Geology Reference
In-Depth Information
dissociation into bicarbonate (HCO 3 ) and carbonate
(CO 3 2− ) ions has already been considered in equations
4.21 and 4.22. The fact that weak acids like carbonic
acid are only partially dissociated gives them a capac-
ity to stabilize a solution against pH change. This
capacity is referred to as buffering .
Imagine two beakers on a laboratory bench, one con-
taining a litre of distilled water and the other a litre of
seawater. Upon adding a few drops of a strong acid
like HCl to the distilled water beaker, while observing
the change in pH using a pH-meter (Appendix B), we
would find that the pH falls from about 7 to perhaps 2
or 3, consistent with the increase in a H + that we would
expect to accompany the addition of strong acid.
The pH of the seawater, on the other hand, although
the beaker is treated exactly the same way, would be
found to remain close to its initial value of 8.1 . The expl-
anation lies in Equation 4.21. 8 The momentary increase
in a H + that occurs when the strong acid is added raises
the ion activity product a H CO
than it actually is. The present level in the oceans is
moderated by exchange reactions between seawater
and ocean-floor basalts, which remove Mg 2+ into min-
erals like chlorite which form as alteration products of
the original ferromagnesian minerals present in basalt.
The oceans do not simply accumulate in solution all of
the solute delivered by continental runoff; the concen-
trations of many elements in seawater are subject to
complex regulatory mechanisms, in many of which
biota also play an important role.
Brines and hydrothermal fluids ( I > 1.0 mol kg -1 )
Near-surface ground waters in the continental crust
are largely of meteoric origin; that is they are derived
ultimately from atmospheric precipitation. They often
have compositions of low ionic strength not very dif-
ferent from river water (Table 4.2), although depend-
ing to some extent on the type of rock through which
they have flowed. In coastal areas there may also be
a  component of seawater present. In deeply buried
sedimentary rocks, however, the pore waters are
connate in origin, originating as seawater trapped
during accumulation of the sediment. Drilling shows
that such waters - oilfield brines, pore waters, 'form-
ation water' and so on - are highly saline (Table 4.4),
having remained in contact with the host lithologies at
elevated temperatures for millions of years.
Hot, hypersaline aqueous fluids are important in
another context: the transport of metals and their
of the seawater,
boosting the ratio aa
well above its equilib-
rium value ( K 1 = 10 −6.4 ). This causes the reverse reaction
in Equation 4.21 to accelerate, driving the equilibrium
back to the left. In other words, the added H + reacts with
HCO 3 to form additional H 2 CO 3 , and this reaction
mops up much of the H + added and restores the pH to
close to its original value. The new a H + actually remains
slightly higher (and the pH slightly lower) than the orig-
inal value, but by no means as high (or the pH as low) as
it would have become had HCO 3 not been present.
This buffering capacity explains why seawater
samples across the globe have pH values lying within
narrow limits of 8.1 to 8.3. Because the dissolved inor-
ganic carbon (often abbreviated to 'DIC') in the oceans
represents a very large reservoir of carbonate, the capac-
ity of this buffer system is huge. Other weak acids pre-
sent in seawater, such as boric (H 3 BO 3 ) and phosphoric
(H 3 PO 4 ) acids, also have a buffering effect but  - being
less abundant - their contribution is much smaller.
Many other compositional aspects of seawater are
regulated by chemical reactions. If all of the Mg 2+ deliv-
ered by rivers were to accumulate in the oceans, for
example, seawater would be many times richer in Mg 2+
Table 4.4 Ground water and brine
Ground water in
Mississippi sandstone
at depth 40 m* (ppm)
Oilfield brine,
Mississippi, at depth
3330 m (ppm)
Na +
K +
Ca 2+
Mg 2+
Fe 2+
SO 4 2−
In principle, equilibrium 4.22 contributes to the buffering
effect too, but at pH ≈ 8 the concentration of CO 3 2− present is
too low (Figure 4.3.1) for it to have significant impact.
*From Todd and Mays (2006).
From Barnes (1979), Table 1.1.
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