Alkali soils

Alkali, or alkaline, soils are clay soils with high pH (> 9), a poor soil structure and a low infiltration capacity. Often they have a hard calcareous layer at 0.5 to 1 metre depth. Alkali soils owe their unfavorable physico-chemical properties mainly to the dominating presence of sodium carbonate which causes the soil to swell. They derive their name from the alkali metal group of elements to which the sodium belongs and that can induce basicity. Sometimes these soils are also referred to as (alkaline) sodic soils.
Alkaline soils are basic, but not all basic soils are alkaline, see: "difference between alkali and base"

Contents 1 Causes 2 Occurrence 3 Agricultural problems 4 Chemistry 5 Solutions 6 Leaching saline sodic soils 7 References

Causes

The causes of soil alkalinity are natural or they can be man-made.

1. The natural cause is the presence of soil minerals producing sodium carbonate (Na₂CO₃) upon weathering.
2. The man-made cause is the application of irrigation water (surface or ground water) containing a relatively high proportion of sodium bicarbonates.

Occurrence

The extent of alkaline soils is not precisely known.^[1]

Agricultural problems

Alkaline soils are difficult to take into agricultural production. Due to the low infiltration capacity, rain water stagnates on the soil easily and, in dry periods, irrigation is hardly possible. Agriculture is limited to crops tolerant to surface waterlogging (e.g. rice, grasses) and the productivity is low.

Chemistry

Soil alkalinity is associated with the presence of sodium carbonates or (soda) (Na₂CO₃) in the soil,^[2] either as a result of natural weathering of the soil particles or brought in by irrigation and/or flood water.

The sodium carbonate, when dissolved in water, dissociates into 2Na⁺ (two sodium cations, i.e. ions with a positive electric charge) and CO3⁼ (a carbonate anion, i.e. an ion with a double negative electric charge).

The sodium carbonate can react with water to produce carbon dioxide (CO₂), escaping as a gas, and sodium hydroxide (Na⁺OH^–), which is alkaline (or rather basic) and gives high pH values (pH>9).^[2]

Notes:

• Water (H₂O) is partly dissociated into H⁺ (hydrogen) and OH^– (hydroxyl) ions. The ion H⁺ has a positive electric charge (+) and the hydroxyl group OH^– has a negative charge (–). In pure, neutral water, the concentration of H⁺ and OH^– ions equals 10^–7 eq/l each (respectively 10^–7 g/l and 17x10^–7 g/l), a very small concentration.
• 1 eq = 1 equivalent weight corresponds to as many grams of the substance as its molecular weight divided by its valence. It is also known as gram-molecule or mole per unit of valence. For the hydrogen ion (H⁺ ) the molecular weight equals 1, and for the hydroxyl group (OH^–) it equals 17. As they are both monovalent (or univalent) their equivalent weights are the same. Substances with the same equivalent weight have equal positive or negative electric charge.
• In neutral water, the pH, being the negative log value of the H⁺ concentration in eq/l, is 7. Similarly, the pOH is also 7. Each unit decrease in pH indicates a tenfold increase of the H⁺ concentration. Similarly, each unit increase in pH indicates a tenfold increase of the OH^– concentration.
• In water with dissolved salts, the concentrations of the H⁺ y OH^– ions may change, but their sum remains constant, namely 7 + 7 = 14. A pH of 7 therefore corresponds to a pOH of 7, and a pH of 9 with a pOH of 5.
• Formally it deserves preference to express the ion concentrations in terms of chemical activity, but this does hardly affect the value of pH.
• Water with excess H⁺ ions is called acid (pH < 6), and water with excess OH^– ions is called alkaline or rather basic (pH > 8). Soil moisture with pH < 4 is called very acid and with pH > 10 very alkaline (basic).

The reaction between Na₂CO₃ and H₂O can be represented as follows:

• 2Na⁺ + CO₃⁼ + 2H⁺ + 2OH^– => 2Na⁺ + 2OH^– + H₂CO₃

The acid H₂CO₃ is unstable and produces H₂O (water) and CO₂ (carbon dioxide gas, escaping into the atmosphere). This explains the remaining alkalinity (or rather basicity) in the form of soluble sodium hydroxide and the high pH or low pOH.

Not all sodium carbonate follows the above chemical reaction. The remaining sodium carbonate, and hence the presence of CO₃⁼ ions, causes CaCO₃ (which is only slightly soluble) to precipitate as solid calcium carbonate (limestone). Hence, the calcium ions Ca⁺⁺ are immobilized.

The presence of abundant Na⁺ ions in the soil solution and the precipitation of Ca⁺⁺ ions as a solid mineral causes the clay particles, which have negative electric charges along their surfaces, to adsorb more Na⁺ in the diffuse adsorption zone (DAZ, see figure, officially called diffuse double layer ^[3] ) and, in exchange, release Ca⁺⁺, by which their exchangeable sodium percentage (ESP) is increased as illustrated in the figure.

Na⁺ is more mobile and has a smaller electric charge than Ca⁺⁺ so that the thickness of the DAZ increases as more sodium is present. The thickness is also influenced by the total concentration of ions in the soil moisture in the sense that higher concentrations cause the DAZ zone to shrink.

Clay particles with considerable ESP (> 16), in contact with non-saline soil moisture have an expanded DAZ zone and the soil swells (dispersion).^[3] The phenomenon results in deterioration of the soil structure, and especially crust formation and compaction of the top layer. Hence the infiltration capacity of the soil and the water availability in the soil is reduced, whereas the surface-water-logging or runoff is increased. Seedling emergence and crop production are badly affected.

Note:

• Under saline conditions, the many ions in the soil solution counteract the swelling of the soil, so that saline soils usually do not have unfavorable physical properties. Alkaline soils, in principle, are not not saline since the alkalinity problem is worse as the salinity is less.

Alkalinity problems are more pronounced in clay soils than in loamy, silty or sandy soils. The clay soils containing montmorillonite or smectite (swelling clays) are more subject to alkalinity problems than illite or kaolinite clay soils. The reason is that the former types of clay have larger specific surface areas (i.e. the surface area of the soil particles divided by their volume) and higher cation exchange capacity (CEC).

Note:

• Certain clay minerals with almost 100% ESP (i.e. almost fully sodium saturated) are called bentonite, which is used in civil engineering to place impermeable curtains in the soil, e.g. below dams, to prevent seepage of water.

Solutions

Alkaline soils with solid CaCO₃ can be reclaimed with grass cultures, ensuring the incorporation of much acidifying organic material into the soil, and leaching of the excess sodium.^[4] Deep plowing and incorporating the calcareous subsoil into the top soil also helps.

In some agricultural areas, the use of subsurface "tile lines" are used to facilitate drainage and leach salts.

It is also possible to reclaim alkaline soils by adding acidifying minerals like pyrite.

Alternatively, gypsum (calcium sulfate, CaSO₄. 2H₂O) can also be applied as a source of Ca⁺⁺ ions to replace the sodium at the exchange complex.^[4] There must be enough natural drainage to the underground, or else an artificial subsurface drainage system must be present, to permit leaching of the excess sodium by percolation of rain and/or irrigation water through the soil profile.

To reclaim the soils completely one needs prohibitively high doses of amendments. Most efforts are therefore directed to improving the top layer only (say the first 10 cm of the soils), as the top layer is most sensitive to deterioration of the soil structure.^[4] The treatments, however, need to be repeated in a few (say 5) years time.

It will be important to refrain from irrigation with poor quality water.

The quality of the irrigation water in relation to the alkalinity hazard is expressed by the following two indexes:

1) The sodium adsorption ratio (SAR,^[2] )

The formula for calculating sodium adsorption ratio is:

                    [Na⁺]                         {Na⁺/23}
SAR = ───────────── = ──────────────
           √[Ca⁺⁺/2 + Mg⁺⁺/2]     √{Ca⁺⁺/40 + Mg⁺⁺/24}

where: [ ] stands for concentration in milliequivalents/liter (briefly meq/l), and { } stands for concentration in mg/l.

It is seen that Mg (Magnesium) is thought to play a similar role as Ca (Calcium).

The SAR should not be much higher than 20 and preferably less than 10;

When the soil has been exposed to water with a certain SAR value for some time, the ESP value tends to become about equal to the SAR value.

2) The residual sodium carbonate (RSC, meq/l,^[2]):

The formula for calculating residual sodium carbonate is:

RSC = [HCO₃^– + CO₃⁼] ‑ [Ca⁺⁺+ Mg⁺⁺]

        = {HCO₃^–/61 + CO₃⁼/30} ‑ {Ca⁺⁺/20 + Mg⁺⁺/12}

which must not be much higher than 1 and preferably less than 0.5.

The above expression recognizes the presence of bicarbonates (HCO₃^–), the form in which most carbonates are dissolved.

Leaching saline sodic soils

Saline soils are mostly also sodic (the predominant salt is sodium chloride), but they do not have a very high pH nor a poor infiltration rate. Upon leaching they are usually not converted into a (sodic) alkali soil as the Na⁺ ions are easily removed. Therefore, saline (sodic) soils mostly do not need gypsum applications for their reclamation.^[5]

References