Soil pH is a measure of the acidity or basicity (alkalinity) of a soil. Soil pH is a key characteristic that can be used to make informative analysis both qualitative and quantitatively regarding soil characteristics. pH is defined as the negative logarithm (base 10) of the activity of hydronium ions ( or, more precisely, ) in a solution. In soils, it is measured in a slurry of soil mixed with water (or a salt solution, such as ), and normally falls between 3 and 10, with 7 being neutral. Acidic soils have a pH below 7 and alkaline soils have a pH above 7. Ultra-acidic soils (pH < 3.5) and very strongly alkaline soils (pH > 9) are rare.
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. 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:
Semantic descriptionpH rangeUltra acidic < 3.5Extremely acidic 3.5–4.4Very strongly acidic 4.5–5.0 Strongly acidic 5.1–5.5Moderately acidic 5.6–6.0Slightly acidic 6.1–6.5Neutral 6.6–7.3Slightly alkaline 7.4–7.8Moderately alkaline 7.9–8.4Strongly alkaline 8.5–9.0Very strongly alkaline > 9.0
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: 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. Calcifuge plants (those that prefer an acidic soil) include Erica, Rhododendron and nearly all other Ericaceae species, many birch (Betula), foxglove (Digitalis), gorse (Ulex spp.), and Scots pine (Pinus sylvestris). Calcicole (lime loving) plants include ash trees (Fraxinus spp.), honeysuckle (Lonicera), Buddleja, dogwoods (Cornus spp.), lilac (Syringa) and Clematis species. This observational method has been used to calibrate Ellenberg's reaction indicator values. Use of an inexpensive pH testing kit, wherein 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 colour 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 electrode 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. In the 2010s, spectrophotometric methods were developed to measure soil pH involving addition of an indicator dye to the soil extract. 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. 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.
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.
Sources of acidity
Many processes contribute to soil acidification. These include: Rainfall: Average rainfall has a pH of 5.6 and is moderately acidic due to dissolved atmospheric carbon dioxide () that combines with water to form carbonic acid (). When this water flows through the soil it results in the leaching of basic cations such as bicarbonates; this increases the percentage of and relative to other cations. Root respiration and decomposition of organic matter by microorganisms release which increases the carbonic acid () concentration and subsequent leaching. Plant growth: Plants take up nutrients in the form of ions (e.g. , , , ), and they often take up more cations than anions, in particular in legumes. However, plants must maintain a neutral charge in their roots. In order to compensate for the extra positive charge, they will release 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 () fertilizers react in the soil by the process of nitrification to form nitrate (), and in the process release 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 , 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: Weathering of silicate, aluminosilicate and carbonate minerals containing , , and 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.
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 or dominate the soluble cations.
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. 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.
Acid rain is the main factor to mobilize aluminium from natural sources and the main reason for the environmental effects of aluminium. However, the main factor of presence of aluminium in saltwater and freshwater are the industrial processes that also release aluminium into air. 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).
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. 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 chemical forms and is a responsible agent for limiting plant 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. Aluminium inhibits root growth. Lateral roots and root tips become thickened, roots lack fine branching, and 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 aluminium 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.
Proton (H+ ion) stress can also limit plant growth. The proton pump, H+-ATPase, of the plasmalemma of root cells works to maintain the near-neutral pH of their cytoplasm. A high proton activity (pH within the range 3.0–4.0 for most plant species) in the external growth medium overcomes the capacity of the cell to maintain the cytoplasmic pH and growth shuts down.
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.
Nutrient availability in relation to soil pH
As discussed above, aluminium toxicity has direct effects on plant growth. However, by limiting root growth, aluminium also reduces the availability of plant nutrients. Because roots are damaged, nutrient uptake is reduced, and deficiencies of the macronutrients (nitrogen, phosphorus, potassium, calcium and magnesium) are frequently encountered in very strongly acidic to ultra-acidic soils (pH<5.0). When aluminium levels increase in the soil, it decreases the pH levels. This does not allow for trees to take up water, meaning they cannot photosynthesize, leading them to die. The trees can also develop yellowish colour on their leaves and veins.
Molybdenum availability is increased at higher pH. This is because the molybdate ion is more strongly sorbed by clay particles at lower pH.
Zinc, iron, copper and manganese show decreased availability at higher pH (increased sorption at higher pH).