Carbon - Freshwater Ecology: Concepts and Environmental Applications - page 232

Increased atmospheric pressure allows greatly increased amounts of
CO 2 to be dissolved in solution. Thus, carbonated beverages stored under
pressure (such as soda, beer, and champagne) lose CO 2 when opened.
These beverages must be slightly acidic, so the equilibrium is forced to the
side of CO 2 . A similar, but catastrophic, release of CO 2 had disastrous con-
sequences in the African Lake Nyos (Sidebar 12.1).
The equilibrium of inorganic carbon also explains the acid neutraliz-
ing or buffering capability of bicarbonate and carbonate. Protons react
with the carbonate as acid is added, so pH changes only a small amount
relative to the concentration of H added to solution. Systems with a sig-
nificant amount of dissolved bicarbonate (e.g., limestone watersheds) are
able to resist the effects of acid precipitation. An opposite response occurs
to addition of base (OH ). The OH ions associate with the H ions so
the equilibrium balances by moving toward the bicarbonate side.
The acid- and base-neutralizing capacity of bicarbonate and the pre-
dominance of bicarbonate ions in many systems have led to using alkalin-
ity and acidity titrations to estimate CO 2 . For alkalinity titrations, acid

can be added with little initial change in pH. After all the bicarbonate and
carbonate have reacted with the added acid, further additions cause pro-
portionally greater decreases in pH per unit of acid added. The alkalinity
is the amount of acid needed to cause these greater decreases in pH.
The dependence of the bicarbonate equilibrium on pH yields plots that
can be used to calculate the relative concentrations of each of the forms of
inorganic carbon when the pH is known (Fig. 12.2). Such data are useful
because CO 2 is the form of inorganic carbon required for photosynthesis.
Many photosynthetic organisms can convert bicarbonate to CO 2 , but CO 2
is still the most easily used form. Thus, knowing the alkalinity and pH of
a solution allows an investigator to calculate the amount of inorganic car-
bon that is immediately available for photosynthesis.
One important precipitate of the bicarbonate equilibrium is calcium bi-
carbonate. The precipitate can form spontaneously when CO 2 is removed
from solution (as a way to balance the equilibrium). This precipitation can
occur when photosynthesis removes CO 2 , when physical factors remove
CO 2 (e.g., degassing of spring waters), or when organisms such as mol-
lusks build their shells (Wetzel, 2001). The resulting precipitate can build
up impressive concretions of whitish calcium bicarbonate on the stems of
the green alga Chara (hence the name stoneworts), in terraced outflows of
hot springs, and in some benthic habitats with photosynthetic microor-
ganisms. Mammoth Hot Springs in Yellowstone National Park is com-
posed of massive carbonate terraces.

Organic Carbon
Organic carbon takes a tremendous variety of forms. The broadest clas-
sifications are dissolved organic carbon (DOC) and particulate organic car-
bon (POC). Stream ecologists further divide the particulate fractions into
fine particulate organic matter (FPOM) and coarse particulate organic mat-
ter (CPOM). A dividing line of 0.5

m has been proposed for the differ-
ence between DOC and FPOM, and a line of 500 m has been proposed

for the division between FPOM and CPOM (Allan, 1995). Further divisions