A 21st-Century Perspective on Calcium Carbonate Formation in Potable Water Systems

Abstract

Formation of calcium carbonate (CaCO₃) has long been a concern in potable water systems. CaCO₃ formation is re-emerging as an important issue due to climate change, higher set-point temperatures in hot water systems, use of scaling and corrosion inhibitors, and self-repair of pipeline leaks. Ironically, actions we are taking to increase the lifespan of distribution systems via reduced corrosion (i.e., adding corrosion inhibitors) might have inadvertently worsened leaks and pipe lifespans, due to interference with natural self-repair from CaCO₃ precipitation. These changes in practice, coupled with knowledge gaps over CaCO₃ formation in water systems, make it important and timely to revisit this topic.

Introduction

Prediction, monitoring, and consequences of calcium carbonate (CaCO₃) formation have long been a key issue in potable water treatment. An early concern was that waters highly undersaturated with CaCO₃ would be corrosive to metallic and concrete pipe infrastructure (Tillmans and Heublein, 1912; Baylis, 1926; Langelier, 1936), and for waters highly supersaturated with CaCO₃, issues related to scaling (i.e., pipe clogging, head loss, higher heating costs) are the primary concerns (Garrett-Price et al., 1985). In the 80⁺ years since the landmark studies examining these issues were published, many changes have occurred that will renew interest in CaCO₃ scaling in potable water systems, including chemical corrosion control, changing water heater set-point temperatures, autogenous repair, global warming, biofilm growth, and erosion corrosion. This review is aimed at summarizing these changes and highlighting their practical implications for water utilities, regulators, building managers, and consumers (Fig. 1).

FIG. 1.

Beneficial and detrimental impacts of CaCO₃ from source to tap. (1) Source water: surface waters equilibrate with CO₂, controlling pH and adding inorganic carbon to the system. Natural inhibitors to CaCO₃ precipitation such as phosphate and NOM enter water through runoff or wastewater discharge, whereas calcium-containing minerals in soils such as limestone can dissolve. (2) Water treatment plant: corrosion inhibitors are also antiscalants that can inhibit CaCO₃ formation. Membrane-softening processes often scale and produce a brine reject stream with high hardness and treated water with lower hardness. (3) Distribution system: CaCO₃ scale forms over distribution lines, protecting them from corrosive waters and potentially sealing leaks, but also increasing the head loss. (4) Premise plumbing: in households and buildings with ion exchange or RO membrane softeners, brine reject streams with high hardness are produced. Water heaters are very susceptible to scaling due to their high temperatures. Scale can easily coat heating elements and clog the narrow channels in on-demand water heaters. Particulate CaCO₃ may accelerate erosion corrosion. CaCO₃, calcium carbonate; CO₂, carbon dioxide; NOM, natural organic matter: RO, reverse osmosis.

Scaling

Scaling is the most obvious manifestation of CaCO₃ precipitation in potable water systems, resulting in flow restrictions by clogging plumbing devices and increased head loss (Krappe, 1940; Crabtree et al., 1999). CaCO₃ scaling is also a problem in granular media filters in lime-softening plants, where its formation can “cement” together the filter bed, making it difficult to regenerate the media (Logsdon et al., 2002). In addition, scaling is a major issue in modern filters, such as reverse osmosis (RO) or ultrafiltration membranes, as the high concentrations of Ca²⁺ and inorganic carbon in the reject stream can lead to CaCO₃ precipitation on the membrane surface (Bremere et al., 1999; Gwon et al., 2003), and it is recommended that feed water pH be lowered to prevent scaling on the membrane and brine discharge system (Seneviratne, 2007).

Increased propensity of CaCO₃ to precipitate at higher temperatures makes scaling a troublesome issue in hot water systems, especially in water heater or water-cooling applications in which scale layers interfere with heat transfer and system efficiency. Water heating is extremely energy intensive. More specifically, the total energy demand for water-heating systems in buildings exceeds the energy demand for the entire water and wastewater utility sector (Brazeau and Edwards, 2011). It is estimated that hard water can lower the heat transfer efficiency of gas storage water heaters by 8.5% and of instantaneous gas water heaters by 30% over a 15-year lifecycle if no softening or deliming is practiced (WQRF, 2011). Softening hard water can reduce the life cycle cost of instantaneous gas water heaters by 22.5% and of gas storage water heaters by 6.6% over a 15-year lifecycle (WQRF, 2011). The annual cost of fouling for industrial operations in the United States alone was estimated to be around $3–10 billion in 1982 (Garrett-Price et al., 1985), including the cost of anti-scaling programs, downtime for maintenance, lost production, cleaning, higher energy costs due to reduced heat transfer, and over-sizing equipment where fouling is expected. The worldwide cost in 2000 was estimated at roughly $26 billion and $8–10 billion in the United States alone (Müller-Steinhagen, 2000).

Autogenous repair

A possible but overlooked benefit of CaCO₃ formation in potable water systems is its ability to seal leaks in a process known as autogenous leak repair. This method may have been first discovered and applied in ancient Rome, where the addition of alkaline wood ash was used to seal leaking pipelines (Pollio 15 BC as translated by Morgan, 1960). Although the mechanism of leak repair is unspecified, the high pH of wood ash solutions of 11.7–13.1 (Pitman, 2006) would also cause solids to precipitate in the water and clog leaks, in addition to the particulates in wood ash that could directly block leaks. In modern times, the self-healing of concrete cracks and leaks is known to be critical to the satisfactory long-term performance of concrete structures and pipelines (Hearn, 1998; Edvardsen, 1999). Autogenous healing via hydration reactions in concrete have been frequently observed (Hearn, 1998). However, CaCO₃ deposition is also a major contributor to concrete repair (Clear, 1985; Edvardsen, 1999). Parks et al. (2010) found that concrete can be repaired by magnesium-silicate precipitates, which also form autogenously at higher pH in certain waters.

Recent Trends Affecting CaCO₃ Formation

Overview

In the past half-century, the science and practice of potable water treatment and distribution have changed markedly in ways that can profoundly alter the extent of CaCO₃ formation and scaling in water main and premise plumbing pipe systems (Fig. 1). Inorganic carbon dissolves into natural water systems from carbon dioxide (CO₂) in the atmosphere, or it is generated through biodegradation of organic matter in ground waters. Ca²⁺ is released in natural waters through dissolution of minerals such as calcite, portlandite, gypsum, and dolomite that come into contact with the water (Table 1). Deicing salt also contributes thousands of tons of Ca²⁺ to water supplies in regions where calcium chloride is used (Houska, 2007).

Table 1.

Solubility of Calcium-Containing Minerals Found in Natural Waters

Mineral	pK_sp (25°C)	Reaction	References
Limestone (calcite + aragonite)	8.34–8.48	\documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\usepackage{upgreek}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document} $$CaC{O_3} \left( s \right) \rightleftharpoons C{a^{2 + }} + CO_3^{2 - }$$ \end{document}	Plummer and Busenberg (1982)
Portlandite	22.8	\documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\usepackage{upgreek}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document} $$Ca{ ( OH ) _2} \left( s \right) \rightleftharpoons C{a^{2 + }} + 2O{H^ - }$$ \end{document}	Brezonik and Arnold (2011)
Dolomite	17.2	\documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\usepackage{upgreek}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document} $$CaMg{ ( C{O_3} ) _2} \left( s \right) + \;{H^ + } \rightleftharpoons MgC{O_3} \left( s \right) + C{a^{2 + }} + HCO_3^ -$$ \end{document}	Sherman and Barak (1999)
Gypsum	4.16	\documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\usepackage{upgreek}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document} $$CaS{O_4} \cdot 2{H_2}O \left( s \right) \rightleftharpoons C{a^{2 + }} + SO_4^{2 - } + 2{H_2}O$$ \end{document}	Brezonik and Arnold (2011)

CaCO₃ formation as a corrosion control method has fallen out of favor

Corrosion has taken a heavy toll on plumbing and distribution systems throughout the United States. The cost of maintaining the aging drinking water infrastructure in the United States alone will exceed $1 trillion over the next 25 years (AWWA, 2012). The original basis for controlling or reducing corrosion in potable water distribution systems was based on the premise that a layer of CaCO₃ formed on pipe surfaces could almost completely protect all pipe materials from corrosive reactants in the water supply, including oxygen and free chlorine. Numerous corrosion indices based on this hypothesis were formulated to monitor and control the supposedly protective carbonate scales, including Langelier saturation index (LSI), Ryznar stability index (RSI), Puckorius scaling index (PSI), Larson–Skold index (LI), and measurement of temporary hardness (Langelier, 1936; Ryznar and Langelier, 1944; Larson and Skold, 1958; Puckorius and Brooke, 1991). Very early on, some of these indices were criticized because they oversimplified the process of scaling (Wiggin et al., 1938; Schneider and Stumm, 1964), but the notion that scales protected pipes from corrosion dominated potable water treatment until about the mid-1980s, when researchers studying corrosion began to notice that iron, lead, and copper corrosion had little to do with CaCO₃ precipitation, and a consensus emerged that the overall approach should be abandoned (Stumm, 1956; Larson and Skold, 1957; Schock, 1989; AWWARF, 1996). However, that consensus neglected to consider the possible benefits of the approach in reducing pipeline leaks and extending the lifetime of assets through autogenous repair processes or clogging (Tang et al., 2013). In practice, many municipal water utilities continue to calculate and target specific levels of CaCO₃ saturation (or undersaturation) to extend the lifetime of cement and mortar-lined surfaces, or in hopes of controlling corrosion of other materials.

CaCO₃ scaling in building heating systems has recently become an increased concern

Many buildings are now beginning to target a higher hot water system temperature set point (>51°C) to reduce growth of pathogens (e.g., Legionella pneumophila). However, at such high temperature set points, scale formation on the surface of the heating element and throughout the system occurs more rapidly. The reason is that as temperature increases, the solubility of CaCO₃ decreases and the rate of precipitation increases. Scaling can create serious problems, including loss of heat transfer efficiency in heating systems, head loss (energy loss) induced by pipe clogging, and, in some cases, erosion corrosion from particles (Krappe, 1940; Crabtree et al., 1999; Dobersek and Goricanec, 2007; Roy and Edwards, 2015). Lower heat transfer efficiency in heating systems can severely impair the proper operation of water heaters in homes and buildings as well as cooling towers and heat exchangers in industrial settings (Garrett-Price et al., 1985).

Tankless on-demand water heaters are emerging as an important energy-conserving device, due to their minimal hot water storage volumes and heat losses due to stagnation. The Department of Energy estimates that a centralized on-demand water heating system can reduce energy consumption by 8–34% (U.S. DoE, n.d.) compared with conventional tank systems. Homeowners with gas-heated on-demand systems are also eligible for federal tax credits (Energy Star, n.d.), and the lack of stored hot water can reduce the risk of opportunistic pathogen growth (Brazeau and Edwards, 2011). Although these factors incentivize home and building owners to install on-demand water heaters, their high surface area-to-volume ratio at heat transfer surfaces, which are inherent to their instant heating design, makes them highly susceptible to CaCO₃ scaling and clogging in scaling waters (Thomas et al., 2006; Brazeau and Edwards, 2011). In some cases, scaling may also reduce or eliminate their energy savings relative to conventional water heaters.

CaCO₃ may have the ability to clog leaks that have formed in pipes and extend the lifetime of distribution systems

It is hypothetically possible that, despite the sound scientific work discrediting Langelier and related indices as a guide to corrosion control, the indices might have actually been “working” as advertised in controlling pipe leaks and extending distribution system lifetimes (DeMartini, 1938) but for different reasons. Specifically, rather than forming a thin CaCO₃ layer to protect pipes from corrosion and associated leaks (Langelier, 1936), it is possible that a high CaCO₃ formation potential could have reduced leak frequency by autogenous repair or clogging of leaks (Tang et al., 2013). Autogenous repair is a process by which incipient existing leaks seal themselves, and CaCO₃ formation is a well-known leak self-repair mechanism for concrete pipes (Edvardsen, 1999; Parks et al., 2010). This mechanism could also repair and heal leaks in metallic pipes before the leak holes grow to catastrophic failure, by clogging of holes, assisting the formation of stronger or thicker pipe scales, or formation of corrosion byproducts (Tang et al., 2013, 2015). More research is needed for correlating actual leaks to CaCO₃ formation potential to test this hypothesis, especially given the poor condition of U.S. distribution system assets and the lack of financial resources for traditional repair and replacement.

CaCO₃ precipitation inhibitors that are naturally present or added at the point of treatment have probably changed the propensity of water to form CaCO₃ scales in the past few decades

Water treatment policies aimed at controlling lead and copper corrosion dramatically increased the use of phosphate and polyphosphate corrosion inhibitors starting around 1990. These inhibitors are known to slow CaCO₃ precipitation kinetics in potable water systems (Lin and Singer, 2005a, 2006; Wang et al., 2017). The use of such corrosion inhibitors might prevent leak repair that occurs naturally for concrete pipe leaks or even possibly for metallic pipe leaks that would have occurred when the inhibitor was absent. Similarly, natural organic matter (NOM), which is a disinfection byproduct precursor in water treatment plants, has been shown to inhibit CaCO₃ precipitation (Lin et al., 2005).

Climate change may play a role in CaCO₃ formation

With rising atmospheric CO₂ levels since the preindustrial era, the potential amount of CaCO₃ to precipitate in potable water systems could increase at temperatures (25–84°C) that are relevant to potable water systems (Fig. 2). With the atmospheric CO₂ concentration increases from the preindustrial era to 2050 (projected), the atmospheric partial CO₂ pressure increases from 280 × 10⁻⁶ atm to 550 × 10⁻⁶ atm. Using MINEQL+ modeling to model waters with an initial amount of 40 mg/L Ca²⁺, the amount of CaCO₃ to precipitate could increase from 11 to 14 mg/L (a 27% increase) at 84°C, which is a temperature that is commonly present at the heating element.

FIG. 2.

Calcite at equilibrium with atmospheric CO₂ levels in the preindustrial era, present day, and those projected for 2050, and temperatures up to 84°C, a temperature commonly present at the heating element. Higher atmospheric CO₂ can increase the precipitation potential. Using Mineql+, waters were equilibrated at [Ca²⁺] = 40 mg/L and P_CO2 = 280 × 10⁻⁶, 390 × 10⁻⁶, and 550 × 10⁻⁶ atm. Using the equilibrium Ca_T, Alk, and pH from the open systems as starting conditions for the closed systems, the potential CaCO₃ formation at equilibrium was calculated for the closed systems.

More frequent and longer droughts caused by global warming can also increase calcium levels in natural waters (Anderson and Faust, 1972; Mosley, 2015). Consequently, the potential amount of CaCO₃ to precipitate in potable water systems may also be rising.

It is worth mentioning that some types of dissolved organic carbon (DOC), which are known to be a CaCO₃ precipitation inhibitor (Lebron and Suarez, 1996), might be increasingly released into natural waters, and they might play a role in CaCO₃ precipitation in potable water systems. Although the major cause of increased DOC levels in natural waters is attributed to a decrease in acid rain deposition (Evans et al., 2006), there is evidence that the increasing CO₂ levels and the rising temperatures due to global climate change (Freeman et al., 2001) might contribute to DOC release. As a result, the increasing DOC release due to climate change might affect CaCO₃ precipitation in potable water systems.

Scaling has been shown to affect biofilm growth in premise plumbing

One study conducted by Fox and Abbaszadegan (2013) found that scales on polyvinyl chloride (PVC) and copper, materials common in premise plumbing, promoted biofilm growth because scaling increased the surface roughness. An increase in biofilm growth was most pronounced on scaled copper surfaces, perhaps because the scale shielded the biofilm from the antimicrobial properties of copper.

Bulk precipitation of CaCO₃ may contribute to erosion corrosion in plumbing systems

Erosion corrosion occurs via mechanical action of fluid and suspended particles, removing protective films on pipe surfaces. Roy and Edwards (2015) investigated erosion corrosion and found that aragonite, which is one polymorph of CaCO₃, with a particle size <2 mm, is capable of creating leaks in copper pipes. Typically, the formation of large CaCO₃ particles is not a concern in potable water systems; however, it has been shown that the CaCO₃ scale can detach from premise plumbing materials (e.g., copper) by shear stress (Royer et al., 2010). The scale that gets released has the potential to cause or accelerate erosion corrosion, especially in hot water recirculating systems that are now required by code in many parts of the country.

The complexity of these phenomena and their profound impact on the infrastructure, energy, and sustainable water nexus make it an ideal time to revisit this important topic (Table 2). The objectives of this review are: (1) to re-evaluate CaCO₃ formation in a modern context, such as global warming, autogenous repair, corrosion inhibitors, and increasing water heater temperatures in buildings; (2) to identify knowledge gaps regarding CaCO₃ formation and control.

Table 2.

Recent Phenomena and Its Effect on Calcium Carbonate Formation

Phenomenon	Effects on CaCO₃ formation	Research priorities
Global warming	Greater CO₂ in atmosphere can raise or lower amount of CaCO₃ that could form	Experimentally alter CaCO₃ formation potential at varying P_CO2
Global warming	More severe droughts could raise Ca²⁺ in natural waters
Corrosion and scaling inhibitors	Phosphates added to drinking water reduce CaCO₃ formation	Find out how CaCO₃ precipitation is affected by higher and lower phosphate concentrations
	Use increased in past 25 years, but pressure on industry to lower the dosage for cost and environmental reasons	Effect of nonphosphate inhibitors on CaCO₃ precipitation
	NOM is a scaling inhibitor, utilities are targeting it for removal because it is a precursor to DBPs	Effect of nonphosphate inhibitors on CaCO₃ precipitation
Raising water heater temperature	Home and building owners are raising water heater temperatures to control opportunistic pathogens. Higher temperatures increase CaCO₃ scaling potential	Improve water heater design to reach higher temperatures but reduce scaling potential
Autogenous repair	CaCO₃ formation could beneficially be used to seal leaks in plumbing and distribution systems	Alter chemistry of drinking water to promote leak repair via CaCO₃
Autogenous repair		Examine the impact that inhibitors have on ability to repair
Biofilm growth	CaCO₃ provides a surface for biofilm to grow on. Shields biofilm from antimicrobial plumbing materials	Determine health implications of CaCO₃-induced biofilm growth in domestic plumbing
Erosion corrosion	Small (<2 mm) CaCO₃ particles and CaCO₃ scale released from pipe surface can cause erosion corrosion	Analyze flow patterns and plumbing design to mitigate erosion corrosion from CaCO₃

CaCO₃, calcium carbonate; CO₂, carbon dioxide; DBPs, disinfection byproducts; NOM, natural organic matter.

Chemistry of CaCO₃ Formation

Background

To fully understand practical aspects of CaCO₃ behavior in potable water systems, it is important to first review the chemistry of CaCO₃ formation and dissolution. CaCO₃ crystals are made of Ca²⁺ and CO₃²⁻ lattice ions, both of which are ubiquitous in natural water systems. It can exist in three crystalline polymorphs as well as in an amorphous phase (ACC) or as a mono- and hexahydrate. The solubility product for each form varies with temperature, as noted in Table 3. CaCO₃ has significantly lower solubility at higher temperatures, causing CaCO₃ formation to be a major issue in many hot water systems even when it is not a problem in cold water systems for waters with moderate hardness and alkalinity (Fig. 3).

FIG. 3.

Amount of potential CaCO₃ formation for all crystalline polymorphs vaterite [(diamond), aragonite (square), and calcite (triangle)] and ACC (circle) under different temperatures in water distribution systems. MINEQL+ simulations on a closed system with alkalinity = 100 mg/L as CaCO₃ at pH 8.5 and (a) initial TOTCa = 40 mg/L and (b) initial TOTCa = 65 mg/L. ACC, amorphous phase.

Table 3.

Temperature Dependence of Solubility Product for Calcium Carbonate Polymorphs

Polymorph	pK_sp at 25°C	K_sp equation	Temperature range	References
ACC	6.04	\documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\usepackage{upgreek}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document} $$\log { { \rm { K } } _ { { \rm { sp } } } } = { \frac { 1247 } { \rm { T } } } - 10.224$$ \end{document}	16–60°C	Clarkson et al. (1992)
Vaterite	7.91	\documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\usepackage{upgreek}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document} $$\log { { \rm { K } } _ { { \rm { sp } } } } = - 172.1295 - 0.077993 { \rm { <sup></sup>T } } + \displaystyle { \frac { 3074.688 } { \rm { T } } } + 71.595 { \rm { <sup></sup>logT } } $$ \end{document}	0–90°C	Plummer and Busenberg (1982)
Aragonite	8.34	\documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\usepackage{upgreek}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document} $$\log { { \rm { K } } _ { { \rm { sp } } } } = - 171.9773 - 0.077993 { \rm { <sup></sup>T } } + \displaystyle { \frac { 2903.293 } { \rm { T } } } + 71.595 { \rm { <sup></sup>logT } } $$ \end{document}	0–90°C	Plummer and Busenberg (1982)
Calcite	8.48	\documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\usepackage{upgreek}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document} $$\log { { \rm { K } } _ { { \rm { sp } } } } = - 171.9065 - 0.077993 { \rm { <sup></sup>T } } + \displaystyle { \frac { 2839.319 } { \rm { T } } } + 71.595 { \rm { <sup></sup>logT } } $$ \end{document}	0–90°C	Plummer and Busenberg (1982

T is temperature in K.

ACC, amorphous phase.

The levels of Ca²⁺ in natural waters are usually governed by regional geology (limestone, dolomite, etc.), whereas the carbonate system is more complex. CO₃²⁻ becomes the dominant form of inorganic carbon at pH 10.3 and higher (Plummer and Busenberg, 1982). \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\usepackage{upgreek}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document} \begin{align*} { \rm{C}}{{ \rm{O}}_2} \left( { \rm{g}} \right) + {{ \rm{H}}_2}{ \rm{O}} \rightleftharpoons { \rm{C}}{{ \rm{O}}_2} \left( {{ \rm{aq}}} \right) + {{ \rm{H}}_2}{ \rm{C}}{{ \rm{O}}_3} \left( {{ \rm{aq}}} \right) \quad \quad{ \rm{p}}{{ \rm{K}}_{ \rm{H}}} = 1.47 \tag{1} \end{align*} \end{document} \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\usepackage{upgreek}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document} \begin{align*} {{ \rm{H}}_2}{ \rm{C}}{{ \rm{O}}_3} \left( {{ \rm{aq}}} \right) \rightleftharpoons {{ \rm{H}}^ + } + { \rm{HCO}}_3^ - \left( {{ \rm{aq}}} \right) \quad \quad { \rm{p}}{{ \rm{K}}_{{ \rm{a}}1}} = 6.4 \tag{2} \end{align*} \end{document} \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\usepackage{upgreek}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document} \begin{align*} { \rm{HCO}}_3^ - \left( {{ \rm{aq}}} \right) \rightleftharpoons {{ \rm{H}}^ + } + { \rm{CO}}_3^{2 - } \left( {{ \rm{aq}}} \right) \quad \quad { \rm{p}}{{ \rm{K}}_{{ \rm{a}}2}} = 10.3 \tag{3} \end{align*} \end{document}

with Ω representing the degree of supersaturation, defined as \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\usepackage{upgreek}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document} \begin{align*} { \rm { \Omega } } = \frac { Q } { { { K_ { sp } } } } \tag { 4 } \end{align*} \end{document}

The ion activity product, Q, for CaCO₃ is the product of the Ca²⁺ and CO₃²⁻ activities, as shown in the equation given next. \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\usepackage{upgreek}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document} \begin{align*} { \rm{Q}} = \left\{ { \left. {{ \rm{C}}{{ \rm{a}}^{2 + }}} \right\} } \right. \left\{ { \left. {{ \rm{CO}}_3^{2 - }} \right\} } \right. \tag{5} \end{align*} \end{document}

When Q < K_sp, the solution is undersaturated, precipitation is not thermodynamically favored, and CaCO₃ solid can dissolve. However, when Q > K_sp, CaCO₃ is supersaturated and it can precipitate.

Precipitation and growth in plumbing and distribution systems

Homogeneous nucleation

Homogeneous nucleation of CaCO₃ occurs less frequently in plumbing and distribution systems than heterogeneous nucleation due to the larger energy barrier of spontaneously forming CaCO₃ nuclei in solution. As a result, homogeneous nucleation is mostly observed at Ω > 7, for experiments conducted at pH 8.50, 25°C, and up to 2 h of the reaction (Lioliou et al., 2007). Even in systems that are stagnant for long periods, homogeneous nucleation is likely to be the less common form of nucleation, because the scales already present or bare surfaces can lower the energy barrier for heterogenous nucleation.

Homogeneous nucleation of CaCO₃ was believed to follow the classical nucleation theory (CNT). CNT states that a nucleus is not stabilized until a critical radius is reached in which the bulk energy and surface energy of the nuclei is balanced (Gebauer et al., 2014). However, the recent discovery of prenucleation clusters (PNCs) undermines CNT (Gebauer et al., 2008). PNCs contain CaCO₃ lattice ions but do not have an explicit phase boundary with the solution. The structural characteristics of PNCs also indicate which polymorph nucleates. More stable PNCs yield ACC with a short-range order corresponding to calcite, whereas less stable PNCs yield ACC with a short-range order corresponding to vaterite (Lam et al., 2007; Gebauer et al., 2014).

Homogeneous CaCO₃ precipitation roughly follows Ostwald's rule of stages. The rule of stages indicates that the least stable polymorph often crystallizes first before transforming into the most stable polymorph (Ng et al., 1996). For CaCO₃, ACC is the first to precipitate, followed by vaterite, before transforming into calcite, which is its most stable form (Wolf et al., 2008). Ogino et al. (1987) observed the same at lower temperatures (14–30°C); whereas at higher temperatures (60–80°C), the transition of ACC to aragonite to calcite was observed. At intermediate temperatures (40–50°C), CaCO₃ crystallization closely follows Ostwald's rule of stages that every polymorph forms before becoming calcite (Ogino et al., 1987).

Heterogeneous nucleation and deposition of scales

Heterogeneous nucleation occurs more readily than homogeneous nucleation because the surface of the foreign solids generally lowers the activation energy barrier for nucleation (Brezonik and Arnold, 2011). As a result, it is likely the most common route of nucleation for CaCO₃ scale in plumbing and distribution systems.

In potable water systems, scaling is more likely related to heterogeneous nucleation combined with deposition. Mathematical modeling has been developed to describe such processes on different surface materials, such as a heat-flux exchanger (Hasson et al., 1968a, 1968b; Chen et al., 2005). These models assume that diffusion of Ca²⁺ and HCO₃⁻ ions from the bulk water toward the scale-liquid interface is the governing mechanism of scale deposition. This is also the reason why many characteristics of surface materials, including surface roughness (Keysar et al., 1994; MacAdam and Parsons, 2004a; Wu et al., 2010; Cheong et al., 2013), waterborne particles, and surface-free energy, influence the scaling.

Various materials may exhibit different susceptibility to CaCO₃ scaling in potable water systems. Common pipe materials include plastic (PVC, chlorinated PVC, or high-density polyethylene), iron (cast, ductile, or steel), copper, and concrete (new and aged) (Tang et al., 2013). Also, nearly all the heating elements in water heaters are made of stainless steel (Whirlpool, 2017), and polypropylene pipes are gaining popularity and market share in building plumbing (Marks, 2017). As for the scaling susceptibly, plastic pipes are generally believed to be less likely to scale than metallic pipes (Githens et al., 1965). For example, copper pipes are scaled in heat exchangers with Ca-containing minerals at a rate twice that of polypropylene (Wu et al., 2009). As for metallic surfaces, aged cast iron would probably have a higher scaling potential than copper or stainless steel, because it has a rougher surface (Jaćimović et al., 2015). As for concrete, the mechanism for scaling is quite different. As concrete ages in water, Ca(OH)₂ leaches out of its micropores and may cause scaling through localized homogeneous crystallization (Tang et al., 2013).

Waterborne particles can enhance the heterogeneous nucleation and subsequent CaCO₃ precipitation in potable water systems (Andritsos and Karabelas, 2003; Pääkkönen et al., 2009). Although suspended calcite crystals are the most effective CaCO₃ seeds (Nancollas and Reddy, 1971; Lin and Singer, 2005b), they are relatively uncommon in natural waters. Some clay particles that contribute to turbidity have been shown to induce CaCO₃ precipitation. For example, montmorillonite, a type of clay, was shown to instantaneously lead to calcite precipitation in a solution with 0.01 mol/L of Ca²⁺ and 0.004 mol/L of total carbonate at pH 10 (Kralj and Vdović, 2000), which is equivalent to Ω = 472 for calcite according to MINEQL+ calculation.

Surface-free energy is believed to influence heterogeneous CaCO₃ nucleation. For CaCO₃ scales, lower surface energy can reduce the scale adhesion strength (Müller-Steinhagen and Zhao, 1997; Förster et al., 1999; Yang et al., 2000, 2002; MacAdam and Parsons, 2004b). As for CaSO₄ scales that are very similar to CaCO₃ scales, Müller-Steinhagen and Zhao (1997) successfully reduced its scaling rate on stainless steel, which is used to make almost all the heating elements in water heaters, by lowering its surface energy via SiF₃⁺ ion implantation. Therefore, a similar result might be expected with CaCO₃ scaling.

Growth rate

Previous models of seeded growth kinetics only considered Ω. Morse (1983) modeled the growth rate by using the equation given next. \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\usepackage{upgreek}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document} \begin{align*} {\rm R} = {\rm k}{ ( { \rm{ \Omega }} - 1 ) ^n} \tag{6} \end{align*} \end{document}

where R is the rate of precipitation normalized to surface area, k is an empirical rate constant, and n is the reaction order.

The growth mechanism of CaCO₃ crystallization, which can be either surface controlled or diffusion controlled, is related to the formation and growth of CaCO₃ in potable water systems. In a surface-controlled mechanism, incorporation of lattice ions into the crystal structure is the rate-limiting process. In a diffusion-controlled mechanism, diffusion of lattice ions from the solution to the crystal surface is the rate-limiting process. As discussed in Heterogeneous Nucleation and Deposition of Scales, scaling in potable water systems is a combined effect of heterogeneous nucleation and scale deposition. Heterogeneous nucleation and subsequent crystal growth is essentially surface-controlled CaCO₃ growth, whereas scaling deposition is governed by diffusion of Ca²⁺ and HCO₃⁻ ions from the bulk water toward the scale-liquid interface, which is essentially diffusion-controlled CaCO₃ growth.

Dissolution

The dissolution rate of calcite (R) follows the rate Equation (7) and three reactions given next [Eqs. (8 )–(10)], based on numerous comprehensive studies on calcite dissolution kinetics or the dissolution of carbonate minerals (Berner and Morse, 1974; Plummer et al., 1978; Dreybrodt, 1980; House, 1981a, 1981b; Chou et al., 1989; Liu and Dreybrodt, 1997; Morse et al., 2007; Mitchell et al., 2010). The first chemical reaction [Eq. (8)] is the largest contributor to dissolution and is often diffusion controlled (Berner and Morse, 1974; Plummer et al., 1978). \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\usepackage{upgreek}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document} \begin{align*} \begin{split}R &= {k_1} \left\{ {H^+} \right\} + {k_2} \left\{ {H_2}CO_3^* \right\} \\ &\quad + {k_3} \left\{ {H_2}O \right\} - {k_4} \left\{ Ca^{2+}\right\} \left\{ HCO_3^-\right\} \end{split} \tag{7}\end{align*} \end{document}

where k₁, k₂, and k₃ are the first-order forward rate constants for the following three reactions that occur simultaneously when dissolution is far from equilibrium. The backward rate constant k₄ is identified as significant only when dissolution is approaching equilibrium. k₁, k₂, and k₃ are dependent on temperature; k₄ is a function of both temperature and P_CO2; and H₂CO₃* is CO_2(aq) + H₂CO₃° (Plummer et al., 1978). \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\usepackage{upgreek}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document} \begin{align*} CaC{O_3} + {H^ + } \rightleftharpoons C{a^{2 + }} + HCO_3^ - \tag{8} \end{align*} \end{document} \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\usepackage{upgreek}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document} \begin{align*} CaC{O_3} + {H_2}CO_3^* \rightleftharpoons C{a^{2 + }} + 2HCO_3^ - \tag{9} \end{align*} \end{document} \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\usepackage{upgreek}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document} \begin{align*} CaC{O_3} + {H_2}O \rightleftharpoons C{a^{2 + }} + HCO_3^ - + O{H^ - } \tag{10} \end{align*} \end{document}

Morphology and polymorph stability

Although calcite is the thermodynamically stable polymorph in most cases, solution temperature can impact the stability and metastability of aragonite and vaterite. At temperatures >60°C, homogeneous precipitation yields aragonite as the dominant polymorph over a period of hours (Wray and Daniels, 1957; Sawada, 1997; Kellermeier et al., 2013). Temperature also affects the relative abundance of each polymorph during the early metastable stage of homogeneous precipitation. Vaterite is the dominant metastable phase at standard temperature; however, at temperatures below 25°C, calcite is the dominant metastable phase and at temperatures above 50°C, aragonite is the dominant metastable phase (Ogino et al., 1987).

The mechanism of CaCO₃ precipitation and calcite formation occurs in stages via redissolution of ACC to vaterite and vaterite to calcite over a period of ∼3 h (Ogino et al., 1987; Sawada, 1997). Knowing the polymorph evolution as CaCO₃ precipitates could have importance in predicting scale formation. Roy and Edwards (2015, 2016) noted that aragonite particles are physically harder (mineral hardness of 3.5–4.0), compared with calcite (mineral hardness of 3.0). However, this contradicts the claim that aragonite is “a removable soft scale,” compared with calcite, which is “a hard scale,” according to studies on magnetic water treatment (Raisen, 1984; Baker and Judd, 1996; Coey and Cass, 2000; Liu et al., 2010). Also, many newly developed water treatment technologies, such as electrically induced precipitation and electromagnetic treatment, have been claiming in their advertisements that the scales formed after such treatment are “soft and can be easily brushed off.” Therefore, we believe that further studies are greatly needed to verify these anecdotal claims of “soft and easily removed aragonite scales,” even though aragonite has a higher mineral hardness compared with calcite. Further, if this claim can be verified, the scale deposition could be predicted or even controlled in some real water systems.

Monitoring CaCO₃ Formation in Potable Water Systems

The problems caused by scaling have led to the development of numerous scaling indices (Table 4) to monitor CaCO₃ formation in potable water systems. Degree of supersaturation (Ω) is the most widely accepted scaling index. Obviously, calculation of Ω is the first and most important step to predicting scaling in potable water systems, as it definitively determines whether scaling is possible. However, Ω does not account for reaction kinetics. As discussed in the previous sections, CaCO₃ precipitation and dissolution can occur very slowly, especially in the presence of natural and artificial scaling inhibitors—thus, calculation of Ω serves best as a first step to predicting scaling in potable water systems.

Table 4.

Scaling Indices Used in Practice

Index or test	Definition	Scaling condition	Accounts for kinetics?
Degree of supersaturation	\documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\usepackage{upgreek}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document} $$ { \rm { \Omega } } = \displaystyle { \frac { \left\{ { C { a^ { 2 + } } } \right\} \left\{ { CO_3^ { 2 - } } \right\} } { { K_ { sp } } } } $$ \end{document}	Ω > 1	No
LSI	LSI = pH_actual − pH_saturated where pH_saturated = pK₂ − pK_s + pCa + pAlk	LSI > 0	No
RSI	RSI = 2·pH_saturated − pH_actual	RSI < 6.2	No
PSI	PSI = 2·pH_saturated − pH_equilibrium where pH_equilibrium = 1.47 + log(Alk) + 4.54	PSI < 6	No
LI	LI = [(Cl⁻) + (SO₄²⁻)]/[(HCO₃⁻) + (CO₃²⁻)]	LI < 0.8	No
CCPP	CCPP = 50,000 ([Alk]–[Alk]_equilibrium) For [Alk]_equilibrium calculation, refer to U.S. EPA (1992)	CCPP > 4 mg/L	No
Temporary hardness	Water is heated until no more minerals precipitate, standard method is unknown	Measured by hardness loss	Yes

CCPP, calcium carbonate precipitation potential; LI, Larson–Skold index; LSI, Langelier saturation index; PSI, Puckorius scaling index; RSI, Ryznar stability index.

The LSI was first described in Langelier (1936) as a way to predict the formation of a corrosion-resistant CaCO₃ coating in distribution systems. In his work, Langelier specifically addressed the degree of supersaturation in a very fundamental way. As a result, this simple index of CaCO₃ precipitation based on fundamental chemical equilibria is still highly relevant to the present day, even if its use to predict corrosion control is out of favor.

The LSI has since been revised to predict scaling in specific systems. However, LSI is overly simplistic in that it assumes that pH and CaCO₃ saturation are the most important factors for corrosion control (Schneider and Stumm, 1964) and it does not account for other dissolved calcium species or other forms of alkalinity (Kutty et al., 1992). Dąbrowski et al. (2010) tested the LSI on a water distribution system and found no correlation between waters with LSI <0 and a water's aggressiveness.

Indices developed after LSI are similar to LSI in that they express CaCO₃ undersaturation or supersaturation, but they take into account and emphasize different factors such as equilibrium pH and noncarbonate hardness. The RSI still uses pH_saturated as does the LSI; however, it recognizes calcium and alkalinity as more important parameters in determining scale formation potential, and, thus, pH_saturated is multiplied by a factor of 2 (Ryznar and Langelier, 1944). The RSI was successfully tested on real waters where it was determined that at an RSI <6.2 scales should form. The PSI was formulated for cooling tower water and includes a term for pH_equilibrium, which depends on alkalinity (Puckorius and Brooke, 1991). PSI stresses the importance of alkalinity in scale formation, because a higher alkalinity increases the water's buffering capacity such that more CaCO₃ is able to precipitate before the water reaches equilibrium pH. Unlike the other indices, the LI does not use pH_saturated in its calculation. LI was originally intended for Great Lakes waters traveling through steel distribution lines (Larson and Skold, 1958), and it is best suited for waters of a similar quality. LI considers the effect of Cl⁻ and SO₄²⁻ on CaCO₃ film formation, since higher Cl⁻ and SO₄²⁻ levels usually indicate a greater presence of noncarbonate hardness (e.g., CaSO₄), which is not conducive to scale formation in potable water systems.

The calcium carbonate precipitation potential (CCPP) is unique in that it attempts to quantify the amount of CaCO₃ that could precipitate. For the purpose of corrosion control, rule-of-thumb guidelines are that the CCPP should have a value of 4–10 mg/L for a thin protective coating to form (Merrill and Sanks, 1977a, 1977b; Merrill and Sanks, 1978; U.S. EPA, 1992). CCPP relies on the estimation of the equilibrium alkalinity, [Alk]_equilibrium, through an iterative process. Rossum and Merrill's (1983) study comparing several scaling indices found that CCPP is the best estimate of a water's ability to deposit or dissolve CaCO₃.

Temporary hardness represents the amount of hardness that precipitates by boiling water, and it was once widely applied as a practical way of describing the scaling potential of water-heating systems. In practice, temporary hardness does not often differentiate the precipitated solids, so it is a measure not only of CaCO₃ precipitation but also of sulfate salts (Greth, 1910; Buswell, 1916; Goudey, 1933). Temporary hardness has also been defined strictly as carbonate hardness, which in the past was measured by titration and soap tests (Hehner, 1883; Norton and Knowles, 1916). Despite its simplicity, the temporary hardness test, or some modern modification, could have value as a quick practical test for scaling potential.

Control and Prevention of Calcium Carbonate Scaling

Control CaCO₃ precipitation through water chemistry quality parameters

To control and prevent CaCO₃ scaling issues, the first and foremost step is to monitor the scaling potential of drinking water (e.g., degree of supersaturation, LSI) at different points of the distribution system, as discussed in Monitoring CaCO₃ Formation in Potable Water Systems. Once CaCO₃ scaling potential is confirmed, different control and prevention methods might be applied for specific environmental conditions and water compositions, according to Table 5.

Table 5.

Control and Prevention of Calcium Carbonate Scaling in Drinking Water Systems

Methods	Likely effects	Rationale
Control CaCO₃ precipitation
Decrease hot water temperature	Effective	Solubility products of CaCO₃ increase with decreasing water temperature.
Decrease water pH	Effective	Saturation ratios of CaCO₃ decrease with decreasing pH.
Remove dissolved Ca	Effective	Saturation ratios of CaCO₃ decrease with decreasing dissolved Ca concentration.
Control CaCO₃ morphology	Sometimes effective	Reduce scaling potential, by preferring the CaCO₃ polymorph with a higher solubility product to form first.
Control the use of surface materials	Sometimes effective	Scaling potentials of CaCO₃ are different on various surface materials.
Control the location of CaCO₃ formation	Sometimes effective	Phosphates delay the formation of CaCO₃, so that it occurs throughout potable water systems.
Delay or prevent CaCO₃ precipitation
Scale inhibitors	Effective	Delay or prevent CaCO₃ precipitation completely, mostly by adsorption of inhibitors onto CaCO₃ nuclei.
Potential beneficial application of CaCO₃
Control CaCO₃ precipitation for autogenous repair of pipe leaks	N/A	Precipitation of CaCO₃ is allowed. However, the precipitated CaCO₃ particles can be used for autogenous repair of the leaks in drinking water pipelines (Tang et al., 2013).
Dutch pellet softening method	Effective	Precipitation of CaCO₃ on calcite pellets removes supersaturation when water passes through a column of calcite pellets. Contact time and surface area are important parameters (Schetters, 2013).

N/A, not available.

Saturation management (pH and dissolved Ca)

Reducing water pH and removing dissolved Ca content are the principle means of saturation management for potable water systems, including the process industry, nanofilters, and RO membranes. Either reducing pH or removing dissolved Ca can reduce the scaling potential by decreasing the ion activity product \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\usepackage{upgreek}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document} $$(Q = \left\{ { \left. {C{a^{2 + }}} \right\} } \right. \left\{ { \left. {CO_3^{2 - }} \right\} } \right.)$$ \end{document} . However, engineers should be careful when considering reducing water pH in drinking water systems, because the decrease in water pH might cause corrosion of metallic pipelines.

Temperature

Reducing water temperature is very effective in reducing CaCO₃ scaling potential in potable water systems. First, the solubility of all CaCO₃ polymorphs (calcite, vaterite, aragonite, and amorphous CaCO₃) is expected to increase with decreasing temperature (Plummer and Busenberg, 1982; Clarkson et al., 1992), causing less CaCO₃ to precipitate at lower temperatures. Second, the reaction kinetics for CaCO₃ formation is much slower at lower temperatures (Morse et al., 2007), thereby forming less CaCO₃ precipitates. Finally, polyphosphates, which are widely present in potable water systems and are stronger inhibitors for CaCO₃ precipitation than orthophosphate (Lin and Singer, 2005), are converted into orthophosphate much faster at a higher temperature through hydrolysis reactions (Holm and Edwards, 2003). Hence, at a lower temperature, more polyphosphates are present and will inhibit CaCO₃ precipitation more. However, the temperature drop for hot water might cause opportunistic premise plumbing pathogens, including L. pneumophila, to grow and transmit in drinking water systems and cause potential human health problems (ASHRAE, 2015).

To further illustrate how effective reducing water temperature is in controlling CaCO₃ precipitation, we compared the scaling potentials of all the potable water samples collected from a national survey at different temperatures relevant to potable water systems (Fig. 4) (U.S. EPA, 2000). In our calculation, we chose calcite saturation ratio (Ω_calcite) to represent the calcite scaling potentials for the 51,846 water samples collected from water utilities in the United States, and we only accounted for solubility change at different temperatures. As the hypothetical set point of the water heater temperature increased from an ambient temperature to higher temperatures, the percentage of water samples that are at risk of scaling (Ω_calcite > 1) is 59% (or 30,538 water samples) at 48°C and 62% (or 31,910) at 60°C, which is much higher than the 47% (or 24,468 water samples) at the ambient temperature (Fig. 4b). This clearly indicates that more waters are at higher scaling risks as temperature increases, due to the decreasing calcite solubility alone. In addition, most of these water samples do not have very high Ca hardness (soft and moderately hard), which might explain why most of these water samples do not have scaling potential at ambient temperatures (Fig. 4c).

FIG. 4.

Scaling potentials of water samples collected from a national survey at different temperatures relevant to potable water systems (U.S. EPA, 2000). (a) Histogram of saturation ratio of calcite (Ω_calcite) at various temperatures, (b) the percentage of water samples at scaling risks (Ω_calcite > 1), and (c) the percentage of water samples at different Ca hardness levels. The saturation ratio of calcite was calculated for 48°C and 60°C, in addition to the ambient temperature provided in the national survey (U.S. EPA, 2000). MATLAB was used to calculate Ω_calcite for 51,846 water samples that have all four parameters of temperature, pH, alkalinity, and Ca hardness. The change of equilibrium constants (K) and solubility products (K_sp) with temperature was accounted for by using the K and K_sp values and ΔH at 25°C from MINEQL+. Ca-containing species considered included Ca²⁺, CaHCO₃⁺, CaOH⁺, and CaCO_3(aq).

Other methods that might be effective in controlling CaCO₃ scaling

Stabilization of metastable CaCO₃ polymorphs over calcite by using different additives might be effective in controlling CaCO₃ scaling potential. Additives, including metal cations, silica, and organic species, can help stabilize metastable CaCO₃ polymorphs, such as aragonite, vaterite, and ACC (Table 6). Zn²⁺ is the only additive that has been reported to help avoid massive calcite scaling when dosed at a concentration that is relevant to drinking water (1.5 × 10⁻⁷–1.5 × 10⁻⁴ M), by preferring aragonite nuclei to form instead of calcite. As a result, Zn can coprecipitate into aragonite nuclei or adsorb on aragonite nuclei and inhibit the further growth of CaCO₃ crystals (Wada et al., 1995; Ghizellaoui and Euvrard, 2008). Other additives must be dosed at concentrations that are much higher than their typical concentrations in drinking water. For example, when silica is added at a concentration of 3.3 × 10⁻³–0.2 M, metastable CaCO₃ polymorphs can be stabilized to coexist with calcite for as long as 2 h at 50°C and 80°C (Kellermeier et al., 2013).

Table 6.

Different Additives and Their Favored Calcium Carbonate Polymorphs

	Species	Concentration	Favored CaCO₃ polymorph	References
Cations	Fe²⁺, Mg²⁺, Ni²⁺ Co²⁺, Cu²⁺	0.2–3.0 M	Aragonite	Wada et al. (1995)
	Zn²⁺	1.5 × 10⁻⁷	Aragonite	Wada et al. (1995); Ghizellaoui and Euvrard (2008)
	Zn²⁺	−1.5 × 10⁻⁴ M	Aragonite	Wada et al. (1995); Ghizellaoui and Euvrard (2008)
Silica	SiO₂	3.3 × 10⁻³−0.2 M	Vaterite, aragonite, and ACC	Kellermeier et al. (2013)
Organic	PAA, PBTCA	5–25 mg/L	Vaterite, aragonite	Yang et al. (2001)
Organic	OMM	8–20 mg/L	Vaterite	Roqué et al. (2004)

OMM, organic matrix macromolecules; PAA, polyacrylic acid; PBTCA, 2-phosphonobutane-1,2,4-tricarboxylic acid.

Surface materials and roughness might affect CaCO₃ scaling potentials as well as the physical and mechanical properties of scales. Since CaCO₃ can form through heterogeneous nucleation on different surface materials, the surface energy and the nucleation energy barrier of different surface materials will affect CaCO₃ scaling potentials. As discussed in Heterogeneous Nucleation and Deposition of Scales, copper has a higher scaling rate compared with steel (Macadam and Parsons, 2004a). In addition, the roughness of the surface materials can affect the physical and mechanical properties of CaCO₃ scales (e.g., porosity and tensile strength). For mild steel surfaces, the rougher the surfaces, the lower porosity and higher tensile strength the CaCO₃ scales possess (Keysar et al., 1994). This might be due to enhanced surface nucleation density and a certain orientation of the calcite structure.

Delay or prevent CaCO₃ precipitation by using scale inhibitors

Scale inhibitors can effectively delay or even completely prevent CaCO₃ scaling problems. Scale inhibitors are widely used to combat severe scaling issues in water processing equipment that handle large volumes of water, including cooling towers, nanofilters, and RO membranes, and production wells in oil and gas fields. These scale inhibitors generally include phosphonates, phosphate esters, polyacrylates, phosphates, and polyphosphates (Matty and Tomson, 1988; Xyla et al., 1991; Shakkthivel and Vasudevan, 2006). However, most of these inhibitors cannot be used in the finished drinking water, except phosphates or polyphosphates.

Inorganic cations Mg²⁺ and Zn²⁺ that are readily present in drinking water systems can also act as scale inhibitors (Table 7). Addition of Mg²⁺ at a concentration of 10⁻³ M and above can significantly inhibit calcite precipitation (Reddy and Wang, 1980; Lin and Singer, 2009). The mechanism behind the inhibiting effect of Mg²⁺ has been confirmed as the adsorption of Mg²⁺ onto calcite active crystal growth sites (or kinks) and subsequent kink blocking (Lin and Singer, 2009; Nielsen et al., 2013), rather than the previously proposed mechanisms of increasing the mineral solubility by incorporating Mg²⁺ into calcite (or incorporation inhibition) (Davis et al., 2000). Addition of Zn²⁺ at a much smaller concentration (∼10⁻⁷ M) can also significantly inhibit calcite precipitation (Ghizellaoui and Euvrard, 2008). It is possible that Zn²⁺ can inhibit CaCO₃ (calcite or aragonite) precipitation through either kink blocking or incorporation inhibition, but the exact mechanism behind this inhibiting effect by Zn²⁺ has not yet been identified. However, no threshold values can be identified for either Mg²⁺ or Zn²⁺, because they reduce the CaCO₃ precipitation rate even at very low concentrations, rather than abruptly halt calcite precipitation once a threshold value is reached.

Table 7.

Effective Concentrations of Calcium Carbonate Scale Inhibitors in Drinking Water Systems

Scale inhibitors	Effective concentration	References
Mg²⁺	6–21 ( × 10⁻³ M)	Lin and Singer (2009)
Zn²⁺	2–7 ( × 10⁻⁷ M)	Ghizellaoui and Euvrard (2008)
Phosphates^a	5 × 10⁻⁸ − 5 × 10⁻⁶ M	Lin and Singer (2005a)
Phosphates^a	5 × 10⁻⁸ − 5 × 10⁻⁶ M	Lin and Singer (2006)
NOM^b	1–5 (mg/L)	Lin et al. (2005)

Phosphates include orthophosphate, pyrophosphate, tripolyphosphate, and hexametaphosphate.

NOM includes Suwannee River fulvic acid, Pacific Ocean fulvic acid, and Williams Lake hydrophobic organic acid.

Other inhibitors that widely exist in drinking water systems are phosphates and NOM (Table 7). Lin and Singer systematically studied the inhibiting effects of orthophosphate (PO₄³⁻), pyrophosphate (P₂O₇⁴⁻), tripolyphosphate (P₃O₁₀⁵⁻), hexametaphosphate (P₆O₁₈⁶⁻), and binary-polyphosphate blends. They suggested that orthophosphate at a concentration of 5 × 10⁻⁶ M, and polyphosphate or polyphosphate blends at a concentration of 5 × 10⁻⁸ M can inhibit calcite precipitation (Lin and Singer, 2005a, 2006). The reason for this inhibiting effect is the adsorption of phosphates onto calcite crystal growth sites and blocking further growth. NOM, such as fulvic acid and humic acid, can also inhibit calcite precipitation through a similar mechanism to phosphates (Hoch et al., 2000; Lin et al., 2005).

Phosphates can be used to advantageously control the location of CaCO₃ formation in potable water systems. Addition of excess lime may only tend to coat the pipeline near the point of addition (Hatch, 1942). However, the addition of orthophosphate (Hatch, 1942) and polyphosphate (McCauley, 1960; Hasson and Karmon, 1981), including hexametaphosphate (Primus and Hunhoff, 1972), to supersaturated waters could delay precipitation or prevent it completely, potentially extending the distance that pipes are coated from CaCO₃ formation.

Potential beneficial application of CaCO₃ precipitation for autogenous repair

Although usually considered harmful, CaCO₃ precipitation might be potentially beneficial when manipulated to repair leaks in drinking water systems via autogenous repair (i.e., leak self-repair) (Clear, 1985; Edvardsen, 1999; Tang et al., 2013). Autogenous repair is a promising and low-cost approach to leak repair that can be achieved by manipulating the water chemistry and formation of CaCO₃ (e.g., calcite) to clog small leaks (holes typically ≈10 μm diameter) (Tang et al., 2013). If successfully applied, autogenous repair through CaCO₃ precipitation might help reduce water losses that cost about $3 billion per year (FHWA, 2002), and extend the lifetime of the critically important and aging drinking water infrastructure that is expected to cost at least $1 trillion to repair, replace, and upgrade through 2035, according to an estimate by the American Water Works Association (Shanaghan, 2012).

With additional research, water utilities might achieve autogenous repair via CaCO₃ precipitation, by either targeting the precipitation of CaCO₃ to occur specifically in leaks (smart precipitation) or clogging the leaks when CaCO₃ precipitates pass through the leaks (physical clogging) (Fig. 5) (Tang et al., 2013). It is even possible that these two mechanisms may be operative simultaneously. More specifically, to achieve smart precipitation, water utilities can appropriately adjust water pH, and “seed” small amounts of CaCO₃ solids into holes by dosing them to water for a short period, or by adding high amounts of CaCO₃ solids to water for a short period after temporarily removing a main from service. Then, water utilities can manipulate the saturation index of the supply water, such that CaCO₃ solids form only within the leak holes to stop leaks, rather than in the bulk water. On the other hand, to achieve smart clogging, water utilities can control the addition of waterborne CaCO₃ particles by pumping a few mg/L of suspended CaCO₃ through the pipeline with the flowing water, such that a portion of the suspended CaCO₃ particles can pass into leaks and seed the holes.

FIG. 5.

Autogenous repair of leaks by CaCO₃ precipitation in drinking water pipelines by smart precipitation (left) and physical clogging (right). Adopted from Tang et al. (2013).

One potential problem of autogenous repair or clogging by CaCO₃ precipitation is the long-term resistance to dissolution and longevity of the repair. It is possible that the repair materials (i.e., CaCO₃) might start to dissolve, once a water that is undersaturated with respect to CaCO₃ is pumped through the drinking water pipelines. However, since CaCO₃ dissolution is not very fast at pH ≥5 (Morse and Arvidson, 2002) and ambient temperature, the repair might be relatively long-lived, even when the water is relatively undersaturated.

Combating scaling in homes and buildings

Various water treatment devices have become commercially available for homeowners to combat scaling issues (Table 8), and many of them have a strong scientific basis. First, ion-exchange softeners are designed to exchange Ca²⁺ with Na⁺, reducing the dissolved Ca²⁺ concentration and scaling potentials in drinking water. Second, membrane treatments (nanofilters and RO) can remove dissolved Ca²⁺ and inorganic carbon to reduce scaling potential, without introducing net extra salinity into the drinking water systems while considering both reject water and treated water streams. Third, surface-catalyzed crystallization induces CaCO₃ precipitation during the treatment process itself, so that future scaling potential can be greatly reduced. Finally, reducing water heater temperature can be easily achieved by homeowners, and it reduces scaling potentials by making all CaCO₃ solids more soluble at lower temperatures.

Table 8.

Effects of Common Scaling Prevention Practices Available for Homes and Buildings

Scaling prevention strategy	Advantages	Disadvantages	Mechanisms	References
Ion-exchange softener	Frequently used; softens water	Regeneration cycles add large amounts of salt to sewage and environment	Remove dissolved Ca	Edzwald (2011)
Nanofilters and RO membranes	High Ca²⁺ rejection (>90%); no net salinity addition to sewage	Energy intensive; large amount of water is wasted	Remove dissolved Ca	Edzwald (2011)
Dutch pellet softening method	More effective than other alternative strategies	Water must initially be supersaturated with CaCO₃	Remove dissolved Ca	Schetters (2013)
Reduce water heater temperature	Lowers scaling potential in water heater; greater savings on energy	Increases likelihood of pathogen formation (i.e., Legionella pneumophila)	Increase the solubility products of CaCO₃ solids	Plummer and Busenberg (1982); ASHRAE (2015)
Magnetic antifouling devices	Inexpensive CaCO₃ control method; little to no maintenance	No agreed-on mechanism of scaling inhibition; questionable credibility	Weak scientific basis	Powell (1998); Limpert and Raber (1985); Al Nasser et al. (2011)

However, researchers are still dubious about the effects of magnetic antifouling, electronic antifouling, and electrostatic devices on treating CaCO₃ scaling issues. These devices have been marketed for half a century, and vendors generally claim that these devices can reduce scaling by applying a magnetic field by using either a permanent magnet or a temporary electromagnetic field generated by electron currents (Powell, 1998). A literature review by Powell (1998) suggests that there are little scientific data to support the claims of these devices, and a systematic experimental study by Limpert and Raber (1985) proved that eight electrostatic, magnetic, and electromagnetic devices did not significantly reduce CaCO₃ scaling. However, Al Nasser et al. (2011) has observed a slight decrease in the scaling rates of CaCO₃ with the usage of electronic antifouling. They suggested that the conversion of calcite to vaterite, which is more soluble at 25°C, might contribute to the decreased scaling rates of CaCO₃.

Knowledge Gaps

This review re-evaluated CaCO₃ formation in potable water systems in the 21st century while considering global warming, autogenous repair/leak clogging, use of corrosion inhibitors, and rising water temperatures in buildings and homes. Because of these emerging issues, research on CaCO₃ formation needs to be directed toward filling the following knowledge gaps: (1) Rising water heater temperatures, calcium levels, and atmospheric CO₂ levels are changing the likelihood of CaCO₃ formation in many potable water systems. (2) Many water treatment plants have relatively recently increased the dosage of inhibitors to combat lead and iron corrosion, and the temperature set point of water heaters is now being raised in many instances to reduce the likelihood of pathogen growth. Therefore, these changes will present new challenges to predicting CaCO₃ scaling. (3) The newly discovered effects of scaling, especially autogenous repair of pipe leaks through precipitation, represents another important knowledge gap. Even though inhibitors, such as phosphates or Zn²⁺, can inhibit CaCO₃ formation, scaling may be beneficial in protecting and repairing potable water pipelines, by forming CaCO₃ in the leak holes to clog the leaks. This phenomenon could potentially alter our perceptions about the desirability of allowing CaCO₃ scaling in potable water systems.

Footnotes

Acknowledgment

The authors acknowledge the financial support of the National Science Foundation under grant CBET-1336616. Opinions and findings expressed herein are those of the authors and do not necessarily reflect the views of the National Science Foundation.

Author Disclosure Statement

No competing financial interests exist.

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A 21st-Century Perspective on Calcium Carbonate Formation in Potable Water Systems

Abstract

Abstract

Introduction

Scaling

Autogenous repair

Recent Trends Affecting CaCO3 Formation

Overview

CaCO3 formation as a corrosion control method has fallen out of favor

CaCO3 scaling in building heating systems has recently become an increased concern

CaCO3 may have the ability to clog leaks that have formed in pipes and extend the lifetime of distribution systems

CaCO3 precipitation inhibitors that are naturally present or added at the point of treatment have probably changed the propensity of water to form CaCO3 scales in the past few decades

Climate change may play a role in CaCO3 formation

Scaling has been shown to affect biofilm growth in premise plumbing

Bulk precipitation of CaCO3 may contribute to erosion corrosion in plumbing systems

Chemistry of CaCO3 Formation

Background

Precipitation and growth in plumbing and distribution systems

Homogeneous nucleation

Heterogeneous nucleation and deposition of scales

Growth rate

Dissolution

Morphology and polymorph stability

Monitoring CaCO3 Formation in Potable Water Systems

Control and Prevention of Calcium Carbonate Scaling

Control CaCO3 precipitation through water chemistry quality parameters

Saturation management (pH and dissolved Ca)

Temperature

Other methods that might be effective in controlling CaCO3 scaling

Delay or prevent CaCO3 precipitation by using scale inhibitors

Potential beneficial application of CaCO3 precipitation for autogenous repair

Combating scaling in homes and buildings

Knowledge Gaps

Footnotes

Acknowledgment

Author Disclosure Statement

References

Recent Trends Affecting CaCO₃ Formation

CaCO₃ formation as a corrosion control method has fallen out of favor

CaCO₃ scaling in building heating systems has recently become an increased concern

CaCO₃ may have the ability to clog leaks that have formed in pipes and extend the lifetime of distribution systems

CaCO₃ precipitation inhibitors that are naturally present or added at the point of treatment have probably changed the propensity of water to form CaCO₃ scales in the past few decades

Climate change may play a role in CaCO₃ formation

Bulk precipitation of CaCO₃ may contribute to erosion corrosion in plumbing systems

Chemistry of CaCO₃ Formation

Monitoring CaCO₃ Formation in Potable Water Systems

Control CaCO₃ precipitation through water chemistry quality parameters

Other methods that might be effective in controlling CaCO₃ scaling

Delay or prevent CaCO₃ precipitation by using scale inhibitors

Potential beneficial application of CaCO₃ precipitation for autogenous repair