The effect of pH on taste and odor production and control of drinking water

Taste and odor (T&O) compounds that can affect the aesthetics of drinking water come from a variety of system inputs, such as surface water, groundwater, source waters impacted by wastewater, water treatment processes, and the water distribution system. The United States Environmental Protection Agency (US EPA) has established guidelines, or Secondary Maximum Contaminant Levels (SMCLs), for T&O compounds in drinking water. One of the SMCL guidelines applies to pH because pH itself can affect the taste of drinking water. This is why the US EPA SMCL for pH is based on its effect on taste. At a pH below 6.5 a bitter taste can sometimes be produced and at a pH above 8.5 a baking soda-like or mouth coating feeling can occur (Burlingame et al. 2007). Thus, pH affects the speciation of carbonates in water which can have a positive or negative impact on the water's taste and mouthfeel.

The desired pH range for drinking water, as influenced by the typical range of human saliva, is between 6.0 and 9.0 (Pedersen et al. 2002). Surface waters in North America can range from 6.5 to 8.5 and groundwaters from 6.0 to 8.5 (Saalidong et al. 2022). Schock et al. (2014) reported that systems are more likely to distribute water with a pH above 8.0 to protect the infrastructure and plumbing systems from corrosion, however, chlorinated systems (systems carrying a free chlorine residual) are more likely to distribute water with a pH between 7.0 and 8.0.

The effect of pH on the actual generation and control of T&O compounds is often overlooked, while there is general agreement that pH is an important parameter in drinking water treatment, corrosion control, and other quality control processes (Burlingame 2012). pH is a complicated variable that directly and indirectly controls dissolved chemical speciation and consequently exerts both subtle and dramatic effects on tastes, odors, mouthfeel, color/particles, and corrosion by-products (Dietrich & Burlingame 2014). This includes odorous organic and inorganic chemicals from piping materials (Khiari et al. 2002). Within each section, this article provides a comprehensive discussion and more in-depth look at the ways pH can be expected to affect T&O occurrence and control. In addition, the impacts of pH on the analyses of various T&O compounds are discussed. It also highlights how research and information gathering could lead to better guidance on optimizing pH for T&O treatment and control.

Water is considered a ‘universal solvent’, i.e., just about everything it comes in contact with becomes dissolved to some degree. Once dissolved in water, a chemical can stay in solution, precipitate out, chemically react, or escape to the atmosphere. Chemicals exist in solution in a state of equilibrium that is affected by their environment, including pH. Changes in pH can make a chemical either more or less reactive, more volatile in the liquid phase, or even change speciation. The pH of about 95% of naturally occurring waters (lakes, reservoirs, streams, groundwater) are within the range of pH 6–9 as controlled by the bicarbonate buffering system. This is the desired pH range for drinking water according to the US EPA SMCL. Compounds whose solubility is inversely related to pH will precipitate out if pH is increased (Snoeyink & Jenkins 1991).

In surface waters, an algal bloom can raise the pH of water from <7 to >9 (Smith 2019; Adams et al. 2022). This change in pH could increase the odor intensity (e.g., fishy, musty, grassy) due to the hydrolysis catalyzed by bases. Changes in pH in source water can also have indirect effects on T&O. For example, altering the pH of the source water containing an algal bloom could indirectly affect the chemicals causing a T&O problem. If pH-sensitive organisms are killed, intracellular T&O compounds are released, which could exacerbate T&O problems (Adams et al. 2021b). Also, anaerobic odors, such as organic sulfides, can appear following the die-off of the algal or cyanobacterial bloom and cause a T&O problem. In turn, the sulfides are themselves affected by the pH of the water (Buerkens et al. 2020). Watson (2010) provides a comprehensive list of these odorous compounds. The fishy odors in untreated surface waters can range from cod liver oil-like, clam shells, dead crab, rotten fish, ammoniacal, to fishy (Suffet et al. 1996), and thus could be affected by changes in pH.

Many odorants in source water are derived from the natural decay of organic matter. The decay of grass in water has been studied because decaying vegetation odors have been a common concern for drinking water supplies (Buffin 1992; Young & Suffet 1997; Khiari et al. 1999). For example, freshly cut grass has a sweet grassy odor due to cis-3-hexenyl acetate. This chemical has an odor threshold of around 50–70 μg/L (Young & Suffet 1997), and occurs upon the decay of plant materials and within some green algae. This can hydrolyze to cis-3-hexen-1-ol (sharp, grassy vegetation odor) which has a much lower odor threshold at around 2 μg/L. At pH 6.1, no hydrolysis occurs, whereas at pH 7.8, hydrolysis can be significant. Thus, hydrolysis can be a significant factor in the speciation of odorant chemicals as well as in their increase or decrease in concentration.

In a wetland or similar environment, diurnal fluctuations in iron can occur according to variations in dissolved oxygen (DO) and oxidation-reduction potential (Goulet & Pick 2001). For example, DO might be lowest at night along with pH due to oxygen consumption from aquatic plants, bacteria, and algae, which encourages the release of iron into solution. Nighttime, anoxic conditions in sediment are favorable for the reduction of ferric iron (Fe³⁺) to ferrous iron (Fe²⁺). In turn, photosynthesis during the day increases oxygen and pH, thus precipitating out iron. Kopáček et al. (2006) proposed that aluminum and iron could be photochemically liberated, wherein soil organic-bound metals enter the aquatic environment. Aluminum and iron are photochemically released as inorganic ions. Then they hydrolyze and precipitate as oxyhydroxide particles, perhaps even binding trace metals and phosphorus along with them. The surface charge of iron oxide particles is affected by the pH of the water which in turn affects the adsorption of other metals and organic matter.

The movement of dissolved metals into and out of sediments occurs from the bulk water, through the sediment–water interface, into the pore water, and onto the solids (and the reverse). Berg et al. (1998) reported that the pore water at the interface can have the highest metal contamination, which is sensitive to many factors and environmental changes. Dissolved Fe²⁺ concentrations can be high in the pore water of sediment when the sediment is anoxic. Biological activity can degrade organics that complex metals and thereby release the metals to the pore water. Therefore, the pH and redox conditions of the pore environment are important to understand. As pH shifts, redox reactions and metal concentrations fluctuate, which can create T&O issues.

Other factors that affect T&O include organic matter degradation; complexation with dissolved organic and inorganic ligands; changes in surface properties of adsorptive particles; formation and dissolution of precipitates; dissolution of hydroxides and release of metals; presence of sulfides (Berg et al. 1998). As dissolved metals are released into the pore water and move out into the interface, they can be scavenged by organic complexes and carbonates. At a lower pH, soluble organic-complexed iron is found (of which two-thirds of the iron is organic complexed). While the organic complexes tend to vary greatly in size, the sulfide complexes are usually smaller and thus more easily stirred up and mobilized. A stream's bulk water interaction with the subsurface stream bed is significant and is where hyporheic exchange occurs (Ren & Packman 2005).

Perhaps the most important reason for understanding iron fate and transport in a waterway (especially a polluted waterway), is that iron has been associated with the fate and transport of other contaminants such as arsenic, lead, and copper and to nutrients such as phosphate. Iron in aquatic sediment is a complex mixture of oxides/hydroxides, carbonates, sulfides, and silicates (Mudroch et al. 1999). Most metals are organically bound and in the silt fraction and iron surfaces have an affinity for the adsorption of other metals. Iron plays an important biological role in nature, influencing algal nitrogen metabolism, chlorophyll synthesis, oxygen cycling, and sulfate reduction (Hutchins 1995).

Iron by itself has no known toxic effects at the levels found in most surface waters, unlike other trace elements (e.g., copper, zinc) that are important micronutrients at low concentrations but harmful at elevated levels. There has been debate regarding the use of iron as a fertilizer in ‘high-nutrient, low-chlorophyll’ areas because it can stimulate phytoplankton growth and create a sink for carbon dioxide (Chisholm et al. 2001). Finally, the precipitation and dissolution of iron–phosphate minerals are important in controlling the mobility of phosphate, which is an important nutrient in surface waters.

Sulfurous and nitrogenous odors and their control have been important in wastewater treatment and, therefore, provide insight into the occurrence and control of odors in source water impacted by wastewater. This includes reclaimed water for potable reuse. Many of the odorous compounds arising in wastewater come from the anaerobic decomposition of sulfurous and nitrogenous organic chemicals. Wastewater also has an abundance of fatty acids, fatty acid esters, and soluble acids that are odorous as described on the T&O wheel (Ginzburg et al. 1999). The ability to strip hydrogen sulfide (H₂S) and ammonia (NH₃) from odorous air emissions during wastewater treatment is dependent in large part on the pH of the wastewater. For example, at acid and neutral pH, NH₃ is ionized to ammonium (NH₄⁺) which as an ion in water cannot be stripped, while above pH 9 or 10, it forms molecular NH₃ and becomes less soluble and more amenable for stripping (Hand et al. 2011). The opposite is true of H₂S which can be easily stripped at low pH, but when pH is adjusted to 8–9 it exists almost solely as HS⁻, eliminating the rotten egg odor because HS⁻ is a non-odorous non-volatile (Qian & Lancaster 2013).

A study to show that low pH favors H₂S odors, while high pH favors the amines and ammoniacal odors was done on samples from four different wastewater processes: preliminary treatment process influent, two different primary sedimentation tank effluents, and a secondary biological treatment process effluent. Sample pH was adjusted to low or high using 6 N sodium hydroxide (NaOH) or sulfuric acid (H₂SO₄). Approximately 1-L samples of wastewater were collected from these four treatment processes and placed on magnetic stirrers for gentle mixing to keep the solids in suspension. Table 1 shows results determined by two trained Flavor Profile Analysis (FPA) panelists. This simple demonstration of the effects of pH was done at the Philadelphia Water Department's Northeast Water Pollution Control Plant.

The primary odor species causing the shift in the odor profiles of the wastewater samples, when the pH was changed, are sulfurous and nitrogenous compounds. Examples of such sulfurous compounds could include the following: allyl and phenyl mercaptans for a garlic odor; dimethyl sulfide, ethyl mercaptan, and methyl mercaptan for a decayed cabbage odor; hydrogen sulfide for a rotten egg odor (WEF 1995). Examples of nitrogenous compounds could include the following: n-butylamine, ethylamine, and ammonia for ammonical odors; dibutylamine, diisopropylamine, methylamine, dimethylamine, and trimethylamine for various types of fishy odors; skatole and indole for fecal type odors (Table 2).

One way to reduce the odorous aspects of a wastewater that has a sulfidy odor problem and a low pH would be to raise the pH. Sodium hydroxide (caustic soda) or calcium hydroxide (lime, Ca(OH)₂) can be added to raise the pH and reduce the odor due to H₂S. In contrast, decreasing the pH by 0.5 units can cause a significant increase in the sulfidy odor (WEF 1995). On the other hand, at pH 11–12, ammonia and amines are in the neutral form and are highly volatile. Therefore, pH control needs to be done in a careful manner. The control of odorous compounds during wastewater treatment will affect the odorous compounds discharged to the receiving waters.

T&O can also be problematic in groundwater sources. Biotic and abiotic processes occur in natural geologic formations, using many different precursors to produce a variety of sulfurous compounds (Brimblecombe 2004). These can be found from minerals (e.g., pyrite, galena, gypsum, barite, etc.) and ores (e.g., copper, zinc, nickel, etc.). When sulfurous compounds arise from biofilms found in the well itself, shock chlorination can be used to control the biofilm. This is more effective at a lower pH as discussed later. Effective removal of sulfides is important so that H₂S does not reform later and so that the biological conversion to organic sulfides does not occur (Burlingame 2020).

Groundwater low in pH and alkalinity is corrosive and is problematic for treatment and distribution if the pH is not adjusted to produce water that is slightly carbonate scale-forming above a pH of 7.5. If left in its corrosive state, metals such as copper and iron can dissolve from plumbing materials into the water and cause staining and flavor problems. This will be discussed in greater detail later (Omur-Ozbek & Dietrich 2011). Iron and manganese are naturally found as groundwater contaminants. Iron can produce a metallic flavor and is found most frequently in groundwater with a pH < 7.0 (Swistock & Sharpe 2019). Sequestration can be used to remove these metals, and as with biofilm control, it is most effective at a lower pH (Kirby 2020). However, oxidation is more effective at a higher pH if sequestration is not desired, so care must be taken to understand how to mitigate the T&O effects of iron and manganese.

Another T&O compound that has been found in groundwater sources is methyl tert-butyl ether (MTBE). An example of the extensive locations of MTBE pollution has been documented in California by Shih et al. (2004). This organic solvent was used as a gasoline additive prior to its phasing out in favor of ethanol in the early 2000s. Groundwater with a low pH (<7) has been shown to have a higher prevalence of MTBE contamination due to the influence of surface runoff (Ayotte et al. 2004). Higher pH (>7) groundwater sources are considered older and thus less impacted by surface runoff. Many state health departments have fact sheets and advisories available for T&O in groundwater, including MTBE and other solvent/fuel/gasoline contaminants of public health concern (WaDoH 2018). In this case, pH may be one indicator of susceptibility of aquifer's contamination.

The change in pH can affect drinking water T&O in three ways:

• Changes can occur in the equilibrium of chemicals, favoring either more or less volatile species.

• Chemical reaction rates during treatment can increase or decrease as pH changes, changing the chemicals, which can result in the production of either more odorous or less odorous products.

• Hydrolysis, the decomposition of a chemical by the chemical reaction with water, can occur and change the chemicals to more odorous or less odorous products.

The effects on water alone are significant since the proton (H⁺) and hydroxyl ion (OH⁻) are common catalysts in aqueous systems. The reactions catalyzed by H⁺ are called acid-catalyzed reactions or acid catalysis, while those catalyzed by OH⁻ are called base-catalyzed reactions or base catalysis. These reactions are impacted by pH (Snoeyink & Jenkins 1991). When it comes to the taste of water, carbonate chemistry can be important. Carbonates and bicarbonates are strongly influenced by pH, and acid or base catalysis.

Since pH is such an important factor in T&O production and control in source water, it has the potential to be important in the treatment of drinking water as well. A sample from a pre-sedimentation basin for a drinking water treatment plant on a river supply was collected and the pH was described in the same way as in Table 1. Table 3 shows odor test results determined by two trained FPA panelists. The analysis was done at the Philadelphia Water Department's water treatment plant's raw water basin (Table 3). This study shows that a higher pH will favor the volatilization of nitrogen-based odors such as those that can impart fishy odors.

Drinking water can be significantly impacted by pollutants that make their way through water treatment or that are changed by water treatment. The Medicinal/Phenolic category of the T&O Wheel includes chlorophenols. When phenols are chlorinated, they form chlorophenols. The formation of odorous chlorophenols favors a pH above 8. Lee (1967) reported that the Threshold Odor Number (TON) increased dramatically upon the chlorination of phenols when the pH was increased from 7 to 8, and then to 9. In many cases, the oxidation of contaminants, including phenols, is enhanced at higher pH values due to the more rapid generation of hydroxyl radicals (Singer & Reckhow 2011). This can be problematic when compounds like chlorophenols are formed more rapidly at higher pH. For example, Adams et al. (2021a) showed that the rapid formation of 2,4,6-trichloropehnol at high pH could be converted to 2,4,6-trichloroanisole by actinomycete colonies in filter media. Potassium permanganate was used as an oxidizing agent to remove the bacterial colonies and T&O, which acts as a stronger oxidizer and is more effective at T&O control at low pH (Adams et al. 2021a).

An example of a synthetic odorant that makes its way through treatment is 2-ethyl-4-methyl-1,3-dioxolane (2-EMD). It is a byproduct of resin manufacturing and has been responsible for odor episodes. This compound has an odor threshold concentration (OTC) at 5–10 ng/L and smells ‘medicinal, sweet’ (Schweitzer et al. 1999). At a pH of 9 it is stable, but at a pH of 3 it hydrolyzes in hours. Thus, at a neutral pH it has limited stability.

Among the chlorinous/ozonous odors in the T&O Wheel, it is well known that pH affects the chlorine species that are formed when water is disinfected. The pH of the water affects both the species of disinfectant that is formed, which can affect the T&O of the water, as well as the disinfection by-products (DBPs) that are formed. The impact of pH on chlorine chemistry is important because most consumer complaints and concerns over tap water T&O relate to the odor of the chlorine residual (Omur-Ozbek 2012). At pH 6, hypochlorous acid (HOCl) predominates, and is a stronger disinfectant than hypochlorite (OCl⁻). HOCl is >80 times more effective for the disinfection of coliforms than OCl⁻ (Snoeyink & Jenkins 1991). At pH > 8.5, OCl⁻ predominates. The OTC for HOCl is about 0.3 mg/L, whereas the OTC for OCl⁻ is about 0.4 mg/L (Krasner & Barrett 1984). With free chlorine, the pK_a is 7.5. This is where HOCl and OCl⁻ are at 50% equilibrium, depending slightly on water temperature and other factors. In the pH range of 6–9, dramatic changes in speciation occur across little changes in pH.

The formation of chloramines as a final residual in drinking water is heavily influenced by the pH of the water (Wang et al. 2017). Monochloramine is the preferred species for purposes of a final residual due to the less preferred T&O associated with dichloramine, and because monochloramine is a more persistent disinfectant residual (Kirmeyer et al. 2004). During chloramination, mostly monochloramine is produced at pH > 7. At pH 5–6, there can be up to 40% of the total residual as dichloramine. At pH 6–7, the ratio decreases to around 20% dichloramine. A target goal is to keep the dichloramine percentage under 20% of the total chlorine residual and less than 0.5 mg/L.

The OTC for monochloramine is around 0.6 mg/L, whereas for dichloramine it is around 0.2 mg/L (Krasner & Barrett 1984). The effect of pH on chloramine formation is very dramatic in the pH range of 6–9. A lower Cl₂/N molar ratio is generally better for minimizing dichloramine (Wang et al. 2017). Kajino et al. (1999) observed a strong correlation of chloramine odor intensity with pH during breakpoint chlorination. They found an odor at pH 6.5, but not at 8.3 during the reaction of chlorine with nitrogenous compounds, including NH₃. The odor intensity increased at pH 3.0. A lower pH results in some species being more soluble, which can also affect the T&O.

A study on dichloramine was conducted for 1 year at the Philadelphia Water Department's water treatment plants using FPA panelists with analysis being done weekly (Table 4). Dichloramine has a stronger chlorinous-type odor than monochloramine, and therefore should be minimized during the formation of chloramine as a residual for the drinking water. Customers can complain about drinking water when the chlorinous odor is too noticeable. Three drinking water treatment plants in Philadelphia were compared on their effluent water quality. Plant #1 was using an elevated pH for corrosion control in the distribution system. This plant also had lower occurrences of chlorinous T&O, and less dichloramine was detected.

The desire to better control the precursors of DBPs has led to requirements for enhanced coagulation. The removal of total organic carbon (TOC) by coagulants, ferric and aluminum based, are optimized at low pHs. Treatment plants that reduce the pH of the incoming water to around 6 and then increase it back to 7 or higher, even over 8, for optimized corrosion control treatment, would be affecting the creation of dichloramine.

If two waters are blended, one with a free chlorine residual of 0.2 mg/L and the other with a chloramine residual containing low concentrations of NH₃ (<0.04 mg/L), an uncontrolled advance to breakpoint chlorination occurs where the ratio of Cl to NH₃ will be greater than 5:1. This favors dichloramine formation. Mixing can also have an effect when a pH change occurs as a result (Health Canada 2016). If a water is predominant in monochloramine and only 5% dichloramine at pH 8.0, a decrease to just 7.5 will result in the dichloramine fraction increasing up to 25% (Evins et al. 1990).

Finally, the conditions of chlorination can affect the control of other issues such as with iron and manganese. The addition of chlorine to water, depending on the form used, can lower or increase the pH depending on alkalinity (Kirmeyer et al. 2002). The reaction of chlorine with dissolved Fe(II) is rapid at pH > 7, so sodium hypochlorite would be preferred.

Another important aspect of pH in water treatment and T&O control is the oxidation of geosmin and 2-methylisoborneol (MIB), the primary causes of earthy–musty odors. Both geosmin and MIB can be degraded by ozone (O₃) more rapidly at a higher pH (>9) (Westerhoff et al. 2006). Hydroxyl radical (·OH) reactions support their degradation. The combination of O₃ and hydrogen peroxide (H₂O₂) can enhance the removal of these T&O compounds. If the source water contains a high level of natural organic matter (NOM) and a high alkalinity (⁠ and ions are scavengers of hydroxyl radicals), O₃ decay is fast at the higher pH, and the decay releases hydroxyl radicals. However, the relative concentration of hydroxyl radicals available for oxidation drastically decreases when the reaction rate of hydroxyl radicals is fast with scavenger ions (⁠⁠) compared to MIB and geosmin. Westerhoff et al. (2006) observed that highly alkaline waters will have a negative impact on the formation of the hydroxyl radical.

As stated in earlier sections, the metals, iron, and copper have significant impacts on taste and flavor-by-mouth. These impacts are common for drinking water distribution systems as corrosion by-products, such as with the release of iron from corroding unlined cast iron pipe, and for building plumbing such as with the corrosion of copper pipe. Thus, drinking water distribution systems and downstream plumbing can be sources of T&O. These metals can appear as reduced or oxidized species in water and pH plays a role.

The taste effect of pH can be driven by changes in ion speciation as pH changes (Burlingame & Anselme 1995). The rusty taste of water is common and is caused mostly by a high concentration of iron at a low pH. At a low pH, ferrous ion [Fe²⁺] is highly soluble even in the presence of DO. However, when the pH is raised above 7, the oxidation of ferrous to ferric [Fe³⁺] iron occurs rapidly. This oxidation rate is dependent on the oxygen dissolution from the gas phase (Snoeyink & Jenkins 1991). Due to its relatively low solubility, Fe³⁺ precipitates out and reduces the metallic taste (Glindemann et al. 2006). The flavor description of metallic is experienced as flavor-by-mouth (Lawless et al. 2004; Shepherd 2006; Tucker et al. 2007). Omur-Ozbek & Dietrich (2011) showed that the metallic sensation has an odor component. Due to the common occurrence of dissolved Fe²⁺ and Cu²⁺ in water systems, and that the solubility of metals is driven by pH, this should be taken into account when studying T&O effects from metals.

Metal ions undergo hydrolysis in water and form complex with OH⁻. As pH increases, so does the percentage makeup of the hydrolyzed species. At pH 6–8, most metal ions are hydrolyzed to some extent. At pH 9, metals (copper, iron, manganese, zinc, aluminum) are significantly hydrolyzed or complexed with OH⁻. This is why a higher pH is often used in coagulation/flocculation (forming ferric or aluminum hydroxides). This is also why a low pH can be problematic with copper and lead dissolution from piping and why some water companies have gone to elevated pH to reduce corrosion (Furatian 2019). The flavor of copper is more noticeable when the copper is soluble, which typically occurs at lower pH values (Dietrich et al. 2008). The backflow of carbonation in restaurants can cause copper toxicity and illnesses and taste complaints because the low pH leaches high levels of copper out of the premise's plumbing (Damaso 2021). This complexation of metals is itself complex and depends on other factors such as organic matter content and alkalinity.

There are many examples of how pH affects the stability of T&O compounds for accurate and reproducible laboratory analyses. This is a critical point to understand. For example, the extraction of halogenated phenols has to be carried out at an acidic pH to maintain their molecular structure (Lee 1967). Hsieh et al. (2012) found, under neutral pH and alkaline conditions, no significant change in the concentrations of geosmin and MIB. Under acidic conditions, however, the dehydration of MIB occurs which results in products with different odor profiles than that of the parent compound (Hsieh et al. 2012). Suffet et al. (1999) reported that there is a pH dependence on the stability of the fishy odorant, trans,trans-2,4-heptadienal. Since this compound is more stable at a pH around 9, samples should be collected and preserved at a high pH and refrigerated prior to analysis. However, Pochiraju et al. (2021) found that this compound is more stable in raw water when preserved at a pH of <2, but more stable in finished water at its original sample pH than when it was preserved at a pH of <2. Beyond collection and preservation, pH can also play a role in the efficiency of extraction of T&O compounds.

Recent advances in chemical analysis of T&O compounds have begun to look at how pH can affect the chemistry of T&O compounds and how changes in pH can affect sample storage and hold times. Pochiraju et al. (2021) studied 18 T&O compounds for sample stability, hold time, and preservation, and provided an example of volatility in relation to changes in pH. This improvement in the understanding of how pH affects T&O compound chemistry has the potential to drive how the water industry applies pH in treatment to control T&O. Zhu et al. (2022) further advanced the current knowledge base by performing a review of biogenic T&O, which included providing known biogenic sources and their corresponding optimal pH ranges for growth and confirmed compound production.

There are T&O compounds that can be affected by pH that are categorized in the T&O wheel (Mallevialle & Suffet 1987; Suffet et al. 1995) for which the updated version can be found in chapter 2 by Suffet et al. (2019), in the book edited by Lin et al. (2019). The Earthy/Musty/Moldy category includes MIB which under acidic conditions can undergo dehydration reactions and convert to non-odor-causing compounds. The Grassy/Hay/Straw/Woody category includes grassy odorants, that at levels of pH typically found for drinking water, undergo hydrolysis. The Fishy/Rancid odors include fishy chemicals that are most stable at high pH, and the formation of aldehydes formed during ozonation at high pH (Suffet et al. 1999).

Therefore, it has been documented that pH affects the stability and speciation of some T&O compounds in water. By providing more specific examples of how pH can affect T&O treatment and control, this information can be applied to enhance the control of T&O compounds in drinking water.

When studying treatment alternatives, the optimization of T&O control using pH should be considered. This can provide an added benefit that makes one alternative better than another. Care must be taken, however, as the positive impact of pH on one treatment objective may compromise another one. It is important for water utilities to obtain a detailed understanding of source water quality that provides for monitoring for any changes in chemistry and biology of the source and preparing for the potential treatment challenges (including T&O events). As an increase is being seen in the occurrences of geosmin/MIB and cyanobacterial blooms, and with wildfires and floods that can change the NOM and T&O character of source water (Wang et al. 2017), consideration of the effects of pH must be taken into account for optimum T&O control as this paper shows. Consequently, more research is needed on how treatment and finished water pH can affect T&O. Therefore, T&O control can be an important factor when making alternative treatment decisions.

The purpose of this article is to understand more completely the various ways that pH can influence T&O production and control for drinking water and to encourage advancement in the state of the science of pH optimization for T&O control. pH can contribute directly and indirectly to production of T&O and to the better control of T&O. In addition, pH can be an indicator of conditions that favor or do not favor T&O production. The information gathered for this paper, for the first time, addresses the effects of pH on T&O control for drinking water in a consistent manner. This paper should serve as a cornerstone for future research to better understand the role and significance of pH in T&O control.

In summation, anytime a treatment process is changed, including with the addition of new sources or for the blending of different waters, steps should be taken to understand how the chemistry will affect the process. pH should be monitored from the source water through the treatment process, and into the distribution system (Kirmeyer et al. 2002). Considerations include the following:

• Choose which oxidant to use that incorporates the benefits of pH.

• Determine whether pH should be adjusted before or after oxidation.

• Determine whether a lower pH in the coagulation process will provide odor control benefits.

H.A. conceptualized and updated the original draft, wrote, reviewed, and edited the article; G.B. conceptualized the original draft, updated, wrote, reviewed, and edited the article; K.I. and L.F. updated the original draft, wrote, reviewed, and edited the article; and I.H.(M.)S. conceptualized the original draft, updated, wrote, reviewed, and edited the article.

All relevant data are included in the paper or its Supplementary Information.

The authors declare there is no conflict.