CHAPTER 6WATER CONTAMINANTS:

OCCURRENCE AND TREATMENT

The introduction of contaminants into water supplies has been shown to berelated to rainfall, the geologic nature of the watershed or underground aquifer,and the activities of nature and the human population. Water contaminants to beexamined in more detail fall in two categories: dissolved matter (Table 6.1) andnonsoluble constituents (Table 6.2). Dissolved gases are included in discussionsof the biological cycles affecting water quality.

As shown in Table 6.1, soluble materials in water are arbitrarily assigned tofive classifications, the first four of which are based on concentration levels, withthe last covering those materials usually transient because continuing reactions inthe aquatic environment change their concentrations.

Many materials are transient because of biological activity. The change in CO2and O2 content with sunlight is one example. Equilibrium between NH3, N2,NO2", and NO3" is another, discussed later in this chapter as part of the nitrogencycle. (See Class 2, Secondary Constituents.)

There are also longer term processes by which nature cycles matter throughliving organisms, which in turn modify the environment and leave their recordsin the rocks. This chapter examines the sources of contaminants in water, manyof which are minerals created by living things. Perhaps the best known are thechalk cliffs of Dover and the coral atolls of the Pacific, both composed OfCaCO3.

Discussing these atolls in his essay on formation of mineral deposits, C. C.Furnace says, "To the casual observer, it would seem that the polyp has builtthese great masses of land out of nothing; but, of course, it cannot do that anymore than man can. It has taken calcium compounds from very dilute solutionsof sea water and built up a shell of calcium compounds to protect itself. In thisprocess of following its preordained metabolic rite, it has concentrated calciumby several thousandfold in the form of an insoluble compound. Insignificant asthe coral polyp may appear, it is one of the most important creatures in changingthe character of the earth's surface."

Many other natural cycles have been operating over countless geologic ages toproduce deposits of sulfur, iron, manganese, silica, and phosphate, to name onlya few. In the village of Batsto, New Jersey, the early American colonists set upthe first blast furnace in the New World. Their source of iron ore was "bogiron"—pure iron oxide precipitated from artesian water by iron-depositingbacteria.

So the presence of many of the mineral constituents in water supplies maysimply represent the return to the aquatic environment of a loan made by theearth to living organisms long ago.

Class 5 Transient constituentsAcidity-alkalinityBiological cycles

Carbon cycle constituentsOrganic C/CH4/CO/CO2/(CH2OyC-tissue

Oxygen cycleO2/CO2

Nitrogen cycle constituentsOrganic N/NHa/NO2-/NO3-/№/amino acids

Sulfur cycle constituentsOrganic S/HS-/SO3

2-/SO4

2-/S°Redox reactions

Oxidizing materialsFrom the natural environment—O2, STreatment residues—Cl2, CrO4

2"Reducing materials

From the natural environment—Organics, Fe2+, Mn2+, HS~

Treatment residues—Organics, Fe2+, SO2, SO3

2'Radionuclides

TABLE 6.2 Nonsoluble Constituents inWater Supplies

Class 1—SolidsFloatingSettleableSuspended

Class 2—Microbial organismsAlgaeBacteriaFungiViruses

SodiumSulfateTotal dissolved solids

PotassiumStrontium

PhosphateZinc

TinTitanium

Class 1

Class 2

Class 3

Class 4

Primary constituents—generally over 5 mg/LBicarbonate MagnesiumCalcium Organic matterChloride Silica

Secondary constituents—generally over 0.1 mg/LAmmonia IronBorate NitrateFluoride

Tertiary constituents—generally over 0.01 mg/LAluminum CopperArsenic LeadBarium LithiumBromide Manganese

Trace constituents—generally less than 0.01 mg/LAntimony CobaltCadmium MercuryChromium Nickel

FIG. 6.1 Theoretical solubilities of carbonate compounds in a water system closed to an externalCO2 environment at 250C. (From Stumm and Morgan, 1970.)

FIG. 6.2 Theoretical solubilites of oxides and hydroxides in water at 250C. (FromStumm and Morgan, 1970.)

As an aid to appreciating the solubilities of the constituents being examined,and thus the limitations of their occurrence in natural water supplies and theresiduals which may be reached in precipitation processes, Figures 6.1, 6.2 and6.3 present solubility characteristics in appropriate locations in the text.2

CLASS 1—PRIMARY CONSTITUENTS

This category includes dissolved solids generally exceeding 5 mg/L, and often sev-eral orders of magnitude above this level.

Bicarbonate (HCO3"—Molecular Weight 61)

The bicarbonate ion is the principal alkaline constituent of almost all water sup-plies. It is generally found in the range of 5 to 500 mg/L, as CaCO3. Its introduc-

Con

cent

rotio

n, m

illim

oles

/lite

r

FIG. 6.3 Theoretical solubilities of oxides and hydroxides in water at 250C, showingamphoteric nature of aluminum and zinc. (From Stumm and Morgan, 1970.)

tion into the water by the dissolving action of bacterially produced CO2 on car-bonate-containing minerals has been explained elsewhere. Normal activities ofthe human population also introduce alkaline materials into water, evidenced bya typical increase of alkalinity of sewage plant effluent of 100 to 150 mg/L abovethe alkalinity of the municipal water supply. Much of this is due to the alkalinityof industrial and domestic detergents.

Alkalinity in drinking water supplies seldom exceeds 300 mg/L. The controlof alkalinity is important in many industrial applications because of its signifi-cance in the calcium carbonate stability index. Alkalinity control is important inboth concentrated boiler water and cooling water in evaporative cooling systems.

Con

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ratio

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Makeup for these systems must often be treated for alkalinity reduction either bylime softening or direct acid addition. Alkalinity is objectionable in certain otherindustries, such as the beverage industry, where it neutralizes the acidity of fruitflavors, and in textile operations, where it interferes with acid dyeing.

Calcium (Ca2+—Atomic Weight 40; Group 11 A, Alkaline Earth Metal, Figures6.1 and 6.3)

Calcium is the major component of hardness in water and usually is in the rangeof 5 to 500 mg/L, as CaCO3, (2 to 200 mg/L as Ca). It is present in many minerals,principally limestone and gypsum. Limestone deposits are often the residue of thefossils of tiny aquatic organisms, such as polyps, that have taken calcium fromthe seawater in which they lived, and used it for their skeletons. This is but oneof many cycles in nature whereby some component of the environment is contin-ually withdrawn by living things and eventually returned directly or indirectly.

Calcium removed from water in softening operations is later returned to theenvironment, often to the watershed, by way of a precipitate or a brine which isthe by-product of the softening reaction. Calcium is a major factor in determiningstability index. Calcium reduction is often required in treating cooling towermakeup. Complete removal is required for many industrial operations, particu-larly for boiler makeup, textile finishing operations, and cleaning and rinsing inmetal finishing operations.

Calcium hardness can be reduced to a level of 35 mg/L as CaCO3 by cold lime-soda softening and to less than 25 mg/L by hot lime-soda softening. It is reducedto less than 1 mg/L by cation exchange methods.

Chloride (Cl —Atomic Weight 35.5; Group VIIA, Halide)

Since almost all chloride salts are highly soluble in water, chloride is common infreshwater supplies, ranging from 10 to 100 mg/L. Seawater contains over 30,000mg/L as NaCl, and certain underground brine wells may actually be saturated,approximately 25% NaCl. Many geologic formations were once sedimentaryrocks in the sea, so it is not surprising that they contain residues of chlorides thatare continually leaching into freshwater sources. The chloride content of sewageis typically 20 to 50 mg/L above the concentration of the municipal water supply,accounting in part for the gradual increase in salinity of rivers as they proceedfrom the headwaters to the sea.

Anion exchange is the only chemical process capable of removing chloridesfrom water; however, physical processes such as evaporation and reverse osmosiscan separate a feedwater into two streams, one with a reduced chloride and theother with an increased chloride content.

The recommended upper limit for chloride in drinking waters is 250 mg/L,based entirely on taste, not on any known physiological hazards.

Magnesium (Mg2+—Atomic Weight 24.3; Group MA, Alkaline Earth Metal, Figs.6.1 and 6.2)

The magnesium hardness of a water is usually about one-third of the total hard-ness, the remaining two-thirds being calcium hardness. Magnesium typically

ranges from 10 to 50 mg/L (about 40 to 200 mg/L as CaCO3). In seawater, mag-nesium concentration is about 5 times that of calcium on an equivalent basis. Theproduction of magnesium hydroxide from seawater is the starting point in themanufacture of magnesium. Magnesium is a prominent component of many min-erals, including dolomite, magnesite, and numerous varieties of clay.

Since magnesium carbonate is appreciably more soluble than calcium carbon-ate, it is seldom a major component in scale except in seawater evaporators. How-ever, it must be removed along with calcium where soft water is required forboiler makeup or for process applications. It may be removed by lime softeningto a residual of 30 to 50 mg/L as CaCO3 cold, or 1 to 2 mg/L as CaCO3 hot. It isalso reduced by ion exchange to less than 1 mg/L as CaCO3.

Organic Matter (Carbon, C4+—Atomic Weight 12; Group IVA, Nonmetal)

Since organic material makes up a significant part of the soil and because it isused by aquatic organisms to build their bodies and produce food, it is inevitablethat water-soluble organic products of metabolism should be present in all watersupplies. There is not much information available on specific organic compoundsin most water sources. (See Table 4.7.) There are literally hundreds of thousandsof known organic compounds, many of which might somehow find their way intothe hydrologic cycle. A complete "organic analysis" of water is impossible. How-ever, one of the by-products of space-age technology has been the development ofnew instruments for organic analysis (see Chapter 7). With these instruments, theanalyst can develop methods of analysis for organic materials of interest—espe-cially those considered by the EPA to be toxic or carcinogenic, such as PCBs(polychlorinated biphenyls) and TTHMs (total trihalomethanes). But unless suchspecific organic compounds are requested at the time a sample is presented, theanalyst uses indirect measures of organics instead (e.g., COD, TOC).

Many waters have a yellowish or tea color due to decayed vegetation leachedfrom the watershed by runoff. These organic materials are broadly classified ashumic substances, further categorized as humic acid (a water-soluble compound),fulvic acid (alkali-soluble material), and humin (high molecular weight, waterinsoluble matter). These organic compounds are molecules having many func-tional groups containing oxygen and hydrogen atoms in various proportions, sothat when organic matter is reported as carbon, as it is in the TOC determination,it is probable that the molecular weight of these humic organic molecules is 2.0to 2.5 times greater than the value reported as carbon. A survey of 80 municipalsupplies in the United States showed an average total organic carbon content inthe finished water of 2.2 mg/L, as C, so the organic matter was probably on theorder of 5 mg/L.

There are a variety of indexes for measuring the gross organic content of water,and there is generally no correlation between them. The organic matter at a sam-pling station on the Mississippi River as determined by these indexes is shown inFigure 6.4. Because some of the functional groups in humic compounds have ionexchange properties, they tend to chelate heavy metals. In spite of this, there isno correlation whatever between the color of a water and its total heavy metalconcentration. A study of the Rhine River showed that humic substances com-prised from 25% (at 1000 m3/s) to 42% (at 3500 m3/s) of the dissolved organicmatter; sulfonic acids ranged from 41% at 1000 m3/s to 17% at 3500 m3/s; a thirdcategory, chloro-organics, ranged from 12% at low flow to 5% at high flow. Theseare refractory, or nonbiodegradable, classes of organic matter. The significance of

FIG. 6.4 Organic matter in the Mississippi River at Cape Girardeau, 1969-1970. (FromUSGS Water Supply Paper 2156.)

this information is simply that each investigator has his or her own purpose instudying organic matter in water and selects the most practical categories to studyand the simplest methods of analysis; there is usually no purpose to identifying30 to 40 specific organic compounds in water—a rather costly procedure—ifrougji indexes, such as TOC or humic substances will suffice for the study.

Some organic materials are truly soluble, but much of it—certainly the humicmatter—is present in colloidal form and can generally be removed by coagula-tion. Alum coagulation at a pH of 5.5 to 6.0 typically reduces color to less than 5APHA units. Organic matter such as is found in domestic sewage often inhibitscalcium carbonate precipitation. If the natural color exceeds about 50, it must bepartially removed for lime soffening to occur. Organic matter may be removedby activated carbon treatment, widely practiced in municipal treatment plantswhen organic matter causes objectionable tastes or odors in the finished water.Generally these tastes and odors are produced by algae, each species having itscharacteristic odor or taste just as with land plants. Also like land plants, algaeproduce organic compounds which may be toxic if enough is ingested by fish oranimals.

Certain organic materials in water polluted by agricultural runoff (e.g., pesti-cide residues) or by industrial wastes in concentrations far below 1 mg/L still exerta significant effect on the biota of the receiving stream. Even when the effect isnot dramatic, as with fish kill, it may have long-term consequences, such as affect-ing reproduction or disrupting the food chain.

Organic matter is objectionable in municipal water chiefly for aesthetic rea-sons. It can be troublesome in industrial supplies by interfering with treatmentprocesses. It is a major factor in the fouling of anion exchange resins, degradingeffluent quality of demineralized water, and requiring early replacement of resin.

Organic <Nitrogen

ColorCOD

Silica (SiO2—Molecular Weight 60; Oxide of Silicon, Group IV, Nonmetal)

Silica is present in almost all minerals, and is found in fresh water in a range of1 to 100 mg/L. The skeletons of diatoms are pure silica, so the silica content ofsurface waters may be affected by seasonal diatom blooms. Silica is considered tobe colloidal because its reaction with adsorbents like MgO and Fe (OH)3 showcharacteristics similar to typical colloids. At high concentrations—over 50 mg/L—the adsorption isotherms (Chapter 3) no longer apply, and it appears thatchemical precipitation occurs instead. There is probably an equilibrium betweenthe silica in colloidal form and the bisilicate (HSiO3") anion. Because of this com-plexity, it is difficult to predict the conditions under which silica can be kept insolution as water concentrates during evaporation.

The term "colloidal silica" is loosely used by water chemists and can be con-fusing. Very little research has been done to categorize the size distribution of thesilica micelles (polymeric groups). It is very clear to the water analyst that thisneeds investigation. The analyst uses a colorimetric test that develops a blue colorto measure silica concentration. Sometimes, particularly with demineralizing sys-tems, there is evidence that some of the silica in water does not produce the bluecolor needed for detection, and this slips through the demineralizer without reac-tion. It seems that some of the silica micelles are too large to react with the chem-ical test reagents and the ion exchange resin, and in this case, the analyst mayreport that colloidal silica is present. A more accurate statement would be thatinert (or nonreactive) silica is present, since for all practical purposes all of thesilica is colloidal, although of differing sizes.

Silica is objectionable at high concentration in cooling tower makeup becauseof this uncertainty about its solubility limits.

It is objectionable in boiler feed water makeup not only because it may form ascale in the boiler itself, but also because it volatilizes at high temperatures andredeposits on turbine blades. Treatment processes that remove silica are: adsorp-tion on magnesium precipitates in the lime softening operation; adsorption onferric hydroxide in coagulation processes using iron salts; and anion exchange inthe demineralization process.

Sodium (Na+-Atomic Weight 23; Group IA, Alkali Metal)

All sodium salts are highly soluble in water, although certain complexes in min-erals are not. The high chloride content of brines and seawater is usually associ-ated with the sodium ion. In fresh waters, its range is usually 10 to 100 mg/L(about 20 to 200 mg/L as CaCO3). Sodium is present in certain types of clay andfeldspar. There is an increase of sodium in municipal sewage of 40 to 70 mg/L inexcess of the municipal water supply. Its concentration is not limited by FederalDrinking Water Standards, so persons on low sodium diets may require specialsources of potable water. The only chemical process for removing sodium iscation exchange in the hydrogen cycle. Evaporation and reverse osmosis alsoreduce sodium, producing a product stream low in sodium and a spent brine highin sodium.

Sulfate (SO42-—Molecular Weight 96; Oxide of Sulfur, Group VIA, Nonmetal)

Sulfate dissolves in water from certain minerals, especially gypsum, or appearsfrom the oxidation of sulfide minerals. Its typical range is 5 to 200 mg/L. The

suggested upper limit in potable water is 250 mg/L, based on taste and its poten-tial cathartic effect. Because calcium sulfate is relatively insoluble—less than 2000mg/L—sulfate may be objectionable in concentrating water high in calcium, as inan evaporative system. High sulfate levels may be reduced measurably by mas-sive lime or lime-aluminate treatment, or in rare cases by precipitation with bar-ium carbonate. It may also be reduced by anion exchange. In the coagulation ofwater with alum, sulfate is introduced at a rate of 1 mg/L SO4 for each 2 mg/Lalum added, while an equivalent amount of alkalinity is neutralized.

Total Dissolved Solids

Since this is the sum of all materials dissolved in the water, it has many mineralsources. Its usual range is 25 to 5000 mg/L. The suggested limit for public watersupplies, based on potability, is 500 mg/L. The principal effect of dissolved solidson industrial processes is to limit the extent to which a water can be concentratedbefore it must be discarded. High concentrations affect the taste of beverages. Therelated electrical conductivity tends to accelerate corrosion processes. A reductionin dissolved solids is achieved by a reduction in the individual components.

CLASS 2—SECONDARY CONSTITUENTS

These are generally present in concentrations greater than 0.1 mg/L and occa-sionally in the range of 1 to 10 mg/L.

Ammonia [NH3—Molecular Weight 17, Usually Expressed as N (Nitrogen); AtomicWeight 14, Group VA, Nonmetal]

Ammonia gas is extremely soluble in water, reacting with water to produceammonium hydroxide. Since this ionizes in water to form NH4

+ + OH~, at highpH, free ammonia gas is present in a nonionized form. At the pH of most watersupplies, ammonia is completely ionized.

NH3 + H2O — NH4OH ̂ NH4+ + OH~ (1)

(The addition of excess OH~ drives the reaction to the left.)Ammonia is one of the transient constituents in water, as it is part of the nitro-

gen cycle and is influenced by biological activity. As is seen in the illustration ofthe nitrogen cycle, Figure 6.5, ammonia is the natural product of decay of organicnitrogen compounds. These compounds first originate as plant protein matter,which may be transformed into animal protein. The return of this protein mate-rial to the environment through death of the organism or through waste elimi-nation produces the organic nitrogen compounds in the environment which thendecay to produce ammonia.

Because this biological process also occurs in sewage treatment plants, ammo-nia is a common constituent of municipal sewage plant effluent, in which its usualconcentration is 10 to 20 mg/L. It also finds its way into surface supplies from theagricultural runoff in areas where ammonia is applied to land as a fertilizer. Ani-

FIG. 6.5 The nitrogen cycle.

mal feed lots also contribute ammonia, which may run off into surface streamsor find its way into underground aquifers.

The effect of sewage plant effluent on the ammonia content of a receivingstream is shown in Figure 6.6. Ammonia is oxidized by bacterial action first tonitrite and then to nitrate, so the concentration is continually being affected bythe input from decay of organic nitrogen compounds and by the output, which isthe uptake by bacteria to convert ammonia to nitrate.

The typical concentration range in most surface supplies is from 0.1 to 1.0 mg/L, expressed as N. It is not usually present in well waters, having been convertedto nitrate by soil bacteria. Certain industrial discharges, such as coke plant wastes,are high in ammonia and account for the ammonia content of some surfacewaters.

The concentration of ammonia is not restricted by drinking water standards.Ammonia is corrosive to copper alloys, so it is of concern in cooling systems andin boiler feedwater.

Ammonia is often deliberately added as the nitrogen source for biologicalwaste treatment systems. This is because the bacteria require nitrogen to produce

NitrateNitrogen

(NOg)

(N+5)

OrganicNitrogen

PlantProteins

OrganicNitrogenAnimai

Proteins

OrganicNitrogen

Compounds

AmmoniaNitrogen

(NH3)

(N'3)

NitriteNitrogen

NOa(N*3)

Atmospheric N

Fixation By AlgaeSoil Bacteria

Denitrificationby BacterialAction

Decomposition

Death

ProteinSynthesis

or Anabolism NitrifyingBacteria

ElectricalDischargeI and I

PhotochemicalFixation

FIG. 6.6 The effect of sewage plant discharge on the Trinity Riverbelow Dallas, Texas, 1969-70. (From USGS Water Supply Paper 215 7.)

protein substances (Figure 6.5), so the nitrogen is usually applied at the ratio ofone part of nitrogen per 20 parts of food, measured as BOD.

Ammonia can be removed by degasification, by cation exchange on the hydro-gen cycle, and by adsorption by certain clays, such as clinoptilolite. It is alsoreduced in concentration by biological activity, as noted above.

Borate [B(OH)4-, Compound of Boron, Atomic Weight 10.8; Group IMA, Nonmetal]

Most of the world's boron is contained in seawater, at 5 mg/L B. Pure supplies ofsodium borate occur in arid regions where inland seas have evaporated to dry-ness, especially in volcanic areas. Boron is frequently present in freshwater sup-plies from these same geologic areas.

It is present in water as nonionized boric acid, B(OH)3. At high pH (over 10),most of it is present as the borate anion, B(OH)4

-. It has little known significancein water chemistry. Its concentration is not limited in municipal waters by potablewater standards. It can be damaging to citrus crops if present in irrigation waterand the irrigation methods tend to concentrate the material in the soil. Althoughboron is in the same group on the periodic chart as aluminum, it behaves morelike silica in aqueous systems; it can be removed by anion exchange and byadsorption.

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Fluoride (F'-Atomic Weight 19; Group VIIA, Halide)

Fluoride is a common constituent of many minerals, including apatite and mica.It is common practice to add fluoride to municipal water to provide a residual of1.5 to 2.5 mg/L, which is beneficial for the control of dental caries. Concentrationsabove approximately 5 mg/L are detrimental, however, usually causing mottled,brittle tooth structure. Because of this, the concentration is limited by drinkingwater standards. High concentrations are present in wastewaters from glass man-ufacture, steel manufacture, and foundry operations. Lime precipitation canreduce this to 10 to 20 mg/L. Fluoride is also reduced by anion exchange and byadsorption on calcium phosphate and magnesium hydroxide. Fluoride forms anumber of complexes, so residuals in fluoride wastes should be analyzed chemi-cally and not by the use of a fluoride electrode.

Iron (Fe2+ and Fe3+—Atomic Weight 55.9; Group VIII, Transition Element,Figures 6.1, 6.2, and 6.3)

Iron is found in many igneous rocks and in clay minerals. In the absence of oxy-gen, iron is quite soluble in the reduced state (as seen in the analysis of well waterscontaining iron). When oxidized in a pH range of 7 to 8.5, iron is almost com-pletely insoluble, and the concentration can be readily reduced to less than 0.3mg/L, the maximum set by drinking water standards. Because iron is so insolublewhen oxidized completely, the actual residual iron after treatment is determinedby how well the colloidal iron has been coagulated and filtered from the water.

Because iron is a product of corrosion in steel piping systems, often the ironfound in water from a distribution system is from this source and does not rep-resent iron left from the treatment process in the water treatment plant.

Nitrate [NO3-—Molecular Weight, 62, Usually Expressed as N (Nitrogen), Atomic

Weight 14; Group VA, Nonmetal]

Nitrate, like ammonia, comes into water via the nitrogen cycle, rather thanthrough dissolving minerals. Its concentration is limited by drinking water stan-dards to 45 mg/L for physiological reasons. There are no reported uses of waterwhere nitrate is a restrictive factor. There is an increase of total nitrogen in sewageplant effluent in the range of 20 to 40 mg/L as N above the level in the watersupply. A great deal of this is ammonia, but some is nitrate. The only chemicalprocess that removes nitrate is anion exchange; nitrate can be converted to nitro-gen in a biological system by the action of the nitrifying bacteria. The nitrate con-tent of well water is usually appreciably higher than surface water.

Potassium (K+—Atomic Weight 39.1; Group IA, Alkaline Metal)

Potassium is closely related to sodium—so much so that it is seldom analyzed asa separate constituent in water analysis. Its occurrence is less widespread innature, and for that reason it is found at lower concentrations than sodium. It hasno significance in public water supplies or in water used for industrial purposes.As with sodium it can be removed chemically only by cation exchange, or by

physical processes such as evaporation and reverse osmosis. Potassium salts arehighly soluble in water (Table 3.5); but as a common constituent of clays, potas-sium is kept from dissolving by the nature of the structure of clay. For that reason,when water-formed deposits contain significant levels of potassium, it is probablycaused by silt, in which case the deposit would also be high in Al2O3 and SiO2.

Strontium (Sr2+—Atomic Weight 87.6; Group HA, Alkaline Earth Metal, Figure6.1)

Strontium is in the same family as calcium and magnesium. Although its solu-bility in the presence of bicarbonate is significant (about half that of calcium), itsoccurrence is usually restricted to geologic formations where lead ores occur, andtherefore its concentration in water is typically quite low. It is completelyremoved by any process used for calcium removal. If not removed by softening,in a scaling water it will be a contributor to the scale problem.

CLASS 3—TERTIARY CONSTITUENTS

This group includes materials genrerally found at concentrations exceeding 0.01mg/L.

Aluminum (Al3+—Atomic Weight 27; Group HIA, Metal)

Although aluminum constitutes a high percentage of the earth's crust as a com-mon component of a wide variety of minerals and clays, its solubility in water isso low that it is seldom a cause for concern either in municipal supplies or indus-

EPMA/ /SEC/VOLT/CM

FIG. 6.7 Variation of particle charge with the nature of theparticle and pH. (From Stumm and Morgan, 1970.)

Alumina (Al9O^)

trial water systems. However, in industrial systems, the carryover of alum floefrom a clarifier may cause deposit problems, particularly in cooling systems wherephosphate may be applied as a stabilizing treatment. Aluminum found in treatedwater systems is usually there because of colloidal residues (alumina, Al2O3) fromthe coagulation of the water if alum or aluminate is used as the coagulant. If theresiduals are objectionable, they can be removed by improved filtration practices.

As shown by the solubility curves (Figure 6.3) aluminum is amphoteric, beingpresent as Al3+ or lower valence hydroxyl forms at low pH and the aluminateanion at higher pH values. As might be expected from this amphoteric nature,alumina particles are positively charged at low pH and negatively charged at highpH, as indicated by Figure 6.7. The effectiveness of alum in precipitating nega-tively charged colloids, such as clay particles from water, is more likely related tothe charge on the precipitated alumina than the charge on the aluminum ion itself,since the aluminum ion is not soluble in the typical coagulation pH range of 5 to7. Its strong negative charge at pH 10.0 to 10.5 helps explain the effectiveness ofsodium aluminate in precipitating magnesium hardness, which is positivelycharged at this pH.

Arsenic (As—Atomic Weight 74.9; Group VA, Nonmetal)

The solubility of arsenic in water is so low that its presence is usually an indicatorof either mining or metallurgical operations in the watershed or runoff from agri-cultural areas where arsenical materials have been used as industrial poisons. Ifin colloidal form, it would be removed by conventional water treatment pro-cesses. Federal regulations limit the content in public water supplies to a maxi-mum of 0.1 mg/L total arsenic. If the material is present in organic form, it maybe removed by oxidation of the organic material and subsequent coagulation, orby an adsorption process, such as passage through granular activated carbon.

Barium (Ba2+—Atomic Weight 137.3; Group UA, Alkaline Earth Metal)

In natural waters containing bicarbonate and sulfate, the solubility of barium isless than 0.1 mg/L, and it is seldom found at concentrations exceeding 0.05 mg/L. Removal to low residuals can be expected in conventional lime treatment pro-cesses. There are instances of barium being added to water for the specific purposeof sulfate reduction. The reaction is hindered because the barium reagent itself isso insoluble that considerable time is needed for the reactions to occur; further-more, sulfate deposition on the surface of the barium reagent makes the processinefficient. Barium is limited in drinking water to a maximum concentration of 1mg/L.

Bromide (Br —Atomic Weight 79.9; Group VIIA, Halide)

Bromine is found in seawater at about 65 mg/L as the bromide ion; some connatewaters produced with oil contain several hundred milligrams per liter and are thesource of commercial bromine. Over 0.05 mg/L in fresh water may indicate thepresence of industrial wastes, possibly from the use of bromo-organo compoundsas biocides or pesticides.

Copper (Cu2+—Atomic Weight 63.5; Group IB Metal)

Copper may be present in water from contact with copper-bearing minerals ormineral wastes from copper production. It is more likely, however, that the cop-per found in water will be a product of corrosion of copper or copper alloy pipingor fittings, or may have been added deliberately to a water supply reservoir foralgae control, as copper sulfate. When copper sulfate is added for algae control,because its solubility is limited, organic chelating materials may be added to thecopper sulfate formulation to keep the copper from precipitating and, therefore,maintain its effectiveness. Drinking water regulations limit the municipal watersupply concentration to 1 mg/L maximum. At higher concentrations, the waterhas an astringent taste. If a water supply is corrosive to copper, the first drawingor tapping of the supply from piping which has been idle overnight may containrelatively high concentrations, and ingestion of this water may cause immediatevomiting. In industrial supplies, the presence of copper can be objectionable as itis corrosive to aluminum. Copper is essential to certain aquatic organisms, beingpresent in hemocyanin in shellfish, the equivalent of hemoglobin in humans.

Lead (Pb2+—Atomic Weight 207.2; Group IV, Metal)

The presence of lead in fresh water usually indicates contamination from metal-lurgical wastes or from lead-containing industrial poisons, such as lead arsenate.However, lead may also appear in water as a result of corrosion of lead-bearingalloys, such as solder. Being amphoteric, lead is attacked in the presence of causticalkalinity.

The limitation on lead in drinking water has been established as 0.05 mg/L,which should be readily achieved with good filtration practice. In wastewaterswhere lead may be complexed with organic matter, it may be solubilized, andoxidation of the organic may be required for complete lead removal.

Lithium (Li+—Atomic Weight 6.9; Group IA Alkali Metal)

This alkaline earth element is rare in nature and seldom analyzed in water. Thereare no records of experience indicating that this material is of concern either inindustrial or municipal water supplies. However, lithium salts are used in psy-chotherapy to combat depression, so there may be a concentration level in waterthat has a psychotropic effect. Lithium salts have a wide variety of uses, but theindustrial consumption is so low that it is not likely to be a significant factor inthe wastewaters from industries using these products.

Manganese (Mn 2+, Mn4+—Atomic Weight 54.9; Group VIIB, Metal)

Manganese is present in many soils and sediments as well as in metamorphicrocks. In water free of oxygen, it is readily dissolved in the manganous (Mn2+)state and may be found in deep well waters at concentrations as high as 2 to 3mg/L. It is also found with iron in acid mine drainage. Wastewaters from metal-lurgical and mining operations frequently contain manganese.

It is an elusive material to deal with because of the great variety of complexesit can form depending on the oxidation state, pH, bicarbonate-carbonate-OHequilibria, and the presence of other materials, particularly iron.

It is limited to 0.05 mg/L maximum by drinking water regulations, becausehigher concentrations cause manganese deposits and staining of plumbing fixturesand clothing. However, concentrations even less than this can cause similareffects, as it may accumulate in the distribution system as a deposit, to be releasedin higher concentrations later if the environment should change, such as bychange in pH, CO2 content, oxidation potential, or alkalinity.

In industrial systems it is as objectionable as iron, particularly in textile man-ufacture or the manufacture of bleached pulp, since small amounts of depositedmanganese can slough off to cause stained products which must be rejected.Reduction to levels as low as 0.01 mg/L are required for certain textile finishingoperations.

In the oxidized state, manganese is quite insoluble, and can be lowered in con-centration, even in an alum coagulation process by superchlorination with ade-quate filtration, even at a pH as low as 6.5. However, the conventional processfor removal of manganese by itself is oxidation plus elevation of the pH toapproximately 9 to 9.5, with retention of approximately 30 min in a reaction ves-sel before filtration. Filters that have gained a coating of manganese oxides canwork very effectively, but may slough manganese if the aquatic environment isradically changed. Manganese is also precipitated by the continuous applicationof potassium permanganate ahead of a manganese form of zeolite.

Organic materials can chelate manganese much as they chelate iron, so

Manganese mg/l

Month

FIG. 6.8 Manganese concentration in the top and bottom layers of a lake, as affected byseasonal changes. (From "Chemistry of Manganese in Lake Mendota, Wisconsin," Envi-ron. Sc. Tech., December 1968.)

Concentra ion at Bottom 23 meters

Fall TurnoverSpring Turnover

Concentration at 10 meters

destruction of the organic matter is often a necessary part of the manganeseremoval process.

Because manganese accumulates in sediments, it is common to find high levelsof manganese in deep water where none may be apparent at the surface. Thisshould be studied in designing the proper intake structure for a plant water sup-ply. An example of this is illustrated in Figure 6.8, showing the manganese con-centrations in a lake at various times of the year and at different depths in thelake.

Phosphate (PO43-, Molecular Weight 95; Compound of Phosphorus P—Atomic

Weight 31)

Phosphorus is found in many common minerals such as apatite, in the form ofphosphate (PO4

3"equivalent weight 31.7). Since phosphate compounds are widelyused in fertilizers and detergents, it is common to find phosphate in silt fromagricultural runoff, with fairly high concentrations being found in municipalwastewater, usually in the range 15 to 30 mg/L as PO4 (about 5 to 10 mg/L P).Since phosphate is commonly blamed as the primary cause of excessive algalgrowths, which lead to eutrophication of lakes and streams, a reduction of phos-phate is being brought about by legislation restricting the amount of phosphate indetergents and also requiring treatment of municipal sewage for phosphateremoval.

Phosphate may be present in water as HPO42" and H2PO4

-, as well as thehigher pH form, PO4

3". The distribution as affected by pH is shown in Figure 6.9.Phosphate can be reduced to very low levels by treatment with alum, sodium

aluminate, or ferric chloride, with a formation of insoluble aluminum phosphateand iron phosphate. It can also be precipitated with lime at a pH over 10 to pro-

FIG. 6.9 The effect of pH on the distribution of various phosphatespecies.

MO

LE

FR

AC

TIO

N

duce residuals less than 2 to 3 mg/L in the form of hydroxyapatite; in a hot pro-cess system, the residuals would be less than 0.5 mg/L.

These phosphate precipitates are often colloidal, and filtration is required toachieve the low residuals specified.

Zinc (Zn2+—Atomic Weight 63.4; Group HB, Metal)

Zinc is a Group UB metal, behaving quite like calcium in solution, although ofconsiderably lower solubility in natural waters with a neutral pH and havingbicarbonate alkalinity. (See Figure 6.3 for its solubility characteristics.) Zinc isseldom found at concentrations over 1 mg/L, with a typical concentration beingapproximately 0.05 mg/L. Because it tends to have an astringent taste, its concen-tration in public water supplies is limited to 5 mg/L maximum.

Zinc may be present in water because of waste discharges from mining, met-allurgical, or metal finishing operations. It may also appear because of corrosionof galvanized steel piping. It is often included in proprietary corrosion inhibitorswhere its effect on steel piping is similar to that of galvanizing.

Zinc would be removed in lime softening operations to residuals well below0.1 mg/L. It can also be removed by cation exchange on either the sodium or thehydrogen cycle.

CLASS 4—TRACE CONSTITUENTS

Materials in this group (Table 6.3) are generally found at concentrations less than0.01 mg/L.

CLASS 5—TRANSIENT CONSTITUENTS

This class includes constituents which change in concentration or activity not bydilution, dissolution, or precipitation, but rather by changes in the aquatic envi-ronment which disturb the equilibrium. These changes may come about from bio-logical activity, oxidation-reduction potential, or radioactive decay.

Acidity-Alkalinity

The typical domain of almost all natural waters is characterized by a pH range of6 to 8, the presence of bicarbonate alkalinity, and some CO2 dissolved in thewater. All waters in contact with limestone, dolomite, or geologic formationsincluding these minerals tend to reach this equilibrium: it is the end result of thechemical reactions that cause the weathering of rocks and of the oxidation-reduc-tion reactions which are mediated by aquatic organisms. Because of this, the fewexceptional streams that contain free mineral acidity (i.e., have a pH below about4.5) usually dissipate this condition by accelerated weathering of the alkalinecomponents of the rocks they contact. Likewise, when the pH exceeds 8 and car-bonate alkalinity begins to appear, this is brought into balance by reaction withcarbon dioxide from the atmosphere or from respiration of aquatic life.

Constituent

Antimony(Sb, group VA,at. wt. 74.9)

Cadmium(Cd, GroupIIB, at. wt.112.4)

Chromium(Cr, groupVIB, at. wt.52)

Cobalt(Co, groupVIII, at. wt.59)

Cyanide(CN~, eq. wt.26)

Mercury(Hg, group IIB,at. wt. 200.6)

Nickel(Ni, groupVIII, at. wt.58.7)

Tin(Sn, groupIVA, at. wt.118.7)

Titanium(Ti, groupIVB, at. wt.47.9)

Occurrence*

Leaching of metallurgicalslags; colloidal hydrousoxides

Plating wastes; limitedsolubility as Cd2+.Note: Restricted inpotable water suppliesto 0.01 mg/Lmaximum.

Plating wastes, coolingtower blowdown.Soluble as CrO4

2~(Cr6+); insoluble asCr3+. Note: Restrictedin potable watersupplies to 0.05 mg/Lmaximum as Cr.

Present in copper- andnickel-bearing oretailings; ceramic wastes.

Wastes from plating shops,coke plants, blastfurnaces, petroleumrefining. Note:Restricted in potablewater supplies to 0.2mg/L maximum.

Wastes from electrolyticNaOH production,leaching of coal ash.Note: Restricted inpotable water suppliesto 0.002 mg/Lmaximum.

Plating wastes; electricfurnace slag or dust; oretailings.

Tinplate waste.

Ilmenite in wellformation.

Behaviorf

Insoluble in aqueous systemscontaining HCO3/CO3/OH.Filterable.

Behaves as Ca and Zn (seeFigure 6.3). Can beprecipitated as carbonate orremoved by cation exchange.

As CrO42+, oxidizing agent.

Can be reduced by SO2 toCr3+ or removed by anionexchange. As Cr3+, colloidalhydrous oxide at neutral pH,filterable.

Behaves as iron.

Behaves as Cl-, NO3", highlysoluble. Can be oxidizedwith Cl2 to CNO ~ and N2 +CO2. Reduced in activatedsludge biodigestion.Removable by anionexchange. Forms complexeswith Cd, Fe.

May be methylated by bacterialactivity and taken up by theaquatic food chain.Removed by closing plantloop or by reduction andfiltration.

Ni2+ behaves as iron, Fe2+.Seldom present as Ni3+.Precipitates as hydroxideand basic carbonate. Sulfideprecipitates are insoluble.Removable by cationexchange.

Converted to hydrous oxidecolloid at neutral pH or inHCO3/CO3 environment.

Colloidal TiO2. Removed byfiltration.

TABLE 6.3 Class 4—Trace Constituents

* Most constituents in class 4 are introduced by industrial waste discharges, and are more generallyfound in surface than in well waters.

Uncontrolled discharge of industrial wastes could wipe out this natural buff-ering effect of equilibrium between the aquatic environment, the atmosphere, andthe lithosphere.

Since legislation now prohibits such discharge, the only circumstances that cancause waters to fall outside the natural conditions are accidental spills of largevolumes of strong chemicals, seepage of acid mine drainage into a stream, or acidrain from air pollution. The mine seepage may not be controllable because ofinability to locate the source or the point of entry into the stream.

Acid rain is caused by the dissolution of acidic gases from the environment,chiefly the oxides of sulfur (SO2, SO3), and perhaps aggravated by nitrogen oxides(NOx). The most prominent source appears to be residues of sulfur from coal-firedboiler plants. When the acid rain falls on alkaline rock or into rivers or lakes inlimestone basins, there may be enough reserve alkalinity in the rock or dissolvedin the water to neutralize the acidity. But often the rain falls in forested areaswhere vegetative litter has developed a soil high in humus. There is no naturalalkalinity then to counteract the acidity of the rainfall, so the runoff is acidic andif the drainage leads to a lake in a granite basin (or a formation free of limestone),the lake itself will become acid and normal aquatic life will disappear.

Areas of the United States and Canada affected by acid rain are shown in Fig-ure 2.11. Possibly at some time in the past, alkaline components in industrial gasdischarges (e.g., alkaline fly ash, cement kiln, or lime kiln dust) alleviated someof this problem, and it is likely that application of limestone to some lakes affectedby rain will be both effective and economically justified. Reduction of sulfuroxides from boiler plant flue gas often utilizes lime as the neutralizing agent. Butcomplete elimination of these sulfur oxides is impossible in the boiler stack. Soalkali treatment of acid-affected lakes may prove more economical and practicalthan further treatment at the boiler stack as further addition of alkali there runsinto diminishing returns. However, this does not correct other deleterious effectsof acid rain, such as damage to the flora of forests and croplands and etching andcorrosion of buildings and other structures.

Carbon Cycle Constituents (Carbon, C—Atomic Weight 12.0; Group IVA,Nonmetal, see Figure 6.10)

Carbon is one of the primary elements of living matter, as shown by the gener-alized formula for biomass. It has been hypothesized that earth's primitive envi-ronment contained carbon in the form of methane plus ammonia, water, andhydrogen gases. Methane is one of the carbon compounds present in the carboncycle (Figure 6.10), produced by fermentation of larger organic molecules. Carbondioxide and bicarbonate-carbonate alkalinity are also prominent in the cycle.These reactions are proceeding in the aquatic environment in the carbon cycle,Figure 6.10.

Methane, a major component of natural gas, is produced by anaerobic decom-position of organic chemical compounds. Methane is given off by anaerobicdecomposition of organic sediments in marsh gases, and the concentration maybecome so high that the swamp gas may ignite. It is more common to find meth-ane in well waters in areas where natural gas is produced than in surface waters.

Oxygen Cycle Constituents

The most common carbon-containing gas is carbon dioxide, discussed in earlierchapters. The carbon dioxide content of surface waters is greatly influenced by

FIG 6.10 The carbon cycle.

bacterial and algal symbiotic existence, illustrated by the oxygen cycle, Figure6.11. During bright sunlight, the photosynthetic reactions proceed so rapidly thatthe water may actually become supersaturated with oxygen, beyond the capacityof the bacteria to utilize. If algae require more carbon dioxide than is availablefrom bacterial respiration, they assimilate carbon dioxide from the bicarbonatealkalinity, producing a trace of carbonate alkalinity. For that reason, carbon diox-ide and oxygen are variable in most surface supplies, as affected by sunlight andthe photosynthetic process (Figure 6.12).

The plant tissue built up by photosynthesis is eventually metabolized by largeraquatic organisms that produce organic compounds and discharge them in theirwastes. Organic compounds are also produced by the death and decay of theaquatic plants and fish life. Organic matter thus produced becomes food for bac-teria and is returned to the cycle as methane by anaerobes and as carbon dioxideby aerobic bacteria.

Deep well water often contains over 25 mg/L CO2, and may be saturated withthis gas at the hydrostatic pressure and temperature in the water table. A drop inpressure as water flows across the well screen may cause the CO2 to come out ofsolution, disrupting the equilibrium and depositing CaCO3 scale.

The theoretical solubility of oxygen exposed to the atmosphere is dependent

CarbonC°

IncompleteCombustion

CarbonMonoxide

COC+2

Oxidation

CarbonDioxideCO2

C+4

Oxidation

IonicDomain

HCOg -CO*

Carbohydrates(CH2O),;

PlantMetabolism

Digestion

AnimalMetabolism

Decayand Death Aerobic

Digestion

OrganicCarbon

CompoundsC+4

MethaneCH4

Direct Photosynthesis

By Chlorophyll - Containing Plants

on the temperature (Figure 6.13). Excess oxygen concentration is due to photo-synthesis, and a deficiency is usually caused by bacterial activity or reducingagents.

Nitrogen Cycle Constituents (Figure 6.5)

Constituents of the nitrogen cycle have been discussed earlier. As was true withthe carbon cycle, the nitrogen cycle is involved with life in the aquaticenvironment.

ALGAL

PHOTOSYNTHESIS

BACTERIAL

RESPIRATION

ATMOSPHERE

O2 CO2

CO2

DISSOLUTION

IN WATER

FIG. 6.11 The oxygen cycle.

FIG. 6.12 Diurnal variation of oxygen andcarbon dioxide in a surface water.

CO2 CONCENTRATION

O2 CONCENTRATION

TEMPERATURE

FIG. 6.13 Solubilities of oxygen and nitrogen inwater at various temperatures.

Sulfur Cycle (Figure 6.14)

Because sulfur is in the same family as oxygen, there are many compounds wheresulfur replaces oxygen in a compound with similar properties. For example,ethanol (CH3CH2OH) and ethyl mercaptan (CH3CH2SH) are analogous com-pounds with only the sulfur and oxygen atoms interchanged.

Certain bacteria can metabolize the sulfur atom in hydrogen sulfide, just asalgae and other plants can metabolize oxygen from water in photosynthesis toproduce free oxygen and carbohydrate. The by-product of the bacterial process ofsplitting H2S is free sulfur. The corresponding chemical equations are:

Algalphotosynthesis

CO2 + 2 H2O > CH2O + O 2 -H H2O (2)

Bacterialaction

CO2 + 2 H2S —> CH2O + 2 S +H2O (3)

Hydrogen sulfide, which is present in some deep well waters and some stagnantsurface waters, is generally produced by the anaerobic decomposition of organic

NITROGEN

OXYGENGA

SS

OLU

BIL

ITY c

c/l

FIG. 6.14 The sulfur cycle.

compounds containing sulfur or by sulfate-reducing bacteria capable of convert-ing sulfate to sulfide.*

All of these biological processes that occur in nature can be put to work undercontrolled conditions by the water specialist, to digest and eliminate undesirableorganic wastes or their by-products.

* In each of these cycles involving biological activity, the changes occurring in surface water are muchmore pronounced than those occurring in deep well waters. The surface supplies are constantly inocu-lated by microbes from the air and from the soil and supplied with solar energy. On the other hand, wellwater usually represents a dead end in the cycle, with the water generally having a long residence time inthe aquifer, with the likelihood that most organisms have been filtered by the porous formation so thatall of the microbes may have been removed from the strata at the well screen. Because of this, it iscommon to find that the constituents of the carbon cycle, particularly CO2, those of the nitrogen cycle,particularly NOj", and those of the sulfur cycle, particularly SO42~ and HS ~, are relatively constant inwell waters. For the same reason, the concentrations OfCO2, NO3", and HS ~ are generally higher indeep well water than in surface supplies.

Sulfateso;2(S+6)

Sulfiteso'2

(S+4)

H2S%HS"1-S"2

(S2)

PlantLife

Processes

AnimalLife

Processes

OrganicSulfur

Compoundss-2

Native Sulfur (S0)Decomposition

(Death)I

Digestion

CombustionOxidation

Reduction(Sulfur

DepositingBacteria

PartialCombustion

Oxidation

BacterialReduction

Oxidation-Reduction Potential

Some materials in water are transient because they tend to oxidize or reduce otherconstituents, either through biological activity, as illustrated by the several bio-logical cycles, or directly. The presence of these materials has already been con-sidered in the text covering the biological cycles and the constituents taking partin these cycles.

The only significant oxidizing materials encountered in the natural environ-ment that can participate in chemical reactions without needing a biological routeare oxygen, which is prevalent in surface supplies, and sulfur, which is encoun-tered by subsurface waters in contact with native sulfur, a special situation. Thereare numerous oxidizing materials that can appear as residues in treated waste-waters—among these are free chlorine and chromate.

The common reducing materials are organic material, ferrous iron (Fe2+),manganese (Mn2+), and bisulfide (HS") from the natural environment; reducingmaterials which may be added as treatment by-products or waste residues includea wide variety of organic matter, ferrous iron from such operations as steel pick-ling, and sulfite, present in certain kinds of pulp mill wastes.

Radionuclides

Water itself is not radioactive, but may contain elements that are. These enter thewater cycle as wastes from nuclear power plants, fallout from nuclear blasts, orthe by-products of metallurgical processing of radioactive materials. In very rarecases, well waters may contain radionuclides—common radioactive elements andtheir isotopes—as natural contaminants.

The water chemist normally deals with concentrations at part per million ormilligrams per liter levels. One mole of any substance represents its molecularweight, so that the concentration of 1 mole/L as calcium carbonate represents 100g/L as CaCO3. Since a mole contains 6 X 1023 molecules, 100 g/L contains thisnumber of molecules, and 100 mg/L contains 6 X 1020 molecules. Even at thevery low concentration of 1 ppb (0.001 mg/L), then, there are still 6 X 1015 mol-ecules of the substance as calcium carbonate in a liter of water. Concentrations ofradioactive substances at this level would be extremely dangerous.

In chemical change, one element reacts with another in a process involving theelectrons surrounding the nucleus, with the nucleus remaining unchanged. Energymay be given off or may be absorbed in the form of heat. With nuclear reactions,on the other hand, only the nucleus is affected, and the products of reaction mayinclude nuclear particles (such as protons, neutrons, or electrons) and energy,including heat and electromagnetic radiation. The typical by-products of nuclearreactions, then, are alpha particles, electrons, and electromagnetic radiations. Thegeneral characteristics of these emanations are shown in Table 6.4.

TABLE 6.4 Types of Radiation

Alpha

Beta

Gamma

Helium nuclei, charge + 2,mass 4

Electrons, charge — 1, mass O

Similar to x-rays

Can penetrate air, but is stopped bysolids.

About 1000 times more penetrating thanalpha radiation. Can penetrate 2-3mm into solids.

Can penetrate air or solids about tentimes deeper than beta radiation.

Some nuclear reactions are extremely rapid, occurring in a split second, butothers may require 1000 years. Because these reactions may never go to comple-tion, the life of a material participating in a nuclear reaction (called a radionu-clide) is expressed as the time required for one-half its energy to be dissipated, itshalf-life. Each nuclear disintegration can be measured by the particles given off.The unit of measurement is the curie, defined as 3.7 X 1010 disintegrations persecond. The levels of radioactive disintegration of concern to the water technol-ogist are so far below this that they are expressed as micromicrocuries, or pico-curies (pCi), a pCi being equal to 2.2 disintegrations per minute.

A variety of radionuclides may be found in waters leaving a nuclear powerplant. The government license to operate such a plant specifies the particularradionuclides which must be identified individually. This requires sophisticatedand painstaking analytical techniques. In municipal supplies, it is not requiredthat individual species be identified because the radiation levels are so low, butan analysis is required for alpha radiation, which is assumed to be contributed by226Ra, beta radiation (assumed to be contributed by strontium, 90Sr) and gross betaactivity. Table 6.5 shows the relative relationship between occupational levels and

TABLE 6.5 Radionuclide Contamination in Water

potable water levels and also relates the radiation intensity to the concentrationof 226Ra or 90Sr in the water. It is obvious that the concentration levels are farbelow those considered significant when dealing with nonradioactive contami-nants in water supplies.

The heavier radionuclides are generally insoluble and may be removed bycoagulation and filtration; the soluble constituents may be removed by ionexchange. These processes become quite complicated when the radionuclides arepresent in very low concentrations and the treatment process may first have toreact with more common contamination, such as calcium and magnesium, atconcentrations far above those of the radionuclides before the radionuclides arereduced to acceptable levels.

METHODS OF ANAL YSIS

It is not in the scope of this handbook to include methods of analysis for each ofthe constituents presented in this chapter. These methods are covered in manualsdevoted exclusively to such analyses (e.g., Standard Methods for Analysis ofWater and Wastewater, ASTM Standards). However, included in Chapter 7 is asummary of the methods of analysis used, detection limits, methods of samplepreservation, and other facts of importance.

Occupational levels(body exposure)pCi/Lppb

Potable levels (ingestion)pCi/Lppb

226Ra

4 X 101

4 X IQ-5

33 X ICT6

90Sr

8 X 102

4 X 10~6

101

5 X 10-8

SUGGESTED READING

Davies, S. N., and DeWiest, R. C. M.: Hydrogeology, Wiley, New York, 1966.Faust, S. J., and Hunter, J. V.: Organic Compounds in the Aquatic Environment, Dekker,

New York, 1961.McKee, J. E., and Wolf, H. W.: Water Quality Criteria, State Water Quality Control Board,

Resources Agency of California, 1963.Stumm, W., and Morgan, J. J.: Aquatic Chemistry, Wiley, New York, 1970.Todd, Dand K.: The Water Encyclopedia, Water Information Center, Port Washington,N. Y., 1970.

U.S. Environmental Protection Agency: National Water Quality Inventory, Report to Con-gress, EPA-440/9-75-014, 1975.

U.S. Environmental Protection Agency: Quality Criteria for Water, EPA-440/9-76-023,1976.

U.S. Environmental Protection Agency: Recommended Uniform Effluent Concentration,U.S. Government Printing Office, 1973.

William, S. L.: "Sources and Distribution of Trace Metals in Aquatic Environments," inAqueous-Environmental Chemistry of Metals (Ruben, J., ed.), Ann Arbor Science, AnnArbor, Mich., 1976.