498
Chemical Oceanography, 4th Edition
can also increase the rates since it has a higher rate of oxidation with sulfide than oxygen.
The Fe 2+ formed from the reaction of dissolved or particulate Fe 3+ can regenerate Fe 2+ to
complete the catalytic cycle. The effect of Fe 3+ on the rates at low concentrations may be
related to the reduction of Fe 3+ to Fe 2+ with HS – and resultant oxidation of Fe 2+ and generation of O 2
– . The effect of the metals on the rates of oxidation of H 2 S below the observable
precipitation of metal sulfides (which may be a slow process) can be attributed to the formation of ion pairs:
M 2+ + HS – → MHS +
(10.85)
The overall rate constant is given by
k[HS – ] T = k HS [HS – ] + k MHS [MHS + ]
(10.86)
where k HS and k MHS are the rate constants for the oxidation of HS – and MHS + , respectively.
If k MHS is greater than k HS , the rate can be increased with the addition of the metal.
The presence of Fe and Mn in natural waters not only increases the rate of oxidation of
sulfide but also can have an effect on the oxidation of intermediates such as sulfite. This
can change the distribution of the products formed during the oxidation.
The final product of the oxidation of sulfide is sulfate, the sulfur compound with the
highest oxidation state and the most stable compound in oxic waters. Various intermediates, such as sulfite and thiosulfate, also can be formed during the course of the reaction.
The products formed from the oxidation of H 2 S in seawater have been studied as a function
of pH, temperature, salinity, and reactant concentration. To examine the mass balance of
sulfur compounds during the oxidation, the experiments were made in pure water where
SO 4
2– formed from oxidation could be measured by an ion chromatographic technique
(Figure 10.86). The major products formed were SO 4
2– , SO 3
2– , and S 2 O 3
2– . Elemental sulfur
log [M]
–8
–7
–6
–5
–4
–3
log (k
H
2 S /k
o
H
2 S )
–1
0
1
2
3
Fe(II)
Fe(III)
Mn(II)
Figure 10.85
The effect of Fe and Mn on the rate of oxidation of H 2 S in seawater at 25°C and pH 8.1.
Chemical Oceanography, 4th Edition
can also increase the rates since it has a higher rate of oxidation with sulfide than oxygen.
The Fe 2+ formed from the reaction of dissolved or particulate Fe 3+ can regenerate Fe 2+ to
complete the catalytic cycle. The effect of Fe 3+ on the rates at low concentrations may be
related to the reduction of Fe 3+ to Fe 2+ with HS – and resultant oxidation of Fe 2+ and generation of O 2
– . The effect of the metals on the rates of oxidation of H 2 S below the observable
precipitation of metal sulfides (which may be a slow process) can be attributed to the formation of ion pairs:
M 2+ + HS – → MHS +
(10.85)
The overall rate constant is given by
k[HS – ] T = k HS [HS – ] + k MHS [MHS + ]
(10.86)
where k HS and k MHS are the rate constants for the oxidation of HS – and MHS + , respectively.
If k MHS is greater than k HS , the rate can be increased with the addition of the metal.
The presence of Fe and Mn in natural waters not only increases the rate of oxidation of
sulfide but also can have an effect on the oxidation of intermediates such as sulfite. This
can change the distribution of the products formed during the oxidation.
The final product of the oxidation of sulfide is sulfate, the sulfur compound with the
highest oxidation state and the most stable compound in oxic waters. Various intermediates, such as sulfite and thiosulfate, also can be formed during the course of the reaction.
The products formed from the oxidation of H 2 S in seawater have been studied as a function
of pH, temperature, salinity, and reactant concentration. To examine the mass balance of
sulfur compounds during the oxidation, the experiments were made in pure water where
SO 4
2– formed from oxidation could be measured by an ion chromatographic technique
(Figure 10.86). The major products formed were SO 4
2– , SO 3
2– , and S 2 O 3
2– . Elemental sulfur
log [M]
–8
–7
–6
–5
–4
–3
log (k
H
2 S /k
o
H
2 S )
–1
0
1
2
3
Fe(II)
Fe(III)
Mn(II)
Figure 10.85
The effect of Fe and Mn on the rate of oxidation of H 2 S in seawater at 25°C and pH 8.1.