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Color removal from a neutral sulfite waste using magnesium coagulation

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Title:
Color removal from a neutral sulfite waste using magnesium coagulation
Creator:
Taylor, James S
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Language:
English

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City of Miami ( local )
Magnesium ( jstor )
Coagulation ( jstor )
Polymers ( jstor )

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University of Florida
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University of Florida
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Copyright James Sherman Taylor. Permission granted to the University of Florida to digitize, archive and distribute this item for non-profit research and educational purposes. Any reuse of this item in excess of fair use or other copyright exemptions requires permission of the copyright holder.
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Full Text










COLOR REMOVAL FROM A NEUTRAL SULFITE
WASTE USING MAGNESIUM COAGULATION







by

JAMES S. TAYLOR





















A DISSERTATION PRESENTED TO THE GRADUATE COUNCIL OF
THE UNIVERSITY OF FLORIDA
IN PARTIAL FULFILLMENT OF THE REQUIREMENTS FOR THE
DEGREE OF DOCTOR OF PHILOSOPHY

UNIVERSITY OF FLORIDA 1976























TO

JANET

JIMMY

AND

BRIT











Acknowledgements


I wish to express my gratitude to my committee chairman, John Zoltek, Jr. for his overall guidance, understanding and friendship in assisting me in my research. I am deeply indebted to T. deS Furman and J. Edward Singley whose technical and exemplary contributions to my education will never be forgotten. The technical insight and timely assistance given me by Ellis D. Verink, Jr. are' sincerely appreciated. I wish to extend my appreciation to H.F. Berger, who, through the National Council for Air and Stream Improvement, made the funding of my research possible. I will always hold his cooperation and patience in high esteem. The contributed research and extensive laboratory work by Gary Christopher and Bevin Beaudet in completing their masters projects is acknowledged and appreciated.

I am sincerely grateful for the sacrifices my wife Janet made and the contributions of my parents in enabling me to pursue my education. The values set forth by my parents years ago came to bear during my research.

I have had many rewarding experiences at the University of Florida and am grateful for the opportunity to have been part of that institution.















iii













Table of Contents

Page

Acknowledgements iii

List of Tables vii

List of Figures ix

Abstract xii

Chapter

1- INTRODUCTION 1
1-1 General Background 1 1-2 Legal Requirements 3 1-3 Purpose of This Work 3

2- COLOR 5
2-1 Color in the Electromagnetic Spectrum 5 2-2 Lignin 7 2-3 Characteristics of Color 9 2-4 Coagulation 13 2-5 Color Removal by Coagulation 23 2-6 Magnesium Coagulation 25 2-7 Color in Pulp Mill Effluents 27

3- LABORATORY PROCEDURES 32
3-1 Feed Solutions 32
3-1.1 Synthetic Waste Solutions 32 3-1.2 Coagulation Chemicals 33 3-1.3 Polymers 33
3-2 Analytical Equipment and Techniques 34
3-2.1 Total Carbon Measurements 34 3-2.2 Color Measurement 34 3-2.3 Incineration 35 3-2.4 Jar Tests 35 3-2.5 Metal Analysis 37 3-2.6 Mobility Measurements 37 3-2.7 pH Measurements 38 3-2.8 Settling Tests 38 3-2.9 Solids Analysis 38 3-2.10 Titration Curves 39
3-3 Experimentation 39
3-3.1 Coagulation Experiments 39 3-3.2 Coagulant Recovery 39 3-3.3 Coagulant Recycle 40




iv












Chapter Page

4- RESULTS 41
4-1 Determination of Coagulation pH and
Coagulant Dose 41 4-1.1 Coagulation pH 41 4-1.2 Coagulant Dose 43
4-1.3 Variation of Coagulation pH with
Coagulant Dose 44 4-1.4 Magnesium Remaining in Solution as
a Function of Final pH 47 4-1.5 Magnesium and Ca(OH) Dose as a
Function of Initial aste Color 49 4-2 Waste Characteristics 51
4-2.1 Untreated Waste Titration Curves 51
4-2.2 Comparison of Untreated and Treated
Waste Titration Curves 58 4-2.3 Waste Content 60
4-3 Color Removal Mechanism 60
4-3.1 Color and Magnesium Titration Curves 60
4-3.2 Magnesium, Calcium, Color and Organic
Carbon Residuals After Coagulation 63 4-3.3 Stoichiometry of Color Removal from
NSSC Waste by Magnesium Coagulation 71 4-4 Settling of Coagulated Wastes 75
4-4.1 Purpose of Settling Tests 75 4-4.2 Sludge Settleability 77 4-4.3 Mechanisms of Sedimentation 82
4-5 Magnesium Recovery and Recycle 89
4-5.1 Recovery Methods 89 4-5.2 Process Reversibility 90 4-5.3 Color-Cation Interaction 90 ++
4-5.4 Chemical Equilibrium of Mg -CO2-H20 96 4-5.5 Sludge Incineration 103 4-5.6 Magnesium Recovery 107 4-5.7 Magnesium Reuse 118

5- DESIGN OF A COLOR REMOVAL PROCESS FOR A NSSC
WASTE USING MAGNESIUM COAGULATION AND RECOVERY 123 5-1 Coagulation 123 5-2 Sedimentation 124 5-3 Vacuum Filtration 127 5-4 Incineration 129 5-5 Carbonation 130

6- COST 134
6-1 Chemical Costs 134 6-2 Capital and Operation Costs 134 6-3 System Costs 137





v











Chapter Page

7- CONCLUSIONS AND RECOMMENDATIONS 145
7-1 Conclusions 145 7-2 Recommendations 147 REFERENCES 149 Biographical Sketch 156













































vi















List of Tables


Table Title Page 2-1 VISIBLE SPECTRUM AND COMPLIMENTARY COLORS 6 4-1 GRAPHIC DETERMINATION OF pKa OF SODIUM BASE NSSC WASTE 57 4-2 UNTREATED AND TREATED NSSC WASTE ANALYSIS 61 4-3 POLYMER DESCRIPTION AND SVI FOR POLYMER ASSISTED
SLUDGES 78 4-4 ELECTROMOBILITY AND ZETA POTENTIAL FOR MAGNESIUM
SLUDGE PRODUCED IN TAP WATER AND NSSC WASTE AT
'VARYING pH 86 4-5 CaCO3 PRECIPITATION IN A NSSC WASTE 94 4-6 CHEMICAL REACTION AND pK VALUES CONSIDERED FOR
Mg++-CO2-H20 SYSTEM 100 4-7 AVERAGE CHARACTERISTICS OF A SLUDGE PREPARED
BY COAGULATING A NSSC WASTE 104 4-8 MgO REACTIVITY AS AFFECTED BY TEMPERATURE 106 4-9 INCINERATED SOLIDS ANALYSIS 108 4-10 CARBONATION OF INCINERATED SLUDGE AT VARYING
CONCENTRATIONS OF NONVOLATILE SOLIDS FOR
MAGNESIUM RECOVERY 113 4-11 COLOR REMOVAL BY LIME-MAGNESIUM COAGULATION
USING THE SAME MAGNESIUM THREE TIMES 119 4-12 COLOR REMOVAL BY LIME-MAGNESIUM COAGULATION
USING THE SAME MAGNESIUM TWICE 121 5-1 SOLIDS LOADING FROM SETTLING BASIN 128 5-2 DESIGN SUMMARY FOR THE TREATMENT OF A NSSC WASTE 132 6-1 CHEMICAL COST TO TREAT A NSSC WASTE '135





vii












Table Title Page


6-2 UNIT OPERATIONS COST SUMMARY
NSSC WASTE COLOR = 5000 139 6-3 PROCESS COST SUMMARY. IN $/1000 GALLONS OF
NSSC WASTE 140 6-4 UNIT OPERATION COST SUMMARY
NSSC WASTE COLOR = 2500 142 6-5 NSSC PRODUCT COST INCREASE DUE TO COLOR
REMOVAL BY MAGNESIUM COAGULATION 144









































viii














List of Figures


Figure Title Page 1.1 NSSC flow diagram 2 2.1 Quinonemethide 8 2.2 Constitution scheme for lignin 10 3.1 Standard Pt-Co color curve 36 4.1 Color residual as a function of final pH 42 4.2 Color residual as a function of final pH 42

4.3 Comparing NaOH and Ca(OH)2 for color removal
via magnesium coagulation 45

4.4 Comparing NaOH and Ca(OH)2 for color removal
via magnesium coagulation 45 4.5 Verification of coagulation pH' 46

4.6 Magnesium remaining in solution as a function
of final pH 48

4.7 Lime dose as a function of initial waste color
for magnesium coagulation 50

4.8 Magnesium dose as a function of initial waste
color using lime 50

4.9 Titration curve of sodium base Mead NSSC waste
with color equal to 2500 52

4.10 Titration curve of sodium base Mead NSSC waste
with color equal to 5000 53

4.11 Titration curve of sodium base Mead NSSC waste
with color equal to 10,000 54

4.12 Titration curve of sodium base Mead NSSC waste
with color equal to 20,000 55

4.13 Titration curve of sodium base Mead NSSC waste
with color equal to 40,000 56


ix












Figure Title Page


4.14 Titration curve of treated and untreated
NSSC waste 59

4.15 Titration curve of raw waste dosed with
magnesium 62

4.16 Organic carbon and color residuals as a
function of final pH 64

4.17 Magnesium and calcium residuals as a function
of final pH 64

4.18 Color, T.O.C., and Mg++ residual after Mg++
coagulation using Ca(OH)2 for pH control 68

4.19 Color, T.O.C., aid Mg++ residuals after Mg++
coagulation using NaOH for pH control 69 4.20 Ratios of [OHJMg++/ Mg++ 72

4.21 Color and pH of a NSSC waste as a function
of Ca(OH)2 concentration 76

4.22 Sludge settling velocity for polymer assisted
and raw sludge 81

4.23 SVI and zeta potential vs. polymer dose for
a nonionic polymer #1905N 85

4.24 SVI and zeta potential vs. polymer dose for
an anionic polymer #837A 85

4.25 Zeta potential of magnesium solids in tap
water and NSSC waste at varying pH 87

4.26 Equilibrium concentrations of Mg++ and Mg(OH)+
with Mg(OH)2 at varying pH 87 4.27 Color reversibility bar graph 91

4.28 Color remaining as a function of CaCO3
precipitation 93

4.29 Color remaining as a function of MgF2
precipitation 93 4.30 Activity ratio diagram for log CT = -1 97



x











Figure Title Page


4.31 Solubility diagram of Mg++ in a CT = 10-1 M
carbonate system 98 4.32 Predominance diagram for log Mg++ = -1 99

4.33 Color/Mg++ ratio as a function of incineration
temperature 109

4.34 Precipitation of MgCO 33H20 by aeration at
various temperatures 111

4.35 pH as a function of carbonation at various
nonvolatile solids concentrations 114

4.36 Magnesium recovered as a function of
carbonation time 115

4.37 % magnesium recovery as a function of
nonvolatile solids concentration 116 5.1 Design data for sedimentation 126

5.2 Flow diagram for lime-magnesium color
removal process 133




























xi










Abstract of Dissertation Presented to the
Graduate Council of the University of Florida
in Partial Fulfillment of the Requirements for the Degree of Doctor of Philosophy


COLOR REMOVAL FROM A NEUTRAL SULFITE WASTE USING MAGNESIUM COAGULATION

By

James S. Taylor

August 1976


Chairman: John Zoltek, Jr.

Major Department: Environmental Engineering Sciences



A color removal process was developed on a laboratory scale that would remove 90% of the initial color of a neutral sulfite semi-chemical (NSSC) pulp waste. The colored waste was coagulated at a pH of 11, with stoichiometric amounts of magnesium and Ca(OH)2. The magnesium and Ca(OH)2 doses were represented by linear equations.

The amount of magnesium required for 90% color removal was reduced 25% when Ca(OH)2 was used for pH control. The reduction in coagulant dose was due to the chelation of the divalent calcium ion and organic acids in the waste. Titrametric techniques demonstrated that the color removal process removed 40% of the acid strength of the NSSC waste, and that 65% of the acids removed had a pK greater than 9.

The zeta potential of the coagulated NSSC waste was -1.00 my at
++
pH 10.3 and zero at pH 12.5. Measurements of the (Mg ), COH ), organic carbon concentration and color removed during the coagulation process indicated that the Mg ion first chelated the organic acids




xii










causing a 35% color increase. The Mg++ ion then formed a precipitate which resulted in color removal. The empirical formula for the precipitate was Mg(OH)1.5R, where R represents the precipitated organic acids. Once the magnesium precipitate formed, the molar ratios of the magnesium removed to the hydroxides removed was 1.5 for varying magnesium doses at constant pH. The consistency of the molar ratio at varying doses indicated color bodies were removed in magnesium coagulation by a chemical reaction.

The color removal process was demonstrated to be completely

reversible by varying the pH. In order to reuse the magnesium, the sludge was incinerated to remove the color from the magnesium solids. The optimum temperature of incineration was found to be 5500C. After incineration, all of the magnesium was recovered by bubbling a 10% CO2-90% air mixture through a slurry containing an incinerated solids concentration of 5318 mg/l. The fraction of magnesium solubilized from the incinerated solids was controlled by the Mg -CO2-H20 system. The controlling solid phase was MgCO33H20.

The same magnesium was used three times to remove 90% of the color from three separate aliquots of NSSC waste. After three uses of the coagulant, 93% of the magnesium was recovered.

The cost of using this process to treat a NSSC waste with an initial color of 2500 and a flow of 10 mgd was estimated to be $0.27/1000 gal.










xiii














CHAPTER 1



INTRODUCTION



1-1 General Background

Pulp and paper manufacturing is one of the largest industries

in the United States. It is also one of theomajor water using industries in the nation, producing from some mills extremely large volumes of highly colored effluents, which are typically discharged to waterways. Color creates a unique problem in a stream. It is readily identifiable in an aesthetic sense and can detract from the natural beauty of a body of water. The amount of light penetrating a stream would be affected by a colored waste discharged to that stream, and could threaten the eco-system in that stream. The National Council for Air and Stream Improvement, a pollution abatement research organization sponsored by member pulp and paper companies, has recognized this problem and has sought for many years to devise economical and effective color removal processes for all pulp and paper plant effluents.

A highly simplified NSSC pulping process diagram is illustrated in Figure 1.1. The wood is prepared for the digestion process by removing the bark and increasing the surface area by a chipping process. The wood-chips are then screened, fed into a digestor and mixed with a sulfite cooking liquor. The function of the sulfite


1




2






Debarking Chipping Screening





Sulfite Cooking Digestion Cooking Liquor







Water Washing





Colored
Waste Bleaching Effulent (chlorine)
(oxygen)






Washing




Further Processing


Fig. 1,1 NSSC flow diagram






3




cooking liquor is to separate the lignin from the wood fiber. After the digestion process, the highly colored water soluble lignin is separated from the pulp by washing. The aqueous washings constitute part of the waste effluent. Depending on the ultimate use of the pulp, additional color removal may occur. Bleaching will further lighten the pulp. After each bleaching operation the pulp is washed, producing additional color in the final waste effluent.



1-2 Legal Requirements

In 1968 the United States Government passed the Clean Water Act. It was amended in 1972 to include all pulp and paper mills using a NSSC production process. In the Federal Register under Pulp, Paper, and Paperboard Point Source Category, Effluent Guidelines and Standards, this law states in summary: All NSSC plants must remove 75% of their effluent color by 1983, and all new NSSC palnts built after 1975 must remove 75% of their effluent color.



1-3 Purpose of This Work

It was the purpose of this research to develop a color removal process for NSSC waste and to give insight into the mechanism by which that color was removed. The first objective was to develop a method of NSSC color removal that could be evaluated for use as a full-scale treatment process. The investigation was limited to jar testing-techniques, with subsequent sludge incineration and coagulant recovery on a laboratory scale. The second objective was to investigate the mechanism by which color removal occurred.




4





Techniques employed in this phase of the research were chemical analyses in conjunction with the determination of the stoichiometric relationships developed in the color removal chemical reactions.












CHAPTER 2



COLOR



2-1 Color in the Electromagnetic Spectrum

Color is a qualitative parameter that does not lend itself to exact engineering measurement. Within the visible region of the spectrum, persons with normal color vision are able to correlate the wavelength of light striking the eye with the subjective sensation of color. Table 2-1 shows the color perceived related to the wavelength.

Objects are seen by either transmitted or reflected light. When "white light," containing the entire spectrum of visible wavelengths, passes through a medium such as a solution of NSSC waste, the medium appears colored to the observer. Since only the transmitted waves reach the observer, their wavelengths determine the color of the medium. Chromophores, or color producing compounds, absorb certain wavelengths of the spectrum depending on the electronic structure of the compound. A change could occur in the electronic configuration of a compound which could change it from a colorless to a colored compound. The oxidation of an alpha-quinone would produce a colorless degradation product, but the self-condensation of the same alpha-quinone would produce a colored product.

Very little evidence has been gathered.on the amount of environmental degradation caused by color. Properties of pollutants such



5





6








TABLE 2-1

VISIBLE SPECTRUM AND COMPLIMENTARY COLORS



Wavelength, mu Color Complimentary Color


400 435 Violet Yellow-green

435 480 Blue Yellow 480 490 Green-blue Orange

490 500 Blue-green Red

500 -.560 Green Purple 560 580 Yellow-green Violet

580 595 Yellow Blue

595 610 Orange Green-blue 610 750 Red Blue-green



Source: Day, R.A., Jr., Underwood, A.L., Quantitative Analysis, Second
Edition, Prentice-Hall, Englewood Cliffs, N.J., (1967).





7




as available nutrients or oxygen demand have been shown to degrade the environment. However, the discharge of highly colored effluent would definitely affect the aesthetic quality of the receiving waters. Color would have a detrimental effect on process water used in he production of highly bleached paper.



2-2 Lignin

Lignin is one of the most abundant natural products on earth, constituting about one-fourth of the woody tissue in plants. It is responsible for most of the color present in natural waters. The natural formation of this cross-linked polymeric material from coniferyl alcohol and related substances is not presently completely understood. Despite considerable research, the structural characterization of lignin has been only partially successful.

Freudenberg (1966) gathered information about lignin structure

from direct oxidation of lignin, from bio-chemical experiments related to alcohols, and from lignin degradation with strong alkali, methylation and oxidation. His experiments enabled an estimation of the relative amount of alcohols which served as building blocks of lignin. Lignification occurs in plant cells when alcohols are liberated and oxidized by natural organic compounds in the presence of air. The free radicals produced then combine and build up lignin. Freudenberg (1966) formed a quinonemethide, as shown in Figure 2.1, by combining radicals that resulted from lignol dehydrogenation. From these experiments he suggested that quinonemethide was a tentative structural unit in lignin. Since the quinonemethide has no opportunity to become





8














H2COH

HC HC





H2COH
IH I OMe HC O HC
I

HC
I







9\ OMe
0



Fig. 2.1 Quinonemethide





9




stabilized by hydrogen migration, it adds on the external electrolytes, particularly hydroxyl compounds and preferably water. Quinonemethide is a chromophore, is yellow and can be easily recognized by its intense color absorption extending into the beginning of the visible range. Quinonemethide can achieve limited stability through polymerization, creating large molecules that can still interact with polar compounds.

It is possible to construct a tentative constitutional scheme for spruce lignin, which probably is similar to other wood lignins. Such a scheme is presented in Figure 2.2. The lignols which originate during lignin formation, together with the hydrolysis products, reveal different ways in which the C6C3 units are combined.

Through natural and industrial processes the lignin is separated from wood fibers and produces chromophores in aqueous solutions. Kirk et al. (1969) prepared lignin by bacterial degradation of wood. The lignin was fractionated by molecular gels into three separate fractions, all of which would absorb light in the visible spectrum. Alder et al. (1966) degraded spruce lignin by acid refluxing in an organic aqueous solution, and was able to separate through fractionation several products that were color producing compounds.



2-3 Characteristics of Color

Many investigators have attributed the color present in water to the natural or induced degradation products of organic matter. Saville (1917) concluded through electrophoretic studies that most organic color was negatively charged and existed in the colloidal size range. Black and Christman (1963a) found that color collected








H2OH H2COH
H HC- 1/2 1/2
-CH H2COH HCO H2COH HC --CH H HC HCHC OMe HC H
OMe I Io
H2COH OH OH
0 CH MeO
I I
HCH-C0 0jC


H2COH OMe I
I HC CH
_ I I
HCO(C6H05)n H HC H2



OHCH 0 HM
HCOH ----
MeO MeO
-0
Me0 0-O OMe
o I
HCOH-" r OMe
CH
HCOH CH COH cO I I I
MeO OMe 0 OH C CH CH

SHCO CH2
CH- Q, H2COH

HC
6 lOMe
H2OH26
Fig. 2:2 Constitttion scheme for lignin











from ten different water samples had similar chemical and physical characteristics. They demonstrated by dialysis that most of the color present in the ten samples resulted from colloidal suspensions. The infrared spectrum for each of the fulvic fractions, the equivalent weights of those fractions, and the concentrations of the fulvic and humic fractions in each colored sample were similar. Black and Christman (1963b) demonstrated that color intensity was pH dependent and would increase with increasing pH. They also found by dialysis that color existed as a colloid, because only 10% of the original sample color could pass a 4.8 micron filter.

Shapiro (1958) found that organic color was mainly dicarboxylic hydroxy aliphatic organic acids of molecular weight 450. He suggested that if phenols were present they were non-color producing organic compounds. He also found that the salts of these acids would pass a cellophane membrane, indicating that they were not colloidal. Shapiro (1958) demonstrated, by chromatographic comparisons, that chemical patterns of color samples taken from different lakes across the country were similar. Any differences that existed in these samples were due to inorganic constituents of the water. Black (1960) suggested that separation techniques used by Shapiro (1958) excluded a portion of the color bodies, and that the excluded portion was in the colloidal size range.

Christman and Ghassemi (1966) isolated seven different phenolic compounds common to wood and water humics. Their organic analysis on wood lignins identified carboxyl and phenol groups as the major building units in color molecules. They described these groups as




12




large aromatic molecules with hydroxyl, methoxy and carboxylic functional groups. Christman and Ghassemi (1966) also found that color extracted from soil would increase with time of soil contact and temperature of the aqueous color medium. Their research showed, as had that of Black and Christman(1966), that color increased with an increase in pH. However, this increase was not linear over the entire pH range.

Taylor and Zoltek (1974), using a kraft effluent treated for

color removal by massive Ca(OH)2 treatment, found that color increase in the waste occurred when the waste was in contact with soil or light. The amount of color increase in the soil-contacted samples was directly proportional to the organic content of the soil. Gjessing and Samdal (1968) studied color fluctuation in a chain of four Norwegian lakes and found that color decreased in all of them except the last lake, where an impoundment occurred. The last lake had a high organic matter content. Gjessing and Samdal (1968) recorded a direct increase in the color of the impounded lake with time of water storage. Their data led to the conclusion that solubilized organic matter produced the color increase in the impounded lake, and the degree of color increase depended on time of impoundment.

Packham (1964) separated color from seven different waters into the same classes as did Black and Christman (1966). He found, based on filtration of the fractions, that the-fulvic acid fraction existed in the colloidal size range and that the humic acid was in the molecular size range. Packham (1968) also revealed that both the fulvic and humic fractions consisted of complex mixtures of many different





13




organic acids. Gjessing and Lee (1967) fractionated the color present in a natural water by gel filtration and found molecular size distributions ranging from greater than 200,000 to as low as 700. They found that the molecular size fraction that contained the largest concentration of organic carbon did not produce the greatest color.

Midwood and Felbeck (1968) purified a yellow color from organic muck and found that t1e organic matter producing color was resistant to chemical or biological degradation. They found that over 80% of the organic carbon was present in the fulvic portion of the color. The infrared spectra showed that aromatic carboxylic acids with aliphatic side groups containing phenolic hydroxyl groups were major components of the color molecules. Day and Felbeck (1974) obtained a yellow water-soluble organic exudate from the domestic waste water fungus Aureobasidium pullulans. The exudate contained no humic acid, although it was yellow and was very homogeneous relative to fulvic acid extracts from soil. Day and Felbeck (1974) demonstrated that fungal activity was one source of color in watersheds, and concluded that watershed management with respect to excess biological activity may help eliminate color problems in watersheds.



2-4 Coagulation

The reader is referred to comprehensive literature reviews on general coagulation that were published by the American Water Works Association (1971) and O'Melia (1972). In this section the emphasis will be on coagulation as it refers to color removal.





14




A colloidal dispersion is electrically neutral, so that the charges on the colloidal surface must be counterbalanced by the charges on the liquid immediately adjacent to the colloidal particle. As a result, an electrical double layer exists at every solidliquid interface. These charged ions are attracted to the colloidal surface electrostatically and repelled due to diffusion. The VerweeyOverbeck model as described by Osipow (1972) stated that the LondonVan der Waals forces were the forces of attraction for colloids in suspension. The forces of repulsion resulted from the electrical repulsion of the separate colloidal double layers. Osipow (1972) demonstrated that this model was.further developed and modified by Guoy, Chapman, Stern and Helmholtz. The essence of the final model was that colloidal suspension would be destabilized if the electrical repulsive forces were reduced such that the London-Van der Waals forces would dominate, -causing coagulation and sedimentation. This concept was supported in some systems by the Schulze-Hardy rule, which states that the critical coagulation concentration of mono-, di- and trivalent ions to coagulate sols of the opposite charge are in the ratio of 100:1.6:0.13. Matijevic et al. (1964a) developed a stabilization-destabilization model for AgBr and AgI suspensions based on neutralization of the electrical double layer with counter ions gained from the hydrolysis of Al(N03)3. Matijevic et al. (1964b) attributed the destabilization of the sols to the Al species on the basis of charge reversal in the coagulation reaction. However, there was a stabilization of the sol which was followed by another sol coagulation. Matijevic et al. (1964b) contributed










the final destabilization to AI(OH)3 precipitation. They presented no explanation for the restabilized sol prior to AI(OH)3 precipitation, since the sol charge remained positive after the first coagulation.

LaMer (1967) developed a bridging theory which provided an

acceptable qualitative model for describing the destabilization of colloids with polymers. The main points of the bridging theory were: 1) the polymer must contain chemical groups that would interact with the colloidal surface; 2) that when this happens only a part of the colloidal surface was covered, and the remainder of the polymer would serve as a bridge upon attachment to another colloid; 3) if no other colloid was available for attachment, or the polymer concentration was too great, the polymer would attach itself to the colloid and restabilize the suspension; 4) intense agitation would sever the polymer bonds to the colloidal surface and possibly restabilize the suspension; 5) the amount of colloidal surface area present was directly proportional to the amount of polymer required for coagulation. The bridging theory explained how chemical interactions between an anionic polymer and negative colloid would produce coagulation.

Packham (1968) studied coagulation of eight different clays by

aluminum hydrolysis and found the coagulant dose continually decreased with increasing concentration. Solubilized calcium and magnesium assisted in lowering the coagulant concentration necessary to destabilize clay suspensions. Packham (1968) demonstrated that the hydrolysis products of alum were important to clay destabilization by





16



zeta potential measurements of clay suspensions dosed with and without alum. Although the alum floc had the same zeta potential for optimum destabilization as did the clay suspension, the zeta potential was not zero. Apparantly electrostatic forces were not controlling destabilization. Schott (1968) studied the deflocculation of water sorping clays by anionic and nonionic surfactants. He found that maximum deflocculation was produced when the surfaces of the clay lattices were completely covered with the nonionic surfactants.

The American Chemistry Society (1968) published data for aluminum hydrolysis in colloidal suspensions showing that the polynuclear species of aluminum were important destabilization factors. A colloidal suspension was destabilized before any floc was formed using alum as the coagulant. They suggested the forces of adsorption between the colloids and the hydrolysis products were responsible for destabilization, because the hydrolyzed species were hydrophobic and were more likely to accumulate at the solid-liquid interface. Another factor leading to colloidal destabilization was that the hydrolysis products had more than one OH- ion that could sorp at the interface. Their data indicated that as the colloidal surface area concentration increased, an increasing coagulant dose was required to destabilize the colloidal suspension.

Langelier and Ludwig (1949) experimented with calcium and alum

flocculation of four different turbid waters varying in exchange capacity. They concluded that the mechanism of colloidal destabilization was controlled by the exchange capacity of the colloids. Michaels (1954), studying the degree of polymer hydrolysis that best promoted coagulation, found that a small amount of hydrolysis was best suited





17




for destabilization. He suggested that the destabilization mechanism was a two step process: 1) polymer sorption onto the colloidal surface and 2) interparticle bridging following polymer sorption to destabilize the colloidal suspension.

Black et al. (1965) evaluated coagulation by anionic polymers and demonstrated destabilization followed by restabilization with excess polymer concentration. Since both the polymer and the colloid were negatively charged, the destabilization was not attributed to coulombic forces, but to the build-up of interparticle bridges through other than electrostatic mechanisms. They also found that a higher velocity gradient for a shorter time period was more effective in destabilization than a lower velocity gradient for a longer time period. Ragunathan et al. (1973) treated turbid waters with alum and concluded, from zeta potential measurements, that the hydrolysis products of alum were controlling destabilization by sorption mechanisms.

Posselt et al. (1968a) examined metal sorption onto a MnO2

anionic sol and found neutral and anionic species did not sorp, a fact supporting an electrostatic mechanism for destabilization. Posselt et al. (1968b) studied Ca sorption onto a negative MnO2 sol and found Ca sorption onto the MnO2 colloidal surface approached a limiting value. The limiting Ca++ sorption indicated a Langmuir monolayer was probably occurring on the MnO2 surface. They restabilized the suspension with more polymer addition, but did not achieve restabilization with increased metal ion concentrations. However, the increased calcium concentration did broaden the optimum range for coagulation.





18



They suggested that a choice of coagulant aid would be based on the potential determining ions of the sol. Robinson et al. (1974) reported that larger increases in the turbidity of a river water increased treatment costs and large quantities of alum were required to produce potable water. Nonionic and cationic polyelectrolytes were found to be more effective than alum, suggesting that for this water an electrostatic mechanism was not controlling destabilization. Aluminum hydrolysis was probably removing turbidity by enmeshment in a sweep floc.

LaMer (1967) defined coagulation as a kinetic process going from a quasi-stable to a more stable phase, and flocculation as the bridging of already coagulated particles that entered into hindered settling. As an example, he cited hydroxyl groups. on flat clay surfaces bonding with hydroxyl radicals of polymers, which allowed metal ions to form insoluble phosphates. LaMer (1967) suggested turbidity, subsidence rate and floc filtration as methods of evaluating destabilization. He also suggested that a negative polymer would best destabilize a negative colloid, because many sites were produced by polymer hydrolysis for bridging.

Birkner and Morgan (1968) measured particle size distribution during coagulation and found stronger floc was produced as floc diameter increased. They demonstrated the rate controlling step was particle agglomeration after coagulation, and that intense agitation was responsible for limited polymer sorption. Dollimore and Horridge (1972) investigated flocculation of China clay using polyacrylamides. They found that the maximum clarity was not coincident with the maximum filtration rate as measured by the Kozeny-Carmen





19




equation. They concluded that the effective length of the flocsupernatant interface was the controlling flocculation parameter. Hahn and Stumm (1968), studying the kinetics of alum hydrolysis for SiO2 sols, determined that there were three steps in the coagulation process: 1) forming polynuclear hydrolysis products; 2) the rate of surface coverage or adsorption of the polymer on the colloidal surface; and 3) the rate of particle transport. The rate limiting step for SiO2 coagulation was shown to be the rate of particle transport. The rate of coagulation was shown to be a function of the collison rate and the collison efficiency.

Tenney and Stumm (1965) demonstrated that hydrolyzing metal ions and organic polymers could be used to successfully coagulate bacteria. A linear relationship was found between the optimum concentration of the polyacrylamide polymer and the bacterial concentration. They also found that phosphates were removed with Al+++ in a chemical reaction, and the optimum pH for the reaction was the same as for optimum bacterial flocculation.

Stumm and Lee (1961) found that the rate of oxidation of ferrous iron was directly proportional to pH. They found an increase of one pH unit near neutral pH resulted in a 100 fold increase in oxidation rate. Schenk and Weber (1968) also determined that the rate of oxidation of ferrous iron increased with increasing pH. They found that silica retarded the hydrolysis of Fe The hydrolysis was not represented by a first order reaction, but approached linearity with time. They suggested that the solubility relationship may have been altered by complexes formed between silica and iron, and that





20




these complexes may have been the mechanism by which activated silica functioned as a coagulant aid. Mohtaoi and Rao (1973) investigated the effects of temperature on aqueous suspensions and concluded that temperature had no perceivable effect on the zeta potential of the sols, or the alum hydrolysis products mixed with cationic, anionic and nonionic polymers. Charge neutralization was determined to be important in destabilizing a colloidal suspension. The neutralization had to be achieved before flocculation occurred. The optimum pH for alum coagulation was found to vary with temperature. However, coagulation with cationic polyelectrolytes was found to be temperature independent of the flocculation rate, optimum pH and coagulant dose.

Stumm (1967) demonstrated that metals acted as Lewis acids and had a tendency to stabilize pH. He described metal ion hydrolysis as a function of pH and metal ion concentration. Stumm (1967) stated that multivalent hydrous oxides were amphoteric and that H+ and OHwere primarily the potential determining ions for such hydrous oxide precipitates. He also stated that metal ions precipitated in the presence of coordinating anions usually as nonstoichiometric mixed precipitates. Stumm et al. (1967) formed polysilicates and classified them into three separate areas: 1) insoluble, 2) stable polymers and 3) the mononuclear wall. They concluded from potential measurements that the interaction between the anionic polymeric phase and the negative sol was due to specific sorption and would overcome electrostatic repulsion. They found optimal destabilization occurred when a fraction of the colloidal surface area was covered and suggested that the mechanism of destabilization for activated silica was the same as for polyelectrolytes.





21




Stumm (1967) published a hydrolysis model for colloidal destabilization that accounted for bridging and electrostatic effects. He postulated that a fraction of the total colloidal surface area must be covered to produce coagulation. He expressed the model mathematically using a Langmiur isotherm by equating the amount of coagulant necessary to produce a certain fractional coverage to the sum of the residual and sorbed coagulant. The fractional surface coverage necessary to destabilize colloidal sols could only be gained from the residual coagulant or the sorbed coagulant. He showed from his model that the required coagulant dosage to produce destabilization could be independent of surface concentration or linearly dependent on surface concentration. In the Stumm model metals first destabilized colloids due to sorption of the hydrolyzed cationic coagulants and restabilized the colloids due to extensive sorption of the hydrolyzed metal coagulants. Finally a precipitation of the metal occurred that destabilized colloids. If the coagulant became attached to the colloidal surface, the coagualnt dose decreased with increasing colloid concentration. In the precipitation zone the coagulant enmeshed the colloids in a sweep floc. If this occurred, the coagulant dose was not a function of the colloidal surface area. If the colloidal concentration was high, the amount of coagulant dosed could be such that initial destabilization by sorption and final destabilization by precipitation would be indistinguishable.

The Stumm model for a large colloidal surface area predicted a large nonstoichiometric coagulant dose that could be reduced if





22




buffering were removed. A system with a medium colloidal surface area required a stoichiometric coagulant dose, and if buffering were present, the zone of coagulation was reduced. If low colloidal surface area were present, a large nonstoichiometric dose would be required to coagulate by precipitation. Stoichiometry could be achieved through alkalinity additions.

Kawamura (1973) reported that Ca(OH)2 .additions should be made after or during alum coagulation for optimum turbidity and color removal. Jeffcoat and Singley (1975) found that Ca(OH)2 addition prior to alum coagulation increased turbidity removal and recommended doing so for optimum coagulation results.

Hannah et al. (1967) measured alum floc size variations with kaolin, polymer and polyphosphate additions. They found kaolin and polymers increased floc size. Polyphosphates hindered floc formation. They recommended that the polyphosphates should be added last in the coagulation process. Hannah.et al. (1967) demonstrated that the order of chemical addition affected the coagulation process.

Olson and Twardrowski (1975) studied the products formed by coagulating high alkalinity waters with ferric hydrolysis and concluded FeCO3(s) may be precipitated instead of Fe(OH)3(s). Guilledge and O'Conner (1973) found arsenic was removed by both alum and ferric chloride hydrolysis. They concluded that adsorption was the removal mechanism. Their results indicated that arsenic was removed better by alum than iron coagulation, the removal was pH dependent and could possibly be the result of a chemical reaction. Stumm and Morgan (1962) found, when doing alkalimetric titrations, that the amount of




23





base required to titrate the aluminum mixture was not increased stoichiometrically in the presence of a pyrophosphoric acid. Their data suggested that phosphate removal by alum coagulation resulted from a chemical reaction. Cornwell (1975), studying alum recovery through liquid-liquid extraction, suggested that phosphate was removed by a chemical reaction producing an aluminum hydroxy phosphate.



2-5 Color Removal by Coagulation

Black et al. (1963) demonstrated that celor present in six different natural waters was removed stoichiometrically by ferric sulfate coagulation. A graph of raw water color verses required coagulant dose was constructed, and the optimum conditions for color removal did not produce a floc that had zero zeta potential. Singley et al. (1967) also found that, to obtain maximum color removal, alkalinity had to be added before coagulation. Ferric sulfate proved to be a better color removing coagulant than alum for the six natural waters tested.

Packham (1965) studied coagulation of organic color that was isolated from river water. He separated the color into humic and fulvic fractions. The mechanisms of alum and ferric coagulation were found to be similar, because stoichiometric amounts of these coagulants were required to remove different concentrations of humic and fulvic acids. Packham (1965) proposed from his data that humic acid was entering into a chemical reaction with aluminum. He determined the empirical formula for such a reaction was AI(OH)2.5R. He found that the fulvic portions were more complex than the humic acid portions





24




and found little evidence of color enmeshment in the A(0OH)3 floc. Packham (1965) did achieve an optimum pH for color removal. Jobin and Ghosh (1972) studied the oxidation of ferrous iron. They found that the addition of humic acid complexed the ferrous iron and retarded the oxidation reaction. Schnitzer (1971) found, at pH 2.5, that insoluble fulvic acid precipitates were formed with aluminum only when more than one metal ion was added for each carboxylic group present. Mangravite et al. (1975) conducted experiments on humic acid removal by alum coagulation. They demonstrated that insoluble aluminum humic precipitation formed at a pH lower than did pure Al(OH)3(s) precipitates. They suggested that color was removed from solution in alum coagulation by a chemical reaction.

Narkis and Rebhum (1975) concluded that the salts of humic and.

fulvic acids acted as anionic polyelectrolytes that reacted chemically with the cationic flocculant, the carboxylate and the phenolate groups. The reaction products formed a colloidal precipitate that could be removed by settling after flocculation. The first step in humic and fulvic acid coagulation was suggested to be a chemical reaction before flocculation by cationic polyelectrolyte addition occurred.

Luner and Dence (1970) determined that the color bodies present in a kraft waste were mostly aromatic and quinoid nuclei with carboxyl or ethylenic groups. The color bodies removed in Ca(OH)2 treatment were carboxylic, phenolic or enolic groups that had precipitated in a chemical reaction with calcium. They found that both the precipitated fractions and the nonprecipitated fractions of kraft waste were acidic, but that the nonprecipitated fractions were more acidic and





25




had a lower average molecular weight than the precipitated fractions. Luner et al. (1970) found, with massive lime treatment, that enolic groups reacted chemically with calcium to produce insoluble precipitates.



2-6 Magnesium Coagulation

Stumm (1968) demonstrated that metal cations such as magnesium, aluminum or calcium could function effectively as coagulants. Magnesium hydrolyzes significantly at pH values encountered in lime softening and produces a voluminous floc which hinders solid handling operations. Eidsness and Black (1957) reduced the volume of sludge produced in water softening operations at Dayton, Ohio and Gainesville, Florida by bubbling CO2 into the sludge to dissolve Mg(OH)2. Sixty per cent of the Mg(OH)2 was solubilized, but no attempt at optimization of magnesium recovery was made. The sludge settled more readily after carbonation. Eidsness and Black (1957) concluded that because Mg(OH)2 existed as a gelatinous coordination complex it could accept a proton more readily than the lyophobic crystals of CaCO3. This enabled the sludge volume to be reduced. Black (1971) suggested that not all the Mg(OH)2.could be removed from lime softening sludge because of MgCO33H20 precipitation ii the carbonation tank. He proposed the use of a floatation process to remove clay from the lime softening sludge before recalcination.

Thompson et al. (1972a) treated potable water samples in the laboratory using magnesium carbonate as the coagulant. They were able to develop an equation relating coagulant dosage to 'raw water





26



color and turbidity when Ca(OH)2 was used to contrbl pH. Thompson et al. (1972a) proposed a potable water treatment process in which magnesium was used as the primary coagulant, and Ca(OH)2 was used to control pH. The magnesium was recovered from the sludge by carbonation, and the Ca(OH)2 was recovered from the remaining CaCO3 by recalcination. They proposed that sludge handling problems associated with conventional coagulation plants would be greatly reduced utilizing the magnesium carbonate process. Thompson et al. (1972b) compared conventional coagulation systems with the proposed magnesium carbonate system. They also demonstrated that as turbidity and color were reduced, the zeta potential of the residual floc was increased.

Dubose et al. (1973) successfully extended the magnesium carbonate process to treatment of domestic sewage in a pilot plant at Gainesville, Florida. He found that magnesium coagulation reduced the total phosphorous to less than 0.1 mg/l P, and significantly reduced the suspended solids, color and oxygen demand of the domestic wastewater.

Black (1974), in pilot plant studies at Melbourne, Florida, found evidence that the color was released from magnesium sludge upon carbonation. This color release was found to stabilize with time, which implied that color release in magnesium recovery may not be a problem. Studies by Taflin et al. (1975), using CO2 gas to redissolve magnesium solids in a lime softening sludge, were discontinued due to a high color return with the recovered magnesium. The potable water produced by using the recovered magnesium as a





27




coagulant was too colored to be acceptable.

Predali and Cases (1973) investigated zeta potential of magnesium carbonates in electrolytes and found OH" and H+ to be the potential determining ions for Mg(OH)2. The Mg(OH)2(s) colloids had a zero zeta potential at the same pH for varying ionic strength aqueous solutions. They concluded from kinetic considerations that MgOH+ must have been the source of the positive charge on the Mg(OH)2(s) colloid. Zoganathan and Maier (1975) found that sand and kaolinite colloids in a solution of 0.005 M MgC12 had a positive zeta potential for a pH of 10.3 or greater. They attributed the positive zeta potential to the increasing percentage of MgOH+ relative to the total species of soluble magnesium.



2-7 Color in Pulp Mill Effluents Fitzgerald, Clemens and Riley (1970) demonstrated that while polymers could destabilize colloids in pulp waste, they would not neutralize the electrical double layer. Zettlemeyer, Micale and Dole (1968) studied sludges from pulp mills and found that flocculation kinetics varied with pH for organic carbohydrate base sludge, but not with an inorganic primary sludge from a newsprint mill. Their data indicated that most of the colloidal bound water was interstital and was not chemically held due to the solid-liquid interface. The National Council for Air and Stream Improvement (1971) studied surface properties of hydrogels resulting from treatment of pulp mill waste and found anionic polymers destabilized negative colloids. They suggested that the polymer sorption onto the negative surface was nonstoichiometric.





28



Davis (1972), using Ca(OH)2 coagulation at Riceboro, Georgia

to remove color from a kraft waste, found calcium solubility decreased as the sodium concentration from the digestion operation was increased. Davis demonstrated that organic carbon, color and calcium concentrations after Ca(OH)2 treatment were related to the initial sodium concentration of the waste. Berger (1964) found that a large Ca(OH)2 dose (15,000-25,000 mg/l) produced a very settleable floc that removed 90% of original color from caustic bleach effluent. The Domitar Limited Research Center (1974) reported that Ca(OH)2 coagulation was not effective for removing color from sulfite liquors. The Interstate Paper Corporation in Riceboro, Georgia used a smaller chemical dose of Ca(OH)2 (1500-2500 mg/l) to remove in excess of 90% of the initial color in a kraft waste. However, the lime dose did exceed the solubility product of Ca(OH)2 and formed a precipitate. Othof and Eckenfelder (1974) studied color removal from three kraft mill effluents by separate coagulation with Ca(OH)2, ferric sulfate and alum. They suggested that ferric sulfate was the better coagulant because of lower coagulant dose and less voluminous sludge volumes. Gould (1973) reported that the effluent from the caustic extract stage of a kraft bleach plant, when treated with Ca(OH)2, would form a metal organo precipitate that removed 90% of the initial color. Approximately 80% of the Ca(OH)2 was recovered in the sludge. Spruill (1973) found Ca(OH)2 treatment was very effective for reducing color in kraft wastes, but was ineffective for removing color from sulfite waste. Leszczynski (1972) concluded that of the many processes proposed for color removal from kraft wastes, only Ca(OH)2





29




precipitation was feasible. Kabeya et al. (1972)'found the rate of absorption of kraft mill lignins on activated carbon to be very low. Katoh and Kimura (1972) found fly ash to be almost as effective as activated carbon in sorping lignin from kraft mill effluents.

The National Council for Air and Stream Improvement (1974)

studied the. mechanism of color removal on activated carbon and found most color bodies existed in the high molecular weight (15,000) range. TOC and color were not removed in equal proportions. They concluded that color removal by activated carbon was not a chemical process, but was due to sorption. Swanson et al. (1973) did a detailed study on Ca(OH)2 treatment of kraft waste and found 86% color reduction, 57% TOC reduction and 17% sugar.reduction. There was no removal of material with molecular weights less than 400. Material with molecular weights greater than 5000 was completely removed, and partial removal was observed for material with molecular weights ranging from 400 to 5000. Swanson (1973) suggested color bodies were aromatic groups that carried a negative charge. Tejera and Davis (1970) used alum, AlCl3 and FeC13 as coagulants in color removal studies on caustic extraction waste and chlorinated waste. They determined both FeCl3 and A1Cl3 were capable of removing 96% of the color from a kraft mill caustic extraction waste, but both coagulants were hampered in the removal of color from the chlorinated waste.

Collins et al. (1969) separately concentrated chlorinated and alkaline extraction bleach effluents from a sulfite and a kraft process by reverse osmosis. Lignosulfonic acids with molecular weights in excess of 10,000 were found in the sulfite waste liquor. Jensen





30



et al. (1964) fractionated spent sulfite waste liquors by gel filtration and ion exclusion and found six different components. Saccahrides and weak organic acids at pH 4 were present in the lower molecular weight range. Fractions above a molecular weight of 40,000 were aromatic lignosulfonic acids and were responsible for most of the color in the waste.

Smith and Christman (1969) treated kraft and sulfite waste with A1C13 and FeCl3 coagulants and found either coagulant would remove 90% of the initial color in the kraft waste. Treatment of the sulfite waste by FeC13 reduced the organic carbon 50%, but increased the color of the sulfite waste. Alum reduced the color of the sulfite waste 67%. Smith and Christman (1969) proposed that the kraft waste had sulhydryl groups on lignin chains and that these groups formed insoluble sulfides during coagulation. The sulfite waste had sulfonate groups in the lignin chain which acted as strong acids and formed hydrolysis products. The mechanism for color removal in the kraft waste was suggested to be a chemical reaction, whereas the mechanism in the sulfite waste was suggested to be sorption on Al(OH)3 surfaces.

Rapson et al. (1971) used seawater as a source of soluble magnesium along with Ca(OH)2 to remove color from a kraft waste. Increased color removal was accompanied by the formation of a floc with a larger surface area than the original Ca(OH)2 floc. A 20% seawater mixture did not remove any more color from the kraft waste than did a 10% seawater mixture. Less color removal was observed when a cardboard effluent was treated, which indicated that different mechanisms





31



might have been responsible for color removal for different kraft effluents. 'The Canadian Pollution Abatement Research Program (1974) used Ca(0H)2 and MgCl2 to treat sulfite waste for color removal. They obtained an 86% reduction of NSSC waste using Ca(OH)2 and MgC12 and a 65% reduction of color using MgC12 without Ca(OH)2. They were able to remove 86% of the color from a bleach-kraft, unbleached kraft, combined bio-kraft, NSSC-NH3 base and a bio-NSSC waste. They did not optimize pH or coagulant dose in the coagulation process.
















CHAPTER 3



LABORATORY PROCEDURES



3-1 Feed Solutions



3-1.1 Synthetic Waste Solutions

All wastes were made by diluting a concentrated color source

with tap water to the desired color concentration. The color source was a stored semichemical neutral sulfite liquor which was taken from a NSSC plant digestor after the cooking operation had been completed. This liquor contained the dissolved constituents of the wood. It was referred to as "sulfite waste liquor," which can be the major source of color in the waste stream of a neutral sulfite semichemical pulp plant. The sulfite waste liquor was obtained from plants located in Harriman, Tennessee and Hartsville, South Carolina, owned by Mead Corporation and Sunoco Products Company respectively. The Mead Corporation supplied soldium base spent sulfite waste liquors and ammonium base spent sulfite waste liquors that were used as a waste source. The Sunoco Products Company supplied a sodium base spent sulfite waste liquor which was also used as a waste source.

There are different processes and many different types of hardwood trees used in neutral sulfite semichemical pulping. Because of this, it was decided at the beginning of this research to determine


32





33



if color could be removed from different semichemical neutral sulfite wastes by magnesium coagulation, but only to use the sodium base waste from the Hartsville, South Carolina plant to study mechanism of color removal by magnesium coagulation.

The color of the stored NSSC spent liquor varied from 250,000 to 500,000 Pt-Co color units. Consequently, to achieve a working color of 5000 Pt-Co color units, a dilution ratio of 50/1 to 100/1 was required.



3-1.2 Coagulation Chemicals

Magnesium sulfate, MgSO4-7H20, was used as the source of magnesium ions for the color removal process. A stock magnesium solution of 50 mg/ml as Mg+ was made in order to minimize the volume of coagulant feed dosed in the process. This was achieved by dissolving 532.6 grams of MgSO4'7H20 in a liter of distilled-deionized water.

Calcium hydroxide and sodium hydroxide were used for pH adjustment during the coagulation reaction. Calcium hydroxide was slurried in a small beaker before it was used, whereas sodium hydroxide was added from previously prepared 10 N and 1 N solutions. When necessary, sulfuric acid and hydrochloric acid were used to adjust the pH downward.



3-1.3 Polymers

Cationic, anionic and nonionic polymers were prepared from

commercial liquids and powders supplied by American Cyanamid Company. The polymers were made from polyacrylamide and amine bases. Stock





34




solutions of 2000 to 3000 mg/l were prepared from the solid based polymers by choosing a weighed amount and dissolving it in an aqueous solution by magnetically stirring it overnight. The liquid based polymers at the same concentrations required only one hour of stirring for stock preparations. For the colloidal. acids, an activator supplied by American Cyanamid (N-478) was required for stock preparation.



3-2 Analytical Equipment and Techniques



3-2.1 Total Carbon Measurements

Total carbon measurements were determined on a Beckman Model 915 Total Carbon Analyzer in conjunction with a Beckman Model 865 Nondispersive Infrared Analyzer. A three microliter sample was used for analysis. The readout was registered on a 0 to 100 scale and was compared to a standard curve. The carbon standards were prepared from potassium biphthalate for organic carbon, and sodium carbonate or sodium bicarbonate for inorganic carbon.



3-2.2 Color Measurement

All color measurements were determined according to NCASI Technical Bulletin 253. This procedure requires all samples for color measurement to be filtered through a 0.80 micron Millipore filter. The pH of the sample was then regulated to 7.6 before the amount of absorbance at a wavelength of 465 millimicrons was recorded. The sample color was then calculated by locating the sample absorbance





35




on a standard curve relating color to absorbance. If the sample had too great a color to be directly measured, the sample was diluted after filtration.

A standard curve was prepared by dissolving 1.246 grams potassium chloroplatinate, K2PtCl6 (equivalent to 0.500 g metallic platinum) and one gram crystallized cobaltous chloride, CoC12*6H20 (equivalent to 0.25 grams metallic cobalt) in distilled water with 100 ml concentrated HC1. This solution was diluted to 1 liter with distilled water. This stock solution was defined as having a standard color of 500 Pt-Co units. A standard curve was prepared and is shown in Figure 3.1. This curve fits the equation:

Color = (2183.4)(absorbance) 4.4 (3-1)



3-2.3 Incineration

Sludge incineration was determined in a Thermodyne furnace, Model F-A1730. Sludge samples were dried at 1030C and filtered through a Buchner funnel on a Whatman no. 40 ashless filter before incineration in the furnace. Incineration temperatures were varied from 1800C to 8500C. Times of incineration were varied from 15 minutes to 120 minutes.



3-2.4 Jar Tests

Jar tests were performed on a Florida Jar Test Machine. Chemicals were dosed simultaneously to four 1 liter beakers. Rapid mixing took place at 100 rpm for three minutes. The Florida Jar Tester was capable of 145 rpm, but due to the heavy floc formed when treating




36


















0.2









0














Color Fig. 3.1 Standard Pt-Co color curve




37




the highly colored waste, a stirring rate of 100 rpm was the maximum that could be attained. Beyond 100 rpm, the magnetic couple between the jar stirrers and the machine was broken by the stirrer over-turning. Slow mix took place at 35 rpm for 15 minutes. The pH was adjusted through both the slow mix and the rapid mix cycles to maintain a constant pH during the coagulation reaction. Floc was allowed to settle for 30 minutes before samples were taken for analysis. If the coagulating mixture had not developed a clear supernatant, a sample was taken and filtered through a no. 40 Whatman filter in order to simplify the required filtering step through the

0.80 micron Millipore filter before color measurement.

The G levels of the rapid mix and the slow mix cycles were 110 sec-1 and 30 sec-1 respectively. The mixing level in the flocculation stage had a Gt valhe of 27,000, which was approximately the low end of the range specified in Waste Treatment Plant Design (1971).



3-2.5 Metal Analysis

Metal analyses were determined on a Varian Techtron Model 1200 Atomic Absorption Spectrophotometer. Magnesium measurements were made at a wavelength of 202.5 nanometers. Calcium measurements were made at a wavelength of 422.7 nanometers. All samples were filtered through a 0.80 micron Millipore filter and treated with 1 ml of 17% lanthium-HC1 solution per 10 ml of sample before calcium and magnesium values were measured.



3-2.6 Mobility Measurements

All mobility measurements were made with a Zeta Meter. The Zeta





38



Meter was used in conjunction with a Riddick cell and a stereoscopic microscope. The multiscale 15X ocular micrometer was used in the Zeta Meter for mobility measurements. A platinum-iridium cathode coupled with a molybdenum anode were used in the Riddick cell.



3-2.7 pH Measurements

All pH measurements were made on a Corning Model 12 expanded scale pH meter. A Corning silver-silver chloride reference electrode in conjunction with a glass electrode were used for all pH measurements.



3-2.8 Settling Tests

All settling- tests were conducted in a standard 1000 ml graduated- cylinder. One liter of waste was coagulated in a jar and immediately transferred to a graduated cylinder where the height'of the sludge-supernatant interface was recorded. The following formula was used to calculate the Sludge Volume Index, SVI:

SVI = ml settled sludge x 1,000
mg/l suspended solids



3-2.9 Solids Analysis

All suspended solids analyses were determined on samples that

were filtered through a no. 40 Whatman filter and dried at 1030C for one hour. Nonvolatile and volatile solids were by filtering the samples through a no. 40 ashless Whatman filter, drying at 1030C for one hour, and recording the weight. The sample was then ignited at 5500C for 60 minutes, after which it was weighed to determine




39




nonvolatile solids. Samples were weighed immediately after drying at i030C, but were cooled for one hour in a dessicator after igniting at 550oC before weighing.



3-2.10 Titration Curves

0.01 and 0.1 N H2SO4 and NaOH were used for determining the acid-base strength of the samples. The volume of waste titrated varied from 50 to 200 ml. One minute was allowed for pH stabilization each time the titrant was added to the sample. A Teflon covered magnetic bar in conjunction with a magnetic stirrer were used to mix the solution during titration.



3-3 Experimentation



3-3.1 Coagulation Experiments

The color removal experiments began by mixing the waste to the desired color concentration and measuring the color as previously described. The next step was regulation of the waste solution pH with Ca(OH)2 or NaOH. The coagulant was then added and the reaction pH was adjusted. Samples for analysis were taken after 30 minutes of settling. Organic carbon measurements and color values were determined immediately after coagulation. The samples to be analyzed for metal concentration were acidified immediately following coagulation.



3-3.2 Coagulant Recovery

Several methods were used to recover the magnesium coagulant.




40




Following coagulation the resulting sludge was filtered through a Buchner funnel and was then dried at 1030C for one hour. The dry sludge was ignited at 5500C, and the resulting nonvolatile solids were placed in contact with a 10% CO2 gaseous stream or stabilized with H2S04 to recover the oxidized magnesium. Two 40 liter volumes of waste were treated in order to produce a large quantity of sludge. These wet solids were then heated until they achieved a constant weight at 1030C. The dried solids were ignited at 5500C to remove the coagulated color.



3-3.3 Coagulant Recycle

The magnesium was recycled to determine the effectiveness of reusing the same magnesium as the primary coagulation in the color removal process. The niethod of recycling the magnesium consisted of incinerating the sludge produced in the coagulation reaction at 5500C. The resulting nonvolatile solids were carbonated for 45 minutes with a 10% C02-90% air gaseous mixture. The recovered magnesium was recycled with and without the nonvolatile solids that were not dissolved during carbonation. The two different techniques of recycling the magnesium determined the effectiveness of the remaining nonvolatile solids in the color removal process.










CHAPTER 4



RESULTS



4-1 Determination of Coagulation pH and Coagulant Dose

Development of the color removal using magnesium coagulation required that the pH control agent, coagulation pH and coagulant dose be determined. The chemicals selected for pH control were Ca(OH)2 and NaOH because they were inexpensive and commercially available. The coagulation pH and- coagulant dose were defined as the minimum pH and dose that resulted in a 90% reduction of the initial color.

A three step technique was used to determine the coagulation

pH and coagulant dose for the color removal process. First the coagulation pH was found by determining the reaction pH where maximum color removal occurred for a constant magnesium dose. The second step was to determine the magnesium dose that removed 90% of the color at the coagulation pH. Finally the stability of the coagulation pH was verified by repeating the first step for a magnesium dose other than the coagulation dose. If the coagulation pH did not shift, then the coagulation pH and coagulant dose were acceptable. If the shift in the coagulation pH occurred, then the coagulation pH had to be determined as a function of both the coagulant dose and the initial color.



4-1..1 Coagulation pH

Figures 4.1 and 4.2 show the curves from which the coagulation pH can be determined using Ca(OH)2 or NaOH for a NSSC waste with an 41




42





4000
Color = 2500 0o Mg = 150 mg/I- Ca(OH)2 3000 Mg = 750 mg/I- NaOH

3

S2000



!000


O2 I 00

9.5 10.0 10.5 11.0 11.5 12.0 12.5 pH

Fig. 4.1 Color residual as a function of final pH



8000
Color = 5000 Mg = 300 mg/I Ca(OH)2 6000 Mg = 400 mg/I- NaOH 6000



4000







0oo
0OO0

9.5 10.0 10.5 11.0 11.5 12.0 12.5 pH

Fig. 4.2 Color residual as a function of final pH





43




initial color of 2500 or 5000. The constant magnesium dose used for each curve is indicated on the figures. The minimum pH at which 90% color removal was achieved was 10.6. The final color was dependent of the pH control agent. At pH 10 when Ca(OH)2 was used, the final color of the waste was increased approximately 30% more than the initial color. This did not occur with NaOH. The same degree of color removal was obtained from pH 10.6 to pH 11.4 using Ca(OH)2 or NaOH. The degree of color removal decreases past 11.4 when NaOH was used. This did not occur with Ca(0OH)2'

These figures show that a 90% reduction of the original color

of the waste was first reached at pH 10.6 for each color. This indicated that the coagulation pH was independent of the variability in the waste color. From the data presented in Figures 4.1 and 4,2 it was concluded the coagulation pH was 10.6.


4-1.2 Coagulant Dose

A NSSC waste with a color of 2500 or 5000 was coagulated at pH 10.6 with a varying magnesium dose. These data are presented in Figures 4.3 and 4.4. The pH control agents were Ca(OH)2 and NaOH. Magnesium coagulation of the NSSC waste with either pH control agent was able to remove 90% of the color. The required magnesium dose for the waste with a color of 2500 was 100 mg/l when Ca(OH)2 was used for pH control. The magnesium dose was 200 mg/l when NaOH was used for pH control. The required magnesium dose for the waste with a color of 5000 was 200 mg/l when Ca(OH)2 was used for pH control, and was 400 mg/l when Na0H was used for pH control. The required coagulant dose was directly proportional to the color of the NSSC waste.





44




The same color reduction was achieved when Ca(OH)2 or NaOH was used for pH control. As is indicated in either Figure 4.3 or 4.4, a larger coagulant dose was required to achieve 90% color reduction when NaOH was used to control pH. The coagulant dose was approximately 50% less when Ca(OH)2 was used to control pH.



4-1.3 Variation of Coagulation pH with Coagulant Dose

All NSSC waste solutions were made by diluting a stored NSSC liquor with tap water. A waste color of 5000owas tested to determine if the coagulation pH was dependent on coagulant dose. Figure

4.5 shows for a color of 5000, the optimum color removal again occurred at pH 10.6. The magnesium dose was 400 m-/l and 'Ca(OH)2 was used to control pH. This was the same pH at which maximum color removal occurred at a color of 5000 for a magnesium dose of 300 mg/l in conjunction with Ca(OH)2. It was therefore concluded that the coagulation pH did not vary with coagulant dose. The only significant difference between the curves in Figures 4.3 and 4.4 was the presence of the calcium ion when Ca(OH)2 was used for pH control. The stored NSSC liquor was made in a digestion process which used NaOH and contained a large concentration of sodium. The additional increase in sodium concentration was therefore not significant when NaOH was used to adjust pH. The sodium concentration of wastes with colors of 2500 and 5000 was 40,000 mg/l and 80,000 mg/l respectively. The amount of sodium increase when NaOH was added to adjust pH was always less than 1000 mg/l, or less than 2.5%. Monovalent ions, such as. sodium, generally do not complex organic compounds to the same extent as divalent ions.





45




Color = 2500 1000 pH = 10.6



750

0


6 500 NaOH

50


250 Ca(OH)2


II I I I
0 100 200 300 400 500 600 Mg mg/I Fig. 4.3 Comparing NaOH and Ca(OH)2 for color removal via magnesium coagulation Color = 5000 .200C pH = 10.6



S1500

o

, 1000 NaH



500 Ca(OH)2
0



0 IU0 200 300 400 500 600 Mg mg/I
Fig. 4.4 Comparing NaOH and Ca(OH)2 for color removal via magnesium coagulation




46



















Color = 5000 4000 Mg 400 ml Ca(OH)2 3000



2000 "

1000




9.5 10.0 10.5 11.0 11.5 12.0 12.5 pH

Fig. 4..5 Verification of coagulation pH





47




The additional sodium added to the waste to adjust pH was not a significant increase in sodium concentration, and did not extensively form any complexes.

The total calcium concentration in the untreated waste was

approximately 40 mg/l as Ca. When Ca(OH)2 was used to adjust pH the calcium concentration in the waste was increased to 600 mg/l. This was significant because the calcium concentration increased and probably did extensively complex the organic compounds.

Figures 4.3 and 4.4 show that the magnesium required to remove 90% of the color.was reduced when Ca(OH)2 was used for pH control rather than NaOH. Both calcium and magnesium are divalent ions and will form common complexes with the organic compounds in pulp wastes. When Ca(OH)2 was added to control pH, Ca+ complexed many organics that Mg++ would have normally complexed in the absence of the added Ca++. Therefore Ca(OH)2 reduced the required coagulant dose.

A magnesium dose of 100 mg/l removed 90% of the color from a

NSSC waste with a color of 2500 when Ca(OH)2 was used for pH control. A magnesium dose of 200 mg/1 was necessary for 90% color removal when NaOH was used to control pH. The complexing ability of the calcium ion was responsible for a 50% reduction in the coagulant dose. Since NaOH is more expensive than Ca(OH)2 and does not reduce the coagulant dose, Ca(OH)2 was chosen as the pH control agent.



4-1.4 Magnesium Remaining in Solution as a Function of Final pH

The magnesium remaining in the treated waste after color removal as a function of final pH is presented in Figure 4.6. For both




48









300

Ca(OH)2

Color = 2500 250
o Color = 5000




200




* 150





100





50





0 10.0 10.5 11.0 11.5 12.0 pHf Fig. 4.6 Magnesium remaining in solution as a function of final pH




49




colors tested, 35-40 mg/l of magnesium remained in solution when the final pH was 10.6. The amount of magnesium remaining in solution was reduced to 4-10 mg/l when the final pH was increased to 11.0. Magnesium in solution remained approximately constant past 11. Coagulation at pH 11 saved approximately 30 mg/l of magnesium from being wasted in the treated effluent. The Ca(OH)2 dose was increased approximately 50 mg/l to raise the coagulation pH to 11. The magnesium saved was worth more than the Ca(OH)2 used to increase the pH. Therefore, it was decided to increase the coagulation pH to 11.

Increasing the coagulation pH to 11 accomplished two things. First, enough magnesium was recovered to make the coagulation process less expensive. Second, the allowable fluctuation in coagulation pH was increased. The per cent color removed was significantly less at any pH less than 10.6. But when the coagulation pH was 11, a reduction of 0.4 pH units would not significantly affect color removal.



4-1.5 Magnesium and Ca(OH)2 Dose as a Function of Initial Waste Color

The waste effluent from a semichemical neutral sulfite plant

varies in color intensity. Because of this variability, the amount of Ca(OH)2 and magnesium to remove 90% of the initial NSSC color was determined as a function of the initial color of the waste. The Ca(OH)2 required is presented in Figure 4.7. The magnesium requirement is presented in Figure 4.8. Both the Ca(OH)2 and the magnesium requirements. were directly dependent on the initial color of the waste. This suggested a stoichiometric relationship between color and coagulant dose.





50






10 mgl/I maximum Mg residual

Lime mg/l= 750 t 0.10 (color) 7500



- 5000



2500




0 500 1000 1500 2000 Ca(OH)2 mg/I Fig. 4.7 Lime dose as a function of initial waste color for magnesium coagulation 7500
90% minimum color removal
Mg mg/I 0.060(color)


5000
0
C-)


O0
o 2500






0 100 200 300 400 Mg" mg/I Fig. 4.8 Magnesium dose as a function of initial waste color using lime









4-2 Waste Characteristics



4-2.1 Untreated Waste Titration Curves

The acid strength of the untreated NSSC waste was determined

by titrating 50 ml samples with varying colors with 1.0 N H2SO4. The acid strength of the waste was defined as the milliequivalents of acid required to change the waste pH from 12 to 2. The results of these titrations are presented in Figures 4.9 through 4.13 for NSSC waste colors of 2,500 to 40,000. The NSSC waste was obtained from the Sunoco Products Corporation in Hartsville, South Carolina and is denoted as sodium base Sunoco NSSC waste in Figures 4.9 through 4.13. As the color of the NSSC waste was increased, the acid strength of the NSSC waste also increased. This indicated that color was acidic, and an increase in color would increase the acidity of the waste.

The titration curves indicated the acidity of the waste was

gained from two functional groups or mixtures of functional groups. These functional groups had pK values in the range of carboxylic acids and phenols or enols. The equilibrium constants for the NSSC wastes were approximated graphically and are presented in Table 4-1. The pK values were identified by locating the inflection points on the titration curves. Approximately 66% of the data points are not represented in Figures 4.9 through 4.13 in order that the titration curves would be uncluttered and clear. These points occurred in two areas, both of which were identified by slight humps on the titration curves. These pK values are approximately 4.6 and 9.8. They differed by four orders of magnitude, which was a large enough separation to allow graphical approximation of pK values.





52















14 Waste = Sodium base sunoco NSSC
V0 = 50 ml
12 Color 2500


10



pH
6


4 2 0
10 8 6 4 2 0 2 4 6 8 10
meq acid X 10 meq base X 10



Fig. 4.9 Titration curve of sodium base sunoco NSSC waste with color equal to 2500





53















14 Waste= Sodium base
V0 = 50 ml

12 Color= 5000


10 8
pH
6


4 2

0 I1111

10 8 6 4 2 0 2 4 6 8 10
meq acid X 10 meq base X 10



Fig. 4.10 Titration curve of sodium base sunoco NSSC waste
with color equal to 5000





54
















14 Waste = Sodium base sunoco NSSC
Vo = 50 ml

12 Color = 10,000


10 8
pH

6


4


2


0
10 8 6 4 2 0 2 4 6 8 10 meq acid meq base



Fig. 4.11 Titration curve of sodium base sunoco NSSC waste with color equal to 10,000





55















14 Waste = Sodium base sunoco NSSC
V0 = 50 ml
12 Color = 20,000


10


8
pH

6


4 2



10 8 6 4 2 0 2 4 6 8 10 meq acid meq base



Fig. 4.12 Titration curve of sodium base sunoco NSSC waste with color equal to 20,000





56
















14 Waste= Sodium base sunoco NSSC
Vo = 50 ml
12 Color= 40,000


10 8
pH
6


4 2



10 8 6 4 2 0 2 4 6 8 10 meq acid meq base




Fig. 4.13 Titration curve of sodium base sunoco NSSC waste with color equal to 40,000




57













-TABLE 4-1

GRAPHIC DETERMINATION OF pKa OF
SODIUM BASE NSSC WASTE



Color pK2 pK1 40,000 9.6 4.6 20,000 9.7 4.5 10,000 10.0 4.7 5,000 10.1 4.7 2,500 9.5 4.5


Average 9.7 4.6





58




4-2.2 Comparison of Untreated and Treated Waste Titration Curves

The acid strength of the untreated NSSC waste was compared with the acid strength of the treated NSSC waste. This was done in order to determine if any reduction in acidity occurred during the color removal process. The previous titration curves of the untreated waste revealed significant acidity in the weak and very weak acid range. These acids would be ionized at pH 11 and available to participate in a chemical reaction. Magnesium as Mg++ is a Lewis acid and is capable of reacting chemically with the ionized anions from the waste acids.

The titration curves ate presented in Figure 4.14 for the same NSSC waste before and after treatment. The acid strength of the NSSC waste was reduced by the color removal process. Before treatment, 0.8 meq of base was required to titrate the waste from pH 12 to pH 9. After treatment, only 0.4 meq of base was required to produce the same change in pH. The amount of base to change the waste from pH 12 to pH 3 before and after treatment was 1.65 and 1.0 meq respectively. Very weak acids have pK values ranging from approximately 8 to 10. Weak acids have pK values of approximately 3 to 5. The total reduction in acid strength during the color removal process was 0.65 meq. Approximately 0.4 meq of this reduction occurred in the very weak acid range from pH 12 to pH 9. This was approximately 60% of the total reduction in acid strength. From the titration data, it was concluded that color reduction by magnesium coagulation does result in at least a reduction of the acids present in the NSSC waste.





59















12

I I

10

9 Untreated NSSC waste Color = 5000
8 Vol= 50 ml pH
7

6

Treated NSSC
waste
4 -Color = 485
Vol= 50 mli
3
0 5 10 15 meq X 10

Fig. 4.14 Titration curve of treated and untreated NSSC waste





60




4-2.3 Waste Content

The treated and untreated waste analyses shown in Table 4-2

were done by the United States Air Force Environmental Health Laboratory at Kelly Air Force Base in Amarillo, Texas. The untreated waste was a sodium base NSSC waste that was prepared from a stored liquor obtained from the Sunoco Products Company. Hydrochloric acid and sulfuric acid were used to adjust the pH to 7.6 before shipment.



4-3 Color Removal Mechaiism


4-3.1 Color and Magnesium Titration Curves

A volume of 200 mls of NSSC waste was dosed with 80 mg (400 mg/l) of magnesium as Mg++. This solution was titrated with a 1.0 N NaOH solution to a pH of 12.0. The titration curve for this experiment is presented in Figure 4.15. In the first portion of the curve, an inflection point was present at pH 9.6. This was approximately the second pK determined earlier from the NSSC waste titration curves. The solution was slightly buffered by the colored NSSC waste at this point in the titration. At pH values higher than 9.6, the acids in the NSSC waste were ionized.

In the second portion of the curve, another inflection point

was found at pH 10.8. Coagulation and 90% color removal occurred at all pH values greater than or equal to 10.6. At pH 10.8 in the titration curves, magnesium was acting as a buffer by hydrolyzing and precipitating out of solution. Color removal was accomplished when the buffering capacity of the colored waste and the magnesium were exceeded. The acids were ionized and were capable of an acid-base reaction




61







TABLE 4-2

UNTREATED AND TREATED NSSC WASTE ANALYSIS. TREATMENT WAS
WITH 150 mg/1 Mg++ AND Ca(OH)2 TO ADJUST pH TO 11.0.



Item Lab Analysis
(mg/1 unless noted)

Untreated Treated


1. Color Pt-Co units 2000.000 150.000 2. Turbidity JTU's 3.000 4.000 3. Chemical oxygen demand 1510.000 784.000 4. Total suspended matter 0.000' 0.000
5. Volatile and fixed
suspended matter 0.000 0.000 6. Oils and greases 0.800 0.500 7. Surfactants as mg/1 LAS 0.800 1.600 8. Chlorides 56.000 920.000 9. Flourides 1.100 0.500 10. Phosphates 0.500 0.300 11. Sulfates 200.000 1150.000 12. Cadmium .01 .02 13. Chromium (hexavalent) .01 .01 14. Chromium (total) .03 .05 15. Copper .05 .03 16.- Cyanides .01 .02 17. Iron .72 .1 18. Lead .05 .07 19. Manganese .28 .05 20. Silver .01 .02 21. Zinc .1 .05 22. Mercury .005 .005 23. Total organic carbon 530.000 350.000 24. Nitrite nitrogen .06 .02 25. Ammonia nitrogen .8 .2





62






















II Initial color= 5000 Vol = 200 mi Mg = 400 mg/I I N NoOH 10
Buffering due to magnesium

pH

9


Slight buffering due to color

8


0 2 4 6 8 10 12 14 16 meq

Fig. 4.15 Titration curve of raw waste dosed with magnesium




63




with magnesium. This reaction could involve the color in a formation of an insoluble precipitate. The removal of this precipitate would remove the color bodies.



4-3.2 Magnesium, Calcium, Color and Organic Carbon Residuals After
Coagulation

A NSSC waste with an initial.color of 2,500 was coagulated with a constant magnesium dose of 150 mg/l. The final pH of coagulation was varied from 10 to 11.5 using Ca(OII)2 to control pH. The magnesium, calcium, organic carbon and color residuals were determined after coagulation. The total organic carbon concentrations and color intensities after coagulation are presented in Figure 4.16. The magnesium and calcium concentrations remaining after coagulation are presented in Figure 4.17.

The total organic carbon concentration was reduced 40% when the NSSC waste was coagulated at pH 11.5. The residual color at this point was 157 Pt-Co color units. When the waste was coagulated at any pH from 10.6 to 11.2, approximately 34% of the total organic carbon was removed. The average residual color in this pH range was 197 Pt-Co color units. Increasing the coagulation pH to 11.5 would only remove an additional 1.8% of the initial color. Coagulation at any pH from 10.6 to 11.2 removed 92% of the initial color.

When the waste was coagulated at pH 11.5, an additional 6%

(27 mg/l) of organic carbon was removed. As was noted, the additional color reduction at pH 11.5 was 1.8%. When the waste was coagulated at any pH from 10.6 to 11.2, the organic carbon was reduced 154 mg/l




64






Initial Concentrations
400C
400---T.O.C. T.O.C.= 455 mg/I 3000 Color= 2500

300 C E 2000 o
4
S200 ---Color


1000
100




0 O 9.5 I0.0 10.5 11.0 11.5 12.0 pHf
Fig.4.16 Organic carbonand color residuals as a. function of final pH

Initial Concentrations 150 Mg= II mg/I 600 Ca = 15 mg/i

0 0


100 aCa remaining 400


E E

50-Mg remaining 2000






9.5 10.0 10.5 11.0 11.5 I2.0 PHf
Fig. 4. 17 Magnesium and calcium residuals as a function of final pH Coagulant dose = 150 mg/i Mg




65




and 2,300 Pt-Co color units were removed. In this pH range, color was reduced by 15 Pt-Co color units for every mg/l of organic carbon removed. For each additional mg/l of organic carbon removed by coagulation at pH 11.5, the color was reduced by only 1.5 Pt-Co color units. From these data it was concluded that not all the organic carbon in the waste contributed equally to the waste.color.

The residual color increased 32% when the coagulation experiment was attempted at pH 10.0. The magnesium residual curve in Figure 4.17 shows that no magnesium precipitated out of solution at pH 10. All of the magnesium was therefore available to form chelates with the NSSC waste. Calcium ions causing increases in the color of a kraft waste due to chelation were reported by Luner and Dence (1971). Color increasing chelates formed by magnesium and quinones have been reported by Day and Underwood (1967), Aromatic quinones are an integral part of basic lignin structure, and lignin is responsible for color in pulp waste. The color increase at pH 10 was probably due to the chelation of lignin building units, possibly direct chelation with quinones.

None of the 150 mg/l of magnesium was removed after coagulation at a final pH of 10.0. The total amount of magnesium available for coagulation was the magnesium dosage and the magnesium present in the waste. For the data presented in Figure 4.17, the total amount was 161 mg/l magnesium. When magnesium precipitation began, a corresponding drop in color intensity was observed, as was a corresponding drop in organic carbon. The concentrations of magnesium, organic





66




carbon, and color decreased with increasing pH, indicating that the decreases in these three parameters were related. Corresponding decreases in magnesium, color and organic carbon occurred simultaneously. A possible relationship for these simultaneous reductions could be a chemical reaction between the color producing organic compounds and the magnesium ions. This relationship would result in the chelation and precipitation of a magnesium organic color-body complex. A second possibility could be the adsorption of the chelated organics onto the voluminous magnesium hydroxide floc.* The data presented in Figures 4.16 and -4.17 could conform to either of these postulated mechanisms.

Ca(OH)2 was used as the pH control agent in these experiments.

As the Ca0OH)2 dissolved, the pH and calcium concentration increased. The increasing pH probably caused the solution to become supersaturated with respect to the magnesium-color body compound. This hypothesis is supported by the data presented in Figure 4.17. As the pH increased from 10.0 to 11.5, the calcium in solution increased from 237 to 681 mg/l as Ca++. The magnesium in solution decreased from 161 to 7 mg/l as Mg++. This increase in calcium and decrease in magnesium concentrations can be visualized as a reaction between the dissolved Ca(OH)2 and Mg++. Such a reaction is shown in Equation 4-1. The value of AGo is -12.601 kcal/mole with the reaction proceeding from left to right.

++ ++ ++
Ca + 20H + Mg = Mg(OH)2(s) + Ca (4-1)

Total organic carbon, magnesium and color residuals were





67




determined after coagulation with a varying magnesium dose. In these experiments, the pH was held constant at 10.6. The pH controls used were NaOH and Ca(OH)2. The data from these experiments are presented in Figures 4.18 and 4.19. In these figures the residual color, organic carbon, and soluble magnesium are plotted as functions of the total millimoles of magnesium available. The total millimoles of magnesium available consists of the initial magnesium plus the coagulant dose.

The initial color of the NSSC waste used to plot Figure 4.18 is

approximately half that of the initial color of the waste used to plot Figure 4.19. The scales in Figure 4.18 are one-half the scales in Figure 4.19. This was done so the two figures could be directly compared without being misleading.

In both figures, a decrease in organic carbon was accompanied by a decrease in color removal. The beginning of floc formation is identified by the dashed lines in Figures 4.18 and 4.19. This point was identified when the floc became visible to the naked eye. No magnesium was removed from solution until floc formed. Once floc formed, the color was reduced below the initial color level of the waste. Before this point, a color increase had occurred due to the chelation of magnesium and calcium with the NSSC waste. After formation of floc, the color, magnesium and organic carbon concentrations were decreased. The floc formed at a smaller magnesium dose using Ca(OH)2 compared to using NaOH for two reasons. First, the initial waste color treated with NaOH was higher. Floc was formed with Ca(OH)2 when the total amount of available magnesium was 2.19 millimoles. This point was not reached with NaOH until the total millimol-es of Mg++




68














3000' 500 Vol= I liter pH= 10.6
o = Color remaining 75 2500 --75 a = Mg remaining 375 = T.O.C.
2000 - E o E
oi 50 a- 1500 250


1000
25
125 25
500


0- 0
0 2.5 5.0 7.5 10.0 Total m mol Mg available Fig. 4.18 Color, T.O.C., and Mg" residual after Mgf coagulation using Ca(OH)2 for pH control




69













6000 1000
I Vol= Iliter 5000 p- H = 10.6 150 0o Color remaining 750
=0 Mg remaining E S4000 E E* = T.O.C. oo
S3000 E 500 -100



2000 E
250 5

1000

I
0 i 0
0 5 10i 15 20 Total m mol Mg available

Fig. 4.19 Color, T.O.C., and Mg residuals after Mg coagulation using NaOH for pH control




70




available were 9.46. If the initial waste color was the only reason for the larger magnesium requirement with NaOH, then the amount of available magnesium required to form floc would increase in proportion to the color increase. This was not the case. The initial color of the waste treated with NaOH was approximately twice that treated with Ca(OH)2. If the magnesium dose was directly proportional to the initial color, then extrapolating the Ca(OH)2 treatment dose would give 4.38 mM/l as the necessary magnesium dose for the NaOH treatment. It was found that 5.08 mM/1 more of magnesium was necessary to form floc using NaOH. This difference was due to the presence of calcium ions when Ca(OH)2 was used to control pH. Floc formation did not occur with the first magnesium additions with either pH control agent. Calcium ions from Ca(OH)2 complexed some of the organics in the waste that would have been complexed by the magnesium had Ca(OH)2 not been used. This enabled floc formation to occur at a smaller magnesium concentration.

Two different chemical reactions, chelation and precipitation, have been identified in the color removal process. First, chelation occurred between the divalent metal ions and the ligands present in the NSSC waste. The chelation demand of the NSSC waste was satisfied before color removal occurred. The magnesium chelates were partly reduced by using Ca(OH)2 to adjust pH. After the chelation demand was satisfied, floc formation and color removal occurred as shown in Figures 4.18 and 4.19.




71



4-3.3 Stoichiometry of Color Removal from NSSC Waste by Mangesium
Coagulation

A possible mechanism of color removal was adsorption of the color bodies on magnesium hydroxide floc. The colored organics would not have been involved in a chemical reaction that formed a magnesium compound, but would have become attached to the floc by Van der Waals forces or hydrogen bonding. If magnesium hydroxide floc were formed, two moles of hydroxide would be required for every mole of magnesium removed from solution. If an insoluble precipitate formed that was a chemical compound consisting of magnesium, hydroxide and organic ions, the moles of magnesium ions removed divided into the moles of hydroxide ions removed would be less than two.

The moles of magnesium removed divided into the moles of hydroxide removed is presented as a ratio in Figure 4.20. There are three different experiments represented in Figure 4.20. In the first two, Ca(OH)2 was used to control pH for color removal from NSSC wastes with initial colors of 2500 and 5000. In the third experiment, NaOH was used to control pH for color removal from an NSSC waste with an initial color of 2500.

The moles of magnesium removed were found by measuring the magnesium concentrations before and after color removal. The moles of hydroxide removed were found by difference. First the moles of hydroxide necessary to raise the pH to 10.6 were found. Then this amount was subtracted from the moles of hydroxides required to raise the pH to 10.6 after the magnesium dose was added. The difference was the hydroxide.demand of the magnesium used to coagulate the color, and was represented as OH]Mg++ in Figure 4.20. If.NaOH was




72

















u= Initial color 2500 coagulated by Mrgt using time So= Initial color 5000 coagulated by Mg11 using lime
7 = initial color 5000 coagulated by Mg using NaOH I71
Pl 6

1 5
4
L,

3

4-o


II, I
0 5 10 15 20 25 m mmol Mg available Fig. 4.20 Ratios of COH]Mgl / [g"t] OH]Mpg" is the moles of hydroxides required by the magnesium for color removal, EMg*] is the moles of Mg required for color removal.





73




used to control pH, the moles of hydroxide were measured from the direct addition of a 1.0 N NaOH solution. If Ca(OH)2 was used to control pH, the increase in calcium concentration before and after coagulation was measured. The calcium increase was doubled to determine the moles of hydroxide required in the color removal process.

The curves shown in Figure 4.20 represent magnesium to hydroxide molar ratios for the floc formed in the color removal process. The coagulation pH was 10.6. Both Ca(OH)2 and NaOH were used to control pH. All of the data points in Figure 4.20 represent some degree of color removal.

When NaOH instead of Ca(OH)2 was used to control pH, a greater magnesium concentration was required before any floc was formed. This was due to chelation and is shown by the separation of the two curves for wastes of equal initial color in Figure 4.20.

The initial points on each of the curves in Figure 4.20 represent the beginning of floc formation. As color removal and floc formation increased, the curves eventually stabilized at 1.5. At the low magnesium doses used initially this ratio was not stable because of the chelation demand of the waste. Once this demand was exceeded, the ratio stabilized at 1.5 and remained there for all subsequent magnesium doses.

A ratio of 1.5 hydroxide ions to 1.0 magnesium ion does not produce an electrically neutralized compound. Another anion had to contribute one-quarter of the total negative charge for the precipitate to be electrically neutral. Color bodies are negatively charged and color was removed as magnesium ions were precipitated. If the color




74




bodies were involved in a chemical reaction with the magnesium and hydroxide ions, then the molar ratio would be less than two. The negatively charged color bodies would electrically neutralize the precipitate. The ratio of 1.5 indicates that an insoluble precipitate was formed in the ratio of 30H-: 2Mg++: 1R" where R- represents the color body.

The formation of an insoluble precipitate is further supported by the stability of the ratio. If the magnesium-color body complex became enmeshed in a Mg(OH)2 precipitate, the overall OH/Mg ratio would be less than two. Some of the magnesium removed would be attributable to the enmeshed chelate and some to the Mg(OH)2 floc. The ratio, however would not be stable for an increasing coagulant dose. As the coagulant dose would increase, the ratio would approach two because mostly Mg(OH)2 would be formed after the chelation demand of the waste was satisfied. As shown in Figure 4.20, the molar ratio does not vary after becoming stabilized at 1.5. From the data presented, it was concluded that a chemical reaction was the mechanism by which color was removed from solution.

Calcium hydroxide has been used successfully to remove color from a kraft waste at a pulp plant in Riceboro, Georgia. Dissolved Ca(OH)2, used in magnesium coagulation, could precipitate as Ca(OH)2 or some other compound and remove NSSC color bodies. However, based on solubility product calculations, no Ca(OH)2 would precipitate at the Ca(OH)2 doses and pH required by magnesium coagulation.

An experiment was done to determine the color removal capability of Ca(OH)2 with reference to a NSSC waste. The residual color and pH





75




were determined in a NSSC waste after Ca(OH)2 addition. These results are shown in Figure 4.21. An increase in the residual color of the waste was noted for the initial doses of Ca(OH)2. This was due to chelation of calcium ions with the organic compounds in the waste. The maximum Ca(OH)2 dose was 2000 mg/l, with a resultant pH of 12 and a color reduction of approximately 45%. No. color was removed until the pH was 11.2. The pH used in magnesium coagulation was 11. At pH 11 the use of Ca(OH)2 alone slightly increased the residual color of the waste. This is shown in Figure 4.21. These results indicate that Ca(OH)2 does not remove any color in the magnesium coagulation process when the coagulation pH is 11.



4-4 Settling of Coagulated Wastes



4-4.1 Purpose of Settling Tests

The purpose of the settling tests was to minimize the volume of sludge and gain some knowledge of the factors governing the settling process. The Sludge Volume Index, SVI, was determined after each settling test on all of the sludges produced during coagulation.

The settling tests were conducted as described in Section 3-2.8 of Chapter 3. Settling tests were performed on .the coagulated wastes and on polymer treated coagulated wastes. Cationic, nonionic, and anionic polymers were used as settling aids in the tests. All of the polymers tested were supplied by the American Cyanamid Company.

Since the magnesium dose depended on the initial color of the waste, a constant concentration of suspended solids was produced in





76








14 NSSC waste 5000 o = Color
a= pH


13 4000





12 3000 pH Color


11 2000





I0 1000





9 0
0 500 1000 1500 2000 Ca(OH)2 mg/I


Fig. 4.21 Color and pH of a NSSC waste as a function of Ca(OH)2 concentration





77




all wastes with the same initial color. The use of polymers as a settling aid was found to have a negligible effect on the suspended solids produced. A coagulated waste with an initial color of 5000 was found to have a suspended solids concentration of approximately 1800 mg/l. If the initial color was reduced by half, the suspended solids produced by coagulation also were reduced by half.



4-4.2 Sludge Settleability

The type, functional group, charge and approximate molecular weight of the polymers used in the settling tests is presented in Table 4-3. The Sludge Volume Index is also presented in Table 4-3. as a function of polymer dose and polymer type.

The SVI of the raw sludge was 352. This was increased to approximately 550 when a cationic polymer was used as a settling aid. The floc was still completely suspended after 30 minutes. No sludgesupernatant interface had developed when any concentration of cationic polymer was added. The floc was very small, completely dispersed and appeared to be in a state of compression during the entire settling test.

When a nonionic polyacrylamide was added to the sludge, the SVI was reduced. When 5.0 mg/l of a nonionic polymer was used as a settling aid, the SVI was reduced to 178. The physical appearance of the floc changed very little. It appeared very small but the degree of dispersion was less than the dispersion of the raw sludge. The floc appeared more dense. The high molecular weight and large size of the nonionic polymer was effective in consolidating the floc.




78








TABLE 4-3

POLYMER DESCRIPTION AND SVI FOR POLYMER ASSISTED SLUDGES PRODUCED FROM AN INITIAL COLOR OF 5000, Mg++ = 350 mg/l, Ca(OH)2 = 1500 mg/l, pH = 10.8



Polymer Type Subunit Charge Molecular Weight


575C Cationic Amine High 500,000 1905N Nonionic Polyacrylamide Zero 15,000,000 1838A Anionic Polyacrylic Acid High 15,000,000 837A Anionic Hydrolyzed Polyacrylamide 5% Low 15,000,000 835A Anionic Hydrolyzed Polyacrylamide 25% High 15,000,000


Polymer Sludge Volume Index Dose mg/1

None 575C 1838A 1905N 837A 835A


00 352
.03 544 342 408
.05 .547 364
.10 544
.30 362 303
.50 547 250
1.00
1.50 203 198
1.80 422
3.00 547 97
5.00 294 178 81 83 10.00 547 294 67 15.00 275 75 20.00 233 67 25.00 230 64




79




The smaller size molecular weight of the cationic polymer was ineffective in consolidating the floc. The smaller polymer created repulsive forces among the floc particles, probably due to its size and positive charge. The nonionic polymer was larger and not charged. Increased floc settleability resulted from the polymer-floc interaction.

Three anionic polymers were investigated as settling aids. The

first of these was a colloidal polyacrylic acid (1838A) which required an activator before use. Once activated, the polymer formed small spheres approximately 0.5 mm in diameter. The polyacrylic acid decreased the SVI to 230 at a concentration of 25 mg/l. This was a significant reduction in SVI but an excessive polymer dose was required. The available surface area for floc interaction was much less when the activated spheres of polyacylamide were formed. The nonionic polymer was dosed as a clear liquor. It was completely soluble in the coagulated mixture and rendered more available surface area to the floc.

Two additional negatively charged polyacrylamide polymers were investigated as settling aids. These polymers were added to the coagulated mixture as clear liquors and were completely soluble. A large amount of polymer surface area was available for floc interaction. Within one minute of the 30 minute settling test for a polymer dose of 5.0 mg/l, the sludge volume had been reduced 85% with either of the polyacrylamides. The floc changed from a light welldispersed floc to a heavy dense floc. The average size of the floc particles changed from approximately one micron to approximately one centimeter.




80




The interaction between the anionic polyacrylamide and the floc was quite rapid. The rapid interaction is shown in Figure 4.22. The sludge interface is presented as a function of settling time. The solids concentration in the sludge was increased approximately seven:fold due to the addition of 3.0 mg/l of a 5% hydrolyzed polyacrylamide.

The negatively charged polyacrylamide was the most effective

settling aid. A high degree of negative charge was not required on the polymer. This was shown by the identical effectiveness of 837A and 835A. A polymer is negatively charged by hydrolysis. The greater the degree of hydrolysis on the polymer, the greater the polymer charge. The 837A polymer was 5% hydrolyzed, and the 835A polymer was 25% hydrolyzed. A 5% hydrolyzed polyacrylamide means that 95 out of every 100 monomer units are uncharged acrylamide groups; the remaining 5 monomer units are negatively charged acrylic groups.



Uncharged polyacrylamide:



C-0

L NH2



Charged polyacrylamide:

-CH 2-CH-CH2-CHC=O C=0

O- NH2
r- H- n .,"





81








1000 Color= 5000 Mg = 350 mg/I 900 Lime 1500 mg/I


800


700 N o Polymer E0
600


= 500


400


300
3.0 mg I Hydrolyzed Polyacrylamide 200 Anionic Polymer


100

I I I I I
0 10 20 30 40 50 Settling time- minutes Fig. 4.22 Sludge settling velocity for polymer assisted and raw sludge





82




The completely uncharged polyacrylamide was not as effective as either charged polyacrylamide indicating the need for a negatively charged polymer during sedimentation. This need was met by a small degree of hydrolysis. Some degree of interaction between the charged carboxylic functional group and the magnesium floc was necessary for optimum settling.



4-4.3 Mechanisms of Sedimentation

There are two main areas of thought about the-mechanisms of

destabilization of colloids. One area deals with the colloidal stability introduced through the mutually repulsive electrical double layers present on similar colloids. The electrical charge on the colloid surface will attract counterions, and if a sufficient number of counterions are available, colloidal destabilization or sedimentation will result. When this occurs the Van der Waals forces of attraction overcome the electrostatic repulsion and the colloids then can agglomerate and settle. However, there are many possibilities where the electrostatic energy involved in a colloid-counterion interaction will be far less than the energy from chemical bonds between colloid-coagulant interactions. Lamer et al. (1967) have developed a bridging theory in which polymers of high molecular weight can destabilize colloidal suspensions. If the polymer contains chemical groups which can interact with the colloids, then the polymer can destabilize the colloids. Once the colloids begin interacting with the polymer, a bridge is formed between the colloids by the polymer. As an increasing number of colloids become attached to the polymer bridge, the likelihood of destabilization increases. For





83




colloidal destabilization to occur by this model, the polymer dose must be coordinated with the colloidal concentration. It is possible to restabilize a colloidal suspension by too great a polymer dose or by shearing the polymer with too high a mixing energy. There are many instances in wastewater treatment where negatively charged colloids are destabilized by anionic polymers. This phenomena can be explained by an interaction between the functional groups and the colloids, as in the bridging model.

The electrophoretic mobility was measured on the floc particles to determine if the settleability of the floc particles increased as the floc charge decreased. The Helmholtz-Smoluchowski (H-S) formula was used to determine the electrophoretic mobility. Riddick (1974) specified the applicability of different zeta potential formulas based on normality of suspending solution and particle diameter. He recommended the Helmholtz-Smoluchowski formula to measure the electrophoretic mobility of any particle suspended in a 1.ON solution whose diameter was 0.8 microns or greater. The floc particles produced in the NSSC waste by coagulation met these specifications.

The H-S formula for determining zeta potential is as follows: ZP = 113,000(Vt/Dt)EM (4-2) ZP = Zeta potential in millivolts EM = Electrophoretic mobility in microns cm/sec volt Vt = Viscosity of suspending liquid at a given temperature in poises

Dt = Dielectric constant of the suspending liquid




84





Figures 4.23 and 4.24 present the SVI and zeta potential as a function of polymer dose for a nonionic and anionic polymer. It will be shown later that a negative potential occurred on the floc as it was formed in the absence of any polymeric settling aid. If electrostatic reduction was the major mechanism of enhanced settling of the negatively charged floc, then the cationic polymer would have been the most effective settling aid. As Table 4-3 shows, the cationic polymer stabilized the floc and severely hindered settling. Conversely, the anionic polymer was seen to be an effective settling aid. Restabilization of the floc was not achieved at the polymer doses tested. The zeta potential was observed to increase from -13 my to -10 my when the polymer dose was varied from 0 to 5.0 mg/l of 837A. It did not approach zero although the SVI of the sludge changed from 352 to 83. The total change in ZP as settleability increased indicated that decreasing electrostatic repulsion was not the major mechanism for floc destabilization. The controlling mechanism was probably polymer bridging.

For clarification of the coagulation reaction between magnesium and NSSC waste, electrophoretic mobilities were determined on magnesium floc produced at varying pH's in tap water and in NSSC waste. The data for these experiments are presented in Table 4-4. The zeta potential as a function of pH is graphed in Figure 4.25. A magnesium concentration of 350 mg/l was used to produce floc in both the tap water and the NSSC waste.

The zeta potential of the magnesium hydroxide floc produced in the tap water was positive, and increased with increasing pH. The









0 1 0
0
0 >

-I0 -I0 0 , o



4 -20 2 -20
o .


500 500400 400

a

300 c 300
4-








200> 200
-o )


i0O0 I 00



0 1.0 2.0 3.0 4.0 5.0' 0 1.0 2.0 3.0 4.0 5.0 Polymer dose mg/l Polyer dose mq Fig. 4.23 SVI and zeta potential vs. poly- Fig. 4.24 SVI and zeta potential vs. polymer dose for a nonionic polymer mer dose for an anionic polymer # 1905 N # 837A











TABLE 4-4

ELECTROMOBILITY AND ZETA POTENTIAL FOR MAGNESIUM SLUDGE PRODUCED
IN TAP WATER AND NSSC WASTE AT VARYING pH. 10 cm BETWEEN ELECTRODES.
AVERAGE TEMPERATURE 23.50C IN TAP WATER, 250C IN NSSC WASTE.


Approx. Specific pH Volts Colloid Samples Aver. ZP EM Colloid Conductance Direction Time my Micron cm
Dia. micro mhos and sec 25oC sec volt microns Distance 25 C Transversed
microns


Tap Water


1.4 4000 12.5 67 +40 5 4.3 17.74 +1.36 1.3 4000 11.5 67 +40 5 5.5 13.83 +1.06 1.1 4000 11.2 67 +30 7 4.9 11.81 +0.91 0.9 4000 11.0 100 .+30 2 6.6 5.87 +0.45 0.5 4000 10.8 100 +30 9 8.5 4.57 +0.35 0.4 4000 10.6 100 +40 2 12.1- 4.31 +0.33 0.2 4000 10.4 100 +40 2 0 11.9 4.44 +0.34

NSSC Waste


2.0 5000 12.5 20 0 - 0 0 1.8 4000 11.5 67 -40 5 9.4 -8.35 -0.64 1.7 3200 11.2 20 -30 3 14.43 -9.00 -0.69 1.7 3000 11.0 40 -40 7 15.1 -8.61 -0.66 1.5 3000 10.8 30 -15 5 8.0 -8.22 -0.63 0.9 3000 10.6 50 -40 9 9.2 -11.35 -0.87 0.6 3000 10.4 50 -40 6 8.7 -12.00 -0.92 0.2 3000 10.2 67 -40 7 5.8 -13.44 -1.03
011





87







S1.00 Mg(OH)2 in top water



+ .5






ao 10.5 11.0 11,5 12.0 1 2.5

5 pH
N
Mg(OH)1.5R in NSSC waste
-.5





-1.0

Fig. 4.2 5 Zeta potential of magnesium solids in
tap water and NSSC waste at varying pH


2 3



4-J o5

Mg(OH)
6 7

I I I I I
10.5 I1.0 11.5 12.0 12.5 pH Fig. 4.26 Equilibrium concentrations of Mg* and Mg(OH)t with Mg(OH)2 at varying pH




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COLOR REMOVAL FROM A NEUTRAL SULFITE WASTE USING MAGNESIUIvl COAGULATION by JAMES S. TAYLOR A DISSERTATION PRESENTED TO THE GRADUATE COUNCIL OF THE UNIVERSITY OF FLORIDA IN PARTIAL FULFILLMENT OF THE REQUIREMENTS FOR THE DEGREE OF DOCTOR OF PHILOSOPHY UNIVERSITY OF FLORIDA 1976

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TO JANET JIMMY o AND BRIT

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Acknowledgements I wish to express my gratitude to my committee chairman, John Zoltek, Jr. for his overall guidance, understanding and friendship in assisting me in my research. I am deeply indebted to T. deS Furman and J. Edward Singley whose technical and exemplary contributions to my education will never be forgotten. The technical insight and timely assistance given me by Ellis D. Verink, Jr. are' sincerely appreciated. I wish to extend my appreciation to H.F. Berger, who, through the National Council for Air and Stream Improvement, made the funding of my research possible. I will always hold his cooperation and patience in high esteem. The contributed research and extensive laboratory work by Gary Christopher and Bevin Beaudet in completing their masters projects is acknowledged and appreciated. I am sincerely grateful for the sacrifices my wife Janet made and the contributions of my parents in enabling me to pursue my education. The values set forth by my parents years ago came to bear during my research. I have had many rewarding experiences at the University of Florida and am grateful for the opportunity to have been part of that institution. 1X1

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Table of Contents Page 111 Acknowledgements List o£ Tables ^^^ List of Figures Abstract Chapter IX xii 1INTRODUCTION • 1 1-1 General Background 1 1-2 Legal Requirements 3 1-3 Purpose of This Work 3 2COLOR 5 2-1 Color in the Electromagnetic Spectrum 5 2-2 Lignin 7 2-3 Characteristics of Color 9 2-4 Coagulation 13 2-5 Color Removal by Coagulation 23 2-6 Magnesium Coagulation 25 2-7 Color in Pulp Mill Effluents 27 3LABORATORY PROCEDURES 32 3-1 Feed Solutions 32 3-1.1 Synthetic Waste Solutions 32 3-1.2 Coagulation Chemicals 33 3-1.3 Polymers 33 3-2 Analytical Equipment and Techniques 34 3-2.1 Total Carbon Measurements 34 3-2.2 Color Measurement 34 3-2.3 Incineration 35 3-2.4 Jar Tests 35 3-2.5 Metal Analysis 37 3-2.6 Mobility Measurements 37 3-2.7 pH Measurements 38 3-2.8 Settling Tests 38 3-2.9 Solids Analysis 38 3-2.10 Titration Curves 39 3-3 Experimentation 39 3-3.1 Coagulation Experiments 39 3-3.2 Coagulant Recovery 39 3-3.3 Coagulant Recycle 40 lY

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Chapter ^^g 4RESULTS ^^ 4-1 Determination of Coagulation pH and Coagulant Dose ^^ 4-1.1 Coagulation pH ^1 4-1.2 Coagulant Dose 43 4-1.3 Variation of Coagulation pH with Coagulant Dose 4.4 4-1.4 Magnesium Remaining in Solution as a Function of Final pH 47 4-1.5 Magnesium and Ca[OH)„ Dose as a Function of Initial Waste Color 49 4-2 Waste Characceristics 51 4-2.1 Untreated Waste Titration Curves 5,1 4-2.2 Comparison of Untreated and Treated Waste Titration Curves 58 4-2.3 Waste Content 60 4-3 Color Removal Mechanism 60 4-3.1 Color and Magnesium Titration Curves 60 4-3.2 Magnesium, Calcium, Color and Organic Carbon Residuals After Coagulation 63 4-3.3 Stoichiometry of Color Removal from NSSC Waste hy Magnesium Coagulation 71 4-4 Settling of Coagulated Wastes 75 4-4.1 Purpose of Settling Tests 75 4-4.2 Sludge Settleability 77 4-4.3 Mechanisms of Sedimentation 82 4-5 Magnesium Recovery and Recycle 89 4-5.1 Recovery Methods 89 4-5.2 Process Reversibility 90 4-5.3 Color-Cation Interaction ^^ 90 4-5.4 Chemical Equilibrium of Mg -CO2-H2O 96 4-5.5 Sludge Incineration 103 4-5.6 Magnesium Recovery 107 4-5.7 Magnesium Reuse 118 5DESIGN OF A COLOR REMOVAL PROCESS FOR A NSSC WASTE USING MAGNESIUM COAGULATION AND RECOVERY 123 5-1 Coagulation 123 5-2 Sedimentation • 124 5-3 Vacuum Filtration 127 5-4 Incineration 129 5-5 Carbonation 130 6COST134 6-1 Chemical Costs 134 6-2 Capital and Operation Costs 134 6-3 System Costs 137

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Chapter P^S 7CONCLUSIONS AND RECOMMENDATIONS 145 7-1 Conclusions 145 7-2 Reconmendations 147 REFERENCES 149 Biographical Sketch 156 vi

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List of Tables Table Title Page 2-1 VISIBLE SPECTRUM AND COMPLIMENTARY COLORS 6 4-1 GRAPHIC DETERMINATION OF pK^ OF SODIUM BASE NSSC WASTE 57 4-2 UNTREATED AND TREATED NSSC WASTE ANALYSIS 61 4-3 POLYMER DESCRIPTION AND SVI FOR POLYMER ASSISTED SLUDGES 78 4-4 ELECTROMOBILITY AND ZETA POTENTIAL FOR MAGNESIUM SLUDGE PRODUCED IN TAP WATER AND NSSC WASTE AT •VARYING pH 86 4-5 CaCO. PRECIPITATION IN A NSSC WASTE 94 4-6 CHEMICAL REACTION AND pK VALUES CONSIDERED FOR Mg"*'"^-C02-H20 SYSTEM 100 4-7 AVERAGE CHARACTERISTICS OF A SLUDGE PREPARED BY COAGULATING A NSSC WASTE 104 4-8 MgO REACTIVITY AS AFFECTED BY TEMPERATURE 106 4-9 INCINERATED SOLIDS ANALYSIS 108 4-10 CARBONATION OF INCINERATED SLUDGE AT VARYING CONCENTRATIONS OF NONVOLATILE SOLIDS FOR MAGNESIUM RECOVERY 113 4-11 COLOR REMOVAL BY LIME-MAGNESIUM COAGULATION USING THE SAME MAGNESIUM THREE TIMES 119 4-12 COLOR REMOVAL BY LIME -MAGNESIUM COAGULATION USING THE SAME MAGNESIUM TWICE 121 5-1 SOLIDS LOADING FROM SETTLING BASIN 128 5-2 DESIGN SUMMARY FOR THE TREATMENT OF A NSSC WASTE 132 6-1 • CHEMICAL COST TO TREAT A NSSC WASTE 135 Vll

PAGE 8

Table Title Page 6-2 UNIT OPERATIONS COST SUMMARY NSSC WASTE COLOR = 5000 139 6-3 PROCESS COST SUMMARY. IN $/1000 GALLONS OF NSSC WASTE 140 6-4 UNIT OPERATION COST SUMMARY NSSC WASTE COLOR = 2500 142 6-5 NSSC PRODUCT COST INCREASE DUE TO COLOR REMOVAL BY MAGNESIUIvl COAGULATION 144 Vlll

PAGE 9

List of Figures Figure Title 1.1 NSSC flow diagram 2.1 Quinonemethide 2.2 Constitution scheme for lignin 3.1 Standard Pt-Co color curve 9 4.1 Color residual as a function of final pH 4.2 Color residual as a function of final pH 4.3 Comparing NaOH and Ca(0H)2 for color removal via magnesium coagulation 4.4 Comparing NaOH and CaCOH)^ for color removal via magnesitim coagulation 4.5 Verification of coagulation pH' 4.6 Magnesium remaining in solution as a function of final pH 4.7 Lime dose as a function of initial waste color for magnesium coagulation 4.8 Magnesium dose as a function of initial waste color using lime 4.9 Titration curve of sodiiom base Mead NSSC waste with color equal to 2500 4.10 Titration curve of sodium base Mead NSSC waste with color equal to 5000 4.11 Titration curve of sodium base Mead NSSC waste with color equal to 10,000 4.12 Titration curve of sodium base Mead NSSC waste with color equal to 20,000 4.13 Titration curve of sodium base Mead NSSC waste with color equal to 40,000 Page 2 8 10 36 42 42 45 45 46 48 50 50 52 53 54 55 56 IX

PAGE 10

Figure Title Page 4.14 Titration curve of treated and untreated NSSC waste ^^ 4.24 SVI and zeta potential vs. polymer dose for an anionic polymer #837A 62 64 64 68 4.15 Titration curve of raw waste dosed with magnesixim 4.16 Organic carbon and color residuals as a function of final pH 4.17 Magnesium and calcium residuals as a function of final pH 4.18 Color, T.O.C., and Mg"^ residual after Mg coagulation using Ca(0H)2 for pH control 4.19 Color, T.O.C., and Mg"^"^ residuals after Mg • coagulation using NaOH for pH control 69 4.20 Ratios of [OHJ^, ++/ Mg*"^ 72 '-"Mg 4.21 Color and pH of a NSSC waste as a function of CaCOH) concentration ^6 4.22 Sludge settling velocity for polymer assisted and raw sludge ^'• 4.23 SVI and zeta potential vs. polymer dose for a nonionic polymer #1905N 85 85 4.25 Zeta potential of magnesium solids in tap water and NSSC waste at varying pH 87 4.26 Equilibrium concentrations of Mg and MgCOH^ with Mg(0H)2 at varying pH 87 4.27 Color reversibility bar graph 91 4.28 Color remaining as a function of CaCO precipitation • ^-^ 4.29 Color remaining as a function of MgF2 precipitation "^ 4.30 Activity ratio diagram for log C = -1 97

PAGE 11

Figure Title Page 4.31 Solubility diagram of Mg"^"^ in a C-j. = lO""^ M carbonate system 98 4.32 Predominance diagram for log Mg = -1 99 4.33 Color/Mg'^''' ratio as a function of incineration temperature 109 4.34 Precipitation of MgCO^ -3^120 by aeration at various temperatures • 111 4.35 pH as a function of carbonation at various nonvolatile solids concentrations 114 4.36 Magnesium recovered as a function of carbonation time 115 4.37 % magnesium recovery as a function of nonvolatile solids concentration 116 5.1 Design data for sedimentation 126 5.2 Flow diagram for lime -magnesium color removal process 133 XI

PAGE 12

Abstract of Dissertation Presented to the Graduate Council of the University of Florida in Partial Fulfillment of the Requirements for the Degree of Doctor of Philosophy COLOR REMOVAL FROM A NEUTRAL SULFITE WASTE USING MAGNESIUM COAGULATION By James S. Taylor August 1976 Chairman: John Zoltek, Jr. Major Department: Environmental Engineering Sciences A color removal process was developed on a laboratory scale that would remove 90% of the initial color of a neutral sulfite semi-chemical (NSSC) pulp waste. The colored waste was coagulated at a pH of 11, with stoichiometric amounts of magnesium and Ca(OH) The magnesium and Ca(OH) doses were represented by linear equations. The amount of magnesium required for 90% color removal was reduced 25% when Ca(OH) was used for pH control. The reduction in coagulant dose was due to the chelation of the divalent calcium ion and organic acids in the waste. Titrametric techniques demonstrated that the color removal process removed 40% of the acid strength of the NSSC waste, and that 65% of the acids removed had a pK greater than 9. The zeta potential of the coagulated NSSC waste was -1.00 mv at ++ pH 10.3 and zero at pH 12.5. Measurements of the (Mg ), [OH ), organic carbon concentration and color removed during the coagulation pro++ cess indicated that the Mg ion first chelated the organic acids Xll

PAGE 13

causing a 35% color increase. The Mg ion then formed a precipitate which resulted in color removal The empirical formula for the precipitate was Mg(OH) R, where R represents the precipitated organic acids. Once the magnesium precipitate formed, the molar ratios of the magnesium removed to the hydroxides removed was 1.5 for varying magnesium doses at constant pH. The consistency of the molar ratio at varying doses indicated color bodies were removed in magnesium coagulation by a chemical reaction. The color removal process was demonstrated to be completely reversible by varying the pH. In order to reuse the magnesium, the sludge was incinerated to remove the color from the magnesium solids. The optimum temperature of incineration was found to be 550 C. After incineration, all of the magnesium was recovered by bubbling a 10% CO^-90% air mixture through a slurry containing an incinerated solids concentration of 5318 mg/1. The fraction of magnesium solubilized from the incinerated solids was controlled by the Mg -CO^-H„0 system. The controlling solid phase was MgC0_'3H^0. The same magnesium was used three times to remove 90% of the color from three separate aliquots of NSSC waste. After three uses of the coagulant, 93% of the magnesium was recovered. The cost of using this process to treat a NSSC waste with an initial color of 2500 and a flow of 10 mgd was estimated to be $0.27/1000 gal. Xlll

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CHAPTER 1 INTRODUCTION 1-1 General Background Pulp and paper manufacturing is one of the largest industries in the United States. It is also one of the^ajor water using industries in the nation, producing from some mills extremely large volumes of highly colored effluents", which are typically discharged to waterways. Color creates a unique problem in a stream. It is readily identifiable in an aesthetic sense and can detract from the natural beauty of a body of water. The amount of light penetrating a stream would be affected by a colored waste discharged to that stream, and could threaten the eco-system in that stream. The National Council for Air and Stream Improvement, a pollution abatement research organization sponsored by member pulp and paper companies, has recognized this problem and has sought for many years to devise economical and effective color removal processes for all pulp and paper plant effluents. A highly simplified NSSC pulping process diagram is illustrated in Figure 1.1. The wood is prepared for the digestion process by removing the bark and increasing the surface area by a chipping process. The wood chips are then screened, fed into a digestor and mixed with a sulfite cooking liquor. The function of the sulfite

PAGE 15

Sulfite Cooking Liquor Water Debarking Chipping Screening 1 Digestion 1 Washing t Bleaching (chlorine) (oxygen) < Washina I Colored Waste Effulent Further Processing Fig. I.I NSSC flow diagram

PAGE 16

cooking liquor is to separate the lignin from the wood fiber. After the digestion process, the highly colored water soluble lignin is separated from the pulp by washing. The aqueous washings constitute part of the waste effluent. Depending on the ultimate use of the pulp, additional color removal may occur. Bleaching will further lighten the pulp. After each bleaching operation the pulp is washed, producing additional color in the final waste effluent. 1-2 Legal Requirements In 1968 the United States Government passed the Clean Water Act. It, was amended in 1972 to include all pulp and paper mills using a NSSC production process. In the Federal Register under Pulp, Paper, and Paperboard Point Source Category, Effluent Guidelines and Standards, this law states in summary: All NSSC plants must remove 75% of their effluent color by 1983, and all new NSSC paints built after 1975 must remove 75% of their effluent color. 1-3 Purpose of This Work It was the purpose of this research to develop a color removal process for NSSC waste and to give insight into the mechanism by which that color was removed. The first objective was to develop a method of NSSC color removal that could be evaluated for use as a full-scale treatment process. The investigation was limited to jar testingtechniques, with subsequent sludge incineration and coagulant recovery on a laboratory scale. The second objective was to investigate the mechanism by which color removal occurred.

PAGE 17

Techniques employed in this phase of the research were chemical analyses in conjunction with the determination of the stoichiometric relationships developed in the color removal chemical reactions.

PAGE 18

CHAPTER 2 COLOR 2-1 Color in the Electromagn.etic Spectrum Color is a qualitative parameter that does not lend itself to exact engineering measurement. Within the visible region of the spectrum, persons with normal color vision are able to correlate the wavelength of light striking the eye with the subjective sensation of color. Table 2-1 shows the color perceived related to the wavelength. Objects are seen by either transmitted or reflected light. When "white light," containing the entire spectrum of visible wavelengths, passes through a medium such as a solution of NSSC waste, the medium appears colored to the observer. Since only the transmitted waves reach the observer, their wavelengths determine the color of the medium. Chromophores or color producing compounds, absorb certain wavelengths of the spectrum depending on the electronic structure of the compound. A change could occur in the electronic configuration of a compound which could change it from a coloirless to a colored compound. The oxidation of an alpha-quinone would produce a colorless degradation product, but the self -condensation of the same alpha-quinbne would produce a colored product. Very little evidence has been gathered on the amount of environmental degradation caused by color. Properties of pollutants such

PAGE 19

TABLE 2-1 VISIBLE SPECTRUM AND COMPLIMENTARY COLORS Wavelength, mu 400 435 455 480 480 490 490 500 500 -560 560 580 580 595 595 610 610 750 Color Violet Blue Green-blue Blue-green Green Yellow-green Yellow Orange Red Complimentary Color Yellow-green Yellow Orange Red Purple Violet Blue Green-blue Blue-green Source: Day, R.A., Jr., Underwood, A.L., Quantitative Analysis Second Edition, Prentice-Hall, Englewood Cliffs, N.J., (1967).

PAGE 20

as available nutrients or oxygen demand have been shown to degrade the environment. However, the discharge of highly colored effluent would definitely affect the aesthetic quality of the receiving waters. Color would have a detrimental effect on process water used in the production of highly bleached paper. 2-2 Lignin Lignin is one of the most abundant natural products on earth, constituting about one-fourth of the woody tissue in plants. It is responsible for most of the color present in natural waters. The natural formation of this cross-linked polymeric material from coniferyl alcohol and related substances is not presently completely understood. Despite considerable research, the structural characterization of lignin has been only partially successful. Freudenberg (1966) gathered information about lignin structure from direct oxidation of lignin, from bio-chemical experiments related to alcohols, and from lignin degradation with strong alkali, methylation and oxidation. His experiments enabled an estimation of the relative amount of alcohols which served as building blocks of lignin. Lignification occurs in plant cells when alcohols are liberated and oxidized by natural organic compounds in the presence of air. The free radicals produced then combine and build up lignin. Freudenberg (1966) formed a quinonemethide, as shown in Figure 2,1, by combining radicals that resulted from lignol dehydrogenation. From these experiments he suggested that quinonemethide was a tentative structural unit in lignin. Since the quinonemethide has no opportunity to become

PAGE 21

HgCOH HC. HC HoCOH HC I! HC \ OMe HC \ OMe Fig. 2.1 Quinonemethide

PAGE 22

stabilized by hydrogen migration, it adds on the external electrolytes, particularly hydroxyl compounds and preferably water. Quinonemethide is a chromophore, is yellow and can be easily recognized by its intense color absorption extending into the beginning of the visible range. Quinonemethide can achieve limited stability through polymerization, creating large molecules that can still interact with polar compounds. It is possible to construct a tentative constitutional scheme for spruce lignin, which probably is similar to other wood lignins. Such a scheme is presented in Figure 2.2. The lignols which originate during lignin formation, together with the hydrolysis products, reveal different ways in which the C^C_ units are combined. Through natural and industrial processes the lignin is separated from wood fibers and produces chromophores in aqueous solutions. Kirk et al. (1969) prepared lignin by bacterial degradation of wood. The lignin was fractionated by molecular gels into three separate fractions, all of which would absorb light in the visible spectrum. Alder et al. (1966) degraded spruce lignin by acid refluxing in an organic aqueous solution, and was able to separate through fractionation several products that were color producing compoxmds. 2-3 Characteristics of Color Many investigators have attributed the color present in water to the natural or induced degradation products of organic matter. Saville (1917) concluded through electrophoretic studies that most organic color was negatively charged and existed in the colloidal size range. Black and Christman (1963a) found that color collected

PAGE 23

10 H2COH — CH I HC I o HgCOH U I ^^^ — CH OMe HpCOH — CH 1 HC HoCOH OMe HCI HCO(C6H,o05)nH H;>COH 3 CH I HCOH I MeO H2COHCO MeOjj'^H.OMe HCI I •0 HoCOH I HC — CO HpCOH I HC — I HCOH I OMe I OH 1/2 1/2 HCO HpCOH i I HC HC — HC CO OH MeO HgC HC I HC i CH CH I CH2 OMe H2COH -CH HCJ y — .MeO O-OH 0— OMe I OH 0-0 /OOH I HCOH i CH I CHOMe HgCOH CH I CH I CO COH I CH I HC. HC^^^i HoioH \ /"' OMe Fig, 2:2 Constitution scheme for lignin

PAGE 24

11 from ten different water samples had similar chemical and physical characteristics. They demonstrated by dialysis that most of the color present in the ten samples resulted from colloidal suspensions. The infrared spectriom for each of the fulvic fractions, the equivalent weights of those fractions, and the concentrations of the fulvic and humic fractions in each colored sample were similar. Black and Christman (1963b) demonstrated that color intensity was pH dependent and v/ould increase with increasing pH. They also found by dialysis that color existed as a colloid, because only 10% of the original sample color could pass a 4.8 micron filter. Shapiro (1958) found that organic color was mainly dicarboxylic hydroxy aliphatic organic acids of molecular weight 450. He suggested that if phenols were present they were non-color producing organic compounds. He also found that the salts of these acids would pass a cellophane membrane, indicating that they were not colloidal. Shapiro (1958) demonstrated, by chromatographic comparisons, that chemical patterns of color samples taken from different lakes across the country were similar. Any differences that existed in these samples were due to inorganic constituents of the water. Black (1960) suggested that separation techniques used by Shapiro (1958) excluded a portion of the color bodies, and that the excluded portion was in the colloidal size range. Christman and Ghassemi (1966) isolated seven different phenolic compounds common to wood and water humics. Their organic analysis on wood lignins identified carboxyl and phenol groups as the major building units in color molecules. They described these groups as

PAGE 25

12 large aromatic molecules with hydroxyl, methoxy and carboxylic functional groups, Christman and Ghassemi (1966] also found that color extracted from soil would increase with time of soil contact and temperature of the aqueous color medium. Their research showed, as had that of Black and Christman(1966) that color increased with an increase in pH. However, this increase was not linear over the entire pH range. Taylor and Zoltek (1974) using a kraft effluent treated for color removal by massive CaCOH)^ treatment, found that color increase in the waste occurred when the waste was in contact with soil or light. The amount of color increase in the soil-contacted samples was directly proportional to the organic content of the soil. Gjessing and Samdal (1968) studied color fluctuation in a chain of four Norwegian lakes and found that color decreased in all of them except the last lake, where an impoundment occurred. The last lake had a high organic matter content. Gjessing and Samdal (1968) recorded a direct increase in the color of the impounded lake with time of water storage. Their data led to the conclusion that solubilized organic matter produced the color increase in the impounded lake, and the degree of color increase depended on time of impoundment. Packham (1964) separated color from seven different waters into the same classes as did Black and Christman (1966). He found, based on filtration of the fractions, that the fulvic acid fraction existed in the colloidal size range and that the humic acid was in the molecular size range. Packham (1968) also revealed that both the fulvic and humic fractions consisted of complex mixtures of many different

PAGE 26

13 organic acids. Gjessing and Lee (1967) fractionated the color present in a natural water by gel filtration and found molecular size distributions ranging from greater than 200,000 to as low as 700. They found that the molecular size fraction that contained the largest concentration of organic carbon did not produce the greatest color. Midwood and Felbeck (1968) purified a yellow color from organic muck and found that ths organic matter producing color was resistant to chemical or biological degradation. They found that over 80% of the organic carbon was present in the fulvic portion of the color. The infrared spectra showed that aromatic carboxylic acids with aliphatic side groups containing phenolic hydroxyl groups were major components of the color molecules. Day and Felbeck (1974) obtained a yellow water-soluble organic exudate from the domestic waste water fungus Aureobasidium pullulans The exudate contained no himic acid, although it was yellow and was very homogeneous relative to fulvic acid extracts from soil. Day and Felbeck (1974) demonstrated that fungal activity was one source of color in watersheds, and concluded that watershed management with respect to excess biological activity may help eliminate color problems in watersheds. 2-4 Coagulation The reader is referred to comprehensive literature reviews on general coagulation that were published by the American Water Works Association (1971) and O'Melia (1972). In this section the emphasis will be on coagulation as it refers to color removal.

PAGE 27

14 A colloidal dispersion is electrically neutral, so that the charges on the colloidal surface must be counterbalanced by the charges on the liquid immediately adjacent to the colloidal particle. As a result, an electrical double layer exists at every solidliquid interface. These charged ions are attracted to the colloidal surface electrostatically and repelled due to diffusion. The VerweeyOverbeck model as described by Osipow (1972) stated that the LondonVan der Waals forces were the forces of attraction for colloids in suspension. The forces of repulsion resulted from the electrical repulsion of the separate colloidal double layers. Osipow (1972) demonstrated that this model was. further developed and modified by Guoy, Chapman, Stern and Helmholtz. The essence of the final model was that colloidal suspension would be destabilized if the electrical repulsive forces were reduced such that the London -Van der Waals forces would dominate, causing coagulation and sedimentation. This concept was supported in some systems by the Schulze-Hardy rule, which states that the critical coagulation concentration of mono-, diand trivalent ions to coagulate sols of the opposite charge are in the ratio of 100:1.6:0.13. Matijevic et al. (1964a) developed a stabilization-destabilization model for AgBr and Agl suspensions based on neutralization of the electrical double layer with counter ions gained from the hydrolysis of A1(N02)2. Matijevic et al. (1964b) attributed the destabilization of the sols to the Al species on the basis of charge reversal in the coagulation reaction. However, there was a stabilization of the sol which was followed by another sol coagulation. Matijevic et al (1964b) contributed

PAGE 28

15 the final destabilization to Al (OH) precipitation. They presented no explanation for the restabilized sol prior to Al (OH) precipitation, since the sol charge remained positive after the first coagulation. LaMer (1967) developed a bridging theory which provided an acceptable qualitative model for describing the destabilization of colloids with polymers. The main points of the bridging theory were: 1) the polymer must contain chemical groups that would interact with the colloidal surface; 2) that when this happens only a part of the colloidal surface was covered, and the remainder of the polymer would serve as a bridge upon attachment to another colloid; 3) if no other colloid was available for attachment, or the polymer concentration was too great, the polymer would attach itself to the colloid and restabilize the suspension; 4) intense agitation would sever the polymer bonds to the colloidal surface and possibly restabilize the suspension; 5) the amount of colloidal surface area present was directly proportional to the amount of polymer required for coagulation. The bridging theory explained how chemical interactions between an anionic polymer and negative colloid would produce coagulation. Packham (1968) studied coagulation of eight different clays by aluminum hydrolysis and found the coagulant dose continually decreased with increasing concentration. Solubilized calcium and magnesium assisted in lowering the coagulant concentration necessary to destabilize clay suspensions. Packham (1968) demonstrated that the hydrolysis products of cilusa were important to clay destabilization by Tft'fllW tUd" tfW I

PAGE 29

16 zeta potential measurements o£ clay suspensions dosed with and without alum. Although the alum floe had the same zeta potential for optimum destabilization as did the clay suspension, the zeta potential was not zero. Apparantly electrostatic forces were not controlling destabilization. Schott (1968) studied the deflocculation of water sorping clays by anionic and nonionic surfactants. He found that maximum deflocculation was produced when the surfaces of the clay lattices were completely covered with the nonionic surfactants. The American Chemistry Society (1968) published data for aluminum hydrolysis in colloidal suspensions showing that the polynuclear species of aluminum were important destabilization factors. A colloidal suspension was destabilized before any floe was formed using alum as the coagulant. They suggested the forces of adsorption between the colloids and the hydrolysis products were responsible for destabilization, because the hydrplyzed species were hydrophobic and were more likely to accumulate at the solid-liquid interface. Another factor leading to colloidal destabilization was that the hydrolysis products had more than one OH" ion that could sorp at the interface. Their data indicated that as the colloidal surface area concentration increased, an increasing coagulant dose was required to destabilize the colloidal suspension. Langelier and Ludwig (1949) experimented with calciiom and alum flocculation of four different turbid waters varying in exchange capacity. They concluded that the mechanism of colloidal destabilization was controlled by the exchange capacity of the colloids. Michaels (1954), studying the degree of polymer hydrolysis that best promoted coagulation, found that a small amount of hydrolysis was best suited

PAGE 30

17 for destabilization. He suggested that the destabilization mechanism was a two step process: 1) polymer sorption onto the colloidal surface and 2) interparticle bridging following polymer sorption to destabilize the colloidal suspension. Black et al (1965) evaluated coagulation by anionic polymers and demonstrated destabilization followed by restabilization with excess polymer concentration. Since both the polymer and the colloid were negatively charged, the destabilization was not attributed to coulombic forces, but to the build-up of interparticle bridges through other than electrostatic mechanisms. They also found that a higher velocity gradient for a shorter time period was more effective in destabilization than a lower velocity gradient for a longer time period. Ragunathan et al (1973) treated turbid waters with alum and concluded, from zeta potential measurements, that the hydrolysis products of alum were controlling destabilization by sorption mechanisms. Posselt et al. (1968a) examined metal sorption onto a Mn02 anionic sol and found neutral and anionic species did not sorp, a fact supporting an electrostatic mechanism for destabilization. Posselt ++ et al. (1968b) studied Ca sorption onto a negative Mn02 sol and found Ca sorption onto the Mn02 colloidal surface approached a limiting value. The limiting Ca sorption indicated a Langmuir monolayer was probably occurring on the Mn02 surface. They restabilized the suspension with more polymer addition, but did not achieve restabilization with increased metal ion concentrations. However, the increased calcium concentration did broaden the optimum range for coagulation.

PAGE 31

18 They suggested that a choice of coagulant aid would be based on the potential determining ions of the sol. Robinson et al. (1974) reported that larger increases in the turbidity of a river water increased treatment costs and large quantities of alum were required to produce potable water. Nonionic and cationic polyelectrolytes were found to be more effective than alum, suggesting that for this water an electrostatic mechanism was not controlling destabilization. Aluminum hydrolysis was probably removing turbidity by enmeshment in a sweep floe. LaMer (1967) defined coagulation as a kinetic process going from a quasi-stable to a more stable phase, and flocculation as the bridging of already coagulated particles that entered into hindered settling. As an example, he cited hydroxyl groups on flat clay surfaces bonding with hydroxyl radicals of polymers, which allowed metal ions to form insoluble phosphates. LaMer (1967) suggested turbidity, subsidence rate and floe filtration as methods of evaluating destabilization. He also suggested that a negative polymer would best destabilize a negative colloid, because many sites were produced by polymer hydrolysis for bridging. Birkner and Morgan (1968) measured particle size distribution during coagulation and found stronger floe was produced as floe diameter increased. They demonstrated the rate controlling step was particle agglomeration after coagulation, and that intense agitation was responsible for limited polymer sorption. Dollimore and Horridge (1972) investigated flocculation of China clay using polyaerylamides. They found that the maximum clarity was not coincident with the maximum filtration rate as measured by the Kozeny-Carmen ^l1 if ^'R M 4^f^1r-m' O ^V9'tmm4 ^ -maire^^—,~Urjii:^..^:iL^.-^

PAGE 32

19 equation. They concluded that the effective length of the flocsupernatant interface was the controlling flocculation parameter. Hahn and Stumm (1968) studying the kinetics of alum hydrolysis for SiO sols, determined that there were three steps in the coagulation process: 1) forming polynuclear hydrolysis products; 2) the rate of surface coverage or adsorption of the polymer on the colloidal surface; and 3) the rate of particle transport. The rate limiting step for Si02 coagulation was shown to be the rate of particle transport. The rate of coagulation was sho\m to be a function of the collison rate and the collison efficiency. Tenney and Stumm (1965) demonstrated that hydrolyzing metal ions and organic polymers could be used to successfully coagulate bacteria. A linear relationship was found between the optimum concentration of the polyacrylamide polymer and the bacterial concentration. They also found that phosphates were removed with Al in a chemical reaction, and the optimum pH for the reaction was the same as for optimum bacterial flocculation. Stumm and Lee (1961) found that the rate of oxidation of ferrous iron was directly proportional to pH. They found an increase of one pH unit near neutral pH resulted in a 100 fold increase in oxidation rate. Schenk and Weber (1968) also determined that the rate of oxidation of ferrous iron increased with increasing pH. They found that silica retarded the hydrolysis of Fe The hydrolysis was not represented by a first order reaction, but approached linearity with time. They suggested that the solubility relationship may have been altered by complexes formed between silica and iron, and that jj^i i iw i i iii*iiiia

PAGE 33

20 these complexes may have been the mechanism by which activated silica functioned as a coagulant aid. Mohtaoi and Rao (1973) investigated the effects of temperature on aqueous suspensions and concluded that temperature had no perceivable effect on the zeta potential of the sols, or the alum hydrolysis products mixed with cationic, anionic and nonionic polymers. Charge neutralization was determined to be important in destabilizing a colloidal suspension. The neutralization had to be achieved before flocculation occurred. The optimum pH for alum coagulation was found to vary with temperature. However, coagulation with cationic polyelectrolytes was found to be temperature independent of the flocculation rate, optimum pH and coagulant dose. Stumm (1967) demonstrated that metals acted as Lewis acids and had a tendency to stabilize pH. He described metal ion hydrolysis as a function of pH and metal ion concentration. St-umm (1967) stated that multivalent hydrous oxides were amphoteric and that H'*' and OH" were primarily the potential determining ions for such hydrous oxide precipitates. He also stated that metal ions precipitated in the presence of coordinating anions usually as nonstoichiometric mixed precipitates. Stumm et al. (1967) formed polysilicates and classified them into three separate areas: 1) insoluble, 2) stable polymers and 3) the mononuclear wall. They concluded from potential measurements that the interaction between the anionic polymeric phase and the negative sol was due to specific sorption and would overcome electrostatic repulsion. They found optimal destabilization occurred when a fraction of the colloidal surface area was covered and suggested that the mechanism of destabilization for activated silica was the same as for polyelectrolytes.

PAGE 34

21 Sturam (1967) published a hydrolysis model for colloidal destabilization that accounted for bridging and electrostatic effects. He postulated that a fraction of the total colloidal surface area must be covered to produce coagulation. He expressed the model mathematically using a Langmiur isotherm by equating the amount of coagulant necessary to produce a certain fractional coverage to the sum of the residual and sorbed coagulant. The fractional surface coverage necessary to destabilize colloidal sols could only be gained from the residual coagulant or the sorbed coagulant. He showed from his model that the required coagulant dosage to produce destabilization could be independent of surface concentration or linearly dependent on surface concentration. In the Stumm model metals first destabilized colloids due to sorption of the hydrolyzed cationic coagulants and restabilized the colloids due to extensive sorption of the hydrolyzed metal coagulants. Finally a precipitation of the metal occurred that destabilized colloids. If the coagulant became attached to the colloidal surface, the coagualnt dose decreased with increasing colloid concentration. In the precipitation zone the coagulant enmeshed the colloids in a sweep floe. If this occurred, the coagulant dose was not a function of the colloidal surface area. If the colloidal concentration was high, the amount of coagulant dosed could be such that initial destabilization by sorption and final destabilization by precipitation would be indistinguishable. The Stumm model for a large colloidal surface area predicted a large nonstoichiometric coagulant dose that could be reduced if

PAGE 35

22 buffering were removed. A system with a medium colloidal surface area requir^ed a stoichiometric coagulant dose, and if buffering were present, the zone of coagulation was reduced. If low colloidal surface area were present, a large nonstoichiometric dose would be required to coagulate by precipitation. Stoichiometry could be achieved through alkalinity additions. Kawamura (1973) reported that Ca (OH) 2 additions should be made after or during alum coagulation for optimum turbidity and color removal. Jeffcoat and Singley (1975) found that Ca(OH) addition prior to alum coagulation increased turbidity removal and recommended doing so for optimum coagulation results. Hannah et al. (1967) measured alum floe size variations with kaolin, polymer and pol>'phosphate additions. They found kaolin and polymers increased floe size. Polyphosphates hindered floe formation. They recommended that the polyphosphates should be added last in the coagulation process. Hannah. et al. (1967) demonstrated that the order of chemical addition affected the coagulation process. Olson and Twardrowski (1975) studied the products formed by coagulating high alkalinity waters with ferric hydrolysis and concluded FeC03Cs) may be precipitated instead of Fe(0H)2(s). Guilledge and 0' Conner (1973) found arsenic was removed by both alum and ferric chloride hydrolysis. They concluded that adsorption was the removal mechanism. Their results indicated that arsenic was removed better by alum than iron coagulation, the removal was pH dependent and could possibly be the result of a chemical reaction. Stumm and Morgan (1962) found, when doing alkalimetric titrations, that the amount of

PAGE 36

23 base required to titrate the aluminiim mixture was not increased stoichiometrically in the presence of a pyrophosphoric acid. Their data suggested that phosphate removal by alum coagulation resulted from a chemical reaction. Cornwell (1975), studying alum recovery through liquid-liquid extraction, suggested that phosphate was removed by a chemical reaction producing an aluminum hydroxy phosphate. 2-5 Color Removal by Coagulation Black et al, (1963) demonstrated that C(5lor present in six different natural waters was removed stoichiometrically by ferric sulfate coagulation. A graph of raw water color verses required coagulant dose was constructed, and the optimum conditions for color removal did not produce a floe that had zero zeta potential. Singley et al (1967) also found that, to obtain maximum color removal, alkalinity had to be added before coagulation. Ferric sulfate proved to be a better color removing coagulant than alum for the six natural waters tested. Packham (1965) studied coagulation of organic color that was isolated from river water. He separated the color into humic and fulvic fractions. The mechanisms of alum and ferric coagulation were found to be similar, because stoichiometric amounts of these coagulants were required to remove different concentrations of humic and fulvic acids. Packham (1965) proposed from his data that humic acid was entering into a chemical reaction with aluminum. He determined the empirical formula for such a reaction was Al(OH) • R. He found that the fulvic portions were more complex than the humic acid portions

PAGE 37

24 and found little evidence o£ color enmeshment in the Al (OH) £loc. Packham (1965) did achieve an optimum pH for color removal Jobin and Ghosh (1972) studied the oxidation of ferrous iron. They found that the addition of humic acid complexed the ferrous iron and retarded the oxidation reaction. Schnitzer (1971) found, at pH 2.5, that insoluble fulvic acid precipitates were formed with aluminum only when more than one metal ion was added for each carboxylic group present. Mangravite et al (1975) conducted experiments on humic acid removal by alum coagulation. They demonstrated that insoluble aluminum humic precipitation formed at a pH lower than did pure Al(OH),(s) precipitates. They suggested that color was removed from solution in alum coagulation by a chemical reaction. Narkis and Rebhum (1975) concluded that the salts of humic and. fulvic acids acted as anionic polyelectrolytes that reacted chemically with the cationic flocculant, the carboxylate and the phenolate groups. The reaction products formed a colloidal precipitate that could be removed by settling after flocculation. The first step in humic and fulvic acid coagulation was suggested to be a chemical reaction before flocculation by cationic polyelectrolyte addition occurred. Luner and Dence (1970) determined that the color bodies present in a kraft waste were mostly aromatic and quinoid nuclei with carboxyl or ethylenic groups. The color bodies removed in Ca(0H)2 treatment were carboxylic, phenolic or enolic groups that had precipitated in a chemical reaction with calcium. They found that both the precipitated fractions and the nonprecipitated fractions of kraft waste were acidic, but that the nonprecipitated fractions were more acidic and

PAGE 38

25 had a lower average molecular weight than the precipitated fractions. Luner et al, (1970) found, with massive lime treatment, that enolic groups reacted chemically with calcium to produce insoluble precipitates. 2-6 Magnesium Coagulation Stumm (1968) demonstrated that metal cations such as magnesium, aluminum or calcium could function effectively as coagulants. Magnesium hydrolyzes significantly at pH values encountered in lime softening and produces a voluminous floe which hinders solid handling operations. Eidsness and Black (1957) reduced the volume of sludge produced in water r-oftening operations at Dayton, Ohio and Gainesville, Florida by bubbling CO2 into the sludge to dissolve Mg(OH)T Sixty per cent of the Mg(OH) was solubilized, but no attempt at optimization of magnesium recovery was made. The sludge settled more readily after carbonation. Eidsness and Black (1957) concluded that because Mg(OH)„ existed as a gelatinous coordination complex it could accept a proton more readily than the lyophobic crystals of CaCOj. This enabled the sludge volume to be reduced. Black (1971) suggested that not all the Mg(0H)2 could be removed from lime softening sludge because of MgC0_*3H20 precipitation in the carbonation tank. He proposed the use of a floatation process to remove clay from the lime softening sludge before recalcination. Thompson et al (1972a) treated potable water samples in the laboratory using magnesium carbonate as the coagulant. They were able to develop an equation relating coagulant dosage to raw water

PAGE 39

26 color and turbidity when Ca(0H)2 was used to control' pH. Thompson et al. (1972a) proposed a potable water treatment process in which magnesium was used as the primary coagulant, and Ca(0H)2 was used to control pH. The magnesium was recovered from the sludge by carbonation, and the Ca(0H)2 was recovered from the remaining CaCO^ by recalcination. They proposed that sludge handling problems associated with conventional coagulation plants would be greatly reduced utilizing the magnesium carbonate process. Thompson et al. (1972b) compared conventional coagulation systems with the proposed magneslum carbonate system. They also demonstrated that as turbidity and color were reduced, the zeta potential of the residual floe was increased. Dubose et al (1973) successfully extended the magnesium carbonate process to treatment of domestic sewage in a pilot plant at Gainesville, Florida. .He found that magnesium coagulation reduced the total phosphorous to less than 0.1 mg/1 P, and significantly reduced the suspended solids, color and oxygen demand of the domestic wastewater. Black (1974), in pilot plant studies at Melbourne, Florida, found evidence that the color was released from magnesium sludge upon carbonation. This color release was found to stabilize with time, which implied that color release in magnesium recovery may not be a problem. Studies by Taflin et al. (1975), using CO2 gas to redissolve magnesium solids in a lime softening sludge, were discontinued due to a high color return with the recovered magnesium. The potable water produced by using the recovered magnesium as a

PAGE 40

27 coagulant was too colored to be acceptable. Predali and Cases (1973) investigated zeta potential o£ magnesiiim carbonates in electrolytes and found OH' and H"*" to be the potential determining ions for MgC0H)2. The Mg(0H)2(s) colloids had a zero zeta potential at the same pH for varying ionic strength aqueous solutions. They concluded from kinetic considerations that MgOH must have been the source of the positive charge on the Mg(0H)2Cs) colloid. Zoganathan and Maier (1975) found that sand and kaolinite colloids in a solution of 0.005 M MgCl2 had a positive zeta potential for a pH of 10.3 or greater. They attributed the positive zeta potential to the increasing percentage of MgOH"^ relative to the total species of soluble magnesium. 2-7 Color in Pulp Mill Effluents Fitzgerald, Clemens and Riley (1970) demonstrated that while polymers could destabilize colloids in pulp waste, they would not neutralize the electrical double layer. Zettlemeyer, Micale and Dole (1968) studied sludges from pulp mills and found that flocculation kinetics varied with pH for organic carbohydrate base sludge, but not with an inorganic primary sludge from a newsprint mill. Their data indicated that most of the colloidal bound water was interstital and was not chemically held due to the solid-liquid interface. The National Council for Air and Stream Improvement (1971) studied surface properties of hydrogels resulting from treatment of pulp mill waste and found anionic polymers destabilized negative colloids. They suggested that the polymer sorption onto the negative surface was nonstoichiometric.

PAGE 41

28 Davis (1972), using CaCOH)^ coagulation 'at Riceboro, Georgia to remove color from a kraft waste, found calcium solubility decreased as the sodium concentration from the digestion operation was increased. Davis demonstrated that organic carbon, color and calcium concentrations after Ca(0H)2 treatment were related to the initial sodium concentration of the waste. Berger (1964) found that a large Ca(OH) dose (15,000-25,000 mg/1) produced a very settleable floe that removed 90% of original color from caustic bleach effluent. The Domitar Limited Research Center (1974) reported that Ca(0H)2 coagulation was not effective for removing color from sulfite liquors. The Interstate Paper Corporation in Riceboro, Georgia used a smaller chemical dose of Ca(0H)2 (1500-2500 mg/1) to remove in excess of 90% of the initial color in a kraft waste. However, the lime dose did exceed the solubility product of Ca(0H)2 and formed a precipitate. Othof and Eckenfelder (1974) studied color removal from three kraft mill effluents by separate coagulation with Ca(0H)2, ferric sulfate and alum. They suggested that ferric sulfate was the better coagulant because of lower coagulant dose and less voluminous sludge volumes. Gould (1973) reported that the effluent from the caustic extract stage of a kraft bleach plant, when treated with Ca(0H)2, would form a metal organo precipitate that removed 90% of the initial color. Approximately 80% of the Ca(0H)2 was recovered in the sludge. Spruill (1975) found Ca(OH) treatment was very effective for reducing color in kraft wastes, but was ineffective for removing color from sulfite waste. Leszczynski (1972) concluded that of the many processes proposed for color removal from kraft wastes, only Ca(0H)2

PAGE 42

29 precipitation was feasible. Kabeya et al. (1972)' found the rate of absorption of kraft mill lignins on activated carbon to be very low. Katoh and Kimura (1972) found fly ash to be almost as effective as activated carbon in sorping lignin from kraft mill effluents. The National Council for Air and Stream Improvement (1974) studied the mechanism of color removal on activated carbon and found most color bodies existed in the high molecular weight (15,000) range. TOC and color were not removed in equal proportions. They concluded that color removal by activated carbon was not a chemical process, but was due to soi-ption. Swanson et al. (1973) did a detailed study on Ca(6H) treatment of kraft waste and found 86% color reduction, 57% TOC reduction and 17% sugar reduction. There was no removal of material with molecular weights less than 400. Material with molecular weights greater than 5000 was completely removed, and partial removal was observed for material with molecular weights ranging from 400 to 5000. Swanson (1973) suggested color bodies were aromatic groups that carried a negative charge. Tejera and Davis (1970) used alum, AlCl^ and FeCl^ as coagulants in color removal studies on caustic extraction waste and chlorinated waste. They determined both FeCl and AlCl were capable of removing 96% of the color from a kraft mill caustic extraction waste, but both coagulants were hampered in the removal of color from the chlorinated waste. Collins et al (1969) separately concentrated chlorinated and alkaline extraction bleach effluents from a sulfite and a kraft process by reverse osmosis. Lignosulfonic acids with molecular weights in excess of 10,000 were found in the sulfite waste liquor. Jensen

PAGE 43

30 et al. (1964) fractionated spent sulfite waste liquors by gel filtration and ion exclusion and found six different components. Saccahrides and weak organic acids at pH 4 were present in the lower molecular weight range. Fractions above a molecular weight of 40,000 were aromatic lignosulfonic acids and were responsible for most of the color in the waste. Smith and Christman (1969) 'treated kraft and sulfite waste with AlCl^ and FeCl_ coagulants and found either coagulant would remove 90% of the initial color in the kraft waste. Treatment of the sulfite waste by FeCl^ reduced the organic carbon 50%, but increased the color of the sulfite waste. Alum reduced the color of the sulfite waste 67%. Smith and Christman (1969) proposed that the kraft waste had sulhydryl groups on lignin chains and that these groups formed insoluble sulfides during coagulation. The sulfite waste had sulfonate groups in the lignin chain which acted as strong acids and formed hydrolysis products. The mechanism for color removal in the kraft waste was suggested to be a chemical reaction, whereas the mechanism in the sulfite waste was suggested to be sorption on Al (OH) ^ surfaces Rapson et al (1971) used seawater as a source of soluble magnesium along with Ca(OH) to remove color from a kraft waste. Increased color removal was accompanied by the formation of a floe with a larger surface area than the original Ca(OH) floe. A 20% seawater mixture did not remove any more color from the kraft waste than did a 10% seawater mixture. Less color removal was observed when a cardboard effluent was treated, which indicated that different mechanisms

PAGE 44

31 might have been responsible for color removal" for different kraft effluents. The Canadian Pollution Abatement Research Program (1974) used Ca(OH)' and MgCl^ to treat sulfite waste for color removal. They obtained an 86% reduction of NSSC waste using Ca(0H)2 and MgCl2 and a 65% reduction of color using MgCl2 without Ca(0H)2. They were able to remove 86% of the color from a bleach-kraft, unbleached kraft, combined bio-kraft, NSSC-NH3 base and a bio-NSSC waste. They did not optimize pH or coagulant dose in the coagulation process.

PAGE 45

CHAPTER 3 LABORATORY PROCEDURES 3-1 Feed Solutions 3-1.1 Synthetic Waste Solutions All wastes were made by diluting a concentrated color source with tap water to the desired color concentration. The color source was a stored semichemical neutral sulfite liquor which was taken from a NSSC plant digestor after the cooking operation had been completed. This liquor contained the dissolved constituents of the wood. It was referred to as "sulfite waste liquor," which can be the major source of color in the waste stream of a neutral sulfite semichemical pulp plant. The sulfite waste liquor was obtained from plants located in Harriman, Tennessee and Hartsville, South Carolina, owned by Mead Corporation and Sunoco Products Company respectively. The Mead Corporation supplied soldium base spent sulfite waste liquors and ammonium base spent sulfite waste liquors that were used as a waste source. The Sunoco Products Company supplied a sodium base spent sulfite waste liquor which was also used as a waste source. There are different processes and many different types of hardwood trees used in neutral sulfite semichemical pulping. Because of this, it was decided at the beginning of this research to determine 32

PAGE 46

33 if color could be removed from different semichemical neutral sulfite wastes by magnesium coagulation, but only to use the sodium base waste from the Hartsville, South Carolina plant to study mechanism of color removal by magnesium coagulation. The color of the stored NSSC spent liquor varied from 250,000 to 500,000 Pt-Co color units. Consequently, to achieve a working color of 5000 Pt-Co color units, a dilution ratio of 50/1 to 100/1 was required. 3-1.2 Coagulation Chemicals Magnesium sulfate, MgS0.-7H„0, was used as the source of magnesium ions for the color removal process. A stock magnesium solution ++ of 50 mg/ml as Mg was made in order to minimize the volume of coagulant feed dosed in the process. This was achieved by dissolving 532.6 grams of MgSO^ "7^120 in a liter of distilled-deionized water. Calcium hydroxide and sodium hydroxide were used for pH adjustment during the coagulation reaction. Calcium hydroxide was slurried in a small beaker before it was used, whereas sodium hydroxide was added from previously prepared 10 N and 1 N solutions. When necessary, sulfuric acid and hydrochloric acid were used to adjust the pH downward. 3-1.3 Polymers Cationic, anionic and nonionic polymers were prepared from commercial liquids and powders supplied by American Cyanamid Company. The polymers were made from polyacrylamide and amine bases. Stock

PAGE 47

34 solutions of 2000 to 3000 mg/1 were prepared from the solid based polymers by choosing a weighed amount and dissolving it in an aqueous solution by magnetically stirring it overnight. The liquid based polymers at the same concentrations required only one hour of stirring for stock preparations. For the colloidal, acids, an activator supplied by American Cyanamid CN-478) was required for stock preparation. 3-2 Analytical Equipment and Techniques 3-2.1 Total Carbon Measurements Total carbon measurements were determined on a Beckman Model 915 Total Carbon Analyzer in conjunction with a Beckman Model 865 Nondispersive Infrared Analyzer. A three microliter sample was used for analysis. The readout was registered on a to 100 scale and was compared to a standard curve. The carbon standards were prepared from potassium biphthalate for organic carbon, and sodium carbonate or sodium bicarbonate for inorganic carbon. 3-2.2 Color Measurement All color measurements were determined according to NCASI Technical Bulletin 253. This procedure requires all samples for color measurement to be filtered through a 0.80 micron Millipore filter. The pH of the sample was then regulated to 7.6 before the amount of absorbance at a wavelength of 465 millimicrons was recorded. The sample color was then calculated by locating the sample absorbance

PAGE 48

35 on a standard curve relating color to absorbance. If the sample had too great a color to be directly measured, the sample was diluted after filtration. A standard curve was prepared by dissolving 1.246 grams potassium chloroplatinate, K2PtClg (equivalent to 0.500 g metallic platinum) and one gram crystallized cobaltous chloride, C0CI2 '61120 (equivalent to 0.25 grams metallic cobalt) in distilled water with 100 ml concentrated HCl. This solution was diluted to 1 liter with distilled water. This stock solution was defingd as having a standard color of 500 Pt-Co units. A standard curve was prepared and is shown in Figure 3.1. This curve fits the equation: Color = (2183.4) (absorbance) -^ 4.4 (3-1) 3-2,3 Incineration Sludge incineration was determined in a Thermodyne furnace. Model F-A1730. Sludge samples were dried at 103C and filtered through a Buchner funnel on a Whatman no. 40 ashless filter before incineration in the furnace. Incineration temperatures were varied from 180C to 850C. Times of incineration were varied from 15 minutes to 120 minutes. 3-2.4 Jar Tests Jar tests were performed on a Florida Jar Test Machine. Chemicals were dosed simultaneously to four 1 liter beakers. Rapid mixing took place at 100 rpm for three minutes. The Flo.rida Jar Tester was capable of 145 rpm, but due to the heavy floe formed when treating

PAGE 49

36 Color Fig. 3.1 Standard Pt-Co color curve

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37 the highly colored waste, a stirring rate o£ 100 rpm was the maximum that could be attained. Beyond 100 rpm, the magnetic couple between the jar stirrers and the machine was broken by the stirrer over-turning. Slow mix took place at 35 rpm for 15 minutes. The pH was adjusted through both the slow mix and the rapid mix cycles to maintain a constant pH during the coagulation reaction. Floe was allowed to settle for 30 minutes before samples were taken for analysis. If the coagulating mixture had not developed a clear supernatant, a sample was taken and filtered through a no. 40 Whatman filter in order to simplify the required filtering step through the 0.80 micron Millipore filter before color measurement. The G levels of the rapid mix and the slow mix cycles were 110 sec"-^ and 30 sec"-*" respectively. The mixing level in the flocculation stage had a Gt value of 27,000, which was approximately the low end of the range specified in Waste Treatment Plant Design (1971) 3-2.5 Metal Analysis Metal analyses were determined on a Varian Techtron Model 1200 Atomic Absorption Spectrophotometer. Magnesium measurements were made at a wavelength of 202.5 nanometers. Calcium measurements were made at a wavelength of 422.7 nanometers. A.ll samples were filtered through a 0.80 micron Millipore filter and treated with 1 ml of 17% lanthium-HCl solution per 10 ml of sample before calcixim and magnesium values were measured. 3-2.6 Mobility Measurements All mobility measurements were made with a Zeta Meter. The Zeta

PAGE 51

38 Meter was .used in conjunction with a Riddick cell and a stereoscopic microscope. The multiscale 15X ocular micrometer was used in the Zeta Meter for mobility measurements. A platinum-iridium cathode coupled with a molybdenum anode were used in the Riddick cell. 3-2.7 pH Measurements All pH measurements were made on a Corning Model 12 expanded scale pH meter. A Corning silver-silver chloride reference electrode in conjunction with a glass electrode were used for all pH measurements. 3-2.8 Settling Tests _. All settling tests were conducted in a standard 1000 ml graduated cylinder. One liter of waste was coagulated in a jar and imme. diately transferred to a graduated cylinder where the height of the sludge-supernatant interface was recorded. The following formula was used to calculate the Sludge Volume Index, SVI: SVI = ml settled sludge x 1,000 mg/1 suspended solids 3-2.9 Solids Analysis All suspended solids analyses were determined on samples that were filtered through a no. 40 Whatman filter and dried at 103 C for one hour. Nonvolatile and volatile solids were by filtering the samples through a no. 40 ashless Whatman filter, drying at 103 C for one hour, and recording the weight. The sample was then ignited at 550C for 60 minutes, after which it was weighed to determine

PAGE 52

39 nonvolatile solids. Samples were weighed immediately after drying at i03C, but were cooled for one hour in a dessicator after igniting at 550C before .weighing. 3-2.10 Titration Curves 0.01 and 0.1 N H2SO4 and NaOH were used for determining the acid-base strength of the samples. The volume of waste titrated varied from 50 to 200 ml. One minute was allowed for pH stabilization each time the titrant was added to the ^sample. A Teflon covered magnetic bar in conjunction with a magnetic stirrer were used to mix the solution during titration. 3-3 Experimentation 3-3.1 Coagulation Experiments The color removal experiments began by mixing the waste to the desired color concentration and measuring the color as previously described. The next step was regulation of the waste solution pH with CaCOH)or NaOH. The coagulant was then added and the reaction pH was adjusted. Samples for analysis were taken after 30 minutes of settling. Organic carbon measurements and color values were determined immediately after coagulation. The samples to be analyzed for metal concentration were acidified immediately following coagulation. 3-3.2 Coagulant Recovery Several methods were used to recover the magnesium coagulant.

PAGE 53

40 Pollowing coagulation the resulting sludge was filtered through a Buchner funnel and was then dried at 103C for one hour. The dry sludge was ignited at 550C, and the resulting nonvolatile solids were placed in contact with a 10% CO2 gaseous stream or stabilized with H SO to recover the oxidized magnesium. Two 40 liter volumes 2 4 of waste were treated in order to produce a large quantity of sludge. These wet solids were then heated until they achieved a constant weight at 103C. The dried solids were ignited at 550C to remove the coagulated color. 3-3,3 Coagulant Recycle The magnesium was recycled to determine the effectiveness of reusing the same magnesium as the primary coagulation in the color removal process. The method of recycling the magnesium consisted of incinerating the sludge produced in the coagulation reaction at 550C. The resulting nonvolatile solids were carbonated for 45 minutes with a 10% CO2-90% air gaseous mixture. The recovered magnesium was recycled with and without the nonvolatile solids that were not dissolved during carbonation. The two different techniques of recycling the magnesium determined the effectiveness of the remaining nonvolatile solids in the color removal process.

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CHAPTER 4 RESULTS 4-1 Determination of Coagulation pH and Coagulant Dose Development o£ the color removal using magnesium coagulation required that the pH control agent, coagulation pH and coagulant dose be determined. Thu chemicals selected for pH control were Ca(OH) and NaOH because they were inexpensive and commercially available. The coagulation pH andcoagulant dose were defined as the minimum pH and dose that resulted in a 90% reduction of the initial color. .; A three step technique was used to determine the coagulation pH and coagulant dose for the color removal process. First the coagulation pH was found by determining the reaction pH where maximum color removal occurred for a constant magnesium dose. Tlie second step was to determine the magnesium dose that removed 90% of the color at the coagulation pH. Finally the stability of the coagulation pH was verified by repeating the first step for a magnesium dose other than the coagulation dose. If the coagulation pH did not shift, then the coagulation pH and coagulant dose were acceptable. If the shift in the coagulation pH occurred, then the coagulation pH had to be determined as a function of both the coagulant dose and the initial color. 4-1. .1 Coagulation pH Figures 4.1 and 4.2 show the curves from which the coagulation pH can be determined using Ca(0H)2 or NaOH for a NSSC waste with an 41

PAGE 55

42 4000 Color = 2500 1 I 0= Mg"^= 150 mg/lCa(0H)2 3000 \ •= Mg* = 750 mg/l NaOH 1 ( \ S i. 2000 Q. \ o o o !000 n '," ".'.-' 1 .. 9.5 1 0.0 10.5 1 1.0 11.5 pH 12.0 12.5 Fig. 4.1 Color residual as a function of final pH 8000 .| 6000 o o 4000 2000 T 9 Color = 5000 o= Mg^ = 300 mg./l Ca{0H)2 \ o^*t— •=Mg* = 400 mg/lNaOH 1 L f^^^"^ 7 1 1 9.5 10.0 10.5 1 1.0 pH 11.5 12.0 12.5 Fig. 4.2 Color residual as a function of final pH

PAGE 56

43 initial color o£ 2500 or 5000. The constant magnesium dose used for each curve is indicated on the figures. The minimum pH at which 90% color removal was achieved was 10.6. The final color was dependent of the pH control agent. At pH 10 when CaCOH)^ was used, the final color of the waste was increased approximately 30% more than the initial color. This did not occur with NaOH. The same degree of color removal was obtained from pH 10.6 to pH 11.4 using CaCOH)^ or NaOH. The degree of color removal decreases past 11.4 when NaOH was used. This did not occur with Ca(OH) These figures show that a 90% reduction of the original color of the waste was first reached at pH 10.6 for each color. This indicated that the coagulation pH was independent of the variability in the waste color. From the data presented in Figures 4.1 and 4.2 it was concluded the coagulation pH was 10.6. 4-1.2 Coagulant Dose A NSSC waste with a color of 2500 or 5000 was coagulated at pH 10.6 with a varying magnesium dose. These data are presented in Figures 4.3 and 4.4. The pH control agents were CaCOH)^ and NaOH. Magnesium coagulation of the NSSC waste with either pH control agent was able to remove 90% of the color. The required magnesium dose for the waste with a color of 2500 was 100 mg/1 when CaC0H)2 was used for pH control. The magnesium dose was 200 mg/1 when NaOH was used for pH control. The required magnesium dose for the waste with a color of 5000 was 200 mg/1 when CaCOH)^ was used for pH control, and was 400 mg/1 when NaOH was used for pH control. The required coagulant dose was directly proportional to the color of the NSSC waste.

PAGE 57

44 The same color reduction was achi^eved when Ca(0H3 or NaOH was used for pH control. As is indicated in either Figure 4.3 or 4.4, a larger coagulant dose was required to achieve 90% color reduction when NaOH was used to control pH. The coagulant dose was approximately 50% less whenCa(0H)2 wa^s used to control pH. 4-1.3 Variation of Coagulation pH with Coagulant Dose All NSSC waste solutions were made by diluting a stored NSSC liquor with tap water. A waste color of 5000was tested to determine if the coagulation pH was dependent on coagulant dose. Figure 4.5 shows for a color of 5000, the optimum color removal again occurred at pH 10.6. The magnesium dose was 400 m^/1 and Ca(0H)2 was used to control pH. This was the same pH at which maximum color removal occurred at a color of 5000 for a magnesium dose of 300 mg/1 in conjunction with Ca(OH) It was therefore concluded that the coagulation pH did not vary with coagulant dose. The only significant difference between the curves in Figures 4.3 and 4.4 was the presence of the calcium ion when Ca(OH) was used for pH control. The stored NSSC liquor was made in a digestion process which used NaOH and contained a large concentration of sodium. The additional increase in sodium concentration was therefore not significant when NaOH was used to adjust pH. The sodium concentration of wastes with colors of 2500 and 5000 was 40,000 mg/1 and 80,000 mg/1 respectively. The amount of sodium increase when NaOH was added to adjust pH was always less than 1000 mg/1, or less than 2.5%. Monovalent ions, such as. sodium, generally do not complex organic compounds to the same extent as divalent ions.

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45 1000 (O c 3 750 o o k 500 o o 250 h Ca(OH) Color = 2500 pH = 10.6 100 200 300 Mg mg/l 400 500 600 Fig. 4.3 Comparing NaOH and Ca(0H)2 for color removal via magnesium coagulation 200C Color = 5000 pH = l0.6 I500 'E 3 o o I s: 1000 V. O O O 500 Ca(0H) lUO 200 300 Mg mg/l 400 500 600 Fig. 4.4 Comparing NaOH and Ca(0H)2 for color removal via magnesium coagulation

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3000 o V2OOO o o o 1000 J2 O. 46 4000 Color5000 Mg*^= 400 ma^l Ca(0H)2 9.5 10.0 10.5 I PH 11.5 12.0 12.5 Fig. 4.5 Verification of coagulation pH

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47 The additional sodium added to the waste to adjust pH was not a significant increase in sodiim concentration, and did not extensively form any complexes. The total calcium concentration in the untreated waste was approximately 40 rag/1 as Ca"^*. When CaCOH)^ was used to adjust pH the calcium concentration in the waste was increased to 600 mg/1. This was significant because the calcium concentration increased and probably did extensively complex the organic compounds. Figures 4.3 and 4.4 show that the magnesium required to remove 90% of the color was reduced when CaCOH)^ was used for pH control rather than NaOH. Both calciijm and magnesium are divalent ions and will form common complexes with the organic compounds in pulp wastes. When Ca(OH) was added to control pH, Ca*"* complexed many organics that Mg"^"^ would have normally complexed in the absence of the added Ca"^"^. Therefore Ca(0H)2 reduced the required coagulant dose. A magnesium dose of 100 mg/1 removed 90% of the color from a NSSC waste with a color of 2500 when Ca(0H)2 was used for pH control. A magnesium dose of 200 mg/1 was necessary for 90% color removal when NaOH was used to control pH. The complexing ability of the calcium ion was responsible for a 50% reduction in the coagulant dose. Since NaOH is more expensive than Ca(0H)2 and does not reduce the coagulant dose, CaC0H)2 was chosen as the pH control agent. 4-1.4 Magnesium Remaining in Solution as a Function of Final pH The magnesium remaining in the treated waste after color removal as a function of final pH is presented in Figure 4.6. For both

PAGE 61

48 E < 300 \ Ca(0H)2 • Color = 2500 2 50 A Color = 5000 200 ( 150 l^ \ 100 50 1 '. 1 LJ l.„ 10.0 las II.O 11.5 12.0 pHf Fig. 4.6 Magnesium remaining in solution as a function of final pH

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49 colors tested, 35-40 mg/1 of magnesium remained in solution when the final pH was 10.6, The amount of magnesiirai remaining in solution was reduced to 4-10 mg/1 when the final pH was increased to 11.0. Magnesium in solution remained approximately constant past 11. Coagulation at pH 11 saved approximately 30 mg/1 of magnesium from being wasted in the treated effluent. The Ca(0H)2 dose was increased approximately 50 mg/1 to raise the coagulation pH to 11. The magnesium saved was worth more than the Ca(OH) used to increase the pH. Therefore, it was decided to increase the coagulation pH to 11. Increasing the coagulation pH to 11 accomplished two things. First, enough magnesium was recovered to make the coagulation process less expensive. Second, the allowable fluctuation in coagulation pH was increased. The per cent color removed was significantly less at any pH less than 10.6. But when the coagulation pH was 11, a reduction of 0.4 pH units would not significantly affect color removal, 4-1.5 Magnesium and Ca(0H)2 Dose as a Function of Initial Waste Color The waste effluent from a semichemical neutral sulfite plant varies in color intensity. Because of this variability, the amount of CaCOH) and magnesium to remove 90% of the initial NSSC color was determined as a function of the initial color of the waste. The Ca(OH)required is presented in Figure 4.7. The magnesium requirement is presented in Figure 4.8. Both the CaCOH) 2 and the magnesium requirements, were directly dependent on the initial color of the waste. This suggested a stoichiometric relationship between color and coagulant dose.

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50 o o I o o o lopoo 10 mg/l maximum Mg residual Lime mg/l = 750 t 0.10 (color) / 7500 / 5000 / 2500 sf 1 I 7500 3 5000 O u 1 a. o o 2500 500 1000 1500 2000 Ca(0H)2 mg/l Rg. 4.7 Lime dose as a function of initial waste color for magnesium coagulation 90% minimum color removal Mg mg/l = 0.060 (color) 100 200 Mg^ mg/l 300 400 Fig. 4.8 Magnesium dose as a function of initial waste color using lime

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51 4-2 Waste Characteristics 4-2.1 Untreated Waste Titration Curves The acid strength o£ the untreated NSSC waste was determined by titrating 50 ml samples with varying colors with 1.0 N H2SO4. The acid strength o£ the waste was defined as the milliequivalents of acid required to change the waste pH from 12 to 2. The results of these titrations are presented in Figures 4.9 through 4.13 for NSSC waste colors of 2,500 to 40,000. The NSSC waste was obtained from the Sunoco Products Corporation in Hartsville, South Carolina and is denoted as sodium base Sunoco NSSC waste in Figures 4.9 through 4.13. As the color of the NSSC waste was increased, the acid strength of the NSSC waste also increased. This indicated that color was acidic, and an increase in color wo'uld increase the acidity of the waste. The titration curves indicated the acidity of the waste was gained from two functional groups or mixtures of functional groups. These functional groups had pK values in the range of carboxylic acids and phenols or enols. The equilibrium constants for the NSSC wastes were approximated graphically and are presented in Table 4-1. The pK values were identified by locating the inflection points on the titration curves. Approximately 66% of the data points are not represented in Figures 4.9 through 4.13 in order that the titration curves would be uncluttered and clear. These points occurred in two areas, both of which were identified by slight humps on the titration curves. These pK values are approximately 4.6 and 9.8. They differed by four orders of magnitude, which was a large enough separation to allow graphical approximation of pK values.

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52 14 12 r PH 8 6 4 Waste = Sodium base sunoco NSSC Vo = 50 ml Color ^ = 2500 ^ / _J 1 1 1 1 u 10 8 6 4 meq acid X 10 2 4 6 8 meq base X 10 10 Fig. 4.9 Titration curve of sodium base sunoco NSSC waste with color equal to 2500

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53 14 12 Waste = Sodium base PH 10 8 6 4 8 6 4 meq acid X 10 2 4 6 8 meq base X 10 10 Fig. 4.10 Titration curve of sodium base sunoco NSSC waste with color equal to 5000

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54 14 L Waste = Sodium base sunoco NSSC PH 12 10 8 6 • 4 2 10 8 6 4 meq acid 4 6 meq base 8 10 Fig. 4.11 Titration curve of sodium base sunoco NSSC waste with color equal to 10,000

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55 14 12 10 8 Waste = Sodium base sunoco NSSC Vq = 50 ml Color = 20,000 PH 8 6 4 meq acid 4 6 meq base 8 Fig. 4.12 Titration curve of sodium base sunoco NSSC waste with color equal to 20,000

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56 Waste = Sodium base sunoco NSSC PH 12 10 8 6 2 Fig. 4.13 Titration curve of sodium base sunoco NSSC waste with color equal to 40,000

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TABLE 4-1 GRAPHIC DETERMINATION OF pK^ OF SODIUM BASE NSSC WASTE 57 40,000 20,000 10,000 5,000 2,500 9.6 9.7 10.0 10.1 9.5 4.6 4.5 4.7 4.7 4.5 Average 9.7 4.6

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58 4-2.2 Comparison of Untreated and Treated Waste Titration Curves The acid strength o£ the untreated NSSC waste was compared with the acid strength of the treated NSSC waste. This was done in order to determine if any reduction in acidity occurred during the color removal process. The previous titration curves of the untreated waste revealed significant acidity in the weak and very weak acid range. These acids would be ionized at pH 11 and available to participate in a chemical reaction. Magnesium as Mg is a Lewis acid and is capable of reacting chemically with the ionized anions from the waste acids. The titration curves are presented in Figure 4.14 for the same NSSC waste before and after treatment. The acid strength of the NSSC waste was reduced by the color removal process. Before treatment, 0.8 meq of base was required to titrate the waste from pH 12 to pH 9. After treatment, only 0.4 meq of base was required to produce the same change in pH. The amount of base to change the waste from pH 12 to pH 3 before and after treatment was 1.65 and 1.0 meq respectively. Very weak acids have pK values ranging from approximately 8 to 10. Weak acids have pK values of approximately 3 to 5. The total reduction in acid strength during the color removal process was 0.65 meq. Approximately 0.4 meq of this reduction occurred in the very weak acid range from pH 12 to pH 9. This was approximately 60% of the total reduction in acid strength. From the titration data, it was concluded that color reduction by magnesium coagulation does result in at least a reduction of the acids present in the NSSC waste.

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59 PH 6 5 4 Jreoted NSSC waste Color = 485 Vol=50ml Untreated NSSC waste Color = 5000 Vol = 50 ml 10 meq X 10 15 Fig. 4.14 Titration curve of treated and untreated NSSC waste

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60 4-2.3 Waste Content The treated and untreated waste analyses shown in Table 4-2 were done by the United States Air Force Environmental Health Laboratory at Kelly Air Force Base in Amarillo, Texas. The xintreated waste was a sodium base NSSC waste that was prepared from a stored liquor obtained from the Sunoco Products Company. Hydrochloric acid and sulfuric acid were used to adjust the pH to 7.6 before shipment. 4-3 Color Removal Mechajiism 4-3.1 Color and Magnesiiim Titration Curves A volume of 200 mis of NSSC waste was dosed with 80 mg (400 mg/1) of magnesixm as Mg This solution was titrated with a 1.0 N NaOH solution to a pH of 12.0. Tne titration curve for this experiment is presented in Figure 4.15. In the first portion of the curve, an inflection point was present at pH 9.6. This was approximately the second pK determined earlier from the NSSC waste titration curves. The solution was slightly buffered by the colored NSSC waste at this point in the titration. At pH values higher than 9.6, the acids in the NSSC waste were ionized. In the second portion of the curve, another inflection point was found at pH 10.8. Coagulation and 90% color removal occurred at all pH values greater than or equal to 10.6. At pH 10.8 in the titration curves, magnesium was acting as a buffer by hydrolyzing and precipitating out of solution. Color removal was accomplished when the buffering capacity of the colored waste and the magnesium were exceeded. The acids were ionized and were capable of an acid-base reaction

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61 TABLE 4-2 UNTREATED AND TREATED NSSC WASTE ANALYSIS. TREATMENT WAS WITH 150 mg/1 Mg"""*" AND Ca(0H)2 TO ADJUST pH TO 11.0. Item Lab Analysis (mg/l unless noted) Untreated Treated 1. Color Pt-Co units 2000.000 150.000 2. Turbidity JTU's 3.000 4.000 3. Chemical oxygen demand 1510.000 784.000 4. Total suspended matter 0.0000.000 5. Volatile and fixed suspended matter 0.000 0.000 6. Oils and greases 0.800 0.500 7. Surfactants as mg/1 LAS 0.800 1.600 8. Chlorides 56.000 920.000 9. Flourides 1.100 0.500 10. Phosphates 0.500 0.300 11. Sulfates 200.000 1150.000 12. Cadmium .01 .02 13. Chromium (hexavalent) .01 .01 14. Chromium (total) ,03 .05 15. Copper .05 .03 16. Cyanides .01 .02 17. Iron .72 .1 18. Lead .05 .07 19. Manganese .28 .05 20. Silver .01 .02 21. Zinc .1 .05 22. Mercury .005 .005 23. Total organic carbon 530.000 350.000 24. Nitrite nitrogen .06 .02 25. Ammonia nitrogen .8 .2

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62 PH Buffering due to magnesium Initial color = 5000 Vol = 200 ml Mg** = 400 mg/l I N NaOH Slight buffering due to color 1 1 I 2 4 6 8 10 12 14 16 meq Fig. 4.15 Titration curve of raw waste dosed with magnesium

PAGE 76

63 with magnesium. This reaction could involve the color in a formation of an insoluble precipitate. The removal of this precipitate would remove the color bodies. 4-3.2 Magnesium, Calcium, Color and Organic Carbon Residuals After Coagulation A NSSC waste with an initial • color of 2,500 was coagulated with a constant magnesium dose of 150 mg/1. The final pH of coagulation was varied from 10 to 11.5 using Ca(0II)2 to control pH. The magnesium, calcium, organic carbon and color residuals were determined after coagulation. The total organic carbon concentrations and color intensities after coagulation are presented in Figure 4.16. The magnesium and calcium concentrations remaining after coagulation are presented in Figure 4.17. The total organic carbon concentration was reduced 40% when the NSSC waste was coagulated at pH 11.5. The residual color at this point was 157 Pt-Co color units. When the waste was coagulated at any pH from 10.6 to 11.2, approximately 34% of the total organic carbon was removed. The average residual color in this pH range was 197 Pt-Co color units. Increasing the coagulation pH to 11.5 would only remove an additional 1.8% of the initial color. Coagulation at any pH from 10.6 to 11.2 removed 92% of the initial color. When the waste was coagulated at pH 11.5, an additional 6% (27 mg/1) of organic carbon was removed. As was noted, the additional color reduction at pH 11.5 was 1.8%. When the waste was coagulated at any pH from 10.6 to 11.2, the organic carbon was reduced 154 mg/1

PAGE 77

64 400 _300 >^ o> E q S 200 100 9.5 Initial Concentrations T.0.C.=455 mg/l Color= 2500 100 10.5 11.0 11.5 pHf Fig. 4.16 Organic carbon and color residuals as a. function of finol pH 3000 c 3 2000 5 o o O 1000 12.0 150 100 E 50 •\ \ Initial Concentrations Mg* = !l mg/l ^ 00^^ = 15 mg/l \ y y^Z^ remaining ^ A \f \ (Vig remaining u ^*^^....^^_^ • 1 1 '9.5 10.0 Fig. 4. 1 7 10.5 pH 11.0 f 11.5 Magnesium and calcium residuals as a function of final pH Coagulant dose = 150 mg/i Mg Vt 600 400 200 E o 12.0

PAGE 78

65 and 2,300 Pt-Co color units were removed. In this pH range, color was reduced by 15 Pt-Co color units for every mg/1 of organic carbon removed. For each additional mg/1 of organic carbon removed by coagulation at pH 11.5, the color was reduced by only 1.5 Pt-Co color units. From these data it was concluded that not all the organic carbon in the waste contributed equally to the waste. color. The residual color increased 32% when the coagulation experiment was attempted at pH 10.0. The magnesium residual curve in Figure 4.17 shows that no magnesium precipitated out of solution at pH 10, All of the magnesium was therefore available to form chelates with the NSSC waste. Calciirai ions causing increases in the color of a kraft waste due to chelation were reported by Luner and Dence (1971), Color increasing chelates formed by magnesiiom and quinones have been reported by Day and Underwood (1967) Aromatic quinones are an integral part of basic lignin structure, and lignin is responsible for color in pulp waste. The color increase at pH 10 was probably due to the chelation of lignin building units, possibly direct chelation with quinones. None of the 150 mg/1 of magnesium was removed after coagulation at a final pH of 10.0. The total amount of magnesixom available for coagulation was the magnesium dosage and the magnesium present in the waste. For the data presented in Figure 4.17, the total amount was 161 mg/1 magnesium. When magnesium precipitation began, a corresponding drop in color intensity was observed, as was a corresponding drop in organic carbon. The concentrations of magnesium, organic

PAGE 79

66 carbon, and color decreased with increasing pH, indicating that the decreases in these three parameters were related. Corresponding decreases in magnesium, color and organic carbon occurred simultaneously,. A possible relationship for these simultaneous reductions could be a chemical reaction between the color producing organic compounds and the magnesium ions. This relationship would result in the chelation and precipitation of a magnesium organic color-body complex. A second possibility could be the adsorption of the chelated organics onto the voluminous magnesium hydroxide floe." The data presented in Figures 4.16 and 4.17 could conform to either of these postulated mechanisms Ca(OH)_ was used as the pH control agent in these experiments. As the Ca(OH)„ dissolved, the pH and calcium concentration increased. The increasing pH probably caused the solution to become supersaturated with respect to the magnesium-color body compound. This hypothesis is supported by the data presented in Figure 4.17. As the pH increased from 10.0 to 11.5, the calcium in solution increased from 237 to 681 mg/1 as Ca The magnesium in solution decreased from 161 to 7 mg/1 as Mg This increase in calciiom and decrease in magnesium concentrations can be visualized as a reaction between the dissolved Ca(0H)2 and Mg"'"''. Such a reaction is shown in Equation 4-1. The value of AG is -12.601 kcal/mole with the reaction proceeding from left to right. Ca"^"^ + 20H" + Mg"^"^ = Mg(0H)2(s) + Ca** (4-1) Total organic carbon, magnesium and color residuals were

PAGE 80

67 determined after coagulation with a varying magnesium dose. In these experiments, the pH was held constant at 10.6. The pH controls used were NaOH and Ca(OH) The data from these experiments are presented in Figures 4.18 and 4.19. In these figures the residual color, organic carbon, and soluble magnesium are plotted as functions of the total millimoles of magnesium available. The total millimoles of magnesium available consists of the initial magnesiim plus the coagulant dose. The initial color of the NSSC waste used to plot Figure 4.18 is approximately half that of the initial color of the waste used to plot Figure 4.19. The scales in Figure 4.18 are one-half the scales in Figure 4.19. This was done so the two figures could be directly compared without being misleading. In both figures, a decrease in organic carbon was accompanied by a decrease in color removal. The beginning of floe formation is identified by the dashed lines in Figures 4.18 and 4.19. This point was identified when the floe became visible to the naked eye. No magnesium was removed from solution until floe formed. Once floe formed, the color was reduced below the initial color level of the waste. Before this point, a color increase had occurred due to the chelation of magnesium and calcium with the NSSC waste. After formation of floe, the color, magnesium and organic carbon concentrations were decreased. The floe formed at a smaller magnesium dose using CaC0H)2 compared to using NaOH for two reasons. First, the initial waste color treated with NaOH was higher. Floe was formed with Ca(0H)2 when the total amount of available magnesium was 2.19 millimoles. This point was not reached with NaOH until the total millimoles of Mg

PAGE 81

6S 3000 2500 I 2000 o o I o o o 1500 1000 500 500 Vol= 1 liter pH= 10.6 o = Color remaining HD = Mg remaining 375 \ •^'^--. • • = T.OC. o ~ "^^ A • ^^-.....^^^^ en E •i\ -^^..^^^^ d \ ^'""••s^ d ,^\ g 250 ^^ ill \ ^y^^ 125 ye |^^~^^--a -^---^"^^ >^..___^ 1 1. ... .... .. 1 1 1^_ 75 E 50 : c c E a> 25 2.5 5.0 7.5 Total m mo! Mg available Fig. 4.18 Color, T.O.C, and Mg"^ residual after Mg^ coagulation using Ca(0H)2 for pH control 10.0

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69 6000 1000 5000 ^4000 2000 1000 750 c 3 o O o 1 cn a: 3000 b 500 k. d o o d o h250 Vol = I liter pH = 10.6 o = Color remaining D = Mg^' remaining • = lO.C. 150 E 100 : c c "c 'o £ OC 50 10 •tt 20 Fig. 4.19 Total m mol Mg available Color, T.O.C., and Mg residuals after Mg coagulation using NaOH for pH control

PAGE 83

70 available were 9.46. If the initial waste color was the only reason for the larger magnesium requirement with NaOH, then the amount of available magnesium required to form floe would increase in proportion to the color increase. This was not the case. The initial color of the waste treated with NaOH was approximately twice that treated with Ca(0H)2If the magnesium dose was directly proportional to the initial color, then extrapolating the Ca(0H)2 treatment dose would give 4.38 mM/1 as the necessary magnesium dose for the NaOH treatment. It was found that 5.08 mM/1 more of magnesium was necessary to form floe using NaOH. This difference was due to the presence of calcium ions when Ca(0H)2 was used to control pH. Floe formation did not occur with the first magnesium additions with either pH control agent. Calcium ions from Ca(0H)2 complexed some of the organics in the waste that would have been complexed by the magnesium had Ca(0H)2 not been used. This enabled floe formation to occur at a smaller magnesium concentration. Two different cliemical reactions, chelation and precipitation, have been identified in the color removal process. First, chelation occurred between the divalent metal ions and the ligands present in the NSSC waste. The chelation demand of the NSSC waste was satisfied before color removal occurred. The magnesium chelates were partly reduced by using Ca(0H)2 to adjust pH. After the chelation demand was satisfied, floe formation and color removal occurred as shown in Figures 4.18 and 4.19.

PAGE 84

71 4-3.3 Stoichiometry of Color Removal from NSSC Waste by Mangesiv un Coagulation A possible mechanism of color removal was adsorption of the color bodies on magnesium hydroxide floe. The colored organics would not have been involved in a chemical reaction that formed a magnesium compound, but would have become attached to the floe by Van der Waals forces or hydrogen bonding. If magnesium hydroxide floe were formed, two moles of hydroxide would be required for every mole of magnesium removed from solution. If an insoluble precipitate formed that was a chemical compound consisting of magnesium, feydroxide and organic ions, the moles of magnesium ions removed divided into the moles of hydroxide ions removed would be less than two. The moles of magnesium removed divided into the moles of hydroxide removed is presented as a ratio in Figure 4.20. There are three different experiments represented in Figure 4.20. In the first two, CaCOH)^ was used to control pH for color removal from NSSC wastes with initial colors of 2500 and 5000. In the third experiment, NaOH was used to control pH for color removal from an NSSC waste with an initial color of 2500. The moles of magnesium removed were found by measuring the magnesium concentrations before and after color removal. The moles of hydroxide removed were found by difference. First the moles of hydroxide necessary to raise the pH to 10.6 were found. Then this amount was subtracted from the moles of hydroxides required to raise the pH to 10.6 after the magnesium dose was added. The difference was the hydroxide demand of the magnesium used to Qoagulate the color, and was represented as [ohJj^ ++ in Figure 4.20. If. NaOH was

PAGE 85

72 -H8 3 o 2 I 500 coagulated by Mg using lime o = Initial color 5000 coagulated by Mg* using lime ti = initial color 5000 coagulated by Mg using NaOH 10 15 20 25 fr mmol Mg available Fig. 4.20 Ratios of [Oillivig-"-/ Qvlg^l [OH]w it is the moles of hydroxides required by the magnesium for color reMg J is the moles of Mg required for color removal.

PAGE 86

73 used to control pH, the moles o£ hydroxide were measured from the direct addition of a 1.0 N NaOH solution. If Ca(0H)2 was used to control pH, the increase in calcium concentration before and after coagulation was measured. The calcium increase was doubled to determine the moles of hydroxide required in the color removal process. The curves shown in Figure 4.20 represent magnesium to hydroxide molar ratios for the floe formed in the color removal process. The coagulation pH was 10.6. Both Ca(0H)2 and NaOH were used to control' pH. All of the data points in Figure 4.20 represent some degree of color removal When NaOH instead of Ca(0H)2 was used to control pH, a greater magnesium concentration was required before any floe was formed. This was due to chelation and is shown by the separation of the two curves for wastes of equal initial color in Figure 4.20. The initial points on each of the curves in Figure 4.20 represent the beginning of floe formation. As color removal and floe formation increased, the curves eventually stabilized at 1.5. At the low magnesium doses used initially this ratio was not stable because of the chelation demand of the waste. Once this demand was exceeded, the ratio stabilized at 1.5 and remained there for all subsequent magnesium doses. A ratio of 1.5 hydroxide ions to 1.0 magnesitmi ion does not produce an electrically neutralized compound. Another anion had to contribute one-quarter of the total negative charge for the precipitate to be electrically neutral. Color bodies are negatively charged and color was removed as magnesium ions were precipitated. If the color

PAGE 87

74 bodies were involved in a chemical reaction with the magnesium and hydroxide ions, then the molar ratio would be less than two. The negatively charged color bodies would electrically neutralize the precipitate. The ratio of 1.5 indicates that an insoluble precipitate was formed in the ratio of 30H~ : 2Mg : 1R~ where R represents the color body. The formation of an insoluble precipitate is further supported by the stability of the ratio. If the magnesium-color body complex became enmeshed in a Mg(0H)2 precipitate, the overall OH/Mg ratio would be less than two. Some of the magnesium removed would be attributable to the enmeshed chelate and some to the Mg(0H)2 floe. The ratio, however would not be stable for an increasing coagulant dose. As the coagulant dose would increase, the ratio would approach two because mostly Mg(0H)2 would be formed after the chelation demand of the waste was satisfied. As shown in Figure 4.20, the molar ratio does not vary after becoming stabilized at 1.5. From the data presented, it was concluded that a chemical reaction was the mechanism by which color was removed from solution. Calcium hydroxide has been used successfully to remove color from a kraft waste at a pulp plant in Riceboro, Georgia. Dissolved Ca(0H}2, used in magnesium coagulation, could precipitate as Ca(OH)-, or some other compound and remove NSSC color bodies. However, based on solubility product calculations, no Ca(0H)2 would precipitate at the Ca(OH)„ doses and pH required by magnesium coagulation. An experiment was done to determine the color removal capability of Ca(0H)2 with reference to a NSSC waste. The residual color and pH

PAGE 88

75 were determined in a NSSC waste after Ca(0H)2 addition. These results are shown in Figure 4,21. An increase in the residual color of the waste was noted for the initial doses of CaC0H)2This was due to chelation of calcium ions with the organic compounds in the waste. The maximum Ca(0H)2 dose was 2000 mg/1, with a resultant pH of 12 and a color reduction of approximately 45%. No. color was removed until the pH was 11.2, The pH used in magnesium coagulation was 11. At pH 11 the use of Ca(0H)2 alone slightly increased the residual color of the waste. This is shown in Figure 4.21. These results indicate that CaC0H)2 does not remove any color in the magnesium coagulation process when the coagulation pH is 11. 4-4 Settling of Coagulated Wastes 4-4.1 Purpose of Settling Tests The purpose of the settling tests was to minimize the voliime of sludge and gain some knowledge of the factors governing the settling process. The Sludge Volume Index, SVI, was determined after each settling test on all of the sludges produced during coagulation. The settling tests were conducted as described in Section 3-2.8 of Chapter 3. Settling tests were performed on the coagulated wastes and on polymer treated coagulated wastes. Cationic, nonionic, and anionic polymers were used as settling aids in the tests. All of the polymers tested were supplied by the American Cyanamid Company. Since the magnesium dose depended on the initial color of the waste, a constant concentration of suspended solids was produced in

PAGE 89

76 500 Fig. 4.21 1000 1500 Ca(0H)2 mg/l 2000 Color and fxH of a NSSC waste as a function of Ca(0H)2 concentration 5000 4000

PAGE 90

77 all wastes with the same initial color. The use of polymers as a settling aid was found to have a negligible effect on the suspended solids produced. A coagulated waste with an initial color of 5000 was found to have a suspended solids concentration of approximately 1800 rag/1. If the initial color was reduced by half, the suspended solids produced by coagulation also were reduced by half. 4-4.2 Sludge Settleability The type, functional group, charge and approximate molecular weight of the polymers used in the settling tests is presented in Table 4-3. The Sludge Volume Index is also presented in Table 4-3 as a function of polymer dose and polymer type. The SVI of the raw sludge was 352. This was increased to approximately 550 when a cationic polymer was used as a settling aid. The floe was still completely suspended after 30 minutes. No sludgesupernatant interface had developed when any concentration of cationic polymer was added. The floe was very small, completely dispersed and appeared to be in a state of compression during the entire settling test. When a nonionic polyacrylamide was added to the sludge, the SVI was reduced. When 5.0 mg/1 of a nonionic polymer was used as a settling aid, the SVI was reduced to 178. The physical appearance of the floe changed very little. It appeared very small but the degree of dispersion was less than the dispersion of the raw sludge. The floe appeared more dense. The high molecular weight and large size of the nonionic polymer was effective in consolidating the floe.

PAGE 91

78 TABLE 4-3 POLYMER DESCRIPTION AND SVI FOR POLYMER ASSISTED SLUDGES PRODUCED FROM AN INITIAL COLOR OF 5000, ++ Mg = 350 rag/l, Ca(0H)2 = 1500 mg/1, pH = 10.8 Polymer Type Subunit Charge Molecular Weight 575C Cationic Amine High 500,000 1905N Nonionic Polyacrylamide Zero 15 ,000,000 1838A Anionic Polyacrylic Acid High 15 ,000,000 837A Anionic Hydrolyzed Polyacrylamide 5% Low15 ,000,000 835A Anionic Hydro lyzed Polyacrylamide 25% High 15 ,000,000 Polymer Sludge Volume Ind ex Dose mg/1 None • 575C 1838A 1905N 837A 835A 00 352 .03 544 342 408 .05 547 364 .10 544 .30 • 362 303 .50 547 250 1.00 1.50 203 198 1.80 422 3.00 547 97 5.00 294 178 81 83 10.00 547 294 67 15.00 275 % 75 20.00 233 67 25.00 230 64

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79 The smaller size molecular weight of the cationic polymer was ineffective in consolidating the floe. The smaller polymer created repulsive forces among the floe particles, probably due to its size and positive charge. The nonionic polymer was larger and not charged. Increased floe settleability resulted from the polymer-floe interaction. Three anionic polymers were investigated as settling aids. The first of these was a colloidal polyacrylic acid (1838A) which required an activator before use. Once activated, the polymer formed small spheres approximately 0.5 mm in diameter. The polyacrylic acid decreased the SVI to 230 at a concentration of 25 mg/1. This was a significant reduction in SVI but an excessive polymer dose was required. The available surface area for floe interaction was much less when the activated spheres of polyacylamide were formed. The nonionic polymer was dosed as a clear liquor. It was completely soluble in the coagulated mixture and rendered more available surface area to the floe. Two additional negatively charged polyacrylamide polymers were investigated as settling aids. These polymers were added to the coagulated mixture as clear liquors and were completely soluble. A large amount of polymer surface area was available for floe interaction. Within one minute of the 30 minute settling test for a polymer dose of 5.0 mg/1, the sludge volume had been reduced 85% with either of the polyacrylamides. The floe changed from a light welldispersed floe to a heavy dense floe. The average size of the floe particles changed from approximately one micron to approximately one centimeter.

PAGE 93

80 The interaction between the anionic polyacrylamide and the floe was quite rapid. The rapid interaction is shown in Figure 4.22. The sludge interface is presented as a function of settling time. The solids concentration in the sludge was increased approximately sevenfold due to the addition of 3.0 mg/1 of a 5% hydrolyzed polyacrylamide. The negatively charged polyacrylamide was the most effective settling aid. A high degree of negative charge was not required on the polymer. This was shown by the identical effectiveness of 837A and 835A. A polymer is negatively charged by hydrolysis. The greater the degree of hydrolysis on the polymer, the greater the polymer charge. The 837A polymer was 5% hydrolyzed, and the 835A polymer was 25% hydrolyzed. A 5% hydrolyzed polyacrylamide means that 95 out of every 100 monomer units are uncharged acrylamide groups; the remaining 5 monomer units are negatively Charged acrylic groups. Uncharged polyacrylamide: -CH2-CHC-0 NHn Charged polyacrylamide: -CH^-CH-CH^-CH2 I 2 c=o c=o / / 0" NH^

PAGE 94

81 E I 600 x: f 500 o a *w C 400 0) D 300 to 200 100 Color = 5000 Mg^ = 350 mg/i Lime = 1500 mg/i N Polymer -3.0 mg/i Hydrolyzed Polyacrylamide Anionic Polymer 20 30 Settling time— minutes 40 50 Fig. 4.22 Sludge settling velocity for polymer assisted and raw sludge t^
PAGE 95

82 The completely uncharged polyacrylamide was not as effective as either charged polyacrylamide indicating the need for a negatively charged polymer during sedimentation. This need was met by a small degree of hydrolysis. Some degree of interaction between the charged carboxylic functional group and the magnesium floe was necessary for optimum settling. 4-4.3 Mechanisms of Sedimentation There are two main areas of thought about the mechanisms of destabilization of colloids. One area deals with the colloidal stability introduced through the mutually repulsive electrical double layers present on similar colloids. The electrical charge on the colloid surface will attract counterions, and if a sufficient number of counterions are available, colloidal destabilization or sedimentation will result. When this occurs the Van der Waals forces of attraction overcome the electrostatic repulsion and the colloids then can agglomerate and settle. However, there are many possibilities where the electrostatic energy involved in a colloid-counterion interaction will be far less than the energy from chemical bonds between colloid-coagulant interactions. Lamer et al. (1967) have developed a bridging theory in which polymers of high molecular weight can destabilize colloidal suspensions. If the polymer contains chemical groups which can interact with the colloids, then the polymer can destabilize the colloids. Once the colloids begin interacting with the polymer, a bridge is formed between the colloids by the polymer. As an increasing number of colloids become attached to the polymer bridge, the likelihood of destabilization increases. For

PAGE 96

83 colloidal destabilization to occur by this model, the polymer dose must be coordinated with the colloidal concentration. It is possible to restabilize a colloidal suspension by too great a polymer dose or by shearing the polymer with too high a mixing energy. There are many instances in wastewater treatment where negatively charged colloids are destabilized by anionic polymers. This phenomena can be explained by an interaction between the functional groups and the colloids, as in the bridging model. The electrophoretic mobility was measured on the floe particles to determine if the settleability of the floe particles increased as the floe charge decreased. The Helmholtz-Sraoluchowski (H-S) formula was used to determine the electrophoretic mobility. Riddick (1974) specified the applicability of different zeta potential formulas based on normality of suspending solution and particle diameter. He recommended the Helmholtz-Smoluchowski formula to measure the electrophoretic mobility of any particle suspended in a 1 ON solution whose diameter was 0.8 microns or greater. The floe particles produced in the NSSC waste by coagulation met these specifications. The H-S formula for determining zeta potential is as follows: ZP = 113,000(V^/D^)EM (4-2) ZP = Zeta potential in millivolts EM = Electrophoretic mobility in microns cm/sec volt V = Viscosity of suspending liquid at a given temperature in poises D^ = Dielectric constant of the suspending liquid

PAGE 97

84 Figures 4.23 and 4.24 present the SVI and zeta potential as a function of polymer dose for a nonionic and anionic polymer. It will be shown later that a negative potential occurred on the floe as it was formed in the absence of any polymeric settling aid. If electrostatic reduction was the major mechanism of enhanced settling of the negatively charged floe, then the cationic polymer would have been the most effective settling aid. As Table 4-3 shows, the cationic polymer stabilized the floe and severely hindered settling. Conversely, the anionic polymer was seen to be an effective settling aid. Restabilization of the floe was not achieved at the polymer doses tested. The zeta potential was observed to increase from -13 mv to -10 rav when the polymer dose was varied from to 5.0 mg/1 of 837A. It did not approach zero although the SVI of the sludge changed from 352 to 83. The total change in ZP as settleability increased indicated that decreasing electrostatic repulsion was not the major mechanism for floe destabilization. The controlling mechanism was probably polymer bridging. For clarification of the coagulation reaction between magnesium and NSSC waste, eleetrophoretic mobilities were determined on magnesium floe produced at varying pH's in tap water and in NSSC waste. The data for these experiments are presented in Table 4-4. The zeta potential as a function of pH is graphed in Figure 4.25. A magnesium concentration of 350 mg/1 was used to produce floe in both the tap water and the NSSC waste. The zeta potential of the magnesium hydroxide floe produced in the tap water was positive, and increased with increasing pH. The

PAGE 98

85 ^^ il < I I aiH i 1 1 o I o o o O O o o o O O in t ro CNJ o Q E O I rd C o > CO E a o c o o N o TJ £ (VJ < ro 00 SJIOAJIIJUJ -|D!;U8|0d d;3Z xspui euiniOA aSpnis o I o CM I O o in o 1 Jm m O F o. 1 en o > o ^ o c Vo •^ ^. c fD> H) o E o c o to 1 CL o O M o O N <40) T3 CO o (U c q OJ E a a o ^ 0) CL fc o (VJ 2 m o S||OAj||jUi -|Dj^U9;0d D43Z xgpu] 9Uin|0A sBpnis

PAGE 99

86 Cfl w Q Q m O u OS • 3 H PJ Q U H O PJ CO b: J < Dh PJ s pj z u a m c/3 o PJ c/:) 3 S 2 J f-H CO W 2 CO t-H g B u 1— 1 uo en in M O H h1 Di C4 H < 2 J 2 > M ca aj < H H u H O <:o a. un PJ • < H to H C/3 CN w < M S PJ o; Q CO H •^ CO <; 2 cS >PJ H Q O, 2 S < PJ l-H H CQ a; O W PJ H U Q < < Di S Di EPJ U &, > w <; <: ij H PJ 2 s o t^ c o u ^ o >o PJ u LO o U CM H o tSl S Lrt CM -H -d .-I o C t-H 0) Hi O JH U -H Q 13 0) (D 1/5 t/l > C C O u •H s o > E P< 0) o o t/l •H C o M-i nJ A •H 4-> s O U 01 p o OhX) f-l c/5 C o o •H u £ • T3 X -H O O Cflf-l 1— I -H Oh O < u +-> CTi E\D vO 1— t LO LO to "* to O CTi -* to to fO 1— I t— I O O O O O + + + + + + + ^ to 1— I t~^ r^ 1— I "* [~00 00 00 LO to t^ ^0 i-H LO ^ 'd" 'd" to LO CTi \0 LO r— I CTl •^ LO 'y ^ 00 (N r— I un Lo r^ o) CT) oi CN o o o o o o o 'y to to to ^ -y + + + + + + + t^ r-^ r^ o o o o \0 ^ ^ o o o o LO LO CN O CO \o CN I— I .— I t— I O O O o o o o o o o o o o o o o o o o o o o o o ^ ^ -nJ>* ^ ^ •* ^ to >— I O^ LO -* CN 1— I r— I r-l O O O O in C/D C/5 2 rjO^ \0 to t->. (N to O^OvOvO^OOOCTlO O O O O O O i-H I I I I I I I OLnOr-HCNLOO-^ to O vO CM to o •^ 00 CTi 00 00 >— I (N to I I I I 1— I rH rH I I I 'y -vf rH O Cvl r^ 00 en •* i-o 00 o^ 00 LO I LO to r^ LO CT^ ^ t~^ OOOOLOOOO TJto "^ — I ^ rt "^ I I I I I I I oc^ooooor^ tNvOCN'^tOLOLnvD lOLOCNOcOvO-^CN CN>— It— (i—tOOOO oooooooo oooooooo OOCNOOOOC LO-^tOlOtOtOtOtO Ooot^i~~Lnc7i>.orM CNl— It— It— IrHOOO

PAGE 100

+ 1.0 + .5 > E I Q. O N Mg(0H)2 in tap water -.5 10.5 1 1.0 Ili5 PH Mg{OH)| gR in NSSC waste 12.0 -1.0*Fig. 4.2 5 Zeta potential of magnesium solids in tap water and NSSC waste at varying pH 87 10.5 Fig. 4.2 6 Equilibrium concentrations of Mg"* and Mg{OH)'*" with Mg(0H)2 ai varying pH

PAGE 101

88 greatest change in zeta potential with respect to pH occurred at pH 11.1, which was close to 10.8, the pK for magnesium hydroxide. The positive charge on colloidal material formed in the presence of Mg in aqueous solution has been attributed by Loganathan and Maier (1975) to the formation of Mg(OH) The equilibrium concentrations of Mg(0H3 and Mg are presented in Figure 4.26 from the equilibrium expressions given in Equations 4-3 and 4-4. The equilibrium concentrations were equated to the activity of these species for these calculations. ++ 2 -10.8 Kgp = {Mg^^XOH"} = 10 C4-3) • + ++ -2.4 K = {Mg(OH) }/{Mg }{0H} = 10 (4-4) The zeta potential as shown in Figure 4.25 increases with increasing + ++ pH. The ratio of the singularly hydroxylated species Mg(OH) to Mg also increased with increasing pH. However, the equilibrium concentration of Mg(OH) decreased. As the concentration of Mg(OH) and Mg decreased, more magnesium hydroxide precipitated from solution. From the data in Figure 4.25 and the equilibrium graphs in Figure 4.26, one can suggest the floe was positively charged in tap water due to its own nature and not due to the sorption of the singularly hydroxylated species. The zeta potential of the magnesium solids produced during the coagulation of the NSSC waste changed significantly to pH 10.7 and past pH 11.7. The first pH was approximately the minimum pH at which 90% color removal was obtained using magnesium coagulation. In the first portion of the zeta potential curve for the treated NSSC waste a

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89 chemical compound was formed that v.as negatively charged. The colloids maintained a relatively constant negative zeta potential through the pH range of maximiM color removal, pH 10.6 to 11.5. When the pH was raised to 12.5, the zeta potential increased to zero as a result of the formation of the neutral magnesium hydroxide. As previously shown, color removal decreased past a coagulation pH of 11.5 when NaOH was used to control pH. These data indicated as the pH exceeded 11.5, hydroxide ions probably displaced color bodies from the magnesium floe. The ability of the hydroxide ion to successfully compete for the reaction sites on the magnesium floe was probably reduced when Ca(OH) was used to control pH due to the complexing ability of the calcixrai ion. 4-5 Magnesium Recovery and Recycle 4-5.1 Recovery Methods Experiments were conducted to find a recovery process that would recover the magnesium in a usable form. One method of recovering the magnesixM was to acidify the magnesium sludge with CO2 gas. This would have returned the magnesium to solution but would have possibly returned the color to solution also. Taflin et al. (1975) initially recovered magnesium from sludge by carbonation in Minneapolis. However, they had to discontinue magnesiiim recovery because of the color build-up in the recycled magnesium feed. One method of removing color from the magnesium was to incinerate the sludge before the magnesium was recycled. This would have oxidized the color bodies to CO2 and left magnesium in a usable form. The CO2 gas from incineration could be used to acidify the

PAGE 103

90 incinerated magnesium sludge. 4-5.2 Process Reversibility A color release experiment was conducted on a one liter sample of NSSC waste with an initial color of 2500. This waste was treated with the optimum design doses of 150 mg/1 magnesium and 1000 mg/1 Ca(OH) The final pH of the coagulated waste was 11.0, and the final color was 200. The pH was lowered to 9.0, the color was measured, and then the pH was raised to 11.0 and the color was again measured. This oscillation of pH completely dissolved and reformed the sludge. When the 'sludge was dissolved, the color returned to 2500. When the sludge was reformed, the color returned to 200. These data are presented in Figure 4.27. The reversibility presented a problem for magnesium recovery by carbonation. Direct recycle with only carbonation would be feasible if the precipitated color remained on the CaCO, floe during carbonation. 4-5.3 Color-Cation Interaction The carbonate alkalinity was low in the NSSC waste because it was prepared from low alkalinity tap water. An experiment was designed to determine if CaCO^ precipitation in situ with the colored waste would remove the NSSC color. If this happened, incineration was not required for magnesium recovery. If CO2 gas was used to recover the magnesium in actual plant operation, the return stream would contain a large amount of HCO^ alkalinity. The exact amount of HCOj-alkalinity would be controlled by the chemical

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91 3000 2000 (0 o o I a: 2 1000 o o Color = 2500 Mg^ = 150 mg/l NaOH Color measured at pH 7.6 after coagulated NSSC waste adjusted to pH 9 ^Color measured at pH 7.6 after coagulated NSSC waste adjusted to pH It 4 5 6 pH cycles 8 Fig. 4.27 Color reversibility bar graph

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92 interaction of magnesium, CO2, water and the organic acids present. In the CaCOj precipitation experiment, Na^CO^ and CaCOH)^ were dosed in equimolar amounts to produce CaCO^ sludge. The amount of CaCOprecipitated per liter ranged from 0.2 to 260 milliraoles. Since NaXO„ and Ca(OH)were the source materials, the amount of 2 3 Z CaCO, precipitated was determined by the calcium difference and checked by the carbonate difference before and after coagulation. The carbonate difference was measured by determining the total inor• ganic carbon before and after coagulation. The residual color change as a function of the millimoles of CaCO^ precipitated is presented in Figure 4.28. The dose data and the change in TOC concentration is summarized in Table 4-5. The color of the NSSC waste was reduced 65% when 265.5 millimoles of CaCO, were precipitated from solution. This amount of calcium was not available from the Ca(0H)2 used to control the pH in the color removal process. The maximum amount of calcium available from the Ca(OH)^ dose for a color of 5000 would be 16.89 millimoles. If this amount of calcium was precipitated as CaCO^, there would be approximately a 4% color removal based on the data presented in Figure 4.28. It has been shown previously that the color removal process was reversible. The precipitation of CaCO^ did not remove a significant amount of color in this process. The color would still be present in the recovered magnesium solution. It was concluded that the color had to be removed from the magnesium floe by incineration in order to reuse the magnesium.

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93 6000 < Initial color increase due to chelation 50 100 150 200 CaC03 precipitatedmmol/l Fig. 4.28 Color remaining as a function of CoCO^ precipitation 3000 o. 2000 c 'c "5 E I 1000 o 5 .10 15 20 MgF2 precipitatedmmol/l Fig. 4.29 Color remaining as a function of MgFg precipitation

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94 eo < a: (/) o CO EO 2 WD O < ^ a < II o E2 O HH Q EU m OS D, to O u U o o u en -J =) <; E EO 1-1 cd 2 2 M O CO CM X + + O U •H ^-^ O 03 — S Oill to O -o tOT3 0) O -H s U ^ f-l rt O o u w <4A >-< f-l ctf O C 'H •rH O tt. u T— I c o --r^ t^ bj:i C •H U — I MO --^ H [M fH S O 1 to > -H o o ^-, + <— 1 Oj o *-^^ U 6 M
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95 A second color-cation experiment was conducted to further investigate the degree o£ magnesium-color interaction. To investigate this, the amount o£ color removed by magnesium was determined for a magnesium precipitate that was not a hydrolysis product. MgF2 was selected as the precipitate because of its low solubility and pH independence. Different jars of NSSC waste with an initial color of 2950 were dosed with F from NaP and Mg from MgSO. The Mg "''"'' dose was 175 mg/1 for each jar. This was the same magnesium dose that would have been used for color removal by magnesium coagulation at high pH. All jars were in a state of MgF2 supersaturation. The general definition of the supersaturation ratio (S) is S = (Q/K) where Q is the ion product, K is the equilibrium constant, and n is the number of ions in the neutral molecule (n = 3 for MgF^) The jars were allowed to stand for 24 hours before samples were taken. The floe formed was very small and not nearly as voluminous as the magnesium floe produced at high pH. The residual color as a function of the millimoles/liter of MgF2 precipitated is presented in Figure 4.29. The initial color of the NSSC waste was reduced 55% when the initial MgF2 supersaturation ratio was approximately 16. At this point all of the magnesium dose, 7.3 mmole, had been precipitated from solution. In the MgF^ experiment, the precipitation of 1 mmole of magnesium removed 231 Pt-Co color units. If more magnesium had been used, the amount of color per mmole of magnesium precipitated might have been larger. However, a significant magnesium-color interaction had been demonstrated in a non pH dependent chemical reaction involving magnesium precipitation.

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96 The data in Table 4-5 shows that the organic carbon level was reduced 40% and the color was reduced 66% when 266 iranole/1 of CaCO_ were precipitated. The molar ratio of Ca removed to CO, removed was consistently less than one as shown in Table 4-5. This indicated that the color removal mechanism with Ca precipitation was similar ++ to Mg precipitation. The negatively charged color body acted as a contributing anion in a chemical reaction. When CaCO, was precipitated from a NSSC waste, the removal of one mmole of calcium reduced the color 12.4 Pt-Co units. In the lime-magnesium color removal process the precipitation of one mmole of magnesium removed approximately 360 color units from solution. Comparison of the ratios of color removed per mmole of magnesium precipitated demonstrates that magnesium is more color reactive than calcium. The magnesium sludge produced at high pH was more effective than the sludge produced in the non pH dependent MgF^ experiment. The larger floe produced by magnesium coagulation at high pH possibly provided more sites for magnesium-color intei-action. From these data it was concluded that magnesium was responsible for color removal in the lime-magnesiiom color removal process. 4-5.4 Chemical Equilibrium of Mg •-CO^-H The chemical equilibrium of the Mg -CO^-HnO system was described by constructing an activity diagram, a solubility diagram and a CO2 predominance diagram. These diagrams are presented in Figures 4.30, 4.31 and 4.32. The equations used to construct these diagrams are presented in Table 4-6. A predominance diagram was constructed for

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97 8 PH 10 II Fig. 4.30 Activity ratio diagram for iogCj=-l. Equations defining relative activities ore given in Table 4.6. ForOMglOH)^, (DMgC03, @MgC03SHgO, and (DMg4(C03)3(0H)2-3H20 which are all solid forms.

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98 o o ^ H. Fig. 4.31 Solubility diagram of Mg in a Cj= iO~ M carbonate system. Equations defining relative activities for all species shown are given in table 4.6. MgCO^SH^O, Mg^(C02)2(OH)2-3H20 and MglOH) are solid forms

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99 \ -1 1 1 1 1 \^ MgCOj3H2O cv 2 \ O \ o H\ a. 4 \ o -J 1 Mg \Mg4(C03)2(0H)2-3H20 6 8 1 1 _i 1 Mg{0H)2 7 8 PH 10 HFig. 4.32 Predominance diagram for log Mg MgC03-3H20 is only stable thermodynomicQlly at high Pqq MgCO -SH 0, Mg^{C02)2(0H)2-3H20 apd MglOH)^ are solid forms

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TABLE 4-6 CHEMICAL REACTION AND pK VALUES CONSIDERED FOR Mg'-'-C02-H20 SYSTEM HCO3" = H"" + CO3 Mg4(C03)3(0H)2-3H20(s) = 4Mg*'' + 3CO3 + 20H" + 5H2O 2' 100 0, „ ^. „ pK at 25 C Reaction ^ H20=H^-OH14.0 a H,CO, = 2H" + Co! 16.6 a 2 3 aq ^ 10.3 a H2CO3* = H"^ + HCO36-2 ^ 1.47 b 29.5 b 4v.--3^3v MgC03-3H20(s) = Mg""* + CO3 + 3H2O 5.4 b ++ ^^= 4 9 b MgC03(s) = Mg + CO3 Mg(OH)„(s) = Mg"^ + 20H' 10-85 "" a Meites, L. Handbook of Analytical Chemistry McGraw-Hill, N.Y., N.Y., (1962) b Stumm, W., Morgan, J.J., Aquatic Chemistry John Wiley § Sons, N.Y., N.Y., (1970) c Day, R.A., Underwood, A.L., Quantitative Analysis Prentice HalL Englewood Cliffs, N.J., (1967)

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101 a Mg** concentration of lO""-^ M. The activity ratio" and solubility diagrams were constructed for a total alkalinity o£ 10 M. The stable species were determined as a function of pH from the activity ratio and solubility diagrams in Figures 4.30 and 4.31. The activity ratio diagram was drawn using Mg"*"* as the reference state. MgCOH)(s) was the stable species above pH 11.2. From pH 11.2 to pH 8.5 the controlling species was Mg^ (00^)3 (OH) ^ 3H2O (s) Below pH 8.5, MgCOj-SH 0(s) was the stable species. These pH ranges of the stable species will shift with variations in the total alkalinity. The practical significance of these diagrams was to identify the species controlling the solubility of Mg"^"^. For a Cj of 10"-^ M, this species is MgC0^-3H20(s) for an open system using CO^ gas as the proton source. The dominance of MgCO^-SH^OCs) is seen in the solubility diagram in Figure 4.31. For a C-p of 10" M an equilibrium point was reached at pH 7.2. This is identified by the intersection of the equilibrium lines in the activity diagram of HCO3" and MgC03-3H20(s) The bicarbonate ions from the solubilizing of CO2 gas are in equilibrium with the solid MgCO •3H20(s) The solid was dissolved by reaction with a proton from H2CO3 This was represented by the following reaction and pH calculations. MgC03'3H20(s) + H"" = Mg""" + HCO3" + 3H2O (4-5) K = 10"''^ -^^ = (Mg"""") (HC03~) (4-6) (H*) (HCO3") = a^C^ (-4_7-, at point of maximum solubility 2(Mg++) = (HCO3-)

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102 104-9 = (a^C^)2/2(H"') (4-8) pH = 7.2 (4-9) It was interesting to note from the activity ratio diagram in Figure 4.30 that the species MgCO,(s) never does exist at equilibrium. A predominance diagram is presented in Figure 4.32. In the predominance diagram the species in equilibrium with CO2 gas is shown. For the CO2 concentration in flue gas (10%-14% CO2) the equilibriumspecies was MgCO„-3H20(s) The magnesium concentration in the recovery tank depends on the pressure of CO2 gas and the equilibrium species, MgC02'3H20(s) For a 10% CO2 gaseous mixture in an aqueous solution, 3.89 x 10 mols of CO2 are dissolved at 20C. A solution containing solid Mg(0H)2(s) mixed with a 10% C0„ gas would come to equilibrium at pH 7.5, The -2 total alkalinity calculated at equilibrium was 6.0 x 10 M, with 94% of it in the HCO-~ form. The maximum calculated magnesiiora solubility considering MgC02'3H20(s) as the dominant species was 863 mg/1. The laboratory system employed was never at true equilibrium. The source of the magnesium was incinerated MgO(s) or Mg(0H)2(s). If Mg dissolution from Mg(OH)„(s) was kinetically favored as compared to MgC0„*3H20(s) formation, a supersaturated system with respect to MgC0_-3H20(s) could exist. The presence of organic acids in the recarbonation basin would increase the Mg solubility, and the species controlling Mg solubility may be an organo-metallic compound. This would be more likely when incineration was not used to remove color, because the organics would not have been converted to CO

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103 4-5.5 Sludge Incineration The colored sludge that was to be incinerated was prepared by treating two separate 40 liter volumes o£ NSSC waste with 350 mg/1 Mg** and 1250 mg/1 of Ca(0H)2. An average of four separate analyses of NSSC waste and sludge is presented in Table 4-7. All sludge samples were dried at 103C until constant weights were obtained before the sludge was incinerated. The magnesium dose plus the magnesium in the NSSC waste was 362 mg/1. The magnesium concentration of the supernatant was 14 mg/1. 1082 mg of solids resulted from the incineration of the sludge produced in one liter of waste. The incinerated solids contained 32.2% magnesium, A mass balance on magnesium revealed that 362 mg were available before color removal and 362 mg of magnesium were accounted for after color removal. The total recovered magnesium in the solids was 96%. Some magnesium was lost in the supernatant. This was probably due to the filtering of the sludge through a no. 40 Whatman ashless filter which passed some very small particles of solid magnesium salts. These would have been retained by the 0.80 micron Millipore filtering apparatus used in the color determination. A chemical representation of the incineration of the colored sludge is as follows: Mg, (OH), _R = NfeO + COo + other gases (4-10) 11.5 ^ where R is the symbol for the color bodies Since the color bodies would be oxidized to CO2 and other gases by incineration, an investigation of the removal rate of color by incineration at varying times and temperatures was implemented. The

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104 TABLE 4-7 AVERAGE CHARACTERISTICS OF A SLUDGE PREPARED BY COAGULATING A NSSC WASTE AT pH 11 WITH 350 mg/1 OF Mg"""" AND 1250 mg/1 OF CaC0H)2 INITIAL COLOR OF NSSC WASTE = 4925 Parameter Value Suspended solids 1800 mg/1 Nonvolatile suspended solids 1082 rag/1 Volatile suspended solids 718 mg/1 % Magnesium in NVS 32.2% SVI 352

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105 primary intent was to find the minimum amount of time and minimum temperature required to remove the color bodies from the magnesium solids. However, a limiting factor influencing the recovery of the magnesium would be the chemical species formed during incineration. The specific gravity of magnesium varied considerably with the temperature and. time of incineration. This is presented in Table 4-8. A decrease in reactivity of MgO was paralleled by an increase in the density of MgO resulting from increasing calcination temperatures. According to Harper (1967) the freshly formed MgO had a high surface area. This area was reduced as heating temperatures were increased. The porosity of the oxide was reduced until, at a sufficiently high temperature, dead-burned magnesia resulted. The dead-burned magnesia resulted from compounds incinerated at temperatures in excess of 900C and was very unreactive. Harper (1967) also found MgO prepared in the range of 400-900. C, called caustic burned magnesia, was readily soluble in acid and rapidly hydrated in cold water. Standard Methods (1971) reports that it has been found that wastewater and effluent residues usually obtain constant weight after 15-20 minutes of ignition at 550 C. It was decided to ignite the sludge samples beginning at a temperature of 150 C and progressing to a final temperature of 850C in increments of 100 C. Magnesia incinerated at temperatures more than 900C would not have been soluble in a carbonated solution. Three different time increments were investigated at 15, 30 and 60 minutes. The 30 and 60 minute tests were discontinued for temperatures past 550 C. This was because the incineration of color was complete after 15 minutes at the high temperatures.

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106 TABLE 4-8 MgO REACTIVITY AS AFFECTED BY TEMPERATURE Temperature C Specific Gravity 600 2.94 700 3.04 850 3.22 1000 3.39 1200 3.48 1400 3.52 1500 3.56 Source : Kirk J E Othmer D F Encyclopedia of Chemical Technology Second Edition, 12, 'John Wiley and Sons, N.Y., 1967,

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107 The results o£ this experiment are presented in Table 4-9. Since the object of incineration was to remove color bodies in order to obtain reusable magnesium, a color/Mg ratio was calculated and is presented in Figure 4.33 as a function of incineration time and temperature. The incinerated sludge was resolubilized with hydrochloric acid. The color and magnesium concentration in the resolubilized sludge were then determined. These two parameters were divided to give a color/Mg ratio. A small ratio (less than one) would indicate that the magnesium was in a reusable form for color removal. Figure 4.33 shows the color/Mg ratio was minimized at 550 C for all incineration times tested. A ratio of 0.25 meant that over 98% of the color on the magnesium solids was removed by incineration and the magnesiiim could be recycled without a significant color concentration in the recycle stream. 4-5.6 Magnesium Recovery Magnesium recovery by recarbonation was selected as the recovery method because of the availability of free CO^ in flue gas, which usually is available at pulp and paper plants. The cost of recovering the magnesium with H2SO4 was $0.24/1000 gal. It was concluded, therefore, that H-SO. was too expensive to use as a means of recovery. In the laboratory study samples of nonvolatile solids were slurried in 200 ml of deionized water in a 400 ml beaker. The beaker was placed on top of a magnetic stirrer and the slurry was agitated by both the stirring bar and a 10% CO2-90% air gaseous mixture. The

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108 TABLE 4-9 INCINERATED SOLIDS ANALYSIS Incineration Temperature Incineration Time Minutes % NVS % VS % Magnesium of NVS Color Pt-Co Cone. NVS mg/1 150 15 30. 60 99.12 97.51 99.02 0.88 2.49 0.98 15.4 15.7 15.5 1157 1419 1325 378.8 379.2 390.0 250 15 30 60 92.72 90.77 87.06 12.94 9.23 7.28 16.5 16.9 17.6 1118 793 368 320.4 280.0 296.3 350 15 30 60 85.17 85.64 81.22 14.83 14.36 18.78 18.0 17.9 18.8 402 357.4 303.5 309.4 285.1 209.3 450 15 30 60 62.59 59.12 57.24 37.41 40.88 42.76 24.4 25.9 26.70 197.0 39.3 26.2 61.4 244.3 282.7 550 15 30 60 54.94 53.28 55.31 45.06 46.72 44.69 27.80 28.70 27.70 13.1 21.9 11.0 264.3 235.3 259.6 650 15 50.72 49.28 25.90 24.05 223.8 750 15 47.43 52.57 36.67 41.52 705.25 850 15 42.13 57.87 39.70 48.07 307.3

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109 AT o 15 min. • 30 min. D 60 min. at= Incineration time 350 450** 550'' 650* Incineration temperature {C) F\q. 4.33 Coior/Mg'^ ratio as a function of incineration temperature

PAGE 123

110 gaseous mixture was introduced into the slurry through a small porous stone diffuser. The 10% CO -90% air gaseous mixture was controlled by rotameters connected to a 100% CO tank and a laboratory air source. The total flow of the gaseous mixture was 2122 ml/min. This was the maximum flow rate that could have been implemented without losing some of the slurry due to turbulance. The pH of the slurry was continually monitored for the total recarbonation period. The temperature of the slurry was approximately 20C. As illustrated in section 4-5.4, the amount of magnesium resolubilized from the incinerated solids by carbonation at equilibrium was controlled by the concentration of 'CO2 gas, not the amount of CO gas. However, the amount of CO^ gas could affect the rate of the dissolution reaction due to surface area contact. A decrease in the temperature of the. carbonation reaction or a decrease in the CO pressure would decrease the soluble magnesium. This can be seen from the thermodynamic expression for the equilibrium constant and the chemical expression of activities for the equilibrium constant. log K = AG 2.303 RT (4-11) Because log K for MgC02"3H20(s) is negative for the recovery reaction, a decrease in temperature would decrease the amount of magnesium in solution at equilibrium. An increase in the CO pressure would increase the CO activity which would increase the amount of ++ Mg and HCO^at equilibrium. The temperature effects on the rate of MgCO^-SH^OCs) precipitation are illustrated in Figure 4.34 in data

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Ill 20 80 100 40 60 Time minutes Fig. 4.34 Precipitation of IV1gC03-3H20 by aeration at various temperatures. Source Blaci<,A.P, EPA #12120 HMZ, Sept., 1 974. "Full Scale Studies of the Magnesium Carbonate Water Treatment Process at Montgomery, Alabama and Melbourne, Florida.

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112 gathered by Black and coworkers in Dayton, Ohio and Montgomery, Alabama. The effect of increasing CO2 pressure is illustrated in the predominance diagram presented earlier in Figure 4.32. The magnesium recovery data are presented in Table 4-10. In each recovery experiment, the time of carbonation was 120 minutes. The pH readings and magnesium concentrations were measured at the time intervals of 0, 5, 15, 30, 45, 60, 90 and 120 minutes. The nonvolatile solids content of the slurry was varied from 5318 mg/1 to 119,692 mg/1. The plot of pH verses carbonation time in Figure 4.35 indicates that an equilibrium pH was approached in the laboratory studies. The ++ actual pH was close to that predicted by the Mg -CO2-H2O system m equilibrium wit}i a 10% CO2 gaseous stream. In this system the solid formed at equilibrium was MgC0„-3H20fs) The theoretically predicted equilibrium pH was 7.6 after a carbonation time of 120 minutes for the laboratory studies. The impurities present and the lack of a true equilibrium condition could have accounted for the difference in pH values. The agreement between the equilibriiMi pH and the actual pH supports the formation of a magnesium carbonate compound in the recovery process. In Figure 4.36 the magnesium concentration after the carbonation of nonvolatile solids is shown as a function of .carbonation time. Figure 4.37 presents the degree of magnesitmi recovery during carbonation as a function of the nonvolatile solids concentration. The data in Figure 4.36 indicate that if the nonvolatile solids concentration was increased, more magnesium was recovered during carbonation. However, the degree of the magnesium recovered from the

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113 TABLE 4-10 CARBONATION OF INCINERATED SLUDGE AT VARYING CONCENTRATIONS OF NONVOLATILE SOLIDS FOR MAGNESIUM RECOVERY Carbonation Test Nonvolatile Solids mg/1 5,318 22,950 33,561 74,550 119,692 Total Available Mg"^"^ mg/1 1,300 5,000 7,317 16,252 26,095 ++ Final Mg Concentration mg/l 1,300 4,803 5,724 9,425 12,776 Final Color Pt-Co 234 2,480 4,368 7,862 6,552 Color/mg Mg""* Ratio 0.18 0.56 0,60 0.48 0.25

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114 Mg"" mg/l •300 5000 7317 16,252 26,095 25 SO 75 100 Carbonation time (minutes) Fig. 4.35 pH as a function of carbonation at various nonvolatile solids concentrations 125

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115 14,000 12,000 E o 10,000 c o c > o 8000 6000 E Jo §> 4000 o 2000 NVS-mg/ Mg-mg/l • 5318 1300 22,950 5000 33,561 7317 n 74,550 16,252 1 1 9,692 26,095 50 75 Carbonation time (minutes) Fig. 4.36 Magnesium recovered as a function of carbonation time

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116 Timeminutes o 120 2.5 m 7.5 lao Non-volatile solids mg/l X 10""* Fig. 4.37 % Magnesium recovery as a function of nonvolatile solids concentration 12.5

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117 nonvolatile solids was reduced as the nonvolatile solids concentration was increased. After 120 minutes of carbonation, the soluble magnesium concentration increased from 1300 to 12,776 mg/1 when the nonvolatile solids concentration was increased from 5318 to 119,692 mg/1. The degree of magnesium recovery decreased from 100 to 49 per cent for these two experiments. The magnesium concentration after carbonation was never limited by the formation of a solid magnesium compound for the carbonation times tested. This would have been apparent if the magnesium concentration had become constant during carbonation and some magnesivun still remained in the nonvolatile solids. As shown in Figure 4.36, the magnesium concentration was always increasing during the carbonation process if there was any magnesium remaining in the nonvolatile solids. The maximum theoretical magnesiiim concentration at equilibrium is 866 mg/1 if MgC02-3H20(s) was the controlling solid phase. This was the controlling solid phase predicted by the theoretical Mg*"'-C02-H20 system presented in Figure 4.32. The controlling solid phase was not determined by these experiments. It might have been possible for the system to become supersaturated with respect to MgCO,'3H^OCs) since the solid magnesium compound used for carbonation was not MgC02-3H20(s) Supersaturation could occur if the dissolution rate of MgO(s) or MgC0H)2(s) was kinetically favored with respect to niicleation of MgCO^-SH 0(s) Another possibility was that other species in the incinerated solids were involved in the controlling solid phase and equilibrium was not achieved during carbonation.

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118 4-5.7 Magnesium Reuse Incineration and carbonation processes were used to recover the magnesium used in the color removal process. Each of these processes affected the chemical form of the magnesium. A magnesium reuse experiment was performed in order to determine if the recovered magnesium could be successfully reused in the coagulation process. The magnesium was recycled twice. Following each use, the magnesium was recovered by incineration at 550C for 15 minutes and then carbonated with a 10% C0„ gas. The nonvolatile solids concentration was 5318 mg/1 during carbonation. Some of the nonvolatile solids were not dissolved during carbonation. These remaining solids were recycled with each magnesium reuse. The magnesium dose was based on the soluble magnesium in the carbonated liquor and the coagulant dose relationship shown in Figure 4.8. The data from these magnesium reuse experiments are summarized in Table 4-11. After incineration the sludge remaining after the first use of the magnesium contained 99.6% of the initial magnesium. However, only 66% of the magnesium was in a soluble form after the sludge from the first use was carbonated. The color of this solution was 600. The remaining magnesium was still in the nonvolatile solids that were not dissolved during carbonation. When the sludge from the second magnesium use was carbonated, 92% of the magnesium was in a soluble form. The color of the recovered liquor was 80. The difference between the per cent of magnesiumsolubilized after the first and second carbonation processes

PAGE 132

119 TABLE 4-11 COLOR REMOVAL BY LIME -MAGNESIUM COAGULATION .USING THE SAME MAGNESIUM THREE TIMES. MAGNESIUM RECOVERY WAS ACCOMPLISHED BY INCINERATION AND CARBONATION FOLLOWING COAGULATION. NONVOLATILE SOLIDS NOT DISSOLVED BY CARBONATION WERE RECYCLED WITH THE RECOVERED MAGNESIUM. Use of Per cent o£ Per cent o£ Magnesium Original Magnesium Color Removed In Soluble Form First 100 92 Second '66 -92 Third 92 91

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120 was probably due to the incomplete combustion o£ the color bodies after the first magnesium use. The incinerated sludge contained 77% nonvolatile solids after the first magnesium use. The average nonvolatile solids reduction in the previous experiments reported in Table 4-9 for 550C and 15 minutes was 55%. The solids loading rate for the incineration process was never optimized. Practically all (99.6%) of the magnesium had been recovered after it had been used in the color removal process. Greater than 90% of the initial color was removed with each of the three magnesium uses. The magnesium dose was the same for each use. The recovered magnesium and the undissolved solids removed as much color in the second ard third uses as did the fresh magnesium in the first use. It was concluded that the magnesium could be successfully recovered and recycled in the lime -magnesium color removal process after the incineration and recovery processes. A second set of magnesiiim reuse experiments was conducted to determine the effectiveness of the recycled solids that were not dissolved in the carbonation process. The recovered magnesixim solution was filtered through a 0.80 micron Millipore filter to remove the solids. The per cent of color removed by the used magnesium with no solids was compared to the per cent of color removed by an equivalent amount of unused magnesium. The data from this experiment are presented in Table 4-12. The reused magnesium that contained no suspended solids removed 13% less color than did an equivalent amount of an unused magnesium when both were used separately in the color removal process. The

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121 TABLE 4-12 COLOR REMOVAL BY LIME-MAGNESIUM COAGULATION USING THE SAME MAGNESIUM TWICE. MAGNESIUM RECOVERY WAS ACCOMPLISHED BY INCINERATION, CARONATION AND FILTRATION FOLLOWING COAGULATION. NO SUSPENDED SOLIDS WERE RECYCLED WITH THE RECOVERED MAGNESIUM. INITIAL COLOR = 5000. Use of Magnesium Final % Color Magnesium Dose mg/1 Color Reduction First 300 550 8d Second 300 1190 76 Second 343 1096 78 Second 395 943 81

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122 presence of the incinerated solids was a significant aid to the reused magnesium in the color removal process. These solids probably provided nucleating surfaces for the forming solids phase.

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CHAPTER 5 DESIGN OF A COLOR REMOVAL PROCESS FOR A NSSC WASTE USING MAGNESIUM COAGULATION AND RECOVERY The color removal process consists o£ several different unit operations. These are coagulation, sedimentat^ion, vacuum filtration, incineration and carbonation. The design of each of these unit operations was considered separately in this chapter. The design parameters were based on the treatment of a NSSC waste with an initial color of 5000 and a flow of 1 mgd. 5-1 Coagulation A velocity gradient of 1000 sec"! was recommended by the AWWA (1969) to achieve adequate coagulant dispersion during the rapid mix process in a contact time of 20 seconds. The tank volume and energy requirements were determined from Equation 5-1. G = h 550 P Vu (5-1) where: G = velocity gradient sec P = water horsepower hp 3 V = tank volume ft 2 • u = viscosity lb sec/ft The tank volume and energy required were 30 ft^ and 1.5 horsepower. 123

PAGE 137

124 The amount of magnesium and CaCOH)™ to treat the waste was determined from the optimum dose equations presented earlier in Figures 4.7 and 4.8. The CaC0H)2 requirement was 10,425 Ibs/mgal, and the magnesium requirement was 2502 Ibs/mgal. S-2 Sedimentation The information required to develop basic design criteria for a secondary sedimentation tank was obtained from the batch settling tests and Equations 5-2, 5-3 and 5-4. A^ = Q/Vj^ C5-2) 2 A = Area required for clarification ft V,= Hindered settling velocity fpm Q = Flow cfm At = Qt^^/Ho (5-3) 2 A. = Area required for thickening ft t^^ = Time required to reach desired concentration min Hq = Initial height of sludge interface ft HqCo = HuCu C5-4) Cq = Initial solids concentration mg/1 Hu = Final height of sludge interface ft Cy = Desired solids concentration mg/1

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125 The settling velocity was determined to be 0.187 £t/rain from the hindered portion of the settling curve in Figure 5.1. The initial and final heights of the sludge interface were 1.12 ft and 0.20 ft. Equation 5-2 gave 610 ft^ as the area required for clarification. The area required for thickening the sludge to a final concentration of 11,520 mg/1 was found from Equations 5-3 and 5-4 to be 1020 ft^. The area required for thickening the sludge to a final concentration of 8000 mg/1 was larger than the area required for clarification and therefore controlled the area for settling. A design area of 1020 ft would provide adequate. settling area for the treatment of any NSSC waste with an initial color of 5000 or less. This was because the settling area was determined by the amount of solids produced during treatment and a smaller amount of solids would be produced from the treatment of a weaker waste. The solids loading that would be passed onto the vacuum filter would be 15,012 Ibs/mgal. If carbonation were used to recover the magnesium, the calcium present in the Ca(0H)2 would precipitate as CaC02 during coagulation. This would increase the mass of solids passed onto the vacuum filter. Since 1250 mg/1 Ca(OH3 2 was used to adjust the waste to pH 11, 16.9 mM of Ca** was available to precipitate as CaCO^. The magnesium concentration in the recycle stream was 1300 mg/1. To achieve the required magnesium dose of 300 rag/1, the recirculation ratio of the waste stream to the recycle stream would have to equal 0.23. The total carbonate concentration in the recycle stream was not determined. For design purposes, all of the calcium from the Ca(0H)2 dose was assumed to precipitate as CaCO,.

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126 JO E .^ 600 S 500 C Initial Concentrations Sus. solids = 1800 nng/l Color = 5000 Mg*= 350 mg/l Lime= i?.50 mg/l Polymer = 3.0 mg/l Hydrolyzed Polyacryiamide (Anionic) 10 20 Settling time minutes Fig. 5.1 Design data for sedimentation

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127 This assumption is theoretically valid because electroneutrality requires that two moles of HCO," are available for every one mole of Mg'^'*' solubilized in the recovery process. Approximately 107 mM of HCO^" would be available in the recycle stream. In the coagulation tank at pH 11, this would be converted to CO3 and precipitate approximately all of the calcium as CaCO Approximately 16.9 mM of CaCO, or 1690 mg/1 would be added to the solids passed onto the vacuum filter. Because the CaCOj formed acts as a settling aid, no additional allowances were made in the settling area calculations. The solids loading data are summarized in Table 5-1. The Ca(OH32 dose in Table 5-1 is the design dose determined for color removal from a NSSC waste by magnesium coagulation. The solids content of the sludge without the CaCOwas determined in Chapter 4. The total solids represented the sum of the solids from the precipitation of CaCO,. The volume of sludge coming from the thickening operation was determined for a color of 5000, and was 175 ml per liter of waste treated. The overflow rate from the sedimentation basin would be 1034 gpd/ft^. The per cent solids in the settled sludge including the CaCOprecipitate would be 1.9.9%. 5-3 Vacuum Filtration Vacuum filtration studies on magnesium sludge were performed at Melbourne, Florida by Black (1974) Sludge was produced by treating a surface water source for color removal. The solids content in the sludge after vacuum filtration was 45%. In those studies.

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128 TABLE 5-1 SOLIDS LOADING FROM SETTLING BASIN Color Pt-Co 5000 Lime dose mg/1 1250 Ca"*"*" mM from lime 16.9 CaC03 mM precipitated 16.9 CaCOj mg/1 1690 Solids mg/1 1800 Total solids mg/1 3490 Ibs/mgal 29,107

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129 2 filtration rates ranged from 11.7 to 20.0 lbs/ft /hr. The sludge produced after NSSC waste treatment would have a lower CaCO^ concentration than the Melbourne sludge, thereby producing a lower filtration rate, Liptak (1974) found that compounds precipitated by 2 Ca(0H)2 can be filtered at a rate of 2 to 6 lbs/ft /hr on a rotary • vacuum filter and the solids content in such a filter cake would vary from 20 to 30%. A design rate of 6 Ibs/ft^/hr and 20% solids content was selected for vacuum filtration of the NSSC sludge following coagulation. Based on a design rate of 6 lbs/ft /hr, the total amount of solids processed would be 144 lbs/ ft /day. For each million gallons of waste treated, 29,107 lbs of solid has to be vacuum filtered. The 2 total area required for vacuum filtration was 202 ft A second vacuum filter would be required if recalcination was used to recover the Ca(0H)2. The solids filtered would be the CaCO^ solids produced in the coagulation basin from the reaction between Ca(0H)2 and the carbonated recycled stream. These solids would amount to 14,094 Ibs/mgal. A design rate of 12 Ibs/ft^/hr was used to determine the size of the vacuum filter for the CaCO, solids. The 2 area required to remove the CaCO^ before lime recovery was 49 ft 5-4 Incineration To estimate the energy requirement of the incineration process, it was necessary to consider both the sensible and latent heat requirements of the NSSC sludge. Lignin sludge has a 8,000 to 10,000 Btu/lb fuel value and contains 70% volatile solids. The fuel value

PAGE 143

130 of the NSSC sludge was estimated at 5140 Btu/lb based on a 45% volatile solids content in the sludge. The specific heat of the solids was estimated to be 0.75 Btu/lb F from Liptak (1974) The sensible heat requirement was determined to be 18,540,516 Btu from the specific heat, material weight and 550C incineration temperature of the sludge. A latent heat requirement of 139,483,500 Btu was determined from the amount of v/ater that had to be evaporated during the incineration process. The amount of available energy was determined from the fuel value of the sludge to be 147,151,500 Btu. The total heat requirement exceeded the available heat by 7,688,000 Btu. This heat would have to be supplied ^ly use of a fuel such as oil or natural gas. 5-5 Carbonation The operating conditions of the carbonation process were selected from the recovery experiments described in Chapter 4. These operating conditions could be further optimized to increase efficiency. In the recovery experiments, 100% magnesium recovery was only obtained after 0.957 grams of incinerated solids was carbonated in 180 mis for 45 minutes with a 10% C02-90% air gaseous mixture flowing at 2122 ml/min. The solids from the recovery experiments were increased by the theoretical amount of CaCO_ that would be added to the sludge when carbonation would be used to recover the magnesium. The additional CaCO_ would probably not affect the amount of magnesium recovered since it would not enter into the Mg"^'^-H20-C02 equilibrium system.

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131 The following design parameters were developed for a 1 mgd flow into the color removal process using the operating conditions that were previously selected. The daily total solids including the additional CaCO, input to the carbonation tank would be 22,351 lbs. Since a recirculation ratio of 0.23 was determined in section 5-2, 230,000 gallons of carbonated recovery liquor would be produced per day. The recovery process required approximately 161,435 lbs of CO2 per day. The most economical source of CO2 gas is flue gas. The flow rate of the flue gas into the carbonation tank would be 9167 cfm if the flue gas contained 10% CO^. A design summary is presented for the lime-magnesium color removal process in Table 5-2, and a flow chart of this process is presented in Figure 5.2.

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132 TABLE 5-2 DESIGN SUMMARY FOR THE TREATMENT OF A NSSC WASTE WITH AN INITIAL COLOR OF 5000 AND A FLOW OF 1 MGD BY THE LIME -MAGNESIUM COLOR REMOVAL PROCESS Unit Process Design Parameter Comments Rapid Mix 20 sec G = 1000 sec -1 Coagulation and Sedimentation 10,425 lbs Ca(0H)2 2502 lbs Mg"*"* 25 lbs 835A 1020 ft^ 29,107 lbs solids 5% hydrolyzed polyacrylamide polymer Sludge thickened to 1.99% Vacuum Filtration 6 lbs/£t /hr 202 ft^ 12 Ibs/ft^/hr 49 ft^ Magnesium sludge CaCOj sludge Incineration T = 550"C 22,298 lbs NVS remaining Carbonation 9167 cfm 161,435 lbs CO2 Flue gas at 10% CO2 ++ 100% Mg recovery

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133 o o o .2 S £o o OO u a> f f /\ CM o_ o ^ IS 6 /\ r~\ ^ o a> c D O o o c o c CO £ 2 >> CL a o ^ o = < PJ TTs ^ tn "o CO T^ £ 8 3 a > LiCM O en O a' ^ CM o u ss o M— O t> CP to _3 .c CM en CM /\ c o c: o 5VI J3 in ro OJ" CM ^ ^ _\/ c _o o c o o o J cc y\ cc: ^ Pfs: a X^ If) 5 9 o o o in \/ a f e 3 Jn 0) g" > 3 O O TD ^ < o ^?cyLcD 5 .2 > 0) ro "^ .-= Si> tr CM OT3 >u o o > o E o u E 3 "55 V c o> o E E E o k. o o o CM in ii. e o o

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CHAPTER 6 COST 6-1 Chemical Costs As o£ June 1976, magnesium sulfate was the most economical source of magnesium found by the suthor, and was available commercially at $120/ton. The cost of 300 mg/1 of Magnesium for treating a NSSC waste with an initial color of 5000 was $0,743/1000 gal. Lime was available commercially at $50/ton as of June 1976. The cost of 1250 mg/1 of lime to treat a NSSC waste with an initial color of 5000 was $0,260/1000 gal. The cost of the polymer 835A from American Cyanamid was $1.50/lb as of June 1976. The polymer dose was 3 mg/1 and added $0,037/1000 gal to the cost of the lime -magnesium color removal process for NSSC waste. The CO2 gas was free since it would be taken from flue gas which is abundant at pulp and paper mills. The summarized chemical cost is presented in Table 6-1 and assumes no chemical recovery. 6-2 Capital and Operation Costs All capital and operating costs were estimated from Liptak (1974) unless otherwise specified. The area required for settling treated 2 waste was 1020 ft per mgd of incoming waste. The capital cost of a settling basin was $18/ft^. The total capital cost for a 1, 5 and 10 134

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135 TABLE 6-1 CHEMICAL COST TO TREAT A NSSC WASTE WITH AN INITIAL COLOR OF 5000 Chemical $/1000 gal MgSO^ 0.743 Ca(OH32 0.260 835A 0.037 Total 1 040

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136 mgd waste flow would be $18,360, $91,800 and $183,600 respectively. The operating cost o£ vacuum filtration without heat treatment was estimated at $4/ton of dry solids for the sludge produced in the color removal process. This cost was $0,058/1000 gal on a unit flow basis and constant for any plant flow. The capital costs for DorrOliver vacuum filters to concentrate the solids resulting from the treatment of a 1, 5 and 10 mgd flow were $100,000, $200,000 and $300,000 respectively. Approximately 3.4 tons of solids would be lost during incineration for every million gallons of waste treated. The capital costs of equipment capable of incinerating the solids from treating a 1, 5 and 10 mgd flow was estimated to be $455,000, $700,000 and $750,000 respectively. The operating cost of the incineration process decreased as the amount of material processed increased. Approximately 3.4, 17.0 and 34.0 tons of solids would be incinerated when respective waste flows of 1, 5 and 10 mgd were treated. The cost per ton of dry solids incinerated was estimated to be $8.52, $3.80 and $3.50. The resulting unit costs per 1000 gallons of treated waste were calculated to be $0,029, $0,013 and $0,012. The lime could be recovered from the carbonation process by passing the remaining slurry onto a vacuiun filter and then to a lime kiln. The CaCO_ would not dissolve in the carbonated slurry and would be available for recalcination. Only CaCO-r solids would remain after carbonation. After carbonation 14,094 lbs of CaCO^ sludge would be vacuum filtered per million gallons of treated waste. The filtration rate of the CaCO^ sludge was estimated to be twice that

PAGE 150

137 o£ the sludge produced in the coagulation reaction. Vacuum filtration o£ this sludge would cost $2.50/ton or $0,018/1000 gal on a unit flow basis. The costs for the capital equipment required to vacuum filter the CaCO„ sludge produced by treating a 1, 5 and 10 mgd waste flow were estimated to be $40,000, $80,000 and $150,000 respectively. The capital cost for the magnesium recovery phase was estimated by the author to be $2,500, $7,500 and $10,000 for waste flows of 1, 5 and 10 mgd respectively. No estimate, was made for the operating cost of the magnesium recovery phase. These costs were covered in a miscellaneous estimate that -will be discussed later. The cost for recalcination at a 1 mgd plant was • estimated at $0.07/1000 gal for operating cost and $200,000 capital cost. A 5 mgd recalcination plant was estimated to cost $500,000 for capital expenditures and $0.05/1000 gal for operating cost. The estimate for a 10 mgd recalcination plant was $600,000 for capital expenditures and $0,035/1000 gal for operating cost. A final miscellaneous cost was estimated for the capital and operating costs for all other equipment necessary to install the color removal process. For a 1, 5 and 10 mgd effluent, the miscellaneous capital and operating costs respectively were $40,000 and $0,030/1000 gal, $100,000 and $0,020/1000 gal, and $150,000 and $0,015/1000 gal. 6-3 System Costs In this section of Chapter 6 the appropriate cost of .different

PAGE 151

138 color removal systems is presented. A complete system designates a color removal scheme using both Ca(OH) and magnesium recovery in the color removal process. All of the unit operations listed in Table 6-2 were used in the complete system as represented on the process flow chart in Figure 5.2. The other color removal systems differed only in the degree of chemical recovery. The costs of all systems considered are presented in Table 6-3. If Ca(QH32 recovery was eliminated, the costs associated with a vacuum filter for CaCO^ sludge and a lime kiln would be eliminated. If magnesium recovery was eliminated, the cost of vacuum filtration, incineration and carbonation would be eliminated. However, the chemical costs associated with both of these systems would be increased. Complete chemical recovery was assumed in the cost calculations. The separate entries in Table 6-2 represent each unit operation used in the color removal process. The capital cost was calculated at 8% interest compounded annually for 25 years. The unit cost was determined by dividing the amount of waste treated in a 25 year period into the capital cost. The unit cost per 1000 gallons of treated waste for each $100,000 of capital equipment required to treat a 1 mgd, 5 mgd and 10 mgd flow was $0.0761, $0.0152 and $0.0076. For a color removal process treating 1 mgd, the process cost siammary presented in Table 6-3 indicates that the most economical system required only magnesium recovery. The unit cost was $0,094 per 1000 gallons. However, at 5 and 10 mgd, a process employing both magnesium and Ca(OH) recovery was the most economical system.

PAGE 152

139 O O O T3 e s 03 O bO O o o CM I M CQ O o o o o Eo CM o o r-H o o 1— H O o r— 1 (—1 o 1 — i o o o o o o o o o o o o • O o o o o o o o o o o o o o o o o o 00 f-H o o to o LO o I— 1 o LO (— 1 o o o LO 1— 1 I— H feO^T o r-r-1 IM \U LO lO t— 1 lo o O r-H t^ 1 — 1 LO o o r-H o O o o OJ o o o o O o o o o o CO o o o o o o o o LO o o o c o o o o o o o ^0 1-H o o o o r-o CO o o LO o o i-H 1 — 1 I— ( o o to o o o to o LO 1— 1 o to o o LO o o o o o o o o CO LO J o r-i o 00 1— 1 o LO to o LO 1— 1 o to r-H o o o o o o oo to CO O o LO r-H r— I LO CN 1 O O I O O O o o o o o LO o o o o o o o o ^ o o o o o o \D to o o LO o o o CO CO o LO CN o o o LO I— t o LO 'go <* LO >— 1 ^ (Ni CO to to o O y U 03 (3 u 1 u s (=: fi 1 o o o o •ri •H •H •H -y +-> o J-1 J-i Ri cd !/5 oj Cd tfl u c ^H 3 !h c !4 3 +-> o +-) o +-> o P o 1—1 •H r-H (D p i—H •H r— ( 3 (U o p: -H 3 c > 3 CD O P 3 • H O 3 0) o H o P 3 -H O p 0) o E-. 1—1 P o o o o e !/l •H P O O u o B tfl 03 CD Oj c 1— 1 Qi > hJ s CO > 1 — 1 oi > hJ s & 0) ni a, u o 00 LO 1 o en CN O 00 t— 1 1 o o o o to o LO CM o o o o o o

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140 TABLE 6-3 PROCESS COST SUMMARY IN $/1000 GALLONS OF NSSC WASTE WITH AN INITIAL COLOR OF 5000 Costs Flow mgd Complete System No CaC0H)2 Recovery No Magnesium Recovery No Ca(0H)2 or Magnesium Recovery Capital 1 0.650 0.468 0.226 0.044 5 0.255 0.167 0.117 0.029 10 0.163 0.106 0.082 0.025 Operating 1 0.205 0.117 0.118 0.030 5 0.159 0.091 0.088 0.020 10 0.137 0.084 0.067 0.015 Chemical 1 0.037 0.297 0.780 1.040 5 0.037 0.297 0.780 1.040 10 0.037 0.297 0.780 1.040 Total 1 0.892 0.882 1.124 1.114 5 0.451 0.555 0.985 1.089 10 0.337 0.487 0.929 1.080

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141 The cost estimate for 5 and 10 mgd was $0,445/1000 gai and $0,322/1000 gal respectively. In the color removal process, both a decreasing effluent color or an increasing plant flow will decrease unit cost. Most NSSC plant effluents range between 5 and' 10 mgd with colors varying from 2500 to 5000. A cost estimate was made for the treatment of a NSSC waste with an initial color of 2500 for flows of 5 mgd and 10 mgd. These cost estimates were determined in the identical manner as those determined for the stronger NSSC waste. They were based on the decreased solids loading produced from treating a weaker waste. These calculations are summarized in Table 6-4, As can be '=:een from the summaries presented in Tables 6-3 and 6-4, the unit cost in $/1000 gal will decrease with increasing plant flow and decreasing waste color. The unit cost estimates in $/1000 gal for treating a 5 and 10 mgd flow of NSSC waste with an initial color of 2500 were calculated to be 0.371 and 0.273 $/1000 gal. When the color was assumed to be 5000, these estimates increased to 0.451 and 0.337 $/1000 gal respectively. Comparison of these cost estimates indicates that flow has a greater effect on cost than does waste color. This implies that diluting a highly colored waste is not cost effective. Color streams should be concentrated whenever possible for the most economical treatment. A more economical unit cost was obtained from treating a 10 mgd waste at a color of 5000 than a 5 mgd waste at a color of 2500. The most economical unit cost was, however, for the higher flow at the lower color. Approximately 20,000 gallons of NSSC waste are discharged for

PAGE 155

142 Pi o (-> O !-> P O P o 0) r-l •H i-H 3 3 3 •H o 3 (D O •H P 3 • H o 3 03 H iH OJ OS c 1 — 1 a > J ^ f-l CO > I-H a > ,-4 S e & PL, (D
PAGE 156

143 every ton of NSSC product produced. The approximate cost to manufacture a ton of NSSC product is $200. The per cent cost increase of NSSC product if the magnesium color removal process was employed is given in Table 6-5.

PAGE 157

144 TABLE 6-5 NSSC PRODUCT COST INCREASE DUE TO COLOR REMOVAL BY MAGNESIUM COAGULATION

PAGE 158

CHAPTER 7 CONCLUSIONS AND RECOMMENDATIONS 7-1 Conclusions A process for removing in excess of 90% of the initial color in a NSSC waste has been developed based on laboratory experiments. The lime-magnesium color removal process mainly involves color removal by coagulation, .incineration of the color bodies and solubilizing the magnesium solids with protons from dissolved CO2 gas. The untreated NSSC waste was shown to have a significant acid strength. The amount of magnesium and Cam) 2 ^""^^^ ^ ^^^ ^^^^ waste in order to remove 90% of the color was directly proportional to the initial NSSC waste color. Increasing coagulant dose did not shift the optimum coagulant pH. When Ca(0H)2 was used for pH control instead of NaOH, the amount of magnesium required for 90% color removal from the NSSC waste was significantly decreased. Color removal achieved by magnesium precipitation in pH dependent and in pH independent chemical reactions demonstrated that magnesium was responsible for color removal in the lime -magnesium color removal process. The color bodies and the magnesium buffered the NSSC waste during the coagulation process. The magnesium ions first chelated the organic acids and increased the color, but eventually formed an insoluble organo-metallic precipitate that removed 90% ofthe initial 145

PAGE 159

146 color. This was demonstrated by the increase in residual color for magnesium additions of 10-50 mg/1 to the waste solution. The color removal reaction proceeded by accepting the hydroxides from Ca(0H)2 or NaOH. A precipitate was eventually produced that removed organic carbon, color and the magnesium from the NSSC waste. A 90% reduction in the initial color of the waste was accompanied by a 34% reduction in the organic carbon concentration. Approximately a 40% decrease of the acid strength of the untreated NSSC waste occurred during the color removal by magnesium coagulation. The jiajority of the reduction in acid strength occurred in the weak acid range, pK = 9, which suggested that the weak acids were responsible for most of the color in the NSSC waste. The color was removed from the NSSC waste by a chemical reaction that involved color bodies, hydroxide and magnesium ions which resulted in the formation of an insoluble precipitate. The empirical formula of the precipitate was MgCOH). R, where R represented the organic color bodies. Color removal from a NSSC waste was achieved by magnesium precipitation in a pH dependent or a pH independent chemical reaction. Ca(OH)^ alone at the coagulation pH of the lime-magnesium color removal process did not remove any color from the NSSC waste. It was concluded that magnesium was responsible for color removal in the lime-magnesium color removal process. The sedimentation of the anionic colored floe was greatly assisted by the addition of an anionic polymer. The settling of the colored floe particles was controlled by polymer-floc bridging when the anionic polymer was present, not by electrostatic repulsion. Cationic polymers did not aid the settling of the colored floe.

PAGE 160

147 The color removal process was reversible. Removal of the color bodies from the sludge hy incineration was required in order to reuse the magnesium for color removal. The optimum temperature of incineration was 550C. Increasing the incineration temperature from 550C to 850C did not remove any more color from the magnesium, but did remove some additional solids. The recovery experiments demonstrated that magnesium could be successfully reused in the lime -magnesium color removal process. In order to reuse the magnesium it was necessary^ to remove the color from the sludge by incineration and to dissolve the incinerated magnesium by carbonation. Approximately 93% of the original magnesium was recovered following three uses of the same magnesium in the limemagnesiiim color removal process. There was no difference in the color reduction achieved by unused and recovered magnesium when the recovered magnesium was recycled with the solids remaining after carbonation. The unit cost of treating a NSSC waste by magnesium coagulation decreased as the volume of waste increased and the initial color of the waste decreased. However, because of the high capital expense of the equipment involved, the unit cost was more sensitive to the volume of waste treated than to the initial color of the waste. 7-2 Recommendations Further research needs to be conducted on possible polymers that can serve as settling aids. Activated silica and starch are two anionic polymers that are more economical and may function as well as the partially hydrolyzed polyacrylamides. The NSSC waste should be

PAGE 161

148 fractionated by gel filtration before and after magnesium coagulation to determine what fractions were removed during treatment. The fractions and functional' groups that contributed most to the color should be determined for the NSSC waste. Different pulp wastes could be treated by magnesium coagulation to determine if the color removal process can be adaptable to other pulp wastes. Possible areas of investigation for additional research would be the structure of the Mg(OH) (s) colloids and the rate of formation of Mg(0H3"^. Optimum rates of energy input into the rapid mix process should be determined using magnesium coagulation in different environments. The fuel value of the sludge produced in the coagulation process should be determined and necessary experiments conducted to design an incinerator for the color removal process. The solids species controlling magnesium solubility in the carbonation process should be identified and the optimiun conditions for coagulant recycle determined. The lime-magnesium color removal process should be studied on a pilot plant scale. The design and operational parameters for this process can only be determined from such a study. The legal requirements of the Water Quality Act and success of the lime-magnesiiom color removal process on a laboratory scale are good reasons to implement this study.

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REFERENCES Alder, E.; Lundquist, K.; Miksche, E. "The Structures and Reactivity of Lignin." Advances in Chemistry Series 59 R.F. Gould, ed. (1966). American Chemistry Society. "The Role of Adsorption of Hydrolyzed Aliiminum in the Kinetics of Coaglation-Adsorption from Aqueous Solutions." Advances in Chemistry Series 79(1968). American Water Works Association Committee Report. "State of the Art of Coagulation." JAWWA 63: 2: 99 (1971). Berger, H.F. U.S. Pat. 3,120,464 (Feb. 1964). Birkner, F.B.; Morgan, J.J. "Polymer Flocculation Kinetics of Dilute Colloidal Suspensions." JAWWA 60: 175 (Feb. 1968). Black, A. P. "Full Scale Studies of the Magnesium Carbonate Water Treatment Process at Montgomery, Alabama and Melbourne, Florida." EPA Project #12120 HMZ (Sept. 1974). Black, A. P. "Some Applications of the Principles of Colloidal Behavior to Water Treatment." Proc. Rudolphs Res. Conf Rutgers State University, New Brunswick, N.J. (June 1960). Black, A. P.; Birkner, F.B.; Morgan, J.J. "Destabilization of Dilute Clay Suspensions with Labeled Polymers." JAWWA, Tl^. 175 (Dec. 1965). Black, A. P.; Christman, R.F. "Characteristics of Colored Surface Waters." JAWWA 55_: 753 (1963a). Black, A. P.; Christman, R.F. "Chemical Characteristics of Fulvic Acids." JAWWA, 5£: 897 (1963b). Black, A. P.; Shuey, B.S.; Fleming, P.J. "Recovery of Calcium and Magnesium Values from Lime-Softening Sludges." JAWWA, 63_: 616 (Oct. 1971). Black, A. P.; Singley, J.E.; Whittle, G.P.; Maulding, J.S. "Stoichiometry of Coagulation of Color Causing Organic Compounds with Fe2(S0^)g." JAWWA 5^: 1347 (Oct. 1963). Black, A. P.; Willems, D.G. "Electrophoretic Studies of Coagulation for the Removal of Organic Color." JAWWA 53: 589 (1961). Canadian Pollution Abatement Research Program. "Color Removal from Biologically Treated Pulp and Paper Mill Effluents." CPAR Project Report 21 D-1 (1974). 149

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150 Christman, R.F.; Ghassemi, M. "Chemical Nature o£ Organic Color in Water." JAWWA 58_: 723 (1966). Collins, J.W.; Webb, A. A.; Didwania, H.P.; Lueck, B.F. "Components of Wood Pulp Effluents." Env. Sci. ^ Tech. Z: 371 (1969). Cornwell, D.A. "Recycling of Alum Used for Phosphate Removal in Domestic Wastewater Treatment." Ph.D. Dissertation, University of Florida (1975) Davis, C. "The Effect of Sodium -Ion Concentration on the Removal of Color from Kraft Linerboard Mill Effluents." Paper presented at the 20th anniversary meeting of Southeastern Tappi, Jacksonville, Florida (1972). Day, H.R.; Felbeck, G.T. "Production and Analyses of a Humic AcidLike Exudate from the Aquatic Fungus Aureobasidium pullulans ." JAWWA 66^: 8: 484 (Aug. 1964). Day, R.A., Jr.; Underwood; A.L. Quantitative Analysis Prentice-Hall, Englewood Cliffs, N.J. (1967). Dollimore, D.; Horridge, T.A. "The Optimum Flocculant Concentration for Effective Flocculation of China Clay in Aqueous Suspensions." Water Research Vol. b. Pergamon Press (1972). Domitar Limited Research Center. "Colour Removal from Biologically Treated Pulp and Paper Mill Effluents." CPAR Project Report 210-1 (Mar. 1974). Dubose, A. "The Effect of Magnesium Concentration on Municipal Wastes." Ph.D. Dissertation, University of Florida (1973). Eidsness, F.A.; Black, A. P. "Carbonation of Water Softening Plant Sludge." JAWWA 49^: 1343 (Oct. 1957). Fitzgerald, C.L.; Clemens, M.M.; Riley, P.B. "Coagulants for Waste Water Treatment." Chem. Eng. Progress 66 : 36 (Jan. 1970). Freudenberg, K. "Analytical and Biochemical Background of a Constitutional Scheme of Lignin." Lignin Structure and Reactions, Advances in Chemistry Series 59 R.F. Gould, ed. (1966). Gjessing, E.T.; Lee, G.F. "Fractionation of Organic Matter in Natural Waters on Sephadex Columns." Env. Sci. & Tech. 7_: 631 (1967). Gjessing, E;T.; Samdal, J.E. "Humic Substances in Water and the Effect of Impoundment." JAWWA, 6£: 451 (1968). Gould, M. "Color Removal from Kraft Mill Effluent by an Improved Lime Process." Tappi 56_: 79 (1973).

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151 Guilledge, J.H.; O'Conner, J.T. "Removal of Arsenic (V) from Water by Adsorption on AlLiminum and Ferric Hydroxides." JAWWA, 65_: 8: 548 (Aug. 1973). Hahn, H.N.; Stumra, w'. "Kinetics of Coagulation with Hydrolyzed Alum." Jour, of Coll. ^ Sci. 288 : 134 (Sept. 1968). Hall, E.S.; Packham, R.F. "Coagulation of Organic Color with Hydrolyzing Coagulants." JAWWA, 57: 1149 (July 1965). Hannah, S.A.; Cohen, J.M.; Robeck, G.G. "Measurement of Floe Strength by Particle Counting." JAWWA, 59: 843 (July 1967). Harper, F.C. "Effect of Calcination Temperature on the Properties of Magnesium Oxides for Use in Magnesium Oxychloride Cements." Jour, of Applied Chem. ]J: 1 (Jan. 1967). Herbet, A.J. "A Process for the Removal of Color from Bleached Kraft Effluents through Modification of the Chem-Recovery System." NCASI, Technical Bulletin No. 157 (June 1962). Interstate Paper Corp, "Color Removal from Kraft Pulping Effluent by Lime Addition." Water Pollution Control Research Series 12040 ENC (Dec. 1971). """ Jeffcoat, B.W.; Singley, J.E. "The Effect of Alum Concentration and Chemical Addition Times on Coagulation." JAWWA, 67_: 4: 177 (Apr. 1975). Jensen, W.; Fremer, K.E.; Forss, K. "The Separation of Components Spent Sulfite Liquor." Tappi 45: 122 (1964). Jobin, R.; Ghosh, M.M. "Effect of Buffer Intensity and Organic Matter on the Oxygenation of Ferrous Iron." JAWWA 6£: 7: 590 (Sept. 1972). Kabeya, H.; Fujii, T.; Kubo, T.; Kimura, Y.; Urano, K. "Renovation of Pulp Mill Waste Water Adsorption Characteristics of Kraft Pulp Lignin on Activated Carbon." Kanipa Gikyoshi (Jap.), 26_: 3: 125. Chem. Abs. 77_: 24541t (1972). Kawamura, S. "Coagulation Considerations." JAWWA 65_: 6: 417 (June 1973) Kirk, T.E.; Brown, W. ; Cowling E.B. "Preparative Fractionation of Lignin by Gel -Permeation Chromatography." Biopolymers 1_ (1969). Kirk, J.E.; Othmer, D.F. Encyclopedia of Chemical Technology 2nd Edition, 12_. John Wiley § Sons, New York (1967).

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152 Kotoh, S.; Kimura, Y. "Study of Renovation of Pulp Mill Waste Water by Treatment with Fly Ash." Jour. Jap. Tappi 25^: 4: 168 (1972). Abs. Bull. Inst. Paper Chem. 42: 8176 (1972). LaMer, V.K. "Coagulation vs. Flocculation of Colloidal Dispersions by High Polymers." Principles fi Applications of Water Chemistry. (4th Rudolphs Conf.). Faust, S.D.; Hunter, J.V. Principles and Applications of Water Chemistry Wiley Interscience (1967) Langelier, W.F.; Ludwig, N.F. "Mechanism of Flocculation in Clarification of Turbid Waters." JAWWA, 41: 163 (Feb. 1949). Leszczynski, C. "Decolorization of Kraft Mill Effluents." Prezegl. Papier (Pol.), 28_: 3: 88 (1972). Abs. Bull. Inst. Paper Chem. 43_: 4154 (1972). Luner, P.; Dence, C. "Mechanisms of Color Removal in the Treatment of Pulping and Bleaching Effluents with Lime." Vol. I. Treatment of Caustic Extraction Stage Bleaching Effluent Tech Bulletin 239 (July 1970). Luner, P.; Dence, C; Bennett, 0.; Ota, M. "The Mechanism of Color Removal in the Treatment of Pulp and Bleaching Effluents with Lime." NCASI, Technical Bulletin No. 242 (Dec. 1970). Mangravite," F.J., Jr.; Buzzell, T.D.; Cassell, E.A.; Matijevic, E.; Saxton, G.B. "Removal of Humic Acid by Coagulation and Microflotation." JAWWA, 6^: 288 (Feb. 1975). Matijevic, E.; Janauer, G.E.; Kerker, M. "Reversal of Charge of Lyphobic Colloids by Hydrolyzed Metal Ions." Jour, of Coll. ^ Sci. 19_: 333 (1964) Maulding, J.S.; Harris, R.H. "Effect of Ionic Environment and Temperature on Coagulation of Color-Causing Organic Compounds with Ferric Sulfate." JAWWA 60: 460 (1968). Michaels, A.S. "Aggregation of Suspensions by Polyelectrolytes." Ind. S Eng. Chem. 46 : 1485 (July 1954) Midwood, R.B.; Felbeck, G.T. "Analysis of Yellow Organic from Fresh Water." JAWWA 60: 352 (1968). Mohtaoi, M.F.; Roa, P.N. "Effect of Temperature on Flocculation of Aqueous Dispersions." Water Research 17 Pergamon Press (1973). Morrow, J.J..; Rausch, E.G. "Colloidal Destabilization with Cationic Polyelectrolytes as Affected by Velocity Gradients." JAWWA 66 : 11: 646 (Nov. 1974).

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153 Narkis H • Rebhum. "The Mechanism of Flocculation Processes in the Presence of Humic Substances." JAWWA, 67_: 2: 101 (Feb. 1975]. NCASI Tech. Bulletins 249 and 212. "Surface Properties of Hydrogels Resulting from Treatment of Pulp and Papermill Effluents." Parts I and II (Sept. 1971). NCASI Tech Bulletin 273. "Studies on Adsorption of Spent Chlorination and Spent Caustic Extraction Stage Liquor Color and Organic Carbon on Activated Carbon." (1974). Oldham, W.F.; Gloyna, E.F. "Effect of Colored Organics on Iron Removal." JAWWA 61_: 611 (1969). Olthof, M.G.; Eckenfelder, W.W., Jr. "A Laboratory Study of Color Removal from Pulp and Paper Wastewaters by Coagulation." Tappi, 57^: 8: 55 (1974). O'Melia C R "Coagulation and Flocculation." Physiochemic al Processes for Water Quality Control W.J. Weber, Jr., ed. Wiley Interscience, New York (1972) Osipow, L.I. Surface Chemistry Theory and Indust rial Application. Robert E. Krieger Publishing Co., Huntington, New York (1972). Packham, R.F. "Some Studies of Coagulation of Dispersed Clays with Hydrolyzing Salts." Jour, of Coll. § Sci. 20_: 81 (1965). Packham, R.F. "Studies of Organic Color in Natural Water." Proc. Soc. Water Treat. Exam. 13^: 316 (1964) Posselt, H.S.; Anderson, F.J.; Weber, W.J., Jr. "Cation Sorption on Colloidal Hydrous Manganese Dioxide." Env. Sci. 5 Tech. 2_: 1087 (1968b) Posselt, H.S.; Reides, A.H.; Weber, W.J., Jr. "Coagulation of Colloidal Hydrous Manganese Dioxide." JAWWA 60_: 48 (1968a). Predali, J.J.; Cases, J.M. "Ze_at, Potential of Magnesium Carbonates in Inorganic Electrolytes." Jour, of Coll. S Inter. Sci. 45: 3 (Dec. 1973). Pruill, E.R.S. "Color Removal and Sludge Disposal for Kraft Mill Effluents." Paper Trade Journal (Aug. 1974). Ragunathan, P.; Cleasby, J.L.; Cerwick, J. A. "Coagulation, Cake Filtration and Filterability." JAWWA, 65: 3: 202 (Mar. 1973). Rapson, B.; Sullivan, D.P.; Brothers, J.S. "The NRSF Sea Water-Lime Clarification Process for Kraft Mill Effluents." Presented at the 60th Annual Meeting of Technical Section C.P.P.A., Montreal, Canada (1971).

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Biographical Sketch James Sherman Taylor was born August 24, 1941, in Miami, Oklahoma. He graduated from secondary school at Miami High School in Miami, Oklahoma and attended Oklahoma State University in Stillwater, Oklahoma on a football scholarship. He received the degree of Bachelor of Science in Industrial Engineering and Management in August 1965, from Oklahoma State University. He accepted a position as process and industrial engineer with 3M Company in Hastings, Minnesota from September 1966 until May 1967. He then worked as a research engineer on the Saturn program in Cape Canaveral, Florida for one year. In June 1968, he accepted a position as a senior engineer with Radiation Incorporated in Melbourne, Florida. He left Radiation in January 1971 to pursue a Masters degree in engineering at the University of Florida in Gainesville, Florida. In June 1972 he received a Master of Engineering degree in Environmental Engineering from the University of Florida. He entered the Ph.D. program in the Department of Environmental Engineering Sciences at the University of Florida in June 1972. He became a registered engineer in the state of Florida in April 1974. He is married to Janet Louise Taylor, formerly of Melbourne, Florida, and has two children, James Sherman Taylor II, age 11, and Briton Ashley Taylor, age 5. The author presently is an Assistant Professor of Environmental Engineering Sciences at the Florida Institute of Technology in Melbourne, Florida. 156

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I certify that I have read this study and that in my opinion it conforms to acceptable standards of scholarly presentation and is fully adequate, in scope and quality, as a dissertation for the degree of Doctor of Philosophy. k Zoltek, 'jy. Chairmiin Associate Professor of Environmental Engineering Sciences I certify that I have read this study and that in my opinion it conforms to acceptable standards of scholarly presentation and is fully adequate, in scope and quality, as a dissertation Lor tne degree of Doctor of Philosophy. / U.^tv---^^^ r. ( •cc T. deS. Furman Professor of Environmental Engineering Sciences I certify that I have read this study and that in my opinion it conforms to acceptable standards of scholarly presentation and is fully adequate, in scope and quality, as a dissertation for the degree of Doctor of Philosophy. :>ir'. Edward Sing ley Professor of Environmental Engineering Sciences I certify that I have read this study and that in my opinion it conforms to acceptable standards of scholarly presentation and is fully adequate, in scope and quality, as a dissertation for the degree of Doctor of Philosophy. Ellis D. Verink, Jr. / Professor and Chairman of Material Science and Engijieering

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This dissertation was submitted to the Graduate Faculty o£ the College of Engineering and to the Graduate Council, and was accepted as partial fulfillment of the requirements for the degree of Doctor of Philosophy. August 1976 ( ;'<^-e.V''''