CHAPTER 6
WATER CONTAMINANTS:
OCCURRENCE AND TREATMENT
The introduction of contaminants into water supplies has been shown to be
related to rainfall, the geologic nature of the watershed or underground aquifer,
and the activities of nature and the human population. Water contaminants to be
examined in more detail fall in two categories: dissolved matter (Table 6.1) and
nonsoluble constituents (Table 6.2). Dissolved gases are included in discussions
of the biological cycles affecting water quality.
As shown in Table 6.1, soluble materials in water are arbitrarily assigned to
five classifications, the first four of which are based on concentration levels, with
the last covering those materials usually transient because continuing reactions in
the aquatic environment change their concentrations.
Many materials are transient because of biological activity. The change in CO2
and 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 nitrogen
cycle. (See Class 2, Secondary Constituents.)
There are also longer term processes by which nature cycles matter through
living organisms, which in turn modify the environment and leave their records
in the rocks. This chapter examines the sources of contaminants in water, many
of which are minerals created by living things. Perhaps the best known are the
chalk 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 built
these great masses of land out of nothing; but, of course, it cannot do that any
more than man can. It has taken calcium compounds from very dilute solutions
of sea water and built up a shell of calcium compounds to protect itself. In this
process of following its preordained metabolic rite, it has concentrated calcium
by several thousandfold in the form of an insoluble compound. Insignificant as
the coral polyp may appear, it is one of the most important creatures in changing
the character of the earth's surface."
Many other natural cycles have been operating over countless geologic ages to
produce deposits of sulfur, iron, manganese, silica, and phosphate, to name only
a few. In the village of Batsto, New Jersey, the early American colonists set up
the first blast furnace in the New World. Their source of iron ore was "bog
iron"—pure iron oxide precipitated from artesian water by iron-depositing
bacteria.
So the presence of many of the mineral constituents in water supplies may
simply represent the return to the aquatic environment of a loan made by the
earth to living organisms long ago.
Class 1
Primary constituents—generally over 5 mg/L
Bicarbonate
Magnesium
Calcium
Organic matter
Chloride
Silica
Class 2
Secondary constituents—generally over 0.1 mg/L
Ammonia
Iron
Borate
Nitrate
Fluoride
Class 3
Tertiary constituents—generally over 0.01 mg/L
Aluminum
Copper
Arsenic
Lead
Barium
Lithium
Bromide
Manganese
Class 4
Class 5
Trace constituents—generally less than 0.01 mg/L
Antimony
Cobalt
Cadmium
Mercury
Chromium
Nickel
Transient constituents
Acidity-alkalinity
Biological cycles
Carbon cycle constituents
Organic C/CH4/CO/CO2/(CH2OyC-tissue
Oxygen cycle
O2/CO2
Nitrogen cycle constituents
Organic N/NHa/NO2-/NO3-/№/amino acids
Sulfur cycle constituents
Organic S/HS-/SO32-/SO42-/S°
Redox reactions
Oxidizing materials
From the natural environment—O2, S
2
Treatment residues—Cl2, CrO4 "
Reducing materials
From the natural environment—
Organics, Fe2+, Mn2+, HS~
Treatment residues—
2+
2
Organics, Fe , SO2, SO3 '
Radionuclides
TABLE 6.2 Nonsoluble Constituents in
Water Supplies
Class 1—Solids
Floating
Settleable
Suspended
Class 2—Microbial organisms
Algae
Bacteria
Fungi
Viruses
Sodium
Sulfate
Total dissolved solids
Potassium
Strontium
Phosphate
Zinc
Tin
Titanium
FIG. 6.1 Theoretical solubilities of carbonate compounds in a water system closed to an external
CO2 environment at 250C. (From Stumm and Morgan, 1970.)
Concentrotion, millimoles/liter
FIG. 6.2 Theoretical solubilites of oxides and hydroxides in water at 250C.
Stumm and Morgan, 1970.)
(From
As an aid to appreciating the solubilities of the constituents being examined,
and thus the limitations of their occurrence in natural water supplies and the
residuals which may be reached in precipitation processes, Figures 6.1, 6.2 and
6.3 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 several orders of magnitude above this level.
Bicarbonate (HCO3"—Molecular Weight 61)
The bicarbonate ion is the principal alkaline constituent of almost all water supplies. It is generally found in the range of 5 to 500 mg/L, as CaCO3. Its introduc-
Concentration, millimoles/liter
FIG. 6.3 Theoretical solubilities of oxides and hydroxides in water at 250C, showing
amphoteric nature of aluminum and zinc. (From Stumm and Morgan, 1970.)
tion into the water by the dissolving action of bacterially produced CO2 on carbonate-containing minerals has been explained elsewhere. Normal activities of
the human population also introduce alkaline materials into water, evidenced by
a typical increase of alkalinity of sewage plant effluent of 100 to 150 mg/L above
the alkalinity of the municipal water supply. Much of this is due to the alkalinity
of industrial and domestic detergents.
Alkalinity in drinking water supplies seldom exceeds 300 mg/L. The control
of alkalinity is important in many industrial applications because of its significance in the calcium carbonate stability index. Alkalinity control is important in
both concentrated boiler water and cooling water in evaporative cooling systems.
Makeup for these systems must often be treated for alkalinity reduction either by
lime softening or direct acid addition. Alkalinity is objectionable in certain other
industries, such as the beverage industry, where it neutralizes the acidity of fruit
flavors, and in textile operations, where it interferes with acid dyeing.
Calcium (Ca2+—Atomic Weight 40; Group 11 A, Alkaline Earth Metal, Figures
6.1 and 6.3)
Calcium is the major component of hardness in water and usually is in the range
of 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 the
fossils of tiny aquatic organisms, such as polyps, that have taken calcium from
the seawater in which they lived, and used it for their skeletons. This is but one
of many cycles in nature whereby some component of the environment is continually withdrawn by living things and eventually returned directly or indirectly.
Calcium removed from water in softening operations is later returned to the
environment, often to the watershed, by way of a precipitate or a brine which is
the by-product of the softening reaction. Calcium is a major factor in determining
stability index. Calcium reduction is often required in treating cooling tower
makeup. Complete removal is required for many industrial operations, particularly for boiler makeup, textile finishing operations, and cleaning and rinsing in
metal finishing operations.
Calcium hardness can be reduced to a level of 35 mg/L as CaCO3 by cold limesoda softening and to less than 25 mg/L by hot lime-soda softening. It is reduced
to 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 in
freshwater supplies, ranging from 10 to 100 mg/L. Seawater contains over 30,000
mg/L as NaCl, and certain underground brine wells may actually be saturated,
approximately 25% NaCl. Many geologic formations were once sedimentary
rocks in the sea, so it is not surprising that they contain residues of chlorides that
are continually leaching into freshwater sources. The chloride content of sewage
is 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 proceed
from the headwaters to the sea.
Anion exchange is the only chemical process capable of removing chlorides
from water; however, physical processes such as evaporation and reverse osmosis
can separate a feedwater into two streams, one with a reduced chloride and the
other 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 hardness, 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, magnesium concentration is about 5 times that of calcium on an equivalent basis. The
production of magnesium hydroxide from seawater is the starting point in the
manufacture of magnesium. Magnesium is a prominent component of many minerals, including dolomite, magnesite, and numerous varieties of clay.
Since magnesium carbonate is appreciably more soluble than calcium carbonate, it is seldom a major component in scale except in seawater evaporators. However, it must be removed along with calcium where soft water is required for
boiler makeup or for process applications. It may be removed by lime softening
to a residual of 30 to 50 mg/L as CaCO3 cold, or 1 to 2 mg/L as CaCO3 hot. It is
also 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 is
used by aquatic organisms to build their bodies and produce food, it is inevitable
that water-soluble organic products of metabolism should be present in all water
supplies. There is not much information available on specific organic compounds
in most water sources. (See Table 4.7.) There are literally hundreds of thousands
of known organic compounds, many of which might somehow find their way into
the hydrologic cycle. A complete "organic analysis" of water is impossible. However, one of the by-products of space-age technology has been the development of
new instruments for organic analysis (see Chapter 7). With these instruments, the
analyst can develop methods of analysis for organic materials of interest—especially those considered by the EPA to be toxic or carcinogenic, such as PCBs
(polychlorinated biphenyls) and TTHMs (total trihalomethanes). But unless such
specific organic compounds are requested at the time a sample is presented, the
analyst uses indirect measures of organics instead (e.g., COD, TOC).
Many waters have a yellowish or tea color due to decayed vegetation leached
from the watershed by runoff. These organic materials are broadly classified as
humic substances, further categorized as humic acid (a water-soluble compound),
fulvic acid (alkali-soluble material), and humin (high molecular weight, water
insoluble matter). These organic compounds are molecules having many functional groups containing oxygen and hydrogen atoms in various proportions, so
that 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 2.0
to 2.5 times greater than the value reported as carbon. A survey of 80 municipal
supplies in the United States showed an average total organic carbon content in
the finished water of 2.2 mg/L, as C, so the organic matter was probably on the
order 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 sampling station on the Mississippi River as determined by these indexes is shown in
Figure 6.4. Because some of the functional groups in humic compounds have ion
exchange properties, they tend to chelate heavy metals. In spite of this, there is
no correlation whatever between the color of a water and its total heavy metal
concentration. A study of the Rhine River showed that humic substances comprised from 25% (at 1000 m3/s) to 42% (at 3500 m3/s) of the dissolved organic
matter; sulfonic acids ranged from 41% at 1000 m3/s to 17% at 3500 m3/s; a third
category, chloro-organics, ranged from 12% at low flow to 5% at high flow. These
are refractory, or nonbiodegradable, classes of organic matter. The significance of
Organic <
Nitrogen
Color
COD
FIG. 6.4 Organic matter in the Mississippi River at Cape Girardeau, 1969-1970.
USGS Water Supply Paper 2156.)
(From
this information is simply that each investigator has his or her own purpose in
studying organic matter in water and selects the most practical categories to study
and the simplest methods of analysis; there is usually no purpose to identifying
30 to 40 specific organic compounds in water—a rather costly procedure—if
rougji indexes, such as TOC or humic substances will suffice for the study.
Some organic materials are truly soluble, but much of it—certainly the humic
matter—is present in colloidal form and can generally be removed by coagulation. Alum coagulation at a pH of 5.5 to 6.0 typically reduces color to less than 5
APHA units. Organic matter such as is found in domestic sewage often inhibits
calcium carbonate precipitation. If the natural color exceeds about 50, it must be
partially removed for lime soffening to occur. Organic matter may be removed
by activated carbon treatment, widely practiced in municipal treatment plants
when organic matter causes objectionable tastes or odors in the finished water.
Generally these tastes and odors are produced by algae, each species having its
characteristic odor or taste just as with land plants. Also like land plants, algae
produce organic compounds which may be toxic if enough is ingested by fish or
animals.
Certain organic materials in water polluted by agricultural runoff (e.g., pesticide residues) or by industrial wastes in concentrations far below 1 mg/L still exert
a significant effect on the biota of the receiving stream. Even when the effect is
not dramatic, as with fish kill, it may have long-term consequences, such as affecting reproduction or disrupting the food chain.
Organic matter is objectionable in municipal water chiefly for aesthetic reasons. It can be troublesome in industrial supplies by interfering with treatment
processes. It is a major factor in the fouling of anion exchange resins, degrading
effluent quality of demineralized water, and requiring early replacement of resin.
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 of
1 to 100 mg/L. The skeletons of diatoms are pure silica, so the silica content of
surface waters may be affected by seasonal diatom blooms. Silica is considered to
be colloidal because its reaction with adsorbents like MgO and Fe (OH)3 show
characteristics similar to typical colloids. At high concentrations—over 50 mg/
L—the adsorption isotherms (Chapter 3) no longer apply, and it appears that
chemical precipitation occurs instead. There is probably an equilibrium between
the silica in colloidal form and the bisilicate (HSiO3") anion. Because of this complexity, it is difficult to predict the conditions under which silica can be kept in
solution as water concentrates during evaporation.
The term "colloidal silica" is loosely used by water chemists and can be confusing. Very little research has been done to categorize the size distribution of the
silica micelles (polymeric groups). It is very clear to the water analyst that this
needs investigation. The analyst uses a colorimetric test that develops a blue color
to measure silica concentration. Sometimes, particularly with demineralizing systems, there is evidence that some of the silica in water does not produce the blue
color needed for detection, and this slips through the demineralizer without reaction. It seems that some of the silica micelles are too large to react with the chemical test reagents and the ion exchange resin, and in this case, the analyst may
report that colloidal silica is present. A more accurate statement would be that
inert (or nonreactive) silica is present, since for all practical purposes all of the
silica is colloidal, although of differing sizes.
Silica is objectionable at high concentration in cooling tower makeup because
of this uncertainty about its solubility limits.
It is objectionable in boiler feed water makeup not only because it may form a
scale in the boiler itself, but also because it volatilizes at high temperatures and
redeposits on turbine blades. Treatment processes that remove silica are: adsorption on magnesium precipitates in the lime softening operation; adsorption on
ferric hydroxide in coagulation processes using iron salts; and anion exchange in
the demineralization process.
Sodium (Na+-Atomic Weight 23; Group IA, Alkali Metal)
All sodium salts are highly soluble in water, although certain complexes in minerals are not. The high chloride content of brines and seawater is usually associated 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 and
feldspar. There is an increase of sodium in municipal sewage of 40 to 70 mg/L in
excess of the municipal water supply. Its concentration is not limited by Federal
Drinking Water Standards, so persons on low sodium diets may require special
sources of potable water. The only chemical process for removing sodium is
cation exchange in the hydrogen cycle. Evaporation and reverse osmosis also
reduce sodium, producing a product stream low in sodium and a spent brine high
in sodium.
Sulfate (SO42-—Molecular Weight 96; Oxide of Sulfur, Group VIA, Nonmetal)
Sulfate dissolves in water from certain minerals, especially gypsum, or appears
from 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 potential cathartic effect. Because calcium sulfate is relatively insoluble—less than 2000
mg/L—sulfate may be objectionable in concentrating water high in calcium, as in
an evaporative system. High sulfate levels may be reduced measurably by massive lime or lime-aluminate treatment, or in rare cases by precipitation with barium carbonate. It may also be reduced by anion exchange. In the coagulation of
water with alum, sulfate is introduced at a rate of 1 mg/L SO4 for each 2 mg/L
alum 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 mineral
sources. Its usual range is 25 to 5000 mg/L. The suggested limit for public water
supplies, based on potability, is 500 mg/L. The principal effect of dissolved solids
on industrial processes is to limit the extent to which a water can be concentrated
before it must be discarded. High concentrations affect the taste of beverages. The
related electrical conductivity tends to accelerate corrosion processes. A reduction
in 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 occasionally in the range of 1 to 10 mg/L.
Ammonia [NH3—Molecular Weight 17, Usually Expressed as N (Nitrogen); Atomic
Weight 14, Group VA, Nonmetal]
Ammonia gas is extremely soluble in water, reacting with water to produce
ammonium hydroxide. Since this ionizes in water to form NH4+ + OH~, at high
pH, free ammonia gas is present in a nonionized form. At the pH of most water
supplies, 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 nitrogen cycle and is influenced by biological activity. As is seen in the illustration of
the nitrogen cycle, Figure 6.5, ammonia is the natural product of decay of organic
nitrogen compounds. These compounds first originate as plant protein matter,
which may be transformed into animal protein. The return of this protein material to the environment through death of the organism or through waste elimination produces the organic nitrogen compounds in the environment which then
decay to produce ammonia.
Because this biological process also occurs in sewage treatment plants, ammonia is a common constituent of municipal sewage plant effluent, in which its usual
concentration is 10 to 20 mg/L. It also finds its way into surface supplies from the
agricultural runoff in areas where ammonia is applied to land as a fertilizer. Ani-
Organic
Nitrogen
Compounds
Ammonia
Nitrogen
(NH3)
Organic
Nitrogen
Animai
Proteins
(N'3)
Decomposition
Death
Protein
Synthesis
or Anabolism
Organic
Nitrogen
Plant
Proteins
Atmospheric N
Fixation By Algae
Soil Bacteria
Nitrifying
Bacteria
Electrical
Discharge
I and I
Photochemical
Fixation
Denitrification
by Bacterial
Action
Nitrite
Nitrogen
NOa
(N*3)
Nitrate
Nitrogen
(NOg)
(N+5)
FIG. 6.5 The nitrogen cycle.
mal feed lots also contribute ammonia, which may run off into surface streams
or find its way into underground aquifers.
The effect of sewage plant effluent on the ammonia content of a receiving
stream is shown in Figure 6.6. Ammonia is oxidized by bacterial action first to
nitrite and then to nitrate, so the concentration is continually being affected by
the input from decay of organic nitrogen compounds and by the output, which is
the 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 converted
to nitrate by soil bacteria. Certain industrial discharges, such as coke plant wastes,
are high in ammonia and account for the ammonia content of some surface
waters.
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 and
in boiler feedwater.
Ammonia is often deliberately added as the nitrogen source for biological
waste treatment systems. This is because the bacteria require nitrogen to produce
O
10,000
8000
Ammonia - PPM as N ®
1.0 2.0 3.0 4.0
5.0
Total Phosphorus as P-PPM® and Dissolved O2-PPM*
Flow CFS x
4000
6000
2000
O
FIG. 6.6 The effect of sewage plant discharge on the Trinity River
below 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 of
one part of nitrogen per 20 parts of food, measured as BOD.
Ammonia can be removed by degasification, by cation exchange on the hydrogen cycle, and by adsorption by certain clays, such as clinoptilolite. It is also
reduced 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 of
sodium borate occur in arid regions where inland seas have evaporated to dryness, especially in volcanic areas. Boron is frequently present in freshwater supplies 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 significance
in water chemistry. Its concentration is not limited in municipal waters by potable
water standards. It can be damaging to citrus crops if present in irrigation water
and the irrigation methods tend to concentrate the material in the soil. Although
boron is in the same group on the periodic chart as aluminum, it behaves more
like silica in aqueous systems; it can be removed by anion exchange and by
adsorption.
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 of
1.5 to 2.5 mg/L, which is beneficial for the control of dental caries. Concentrations
above approximately 5 mg/L are detrimental, however, usually causing mottled,
brittle tooth structure. Because of this, the concentration is limited by drinking
water standards. High concentrations are present in wastewaters from glass manufacture, steel manufacture, and foundry operations. Lime precipitation can
reduce this to 10 to 20 mg/L. Fluoride is also reduced by anion exchange and by
adsorption on calcium phosphate and magnesium hydroxide. Fluoride forms a
number of complexes, so residuals in fluoride wastes should be analyzed chemically 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 oxygen, iron is quite soluble in the reduced state (as seen in the analysis of well waters
containing iron). When oxidized in a pH range of 7 to 8.5, iron is almost completely insoluble, and the concentration can be readily reduced to less than 0.3
mg/L, the maximum set by drinking water standards. Because iron is so insoluble
when oxidized completely, the actual residual iron after treatment is determined
by 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 iron
found in water from a distribution system is from this source and does not represent 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 than
through dissolving minerals. Its concentration is limited by drinking water standards to 45 mg/L for physiological reasons. There are no reported uses of water
where nitrate is a restrictive factor. There is an increase of total nitrogen in sewage
plant effluent in the range of 20 to 40 mg/L as N above the level in the water
supply. A great deal of this is ammonia, but some is nitrate. The only chemical
process that removes nitrate is anion exchange; nitrate can be converted to nitrogen in a biological system by the action of the nitrifying bacteria. The nitrate content 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 as
a separate constituent in water analysis. Its occurrence is less widespread in
nature, and for that reason it is found at lower concentrations than sodium. It has
no 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 are
highly soluble in water (Table 3.5); but as a common constituent of clays, potassium 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 probably
caused 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, Figure
6.1)
Strontium is in the same family as calcium and magnesium. Although its solubility in the presence of bicarbonate is significant (about half that of calcium), its
occurrence is usually restricted to geologic formations where lead ores occur, and
therefore its concentration in water is typically quite low. It is completely
removed 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 0.01
mg/L.
Aluminum (Al3+—Atomic Weight 27; Group HIA, Metal)
Although aluminum constitutes a high percentage of the earth's crust as a common component of a wide variety of minerals and clays, its solubility in water is
so low that it is seldom a cause for concern either in municipal supplies or indusEPM
A/ /SEC/VOLT/CM
Alumina (Al 9 O^)
FIG. 6.7 Variation of particle charge with the nature of the
particle and pH. (From Stumm and Morgan, 1970.)
trial water systems. However, in industrial systems, the carryover of alum floe
from a clarifier may cause deposit problems, particularly in cooling systems where
phosphate may be applied as a stabilizing treatment. Aluminum found in treated
water systems is usually there because of colloidal residues (alumina, Al2O3) from
the coagulation of the water if alum or aluminate is used as the coagulant. If the
residuals are objectionable, they can be removed by improved filtration practices.
As shown by the solubility curves (Figure 6.3) aluminum is amphoteric, being
present as Al3+ or lower valence hydroxyl forms at low pH and the aluminate
anion at higher pH values. As might be expected from this amphoteric nature,
alumina particles are positively charged at low pH and negatively charged at high
pH, as indicated by Figure 6.7. The effectiveness of alum in precipitating negatively charged colloids, such as clay particles from water, is more likely related to
the 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 to
7. Its strong negative charge at pH 10.0 to 10.5 helps explain the effectiveness of
sodium aluminate in precipitating magnesium hardness, which is positively
charged 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 indicator
of either mining or metallurgical operations in the watershed or runoff from agricultural areas where arsenical materials have been used as industrial poisons. If
in colloidal form, it would be removed by conventional water treatment processes. Federal regulations limit the content in public water supplies to a maximum of 0.1 mg/L total arsenic. If the material is present in organic form, it may
be removed by oxidation of the organic material and subsequent coagulation, or
by 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 is
less 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 processes. There are instances of barium being added to water for the specific purpose
of sulfate reduction. The reaction is hindered because the barium reagent itself is
so insoluble that considerable time is needed for the reactions to occur; furthermore, sulfate deposition on the surface of the barium reagent makes the process
inefficient. Barium is limited in drinking water to a maximum concentration of 1
mg/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 connate
waters produced with oil contain several hundred milligrams per liter and are the
source of commercial bromine. Over 0.05 mg/L in fresh water may indicate the
presence of industrial wastes, possibly from the use of bromo-organo compounds
as biocides or pesticides.
Copper (Cu2+—Atomic Weight 63.5; Group IB Metal)
Copper may be present in water from contact with copper-bearing minerals or
mineral wastes from copper production. It is more likely, however, that the copper found in water will be a product of corrosion of copper or copper alloy piping
or fittings, or may have been added deliberately to a water supply reservoir for
algae control, as copper sulfate. When copper sulfate is added for algae control,
because its solubility is limited, organic chelating materials may be added to the
copper sulfate formulation to keep the copper from precipitating and, therefore,
maintain its effectiveness. Drinking water regulations limit the municipal water
supply concentration to 1 mg/L maximum. At higher concentrations, the water
has an astringent taste. If a water supply is corrosive to copper, the first drawing
or tapping of the supply from piping which has been idle overnight may contain
relatively high concentrations, and ingestion of this water may cause immediate
vomiting. In industrial supplies, the presence of copper can be objectionable as it
is corrosive to aluminum. Copper is essential to certain aquatic organisms, being
present 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 metallurgical 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-bearing
alloys, such as solder. Being amphoteric, lead is attacked in the presence of caustic
alkalinity.
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 wastewaters
where lead may be complexed with organic matter, it may be solubilized, and
oxidation 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. There
are no records of experience indicating that this material is of concern either in
industrial or municipal water supplies. However, lithium salts are used in psychotherapy to combat depression, so there may be a concentration level in water
that has a psychotropic effect. Lithium salts have a wide variety of uses, but the
industrial consumption is so low that it is not likely to be a significant factor in
the 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 metamorphic
rocks. In water free of oxygen, it is readily dissolved in the manganous (Mn 2+ )
state and may be found in deep well waters at concentrations as high as 2 to 3
mg/L. It is also found with iron in acid mine drainage. Wastewaters from metallurgical and mining operations frequently contain manganese.
It is an elusive material to deal with because of the great variety of complexes
it can form depending on the oxidation state, pH, bicarbonate-carbonate-OH
equilibria, and the presence of other materials, particularly iron.
It is limited to 0.05 mg/L maximum by drinking water regulations, because
higher concentrations cause manganese deposits and staining of plumbing fixtures
and clothing. However, concentrations even less than this can cause similar
effects, as it may accumulate in the distribution system as a deposit, to be released
in higher concentrations later if the environment should change, such as by
change in pH, CO2 content, oxidation potential, or alkalinity.
In industrial systems it is as objectionable as iron, particularly in textile manufacture or the manufacture of bleached pulp, since small amounts of deposited
manganese 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 finishing
operations.
In the oxidized state, manganese is quite insoluble, and can be lowered in concentration, even in an alum coagulation process by superchlorination with adequate filtration, even at a pH as low as 6.5. However, the conventional process
for removal of manganese by itself is oxidation plus elevation of the pH to
approximately 9 to 9.5, with retention of approximately 30 min in a reaction vessel before filtration. Filters that have gained a coating of manganese oxides can
work very effectively, but may slough manganese if the aquatic environment is
radically changed. Manganese is also precipitated by the continuous application
of potassium permanganate ahead of a manganese form of zeolite.
Organic materials can chelate manganese much as they chelate iron, so
Manganese mg/l
Concentra ion at Bottom 23 meters
Fall Turnover
Spring Turnover
Concentration at 10 meters
Month
FIG. 6.8 Manganese concentration in the top and bottom layers of a lake, as affected by
seasonal changes. (From "Chemistry of Manganese in Lake Mendota, Wisconsin," Environ. Sc. Tech., December 1968.)
destruction of the organic matter is often a necessary part of the manganese
removal process.
Because manganese accumulates in sediments, it is common to find high levels
of manganese in deep water where none may be apparent at the surface. This
should be studied in designing the proper intake structure for a plant water supply. An example of this is illustrated in Figure 6.8, showing the manganese concentrations in a lake at various times of the year and at different depths in the
lake.
Phosphate (PO43-, Molecular Weight 95; Compound of Phosphorus P—Atomic
Weight 31)
MOLE FRACTION
Phosphorus is found in many common minerals such as apatite, in the form of
phosphate (PO43"equivalent weight 31.7). Since phosphate compounds are widely
used in fertilizers and detergents, it is common to find phosphate in silt from
agricultural runoff, with fairly high concentrations being found in municipal
wastewater, 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 algal
growths, which lead to eutrophication of lakes and streams, a reduction of phosphate is being brought about by legislation restricting the amount of phosphate in
detergents and also requiring treatment of municipal sewage for phosphate
removal.
Phosphate may be present in water as HPO42" and H2PO4-, as well as the
higher pH form, PO43". 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 phosphate
and 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 phosphate
species.
duce residuals less than 2 to 3 mg/L in the form of hydroxyapatite; in a hot process system, the residuals would be less than 0.5 mg/L.
These phosphate precipitates are often colloidal, and filtration is required to
achieve 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 of
considerably lower solubility in natural waters with a neutral pH and having
bicarbonate alkalinity. (See Figure 6.3 for its solubility characteristics.) Zinc is
seldom found at concentrations over 1 mg/L, with a typical concentration being
approximately 0.05 mg/L. Because it tends to have an astringent taste, its concentration in public water supplies is limited to 5 mg/L maximum.
Zinc may be present in water because of waste discharges from mining, metallurgical, or metal finishing operations. It may also appear because of corrosion
of galvanized steel piping. It is often included in proprietary corrosion inhibitors
where its effect on steel piping is similar to that of galvanizing.
Zinc would be removed in lime softening operations to residuals well below
0.1 mg/L. It can also be removed by cation exchange on either the sodium or the
hydrogen cycle.
CLASS 4—TRACE CONSTITUENTS
Materials in this group (Table 6.3) are generally found at concentrations less than
0.01 mg/L.
CLASS 5—TRANSIENT CONSTITUENTS
This class includes constituents which change in concentration or activity not by
dilution, dissolution, or precipitation, but rather by changes in the aquatic environment which disturb the equilibrium. These changes may come about from biological activity, oxidation-reduction potential, or radioactive decay.
Acidity-Alkalinity
The typical domain of almost all natural waters is characterized by a pH range of
6 to 8, the presence of bicarbonate alkalinity, and some CO2 dissolved in the
water. All waters in contact with limestone, dolomite, or geologic formations
including these minerals tend to reach this equilibrium: it is the end result of the
chemical reactions that cause the weathering of rocks and of the oxidation-reduction reactions which are mediated by aquatic organisms. Because of this, the few
exceptional streams that contain free mineral acidity (i.e., have a pH below about
4.5) usually dissipate this condition by accelerated weathering of the alkaline
components of the rocks they contact. Likewise, when the pH exceeds 8 and carbonate alkalinity begins to appear, this is brought into balance by reaction with
carbon dioxide from the atmosphere or from respiration of aquatic life.
TABLE 6.3
Class 4—Trace Constituents
Constituent
Occurrence*
Behaviorf
Antimony
(Sb, group VA,
at. wt. 74.9)
Cadmium
(Cd, Group
IIB, at. wt.
112.4)
Leaching of metallurgical
slags; colloidal hydrous
oxides
Plating wastes; limited
solubility as Cd2+.
Note: Restricted in
potable water supplies
to 0.01 mg/L
maximum.
Plating wastes, cooling
tower blowdown.
Soluble as CrO42~
(Cr6+); insoluble as
Cr3+. Note: Restricted
in potable water
supplies to 0.05 mg/L
maximum as Cr.
Present in copper- and
nickel-bearing ore
tailings; ceramic wastes.
Insoluble in aqueous systems
containing HCO3/CO3/OH.
Filterable.
Behaves as Ca and Zn (see
Figure 6.3). Can be
precipitated as carbonate or
removed by cation exchange.
Chromium
(Cr, group
VIB, at. wt.
52)
Cobalt
(Co, group
VIII, at. wt.
59)
Cyanide
(CN~, eq. wt.
26)
Mercury
(Hg, group IIB,
at. wt. 200.6)
Nickel
(Ni, group
VIII, at. wt.
58.7)
Wastes from plating shops,
coke plants, blast
furnaces, petroleum
refining. Note:
Restricted in potable
water supplies to 0.2
mg/L maximum.
Wastes from electrolytic
NaOH production,
leaching of coal ash.
Note: Restricted in
potable water supplies
to 0.002 mg/L
maximum.
Plating wastes; electric
furnace slag or dust; ore
tailings.
Tin
(Sn, group
IVA, at. wt.
118.7)
Tinplate waste.
Titanium
(Ti, group
IVB, at. wt.
47.9)
Ilmenite in well
formation.
As CrO42+, oxidizing agent.
Can be reduced by SO2 to
Cr3+ or removed by anion
exchange. As Cr3+, colloidal
hydrous oxide at neutral pH,
filterable.
Behaves as iron.
Behaves as Cl-, NO3", highly
soluble. Can be oxidized
with Cl2 to CNO ~ and N2 +
CO2. Reduced in activated
sludge biodigestion.
Removable by anion
exchange. Forms complexes
with Cd, Fe.
May be methylated by bacterial
activity and taken up by the
aquatic food chain.
Removed by closing plant
loop or by reduction and
filtration.
Ni2+ behaves as iron, Fe2+.
Seldom present as Ni 3+ .
Precipitates as hydroxide
and basic carbonate. Sulfide
precipitates are insoluble.
Removable by cation
exchange.
Converted to hydrous oxide
colloid at neutral pH or in
HCO3/CO3 environment.
Colloidal TiO2. Removed by
filtration.
* Most constituents in class 4 are introduced by industrial waste discharges, and are more generally
found in surface than in well waters.
Uncontrolled discharge of industrial wastes could wipe out this natural buffering effect of equilibrium between the aquatic environment, the atmosphere, and
the lithosphere.
Since legislation now prohibits such discharge, the only circumstances that can
cause waters to fall outside the natural conditions are accidental spills of large
volumes of strong chemicals, seepage of acid mine drainage into a stream, or acid
rain from air pollution. The mine seepage may not be controllable because of
inability 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-fired
boiler plants. When the acid rain falls on alkaline rock or into rivers or lakes in
limestone basins, there may be enough reserve alkalinity in the rock or dissolved
in the water to neutralize the acidity. But often the rain falls in forested areas
where vegetative litter has developed a soil high in humus. There is no natural
alkalinity then to counteract the acidity of the rainfall, so the runoff is acidic and
if 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 Figure 2.11. Possibly at some time in the past, alkaline components in industrial gas
discharges (e.g., alkaline fly ash, cement kiln, or lime kiln dust) alleviated some
of this problem, and it is likely that application of limestone to some lakes affected
by rain will be both effective and economically justified. Reduction of sulfur
oxides from boiler plant flue gas often utilizes lime as the neutralizing agent. But
complete elimination of these sulfur oxides is impossible in the boiler stack. So
alkali treatment of acid-affected lakes may prove more economical and practical
than further treatment at the boiler stack as further addition of alkali there runs
into diminishing returns. However, this does not correct other deleterious effects
of acid rain, such as damage to the flora of forests and croplands and etching and
corrosion 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 generalized formula for biomass. It has been hypothesized that earth's primitive environment contained carbon in the form of methane plus ammonia, water, and
hydrogen gases. Methane is one of the carbon compounds present in the carbon
cycle (Figure 6.10), produced by fermentation of larger organic molecules. Carbon
dioxide 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 decomposition of organic chemical compounds. Methane is given off by anaerobic
decomposition of organic sediments in marsh gases, and the concentration may
become so high that the swamp gas may ignite. It is more common to find methane 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 earlier
chapters. The carbon dioxide content of surface waters is greatly influenced by
Organic
Carbon
Compounds
C+4
Animal
Metabolism
Decay
and Death
Methane
CH4
Aerobic
Digestion
Oxidation
Digestion
Plant
Metabolism
Carbon
C°
Carbon
Monoxide
CO
C+2
Incomplete
Combustion
Oxidation
Carbohydrates
(CH2O),;
Direct Photosynthesis
By Chlorophyll - Containing Plants
Carbon
Dioxide
CO2
C +4
Ionic
Domain
HCOg -CO*
FIG 6.10 The carbon cycle.
bacterial and algal symbiotic existence, illustrated by the oxygen cycle, Figure
6.11. During bright sunlight, the photosynthetic reactions proceed so rapidly that
the water may actually become supersaturated with oxygen, beyond the capacity
of the bacteria to utilize. If algae require more carbon dioxide than is available
from bacterial respiration, they assimilate carbon dioxide from the bicarbonate
alkalinity, producing a trace of carbonate alkalinity. For that reason, carbon dioxide and oxygen are variable in most surface supplies, as affected by sunlight and
the photosynthetic process (Figure 6.12).
The plant tissue built up by photosynthesis is eventually metabolized by larger
aquatic organisms that produce organic compounds and discharge them in their
wastes. Organic compounds are also produced by the death and decay of the
aquatic plants and fish life. Organic matter thus produced becomes food for bacteria and is returned to the cycle as methane by anaerobes and as carbon dioxide
by aerobic bacteria.
Deep well water often contains over 25 mg/L CO2, and may be saturated with
this gas at the hydrostatic pressure and temperature in the water table. A drop in
pressure as water flows across the well screen may cause the CO2 to come out of
solution, disrupting the equilibrium and depositing CaCO3 scale.
The theoretical solubility of oxygen exposed to the atmosphere is dependent
ALGAL
PHOTOSYNTHESIS
ATMOSPHERE
O2
CO2
CO2
BACTERIAL
RESPIRATION
DISSOLUTION
IN WATER
FIG. 6.11 The oxygen cycle.
O2 CONCENTRATION
CO2 CONCENTRATION
FIG. 6.12 Diurnal variation of oxygen and
carbon dioxide in a surface water.
on the temperature (Figure 6.13). Excess oxygen concentration is due to photosynthesis, and a deficiency is usually caused by bacterial activity or reducing
agents.
Nitrogen Cycle Constituents (Figure 6.5)
Constituents of the nitrogen cycle have been discussed earlier. As was true with
the carbon cycle, the nitrogen cycle is involved with life in the aquatic
environment.
GASSOLUBILITY cc/l
NITROGEN
OXYGEN
TEMPERATURE
FIG. 6.13 Solubilities of oxygen and nitrogen in
water at various temperatures.
Sulfur Cycle (Figure 6.14)
Because sulfur is in the same family as oxygen, there are many compounds where
sulfur replaces oxygen in a compound with similar properties. For example,
ethanol (CH3CH2OH) and ethyl mercaptan (CH3CH2SH) are analogous compounds with only the sulfur and oxygen atoms interchanged.
Certain bacteria can metabolize the sulfur atom in hydrogen sulfide, just as
algae and other plants can metabolize oxygen from water in photosynthesis to
produce free oxygen and carbohydrate. The by-product of the bacterial process of
splitting H2S is free sulfur. The corresponding chemical equations are:
Algal
photosynthesis
CO2 + 2 H2O
> CH2O + O 2 -H H2O
(2)
Bacterial
action
CO2 + 2 H2S —> CH2O + 2 S +H2O
(3)
Hydrogen sulfide, which is present in some deep well waters and some stagnant
surface waters, is generally produced by the anaerobic decomposition of organic
Organic
Sulfur
Compounds
2
s-
Animal
Life
Processes
Oxidation
Decomposition
(Death)
I
H2S%HS"1-S"2
(S 2 )
Partial
Combustion
Digestion
Plant
Life
Processes
Bacterial
Reduction
Native Sulfur (S0)
Combustion
Oxidation
Reduction
(Sulfur
Depositing
Bacteria
Sulfite
so'2
(S+ 4 )
Sulfate
2
so;
(S+ 6 )
FIG. 6.14 The sulfur cycle.
compounds containing sulfur or by sulfate-reducing bacteria capable of converting sulfate to sulfide.*
All of these biological processes that occur in nature can be put to work under
controlled conditions by the water specialist, to digest and eliminate undesirable
organic wastes or their by-products.
* In each of these cycles involving biological activity, the changes occurring in surface water are much
more pronounced than those occurring in deep well waters. The surface supplies are constantly inoculated by microbes from the air and from the soil and supplied with solar energy. On the other hand, well
water usually represents a dead end in the cycle, with the water generally having a long residence time in
the aquifer, with the likelihood that most organisms have been filtered by the porous formation so that
all of the microbes may have been removed from the strata at the well screen. Because of this, it is
common 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 in
well waters. For the same reason, the concentrations OfCO 2 , NO3", and HS ~ are generally higher in
deep well water than in surface supplies.
Oxidation-Reduction Potential
Some materials in water are transient because they tend to oxidize or reduce other
constituents, either through biological activity, as illustrated by the several biological cycles, or directly. The presence of these materials has already been considered in the text covering the biological cycles and the constituents taking part
in these cycles.
The only significant oxidizing materials encountered in the natural environment that can participate in chemical reactions without needing a biological route
are oxygen, which is prevalent in surface supplies, and sulfur, which is encountered by subsurface waters in contact with native sulfur, a special situation. There
are numerous oxidizing materials that can appear as residues in treated wastewaters—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; reducing
materials which may be added as treatment by-products or waste residues include
a wide variety of organic matter, ferrous iron from such operations as steel pickling, and sulfite, present in certain kinds of pulp mill wastes.
Radionuclides
Water itself is not radioactive, but may contain elements that are. These enter the
water cycle as wastes from nuclear power plants, fallout from nuclear blasts, or
the by-products of metallurgical processing of radioactive materials. In very rare
cases, well waters may contain radionuclides—common radioactive elements and
their isotopes—as natural contaminants.
The water chemist normally deals with concentrations at part per million or
milligrams per liter levels. One mole of any substance represents its molecular
weight, so that the concentration of 1 mole/L as calcium carbonate represents 100
g/L as CaCO3. Since a mole contains 6 X 1023 molecules, 100 g/L contains this
number of molecules, and 100 mg/L contains 6 X 1020 molecules. Even at the
very low concentration of 1 ppb (0.001 mg/L), then, there are still 6 X 1015 molecules of the substance as calcium carbonate in a liter of water. Concentrations of
radioactive substances at this level would be extremely dangerous.
In chemical change, one element reacts with another in a process involving the
electrons surrounding the nucleus, with the nucleus remaining unchanged. Energy
may 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 may
include nuclear particles (such as protons, neutrons, or electrons) and energy,
including heat and electromagnetic radiation. The typical by-products of nuclear
reactions, then, are alpha particles, electrons, and electromagnetic radiations. The
general characteristics of these emanations are shown in Table 6.4.
TABLE 6.4 Types of Radiation
Alpha
Beta
Helium nuclei, charge + 2,
mass 4
Electrons, charge — 1, mass O
Gamma
Similar to x-rays
Can penetrate air, but is stopped by
solids.
About 1000 times more penetrating than
alpha radiation. Can penetrate 2-3
mm into solids.
Can penetrate air or solids about ten
times deeper than beta radiation.
Some nuclear reactions are extremely rapid, occurring in a split second, but
others may require 1000 years. Because these reactions may never go to completion, the life of a material participating in a nuclear reaction (called a radionuclide) is expressed as the time required for one-half its energy to be dissipated, its
half-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 per
second. The levels of radioactive disintegration of concern to the water technologist are so far below this that they are expressed as micromicrocuries, or picocuries (pCi), a pCi being equal to 2.2 disintegrations per minute.
A variety of radionuclides may be found in waters leaving a nuclear power
plant. The government license to operate such a plant specifies the particular
radionuclides which must be identified individually. This requires sophisticated
and painstaking analytical techniques. In municipal supplies, it is not required
that individual species be identified because the radiation levels are so low, but
an analysis is required for alpha radiation, which is assumed to be contributed by
226
Ra, beta radiation (assumed to be contributed by strontium, 90Sr) and gross beta
activity. Table 6.5 shows the relative relationship between occupational levels and
TABLE 6.5 Radionuclide Contamination in Water
226
Occupational levels
(body exposure)
pCi/L
ppb
Potable levels (ingestion)
pCi/L
ppb
Ra
90
Sr
4 X 101
4 X IQ- 5
8 X 102
4 X 10~6
3
3 X ICT6
101
5 X 10-8
potable water levels and also relates the radiation intensity to the concentration
of 226Ra or 90Sr in the water. It is obvious that the concentration levels are far
below those considered significant when dealing with nonradioactive contaminants in water supplies.
The heavier radionuclides are generally insoluble and may be removed by
coagulation and filtration; the soluble constituents may be removed by ion
exchange. These processes become quite complicated when the radionuclides are
present in very low concentrations and the treatment process may first have to
react with more common contamination, such as calcium and magnesium, at
concentrations far above those of the radionuclides before the radionuclides are
reduced to acceptable levels.
METHODS OF ANAL YSIS
It is not in the scope of this handbook to include methods of analysis for each of
the constituents presented in this chapter. These methods are covered in manuals
devoted exclusively to such analyses (e.g., Standard Methods for Analysis of
Water and Wastewater, ASTM Standards). However, included in Chapter 7 is a
summary of the methods of analysis used, detection limits, methods of sample
preservation, and other facts of importance.
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 Congress, 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," in
Aqueous-Environmental Chemistry of Metals (Ruben, J., ed.), Ann Arbor Science, Ann
Arbor, Mich., 1976.