CHAPTER 17 INTERNAL CORROSION AND DEPOSITION CONTROL1 Michael R Schock, Chemist U.S Environmental Protection Agency Water Supply and Water Resources Division Cincinnati, Ohio Corrosion is one of the most important problems in the drinking water industry It can affect public health, public acceptance of a water supply, and the cost of providing safe water Deterioration of materials resulting from corrosion can necessitate huge yearly expenditures of resources for repairs, replacement, and system Many times the problem is not given the attention it needs until expensive changes or repairs are required Corrosion tends to increase the concentrations of certain metals in tap water Two potentially toxic metals (lead and cadmium) are attributable almost entirely to leaching caused by corrosion Three other metals—copper, iron, and zinc—cause staining of fixtures, or metallic taste, or both Low levels of tin and antimony can be caused by the corrosion of lead-free solders (Herrera, Ferguson, and Benjamin, 1982; Subramanian, Connor, and Meranger, 1991; Subramanian, Sastri, and Connor, 1994) Nickel has sometimes been mentioned as a potential contaminant from the plating of decorative plumbing fixtures The promulgation of the Lead and Copper Rule by the U.S Environmental Protection Agency (USEPA) in 1991 has created an emphasis on corrosion control in distribution systems, as well as domestic, public, and institutional plumbing systems (Federal Register, 1991a,b, 1994a) Corrosion products attached to pipe surfaces or accumulated as sediments in the distribution system can shield microorganisms from disinfectants (see Chapter 18) These organisms can reproduce and cause problems such as bad tastes, odors, slimes, and additional corrosion Several researchers have recently promoted corrosion control within the distribution system as an effective way to maintain water quality and adequate disinfection (Rompré et al., 1996; Schreppel, Frederickson, and Geiss, 1997; Camper, 1997; Kiéné, Lu, and Lévy, 1996; Norton et al., 1995; Olson, 1996) The views expressed in this paper are those of the author and not necessarily reflect the views or policies of the U.S Environmental Protection Agency 17.1 17.2 CHAPTER SEVENTEEN The term corrosion is also commonly applied to the dissolution of cement-based materials, and the leaching of their free lime component The most common manifestation of this problem is the increase in pH, which can be detrimental to disinfection and the aesthetic quality of the water, as well as reducing the effectiveness of phosphate corrosion inhibitor chemicals intended to control the corrosion of metals The release of asbestos fibers is of regulatory concern, and in extreme cases, the chemical attack on the pipe by the water may cause a reduction of structural integrity and, ultimately, failure Even when a water system passes all regulatory requirements, the release of corrosion by-products by miles of distribution system and domestic piping, and the application of corrosion inhibitor chemicals containing metals such as zinc, can be significant sources of metal loading of wastewater treatment plants This contamination source can affect their ability to meet discharge or sludge disposal limits Phosphate-based corrosion inhibitors can provide unwanted nutrients to wastewater plants and can cause violations of wastewater or other discharge regulations, or water quality problems in ecosystems receiving the water Corrosion-caused problems that add to the cost of water include the following: Increased pumping costs caused by tuberculation and hydraulic friction Loss of water and water pressure caused by leaks Water damage to the dwelling, requiring that pipes and fittings be replaced Replacing hot water heaters Customer complaints of “colored water,” “stains,” or “bad taste,” for which the response may be expensive both in terms of money and public relations Increased wastewater and sludge treatment and disposal costs Increased dosage of chlorine to maintain a distribution system residual Corrosion is the deterioration of a substance or its properties because of a reaction with its environment In the waterworks industry, the “substance” that deteriorates may be a metal pipe or fixture, the cement mortar in a pipe lining, or an asbestoscement (A-C) pipe For internal corrosion, the “environment” of concern is water All waters are corrosive to some degree A water’s corrosive tendency will depend on its physical and chemical characteristics Also, the nature of the material with which the water comes in contact is important For example, water corrosive to galvanized iron pipe may be relatively noncorrosive to copper pipe in the same system Corrosion inhibitors added to the water may protect a particular material, but may either have no effect or may be detrimental to other materials Physical and chemical interactions between pipe materials and water may cause corrosion An example of a physical interaction is the erosion or wearing away of a pipe elbow from high flow velocity in the pipe An example of a chemical interaction is the oxidation or rusting of an iron pipe Biological growths in a distribution system (Chapter 18) can also cause corrosion by providing an environment in which physical and chemical interactions can occur The actual mechanisms of corrosion in a water distribution system are usually a complex and interrelated combination of these physical, chemical, and biological processes They depend greatly on the materials themselves, and the chemical properties of the water The purpose of this chapter is to provide an introduction to the concepts involved in corrosion and deposition phenomena in potable waters Each material that can corrode has a body of literature devoted to it Detail on the form of corrosion of each metal or piping material and specific corrosion inhibi- INTERNAL CORROSION AND DEPOSITION CONTROL 17.3 tion practices that might be employed can be found in a comprehensive text (AWWARF-TZW, 1996; Trussell, 1985; Snoeyink and Kuch, 1985) and water treatment journal articles Table 17.1, modified slightly from the original source (Singley, Beaudet, and Markey, 1984; AWWA, 1986), briefly relates various types of materials to corrosion resistance and the potential contaminants added to the water In general, plastic plumbing materials are more corrosion-resistant, but they are not without their own potential problems CORROSION, PASSIVATION, AND IMMUNITY Electrochemical Reactions Metal species can be released into water either from the simple dissolution of existing scale materials, or actual electrochemical corrosion followed by dissolution In some cases, scale materials formed from corrosion by-products may be eroded from the pipe surfaces Almost all mineral salts dissolve in water to some extent, from insignificant traces to high concentrations in seawater This section will provide a general overview of some aspects of the electrochemistry of metallic corrosion as it applies in the context of drinking water treatment However, many specialized texts on electrochemistry and electrochemical corrosion are available (Piron, 1991; Ailor, 1970; Bockris and Reddy, 1973; Butler and Ison, 1966; NACE, 1984; Pourbaix, 1966, 1973; Thompson, 1970) and should be consulted by readers who are interested in a comprehensive examination of the subject For corrosion of any type to occur, all of the components of an electrochemical cell must be present These include an anode, a cathode, a connection between the anode and cathode for electron transport, and an electrolyte solution that will conduct ions between the anode and cathode The anode and cathode are sites on the metal that have different electrical potential Differences in potential may arise because metals are not completely homogeneous If any one of these components is absent, a corrosion cell does not exist and corrosion will not occur Oxidation and dissolution of the metal takes place at the anode The electrons generated by the anodic reaction migrate to the cathode, where they are discharged to a suitable electron acceptor, such as oxygen The positive ions generated at the anode will tend to migrate to the cathode, and the negative ions generated at the cathode will tend to migrate to the anode Migration occurs as a response to the concentration gradients and to maintain an electrically neutral solution At the phase boundary of a metal in an electrolyte solution an electrical potential difference exists between the solution and the metal surface This potential is the result of the tendency of the metal to reach chemical equilibrium with the electrolyte solution This oxidation reaction, representing a loss of electrons by the metal, can be written as Me ⇔ Me z+ + ze− (17.1) Equation 17.1 indicates that the metal corrodes, or dissolves, as the reaction goes to the right This reaction will proceed until the metal is in equilibrium with the electrolyte containing ions of this metal The current that results from the oxidation of the metal is called the anodic current In the reverse reaction, the metal ions are chemically reduced by combining with electrons.The current resulting from the reduction (the reaction, Eq 17.1, going TABLE 17.1 Corrosion Properties of Materials Frequently Used in Water Distribution Systems* Plumbing material Copper Corrosion resistance Good overall corrosion resistance; subject to corrosive attack from high flow velocities, soft water, chlorine, dissolved oxygen, low pH, and high inorganic carbon levels (alkalinities) May be prone to “pitting” failures Lead Corrodes in soft water with pH < 8, and in hard waters with high inorganic carbon levels (alkalinities) and pH below ∼7.5 or above ∼8.5 Mild steel Subject to uniform corrosion; affected primarily by high high dissolved oxygen and chlorine levels, and poorly buffered waters Cast or ductile iron Can be subject to surface erosion by aggressive waters (unlined) and tuberculation in poorly buffered waters Galvanized iron or steel Subject to galvanic corrosion of zinc by aggressive waters, especially of low hardness; corrosion is accelerated by contact with copper materials; corrosion is accelerated at higher temperatures as in hot-water systems; corrosion is affected by the workmanship of the pipe and galvanized coating Asbestos-cement, concrete, Good corrosion resistance; immune to electrolysis; cement linings aggressive (soft) waters can leach calcium from cement; polyphosphate sequestering agents can deplete the calcium and substantially soften the pipe Plastic Resistant to corrosion Brass Good overall resistance; different types of brass respond differently to water chemistry; subject to dezincification by waters of pH > 8.3 with high ratio of chloride to carbonate hardness Conditions causing mechanical failure may not directly correspond to those promoting contaminant leaching Primary contaminants from pipe Copper and possibly iron, zinc, tin, antimony, arsenic, cadmium, and lead from associated pipes and solder Lead Iron, resulting in turbidity and red-water complaints Iron, resulting in turbidity and red-water complaints Zinc and iron; cadmium and lead (impurities in galvanizing process) Asbestos fibers; increase in pH, aluminum, and calcium Some pipes contain metals in plasticizers, notably lead Lead, copper, zinc * Source: Adapted from Singley, J E., B A Beaudet, and P H Markey, “Corrosion manual for internal corrosion of water distribution systems,” U.S Environmental Protection Agency, EPA/570/9-84-001 Prepared for Office of Drinking Water by Environmental Science and Engineering, Inc., Gainesville, FL, 1984 17.4 INTERNAL CORROSION AND DEPOSITION CONTROL 17.5 to the left) is called the cathodic current At equilibrium, the forward reaction proceeds at the same rate as the reverse reaction, and the anodic current is equal to the cathodic current Thus, no net corrosion is occurring at equilibrium The velocity of an electrochemical reaction, unlike that of a normal chemical reaction, is strongly influenced by the potential itself Corrosion results from the flow of electric current between electrodes (anodic and cathodic areas) on the metal surface These areas may be microscopic and in very close proximity, causing general uniform corrosion Alternatively, they may be large and somewhat remote from one another, causing pitting, with or without tuberculation Electrode areas may be induced by various conditions, some because of the characteristics of the metal and some because of the character of the water at the boundary surface Especially significant are variations in the composition of the metal or the water from point to point on the contact surface Impurities in the metal, sediment accumulations, adherent bacterial slimes, and accumulations of the products of corrosion are all related either directly or indirectly to the development of electrode areas for corrosion circuits Figure 17.1 shows an example of corrosion reactions taking place on a fresh pipe surface with proximate anodic and cathodic areas (Snoeyink and Jenkins, 1980) In almost all forms of pipe corrosion, the metal goes into solution at the anodic areas As the metal dissolves, a movement of electrons occurs and the metal develops an electric potential Electrons liberated from the anodic areas flow through the metal to the cathodic areas where they become involved in another chemical reaction, and the metal develops another electric potential The focus of corrosion control by water treatment methods is usually attempting to retard either or both of the primary electrode reactions The Nernst Equation The Nernst equation is a relationship that allows the driving force of the reaction to be computed from the difference in free energy levels of corrosion cell components The free energy difference under such conditions depends on the electrochemical potential, which, in turn, is a function of the type of metal and the solid- and aqueousphase reaction products Electrons (electricity) will then flow from certain areas of a metal surface to other areas through the metal A metal may go into solution as an ion, or may react in water with another element or molecule to form a complex, an ion pair, or insoluble compound FIGURE 17.1 Adjoining anodes and cathodes during the corrosion of iron in acidic solution (Source: Water Chemistry, V L Snoeyink and D Jenkins Copyright © 1980, John Wiley & Sons, Inc Reprinted by permission of John Wiley & Sons, Inc.) 17.6 CHAPTER SEVENTEEN The equilibrium potential of a single electrode can be calculated by using the Nernst equation for the general reaction (Eq 17.1): RT z+ z+ − ᎏ ln {Me } EMe/Mez+ = EMe/Me zF (17.2) where EMe/Mez+ = the potential (volts) E 0Me/Mez+ = the standard potential (volts), a constant that can be obtained from tables of standard reduction potentials { } denote activity of the ion Mez+ R = the ideal gas constant (about 0.001987 Kcal/deg⋅mol−1) T = the absolute temperature (°K) F = the Faraday constant (23.060 Kcal/V) z = the number of electrons transferred in the reaction The Me/Mez+ subscript indicates the reaction written as a reduction The Nernst equation, written as Eq 17.3, is for a single electrode, assuming that the electrode is coupled with the normal hydrogen electrode, at which the reaction 2H+ + 2e− → H2(g) E0 = 0.00 (17.3) takes place, and reactants and products are assumed by thermodynamic convention to equal (Stumm and Morgan, 1981; Garrels and Christ, 1965) The driving force computed by the Nernst equation is directly related to the Gibbs free energy for the overall reaction, through the relationship ∆Gr0 = −zFE, where ∆Gr0 is the freeenergy change for the complete reaction (Snoeyink and Jenkins, 1980; Stumm and Morgan, 1981) In this convention, standard “half-cell” potentials are tabulated with the reactions written as reductions It is usually more useful to use the Nernst equation in a general form for the balanced net reaction of two half-cells, each being of the form ox + ze− ⇔ red (17.4) where “ox” and “red” indicate “oxidized” and “reduced” species, respectively (Snoeyink and Jenkins, 1980; Stumm and Morgan, 1981, 1996): {red} RT Ered/ox = E0red/ox − ᎏ ln ᎏ zF {ox} (17.5) The subscript “red/ox” indicates overall cell potentials for the total balanced reaction At 25°C, and with the conversion to base-10 logarithms for convenience of calculation, Eq 17.5 can be rewritten as 0.0591 Ered/ox = E0red/ox − ᎏ log Q z (17.6) where Q is the reaction quotient ({red}/{ox}) At equilibrium, no electrochemical current is generated, and the oxidants and reductants are at their equilibrium activities Thus, the reaction quotient Q becomes equal to the equilibrium constant K for the overall reaction (Snoeyink and Jenkins, 1980) In drinking water systems, the oxidation half-cell reaction of a metal, such as iron, zinc, copper, or lead, is coupled with the reduction of some oxidizing agent, such as dissolved oxygen or chlorine species Example half-cell reactions are the following: INTERNAL CORROSION AND DEPOSITION CONTROL 17.7 O2 + 2H2O + 4e− ⇔ 4OH− (17.7) HOCl + H+ + e− ⇔ (1/2) Cl2 (aq) + H2O (17.8) ⁄2 Cl2 (aq) + e ⇔ Cl − − (17.9) Equations 17.8 and 17.9 can be combined to yield the net reaction: HOCl + H+ + 2e− ⇔ Cl− + H2O (17.10) which represents a significant oxidizing half-cell reaction for metals in drinking water If Eq 17.7 is written in the Nernst form (Eq 17.5), and the ionic strength of the water is low enough that it can be assumed to have unit activity ({H2O} = 1), then 0.0591 {OH−} EO2 /OH − = E 0O2 /OH − − ᎏ log ᎏ {O2} (17.11) The oxidation potential for this reaction clearly depends on pH, because the [OH−] is raised to the fourth power in the numerator Similarly, the oxidation potential of the hypochlorous acid reaction is directly related to pH: {Cl−} 0.0591 EHOCl/Cl − = E 0HOCl/Cl − − ᎏ log ᎏᎏ {HOCl}{H+} (17.12) because of the [H+] term in the denominator By thermodynamic definition, the corrosion (and, hence, dissolution of metals from plumbing materials) can only occur if the overall cell potential exceeds the equilibrium cell potential (Pourbaix, 1973; Snoeyink and Jenkins, 1980; Stumm and Morgan, 1981) Reactions such as Eq 17.7 and Eq 17.10 can be combined with metal oxidation half-cells (Eq 17.1) to show overall corrosion reactions likely to occur in drinking water Examples are 2Pb (metal) + O2 + 2H2O ⇔ 2Pb2+ + 4OH− (17.13) 2Fe (metal) + O2 + 2H2O ⇔ 2Fe2+ + 4OH− (17.14) Pb (metal) + HOCl + H ⇔ Pb + Cl + H2O (17.15) Fe (metal) + OCl− + 2H+ ⇔ Fe2+ + Cl− + H2O (17.16) + 2+ − As an example, the overall Nernst expression for the last preceding equation, at 25°C, is {Fe2+}{Cl−} 0.0591 EFe2+/OCl − = E 0Fe2+/OCl − − ᎏ log ᎏᎏ {OCl−} {H+}2 (17.17) Note that in Eq 17.17, the activities are for the free aqueous species, not the total concentrations The overall potential of the reaction will depend on several factors: the ionic strength; the temperature; the degree of hydroxide complexation of the metal; the presence of complexing agents for the metal; side reactions of any ligands with other species; and the limits on the free metal ion caused by solubility, such as the formation of corrosion product solids The rates and possibilities for some reactions can also be limited by a barrier to the diffusion of oxidants to the surfaces of 17.8 CHAPTER SEVENTEEN the materials, where “fresh” metal is available to oxidize Anodic and cathodic electrode areas may be induced by various conditions, some because of the characteristics of the metal, and some because of the character of the water at the boundary surface Impurities in the metal, sediment accumulations, adherent bacterial slimes, and accumulations of the products of corrosion are all related in some way to the development of electrode areas that can enable the operation of corrosion circuits Corrosion Products on Pipe Surfaces Metal surfaces may be protected either by their being “immune” or by rendering them “passive.” If a metal is protected by immunity, the metal is thermodynamically stable, and is therefore incorrodible (Pourbaix, 1973) For some metals, such as copper, this can occur in groundwaters that are somewhat anoxic (Lytle et al., 1998) Sometimes, this region of electrochemical behavior is only possible when water itself is not chemically stable, so it is only encountered in potable water systems when the consumption of externally supplied energy (cathodic protection) occurs Passivation occurs when the metal is not stable, but becomes protected by a stable film The protection can be perfect or (more usually) imperfect, depending upon whether the film effectively shields the metal from contact with the solution (Pourbaix, 1973) True passivation films must satisfy several requirements to effectively limit corrosion Particularly, they must be electrically conductive, mechanically stable (neither flaking nor cracking), and continuous Analysis of corrosion problems is complicated by the variety of chemical reactions that take place across the surface For example, consider the reactions at an iron or steel surface, in water where oxygen is the only oxidant, and aqueous iron complexation is negligible The primary reaction occurring at the anodic sites is: Fe(s) → Fe2+ + 2e− (17.18) The Fe2+ may then diffuse into the water, or it may undergo a number of secondary reactions Fe2+ + CO2− ⇔ FeCO3(s) (siderite) Fe2+ + 2OH− ⇔ Fe(OH)2(s) 2Fe + 1/2 O2 + 4OH ⇔ 2FeOOH(s) + H2O 2+ − (17.19) (17.20) (17.21) The hydrated ferric oxides that form from reactions similar to that shown in Eq 17.20 are reddish and, under some conditions, may be transported to the consumer’s tap Tertiary reactions may also occur at the surface Possibilities are: 2FeCO3(s) + 1/2 O2 + H2O ⇔ 2FeOOH(s) + 2CO2 3FeCO3(s) + 1/2 O2 ⇔ Fe3O4(s) (magnetite) + CO2 (17.22) (17.23) Reactions such as Eqs 17.21 to 17.23 can reduce the rate of oxygen diffusing to the anode, thus the formation of oxygen concentration cells Other reactions that affect corrosion may take place, depending on the composition of the water and the type of metal At the same time as the anodic reactions are taking place, a variety of cathodic reactions may be occurring Perhaps the most common in drinking water distribution systems is the acceptance of electrons by O2 e− + 1/4 O2 + 1/2 H2O ⇔ OH− (17.24) INTERNAL CORROSION AND DEPOSITION CONTROL 17.9 This reaction causes an increase in pH near the cathode and triggers the following additional reactions: OH− + HCO3− ⇔ CO32− + H2O (17.25) Ca + CO ⇔ CaCO3(s) (17.26) 2+ 2− These reactions can cause CaCO3(s) to precipitate from some waters in which the bulk solutions are undersaturated with this solid, because the pH increase in the vicinity of the cathode forms enough CO32− to cause supersaturation with respect to CaCO3(s) Several studies have shown that the pH at the surface of pipe can be significantly different from that in the bulk solution (Snoeyink and Wagner, 1996), although many studies reporting extremely high pHs at the surface have neglected to adequately include consideration of buffering by the carbonate system in the water, and the possible role of solids such as CaCO3(s) and Mg(OH)2(s) (Dexter and Lin, 1992; Lewandowski, Dickinson, and Lee, 1992; Watkins and Davies, 1987) The deposits that form on pipe surfaces may be (1) a mixture of corrosion products that depend both on the type of metal that is corroding and the composition of the water solution [e.g., FeCO3(s), Fe3O4(s), FeOOH(s), Pb3(CO3)2(OH)2(s), Zn5(CO3)2(OH)6(s)]; (2) precipitates that form because of pH changes that accompany corrosion [e.g., CaCO3(s)]; (3) precipitates that form because the water entering the system is supersaturated [e.g., CaCO3(s), SiO2(s), Al(OH)3(s), MnO2(s)]; and (4) precipitates or coatings that form by reaction of components of inhibitors, such as silicates or phosphates with the pipe materials (e.g., lead or iron) The nature of the scales or deposits that form on metals is very important because of the effect that these scales have on the corrosion rate The formation of scales, such as CaCO3(s) and iron carbonates on corroding iron or steel, are normally thicker and have higher porosity than the passivating films Deposits and scales not decrease the corrosion rate as much as true oxide films do, and the same corrosion current-potential relationship for passivating films does not occur for such scales The complex interactions can be illustrated by the case with steel corrosion Scale formation on steel by minerals such as calcium carbonate reduces the corrosion rate by decreasing the rate of oxygen transport to the metal surface, thereby decreasing the rate of the cathodic reaction Passivating iron oxide films on steel cause an anodic-controlled corrosion reaction A very long time, from many months to years, may be required for the corrosion rate of iron and steel to stabilize because of the complex nature of the scales.A much shorter time may be sufficient for other metals If a scale reduces the rate of corrosion, it is said to be a protecting scale; if it does not, it is called nonprotecting The importance of scale is also demonstrated by the phenomenon of erosion corrosion, observed at points in the distribution system or in domestic plumbing systems where a high-flow velocity or an abrupt change in direction of flow exists The more intense corrosion that often is observed at such locations can be attributed to the abrasive action of the fluid (caused by turbulence, suspended solids, and so forth) that scours away or damages the scale, and to the velocity of flow that carries away corrosion products before they precipitate and that facilitates transport of corrosion reactants more efficiently (Snoeyink and Wagner, 1996) Changes in water treatment or source water chemistry over time can produce successive layers of new solid phases, remove or change the nature of previously existing deposits, or both Figure 17.2 shows an example of the complex nature of scale on a cast-iron distribution pipe (Singley et al., 1985; Benjamin, Sontheimer, and Leroy, 1996) Scales of similar chemical composition can have a significantly different impact on corrosion and metal protection because properties such as uniformity, 17.10 CHAPTER SEVENTEEN FIGURE 17.2 Schematic of scale on a cast-iron distribution pipe, showing complex layered structure (Source: Internal Corrosion of Water Distribution Systems, 2nd ed., American Water Works Association Research Foundation, Denver, CO, 1996.) adherability, and permeability to oxidants can vary depending upon such factors as trace impurities, presence of certain organics, temperature of deposition, length of time of formation, and so forth Scales that form on pipes may have deleterious effects in addition to the beneficial effect of protecting the metal from rapid corrosion or limiting the levels of toxic metals (such as lead) in solution Water quality should be controlled so that the scale is protective but as thin as possible, because as bulk of scale increases, the capacity of the main to carry water is reduced The formation of uneven deposits such as tubercles increases the roughness of the pipe surface, reducing the ability of the mains to carry water, and may provide shelters for the growth of microorganisms To properly interpret field and laboratory data from corrosion control studies, it is important to understand that there may be significantly different reactions occurring between the water constituents and the surfaces of “new” pipes compared with “old” pipes Conceptually, this is illustrated in Figure 17.3 for lead pipe On the new surface [Figure 17.3(a)], the full corrosion reaction can occur, with oxidation of lead followed by the development of a passivating film Once the film is sufficiently developed [Figure 17.3(b)], the oxidants in the water no longer can directly contact the metal of the pipe material itself Therefore, the oxidation step will not occur, and metal release will become a function of the physical adherence or the solubility of the surface deposit, unless water conditions become anoxic and the metal(s) in the surface deposit become electrochemically reduced Thus, corrosion inhibitor chemicals that stifle reactions occurring at cathodic surface sites may appear much better in tests using new metal surfaces than they may operate when applied to distribution system pipes covered by thick scales or corrosion deposits With well-developed surface deposits present, the solubility and surface sorption chemistry of the existing scales is much more important in developing water treatment targets than predictions based on the pure corrosion chemistry of the metal INTERNAL CORROSION AND DEPOSITION CONTROL 17.95 water to their customers and (2) to meet regulatory requirements For the purpose of holistically controlling the corrosion of the multiple materials present in the water system, water samples should be targeted toward the entry points and throughout the distribution system Any “mixing” or hydraulically isolated zones should be covered by samples targeted in time and space to determine water quality changes that might result from corrosion (e.g., iron, calcium, or aluminum release, turbidity, pH changes, disinfectant residual, corrosion inhibitor, or dissolved oxygen depletion), microbial activity (pH change from nitrification, sulfate reduction), or other causes It is also useful to determine if the metal release from corrosion comprises mainly soluble or particulate chemical species To represent conditions at the customer’s tap, “standing” samples should be taken from an interior faucet in which the water has remained for a specific number of hours (i.e., overnight) The age of the plumbing in which the sampled water is in contact can be very important information, especially for copper The stagnation curves representing metal levels versus time are very complex, and differ somewhat for different metals For example, copper levels have been observed to increase for as much as 72 hours or more, in the continued presence of chlorine or dissolved oxygen (Schock, Lytle, and Clement, 1995a,b; Lytle and Schock, 1996, 1997; Werner, Groß, and Sontheimer, 1994) However, in the absence of oxidants, or after the depletion of oxidants, the copper levels have often been seen to actually peak and then go down The sample should be collected as soon as the tap is open If the contribution of the faucet to metal levels is of interest, a small sample must be collected (i.e., 60 to 125 mL) in addition to samples from the rest of the plumbing system Larger volumes are necessary to include the pipe contribution Sample volumes may be thought of to represent linear plumbing distance, so contamination sources can be characterized by carefully planned sequences of samples of specific volumes appropriate to the given study (Lytle, Schock, and Sorg, 1994; Wysock, Schock, and Eastman, 1991) Many important decisions are likely to be made based on the sampling and chemical analyses performed by a utility Therefore, care must be taken during the sampling and analysis to obtain the best data Handling of samples for pH, alkalinity, and CO2 analyses often requires special precautions (Schock and Schock, 1982) Samples should be collected without adding air and with minimal agitation, as air tends to remove CO2 and also affects the oxygen content in the sample To collect a sample without additional air and to minimize exchange of volatile gases (such as CO2 in waters that are frequently out of equilibrium with the atmosphere), fill the sample container to the top so that a convex dome is formed at the opening and no bubbles are present If possible, the sample bottle should be filled below the surface of the water using tubing, so that the water is not contaminated by the faucet material Cap the sample bottle as soon as possible For general purposes, high-density linear polyethylene bottles are very suitable for metals and most other water quality constituents For pH measurements off-site, the bottles should be glass or a material assured to be impermeable, especially toward carbon dioxide, and to a lesser extent, oxygen Most plastic bottle manufacturers can provide tables of gas permeabilities for the plastics used in their products, which will be helpful for the selection of the proper bottles for the purpose When available, a direct analysis of DIC concentration (or TIC concentration, if unfiltered) is generally more accurate than computing DIC from pH and alkalinity measurements, unless particular care is taken in the analyses and sophisticated equivalence point and pH stabilization point techniques are employed (Schock and George, 1991; Schock and Lytle, 1994) Glass, or other material impermeable to CO2, must be used for the container, and the cap must allow no air space 17.96 CHAPTER SEVENTEEN To determine if metal release is primarily particulate or dissolved, frequently filtrations are used This is a deceptively difficult procedure, as there are numerous opportunities for erroneous data and biases that can significantly affect conclusions Sorption losses are especially acute on glass materials and with cellulose acetate membranes, so polycarbonate, teflon, or stainless-steel apparatus and filter supports are preferable Polycarbonate filter membranes have generally been shown to cause less bias than membranes made of most other materials Some cleaning and filter apparatus preparation precautions have been described in some studies (Schock, Lytle, and Clement, 1995b; Schock and Gardels, 1983) Before using data from sample filtration, a careful laboratory and field test evaluation should be conducted of all materials used and all steps in the procedure to determine if there is a potential for losses of dissolved constituents, or contamination from the materials The constituents that should be analyzed in a thorough corrosion-monitoring program depend to a large extent on the materials present in the system’s distribution, service, and household plumbing lines Table 17.6 summarizes parameters recommended to be analyzed in a thorough corrosion-monitoring program Temperature and pH should be measured in situ (in the field) with proper precautions for atmospheric CO2 exchange Frequency of analysis depends on the extent of the corrosion problems experienced in the system, the degree of variability in source and finished-water quality, the type of treatment and corrosion control practiced by the water utility, and cost considerations When phosphates or silicates are added to the water, samples should be collected at the far reaches of the system and analyzed for polyphosphates, orthophosphates, and silicate, as appropriate If no residual phosphate or silicate is TABLE 17.6 Recommended Analyses for a Thorough Corrosion Monitoring Program General parameters for all investigations In situ measurements Dissolved gases, oxidants Parameters required to calculate CaCO3-based indices Parameters for A-C pipe Background parameters for metal pipe Iron or steel pipe Lead pipe or lead-based solder Copper pipe Galvanized iron pipe Brass (faucets and valves) All metal pipes Corrosion “inhibitor” constituents (all metal pipes) pH, temperature, CO2 if low-pH groundwater Oxygen, hydrogen sulfide,* free chlorine residual, total chlorine residuals (if ammonia present or used) Calcium, total hardness (or magnesium), alkalinity (or DIC), total dissolved solids (or conductivity)† Add to general parameters Fiber count, iron, zinc, silica, polyphosphate, aluminum, manganese Add to general parameters Iron, chloride, sulfate Lead, copper, chloride, sulfate Copper, lead, sulfate, chloride Zinc, iron, cadmium, lead, chloride Zinc, copper, lead, sulfate Add to general parameters Orthophosphate, polyphosphate,‡ silica * It only needs to be analyzed when suspected † If a complete water analysis is done, these can be neglected because ionic strength can be directly computed ‡ For most potable waters; derived from analyzing total phosphate and subtracting orthophosphate INTERNAL CORROSION AND DEPOSITION CONTROL 17.97 found, the feed rate should be increased When calcium carbonate precipitation is practiced, monitoring of the parameters necessary to compute the LSI or the CCPP in the far reaches of the system is also necessary Continual monitoring of pH and temperature are important, because they are so interdependent An important quality assurance practice in corrosion studies is to frequent complete chemical analyses of the finished water and in many locations throughout the distribution system, and carefully examine the data for ion balance errors and internal consistency of trends and interrelationships among constituents that should follow known behavior and concentration patterns Chemical equilibrium modeling computer programs provide a very good tool to this evaluation (APHA-AWWA-WEF, 1992, 1995; Parkhurst, Thorstenson, and Plummer, 1980; Allison, Brown, and Novo-Gradac, 1991; Schecher and McAvoy, 1994a,b, 1998) When corrosion control studies obtain good baseline data and have a comprehensive, well-designed monitoring program in place throughout their corrosion control effort, the data are invaluable to other utilities and corrosion scientists They can then work with a larger body of knowledge of drinking water chemistry and treatment, and implement future corrosion control and public health protection strategies much more effectively and efficiently ACKNOWLEDGMENTS The author gratefully acknowledges the extensive contributions of all of the writers of the AWWA publication Corrosion Control for Operators, the AWWARF/DVGWForschungsstelle/TZW manuals Internal Corrosion of Water Distribution Systems, and the late Dr T E Larson, whose works were heavily cited and directly included in this chapter The author is 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CORROSION AND DEPOSITION CONTROL 17. 7 O2 + 2H2O + 4e− ⇔ 4OH− (17. 7) HOCl + H+ + e− ⇔ (1/2) Cl2 (aq) + H2O (17. 8) ⁄2 Cl2 (aq) + e ⇔ Cl − − (17. 9) Equations 17. 8 and 17. 9 can be combined to yield... Kolle, and Snoeyink, 1981) Figures 17. 13 and 17. 14 are revised potential-pH diagrams based on the same species as were used to create Figures 17. 4, 17. 5, 17. 7, and 17. 8 Areas of immunity and passivation... at 25°C, is {Fe2+}{Cl−} 0.0591 EFe2+/OCl − = E 0Fe2+/OCl − − ᎏ log ᎏᎏ {OCl−} {H+}2 (17. 17) Note that in Eq 17. 17, the activities are for the free aqueous species, not the total concentrations The