Geology Reference
In-Depth Information
Note that the simple quantitative concept of solubil-
ity applicable to single-salt solutions ceases to have
any meaning in natural waters, where the activity of
each ionic species can include contributions from a
number of dissolved salts. This stresses the value of
expressing solubility in the form of an equilibrium
constant, the solubility product
K
. To summarize:
Increasing the air pressure or the concentration of
CO
2
in air would shift equilibrium 4.19 to the right,
increasing the solubility of CO
2
in water, as expressed
by
a
HCO
2
. Opening a can of fizzy drink, on the other
hand, releases the gas pressure inside, and because the
dissolved CO
2
has suddenly become supersaturated it
forms bubbles of the gas phase, a process similar to the
formation of vesicles in a molten lava.
Most gases (including CO
2
) are less soluble in hot
water than cold (Box 4.2).
3
Ionactivityproduct
>
K
meansasupersaturated
solution
meansas
;
=
K
aturated solution in
equilibrium with solid phase
meansanunde
;
Dissociation of weak acids
Equation 4.19 is in fact a slight simplification. The car-
bonic acid term (H
2
CO
3
) on the right-hand side actu-
ally represents the sum of
three
dissolved carbonate
species occurring in natural aqueous solutions: H
2
CO
3
0
,
HCO
3
−
and CO
3
2−
. (The superscripts 0, − and 2− merely
indicate the charge - or lack of it - on each of these
molecules.) In what relative proportions do these car-
bonate species occur?
Carbonic acid is an example of a weak acid
(Appendix B), meaning that, unlike familiar strong
laboratory acids such as HCl (hydrochloric acid) which
dissociate (ionize) completely in aqueous solution,
carbonic acid dissociates into ions only to a small
extent. This takes place in two successive stages:
<
K
rsaturated
solution.
Other kinds of equilibrium constant
Solubility of a gas
Water can dissolve gases as well as solids. A geologi-
cally important example is carbon dioxide (CO
2
),
which on dissolving forms a
weak acid
called
carbonic
acid
(H
2
CO
3
). Water in equilibrium with carbon dioxide
in the air is therefore always slightly acidic, a property
relevant to many chemical weathering processes.
The dissolving of CO
2
in water can be written as a
chemical reaction:
aa
⋅
CO HO HCO
2
+
(4.19)
+
−
H CO
+
+
-
10
64
−
.
1.
HCOHHCO
K
=
=
3
2
2
3
2
3
3
1
a
gas
solution
solution
carbonic
acid
bicarbonate
ion
HCO
2
3
(4.21)
Using the appropriate ways of recording concentra-
tions in these phases, the equilibrium constant is:
aa
⋅
+
2
−
H O
-
+
2
−
−
1
0.
2.
HCOHCO
+
K
=
=
10
3
a
pX
a
3
3
2
a
HCO
HCO
K
=
≈
(4.20)
bicarbonate
carbonate
2
3
2
3
-
HCO
3
HCO
⋅
p
2
3
(4.22)
CO
HO
CO
2
2
2
since the mole fraction of water
X
HO
2
in dilute solution
is very close to 1.00.
p
CO
2
is the
partial pressure
of CO
2
in the atmosphere, which in normal air at the present
time
3
is equal to 39.1 Pa (=391 ppmv).
This equilibrium constant, although relating to the sol-
ubility of a solute species, is quite different in form from
the solubility product of Equation 4.14. The reason is that
the behaviour of CO
2
in a solution, as reaction 4.19 shows,
is unlike that of ionic compounds, and this is reflected
in the mathematical form of an equilibrium constant.
It is these dissociation reactions, and the accompany-
ing release of hydrogen H
+
ions, that make aqueous
solutions of CO
2
slightly acidic (Appendix B). Indeed,
knowing the equilibrium constants
K
1
and
K
2
for reac-
tions 4.21 and 4.22, it is possible to calculate the
pH
(acidity) of pure water that has equilibrated with
atmospheric carbon dioxide (Exercise 4.3).
K
1
and
K
2
represent a class of equilibrium constant
known as
dissociation constants
. One finds that, with
all
polyprotic
acids such as H
2
CO
3
(Appendix B),
K
1
is much larger than
K
2
. The acidity that arises from dis-
solution of CO
2
in pure water is therefore almost
entirely due to reaction 4.21 alone.
3
The annual average value of
p
CO
2
has risen from 31.5 Pa in
1960 to 39.1 Pa today (Figure 9.2).
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