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We should first understand the source of the negative charges which attract the hydrogen
ions of the soil. First and most obvious are the negative charges within the clay crystals
due to isomorphic 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). Other negative
charges occur as a result of dissociation of the hydrogen in various hydroxyl groups which
include SiOH and AlOH groups as well as COOH (carboxyl) and phenyl-OH groups. The tendency
for dissociation increases with an increase in pH. This leaves a negative charge, and an
exchange takes place.
The mode of hydrogen release is different at varying pH levels; it is necessary to discuss
them separately. 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 ions, but are 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 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 bases. The source of the hydroxyls
is the replacement of the hydrogen and aluminum by base cations. The adsorption of the
bases releases hydroxyls to the soil solution. In addition, when the clay micelle contains
bases, dissociated water reacts so that the hydrogen ions exchange with base 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 on page
1 of 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 the 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 there 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 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 whichdetermine buffering capacity include: organic matter content, texture or clay
content, texture or clay content and type of clay, and percentage of base 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 (see Figure 9.8 in Brady and Weil). 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 evidence is unclear concerning the effect of drastic pH changes on the
higher plants.
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 is somewhat
limited but is adequate for most field applications of the procedure is carefully
followed. Sloppy technique produces sloppy results. 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 dig down about 7 or 8 inches and place a
vertical slice on a clean plastic sheeting. After all the points have been dug, thoroughly
mix the composited sample on the plastic sheet. Then quarter it and remove a small portion
from each quarter. You should have a pint of soil in the quartered sample. If you are
going to sample a large field it is best to subdivide the field and take several composite
samples. This way you will obtain a measure of the variation of pH over the field and
perhaps the parts will need to be treated differently.
Remember that pH is highly variable over space and time. The results of one sample are
usually not adequate for a measure of pH over a large area. The farmer is generally
interested in pH for only the surface 8 to 10 inches as this is the depth of most of the
feeding roots of field crops. 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.
After collecting and compositing a soil sample 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. he depression
should be filled level, the excess being carefully scraped away. A plastic rod may be used
to get the soil in place. Do not use a knife blade or similar object as the steel may
contaminate the sample, or do not use your thumb of forefinger to tamp the soil. Fill 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. For our
purposes we will omit the Duplex step and most of you may begin directly with the
Bromcresol Green indicator. Take the dropper bottle and 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. If you accidentally flood the sample,
start over again because youll only detect the color of the indicator if you proceed any
further. 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.
{Next slide) Compare the colors under a bright light. In the
picture the value is between 5 and 6. (Next slide) The sample
procedure is used to determine the pH with the more precise indicators. Use the
appropriate color chart and compare the colors under a bright light. 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 instrument consists of the glass electrode on the right,
operating switches to the front, with a digital readout for the measurement value. 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, the sample is stirred again, 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. The switches
of the meter may differ from the one you use today, but the operation is very similar.
When the electrode is properly immersed, simply turn the measure switch to read or measure
the pH and note the readout of the meter. Do not use the electrode as a stirring rod.
Switch back to standby or zero position and remove the electrode from the soil sample. Do
not remove the electrode from the sample until the switch has been turned from the measure
position. Damage to the electrode or the meter may result, so please observe this
precaution. After the electrode has been removed from the
sample it is rinsed 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.
Earlier I mentioned the spatial and temporal variation in pH in the field. Because of this
variation, a single or even several pH values only give general indications of soil
conditions. In the past, and to a certain extent today, an inordinate amount of confidence
has been placed on pH for site or community classification. This arises because pH is so
easy to measure, and, it has been well established that there are definite correlations
between pH and other soil conditions. But, pH remains so variable and dependent upon
factors not measurable on site that its use should be in conjunction with many other soil
variables.