5
Bacteria and Marine Biogeochemistry
182
By this reaction, a single electron is transferred
from each Fe
2+
ion to the O 2 molecule. The ferrous
iron is thereby oxidized to ferric iron, whereas the
oxygen is reduced to water. In the geochemical
literature, O 2 in such a reaction is termed the
oxidant and Fe
2+
the reductant. In the biological
literature, the terms electron acceptor for O 2 and
electron donor for Fe
2+
are used.
5.4.1
Free Energy
Chemical reactions catalyzed by microorganisms
yield highly variable amounts of energy and some
are directly energy consuming. The term free
energy, G, of a reaction is used to express the
energy released per mol of reactant, which is
available to do useful work. The change in free
energy is conventionally expressed as ∆G
0
, where
the symbol ∆ should be read as ‘change in’ and
the superscript
0
indicates the following standard
conditions: pH 7, 25°C, a 1 M concentration of all
reactants and products and a 1 bar partial
pressure of gases. The value of ∆G
0
for a given
reaction is expressed in units of kilojoule (kJ) per
mol of reactant. If there is a net decrease in free
energy (∆G
0
is negative) then the process is
exergonic and may proceed spontaneously or
biologically catalyzed. If ∆G
0
is positive, the
process is endergonic and energy from ATP or
from an accompanying process is required to drive
the reaction. The change in free energy of several
metabolic processes in prokaryotic organisms is
listed in Table 5.2.
A thorough discussion of the theory and
calculation of ∆G
0
for a variety of anaerobic
microbial processes is given by Thauer et al.
(1977). As a general rule, processes for which the
release of energy is very small, < ca 20 kJ mol
-1
, are
insufficient for the formation of ATP and are thus
unable to serve the energy metabolism of
microorganisms (Hoehler et al. 1998). It is
important to note, however, that standard
conditions are seldom met in the marine
environment and that the actual conditions of pH,
temperature and substrate/product concentrations
must be known before the energetics of a certain
reaction can be realistically calculated. Several
processes, which under standard conditions
would be endergonic, may be exergonic in the
normal marine sediment.
An important example of this is the formation
of H 2 in several bacterial fermentation processes
which is exergonic only under low H 2 partial
pressure. The hydrogen cycling in sediments is
therefore dependent on the immediate consumption of H 2 by other organisms, such as the sulfate
reducing bacteria, which keep the partial pressure
of H 2 extremely low. The H 2 -producing and the H 2 -
consuming bacteria thus tend to grow in close
proximity to each other, thereby facilitating the
diffusional transfer of H 2 at low concentration
from one organism to the other, a so-called ‘interspecies hydrogen transfer’ (Conrad et al. 1986;
Schink 1997).
5.4.2
Reduction-Oxidation Processes
The electron transfer in redox processes is often
accompanied by a transfer of protons, H
+
. The
simplest example is the oxidation of H 2 with O 2 by
the so-called ‘Knallgas-bacteria’, which occur
widespread in aquatic sediments:
H 2 + ½O 2 → H 2 O
(5.14)
In the energy metabolism of cells, an intermediate
carrier of the electrons (and protons) is commonly
required. Such an electron carrier is, for instance,
NAD
+
(nicotinamid adenin dinucleotide), which
formally accepts two electrons and one proton
and is thereby reduced to NADH. The NADH may
give off the electrons again to specialized electron
acceptors and the protons are released in the cell
sap. Thereby, the NADH, which must be used
repeatedly, is recycled.
A redox process such as Eq. 5.14 formally
consists of two reversible half-reactions. The first
is the oxidation of H 2 to release electrons and
protons:
H 2 ⇔ 2e - + 2H +
(5.15)
The second is the reduction of oxygen by the
transfer of electrons (and protons):
½O 2 +2e - + 2H + ⇔ H 2 O
(5.16)
Compounds such as H 2 and O 2 vary strongly in
their tendency to either give off electrons and
thereby become oxidized (H 2 ) or to accept
electrons and thereby become reduced (O 2 ). This
tendency is expressed as the redox potential, E 0
'
,
of the compounds. This potential is expressed in
volts and is measured electrically in reference to
a standard compound, namely H 2 . By convention,
redox potentials are expressed for half reactions
Bacteria and Marine Biogeochemistry
182
By this reaction, a single electron is transferred
from each Fe
2+
ion to the O 2 molecule. The ferrous
iron is thereby oxidized to ferric iron, whereas the
oxygen is reduced to water. In the geochemical
literature, O 2 in such a reaction is termed the
oxidant and Fe
2+
the reductant. In the biological
literature, the terms electron acceptor for O 2 and
electron donor for Fe
2+
are used.
5.4.1
Free Energy
Chemical reactions catalyzed by microorganisms
yield highly variable amounts of energy and some
are directly energy consuming. The term free
energy, G, of a reaction is used to express the
energy released per mol of reactant, which is
available to do useful work. The change in free
energy is conventionally expressed as ∆G
0
, where
the symbol ∆ should be read as ‘change in’ and
the superscript
0
indicates the following standard
conditions: pH 7, 25°C, a 1 M concentration of all
reactants and products and a 1 bar partial
pressure of gases. The value of ∆G
0
for a given
reaction is expressed in units of kilojoule (kJ) per
mol of reactant. If there is a net decrease in free
energy (∆G
0
is negative) then the process is
exergonic and may proceed spontaneously or
biologically catalyzed. If ∆G
0
is positive, the
process is endergonic and energy from ATP or
from an accompanying process is required to drive
the reaction. The change in free energy of several
metabolic processes in prokaryotic organisms is
listed in Table 5.2.
A thorough discussion of the theory and
calculation of ∆G
0
for a variety of anaerobic
microbial processes is given by Thauer et al.
(1977). As a general rule, processes for which the
release of energy is very small, < ca 20 kJ mol
-1
, are
insufficient for the formation of ATP and are thus
unable to serve the energy metabolism of
microorganisms (Hoehler et al. 1998). It is
important to note, however, that standard
conditions are seldom met in the marine
environment and that the actual conditions of pH,
temperature and substrate/product concentrations
must be known before the energetics of a certain
reaction can be realistically calculated. Several
processes, which under standard conditions
would be endergonic, may be exergonic in the
normal marine sediment.
An important example of this is the formation
of H 2 in several bacterial fermentation processes
which is exergonic only under low H 2 partial
pressure. The hydrogen cycling in sediments is
therefore dependent on the immediate consumption of H 2 by other organisms, such as the sulfate
reducing bacteria, which keep the partial pressure
of H 2 extremely low. The H 2 -producing and the H 2 -
consuming bacteria thus tend to grow in close
proximity to each other, thereby facilitating the
diffusional transfer of H 2 at low concentration
from one organism to the other, a so-called ‘interspecies hydrogen transfer’ (Conrad et al. 1986;
Schink 1997).
5.4.2
Reduction-Oxidation Processes
The electron transfer in redox processes is often
accompanied by a transfer of protons, H
+
. The
simplest example is the oxidation of H 2 with O 2 by
the so-called ‘Knallgas-bacteria’, which occur
widespread in aquatic sediments:
H 2 + ½O 2 → H 2 O
(5.14)
In the energy metabolism of cells, an intermediate
carrier of the electrons (and protons) is commonly
required. Such an electron carrier is, for instance,
NAD
+
(nicotinamid adenin dinucleotide), which
formally accepts two electrons and one proton
and is thereby reduced to NADH. The NADH may
give off the electrons again to specialized electron
acceptors and the protons are released in the cell
sap. Thereby, the NADH, which must be used
repeatedly, is recycled.
A redox process such as Eq. 5.14 formally
consists of two reversible half-reactions. The first
is the oxidation of H 2 to release electrons and
protons:
H 2 ⇔ 2e - + 2H +
(5.15)
The second is the reduction of oxygen by the
transfer of electrons (and protons):
½O 2 +2e - + 2H + ⇔ H 2 O
(5.16)
Compounds such as H 2 and O 2 vary strongly in
their tendency to either give off electrons and
thereby become oxidized (H 2 ) or to accept
electrons and thereby become reduced (O 2 ). This
tendency is expressed as the redox potential, E 0
'
,
of the compounds. This potential is expressed in
volts and is measured electrically in reference to
a standard compound, namely H 2 . By convention,
redox potentials are expressed for half reactions