EPA-R2-72-064
DECEMBER 1972
Environmental Protection Technology Series
Calcium Phosphate Precipitation
in Wastewater Treatment
Office of Research and Monitoring
U.S. Environmental Protection Agency
Washington. DC. 20460
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RESEARCH REPORTING SERIES
Research reports of the Office of Research and
Monitoring, Environmental Protection Agency, have
been grouped into five series. These five broad
categories were established to facilitate further
development and application of environmental
technology. Elimination of traditional grouping
was consciously planned to foster technology
transfer and a maximum interface in related
fields. The five series are:
1. Environmental Health Effects Research
2. Environmental Protection Technology
3. Ecological Research
4. Environmental Monitoring
5. Socioeconomic Environmental studies
This report has been assigned to the ENVIRONMENTAL
PROTECTION TECHNOLOGY series. This series
describes research performed to develop and
demonstrate instrumentation, equipment and
methodology to repair or prevent environmental
degradation from point and non-point sources of
pollution. This work provides the new or improved
technology required for the control and treatment
of pollution sources to meet environmental quality
standards.
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EPA-R2-72-064
December 1972
CALCIUM PHOSPHATE PRECIPITATION
IN WASTEWATER TREATMENT
by
A. B. Menar
D. Jenkins
Grant #17080 DAR
Project Officer
Warren A. Schwartz
U.S. Environmental Protection Agency
National Environmental Research Center
Cincinnati, Ohio 45268
for the
OFFICE OF RESEARCH AND MONITORING
U.S. ENVIRONMENTAL PROTECTION AGENCY
WASHINGTON, D.C. 20460
For sale by the Superintendent ol Documents, U.S. Government Printing Office, Washington, D.C. 20402 - Price $1.60
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EPA Review Notice
This report has been reviewed by the Environmental Protection Agency,
and approved for publication. Approval does not signify that the
contents necessarily reflect the views and policies of the Environmental
Protection Agency, nor does mention of trade names or commercial products
constitute endorsement or recommendation for use.
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ABSTRACT
This investigation examined the precipitation of calcium phosphate both
from chemically-defined solutions representative of wastewater composition
and from wastewater. The steady state solid phase that controlled
dissolved phosphate residual was an amorphous tricalcium phosphate. The
solubility of this phase, determined from chemically-defined systems,was
used with success to predict dissolved phosphate residuals from both
chemically-defined systems and wastewaters. Suspension recycle was found
to result in lower dissolved phosphate residuals, but poor suspension
settling below pH 10 made this process difficult to maintain. Suspension
settling was enhanced by Mg(OH)2 precipitation but not by CaC03 precipi-
tation. In wastewater of moderate alkalinity and hardness, a phosphate
removal in excess of 80% was consistently achieved at pH 9.5 with lime
doses of, at the most, 200 mg/ji as CaC03. The overall phosphate removal
performance was dictated by the performance of the precipitation reactor
and its ensuing sedimentation basin. Phosphate-containing particles that
escaped sedimentation could not be removed by filtration because they
dissolved rapidly during the recarbonation process that necessarily
precedes the filtration step.
This report was submitted in fulfillment of Project Number 17080 DAR
under the sponsorship of the environmental Protection Agency by the
Sanitary Engineering Research Laboratory, University of California,
Berkeley, California 94720.
in
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CONTENTS
Secti on Page
I Conclusions 1
II Recommendations 5
III Introduction 7
IV Phosphate Removal from Wastewaters by Calcium
Phosphate Precipitation 9
V Mechanism of Calcium Phosphate Precipitation 15
VI Approach Rationale 21
VII Experimental Equipment and Procedures 23
VIII Precipitation of Calcium Phosphate in
Chemically-Defined Systems 31
IX Precipitation of Phosphate in Wastewater 57
X Acknowledgments 71
XI References 73
XII Glossary 77
XIII Appendices 79
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FIGURES
Page
1 Lime Requirement to Reach pH 11 as a Function of
Wastewater Alkalinity 10
2 Schematic Flowsheet for Experiments on Chemically-
Defined Systems 24
3 Schematic of Calcium Phosphate Precipitation with
Suspension Recycle 25
4 Treatment Systems Used in Studies at SERL Pilot Plant . 27
5 Dependence of Steady State PT Concentration on
Solution Composition and pH in Batch Reactors .... 33
6 Dependence of Steady State Pj Concentration on
Detention Time in CSTR Experiments 34
7 Effect of pH on pA for Various Calcium Phosphate
Solids 36
8 Effect of pH on pA of Tricalcium Phosphate 37
9 Critical Ionic Concentrations in Batch and CSTR
Experiments 38
10 Effect of Mg on pA Ca3(PQk)2 40
11 Influence of Mg on Steady State Concentrations of
Ca, Mg, CT, and PT at pH 9.5 42
12 Effect of CT on pA Ca3(POl|)2 44
13 Effect of CT on Steady State Concentrations of
Ca, Mg, and Py and on Amount of Precipitated
Carbonate at pH 9.5 45
14 Effect of Cj on Steady State Concentrations of
Ca, Mg, and Py and on Amount of Precipitated
Carbonate at pH 11.0 46
i i
15 Effect of pH on Ca/Pj Mole Ratio of the Precipitate
and on Carbonate Incorporation Into the Precipitate . 47
16 Effect of Suspension Concentration on Effluent
Phosphate Concentration in Precipitate Recycle
Experiments 53
vi
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Page
17 Steady State Ca and Cj Concentrations During
Precipitate Recycle Experiments ............ 53
18 Effect of Suspension Concentration on Rate of Calcium
Phosphate Precipitation in Batch Experiments ..... 54
19 Effect of Suspension Concentration on pA Ca3(POlt)2
in Precipitate Recycle Experiments .......... 55
20 Phosphorus Removal Performance of the Entire Process
Train During Pilot Plant Runs on Wastewater ...... 60
21 Comparison of Lime Slurry Addition to Dissolved
Lime Concentration in Wastewater Experiments ..... 62
22 Change in Dissolved Ca and Py Concentrations in
Reactor Compartments - Wastewater Experiments ..... 63
23 Dependence of pA Ca3(POif)2 on pH f°r Wastewater
Experiments ...................... 64
24 Settling Properties of Calcium Phosphate-Carbonate
Suspensions (Series I) ................ 67
25 Settling Properties of Calcium Phosphate-Carbonate
Suspensions (Series II) ................ 68
26 Critical Ionic Concentrations as Determined by Walton
et jj]_. and by This Investigation ........... 84
27 Effect of Unit Processes on pH .............. 90
28 Effect of Unit Processes on Total Phosphate
Residuals (Runs 1, 2, and 3) ............. 91
29 Effect of Unit Processes on Total Phosphate
Residuals (Runs 4 and 5) ............... 92
30 Effect of Unit Processes on Dissolved Phosphate
Residuals (Runs 1 and 2) ............... 92
31 Effect of Unit Processes on Dissolved Phosphate
Residuals (Runs 3, 4, and 5) ............. 93
32 Effect of Unit Processes on Dissolved Calcium
Concentration (Runs 1, 2, and 3) ........... 94
33 Effect of Unit Processes on Dissolved Calcium
Concentration (Runs 4 and 5) ............. 95
34 Effect of Unit Processes on Alkalinity .......... 96
vii
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TABLES
Page
1 Component Concentrations in Wastewater
(From Ferguson [3]) 21
2 Initial Composition of Batch Systems Used to
Define Time to Reach Steady State 32
3 Initial Concentrations in CSTR Experiments to
Determine Effect of Magnesium 41
4 Initial Conditions for CSTR Experiments to Determine
Effect of Carbonate (CT) 43
5 Initial Conditions for CSTR Experiments to Determine
Effect of Carbonate on Phosphate and Other
Component Residuals 43
6 Comparison of Predicted and Experimental Residual
Phosphate Values in Chemically-Defined Systems .... 50
7 Comparison of Predicted and Experimental Residual
Phosphate Values in Chemically-Defined Systems
Corrected for Calcium Carbonate Precipitation 51
8 Operating Conditions and Results for Wastewater
Experiment 1 58
9 Operating Conditions and Results for Wastewater
Experiment 2 58
10 Operating Conditions and Results for Wastewater
Experiment 3 59
11 Operating Conditions and Results for Wastewater
Experiment 4 59
12 Operating Conditions and Results for Wastewater
Experiment 5 60
13 Comparison of Predicted and Average Experimental
Dissolved Phosphate Residuals for Wastewater
Experiments 63
14 Initial Component Concentration for Precipitate
Separation Experiments, Series I 66
15 Initial Component Concentration for Precipitate
Separation Experiments, Series II 66
vi ii
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SECTION I
CONCLUSIONS
The general objective of this investigation was to develop a method of
predicting the performance of phosphate removal processes that involve
the precipitation of calcium phosphate. The development of this predic-
tive method was based on a steady state precipitation model. The model
was developed by conducting experiments in chemically-defined systems and
then examining the utility of these results in predicting dissolved
phosphate residuals obtained by precipitation of calcium phosphate from
wastewaters.
It became evident as the work progressed that a major determining factor
in the efficiency of calcium phosphate precipitation processes was the
ability to separate the insolubilized phosphate. The objectives of the
work were therefore expanded to include an examination of some of the
chemical factors that influenced precipitate separation by sedimentation,
and of some of the process considerations that determine the efficiency
of particulate phosphate removal.
CHEMICALLY-DEFINED SYSTEMS
The effect of common wastewater components (calcium, carbonate, magnesium)
and precipitation conditions (pH and reaction time) on dissolved phosphate
residual was investigated in chemically-defined systems in both CSTR and
batch reactors. At chemical component concentrations typical of waste-
waters, steady state levels of dissolved phosphate were reached in
continuous flow reactors after nominal residence times of 10 min and
were maintained for at least 200 min - an average residence time range
typical of wastewater precipitation processes.
The data from chemically-defined systems suggested that the nature of
the steady state phase was one with the stoichiometric and solubility
characteristics of tricalcium phosphate Ca3(POit)2(pA = 23.56). This
solid was amorphous and did not have a distinct X-ray diffraction
pattern. However, later calcium phosphate precipitation experiments in
chemically-defined solutions in which precipitate recycle was conducted
showed the presence of a solid with the X-ray diffraction pattern of
tricalcium phosphate (Ca3(P04)2-4H20). Using the solubility of trical-
cium phosphate (pA = 23.56) to predict residual dissolved phosphate
concentration in chemically-defined systems gave agreement within 25%
of experimental values for systems in which the initial magnesium concen-
tration was less than 2 mM (200 mg/£ as CaC03) and the initial carbonate
concentration was less than 4 mM (400 mg/£ as CaC03) - neither of which
is rarely exceeded in wastewaters.
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the recycling of precipitate suspension in chemically-defined systems
was found to enhance the removal of dissolved phosphate and therefore
lower the phosphate residual for a given pH and solution composition.
Difficulties were encountered, however, in these experiments because of
poorly settling suspensions which resulted in the inability to maintain
high suspension concentrations for long periods. Studies in the critical
area of the settling properties of calcium phosphate-carbonate suspensions
conducted in chemically-defined systems revealed that initial magnesium
concentration and the pH of precipitation were the most important factors
Influencing the separation of these suspensions by sedimentation. Increases
1n the pH to between 9.0 and 11.0 improved the settling of the suspension.
The beneficial influence of magnesium concentration on suspension settling
was exerted most strongly at pH values of greater than 10.5 where
precipitation of gelatinous Mg(OH)2 was likely. Suggestions in the
literature that calcium carbonate precipitation aids in precipitate
separation by producing a dense suspension were not supported by this
investigation.
In light of the stated findings it appears that high (>90%) phosphate
removal by calcium phosphate precipitation followed by suspension
settling is only possible at pH values of greater than 10. It would
also appear that to achieve the same degree of phosphate removal at a
pH of 9 to 10 the precipitate must be coagulated either by cationic
polyelectrolytes, ferric chloride, or alum.
PHOSPHATE REMOVAL FROM WASTEWATERS
Investigations of phosphate removal from wastewaters by calcium phosphate
precipitation were conducted at the Sanitary Engineering Research
Laboratory (SERL) wastewater treatment facility. The phosphate removal
performance of two treatment schemes was examined. The first involved
primary sedimentation, activated sludge, lime precipitation, recarbonation,
and clinoptilolite sorption; the second, lime precipitation following
primary sedimentation.
Values of residual dissolved phosphate in wastewater precipitated with
lime at pH values between 9.5 and 11.0 could be predicted to within
25% of their experimental values if these predictions were based on the
initial concentrations of dissolved lime rather than on the total lime
added and if a correction for calcium complexation was made. This was
because the efficiency of lime dissolution in the slaking operation was
poor and lime continued to dissolve throughout the several compartments
of the precipitation unit. ,
An overall phosphate removal of greater than 80%, to achieve a residual
of less than 2 mg P/a, could be consistently achieved from this waste-
water, which was one of average alkalinity (240 mg/j> as CaC03) and a
typical Ca/Mg mole ratio of 3, by precipitating activated sludge effluent
or primary effluent with lime doses of 200 mg/t, as CaC03 to achieve a
pH of 9.5.
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The overall phosphate removal performance of the entire treatment train
was dictated by the combined performance of the precipitation reactor
and the ensuing precipitate separator (in this case a sedimentation
basin). Any calcium phosphate particles that escaped sedimentation were
dissolved in the low average residence time (5 min) recarbonation basin
that necessarily preceded the clinoptilolite sorption columns. This
observation has important ramifications in the removal of phosphate by
calcium phosphate precipitation. It means that improvement of the overall
phosphate removal of a calcium phosphate precipitation process cannot be
achieved by post-filtration of the effluent from the solids separator
of such a process. Filtration of such an effluent will require a prior
downward pH adjustment, but this pH adjustment will cause a rapid
dissolution of calcium phosphate particles making it impossible to remove
them by filtration. These observations emphasize the importance of
solids separation in calcium phosphate removal and suggest this general
topic as a fruitful area for investigation.
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SECTION II
RECOMMENDATIONS
The work reported herein has demonstrated that short-term calcium phos-
phate precipitation processes can be modelled successfully on the basis
of the formation of an amorphous tricalcium phosphate. Indications that
this steady state solid phase can be made to transform and grow into a
more insoluble calcium phosphate phase were obtained from precipitate
recycle experiments. It is recommended that these experiments be further
pursued to determine the factors that influence the formation and growth
of more insoluble calcium phosphate solids in wastewaters since the
formation of these solids will allow the attainment of lower dissolved
phosphate residuals. Investigations of these phenomena should proceed
along several lines including: 1) the role of recycled precipitate in
providing crystal growth opportunity, 2) the role of precipitate in a
sludge blanket clarifier in providing sites for crystal growth, and
3) the possible use of calcium phosphate clinker from lime regenerated
by recalcining to provide material on which calcium phosphate crystal
growth may occur.
An important conclusion of this investigation is that the phosphate
residuals in effluents of calcium phosphate precipitation processes
cannot be reduced by post-filtration because prior downward pH adjustment
will dissolve the phosphate-containing particles. It is important
therefore to devote considerable effort to improving precipitate removal
by sedimentation processes, especially since it is possible to produce
low dissolved phosphate residuals at pH values of 8.5 -9.5 (but at
these pH values precipitate separation is difficult). The factors that
influence the flocculation and settling properties of calcium phosphate-
carbonate suspensions produced at pH values below 10 should be investigated.
Such an investigation should consider the surface properties of such
suspensions and their possible modification by coagulants to improve
their settling properties.
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SECTION III
INTRODUCTION
PHOSPHATES IN WASTEWATERS
The conventional combination of primary sedimentation and secondary
biological treatment processes with anaerobic digestion for sludge treat-
ment can be relied upon to remove between 2 and 3 mg P/£ from domestic
wastewater of average strength. With the current domestic wastewater
total phosphate content averaging about 10 mg P/a, such removals typically
account for some 20% to 30% of the incoming phosphate in domestic waste-
water. About half of the total phosphate in domestic sewage is derived
from synthetic detergent builders so that even if the phosphates in these
products were to be completely replaced by nonphosphate containing
compounds, a sewage with some 5 mg P/Jl would result. Treatment of such
a wastewater by current primary and secondary treatment schemes would be
expected to leave a residual of some 2 to 3 mg P/i. In instances where
phosphate has been identified as a nutrient limiting the growth of aquatic
photosynthetic organisms it is generally agreed that growth control over
these organisms by controlling phosphate concentration may only be exerted
when phosphate levels on the order of 50 yg/£ or less are reached. It
would therefore appear that in such instances treatment further than the
conventional primary and secondary biological methods would be needed to
produce such phosphate levels. Since complete elimination of phosphate
from synthetic detergents is neither an impending nor a likely event and
would not reduce raw sewage phosphate to less than 5 mg P/a in any case,
the development of processes that achieve high phosphate removals is
justified.
Of the treatment processes suggested for increasing phosphate removal to
a level greater than that possible by primary and secondary treatment,
those using precipitation with metal ions are the only ones to have found
wide application and to be economically feasible. Of the chemical
precipitants commonly used (ferric and ferrous iron, aluminum and calcium
salts) lime has been the most common, possibly because of a combination
of its cheapness, the capability for its regeneration, and its familiarity
in the field. Most calcium phosphate precipitation schemes involve
raising the pH of the waste stream to 11 or higher. Under these conditions
low phosphate residuals (<1 mg P/a) and readily settleable suspensions
(supposedly because of concomitant Mg(OH)2 precipitation) are obtained.
However, to achieve these high pH values on a variety of typical waste-
waters Nesbitt [1] has reported that lime doses of between 280-720 mg/£
as Ca(OH)2 are required since the lime dose is largely determined by the
wastewater alkalinity. Moreover, pH adjustment of the treated effluent
(commonly by recarbonation) is mandatory following these high pH
precipitation processes.
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Several authors (Menar and Jenkins [2], Ferguson [3], Ferguson et al.
[4], and Jenkins et al_. [5]) have suggested and demonstrated experimentally
that phosphate residuals of 1 mg/fc or less can be obtained in wastewaters
and synthetic systems representing wastewaters under some conditions at
pH values far below 11. Most of these experiments have, however, been
conducted in batch systems with reaction times and concentration ranges
different from those that exist in a typical lime precipitation unit for
removing phosphate from wastewater.
It was the purpose of this investigation to determine the behavior of
chemically-defined systems representative of wastewater when subjected
to phosphate precipitation by the addition of calcium salts (including
lime) at various conditions of pH and under the physical constraints of
reaction time and reactor design typical of a wastewater phosphate
precipitation process. The overall aim of such experiments was to develop
a predictive method that would allow the determination of chemical dose
and phosphate residuals from chemically-defined systems and test it on
wastewater.
OBJECTIVES
The general objective of this research was to develop a method for
predicting the performance of phosphate precipitation processes involving
the addition of calcium as a precipitant. Such processes include but are
not necessarily restricted to: lime addition; the addition of calcium
salts plus strong base; and the use of existing wastewater calcium hard-
ness as a phosphate precipitant accompanied by aeration for upward pH
adjustment. The general objective was reached by the fulfillment of the
two following specific objectives: 1) the development of a predictive
method based on calcium phosphate precipitation experiments in chemically-
defined systems containing calcium, magnesium, orthophosphate, carbonate,
hydrogen ion and water, and 2) the testing of the results of these experi-
ments on domestic wastewaters under differing conditions of precipitation
to determine the utility of the predictive method.
It became evident as the work progressed that a major determining factor
in the efficiency of calcium phosphate precipitation processes was the
ability to separate the insolubilized calcium phosphate. The objectives
of the work were therefore expanded to include an examination of some of
the chemical factors that influenced precipitate separation by sedimentation,
8
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SECTION IV
PHOSPHATE REMOVAL FROM WASTEWATERS BY
CALCIUM PHOSPHATE PRECIPITATION
The precipitation of calcium phosphate from wastewater has most commonly
been conducted by the addition of unslaked or slaked lime to provide
simultaneous increases in both calcium and hydroxyl ion. Most calcium
phosphate precipitation schemes involve raising the pH of the waste stream
to at least 10.5 (and often above pH 11) because in this pH range experi-
ence has shown that low dissolved phosphate residuals (<1 mg P/£) are
obtained together with a settleable precipitate.
Perhaps the earliest observation of the use of lime for phosphate preci-
pitation in the waste treatment field was by Rudolfs [6]. He concluded
that phosphate removals exceeding 90% — to reach residuals of approximately
0.3 mg P/a — could be obtained together with excellent flocculation by
lime addition to reach a pH of higher than 10.1.
Because phosphate precipitation processes using calcium have been conducted
at high pH values, lime requirements have been largely dictated by the
alkalinity of the wastewaters and have borne little or no relationship to
its phosphate concentration. Thus Wuhrmann [7] states that the lime
requirement for phosphate removal at pH 10.5 to 11 is equal to 1.5 times
the carbonate hardness. Mulbarger et^ a]_. [8] and Buzzell and Sawyer [9]
present a summary figure (Figure 1) of lime dose required to attain pH
11.0 as a function of wastewater alkalinity. The wide variation in lime
doses to produce low dissolved phosphate residuals (0.4-2.4 mg P/s,) is
indicated by Nesbitt's [1] summary of the literature in which it is
revealed that doses ranging from 280 to 720 mg/i Ca(OH)2 were employed
by various investigators. »
From a survey of operating data in tertiary lime precipitation plants at
Blue Plains, Washington, D. C.; Pomona, California; Lebanon, Ohio;
Las Vegas, Nevada; and S. Lake Tahoe, California, Seiden and Patel [10]
have concluded that residual dissolved phosphate concentration will be
about 0.15 rag P/A at pH 11. Notwithstanding the conclusion that high pH
operation produces readily settleable precipitates, it is Nesbitt's
opinion that a minimum phosphate residual can only be obtained if a
filtration step (necessarily requiring recarbonation) follows lime
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500
400
5 300
S
CM
200
100
• Boston, Mass.
o Marlboro, Mass.
A Clinton, Mass.
A Leominster, Mass.
• Hudson, Mass.
a Schenectady, N. Y.
After Buzzell & Sawyer [9]
50
100
150 200 250
ALKALINITY, mg CaC03/fc
300
350
400
FIGURE 1. LIME REQUIREMENT TO REACH pH 11 AS A FUNCTION OF WASTEWATER ALKALINITY
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precipitation. It is interesting, however, to note that if such a
filtration step is contemplated for the removal of calcium phosphate
particles from the effluent of a high pH precipitation unit it will be
necessary to lower the pH value of the effluent (possibly by recarbonation)
prior to filtration. The effect that the fall in pH has on the dissolu-
tion of calcium phosphate particles has not been discussed by authors who
suggest filtration following lime precipitation to enhance phosphate
removal.
Wuhrmann [7] has indicated that the calcium carbonate and apatite preci-
pitates that form at high pH values do not settle readily, and he suggests
the addition of 1 to 2 mg/Jl Fe(III) to aid flocculation of the precipi-
tates. In pilot plant post-precipitation or tertiary treatment of acti-
vated sludge effluent by this method (268 mg/s, Ca(OH)2, 1 to 2 mg/£
Fe(III), pH 11.1) phosphate reductions from 3.2 to 0.4 mg P/a were
obtained. Results that indicate the carryover of phosphate in particles
were obtained by Owen [11] who precipitated phosphate from a high-rate
trickling filter effluent. Using 720 mg/2, Ca(OH)2, a dissolved phosphate
residual of 0.13 mg P/a was obtained, while the total phosphate residual
was 1.7 mg P/a. Reducing the lime dose to 360 mg/s, produced a soluble
phosphate residual of 1.5 mg P/£ and a total residual of 2.6 mg P/i.
The use of alum and lime at pH of 10 to 11 has been reported by Spiegel
and Forrest [12] at Amarillo, Texas, where the wastewater phosphate
concentration was reduced to at least 0.4 mg P/a.
Experience at the South Tahoe Public Utilities District Plant [13,14] has
shown that the addition of 400 mg/a CaO to activated sludge effluent to
raise the pH to 11.5 will produce a phosphate residual of 0.3 mg P/a by
plain settling of the precipitate. However, the process flow sheet has
included multimedia filtration subsequent to the sedimentation of the
precipitate.
A Densatop pilot plant operated at the Los Angeles County Sanitation
District's Pomona facility for conditioning activated sludge effluent
prior to treatment by ion exchange showed the beneficial effect of
filtration in decreasing phosphate residuals [15]. When operated between
pH 10 and 11 the Densator produced an effluent containing an average total
phosphate of 1.85 mg P/a and an orthophosphate of 0.2 mg P/a. Filtration
changed these values to 0.11 and 0.09 mg P/4, respectively. Operation at
pH values in the range 9.1 to 9.8 produced mean unfiltered residuals of
5.5 mg P/a total phosphate and 0.86 mg P/a orthophosphate, which were
reduced to 0.65 and 0.48 mg P/a by filtration.
In summary, high pH (>10.5) lime precipitation processes, therefore,
appear to produce low residuals (0.5 mg P/a}, but separation of the
colloidal precipitate may be difficult and require the use of coagulants.
There are several processes and suggested processes for calcium phosphate
precipitation that involve the use of pH values lower than 10.5. In one
11
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of these, the phosphate extraction process (PEP) proposed by Albertson
and Sherwood [16], a key feature is the recycle of previously formed
chemical sludge. The process involves lime addition to raw wastewater
in a flocculator-clarifier to achieve a pH of 9.5 to 10.5, and recycle
of clarifier underflow to maintain a solids concentration of 500 to
2000 mg/Si in the flocculator. The clarifier effluent which contains
residual total phosphate of about 2 to 3 mg P/a is fed to a completely
mixed activated sludge aeration basin which further reduces the phosphate
content to 1-2 mg P/a and is claimed to provide an overall process BOD
and suspended solids removal of 90% to 95%. Chemical sludge recovery by
thickening, dewatering, combustion (calcining), and scrubbing is proposed.
It is also proposed that pH adjustment following lime treatment to a
range that is suitable for activated sludge operation can be achieved by
the biologically produced C02 in the completely mixed aeration basin.
This proposition is supported by studies at the South Tahoe Public
Utilities District Plant [14] which indicate that when recalcined lime
was added to the primary clarifier to produce a primary effluent of
pH 11 the aeration basin contents some 5 ft into the aeration basin had
a pH of 8.5 to 8.6.
Buzzell and Sawyer £9] expressed doubts that an activated sludge unit
could be effectively operated on a primary effluent that has been treated
with lime to a pH of 10 to 11 since insufficient degradable organic carbon
would be present. In addition, it is possible that the precipitation unit
might reduce phosphate concentrations to the point that the ensuing
biological process would be phosphorus limited; it also is possible that
carryover of solids from the flocculator process could significantly
decrease the phosphate removal of the process. Humenick and Kaufman [17]
have used the foregoing rationale to suggest that a chemical-biological
process combination for achieving high BOD and phosphate removals should
use a high-rate biological treatment prior to a chemical precipitation
unit.
The recirculation of chemical sludge in the PEP process besides aiding
the flocculation of wastewater particulates, enables total phosphate
residuals of about 2 mg P/i to be achieved at approximately one pH unit
lower than without precipitate recycle and consequently with about one
half the lime dose of a conventional lime precipitation process. The
presence of recycle undoubtedly increases the rate of calcium phosphate
precipitation by providing nuclei or growth sites for precipitate
formation and growth. -The provision of nuclei for the precipitation of
apatite has been shown by Stumm and Leckie [18] and Ferguson e_t a\_. [4]
to strongly influence the rate of phosphate removal under similar condi-
tions in chemically-defined systems.
The significance of sludge recycle has been demonstrated by work at. the
South Tahoe plant [14] in which introduction of lime mud recycle from the
chemical clarifier to the lime rapid mix reduced the mean effluent
phosphate concentration from 0.22 to 0.16 mg P/£. Recycling of scrubber
water from the sludge and recalcining furnaces (which contained a total
phosphate concentration of 12.8 mg P/z, of which 12.3 mg P/a was parti-
culate) caused a reduction of the effluent orthophosphate from 0.16 to
12
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0.09 mg P/X,. When these recycle streams were removed from the rapid lime
mix basin and again introduced to the primary clarifier, the average
effluent phosphate concentration increased from 0.09 to 0.31 mg
Schmidt and McKinney [19] conducted laboratory tests of a chemical-
biological scheme that is essentially similar to the PEP process except
that it includes no provision for precipitate recycle. These workers
used a lime dose of about 150 mg/x. Ca(OH)2 to reach a pH of 9.5 in raw
sewage precipitation. At this pH these authors claim that polyphosphates
present in the raw sewage prevent the precipitation of calcium carbonate
by adsorbing onto calcium carbonate nuclei. Since the major consumption
of lime in a precipitation process is devoted to calcium carbonate preci-
pitation, a reduction in lime dose will be obtained by this phenomenon.
No full-scale or pilot-plant application of this work has been reported.
There are several activated sludge treatment plants in the United States
at which phosphate removals are achieved that are far in excess of those
predictable by normal biological means. There are two general schools of
thought concerning the mechanism of these removals — that which proposes
a biological mechanism ("luxury uptake") and that which proposes a chemical
precipitation mechanism. Recent studies by Ferguson e_t ^1_. [20] have
demonstrated that the high removals of phosphate experienced at the
Rilling Road Treatment Plant of San Antonio, Texas can be explained on
the basis of a combination of calcium phosphate precipitation and reactor
configuration (aeration basin design). This demonstration was based on
the previous work of Menar and Jenkins [2] who explained the so-called
luxury uptake of phosphate in terms of calcium phosphate precipitation.
They proposed that precipitation occurred in an activated sludge aeration
basin at a point where C02 stripping caused the pH to increase to the
point that calcium phosphate precipitated. The behavior observed at
San Antonio [21] could not be reproduced by a pilot plant operated by
Menar and Jenkins [2] at San Ramon, California. Thus, at San Antonio
soluble phosphate residuals of 1 to 2 mg P/2. were obtained at pH 7.8
to 7.9 while at San Ramon, even at pH values above 8.5, soluble phosphate
concentrations of greater than 10 mg P/l remained. Ferguson and McCarty
[22] explained this difference in behavior based on the effects of
carbonate and magnesium on phosphate precipitation proposing that in the
lower alkalinity and magnesium concentrations at San Antonio (compared
to San Ramon) more rapid calcium phosphate precipitation to form a more
insoluble solid occurred. Ferguson et_ al_. [20] proposed that apatite
formation occurred significantly only during residence times typical of
activated sludge treatment in the plug flow aeration basin of the Rilling
Road activated sludge plant at San Antonio and not in the more back-
mixed basins of the East and West plants.
Calcium phosphate precipitates formed in primary, secondary, or tertiary
treatment units will tend to redissolve at the lower pH values encountered
in anaerobic digestion. Indeed, redissolution of calcium phosphates
precipitated in activated sludge aeration basins is found even in the
slightly lower pH (higher C02) environment encountered in the sludge
13
-------
blanket of a secondary clarifier [2]. Release of phosphorus has been
generally observed from all activated sludges claimed to be showing
"luxury uptake" when the sludges were exposed to anaerobic conditions
and lowering of the pH.
14
-------
SECTION V
MECHANISM OF CALCIUM PHOSPHATE PRECIPITATION
INTRODUCTION
Although the use of lime for precipitating calcium phosphate from waste-
water has been suggested and tested widely, the mechanism of calcium
phosphate precipitation from this medium is not well defined. For the
conditions commonly encountered in wastewater, hydroxyapatite (Ca5OH(POit)3)
is the thermodynamically stable solid calcium phosphate phase. The
concentrations of dissolved calcium and phosphate in lime-treated domestic
wastewaters, however, far exceed the equilibrium values of crystalline
hydroxyapatite. It is therefore apparent that either precipitation
kinetics or the presence of more soluble calcium phosphate phases control
phosphate residuals in wastewaters treated with lime.
PRECIPITATION OF CALCIUM PHOSPHATE FROM
CALCIUM PHOSPHATE SOLUTIONS
A considerable body of experimental work in chemically-defined solutions
has shown that the initial solid phase that appears in the rapid preci-
pitation of calcium orthophosphate from basic solutions is structurally
noncrystalline. The noncrystalline structure is metastable and, if
allowed to remain in contact with the preparative solution, it will
spontaneously convert into a crystalline product with a chemical
stoichiometry and X-ray diffraction pattern typical of the mineral
apatite.
Recent kinetic studies by Stumm and Leckie [18] have defined three phases
in calcium phosphate precipitation. These were: nucleation —the
formation of amorphous calcium phosphate from solution; phase transfor-
mation — the slow transformation of amorphous calcium phosphate into
crystalline apatite; and crystal growth of crystalline apatite. Phosphate
is removed from solution during nucleation and crystal growth but not
during phase transformation. Since the dissolved phosphate residual in
equilibrium with crystalline apatite will be lower than that coexistent
with the amorphous nucleating phase, Stumm and Leckie [18] suggested that
low wastewater phosphate residuals might be obtained by recycle of
preformed crystalline apatite to a precipitation reactor. Their sugges-
tion was based on the observation that the presence of apatite crystals
largely eliminated the phase transformation step between nucleation and
crystal growth. Ferguson ejt al_. [4] recently have demonstrated that the
addition of such preformed precipitate to calcium phosphate carbonate
15
-------
solutions supersaturated with respect to calcium phosphate, eliminates
lag periods prior to precipitation that were observed in batch reactors
not containing added precipitate.
There have been various reports of the nature of the calcium phosphate
phase that precedes crystalline apatite. Strates jst ^1_. [23] suggest
that dicalcium phosphate dihydrate (CaHPOi+-2H20) first crystallizes and
rapidly converts into apatite. McGregor and Brown [24] maintain that
octacalcium phosphate CattH(P0lt)3 is the metastable precursor of apatite.
Booth and Coates [25] suggest that dicalcium phosphate dihydrate is the
first crystalline product to appear and it transforms into octacalcium
phosphate, Ca^HKPOiJa, which itself then transforms into apatite.
Walton et_ al . [26] found that over the pH range 6.3 to 9.04 calcium
phosphate "[presumably apatite) was preceded by a metastable precursor
of the stoichiometry but not the structure of tricalcium phosphate,
Ca3(POj2.
The stable crystalline apatite phase formed from pure calcium phosphate
solutions, while yielding X-ray diffraction patterns typical of apatite
(Ca5(P0lt)3OH) , may deviate significantly in stoichiometry from the Ca/P
mole ratio of 1.67 predicted from this formula. Such calcium deficient
and other nonstoichiometric apatites have been reported over the entire
range of Ca/P mole ratios from 1.33 to 2 [27] with crystal structures
indistinguishable by X-ray diffraction from well-crystallized apatite.
Several theories have been advanced to explain the existence of these
nonstoichiometric apatites - thus some workers [28] explain low Ca/P
ratios by postulating the presence of adsorbed phosphate on the solid
surface, possibly in the form of surface complexes. Rootare et al .
[29] believe that these surface complexes control the phosphate concen-
tration in a solution in equilibrium with apatite and are formed by
hydrolysis of the apatite surface thus:
Ca10(POit)6(OH)2 + 6H20 $ 2Ca2+ + 2HPOJ" + 4[Ca2HPOtf(OH)2] (1)
Ca2HP0lt(OH)2 * 2Ca2+ + HPOJ" + 20H" (2)
The evidence for the existence of such a complex is indirect — and
largely based on a change in Ca/P mole ratio in solution from 1 to 1.67
as the concentration of an apatite suspension is decreased.
McConnell [30] suggests the simultaneous precipitation of lower Ca/P
ratio solids together with apatite, e.g., brushite (CaHPO^, Ca/P =: 1)
and octacalcium phosphate (CaitH(P0lf)3, Ca/P = 1.33). However, this
suggestion would not be valid if the apatite formed had no detectable
phase impurities. Other investigators suggest the formation of apatite
with a calcium deficient, defective lattice [31] as well as the isomorphous
substitution of ions or groups for those normally present in the apatite
lattice.
16
-------
PRECIPITATION OF CALCIUM PHOSPHATE FROM SOLUTIONS
CONTAINING MAGNESIUM AND CARBONATE
Domestic wastewaters contain a wide variety of organic and inorganic
compounds, many of which might influence the precipitation of calcium
phosphate. However, hydrogen ion (pH), magnesium, carbonate, and, to
a lesser extent, fluoride (because of its low concentration in waste-
waters), appear to be the most significant. Their effects will be
reviewed in the context of the three steps of precipitation (nucleation,
phase transformation, and crystal growth) defined by Stumm and Leckie [18].
Since pH influences the relative abundance of the various phosphate
species and components of the carbonate system, it is important in
determining the range of stability of various calcium phosphate solids.
However, for pH values above 7 with solution Ca/P ratios of >0.5 the
precipitation of monocalcium phosphate (Ca(H2POit)2) and its monohydrate
(Ca(H2POi.)2'H20) or dicalcium phosphate (monetite, CaHPOiJ and its
hydrate (brushite, CaHPOi+'2H20) are not important. In fact for commonly
encountered wastewater Ca/P ratios, hydroxyapatite is the stable phase
at all commonly encountered process pH values.
pH influences the stability of formed precipitates in the sense that at
high pH values rapid precipitation of unstable solids results, whereas
at lower pH values slower precipitation of more stable solids may occur.
Walton et_ aj_. [26] reports that hydroxyl ion is directly involved in the
formation of the critical nucleus in calcium phosphate precipitation and
that the rate of nucleation is a function of [OH~]2. Eanes ejt ail_. [32]
report that a pH increase causes a reduction in the rate of crystalliza-
tion from the amorphous phase. They report that conversion to apatite
was completed in 2.75 hr at an initial pH of 8.0, while at a pH of 9.8
complete conversion took place in 6.5 hr. The apparent stabilizing
effect of the higher pH upon the noncrystalline phase is probably due to
the incorporation of impurities into the precipitate during the more
rapid high pH nucleation phase.
Magnesium has been reported to influence both the rate of calcium phos-
phate precipitation and the nature of the precipitated solid. It is not
likely, however, that magnesium phosphate or carbonate solids will form
for the ranges of magneisum, carbonate, and phosphate concentrations
typical of wastewaters. Magnesium inhibits the nucleation of calcium
phosphate, possibly due to the competition between magnesium and calcium
ions and to the formation of the strong magnesium phosphate complex
MgHPOlf[33]. The presence of magnesium at calcium phosphate crystal
growth sites has been postulated to slow down the transformation of
amorphous calcium phosphate phases into apatite [4]. Newesely [34] has
reported that anhydrous tricalcium phosphate and Mooney and Aia [35]
have reported that beta tricalcium phosphate (whitlockite) form instead
of apatite in the presence of magnesium. Whitlockite is often found
with magnesium as a minor component -with contents of about 6 to 8
atom percent. Other metals, e.g. manganese and ferrous iron, also
17
-------
stabilize whitlockite. Ferguson [3] has proposed that the presence of
Mg/Ca mole ratios of greater than 5 stabilize beta tricalcium phosphate
in sludge digester liquors and proposed that in wastewaters of differing
Ca/Mg ratios either apatite or a stabilized beta tricalcium phosphate
may control the dissolved phosphate concentration.
In systems containing carbonate and phosphate the effect of Mg on calcium
phosphate and carbonate precipitation appears to be complex and pH
dependent. Magnesium is reported to interfere with the precipitation of
the calcium carbonate solids, calcite and aragonite, by inhibition of
nucleation and lattice distortion. Calcites containing magnesium as a
minor component exist. These magnesian calcium carbonates are more
soluble than calcite and their solubility increases with their magnesium
content over the range 0-20%, from pKSD 8.4 to pK<-D 6.3 (Chave et al.
[36]). SP
Recent work in laboratory batch systems at phosphate, carbonate, and
magnesium concentrations typical of wastewaters [4] showed that besides
forming calcium carbonate and phosphate solids, carbonate and phosphate
interacted in the formation of these respective solids. Carbonate has
been reported by Ferguson e_t al_. [4] to inhibit both the nucleation and
phase transformation of calcium phosphates. Newesely [34] reported that
carbonate at 0..8 mM disturbed the cyrstallization of calcium phosphate,
and Ferguson [3] reported that no crystalline calcium phosphate is formed
from solutions of typical wastewater constituent concentrations containing
10 mM carbonate. However, the presence of carbonates in bone, dentine,
and enamel (all of which are apatite) and the existence of minerals such
as francolite and cellophane suggest the possibility of forming mixed
calcium carbonate—phosphate compounds [28]. Rapid precipitation, as
encountered in wastewater calcium phosphate precipitation processes,
results in mixed calcium phosphate—carbonate solids because the discrim-
.ination between these radicals becomes difficult and recrystallization
is a slow process. Simpson [37] has reported that high partial pressures
of C02 destabilize apatite and has observed the formation of octacalcium
phosphate at C02 partial pressures of between 0.01 and 0.1 atm. and at
pH values below 8.1.
SURFACE PROPERTIES OF CALCIUM PHOSPHATE SUSPENSIONS
It is generally observed that calcium phosphate-carbonate suspensions
settle well whe'h produced at pH values greater than about 10 but that
the opposite is true for suspensions produced at lower pH values than
this. 'The stability of colloidal calcium phosphate—carbonate suspensions
can be attributed tjo electrpstatic repulsive forces created by' a surface
charge that originates from ionization of various groups on the solid
side of the solid-liquid interface by isomorphous substitution, or by
ion adsorption. Somasundaran [38] has suggested surface hydrolysis as a
possible mechanism of.charge development on apatite:
18
-------
P0j~ + H20 J HPOJ" + OH" (3)
HPOif + H20 £ H2PO; + OH" (4)
H2PO; + H20 Z H3PO^ + OH" (5)
Ca2"1' + OH" Z CaOH+ (6)
CaOH+ + OH" t Ca(OH)2 (7)
At low pH values reactions (3), (4), and (5) will proceed in the forward
direction and (6) and (7) will proceed in the reverse direction, while
at high pH values the reverse is true. It is also evident from reactions
(3) through (7) that the charge on the solid surface depends on pH of the
solution. Somasundaran [38] found the pH at the zero point of charge
(PHzpc) for apatite to be 6. Stumm and Morgan [33] argue that the pHZpc
of salt-type minerals depends in a complicated way on the pH and also on
the concentration of all potential-determining ions. The presence of
precipitate impurities changes the pHzpc because the zpc of solids should
correspond to the pH of the charge balance (electroneutrality) of potential-
determining ions [39].
Calcium phosphate precipitation in wastewater treatment normally involves
raising the pH to 10.5 or higher. At these pH values Mg(OH)2 also preci-
pitates and its formation has been proposed as a factor in improving the
settling properties of the calcium phosphate-carbonate suspension. It
is possible that specific adsorption of Mg(OH)2 on the negative suspension
surface [33] takes place. It is also possible that bridging occurs
between Mg(OH)2 and the surface phosphate groups leading to a three-
dimensional network [40].
SUMMARY
Precipitates are formed in solutions as a result of nucleation followed
by crystal growth. In calcium phosphate precipitation, however, crystal
growth does not usually follow the nucleation process immediately, but
is preceded by a lengthy phase transformation step. During the later
step, the nucleated amorphous calcium phosphate is transformed into a
crystalline phase with no net phosphate removal from solution.
Hydrogen ion concentration exerts a significant effect on each of the
three phases of calcium phosphate precipitation. An increase in pH
speeds up nucleation, while at the same time appears to cause a reduction
19
-------
in the rate of crystallization. In addition, pH exerts a significant
effect on the surface properties of the suspension.
Magnesium appears to inhibit the nucleation of calcium phosphate, to
slow down the formation of apatite from amorphous calcium phosphate, and
to stabilize the formation of tricalcium phosphate. It also tends to
interfere with the precipitation of calcium carbonate. At pH values
above 10.5, Mg(OH)2 precipitates and appears to be involved in producing
a readily flocculable calcium carbonate-phosphate suspension.
Carbonate appears to inhibit both the nucleation and the phase transfor-
mation of calcium phosphate. A high partial pressure of CO^ reportedly
destabilizes apatite and enhances the formation of octacalcium phosphate.
20
-------
SECTION VI
APPROACH RATIONALE
The approach to the problem of predicting the dissolved phosphate resi-
duals obtainable by calcium phosphate precipitation from wastewater was
first to study a model or chemically-defined system representative of
the relevant components of wastewater. The results of this study were
then applied to the prediction of dissolved phosphate residuals obtained
experimentally from wastewater.
The complex and variable nature of domestic wastewater makes the use of
simple models difficult though necessary. This work used the model
system, Ca - Mg - Pj - Cj - H+ - H2Q^ where CT = H2C03 + HCOg + COi"
and PT = H3P04 + H2PO; + HPOif + POij with which previous authors [3,4,
22] have had some success in modelling waste treatment situations for
calcium phosphate precipitation equilibria and kinetics. It must be
realized, however, that since domestic sewage contains species such as
polyphosphate, fluoride, and organic acids, as well as clay minerals and
organic particulate solids, the use of a simplified system excluding
these cannot provide an exact representation of the actual situation.
The components of the model system vary in domestic sewage over the
approximate ranges indicated in Table 1. The pH in domestic wastewater
is typically within the range 6.5 to 8, averaging about 7.5, but wastewaters
subject to calcium phosphate precipitation may have pH values of up to
11.5. In the model chemically-defined system, the range of chemical
conditions studied were purposely made wider than those cited in Table 1.
TABLE 1
COMPONENT CONCENTRATIONS IN WASTEWATER
(from Ferguson [3])
Component
Calcium
Magnesium
Carbonate
Phosphate
Range
mM
0.
0.
2
0.
5-5
25 - 1
- 8
1 -0.5
Typical
Value
mM
1.5
0.5
4.0
0.3
21
-------
The precipitation model described by Stumm and Leckie [18] was adopted
as the basis for study of calcium phosphate precipitation. It was postu-
lated, on the basis of the literature on calcium phosphate precipitation
and a knowledge of the reaction times used in wastewater calcium phos-
phate precipitation processes, that a steady state amorphous calcium
phosphate phase controlled the residual dissolved phosphate concentration.
The validity of this steady state postulate was first tested and then the
nature of the steady state phase was investigated.
The activity product variation of various phases with pH and solution
composition was tested. The stoichiometry of the steady state phase was
determined. Attempts were made to obtain crystallographic information
from X-ray powder diffraction patterns of the aged precipitate.
The results obtained from the chemically-defined systems were applied to
the precipitation of calcium phosphate by lime from wastewaters. Since
it became evident, as the work progressed, that precipitate separation
was a major factor in determining the total phosphate residual obtainable,
the study devoted some attention to the chemical factors that influence
the removal of precipitate by sedimentation.
Limited attention was also directed toward the effect of precipitate
recycle on dissolved phosphate residual. This was prompted by the
precipitation model and results of previous investigators who attributed
low dissolved phosphate residuals to the growth of crystalline calcium
phosphates.
22
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SECTION VII
EXPERIMENTAL EQUIPMENT AND PROCEDURES
EXPERIMENTAL EQUIPMENT
The experimental phases of the study involved runs both in chemically-
defined systems containing reagent grade chemicals dissolved in distilled
water and in domestic wastewater from the SERL pilot plant. For the
chemically-defined systems both continuous flow and batch reactors were
employed. The experiments with sewage used a continuous flow system
onlv.
Chemically-Defined Systems —Continuous Flow
Experi ments
Continuous flow experiments on chemically-defined systems were conducted
in a single acrylic plastic CSTR with inside dimensions of 10 in. x 10
in. x 10 in. The reactor contents were mixed at 71 rpm by a stainless
steel stake and stator stirrer mounted vertically and driven by a 1/8 hp
DC motor (Bodine Co.) controlled by a Speed Controller (Minarik Co.).
Power input was measured by a torquemeter that could be attached to the
vertical stirrer shaft. Power input was such that an average velocity
gradient of approximately 75 sec"1 was maintained in the experiments.
To prevent precipitation in the feed lines and feed storage vessels prior
to the reactor, it was necessary to prepare the chemical solutions in
three separate batches and then mix these in a well-defined fashion just
preceding their entry into the reactor. The bulk of the feed, which was
either stored in 55-gal stainless steel drums or in a 1000-gal redwood
tank lined with a plastic sheet contained the calcium, magnesium, and
orthophosphate components. This solution was prepared by adding stock
solutions of reagent grade CaCl2, MgCl2, and H3P04 to distilled water
(Figure 2). '
After 30 min aeration to mix the solution and to produce a uniform pH
of 3.5, the solution was pumped to a constant head tank using a stainless
steel centrifugal pump. The solution then flowed by gravity through a
1/2-in. PVC pipe and a Brooks Full View Rotameter into the reactor. A
concentrated NaOH solution (for pH adjustment) was pumped into the
influent stream through a "T-joint" 20 in. ahead of the reactor inlet.
At a point 2 in. before the reactor inlet the carbonate component (Cj)
was added as a concentrated solution of sodium carbonate.
The reactor pH was monitored continuously using a Radiometer Model 22
pH meter and a combined calomel-glass electrode. The reactor effluent
23
-------
ro
P.
1
CaCl2
MgCl2
Solutions
Constant
Head Tank
$r
\
k. r
— i
Pump
Pumps
Stake and
Stator
Stirrer
CSTR
Rotameter
Effluent
Sedimentation Basin
Sediment
Drai n
FIGURE 2. SCHEMATIC FLOWSHEET FOR EXPERIMENTS ON CHEMICALLY-DEFIED SYSTEMS
-------
discharged through 1 1/2-in. PVC pipe. Samples for analysis of dissolved
Pp C-j-, Ca, and Mg were taken from the reactor effluent pipe after steady
state had been reached and filtered immediately through 0.45-y membrane
filters (Mi Hi pore HA). When it was desired to collect precipitate
samples for X-ray diffraction analysis, the reactor effluent was collected
in a 30-£ sedimentation basin.
In continuouous flow experiments employing a recycle of suspension, a
circular 12-4 concave-bottomed sedimentation basin followed the completely
mixed reactor in the flow-scheme illustrated in Figure 3. Suspension
was recycled to the reactor from the underflow of the sedimentation basin
using a peristaltic pump (Sigmamotor Model T65). In these experiments the
mean hydraulic residence times were 42 min in the reactor and 34 min in
the sedimentation basin. Sufficient feed solution for several days of
continuous operation was prepared by adding reagent grade CaCl2, MgCl2,
and H3P04 to 3000 a of well-mixed tap water held in a 1000-gal polyethy-
lene-lined redwood tank. The performance of the recycle reactor was
monitored on the basis of grab samples taken from the sedimentation basin
overflow.
3,000-4
Feed Water
Tank
Effluent
12-4
Clarifier
Suspension Recycle
FIGURE 3. SCHEMATIC OF CALCIUM PHOSPHATE PRECIPITATION
WITH SUSPENSION RECYCLE
Chemically-Defined Systems - Batch Experiments
Batch experiments were conducted on chemically-defined systems when it
was demonstrated that they could be expected to yield information that
was identical to that obtained in continuous flow, steady state system.
In these batch experiments the reactor was a magnetically stirred 4-4
beaker. Predetermined
mixed rapidly into 4 4
amounts of Na2C03 and
of distilled water in
PQk stock solutions were
batch reactor and the pH
25
-------
of the mixture adjusted immediately to the desired value with NaOH
solution. A mixture of the desired amounts of CaCl2 and MgCl2 stock
solutions were then added rapidly to the stirred solution and at that
time the measurement of reaction time was started. A 15-min reaction
time was employed in these experiments on the basis of results from
continuous flow studies. During this reaction time the pH was monitored
and maintained constant by NaOH addition. Samples for the analysis of
dissolved Cj, Pj> Mg» and Ca were withdrawn after 15 min and immediately
filtered through 0.45-y membrane filters.
Wastewater Experiments
Continuous flow wastewater precipitation experiments were conducted in
a 3-compartment epoxy-coated galvanized steel reactor of which each
compartment was 15 in. x 15 in. x 15.8 in. water depth. The contents
of each compartment were stirred by straight blade turbine paddles mounted
on a vertical shaft. The precipitation unit was an integral part of a
chemical-biological process train that, in addition to lime addition,
precipitation reactors, and two sedimentation basins, contained a
recarbonation basin and clinoptilolite columns for ammonia removal
(Figure 4).
Sufficient lime slurry for 24-hr operation was prepared by adding
quicklime to 100 a of tap water in a 30-gal plastic bin. The lime
slurry was made up to a concentration of either 300 mM or 600 mM, depending
on what pH of operation was desired. Lime slurry batches were analyzed
for CaO immediately after preparation and immediately prior to the
preparation of a fresh batch. Lime doses were computed on the basis of
these analyses combined with measurement of the volume of lime slurry
pumped. The lime slurry was introduced into the wastewater feed at a
point 2 in. ahead of the first reactor and at a rate that was determined
by a pH controller in the reactor, preset to the desired pH value.
After passing through the three reactor compartments, the wastewater
entered a circular concave-bottom sedimentation basin. The performance
of the precipitation unit was monitored by analysis of samples which
were collected continuously from the effluent of the sedimentation basin.
Samples of both influent and effluent were stored at 4°C between sampling
and analysis. In addition to the analyses conducted on the chemically-
defined systems, measurements of total orthophosphate and dissolved
hydrolyzable phosphate plus orthophosphate were made on samples of
influent, lime precipitated effluent, the recarbonated effluent, and the
clinoptilolite column effluent. Each experimental run with the waste-
water treatment plant lasted for at least four days.
26
-------
Lime
1,
1
A
A
Primary Activated rl , ..
Sedimentation Sludae Flocculation
<
Sedi
menta
tion
Recar
bona-
tion
Sedi-
menta-
tion
Ammonia
Sorption
Primary
Sedimentation
Activated
Sludge
ro
—i
Lime
Flocculation tion
Sedimenta^ """ Floccu- Sedl"-
Recar- lation menta-
bona- tion
tion
Ammonia
Sorption
FIGURE 4. TREATMENT SYSTEMS USED IN STUDIES AT SERL PILOT PLANT
-------
ANALYTICAL PROCEDURES
Sampling Procedures
Sampling procedures varied with the system under investigation and have
been discussed previously with descriptions of the experimental systems.
Analytical Methods
Chemical and physical analyses were generally performed in accordance
with Standard Methods [41], the FWPCA Methods for Chemical Analysis
[42], or the SERL Analytical Methods [43]. The following brief descrip-
tion of analytical methods is presented to indicate method selection
and any novel method modification employed.
Soluble Calcium. Soluble calcium was determined by EDTA titration
using hydro>
-------
. Six min reaction time was used instead of 10 min. The two
latter modifications were made to reduce the possibility of
polyphosphate hydrolysis during the orthophosphate analysis.
Qrtho + Hydrolyzable Phosphate. A 50-ms, sample was placed in a
125-m«, Erlenmeyer flask and 1 ml of strong acid solution (310 ml cone.
\\2SQk to 600 mil distilled water) was added. The sample was boiled gently
for 30 to 40 min. Phenolphthalein (3 drops) was followed by 1 N NaOH
until a pink color developed. This was discharged by the addition of
strong acid solution. The sample was cooled and diluted to 50 mi and
the analysis continued as for orthophosphate.
Inorganic Carbon. Inorganic carbon was determined with a Beckman
Carbon Analyzer, Model 215A.
A1 ka 1i ni ty . Alkalinity was determined by titration with standard acid
to a pH of 4.3.
Suspended Solids. Suspended solids were determined by the membrane
filter technique of Winneberger ejb al_. [44] using 0.45-y membrane filters.
Suspension Settling Tests. Settling tests were conducted on reactor
effluent by collecting it in a 4-& beaker and allowing it to settle
quiescently for 30 min. The settling properties of the suspension were
evaluated in terms of the total phosphate removal achieved during this
period by taking samples for total phosphate analysis before and after
settling, using a pipette to sample 1 in. below the surface of the
reactor effluent.
X-Ray Powder Diffraction Analysis. Suspensions for X-ray powder
diffraction analysis were removed from the bottom of the sedimentation
basin following the CSTR and centrifuged at about 1000 rpm. After
decanting the supernatant, the suspension was quickly frozen by immersion
in a dry ice-n-butyl alcohol mixture. The frozen suspension was
lyophilized for a period of 48 hr. The X-ray powder diffraction pattern
was determined using a Norelco diffractometer and scanning the spectrum
from 4 degrees to 60 degrees. Because early experience demonstrated
the presence of noncrystalline materials, lyophilization was later
replaced by air-drying.
Lime Purity. A 0.5-gram finely pulverized sample was dissolved in
100 ma of a 10% sugar solution after one hour on a shaking machine.
The solution was filtered and a 25-ms, aliquot was titrated with 0.1 N
sulfuric acid, using phenolphthalein indicator.
Total Ca(OH)2. Total Ca(OH)2 was determined on a sample of lime
siurry after ju min refluxing with HC1. A suitable sample volume
(usually 10 ml) was added to 25.00 mi of 1 N HC1 solution and the
mixture was refluxed for 30 min. After cooling to room temperature the
excess HC1 was back-titrated with 1 N NaOH using phenolphthalein
indicator.
29
-------
SECTION VIII
PRECIPITATION OF CALCIUM PHOSPHATE IN
CHEMICALLY-DEFINED SYSTEMS
INTRODUCTION
Batch and CSTR work in chemically-defined systems took the form of a
series of experiments in which one or another of the components of the
Ca - Cj - Pj H+ - Mg - H20 system was varied singly to determine its
effect on the residual phosphate or other component concentrations. The
range over which individual component concentrations were varied was
selected to cover the range commonly encountered in wastewater treatment.
Nominal residence times of 15 min were generally used in CSTR experiments
and a reaction time of 15 min was used in the batch experiments. Some
kinetic studies employed variable detention (CSTR) and reaction (batch)
times.
TIME TO ESTABLISH STEADY STATE
Prior to conducting experiments to determine the effect of various solu-
tion components on the phosphate residuals obtained by calcium phosphate
precipitation, it was necessary to determine the time required to reach
steady state values of phosphate residual that were representative of
these precipitation processes. From the previous discussion of the phases
involved in the precipitation of calcium phosphate from aqueous solution,
it might be surmized that steady state dissolved phosphate residuals
might be reached rather rapidly — at least within a few minutes of mixing
reactants to produce a supersaturated solution. The period over which
this steady state dissolved phosphate level is maintained is of importance
since it is necessary to know whether measurements made soon after the
initial attainment of steady state are representative of the conditions
that might exist in phosphate, precipitation reactors of longer residence
times.
To establish the time required to reach steady state and the duration of
this steady state condition, a series of batch and CSTR experiments were
conducted under conditions typical of wastewaters. Three series of batch
experiments at 23°C were conducted for systems containing Ca and Pj only,
Ca, Pj, and Cy, and Ca, Pj, Cj, and Mg. The initial concentration of
each of these components is summarized in Table 2. Each of these systems
was examined at pH values of 8.0, 9.0, 10.0, and 11.0 -values that were
maintained constant throughout the experiments by addition of NaOH
31
-------
TABLE 2
INITIAL COMPOSITION OF BATCH SYSTEMS USED TO
DEFINE TIME TO REACH STEADY STATE
Run
I
II
III
mM
Ca
2.0
2.0
2.0
PT
0.5
0.5
0.5
CT
-
2.0
2.0
Mg
-
-
2.0
Figure 5 shows that in each of the batch experiments steady state was
reached in less than 10 min. Neither the introduction of carbonate
(Run II) nor carbonate and magnesium (Run III) into the system lengthened
the time required to reach steady state, although changes in the levels
of steady state phosphate concentration were observed.
CSTR experiments at pH 8 were conducted with nominal reactor detention
times varying from 3 to 200 min and with the initial concentrations of
reactants reported in Figure 6. Because of their time-consuming nature
and because of the results of the previous batch experiments, the CSTR
runs were only conducted for a limited range of conditions. Thus only
solutions containing Ca and Pj were tested in CSTR runs because batch
experiments had indicated that the presence of Mg and Cj did not affect
the time taken to attain steady state. Also since the batch experiments
indicated a slightly slower rate of attainment of steady state at pH 8,
the CSTR experiments were conducted at this pH value.
In all of the CSTR experiments, steady state dissolved phosphate concen-
trations were reached in less than a 10-min mean residence time and main-
tained for up to at least 200 min — a residence time typical of a phosphate
precipitation process. On the basis.of these results, reaction times of
15 min (batch or CSTR) were selected as being truly representative of the
steady state conditions existing in nonrecycle wastewater calcium phos-
phate precipitation processes.
THE NATURE OF THE STEADY STATE SOLID PHASE
It is important to determine the solubility characteristics and the nature
of the steady state solid phase that forms under conditions representative
of wastewater calcium phosphate precipitation processes. An insight into
the stoichiometry and solubility behavior of this solid was obtained by
32
-------
10
-1
pH 8.0
Run No.
I 0
II Q
III A
Initial Conditions"
mM
Ca
2.0
2.0
2.0
PH
PT
0.5
0.5
0.5
CT
2.0
2.0
Mg
2.0
9.0
S-:^—.-!).-
J-
Batch, 23°C
;-%_
10
-2
_.Q El-
pH 10.0
J§ 1
a
oVnt
PH n.o
to -a -Q—
E
10
-3
15
30
45
TIME, min
FIGURE 5.
DEPENDENCE OF STEADY STATE PT CONCENTRATION
ON SOLUTION COMPOSITION AND pH IN BATCH
REACTORS
33
-------
0.6
0.5
0.
0.4
0.3
0.2
0.1
1 1
V D D
a
V ° o
I
Run
A
o
a
0
Initial Conditions, mM
Ca Pj Mg CT
3.7 to 4.2 1.0 to 1.1 0.0 0.0
3.7 to 4.0 0.5 to 0.6 0.0 0.0
1.8 to 2.1 1.0 to 1.1 0.0 0.0
2.1 to 2.2 0.5 to 0.6 0.0 0.0
0 " ° 20° to 24°C
£* CSTR
V V A &
i i
—
<& <
\
50
100
DETENTION TIME (V/Q), min
150
200
FIGURE 6. DEPENDENCE OF STEADY STATE PT CONCENTRATION ON DETENTION TIME
IN CSTR EXPERIMENTS
-------
analysis of both steady state solid and liquid composition of batch and
CSTR experiments.
CSTR experiments were conducted between 19° and 23°C with chemically-
defined solutions in which initial component concentrations were in the
ranges: Ca, 2.1 to 2.4 mM; Mg, 0 to 0.6 mM; PT, 0.3 to 0.5 mM; CT, 0 to
14.8 mM; pH 8 to 11. The initial mole ratios of the various components
varied as follows: Ca/P, 3.9 to 6.7; Ca/Mg, 2.1 to 4.8; Ca/Cj, 0.2 to
2.4.
Batch experiments with 15 min reaction time were conducted at 19° to
23 C using initial solution concentrations in the range: Ca, 0.4 to
5.5 mM; Mg, 0.1 to 0.5 mM; PT, 0.3 to 0.4 mM; CT, 0.8 to 2.7 mM; pH 8
to 11. The initial mole ratios of the various components varied as
follows: Ca/PT, 1.1 to 17.2; Ca/Mg, 6.9 to 12.9; Ca/Cj, 2.0 to 7.3.
In all of these experiments nucleation proceeded without detectable lag
period and steady state dissolved phosphate levels were reached within a
10-min period. The nature of the steady state solid can be circumstan-
tially deduced, and a useful tool to predict steady state dissolved
phosphate residual can be obtained by computing the activity product of
various postulated solids from the steady state dissolved component data.
This procedure was initially performed for selected experimental data for
the several solids that have been postulated as apatite precursors and
therefore might be expected to be present under steady state conditions
prior to transformation to an equilibrium solid. Figure 7 is a plot of
the negative logarithm of the activity product (pA) versus pH for the
solids, dicalcium phosphate (DCP) (CaHPOiJ, octacalcium phosphate (OCP)
(Ca^FKPOiJo), hydroxyapatite (HAP) (Ca5(OH)(POlf)3) and tricalcium phos-
phate (TCP) (Ca3(POlt)2). It is evident from Figure 7 that only the
activity product for TCP appears to be constant over the entire pH range
(8 to 11) of the experiments. Following this evidence further, Figure 8
presents a plot of pA of TCP for all data points (both batch and CSTR)
collected in chemically-defined systems. The regression line fitted to
these data points has the form:
pA = -0.064 pH + 24.15
(8)
indicating the virtual constancy of the activity product with pH over the
range pH 8 to 11.
This value of the negative logarithm of the activity product of TCP which
averages 23.56 is somewhat lower than values reported in the literature
which range from approximately 25 to 29 [45-47]. By using the plotting
technique of Walton ejt al_. [26] (Appendix A), which determines the
stoichiometry of a nucleating solid phase, Figure 9 shows that the Ca/P
mole ratio in the precipitate is 1.44. This value is very close to the
stoichiometry of TCP which would give a theoretical Ca/P mole ratio of
1.5
35
-------
38
36
34
32
30
28
26
24&
22
9
pH
11
FIGURE 7.
EFFECT OF pH ON pA FOR VARIOUS CALCIUM
PHOSPHATE SOLIDS
36
-------
25
1 1
Reactor
Batch o
CSTR A
Initial Conditions, mM
Ca
0.4-0.5
2.1-2.4
Mg
0.1-0.5
0.0-1.9
PT
0.3-0.4
0.3-0.4
CT
0.8-2.7
0.0-14.8
L
243-pA = -0.064 pH + 24.15 ©
o
Q.
CO
(O
O
-------
co
CO
n
©O
12
19° to 23°C
15 min
&A3&A
Reactor
Batch o
CSTR A
Initial Conditions, mM x 103
Ca
0.4-5.5
2.1-2.4
Mg
0.1-0.5
0.0-0.6
PT
0.3-0.4
0.3-0.4
CT
0.8-2.7
0.0-14.8
13
Ca/P = 1.44
i
A
A
A
A
18
19
20 21
FIGURE 9. CRITICAL IONIC CONCENTRATIONS IN BATCH AND CSTR EXPERIMENTS
-------
X-ray powder diffraction analysis of the solids produced in CSTR experi-
ments failed to reveal the presence of any crystalline calcium phosphate.
For experiments conducted at pH 10 and higher, crystalline.calcite (CaC03)
was detected. The X-ray analyses indicate that the steady state calcium
phosphate solid is amorphous in nature - a fact that is supported by the
lower pA values obtained here for TCP in comparison to those values
reported in the literature.
EFFECT OF SOLUTION COMPOSITION ON STEADY STATE
PHOSPHATE RESIDUAL
Even though the evidence produced in the previous section tends to support
the presence of TCP as a steady state solid phase, there is a considerable
range of variation in the values of the activity product of TCP at each
of the pH values examined. While such variation (over about 1 to 2 orders
of magnitude in pA) is'common in equilibrium data derived from hetero-
geneous systems, it is necessary to determine whether these variations
are in any way related to the concentrations of other important solution
components — especially to the concentration of Cj and Mg. In the data
analysis that follows, the effect of both of these components on the
activity product of TCP and composition of the steady state solid phase
is examined.
Effect of Magnesium
Previous reports have indicated that magnesium concentration is an impor-
tant factor in calcium phosphate precipitation under conditions typical
of wastewaters. Its effects have been classified as equilibrium effects
(in which the solid calcium phosphate phase is altered by the presence of
magnesium) and kinetic (in which the rate of attainment of equilibrium
or steady state phosphate residuals is influenced by magnesium concen-
tration). An additional "process effect" can be visualized, i.e., in
which magnesium influences the separability of precipitated solids.
The influence of magnesium on steady state dissolved phosphate residual
was tested in five series of CSTR experiments for which the initial
component concentrations are listed in Table 3.
For the range of magnesium concentrations (more correctly Ca/Mg ratios)
commonly encountered in wastewaters, the pA of TCP does not appear to be
affected by magnesium concentration. Thus Figure 10 shows that only when
initial Ca/Mg ratios were less than 0.6 (commonly most wastewaters have
Ca/Mg ratios in the range of 1.5 to 6) did the activity product of TCP
increase significantly. These results were supported by an additional
series of CSTR experiments in which the steady state levels of dissolved
PT, CT, Ca, and Mg were measured as a function of varying initial
magnesium concentration in the range 0.5 to 10.9 mM. In this series of
experiments the pH was maintained constant at 9.5, and the initial
component concentrations were constant at Ca, 2.2 mM; Py, 0.38 mM, and
39
-------
25
24
O
Q-
co
to
O
23
22
Run
A
E
O
e
Initial Conditions
mM x 103
Ca
2.2
2.1
2.2
2.2
Mg
0.6
2.3
3.6
10.8
PT
0.4
0.4
0.4
0.4
CT
2.7
4.1
2.6
2.4
Ca/Mg
3.7
0.9
0.6
0.2
19° to 23°C
15 nrin
CSTR
10
11
pH
FIGURE 10. EFFECT OF Mg ON pA Ca3(P04)2
40
-------
TABLE 3
INITIAL CONCENTRATIONS IN CSTR EXPERIMENTS
TO DETERMINE EFFECT OF MAGNESIUM
Experiment
No.
1
2
3
4
mM
Ca
2.2
2.1
2.2
2.2
PT
0.4
0.4
0.4
0.4
CT
2.7
4.1
2.6
2.4
Mg
0.6
2.3
3.6
10.8
Mole Ratio
Ca/PT
6.0
5.8
6.1
6.1
Ca/Mg
3.7
0.9
0.6
0.2
Ca/CT
0.8
5.1
0.8
0.9
Cj, 2.5 mM. The data in Figure 11 indicate that no significant increase
in residual dissolved phosphate occurred until the initial magnesium
concentration was greater than 2 mM resulting in an initial Ca/Mg mole
ratio of less than 1.
As in previous experiments, X-ray powder diffraction analyses failed to
detect crystalline calcium phosphate in any of the solids derived from
these CSTR experiments. Crystalline calcite (CaC03) was detected in
suspensions produced from solutions containing initial calcium concen-
trations of 2.1 to 2.2 mM and initial magnesium concentrations of 0.5 mM.
Suspensions produced from solutions in which the initial calcium concen-
tration was 2.1 mM and the initial magnesium concentration was 2.3 mM
failed to show the presence of crystalline calcite even after 1000 hr of
contact between the suspension and the solution from which it was produced.
After 4000 hr of solution-suspension contact a crystalline calcium
carbonate, designated as Calcite III, was detected in the suspension.
Effect of Carbonate
The effect of carbonate on the precipitation of calcium phosphate from
wastewaters has been interpreted both in terms of a competition between
the calcium phosphate and calcium carbonate solids for calcium ion and
by the modification of calcium phosphate solids such as apatites by the
incorporation of carbonate. Both of these interpretations predict that
the effect of increasing carbonate concentration will generally be to
increase dissolved phosphate residual at a given calcium dose or conversely
to require a higher calcium dose to reach the same phosphate residual.
Experiments on chemically-defined solutions were undertaken in CSTR
systems with 15 min nominal residence time to determine the significance
of the effect of carbonate on dissolved phosphate residuals at concen-
trations and under conditions that typified wastewater calcium phosphate
41
-------
12
10
8
6
4
2
0
Initial Conditions, mM
Ca
2.2
Mg
0.5-10.9
PT
0.38
CT
2.5
15 min
23°C
CSTR
pH 9.5
0123
4567
Influent Mg, nM
Mq
8 9 10 11
FIGURE 11. INFLUENCE OF Mg ON STEADY STATE
CONCENTRATIONS OF Ca, Mg, CT, AND
PT AT pH 9. 5
42
-------
precipitation processes. Initially four series of CSTR experiments were
conducted with the initial conditions as outlined in Table 4. The results
of these experiments, presented in Figure 12, show that up to. initial Cj
concentrations of 10.6 mM no significant effect on the activity product
of TCP was detectable. At initial Cj levels of between 10.6 and 14.8 mM
the activity product of TCP increased significantly for pH values above
9.5.
TABLE 4
INITIAL CONDITIONS FOR CSTR EXPERIMENTS TO
DETERMINE EFFECT OF CARBONATE (CT)
Exp.
No.
1
2
3
4
pH
Range
8 -11
8 -11
8 -11
8.5-11
mM
Ca
2.1-2.4
2.2
2.1-2.3
2.1-2.3
Mg
0-0.5
0.6
0-0.6
0.6
PT
0.3-0.5
0.4
0.4-0.5
0.4
CT
0
2.5-2.7
5.0-7.2
10.6-14.8
Mole Ratio
Ca/Mg
4.8-6.0
3.7
3.8-5.2
3.5-3.9
Ca/Cj
-
0.8-0.9
0.4
0.2
Ca/PT
3.9-6.5
6
4.2-6.2
3.9-6.5
The effect this increase in activity product had on phosphate residual
at pH 9.5 and 11 in systems containing calcium, magnesium, phosphate,
and carbonate is illustrated in Figures 13 and 14 which are the results
of 15-min nominal residence time CSTR experiments conducted at the range
of initial experimental conditions depicted in Table 5.
TABLE 5
INITIAL CONDITIONS FOR CSTR EXPERIMENTS TO
DETERMINE EFFECT OF CARBONATE ON PHOSPHATE
AND OTHER COMPONENT RESIDUALS
PH
9.5
11.0
mM
Ca
2.0-2.2
2.2-2.3
Mg
0.6
0.6
PT
0.4
0.4
CT
2.6-14.4
2.5-14.8
Mole Ratio
Ca/Mg
3.4-3.7
3.7-3.8
Ca/Cj
0.1-0.9
0.1-0.9
Ca/PT
4.4-6.1
6.0-6.1
43
-------
25
Run
A
Q
O
O
Initial Conditions, mM x 103
Ca
2.1-2.4
2.2
2.1-2.3
2.1-2.3
Mg
0.0-0.5
0.6
0.0-0.6
0.6
PT
0.3-0.5
0.4
0.4-0.5
0.4
CT
0.0
2.5-2.7
5.0-7.2
10.6-14.8
22
11
FIGURE 12. EFFECT OF CT ON pA Ca3(P04)2
44
-------
ro
O)
1.5
1.0
0.5 -
0.0
1.6
1.4
1.2
1.0
0.8
Initial Conditions, mM
Ca
2.0-2.2
Mg
0.6
PT
0.4
CT
2.6-14.4
8 10
INFLUENT CT, mM
12
CSTR
pH 9.5
20° to 23°C
15 min
18.0
16.0
14.0
12.0
CM
10.0 °
X
8.0 I
6.0 ^
4.0
2.0
0.0
14
16
FIGURE 13. EFFECT OF CT ON STEADY STATE CONCENTRATIONS OF Ca, Mg, AND PT
AND ON AMOUNT OF PRECIPITATED CARBONATE AT pH 9.5
-------
4.0
3.0
2.0
1.0
0.0
0.7
0.6
O)
-a
c
to
10
0.5
0.4
0.3
Initial Conditions, mM
Ca
2.2-2.3
Mg
0.6
PT
0.4
CT
2.5-14.8
19° to 20°C
0.9
0.7
CM
O
0.5
0.3
0.1
8 10
INFLUENT CT, mM
12
14
16
FIGURE 14. EFFECT OF CT ON STEADY STATE CONCENTRATIONS OF Ca, Mg, AND PT
AND ON AMOUNT OF PRECIPITATED CARBONATE AT pH II.0
-------
At both pH values there is a very gradual increase in residual dissolved
phosphate up to initial carbonate concentrations of 10 mM. Thereafter an
extremely rapid increase in residual dissolved phosphate occurs at both
pH values. These data are consistent then with the effects of carbonate
on the activity product of TCP. Evidence exists from these data that
carbonate is incorporated into the solids increasingly with increasing
initial carbonate concentration. There is consistently higher carbonate
incorporation into the solids at pH 11.0 than at pH 9.5 for all initial
carbonate concentrations. Evidence that calcium carbonate precipitation
occurred in both of these experiments is offered by additional CSTR data
depicted in Figure 15. Here the ratio of Ca/PT in the precipitate is
plotted versus pH on the same graph as the CT removed from solution. The
curves have almost identical shapes and support the observation that
UJ
UJ
Cf-
CL
D.
O
i—i
I—
-------
calcium carbonate precipitation occurs at pH 9.5 and increasingly there-
after up to pH 11.0. These data also indicate that below pH 9.0 to 9.5
in systems with typical wastewater component concentrations, calcium
carbonate precipitation is not significant.
Examination of the amount of carbonate and calcium removed from solution
in these experiments suggests that more carbonate disappears from solution
than can be accounted for by the precipitation of calcium carbonate (with
a Ca/Cj mole ratio of 1). It must therefore be assumed, as has been
indicated in the literature, that incorporation of the carbonate into
the calcium phosphate solid occurs in such systems. This phenomenon has
been reported to produce calcium phosphate solids whose solubility
increases with their carbonate content. It is possible that the gradual
increase of residual dissolved phosphate with increasing initial carbonate
concentration, noted in Figures 12 and 13, and the gradual increase in TCP
activity product, noted in Figure 11, is caused by carbonate inclusion
in the calcium phosphate solid.
In these CSTR experiments Figures 13 and 14 demonstrate that for the entire
range of initial carbonate concentrations and for both pH values of 9.5
and 11 there appears to be very little if any incorporation of magnesium
into the solid.
It might be concluded then, that for the purpose of predicting dissolved
phosphate residuals, the activity product of TCP derived in these experi-
ments can be used satisfactorily for waters containing carbonate concen-
trations of below 4 mM (or alkalinities of below 400 mg CaC03/£) - values
of alkalinity that are rarely exceeded in wastewaters.
METHOD OF PREDICTING DISSOLVED PHOSPHATE RESIDUAL
CSTR and batch experiments have shown that for the range of component
concentrations commonly found in wastewaters, dissolved phosphate resi-
duals may be predicted from the activity product of TCP as follows:
[Ca2+]|s [PO;-]S2S = A (8)
where the subscript ss stands for steady state. The mean value for the
negative logarithm of the activity product was determined to be 23.56.
From a materials balance [Ca2+]ss can be replaced by
using the fact developed from Figure 9 that the Ca/P mole ratio of the
precipitation approximates 1.5. Introducing a value of lo~23-56 for the
48
-------
activity product and substituting Equation (9) into Equation (8) with
the elimination of [Ca2+]ss, the logarithmic form of the equation becomes
3 log {[Ca2+]in - 1.5 [P1n
2 log
= -23.56
(10)
In this equation the only unknown is Pss, the steady state concentration
of total dissolved phosphate, since the value of PO^" can be expressed
as a function of pH, total dissolved phosphate, and equilibrium constants
as follows:
ss
[H+]/k3 + [H+?/k2k3 + [H+]3/k1k2k3]
(11)
where kl5 k^, and k3 are the first, second, and third equilibrium constants
of phosphoric acid. Equation (10) can be solved manually for Pss as
illustrated in Appendix B.
Using this equation to predict phosphate residuals from the CSTR and
batch experiments in chemically-defined systems gave excellent results
(within 25% of experimental values) for systems with pH between pH 8 and
10, with initial magnesium up to 2 mM, and initial carbonate up to 4 mM
(Table 6).
At pH 11.0 an error of 68% between predicted and experimental values was
obtained for initial carbonate concentrations up to 3 mM. This error
increased slightly (the predicted values were lower than experimental
values) to the neighborhood of 73% when the initial carbonate concen-
tration was increased to 6 mM. The increasing error in predicted values
at both increased pH and carbonate concentration is undoubtedly caused
by calcium carbonate precipitation which reduces the effective initial
value of calcium. These errors were reduced so that predicted and
experimental values of dissolved phosphate residual were within 25%
when the precipitation of calcium carbonate was taken into consideration
(Table 7).
CALCIUM PHOSPHATE PRECIPITATION AT pH
SUSPENSION RECYCLE
8 WITH
As previously indicated, the calcium phosphate precipitation model of
Stumm and Leckie [18] visualized precipitate formation in three steps —
nucleation of an amorphous phase, phase transformation, and growth of a
crystalline phase. In the short-term calcium phosphate precipitation
processes examined in the previous part of this chapter it was postulated
that the solid present in the phase transformation step would control the
49
-------
TABLE 6
COMPARISON OF PREDICTED AND EXPERIMENTAL RESIDUAL
PHOSPHATE VALUES IN CHEMICALLY-DEFINED SYSTEMS
PH
8.0
9.0
9.5
10.0
n.o
pA Ca3(P0lt)2
Range
High 22.92
Low 24.35
Mean 23.69
High 22.80
Low 23.68
Mean 23.49
High 22.96
Low 24.08
Mean 23.71
High 22.77
Low 24.03
Mean 23.74
High 22.44
Low 24.44
• Mean , 23.55
Dissolved Py, mM x lO1*
Experimental
2100
4200
4500
1200
620
550
280
no
34
130
230
48
94
3.3
6.5.
Predicted
1000
1300
4280
500
530
500
185
177
31
43
222
52
13.5
5.1
6.4
Error3
%
-57.2
-69.0
- 4.4
-58.4
-14.5
- 9.1
-34.0
+60.7
- 8.8
-67.0
+ 3.5
+ 8.3
-85.6
+54.6
- 1.5
% Error =
x 100.
dissolved phosphate residual. It was found that at residence times
typical of the processes this steady state solid has the characteristics
of an amorphous tricalcium phosphate.
There is considerable evidence [3,4,18] that the phase transformation
step can be shortened (or even eliminated) by the presence of preformed
crystalline material in the precipitating medium. Indeed the reported
50
-------
TABLE 7
COMPARISON OF PREDICTED AND EXPERIMENTAL RESIDUAL
PHOSPHATE VALUES IN CHEMICALLY-DEFINED SYSTEMS
CORRECTED FOR CALCIUM CARBONATE PRECIPITATION
pH
11.0
11.0
11.0
CT
mM
1.5
3.0
6.0
Dissolved PT, mM x TO1*
Predicted
Uncorrected
for CaC03
Precipi-
tation
5.16
5.16
5.11
Corrected
for CaC03
Precipi-
tation
3.90
14.1
13.4
Experi-
mental
>
3.23
16.1
18.7
Error9, %
Uncorrected
for CaC03
Precipi-
tation
+60
-68
-73
Corrected
for CaC03
Precipi-
tation
+22
-13
-25
a% Error =
P - E
x 100.
success of the Phosphate Extraction Process [16], in which precipitated
solids are recycled, is attributed to providing a high concentration of
solids to encourage more rapid crystal growth of calcium phosphate.
Because of these observations, a limited investigation was conducted on
chemically-defined systems to determine the feasibility of obtaining low
phosphate residuals at pH 8 through the recycle of preformed precipitate.
Two CSTR experiments, each of about 2 weeks duration, were
a single CSTR followed by a sedimentation basin from which
pension could be recycled to the CSTR (Figure 3).
conducted using
settled sus-
In the first experiment an attempt was made to build up precipitate by
recycling sedimentation basin underflow from a solution containing the
following initial component concentrations: Ca, 2.2 mM; Mg, 0.5 mM;
PT, 0.3 mM; CT, 3.2 mM; pH 8. After a period of 13 days i,t became evident
that it would not be possible to build up a high precipitate level in the
reactor because of the poor settling characteristics of the precipitate.
Indeed, throughout the almost two weeks of continuous operation the
reactor suspension concentration did not exceed 6 mg/£ and no significant
phosphate removal was achieved. ' . _t
In the second experiment an aged calcium phosphate-carbonate suspension
was preformed in the reactor by adding 3 a, 10 M CaCl2 solution to 10 £
51
-------
distilled water followed by 750 mi cone. H3POi,. and then sufficient 50%
NaOH solution to achieve a pH of 8. After aging the stirred suspension
for 2 days the CSTR was continuously fed with a solution whose initial
composition was: Ca, 2.2 mM; Mg, 0.5 mM; Pj, 0.3 mM; Cj, 3.2 mM; pH 8.2.
Settling problems were again encountered and the initial suspension
concentration of 5900 mg/fc fell rapidly during the first 2 days of opera-
tion to a level of approximately 1500 to 2000 mg/£. Over the one week
period that suspension concentrations of 1500 to 2000 mg/£ were maintained,
it was possible to remove 60% of the incoming phosphate. Following this
period of operation, continuing difficulties in suspension separation
caused a further decrease in reactor suspension concentration to about
50 mg/x. at the termination of the experiment (Figures 16 and 17). At
this time the dissolved phosphate removal had decreased to 20%.
At the same time that these CSTR experiments were proceeding, batch
experiments were conducted on solutions of identical composition to the
CSTR feed. Suspension from the CSTR reactor was introduced into two 4-a
beakers containing CSTR feed solution to achieve initial levels of 3 and
3000 mg/i suspension, respectively. The batch reactors were stirred for
29 hr. Figure 18 shows that the dissolved phosphate concentration in the
high suspension concentration reactor decreased more rapidly and to a
lower level than the dissolved phosphate concentration in the low suspen-
sion concentration reactor. A further illustration that the presence of
calcium carbonate-phosphate suspension enhances phosphate removal was
provided by a series of experiments in which solutions of composition
identical to the CSTR feed were filtered through layers of suspension
taken from the CSTR reactor that had been deposited on the surface of
membrane filters in a 2-cm deep bed. When the filtrate composition was
analyzed for phosphate it was revealed that one passage of CSTR feed
solution through such a precipitate bed effected a 54% removal of phosphate;
this was increased by successive refiltrations to 56% and 60% removal.
These results are of significance because the conditions in such a
refiltration experiment might be thought of as being somewhat similar
to those existing in the sludge blanket of an upflow clarifier. It
might be surmised that such an expanded bed of precipitate could have
many of the important features (albeit with a lower residence time) of a
reactor to which precipitate is recycled.
X-ray powder diffraction analysis of the suspensions from continuous,
batch, and refiltration experiments provided an important piece of evi-
dence in support of the postulate that the steady state solid was a
tricalcium phosphate. The presence of Ca3(POit)2-nH20 was detected by
these analyses and by chemical analysis of the solid. The value of n
was found to be 4.
i i
The effect of recycled solids concentration on the activity product of
TCP is shown in Figure*19. The close correlation between activity product
and suspended precipitate concentration serves to emphasize the catalytic
effect of solids in the formation of more insoluble phases.
These experiments clearly indicate that low phosphate residuals can be
achieved at pH 8 from waters with a mineral composition typical of many
52
-------
0.5
0.4
0.3
0
0.1
Feed
Run 2
Ca - 2.2
Mg 0.5
Pj = 0.32
CT = 3.2
inM
1 T
pH = 8 Preformed Solids
21°C
^"t*"^
6,000
5,000 g
4,000 o
3,000 o
2,000
1,000
0.
co
60
0 1 23 4 567 8 9 10 11 12 13 14
TIME, days
FIGURE 16. EFFECT OF SUSPENSION CONCENTRATION
ON EFFLUENT PHOSPHATE CONCENTRATION
IN PRECIPITATE RECYCLE EXPERIMENTS
01 23 4 5678 9 10 11 12 13 14
FIGURE 17. STEADY STATE Ca AND CT CONCENTRATIONS
DURING PRECIPITATE RECYCLE EXPERIMENTS
53
-------
en
•£»
I I I ' I I I I
Suspended Solids, 3 mg/8.
Suspended Solids, 3000 mg/fc
pH = 8.2
20°C
8
10 12 14 16 18 20
TIME, nr
24 26 28 30
FIGURE 18. EFFECT OF SUSPENSION CONCENTRATIONS ON RATE OF CALCIUM
PHOSPHATE PRECIPITATION IN BATCH EXPERIMENTS
-------
Run 2
Feed, Ca, 2.2 rrfl
Mg, 0.5
PT, 0.32
CT> 3.2
4 5 6 7 8 9 10 11 12 13 14
6,000
4,500
3,000
Q
LU
a
1,500
FIGURE 19. EFFECT OF SUSPENSION CONCENTRATION ON
pA Ca3(P04)2 IN PRECIPITATE RECYCLE
EXPERIMENTS
domestic wastewaters by the use of precipitate recycle to produce high
suspension concentrations for calcium phosphate crystal growth. The
experiments also emphasize the importance of precipitate separation in
calcium phosphate precipitation processes and the absolute necessity for
development of methods to enhance the settling properties of the calcium
phosphate-carbonate suspensions produced at pH values below 10.
55
-------
SECTION IX
PRECIPITATION OF PHOSPHATE IN WASTEWATER
Five series of phosphate precipitation experiments were conducted on
wastewater in the SERL wastewater treatment facility. Each experimental
run lasted for at least 4 days. In the first four experiments the phos-
phate removal performance of each individual unit of the entire treatment
train was examined, including primary sedimentation, activated sludge,
lime precipitation, recarbonation, and clinoptilolite sorption. In the
fifth experimental run the phosphate removal performance of a lime
precipitation unit operating on the primary effluent was assessed.
OVERALL PHOSPHATE REMOVAL PERFORMANCE
The phosphate removal performance over the entire experimental period
for the whole process train is summarized in Figure 20. Performance
data are given for actual phosphate removal as well as maximum possible
phosphate removal —i.e., the removal of phosphate that would occur if
perfect separation of particulates were possible.
Daily experimental values derived from 24-hr composite samples for pH,
Ca, Mg, alkalinity and dissolved and total phosphate are presented in
Appendix C. Mean values of these parameters, together with the mean
values of imposed operating conditions, are presented in Tables 8 through
12.
In experiment 1 lime slurry was added to activated sludge effluent and
precipitation conducted in a 3-compartment mixer-flocculator with nominal
residence time of 46.5 min in a run of 8 day's duration. Lime doses of
2.9 mM (290 mg as CaCO^) raised the pH of the activated sludge effluent
feed from 7.2 to 10.8 and reduced the total phosphate from an average
influent value of 11.1 mg P/i to a residual of 0.5 mg P/s, in the effluent
from the clinoptilolite column (Table 8). In this experiment, as was the
case in all of the wastewater runs, the dissolved phosphate concentration
leaving the clarifier following the precipitation unit was lower than
that following recarbonation. In this experiment, average dissolved
phosphate levels of 0.12 mg P/a were present in the settled effluent from
precipitation while the dissolved phosphate following recarbonation of
this effluent was 0.50 mg P/£. Subsequent experiments showed that virtually
all of the particulate phosphate that escaped sedimentation redissolved
in the lower pH medium of the recarbonation basin within the 5-min average
detention time of this basin. Indeed, following recarbonation only an
insignificant fraction of the phosphate was present in the effluent as
particulate matter. It is this factor that accounts for the observation
57
-------
TABLE 8
OPERATING CONDITIONS AND RESULTS FOR WASTEWATER
EXPERIMENT I
Sample Location
Primary Effluent
Activated Sludge
Effluent
Lime Precipitation
Effluent
Recarbonation
Effluent
Clinoptilolite
Effluent
P
Total
roM
0.37
0.36
0.017
Dissolved
raM
0.31
0.30
0.00389
0.016
0.016
Ca
Dissolved
mM
1.38
1.39
1.93
-
1.72
Mg
Dissolved
mM
0.37
0.37
0.26
0.26
0.25
CT
Dissolved
mM
-
1.97
0.67
Alkalinity
Dissolved
mM
2.65
0.85
1.50
1.75
1.76
PH
7.4
7.2
10.8
8.5
8.5
Feed to precipitation unit: Activated sludge effluent; Number of compartments: 3; Nominal
residence time: 46.5 min; Total lime added: 2.9 mM; Dissolved lime added: 0.1 mM.
TABLE 9
OPERATING CONDITIONS AND RESULTS FOR WASTEWATER
EXPERIMENT 2
Sample Location
Primary' Effluent
Activated Sludge
Effluent
Lime Precipitation
Effluent
Recarbonation
Effluent
Clinoptilolite
Effluent
P
Total
mM
0.39
0.35
0.053
0.056
0.056
Dissolved
mM
0.34
0.0184
0.053
0.054
Ca
Dissolved
mM
1.36
1.37
1.76
1.72
1.14
Mg
Dissolved
mM
0.42
0.41
0.36
0.36
0.37
CT
Dissolved
mM
2.10
1.95
Alkalinity
Dissolved
mM
2.48
0.84
1.77
1.79
1.78
PH
7.4
7.3
918
8.2
8.4
Feed to precipitation unit: Activated sludge effluent; Number of compartments: 3; Nominal
residence time: 46.5 min; Total lime added: 1.6 mM; Dissolved lime added: 0.1 rtM.
58
-------
TABLE 10
OPERATING CONDITIONS AND RESULTS FOR WASTEWATER
EXPERIMENT 3
Sample Location
Primary Effluent
Activated Sludge
Effluent
L1me Precipitation
Effluent
Recarbonation
Effluent
Clinoptilolite
Effluent
P
Total
mM
0.37
0.0089
0.0076
0.0076
Dissolved
mM
0.28
0.00323
0.007
0.007
Ca
Dissolved
nfl
1.20
1.51
1.90
1.41
Mg
Dissolved
mM
0.38
0.24
0.26
0.26
CT
Dissolved
ntt
3.8
0.6
-
5.0
Alkalinity
Dissolved
mM
2.35
2.46
2.94
2.97
pH
7.4
11.0
7.7
7.8
Feed to precipitation unit: Primary effluent; Number of Compartments: 3; Nominal residence
time: 46.5 min; Total lime Added: 4.9 mM; Dissolved lime added: 0.2 mM.
TABLE 11
OPERATING CONDITIONS AND RESULTS FOR WASTEWATER
EXPERIMENT 4
!•
Sample Location
Primary Effluent
Activated Sludge
Effluent
Lime Precipitation
Effluent
Recarbonation
Effluent
Clinoptilolite
Effluent
P
Total
mM
0.35
0.018
0.027
0.021
Dissolved
mM
0.28
-
0.0074
0.023
0.019
Ca
Dissolved
nti
1.10
-
1.07
1.26
1.24
Mg
Dissolved
mM
0.37
0.31
0.30
0.33
CT
Dissolved
mM
3.2
1.4
4.0
4.3
Alkalinity
Dissolved
mM
-
-
PH
7.4
-
10.2
Feed to precipitation unit: Primary effluent; Number of compartments: 3; Nominal residence
time: 46.5 min; Total lime added: 2.8 mM; Dissolved lime added: 0.1 mM.
59
-------
TABLE 12
OPERATING CONDITIONS AND RESULTS FOR WASTEWATER
EXPERIMENT 5
Sample Location
Primary Effluent
Activated Sludge
Effluent
Lime Precipitation
Effluent
Recarbonation
Effluent
Clinoptilolite
Effluent
P
Total
mM
0.40
.
0.026
.
Dissolved
mM
0.33
0.019
Ca
Dissolved
nW
1.40
1.71
-
Mg
Dissolved
mM
0.42
0.36
-
CT
Dissolved
mM
3.60
3.00
Alkalinity
Dissolved
mM
2.46
3.06
-
PH
7.5
9.6
-
Feed to precipitation unit: Primary effluent; Number of compartments: 4; Nominal residence
time: 42 min; Total lime added: 2.6 mM; Dissolved lime added: 0.2 mM.
100
Run 1
i+*
LjJ C
ee.
-------
than no phosphate removal was effected by the clinoptilelite columns -
since removal by these columns can only be expected by filtration and
filtration is only effective in removing particles.
In experiment 2 activated sludge effluent was treated with lime in a 3-
compartment mixer-flocculator with average nominal residence time of 46.5
min for an experimental period of 5 days. Lime doses of 1.6 mM (160 mg
as CaC03/)i) raised the pH of the activated sludge effluent from 7.3 to
9.8 and reduced the influent total phosphate from 10.8 mg P/£ to a value
of 1.7 mg P/£ following treatment by the whole process stream including
recarbonation and clinoptilolite columns (Table 9). Dissolution of parti-
culate phosphate occurred during recarbonation (causing an increase in
dissolved phosphate from 0.6 mg P/n to 1.7 mg P/£) as the pH was reduced
from 9.8 to 8.2.
In experiment 3 primary effluent was treated in a 3-compartment mixer-
flocculator with average nominal residence time of 46.5 min for an experi-
mental period of 10 days. Lime doses of 4.9 mM (490 mg as CaC03/£)
raised the pH of the primary effluent to 11.0 and reduced total phosphate
from a value of 11.5 mg P/2. to a residual of 0.23 mg P/2. following the
process stream of recarbonation and clinoptilolite treatment (Table 10).
Recarbonation resulted in an increase of dissolved phosphate from 0.1 mg
P/£ to about 0.2 mg P/fc as the pH was adjusted downward from 11.0 to 7.7.
In run 4 the experimental conditions were similar to those used in experi-
ment 3 with the exception that, over a five-day period, primary effluent
was dosed with 2.8 mM lime (280 mg as CaC03/£) to reach a precipitation
pH of 10.2. Final process total phosphate levels averaged 0.65 mg P/a —
levels that were increased by recarbonation from the 0.56 mg P/e. achieved
following sedimentation of the precipitated effluent.
In the final wastewater precipitation experiment (experiment 5) primary
effluent was treated in a 4-compartment mixer-flocculator with an average
residence time of 42 min for an experimental period of 4 days. At the
precipitation pH of 9.6, achieved by a lime dose of 2.6 mM (260 mg as
CaC03/Jl), the precipitation-sedimentation unit reduced the total phosphate
from 12.4 mg P/£ to 0.81 mg P/i.
COMPARISON OF PREDICTED AND EXPERIMENTAL
PHOSPHATE RESIDUALS
Good prediction of effluent dissolved phosphate concentration using
Equation (10) derived from the precipitation model developed from chemically-
defined solutions was obtained when based on the concentrations of dis-
solved constituents and with corrections being made for complex!ng of
calcium and for calcium carbonate precipitation. This procedure was made
necessary largely by the extremely poor efficiency of lime dissolution
obtained in the pilot plant. Thus Figure 21 indicates that of the lime
added to the lime dissolution unit some amount on the order of 10$ or
61
-------
LU
Experiment
1
FIGURE 21. COMPARISON OF LIME SLURRY ADDITION TO
DISSOLVED LIME CONCENTRATION IN
WASTEWATER EXPERIMENTS
less became dissolved in the wastewater in the first compartment of the
precipitation reactor. Additional complications arose because the lime
continued to dissolve in the subsequent reactor compartments. Thus
Figure 22 shows the increase in dissolved calcium caused by lime dis-
solution as one progresses down the precipitation reactors. ' !
When the dissolved calcium concentration was used the model produced
predictions of dissolved phosphate residuals that were within 25% of
experimental values (Table 13). Previous attempts to predict residual
62
-------
(O
CJ
o
to
CO
Point of
-Lime Addition
1 2
COMPARTMENT NO.
FIGURE 22. CHANGE IN DISSOLVED Ca AND PT
CONCENTRATIONS IN REACTOR COMPART-
MENTS -WASTEWATER EXPERIMENTS
phosphate concentrations in wastewater as a result of chemical dosing have
been in error by a factor of some 5 times [3].
TABLE 13
COMPARISON OF PREDICTED AND AVERAGE EXPERIMENTAL
DISSOLVED PHOSPHATE RESIDUALS FOR
WASTEWATER EXPERIMENTS
Experiment
No.
1
2
3
4
5
Dissolved Phosphate
Residual, mM x 10^
Predicted
47.5
208
36.2
92.2
230
Experimental
38.9
184
32.3
74.0
190
Error9
%
+22
+13
+17
+25
+21
a0/ F _ Predicted - Experimental 1QO
/0 hrror Experimental x IUU'
63
-------
Figure 23 shows the variation of the activity product of TCP with pH in
comparison with the values of this constant from chemically-defined sys-
tems. In computing the activity product of TCP for wastewaters the
complexing of calcium was taken into account (Appendix B). The line for
the wastewater points was fitted by least squares analysis. The agreement
between the values of this constant from the two types of system is good
at pH values below 10.5. At higher pH values the pA value of TCP in the
wastewater systems (where lime was used as a precipitant) is significantly
lower than for chemically-defined systems where dissolved calcium salts
were used as precipitants.
24
to
o
22
I
0
Chemically-Defined
_^.Systems ©
Experiment
0
B
A
^
0
1
2
3
4
5
I
9.0
9.5
10.0 10.5
PH
11.0
11.5
FIGURE 23. DEPENDENCE OF pA Ca^PO^ ON pH
FOR WASTEWATER EXPERIMENTS
OVERALL PHOSPHATE REMOVAL PERFORMANCE
In the wastewater tested(which is one with an average alkalinity of 2.4
mM and a typical Ca/Mg mole ratio of approximately 3 to 1), a phosphate
removal in excess of 80% was consistently achieved at pH 9.5 with lime
doses of, at the most, 2 mM (200 mg/2 as CaC03). Indeed, for most of
64
-------
the experimental period phosphate removals exceeded 90%. It should be
noted, however, that the overall phosphate removal performance (total
rather than the removal based on dissolved phosphate concentration) of the
entire treatment train was dictated by the performance of the precipi-
tation reactor and its ensuing sedimentation basin. It was demonstrated
that any phosphate-containing particles that escaped sedimentation were
rapidly dissolved during recarbonation (Tables 8 through 12). This
observation raises the important issue of the efficacy of filtration for
improving the phosphate removal for wastewater treated by lime precipi-
tation. Effluents from all such processes will require downward pH
adjustment by some process such as recarbonation prior to filtration.
As these experiments have demonstrated, however, even short-term (5 min)
recarbonation or other low pH environments will cause the rapid and
complete dissolution of the particulate phosphate, thus making it non-
removable by filtration. It must be concluded, therefore, that filtration
cannot be used to reduce the phosphate residual in lime-treated wastewater
and that the effective phosphate residual of a lime precipitation phos-
phate removal process is determined by the efficiency of the sedimentation
process following precipitation.
Because of this observation, which emphasized the importance of precipitate
removal by sedimentation, some preliminary experiments were conducted on
chemically-defined systems to investigate the effect of solution composi-
tion on the behavior of calcium phosphate—carbonate suspensions during
sedimentation.
Factors Influencing Separation of Precipitates
Previous literature has attributed the inability to operate calcium
phosphate precipitation processes efficiently below pH values in the
10.5 to 11 region to the production of poorly flocculated precipitates
that cannot be readily separated. Operation of these processes at pH
values of 10.5 or greater was thought to be beneficial from a precipitate
separation standpoint because of the formation of CaC03 (which supposedly
increases suspension density) and of Mg(OH)2 (which reportedly is gela-
tinous and acts as a suspension binder).
Two series of CSTR experiments were conducted on chemically-defined solu-
tions with the objective of determining the effect of precipitating
conditions on the settling properties of the precipitates that were formed.
Variables investigated included pH; initial concentrations of calcium,
magnesium, and carbonate; and the initial Ca/Mg ratios. Both series of
experiments employed single CSTR's at 23°C, with 15 min nominal residence
time and with mean velocity gradient values of 76 sec"1. The settling
properties of the suspension were assessed by quiescent sedimentation
for 30 min following the CSTR. Initial component concentrations for the
two series of experiments are presented in Tables 14 and 15.
Of the variables studied, initial magnesium concentration and pH exerted
the most influence on the separation by sedimentation of calcium
65
-------
TABLE 14
INITIAL COMPONENT CONCENTRATION FOR PRECIPITATE
SEPARATION EXPERIMENTS, SERIES I
Run No.
1
2
3
4
Ca
mM
14.2
2.2
2.2
2.2
Mg
mM
3.7
0.6
3.6
10.8
PT
mM
0.4
0.4
0.4
0.4
CT
mM
2.4
2.4
2.6
2.4
pH
(Range of
Variation)
8 -11
8 -11
8.5-10.9
8.5-11
Ca/Mg
Mole
Rati o
3.8
3.7
0.6
0.2
Ca/CT
Mole
Ratio
5.9
0.9
0.8
0.9
Ca/PT
Mole
Rati o
4.0
6.1
6.1
6.0
TABLE 15
INITIAL COMPONENT CONCENTRATION FOR PRECIPITATE
SEPARATION EXPERIMENTS, SERIES II
Run No.
1
2
3
4
Ca
mM
2.2
2.2
2.2
2.2
Mg
mM
0.6
0.6
0.6
0.6
mM
0.4
0.4
0.4
0.4
CT
mM
(Range of
Variation)
2.7-12.7
2.7-12.7
2.7-12.7
2.7-12.7
PH
9.0
9.5
10.0
11.0
Ca/Mg
Mole
Rati o
3.7
3.7
3.7
3.7
Ca/CT
Mole Ratio
(Range of
Variation)
0.2-0.8
0.2-0.8
0.2-0.8
0.2-0.8
Ca/PT
Mole
Ratio
6.1
6.1
6.1
6.1
phosphate-carbonate suspensions. Examination of Figures 24 and 25 shows
that increasing the pH of suspensions in general improved their settling
characteristics. Indeed, if one were to select the very modest criterion
that 90% removal of precipitate by sedimentation was the minimum acceptable
performance for a precipitate separation process, then only operation at
a pH value of greater than 10.5 would consistently achieve this objective.
In the pH range between 8 and 10 the suspensions generally settled poorly.
The influence of magnesium concentration on suspension settling is
66
-------
Run
No.
1
2
3
4
Initial Concentration
mM
PT
0.4
0.4
0.4
0.4
CT
2.4
2.4
2.6
2.4
Ca
14.2
2.2
2.2
2.2
Mg
3.7
0.6
3.6
10.8
CSTR (G = 76 sec'1,
T = 900 sec) followed by
30 m1n quiescent settling
FIGURE 24.
pH
SETTLING PROPERTIES OF CALCIUM
PHOSPHATE-CARBONATE SUSPENSIONS
(SERIES I)
67
-------
uj z —i 0.4
i— •-'i-
-------
concentrations of a positive specie such as Ca2+ and uninfluenced by the
presence of negatively charged species such as HCC>3 or CO*'.
These preliminary observations on the separation characteristics of
calcium phosphate-carbonate suspensions point to the importance of
suspension surface properties in the removal of phosphate by precipi-
tation. They also indicate that a fruitful area for investigation exists
in the physical characterization of these suspensions to provide infor-
mation that will lead to effective methods of precipitate separation at
pH values below 10.5. At this stage one might conclude that high (>90%)
removals of phosphate by calcium phosphate precipitation and sedimenta-
tion of solids are only possible at pH values of above 10. To achieve
80% overall phosphate removal at pH values of 10 or lower, the precipitate
must be coagulated with a cationic material, and in these experiments it
appeared that a large excess calcium dose achieved this objective. It
would be better to achieve the same results with lower concentrations of
alternative flocculants such as polyelectrolytes or alum or ferric salts
which may have both an electrostatic and bridging function.
69
-------
SECTION X
ACKNOWLEDGMENTS
This investigation was supported in part by Grant No. 17080 DAR from the
Environmental Protection Agency. The work was performed by Mr. Arnold
B. Menar'un'der the direction of Professor David Jenkins. Mr. Warren
Schwart? served as Project Officer for the Environmental Protection
Agency. The assistance of Miss Christine Tarr and the late Mr. Philip
Palmerlperforming analyses and conducting experimental work throughout
this study is appreciated. The generosity of Mr. Robert Jack for pro-
viding X-ray crystallographic facilities is acknowledged. Portions of
the experimental phase of this study were conducted with other work
performed concurrently on this project. The assistance of Drs. John H.
Koon and Larry A. Esvelt in this regard is greatly appreciated.
71
-------
SECTION XI
REFERENCES
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Pollution Control Fed., 41, p. 708 (1969).
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Treatment Processes. Part II -Mechanism of Enhanced Phosphate
Removal by Activated Sludge, SERL Report 68-6, Berkeley Sanit.
Eng. Research Lab., University of California (1968).
3. Ferguson, J. F., "The Precipitation of Calcium Phosphates from Fresh
Waters and Wastewaters," Ph.D. Dissertation, Stanford University,
Stanford, California (1970).
4. Ferguson, J. F., Jenkins, D., and Stumm, W., "Calcium Phosphate
Precipitation in Wastewater Treatment," Water - 1970, American
Institute of Chemical Engineers.
5. Jenkins, D., Ferguson, J. F., and Menar, A. B., "Chemical Processes
for Phosphate Removal," Wastewater Reclamation and Reuse Workshop,
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7. Wuhrmann, K., "Objective, Technology and Results of Nitrogen and
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8. Mulbarger, M. C., Grossman III, E., Dean, R. B., Grant, 0. L., "Lime
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10. Seiden, L., and Patel, K., Mathematical Model of Tertiary Treatment by
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American Research Division (1969).
11. Owen, R., "Removal of Phosphate from Sewage Plant Effluents with Lime,"
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73
-------
12. Spiegel, M., and Forrest, T. H., "Phosphate Removal: Summary of
Papers," J. Sanit. Eng. Div., Proc. ASCE. 95 (SA5), pp. 803-827
(1969).
13. Suhr, L. G., Nutrient Removal at Lake Tahoe - Costs and Criteria,
report by Cornell, Howland, Hayes, & Merryfield, Engineers and
Planners, Corvallis, Oregon (n/d).
14. Culp, R. L., Advanced Wastewater Treatment as Practiced at South
Tahoe, Final Report to EPA on R & D Project 17010 ELQ (WRPD 52-01-
67TT1971).
15. Garland, C. F., "High Density, Solids Contact," paper presented at
FWQA sponsored seminar, Chicago, 111., June 26-27 (1968).
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17. Humenick, M. J., and Kaufman, W. J., "An Integrated Biological-
Chemical Process for Municipal Wastewater Treatment," paper presented
at the 5th International Water Pollution Research Conference, San
Francisco (1970).
18. Stumm, W., and Leckie, J. 0., "Phosphate Exchange with Sediments;
Its Role in the Productivity of Surface Waters," paper presented at
the 5th International Water Pollution Research Conference, San
Francisco (1970).
19. Schmidt, L. A., and McKinney, R. E., "Phosphate Removal by a Lime-
Biological Treatment Scheme," J. Water Pollution Control Fed., 41,
p. 1259 (1969).
20. Ferguson, J. F., Jenkins, D., Eastman, J. D., "Calcium Phosphate
Precipitation at Slightly Alkaline pH Values," (in press) J. Water
Pollution Control Fed.
21. Vacker, D., Connell, C. H., and Wells, W. N., "Phosphate Removal
Through Municipal Wastewater Treatment at San Antonio, Texas,"
J. Water Pollution Control Fed., 39, p. 750 (1967).
22. Ferguson, J. F., and McCarty, P. L., The Precipitation of Phosphates
From Fresh and Waste Waters, Technical Report No. 120, Dept. of
Civil Engineering, Stanford University, Stanford, Calif. (1969).
23. Strates, B. S., Neuman, W. F., and Levinskas, G. J., "The Solubility
of Bone Mineral. II. Precipitation of Near Neutral Solutions of
Calcium and Phosphate," J. Physical Chem., 61. p. 279 (1957). ,
24. MacGregor, J., and Brown, W. E., "Blood:Bone Equilibrium in Calcium
Homeostasis," Nature, 205, p. 359 (1965).
74
-------
25. Booth, D. H., and Coates, R. V., "The Stability of Calcium Hydrogen
Phosphate Precipitated from Solutions of Calcium Nitrate and Phos-
phoric Acid," J. Chem. Soc.. p. 4914 (1961).
26. Hal ton, A. G., Bodin, W. J., Furedi, H., and Schwartz, A., "Nucleation
of Calcium Phosphate from Solution," Canadian J. Chem, 45, p. 2695
27. Van Wazer, J. R., Phosphorus and Its Compounds. Volume I, Inter-
science (1961).
28. Hendricks, S. B., and Hill, W. L., "The Nature of Bone and Phosphate
Rock," Proceedings National Academy of Science of US, 36, p. 731
(1950).
29. Rootare, H. M., Deitz, V. P., and Carpenter, F. G., "Solubility
Product Phenomena in Hydroxyapatite-Water Systems," J. Colloid
Sci.. 17, p. 179 (1962).
30. McConnell, D., "Crystal Chemistry of Hydroxyapatite — Its Relation
to Bone Mineral," Arch. Oral Biol., 10, p. 421 (1965).
31. Berry, E. E., "The Structure and Composition of Some Calcium-Deficient
Apatites," J. Inorg. Nucl. Chem., 29, p. 317 (1967).
32. Eanes, E. D., Gillessen, I. H., and Posner, A. S., "Mechanism of
Conversion of Non-Crystalline Calcium Phosphate to Hydroxyapatite,"
Proceedings International Conference on Crystal Growth, New York
(1967).
33. Stumm, W., and Morgan, J. J., Aquatic Chemistry, Wiley-Interscience
(1970).
34. Newesely, H., "Changes in Crystal Types of Low Solubility Calcium
Phosphates in the Presence of Accompanying Ions," Arch. Oral Biol.,
Special Supplement, 6. p. 174 (1961).
35. Mooney, R. W., and Aia, M. A., "Alkaline Earth Phosphates," Chem.
Rev., p. 433 (1961).
36. Chave, K. E., Deffeyes, K. S., Weyl, P. K., Garrels, R. M., and
Thompson, M. E., "Observations on the Solubility of Skeletal Carbon-
ates in Aqueous Solutions," Science, 137, p. 33 (1962).
37. Simpson, D. R., "Apatite and Octa-Calcium Phosphate: Effects of
Carbon Dioxide and Halogens on Formation," Science. 154, p. 1660
(1966).
38. Somasundaran, P., "Zeta Potential of Apatite in Aqueous Solutions
and Its Change During Equilibration," J. Colloid and Interface Sci..
27, p. 659 (1968).
75
-------
39. Parks, G. A., "The Isoelectric Points of Solid Oxides, Solid Hydroxides,
and Aqueous Hydroxo Complex Systems," Chem. Rev., 65, p. 177 (1965).
40. LaMer, V. K., and Smellie, R. H., "Flocculation, Subsidence and
Filtration of Phosphate Slimes," J. Colloid Sci., 11, p. 704 (1956).
41. American Public Health Association, Standard Methods for the Examina-
tion of Water and Wastewater, llth Edition (1960) and 10th Edition
(1955).
42. FWPCA Methods for Chemical Analysis of Water and Wastes. Federal
Water Quality Administration, Cincinnati, Ohio (1969).
43. Jenkins, D., Analytical Methods. Berkeley: Sanit. Eng. Research
Lab., University of California (1966).
44. Winneberger, J. H., Austin, J. H., and Klett, C. A., "Membrane Filter
Weight Determinations," J. Water Pollution Control Fed.. 35. p. 807,
(1963).
45. Basset Jr., H., "Beitrage zum Studium der Calciumphosphate," ~L._
Anorg. Chemie, 59, p. 1 (1908).
46. Bard, A. J., Chemical Equilibrium, Harper and Row, New York (1966).
47. Sillen, L. G., and Martell, A. E., Stability Constants of Metal-Ion
Complexes, London Chem. Soc., Special Publication No. 17 (1964).
48. Somasundaran, P., and Agar, G. E., "The Zero Point of Change of
Calcite," J. Colloid and Interface Sci.. 24. p. 433 (1967).
76
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SECTION XII
GLOSSARY
Symbol Definltion
[Ca2+]in Influent calcium ion concentration
[Ca2+]ss Steady state calcium ion concentration
[CaTJ Total dissolved calcium concentration
CSTR Continuously stirred tank reactor
CT Total carbon = H2C03 + HCOg + COg"
DCP Dicalcium phosphate
HAP Hydroxyapatite
hp Horsepower
kj First equilibrium constant of phosphoric acid
k2 Second equilibrium constant of phosphoric acid
k3 Third equilibrium constant of phosphoric acid
a 1i ter
M9(OH)2 Magnesium hydroxide
mg/a milligram per liter
nm nanometers
OCR Octacalcium phosphate
P. Influent phosphate concentration
[P0^~] Steady state P0^~ concentration
P Steady state phosphate concentration
PT Total orthophosphate = H3POl( + H2PO^ + HPOj" + POj"
PVC Polyvinyl chloride
77
-------
Symbol Definition
pA Negative logarithm of a solubility product
pH pH of zero point of charge
zp c*
sec'1 Reciprocal seconds
SERL Sanitary Engineering Research Laboratory
TCP Tricalcium phosphate
zpc Zero point of charge
y Mi cron
78 .
-------
SECTION XIII
APPENDICES
Page
A. Walton Technique for Determining Stoichiometry'
of Nucleating Solid Phase 81
B. Example of Manual Computation of Residual
Phosphate Values 85
C. Daily Wastewater Experimental Data 89
79
-------
APPENDIX A
WALTON TECHNIQUE FOR DETERMINING
STOICHIOMETRY OF NUCLEATING
SOLID PHASE [26]
81
-------
Walton et^ al. [26] state that the energy barrier to nucleation, AQ°
(nucleatiorTJ" is given by the following expression:
AG° (nucleation) = AG° (cluster) - AG°sp (solubility) . (12)
The terms on the right-hand side of Equation (12) can be evaluated ,as
follows: The standard free energy of the cluster, AG° (cluster), depends
on th'e type and number of interacting ionic species. In general, if
ionic specie Aa+ interacts with ionic specie B^" to form a neutral ionic
aggregate, the metastable equilibrium is given by the equation:
xbAa+ + xaBb" + (Aba+Bab~) . (13)
/\
The condition for equilibrium is AGX = 0, or
o
AGXQ = bxyA + axpB - xy AB - Za-jj =0 (14)
where:
yA» ^B' an<^ yAB are ^e cnemi'ca^ potentials of species Aa , B ~, and
(Ag+ Bb-)x, respectively,
oij is the interfacial energy of the interface between the i and j
phases.
After equating U/\B = ^°AB (solid). Walton defines the total standard
energy of the critical cluster AG x as:
r AG°Xo (cluster) = -RT ln[A]b[B]a + z ^ . (15) .
Also
AG°sp (solubility) = -RT In K^ (16)
where KS is the solubility product.
Now, substituting Equations (15) and (16) in Equation (12) we obtain
Equation (17):
82
-------
AG° (nucleation) = -RT ln[A]b[B]a + z -& - RT In K . (17)
x s p
Rearranging Equation (17)
ln[A]b[B]a = z ii + In K - AG° (nucjeation) _ (1Q)
XKI S p KI
The right-hand side of Equation (18) is constant at constant temperature
for any one critical size nucleus of constant composition and interfacial
energy. Thus:
ln[A] = - | ln[B] + constant . (19)
Thus, if the above criteria hold, a plot of the logarithm of the critical
concentration of the cation A against that of the anion B should yield a
straight line. The slope of this line gives the stoichiometric ratio of
the initial phase.
The above derivation does not depend upon the detailed nature of the
nucleation process and similar reasoning can be applied to systems in
which the nucleus contains more than two ionic species.
Walton and coworkers [26] visualize the formation of a critical nucleus
and the beginning of calcium phosphate precipitation as follows:
ax HgPO^ + bx Ca(OH)2 + x (c - 2b) H20
[CabH3a_2b (POJ3C H20jx (critical nucleus)
precipitation (20)
At the onset of nucleation, the relation between the concentrations of
the species is given by Equation (20) i.e.,
log ([Ca2+][OH"]2)b ([H+][H2PO;])a = constant . (21)
Equation (21) could be rearranged
b log ([Ca2+][OH~]2) + a log ([H+][H2PO;;] \ = constant (22)
log ([Ca2+][OH~]2) = - (£) log ([H+][H2PO;])+ constant (23)
83
-------
Thus if log ([Ca2+][OH"]2)is plotted against log ([H+][H2POi;] )for a
series of experiments, a straight line should result with a slope of
a/b that gives the ratio of phosphate to calcium in the nucleating or
steady state phase. Data from Walton et_ aj_. [26], together with data
from the chemically-defined systems of this work, are presented in
Figure 26. The data from the chemically-defined systems appear in the
text of this report (Figure 9) on a somewhat more expanded scale.
18
16
'= 14
o
(O
O
12
10
Data from Walton et
Ca/P =1.49
Ca/P =1.44
Run
0
A
°C
37
1.9-23
Time
min
1,440
15
Reactor
BATCH
CSTR
Initial Conditions
M x 103
Ca
.04-2.0
2.1-2.4
Mg
0.0
0.0-0.6
PT
1.0
0.3-0.5
°r
0.0
0.0-14.8
n
FIGURE 26.
13
15
17
19
21
CRITICAL IONIC CONCENTRATIONS AS
DETERMINED BY WALTON et_aL AND
BY THIS INVESTIGATION
84
-------
APPENDIX B
EXAMPLE OF MANUAL COMPUTATION OF
RESIDUAL PHOSPHATE VALUES
85
-------
In this presentation the manual computation of residual dissolved phos-
phate concentrations will be illustrated by the following example:
Initial Conditions: Ca, 2.19 mM; Pj, 0.38 mM; Mg, 0.59 mM; CT? 5.95 mM.
Conditions of Precipitation: pH 10.
Predictive Equation:
3 log{[Ca2+]in - 1.5 [P1f) - PS$]} + 2 log[POj~] = - 23.56 (10)
Equation (10) is solved for Pss by trial and error. Thus, assume
[Pin - P$SJ = 0.9 Piri, i.e., PSS = 3.8 x 10~5M. Now, compute
[Ca2+]ss = (2.19 - 0.9) x 1.5 x 0.38 x 10"3
1.68 x 10"3M .
Substituting this value for [Ca2"1"]-,.. in Equation (10) and solving for
[P0j~]ss 2.34 x 10"8M .
The fraction that [P0ij~]ss is of Pss at any pH can be computed from
equilibrium constants of phosphoric acid and [H+].
Thus,
PSS = 4.9 x 10"6M
This value of PSg is not within 5% of the assumed PSs value. Therefore,
the computation is repeated using,
P._ = 4.9 x 10"6M
j o
This yields
= 5.26 x 10"6M.
The difference between this value and the assumed value is again greater
than 5% so that a third iteration is necessary using,
86
-------
PSS = 5.26 x 10~6M .
The computation now yields
PSS = 5.26 x 10"6M
an identical value to that assumed.
The computation of residual dissolved phosphate from Equation (10) can
be modified when necessary to account for the calcium that is required
to precipitate calcium carbonate. The solubility product chosen for
calcium carbonate was
[Ca2+][COf ] = 2.6 x 10'6 (24)
a value typical of a destabilized calcium carbonate formed under the
precipitation conditions encountered in wastewater precipitation
processes [10,22].
Now, setting the [Ca2+] incorporated into calcium carbonate precipitate
equal to x, we have
[Ca2+ - x][CO§" - x] = 2.6 x 10"6 • (25)
This quadratic equation can be solved and the amount of calcium removed
as calcium carbonate precipitate determined.
Additional corrections for the complexation of calcium become significant
at pH values of greater than pH 10. The various complexes of calcium
can be regarded as removing the calcium from the sphere of activity in
precipitating phosphate. The complexes considered were: CaOH+,
CaC03(aq), CaHC03, and CaPOi;. Thus,
[CaTJ = [Ca2+l + [CaOH+] + [CaC03(aq)] + [CaMCO^] + [CaPO^] (26)
where [Cay] is the total dissolved calcium concentration. The equili-
brium constants for these complexes are as follows:
Ca2+ + OH" = CaOH+ KCaOH+ = 23-5
Ca2+ + CO2' = CaC03(aq) KCaC03(aq) = 1>59 *
Ca2+ + HC03 = CaHC03 KCaHCO" = 18'2
Ca2+ + POr = CaPOl Kr,Dn- = 2.89 x 104
87
-------
From the mass balance equation for total dissolved calcium, Equation
(26), and the equilibrium constants for calcium complexes the following
expression for [Ca2+] is obtained:
[Ca2+] (27)
[Ca?]
1 + KCaOH+ [0irl
+ KCaC03(aq)
[C0|-] H
h KCaHCO| [F
1C°3] + KCaPOi; Cp°'"l
The most significant terms in this equation are those for the CaC03(aq)
and CaPOi; complexes. The CaOH+ and CaHCOlj complexes can be neglected
at pH values below 11. Thus a simplified form of Equation (27) can be
used
[CaT]
[Ca2+] = - ! - ' (28)
1 + KCaC03(aq) "
For the ranges of concentrations and pH conditions in the experiments
conducted here, neglecting the CaPO£ complex, would introduce an error
of 52-102 in values of [Ca2 ]. Thus Equation (28) could be further
simplified to
[CaT]
[Ca2+] = - 1 - — . (29)
-
1 + KCaC03(aq)
Computation of the initial [Ca2+] available for precipitating calcium
phosphate, therefore, consists of calculating the amount of calcium
removed by calcium carbonate precipitation and then converting the
residual dissolved calcium value to [Ca2+] by correcting for calcium
complexes by one of the Equations (27, 28, or 29).
88
-------
APPENDIX C
DAILY WASTEWATER EXPERIMENTAL DATA
89
-------
8 -
7*
S A V\ S\—
> ^ -cj tr
_ Run 2
>-^ — -O— —
trr^fc— — gt^r5^
*- ^ If -V
$—==£ §_
A ' ' i '
-ti
T7
0 i i i , ,i ,
TO
10
I
9
8
7
9
8
Run 3
Run 4
_L
I
'i
T I T
. .Q. A
Run 5
I I I
'ilot Plant Streams
o Primary Effluent
a Activated Sludge Effluent
A Lime Precipitation Effluent
S7 C02 Recarbonation Effluent
0 ClrinoDtilelite Effluent.i ,
5 6
TIME, days
8
9 /'
10
FIGURE 27. EFFECT OF UNIT PROCESSES ON pH
90
-------
0.5
0.4 -
0.0
„ "••» s
t! C
I 0.3
Q-
1 0.2
P "''l
1
0.0
] "^ -n ^* -^^.j Pilot Plant Streams
"- o Primary Effluent
Q Activated Sludge Effluent
A Lime Precipitation Effluen.
Run 2 v C02 Recarbonation Effluent
0 Clinoptilolite Effluent
* i i i i i i i
10
TIME, days
FIGURE 28. EFFECT OF UNIT PROCESSES ON TOTAL
PHOSPHATE RESIDUALS (RUNS 1, 2 AND 3)
91
-------
0.4
A
LU
O
Q.
_l
1
0
0
u
o
(
.3
.2
.1
0*
Primary E
Run
^fluent -
4
"Lime Precipitated,
bonation, and
-Effluents
!h w» ^
y ^i ^
•° 1 2
TIME,
C02 Recar
Clinoptilol
?»
Jl
3
days
g
*P
4
ite
—
r
"5
1
LU
1—
D-
o
o.
^
o
5
0
0
0
0
<
.3
.2
.1
.0<
Primary
Run
Effluent -
5
Lime Precipitation"
Effluent
-
__ — • a—
r* i
1 2
TIME,
-
& y
i *\
3 '
days
FIGURE 29. EFFECT OF UNIT PROCESSES ON TOTAL
PHOSPHATE RESIDUALS (RUNS 4 AND 5)
o
40
co
Cllnopt. Eff.-
j r i
2468
TIME, days
Activated Sludge
Effluent
2468
TIME, days
2468
TIME, days
Lime Precipitated C02 Recarbonated and
Effluent Clinoptilolite Effluents
Activated Sludge Effluent
0.005- C02 Rec.
Clinoptilolite
Eff.
Run 2
Lime Ppt. Eff.
i i
234
TIME, days
234
TIME, days
FIGURE 30. EFFECT OF UNIT PROCESSES ON DISSOLVED
PHOSPHATE RESIDUALS (RUNS 1 AND 2)
92
-------
8
10
0.012
0.010
0.008^
0.006
1 0.004
*
a 0.002
Run 3
C02 Recarbonated Effluent
Lime Precipitation
"Effluent
Clinoptilelite
Effluent
o
GO
I-H
o
1
0.3
0.2
0.1
0.0
0.4
Run 4
Primary Effluent
1
0.2
Run 5
Primary Effluent
i i
1 2 3
TIME, days
0.002
0.001
Run 5
Lime Precipitated
Effluent
1
2 3
TIME, days
FIGURE 31. EFFECT OF UNIT PROCESSES ON DISSOLVED
PHOSPHATE RESIDUALS (RUNS 3, 4 AND 5)
93
-------
o
o
CO
I—I
o
Pilot Plant Streams
o Primary Effluent
n Activated Sludge Effluent -
A Lime Precipitation Effluent
v C02 Recarbonation Effluent-
0 Clinoptilolite Effluent
456
TIME, days
8
10
FIGURE 32.
EFFECT OF UNIT PROCESSES ON DISSOLVED
CALCIUM CONCENTRATION (RUNS 1, 2 AND 3)
94
-------
1.8
Pilot Plant Streams
• o Primary EFfluent
& Lime Precipitation Eff
^7 C02 Recarbonation Eff.
234
TIME, days
Run 5
2 3
TIME, days
FIGURE 33. EFFECT OF UNIT PROCESSES ON DISSOLVED
CALCIUM CONCENTRATION (RUNS 4 AND 5)
95
-------
Run 1
Run 2
2.0
_C02 Recarb. Eff.
12345678
12345
3.0
i r
Prim. Eff.
Clinopt. Eff.
C02 Rec. Eff.
Clinipt. Eff.
C02 Recarb. Eff.
Lime Ppt. Eff.
2.0
5 6 7
TIME, days
o
Run 4
Run 5
3.1
^ I T
Lime Ppt. Eff.
Prim. Eff.
I I
I
2.5 -
12345
TIME, days TIME, days I
FIGURE 34. EFFECT OF UNIT PROCESSES ON ALKALINITY
96
-------
SELECTED WATER
RESOURCES ABSTRACTS
INPUT TRANSACTION FORM
;. Report No.
w
' Calcium Phosphate Precipitation in
Waste-water Treatment
Menar; Arnold B. and Jenkins, David
California University
Sanitary Engineering Research Laboratory
.5., KsportDate
-5 June 1972
8. Performing Organization
72-6
.. ...
EPA WQO 1 7080 DAR
Is Type , f Repotl aad
Period Covered
12. 'Srinsorir.r Organisation
Environmental Protection Agency report
number. EPA-R2-72-064, December 1972.
.,.-• , • \_ i
This investigation examined the precipitation of calcium phosphate both
from chemically-defined solutions representative of wastewater composition and from
wastewater. The steady state solid phase that controlled dissolved phosphate residual
was an amorphous tricalcium phosphate. The solubility of this phase, determined
from chemically-defined systems, was used with success to predict dissolved
phosphate residuals from both chemically-defined systems and wastewaters. Sus-
pension recycle was found to result in lower dissolved phosphate residuals, but poor
suspension settling below pH 10 made this process difficult to maintain. Suspension
settling was enhanced by Mg(OH)2 precipitation but not by CaCOS precipitation. In
wastewater of moderate alkalinity and hardness, a phosphate removal in excess of
80% was consistently achieved at pH 9. 5 with lime doses of, at the most, 200 mg/
liter as CaCO3. The overall phosphate removal performance was dictated by the
performance of the precipitation reactor and its ensuing sedimentation basin.
Phosphate-containing particles that escaped sedimentation could not be removed by
filtration because they dissolved rapidly during the recarbonation process that
necessarily precedes the filtration step.
173. Descriptors
* Waste Treatment — Sewage Treatment
I7b. Identifiers
Calcium phosphate, Tricalcium phosphate, Calcium carbonate,
Calcite, Magnesium hydroxide
;;.. fOlVRR Fit-Ill
,! 05D
------- |