Geology Reference
In-Depth Information
But the pH of natural waters can vary widely
(Figure 4.1). Box 4.3 shows how pH affects the relative
stability and occurrence of the three natually occurring
forms of dissolved carbonate.
The carbonic acid present in natural waters deter-
mines whether they will dissolve carbonates (lime-
stone) or precipitate them:
local increase in atmospheric CO
2
content (as occurs in
soil pore spaces owing to the oxidation of organic mat-
ter, for example), or an increase in total pressure leads
to a higher H
2
CO
3
concentration which 'shifts the equi-
librium to the right' (meaning more CaCO
3
is dissolved
to form bicarbonate).
4
An increase in temperature, on
the other hand, expels dissolved CO
2
(Box 4.2 and pre-
vious section) and so makes CaCO
3
less soluble.
weathering
deposition
2
+
+
-
CaCO HCO
3
+
Ca
HCO
bicarbonate
2
(4.23)
2
3
3
solid
solution
solution
4
Some spring waters, having become saturated with calcium car-
bonate through equilibration with limestone at depth (i.e. under
elevated pressure), discharge CO
2
during ascent and deposit at
the surface a distinctively porous form of limestone called
tufa
.
Physical conditions influence this equilibrium chiefly
through the amount of dissolved CO
2
(as H
2
CO
3
). A
Box 4.3 Speciation of carbonic acid: the influence of pH
equations 4.21 and 4.22 suggest that potentially h
2
CO
3
,
hCO
3
-
and CO
3
2-
might
all
coexist in solution. Furthermore,
since
a
h
+
appears in both equilibrium constants, their rela-
tive proportions must vary with ph. that being so, which of
these three carbonate species would predominate in which
part of the ph range? We can answer that question by rear-
ranging equations 4.21 and 4.22 as follows:
1.
a
a
a
KK
−
pH
10
HCO
+
==
H
(4.3.1)
23
-
1
1
HCO
3
a
a
a
KK
−
pH
10
-
HCO
+
==
H
2.
(4.3.2)
3
2
-
2
2
CO
3
Figure 4.3.1
the relative abundance and stability ranges
of undissociated carbonic acid (h
2
CO
3
), bicarbonate ion
(hCO
3
−
) and carbonate ion (CO
3
2−
) - all represented as
mole % of total carbonate present - as a function of ph.
(Note that
a
CO
3
2
−
is negligible at low ph when h
2
CO
3
is
present and
a
HC
23
is insignificant at high ph when CO
3
2−
is present.)
since ph = −log
10
(
a
h
+
) and therefore
a
h+
= 10
−ph
(see
appendix a).
these rearranged equations tell us that the activity
ratios
a
a
and
a
a
-
HCO
HCO
are ph-dependent. We can see
23
3
-
2
-
HCO
CO
3
3
from equation 4.3.1 that h
2
CO
3
will be most abundant at
low values of ph (= high values of 10
−ph
), i.e. in acid solu-
tions - as shown in Figure 4.3.1 - whereas equation 4.3.2
tells us that CO
3
2−
will be found only in alkaline solutions
(high ph = low values of 10
−ph
). to calculate
a
HCO
3
Ḥ
as a per-
centage of
a
Figure 4.3.1 shows how the abundance of each species
varies with ph. the term
speciation
is often used to describe
the identity and relative abundance of the different chemi-
cal forms that an element or compound may adopt under
different conditions (for example, at various ph values).
Similar plots (called Bjerrum plots)
5
can be drawn for
other polyprotic acids present in seawater - such as boric
acid (h
3
BO
3
), phosphoric acid (h
3
pO
4
) and silicic acid
(h
2
SiO
3
) - which, like h
2
CO
3
, occur in a number of ph-
dependent dissociated forms (see Libes, 2009, for
Bjerrum plots of these acids).
(
)
+
a
:
HCO
-
HCO
23
3
a
+
a
a
a
−
pH
10
HCO
-
HCO
HCO
23
=
+
1
=
+
1
3
23
a
K
-
-
1
HCO
HCO
3
3
taking the inverse and multiplying by 100%:
−
1
a
−
pH
10
HCO
−
=× +
100
×
100
1
%
(4.3.3)
3
1
a
+
a
K
5
after Danish chemist Niels Bjerrum.
−
HCO
HCO
23
3
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