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An acceptable range for drinking water pH is from 6.5 to 8.5. Corrosion effects may become significant below pH 6.5, and the frequency of incrustation and scaling problems may be increased above pH 8.5. With increasing pH levels, there is also a progressive decrease in the efficiency of chlorine disinfection processes.
The p value of any entity is defined as the negative common logarithm of that entity. Thus,
pM = - log10M
where M is an activity, concentration or equilibrium constant, for example. The pH of a solution is the negative common logarithm of the hydrogen ion activity, aH+:
pH = - log10 (aH+)
In a dilute solution, the hydrogen ion activity is approximately equal to the concentration of hydrogen ion.
The pH of an aqueous sample is usually measured electrometrically with a glass electrode.Footnote 1,Footnote 2 The pH may therefore be operationally defined in terms of E, the electromotive force in volts between a glass electrode and a reference electrode when the electrodes are immersed in the sample, and EB, the electromotive force obtained when the electrodes are immersed in a reference buffer solution, the assigned pH of which is designated pHB.Footnote 2 Thus,
pH = pHB + [(E - EB)F/2.3026RT]
where F is the faraday, R is the gas constant and T is the absolute temperature. The reference buffer solution is thus an integral part of the standard method used.
Temperature exerts two significant effects on pH measurement.Footnote 1,Footnote 2 The change in potential per pH unit varies (according to the term in the defining equation) from about 55 mV/pH unit at 5°C to about 66 mV/pH unit at 60°CFootnote 2; this can be overcome electronically by a temperature compensation adjustment provided on the better commercial instruments. The constants of chemical equilibria existing in a buffer solution vary and therefore affect pHB; pHB values must accordingly be corrected for temperature. The pH values at various temperatures of standard buffer solutions are known. Changes of temperature also affect the ionization equilibria of any weak acids and bases that are present in a water sample. The magnitude of this effect depends, to a large extent, on the alkalinity of the sample. The temperature of the sample should therefore be recorded together with the pH measurement.
At high pH levels, glass electrodes suffer from interference by alkali metal and alkaline earth metal cations. This is commonly referred to as the "sodium ion effect." For normal general-purpose glass electrodes, this effect becomes noticeable at pH levels above about 10.5, when a positive error is observed.Footnote 3 To reduce the possibility of such errors, measurements above pH 10 should be made with a high-alkalinity type of electrode.Footnote 1,Footnote 2
The pH of an aqueous system is a measure of the acid-base equilibrium achieved by various dissolved compounds and, in most natural waters, is controlled by the carbon dioxide - bicarbonate - carbonate equilibrium system. The following equilibria are involvedFootnote 4:
H2O ↔ H+ + OH-, pKw = 14.0
CO2 (g) ↔ CO2 (aq)
CO2 (aq) + H2O ↔ H2CO3, pKa ≈ 2.8
H2CO3 ↔ H+ + HCO3-, pK1 = 6.35
HCO3- ↔ H+ +CO32-, pK2 = 10.3
where the pK values are those at 25°C.
All these component equilibria are affected by temperature; Kw exhibits the largest variation. In pure water, a decrease in pH of about 0.45 occurs as the temperature is raised by 25°C.Footnote 5 In waters with a buffering capacity that is imparted by bicarbonate, carbonate and hydroxide ions, this temperature effect is modified.Footnote 5
The pH of most raw water sources lies within the range 6.5 to 8.5.Footnote 6 In some soft water areas, however, the pH can be considerably lower as a result of the leaching of organic acids from decaying vegetationFootnote 7 and the presence of dissolved carbon dioxide.Footnote 8 In some groundwater sources, carbon dioxide, generated by bacteriological oxidation, is unable to escape to the atmosphere, and even lower pH levels will result.Footnote 9
Hydrogen ion concentration may be significantly altered during water treatment. Chlorination tends to lower the pH, whereas water softening using the excess lime/soda ash process raises pH levels. In a survey of the water supplies of the 100 largest U.S. cities,Footnote 10 it was found that although all the cities used a raw water supply with a pH less than 9, 17 cities provided drinking water with pH levels greater than 9. The range of pH in finished water encountered in this survey was from 5.0 to 10.5, with a median level of 7.5.
The economic loss caused by corrosion in water mains and water treatment facilities has been estimated at $375 million annually in the United States.Footnote 11 In contrast to the corrosion problem is the loss of distribution capacity and the concomitant increase in pumping costs that result from calcium carbonate deposition.Footnote 12
Metals used in distribution systems, such as cast iron, steel and copper, tend to corrode in contact with water because of their thermodynamic instability. The deterioration of concrete, asbestos-cement and cement-lined cast iron pipe, all of which are commonly used in distribution systems, may also occur. Natural waters contain gases, colloidal matter and a variety of electrolytes and non-electrolytes that, together with pH, determine the extent to which corrosion is possible in a given situation.Footnote 7 In general, the presence of anions that form soluble compounds with a metal increases the corrosiveness of the water with respect to that metal, whereas the presence of anions that form insoluble compounds may increase its passivity.
The role of pH in the corrosion of metals used in water distribution has been summarized by DraneFootnote 7 as follows:
1. Steel corrodes at approximately the same rate at all pH levels commonly found in natural waters. The form that the corrosion takes is, however, affected by pH. At values between 7.5 and 9.0, there is a tendency for the corrosion products to adhere in a hard, crusty deposit. At lower pH values, adherent corrosion products are not so evident, although a very hard form of deposit is sometimes seen in pipes that have been in service for some years. Loss of head, owing to scaling of a pipe, is more commonly found in the higher pH range; at lower pH values, "red water" complaints arising from corrosion products in suspension are more common. Cast iron behaves in a manner similar to steel at alkaline pH values, but at lower pH values it is subject to graphitization.
2. Copper is markedly affected by pH. In aggressive waters, slight corrosion occurs, and the small amount of copper in solution may cause staining of fabrics and plumbing fixtures. In addition, redeposition of copper on aluminum or galvanized surfaces sets up electrochemical cells resulting in pitting of these metals. In most waters, the critical pH value is about 7.0, but in soft waters containing organic acids it may be higher.
3. Lead corrosion is affected by carbonate content, pH and mineral constituents. The simplest method of control is usually to increase the pH by adding alkali. Few waters are plumbo-solvent if the pH is above 7.0.
Zinc coatings on iron and galvanized steel are attacked in the same way as iron, but usually more slowly. Very alkaline waters, above about pH 10.5, can be aggressive to zinc(13) and will often remove galvanized coatings.
Corrosion control may result from calcium carbonate deposition. The factors affecting this process are temperature, pH, total dissolved solids, hardness, carbon dioxide and alkalinity. A rigorous treatment of the calcium carbonate - bicarbonate equilibrium is invariably impossible under practical conditions. Accordingly, a number of semi-empirical and empirical relationships using easily measured parameters have been developed.
One of the earliest relationships was developed by Langelier,Footnote 5,Footnote 14 who studied the mathematical relationship between calcium carbonate, calcium bicarbonate and carbon dioxide in water. The simpler form of his equation, which is applicable when pHs lies between 6.5 and 9.5, isFootnote 15:
pH<s = pCa2+ + pAlk + (pK2 - pKs)
where pHs (the saturation pH) is the pH at which, with no change of alkalinity, calcium content or dissolved solids, the water would neither deposit nor dissolve calcium carbonate; the other terms are as follows:
pCa2+ : is the negative logarithm of the calcium concentration, expressed as mg CaCO3/L,</
pAlk : is the negative logarithm of the alkalinity to methyl orange, expressed as mg CaCO3/L,
pK2 : is the negative logarithm of the ionization constant of HCO3-,
pKs : and is the negative logarithm of the solubility product of CaCO3.
The term (pK2 - pKs) is a function of ionic strength and temperature. Nomograms relating this quantity to readily measured water parameters -- for example, pH, total dissolved solids, alkalinity and temperature -- have been published.Footnote 5,Footnote 6,Footnote 14.
The Langelier saturation index (SI) is defined as the difference between the actual pH of the water and the saturation pH:
SI = pH - pHs
A positive value for SI indicates scale-forming tendencies, and a negative value indicates scale-dissolving or aggressive qualities.
Although the solubility product of calcium carbonate decreases with increasing temperature, leading to increased deposition at higher temperatures under most conditions, this is not true for waters of low alkalinity (<50 mg/L as calcium carbonate). In such waters, the decrease in pH with temperature outweighs the decrease in solubility of calcium carbonate. This decrease in pH at temperatures above about 55°C actually increases the solubility of calcium carbonate. The effect on saturation index is particularly acute below 25 mg/L alkalinity in hot water systems, because there is insufficient alkalinity to buffer the effect of temperature on pH.Footnote 16 Ryznar introduced an empirical relationship, based on the Langelier saturation pH, named the stability indexFootnote 17:
stability index = 2 pHs - pH
Values of the index greater than about 7.0 indicate corrosive water, and values less than 7.0 indicate scale-forming tendencies. This index is of interest in evaluating waters of widely different compositions.
An aggressivity index (AI) has been established by the American Water Works Association for use with asbestos-cement pipeFootnote 18:
AI = pH + log (Alk) (Ca2+)
A value of less than 10 indicates that the water is unsuitable for use in such pipes.
Numerous other indices are also in use -- for example, momentary excess, driving force index, marble test, Enslow's stability indicator -- all of which are based on the calcium carbonate - bicarbonate equilibrium system. These have been discussedFootnote 15 with respect to water softening.
The effectiveness of corrosion protection by altering pH and calcium content (normally by using lime and/or carbon dioxide) depends on judicious balance of the carbonate-bicarbonate equilibrium system. A water that is exactly in equilibrium -- i.e., just stabilized -- with respect to calcium carbonate is normally corrosive to iron and steel because it has no power to form a calcium carbonate deposit. Supersaturated water, on the other hand, will form substantial scale unless suitably treated. This scale may or may not inhibit corrosion, depending on how well it adheres to the metal and on its porosity. The relationship between pH and the quality of deposited carbonate films has been studied by McClanahan and Mancy.Footnote 12
Ideal conditioning of water with respect to inhibition of corrosion and incrustation is given by the following characteristicsFootnote 19:
Waters that are excessively hard do not usually lead to severe corrosion problems, but they are prone to excessive incrustation and also reduce the effectiveness of soaps. Hardness is usually removed in water treatment by precipitation of the hardness-producing cations -- calcium as the carbonate and magnesium as the hydroxide. Lime is added, usually in a calculated excess, precipitating most of the calcium carbonate and magnesium hydroxide; this is followed by soda ash addition to remove the excess lime and any non-carbonate hardness. Water treated in this manner will have a pH of the order of 10.9 and will have scale-forming tendencies.Footnote 15 Stabilization can be achieved by recarbonation (addition of carbon dioxide) to a pH of 9.7 to 10 and by the addition of 0.25 to 0.5 mg sodium polyphosphate per litre.Footnote 15 However, other authoritiesFootnote 20 have recommended recarbonation to pH 8.6, to stabilize the water against subsequent excessive calcium carbonate deposition in the distribution system.
There is one final area that should be mentioned with respect to corrosion and incrustation, and that is the buildup of either beneficial or harmful biological slime on distribution pipe surfaces. The slime may serve to prevent the removal of oxidation products from and the penetration of oxygen to the pipe walls, thus inhibiting corrosion. Alternatively, excessive growths could create regions of locally low pH at the pipe surface by generating carbon dioxide. This could lead to localized corrosion, even though the bulk water might possess favourable stability and aggressivity indices.Footnote 21 Red water complaints are often the result of sudden growth of iron bacteria, which produce ferric hydroxide as the metabolic end product. The development of iron bacteria may be such that severe blocking of water pipes can occur in a matter of weeks. The growth of iron bacteria is very pH dependent, occurring over the range 5.5 to 8.2, with an optimum pH of about 6.5.Footnote 22
Insofar as aqueous chemical equilibria invariably involve hydrogen (and hydroxyl) ions, pH will be related, in one or more of several different ways, to almost every other water quality parameter.
Taste and odour in drinking water arise from a wide variety of sources, and no generalizations as to the effect of pH on these parameters can be made. In waters prone to sulphur contamination, the formation of gaseous hydrogen sulphide, which leads to "bad egg" odours, is thermodynamically favoured at pH values less than about 7.0.Footnote 13 Nitrogen trichloride, which has an objectionable pungent odour,Footnote 23 tends to be formed in greater concentrations at low pH values (<pH 7) during the chlorination process.Footnote 24 It has also been reported that at high pH levels drinking water acquires a bitter taste.Footnote 25
The colour intensity in a given water sample is increased by raising the pH.Footnote 26 This effect, known as the "indicator effect," has led to the suggestion that all colour measurements for water quality control be carried out at a standard pH of 8.3.Footnote 27
Turbidity, taste- and odour-producing compounds, micro-organisms and colour can be removed by a combination of coagulation, flocculation and filtration. The efficiencies of coagulation and flocculation processes are markedly dependent on pH, and it is standard practice in water treatment to adjust pH so that optimum floc formation is achieved.Footnote 28,Footnote 29 In certain instances, filtration efficiency is also sensitive to pH.Footnote 30
Of greater importance to the microbiological quality of water is the influence of pH on the effectiveness of chlorine disinfection. The germicidal efficiency of chlorine in water is lower at higher pH values; this has been attributed to the reduction in hypochlorous acid concentration with increasing pH.Footnote 33-Footnote 36 Hypochlorous acid has a germicidal effectiveness about 100 times greater than that of the hypochlorite ion.Footnote 37 Most natural waters, however, contain ammonia-nitrogen, which reacts with chlorine and hypochlorous acid to form mono-, di- and trichloramines (combined available chlorine), the relative amounts of which depend on pH. In many, if not most, treatment plants using chlorine disinfection, sufficient chlorine is added to oxidize all the ammonia and to provide excess free chlorine (break-point chlorination). Under these conditions, the hypochlorous acid concentration reaches a maximum at a pH of about 7.5, with lower concentrations at lower and higher pH values.Footnote 24
Chlorination in water treatment serves two purposes. The first is to inactivate pathogenic organisms in the water before entry into the distribution system. The second is to ensure that a free chlorine residual persists to the user's tap. It may be argued that high pH, with the attendant reduction in the rate of germicidal action, is detrimental to the effectiveness of free chlorine in the distribution system. However, it must be remembered that the hypochlorous acid - hypochlorite ion system is a chemical equilibrium, and removal of hypochlorous acid by reaction with micro-organisms wil lead to further formation of the acid, provided that there is a free chlorine residual. The reaction rate of hypochlorous acid disinfection is slower at lower pH levels, but this may be compensated for by longer contact times. The important parameter is the total available chlorine (as both HOCl and OCl-).
Ozone, which is used in more than 20 water treatment facilities in Quebec,Footnote 38 and chlorine dioxide, which is primarily employed in Canada for taste and odour control, are alternative disinfection agents. The effectiveness of both is unchanged by pH within the range of pH values encountered in drinking water. Chlorine dioxide has a germicidal effectiveness similar to that of hypochlorous acid, whereas that of ozone is considerably greater.Footnote 24
A major source of metal contamination in drinking water is corrosion in the water supply system.Footnote 39 Two of the potentially most troublesome metals are lead and cadmium. Lead is immune to corrosion at all pH levels above 6 in pure water; in the presence of carbonates and bicarbonates, lead is passive between about pH 4 and 12, but it becomes subject to corrosion above pH 12.Footnote 12 In a study of drinking water with a low alkalinity and a fairly low pH, high levels of lead were found in the drinking water of households that had lead plumbing.Footnote 40 Cadmium, in pure water, is passive between about pH 9 and 13.5, but experimental data show that corrosion is significant only below pH 6.Footnote 13
Water with an aggressivity index below 10 has been shown, in laboratory studies, to promote the release of asbestos fibres from asbestos-cement pipes.Footnote 40 It was also pointed out that pipes coated with coal tar enamel might be a source of polynuclear aromatic hydrocarbons (PAHs) owing to the leaching action of water.Footnote 40
Trihalomethanes are formed during the chlorination of waters with an organic carbon content. Postulated mechanisms for the formation of trihalomethanes involve initial reaction of hypochlorous acid with the organic precursors: electrophilic aromatic substitution of positive chlorine species and electrophilic addition of positive chlorine to appropriately activated double bonds,Footnote 41 and, more specifically, the reaction of chlorine with m-dihydroxybenzene structures within a fulvic acid lattice.Footnote 42 Such reactions, like the simpler haloform reaction, are base catalysed. It has been shown, both in the laboratoryFootnote 43 and under practical conditions in a water treatment plant,Footnote 44 that, for a given organic carbon content and chlorine dose, higher concentrations of trihalomethanes are formed at higher pH values.
Because pH is related to a variety of other parameters, it is not possible to determine whether pH has a direct relationship with human health. Insofar as pH affects the unit processes in water treatment that contribute to the removal of viruses, bacteria and other harmful organisms, it could be argued that pH has an indirect effect on health. The destruction of viruses by the high pH levels encountered in water softening by the lime/soda ash process could also be considered beneficial. On the other hand, the increased yield of trihalomethanes at high pH values may be detrimental.
In one of the few epidemiological studies carried out on drinking water supplies in which pH was one of the parameters considered, Taylor and co-workersFootnote 45 were unable to obtain any significant correlation between the incidence of infectious hepatitis and finished water pH. Sixteen U.S. cities that used surface water as a source of drinking water were considered in the study.
1. There are no specific health effects on which to base limits for the pH of drinking water. The main purpose in controlling pH is to produce water in which corrosion and incrustation are minimized. These processes, which can cause considerable damage to the water supply system, result from complex interactions between pH and other parameters such as dissolved solids, dissolved gases, hardness, alkalinity and temperature.
2. As a generalization, metal corrosion may become significant below a pH of about 6.5; incrustation and scaling problems are most commonly encountered above about pH 8.5.
3. The acceptable range for drinking water pH is therefore from 6.5 to 8.5. In general, waters with a pH within this range can be stabilized with respect to corrosion and incrustation by simple pH adjustment. By keeping the pH below 8.5, the rate of chlorine disinfection is increased and the production of trihalomethanes is reduced.
American Public Health Association/American Water Works Association/Water Pollution Control Federation. Standard methods for the examination of water and wastewater. 14th edition. Washington, DC (1976).
American Society for Testing and Materials. Standard method of test for pH of water and waste water. In: Annual book of ASTM standards. Part 31. Philadelphia, PA. p. 178 (1976).
Feldman, I. Use and abuse of pH measurements. Anal. Chem.,28: 1859 (1956).
Goldman, J.C., Porcella, D.B., Middlebrooks, E.J. and Toerien, D.F. Review paper. The effect of carbon on algal growth -- its relationship to eutrophication. Water Res., 6: 637 (1972).
Langelier, W.F. Effect of temperature on the pH of natural waters. J. Am. Water Works Assoc., 38: 179 (1946).
Webber, W.J., Jr., and Stumm, W. Mechanism of hydrogen ion buffering in natural waters. J. Am. Water Works Assoc., 55: 1553 (1963).
Drane, C.W. Natural waters. In: Corrosion. 2nd edition. L.L. Shreir (ed.). Newnes-Butterworths, London, U.K.
Butler, G. and Ison, W.C.K. Corrosion and its prevention in waters. Reinhold, New York, NY. p. 25 (1966).
Sawyer, C.N. and McCarty, P.L. Acidity. In: Chemistry for sanitary engineers. 2nd edition. McGraw-Hill Series in Sanitary Science and Water Resources Engineering, McGraw-Hill, Toronto. p. 320 (1967).
Durfor, C.N. and Becker, E. Selected data on public supplies of the 100 largest cities in United States, 1962. J. Am. Water Works Assoc., 56: 237 (1964).
Hudson, H.E., Jr., and Gilcreas, F.W. Health and economic aspects of water hardness and corrosiveness. J. Am. Water Works Assoc., 68: 201 (1976).
McClanahan, M.A. and Mancy, K.H. Effect of pH on quality of calcium carbonate film deposited from moderately hard and hard water. J. Am. Water Works Assoc., 66: 49 (1974).
Pourbaix, M. Atlas of electrochemical equilibria in aqueous solutions. 2nd edition. National Association of Corrosion Engineers, Houston, TX (1974).
Langelier, W.F. Chemical equilibria in water treatment. J. Am. Water Works Assoc., 38: 169 (1946).
Dye, J.F. and Tuepker, J.L. Chemistry of the lime-soda process. In: Water quality and treatment; a handbook of public water supplies. 3rd edition. American Water Works Association. McGraw-Hill, Toronto. p. 313 (1971).
Larson, T.E. Corrosion phenomena -- causes and cures. In: Water quality and treatment; a handbook of public water supplies. 3rd edition. American Water Works Association. McGraw-Hill, Toronto. p. 295 (1971).
Ryznar, J.W. A new index for determining amount of calcium carbonate scale formed by a water. J. Am. Water Works Assoc., 36: 472 (1944).
American Water Works Association. AWWA Standard C400-77 revision of C400-75. Denver, CO (1977).
Merrill, D.T. and Sanks, R.L. Corrosion control by deposition of CaCO3 films: Part 1. A practical approach for plant operators. J. Am. Water Works Assoc., 69: 592 (1977).
Sawyer, C.N. and McCarty, P.L. Residual chlorine and chlorine demand. In: Chemistry for sanitary engineers. 2nd edition. McGraw-Hill Series in Sanitary Science and Water Resources Engineering, McGraw-Hill, Toronto. p. 363 (1967).
O'Connor, J.T., Hash, L. and Edwards, A.B. Deterioration of water quality in distribution systems. J. Am. Water Works Assoc., 67: 113 (1975).
Shair, S. Iron bacteria and red water. Ind. Water Eng. (March- April): 16 (1975).
American Water Works Association Research Foundation. Handbook of taste and odor control experiences in the U.S. and Canada. American Water Works Association, Denver, CO (1976).
Morris, J.C. Chlorination and disinfection -- state of the art. J. Am. Water Works Assoc., 63: 769 (1971).
U.S. Environmental Protection Agency. Statement of basis and purpose for the National Secondary Drinking Water Regulations. March (1977).
Black, A.P. and Christman, R.F. Characteristics of coloured surface waters. J. Am. Water Works Assoc., 55: 753 (1963).
Singley, J.E., Harris, R.H. and Maudling, J.S. Correction of color measurements to standard conditions. J. Am. Water Works Assoc.,58: 455 (1966).
Sawyer, C.N. and McCarty, P.L. Chemical coagulation of water. In: Chemistry for sanitary engineers. 2nd edition. McGraw-Hill Series in Sanitary Science and Water Resources Engineering, McGraw-Hill, Toronto. p. 341 (1967).
Maudling, J.S. and Harris, R.H. Effect of ionic environment and temperature on the coagulation of color-causing organic compounds with ferric sulfate. J. Am. Water Works Assoc., 60: 460 (1968).
AWWA Research Committee on Coagulation. Coagulation-filtration practice as related to research. A committee report. J. Am. Water Works Assoc., 66: 502 (1974).
Davis, B.D., Dulbeco, R., Eisen, H.N., Ginsberg, H.S. and Wood, W.B., Jr. Microbiology. 2nd edition. Harper and Row, New York, NY. p. 92 (1973).
Rudolfs, W., Falt, L.L. and Ragatzkie, R.A. Literature review on the occurrence and survival of enteric, pathogenic, and relative organisms in soil, water, sewage, and sludges, and on vegetation. Sewage Ind. Wastes, 22: 1261 (1950).
Butterfield, C.T., Wattie, E., Megregian, S. and Chambers, C.W. Influence of pH and temperature on the survival of coliforms and enteric pathogens when exposed to free chlorine. Public Health Rep., 58: 1837 (1943).
Smith, W.W. and Bodkin, R.E. The influence of hydrogen ion concentration on the bactericidal action of ozone and chlorine. J. Bacteriol., 47(A17): 445 (1944).
Scarpino, P.V., Berg, G., Chang, S.L., Dahling, D. and Lucas, M.A comparative study of the inactivation of viruses in water by chlorine. Water Res., 6: 959 (1972).
Kott, Y., Nupen, E.M. and Ross, W.R. The effect of pH on the efficiency of chlorine disinfection and virus enumeration. Water Res.,9: 869 (1975).
White, G.C. Chlorination and dechlorination: a scientific and practical approach. J. Am. Water Works Assoc., 60: 540 (1968).
Federal, Provincial and Municipal Government Agencies, and the Federation of Associations on the Canadian Environment. National Inventory of Municipal Waterworks and Wastewater Systems in Canada 1975 (1975).
Craun, G.E. and McCabe, L.J. Problems associated with metals in drinking water. J. Am. Water Works Assoc., 67: 593 (1975).
McFarren, E.F., Buelow, R.W., Thurnau, R.C., Gardels, M., Sorrell, R.K., Snyder, P. and Dressman, R.C. Water quality deterioration in the distribution system. Water Quality Technology Conference, Kansas City, MO, December (1977).
Morris, J.C. and McKay, G. Formation of halogenated organics by chlorination of water supplies. U.S. Environmental Protection Agency Rep. EPA 600/1-75-002 (1975).
Rook, J.J. Chlorination reactions of fulvic acids in natural waters. Environ. Sci. Technol., 11: 478 (1977).
Stevens, A.A., Slocum, C.J., Seeger, D.R. and Robeck, G.G. Chlorination of organics in drinking water. J. Am. Water Works Assoc.,68: 615 (1976).
Harms, L.L. and Looyenga, R.W. Chlorination adjustment to reduce chloroform formation. J. Am. Water Works Assoc., 69: 258 (1977).
Taylor, F.B., Eagen, J.H., Smith, H.F.D., Jr., and Coene, R.F. The case for water-borne infectious hepatitis. Am. J. Public Health, 56: 2093 (1966).