Soil pH

Soil pH is a measure of the

Soil pH is considered a master variable in soils as it affects many chemical processes. It specifically affects plant nutrient availability by controlling the chemical forms of the different nutrients and influencing the chemical reactions they undergo. The optimum pH range for most plants is between 5.5 and 7.5;^[3] however, many plants have adapted to thrive at pH values outside this range.

Classification of soil pH ranges

The United States Department of Agriculture Natural Resources Conservation Service classifies soil pH ranges as follows:^[4]

0 to 6=acidic,7=neutral and 8 and above alkalinity

Determining pH

Methods of determining pH include:

• Observation of soil profile: certain profile characteristics can be indicators of either acid, saline, or sodic conditions. Examples are:^[5]
  • Poor incorporation of the organic surface layer with the underlying mineral layer – this can indicate strongly acidic soils;
  • The classic podzol horizon sequence, since podzols are strongly acidic: in these soils, a pale eluvial (E) horizon lies under the organic surface layer and overlies a dark B horizon;
  • Presence of a caliche layer indicates the presence of calcium carbonates, which are present in alkaline conditions;
  • Columnar structure can be an indicator of sodic condition.
• Observation of predominant flora.
  Cornus spp.), lilac (Syringa) and Clematis
  species.
• Use of an inexpensive pH testing kit, where in a small sample of soil is mixed with indicator solution which changes colour according to the acidity.
• Use of
  litmus paper
  . A small sample of soil is mixed with distilled water, into which a strip of litmus paper is inserted. If the soil is acidic the paper turns red, if basic, blue.
• Certain other fruit and vegetable pigments also change color in response to changing pH. Blueberry juice turns more reddish if acid is added, and becomes indigo if titrated with sufficient base to yield a high pH. Red cabbage is similarly affected.
• Use of a commercially available electronic pH meter, in which a glass or solid-state electrode is inserted into moistened soil or a mixture (suspension) of soil and water; the pH is usually read on a digital display screen.^[6]
• In the 2010s, spectrophotometric methods were developed to measure soil pH involving addition of an indicator dye to the soil extract.^[7] These compare well to glass electrode measurements but offer substantial advantages such as lack of drift, liquid junction and suspension effects.

Precise, repeatable measures of soil pH are required for scientific research and monitoring. This generally entails laboratory analysis using a standard protocol; an example of such a protocol is that in the USDA Soil Survey Field and Laboratory Methods Manual.^[8] In this document the three-page protocol for soil pH measurement includes the following sections: Application; Summary of Method; Interferences; Safety; Equipment; Reagents; and Procedure.

Summary of Method

The pH is measured in soil-water (1:1) and soil-salt (1:2 ${\displaystyle {\ce {CaCl2}}}$) solutions. For convenience, the pH is initially measured in water and then measured in ${\displaystyle {\ce {CaCl2}}}$. With the addition of an equal volume of 0.02 M ${\displaystyle {\ce {CaCl2}}}$ to the soil suspension that was prepared for the water pH, the final soil-solution ratio is 1:2 0.01 M ${\displaystyle {\ce {CaCl2}}}$.
A 20-g soil sample is mixed with 20 mL of reverse osmosis (RO) water (1:1 w:v) with occasional stirring. The sample is allowed to stand 1 h with occasional stirring. The sample is stirred for 30 s, and the 1:1 water pH is measured. The 0.02 M ${\displaystyle {\ce {CaCl2}}}$ (20 mL) is added to soil suspension, the sample is stirred, and the 1:2 0.01 M ${\displaystyle {\ce {CaCl2}}}$ pH is measured (4C1a2a2).

— Summary of the USDA NRCS method for soil pH determination^[8]

Factors affecting soil pH

The pH of a natural soil depends on the mineral composition of the parent material of the soil, and the weathering reactions undergone by that parent material. In warm, humid environments, soil acidification occurs over time as the products of weathering are leached by water moving laterally or downwards through the soil. In dry climates, however, soil weathering and leaching are less intense and soil pH is often neutral or alkaline.^[9][10]

Sources of acidity

Many processes contribute to soil acidification. These include:^[11]

• Rainfall: Average rainfall has a pH of 5.6 and is moderately acidic due to dissolved atmospheric carbon dioxide (CO
  ₂) that combines with water to form carbonic acid (H
  ₂CO
  ₃). When this water flows through the soil it results in the leaching of basic cations as bicarbonates; this increases the percentage of Al³⁺
  and H⁺
  relative to other cations.^[12]
• Root respiration and decomposition of organic matter by microorganisms release CO
  ₂ which increases the carbonic acid (H
  ₂CO
  ₃) concentration and subsequent leaching.
• Plant growth: Plants take up nutrients in the form of ions (e.g. NO⁻
  ₃, NH⁺
  ₄, Ca²⁺
  , H
  ₂PO⁻
  ₄), and they often take up more
  anions
  . However, plants must maintain a neutral charge in their roots. In order to compensate for the extra positive charge, they will release H⁺
  ions from the root. Some plants also exude organic acids into the soil to acidify the zone around their roots to help solubilize metal nutrients that are insoluble at neutral pH, such as iron (Fe).
• Fertilizer use: Ammonium (NH⁺
  ₄) fertilizers react in the soil by the process of nitrification to form nitrate (NO⁻
  ₃), and in the process release H⁺
  ions.
• Acid rain: The burning of fossil fuels releases oxides of sulfur and nitrogen into the atmosphere. These react with water in the atmosphere to form sulfuric and nitric acid in rain.
• Oxidative weathering: Oxidation of some primary minerals, especially sulfides and those containing Fe²⁺
  , generate acidity. This process is often accelerated by human activity:
  • Mine spoil: Severely acidic conditions can form in soils near some mine spoils due to the oxidation of pyrite.
  • Acid sulfate soils formed naturally in waterlogged coastal and estuarine environments can become highly acidic when drained or excavated.

Sources of alkalinity

Total soil alkalinity increases with:^[13][14]

• Weathering of silicate, aluminosilicate and carbonate minerals containing Na⁺
  , Ca²⁺
  , Mg²⁺
  and K⁺
  ;
• Addition of silicate, aluminosilicate and carbonate minerals to soils; this may happen by deposition of material eroded elsewhere by wind or water, or by mixing of the soil with less weathered material (such as the addition of limestone to acid soils);
• Addition of water containing dissolved bicarbonates (as occurs when irrigating with high-bicarbonate waters).

The accumulation of alkalinity in a soil (as carbonates and bicarbonates of Na, K, Ca and Mg) occurs when there is insufficient water flowing through the soils to leach soluble salts. This may be due to arid conditions, or poor internal soil drainage; in these situations most of the water that enters the soil is transpired (taken up by plants) or evaporates, rather than flowing through the soil.^[13]

The soil pH usually increases when the total alkalinity increases, but the balance of the added cations also has a marked effect on the soil pH. For example, increasing the amount of sodium in an alkaline soil tends to induce dissolution of calcium carbonate, which increases the pH. Calcareous soils may vary in pH from 7.0 to 9.5, depending on the degree to which Ca²⁺
or Na⁺
dominate the soluble cations.^[13]

Effect of soil pH on plant growth

Acid soils

High levels of aluminium occur near mining sites; small amounts of aluminium are released to the environment at the coal-fired power plants or incinerators.^[15] Aluminium in the air is washed out by the rain or normally settles down but small particles of aluminium remain in the air for a long time.^[15]

Acidic precipitation is the main natural factor to mobilize aluminium from natural sources^[16] and the main reason for the environmental effects of aluminium;^[17] however, the main factor of presence of aluminium in salt and freshwater are the industrial processes that also release aluminium into air.^[16] Plants grown in acid soils can experience a variety of stresses including aluminium (Al), hydrogen (H), and/or manganese (Mn) toxicity, as well as nutrient deficiencies of calcium (Ca) and magnesium (Mg).^[18]

Aluminium toxicity is the most widespread problem in acid soils. Aluminium is present in all soils to varying degrees, but dissolved Al³⁺ is toxic to plants; Al³⁺ is most soluble at low pH; above pH 5.0, there is little Al in soluble form in most soils.^[19][20] Aluminium is not a plant nutrient, and as such, is not actively taken up by the plants, but enters plant roots passively through osmosis. Aluminium can exist in many different forms and is a responsible agent for limiting growth in various parts of the world. Aluminium tolerance studies have been conducted in different plant species to see viable thresholds and concentrations exposed along with function upon exposure.^[21] Aluminium inhibits root growth; lateral roots and root tips become thickened and roots lack fine branching; root tips may turn brown. In the root, the initial effect of Al³⁺ is the inhibition of the expansion of the cells of the rhizodermis, leading to their rupture; thereafter it is known to interfere with many physiological processes including the uptake and transport of calcium and other essential nutrients, cell division, cell wall formation, and enzyme activity.^[19][22]

Proton (H⁺ ion) stress can also limit plant growth. The

In soils with a high content of manganese-containing minerals, Mn toxicity can become a problem at pH 5.6 and lower. Manganese, like aluminium, becomes increasingly soluble as pH drops, and Mn toxicity symptoms can be seen at pH levels below 5.6. Manganese is an essential plant nutrient, so plants transport Mn into leaves. Classic symptoms of Mn toxicity are crinkling or cupping of leaves.^[24]

Nutrient availability in relation to soil pH

Soil pH affects the availability of some plant nutrients:

As discussed above, aluminium toxicity has direct effects on plant growth; however, by limiting root growth, it also reduces the availability of plant nutrients. Because roots are damaged, nutrient uptake is reduced, and deficiencies of the

Molybdenum availability is increased at higher pH; this is because the molybdate ion is more strongly sorbed by clay particles at lower pH.^[28]

Zinc, iron, copper and manganese show decreased availability at higher pH (increased sorption at higher pH).^[28]

The effect of pH on phosphorus availability varies considerably, depending on soil conditions and the crop in question. The prevailing view in the 1940s and 1950s was that P availability was maximized near neutrality (soil pH 6.5–7.5), and decreased at higher and lower pH.^[29][30] Interactions of phosphorus with pH in the moderately to slightly acidic range (pH 5.5–6.5) are, however, far more complex than is suggested by this view. Laboratory tests, glasshouse trials and field trials have indicated that increases in pH within this range may increase, decrease, or have no effect on P availability to plants.^[30][31]

Water availability in relation to soil pH

Strongly alkaline soils are

Many strongly acidic soils, on the other hand, have strong aggregation, good internal drainage, and good water-holding characteristics. However, for many plant species, aluminium toxicity severely limits root growth, and moisture stress can occur even when the soil is relatively moist.^[19]

Plant pH preferences

In general terms, different plant species are adapted to soils of different pH ranges. For many species, the suitable soil pH range is fairly well known. Online databases of plant characteristics, such as USDA PLANTS^[35] and Plants for a Future^[36] can be used to look up the suitable soil pH range of a wide range of plants. Documents like Ellenberg's indicator values for British plants^[37] can also be consulted.

However, a plant may be intolerant of a particular pH in some soils as a result of a particular mechanism, and that mechanism may not apply in other soils. For example, a soil low in molybdenum may not be suitable for soybean plants at pH 5.5, but soils with sufficient molybdenum allow optimal growth at that pH.^[26] Similarly, some calcifuges (plants intolerant of high-pH soils) can tolerate calcareous soils if sufficient phosphorus is supplied.^[38] Another confounding factor is that different varieties of the same species often have different suitable soil pH ranges. Plant breeders can use this to breed varieties that can tolerate conditions that are otherwise considered unsuitable for that species – examples are projects to breed aluminium-tolerant and manganese-tolerant varieties of cereal crops for food production in strongly acidic soils.^[39]

The table below gives suitable soil pH ranges for some widely cultivated plants as found in the USDA PLANTS Database.

In natural or near-natural plant communities, the various pH preferences of plant species (or ecotypes) at least partly determine the composition and biodiversity of vegetation. While both very low and very high pH values are detrimental to plant growth, there is an increasing trend of plant biodiversity along the range from extremely acidic (pH 3.5) to strongly alkaline (pH 9) soils, i.e. there are more calcicole than calcifuge species, at least in terrestrial environments.^[40][41] Although widely reported and supported by experimental results,^[42][43] the observed increase of plant species richness with pH is still in need of a clearcut explanation. Competitive exclusion between plant species with overlapping pH ranges most probably contributes to the observed shifts of vegetation composition along pH gradients.^[44]

pH effects on soil biota

The opposition between acido-tolerance and acido-intolerance is commonly observed at species level within a genus or at genus level within a family, but it also occurs at much higher taxonomic rank, like between soil fungi and bacteria, here too with a strong involvement of competition.^[54] It has been suggested that soil organisms more tolerant of soil acidity, and thus living mainly in soils at pH less than 5, were more primitive than those intolerant of soil acidity.^[55] A cladistic analysis on the collembolan genus Willemia showed that tolerance to soil acidity was correlated with tolerance of other stress factors and that stress tolerance was an ancestral character in this genus.^[56] However the generality of these findings remains to be established.

At low pH, the

Effects of pH on soil biota can be mediated by the various functional interactions of soil foodwebs. It has been shown experimentally that the collembolan Heteromurus nitidus, commonly living in soils at pH higher than 5, could be cultured in more acid soils provided that predators were absent.^[58] Its attraction to earthworm excreta (mucus, urine, faeces), mediated by ammonia emission,^[59] provides food and shelter within earthworm burrows in mull humus forms associated with less acid soils.^[60]

Effects of soil biota on soil pH

Soil biota affect soil pH directly through excretion, and indirectly by acting on the physical environment. Many soil fungi, although not all of them, acidify the soil by excreting oxalic acid, a product of their respiratory metabolism. Oxalic acid precipitates calcium, forming insoluble crystals of calcium oxalate and thus depriving the soil solution from this necessary element.^[61] On the opposite side, earthworms exert a buffering effect on soil pH through their excretion of mucus, endowed with amphoteric properties.^[62]

By mixing organic matter with mineral matter, in particular clay particles, and by adding mucus as a glue for some of them, burrowing soil animals, e.g. fossorial rodents, moles, earthworms, termites, some millipedes and fly larvae, contribute to decrease the natural acidity of raw organic matter, as observed in mull humus forms.^[63][64]

Changing soil pH

Increasing pH of acidic soil

Finely ground

The pH of an alkaline soil can be reduced by adding acidifying agents or acidic organic materials. Elemental sulfur (90–99% S) has been used at application rates of 300–500 kg/ha (270–450 lb/acre) – it slowly oxidizes in soil to form sulfuric acid. Acidifying fertilizers, such as ammonium sulfate, ammonium nitrate and urea, can help to reduce the pH of a soil because ammonium oxidises to form nitric acid. Acidifying organic materials include peat or sphagnum peat moss.^[69]

However, in high-pH soils with a high calcium carbonate content (more than 2%), it can be very costly and/or ineffective to attempt to reduce the pH with acids. In such cases, it is often more efficient to add phosphorus, iron, manganese, copper and/or zinc instead, because deficiencies of these nutrients are the most common reasons for poor plant growth in calcareous soils.^[70][69]

See also

• Acid mine drainage
• Acid sulfate soil
• Cation-exchange capacity
• Fertilizer
• Liming (soil)
• Organic horticulture
• Redox gradient

References