Photosynthesis ensures a constant renewal of elemental oxygen, the main oxidant
in the geosphere, and produces organic compounds, which enter the soil and water
after the biota dies. These organic substances then act as reducing agents through
respiration or decay (natural decomposition of organic substances). This process is
actually the above-mentioned photosynthetic reaction (Eq. 1.2) in reverse; oxygen is
consumed releasing carbon dioxide, water, and energy.
Another very important component of the natural geochemical system is iron.
Iron is the most widespread transient metal element and the second most widespread
metal on Earth. The content of iron in the Earth’s crust is 62 g/kg. In nature, it is
mainly found in valence states of Fe
2+ and Fe
3+ . Under oxidation conditions, stable
iron is in the valence state of Fe
3+ ; under reducing conditions, it is in the valence
state of Fe
2+ , and minerals containing Fe
3+ are also common. Besides oxidation of
organic matter, a change in the valence state of Fe is the main process that generates
electron transfer in nature. Electrons are generated during oxidation and consumed
during reduction. Fe
4+ compounds are very unstable and have no practical significance, Fe
5+ and Fe
6+ compounds are used as very strong oxidizing agents but are not
found in nature. Pure iron rarely occurs in nature (in volcanites). Iron in low valence
states (Fe
0 ) can be used as a reducing agent.
In the case of iron (ZVI nanoparticles), it is known that water and dissolved
oxygen cause iron oxidation, which leads to corrosion (rusting) by the following
Eq. (1.3):
Fe
0
þ ½O 2 þ H 2 O ! Fe
2þ
þ 2OH
À
ð1:3Þ
The very name of the redox process—oxidation-reduction reaction—implies that
there are two processes occurring—oxidation and reduction, which are inseparable,
simultaneous, and dependent. This is a contrast to acid-base reactions, where protons
(H
+ ions) are released by, e.g., acid dissolution in water, solvated and stable in a
water environment.
Oxygen participates in the above process in the reduction part in which electrons
are consumed and the oxidation state of oxygen reduces according to the formal
reaction (Eq. 1.4).
½O
2
0 þ H
þ
þ 2e
À
! O
ÀII H
À
ð1:4Þ
Reversely, iron participates in the oxidation process in which electrons are
released and the iron oxidation state increases.
Fe
0
! Fe
2þ
þ 2e
À
ð1:5Þ
By combining these two equations—reduction (Eq. 1.4) and oxidation (Eq. 1.5)
respectively—we obtain the final equation for iron rusting, where two electrons are
exchanged between the iron and oxygen. For such a heterogeneous process, the
surface area of the solid reagent is the key parameter for the reaction rate. This has a
1 Geochemical Principles of Reductive Remediation Processes
5
in the geosphere, and produces organic compounds, which enter the soil and water
after the biota dies. These organic substances then act as reducing agents through
respiration or decay (natural decomposition of organic substances). This process is
actually the above-mentioned photosynthetic reaction (Eq. 1.2) in reverse; oxygen is
consumed releasing carbon dioxide, water, and energy.
Another very important component of the natural geochemical system is iron.
Iron is the most widespread transient metal element and the second most widespread
metal on Earth. The content of iron in the Earth’s crust is 62 g/kg. In nature, it is
mainly found in valence states of Fe
2+ and Fe
3+ . Under oxidation conditions, stable
iron is in the valence state of Fe
3+ ; under reducing conditions, it is in the valence
state of Fe
2+ , and minerals containing Fe
3+ are also common. Besides oxidation of
organic matter, a change in the valence state of Fe is the main process that generates
electron transfer in nature. Electrons are generated during oxidation and consumed
during reduction. Fe
4+ compounds are very unstable and have no practical significance, Fe
5+ and Fe
6+ compounds are used as very strong oxidizing agents but are not
found in nature. Pure iron rarely occurs in nature (in volcanites). Iron in low valence
states (Fe
0 ) can be used as a reducing agent.
In the case of iron (ZVI nanoparticles), it is known that water and dissolved
oxygen cause iron oxidation, which leads to corrosion (rusting) by the following
Eq. (1.3):
Fe
0
þ ½O 2 þ H 2 O ! Fe
2þ
þ 2OH
À
ð1:3Þ
The very name of the redox process—oxidation-reduction reaction—implies that
there are two processes occurring—oxidation and reduction, which are inseparable,
simultaneous, and dependent. This is a contrast to acid-base reactions, where protons
(H
+ ions) are released by, e.g., acid dissolution in water, solvated and stable in a
water environment.
Oxygen participates in the above process in the reduction part in which electrons
are consumed and the oxidation state of oxygen reduces according to the formal
reaction (Eq. 1.4).
½O
2
0 þ H
þ
þ 2e
À
! O
ÀII H
À
ð1:4Þ
Reversely, iron participates in the oxidation process in which electrons are
released and the iron oxidation state increases.
Fe
0
! Fe
2þ
þ 2e
À
ð1:5Þ
By combining these two equations—reduction (Eq. 1.4) and oxidation (Eq. 1.5)
respectively—we obtain the final equation for iron rusting, where two electrons are
exchanged between the iron and oxygen. For such a heterogeneous process, the
surface area of the solid reagent is the key parameter for the reaction rate. This has a
1 Geochemical Principles of Reductive Remediation Processes
5