Agriculture Reference
In-Depth Information
the autocatalysis of Fe
2
+
oxidation by adsorption of Fe
2
+
on ferric hydroxide
formed in the reaction. The adsorption is likely to be pH-dependent, a decrease
in pH tending to decrease sorption and increase the concentration of Fe
2
+
in
solution. Hence there may be complex interactions between the mobility of Fe
2
+
,
the rate of oxidation, and pH changes caused by the reaction. Such interactions
can produce banded distributions of iron around an O
2
source, as found, for
example, by Saleque & Kirk (1995) for the distribution of iron near rice roots
and calculated by Kirk
et al
. (1990) with a model of the coupled diffusion and
reaction of O
2
,Fe
2
+
and acidity in soil. This is an example of the Liesegang
phenomenon (Stern, 1954; Keller, 1980).
KINETICS OF Fe
2
+
OXIDATION
4.4.1
Aqueous Solution
The reaction between Fe
2
+
and O
2
to form insoluble ferric hydroxide can be
written
4Fe
2
+
+
O
2
+
10H
2
O
=
4Fe(OH)
3
+
8H
+
(
4
.
36
)
Equation (4.36) shows that two H
+
ions are produced for each mole of Fe
2
+
oxidized, i.e. the reaction is accompanied by acidification. In aqueous solution,
the rate is found to be very sensitive to pH and at near neutral pH the reaction
is accelerated 100-fold if the pH is raised by one unit. The following empirical
rate law applies in the pH range 5 - 8 (Stumm and Lee, 1961; Wehrli, 1990)
−
d[Fe(II)]
/
d
t
=
k
[O
2
][OH
−
]
2
[Fe(II)]
(
4
.
37
)
where
k
≈
2
×
10
14
mol
3
dm
−
9
s
−
1
at 25
◦
C and [Fe(II)] is the sum of the
concentrations of Fe(II) species present — Fe
2
+
and its hydroxy complexes,
FeOH
+
and Fe(OH)
2
, for which the formation constants are 10
−
4
.
5
mol
−
1
dm
3
and 10
−
7
.
4
mol
−
2
dm
6
, respectively. Therefore [Fe(II)]
[Fe
2
+
], but the pH
dependence of the rate is due to the parallel oxidation of the three species.
At [O
2
]
≈
0
.
28 mM (i.e. in equilibrium with atmospheric
P
O
2
), the half time for
the reaction is 0.34 h at pH 7 and 143 days at pH 5.
As discussed in Section 4.1, most redox reactions reach equilibrium only
slowly if they are not catalysed. Oxidation of Fe
2
+
is catalysed by adsorption of
Fe
2
+
onto Fe(OH)
3
formed in the reaction, so Equation (4.36) only holds for the
initial rates of reaction. Tamura
et al
. (1976) studied the oxidation of a solution
of Fe
2
+
at different controlled pHs near neutral and with varying additions of
Fe(OH)
3
. The reaction obeyed the rate law
−
d[Fe
2
+
]
/
d
t
=
k
[O
2
][OH
−
]
2
[Fe
2
+
]
+
k
S
[O
2
][Fe
2
+
]
ad
=
(
4
.
38
)
where [Fe
2
+
] is the concentration in solution, [Fe
2
+
]
ad
the concentration adsorbed
on Fe(OH)
3
and
k
S
the rate constant for oxidation of adsorbed Fe
2
+
(
=
73 mol
−
1
