Na
+ are enriched in seawater. This because the strong
hydration prevents absorption on clay minerals which
usually have a negative surface charge. K
+
, Rb
+ and
Cs
+
, on the other hand, have larger ionic radii and
consequently are less strongly hydrated. This leaves
them with a more effective positive surface charge
which facilitates their adsorption onto clay minerals,
etc.
This is demonstrated in nature during weathering
and transport. While similar amounts of potassium and
sodium are dissolved during weathering of basement
rocks, the potassium concentration in the sea is much
lower (K/Na ratio of only 1:30). This is because K
+ is
more effectively removed by adsorption because it is
less protected by hydration. The same is true to an
even greater extent for Rb
+ and Cs
+
, which are
adsorbed even more readily. These ions therefore
have a relatively short residence time in seawater,
between being delivered by rivers and then removed
by accumulating sediment. The ratio between the
average concentration of an ion in river water and
the concentration in average sea water is an expression
of the residence time.
The large potassium ions can only fit in to certain
crystal structures. In mica and illite each potassium ion
is surrounded by twelve oxygen or OH groups.
With regard to Group 2 elements, Mg
++ for example
will be more strongly hydrated than Ca
++ because it is
a smaller ion. As a result, Mg
++ has a greater tendency
to remain in solution in seawater and has a residence
time which is 13 times that of calcium. However,
despite the fact that the Mg/Ca ratio in seawater is 5,
it is calcium carbonate which is the first to form
through chemical and biological precipitation. Dolomite or magnesite do not precipitate directly from
seawater and this is in part due to the strong hydration
of Mg
++
. Normally, if we had naked (unhydrated) ions,
MgCO 3 and FeCO 3 would be more stable than CaCO 3
because Mg
++ and Fe
++ have greater ionic potentials
and stronger bonding to the CO
2À
3 ion. However with
increasing temperature the hydration declines because
the bonds with the dipole of the water molecules
become weaker. Mg
++ is then more likely to be
incorporated into the carbonate mineral structures.
Therefore during diagenetic processes at 80–100
C,
magnesium carbonates precipitate more readily even
if the Mg
þþ
=Ca
þþ and Fe
þþ
=Ca
þþ ratios are low.
Even if Mg is preferred in the carbonate structure and
also in the clay minerals, very little magnesium is
usually available in the porewater in the deeper parts
of sedimentary basins except in the presence of
evaporites with Mg salts.
3.2
Redox Potentials (Eh)
Oxidation potential (E) is an expression of the tendency
of an element to be oxidised, i.e. to give up electrons so
it is left with a more positive charge. This potential can
be measured by recording the potential difference (positive or negative) which arises when an element
functions as one electrode in a galvanic cell. The other
electrode is a standard one, normally hydrogen. The
oxidation potential of the reaction H 2 ¼ 2H
þ
þ 2e
(electrons) is defined as E
0
¼ 0:0 V at 1 atm and H
+
concentration of 1 mol/l at 20
C. Different conventions
have been used to assign plus and minus values. In
geochemical literature, metals with a higher reducing
potential than hydrogen are assigned negative values,
e.g. Na ¼ Na
þ
þ e
À
¼ À2:71V, while strongly
oxidising elements are given a positive sign, e.g.
2F
À
¼ F 2 þ 2e ¼ 2:87V. A list of redox potentials
shows which elements will act as oxidising agents,
and which will be reducing agents. Reactions which
result in a negative oxidation potential (E) will proceed
spontaneously, while those which have positive voltage
will be dependent on the addition of energy from an
outside source. We can predict whether a redox reaction
will occur by using Nernst’s Law (see chemistry
textbooks).
3.3
pH
The
ionisation
product
for
water
is
H
þ
½ Á OH
À
½
¼10
À14 . The concentration of H
+ in
neutral water will be 10
À7 . pH is defined as the negative logarithm of the hydrogen ion concentration activity, and is therefore 7 for neutral water (at 25
C).
However, the ionisation constant varies with temperature, e.g. at 125
C the ionisation constant for water is
½H
þ
Á ½OH
À
¼ 10
À12 . In other words, neutral water
then has a pH of 6. It is important to remember this
when considering the pH of hot springs or in deep
wells, for example oil wells.
94
K. Bjørlykke
+ are enriched in seawater. This because the strong
hydration prevents absorption on clay minerals which
usually have a negative surface charge. K
+
, Rb
+ and
Cs
+
, on the other hand, have larger ionic radii and
consequently are less strongly hydrated. This leaves
them with a more effective positive surface charge
which facilitates their adsorption onto clay minerals,
etc.
This is demonstrated in nature during weathering
and transport. While similar amounts of potassium and
sodium are dissolved during weathering of basement
rocks, the potassium concentration in the sea is much
lower (K/Na ratio of only 1:30). This is because K
+ is
more effectively removed by adsorption because it is
less protected by hydration. The same is true to an
even greater extent for Rb
+ and Cs
+
, which are
adsorbed even more readily. These ions therefore
have a relatively short residence time in seawater,
between being delivered by rivers and then removed
by accumulating sediment. The ratio between the
average concentration of an ion in river water and
the concentration in average sea water is an expression
of the residence time.
The large potassium ions can only fit in to certain
crystal structures. In mica and illite each potassium ion
is surrounded by twelve oxygen or OH groups.
With regard to Group 2 elements, Mg
++ for example
will be more strongly hydrated than Ca
++ because it is
a smaller ion. As a result, Mg
++ has a greater tendency
to remain in solution in seawater and has a residence
time which is 13 times that of calcium. However,
despite the fact that the Mg/Ca ratio in seawater is 5,
it is calcium carbonate which is the first to form
through chemical and biological precipitation. Dolomite or magnesite do not precipitate directly from
seawater and this is in part due to the strong hydration
of Mg
++
. Normally, if we had naked (unhydrated) ions,
MgCO 3 and FeCO 3 would be more stable than CaCO 3
because Mg
++ and Fe
++ have greater ionic potentials
and stronger bonding to the CO
2À
3 ion. However with
increasing temperature the hydration declines because
the bonds with the dipole of the water molecules
become weaker. Mg
++ is then more likely to be
incorporated into the carbonate mineral structures.
Therefore during diagenetic processes at 80–100
C,
magnesium carbonates precipitate more readily even
if the Mg
þþ
=Ca
þþ and Fe
þþ
=Ca
þþ ratios are low.
Even if Mg is preferred in the carbonate structure and
also in the clay minerals, very little magnesium is
usually available in the porewater in the deeper parts
of sedimentary basins except in the presence of
evaporites with Mg salts.
3.2
Redox Potentials (Eh)
Oxidation potential (E) is an expression of the tendency
of an element to be oxidised, i.e. to give up electrons so
it is left with a more positive charge. This potential can
be measured by recording the potential difference (positive or negative) which arises when an element
functions as one electrode in a galvanic cell. The other
electrode is a standard one, normally hydrogen. The
oxidation potential of the reaction H 2 ¼ 2H
þ
þ 2e
(electrons) is defined as E
0
¼ 0:0 V at 1 atm and H
+
concentration of 1 mol/l at 20
C. Different conventions
have been used to assign plus and minus values. In
geochemical literature, metals with a higher reducing
potential than hydrogen are assigned negative values,
e.g. Na ¼ Na
þ
þ e
À
¼ À2:71V, while strongly
oxidising elements are given a positive sign, e.g.
2F
À
¼ F 2 þ 2e ¼ 2:87V. A list of redox potentials
shows which elements will act as oxidising agents,
and which will be reducing agents. Reactions which
result in a negative oxidation potential (E) will proceed
spontaneously, while those which have positive voltage
will be dependent on the addition of energy from an
outside source. We can predict whether a redox reaction
will occur by using Nernst’s Law (see chemistry
textbooks).
3.3
pH
The
ionisation
product
for
water
is
H
þ
½ Á OH
À
½
¼10
À14 . The concentration of H
+ in
neutral water will be 10
À7 . pH is defined as the negative logarithm of the hydrogen ion concentration activity, and is therefore 7 for neutral water (at 25
C).
However, the ionisation constant varies with temperature, e.g. at 125
C the ionisation constant for water is
½H
þ
Á ½OH
À
¼ 10
À12 . In other words, neutral water
then has a pH of 6. It is important to remember this
when considering the pH of hot springs or in deep
wells, for example oil wells.
94
K. Bjørlykke