unburned layer underneath and initiates combustion.
A candle with an average volume of 11 in3, a diameter of 6 3/8 in, and height of 1.9 in produced approximately
25 ft3 02 STP and burned ten minutes. This oxygen contained
0.2 ppm C12, 5 ppm CO, and at least 6 grains of NaCl. The NaCl was collected on a filter. The maximum rate of gas evolution was 5 scfm for the entire candle and the overall average rate was 2.5 scfm.
Chlorate candles have long shelf life and are instantly available when needed. For maximum safety in the storage of candles, no organic material should be near by. Candles heated to the point of ignition are not dangerous by them selves, but organic matter (wood, cardboard, tubing, etc.) in contact with the candles can create a vigorous fire which might ignite other candles.
2. Alkali metal superoxides and peroxides
(See Carbon Dioxide Removal Section II, 1 & 2)
3. Decomposition of hydrogen peroxide
Hydrogen peroxide, H202, is available commercially at a concentration of 90% by weight percent or lower. Higher concentrations are desirable since they have a higher content
24 of oxygen and greater stability. To generate oxygen, hydrogen peroxide decomposes according to the equation (Rousseau, 1963):
~ H202 (liq.) H20 (liq.) + 1/202 + 1242 Btu/lb.. (30) Catalysts are required for smooth and rapid decomposition of hydrogen peroxide. Silver screen packs appear to be a good catalyst for the reaction. Disadvantages of hydrogen peroxide include its high toxicity. Concentrated peroxide blisters the skin on contact. Vapors and aerosals entrained with the oxygen are deleterious to the respiratory system.
4. Miscellaneous
A. Oxygen generation by electrolysis of water
The conventional electrolysis cell uses a potassium hydroxide (35% KOH in water) electrolyte which is contained in an asbestos matrix held between electrodes. 02 and H2 are generated at the anode and cathode respectively. The unit normally operates at about 150°F and consists of a condenser separator, current controllers, back pressure regulators and necessary plumbing, wiring and controlling devices.
Conventional processes for electrolysis of water require its initial condensation to the liquid state and separation from other gases. A possible method for achieving, in one step, the separation and electrolysis of water utilizes phosphoric anhydride (P205) as a desiccating matrix positioned between two electrodes. The air stream carrying the moisture passes over the anode side, and moisture is continuously adsorbed by the exposed phosphoric anhydride. Reaction
25 between the water and P20S can be expressed by:
P2 0S + H20 ~ 2HP03 and results in ionizable products. The matrix then becomes conductive, and application of a voltage between the electrodes causes migration of ions.
Electrolysis occurs at the anode and cathode converting water into oxygen and hydrogen and regenerating the phosphoric anhydride; the overall reaction being represented by:
2HP03 ~ H2 + 1/2 02 + P20S The high power requirements of the electrolysis cell is a disadvantage for submersible operations.
B. Photosynthetic gas exchange
Algae cells (green plants) use light energy to convert CO2 and water into oxygen and organic compounds required for the formation of new cell material (Rousseau,
1964; Miller, 1966). The system still requires a great deal of developmental work. The possibility of a compact, efficient algal system is still remote.
26 TRACE CONTM~INANTS REMOVAL
The source of contaminants in submersibles are man and his activities, materials and outgassing, equipment and processes, and finally, malfunctions and emergencies.
Principal sources of contaminants from man are expired air, urine, feces, flatus and perspiration (Roth, 1964;
Slonim, 1966). The primary components of expired air influencing toxicity are CO 2, H20, and carbon monoxide (Sjostrand, 1952; Tiunov, 1966; Wang, 1975, 1978). Compounds of relatively high vapor are outgassed from solid materials and from the hydrocarbon lubricants and operating fluids of machines. They originate from such sources as plastics, lUbricating compounds, insulations, paints, cements, and residual solvents from degreasing treatment. Aliphatic and olefinic hydrocarbons may originate as impurities in breath ing oxygen. The presence of these compounds in compressed gases stems from the cracking of hydrocarbon compressor oils.
By far the main contaminants of diving gases are those found in compressed air. These are primarily CO, CO 2, oxides of nitrogen, oil mist, and gaseous hydrocarbons. The impor tance of trace contaminants is not completely clear. The effect of CO is definite and clearly detrimental, as is CO 2 present in physiologically active amounts. Oil mist can cause lung damage. Freon-breakdown products, oxides of nitrogen, S02 and other irritant gases can cause serious responses if present in large enough amounts.
27 Various chemical substances for controlling these
contaminants have been considered. The substances suggested
fall into several classifications depending upon the nature of the contaminants and on the removal process.
I. Contaminant Sorption by Solids
1. Active charcoal
Activated carbons can be made from a variety of
carbonaceous raw materials. Generally the process consists
of selection of an organic raw material, dehydration and
carbonization, and finally, activation to produce a highly
porous structure. The ability of activated carbon to adsorb materials from gases stems from its highly porous structure.
Each particle consists of a vast network of interconnecting pores of a variety of sizes. The highly porous structures
result in a very large surface area, providing many sites
upon which adsorption of molecules can take place. Normally,
adsorption on activated carbon is the result of physical attraction of molecules to the carbon surface by van der Waals forces. Molecules with higher molecular weights normally experience greater forces of attraction than materials of lower molecular weights. Hence, activated carbons, aside from the effects of molecular screening due to the sizes of the pores, have a preference for higher molecular weight substances.
Generally, activated carbon has a high adsorption capacity for many of the anticipated trace contaminants. It adsorbs effectively the vapors of most materials that are
28 liquid at temperatures of OaF or above. This includes many hydrocarbons, alcohols, ketones, aldehydes, mercaptans, organic acids, and halogenated materials. Most of the odor causing and toxic materials expected to accumulate in a submersible atmosphere are effectively adsorbed by charcoal.
Activated charcoal, however, is a poor adsorbent of methane, hydrogen, carbon monoxide, carbon dioxide, ammonia, sulfur dioxide and nitrogen dioxide. Water is adsorbed but not stained. Activated charcoal is sometimes impregnated with phosphoric acid or loaded 60% with NiC12 for the removal of ammonia and carbon dioxide, respectively. Essentially com plete removal of 802 can also be achieved using activated charcoal loaded to 10% with LiOR (Moore, 1974).
2. Molecular sieve
The molecular sieve materials are alkali metal aluminosilicates, quite similar to many natural clays. The sieves are characterized by a large number of uniform pores, and the pore size is controlled by varying the natural of the metal ion included in the crystal lattice (see Carbon
Dioxide Removal IV).
Although synthetic zeolites are available which will adsorb hydrocarbons and other contaminants from the air, they have less capacity for these materials than activated carbon.
They do not adsorb hydrogen, carbon monoxide, or methane effectively at room temperature. Furthermore, since a polar substance such as water displaces less polar materials from
29 molecular sieves, the processed alr flowing through the molecular sieve must be thoroughly dry. Molecular sieves do not appear to perform as satisfactorily as activated carbon for trace contaminant control systems.
3. Silica gel, activated alumina
Other solid adsorbents for gases, such as silica gel and activated alumina, are much less effective than activated charcoal for the adsorption of the trace contaminants expected to occur in the submersible.
II. Chemical Conversion of Contaminants
1. Hopaclite
Hopaclite is basically a mixture of manganese oxide
(78.3%), cupricoxide (13.1%), and traces of other elements
(Christian, 1963). It has been used for many years at room temperature or at lower temperature for catalyzing the oxida tion of carbon monoxide. At 250°F, the Hopaclite is relatively immune to the effects of water vapor, whereas at lower temperatures it must be well protected from water vapor.
CO removal efficiency was found to be reduced to about 60% when the atmosphere was saturated with water vapor at room temperature (Thomas, 1969).
Hopaclite is also a good catalyst for conversion of hydrocarbons and other organic materials at low concentration in air to carbon dioxide (Thomas, 1970; Johnson, 1967). At temperatures of 300 to 400°C, both aromatic compounds and aliphatic hydrocarbons except methanol gave about 90% or
30 better conversion to CO 2. At 200°C, combustion of some hydrocarbons was incomplete. The combustion of methane over the catalyst was very inefficient. It produced only S% of the theoretical CO 2 at 3S0°C and about 30% at 400°C. Hopaclite also catalytically converts certain organic com pounds to other toxic species. Ammonia and organic nitrogen compounds yield nitrous oxide on oxidation and Freon-II decomposes by Hopaclite to other acid vapors.
2. Purafil-permanganated (3%) alumina
Purafil is a combination of a high surface area substrate (activated alumina) and a broad spectrum oxidant
(potassium permanganate). The substance is non-toxic and non flammable and can be used to control gaseous contaminants such as CO, ammonia, oxides of nitrogen, mercaptans and amines, ethylene and other sulfur compounds.
At room temperature, Purafil has no effect on CO, either retentive or catalytic (Moore, 1974). However, CO is
aC quantitatively converted to CO 2 at lSO°C. NH 3 at l80 was found 100% converted to another gas, presumably N2. Although Purafil does appear to be highly promising, some caution should be exercised; it may behave very similarly to Hopaclite.
It catalytically converts certain organic compounds to highly toxic species.
31 GAS PROPERTIES
I. Carbon Dioxide (C0 2) Carbon dioxide is a nonflammable, colorless and odorless
gas which is about 1.5 times as heavy as air. It has a mildly
acid taste due to the formation of carbonic acid in the mouth.
One volume of CO 2 will dissolve in approximately one volume of water at atmospheric pressure and 15°C. The gas is readily
liquified by compression because its critical temperature is
relatively high (3l.10C). When the liquid is allowed to
evaporate, it freezes to a snow-like solid at -56.2°C. The
solid vaporizes without melting (sublimes) because its vapor
pressure is one atmosphere at -78.5°C. If a carbon dioxide
concentration builds up in a breathing medium, harmful effects result (carbon dioxide excess).
Carbon dioxide is not a chemically active compound as
such and high temperatures are generally required to promote
its reactions. Carbon dioxide is stable under normal con ditions but at temperatures above 1700°C, it dissociates into oxygen and carbon monoxide to an extent (15.8% at 2227°C)
according to the following formula: 2 CO 2 + 2 CO + 02. For thermodynamic and detailed physical data, see Matheson
Gas Data Book (1971) and u.S. Navy Diving Gas Manual (1971).
II. Nitrogen (N2) Under ordinary conditions nitrogen is a colorless,
odorless and tasteless gas. Its density under standard
conditions is 1.2506 g/£ which is slightly less dense than
32 air. Under standard conditions, 100 m~ of water dissolves
2.4 m~ of nitrogen. Nitrogen comprises approximately 79% by volume of the air. It will not burn and will not support combustion.
Nitrogen is extremely inert except when heated to very high temperatures where it reacts with certain active metals to form nitrites. At high temperatures nitrogen combines with hydrogen forming ammonia (NH 3) and with oxygen forming nitric oxide (NO). Under high pressure nitrogen has an intoxicating effect (nitrogen narcosis). Thermodynamic and compressibility data be seen in Matheson Gas Data Book (1971) and U.S. Navy Diving Gas Manual (1971).
III. Oxygen (02 ) Oxygen exists in a free state in the atmosphere, of which it forms approximately 21% by volume. It is colorless, tasteless and odorless. Matter cannot burn unless oxygen is present, but oxygen itself is not flammable. Oxygen alone is capable of supporting life and in some instances is used instead of air as a breathing medium. When breathed too long under increased pressure, oxygen has a harmful effect on the body, known as oxygen poisoning. Oxygen is an active element that combines either directly or indirectly with nearly all the other elements, the only exception being the inert gases.
For thermodynamic data and its reactivity, see Matheson Gas
Data Book and U.S. Navy Diving Gas Manual.
33 IV. Air
Air is a mixture. Nitrogen, oxygen, and the rare gases are present in the atmosphere in almost constant pro portion. The percentage composition of dry air does not vary much with location on the earth's surface or with altitude.
However, the density of air, as reflective in pressure,
varies greatly with altitude. The average pressure at sea
level and 45° latitude os 760 mmHg. The reactivity of alr is
due to its content of oxygen. Thermodynamic data is also
shown in Matheson Gas Data Book (1971) and u.s. Navy Diving
Gas Manual (1971).
V. Helium (He)
Helium is colorless, odorless, tasteless and inert.
It is exceptionally light, nontoxic, nonexplosive and is
only slightly soluble in water (0.80 parts in 100 parts).
It will not react with other elements or compounds under ordinary conditions. Helium when mixed with the proper proportion of oxygen, forms an artificial atmosphere which
is less dense and less narcotic under pressure than air.
Helium conducts heat much more rapidly than air. Its
thermodynamic properties and compressibility data are listed
in Matheson Gas Data Book and U.S. Navy Gas Manual.
VI. Carbon Monoxide (CO)
Carbon monoxide is another important harmful gas. It
is colorless, odorless, tasteless and highly poisonous.
Carbon monoxide is produced by incomplete combustion of
34 carbon-bearing materials. Carbon monoxide is found in dangerous concentration in engine exhausts and in closed compartments where paint has been deteriorating. If it contaminates an air supply, serious consequences can result (carbon monoxide poisoning).
35 Table 1. CO 2 Capacity of Alkali and Alkaline Earth Metal Oxides and Hydroxides
Formula CO 2 Capacity Compound Weight (lbs CO 2/lb Compound) Li20 29.9 1. 473 Na20 62.0 0.710 K20 94.2 0.467 MgO 40.3 1. 091 CaO 56.1 0.785 SrO 103.4 0.425 BaO 153.4 0.287 Ag20 231. 6 0.189
LiOH 23.9 0.919 NaOH 40.0 0.550 KOH 56.1 0.392 Mg(OH)2 58.3 0.754 Ca(OH)2 74.1 0.594 Ba(OH)2 171. 4 0.257 Soda lime 46.8 0.488 Baralyme 87.5 0.503
36 Table 2. CO 2 Capacity of Peroxides and Superoxides
Formula CO 2 Capacity Available Oxygen Compound Weight (lbs C02/1b Compound) (in Wt %)
0.959 34.9 Li202 45.9 Na202 78.0 0.564 20.5
K202 110.2 0.399 14.5 Na02 55.0 0.400 43.6 K02 71.1 0.309 33.8
37 Table 3. Properties of C02 Absorbents
Chemical Purity Causticity as m~ of O.lON Type of Water of Dehydrated Porosity HC~ to Titrate Dust From Absorbent Content(%) Absorbent (%) ( %) 10 g of Material
Anhydrous a Lithium Hydroxide 0.00 101 14.8 0.65
Partially Hydrated Lithium Hydroxideb 5.45 96.9 8.76 26.8
Sodasorb Hpc 14.8 107e 7.69 4.85
Sodasorbc 15.5 104e 2.86 6.45 co C") Baralymed 17.5 98.8f 3.85 3.20 a - a product of Foote Mineral Corporation b - a product of Lithium Corporation of America c - a product to W.R. Grace & Company d - a product of National Cylinder Gas e - basis on 84% Ca(OH)2 + 1.65% NaOH + 2.89 KOH ( 8 ) f - basis on 79.5% Ca(OH)2 + 3.7% Ba(OH)2 + 0.42% NaOH + 5.19%KOH (8) Table 4. CO Absorption Capacity of Ion Exchange Resins 2 (after Tepper et al., 1969)
2 Dynamic C02 Capacity H Life l 2O (mg CO2/g Sample Description (%) (min) sorbent)
IR-45 Weak Base, 2° Amine 56 11 3 XE-236 Weak Base, 2° Amine 50 67 18 XE-233 Weak Base, 2° Amine 40 112 20 IRA-68 Weak Base, 3° Amine 60 1 0.3 IRA-93 Weak Base, 3° Amine 56 26 7 0"1 Epon M 812/DET Epoxide/2° Amine 48 188 59 AA/TEP Acrylic Acid/2° Amine 65 116 47 IRA-400 Strong Base, quaternary 47 260 80 IRA-910 Strong Base, quaternary 65 346 109
lTime for effluent to reach 0.2% CO 2.
2 C0 2 removed until 0.2% CO2 is obtained in effluent.
Note: sample bed depths were 50-150 rom, according to particle size and density. Table 5. Physical and Chemical Properties of IR-45 Resin
(after Martin et al., 1970)
Physical Characteristics
Physical form Uniform, beadlike particles
Density 39-43 Ib/ft3
Moisture content 37-45%
Screen grading 20-50 mesh
Effective size 0.35-0.50rnrn
Uniformity coefficient 1. 6 max
Voids 35-45%
Chemical Characteristics 3 Exchange capacity 43 Kg/ft max as CaCo3 (27 Kg dynamic capacity)
pH Range 0-7
Chemical stability Excellent; completely insoluble
and inert in strong acids (except
nitric) ,conc'd alkalies,
aliphatic and aromatic hydrocarbons,
alcohols, ethers and all other
cornmon solvents; prolonged exposure
to strong oxidizing agents should
be avoided
Stability at elevated
temperatures Outstanding; exchange capacity
unchanged after prolonged
exposure to boiling water
40 Table 6. Comparison of Oxygen-producing Chemicals
[after Coe et ale (1953) and Petrocelli (1965)]
K0 Li LiN0 2 Na02 202 Na03 3 LiCl04 NaCI03 H202
Available 02 (theoreti-
cal), weight percent 33.8 43.6 34.8 56.3 23.2 60.1 45.1 47.1
Purity - 0.90 (a) - 1.00 1.00 - 0.90
Available 02' Ib/lb 0.32 0.392 0.375 0.56 0.232 0.601 0.40 0.423 ~ Density, Ib/in. 3 0.0237 - 0.0774 - 0.0861 0.0878 0.0815 0.0502
Heat of reaction,
b c d Btu/lb 4 1 5 6 3 5 e-36 3 +1515 -488 -596 +422 +1106
H balance, Ib/lb -0.0207 -0.0246 -0.136 0 0 0 +0.577 20 - H balance, Ib/lb 02 -0.0862 -0.0862 -0.225 0 20 - 0 0 +1.34 0.31 0.40 CO2 balance Ib/lb 0.96 0.31 - - - -
a 10 percent Li204. b + indicates exothermic reaction; - indicates endothermic reaction.
c 2 K0 + 1.23 CO + 0.23 H 0.77 K + 0.46 KHC0 + 1.5 0:;J' 2 2 20 2C03 3 d 2 Na0 + 1.23 CO + 0.23 H 0.77 Na2C03 + 0.46 NaHC0 + 1.5 " 2 2 20 3 °2 e Li 202 0.7 tJ1 <, tJ1 0.6
~ e..... ~ +J ~ .,-i 0.5 .,-, ~ o I:.T ro V ~ P., ~ ro j ..,. o 0.4 i'n.. ~ ~ 0 " .,-i +J P., 0.3 H 0 lJl [;J - 10° e .Q ,:x; 0 - looe N 0.2 8. - 0 27° e o 0.1
o o 10 20 J O 11.0 so GO 70 8 0 90 100 Relative Humidity ( %) Figure 1. CO~ absor?tion capacity of ~nhydrous Lithium "- Hydroxide at various temperatures and humidities
(after Wang, 1979)
42 0.8
t» <, b' 0.7 ~ ~ r-, ~ I- ~. ~ ...... +J 0.6 I'...... :. 'M ~ U ""'"'" eu ~ - 0.. 'l A eu -- re u 0.5 s:: 0 ·M +J ~ P.. 0.4 H El- 10° C 0 , Ul 0- lS· C ~ 6,- 27°'C 0.3 N u0 0.2
0.1 o 10 20 30 40 50 60 70 80 90 100
Relative Humidity (%)
Figure 2. CO absorption capacity of partly hydrated 2 lithium hydroxide at various temperatures and humidities (after Wang, 1979)
43 -A ~~ 0.25 , »> -- ~ ~ ~ A 0.20 r /' -: 1:? - 95° F /}; - 80° F 0.15 P o - 65° F If o - 50° F
u> ~ .r.'L 0.10 Ii j V .L r> N o o ~ 0.05 'tt """ ""
0.00 o 10 20 30 40 50 60 70 80 90 100 Relative Humidity (%)
Figure 3. CO 2 absorption capacity (gig) ot So~aso~b «t various temperatures and humidities (after
Wang, 1980)
" 44 0.30 ~ 01 -' <, 1-1" 1\ ~ tJ' -- ~ :>t 0.25 ~ +J r ~ .r-! U Itl V 0.. Itl u 0.20 ~V ~ ~c> v ~ 0 .r-! 95° F +J *- 0.. 8 - 80° F I-l / -: 0 0.15 I o - 65° F 1Il ~ / V o - 50° F N 0 ~ o . . 0.10 ~ .rx:.... - ~ . .,..
0.00 o 10 20 30 40 50 60 70 80 90 100
Relative Humidity (%)
(g/~of Figure 4. CO 2 absorption capacity Sodasorb HP at various temperatures and humidity (after Wang, 1980)
45 I 1 ill - 95° F 0.25 t::. - 80° F . 0 - 65· F V e - 50· F tJ'l <, 0.20 tJ'l I >t +J 0..-i o 0.15 ~ /' ro / A 0-ro --- ~ A/ o / --- -It.{" V c 0 0..-i 0.10 ~ / +J ~~ 0- A H V 0 (/) V ~ ~ J:\.- ~ 0.05 f':\. ...n I", N 0 -- U -
0.00
o 10 20 30 40 50 60 70 80 90 100
Relative Humidity (%)
Figure 5. CO2 absorption capacity (gig) of Bara1yme at various temperatures and humidities (after Wang, 1980)
46 1\ • I I I ~t--.....l .. I I I -".';2 11. - --..;".'2 0.55 J - b..- ~ - V ~ ,/ ~ 0.50 .0/'./ M:1. - - AQhydrous LiOH 0.45 I *0 - Partially 'hydr ated LiOH 8- - Sodasorb HP b'1 <, e - Sodasorb b'1 0.40 1/ II 0 - Baralyme ~ +J ·M j U n:l 0.35 o, I on:l 0.30 s:: .»: 0 ·M V +J o, ~ ) ~ ~ 0.25 I 0 til tf ~ V 0.20 / I;) N I I ..... o0 V 0.15 I /; I ~~ 0.10 ~V ... -? .,I2r
0.05 I I I I I I I
0.2 0.4 0.6 0.8 1.0 1.2 1.4 1.6
Residence Time (seconds)
Figure 6. CO 2 absorption capaci.t.y vs. residence time at 80°F and
90% relative humidity (after Wang, 1981)
47 45
40 +Js:: Q) TEMPERATURE 77°F ..0 H 0 en 35 ~ >. H 30 Cl 4-l 0 .0 25 r-f
0 0 r-f <, 20 '0 Q) ..0 H 0 en 15 '0 ..:x: H Q) +J eu 10 3= 4-l 0 5 .0 r-f
0 0 20 40 60 80 100
Relative Humidity %
Figure 7. Equilibrium water-
adsorption capacities
(after Griesmer, et al., 1966)
48 15 14 +'s:: Q) 13 / .Q H 0 12 / ) Ul '0 V ..-x: 11 / j- ~ H Cl 10 / /'1 ~ V / 0 9 TYPE 5A/ / / .Q r-l £/4A / 0 8 0 r-l V/ ~X <, 7 .J '0 Q) .Q H 6 / / 0 V Ul '0 .;.: V ..-x: 5 /
N 0 4 / o / ~ 3 /7 / ~ 0 ",P l7 V .Q V/ SILICA~ ,-I 2 I-~ I,.--- I V ALU~ 1 ./ .--l.-----' I _ ~ ~ 0 V ACTIVATED C~RBON
0.2 0.5 1 2 5 10 20 50
Carbon Dioxide Pressure, rom Hg
Figure 8. Comparison of adsorption capacity of
three types of mo1ecu1ea~ sieves at
77°F (after Rousseau, 1963)
49 28
26 .j..J ~ Q) 24 .0 H 0 til 22 '0 .cx: ~ 20 H t:l 18 4-1 0 .0 16 r-l ...... ~ - 0 14 »> ~ 0 r-l <, ,~ V -o 12 /'" Q) TEMP: 32° F ~ ~- .0 ~ H ./-: 0 10 til '0 V .cx: 8 .,.. ,V ./ ". N V 0 77 -: 6 -"V /'" o /" 4-1 ..------0 4 ~ ...... '.0 ~ V ~ V r-l ------~ 2 122 - - - 212 -- 0 1 2 3 5 10---20 40 60
Carbon Dioxide Pressure, rom Hg
Figure 9. Adsorption isotherms for Type 5A
mo1ecu1ear sieve (after Rousseau,
1963)
50 45 I I 11 I I I I II I I I 10- - I- 1:110-- 40 ...,- ~~ ,I:: 10- ,/ "" :/1:) 10- .~ - 35 MEA"'} ~( ~ / L INSOLUBLE PPT 0 / APPEARS - .,-l - .jJ - - ::l ..-l 30 V/ 0 1:)/ (f) f-- K.ALKAZID M - (j) /1 e f-- - ::l 25 1-:- i J>e ..-l 0 :> - 11.~ - <, / '0 - - (j) , P on 20 r H 0 Ul - - ~ f-- - 15 N /I 0 u f-- l~ - (j) BOTH SOLUTIONS 3.61 N V 0 - e 10 - CONTACT TEMP. 88 F ::l A ..-l BE> 0 :> - - 5 =1is - - o ~I II II II II I II o 30 60 90 120 150 090 420
Contact Time (Min)
Figure 10. Carbon dioxide absorption rates and capacity (after Goan, 1960)
51 4B I I I I I I I I I c 0 .... - ..-1 ~ 40 .\ ::l 3.61 N. SOLUTION ..--l 0 0 DISAPPEARS TEMP. 100 0 U) ~PT ATMOSPHERIC PRESSURE - OJ 8 32 r-l 0 - :> ~\~ <, '0 24 OJ .0 ~ ~ALKAZID M H - 0 ~ - lJ) ~ 16 <, ..,. ~ 0~ '" N 0 - " MEA~ - U 0 8 0 OJ - 8 ---- 0 r--- ..--l ~ - 0 :> II II I I I I I o 8 16 24 32 40 48 56 64 72
Reboiling Time (Min.)
Figure 11. Carbon dioxide de~orption r a t e s (after
Goan, 1960)
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