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Soil pH (Wiki)

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  (Redirected from Soil acidity)

The soil pH is a measure of the acidity or basicity in soils. pH is defined as the negative logarithm (base 10) of the activity of hydronium ions (H+ or, more precisely, H3O+aq) in a solution. It ranges from 0 to 14, with 7 being neutral. A pH below 7 is acidic and above 7 is basic. Soil pH is considered a master variable in soils as it controls many chemical processes that take place. It specifically affects plant nutrient availability by controlling the chemical forms of the nutrient. The optimum pH range for most plants is between 5.5 and 7.0,[1] however many plants have adapted to thrive at pH values outside this range.

Global variation in soil pH. Red = acidic soil. Yellow = neutral soil. Blue = alkaline soil. Black = no data.

Classification of soil pH ranges

The United States Department of Agriculture Natural Resources Conservation Service, formerly Soil Conservation Service classifies soil pH ranges as follows: [2]

Denomination pH range
Ultra acid < 3.5
Extreme acid 3.5–4.4
Very strong acid 4.5–5.0
Strong acid 5.1–5.5
Moderate acid 5.6–6.0
Slight acid 6.1–6.5
Neutral 6.6–7.3
Slightly alkaline 7.4–7.8
Moderately alkaline 7.9–8.4
Strongly alkaline 8.5–9.0
Very strongly alkaline > 9.0

Sources of soil pH

Soil acidity

[3] Acidity in soils comes from H+ and Al3+ ions in the soil solution and sorbed to soil surfaces. While pH is the measure of H+ in solution, Al3+ is important in acid soils because between pH 4 and 6, Al3+ reacts with water (H2O) forming AlOH2+, and Al(OH)2+, releasing extra H+ ions. Every Al3+ ion can create[clarification needed] 3 H+ ions. Many other processes contribute to the formation of acid soils including rainfall, fertilizer use, plant root activity and the weathering of primary and secondary soil minerals. Acid soils can also be caused by pollutants such as acid rain and mine spoilings.

  • Rainfall: Acid soils are most often found in areas of high rainfall. Excess rainfall leaches base cation from the soil, increasing the percentage of Al3+ and H+ relative to other cations. Additionally, rainwater has a slightly acidic pH of 5.7 due to a reaction with CO2 in the atmosphere that forms carbonic acid.
  • Fertilizer use: Ammonium (NH4+) fertilizers react in the soil in a process called nitrification to form nitrate (NO3−), and in the process release H+ ions.
  • Plant root activity: Plants take up nutrients in the form of ions (NO3, NH4+, Ca2+, H2PO4, etc.), and often, they take up more cations than 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 will 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).
  • Weathering of minerals: Both primary and secondary minerals that compose soil contain Al. As these minerals weather, some components such as Mg, Ca, and K, are taken up by plants, others such as Si are leached from the soil, but due to chemical properties, Fe and Al remain in the soil profile. Highly weathered soils are often characterized by having high concentrations of Fe and Al oxides.
  • Acid Rain: When atmospheric water reacts with sulfur and nitrogen compounds that result from industrial processes, the result can be the formation of sulfuric and nitric acid in rainwater. However the amount of acidity that is deposited in rainwater is much less, on average, than that created through agricultural activities.
  • Mine Spoil: Severely acidic conditions can form in soils near mine spoils due to the oxidation of pyrite.
  • decomposition of organic matter by micro organisms releases co2 which when mixed with soil water forms weak carbonic acid (H2CO3)

Sources of basicity

Basic soils have a high saturation of base cations (K+, Ca2+, Mg2+ and Na+). This is due to an accumulation of soluble salts are classified as either saline soil, sodic soil, saline-sodic soil or alkaline soil. All saline and sodic soils have high salt concentrations, with saline soils being dominated by calcium and magnesium salts and sodic soils being dominated by sodium. Alkaline soils are characterized by the presence of carbonates.[clarification needed]

Effect of soil pH on plant growth

Acid affected soils

[4] Plants grown in acid soils can experience a variety of symptoms including aluminium (Al), hydrogen (H), and/or manganese (Mn) toxicity, as well as potential nutrient deficiencies of calcium (Ca) and magnesium (Mg).
Aluminium toxicity is the most widespread problem in acid soils. Aluminium is present in all soils, but dissolved Al3+ is toxic to plants; Al3+ is most soluble at low pH, above pH 5.2 little aluminum is 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 damages roots in several ways: In root tips and Aluminium interferes with the uptake of Calcium, an essential nutrient, as well as bind with phosphate and interfere with production of ATP and DNA, both of which contain phosphate. Aluminium can also restrict cell wall expansion causing roots to become stunted.

Below pH 4, H+ ions themselves[clarification needed] damage root cell membranes.

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

Nutrient availability in relation to soil pH

[5] Nutrients needed in large amounts by plants are referred to as macronutrients and include nitrogen (N), phosphorus (P), potassium (K), calcium (Ca), magnesium (Mg) and sulfur (S). Elements that plants need in trace amounts are called trace nutrients or micronutrients. Trace nutrients are not major components of plant tissue but are essential for growth. They include iron (Fe), manganese (Mn), zinc (Zn), copper (Cu), cobalt (Co), molybdenum (Mo), and boron (B). Both macronutrient and micronutrient availability are affected by soil pH. In slightly to moderately alkaline soils, molybdenum and macronutrient (except for phosphorus) availability is increased, but P, Fe, Mn, Zn Cu, and Co levels are reduced and may adversely affect plant growth. In acidic soils, micronutrient availability (except for Mo and Bo) is increased. Nitrogen is supplied as ammonium (NH4) or nitrate (NO3) in fertilizer amendments, and dissolved N will have the highest concentrations in soil with pH 6–8. Concentrations of available N are less sensitive to pH than concentration of available P. In order for P to be available for plants, soil pH needs to be in the range 6.0 and 7.5. If pH is lower than 6, P starts forming insoluble compounds with iron (Fe) and aluminium (Al) and if pH is higher than 7.5 P starts forming insoluble compounds with calcium (Ca).

Nutrient availability in relation to soil pH[6]

Most nutrient deficiencies can be avoided between a pH range of 5.5 to 6.5, provided that soil minerals and organic matter contain the essential nutrients to begin with.

Determining pH

Methods of determining pH include:

  • Observation of soil profile: Certain profile characteristics can be indicators of either acid, saline, or sodic conditions. Strongly acidic soils often have poor incorporation of the organic surface layer with the underlying mineral layer. The mineral horizons are distinctively layered in many cases, with a pale eluvial (E) horizon beneath the organic surface; this E is underlain by a darker B horizon in a classic podzol horizon sequence. This is a very rough gauge of acidity as there is no correlation between thickness of the E and soil pH. E horizons a few feet thick in Florida usually have pH just above 5 (merely "strongly acid") while E horizons a few inches thick in New England are "extremely acid" with pH readings of 4.5 or below.[1][2] [3] In the southern Blue Ridge Mountains there are "ultra acid" soils, pH below 3.5, which have no E horizon.[4] Presence of a caliche layer indicates the presence of calcium carbonates, which are present in alkaline conditions. Also, columnar structure can be an indicator of sodic condition.[7]
  • 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.
  • Use of an inexpensive pH testing kit based on barium sulphate in powdered form, where in a small sample of soil is mixed with water which changes colour according to the acidity/alkalinity.
  • 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 alkaline, blue.
  • Use of a commercially available electronic pH meter, in which a rod is inserted into moistened soil and measures the concentration of hydrogen ions.

Examples of plant pH preferences

This section may contain original research. Please improve it by verifying the claims made and adding inline citations. Statements consisting only of original research may be removed. (November 2010)

Changing soil pH

Increasing pH of acidic soil

The most common amendment to increase soil pH is lime (CaCO3 or MgCO3), usually in the form of finely ground agricultural lime. The amount of lime needed to change pH is determined by the mesh size of the lime (how finely it is ground)and the buffering capacity of the soil. A high mesh size (60–100) indicates a finely ground lime, that will react quickly with soil acidity. Buffering capacity of soils is a function of a soils cation exchange capacity, which is in turn determined by the clay content of the soil, the type of clay and the amount of organic matter present. Soils with high clay content, particularly shrink–swell clay, will have a higher buffering capacity than soils with little clay. Soils with high organic matter will also have a higher buffering capacity than those with low organic matter. Soils with high buffering capacity require a greater amount of lime to be added than a soil with a lower buffering capacity for the same incremental change in pH.
Other amendments that can be used to increase the pH of soil include wood ash, industrial CaO (burnt lime), and oyster shells. White firewood ash includes metal salts which are important for processes requiring ions such as Na+ (sodium), K+ (potassium), Ca2+ (calcium), which may or may not be good for the select flora, but decreases the acidic quality of soil.
These products increase the pH of soils through the reaction of CO32− with H+ to produce CO2 and H2O. Calcium silicate neutralizes active acidity in the soil by removing free hydrogen ions, thereby increasing pH. As its silicate anion captures H+ ions (raising the pH), it forms monosilicic acid (H4SiO4), a neutral solute.

Decreasing pH of basic soil

See also


  1. ^ Perry, Leonard. "PH for the Garden". Retrieved 11 December 2012.
  2. ^ Soil Survey Division Staff. "Soil survey manual.1993. Chapter 3, selected chemical properties.". Soil Conservation Service. U.S. Department of Agriculture Handbook 18. Retrieved 2011-03-12.
  3. ^ Sparks, Donald; Environmental Soil Chemistry. 2003, Academic Press, London, UK
  4. ^ Brady, N. and Weil, R. The Nature and Properties of Soils. 13th ed. 2002
  5. ^ http://www.extension.org/pages/9875/soil-ph-and-nutrient-availability
  6. ^ Finck, Arnold (1976). Pflanzenernährung in Stichworten. Kiel: Hirt. p. 80. ISBN 3-554-80197-6.
  7. ^ Buol, S. W., R. J. Southard, R.C. Graham and P.A. McDaniel. Soil Genesis and Classification. (5th) Edition, Ia. State Press p. 494. 2002
  8. ^ Brady, N. and Weil, R. The Nature and Properites of Soils. 13th ed. 2002

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