Environmental Engineering Reference
In-Depth Information
change) undoubtedly will influence aquatic ecosystems (Hutchin
et al.,
1995; Magnuson
et al.,
1997; Megonigal and Schlesinger 1997; Tobert
et al.,
1996).
When CO
2
is dissolved in water, it can exist in a variety of forms, de-
pending on pH. The forms are
carbon dioxide, carbonic acid, bicarbonate,
and
carbonate.
The sum of the concentrations of all these forms is the in-
organic carbon concentration and is signified as
CO
2
. Under most con-
ditions in aquatic systems, CO
2
is rapidly converted to carbonic acid so
they will be considered the same. The chemical conversions among these
forms are referred to as the
bicarbonate equilibrium
. Understanding this
series of chemical reactions is necessary to comprehend how aquatic ecosys-
tems are buffered against changes in pH and how CO
2
becomes available
for photosynthesis (Butler, 1991). The bicarbonate equilibrium can be rep-
resented as
H
HCO
3
2H
CO
3
CO
2
H
2
O
⇔
H
2
CO
3
⇔
⇔
carbon dioxide
carbonic acid
bicarbonate
carbonate
The “
” symbol indicates an equilibrium reaction. Adding or taking away
chemicals at any part of the reaction can force the reaction. For example,
if acid (H
) is added to a bicarbonate solution, the equilibrium is weighted
too heavily to the right-hand side of the equation, so the bicarbonate will
convert spontaneously to carbonic acid or carbon dioxide. This is demon-
strated easily by adding an acid such as vinegar to a solution of the sodium
salt of bicarbonate (baking soda). Adding the acid will cause production
of CO
2
as the equilibrium is reestablished. Because the CO
2
gas has a lim-
ited solubility in acidic water, it will bubble out. Thus, as pH changes so
do the relative amounts of bicarbonate, carbonate, and carbonic acid (Fig.
12.2).
⇔
1.0
HCO
3
-
CO
2
(H
2
CO
3
)
CO
3
2-
0.8
0.6
0.4
0.2
0.0
3
4
5
6
7
8
9
10
11
12
pH
FIGURE 12.2
The relative concentrations of inorganic compounds involved in the bicar-
bonate equilibrium as a function of pH.
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