Encyclopedia of Environmental Science and Engineering, Volume I and II

WATER CHEMISTRY 1265

The relations in Figure 5 do not fully describe the solu-

bility of the corresponding oxides and hydroxides, since

in addition to free metal ions, the solution may contain

hydrolyzed species (hydroxo complexes) of the form. The

solubility of the metal oxide or hydroxide is therefore

expressed more rigorously as

MeT Mez Me OHnzn

n

[][()].

1

∑ (25)

Plots of this equation as a function of pH are given in Figure 6

for ferric hydroxide, zinc oxide, and cupric oxide.

Solubility of Carbonates

The maximum soluble metal ion concentration is a function

of pH and concentration of total dissolved carbonate species.

Calculation of the equilibrium solubility of the metal ion for a

given carbonate for a water of a specific analytic composition

discloses whether the water is over-saturated or undersaturated

with respect to the solid metal carbonate. In the case of calcite

[]

[]

Ca

CO

2 0

3

2

0

2





KK

C

ss

Ta

. (26)

Since a 2 is known as a function of pH, Eq. (26) gives the

equilibrium saturation value of Ca^2 ^ as a function of C T

and pH. An analogous equation can be written for any

metallic cation in equilibrium with its solid metallic car-

bonate. These equations are amenable to simple graphical

representation in a log concentration versus pH diagram as

illustrated in Figure 7.

† Note that this solubility product is expressed for activities, as

represented by {}.

1 2 3 4 5 6

2 4 6810 12

Fe3+ Al3+

Cu2+ CuO(s)

Cu2+ Zn2+Fe2+Cd2+ Mg2+ Ag+

CO2+

–log [Me

z+

]

pH

FIGURE 5 Solubility of oxides and hydroxides: free metal ion concentration

in equilibrium with solid oxides ore hydroxides. As shown explicitly by the equi-

librium curve for copper, free metal ions are constrained to concentrations to the

left of (below) the respective curves. Precipitation of the solid hydroxides and

oxides commences at the saturation concentrations represented by the curves. The

formation of hydroxo metal complexes must be considered for the evaluation of

complete solubility of the oxides or hydroxides. Ref.: Stumm, W. and J. Morgan,

Aquatic Chemistry, Wiley-Interscience, New York, 1970, p. 171.

Control of Solubility

Solubility calculations, such as those exemplified above, give

thermodynamically meaningful conclusions, under the speci-

fied conditions (e.g., concentrations, pH, temperature and pres-

sure), only if the solutes are in equilibrium with that solid phase

for which the equilibrium relationship has been formulated.

For a given set of conditions the solubility is controlled by the

solid giving the smallest concentration of solute. For example,

within the pH range of carbonate bearing natural waters, the

stable solid phases regulating the solubility of Fe(II), Cu(II),

and Zn(II) are, respectively, FeCO 3 (siderite), CuO (tenorite)

and Zn (OH) (CO 3 ) (hydrozincite).

Unfortunately, it has not yet been possible to determine

precise solubility data for some solids important in the reg-

ulation of natural waters. Among these are many clays and

dolomite (CaMg(CO 3 ) 2 ), a mixed carbonate which con-

stitutes a large fraction of the total quantity of carbonate

rocks. The conditions under which dolomite is formed in

nature are not well understood and attempts to precipitate

it in the laboratory from solutions under atmospheric con-

ditions have been unsuccessful. These difficulties in ascer-

taining equilibrium have resulted in a diversity of published

figures for its solubility product, ({Ca^2 }{Mg^2 }{Co 32 }2}†,

ranging from 10^ 16.5 to 10^ 19.5 (25C).

The Activity of the Solid Phase

In a solid-solution equilibrium, the pure solid phase is defined

as a reference state and its activity is, because of its constancy,

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