Handbook for Water Treatment Operation PDF

HANDBOOK FOR THE OPERATION OF WATER TREATMENT WORKS The Water Research Commission The Water Institute of Southern Africa Editor: Frik Schutte Water Utilisation Division Department of Chemical Engineering University of Pretoria TT 265/06 MARCH 2006 Obtainable from: Water Research Commission Private Bag X03 Gezina 0031 The publication of this report emanates from a project entitled: Development of a handbook for the operation of water treatment works (Consultancy No 559) DISCLAIMER This report has been reviewed by the Water Research Commission (WRC) and approved for publication. Approval does not signify that the contents necessarily reflect the views and policies of the WRC, nor does mention of trade names or commercial products constitute endorsement or recommendation for use. ISBN 1-77005-428-6 Printed in the Republic of South Africa Cover photographs courtesy of Rietvlei Water Treatment Works TABLE OF CONTENTS PURPOSE OF THE HANDBOOK....................................................................... 1 WHO SHOULD USE THE HANDBOOK........................................................... 2 PART A: BACKGROUND 1. Introduction to water quality and treatment.................................................... 3 2. Drinking water quality..................................................................................... 25 3. Introduction to Water Chemistry..................................................................... 38 4. Water treatment calculations........................................................................... 58 PART B: TREATMENT PROCESSES 1. Coagulation-flocculation................................................................................. 70 2. Sedimentation and flotation............................................................................. 87 3. Sand filtration.................................................................................................. 98 4. Disinfection..................................................................................................... 117 5. Chemical stabilisation..................................................................................... 145 6. Fluoridation..................................................................................................... 158 7. Residuals handling and treatment.................................................................... 164 8. Advanced processes......................................................................................... 172 PART C: MANAGEMENT AND CONTROL 1. General management concepts........................................................................ 193 2. Treatment plant control................................................................................... 200 3. Maintenance and trouble shooting................................................................... 219 4. Safety............................................................................................................... 227 PART D: APPENDICES AND REFERENCES References................................................................................................................ 232 iii iv PURPOSE OF THE HANDBOOK The treatment and supply of drinking water is a challenging task that includes many diverse aspects ranging from water quality management in catchments, operation and control of water treatment plants, distribution of treated water, community participation, project management, etc. The purpose of this handbook is to provide comprehensive information specifically on all aspects related to the treatment of water for domestic use. The focus is on the operational aspects of treatment plants and processes and not on process design. Sufficient background and process descriptions are provided to enable a proper understanding of the functioning of the different processes and on aspects such as the suitability of processes for different types of water and the limitations of different processes. The purpose of this Handbook is not to provide an instruction manual or task list for process controllers or operators to operate unit processes or a treatment plant. A handbook in the form of an illustrated operational guide is being prepared for the WRC for this purpose. The objective of this Handbook is to provide the plant supervisor and process controller with sufficient knowledge and insight to: Assess raw water quality as well as the quality of water from individual unit processes and the treatment plant as a whole to ensure that final water of the required quality is produced Understand the implications to consumers and other stakeholders if sub-standard water is produced and supplied Calculate and make adjustments to dosages and operating parameters in response to changes in raw water quality or other requirements Assess the performance of unit processes and the plant as a whole Identify potential causes of poor performance of unit processes Optimise the performance of unit processes and the plant as a whole Perform basic management tasks including water loss assessment and control and safety management The authors recognise that the majority of plant operators currently in operating positions might find the level of material in this Handbook too advanced and would encourage them to use the illustrated handbook mentioned above. However, the authors feel that wherever possible plant operators should be encouraged to improve their knowledge and qualifications and this Handbook should be useful for that purpose. The layout of this Handbook is such that some aspects are repeated in different sections. Many aspects covered in Part A (e.g. in Overview of Treatment Processes and in Water Chemistry) are discussed again in Part B (Treatment Processes). This duplication is beneficial to the reader since the focus of the different sections is different and the same material is presented from different perspectives. 1 This is a first attempt at developing a Handbook specifically for the operation of drinking water treatment plants to match the Handbook for the Operation of Wastewater Treatment Works. The initiatives of the Water Research Commission (WRC) and the Water Institute of Southern Africa (WISA) are recognised to have the Handbook developed. Some sections in this Handbook are similar to certain sections in the Wastewater Handbook because the subject material of the basic aspects is very similar. A number of reference books has been used in the compilation of the subject material and they are listed in the reference section. The study material has been compiled by experts in the field with many years of experience in the field of drinking water treatment. However, it is recognised that the Handbook may have shortcomings from the perspective of process controllers and operators who may have specific requirements not adequately addressed by the book. For this reason Handbook users are requested to submit any suggestions for improvement of future editions to the Water Research Commission or WISA or any of the authors. WHO SHOULD USE THE HANDBOOK? The level of presentation in this Handbook is on the BSc. and BTech. level, i.e. suitable for training of process controllers and treatment plant operators at tertiary training level. Some parts of the Handbook will be useful for in-house training of treatment plant operators and process controllers. The Handbook provides basic information on water treatment processes and water supply that should also be useful for other people involved in water treatment and water supply, including engineers and scientists. The groups who may find the Handbook useful include: Water treatment plant supervisors and managers Water treatment plant process controllers Engineers and scientists involved in water treatment and supply Educators and students 2 PART A: BACKGROUND CHAPTER A1: INTRODUCTION TO WATER QUALITY AND TREATMENT Frik Schutte INTRODUCTION The term “water quality” describes the physical, chemical and microbiological characteristics of water. These properties collectively determine the overall water quality and the fitness of the water for a specific use. These properties are either intrinsic to the water or are the result of substances that are dissolved or suspended in the water. Water quality is only meaningful when evaluated in relation to the use of the water. The reason is that water of a certain quality may be fit for a specific use, but completely unfit for another use. For example, water that is fit for human consumption may not be fit as boiler feed water because the dissolved inorganic salts that are acceptable in drinking water, are not tolerated in boiler feed water, since they may precipitate and cause blockages in the boiler equipment. Water that is fit for domestic use (drinking water) must comply with specific requirements. The most important requirement is that it must be safe to drink. Many raw water sources contain harmful micro-organisms or other substances in concentrations that make the water unsafe to drink or in other ways unfit for domestic use. These organisms and substances must be removed from the water by means of treatment processes to make the water fit for domestic use. In addition to the requirement that water must be safe to drink, water for domestic use must also be aesthetically pleasing (have a clean appearance, taste and odour) and it must furthermore be chemically stable (i.e. it must not cause corrosion or form deposits in pipes or fixtures such as geysers). The principal objective therefore of water treatment is to produce water that is fit for domestic use reliably and consistently from a raw water source at a cost that is reasonable to the consumers. A water treatment plant employs many individual treatment processes (sometimes called unit processes and unit operations) that are linked in a process train to produce water of the desired quality. In this chapter a brief overview is given of historic developments in water treatment, followed by a discussion of water quality aspects that are relevant to drinking water treatment. This is followed by an overview of the different treatment processes commonly used for water treatment and the processing of residuals. Each of these processes is discussed in detail in subsequent chapters in the Handbook. 3 HISTORIC DEVELOPMENTS The first references to clean water or ‘sweet water’ and for water to be good for use after passage over a certain number of stones date back about 3000 years to Biblical times. The Roman aqueducts are well-known later landmarks in Europe as testimony of conveying clean water to cities. By the eighteenth century the removal of particles from water by filtration was known as an effective way of clarifying water and the first municipal water filtration plant started operating in Scotland in 1832. However, the main objective at that time was simply to supply clear water because the germ theory and the knowledge that diseases could be spread by water was still unknown. It was only in 1855 that Dr John Snow, an epidemiologist showed empirically that a cholera outbreak in London was caused by drinking water contaminated by faecal wastes from a cholera patient. However, the concept of disinfection as a disease preventing measure and a practical disinfection process only developed much later. Pasteur demonstrated his germ theory only in the 1880’s and chlorination as treatment process was developed after 1905. By the early 1900’s the large increase in the number of water supply systems without proper treatment in the USA contributed to major outbreaks in water-borne diseases. However, it was only with the introduction of chlorine as a treatment process to disinfect water in 1908 that the spreading of diseases through contaminated water could be controlled. Chlorination was rapidly accepted as an essential part of water treatment and this resulted in a substantial decline in the number of deaths due to water-borne diseases. Research on coagulation-flocculation, sedimentation and filtration as basic water treatment processes during the early part of the previous century contributed to a better understanding of these processes and much improved performance. New processes were also developed during that time in Europe. The use of ozone for disinfection and taste- and colour enhancement was introduced early in the century in France and Germany. The most significant process development since the introduction of chlorine during the previous century was the development of synthetic membranes as treatment process. The first practical reverse osmosis membranes for the desalination of seawater were developed in the 1960’s. Later, other types of membranes were developed, including nanofiltration (NF), ultrafiltration (UF) and microfiltration (MF) membranes. These membranes find application in water treatment other than just desalination. For example, NF and UF membranes are used to replace some conventional treatment processes for removal of natural organic substances and micro-organisms from water. In South Africa water treatment and supply followed the pattern of the western world and in some aspects South Africa actually lead the way. For example, the concepts of water reclamation and reuse were pioneered in South Africa and Namibia. South Africa is also known as a world leader in the field of biological nutrient removal in advanced wastewater treatment. Sadly, the supply of clean and safe water in South Africa was until recently limited to formal municipalities. Most rural areas and many townships had no (or limited) water 4 supply and poor sanitation. In many rural areas in the country water is abstracted from a river, stream, borehole or well and consumed without any treatment or with only limited treatment. In these situations, the health of consumers is often at risk. Often the erection of a conventional treatment plant and the effective running of such processes may not be possible in many rural areas and alternative approaches have to be followed. The situation with respect to water supply has however, improved dramatically over the last number of years and it is a clear objective of government (and it is stated as a basic human right in the RSA Constitution) that every citizen must have access to clean water and basic sanitation. Many ambitious programmes are currently under way for water supply and sanitation in previously disadvantaged communities and in rural areas. The result of these developments is that the need for engineers and trained operators and process controllers has increased and special efforts are needed to provide appropriately trained people at all levels in the water treatment and supply industry. The main challenges facing the water industry today include: Deterioration in the quality of many raw water sources Removal of potentially harmful synthetic organic substances in water sources Removal of resistant micro-organisms from water Improved training of process controllers for new processes and process optimisation Demands for process integration and flexibility WATER QUALITY General aspects of water quality Water is a unique substance and one of its unique characteristics is its capacity to dissolve a variety of substances. As water moves through its cycle, called the hydrological cycle, comprising of rainfall, runoff, infiltration, impounding, use and evaporation, it comes into contact with many different substances that may be dissolved by the water to a greater or lesser extent or that may be suspended in the water. The type and amount of the dissolved substances together with suspended and colloidal substances (very small suspended particles) collectively determine the overall quality of the water and its fitness for domestic use. The types of contaminants or substances of concern that may occur in water sources vary over a wide spectrum and include inorganic salts, micro-organisms, clay particles and organic material. Those with similar characteristics that can be treated by the same type of treatment process are normally grouped together for design purposes and for general discussion. It is normally not possible to consider each individual substance of concern with the view to treatment. There are exceptions, however, for example the removal of a toxic substance from water is often specific for the particular substance. 5 The substances of concern in water can be categorised in different ways, e.g. as dissolved or suspended, as inorganic or organic, as macro or micro substances, as natural or synthetic substances, suspensions of micro-organisms etc. For the purpose of this discussion the characteristics of different groups of substances that are treated by he same types of processes will be considered briefly Dissolved substances Most substances are to a greater or lesser extent dissolved by water. Substances that are dissolved by water include gasses such as oxygen (O2), carbon dioxide (CO2) and ammonia (NH3), inorganic compounds such as sodium chloride (NaCl) and calcium sulphate (CaSO4) and organic substances such as humic acids and carbohydrates. Dissolved substances are generally more difficult to remove from water than suspended substances, since they must either be converted into the solid form by means of precipitation, or to the gas form by means of oxidation so that the gas can escape or be stripped from water. A further possibility to remove dissolved substances is by using advanced processes such as reverse osmosis or activated carbon adsorption. Suspended and colloidal substances In addition to the substances that are dissolved in water, some substances may not dissolve in water but remain in suspension as very small suspended or colloidal particles. Suspended solids are defined as solids that are relatively large and settle easily under quiescent conditions. Suspended solids are normally determined by filtering the suspended solids from a water sample of known mass, and determining the mass of the dried solids. Colloidal particles on the other hand are too small to settle and they also carry an electrical charge that prevents them from settling. They can actually remain in suspension for days without settling. A colloidal system is defined as a system in which particles in a finely divided state are dispersed in a continuous medium. Colloidal particles are not limited to any particular group of substances but are defined by size. The colloidal size range is generally regarded to extend in size from about 10 nanometer (nm) to 1 micrometer (μm). Colloidal particles impart undesirable properties to water: Turbidity is most often caused by inorganic clay minerals in surface water. Most turbidity particles are hydrophobic (water repelling) and range in size from 0.2 to 10 μm. Turbidity can be readily removed from water by coagulation-flocculation and separation. Colloidal organic substances, i.e. humic and fulvic acids with molecular mass ranging between 800 and 50000 Daltons (Mol Mass units) generally cause colour in natural water. Colloidal metal hydroxides (e.g. iron) also cause colour in water. Most of the particles responsible for colour are hydrophilic (water attractive) and more difficult to remove by coagulation than turbidity particles. 6 Bacteria, viruses and micro-algae are also colloidal in nature. They consist of polar organic molecules, are hydrated and hydrophilic. Certain complex organic compounds in treated industrial wastewater can also be considered as colloidal. Colloidal suspensions are stable and must be destabilised before it is possible to aggregate them into bigger floc particles that can be removed by sedimentation and filtration. Destabilisation is effected by coagulation, and aggregation by flocculation. These processes are discussed in detail in subsequent chapters. CHEMICAL WATER QUALITY Dissolved organic and inorganic substances determine the chemical quality of water. These substances have a wide range of effects on the chemical properties of the water. For example, some of these substances can be toxic (chromium, arsenic), while other cause the water to be hard or scale forming (calcium carbonate), and other chemical compounds may affect the taste and odour of the water (sodium chloride, geosmin). A large range of inorganic chemical compounds can be present in water. These compounds such as sodium chloride, NaCl and calcium sulphate, CaSO4 dissolve in water in the form of the respective ions, i.e. Na+, Ca2+, Cl- and SO42-. A chemical analysis of the water gives the concentration of individual ions, normally in mg/l. The total quantity of dissolved inorganic compounds in water is expressed as the concentration of Total Dissolved Inorganic Solids (TDIS or more commonly TDS) in mg/l. An indication of the TDS is given by the Electrical Conductivity (EC) of the water, as is discussed below under physical properties. A full inorganic chemical analysis of raw water and final treated must be performed at regular intervals to determine whether any compounds are present in the raw water at concentrations that may be a cause for concern. A full analysis of the raw water must be performed at least once per year while the analysis on treated water must be in accordance with the specification or guidelines prescribed for the treatment plant (SABS 241, DWAF Water Quality Guidelines for Domestic Use) A very large variety of organic substances can be present in water. These substances may either be natural substances such as decaying plant material, algal or bacterial by- products and carbohydrates, or synthetic compounds such as pesticides, herbicides and solvents as well as products formed during water treatment such as chloroform and other chlorinated products. There are many thousands of organic compounds that have been identified in water, most of them at very low concentrations. Organic compounds have carbon as a main element in their composition and most of them do not go into solution as ions but go into solution as molecules of the compound. Organic compounds are determined either collectively by means of oxidation, as individual compounds or as groups of compounds. An indication of the general organic quality of the water can be obtained by means of the determination of cumulative parameters such as: 7 Total organic carbon (TOC) Dissolved organic carbon (DOC) Chemical oxygen demand (COD) Biological oxygen demand (BOD) Individual organic compounds or groups of compounds are generally determined by gas chromatography and mass spectrohotometry (GC-MS) or by other specialised methods. One of the organic groups of compounds of particular interest in drinking water treatment is the so-called Trihalomethanes, expressed in Pg/l THM. This group includes compounds such as chloroform (CHCl3 ) and related compounds. Other chemical water quality parameters There are a number of collective water chemical parameters specific to water treatment. These include alkalinity, hardness, chemical stability, free available - and combined chlorine species. Alkalinity of water is a measure of its acid-neutralising capacity. Alkalinity plays an important role to buffer water and prevent changes in pH due to addition of acid, or acid-producing chemicals such as ferric chloride. Alkalinity is determined by the concentration of carbonate, bicarbonate and hydroxide species in the water and by the pH. Different forms of alkalinity are distinguished by titrating the water to different specific end points. Total alkalinity (also known as the Carbonic Acid Alkalinity) is determined by titrating with a strong acid to a pH end point of 4,5. Phenolphtalein alkalinity (also known as bicarbonate alkalinity) is determined by titrating a water sample to the phenolphthalein end point of pH 7.8. (See Chapter A3 Introduction to Water Chemistry for a detailed discussion). The hardness of water is determined by the concentration of divalent cations in the water, mostly calcium and magnesium and is expressed as mg/l CaCO3. Hardness affects the lather-forming ability of water with soap. Different forms of hardness can be distinguished (see Chapter A3) all expressed as mg/l CaCO3: carbonate or temporary hardness, which is caused by calcium and magnesium associated with bicarbonate in the water non-carbonate or permanent hardness, which is caused by calcium and magnesium associated with ions other than bicarbonate such as chloride and sulphate, calcium hardness, caused by all the calcium ions in solution, magnesium hardness, caused by all the magnesium ions in solution, and total hardness, which is the sum of calcium and magnesium hardness. 8 Table A1.1 gives an indication of classification of waters in terms of hardness: TableA1.1: Hardness classification Hardness classification Total hardness as mg/l CaCO3 Soft Less than 50 Reasonably soft 50 to 100 Slightly hard 100 to 150 Reasonably hard 150 to 250 Hard 250 to 350 Very hard More than 350 The chemical stability of water is a very important characteristic since it determines whether water will be chemically stable, aggressive-corrosive or scale forming. This has very important cost implications for maintenance of distribution systems. If water is supersaturated with respect to calcium carbonate, the calcium carbonate will precipitate and form a layer of chemical scale on the surface of pipes and fixtures. A thin layer of calcium carbonate affords protection against corrosion, while excessive precipitation reduces the carrying capacity of pipes and may even lead to blocking of pipes in extreme cases. On the other hand, if water is under-saturated with respect to calcium carbonate, any layer that may have precipitated will dissolve leaving the metal or other pipelining material exposed and subject to chemical attack (e.g. corrosion). It is therefore advisable to treat water to a slight super-saturation for protection against corrosion. There are different methods to express chemical stability. The indices that have been used generally are the Langelier saturation index (LSI) and the Ryznar stability index (RI). These indices are qualitative in nature and therefore do not provide adequate information. A quantitative, and therefore more satisfactory way of determining the chemical stability of water is to calculate the calcium carbonate precipitation potential of the water. Calcium Carbonate Precipitation Potential (CCPP) is a parameter that gives the actual mg/l of CaCO3 that would theoretically precipitate from the water. A positive CCPP of about 4 mg/l has been shown to give adequate protection against corrosion without excessive CaCO3 precipitation. The determination of CCPP has been made very simple by using the Stasoft computer program that is available from the Water Research Commission (See Chapter on Stabilisation). Residual chlorine Chlorine is the most generally used agent to disinfect water. Stringent control of the amount of chlorine dosed and the residual chlorine concentration after a certain contact time is necessary to ensure microbiologically safe water. Chlorine gas (Cl2) dissolves in water to form hypochlorous and hydrochloric acid. The actual disinfecting agent is hypochlorous acid which dissiociates to form the hypochlorite ion, OCl-. 9 The amount of hypochlorous acid (HOCl) together with hypochlorite ion (OCl-) is termed free available chlorine. The residual of free available chlorine must generally be 0,5 mg/l after a contact time of 30 minutes to ensure properly disinfected water. The C.t concept (residual concentration multiplied by contact time) is discussed in detail in the chapter on Disinfection. Chlorine in the form of monochloramine (together with other chloramine species) termed combined available chlorine) is also used for water disinfection. It is formed when chlorine is added to water that contains a small amount of ammonia. The ammonia reacts with HOCl to form monochloramine, NH2Cl. It is much less effective as a disinfectant than HOCl (the same order of effectiveness as chlorite ion). However, it has the advantage of being much more stable in water than free available chlorine. For this reason it is often used to provide residual protection in larger distribution systems. PHYSICAL WATER QUALITY The physical quality of water is determined by intrinsic characteristics as well as by dissolved and colloidal substances in the water. Intrinsic physical properties include temperature, viscosity, and surface tension. Other physical properties such as electrical conductivity, colour, taste and odour are determined by the presence of dissolved and colloidal substances in the water. Some characteristics of water are often indicated as physical characteristics, while they are in actual fact chemical in nature, pH being an example. These properties can also be called physico-chemical properties. The general physical properties of water that play a role in treatment are discussed in the following section. Turbidity gives an indication of the concentration of colloidal particles in water. Turbidity is expressed in nephelometric turbidity units, NTU. It is determined in a Nephelorometer by comparing the intensity of light scattered by the water sample to the intensity of light scattered by a standard reference in the turbidity meter. The turbidity of raw water can be as low as a 1 or 2 NTU in groundwater and up to several hundred in turbid surface water, e.g. after a rain storm. The turbidity of drinking water should be 90 % Poor > 97 % > 85 % > 90 % NOTE If the compliance frequency targets in respect of microbiological and chemical requirements are in conflict with one another, take a conservative approach and classify performance according to the lower category. 36 Water quality performance should be established on an annual basis indicating compliance (as a percentage) to each requirement listed in table 1 (microbiological safety requirements) and water quality determinants listed and tested for in respect of each class as indicated in table 2 (physical, organoleptic and chemical requirements). It should be noted that only determinants with health implications should be considered for compliance and the other only included for information and operational purposes. It is also recommended that analyses not performed should be documented. Table A2.7: Operational water quality alert values. 1 2 3 Determinant Unit Alert value Turbidity NTU 5 Residual chlorine mg/ l < 0,5a Heterotrophic plate countb count/ml 5 000 [Total] coliform bacteriac count/100 ml 10 Somatic coliphagesd count/100 ml 1 Cytopathogenic virusese count/100 ml 1 Protozoan parasitese count/100 ml 1 (Giardia/Cryptosporidium) a Dependent on network characteristics and chlorine demand. A residual of 0,5 mg/L applies to the waterwork’s final water. The appropriate level in distribution system is 0,2 mg/L. Where other disinfectants are used, appropriate alert levels should be selected. b Process indicator that provides information on treatment efficiency and aftergrowth in distribution networks. c Indicates potential faecal pollution and provides information on treatment efficiency and aftergrowth. d Process indicator that provides information on treatment efficiency and could serve as a model for human enteric viruses. e Confirms a risk of infection and faecal pollution, and provides information on treatment efficiency. The detection of selected viruses confirms faecal pollution of human origin. 37 CHAPTER A3: INTRODUCTION TO WATER CHEMISTRY Frik Schutte INTRODUCTION Water Chemistry forms the basis of most water treatment processes (unit processes) while physical forces and phenomena form the basis of unit operations. However, most water treatment processes involve both chemical and physical forces and for the purpose of this Handbook both physical and chemical changes are referred to as processes. Most processes are based on reactions between chemicals added to water and the substances dissolved or suspended in water, or with water itself. A working knowledge of basic chemical concepts and the reactions that take place in water treatment is therefore essential knowledge for water treatment process controllers and operators. Water Chemistry is a complex area of study and for the purpose of this chapter only the more basic aspects are considered. Some general chemistry aspects are covered in the first part of the chapter and this is followed by a detailed discussion of concentration units. General water treatment chemical reactions and equations are considered next and these concepts are illustrated by examples and calculations. Acid-base reactions, precipitation equilibria, the carbonate system and oxidation- reduction reactions form the rest of the chapter. Basic mass balance and flow of material concepts and examples are also provided. Basic chemical aspects All the substances that are dissolved or suspended in water, like all physical objects, are composed of the following elementary particles: electrons protons neutrons These elementary particles occur in various combinations to form atoms. The atom is the smallest unit of matter that have unique chemical characteristics. Protons and neutrons occur in the nucleus of the atom, while electrons orbit the nucleus. There are 92 different, naturally occurring atoms, each one called an element. An element is defined by its atomic number (the number of protons in the nucleus). The atomic mass of an element is equal to the sum of neutrons and protons in the nucleus. The atomic mass is an important entity because it is used to calculate the molar or formula mass of compounds and the molar concentration of solutions. Elements occur in “families” with similar properties but with different atomic masses. For example, calcium, strontium and barium are all found together as metal carbonates, while chlorine, bromine and iodine are all reactive volatile elements, readily forming halide salts such as NaCl, NaBr, NaI. The Periodic Table consists of all the elements arranged according to their atomic masses and in their “families”. 38 Ions. All atoms are electrically neutral because the number of electrons with negative charge equals the number of protons with positive charge. However, an atom may gain or loose one or more electrons, in which case it acquires a net electrical charge and form an ion. If the overall charge is positive the ion is called a cation (e.g. Na+, Ca2+, Al3+); if the overall charge is negative it is called an anion (e.g. Cl-, O2- ). Ions may also be formed by a combination of atoms with an overall charge, called polyatomic ions (e.g. NH4+, OH-, SO42- ). When doing an analysis, the concentration of the ions is determined and reported rather than the concentration of the inorganic compounds. Molecules are formed through a linking of atoms by means of different kinds of bonding. The water molecule H2O is formed through covalent bonding of 2 hydrogen atoms with 1 oxygen atom. Salts. Molecules of salts such as table salt, NaCl or CaSO4 do not exist individually. They join together to form visible solid crystals. The crystals are formed through ionic bonding between the positively charged sodium and negatively charged chloride. However when NaCl crystals are added to water, each sodium ion and each chloride ion is surrounded by a shell of water molecules. These “hydration shells” keep the ions separated or dissociated and allow the salt to dissolve readily in water to form individual ions. Other substances such as sugar also dissolve readily in water although they contain no ionic bonds. The reason is that these organic substances are made up of polar molecules containing chemical groups with a net electric charge, or polarity. Organic compounds such as sugar do not form ions when they dissolve, but the molecules as such go into solution. Molar mass The gram mole or mole is derived from the concept of the “chemical amount” of a substance. It is convenient for the purpose of calculating the mass of substances in a reaction to group atoms or molecules in so-called counting units that contain the same number of atoms or molecules. This counting unit is called the mole. One mole of a compound e.g. water contains 6,023 x 10 23 molecules and is a quantity with a mass in gram equal to the molecular mass or formula mass of that compound. For example the formula mass of water is 18,02 g/mol, calculated as follows. Atomic mass of hydrogen: H = 1.008 g/mol (from chemical tables) Atomic mass of oxygen O = 15,9994, usually rounded to 16g/mol Formula mass of H2O = (2 x 1,008) + (1 x 16) = 18,02 It is significant that one mole contains the same number of molecules, atoms or ions whatever the compound or element. This number of molecules is called Avogadro’s number and is approximately equal to 6,023 x 10 23. By definition one mole of a substance contains an Avogadro’s number of atoms, molecules, or ions. 39 23 Therefore, 6,023 x 10 atoms of oxygen is equal to 15,9994 g O (usually taken as 16); and 6,023 x 10 23 molecules of oxygen is equal to 31,998 g O2 (usually taken as 32). Similarly, 6 x 10 23 molecules of water is equal to 18,02 g H2O (usually taken as 18). The molecular mass of a substance (MM) is the mass in grams of one mole of particles (atoms, molecules or ions) of that substance (element, compound or ion). A one molar solution consists of one mole of a substance dissolved in water to make a solution of one litre and is indicated as 1 M. A one molal solution on the other hand, consists of one mole of a substance dissolved in 1 kg of pure water. Note: Molal is 1 mole + 1 kg water; Molar is 1 mole made up to 1 litre of water Organic compounds have carbon as a main element in their composition. They behave differently when they go into solution. They mostly do not dissolve as ions but go into solution as molecules of the compound. A variety of organic compounds (organics) can be present in water. These include natural organic compounds such as algae- and bacterial by-products, carbohydrates and proteins, synthetic organic compounds such as pesticides and herbicides, and products formed during water treatment such as chloroform and other chlorinated products. These organic compounds are usually present in natural waters at very low concentrations, but they may be harmful even at low concentrations. It is not always possible to determine each individual organic compound in water and it would also be very costly because of the large variety of compounds that may be present at very low concentrations. The determination of aggregate parameters such as Dissolved Organic Carbon (DOC) or Chemical Oxygen Demand (COD) gives an indication of the general organic quality of the water. CONCENTRATION UNITS It is extremely important that an analysis, or the dosage of treatment chemicals, or the result of a calculation must be reported as the value together with the relevant units. There are different units in which concentration or dosages can be reported and if the units are not stated the result can be interpreted using the wrong units. SI units are to be used where possible. However, non-SI units are often used in the literature and some are also useful for specific purposes. It is therefore important to be aware of the different non-SI units as well. The amount of dissolved substances in water is most commonly expressed in concentration units, i.e. mass per volume, mostly in milligram per litre (mg/l). There are however, also other concentration units that are used for certain purposes and for certain constituents as shown below. 40 Concentration units in (physical) mass per volume Milligram per litre, mg/l – most commonly used concentration unit for most substances in water. Parts per million, ppm - for dilute solutions ppm is practically identical to mg/l, but it is not part of the SI system and should preferably not be used. Microgram per litre, Pg/l - used for very low concentrations. 1000 Pg/l = 1 mg/l. Parts per billion, ppb –for dilute solutions ppb is practically identical to Pg/l, but is not part of the SI system and should preferably not be used. Gram per litre, g/l – used for high concentrations. 1000 mg/l = 1 g/l. Percentage, % - used for high concentrations, e.g. chemicals. Percentage is similar to parts per hundred, or pph or g/100g. Concentration units in molar mass units per volume For certain purposes it is essential to express concentration in chemical mass units, rather than physical mass units. For example, for calculations involving titration, solubility- and reaction equilibria the units of moles per volume must be used. Molarity Number of moles of solute per litre of solution, mol/l. One mole of a substance is defined as an amount consisting of 6,023 x 1023 particles of the substance (atoms, ions molecules or formula units). The mass of one mole of a substance can conveniently be calculated as the molecular or formula mass expressed in grams. Concentration in mol/l is known as the molar concentration or molarity of the solution. Millimoles per litre, mmol/l where one millimole is 1/1000 of a mole or the mass of one millimole calculated as the molecular or formula mass expressed in mg. Example: Calculate the molar mass of hydrated lime, Ca(OH)2 Molar mass of Ca(OH)2 = 40 + 2(16+1) = 74 g/mol (or mg/mmol) A 1 Molar (1M) solution of Ca(OH)2 will therefore contain 74 g/l of Ca(OH)2 A solution containing 74 mg/l of Ca(OH)2 has a concentration of 0,001 M Concentration units in chemical equivalents per volume The Normality (concentration expressed in chemical equivalents per litre) is a measure of the reacting power of a solution. One equivalent of one substance always reacts with one equivalent of another substance. 41 The equivalent mass of a compound is that mass of the compound which contains one mole of available hydrogen or its chemical equivalent. The equivalent mass of a compound can be determined as follows: Equivalent mass = Molar mass / z, where z is a factor which depends on the chemical reaction involved. For acids the value of z is equal to the number of moles of H+ displaceable from one mole of acid, e.g. for HCl, z = 1, while for H2SO4, z = 2. For bases, the value of z is equal to the number of moles of H+ with which the base will react. For NaOH, z = 1 and for Ca(OH)2 , z = 2. For oxidation/reduction reactions, the value of z equals the change in oxidation number of the particular compound involved in the reaction. For other compounds such as salts that are not involved in oxidation/reduction reactions, the value of z equals the valence (bonding ability) of the compound. Concentration in equivalents of solute per litre of solution, eq/l. This is known as the normality of the solution. A 1 Normal solution contains 1 equivalent of solute per litre of solution. A 0,001 N solution contains 1 milli-equivalent of solute per litre of solution, meq/l, where one milli-equivalent is equal to 1/1000 of an equivalent. Example Calculate the equivalent mass of hydrated lime, Ca(OH)2 Molar mass of Ca(OH)2 = 40 + 2x(16+1) = 74 g/mole There are 2 (OH) groups, z = 2 Equivalent mass of Ca(OH)2 = 74/2 = 37 g/mole A one Normal solution (1N) contains 1 equivalent or 37 g of Ca(OH)2/l of water. A solution containing 37 mg of Ca(OH)2/l of water has a concentration of 0,001N Concentration units in mass per volume expressed on a defined basis The concentration of a substance may be expressed on the basis of an element or ion in the compound. The concentration of nitrate (NO3-), for example, may be expressed either as mg/l NO3- or as mg/l N (NO3--N). Similarly, the concentration of phosphate may be expressed as mg/l PO43- or mg/l P. The conversion between the different nitrogen units from mg/l NO3- to mg/l as N or from mg/l N to mg/l NO3- is done by multiplying the measured concentration by the ratio of the equivalent mass of N/ NO3- or NO3-/N respectively. 42 Example Express 25 mg/l NO3- as mg/l N is as follows: 25 mg/l NO3- x (Equivalent mass of N / Equivalent mass of NO3- ) = 25mg/l NO3- x 14,01/62.01 = 5,65 mg/l N, where 14,01 is the equivalent mass of nitrogen and 62,01 is the equivalent mass of NO3-. To convert 5 mg/l N to mg/l NO3- 5 mg/l N x 62,01/14,01 = 22,13 mg/l NO3- The concentration in milligram per litre of one compound can also be expressed in terms of mg/l of another compound. In the case of hardness, alkalinity and the concentration of chemicals used for softening of water, concentrations are conventionally expressed in terms of mg/l calcium carbonate (mg CaCO3/l). This means that the concentration given is not that of the element, ion or compound itself but it is converted to equivalent amounts of CaCO3 units. Example Convert 75mg/l calcium as Ca2+ to mg/l as CaCO3. This is done by multiplying 75 mg/l Ca2+ by the ratio of the equivalent mass of CaCO3 to the equivalent mass of Ca2+: 75 mg/l Ca2+ x (50/20) = 187,5 mg/l CaCO3, Where 50 is the equivalent mass of CaCO3 (100/2), and 20 is the equivalent mass of Ca2+ (40/2). Table A3.1 gives a summary of the conversions between different sets of units for chemicals normally involved in water treatment. Table A3.1 Chemical. Molar scale. Equivalent scale. Parameter Concentration of Concentration in X mmol/l Meq/l parameter in mg/l mg/l CaCO3 Ca(OH)2 X/74 X/37 Ca2+ or OH- X * 50/37 Lime CO2 X/44 X/22 CO2 X * 50/22 Carbon dioxide Na2CO3 X/106 X/53 CO32- X * 50/53 Soda ash NaOH X/40 X/40 OH- X * 50/40 Caustic soda H2SO4 X/98 X/49 H+ X * 50/49 Sulphuric acid HCl X/36 X/36 H+ X * 50/36 Hydrochloric acid 43 CaCO3 X/100 X/50 Ca2+ or CO32- X * 50/50 Calcium carbonate Cl- X/35 X/35 Cl- X * 50/35 Chloride SO42- X/96 X/48 SO42- X * 50/48 Sulphate Cation / Anion Balance A cation-anion balance is an easy way to check the completeness or accuracy of a water analysis. Since water cannot have a net electrical charge, the total amount of cations must equal the total amount of anions, when expressed on a specific basis. The balance can only be done in terms of equivalents or milli-equivalents per litre, or when concentration is expressed as mg/l CaCO3. These two ways of expression of the concentration of a solution takes the reacting capacity into account, which is not the case for mg/l as ion or moles/l. Example: A water analysis giving the major inorganic constituents (cations and anions) and their concentrations in mg/l is shown in the first two columns of the Table below. Columns 1 and 2 give the ion and concentration in mg/l of the ion. Column 3 shows the molar, or formula mass of each ion in g/mole or mg/mmole, given in the Periodic Table of Elements and in Chemistry handbooks. Column 4 shows the molarity (in mmol/l) calculated by dividing mg/l of ion by the molar mass. For example, for Ca2+ : (107 mg/l) / (40,078 mg/mmole) = 2,68 mmol/l. Column 5 shows the equivalent mass calculated by dividing the molar mass by the valence of the ion (z). For example, for Ca2+: 40,078 / 2 = 20,039 g/eq, or mg/meq. Column 6 shows the normality in meq/l calculated by dividing mg/l of ion by the equivalent mass. For Ca2+: (107 mg/l) / (20,039 mg/meq)= 5,35 meq/l. Column 7 shows the concentration expressed as mg/l CaCO3 calculated by multiplying mg/l of ion by the ratio of the equivalent mass of CaCO3 (50,044 mg/meq) to equivalent mass of ion. For example, Ca2+: 107 x 50,004 / 20,039 = 267 mg/l as CaCO3. 44 Cations mg/l as Molar mass Molarity Equivalent Normality mg/l as ion g/mole mmoles/l mass g/eq meq/l CaCO3 Ca2+ 107 40,078 2,67 20,039 5,34 267 Mg2+ 20 24,305 0,82 12,152 1,64 82 Na+ 50 22,990 2,17 22,990 2,17 109 K+ 15 39,098 0,38 39,098 0,38 19 Total cations 9,53 477 mg/l as Molar mass Molarity Equivalent Normality mg/l as Anions ion g/mole mmoles/l mass g/eq meq/l CaCO3 Total 260 61,017 4,26 61,017 4,26 213 Alkalinity as HCO3- SO42- 117 96,064 1,22 48,032 2,44 122 Cl- 90 35,453 2,54 35,453 2,54 127 NO3- 20 62,005 0,32 62,005 0,32 16 Total anions 9,56 478 % Difference= 100 (Total anions- 0,3% 0,2% total cations)/ Total cations A check on the accuracy of the analysis is done by comparing the total cation concentration in meq/l to the total anion concentration in meq/l (or the concentrations expressed as mg/l CaCO3 ). In this example the totals differ by less than 2 % and the analysis is therefore accepted as accurate. The composition of the water can be represented visually by a bar diagram. The diagram is constructed with cations in the top part starting with calcium, then magnesium and then other cations. The anions are shown in the bottom part starting with alkalinity. The length of each block representing a particular ion is drawn to scale in meq/l or in mg/l as CaCO3. Ca 267 Mg 82 Na 109 19 Alk 213 SO4 122 Cl 127 16 This shows that: Total hardness = 267 +82 = 349 mg/l CaCO3 Calcium carbonate hardness = 213 mg/l CaCO3 (Ca associated with alkalinity) Calcium non-carbonate hardness = 267 – 213 = 54 mg/l CaCO3 (Ca assoc with other ions) Magnesium carbonate hardness = 0 (Mg assoc with alkalinity) Magnesium non-carbonate hardness = 82 mg/l CaCO3 (Mg assoc with other anions) 45 WATER TREATMENT REACTIONS AND EQUATIONS The chemical concepts discussed above form the basis of chemical reactions in water treatment. The reactions are presented by chemical equations that provide a ‘chemical picture’ of the reaction. The equations also allow one to do basic calculations of chemical dosages, the quantity of active species, quantity of products, etc. The purpose of a chemical equation is to express what happens during a chemical reaction. Since matter cannot be created or destroyed but converted from one form to another, the same number of atoms and the same mass of material must be present at the end of a reaction compared to the beginning of the reaction. This means that all chemical reactions must be balanced, i.e. all the materials that are shown on the left side of an equation must also be shown on the right side. In this part the following aspects are considered: general reactions, acid-base reactions, oxidation reactions, precipitation reactions, and the carbonate system in water. General chemical reactions A general chemical reaction is one in which no oxidation-reduction reaction takes place. An example of a general reaction in water treatment is the addition of a coagulant such as ferric chloride to water. The ferric chloride reacts with the water and forms certain products. The reaction can be represented by the following equation: FeCl3 + 3H2O → Fe(OH)3 + 3HCl (Equation 1) This equation is balanced because the same number of atoms of each element appears on the left and right side of the equation. This equation means that 1 mole of ferric chloride reacts with 3 moles of water to form 1 mole of ferric hydroxide and 3 moles of hydrochloric acid. The ferric hydroxide will normally precipitate as a solid and this can be indicated by adding a vertical arrow next to it, Fe(OH)3↓. The HCl is a strong acid and will react with the alkalinity in the water represented as calcium bicarbonate: 2HCl + Ca(HCO3)2 → CaCl2 + 2H2O + 2CO2 (Equation 2) This means that 2 moles of hydrochloric acid react with 1 mole of alkalinity (destroy 1 mole of alkalinity) and produce 2 moles of CO2 and 1 mole of calcium chloride. These two reactions take place simultaneously and we can therefore add them to give one overall reaction. Equation 1 x 2: 2FeCl3 + 6H2O → 2Fe(OH)3 + 6HCl (Eq 3) Equation 2 x 3: 6HCl + 3Ca(HCO3)2 → 3CaCl2 + 6H2O + 6CO2 (Eq 4) Eq 3 + 4: 2FeCl3 + 3Ca(HCO3)2 → 2Fe(OH)3 + 3CaCl2 + 6CO2 46 The overall equation shows that 2 moles of ferric chloride react with 3 moles of alkalinity and produce 2 moles of ferric hydroxide and 3 moles of calcium chloride and 6 moles of carbon dioxide. Similar equations can be written for other coagulants. For aluminium sulphate (Alum, Al2(SO4)3. 18 H2O ): This indicates that the 18 molecules of water are attached to the Al2(SO4)3 but do not participate in the chemical reaction. The water molecules must, however, be taken into account when a solution of Al2(SO4)3 with a specific concentration is prepared. Al2(SO4)3. 18 H2O + 3Ca(HCO3)2 → 2Al(OH)3 + 3CaSO4 + 6CO2 + 18 H2O For hydrated lime Ca(OH)2 : Ca(HCO3)2 + Ca(OH)2 → 2CaCO3 + 2H2O These balanced equations provide information on what happens during the chemical reactions and they can be used to calculate different items such as such as the mass of calcium carbonate consumed as a result of a certain dosage of ferric chloride, or the mass of ferric hydroxide sludge produced in the reaction. It must be noted however, that the equations given above are actually simplified representations since there are different intermediate species formed during the reactions depending on conditions such as pH. EQUILIBRIUM CHEMISTRY Acids such as HCl and H2SO4 are termed strong acids and dissociate completely when added to water. HCl H+ + Cl- This means that when HCl is added to water, there is virtually no HCl in solution, only H+ and Cl-. However acids such as carbonic acid (H2CO3) and acetic acid (CH3COOH) are weak acids and dissociate only partially when added to water. H2CO3 ↔ H+ + HCO3- This means that there is an appreciable amount of H2CO3 in solution in addition to H+ and HCO3. The extent of the reaction or the degree to which the reaction proceeds, is given by the dissociation constant: K = [H+] [HCO3-] / [H2CO3] = 10-6,3 mole/l Where the [ ] indicates concentration in mole/l , and the value of K is constant for the specific substance at a certain temperature. K values for most weak acid-base systems are given in textbooks. 47 Water is also a weak acid. H2O ↔ H+ + OH- K = [H+][OH-] / [H2O] = 1,8 x 10–16 The concentration of water is (1000/18) = 55,5 mol/l, which gives Kw = [H+][OH-] = 10-14 mol/l Where [H+] and [OH-] are the molar concentrations of H+ and OH-. This indicates that the product [H+][OH-] must always be equal to 10-14 mol/l. So, if H+ is added to water (addition of acid) the concentration of OH- will drop so that the product of the two concentrations will remain at Kw = 10-14 mol/l Since the concentration of [H+] and [OH-] can vary over the very wide range of 10 to 10-14 it is convenient to use a logarithmic scale to express the concentration. For this purpose the pH function was introduced as: pH = - log [H+ ] This means that a [H+] concentration of 10-7 mole/l is expressed as a pH of 7. At this pH the concentration of [OH-] must therefore also be 10-7 mol/l in order to maintain the value of Kw. This value is therefore referred to as the neutral pH because the concentrations of both ions are equal. Similarly a [H+] concentration of 10-3 mol/l is expressed as a pH of 3, and the [OH-] concentration must then be 10-11 mol/l. All pH values below 7 are acidic and those above 7 are alkaline. Because the pH scale is logarithmic, it means that a change of one pH unit is equal to a 10 times increase in concentration of the one type of ion and a 10 times decrease in concentration of the other type. The carbonate system Water contains various dissolved species and ions. These species interact with one another and with water. One of the most important systems in water that has an important effect on water treatment is the carbonate system. The species that make up the carbonate system are the following: CO2, H2CO3, HCO3-, CO32-, H+ and OH-. Carbon dioxide is present in the atmosphere and it dissolves in water to form the weak acid, carbonic acid, which then dissociates into the hydrogen and bicarbonate ions according to reaction. CO2 + H2O ↔ H2CO3 ↔ H+ + HCO3- Ka = [H+] [HCO3-] / [H2CO3] = 10-6,3 mole/l 48 The bicarbonate ion in turn, dissociates to form H+ and the carbonate ion CO32- HCO3- ↔ H+ + CO32- Kb = [H+] [CO32-] / [HCO3-] = 10-10,3 mole/l The total carbonic species in solution is represented by Ct and is: Ct = [H2CO3] +[ HCO3-] +[ CO32-] The relative amount of each species is a function of the pH of the water. The equations above can be used to calculate the relative amounts and plot these values on a so-called pH-pC diagramme. This type of diagramme shows the log of the concentration (pC) of the different species as a function of pH. Figure A3.1 shows an example of the log C vs pH plot for the carbonate system. This shows how the concentration of the different species (as log) varies as a function of pH. It shows that at low pH levels the dominant species is H2CO3 up to a pH of about 6. At pH 6,3 the concentrations of H2CO3 and HCO3- are equal and thereafter the HCO3- species is dominant up to a pH of about 10. At pH of 10,3 the concentrations of HCO3- and CO32- are equal and at higher concentrations CO32- dominates. The figure also shows how the concentrations of H+ and OH- vary as a function of pH. Figure A3.1 Log [concentration] - pH for carbonate system 49 Alkalinity and acidity As is indicated above, alkalinity reacts with hydrochloric acid that forms upon addition of ferric chloride to water. Alkalinity can therefore be regarded as the acid neutralising capability of water. Alkalinity is determined by titration of a water sample with a standardised strong acid to a specific pH end point. The end point to which titrations are carried out, are selected to give an indication of the type of alkalinity (or the carbonate species) that is present. Total Alkalinity is determined by titration to the end point where all the species contributing to alkalinity (OH-, CO32- and HCO3-) have been neutralised. This end point is the methyl-orange end point, which is approximately pH 4,5 depending on the initial condition of the water. This alkalinity is termed Total Alkalinity or M- Alkalinity, or H2CO3 Alkalinity. Titration to the phenolphthalein endpoint of pH 8,4 gives the contribution of hydroxides and carbonates to alkalinity. This is termed P- or Phenolphtalein alkalinity or HCO3- Alkalinity. Total Alkalinity can also be calculated from the following relationship: T Alkalinity = (OH-) + (HCO3-) + (CO32-) – (H+) The buffer capacity is closely related to the amount of alkalinity in the water. Buffer capacity is the ability of the water to absorb acid without a substantial decrease in pH. Acidity is defined as the base neutralising capacity of water and is determined by titration of the water with a standardised strong base to certain pH end points. Acidity can be calculated from the following expression: Acidity = (H2CO3)+ (HCO3-) + (H+)- (OH-) Where these terms are expressed in mg/l as CaCO3. This end point is more difficult to determine and acidity is therefore only used for specific purposes. SOLUBILITY EQUILIBRIA The solubility of salts in water varies widely. Certain salts have very high solubilities, e.g. NaCl, while others have very low solubilities. This characteristic of low solubility is often used to separate such a compound from water by precipitating it as an insoluble precipitate. When the concentration of ions of a sparingly soluble salt is increased beyond a certain point in solution the salt will start to precipitate from solution. For example when the Ca2+ and CO32- concentration is increased beyond the maximum solubility, solid CaCO3 will precipitate. Ca2+ + CO32- ↔ CaCO3↓ 50 This reaction forms an equilibrium that can be quantified by the following expression: Ks = [Ca 2+] [CO32-] Where Ks is the equilibrium constant or the solubility product for CaCO3, and [Ca2+] and [CO32-] represent the molar concentrations of the cation and anion respectively. Ks represents the maximum value the product of the ion concentrations can have for a given set of conditions. If [Ca2+] [CO32-] < Ks the solution is under-saturated and no precipitate will form. If [Ca2+] [CO32-] > Ks the solution is super-saturated and CaCO3 will precipitate until the ion product just equals Ks. The smaller the value of Ks the lower the solubility of the compound and the easier it is to remove the compound to a very low concentration. For example the Ks value for Fe(OH)3 is equal to 10-37 mol/l. This very small value means that ferric hydroxide will precipitate almost completely from water. However, the solubility depends greatly on pH and it is therefore important to maintain the pH of water within the correct range to ensure as complete as possible precipitation of a substance. Example Calculate the concentration of Ca2+ (aq) in a saturated CaSO4 solution in water in mg/l if the Ksp = 1.9 x 10-4 CaSO4 (s) ↔ Ca2+ (aq) + SO42- (aq) For each mole of CaSO4 (s) that dissolves, one mole of Ca2+ (aq) and one mole of SO42- (aq) are formed Ksp = [Ca2+] [SO42-] = 1.9 x 10-4 Say the molar concentration of each ion = y, then y2 = 1.9 x 10-4 y = 0.0138 mole/l 1 mole of Ca2+ = 40 g/mole, so y = 0.0138 mole/l x 40 g/mole = 0.552 g/l or 552 mg/l Common ion effect The solubility behaviour of a precipitate in a solution that contains a common chemical species is important in process chemistry. It forms the basis of precipitation processes for removal of heavy metals and for softening of hard waters. 51 In the reaction Ca2+ + CO32- ↔ CaCO3 the addition of either Ca2+ or CO32- ions will result in the value of the ion product being greater than Ks. [Ca2+] [ CO32-] > Ks and precipitation will occur until a new equilibrium is established where the ion product is just equal to Ks. Ks = [Ca2+] [CO3]2- = 4,7 x 10-9 The concentration of calcium and carbonate at equilibrium can be calculated as follows: Since 1 mole of calcium reacts with 1 mole of carbonate we can write: [Ca2+] [CO32-] = [Ca2+]2 = 4,7 x 10-9 [Ca2+] = 0,0000686 mol / l Ca2+ = 0,0000686 x 40,08 x 1 000 mg/l Ca2+ = 2,75 mg/l, which is the theoretical maximum value of the ions in solution before precipitation will commence. In practice however, these low values are not easily achieved. The solubility behaviour of most slightly soluble salts is much more complicated than is suggested by this example. Complex formation, hydrolysis, pH and other phenomena have important effects on solubility behaviour. Example Calculate the residual magnesium concentration that exists in a saturated magnesium hydroxide solution if enough sodium hydroxide has been added to the solution to increase the equilibrium pH to 11,0. Mg(OH)2(s) ¾ Mg2+ + 2OH- The solubility product constant for this reaction Ksp = 1,2 x 10-11. Determine the hydroxide ion concentration: Kw = [H+][OH-] = 10-14 at 25°C Because the pH is 11, [H+] = 10-11 mole/l we know that [OH-] = 10-14 / 10-11 = 10-3 mole/l 52 Establish the solubility product constant expression, and solve for the magnesium ion concentration: Ksp = [Mg2+][OH-]2 [Mg2+] = (1,2 x 10-11) / (10-3)2 = 1,2 x 10-5 mol/l or 0,29 mg/l Since hardness ion concentrations are normally expressed as mg/l CaCO3, multiply the concentration by the ratio of the equivalent weights: 0,29 x 50 / 12.2 = 1,2 mg/l as CaCO3 Softening. The concepts of solubility, common ion effect, hardness and the carbonate system can be illustrated by the reactions that take place during chemical softening of water. Softening involves the removal of calcium and magnesium ions from the water. Softening can be achieved by means of ion exchange or nanofiltration or by means of chemical precipitation. In the chemical precipitation process calcium is precipitated as calcium carbonate, CaCO3 and magnesium as magnesium hydroxide, Mg(OH)2, since these compounds have low solubility product values of 4,7 x 10-9 for CaCO3 and 8,9 x 10 –12 for Mg(OH)2 respectively. In order to precipitate CaCO3 the solubility product must be exceeded. This means that the concentration of calcium and carbonate ions must be increased. This can be achieved by the addition of hydrated lime, Ca(OH)2 which will increase the Ca2+ concentration and at the same time increase CO32- concentration (for calcium carbonate hardness). The CO32- concentration is increased as a result of the increase in pH to about 10,3 (due to the addition of OH- ions) which causes the bicarbonate ions HCO3- to be converted to CO32- as can be seen in the pH-pC diagram of the carbonate system. Ca(HCO3)2 + Ca(OH)2 → 2CaCO3↓ + 2H2O In the case of calcium non-carbonate hardness the ions associated with calcium cannot be converted to carbonate ions and a different strategy must be used. In this case lime is first added to increase the pH and add calcium to precipitate whatever carbonate hardness is present. In order to utilise the common ion effect a source of carbonate ions must be added and sodium carbonate, Na2CO3, also called soda ash is used. Ca SO4 + Na2CO3 → CaCO3↓ + Na2SO4 In the case of magnesium hardness the magnesium is precipitated as magnesium hydroxide, Mg(OH)2. The minimum solubility of magnesium hydroxide is at a pH of about 11, so more lime must be added to increase the pH and at the same time add more OH- ions in order to exceed the solubility product. Mg(HCO3)2 + 2Ca(OH)2 → 2CaCO3↓ + Mg(OH)2↓ +2H2O 53 In the case of magnesium non-carbonate hardness lime must be added to increase the pH and supply OH- ions and then soda ash must be added to precipitate the excess calcium ions added. MgSO4 + Ca(OH)2 → Mg(OH)2↓ + Ca SO4 Ca SO4 + Na2CO3 → CaCO3↓ + Na2SO4 Recarbonation. When water has been treated by the addition of lime and soda ash the pH is normally around 11 and the water has to be stabilised before distribution. This is normally achieved through the addition of CO2 to reduce the pH to the level where the water is stable with respect to CaCO3 (normally at a CCPP of 4 mg/l and a pH level of 7 to 8,5). Recarbonation involves the addition of CO2 to neutralise excess OH- ions and to convert carbonate ions to the bicarbonate form to reduce the scale-forming potential of the water. This can be done in a two-stage or one-stage process. In a two-stage process CO2 addition takes place to neutralise Ca(OH)2 and precipitate excess CaCO3 which is removed in a settler. Further addition of CO2 converts CO32- to HCO3- to prevent precipitation of CaCO3 on filter sand or in distribution systems. Ca(OH)2 + CO2 → CaCO3↓ + H2O CaCO3 + CO2 + H2O → Ca(HCO3)2 OXIDATION – REDUCTION REACTIONS In oxidation-reduction or redox reactions some of the atoms or ions undergo a change in oxidation number, which implies that electrons are gained or lost by the atoms. Oxidation is the loss of electrons resulting in an increase in oxidation number of one or more atoms. Reduction is a gain in electrons resulting in a decrease in oxidation number. The two processes must always occur simultaneously. An oxidising agent is an atom, ion or molecule that takes up electrons from other substances. An example of a redox reaction is the oxidation by chlorine of soluble manganese (ii) and the subsequent removal of manganese by precipitation of manganese (iv) dioxide. Some elements such as iron and manganese may have different oxidation numbers depending on the compound or the reaction in which it is involved. The oxidation number is calculated on the basis that the net charge of a molecule must be zero, and that of an ion must be the charge that the ion carries. For example, the oxidation number of manganese in MnO2 is +4 because the oxidation number of oxygen is always –2 and since there are 2 O atoms for each Mn atom [2 x (-2) = (-4)] Cl2(g) + Mn2+(aq) + 2H2O(l) MnO2(s) + 2Cll-(aq) + 4H+(aq) ox number: 0 +2 +1 -2 +4 -2 -1 +1 54 Electrons are negatively charged entities. This means that when an atom looses one electron, it looses one negative charge, or gains one positive charge. Manganese looses 2 electrons per Mn atom when changing from the +2 state in Mn2+ to the +4 state in MnO2. Since oxidation is defined as a loss of electrons Mn2+ is oxidized to the +4 state. Chlorine on the other hand gains one electron per atom when changing from molecular chlorine to the chloride ion and is therefore reduced in the process. Chlorine is the oxidising agent that oxidises manganese and in the process chlorine is reduced. Since electrons cannot be destroyed, the total gain in electrons must always equal the total loss in electrons. The 2 electrons lost by Mn2+ equal the 2 electrons gained by the 2 chlorine atoms, each gaining one electron. Oxidation-reduction reactions play an important role in water treatment. Examples include the oxidation of ammonia by chlorine in the breakpoint reaction. Another example is the oxidation of iron and manganese by oxygen in the air, or by chlorine, or potassium permanganate. A further example is the oxidation of organic material causing taste and odour in the water by chlorine or ozone. The oxidising agents (or oxidants) used in water treatment include oxygen, chlorine, ozone, chlorine dioxide, potassium permanganate, hydrogen peroxide. Chlorine chemistry Chlorine is a strong oxidising agent and it reacts and oxidises some of the essential systems of micro-organisms thereby inactivating or destroying them. The different forms in which chlorine is used for disinfection, have different oxidising powers and this must be taken into account to ensure effective disinfection. Chlorine is normally used for disinfection on large plants in the gaseous form, but calcium hypochlorite and sodium hypochlorite are two other chlorine compounds that can also be used for disinfection. Chlorine gas, Cl2 dissolves in water to form hypochlorous and hydrochloric acid. Cl2 + H2O → HOCl +HCl The actual disinfecting agent is hypochlorous acid which dissiociates as follows: HOCl H+ + OCl- The H+ formed in the process reacts with alkalinity and leads to a reduction in alkalinity, and may lead to a reduction in pH if insufficient alkalinity is available. The chlorine species in the form of hypochlorous acid, HOCl plus the hypochlorite ion, OCl- are termed free available chlorine. Chlorine in the form of monochloramine (together with other chloramine species) is termed combined available chlorine (see below under break point chlorination). 55 HOCl is much more effective for disinfection than the hypochlorite ion - about 80 (or even more) times more effective. The relative quantities of these two species are determined by the pH of the water. At pH below 7, HOCl is the predominant species while at pH above about 7,5 hypochlorite ion, OCl- predominates. It is therefore important that the pH of the water be taken into account when determining the required chlorine dosage for disinfection. Calcium hypochlorite (commonly known as HTH) dissolves in water as follows: Ca(OCl)2 = Ca2+ + 2 OCl- the hypochlorite ion hydrolises to form HOCl: OCl- + H2O → HOCl + OH- The OH- formed in the process results in an increase in pH (in contrast to chlorine gas where H+ is formed). Sodium hypochlorite, NaOCl (commonly known as household bleach under different brand names) is available as a solution. Water treatment sodium hypochlorite contains 12 to 13% of hypochlorite, which is equivalent to 10 - 12 % available chlorine. Bleach contains about 6 – 8% free available chlorine. Sodium hypochlorite is relatively unstable and deteriorates fairly rapidly, especially when exposed to sunlight. NaOCl → Na+ + OCl- OCl- + H2O → HOCl + OH- Monochloramine (so-called combined available chlorine) is also used for water disinfection. It is formed when HOCl is added to water that contains a small amount of ammonia. The ammonia reacts with HOCl to form monochloramine, NH2Cl. It is much less effective as a disinfectant than HOCl (the same order of effectiveness as chlorite ion). However, it has the advantage of being much more stable in water than free available chlorine. For this reason it is often used to provide residual protection in larger distribution systems. An important concept in disinfection by means of chlorine is that of breakpoint chlorination. Breakpoint chlorination refers to the reaction between chlorine and ammonia in water. When chlorine is added to water it reacts with ammonia. In this process chloramines are formed. When more and more chlorine is added it breaks down (oxidises) the chloramines and then forms free available chlorine and nitrogen. The point where all chloramines have been oxidised is termed the breakpoint. The important point is that only after the break point has been reached, free available chlorine is formed resulting in effective disinfection. Chlorine also reacts with other compounds that exercise a chlorine demand such as certain organics, iron and manganese. These reactions do not participate in the breakpoint reaction. Figure A3.2 shows a typical breakpoint curve. 56 The following reactions take place when chlorine is added to water containing ammonia: HOCl + NH3 → NH2Cl + H2O (monochloramine) HOCl + NH2Cl → NHCl2 + H2O (dichloramine) HOCl + NHCl2 → NCl3 + H2O (trichloramine or nitrogen trichloride) 2 NHCl2 + HOCl → N2 + 3 HCl + H2O Figure A3.2 Breakpoint curve Zone 1 Zone 2 Zone 3 Ammonia-N concentration Chlorine concemtration Total Chlorine applied Ammonia-N Conc. Hump Measure Chlorine residual Breakpoint 0.0 1.0 2.0 3.0 4.0 5.0 6.0 7.0 8.0 9.0 Cl2:NH4+-N Mass Ratio 57 CHAPTER A 4: WATER TREATMENT CALCULATIONS Frik Schutte INTRODUCTION There are many types of calculations that treatment plant operators and process controllers must be able to perform. These include: Calculation of mass of chemicals to make up chemical solutions of a specified concentration Calculation of chemical dosages Calculation of species concentration Calculation of the mass and volume of wastes produced Different types of problems and calculations require different approaches and different techniques. For example, to calculate the mass or concentration of certain species in a reaction, the starting point is to compile a balanced equation of the chemical reaction. For the calculation of species concentration in precipitation reactions, equilibrium equations and expressions are the starting point. For calculation of dosages and material flow, the starting point may be to compile a mass balance of the system. The different approaches are discussed in the following sections and are illustrated by examples. CALCULATIONS FROM CHEMICAL EQUATIONS Chemical treatment processes can be represented by chemical reactions and these reactions form the basis of many chemical calculations. Reactions are presented by chemical equations that provide a ‘chemical picture’ of the reaction. The equations allow one to do basic calculations of chemical dosages, the quantity of active species, quantities of products, etc. The purpose of a chemical equation is to express what happens during a chemical reaction. Since matter cannot be created or destroyed but only be converted from one form to another, the same number of atoms and the same mass of material must be present at the end of a reaction compared to the beginning of the reaction. This means that all chemical reactions must be balanced, i.e. all the materials that are shown on the left side of an equation must also be shown on the right side. As an example let us review the reaction when ferric chloride is added to water as coagulant: FeCl3 + 3H2O → Fe(OH)3 + 3HCl (Equation 1) 2HCl + Ca(HCO3)2 → CaCl2 + 2H2O + 2CO2 (Equation 2) Equation 1 x 2: 2FeCl3 + 6H2O → 2Fe(OH)3 + 6HCl (Eq 3) Equation 2 x 3: 6HCl + 3Ca(HCO3)2 → 3CaCl2 + 6H2O + 6CO2 (Eq 4) 58 These two equations can be added to give the overall reaction: Eq 3 + 4: 2FeCl3 + 3Ca(HCO3)2 → 2Fe(OH)3 + 3CaCl2 + 6CO2 The overall equation shows that 2 moles of ferric chloride react with 3 moles of alkalinity and produce 2 moles of ferric hydroxide and 3 moles of calcium chloride and 6 moles of carbon dioxide. Similar equations can be written for other coagulants. For aluminium sulphate (Alum, Al2(SO4)3.18 H2O ): Al2(SO4)3.18 H2O + 3Ca(HCO3)2 → 2Al(OH)3 + 3CaSO4 + 6CO2 + 18 H2O For lime Ca(OH)2 : Ca(HCO3)2 + Ca(OH)2 → 2CaCO3 + 2H2O These equations are the starting place for most of the general types of calculations performed on water treatment chemicals and related aspects. The following examples illustrate these types of calculations. Example 1: How much alkalinity is consumed by the addition of 15 mg/l FeCl3 to water? From the balanced equation above: 2FeCl3 + 3Ca(HCO3)2 → 2Fe(OH)3 + 3CaCl2 + 6CO2 3 moles of Ca(HCO3)2 are consumed by 2 moles of FeCl3 Mol Mass of FeCl3 = 55,8 + (3 x 35,5) = 162,3 mg/mmol Mol Mass of Ca(HCO3)2 = 40 + 2[1 + 12 + (3 x 16)] = 162 mg/mmol That is: 3 x 162 = 486 mg/l Ca(HCO3)2 in solution is consumed by 2 x 162,3 = 324,6 mg/l FeCl3 That is: 486/ 324,6 = 1,5 mg/l Ca(HCO3)2 is consumed per mg/l FeCl3 For 15 mg/l FeCl3 : 15 x 1,5 = 22,5 mg/l alkalinity as Ca(HCO3)2 is consumed. Alkalinity is normally expressed as CaCO3, and the conversion can be done as follows: 59 22,5 mg/l as Ca(HCO3)2 = 22,5 x Eq mass CaCO3 / Eq mass Ca(HCO3)2 = 22,5 x 50/81 = 13,9 mg/l CaCO3 The implication is that if the water contains less alkalinity than 22,5 mg/l as Ca(HCO3)2 or 13,9 mg/l as CaCO3 , the pH will drop as a result of the dosing of 15 mg/l FeCl3. Example 2: Water with a low alkalinity of 12 mg/l as CaCO3 is to be treated with alum together with lime to add alkalinity. Determine the amount of lime as CaO to be added to prevent the pH from dropping if the alum dosage is 55 mg/l. Al2(SO4)3 reacts with lime as follows: Al2(SO4)3. 18 H2O + 3 Ca(OH)2 → 2 Al(OH)3 + 3 CaSO4 + 18 H2O Al2(SO4)3 reacts with natural alkalinity as follows: Al2(SO4)3. 18 H2O + 3 Ca(HCO3)2 → 2 Al(OH)3 + 3 CaSO4 + 6 CO2 + 18 H+ Amount of alum that will react with natural alkalinity: Alkalinity = 12 mg/l as CaCO3 x 162,1 mg/l Ca(HCO3)2/ 100 mg/l CaCO3 = 19,4 mg/l Ca(HCO3)2 MW of alum = 27 x 2 + 3[32 + (16 x 4)] + 18(2 + 16) = 666 MW of Ca(HCO3)2 = 162,1 1 mole of alum reacts with 3 moles of alkalinity Therefore the mass of alum to react with alkalinity = 19,4 x 666/(3x162,1) = 26,6 mg/l alum 60 The amount of alum remaining = 55 – 26,6 = 28,4 mg/l MW of Ca(OH)2 = 40 + 2(16 +1) = 74 MW of CaO = 40 + 16 = 56 1 mole of alum reacts with 3 moles of Ca(OH)2 Ca(OH)2 required = 28,4 mg/l alum x (3 x 74)/666 = 9,48 mg/l Ca(OH)2 9,8 mg/l Ca(OH)2 x 56/74 = 7,2 mg/l CaO Example 3: The FeCl3 dosage determined from jar tests for a certain water source is 20 mg/l. What must the dosage rate be in ml/min of FeCl3 for a dosage of 20 mg/l FeCl3. The FeCl3 is provided as a 43% solution (mass FeCl3/mass solution) with a density of 1,48 kg/l. The flow to the plant is 5000 m3/h. Approach: The dosage in mg/l (mass of pure FeCl3/volume water) must be converted to dosage in ml/min (volume of 43% FeCl3 solution/time). This means the mass pure FeCl3 /volume dosage must first be converted to mass of 43% solution and then to volume of 1,48 kg/l solution. This step is followed by converting volume/volume to volume/time by multiplying with the volumetric flow rate. Step 1: Convert the dosage in mg/l pure FeCl3 to mg/l FeCl3 solution. 20 mg FeCl3 is contained in 20/0,43 mg solution = 46,5 mg solution (46,5 x 43/100=20) Step 2: Convert mg solution to ml solution. 46,5 mg solution = 46,5mg x 1/1,48ml/g x 1/1000 g/mg =0,0314 ml solution. 20 mg/l FeCl3 =46,5 mg/l solution = 0,0314 ml/l solution Step 3: Convert ml solution/l water to ml solution per minute. This means we have to take into account conversion of m3 to litre (1 m3 x 1000 l/m3) and conversion of ml to litre (1 ml x 1/1000 l/ml) and conversion of hours to minutes (1 hr x 60 min/hr). It is most important to make sure that the units are consistent and that they cancel out to give the required units. 0,0314 ml solution/ l water x 1000 l/m3 water = 31.4 ml solution/m3 water 31,4 ml solution/m3 water x 5000 m3 water/ hr x 1/60 hr/min = 2616.7 ml solution/min Dosage is 2616.7 ml FeCl3, solution per minute for a flow of 5000m3/hr 61 Step 2: Convert the dosage in mg FeCl3 solution per litre water to ml FeCl3 solution per litre H2O 101,46 mg/l solution = [101,46 mg solution/l H2O] x [1 ml soln./ 1 480 mg soln.] = 0,069 ml soln./litre H2O soln. = solution Alternative The problem may also be stated that a dosage of 10 mg/l as Fe3+ is required for coagulation. Calculate the dosage for the same conditions given above. Approach: In this case one first has to convert mg/l Fe3+ to mg/l FeCl3 and then follow the same procedure as for the example above. Step 1: Convert mg/l Fe3+ to mg/l FeCl3 : FeCl3 Fe3+ + 3 Cl- [1 mmol FeCl3 (Mol mass 162,3) gives 1 mmol Fe3+ (Mol mass 55,8)] 12 mg Fe3+/l H2O = (12 mg Fe3+/ l H2O) x (1 mmol FeCl3 / 1 mmol Fe3+) x (1 mmol Fe3+/55,8 mg Fe3+) x (162,3 mg FeCl3/ 1 mmol FeCl3 ) = 34,9 mg FeCl3 /l H2O Step 2: Convert mg FeCl3 per litre water to mg FeCl3 solution per litre water as above. 34,9 mg FeCl3 / l H2O] x [100 mg solution/43 mg FeCl3] = 81,16 mg solution per l water Step3: Convert mg solution per litre water to ml solution per litre H2O 81,16 mg solution/l H2O x 1 ml sol./ 1 480 mg sol = 0,0548 ml sol./litre H2O Step 4: For 5000 m3/h: 0,0548 ml sol./litre H2O x 5000 m3/h x 1000 l/ m3 x 1hr/60 min = 4569,9 ml soln./min 62 Example 4: Calculate the solubility of Ca2+ (aq) in a saturated CaSO4 solution in water in mg/l if the Ksp = 1.9 x 10-4 CaSO4 (s) ↔ Ca2+ (aq) + SO42- (aq) For each mole of CaSO4 (s) that dissolves, one mole of Ca2+ (aq) and one mole of SO42- (aq) are formed Ksp = [Ca2+] [SO42-] = 1.9 x 10-4 Say the molar concentration of each ion = y, then y2 = 1.9 x 10-4 y = 0.0138 mole/l 1 mole of Ca2+ = 40 g/mole, so y = 0,0138 mole/l x 40 g/mole = 0,552 g/l or 552 mg/l The concentration of calcium and carbonate at equilibrium can be calculated as follows: Ks = [Ca2+] [CO3]2- = 4,7 x 10-9 Since 1 mole of calcium reacts with 1 mole of carbonate we can write: [Ca2+] [CO32-] = [Ca2+]2 = 4,7 x 10-9 [Ca2+] = 0,0000686 mol / l Ca2+ = 0,0000686 x 40,08 x 1 000 mg/l Ca2+ = 2,75 mg/l This is the theoretical maximum value of the ions in solution before precipitation will commence. In practice however, these low values are not easily achieved. The solubility behaviour of most slightly soluble salts is much more complicated than is suggested by this example. Complex formation, hydrolysis, pH and other phenomena have important effects on solubility behaviour. 63 Example 5: Calculate the residual magnesium concentration that exists in a saturated magnesium hydroxide solution if enough sodium hydroxide has been added to the solution to increase the equilibrium pH to 11.0. Mg(OH)2(s) ↔ Mg2+ + 2OH- The solubility product constant for this reaction Ksp = 1,2 x 10-11. Determine the hydroxide ion concentration: Kw = [H+][OH-] = 10-14 at 25°C Because the pH is 11, [H+] = 10-11 mol/l we know that [OH-] = 10-14 / 10-11 = 10-3 mol/l Establish the solubility product constant expression, and solve for the magnesium ion concentration: Ksp = [Mg2+][OH-]2 [Mg2+] = 1.2 x 10-11 / (10-3)2 = 1,2 x 10-5 mol/l or 0,29 mg/l Since hardness ion concentrations are normally expressed as CaCO3, multiply the concentration by the ratio of the equivalent weights: 0,29 x 50 / 12,2 = 1.2 mg/l as CaCO3 An example of a redox reaction is the oxidation by chlorine of soluble manganese (ii) and the subsequent removal of manganese by precipitation of manganese (iv) dioxide. Some elements such as iron and manganese may have different oxidation numbers depending on the compound or the reaction in which it is involved. The oxidation number is calculated on the basis that the net charge of a molecule must be zero, and of an ion must be the charge that the ion carries. For example, the oxidation number of manganese in MnO2 is +4 because the oxidation number of oxygen is always –2 and since there are 2 O atoms for each Mn atom [2 x (-2) = (-4)] Cl2(g) + Mn2+(aq) + 2H2O(l) MnO2(s) + 2Cl-(aq) + 4H+(aq) Oxidation number: 0 +2 +1 -2 +4 -2 -1 +1 Electrons are negatively charged entities. This means that when an atom looses one electron, it looses one negative charge, or gains one positive charge. Manganese looses 2 electrons per Mn atom when changing from the +2 state in Mn2+ to the +4 state in MnO2. Since oxidation is defined as a loss of electrons Mn2+ is oxidized to the +4 state. Chlorine on the other hand gains one electron per atom when changing from molecular chlorine to the chloride ion and is therefore reduced in the process. Chlorine is the oxidising agent that oxidises manganese and in the process chlorine is reduced. Since electrons cannot be destroyed, the total gain in electrons must always 64 equal the total loss in electrons. The 2 electrons lost by Mn2+ equal the 2 electrons gained by the 2 chlorine atoms, each gaining one electron. Example 6: Chlorine gas (Cl2) is often used as oxidant in water treatment for oxidation and removal of iron and manganese. The equations for the oxidation and reduction reactions taking place when Mn2+ is oxidised to MnO2 and Fe2+ is oxidised to Fe3+ respectively by Cl2 are standard half reactions for the oxidation and reduction reactions. The overall balanced reactions can be written by adding the respective half reactions: ½ Mn2+ + H2O = ½ MnO2 + 2H+ + e- ½ Cl2 + e- = Cl- ½ Mn2+ + ½ Cl2 + H2O = ½ MnO2 + 2H+ + Cl- Fe2+ = Fe3+ + e- ½ Cl2 + e- = Cl- Fe2+ + ½ Cl2 = Fe3+ + Cl- Calculate the theoretical amount of Cl2 required to oxidise Fe2+ to Fe3+. Give the answer in mg/l Cl2 per mg/l Fe2+. 0,5 mole Cl2 per 1 mole of Fe2+ 0,5*2*(35,45) mg/l Cl2 per 55,85 mg/l Fe2+ = 35,45 mg/l Cl2 per 55,85 mg/l Fe2+ = 0,6347mg/l Cl2 per mg/l Fe2+ Calculate the amount of acidity produced in the reaction as mg/l H+ per mg/l Mn2+ oxidised. The H+ reacts with natural alkalinity in the water (HCO3-). Write the equation for this reaction and calculate the amount of alkalinity destroyed by the oxidation of Mn2+. Give the answers for acidity produced and alkalinity destroyed in mg/l CaCO3 per mg/l Mn2+ oxidised. 4 moles H+ produced per mole Mn oxidised 4 mg/ l H+ per 54,94 mg/l Mn oxidised = 0,0728 mg/l H+ per mg/l Mn = 0,0728 * 50/1 = 3,64 mg/l CaCO3 per mg/l Mn H+ + HCO3- = H2CO3 1 mol HCO3- per mol H+ = 61mg/l HCO3- per mg/l H+ * 0,0728 mg/l H+ per mg/l Mn = 4,44 mg/l HCO3- per mg/l Mn = 4,44 * 50/61 = 3,64 mg/l CaCO3 per mg/l Mn2+ 65 Experimental studies showed that a dosage of 2,1 mg/l as HOCl is required for effective oxidation of iron and manganese in a groundwater. The chlorine is purchased as a solution containing 15% NaOCl with a specific gravity of 1,12 g/ml. Determine the dosage rate (ml/min) for a dosage of 2,1 mg/l HOCl for a small treatment plant with a flow of 1200 m3/day if the NaOCl is diluted 1:5 (volume basis) before dosage. NaOCl = Na+ + OCl- + H+ = HOCl 1 mol HOCl = 1 mol NaOCl 52,45 mg/l HOCl = 74,45 mg/l NaOCl 2,1 mg/l HOCl = 2,1 mg HOCL/l * 74,45 mg NaOCl/ 52,45 mg HOCl/l = 2,98 mg/l NaOCl 2,98 mg NaOCl/l * 100mg soln/ 15mg NaOCl * 1200 m3/day * 1day/24hr *1hr/60min*1000l/m3 * 1000mg/g =16,556 g NaOCl/min =16,556 gNaOCl/min *1ml soln/1,12g soln =14,782 ml soln/min undiluted =14,782 + 5*14,782 = 88,69 ml/min diluted NaOCl MASS BALANCE CALCULATIONS Introduction In water treatment processes and plant the flow of water and of material is of primary importance. The reason is that it is important to account for water and chemicals usage and production of wastes. It is even more important in industry to account for the raw material and water used in order to determine how much material is converted to product and how much ends up as waste and how the measured concentration of waste streams compare with measured values. The only way that these types of questions can be answered is by means of compiling mass and water balances. The principles of conservation of mass form the basis of mass balance calculations. Material cannot be created or destroyed; it only changes form. The transport of material may be expressed as volumetric flow rate, such as mega litres per day or litres per hour, or in terms of mass flow, such as tons per hour, grams per second, or kg moles per hour. It is very important in mass balance calculations to ensure that units are consistent. This is normally achieved by writing down the units of every component as well of units to convert different sets of units to the same basis. This is illustrated in the examples below. CONSERVATION OF MASS The law of conservation of mass implies that material is neither created nor destroyed during normal activities of man and nature (with exception of nuclear reactions). All the mass that enters an enclosed region must either leave the region or accumulate in it. This is expressed as follows: (total mass in) - (total mass out) = (accumulation of mass in the region) 66 If there is no conversion of a particular species into other species by chemical or biological transformation, the conservation of that species can be expressed as: (total species i in) - (total species i out) = (total accumulation of species i) It is often convenient to deal with mass flow rates. The conservation laws then take on the form: (total rate of mass flow in) - (total rate of mass flow out) = (rate of accumulation of mass in the region) (total rate of flow of species i in) - (total rate of flow of species i out) = (rate of accumulation of species i) The special case where there is no accumulation or depletion is called steady state: (total rate of mass flow in) = (total rate of mass flow out) GUIDELINES TO DEVELOP A MATERIAL BALANCE The following can be used as general guidelines to solve a mass balance problem: Draw a diagram or a flow sheet for the problem and show all the known quantities and streams as well as flows and concentrations. Decide on the basis for solving the problem, e.g. a time interval such as one hour or one day, or a specific input such as 1000 m3 or 1000 kg. Identify the components. A tie component is an element, molecule, or kind of material that goes through the system unchanged. Dry air, water, or an inert solid can be used as a tie component. Select the boundaries of the region over which the material balances will be made. It is often necessary to select more than one region with different boundaries in order to calculate one component in one region that can be used as input to another region. The following examples show different mass balance applications: Example 7: A simple mass balance problem is one that only involves one component such as water. In this example we consider only one component involving the change in water level of a storage tank over time. A storage tank with a surface area of 25 m2 (5m x 5m) receives a steady inflow of 80 l/s, while the average demand (outflow) is 20m3 per day. By how much will the water level change after 5 days? 5m Flow sheet: 80 l/s 20 m3 /d 5m 67 Boundaries are from inlet to outlet Basis is 5 days Inflow – outflow = accumulation We have to convert the different units to the same basis to allow addition and subtraction of volumes. Inflow = 2 l/s, convert this to m3 in 5 days = 2 l/s * 1/1000 m3/l * 60/1 s/min * 60/1 min/h *24 h/d * 5d = 864 ( l * m3 * s * min * h * d) / (s * l * min * h *d) = 846 m3 Outflow = 200 m3 /d * 5d = 1000 m3 Accumulation = inflow (846 m3) – outflow (1000 m3) = - 136 m3 (negative accumulation is depletion, i.e. a drop in level) We know the loss in volume and the area of the tank (5m x 5m = 25 m2) and can therefore calculate the height, i.e. drop in level Drop in level = 136 m3 / 25 m2 = 5,44 Example 8: If we wish to calculate the concentration of dissolved components in the water, the problem becomes more complex. In this example we calculate the area of an evaporation dam into which the concentrate from a reverse osmosis (RO) plant is disposed of and we calculate the concentration of the brine after a certain period. An RO plant produces 150 m3 of brine per day that is disposed of in an evaporation dam. The average precipitation (rainfall) for the area is 800 mm per year and the average rate of evaporation is 1400 mm per year. What must the area of the evaporation dam be, in order that the level of brine does not rise by more than 500 mm per year. If the TDS concentration in the brine is 8500 mg/l, what will the concentration be after two years? Boundaries are the evaporation dam Basis is 1 year (could also be 1 day, but rainfall and evaporation figures are normally given as averages per year) Inflow = br

Handbook for Water Treatment Operation PDF

Document Details

Tags

Related

Summary

Full Transcript

Upgrade to continue