The soil pH is closely linked to the concepts alkalinity and acidity (see acid neutralizing capacity). A neutral solution has pH 7 while an acid solution has pH less than 7 (more H+ than OH-) and a basic solution pH larger than 7 (more OH- than H+) but there is, contrary to popular belief, no exact limit to the pH range. In natural soils and surface waters buffer systems make pH levels below 3 uncommon, but not impossible. The exposure of the soil to sunlight does not usually affect the pH of the soil.
(NOTE: While a basic solution always has a pH larger than 7, an alkaline solution (i.e. a solution with positive acid neutralizing capacity) does not necessarily have a pH larger than 7. For details on the relation between pH and ANC, see acid neutralizing capacity)
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.
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.
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 wet), 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 realize 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 sea salt 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 CO2 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 CaCl2, overhead shaking of soil sample with added water, etc. The CaCl2 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 (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 deficiency 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.
To understand how acid soils are formed, take a simple walk through a woodland. Rainfall filters through trees and into the ground, where it dissolves limestone sediment and other alkaline minerals that help neutralize soil acidity. The woodland floor is carpeted in needles of conifers, leaves of hardwood trees, and other dead plant matter, all of which increase soil acidity as they decompose. Unless this woodland is on top of a huge deposit of gypsum or other alkaline minerals, the soil will tend to be acid.
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 held 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.
The pH value of a soil is influenced by the kinds of parent materials from which the soil was formed. Soils developed from basic rocks generally have higher pH values than those formed from acid rocks.
Rainfall also affects soil pH. Water passing through the soil leaches basic nutrients such as calcium and magnesium from the soil. They are replaced by acidic elements such as aluminum and iron. For this reason, soils formed under high rainfall conditions are more acidic than those formed under arid (dry) conditions.
Human distractions like pollution alter the pH of soil. Researches have also revealed that soil pH is affected by the vehicular and ongoing traffic. This largely hampers the soil pH and in turn the primary productivity by compacting the soil and decreasing its friability.
Application of fertilizers containing ammonium or urea speeds up the rate at which acidity develops. The decomposition of organic matter also adds to soil acidity.
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 the decay of organic matter, which immeasurably benefit the soil. These processes also prefer near-neutral conditions.
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.
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:
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:
The pH of an alkaline soil is lowered by adding sulphur, 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.