Soil Acidity and Alkalinity

Quantitative Insignificance

Active acidity is an extremely small pool. A simple calculation is enough to confirm this. Recall that pH is the negative logarithm of H⁺ concentration. For example, in soil with pH 4, the H⁺ concentration is 10⁻⁴ mol/L, which equals 0.0001 mol/L. If we consider a topsoil layer 15 cm thick over 1 hectare (about 2 million kg of soil) at 20% moisture (i.e., roughly 400,000 liters of water), the total amount of active H⁺ in this volume would be negligibly small. Estimates suggest that neutralizing all the active acidity in such a layer would require only about 2 kg of pure calcium carbonate (CaCO₃) (Weil, 2017). However, in practice, raising the pH of such soil from 4 to 6.5 may require 5–10 tons or more of lime. This vast difference indicates that the bulk of the acidity is hidden not in the solution but in the solid phase—this is the so-called potential (reserve) acidity, which we will turn to in the next chapter.

Measurement Methods and Their Interpretation

Determining the pH of a water suspension (pH_H2O) is the most common method, but it has significant drawbacks. Water is a weak electrolyte, and its ionic strength is low. Adding water to soil dilutes the soil solution, shifting the equilibrium between ions in the solution and on the colloid surface. As a result, some H⁺ and Al³⁺ held on exchange sites move into the solution, and the measured pH may be somewhat higher than that existing under natural field moisture conditions.

To obtain a more stable and reproducible result, especially in soils with high salt content or variable charge, pH is measured in a suspension with a salt background—most often in 0.01 M CaCl₂ (pH_CaCl2) or in 1 M KCl (pH_KCl) (Weil, 2017; White, 2006). The cations of the salt solution (Ca²⁺ or K⁺) displace some of the exchangeable H⁺ and Al³⁺ from the colloids into the solution, leading to a decrease in pH compared to the water extract. Typically, pH_CaCl2 is 0.3–0.6 units lower than pH_H2O (Essington, 2015). The difference between these indicators indirectly indicates the magnitude of the exchangeable (potential) acidity. In some highly weathered acidic soils where iron and aluminum oxides with a positive charge (anion exchangers) predominate, pH_KCl may, conversely, be higher than pH_H2O due to the displacement of hydroxyl ions (OH⁻) (Bloom & Skyllberg, 2012; van Raij & Peech, 1972). This is a vivid example of how different colloid compositions affect the interpretation of the same indicator.

Biological Significance of Active Acidity

Despite its quantitative smallness, active acidity is the habitat for plant roots and soil microorganisms. It is the concentration of H⁺ in the solution that determines:

• Nutrient Availability: In acidic environments (pH < 5.5), many cations (Fe, Mn, Zn, Cu) become more soluble and can reach toxic concentrations, while the availability of phosphorus, molybdenum, and calcium drops sharply (Foth, 1990; Weil, 2017). At high pH (alkaline), the opposite is observed—microelement deficiency.
• Aluminum Toxicity: At pH below 5.0–5.2, phytotoxic forms of aluminum (primarily Al³⁺ and AlOH²⁺) appear in the solution, damaging root systems (Sparks, 2003; Eash et al., 2016). This is one of the main reasons for plant depression on acidic soils.
• Microbial Activity: Most bacteria, including nitrogen fixers and nitrifiers, prefer a neutral or slightly alkaline environment; in acidic soils, their activity is suppressed, and fungi begin to play a dominant role (Weil, 2017).

However, it is crucial to understand: the H⁺ concentration itself is rarely a direct toxic factor (except for extremely low pH < 4.0). Most often, the harmful effect of an acidic reaction is mediated through aluminum, manganese, and nutritional disturbances. Active acidity is merely a "signal" indicator, but diagnosis begins with it.

Chapter Conclusion

Active acidity is the measured concentration of H⁺ in the soil solution. It is extremely small but has a decisive influence on the vital activity of roots and microorganisms. However, it does not provide a complete picture of the soil's acidity as a system. The main "reserve" of acidity is associated with the solid phase. That is why soils with the same pH can require vastly different amounts of lime for neutralization—their buffering capacity differs. We turn to understanding this buffering capacity and its constituent pools in the next chapter.

Definition and Essence of the Phenomenon

Potential (or reserve) acidity is the part of the total soil acidity associated with the solid phase and not directly manifested in the soil solution. It is "conserved" in the form of hydrogen and aluminum ions held on the surface of soil colloids (clay minerals and humus) in an exchangeable or non-exchangeable state (Foth, 1990; Weil, 2017).

It is this hidden form of acidity that determines the buffering capacity of the soil—its ability to resist changes in pH. Soils with a large reserve of potential acidity (high clay and humus content) are like a massive flywheel: moving their pH requires significant effort in the form of tons of lime. Conversely, sandy soils with a small reserve of potential acidity react to acid or alkali additions sharply and quickly.

Potential acidity is hundreds to thousands of times greater than active acidity. If active acidity can be neutralized by kilograms of lime per hectare, neutralizing potential acidity requires tons (Weil, 2017; Eash et al., 2016).

Composition of Potential Acidity: Two Main Components

Traditionally, potential acidity is divided into two main forms, differing in how tightly the acidic ions are bound to the colloids and by which reagent they can be displaced.

1. Exchangeable (or Salt-Displaceable) Acidity

This is the part of potential acidity represented by H⁺ and Al³⁺ cations held on the exchange sites of colloids by electrostatic forces (Eash et al., 2016; Foth, 1990). These ions are in dynamic equilibrium with the soil solution and can be relatively easily displaced (replaced) by other cations, for example, by treating the soil with a neutral salt extract (1 M KCl or 0.01 M CaCl₂).

Crucial Fact: In acidic mineral soils, exchangeable acidity is overwhelmingly due to aluminum (Al³⁺) rather than hydrogen (H⁺) (Sparks, 2003; White, 2006). This is one of the most important discoveries in soil chemistry, which radically changed the understanding of the nature of acidity. As Thomas (1977) wrote, there were heated debates for a long time about whether H⁺ or Al³⁺ was the main "culprit." It was only in the mid-20th century that it was definitively proven that in acidic soils, exchange sites are predominantly occupied by aluminum. We will discuss why this is so in detail when we talk about the sources of acidity.

Methodological nuance: Exchangeable acidity is determined in the laboratory by displacing ions from the soil with a KCl solution and then titrating the resulting extract with alkali. The obtained value is expressed in cmol(equiv)/kg and is called "exchangeable acidity" or "KCl-extractable acidity." In organic soils and raised bogs, the contribution of H⁺ to exchangeable acidity can be significant, but in mineral soils, Al³⁺ dominates (Weil, 2017).

2. Hydrolytic (or Residual) Acidity

This is a "deeper" and more tightly bound form of acidity. It is caused by hydrogen and aluminum ions that are part of the colloidal particles themselves. This includes:

• H⁺ covalently bonded to functional groups of humus (carboxyl -COOH, phenolic -OH) (Bloom & Skyllberg, 2012).
• Non-exchangeable forms of aluminum—hydroxy-aluminum polymers [Al(OH)₂⁺, Al(OH)³⁺, etc.] that may be located in the interlayer spaces of 2:1 clay minerals (e.g., in vermiculite) or strongly chemisorbed on the surface of oxides. These forms are not displaced by neutral salts (Sparks, 2003; White, 2006).

Hydrolytic acidity is revealed only by treating the soil with hydrolytically alkaline salts (e.g., sodium acetate CH₃COONa) or by direct titration of the suspension with alkali to a certain pH (usually 8.2). It is with this form that the concept of "acidity titratable to pH 8.2" is associated, which, together with exchangeable acidity, gives an idea of the total acidity of the soil.

Interrelation of Forms and Practical Significance

All three forms of acidity (active, exchangeable, and hydrolytic) are in dynamic equilibrium (see diagram in Weil, 2017). If we add lime (base) to the soil, it primarily neutralizes the active acidity in the solution. But as soon as the H⁺ concentration in the solution drops, the equilibrium shifts: some H⁺ and Al³⁺ ions move from exchange sites into the solution to restore the lost equilibrium. This continues until the entire exchangeable pool is exhausted. Only after this does the hydrolytic (residual) acidity come into play. It is this cascading process that provides buffering.

Why is it important to distinguish these pools?

1. For Diagnosis: Knowing only the pH (active acidity) is insufficient. To give a sound liming recommendation, an agrochemical laboratory must determine the amount of exchangeable and/or hydrolytic acidity.

2. For Understanding Toxicity: The toxic effect on plants in acidic soils is associated not with H⁺ itself, but with dissolved and exchangeable aluminum (Al³⁺). The higher the exchangeable acidity, the greater the potential risk of aluminum toxicity (Sparks, 2003). Therefore, in tropical and subtropical soils, the lime requirement is often calculated not based on the target pH, but on the required reduction of aluminum saturation of the SAC to a safe level (e.g., < 20–30% of CEC) (Foth, 1990; Weil, 2017).

3. For Prediction: High humus content and clay minerals like montmorillonite provide a huge reserve of potential acidity. Such soils are well-buffered. Conversely, soils dominated by kaolinite or iron and aluminum oxides, as well as sandy soils, have low buffering capacity, and their pH is easily "knocked off" even by small amounts of fertilizers.

Key takeaway of this chapter: Potential acidity is the "reservoir" of H⁺ and Al³⁺ ions associated with the solid phase of the soil. It is hundreds to thousands of times greater than active acidity, and it is precisely what determines the buffering properties of the soil and its actual liming requirement.

We have analyzed the composition of potential acidity. Now we are ready to move on to the next fundamental concept—buffering (Chapter 4), where we will synthesize knowledge about active and potential acidity and answer the main question of the lecture.

Definition and Essence of the Phenomenon

Buffering (or buffer capacity) is the soil's ability to resist changes in pH when acid or alkali is added (Eash et al., 2016; Weil, 2017; Bloom & Skyllberg, 2012). In other words, it is the "safety cushion" that prevents the soil solution from sharply acidifying or alkalizing with every rainfall, fertilizer application, or root respiration.

It is buffering that directly answers our key question. Two soils with the same pH (e.g., 5.0) can have completely different buffering capacities. In one (sandy, low in humus), the pH will drop sharply to 4.0 from a small dose of acid fertilizer. In the other (clayey, rich in organic matter), the same dose will cause only a barely noticeable change. Therefore, changing the pH of a well-buffered soil requires significantly more lime (or sulfur) than a weakly buffered one.

Mechanism of Buffering: Equilibrium Between Three Pools

We already know about the three pools of acidity: active (soil solution), exchangeable (H⁺ and Al³⁺ ions on the colloid surface), and hydrolytic (tightly bound forms in the colloid and humus structure). Buffering is ensured by the dynamic equilibrium between these pools.

A classic illustration often used in textbooks is a coffee machine (Weil, 2017). Active acidity is the coffee in the small transparent tube outside the machine. There is very little of it. Exchangeable and hydrolytic acidity is the huge coffee tank inside. If we pour out some coffee from the tube (neutralize the active acidity), the level in it will drop quickly. But the machine will immediately start pouring coffee from the tank into the tube until the levels equalize. Externally, the level in the tube has hardly changed, but we have spent part of the "reserve" from the tank. Similarly, in soil: when lime (alkali) is added, it first reacts with active H⁺ in the solution. As soon as their concentration drops, the transition of H⁺ and Al³⁺ from exchange sites into the solution begins, maintaining the pH at an almost constant level until the exchangeable pool is exhausted. When it too is depleted, hydrolytic acidity comes into play, and the process continues (Weil, 2017).

Quantitative Assessment of Buffering

The buffering capacity of a soil can be assessed in several ways:

1. By the magnitude of the cation exchange capacity (CEC). The higher the CEC, the more ions can be held on the colloid surface, and the larger the "reservoir" for exchange. High humus content and clay minerals with high CEC (e.g., montmorillonite) provide maximum buffering. Low CEC (sandy soils, kaolinite, oxides) means low buffering (Eash et al., 2016; Foth, 1990).

2. By titration curves. These are graphs showing how the soil pH changes upon the addition of certain amounts of acid or alkali. The classic study by Magdoff and Bartlett (1985) showed that for soils with high organic matter content, the titration curve has a gentler slope in the pH range 4.5–6.5, indicating high buffering. Conversely, soils with low organic matter content react to acid or base addition with a much sharper pH change. Moreover, the authors showed that when titration data are recalculated per unit of organic matter, the curves for different soils practically coincide, indicating the key role of organic matter in the buffering of surface horizons.

3. Through buffer pH. In agrochemical laboratories, special buffer solutions are used to calculate lime requirements (e.g., the Adams-Evans method or SMP buffer). By mixing soil with such a buffer having a fixed pH (e.g., 8.0 or 7.5) and measuring how much the soil "knocked down" its pH, the capacity of potential acidity and, consequently, the soil's buffering is judged (Essington, 2015; Weil, 2017).

Factors Determining Buffering

As we have understood, buffering depends on the composition and properties of the soil's solid phase:

• Organic matter. Possesses a huge number of functional groups (-COOH, -OH) capable of binding and releasing H⁺. It is humus that is the main source of buffering in the upper horizons of most soils.
• Composition of clay minerals. 2:1 minerals (smectites, vermiculites) with high CEC and ability to expand interlayers create greater buffering than 1:1 minerals (kaolinite) or iron and aluminum oxides (Foth, 1990; White, 2006).
• Carbonate content. In alkaline and neutral soils, the presence of CaCO₃ provides powerful buffering in the pH range 7–8.5 (Bloom & Skyllberg, 2012). As long as solid calcium carbonate is present in the soil, its pH cannot drop below a certain level, as it will dissolve, neutralizing incoming acid.

Why This Matters in Practice

Understanding buffering is critically important for decision-making in agriculture and ecology:

• Liming. In well-buffered soils, significantly more lime is required to achieve the target pH. Ignoring this fact leads to under-liming and reduced yields. In weakly buffered soils, on the other hand, it is easy to overdo lime (over-liming), causing microelement deficiencies.
• Acid Rain. It is the buffering capacity of the soil that determines how severely an ecosystem will suffer from acid precipitation. Poor, sandy soils with low buffering (e.g., Podzols on granite) are the most sensitive and acidify quickly (Weil, 2017; Bloom & Grigal, 1985).
• Fertilizers. Systematic application of physiologically acid fertilizers (ammonium nitrate, ammonium sulfate) leads to gradual acidification. In well-buffered soils, this process is slow, and pH can remain stable for decades, while in weakly buffered soils, it drops rapidly, requiring frequent liming.

Key takeaway of this chapter: Buffering is an integral property of the soil, determined by the equilibrium between active and potential forms of acidity. It depends on the composition of the SAC (humus, clay type) and the presence of carbonates. It is buffering that explains why soils react differently to external influences, and it is precisely this that must be taken into account when calculating amendment rates.

We have covered what buffering is and what it depends on. Now it is logical to move on to the question: where does this acidity come from in the first place? What natural and anthropogenic processes generate H⁺ in the soil? Let's move on to the next chapter, 5. Sources of Acidity.

1. Carbonic Acid (H₂CO₃)

This is the most universal and constantly acting source of acidity. Plant roots and soil microorganisms continuously respire, releasing carbon dioxide (CO₂) into the soil air. Its concentration in the soil can be tens to hundreds of times higher than atmospheric (0.03–0.1 atm and above) (Weil, 2017; Huang, 2012). Dissolving in soil moisture, CO₂ forms weak carbonic acid:

This process is the main driver of chemical weathering in nature. It provides a constant, albeit weak, influx of H⁺ into the soil solution. However, carbonic acid is a weak electrolyte (its pK_a = 6.35), so it is effective as a source of acidity only at pH > 5.0. In more acidic environments, its share in the total H⁺ balance decreases.

2. Organic Acids

During the decomposition of plant residues and humus by microorganisms, a wide range of organic acids is formed: from low molecular weight (oxalic, citric, malic) to high molecular weight humic acids (fulvic acids) with carboxyl (-COOH) and phenolic (-OH) groups (Bloom & Skyllberg, 2012; Weil, 2017). These acids are significantly stronger than carbonic acid (their pK_a lies in the range 3–5), and they play an important role in acidifying the upper soil horizons, especially under forest litter and in organic horizons.

The general process can be represented as the oxidation of organic matter:

It is important to note that the process of organic matter accumulation itself also contributes to acidification, as organic acids form soluble complexes with base cations (Ca²⁺, Mg²⁺, K⁺), facilitating their removal from the soil (Weil, 2017).

3. Nitrification (Oxidation of Nitrogen)

This is one of the most powerful anthropogenically enhanced sources of acidity in arable soils. Ammonium nitrogen (NH₄⁺), coming from mineral fertilizers or formed during the mineralization of organic matter, undergoes biological oxidation to nitrates (NO₃^-) in a two-stage nitrification process:

Overall:

For every mole of oxidized ammonium, two moles of H⁺ are formed (Eash et al., 2016; Weil, 2017). If the resulting nitrates are leached from the soil (rather than taken up by plants), this acidity remains in the system, causing its acidification. This is why the systematic application of ammonium fertilizers without liming leads to progressive soil acidification.

4. Oxidation of Sulfur

Sulfur can be present in the soil in various forms: organic (in proteins), elemental (S⁰), and mineral (in sulfides, e.g., pyrite FeS₂). When these reduced forms are oxidized by microorganisms (e.g., Thiobacillus spp.), sulfuric acid (H₂SO₄) is formed:

This process underlies the formation of acid sulfate soils (e.g., during drainage of coastal mangrove swamps or coal mining). The pH of such soils can drop to values of 2–3, which is detrimental to plants and causes the formation of so-called "acid mine drainage" (Weil, 2017; Essington, 2015).

5. Acid Rain

Industrial emissions of SO₂ and NOx into the atmosphere lead to the formation of sulfuric and nitric acids, which return to the surface with rain and snow. The pH of such precipitation can drop to 3.0–4.0. However, in most agricultural soils, the contribution of acid rain to total acidification is significantly less than that of fertilizer nitrification, except in weakly buffered (carbonate-poor) forest ecosystems (Weil, 2017; Sparks, 2003).

6. Plant Uptake of Cations

Plant roots, absorbing cations from the soil solution (K⁺, Ca²⁺, Mg²⁺, NH₄⁺), must maintain electroneutrality. If cation uptake exceeds anion uptake, the roots release excess H⁺ into the rhizosphere (Weil, 2017; Eash et al., 2016). This local process can significantly lower the pH in the immediate vicinity of the roots, especially under ammonium nutrition.

The Special Role of Aluminum as an Acidity Generator

We have already mentioned that exchangeable acidity in mineral soils is almost entirely due to aluminum. Where does it come from? This process is inextricably linked to the mineral composition of the soil.

When additional H⁺ enter the soil solution (from any of the sources listed above), they attack the crystal lattice of clay minerals and other aluminosilicates. An exchange occurs: H⁺ displaces Al³⁺ from the mineral structure into the solution, where it becomes hydrated and becomes an exchangeable cation (Foth, 1990; White, 2006; Sparks, 2003):

However, the process does not end there. Once in the solution, the hydrated aluminum ion Al(H₂O)₆³⁺ begins to hydrolyze—reacting with water, splitting off OH⁻ ions and releasing additional H⁺ into the solution (Sparks, 2003; White, 2006; Essington, 2015):

Thus, each Al³⁺ ion that enters the solution can generate up to three additional H⁺! This is the key feature of aluminum as a source of "self-sustaining" acidity. The more acidic the environment, the more actively aluminosilicates dissolve, and the more aluminum enters the solution, which, in turn, further acidifies the environment. This positive feedback loop is the main reason why acidic soils tend to become even more acidic.

Natural Counterbalance: Neutralization Processes

It is important to remember that processes consuming H⁺ and counteracting acidification exist in nature (Foth, 1990; Weil, 2017):

• Weathering of primary minerals: Consumes H⁺ with the formation of secondary clay minerals and the release of base cations (Ca²⁺, Mg²⁺, K⁺, Na⁺).
• Denitrification: Reduction of nitrates to gaseous nitrogen (N₂, N₂O) under anaerobic conditions consumes H⁺.
• Plant uptake of anions: If plants absorb more anions (NO₃^-, SO₄^2-) than cations, they release OH⁻ or HCO₃^- into the rhizosphere, raising the pH.
• Reduction of iron and manganese: In waterlogged soils, the reduction of Fe³⁺ to Fe²⁺ and Mn⁴⁺ to Mn²⁺ consumes H⁺.

The balance between these processes determines whether the soil will be acidic, neutral, or alkaline. In humid regions, where base leaching dominates, acid-forming processes prevail, and soils evolve towards acidification.

Key takeaway of this chapter: Soil acidity has many sources: from natural (carbonic acid from respiration, organic acids from decomposition) to anthropogenically enhanced (fertilizer nitrification, acid rain). Aluminum plays a key role in the development and maintenance of acidity; released from minerals under the action of H⁺, it itself becomes a powerful generator of new H⁺ through hydrolysis.

We have covered where acidity comes from. Now it is logical to move on to its "antipode"—alkalinity. Which processes lead to the accumulation of bases in the soil? Let's move to Chapter 6. Alkalinity.

Excellent, we have come to the other side of the coin—alkalinity. Understanding its nature is no less important than knowing about acidity, especially in the context of arid regions and specific soil conditions.

Definition and Boundaries

Alkalinity is a property of the soil solution caused by the presence of hydroxyl ions (OH⁻) or anions of weak acids that can hydrolyze to form OH⁻. Quantitatively, alkalinity manifests at pH > 7.0 (Foth, 1990; Weil, 2017). Recall that pH is a logarithmic measure of H⁺ activity. At pH > 7, the H⁺ concentration becomes less than the OH⁻ concentration. However, as with acidity, it is important to distinguish between active alkalinity (solution pH) and reserve (buffer) alkalinity associated with the solid phase of the soil.

Pure water in equilibrium with atmospheric CO₂ has a pH of about 5.6 due to the formation of carbonic acid. Therefore, pH > 7 in soil always indicates the presence of buffer systems capable of binding H⁺ and accumulating bases. The main such system in most alkaline soils is carbonates and bicarbonates (Essington, 2015; Bloom & Skyllberg, 2012).

The Main Source: Carbonates and Hydrolysis

The main cause of alkalinity in most soils is the presence of calcium and magnesium carbonates (CaCO₃, MgCO₃) or their mixtures. These minerals are widespread in arid and semi-arid areas where lack of moisture prevents their leaching from the soil profile. Additionally, carbonates can occur in humid zones on carbonate rocks (limestones, chalks, marls) or young alluvial deposits.

The hydrolysis of calcium carbonate in aqueous solution proceeds as follows (Foth, 1990; Weil, 2017):

The resulting bicarbonate ion (HCO₃^-) and hydroxide ion (OH⁻) cause the alkaline reaction. The equilibrium of this reaction is strongly dependent on the partial pressure of CO₂ in the soil air. In the presence of CO₂, the reaction shifts to the right, enhancing carbonate dissolution. However, when the pH rises above 8.3, bicarbonates begin to convert to carbonates (CO₃²⁻), which reduces the solubility of CaCO₃ and limits further pH increase (Essington, 2015). Therefore, in soils saturated with calcium carbonates, the pH rarely exceeds 8.3–8.5—this is the so-called "carbonate buffer."

Soda Alkalinity and the Role of Sodium

If sodium, rather than calcium, is present in the soil as sodium carbonate (Na₂CO₃) or bicarbonate (NaHCO₃), the pH can rise much higher—up to 9.5 and even 10.5. This is because:

• Na₂CO₃ is more soluble than CaCO₃;
• during hydrolysis of Na₂CO₃, a strong base NaOH is formed:

that is, an excess of OH⁻ appears in the solution, not rigidly bound by calcium equilibrium.

Such soils are called soda-saline or soda-type solonchaks (Weil, 2017). They are characteristic of closed, drainage-less depressions in arid areas where salts accumulate. Sodium, adsorbed on the colloid surface, greatly alters their physicochemical properties, which we will discuss shortly.

The Role of Colloids: Exchangeable Sodium and Dispersion

Unlike acidity, where aluminum plays a key role, in alkalinity, exchangeable sodium (Na⁺) held on the negatively charged surfaces of clay and humus occupies a special place. When the proportion of sodium in the SAC exceeds 15% of the cation exchange capacity, the soil acquires special properties and is called solonetzic (or sodium-affected). This manifests as follows:

• Colloid dispersion. Sodium, having a large hydrated radius and low adsorption energy, weakly draws negatively charged particles together. As a result, colloids (clay particles) repel each other, forming a suspension rather than aggregates. This leads to soil slumping, crust formation, and a sharp decrease in water permeability (Eash et al., 2016; Foth, 1990). The soil becomes practically impermeable to water, which is vividly illustrated by the figure in the textbook by Eash et al. (2016), comparing soil aggregation when Ca²⁺ dominates (good structure) versus Na⁺ (dispersion).
• High pH. Sodium saturation of the SAC leads to an increase in pH due to the hydrolysis of exchangeable Na⁺, which replaces H⁺ in the solution, promoting the accumulation of OH⁻ and the formation of an alkaline environment (Foth, 1990).

Buffer Systems in Alkaline Soils

As in acidic soils, buffering exists in alkaline soils. The main buffer pairs are:

1. Carbonate-bicarbonate: CO₃^2- / HCO₃^- — manifests in the pH range 8.3–10.2.

2. Calcium carbonate: CaCO₃ / Ca²⁺ + CO₃^2- — limits pH from above (_{8.3–8.5) in the presence of solid CaCO}3~.

3. Silicate (to a lesser extent): due to silicate hydrolysis.

Thanks to these systems, alkaline soils, like acidic ones, resist pH changes, but the mechanism is different—not cation exchange, but carbonate dissolution/precipitation. This is especially important when applying acid fertilizers: as long as there is a reserve of CaCO₃, the pH will remain around 8.0–8.3, neutralizing added H⁺ (Bloom & Skyllberg, 2012).

Biological Significance of Alkalinity

For plants and microorganisms, an alkaline environment creates its own set of problems, contrasting with acid stress (Foth, 1990; Weil, 2017):

• Microelement deficiency. At high pH (especially > 7.5), the solubility and availability of iron (Fe), manganese (Mn), zinc (Zn), copper (Cu), and boron (B) drop sharply. This leads to chlorosis (leaf yellowing) and reduced yields.
• Reduced phosphorus availability. In an alkaline environment, phosphorus binds with calcium to form sparingly soluble calcium phosphates (apatites, etc.), limiting plant phosphorus nutrition.
• Increased molybdenum availability. Unlike most elements, molybdenum (Mo) becomes more available at high pH, which in rare cases can lead to toxicity.
• Reduced root activity. High pH can directly damage root hairs, especially in plants not adapted to alkaline conditions.

Difference from Acidity: The Nature of "Hidden" Alkalinity

In acidic soils, the hidden reserve is exchangeable H⁺ and Al³⁺. In alkaline soils, the analogue is primarily carbonate buffering and exchangeable sodium. However, if exchangeable sodium does not manifest itself as a direct source of OH⁻ (it only promotes hydrolysis), then carbonates are a direct reserve for acid neutralization. Therefore, when determining alkalinity, one speaks not of "potential alkalinity" but of "reserve alkalinity," measured, for example, by carbonate content or buffer capacity towards acid. In practice, this is often defined as "carbonate alkalinity" or "total alkalinity" (Essington, 2015).

Key takeaway of this chapter: Soil alkalinity is primarily due to the presence of carbonates (Ca, Mg) and, in more extreme cases, sodium carbonates. It determines the availability of micronutrients and soil structure through the influence of exchangeable sodium on colloid dispersion. Buffering in an alkaline environment is related to the solubility of carbonates and is balanced by the partial pressure of CO₂.

We have examined both sides of the soil reaction—acidity and alkalinity. Now it is time to directly link these phenomena to soil colloids, as was stated as the main goal of the lecture. Let's move on to Chapter 7. The Influence of Colloids on pH. A Very Modern Connection.

Two Types of Charge: Permanent and Variable

To understand the influence of colloids on pH, we need to grasp two key concepts introduced in previous lectures (White, 2006; Eash et al., 2016):

1. Permanent (permanent) charge. Arises in the crystal lattice of 2:1 clay minerals (smectites, vermiculite, illite) as a result of isomorphous substitution—the replacement of one cation in the structure by another with a lower charge. For example, replacing Si⁴⁺ with Al³⁺ in the tetrahedral sheet or Al³⁺ with Mg²⁺ in the octahedral sheet creates a deficit of positive charges, which manifests as a permanent negative charge on the surface. This charge is independent of pH. It exists always and determines the base cation exchange capacity (CEC) of such minerals.

2. Variable (pH-dependent) charge. Arises on surfaces where there are functional groups capable of attaching or releasing protons (H⁺). Such surfaces include:

• Humus (organic matter). The carboxyl (-COOH) and phenolic (-OH) groups of humic acids can dissociate by losing H⁺, creating a negative charge (R-COO^-). The higher the pH, the more groups dissociate, and the higher the negative charge (Bloom & Skyllberg, 2012; Huang, 2012).
• Iron and aluminum oxides and hydroxides (Fe₂O₃, Al₂O₃, FeOOH, Al(OH)₃). Hydroxyl groups (=Al-OH, =Fe-OH) are present on their surfaces, which can behave as amphoteric compounds. At low pH, they protonate (attach H⁺) and acquire a positive charge (=Al-OH₂⁺). At high pH, they deprotonate (lose H⁺) and acquire a negative charge (=Al-O⁻).
• Edge surfaces (edges) of layered silicates. In 1:1 minerals (kaolinite) and on the edges of 2:1 minerals, hydroxyl groups are also present, exhibiting pH-dependent charge (White, 2006; Sparks, 2003).

Why Does This Matter for pH?

The ratio of permanent to variable charge in the soil determines its point of zero charge (PZC). This is the pH at which the total surface charge of the colloids is zero (White, 2006; Bloom & Skyllberg, 2012).

• If the soil pH is above the PZC, the colloid surface carries a net negative charge and acts as a cation exchanger.
• If the soil pH is below the PZC, the surface carries a net positive charge and acts as an anion exchanger.

This is precisely what explains why soils with different colloidal compositions behave so differently!

Let's consider three main types of soil colloids from the perspective of their influence on pH.

1. 2:1 Clay Minerals (Smectites, Vermiculite, Illite)

These minerals have a high permanent negative charge (80–150 cmol(equiv)/kg). The variable charge on the edges constitutes only a small fraction of the total. Therefore:

• Their CEC is weakly dependent on pH. The main charge is always present.
• They are powerful cation exchangers. This ensures high buffering capacity, as they can hold a large number of exchangeable cations (including H⁺ and Al³⁺) (Foth, 1990).
• In an acidic environment, they become "aluminum" ones. Since they strongly retain Al³⁺, this cation becomes the main component of exchangeable acidity (Sparks, 2003; White, 2006).

2. Humus (Organic Matter)

Humus carries exclusively pH-dependent charge created by carboxyl and phenolic groups (Huang, 2012; Bloom & Skyllberg, 2012). Its CEC can reach 200–400 cmol(equiv)/kg at high pH but drops sharply as pH decreases.

• It is the main "buffer" in the upper horizons. As shown by Magdoff and Bartlett (1985), it is organic matter that determines the buffering of many surface soils in the pH range 4.5–6.5.
• The key role of aluminum. Under acidic conditions (pH < 5), humus actively binds Al³⁺ into strong complexes that are not displaced by neutral salts (Bloom et al., 2005; Hargrove & Thomas, 1982). This "organic" aluminum does not manifest as exchangeable acidity, but it is very important for buffering, because when the pH rises, it hydrolyzes, releasing H⁺.

A Modern View of Aluminum on Organic Matter

This is a very interesting and important point. When Al³⁺ binds to humus, it actually "neutralizes" part of the negative charge of the organic matter. As a result:

• The CEC of humus decreases. Sites occupied by aluminum cease to be exchangeable.
• Acidity changes. Aluminum bound to humus is a weaker acid (i.e., it releases H⁺ at a higher pH) than H⁺ bound to the same groups. Therefore, in soils rich in organic matter, at the same pH value, there can be significantly less exchangeable acidity than in mineral soils (Weil, 2017). This is one reason why peat soils are less toxic to plants than mineral soils with the same pH—aluminum there is firmly "locked" in organic complexes.

3. Iron and Aluminum Oxides (and Kaolinite)

These minerals are dominated by variable charge. Their PZC lies in the pH range 7–9 (for oxides) and around pH 2–3 for kaolinite edges (White, 2006; Bloom & Skyllberg, 2012).

• In acidic soils, they are anion exchangers. At pH below the PZC, their surface is positively charged. This means they will retain anions (e.g., H₂PO₄⁻, SO₄^2-) and repel cations, including H⁺ and Al³⁺ (Sparks, 2003; Weil, 2017).
• They do not create exchangeable acidity. Since they do not retain Al³⁺ in exchangeable form, their contribution to potential acidity is minimal.
• Their CEC is strongly pH-dependent. As pH increases, their negative charge increases, leading to a sharp increase in CEC. This is one reason why highly weathered tropical soils (Oxisols) have low CEC in their natural acidic state, but upon liming, their CEC can increase substantially.

Practical Conclusion: Two Types of Soil pH

From all that has been said, a crucial conclusion for practice emerges: soils dominated by permanently charged colloids (2:1 clays) and soils dominated by variably charged colloids (organic matter, oxides, kaolinite) behave differently.

• In the first case (2:1 clays), the main problem is high exchangeable acidity associated with Al³⁺. Lime is required for its neutralization, and the effect of liming will be long-lasting due to high CEC.
• In the second case (organic and highly weathered soils), the main factor is pH-dependent charge. Liming not only neutralizes acidity but also sharply increases CEC, improving cation retention. However, in such soils, the risk of "over-liming" is higher, and the target pH is often set lower (around 5.5) to avoid microelement deficiencies and sharp changes in colloid chemistry (Foth, 1990; Weil, 2017).

Effective Cation Exchange Capacity (ECEC)

It is precisely to account for the influence of pH on colloids that the concept of effective cation exchange capacity (ECEC) was introduced. ECEC is the sum of exchangeable cations (Ca²⁺, Mg²⁺, K⁺, Na⁺, Al³⁺, H⁺) determined at the natural pH of the soil (White, 2006). Unlike CEC measured at pH 7 (potential CEC), ECEC reflects the real exchange capacity under given conditions. In acidic soils with variable charge, ECEC is significantly lower than CEC at pH 7. This is a critically important indicator for diagnosing the fertility of acidic soils.

Key takeaway of this chapter:

Soil pH is not just the result of "acidic" or "alkaline" processes. It is a direct consequence of the composition and properties of soil colloids. The type of colloids (2:1 clays, humus, oxides) determines which charge (permanent or variable) dominates, how the soil will retain Al³⁺ and other cations, and, ultimately, how buffered it will be. Understanding this is the key to proper soil fertility management.

We have covered how exactly colloids influence pH. Now we are ready to move on to the logical conclusion—the systematization of all knowledge and the description of specific soil buffer systems. In the next chapter, we will bring everything together.

1. Buffer System "Aluminosilicate Dissolution" (Extremely Acidic Soils, pH < 3.5–4.0)

This is the "deepest" and, in a sense, emergency system. It is triggered when the pH drops below 3.5–4.0, which is typical for acid sulfate soils (pyrite dumps, drained mangroves), as well as soils subjected to prolonged exposure to strong acids.

Mechanism: At such a high concentration of H⁺ in the solution, intensive destruction (hydrolysis) of the crystal lattice of clay minerals begins, primarily 2:1 layered silicates (Bloom & Skyllberg, 2012; Huang, 2012). Protons attack the alumino-oxygen and silicate bonds, releasing Al³⁺, Fe³⁺, and silicic acid into the solution. For example, for montmorillonite:

This process consumes a huge amount of H⁺ (up to 6 moles per mole of structural cations), so the pH remains at an extremely low level, but its further fall is sharply slowed down. However, this buffer system has serious consequences:

• Destruction of the SAC. Clay minerals lose their structure, and soil CEC decreases.
• Release of toxic ions. Al³⁺ and Fe³⁺ in such concentrations are harmful to plants.

This system is a signal of soil degradation, and the restoration of such soils requires drastic measures.

2. Aluminum Buffer System (Acidic Soils, pH 4.0–5.5)

This is the main buffer system for most acidic mineral soils (Podzols, Brown Forest soils, Red Earths, many soils of the taiga and humid tropics). In this pH range, aluminum, present in several forms, is the main buffer.

Main mechanisms:

• Hydrolysis of exchangeable Al³⁺. As we already know (Chapter 5), Al³⁺ ions held on the exchange sites of clay minerals and humus hydrolyze, releasing H⁺ (White, 2006; Sparks, 2003; Weil, 2017):

Each of these steps shifts right or left depending on pH, creating a powerful buffering effect. When alkali (lime) is added, these reactions shift to the right, consuming OH⁻, and the pH remains almost unchanged for a long time until all the "exchangeable" aluminum is used up.

• Formation and dissolution of hydroxy-aluminum polymers. In the same pH range, polymeric forms of aluminum such as [Al₁₃O₄(OH)₂₄(H₂O)₁₂]⁷⁺ (the so-called Al₁₃ polymer) and others can exist (Sparks, 2003; Bertsch & Parker, 1996). These large polycations can become fixed in the interlayer spaces of 2:1 clay minerals (forming so-called hydroxy-interlayered layers) as well as on oxide surfaces. They also participate in buffer reactions, donating or accepting H⁺ as pH changes.

• Aluminum on organic matter. As we discussed in Chapter 7, in acidic soils, a significant portion of aluminum is bound to humus in strong complexes (Bloom et al., 2005). This aluminum is not displaced by salts, but it participates in buffer reactions, especially in the pH range 5–8, when upon alkalization it hydrolyzes and precipitates as Al(OH)₃ (Hargrove & Thomas, 1982; Weil, 2017).

It is the aluminum buffer system that is the main reason why the pH of acidic soils with high clay and organic matter content is difficult to shift to the neutral side. It is on it that all modern methods for calculating lime rates based on exchangeable acidity and base saturation are founded.

3. Exchangeable (Cation Exchange) Buffer System (Slightly Acidic and Neutral Soils, pH 5.5–7.0)

This is the main buffer system for most cultivated and natural soils dominated by 2:1 clay minerals (Chernozems, Gray Forest soils, Soddy-Podzolic soils with high CEC). In this pH range, exchangeable aluminum is practically absent (it is precipitated as Al(OH)₃), and the main role is played by cation exchange between the solution and the colloid surface.

Mechanism: The colloid surface (clay + humus) carries a negative charge and holds exchangeable cations: Ca²⁺, Mg²⁺, K⁺, Na⁺, as well as some H⁺. When acid (H⁺) is added to the solution, it displaces these cations from exchange sites, binding to the colloid. When alkali (OH⁻) is added, the reverse process occurs: OH⁻ neutralizes H⁺ in the solution, the equilibrium shifts, and some H⁺ from the colloids move into the solution, while their places are taken by base cations (e.g., Ca²⁺ from lime).

This process is described by classical cation exchange equations (Gapon's, Nikolsky's, etc.), which we considered in the lecture on the SAC. The buffer capacity of this system is directly proportional to the soil CEC: the higher the CEC, the more ions can be exchanged, and the higher the buffering.

Features:

• This system is very sensitive to the composition of exchangeable cations. If divalent cations (Ca²⁺, Mg²⁺) predominate, buffering is higher, as they are more strongly retained on the surface. If there are many monovalent ones (Na⁺, K⁺), buffering decreases.
• This system operates quickly and reversibly, ensuring short-term pH stability.

4. Carbonate Buffer System (Alkaline and Neutral Soils, pH 7.0–8.5)

This is the main buffer system for soils in arid and semi-arid regions, as well as for carbonate soils (Rendzinas, Chernozems on carbonate rocks, Sierozems). It is determined by the presence of solid calcium and magnesium carbonates (CaCO₃, MgCO₃) (Bloom & Skyllberg, 2012; Essington, 2015; Weil, 2017).

Mechanism: Equilibrium between solid carbonate, dissolved ions, and CO₂ in the soil air:

When acid (H⁺) is added, the equilibrium shifts to the right: carbonate dissolves, binding H⁺ and forming bicarbonates, thus neutralizing the acid. When alkali (OH⁻) is added, the equilibrium shifts to the left: carbonate precipitates, binding excess OH⁻.

Features:

• This is a powerful buffer system. Its capacity is determined by the amount of carbonates in the soil, which can reach tens of percent by mass.
• It limits pH from above. As long as solid CaCO₃ is present in the soil, the pH cannot rise above approximately 8.3–8.5 (under normal partial pressure of CO₂). Further pH increase is possible only in the presence of more soluble carbonates (e.g., Na₂CO₃).
• This system determines alkaline buffering and is the reason why carbonate salinization suppresses root activity and microelement availability.

5. Silicate and Organic Buffer System (High pH > 8.5)

These are additional buffer systems operating under very alkaline conditions (soda solonchaks, pH > 9). They are associated with:

• Hydrolysis of silicates. In a strongly alkaline environment, hydrolysis of silicate minerals occurs with the release of silicic acid and cations, which consumes OH⁻.
• Dissociation of phenolic groups of humus. At pH > 8.5, the phenolic groups of organic matter (pK_a ~ 9–10) begin to actively dissociate, creating additional negative charge and buffering.

However, these systems are of lesser practical importance for most agricultural soils.

Generalizing Scheme: Hierarchy of Buffer Systems

Thus, we can represent the hierarchy of soil buffer systems depending on pH:

Practical Significance of the Hierarchy

1. Diagnosis. Knowing the soil pH, we can understand which buffer system dominates in it. This determines the choice of analysis methods (e.g., for acidic soils, exchangeable acidity needs to be determined, not just pH) and reclamation methods.

2. Prediction. Understanding the hierarchy allows predicting how the soil will respond to long-term impacts (e.g., acid rain or systematic fertilizer application). Will the pH decrease slowly (thanks to the aluminum buffer system) or drop sharply when it is exhausted?

3. Calculation of amendment rates. For liming, we must neutralize not only the active acidity but also the entire reserve of exchangeable acidity. Therefore, the lime rate is calculated based on the amount of exchangeable acidity (or hydrolytic acidity for complete neutralization), not on pH. Conversely, to acidify alkaline soils, the carbonate buffering must be overcome, which requires significant amounts of sulfur.

Key takeaway of this chapter:

Soil buffering is not a uniform property but a cascade of parallel and sequential systems, each dominating in a specific pH range. Understanding this hierarchy (from aluminum to carbonate) is the key to proper soil fertility management and preventing soil degradation.