I'm taking a soil science class for school and one of the chapter topics on pH i found could be beneficial for new growers and experienced growers alike. This research has been provided by the Pennsylvania State University school of Agriculture and Soil Sciences, this is not my work. [FONT="]Soil pH, Acidity and Alkalinity[/FONT] [FONT="]Soil is a mixture of acid/base systems. The generated H+ ions are neutralized by bases. The compounds that generate H+ ions are called acids, and those that accept H+ ions are called bases. Another name for H+ ions is proton. In aqueous soil phase, a proton is hydrated. We call a hydrated proton a hydronium (HO3+) or solvated proton. [/FONT] [FONT="]The process of generating H+ ions in soil is called soil acidification process. Soil acidification can be defined as the decrease in the acid neutralization capacity (ANC) of a soil. The opposite of soil acidification is soil alkalinization which is the increase in ANC.[/FONT] [FONT="]The degree of soil acidity or soil alkalinity (basicity) is characterized in terms of pH.[/FONT] [FONT="]Importance of Soil pH, Acidity and Alkalinity[/FONT] [FONT="]There are countless chemical reactions necessary for plant life and environmental protection that can only occur within a very specific pH range, thus the understanding and management of soil acidity and alkalinity will provide checks and balances to maintain pH within a narrow desirable range.[/FONT] [FONT="]The aim in managing soil acidity and alkalinity is not to achieve a particular pH value, but to adjust the acidity or alkalinity to the point where there are no toxic metals in soil solution and to the point were the availability of nutrients is at its optimum. This condition is usually achieved when the soil pH is between 5.8 and 6.5, however some plants have special acidity requirements [/FONT] [FONT="]Specific Importance of Soil pH[/FONT] [FONT="]Some specific examples of the importance of studying pH.[/FONT] [FONT="]Soil pH influences soil mineral weathering and decomposition of organic matter and the release, solubility, mobility, and hence the concentration and bioavailability of essential and toxic ions. For example aluminum is more active and toxic to plant roots at pH less than 5.5. Therefore, the main cause of reduced plant growth in acid soils is the Al damage to roots.[/FONT] [FONT="]Soil pH affects biological composition of an ecosystem, for example, spruce prevails on soil pH of 4 to 5; hemlock, 5-6; Black Walnut, 6 to 8; standard crested wheatgrass 6.6 to 8.5.[/FONT] [FONT="]Soil pH influences the microbiological compositions and function of microorganisms and related biochemical reactions. For example Rhizobium nitrogen fixing bacteria do not function well under acid conditions such as pH levels less than 4.5[/FONT] [FONT="]Soil pH affects the performance/carryover and activity of some pesticides and herbicides. Hydrolysis, an important non-microbial herbicides breakdown mechanism, slows significantly at soil pH values near 7.0. For example, degradation of some triazine and sulfonylurea herbicides in high pH soil can be reduced enough to result in carryover.[/FONT] [FONT="]Soil pH affects the cation exchange capacity (CEC) of organic matter and soil minerals such as kaolinite, hematite, and goethite. These minerals and organic functional groups have pH-dependent charges.[/FONT] [FONT="]Soil pH is one of several properties used as a general indicator of soil corrosivity. Generally, soils that are either highly alkaline or highly acid are likely to be corrosive to steel. Soils that have pH of 5.5 or lower are likely to be highly corrosive to concrete.[/FONT] [FONT="]Soil pH[/FONT] [FONT="]What is the definition of soil pH?[/FONT] [FONT="]The pH of a solution is defined as the negative logarithm to base 10 of the H+ ion activity or the logarithm of the reciprocal of the H+ ion activity (not the concentration of H+).[/FONT] [FONT="]Mathematically[/FONT] [FONT="]pH = -log10 [H+][/FONT][FONT="]γ[/FONT][FONT="]H+[/FONT] [FONT="]or[/FONT][FONT="][/FONT] [FONT="]pH = -log H+ activity in moles H+/liter[/FONT] [FONT="]or[/FONT][FONT="][/FONT] [FONT="]pH = log (1/H+) activity in miles H+/liter[/FONT] [FONT="]Where[/FONT] [FONT="]A = activity[/FONT] [FONT="][ ] = concentration[/FONT] [FONT="]γ[/FONT][FONT="] = activity coefficient [/FONT] [FONT="]log = Logarithm, which is defined as the power (exponent) to which a base, such as the common logarithm (base 10)[/FONT] [FONT="]pH Scale Concept[/FONT] [FONT="]What is the difference in moles per liter of H+ between a pH 6 solution versus a pH 7 solution?[/FONT] [FONT="]Because the pH of a solution is defined as the negative logarithm to base 10 of the H+ ion activity, a change of one pH unit represents a 10-fold change in concentration of hydrogen ion. For example, log of 100 to base 10 is 2. That is, to what power should we raise 10 to get 100? The answer is 2 because 102 = 100. The negative log is -2. The negative log represented by letter p is used to represent small concentrations. Now suppose the activity of H+ was 0.0001 M or moles per liter. We can write the activity using scientific notation as 10-4 (10 to the power of minus four). Negative log to base 10 of 0.00001 = -log10(0.0001) = 4. Note that Log to base 10 of 0.00001 = log10(0.0001) = -4.[/FONT] [FONT="]As a rule of thumb if the concentration of charged ions present in the solution is below 0.001 M we do not have to be concerned about activities and we can use classic pH definition:[/FONT] [FONT="]pH = minus logarithm base 10 of [H+] = -log[H+][/FONT] [FONT="]The pH value is an expression of the ratio of [H+] to [OH-] (hydroxide ion concentration). Hence, if the [H+] is greater than [OH-], the solution is acidic. Conversely, if the [OH-] is greater than the [H+], the solution is basic.[/FONT] [FONT="]Concentration pH Scale[/FONT] [FONT="][/FONT][FONT="][/FONT] [FONT="]pH expresses the acidity of alkalinity of a solution[/FONT][FONT="][/FONT] [FONT="]pH Range of Soils[/FONT] [FONT="]The concentration of H+ is soils is usually confined to 1 - 10-14M range. Thus soil pH is measured on a scale containing values falling between 0 and 14. On the pH scale pure water has pH 7, and air is slightly acidic - with pH of about 5.7. All values on the pH scale lower than 7 denote soils that are acidic - the lower the pH, the more acidic the soil solution. On the contrary soils with pH above 7 are basic - the higher the pH the more basic the soil solution. pH range in soil is usually from 3 to 8 which is 0.001 M H+ to 0.00000001 M H+ or 10-3 to 10-8 moles of H+ ions per liter of soil solution.[/FONT] [FONT="]As pH changes from 3 to 8, by what factor does the H+ ions decrease? Is the decrease linear or logarithmic?[/FONT] [FONT="]The H+ ions decrease by a factor of 100,000! The decrease is logarithmic and not linear[/FONT] [FONT="]Look at the table below on soil orders; is there any association between the soil characteristics and pH values?[/FONT] [FONT="]pH of Soil Orders[/FONT] [FONT="]What factors may be contributing to the differences in pH of the soil orders?[/FONT] [FONT="]Soil Order[/FONT] [FONT="]Brief Characteristics[/FONT] [FONT="]pH[/FONT] [FONT="]0xisols[/FONT] [FONT="]Very “old†soils-humid tropics, rich in iron and aluminum oxides called "sesquioxides". [/FONT] [FONT="]??[/FONT] [FONT="]Andisols[/FONT] [FONT="]Volcanic ash rich soils[/FONT] [FONT="]??[/FONT] [FONT="]Gelisols[/FONT] [FONT="]Soils on permafrost and consist of mineral or organic material, or both and experienced cryoturbation (frost churning)[/FONT] [FONT="]??[/FONT] [FONT="]Spodosols[/FONT] [FONT="]Soils found in cool, moist environments under coniferous forest vegetation with an eluviated (E) horizon (zone of leaching).[/FONT] [FONT="]4.93[/FONT] [FONT="]Histosols[/FONT] [FONT="]very high content of organic matter[/FONT] [FONT="]5.50[/FONT] [FONT="]Ultisols[/FONT] [FONT="]Highly weathered soils, often red/yellow in color reflecting the oxidation of iron and aluminum[/FONT] [FONT="]5.6[/FONT] [FONT="]Alfisols[/FONT] [FONT="]Soils developed under temperate forests of the humid mid-latitudes[/FONT] [FONT="]6.08[/FONT] [FONT="]Inceptisols[/FONT] [FONT="]Soils just starting to show horizon[/FONT] [FONT="]6.0[/FONT] [FONT="]Mollisols[/FONT] [FONT="]Soils with dark brown to black organic rich surface layers[/FONT] [FONT="]6.51[/FONT] [FONT="]Vertisols[/FONT] [FONT="]Soils rich in expandable (high shrink-swell) clay minerals[/FONT] [FONT="]6.72[/FONT] [FONT="]Aridisols[/FONT] [FONT="]Soils of arid and semiarid environments high in pedogenic carbonates[/FONT] [FONT="]7.26[/FONT] [FONT="]Entisols[/FONT] [FONT="]Soil lacking horizons, just formed[/FONT] [FONT="]7.32[/FONT] [FONT="]Types Of Soil Acidity[/FONT] [FONT="]How do the following types of acidity differ: active, exchangeable, residual, potential, and total acidity.[/FONT] [FONT="]Active acidity[/FONT][FONT="] is due to free H+ and Al3+ ions and is the actual H+ activity of the soil solution estimated by pH measurement at soil to water ratios of 1:1, 1:2.5 or 1:5 in most cases. This ratio widens to 1:25 for peaty soils and peat. The pH values estimate the degree of acidity of the soil- solution, slurry, suspension, or extract. [/FONT] [FONT="]Exchangeable or salt-replaceable acidity[/FONT][FONT="] is the aluminum and hydrogen that can be replaced from an acid soil by an unbuffered salt solution such as KCl or NaCl. It is found by means of displacing H+ and Al3+ ions (acidic cations) from the soil adsorption complex by solutions of neutral salts. Usually 1 M KCl solution is used to extract exchangeable acidity.[/FONT] [FONT="]Non-exchangeable acidity or Residual acidity[/FONT][FONT="] is non-exchangeable but titratable acidity that is neutralized by lime or a buffered salt solution to raise the pH to a specified value (usually 7.0 or 8.0) but which cannot be replaced by an unbuffered salt solution. It can be calculated by subtraction of exchangeable acidity from total acidity.[/FONT] [FONT="]Potential acidity[/FONT][FONT="] is the reserved acidity which includes the exchangeable and non-exchangeable but titratable acidity as well as the acidity arising from oxidation of sulfur compounds in acid sulfate soils. Potential acidity encompasses the buffering capacity of a soil. [/FONT] [FONT="]Total acidity[/FONT][FONT="] includes exchangeable and non-exchangeable acidity and is calculated by subtraction of exchangeable bases from the cation exchange capacity determined by ammonium exchange at pH 7.0. It can also be determined directly using pH buffer-salt mixtures (e.g. BaCl2 plus triethanolamine, pH 8.0 or 8.2) and titrating the basicity neutralized after reaction with a soil.[/FONT] [FONT="]Soil Acid Forming Reactions[/FONT] [FONT="]Many different reactions or processes occur to make a soil acid. Following are some examples:[/FONT] [FONT="]Carbonation:[/FONT][FONT="] Of all the important substances responsible for actual acidity, carbonic acid is the most important in natural soil systems.[/FONT] [FONT="]CO2 + H2O = HCO3- + H+ [/FONT] [FONT="]HCO3- + H2O = H2CO3 + H+ [/FONT] [FONT="]Leaching of basic cations:[/FONT][FONT="] As basic cations get leached downward within the soil profile, percent acid saturation increases, hence pH decreases.[/FONT] [FONT="]% acid saturation = {[Al3+ + H+]/[Na+ + K+ + Mg2+ + Ca2+ + Al3+ + H+]}*100[/FONT] [FONT="]Worldwide, the most common causes of acidity in nonagricultural soils are high rainfall and leaching.[/FONT] [FONT="]Removal of crops and crop residues:[/FONT][FONT="] Removal of crops or trees or any plant material amounts to removal of basic cations and hence an increase in acid saturation. [/FONT] [FONT="]Uptake of cations by plants: [/FONT][FONT="]Plants maintain electrical charge balance in their cells. If a plant takes in more cations, e.g. NH4+, it has to exude the H+ to maintain electrical balance. Thus NH4+ + Root-H+ = Root-NH4+ + H+. This explains why the rhizosphere pH is relatively lower than the surrounding soil pH.[/FONT] [FONT="]Soil Parent Material:[/FONT][FONT="] Soils derived from parent materials low in bases (e.g. granite and obsidian) become more acid easily than basic materials (e.g. limestone and basalt).[/FONT] [FONT="]The Weathering Process:[/FONT][FONT="] The weathering process decomposes 2:1 clay minerals to 1:1 minerals. Soils characterized by 1:1 minerals are rich in Al3+ and Fe3+ and through hydrolysis of these ions, pH decreases due to the regeneration of H+. [/FONT] [FONT="]KAlSi3O8 + 4H2O + 4H+ = K+ + Al3+ + 3Si(OH)4[/FONT] [FONT="]Al3+ + 3H2O = 3H+ + Al(OH)3 (gibbsite)[/FONT] [FONT="]Oxidation:[/FONT][FONT="] Oxidation of pyrite, atmospheric SO2 and NO, and organic sulfur to release H+.[/FONT] [FONT="]Acid rain:[/FONT][FONT="] Acid precipitation (rainfall) that has a pH level of less than 5.6. Two primary sources of acid rain are sulfur dioxide (SO2), and oxides of nitrogen (NOx). Nitrogen monoxide and nitrogen dioxide are all oxides of nitrogen.[/FONT] [FONT="]SO2(g)+O2(g) -> SO3(g) [/FONT] [FONT="]SO3(g)+H2O(l) -> H2SO4(aq) = 2H+ + SO42+[/FONT] [FONT="]2NO2(g) + H2O(l) -> 2H+ + NO3- + NO2[/FONT] [FONT="]NO(g) + NO2(g) + H2O(l) -> 2H+ + 2NO2[/FONT] [FONT="]3NO2(g) + H2O(l) -> 2H+ + 2NO3- + NO(g)[/FONT] [FONT="]Fertilization:[/FONT][FONT="] e.g NH4 based fertilizers[/FONT] [FONT="]NH4+ + 2O2 = NO3- + 2H+ + H2O[/FONT] [FONT="]Decomposition/mineralization of soil organic matter:[/FONT][FONT="] When fresh organic matter decomposes, organic acids are formed.[/FONT] [FONT="]Materials to Raise Soil pH[/FONT] [FONT="]In some situations, it may be necessary to lime a soil to raise pH to a desirable level. Liming materials include:[/FONT] [FONT="]Name[/FONT] [FONT="]Formula[/FONT] [FONT="]*CCE[/FONT] [FONT="]Source[/FONT] [FONT="]Shell meal[/FONT] [FONT="]CaCO3[/FONT] [FONT="]95[/FONT] [FONT="]natural shell deposits[/FONT] [FONT="]Limestone[/FONT] [FONT="]CaCO3[/FONT] [FONT="]100[/FONT] [FONT="]Pure form, fine ground[/FONT] [FONT="]Hydrate (slaked) lime[/FONT] [FONT="]Ca(OH)2[/FONT] [FONT="]120-135[/FONT] [FONT="]Steam burned[/FONT] [FONT="]Quick Lime[/FONT] [FONT="]CaO[/FONT] [FONT="]150-175[/FONT] [FONT="]Kiln burned[/FONT] [FONT="]Dolomite[/FONT] [FONT="]CaCO3MgCO3[/FONT] [FONT="]100[/FONT] [FONT="]Natural deposits[/FONT] [FONT="]Sugarbeet lime[/FONT] [FONT="]CaCO3[/FONT] [FONT="]80-90[/FONT] [FONT="]Sugarbeet byproduct[/FONT] [FONT="]Calcium Silicate[/FONT] [FONT="]CaSiO3[/FONT] [FONT="]60-80[/FONT] [FONT="]Slag, a byproduct of iron industry[/FONT] [FONT="]*CCE = acid neutralizing capacity relative to pure CaCO3[/FONT][FONT="][/FONT] [FONT="]The basic reactions that take place in the soil are:[/FONT] [FONT="]CaO + H2O + = Ca(OH)2 Ca(OH)2 + 2H+ = Ca2+ + 2H2O[/FONT] [FONT="]CaCO3 + H2O = Ca2+ + HCO3-+ OH- HCO3- + H2O = H2CO3 + OH- H2CO3 = H2O + CO2[/FONT] [FONT="]2H+colloid + CaCO3 ---> Ca2+colloid + CO2↑ + H2O[/FONT] [FONT="]Ca2+ from lime replaces two H+ ions on the cation exchange complex. The H+ ions combine with OH- (hydroxyl) ions to form water. pH increases because the acidity source (H+) has been reduced.[/FONT] [FONT="]Ca2+ does not raise the pH.[/FONT] [FONT="]Materials to Lower Soil pH[/FONT] [FONT="]In some situations, it may be necessary to acidify a soil to lower pH to a desirable level. Acidifying materials include:[/FONT] [FONT="]Elemental sulfur[/FONT][FONT="] – the cheapest material to use for acidifying soil but may require up to four months for the conversion of sulfur to the sulfate ion by certain soil microorganisms.[/FONT] [FONT="]2S + 3O2 + 2H2O = 2H2SO4 = 2SO42- + 4H+[/FONT] [FONT="]Iron sulfate[/FONT][FONT="] – Though more expensive, it lowers pH within weeks instead of months. However, about 7 times more iron sulfate than elemental sulfur is required to lower a soil´s pH.[/FONT] [FONT="]Fe2(SO4)3 + 6H2O = 2Fe(OH)3 + H2SO4[/FONT] [FONT="]Sphagnum moss peat [/FONT][FONT="]– Though rather expensive and requires large amounts (1 part peat to 1 part soil by volume), it has a pH of 4.0-4.5 and will lower pH for a few years after incorporation.[/FONT] [FONT="]Ammonium sulfate [/FONT][FONT="]– Is the most acidifying nitrogen fertilizer and is recommended to apply as nitrogen source based on nitrogen requirements of the plant.[/FONT] [FONT="](NH4)2SO4 + 4O2 = 2NO3- + SO42+ + 4H+ + 2H2O[/FONT] [FONT="]Aluminum sulfate [/FONT][FONT="]– Is similar to iron sulfate for lowering pH but addition of Al may lead to plant toxicity.[/FONT] [FONT="]Al2(SO4)3 + 6H2O = 2Al(OH)3 + H2SO4[/FONT] [FONT="]Soil pH And Acidity Determination[/FONT] [FONT="]A pH measurement is normally made by means of either a colorimeter or an electrometric method. The colorimeter method involves suitable dyes or acid-base indicators, the colors of which change with H-ion activity. The electrode method involves a glass, H+ - sensing (indicator) electrode paired with a reference electrode attached to a suitable meter for measuring electromotive force (emf), which is proportional to pH.[/FONT] [FONT="]Soil solution pH = -log10 [H+] of a soil solution.[/FONT] [FONT="]Amounts of total soil acidity are most often measured in one of the following ways:[/FONT] [FONT="]Titration with base or equilibration with lime.[/FONT] [FONT="]Leaching with a buffered solution, followed by titration of the leachate.[/FONT] [FONT="]Subtracting the sum of exchangeable bases from CEC.[/FONT] [FONT="]Equilibration with a buffered solution and estimation of acidity by pH change.[/FONT] [FONT="]Extraction with an unbuffered solution neutral of a salt.[/FONT]