pH and Bicarbonate Buffering
The pH of a sample of water is a measure
of the concentration of hydrogen ions. The term pH is derived from the
French puissance d’hydrogene, meaning “strength of hydrogen,” because
the hydrogen ion controls acidity (Horne and Goldman, 1994). It is
calculated on a exponential scale of 1 to 14. Mathematically, pH is
the negative logarithm of the
hydrogen ion (H+) concentration. What this means to those of us who
are not mathematicians is that at higher pH,
there are fewer free hydrogen ions, and that a change of one pH unit
reflects a tenfold change in the concentration
of the hydrogen ion. For example, there are 10 times as many hydrogen
ions available at a pH of 7 than at a pH of 8. The pH scale ranges from
0 to 14. A pH of 7 is considered to be neutral. Substances with pH of
less that 7 are acidic; substances with pH greater than 7 are basic.
The pH of water determines the solubility (amount that can be dissolved
in the water) and biological availability (amount that can be utilized
by aquatic life) of chemical constituents such as nutrients (phosphorus,
nitrogen, and carbon) and heavy metals (e.g., lead, copper, cadmium).
For example, in addition to affecting how much and what form of phosphorus
is most abundant in the water, pH may also determine whether aquatic
organisms can make use of the phosphorus. In the case of heavy metals,
the degree to which they are soluble determines their toxicity. Metals
lower pH and thus more toxic.
Reasons for Natural Variation
Photosynthesis uses up dissolved carbon dioxide. CO2 removal,
in effect, reduces the acidity of the water and so pH increases. In contrast,
of organic matter
produces CO2, which dissolves in water as carbonic acid, thereby
lowering the pH. For this reason, pH may be higher during daylight hours
the growing season, when photosynthesis is at a maximum. Respiration
and decomposition processes lower pH. Like dissolved oxygen concentrations,
pH may change with depth in a lake, due again to changes in photosynthesis
and other chemical reactions. There is typically a seasonal decrease
in pH in the lower layers of a stratified lake because CO2 accumulates.
There is no light for plants to fix CO2 and decomposition
River water is complex; it is full of chemical "shock absorbers" that
prevent major changes in pH. Small or localized changes in pH are quickly
modified by various chemical reactions, so little or no change may be
measured. This ability to resist change in pH is called buffering capacity.
Not only does the buffering capacity control would-be localized changes
in pH, but it controls the overall range of pH change under natural conditions.
The pH scale may go from 0 to 14, but the pH of natural waters hovers
between 6.5 and 8.5.
When carbon dioxide dissolves in water a small fraction is hydrated
to form carbonic acid.
Because it is acidic, carbonic acid dissolves calcium carbonate
rocks, neutralizing the soil and stream water, and forming calcium bicarbonate.
As a result
the calcium bicarbonate content of streams largely determines its pH
(balance between hydrogen and hydroxide ions). Buffering is then due
presence of carbon dioxide, carbonic acid, bicarbonate ions, and carbonate
in freshwater and effectively promotes resistance to changes in pH.
If acid (hydrogen ions) is added to this buffer solution (water with
dissolved calcium bicarbonate), the equilibrium is shifted and carbonate
ions combine with the hydrogen ions to form bicarbonate. Subsequently,
the bicarbonate then combines with hydrogen ions to form carbonic acid,
which can dissociate into carbon dioxide and water. Thus the system pH
is unaltered (buffered) even though acid was introduced.
Natural waters vary in acidity and alkalinity from local geology and
other natural causes, as well as, anthropogenic inputs. Normally, pH
values below 5 and exceeding 9 are harmful to organisms, thus making
buffering capacity in water of critical importance.
Influence on Biota
The influence of pH and buffering on organisms is
often difficult to pin down specifically due to the interrelatedness
of many factors. Generally,
organisms appear to thrive in waters of increasing hardness (calcium
and magnesium concentration). Surveys of species abundance and species
richness have shown that stream water of very low ionic concentrations
has reduced species abundance and species richness. This appears to
true for trout (McFadden and Cooper, 1962) in terms of growth rate and
upper size limit. Molluscs, crustaceans, and leeches also show responses
to varying ionic concentrations (Hynes, 1970; Macan, 1974). Russel-Hunter
et al. (1967) showed that Ca2+ hardness could be associated with estimates
of species richness among mollusks for geographic regions.
||% of Potential Molluscan Species
|> 25 mg/l
||up to 100%
Often clear trends are not present when relating species abundance
or species richness to ion concentrations in streams. Factors such as
substrate complexity and stream temperature can interact with water chemistry
and, in many cases, may be more important in determining the suitability
of a stream´s environment to a particular species or taxonomic