Chemistry of inorganic compounds – Treating mixture to obtain metal containing compound – Platinum group metal
Reexamination Certificate
1999-10-29
2002-11-26
Bos, Steven (Department: 1754)
Chemistry of inorganic compounds
Treating mixture to obtain metal containing compound
Platinum group metal
C423S042000, C423S043000, C423S050000, C423S087000, C423S092000, C423S127000, C423S140000
Reexamination Certificate
active
06485696
ABSTRACT:
FIELD OF THE INVENTION
The present invention relates to a method for recovering and/or removing metallic elements from waste water.
BACKGROUND OF THE INVENTION
Environmental Hazard
Runoffs from a variety of industrial operations such as electrical power plants, steel plants, and mines are known to be contaminated with various metal compounds including iron, manganese, aluminum, zinc, copper, lead, arsenic, and chromium. These contaminants pose a serious environmental problem, as these runoffs cannot be safely discharged into the environment. Previously used methods to remove these contaminants involved adding lime, soda ash, or other neutralizing agents, and treating the runoff in a holding pond or clarifying tank. However, these methods have not been satisfactory because of the long periods of time required to effect treatment.
In particular, acid mine drainage from active and abandoned coal and metal mines is a serious environmental problem that affects thousands of miles of streams in the United States and elsewhere. Acid mine drainage, as well as mine tailings and refuse piles, also discolors streams. The metal sulfides that are part of many coal beds and ore deposits are oxidized when mining operations bring them within reach of oxidizing conditions created, directly or indirectly, by atmospheric free oxygen. This oxidation produces dissolved sulfate ion and metal ions. Because acidic wastewater, including acid mine drainage, discolors streams, damages ecological systems, and harms wildlife, federal and state limits for waste-water effluent require that the discharge pH be between 6 and 9, and that the concentration of common metallic elements such as iron and manganese be less than 2 milligrams per liter (mg/L) [cf. U.S. Environmental Protection Agency, 1994]. More stringent limits exist for toxic elements such as lead, arsenic, cadmium, and mercury.
Slow Oxidation by Oxygen and Extensive Ferric Iron Stains
Thermodynamically, aerated water should be capable of oxidizing ferrous iron (Fe
2+
) completely to ferric iron (Fe
3+
), with a equilibrium redox potential (also referred to as Eh when referenced to the standard hydrogen electrode, or ORP if not so specified) of:
O
2
+4H
+
+4
e
−
=2H
2
O
Eh=
1.23−0.059 pH+0.015 log (pO
2
) volts [Eq. 1]
This Eh value is shown as line [b] in
FIG. 1
, which is an Eh-pH diagram for the Fe—H
2
O system.
At sea level, pO
2
=0.21 bars, and [Eq. 1 reduces to:
Eh=
1.22−0.059 pH volts [Eq. 1′].
In acid solutions, the ratio of ferrous ion and ferric ion concentrations is related to Eh as
Fe
3+
+e
−
=Fe
2+
Eh=
0.77+0.059 log(Fe
3+
/Fe
2+
) volts [Eq. 2]
By equating [Eq. 2] with [Eq. 1′], one can readily see that, at equilibrium, only a trace of Fe
2+
should remain in solution, even when the pH is 3. When the pH increases, Fe
3+
is precipitated as Fe(OH)
3
:
Fe
3+
+3H
2
O=Fe(OH)
3
+3H
+
,
log(Fe
3+
)=4.84−3 pH pH unit [Eq. 3]
and Fe
3+
is precipitated as Fe(OH)
3
:
Fe(OH)
3
+3H
+
+2
e
−
=Fe
2+
+3H
2
O
Eh=
1.06−0.177 pH−0.059 log(Fe
2+
) volts [Eq. 4]
These equations show that, upon exposure to air or meteoric water saturated with air; iron should be precipitated as a solid phase, even in mildly acidic discharges, if true thermodynamic equilibrium were attained.
In reality, however, the Eh values of aerated waters are about 0.4 volts lower than the value calculated in [Eq.1′]. Sato [1960] measured the in situ Eh-pH values of mine waters at various depths, and found that, regardless of the type of deposits, the values in the oxidized zone were distributed within two parallel lines on an Eh-pH diagram, shown in FIG.
1
. These lines are defined by the reaction:
O
2
+2H
+
+2
e
−
=H
2
O
2
Eh=
0.68−0.059 pH+0.0295 log[(pO
2
)/(H
2
O
2
)] volts [Eq. 5]
The lower parallel line, marked [c] in
FIG. 1
, corresponds to the (pO
2
)/(H
2
O
2
) ratio of unity:
Eh=
0.68−0.059 pH volts [Eq. 5′]
The upper parallel line, marked. [d] in
FIG. 1
, corresponds to the ratio of 10
6
:
Eh=
0.86−0.059 pH volts [Eq. 5″]
In laboratory oxidation of both Fe
2+
and Mn
2+
solutions, Sato [1960] showed clearly that the rate of Eh increase dropped drastically at the Eh value of [Eq. 5′] and never went over the value of [Eq. 5″] even after prolonged aeration. A plausible explanation is that O
2
is somehow reluctant (i.e., a high activation energy barrier exists) to being split up in one step, and the faster reaction path is to form hydrogen peroxide as an intermediate. However, in the presence of iron, manganese, or a similar multivalent element, hydrogen peroxide is catalytically decomposed to oxygen and water, as discussed by Latimer [1952]. The result of this cyclic process is that oxygen becomes incapable of raising Eh beyond [Eq. 3″] in acidic solutions. Alternative paths may be provided by some aerobic microorganisms, but even with such help, the process of oxidation by oxygen is still relatively slow at surface temperatures.
The extremely slow rate of oxidation by oxygen beyond the Eh of [Eq. 3″] is the primary cause for the phenomenon of red iron stains formed for miles in the downstream direction from both treated and untreated acid mine discharge sites. When the pH is 2.5 to 4, ferric hydroxide can be precipitated by aeration alone, albeit slowly. The slow reaction ensures that a large fraction of the iron remains in solution as ferrous ion for a long time. Furthermore, the precipitation of iron as ferric hydroxide releases free sulfuric acid, which then becomes a secondary acidification step. Therefore, the pH often decreases downstream upon exposure to air even after neutralization treatment by anoxic limestone drains or by lime or caustic soda, which are bases used to bring the pH to more than 6 to meet discharge regulations.
The above pH range partially overlaps the range of active precipitation of gelatinous aluminum hydroxide (pH 4.5 to 6, Nordstrom and Ball [1986], and Hemingway [1982]; see
FIG. 2
) and aluminite, Al
2
(SO
4
) (OH)
4
.7H
2
O (pH 4.0, Robbins et al, [1996]). These poorly crystalline aluminum hydroxide and hydroxy-sulfate compounds and ferric hydroxide typically coat the limestone used in treatment, slowing down the neutralization process.
In
FIG. 1
, the range of pH was limited in this diagram to 1 to 11, because that is the pH range which is relevant to the present invention. The solid compound phases of iron are the most unstable hydroxide phases, Fe(OH)
3
and Fe(OH )
2
, because these are the phases that actually precipitate upon addition of base, and also upon oxidation in the case of Fe(OH)
2
. The shaded area around these hydroxides is the region of active precipitation of iron as a solid from its ionic state in aqueous solution. The border on the ionic side is defined by the activity (concentration) of 10
−2
molar (560 mg/L Fe) and that on the solid side by the activity of 10
−5
molar (0.56 mg/L Fe) of Fe
2+
or Fe
3+
ion. The horizontal line that crosses the diagram at an Eh value of 770 mV indicates equal activities of ferrous and ferric ions. The four parallel lines with a slope of 59 mV per pH marked with a bracketed letter are as follows:
[a] the standard-hydrogen electrode i.e., H
2
—H
2
O redox couple;
[b] the standard potential of the H
2
O—O
2
couple;
[c] the standard potential of the O
2
—H
2
O
2
couple; and
[d] the empirical limiting potential of oxid
Robbins Eleanora I.
Sato Motoaki
Hunt, Jr. Ross F.
The United States of America as represented by the Secretary of
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