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Soil pH

2007 Schools Wikipedia Selection. Related subjects: Geology and geophysics

   Soil pH is an indication of the alkalinity or acidity of soil. It is
   based on the measurement of pH, which is based in turn on the activity
   of hydrogen ions (H^+) in a water or salt solution.

   When in balance (pH 7) the soil is said to be neutral. The pH scale
   covers a continuum ranging from 0 (very acidic) to 14 (very alkaline or
   basic). It is however uncommon to find soils at either extreme of this
   range. Under many conditions soils tend to become more acid or alkaline
   over time if steps are not taken to maintain a balance.

   (NOTE: Alkaline and basic are not interchangeable. However, aside from
   uncommon examples such as Ammonias, testing Alkaline and testing pH
   bring about similar results. The higher the Alkalinity, the greater the
   tendency towards a base.)

   Soil pH is an important consideration for farmers and gardeners for
   several reasons, including the fact that many plants and soil life
   forms prefer either alkaline or acidic conditions, that some diseases
   tend to thrive when the soil is alkaline or acidic, and that the pH can
   affect the availability of nutrients in the soil.

Nutrient availability in relation to soil pH

   The majority of food crops prefer a neutral or slightly acidic soil.
   Some plants however prefer more acidic (e.g., potatoes, strawberries)
   or alkaline ( brassicas) conditions.
                  Acid                  Neutral              Alkali
                  4   4.5 5   5.5 6   6.5 7   7.5 8   8.5 9   9.5 10
   nitrogen, N
   phosphorus, P
   potassium, K
   calcium, Ca
   magnesium, Mg
   sulfur, S
   iron, Fe
   manganese, Mn
   boron, B
   copper, Cu
   zinc, Zn
   molybdenum, Mo

   The above table gives a guide to the availability of several nutrients
   at various pH values

   During the acidification process the decrease in pH results in a
   release of positively charged ions (cations) from the cation exchange
   surfaces (organic matter and clay minerals). In the short term
   acidification thus increases the concentration of potassium (K),
   magnesium (Mg), and calcium (Ca) in soil solution. Once the cation
   exchange surface has become depleted of these ions, however, the
   concentration in soil solution can be quite low and is largely
   determined by the weathering rate. The weathering rate in turn is
   dependent on such things as mineralogy (e.g. presence of easily
   weathered minerals), surface area (i.e. the soil texture), soil
   moisture (i.e. how large a fraction of the mineral surface area that is
   wetted), pH, concentration of base cations such as Ca, Mg and K as well
   as concentration of Aluminium. The amount of plant available nutrients
   is a much more difficult issue than soil solution concentrations. The
   term plant available nutrients usually include pools other than soil
   solution but which are supposed to replenish soil solution pretty fast
   e.g. through cation exchange. One reason for including such pools is
   the plants capability of releasing organic acids which increase the
   total soil solution concentration of some cation nutrients that are
   important for the plant.

   It is thus important to realise that there exists no simple relation
   between soil solution concentration of Ca, Mg and K and reasonable
   pH-values. The reason for this is that Ca, Mg and K are base cations,
   i.e. cations of strong bases and strong bases are fully dissociated at
   the pH-ranges occurring in most natural waters. However, as the soil
   solution pH is dependent on mineral weathering and mineral weathering
   increase pH by releasing Ca, Mg and K a soil which is rich in easily
   weatherable minerals tends to have both a higher pH and higher soil
   solution concentration of Ca, Mg and K. On the other hand deposition of
   sulphate, nitrate and to some extent ammonia decrease pH of soil
   solution essentially without affecting Ca, Mg and K concentrations
   whereas deposition of seasalt increases Ca, Mg and K concentrations
   without having much of an effect on soil solution pH.

   When interpreting soil solution pH values it is essential to take into
   account the method by which pH has been measured. Depending on whether
   or not the water has been equilibrated with ambient CO[2] pressure or
   not the pH reported from the same site may be either high or low. This
   is simply because the carbon dioxide pressure deep down in the soil
   might be 10–20 times higher than the ambient pressure due to
   decomposition of organic material. The higher carbon dioxide pressure
   result in more carbonic acid and hence a lower pH. Furthermore, soil
   solution can be extracted from the soil in many ways, e.g. by
   lysimeters, zero-tension lysimeters, centrifugation, extraction with
   CaCl[2], overhead shaking of soil sample with added water, etc. The
   CaCl[2] extraction method do not give the actual soil solution pH but
   rather a mix between soil solution pH and what is easily available e.g.
   through cation exchange. Also when mixing soil samples with water and
   using overhead shakers (or similar) the result is a mix between actual
   soil solution and cation exchange, although the hope is that the
   extracted water will be similar to the actual soil solution in most
   respects. If centrifugation or pressurised lysimeters are used, care
   must be taken that the extracted water do not include water that is not
   readily available (think wilting point and crystal water). Naturally,
   taking a sample introduces a disturbance of the system, which can e.g.
   result in a change in nutrient uptake and decomposition rates (e.g. due
   to cutting of fine roots when placing the lysimeter).

   Many nutrient cations such as zinc (Zn^2+), aluminium (Al^3+), iron
   (Fe^2+), copper (Cu^2+), cobalt (Co^2+), and manganese (Mn^2+) are
   soluble and available for uptake by plants below pH 5.0, although their
   availability can be excessive and thus toxic in more acidic conditions.
   In more alkaline conditions they are less available, and symptoms of
   nutrient defficiency may result, including thin plant stems, yellowing
   ( chlorosis) or mottling of leaves, and slow or stunted growth.

   pH levels also affect the complex interactions among soil chemicals.
   Phosphorus (P) for example requires a pH between 6.0 and 7.0 and
   becomes chemically immobile outside this range, forming insoluble
   compounds with iron (Fe) and aluminium (Al) in acid soils and with
   calcium (Ca) in calcareous soils.

Soils and acidity

   Under conditions in which rainfall exceeds evapotranspiration
   (leaching) during most of the year, the basic soil cations (Ca, Mg, K)
   are gradually depleted and replaced with cations helds in colloidal
   soil reserves, leading to soil acidity. Clay soils often contain Fe and
   hydroxy Al, which affect the retention and availability of fertilizer
   cations and anions in acidic soils.

   Soil acidification may also occur by addition of hydrogen, due to
   decomposition of organic matter, acid-forming fertilizers, and exchange
   of basic cations for H^+ by the roots.

   Soil acidity is reduced by volatilization and denitrification of
   nitrogen. Under flooded conditions, the soil pH value increases. In
   addition, the following nitrate fertilizers -- calcium nitrate,
   magnesium nitrate, potassium nitrate and sodium nitrate -- also
   increase the soil pH value.

   Some alkaline soils have Calcium in the form of limestone that is not
   chemically available to plants. In this case sulfuric acid or Sulfur
   may be added to reclaim the soil.

Soil life and pH

   A pH level of around 6.3-6.8 is also the optimum range preferred by
   most soil bacteria, although fungi, molds, and anaerobic bacteria have
   a broader tolerance and tend to multiply at lower pH values. Therefore,
   more acidic soils tend to be susceptible to souring and putrefaction,
   rather than undergoing the sweet decay processes associaeed the decay
   of organic matter, immeasurably benefitting the soil, also prefer these
   near-neutral conditions.

pH and plant diseases

   Many plant diseases are caused or exacerbated by extremes of pH,
   sometimes because this makes essential nutrients unavailable to crops
   or because the soil itself is unhealthy (see above). For example,
   chlorosis of leaf vegetables and potato scab occur in overly alkaline
   conditions, and acidic soils can cause clubroot in brassicas.

Determining pH

   PH is not constant in soil or water, but varies on a seasonal or even
   daily basis due to factors such as rainfall, biological growth within
   the soil, and temperature changes. Rather, a map of the pH level is a
   mosaic, varying according to soil crumb structure, on the surface of
   colloids, and at microsites. The pH also exhibits vertical gradients,
   tending to be more acidic in surface mulches and alkaline where
   evaporation, wormcasts, and capillary action draw bases up to the soil
   surface. It also varies on a macro level depending on factors such as
   slope, rocks, and vegetation type.Therefore the pH should be measured
   regularly and at various points within the land in question.

   Methods of determining pH include:
     * Observation of predominant flora. Calcifuge plants (those that
       prefer an acidic soil) include Erica, Rhododendron and nearly all
       other Ericaceae species, many Betula ( birch), Digitalis (
       foxgloves), gorse, and Scots Pine. Calcicole (lime loving) plants
       include Fraxinus ( Ash), Honeysuckle (Lonicera), Buddleia, Cornus
       spp ( dogwoods), Lilac(Syringa) and Clematis spp.
     * Observation of symptoms that might indicate acidic or alkaline
       conditions, such as occurrence of the plant diseases mentioned
       above or salinisation of alkaline soils. The house hydrangea
       (Hydrangea macrophylla) produces pink flowers at pH values of 6.8
       or higher, and blue flowers at pH 6.0 or below.
     * Use of an inexpensive pH testing kit based on barium sulfate in
       powdered form, wherein 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 acid 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.

Altering soil pH

   The aim when attempting to adjust soil acidity is not so much to
   neutralise the pH as to replace lost cation nutrients, particularly
   calcium. This can be achieved by adding limestone to the soil, which is
   available in various forms:
     * Agricultural lime (ground limestone or chalk). These are natural
       forms of calcium carbonate which are extracted in the UK from areas
       such as the Mendips and Salisbury Plain. This is probably the
       cheapest form of lime for gardening and agricultural use and can be
       applied at any time of the year. These forms are slow reacting,
       thus their effect on soil fertility and plant growth is steady and
       long lasting. Ground lime should be applied to clay and heavy soils
       at a rate of about 500 to 1,000 g/m² (1 to 2 lb/yd² or 4,500 to
       9,000 lb/ac).
     * Quicklime and slaked lime. The former is produced by burning rock
       limestone in kilns. It is highly caustic and cannot be applied
       directly to the soil. Quicklime reacts with water to produce
       slaked, or hydrated, lime, thus quicklime is spread around
       agricultural land in heaps to absorb rain and atmospheric moisture
       and form slaked lime, which is then spread on the soil. Quicklime
       should be applied to heavy clays at a rate of about 400 to 500 g/m²
       (0.75 to 1 lb/yd² or 3,600 to 4,500 lb/ac), hydrated lime at 250 to
       500 g/m² (0.5 to 1 lb/yd²). However, quicklime and hydrated lime
       are very fast acting and are not suitable for inclusion in an
       organic system. Their use is prohibited under the standards of both
       The Soil Association and the Henry Doubleday Research Association.
     * Calcium sulfate, also known as gypsum can not be used to amend soil
       acidity. It is a common myth that gypsum effects soil acidity.

   The pH of an alkaline soil is lowered by adding sulfur, iron sulfate or
   aluminium sulfate, although these tend to be expensive, and the effects
   short term. Urea, urea phosphate, ammonium nitrate, ammonium
   phosphates, ammonium sulfate and monopotassium phosphate also lower
   soil pH.
   Retrieved from " http://en.wikipedia.org/wiki/Soil_pH"
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