The balance of hydrogen ions and basic ions determines soil acidity.
A clay particle with abundant hydrogen ions acts as a weak acid,
whereas if fully charged with bases (such as calcium, Ca) it has a
neutral or alkaline reaction. In practice, soil pH is usually regulated by
the presence of calcium cations; soils become more acid as calcium
is leached from the soil
faster than it is replaced. This is the tendency
in temperate areas where rainfall (carbonic acid see) exceeds
evaporation over the year.
Hydrogen ions take over the soil’s cation exchange sites
and the pH falls. Soils with large reserves of calcium (containing
pieces of chalk or limestone) do not become acid because they are
kept base-saturated. In contrast, calcium ions are readily leached from
free-draining sands in high rainfall areas and these soils tend to go acid
rapidly (see podsols). In addition to the carbonic acid in rainfall,
there are several other sources of acid that affect the soil:
- Acid rain (polluted rain and snow) is directly harmful to vegetation,
but also contributes to the fall in soil pH.
- Organic acids derived from the microbial breakdown of organic
matter, e.g. humic acids, also lead to an increase in soil acidity.
- Fertilizers. The bacterial nitrification of ammonia to nitrate yields
acid hydrogen ions. Consequently fertilizers containing ammonium
salts prevent calcium from attaching to soil colloids and cause
calcium loss in the drainage water. Other fertilizers have much less
- Crop removal. Calcium and magnesium are plant nutrients and the
soil’s lime reserves are therefore gradually reduced by the removal of
In climates where the evaporation exceeds rainfall over the year, the
dissolved salts are brought to the surface. As the water is lost from the
soil by evaporation, the dissolved salts accumulate on the surface. These
are usually basic (alkaline) in action so the soil pH rises. An extreme
example of this is the salt (sodic) desert, e.g. Utah Salt Flats.
is the ability of water to maintain a stable pH.
Pure water has no buffering capacity; the addition of minute quantities
of acid or alkali has an immediate effect on its pH. In the laboratory buffer tablets can be added to water to enable the solution to be
maintained at a specified pH which would resist change despite the
addition of some acid or alkali. This is useful for standardizing a pH
meter, usually setting the instrument at precisely pH 7 and pH 4 for
work on soil pH.
The buffering capacity of soil water reduces the effect of acidity coming
from rainfall or from pollution, e.g. acid rain. Chalky or limestone soils,
for instance, are very alkaline and can neutralize acids more effectively
than acid peat soils. The cation exchange capacity of clays reduces the
effect because the hydrogen ions exchange with calcium ions on the clay’s
colloid surface. Since the number of hydrogen ions being released or
absorbed is small compared with the clay’s reserve, the pH changes very
little. High humus soils similarly have the advantage of a high buffering
capacity. A related buffer effect is seen when acids, such as the carbonic
acid of rain, are incorporated into soils with ‘free’ lime present; the acid
dissolves some of the carbonate with no accompanying change in pH.