The Effects of Urine Storage Conditions on Struvite Recovery

THE EFFECTS OF URINE STORAGE CONDITIONS ON STRUVITE RECOVERY

ELIZABETH ANNE TILLEY

B.A.Sc, (Environmental Engineering) University of Waterloo, 2003

. A THESIS SUBMITTED IN PARTIAL FULFILLMENT OF THE REQUIREMENTS FOR THE DEGREEE OF

MASTER OF APPLIED SCIENCE

THE FACULTY OF GRADUATE STUDIES

(CIVIL ENGINEERING)

THE UNIVERSITY OF BRITISH COLUMBIA

September 2006

Phosphorus, like oil, is a non-renewable resource that must be harvested from finite resources in the earth's crust. An essential element for life, phosphorus is becoming increasingly scarce, contaminated, and difficult to extract. Struvite, or magnesium ammonium phosphate (MgNFLPC^ 6H2O) is a white, crystalline phosphate mineral that can be used as a bioavailable fertilizer and can be recovered from aqueous solutions such as digestor supernant.

In response to diminishing water resources, increasing nutrient pollution, and largely unaffordable centralized treatment, a paradigm of Ecological Sanitation (EcoSan) has emerged. A central tenant of EcoSan technology is nutrient recovery; by separating urine from feces in the absence of water, urine can be used as a clean, concentrated nutrient source. Urine harvested in this manner is used as a liquid fertilizer with varying degrees of success and acceptance.

This research examines the potential of urine to be a feedstock for struvite recovery. By recovering a sustainable source of phosphorus from urine, the prospect of appropriate sanitation and closed-loop nutrient systems, may move closer to reality.

In laboratory experiments using synthetic and real human urine, different methods of preparing urine to be used as a feedstock for struvite recovery, were examined. The effect of temperature, faecal contamination, dilution, and headspace on stored nutrient levels was examined. The effect of adding different quantities of magnesium, at different times, on the amount of phosphorus that could be removed from solution, was also examined.

An average of 70% of phosphorus could be removed from real urine in the form of struvite when magnesium was added to the urine solution after ureolysis had forced the precipitation of calcium and magnesium minerals; magnesium added before ureolysis began retarded the process. Dilution and the presence of wastewater were found to affect the rate of ureolysis but not the purity of the struvite recovered; recovered struvite was approximately 99% pure regardless of dilution or contamination.

Based on a comparison of the results, synthetic urine was found to be representative of the general behaviour of real urine during struvite formation. Table of Contents

Abstract • • — 11 Table of Contents »v

List of Tables »vii

List of Figures viii

Acknowledgements 'x Dedication • x Co-Authorship Statement...... x'

CHAPTER 1: Background • 1 1.1 Literature Review 1.1.1 Water and Sanitation 3 1.1.2 Sanitation Systems 4 1.1.2.1 Current Western model 4 1.1.2.2 Ecosan -4 1.1.3 Source Separated Urine 6 1.1.3.1 Composition and Behaviour 6 1.1.3.2 Contamination 8 1.1.3.3 Experiences and Attitudes 10 1.1.3.4 Life Cycle Analysis U 1.1.3.5 Options for Nutrient Recovery 11 1.1.4 Summary 14 1.2 Phosphorus 14 1.2.1 Phosphorus as a Resource 14 1.2.2 Phosphorus Chemistry • 16 1.2.2.1 Solubility Constant 16 1.2.2.2 Conditional Solubility Product 18 1.2.2.3 Supersaturation Ratio 19 1.2.3 Struvite : 19 1.2.3.1 Struvite Morphology 19 1.2.3.2 Struvite Crystallization : 20 1.2.3.3 Struvite Formation 21 1.2.3.4 The Solubility Product of Struvite 21 1.2.3.5 Chemistry of Struvite Formation :..22 1.2.4 Calcium Phosphates 26 1.2.4.1 Mineral Conversion 27 1.2.4.2 Magnesium as an Impurity 28 1.2.4.3 Struvite and Calcium Phosphates ; : 28 1.2.5 Urine 29 1.2.5.1 Urine and Struvite 30 1.2.5.2 Urease.. . 32

iv 1.2.6 Summary 35 1.3 Theme and Objectives 35 1.4 References 37

CHAPTER 2: Preparation of urine for optimal struvite recovery 48 2.1 Introduction 48 2.2 Materials and Methods 50 2.2.1 Part A 51 2.2.2 PartB 52 2.2.3 PartC .....53 2.3 Results and Discussion 55 2.3.1 Part A 55 2.3.1.1 Temperature 55 2.3.1.2 Headspace 60 2.3.1.3 Dilution , 61 2.3.1.4 Faecal Contamination '. 63 2.3.1.5 Precipitate : 64 2.3.2 PartB 64 2.3.2.1 Effect of Magnesium on Ureolysis 64 2.3.3 PartC 69 2.3.3.1 pH 70 2.3.3.2 Phosphorus 71 2.4 Conclusions 76 2.5 References 78

CHAPTER 3: Recovery of struvite from stored urine 81 3.1 Introduction : 81 3.2 Materials and Methods. 83 3.2.1.1 Urine Collection 83 3.3 Experimental Design 84 3.3.1 Spontaneous Precipitation •. v. 86 3.3.2 Struvite Precipitation ; 86 3.3.3 Chemical Analysis .' 87 3.4 Results. : 87 3.4.1 Composition 87 3.5 Ammonia and pH • 88 3.6 Spontaneous Precipitation 94 3.7 Calcium Removal -. -. 94 3.8 Struvite Recovery 96 3.8.1 Phosphorus 96

v 3.8.2 Calcium 98 3.8.3 Timing of Magnesium Dosage 99 3.8.4 Purity of Crystals 101 3.8.5 Comparison 101 3.8.6 Ureolysis 102 3.8.6.1 Mineral Precipitation 103 3.9 Conclusions 105 3.10 References 108

CHAPTER 4: Discussion and Conclusions 110 4.1 Summary HO 4.2 Strengths and Weaknesses Ill 4.3 Recommendations for Future Work 113 4.4 References 116

Appendices 117 Appendix I 117 Appendix II 150 Appendix III 172

vi List of Tables

Table 1.1 Composition of synthetic urine 30 Table 2.1. Composition of synthetic urine 50 Table 2.2. Matrix of solutions in Part A 51 Table 2.3. Matrix of solutions in Part B 53 Table 2.4. Matrix of solutions in Part C 54 Table 2.5. Dissociation constants for ammonia 57 Table 2.5. Net loss of ammoniacal nitrogen after 30 days 60 Table 2.6. Percentage (%) loss of ammoniacal nitrogen 61 Table 2.7. Calcium content in struvite crystals 76 Table 3.1. Summary statistics of urine collection 83 Table 3.2. Typical and measured concentrations (mg/L) of selected urine constituents 87 Table 3.3. Contamination as a percentage of struvite mass metals 101 Table 3.4. Comparison of mineral formation in synthetic and real urine 103

vii List of Figures

Figure 2.1. pH change with storage at a) 4°C and b) 23°C 56 Figure 2.2. Changes in ammonia in solutions stored at a) 4°C and b) 23°C 58 Figure 2.3. Changes in pH vs. ammonia concentration 59 Figure 2.4. Effects of dilution on pH for solutions at room temperature 62 Figure 2.5. Effect of wastewater on pH 63 Figure 2.6. Effect of magnesium on pH in a) full strength and b) diluted urine solutions 65 Figure 2.7. Ammonia generation in a) full strength and b) dilute urine solutions 67 Figure 2.8. Phosphate changes in magnesium dosed urine solutions 68 Figure 2.9. Effect of magnesium on pH 70 Figure 2.10. Changes in ammonia concentration with time 71 Figure 2.11. Changes in phosphorus with storage 72 Figure 2.12. Percent loss of phosphorus, calcium, and magnesium 73 Figure 2.13. Allocation of phosphorus 74 Figure 2.14. Allocation of calcium 75 Figure 3.1. Schematic of experimental design 85 Figure 3.2. pH changes in solutions 89 Figure 3.3. pH changes in solutions that were dosed with magnesium 90 Figure 3.4 Changes in ammonia with storage time 92 Figure 3.5. Ammonia changes in solutions dosed with magnesium 93 Figure 3.6. Spontaneous removal of calcium, magnesium and phosphorus 95 Figure 3.7. Allocation of phosphorus : 97 Figure 3.8. Concentration of phosphorus remaining in solution 98 Figure 3.9. Change in calcium over time 99 Figure 3.10. Allocation of phosphorus in solutions #3 and 7 100

viii Acknowledgements

I cannot offer enough thanks to the following people who have helped me navigate the world of graduate school, academic research, and life in general:

• My supervisors Jim Atwater and Don Mavinic who kept me motivated, inspired and on track for two years, which is no easy feat;

• Paula Parkinson and Susan Harper who endured me and my foibles on a daily basis for several months, and yet were never too busy to answer my questions;

• Fred Koch who helped me come up with this idea and convinced me that I was the one that should study urine;

• Imelda Wong and Kay Teschke who were not only wonderful at answering my questions, but very supportive and helpful when it came to matters more complicated than research;

• The students and mentors of the Bridge program who were always enthusiastic, inspiring, and never too busy to help, talk, or advise.

Above all though, my friends and parents were the ones who always had perspective, coffee, care packages, and phone calls to help me stay focused on what I wanted to do, and why I wanted to do it. Thank-you.

ix To Susan Smith and the courageous men and women of Tlamacazapa, whose daily struggle have taught me the true value of water Co-Authorship Statement

I alone conducted the research, data analysis, and manuscript preparation. Prof. James

Atwater and Dr. Don Mavinic assisted me in the design of the research program, suggested areas for further investigation, and offered valuable critiques of the completed manuscript.

Dr. Kourosh Afshar assisted me with the logistics of urine collection and storage. CHAPTER 1: Background

"Current waste management policies and practices are abusive to human well-being, economically unaffordable, and environmentally unsustainable "

-Bellagio Statement issued in February 2000 by Environmental Sanitation Working Group of the Water Supply and Sanitation Collaborative Council (Bellagio, 2000).

As inspired by this quote, the goal of this thesis to contribute to the growing field of work being done to reshape, rebuild, and rethink our relationship with the wastes we produce and the water we use as we strive for sustainable human and environmental health.

The current sanitation paradigm is based on water; fresh, potable water is used as a vehicle to dilute and transport wastes. Energetically and chemically intensive practices then re- concentrate harmful pathogens along with valuable nutrients in a mixed, foul sludge. Despite an increasing shortage of finite phosphorus reserves, this life-essential element is most often removed and disposed of along with a suite of biological and chemical contaminants.

The alternative to this paradigm is one in which wastes are not diluted, but separated, and are treated as resources, rather than wastes. In this way, water is conserved, pathogen spread is limited, and nutrients can more easily be recovered from concentrated resources. It is within this paradigm of source-separation and nutrient recovery, that this research is based.

In this work, I investigate the recovery of a valuable resource, phosphorus, from what was once considered a waste product, urine. By binding phosphorus in the form of a profitable bioavailable fertilizer known as struvite, an impetus to re-evaluate our sanitation systems and our definition of "wastes" is created. I examine how different storage conditions affect the composition of urine, and subsequently, the quality of struvite that is recovered.

This work draws on diverse fields of research: urology, crystallography, geochemistry, ecological engineering, and classic wastewater treatment technology. Each of these areas contributes to the larger picture and a holistic understanding of how phosphorus is removed from the body, how struvite can form in the body, how struvite can form in nature, and how

1 from the body, how struvite can form in the body, how struvite can form in nature, and how engineered design can cause the accidental, or intentional formation of struvite in infrastructure.

This thesis is presented in a manuscript-based format, which contains a literature review, two articles that will be submitted to peer-reviewed journals, and a concluding chapter, which contains both a summary of the work completed and recommendations for further research.

Chapter 1 is a literature review and is divided into two sections; the first section, Water and Sanitation, summarizes the state of pioneering sanitation research, the barriers to urine collection, and the resource potential of urine obtained in this manner. The second section, Phosphorus, addresses the need for phosphorus recycling, the chemistry of phosphorus mineral formation, the importance of phosphorus in urology, and the state of struvite research, including a summary of why recent work to recover struvite from urine is inadequate for wide-spread application. The literature reviewed covers the social, technical, and scientific aspects that would be expected in a multi-disciplinary study, such as this.

The first article (Chapter 2), "Preparation of urine for optimal struvite recovery" examines the changes that occur in synthetic urine and consequently, how it can be prepared such that an optimal amount and quality of struvite can be recovered. Specifically, the goal of this work was to understand how nutrient levels in stored urine behave, how varying the amount of magnesium added at the time that magnesium is added affects the solution, and when struvite can be best harvested from urine solutions.

The second article (Chapter 3), "Recovery of struvite from stored urine", presents the findings from struvite recovery studies using real human urine. As well, the feasibility and suitability of using synthetic urine versus real urine for nutrient recovery studies are compared.

The concluding chapter (Chapter 4) summarizes the key findings from the two research papers, highlights the strengths and weaknesses of the work, discusses the significance of

2 this work in the context of the research field, offers suggestions and directions for future work.

Appendices are found at the end of the thesis. Appendix I contains the data used in Chapter 2, Appendix II contains data the data used in Chapter 3, Appendix III contains the Certificate of Approval from the Clinical Ethical Review Board and Biohazard Approval Certificate, and Appendix IV contains the abstracts that accompany the manuscripts in Chapter 2 and 3.

The adoption of struvite recovery from urine will not be immediate, but the planning for a dramatic change in the way we use resources must begin immediately. It is my hope that this research will add momentum to the gradual shift towards a sustainable, healthy, equitable future.

1.1 Literature Review

1.1.1 Water and Sanitation The Millennium Declaration, which was established by the United Nations in 2000, set eight goals for development to address the increasing disparities in health, education, and equality. Currently, one billion people lack access to safe drinking water and 2.4 billion do not have adequate sanitation. One of these goals, Goal 7: "Ensure environmental sustainability", includes the condition that by 2015, half of the lacking population will receive access to safe drinking water. That means improvements will have to be made for 274,000 people/day until 2015 (UNESCO, 2006). This goal, while seemingly unachievable, is of dire importance given that around 4 million people die from polluted water every year; water that in most cases is polluted by lack of sanitation services (Otterpohl, 2001).

From a technical standpoint we understand that safe, abundant water is essential to good health, but what we often overlook are the social implications. Deaths and diseases are quantifiable; insecurity, poverty, oppression, violence, and despair, are the incalculable and dehumanizing consequences of a world in which water has increasingly become a right of a few (Smith and Marin, 2005). The importance of clean water then, must be recognized as

3 more than just a key to health and hygiene, but as a fundamental component of equality, and as a human right (Klawitter, 2005).

1.1.2 Sanitation Systems

1.1.2.1 Current Western model The Bellagio Statement, issued in February 2000 by the Environmental Sanitation Working Group of the Water Supply and Sanitation Collaborative Council asserts that current waste management policies and practices are abusive to human well-being, economically unaffordable and environmentally unsustainable, and that radical changes must be made (Bellagio, 2000).

The current sanitation system is a linear, end of pipe solution with no provisions for water conservation or nutrient reclamation. The system is based on the assumptions that a) human excreta is a waste, and there is little value in it b) waste is best diluted with water for treatment and transportation and c) treatment is best done at a centralized facility. In this way, vast quantities of water are chemically and biologically contaminated, the potential to recovery valuable nutrients is lost, large energetic and chemical requirements are used to treat the waste centrally, and water that could be otherwise left for environmental capital is. used as a transport medium to convey and disperse pathogens widely (Larsen et al., 2001).

1.1.2.2 Ecosan The average person (in the western world) produces about 500 L of urine and 50 L of faeces a year. Thus, by using a traditional flush toilet a person can use and contaminate between 20,000 to 100,000L of water (Otterpohl, 2001). Not only is this a gross waste of water resources, it is an unfortunate waste of nutrients and energy. Ecological sanitation (Ecosan) is a paradigm of sanitation that emphasizes resource reclamation rather than waste treatment. The focus of Ecosan is to reduce water use, reduce energy use, recover nutrients and maintain human and environmental health. In this framework, there is no such thing as "wastewater" because neither water nor nutrient rich human by-products are wastes. By shifting from isolated problem-solving (i.e. end of pipe treatment) to a holistic, integrated approach, drinking water, surface water, energy and nutrients can be used sustainably while maintaining a high degree of integrated public health (Wilsenach et al., 2003).

4 1.1.2.2.1 Infrastructure One of the hallmarks of Ecosan is decentralization, but decentralization need not mean a loss of comfort or an arduous amount of maintenance. The Ecosan technologies available range from basic stand-alone units to state-of-the art systems with semi-centralized treatment. Composting toilets are water-free units that compost human and plant material (for C:N control) to produce a nutrient rich, sanitized compost. Inexpensive, simple composting toilets are an excellent sanitation solution for developing and arid regions, but highly engineered, showroom style units are also available for the conscientious, hands-off homeowner (Jenkins, 1999).

Vacuum toilets (like those found on airplanes) are not appropriate for undeveloped areas as they are heavily engineered, but they use very little water and can be installed at any angle or location in a building (since gravity does not affect the plumbing). In Germany, a vacuum toilet system that was installed for a housing unit of 350 was locally centralized in order to produce biogas, energy, and fertilizer for a local farm (Otterpohl et al., 1997; Otterpohl, 2001) .

Separating, or Dry Toilets, conduct the liquid and solid matter to separate storage units. By separating the resource streams, urine can be harvested for use, and faecal matter can be dried and sanitized. In a developing context, dry toilets can be operated with no water and the solid matter can either be composted or sun dried (Otterpohl, 2001; Peasey, 2000). Development agencies and NGOs have successfully promoted dry toilets in many parts of Latin America and Africa, but as one researcher in Tanzania points out, the systems should ... "not be regarded merely as a second-rate solution for poor people" (Mashauri and Senzia, 2002) . Externally, the NoMix toilet is indistinguishable from a classic toilet, except that functionally, it is completely different. Separated urine is stored in a tank and the faeces are flushed with a small amount of water into the sewer system. Stored urine can be collected and spread on fields, or bled into the sewer system to reduce peak loads and decrease the risk of overflows (Larsen et al., 2001). With a low-flush separating toilet the amount of water used for a flush is dependent on the type of waste to be flushed; a flush for urine is only 200 mL and a flush for faecal matter is a little larger but can vary. Because so little water is used,

5 separating toilets can reduce total water use to less than 10 L/person/day (Berndtsson, 2006; Otterpohl et al., 1997). Many new toilets are low flush and water saving, but since they don't separate they are still diluting and dispersing nutrients, which makes recovery difficult. The hygienic danger associated with sewage, or blackwater, is almost entirely from the faecal matter, whereas most nutrients (N, P, K) are from urine (Otterpohl, 2001).

1.1.3 Source Separated Urine 1.1.3.1 Composition and Behaviour The study of urine has classically been relegated to the urological domain, however with the advent of Ecosan technology, engineers have taken it upon themselves to investigate the behaviour and properties of urine in the context of population production; a scale unheard of in the health sciences.

The average mass of N and P in fresh urine was found to be 13 and 1 g/person/day respectively, but it was also found that the values fluctuate over the course of the day (Hellstrom and Karrman, 1996). Urine contributes the greatest percentage of phosphorus and nitrogen to the waste stream: 80-90% of the nitrogen and 55-70% of the phosphorus (Berndtsson, 2006; Hanaeus et al., 1997). So, considering that 4/5 of flushes are for urine only, it is logical that non-diluted, source separated urine is a potentially concentrated source of nutrients (Berndtsson, 2006; Wilsenach and Van Loosdrecht, 2004).

Research into the dynamic nature of urine has shown that increased temperature, dilution, and the presence of wastewater all increase the rate of urea conversion, which results in elevated pH values and ammonia concentrations. Stored urine has been found to be stratified with respect to nutrients, pH and pathogens; the highest levels of each being found within the bottom 5 cm of the tank (Hoglund et al., 2000). Acidification, either with sulphuric or acetic acid, was found to retard or stop the conversion of urea, but sulphuric acid was more effective and acid addition was significantly more effective when added before conversion began (Hanasus et al., 1996; Hellstrom et al., 1999; Hellstrom and Karrman, 1996). Thus, the temperature at which urine is stored, the length of time for which it is stored, the degree of

6 contamination, and the amount of water with which the urine is diluted, can significantly impact the composition of stored urine, and thus, its suitability for reuse. One of the unfortunate limitations of urine storage is the formation of mineral deposits that can sometimes clog parts of the system. Work by Udert et al. (2003), more than any other, has contributed to the understanding of precipitation dynamics, precipitate composition, and the implications of precipitation in urine collecting systems. In a series of experiments, precipitates from 4 different urine collection/storage systems were collected in order to determine how ureoloysis and dilution affect the amount and type of precipitates that accumulate. To verify the results, a model was developed to calculate the thermodynamic equilibrium of the urine system by considering the solubility equilibria of 15 different solutes (urea, calcium, magnesium potassium, sodium, ammonia, carbonate, phosphate, sulphate, chloride, citrate, oxalate, protons, hydroxyl ions) and their respective complexes (Udert et al., 2003a; Udert et al., 2003b; Udert et al., 2003c). The results showed that: - calcite, struvite and hydroxyapatite were the only precipitates to form; - dilution influences that type of precipitates formed; and - dilution with rainwater and tapwater were different.

Four different real collection/storage systems were examined with dilution factors (DF=total volume/volume of urine) of approximately 600, 30, 4, and 1. Each level of dilution was found to produce a unique composition of precipitates. Only calcite was found in the highly dilute systems (DF=600) whereas HAP was found consistently at all lower dilutions. Struvite was found in the same dilution range as HAP, although the proportion of struvite increased as dilution decreased until it effectively dominated at low (1 and 4) dilution factors (Udert et al., 2003b). The effect of dilution was examined in detail with the use of a model, which was used to estimate the precipitation potential and composition of different complexes. Using dilution factors up to 30, the model results indicate that struvite and HAP would coexist up to a DF of 18, with struvite decreasing from above 60% down to zero with increasing dilution. Calcite formation began at a DF of 18 and increased continuously as HAP content decreased accordingly. However, these results were shown to vary when rain water, i.e. water with low levels of calcium and magnesium. Modelled results indicated that

7 In conjunction with dilution, the degree of ureolysis was found to play an important role in determining the precipitation. Using the model developed, Udert shows how the saturation indices for calcite, struvite and HAP vary with the percentage of hydrolyzed urea. Calcite and HAP are saturated initially; however, struvite is undersaturated until 2.5% of urea is hydrolyzed. After 20% of urea is hydrolyzed, all three minerals achieved maximum saturation and there was no change in saturation index or pH after this point (Udert et al., 2003b). Moreover, free urease was found to be responsible for 23% of the ureolysis (Udert et al., 2003a).

Phosphate is removed from urine in the form of HAP or MAP. The form and amount of precipitate is influenced by the dilution factor and the type of water used for dilution. Modelled results showed that a minimum amount of phosphate, 28%, is lost in undiluted urine and up to 100% when the urine is diluted 20 times. This was true however only for tap water dilution- phosphate removal actually decreased with increased rainwater dilution. Since calcium is limiting to the system, tap water provided a source of calcium, and thus, allowed phosphate to be removed completely. Calcium was limiting in rainwater and after calcium was exhausted, no further P removal occurred (Udert et al., 2003c).

Based on the model, 100% of calcium was removed from urine diluted with either tap or rain water at all dilutions less than 30, however removals in the actual system ranged from 50% to 95%. It is interesting that the actual amount of calcium measured increased with decreasing dilution, i.e. the less dilute the system, the higher the recovered calcium percentage (Udert et al., 2003c).

1.1.3.2 Contamination If urine is to be used as a fertilizer, the quality of urine, not only the nutritive value of it, but the presence of contaminants, will be scrutinized. Classically, farmers have worried about heavy metals contaminating synthetic fertilizers, but now it is the pharmaceuticals and trace chemicals that worry them: 30% of Swiss farmers are concerned about micropollutants in urine (Lienert et al., 2003). In 1995 the average amount of cadmium in Swedish fertilizer was 26 mg/kg P; a study of metals revealed that source separated urine had on average, <3.2 mg cadmium/kg P. Overall, heavy and trace metals (Pb, Hg, Cr, Ni, Cu, Zn) in stored urine were

8 found to be either non-detectable or well below Swedish fertilizer standards (Jonsson et al., 1997).

Urine is generally considered to be sterile and clean, although a few pathogens are carried by urine. In practice however, it is faecal contamination that is the primary concern because it is easy for faecal material to enter the urine collection system and contaminate the urine solution (Hoglund et al., 1998; Jonsson et al., 1997). The common practice is to store urine for a sufficiently long time, such that it is sanitized; the recommended time in Sweden is 6 months but tanks are usually emptied every 3-12 months. It has been shown that although faecal material was low (about 1 mg/L), a variety of pathogens may exist in urine (Hoglund et al., 1998; Jonsson et al., 1997). One study investigated 9 organisms and found low levels of total coliforms, rare instances of E. coli, high concentrations of faecal streptococci, and

varying amounts of Clostridia. Depending on pH and dilution, most organisms die in less than 1 day, but others like CI. Perfringens, salmonellaphages, or faecal streptococci are persistent after more than 30 days. It was suggested that, since some faecal indicator bacteria had a short survival time, they are not good indicators of contamination, so faecal sterols, especially coprostanol could be used as a proxy indicator instead (Hoglund et al., 1998). Also, pathogen levels in urine are not constant; they vary with depth, and are especially concentrated within the bottom 5cm of tanks, likely because of adsorption to settled particles (Hoglund et al., 2000).

Micropollutants, i.e. the endocrine disrupting hormonal and pharmaceutical contaminants, are appearing in increasing quantities in waterways. Because most compounds occur in such trace quantities, traditional wastewater treatment does not have the robustness or specificity to treat an array of unidentified compounds hidden in the mix. The simple process of storage is known to facilitate degradation; ibuprofen for example, will degrade at 15 degrees in 2-5 days. Unfortunately, WWTPs do not usually have the retention times necessary to achieve complete degradation, especially for compounds that may require 15-30 days of storage. By source separating however, micropollutants can be dealt with at the source and dealt with in concentrations 50-100 times higher than in sewage. High concentrations treated with longer retention times in a urine solution with low soluble organics will likely make it easier

9 to apply specific treatments capable of preventing the ultimate discharge of the hazardous chemicals (Larsen et al., 2004).

1.1.3.3 Experiences and Attitudes By 1999, Sweden had installed about 3000 urine-separating toilets in both private homes and public buildings. Operationally, the separated urine from the toilets in a housing development or high density building, is fed into a central storage tank. The storage tank is sized based on average flush values (0- 0.7 L to flush urine, 3.5 L to flush faeces), and to give the urine a sufficient amount of time to become sanitized before it is used as a fertilizer (maximum 8 months) (Berndtsson, 2006; Hanaeus et al., 1997). One study estimated that, for a high- density building with 125 people, a urine collection/storage system would, despite water savings, cost an additional 7 Euros per person per year. The extra costs were blamed on the expense of trucking the urine 10 times a year; increased maintenance costs were cited

(Berndtsson, 2006). Unfortunately, stored urine was often found to be overly dilute either because of infiltration or because users selected the large volume flush, even when it was not needed (Berndtsson, 2006; Hanaeus et al., 1997).

Socially, the proper functioning of a separating system is highly dependent on awareness and acceptance (Berndtsson, 2006; Hanaeus et al., 1997; Pahl-Wostl et al., 2003). Some earlier systems had debilitating problems but designs have improved and systems now are generally functioning well (Fittschen and Niemczynowicz, 1997). When users and citizens were posed with the concept of urine separation, they were generally in favour of urine-separating toilets

(77% and 80 % respectively) although those who did not have the technology already were not willing to accept any additional costs or effort (Berndtsson, 2006; Pahl-Wostl et al.,

2003). Citizens were also positive towards nutrient recycling: 80% would rather have vegetables fertilized with urine rather than artificial fertilizer (Pahl-Wostl et al., 2003).

Farmers, as well, had positive attitudes towards using urine on their fields: 57% of Swiss farmers thought that urine-based fertilizers were a good or very good idea, and only 33% thought it was a bad idea. Even though nearly 100% of respondents said they would use hog manure, the majority admitted that they would prefer a grainy, rather than liquid fertilizer—

42% of farmers would even purchase a urine-based fertilizer if it were available (Lienert et al., 2003).

10 1.1.3.4 Life Cycle Analysis To determine if urine separation is sustainable—financially, energetically, or resource- wise—several analyses have been completed.

Energy is finite and can only be converted, whereas exergy, the part of energy that can be converted into other forms of energy, is systematically degraded. Essentially, it can be used to define the quality of a resource. Exergy analysis of different waste treatment systems (traditional WWTP, septic tanks and filter beds, and urine separation with biogas recovery), found that the exergy consumption of a urine separation system was lower than more conventional methods. This was mostly due to the lower consumption of electricity and the beneficial recovery of high-exergy nutrients (Hellstrom and Karrman, 1997). The exergy balance of treating organics and nutrients was also calculated by taking into account operation, methane production, N and P recovery, transport, and spreading. Conventional treatment with reverse osmosis was found to consume the most exergy, while urine separation consumed the least (Hellstrom, 2003).

Various life cycle assessments have been performed in Sweden (in Swedish) and by and large, they all indicate that urine separation is far preferable to conventional sewage systems (Bengtsson et al., 1997; Jernlid and Karlsson, 1997; Tillman et al., 1996). A life cycle analysis including energy, materials and environmental considerations was conducted to evaluate 2 methods of preventing nitrogen discharge. Compared to denitrification, urine separation and spreading consumed only 42% of the energy and had only 52% of the greenhouse potential (i.e. CO2 equivalent) (Maurer et al., 2003). A similar study concluded that urine separation can have energy savings up to 36 % depending on how far the urine must be transported, but up to 221 km, energy savings could be realized (Jonsson, 2002)

1.1.3.5 Options for Nutrient Recovery

1.1.3.5.1 Faecal Matter

Because faecal matter has far fewer nutrients than urine, research into the reuse of solid matter has not been as vigorous as that of urine. Still, as part of a holistic approach to waste management, the need to treat faecal matter cannot be ignored. In a urine separating system,

11 the standard method is to flush it but other options have been investigated such as thermal composting (Vinneras et al., 2003a) or disinfection with urea (Vinneras et al., 2003b), both of which are quite effective at inactivating most pathogens. A whirlpool surface tension separator has been shown to effectively recovery the majority of nutrients from faecal matter when the system is combined with urine separation (Vinneras, 2004). Peasey presents a comprehensive summary of the health-outcomes and issues associated with numerous working urine-separation systems in Mexico and South America (Peasey, 2000).

1.1.3.5.2 Urine Spreading raw urine on farmers' fields is not the only way to use source-separated urine. Various methods of using, transporting, and transforming it have all been investigated to encourage the viability of separation.

Larsen (Larsen and Gujer, 1996) suggested near the onset of separation technology that urine, or Anthropogenic Nutrient Solution (ANS), could be collected and sent to a central treatment facility to be treated separately. By removing urine from the dilute waste stream, he argued, nitrification and BNR wouldn't be necessary and energy and oxygen requirements would decrease. Models of this concept have been developed and have shown variable results. Model results from a modified UCT system show that complete nutrient removal could be possible with 75% urine separation (Wilsenach and van Loosdrecht, 2003), while another study showed that only nitrate removal increased with increasing separation (60%) (Wilsenach and Van Loosdrecht, 2004). A more recent investigation concluded that the nutrient removal gains were good but were far less important than the resource savings and energy production possibilities that come with urine diversion (Wilsenach and van Loosdrecht, 2006). A stochastic model was developed and used in a virtual case study; by varying urine discharges to the treatment plant, peak dry weather ammonia loads could be reduced by 30%. Furthermore, 50% of the annual urine released as combined sewer overflows could be prevented (Rauch et al., 2003). Although the separation of urine as a "green" concept is attractive, it is still a long way off. However, by considering operational improvements to overall wastewater management, a policy window for more widespread adoption could be realized.

12 As shown in various LCAs and energy balances, transport is often the most significant cost associated with urine reuse. Freezing urine was investigated as a way of reducing the water content and concentrating nutrients into a smaller volume for easier, more efficient transport (Gulyas et al., 2004; Lind et al., 2001). Because the process of freezing excludes ions from the ice lattice, about 80% of the nutrients in urine could be concentrated into 25% of the , volume at a temperature of-14°C (Lind et al., 2001). However, because the energy inputs for freezing were so high, freezing would only be cost effective when the urine had to be transported more than 75 km, otherwise, the energy to freeze the urine would greater than that of the transport (Gulyas et al., 2004).

More unconventional research has focused less on phosphorus recovery and more on nitrogen recovery. Isobuthylaldehyde di urea (IBDU) is a slow release fertilizer that is currently produced by the chemical industry from aqueous urea solutions. An attempt to extract IBDU from urine was successful but limited by the energy inputs required (Behrendt et al., 2002). Electrodialysis, as a means of concentrating solutes for reuse, was found to have satisfactory results (Pronk et al., 2006) and natural zeolites have been used to extract ammonia from urine for further reuse (Beler-Baykal et al., 2004).

Still another recovery process is the formation of struvite (magnesium ammonium, phosphate) from urine; given the high quantity of phosphate and ammonia in urine the recovery of struvite seems only logical. To date though, only two studies have been done to determine the feasibility of recovery struvite from urine. Larsen and Gujer (Larsen and Gujer, 1996) mentioned the idea in 1996, but experimental results did not emerge until 2000. The work examined the co-removal of struvite and ammonia with zeolite adsorption from fresh or synthetic urine with the use of MgO as a magnesium source. Even though struvite was recovered, it precipitated in conjunction with brucite and montgomerite which have unknown value. The magnesium oxide used as a source of magnesium, also affects pH, thus the control of one parameter is inextricable from the other. The presence of zeolites had potentially synergistic effects and likely leached metals into the solution which contaminated the struvite with trace metals (Ban and Dave, 2004; Lind et al., 2000). As a model for the recovery of struvite from source separated urine, this work was not entirely appropriate.

13 Stored urine has at least 28% less phosphorus than fresh urine (Udert et al., 2003b) and in a realistic setting, fresh urine would be next to impossible to obtain.

1.1.4 Summary Not only is Ecosan technology a way of providing a reliable sanitation solution for the billions whose health is in jeopardy because of lacking hygiene services, it has the potential to provide a more efficient alternative to replace the current wasteful systems. Research in this field is growing rapidly and many experimental methods and systems for waste treatment and reuse are being tested. The work that was done to recover struvite from urine opened up a new field of research and laid a foundation for other researchers to examine the possibility of recovering phosphorus from urine.

1.2 Phosphorus

1.2.1 Phosphorus as a resource Biological production on earth is limited by water and nutrients. Increased demand, intensifying pollution, shifting resource allocations and dwindling energy supplies threaten both. Nitrogen, which is one of the three essential nutrients (along with phosphorus and potassium) is constantly cycled through the water, soil, and atmosphere by various physical and biochemical processes, ensuring a constant supply for primary production. Phosphorus on the other hand, has no natural cycle to replenish degraded stocks. Once a deposit is removed and dispersed, there is no natural mechanism for recovering and reconstituting that mineral other than the eternity it takes for the natural process of sedimentary deposition (Krauskopf, 1967).

Economically exploitable reserves are running out; complete depletion is expected in 60-130 years and half of the currently economically recoverable phosphate resources will be exhausted in 60-70 years (Steen, 1998). Since global agriculture depends heavily on phosphorus-based fertilizers, (mineral fertilizers consume (80%) of global reserves (Steen, 1998). biological production could be severely impacted as phosphorus becomes more scarce. Dwindling supplies mean more than just finding the last remaining deposits:

• Strip-mining damage. More than 75% of phosphate rock is surfaced mined in a process which has tremendous environmental effects not only because of the

14 surface scarring, but also because of the waste associated with disposing of the 60-95% of the rock that does not contain phosphate. • Increasing contamination. Because it's highly reactive, phosphorus never occurs in its pure form. Phosphorus is often bound to innocuous metals like calcium and sodium, but it also readily binds with heavy metals like cadmium. Thus, mined phosphates must be heavily processed to remove impurities (and processing = resource use), • Increasing energy investments. As phosphate reserves diminish, the energy needed to mine and refine them, will increase. Concentrated, uncontaminated reserves will deplete to the point where the energy costs associated with recovery of low-grade reserves will prove uneconomical (Steen, 1998).

Excess phosphorus (from sewage or fertilizer runoff) can eutrophy water bodies. As such, great progress has been made to decrease phosphorus inputs to the environment (Randall, 2003). Unfortunately, most technologies developed simply convert dissolved phosphorus into a solid, non-usable form. The precipitation of phosphorus from wastewater with iron or alum, the accumulation of P in a microbial mass, or the uptake of P in wetlands are all useful at preventing point discharges, but the P recovered is usually contaminated or inappropriate for land application (de-Bashan and Bashan, 2004).

In summary, raw phosphorus resources are diminishing while phosphorus pollution is increasing. What is needed then, is a phosphorus source that is a renewable, sustainable and has little environmental affect. Toilet wastes, urine, wastewater sludge, incinerated sludge, and wastewater have been examined as potential source of P, but in most cases, the appropriate, cost effective technology does not exist (Balmer, 2004). One exception is the recovery of phosphate from human sewage, in the form of struvite. For many years, the mineral struvite (magnesium ammonium phosphate) has been known as an effective, phosphorus based fertilizer (Bridger et al., 1962; Lunt et al., 1964). Recent technologies developed in Canada and Japan have made the recovery of struvite from wastewater possible, signifying the beginning of the first large-scale, recycled phosphorus industry (Abe, 1995; Adnan et al., 2004; Shimamura et al., 2003). It will be these and other developing

15 technologies that will provide sustainable, economically viable sources of phosphorus to buoy and maintain biomass production (i.e. for food, ethanol production, fibres, etc) through the impending phosphorus shortage.

1.2.2 Phosphorus Chemistry Phosphates can occur in several different forms. An orthophosphate molecule has a tetrahedral shape with the P atom in the centre surrounded by oxygen atoms at the corner points. Polyphosphates and metaphosphates are formed when two or more orthophosphate groups join together such that there is a P-O-P chain. As a group, polyphosphates are linear, while metaphosphates are cyclic (Snoeyink and Jenkins, 1980).

Phosphates are naturally found with metals such as calcium and magnesium. This section will discuss the chemistry of phosphate minerals, specifically calcium and magnesium phosphate minerals, with special attention paid to struvite.

1.2.2.1 Solubility Constant The solubility constant describes the constituent ion concentrations of a mineral dissolved in pure water at equilibrium; theoretically, it is valid for any pH. (Doyle and Parsons, 2002; Snoeyink and Jenkins, 1980).

The solubility constant or (solubility product) is derived from the general equation:

y+ z (1) A2By--> zA +yB ' such that

(Snoeyink and Jenkins, 1980) Curly brackets denote activity, rather than molar concentration. Activity is the product of the molar concentration and the activity coefficient:

(4) {C}=Y[C]

16 To determine the solubility product, an accurate measurement of the ionic species in solution is required. By introducing the ionization fraction into the equation, the solubility product can be calculated with total analytical concentrations of the ions:

(5) K=(YMg2+CTMgCt Mg2±)(YNH4+CrNH3Cl NH4.+)(Y P043-CTP04CX P043) where a is the ionization fraction, y is the activity coefficient, and Cj is the total analytical concentration (of magnesium, ammonia nitrogen and orthophosphate) (Snoeyink and Jenkins, 1980). A discussion of the two coefficients follows.

1.2.2.1.1 Activity Coefficient Activity describes the "effective" concentration of an ion, i.e. the "part of the analytical concentration that determines its behaviour towards other ions" (Krauskopf, 1967). In very dilute solutions where ions behave independently of each other, activity can be approximated with molarity (mol/L) by setting the activity coefficients equal to one. However, as the ionic strength of the solution increases, electrostatic interactions between ions increase and the activity of ions becomes less than the measured concentration. Thus activity coefficients must be employed to lower the "measured concentration" to reflect the true activity of the ion (Snoeyink and Jenkins, 1980).

Activity coefficients are a function of ionic strength, u., which is used to "describe the intensity of the electric field in a solution" where

(6) u=l/2 2(CiZi2) and C= the molar concentration of the ion and Z is the charge of the ion. There are several methods of calculating an activity coefficient, and each method is appropriate for a different type of solution. • DeBye-Huckel limiting law is valid for ionic strengths less than 5x10" and is given

by (7) -log(YO=0.5ZiV1/2

• The Extended DeBye-Huckel (DH) equation is applicable for ionic strengths less than 0.1 and is given by the following equation.

(8) -log(Yi)=AZi2u.1/2/l+BaiU.1/2

17 where A and B are constants that relate to the solvent temperature and ctj is a constant that relates to the diameter of the hydrated ion. • The Guntelberg approximation is used to calculate y for ions at temperatures other than 15°C: (9) -logY=0.5Ziy/2/l+u,'/2 which is just the extended DH theory with A=0.5 and Baj=l (Snoeyink and Jenkins, 1980). • The Davies modification is used for high ionic strength solutions. Generally it can be used for solutions with ionic strengths, less than 0.5M although it is unclear whether it should be used for solutions over 0.1M (Pitzer, 1979). It is given by: (10) -logY=0.5ZiV'/2/(l+^1/2)-0.2n (Stumm and Morgan, 1981)

Unfortunately, there is no good theory for activity coefficients at ionic strengths greater than 0.5. The ionic strength of urine is typically 0.33 M and thus, the Davies modification is most appropriate although questionable depending on some critics (Babic-Ivancic et al., 2004).

1.2.2.1.2 Ionization Fraction The ionization fraction, denoted by a, is independent of the measured concentration, but is pH dependent. Although, it is usually ignored to simplify calculations, it defines the fraction of the free ions to the total dissolved species that exists at a given pH. The ionization fraction is important because it can be used to alter the analytical concentration to more accurately reflect the presence of reaction-available ions (Snoeyink and Jenkins, 1980).

1.2.2.2 Conditional Solubility Product Accurate determinations of the solubility product are made difficult by the limits of available analytical techniques. The conditional solubility product (Ps) is dependent on pH and solution conditions but is useful because it makes use of total rather than free ion concentrations; thus, it can provide a solution-specific solubility constant that is far easier to measure.

(11) PS= CT,MgCT,NH3Cr,P04 where CT, is the total concentration of the ion or anion in all its complex forms.

The equilibrium conditional solubility product (Pseq) can also be defined for specific solutions and is given by the following:

18 (12) Pseq-K/(YMg2+a Mg2+)(YNH4+« NH4.+)(Y P043-01 P043r) If the conditional solubility product is the same as the equilibrium then the system is at

equilibrium (Ps= Pseq).

1.2.2.3 Supersaturation Ratio The supersaturation ratio (SSR) defines the degree to which a solution is supersaturated with a given compound. The supersaturation ratio is defined as the ratio of conditional solubility product (Ps) and the equilibrium conditional solubility product (Pseq). The SSR is given as:

eq (13) SSR= Ps/ Ps When the ratio is greater than 1, the potential for precipitation exists; conversely if the-value is less than 1, conditions for dissolution exist (Stumm and Morgan, 1981).

1.2.3 Struvite Since Borgerding (Borgerding, 1972) discovered struvite in digested sludge lines at an activated sludge plant, concerted efforts have been made to not only prevent its accidental formation, but to encourage its formation under controlled conditions such that it can be reclaimed as a renewable source of phosphorus.

1.2.3.1 Struvite Morphology Struvite is a whitish, crystalline mineral that is comprised of magnesium, ammonium, and phosphate in equal molar amounts (Doyle and Parsons, 2002). Struvite crystallizes in an orthorombic system built with regular PO4 " tetrahedra, distorted Mg(H20)6 octahedra and

NH4"1" groups that are all held together with hydrogen bonds (Abbona and Boistelle, 1979). Abbona, a leading mineral researcher who has published numerous papers on struvite over the past 30 years, showed how struvite habit is variable and dependent on growth conditions. Pure struvite crystals do not take one form, but rather assume a unique habit depending on the conditions under which they were grown. By growing crystals from pure solutions, Abbona and Boistelle showed that different growth kinetics would either promote flat crystals with large faces or stick like crystals, with gradations of face size in between (Abbona and Boistelle, 1979). Impurities, pH and supersaturation were also identified as factors which influence struvite habit (Abbona et al., 1984). In an equal molar solution where [Mg]=[P] between 2x10" - 5x10" and over a range of pH 5.5-10, it was shown that individual crystals formed at lower pH values than aggregates and at increasingly higher pHs

19 as the concentration decreased (Abbona and Boistelle, 1985). At higher concentrations, aggregates form too easily to detect single crystals. The line that can be drawn between the region where no crystals form and where single crystals form is the boundary of the metastable limit.

1.2.3.2 Struvite Crystallization From a supersaturated solution, struvite formation occurs in three steps: nucleation, crystal growth, and agglomeration and ripening. However, supersaturation alone does not predict precipitation; a supersaturated solution can be stable indefinitely. Kinetic factors and the presence of fine particles also affect when and how struvite will form.

1.2.3.2.1 Nucleation Homogenous nucleation is the formation of a crystal nucleus from the constituent ions. This type of nucleation is rare because of the energy requirements needed to form an organized structure from a random solution. Heterogeneous nucleation—the use of foreign particles as nuclei—is less energy intensive, and more common (Snoeyink and Jenkins, 1980).

1.2.3.2.2 Growth

Once a nucleus is formed, a crystal begins to grow. The growth rate is described by :

(14) ^£m-_fcS(C-C*)" where C* is the saturation concentration (mol/L), C=actual concentration of limiting ion (mol/L), k=rate constant, S=surface area available for precipitation, and n=constant to describe the diffusion of ions to the surface (when the diffusion growth rate, the value of N is unity). Growth continues until a state of equilibrium is reached (Snoeyink and Jenkins, 1980).

1.2.3.2.3 Agglomeration and Ripening Any number of nuclei may form and grow, but crystals do not necessarily persist. A crystal that forms may be unstable and may transform over time. The term "ripening" describes the change in crystal size and "aging" describes a change in crystal structure. As large particles grow, smaller particles may either dissolve because the saturation of the solution has been reduced or they will agglomerate to form large particles (Snoeyink and Jenkins, 1980).

20 1.2.3.3 Struvite Formation The most common equation given for the formation of struvite assumes that free ions bind with 6 moles of water to form fmol of struvite:

2+ + 3 (15) Mg +NH4 +PO4 "+6H20 MgNH4P04 6H20 (Doyle and Parsons, 2002) However, an alternative formula, used by mostly Japanese researchers (Mitani et al., 2003; Shimamura et al., 2003; Yoshino et al., 2003), considers the speciation of phosphate at the pH range necessary for struvite formation:

2 + 2+ (16) HP04 " +NH4 +Mg +OH- +5H20-» MgNH4P04 6H20 Still another equation is given by Babic (Babic-Ivancic et al., 2002), in which protons are shown to be generated, rather than hydroxyl ions being consumed:

2+ + + (17) Mg + NH4 +H2P04- +6H20-»2H + MgNl^PCU 6H20

This equation is more reflective of the state of phosphate at the working pH; for H2P04"

2 7207 2 -»HP04 " + H+ the dissociation constant is 10" so that above pH 7.202, the HP04 " species is dominant (Abbona and Boistelle, 1985).

1.2.3.4 The Solubility Product of Struvite Since struvite has gained notoriety in the wastewater industry, a clear understanding of struvite chemistry has become increasingly important. One of the most fundamental and debated points of struvite formation is the solubility constant, Ksp. Predictions of the struvite solubility constant have been made either by forming or dissolving struvite in distilled water, however particle size, aging, complex formation, and impurities all alter the experimental value (Snoeyink and Jenkins, 1980). Calculating an accurate Ksp value is difficult since ionic strength and activity must be calculated based on accurate measurements of conductivity, pH, temperature, and concentration. Another complicating factor is that of side reactions. All of the struvite-forming ions are capable of forming numerous complexes with other species in solution and the type and number of complexes formed is dependent on the solution matrix and the conditions of the solution. It is important to note how ionic strength, and therefore the choice of equation will change dramatically over the range of solutions. Doyle (Doyle and Parsons, 2002) presents a summary of pKsp values that have been reported; they range from a high of 9.4 (Borgerding, 1972) to a low of 13.26 (Ohlinger et al., 1998). Despite similar methods, each researcher chose a different set of equilibria equations and thus, different results were produced in each case. Ohlinger did not include magnesium-

21 phosphate interactions (Ohlinger et al., 1998), Bouropoulos included magnesium ammonia interactions (Bouropoulos and Koutsoukos, 2000), whereas Babic ignored magnesium- ammonia interactions but included magnesium sulphate complexes (Babic-Ivancic et al., 2002). Aage et al. used P32 tagged struvite to determine a Ksp value of 12.94 but did not consider any magnesium or phosphate complexes (Aage et al., 1997).

1.2.3.5 Chemistry of Struvite Formation Struvite chemistry is complicated; ion concentration, pH, supersaturation, temperature and complexing agents will control when and how struvite will precipitate out of solution (Bouropoulos and Koutsoukos, 2000). Magnesium, ammonium and phosphate are required in a 1:1:1 molar ratio, but ions are involved in side reactions, which will limit availability. Also, the relative presence of the reactive species is affected by pH: ammonia decreases and phosphate increases with increasing pH. Over the range of pH 7-9 the percentage of ammoniacal nitrogen in the NH4+ form decreases from 99% to 64% until the pK value of 9.3 is achieved, at which point ammonia and ammonium are in equilibrium. Over the same range, the fraction of PO4-P present as PO43" increases from nearly 0 to over 60% (Booker et al., 1999; Stumm and Morgan, 1981). Clearly, the pH plays a strong role in determining the availability of NH4 and PO4 for struvite formation.

1.2.3.5.1 Effects of pH Struvite has been shown to precipitate out of solution at pH values between 7 and 11 (Doyle and Parsons, 2002). The controlling factor is not entirely pH but, in fact, whether or not the solution is in the metastable phase, which is partially controlled by pH. In full-scale operations a caustic, either sodium hydroxide or lime, is added to increase the pH to the target working range (between 8 and 10) but other methods such as aeration have been shown to be effective at increasing the pH up to 9.1 (Battistoni et al., 2002; Suzuki et al., 2002). Fujimoto (Fujimoto et al., 1991) found that NaOH was more effective than lime or MgOHb, and generally, NaOH is the preferred caustic in practice (Adnan et al., 2003a; Adnan et al., 2003b; Britton et al., 2005).

Although struvite can form over a range of pH values it does not form in the same way at all pH values. Abonna et al. showed that increasing pH decreased induction time (i.e. time for

22 nucleation to occur (Abbona et al., 1982) and it was also shown that if the pH was below 8 then the struvite precipitation was slow and could be delayed by several days (Battistoni et al., 1997). Using anaerobic swine lagoon liquid, Nelson showed that as pH increased up to a maximum of 8.9 or 9.25 (depending on the mixture), phosphate removal also increased (Nelson et al., 2003). Using human wastewater, different researchers identified pH ranges over which marked increases in phosphate removal were observed: between 7.6 and 8.8 (Britton et al., 2005), between 8 and 9 (Katsuura, 1998), and between pH 8.1 and 8.8 (Adnan et al., 2004). These have been identified as preferred operating ranges. One researcher found that beyond a pH of 8.5-9, there was essentially no change in P removal (Stratful et al, 2001). All of these data indicate that, in general, high pH values (>8) are preferable for struvite formation and that optimal pH ranges are very solution specific.

In a working reactor, crystal quality was impacted at higher pH values, i.e. more dendritic crystals and fines were produced (Adnan et al., 2004). Other data shows how morphology changes as a function of pH and magnesium concentration. At higher pH values there are more rod-like crystals, lower pHs yield more dendrites, and the lowest pHs yield rhombohedral crystals (however these data are not independent of magnesium concentration (Babic-Ivancic et al., 2002).

1.2.3.5.2 Minimum Solubility With increasing pH, struvite solubility decreases until a minimum solubility point is reached (Doyle and Parsons, 2002). In fact, minimum solubility is the point where the product of the constituent ions is a minimum (Snoeyink and Jenkins, 1980). Estimates of minimum solubility range from 9 (Munch and Barr, 2001; Nelson et al., 2003) to 10.3 (Ohlinger et al., 1998) to 10.7 (Stumm and Morgan, 1981). The differences have been attributed to CO2 and the effect of ionic strength, which decreases the effective concentration of ions (Babic- Ivancic et al., 2002; Ohlinger et al., 1998).

1.2.3.5.3 Effects of Temperature Temperature has a strong influence on struvite precipitation: it affects the solubility constant, the reaction rate, and the solubility, which increases with increased temperature (Snoeyink and Jenkins, 1980).

23 Aage showed that, beyond 10°C, solubility increases until it reaches a maximum at 50°C. After this point, solubility decreases until 64°C at which point the crystal structure appears to change (Aage et al., 1997). This is in agreement with Burns and Finlayson (Burns and Finlayson, 1982) who determined the solubility constants at 25, 35, 38 and 45 degrees and found that solubility decreases with decreased temperature.

Constructed morphology diagrams for struvite at both 25 and 37 degrees show that both the boundary of precipitation and area of different morphology types varied between the two temperatures tested (Babic-Ivancic et al., 2002).

In synthetic feedwater, the struvite growth rate was greater at 15 degrees than at 25 degrees and furthermore, the crystals formed at the lower temperature were generally larger, and thus, more useable. (Adnan et al, 2004)

1.2.3.5.4 Supersaturation Ratio Theoretically, when the SSR is greater than 1, struvite can precipitate, although mixing energy may also affect the point at which precipitation occurs. The exact effect is unknown, but mixing energy primarily controls growth rates and habit, whereas the actual nucleation is controlled by supersaturation (Ohlinger et al., 1999).

No variation in morphology was found over an SSR range from 1.13 to 3.33, although the induction time was found to be inversely related to the SSR (Bouropoulos and Koutsoukos, 2000). Similar results showed that there was an optimal supersaturation range over which the crystals produced were neither too small nor too rough, but an excessive supersaturation resulted in the production of fines (Hirasawa et al., 2002). In terms of P recovery, phosphate removal was shown to increase over a SSR range of 0-10, after which point, no change occurred (Britton et al., 2005).

1.2.3.5.5 Ammonia Increased ammonia concentration has been found to increase the purity of precipitated struvite. Working with an initial PO4 concentration of 742 mg/L, an essentially pure struvite product could be formed by increasing the ammonium:magnesium ratio beyond 15:1

24 (Stratful et al., 2001). In the presence of excess ammonium (NH4), P removal is also increased (Katsuura, 1998). Stratful et al propose that ammonium acts as a stabilizer for pH, i.e. it acts as a buffer which stabilizes the system when the pH drops due to struvite formation (see Equations 16 and 17). The authors also propose that the buffering effect is likely to be more significant at lower concentrations of phosphate when there would be a less noticeable pH drop from struvite formation (Stratful et al., 2001).

1.2.3.5.6 Magnesium Either magnesium chloride (MgCh) or magnesium hydroxide (Mg(OH)2) can be added to solution to provide the required magnesium, although each has specific advantages and disadvantages. Magnesium chloride is more expensive but it dissociates faster, whereas magnesium hydroxide is cheaper but affects pH such that it cannot be adjusted independently (Munch and Barr, 2001). In a comparison study using anaerobic supernatant, magnesium chloride was shown to work better (Wu and Bishop, 2004).

Several authors working with swine waste have reported that increased Mg: P ratios ranging from 1.1-1.6 result in increased P removal (Beal et al., 1999; Maekawa et al., 1995; Nelson et al., 2003). An increase in Mg:P ratio lowers the pH needed for formation, which practically, can reduce the need for caustic. Dastur (Dastur, 2001) showed that an Mg:P ratio of 1.3:1 was optimal for producing struvite in a crystallizer while other work showed that 3.3:1 produced crystals bigger than ones formed at 1.5:1 (Huang et al., 2006). However, this was contrary to other results which showed that higher ratios (1.3:1 and 4.59:1) produced unacceptable crystals and that a ratio of 0.34:1 produced far superior crystals (Adnan et al., 2004). The isoelectric point for magnesium is pMg=1.75 which indicated that formation may be favoured at higher magnesium concentrations (Bouropoulos and Koutsoukos, 2000)

1.2.3.5.7 Effect of Calcium Struvite formation can be inhibited by the presence of foreign ions and none is as important as calcium, because of its ability to compete for ions and form complexes (Momberg and Oellermann, 1992).

25 When the calcium:magnesium ratio is less than 0.25, struvite, rather than apatite, will be formed (Mitani et al., 2003). Because calcium phosphates precipitate under similar conditions as magnesium phosphates, the main impurities in crystallized struvite are often calcium and carbonate, about 3.4 and 5% respectively, of the crystal mass (Huang et al., 2006). One study showed that when the concentration of magnesium was lower than calcium, HAP or calcite would be the dominant minerals (Battistoni et al., 2002). In a study using swine waste (Ca: Mg= 188:129 mg/L) only 26% of the P removed was in form of HAP; even with an excess of calcium, MAP still formed preferentially (Suzuki et al., 2002). Since calcium can affect struvite formation and chemistry, it is important to examine the calcium phosphate system in order to better understand its potential for interfering with struvite recovery.

1.2.4 Calcium phosphates The calcium phosphate system is very complex. The number of compounds that can form, the competing side reactions, the ripening and aging with time, and the kinetic factors needed for nucleation all contribute to the complexity. The order of calcium phosphates in terms of increasing solubility is (Abbona et al., 1986):

o Hydroxyapatite (HAP): Ca5(P04)3OH

o Whitlockite: Ca3(P04)2

o Octacalcium phosphate: (OCP) Ca4H(P04)3 2.5 H20

o Monetite: CaHP04

o Brushite: CaHPO42H20 Theoretically, the minerals, in a supersaturated solution, should precipitate in reverse order, i.e. brushite then monetite, etc. However, the order of precipitation isn't solely based on solubility constant; kinetic factors and impurities will significantly impact the presence and composition of the calcium phosphates that precipitate (Abbona et al., 1986).

Abbona et al. (Abbona et al., 1986) investigated different ratios of calcium: magnesium in concentrated phosphate solutions (P=500-10 mM). After 24 hours they found that only brushite and amorphous calcium phosphate (ACP) had precipitated. This was unexpected because in the specific system OCP, HAP, whitlockite and monetite were more supersaturated than brushite, but still, they did not precipitate. Because brushite has the

26 simplest chemical composition and the fewest atoms, it is easiest to form; the others require multiple phosphate units and require more energy to nucleate. The calcium phosphates which precipitate first are not necessarily the least soluble or the ones that are most supersaturated, but the ones that have the highest nucleation rates (Abbona et al., 1988) and the ones that are the most hydrated with the simplest structure (Boistelle et al., 1993). The solubility constants of most minerals are known but most researchers agree that the kinetics ultimately govern the precipitation sequence (Abbona and Franchiniangela, 1990; Boistelle et al., 1993). Unfortunately, the kinetics of calcium phosphate precipitation are largely unknown. The order of mineral phosphate formation is dependent not on any one fact, but a combination of solubility constants, kinetics, concentration, pH, supersaturation and Mg:Ca and time (Abbona et al., 1986).

1.2.4.1 Mineral Conversion To better understand the aging process of calcium phosphates, crystals were examined after 1-15 months. Generally, the amorphous compounds that had precipitated early were converted to brushite crystals. Even though HAP was very saturated in solution, it did not occur as a final product. This could be explained by the fact that the critical supersaturation value had not been met. A critical value—one that is very high for HAP—must be achieved before precipitation can occur (Abbona et al., 1988). Apparently, supersaturation alone is not sufficient to induce nucleation- especially not for a complex mineral like HAP. There is a great deal of controversy over the formation of HAP. Some argue it converts directly from ACP and some believe there is an OCP intermediary; Abbona showed that depending on conditions, HAP will either evolve from ACP or OCP or nucleate directly from solution (Abbona and Baronnet, 1996).

As a phosphate crystal forms, the solution pH drops. The pH continues to drop over the course of the crystal's formation, but the steepest pH drop is within the first 1-5 hours. Also, the greatest pH drops were in those solutions with the highest SSR. Thus, as a crystal precipitates, the conditions for further precipitation are altered (Abbona et al., 1986).

27 1.2.4.2 Magnesium as an Impurity The presence of impurities will alter the order of phosphate mineral precipitation: at low concentrations, magnesium can alter the order of formation, or delay or prevent the conversion of an amorphous phase to a crystalline phase (Abbona et al., 1986). Magnesium has two effects on calcium phosphate formation: it forms complexes with phosphates and it can adsorb onto the precipitate, Which inhibits nucleation and growth (Abbona et al., 1988). Magnesium is a well-known inhibitor of calcium phosphates, specifically HAP (Boistelle et al., 1993; Wild et al., 1996). HAP formation is inhibited because magnesium stabilizes the ACP precursor, which results in smaller, shorter, more defected, and more irregularly shaped crystals (Abbona and Baronnet, 1996; Salimi et al., 1985). Magnesium will prolong the induction period of HAP by up to 4 times as well as disfigure the shape and decrease the size (Abbona and Baronnet, 1996).

It was found that magnesium depressed the supersaturation conditions of HAP, OCP and whitlockite more than it did for brushite, although brushite was inhibited because of the

increased ionic strength, the loss of HPO42" ions to magnesium complexes and the increase in acidity (Abbona and Franchiniangela, 1990). Most calcium phosphates are impaired in some way although dicalcium phosphate dihydrate is not noticeably affected (Salimi et al., 1985).

It is not however, the magnesium that solely affects calcium phosphates: Abbona found that calcium has more of an effect on the formation of magnesium phosphates than magnesium has on calcium phosphates (Abbona et al., 1986). After a certain point, magnesium ceases to be a calcium phosphates impurity, but becomes dominant such that magnesium phosphates, e.g. struvite, are formed.

1.2.4.3 Struvite and Calcium Phosphates Calcium phosphates usually precipitate before magnesium phosphates, so it is interesting to know what concentration, pH, and time conditions will favour the relative formation of each (Abbona etal., 1986).

Over a continuum of calcium and magnesium fractions (such that 0

28 (i.e. in the absence of other calcium phosphates) was only formed when xMg>0.8. Thus, struvite formation is facilitated by low calcium content (Abbona et al., 1986). So long as the supersaturation condition is met, struvite can form in concentrated solutions (i.e. P=500 mM) at pH values just above 6. In less concentrated systems (i.e. P=10 mM) pure struvite will not form until almost pH 7. With increasing concentration, the pH where struvite can nucleate is decreased- i.e. the more dilute the solution, the higher the pH that is needed to nucleate (Abbona et al., 1986). In urine-strength solutions, struvite was found to precipitate at high pH values (>9) but never alone- either with HAP or ACP which is expected since xMg

There does not appear to be a comprehensive method for determining the final composition of calcium or magnesium phosphates since so many factors are at play, not the least of which is the fact that each inhibits the others formation. It can be said then that the surest way to achieve a pure end member is to rule out complicating factors and to begin work with a solution free of the unwanted ions.

1.2.5 Urine Urine is a highly complex mixture of organic and inorganic compounds that varies in composition with diet, gender, age, hemodynamic function, endocrine status, body position, and time of day among other things (Diem and Lentner, 1970; Griffith and Dunn, 1978). The Documenta Geigy is the standard reference on urine composition; it itemizes, in detail, the metals, salts, amino acids, and other trace compounds that occur in urine (Diem and Lentner, 1970). Some stabilizing compounds have been tested that allow real urine to be stored until it is examined, but depending on the analysis, not all preservatives may be appropriate (Griffith and Dunn, 1978). Because urine is spatially and temporally variable, artificial urine is sometimes preferable for urological research when a standardized and replicatable solution is needed. However, reproducing urine in its entirety, i.e. including all metabolites and trace amino acids- would be impossible. Simplified "recipes" for synthetic urine have been developed which are solutions that include only the most important and prevalent components. There are likely hundreds of different synthetic urines in use, and often authors 1 simply create a research specific mixture (Grases et al., 1999) while others are still trying to

29 propose a universal solution (Brooks and Keevil, 1997). Burns and Finlayson (Burns and Finlayson, 1980) present a summary of the most popular recipes of the time, and present their own standard reference solution. Still, it is Griffith's simple solution of 11 organic and inorganic solutes devised in 1976 that is predominantly used by researchers, even 30 years later (Griffith et al., 1976). Griffith's recipe is based on the landmark work of Robertson et al. who compiled the data of 60 healthy and 60 stone-affected men to arrive at representative values (Robertson et al., 1968b). Griffith's recipe is given in Table 1.1.

Table 1.1 Composition of synthetic urine (adapted from Griffith et al., 1976) Species Concentration (g/L) Concentration (mmol/L)

Ca: 4.3 CaCl2-H20 0.65

MgCl2-6H20 0.651 Mg: 3.2 NaCl 4.6

2.3 S04: 16 Na2S04

Na3C6H807*2H20 0.65 Citrate: 2.3 (sodium citrate dihydrate)

Na2C204 (sodium oxalate) 0.020 Oxalate: 0.149

KH2P04 2.8 P04: 20.5 (potassium phosphate monobasic) KC1 1.6 NH : 19 NH4C1 1 4

CO(NH2)2 (urea) 25

C4H7N30 (Creatinine) 1.1 Total Na=118mEq Total K= 42 mEq

PH= 5.8 :

This solution includes the major ions, the stone-forming oxalate, the inhibitory citrate, the metabolite creatinine, and a significant amount of urea. Because of its general and comprehensive composition, it is appropriate for most types of urological research.

1.2.5.1 Urine and Struvite Uroliths, or "stones" which occur in the urinary tract, have been around for thousands of years. Historically, incidence rates have been as high as 90% (Hesse and Siener, 1997) but over the past century, the rate has dropped to around 4-15% in industrialized countries (Grases, 1998). In the past, infection stones (that result from urinary tract infections) were predominant, but since World War II, calcium oxalate stones (that are primarily the result of

30 dietary disorders) have become more prevalent and now account for 70% of urinary stones (Grases, 1998; St. Joseph's Hospital, 2006). Infection stones are phosphate-containing calculi; approximately 10-20% of them are composed of struvite with various amounts of HAP and carbonate apatite (Babic-Ivancic et al., 2004).

Another urological problem is the encrustation and blockage of indwelling urethral catheters. The crust of struvite that builds up is formed under similar conditions as infection stones and is compositionally very similar (Morris and Stickler, 1998).

Because the internal build-up of struvite has plagued humanity for so long, much research has been undertaken to understand the causes and mechanisms of deposition. What was likely the first solubility product (pK) for struvite in urine was reported in 1910 as 12.60 (Bube, 1910). Subsequent research put the value in the range of 12.41 -12.67 depending on the pH and ionic strength (Elliot et al., 1959). The value of 12.6 was later corroborated (Robertson et al., 1968a), although recent unpublished work suggests that the value is closer to 13.4 (pers. comm., Ronteltap et al., 2005). Determining a single solubility constant for struvite from a saturated solution with an ionic strength of 0.33 mol/dm (Sohnel and Grases, 1995) and hundreds of competing reactions, is a puzzle that may never be solved.

1.2.5.1.1 Inhibitors Calcium and magnesium phosphate inhibitors are ubiquitous in most body fluids- especially serum, saliva and urine (Eidelman et al., 1987; Garnett and Dieppe, 1990; Howard and Thomas, 1958). Inhibitors include phosphorylated proteins like statherin, phosphocitrate or citrate, pyrophosphate, peptide, glycoproteins, glycosaminoglycans like chondroitin sulphate, heparin sulphate and sulphate (Babic-Ivancic et al., 2004; Sharma et al., 1992; Suller et al., 2005). Urine is supersaturated for most of the stone forming constituents (i.e. calcium or magnesium phosphates); therefore, the study of inhibitors, especially citrate (NaaCeHjOv) and phosphocitrate (NasCe^OioP), are of great importance (Babic-Ivancic et al., 2004).

Phosphocitrate has been shown to be a strong inhibitor of HAP (Reddi et al., 1980; Tew et al., 1980; Williams and Sallis, 1979; Williams and Sallis, 1982). A comparison of both citrate and phosphocitrate on OCP showed phosphocitrate to be a much better inhibitor than

31 citrate, although the inhibitory powers of the two together are greater than the sum of the parts. This is important because OCP is a precursor to HAP and thus phosphocitrate can also inhibit HAP formation from OCP (Sharma et al., 1992).

Struvite has been shown to be inhibited by both citrate and phosphocitrate (McLean et al., 1990; McLean et al., 1991; Ogawa et al., 2000; Stevenson et al., 2000; Wierzbicki et al., 1997). In vivo, citrus-sourced citrate was shown to lower the pH at which phosphates precipitated from urine, although the internal mechanism remained unclear (Suller et al., 2005). Wierzbicki showed that far more than citrate, phosphocitrate changed the physical crystal structure of struvite and stopped growth all together when the concentration was high enough (Wierzbicki et al., 1997). The struvite developed an arrowhead morphology, and crystals that formed were much smaller. Phosphocitrate has an affinity for the growing faces and because of the negatively charge phosphate group which is needed in the crystal, an ability to incorporate directly into the crystal lattice (McLean et al., 1991; Wierzbicki et al., 1997).

1.2.5.2 Urease Infection stones are caused by urinary tract infections with ureolithic microorganisms (Babic-Ivancic et al., 2004). Human urine is an excellent culture medium; it can support high levels of bacterial growth such as Proteus spp and Pseudomonas spp, which account for about 70% of infections (Brooks and Keevil, 1997). Urine is undersaturated with respect to struvite and thus, a change in urine chemistry must occur to saturate the urine such that precipitation can occur (Babic-Ivancic et al., 2004; Griffith et al., 1976; Robertson et al., 1968b).

Griffith's early work showed that struvite and apatite crystals form when urine becomes supersaturated as a result of urea hydrolysis by bacterial urease (Griffith et al., 1976; Griffith et al., 1976; Griffith, 1979). Urease-positive bacteria were thought to be the sole agents responsible for the hydrolysis of urea and subsequent supersaturation of struvite (Griffith et al., 1976). Later work however, demonstrated that struvite precipitation could also be induced by urease negative bacteria, i.e. bacteria that can generate ammonium ions and

32 increase the pH of the urine by mechanisms other than urease activity (Rivadeneyra et al., 1999).

Reports vary as to what conditions are necessary for struvite stone formation. Early work on struvite showed it would only form in urine at a pH above 7.1 (Elliot et al., 1959) although subsequent work has shown struvite to crystallize at 6.5 (Grases, 1998), 7.4 (Griffith, 1973) between 7.5 and 8, (Grases et al., 1996; Perez-Garcia and Rivadeneyra, 1989), and 8 (Griffith et al., 1976). The discrepancies may be attributable to differences in urine composition, but other variables have been suggested as well. Mucin, a nitrogenous protein in mucous, appears to help struvite nucleation because at a similar pH with the absence of mucin, brushite will be the lone crystal to precipitate (Grases et al., 1996). The type of bacteria may ^ also play a role; the lipopolysaccharide of different species of proteus may enhance or inhibit struvite stone formation depending the ability of nucleus forming cations to bind (Torzewska et al., 2003). •

Although the conditions of struvite stone formation are important to understand, the task of the urologist is to prevent stone formation. Understanding when struvite will dissolve describes the other half of the system and is equally important. Struvite stones have been shown to dissolve in undersaturated urine (far more readily than apatite stones) (Griffith et al., 1976). By lowering the urine pH from 6.5-5.75, the dissolution rate of struvite stones can be increased by 35% (Jacobs et al., 2001). To lower the conditional solubility product and prevent stone formation, it has been suggested to dilute urine by increasing fluid intake (Suller'et al., 2005). Organic inhibitors, whilst quife effective at controlling other types of stones, are not particularly useful at preventing infection stone formation (Elliot et al., 1959; Grases et al., 1996; Griffith et al., 1976). Few data are available to show the relative amounts of calcium and magnesium that will favour struvite formation, although in general, low calcium levels will favour struvite (Elliot et al., 1959) and HAP will form preferentially when magnesium is abnormally low (Grases et al., 1996).

Enzymes are a group of proteins that are responsible for catalyzing specific cellular reactions; spontaneous degradation of urea occurs with a half-life of 3.6 years, but with the

33 aid of the enzyme urease, hydrolysis is 10,000 times faster (Amtul et al, 2004). Urease, which is also known as urea amidohydrolase was, in 1926, the first enzyme to be crystallized (Brown, 1976). Urease catalyzes the hydrolysis of urea to ammonia and carbon dioxide by the simplified reaction:

(18) CO(NH2)2+H20 ^ 2NH3 +C02 Or more properly as the following series of reactions:

(19) CO(NH2)2 + H20 NH3 + NH2CO2H

After carbamate (NH2CO2H) is formed it spontaneously dissociates to carbonic acid to release a second molecule of ammonia

(20) NH2CO2H + H20 -» NH3 +H2CO3 The two molecules of ammonia and carbamate are in equilibrium with their protonated and deprotonated forms and the end result is a pH increase:

+ (21) H2C03^ H +HC03"

+ (22) 2NH3 +2H20 -» 2NH4 + 20H" (Amtul et al., 2002; Mobley et al., 1995)

Urease is significant because of its ability to transform environmental nitrogen, to recycle nitrogenous waste in the rumens of domestic animals, and because of its undesirable consequences in urinary infections (Mobley et al., 1995). Essentially, urease enables organisms to use urea as a nitrogen source (Amtul et al., 2002). Humans however, do not need to use urea as a source of nitrogen and therefore, cannot convert it. We produce about 10 kg of urea a year, but for humans it is a waste product that is best left as urea and should not be converted internally; urease is foreign to the urinary tract of humans and its presence is the sign of infection (Amtul et al., 2002).

Although over 200 species of gram-negative and gram-positive bacteria have been found capable of urease activity, it is the genus Proteus, pathogenic rod bacteria, that is commonly linked to urinary tract infections. There are four species: Proteus vulgaris, P. mirabilis, P. penneri and P. myxofaciens, of which P. mirabilis is the third most common cause of urinary tract infections and the second most common cause of catheter encrustation, although both P. vulgaris and P. penneris are virulent. The ability to hydrolyze urea to ammonia is the

34 hallmark of Proteus, which explains the unfortunate outcome of Proteus infected urinary tracts (Rozalski et al., 1997).

1.2.6 Summary Recycled phosphorus will be the solution to the phosphorus shortage and as a means of binding phosphorus in a safe, useable form, it seems that struvite may be the answer. Struvite from urine is something that nature invented millennia ago; by replicating the same processes in vitro, a potential huge source of phosphorus can be harvested.

1.3 Theme and Objectives There is no doubt that current sanitation services will need to be redesigned to meet increasing population and resource demands in the coming years. Clearly, there are many reasons to separate urine using EcoSan technology, and, as such, there is a potentially large, un-tapped resource of phosphorus-rich urine that can be used for the production of struvite. Struvite technology is developing quickly, but to date, large-scale struvite recovery has only made use of different treatment side streams. Limited research has been conducted using fresh, not stored, urine as a feedstock for struvite production.

This thesis examines the possibility of recovering struvite from stored urine. The research is divided into two parts. The goal of the first part is to understand the behaviour of urine in different storage conditions, and the goal of the second part is to make use of those changes to recover pure struvite from stored urine.

To better understand the behaviour of urine, the specific aims are as follows: • Understand how dilution, temperature, and faecal contamination, affect urine over time (i.e. with storage) • Determine how pH and levels of calcium, magnesium, phosphorus and ammonia change as a function of being stored in different conditions (as above) • Identify how and when spontaneous precipitation occurs and how it affects the composition of urine

35 In pilot and full-scale struvite production settings, the majority of the chemical cost can be attributed to caustic (NaOH); Jaffer et al. estimate that sodium hydroxide accounted for 97% of the chemical cost (Jaffer et al., 2002). Therefore, it was a primary goal of this research to determine if, and how, struvite could be recovered without the use of caustic. The findings from the first part indicated that the pH rise caused by ureolysis would be sufficient to remove the requirement for a caustic addition; the pH of the solution could be naturally raised to an optimal working range for struvite recovery simply by allowing urea to be hydrolyzed. Thus, the first part served as a starting point for research in the second part; stored urine that was shown to have properties conducive to struvite recovery was studied further. The goals of the struvite recovery stage are as follows: • Produce struvite without the addition of a caustic • Determine when the best time to dose urine with magnesium to induce struvite • Maximize the amount of phosphorus that is recovered in the form of struvite • Minimize the amount of phosphorus that is lost from urine before struvite recovery • Maximize the amount of calcium that is removed from solution before struvite recovery to ensure a calcium-free struvite product

Based on published literature that pertains to different aspects of this research, the following hypotheses have been formulated: • The pH increase as a result of ureolysis will raise the solution pH to a point that will be optimal for struvite recovery. • Adding magnesium to fresh urine will cause a variety of compounds to precipitate • By dosing urine after spontaneous precipitation has removed calcium from solution, purer struvite will be recovered; the more calcium removed, the better quality the struvite that will be recovered. • There will be an optimal amount of dilution and faecal contamination that will produce the best quality struvite.

By addressing the research objectives outlined above, this timely research will contribute to an important, and quickly growing, body of knowledge.

36 1.4 References

Aage, H. K., Andersen, B. L., Blom, A., Jensen, I., 1997. The solubility of struvite. J.

Radioanal. Nucl. Chem., 223 (1-2),(213-215.

Abbona, F., Boistelle, R., 1985. Nucleation of struvite (Mgnh4po4.6h2o) single-crystals and aggregates. Cryst. Res. Technol, 20 (2), 133-140.

Abbona, F., Boistelle, R., 1979. Growth morphology and crystal habit of struvite crystals (Mgnh4po4.6h2o). J.Cryst.Growth, 46 (3), 339-354.

Abbona, F., Calleri, M., Ivaldi, G., 1984. Synthetic struvite, Mgnh4po4.6h2o - correct polarity and surface-features of some complementary forms. Acta Crystallographica Section B-Structural Science, 40 (JUN), 223-227.

Abbona, F., Franchiniangela, M., 1990. Crystallization of calcium and magnesium phosphates from solutions of low concentration. J.Cryst.Growth, 104 (3), 661-671.

Abbona, F., Madsen, H. E. L., Boistelle, R., 1988. The final phases of calcium and magnesium phosphates precipitated from solutions of high to medium concentration. J.Cryst.Growth, 89 (4), 592-602.

Abbona, F., Madsen, H. E. L., Boistelle, R., 1982. Crystallization of 2 magnesium phosphates, struvite and newberyite - effect of pH and concentration. J.Cryst.Growth, 57 (1), 6-14.

Abbona, F., Madsen, H. E. L., Boistelle, R., 1986. The initial phases of calcium and magnesium phosphates precipitated from solutions of high to medium concentrations. J.Cryst.Growth, 74 (3), 581-590.

Abbona, F., Baronnet, A., 1996. A XRD and TEM study on the transformation of amorphous calcium phosphate in the presence of magnesium. J.Cryst.Growth, 165 (1-2), 98-105.

Abe, S., 1995. Phosphate removal from dewatering filtrate by MAP process at Seibu treatment plant in Fukuoka City. Sewage Works in Japan, 43, 59-64.

Adnan, A., Dastur, M., Mavinic, D. S., Koch, F. A., 2004. Preliminary investigation into factors affecting controlled struvite crystallization at the bench scale. J. Environ. Eng. Sci., 3 (3), 195-202.

Adnan, A., Koch, F. A., Mavinic, D. S., 2003a. Pilot-scale study of phosphorus recovery through struvite crystallization - II: Applying in-reactor supersaturation ratio as a process control parameter. J. Environ. Eng. Sci., 2 (6), 473-483.

Adnan, A., Mavinic, D. S., Koch, F. A., 2003b. Pilot-scale study of phosphorus recovery through struvite crystallization - examining the process feasibility. J. Environ. Eng. Sci., 2(5), 315-324.

37 Amtul, Z., Atta-ur-Rahman, Siddiqui, R. A., Choudhary, M. I., 2002. Chemistry and mechanism of urease inhibition. Curr.Med.Chem., 9 (14), 1323-1348.

Amtul, Z., Rasheed, M., Choudhary, M. I., Rosanna, S., Khan, K. M., Atta-ur-Rahman, 2004. Kinetics of novel competitive inhibitors of urease enzymes by a focused library of oxadiazoles/thiadiazoles and triazoles. Biochem.Biophys.Res.Commun., 319 (3), 1053- 1063.

Babic-Ivancic, V., Kontrec, J., Kralj, D., Brecevic, L., 2002. Precipitation diagrams of struvite and dissolution kinetics of different struvite morphologies. Croat. Chem. Acta., 75 (1), 89-106.

Babic-Ivancic, V., Kontrec, J., Brecevic, L., 2004. Formation and transformation of struvite and newberyite in aqueous solutions under conditions similar to physiological. Urol. Res., 32 (5), 350- 356.

Balmer, P., 2004. Phosphorus recovery - an overview of potentials and possibilities. Water Sci.Technol., 49 (10), 185-190.

Ban, Z., Dave, G., 2004. Laboratory studies on recovery of N and P from human urine through struvite crystallisation and zeolite adsorption. Environ.Technol., 25 (1), 111- 121,

Battistoni, P., De Angelis, A., Prisciandaro, M., Boccadoro, R., Bolzonella, D., 2002. P removal from anaerobic supernatants by struvite crystallization: Long term validation and process modelling. Water Res., 36 (8), 1927-1938.

Battistoni, P., Fava, G., Pavan, P., Musacco, A., Cecchi, F., 1997. Phosphate removal in anaerobic liquors by struvite crystallization without addition of chemicals: Preliminary results. Water Res., 31 (11), 2925-2929.

Beal, L., Burns, R. T., Stalder, K., 1999. Effect of anaerobic digestion on struvite production for nutrient removal from swine waste prior to land application. Proceedings of the 1999 ASAE Annual International Meeting, Toronto, Canada, 994042

Behrendt, J., Arevalo, E., Gulyas, H., Niederste-Hollenberg, J., Niemiec, A., Zhou, J., Otterpohl, R, 2002. Production of value added products from separately collected urine. Water Sci.Technol., 46 (6-7), 341-346.

Beler-Baykal, B., Bayram, S., Akkaymak, E., Cinar, S., 2004. Removal of ammonium from human urine through ion exchange with clinoptilolite and its recovery for further reuse. Water Sci.Technol., 50 (6), 149-156.

Bellagio Statement: Clean, Healthy and Productive Living: A New Approach to Environmental Sanitation, (2000).

Bengtsson, M., Lundin, M., & Molander, S. (1997). Life cycle assessment of wastewater systems-Case studies of conventional treatment, urine sorting and liquid composting in

38 three Swedish municipalities. Goteborg, Sweden: Technical Environmental Planning, Chalmers University of Technology.

Berndtsson, J. C, 2006. Experiences from the implementation of a urine separation system: Goals, planning, reality. Build.Environ., 41 (4), 427-437.

Boistelle, R., Lopezvalero, I., Abbona, F., 1993. Calcium-phosphate crystallization with magnesium. Nephrologie, 14 (6), 265-269.

Booker, N. A., Priestley, A. J., Fraser, I. H., 1999. Struvite formation in wastewater treatment plants: Opportunities for nutrient recovery. Environ.Technol., 20 (7), 777-782.

Borgerding, J., 1972. Phosphate deposits in digestion systems. J.Wat.PollutControl.Fed., 44 (5), 813-819.

Bouropoulos, N. C, Koutsoukos, P. G., 2000. Spontaneous precipitation of struvite from aqueous solutions. J. Cryst. Growth, 213 (3-4), 381-388.

Bridger, G., Salutsky, M., Starostka, R., 1962. Micronutrient sources, metal ammonium phosphates as fertilizers. J.Agric.Food Chem., 10 (3), 181-188.

Britton, A., Koch, F., Mavinic, D., Adnan, A., Oldham, W., Udala, B., 2005. Pilot-scale struvite recovery from anaerobic digester supernatant at an enhanced biological phosphorus removal wastewater treatment plant. J.Environ.Eng.Sci, 4, 265-277.

Brooks, T., Keevil, C. W., 1997. A simple artificial urine for the growth of urinary pathogens. Lett.Appl.Microbiol., 24 (3), 203-206.

Brown, W. H. (1976). Introduction to organic and biochemistry. Willard Grant Press, Boston.

Bube, K. J., 1910. Uber magnesiumammoniumphosphat. 49 (9), 525-596.

Burns, J. R., Finlayson, B., 1982. Solubility product of magnesium ammonium phosphate hexahydrate at various temperatures. J. Urology, 128 (2), 426-428.

Burns, J. R., Finlayson, B., 1980. A proposal for a standard reference artificial urine in in vitro urolithiasis experiments. Invest.Urol., 18 (2), 167-169.

Dastur, M. (2001). Investigation into the factors affecting controlled struvite crystallization at the bench-scale. (MASc thesis, Department of Civil Engineering, The University of British Columbia). de-Bashan, L. E., Bashan, Y., 2004. Recent advances in removing phosphorus from wastewater and its future use as fertilizer (1997-2003). Water Res., 38 (19), 4222-4246.

Diem, K., & Lentner, C. (Eds.). (1970). Documenta Geigy: Scientific Tables. Basle Geigy.

39 Doyle, J. D., Parsons, S. A., 2002. Struvite formation, control and recovery. Water Res., 36 (16), 3925-3940.

Eidelman, N., Chow, L. C, Brown, W. E., 1987. Calcium phosphate saturation levels in ultrafiltered serum. Calcif.Tissue Int., 40 (2), 71-78.

Elliot, J. S., Sharp, R. F., Lewis, L., 1959. The solubility of struvite in urine. J.Urol., 81 (3), 366-368.

Fittschen, I., Niemczynowicz, J., 1997. Experiences with dry sanitation and greywater treatment in the ecovillage Toarp, Sweden. Water Sci. Technol., 35 (9), 161-170.

Fujimoto, N., Mizuochi, T., Togami, Y., 1991. Phosphorus fixation in the sludge treatment system of a biological phosphorus removal process. Water Sci. Technol., 23 (4), 635- 640.

Garnett, J., Dieppe, P., 1990. The effects of serum and human albumin on calcium hydroxyapatite crystal growth. Biochem.J., 266 (3), 863-868.

Grases, F., Costa-Bauza, A., Garcia-Ferragut, L., 1998. Biopathological crystallization: A general view about the mechanisms of renal stone formation. Adv.Colloid Interface Sci., 74, 169-194.

Grases, F., Sohnel, O., Vilacampa, A. I., March, J. G., 1996. Phosphates precipitating from artificial urine and fine structure of phosphate renal calculi. Clin. Chim. Acta, 244 (1), 45-67.

Grases, F., Villacampa, A. I., Costa-Bauza, A., 1999. Ammonium and sodium urates precipitating from synthetic urine and fine structure of urate renal calculi. Urol.Res., 27 (2), 141-147.

Griffith, D. P., 1979. Urease stones. Urol.Res., 7 (3), 215-221.

Griffith, D. P., Musher, D. M., 1973. Prevention of infected urinary stones by urease inhibition. Invest.Urol., 11 (3), 228-233.

Griffith, D. P., Bragin, S., Musher, D. M., 1976. Dissolution of struvite urinary stones - experimental studies invitro. Invest.Urol., 13 (5), 351-353.

Griffith, D. P., Dunn, D., 1978. Collection and preservation of urine for biochemical analyses. Invest.Urol., 15 (6), 459-461.

Griffith, D. P., Musher, D. M., Itin, C, 1976. Urease - primary cause of infection-induced urinary stones. Invest.Urol., 13 (5), 346-350.

Gulyas, H., Bruhn, P., Furmanska, M., Hartrampf, K., Kot, K., Luttenberg, B., Mahmood, Z., Stelmaszewska, K., Otterpohl, R., 2004. Freeze concentration for enrichment of nutrients in yellow water from no-mix toilets. Water Sci.Technol., 50 (6), 61-68.

40 Hanaeus, A., Hellstrom, D., Johansson, E., 1996. Conversion of urea during storage of human urine. Vatten, 52 , 263-270.

Hanaeus, J., Hellstrom, D., Johansson, E., 1997. A study of a urine separation in an ecological village in northern Sweden. Water Sci. Technol., 35 (9), 153-160.

Hellstrom, D., 2003. Exergy analysis of nutrient recovery processes. Water Sci. Technol., 48 (1), 27-36.

Hellstrom, D., Karrman, E., 1996. Nitrogen and phosphorus in fresh and stored urine. Environ.Res.Forum, 5 (6), 221-226.

Hellstrom, D., Johannson, E., Grennberg, K., 1999. Storage of human urine: Acidification as a method to inhibit decomposition of urea. Ecol.Eng., 12 (3-4), 253-269.

Hellstrom, D., Karrman, E., 1997. Exergy analysis and nutrient flows of various sewerage systems. Water Sci. Technol., 35 (9), 135-144.

Hesse, A., Siener, R., 1997. Current aspects of epidemiology and nutrition in urinary stone disease. World J.Urol., 15 (3), 165-171.

Hirasawa, I., Kaneko, S., Kanai, Y., Hosoya, S., Okuyama, K., Kamahara, T., 2002. Crystallization phenomena of magnesium ammonium phosphate (MAP) in a fluidized- bed-type crystallizer. J.Cryst.Growth, 237, 2183-2187.

Hoglund, C, Stenstrom, T. A., Jonsson, H., Sundin, A., 1998. Evaluation of faecal contamination and microbial die-off in urine separating sewage systems. Water Sci. Technol., 38 (6), 17-25.

Hoglund, C, Vinneras, B., Stenstrom, T. A., Jonsson, H., 2000. Variation of chemical and microbial parameters in collection and storage tanks for source separated human urine. J Environ. Sci. Heal. Part A-Toxic/Hazardous Substances & Environmental Engineering, 35 (8), 1463-1475.

Howard, J., Thomas, W., 1958. Some observations on rachitic rat cartilage of probable significance in the etiology of renal calculi. Trans.Am.Clin.Climatol.Assoc, 70 , 94-102.

Huang, H., Mavinic, D., Lo, K., Koch, F., 2006. Production and basic morphology of struvite crystals from a pilot-scale crystallization process. Environ.Technol., 27 (3), 233-245.

Jacobs, D., Heimbach, D., Hesse, A., 2001. Chemolysis of struvite stones by acidification of artificial urine - an in vitro study. Scand.J.Urol.Nephrol., 35 (5), 345-349.

Jaffer, Y., Clark, T. A., Pearce, P., Parsons, S. A., 2002. Potential phosphorus recovery by struvite formation. Water Res., 36 (7), 1834-1842.

Jenkins, J. (1999). The Humanure Handbook: A guide to composting human manure (3rd edition ed.). USA: Jenkins Publishing.

41 Jernlid, A., Karlsson, K. (1997). Vata toalettsystem-driftstudier och utvardering med livscykelanalys i Sodra Valsangs ekoby. Examensarbete 1997:6. Goteborg.: VA-teknik, Chalmers tekniska hogskola..

Jonsson, H., 2002. Urine separating sewage systems - environmental effects and resource usage. Water Sci.Technol., 46 (6-7), 333-340.

Jonsson, H., Stenstrom, T. A., Svensson, J., Sundin, A., 1997. Source separated urine- nutrient and heavy metal content, water saving and faecal contamination. Water Sci. Technol., 35 (9), 145-152.

Katsuura, H. (1998). Phosphate recovery from sewage by granule forming process (full scale struvite recovery from a sewage works at shimane prefecture, japan). International Conference on Phosphate Recovery from Sewage and Animal Waste, UK. Warwick University.

Klawitter, S., Qazzaz, H., 2005. Water as a human right: The understanding of water in the arab countries of the middle east. Int. J. Water Resour.Dev., 21 (2), 253-271.

Krauskopf, K. B. (1967). Introduction to geochemistry. New York, NY: McGraw-Hill.

Larsen, T. A., Gujer, W., 1996. Separate management of anthropogenic nutrient solutions (human urine). Water Sci. Technol., 34 (3-4), 87-94.

Larsen, T. A., Lienert, J., Joss, A., Siegrist, H., 2004. How to avoid pharmaceuticals in the aquatic environment. J.Biotechnol., 113 (1-3), 295-304.

Larsen, T. A., Peters, I., Alder, A., Eggen, R., Maurer, M., Muncke, J., 2001. Re-engineering the toilet for sustainable wastewater management. Environ.Sci.Technol., 35 (9), 192A- 197 A.

Lienert, J., Haller, M., Berner, A., Stauffacher, M., Larsen, T. A., 2003. How farmers in Switzerland perceive fertilizers from recycled anthropogenic nutrients (urine). Water Sci. Technol., 48 (1), 47-56.

Lind, B. B., Ban, Z., Byden, S., 2001. Volume reduction and concentration of nutrients in human urine. Ecol.Eng., 16 (4), 561-566.

Lind, B. B., Ban, Z., Byden, S., 2000. Nutrient recovery from human urine by struvite crystallization with ammonia adsorption on zeolite and wollastonite. Bioresour.Technol., 73 (2), 169-174.

Lunt, O., Kofranek, A., Clark, S., 1964. Nutrient availability in soil, availability of minerals from magnesium ammonium phosphates. J.Agric.Food Chem., 12 (6), 497-504.

Maekawa, T., Chung-Min, L., Xing-Dong, F., 1995. Nitrogen and phosphorus removal for swine wastewater using intermittent aeration batch reactor followed by ammonium crystallization process. Water Res., 29 (12), 2643-2650.

42 Mashauri, D., & Senzia, M. (2002). Reuse of nutrients from ecological sanitation toilets as a source of fertiliser. The Proceedings of the 3rdlNternational Conference on Integrated Environmental Management in Southern Africa, Johannesburg, 27-30.

Maurer, M., Schwegler, P., Larsen, T. A., 2003. Nutrients in urine: Energetic aspects of removal and recovery. Water Sci. Technol., 48 (1), 37-46.

McLean, R. J., Downey, J., Clapham, L., Nickel, J. C, 1990. Influence of chondroitin sulfate, heparin sulfate, and citrate on proteus mirabilis-induced struvite crystallization in vitro. J.Urol., 144 (5), 1267-1271.

McLean, R. J., Downey, J., Clapham, L., Wilson, J. W., Nickel, J. C, 1991. Pyrophosphate inhibition of proteus mirabilis-induced struvite crystallization in vitro. Clin.Chim. Acta, 200(2-3), 107-117. '

Mitani, Y., Sakai, Y., Mishina, F., Ishiduka, S., 2003. Struvite recovery from wastewater having low phosphate concentration. J. Wat. Env. Tech., 1 (1), 13-18.

Mobley, H. L., Island, M. D., Hausinger, R. P., 1995. Molecular biology of microbial ureases. Microbiol.Rev., 59 (3), 451-480.

Momberg, G. A., Oellermann, R. A., 1992. The removal of phosphate by hydroxyapatite and struvite crystallization in south africa. Water Sci. Technol., 26 (5-6), 987-996.

Morris, N. S., Stickler, D. J., 1998. The effect of urease inhibitors on the encrustation of urethral catheters. Urol.Res., 26 (4), 275-279.

Munch, E. V., Barr, K., 2001. Controlled struvite crystallisation for removing phosphorus from anaerobic digester sidestreams. Water Res., 35 (1), 151-159.

Nelson, N. O., Mikkelsen, R. L., Hesterberg, D. L., 2003. Struvite precipitation in anaerobic swine lagoon liquid: Effect of pH and mg:P ratio and determination of rate constant. Bioresour.Technol., 89 (3), 229-236.

Ogawa, Y., Sugaya, K., Koyama, Y., Hatano, T., 2000. Impact of citrate therapy on the circadian rhythm of urinary magnesium ammonium phosphate saturation in normal individuals. Int. J. Urol., 7 (8), 287-291.

Ohlinger, K. N., Young, T. M., Schroeder, E. D., 1999. Kinetics effects on preferential struvite accumulation in wastewater. J. Environ. Eng. Sci., 125 (8), 730-737.

Ohlinger, K., Young, T., Schroeder, E., 1998. Predicting struvite formation in digestion. Water Res., 32 (12), 3607-3614.

Otterpohl, R. (2000). Design of highly efficient source control sanitation and practical experiences. EURO-Summer School DESAR, Wageningen, Netherlands.

43 Otterpohl, R., Grottker, M., Lange, J., 1997. Sustainable water and waste management in urban areas. Water Sci. Technol., 35 (9), 121-133.

Pahl-Wostl, C., Schonborn, A., Willi, N., Muncke, J., Larsen, T. A., 2003. Investigating consumer attitudes towards the new technology of urine separation. Water Sci. Technol., 48 (1), 57-65.

Peasey, A. (2000). Health aspects of dry sanitation with waste reuse. Task No. 324. Water and Environmental Health at London and Loughborough.

Perez-Garcia, I., Rivadeneyra, M., 1989. Influence of pH on struvite formation by bacteria. Chemosphere, 18 (7-8), 1633-1638.

Pitzer, K. (1979). Theory: Ion interaction approach. Activity in electrolyte solutions (pp. 209- 265). Boca Raton: CRC Press.

Pronk, W., Biebow, M., Boiler, M., 2006. Electrodialysis for recovering salts from a urine solution containing micropollutants. EnvironSci.Technol., 40 (7), 2414-2420.

Randall, C. W., 2003. Potential societal and economic impacts of wastewater nutrient removal and recycling. Water Sci. Technol., 48 (1), 11-17.

Rauch, W., Brockmann, D., Peters, I., Larsen, T. A., Gujer, W., 2003. Combining urine, separation with waste design: An analysis using a stochastic model for urine production. Water Res., 37(3), 681-689.

Reddi, A. H., Meyer, J. L., Tew, W. P., Howard, J. E., Lehninger, A. L., 1980. Influence of phosphocitrate, a potent inhibitor of hydroxyapatite crystal growth, on mineralization of cartilage and bone. Biochem.Biophys.Res.Commun., 97 (1), 154-159.

Rivadeneyra, M. A., Gutierrez-Calderon, A., Rivadeneyra, A. M., Ramos-Cormenzana, A., 1999. A study of struvite precipitation and urease activity in bacteria isolated from patients with urinary infections and their possible involvement in the formation of renal calculi. Urol. Int., 63 (3), 188-192.

Robertson, W. G., Peacock, M., Nordin, B. E. C, 1968a. Activity products in stone-forming and non-stone-forming urine. Clin.Sci., 34 (3), 579-594

Robertson, W. G., Peacock, M., Nordin, B. E., 1968b. Activity products in stone-forming and non-stone-forming urine. Clin.Sci., 34 (3), 579-594.

Ronteltap, M., Biebow, M., Maurer, M., & Gujer, W. (2006). Thermodynamics of struvite precipitation in source separated urine. Unpublished manuscript.

Rozalski, A., Sidorczyk, Z., Kotelko, K., 1997. Potential virulence factors of proteus bacilli. Microbiol. Mol. Biol. R., 61 (1), 65-89.

44 Salimi, M., Heughebaert, J., Nancollas, G., 1985. Crystal growth of calcium phosphates in the presence of magnesium ions. Langmuir, 1 (1), 119-122.

Sharma, V., Johnsson, M., Sallis, J., Nancollas, G., 1992. Influence of citrate and phosphocitrate on the crystallization of octacalcium phosphate. Langmuir, 8 (2), 676- 679.

Shimamura, K., Tanaka, T., Miura, Y., Ishikawa, H., 2003. Development of a high-efficiency phosphorus recovery method using a fluidized-bed crystallized phosphorus removal system. Water Sci. Technol., 48 (1), 163-170.

Smith, S. E., Marin, L. E., 2005. Water and the rural poor in Latin America: The case of Tlamacazapa, Guerrero, Mexico. Hydrogeol.J., 13 (1), 346-349.

Snoeyink, V. L., & Jenkins, D. (1980). Water chemistry. New York, NY: John Wiley and Sons, Inc.

Sohnel, O., Grases, F., 1995. Calcium oxalate monohydrate renal calculi, formation and development mechanism. Adv.Colloid Interface Sci., 59, 1-17.

St. Joseph's Hospital. (2006). Metabolic stone clinic. Retrieved June/23, 2006 from http://www.sjhc.london.on.ca/sjh/programs/urology/stone.htm

Steen, I., 1998. Phosphorus availability in the 21st century. September-October, 25-31.

Stevenson, A. E., Wrigglesworth, D. J., Smith, B. H., Markwell, P. J., 2000. Effects of dietary potassium citrate supplementation on urine pH and urinary relative supersaturation of calcium oxalate and struvite in healthy dogs. Am.J.Vet.Res., 61 (4), 430-435.

Stratful, I., Scrimshaw, M. D., Lester, J. N., 2001. Conditions influencing the precipitation of magnesium ammonium phosphate. Water Res., 35 (17), 4191-4199.

Stumm, W., Morgan, J. J. (1981). Aquatic chemistry: An introduction emphasizing chemical equilibria in natural waters. New York: Wiley.

Suller, M. T., Anthony, V. J., Mathur, S., Feneley, R. C, Greenman, J., Stickler, D. J., 2005. Factors modulating the pH at which calcium and magnesium phosphates precipitate from human urine. Urol. Res., 33 (4), 254-260.

Suzuki, K., Tanaka, Y., Osada, T., Waki, M., 2002. Removal of phosphate^ magnesium and calcium from swine wastewater through crystallization enhanced by aeration. Water Res., 36 (12), 2991-2998.

Tew, W. P., Mahle, C, Benavides, J., Howard, J. E., Lehninger, A. L., 1980. Synthesis and characterization of phosphocitric acid, a potent inhibitor of hydroxylapatite crystal growth. Biochemistry, 19 (9), 1983-1988.

45 Tillman, A. M., Lundstrbm, H., & Svingby, M. (1996). Livscykelanalys av alternativa avloppssystem i bergsjon och hamburgsund. Delrapport fran ECO-GUIDE- projektet.Databilaga.

Torzewska, A., Staczek, P., Rozalski, A., 2003. Crystallization of urine mineral components may depend on the chemical nature of proteus endotoxin polysaccharides. J.Med.Microbiol., 52 (6), 471-477.

Udert, K. M., Larsen, T. A., Biebow, M., Gujer, W., 2003a. Urea hydrolysis and precipitation dynamics in a urine-collecting system. Water Res., 37 (11), 2571-2582.

Udert, K. M., Larsen, T. A., Gujer, W., 2003. Biologically induced precipitation in urine- collecting systems. Water Sci. Technol.: Water Supply, 3 (3), 71-78.

Udert, K. M., Larsen, T. A., Gujer, W., 2003c. Estimating the precipitation potential in urine- collecting systems. Water Res., 37 (11), 2667-2677.

UNESCO. (2006). World water assessment programme. Retrieved June 26, 2006 from http.7/www.unesco.org/water/wwap/facts_figures/mdgs.shtml

Vinneras, B., 2004. Separation efficiency in a whirlpool surface tension separator, separating faeces and toilet paper for nutrient recovery—pilot-scale study. Water Sci.Technol., 50 (6), 115-121.

Vinneras, B., Bjorklund, A., Jonsson, H., 2003a. Thermal composting of faecal matter as treatment and possible disinfection method— laboratory-scale and pilot-scale studies. Bioresour.Technol., 88 (1), 47-54.

Vinneras, B., Holmqvist, A., Bagge, E., Albihn, A., Jonsson, H., 2003b. The potential for ' disinfection of separated faecal matter by urea and by peracetic acid for hygienic nutrient recycling. Bioresour.Technol., 89(2), 155-161.

Wierzbicki, A., Sallis, J. D., Stevens, E. D., Smith, M., Sikes, C. S., 1997. Crystal growth and molecular modeling studies of inhibition of struvite by phosphocitrate. Calcif.Tissue Int., 61 (3), 216-222.

Wild, D., Kisliakova, A., Siegrist, H., 1996. P-fixation by mg, ca and zeolite A during stabilization of excess sludge from enhanced biological P-removal. Water Sci. Technol., 34 (1), 391-398.

Williams, G., Sallis, J. D., 1982. Structural factors influencing the ability of compounds to inhibit hydroxyapatite formation. Calcif.Tissue Int., 34 (2), 169-177.

Williams, G., Sallis, J. D., 1979. Structure-activity relationship of inhibitors of hydroxyapatite formation. Biochem.J., 184 (1), 181-184.

Wilsenach, J. A., Maurer, M., Larsen, T. A., van Loosdrecht, M. C. M., 2003. From waste treatment to integrated resource management. Water Sci. Technol., 48 (1), 1-9.

46 Wilsenach, J. A., van Loosdrecht, M. C. M, 2006. Integration of processes to treat wastewater and source-separated urine. J. Environ. Eng. Sci., 132 (3), 331-341

Wilsenach, J., van Loosdrecht, M., 2003. Impact of separate urine collection on wastewater treatment systems. Water Sci.Technol., 48 (1), 103-110.

Wilsenach, J. A., Van Loosdrecht, M. C. M., 2004. Effects of separate urine collection on advanced nutrient removal processes. Environ.Sci.Technol., 38 (4), 1208-1215.

Wu, Q. Z., Bishop, P. L., 2004. Enhancing struvite crystallization from anaerobic supernatant. J. Environ. Eng. Sci., 3 (1), 21-29.

Yoshino, M., Yao, M., Tsuno, H., Somiya, I., 2003. Removal and recovery of phosphate and ammonium as struvite from supernatant in anaerobic digestion. Water Sci. Technol., 48 (1), 171-178.

47 CHAPTER 2: Preparation of urine for optimal struvite recovery 1

2.1 Introduction Phosphorus, one of the three essential nutrients for plant growth (along with nitrogen and potassium), is a finite resource that is mined from natural deposits^deposits that are quickly disappearing. It has been projected that half of the currently economically recoverable phosphate resources will be exhausted in 60-70 years (Steen, 1998). If global agriculture is to meet the ever-growing demands of an exploding population, new sources of phosphorus must quickly be found.

To address the coming shortage, innovative technologies have been developed to recover phosphorus from the human sewage—the same sewage that is often responsible for the eutrophication of water bodies (Randall, 2003). By raising the pH with a caustic and adding a source of magnesium, phosphorus can be precipitated in the form of struvite.

Struvite (MgNF^PCM H20) is a bioavailable phosphate fertilizer that has been recovered from wastewater and sold in Japan since the 1990s (Ueno and Fujii, 2001). A similar technology developed at UBC crystallizes small pellets of garden-ready struvite, but to date, struvite has only been harvested from dilute waste streams.

Urine, which accounts for only 1% of the volume of wastewater (Berndtsson, 2006), actually contains over 70% of the N and over 90% of the P excreted by humans; this would make it an ideal feedstock for struvite recovery (Hanaeus et al., 1997). Humans have been making struvite from urine for millennia; unfortunately, however, in the form of renal calculi, or kidney stones, that form as a result of urinary tract infections (Babic- Ivancic et al., 2004; Brooks and Keevil, 1997; Hesse and Siener, 1997). Struvite crystals form when urine becomes supersaturated as a result of urea hydrolysis by bacterial urease (Griffith et al., 1976a; Griffith et al., 1976b; Griffith, 1979). The enzyme urease catalyzes the hydrolysis of urea to ammonia and carbon dioxide by the simplified reaction:

urease CO(NH2)2+H20 • 2NH3 +C02

1 A version of this chapter has been/will be submitted for publication.

48 which results in increased solution saturation that can lead to struvite precipitation (Amtul et al., 2002; Mobley et al., 1995). Instead of biologically induced changes, struvite is recovered from the newly developed crystallizers with the use of an alkaline additive, such as NaOH (Adnan et al., 2003; Britton et al., 2005).

Urine is generally considered to be sterile; however, faecal material can easily contaminate the urine collection system with urease, causing a rapid increase in pH and ammonia (Hoglund et al., 1998; Jonsson et al, 1997). The supersaturation conditions that are induced then cause the spontaneous precipitation of calcium and magnesium phosphates. Operationally, accumulated minerals can clog the components of separating toilets and impair performance (Udert et al., 2003a; Udert et al., 2003b). This results in the removal of phosphorus from solution, which is then unavailable for struvite recovery.

The idea to recover struvite from urine was proposed at least as early as 1996 (Larsen and Gujer, 1996), but only now is the technology catching up. Toilets that allow for the collection of urine separately from faecal matter are gaining acceptance, as is the idea that urine could be used a fertilizer (Berndtsson, 2006; Lienert et al., 2003; Pahl-Wostl et al., 2003).

Some work has addressed the influence that different storage conditions (e.g. temperature, faecal contamination) have on the dynamics of urine, but none have examined the storage of urine in the context of struvite recovery, i.e. how do different storage conditions affect the quality of struvite that can be recovered (Hanaeus et al., 1996; Hellstrom et al., 1999; Hellstrom and Karrman, 1996). As well, some researchers have recovered struvite from fresh urine, but for most separating and collection systems, fresh urine would be difficult, if not impossible to obtain (Ban and Dave, 2004; Lind et al., 2000).

The goal of this work was to better understand the potential of urine to be used as a feedstock for struvite recovery. Different storage conditions and magnesium dosing regimes were evaluated as ways to optimally prepared urine solutions, such that a maximum amount of pure struvite could be recovered. To do this, three specific studies were conducted to more

49 closely examine different facets of urine preparation and storage. The goal of Part A was to verify the effects of temperature, contamination, dilution, and headspace on stored urine; in Part B, the effect of adding magnesium to fresh urine before ureolysis, was determined, and in Part C, the effect of variable magnesium additions on urine that had undergone ureolysis was determined. An important element of this work was the attention paid to calcium. Calcium is known to be one of the main impurities in struvite (Huang et al., 2006), and its presence can inhibit the formation of magnesium phosphates (Abbona et al., 1986). Thus, by working to maximize phosphorus and minimize calcium in urine solutions, the best conditions for struvite formation can be produced

2.2 Materials and Methods All experiments were conducted using synthetic urine as per the method given by Griffith (Griffith et al., 1976b). Table 2.1 lists the quantities and constituents used for synthetic urine.

Table 2.1. Composition of synthetic urine (Griffith et al, 1976b) Species Concentration (g/L) Concentration (mmol/L)

CaCl2-H20 0.65 Ca: 4.3

MgCl2-6H20 0.651 Mg: 3.2 NaCl 4.6

Na2S04 2.3 S04: 16

Na3C6H807*2H20 (sodium 0.65 Citrate: 2.3 citrate dihydrate)

Na2C204 (sodium oxalate) 0.020 Oxalate: 0.149

KH2P04 (potassium phosphate 2.8 P04: 20.5 monobasic) KC1 1.6

NH4C1 1 NH4: 19

CO(NH2)2 (urea) 25

C4H7N30 (Creatinine) 1.1 Total Na=l 18 mEq Total K= 42 mEq pH= 5.8

Synthetic urine is commonly used in urological research when experimental constraints do not permit the use of real urine. Because these investigations were preliminary, synthetic urine was used since the collection of urine from human subjects is both arduous and time consuming, as well as being subject to ethical and safety review processes.

50 2.2.1 Part A The goal of Part A was to determine the behaviour and composition changes that occur under different storage conditions. To examine the effect of temperature, headspace, dilution, and faecal contamination, a 32- solution matrix was used. Temperature affects biological activity (i.e. urease production) and the rate of ureolysis (Moyo, 1989). The amount of headspace influences the partial pressure of ammonia and may alter the amount that volatilizes (Hanaeus et al., 1996). Dilution is known to affect the final composition of stored urine (Udert et al., 2003a) and highly concentrated solutions are sometimes slower to undergo reactions due to the shielding effect of ions. Urine is easily contaminated by faecal matter during collection; here, wastewater serves as a proxy for faecal matter (Hoglund et al., 1998). Two levels of temperature (4°C and 23°C), two levels of dilution (full and half strength), two types of headspace systems (open jar and closed jar) and four levels of faecal contamination (0%, 5%, 10% and 25% wastewater (WW) by volume) were used. The matrix of urine solutions is shown in Table 2.2.

Table 2.2. Matrix of solutions in Part A Solution # Solution # Nominal Open/Closed WW (mL) (4°C) (23°C) Dilution 1 17 0 open 0 2 - 18 0 open 50 3 19 0 open 100 4 20 0 open 250 5 21 0 closed 0 6 22 0 closed 50 7 23 0 closed 100 8 24 0 closed 250 9 25 1/2 open 0 10 26 1/2 open 50 11 27 1/2 open 100 12 28 1/2 open .. 250 13 29 1/2 closed 0 . 14 30 1/2 closed 50 15 31 1/2 closed 100 16 32 1/2 closed 250

Thirty-two different solutions are shown in Table 2.2; solutions #1-16 were stored at 4°C and solutions #17-32 were stored at 23°C. Both sets of sixteen were prepared in the exact same

51 manner. Each solution was prepared in a 1-L Nalgene, wide-mouth jar with a screw-top lid. The jars were soaked in bleach and rinsed with distilled water to remove any biological contaminants. Half of the solutions were used at full-strength (0 dilution), and half of the solutions were diluted; distilled water was used to dilute the urine by 50%. Solutions that were "closed" were stored with lids on, except during sampling, and "open" solutions were left open for the duration of the experiment. Primary effluent, (or "wastewater") from the UBC pilot plant was used as a source of faecal contamination. Different amounts of wastewater (0, 50, 100, or 250 mL) of the total volume (IL) were replaced with wastewater, such that each solution had the same final volume. Every attempt was made to keep the wastewater-free solutions sterile (i.e. 0%), but since urease is ubiquitous, especially in an Environmental lab, and especially for those solutions that were stored open, contamination was inevitable.

Unless stated otherwise, the following methods were used. The pH of all solutions was measured with a pH probe calibrated with standard solutions (Fisherbrand) and the probe was rinsed with distilled water between samples. For chemical analyses, approximately 5 mL of sample was withdrawn with a syringe and filtered into sample vials with a 0.45 mm

Millipore filter. Samples were preserved to a pH of 2 with 5% H2SO4 and stored at 4°C until they could be analyzed. Ammonia and phosphate were measured with a Lachat QuikChem 8000 flow injection instrument. Calcium, magnesium, aluminium, iron and potassium samples were prepared with the appropriate matrix modifiers as per the instrument manual; analyses were made using a Varian Inc. SpectrAA220 Fast Sequential Atomic Absorption Spectrophotometer. Atomic absorption was used to measure all of the elements except for potassium, which was measured with atomic emission. A Bruker D8 Advance powder X-ray diffractometer, equipped with copper radiation and a graphite monochromator, was used to identify the crystal structure. Crystal constituents were measured by dissolving 100 mg of struvite crystals in 50 mL 0.5% nitric acid and analyzing the solution using atomic absorption, as described above.

2.2.2 Part B \ Eight of the solutions from Part A were replicated for this part of the study. The aim was determine the effect of magnesium on fresh urine that had not been hydrolyzed and to see if

52 struvite could be formed as the urine underwent ureolysis. All eight solutions examined were stored at room temperature with lids on. Solutions #1-4 were full strength and solutions #5-8 were diluted, so the effect of dilution could be examined (the eight solutions examined were the same as of #21-24 (full strength), and #29-32 (half strength) from Part A). Table 2.3 is a list of solutions used in Part B.

Table 2.3. Matrix of solutions in Part B Nominal Solution # (23°C) Open/Closed WW (mL) Dilution 1 0 closed 0 2 0 closed 50 3 0 closed 100 4 0 closed 250 5 1/2 closed 0 6 1/2 closed 50 7 1/2 closed 100 8 1/2 closed 250

Each of the 8 solutions in the set was spiked with 1000 mg of magnesium. Samples were taken just before magnesium was added, and approximately 30 minutes after the magnesium had been mixed and allowed to react. Samples were then taken daily over the course of the storage time as above.

2.2.3 Parte Eighteen solutions were examined in this part of the study. The aim was to determine the effect of magnesium on urine solutions that had been hydrolyzed, and to see if struvite could be formed in urine after ureolysis had set in. Urine that had been hydrolyzed was defined as having an elevated pH (>8) and having surrendered nutrients and metals in the form of insoluble minerals (spontaneous precipitation). In this study, only three levels of contamination were investigated (0,100, and 250 mL wastewater) but an extra level of dilution was investigated (full strength, 14 strength, lA strength). Table 2.4 shows the solutions used in Part C.

53 Table 2.4. Matrix of solutions in Part C Solution # Solution # Nominal Open/Closed WW (mL)

(1:1 Mg:P04) (2:lMg:P04) Dilution 1 10 0 closed 0 2 11 0 closed 100 3 12 0 closed 250 4 13 1/2 closed 0 5 14 1/2 closed 100 6 15 1/2 closed 250 7 16 1/4 closed 0 8 17 1/4 closed 100 9 18 1/4 closed 250

All eighteen solutions were stored at room temperature with lids on. Solutions # 1-9 were identical to solutions #10-18 so that a comparison could be made between two different amounts of magnesium. From Part A, it was known that a precipitate formed when the solution pH passed 8. When a precipitate formed, the solution was filtered using a Whatman 934-AH (1.5mm) glass filter. The precipitate was saved and the filtrate was re-sampled for all parameters. The remaining filtrate (500 mL) was dosed with a concentrated magnesium solution to achieve either a 1:1 P04:Mg ratio (solutions #1-9) or 1:2 (solutions #10-18). Magnesium additions were based on the assumption that the PO4-P concentration in full strength urine was 650 mg/L and that dilutions were perfectly linear.

After magnesium was added, the mixture was shaken vigorously for 1 minute and let to react for 1 hour, at which point all the solution was re-sampled for all parameters. The precipitate that formed was filtered out and the remaining filtrate sampled for a final time.

Because the precipitates were of interest, the wastewater was filtered with a FisherBrand P8 large pore filter, to remove large particulate matter in an attempt to minimize particulate matter that attached to precipitates. Each solution (#1-18) was monitored daily for changes in pH and sampled for ammonia, phosphate, magnesium and calcium.

54 2.3 Results and Discussion

2.3.1 Part A

2.3.1.1 Temperature The most pronounced difference between the different urine solutions was a result of storage temperature. All urine solutions stored at room temperature reached a maximum pH (about 9.3) within 19 days, and depending on the amount of wastewater present, many achieved a maximum pH in 4 days. No urine solution stored at 4 degrees reached a maximum pH above 9 before 9 days; half of the mixtures were stable for 23 days and after 30 days of monitoring, 5 of the 16 mixtures were still hovering near the initial pH. The inhibitory effects of temperature are shown in Figure 2.1.

To completely assess the impact of temperature on ureolysis in urine, more than a comparison of pH values is needed. Ammonia is a product of ureolysis and is a good indicator of urease action. Ammonia will raise the pH to a set point (that corresponds to the pK for a given temperature) but as urease continues to cleave urea, the pH will remain stable while ammonia will continue to increase.

It is clear, when comparing Figures 2.1a and 2.1b, that storage at 23°C significantly increases the rate at which pH increases. From Figure lb, it can be seen that all solutions uniformly reached a maximum pH of about 9.3, whereas the solutions stored at 4°C, did not uniformly reach a maximum pH; the values were more stratified. As well, some solutions reached higher peak values (i.e. 9.5) than the corresponding solution that was stored at 23°C. This discrepancy can be explained by the higher pK values for ammonia at lower temperatures, as summarized in Table 2.5 (Perrin, 1982).

Table 2.5. Dissociation constants for ammonia pK Temperature (°C) 10.081 0 9.903 5 9.730 10 9.564 15 9.401 20 9.246 25

Thus, as the temperature decreases, the pH at which equilibrium between ammonia and ammonium is attained, increases. At 23°C, the pK is approximately 9.3, which corresponds to the pH observed at equilibrium, whereas the solutions stored at 4°C should have reached a maximum pH near 9.9. However, since the temperature and headspace were not held constant during sampling, the system did not reach the theoretical limit.

Figure 2.2 shows the dramatic effect that temperature has on the ammonia concentration (measured in mg NH3-N/L) in the urine mixtures over the course of 30 days.

57 12500 a)

10000

i Sample #

« 7500 -*-4 -m~s

V - 6 -4—7 8 9 -•-10 -*-12 -*-13 -*-14 15 • 16

Day

12500

Figure 2.2. Changes in ammonia in solutions stored at a) 4°C and b) 23°C Every solution stored at room temperature closely approached, or exceeded 5000 mg NH3- N/L, yet not a single solution stored at 4°C reached 5000 mg/L. The stratified values seen in Figure 2.2b are due to the effects of dilution with distilled water and wastewater. Every line plotted below 7500 mg/L represents a dilute solution, and the lines above this point represent solutions that are full strength urine. Stratification within in these groupings is due to the replacement of a volume (0-25%) with wastewater.

The relation between pH and ammonia is further demonstrated in Figure 2.3

14000

12000

X 10000 m + • X Z 8000 Ol E S 6000 c o £ E < 4000

2000 — . • »t>

A • 5.5 6.5 7.5 8.5 9 9.5 10 pH

Figure 2.3. Changes in pH vs. ammonia concentration

When pH is plotted against ammonia concentration for all 32 solutions for all sample points, a backwards "L" shape is formed. The plot highlights several interesting features: ammonia levels do not significantly increase (above 2000 mg/L) until after a pH of 9 is reached, pH never exceeds 9.5 regardless of ammonia concentration, and the gap between pH 7 and 9 indicates that the rise in pH between these two points is rapid—so much so, that in most cases, no data could be captured with the sampling strategy. Thus, the potential for ammonia generation in urine that reaches a pH of 9 or higher is great.

59 2.3.1.2 Headspace In a closed system, ammonia and ammonium will stay in equilibrium, but ammonia is volatile and in an open system, it will volatilize; this will result in the overall loss of ammoniacal nitrogen from the system. To determine what effect the loss of ammonia had on the dissolved ammonia concentrations, solutions were either stored with a lid (closed) or no lid (open); however, it should be noted that lids were removed approximately once every three days, for about 10 minutes, for sampling. To calculate losses, the difference in measured ammoniacal nitrogen concentrations for each solution pair was calculated. Solutions stored at room temperature reached steady-state ammonia levels and so the average of the last three measurements was used to calculate the difference between initial and final levels. No comparison can be made for solutions stored at 4°C, because the temperature of the headspace increased during measurements, thus creating a vacuum and pulling ammonia out of solution. The net ammonia loss between solution pairs stored at 23°C is shown in Table 2.5. For each set of conditions there are 4 solution pairs—that is four pairs of solutions, each with closed and open solution (e.g. solution #1-5, 2-6, etc).

Table 2.5. Net loss of ammoniacal nitrogen after 30 days (mg NH3-N/L). %ww Full strength Dilute (50%) 0 1150* 510 5 1403 983 10 1143 740 25 1360 877 * based on the average of 2 highest values

Based on the values presented in Table 2.5, it is obvious that the full strength urine was more prone to ammonia loss than the dilute urine. For comparison, the average percentage loss is presented in Table 2.6. The percent loss was calculated for both the solutions in the solution pair (since each had a different maximum value), and the two percentages were averaged. As above, the maximum value was taken as the average of the last three values.

60 Table 2.6. Percentage (%) loss of ammoniacal nitrogen % WW Full strength Dilute (50%) 0 8* 9 5 13 19 10 11 14 25 16 21 * based on the average of 2, rather than 3, highest values

Compared to the raw losses shown in Table 2.5, the normalized data shown in Table 2.6 allow for a more accurate comparison of ammonia losses as a percentage of the total generated. Data from Table 2.5 indicate that full strength urine lost, as net mass, more nitrogen than dilute urine, but these data indicate that as a percentage, dilute urine is more prone to nitrogen loss. There does not appear to be a linear correlation with the amount of wastewater in solution, although generally, there is a greater loss with an increased amount of wastewater

2.3.1.3 Dilution The effect of dilution is best demonstrated by examining the pH changes of the solution; ammonia changes are not appropriate since ammonia generation is a function of initial urea concentration, which is obviously related to dilution. Compared to full strength solutions, dilute solutions reacted, and reached a steady state pH, more quickly. Figure 2.4 shows both the dilute and full strength solutions that were kept at room temperature. Solutions kept at 4 degrees were generally non-reactive and are not useful for comparison.

61 10

5 I 1 1 1 : 1 1 1 0 5 10 15 20 25 30 35 Day

Figure 2.4. Effects of dilution on pH for solutions at room temperature. Solid lines=dilute solutions, dashed lines=full strength solutions.

Most full strength solutions began to increase in pH later than dilute solutions and reached a pH value of at least 9, after the corresponding dilute solution. The difference was more pronounced in the solutions that had not been dosed with wastewater. The two sets of lines to the right of the graph represent the "sterile" solutions, i.e. the solutions to which no wastewater had been intentionally added, but to which urease is eventually introduced from the environment. Because the dilute solutions reacted more quickly, it appears that less contamination is needed for ureolysis. Since the solutions were not mixed, diffusion was the only means by which urease could come in contact with urea in order to hydrolyze it. Full strength solutions are quite saline and the concentration of solutes would be such that diffusion may be inhibited; in dilute solutions, diffusion would occur more readily and thus, a small amount of urease would more easily diffuse throughout the solution, and therefore, have more of an effect.

62 2.3.1.4 Faecal Contamination Urease is ubiquitous in nature, so although measures were taken to reduce cross contamination, it was difficult to prevent foreign urease from contaminating "sterile" solutions. Despite the fact that contamination did occur, the solutions were not contaminated for about the first two weeks of storage. The effect of varying quantities of wastewater is shown in Figure 2.5. For maximum clarity, only a subset of data (full strength solutions stored at 23°C) is shown. Figure 2.5 shows the pH variation with time for the 8 solutions.

to i 1

9.5

5.5

5 1. , , , 0 5 10 15 20 25 30 35

Day

Figure 2.5. Effect of wastewater on pH. Solid black=0%WW (sterile), dashed grey=5%WW, solid grey=10%, dashed black=25% WW.

Figure 2.5 shows that the onset of urea conversion is related to the amount of wastewater, and therefore urease, present in solution: solutions with 25%WW react almost immediately, while solutions with only 5% are delayed several days for complete conversion. To highlight the trend, only one quarter (8/32 solutions) of the data is presented in Figure 2.5, although over the entire data set, the same trends can be observed. Within a given level of contamination, e.g. 5%, the data is stratified such that diluted solutions react before full strength solutions (with the same degree of contamination), but after solutions with higher

63 degrees of concentration. There is little difference between any of the solutions contaminated by 25%; urea hydrolysis is nearly instantaneous and no clear pattern can be discerned.

2.3.1.5 Precipitate At various points during the 30 days of monitoring, some solutions precipitated a white/beige solid that accumulated in the bottom of the jars. After 30 days, the precipitate was filtered onto Whatman 934-AH filter and dried. The dried solid was then analyzed using x-ray powder crystal diffraction. Unfortunately, the crystal structures appear to have been destroyed by high temperatures and therefore, the solids cannot be accurately identified.

2.3.2 PartB Experiments in Part B examined the effect of adding magnesium to urine solutions before the onset of ureolysis, to determine if pre-hydrolyzed urine was appropriate as a feedstock for struvite recovery. The results from Part A indicated that solutions stored at 4°C were generally non-reactive and that there was little difference between solutions that were stored with or without lids. Thus, the solutions examined in this part were stored at room temperature with lids on (refer to Table 2.2 to a description of each solution)

2.3.2.1 Effect of magnesium on ureolysis It was originally thought that, by adding magnesium to the urine solutions, precipitation would occur. Although a small amount of precipitate was form, what was more interesting was the way in which the magnesium altered the behaviour of the solution. Effectively, the magnesium retarded the hydrolysis of urea, which, in turn, retarded the rate of pH increase and of ammonia production. Figure 2.6 shows how the pH in solutions with magnesium added is significantly lower than in the solutions examined without magnesium in Part A.

64 5 J , , , , , 1 0 20 40 60 80 100 120 Hour

5 i , ; , ; 1 0 20 40 60 80 100 120 Hour

Figure 2.6. Effect of magnesium on pH in a) full strength and b) diluted urine solutions

For both full strength and dilute solution sets, the solutions spiked with magnesium are slower to increase in pH, although none of the sterile solutions in either set show any increase. Magnesium is known to depress the supersaturation conditions for various calcium phosphate minerals and prolong the induction period of HAP by up to 4 times (Abbona and Franchiniangela, 1990; Abbona and Baronnet, 1996). Because of the delay in ureolysis, the onset of calcium precipitation was also delayed. The delay in ureolysis was likely due to the increased presence of ions in solution, which limit the ability of urease to attach to urea and catalyze the hydrolysis. The full strength solutions with magnesium never reached a steady state and, although the dilute solutions appear to be in equilibrium, the pH value attained is not as high as the non-spiked solution. It should be noted though, that the two experiments were run almost 2 months apart and, therefore, the wastewater used was not from the same batch (although it was from the same source and it should be of similar quality). The same wastewater was used in Part C and affected the urine the same way that it did in Part A, indicating that the delay observed in Part B was due entirely to the magnesium addition.

The evolution of ammonia was also delayed in the presence of magnesium. Figure 2.7 shows the difference in ammonia concentration for the full strength and dilute solutions.

66 Figure 2.7. Ammonia generation in a) full strength and b) dilute urine solutions Because of the short time scale and the erratic behaviour of ammonia during this time, it is difficult to quantify the effect that magnesium has on the generation of ammonia. However, from Figure 2.7, it appears that the generation of ammonia in both solution sets is affected, with the dilute set showing the greatest discrepancy.

Looking at Figure 2.6 again more closely, one can see that there is an immediate drop in pH as soon as the magnesium is added. The concentration of phosphate exhibits the same trend as shown in Figure 2.8.

350 -I 1

300 •— — ------—

Hour

Figure 2.8. Phosphate changes in magnesium dosed urine solutions. Dilute=dashed line, full strength =solid line

For unknown reasons, the phosphorus values displayed are almost exactly Vi what they should be. Likely, the dilution factor was noted incorrectly, but it cannot be proven that this was the actual reason for the discrepancy. Still, the pattern is consistent and, for that reason, the data are included.

There are several interesting features to notice in Figure 2.8. As discussed above, each solution exhibits a slight drop in concentration when the magnesium is added. Figure 2.8

68 shows that the small drop is then followed by a significant decrease, which removes almost all phosphate from solution. Abbona et al documented this behaviour when crystals were precipitated from a saturated phosphate solution (Abbona et al., 1986). The researchers noted that a steep pH drop occurs within the first 1-5 hours, which is followed by a subsequent pH drop over the course of the crystals formation and aging. When magnesium is added, the average loss is about 16 mg/L and 6 mg/L for full strength and dilute solutions, respectively. Compounds that are very insoluble and are close to being supersaturated are able to precipitate with only a slight increase in the ion activity product. The magnesium addition saturates the solution for some very soluble complexes and forces them to precipitate. Since so few compounds are saturated at this point, precipitated minerals consume very little phosphate. Later on, all solutions, except for the sterile solutions, then go on to lose a larger amount of phosphorus. The dilute solutions precipitate first, then in order of decreasing wastewater, and are followed by full strength solutions, which precipitate in the identical order (with respect to wastewater). The more pronounced drop in phosphorus that happens after at least 20 hours is a result of increased solution saturation with respect to phosphate minerals. By hour 20, pH and ammonia levels are elevated, which also cause an increase in the supersaturation ratio for some minerals. Elevated pH, magnesium and ammonium concentrations in solution force the solubility constants of phosphate minerals to be exceeded, resulting in the precipitation of solids and the loss of phosphate from solution. Unfortunately, the composition of the minerals that precipitated is not known since high storage temperatures altered the structure to the point that the crystal could not be identified.

2.3.3 Parte The effect of adding different quantities of magnesium to urine, after it had been hydrolyzed, was examined in Part C (refer to Table 2.3 for a description of the solutions used). The results of Part C show that phosphate minerals precipitate out of solution as a result of ureolysis. The spontaneous precipitation rids the solution of calcium, which can interfere with struvite formation, and produces a urine matrix that has all of the conditions necessary for struvite formation: high pH (9.3), high ammonia concentrations (>1000 mg/L) and low calcium content. Other work with real collection systems indicates that calcite, HAP, and struvite are the phosphate minerals that precipitate in urine collection systems (Udert et al.,

69 2003a). It is shown here that by adding magnesium to the remaining urine solution, pure struvite could be harvested.

2.3.3.1 rjH The 18 solutions examined in this part yielded similar results as the solutions examined in Part A Figure 2.9 shows the urease induced pH increase for each of the nine solution pairs, as well as the pH drop incurred when magnesium was added to the high pH solution.

9.5

0 24 48 72 96 120 144 168 Hours

Figure 2.9. Effect of magnesium on pH

The matched solution pairs behave almost identically, as they should, since they are subjected to the exact same conditions except for the magnesium dosage (recall that solutions #1-9 were dosed in a 1:1: Mg:PC»4 ratio and solutions #10-18 were dosed in a 2:1 ratio). It is interesting to note that the pH drop shown in Figure 2.9 indicates the addition of magnesium. Also, this drop is much larger than the pH drop shown in Figure 2.6, when magnesium is added to fresh urine solutions. This is predictable since only a small quantity of precipitate, and thus a small number of protons, were removed from solution in Part B (Figure 2.6), as

70 compared to the large volume of solids settled out when magnesium causes struvite to form in Part C. Just as the pH drops when magnesium is added to hydrolyzed urine, a similar precipitation-induced decrease in ammonia concentration is shown in Figure 2.10.

2000 1

1800

Hours

Figure 2.10. Changes in ammonia concentration with time

Again, the solution pairs are almost identical and there is very little difference between the two solutions, despite the fact that one is dosed with twice as much magnesium as its pair. Also, when comparing the values presented here to the values presented in Figure 2.2, it is clear that the ammonia concentration required to induce spontaneous precipitation (i.e. just before magnesium is added) is much lower.

2.3.3.2 Phosphorus The spontaneous precipitation of solids from solution is shown to consistently happen when the urine mixture, regardless of composition, reaches a pH greater than or equal to 8. The decrease in phosphorus concentrations, due to spontaneous precipitation, is shown in Figure 2.11.

71 900

0 24 48 72 96 120 144 168 Hours

Figure 2.11. Changes in phosphorus with storage

Phosphorus is almost completely removed from solution when supersaturated conditions are sufficient to release phosphate minerals. The precipitate was collected but, because it was a mixture of settled wastewater, organic residue, and crystals, the composition of it could not be determined. Although the minerals formed were not identified, the amount of calcium, magnesium and phosphorus that formed minerals is known. Figure 2.12 compares the percentages of calcium, magnesium and phosphorus removed due to spontaneous precipitation. Solutions that did not precipitate spontaneously, i.e. the sterile solutions that were not affected by ureolysis, are not shown.

72 SI Magnesium

Solution #

Figure 2.12. Percent loss of phosphorus, calcium, and magnesium to spontaneous precipitation

The above figure shows both data sets (solutions #1-9 are the same as solutions #10-18 except for the magnesium dosage) presented here for the sake of comparison. An average of 89% of magnesium, 83% of calcium, and 31% of phosphorus are removed due to spontaneous precipitation. Consequently, about 70% of phosphorus remains in solution to be harvested in the form of struvite. The solution that remains (after spontaneous precipitation) is used for struvite recovery as it contained only about 15% of the calcium that can interfere with optimal crystal formation. Work done with calcium and magnesium phosphates indicates that, unless the fraction of magnesium is greater than the fraction of calcium, struvite will not preferentially form (Abbona et al., 1986).

The mixed-mineral precipitate was filtered out, and the remaining urine solution was dosed with magnesium in either a 1:1 or a 2:1 Mg:P04 ratio, based on the assumed value of 650 mg PO4-P/L (the magnesium dose took into account dilution and was adjusted accordingly). When magnesium was added to the urine solutions, a white precipitate quickly formed and settled. Figure 2.13 shows the phosphorus balance, i.e. how much phosphorus was lost to

73 spontaneous precipitation, bound in struvite, or left in solution. Because the precipitated struvite stuck to the jar, it was difficult to obtain an accurate total mass of precipitate. The

amount allocated to struvite formation was assumed to be the difference between the

concentration in solution before and after magnesium addition.

B in solution LI struvite formation • spontaneous precip

2 3 5 6 8 9 11 12 14 15 17 18 Solution #

Figure 2.13. Allocation of phosphorus

74 100%

90%

80%

70%

60%

c • in solution B struvite formation H 50% • spontaneous preclp a 40%

2 3 5 6 8 9 11 12 14 15 17 18 Solution #

Figure 2.14. Allocation of calcium

The majority of phosphorus was removed from solution in the form of struvite, while the majority of calcium was removed via spontaneous precipitation. Figure 2.14 indicates that a small percentage (up to 8%) of calcium was formed into struvite; however, this may not be entirely true; subsequent data show that very little calcium contaminated the struvite. These data simply indicate that measurements of calcium in solution, before and after struvite formation, had a margin of error that may have amounted to several percent. The second set of nine solutions (solutions #10-18) was dosed with twice as much magnesium as the first nine; however there is no significant difference between the sets. The removal of phosphorus and calcium varies primarily with dilution; the fifth solution pair (# 5 and #14) seems to have the best removals of phosphorus to struvite; this pair is diluted by half and is 10% wastewater.

The crystals that formed in the relatively calcium-free urine matrix that was dosed with magnesium were analyzed with X-ray crystallography and found to be struvite. The struvite crystals were then analyzed for calcium (as described above) to determine the purity.

75 As shown in Table 2.7, the crystals contained very little calcium.

Table 2.7. Calcium content in struvite crystals Solution # % Calcium (by mass) Crystal Purity (%) 2 0.04 99.96 3 0.03 99.97 5 0.05 99.95 6 0.12 99.88 8 0.05 99.95 11 N.D* >99.98 12 N.D* >99.98 15 N.D* >99.98 16 0.04 99.96 17 N.D* >99.98 18 N.D* >99.98 * ND= Non-Detect value below 0.25 mg/L

Every crystal was at least 99.95% pure and there was no significant difference between the crystals from different solutions.

2.4 Conclusions Current struvite technology utilizes digester supernatant as a feedstock, to which a caustic and a magnesium source are added to increase the pH and supersaturation ratio to induce precipitation. This technique is effective, although chemical additions are costly and the concentration of phosphorus in the feedstock is low, thus, the amount of struvite recovered minimal. This research investigated the possibility of using phosphorus-rich urine as a feedstock for struvite, and found that pure struvite could, in fact, be recovered with minimal chemical additions. It was found that the natural changes in urine that result from ureolysis, can be used to optimally prepare stored urine for struvite recovery. Urine that has undergone ureolysis has a high pH (9.3), elevated ammonia levels and low levels of calcium (that can interfere with struvite formation). By using ureoloysis to raise the pH, rather than a chemical addition, significant savings at the full scale could be realized. The spontaneous precipitation that occurs naturally removes almost all calcium from solution, leaving a urine matrix that is still high in phosphate and ammonia and suitable for magnesium dosing and struvite formation. For dilutions greater than %, there is essentially no difference in the quality of struvite that is recovered; all solutions have a significant amount of phosphate remaining and

76 very low concentrations of other interfering metals. Struvite recovered was over 99.5% calcium free and different magnesium dosage ratios (1:1 or 2:1 Mg: P04) had no noticeable effect on the quality of struvite.

Furthermore, several other conclusions related to the storage and preparation of urine were drawn: • Low temperatures significantly slow the rate of ureolysis and consequently the rate of ammonia generation and pH decrease. • Nitrogen is lost through the volatilization of ammonia when urine is stored in the open although, compared to the total amount of ammonia present, the loss is relatively small. • Dilute urine is affected by ureolysis and reaches a maximum pH more quickly than full strength urine. • The onset of ureolysis is related to the amount of faecal matter that contaminates the solution; the more urease that contaminates the urine, the faster it will react and reach a maximum pH of approximately 9.3. • When magnesium is added to sterile urine before it has reacted, ureolysis is delayed; it is delayed more in full strength urine than in dilute urine. • Phosphate minerals will precipitate from stored urine that undergoes ureolysis; an average of 31% of phosphorus, 83% of calcium and 89% of magnesium will be lost to the spontaneous precipitation of minerals from solution.

By storing urine, allowing it to hydrolyze, isolating it from precipitated minerals, and dosing it with magnesium following this process, pure struvite can be recovered, regardless of dilution.

77 2.5 References

Abbona, F., Franchiniangela, M., 1990. Crystallization of calcium and magnesium phosphates from solutions of low concentration. J.Cryst.Growth, 104 (3), 661-671.

Abbona, F., Madsen, H. E. L., Boistelle, R., 1986. The initial phases of calcium and magnesium phosphates precipitated from solutions of high to medium concentrations. J.CrystGrowth, 74 (3), 581-590.

Abbona, F., Baronnet, A., 1996. A XRD and TEM study on the transformation of amorphous calcium phosphate in the presence of magnesium. J.Cryst.Growth, 165 (1-2), 98-105.

Adrian, A., Koch, F. A., Mavinic, D. S., 2003. Pilot-scale study of phosphorus recovery through struvite crystallization - II: Applying in-reactor supersaturation ratio as a process control parameter. J. Environ. Eng. Sci., 2 (6), 473-483.

Amtul, Z., Atta-ur-Rahman, Siddiqui, R. A., Choudhary, M. I., 2002. Chemistry and mechanism of urease inhibition. Curr.Med.Chem., 9 (14), 1323-1348.

Ban, Z., Dave, G., 2004. Laboratory studies on recovery of N and P from human urine through struvite crystallisation and zeolite adsorption. Environ.Technol., 25 (1), 111- 121. •

Berndtsson, J. C, 2006. Experiences from the implementation of a urine separation system: Goals, planning, reality. Build.Environ., 41 (4), 427-437.

Brooks, T., Keevil, C. W., 1997. A simple artificial urine for the growth of urinary pathogens. Lett.Appl.Microbiol., 24 (3), 203-206.

Griffith, D. P., 1979. Urease stones. Urol.Res., 7 (3), 215-221.

Griffith, D. P., Bragin, S., Musher, D. M., 1976a. Dissolution of struvite urinary stones - experimental studies invitro. Invest.Urol., 13 (5), 351-353.

Griffith, D. P., Musher, D. M., Itin, C, 1976b. Urease - primary cause of infection-induced urinary stones. Invest.Urol., 13 (5), 346-350.

Hanaeus, A., Hellstrom, D., Johansson, E., 1996. Conversion of urea during storage of human urine. Vatten, 52 , 263-270.

78 Hanaeus, J., Hellstrom, D., Johansson, E., 1997. A study of a urine separation in an ecological village in northern Sweden. Water Sci. Technol., 35 (9), 153-160.

Hellstrom, D., Karrman, E., 1996. Nitrogen and phosphorus in fresh and stored urine. Environ. Res. Forum, 5 (6), 221-226.

Hellstrom, D., Johannson, E., Grennberg, K., 1999. Storage of human urine: Acidification as a method to inhibit decomposition of urea. Ecol.Eng., 12 (3-4), 253-269.

Hesse, A., Siener, R., 1997. Current aspects of epidemiology and nutrition in urinary stone disease. World J.Urol., 15 (3), 165-171.

Hoglund, C, Stenstrom, T. A., Jonsson, H., Sundin, A., 1998. Evaluation of faecal contamination and microbial die-off in urine separating sewage systems. Water Sci. Technol., 38(6), 17-25. •

Huang, H., Mavinic, D., Lo, K., Koch, F., 2006. Production and basic morphology of struvite crystals from a pilot-scale crystallization process. Environ.Technol., 27 (3), 233-245.

Jonsson, H., Stenstrom, T. A., Svensson, J., Sundin, A., 1997. Source separated urine- nutrient and heavy metal content, water saving and faecal contamination. Water Sci. Technol., 35 (9), 145-152.

Larsen, T. A., Gujer, W., 1996. Separate management of anthropogenic nutrient solutions (human urine). Water Sci. Technol., 34 (3-4), 87-94.

Lienert, J., Haller, M., Berber, A., Stauffacher, M., Larsen, T. A., 2003. How farmers in Switzerland perceive fertilizers from recycled anthropogenic nutrients (urine). Water Sci. Technol., 48 (1), 47-56.

Lind, B. B., Ban, Z., Byden, S., 2000. Nutrient recovery from human urine by struvite crystallization with ammonia adsorption on zeolite and wollastonite. Bioresour.Technol., 73 (2), 169-174.

Mobley, H. L., Island, M. D., Hausinger, R. P., 1995. Molecular biology of microbial ureases. Microbiol.Rev., 59 (3), 451-480.

Moyo, C, Kissel, D., Cabrera, M., 1989. Temperature effects on soil urease activity. Soil Biol.Biochem., 21 (7), 935-938.

Pahl-Wostl, C, Schonborn, A., Willi, N., Muncke, J., Larsen, T. A., 2003. Investigating consumer attitudes towards the new technology of urine separation. Water Sci. Technol., 48 (1), 57-65.

Perrin, D. D. (1982). Ionisation constants of inorganic acids and bases in aqueous solution. Oxford: Pergamon.

79 Randall, C. W., 2003. Potential societal and economic impacts of wastewater nutrient removal and recycling. Water Sci. Technol., 48 (1), 11-17.

Steen, I., 1998. Phosphorus availability in the 21st century: Management of a non-renewable resource. Phosphorus and Potassium, British Sulphur Publishing, no. 217, September October, 25-31.

Udert, K. M., Larsen, T. A., Gujer, W., 2003. Biologically induced precipitation in urine- collecting systems. Water Sci. Technol.: Water Supply, 3 (3), 71-78.

Udert, K. M., Larsen, T. A., Gujer, W., 2003b. Estimating the precipitation potential in urine- collecting systems. Water Res., 37 (11), 2667-2677.

Ueno, Y., Fujii, M., 2001. Three years experience of operating and selling recovered struvite from full-scale plant. Environ.Technol., 22 (11), 1373-1381.

80 3 CHAPTER 3: Recovery of struvite from stored urine2

3.1 Introduction

Source separation, that is, the separation of urine from faeces at the point of collection, has emerged as one of the most promising ways of providing sanitation services in arid and developing regions, reducing water pollution and disease, and reclaiming valuable nutrients for agriculture (Berndtsson, 2006; Peasey, 2000; Rauch et al., 2003).

Urine contributes very little volume to mixed waste streams, but accounts for over 70% of the nitrogen and over 90% of the phosphorus that is expelled by humans (Hanaeus et al., 1997). In the absence of pathogenic faecal material, a clean, concentrated nutrient source can be harvested. Phosphorus is a finite resource, and half of the currently economically recoverable phosphate resources are expected to be exhausted in 60-70 years. Thus, the need for a sustainable source of phosphorus coupled with the need to reduce eutrophication caused by phosphorus pollution has further strengthened the case for urine separation.

As interest in urine separation grows, so does the need to better understand complex urine dynamics. Until recently, the chemistry and kinetics of urine were the exclusive domain of urologists, but increasingly it is sanitary engineers and agricultural researchers who are investigating the properties of urine in a new light. When urine is stored, it will inevitably undergo urease driven hydrolysis (ureolysis), such that the solution pH and ammonia concentration will increase (Hanaeus et al., 1996; Hellstrom et al., 1999; Hellstrom and Karrman, 1996; Udert et al., 2003a). As a consequence of this conversion, hydroxyapatite or

HAP (Caio(P04)6(OH)2), struvite (MgNH4P04 6H20), and/ or calcite (CaC03) precipitate from urine solutions and accumulate in collection and storage systems. The type and amount of mineral that precipitates is a function of the amount and type of dilution water that is used to flush the toilet or urinal (Udert et al., 2003b; Udert et al., 2003c). This spontaneous precipitation has the effect of lowering calcium, magnesium and phosphate concentrations in solution. Depending on the amount of water used, between 20-80% of phosphate will be removed in the form of mineral deposits (Udert et al, 2003b).

A version of this chapter has been/will be submitted for publication

81 Over the past decade, struvite crystallization from wastewater has emerged as an industry capable of producing commercial-quality struvite pellets from mixed sewage side-streams. In doing so, wastewater treatment plants equipped with crystallizer technology are able to supply a sustainable source of phosphorus to the market, while decreasing phosphorus discharge and waste (Abe, 1995; Britton etal., 2005). Depending on the composition of the feedstock and the amount of calcium in solution, HAP may precipitate along with, or instead of, struvite (Battistoni et al., 2000). Even if HAP does not form, calcium is known to interfere with the formation of magnesium phosphates (Abbona et al., 1986). Although different types of phosphates are potentially useful minerals, the current recycled phosphorus market is based on struvite, and a homogenous, pure, product is expected.

Struvite is comprised of magnesium, ammonium, and phosphate in equal molar amounts; since Urine is rich in phosphate and ammonia, the recovery of struvite from urine is logical. Researchers have recovered struvite, along with other minerals, from fresh urine that had been dosed with a caustic for pH control (Ban and Dave, 2004; Lind et al., 2000). However, fresh urine as a feedstock for struvite recovery is a virtual impossibility, as it undergoes ureolysis soon after entering the collection tank, and as mentioned earlier, changes in composition (Udert et al., 2003a).

This research examines the potential of stored urine, i.e. urine that has undergone ureolysis, to be used as a feedstock in struvite recovery. Stored urine has a high pH (>8), which is ideal for struvite recovery (Munch and Barr, 2001) and does not require the caustic addition that other feedstocks do. Moreover, calcium has been precipitated out of solution, leaving a urine matrix that is preferable for struvite formation. By working with urine that is diluted less than 25%, a maximum of phosphorus is available for struvite recovery.

The goal of this research is to determine how the quality of struvite recovered from urine that is prepared and stored in different ways, is affected. Additionally, the results obtained from previous work using synthetic urine are compared to the results presented here (using real urine) to determine how effective a proxy synthetic urine can be, in the study of urine-based nutrient recovery.

82 3.2 Materials and Methods Preliminary investigations utilized synthetic urine prepared as per the method given by Griffith (Griffith et al., 1976). Synthetic urine can be prepared on demand and in large volumes. Also, it is easy to prepare and pleasant to work with. For that reason, it was an adequate surrogate for bench scale investigations. The results of that Work are presented elsewhere (See Chapter 2). The results presented here are part of a larger investigation intended to replicate and verify previous results with real urine, in order to account for the presence of organic and non-organic compounds that could alter the behaviour of the matrix. Although synthetic urine is commonly used in urological research, the solution used contained only the eleven most prevalent constituents in urine. Presented here are the results of struvite recovery work done with real urine, and based on that work, an evaluation of synthetic urine as a proxy for real urine in nutrient recovery research.

3.2.1.1 Urine Collection In accordance with the requirements of the UBC Clinical Ethics Research Board, twelve volunteers were recruited to provide a 24-hour urine sample. The participants collected their urine over the course of 24 hours and returned the samples the following morning. The urine collected from the subjects had no identifying traits and could not be linked to the participant. The age and sex of the participant was recorded and before the unlabelled samples were mixed, the volume of each was recorded. Table 3.1 shows the summary statistics of the sample population.

Table 3.1. Summary statistics of urine collection Number of male participants 8 Number of women participants 4 Average age of men: 29 Average age of women 31 Average volume of 24-hour sample (L) 1.5 Maximum 24-hour sample volume (L) 2.9 Minimum 24-hour sample volume (L) 0.4 Total volume collected (L) 18.3

The collected urine was mixed into a composite sample. The composite urine mixture was

used immediately.

83 3.3 Experimental Design A factorial design was used to examine the effect of flushing water (dilution), faecal contamination (prevalence of urease), and storage (reaction time).

The effects of dilution were examined for two reasons. Practically, dilution water will always be present in stored urine either whether it is from leaks, cleaning, or from flushing. Secondly, the degree of dilution directly influences the ionic strength of the solution, which in turn, effects its precipitation potential and the amount of phosphorus that will be available for struvite recovery.

After several days of storage, different calcium and phosphate minerals will precipitate out of. urine. Since no anthropogenic additions are needed for this type of precipitation, it will be referred to as 'spontaneous precipitation'. Previous work indicated urine that was diluted by more than lA, lost nearly 80% of phosphorus to spontaneous precipitation, and thus, dilute urine would not be an appropriate feedstock for struvite recovery (Udert et al., 2003b). In this work, three levels of dilution were investigated: full strength (FS), half strength (1/2 S), and quarter strength (1/4 S). The urine was diluted with distilled water.

To induce ureolysis, urease—the urea-specific enzyme that catalyzes the hydrolysis of urea to ammonia—must be present. To account for different amounts of urease, three levels of faecal contamination were examined (since faecal material is rich in urease). Primary _ effluent from the UBC Pilot plant was used as a source of faecal urease. Primary effluent was chosen because it could easily be obtained without special permission (as opposed to solid faecal matter) and because it had fewer solids than raw sewage, but retained a significant amount of urease compared to the treated effluent. By adding varying amounts of primary effluent (wastewater) to urine, the amount of faecal contamination that could be used to prepare urine for struvite recovery, could be determined. Three levels of contamination were investigated: no contamination (pure urine), 10%, and 25% wastewater, whereby 10% or 25% of the total volume of the urine solution (regardless of its dilution) was replaced by wastewater. Figure 3.1 is a schematic diagram of the solution matrix.

84 Increasing dilution

Figure 3.1. Schematic of experimental design

Each 1-L mixture was prepared in a 1-L Nalgene, wide-mouth jar with a screw-top lid and was stored at room-temperature (approximately 22°C). Lids were left on the samples when they were not being examined and were open for the duration of the testing and sampling.

Two sets of the nine solutions were prepared. The experiment was conducted using the first set, set "a" (solutions #la-9a). The second set, set "b" (solutions #lb-9b) were prepared in the exact same manner as set "a". Both sets were allowed to undergo ureolysis and spontaneously precipitate. The precipitate from set "a" was then removed and the remaining solution was dosed with magnesium to induce struvite precipitation. The precipitate from set "b" was not removed and the solution was simply left to simulate continued storage. However, solution #2b was used as a replicate for #2a, i.e. it underwent all of the same treatment stages as solution #2a. Also, solutions #3b and #7b were used to determine the effect of adding magnesium before spontaneous precipitation. These two solutions were used to determine the effect of adding magnesium to a urine solution that had a lower pH and a significant concentration of calcium in solution.

85 3.3.1 Spontaneous Precipitation All urine solutions were monitored daily for changes in pH, and samples were taken daily (approximately 10 mL) to measure ammonia, phosphate, magnesium and calcium. Samples were filtered with a Millipore 0.45 um filter and preserved with 5% H2SO4 to pH 2. Samples were stored at 4°C until they could be analyzed. When the solution pH reached or exceeded - 8, spontaneous precipitation occurred and a solid accumulated on the bottom of the jar. Once a precipitate had formed, the supernatant was decanted and filtered with a Whatman #52 filter. The filtrate was re-sampled for all the parameters described above. It should be noted, however, that only solutions #la-9a were filtered; solutions #lb-9b were left untouched (i.e. the precipitate stayed in solution) but they were monitored continuously unless otherwise noted.

3.3.2 Struvite Precipitation Following the spontaneous precipitation of minerals from solution, the supernatant was decanted and filtered to obtain a clear urine solution that was free of solids. Magnesium was then added to the urine solution to induce struvite precipitation. The majority of magnesium that is naturally present in urine is removed from solution during spontaneous precipitation in the form of insoluble minerals and is, therefore, not available for struvite formation. Struvite is also known as MAP, or magnesium ammonium phosphate (MgNH4P04 6H2O) and therefore requires equal parts of the three species. Although magnesium is theoretically in a 1:1 molar ratio with phosphate, it has been shown that struvite formation (and phosphorus removal) is optimized when magnesium is added in excess, at a ratio of 1.3:1 (Dastur, 2001). Magnesium was added in excess based on the phosphate concentration previously observed in synthetic urine (approximately 650 mg PO4-P/L); however, because the actual amount of phosphate present was about 35% lower than expected, the ratio of magnesium added to phosphate was actually 1.7:1. Magnesium solutions were prepared using MgCl6 H2O and distilled water so that the solution would have a negligible effect on the pH. After the magnesium addition, the solution was mixed for 60 seconds and left to react for 1 hour. After the reaction phase, all parameters were measured and samples were taken. The precipitate (struvite) was then filtered with a Whatman #52 filter. The precipitate was saved for analysis and the filtrate was reanalyzed.

86 3.3.3 Chemical Analysis Ammonia and phosphate were measured with a Lachat QuikChem 8000 flow injection instrument. Calcium, magnesium, aluminium, iron and potassium were measured with a Varian Inc. SpectrAA220 Fast Sequential Fast Sequential Atomic Absorption Spectrophotometer. Atomic absorption was used to measure all of the elements except for potassium, which was measured with atomic emission. A Bruker D8 Advance powder X-ray diffractometer equipped with copper radiation and a graphite monochromator was used to identify the crystal structure. Contaminants in the crystals were measured by dissolving 100 mg of struvite crystals in 50 mL 0.5% nitric acid and analyzing the solution using atomic absorption as described above.

3.4 Results

3.4.1 Composition Compared to the typical values for urine in the Documenta Geigy (Diem and Lentner, 1970) (a standard reference text on biological constituents) and the synthetic mixture developed by Griffith, the average values of the constituents measured (ammonia, phosphate, calcium, magnesium) were consistently lower than either one or both sets of typical values. Table 3.2 presents a comparison between the three sets of data.

Table 3.2. Typical and measured concentrations (mg/L) of selected urine constituents

PO4-P NH3-N Mg Ca Documenta 760§ 480* 120* 230* Griffith 638 264 78 202 measured 450 300 70 70 1 '- 95% of lowest value; ""-average of the meaion irnliiovalueps rV\rfor- twomen annund \women l ;/-v t"v\ an; mean value for adults

The measured values are consistently lower than the other standard references. According to the Documenta Geigy, each person excretes approximately IL of urine per day. Since several of the participants submitted large volumes of urine (see Table 3.1) the composite sample was likely more dilute than average. Because work with urine is time sensitive and analysis cannot be done immediately, the composition was not known until a few days into the experiment. The values, although low, are not completely unrealistic and thus, the urine obtained was considered adequate for experimental use.

87 3.5 Ammonia and pH When urease hydrolyzes urea, ammonia is produced, which in turn dissociates, increasing the pH of the solution:

CO(NH2)2+H20-» C02+2NII3

The supersaturation of the solution is increased and spontaneous precipitation occurs. The following figures show the increase of pH with time. Figure 3.2 shows the changes that occurred in solutions (#la-9a) that were dosed with magnesium following spontaneous precipitation and Figure 3.3 shows the changes that occurred in solutions that were left unaltered.

88 Figure 3.2. pH changes in solutions with a) no added wastewater and b) 10% and 25% wastewater 90 From Figure 3.2 it can be seen that there is a significant difference rate of pH change between the solutions that had no urease added (Figure 3.2a) and those that did (Figure 3.2b). Those solutions that had wastewater added reacted quickly; however, the actual amount of wastewater that was added (10% or 25%) did not significantly affect the rate of pH increase. Solutions with the most wastewater did not necessarily react first because dilution also affected the rate. Rather than increasing immediately, the pH of the solutions (shown in Figures 3.2a and 3.3a) slowly dropped and then slowly increased again before spiking. The solutions that were dosed with magnesium (Figure 3.2) display a sharp drop in pH when magnesium was added; this indicates the formation of struvite. Although magnesium was added following spontaneous precipitation, it is not possible to tell where spontaneous precipitation occurred from these graphs, only where magnesium was added. The solutions that were left to undergo ureolysis for the duration of the storage time (Figure 3.3), reached a

maximum pH value of about 9. It is worth noting that solutions #2b and 3b (Figure 3.3b) were dosed with magnesium before spontaneous precipitation. This explains why the solution pH dropped and why the solutions were not able to reach a stable pH value.

Similarly, the concentration of ammonia in solution increased with increased storage. Figures 3.4 and 3.5 show the changes that occurred as a result of ureolysis.

91 Figure 3.4 Changes in ammonia with storage time in solutions with a) no added urease b) 10% and 25% wastewater added Time (hours)

Figure 3.5. Ammonia changes in solutions dosed with magnesium and a) no added urease, 10% and 25% wastewater Figure 3.5 shows how ammonia levels in solutions spiked with wastewater increased steadily after approximately 24 hours, while solutions without added wastewater did not begin to increase until after several days. Similarly, the concentrations increased, as in Figure 3.4; however, there is a sharp decrease in ammonia when the solutions are dosed with magnesium. Some of the solutions in Figure 3.5 reached steady state concentrations; the difference between the maximum ammonia levels is simply a function of the initial dilution with water and wastewater (i.e. less urea to hydrolyze).

3.6 Spontaneous Precipitation Precipitation consistently happened when the urine mixture, regardless of composition, reached a pH greater than or equal to 8. Spontaneous precipitation occurred in a predictable fashion that varied first with dilution and then with wastewater (urease). The first mixtures to precipitate were the most dilute with the higher percentages (10 and 25%) of wastewater: solutions #6, 8 and 9. The second set of mixtures to precipitate were those that were less dilute but still with a high proportion of urease: solutions #2, 3, and 5. The urine mixtures with no added urease were last to precipitate: solutions #1,4, and 7. Although equipment was sanitized with bleach and steps were taken to keep the urine sterile, urease was likely present from the time of collection and thus, the urine could not be kept sterile.

3.7 Calcium Removal The amount of calcium removed via spontaneous precipitation was calculated by balancing the amount of calcium in solution before and after the precipitate had formed. The data presented in Figure 3.6 shows the amount of calcium, magnesium, and phosphorus lost from solution when spontaneous precipitation occurred. Figure 3.6a shows the amount lost in terms of a percent of the initial amount, and Figure 3.6b shows the absolute change in solution concentration. The data shown are averages of the two solutions.

94 120

• Phosphorus • Calcium B Magnesium

Solution #

Figure 3.6. Spontaneous removal of calcium, magnesium and phosphorus in terms of a) percent and b) concentration

From Figure 3.6 it can be seen that the removals of calcium, magnesium and phosphate all follow different patterns. As a percent of the total (Figure 3.6a), magnesium is removed

95 consistently at an average of 80%, as is phosphorus, at an average of 23%. The amount of calcium removed, however, is highly solution dependent; a maximum of 80% down to a minimum of 13% is removed, depending on the solution. In general, the percent of calcium removed decreases with increasing dilution. However, as shown in Figure 3.6b, the concentration of each of the elements generally decreased with increasing dilution. Although over 100 mg of phosphorus can be lost (solution #2), solutions that are less dilute retain a higher percentage of phosphorus and a minimal amount of calcium.

3.8 Struvite Recovery

3.8.1 Phosphorus There is a near total removal of magnesium from solution due to spontaneous precipitation, indicating struvite formation (since struvite is the only magnesium based compound known to precipitate from urine). However, because a variety of compounds precipitate together, it is not possible to recover struvite independently from the other minerals. Also, the minerals that form during spontaneous precipitation are mixed with organic sludge and are not in a readily useable form. To recover the remaining 70% of phosphorus n the form of struvite, additional magnesium (in the form of magnesium chloride) must be added to post- precipitation urine to induce precipitation. In this way, struvite, and only struvite, can be harvested from the urine solution that has lost the majority of the magnesium and calcium and only 26% of the phosphorus. Figure 3.6 compares the fate of phosphorus in both data sets "a" and "b" (i.e. solutions dosed with magnesium and solutions not dosed).

96 ~~ In solution • struvite formation • spontaneous precip

la 2a 4a 5a 6a 8a 9a lb 2b 4b 5b 6b 8b 9b Solution #

Figure 3.7. Allocation of phosphorus

Since the real goal of struvite recovery is phosphorus recovery, it is important to minimize the amount left in solution or in mixed mineral precipitate, and maximize the amount bound into struvite. An average of 70% of the phosphorus was recovered in the form of struvite, and only an average of 6% was left in solution. The most dilute solutions (# 8a and 9a) had the greatest amounts of phosphorus remaining in solution, while the other solutions had very low percents (i.e. 3%) still in solution. Since solutions # 3b and 7b were dosed with magnesium before spontaneous precipitation, the results are significantly different (as presented later in Figure 3.10). In general, it appears that there is an increasing percent of phosphorus remaining in solution with increasing dilution. Figure 3.8 shows the concentrations of phosphorus remaining in solution.

97 <

E 20

o IB Phosphorus

15 c o u 10

int.;* I—v. • ^1 "•VI Hi La. Id

la 2a 3a 4a 5a 6a 7a 8a 9a Solution #

Figure 3.8. Concentration of phosphorus remaining in solution

It appears that there is a threshold concentration for phosphorus: concentrations below 4 mg/L were not obtained. However, in a full-scale operation, increased mixing and recycle may be effective at lowering the threshold limit. It is also interesting to note that the concentration of phosphorus in solution begins to increase with increasing dilution and wastewater. The concentration peaks at solution #7 and then decreases slightly in the solutions that have increased amounts of wastewater.

3.8.2 Calcium By removing calcium and other contaminant ions through spontaneous precipitation, a "cleaner" matrix is available for struvite formation. Figure 3.5 shows that up to 80% of calcium was precipitated from solution, depending on the dilution. However, when magnesium was added to solution, an interesting phenomenon occurred: calcium that was thought to have precipitated in mineral form, returned to solution. Figure 3.9 shows how calcium levels dropped and then sharply increased again when magnesium is added.

98 80

Solution #

—K-4a

:~- 6a

-B-8a

0 0 24 48 72 96 120 144 168 Time (hour)

Figure 3.9. Change in calcium over time

The sharp increase in calcium was unexpected, but there is a plausible explanation. Following the spontaneous precipitation, the urine solution was filtered using a Whatman #52 filter. This filter was chosen because, after several hours of using a finer Whatman 934- AH (1.5mm) filter, very little of the liquid had passed through. The #52 filter appeared to retain all of the solids while allowing the liquid to pass through in a timely manner. Likely, this filter allowed tiny crystals to pass through so that, when the magnesium was added, some mineral-bound calcium could be replaced by magnesium (which liberated calcium ions to solution). Practically, extended settling periods, filtration, and/or centrifuge could be used to minimize the amount of fine calcium crystals present in solutions used for struvite recovery.

3.8.3 Timing of Magnesium Dosage Two solutions from set "b" were dosed with magnesium prior to spontaneous precipitation to determine the effect of magnesium addition on urine that still had high concentrations of calcium. Solutions # 3a and 7a were dosed with magnesium after spontaneous precipitation, and #3b and 7b were dosed with magnesium before spontaneous precipitation. By comparing

99 these two sample pairs, what effect the magnesium addition timing had on the type and quality of mineral recovered could be determined. Figure 3.10 shows how phosphorus was allocated for both solution pairs.

U in solution • struvite formation • spontaneous preclp

3a 7a 3b 7b Solution #

Figure 3.10. Allocation of phosphorus in solutions #3 and 7

The results from this comparison experiment were not predicted. Because of the elevated magnesium and calcium levels and because of the relatively low pH values (i.e. <9), struvite was not expected to form. The presence of calcium was expected to favour the formation of HAP or at least prevent the formation of struvite. However, struvite was formed.

The first thing to note in Figure 3.10 is the fact that neither solution #3b nor 7b lost any phosphorus to spontaneous precipitation, while the corresponding pair did. More phosphorus remained in solutions #3b and 7b than in the corresponding pair; however, more phosphorus was actually recovered as struvite in solution #3b than in #3a. In solution #3b, nearly 90% of the phosphorus was recovered as struvite, which would indicate that perhaps an early magnesium addition is more effective than one following spontaneous precipitation. The same cannot be said about solutions #7a and # 7b: far more phosphorus was recovered as

100 struvite from solution #7a, when the solution was dosed with magnesium after spontaneous precipitation. Based on these few, inconsistent data points, it is difficult to determine what the real effect of an early magnesium addition was.

3.8.4 Purity of Crystals The recovered struvite crystals were analyzed for four of the contaminants most commonly encountered in struvite (Huang et al., 2006) and the results are presented in Table 3.3. The level of contamination is given as a percentage of the struvite mass.

Table 3.3. Contamination as a percentage of struvite mass metals Struvite from mixture # %Fe % Al %K %Ca la n.d. n.d. n.d. 0.23 2a n.d. n.d. 0.44 0.22 2b n.d. n.d. 0.42 0.22 3a n.d. n.d. 0.38 0.18 3b n.d. 0.12 0.50 0.44 4a n.d. n.d. 0.43 0.31 5a 0.0025 n.d. 0.40 0.23 8a n.d. n.d. 0.43 0.21 MDL as % metal 0.0002 0.01 0.06 0.00 PLQ as % metal 0.0012 0.10 0.32 0.02

Generally, iron and aluminium were not present in measurable quantities, and potassium and calcium were constituted less than 0.5% of the total mass. Since these are the most common contaminants, the struvite produced is likely over 99% pure. The struvite produced from solution #7b was unavailable for analysis and so is not included; however, the results from solution #3b are not significantly different than those from solution #3a, which indicates that the calcium in solution did not affect the purity of the struvite. Further work is needed however, since nothing conclusive can be said about this single result.

3.8.5 Comparison The data presented here, using real human urine, are the verification of work done previously with synthetic urine. Synthetic urine was used in preliminary studies for several reasons. The use of human urine is subject to both ethical and safety approval, it must be used right away, collection requires detailed coordination among volunteers, and the recruitment of volunteers is not always easy. Synthetic urine is a recognized substitute for real urine and is

101 used frequently in urological research. Obviously, the two solutions are different, and after conducting the same experiment with both real and synthetic urine, those differences, and the importance of them, can be more accurately judged to determine whether or not synthetic urine is an acceptable surrogate in the study of nutrient recovery from urine.

3.8.6 Ureolysis The first and most noticeable difference between real and synthetic urine is the smell. As synthetic urine undergoes ureolysis it gives off a strong, but clean, ammonia smell; similar to an ammonia-based cleaning product. When real urine undergoes ureolysis, it gives off a repulsive, pungent odour, that smells more of festering sewage than ammonia. Work with synthetic urine was conducted on a laboratory bench in a laboratory shared with several other researchers, and the smell was not offensive. To work with real urine this researcher had to use a eucalyptus salve (Vies VapoRub™) under her nose, while wearing a facemask and working under a fume hood to keep from becoming nauseous. The personal and societal reasons for working with synthetic urine are obvious.

The differences in ureolysis were also noticeable in the behaviour of pH and ammonia levels. The pH in synthetic solutions increased steadily, in contrast to the dip in pH displayed initially by real solutions. Also, the synthetic sterile solutions stayed "sterile" for at least 10 days but the real solutions became contaminated and reached a maximum pH just before 7 days. Similarly, ammonia levels in synthetic solutions, on average, reached a steady state level after 15 days, while similar real solutions reached a maximum in 2-5 days. In terms of ureolysis, it seems that the lack of organic matter in the synthetic urine retards the onset and progression of ureolysis; likely, the lack of organic foodstuff for urease-producing organisms is a limiting factor and results in a lower level of urease. Also, since all of the real "sterile" solutions became contaminated, it seems that only a very slight amount of urease contamination is needed to induce ureolysis in real urine, whereas slight, if not significant, contamination can be endured by synthetic solutions with little effect. All of this implies that in an experimental setting, it may require more time to achieve steady-state for synthetic urine, especially if only a minimal amount of wastewater is added.

102 3.8.6.1 Mineral Precipitation As shown in Table 3.2, the composition of the real urine in this work and the composition of the synthetic urine are not entirely comparable; the real urine collected was more dilute, but the dilution was not linear with respect to every ion. Still, the real urine, albeit more dilute, did not behave the same as dilute synthetic urine, which indicates a difference that is caused by more than just dilution. Table 3.4 is a summary of the average percent removals of calcium and phosphorus at different stages of processing, for both synthetic and real urine.

Table 3.4. Comparison of mineral formation in synthetic and real urine Element Average % removed' Synthetic Real % removed in spontaneous precipitation. 31 24 Phosphorus % removed in struvite precipitation 62 70 % remaining in solution 7 6 % removed in spontaneous precipitation. 83 40 Calcium % removed in struvite precipitation 15 -6 % remaining in solution 1 66 Magnesium % removed in spontaneous precipitation. 89 80 Number of values averaged: n= 12 14

Generally, the removal of phosphorus is comparable between synthetic and real urine. It is interesting, that despite being a more dilute, and complex solution, more phosphorus was removed from real urine than synthetic urine. The real urine lost an average of 7% less phosphorus to spontaneous precipitation and an average of 8% more phosphorus was subsequently recovered in the form of struvite. It appears then, that the higher recovery of struvite may be directly related to the amount of phosphorus that remains in solution following spontaneous precipitation, however, because the two types of urine had different concentrations and ionic ratios (e.g. Mg:Ca), no single factor can account for the differences observed.

Figure 3.6 shows that calcium removal is highly dependent on the urine dilution; Table 3.4 shows that calcium removal is also dependent on urine type. Table 3.4 shows that less than half of the calcium removed by spontaneous precipitation in the synthetic urine was removed in the real urine. As discussed earlier, a coarser filter had to be used when filtering precipitate from the urine solution because of the thick yellow sludge that formed along with

103 the precipitate and prevented its easy removal (the precipitate formed in the synthetic urine was very crystalline and powdery, and filtered easily on a fine filter). Thus, the difference in filter can account for the increase in calcium that occurred when magnesium was added, but it cannot account for the other discrepancies. Not only were calcium removals due to spontaneous precipitation much higher in synthetic urine, the trend of precipitation was entirely different. Calcium removal in synthetic urine showed an arching trend; with increasing dilution, removal percentages up to a mid-point, and then fell to levels slightly below those of undiluted samples. In real urine, calcium removals decreased steadily from full strength down to the most dilute solution. It should be noted however, that since the real urine was more dilute, the downward trend could be an extension of the downside of the curve exhibited by synthetic urine. The highest calcium removal percentages in synthetic urine are comparable to the removal percentages seen in the most dilute real solutions, which could be expected. In both cases, however, the struvite that was recovered was essentially pure.

The sludge that was produced by spontaneous precipitation from real urine was a yellow, viscous sludge with white crystals embedded in it, whereas only white powder was recovered when spontaneous precipitation occurred in synthetic urine. The concurrent precipitation of minerals and sludge from real urine solutions made the identification of crystals difficult.

Analytically, real urine proved to be more difficult to work with than synthetic urine. The presence of organics and unknown compounds made the urine more difficult to analyze. When analyzing for metals with the atomic absorption instrument, the burner constantly became encrusted with salts and had to be removed and cleaned after ever 30 samples. Also, there were no standard methods provided with the instrument or in standard analytical texts that described how to prepare urine for analysis, so the samples were often diluted and treated like water samples, which is not entirely correct.

Although real urine contained organic matter and was more complex, the behaviour was generally the same as that of synthetic urine. Despite minor differences in the way that the solutions underwent ureolysis and the formation of organic sludge, the solutions were, for

104 this purpose, interchangeable. However, to obtain the most descriptive results, the real urine feedstock should be characterized, so that a synthetic solution with the same composition can be prepared.

3.9 Conclusions Unless urine is stored under very sterile conditions, spontaneous precipitation will occur. Increased levels of urease and dilution were shown to induce spontaneous precipitation more quickly.

Calcium removal, due to spontaneous precipitation, was found to decrease from a maximum of 80% to a minimum of 13%, with increasing dilution and increasing contamination. In general, the percent of calcium removed decreases with increasing dilution. When magnesium was added to the urine liquor following the removal of precipitated minerals, the concentration of calcium in solution was found to increase, possibly as a result of magnesium replacing calcium in tiny crystals that had not been filteredout . In scaled-up recovery processes then, urine liquors to be used for struvite recovery should be carefully decanted and/or filtered so that fineparticle s are not re-suspended into solution before magnesium dosing.

Approximately 24% of the phosphorus in solution spontaneously precipitated as phosphate minerals; the removal was generally independent of dilution (at low dilution levels

Organic constituents in the real urine likely accelerated the onset and the rate of ureolysis, but otherwise, it was difficult to determine if organics had any effect on the precipitation potential of a given solution. Certainly, the organic sludge that settled out of the real urine hindered the recovery of any calcium phosphate minerals, but in the absence of that sludge, struvite was formed and harvested with ease. Struvite was recovered from the urine matrix that remained following spontaneous precipitation and approximately 70% of the phosphorus was recovered in the form of struvite although the percentage and mass of phosphorus recovered as struvite decreased with increased dilution. High quality struvite (-99% pure)

105 was recovered from human urine when post-precipitation urine was dosed with magnesium in a ratio of 1.7:1 (Mg:P04). In a single solution comparison, the struvite recovered from a urine solution that had not undergone spontaneous precipitation, was essentially as pure as struvite recovered from a calcium-reduced urine matrix.

Human urine is variable; it varies with time, diet, gender, person, and a host of other factors. In this work, the urine was found to be very dilute compared to published values and synthetic urine. Despite the fact that the sample used was more dilute, the general behaviour was consistent with previous work using synthetic urine. Most importantly, the amounts of phosphorus removed were similar and the purity of the struvite recovered was equally high. For scaled-up processing (e.g. in a struvite crystallizer), synthetic urine can be used as a proxy for urine however the synthetic urine should be prepared to reflect the specific composition of the urine feedstock that can be expected and real urine should be used for final verification. Although urine composition varies with age, gender, environment, and a suite of other factors, representative population studies should be done to deterine the general range of constituents, concentrations, and ionic ratios that would normally be found in collected urine. Although the amount of calcium in solution does not affect the purity of the struvite, it does seem to affect the amount of phosphorus that can be recovered as struvite. Other factors, such as the amount of magnesium required, the amount of sludge generated, and the value of the struvite that can be recovered, will be calculated most accurately when the synthetic urine most closely matches the qualities of the real feedstock to be used.

In the process of recovering struvite from urine two different waste streams are created: a high pH, high ammonia urine solution that has few nutrients remaining, and a sludge that is a mixture of organics and precipitated minerals. Full-scale operations must consider how to dispose of or, preferably, beneficially use these waste streams.

By allowing stored urine to increase in pH naturally, a urine matrix that is low in calcium and has a composition favourable for struvite recovery can be produced. The most concentrated urine is the best feedstock for struvite recovery because the most phosphorus can be harvested from solution following spontaneous precipitation. By simplifying the struvite

106 recovery process, while reducing the cost of chemical and infrastructure controls, the possibility of widespread struvite recovery from urine comes closer to reality. 3.10 References

Abe, S., 1995. Phosphate removal from dewatering filtrate by MAP process at seibu treatment plant in Fukuoka city. Sewage Works in Japan, 43, 59-64.

Ban, Z., Dave, G., 2004. Laboratory studies on recovery of N and P from human urine through struvite crystallisation and zeolite adsorption. Environ.Technol., 25 (1), 111- 121.

Battistoni, P., Pavan, P., Prisciandaro, M., Cecchi, F., 2000. Struvite crystallization: A feasible and reliable way to fix phosphorus in anaerobic supernatants. Water Res., 34 (11), 3033-3041.

Berndtsson, J. C, 2006. Experiences from the implementation of a urine separation system: Goals, planning, reality. Build.Environ., 41 (4), 427-437.

Dastur, M. (2001). Investigation into the factors affecting controlled struvite crystallization at the bench-scale. (MASc thesis, Department of Civil Engineering, The University of British Columbia).

Diem, K., & Lentner, C. (Eds.). (1970). Documenta Geigy: Scientific Tables. Basle Geigy.

Griffith, D. P., Musher, D. M., Itin, C, 1976. Urease - primary cause of infection-induced urinary stones. Invest.Urol., 13 (5), 346-350.

Hanasus, A., Hellstrom, D., Johansson, E., 1996. Conversion of urea during storage of human urine. Vatten, 52 , 263-270.

Hanaeus, J., Hellstrom, D., Johansson, E., 1997. A study of a urine separation in an ecological village in northern Sweden. Water Sci. Technol., 35 (9), 153-160. Hellstrom, D., Karrman, E., 1996. Nitrogen and phosphorus in fresh and stored urine. Environ.Res.Forum, 5 (6), 221-226.

Hellstrom, D., Johannson, E., Grennberg, K., 1999. Storage of human urine: Acidification as a method to inhibit decomposition of urea. Ecol.Eng., 12 (3-4), 253-269.

Huang, H., Mavinic, D., Lo, K., Koch, F., 2006. Production and basic morphology of struvite crystals from a pilot-scale crystallization process. Environ.Technol., 27 (3), 233-245.

108 Lind, B. B., Ban, Z., Byden, S., 2000. Nutrient recovery from human urine by struvite crystallization with ammonia adsorption on zeolite and wollastonite. Bioresour.Technol., 73 (2), 169-174.

Munch, E. V., Barr, K., 2001. Controlled struvite crystallisation for removing phosphorus from anaerobic digester sidestreams. Water Res., 35 (1), 151-159.

Peasey, A. (2000). Health aspects of dry sanitation with waste reuse. Task No. 324, Water and Environmental Health at London and Loughborough.

Rauch, W., Brockmann, D., Peters, I., Larsen, T. A., Gujer, W., 2003. Combining urine separation with waste design: An analysis using a stochastic model for urine production. Water Res., 37 (3), 681-689.

Udert, K. M., Larsen, T. A., Biebow, M., Gujer, W., 2003a. Urea hydrolysis and precipitation dynamics in a urine-collecting system. Water Res., 37 (11), 2571-2582.

Udert, K. M., Larsen, T. A., Gujer, W., 2003. Biologically induced precipitation in urine- collecting systems. Water Sci. Technol.: Water Supply, 3 (3), 71-78.

Udert, K. M., Larsen, T. A., Gujer, W., 2003c. Estimating the precipitation potential in urine- collecting systems. Water Res., 37 (11), 2667-2677.

109 CHAPTER 4: Discussion and Conclusions

4.1 Summary The results contained in this thesis demonstrate that pure struvite can be recovered from stored urine. The work was divided into two sections and presented in two separate, but related papers. The first paper summarizes the work done to determine the effect of different storage and magnesium dosing regimes on struvite recovery from synthetic urine. The second paper summarizes the work done to verify the results from the first paper with real urine and compares the two sets of results.

Under the main goal of struvite recovery, several specific research questions (as outlined in Chapter 1) were answered. Increased levels of dilution, temperature, and faecal contamination allow ureolysis to set in more quickly and increase the rate at which it occurs. Calcium, magnesium, and phosphorus are all removed from solution in the)form of insoluble minerals whereas ammonia concentrations continue to increase until equilibrium is achieved, at which point a constant level is maintained. The spontaneous precipitation of minerals occurs when the pH reaches 8. The amount of calcium removed is dependent on dilution, while the amount of phosphorus removed, is not.

Struvite was recovered from stored urine without the addition of a caustic; since caustic can account for up to 97% of the chemical costs in struvite production (Jaffer et al., 2002), the importance of this cannot be overstated. The natural pH increase due to ureolysis was sufficient to achieve a pH suitable for struvite recovery and remove the need for caustic addition. Magnesium added to urine that had not undergone spontaneous precipitation retarded ureolysis. Magnesium that was added after ureolysis-induced spontaneous precipitation immediately caused the precipitation of calcium-free struvite. The maximum mass of phosphorus was recovered from full strength urine; the maximum mass of calcium was also removed from full-strength urine due to spontaneous precipitation, thus leaving a minimum amount available in solution. It was not entirely clear how affected by calcium, struvite would be, if calcium was not removed from solution before struvite recovery.

110 The most important finding is that characterization and dilution are the most important factors that influence the quantity of struvite that can be recovered. The strength of the urine (i.e. the concentration of phosphorus, calcium, and magnesium) affects the amount of ions that are lost to precipitation and ultimately the quantity that remain for struvite formation. Stored urine should be stored with as little dilution water as possible to ensure that a maximum amount of phosphorus is available for struvite recovery.

The most important findings of this research are that a) struvite can be crystallized from stored urine without the addition of caustic, and b) that within the range examined, faecal contamination and dilution do not affect the quality of struvite that can be recovered. By allowing urine to achieve high levels of ammonia and low levels of calcium with natural storage, an ideal feedstock for struvite formation can be prepared. A follow-up study utilizing stored urine as a feedstock in a fluidized bed reactor will verify this hypothesis at the full-scale.

4.2 Strengths and weaknesses This research was groundbreaking; prior to this work, no one had researched the resource potential of urine in the Department of Civil Engineering, dr at UBC, for that matter. In fact, apart from personal communication with one researcher at the University of Waterloo, this researcher has not found evidence that similar research is taking place anywhere in Canada. Furthermore, there have only been two papers published (in English) that investigate the recovery of struvite from fresh urine, and this work appears to be the first to address struvite recovery from stored urine. While this work is innovative, it is limited by a lack of infrastructure and resources. This type of research (nutrient recovery/Ecosan) is developing rapidly in countries like Sweden, Switzerland, and the Netherlands, where research projects and large-scale Ecosan infrastructure (e.g. public buildings, dormitories, etc.) have been functional for decades. Work done in these settings makes use of the real urine collected in collection tanks, collected from hundreds of people, that is stored under real conditions. This researcher could only attempt to replicate these conditions at the bench scale, in a laboratory. Because of these physical constraints, the research was limited in some respects. One-litre volumes were used in all the work, so the effects of stratification, and other scale factors were likely not detected. Storage tanks are added to constantly over the course of a day;

111 fresh urine is quickly mixed with stored urine and therefore fresh urine will never remain sterile. It was, however, the intent of this research to understand the initial phases of urine conversion, and so by beginning with sterile conditions, the factors that influence conversion could more easily be identified. Urease is ubiquitous in storage tanks, whereas in experiments, it was introduced artificially by using primary effluent from the wastewater treatment plant. Wastewater would likely never contaminate stored urine; however, the effluent used was easy to obtain, it could be used without special approval (that would not be the case with solid faecal matter) and it was a guaranteed source of urease. So, while it was an adequate substitute, the actual amount of urease in it was never measured. The concentration of urease could not be related to any of the other parameters; only a general relation between the volume of wastewater added and the other parameters could be developed.

It is well known that struvite turn into a different mineral if ambient conditions change (Babic-Ivancic et al., 2004; Boistelle et al., 1983). For that reason, the struvite was removed from solution after only 1 hour of reaction time, since the ammonia concentration and pH were still steadily increasing. It is not known if more or less struvite would have formed if the system had been allowed to reach equilibrium. If magnesium had been added to the urine solution once it had reached equilibrium, the struvite may not have been as susceptible to changes. For most experiments, daily samples for pH, ammonia, calcium, phosphorus, and magnesium were taken, and the density of this level of sampling was found to be more than adequate. If, however, pH measurements were taken more densely, especially during precipitation events, more information could have been captured about the conditions that are responsible for mineral precipitation.

The rate that the urine (both artificial and real) went through the filters was unpredictably slow. In some instances, so much precipitate accumulated on the filter that no further liquid could pass. In these instances, the filter had to be removed and replaced with a second filter to be used for the remainder of the filtering. Due to the rapid pace and the time-sensitive nature of this process, the mass of the precipitate (summed over numerous filters) was not

112 obtained. The purpose of this work was to examine primarily the quality of the struvite obtained, not the quantity, and so this missing data did not affect the objectives of the study.

This work was a comprehensive investigation into the behaviour of synthetic and real urine, and the subsequent effects on the quality of recovered struvite. This body of research has laid the foundation for what will hopefully be continued research into the recovery of struvite from urine.

4.3 Recommendations for Future Work From this foundation work, there are many possible directions for follow up studies that will significantly further the state of knowledge on this topic.

To continue this line of research, the first step that should be taken is using the UBC struvite crystallizer with stored urine as a feedstock. Producing useable struvite crystals will add significantly the fields of struvite recovery and urine re-use. However, prior to using the crystallizer, continued bench scale work should be done to determine the changes in urine until it comes to equilibrium. Based on this information, the ideal storage time can be determined. Since the degree of contamination makes no discernable difference to the quality of struvite, the effect of different levels of dilution should be further examined. Tied to this should be an examination of the effect of calcium in solution; the calcium concentration range that can be tolerated in solution without affecting the purity of the recovered struvite should be determined.

A follow up to crystallization should be an investigation into the use of "spent" urine liquor, i.e. urine from which struvite has been recovered. This urine will have little to no phosphorus, but will still have very high levels of ammonia and a pH around 9. One possible use is to recover the remaining ammonia using a zeolite (Booker et al., 1996; Cooney et al., 1999a; Cooney et al., 1999b) or stripping the ammonia and recovering it in the presence of phosphoric acid, to generate another type of phosphorus-based fertilizer. Alternatively, the nitrogen rich solution could be used as an additive at WWTPs to improve processes or in struvite recovery processes, using digester supernatant that is low in pH or ammonia.

113 As was noted in the research using real urine, the calcium concentration increased when magnesium was added because of a process of substitution. This phenomenon and its effects on scaled-up struvite recovery should be investigated further.

Different levels ofmagnesium additions were examined to determine what, if any, effect excess magnesium had on struvite quality, but no significant effect was noticed. A systematic study of the effect of magnesium on the quality of struvite would allow for the quality, and the cost of, struvite to be optimized. As well, a study of the relationship between magnesium and both calcium and phosphorus would be useful in determining the mechanisms that favour struvite formation.

An interesting, if not pressing, field of research would be an examination of the effect of urease inhibitors. Several compounds are known to impair or impede the ability of urease to cleave urea (Amtul et al., 2002; Amtul et al., 2004; Griffith et al., 1978; Morris and Stickler, 1998; Varel et al., 1999). The effect of urease inhibitors on stored urine and the subsequent effect on struvite could be useful, allowing urine to be stored for longer, and thus used more efficiently with a minimal loss of ammonia.

One of the most pressing concerns associated with struvite recovered from urine pertains to the presence of pharmaceutical chemicals (Lienert et al., 2003). More work needs to be done to determine a) the quantities and types of chemicals present in stored urine, b) the degradation rate of pharmaceuticals and c) the potential of struvite as a sink for chemicals.

In terms of urine separation and its potential to achieve health and sanitation goals in underdeveloped areas, a comprehensive epidemiological study should be conducted to determine the direct health benefits of urine-separating facilities, with a focus on the workers who handle and treat the products.

To date, there are no provincial or national standards (building, plumbing, health) that address source-separating toilets. A comprehensive study of the regulations and the economics of source-separation in Canada would be the first study of its kind in Canada.

114 Since citizens of Vancouver do not currently have water meters, special attention should be paid to where and how the economic benefits of water-saving and/or nutrient recovery could be motivations for adoption.

A model of a urine-separation and struvite recovery system would be instrumental in determining how different parameters such as users, storage volume, transport distances, magnesium dosage, etc., would affect the feasibility of such a system in Canada.

An investigation into the predicted and presumed health effects of urine separation in single- family dwellings would be useful to better understand the conflict between different governing authorities (i.e. health, environment, building).

115 4.4 References

Amtul, Z., Atta-ur-Rahman, Siddiqui, R. A., Choudhary, M. I., 2002. Chemistry and mechanism of urease inhibition. Curr.Med.Chem., 9 (14), 1323-1348.

Amtul, Z., Rasheed, M, Choudhary, M. I., Rosanna, S., Khan, K. M., Atta-ur-Rahman, 2004. Kinetics of novel competitive inhibitors of urease enzymes by a focused library of oxadiazoles/thiadiazoles and triazoles. Biochem.Biophys.Res.Commun., 319 (3), 1053- 1063.

Boistelle, R., Abbona, F., Madsen, H. E. L., 1983. On the transformation of struvite into newberyite in aqueous systems. Phys. Chem. Miner., 9 (5), 216-222.

Booker, N., Priestley, A., Fraser, I., 1999. Struvite formation in wastewater treatment plants: Opportunities for nutrient recovery. Environ.Technol., 20 (7), 777-782.

Cooney, E. L., Booker, N. A., Shallcross, D. C, Stevens, G. W., 1999a. Ammonia removal from wastewaters using natural australian zeolite. I. characterization of the zeolite. Sep.Sci.Technol., 34 (12), 2307-2327.

Cooney, E. L., Booker, N. A., Shallcross, D. C, Stevens, G. W., 1999b. Ammonia removal from wastewaters using natural australian zeolite. II. pilot-scale study using continuous packed column process. Sep.Sci.Technol., 34 (14), 2741-2760.

Griffith, D. P., Gibson, J. R., Clinton, C. W., Musher, D. M., 1978. Acetohydroxamic acid - clinical studies of a urease inhibitor in patients with staghorn renal calculi. J.Urol., 119 (1), 9-15.

Jaffer, Y., Clark, T. A., Pearce, P., Parsons, S. A., 2002. Potential phosphorus recovery by struvite formation. Water Res., 36 (7), 1834-1842.

Morris, N. S., Stickler, D. J., 1998. The effect of urease inhibitors on the encrustation of urethral catheters. Urol.Res., 26 (4), 275-279.

Varel, V. H., Nienaber, J. A., Freetly, H. C, 1999. Conservation of nitrogen in cattle feedlot waste with urease inhibitors. J.Anim.Sci., 77 (5), 1162-1168.

116 APPENDIX I

117 Matrix set-up Part A

Sample # WW % Temperature Dilution Open/Closed Vol urine (mL) Vol WW 1 0 4 0 open 1000 0 2 5 4 0 open 950 50 3 10 4 0 open 900 100 4 25 4 0 open 750 250 5 0 4 0 closed 1000 0 6 5 4 0 closed 950 50 7 10 4 0 closed 900 100 8 25 4 0 closed 750 250 9 0 4 50 open 500 0 10 5 4 50 open 475 50 11 10 4 50 open 450 100 12 25 4 50 open 375 250 13 0 4 50 closed 500 0 14 5 4 50 closed 475 50 15 10 4 50 closed 450 100 16 25 4 50 closed 375 250 17 0 25 0 open 1000 0 18 5 25 0 open 950 50 19 10 25 0 open 900 100 20 25 25 0 open 750 250 21 0 25 0 closed 1000 0 22 5 25 0 closed 950 50 23 10 25 0 closed 900 100 24 25 25 0 closed 750 250 25 0 25 50 open 500 0 26 5 25 50 open 475 50 27 10 25 50 open 450 100 28 25 25 50 open 375 250 29 0 25 50 closed 500 0 30 5 25 50 closed 475 50 31 10 25 50 closed 450 100 32 25 25 50 closed 375 250

TOTAL VOL 21.6 1 3.2

Chapter 2 Data- Appendix I 118 Schematic of Experimental Set-up Part A

Each circle represents one 1 -L mixture. A black bar indicates a lid- the absence of a black bar indicates no lid. Fully shaded circles indicate that urine is full strength; half shaded circles indicate that urine is diluted by half.

Chapter 2 Data- Appendix I 119 pH Data Part A

Sample # 29-Aug 30-Aug 1-Sep 2-Sep 6-Sep 8-Sep 12-Sep 16-Sep 20-Sep 27-Sep 1 5.71 5.91 5.96 6.07 6.12 6.17 6.25 6.26 6.34 6.39 2 5.76 5.99 6.06 6.16 6.21 6.21 6.35 6.4 6.5 6.64 3 5.8 6.07 6.15 6.25 6.34 6.36 6.53 6.33 6.28 8.84 4 5.89 6.25 6.41 6.57 6.8 7.01 8.98 9.17 9.09 9.14 5 5.73 5.93 5.97 6.04 6.03 6.05 6.12 6.11 6.16 6.23 6 5.77 6.01 6.07 6.13 6.17 6.12 6.28 6.29 6.35 6.44 7 5.81 6.07 6.15 6.24 6.32 6.32 6.47 6.54 6.61 7.06 8 5.91 6.22 6.43 6.56 6.77 6.89 8.57 9.18 9.12 9.18 9 5.8 6.1 6.17 6.28 6.32 6.37 6.5 6.55 6.62 9.08 10 5.88 6.21 6.34 6.43 6.69 7.62 9.21 9.33 9.25 9.22 11 5.92 6.31 6.53 6.65 7.33 9.03 9.33 9.43 9.38 9.4 12 6.08 6.59 6.92 7.18 9.2 9.27 9.37 9.54 9.46 9.48 13 5.8 6.07 6.11 6.18 6.21 6.21 6.31 6.32 6.46 9.17 14 5.88 6.15 6.26 6.34 6.54 6.81 9.08 9.31 9.25 9.28 15 5.92 6.24 6.43 6.52 6.86 8.41 9.22 9.37 9.31 9.36 16 6.09 6.57 6.86 7.1 9.15 9.29 9.4 9.5 9.44 9.47 17 5.73 5.82 5.96 6.13 6.34 6.51 7.03 9.34 9.3 9.27 18 5.78 5.93 6.21 6.65 9.36 9.25 9.28 9.38 9.27 9.25 19 5.81 6.02 6.66 8.82 9.35 9.24 9.26 9.37 9.27 9.26 20 5.92 6.41 8.97 9.21 9.32 9.23 9.25 9.34 9.26 9.24 21 5.74 5.81 5.87 5.92 6.04 6.08 6.07 9.28 9.27 9.24 22 5.78 5.91 6.15 6.57 9.31 9.22 9.24 9.34 9.24 9.23 23 5.82 6.02 6.52 8.59 9.32 9.23 9.24 9.33 9.24 9.22 24 5.92 6.37 8.89 9.2 9.31 9.21 9.23 9.31 9.22 9.21 25 5.81 5.97 6.18 6.35 7.62 9.13 9.25 9.36 9.24 9.22 26 5.87 6.18 8.96 9.1 9.3 9.23 9.24 9.33 9.25 9.18 27 5.92 6.41 9.07 9.19 9.3 9.24 9.25 9.35 9.26 9.22 28 6.09 6.99 9.19 9.22 9.28 9.22 9.23 9.33 9.24 9.19 29 5.81 5.92 6.09 6.12 6.27 6.65 9.24 9.33 9.23 9.2 30 5.89 6.17 8.91 9.13 9.29 9.22 9.23 9.33 9.22 9.19 31 5.93 6.43 9 9.22 9.29 9.21 9.22 9.32 9.23 9.2 32 6.46 9.16 9.2 9.27 9.21 9.21 9.32 9.21 9.17

Chapter 2 Data- Appendix I Concentration Data Part A: Ammonia (mg NH3-N/L)

Sample # 29-Aug 1-Sep 6-Sep 13-Sep 20-Sep 27-Sep 1 262 369 266 214 263 302 2 239 232 275 224 236 338 3 228 233 269 232 288 857 4 199 231 324 703 893 1520 5 252 239 272 212 243 317 6 242 238 252 235 248 363 7 234 238 294 263 308 580 8 206 243 342 718 1110 1760 9 134 122 131 114 125 741 10 134 135 173 660 825 1250 11 120 143 261 830 1358 2670 12 107 159 545 938 2725 4550 13 128 127 129 108 139 1200 14 122 131 159 650 1060 1460 15 118 136 218 775 1128 2070 16 104 174 593 1153 2040 4290 17 254 237 283 847 10300 11200 18 250 267 5950 10500 9920 8870 19 234 379 6020 10100 9490 9780 20 202 1650 5340 8230 7870 7690 21 255 259 281 395 12100 11700 22 247 298 5890 11100 11600 10800 23 228 372 6080 10600 10800 11400 24 207 1500 5860 9180 9370 9320 25 127 134 292 5780 5430 5600 26 125 999 4490 5340 4880 4350 27 123 1260 4390 5020 4560 4730 28 114 26100 44600 4080 3800 3580 29 131 129 153 5910 6250 6180 30 124 1020 4590 5820 5810 5890 31 125 1290 4530 5600 5200 5730 32 111 2300 4030 4650 4800 4640

Chapter 2 Data- Appendix I Concentration Data Part A: Phosphorus (mg P04-P/L)

Sample # 29-Aug-05 1-Sep 6-Sep 1 626 593 635 2 622 562 658 3 538 550 573 4 445 440 479 5 610 605 637 6 573 579 571 7 546 543 596 8 464 441 449 9 310 304 301 10 302 303 286 11 263 262 250 12 217 211 157 13 299 297 93 14 276 275 280 15 253 253 257 16 209 212 156 17 614 573 523 18 598 519 426 19 545 471 398 20 452 303 323 21 616 590 582 22 594 572 421 23 537 509 409 24 463 316 348 25 296 281 259 26 282 233 200 27 269 203 195 28 227 1820 1900 29 304 293 283 30 274 216 198 31 275 193 198 32 219 160 160

Chapter 2 Data: Appendix I pH Data Part B Solution # Time Day Hour Cum. Time (hr) 1 2 3 4 5 6 7 8 10/19/05 16:00 0 5.88 5.91 5.93 6.02 5.93 5.98 6.02 6.19 10/20/05 10:00 20 10 0 5.89 5.93 5.97 6.12 5.92 6.02 6.1 6.33 10/20/05 11:00 20 11 1 5.68 5.7 5.74 5.86 5.66 5.73 5.8 6 10/21/05 15:00 21 15 29 5.71 5.79 5.83 6.02 5.82 5.93 6.02 10/22/05 14:30 22 14 52 5.80 5.96 6.02 6.59 5.91 6.44 6.91 8.15 10/23/05 13:30 23 13 75 5.91 6.2 6.4 8.27 6.03 8.44 8.69 8.79 10/24/05 14:00 24 14 100 6.06 7.25 7.8 8.63 6.15 8.88 8.95 8.93 10/25/05 9:30 25 9 119 6.18 8.43 8.72 9.01 6.20 8.95 9.06 9.03

Chapter 2 Data- Appendix I 123 Concentration Data Part B- Ammonia (mg NH3-N/L) Solution # Time Day Hour Cum. Time (hr) 1 2 3 4 5 6 7 8 10/19/05 16:00 0 241 243 229 192 122 125 117 105 10/20/05 10:00 20 10 0 252 242 232 194 127 120 113 103 10/20/05 11:00 20 11 1 241 231 220 183 122 120 111 99 10/21/05 15:00 21 15 29 238 224 219 200 119 123 117 110 10/22/05 14:30 22 14 52 241 234 231 255 118 144 177 229 10/23/05 13:30 23 13 75 238 255 268 506 124 379 520 565 10/24/05 14:00 24 14 100 248 382 477 1180 128 810 945 770

Chapter 2 Data- Appendix I 124 Concentration Data Part B- Phosphorus (mg P04-P/L) - Solution # Time Day Hour Cum. Time (hr) 1 2 3 4 5 6 7 8 10/19/05 16:00 0 271 272 254 210 121 125 117 98 10/20/05 10:00 20 10 0 289 274 259 204 135 122 116 95 10/20/05 11:00 20 11 1 271 255 242 193 123 120 107 90 10/21/05 15:00 21 15 29 267 248 235 201 124 121 no 91 10/22/05 14:30 22 14 52 267 250 241 194 123 116 111 38 10/23/05 13:30 23 13 75 257 251 233 20 120 22 16 16 10/24/05 14:00 24 14 100 253 170 29 16 110 15 15 15

Chapter 2 Data- Appendix I 125 Concentration Data Part B: Magnesium (mg/L) Solution # Time Dav Hour Cum. Time (hr) 1 2 3 4 5 6 7 8 10/19/05 16:00 0 62 58 59 48 7 29 30 24 10/20/05 10:00 20 10 0 56 57 58 47 29 28 32 25 10/20/05 11:00 20 11 1 1021 1097 822 849 860 1079 977 890 10/21/05 15:00 21 15 29 1099 1024 975 1077 964 979 891 820 10/22/05 14:30 22 14 52 1133 1083 1026 1205 820 951 859 747 10/23/05 13:30 23 13 75 979 1152 1018 765 991 818 797 734 10/24/05 14:00 24 14 100 999 916 828 693 867 759 758 688

Chapter 2 Data- Appendix I 126 pH Data Part C Solution # Time Day Hour Cum. Time fhr) 1 2 3 4 5 6 7 8 9 1/9/06 2:00 PM 9 14 0 5.65 5.71 5.78 5.69 5.79 5.94 5.74 5.91 6.14 1/10/06 10:00 AM 10 10 20 5.61 5.7 5.83 5.68 5.84 6.04 5.71 6.02 6.27 1/11/06 10:00 AM 11 10 44 5.62 5.78 5.95 5.71 6.02 6.32 5.86 6.24 6.66 1/12/06 3:00 PM 12 15 73 5.62 5.94 5.14 5.85 6.62 8.06 6.08 7.33 8.55 1/12/06 5:00 PM 12 17 75 8.66 9.28 1/12/06 6:30 PM 12 18 76 7.88 8.61 1/12/06 7:00 PM 12 19 77 7.89 8.59 1/13/06 10:00 AM 13 10 92 5.75 6.19 6.71 6.01 8.95 6.25 9.04 1/13/06 4:00 PM 13 16 98 8.92 9.01 1/13/06 4:30 PM 13 16 98 8.71 8.93 1/13/06 5:30 PM 13 17 99 8.7 8.87 1/14/06 2:00 PM 14 14 120 5.81 6.37 8.96 6.2 6.35 1/14/06 4:00 PM 14 16 122 8.97 1/14/06 5:00 PM 14 17 123 8.73 1/14/06 6:00 PM 14 18 124 8.71

1/15/06 3:00 PM 15 15 . 145 5.82 8.2 6.3 6.66 1/15/06 5:30 PM 15 17 147 8.25 1/15/06 6:30 PM 15 18 148 7.02 1/15/06 6:45 PM 15 18 148 7.08

1/16/06 12:00 PM 16 12 166 5.9 6.42 6.76

Chapter 2 Data- Appendix I 127 pH Data Part C Solution # Time Dav Hour Cum. Time fhr) 10 11 12 13 14 IS 16 17 18 1/9/06 2:00 PM 9 14 0 5.66 5.7 5.81 5.7 5.79 5.95 5.73 5.91 6.15 1/10/06 10:00 AM 10 10 20 5.63 5.7 5.83 5.68 5.85 6.04 5.72 6.03 6.28 1/11/06 10:00 AM 11 10 44 5.64 5.79 5.95 5.76 6.05 6.29 5.85 6.3 6.68 1/12/06 3:00 PM 12 15 73 5.67 5.95 6.15 5.88 6.62 7.17 6.07 7.27 8.65 1/12/06 5:00 PM 12 17 75 8.92 1/12/06 6:30 PM 12 18 76 8.72 1/12/06 7:00 PM 12 19 77 8.69 1/13/06 10:00 AM 13 10 92 5.77 6.17 6.28 6.05 9 9.01 6.25 9.04 1/13/06 4:00 PM 13 16 98 9.02 8.97 9.03 1/13/06 4:30 PM 13 16 98 8.81 8.74 8.91 1/13/06 5:30 PM 13 17 99 8.79 8.71 8.87 1/14/06 2:00 PM 14 14 120 5.84 6.39 8.96 6.16 6.33 1/14/06 4:00 PM 14 16 122 8.98 1/14/06 5:00 PM 14 17 123 8.66 1/14/06 6:00 PM 14 18 124 8.59 1/15/06 3:00 PM 15 15 145 5.9 8.29 6.27 6.48 1/15/06 5:30 PM 15 17 147 8.42 1/15/06 6:30 PM 15 18 148 6.91 1/15/06 6:45 PM 15 18 148 7.02 1/16/06 12:00 PM 16 12 166 5.98 6.37 7.85

Chapter 2 Data- Appendix I 128 Conductivity Data Part C (mS/cm) Solution # Time Day Hour Cum. time (hr) 1 2 3 4 ' 5 6 7 8 9 >'> *' - «i >,<5^ v-rtV£u"3' .,v,-... :io.9tf "- . '-9.S9 . .8 45 5 38 1- Wtf U 59 - ' -1 >J- '-15,62 •-.•.••:TS.92 .'v' ' •>• ';20 — "15.'82 «.*./• 10.-96 398 8 49 •^'•s:?8 \» ."4 47 1/11/06 10:00 AM ii 10 44 22.3 20.4 17.15 12.07 11.13 9.52 6.52 6.06 5.32 1/12/06 3:00 PM 12 15 73 22.4 20.6 17.17 12.1 11.67 10.98 6.59 6.74 6.43 1/12/06 5:00 PM 12 17 75 1/12/06 6:30 PM 12 18 76 1/12/06 7:00 PM 12 19 77 1/13/06 10:00 AM 13 10 92 22.4 20.7 18.05 12.07 15.23 6.66 9.25 1/13/06 4:00 PM 13 16 98 1/13/06 4:30 PM 13 16 98 1/13/06 5:30 PM 13 17 99 1/14/06 2:00 PM 14 14 120 22.4 19.88 24 12.27 6.6 1/14/06 4:00 PM 14 16 122 1/14/06 5:00 PM 14 17 123 1/14/06 6:00 PM 14 18 124 1/15/06 3:00 PM 15 15 145 21 23.7 12.28 6.88 1/15/06 5:30 PM 15 17 147 1/15/06 6:30 PM 15 18 148 1/15/06 6:45 PM 15 18 148 1/16/06 12:00 PM 16 12 166 22.5 12.42 6.51

IWWWWj^WWSSIiai * Samples taken with an Instrument that was Improperly calibrated

Chapter 2 Data- Appendix I Conductivity Data Part C (mS/cm) Solution # Time Hour Cum. time (hr) 10 11 12 13 14 15 16 17 18 if»06?-Q0+M ..,--=8^0 .20 3: .18.48 W*#psi8' • ••ifl.9 -••-9;9l 8.44 . 5 88 ! .'. /- -4.55 ' .'Aaa -'.-18.64 h#<¥M8}48" 15 82 . .-. 10.-84 . " 7.62 5 86 U% « \4,6 1/11/06 10:00 AM 11 10 44 22.4 20.4 17.53 12.02 11.05 9.46 6.5 6,01 5.16 1/12/06 3:00 PM 12 15 73 22.4 20.6 17.7 11.98 11.47 10.46 6.51 6.59 6.57 1/12/06 5:00 PM 12 17 75 1/12/06 6:30 PM 12 18 76 1/12/06 7:00 PM 12 19 77 1/13/06 10:00 AM 13 10 92 22 20.7 18.54 12.11 15.65 13.81 6.59 8.04 1/13/06 4:00 PM 13 16 98 1/13/06 4:30 PM 13 16 98 1/13/06 5:30 PM 13 17 99 1/14/06 2:00 PM 14 14 120 22.2 20.6 24.1 12.19 6.62 1/14/06 4:00 PM 14 16 122 1/14/06 5:00 PM 14 17 123 1/14/06 6:00 PM 14 18 124 1/15/06 3:00 PM 15 15 145 22.5 23.7 12.26 6.71 1/15/06 5:30 PM 15 17 147 1/15/06 6:30 PM 15 18 148 1/15/06 6:45 PM 15 18 148 1/16/06 12:00 PM 16 12 166 22.4 12.31 7.58

* Samples taken with an Instrum : Improperly calibrated

Chapter 2 Data- Appendix I Concentration Data Part C: Ammonia (mg NH3-N/L) Solution # Time Dav Hour Cum. time (hr) 1 2 3 4 5 6 7 8 9 1/9/06 2 00 PM 9 14 0 317 278 239 162 144 128 79.3 75.1 69.5 1/10/06 10 00 AM 10 10 20 313 292 246 159 145 125 85.3 76.9 72.9 1/11/06 10 00 AM 11 10 44 366 293 250 160 150 141 83.3 83 93.9 1/12/06 3 00 PM 12 15 73 310 305 268 165 232 369 87.9 178 276 1/12/06 5 00 PM 12 17 75 512 375 1/12/06 6 30 PM 12 18 76 428 343 1/12/06 7 00 PM 12 19 77 402 352 1/13/06 10 00 AM 13 10 92 303 300 374 158 979 94.3 696 1/13/06 4 00 PM 13 16 98 1070 811 : 1/13/06 4 30 PM 13 16 98 967 707 1/13/06 5 30 PM 13 17 99 942 699 1/14/06 2 00 PM 14 14 120 293 341 1710 168 92.4 1/14/06 4 00 PM 14 16 122 1730 1/14/06 5 00 PM 14 17 123 1500 1/14/06 6 00 PM 14 18 124 1430 1/15/06 3 00 PM 15 IS 145 321 834 185 121 1/15/06 5 30 PM 15 17 147 893 1/15/06 6 30 PM 15 18 148 751 1/15/06 6 45 PM 15 18 148 718 1/16/06 12 00 PM 16 12 166 313 191 143

Chapter 2 Data- Appendix 1 Concentration Data Part C: Ammonia (mg NH3-N/L) Solution # Time Dav Hour Cum. time (hr) 10 11 12 13 14 15 16 17 18 1/9/06 2:00 PM 9 14 0 312 286 241 155 152 132 88.5 75.8 68.8 1/10/06 10:00 AM 10 10 20 322 298 243 163 151 128 86.4 77.5 70.8 1/11/06 10:00 AM 11 10 44 338 310 255 159 157 138 79.5 82.6 93.5 1/12/06 3:00 PM 12 15 73 320 303 273 166 235 276 87.6 162 317 1/12/06 5:00 PM 12 17 75 497 1/12/06 6:30 PM 12 18 76 425 1/12/06 7:00 PM 12 19 77 416 1/13/06 10:00 AM 13 10 92 302 294 408 162 1160 1070 83.5 694 1/13/06 4:00 PM 13 16 98 1360 1080 818 1/13/06 4:30 PM 13 16 98 1200 924 728 1/13/06 5:30 PM 13 17 99 1160 926 735 1/14/06 2:00 PM 14 14 120 284 361 1740 170 81.3 1/14/06 4:00 PM 14 16 122 1720 1/14/06 5:00 PM 14 17 123 1480 1/14/06 6:00 PM 14 18 124 1450 1/15/06 3:00 PM 15 15 145 289 920 174 111 1/15/06 5:30 PM 15 17 147 947 1/15/06 6:30 PM 15 18 148 788 1/15/06 6:45 PM 15 18 148 791 1/16/06 12:00 PM 16 12 166 301 195 221

Chapter 2 Data- Appendix I Concentration Data Part C: Phosphorus (mg P04-P/L) Solution # Time Day Hour Cum. time (hr) 1 2 3 4 5 6 7 8 9 1/9/06 2:00 PM 9 14 0 728 647 541 362 313 266 172 150 126 1/10/06 10:00 AM 10 10 20 714 657 527 358 317 254 177 152 122 1/11/06 10:00 AM 11 10 44 811 652 535 361 311 256 178 151 122 1/12/06 3:00 PM 12 15 73 705 654 552 362 313 225 166 148 95.6 1/12/06 5:00 PM 12 17 75 191 91.1 1/12/06 6:30 PM 12 18 76 21 13.9 1/12/06 7:00 PM 12 19 77 17.2 13.3 1/13/06 10:00 AM 13 10 92 657 596 492 338 221 168 111 1/13/06 4:00 PM 13 16 98 207 103 1/13/06 4:30 PM 13 16 98 11.6 17.6 1/13/06 5:30 PM 13 17 99 9.6 12.5 1/14/06 2:00 PM 14 14 120 613 561 344 334 161 1/14/06 4:00 PM 14 16 122 343 1/14/06 5:00 PM 14 17 123 8.61 1/14/06 6:00 PM 14 18 124 16.3 1/15/06 3:00 PM 15 15 145 630 421 329 160 1/15/06 5:30 PM 15 17 147 400 1/15/06 6:30 PM 15 18 148 149 1/15/06 6:45 PM 15 18 148 101

Chapter 2 Data- Appendix I Concentration Data Part C: Phosphorus (mg P04-P /L) Solution # Time Day Hour Cum. time (hr) 10 11 12 13 14 15 16 17 18 1/9/06 2:00 PM 9 14 0 706 640 532 350 328 273 180 155 130 1/10/06 10:00 AM 10 10 20 724 661 519 361 324 258 178 148 119 1/11/06 10:00 AM 11 10 44 751 681 535 350 326 252 172 154 120 1/12/06 3:00 PM 12 15 73 714 648 542 360 308 259 171 153 92.5 1/12/06 5:00 PM 12 17 75 87.4 1/12/06 6:30 PM 12 18 76 10.1 1/12/06 7:00 PM 12 19 77 9.57 1/13/06 10:00 AM 13 10 92 654 581 493 344 216 180 169 107 1/13/06 4:00 PM 13 16 98 199 168 100 1/13/06 4:30 PM 13 16 98 8.02 9.24 9.81 1/13/06 5:30 PM 13 17 99 9.59 10.2 10.4 1/14/06 2:00 PM 14 14 120 640 570 349 329 175 1/14/06 4:00 PM 14 16 122 338 1/14/06 5:00 PM 14 17 123 9.98 1/14/06 6:00 PM 14 18 124 10.4 1/15/06 3:00 PM 15 15 145 631 421 332 179 1/15/06 5:30 PM 15 17 147 393 1/15/06 6:30 PM 15 18 148 49.4 1/15/06 6:45 PM 15 18 148 28.8

Chapter 2 Data- Appendix I Concentration Data Part C: Magnesium (mg/L) Solution # Time Day Hour Cum. time fhr) 1 2 3 4 5 6 7 8 9 1/9/06 2:00 PM 9 14 0 79.58 66.15 58.93 39.66 35.35 27.70 18.30 14.64 13.96 1/10/06 10:00 AM 10 10 20 1/11/06 10:00 AM 11 10 44 80.98 59.02 47.91 30.82 26.50 23.65 16.15 13.88 9.72 1/12/06 3:00 PM 12 15 73 72.07 63.13 48.93 29.52 25.86 13.42 11.59 8.61 4.32 1/12/06 5:00 PM 12 17 75 5.78 2.28 1/12/06 6:30 PM 12 IS 76 94.37 42.55 1/12/06 7:00 PM 12 19 77 89.69 42.20 1/13/06 10:00 AM 13 10 92 80.52 68.40 226.99 34.77 0.00 16.44 0.00 1/13/06 4:00 PM 13 16 98 0.25 0.40 1/13/06 4:30 PM 13 16 98 56.86 47.08 1/13/06 5:30 PM 13 17 99 57.19 40.44 1/14/06 2:00 PM 14 14 120 75.70 66.73 1.18 36.19 16.37 1/14/06 4:00 PM 14 16 122 0.39 1/14/06 5:00 PM 14 17 123 208.40 1/14/06 6:00 PM 14 18 124 212.56 1/15/06 3:00 PM 15 15 145 87.58 9.69 34.43 15.91 1/15/06 5:30 PM 15 17 147 1.17 1/15/06 6:30 PM 15 18 148 269.06 1/15/06 6:45 PM 15 18 148 237.43

Chapter 2 Data- Appendix 1 Solution # Time Dav Hour Cum. time (hr) 10 11 12 13 14 15 16 17 18 1/9/06 2:00 PM 9 14 0 72.19 60.12 50.27 31.72 30.35 25.08 15.58 12.21 6.18 1/10/06 10:00 AM 10 10 20 1/11/06 10:00 AM 11 10 44 69.78 63.70 49.24 29.70 22.37 19.48 10.30 6.64 5.44 1/12/06 3:00 PM 12 15 73 70.25 63.52 53.08 32.31 29.16 21.39 12.25 8.89 3.29 1/12/06 5:00 PM 12 17 75 2.01 1/12/06 6:30 PM 12 18 76 101.36 1/12/06 7:00 PM 12 19 77 101.22 1/13/06 10:00 AM 13 10 92 76.45 69.54 55.47 36.14 1.05 0.75 14.96 0.78 1/13/06 4:00 PM 13 16 98 0.09 0.25 0.50 1/13/06 4:30 PM 13 16 98 233.18 299.57 122.27 1/13/06 5:30 PM 13 17 99 236.83 291.48 130.69 1/14/06 2:00 PM 14 14 120 73.03 68.17 2.73 35.60 18.36 1/14/06 4:00 PM 14 16 122 0.00 1/14/06 5:00 PM 14 17 123 634.42 1/14/06 6:00 PM 14 18 124 662.63 1/15/06 3:00 PM 15 15 145 73.99 0.72 35.94 19.71 1/15/06 5:30 PM 15 17 147 0.00 1/15/06 6:30 PM 15 18 148 669.34 1/15/06 6:45 PM 15 18 148 667.74

Chapter 2 Data- Appendix I 136 Concentration Data Part C: Calcium (mg /L) Solution # Time Dav Hour Cum. time (hr) 1 2 3 4 5 6 7 8 9 1/9/06 2:00 PM 9 14 0 147.5 133.2 105.5 78.6 69.7 52.9 42.5 33.8 39.1 1/10/06 10:00 AM 10 10 20 1/11/06 10:00 AM 11 10 44 180.1 143.3 115.8 79.2 70.4 62.1 44.0 11.7 23.3 1/12/06 3:00 PM 12 15 73 149.6 129.9 108.0 72.4 61.7 19.0 33.4 26.6 9.3 1/12/06 5:00 PM 12 17 75 2.1 10.5 1/12/06 6:30 PM 12 18 76 9.3 10.5 1/12/06 7:00 PM 12 19 77 8.4 10.8 1/13/06 10:00 AM 13 10 92 152.0 139.4 124.5 81.6 3.8 39.9 3.7 1/13/06 4:00 PM 13 16 98 2.9 3.1 1/13/06 4:30 PM 13 16 98 2.7 1.6 1/13/06 5:30 PM 13 17 99 2.5 3.1 1/14/06 2:00 PM 14 14 120 146.4 118.7 6.0 80.2 44.9 1/14/06 4:00 PM 14 16 122 9.3 1/14/06 5:00 PM 14 17 123 9.3 1/14/06 6:00 PM 14 18 124 8.6 1/15/06 3:00 PM 15 15 145 138.7 21.4 77.7 42.5 1/15/06 5:30 PM 15 17 147 18.3 1/15/06 6:30 PM 15 18 148 20.4 1/15/06 6:45 PM 15 18 148 18.8

Chapter 2 Data- Appendix I 137 Concentration Data Part C: Calcium (mg /L)

17 18 Time Dav Hour Cum. time (hr) 10 11 12 13 14 15 16 1/9/06 2:00 PM 9 14 0 157.5 140.2 118.8 80.9 71.9 60.8 40.8 34.9 23.1 1/10/06 10:00 AM 10 10 20 1/11/06 10:00 AM 11 10 44 148.1 139.9 118.5 77.2 62.1 54.0 29.5 21.0 13.2 1/12/06 3:00 PM 12 15 73 136.8 131.6 110.9 70.8 63.3 51.0 34.8 25.7 10.2 1/12/06 5:00 PM 12 17 75 9.3 1/12/06 6:30 PM 12 18 76 9.0 1/12/06 7:00 PM 12 19 77 9.8 1/13/06 10:00 AM 13 10 92 148.7 135.3 110.4 79.0 5.6 6.0 42.1 5.9 1/13/06 4:00 PM 13 16 98 3.2 3.6 3.8 1/13/06 4:30 PM 13 16 98 3.7 4.3 2.7 1/13/06 5:30 PM 13 17 99 4.1 4.6 2.4 1/14/06 2:00 PM 14 14 120 150.2 116.8 7.5 86.3 44.1 1/14/06 4:00 PM 14 16 122 7.6 1/14/06 5:00 PM 14 17 123 7.5 1/14/06 6:00 PM 14 18 124 7.7 1/15/06 3:00 PM 15 15 145 138.8 21.8 81.2 46.5 1/15/06 5:30 PM 15 17 147 16.3 1/15/06 6:30 PM 15 18 148 17.4 1/15/06 6:45 PM 15 18 148 17.6

Chapter 2 Data- Appendix I 138 Part C Calculations Solution # • 5 9.8 A PRE-PRECIP AVERAGE 628 529 314 259 150 123 B POST-Ca Precip AVERAGE 411 344 214 208 107 93 C Post Struvite Precip Average 125 12 11 19 15 14 «, «•„'•> * * ^ A-B ' * vii % 4**.«6fen /9^217 v . ^100 51 a >u w43 30 f mass difference »-C/ 'iJf %?#AV " 503 517 i *• '303 * '^240 , * *\ *s V135 * 110 .,^14 ^ •> B-C>* > , < * 331 A ;^203 ri i> * tl89 \'V t 92 80 % removed In Ca precip 35 35 32 20 29 24 % difference % removed in struvite precip 45 63 65 73 61 65 % remaining 20 2 3 7 10 11

CALCIUM 2 3 1 5 6 *«^l«*8 A PRE-PRECIP AVERAGE 133 113 67 58 24 31 B POST-Ca Precip AVERAGE 20 8 3 11 3 10 C Post Struvite Precip Average 20 9 3 9 2 11

l 1 f ' " •. ' A-B '"*' * „\ rt "'^V . \ 113 106 i" *Srf« «t64 21 * ^< r^2i mass difference j A-C «/ v i-Kia „'^104 >. V , <* -*H65 * ,>W49 22 B-C - * , > - 0 -1 1 •"•n \ > 1 % removed In Ca precip 85 93 95 82 86 68 °/o difference % removed in struvite precip 0 -1 1 3 4 -2 % remaining 15 8 4 15 10 34

Chapter 2 Data- Appendix I Part C Calculations Solution # il?(fil@SliVliCi RUS * 11 .. " . * .12 a* <# 15 A PRE-PRECIP AVERAGE 630 524 322 261 153 123 B POST-Ca Precip AVERAGE 407 344 208 174 104 90 C - Post Struvite Precip Average 39 10 9 10 10 10 >' t t «v * V\W H4« £4-# ".is* < K/ • '.j.^-5 -,.1^591' tw"3i3 ^251 -/'n,l42 V.2> "-MU13 H L •* 1 A.M 14368 ^ * > 333 •164 ' > 93 B-C^ W 'W*# *» *,"*«i*lr % removed In Ca precip 35 34 35 33 32 27 % difference % removed In struvite precip 58 64 62 63 61 65 % remaining 6 2 3 4 7 8

l l W-f'^v.'-i '•»*• „•'.-, [CALCIUM ."!ff:> V " '• * -.\ .- «• 12 >»*•>. tVWl5 . 17 / 18 A PRE-PRECIP AVERAGE 133 115 66 55 27 18 B POST-Ca Precip AVERAGE 19 8 4 5 5 10 C Post Struvite Precip Average 18 8 4 4 3 9

114 107 ' « „ ^**61 V <* 8 ArB> < 'Vlf.^n. fi V*V -^T f v»* %# «, 50 - 1 mass.difference age c 'i« 115 62 51 2 **• .? -1 r "0 11 2 BSC -1 "'W/? * - % removed In Ca precip 86 93 93 91 82 46 °/o difference % removed in struvite precip 1 0 1 1 8 2 % remaining 13 7 6 8 9 52

Chapter 2 Data- Appendix I Crystal Analysis Part C

Solution # Mass of crystals (g) [struvite]- mg/L. Ca ( mg/L) % Ca Purity (%) 2 0.0997 997 0.381 0.04 99.96 3 0.0996 996 0.280 0.03 99.97 5 0.0998 998 0.536 0.05 99.95 6 0.1004 1004 1.178 0.12 99.88 8 0.1005 1005 0.550 0.05 99.95 11 0.1002 1002 0.211 0.02 99.98 12 0.0999 999 0.090 0.01 99.99 15 0.1007 1007 -0.035 N.D >99.98 16 0.1004 1004 0.361 0.04 99.96 17 0.0998 998 -0.315 N.D >99.98 18 0.1007 1007 -0.451 N.D >99.98

Chapter 2 Data- Appendix I 141 Part C Crystal Analysis- Solution #5

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——•—i— r~n i—i—i—r i—i—i—r i—i—|—i—i—i— 10 20 30 40 50 60 7C 2-Theta-Scale

0 0 0 fflRle #5raw- Tvoe 2TrrTh locked- Slat 500 ° - End: 70.00 - 9ec: 0.02 - Steo time Q1 s- Term: 25 °C (Poem) - Time Started: 1133820736 s- 2-Theta 5.00 ° - Theta 250 ° - Chi: 0.00 - Phi: 0.00 BoO-015-0762 (*) - Struvite. svn- |vlH4(vtPC4-6H20- Y: 50.91 %- d xbv: 1. - WL 1.5406- (^hcrhorbic-a 6.94500- b 11.20800- c 6.13550- alcha 90.000 - beta 90.000 -rjarrma 90.000- Rimtive - P

142 Part C Crystal Analysis- Solution #6

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^ ilJ • I I—i—i—i—i—m—l—| i i—r 6 10 20 2-Theta-Scale BlRle #6raw- Tvoe 2WTh locked - Start 5.00 ° - End: 70.00 ° - Steo: 0.02 ° - Steo tirre 0.1 s - Term: 25 °C (Rxm) - Tims Started: 1138821632 s - 2-Theta 5.00 ° - Theta 250 0 - Chi: 0.00 ° - Phi: 0.00 ®00-01S0762 (*) - Struvite. svn - NhWfvtPC4-6H20- Y: 50.91 %- d x bv: 1. - WL 1.5406 - Qthcrhcrrbic - a 6.94500 - b 11.20800 - c 6.13550 - al dna 90.000 - beta 90.000 - oarrma 90.000 - R-irritive - P

143 Part C Crystal Analysis- Solution #8

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• •

—i—i—i—i—i—i——i—i—i—i—r—r 1—p*1—I—r 1 1 1 II i" i -|- i—i—rn—|—1—i—i—i—|—r ' i I T I 10 20 30 40 50 60 7( 2-Theta-Scale B3Rle #araw- Tvoe 2TrrTh locked - Start 5000 - Erri: 70.000 - Steo: 0.020 - Steo tirre 0.1 s- Term: 25 °C (Fbcrrt - Tirre Sbrted: 1138823296 s- 2-Theta 5.00 ° - Theta 250 ° - Chi: 0.00 ° - Fhi: 0.00 Hoacri5-rtf62n-Stnj\«te.sv^ %-dxbv:1.-VvL 1.5406-OthcrrKTtt»ic-a6S4500-b 11.20800-c 6.13550-aldia90.0rx)-beta 90.m0-oaiTrra 90.000-Prim^

144 Part C Crystal Analysis- Solution #9

•

r "i —|—f 1 r -i—r I 1 1 ' II -|—i—i—i—i—|—ri—i—i—|—i i—r — — ^ I 7( 6 10 20 40 50 60 2-Theta-Scale Stale: #9.raw- Tvoec 2ThrTh locked - Start 5.00 ° - End: 70.00 ° - Seo: 0.02 ° - Steo time 0.1 s- Term: 25 °C (Poem) - Time Started: 1138822912 s- 2-Theta 5.00 ° - Theta 250 ° - Chi: 0.00 ° - Fhi: 0.00 EoO-01SC762n-Sru\^ %-dxbv: 1.-WL 1.54C6-Crthcrhcrnbic-a6.94500-b11.20c^^

145 2300 2200 2100 Part C Crystal Analysis- Solution #14 2000 1900 1800 1700 1600 1500 1400 1300 1200 1100 1000 900 800 700 600 500 400 ' 300 • 200 100 •

0 • I I I I 10 20 30 40 50 60 7C

2-Theta-Scale Etals #14.raw-TvDe 2Th/Th locked- Sart: 5.00 ° - End: 70.00 ° - Steo: 0.02 ° - Sec time: 0.1 s - Term: 25 °C(Rocrrt- Tirre Started: 1138822528 s - 2-Tneta: 5.000 - Theta 2.50 ° - Oi: 0.000 - PH: 0.0 HOO-OI5-0762 C) - Struvite. svn - NH4fvtPC4'6H20- Y: 50.91 %-d xbv: 1.-WL 1.5406 - Crthcmorrbic - a 6.94500 - b 11.20800 - c 6.13550 - a! cha 90.000 - beta 90.000 - aarrrra 90.000 - Prirritrve - P

146 1900

1800

1700 ~i Part C Crystal Analysis- Solution #15

1600

1500

1400

1300

1200 -

1100 -

1000 -

900 -

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I I i i i i ^ i i i i | rn r—i | n i i |—i—i—i—i—y^-i—i—i—i—r-n—i—i—r-"i

10 40 50 60 2-Theta-Scale EJFile #15.rav-Tvrje: 2Th/Th locked - Start: 5.00" - End: 70.00 ° - Stea 0.020 - Steo time 0.1 s - Term: 25 °C( Perm)-Tirre Started: 1138822016 s - 2-Theta: 5.00 ° - Theta 2.50° - Ch: 0.00 ° - Ph: 0.0 llJcO-015-0762n-aruvte.svn-NH4lv^^ 1.5406-Qthcmcrttjic-a 6.94500-b11.20800-c 6.13550-a1*^

147 Part C Crystal Analysis- Solution #17

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10 20 30 40 50 60 2-Theta-Scale

E3file #17.raw-Tvce: 2Th/Th locked - Sat 5.00 ° - End 70.00 ° - Stea 0.02° - SeD time 0.1 s - Term.: 25 "Cf Peon)-Time Started: 1138820052 s - 2-Theta: 5.000 - Theta 2.500 - Or: 0.00 °- PH: 0.0 l!ba01S0762n-aruvite.svn-r^^ 1.54fJ6-Crthcrhcritiic-a6.94500-b11.2DcrjO-c6.13550-a

148 Part C Crystal Analysis- Solution #18

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T—r—|—i—i I—r 10 20 30 40 2-Theta-Scale Efele: #18.raw-Tvrje: ZTh/Tri locked - Start: 5.00 ° - End 70.00 °- Stea 0.02 ° - Seo time: 0.1 s - Term: 25 °Cf RocrnV Time aarted: 1138821120 s - 2-Tneta: 5.000 - Theta 2.50 ° - Ch: 0.00 ° - Pri: 0.0 ®TXKri50762n-aru\ite.Svn-N^ %-dxbv: 1.-WL 1.5406-Crthcmorrbic-a 6.94^

149 APPENDIX II

150 pH data Solution # Time Dav Hour Cum. time (hr) 1 2 3 4 5 6 7 8 9 3/29/06 5:00 PM 29 17 0 6.36 6.41 6.45 6.45 6.49 6.55 6.51 6.57 6.66 3/30/06 12:30 PM 30 12 19 6.01 6.11 6.25 6.12 6.17 6.43 6.24 6.37 7.11 3/31/06 12:30 PM 31 12 43 5.93 6.81 7 6.08 6.76 8.14 6.16 8.12 8.95 3/31/06 12:30 PM 31 12 43 8.53 8.49 7.82 3/31/06 1:30 PM 31 13 44 8.15 8.28 8.97 3/31/06 2:30 PM 31 14 45 8.19 8.27 9.38 4/1/06 1:30 PM 1 13 68 5.92 8.64 8.62 6.37 8.58 6.26 4/1/06 2:30 PM 1 14 69 8.65 8.65 8,65 4/1/06 3:30 PM 1 15 70 8.49 8.43 8.43 4/1/06 4:30 PM 1 16 71 8.47 8.49 8.5 4/2/06 4:00 PM 2 16 95 6.57 8.61 6.79 4/2/06 5:00 PM 2 17 96 8.68 4/2/06 6:00 PM 2 18 97 8.52 4/2/06 7:00 PM 2 19 98 8.53 4/3/06 3:00 PM 3 15 118 8.68 7.69 4/3/06 4:00 PM 3 16 119 8.73 4/3/06 5:00 PM 3 17 120 8.37 4/3/06 6:00 PM 3 18 121 8.41 4/4/06 11:30 AM 4 11 138 8.25 4/4/06 12:30 PM 4 12 139 8.31 4/4/06 1:30 PM 4 13 140 7.85 4/4/06 2:30 PM 4 14 141 7.87 4/5/06 10:30 AM 5 10 161

l> -• -1 missed sample- replaced with value of pair

Chapter 3 Data- Appendix II pH data Solution # Time Dav Hour Cum. time (hr) 10 11 12 13 14 15 16 17 18 3/29/06 S 00 PM 29 17 0 6.38 6.39 6.45 6.46 6.49 6.57 6.53 6.59 3/30/06 12 30 PM 30 12 19 6.07 6.08 6.25 6.14 6.17 6.47 6.28 6.36 7.14 3/31/06 12 30 PM 31 12 43 5.89 6.92 7.17 6.04 6.86 8.12 6.06 8.01 9 3/31/06 12 30 PM 31 12 43 7 3/31/06 1 30 PM 31 13 44 7.49 3/31/06 2 30 PM 31 14 45 4/1/06 1 30 PM 1 13 68 5.8 8.64 5.96 8.6 8.9 6.03 8.9 8.92 4/1/06 2 30 PM 1 14 69 8.72 6.11 4/1/06 3 30 PM 1 15 70 8.6 6.22 4/1/06 4 30 PM 1 16 71 8.69 4/2/06 4 00 PM 2 16 95 5.96 6.08 8.97 9.04 8.98 8.91 4/2/06 S 00 PM 2 17 96 4/2/06 6 00 PM 2 18 97 4/2/06 7 00 PM 2 19 98 4/3/06 3 00 PM 3 15 118 6.65 6.25 9.05 9.02 8.99 8.89 4/3/06 4 00 PM 3 16 119 4/3/06 5 00 PM 3 17 120 4/3/06 6 00 PM 3 18 121 4/4/06 11 30 AM 4 11 138 8.16 6.55 9.05 9.01 8.96 8.86 4/4/06 12 30 PM 4 12 139 4/4/06 1 30 PM 4 13 140 4/4/06 2 30 PM 4 14 141 4/5/06 10 30 AM 5 10 161 8.65 8.66 9.05 8.98 8.94 8.86

I "' Jj missed sample- replaced with value of pair

Chapter 3 Data- Appendix II Conductivity Data (mS/cm)

Time Day Hour Cum. time (hr) 1 2 3 4 5 6 7 8 9 3/29/06 5:00 PM 29 17 0 15.61 14.26 12.21 8.33 7.65 6.55 4.37 4.04 3.54 3/30/06 12:30 PM 30 12 19 15.87 14.51 12.65 8.41 7.79 7 4.4 4.18 4.29 3/31/06 12:30 PM 31 12 43 11.24 11.61 15.33 8.58 9.08 8.37 4.53 5.99 8.26 3/31/06 12:30 PM 31 12 43 10.23 6.37 9.19 3/31/06 1:30 PM 31 13 44 9.69 6.76 9.4 3/31/06 2:30 PM 31 14 45 11.05 6.76 9 4/1/06 1:30 PM 1 13 68 15.96 22.6 19.92 9.11 12.37 4.63 4/1/06 2:30 PM 1 14 69 22.9 20.3 12.68 4/1/06 3:30 PM 1 15 70 22.5 21.8 13.44 4/1/06 4:30 PM 1 16 71 24 20.8 13.47 4/2/06 4:00 PM 2 16 95 17.15 13.11 5.03 4/2/06 5:00 PM 2 17 96 13.29 4/2/06 6:00 PM 2 18 97 14.16 4/2/06 7:00 PM 2 19 98 14.15 4/3/06 3:00 PM 3 15 118 23 5.73 4/3/06 4:00 PM 3 16 119 23.1 4/3/06 5:00 PM 3 17 120 24.3 4/3/06 6:00 PM 3 18 121 24.4 4/4/06 11:30 AM 4 11 138 6.32 4/4/06 12:30 PM 4 12 139 6.33 4/4/06 1:30 PM 4 13 140 6.7 4/4/06 2:30 PM 4 14 141 6.63 4/5/06 10:30 AM 5 10 161

Issed sample- replaced with value of pali

Chapter 3 Data- Appendix II 153 conductivity Data (ms/cm; Solution # 16 17 18 Time Dav Hour Cum. time Thr) 10 11 12 13 14 15 6.54 4.32 4.06 3/29/06 5:00 PM 29 17 0 15.6 14.24 12.17 8.32 7.62 4.25 3/30/06 12:30 PM 30 12 19 15.9 14.53 12.65 8.43 7.83 7.03 4.35 4.28 3/31/06 12:30 PM 31 12 43 16.03 17.34 15.46 8.53 9.13 9.74 4.4 5.9 8.79 3/31/06 12:30 PM 31 12 43 17.93 3/31/06 1:30 PM 31 13 44 17.97 3/31/06 2:30 PM 31 14 45 9.33 4/1/06 1:30 PM 1 13 68 15.95 22.9 8.5 12.69 14.45 4.42 9.4 4/1/06 2:30 PM 1 14 69 23.6 5.18 4/1/06 3:30 PM 1 15 70 23 5.18 4/1/06 4:30 PM 1 16 71 22.3 10.87 9.64 4/2/06 4:00 PM 2 16 95 16.02 8.55 16.66 17.33 4/2/06 5:00 PM 2 17 96 4/2/06 6:00 PM 2 18 97 4/2/06 7:00 PM 2 19 98 9.92 4/3/06 3:00 PM 3 15 118 17.51 8.74 19.86 12.07 11.16 4/3/06 4:00 PM 3 16 119 4/3/06 5:00 PM 3 17 120 4/3/06 6:00 PM 3 18 121 4/4/06 11:30 AM 4 11 138 20.9 9.23 20.5 17.83 11.37 10.04 4/4/06 12:30 PM 4 12 139 4/4/06 1:30 PM 4 13 140 4/4/06 2:30 PM 4 14 141 4/5/06 10:30 AM 5 10 161 23.1 12.24 20.5 17.89 11.46 10.1

I .I missed sample- replaced with value of pair

Chapter 3 Data- Appendix II 154 Concentration data: NH3-N (mg/L) Solution # Time Dav Hour Cum. time Thr) 1 2 3 4 5 6 7 8 9 3/29/06 5:00 PM 29 17 0 298 260 222 144 134 120 75.9 76.1 67.3 3/30/06 12:30 PM 30 12 19 295 289 264 151 148 161 77.9 89 152 3/31/06 12:30 PM 31 12 43 314 639 670 179 336 554 99.3 336 974 3/31/06 12:30 PM 31 12 43 607 411 1050 3/31/06 1:30 PM 31 13 44 347 1000 3/31/06 2:30 PM 31 14 45 348 863 4/1/06 1:30 PM 1 13 68 305 1690 1450 254 782 114 4/1/06 2:30 PM 1 14 69 •&«*.'. >f«i32" 4/1/06 3:30 PM 1 15 70 3 93 ~* , 3 68 »>«,HWV43 4/1/06 4:30 PM 1 16 71 3 85 « . .* !3 53 4/2/06 4:00 PM 2 16 95 555 973 162 4/2/06 5:00 PM 2 17 96 996 4/2/06 6:00 PM 2 18 97 1080 4/2/06 7:00 PM 2 19 98 912 4/3/06 3:00 PM 3 15 118 1710 273 4/3/06 4:00 PM 3 16 119 1690 4/3/06 5:00 PM 3 17 120 1340 4/3/06 6:00 PM 3 18 121 1310 4/4/06 11:30 AM 4 11 138 331 4/4/06 12:30 PM 4 12 139 328 4/4/06 1:30 PM 4 13 140 278 4/4/06 2:30 PM 4 14 141 274 4/5/06 10:30 AM 5 10 161

S^SKJBSre-runqsampte

Chapter 3 data- Appendix II Concentration data: NH3-N (mg/L) Solution #

Time Day Hour Cum. time (hr) 10 11 12 .13 14 15 16 17 18 3/29/06 5:00 PM 29 17 0 280 268 226 151 140 120 75.5 73.4 67 3/30/06 12:30 PM 30 12 19 292 278 272 156 152 169 76.3 91.6 162 3/31/06 12:30 PM 31 - 12 43 314 779 721 161 343 596 87.6 327 1090 3/31/06 12:30 PM 31 12 43 669 3/31/06 1:30 PM 31 13 44 640 3/31/06 2:30 PM 31 14 45 4/1/06 1:30 PM 1 13 68 280 1710 179 879 1660 128 944 982 4/1/06 2:30 PM 1 14 69 •W*'ffi0uO 90.2 4/1/06 3:30 PM 1 15 70 wimmtm 76.9 4/1/06 4:30 PM 1 16 71 4/2/06 4:00 PM 2 16 95 329 177 1910 2290 1350 1100 4/2/06 5:00 PM 2 17 96 4/2/06 6:00 PM 2 18 97 4/2/06 7:00 PM 2 19 98 4/3/06 3:00 PM 3 15 118 515 183 2610 2280 1370 1090 4/3/06 4:00 PM 3 16 119 4/3/06 5:00 PM 3 17 120 4/3/06 6:00 PM 3 18 121 4/4/06 11:30 AM 4 11 138 1050 250 2560 2150 1310 1130 4/4/06 12:30 PM 4 12 139 4/4/06 1:30 PM 4 13 140 4/4/06 2:30 PM 4 14 141 4/5/06 10:30 AM 5 10 161 1730 757 2580 2200 1320 1060

Chapter 3 data- Appendix II Concentration Data: P04-P (mg/L) Solution #

Time Dav Hour Cum. time (hr) 1 2 3 4 5 6 7 8 9 3/29/06 5:00 PM 29 17 0 451 394 333 216 194 162 105 96.8 81.7 3/30/06 12:30 PM 30 12 19 441 405 318 220 194 162 107 93.7 78 3/31/06 12:30 PM 31 12 43 462 397 327 220 198 131 103 87.6 57.9 3/31/06 12:30 PM 31 12 43 116 77.1 50.8 3/31/06 1:30 PM 31 13 44 11.6 15.3 10.4 3/31/06 2:30 PM 31 14 45 13.7 9.94 4/1/06 1:30 PM 1 13 68 440 293 233 206 147 105 4/1/06 2:30 PM 1 14 69 4/1/06 3:30 PM 1 15 70 -.£r>3.SB 4/1/06 4:30 PM 1 16 71 V •m 4/2/06 4:00 PM 2 16 95 418 186 114 4/2/06 5:00 PM 2 17 96 169 4/2/06 6:00 PM 2 18 97 4.6 4/2/06 7:00 PM 2 19 98 4.29 4/3/06 3:00 PM 3 15 118 350 118 4/3/06 4:00 PM 3 16 119 326 4/3/06 5:00 PM 3 17 120 6.85 4/3/06 6:00 PM 3 18 121 4.11 4/4/06 11:30 AM 4 11 138 99.5 4/4/06 12:30 PM 4 12 139 99.4 4/4/06 1:30 PM 4 13 140 31 4/4/06 2:30 PM 4 14 141 31.2 4/5/06 10:30 AM 5 10 161

mnVjWT Wasm rerun sample I

Chapter 3 Data- Appendix II Concentration Data: P04-P (mg/L) Solution # Time Dav Hour Cum. time (hr) 10 11 12 13 14 15 16 17 IS 3/29/06 5:00 PM 29 17 0 440 400 333 223 200 169 106 96.6 81.7 3/30/06 12:30 PM 30 12 19 427 388 337 220 199 159 103 95.2 76 3/31/06 12:30 PM 31 12 43 434 406 308 214 197 129 104 88.1 57.2 3/31/06 12:30 PM 31 12 43 47.7 3/31/06 1:30 PM 31 13 44 26 3/31/06 2:30 PM 31 14 45 4/1/06 1:30 PM 1 13 68 422 288 213 142 125 105 72.5 55.4 4/1/06 2:30 PM 1 14 69 291 90.2 4/1/06 3:30 PM 1 15 70 84.1 76.9 4/1/06 4:30 PM 1 16 71 73.5 4/2/06 4:00 PM 2 16 95 407 232 151 126 77.8 57.1 4/2/06 5:00 PM 2 17 96 4/2/06 6:00 PM 2 18 97 4/2/06 7:00 PM 2 19 98 4/3/06 3:00 PM 3 15 118 393 225 154 126 80.2 56.9 4/3/06 4:00 PM 3 16 119 4/3/06 5:00 PM 3 17 120 4/3/06 6:00 PM 3 18 121 4/4/06 11:30 AM 4 11 138 331 213 162 122 74.7 58 4/4/06 12:30 PM 4 12 139 4/4/06 1:30 PM 4 13 140 4/4/06 2:30 PM 4 14 141 4/5/06 10:30 AM 5 10 161 326 183 149 123 74.2 54.8

Chapter 3 Data- Appendix II Concentration Data: Mg (mg/L) Solution # Time Dav Hour Cum. time (hr) 1 2 3 4 5 6 7 8 9 3/29/06 5:00 PM 29 17 0 68.98 61.66 52.98 31.76 30.18 28.48 16.04 15.36 16.02 3/30/06 12:30 PM 30 12 19 3/31/06 12:30 PM 31 12 43 61.58 52.24 43.08 29.16 28.66 7.62 13.62 12.78 2.56 3/31/06 12:30 PM 31 12 43 4.66 8.72 1.9 3/31/06 1:30 PM 31 13 44 129.08 62 72.48 3/31/06 2:30 PM 31 14 45 60.66 63.16 4/1/06 1:30 PM 1 13 68 61.64 3.44 1.82 29.98 6.22 14.5 4/1/06 2:30 PM 1 14 69 6.0B 4/1/06 3:30 PM 1 15 70 196.26 237.82 123.86 4/1/06 4:30 PM 1 16 71 220.54 224.32 121.1 4/2/06 4:00 PM 2 16 95 62.58 9.3 15.78 4/2/06 5:00 PM 2 17 96 3.42 4/2/06 6:00 PM 2 18 97 101.3 4/2/06 7:00 PM 2 19 98 100.44 4/3/06 3:00 PM 3 15 118 19.48 15.72 4/3/06 4:00 PM 3 16 119 4.52 4/3/06 5:00 PM 3 17 120 206.28 4/3/06 6:00 PM 3 18 121 214.3 4/4/06 11:30 AM 4 11 138 17.06 4/4/06 12:30 PM 4 12 139 11.62 4/4/06 1:30 PM 4 13 140 76.76 4/4/06 2:30 PM 4 14 141 76.22 4/5/06 10:30 AM 5 10 161

neaatlve value ISOLD below PLO

Chapter 3 Data- Appendix II Concentration Data: Mg (mg/L)

16 17 18 Tlm« Dav Hour Cum. time (hr) 10 11 12 13 14 15 14.72 15.38 15.4 3/29/06 5 00 PM 29 17 0 56.98 54.24 47.7 29.6 28.52 26.14 3/30/06 12 30 PM 30 12 19 2.18 3/31/06 12 30 PM 31 12 43 57.32 50.48 38.8 29.94 67.5 6.76 15.74 13.66 3/31/06 12 30 PM 31 12 43 292.28 3/31/06 1 30 PM 31 13 44 260.18 3/31/06 2 30 PM 31 14 45 13.32 1.54 0.36 4/1/06 1 30 PM 1 13 68 58.3 0.72 28.58 1.48 129.04 4/1/06 2 30 PM 1 14 69 2.26 wmmm 134.84 4/1/06 3 30 PM 1 15 70 3.74 4/1/06 4 30 PM 1 16 71 1.84 1.62 1.76 1.66 4/2/06 4 00 PM 2 16 95 57.76 30.76 1.98 4/2/06 5 00 PM 2 17 96 4/2/06 6 00 PM 2 18 97 4/2/06 7 00 PM 2 19 98 31.06 2.66 2.22 2.68 2.54 4/3/06 3 00 PM 3 15 118 59.54 4/3/06 4 00 PM 3 16 119 4/3/06 5 00 PM 3 17 120 4/3/06 6 00 PM 3 18 121 2.78 2.36 2.64 2.18 4/4/06 11 30 AM 4 11 138 12.46 30.42 4/4/06 12 30 PM 4 12 139 4/4/06 1 30 PM 4 13 140 4/4/06 2 30 PM 4 14 141 2.06 1.9 2.18 2.3 4/5/06 10 30 AM 5 10 161 3.46 13.38

negative value below PLO

Chapter 3 Data- Appendix II 160 Concentration Data: Ca fmq/U

3 4 5 6 7 8 9 Dav Hour Cum. time (hr) 1 2 Time 16.04 51.92 35.86 30.88 29.4 19.3 17.86 3/29/06 5:00 PM 29.00 17 0.00 68.98 61.38 3/30/06 12:30 PM 30.00 12 19.00 34.5 34.16 24.68 17.96 16.28 13.9 V31/flfi 12-30 PM 31.00 12 43.00 69.24 59.14 50.56 22.78 14.74 12.66 3/31/06 12*30 PM 31.00 12 43.00 21.14 12.88 12.72 3/31/06 1:30 PM 31.00 13 44.00 12.44 11.14 3/31/06 2:30 PM 31.00 14 45.00 34.96 19.46 30.86 15.2 16.04 4/1/06 1:30 PM 1.00 13 68.00 50.92 16.38 4/1/06 2:30 PM 1.00 14 69.00 34.5 22.32 4/1/06 3:30 PM 1.00 15 70.00 31.3 37.48 23.64 4/1/06 4:30 PM 1.00 16 71.00 46.12 15.44 16.92 4/2/06 4:00 PM 2.00 16 95.00 59.26 14.68 4/2/06 5:00 PM 2.00 17 96.00 26.66 4/2/06 6:00 PM 2.00 18 97.00 27.22 4/2/06 7:00 PM 2.00 19 98.00 18.62 4/3/06 3:00 PM 3.00 15 118.00 14.46 4/3/06 4:00 PM 3.00 16 119.00 12.48 4/3/06 5:00 PM 3.00 17 120.00 19.66 4/3/06 6:00 PM 3.00 IS 121.00 19.34 17.36 4/4/06 11:30 AM 4.00 11 138.00 16.48 4/4/06 12:30 PM 4.00 12 139.00 15.76 4/4/06 1:30 PM 4,00 13 140.00 16.08 4/4/06 2:30 PM 4.00 14 141.00 4/5/06 10:30 AM 5.00 10 161.00

ggSSBPBia«SS^ negative value BOLD Ire-run samples

Chapter 3 Data- Appendix II Concentration Data: Ca fma/U Solution # Time Dav Hour Cum. time rhr) 10 11 12 13 14 15 16 17 18 3/29/06 5:00 PM 29.00 17 0.00 64.76 58.56 52.16 36.46 31.92 29.04 19.86 19.88 15.88 3/30/06 12:30 PM 30.00 12 19.00 3/31/06 12:30 PM 31.00 12 43.00 65.46 57.42 47.08 33.9 31.54 24.12 19.68 16.46 14.18 3/31/06 12:30 PM 31.00 12 43.00 40.62 3/31/06 1:30 PM 31.00 13 44.00 38.44 3/31/06 2:30 PM 31.00 14 45.00 4/1/06 1:30 PM 1.00 13 68.00 57.6 18.18 30.02 13.84 18.56 17.3 13.1 13.08 4/1/06 2:30 PM 1.00 14 69.00 17.74 16.6 4/1/06 3:30 PM 1.00 15 70.00 20.66 16.82 4/1/06 4:30 PM 1.00 16 71.00 22.66 4/2/06 4:00 PM 2.00 16 95.00 59.38 33.92 14.6 18.62 14.24 13.42 4/2/06 5:00 PM 2.00 17 96.00 4/2/06 6:00 PM 2.00 IS 97.00 4/2/06 7:00 PM 2.00 19 98.00 4/3/06 3:00 PM 3.00 15 118.00 62.2 32.62 15.08 19.54 15.74 13,32 4/3/06 4:00 PM 3.00 16 119.00 4/3/06 5:00 PM 3.00 17 120.00 4/3/06 6:00 PM 3.00 18 121.00 4/4/06 11:30 AM 4.00 11 138.00 27.58 32.92 14.12 18.6 14.46 13.18 4/4/06 12:30 PM 4.00 12 139.00 4/4/06 1:30 PM - 4.00 13 140.00 4/4/06 2:30 PM 4.00 14 141.00 4/5/06 10:30 AM 5.00 10 161.00 15.26 12.62 12.38 18.26 14.34 13.02

mSmmmSSS negative value' BOLD ire-run samples

Chapter 3 Data- Appendix II Concentration data: P04-P, Ca (mg/L) Solution # , . . 2a 3a :.- •• sa •»«Sf8a: PHOSPHORUS ' i 2 C ,** f »a.3 ' iiaf«;211 mmmsoSWSfSHS ®78 iwsisfiiwja'st: '-.»v:-..-.>.!r42g «0»ii»:S»:68 24 27 29 18 mmmm*29 mssmmm 24 mmmms$ 8 14 32 °/o removed In Ca precip % difference 75 73 70 80 69 69 63 71 55 °/o removed In struvite prec! % remainina 1 1 1 2 2 7 29 15 13

CALCIUM "-<'3 - - * . 4 • .v A PRE-PRECIP AVERAGE -ir6a6 60 51 34 33 29 18 18 16 B POST-Ca PreclD AVERAGE 13 35 19 15 16 24 17 16 13 C Post Struvite Precip Averaqe 20 39 36 27 23 21 16 13 12 AkB*f«'!?j*it*«i' "1 i '. • • '. 52 ' 25 19 SS»5S»JWi7, '2 massitllfference s>' ,A-C*f k, .its* '!V.i f 15 MWWW2 4

1 > ti.-K^' 22 WMimi vmtmms iB?C-*.s - "t * -6 .a -17 sseswsasBasa S^)^S^!iraEfl- % removed In Ca precip 80 » i 42 62 Immmm55 smsmsm 51 19 5 13 17 % difference -9 -6 -32 -36 -22 9 6 16 8 °/o removed in struvite precl 30 64 70 80 71 72 90 71 74 °/o remainlnq

Chapter 3 data- Appendix II Concentration data: P04-P, Ca (mg/L) Solution * . .1.-; . • }. lb i 2b 3b - 4b - -r 5b 6b 7b 8b 59b"

PHOSPHORUS 10 ••' , ^M.A 11 15 16 «„wsw.-*-t*™l'7. A PRE-PRECIP AVERAGE 421 398 &r«.-,..«,««,326« 220 199 164 105 93 79 B POST-Ca Precip AVERAGE 329 290 326 183 152 125 105 76 57 C Post Struvite Precip Averaac 329 79 37 183 152 125 84 76 57 A D 92 103 0 » .! >.*-> 37 ,47 39 0 W 17 \4 22 \ \ ,i mass difference!;" >.. v *•* v i\*iH> .< ^19.r-V w .\«289 " \ - 37 «" . ' 47 , ^39 .'21 > - 17 22 B-C r t ' * "tn- st "ty , f <' "«> (v - -i if 0 211 : 289 v ^ (J .".o » . ,21 0 t t 0 % removed In Ca precip 22 27 0 17 24 24 0 19 28 % difference % removed in struvite precl 0 53 89 0 0 0 20 0 0 % remainina 78 20 11 83 76 76 80 81 72

-• • . '"-v. cAffihiM. ••> A-B ' • 0 . 21 t 0 4 3 >.«%; -* is J »C - ' ' 40 «&.,;>?». »36 21 •i" - *18 .,-f- • -,.,9 i, ..2 4 - 3 mass difference.'-.!'-- •f ^1 ,.*\ io * , >•. "\ issB- C • • •* .4 l ,f 10 -i't V -'0 \ 2 *| 0 V 0 % removed In Ca precip 65 69 0 62 56 32 0 21 16 % difference °/o removed in struvite precl 0 -6 20 0 0 0 12 0 0 % remaining 35 37 80 38 44 68 88 79 84

Chapter 3 data- Appendix II Crystal Analysis Metal: Fc Struvite # AA reading RSD % SD Mass of crystal (mg) [struvite]- mg Cone. Metal % Metal 1 - 18.4 0.00000 100.1 2002 - 2 - 64.6 0.00000 99.9 1998 - 3 - 100 0.00000 99.5 1990 - 4 - 20.2 0.00000 100 2000 - 5 0.05 26.5 0.01325 100 2000 0.05 0.0025 8 - 100 0.00000 100.3 2006 - 11 - 60.2 0.00000 100.3 2006 - 12 - 100 0.00000 100.4 2008 - blank - 17.9 0.00000 SD 0.00147 MDL 0.00462 PLQ 0.02311 MDL as % metal 0.00023 PLQ as % metal 0.00116 % values that are lower are ND

Metal: K Struvite # AA reading RSD % SD Mass of crystal (mg) [struvite]- mg Cone. Metal % Metal 1 3.22 100 3.22000 100.1 2002 3.22 0.161 2 8.71 0.4 0.03484 99.9 1998 8.71 0.436 3 7.56 0.6 0.04536 99.5 1990 7.56 0.380 4 8.6 0.7 0.06020 100 2000 8.6 0.430 5 8.09 1.4 0.11326 100 2000 8.09 0.405 8 8.63 0.5 0.04315 100.3 2006 8.63 0.430 11 8.5 0.9 0.07650 100.3 2006 8.5 0.424 12 9.99 0.5 0.04995 100.4 2008 9.99 0.498 blank 0.11 1.2 0.00132 SD 0.40495 MDL 1.27155 PLQ 6.35777 MDL as % metal 0.06358 PLQ as % metal 0.31789 % values that are lower are ND

Metal: Al Struvite # AA reading RSD % SD Mass of crystal (mg) [struvite]- mg Cone. Metal % Metal 1 1.16 8.4 0.09744 100.1 2002 1.16 0.058 2 0.76 6.3 0.04788 99.9 1998 0.76 0.038 3 0.92 6.6 0.06072 99.5 1990 0.92 0.046 4 1.16 2.4 0.02784 100 2000 1.16 0.058 5 1.05 8.4 0.08820 100 2000 1.05 0.053 8 1.13 4.9 0.05537 100.3 2006 1.13 0.056 11 1.23 3.1 0.03813 100.3 2006 1.23 0.061 12 2.36 3 0.07080 100.4 2008 2.36 0.118 blank 1.21 8 0.09680 SD 0.06480 MDL 0.20347 PLQ 2.03465 MDL as % metal 0.01017 PLQ as % metal 0.10173 % values that are lower are ND

Metal: Ca Struvite # AA reading RSD % SD Mass of crystal (mg) [struvite]- mg Cone. Metal % Metal 1 4.52 0.4 0.01808 100.1 2002 4.52 0.226 2 4.46 0.4 0.01784 99.9 1998 4.46 0.223 3 3.68 0.7 0.02576 99.5 1990 3.68 0.185 4 6.19 1.4 0.08666 100 2000 6.19 0.310 5 4.65 0.2 0.00930 100 2000 4.65 0.233 8 4.23 0.4 0.01692 100.3 2006 4.23 0.211 11 4.32 0.4 0.01728 100.3 2006 4.32 0.215 12 8.84 0.1 0.00884 100.4 2008 8.84 0.440 blank 0.16 0.9 0.00144 average 0.02246 MDL 0.07052 PLQ 0.35259 MDL as % metal 0.00353 PLQ as % metal 0.01763 % values that are lower are ND

Chapter 3 Data- Appendix II 1 Crystal Analysis- Solution #1

71 2-Theta-Scale Ekruvitefrcrn urine aoril 11 - Rle sarnie #1 .raw- Tvoe ZTh/Th locked - Start: 5.00 ° - End 70.000 - Stea 0.020 - Sen tirre: 0.8 s - Term.: 25 °C (Ftocrrt - Time Sated: 1144780760 s - 2-Theta 5.00 ° - LHoo-01 5-0762 O - Struvite. svn - r>l-MlvtPC4-6H20- Y: 41.66 %- d xbv: 1. - WL 1.5406 - Crthcrhorrbic -a 6.94500 - b 11.20800 - c 613550 - a cha 90.000 - beta 90.000 - aamrm 90 000 - Frirritive - P

166 1500

Crystal Analysis- Solution #4

4>M i i I 'i I i I I i i i r | i i i i—i—i—r

10 20 40 50 60 71 2-Theta-Scale fflstruvrtefrtrn urine aoril 11-Rla sarDle#t/aiw-T\oe:2TrvThlccked- Start: 5.00 ° - End 70.00 °-Stea 0.02° -Steo time: 04 s-Terro.: 25 °C(Rxm^-Tirre Started: 11W 5.00 °- Ebo-015-0762 n - Struvite. svn - NH4rvtf>C4-6H20- Y: 41.66 %-dxbv: 1.-WL 1.5406 - Ctthcrhcrrbic - a 6.94500 - b 11.20800 - c 6.13550 - a cha 90.000 - beta 90.000 - aarrma 90.000 - R-irritive - P

167 Crystal Analysis- Solution #5

2-Theta-Scale r^ruvftefrcrn urine aoril 11 - Rla sarmle #5.ra/v- Tvoe: 2Th/Th locked - Siart: 5.00 ° - End 70.00 ° - Stea 0.02 ° - Steo time: 02 s - Terrn.: 25 *C (Room) - Tirre Started: 1144791168 s - 2-Theta 5.00 ° - Bo^OlS07c2n-9ruMte.Svn-.M-W 1.-WL 1.54CI6-Crthcrrvorrbic-a 6.94500-b11.20800-c6.13550-alriia 90.000-beta 90.030-aartTra

168 Crystal Analysis- Solution #8

i i—i—r "i—I—i—i—i—i—T I—i—i—i—i—i—rn—i—i—|—i—i—i—i—|—i—i—i—i— 10 20 40 50 60

2-Theta-Scale fflstruvitefrcm urine aoril 11-File sarrole*8.ra«-Tvoe: 2Th/Th locked - Start: 5.00 "-End: 70.00 °-Steo: 0.02 °-Steo time: 02s-Terrt>.:25 t:(Rocm1-Time Started: 1144792064 s - 2-Theta 5.00 °- ®00-0150762 n - Sruvi te. svn - M-WM3PC4-6H20 - Y: 41.66 % - d x bv: 1. - WL 1.5406 - Crthcrhcrrbic - a 6.94500 - b 11.20800 - c 6.13550 - al cha 90.000 - beta 90.000 - carrrrB 90.000 - Prirritive - P

169 Crystal Analysis- Solution #11

I—r 1—r~T^—i—r

6 10 20 30 2-Theta-Scale

EdWruvttefrcrn urine aoril 11 - Rle sarrDle #1 l.raw- Tvoe 2ThTh locked - Start 500 ° - End: 70.00 ° - Steo: 0.02 ° - Steo tirre: 0.2 s - Term: 25 °C fRxrrt - Time Started: 1144792832 s - 2-Theta 5.00' BoO-01 5-0762 (*) - Struvite. svn - WWtPOA-QrQO- Y: 41.66 %-dxbv. 1.-WL 1.5406 - CHhcitKrrbic - a 6.94500 - b 11.20800 - c 6.13550 - alcha 90.000 - beta 90.000 - carrma 90.000 - R-irritive - P

170 Crystal Analysis- Solution #12

y -r-r-r "i—[~n—i—r T—i—i—r 10 20 30 40 50 60 2-Theta-Scale Estruvitefrtrn urine aoril 11 - Rle sarrole #12raw- Tvoe 2TrrTh locked - Start 500 ° - End: 70.006 - Steo: 0.02 ° - Steo time 0.2 s - Term.: 25 °C (Rxm) - Time Started 1144793600 s - 2-Theta 5.00 ° Ebo-01 S0762 (*) - Struvite. svn - M-l4IVtPC4-6H20- Y: 41.66 %- d xbv: 1. - WL 1.5406 - Crthcmcrrbic -a 6.945D0- b 11.20800- c 6.13550 - a cha 90.000 - beta 90.000 - oarrma 90.000 - R-irrftive - P

171 APPENDIX III

172 SUBJECT INFORMATION AND CONSENT FORM

Recovery of Struvite from Source Separated Urine

Principal Investigator: James Atwater, P.Eng Department of Civil Engineering University of British Columbia

Co-Investigator: Elizabeth Tilley, MASc Candidate Department of Civil Engineering University of British Columbia

Sponsor: None

Emergency Telephone Number: N/A 1. INTRODUCTION

You are being invited to take part in this research study because you are healthy, free of urinary tract disease and are able to provide urine that is typical of an adult who eats a balanced diet.

2. YOUR PARTICIPATION IS VOLUNTARY

Your participation is entirely voluntary, so it is up to you to decide whether or not to take part in this study. Before you decide, it is important for you to understand what the research involves. This consent form will tell you about the study, why the research is being done, what will happen to you during the study and the possible benefits, risks and discomforts.

If you wish to participate, you will be asked to sign this form. If you do decide to take part in this study, you are still free to withdraw at any time and without giving any reasons for your decision.

If you do not wish to participate, you do not have to provide any reason for your decision not to participate nor will your university education be affected in any way.

Please take time to read the following information carefully and if you wish, discuss it with your family, friends, and doctor before you decide.

3. WHO IS CONDUCTING THE STUDY?

This study is being conducted as part of Masters thesis to be submitted to the Department of Civil Engineering at the University of British Columbia.

4. BACKGROUND

Phosphorus is mined from natural deposits but it is becoming increasingly rare and in the future we will need to find new sources. Finding a new source is important because phosphorus is an essential nutrient for agriculture

Struvite is a white crystal that contains phosphorus and nitrogen and it is a very good fertilizer. Studies have shown that struvite can be recovered from urine however, the process and way that it can be recovered is not very clear. More research is needed to understand if we are ever to recover a phosphorus-based fertilizer from urine.

Preliminary research has shown promising results utilizing synthetic urine. The next step is to verify the results using real urine.

177 5. WHAT IS THE PURPOSE OF THE STUDY?

The purpose of this study is to recover phosphorus from human urine. The research will seek to identify ways of storing and treating urine such that the maximum amount of phosphorus can be removed in the form of struvite

6. WHO CAN PARTICIPATE IN THE STUDY?

Anyone who is in good health, eats a well-balanced diet and can provide approximately 1 L of urine over the course of 24 hours can participate in the study.

7. WHO SHOULD NOT PARTICIPATE IN THE STUDY?

Anyone who is ill, taking drugs or medication, and those who plan to fast during the study period should not participate in this study. As well, anyone who would feel uncomfortable collecting, storing and transporting his or her own urine over the course of 24 hours, should not participate.

8. WHAT DOES THE STUDY INVOLVE?

This study will take place at the University of British Columbia and will involve approximately 20 people.

All subjects in the study will be asked to collect their urine over the course of 24 hours. It is important to collect the urine over a full day because the urine characteristics vary over the course of a day. Urine that is discharged first thing in the morning is much more concentrated and has a darker colour because the body has not had water for the whole night. Urine from midday is much lighter in colour and has a different composition because the body has consumed food and water. By collecting urine for a full day, an average urine sample will be obtained.

Each subject will be provided with enough containers to collect their urine over 24 hours. Subjects will also be given dark coloured bags to store the containers in so that the contents are not visible. Each subject will be asked to use a container to capture any and all urine that is expelled during the day. Only urine is to be collected; a regular toilet should be used for fecal material. Although it is important to collect all the urine that is generated over a day, it is not necessary to collect urine that may be released during a bowel movement.

After a full day of urine collection, the subject will be asked to return the containers to the research laboratory. When the subject returns the containers, the subject will also return a form that they received along with the collection containers. The form will require the subject to state their gender and age (not their name). Also, the subject will be required to state whether or not they were able to collect all of their urine. If some urine is missing, the subject will be asked to

178 estimate the fraction that is missing (e.g. only one half of urine was returned because the subject did not feel comfortable collecting urine at the workplace). All of the urine that is returned by all of the subjects will be mixed in a single storage container and used for experimentation.

Overview of the Study

The goal of the study is to remove the phosphorus from urine in the form of struvite. The urine that is provided by the subjects will be used to conduct experiments that will help the researchers determine how struvite can be recovered. In order for struvite to form, magnesium must be added; different amounts of magnesium will be added to the urine and the urine will be examined to see what is produced.

If You Decide to Join This Study: Specific Procedures

If you agree to take part in this study, the procedures you can expect will include the following:

Before You Begin the Study You will be given enough containers for you to be able to use one for each urination (approx 5-10 depending on your expressed need). You will also be given dark coloured bags that you can use to carry the full and empty containers. You will be informed of the first day of collection and asked to being collecting your urine on that day.

Urine collection On the first day of the collection you will carry on with your day as normal, but instead of urinating into a toilet, you will instead urinate into a collection bottle. You will still be able to use washrooms and will position the urine collection bottle under the urine stream to collect it. After you have emptied your bladder, you will seal the container with the lid and will place the container in one of the bags provided for your privacy. You will collect urine for 24 hours, which means that you will collect urine from the first thing in the morning, all through the day, and at night.

Returning the urine The morning after you have completed the 24-hour collection period, you will be asked to return the collection bottles that contain the urine as well as the ones that you did not use, to a laboratory at the University of British Columbia. You will be told the time and room number before you begin the study.

Number of collection periods The study will run from March 01, 2006 until June 01, 2006. You may be contacted up to three times during that period to participate in a 24-hour collection period.

9. WHAT ARE MY RESPONSIBILITIES?

The subject is responsible for eating and drinking a normal healthy amount during the day. The subject is responsible for accurately reporting the fraction of the day's urine that was collected.

179 The subject is responsible for returning the collected urine in a timely manner, as the nature of the research is time sensitive.

10. WHAT ARE THE POSSIBLE HARMS AND SIDE EFFECTS OF PARTICIPATING?

There are no known risks associated with participating in this research although some subjects may feel anxious or embarrassed when having to enter or leave a washroom with a collection container.

11. WHAT ARE THE BENEFITS OF PARTICIPATING IN THIS STUDY?

There are no benefits from participating in this study.

We hope that the information learned from this study can be used in the future for the betterment of society.

The published results will be available, should you have an interest in reviewing our findings.

12. WHAT ARE THE ALTERNATIVES TO THE STUDY TREATMENT?

Previous research has been conducting using synthetic urine, but we wish to replicate our results using real urine to ensure the validity of the results.

13. WHAT HAPPENS IF I DECIDE TO WITHDRAW MY CONSENT TO PARTICIPATE?

Your participation in this research is entirely voluntary. You may withdraw from this study at any time. If you decide to enter the study and to withdraw at any time in the future, there will be no penalty

If you choose to enter the study and then decide to withdraw at a later time, all data collected about you during your enrolment in the study will be retained for analysis. By law, this data cannot be destroyed

180 14. WHAT HAPPENS IF SOMETHING GOES WRONG?

Signing this consent form in no way limits your legal rights against the sponsor, investigators, or anyone else.

15. CAN I BE ASKED TO LEAVE THE STUDY?

If you are not complying with the requirements of the study or for any other reason, the researcher may withdraw you from the study

16. AFTER THE STUDY IS FINISHED

After the study has completed you will not be contacted.

17. WHAT WILL THE STUDY COST ME?

There will be no reimbursement for study related expenses if that is the case, nor will you be paid for participating

18. WILL MY TAKING PART IN THIS STUDY BE KEPT CONFIDENTIAL?

Your confidentiality will be respected. No information that discloses your identity will be released or published without your specific consent to the disclosure. However, research records may be inspected in the presence of the Investigator or his or her designate by representatives of Health Canada and the UBC Research Ethics Board for the purpose of monitoring the research. However, no records, which identify you by name or initials, will be allowed to leave the Investigators' offices.

19. WHO DO I CONTACT IF I HAVE QUESTIONS ABOUT THE STUDY DURING MY PARTICIPATION?

If you have any questions or desire further information about this study before or during participation, you can contact Elizabeth Tilley

181 20. WHO DO I CONTACT IF I HAVE ANY QUESTIONS OR CONCERNS ABOUT MY RIGHTS AS A SUBJECT DURING THE STUDY?

If you have any concerns about your rights as a research subject and/or your experiences while participating in this study, contact the 'Research Subject Information Line in the University of British Columbia Office of Research Services' at 604-822-8598

182 21. SUBJECT CONSENT TO PARTICIPATE

a. I have read and understood the subject information and consent form. b. I have had sufficient time to consider the information provided and to ask for advice if necessary. c. I have had the opportunity to ask questions and have had satisfactory responses to my questions. d. I understand that all of the information collected will be kept confidential and that the result will only be used for scientific objectives. e. I understand that my participation in this study is voluntary and that I am completely free to refuse to participate or to withdraw from this study at any f. I agree to being contacted up to three (3) times to provide specimens and I understand that I may change my mind at any point without repercussions g. I understand that I am not waiving any of my legal rights as a result of signing this consent form. h. I understand that this study will not provide any benefits to me. i. I have read this form and I freely consent to participate in this study. j. I have been told that I will receive a dated and signed copy of this form.

By signing this form you are consenting to participate in the study as detailed above.

SIGNATURES

Printed name of subject Signature Date

Printed name of witness Signature Date

Printed name of principal investigator/ Signature Date designated representative

183