I N T E R N A C O17.O S I O N A N D D E P O S I T I O N CCHAPTER S E V E N T E E N RR ONTROL INTERNALLCORROSION AND DEPOSITION CONTROL R SEVENTEEN CHAPTE

17. 17.

C H A P T E R 1 7 ___________

INTERNAL CORROSION AND DEPOSITION CONTROL 1Michael R. Schock, ChemistU.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 metalscopper, iron, and zinccause 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 (Rompr et al., 1996; Schreppel, Frederickson, and Geiss, 1997; Camper, 1997; Kin, Lu, and Lvy, 1996; Norton et al., 1995; Olson, 1996).

1 The views expressed in this paper are those of the author and do not necessarily reflect the views or 17 . Agency. policies of the U.S. Environmental Protection1

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: 1.Increased pumping costs caused by tuberculation and hydraulic friction 2.Loss of water and water pressure caused by leaks 3.Water damage to the dwelling, requiring that pipes and fittings be replaced 4.Replacing hot water heaters 5.Customer complaints of colored water, stains, or bad taste, for which the response may be expensive both in terms of money and public relations 6.Increased wastewater and sludge treatment and disposal costs 7.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 waters 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-

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 with-out their own potential problems. CORROSION, PASSI VATION, AND IMMUNITYElectrochemical 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 ~ Mez+ + 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

Primary contaminants from pipe

Good overall corrosion resistance; subject to corrosive Copper and possibly iron, zinc, tin, antimony, arsenic, attack from high flow velocities, soft water, chlorine, cadmium, and lead from associated pipes and dissolved oxygen, low pH, and high inorganic carbon solder. levels (alkalinities). May be prone to pitting failures. Lead Corrodes in soft water with pH 8.3 with high ratio of chloride to carbonate hardness. Conditions causing mechanical failure may not directly correspond to those promoting contaminant leaching.* 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.

INTERNAL CORROSION AND DEPOSITION CONTROL

17.

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.

1 7 . 1 Adjoining anodes and cathodes during the corrosion of iron in acidic solution. ( S o u r c e : W a t e r C h e m i s t r y , V. L. Snoeyink and D. Jenkins. Copyright 1980, John Wiley & Sons, Inc. Reprinted by permission of John Wiley & Sons, Inc.) FIGURE

17.

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 ~ E M e / M e z+ = E 0M e / Me

z+

-

ln { M ezF

z+}

(17.2)

where E Me/ Mez + = the potential (volts) E 0 Me/ Me = the standard potential (volts), a constant that can be obtained from tables of standard reduction potentials { } denote activity of the ion M e z + 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 + + 2 e - ~ H 2 (g) E 0 = 0.00 (17.3) takes place, andz+

reactants and products are assumed by thermodynamic convention to equal 1 (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 ~G r 0 = -z F E , where ~G r 0 is the free-energy 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 + z e - ~ red (17.4) where ox and red indicate oxidized and reduced species, respectively (Snoeyink and Jenkins, 1980; Stumm and Morgan, 1981, 1996):R T {red} ~ ~ ln (17.5) zF {ox} 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 ~ E r e d / o x = E 0r e d / o x log Q Ered/ox= E0red/ox -

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

O 2+ 2H 2O + 4 e- ~ 4OH HOCl + H + + e- ~ (1/2) Cl 2 (aq) + H 2O1

(17.7) (17.8) (17.9) (17.10)

Cl 2 (aq) + e- ~ Cl 2

Equations 17.8 and 17.9 can be combined to yield the net reaction: HOCl + H + + 2 e- ~ Cl- + H 2O

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 ({H 2O} = 1), then 0.0591 {OH - } ~~ EO /OH- = E0 O /OH- log (17.11) 4 {O 2}22

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: 0.0591 {Cl - } (17.12) ~ ~~ EHOCl/Cl - = E0 HOCl/Cl - log 2 {HOCl}{H + } 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) +O 2 + 2H 2O ~ 2Pb 2+ + 4OH 2Fe (metal) +O 2 + 2H 2O ~ 2Fe 2+ + 4OH Pb (metal) + HOCl + H + ~ Pb 2+ + Cl- + H 2O Fe (metal) + OCl + 2H + ~ Fe2+

(17.13) (17.14) (17.15) (17.16)

+ Cl- + H 2O

As an example, the overall Nernst expression for the last preceding equation, at 25 C, is (17.17) EFe /OCl - = E0 Fe /OCl - - 0.0591 log {Fe 2+}{Cl-} ~ ~~ 2 {OCl -} {H +} 2 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 of2+ 2+

17.

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) ~ Fe 2+ + 2 e 2+

(17.18)

The Fe may then diffuse into the water, or it may undergo a number of secondary reactions. Fe 2+ + CO 3 2- ~ FeCO 3(s) (siderite) Fe2+

(17.19) (17.20) (17.21)

+ 2OH ~ Fe(OH) 2(s)

-

2Fe 2+ +1/2 O 2 + 4OH - ~ 2FeOOH(s) + H 2O

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 consumers tap. Tertiary reactions may also occur at the surface. Possibilities are: 2FeCO 3(s) + 1/2 O 2+ H 2O ~ 2FeOOH(s) + 2CO 2 3FeCO 3(s) + 1/2 O 2 ~ Fe 3O 4(s) (magnetite) + CO 2 (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 O 2. e- + 1/4 O 2 + 1/2 H 2O ~ OH(17.24)

INTERNAL CORROSION AND DEPOSITION CONTROL

17.

This reaction causes an increase in pH near the cathode and triggers the following additional reactions: OH- + HCO 3 - ~ CO 3 2- + H 2O (17.25) Ca 2+ + CO 3 2- ~ CaCO 3(s) (17.26) These reactions can cause CaCO 3(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 CO 3 2- to cause supersaturation with respect to CaCO 3(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 CaCO 3(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., FeCO 3(s), Fe 3O 4(s), FeOOH(s), Pb 3(CO 3) 2(OH) 2(s), Zn 5(CO 3) 2(OH) 6(s)]; (2) precipitates that form because of pH changes that accompany corrosion [e.g., CaCO 3(s)]; (3) precipitates that form because the water entering the system is supersaturated [e.g., CaCO 3(s), SiO 2(s), Al(OH) 3(s), MnO 2(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 CaCO 3(s) and iron carbonates on corroding iron or steel, are normally thicker and have higher porosity than the passivating films. Deposits and scales do not decrease the corrosion rate as much as true oxide films do, and the same corro sion 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 corro sion, 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.

CHAPTER SEVENTEEN

1 7 . 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.) FIGURE

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 occur-ring 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.

New Lead Surface Both anodic and cathodic sites Example Anodic site reactions Pb 5Pb2+

Old Lead Surface No anodic and cathodic sites

Pb3 4

2+

+ 2e4 3

Pb3 (CO3 )2 (OH)2 + 2PO4 3 5[Pb (CO ) (OH) ] + 9PO3 3 2 2 3 4

Pb3 (PO4 )2 + 2CO35 4 3

2

+ 2OH2 3

+ 3PO

+ H2 O

Pb (PO ) OH + H+5

3[Pb (PO ) OH] + 1 0CO

+ 7OH

Example Cathodic site reactions HOCI + 2e 5Zn2+

CI + OH2 3

+ 2CO

+ 6OH

Zn (CO ) (OH) Ca5 3 2 6

2+

+ HCO

3

+ OH CaCO + H2 O3

(a)FIGURE 17.3

(b)

Schematic representation of different surface reactions and their relation to passivation between (a) new and (b) old lead pipe surfaces. Corrosion Kinetics

A three-step process is involved in governing the rate of corrosion of pipe: (1) transport of dissolved reactants to the metal surface, (2) electron transfer at the surface, and (3) transport of dissolved products from the reaction site (Trussell, 1985). When either or both of the transport steps are the slowest, rate-limiting step, the corrosion reaction is said to be under transport control. When the transfer of electrons at the metal surface is rate-limiting, the reaction is said to be under activation control. The formation of solid natural protective scales that inhibit transport are often an important factor in transport control. This section will only present an overview of the concepts involved in the rates of corrosion reactions in potable water systems. Numerous reference articles and texts exist that present a detailed development of the theories that are the basis for many direct electrochemical rate-measuring techniques. Corrosion is often described in terms of numerous tiny galvanic cells on the surface of the corroding metal. Such localized anodes and cathodes as those described are not fixed on the surface, but are statistically distributed on the exposed metal over space and time. The electrochemical potential of the surface is determined by the mixed contributions to potential of both the cathodic and anodic reactions, averaged over time and over the surface area. Both the individual anodic and cathodic half-reactions are reversible and occur in both directions at the same time. When the electrode is at its equilibrium, the rates of reaction in both the cathodic and anodic half-cells are equal. Given the thermodynamic basis for corrosion described above, and the body of knowledge about kinetic factors that affect the rate of corrosion of metals, several properties of the water passing through a pipe or device that influence the rate of corrosion can be identified. Some of the water-related properties are: (1) concentration of dissolved oxygen, (2) pH, (3) temperature, (4) water velocity, (5) concentration and type of chlorine residual, (6) chloride and sulfate ion concentration, and (7) concentration of dissolved inorganic carbon (DIC) and calcium. These properties interrelate, and their effect depends on the plumbing material as well as the overall water quality. References specific to the type of plumbing situation (pipe, soldered joint, galvanic connection, faucet, or flow-control device) and the material of interest should be consulted for the most appropriate information on corrosion rate control. Some generalizations will be considered in a later section.

17.

CHAPTER SEVENTEEN

Solubility Diagrams

The solubility of passivating films on the pipe surface is the most important factor in determining whether a given water quality can meet many drinking water regulations that are based on health effects of ingested metals. Asbestos fibers are frequently released into the water as a result of the dissolution of the cement matrix originally holding them. Solubility places a lower limit on the level to which metals can be controlled by modifying water chemistry at the treatment plant. Solubility reflects an ideal equilibrium condition, does not correspond directly to tap water levels, and should not be expected to (Schock, 1990a; Britton and Richards, 1981). Other factors are important, such as the physical location of the plumbing materials relative to the sample collected, the release of particulates from the deposits on the pipes, the rate of the chemical reactions that affect the mobilization of the metals relative to the standing time of the water before collection, and many other variables. However, it is widely believed that trends in the response of tap water metal levels to changes in key water chemistry variables usually follow the predictions of solubility models, and that has been recently verified in principle by some recent surveys of over 2500 U.S. utilities of all sizes (Edwards, Jacobs, and Dodrill, 1999). To display solubility relationships in a relatively simple two-dimensional manner, the total solubility of a constituent is plotted as a function of a master variable (Snoeyink and Jenkins, 1980; Stumm and Morgan, 1981, 1996; Bard, 1966). Other solution parameters can affect solubility such as the ionic strength, or the concentration of dissolved species (ligands) that can form coordination compounds or complexes with the metal. The solubility diagram does not necessarily give the information needed to minimize the rate of corrosion, because rate is affected by kinetic parameters that depend on the relative rates of the oxidation, dissolution, diffusion, and precipitation reactions. It also cannot predict the ability of a pipe coating to adhere to the pipe surface, or the permeability of a coating to oxidants from the water solution or pitting agents. It does, however, give important information for estimating how water quality affects attainment of maximum contaminant levels (MCLs) or action levels (ALs) for drinking water, the potential for precipitating passivation films on pipe surfaces, or the deposition of other solids important in water treatment, such as calcium carbonate, octacalcium phosphate, and aluminum or ferric hydroxide. To construct this type of diagram, an aqueous mass balance equation must be written for the metal (or other constituent of interest), with the total solubility (ST ) as the unknown. The mass balance expression should include the concentration of the uncomplexed species (free-metal ion), along with all complexes to be included in the model. For lead (II), the simple relationship for the free ion, plus only the hydroxide and carbonate complexes (the simplest system) is:ST , P b ( I I ) , C O = [Pb ] + [Pb(OH) 2 ] + [Pb(OH) 3 ] + [Pb(OH) 4 ] + 2[Pb 2 (OH) ] +2+ 0 23+ 3

3[Pb 3 (OH) 4 ] + 4[Pb 4 (OH) 4 ] + 6[Pb 6 (OH) 8 ] + [PbHCO 3 ]2+ 4+ 4+ +

+ [PbCO 3 ] + [Pb(CO 3 ) 2 ]0 2-

(17.27)

If the concentration of a complex is going to be relatively negligible compared with the total solubility across the pH range of interest, it can be excluded from the model for simplicity. Frequently, only four or five species are significant in a system. An example of two such complete mass balance expressions are the following for lead (Schock, 1980, 1981b) and copper (Schock, Lytle, and Clement, 1995a,b) systems

INTERNAL CORROSION AND DEPOSITION CONTROL

17.

only containing carbonate species in addition to hydrolysis species, which is the minimum system composition for potable waters. 1,1 [Pb 2 + ] 1 2[Pb 2 + ] 1 3[Pb 2 + ] 1 4[Pb 2 + ] ~ ~, , ,

~ S T,Pb(II),CO3 = [Pb 2+ ] +___________+

~

~

~

~

~

+___________+___________

[H + ] 2, 1[Pb ]~ ~2+ 2

[H + ] 2 3, 4[Pb ]2+ 3

[H + ] 3 4, 4[Pb ]~ ~2+ 4

[H + ] 4 6, 8[Pb 2 + ] 6 + 2_____+ 3 + 6___________

~

~

~

~

+4

[H + ]+ 1,1,1[Pb 2+

[H + ] 4] [H+ 2-

[H + ] 4 2+

[H + ] 8] [CO 3 2-]

] [CO 3 ] + 1,0,1[Pb2+

(17.28)

1, 1[Cu ]~ ~ ~ ~

2+

1, 2[Cu ]~ ~

+ 1,0,2[Pb 2+ ] [CO 32- ] 2 1, 3[Cu 2 + ] 1, 4[Cu 2 + ]~ ~ ST,Cu(II),CO3 =

[Cu 2+ ] +_____________________+ +___________+__________ [H + ] [H + ] 2 [H + ] 3 [H + ] 4 2+ 2 2+ 3 2 2[Cu ] 3 4[Cu ] +2 +3 + 1,1,1[Cu 2+ ] [H + ] [CO 32- ] ~ ~ (17.29), ,

~

~

[H + ] 2 [H + ] 4 In these equations, m,h,c represents the formation constant for the complex having a stoichiometry of m metal ions, h hydrogen ions, and c carbonate ions, corrected for ionic strength and temperature (Stumm and Morgan, 1981, 1996). To obtain the predicted solubility, the solubility constant expression for each solid of interest is rearranged and solved to isolate the free species concentration, and is substituted into each term of the mass balance equation (e.g., Eq. 17.28 or 17.29). Sometimes, an iterative technique is used to simultaneously solve several related mass balances, depending upon exactly the type of diagrams desired and assumptions about the availability of components involved in the precipitation and dissolution reactions in the system. A diagram is then constructed for each solid, and the curves are superimposed.The points of minimum solubility are then connected, giving the final diagram. This procedure is discussed in more detail by Schock (1980, 1981b) for lead (II); Schock, Lytle, and Clement for copper (1995a,b); and Snoeyink and Jenkins (1980) for iron (II). When additional aqueous species are present, such as orthophosphate or sulfate, they are just added to the basic mass balance expressions, such as those shown above. For example, the contribution to lead(II) solubility from aqueous orthophosphate species is represented by: ST, Pb(II), PO4 = [PbHPO 40 ] + [PbH 2PO 4+ ] (17.30) So, the total lead

solubility ST, Pb(II) = ST, Pb(II),CO + ST, Pb(II), PO , or3 4

ST, Pb(II) = [Pb 2+ ] + [Pb(OH) 2

0] +

[Pb(OH) 3 -] + [Pb(OH) 4 2-] + 2[Pb 2(OH) 3+ ] +

3[Pb 3(OH) 4 2+] + 4[Pb 4(OH) 4 4+] + 6[Pb 6(OH) 8 4+] + [PbHCO 3 +]+ [PbCO 3 0] + [Pb(CO 3) 22-] +

[PbH 2PO 4 +] + [PbHPO 4 0]

(17.31)

Similarly, other ligands (such as chloride and sulfate) can simply be added to the expressions.

17.

CHAPTER SEVENTEEN

Because these diagrams are inherently two-dimensional (solubility on the y-axis, pH on the x-axis), there are often additional variables that have an important impact on solubility (such as temperature, ionic strength, carbonate concentration or alkalinity, and orthophosphate concentration), these diagrams must display solubility and species concentration lines for fixed concentrations or values of the other variables. For instance, Figure 17.4 displays a solubility diagram for lead in the system

A solubility diagram showing dissolved lead (II) species in equilibrium with lead (II) solids in a pure system containing 3 mg C/L DIC (2.5 10- 4 mol/L) at 25C and I = 0.005 mol/L. ( S o u r c e : data from Schock and Wagner, 1985.)FIGURE 17.4

INTERNAL CORROSION AND DEPOSITION CONTROL

17.

containing only carbonate species and water, showing the important aqueous species and the stability domains of two lead solids. In this figure, the dissolved inorganic carbon (DIC) concentration was fixed at 0.00025 mol/L (3 mg C/L), with a temperature of 25 C and an assumed ionic strength of 0.005. Figure 17.5 shows the same basic system, but with a total DIC concentration of 0.0025 mol/L (30 mg C/L). In both diagrams, the implicit assumption was also made that the solution redox potential was not high enough to cause the formation of Pb(IV) aqueous or solid

A solubility diagram showing dissolved lead (II) species in equilibrium with lead (II) solids in a pure system containing 30 mg C/L DIC (2.5 10- 3 mol/L) at 25C and I = 0.005 mol/L. ( S o u r c e : data from Schock and Wagner, 1985.)FIGURE 17.5

17.

CHAPTER SEVENTEEN

species, or a different diagram would apply for the higher oxidation state species of lead. Additionally, when metal solubility becomes high, care must be taken to ensure that computational constraints on the system, such as the fixed ionic strength and the total ligand concentration (mass balance), are not violated by the presence of high concentrations of dissolved metal species. Note that if the same complex formation constants were used, but a less soluble constant was used for PbCO 3(s) (cerussite) solubility, the simple carbonate solid would be predicted to be stable over a wider pH range. Likewise, a less soluble constant for Pb 3(CO 3) 2(OH) 2 (hydrocerussite) would expand its stability field relative to that of cerussite. Analogous information can be obtained through the careful construction and study of solubility diagrams for other metals, such as copper and zinc. When constructing solubility diagrams for any metal, the selection of solid and aqueous species must truly represent the system to be modeled, or very erroneous conclusions can result. For example, the aragonite form of calcium carbonate is frequently found in deposits formed in systems having galvanized pipe, rather than the calcite form. Also, the ferric iron deposits formed in mains are frequently a relatively soluble hydroxide or oxyhydroxide form [Fe(OH) 3 or FeOOH] rather than an ordered form such as hematite (Fe 2O 3). To go along with this concept, the equilibrium constants must also be accurate to give realistic concentration estimates, and knowledge of changes in the equilibrium constants with temperature is essential, especially when projections of depositional tendency have to be made into hot or very cold water-piping systems. Critical evaluation of data appearing in handbooks and published papers is necessary to avoid using incorrect values, and occasionally review articles or major works by rigorous researchers can be consulted for reliable values. An important assumption behind the diagrams is that the system must reach thermodynamic equilibrium for the calculations to be truly valid, unless kinetic factors are incorporated into the model. Sometimes improvements can be made in predictions by using metastable species in the calculations, although it is not thermodynamically rigorous to do so. Metastable solids have been found to govern copper (Schock, Lytle, and Clement, 1994, 1995a,b; Edwards, Meyer, and Schock, 1996) and sometimes lead (Schock, Wagner, and Oliphant, 1996) levels in drinking water.

1 7 . 6 Solubility and saturation. A schematic solubility diagram showing concentration ranges versus pH for supersaturated, metastable, saturated, and undersaturated solutions. ( S o u r c e : A q u a t i c C h e m i s t r y , W. Stumm and J. J. Morgan. Copyright 1980, John Wiley & Sons, Inc. Reprinted by permission of John Wiley & Sons, Inc.) FIGURE

INTERNAL CORROSION AND DEPOSITION CONTROL

17.

Stumm and Morgan have discussed the formation of precipitates and how this is related in a conceptual way to solubility diagrams, as illustrated by Figure 17.6 (Stumm and Morgan, 1981,1996). They define an active form of a compound as one that is a very fine crystalline precipitate with a disordered lattice. It is generally the type of precipitate formed incipiently from strongly oversaturated solutions. Such an active precipitate may persist in metastable equilibrium with the solution and may convert (age) slowly into a more stable, inactive form. Measurements of the solubility of active forms give solubility products that are higher than those of the inactive forms. The formation of some of the iron hydroxide or oxyhydroxide solids in pipe deposits mentioned previously provides an example of this phenomenon. Hydroxides and sulfides often occur in amorphous and several crystalline modifications. Amorphous solids may be either active or inactive. Initially formed amorphous precipitates or active forms of unstable crystalline modifications may undergo two kinds of changes during aging. Either the active form of the unstable modification becomes inactive or a more stable modification is formed. With amorphous compounds, deactivation may be accompanied by condensation or dehydration. When several of the processes take place together, nonhomogeneous solids can be formed upon aging. Similar phenomena can occur with basic carbonates, such as a transition from one form to another with changes in pH or DIC over time. Rather than construct a different detailed diagram for each level of a secondary variable, such as DIC, frequently multiple lines representing the different levels of this variable are added to a single diagram, and the aqueous species are omitted. For metals (such as zinc, copper, and lead) with solubilities that tend to be influenced by complexation, the expansion of the diagram to include a third dimension is often useful. For a qualitative, conceptual understanding, a three-dimensional (3-D) surface can be constructed that can show multiple trends in a complex system at a glance. Figure 17.7 shows the two troughs representing PbCO 3(s) and Pb 3(CO 3)(OH) 2(s), the different trend of solubility with DIC concentration for each solid, and the dis-

Three-dimensional representation of the effect of DIC concentration and pH on lead (II) solubility, assuming ionic strength = 0.005, and temperature = 25C. ( S o u r c e : Schock and Wagner, 1985.)FIGURE 17.7

17.

CHAPTER SEVENTEEN

tinct solubility minimum in the system (Schock, 1980, 1981b; Schock, Wagner, and Oliphant, 1996; Schock and Gardels, 1983). To obtain a better quantitative estimate of trends of solubility resulting from interrelationships between two major variables, diagrams such as Figure 17.8 can be constructed. Operating on the same principles as topographic maps, such contour diagrams present a map view of surfaces such as Figure 17.7. The diagrams are derived by interpolating levels of constant concentration within a three-dimensional array of computed solubilities at different combinations of the other two master variables (e.g., pH and DIC). Several different mathematical algorithms are widely used in commercially available computer software, and considerable care must be exercised in selecting algorithms and data point spacing to prevent the creation of erroneous contouring artifacts. Rapid changes in solubility with respect to a master variable (here, pH or DIC) are shown by closely spaced contour lines. A series of these diagrams at levels of a third master variable (such as orthophosphate concentration, temperature, and so forth) can be useful to help display multiple interactions with a minimum of diagrams. They also enable a direct reading of estimated solubilities, without having to guess from an indirect perspective, such as with Figure 17.7. One problem with this type of diagram, as well as with the 3-D surface plots, is that most metals can undergo a change insolubility of three orders of magnitude or even more, over the range of conditions that might be reasonable for potable waters. Thus, logarithmic scales are often necessary for the metal concentrations, which can be somewhat confusing to read.

Contour diagram for lead (II) solubility assuming the formation of PbCO and Pb (CO ) (OH) , computed for I =0.02, and temperature = 25C. ( S o u r c e : equilibrium constant data from Schock and Wagner, 1985.)FIGURE 17.83 3 3 2 2

INTERNAL CORROSION AND DEPOSITION CONTROL

17.

Pourbaix or Potential/pH Diagrams

Using the Nernst equations for appropriate electrochemical half-reactions, it is possible to construct potential-pH diagrams, which are also called Eh-pH, or Pourbaix, diagrams. These diagrams have been popularized by Pourbaix and his coworkers in the corrosion field (Pourbaix, 1966, 1973; Obrecht and Pourbaix, 1967), by Garrels and Christ, and by Stumm and Morgan in geochemistry (Garrels and Christ, 1965; Stumm and Morgan, 1981). A similar type of diagram uses the concept of electron activity, pE, which is analogous to the concept of pH (Snoeyink and Jenkins, 1980; Stumm and Morgan, 1981). The Pourbaix diagrams include the occurrence of different insoluble corrosion products of the dissolved metal that limit the concentration of the free metal ion. These diagrams mainly give information about thermodynamically stable products under different conditions of electrochemical potential. The position of the boundaries of each region is also a function of the aqueous concentrations (activities) of ions that participate in the half-cell reactions. Potential-pH diagrams are particularly useful to study speciation in systems that could contain species of several possible valence states within the range of redox potential normally encompassed by drinking water, such as manganese, iron, arsenic, and copper. Obtaining an accurate estimate of the redox potential of the drinking water is usually an important limitation in using potential-pH diagrams. The diagrams are also useful to gauge the possible reliability of electrochemical corrosionrate measurement techniques. Measurement methods that rely on the imposition of a potential to the pipe surface may shift the pipe surface into the stability domain of a solid that would not normally form when freely corroding. The imposed potential might also serve to alter the nature of the surface phase, leading to an erroneous identification of the dominating corrosion or passivation reactions as the result of a surface compound analysis. the solubility versus pH plots discussed earlier. This relationship is shown schematically in Figure 17.9. If a conventional two-dimensional solubility diagram is considered as a vertical plane, a potential-pH diagram may be thought of as a slice that is perpendicular to the solubility plane, which cuts through the plane at a single concentration of the metal. In actuality, potential-pH diagrams are usually computed in terms of activities rather than concentrations, but the difference is usually not important for practical purposes. The activities of all aqueous and solid species must be fixed, while the electrical potential of the solution and the pH become the master variables. If a given water constituent (such as calcium) does not F I G U R E 1 7 . 9 Schematic relationship between potential pH and conventional solubility versus pH oxidize or reduce under meaningful physical conditions, no additional useful diagrams. A potential-pH diagram is related to information is gained (beyond

17.

CHAPTER SEVENTEEN

that directly available with a solubility diagram) by constructing a potential-pH diagram. Figure 17.10 is an example of a potential-pH diagram for iron in water (AWWARFDVGW, 1985). This particular diagram assumes Fe2O3 and Fe3O4 to be the solid phases that can control iron solubility. The diagram in Figure 17.10 shows that iron and water are never thermodynamically stable simultaneously, because the iron metal field (Fe) falls below the line where water is reduced to H 2 gas. At such a low electrode potential, the iron will not corrode (i.e., it is immune). The iron potential is reduced to the immune region, for example, by cathodic protection. To accomplish this, the iron must be coupled with another, more easily corrodible material, such as magnesium. At low pH (