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Soil pH

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 concentration 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.

pH is important for the organic gardener for several reasons, including the fact that many plants and soil life forms prefer either acid or alkaline 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.

Table of contents

Nutrient availability in relation to soil pH

The majority of food crops prefer a neutral or slightly acidic soil, because the solubility of most nutrients necessary for healthy plant growth is highest at at pH 6.3-6.8. Some plants however prefer more acidic (e.g., potatos, strawberries) or alkaline (brassicas) conditions.

When the pH falls below 5.5, most major plant-nutrient minerals (those needed in substantial quantities to promote healthy plant growth include nitrogen (N), phosphorus (P), potassium (K), sulphur (S), magnesium (Mg), and calcium (Ca)) and some micronutrients[?] (elements important to plant growth in very small amounts) become insoluble and hence unavailable for uptake by plant roots.

Many cationic (positively charged) nutrients such as zinc (Zn2+), aluminium (Al3+), iron (Fe2+), copper (Cu2+), cobalt (Co2+), and manganese (Mn2+) 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.

This table indicates the availability of several nutrients at various pH values:

  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    

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 fertilisers, 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, nitrate fertilizers containing cations such as Ca, K, or Na also increase the soil pH value.

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 associated with a healthy, living soil. Earthworms, whose feeding and tunnelling activities aerate the soil and speed 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 varys 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 Betula (birch), Digitalis (foxgloves), ericas (heath), azaleas, gorses, and Scots Pine. Calcicole (lime loving) plants include Fagus spp (Beech), Fraxinus (Ash), Honeysuckle (Lonericera), Buddleia, Cornus spp (dogwoods[?]), and Corylus spp (hazel[?]).
  • Observation of symptoms that might indicate acidic or alkaline conditions, such as occurrence of the plant diseases mentioned above or salinisation[?] of alkaline soils.
  • Use of an inexpensive pH testing kit based on barium sulphate 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:

  • Ground limestone and ground 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 1-2 lb per square yard.
  • 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 0.75 to 1 lb per square yard, hydrated lime at 0.5 to 1 lb. 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 sulphate, known as gypsum can be used to amend[?] soil acidity and is also useful for lightening the structure of heavy clays. Gypsum can be purchased or can sometimes be obtained from old domestic plaster.

The pH of an alkaline soil is lowered by the adding sulphur, although this tends to be expensive, and the effects short term.

See also



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