How to acidify soil

The Four Things You Need to Know About Soil pH

Don’t be too quick to blame horrendous-sounding afflictions like “verticillium” and “fusarium” or any other diseases for the sickly yellowing of your pin oak’s or geranium’s leaves. The problem may be that your soil’s pH is out of whack. Every plant has its preferred range of soil acidity, and when the pH level is out of that range, a host of ills may follow. A basic understanding of pH will not only help keep your garden healthy but also assist you if things go bad. Here is what you need to know to make smart decisions about managing your soil’s pH.

1. What is pH?
The acidity or alkalinity of a substance is measured in pH units, a scale running from 0 to 14. A pH of 7 is neutral. As numbers decrease from 7, the acidity gets higher. As numbers increase from 7 so does the alkalinity. Soils generally range from an extremely acidic pH of 3 to a very alkaline pH of 10. This range is a result of many factors, including a soil’s parent material and the amount of yearly rainfall an area receives. Most cultivated plants enjoy slightly acidic conditions with a pH of about 6.5. Pin oak, gardenia, blueberry, azalea, and rhododendron are among the plants that demand a very acidic pH of 4.5 to 5.5.

2. What does pH do? Soil pH has indirect yet far-reaching effects on plants. Plant nutrients become available or unavailable according to the soil’s pH level (chart, right). Yellowing between the veins of young leaves indicates an iron deficiency, a condition arising not from a lack of iron in the soil but from insufficient soil acidity to put iron into a form that a plant can absorb. Most plants thrive in slightly acidic soil because that pH affords them good access to all nutrients.

The darker side of soil pH is plant poisoning. Too low a pH level can render the plant nutrient manganese available at toxic levels; geraniums are particularly sensitive to this, showing their discomfort with yellowed, brown-flecked, or dead leaves. A pH level that is too low also liberates aluminum—not a plant nutrient—in amounts that can stunt root growth and interfere with a plant’s uptake of nutrients. At a high pH level, the plant nutrient molybdenum becomes available in toxic amounts.

Soil pH also influences soil-dwelling organisms, whose well-being, in turn, affects soil conditions and plant health. The slightly acidic conditions enjoyed by most plants are also what earthworms like, as do microorganisms that convert nitrogen into forms that plants can use.

3. How do you adjust your pH?
Before attempting to change your soil’s pH, you must know its current level. This will determine how much you need to raise or lower it, if at all. A simple soil test can be done at home or by a soil-testing laboratory. You must also know your soil’s texture, be it clay, sand, or something in between. More material is needed to change the pH level of a clay soil than for a sandy soil because the charged surfaces of clays make them more resistant to pH changes than the uncharged surfaces of sand particles.

Generally, limestone is used to raise a pH level, and sulfur is used to lower it. Limestone is relatively pure calcium carbonate, but dolomitic limestone is a mix of calcium carbonate and magnesium. Pound for pound, dolomitic limestone neutralizes more acidity than pure limestone and adds magnesium to the soil, perfect for those who garden in the East or the Pacific Northwest where this nutrient is naturally low.

Limestone and sulfur are available in powdered or pelletized form, with the latter being easier to spread uniformly and causing less of a health hazard from dust. Avoid using powdered sulfur sold as a fungicide because it is finer and more expensive than needed for acidifying soil. Neither limestone nor sulfur is soluble in water, so mix these materials thoroughly into the top 6 inches of soil when quick action is needed. Otherwise, just lay the material on top of the ground, and let it gradually work its way down.

4. Why should you monitor your pH? Once the pH level is adjusted for the plants you are growing, do not put it out of your mind. Maintaining the correct pH level for your soil is an ongoing task, especially in the naturally acidic soils of the East and the Northwest, where rainfall leaches out calcium and other alkaline-forming elements. Naturally alkaline soils will keep shifting up the pH scale because of the rock minerals from which they were formed. In some cases, acidifying these soils is unfeasible. Even fertilizers can shift your soil pH over time, with materials such as ammonium sulfate and ammonium nitrate pushing the pH level lower and potassium nitrate or calcium pushing the value higher. Hence, there’s a need for regular additions of limestone or sulfur.

Soil pH

A. pH

Soil pH influences solubility, concentration in soil solution, ionic form, and mobility of micronutrients in soil, and consequently acquisition of these elements by plants (Fageria, Baligar and Edwards, 1990; Fageria, Baligar, and Jones, 1997). As a rule, the availability of B, Cu, Fe, Mn, and Zn usually decreases, and Mo increases as soil pH increases. These nutrients are usually adsorbed onto sesquioxide soil surfaces. Table IV summarizes important changes in micronutrient concentrations as influenced by soil pH and consequent acquisition by plants. Table V has been provided to show acquisition of Cu, Fe, Mn, and Zn by rice grown at various soil pH values.

Table IV. Influence of Soil pH on Micronutrient Concentrations in Soil and Plant Uptakea

Element Influence on concentration/uptake
B Increasing soil pH favors adsorption of B. This element generally becomes less available to plants. Availability and uptake of B decrease dramatically at pH > 6.0.
Cl Chloride is bound tightly by most soils in mildly acid to neutral pH soils and becomes negligible to pH 7.0. Appreciable amounts can be adsorbed with increasing soil acidity, particularly by Oxisols and Ultisols, which are dominated by kaolinitic clay. Increasing soil pH generally increases Cl uptake by plants.
Cu Solubility of Cu2+ is very soil pH dependent and decreases 100-fold for each unit increase in pH. Plant uptake also decreases.
Fe Ferric (Fe3+) and ferrous (Fe2+) activities in soil solution decrease 1000-fold and 100-fold, respectively, for each unit increase in soil pH. In most oxidized soils, uptake of Fe by crop plants decreases with increasing soil pH.
Mn The principal ionic Mn species in soil solution is Mn2+, and concentrations decrease 100-fold for each unit increase in soil pH. In extremely acid soils, Mn2+ solubility can be sufficiently high to induce toxicity problems in sensitive crop species.
Mo Above soil pH 4.2, MoO42− is dominant. Concentration of this species increases with increasing soil pH and plant uptake also increases. Water-soluble Mo increases sixfold as pH increases from 4.7 to 7.5. Replacement of adsorbed Mo by OH− is responsible for increases in water-soluble Mo as soil pH increases.
Zn Zinc solubility is highly soil pH dependent and decreases 100-fold for each unit increase in pH, and uptake by plants decreases as a consequence.
Ni Ni2+ is relatively stable over wide ranges of soil pH and redox conditions. However, availability is usually higher in acidic than in alkaline soils. At pH 7 and higher, retention and precipitation increase. Increasing the pH of serpentine soils through liming from 4 to 7 reduced Ni in plant tissue.
Co Solubility and availability of Co decrease with extreme soil pH. Presence of CaCO3, and high Fe, Mn, SOM, and moisture.

a Adriano (1986), Fageria, Baligar, and Jones (1997), and Tisdale et al. (1985).

Table V. Influence of Soil pH on Acquisition of Cu, Fe, Mn, and Zn by Upland Rice Grown in an Oxisol of Brazila

a Fageria (2000c). b P < 0.05. c P < 0.01.

Boron is the only micronutrient to exist in solution as a nonionized molecule over soil pH ranges suitable for the growth of most plants. Increasing soil pH decreases B availability by increasing B adsorption onto clay and Al and Fe hydroxyl surfaces, especially at high soil pH (Keren and Bingham, 1985). The highest availability of B was at pH 5.5–7.5, and the availability decreased below or above this pH range. In other studies, B adsorption increased from pH 3 to 8 on kaolinite, montmorillonite, and two arid zone soils with peak adsorption at pH 8–10 and decreases from pH 10 to 12 (Goldberg, Forster, Lesch et al., 1996). Reduced B availability occurs from liming (called “B fixation”)(Fleming, 1980) as CaCO3 acts as an adsorption surface. As such, B deficiency may occur in plants grown in limed acid soils.

Chloride is bound only lightly by most soil-exchange sites in acid to neutral soils and becomes negligible to pH 7.0. Chloride is easily leached from soil.

Considerable soil Cu is specifically adsorbed as pH increases. For example, increasing the pH from 4 to 7 increased Cu adsorption (Cavallaro and McBride, 1984), and Cu was adsorbed on inorganic soil components and occluded by soil hydroxide and oxides (Martens and Westermann, 1991). Increases in soil pH above 6.0 induces hydrolysis of hydrated Cu which can lead to stronger Cu adsorption to clay minerals and OM. Readily soluble sources of Cu (exchangeable or sorbed) were highly toxic to citrus, and Cu concentrations decreased considerably with soil pH increases above 6.5 (Alva et al., 2000). Over-liming acid soils may also lead to Cu deficiency. SOM is a primary constituent for Cu adsorption and readily complexes Cu. As the pH increases, the sizes of organic colloids of high molecular weight diminish, thus increasing the surfaces where Cu can be adsorbed (Geering and Hodgson, 1969).

The solubility of Fe decreases by ∼1000-fold for each unit increase of soil pH in the range of 4 to 9 compared to ∼100-fold decreases in the activity of Mn, Cu, and Zn (Lindsay, 1979). Iron exists in Fe0 (metallic), Fe2+ (ferrous), and Fe3+ (ferric) forms. Under acidic conditions, Fe0 readily oxidizes to Fe2+, and Fe2+ oxidizes to Fe3+ as the pH increases above 5. Ferric Fe (Fe3+) is reduced to Fe2+ and is readily available to plants in acidic soils, but precipitates in alkaline soils. Iron oxides are dominant in governing Fe solubility in soils. Minimum Fe solubility occurs between pH 7.5 and 8.5, which is the pH range of many calcareous soils (Lindsay, 1991). The increases in soil pH or Eh shift Fe from exchangeable organic forms to water-soluble and Fe oxide forms. The solubility of Fe in well-aerated soils is controlled by dissolution and precipitation of Fe3+ (Moraghan and Mascagni, 1991). Decreasing rhizosphere pH with added N (NH4–N) and/or K (KCl and/or K2SO4) was effective for increasing Fe uptake by plants (Barak and Chen, 1984). Applying FeSO4 with acid-forming fertilizer also increased Fe availability to plants (Moraghan and Mascagni, 1991).

Soil pH affects solubility, adsorption, desorption, oxidation of Mn, and reduction of Mn oxides in soil. As the pH decreases, Mn is mobilized from various fractions and increases Mn soil solution concentrations and availability. Exchangeable Mn (plant available form) was high at low soil pH (<5.2), while organic and Fe oxide fractions of Mn (low availability form) were high at high pH (Sims, 1986). In sandy soil, increasing pH also increased organic fractions of Mn (Shuman, 1991). Increasing soil pH with Mg applications on peanut decreased Mn toxicity and leaf and stem Mn concentrations (Davis, 1996). The reduction of Mn4+ to Mn2+ is greatest at low soil pH, and acid soil conditions (<5) lead to Mn toxicities for many sensitive plant species (Mortvedt, 2000). In addition, high-molecular-weight organic colloids diminish as soil pH increases to increase surfaces where Mn as well as Cu and Fe can be adsorbed (Geering and Hodgson, 1969). Soil solution Mn increased 1.6-fold for each unit decrease in pH in a well-drained Mollisol acidified with high N fertilizer, indicating that soil acidity and aeration are important for Mn availability (Fageria and Gheyi, 1999). Manganese, Cu, and Fe are generally more available under conditions of restricted drainage or in flooded soils (Ponnamperuma, 1972).

Molybdenum is the only micronutrient whose availability normally increases with increases in soil pH. The active form of Mo is normally MoO42−, which tends to polymerize when in solution. This condition is enhanced by acidification which could partially explain the low availability of Mo in some acid soils (Kabata-Pendias and Pendias, 1984). The solubility of CaMoO4 and H2MoO4 (molybdic acid) increases with increases in soil pH. Molybdenum sorption on Fe oxides increased with decreases in soil pH in the range of 7.8 to 4.5 (Hodgson, 1963). Adsorption of Mo on Al and Fe oxides was maximum at pH <5, and decreased as the pH increased >5 with little or no adsorption at pH 8 (Goldberg, Forster, and Godfrey, 1996). Soil pH had pronounced effects on Mo adsorption between 3 and 10.5 with virtually no adsorption at pH 8 (Goldberg and Foster, 1998). Adsorption of Mo on hydrous Fe and Al oxides decreased as soil pH increased, and the addition of lime to soil normally increased Mo solubility and a cquisition by plants (Williams and Thornton, 1972). In addition, maximum Mo adsorption on Al and Fe oxides was at pH 4–5, but adsorption was maximum at pH 3.5 with humic acid and decreased as soil pH increased (Biback and Borggaard, 1994). Different mechanisms were apparent for Mo adsorption with humic acid compared to Al/Fe oxides, which involved complex formation between carboxyl and phenolic groups. Harmful effects occasionally arise for legumes grown in acid soils, as Mo deficiency may be more dominant than Al toxicity (Bohn et al., 1979). In some cases, both lime and Mo applications may be needed to provide adequate Mo to plants (Lindsay, 1991).

Soil pH is more important than any other single property for controlling Zn mobility in soils (Anderson and Christensen, 1988). Increasing soil pH generally decreased Zn availability to plants (Saeed and Fox, 1977), and such decreases were usually due to higher adsorption of Zn. As soil pH increases above pH 5.5, Zn is adsorbed on hydrous oxides of Al, Fe, and Mn (Moraghan and Mascagni, 1991). However, the extent to which Zn is retained on Fe and Al hydrous oxides is influenced by the nature of clay minerals, surface conditions, and pH (Harter, 1991). In some cases, a soil pH higher than 7 may increase soil solution Zn due to solubilization of OM and also forms Zn(OH)+ and increased complexation of Zn with a lower positive charge (Barber, 1995). Gradual decreases in Zn activity as soil pH increases have been attributed to increased cation-exchange capacity (Stahl and James, 1991). Thirtyfold decreases in Zn concentration in acid soil have been reported for each unit increase in soil pH between 5 and 7 (McBride and Blasiak, 1979). Zinc was preferentially adsorbed over Cu on exchange sites indicating that chemisorption of hydrolyzed Zn occurs. Zinc adsorption is a major factor contributing to low concentrations of solution Zn in Zn-deficient soils. Soil pH affected Zn adsorption either by changing the number of sites available for adsorption or by changing the concentration of Zn species that is preferentially absorbed by plants (Barrow, 1986). Over-liming of soil may induce Zn deficiency and decrease Zn availability, especially at a high soil pH. Zinc absorption by wheat decreased as H+ concentrations increased, presumably because of the direct effects of H+ toxicity and the indirect effects of competition between Zn2+ and H+ for uptake sites on root surfaces (Chairidchai and Ritchie, 1993). The effect of pH may also be modified by organic ligands, and these ligands may decrease Zn uptake by plants as soil pH increases. Zinc deficiency may be expected in slightly acid and particularly in alkaline soils where inorganic Zn in equilibrium with soil Zn decreases between 10−8 and 10−10 M (Lindsay, 1991).

Chemisorption of Ni on oxides, noncyrstalline alumino silicates, and layer silicate clays is favored at soil pH > 6, but exchangeable and soluble Ni2+ is favored under lower pH conditions (McBride, 1994). The mobility of Ni is moderate in acid soils and becomes low in neutral and alkaline soils. Cobalt solubility decreases with increases in soil pH because of increased chemisorption on oxides and silicate clay, complexation by OM, and possible precipitation of Co(OH)2 (McBride, 1994). Cobalt is somewhat mobile in acid soils, but reduces as soil pH approaches neutrality.

Acid soils

Note Number: AG1182
Updated: April 2005

Why worry about acid soils?

Soil acidity is a natural and induced chemical condition of soils that can:

  • decrease the availability of essential nutrients;
  • increase the impact of toxic elements;
  • decrease plant production and water use;
  • affect essential soil biological functions like nitrogenfixation; and
  • make soil more vulnerable to soil structure decline and erosion.

The process of soil acidification is a potentially serious land degradation issue. Without treatment, soil acidification will have a major impact on agricultural productivity and sustainable farming systems and acidification can also extend into subsoil layers posing serious problems for plant root development and remedial action.

In some regions, there has been a drop of one pH unit over the last 20 to 30 years. Already, some farming areas have lost the ability to grow preferred agricultural species such as phalaris and lucerne simply because, without lime, the soil is too acid.

Understanding soil acidity

Figure 1: The Causes of Soil Acidity

Soil acidity occurs naturally in higher rainfall areas and can vary according to the landscape geology, clay mineralogy, soil texture and buffering capacity. Soil acidification is a natural process, accelerated by some agricultural practices (Figure 1).

When plant material is removed from the paddock, alkalinity is also removed. This increases soil acidity. When grain, pasture and animal products are harvested from a paddock, the soil is left more acid. Hay removal is particularly acidifying because large amounts of product are removed.

More significantly, soil acidification is most often a result of nitrate leaching. Nitrogen is added to the soil in a number of ways:

  • nitrogen fixed by legume-based plants;
  • as nitrogen based fertilisers;
  • from breakdown of organic matter; and
  • dung and urine.

Acidification occurs in agricultural soils as a result of the:

  • removal of plant and animal products;
  • leaching of excess nitrate;
  • addition of some nitrogen based fertilisers; and
  • build-up in mostly plant-based organic matter.

Figure 2: Relationship between pH measured in Calcium Chloride and Water

Soil pH is a measure of acidity or alkalinity. A pH of 7 is neutral, above 7 is alkaline and below 7 is acid. Because pH is measured on a logarithimc scale, a pH of 6 is 10 times more acid than a pH of 7. Soil pH can be measured either in water (pHw) por in calcium chloride (pHCa) and the pH will vary depending on the method used. As a general rule, pH measured in calcium chloride is 0.7 of a pH unit lower than pH measured in water (Figure 2). When a laboratory measures your soil’s pH it is important that they specify which method (water or calcium chloride) was used.

For most acid soils, the most practical management option is to add lime to maintain current soil pH status or increase surface soil pH.

The acid attack

Acidity itself is not responsible for restricting plant growth. The associated chemical changes in the soil can restrict the availability of essential plant nutrients (for example, phosphorus, molybdenum) and increase the availability of toxic elements (for example, aluminium, manganese). Essential plant nutrients can also be leached below the rooting zone. Biological processes favourable to plant growth may be affected adversely by acidity.

Bacterial populations generally prefer a slightly acid environment. However highly acidic soils can inhibit the survival of useful bacteria, for example the rhizobia bacteria that fix nitrogen for legumes. As the soil acidifies, the favorable environment for bacteria, earthworms and many other soil organisms is degraded. Acid soils have a major effect on plant productivity once the soil pHCa falls below 5:

  • pHCa 6.5 – optimum for most plant growth; neutral soil conditions; some trace elements may become unavailable.
  • pHCa 5.5 – balance of major nutrients and trace elements available.
  • pHCa 5.0 – aluminium may become soluble in the soil depending on soil type; phosphorus combines with aluminium and may be less available to plants.
  • pHCa 4.5 – manganese becomes soluble and toxic to plants in some soils; molybdenum is less available; soil bacterial activity slows down; aluminium becomes soluble in toxic quantities.
  • pHCa 4.0 – soil structural damage begins to occur.

Soil pH will influence both the availability of soil nutrients to plants and how the nutrients react with each other. At a low pH many elements become less available to plants, while others such as iron, aluminum and manganese become toxic to plants and in addition, aluminum, iron and phosphorus combine to form insoluble compounds. In contrast, at high pH levels calcium ties up phosphorus, making it unavailable to plants, and molybdenum becomes toxic in some soils. Boron may also be toxic at high pH levels in some soils.

The relative availability of 12 essential plant nutrients in well-drained mineral soils in temperate regions in relation to soil pH is shown in Figure 3. A pHCa range between 5 and 6 (between heavy lines) is considered ideal for most plants.

Understanding soil pH by testing

Soil pH is one of the most routinely measured soil parameters. It is used as a benchmark to interpret soil chemical processes and governs the availability of many essential or toxic elements for plant growth.

Soil pH is a common measure of the soil’s acidity or alkalinity because:

  • testing is relatively easy; and
  • field equipment to measure pH is relatively inexpensive.

Figure 3: Effect of pHca on the availability of plant elements.

Field test kits are available that use colour to indicate pH levels. The kits are inexpensive, easy to use and will test a lot of samples but should not be relied on for decisions such as rates of lime application. Test kits will only tell you whether your soil is acid or alkaline.

A number of compact testing meters that can be used out in the paddock are available, most of which are capable of giving accurate results if used correctly. Professional soil analysis is recommended and sending soil samples to a recognised laboratory ensures the most accurate results.

Testing of both topsoil and subsoil is recommended. When interpreting plant responses based on soil pH, the surface (A horizon) and sub-surface (B horizon) need to be considered.

The soil pHW is considered to be closer to the pH that the plant roots experience in the soil. But it is subject to large variation within the paddock because of seasonal changes in soil moisture and the ionic concentration of the soil solution that is related to the amount of total salts in the soil.

Research has shown that seasonal variation of pHW can vary up to 0.6 of a pH unit in any one year. In comparison, the measurements of soil pHCa is less affected by seasons.

Farmers can take soil samples at different times during the year without affecting the final diagnosis or interpretation.

Soil pHCa measurements in Australia vary from pHCa 3.6 to pHCa 8 for a range of different soil textures (sandy loams to heavy clays). Soil pHW values lie between pHW 4 and pHW 9.

Higher pHW values to around 10 may be associated with alkali mineral soils containing sodium carbonates and bicarbonates.

Useful tips

  • Soil pH is measured in either water or in calcium chloride. When measured in calcium chloride, the result is lower than pH measured in water.
  • The pHW may be higher by 0.6 to 1.2 in low salinity soils and higher by 0.1 to 0.5 in high salinity soils. Research has shown a difference of 0.7 for a wide range of soils.
  • Soil testing will tell you the current acidity status of your paddock. If your soil pHCa is above 5.5 then there is little immediate risk of acidity.
  • Lime can restore productivity in acid soils and should be considered once the pH drops below pHCa 5.0 if sensitive species are to be grown successfully.
  • You are unlikely to get responses to lime if other nutrients are lacking. This should show up in a soil test or plant tissue analysis and should be corrected. Conversely, you may not get a response to some nutrients if the soils are too acid. A holistic balanced approached is necessary.
  • Lime responses are generally seen in the first and second year for cropping systems, but can take up to five years depending on soil type, rainfall and lime quality for permanent pasture systems.
  • It is necessary to re-lime your paddock about every 10 years, depending on the rate of re-acidification.
  • If paddocks with an acidity problem are not limed, the soil pH will continue to fall and settle at pHCa3.8 to 4.2.
  • The amount of lime you need to apply varies according to soil type. Field experiments have shown that up to 5 tonnes a hectare on clay loams and 1.5 tonnes a hectare on sandy soils is needed to increase pH by one unit.
  • Lime moves slowly (0.5 to 1cm per year) through the soil profile via the soil macropore structure. Incorporation into the soil profile, where possible, will assist effective treatment.
  • In permanent pasture situations, spreading the lime on the surface and allowing it to work its way into the soil is acceptable. Surface application is better than no application.

The following Agnotes may assist landholders with field sampling procedures for soils:

Agnote AG0375: Sampling soils for growing pastures, field and fodder crops.

Agnote AG0376: How to sample soils used for flower, fruit, grape and vegetable production.

Agnote AG0889: Guidelines for sampling soils, fruits, vegetables and grains for residue testing.

Acknowledgements

This information note was developed by Carole Hollier and Michael Reed. Rutherglen. April 2005.

Appendix C. Acid And Basic Fertilizers

Fertilizers may alter the soil pH by either adding or removing acidity in the soil. The degree to which the pH changes is determined by cation exchange, in which an equivalent amount of one cation is exchanged for an equivalent amount of another.

Cation Equivalents

Cation exchange is an electrostatic phenomenon. It arises from an attraction between positively charged cations and negatively charged soil particles, or micelles. The soil particles may be either clay or organic matter. Exchange occurs in the displacement of cations of one species by those of another. Of the major cations, hydrogen and potassium have a charge or valence of +1, calcium and magnesium a valence of +2, and aluminum a valence of +3. A micelle has a high negative charge and attracts many cations. A calcium ion will neutralize two of those negative charges, potassium only one. Hence if the calcium drifts away from the micelle, its place can be occupied by two potassium ions or two hydrogen ions or one magnesium ion.

A mole is a unit of measure denoting a fixed number of ions, about 0.6 trillion trillion (known as Avogadro’s number). A mole of potassium ions has the same number of ions as a mole of magnesium ions. The conversion factor that relates the weight of a cation to the number of moles is the formula weight, which is the same as the atomic weight for chemical elements. The number of moles of a cation is equal to the actual weight divided by the formula weight. Calcium has a formula weight of 40, and so the conversion factor for calcium is 40 gm/mole. Thus, 1000 grams of calcium contains (1000 gms)/(40 gms/mole) = 25 moles.

A mole, however, is not a good unit of measure. In cation exchange, a mole of potassium ions will not displace a mole of calcium ions; since it has only half as much electrostatic charge of calcium, the potassium will only displace half a mole of calcium. A better unit of measure is an equivalent. One equivalent of potassium ions will displace one equivalent of calcium ions. The number of equivalents of a cation is the number of moles multiplied by the ionic valence. So 1000 grams of calcium is 25 moles, or 50 equivalents.

If a fertilizer has a liming effect, it is able to neutralize some of the acidity in the soil. We can determine the liming value by calculating the number of equivalents of hydrogen which are neutralized. If the fertilizer has an acidifying effect, we can calculate the number of equivalents of hydrogen which the fertilizer adds to the soil.

Liming Fertilizers

Limestone

Since calcitic limestone is the most common material for neutralizing acid soil, it will be the reference. The calculations are easier with negligible error by assuming that it is pure calcium carbonate.

In neutralizing acidity, the carbonate in limestone reacts with hydrogen:

CaCO3 + 2H+ → Ca++ + H2O + CO2

The calcium replaces the hydrogen at a micelle. Since one molecule of limestone reacts with two of hydrogen, and since a mole of limestone and a mole of hydrogen ions contains the same number (Avogadro’s number), it follows that one mole of limestone neutralizes two moles of hydrogen.

The formula weight of calcium carbonate is 100 grams/mole. One pound of limestone, or 454 grams, contains 4.54 moles. One pound of limestone will then neutralize 4.54 X 2, or about 9 equivalents of acidity.

Wood Ashes

The liming value of wood ashes is in its oxides and carbonates. Since this is only an estimate, we can assume that all of the calcium, magnesium and potassium exist as oxides.

One pound of wood ashes, or 454 grams, then contains 105 grams of calcium ions, 9.5 grams of magnesium ions, and 18.6 grams of potassium ions. The formula weights of calcium, magnesium and potassium are 40, 24 and 78 grams/mole, respectively, and their ionic valences are +2, +2, and +1, respectively. The liming equivalents of one pound of average ashes is then, approximately:

One pound of wood ashes neutralizes about 6 equivalents of acidity. Since one pound of limestone neutralizes 9 equivalents of acidity, a pound of ashes has the same liming value as about 2/3 pounds of limestone.

Sodium Nitrate

Experience shows that sodium nitrate has a liming effect, but the reason is not clear, because sodium nitrate is a neutral salt. Sodium ions have an alkaline effect, forming sodium hydroxide with water, and nitrate ions have an acidifying effect, forming nitric acid with water:

Na2NO3 + 2H2O → 2NaOH + H2NO3

One possible explanation for an imbalance is that the sodium hydroxide reacts with carbonic acid in the soil to form sodium carbonate:

2NaOH + H2CO3 → Na2CO3 + 2H2O

Sodium carbonate has a low solubility. What may be happening is that while the nitrate ions are taken up by plants or organisms or lost by leaching or denitrification, the associated hydrogen ions leach from the soil, and the sodium carbonate remains behind .

If this does occur, the sodium carbonate eventually dissolves and has a liming effect according to the reaction:

Na2CO3 + 2H++ → 2Na+ + Na2CO2 + H2O

Each mole of sodium ions replaces one mole of hydrogen ions. Since sodium nitrate contains one sodium ion, a mole of sodium nitrate should lead to the neutralization of one mole of hydrogen ions. Sodium nitrate has a formula weight of 85, and so a pound (454 gms) contains about 5.3 moles, and it can neutralize 5.3 equivalents of acidity. Since one pound of limestone neutralizes about 9 equivalents of hydrogen ions, one pound of sodium nitrate should have a liming capacity somewhat more than 1/2 pound of limestone.

According to the official method for determining the acidifying effect of fertilizers, the liming value of sodium nitrate is actually about 1/3 pound of limestone for each pound of sodium nitrate . In the absence of any other plausible hypothesis to explain the liming effect of sodium nitrate, this model gives a reasonable explanation, except perhaps that a significant fraction of the nitrate and hydrogen ions remain in the soil.

Bone Meal and Rock Phosphate

The claim sometimes made for the liming capability of rock phosphate is exaggerated. It is based on an artificial chemical reaction splitting tricalcium phosphate into two products:

Ca3(PO4)2 → P2O5 + 3CaO

There are several problems with this:

  • rock phosphate is not tricalcium phosphate but rather a complex mixture which includes either fluoride, chloride or hydroxyl; in order to account for these variations, the chemical formula may be be written as Ca5(PO4)3(F,Cl,OH).
  • tricalcium phosphate is stable and doesn’t split spontaneously
  • P2O5 is not part of any normally occurring reaction: it is prepared by igniting pure phosphorus, which doesn’t exist naturally by itself.

Nevertheless, in an effort to at least estimate an upper bound to the liming capability of rock phosphate, it may be worthwhile to assume that it is indeed tricalcium phosphate.

Tricalcium phosphate can react with water in three different ways to produce calcium oxide:

Ca3(PO4)2 + H2O → 2CaHPO4 + CaO

Ca3(PO4)2 + 2H2O → Ca(H2PO4)2 + 2CaO

Ca3(PO4)2 + 3H2O → 2H2PO4 + 3CaO

The first reaction produces dicalcium phosphate, whereby one of the calcium moles has a liming effect; the second produces monocalcium phosphate with two moles of calcium having a liming effect; and the third phosphoric acid with all three moles of calcium having a liming effect.

In practice, all three reactions should take place to some extent. But the degree to which they occur depends on the soil pH. Dicalcium phosphate predominates when the soil is alkaline, monocalcium phosphate when the soil is acidic, and phosphoric acid only when the soil is extremely acidic. The last condition is unlikely in agricultural soils, and we shall ignore the phosphoric acid option.

At a pH of 7, monocalcium and dicalcium phosphate exist in approximately equal amounts. If liming were necessary at a pH of 7, the net liming capability of one mole of the combination would be equivalent to about 1-1/2 moles of lime. So at a ph where lime is adviseable, the liming value of the mixture of phosphates should lie in the range between 1-1/2 and 2 moles of lime; maybe a reasonable value is the average, or 1-3/4 moles.

Since a mole of tricalcium phosphate contains three moles of calcium, the portion of calcium contributing to a liming effect is 1-3/4 / 3, or about 60%.

Colloidal rock phosphate has a stated CaO content of about 20%. From Appendix A. Conversion Factors , the actual calcium content is 20 / 1.4, or about 14%. Of this, 60%, or about 8% of the calcium, should be associated with a lime value. A pound of colloidal phosphate then contains 454 X 0.08, or about 36 grams of calcium having a liming value. Since the formula weight of calcium is 40 and its ionic valence 2, one pound of colloidal rock phosphate neutralizes 36 X 2 / 40 = 1.8 equivalents of acidity. Thus a pound of colloidal phosphate has a liming value of almost 1/5 pound of limestone. A normal application rate of 1 ton/acre will supply the equivalent of about 400 lbs of limestone/acre. As stated above, however, this is an upper limit; some of the calcium is associated with one or more of the variables in apatite – most likely fluoride, which is common in colloidal phosphate from Florida, a principal source in the U.S.

Hard rock phosphate may have a stated CaO content of about 35%. A similar calculation predicts a liming value – also an upper limit – of about 1/3 pound of limestone.

Bone meal does contain limestone. Theoretically, 1/4 of its calcium is in limestone and 3/4 in tricalcium phosphate. If bone meal has a specified CaO content of 28%, 7% should be in the form of limestone and 21% of tricalcium phosphate. With this division, one pound of bone meal has a liming value similar to that of hard rock phosphate, or about 1/3 pound of limestone.

Magnesia

Magnesia is magnesium oxide. We can determine its liming value using the same reasoning as we did for wood ashes, because in that example we assumed that wood ashes contain 3.5% magnesium oxide. Now consider a pound of magnesium oxide, or 454 grams. The number of equivalents in a pound is 454 gms X 2 equiv/mole / 24 gm/mole, or 38 equivalents. Then one pound of magnesia should have the same liming value as approximately 4 pounds of limestone.

Acidifying Fertilizers

Ammonium Sulfate

Chemically, ammonium sulfate is a neutral salt. In soil, however, the ammonium is oxidized by various bacteria to nitrite and thence to nitrate. The net reaction is:

(NH4)2SO4 + 4O2 → 4H+ + 2NO3- + SO4– + 2H2O

One mole of ammonium sulfate produces 4 moles of hydrogen ions.

The formula weight of ammonium sulfate is 132 gm/mole. One pound of ammonium sulfate contains 3.44 moles. Since each mole of ammonium sulfate produces 4 moles of hydrogen ions, one pound of ammonium sulfate will produce 3.44 X 4 = 13.8 equivalents of acidity. So each pound of ammonium sulfate theoretically requires about 1-1/2 pounds of limestone to neutralize its acidity.

According to the official method used to determine the acidifying effect of fertilizers, one pound of ammonium sulfate requires only 1.1 pounds of limestone to neutralize it . So something else is also occurring which remains unknown (at least to me; maybe the oxidation of nitrite is more complex than is assumed in the above chemical reaction).

Muriate of Potash (Potassium Chloride)

Theoretically, potassium chloride is a neutral salt; it should not affect the soil pH. Whether it actually does is a matter of opinion. The official method used to determine the acidity of fertilizers predicts that potassium chloride has no acidifying effect, but in the field an effect has been noticed.

Potassium chloride has a potential to acidify a soil wherever leaching is prominent. The potassium is absorbed by the plant or fixed in the soil, and the chloride leaches out, taking with it any available cations, principally calcium and magnesium. Their place at the cation exchange micelle is taken up by hydrogen and aluminum; the result is a drop in the soil pH. This activity is, in fact, the predominant way in which soils in humid areas become acid, so it is a reasonable model for estimating the acidifying effect of muriate of potash.

Potassium chloride has a formula weight of about 75 grams/mole. One pound of potassium chloride then contains about 6 moles, and each mole produces one equivalent of acidity. So a pound of muriate of potash has a potential acidity which can be neutralized by about 2/3 pounds of limestone.

Sulfur

Soil bacteria oxidize sulfur with a net reaction:

2S +3O2 + 2H2O → 4H+ +2SO4-

The result is a release of two hydrogen ions for each sulfur atom.

The formula weight of sulfur is 32. One pound contains 454 / 32, or about 14 moles, and produces 28 equivalents of acidity. The only reason for adding sulfur to the soil in substantial quantities is to neutralize excessive quantities of limestone. Sulfur is sometimes dusted on trees as a fungicide, but the amount usually used is too small to have a significant effect on soil acidity. One pound of sulfur will neutralize 28 / 9, or about 3 pounds of limestone.

Acid Rain

Although it may supply nitrogen and sulfur, acid rain is not usually considered as a fertilizer. Its influence on our environment, however, is of great interest, and estimating its effect on the soil should be worth the effort.

The calculation is based upon the assumption that rain is unbuffered. This means that there is no reservoir of acidity: it is all in solution. Consequently, the acidity is the hydrogen ion concentration in solution. It is related to the pH by the equation:

acidity = 10-pH moles/liter

A neutral solution having a pH of 7 will then contain 10-7, or 1/10,000,000 moles of hydrogen ions/liter. For our purpose we want to know the hydrogen ion concentration in terms of equivalents/gm of water. This can be determined, since we know that the density of water is 1 gm/ml, and one mole of hydrogen ions is equal to an equivalent of hydrogen. The result is that the hydrogen ion concentration is given by:

acidity = 10-pH-3 equivalents/gm

Consider one inch of rain. Its weight over an acre is:

W = 1 gm/cm3 * 2.54 cm/inch * 43560 ft2/acre * (30.51 cm/ft)2

or

W = 103,000,000 gm/acre per inch of rain

Consequently the acidity of 1 inch of rain falling on an acre is:

acidity = 103 * 106 * 10-pH-3

As a reference, the liming value of limestone is 9 equivalents/pound. The limestone requirement (L.R.) to neutralize 1 inch of acid rain is

L.R. = 103 * 103 * 10-pH

L.R. = 11000 * 10-pH

For example, the lime requirement to neutralize rain with a pH of 4 is about l lb/acre/inch of rain. A season of 100 inches of this rain will need about 100 lbs of limestone/acre to neutralize it. So much acid rain is rare, but even then the resulting acidity is not much when we are accustomed to spreading a ton of limestone every few years.

The prediction that acid rain has a negligible effect on soil is due to the assumption that rain is unbuffered. It should, however, be reasonable, because otherwise the rain would contain a high content of buffering chemicals, which is unlikely.

This is not meant to infer that acid rain is not influencing our environment. There is little doubt of its impact on unbuffered lakes and streams, and it appears to affect plant foliage by direct contact. But its influence on agricultural soils should be insignificant.

Superphosphate and Triple Phosphate

In principle the acidifying tendency of synthetic phosphates can be calculated. In practice the effort is not worthwhile. Phosphate fertilizers contain several forms of calcium phosphate, each with its own chemical behavior. An experimental determination seem more feasible. I did this with a single sample of superphosphate and triple phosphate, by measuring the amount of alkali needed to neutralize a water mixture. The results were as follows:

  • One pound of the superphosphate sample required about 1/5 pound of limestone to neutralize its acidity
  • One pound of the triple phosphate sample required about 5 pounds of limestone to neutralize its acidity.

Apparently the triple phosphate was treated with a large excess of acid in order to increase its solubility. Note that this test was done with only one sample of each fertilizer, and the results may be not valid for all samples.

Executive Summary

– Soil acidification is a natural process in high rainfall environments where leaching slowly acidifies soil over time.

– Intensive agriculture can speed up soil acidification through many processes – increasing leaching, addition of fertilizers, removal of produce and build-up of soil organic matter.

– Of all the major fertilizer nutrients, nitrogen is the main nutrient affecting soil pH, and soils can become more acidic or more alkaline depending on the type of nitrogen fertilizer used.

– Nitrate-based products are the least acidifying of the nitrogen fertilizers, while ammonium-based products have the greatest potential to acidify soil.

– Soil acidification due to use of phosphorus fertilizers is small compared to that attributed to nitrogen, due to the lower amounts of this nutrient used and the lower acidification per kg phosphorus. Phosphoric acid is the most acidifying phosphorus fertilizer.

– Potassium fertilizers have little or no effect on soil pH.

Background

Soil acidification is a widespread natural phenomenon in regions with medium to high rainfall, and agricultural production systems can accelerate soil acidification processes through perturbation of the natural cycles of nitrogen (N), phosphorus (P) and sulfur (S) in soil, through removal of agricultural produce from the land, and through addition of fertilizers and soil amendments that can either acidify soil or make it more alkaline (Kennedy 1986). Changes in soil pH may be advantageous or detrimental depending on the starting pH of the soil and the direction and speed of pH change – for example decreases in soil pH in alkaline soils may be advantageous for crop production due to benefits in terms of the availability of P and micronutrients e.g. zinc (Zn) (Mitchell et al. 1952). On the other hand, decreases in soil pH for a highly acidic soil may be detrimental in terms of increasing crop susceptibility to toxicity induced by increased solubility of aluminium (Al) or manganese (Mn) as soil pH falls (Wright 1989).

Key processes and reasons for changes in soil pH in agricultural systems are described below.

Fertilizer use

Use of mineral or organic fertilizers in agriculture increases inputs of nutrients to soils, and the form in which the nutrients are applied and their fate in the soil-plant system determine the overall effects on soil pH. Macronutrients (N, P, potassium (K), and S) have the major effects on pH as they are added in much larger quantities to soil than micronutrients.

The form of N and the fate of N in the soil-plant system is probably the major driver of changes in soil pH in agricultural systems.

The key molecules of N in terms of changes in soil pH are the uncharged urea molecule (0), the cation ammonium (NH₄+) and the anion nitrate (NO₃-). The conversion of N from one form to the other involves the generation or consumption of acidity, , and the uptake of urea, ammonium or nitrate by plants will also affect acidity of soil (Figure 1).

Figure 1. Soil acidity and nitrogen fertilizers (modified from (Davidson 1987)). MAP = monoammonium phosphate, DAP = diammonium phosphate, SoA = sulfate of ammonia, CAN = calcium ammonium nitrate, sodium nitrate

It can be seen in Figure 1 that ammonium-based fertilizers will acidify soil as they generate two H⁺ ions for each ammonium molecule nitrified to nitrate. The extent of acidification depends on whether the nitrate produced from ammonium is leached or is taken up by plants. If nitrate is taken up by plants the net acidification per molecule of ammonium is halved compared to the scenario when nitrate is leached. This is due to the consumption of one H⁺ ion (or excretion of OH⁻) for each molecule of nitrate taken up – this is often observed as pH increases in the rhizosphere (Smiley and Cook 1973). Anhydrous ammonia and urea have a lower acidification potential compared to ammonium-based products as one H⁺ ion is consumed in the conversion to ammonium. Nitrate-based fertilizers have no acidification potential and actually can increase soil pH as one H⁺ ion is absorbed by the plant (or OH⁻ excreted) in the uptake of nitrate.

The form of P fertilizer added to soil can affect soil acidity, principally through the release or gain of H⁺ ions by the phosphate molecule depending on soil pH (Figure 2). If phosphoric acid (PA) is added to soil, the molecule will always acidify soil as H⁺ ions will be released – one H⁺ ion if the soil pH is less than ~6.2 and two H⁺ ions is the soil pH is above 8.2. Monoammonium phosphate (MAP), single superphosphate (SSP) and triple superphosphate (TSP) all add P to soil in the form of the H₂PO₄⁻ ion, which can acidify soil with a pH greater than 7.2 but has no effect on soil pH in acidic soils. The form of P in diammonium phosphate (DAP) is HPO₄²⁻ which can make acidic soils (pH<7.2) more alkaline but has no effect on soil with a pH>7.2. The hydrolysis of ammonium polyphosphate (APP), where the P present as the P₂O₇⁴⁻molecule converts to HPO₄²⁻, is pH neutral and hence any acidification due to adding P can be regarded as similar to DAP. SSP or TSP are sometimes declared to cause soil acidification due to reaction products being very acidic;

Ca(H₂PO₄)₂+ ₂H₂O -> CaHPO₄ + H⁺ + H₂PO₄⁻

but in soils with pH values less than 7.7 the following reaction neutralizes the acidity produced so that there is no net acidification;

CaHPO₄ + H₂O -> Ca₂+ + H₂PO₄⁻ + OH⁻

In high pH soils (pH >7.2), dissociation of H+ ion from the H₂PO₄⁻ molecule will generate some acidity.

Crop uptake of P has little effect on soil acidity due to the small amounts of fertilizer P taken up in any one year – hence fertilizer chemistry dominates pH changes and significant differences in rhizosphere pH have not been observed for uptake of different orthophosphate ions.

Figure 2. Soil acidity and P fertilizers. MAP = monoammonium phosphate, DAP = diammonium phosphate,

SSP = single superphosphate, TSP = triple superphosphate, APP = ammonium polyphosphate.

The form of S fertilizer added to soil can affect soil acidity, principally through the release of H⁺ ions by the addition of elemental S (S⁰) or thiosulfate (S₂O3²⁻, in ammonium thiosulfate – ATS) (Figure 3). However, the amounts of S added to soil and taken up by plants are generally small in comparison to N.

Figure 3. Soil acidity and S fertilizers. S⁰ = elemental S, ATS = ammonium thiosulfate, SoA = sulfate of ammonia.

For each molecule of S⁰ added to soil, two H⁺ ions will be generated, and these can be balanced through plant uptake by either uptake of H⁺ (same as excretion of OH⁻ ions) or the generation of OH⁻ (effectively organic anions) within the plant to form alkaline plant material (“ash alkalinity”). Where produce is removed (which is often the case in agricultural systems) net acidification of soil will occur if S⁰ or ATS are used.

Potassium

The form in which K is added to soil – either muriate of potash (KCl) or sulfate of potash (K₂SO₄) – has no effect on soil acidification.

Acidification by Microessentials products

Information On How To Raise Acid Level In Soil

For gardeners growing an acid loving plant like blue hydrangea or azalea, learning how to make soil acidic is important to its overall health. If you don’t already live in an area where the soil is acidic, making soil acidic will involve adding products that lower the soil pH. Soil pH measures the alkalinity or acidity levels, which range from 0 to 14 on the pH scale. The middle (7) is considered neutral while levels falling below 7 are acidic and those above that number are alkaline. Let’s take a look at how to raise acid level in soil.

What Types of Plants Grow in Acidic Soil?

While most plants grow best in soils between 6 and 7.5, others are favorable to more acidic conditions. Some of the most common and sought-after plants actually prefer acidic soil, even though many of them may be grown in a wide range of growing conditions.

The acid-loving plants that you can grow in acidic soil include:

  • azaleas and rhododendrons
  • hydrangea
  • gardenias
  • camellias
  • wood anemone
  • bleeding heart
  • various carnivorous plants
  • holly shrubs
  • crepe myrtle
  • calla lilies
  • pine trees

Even blueberries thrive in this type of soil pH.

How Do I Make My Soil More Acidic?

If your plants aren’t growing in your soil conditions because of too much alkalinity, then it may be necessary to learn more about how to raise acid level in soil pH. Before making soil acidic, you should first perform a soil test, which your local County Extension Office can assist you with, if needed.

One of the easiest ways to make soil more acidic is to add sphagnum peat. This works especially well in small garden areas. Simply add an inch or two of peat to the topsoil in and around plants, or during planting.

For another quick fix, water plants several times with a solution of 2 tablespoons vinegar to a gallon of water. This is a great way to adjust pH in container plants.

Acidifying fertilizers can also be used to help raise acidity levels. Look for fertilizer containing ammonium nitrate, ammonium sulfate, or sulfur-coated urea. Both ammonium sulfate and sulfur-coated urea are good choices for making soil acidic, especially with azaleas. However, ammonium sulfate is strong and can easily burn plants if not used carefully. For this reason, you should always read and follow label instructions carefully.

In some instances, applying elemental sulfur (flowers of sulfur) is effective. However, sulfur is slow acting, taking several months. This is also most often used by large-scale growers rather than the home gardener. Granular sulfur is deemed safe and cost effective for smaller garden areas, with applications of no more than 2 pounds per 100 square feet.

Sometimes recommended as a method of lowering the pH enough to turn hydrangea blooms from pink to blue is iron sulfate. Iron sulfate acts more quickly (two to three weeks) but should not be used on a regular basis as heavy metals accumulate in the soil, becoming harmful to the plants.

Soil Acidity

SOIL ACIDITY

Background

  • pH 7 is neutral.
  • Soil with pH levels above 7 are alkaline; those of less than 7 are acidic.
  • The lower the pH, the more acidic is the soil.
  • Soils in humid regions tend to be acidic; those in semiarid and arid regions tend to be around neutral or alkaline.
  • Acidification is a natural process.
  • Most commercial nitrogen fertilizers are acid forming, but many manures are not.
  • Crops have different pH needs—probably related to nutrient availability or susceptibility to aluminum toxicity at low pH.
  • Organic acids on humus and aluminum on the CEC account for most of the acid in soils.

Management

  • Use limestone to raise the soil pH (if magnesium is also low, use a high-magnesium—or dolomitic— lime).
  • Mix lime thoroughly into the plow layer.
  • Spread lime well in advance of sensitive crops if at all possible.
  • If the lime requirement is high—some labs say greater than 2 tons; others say greater than 4 tons— consider splitting the application over two years.
  • Reducing soil pH (making soil more acid) for acidloving crops is done best with elemental sulfur (S).

Background

Many soils, especially in humid regions, were acidic before they were ever farmed. Leaching of bases from soils and the acids produced during organic matter decomposition combined to make these soils naturally acidic. As soils were brought into production and organic matter was decomposed (mineralized), more acids were formed. In addition, all the commonly used N fertilizers are acidic—needing from 4 to 7 pounds of agricultural limestone to neutralize the acid formed from each pound of N applied to soils.

Plants have evolved under specific environments, which in turn influence their needs as agricultural crops. For example, alfalfa originated in a semiarid region where soil pH was high; alfalfa requires a pH in the range of 6.5 to 6.8 or higher (see figure 20.1 for common soil pH levels). On the other hand, blueberries, which evolved under acidic conditions, require a low pH to provide needed iron (iron is more soluble at low pH). Other crops, such as peanuts, watermelons, and sweet potatoes, do best in moderately acid soils in the range of pH 5 to 6. Most other agricultural plants do best in the range of pH 6 to 7.5.

Several problems may cause poor growth of acidsensitive plants in low pH soils. The following are three common ones:

  • aluminum and manganese are more soluble and can be toxic to plants;
  • calcium, magnesium, potassium, phosphorus, or molybdenum (especially needed for nitrogen fixation by legumes) may be deficient; and
  • decomposition of soil organic matter is slowed and causes decreased mineralization of nitrogen.

The problems caused by soil acidity are usually less severe, and the optimum pH is lower, if the soil is well supplied with organic matter. Organic matter helps to make aluminum less toxic, and, of course, humus increases the soil’s CEC. Soil pH will not change as rapidly in soils that are high in organic matter. Soil acidification is a natural process that is accelerated by acids produced in soil by most nitrogen fertilizers. Soil organic matter slows down acidification and buffers the soil’s pH because it holds the acid hydrogen tightly. Therefore, more acid is needed to decrease the pH by a given amount when a lot of organic matter is present. Of course, the reverse is also true—more lime is needed to raise the pH of high-organic-matter soils by a given amount (see “Soil Acidity” box).

Limestone application helps create a more hospitable soil for acid-sensitive plants in many ways, such as the following:

  • by neutralizing acids;
  • by adding calcium in large quantities (because limestone is calcium carbonate, CaCO3);
  • by adding magnesium in large quantities if dolomitic limestone is used (containing carbonates of both calcium and magnesium);
  • by making molybdenum and phosphorus more available;
  • by helping to maintain added phosphorus in an available form;
  • by enhancing bacterial activity, including the rhizobia that fix nitrogen in legumes; and
  • by making aluminum and manganese less soluble.

Almost all the acid in acidic soils is held in reserve on the solids, with an extremely small amount active in the soil water. If all that we needed to neutralize was the acid in the soil water, a few handfuls of lime per acre would be enough to do the job, even in a very acid soil. However, tons of lime per acre are needed to raise the pH. The explanation for this is that almost all of the acid that must be neutralized in soils is reserve acidity associated with either organic matter or aluminum.

pH Management

Soil testing labs usually use the information you provide about your cropping intentions and integrate the three issues (see the discussion under “pH Management” of the three pieces of information needed) when recommending limestone application rates. Laws govern the quality of limestone sold in each state. Soil testing labs give recommendations based on the use of ground limestone that meets the minimum state standard.

Increasing the pH of acidic soils is usually accomplished by adding ground or crushed limestone. Three pieces of information are used to determine the amount of lime that’s needed:

  • What is the soil pH? Knowing this and the needs of the crops you are growing will tell you whether lime is needed and what target pH you are shooting for. If the soil pH is much lower than the pH needs of the crop, you need to use lime. But the pH value doesn’t tell you how much lime is needed.

  • What is the lime requirement needed to change the pH to the desired level? (The lime requirement is the amount of lime needed to neutralize the hydrogen, as well as the reactive aluminum, associated with organic matter.) A number of different tests used by soil testing laboratories estimate soil lime requirements. Most give the results in terms of tons per acre of agricultural grade limestone to reach the desired pH.
  • Is the limestone you use very different from the one assumed in the soil test report? The fineness and the amount of carbonate present govern the effectiveness of limestone—how much it will raise the soil’s pH. If the lime you will be using has an effective calcium carbonate equivalent that’s very different from the one used as the base in the report, the amount applied may need to be adjusted upward (if the lime is very coarse or has a high level of impurities) or downward (if the lime is very fine, is high in magnesium, and contains few impurities).

Soils with more clay and more organic matter need more lime to change their pH (see figure 20.2). Although organic matter buffers the soil against pH decreases, it also buffers against pH increases when you are trying to raise the pH with limestone. Most states recommend a soil pH of around 6.8 only for the most sensitive crops, such as alfalfa, and of about 6.2 to 6.5 for many of the clovers. As pointed out above, most of the commonly grown crops do well in the range of pH 6.0 to 7.5.

There are other liming materials in addition to limestone. One commonly used in some parts of the U.S. is wood ash. Ash from a modern airtight wood-burning stove may have a fairly high calcium carbonate content (80% or higher). However, ash that is mainly black— indicating incompletely burned wood—may have as little as 40% effective calcium carbonate equivalent. Lime sludge from wastewater treatment plants and fly ash sources may be available in some locations. Normally, minor sources like these are not locally available in sufficient quantities to put much of a dent in the lime needs of a region. Because they might carry unwanted contaminants to the farm, be sure that any new by-product liming sources are field tested and thoroughly evaluated for metals before you use them.

“Overliming” injury. Sometimes problems are created when soils are limed, especially when a very acidic soil has been quickly raised to high pH levels. Decreased crop growth because of “overliming” injury is usually associated with a lowered availability of phosphorus, potassium, or boron, although zinc, copper, and manganese deficiencies can be produced by liming acidic sandy soils. If there is a long history of the use of triazine herbicides, such as atrazine, liming may release these chemicals and kill sensitive crops.

Need to lower the soil’s pH? When growing plants that require a low pH, you may want to add acidity to the soil. This is probably only economically possible for blueberries and is most easily done with elemental sulfur (S), which is converted into an acid by soil microorganisms over a few months. For the examples in figure 20.2, the amount of S needed to drop the pH by one unit would be approximately 3/4 ton per acre for silty clay loams, 1/2 ton per acre for loams and silt loams, 600 pounds per acre for sandy loams, and 300 pounds per acre for sands. Sulfur should be applied the year before planting blueberries. Alum (aluminum sulfate) may also be used to acidify soils. About six times more alum than elemental sulfur is needed to achieve the same pH change. If your soil is calcareous—usually with a pH over 7.5 and naturally containing calcium carbonate—don’t even try to decrease the pH. Acidifying material will have no lasting effect on the pH because it will be fully neutralized by the soil’s lime.

Cation Exchange Capacity Management | Top | Remediation of Sodic (Alkali) and Saline Soils

#1 Soil Tester
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Soil is the foundation on which your garden grows. Its composition can help or hinder plant growth.

Soil is composed of clay, sand and silt and the ratio varies from area to area. The type of soil you have also determines its aeration, drainage and water holding capacity.

The four major components of soil are:

  • Rocks and minerals (the solid portion)
  • Organic matter, which includes microorganisms and plants (living or dead)
  • Water
  • Air — yes, your soil does contain air. The amount varies by type of soil. Sand is highly aerated with clay being the exact opposite.

Amounts of each of these components determine soil quality and whether or not your plants will grow and thrive. Ideal garden soil consists of 25% air, 25% water, 40% mineral matter and 10% organic matter. It is dark-colored, smells kind of sweet, compresses into a loose lump in your hand when moist and is chock-full of earthworms. As you already know, this almost never happens!

Get your gardens off to a great start and keep them productive with premium quality soil amendments. Need advice? Our Soils Blog provides the ideas, information and practical experience you need to get the job done right.

Improving Soil

You could say, building soil is one of the major concepts of organic gardening. Feed the soil and the soil feeds the plants. Healthy soil promotes strong vigorous plants that are more capable of resisting insect and disease problems. But what does your soil need? There’s only one way to know for sure: get your soil tested (see Guidelines for Choosing a Soil-Testing Laboratory – PDF). The results of your test will tell you the soil’s pH and what nutrients are in abundant or short supply.

IT’S ORGANIC!

A soil test can cost anywhere from $10 to $40 per sample and should be done every two to three years (contact your local extension agent for a soil test laboratory in your area). Don’t want the hassle? Purchase a simple, yet accurate soil test kit that uses a “color comparator” and capsule system for under $20.00. The Rapitest Soil Test Kit contains 40 tests (10 each for pH, nitrogen, phosphorus and potassium) and includes instructions for adjusting your soil conditions.

Adding Organic Matter

Mixing organic material, preferably compost and natural soil amendments, with your soil is of great importance to your success as a gardener. It improves soil structure, texture and aeration; helps maintain a neutral pH, adds needed nutrients for plant growth and allows the soil to hold more water which is a good thing if you live out here in the arid West.

Soil amendments should be considered as an investment in your garden soil which will pay off year after year. When these organic materials are added, they not only act in varying degrees as fertilizers, providing a mix of nutrients to plant roots, but will also build the structure of the soil by increasing its organic content. You can choose from a variety of soil amendments, like bone meal, greensand or Azomite, which are derived from natural sources and each suited to a particular need.

The pH of your soil is a measure of acidity or alkalinity. Soil pH is extremely important and directly affects nutrient availability. The pH scale ranges from 0 to 14, with 7 as neutral. Numbers less than 7 indicate acidity while numbers greater than 7 indicate alkalinity.

The pH of a soil is one of a number of environmental conditions that affects the quality of plant growth. Plants can only absorb nutrients that are in water soluble form and if the soil pH is too high, or too low, the needed elements and compounds may remain insoluble or incapable of being dissolved (see Changing the pH of Your Soil). Most garden vegetables, grasses and ornamentals do best in a slightly acidic soil with a pH between 5.8 and 6.8. Within this range, roots can absorb and process available nutrients. Azaleas, rhododendrons, blueberries and conifers prefer acidic soils (pH 5.0 to 5.5).

INCREASES pH

Increasing Soil pH. To correct acid soil, you add alkaline material (no kidding!). Doing this is called liming, probably because the most common solution is to add limestone. The finer the lime particles, the more rapidly it becomes effective. Different soils will require a different amount of lime to adjust the pH value. The texture of the soil, organic content and the plants to be grown are all factors to consider when adjusting the pH value. Click on the Limestone Calculator to determine how much lime you will need. Wood ashes will also raise soil pH. However, they break down very quickly, so use with caution. Over applying can cause serious soil imbalances.

Finely ground Oyster Shell Lime is an all-natural soil amendment composed of… well, oyster shells from the seafood industry. Use to raise pH in acidic soils and to correct calcium deficiencies (contains up to 39% calcium). Apply 2-4 Tbsp per plant or 50 lbs per 1,000 square feet, depending on soil analysis and crop grown. Breaks down quickly and does not pose the health risks associated with hydrated lime products.

Decreasing Soil pH. To correct alkaline soil, a source of acid is needed. Elemental sulfur is most commonly used by organic gardeners. However, sulfur requires some time before it is converted to sulfuric acid with the aid of soil bacteria. This conversion rate is dependent on the particle size of the sulfur, the amount of soil moisture, soil temperature and the presence of the bacteria. As a result, it can take several months to decrease the pH value. Click on the Sulfur Calculator to determine how much sulfur you will need. Do not attempt to change pH by more than 1 pH unit per year.

Approved for organic use, Yellowstone Brand® Elemental Sulfur or “split pea” sulfur lowers pH in alkaline soils and helps acid-loving plants to achieve optimum growth. Use 1 Tbsp per 4″ of pot diameter or broadcast approximately 10 lbs per 1,000 square feet and work into the soil. Do NOT exceed more than two applications per year.

Garden Myths – Learn the truth about gardening

There is a lot of advice on how to make make acidic soil both in print and on the net. You can use coffee grounds, pine needles, and sulfur to name a few. This advice has two problems. Firstly, the recommended product may not actually acidify soil. For example in Do Pine Needles Acidify Soil I show that pine needles do not make acidic soil. Coffee grounds don’t acidify soil either. The second problem is that before such advice is given it is important to know the soil types (ie soil texture) being treated. Let’s take a closer look at this.

Soil texture is important when trying to acidify soil

Soil Texture

Your soil has been made over millions of years using the rocks that were present at your location. It might have a lot of sand, or a lot of clay. It will also contain minerals based on the type of rock that was degraded to make your soil. The ability for any soil amendment to change soil acidity depends very much on the soil type you have. Let’s look at a couple of examples.

Soil that is very sandy usually does not contain much in the way of minerals. If you add a small amount of acidic material to the soil it will become acidic, at least for a short period of time. The problem with sand is that minerals and added acid leach away quickly; so the acidification of sand is a short term event – your soil will not stay acidic for long.

If your soil contains significant amounts of loam or clay, the soil could be naturally acidic or alkaline. It will contain minerals that will react with the added acid. Any acidification of the soil depends very much on the composition of these minerals. The minerals may be able to neutralize, or buffer the added acid. The importance of this buffering ability is discussed in Liming Acidic Soil. Soil testing is the only way to determine the pH buffer value.

Most of Southern Ontario is a clay loam. The base rock here is limestone and there is a lot of limestone, both as rocks and as minerals in the soil. These minerals are able to neutralize any acid that is added.

Consider this fact. Rain dissolves CO2 from the air as it falls to earth producing carbonic acid (this is not due to pollution). This rain, even without the added pollution has a pH of 5.5. This acidic rain has been falling in Ontario for millions of years and even after such a long time of ‘acidifying the soil’ our soil pH is still 7.4. How can this be? Our soil contains a lot of neutralizing minerals due to the limestone. As soon as an acidic material is added to the soil it is quickly neutralized so that it has no net effect on the soil pH.

In Northern Ontario and Quebec, the base rock is granite, not limestone. Granite is very stable and hardly reacts with acid. The soils in these areas are generally acidic and the addition of more acidic material will make the soil pH more acidic. In fact the pollution over the last 50 years has made the rain more acidic (ie pH lower than 5.5), and this has resulted in the soil in some areas becoming more acidic.

I have split Ontario into two parts, northern and southern, but in each area there are exceptions to the above statements. You can find acidic soil in the south and alkaline soil in the north.

Conclusion:

The acidification effect of any material on soil depends very much on the soil types you have. Simply saying that “material ABC” acidifies soil is not correct. It may acidify some soils and not others.

1) Photo Source: Mikenorton

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It’s no secret that soil is the most important part of the garden. Not only does it create healthy plants, or lead to their demise, but it is also full of information that can help us grow a better garden. This at-home soil pH test will give you a general idea of the pH of your soil. If you want to find out the exact pH level, you will need a test kit.

Materials:

  • Distilled Water (because it has a neutral pH. You can use regular water, but it could affect the outcome)
  • White vinegar (an acid)
  • Baking soda (a base or alkaline)
  • A bowl and spoon

Let’s Test Soil!

Scoop up a small amount of soil from an area in your garden.

Mix in a bit of water to the soil: enough to make a loose mud.

Pour a little bit of vinegar to the bowl. If it fizzes up, the soil is alkaline. As you can see, there was no fizz in my soil pH test, which would suggest that my soil is acidic.

To double-check the results, grab another scoop of soil, wet it with the water and mix again. Then sprinkle baking soda in it and mix. If it fizzes, the soil is acidic.

On the second soil pH test, my soil did fizz up, which means the soil is acidic.

You certainly do not have to perform both tests to determine the pH of your soil. Just one will suffice, but you can try both to confirm the results if you like. To be honest, I already knew that my soil is acidic, but in the name of garden science I had to confirm!

Now that you are armed with this basic knowledge about your soil, you can use it to do cool things like change the color of your hydrangea!

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