SHORT COURSE

WATER PURIFICATION AND TREATMENT

Everpure, Inc. 660 North Blackhawk Drive Westmont, IL 60559-9005

Phone: 630-654-4000

Fax: 630-654-1115 E-mail: info@everpure.com

Prepared by

William H. Beauman

Everpure, Inc. 1998

Updated January, 2001

EVERPURE, QC4 and MICRO-PURE are registered trademarks of Everpure, Inc.

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TABLE OF CONTENTS

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SECTION I: INTRODUCTION 5 A. General Information About Water5 1. The Water Cycle ...........................................................................5 2. The Basics of Water Chemistry.....................................................6 a) Dissolved, Particulate, or Colloidal .........................................6 i) Turbidity...........................................................................7 ii) Discussion of Sizes and dimensions...............................7 b) Organic or Inorganic ...............................................................8 c) Biological or Mineral ...............................................................9 d) Ionized or not..........................................................................9 i) Acids, Bases, and Salts ...................................................9 e) Aesthetic, Health-related or Nuisance ..................................10 f) Adsorption .............................................................................10 I) Carbon Adsorption of Chlorine ......................................11 ii) Activated Carbon ..........................................................12 iii) Freundlich Carbon Isotherms.......................................14 g) Chemical Equilibrium.............................................................15 i) The Water Equilibrium ...................................................16 ii) The Carbonic Acid Equilibrium......................................16 iii) About Buffers ...............................................................17 h) Oxidation and Reduction .......................................................18 i) Oxidation-Reduction Potential ........................................19 ii) Electromotive Series ......................................................21 B. Drinking Water Quality Standards ......................................................21 C. Centrally Treated Water Supplies22 D. Individual Water Supplies...................................................................22 SECTION II: DESCRIPTION OF WATER PROBLEMS................................23 A. Microbiological Problems..................................................................23 1. Viruses .........................................................................................23 2. Bacteria ........................................................................................23 i) Silver Filters ............................................................................24 ii) Bacteriological Testing..........................................................25 3. Algae ...........................................................................................26 4. Fungi (molds) ...............................................................................26 5. Protozoa and larger parasites ......................................................26 B. Iron Water, including Manganese27 C. Sulfur Water28 D. Acid Water28 E. Alkaline Water and Excessive Alkalinity29 F. Hard Water and Scale............................................................30 i) The Langelier Index.......................................................................31

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ii) Polyphosphates and Threshold Treatment32 G. Brackish Water and Excessive Total Dissolved Solids ....................34 i) Osmotic Pressure............34 H. Turbidity 35 I. Taste & Odor..................................................................................36 i) Chlorine T&O................................................................................36 ii) Earthy-musty-moldy-mildewy-fishy T&O......................................37 iii) Chlorophenols38 J. Color ...........................................................................................39 K. Toxic Organic Contamination40 SECTION III: TYPES OF WATER TREATMENT AND THE EQUIPMENT USED .................................................................. A. Mechanical Filtration41 1. Pressure Drop .............................................................................42 2. Hydraulic Capacity.......................................................................42 3. Everpure Precoat Filters..............................................................43 4. Carbon Blocks .............................................................................45 B. Adsorption46 C. Ion Exchange46 1. Softening (Conditioning)...............................................................47 2. Dealkalization ...............................................................................48 3. Demineralization...........................................................................49 D. Oxidation Filtration..............................................................................50 1. Manganese Greensand...............................................................50 2. Granular Brass51 E. Chlorination, Coagulation, and pH adjustment .................................52 F. Membrane Systems...........................................................................56 1. Ultrafiltration .................................................................................56 2. Nanofiltration ................................................................................56 3. Reverse Osmosis .........................................................................57 G. Ultra-Violet Irradiation........................................................................58 A note on typefaces: Underlining shows emphasis. Italics identifies foreign terms, including the Latin names of organisms. Bold-face sets apart important terms or industry jargon to remember.

SECTION I

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INTRODUCTION Congratulations, and welcome! Your initiative in enlisting in this Short Course on potable water treatment should be rewarded with sufficient expertise in the field to understand most water problems and their solutions and to design effective treatment systems. There will always be much more to learn than can be given here, but this course should provide the minimum background for a water treatment professional. A. General Information About Water Water is arguably the single most important substance in our lives. It has been called the “universal solvent,” and it is the fundamental requirement of all life. Virtually all of the water in existence today has been on Earth for billions of years; any “new” water comes only from comets. All the rest is continually being recycled. More than 80% of the Earth’s surface is covered by water, but less than 1% of it is drinkable, non-saline water. Water comprises about 2/3 of our bodies. Most of us could go without food for weeks, but we would die of thirst after only a couple of days. People in the developed nations each use some 123 gallons (466 L) of water every day, and we each drink an average of about a liter a day. The U.S. EPA bases regulatory matters on the assumed consumption of two liters per day per person. The Water Cycle 1. The water cycle begins with water vapor—moisture in the air from evaporation of water from the surfaces of oceans, lakes, etc.; transpiration from ground water through plants into the air from their leaves; and water vapor belched up from volcanoes. The vapor condenses under the influence of reduced temperature to form clouds, and then condenses further onto the surface of particles to form precipitation. Rain washes much of the air pollution out of the air, making rain a significant cause of water pollution. Acidic gases from combustion sources cause acid precipitation (“acid rain”), but even without any man-made acids in the air, normal levels of carbon dioxide (about 0.04%) dissolve in the rain to produce carbonic acid, H2CO3. 2. This slightly acidic rainwater lands on the ground and percolates through the soil, where it begins to dissolve grains of limestone (CaCO3) and dolomite (Ca/MgCO3). The resulting solution then contains calcium and magnesium ions(Ca+2 and Mg+2), called hardness; plus carbonate (CO3

−−−−2) and bicarbonate (HCO3

−−−−) ions, together called alkalinity. These are the dominant ions in most waters, and they form the basis of water chemistry. The water also begins to dissolve many other minerals and organic matter in the soil, and excess acidity can cause leaching of valuable nutrients, which then pollute streams. As the rainwater soaks into the soil and moves downward through many layers of Earth, it is filtered by particles of dirt, which also adsorb many dissolved contaminants, becoming relatively pure ground water. If it penetrates as deep as 10 meters,

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this “natural purification” may be sufficient to render the water potable without any further treatment, but wells from shallower depths cannot be considered free of “influence from the surface” (penetration of pathogens from surface contami-nation, especially protozoan cysts). The water collects above impervious layers of clay or stone, often extending over vast regions, and moves very slowly through pores, cracks, and other channels in the ground as an aquifer. Most are very ancient, and it may have been 50,000-250,000 years or more since the water there fell as rain. Thus, well water from deep aquifers is considered to be a non-renewable resource, like petroleum. Pumping out water faster than it can be renewed can lead to a significant lowering of the surface terrain, called subsidence (pronounced “sub-SI-dense”). Areas with varied geologic histories may have several aquifers stacked on top of one another, each receiving water from its own regional source and having its own unique chemical composition. The upper level of an aquifer is called the water table, and in most places the water table of the uppermost aquifer coincides with the water level of the lakes and rivers in the region. That water flows directly to the surface, but water from the deeper aquifers may be pumped out by wells, spewed out of volcanoes, or drawn into the roots of plants by osmotic pressure and transpired out of the leaves to become water vapor again, completing the cycle. The Basics of Water Chemistry Some understanding of basic chemistry is needed to understand water chemistry, but the only background needed to begin this course is the ability to recognize the symbols of the elements—C for carbon, H for hydrogen, Ca for calcium, etc. The rest is explained below. The various impurities that can be found in water can be categorized in several different ways: Are they dissolved, particulate, or colloidal? Are they organic or inorganic? Are they biological or mineral? Are they ionized or not? Are they aesthetic contaminants, health-related contaminants, or just process nuisances? Each of these will be discussed, along with the concepts of oxidation/reduction, adsorption, and chemical equilibrium, which will be explained and demonstrated with the pH scale and the alkalinity continuum. a) Dissolved, Particluate or Colloidal: Water molecules are very tiny, and they are also “polar” in nature, which means the molecules are not symmetrical: one end is bigger than the other, so they’re pointed, and the big end also has a “partial” positive electrical charge, while the small end is partially negative. They look a bit like a Mickey Mouse head in silhouette, with the ears being the two hydrogen atoms. Their size, shape, and electrical polarity enable water molecules to be invasive and squeeze into the smallest spaces between the molecules of solids and pry them apart. This is dissolution, and the molecules, atoms, or small clusters that are liberated into the water (called a solution now) are said to be “dissolved” if they are too small ever to settle under the influence of gravity. Anything solid that is suspended in water but is large enough to settle eventually is a particle. Molecules and clusters of intermediate size are called

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colloids. Examples of colloidal materials include biological polymers such as protein molecules, tannins and lignins leached from vegetation, tiny aggregates of rust and clay, fragments of microbes and larger organisms, even complete viruses, asbestos fibers, etc., up to a size of a few tenths of a micron in diameter. This is also the same size range as the wavelengths of visible light (violet light is about 0.40 um or 400 nm; red light, about 0.77 um or 770 nm), and that is why particles of this size cause haziness or cloudiness in water, called turbidity. Turbidity is measured by shining light through a test tube of the water and measuring the intensity of the light that comes out the other side. However, unlike ordinary photometers used in chemistry labs (which have the light source, sample, and detector all in a line), the detector for turbidity measurements is positioned at a 90° angle from the incoming light beam, so that only the reflected haze is measured. This is called nephelometry, and that is the origin of the Nephelometric Turbidity Unit, NTU. One NTU is not detectable by eye, but a reading of 15 NTU is noticeably cloudy, and murky river water may have a turbidity of several hundred. Public health or environmental regulations every-where require the turbidity to be low at the time of disinfection, because patho-gens can be shielded by particles. In the U.S., the requirement is <1 NTU 95% of the time, (<0.5 NTU for those systems using “conventional” or “direct” filtra-tion), with an absolute maximum of 5 NTU. When this is exceeded, Boil Water Orders may be required by law. That is mostly because of protozoan cysts in surface waters, which are difficult or impossible to kill. The large urban water-works using the most advanced water treatment methods are able to produce water with less than 0.5 NTU, but it invariably picks up millions of particles from the pipes and mains of the distribution system and may reach the point of use with several units of turbidity. Discussion of Sizes and Dimensions: How big is 0.40 um or 400 nm? The fundamental length or distance is the meter (symbol: m), but the unit most used in describing particles in water is the micro-meter (one-millionth of a meter) (symbol: um or µm), commonly called a micron. A meter is about a yard long; a millimeter (one-thousandth of a meter) (symbol: mm) is about 1/32 inch, and a micron is one-thousandth of that: 0.00003937 in. Also, many people think in terms of thousandths of an inch, or “mils” for very tiny measurements: a mil or 0.001 in. is 25.4 microns. A particle 40 microns in diameter is just barely visible to people with 20/20 vision—smaller than 40 um is microscopic. The table below will help illustrate the range. Comparison of Particle Sizes 1-inch ball 25.4 mm (millimeters) = 25,400 um or µm or “microns” (micro-meters) pollen 10 - 100 microns smallest item visible to naked eye 40 microns 1 mil or 0.001 inch 25.4 microns

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fog droplet 2 - 50 microns “dirt” 40+ microns silt and clay ½ - 20 microns pathogenic protozoan cysts 3 - 20 microns Cryptosporidium oocysts 3 - 7 microns Cyclospora cysts 8 - 10 microns Giardia cysts 8 - 12 microns Entamoeba cysts 12 - 20 microns red blood cells 7 ½ microns most bacteria and algae ½ - 5 microns “turbidity” 0.1 - 5 microns colloids 0.1 - 5 microns wavelengths of visible light 0.40 (blue) – 0.77 (red) microns = 400 (blue) – 770 (red) nm (nanometers) cigarette smoke 10 – 1000 nm viruses 10 – 250 nm protein molecules 2 – 50 nm individual atoms 0.05 – 0.25 nm = 0.5-2.5 Angstroms = 50 – 250 pm (picometers) b) Organic or Inorganic: the term, “organic” originally meant “made by an organism,” but now it refers to chemical compounds containing the element carbon, in some form other than the simple carbonate minerals such as lime-stone. Inorganic generally refers to minerals dissolved in the water—salts, metals, “hardness,” corrosion products, etc. Over the years since pollution became an issue, organic came to mean man-made in factories—synthetic, and therefore unnatural and dangerous. Organic foodstuffs are those grown without the aid of synthetic pesticides and fertilizers. These trends in the language are dangerous, because they ignore significant inorganic hazards that are also the result of industrialized life, such as toxic corrosion products and asbestos. In water chemistry, the organic contaminants are either named individually or grouped with acronyms such as: TOC = Total Organic Carbon NPTOC = Non-Purgeable TOC (purgeable CO2 [carbonate] removed first) TOX = Total Organic Halide (containing chlorine, bromine, and/or iodine) THM = Tri-Halogenated Methane VOC = Volatile Organic Chemical AOC = Assimilable Organic Carbon (digestible by microbes) NOM = Naturally-occurring Organic Matter DBP = Disinfection By-Products c) Biological or Mineral: This is nearly the same as organic or inorganic, but a distinction needs to be made for films on surfaces—biofilms versus scale deposits. Usually the two are mutually exclusive, because the high tempera-tures that often lead to scale formation usually kills micro-organisms. However, commercial ice makers of the recirculating cascade or spray types are suscep-tible to both, simultaneously. Scale develops because the minerals in the water become progressively more concentrated as the ice grows, but psychrophilic

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(cold-loving) bacteria and molds also form complex slimes that require regular cleanup. d) Ionized or not: When dissolved in water, some chemical substances split up as unequal fragments with electrical charges, positive and negative. Each charged atom or group is called an ion. Opposite charges attract and like charges repel, just like the North and South poles of a magnet. Positive ions are called cations because they migrate toward the cathode in an electric field, and negative ions are called anions because they go toward the anode. (These are pronounced CAT-ions and AN-ions.) As a generalization, inorganic chemicals are most usually “ionic” compounds that ionize when dissolved, and organic chemicals are usually non-ionic, but there are many exceptions. Three chemical types are always ionic, whether organic or inorganic, and those are acids, bases, and salts: Acids: chemicals that liberate a hydrogen ion, H+, when dissolved in water. The H+ is one charged fragment of the original molecule, and the remainder becomes negatively-charged as a result. Example: citric acid: C6H8O7 ⇔ C6H7O7

−−−− + H+ citric acid citrate hydrogen molecule ion ion Bases: chemicals that liberate a hydroxide ion, OH−−−−, when dissolved in water. The OH−−−− is one charged fragment of the original molecule, and the remainder becomes positively-charged as a result. Example: sodium hydroxide: NaOH → Na+ + OH−−−− sodium sodium hydroxide hydroxide ion ion Salts: chemicals produced by mixing an acid solution and a base solution and then crystallizing or evaporating to dryness. Two things combine: the H+ and OH−−−− combine to make water, and the other two oppositely- charged fragments combine to make a salt. When salts are dissolved again, they always dissociate into + and −−−− ions. Example: sodium citrate: Na+ + C6H7O7

−−−− ↔↔↔↔ NaC6H7O7 sodium citrate sodium citrate ion ion molecule e) Aesthetic Contaminants, Health-Related Contaminants, Process Nuisances: These are not as obvious as they might seem. Regulatory agencies the world

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over recognize that the offensive taste of iron in water is harmless and that toxic levels of lead or a pesticide may have no taste or odor at all, and regulations are always separated into mandatory health-related requirements and other aesthetic requirements which are only recommended. However, there are cross-overs, and others that cause problems only for equipment rather than for people. The recommended pH range of 6.5-8.5 becomes mandatory if there is a problem with corrosion of lead, copper, and cadmium plumbing materials. Excess sulfate or TDS (Total Dissolved Solids) causes short-term diarrhea in visitors unaccus-tomed to a water supply, but this is considered only an aesthetic problem. Like-wise, excessive silver causes a deathly graying of the skin and the whites of the eyes, and excessive fluoride ion causes ugly stains and even malformations of bones and teeth, but these are not considered to be health problems. Too much zinc can cause vomiting, but since there is no permanent damage, zinc is also only an aesthetic contaminant. As for process nuisances, there is nothing toxic about silica or hardness and alkalinity, and there are no regulations for them, but in excess they can lead to very damaging and costly scale buildup. Similarly, turbidity is very important at the time of disinfection (when it must be very low), but its importance to our industry lies in its ability to scratch valves, add to scale, and plug filters. f) Adsorption: Not to be confused with absorption (which is what a sponge does), adsorption is the attraction of tiny particles or dissolved molecules to a solid surface and holding them there by weak intermolecular forces. It is similar in concept to magnetism and the attraction due to static electricity, but much weaker. In theory, every atom in the universe has some degree of affinity for every other atom in the universe, just like gravity. But, just as gravity requires enormous masses like planets and stars to show its effects, adsorption requires extremely tiny distances to show its effects. In adsorption, the particle in question is randomly bounced around the solution by collisions with water molecules and other molecules in the water. (This is called Brownian motion. It is estimated that an atom or molecule in water is involved in a million-billion-trillion or 1027 collisions with other atoms or molecules every second. This is part of the definition of temperature.) Eventually, by chance, it will be bounced so close to the surface of a wall or another larger particle that there are very few water molecules separating it from the surface. When that happens, those few molecules produce only a few collisions from that side, and the particle is overwhelmed by collisions from the other sides and tends to become “plastered” to the surface by a continual barrage of collisions from the solution. This is the “physical” half of adsorption. The “chemical” half occurs if there is any chemical affinity between the particle and the material of the surface. If there is, the particle will become attached (adsorbed) and stay there; if not, it will bounce off right away or just diffuse away, later. The adsorptive forces (called van der Waals or London forces) are so weak that adsorbed substances can become desorbed rather easily—by adding certain acids, by heating the system, or by merely removing the contaminant from the

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influent water. For example, activated carbon filters or ion exchange beds nearing exhaustion are subject to desorption if the water quality suddenly changes for the better. That shows that these treatment techniques are equilibrium (balance) phenomena in which sorption and desorption both occur and achieve an average condition, like a well-matched tug-of-war. Since adsorption requires a surface, commercial adsorbent materials have very large surface areas and are exemplified by activated carbon, activated alumina, and fine powders such as baking soda. But many substances are so very insoluble or otherwise so readily adsorbable that even small surface areas can make a big difference. For example, most heavy metal ions (lead, mercury, copper, cadmium, silver, chromium) adsorb so strongly to the walls of both glass and plastic sample bottles that more than half of the total contamination can be missed in an analysis if the sample bottles are not treated with nitric acid first, to cause desorption. Similarly, many chlorinated hydrocarbons like the poly-chlorinated biphenyls (PCBs) adsorb so readily to both metal and plastic plumbing and filter materials that even coarse prefilters remove them very well. The adsorption and reduction of disinfectant chlorine by activated carbon is a special case. Activated carbon is a mild reducing agent and chlorine is a strong oxidizing agent, so after chlorine becomes adsorbed, it then actually reacts with the carbon. The chlorine is reduced to chloride ion (as in table salt and sea water), one atom of carbon is oxidized to carbon dioxide, and both are released to the solution (desorbed). Meanwhile, most of the spots on the activated carbon where all this took place become “auto-regenerated” back to their original, like-new condition, ready to adsorb again. For free available chlorine (FAC), this takes only about fifteen minutes, which means that a small amount of carbon can achieve an acceptable steady-state condition if the flow rate is slow or intermittent. For “combined chlorine” (monochloramine), the reaction is much slower, and more carbon or more contact time is needed to achieve equivalent reductions. The chemical reactions between activated carbon’s “active sites” (C*) and these forms of chlorine are shown below. Note that any surface oxides on the carbon are recycled when reacted with monochloramine, while they are oxidized to CO2 and lost when reacted with free chlorine. Free Chlorine Cl2 + H2O ⇔ HOCl + H+ + Cl − (forming “aqueous chlorine”) C* + 2Cl2 + 2H2O ⇒ C*O2 + 4H+ + 4Cl− (the overall reaction) C* + HOCl ⇒ C*O + H+ + Cl−−−− C*O + HOCl ⇒ C*O2 + H+ + Cl−−−− Combined Chlorine: Monochloramine C* + NH2Cl + H2O ⇒ C*O + NH3 + H+ + Cl−

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C*O + 2NH2Cl ⇒ C* + N2 + H2O + 2H+ + 2Cl− Finally, most dissolved/suspended particles and molecules in drinking water that are highly adsorbable to something usually do become adsorbed to a larger particle before reaching the point of use. Thus, adsorbable contaminants can often be removed by mechanical fine-filtration because the contaminant in question is already adsorbed to a larger particle. If you remove the particle, you remove the adsorbed contaminants along with it. This commonly applies to heavy metal ions, many pesticides, other chlorinated hydrocarbons, viruses, and asbestos fibers. About Activated Carbon: Granular activated carbon (GAC) and powdered activated carbon (PAC) are the predominant adsorbents used in our industry. They can be made from nearly anything organic: coal, petroleum, wood, coconut shells, peach pits, ion exchange resin beads, fabrics, even waste plastics. The starting material is first charred—heated without air or oxygen, so it doesn’t burn up. Everything that can be vaporized or melted bubbles out as tar or pitch, leaving many holes and channels. Then the charred material is heated further, to above 1000°C (hot enough to melt aluminum and lead), with the introduction of live steam or other activating chemicals. The superheated water vapor is extremely corrosive, etching more holes and extending channels to an amazing degree. Metallic impurities are preferentially attacked and washed out, resulting in a significant purification of the original material. However, the heat of activation does more than extend holes and channels and increase the surface area of carbon; it also changes the fundamental crystal form from amorphous “carbon black” to the perfect crystalline array of graphite plates. The carbon atoms in graphite are arranged in sheets or plates of interlocking six-atom rings that look like slices through a honeycomb. Such a perfect arrange-ment causes the London forces to focus and concentrate at the surface, making activated carbon the best (strongest and most general) adsorbent known. After activation, the carbon may be treated further to produce specific chemical qualities on the surface. For example, an acidic environment produces carbon with maximum capacity for heavy metals but minimal capacity for chlorinated organics, while an alkaline environment does the opposite. Most grades used in our industry are made for organic adsorption. When activation is complete, the carbon is a delicate, airy material that is so full of holes, it can barely hold together. It is crushed to a powder, and then proprietary binders are added to form granules of the desired size. The final product has a total internal and external surface area of more than 1000 square meters per gram, or half a football field inside a piece the size of a pea. Activated carbon adsorption is useful because the material has strong chemical affinities for several important classes of contaminants that are common in water. These are:

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1. Disinfectant chlorine: “Free available chlorine” (FAC) is readily

adsorbed, then chemically reduced, and finally desorbed as chloride ion along with one molecule of carbon dioxide, with auto-regeneration of most of the carbon’s active sites and nearly infinite capacity. “Combined chlorine” (monochloramine) is less easily adsorbed, requiring more carbon or reduced flow rate for equivalent performance.

2. Organic compounds containing chlorine and other halogens: Simple halogenated hydrocarbons are highly adsorbable to activated carbon. This includes a great many pesticides (DDT, Endrin, Lindane, Chlordane, etc.), industrial solvents (trichloroethylene, trichloroethane, tetrachloroethylene, carbon tetrachloride, etc.), and disinfection byproducts (THMs including chloroform, chloral hydrate, etc.).

3. Organic compounds containing benzene rings: These include some of the most toxic chemicals, such as benzene, toluene, dioxins, polychlorinated biphenyls (PCBs), and phthalate esters (plasticizers for vinyls).

4. Heavy metals: Lead, cadmium, and mercury adsorb readily, both as dissolved ions and colloidal oxide or carbonate particles, but the capacity is limited—similar to the capacity for THMs.

5. Taste and Odor (T&O) compounds: The substances produced by microbes that are responsible for the common musty-earthy-mildewy T&O are extremely well adsorbed and with very great capacity.

There is great variation in the adsorbability of dissolved/suspended substances, and also great variability in the adsorptive capacity of different adsorbents. A bed of granular activated carbon (GAC) may be exhausted with respect to chloroform and other volatile organic compounds (VOCs) after only a few hundred bed volumes, yet continue to adsorb PCBs for many thousands more. Different grades and types of activated carbon have different capacities for the same contaminant as well as various contaminants, which means that one must be very careful and specific in making comparisons. Chemists have developed a standard procedure for comparison of adsorption capacities, called an isotherm. A Freundlich carbon isotherm is determined by preparing several identical bottles of powdered activated carbon suspended in water. (In German, “eu” is pronounced “oi,” so Freundlich sounds like “Froindlich.”) Varying amounts of a contaminant are added to the bottles, and all are mixed until adsorption has reached equilibrium under those conditions of temperature and pressure. Then the carbon is filtered out and the solutions are analyzed to find how much contaminant remains unadsorbed in each one. The Freundlich equation is used to calculate the amount of contaminant that was adsorbed per milligram of carbon, and each data point is graphed on logarithmic graph paper with the carbon capacity in mg/g on the Y-axis and the final equilibrium concentration in mg/L on the X-axis. Finally, the “average” line representing all of the data points is drawn. That average line on the graph paper is called the isotherm for that contaminant and that carbon under those conditions. But that line covers a range of capacities; the one value used for

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comparison purposes has been designated by international agreement to be the capacity in mg/g of carbon on the Y-axis that corresponds to the value of 1.0 mg/L on the X-axis. If the isotherm does not cross the 1.0 mg/L point, the line is artificially extended (“extrapolated”) to that level for the purpose. The Freundlich Equation can be represented as:

where qe = concentration of contaminant on the carbon at equilibrium (Y-axis value) Ce = concentration of contaminant in solution at equilibrium (X-axis value) K = a constant 1/n = another constant The Ce is determined by analysis; the qe is calculated using the equation and the two Freundlich constants that were determined by the chemist who published the data. By convention in our industry, chloroform, the main THM, has been selected as the least adsorbable contaminant that activated carbon can be claimed to adsorb effectively—any contaminant that has a Freundlich isotherm qe value less than that for chloroform cannot be said to be removed by carbon adsorption. We “draw the line” at chloroform; anything less adsorbable than chloroform is deemed not worth the trouble. Examples: in one test series, the Freundlich capacity for chloroform is 2.6 mg/g GAC (the isotherm passes through 1.0 mg/L above the X-axis at the point where the Y-axis reads 2.6 mg/g GAC), while the equivalent value for trichloroethylene is 30 mg/g GAC. That is a much higher value, meaning that that particular GAC will adsorb trichloroethylene much more easily than chloroform. However, in the same data set, the capacity value for methylene chloride is only 1.3 mg/g, and thus we say that methylene chloride cannot be adsorbed efficiently by that GAC. The carbon’s capacity for it is too small to make an economically viable product. See the example below.

nee CKq

1

)(=

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It is important to remember that carbon capacity figures derived from Freundlich isotherms are to be used only for comparisons—e.g., carbon A is better than carbon B, or contaminant X is easier to remove than contaminant Y. They should not be used as concrete capacities to calculate how long a filter should last. The reason is that the isotherm data are produced at equilibrium, which may take several days of stirring in the lab to achieve. But a bed of GAC or a filter cartridge operates on a dynamic, flowing basis, and equilibrium conditions may not be achieved even after a weekend downtime. g) Chemical Equilibrium is a special property of the reactions and solutions of many substances. Whenever you see a chemical equation with arrows pointing both forward and backward, it is an equilibrium reaction. In an ordinary reaction, the starting materials simply react until one of them is all gone, and then it stops dead. But an equilibrium reaction does not automatically go to completion. Instead, it grinds to a stall (achieves equilibrium: the reverse reaction equals the forward reaction) at some characteristic point, even though there may still be plenty of starting materials left. Every reaction has its own regular balance-point: one reaction may stall after only 0.0000001% of the starting materials have been used up; another may go to the point of being 99.999999% complete. The end-point characteristic to each equilibrium reaction is defined by a number called the Equilibrium Constant, KEq , which never changes. Whenever the concentration of any of the reactants or products is changed (by reacting or adding or removing something), all of the other substances in the equation change in compensation, and they change in such a way as to preserve the value of the KEq. For example, consider the most important equilibrium of all, which is also the simplest: The Water Equilibrium describes the dissociation of water molecules into their component parts, hydrogen ions and hydroxide ions (H+ and OH−−−−): H2O ⇔ H+ + OH−−−− KEq = 10−−−−14

S a m p le F r e u n d lic h C a r b o n is o th e r m s

0 .0 0 1

0 .0 1

0 .1

1

1 0

1 0 0

0 .0 0 0 1 0 .0 0 1 0 .0 1 0 .1 1 1 0

m g /L a t E q u ilib riu m (C e )

mg/

g C

arbo

n C

apac

ity (q

e)

T ric h lo ro e thy le ne C hlo ro f o rm M e thy le ne C h lo r id e

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10−−−−14 is a very small number, which means the reaction doesn’t “go very far,” and only about one molecule in ten million is actually dissociated at any instant. In pure water with nothing added, H+ and OH−−−− will be identical in concentration, equal in this simple case to the square root of the KEq, or 10−7. This reaction and relationship are so important that a special shorthand notation has been developed, called the pH system: use the exponent only, drop the sign and use “p” as a symbol, and you get pH = 7.0 for pure water. (The pOH = 7.0, also.) All this is important because it’s water we’re talking about, but also because H+ is very important to water chemistry. The hydrogen ion is the smallest and most chemically active piece of ordinary matter known. Hydrogen is the smallest of the elements—a single proton with a single electron whizzing around it. Remove the electron, and you have a single, lone, naked proton in solution that is highly reactive, invasive, and corrosive. (It’s not really naked and alone; sometimes it is shown attached to a water molecule as H3O+.) H+ defines acidity; H+ is acid. So, the water equilibrium defines acidity: pH less than 7 is acidic; pH greater than 7 is alkaline or “basic.” If either H+ or OH−−−− is increased artificially, by adding or removing acid or alkali, the other changes in the same proportion so as to main-tain the value of the KEq, which is “chiseled in stone.” Thus, if acid is added to pure water to make the pH = 3, the equilibrium shifts so that the pOH = 11 and the KEq is preserved as 10−−−−14. The Carbonic Acid-Calcium Carbonate Equilibrium is the other “most important” equilibrium in water chemistry. When water falls as rain it absorbs carbon dioxide gas from the air, forming carbonic acid, which instantly dissociates into bicarbonate and carbonate ions: CO2 + H2O ⇔ H2CO3 ⇔ H+ + HCO3

−−−− ⇔ 2H+ + CO3−−−−2

carbon water carbonic hydrogen bicarbonate carbonate dioxide acid ion ion ion When this slightly acidic water soaks into the ground and contacts limestone (calcium carbonate) or dolomite (mixed calcium and magnesium carbonate), it dissolves some of the rock, making “hardness” and “alkalinity:” 2H+ + Ca/MgCO3 ⇔ Ca+2 + Mg+2 + H2CO3 (which dissociates acid lime/dolomite “hardness” “alkalinity” as shown above) “Hardness” is the sum of all ions that react with soap to inhibit lathering and precipitate soap scum or “bathtub ring.” It happens that they are all metals with more than one “+” charge—mostly Ca+2 and Mg+2 in most water supplies, but also including Fe+2, Cu+2, Zn+2, Mn+2, etc., if present. “Alkalinity” is a confusing term because it came into use before the chemistry was understood. As a word, alkalinity just means the opposite of acidity, or something that consumes or neutralizes acidity. When it was realized that both

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of the carbonic acids and their ions in the two equations just above are the same thing, “alkalinity” became the term in water chemistry to represent the sum of CO2 + H2CO3 + HCO3

−−−− + CO3−−−−2 , and thus the total buffering power of the water

to resist changes in acidity. All the double arrows show that the reactions go both forward and backward, which means that limestone dissolved in one place may deposit as lime scale in another place later, if the equilibrium shifts. The equilibrium can be shifted by adding or consuming any of the constituents involved. Remember that the relationship between reactants and products—the KEq—is chiseled in stone, and any change from the outside will be compensated for, instantly, in a way that preserves the value of the KEq.

The values of the three KEq for the three carbonic acid equilibria (one for each arrow) are intentionally absent from this discussion, to avoid getting you bogged down in numbers. What matters is that you get a “feel” for the concept of equilibrium. It makes sense that adding more reactants (on the left side of an equation) would make a reaction “go” further (to the right). But with equilibrium, you can also make a reaction go to the right by removing some of the final products. It’s as if the removal creates sort of a chemical vacuum which demands to be filled. The reverse is also true: to make a reaction reverse (go to the left), you can either load up the right side (add more final product) or lighten the left side (remove reactants). It is as though the reactants and products are physically connected by the arrow, and you can push or pull on the reaction from either end. And in chemical systems with several reactions connected in series, like the first one, reprinted here, the “connection” goes through all of them. Note that one of the final products of the third reaction (far right) is H+, or acid: CO2 + H2O ⇔ H2CO3 ⇔ H+ + HCO3

−−−− ⇔ 2H+ + CO3−−−−2

If you add more acid, the entire three-reaction continuum will reverse (go to the left), causing some CO2 to bubble away and be lost forever, thus reducing the total alkalinity of the system. This is called dealkalization.

About Buffering: Just above, it was said that the presence of alkalinity “buffers” water against changes in pH. Buffers are chemical stabilizers that use equi-librium to establish and maintain a chemical balance. The common “mineral” acids such as hydrochloric, sulfuric, and nitric acid do not work as buffers because their dissociation to produce H+ ions is not governed by equilibrium: they all become virtually 100% dissociated the instant they are dissolved. Because of this, they are called “strong” acids. Only “weak” acids and their salts like carbonic acid and sodium carbonate can be buffers. The KEq of carbonic acid is about 10−7, which means that normally, only one out of millions of carbonic acid molecules is split into H+ and HCO3

− ions at any instant. Note that the value of the KEq is close to the value of the pH (and also the pOH) when acidity is “neutral:” pH = 7 = pOH. That means carbonic acid is good for buffering water near pH 7. However, note that the bicarbonate ion, HCO3

− , still has one hydrogen to lose, and bicarbonate ion therefore has its own KEq, which

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is about 10−11. That means sodium bicarbonate would be a good buffer for water at pH values near pH 11. For comparison, the KEq of acetic acid (vinegar) is about 10−5, which means that acetic acid would be good for buffering water in the vicinity of pH 5.

So, how do buffers work, exactly? They work by having half of their capability inactive, in reserve, so to speak, from both the acidic and alkaline point of view. When the pH is equal to the exponent on the KEq, 50% of the weak acid molecules in solution (carbonic acid in this case) are whole and 50% are dissociated into H+ and HCO3

− . If extra H+ is artificially added to the solution, it will instantly be incorporated into the carbonic acid-bicarbonate-carbonate “alkalinity continuum,” with very little effect on the pH. Likewise, if some OH− ion is artificially added, it will instantly react with and be neutralized by a H+ ion, but that H+ will instantly be replenished from the alkalinity, again with very little impact on the pH. As long as any total alkalinity remains in the water, it will be able to “absorb” either H+ or OH− in the vicinity of pH 7 without changing the pH very much. Incidentally, it is possible to buffer a solution against changes in ions other than H+ --chloride ion, Cl−−−−, for example.

h) Oxidation and Reduction are the result of an atom or molecule losing or gaining an electron. Before that makes any sense, you need to understand that atoms are made of a “nucleus” containing positively-charged protons and a “cloud” of an equal number of negatively-charged electrons orbiting around it, as shown in the familiar “solar system” icon for atoms and atomic energy. Some of the outermost electrons may be lost rather easily (or not, depending on the element), thus destroying the balance between protons and electrons and changing the character of the atom greatly. When an electron is lost, the protons outnumber the electrons by one + charge, and the atom as a whole then becomes oxidized with a +1 charge, as in the sodium ion, Na+. Conversely, if the electron cloud is capable of accepting another electron, then the electrons would outnumber the protons by one − charge and then the atom would become reduced and have a –1 charge, as in the chloride ion, Cl−−−−. In their elemental forms of sodium metal and chlorine gas, the atoms are balanced and electrically neutral and have a net charge (oxidation state) of zero.

Consider the enormous difference between a piece of sodium metal and a cloud of greenish chlorine gas on the one hand and a spoonful of salt crystals or sea water on the other. Yet they are almost identical—both are a collection of sodium and chlorine atoms, and they both have the same number of electrons over all, but in the salt crystals, the sodium atoms all have one electron too few, and the chlorine atoms all have one electron too many. They balance out in the NaCl crystal, but that is not the same. Na+ ion is simply different from Na0, and Cl− ion is nothing like Cl0. So, electrons are extremely important to chemistry, and when an atom becomes oxidized or reduced, its chemistry changes completely. The iron atoms in the centers of the hemoglobin molecules in our

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red blood cells must be in the +2 oxidation state—ferrous iron, Fe+2—to work properly in carrying oxygen to our tissues. If they get oxidized to ferric iron, Fe+3, by nitrite, cyanide, or carbon monoxide, for example, the hemoglobin doesn’t work properly, and we die. One tiny electron on each iron atom makes all the difference.

Oxidation and reduction always occur together, because they are like the two sides of a coin. When an atom or molecule loses an electron (becomes oxidized), that electron must go somewhere, and the atom or molecule that accepts it becomes reduced. When a “reducing agent” does its job, it becomes oxidized in the process. When an “oxidizing agent” does its job, it becomes reduced in the process. Oxidation is loss of electrons, and reduction is acceptance of those same electrons by a “partner” in the reaction. Every chemical has its own unique tendency to gain or lose electrons, called its oxidation-reduction, or “redox” potential, or ORP, which is expressed in volts and is available from reference books. These voltages can be used to calculate whether a particular reaction will “go” without introducing energy from the outside, but that will not be discussed here. Some examples:

• the simple battery used for teaching electricity, made by dipping a zinc bar and a copper bar into a solution of copper sulfate: Zinc loses electrons more easily than copper does, so the zinc becomes oxidized and the copper becomes reduced. If you simply dip a zinc bar into a solution of copper sulfate, you get a copper-plated zinc bar which is badly corroded, because every atom of copper that “plates out” on the zinc bar is matched by a zinc ion liberated into the solution. If you use both bars and connect them with a wire to form an electrical circuit, the plating process will proceed until the zinc bar disintegrates and breaks the circuit. The reaction below describes both situations:

Cu+2 + Zn0 ⇔ Cu0 + Zn+2

• the oxidation of hydrogen sulfide (“rotten egg” smell) by chlorine: if only the minimum amount of chlorine is used, the result is elemental sulfur, which is a solid that must be filtered afterward. However, it is possible to oxidize the sulfur atom all the way to sulfate if more chlorine is used. The first reaction below shows the sulfur being oxidized from the -2 state in H2S to the zero state in elemental S, and the second one shows the sulfur being oxidized to the +6 state as sulfuric acid.

H2S + Cl2 ⇔ 2 H+ + 2 Cl−−−− + S

H2S + 4 Cl2 + 4 H2O ⇔ 8 H+ + 8 Cl−−−− + H2SO4

Most elements have at least two possible oxidation states. Note chlorine, below, which has more possibilities than most elements:

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Oxidation Chemical State Name Comments ClO4

−−−− 7 perchlorate ion used in rocket fuel ClO3

−−−− 5 chlorate ion used in explosives ClO2 4 chlorine dioxide gas used for oxidation and disinfection ClO2

−−−− 3 chlorite ion used as a bleaching agent ClO−−−− 1 hypochlorite ion in laundry bleach Cl2 0 chlorine gas elemental chlorine Cl−−−− -1 chloride ion as in table salt and sea water This also demonstrates some useful rules of nomenclature: “-ate” = suffix denoting a “high” oxidation state; same as “-ic” suffix for acids* “-ite” = suffix denoting a “low” oxidation state; same as “-ous” suffix for acids* “per-” (short for “hyper-”) = prefix denoting “higher oxidation state than ‘-ate’” “hypo-” = prefix denoting “lower oxidation state than ‘-ite’” “-ide” = suffix denoting the complete absence of oxygen

*names of some acids: nitric acid, nitrous acid, sulfuric acid, sulfurous acid, perchloric acid, chloric acid, chlorous acid, hypochlorous acid

The Electromotive Series is a listing of metallic elements in Potassium descending order of their oxidation-reduction potential or ORP. Sodium One of the characteristics of metals is the free movement of the Magnesium outer “valence” electrons, which makes them good electrical Aluminum conductors. The ORP voltage is a measure of each element’s Zinc readiness to lose electrons (be oxidized). They are arranged Iron at right, with the most reactive element at the top and the least Tin reactive at the bottom. Elemental potassium and sodium are Lead so reactive that they react with water as if it were strong acid, Hydrogen making NaOH or KOH and hydrogen gas. At the other end, Copper everyone knows that silver, platinum and gold are “noble” Mercury metals that are difficult or impossible to corrode. This list can Silver also be used in the reverse: the ions of a metal can cause Platinum the corrosion of any metal above it in the list. For example, Gold if a copper wire is immersed in a solution of silver ions, the Ag+ → Ag (plating) and the Cu → Cu+ (corrosion). Thus, silver metal is relatively non-corrodable, but silver ion is highly corrosive. Alchemists’ old name for silver nitrate (AgNO3) was “lunar caustic”—the Ag+ ion is so caustic and corrosive that it will even react with the proteins in your flesh. Perhaps that is why Ag+ and Cu++ have some weak antibacterial activity, even at low concentrations.

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B. Drinking Water Quality Standards Every sovereign nation adopts its own set of standards, or “norms” for drinking water quality. The most influential and widely-recognized standards are those developed by the United Nations-World Health Organization, the European Union, Japan, and the United States EPA. They are all changed or updated regularly, so any listing here would be out of date quickly. There are always two types of regulations:

Primary requirements relate to health effects, and compliance is mandatory.

Secondary requirements relate to aesthetic effects and are only recommended.

In addition, the U.S. EPA also establishes maximum contaminant level (MCL) goals for some contaminants which for some reason cannot presently be attained. For example, all carcinogens are automatically given an MCLG of zero, even though everyone understands that zero is impossible to achieve.

In the United States, all water systems that regularly supply drinking water to at least 25 people or 15 “service connections” are Public Water Supplies which must comply with the regulations. If there is only one building or subscriber, such as a rural factory or school, it is a Non-Community Public Water Supply. If there is only one site and the 25 people are not necessarily the same 25 people every day, as at a rural restaurant or service station, it is a Non-Community, Transient Public Water Supply.

C. Centrally Treated Water Supplies

As a company, we support the practice of large-scale, central treatment of drinking water by professionals. The municipal waterworks generally do an excellent job, and they make our job easier. But the smaller the customer base, the fewer resources are available, and the smallest systems have the most difficulty meeting the requirements. Also, the smaller the system, the more likely it has a ground water source, and those generally require less treatment than water from lakes and rivers. Surface water often contains “color” (tannins) and high turbidity which must be coagulated with special chemicals before filtering. These tannins are also the major source of unwanted byproducts of chlorine disinfection such as THMs and many other chlorinated organic compounds. Many surface waters are mostly snow-melt or rain water with relatively little total mineral content, called Total Dissolved Solids (TDS). This allows the water to be corrosive, or aggressive to plumbing materials so that asbestos fibers and lead, zinc, cadmium, copper, and iron corrosion byproducts leach into the water from asbestos-cement pipe, brass fittings, lead-based solder, and copper and galvanized pipes. The microbiological quality of tap water is generally very good as the water leaves the waterworks—often less than 10 HPC/mL (Heterotrophic Plate Count), but after a few hours or miles of transit through the distribution

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system, the chlorine dissipates and allows the bacteria that survived disinfection to multiply. In addition, biofilms containing many species build up and flake off into the water. It is important to remember that less than one-half of 1% of the potable tap water produced by large waterworks is actually consumed by people. The rest is used for flushing toilets, washing everything from laundry to streets, watering lawns, fighting fires, etc.

D. Individual Water Supplies

The smallest of all water supply types, whether public or private, is the individual system for a single user. This is usually a private well, because the minimum treatment for a surface supply is often more trouble than is worthwhile: chemical disinfection and fine-filtration to remove parasites are unavoidable, and the addition of flocculants and the holding time needed to promote coagulation may also be needed. We usually recommend that well waters be chlorinated (for disinfection) and that will also change any iron or sulfide to filterable particles which are easy to remove. Well water often requires softening or conditioning (removal or reduction of hardness), and that can also remove small to moderate amounts of iron and manganese by ion exchange.

SECTION II

DESCRIPTION OF WATER PROBLEMS

A. Microbiological Problems (Primary MCL: absence of Fecal Coliforms, viruses, cysts)

Drinking water may contain all types of microbes—viruses, algae, molds, bacteria, and parasites—but they are not all necessarily dangerous to health. Some just cause bad taste and odor or clog filters. Algae and molds seldom have any health significance, but all products and claims relating to the other three microbial types are watched carefully by public health officials. At present, Cyst Reduction is the only permitted microbiological reduction claim, unless bacteria, viruses, and cysts are all killed/removed simultaneously. The require-ment for that is 3-log reduction (99.9%) for Cryptosporidium oocysts, 4-log reduction (99.99%) of two kinds of virus, and 6-log reduction (99.9999%) of Klebsiella terrigena bacteria, which is an environmental coliform. Any product that can document all three abilities can be labeled a “microbiological purifier.”

1. Viruses are the smallest microbes—only a few hundredths of a micron—and they are not truly alive, because they are not fully functional. They are parasites of individual cells and cannot do anything by themselves. Most are very specific and limited in the type of cell they can infect, and the only ones in water that are of interest to humans are those that come from human sewage. Their numbers in contaminated raw waters are usually about one per cent of the fecal coliform

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count (see below, in “bacteria”). They are too small and too few to cause any mechanical or maintenance problems for filters; the only “problem” they cause is disease. They are only a minor problem in developed nations with good sewage and water treatment practices, but rotaviruses in particular are the #1 cause of human death and misery world-wide: infant diarrhea caused by rotaviruses in the drinking water accounts for fully half of all human deaths on Earth each year.

2. Bacteria are the next-smallest microbes, single cells ranging from about 0.2 microns to 10 microns in size. They are the smallest, simplest, and most ancient fully-functional life form. They have only one chromosome, and the genetic material is not even organized in a nucleus. Bacteria are the most numerous and most varied life form on Earth, accounting for more total biomass than any other. Every gram of good, productive topsoil contains some 10 billion bacteria, comprised of thousands of colonies with millions of organisms in each colony. They occupy every known habitat. Most of them are harmless—indeed, most have no interaction with people at all; a few cause inconvenience in the form of bad taste and odor or a tendency to cause clogging of fine-filters; and a very few can cause disease. The most common odor-causing types merit special mention: they are filamentous bacteria resembling mold under the microscope, except that they are much thinner than molds, and are called actinomycetes. Many of them produce antibiotics of the “-mycin” type, in addition to some of the same “earthy-musty-mildewy-moldy-fishy” odor compounds produced by molds and algae. Bacteria are so varied, it is difficult to know how to categorize them, but the following types are important to know about:

-autotrophs: those that make their own food by photosynthesis or oxidation/ reduction of minerals

-heterotrophs: those that must consume organic matter from the environment to live

-HPC organisms: those detected by the Heterotrophic Plate Count procedure

-pathogens: those capable of causing disease

-opportunistic pathogens: those capable of causing disease only if given an unusual “opportunity”—open wounds, burns, defective immunity, etc.

-coliforms: those that “resemble” (biochemically) Escherichia coli, the predominant commensal intestinal organism of mammals—indicators of sewage contamination

-Total Coliforms: an important public health screening test to identify all coliform organisms, both fecal and environmental, belonging to the genera Escherichia, Klebsiella, Enterobacter, Citrobacter, and Serratia

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-Fecal Coliforms: an important public health test to identify E. coli, specifically

-Anaerobic bacteria: those that prefer or require the absence of dissolved oxygen; they often produce “septic” smells, and some produce “rotten egg” smell

-“iron bacteria:” partial autotrophs that oxidize dissolved iron to rust

-“sulfur bacteria:” anaerobic autotrophs that chemically reduce sulfate ion or metabolize organics containing sulfur to produce the stink of rotten eggs, or hydrogen sulfide, H2S, and also organic sulfides, all with offensive T&O.

“Silver Filters” or filters containing silver as an antibacterial agent are very common and deserve special mention. The silver ion, Ag+, has a long history in water treatment, going back to Roman times, and there is no doubt that it can be an effective agent against many bacterial pathogens if the concentration and contact time are sufficient. However, it is illegal in the U.S. to make that claim, because it is not effective against all pathogens. Many bacteria are relatively immune to silver, and it has no effect on viruses or cysts at all, so antimicrobial claims can easily be dangerously misinterpreted. Thus, such products are allowed to claim only that the silver is bacteriostatic, meaning that rapid, uncontrolled bacterial growth inside the filter is inhibited. Further, it is illegal to suggest that silver has any effect on harmful or pathogenic bacteria; a manufac-turer may claim only that silver inhibits the growth of those harmless bacteria that might produce bad taste and odor or premature clogging of filters. Unfortun-ately, silver ion is not very effective against many of those bacteria, so we are in the ridiculous situation of being unable to make an accurate and truthful claim because it could be misinterpreted, and unwilling to make the only permitted claim because it is not very accurate. Everpure, Inc. now promotes the use of its filters with silver in applications where they would be most useful (where they are required by law in other countries, or where bacterial pathogens might be present and users need all the help they can get), while using labeling language that complies with U.S. law.

Bacteriological Testing: For some reason, the bacteria that inhabit drinking water systems are difficult to detect by culturing them. More than 99% of those shown to be present by staining them with dyes and counting them with the aid of a microscope just don’t grow into countable colonies when spread onto a nutrient medium in a petri dish. They are called “non-fermenters” or “viable but nonculturable organisms.” Part of the reason for the error in plate counts may be that many are trapped in a clump of biofilm containing hundreds of individuals, which grows out as a single colony. Or, it may be that their growth requirements (nutrients, pH, temperature, atmosphere, salinity, etc.) are simply too exotic and different from what is provided by the “standard” test conditions. However, the main reason is believed to be the phenomenon of bacterial dormancy: each specific strain of bacteria seems to require a rest of some weeks or months after

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growing actively for a while. At any moment, the bacterial population of a muni-cipal distribution system will be comprised mostly of three to seven specific types, each following its own schedule of growth, decline, dormancy, and regrowth, and almost all of them are dormant most of the time. That means that, regardless of the specific plate count procedure used, the true bacterial numbers are at least 100 times, perhaps as much as 1000 times higher. Thus, plate counts are not very accurate or informative. It is folly to try to compare individual plate count results; only controlled bacteriological testing programs, in which the same trained person takes the samples in the same way, at the same place, at the same time of day, etc., yield scientifically valid data that can be compared. Even then, the variation is so great that the averages must differ by more than a “log” (logarithm) or power of ten to be considered statistically different. For example, an average influent plate count of, say, 350/mL and an average filter effluent plate count of 3450/mL are not significantly different numbers, statistically speaking.

Tests for coliform organisms are done on 100 mL volumes, either by filtering 100 mL through a membrane and then culturing the membrane, or by adding special chemicals to a bottle containing 100 mL of sample. The special nutrient medium and temperature are inhibitory to most non-coliform organisms, but many unwanted HPC organisms may grow anyway, making the true coliforms difficult or impossible to detect. When that happens, the lab analyst may reject the sample with a “TNTC” notation (Too Numerous To Count) and request a new one. In the U.S., any repeat coliform tests are to be analyzed using a “Presence-Absence” test procedure, in which chemicals called “MMO” and “MUG” are added to 100 mL of the sample. If there are any total coliform organisms in the sample, their metabolism will turn the MMO yellow by the next day. If there are any fecal coliform organisms (E. coli) in the sample, they will turn the MMO yellow and also metabolize the MUG to something that fluoresces chartreuse when viewed with UV illumination.

3. Algae used to be considered plants, but the world is no longer divided into “animal, vegetable, or mineral.” Now, phytoplankton (algae) and zooplankton (protozoa) are both placed in a new fifth kingdom called Protoctista , giving us the present classification into bacteria, fungi, protoctists, plants, and animals. Algae are often one-celled and microscopic in size. They produce three types of problem, two of them serious. Several kinds of microscopic one-celled and filamentous algae produce the “musty-earthy-mildewy-moldy-fishy” taste and odor which ranks No. 2 in consumers’ complaints of tap water (after chlorine T&O). Some of the same algal species are also capable of causing serious clogging of filters. During seasonal algae blooms in reservoirs, their numbers may be so enormous that the municipal filtration plant output is less than half of actual production—they use more water backwashing the filters than they put into the distribution system. When they take short-cuts at the waterworks, point-of-use filters also clog quickly. The less serious problem is toxicity. Two or three common species produce both nerve and liver toxins that can be deadly. This is not a serious problem because the water does not become dangerous until after

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it has become so smelly and unsightly as to be disgusting to a person. Thus, pets and farm animals may be affected by algal toxins in ponds and puddles, but hardly ever people. Algal toxins are not presently regulated, but they are being watched carefully.

4. Molds are fungi. Molds in water are mostly microscopic and single-celled, although they can grow into sheets of slime and larger structures like mush-rooms if given the chance. They often contaminate water-using equipment via airborne spores. They are a major source of the “musty-earthy-mildewy-moldy-fishy” T&O compounds that are also produced by algae and filamentous bacteria. With one exception, molds are harmless, and the only “problems” they cause are moldy T&O and ugly slimes. The exception refers to people with damaged immune systems, who can die from some fungal infections.

5. Protozoa and Larger Parasites often escape disinfection processes, and physical removal by mechanical filtration is often the most cost-effective remedy. Boiling for one minute will kill everything except bacterial spores, but that is the most costly approach. No chemical disinfectant has been found reliable, and standard ultra-violet systems are ineffective. It happens that the parasite most difficult to kill or remove is also one of the most prevalent—Cryptosporidium oocysts may be as small as 3 microns in diameter, but most are 4-7 um in size. They can be found in virtually all surface raw water supplies worldwide, and no municipal or regional waterworks anywhere can guarantee killing or removing them all. Most strive for 99.9% reduction using multiple process barriers, but many do not succeed. Thus, every person using a surface water supply needs a fine-filter with 99.9% efficiency (or better) for reduction of 1-micron particles at the point of use. So far, Cryptosporidium is the smallest, and 1-micron filtration is good enough. All other known waterborne parasites—Giardia , Entamoeba, and Cyclospora cysts, various round worms, tapeworms, flukes, and their eggs, Schistosoma larvae, etc.—are much larger. B. “Iron Water” (including Manganese). Secondary MCL: 0.3 mg/L Fe + Mn Iron and manganese may be in the water as it is pumped up from wells, or iron pipes and other equipment may get oxidized (corroded) to form rust. Iron is almost always in the ferrous form, the Fe+2 ion, when fresh from the well or corrosion, but it is easily oxidized further, to the ferric form, Fe+3. Ferrous iron is the one producing metallic taste. Ferric iron in water attracts hydroxide ions so strongly that it will steal them from water if the pH is above about 5.0. The reaction produces iron floc, which is a gooey, rust-colored mass. Fe+3 + 3 H2O ⇔ Fe(OH)3 (solid) + 3 H+ KEq = 10+33 The exceptionally large equilibrium constant indicates that the reaction will go essentially to completion, and hardly even a single atom of Fe+3 in solution will

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escape precipitation. This is the usual cause of consumer complaints of yellow stains on porcelain and laundry. The usual remedy is either oxidation to Fe+3 followed by filtration to remove the iron floc, or removal by ion exchange water softening if the concentration is not too high. Oxidation is usually done by feeding liquid chlorine bleach with a chemical feed pump, followed by a contact tank and filtration system to filter out the floc. Chlorination is preferred because it also disinfects, but it can also be done with oxidizing media such as manganese greensand and granular brass. Actual rust is a combination of the oxides of both ferrous and ferric iron: FeO/Fe2O3, sometimes written Fe3O4. Iron in water from wells with high organic content is sometimes combined with huge tannin and lignin molecules from rotting vegetation, called “color bodies,” to form an even more deeply-colored product called heme iron, which is difficult to remove. (Heme refers to the blood pigment, which has an atom of ferrous iron at the center of an organic complex. But the organic part of heme and hemoglobin is nothing like tannins and lignins.) Some tannin molecules are large enough to be called colloidal, and it is often possible to remove most heme iron by fine-filtration in the sub-micron range. But if the molecular weight is too small, it must be removed either by ultrafiltration or nano-filtration, or by chemical coagulation followed by standard filtration. Manganese usually occurs in combination with iron in well waters, and it makes the stains on laundry and porcelain even darker. When Mn+2 is oxidized by dissolved oxygen or chlorine, the result is manganese dioxide, MnO2, which is dark brown in color. It is removed by the same methods used to remove iron. C. “Sulfur Water.” Secondary MCL for odors: 3.0 TON “Sulfur Water” refers to the stink of hydrogen sulfide, H2S, which is almost always derived from bacterial activity. There are several species from two different families: one family specializes in reducing sulfate ion, SO4

−2, to sulfide, S−−−−2, which then immediately acquires two hydrogen ions to become H2S. The other family consumes organic matter containing sulfur and metabolizes the sulfur to hydrogen sulfide. H2S makes the infamous “rotten egg” smell, and it is also an acid which can cause rapid corrosion of all types of plumbing materials. Therefore, it is always important to treat sulfur water, even if people get accustomed to the smell. Sulfur and iron and the bacteria that oxidize or reduce them often occur together, and when they do, the laundry and plumbing stains are black, due to black ferrous sulfide, FeS. Such bacterial mixtures often produce peculiar “septic” odors that are similar to that of hydrogen sulfide. Analysis of water samples for hydrogen sulfide is possible, but since it is a volatile gas, a special preservative must be used at the time of sampling. However, the nose is a superior detector, and since the odor threshold is so low that even the tiniest concentrations must be removed, an accurate analysis is usually not needed. The only good remedy is chemical oxidation followed by

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filtration to remove the elemental sulfur. Activated carbon can remove low levels for a short time, but it is inefficient and has low capacity. New catalytic carbons are an improvement, but a long contact time (slow flow rate) is still needed. Chlorination is the preferred approach because it also disinfects, but oxidizing media such as manganese greensand and granular brass may also give satisfaction. D. Acid Water. Secondary MCL: pH between 6.5 and 8.5 A water supply may be too acidic because of acid rain, contamination from acid mine drainage, or it may simply be “too pure” (like distilled) and have insufficient alkalinity to buffer the water against atmospheric acids. The secondary MCL becomes mandatory in the U.S. only if it is necessary to control the corrosion of lead- and copper-containing plumbing materials. (Other materials also corrode, but they are less toxic.) The remedy is to add alkalinity to the water, and there are two approaches. If it is a private well that is being or is to be chlorinated, it is a simple matter to add a solution of sodium carbonate (“soda ash”, Na2CO3), or sodium hydroxide (“caustic soda”, NaOH), to the chlorine bleach. If there is no chemical pump, use granular bed filters containing calcite (calcium carbonate: ground limestone or marble) or magnesia (magnesium oxide) media, or a mixture of the two. These dissolve slowly in the acid water, consuming excess acid and providing the water with additional hardness and alkalinity. E. Alkaline Water and Excessive Alkalinity. Secondary MCL: pH between 6.5 and 8.5 Alkaline water and excessive alkalinity are not necessarily the same thing. True alkaline waters are rare and occur mostly in desert areas or regions with geologic deposits of “trona,” bauxite, borax and other alkaline ores. Such waters are often undrinkable due to high TDS or salinity, regardless of the pH. Excessive alkalinity may be paired with high hardness, or it may be due to less extreme levels of the same minerals that produce alkaline waters. The notion of excessive alkalinity is relative, also. Generally, more than 250 mg/L alkalinity can be expected to cause problems with lime scale, but levels as low as 80 mg/L are preferred by some producers of “post-mix” beverages requiring carbonated water for soft drinks. That is because any alkalinity in the water will consume some of the acidity from acidulants such as citric acid in the syrup concentrate that give the necessary tartness to the drinks. In addition, alkalinity in the water fights against the carbonation process itself. Thus, commercial dealkalization technology is very important. All remedies to excessive alkalinity involve adding acid to the water in some way. It is not advisable to add acid to the chlorine solutions being fed for disinfection and/or oxidation, as is often done with alkaline additives used to treat acid

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conditions, because that would produce toxic chlorine gas in the storage tank before it could be fed. Acids must be dosed separately. One way is with a simple “pot-feeder” cartridge similar to those used for many polyphosphate feeders, except that the material being dissolved and fed is a food-grade acidulant such as citric acid, tartaric acid, malic acid, etc. A portion of the incoming water is directed to the feeder, which dissolves some of the acid to form a concentrated solution which is then metered back into the main stream through an orifice that is sized to produce the proper proportional feed. The ion exchange approach is often preferable, especially if the water is also too hard or has excessive TDS. Dealkalization by ion exchange involves the use of a cationic ion exchange resin in “hydrogen form” (loaded up with H+ ion or “acid”): any cations such as the hardness ions (Ca+2 and Mg+2), plus sodium, Na+, potassium, K+, etc. are exchanged for H+ ion, which then immediately neutralizes one H-equivalent of alkalinity. This may be by direct combination with OH−−−− to make water, or the H+ may combine with one of the “alkalinity ions” (carbonate or bicarbonate), driving that equilibrium one step to the left and causing one molecule of CO2 to bubble away and be lost from the system. Either way, this treatment approach represents a “double-whammy:” not only is the excess alkalinity reduced; an equivalent amount of hardness or sodium ion is also removed, and the overall TDS is reduced as well. A degree of added control can be achieved by using a weakly-acidic cation resin (“WAC” resin) instead of the strong-acid cation resin used for water softening. (See the next section for a full discussion of “strong” and “weak” resins.) The result in this case is that such a resin will remove only the amount of hardness that is “balanced” by alkalinity. For example, if a water has 250 ppm of hardness but only 200 ppm of alkalinity, WAC resin will remove only 200 ppm of the hardness, and the H+ that is liberated from the resin will neutralize all of the alkalinity. If the alkalinity level is greater than the hardness level, WAC resin will remove all of the hardness but only an amount of alkalinity equal to the original hardness level. (This kind of arithmetic is permitted only when the hardness and alkalinity values are expressed “as CaCO3.”) F. Hard Water (not regulated) “Hardness” in water is the sum of all ions that react with soap to produce soap scum or “bathtub ring,” and also inhibit lathering. The problem ions are all metals with more than one “+” charge—mostly calcium (Ca+2) and magnesium (Mg+2) in most water supplies, but also including zinc (Zn+2), copper (Cu+2), manganese (Mn+2), and others if present. In addition to interfering with cleaning, hardness also combines with alkalinity to form “lime scale,” which is calcium carbonate, CaCO3. Scale is a problem because it forms hard deposits in water lines and water-using equipment which can scratch valves, insulate the equipment and interfere with heating or chilling operations, and cause clogging. The standard treatment for hard water is by “softening,” or “water conditioning” by “sodium-cycle ion exchange.” Hard water is directed through a bed of cationic

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ion exchange resin in “sodium form” (with Na+ ions attached to the resin), which exchanges the Na+ ions for hardness ions in the water. However, softening is not always desirable, because the Na+ ions put into the water may be unwanted. Even if the possible negative health effects of Na+ on blood pressure are avoided by using K+ as the exchanged ion instead of Na+, both the taste and the TDS effect of either one can make the treated water taste salty and be unsuitable for brewing coffee or making soft drinks. The taste threshold for Na+ is in the 100-150 ppm range for most people, and softening water with a hardness of 217 ppm as CaCO3 will result in the addition of 100 ppm Na+ to the water. Thus, about 15 grains per gallon or 257 ppm as CaCO3 should be considered the maximum hardness that can be treated by softening without producing a salty taste. Sometimes reverse osmosis is used to remove excess sodium and other dissolved minerals with health effects. The Langelier Index (LI) is a system for estimating or predicting the amount or degree of problems with lime scale a particular water supply will cause. It is based on calculating the pH at which the water would reach the saturation point for calcium carbonate (pHs) using the data from a chemical analysis of the water. Specifically, the LI is equal to the actual pH minus the calculated pHs. This produces a number, usually between –3 and +3. If the LI is positive, the water will deposit calcium carbonate; if the LI is negative, the water will dissolve calcium carbonate. To calculate the pHs, you need to know the total dissolved solids (TDS), the concentrations of calcium ion and total alkalinity, and the actual pH of the water. You also need to decide what temperature you’re interested in, and it is useful to have an electronic calculator that will give you the “logs” (logarigthms) of the concentrations of calcium and alkalinity. A short log table is given below, along with tables for special constants derived from the temperature and the TDS. pHs = A + B – log(Ca+2) – log(alkalinity) where A = constant derived from temperature B = constant derived from TDS Note that a negative LI is sometimes misused to predict the corrosiveness of a water to metal plumbing materials. The LI does have some influence on corrosion, but it is only one of many factors, and it is too simplistic an answer to that question. Temp. Constant TDS Constant Ca+2 or alk. °C A mg/L B mg/L as CaCO3 log 0 2.60 0 9.70 10 1.00 4 2.50 100 9.77 20 1.30

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8 2.40 200 9.83 30 1.48 12 2.30 400 9.86 40 1.60 16 2.20 800 9.89 50 1.70 20 2.10 1000 9.90 60 1.78 25 2.00 70 1.84 30 1.90 80 1.90 40 1.70 100 2.00 50 1.55 200 2.30 60 1.40 300 2.48 70 1.25 400 2.60 80 1.15 500 2.70 600 2.78 700 2.84 800 2.90 900 2.95 1000 3.00 Example LI Calculation Given: pH = 8.40; TDS = 400 ppm ; Ca+2 = 200 ppm as CaCO3; Alk. = 300 ppm as CaCO3 How bad will scaling be in an ice maker? (0°C) LI = pH – [A + B – log Ca+2 – log alk.] Interpretation of LI Values +3 = very severe scaling tendency = 8.40 – [2.60 + 9.86 – 2.30 – 2.48] +2 = severe scaling tendency +1 = moderate scaling tendency = 8.40 – 7.68 0 = no scaling tendency -1 = slightly corrosive to lime scale = +0.72 -2 = moderately corrosive to scale -3 = severely corrosive to lime scale When softening is not an option, the scale-forming potential of the water can be limited by reducing the pH or the alkalinity, which are discussed above. But often the simplest and most cost-effective treatment to inhibit lime scale forma-tion and deposition is Everpure’s proprietary InsurIceTM System, which is the combination of fine-filtration and low-level polyphosphate feed. (Regardless of the name, the InsurIce approach also works very well for coffee brewing. The conditions inside steamers are too extreme, so they are treated by dealkaliza-tion.) Polyphosphate treatment is very important and needs its own discussion. Polyphosphates are polymers of the phosphate ion, PO4

−−−−3. The smallest one is pyro-phosphate, P2O7

−−−−4; the next is called tripolyphosphate, P3O10−−−−5, and so on,

up to large molecules of 20 or 30 phosphates linked together in chains, having very large electrical charges. They are useful for inhibiting lime scale in two ways. The first is direct, mechanical interference with the growth of calcium carbonate crystals. The shape of the PO4

−3 group is somewhat similar to the shape of CO3

−−−−2, and this similarity enables a phosphate to take the place of carbonate on the surface of a crystal of growing calcium carbonate. However, it

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is just enough unlike carbonate to prevent any further crystal growth on top of it. Chemists call this type of mechanism “competitive inhibition:” lime scale crystal growth is inhibited when polyphosphate ions compete with carbonate ions for the same Ca++ sites on the surface of growing scale; when the poly-phosphate wins, scaling is stopped at that point for several hours or days, depending on temperature. At boiling water temperature, polyphosphates break apart or “revert” to single PO4 units, or “ortho-phosphate” in less than an hour. If the hardness level is very high, it may cause precipitation of calcium ortho-phosphate, thus negating the original reason for using a polyphosphate. The other mechanism of lime scale inhibition by polyphosphate feed is called “electro-static dispersion.” There are millions of tiny particles of all kinds in water—not just dirt, but also tiny scale particles that use a speck of dirt as a “nucleus” to crystallize upon and grow—and the vast majority of them carry a negative electrical charge. That means they all naturally repel one another to some extent, because, just as in magnets, opposites attract and likes repel. But there is a lot of violent activity in water on a microscopic scale: molecules and particles in water are involved in some million-billion-trillion or an estimated 1027 collisions per second. Sometimes these particles collide with enough momen-tum to overcome the repulsion, and then they stick together to form a larger particle. That is one of the ways scale and other kinds of sediment grow: by agglomeration. Fortunately, the repelling force can be magnified by dosing the water with dissolved ions having many negative charges, and that is exactly what polyphosphates do. The smallest polyphosphate has a -4 charge; many have charges of -20 or -30 or even more. They cluster around particles and add their charges to make them repel each other thousands of times more strongly than before, with the result that very few are able to overcome the repelling force and become agglomerated into larger masses that deposit as scale or sludge. The concentration of polyphosphates needed to be effective is less than 10 ppm, and it does not depend on the hardness level. Neither the competitive inhibition mechanism nor the electrostatic dispersion mechanism operates “chemically,” with one molecule reacting with another molecule one-on-one. Instead, the presence of small amounts of polyphosphate creates a non-scaling environment by physical means, and the overall effect is called threshold treatment. Higher polyphosphate concentrations of 20-100 ppm usually exceed the “threshold” of calcium polyphosphate precipitation, producing massive amounts of phosphate scale and sludge. Still higher concentrations—up to 500 ppm or more—create a “soft” water by sequestering every hardness ion with a polyphosphate molecule as a “complex ion” which remains dissolved. Examples of this approach are high-phosphate laundry detergents and “bath salts,” but it is never used for drinking water because such high levels of polyphosphates can cause diarrhea in sensitive persons. If the water has been pre-filtered by an efficient fine-filter, there will be very few particles present at all, and the enhanced repelling force will be even more effective. Everpure research has shown that fine-filtration alone can reduce

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scale by half or more, just because scale includes a lot of dirt, iron floc, and other particulate matter in addition to calcium carbonate crystals. Pairing efficient fine-filtration with threshold polyphosphate treatment can reliably reduce scaling, and therefore maintenance costs for ice-makers and coffee brewers, by more than 80%. This patented technology is embodied in Everpure InsurIceTM products for ice makers and is effective for water supplies having up to about 15 grains per gallon (GPG) or 250 ppm hardness as CaCO3. Above 20 GPG hardness, dealkalization usually gives better results. G. Brackish Water and Excessive Total Dissolved Solids (Secondary MCL: 500 ppm; 1000 ppm in California) Water can acquire excessive dissolved mineral content by evaporation, as in the case of Colorado River water being directed through the desert in aqueducts to Southern California; by flowing through deposits of soluble salts, as in the true alkaline waters that are full of sodium sulfate or sodium carbonate; and by sea water intrusion, as is seen in all coastal areas. These waters are not merely unpalatable; they are actually unhealthful until the excess dissolved material is removed. In an extreme case, such as drinking sea water, it can be fatal. That is because of osmotic pressure, which deserves its own discussion. Osmotic Pressure is a kind of fluid pressure that develops when solutions of different strength are separated by a semi-permeable membrane—such as the membranes inside our bodies and surrounding all of our cells. [Interestingly, osmotic pressure is one of the few phenomena that depends only on the total number of particles (including atoms, ions and molecules) in solution, and not their size, or type, or chemistry. A tiny hydrogen ion exerts the same effect as a huge protein molecule or microscopic particle of scale. The only other phenomena with this characteristic are freezing point, boiling point, and vapor pressure.] A semi-permeable membrane may be of the osmosis type (which allows only water to pass through) or the dialysis type (which allows only dissolved ions to pass through), or it may actually be a living membrane able to choose specific materials. All are based on diffusion, which is a fundamental physical process in which atoms and molecules that are concentrated in one place tend to spread out and equalize the concentration—like a drop of perfume permeating a room. The “drive” to equalize concentrations is fundamental in nature, and this is the “force” behind osmotic pressure. When solutions of different strength are separated by an osmotic membrane, the system “attempts” to equalize the concentrations, but the only mechanism permitted is to allow water to move one way or the other (that’s the definition of an osmotic membrane). The only way movement of water can equalize the concentrations is by moving through the membrane in the direction of the more concentrated solution, as if in an attempt to dilute it. (If water went the other way the concentrated side would only get more concentrated.) Thus, to look at it another way, concentrated solutions have the ability to draw water across

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membranes toward and into them. A concentrated solution exerts a real force upon a dilute one if they’re connected by a semi-permeable membrane, and the magnitude of the force is a direct function of the TDS. Therefore, if we drink something that is more concentrated than our bodily fluids, like sea water, it draws moisture out of our tissues and into the intestines, causing diarrhea and making us more dehydrated than before. The only way to reverse the process is to apply pressure to the concentrated side—a pressure greater than the osmotic pressure the solution already has—and that is reverse osmosis. Water with a TDS of 2000 ppm exerts an osmotic pressure of about 20 psi, and that is there-fore the usual limit of water quality for RO systems that use only the 40-60 psi line pressure provided by the water works. Water with more than 2000 ppm TDS requires pumps and high-pressure housings for RO to be cost-effective. The ion exchange process for reducing TDS is called demineralization. It requires two types of ion exchange media: a cationic resin or zeolite in “hydrogen form” (regenerated with acid, so that the exchange sites are loaded with H+ ions); and an anionic resin in “hydroxide form” (regenerated with strong base, so that the exchange sites are loaded with OH−−−− ions). Put them together and they exchange everything with an electrical charge for either H+ or OH−−−−, which then combine to form water. However, molecules that are not ionized will not be removed. Examples: organic solvents classed as volatile organic chemicals (VOCs) such as trichloroethylene and benzene, most insecticides and herbicides, most of the organic taste & odor compounds, and a host of other organic chemicals both natural and synthetic. Demineralization leaves in a lot that both reverse osmosis and distillation can remove. H. Turbidity (Primary Standard: less than 1.0 NTU generally, or 0.5 NTU where “direct filtration” is used, 95% of the time during disinfection; maximum of 5.0 NTU at any time) Turbidity is cloudiness or haziness in water, caused by a dispersion or scattering of light by dissolved particles that are the same size as the wavelengths of light used to illuminate them. (The visible part of the light spectrum includes light with wavelengths from about 0.40 microns for violet light to about 0.77 microns for red light.) It is measured by analyzing the scattered light intensity at a right angle or 90° to the path of the test light beam. This technique is called nephelometry, and that is the source of the “N” in the NTU unit of turbidity. The particles that cause turbidity may be clear or opaque, light or dark colored, crystalline or amorphous, made of various chemical compositions. Some particles may refelct light; others may absorb or refract the light, and it is there-fore impossible to convert turbidity values to concentrations of sediment in mg/L. However, most turbidity is just “dirt” and “dust,” composed of alumina or silica or alumino-silicate—the most prevalent minerals in the Earth’s crust and the stuff of nearly all rocks except limestone. Thus, an analysis of turbid water nearly always shows aluminum, and it’s difficult to distinguish between this naturally-

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occurring aluminum and aluminum floc (from the coagulation step in the treatment train used by large water works) that has gotten past the filters. That is unfortunate, because blobs of floc may well contain dangerous levels of toxic or infectious sediment. (That is the purpose of flocculation/coagulation—to concentrate turbidity and dirt by agglomeration so they can be removed more efficiently by the large granular bed filters.) Turbidity is the only contaminant that must be determined daily (every four hours at a minimum; preferably, continuously in “real time”) by waterworks operators, because low turbidity is necessary for effective disinfection. Pathogens are easily shielded from the effects of disinfectants by particles and sediment in the water. But the turbidity measured at the treatment plant bears no resemblance to the turbidity that exists at the far reaches of the distribution system. There, many decades of sedimentation, biological growths, and corrosion products accumulate, become encrusted, and slough off again to produce a steady supply of new particles which clog filters, score valves, and cause foul tastes and odors. This is the turbidity that matters to our products. I: Taste & Odor (Secondary Standard, Odor: <3 TON [Threshold Odor No.]) Most taste is actually odor, and it was found that a Threshold Odor Number of 3 (a 3 : 1 dilution with odor-free water makes the bad T&O disappear) is the maximum most people would tolerate before abandoning a water supply (perhaps in favor of a less safe supply). The predominant bad T&O in potable water these days is that of chlorine in some form. Second is the very common “musty-earthy-moldy-mildewy-fishy” T&O produced by certain algae, bacteria, and molds. Other fairly common offenders are iron and sulfur compounds already discussed above, other strange T&O occasionally produced by micro-organisms, and the strong medicinal T&O of chlorophenols that sometimes plagues systems using marginal disinfection practices. Each of these will be discussed below. Chlorine T&O is familiar to everyone, but there are important differences. All odorous substances are, by definition, volatile molecules, and that means elemental chlorine gas, or Cl2 when we smell Free Available Chlorine or FAC. That includes the hypochlorous acid molecule, HOCl, which is part of the equilibrium. These have a rather “pure” smell when compared with the heavy, dull smell of monochloramine, which is a form of chlorine intentionally produced by many water works trying to limit the production of THMs. And since mono-chloramine is much weaker as a disinfectant than FAC, it is often used at much higher concentrations, making the T&O even worse. It is produced by adding ammonia (NH3) to water that already has a residual of FAC, and if the wrong proportions are used or the pH is off, dichloramine may also be produced. Trichloramine or nitrogen trichloride exists only at very low pH, so it can be ignored. These chloramines can be recognized as the smell of a poorly-managed swimming pool, which has a steady input of ammonia from the

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perspiration and urine of swimmers. If the chlorine level does not keep pace with the ammonia influx, the pool area will stink of chloramines, which is much worse than the smell of ordinary free chlorine. The same is true of drinking water, where chloramines may also be unintended. When raw water is first chlorinated, ammonia and several naturally-occurring amines react first; if there isn’t enough chlorine to oxidize them completely, disinfection suffers. As more chlorine is added, the oxidation of the nitrogen in the amines progresses from mono- to dichloro- to trichloramine, and finally to nitrogen gas, which bubbles away and is lost. The last reaction is known as the breakpoint reaction because that is the one that destroys the last of the chlorine demand and permits FAC to persist so that actual disinfection can begin. The chemical reactions for these chlorine species are all given below: Free Available Chlorine Cl2 + H2O ⇔ HOCl + H+ + Cl−−−− ⇔ OCl−−−− + 2 H+ + Cl−−−− chlorine hypochlorous hypochlorite acid ion Combined Chlorine Cl2 + NH3 ⇔ NH2Cl (sometimes intentionally produced, chlorine ammonia monochloramine sometimes not) + Cl2 ⇔ NHCl2 (never intentional; very strong smell) dichloramine + Cl2 ⇔ NCl3 (only at pH < 3) nitrogen trichloride + Cl2 ⇔ N2 + N2O + HCl nitrogen nitrous gas oxide (the “Breakpoint Reaction”) Earthy-musty-moldy-mildewy-fishy T&O is well known to everybody, even if they don’t associate it with water. It is surprising that molds are not the only source, but algae and certain bacteria also produce some of the same chemical compounds. The differences in smell may simply be due to differences in concentration and mixtures. For example, one of the compounds that smells musty in the concentrations found in water smells like—is the aroma of green bell peppers at lower concentrations. They’re exactly the same molecule. These compounds get into drinking water most often from algae blooms in the source water, which are usually seasonal. Filamentous bacteria called actinomycetes are also common causes, but their growths in reservoirs are more continuous and seldom cause unexpected trouble. Actual mold growth is less common and occurs mostly in long-abandoned plumbing or dead-end mains that have not received fresh, chlorinated water for a long time.

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Further to the discussion of “sulfur water” (which see, above): One of the bacterial families that is able to reduce sulfate ion (SO4

−−−−2) to sulfide ion and hydrogen sulfide (S−−−−2 and H2S) (“rotten egg” smell) can also produce spores, which are tough “survival forms” similar to protozoan cysts, only smaller. The bacteria are anaerobic, meaning they cannot tolerate dissolved oxygen, and they grow only in unused, dead-end mains or other nooks and crannies in the distribution system where all the oxygen has dissipated. They produce spores to survive exposure to oxygen if it ever returns. When water treatment equipment with good mechanical filtration ability is used on supplies carrying these spores, the spores can become lodged in the media. If the equipment is not used enough to keep the water inside them fresh and oxygenated, the spores can grow into active bacteria and begin producing rotten egg smell. It is usually best to replace or re-bed such media, but it may be possible to “vend” them often enough to kill any of the new bacteria before they can make new spores. Other anaerobic bacteria can make a different kind of foul T&O when oxygen is absent, but they do not produce spores and are easier to get rid of. The smell they produce is usually characterized as “septic,” referring to odors from sewage that are different from pure hydrogen sulfide. It is not necessarily sewage bacteria that are responsible for this problem in filters; many anaerobic bacteria can make the same odorous compounds. Chlorophenols are intensely offensive byproducts of disinfection that occur only when the concentration of disinfectant is marginal—that is, barely sufficient to kill pathogens, but without leaving much of a continuing residual to fight any sub-sequent contamination or chemical impurities. The T&O is strong, bitter, and iodine-like and is usually described as “medicinal.” There are several; the worst is called 2,4-dichlorophenol and has an odor threshold in the low parts-per-trillion range. The chemical “phenol” is a six-carbon ring (a benzene ring) with an -OH group attached somewhere. The huge molecules of tannins and lignins that give the brown-yellow color to swamp water (and also to tea) are highly “phenolic” in nature, meaning there are a great many benzene rings with one or several OH groups attached, all linked up into an irregular polymer chain with a molecular weight in the millions. These tannin and lignin molecules are the main source of THMs and other disinfection byproducts when the disinfectant (chlorine, chlorine dioxide, or ozone) chops them into little pieces and then chlorinates the fragments. The fragments that are, or contain, a phenol quickly become chlorophenols, which may easily be further oxidized (to destruction), but only if there is sufficient chlorine or ozone. A chlorophenol problem can also arise if water-using equipment contains phenolic materials such as PPO (poly-phenylene-oxide)-based plastic or rubber compounds that let phenol and derivatives leach into the treated water, and the residual chlorine level is very low. The foul T&O may occur only randomly or rarely and can even plague equipment already protected by filtration systems. All it takes is a few ppt (parts per trillion) each of chlorine and a phenol, both of which are undetectable at that

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level by any standard analytical method. J. Color (Secondary Standard: 15 “Color Units”) “Color” in water refers to the well-known yellowish-brownish tint caused by the presence of humus, a complex and variable collection of organic molecules derived from rotting wood and other plant materials containing cellulose, humic acids, tannins and lignins. It is common in surface waters wherever there is a lot of dead vegetation, as in southern swamps, northern bogs, and areas with logging/paper pulp operations. The final product is the result of centuries of microbial transformations, and every molecule is different, but they are all large polymers of benzene rings with hydroxyl, carboxyl, and methoxy groups such as this illustrative but hypothetical example proposed by Christman and Ghassemi [“Chemical Nature of Organic Color in Water,” JAWWA 58:6, 722-741,1966]: Color is measured by comparison with a special mixture of cobalt and platinum salts, called the Hazen Standard, which has that tint. 15 Color Units is very pale—just barely discernable without side-by-side comparisons. Some can be removed by activated carbon, but only with difficulty. Treatment with lime or other source of calcium ion will precipitate much of it, for removal by fine-filtration. Flocculation and coagulation with “alum” is even better. RO, UF and NF membranes remove color very well. Color and Taste & Odor : Many odorous molecules are produced as fragments of color molecules after they are attacked by chemical disinfecants. In addition to the phenols, these include chemicals with names like geraniol (which smells like geraniums), pinene (from pine trees), camphor, cinnamic alcohol, vanillin, etc., which can produce some very exotic mixtures of odors. Color and Sediment : The vanillin that comes from humus is the same as the vanillin from oak barrels that is so prized in aged wines and spirits. Those products also contain other high molecular weight, humus-like molecules that give an amber hue to brandy and whiskey. The largest of them will precipitate with iron, manganese, calcium, zinc, and other metal ions, which means that highballs made with very hard water may produce ugly sediment in the glass. Cheap liquors whose color comes from charred sugar (caramel) instead of aging in wood do not produce such sediments. K. Contamination by Toxic Organics (Various Specific Primary MCLs)

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Currently there are 57 organic chemicals out of the 83 contaminants regulated under the U.S. Safe Drinking Water Act, with MCLs ranging from 10 mg/L for xylene to only 0.00000003 mg/L for dioxin. It is beyond the scope of this Short Course to discuss each one, or even to try to classify them all in terms of toxicity or solubility, but it is important to know about the sources of many of them and whether they are adsorbable to activated carbon. Agricultural Pesticides: Half (27) of them are insecticides and herbicides, and most of those are chlorinated hydrocarbons which are quite insoluble, and therefore highly adsorbable to particles and to activated carbon. Those containing nitrogen (amines) and phosphorus (phosphates) are more soluble and less adsorbable, and those containing several oxygen atoms (alcohols [hydrox-ides] and organic acids with –COOH groups) are not well adsorbed to anything and are very difficult to remove, except by RO. Industrial Solvents and Byproducts: There are 22 of these. Most industrial solvents are small, volatile, low-molecular-weight molecules that readily penetrate RO membranes and adsorb only poorly to activated carbon. Nine more byproducts of plastics and rubber manufacturing are variable in carbon adsorption: some are adsorbed extremely well, and others extremely poorly. But most of these are larger than the solvents and can be removed by RO membranes. Disinfection Byproducts (DBPs): This refers mostly to byproducts of chlorine disinfection, acting on naturally-occurring humus and “color bodies,” but ozone can produce the same chemicals because it is a powerful oxidizer that can change chloride ion to chlorine. At present, only one chemical “family” is regulated—the THMs or trihalomethanes, as exemplified by chloroform, CHCl3. The MCL is 100 ppb total THMs (TTHMs), but that is expected to change significantly soon. In addition, two more families of DBPs, the halogenated acetic acids (HAAs) and acetic acid nitriles (HANs), will be regulated in the coming years. All are presumed to be carcinogens, and THMs within the presently permitted levels have also been implicated in human miscarriages.

SECTION III

TYPES OF WATER TREATMENT AND WATER TREATMENT EQUIPMENT A. Mechanical Filtration and Filtration Equipment Anything that blocks a fluid’s path and reduces the cross-sectional area through which solid objects can pass is a filter. Two logs across a stream will filter out dead cows. But there is no need to go that far—the word “filter” comes from “felt”, a mat of wool that is a very good filter. In fact, the golden fleece of ancient Greece was real, not myth: gold dust does adhere to wool, and mats of it used in those early ore sluices were one of the first industrial uses of filtration.

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Felt is a good example of one of the two basic types of filter, which is the depth filter: a mass of many random layers of media forcing the water into a tortuous path. Solids get entangled and are removed from the flow. The other filter type is surface filter, such as a fish net or a window screen: a single layer of mesh-work which stops everything larger than a certain size, and lets everything smaller than that size pass. Almost all filters in commerce are depth filters. Another way to categorize filters is by the size of the particles they can remove: Designation Particle Size Significance coarse filtration 30 um and up visible dirt fine-filtration 0.1 um and up colloids, cysts ultra-filtration 500,000 MW and up proteins, viruses nano-filtration 1000 MW and up “membrane softening” hyper-filtration atoms “reverse osmosis” Fine-filtration down to fractions of a micron is the same as “micro-filtration,” and it represents the best mechanical filtration that can be accomplished with ordinary materials and hydraulic designs. These include ceramic “candle” filters, pressed-carbon block filters, precoat filters, sintered glass/metal/plastic powder filters, and “microporous” polymeric membrane filters (usually pleated to maximize surface area). Removing particles smaller than 0.1 micron at reasonable pressures and useful flow rates requires abandoning the fundamental concept of mechanical filtration by physical interference with the flow, and adopting an entirely new and different paradigm: membrane filtration. Not to be confused with “microporous polymeric membrane filters”, the semi-permeable membranes used for ultrafiltration, nanofiltration, and reverse osmosis systems have no “pores” and do not physically remove particles; they exclude them chemically because the membrane material is formulated to be extremely hydrophilic (“water-loving”). Water dissolves into the membrane and makes its way to the other side by diffusion, while most everything that is not water is passively left behind. Ultrafilters are generally designed much like conventional mechanical filters, and those without a continuous waste stream to get rid of accumulations are subject to rapid clogging. However, RO and nanofilter membrane modules are designed with a continuous waste stream and “tangential flow” geometry which permit use for years without clogging. A special note on colloids: these are the largest molecules or the smallest particles—so small that they never settle due to gravity. They are kept in suspension forever by Brownian motion, the random vibration of microscopic particles in water caused by collisions with water molecules and other molecules dissolved in the water. Common colloidal suspensions are blood (both the blood cells and the protein molecules are colloids), tea or swamp water (tannins and lignins are colloids), and tap water containing micro-organisms and very fine turbidity. It is notable that the upper end of the colloidal range coincides with the

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sizes of the wavelengths of visible light: violet light has dimensions of about 0.44 um; red light, about 0.77 um. This size range also overlaps with that of the largest particles that can be removed by physio-chemical adsorption rather than by mechanical filtration. Thus, a mechanical filter efficient in the colloidal range can benefit from simultaneous adsorption of even smaller particles if the filter medium is, or contains, an effective adsorbent such as activated carbon. This is the advantage of precoat carbon filters and carbon block fine-filters made from powdered activated carbon: they filter mechanically down to the limit of that technology, and then there is at least the hypothetical opportunity (depending on flow rate) to remove everything else that is adsorbable, right down to the level of individual atoms and molecules. Pressure Drop: Anything put in the way of flowing water causes a resistance to the flow (by definition), slowing it down and creating a back-pressure known as the differential pressure, delta-P or ∆∆∆∆P, or pressure drop. It is simply the dynamic (flowing) line pressure provided by the municipal waterworks or other pressure source, minus the gauge pressure at the outlet of the filter (or what-ever). It is equal to the sum of all the many little resistances along the way from point A to point B: every centimeter of pipe or tubing, every fitting and valve with a reduced orifice, every particle or fiber of filter medium, even every change of direction with an ell or tee adds to the overall pressure drop of a system. Hydraulic Capacity: The importance or significance of pressure drop is that it determines the hydraulic life (capacity before becoming plugged with sediment) of filters. One end-point of mechanical filtration is the appearance of either particles of the filter media or the filtered solids in the effluent. The force of the flowing water carries particles toward the effluent, where they eventually break through if they can. If the pressure and flow are very regular, this can occur “on schedule,” and the hydraulic life is somewhat predictable. Usually, however, flow surges or water hammers cause the media to shift or flex, so that particles previously removed from the flow (filtered out) are dislodged and stream toward the outlet suddenly. This is known as channeling and dumping, and once a filter has had a channeling and dumping event, it is usually finished as an effective mechanical filter. Some media may be able to “repair” themselves momentarily and thus last a long time with repeating cycles of filtering and dumping, until the medium is either backwashed and rebedded or replaced. Channeling and dumping is both dangerous and disgusting, and it should not be permitted. It is the job of filter designers to assure that filtration will not break down (allow channeling and dumping) until the ∆P exceeds a certain level, usually 40 psi (276 kPa or 2.72 atm). That is because the minimum line pressure in the U.S. is usually about 40 psi, which means that when the ∆P is also 40 psi, there may be very little flow of water. All or nearly all of the pressure provided is consumed by the system, with little or nothing left over to provide “push” to make the water flow. Thus, a ∆P of 40 psi is usually considered the hydraulic end of life. It should be easy to see that in any system designed to transport water, a filter in the line will probably be the major contributor to

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pressure drop. And it is self-evident that: a), finer-grained media provide better mechanical filtration than coarse media, and b), finer-grained media also produce greater pressure drop than coarse media. Thus, it should be easy to see that there is a trend from coarse filters with low pressure drop to fine-filters with higher pressure drop, and that there is a prac-tical limit to the thickness or depth of filter media because of pressure drop. For granular media, the practical limit is about 30 inches (76 cm) of bed depth: beds of granular carbon, ion exchange resins or filter sand deeper than that can be expected to have excessive initial ∆P at the usual flows and pressures. It should also be easy to see that, as the filter media particle size becomes increasingly smaller, the pressure drop for a given media depth increases. Or to put it the other way, when the media particle size is decreased, the media depth must decrease if the same initial ∆P is to be maintained. This trend is taken to the logical limit in precoat filters, which are Everpure’s specialty. Everpure precoat filters use finely-powdered filter media (mostly powdered activated carbon, but also including diatomaceous earth) hydraulically deposited on and retained by a filtration barrier, which is a fabric. Cartridges are shipped with the Micro-Pure media mix dry and powdery in the bottom. When a new cartridge is first activated, the incoming water makes a slurry of the media, which begins to deposit on the fabric septum as soon as the pressure vessel fills with water and the water pressure begins to force water through. The original design uses a relatively coarse, woven polypropylene filament fabric that lets some of the Micro-Pure powder through as initial “black water” during a run-in or activa-tion, which is required for several minutes. Soon, the particles of media form interlocking structures analogous to the construction of arches and domes, creating a stable filtration layer called the precoat cake. Since the 1970s a much finer non-woven, paper-like polyethylene fabric has enabled production of precoat filters with no need for an initial activation to purge the initial filter media fines. These precoat filters all have a media depth of only a few millimeters—less than a quarter-inch—spread out over a large, pleated septum with several square feet of filtration surface area. Their mechanical filtration efficiency is NSF-Certified as >99.9% at the level of 0.5 um. The large surface area compensates for the relatively high pressure drop per unit of area, enabling acceptable overall ∆Ps and flow rates. The importance of filtration surface area is shown by the hydraulic equation:

where ∆Pt is pressure drop at time t (set at 40 psig) K is a constant specific to the conditions q is flow rate in gal/min. (held constant) A is filtration surface area

CtAqKPt 2

2

=∆

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C is concentration of particulate mater (held constant) t is time in minutes If the values of all the factors are supplied and the equation is solved for time, which then converts into gallons or liters, it can be shown that the A2 term, filtration surface area, so dominates the outcome that a much simpler “rule of thumb” can be used to estimate the increase in hydraulic capacity attributable to a given increase in surface area:

new capacity ≈ (original capacity) (filtration surface area increase)2 . For example, if the area is doubled by using two filters instead of one, the simpli-fied equation indicates that the pair should last approximately 22 or 4 times as long as one. Three filters → 32 or 9 times as long; four → 42 or about 16 times as many gallons before they plug up, and so on. (If this seems too fantastic to be true, consider that the flow rate for each filter is divided by 2 or 3 or 4 in multiple filter installations at the same time the area is multiplied, and flow rate, q2, is also a squared term.) The significance of this relationship is seen in the following comparison of the Everpure 12-in. precoat filter and comparably sized (3 in. diam. X 12 in.) granular bed and spool or carbon block filters, all flowing at 0.5 gal/min.: Filter Type Filtration Surface Area GPM/sq.ft. Granular bed 7.1 sq.in. 10.2 Spool/Block 113 sq.in. 0.64 12-in. Precoat 444 sq.in. 0.16 Granular media always channel and dump before the pressure drop reaches 40 psig, so the hydraulic equation, above, and its consequences do not apply to them. Also, many spool and block filters are far too coarse to resist breakdown when the ∆P exceeds 40 psig, and some of them may never plug up. But any filter that can qualify to claim Cyst Reduction (document 99.9% efficiency at removing live Cryptosporidium oocysts or 99.95% reduction of 3-4 um particles) will quickly become a precoat filter as particles from the water accumulate on the leading surface. Thus, it is proper to compare the expected hydraulic lifetimes of the hypothetical 3 in. diam. X 12 in. cyst-removing carbon block filter and the Everpure QC4 filter, using the “rule of thumb.” The fraction 444/113, representing their surface areas expressed as a multiple, squared, gives the precoat 3.92 or more than 15 times the hydraulic capacity of the block. However, in the real world, these carbon blocks are made to fit standard pressure vessels and are actually 2.5 in. diam. X 10 in., so the real-world numbers are (444/78.5)2 ≈ more than 30 times greater capacity. Carbon Block filters can be fabricated by molding or extruding a mixture of

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activated carbon, other special adsorbents, filter aids, and binders into a stable, unmovable structure of nearly any size, shape and porosity with a variety of possible claims. Since the grains of media are held in place by binders, carbon blocks are much less susceptible to hydraulic failure due to channelling and dumping than ordinary granular media. However, that also assures that they cannot avoid total clogging, eventually. The thickness of blocks allows the entrapment and adsorption mechanisms of depth filtration to dominate the removal of fine particles, so that they often achieve considerably greater hydraulic capacity than that predicted by their filtration surface area and the hydraulic equations discussed above. Thus, the 30 : 1 theoretical advantage of precoat filters noted in the previous paragraph should be revised to some lesser factor, say, 15-20 : 1. Carbon block filters can also have excellent adsorptive ability and capacity, since they may contain large amounts of media, sometimes exceeding that of GAC filters of the same size. Thus, unlike precoat filters, they have significant capacity for reduction of “combined chlorine” (monochloramine) and VOCs (Volatile Organic Chemicals) such as THMs and industrial solvents, and it is a simple matter to include special adsorbents for heavy metals in the media mixture. Coarse Pre-filters are used to remove most of the large suspended matter so that fine-filters will last longer. They may be large granular media filters that are backwashable, or smaller, less costly cartridge filters comprised of plastic foams, wound string filters, or fibrous mats of various compositions. They need to be inexpensive enough to permit replacement often. Almost all of them are susceptible to channelling and dumping. Multi-Media Filters use granular media of several (usually three) types that are arranged in layers according to density and particle size. They are more efficient than ordinary granular media bed-filters, but they must be backwashed more often. B. Adsorption and Adsorption Equipment As described in Section I, adsorption is a two-step process (movement from the bulk of the solution to a surface by diffusion, followed by weak chemical bonding at the surface), and that takes time, which requires a minimum contact time between the water and the adsorbent to be effective. Contact time can be increased by using more adsorbent or by reducing the flow rate. The empty bed contact time (EBCT) is an important engineering property of adsorbers, calcula-ted by dividing the volume of the filter bed by the flow rate (e.g., liters divided by liters per minute yields minutes of EBCT). Calculations involving it apply mostly to designing large municipal filters with hundreds of cubic feet of media, and not so much to the small cartridge filters used for point-of-use treatment. Scale-up calculations for these are limited to simple, direct proportions with a maximum

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multiplier or divisor of three. (e.g., if one liter of media lasts for 1000 L of test water, then three liters of media should last for 3000 L. But larger amounts should be tested again.) The granular adsorption media used by this industry are limited to bed depths of 30 inches, or about a meter, because of the pressure drop they produce at useful flow rates. Thus, there is a limit to the number of small filter cartridges that can be plumbed in series to achieve better adsorption performance. C. Ion Exchange and Ion Exchange Equipment The first step of ion exchange is identical to the first stage of adsorption—movement of contaminant molecules from the bulk of the solution to the surface of the media by diffusion and Brownian motion—but the similarities end there. While adsorption almost always involves only complete, electrically neutral (un-ionized) molecules, ion exchange occurs only with ions—materials having a + or – electrical charge. And while the second stage of adsorption (capturing the contaminant and removing it from circulation) depends only on the chemical affinity between the contaminant and the media, the capturing step for ion exchange depends mostly on the electrical charges. And finally, after capture of a contaminant, ion exchange is much more easily reversed than adsorption, enabling many cycles of use and regeneration economically and conveniently. Ion exchange depends on the existence of special media with inherent, permanent electrical charges that attract and hold ions with the opposite charge. Naturally-occurring and synthetic inorganic minerals with this property are called zeolites, which are special crystalline compounds of silicate and aluminate. Those chemical groups are both anions (ions with a negative electrical charge), and therefore zeolites can exchange only cations (ions with a positive electrical charge) such as Na+ or Ca+2. The ion exchange capacity of zeolites is not very great, so synthetic organic ion exchange resins were developed to make water softeners more marketable. Organic ion exchange resins can be formulated with either anionic or cationic ingredients of many types, yielding products with nearly every conceivable functionality and strength. Thus, there are resins formulated specifically to exchange nitrate ion preferentially, or fluoride ion, or uranium and plutonium from nuclear wastes, and so on. (These special resins are rare, and therefore very expensive.) Ion exchange is most efficient when the media is housed in cylindrical beds and the water flows downward through it. The exchange of ions is very rapid, and there is a zone or band of activity that slowly moves down the bed. Above the zone, all of the active sites are exhausted, and below the zone, all are still in completely regenerated form. The thickness or height of the exchange zone is determined by flow rate: the slower the flow, the narrower the zone. Just as the “unit (area) flow rate” in GPM/sq.ft. or L/min/m2 is the important factor for the hydraulics of mechanical filtration, the “unit (volume) flow rate” in GPM/cu.ft. or L/min/m3 is the relevant parameter for the chemical performance of media used

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for adsorption, ion exchange, oxidation, etc. The usual unitary flow rate for ion exchange systems is about 2 GPM/cu.ft. (0.214 L/min/m3 ) or 16 bed volumes per hour. 1. Softening (Conditioning): “Sodium cycle cation exchange” is the most common use of ion exchange, in which “hardness” (mostly ions of calcium and magnesium) is exchanged for sodium ions to produce “soft” or “conditioned” water. It is made possible by the availability of an ion exchange resin with millions of “active sites” made of the chemical group called “sulfonate,” which can be represented as R-SO3

− [M+]. The R represents the connection to the polymer backbone, and the M+ is the balancing cation (called a counter-ion, needed to make an electrically neutral product that doesn’t give a shock when it is touched) at each of the millions of “active sites” in each of the millions of little resin beads. Since the sulfonate is an anion, it can exchange only cations, and it happens that this particular chemical arrangement produces a strong affinity for Ca+2 and Mg+2, and a weak affinity for Na+. That means that, given a choice, the resin would “prefer” to capture Ca+2 and Mg+2 rather than Na+. (There is no real preference or choice. It’s just a matter of stability and what lasts. Just as dissolved molecules are involved in a million-billion-trillion or 1027 collisions per second, attached contaminant molecules or ions are grabbed onto and let go (sorbed and desorbed) millions or billions of times each second. If there is some affinity between the two, then the contaminant will stay attached for a greater percentage of the time, and we can say it has been adsorbed or exchanged, and the fraction of time spent attached describes the strength of the adsorption or exchange.) However, if there is no choice, or if the Na+ concentration is much, much higher than anything else, the resin will “load up” with Na+ ion on all of the millions of active sites. That is exactly what is done. The resin is initially treated with about 5% (50,000 ppm) Na+ + Cl− salt brine, and the 50,000 ppm Na+ completely overwhelms the effect of any other cations that might be in the water, so the resin loads up with Na+ ion. In “service,” this fully regenerated resin is then exposed to water containing Ca+2 and Mg+2 ions. The resin exercises its preference and exchanges two Na+ ions for each Ca+2 or Mg+2 it encounters, as long as there are any Na+ ions left attached to the resin. When they’re all gone, the resin is said to be exhausted and ready for a new regeneration with a fresh dosage of 50,000 ppm of salt brine. The resin still “prefers” to have Ca++ and Mg++ attached rather than Na+, so the strong brine solution needs to flow through the resin bed for 15-30 minutes, continuously. This assures that all of the calcium and magnesium ions on the resin have the opportunity to be displaced by the overwhelming presence of sodium ion. Such timing and the necessary initial back-washing and final rinse of the bed are done automatically by the water conditioner equipment package, so there is no need to go into detail here. It is sometimes said that softening does not change the overall TDS (Total dissolved Solids) of the water. That is not exactly true. The change is small, but there will be a slight increase due to the atomic weights of the elements that are involved, depending on the water supply. The A.W. of calcium is 40; that of

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magnesium is 24; and sodium, 23. Thus, when one calcium is exchanged for two sodium ions, their contribution to the TDS is 46, which is 15% more than the 40 they replaced. If it is magnesium that is exchanged for two sodium ions, the change would be 48 versus 40, or 20% more. Most water supplies have con-siderably more calcium than magnesium, with the average composition being probably about 2/3 calcium and 1/3 magnesium. That would produce an “average atomic weight” of about 35, compared with 46 for the two sodiums. Thus, the TDS of a hard water might increase as much as 30%, from, say, 350 mg/L to 450 mg/L after being softened. Softener resin’s special affinity for Ca++ and Mg++ is actually for “divalent cat-ions,” which means that any other member of their chemical family (the rest of column 2 of the Periodic Table of the Elements: strontium, barium, and radium), plus a great many other metal ions with two charges, will also be efficiently removed by water conditioning. Removal of radium and its radioactivity is a very important ability, as is removal of moderate concentrations of manganese and iron. Softeners also remove lead, but that contaminant comes mostly from plumbing materials that come after the softener. 2. Dealkalization is the addition of acid in some form, to drive the alkalinity equilibrium to the left (see the Chemical Equilibrium discussion in Section I), so that some of the alkalinity is lost as CO2. To do that with ion exchange requires loading the resin’s (or zeolite’s) active sites with H+, which is the embodiment of acidity. That is accomplished by regenerating the bed with a 5% acid solution (50,000 ppm H+) instead of a salt brine. The resin’s active sites have a “preference” for Ca++ and Mg++ and other divalent cations, but the high concen-tration of H+ ions overwhelms the preference and eventually removes all of the calcium and magnesium if the regeneration lasts long enough.. During the service cycle, the resin exercises its preference and readily exchanges each Ca++ and Mg++ ion in the water for two H+ ions. These affect the alkalinity equilibrium instantly and cause some to be lost as CO2 gas. Thus, the overall result is removal of both hardness and alkalinity, and a reduction in TDS as well. Dealkalization can be done with the same type of ion exchange resin as that used for softening, called strong-acid or strongly acidic cation exchange resin, but that is risky: when there is more hardness in the water than alkalinity (a common condition), more H+ will be produced than there is alkalinity to be neutralized, resulting in an acidic, corrosive, over-treated water. Therefore, it is more prudent to use a different type of resin, called weak-acid or weakly acidic cation exchange resin, sometimes referred to as “WAC resin,” for Weak Acid Cation. The difference between them is the strength to overpower the effects of weak acids, or buffers, in the water. (See the discussion of buffers in Section I.) When a WAC resin is regenerated with acid and used to exchange H+ for hard-ness ions, it works only as long as alkalinity remains available to neutralize the H+. If the water has more hardness than alkalinity, WAC resins in H+ form work only up to the point of destroying all of the alkalinity, and then they stop. For

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example, in treating a water supply with 300 ppm of hardness as CaCO3 but only 250 ppm of alkalinity as CaCO3, WAC resin would exchange only 250 of the 300 ppm, leaving 50 ppm of hardness and zero alkalinity, with no risk of over-treating and creating a corrosive condition. If the water has more alkalinity than hard-ness, WAC resin will exchange all of the hardness and neutralize the same amount of alkalinity. There is also a third ion exchange process for dealkalization, and that is using a strong-base or strongly-basic anion exchange resin in chloride form. It is regenerated with NaCl salt brine just like softening; the difference is that the resin exchanges two Cl− ions for CO3

−2 ions instead of two Na+ ions for Ca++ or Mg++ ions. The system is set up and run exactly like a softener. This is hardly ever done, because softener resin is among the cheapest and strong-base anion resin the most costly media available. 3. Demineralization or deionization is the removal by ion exchange of all ions in solution, both cations and anions. That requires use of both cation exchange resin in H+ form and anion exchange resin in OH− form. That means that the cation resin is regenerated with acid (H+ ion), and the anion resin is regenerated with strong base or OH− ion. The cation resin exchanges all cations for H+ ion, and the anion resin exchanges all anions for OH− ion. The two combine to form water, with nothing left over, so that the treated water is similar to distilled water quality. It is not identical to distilled water, because the ion exchange process does not remove electrically neutral molecules such as sugars, alcohols, many pesticides, VOCs, and other un-ionized organic compounds which may be present. The two resin types may be combined in a “mixed bed” or in separate beds or cartridges. It is difficult (but possible) to regenerate a mixed bed, because the two resins must be separated first. This is done by “fluidizing the bed” (back-washing, or operating up-flow) so that the resin types separate by density. Once separated, they are regenerated separately and then re-mixed. The amounts of the resins used must be chemically balanced, so that they both become exhausted at the same time. Both the anion and cation resins may be either strong or weak-acting resins, but the usual (most economical) combination is a strong-acid cation resin paired with a weak-base anion resin. Since a weak-base anion resin cannot exchange the anions of weak acids unless the conditions are already acidic, the strong-acid cation resin in H+ form is always put first. To illustrate: a weak-base anion resin in OH− form will not exchange HCO3

− and CO3−2 ions unless there is free H+

already available to neutralize the OH−. (This is analogous to the situation with WAC resin [see above, under Dealkalization], which will not exchange H+ for hardness unless sufficient alkalinity is already available to balance it.) If the water encounters the cation resin loaded with H + first, then the water will have exchanged all cations for H+ already and be strongly acidic when it enters the

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cartridge/bed containing the weak-base anion resin. D. Oxidation filtration There are two kinds of filter media which can be used in beds or cartridges to oxidize ferrous iron, manganese and sulfide ion/hydrogen sulfide in water. (See the Iron Water and Sulfur Water discussions in Section II.) One is a zeolite called “manganese greensand,” either natural or synthetic, loaded up or exchanged with Mn++ ion and then oxidized (regenerated) with chlorine or potassium permanganate (KMnO4). The Mn++ on the zeolite becomes MnO2, called manganese dioxide, which is a good oxidizing agent. When water containing Fe++ or Mn++ or S−2 flows through a bed of this manganese greensand in MnO2 form, they are readily oxidized to Fe+++ or MnO2 or S, respectively, which are all particles for a fine-filter to remove. At the same time, the MnO2 on the zeolite is reduced back to its original Mn++ ionic form, which must then be re-oxidized (regenerated) back to MnO2 again in a new cycle. (Careful—don’t be confused by the fact that we’re using manganese to remove manganese and iron and hydrogen sulfide. The Mn++ in the water is oxidized to MnO2 and removed by filtration later. The MnO2 on the zeolite gets reduced to Mn++ in the process, but that Mn++ is supposed to stay attached and be recycled back to MnO2.) When manganese greensand is used to oxidize Mn++ ion or Fe++ ion or S−2 ion in water, the reactions are: Zeolite-MnO2 + Mn++ (dissolved) ⇒ Zeolite-Mn++ + MnO2 “ + Fe++ (dissolved) ⇒ “ + Fe+++ [⇒ Fe(OH)3 ] “ + S−2 (dissolved) ⇒ “ + S All three end-products—the MnO2, Fe(OH)3, and S—are filterable particles. Also do not be confused because one of the possible regenerant chemicals is another form of manganese, potassium permanganate. It’s just a coincidence that three different forms of manganese can be involved in this technology. If a bed of manganese greensand is allowed to become completely exhausted, the Mn++ attached by ion exchange to the zeolite may be lost—exchanged—and require being regenerated by ion exchange before oxidizing it to MnO2 for another cycle of oxidation filtration. Therefore, it is important that manganese greensand installations be regenerated promptly on schedule, rather than waiting until the oxidizing function begins to fail. By then, some of the MnO2 will already be gone from the zeolite, and peak efficiency will be lost. The other redox medium is copper-zinc alloy (brass). It was realized only recently that one reason disinfectant chlorine dissipates so quickly in the plumbing overnight is that the chlorine actively corrodes (oxidizes) the surface of copper pipe and brass fittings, which then liberate ions of Cu+ or Cu++ or Zn++ into the water. (This is also the source of nearly all lead contamination in

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water—corrosion of brass and solder containing 8% or more lead, used to make the materials machinable and less brittle. Granular brass filter media is made from pure alloy, so there is no lead hazard.) The presence of copper and zinc atoms next to each other in an alloy allows an easy exchange of electrons, and there is a natural tendency for electrons to flow from zinc to copper. If the two were present as separate bars of zinc and copper metal connected by a wire, as in the example in the discussion of Oxidation and Reduction in Section I, the flow of electrons would do work, such as lighting a light bulb. In an alloy no wire is needed, but electrons do not flow spontaneously because there is no place for them to go. But if the alloy is immersed in water containing contaminants like chlorine that can be reduced, the contaminants accept the electrons from the zinc, which then becomes Zn++ ion in solution. Later, further down in the bed of granules, the Zn++ ions tend to precipitate out as zinc oxide (ZnO) or hydroxide [Zn(OH)2] and slowly coat the alloy, leaving it less and less effective. This can occur slowly or rapidly, depending on the pH, the amount of chlorine, and the level of dissolved oxygen in the water. Some Cu++ ions may be liberated, but this again depends on the levels of chlorine, oxygen, and acidity. The net result is a low but acceptable concentration of both Zn++ and Cu++ ions in the treated water. The U.S. limits are 5 ppm for zinc (recommended) and 1.3 ppm for copper (mandatory). The exact electrochemical mechanism is not quite so clear for oxidation of Fe++ or H2S by granular brass filtration. However, the voltages for both of these two oxidations lie between those for copper and zinc in the Electromotive Series (which see, Section I), which means that electrons can go either way (cause either oxidation or reduction) between them. The reduction of hydrogen sulfide is said to proceed by oxidation of copper and zinc, which then precipitate the sulfide as ZnS and CuS for later removal by mechanical fine-filtration. E. Chlorination, Coagulation, and pH Adjustment Chlorination, coagulation, and pH adjustment are very different activities, but they are often accomplished simultaneously using the same equipment: a chemical feed pump with its reservoir of concentrated chemicals, followed by a contact/mixing tank. If the oxidant/ disinfectant is liquid chlorine bleach (5.25-18% sodium hypochlorite, NaOCl), the coagulant is sodium aluminate (NaAlO2), and the pH reagent is sodium hydroxide (NaOH, “caustic soda”) or sodium carbonate (Na2CO3, “soda ash”), they can all be mixed together and fed with a single pump. The preferred process is a variation of standard practice known as “super-chlorination-dechlorination” because a large excess of chlorine is used (about 10 times the usual dosage) and then removed before drinking. It is based on the concept of the “CT envelope,” which states that the arithmetic product of one combination of the chlorine concentration in ppm, C, times the contact time in minutes, T, is equivalent in disinfection ability to any other combination that

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produces the same total. For example, 30 minutes of contact time with 0.5 ppm FAC is equivalent to 5 minutes at 3 ppm, because they both total a CT of 15, which is the minimum for less-than-ideal conditions. Large waterworks have the time and space to use low chlorine concentrations with long contact times, but it is more economical for small systems to use relatively small contact tanks with ten times as much chlorine as usual, and then waste 95% of it. At a flow rate of 5 gal/min., a 5-minute contact time would only require a 25-gallon tank (if the mixing were perfect), while a 30-minute contact time would require a 150-gallon tank. (Correction for mixing efficiency changes those figures to 89 gallons and 536 gallons. Details on page 55.) Sodium hypochlorite solutions are made by bubbling chlorine gas up through a solution of sodium hydroxide: Cl2 + NaOH → H+ + Cl− + NaOCl (not an equilibrium reaction) The resulting solution is pH 10-12, depending on the local producer, and the free chlorine is present as the sodium salt of hypochlorous acid, which is a “weak” acid with a KEq = 10−7.5. However, NaOCl itself is “strong”—a simple salt like NaCl that is 100% ionized the instant it is dissolved: NaOCl + H2O → Na+ + OCl−−−− (not an equilibrium reaction) Only the acid, HOCl, is weak: OCl− + H+ ⇔ HOCl KEq = 10−7.5 That implies two things: first, that when the chlorine bleach is injected into the water, it will instantly change from being all hypochlorite ion, OCl−, in the bleach to being mostly in the form of the un-ionized hypochlorous acid molecule, HOCl, in the water. Second, hypochlorous acid’s KEq is close enough to the H+ concentration in neutral water to interfere with the pH of the water into which it is being injected. Actually, because of the small amounts involved, it’s the other way around: the pH of the water interferes with the hypochlorous acid equilib-rium. This is extremely important, because only the HOCl molecule disinfects. All forms of free chlorine are equally effective as oxidizing agents for iron and manganese and hydrogen sulfide, and OCl−−−− ion can kill some viruses, but if the pH is too high to allow most of the HOCl to exist as a neutral molecule that can penetrate cell membranes, there will be little disinfection of bacteria. The 50/50 point in the dissociation of hypochlorous acid occurs at pH 7.4, and if the pH rises to 8.5 the HOCl will be 95% dissociated and useless as a bactericide. Therefore, it is important not to raise the pH any higher than necessary if the chlorination system is expected to disinfect as well as oxidize. Sodium Aluminate is a source of “alum,” or aluminum salts that form a fluffy precipitate called “floc” when dissolved in water at ordinary water pH. If the concentration is high enough, the pH is right, and there is enough gentle mixing

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to promote agglomeration of the precipitate into larger clumps which entrap particles of dirt and turbidity (called coagulation), the water can be greatly clarified without the use of fine-filters. In large waterworks there is lots of space and time, and after flocculation and coagulation the floc is allowed to settle, and virtually all of the suspended particles sink to the bottom enmeshed in a solid, jelly-like mass. The clear water is then sent to large granular media filters, but the filters only have to remove the occasional clump of floc that gets stirred up. Everything else has already been removed as sludge. Small, individual water supply systems do not have the luxury of acres of space for water treatment, so there is no settling phase, and the final filters are more important. Aluminum is somewhat unusual in its ability to precipitate as a hydroxide and then to re-dissolve as a “complex ion” with the addition of more strong base: Al+++ + 3OH− ⇔ Al(OH)3 (solid aluminum hydroxide “floc”) + OH− ⇔ AlO2

− (soluble aluminate ion) + H+ + H2O Thus, if a solution of sodium aluminate (NaAlO2) is added to water with a “normal” pH (near pH 7) or lower, the aluminate ion will convert to the Al(OH)3 form and precipitate, or flocculate as a fluffy, gooey solid material that entraps colloids and particles. Sodium aluminate can be purchased as a 10% solution or prepared by dissolving any aluminum salt in water and then adding NaOH until it first precipitates and then redissolves. Note that both the chlorine bleach (sodium hypochlorite) and the flocculant (sodium aluminate) are produced as sodium hydroxide (NaOH) solutions. Both chemicals will affect the pH only slightly, because they are added only at low ppm levels. However, more NaOH can be added to either at any time. That means that the final pH of any water supply that is treated with either or both chemicals can be raised to any desired level by adding more NaOH, and they can all be mixed together and fed as a single solution. As an alternative, sodium carbonate (Na2CO3, “soda ash”) may be used to raise the pH instead of NaOH. This is sometimes a good idea, since NaOH is so dangerous to use, while sodium carbonate is relatively innocuous. However, since it is not as strongly alkaline, Na2CO3 may not be strong enough to keep sodium aluminate from precipitating too early, before it is fed, thus clogging the pump. This is seldom a problem except when the brand of bleach locally available is deficient in NaOH. So, if addition of sodium aluminate to chlorine bleach causes immediate precipi-tation, NaOH is the only recourse. Preparing such a mixed chemical feed solution is unfortunately rather complicated. The best way is to treat each ingredient separately in a “jar test” as described below, and then combine them in the indicated ratio. The setting on the feeder pump is then determined by trial and error.

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Determining the Chlorine Dose: You will need a gallon jug, an ordinary bucket, an eye-dropper, and a chlorine test kit that measures free chlorine (not “total chlorine”). Put a gallon of the water to be treated into the bucket and start adding chlorine bleach drop by drop, counting the drops. After 5 drops, mix and wait 5 minutes, and then test for free chlorine with the test kit. The 5-minute wait is to allow time for the chlorine to react with any iron, sulfide, ammonia, organics, whatever, that may be present and acting as “oxidant demand” or “chlorine demand.” The demand, if any, must be destroyed before any extra chlorine left over (called the “residual”) can be made available to do any disinfection. If the demand is not destroyed, the chlorine that would be measured with a total chlorine test would include any “combined chlorine” consisting of monochlor-amine, dichloramine, and organic chloramines derived from amino acids, etc., which are all very poor disinfectants and also very weak oxidizers. (They are very stable and last a long time in the pipes, which is one reason monochlor-amine is sometimes intentionally produced by large water works with large distribution systems. But that is done only after effective disinfection with free chlorine has already been completed.) The goal in the small chlorinator systems used for private wells is a concentration of 3-5 ppm free chlorine. If it’s already higher than that after 5 drops, dump the bucket and start over with a diluted bleach solution. Repeat the process of adding drops of bleach, mixing, waiting 5 minutes, and testing as many times as it takes to find the total number of drops of bleach needed to destroy the demand and produce a residual of 3-5 ppm free chlorine after a five-minute contact time. Record the number of drops. Determining the Sodium Aluminate Dose: Using the same gallon jug, bucket and eye-dropper (cleaned up, of course!), add and count the drops of sodium aluminate needed to coagulate all of the turbidity or color in the water. After every 5 drops, stir very strongly for about 30 seconds, then back-stir for a moment to quiet the solution, and then let it stand undisturbed for several minutes, until the floc begins to settle. As soon as the clear solution above the cloudiness can be seen, judge for yourself whether it is clear or color-less enough. If not, add more sodium aluminate, mix violently, wait, and look again. Repeat until you find the total number of drops needed per gallon to achieve the effect you want. Putting it All Together: The total number of drops of chlorine bleach and sodium aluminate represent the ratio of the volumes of solution to mix together. For example, if the chlorine dosage is 10 drops and the alum dosage is 20 drops, mix the solutions in the ratio of 1:2. Then try out the mixed solution in the actual installation: set the chemical feed pump at half-speed, turn on the water, and take samples. Adjust the setting as needed to produce a free chlorine residual of 3-5 ppm . But first, we need to discuss the rest of the system. Designing a Chemical Feed System: The starting point is the desired or expected peak flow rate. For most homes and small businesses, 5 gallons per minute or 19 liters per minute is about right. The size of the contact tank or

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mixing tank required is calculated by multiplying the flow rate times the desired contact time and then dividing by a mixing efficiency factor: (5 gal/min.) (5 min.) = 25-gal. tank; ÷ 28% = 89-gallon tank. The efficiency of mixing in ordinary pressure tanks is very poor. Treatment chemicals and treated water “short-circuit” from influent to effluent in contact tanks quickly, providing much less actual contact time than expected. The best that can be done is about 28% efficiency, and even that requires filling the bottom third of an up-flow tank with pea-sized gravel. Thus, at a flow rate of 5 gal/min., a standard 82-gallon (310 L) galvanized steel tank with pea-gravel will provide about 4.6 minutes of contact time; a 120-gal. (450 L) tank, 6.7 min. F. Membrane Systems This section is not about the “microporous polymeric membranes” used to make pleated membrane micro-filters and also to collect and concentrate bacteria in microbiology labs. Those are discussed in Sec. III-A, Mechanical Filtration. Membrane systems are entirely different: they have no “pores,” and they do not “filter” anything. Instead, they are formulated to be so extremely hydrophilic (“water-loving”) that water dissolves into the membrane, and other materials (chemicals, particles, microbes) are simply excluded by their chemical nature. Exceptions—things membrane systems don’t remove well—include small, water-like molecules such as alcohols, and other very small molecules, even if they aren’t much like water, such as THMs and other VOCs. Once water molecules are in the membrane, they find their way to the other side by diffusion and leave everything else behind. The three major technologies using membrane systems are reverse osmosis (RO), nanofiltration, and ultrafiltration. Several membrane types are in use, and the most important ones are the cellulose acetate types, the polyamide types, and the thin-film composite (TFC) types. The polyamide membranes are readily damaged by disinfectant chlorine, but the cellulose acetate membranes are not. TFCs are intermediate in chlorine sensitivity, and both are often protected by activated carbon or granular brass prefilters. Chlorine-resistant cellulose acetate membranes are used where the absence of chlorine would allow degradation of the membrane by bacteria. 1. Ultrafiltration is the coarsest of the membrane technologies, generally being limited to removing colloids the size of viruses and molecules the size of large proteins and “color bodies” derived from rotting vegetation. Ultrafilters are produced as long sheets of membrane that are rolled up into a cylinder. Alter-nating layers of netting provide space for water to penetrate into the interior of the roll. This provides more surface area than the pleats often used for micro-filters. They require good prefiltration by a micro-filter or fine-filter because they clog quickly. Therefore, they do not find much use in water treatment. Some food processing applications provide a “waste” stream that recirculates. This helps delay clogging and enhances the life of the membrane.

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2. Nanofiltration uses membranes that remove particles and molecules hundreds of times smaller than ultrafilters, down to a molecular weight of about 1000. They are produced the same way as ultrafilters—rolled-up membrane cylinders with netting spacers—but the design also includes a “reject” stream that directs most of the water to waste before it can become so concentrated that it causes clogging. This produces a “tangential flow” pattern that is less conducive to clogging. That means there are three plumbing connections for nanofilters: inlet, outlet product water, and outlet waste water. This is inefficient in terms of wasting a lot of water, but highly efficient in terms of saving money on premature replacement of expensive membrane cartridges. Nanofiltration can be used alone to purify many water supplies without using any other equipment or chemicals, because it can remove all particles of all sizes: all colloids, all types of micro-organisms, nearly all pesticides (except for those that are so small that they qualify as VOCs), and nearly all organic material that can act as food for bacteria or serve as the precursors of THMs and other disinfection byproducts. It can even remove 2/3 or more of ordinary hardness and alkalinity, so that nanofiltration is often called “membrane softening.” 3. Reverse Osmosis is similar to nanofiltration, except that the membrane is a thousand times more efficient, with the ability to remove 90% or more of indi-vidual atoms and ions. Ions with only one charge are removed less efficiently than those with multiple charges. Thus, RO removes hardness and alkalinity with 97-99% efficiency, while flouride (F−), nitrate (NO3

− ), cyanide (CN− ), silver (Ag+), etc. are removed only with about 95% efficiency. Mercury (Hg+) is the most difficult, with efficiencies as low as 60% in some cases. As an exception, some membranes are designed specifically for desalination of sea water, and these remove Na+ and Cl− with greater than 99% efficiency. RO unfortunately wastes even more water than nanofiltration, often as much as 90%. Usually, about 8 volumes are wasted to produce one volume of purified water. Ultrafiltration and nanofiltration membranes are too coarse to develop much of an osmotic pressure across them, (which see, Section II), but of course RO membranes do. As a rule of thumb, one psi (6.9 kPa) of osmotic pressure is produced by every 100 ppm of total dissolved solids. That means 2000 ppm of TDS will produce about 20 psi (138 kPa) of osmotic back-pressure to overcome, and that is the usual limit of water quality for RO systems that use only the available water pressure to drive the process. Higher pressures, supplied by pumps, allow RO systems to be both more efficient and more productive. Such larger systems often have elaborate pretreatment to improve efficiency and reduce maintenance. In addition to basic micro-filtration and chlorine removal, this may include softening by ion exchange, injection of anti-scale chemicals, and a use of a hot/cold mixing valve. RO systems produce water so slowly that the product water is always collected in tanks for dispensing later. Large systems with their own pumps usually have a second pump to repressurize the product water, but the small systems that use

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ordinary line pressure often use small “bladder” tanks that are pre-pressurized, usually to 5 psi (34.5 kPa). These have only a single connection to serve as both inlet and outlet. The RO system produces its trickle of water, which collects in the tank as long as it can overcome the 5 psi pressure, or until the back-pressure reaches 2/3 of the line pressure, and RO systems with a special pressure-sensing valve shut themselves off. This saves water that otherwise would run to waste while the tank is full. When water is drawn out for use, it is propelled by the pressure in the tank, which is always at least 5 psi even when empty. The reduction in pressure caused by using some of the water actuates the pressure-sensing valve and allows systems to start up again and begin producing more RO water. Unfortunately, even RO membranes are unable to remove the small, low-molecular-weight, volatile organic compounds (VOCs) such as THMs, solvents, and some pesticides, and these must still be removed by adsorption to activated carbon. Some systems are designed with the carbon adsorber immediately after the membrane, subject to the slow trickle of product water, and others place the adsorber between the bladder tank and the dispensing faucet. Either way, only the product water is so treated. It would be a tremendous waste of carbon to remove THMs, etc. from the influent water and then waste 90% of it as reject water. G. Ultra-Violet Irradiation Electromagnetic radiation (light) covers an enormous range of wavelengths and energies, from very weak but long radio waves many miles in length to extremely powerful X-rays and gamma-rays with wavelengths only tiny fractions of a micron in size. The size of the wave is inversely proportional to the energy it carries, which means that light with the smaller wavelengths can do more damage. The visible portion of the spectrum has wavelengths that range from about 0.40 um or 400 nm for violet light to about 0.77 um or 770 nm for red light. Light with wavelengths shorter than violet (but greater energy) is ultra-violet, and light with wavelengths longer than red (but less energy) is infra-red. Thus, infra-red light is simply “heat,” while visible light can give you a suntan, and UV light can cause terrible burns and blindness. It happens that UV light with a wavelength near 0.254 um or 254 nm is able to produce resonance effects in the DNA of living cells which cause breaks and kinks, leading to cell death. Therefore, UV irradiation is useful as a method of disinfection for water. The energy is measured in watts (micro-watts) per cm2, and the total dosage is the wattage multiplied by the time: uWsec/cm2. Thus, a standard UV bulb that delivers 3800 uW/cm2 will produce a dosage of 38,000 uWsec/cm2 after ten seconds of exposure. That is the minimum dosage required to meet the NSF International standard for Class A “purification” of raw waters that may contain pathogenic bacteria and viruses. A lesser standard of 16,000 uWsec/cm2 can be NSF-Certified for controlling bacterial regrowth in the pipes of

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systems in which the water has already been disinfected, and only non-patho-genic organisms are present. Standard UV systems are not powerful enough to kill protozoan cysts and oocysts and the larger parasites, so these still require physical removal. New pulsed-UV systems may prove to be effective against cysts, but they are not yet approved by any health regulatory agency. Advantages of UV treatment:

1. Treatment is essentially instantaneous. It is not difficult to provide a 10-second contact time (an exposure vessel of less than a liter at a flow rate of 1 gallon/min. (3.79 L/min.)) with modest resources.

2. It’s all electrical, with no moving parts to break down, so the system can be relatively simple in design and maintenance.

3. Because it’s all electrical, fail-safe measures are relatively easy to build in. 4. Water clarity is extremely important, but fine-filtration to remove turbidity

will also remove any protozoan cysts and larger parasites that would not be killed.

Disadvantages of UV treatment:

1. There is no residual activity to protect the water against subsequent contamination, so the overall system design may be complicated by the need to place the UV last.

2. Routine maintenance to clean the optical surfaces is mandatory; the fail-safe system should turn the system off if/when they get dirty.

3. Special meters for monitoring the wavelength and intensity of the UV light are very costly, so complete systems cannot be inexpensive.

4. Treated water should be kept in the dark for 30 minutes, because ordinary sunlight has the ability to activate repair enzymes in many bacteria, and as many of 2/3 of those previously killed may be revived after only a moment’s exposure.

5. There is a special design problem with small systems intended for intermittent use: cold lamps require a warm-up to achieve peak efficiency, but turning the system on and off repeatedly damages the lamp, and leaving it on all the time heats the water and stimulates the growth of the few organisms that survive.

6. UV treatment must not be used in recirculating systems because that can produce a “super-strain” of UV-resistant bacteria that might be dangerous.

7. UV light is damaging to many plastics, so special shielding may be important.