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Chern. Eng. Cornrnun.Vol. 33 pp. 93-1050098-6445/85/3301-0093$20.00/0 -" 1985Gordon and Breach, Science Publishers, Inc.and OPA Ltd. Printed in the U.S.A.

THERMAL DECOMPOSITION OF SODIUMBICARBONATETIMOTHY C. KEENER!

Department of Civil and Environmental Engineering, 71University of Cincinnati, Cincinnati, OH 45221

GEORGE C. fRAZIERChemical, M etullurqical, and Polymer Engineering Department

The University of Tennessee, Knoxville, TN 37996-2200WAYNET. DAVIS

Civil Engineering DepartmentThe University of Tennessee, Knoxville, TN 37996-2200

(Received June 13. 1983; infinal form JUlie 4. 1984)The thermal decomposition of sodium bicarbonate, a candidate material for Iluegas desulfunzation. hasbeen investigated over the temperature range of 225-350F (.380-450K) and over the particle size range of51-140 ,urn. The shrinkingcore model, with chemical reaction as the ratecontrolling step, providesa good fitto the data in the temperature range investigated. However, caution should be exercised in extrapolatingthese results into the range of about 600"F (about 590K) where sintering of this material is reported to occur.The activation energy of the decomposition reaction is 20.5 kcal/mol (85.7 k.l/rnol).KEYWORDS Thermal Decomposition Sodium bicarbonate

INTRODUCTIONThe focus on sodium bicarbonate (NaHC0 3) as a dry additive injection material forflue gas desulfurization (fGD) systems has come about mainly from the abundanceand proximity of nahcolite to oil shale reserves, and the desire to identify an all dryfGD system. Nahcolite is a naturally occurring; form of sodium bicarbonate (70%NaHC0 3 , remainder inerts), found in large quantities in four major geological basinsin the Green River formation. This formation is located in northwestern Colorado,northeast Utah and southeast Wyoming (Knight, 1977) which places it within areasonable transportation range for coal-fired steam plants located in the western partof the United States.

Data from pilot plant experiences (Carson, 1980) on dry injection of nahcolite andpure sodium bicarbonate have shown that continuous injection of NaHCO, issuperior to batch inject ion for S02 removal. Also, it has been shown that sodiumcarbonate (the product of decomposi tion) does not react as readily with S02 as theparent NaHC0 3 .

t To whom correspondence should be addressed.


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94 T.C. KEEN ER, G.C. FRAZIER A N D W.T. DAVISA possible explanation for this behavior may stem from the effect of the heatedf lue gas on thermally decompos ing the NaH CO , in to the product N a 2 C 0 3 while

undergoing simultaneous reaction with SO,.Th e decomposition reaction for sodium bica rbonate is given by:

Th e temperature at which dec omposition o ccurs will depend, of course, on the extentof the carbon dioxide and water vapor back pressures. Perry er al . (1973) reportdecomposition occurring at 518F (270C) while the C hurch a nd Dwight C o. l i teratureindicate partial de compo sition a t temp eratures as low as 100F (38C). Th e differencein these values is presumably due to differences in the back pressures of the gaseousspecies which is not always reported in the literature. Information concerning thedecomposition rate coefficient has not been found in the literature.The thermal decomposition temperature has been reported (H owatson et a / . ,1977)to be an important parameter in pore development in the product sodium carbo nate.The optimum decomposition temperature range has been found to be from 250 to300F (121-149C) for minus 200 mesh (77 pm m aximum size) particles (Knight, 1977).However, the analytical techniqu e providing the dat a on which this conclusion is basedwas not given. Anothe r study (Stern, 1978)showed tha t th egre atest specific surface areawas generat ed between 40 0 an d 600F (204-316C) in a forced draft oven. Loss ofspecific surface area was observed at temperature s abo ve 600F (316C). This loss wasattributed to sintering, which was confirmed through scanning electron microscopic(SE M ) analysis. Th e effect of sintering could therefore negate possible adv anta ges ofsurface area generation by thermal co mm inuti on by injecting this material directly intothe flame zone of a furnace.

E X P E R I M E N T A L A P P R O A C HIt is with these results in mind that th e rate of the thermal decomp osition of N aH C O ,by a heated gas steam w as studied as a function of tem perature an d particle size. Th emethod ology consisted of passing a preheated gas stream over a fixed bed of discrete,sized particles and by m onito ring the ratk of c arb on dioxide evolution with respect t otime, under isothermal conditions. The data were then interpreted in terms of ashrinking core model.A schematic diagram of theexperimental appa ratus is shown in Figure 1. Instrumentgrad e compressed air (relative humidity of 6-10'x at amb ient temp erature) wasmetered via calibrated rotameter C to either one of two dimerential type fixed bedreactors, E. Th e reac tors were 8 in (20.3 cm) in le ngth by 1.75 in (4.45 cm) inside d ia-meter. Tempe rature was monitored 1.5in (3.8cm) upstream of the bed by a type Kthermocouple at point 7:The temperature inside the reactors could be maintained atf F (+ 23C). T he fiber-glass filters used for the solids bed su ppo rt were placed on astainless steel wire mesh implant, F, in the lower portion of the reactor, which wasdisassembled and renewed after each test. Th e carrier gas from the base of the re actorswas then directed through orifice meter H a n d exhausted via vacuum pum p I. The flowrate for all tests was 0.89ft3/min (4.2 x 10-4m3/s ) (STP) and was m oni tored by


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T H E R M A L D E C O M P O S I T IO Nhbt

2entLegend :

A Charcoal bed G C02 analyzerB Flow contro l valve H O r ~ f ~ c ee t e rC Flow mete r I PunlpD Hea ter J Solids injectio n tub eE Reactor P Pressure tapF Fixed bed supp ort T Thermocouple

FIGURE I Schematic diagram of the experimental apparatus.

maintaining constant rotameter readings at C and a constant reactor pressure at P.Orifice meter H was monitored to assure leak-proof operatio n after a test hadcom men ced. All materials of constr uctio n for the system consisted of either 304 or 316stainless steel. Additional details regarding the experimental app ara tus m ay be foundelsewhere (Keener, 1982).The test material used in the study was commercial grade sodium bicarbonatesupplied by the Church and Dwight Corporation. The sodium bicarbonate wasclassified into three discrete particle size ranges by means of double sieving throughsta nda rd Tyler sieves. Th e smallest particle size range studied passed 270 mesh and wasretain ed by 325 mesh. Th e medium size range passed 170 mesh bu t was retained by 200mesh, and the largest particles passed 80 but were retained by 100 mesh. A sample ofthe particles was examined by microscopy. The particles were found to be approx-imately spherical in shap e with a n elongation facl.or (length to breadth measurem entratio) not exceeding 1.25 for each of the th ree sizes. Th e equival ent diam eters (where theequivalent diame ter is defined as the length of the side of a cu be of equal volume a ndthe thickness is assumed equal to the length of on e side) were found to be 51 ,87 an d140 pm, respectively.Th e test proced ure consisted of heating the compressed air at D, Figure 1, until asteady state test condition prevailed, as determined by monitoring thermocouple 7:


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96 T.C. KEE NER , G.C. FRAZIER AN D W.T. DAVISStatic pressure inside the reactors, read a t point P, was maintained a t one atmosphe reby flow con trol valve B. When test co nditions were obtaine d, the reactors were placedunder a slight, momentary vacuum in order to inject the particles through line J.A three gram sample, which was found to give a uniform bed d ept h of approxim ately1/16 in (0.159cm), was used for all tests. A bed of particles could be inserted into thereactor by this method in about two to four seconds. Evolved carbon dioxide levelswere then monitored as a function of t ime by a H oribainfrared analyzer (Model P IR-2000) located at point G. Th e reproducibility of this instrument was better tha n f .5%and the base line drift was less than 1% of full scale in 24 hours. The zero point wasestablished against essentially C0,-free nitrogen and calibration points were deter-mined using standard samples of 1% and 2% CO, in nitrogen. The instrument waszeroed prior to each run.The time required to heat the bed of injected particles to test conditions wasestimated in orde r to insure that the non-isotherm al period (particle injection time plusparticle heat up time) was small compared to the isothermal reaction period. Theme thod employed for this calculation was that of Fu rnas(l 930) , wh o estimated the bedheating rate for calcium carbonate particles undergoing thermal decomposition. Thismethod employs a transient bed heat-up model and requires a knowledge of the beddepth, initial tempera ture of the particles, an d the therm al properties of b oth the fluidand solids. The results from the use of this model on the experimental conditionspresented here suggests that the bed of sodium bicarbonate particles is essentiallyisothermal (9% of the free stream , steady state fluid temperature) in less than abo uttw o seconds, for beds of particles of all three sizes used in this work. A dditionally, th etime required to heat a single particle, modelled as a sphere, was estimated by themethod outlined by Perry (1973).Th e results of this calculation in dicate that the largestparticles of interest here a re essentially isothermal after approxim ately one second, aresult consistent with the previous calculation. Therefore, the total time, injection plusheating, for achieving essentially isothermal cond itions in the particle bed is estimatedto be from 4-6 seconds which, as will be show n later, is negligible com pare d to the timerequired for essentially complete decomposition of the sodium bicarbonate in thetemperature range s tudied. ,R E S U LT S A N D I N T E R P R E T A T I O NThe sodium bicarbonate thermal deconlposition data, fractional conversion, X, a s afunction of time, t , are provided in Figures 3 and 4 for three particle size ranges, and sixtempera tures in the range 225-350F (107-177C). T his temperature range isconsidered particularly relevant for the use of sodium bicarbonate as a dry sorben tmaterial in convention al fabric filters. Th e times required for -95-98% conversionrange from abo ut 5.1 x 10' to 6.2 x l o4 S ,depending on temperature a nd particle s ize.Exploratory tests conducted at temperatures abov e 350F (177C) produced initialCO , compo sitions in the effluent gas which exceeded the range (0-2.5%) of the CO ,analyzer and thus were not used in the determination of the decomposition kineticparameters .Th e fractional conversions, X, were computed from the C O, evolved to t ime t byintegrating under the effluent gas CO , concentration curves, a typical one of which is


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T H E R M A L D E C O M P O S I T IO N

J0 200 400 600 800Decomposition Time, Seconds

F I G U R E 2 Typical carbon dioxide production profile for the decomposition of sodium b i carbona~e .

provided a s Figure 2, and by relating this result to sodium bicarbon ate throug h thestoichiom etry of Eq. (I) and the definition of the fractional con version :

The subscripts o a n d i refer to the initial moles and that at time t , respectively. Theamount reacted at t ime t , the nu mer ator of Eq. (2), is given in terms of the CO ,produced according to:

CAI" - CA l i = flCC1, (3 )The total CO , produced a t complete decomposit ion, [CII, is [A],/a, so the fractionalconversion may also be writ ten as:

x = CCl,ICCl, (4 )The tailing off of the CO, evolution curve, Figure 2, prevented precise determinationof the time, T, for 100% decomposit ion, due to experimental and measurementinaccuracies. These times were therefore estimated by a regression technique describedbelow. The total mass of CO, available from a three gram sample of NaHCO, is0.786 gram, and those tests which did not integrate to within If:5"/, of this value wererejected.

M O D E L D E V E L O P M E N TThe da ta , F igures 3 and 4, were fi t by a model of the shrinking core type (Kunii andLevenspiel, 1969), following the work of Lang muir (l9 16) , wh o showed that thermal


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Decomposition Time , SecondsF I G U R E 3 Test of the reaction controlled, shrinkingcore model for decomposition of sodium bicarbonatein the temperature range of 225-275F.

-"0 200 400 60 0 800 1000 1200 1400 1600 1800

Decomposition Time. SecondsF I G U R E 4 Test of the reaction controlled,shrinkingcoremodel for decomposition of sodium bicarbonatein the temperature range of 300-350F.


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T H E R M A L D E C O M P O S I T I O N 9 9decomposition reactions occur at definite interfaces within the solid. The rate ofmovement of this interface from the outer edge of the decomposing solid toward itscenter is governed by heat transfer from the gas stream to the solid, the thermalprope rties of the solid, the decomp osition reaction rate, or the rate of m igration of t hedecomposit ion gaseous products a way from the interface (Ingraham e t a / . , 1963; Keha tet al . , 1963; Roberts et al., 1978). In simple terms, the decomposition rate may becontrolled by either heat transfer, chemical reaction, or diffusional processes. Forexample, the rate of decomposition of calcium carbonate has been shown to becontrolled by a stage of the process occurring in the interface at calcinationtemp eratures below 850C and to be heat transfer limited at tem peratures ab ove 900C(Ingraham an d M arier , 1963; Kehat and Mark in, 1963; Narsimhan, 1961). Ewing et a / .(1979) examined the n ature of C a O produced by calci te powder decomposit ion, asaffected by back pressures of C O ,. Da rro ud l and Searcy (1981) investigated the effectof C O, pressure on th e rate of deco mpo sition of calcite. Hill and Win ter (1956) reporttherm al dissociation pressures as a function of temp erature for cal ciu n~ arbonate.Ma cDo nald (1951) showed that the decom posit ion of sodium bicarbonate m ay becompared to that of calcium carbonate. However , as this carbonate decomposes atsubstantially lower temperatures than CaCO,, the problem becomes that of establish-ing the tem perature range where, for example, chemical reaction is rate controlling.Th e experimentally determined C O, evolution curves (Figure 2) indicate that thedecom position times for sodium bicarb onate ar e in the range of ten to twelve minu tesor greater in the temperature range of 350 F o r less. Co mp ariso n of this time with thetherma l response time of ab ou t one second o r less for a particle in the experimental sizerange sugg ests that heat transfer is not th e controlling process un der the experimentalcond itions. Diffusion of carb on dioxide o r water vap or (see Eq. (I) ) away from thereaction interface throu gh the pro du ct solid could be a rate controllin g step. However,this has been found no t to be the case in the Ca CO , system below 950C (Hyatt et a / . ,1958), and is also assumed here not t o be rate conlroll ing for the N aH CO , system du eto the relat ively o pen m atr ix produced in the p roduct solid (Howatson et a / . , 1977).The app roach taken here was therefore to interpret the dat a in terms of a shr inkingcore model of the reaction rate control type. The relatively good fi t between theexperimental data a nd the model results supp ort this hypothesis.Following ln grah am an d M arier (1963) the model is developed using the assu n~ pti onthat the reaction rate is proportional to the reactant/solid product interfacial surfacearea. Accordingly, the rate of formatio n of C O, by Eq . (I ) in a particle of sphericalshape with a shr inking core of radius r, is given by:

fiN,Rc =- k(4nr:)/M,.fitThis equatio n m ay be integrated under isotherm,al cond it ions af ter relating the C O,an d sodium bicarbonate reaction rates throug h the stoichiometry of Eq. ( I ) and af tercom puting the moles of sodium bicarbonate in a sph ere of density p, and radius r , . T heresult in term s of the fractional conversion of N aH C O, , X, is (see Levensp iel, 1972, fordetails):


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100 T.C. KEENER , G.C. FR AZIER A ND W.T . DAVISEquation (6) suggests a plot of [ l - 1 - X)L13] (p ,M, r , / aMA) ersus t should be astraight line of slope k . Th e data are plotted in this manner in Figures 3 and 4, alongwith the regression lines for each d at a set. Th e squ are of the correlation coefficients, r 2 ,are greater th an 0.972 in all cases, which suggests that a shrin king core model basedon a chemical reaction rate controlling step is reasonable for this system in thetem perat ure range of 225-350F.Th e values of the rate coefficients extracted from Figures 3 an d 4 are plotted inFigure 5, assuming a n A rrhenius form,

Th e regression line throug h these dat a give r 2 = 0.998. Th e Arrhenius form thereforeappears reasonable, at least over the experimental temperature range. The values forthe activation energy, E, and the pre-exponential factor, A, are as follows:

Reciprocal Temperature, OK-'FIGURE 5 Arrhenius plot for the decomposition reaction


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T H E R M A L D E C O M P O S I T I O N 1 01Th e magn itude of the activation energy is appreciab ly higher than 2-3 kcal/molewhich has been reported (Gordon and Ford, 11372) as the range of the apparentactivation energy for diffusion controlled processes. This result supports the as-sumption, above, that the decomposition reaction is kinetically controlled. However,this may n ot b e the case at higher tempe rature (e.g., 600F ( 3 1YC), or above) , where lossof surface are a (and porosity) has been repo rted to o ccur du e to sintering (Stern, 1978).Finally, Eq. (6) may be used to c om pu te the time, T, for 100% decomposition(X + 1.0) of a N aH C O , particle of radius r, . This time is given by:

This result was verified by calculating values of r , by use of the experimental data,Figures 3 a n d 4, which were then com pared with values comp uted according to Eq. (9)with k given by Eq s. (7) an d (8). Th e experimenta l 7,'s were established by regressing o n[ I - 1 - X)'I3] VS.&IT, s suggested by the form of Eq. (6). As r e was unknown a t theoutset , an i terative procedure was required. The %value f T, accepted was that valuewhich yielded the maxim um value of the correlation coefficient, r . This approach wastaken because the experimental method was not considered sufficiently reliable toprovid e precise values of T,, as explained abov e. The results of this analysis are show nin Fig ure 6, whe re E q. (6) with th e indicated value!; of T, is comp ared with the da ta forsix representative particle size-temperature sets. The r , values are compared w ith thosecomputed by Eq. (9 ) in Figure 7, yielding a value of r2 = 0.96. Equation (9) therefore

1.0X

9 0 . 8I0z-0 0.6CQul A - 51 urn, 2 5 0 '~ , ,=10200s.b 0 - 51 urn, ~O O' F, e=llOOs.0.4 - 87 urn, 22 5 '~ . e=54000s.00 0 87 urn, 35 0 '~ . e=570s .-o V -140um, 2 7 5 ' ~ ~= 1 0 1 0 0 s5 0.2.- -140 urn, 3 5 0 ' ~ . e 9 5 0 5.+u - odel, Eq. 62LL

00 0 2 0.4 0.6 0 8 1.0

Normahzed Decornpos~tlonTime, tireFIGURE 6 Evaluation of th e time required for 10% decomposition of sodium bicarbonate.


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T.C. KEEN ER, G.C. FRAZIER A ND W.T. DAVIS

4Experimental Decomposrt~onTime, 5 , 10 SecondsFIGURE 7 Comparison or the experimentally determined decomposition times with predictions ofEq . (9).

appears adequate for estimating decomposition times, but caution must be taken inapplying i t at tem peratures which are appreciably higher than the range used here.

C 0 2 E V O L U T I O N R A T EA useful equa tion ca n now be developed for compu ting the ca rbo n dioxide release rate,R; , from a bed of particles of size, r,7,and of uniform temperature. The appro ach is toexpress the fraction al conversion, X, Eq. (6 ) , n terms of the radius, r , , of the shrinkingcore and eliminate r , between the result a nd Eq. (5). Th e result is as follows, afterexpressing the param eters in term s of 7 and m,, the intial mass of N aH CO , in the bed:

Thus, R; can be computed once 7 has been estimated according to Eq. (9). Implicitin Eq. (10) is the ass ump tion tha t Eq. ( I ) is irreversible. Such is the case in thoseenvironments in which the product of the CO, and water vapor partial pressure isappreciably lower than the equilibrium partial pressure product.Th e equilibrium partial pressure product may be comp uted from Eq. ( 1 I), in whichthe activities of the solids in Eq. ( I ) are take n as unity a nd t he activities of the gases aretaken as equal to their partial pressures, which restricts application t o those situation swhere the gas phase is essentially ideal:


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T H E R M A L D E C O M P O S I T I O N 103TABLE l

Equilibrium partial pressures of carbon dioxide at selec1t:d temperatures and relative humiditiesPartial Pressure CO,, p, , atmTemperature Equil. Const."

"C K RH = 10% RH = 20%

'Computed by use of the data of Barner and Scheuerman (1978).Th e equilibrium partial pressure of c arbon dioxide is shown in Table I for selectedtemperatu res a nd w ater v apor partial pressures (expressed a s relative humidities). Th evalues of the equilibrium constant, K , are c omp uted from the second of Eqs. (1 l),where the free energy change for Eq. ( I ) , AGO, ir; com pute d by the use of the dat aprovided by Barn er and S cheuerm an (1978).Th e data of Table I suggest, for example, th at for a system a t 150C (302 F) and lo'%,relative humidity, back pressure of CO , appreciably less than the equilibrium partialpress ure of 1.4 x 10, atm should not affect the decomposition rate of sodiumbicarbonate. Just how much less is not know n at this time, although a singleex perimentperformed in this laborator y with a CO , partial pressure of 0.01 atm , a relativehumidity of not more than lo"/,, at a temperature of 300F did not extend thedecomposition time over that measured for essentially zero CO, back pressure.Experim ental limitations prevented a n extension of this work to higher CO , partialpressures. Th e data of Table 1 seem to suggest that water v apor a nd C O, partialpressures normally found in flue gases (no more I.han a few tenths of an atm osph ereeach) should not affect NaH CO , decomp ositior~ imes appreciably, so long as thetemperature is above about 150C. A few experiments to confirm this expectationwould be useful. The rate of decomposition of NaHCO, is of interest because it hasbeen found to influence the rate of reaction between Na H CO , and S O, in simulatedflue gas (Keener, 1982). Th e results of this investigation are re porte d elsewhere.

C O N C L U S I O N S1. T he thermal decomposition of N aH C O , granules is chemical kinetically controlledin the temperature range of 225-350F (380--45010. The shrinking core modelprovides a good fit to the dat a.2. The rate coefficient, k, is of Arrhenius form, with an activation energy of 20.5

kcal/mole (85.8 kJ/mol).3. The decomposition times, 7, can be estimated by use of the formula:T = P A M ~ ~ J ( ~ ~ N ~ )

4. T he decomposition rate of N aH CO , in the tempera ture range of 150C (302F) isnot affected by CO , partial pressures as high as abo ut 0.01 atm. Comp ariso n of C O ,


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104 T.C. KE EN ER, G.C. FRAZIER A N D W.T. DAVISequilibrium partial pressures, computed by the use of available thermodynamicdata, with those normally found in flue gas suggest that the decomposition ofN aH C O , should not be affected by conditions prevailing in flue gas so long as thetemperature is about 150C (302"F), although experimental confirmation of thisexpectation would be useful.

A C K N O W L E D G M E N TThis work wassupported by the U S . Department of Energy, through the G ra nd Fo rksEnergy Technology Center , under contract GFETC DE-RD18-80FC10184 and wascond ucte d in the Civil Engineering Dep artm ent of the University of Tennessee.

N O M E N C L A T U R Emoles of NaH CO ,Arrhenius pre-exponential factor, g/cm2sstoichiome tric coefficient, 2 moles NaH CO ,/mo le CO ,moles of C 0 2Arrhenius a ctivation energy, cal/g-molereaction rate coefficient, g/cm2sequilibrium constan tmass of Na HC O, sample , gmolecular weight of NaHCO,molecular weight of CO,number of particles in samplemoles of CO , per particlepartial pressure of ca rbo n dioxide, atmpartial pressure of water vapor, atmgas constantrate of CO, released per particle per unit time, moles/srate of CO , released by sample, moles/s = I I R ,linear regression correlation coefficientradiu s of un-reacted core, cminitial radius of NaHCO, particle, cmtemperature, Ktime, sfractional conversion of N aH CO , particlesta nda rd free energy change for the reaction, cal/g-mol


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T H E R M A L D E C O M P O S I T IO Np, density of NaHCO,, g/cm3T time for 100% decomposition, s(E q. (9))z e experimentally determined time for loo"/, decomposition, s

R E F E R E N C E SBarner. H.E., and Sc heuer man , R.V., Ilut~rlhool;f Thermochemical Data ur C ompoutlds and Aqueous Species,John Wiley and Sons, New York (1978).Carson, John R., "Removal of Sulfur Dioxide an d N itric Oxide from a F lue Gas S tream by Tw o SodiumAlkalis of Various Sizes", M asters Thesis, University o f Tennessee, Knoxville, August 1980.Church and Dw ight Sodium Bicarbonate (manufacturers brochure). Chur ch and Dwight Co., Inc., New York,

N V. ..Darroudl, T., and Searcy, A.W., J. Phys. Chem., 85,3971-74(1981).Ewing, J., Beruto, D.. and Searcy, A.W., J. Am. Cerum. Soc.,6;!.580-84(1979).Furnas, C.C.,AIC hE Transoctiotls, XXIV, 142 (1930).Go rdo n, A.J., and Ford , R.A., The Chernist's Companion: A lfandhook of Practical Data, Techniques andRejerences, John Wiley and Sons. New York (1972).Hill, K.J., and Winter, E.R.S., J. Phys. Chem.. 60, 1361-62 (1956).Howatson, J., Smith, J. Ward, Ou tka, D.A., and Dewald, H.D., "Nahcolite Properties Afecting Stack Ga sPollutant Absorption", Proceedings of 5th. National Conference on Energy and the Environment,Cincinnati, Ohio, November, 1977.Hyatt, E.P., Culler, I.B., and W adsworth, M.E., J. Am. Ccram. Soc.,41, 70(1958).Ingraham, T.R., and Marier, P., Con. J .Ck en~ . ng., 41, 170(1963).Keener. T.C., "Thermal D eco m po sih n o f Sodium Bicarbonate and Its Effect on the Reaction of SodiumBicarbonate and Sulfur Dioxide in a Simulated Flue Gas", Ph.D. Dissertation, University ofTennessee (1982).Kehat, E., and Markin. A,, Can. J. C h m . Eng., 45.40 (1963).Knight. Joh n H.. The Useof N ahm litefor Rem ooalff ulfur Dioxide and Nifroqen Oxidesfrom F lue Cas, Th eSuperior Oil Company (1977).Kunii, D., an d Levenspiel. 0 . . Fluidizrttion Enqineering, John 'Wiley & Sons, Inc.. New York (1969).Langmu ir, I., J. Am. Chem. Suc..38, 2263 (1916).Levenspiel, 0..hen~ical eaction Engineering. 2nd edition, John Wiley and Sons, New York (1972).Mac Don ald, J.Y., T rans. Farudoy. Soc.. 47, 860 (1951).Narsimhan, G., Chem. Engr. Sci., 16, 7 (1961).Perry, R.H., and Chilton, C.H., Cl~enlicol nqineer's Ilnndhook, 5th edition, McG raw Hill, New York (19731.Roberts, J.A., Jr., Jacobson. N.S., and Searcy, A.W., J. Chem. .Phys., 69, 5562 (1978).Stern. Frederic R., "Bench-Scale Study of Sulfur and Nitrogen Ox ides Ads orption by Nahco lite and Trona" ,Masters Thesis, University of North Dakota (1978).

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