1256WATER CHEMISTRY
AQUATIC CHEMICAL EQUILIBRIAIn this section a few example will be given that demonstrate
how elementary principles of physical chemistry can aid in
the recognition of interrelated variables that establish the
composition of natural waters. Natural water systems usu-
ally consist of numerous mineral assemblages and often of
a gas phase in addition to the aqueous phase; they nearly
always include a portion of the biosphere. Hence, natural
aquatic habitats are characterized by a complexity seldom
encountered in the laboratory. In order to distill the perti-
nent variables out of a bewildering number of possible ones,
it is advantageous to compare the real systems with their
idealized counterparts.
Thermodynamic equilibrium concepts represent the
most expedient means of identifying the variables relevant
in determining the mineral relationships and in establishing
chemical boundaries of aquatic environments. Since mini-
mum free energy describes the thermodynamically stable
state of a system, a comparison with the actual free energy
can characterize the direction and extent of processes that
are approaching equilibrium. Discrepancies between equi-
librium calculations and the available data of real systems
give valuable insight into those cases where chemical reac-
tions are not understood sufficiently, where non-equilibrium
conditions prevail, or where the analytical data are not suf-
ficiently accurate or specific.Alkalinity and Acidity for Aqueous Carbonate
Systems
Alkalinity and acidity are defined, respectively, as the
equivalent sum of the bases that are titratable with strong
acid and the equivalent sum of the acids that are titratable
with strong base; they are therefore capacity factors which
represent, respectively, the acid and base neutralizing
capacities of an aqueous system. Operationally, alkalinity
and acidity are determined by acidimetric and alkalimetric
titrations to appropriate pH end points. These ends points
(equivalence points) occur at the infection points of titra-
tion curves as shown in Figure 1 for the carbonate system.
The atmosphere contains CO 2 at a partial pressure of
3 10 ^4 atmosphere, while CO 2 , H 2 CO 3 , HCO 3 and CO 32
are important solutes in the hydrosphere. Indeed, the carbon-
ate system is responsible for much of the pH regulation in
natural waters.The following equations define for aqueous carbonate
systems the three relevant capacity factors: Alkalinity (Alk),
Acidity (Acy), and total dissolved carbonate species ( C T ): †[]Alk ⎡⎣HCO− (^33) ⎦⎤^2 ⎣⎡CO^2 ⎦⎤ ⎡⎣OH ⎦⎤ ⎣⎡H ⎤⎦ (1)
[]Acy^2 []H CO 23 ⎣⎡HCO (^3) ⎦⎤ ⎣⎡H ⎦⎤ ⎣⎡OH⎦⎤ (2)
Cb[]H CO 23 ⎣⎡HCO 3 ⎤⎦ ⎣⎡CO^23 ⎤⎦ (3)
where [H 2 CO 3 * ] [CO 2 (aq)] [H 2 CO 3 ].
These equations are of analytical value because they
represent rigorous conceptual definitions of the acid neutral-
izing and the base neutralizing capacities of carbonate sys-
tems. The definitions of alkalinity and acidity algebraically
† Brackets of the form [ ] refer to concentration, e.g., in moles per
liter.
FIGURE 1 Alkalinity and acidity titration curve for the aque-
ous carbonate system. The conservative quantities alkalinity
and acidity refer to the acid neutralizing and base neutralizing
capacities of a given aqueous system. These parameters can be
determined by titration to appropriate equivalence points with
strong acid and strong base. The equations given below define
the various capacity factors rigorously. Figure from Stumm, W.
and J. Morgan, Aquatic Chemistry, Wiley-Interscience, New
York, 1970, p. 130.
9
7
5
xyz
pH
[CO 2 –Acy] [CO 3 2––Alk]
[Acy]
Addition of Acid
Addition of Base
[H+–Acy] [Alk]
[OH––Alk]
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