Soil Acidity

Negative charges expressed on surfaces of clay crystals are due to isomorphous substitution. These charges are permanent, do not vary with change in pH, but do vary depending on the type of clay (recall the CEC exercise). Negarive charges also result from dissociation of hydrogen from hydroxyl SiOH and AlOH groups (broken crystal edges). Negative charges expressed on organic colloids surfaces result from dissociation of the hydrogen from carboxyl, phenolic hydroxy, and alcohol hydroxy groups. The tendency for dissociation increases with increasing pH. Cations, including H and Al ions are adsorbed to colloid surfaces.

Hydrolysis of Al generates hydrogen ion, playing a key role in soil solution hydrogen ion concentration (pH, also known as active acidity). The mechanism of hydrogen release varies with pH. Under very acid soil conditions, most aluminum is soluble in the form of Al+++ or Al(OH)++ cations, which become strongly adsorbed in preference to hydrogen. The Al+++ in solution, which is in equilibrium with exchangeable Al, tends to hydrolyze so that Al+++ and water combine to give Al(OH)++ and one hydrogen ion. Adsorbed hydrogen is a second (minor) source of hydrogen since the adsorbed hydrogen ions are in equilibrium with those in soil solution. In moderately acid soils, bases are more abundant, and so Al+++ can no longer can no longer exist as an ion, but is converted to Al(OH)++. Since hydroxyl ions are more abundant in this condition of soil acidity, more complex Al(OH)++ ions are formed, reducing the probability of their hydrolyzing to form hydrogen ions. However, those that remain in the soil solution do hydrolyze to form hydrogen ions, but their numbers are less than in the very strongly acid soil.

In neutral or alkaline soils, the solution or exchange sites are no longer dominated by either hydrogen or aluminum ions. The permanent exchange sites are now occupied primarily by non acid cations (a.k.a. exchangeable bases}. Under these conditions the aluminum hydroxyls are converted to Al(OH)3 - gibbsite - which is insoluble. This reaction is brought about by the abundance of hydroxyl ions. Also, the hydrogen ions of exchange sites have been largely replaced by Ca++, Mg++, and other non acid cations The source of the hydroxyls is the replacement of the hydrogen and aluminum by base cations. The adsorption of the non acid cations releases hydroxyls to the soil solution. In addition, when the clay micelle contains non acid cations, dissociated water reacts so that the hydrogen ions exchange with non acid cations, and the hydroxyl portion of the water is now in solution contributing to the hydroxyl concentration. Acidity of the soil is related to the relative amounts of hydrogen ions and Al+++ adsorbed compared to the adsorbed metallic cations.

Now that we have established the source of hydrogen ions in the soil, consider the heterogeneity of distribution of these ions in the soil. Refer now to Figure 10.1 in the lab manual, which shows a swarm of hydrogen ions adsorbed on the clay micelle to the left. As the distance from the micelle surface increases, the number of hydrogen ions decreases until the limit of colloidal attraction is reached. Hydrogen ions that occur beyond this limit, free to react in solution, are referred to as the active acidity. Those hydrogen ions associated with the micelle either as exchangeable ions (salt replaceable) or as tightly adsorbed Al hydroxyl ions (residual acidity) are referred to as the reserve acidity. The diagram shows only hydrogen ions, but you should keep in mind that many other ions are present in the soil solution and adsorbed on colloid surfaces.

The equilibrium between hydrogen ions in solution and those adsorbed on the clay micelle was mentioned earlier. However, the numbers are quite unequal. The large numbers of adsorbed hydrogen ions is due to a large number of hydrogen ions, with the outer ones moving into solution, and those in solution closest to the micelle being strongly adsorbed. Hydrogen ions are continually changing places due to their motion about the charge source. The number of hydrogen ions in solution is always smaller than the number of exchangeable hydrogen ions. Because of this equilibrium, exchangeable hydrogen will move into solution to replace those hydrogen ions which are neutralized by some reaction. On the other hand, if hydrogen ions are added to the solution, a certain number of the ions are adsorbed on the micelle, thus maintaining the equilibrium between the active and the reserve acidities. The tendency of the soil to maintain this equilibrium (i.e. resist change in pH) is termed buffering.

Hydrogen ions are added to the soil by the formation of organic and inorganic acids when percolating water moves through the organic and surface mineral horizons. Hydrogen ions are removed from the soil by leaching or by the addition of bases which neutralize the hydrogen ions. Bases are added by weathering of parent material and by fertilization. The greater the cation exchange capacity of the soil, the greater its buffering capacity. The factors which determine buffering capacity include: organic matter content, texture or clay content, clay type, and acid saturation.

Buffering of a solution follows a characteristically shaped curve when a base is titrated against an acid, or vice versa. In lecture, we will discuss a paper written in 1985 by Rich Bartlett and Bruce James, published in the Soil Science Society of America Journal [49: 145-148]. They point out that the the relationship between pH and base saturation has been misinterpreted as a buffering curve. Furthermore, that relationship is best described as a straight line. The original graph of that relationship is most commonly cited as an s-shaped curve. This s-shape is a consequence of including calcareous soils in which a CaCO3-CO2-H2O system dominates the pH. "Much more CaCO3 than needed to achieve 100% base saturation must be added to produce a truly calcareous soil".

Analysis of the relationship between soil pH expressed as a function of ammendment of acid/base per gram of organic matter done for 51 soils collected from Vermont produced an s-shaped buffer curve. However, that curve showed that the buffering capacity of Vermont soils was least in the pH range of 5.0 to 5.6, and greatest for soil pH greater than 7 and less than 4. This constrasts directly with the previous interpretation of the older version of the base saturation curve.

The importance of preventing wide fluctuations in soil pH is pointed out in both the text and the exercise. Since pH influences the availability of nutrients by affecting the type of compound formed, which predetermines solubility, it is obvious that any drastic change in pH would create some problems. Also, soil organism growth and activity are significantly influenced by soil pH; drastic changes in pH would have serious consequences on organism activity and availability of nutrients derived from organic matter decomposition.

The pH kit contains several colorimetric pH indicators, porcelain or plastic spot plates to hold soil samples, some distilled water for rinsing the plates, two color charts, directions, and a small handbook. Accuracy (plus or minus 0.2 pH units) is adequate for most field applications. The soil sample must be representative of the soil conditions. If you are interested in obtaining a single pH value for a small area such as your front yard, it is best to obtain a soil sample from several randomly located points within the area. At each point use a punch tube to sample the upper 20 cm at several points. Thoroughly mix the composited sample and select a subsample for analysis.

Remember that pH is variable over space and time. For detailed soil characterization, we would obtain samples of each horizon of the profile and composite them by horizon. The pH will vary widely from horizon to horizon and we can compare the values of one horizon only with a similar horizon from another profile.

Scoop out a small portion of soil and place it in the large depression of the white plate. The plate should be cleaned with distilled water, then dried with a clean tissue. The depression should be filled level, the excess being carefully scraped awayFill a second depression for a duplicate test.

Some kits may have Duplex indicator which is used to bracket the pH value, followed by a more precise determination with a specific indicator. Carefully add indicator drop by drop until the soil appears to be nearly saturated, or when the indicator just barely begins to glisten. Four or five drops will suffice. When you have added sufficient indicator, tilt the plate slightly towards the small depression, and, using the plastic rod, coax the solution across the canal to the small depression. Be careful not to muddy the solution by stirring the soil.

The Duplex indicator is used to determine which specific dye should be used to measure the pH more precisely Compare the colors under a bright light. In the picture pH is between 5 and 6. In this example we are using Chlorophenol Red. The pH appears to be about 6.4 . These charts give the pH to the nearest 0.2 of a pH unit. Though the whole procedure appears to be quite crude, good results can be obtained by this method.

The theory behind the electronic pH meter. glass electrode is explained for you in the laboratory exercise. It is obvious that we can measure soil pH accurately and very precisely with this instrument, but again the value obtained is only as good as the the sample. The soil sample is prepared by placing an air-dry sample in a 50 ml beaker and adding distilled water to obtain a 1:1 ratio of soil to water. After stirring the wetted sample and allowing it to stand for at least one-half hour and the electrode is carefully immersed into the beaker so that the end of the electrode just touches the top of the soil within the supernatant liquid, but with the electrode completely immersed.

When the electrode is properly immersed, press the measure switch and record the value when the reading stabalizes. Do not use the electrode as a stirring rod. After the electrode has been removed from the sample, rinse with distilled water catching the water in a beaker. If another reading is not made immediately, the electrode should be immersed in clean distilled water.

pH is so easy to measure and it is correlated with nutrient availability as well as biologic activity.