An organism has an optimal growth rate when the hydrogen ion concentration is very high

8.5.2.3 pH

The hydrogen ion concentration (pH) influences the growth and metabolism of the microorganisms. It is very critical to monitor the pH of a contaminated site and regulate it to accomplish optimum growth of in situ PAH-degrading microbes. Since the pollutant may also influence the pH of contaminated sites, the indigenous microorganisms at the sites that do have PAH-degrading ability might not transform PAHs under these altered acidic or alkaline conditions. Moreover, many sites contaminated with PAHs are not at the optimal pH for bioremediation. Therefore, the pH of such sites should be adjusted, for instance, by the addition of lime, etc. Sometimes, in situ microorganisms at a contaminated site may be tolerant to the site conditions, but may have the potential to metabolize PAHs in suboptimal conditions. Therefore, it is also necessary to isolate and characterize the indigenous PAH-degrading microorganisms and accordingly adjust the optimum pH at the site.

Questions and Problems

1.

The hydrogen ion concentration of the soil solution from a particular soil is 3 × 10− 6 mol L− 1. What is the pH of the soil solution?

What is the soil textural class of a soil with 20% sand, 60% silt, and 20% clay?

A 100-g sample of a moist soil initially has a moisture content of 15% on a dry weight basis. What is the new moisture content if 10 g of water is uniformly mixed into the soil?

Which factors within this chapter affect the cation-exchange capacity (CEC) of a soil? Explain why.

Which factors can potentially affect the transport of contaminants through soil and vadose zone? Explain why.

How does soil moisture content affect the activity of aerobic and anaerobic soil microorganisms?

Compare and contrast surface soils with the vadose zone.

pH

The activity of hydrogen ions (as measured by pH) is of particular relevance to soil chemical reactions, processes, and soil-water chemical composition. Hydrogen ions are involved in the majority of chemical reactions in the soil environment. The rates of numerous dissolution/precipitation and redox reactions are affected by the presence or absence of H+ ions. For example, the solubility of most metal hydroxides increases with soil acidity. A simple equilibrium formula for slightly soluble solids may be written as follows:

where M and X are water soluble species.

Note that this equation is similar to Equation 13.5. Therefore a solubility product expression can be written as follows:

where Ksp called a solubility product constant.

If we define the molar solubility of each component as S, then the solubility product expression can be simplified as follows:

Note that a more general form of this solubility product expression is needed when the number of moles for each component differs, as in the generalized case of MaXb solid. Thus in case of the solid amorphous iron oxides commonly found in soils, their solubilities in water can be summarized by the following equilibrium reaction, which is the sum of Fe oxide dissolution and the dissociation of water:

Eq. 13.23Fe(OH)3=Fe3++3OH -

Assuming that the activity of the iron oxide solid is ≃ 1, then Equation 13.23 can be written in terms of solubility product expression:

Eq. 13.24[Fe3+ ][OH-]3=K0s≃10-38

Also, assuming that the concentration of all species in solution are the same (Fe3+ ≃ OH–)= nS, where S is the solubility and n is the number of moles of each species, Equation 13.24 can be simplified to:

If the activity of OH–is decreased ten-fold (one pH unit), Equation 13.25 indicates that the activity of Fe+++ species in the soil-water environment will increase by 103 fold because the right side of the equation cannot change in value (Figure 13.4).

A more general relationship to estimate the solubility of most metal hydroxides and oxides as a function of pH can be derived by combining the equilibrium reaction for metal oxides or hydroxides and acid:

Eq. 13.26M(OH)n+nH+=Mn+nH2O

and metal and water stability constants, Ks and Kw = 14, respectively, to yield the general relationship:

Eq. 13.27log[Mez+]=logKs+z(14)-z(pH)

where n is also equal to z, which is the charge of the metal ion, and pH is the soil acidity.

Additionally, most redox reactions are also dependent on the activities of H+ ions. The “Measuring pH” section provides a full description of the principles of pH measurement and a detailed discussion of this measurement.

23.3.2 Hydrogen ion

The concentration of hydrogen ion in water is represented by pH. In Bagmati River, Bhatt et al. (2013) reported pH variation from 6.6 to 7.1 with higher value in the downstream urban areas. The higher value of pH toward downstream is attributed to the higher production of bicarbonate due to greater organic inputs and high residence time. The pH increased after Gaurighat (near Pashupatinath) possibly due to religious rituals. Bhandari et al. (2017) also reported increase in pH as the river moves toward the urban area. In the river, the highest pH was 6.94 in winter season and lowest (6.68) in rainy season (Mehta and Rana, 2017). The permissible range for pH for bathing water is between 6.0 and 9.0 (EEC, 1978) and for the fisheries is between 6.5 and 8.5 (EMECS, 2001). The pH values in the river seem increased over the period of time (Table 23.2).

Table 23.2. Physio chemical parameters at different places at different time period.

3.4.1.5 Hydrogen Ion Concentration

The concentration of hydrogen ions plays a vital role in any biochemical reactions. The growth of microorganisms varies widely, from pH 1 to 2 at the acid end to pH values between 9 and 10 in alkaline lakes and soils. Each microorganism has a definite pH range in which to grow. Acidophiles have their growth optimum between pH 1 and 5.5; neutrophiles between pH 5.5 and 8.0; alkaliphiles prefer a pH in the range of 8.5–11.5 (Figure 9).

Extreme alkaliphiles have their growth optima at pH 10 or higher. Most bacteria are neutrophiles. Most fungi prefer slightly acid surroundings, approximately a pH in the range from 4 to 6; algae also seem to prefer slight acidity (Figure 10). However, there are many exceptions to these generalizations. For example, the archaeon S. acidocaldarius, a native to acidic hot springs, grow well in a pH of approximately 1–3 at a high temperature.

Alkaliphilic microorganisms are usually found in highly basic habitats such as soda lakes and soils high in carbonates. However, the most well-studied alkaliphilic bacteria have been aerobic nonmarine in nature and many are of Bacillus sp. Some alkaliphilic bacteria are also halophilic (salt loving), and most of these are archaea. Some alkaliphiles have found industrial uses because they produce hydrolytic enzymes, such as proteases and lipases, which function well at alkaline pHs and are used as supplements for household detergents.

The range of pH through which a particular enzyme can act effectively is usually quite narrow. Most biological processes employed in wastewater treatment involve the use of soil organisms operating in mixed cultures. The enzyme systems of these organisms are adapted to operating in essentially neutral solutions; therefore, it is important that pH be controlled over a rather narrow range of approximately 6–9. In a batch culture, the pH can change during growth as the result of metabolic reactions that consume or produce acidic or basic substances. Thus, the control of pH is necessary by means of adding chemicals called buffers. Different buffers must be used to buffer at different pH values like K2HPO4 and KH2PO4 for near neutral pH ranges (i.e., pH 6–7.5).

Introduction

Glass electrodes sensitive to hydrogen ions are the most commonly used sensors in chemistry and related disciplines. They belong to the group of potentiometric membrane sensors and are constructed in various configurations, depending on the application. Their basic design and properties are described and discussed in this article. Among their important properties is the selectivity, which depends on the composition of the glass. Glass electrodes are used mainly for pH measurements, but they may compose a part of more sophisticated systems in gas sensors or enzymatic sensors. Recently glasses based on a nonsilica structure have been used in potentiometric measurements.

Cation Exchange

Cation exchange can increase alkalinity whenever hydrogen ions in solution exchange on surfaces for base cations. The effect is generally reversible, and thus the process may not contribute to long-term increases in alkalinity once the cation exchange sites are depleted. In fact, cation exchange can act in reverse if base cations are added from sea spray or road salt to a soil solution, causing temporary acidification and loss of alkalinity. Nevertheless, soils with large cation exchange capacities can act as a large buffering reserve against relatively short-term acidification events.

pH

As the negative log of the hydrogen ion concentration, pH is intimately connected to water in the soil and the soil organisms. With regard to soil, wherever rainfall is abundant, water leaches bases from the soil and the soils are predominantly acid; wherever rainfall is scant, there is not enough water to leach the bases from the soil and the soils are predominantly alkaline. Because soils have a pH-dependent charge originating from organic matter and the broken edges of clays, this charge decreases as the soil becomes more acid, reducing not only the cation exchange capacity of the soil, but also the ability of the soil to retain nutrients. Reducing the pH does not necessarily reduce the entire cation exchange capacity in soil; some charge may be associated with isomorphous substitution in clay minerals and this charge is pH-independent. With regard to soil organisms, when soils are saturated, the hydrogen ion concentration is diluted; when soils are dry, the hydrogen ion concentration is concentrated and this concentration affects the ability of soil organisms to control anions and cations and offset high osmotic potential.

Although some bacteria and soil fauna tolerate low pH (acidophiles; able to grow at pH 3 or lower) or high pH (alkaliphiles; able to grow at pH 9 or higher), most bacteria and soil fauna are neutrophiles and prefer a near-neutral pH range of 6–7. The same is true of soil algae, which again are found in acid and highly alkaline environments, but prefer pH ranges of nearly neutral to slightly alkaline. Many bacteria and algae prefer this pH range because soil nutrients are most available. Furthermore, a neutral pH avoids heavy metal toxicity, because almost all these metals (a density of greater than 5 Mg m−3) are only active in acid soils (pH <5). In contrast, fungi are acid-tolerant (pH 4–7) and tend to be abundant in acid soils, where there is less competition from bacteria.

7.2 pH

pH is a measure of the hydrogen ion concentration in solution and is also referred to as the degree of acidity or alkalinity. As a sample's pH changes, many precipitation, coprecipitation, and sorption processes can occur that alter the sample's chemical composition and reaction rates. Biological processes of a sample are also influenced by its pH. Changes in the dissolved gas content of a sample can alter the pH. Groundwater is generally in equilibrium with CO2 at a partial pressure several times that of the atmosphere. On exposure to the atmosphere, this CO2 escapes and the pH rises. It is therefore important that pH is measured on-site. The pH is generally measured using a combination electrode calibrated using standards of pH 4, 7, and 10. In addition, it is recommended that a dilute acid solution of known pH is used as a control check and that additional control standards are also checked if the pH is outside the range 4–10.

4.3.1.1 Examples of Extractants and Separation of Metals

There are mainly two classes of extractants: acidic and chelating extractants.

Acidic extractants are those in which hydrogen ions or protons of the extractant are exchanged for metal ions, as per the following equation:

(4.17)Mn++nHA↔MAn+nH+

The equation implies, the reaction is governed by the pH of the system. As will be discussed further, pH is the principal control variable, which influences both extractability of metals as well as separation or two metals by acidic extractants. The percent extraction of metals is usually given at the equilibrium pH of the solution. In addition, for the purpose of comparing the extraction of various metals by a specific acidic extractant, the order of extraction is expressed in terms of pH1/2, (also denoted by pH50 or pH0.5) which is defined as the pH at which 50% of the metal in the aqueous phase is extracted into the organic phase. This is derived from dissociation equilibria of acidic extractants (see Ritcey and Ashbrook, 1984, p. 20).