INFRARED SPECTROSCOPIC STUDY OF GAS-SOLID INTERACTIONS
Edwin F. Rissmann
National Air Pollution Control Administration Technical Center
Durham, North Carolina
April 1970
NATIONAL TECHNICAL INFORMATION SERVICE
Distributed , • ,'to foster, serve and promote the
nation's economic development
and technological advancement.
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INFRARED SPECTROSCOPIC
STUDY OF GAS-SOLID INTERACTIONS
FINAL REPORT
APRIL 1970
Process Control Engineering Division
National Air Pollution Control Administration
Consumer Protection and Environmental Health Services
Public Health Service
U. S. Department of Health, Education and Welfare
Contract No. CPA 22-69-59
Submitted by
General Technologies Corporation
A Subsidiary of Cities Service Company
1821 Michael Faraday Drive
Reston, Virginia 22070
by
Edwin F. Rissmann
NATIONAL XECHNKAL
INFORMATION SERVICE
tUKKRO TITLE PjSE
FOB TEOffiCM. RETORTS
. (Upon no.
, APTP-0.586.
Infrared Spectroscopic Study of Gas-Sol 1d'Interactions
f. Authoni)
- Rissmann
9. Perfoniln
•hit Orpnlallon Nan end Arjfttn
General Technologies Corporation
A Subsidiary of Cities Service Company
1821 Michael Faraday Drive
Reston. Virginia 22070
12. Sponsoring Afency HMO i
National A1r Pollution Control Administration Technical Center
411 West Chapel H111 Street
Durham. North Carolina 27701
3. Recipient i Catalog No.
S. Report Dite
iril 1970
organization Code
8. Pertorailnj Organization Rent. No.
10. Project/TMk/Work Unit No.
11. Contract/Grant No.
CPA 22r69-59
13. Type of Report & Period Covered
14. Sponsoring Ajency Code
16. Atutraca .The aim;was to-apply infrared spectroscopic methods Co handle large numbers or
process control samples and to obtain kinetic information to evaluate the effects of var-
iables such as sulfur dioxide concentration on overall process efficiencies and to under-
stand the deadburnlng phenomenon associated with limestone processes for remo"aL of sul-
fur dioxide from flue gases. Major effort was directed toward development of rapid reli-
able methods for analysis of reacted limestones. A solvent system, suitable for dissolving
limestone and capable of being made infrared inactive by the dual cell technique was
developed. The system—tetrasodium EDTA (ethylenediaminetetraacetic acid) saturated water
was found to be the only one of a large number of systems investigated capable of dissol-
ving limestone. Special iirfrare^ liquid cells J-aicions thick were developed." The feasi-
bility was studied of using a D2p based solvent system to determine the. oxide and hydro-
xide contents of reacted limestones. The KBr pellet technique_-utth and without use.-of a
internal standard*'was studied in considerable detail. Ion exchange resin techniques were
also studied in conjunction with the KBr method. Infrared spectroscopic techniques were
qlso applied to studies of gas-solid interactions between the sulfur djasidje. content of
17. Key Wonfe ind Oocunent AnX-nli. (•). DMcrlptora
Air pollution
Infrared spectroscopy
Sulfur dioxide
Limestone
Air pollution control equipment
Process control
Reaction kinetics
Solvents
EDTA
Infrared equipment
Concentration (composition)
pH
Feasibility
Heavy water-
Potassium bromide
Ion exchange resins
Sulfates
Hydroxides
Carbonates
Silicon oxides
Roasting
(over)
Temperature
Hydration
Magnesium oxides
Dolomite (rock)
Absorption
Oxidation
Infrared cells
Deadburning
I7e.
13/02, U/02. 07/02
19. Security cuts (This Report)
UNCLASSIFIED
XLSwtrlty Clut.mil* Page)
UNCLASSIFIED
1. No. of Pages
83
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16. Abstracts (cont'd)
simulated flue gas and various limestone absorbents and to an investigation of the dead-
burning phenomenon. The latter has revealed the strong possibility of1: chemical reactions
with silicate Impurities during high temperature calcination. Electron nlcroprobe data
has revealed that at lower calcination temperature, all of the silicates are pre.sent in
discrete phases In the limestones .-W4th__incr easing calcination temperatures-silicates
were found to diffuse throughout the material. -Studies were also perfo^jgd to confirm
the hydration of HgO in calcined, slaked, dolomltic materials. •"
INFRARED SPECTROSCOPIC
STUDY OF GAS-SOLID INTERACTIONS
FINAL REPORT
APRIL 1970
Process Control Engineering Division
•National Air Pollution Control Administration
Consumer Protection and Environmental Health Services
Public Health Service
12. 5. Department of Health, Education and Welfare
Contract No. CPA 22-69-59
Submitted by
General Technologies Corporation
A Subsidiary of Cities Service Company
1821 Michael Faraday Drive
Reston, Virginia 22070
by
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ABSTRACT
This report constitutes work conducted during the second year of effort to study gas-
sol id interactions by infrared techniques. A report on the first year's work has been
published by the Clearinghouse for Federal Scientific and Technical Information under Public
Health Service Number PB 182-988.
The major effort in this investigation has been directed toward development of rapid
reliable methods for analysis of reacted limestones. A solvent system, suitable for dissolving
limestone and capable of being made infrared inactive by the dual cell technique was
developed during riiij contract, The system—tetrasodium EDTA (ethylenediaminetetraacetic
acid) saturated water was found to be the only one of a large number of systems investi-
gated capable of dissolving limestone., Of the many organic systems studied, only solvents
containing OH groups, which would dissolve EDTA were of any use. Trie aqueous solution
was found to be best in that limestone solubility from it was highest.
To use such a system for infrared studies, jpecial infrared liquid cells 3 microns
thick ha«( (a be developed.^Using these cells and the new solvent system, calibration
xEurves were obtained for various concentrations of carbonate and sulfate and a number of
/ limestone samples were analyzed for sulfate content. The analysis results compared
;' favorably with values obtained by wet chemical methods and a detailed error analysis of
I the new method revealed the importance of solution pH as a variable in the analysis.
' Dissolution of limestone into the saturated tetrasodium EDTA solutions generally caused
i an increase in pHr the amount of which depended on the amount of sample dissolved and
\ ill unreacted oxide content.
Utilizing the thin infrared cells, studies were also made of the feasibilityToF using
a I>2O based solvent system to determine the oxide and hydroxide contents of reacted
Ihnestones. \ From the results obtained, such a procedure does appear to have promise.
••C^A number of other infrared analytical techniques were also investigated under
this contract.'- The KBr pellet technique, with and without use of an internal standard
was studied in considerable detail..Jt was found that these methods yielded results which
were subject to considerable variation in accuracy and precision. Ion exchange resin
techniques were also studied in conjunction with the KBr method. With this approach,
reliable results were obtained for sulfate analysis but analyses for hydroxide and carbonate
proved impossible one reproducible basis.
Infrared spec&oscopic techniques were also appl ied to studies of gas-sol id interactions
between the sulfur dSoxide content of a simulated flue gas and various limestone absor-
bents and to an imimHtjution of the deadbuming phenomenon. , In situ spectroscopic studies
have been conducted in a high temperature infrared cell in the 3?5-475°C range and
kinetic data has been obtained on two of the individual steps for the overall process
reaction. It has been found that at elevated temperatures the absorption reaction:
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with an activation energy of 13.8 Kcal/mole is rate determining. At lower temperature,
the oxidation process yielding sulfate, with an activation energy of 41.9 Kcal/mole
becomes rate controlling.
Investigations into the mechanisms of deadburning hav; revealed the strong possi-
bility of chemical reactions with silicate impurities during high temperature calcination.
Indeed, shifts in position and changes in intensity with rising calcination temperature
have been noted for the Si-O stretches due to silicate impurities in three series of lime-
stones examined. Electron microprobe data has revealed that at lower calcination tem-
perature, all of the silicates are present in discrete phases in the limestones. With
Increasing calcination temperature, silicates were found to diffuse throughout the material.
Studies were also performed during thii contract to confirm the hydration of MgO
in calcined, slaked dolomitic materials.
TABLE OF CONTENTS
INTRODUCTION AND SUMMARY
EXPERIMENTAL RESULTS AND DISCUSSION
A. ANALYTICAL STUDIES
Page
1
4
4
1 . Introduction
2. Pellet Technique Studies
(a)
(b)
Experimental Procedures
Pellet Studies - Results and Discussion
(1) Application of the KBr Technique
to the Quantitative Analysis of
Sulfate
(2) Infrared Quantitative Analysis of
Sulfate Utilizing An Internal
Standard
(3) Application of the KBr Technique
to the Quantitative Analysis of
Carbonate, Hydroxide and Silica
(4) Use of Freeze Drying and Ion Exchange
Techniques
3 . Infrared Solvent System
(a)
(b)
(e)
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c.
TABLE OF CONTENTS (cont.)
1. Introduction
2. Experimental
3. Results and Discussion
(a) Calcination Temperature Studies
(b) Electron Microprobe Studies
(c) Studies of H/dration of Calcined Limestone
HIGH TEMPERATURE KINETIC STUDIES
1.
2.
3.
4.
General Introduction
Experimental
Results
Discussion of High Temperature Study Results
(a) Mechanism of Reaction
(b) Calculation of Pre-Exponential Factors
for the Absorption and Oxidation Steps
(1) Oxidation of Sulfite to Sulfate
(2) Absorption of SO2
III CONCLUSIONS AND RECOMMENDATIONS
IV REFERENCES
V ACKNOWLEDGEMENTS
Page
45
46
46
46
51
53
57
57
57
60
64
64
67
67
67
69
73
75
lv
LIST OF ILLUSTRATIONS
Figure
I Freeze-DryingCell
2 Plot of Corrected Absorbance of 1158 cm"' Sulfate Absorption
Band Versus Wgt. % of CaSO4 in KBr
3 Plot of Corrected Absorbance of 678 cm"'Sulfate Absorption
Band Venus Wgt. % of CaSO4 in KBr
4 Plot of Corrected Absorbance of 1158 cm"' Sulfate Band
Versus Wgt. % CaSO4 in KBr
5 Plot of Ratio of Absorbance of 1158 cm'1 Band of CaSO4
to 2098 Band of Pb (SCN)2 Versus Wgt. % CaSO4 in Standard
6 Plot of Ratio of Absorbance of 678 cm"' Band to 2098 cm"' Band
of Pb (SCN)2 Versus Wgt. % CaSO4 in Standard Mixture
7 Plot of Corrected Absorbance of 1434 cm"' Carbonate Band
Venus Wgt. % CaCO3 in KBr
8 Corrected Absorbance of 878 cm"' Carbonate Band Venus
Wgt.%CaCO3JnKBr
9 Plot of Corrected Absorbance of 3658 cm"1 Ca(OH)2 Band
Venus Wgt. % Ca(OH)2 in KBr
10
11
12
13
14
15
16
Plot of Corrected Absorbance of 1086 cm"1 SiO2 Absorption
Band Venus Wgt. % SiO2 in KBr
Absorbance Venus Concentration for Sulfate at 622 cm"'
from Ion Exchange Method Studies
Carbonate Calibration Curves - Ion Exchange Studies
Sulfate Calibration Curve - Ion Exchange Studies
Spectra of Carbonate and Sulfate in Aqueous Films
Infrared Cell for Use with Limestone Solvent
Calibration Curve with .003 mm Cell for Sulfate Ion at 1110 cm"'
Page
11
'2
13
15
16
17
18
20
21
22
27
28
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LIST OF ILLUSTRATIONS (cent.)
Figure
17 Calibration Curve with .003 mm Call for Carbonate Ion
at 1400cm-1
18 Spectra of EOTA Solution Alone and EDTA Solution
Containing Dissolved Reacted Limestone
19 Percent Error Versus Solution pH for Limestone Analyses
20 Infrared Spectrum of the Disodium Salt of Ethylenediaminetetraacetic
Acid (KBr Pellet)
21 Effect of Sample Sulfate Content on Solution pH Change
22 Percent SOg Variation as a Function of Cell Age
23 Percentage of Total Band Shifts Observed as a Function
of Calcination Temperature
24 Line Analysis Microprobe Run on Series 100 1700°F Particle
25 Electron Microprobe Line Scan for Silicon on 3200°F Series 100
Particle
26 Preferential Reaction of Hydrated MgO with SO2
27 Design of High Temperature Cell for Use in Dual Cell System
28 Diagram of System Employed to Hold Windows in Place
29 Reaction CaO + SO2 Monitoring 900 cm
30 Reaction CoO + SO2 + 1/2O2 Monitoring 1110 cm"1
31 Kinetic Plots of Data for SO2 Absorption and Oxidation Steps
Page
35
36
38
39
41
42
50
52
54
56
58
59
62
63
66
vi
LIST OF TABLES
Toble
I Solubility of Limestone and EDTA
II Classification of Solvents Used by Polar Group*
Spectra of Insoluble Limestone Residues
Analyses of Limestones
Effects of Solution pH on Analysis Errors
Compositions of the Series of Limestones Used in Deadburning Studies
Spectra of Limestones as a Function of Calcination Temperature
Ill
IV
V
VI
VII
VIII
IX
X
Hydration of MgO and Evidence for Preferential Reaction Between
MgO and Sulfur Dioxide
Rate Data for the Reaction CaO + SO2 + 1/2O2—»CaSO4
Kinetic Data of SO2 Absorption
Page
25
25
30
33
37
47
48
55
65
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*-.£,- - - ;
SECTION I
INTRODUCTION AND SUMMARY
The removal of sulfur dioxide from power plant flue gates has recently become a
subject of considerable concern. Several processes have been proposed to satisfactorily
remove this pollutant by converting it into either saleable or non-toxic, easily discard-
able, products. Two potentially economical methods, which result primarily in formation
of non-toxic discardable calcium sulfate, are the wet and dry limestone scrubbing processes.
For these techniques to be adequately studied, several problems must first be solved.
New analyses which are sufficiently rapid so as to be able to handle large numbers of
process control samples must be developed. The classical wet chemical techniques already
in use are simply too time consuming If a number of reacted limestone constituents must be
determined on a routine basis. Also, kinetic information must be obtained so as to be
able to evaluate the effects of variables such as sulfur dioxide concentration on overall
process efficiencies and existing problems, such as the deadburning phenomenon, must be
understood and overcome.
The aim of this program has been to utilize infrared spectroscopic methods to identify
and solve some of the above problems. Specifically, three areas have been studied in
considerable detail:
1. Much effort has been directed, during this study, to the development of a
rapid, accurate infrared method for the routine analysis of the constituents
of reacted limestone.
2. The kinetics of the overall reaction between oxygen, sulfur dioxide and calcined
limestone were investigated in considerable detail in the 357-473°C range.
3. A study was made of the possible influences of silicate impurities in the
deadburning (i.e. loss of reactivity) of limestone as a function of calcination
conditions and the hydration of MgO in slaked calcined dolomitic materials
was confirmed.
For the first task to develop rapid, accurate infrared analytical techniques for
limestone, a number of techniques were studied in considerable detail. The KBr pellet
method, with and without the use of an internal standard was investigated as a technique
for rapid analysis. It was found that the calibration curves obtained for sulfate, carbonate,
silicate and hydroxide contained considerable scatter and that sulfate analyses run on a
number of samples gave results of. In general, poor reproducibil try and accuracy. Attempts
were mode to improve the KBr pellet method by using it in conjunction with standard
freeze drying ion exchange techniques. While reproducible and fairly accurate sulfate
analyses were found to result from this, useable calibration curvet could not be readily
obtained for either carbonate or hydroxide, thus I imiting the usefulness of this approach
to sulfate analysis.
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The use of o solution infrared method which would provide a reliable limestone
analysts was therefore investigated. With this method, a solution of limestone would
be placed .in the sample beam of the spectrometer. A cell of equal thickness filled
with only the solvent would be used in the reference beam, in order to develop such
a method, the first step was the selection of a suitable limestone solvent. Of the many
solvent systems investigated, only tetrasodium EDTA (ethylenediaminetetraacetic acid)
saturated waterl") was found to effectively dissolve lime and limestone in sufficient amounts
to be useful.
An aqueous solvent, however, presented the problem of developing infrared cells
suitable for such work. Water, in .thicknesses of greater than 10 microns is almost totally
absorbing,in most of the infrared region(2) and the thinnest commercially available cells
are about 7 micron* in thickness. In past quantitative studies with aqueous solutions, thin
film cells were prepared by deposition of a film of cellulose acetate from acetone solution
onto the edges of flat salt plates(3). However, while such an approach is excellent for
preparing single cells, it is useless for making a matched pair of cells. The thin film
cells used in our studies were prepared by deposition of a 3 micron silver film onto the
surfaces of flat silver chloride plates.
Using'a matched pair of such cells, cal ibration curves were obtained for sulfate and
carbonate dissolved in saturated tetrasocTium EDTA solution and a number of limestones
were analyzed for their sulfote contents. The results agreed well with those obtained from
classical wet chemical methods of analysis. A detailed error analysis performed on the
data obtained with about one dozen different I imestones revealed the importance of solution
pH as a variable influencing accuracy of the results. ; :f
For the second task kinetic studies were conducted in the 375-475°C range using a
high temperature dual cell arrangement to minimize interference from emissions problems '-
that have been encountered in earlier work(4). In these studies, with both cell com-
partments maintained at the same temperature, the oxidation and absorption of SOj on
limestone were studied in detail. Separate series of kinetic experiments, with and without
oxygen present in the gas phase, were performed to measure the absorption of SO; and the
formation of sulfate. The data collected was obtained by two techniques; monitoring of
a fixed frequency (1110 cm"' for sulfate and 900 cm"' for sulf ite) and scanning of the
1300-700 cm-' region at regular intervals. Both sets of data proved to be in reasonable
agreement with each other. From the sulfate formation rate data, an activation energy
of 41.9 K cat/male was computed for the reaction:
step:
+ 1/202 »CoS04 (1)
From the 900 cm'1 data, a value of 13.8 Kcal/mole was found for the SOj absorption
CaO + SO2 > CaSO3 (2)
' ' From further calculations of pre-exponential facton, It was then demonstrated that at
teiiycrurung above 500°C, the absorption step is the overall rate determining reaction for
the dry limestone process. Below 500°C, the formation of sulfate via reaction I becomes
rote controlling.
For the third task, studies were made in some detail of the possible role silicates
may nave in the deactivation of limestone (deadburning) during high temperature calcin-
ation. Specifically, three series of limestone materials were studied. Spectra were run
by the KBr pellet technique of these materials prepared at various calcination temperature.
Analysis of all the spectra obtained revealed changes in intensity and band positions for
the SI-0 stretching bands due to silicate impurities present. This confirms earl ier pre-
liminary work<4) in which similar band shifts were reported. Aside from the changes in
the sil icate bands, no other significant differences in Infrared spectra between high
and low temperature calcines of the same materials could be noted. Thus, the differences
in Infrared spectra have been interpreted in terms of chemical reactions between the silica
and/or silicates present and the calcined limestone. Studies using the electron microprobe
have found that in the low temperature calcines, essentially all of the silicates are con-
tained in discrete phases within the limestone particles. For the higher temperature calcines
some silicates were found in all phases of the materials.
An investigation aimed at confirming the hydration of magnesium oxide in calcined,
slaked dolomitic materials was also conducted. It was shown that in these unreacted
materials, hydration, indeed, does occur. It was further demonstrated that, upon re-
action with flue gas at 750°C, the hydrated MgO either decomposes or reacts preferen-
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'••-l-SS,
SECTION II
EXPERIMENTAL RESULTS AND DISCUSSION
The work performed under this contract was aimed at the solution of three discrete
sets of problems:
t. Analytical investigations,
2. Studies of deadburning and MgO hydra t ion and,
3. High temperature kinetic studies.
In this report, each of these areas will be treated, in whole, in a separate section. This,
it is felt, will yield a more easily understandable presentation.
A. ANALYTICAL STUDIES / .
1. Introduction
One of the uiins of this program has been to develop infrared methods for
rapid analysis of the constituents of reacted I imestones. Two types of analytical methods
were investigated in considerable detail. The first was the pressed pellet method and the
second was the solvent technique. In studying appl ication of the KBr technique, inves-
tigations were also made of application of freeze drying and ion exchange techniques in
an effort to gain greater accuracy. Investigations of the solvent approach centered
around the choice of a satisfactory limestone solvent and the design and use of an infrared
cell system compatible with he solvent chosen.
In this portion of the report, the pellet and solvent technique studies will
be treated completely in separate sections for ease of reading.
2. Pellet Technique Studies
(a) Experimental Procedures
For studies using the KBr pellet method, the following general pro-
cedures were employed:
Pel I ah containing from 0.01 to 1.0 weight percent of the desired
material were prepared by grinding the appropriate amount of this material together with
potassium bromide in the Wig-L-Bug for from 30 to 60 seconds. The KBr pellets were
then prepared by pressing the samples in an evacuable die by applying a pressure of 2700
pounds for two minutes. The infrared spectra of the pellets were recorded using a beam
attenuator in the reference beam to adjust the 100% base line and calibration curves
were constructed for the various species of interest by plotting the absorbancei at specific
frequencies versus concentration of that species. For all of this work, the Beckman IR 10
Infrared Spectrophotometer was used.
In some studies using the pellet method, a pure KBr disk was used in
the reference beam in place of the attenuator. It may be noted that this approach did
not yield any significant Improvement in the results.
For studies using an internal standard, the above approach was modi-
fied by incorporating a known weighed amount of internal standard along with the material
of interest into the KBr pellets. The internal standard used was lead thiocyanate. In the
application of this technique, the internal standard was prepared by adding 0.23% by
weight of lead thiocyanate (preground for 10 minutes in the Wig-L-Bug) to 4.9327 grams
of KBr which had been dried at 120°C. This mixture was then carefully mixed and ground
in aliquots and then stored in a dessicator over phosphorous pentoxide prior to use. In
preparing calibration plots with the internal standard technique, correlations were made
between the ratio of the absorbance at one frequency of the species of interest to the
absorbance of the thiocyanate bands with the concentration of the desired species in the
KBr pellets.
To obtain a more mixture of samples in KBr, the use of freeze-drying
techniques were employed in some investigations. The basic design of the cells used are
shown in Figure 1. They consisted of 12/30 and 24/40 standard taper glass joints and a
four millimeter bore stopcock. For the freeze drying experiments, the solutions were
frozen using a dry ice acetone mixture and the water removed under vacuo, leaving a
dried uniform mixture which could be pelletized. Infrared spectra of the samples were
measured in the usual manner.
Two variations of the freeze drying technique were investigated. In
the first technique, a standard calcium sulfate solution was prepared by dissolving a known
amount of this material in the minimum amount of 1M hydrochloric acid and diluting the
solution to 100 milliliters. Aliquots of this solution were mixed with a standard solution
containing 0.03323 gram of potassium bromide and 0.000054 gram of potassium azide
per milliliter. The potassium azide was added for use as an internal standard. The mix-
tures were frozen using both a dry ice acetone bath and I iquid nitrogen. The water was
removed under vacuum leaving the resultant mixture in the form of a powder. This powder
was then dried in an oven at 120°C for approximately two hours and potassium bromide
pellets were prepared after grinding the mixtures for 15-30 seconds in the Wig-L-Bug,
The second technique consisted of using ion-exchange resins to dissolve
the calcium culfate. A number of resins in the Na+, K+ and H+ forms were utilized in
these studies. Problems encountered with these materials will be treated in a later
section of this report.
After suitable calibration curves for sulfate had been obtained with
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To Vacuum System
A
GTC 191-8
12/30 Standard Taper
Glass Joint
4 mm Bore Stopcock
24/40 Standard
Toper Gtau Joint
Figure 1. Freeze-Drying Cell
above procedures for sample preparation.
(b) Pellet Studies - Results and Discussion
(1) Application of the KBr Technique to the Quantitative Analysis of
Sulfote
The application of the KBr technique to the quantitative analysis
of limestone samples can be best understood by considering the Beer-Lambert Law given
in Equation 3.
A = act
(3)
where A is the absorbance. In {•A, a is the absorption coefficient, c is the concentration
of calcium sulfate in terms of weight percent in potassium bromide and I is the path length
of infrared radiation through the pellet (thickness of the pellet in mm). Fora particular
substance, for example, calcium sulfate, the absorption coefficient should be constant and
the absorbance depends upon the concentration of the absorbing species and the path
length of infrared radiation through the sample. The standard calibration plot is shown in
Figure 2 for the 1158 cm*' bond of calcium sulfate and the standard calibration plot
using the 678 cm"' band of calcium sulfate is shown in Figure 3. In both of these plots,
the best least square fit line is shown along with the experimental points.
In order to check the accuracy of this technique for the quanti-
tative analysis of sulfate, the absorptivity of the 1158 cm*' and 678 cm*' sulfate ab-
sorption bands were calculated for each of the KBr pellets containing the sample. The
absorptivity is given by Equation 4.
k = A <<>
where k = absorptivity
A = baseline absorbance
c = concentration in weight percent
I - thickness of the KBr disc in millimeters.
For each pellet, the value of k should be constant. However, results showed there was a
deviation in the value of k for both the 1158 cm*' and 678 cm*' absorption bands. The
percent coefficient of variation of the mean was 6.7% for the 1158 cm*' absorption band
and 8% for the 678 cm*' absorption band. The amount of scattering of points is also
indicated in the standard curves for these absorption bands.
Standard calibration curves of calcium sulfate In potassium
bromide were also prepared by correcting the absorbance against a pellet containing
pure potassium bromide which was placed in the reference beam of the spectrophotometer
in place of the beam attenuator. The absorbance of the sulfate band was corrected
-------
GTC 191-2
2.4
•s
,..
0.4
.05
0.1 0.15 0.2
Wgt. % of CaSO4 in KBr
0.25
0.3
0.35
Figure 2. Plot of Corrected ADsorbance of 1158cm*1 SolFate
Absorption Band Venus Wgt. % of CoSO4 in KBr
GTC 191-3
Wgt. %ofG»S04inKBr
0.35
Figure 3. Plot of Corrected Abforbanee of 678cm"1 Sulfor*
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weight of KBr disc in reference beom
weight of KBr disc containing calcium sulfate
The weight of the potassium bromide disc in the reference beam was 706.4 mg. The calibra-
tion curve for the 1158 cm"' sulfate absorption band is shown in Figure 3. There is still,
however, quite a bit of scattering of points.
The calibration plots of Figures 2, 3, and 4 were used to analyze
sample No. BCR 1699 which was exposed to flue gas at 1800°F. The amount of sulfate
found by the first infrared spectroscopic technique was 26.1% using the calibration plot
constructed from the absorbance of the 1158 cm'1 sulfate band and 25.4% using the cali-
bration plot of the 678 cm"' sulfate band. The amount of sulfate found by the second
infrared spectroscopic technique was 19.9% CaSC>4 utilizing the plot of Figure 4. The
percent sulfate found in sample No. BCR 1699 expressed as SC>3 as reported by the National
Center for Air Pollution Control was 16,5%.
(2) Infrared Quantitative Analysis of Sulfate Utilizing An Internal
Standard
Calibration plots obtained using the internal standard method are
shown in'Figures 5 and 6.
The potassium bromide disc technique employing an internal
standard is explained in the following manner. Let the absorbancy of a known material
to be assayed at wavelength, X^, be given by
Ak = ak*kck (5)
where S|<, 4|< and C|< have been defined by Equation 4 and the absorbance of the internal
standard at wavelength, \. is given by
A5 = V*cs
Now, dividing Equation 5 by Equation 6, we have
Ak. = ik-t-kCk
A5 as^cs .
(6)
(7)
The t 's cancel and because Oj, and as are both constants at the wavelengths at which the
measurements are made, and c , the concentration of the internal standard is constant,
these constants can be accumulated in an overall constant, K, and we have
(8)
Therefore, a plot of A^A, versus c|< will give a straight line. In this method, it is not
necessary to determine o|( and ns or even to known es exactly in order to determine a
10
GTC
, , c
Sulfof. Bend V.m. Wgt. % CoSO4 in KBr
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GTC191-5
0.05
0.1 0.15 0.2
Wgt. % CoSO4 in Standard
0.25
Figure 5. Plot of Ratio of Abtorbance of 1158cm"1 bond of CaSO4
to 2098 bond of Pb{SCN)2 Vena Wgt. % CoSO4 in Standard
12
GTC 191-6
0.05
0.1 0.15 0.2
Wgt. % CaSOj in Standard Mixtun
-------
working curve. Also, the method eliminates the need to measure the thickness of the
potassium bromide disc in the analysis. Calibration data obtained with this method are shown i
Figures 5 and 6. Results using this technique to analyze Limestone Sample No. BCR 1699
reacted with flue gas at 1800°F showed a sulfate concentration of 13.4% utilizing the
1158 cm*' sulfate absorption band and 15.5% utilizing the 678 cm"' sulfate absorption
band. These values are in only fair agreement with the 16.5% found by wet chemical
methods of analysis.
(3) Application of the KBr Technique to the Quantitative Analysis of
Carbonate, Hydroxide and Silica
The KBr pellet techniques, without use of an internal standard
developed for sulfate were extended to carbonate analysis. The calibration curves for the
1434 cm"' absorption band and the 878 cm"' absorption band are presented in Figures 7 and
8. From the data obtained, the plots indicated a deviation from the Beer-Lambert Law
since curved lines are obtained instead of straight line plots. The best free-hand curve has
been drawn through the calibration points. The KBr technique was also extended to
analysis for hydroxide and silica. Calibration curves were obtained by the same methods
used for carbonate. They are shown in Figures 9 and 10.
(4) Use of Freeze Drying and Ion Exchange Techniques
Freeze drying techniques were studied both in conjunction with the
use of potassium azide as an internal standard and in combination with ion exchange
resins.
The work involving the use of potassium azide was unsuccessful
due to interactions and/or reactions between the azide and the HCI used in solution
preparations. Indeed, solutions of CaSO^ in dilute HCI, to which were added known
amounts of azide, when freeze dried and later converted into pellets, showed calibration
data of poor quality. Due to these difficulties, the idea of using azide as an internal
standard was not pursued further.
Two forms of the Dowex 50W-X8 sulfonate based ion exchange
resin were used in studies aimed at developing an ion exchange technique for dissolution
of limestone to be used in conjunction with the KBr pellet technique. Two of the resins
forms H+ and Na+ were obtained from commercial sources.
With the H+ form used, problems were encountered in that both
reacted with carbonate containing solutions to liberate CC^. This made this material
of little value for use in solubilizing limestone for analytical studies. The Na+ form of
the resin caused another set of unsuspected problems. When distilled water was treated
with this material and the resin then removed by filtration, the distilled water was found
to contain dissolved species which exhibited, upon infrared examination, bands in the
1450-1400 and 1200-1100 cm"1 regions which would interfere with analytical studies.
Fortunately, this problem of Impurity leaching could be solved by prelcaching of all
14
GTC 191-9
0.05 0.10 0.15 0.20
Weight Percent of Calcium Carbonate in KBr
0.25
Figure 7. Plot of Corrected Abtorbance of 1434 cm'1 Carbonate Band
Venus Wgt.% CoCO3 in KBr
-------
0.
0.6
"i
£
S 0.5
0.4
0.3
U
o 0.2
0.1
GTC 191-10
0.05 0.1 0.15 0.2
Weight Percent of Calcium Carbonate in KBr
Rgure 8. Corrected Absorfaonce of 878 cm"' Carbonate Band
Venus Wgt. % CoCCX, In KBr
16
0.25
GTC 191-11
1.4
1.3
1.2
1.1
1-0
» 0.9
03
•fc 0.8
8
J 0.7
•8
0.6
I 0.5
0.4
0.3
0.2
0.1
0.0
0.05 0.10 0.15 0.20 0.25
Weight Percent of Calcium Hydroxide in KBr
Figure 9. plot of Corrected Abtorbance of 3658 cm'1 Co(OH)2 Band
Venui Wgt. % Ca(OH)2 in KBr
-------
OTC 191-12
0.05 0.10 0.15
Weight Percent of Silloo In KBr
0.20
0.25 0.30
Figure 10. Plot of Corrected Absorbance of 1086 cm'1 SiO. Absorption Bond
Venus Wgt.%S!O2 in KBr
18
resins used. The No+ resins were preleached in warm distilled water for 24 hours twice
before use in analytical work. This procedure was found to be satisfactory in that the
leached resins were found after this treatment to no longer introduce measurable contam-
inants into the solutions being treated. It may be noted that the above leaching procedure
is similar to that recommended by the manufacturers of a similar material, Amberllte
120, prior to use(5).
Using the leached Na+ form Dowex resin, calibration curves
were made for sulfate and carbonate using the 622 and 1158 cm"' bands for sulfate,
and the 1428 and 873 cm"' carbonate absorption peaks. These curves are shown in Figure*
11-13.
The general procedure employed in these studies was to prepare
solutions of CaSOj or CaCC>3, treat them for 24 hours with the leached ion exchange resin
to convert the dissolved salts to their sodium forms, filter out the resin, and then take
aliquots of the solutions and mix them with fixed amounts of the KBr standard solution.
The resulting solutions were taken and freeze dried, the mixed materials were recovered
and pellets were then made and their spectra recorded.
Attempts were made to construct calibration curves for hydroxide
using the above techniques, however, spectra of all pellets analyzed revealed that much
of the hydroxide had, at some stage, been converted to carbonate. Special precautions,
such as prepurging of the Ca(OH)2 solutions with Nn gas did not eliminate the problem.
From studies performed, it was concluded that either the resin reacts with hydroxide to
yield carbonate or COo is absorbed from the atmosphere sufficiently readily to make
meaningful analysis extremely difficult.
The calibration data obtained above was utilized for analysis of
limestone sample No. BCR 1699. 25.8 mg of this material were weighed into a small flask
to which was added 10 cc of distilled water and 2.0933 g of doubly leached Na+ form Dewex
resin. The solution was let stand for 16 hours to complete the ion exchange process and
the reacted resin was then removed by filtration. The resulting solution was then diluted
to 200 ccand two 2 ml aliquots were taken, mixed with a fixed amount of KBr standard
solution and the resulting solutions were then freeze dried, the mixed materials recovered,
pellets made and spectra run. The spectra revealed no hydroxide to be present in either
of the two pellets studies, in agreement with the difficulties found above. Peaks due to
sulfate and carbonate were the only ones of sufficient intensity to analyze. Due to the
problems related above with conversion of hydroxide to carbonate, analysis was carried
out only for the sulfate. The results showed 17.1 and 18.6 percent respectively for
sulfate as determined by the 622 cm'1 band and 17.8 and 18.8 as determined by the
1158 cirri band. An analysis of this material previously provided by NAPCA showed
16.5 percent sulfate as determined by a wet chemical technique.
A brief comparison of all results of this sample is now presented.
-------
.40
cT'30
^
z°
o .25
"E
-------
1.4
1.2
Weight Percent
Figure 13. Sulfate Calibration Curve - Ion Exchange Studies
Method
NAPCA analysis
Pressed pellet
Internal standard method
Freeze drying and ion exchange
% Sulfate
16.5
26.1-19.1%
13.4-15.1%
17.1 - 18.8%, 18.1% average
As can be seen, the freeze drying and ion exchange technique appears to give results of
about the same quality as the other techniques previously studied. The average results
were about 10 percent different from that obtained by non-spectroscopic methods and the
range of results was *6 percent. However, the difficulties encountered in attempting to
analyze for hydroxide and the observed conversion of hydroxide to carbonate seem at
present to limit this technique under practical conditions to determination of sulfate and
silicate. Under far more rigorous conditions, perhaps the other species may be determinate
in a meaningful manner. The range of values (i.e. the data spread) for this method does
not seem to be greatly narrowed, although the calibration plots do appear to exhibit far
less scatter than those reported for the other pellet techniques. It may be noted that for
all of the approaches to development of a pellet method of analysis, either the scatter
in calibration (fata has been too great to yield accurate (* 10%) results, or, as in the case
of the ion exchange method, specific difficulties have been encountered in the analysis
of one or more species of interest. Also, the necessity of freeze drying of solutions and
the time required for use of exchange techniques make the above methods studied of little
use in cases where very rapid, accurate analyses for more than one component are desired.
3. Infrared Solvent System
(o) Selection of a Solvent System
One possible method for rapid analysis of reacted limestones is to
dissolve these materials in a solvent and determine the constitution of the resulting
solutions by infrared quantitative analysis. To properly do this, a satisfactory solvent,
which either is infrared inactive or can be made infrared inactive in the regions of interest
must exist. Dissolution of limestones, in past work, for wet chemical quantitative analyses
has always involved the use of aqueous solutions of various organic or inorganic acids(l).
Unfortunately, this approach is of little use in our case as the carbonate content will be
converted to CC*2- Another aqueous solvent system, based on the complexing properties
with divalent ions of EDTA (ethylenediaminetetraacetic acid) was reported to effectively
dissolve limestone without evaluation of COoCl. Unfortunately, water presents problems,
due to the high extinction coefficients of its bands, in infrared spectroscopic work.
A series of investigations were therefore aimed at finding either a
direct solvent for limestone or a material, to which EDTA could be added with dissolution
and then limestone dissolved via complexing with the EDTA. The solubility of CaSC^,
CaCC>3 and Ca(OH)2 was investigated for each of several materials. In addition, the
solubility of EDTA In these solvents was checked. The general procedure used was to
attempt to dissolve the EDTA or calcium salt with stirring in the solvent in question. After
-------
a few hours, the solutions were filtered and then freeze dried and the amounts of residues
(or dissolvable material) checked. Aside from water, DMSO, and ethylenediamine, both
limestone and EOTA were found to be insoluble in all of the liquids studied. For ethylene-
diamine and DMSO, while solubility of EDTA (ethylenediamine tetroacetic acid) was ob-
served, EDTA saturated solutions failed dissolve any measurable quantities of either calcium
sulfate, hydroxide or carbonate. Indeed, while it was possible to prepare at 20 weight
percent solution of EDTA in ethylenediamine, this solution failed to show any ability to
dissolve the calcium salts. The results of these investigations are shown in Table I. A
listing of the materials studied according to polar groups present is given in Table II. As
can be seen, materials containing all of the principal polar groups were investigated. Thus,
it appears only aqueous EDTA solutions will directly dissolve limestone. In all of the
above studies, the acid and disodium forms of EDTA were used in all cases. In the case of
ethylenediamine, both materials were observed to possess the same high solubility and
inability to complex with calcium salts in this solvent.
An investigation of the solubilities of CaCC>3 and CaSO4 in aqueous
saturated EDTA solutions was also carried out using the disodium form of EDTA. It was
found, however, that the carbonate dissolved in this solution with evolution of CC^.
Accordingly, another saturated solution of disodium EDTA was prepared and to it was
added sufficient NaOH to effect complete formation of the tetrasodium form of EDTA.
Utilizing this solvent and adding CaSC>4 in 50 mg amounts and then stirring to effect
more rapid solution, it was found that, at 25°C, 4.55 g of CaSC>4 could be easily dis-
solved in 100 ml of solvent. Similar studies, using calcium carbonate showed a solubility
of greater than 3.4 grams per 100 ml of solution.
(b) Design of Cell for Use with Aqueous EDTA
After the solvent system had been chosen, investigations of spectra of
sulfote, carbonate, and hydroxide ions in aqueous solutions were undertaken to obtain
calibration curves for limestone analysis. Initial investigations showed that water con-
taining cells as thin as .007 mm completely absorbed over most of the infrared region. As
.007 mm was the smallest commercially available spacer for use in a liquid cell, it was
decided to use on alternative approach. A drop of NaCI saturated water was placed on
one of the silver chloride cell windows and the other window then pressed directly to it,
without the use of a spacer. (The NaCI solution was used instead of distilled water be-
cause this solution was found to spread more easily over the face of the silver chloride
window.) This gave a uniform thin liquid film from which the spectra obtained indicated that
this technique would result in obtaining useful spectra. Only the region from 3500-3100 cm"'
was now completely absorbing. The rest of the spectral region was to some degree useable.
Accordingly, an aqueous solution, saturated with both NaCI and sodium sulfate was
prepared and utilized in the same fashion. Figure 14a shows the resulting spectrum. The
bands at 3300 and 1600 cm"' are due to the water. The 1100 cm"' band is due to sulfate
ion. Spectra of only the water film alone (without dissolved Na2SO4) were compared
with this spectrum and from this an approximate background (Figure 14a) was sketched in.
As can be seen, absorbances (uncorrected) of about 0.8 can be achieved for sulfate ion in
saturated solutions with thli technique. The above techniques were then employed with an
24
Solvent
Table I. Solubility of Limestone and EDTA
EDTA Limestone SolvenH-EDTA+Limestor
Acetonitrile
Acetic anhydride
Dierhylformamide
DMSO
Diethanolamine
Acetylacetone
Methonol
Acetone
Ammonium sulfate
Carbon disulfide
C$2 + diethylformamide
Forma mide
Ethylenediamine
Nitromethane
Chloroform
Carbon tetrachloride
Ethyl ene glycol
Glycerol
Proplylene carbonate
Butyl acetate
Quinoline
Triethanolamine
insoluble insoluble
" reacts to liberate CO2
" insoluble
slightly soluble
insoluble "
insoluble
reacts to liberate CC
insoluble
soluble
insoluble
very soluble
insoluble
insoluble*
insoluble
insoluble*
insol uble
*The sodium salts (i.e. Na2CO3, Na2SO4 and NoOH exhibit some solubility
in these solvents.
Table II. Classification of Solvents Used by Polar Groups
Group Solvents
-OH water, methanol, diethanolamine, triethanolomine, glycerol, ethyl ene glycol
-NH2 or formamide, dimethylformomide, diethylformamide, diethanolamine, triethanol-
-Nl*2 amine, quinoline, ethylenediamine
-C SN acetonitrile
-C = O acetylacetone, acetone, acetic anhydride
C = 5 carbon disulfide, dimethylsulfoxide (DMSO)
Esters proplyene carbonate, butylacetate
Misc. nitromethane, chloroform, carbon tetrachloride
-------
NojCO^ saturated solution containing a small amount of added Nc^SO^j (Figure 14b).
As can be seen, bands due to both carbonate and sulfate are readily seen, the carbonate
absorbing here at ca. 1400 cm"'.
A solution, saturated with the tetrasodium form of EDTA was prepared
and saturated with calcium sulfate. The spectrum of a thin film of this solution (Figure 14c)
clearly shows a useable sulfate band at ca. 1100 cm'' in addition to a series of bands
due to dissolved EDTA. Indeed, from the spectrum present (Figure 14c), sulfate absor-
bances of as large as 0.8 can be achieved with this simple technique demonstrating the
feasibility of this approach.
Estimates of the film thicknesses for which meaningful spectra were
obtained were made from the volume of one drop of water from the dispensing eyedropper
used. From the cell's known width and length, a value of .004 mm was calculated as
the maximum film thickness assuming none of the water escaped via flowing over one of the
salt plate edges. Probably, however, some solution did escape, although it is fairly
certain that at least 50% of that added was utilized as the film. Thus, a thickness of
ca. .003 mm appeared to be optimum for our purposes. The problem, however was to
obtain this thickness each time. As spacers were not available in such small thicknesses,
a special cell had to be designed.
(c) Sulfote Analysis Using the Solvent Technique
(1) General Procedure
Quantitative determinations for sulfate content of reacted
limestones have been conducted by the infrared method using cells such as those shewn
in Figure 15. The cells consisted of two optically flat silver chloride plates. On the
outer portions of the bottom plate was deposited a 0.003 mm thick silver film which
served as a cell spacer. The other plate had two small holes drilled through it for ad-
mission of sample. This assembly was of the proper size that it could be used in con-
junction with commercial cell holders. For this analytical study, a pair of such cells,
equal in thickness to within ±2 percent, were used; one containing a saturated aqueous
solution of tetrasodium EDTA which was placed in the spectrometer reference beam and
the other containing a solution of calcium sulfate or dissolved limestone in this solvent
system which was used in the sample beam. All spectra were run on a Beckman IR10
infrared spectrometer.
It may be noted that this approach is clearly superior to earlier
methods of working with aqueous solutions. Thin film cells used in earlier studies
employed thin polymeric films deposited onto the face of a salt plate from solution (i.e.
cellulose acetate from acetone solutlon)(3,6). fl,js approach was satisfactory for pre-
paring single cells, but is obviously of little use in constructing matched pairs of cells.
Past studies by infrared techniques wing aqueous solutions were largely limited to
qualitative studies^. 10)f however, the recent advances mentioned immediately above
have made possible some quantitative work, some of which was only recently reported(°).
26
a. Spectrum of Sulfate Saturated Water
! I .1 I I
b. Spectrum of Solution Containing Carbonate and Sulfate
c. Spectrum of EDTA
Saturated Solution
figure 14. Spectra of Carbonate and Sulfate in AquiOM Film
-------
.003 mm Film
of Silver De-
posited onto
Outer Portions
Only of
j! ii
i n
1 I I |
f////A Y////^
TopAgCI
Plate
Bottom AgCI
Plate
Bottom AgCI Plate
Figure 15. Infrared Cell for Use with Limestone Solvent
28
The saturated tetrasodium EDTA solutions used were prepared by
dissolving this material in distilled water to saturation and storing the resulting solutions
over an excess of the EDTA salt. These solutions contained ~290 g/liter EDTA.
Solutions of either calcium sulfate or limestone were prepared by
dissolving a known amount of the material of interest in a specified volume of solvent. In
the case of some of the limestone samples, difficulties were encountered in dissolving all
of the materials. In some cases, an insoluble residue was found to remain. This is in
accord with the results of Hill and Goebel(') who found that limestone constituents such
as silica could not be dissolved by EDTA. (These residues were recoveredbyfiltration,
washed free of EDTA and incorporated into KBr pellets (2% by weight residue) for analysis.)
Spectra of residues obtained in our studies revealed them to contain no'sulfate but large
amounts of silicates. The results for all of the samples studied are shown in Table III. Cali-
bration curves for sulfate absorbance were obtained from the data on a number of sulfate
solutions using the dual cell method described above.
In the data analysis, the absorbances at the sulfate band maxi-
mum (1110 cm"') were measured assuming the background over the sulfate band region was
linear. Studies using the dual cells, with only saturated EDTA present in each cell have
shown this assumption to be valid. For the calibration studies, a plot was then made of
absorbance versus concentration for the CaSO^j samples studied. This is shown in Figure 16.
As the same cell thicknesses were employed each time, no corrections had to be made to
convert the absorbance data to absorbance per unit thickness.
The analytical results (Table IV), are expressed as percent SOj.
These numbers were arrived at assuming all of the sulfate in the reacted limestone samples
was present as CaSO*. Such an assumption agrees well with the findings of BorgwardtU 1) <
in studies of the reactions of calcined limestones with flue gases.
Wet chemical analysis of the limestones studied were conducted
by NAPCA to corroborate the infrared results. Specifically, the limestone samples were
ground to a fine powder after drying in an oven for one hour to remove physically absorbed
water. An accurately weighed sample of 0.2-0.3 g was heated in a mixture of distilled
water and cation exchange resin for one hour. The ion-exchange resin(12) serves two
purposes in this procedure. The first is that the equilibrium of CaSO^ is shifted toward
the right by continually removing calcium ions from the solution and replacing them with
hydrogen ions. This releases the sulfate to solution as h^SO^. Other slightly soluble
sulfates are made soluble in the same manner. Simultaneously the resin renders the
solution virtually free of cations which might otherwise interfere with the final Jitrhietric
determination. The solution was filtered through a wad of glass wool into a volumetric
flask with washings from the glass wool. A suitable aliquot was taken, made to 80% with
isopropyl alcohol and titrated with 0.005 NBa (ClO^ usi"9 Thorin indieatoK'3). As
can be seen, good agreement exists for all samples. Duplicate determinations for each of
the samples further show that method reproducibility is quite satisfactory.
Some exploratory work was also conducted for analysis of the
carbonate and dissolvable silicate contents of limestone using the same techniques. However,
-------
Table III. Spectra of Insoluble Limestone Residues
Sample
Number
187
97
199
332
330
Band Position
(cm-1)
3650
1400
1240
1035
970
880
480-570
1400
1025
1250
970
880
490
1420
1250
1080
1025
935
800
475
1420
1080
1000
880
480
1440
1240
1030
950
880
480
Strength
very weak
strong
weak
strong
strong
medium
very weak
very weak
strong
strong
strong
strong
strong
very weak
weak
weak
weak
weak
weak
very weak
strong
weak
medium
medium
Assignment
OH-
COf
?
Si-O
Si-0
Si-O
Si-0
C0,=
Si-0
?
Si-O
Si-O
Si-O
Si-O
Si-O
Si-O
Si-O
Si-O
C03=
Si-O
Si-O
Si-O
Si-O
C03=
?
Si-0
Si-O
Si-O
Si-O
377
337
no residue - sample entirely dissolved
no residue - sample entirely dissolved
30
Sample
Number
1351
Table III. (cont.) Spectra of Insoluble Limestone Residues
Strength
Band Position
(cm"')
1360
Assignment
very weak
very weak
strong
weak
strong
weak
weak
weak
very weak
strong
strong
strong
weak
weak
strong
C03°
?
Si-O
Si-O
Si-O
Si-0
?
Si-0
C03=
Si-O
Si-O
Si-0
Si-O
?
Si-O
1699
1410
1255
995
955
900
840
770
500
1420
1050
940
870
830
765
500
no residue - sample entirely dissolved
-------
Table IV. Analyses of Limestones
Sample
Number
330
199
337
97
377
187
332
BCR 1699
(Solution I*)
BCR 1699
(Solution II*)
BCR 1351
BCR 1360
Wt%SO3
By Infrared
47.26
44.50
28.08
27.61
47.12
46.29
27.88
29.20
11.56
11.50
30.45
29.86
18.37
17.61
17.39
17.63
17.40
17.11
18.45
40.30
41.77
37.76
40.12
Average
45.88
27.84
46.70
28.54
11.53
30.15
17.79
17.52
17.78
4K03
38.94
Average
Wet Chemical
Value
44.5
27.1
42.8
30.4
10.5
28.4
16.5
16.5
16.5
39.7
37.5
Ratio of Infrared to
Wet Chemical Values
1.03
1.03
1.09
.95
1.10
1.06
1.08
1.06
1.08
1.035
1.04
0.4 0.8 1.2 1.6
Concentration (g CaSO^/IOOcc)
2.0
Figure 16. Calibration Curve with .003 mm Cell for Sulfate Ion at 1110 cm'1
•Solutions I and II of sample BCR 1699 were different concentrations of the same limestone
material in saturated fetrasodium EDTA. The second solution, which exhibited the higher
pH, was the more concentrated.
32
-------
the calibration curve obtained for carbonate (Figure 17) exhibited considerable scatter
and carbonates were not further examined.
(2) Sulfote Analyses - Results
Calibrations for sulfate were made using solutions of calcium sulfate
dissolved in the saturated aqueous tetrasodium EOTA solvent. In these experiments, the
test solutions were placed in the cell in the spectrometer sample beam. In the reference
beam was placed a duplicate cell containing the saturated EDTA solution, which acted to
virtually eliminate most of the spectrum due re-water and EDTA. This technique was
completely successful for the sulfate calibrations as can be seen in Figure 16, which shows
little scatter and good reproducibilify using the 1110 cm"' sulfate band.
These techniques were then used to analyze a number of samples
of reacted limestone for sulfate content. In all but the two cases of sample 1699, between
1.4 and 1.8 grams of limestone were dissolved in 100 cc of the EDTA solution to gain con-
centrations in same range as calibration points. For sample 1699, concentrations of
~0.25 and 0.4 g/100 cc were used. Figure 18 shows spectra in the 1750-800 cm"' region
of both dissolved limestone and of the EDTA solvent above using a single thin film cell.
As can be seen, differences due to the limestone exists at ~1440 and 1110 cm"'. Spectra
obtained using a set of cells as described above, show essentially only these bands. Table
IV shows the analysis results along with results obtained by the National Air Pollution
Control Administration's (NAPCA) Cincinnati laboratories for these same samples using
wet chemical.analysis methods. As can be seen, the values obtained by the infrared method
are somewhat high. A discussion of the reasons for this is given below:
(3) EDTA-pH Studies
The pH's of the solutions of limestones and of the pure saturated
tetrasodium EDTA aqueous solution were determined with Grammercy Universal Indicator
using the color charts provided by the manufacturer. Over the pH range 8.5-10.0, this
indicator changes slowly from green to blue to violet to reddish violet. The exact pro-
cedure employed in these investigations for each solution was to add 1 mil of the solution
in question to 5 mil of distilled water and then add a few drops of indicator, shake the
solution to insure proper mixing and observe the resulting color. The results for each of
the solutions used are given in Table V.
A graph of the data given in Table V appears in Figure 19. As
can be seen, there is a direct correlation between deviation in pH, from that of the pure
EDTA solution and percent error in the analytical results. Thus, solutions which had pH's
lower than that of the reference solution gave consistently low analysis results, while
solutions of increased alkalinity gave consistently high values.
The above findings can be readily interpreted by examination of
the spectra of the various forms of EDTA. The spectrum of the disodium salt (Figure 20) shows
a sharp band at ca. 1180 cm"'. Single cell spectra of the tetrasodium Form of EDTA, both
in the presence andabsence of limestone are shown in Figure 18. Comparison of the two
34
0.24
0.20
0.16
0.12 -
0.08
0.04 -
0.8 1.6 2.4 3.2
Concentration (g/CaCOj/lOOcc)
4.0
Figure 17. Calibration Curve with .003 mm Cell for Carbonate ten at 1400 cm'
,-1
-------
1750
1600
1400 1200
Wavenumber cm"'
1000
Figure 18. Spectra of EDTA Solution Alone and EDTA Solution
Containing Dissolved Reacted Limestone
-0
-0.1
-0.2
-0.3
-0.4
-0.5
-0.6
M-7
^0.8
r0.9
PI.O
P2:5o
'lot)
Table V. Effects of Solution pH on Analysis Errors
Solution
Sample
Number
EDTA
330
199
187
97
337
377
332
1351
1360
1699
CD
1699
('2)
-El
9.0
9.0
9.0
9.2
8.8
9.3
9.5
9.6
9.3
9.4
9.3
9.5
% Error In Analysis
3% high
3% high
6% high
5% low
9% high
10% high
8% high
6% high
8% high
3.5% high
4% high
-------
8.6
Figure 19. Percent Error vs. Solution pH for Llmatene Analyses
38
SfSSSSSSSS
^
U
I
£
S
UJ
•s
1
5
J
•s
8.
00
-------
spectra shows this band has shifted to ca. 1135 cm"' and broadened. The ph's in solution'
of the two materials are 4.Sand 9.0 for saturated solutions of the two EDTA forms. By
making the tetrasodium EDTA solutions more alkaline, perhaps, further shifts in he 1135 cm~l
band could occur for the following reasons:
a). The carbon nitrogen stretches for — C — NH2 groups occur
in the region 1200-1000 crrrl, the exact position depending on other groups banded to the
carbon and on the degree of profanation of the NH2 group in the medium in which it is
being examined.
b). For EDTA salts, the structure is:
CH2COO- CH2COO"
>N — CH2 — CH2 — N<
CHjCOO"
CHjCOCT
c). The profanation equilibrium
-NR2
NR2H + OH"
is pH dependent. Indeed, making the solution more alkaline should shift the equilibrium
to the left and thus bring the C— N absorption band closer to the 1110 cm"' region being
examined.
Figure 21 shows a plot of the change in solution pH on limestone
dissolution versus %SO3 (or sulfate content) of the sample. The line drawn is that to
best fit the data. As can be seen, the samples containing the least sulfate (and hence the
largest amount of unreacted oxide and hydroxide) clearly cause the greatest changes in pH
toward alkalinity.
Figure 22 shows a plot of percent $03 variation* between
duplicate runs on the same samples as a function of the order in which the samples were
run. As can be seen, with increasing cell use (rising run number), the data spreads on
individual samples rises. This can be attributed to factors such as cell erosion and general
wear ana/or oxidation of the silver spacers and silver chloride windows. With the cell
arrangement, it is impossible to repolish one of the windows due to the silver film present.
Window repolishing would necessitate evaporation of a new spacer film making such an
approach unrealistic. For future work, an improved cell should either have more wear
•By percent SOj is meant the difference between two duplicate runs on the same sample.
This, for example, if answers of 18 and 20 percent were found, the percent SOg vari-
ation plotted would be 20-18 or 2.0.
40
40
JD
I 30
-------
3.0
Figure 22. Percent $©3 Variation as a
Function of Cell Age
42
resistant material! than silver or AgCI or use a demountable 3 micron spacer so windows
can readily be repolished,
(a) Analysis for Other Limestone Constituent!
Using the same methods as for sulfare analysis, a calibration curve,
as shown earlier was constructed for carbonate using the 1400 cm"' CO$* bond. Unfor-
tunately, the data showed considerable scatter, particularly at lower concentrations of
carbonate. From spectra obtained for the dissolved EOTA, it was ascertained that the
cause of this was small differences in cell thicknesses. EDTA displays a number of intense
bands in the 1400 cm"' region, but not near 1110 cm"'. Thus, analysis for carbonate
can only be conducted, at present using such a procedure, with very concentrated solutions.
Spectra run by the techniques above of sodium silicate dissolved in the
saturated EDTA solutions revealed bands of useable Intensity at ca. 1100, 1020 and 850
cm"'. Thus, dissolved silicates can also be determined by the above techniques.
Spectra run of pure 02O revealed absorption only in the 2600-2000 cnTl
and 1250-1150 air1 regions. Addition of either LiOH, NaOH, or KOH to D2O yield
bands at ca. 3400 cm-' and 1650 cm-' as is expected for OH groups tied to similar species
by hydrogen bonding, as is the case in the liquid phaseO*). In a series of experiments
the percentage of NaOH in DjO was increased over a series of runs. The 3400 and
1650 cm'l bands were both observed to increase with increasing hydroxide concentrations.
These experiments, using DjO as solvent, were performed using a single .003 mm cell
with a wire screen in the reference beam. In another series of studies, spectra were
run of solutions of Ca(OH)2 in EDTA saturated DjO. In these experiments, a thin (.003 mm)
cell filled with EDTA saturated DoO was used in the reference beam. Spectra obtained
revealed strong bread bands at 3400 and 1660 cm"'.
Using one of the thin film cells, spectra were also recorded of solutions
of LiNO3, NaNO2, Na2HPO4 and Na2$O3 to demonstrate that other ions, such as
nitrate and phosphate could also be quantitatively determined in solution by the infrared
method. For these studies, no cell or screen was used in the reference beam. The spectra
for the above salts were obtained in both HjO and t^O. In both solvents, the following
band positions were noted:
NO3-
-
1370 cm"' (S)
1240 cm-1 (S)
1075 cm-1 (S)
920 cm"' (S)
(e) Scale Expansion Technique Studies
830 cm"' (W)
1350 cm-1 (Wsh)
985 cm-1 (M)
Using the scale expansion accessory and an external chart recorder,
efforts were made in this period to determine the minimum amount of dissolved carbonate,
in the absence of EDTA, which would be detectable with the .thin film cell method. For
these studies, a duplicate cell, filled with water, was used in the reference beam. All
-------
of these spectra were run at low scanning speed and fairly low gain to minimize nolle
problems, and the 10X scale expansion setting was used. The results are given below:
Solution
Concentration
O.lOg/lOOcc
0.01 g/100 cc
0.001 g/100 cc
10X Expanded
Absorbance
0.05
0.007
not detectable
Unexpanded
Absorbance
0.005
0.0007
Unexpanded
Noise Level
0.0003
0.0003
0.0003
As can be seen, the minimum detectable amount of dissolved carbonate
observable with our apparatus appears to be about 10~2 g/100 cc. From other calibration
data, a calculated value for the minimum observable concentration can be calculated
assuming an inherent noise level of *0.0003 absorbance. Here, the minimum useful
absorbance observable would be —0.0006, twice the inherent noise level. From the
slope of the carbonate calibration curve reported earlier, an absorbance of 6 x 10"'* would
correspond to a carbonate concentration of ca. 9 x 10~3 g/100 cc, in good agreement with
the experimental value reported above.
As the extinction coefficient for sulfate, as judged by the calibration
curves obtained earlier, is of the same order of magnitude as that for carbonate, the
minimal detectable concentration of sulfate should also be ca. 10-2 g/IOO cc.
(f) Discussion of Results with Thin Film Cells
The methods developed in this study using very thin cells appear to give
satisfactory results for sulfate analysis and probably could also be directly utilized for
silicate and hydroxide determinations. The problem of carbonate analysis, however,
remains to be solved. At the present state of development, our techniques would require
application of additional methods, such as are due to Robinson(15), to sort out contributions of
EDTA absorption to the scatter in the carbonate studies. For the other species, however,
fortunately, such an elaborate procedure should not prove to be necessary.
The studies using D2O have revealed this to be a superior solvent for
analytical purposes to water. Also, it may be possible in D2O based solvents to separately
determine the hydroxide and unreacted oxides of limestone. Thus, if one were to use
NajDj EDTA and Na2H2 EDTA in D2O as solvents, for CaO and Ca(OH)2, the reaction
occurring would be:
1. Na2D EDTA + CaO
2. Na2D2 EDTA + Ca(OH)2-
3. Na2H2 EDTA + CaO
4. H2O + D2O (excess)
5. No2H2 EDTA + Ca(OH)2-
D2O + CaNa2 EDTA
2 DOH + CaNa2 EDTA
H2O + CaNa2 EDTA
2DOH
H2O + CaNa2 EDTA
From the above reactions, one can see that using the deuterated EDTA In
one would measure only the hydroxide present, while using the dihydrogen EDTA, one
could determine oxide plus hydroxide. Thus, from the two measurements one could obtain
the concentrations of oxide and hydroxide present in reacted limestones.
To extend the techniques developed for sulfate analysis further will
require construction of improved cells. As can be seen from Figure 22, the silver chloride
cells possess ugly problems which tend to limit their usefulness. More resistant materials,
clearly, should be used. In future studies, materials such as ZnS or ZnSe could be employed
and a better spacer material, such as Teflon or tantalum could likewise be utilized to improve
cell lifetimes. Also, better ways of filling of the cells should be developed. The standard
filling arrangement for commercial cell holders works well if the sample cavity to be filled
is sufficiently thbk; however, for our cases to fill the cells and have them free from air is
a tedious procedure.
A summary, from all the data presented above, it can be seen that a
rapid accurate, infrared method for sulfate analysis has been developed and that this
general method should also be applicable to other uses of interest. All that will have to be
done to achieve this will be to obtain the necessary calibration curves under controlled
pH conditions.
B. DEADBURNING AND MgO HYDRATION STUDIES
1. Introduction
The dry limestone injection process for sulfur dioxide removal from flue gases
offers a number of advantages based on its low cost and simplicity of operations. However,
in certain cases a number of problems have been formed which are related to incomplete
reaction of the limestone. One of the possible causes of this incomplete reaction is the
phenomenon called deadburning, wherein the calcined limestone becomes relatively
inactive toward either hydration or chemical reaction due to either collection of im-
purities at its surface or to a large loss of surface area and pososity.
A number of studies have been made of the deadburning phenomenon; however,
much disagreement still exists as to the importance of several of the variables involved.
Thus, while some authors have related reactivity to pore volumes and surface areas available
for reaction(16~'8), others have stressed the possible influence of various solid state reactions
which can occur with impurities during calcination. Impurities during calcination. Im-
purities such as SiOo, AI2O3, Fe2O3, phosphates, sulfates and potassium and sodium
salts are known to differ widely with sources(I9), and it is known that above 17509C
reactions between CaO and SiO2 do occur which result in formation of surface silicates.
Indeed, it has been found that limes high in silica may react with water like hard burned
(i.e. low surface area) materials even though extensive sintering did not occur(20).
An extensive study of the pyrochemical reactions occurring in impure limestone
has been conducted by Lee(2l). With increasing calcination temperature several reactions
occur including formation of silicates, ferrites and alumina ferrites, all of which lead to
inert materials which, at high temperature, melt and diffuse readily to the surface causing
particle shrinkage and reduction in porosity. On cooling, it has been found that these
-------
-• s F
V-3-- o
-------
Table VII. Spectra of Limestones as a Function of Calcination Temperature
1700°F
101 Series
2000°F
2300°F
2600°F
3655s
2960 w
2350 w
ISOOw
1440s
1260w
IVZOw
980s
915 w
875 w
850 w
3650s
2920 w
2350 w
1440*
1260 w
1115s
993s
890s
875 w
845 w
3650s
2940 w
2350 w
1440s
1260 w
1120 vw
lOOOc
910 w
875 w
845 vw
% = strong
w = weak
3655s
2955 vw
2340 w
_
1440s
1250w
IllOw
995m
920 w
875 w
850 w
3650s
2930 w
2350 w
1440s
1260 w
1115s
995s
910 w
875 w
845 w
3650s
2940 w
2350 w
1440s
1260 w
1110 vw
1005m
925w
875 vw
860 vw
b
m
3655s
2960 w
2350 w
_
1440s
1235w
IllOw
995 w
940 w
875 w
850 w
3650s
2930 w
2350 w
1440s
1260w
1105m
1005mb
920 vw
875 vw
845 m
3650s
2940 w
2350 w
1440s
1260 w
HlOvw
lOlOw
925 w
875 vw
855 w
= broad
= medium
3650s
2960 w
2350 w
1440s
1235w
llOOw
1000 w
940 w
875 w
855m
102 Series
3650s
2930 w
2350 w
1440s
1260w
I105w
1010 wb
920 vw
875 vw
855m
100 Series
3650s
2940 w
2350 w
1440s
1260w
1110 vw
1010 vw
925 w
„
865 w
vw = very weak
3200°F
3650s
2960 w
2350 w
1440s
1230 w
1095w
1005 w
950 w
875 w
860m
3650s
2930 w
2350 w
1440s
1260 w
1010 wb
920 vw
875 vw
865m
3650s
2940 w
2350 w
1440s
1260w.
1105 vw
1010 vw
925 w
870m
Assignment
OH"
C03=
C03=
CO3=
C03=
?
Si-O
Si-O
Si-O
C03=
Si-O
OH"
C03=
CO3=
C03=
?
Si-O
Si-O
Si-O
C03=
Si-O
OH"
C03=
C03=
C03=
?
Si-O
Si-O
Si-O
C03=
Si-O
48
(2)
(3)
the 1115 cm"' band decreases in intensity per unit
weight of sample with increasing calcination temperature
and shifts to lower frequencies.
the 990 cm-'band shifts to ca. 1010"' with in-
creasing calcination temperature and decreases
strongly in intensity.
(4) the 890 cm-' band shifts to ca. 920 cm"' and
decreases in intensity with increasing calcination
temperature.
(5) the 845 cm"' band shifts to 865 cm"' and increases
in intensity with rising calcination temperature.
(6) Generally, the four above bands broaden considerably
with increasing calcination temperature.
Figure 23 shows a plot of the fraction of total shift observed as a
function of calcination temperatures for the averages of the three bands of principal
interest (i.e. 995, 890 and 845 cm"'). Here, the average band positions and shifts for
each temperature for all samples were calculated and the average band shift is plotted
against calcination temperature. As can be seen, the band at ca. 930 and 1010 cm"'
seem to exhibit the same dependence on calcination temperature, while the 860 cm"'
band exhibits a different dependence.
The cause of these dependencies can be traced to either of two effects,
one of which is the influence of particle size and the other of which is due to changes in
chemical composition. The following arguments can be involved in favor of the in-
fluence of composition.
a).
, _. ...anci noquencies
i corresponding metal sil icotes. Indeed, below are
listed the positions for the corresponding bands for sodium
orthsilicate and metasilicate, and the frequency differences
the corresponding bands.
Sodium Orthosilicate
(cm"')
1080
990
920
840
755
Sodi
ium Metasilicate
(cm"')
1020
950
870
795
700
Difference
(cm"')
60
40
50
45
55
-------
O 8o;0 cm'*
D 930cm-'
A 1007 cnf1
2300 2600 2900
Calcination Temperature °F
3200
Figure 23. Percentage of Total Band Shifts Observed as a
Function of Calcination Temperature
50
Thus, there Is an average shift of ca. 50 cm"' in going
from Na^S 104. The spectra of the two sodium silicates
were run by the KBr pellet technique.
b). There is considerable evidenced) that the reactions:
CaO + SiC>2
CoSiOj + CaO-
CaSiO,
do occur in limestones at the calcination temperature
used in our work. Thus, our average shifts of ca 21 cm"'
could be readily explained by a partial reaction as the
meta and orthosilicate bands, at least for the sodium salts,
do overlap and the maximum positions would be expected
to shift toward the. orthosilicate positions as the reaction
proceeded.
Further confirmation of the presence of meta and orthosil icates would
require the use of other techniques such as X-ray diffraction. Such studies should, for
ease of investigation, be conducted on samples of high silica content. It has not been
possible to show, in this infrared study alone, where, with respect to the surface, the
silicates are located. Indeed, our spectra were obtained by transmission and show only
total silicates present. Evidence for or against preferential diffusion of silicates to the
surface, as has been claimed(21) must be obtained by other techniques such as the use
of Hie electron microprobe or low energy electron diffraction.
(b) Electron Microprobe Studies
Line analysis microprobe determinations for silicon content were run
on a number of particles of the 1700°F and 3200°F calcined samples of the three series
of samples provided by NAPCA. In addition, a small number of analyses were run for
iron an the Series 102 samples, however insufficient data was collected for these .
runs to justify drawing any fair conclusions.
The experimental procedure consisted of making disk type samples in
degassed epoxy resin; allowing the resin to cure and polishing of the disk until flat crass
sections of a number of limestone particles were observed. Care was taken in these steps
to avoid contamination of the surfaces with either silica or iron compounds. The disks
were then taken and inserted into the microprobe apparatus and I ine scam across individual
particle cross sections for silicon were run on a number of individual particles. A few similar
runs were also made for iron.
Line scans across a number of individual particles of the 1700°F calcines
of all three types of limestone for silicon revealed that almost all of the silica was confined
to discrete regions of the particles. In an optical microscope these regions appeared to
be darker and to constitute discrete phases. A typical line analysis run is shown in Figure 24
-------
UOJ40JJU93UO-) U03I|JS
52
fora 1700°F calcined particle of the Series 100 limestone. The trace was run across the
particle at 160 u/min. As can be seen, for the particle, most of the silicon was located
in two small regions.
Similar scans were made on 3200°F calcines. A typical run is shown
in Figure 25 fora Series 100 particle. As can be seen, while discrete regions of high
silicon content are still present, most of the other material (or phases) present now also
contain some silicon (i.e. silica has diffused throughout the limestone). This scan was
made at 160 u/min under the exact same conditions used for the 1700°F samples.
All of the above work can be taken as evidence for diffusion of silicates
and possibly also for chemical reaction. Clearly, what is needed is a detailed analysis
for cajcium, oxygen and silicon for each of the phases (regions) present at different
calcination temperatures to gain complete proof of the occurrence of chemical reactions.
No enough particles were scanned during this brief study to gain much evidence for
surface buildup of silicates. What is clearly needed in future work are microprobe scans
for a number of elements run on a number of particles for different temperature calcines
of each material. Only from such a collection of information can conclusive evidence
for buildup of surface silicate layers be obtained. Also, the scanning rate in the above
experiments was 160 u/min. If the silicate layers are quite thin (i.e. ~l-2u), then a
much slower scanning speed will have to be used to clearly detect their presence and
composition.
(c) Studies of Hydration of Calcined Limestone
In a series of experiments pure magnesium and calcium oxides (Fisher
Analytical Reagent grade) were exposed at 23°C to about 25 mm of water vapor in a
vacuum system for a few hours. The water vapor was then removed by evacuation and
the materials were then mixed with KBr. Pellets were then made and spectra were run.
The results showed a sharp OH band at 3712 cm~l for the hydrated MgO and a sharp
peak at 3655 cm"' for the calcium samples. The presence of these sharp OH stretching
bands confirms the hydration of both CaO and MgO. For perfectly anhydrous samples,
such bands were not observed. These experiments were repeated several times with the
same results.
A number of uncalcined, calcined and reacted materials, provided by
NAPCA, were also examined by the KBr technique. The behavior of the OH bands of
many of these samples, both calcined and reacted, is presented in Table VIII. Those
listed include all samples containing more than 5 percent MgO. For all other samples of
this series examined, in all cases where less that 5 percent magnesium oxide was present,
no band at 3712 cm"' could be detected. As can be seen from the table, in all cases the
3712 cm"' band appears to be removed by exposure to simulated flue gas at elevated
temperature to a far greater amount than the 3650 cm*' band is affected. Indeed, in
many cases, while two bands were observed for the calcined samples, only one band, at
3650 cm*', could be detected for the reacted materials. A typical case is shown in
Figure 26, where the top spectrum, of a calcined but unreacted material, containing 28
percent MgO, is shown to exhibit two strcng bands at 3650 and 3710 cm'. The bottom
-------
I I
i
8
UJ
8
s
UOI4OMU33UO3 UC4I|IS
54
Table VIII. Hydratlon of MgO and Evidence for Preferential
Reaction Between MgO and Sulfur Dioxide
Sample
Number
200
BCR 1697
201
BCR 1351
202
BCR 1352
203
BCR 1684
204
BCR 1342
205
BCR 1361
206
BCR 1360
%MgO
43
28.5
31
39
28
7.2
13
Calcined
3712 cm"' strong
3650 cm"' strong
3712cm"' medium
3650 cm"' strong
3712 cm-1 medium
3650 cm-1 strong
3712 em"' strong
3650 cm'' strong
3712 em-' strong
3650 cm*' strong
3712 cm-' very weak
3650 cm-' strong
3712 cm"' weak
3650 cm"' strong
Reacted
3712 em'' very weak
3650 cm"' strong
3712 em"' absent
3650 em-' strong
3712 cm*' absent
3650 cm"' strong
3712 cm"' very weak
3650 cm"' strong
3712 cm"' absent
3650 cm"' strong
3712 cm"' absent
3650 cm'' strong
3712 em'' absent
-------
» « * s 8 « «M! 38
jj. M-'H*-l---MI'"M".i-|:--'M'hrl
' - * s « » i| * 5»?:
56
spectra of this tame material after reaction with flue gas, can be seen to exhibit only
one OH band at 3650 ctn-1. All of this may be taken as proof of the hydration of
magnesium oxide In slaked dolomitic samples and of either a preferential reaction between
MgO and some constituents of flue gases or a decomposition of magnesium hydroxide
under reaction conditions.
C. HIGH TEMPERATURE KINETIC STUDIES
1. General Introduction
Work on an earlier program has shown that at fairly low temperature (i.e.
~250 Q, the rate determining step in the reaction between calcined limestone, oxygen
and sulfur dioxide is the oxidation of absorbed SO2 (possibly as sulfite) to sulfate(4).
Unfortunately, in these investigations, the individual steps in the overall process were
not separately investigated and complete in situ reaction studies were not conducted using
the infrared method. As a result, the conclusions drawn were somewhat open to question
because the overall mechanism operating was not proven. The high temperature studies
undertaken under this contract were performed to remedy the above information gaps and
to obtain useful kinetic data at temperatures closer to those used under operating process
conditions.
2. Experimental
For all of these investigations, a set of high temperature optical cells were
used. The sample cell, as is shown in Figure 27 contained NRC foil flanges capable of
withstanding repeated heating and cooling cycles from cryogenic temperatures to over
500°C. The cell body was made of stainless steel and aluminum gaskets were used to
achieve an airtight seal between the detachable parts. The exit and entrance gas ports
contained Kovar seals for attachment to glass apparatus.
As it proved impossible to obtain a sealant with which to attach the Irtran 2
windows to the cell, another approach was used. The windows were of a sire that they
fitted reasonably tight into their position. The insertion of tight fitting metal large "O"
ring type spacers behind these windows were found to hold them quite firmly in place.
While such o system was not vacuum tight, outgassing of the samples plated onto the
windows could still be accomplished in a stream of flowing inert gas at elevated temper-
atures. A sketch of this part of the system is shown in Figure 28. Also, as the windows
were tight fitting and the exit gas port from cell was fairly large in cross-sectional area,
gas losses around the windows were expected to total no more than one or two percent
(from cross-sectional area considerations) of the total gas flow. As the sample was exposed
to a constant flowing stream of gas, contamination back diffusion through small window
leaks was shown to be negligible.
For the reference beam cells, a high temperature single cellW, minus its
Viton "O" rings, was used. In this state this cell could be maintained at the same
temperature as the working cell, but could not be evacuated. This cell contained an
-------
<3
I
£
5
&
3
I
E
58
d
s.xx w\\\\\ S S V \ ^S\l
\, IRTRAN Window \.
Long Mete I
Cylindrical Spacer
(tight Rtting)
Figure 28. Diagram of System Employed to Hold Windows in Place
-------
identical set of Irtran 2 windows and for all experiments was maintained at the same
temperature as the sample cell to minimize radiation problems. No gases were flowed
through the reference beam cell.
For the high temperature studies, samples were deposited on the Irtran windows
by spraying a suspension of limestone in acetone onto the windows and allowing the'
acetone to evaporate. Prior to conducting all of the runs, the samples were outgassed
at the run temperature in nitrogen at 400°C. The limestone used in all of the studies was the
1700°F calcined Series 100 material provided by NAPCA. The total sample weight
on both windows was kept at ca. 50 mg for all experiments. Some problems initially
encountered in these studies with variation in flow rates were solved by placing a flowmeter
at the exit of the cell. Temperatures were maintained by the insulated resistance wire heater
around the cells. Temperatures, which were monitored by thermocouples attached to the
windows, could be readily maintained at *ca. 5°Cforall experiments. All gases were fed
into the cells from a gas mixing system described in an earl ier study(^) . Each of the gases
was fed from its tank source, through a calibrated flow meter, into the flow lines, the N2
was presatuiated by two water presatuiators in series at 22 C.
In all experiments, the window; plus samples were carefully weighed before and
after reactions with SO2 to gain the weight increases (and hence amount of product formed).
This information was used to compute kinetic data presented later in this report. The surface
area of the Series 100 limestone used was 3. 1
In these studies, 475°C was the highest temperature investigated. Efforts
were made to gather information at higher temperatures due to problems involving window
oxidation and degradation under such conditions. It was found that at 500 C, a small
amount of window oxidation occurred. For more data to be obtained at such elevated
temperatures, a window material more resistant than Irtran 2 would have to be used.
3. Results
Preliminary scans run on samples during reaction between limestone (Fredonia
1700°F) and a simulated flue gas consisting of 0.3% $03, 2% Oj, 2% H2O, and the
balance nitrogen revealed formation of both sulfite and sulfate at 400°C. The sulfite
and sulfate band maximum were seen at 920 and 1 1 10 cm"', respectively. No bands due
to sulfate or sulfite formation could be detected by passing the flue gas through the
empty cell past the Irtran 2 windows.
During these studies, it was found that the band intensities increased markedly
in cooling of the sample to room temperature. Accordingly, an experiment was perforrnad
with a -jeered sample, wherein the sample beam runs blocked off and the emission spectra
of the 400°C sample wai recorded. The spectra clearly showed emission bands at 1110
and 920 cm"' due to sulfat* and sulfite. As the intensity of the emission bands were
expected to increase with rising temperature, the absorbances of samples should decrease with
increasing temperature for our present arrangement. This did not prove to be a problem at the
temperatures of our study, but probably would cause severe problems at more elevated
temperatures.
60
Using the apparatus (i.e. the dual cell with the reference cell consisting of
two Irtran 2 windows maintained at the same temperature as the sample cell), kinetic
experiments were performed in the 380-475°C range. In one study, a mixture of 0.3%
SO2, 2% water vapor, balance nitrogen was passed over the samples at 800 cc/mln and
the 920 cm"' absorbance was monitored with time using the external recorder. Gas flow
rates were monitored at room temperature. The absorbances were found to increase with
time in a semi-logarithmic manner in all cases, leveling off to a fixed value after finite
times. A kinetic plot of one of these experiments is presented in Figure 29. The vertical
scale,
1 - absorbance - initial absorbance
total absorbance change
can be shown by simple algebra to be equivalent to
1 - const, cone. SOg^f
const, cone. SO;j= final
where cone. SO3~j is the concentration of $03" at time T and cone. SO3~ final is the
final concentration of SC^", which was the concentration of sulfate present when no
further absorbance changes occurred at 920 cm .
Now, KSC>3~ final is proportional to the total number of reactive sites
(i.e. to the total amount of reactive CaO. Thus, KSO3= - KSO3=T is proportional to
the amount of unreacted CaO at time T. Therefore, the above expression becomes:
cone, unreacted CaOj
CaO reactive Tota|
and the vertical axis of Figure 29 may be expressed as log cone. CaO. The linearity
this plot (Figure 29) is of interest in that it demonstrates a first order dependence on
CaO concentration for the reaction:
CaO + SO2 - » CaSO3
Another series of experiments were conducted monitoring the sulfate band at
1110 cm"' . In one series of experiments, the absorbance at 1110 cm"' was monitored with
time and in the other series of experiments, scans were made at ten minute intervals
between 1300 and 800 cm"' . These runs were all conducted between 380 and 450°C using
a gas mixture consisting of 0.3% SO2, 2% water, 2% oxygen and the balance nitrogen.
The flow rates for all these runs were 800 cc/min. Plots of
1 - obsorbonce-r - initial obsorbonce
Total absorbance change
versus time are shown for the 1 1 10 cm*' sulfate band for one run in Figure 30. As con be
-------
1.0
0.9
0.8
0.7
0.6
0.5
0.4
0.3
0.2
-0.1
0.09
0.08
0.07
0.06
0.05
0.04
0.03
0.02
0.01
0-CL
Figure 29. Reaction CaO + SO2 Monitoring 900 em"1
0.5
1.0
Time (min.)
1.5
62
2.0
2.5
3.0
1.
12/18 - scans
40 60
Time (min.)
100
Figure 30. Reaction CaO + SO2 + 1 /2O2 Monitoring 1110 cm"'
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seen, for conversion to sulfate, a zero order kinetic process appears to be operating. The
1 - obsprbaneeT - obsorbonce initial
Total absorbance change
can be shown to be related to both:
and KSO4-F-KSO4=T
where 504 j is the concentration of sulfate at time Tand SO4 p is the final concentration
of sulfate achievable. Now, as SO^ is related to the amount of reactable CaO Initially
present, the above can be rewritten as
KCaOT
K CaO total reactable
Thus, the zero order kinetics reveals that the rate of conversion to sulfate is independent
of the concentration of CaO present. This zero order dependence was exhibited for all
of the runs monitoring 1110 cm"' . Also, good agreement was obtained between both
methods of monitoring absorbance changes. In all the studies, complete conversions to
sulfate were not achieved, as judged by weight gains. In most cases, about 40 percent
of the CaO present was reactable.
4, Discussion of High Temperature Study Results
(a) Mechanism of Reaction
The experiments performed have shown a first order dependence in CaO
for the reaction:
CaO + SO2 - > CaSOs
and a zero order dependence for:
CaO + SO2 + 1/2O2 - > CaSO4
This can be interpreted in terms of the following mechanistic sequence
1 . CaO + SO2 - > CoS03 (fast)
2. CaSO3 + 1/202 - > CoSO4 (slow)
As reaction 1 is far more rapid, the surface becomes rapidly covered by sulfite, which is
only slowly converted to sulfate. Such a scheme would readily explain the observed
kinetic dependence].
64
Relevant rate data extracted from the series of curves obtained moni-
toring 1110 cm*' is presented in Table IX. A plot of this data appears in Figure 31.
From this plot of rate versus reciprocal temperature, an activation energy of 41.9 Kcal
is computed for the oxidation reaction resulting in sulfate formation. A similar data
evaluation was conducted for the information obtained monitoring the 920 cm"' sulfite
band. Table X gives the rate data obtained from these runs and a plot of this data is also
presented in Figure 31. From this graph of rate versus reciprocal temperature, an acti-
vation energy of 13.9 Kcal/mole was computed for the SO, absorption reaction (step 1).
The rate data presented here clearly shows that at temperature below
~500°C, the rate of oxidation of sulfite
(i.e. CaSO3+ 1/2O2 » CoSO^
is the rate determining reaction. Above 500°C by extrapolation of the data, such as can
be seen from Figure 31, the rate of absorption of SO2 becomes rate determining. Indeed, the
activation energy we are reporting for this step (13.8 Kcal/mole) agrees well with the
overall activation energies reported for the dry limestone process at higher temperatures
(co. 900°Q. Values ranging from 8 to 18 Kcal/mole have been reported at these temp-
era turesO 1).
Table IX. Rate Data for the Reaction CaO + SO2 + 1/2O2—»CaSO4
Temperature Time for Complete Reaction
(°O Minutes Seconds Molecules/cm2sec 1/TxlO3
400
400
380
450
445
63
65
110
6.5
8.3
3.78 x103
3.90x 103
6.6 x103
3.9 x 102
4.3 x102
3.08xl012
2.97x 1012
1.97x 1012
3.00x 1013
2.61 x1013
1.50
1.50
1.53
1.38
1.39
Table X. Kinetic Data of SO2 Absorption
Temperature
400
420
450
408
475
Time
(sec)
372
318
204
264
144
Tx 103°K'1
1.50
1.44
1.38
1.48
1.34
Molecules
(cm2 sec)
3.12x1013
3.65x 10'3
5.70x 10'3
4.40x 10'3
8.05xl0'3
It has been reported at higher temperatures that the overall absorption
and oxidation rate of SO2 is dependent in first order on the SO2 concentration^1). This
finding strongly supports the above conclusions. Indeed, from the above we are postulating
-------
O Oxidation to Sulfate
SO2 Absorptl
1.34 1.36 1.38 1.40 1.42 1.44 1.46 1.48 1.50 1.52 1.54
Figure 31. Kinetic Plots of Data for SO? Absorption
as the rate determining reaction at elevated temperatures
CaO + SO2 - » CaSO3
Such a reaction is expected to be first order in both CaO surface concentrations and in S
(b) Calculation of Pre-Exponential Factors for the Absorption and Oxidation
(1) Oxidation of Sulfite to Sulfote
Using the equation A = A<,e" 4H*/RT, where A is the number of
molecules/site see at a given temperature, Ao is the pre-exponential factor, A H* is the •
activation energy, R Is the gas constant and T the absolute temperature, and rearranging, •
we have:
Using the 450°C data presented in Table IX for sulfate formation, we have, assuming
10-15 crn2/site (i.e. each site = 10 A2):
Ao = 3 x 10-2 molecules/site sec e 41, WP
2(723)
Ao= 9.51 x ID10 molecules/site second
= 9.51 x 1025 molecules/cm2 sec
= 1.58x 102 moles/cm2sec
(2) Absorption of SO2
Using the same equations above and data presented in Table X
we have:
= 7.07 x 102 molecules/site sec
= 7.07 x 10'7 molecules/cm2 sec
Ao = 1 . 18 x 10-* mole«/cm2 Me
A comparison of the above Ao values with each other clearly
demonstrates what was stated earlier in this report (i.e. that at temperatures at which the
dry limestone process is operated (900°Q the absorption of SOj is clearly the overall
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rate limiting step. Indeed, a surface reaction of the type CaO + SO2
would be expected to be first order in both reactants. While time did not allow us to
fully confirm this point on this program, a few experiments could really check the
overall validity of the above postulated kinetic picture. Rates could be measured at
several SC>2 partial pressures using the Series 100 materials. All of these runs, of
course, would be conducted at the same temperature. Also, data obtained by monitor-
ing the rate of sulfate formation at 1110 cm"' taken above 500°C should prove most
revealing when correlated with data obtained above. Indeed, if SC>2 absorption is
rate limiting, then the rate curve for sulfate formation (Figure 31) should exhibit a bend
and follow the same line as the SC>2 absorption curve as, for our experiments, sulfate
cannot be formed raster than SC>2 is absorbed. Indeed, if SC>2 absorption is rate limiting
and the rate of oxidation to sulfate is much faster than the absorption reaction (as is
expected to be the case above 500°Q, then the rates of SC>2 absorption (as determined by
the formation of the 920 cm"' band in the absence of C>2) should be the same. These point;
could readily be checked to give a complete proof of the mechanism postulated in this
report.
68
SECTION III
CONCLUSIONS AND RECOMMENDATIONS
This study has illustrated the applicability of infrared methods to the study of
problems involved in various limestone scrubbing processes. Using infrared spectros-
copy it has been possible to show the mechanism by which limestone removes SO2 from
flue gases in the presence of oxygen, to demonstrate that high temperature reactions with silic
or silicates occur in impure limestone during calcination, and to developa rapid method for
analysis of the sulfate content of reacted limestone which may be extendable to analysis
for other species of interest as well.
The results of this program may be summarized as follows:
1. Pressed pellet techniques, either when used alone as in conjunction with
freeze drying and ion exchange methods, do not give as precise results
as can be achieved by wet chemical means. The pressed pellet method
alone gives calibration curves which contain considerable scatter. When
used with freeze drying techniques, the pellet method can give good
results, but the process becomes very time consuming. Application of
ion exchange techniques does, apparently, lead, in conjunction with
the pellet technique, to a reliable method for sulfate analysis, but this
approach cannot be extended without great difficulty to analysis for
carbonate and hydroxide.
2. A new type of thin film cell has been developed which will allow inves-
tigators to make quantitative studies in water based solvent systems by
infrared methods. The new cell, .003 mm thick consists of two optically
flat silver chloride plates. On the outer portions of the face of one of these
plates is deposited a uniform 3 micron silver film.
3. Studies using a matched pair of such cells (one containing the water
based solvent which was placed in the reference beam and the other
containing test solutions, which was inserted in the sample beam) have
shown that, using saturated tetrasodium EDTA aqueous solution as a
solvent and a differential technique, quantitative analysis of the
sulfate contents of reacted limestone can be readily conducted.
4. Investigations have shown that the techniques developed for sulfate
analysis in solution can be extended to studies of species such as
carbonate, nitrate, nitrite, dissolvable silicate, and phosphate. If
D2O is used in place of water as the solvent, then such technique!
may also be extendable to hydroxide and oxide.
5. Studies, using the KBr technique, have shown that shifts and changes
in intensity per unit weight of sample occur for the silicate bands in
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impure limestones with increasing calcination temperature. This evidence
can be correlated with data on different kinds of silicates and it can be
demonstrated that reactions are occurring with silica or silicates within
the impure limestones during calcination at higher temperatures. Some
electron microprobe data obtained on the same materials as were used for
the infrared studies has revealed that after calcination at low tempera-
tures (1700°F) essentially all the silicates are present in discrete phases
within the limestones. After calcination at higher temperatures (3200°F)
some silicate is present in all -phases of the limestone.
. 6. The hydration of MgO in calcined, slaked dolomitic materials has been
confirmed and evidence has been obtained to indicate either a pref-
erential reaction of the h/drated MgO with some flue gas constituent or
a decomposition of Mg(OH)2 under conditions of reaction of dolomites with
flue gases.
7. Infrared studies of the reaction of SC>2 with limestone in the 380-475°C
temperature range have provided a large volume of useful kinetic data. •
. It has been demonstrated that the absorption reaction, CaO + SO2 >
CaSC>3, with an activation energy of 13.8 Keal/mole is the rate deter-
mining reaction above 500°C. At lower temperatures, the oxidations
reaction, CaSC>3 + l/2Oo——> €0504, with an activation energy
of 41.9 Kcal/mole is rate determining. Pre-exponential factors have
also been obtained for the above reactions.
To further utilize the techniques developed in this study and to gain further needed
information by infrared methods on some of the problems involved in limestone scrubbing
the following set of recommendations are made:
1. Thin film type cells of the variety used in the development of the
method for sulfate analyses should be developed from more resistant
materials than silver and silver chloride. An investigation should be
made of other window and film materials to determine the best com-
binations of useable compounds. Specifically, materials such as .zinc
sulfate, zinc selenide, silicon, KRS-S, germanium and arsenic trisulfide
glasses, which are known to be water resistant should be investigated
for use as possible window materials to be used with the alkaline (pH ~9)
EDTA solutions. A study should also be made of the use of thtn films of
materials such as Teflon and tantalum, which are both corrosion and
.wear resistant, for use as film or spacers.
2. The design of the thin film cells used in this study, while novel, presented
some difficulties. With the thickness of the cells being only 3 microns,
hov«ver, problems have been encountered in filling of these cells. In
future work, studies should be made with the aim of developing a more
readily fillable cell. One possible approach to this problem lies in
the design of an evacuable cell, where the test solutions can be drawn
into and through the cells via evacuation.
70
3. The methods and techniques developed for rapid analysis of sulfate in an
aqueous based solution should be extended to development of procedures
for rapid analyses for species such as nitrate, nitrite, carbonate, bicar-
bonate, sulflte, and bisulfite. Efforts should be made to determine which
of these species and sulfate can be determined in the presence of one
another in solution. Perhaps, it may be possible, from one spectrum,
to determine the concentrations of several species of interest.
4. Work with the methods developed for sulfate analysis should be
extended to the use of D2
-------
c. Volumetric experiments to determine the rote of SC>2 absorption
at various SOo partial pressures in the same temperature range
as the kinetic studiec presented in this report. This additional
data will provide enough information to either prove or disprove the
mechanisms proposed in this report.
72
SECTION IV
REFERENCES
1. Hill, W. E. and Goebel, E. D., State Geological Survey of Kansas Bulletin 165
Part/, 1963.
2. White, R. G., Handbook of Industrial Infrared Analysis, Plenum Press, N. Y.,
p. 107, 1964.
3. Thompson, W. K., Trans. Faraday Soe. 61, 1965, p. 2635.
4. Burton, J. S., Final Report, Contract No. PH 86-68-78, January 1969.
5. Technical Bulletin, Amberlite 120, Rohm and Haas Company, Philadelphia, Pa.
6. Senior, W. A. and Vemall, R. E., J. Phys. Chem. 73, 1969, p. 4242.
7. Jones, R. N. and Sandorfy, C., "Infrared and Raman Spectra, Applications in
Chemical Applications of Spectroscopy", W. West(ed.), Interscience, N. Y,, 1956,
p. 246.
8. Nachod, F. C. and Martin, C. M., Appl. Spectroscopy 13, 1959, p. 45.
9. Stemglanz, H., Appl Spectroscopy 10, 1956, p.77.
10. Kulbom, S. D. and Smith, H. F., Anal. Chem. 35, 1963, p. 912.
11. Borgwardt, R. H., Environmental Science and Technology, 4, 1970, p. 59.
12. Schafer, H. N. S., Anal. Chem. 35, 1963, p. 53.
13. Fritz, J. S. andYamamura, S. S., Anal. Chem. 27, 1955, p. 1461.
14. Serratosa, J. M. and Bradley, W. F., JACS 63, 1958, p. 1164.
15. Robinson, D. Z., Anal. Chem. 24, 1952, p. 619.
16. Sulfur Oxide Removal from Power Plant Stock Gas - Tennessee Volley Authority, 1968.
17. Hotfield, J. D., Report for 3rd Limestone Symposium, Clearwater, Florida, Dec. 1967.
18. Porter, A. E., Report for 3rd Limestone Symposium, Clearwater, Florida, Dec. 1967.
19. Boynton, R. S., "Chemistry and Technology of Lime and Limestone, New York, 1965,
pp. 17-20.
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20. Boynton, R. S., "Chemistry and Technology of Lime and Limestone, New York, 1965,
pp. 17-20.
21. Le«, H., "Refractories from Ohio Dolomite", Ohio State Union Eng. Ext. Sta. News.
XIX, April 1947, p. 2.
74
SECTION V
ACKNOWLEDGEMENTS
The author wishes to thank Dr. Joshua Bowen, Dr. Dennis Drehmel, and Mr. Robert
Larkin of the National Air Pollution Control Administration, Fairfax Facility, Cincinnati,
Ohio for helpful discussions and for providing some of the samples used in this study.
The author also wishes to thank Dr. Robert G. Shaver of General Technologies
Corporation for helpful discussions and Messrs. Leon Ferguson and Frank Lysy, also of
GTC for excellent experimental assistance on parts of this program.
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