497
Processes in the Oceans
k = (k 0 + k 1 K 1 /[H+])/(1 + K 1 /[H + ])
(10.77)
where k 0 = 80 kg H 2 O mol –1 h –1 for the oxidation of H 2 S and k 1 = 344 kg H 2 O mol –1 for the
oxidation of HS – :
k 0
H 2 S + O 2 → Products
(10.78)
k 1
HS– + O 2 → Products
(10.79)
The value of K 1 is the dissociation constant for the ionization of H 2 S. The effects of temperature and ionic strength on the rate constants k 0 and k 1 have been given by
log k 0 = 9.22 – (2.4 × 10 3 )/T
(10.80)
log k 1 = 10.50 + 0.16 pH – (3.0 × 10 3 )/T + 0.44 I 1/2
(10.81)
These equations are valid from pH = 4 to 8, t = 5 to 65°C, and I = 0 to 6 M.
Field measurements made on the oxidation of H 2 S in the Black Sea (Millero, 1991c), the
Framvaren Fjord (Millero, 1991a), the Chesapeake Bay (Millero, 1991b), and the Cariaco
Trench (Zhang and Millero, 1993a,b) yielded half- times that were much faster than those
determined in the laboratory on Gulf Stream seawater (see Figure 10.61). To determine if
this increase was due to trace metals, we measured the rates of oxidation of H 2 S in seawater with added transition metals (Vazquez, Zhang, and Millero, 1989). These studies show
that at concentrations below 300 nM, the rates were only affected by Fe 2+ , Cu 2+ , and Pb 2+ .
At higher metal concentrations, the rates of oxidation of H 2 S increased for all the metals
except Zn 2+ . The order of the increase in the rates at higher concentrations for these metals is Fe 2+ > Pb 2+ > Cu 2+ > Fe 3+ > Cd 2+ > Ni 2+ > Co 2+ > Mn 2+ .
Only Fe 2+ and Mn 2+ have levels in anoxic basins high enough to affect the oxidation
of H 2 S. The effect of metals on the oxidation of H 2 S with oxygen can be estimated from
(Figure 10.85)
log(k/ k 0 ) = a + b log[M]
(10.82)
where
a = 6.55, b = 0.820 for Fe(II)
a = 5.18, b = 0.717 for Fe(III)
a = 1.68, b = 0.284 for Mn(II)
These equations are valid, respectively, from 10 –8 to 10 –5.3 , 10 –7.2 to 10 –3.3 , and 10 –5.9 to 10 –3.3 M.
The larger effect of Fe 2+ is probably related to the formation of dissolved Fe 3+ from the
rapid oxidation of Fe 2+ with oxygen. The oxidation of Fe 2+ provides a higher initial concentration of Fe 3+ than can be added from a stock solution of Fe 3+ (which may be locally
supersaturated). The peroxide generated by the oxidation of Fe 2+ ,
O 2
– + H + = HO 2
(10.83)
HO 2 + HO 2 = H 2 O 2 + O 2
(10.84)
Processes in the Oceans
k = (k 0 + k 1 K 1 /[H+])/(1 + K 1 /[H + ])
(10.77)
where k 0 = 80 kg H 2 O mol –1 h –1 for the oxidation of H 2 S and k 1 = 344 kg H 2 O mol –1 for the
oxidation of HS – :
k 0
H 2 S + O 2 → Products
(10.78)
k 1
HS– + O 2 → Products
(10.79)
The value of K 1 is the dissociation constant for the ionization of H 2 S. The effects of temperature and ionic strength on the rate constants k 0 and k 1 have been given by
log k 0 = 9.22 – (2.4 × 10 3 )/T
(10.80)
log k 1 = 10.50 + 0.16 pH – (3.0 × 10 3 )/T + 0.44 I 1/2
(10.81)
These equations are valid from pH = 4 to 8, t = 5 to 65°C, and I = 0 to 6 M.
Field measurements made on the oxidation of H 2 S in the Black Sea (Millero, 1991c), the
Framvaren Fjord (Millero, 1991a), the Chesapeake Bay (Millero, 1991b), and the Cariaco
Trench (Zhang and Millero, 1993a,b) yielded half- times that were much faster than those
determined in the laboratory on Gulf Stream seawater (see Figure 10.61). To determine if
this increase was due to trace metals, we measured the rates of oxidation of H 2 S in seawater with added transition metals (Vazquez, Zhang, and Millero, 1989). These studies show
that at concentrations below 300 nM, the rates were only affected by Fe 2+ , Cu 2+ , and Pb 2+ .
At higher metal concentrations, the rates of oxidation of H 2 S increased for all the metals
except Zn 2+ . The order of the increase in the rates at higher concentrations for these metals is Fe 2+ > Pb 2+ > Cu 2+ > Fe 3+ > Cd 2+ > Ni 2+ > Co 2+ > Mn 2+ .
Only Fe 2+ and Mn 2+ have levels in anoxic basins high enough to affect the oxidation
of H 2 S. The effect of metals on the oxidation of H 2 S with oxygen can be estimated from
(Figure 10.85)
log(k/ k 0 ) = a + b log[M]
(10.82)
where
a = 6.55, b = 0.820 for Fe(II)
a = 5.18, b = 0.717 for Fe(III)
a = 1.68, b = 0.284 for Mn(II)
These equations are valid, respectively, from 10 –8 to 10 –5.3 , 10 –7.2 to 10 –3.3 , and 10 –5.9 to 10 –3.3 M.
The larger effect of Fe 2+ is probably related to the formation of dissolved Fe 3+ from the
rapid oxidation of Fe 2+ with oxygen. The oxidation of Fe 2+ provides a higher initial concentration of Fe 3+ than can be added from a stock solution of Fe 3+ (which may be locally
supersaturated). The peroxide generated by the oxidation of Fe 2+ ,
O 2
– + H + = HO 2
(10.83)
HO 2 + HO 2 = H 2 O 2 + O 2
(10.84)