EPA R? 79 flfiQ
NOVEMBER 1972 ENVIRONMENTAL PROTECTION TECHNOLOGY SERIES
Oxidation of Pyrites in Chlorinated
Solvents
Office of Research and Monitoring
U.S. Environmental Protection Agency
Washington, D.C. 20460
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RESEARCH REPORTING SERIES
Research reports of the Office of Research and
Monitoring, Environmental Protection Agency, have
been grouped into five series. These five broad
categories were"established to facilitate further
development and application of environmental
technology. Elimination of traditional grouping
was consciously planned to foster technology
transfer and a maximum interface in related
fields. The five series are:
1. Environmental Health Effects Research
2. Environmental Protection Technology
3. Ecological Research
4. Environmental Monitoring
5. Socioeconomic Environmental Studies
This report has been assigned to the ENVIRONMENTAL
PROTECTION TECHNOLOGY series. This series
describes research performed to develop and
demonstrate instrumentation, equipment and
methodology to repair or prevent environmental
degradation from point and non-point sources of
pollution. This work provides the new or improved
technology required for the control and treatment
of polluto.cn sources to meet environmental quality
standards..
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EPA-R2-72-069
November 1972
OXIDATION OF PYRITES IN CHLORINATED SOLVENTS
By
Joseph C. Troy
Joseph A. Boros
Donald R. Brenneman
Contract No. lU-12-897
Project 1^010 FMM
Project Officer
Dr. James M. Shackelford
Pollution Control Analysis Branch
Environmental Protection Agency
Washington, D.C. 20h60
Prepared for
OFFICE OF RESEARCH AND MONITORING
U.S. EWvTRONMENTAL PROTECTION AGENCY
WASHINGTON, D.C. 20^60
For sale by the Superintendent of Documents, U.S. Government Printing Office, Washington, D.C. 20402 - Price $1.00
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EPA Review Notice
This report has been reviewed by the Environmental Pro-
tection Agency and approved for publication. Approval does
not signify that the contents necessarily reflect the views
and policies of the Environmental Protection Agency, nor
does mention of trade names or commercial products constitute
endorsement or recommendation for use.
11
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ABSTRACT
The purpose of this study was to develop methods for ex-
tracting sulfur and iron compounds from pyritic waste mater-
ials, producing a final refuse that is incapable of causing
water pollution, and at the same time, conserving useful
mineral resources. The primary method under investigation
was the chlorination of pyrites in non-aqueous solvents.
An attractive feature of the proposed method was its cyclic
nature. Chlorine would serve as the original driving force
for the expected reactions, but would be replaced as pri-
mary oxidant by two products of the chlorination reaction,
ferric chloride and sulfur dichloride.
Test results indicated that the successful extraction of
sulfur and iron compounds depended upon the continuous
addition of chlorine gas to the system. The expected re-
placement of chlorine by ferric chloride and sulfur di-
chloride was not achieved under conditions of this study.
The reaction rates with chlorine gas varied with the parti-
cle size and source of the pyritic material; with the
choice of solvent; and with the means of providing inti-
mate contact between pyrites and chlorine-saturated solvents
Recycling solvents through beds of pyrite significantly
improved rates of reaction, while elevated temperatures and
treatment in the presence of ultra-violet radiation yielded
slight improvements in reaction rate. This report contains
20 references.
This report was submitted in fulfillment of Contract No.
14-12-897 under the sponsorship of the Water Quality Office,
Environmental Protection Agency.
111
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CONTENTS
Section Page
I Conclusions 1
II Recommendations 3
III Introduction ~"
IV Description of Test Apparatus and Procedures 9
V Results - Unchlorinated Solvents 1°
VI Results - Chlorinated Solvents 21
VII Recycling of Solvents "
VIII Other Pyrite Decomposition Tests ^5
IX Acknowledgements ^9
X References 51
v
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FIGURES
No. Page
1 Flow Diagram - Cyclical Process 7
2 Plastic Reaction Cell 11
3 Solvent Recycle System 12
4 Glass Reaction Cell (Modified) 13
5 Apparatus for In-Situ Reactions with Oxygen 14
6 Recycle Tests Using Benzene 26
7 Recycle Tests Using Carbon Tetrachloride 27
8 Recycle Tests Using Chloroform 28
9 Recycle Tests Using 1,2 Dichloroethane 29
10 Recycle Tests Using Dimethylformamide 30
11 Recovery of Iron Using Chlorinated Methanol 36
VI
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TABLES
Nos. Page
1 Comparison of Pyritic Materials 16
2 Unchlorinated Solvents 20
3 Chlorinated Solvents 22
4 Four-Hour Recycle Tests 31
5 Long-Term Recycle Run Using Chlorinated Methanol 33
6 Long-Term Recycle Run Using Chlorinated Methanol
with Electrolysis 35
7 Four-Hour Recycle Runs Using Chlorinated Solvent
Combinations 38
8 Four-Hour Recycle Runs at Elevated Temperatures 39
9 Four-Hour Recycle Runs Under Ultra-Violet
Radiation 41
10 Four-Hour Recycle Runs from Different Sources
and Sizes of Pyrite 43
11 Solubility of Pyrites in Chlorinated Solvents with
Abrasion 46
Vll
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SECTION I
CONCLUSIONS
1. Decomposition of pyrites in chlorinated solvents is a
possible means for converting certain solid wastes from
mining operations into usable sulfur and iron products.
Rates of reactions causing such decomposition are ad-
versely influenced by a variety of factors, including
the deposition of solid reaction products on unreacted
pyrite surfaces.
2. The organic solvents tested showed varying tendencies to
sustain reaction rates, due to differences in their
ability to dissolve the reaction products.
3. Pyritic materials from different sources varied in
their susceptibility to decomposition by chlorinated
solvents.
4. Reaction rates were significantly improved by recycling
chlorinated solvents through the pyritic mass.
5. Reaction rates were moderately improved by vigorous
abrasion of pyrite particles, thus continuously
exposing fresh surfaces to the chlorinated solvents.
6. Reaction rates were slightly improved under ultra-violet
radiation, and at temperatures higher than ambient.
7. All successful attempts at decomposition of pyrites de-
pended upon the continuous addition of chlorine to
the system. The desired reaction utilizing the ferric
chloride produced by the initial chlorination of pyrites
as the oxidant for continuing decomposition was not
achieved under conditions of this test. A continuous
chlorine consumption of this magnitude would place
severe economic limitations upon use of this process.
8. No measurable oxidation of pyrites or reaction products
was achieved using oxygen or air in non-aqueous solu-
tions, even under oxygen partial pressures exceeding
one atmosphere.
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SECTION II
RECOMMENDATIONS
Further study of the basic reactions involved in pyrite
decomposition is required. In particular, secondary reac-
tions wherein ferric and sulfur chlorides are utilized as
oxidants must be optimized. Otherwise, the continuous
chlorine demand of the system will render such treatment
techniques undesirable from an economic standpoint.
A system for classifying pyritic material according to
its susceptibility to decomposition by chlorination,
oxidation, acidification, or any other treatment method
should be developed. The parameters that cause one pyrite
to react while a similar pyrite is nearly inert should be
defined.
Improved means for applying metallurgical techniques to the
pyrite problem should be devised. These need not necessar-
ily require smelting at high temperatures. Pressure leach-
ing at temperatures between 125 and 210°C has proved
successful.
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SECTION III
INTRODUCTION
Whenever coal is removed from the earth, whether by surface
mining or deep mining, it is a necessary and common prac-
tice to separate unwanted materials from the coal. These
materials — shale, clay, pyrite, earth, coal fines, etc.—
are found in reject piles near most coal v/asheries, crush-
ing and screening stations and power plants, in addition to
active and abandoned coal mines throughout Appalachia.
The total weight of refuse associated with bituminous and
anthracite coal mining throughout the United States is
estimated in the billions of tons.
Such waste materials, on exposure to air, often ignite spon-
taneously, causing serious pollution of the atmosphere with
noxious fumes of sulfur dioxide. Even in the absence of
combustion, exposure of these refuse piles to air and rain-
fall over long periods of time leads to reactions which
generate iron sulfates and sulfuric acid in the drainage from
such areas. The poor quality of this run-off leads to con-
tamination of receiving streams in precisely the same manner
as acid mine drainage.
Processes have been proposed to recover coal fines from re-
fuse piles as a means of eliminating the hazard of spon-
taneous combustion. This reprocessing produces a final
refuse enriched in pyrites. If additional treatment methods
were found which could extract usable sulfur and iron com-
pounds from pyrite, a final refuse could be produced that is
neither combustible nor capable of causing water pollution,
and at the same time, the conservation of mineral resources
could be enhanced. The chlorination of pyrites in organic
solvents is proposed as the possible means for recovery of
sulfur and iron compounds.
The chlorination of pyrite in chlorinated carbon tetra-
chloride or dichloroethane was reported by McElroy and
Peters(D in 1968. They found that iron pyrite reacts with
chlorine dissolved in either solvent to produce ferric
chloride and sulfur dichloride at ambient temperatures
according to the following reaction:
2FeS2^ + 7C12 • rn14 ) 2FeCl3 + 4SC12 (1)
Both products of this initial reaction may in turn be used
as reactants with additional pyrites to produce ferrous
chloride and elemental sulfur as follows:
FeS2 + 2FeCl3 re1.1 4) 3FeCl2 + 2S° (2)
FeS2 + SC12 rn 4) FeCl2 + 3S° (3)
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The overall reaction of importance may be summarized as
follows :
FeS2 + C12 - > FeCl2 + 2S° (4)
The continuance of reaction (2) depends upon the successful
oxidation of some of the FeCl2 produced in reaction (4)
back to FeCl3 according to the following reaction:
12FeCl2 + 302 - ) 8FeCl3 + 2Fe2O3 (5)
Reaction (5) is critical from the standpoint of economics.
If the ferrous chloride generated by the other reactions
can be converted to ferric chloride and ferric oxide, this
ferric chloride may be recycled as in reaction (2) . The
overall reaction using (2) and (5) eliminates the use of
chlorine except as a starter to produce a sufficient in-
ventory of FeCl3 required for reaction (2) . In this way a
cyclical process (Figure 1) can be set up whereby pyrites
will be reacted in a non-aqueous media with ferric chloride
and sulfur dichloride , producing elemental sulfur for sep-
aration as a solid or slurry, and ferrous chloride for re-
cycling. The ferrous chloride produce may be oxidized to
yield iron oxides for separation as solids, and ferric
chloride for further reaction with additional pyrites.
Overall, if reactions (2) through (5) prove feasible, the
entire cyclical process can be summarized by the following
combined reaction:
4FeS2 + 302 - - - > 2Fe2°3 + 8S° (6)
From reaction (6) it may be inferred that one (1) ton of
oxygen may be reacted with five (5) tons of pyrite to
produce 3.33 tons of iron oxide and 2.67 tons of sulfur.
This study was designed to provide information on the above
reactions to determine their usefulness as a means for re-
covery of sulfur and iron compounds from coal mining refuse
material. Key questions to be answered include:
1. The solubilities of the principal reactants and reac-
tion products in selected chlorinated solvents at
ambient and elevated temperatures.
2. The effect on rates of reaction achieved by:
a. Using other non-aqueous organic solvent.
b. Using pyrites from different sources.
c. Using pyrites of various sizes and surface areas.
d. Using elevated temperatures (up to the solvent's
boiling point) .
e. Using ultra-violet radiation.
f. Using pure oxygen or air as the oxidant for ferrous
chloride .
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FeS-
Cl-
i
i
REACTION (1) STARTER ONLY
solvent
2FeS2+7Cl2
2FeCl3+4SCl2
FeS,
I
I
FeS-
FeClo inSolvent SC10 in Solvent
2
REACTION (2)
solvent
REACTION (3)
solvent
FeS+SCl
I U2 I
FeCl2 in Solvent I FeCl2 in Solvent
Sulfur ^_ X X Sulfur
FeCl3 in
Solvent
REACTION (5)
solvent
12FeCl
FLOW DIAGRAM
Iron Oxide
CYCLICAL PROCESS
FIGURE NO.
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SECTION IV
DESCRIPTION OF TEST APPARATUS AND PROCEDURES
Efforts to convert pyrites to usable forms of iron and sulfur
were limited to a series of bench scale tests in appropri-
ate laboratory glassware. The solubility measurements were
made by adding excess amounts of each reactant and reaction
product to separate portions of solvent in a constant temp-
erature bath. The bath temperature was raised slightly
then allowed to fall back to equilibrium at 25°C and at
selected higher temperatures, depending on the solvent.
Constant stirring was maintained throughout a 24-hour
period. Aliquots were withdrawn at the end of this time,
and analyzed for dissolved matter by evaporating the solvent.
Additional analytical tests were performed on residues after
evaporation to determine iron or sulfur content where
applicable.
The preliminary reaction, equation (1), between pyrite and
chlorine in non-aqueous solvents was evaluated by sparging
chlorine gas through a fine suspension of 200 mesh pyrite
in solvent. After 24 hours, aliquot portions were with-
drawn, filtered and analyzed for iron and sulfur in the
residues after evaporation of the solvents. The presence
of either or both elements in the filtered solvent was taken
as an indication that reaction (1) or some similar reaction
involving the decomposition of pyrites had actually occurred.
The pyrites rema'inina in suspension were allowed to settle
to the bottom of the reaction vessel. The excess solvent
was transferred to polyethylene containers. Within hours,
some specimens exhibited vapor pressures greatly exceeding
those expected from any entrapped solvents alone. Upon
visual examination these pyrites were exuding a fuming vapor
with the characteristic odor and color of sulfur dichloride,
one of the reaction products reported by McElroy and Peters.
Although this finding is encouraging from the standpoint
of confirming the success of reaction (1), the fact that
the sulfur dichloride tended to remain on the pyrite sur-
faces may indicate problems in effecting the transfer of
the reaction products from the sites where they form.
Since this transport problem arose early in the test pro-
gram, it was decided that some form of recycle system for
bringing chlorinated solvents to the reaction site, and
for carrying away the products of reaction had to be
provided. A first effort involved the use of a peristaltic
-------
pump and an acrylic reaction chamber with return lines
carrying solvent back to a reservoir. In the event that
later tests utilizing electrolysis proved beneficial, a
cover with permanently mounted electrode pairs was added
to the reaction cell. Figure 2 shows a sketch of the orig-
inal cell. Solvent was pumped through the pyrite bed at
100 ml/minute, collected in the reservoir, then pumped
through the cell again. In this manner, any beneficial
effects from FeCl3 or SC12 formed by the initial chlor-
ination of pyrites would be utilized by recycling.
Although this system as designed worked well with certain
solvents, e.g., CC14/ the plastic parts deteriorated rapid-
ly in the presence of other chlorine-saturated solvents.
Also, the tubing in the peristaltic pump had poor life
expectancies, sometimes rupturing at the flexing point dur-
ing relatively short recycle runs. For these reasons, the
entire recycle system was modified as shown in Figures 3
and 4. The plastic reaction cell was replaced with a 1"
I.D. x 6" long heavy-walled glass tube with polypropylene
end caps. The peristaltic pump was replaced by a reversible
motor operating a traveling reservoir to provide gravity
flow of solvents through the pyrite bed in the reaction
cell. The motor was operated at a rate of speed which
caused solvent to flow through the cell at 100 ml/minute.
This apparatus functioned well through the remainder of
the tests, requiring only periodic replacement of the flex-
ible tubing connecting the traveling reservoir to the reac-
tion cell. Elevated temperatures were attained by wrapping
a heating tape around the reaction cell. A polypropylene
adaptor was designed to accomodate an ultraviolet lamp
for use during that phase of testing. In all recycle tests,
the objective was to measure the weight loss of pyrites
during the test run, and to determine iron and sulfur con-
centrations from aliquots of the recycling solvents as
the test progressed.
The in-situ oxidation of pyrites using pure 02 was studied
in a 1000 ml reaction flask fitted for gas-tight opera-
tion as shown in Figure 5. The flask was equipped with
a stirrer and with inlet ports to admit pyrite, chlorinated
solvents and oxygen under constant pressure. An initial
gas pressure of 0.433 PSI (12 inches of water) was used
for most tests. The objective was to monitor the utili-
zation of oxygen in terms of volumetric displacement and
pressure changes.
10
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PYC. SCREEN-
o
o
'|6 ACRYLIC FRAME WITH-
20-'8" DIA. HOLES
c/1
O
"4" ACRYLIC
4-ELECTRODES
SECTION-AA
COVER 8 ELECTRODE HOLDER-
PYRITE BED
C-3
tT]
»T-:H
+—i-
A
RETURN
'4" ACRYLIC
TUBE (TYR)
SOLVENT AND
CHLORINE INLET
PLASTIC REACTION CELL
FIGURE NO. 2
11
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VENT-
CLAMP
GLASS REACTION
CELL(DETAIL IN
FIG. 4)
-FIXED RESERVOIR FOR SOLVENTS
-PULLEY
CONNECTION TO CHLORINE GAS
SOURCE
NYLON CORD
PULLEY
STOP
TRAVELING^
RESERVOIR
FOR
MERCURY
SWITCH
REVERSIBLE
MOTOR
REVERSING CONTROL-
SOLVENT RECYCLE SYSTEM
FIGURE NO. 3
12
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ELECTRODE
CONNECTION
POROUS STONE SPARGER
FOR CHLORINE INLET
THREADED-
CONNECTOR
ELECTRODE CONNECTION
CONNECTION TO
FIXED RESERVOIR
THREADED CONNECTOR
I I.D. GLASS TUBING
50 GRAMS OF PYRITE
CONNECTION TO
TRAVELLING RESERVOIR
CHLORINE GAS INLET
GLASS REACTION CELL (MODIFIED)
FIGURE NO. 4
13
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GAS
DISPLACEMENT
CYLINDER
HEAD TANK
FROM GAS
CYLINDER
SUPPLY
1000 ML.
REACTION
FLASK
APPARATUS FOR IN-SITU REACTIONS
FIGURE NO. 5
WITH OXYGEN
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Analytical Procedures Used to Identify Products
No satisfactory procedure for analyzing the iron or sulfur
reaction products directly in aliquots of the chlorine
saturated solvents was found. Each determination was made
on residues after evaporation of the solvents, with the
attendant possihility of changes in the oxidation state.
For this reason, most determinations were reported as total
iron and total sulfur. Iron was measured either using the
0 - phenanthroline colorimetric method ^ ' or the atomic
absorption technique (3) after extraction of the dried
residue with hot hydrochloric acid. Sulfate was found
using the turbidimetric method ^ ' . Elemental sulfur in
residues was identified by a technique described by GyenesH)
and by Kucharsky and Safarik^5). This latter method is less
well-known, so it is described in more detail. Extract
the sample in a vibrator with 80 ml of benzene for one
hour. Filter the solution into a 100 ml volumetric flask.
Dilute to the mark with acetone. Mix 20 ml of diluted sam-
ple with 1 ml of water and 3 drops of bromothymol blue indi-
cator. Titrate with 0.05 M KCN solution in isopropanol,
taking care to maintain a temperature of 50 - 60°C through-
ought the titration. The end-point is attained when a blue
color persists. Warm the sample and if the color changes
from blue to yellow, add titrant until the color change is
permanent. The KCN solution is originally standardized
against a pure elemental sulfur solution containing 1.6033g
of purified, dried sulfur in 800 ml of benzene, then diluted
to 1 - liter with acetone. The same titration procedure
is used as in determinina the unknown sulfur concentration.
The analyses on solid pyrite samples, both before and after
reaction with chlorine-saturated solvents, were performed
using x-ray spectrographic methods for determining element-
al compositions and x-ray diffraction to identify crystalline
structures. Surface area measurements for selected pyrite
specimens were performed by the Materials Analysis Laboratory
of Micromeritics Instruments Corporation, following the
standard multipoint B.E.T. technique (b) using krypton gas
adsorption on an Orr surface area pore-volume analyzer.
The pyrites chosen for these tests were obtained from two
sources: coal pyrites from the Shawville Power Station of
the Pennsylvania Electric Company near Clearfield, Penn-
sylvania; and pyrite concentrate from the Metal Mining Divi-
sion of Kennecott Copper Corporation's concentrate plant
near Salt Lake City, Utah. The Shawville Coal pyrite showed
a weathered appearance and fragile structure, and was
brought to the surface at least two years ago. The Utah
pyrite concentrate, on the other hand, was smooth and hard,
being above ground less than two (2) months. Both samples
were chemically similar, but varied considerably in physical
structure. A comparison of the two pyritic materials appears
in Table 1.
15
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TABLE 1
All results in per cent, except as noted
Shawville Utah Pyrite
Coal Pyrite Concentrate
Iron 30.8 33.6
Sulfur 31.2 36.8
Silica as Si02 1-1 2.0
Calcium 0.6 0.5
Magnesium 1.1 0.6
Aluminum 0.7 0.4
Copper 0.4 0.7
Crystalline Matter 40% FeS- 70% FeS2
Loss on Ignition 36.0 33.0
Specific Gravity 3.70 5.05
Surface Area (-60+200 mesh frac.)3.28M2/Gram 0.17M^/Gram
16
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Chemically, both pyrites showed minor differences, but the
Utah concentrate contained nearly twice as much crystalline
material as the Shawville. This difference was also
reflected in the respective specific gravities, indicating
that Shawville was typical of coal pyrites, where Utah
was more similar to "museum-grade" pyrites. Both values
agree with those reported by Clark ('"), 3.2-3.4 and 5.0
for coal pyrites and "pure" pyrites respectively. Gross
differences in surface areas were also cited by Clark,
and by Braley^8), who reported surface areas of 4.4 and
0.4M^/gram for sulfur ball pyrites and "museum-grade"
pyrites respectively.
17
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SECTION V
RESULTS
Solubility Measurements - Unchlorinated Solvents
Ten grams of each solid material (ferric chloride, elemental
sulfur, and Shawville pyrite) in powder form were placed in
separate 400 ml Erlenmeyer flasks. The pyrites were pre-treated
by soaking in 0.3N HC1 overnight, rinsing 5 times with de-
ionized water and 2 times with acetone, then drying for 4
hours at 40°C. A 150 ml portion of each solvent under test
was added and the flasks were stoppered to prevent evapora-
tion. The mixtures were shaken carefully, and maintained
at a constant temperature of 28°C for 18 hours. More ferric
chloride or sulfur was added as required. When no further
additions of solutes could be put into solution, the temper-
ature was allowed to fall back to 25°C and was maintained
at this level for 24 hours. Continuous agitation was pro-
vided throughout the period. A 50 ml portion was withdrawn
from each sample and drawn through an 0.45 micron membrane
filter to remove undissolved matter. The solvent was driven
off by careful evaporation in a tared distilling flask,
and the residue on evaporation was weighed. P. solvent blank
correction was applied by carryina 50 ml portions of each
pure solvent through the sam.e procedures, then subtracting
the weight of residue thus obtained, if any, from the
corresponding solvent/solid mixture. Solubilities are
expressed in Table 2 as grams of solute/100 grams of solvent
at 25°C, following the application of appropriate factors
based on each solvent's specific gravity.
19
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TABLE 2
UNCHLORINATED SOLVENTS
Solvent
Grams of Solute/100 Grams of Solvent @ 25°C
FeCloSShawville Pyrite
Acetone
63*
2.1*
0.061
Benzene
0.42
1.0*
JO. 95 0.027
^50*
Carbon
Carbon
Disulfide <0.01 )48.8 0.018
Tetrachloride Insoluble
Chloroform 1.54
1,2 Dichloroethane 0.29
Dimethylf ormamide 20
Methanol 155*
0.831*
[0.82 Insoluble
, — '
.1.23 0.019
0.66 0.019
<0.001
0.3*
0.31 0.370
Nitrobenzene
Sulfur Monochloride
1.70 0.09 0.007
0.1 Very Soluble* Reacts
*Data reported in various solubility tables
values shown were measured experimentally.
All other
20
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SECTION VI
RESULTS
Solubility Measurements - Chlorinated Solvents
Attempts to repeat tile above experiments using chlorine-
saturated solvents proved to be more difficult because:
1. Certain solvents reacted with chlorine to form differ-
ent solvents or mixtures, e.g., sulfur monochloride
was chlorinated to sulfur dichloride.
2. Elemental sulfur reacted with chlorine in most solvents,
changing to sulfur chlorides rather than remaining in
elemental form.
3. Residues on evaporation of chlorinated solvents con-
taining excess sulfur and pyrites rarely resembled the
original solutes, indicating that reactions had occurred,
rather than simply reaching a solution equilibrium.
Confirmation of this observation was made by repeating
the measurements on pyrite/solvent mixtures two weeks
after the original solubility measurements were made.
Table 3 shows this comparison.
All solvents in Table 3, except carbon disulfide, and possi-
bly dimethylformamide continued to react slowly after the
24 hour measurements were made. CS2 and DMF reaches some sort
of equilibria in 24 hours, although both solvents changed
color slightly upon chlorination, even prior to the addition
of pyrites. Methanol had a twelve-fold greater ability to
"dissolve" pyrite in the unchlorinated stage (see Table 2),
indicating a reaction between chlorine and this solvent, also.
Two of the 15 day specimens, benzene and 1,2 dichloroethane,
showed evidence of the formation of elemental sulfur and
sulfur chlorides, products of a reaction and not a true solu-
tion. Although the quantities of elemental sulfur determined
by the technique described by Gyenes'^) were very small
(0.020 and 0.035 grams of sulfur/100 grams of solvent for
benzene and 1,2 dichloroethane respectively), this sulfur was
present as a finely-divided solid. Table 2 indicates sulfur
solubilities 20 to 50 times greater than the concentrations
reported. Even if one assumes that the sulfur was formed
from reactions between FeCl3 or SC12 and FeS2 as in equations
(2) and (3) after all Cl2 was depleted, the quantities of
sulfur thus formed would be expected to dissolve in the sol-
vents. Otherwise, such reaction products could become
rate-limiting by adhering to unreacted pyrite surfaces.
21
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TABLE 3
CHLORINATED SOLVENTS
Grams of Residue/100 Grams of Solvent @ 25°C
Solvent 24 Hour Contact Time15 Day Contact Time
Acetone 0.070 0.14
Benzene 0.052 0.48
Carbon Disulfide 0.036 0.036
Carbon Tetrachloride 0.074 0.22
Chloroform 0.048 0.23
1,2 Dichloroethane 0.13 0.59
Dimethylformamide 0.001 0.001
Methanol 0.03 0.11
Nitrobenzene 0.05 0.13
22
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Since the pyrite "solubility" tests had become in fact reac-
tion rate tests, it was decided at this point to evaluate
equation (1) using <200 mesh Shawville pyrite in chlorine-
saturated carbon tetrachloride. Twenty grams of pyrite fines
were added to 300 ml of chlorine-saturated CC14 in a 1000-ml
reaction flask. Chlorine gas was continuously sparged through
the mixture at a rate of 15ml/minute. After 24 hours, two
100-ml portions of the solvent were withdrawn and filtered
through 0.45 micron membrane filters. The filtrates were
evaporated to dryness and dissolved solids concentrations
were determined. The two aliquots showed good agreement
at 0.1189 and 0.1196g/100 grams of solvent. These residues
in turn were separately analyzed for total iron and total
sulfur, following digestion in hydrochloric acid. Quantities
found in this residue were 0.0146 grams of iron as Fe, and
0.048 grams of sulfur as SO^. If the iron were present
entirely in the FeCl3 form, the 0.0146 grams would be equiv-
alent to 0.0425 grams of FeCl3 produced. .And if all of the
sulfur were present as SC12 , the 0.048 grams as SO4 would
represent 0.0447 grams of SC12. From equation (1) two (2)
moles of SC12 are produced for each mole of FeCl3 , and the
ratio between the two theoretical products above was found
to be 1.05:1, indicating that other forms of iron or sulfur
are present, not simply FeCl3 and SC12. In addition, the
sum of the theoretical weights above is 0.0872 grams, leaving
a weight of 0.032 grams unaccounted for in the residue on
evaporation. At least, the presence of iron and sulfur
compounds in the filtrate was an indication that pyrite
decomposition had actually occurred.
The unreacted pyrites remaining in suspension were recovered
by decanting the excess solvent, and drying in a warm air
current. An attempt was made to weigh this material, but
constant weight was not achieved at first. The material was
transferred to a capped polyethylene bottle for storage
overnight, but within a few hours the sides of the bottle
bulged under a vapor pressure far in excess of that ex-
pected from either the reactants or the chlorine gas which
may have permeated the porous pyrite bodies. Upon removing
the bottle cap, a dark red fuming vapor with a character-
istic sulfur chloride odor escaped to the air. The fact
that this reaction product was not removed from the reaction
site by the carbon tetrachloride raised questions about the
rate-limiting capabilities of this product.
23
-------
SECTION VII
RECYCLING OF SOLVENTS
At this point in the project it was decided to pursue the
problem of transporting reaction products from the pyrite
surfaces, and thus continually expose fresh surfaces for
further reaction. One approach tested involved continu-
ous recycling of the solvent. If product solubilities were
sufficient, the flowing solvent stream would wash away the
deposits and enable the reaction to continue at its maximum
rate. In addition, by recycling the solvents, any benefits
to be derived from the oxidation of pyrite by FeCl3 or SCl^
would simultaneously be attained. Since carbon tetrachloride
showed relatively low solvating capabilities on tests per-
formed thus far, recycle tests were run using a variety of
chlorinated solvents and mixtures of solvents.
A typical test run began with an initial charge of 50.0
grams of Shawville pyrite in the glass reaction cell (Fig-
ure 4). A 500ml charge of the solvent to be tested was
placed in the traveling reservoir, then forced to flow back
and forth through the pyrite bed at a rate of lOOml/minute.
Chlorine gas was introduced to the reaction cell via a
porous stone sparger at a rate of 25-27ml/minute. At 1, 2
and 4 hours into the test, 50 ml aliquots of recycling sol-
vent were withdrawn and analyzed for total iron concentra-
tion. The solvent volumes thus withdrawn were replaced
with equivalent volumes of pure solvent. At the end of four
hours, the pyrites remaining in the chamber were rinsed with
pure solvent, washed with deionized water, and dried. The
total weight lost by the pyrites was determined. Runs for
the most promising solvents were repeated with the addition
of one (1) part dimethylformanide (DMF) to nineteen (19)
parts of solvent, and with this solvent pair under ultra-
violet light. Graphs for four (4) hour test runs with
benzene; carbon tetrachloride; chloroform; 1,2 dichloroethane;
and dimethylformamide are shown in Figures 6 through 10,
and four-hour pyrite weight losses for these solvents and
for other, less successful runs are listed in Table 4
which follows.
The behavior of DMF was most interesting. Although a poor
solvent when used alone, it improved the solvating
capabilities of most organics significantly when blended
in at one (1) part DMF:19 parts solvent. In the case of
1,2 Dichloroethane the pyrite decomposition rate was changed
from the poorest to one of the most effective. Tests were
repeated using 1:3 and 1:1 blends of DMF and 1,2 Dichloro-
ethane, but increasing DMF concentration led to decreasing
pyrite decomposition rates. The pyrite weight losses were
4.48 and 2.84 grams respectively.
25
-------
RECYCLE TESTS USING BENZENE
ORIGINAL WEIGHT OF SHAWVILLE PYRITES
= 50 GRAMS
ILL
-
...1
- T-
> WEIGHT LOST BY PYRIT]
FROM SOLVENT
CgHg+DMF (19:1) & UV LIGHT
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jPPH^)i)^^^|
O
H
EH
ffi
W
0
I 2 3
RECYCLE TIME IN HOURS
FIGURE NO. 6
26
-------
RECYCLE TESTS USING CARBON TETRACHLORIDE
4 |
C3
EH
O
H
w
ORIGINAL WEIGHT OF SHAWVILLE PYRITES
» 50 GRAMS
: WEIGHT LOST BY PYRIC
: WEIGHT OF IRON RECO^
; FROM SOLVENT
-j-OL Mil: 1 1 | | ,| 1 | , | | [iM-J^MMffF
£ CC14 ~ DMF (19:1) & UV LIC
^i:^:g::::|:-|::::ll|::|:P
2 2
g|i-ij:ip-|'"T "|;: = ^ E::|":::|:
2
' ill .. Li-i. J. ,.]
,E .EEE;;EE|E|E|E:::EEEE;EEEEE-:EE;
/ERED ::::|:|g:--^|:ffi|:"::":
- IMMN i — -"[i:??:<' f 1^::::::i:::::::
W^zS::::-^-'!::^!::::::::^:^:
3HT— f:p::;!-z?:|:|:"Efc:"|:g:p:|
:|:'>s=i CC1 + DMF (19:1) :
4-j-I p • ft
1
3 4
RECYCLE TIME IN HOURS
FIGURE NO. 7
27
-------
RECYCLE TESTS USING CHLOROFORM
ORIGINAL WEIGHT OF SHAWVILLE PYRITES
= 50 GRAMS
± TATFTnHT1 T.nQrp nv nVDTT'-n'C
i --•
1
i CH(
*
T.
E: c
4_j- J
-+i-5
~lt
. _(_^ — L
Hrr
:i^S
-T —
WEIGHT OF IRON RECOVEI
FROM SOLVENT
-|ilt^-Li--±— ^--^i:|J— "g%
:13+DMF (19:1) &':UV LIGHT-
HC13+DMF (19:l)-: = = p::::|:::|:
EE;||B;;EEEEE|EE;|EEEEEEE;EEiE^EE;EE^]z?
wH^ttfrnffl i l| }fMm^
-;S ±-- -± -^-/±^ — \
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P T IcFl IT i 1 WtyffllT lillfH^t
_- -^ JJ_,! i -t- _^i L [ .__
i
*ED :::::::::::::::::::::::::::::::::::
•* "'- :~~z::^~: :":::::::: ::"~:::E::
\y\\\ mil
:::::::::^:::::::::::::::::|::|::::|::
L_ _L_
::::::||::L::::::::|:::::::±:::::|
_ . ... _ . J . . . .
. ^f J~__
_a"
; •"-11 ?-:-":=-
J UilU- CHC13 '
i^EiEEiiEEEiEEEEiiEEEEEEEEiiiEEEEEE;
4
EH
ffi
a
H
0
4
RECYCLE TIME IN HOURS
FIGURE NO. 8
28
-------
RECYCLE TESTS USING 1, 2 DICHLOROETHANE
ORIGINAL WEIGHT .OF SHAWVILLE PYRITES
=50 GRAMS
o
z
EH
ffi
WEIGHT LOST BY PYRITES
WEIGHT OF IRON RECOVERED
FROM SOLVENT
CH3CHC12+DMF(19:1) & UV LIGHT
CH3CHC12 + DMF (19 :
CH3CHC12
IN HOURS
0
RECYCLE TIME
FIGURE NO. 9
29
-------
RECYCLE TESTS USING DIMETHYLFORMAMIDE
ORIGINAL, WEIGHT OF SHAWVILLE PYRITES
= 50 GRAMS
WEIGHT LOST BY PYRITES
WEIGHT OF IRON RECOVERED
FROM SOLVENT
nttfttitfi+]
EH
a
o
0
DMF & UV LIGHT i
2 3
RECYCLE TIME IN HOURS
FIGURE NO. 10
30
-------
TABLE 4
FOUR-HOUR RECYCLE TESTS
Shawville Pyrite Weight Loss in Grams
(original weight = 50.0 grams)
Solvent Used
Solvent
Alone
Solvent + DMF
19:1
Solvent + DMF +
UV Light
19:1
Benzene 2.80
Carbon Tetra-
chloride 0.72
Chloroform 1.92
1,2 Dichloro-
ethane 0.28
DimethyIforma-
mide 0.56
Methanol 3.42
Nitrobenzene 1.72
3.72
4.40
5.08
4. 88
0. 84
1.08
3.78
4.56
5.32
5.20
31
-------
In reviewing the graphs in Figures 6 through 10, it is
apparent that recovery of iron proceeds in a linear fashion,
especially after the first hour of recycle. At the end of
each 4-hour run, the final concentration of iron in the
solvent ranged between 36 and 52% of the pyrite weight lost
during that time. I-f all of the iron in pure FeS2 were
recovered, the iron would weigh 46.6% of the total weight.
But for the Shawville pyrites used for these recycle tests
the original iron content was only 30.8%. One possible
explanation for this observation may involve possible
weight gains by the "unreacted" pyrites weighed at the end
of a run to determine weight losses. The surface area of
Shawville pyrite is so great (3.28M2/gram) that it is very
difficult to thoroughly dry the sample. Solvents, reaction
products, or moisture from exposure to air all may be en-
trapped and weighed as pyrites. Other approaches to moni-
toring the reaction rate while continuing the tests were
considered. Singer(9) determined FeS2 decomposition rates
by measuring the sulfate uptake in main water, but in our
tests, sulfur would be present in too many other forms. In
Sato's work(10) he reported polysulfides, thiosulfates,
polythionates, sulfites, and elemental sulfur in addition to
the formation of sulfates as intermediate sulfur states.
Sherman and Strickland(H) determined dissolution rates in
acid chlorine solutions by monitoring the consumption of
chlorine gas, but for our purposes, interactions between
chlorine and the various solvents make this approach less
usable. After considering these alternatives, it was de-
cided to continue to use the recovery of iron as a measure
/ -1 O \
of pyrite decomposition, as McKay and Halpern ^-L ' did in
their kinetic study of pyrite oxidation rates in aqueous
suspensions.
Long-Term Recycle Runs
In an effort to decompose the major portion of a 50 gram
Shawville pyrite mass, long-term recycle runs were attempted,
using methanol as the solvent due to its ability to dissolve
FeCl3 readily. Other than the change in test duration, all
conditions were the same as those for a typical 4-hour run,
except that iron concentrations would be monitored hourly
during the normal working day. Results obtained are shown
in Table 5.
32
-------
TABLE 5
LONG-TERM RECYCLE RUN USING CHLORINATED METHANOL
Net Gain in Overall Iron
Elapsed Grams of Iron Grams of Iron Recovery Rate
Time in Hours in Solvent Within Past Hour Grams/Hour
1 1.1 1.1 1.1
2 2.4 1.3 1.2
3 2.9 0.5 1.1
4 3.4 0.5 0.9
5 3.8 0.4 0.8
6 4.1 0.3 0.7
7 4.4 0.3 0.6
8 4.8 0.4 0.6
9 5.2 0.4-0.6
10-24 9.3 O.S^hr.^ 0.4
25 9.4 °'1r^
26 9.4 0.0^
27 9.5 0.1©
28 10.2 0.7 0.4
29 11.2 1.0 0.5
30 11.8 0.6 0.5
31 12.2 0.4 0.5
32 12.4 0.2 0.4
Chlorine addition and recycling continued throughout the
night.
Chlorine shut-off for three (3) hours. Overall iron
recovery rate based only on those hours of chlorine
addition.
33
-------
Original Weight of Pyrites = 50.0 Grams
Final Weight of Pyrites = 21.4 Grains
Net Loss in Weight = 28.6 Grams
Since the iron content of Shawville pyrite was found to be
30.8%, there were 15.4 grams of iron available in the
50 gram sample. At the end of the 32-hour run, 12.4 grams
of iron was found in the solvent, indicating an 80.6% re-
covery of iron, while decomposing 57.2% of the pyrite.
Another 30+ hour recycle run was planned duplicating the
earlier run, except that continuous electrolysis was pro-
vided throughout the run. In small scale tests, electrolysis
tended to enhance pyrite decomposition rates for certain
solvents. As it developed in this test, electrolysis in-
hibited the reaction considerably. After the unreacted
pyrites were removed from the reaction cell at the conclusion
of the test, the reason became apparent. A fine iron sus-
pension was tightly attached to pyrite particles throughout
the bed, reducing available surface areas to a fraction of
their original size. Results for this test are shown in
Table 6, and together with the previous run without elec-
trolysis, are presented in graph form in Figure 11.
34
-------
TABLE 6
LONG-TERM RECYCLE RUN USING CHLORINATED METHANOL
WITH ELECTROLYSIS
Net Gain in Overall Iron
Elapsed Grams of Iron Grams of Iron Recovery Rate
Time in Hours in Solvent Within Past Hour Grams/Hour
1 0.3 0.3 0.3
2 0.7 0.4 0.4
3 0.9 0.2 0.3
4 1.1 0.2 0.3
5 1.3 0.2 0.3
6 1.5 0.2 0.3
7 1.7 0.2 0.2
8 2.0 0.3 0.3
9 2.3 0.3 0.3
10-24 2.6 0.02/hr.* - *
25 2.8 0.2 0.3
26 3.0 0.2 0.3
27 3.3 0.3 0.3
28 3.8 0.5 0.3
29 4.3 0.5 0.3
30 4.9 0.6 0.3
31 5.3 0.4 0.3
32 5.5 0.2 0.3
33 5.6 0.1 0.3
* No chlorine added overnight. Overall iron recovery
rate based only on those hours of chlorine additon.
35
-------
RECOVERY OF IRON USING CHLORINATED
METHANOL
WITH ELECTROLYSIS
j-i-UJ-aUJ±amill I! i IthittrHtffl-M-
X * WITHOUT ELECTROLYSIS
CHLORINE
SUPPLY -
SHUT-OFF
TURNED ON
CHLORINE
SUPPLY
SHUT-OFF
CHLORINE
SUPPLY
TURNED ON
10 20
ELAPSED TIME IN HOURS
FIGURE NO. 11
36
-------
Original Weight of Pyrites = 50.0 Grams
Final Weight of Pyrites = 38.6 Grams
Net Loss In Weight. = 11.4 Grams
Once again there were 15.4 grams of iron available in the
50 gram pyrite sample, but at the end of 33 hours, only
5.6 grams were recovered. Thus, iron recovery was 36.4%
while pyrite decompositon decreased to 22.8%.
Both curves in Figure 11 show the effect on pyrite decom-
position rate when the fresh chlorine gas supply is shut
off. Iron recovery rates level off to 0.02-.03 grams/hour,
less than 10% of their normal rates with chlorine addition.
Upon resumption of chlorination, both tests showed a surge
in iron recovery before returning to the 0.3-0.5 grams/hour
level.
Recycle Tests Using Combinations of Solvents
Since the addition of one (1) part DMF to nineteen (19)
parts of various solvents did improve pyrite decomposition
rates significantly, a limited number of pairs of solvents
were tested in the recycle system. The tests were necessar-
ily limited, since the number of possible combinations,
and the ratios between pairs of solvents within each com-
bination provide possibilities outside the scope of this
project. In addition, some tests were run with added ele-
mental sulfur, as recommended by McElroy and Peters during
investigations of sulfur chlorides as the principal chlor-
inating and solvating agent in the decomposition of pyrites,
Fifty gram portions of pyrite were reacted with the solvent
combinations over a 4-hour recycle period. Results appear
in Table 7.
Recycle Tests at Elevated Temperatures
The four-hour recycle tests were performed at higher temper-
atures on six of the chlorine saturated solvents listed in
Table 4. The temperatures selected for each reaction were
dependent upon the solvent boiling point. Results were
compared with earlier tests, and are summarized in Table
8 following. The three best tests using DMF additives were
also reported.
Benefits derived from elevated temperature recycle runs
were somewhat less than anticipated. McElroy and Peters
reported that the chlorination rate increased 2.5 times
37
-------
TABLE 7
FOUR-HOUR RECYCLE RUN USING
CHLORINATED SOLVENT COMBINATIONS
Solvent Combination
Additive
Shawville Pyrite
Weight Loss in Grams
This Best Previous Test
Test With Either Solvent
Acetone/1,2 Dichloro-
ethane (1:1)
CS^/Nitrobenzene (1:1)
CC14/DMF (19:1)
CC14/DMF (19:1)
CC14/S2C12/DMF
(17:2:1)
CC14/S2C12/DMF
(17:2:1)
5 g S°
2.5 g S°
5 g S°
2.5 g S°
0.404
0.836
3.22
4.16
1.88
3.68
0.52
1.
4
78
40
4.40
4.40
4.40
In every example above, some previous run using solvents
alone or in 19:1 ratios with DMF yielded better pyrite
decomposition rates than the combination tested. Optimiz-
ing the possible combinations and concentrations is an
area for further exploration.
38
-------
TABLE 8
FOUR-HOUR RECYCLE RUNS AT ELEVATED TEMPERATURES
Solvent
Boiling Test
Point Temp.
C° C°
Shawville Pyrite
Weight Loss in Grams
Hi-Temp.Ambient
Benzene 80.1
Carbon Tetrachloride 76.8
Carbon Tetrachloride/
DMF (19:1)
Chloroform 61.2
Chloroform/DMF (19:1)
1,2 Dichloroethane 83.7
1,2 Dichloroethane/
DMF/ (19:1)
Methanol 64.7
Nitrobenzene 210.9
60°
60°
60°
50U
50°
60"
60°
50°
90°
3.04
0.96
5.08
2.28
6.16
0.44
4.56
4.88
0. 82
2.80
0.72
4.40
1.92
5.08
0.28
4.88
3.42
1.72
39
-------
with CCl^ as the solvent when the temperature was raised from
25°C to 41°C. Only about half of that rate of increase was
observed during our 4-hour recycle test at 60°C. Nelson^13),
et al, reported that oxidation rates for pyrites increased
as temperature increased, but chlorination tended to increase
as temperature decreased. The reaction may be dependent on
the amount of chlorine which can be dissolved in solvents,
a quantity which becomes smaller as temperatures go higher.
For maximum benefits from increased temperatures, pressure
leaching(^' or roasting techniques (15; are used.
Recycle Tests Under Ultra-violet Radiation
During earlier tests it was found that even the best of the
decomposition rates using nineteen (19) parts of solvent
with one (1) part DMF could be improved slightly in the
presence of an ultra-violet radiation source. Since no
prior evaluations were made on solvents and UV-radiation in
the absence of DMF, four such recycle runs using different
chlorine-saturated solvents were performed. The chlorinated
solvents chosen were benzene, carbon tetrachloride, chloro-
form, and 1,2 dichloroethane. All tests were run at ambient
temperatures using 50 grams of Shawville pyrite and a
four-hour recycle time. Findings are listed in Table 9
following.
The introduction of ultra-violet radiation had slight effect
on reaction rates using the more efficient solvents, benzene
and chloroform, increasing the rate of pyrite decomposition
by 6 and 17%. For less efficient solvents, carbon tetra—
chloride and 1,2 dichloroethane, the improvement was sig-
nificant at 86 and 186% respectively. Despite these size-
able gains, the less efficient solvents with UV-radiation
still tended to decompose pyrites at a slower rate than the
more efficient solvents without UV-radiation.
Effect of Particle Size, Surface Area, and 'Source
All tests performed thus far, both static and recycle, were
run using Shawville coal pyrites from one homogeneous sample
ranging in particle size downward from 1/8" to dust. To
evaluate the effect of particle size, surface area and source,
the samples of Shawville pyrites and Utah pyrite concen-
trates were sieved into two fractions: -18+60 and -60+200.
Samples of the finer of these two fractions (-60+200) were
analyzed for surface areas. These and other chemical and
plysical parameters were shown in Table 1. Eight 4-hour
recycle runs using 50 gram portions of each size fraction
and from each source were run using chlorine-saturated car-
bon tetrachloride and chloroform. Data from these eight runs
40
-------
TABLE 9
FOUR-HOUR RECYCLE RUNS UNDER ULTRA-VIOLET RADIATION
Shawville Pyrite
Weight Loss in Grams at 25° C
Solvent Without UVWith UV
Benzene 2.80 2.96
Carbon Tetrachloride 0.72 1.34
Chloroform 1.92 2.26
1,2 Dichloroethane 0.28 0.80
41
-------
are summarized in Table 10, together with results obtained
previously on the homogeneous Shawville sample. The Utah
concentrates did not have sufficient +18 sieve material
to run that comparison.
From the data it is apparent that Utah concentrates were
more easily decomposed than Shawville coal pyrites. Only
Shawville fines in chloroform approached the reaction rate
of the western pyrites. This finding is especially inter-
esting, since the surface area of Shawville pyrite is nine-
teen (19) times greater than the surface area of Utah
concentrates. Many researchers(7) (8) (13) d6) have found
direct correlations between surface area and reaction rates
for pyritic oxidation in natural environments. One possi-
ble explanation for this apparent contradiction may lie in
the higher crystalline pyrite concentration of the Utah
concentrates. This hard, smooth-surfaced structure would be
easier to wash with solvent than the multi-pore surface of
Shawville coal pyrites. The latter's pores could become
coated with reaction products, inhibiting transfer of these
products from the reaction sites. Other observers have
reported similar peculiarities in working with pyrites from
various sources. Clark(7) found that some coal pyrites
/ -I -7 \
oxidize rapidly, while others are nearly inert. Hanna ^-L/ ' '
et al, proved that coal seam pyrites behaved very differ-
ently from museum-grade pyrites. Mapstone^ ' found wide
variations in pyrites from the same source, depending on the
effects of weathering.
For a given source, and for each solvent, the effect of
particle size was predictable. The finer particle sizes
yielded the higher decomposition rates. The two runs utiliz-
ing Utah concentrate fines with each solvent were allowed
to recycle past the 4-hour time limit, but the chlorine
supply was shut off. It was hoped that the decomposition
of pyrites would be continued by reactions (2) and (3), but
as in the long-term recycle tests with methanol, the reac-
tions slowed to undetectable levels within two hours. In
their work on the chemical aspects of acid mine drainage,
Barnes and Romberger(19) found that ferric ion was not a
major contributor to the oxidation of pyrites unless a very
high Fe+3/FeS2 ratio was attained. Their work emphasized
the importance of bacterial oxidation of pyrites. These
useful organisms were reported to have accomplished 80%
complete pyrite oxidation within 4-5 days in studies con-
ducted by Silverman, Rogoff, and Wender^20). Unfortunately,
none of these bacteria could tolerate the chlorine-saturated
solvent environments required for these tests.
42
-------
TABLE 10
FOUR-HOUR RECYCLE RUNS FROM DIFFERENT SOURCES AND
SIZES OF PYRITE
Pyrite Weight Loss in Grams
Size Fraction in CCJ4 in
Utah Concentrate -60+200 12.0 15.9
-18+60 9.6 11.7
Shawville -60+200 1.22 11.4
-18+60 1.04 5.8
-1/8" to dust 0.72 1.9
43
-------
Recycle Test Using Chlorine-Saturated Water
One final 4-hour recycle test was run using 50 grains of
Shawville pyrite and 500 ml of chlorine-saturated, deionized
water as the solvent. The total weight loss measured by
weighing the pyrite at the end of the test was less than
0.1 grams, while the iron recovered in the recycling sol-
vent weighed only 0.02 grams. Pyrites were not decomposed
by reaction with chlorinated water.
44
-------
SECTION VIII
OTHER PYRITE DECOMPOSITION TESTS
Static Tests Using Vigorous Abrasion
As an alternative to recycling of solvents, an attempt was
made to provide transport of reaction products away from
the pyrite surfaces by vigorous abrasion in a Waring Blender.
Twenty grams of Shawville pyrite fines (<200 mesh) were
used in tests similar to the solubility measurements report-
ed in Table 3, except that contact times were shortened to
four (4) hours. Tests were run with chlorine-saturated
carbon tetrachloride, chloroform, and 1,2 dichloroethane.
Chlorine gas was sparged into each charge for 10 minutes
prior to the test, then for five (5) minutes at half-hour
intervals during the test. At all other times the pyrite
suspension was agitated at 1450 RPM in 300 ml of solvent.
Samples were withdrawn at hourly intervals for determina-
tion of dissolved matter. Results are shown in Table 11.
With 1,2 dichloroethane showing the most improvement,
the dissolution of pyrites was found to proceed at an
accelerated rate when compared with earlier solubility mea-
surements run with no provisions for vigorous abrasion.
Using abrasion, 50% more solutes were found in that solvent
in a four-hour contact time than were originally found in
24 hours. However, rates were still somewhat slower than
those obtained in recycle tests run for the same period of
time.
Evaluation of Reaction (2) - FeCl-j as an Oxidant
A saturated solution of FeCl3 in unchlorinated methanol was
prepared by dissolving 375 grams of FeCl3 in 300 ml of meth-
anol. The absence of chlorine was required to fully evaluate
the oxidation of pyrites by ferric chloride. Methanol was
chosen as the solvent to insure the high ferric ion/pyrite
ratio recommended by Barnes and Romberger^ ' to accomplish
any oxidation at all. A 150 ml aliquot of this saturated
solution was diluted to 300 ml with pure methanol, giving
a 50% saturated solution. Chunks of Shawville pyrite 0.75
to 1.5 grams in weight were immersed into each mixture over-
night. Pyrite weight losses following thorough rinsing and
drying of the specimens were 0.046 grams in the 100% satur-
ated solution, and 0.076 mg in the 50% saturated solution.
Both losses were considerably less than the 0.371 grams of
residues reported for unchlorinated methanol in previous
solubility tests. The high concentrations of iron salts
present may have inhibited the action of unchlorinated
methanol on pyrites.
45
-------
TABLE 11
SOLUBILITY OF SHAWVILLE PYRITES IN
CHLORINATED SOLVENTS WITH ABRASION
Elapsed" Grams of Residue/100 Grams of Solvent @ 25°C
Time CC14 CHC1_3 CH CHC12
1 Hour 0.010 0.006 0.059
2 Hours 0.022 0.016 0.112
3 Hours 0.035 0.028 0.170
4 Hours 0.048 0.039 0.210
Static Solubility Tests
Run earlier: (no abrasion)
24 Hours 0.074 0.048 0.13
15 Days 0.22 0.23 0.59
46
-------
In-situ Oxidation of Pyrites with Oxygen
The apparatus shown in Figure 5 was used to evaluate the
possibility of direct oxidation of pyrites in chlorinated
carbon tetrachloride. Tests were run using Shawville
pyrite fines, and using 1/4" mesh size lumps; at ambient
temperatures and at 41°C; and at one atmosphere and at two
atmospheres pressure. In all cases, samples were maintained
under 100% Cu atmospheres for 72 hours. All tests led to
the same conclusion—no oxygen consumption by any reactant
or reaction product had occurred during any part of the 72
hours.
47
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SECTION IX
ACKNOWLEDGEMENTS
The pyritic materials used for this project were obtained
through the co-operation of the management of Pennsylvania
Electric Company's Shawville Power Station near Clearfield,
Pennsylvania, and th-rough the assistance of Mr. K. M.
Ogilvie, Combustion Engineer, Kennecott Copper Corporation's
Metal Mining Division, Salt Lake City, Utah.
The support of the project by the Water Quality Office,
Environmental Protection Agency, and the help provided by
Mr. Donald J. 0'Bryan, Jr., Dr. James M. Shackelford, the
Grant Project Officer and Ernst P. Hall, Chief of the
Pollution Control Analysis Branch, is acknowledged with
sincere thanks.
Mr. Robert J. Ondof and Mr. Richard C. Rice of the Cyrus Win.
Rice Division - NUS CORPORATION, designed, built, and opera-
ted the apparatus used during the course of this project.
Mr. Dennis A. Clifford, formerly of the Cyrus Wm. Rice
Division - NUS CORPORATION, took part in preparation of this
report.
49
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SECTION X
REFERENCES
1. McElroy, R., and Peters, E. E., "Chlorination of Pyrites
In Chlorinated Hydrocarbon Solvents," Paper presented
at the Seventh Annual Conference of Metallurgists,
Vancouver, British Columbia (1968).
2. American Public Health Association, Inc., Standard Methods
for the Examination of Water and Waste Water,
Thirteenth Edition, American Public Health Association,
New York, N. Y, (1971).
3. Environmental Protection Agency, Water Quality Office,
Analytical Quality Control Laboratory, Methods for
Chemical Analysis of Water and Wastes. 16020 . . .
7/71, Cincinnati, Ohio (1971).
4. Gyenes, I., Titration in Non-Aqueous Media, D. Van
Nostrand Company, Inc., Princeton, N. J. (1967).
5. Kucharsky, J., and Safarik, L., Titrations in
Non-Aqueous Solvents, Elsevier Publishing Company,
New York (1965) .
6. Burnauer, S., Emmett, P. H., and Teller, E., "The Ad-
sorption of Gases in Multimolecular Layers," Journal
Am. Chem. Soc., 6_0_, pp. 309-316 (1938).
7. Clark, C. S., "Oxidation of Coal Mine Pyrite," Journal
of the Sanitary Engineering Division, ASCE, 92, No. SA-2,
Proceedings Paper #4802, pp. 127-145(1966)"
8. Braley, S. A., "Summary Report on Commonwealth of Penn-
sylvania Department of Health Industrial Fellowship
Nos. 1 to 7 inclusive," Mellon Institute Fellowship
No. 326B, February, 1964
9. Singer, P. C., "Oxygenation of Ferrous Iron," Water
Pollution Control Research Report No. 10410 06/69,
Harvard University (1970).
10. Sato, M., "Oxidation of Sulfide Ore Bodies, II. Oxida-
tion Mechanisms of Sulfide Minerals at 25°C.," Economic
Geology, 55, pp. 1202-1231 (1960).
11. Sherman, M. I., and Strickland, J. D. H., "Dissolution
of Pyrite Ores in Acid Chlorine Solutions," Journal of
Metals, Transactions of AIME, 9, pp. 1386-1388 (1957).
51
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12. McKay, D. R., and Halpern, J., "A Kinetic Study of the
Oxidation of Pyrite in Aqueous Suspension," Transactions
of the Metallurgical Society of AIME, 212, pp. 301-309
(1958).
13. Nelson, N. W., Snow, R. D., and Keys, D. B. , "Oxidation
of Pyritic Sulfur in Bituminous Coal," Industrial and
Engineering Chemistry, 25, No. 12, pp. 1355-1358 (1933).
14. Warren, I. H., "The Generation of Sulphuric Acid from
Pyrite by Pressure Leaching," Australian Journal of
Applied Science, 7_, pp. 346-358 (1956) .
15. LaRosa, P. J., and Michaels, H. J., "Study of Sulfur Re-
covery from Coal Refuse," Environmental Protection Agency,
Water Quality Office, Water Pollution Control Research
Series, 04010 FYY 09/71, Black, Sivalls, and Bryson,
Incorporated (1971).
16. Lorenz, W. C., and Tarpley, E. C., "Oxidation of Coal
Mine Pyrites," U. S. Department of the Interior, Bureau
of Mines, Report of Investigations No. 6247 (1963).
17. Hanna, G. P. Jr., Lucas, J. R., Randies, C. I., Smith,
E. E., and Brant, R. A., "Acid Mine Drainage Research
Potentialities," Journal of the Water Pollution Control
Federation, 35, pp. 275-294 (March 1963) .
18. Mapstone, G. E., "The Weathering of Pyrite," Chemistry
and Industry, 73, pp. 577-578 (1954).
19. Barnes, H. L., and Romberger, S. B., "Chemical Aspects
of Acid Mine Drainage," Journal of the Water Pollution
Control Federation, 40, No. 3, pp. 371-384 (1968).
20- Silverman, M. P., Rogoff, M. H., and Wender, I., "Re-
moval of Pyritic Sulphur from Coal by Bacterial Action,"
Fuel, 42, pp. 113-124 (1963) .
52
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Fu-M &, Gri'i/p
05B, 05G
SELECTED WATER RESOURCES ABSTRACTS
INPUT TRANSACTION FORM
Environmental Protection Agency, Water Quality Office (Federal Office)
Cyrus Wm. Rice Division, NUS Corporation (Contractor)
Title
Oxidation of Pyrites in Chlorinated Solvents
10
Authors)
Boros, Joseph A.
Brenneman, Donald R.
Troy, Joseph C.
16
EPA Contract No. 14-12-897
21 N
22
Citation
Environmental Protection Agency report
number EPA-R2-T2-069, November' 1972.
23
Descriptors (Starred First)
*Pyrites, *Chlorination, *Coal Mine Wastes, *Iron Compounds, *Sulfur
Compounds, Waste Dumps, Ultimate Disposal, Acid Mine Water, Water
Pollution Sources, Sulfides, Pollution Abatement.
25
Identifiers (Starred first)
*Non-Aqueous Chlorination, Chlorinated Solvents
27 Abstract The purpose of this study was to develop methods for extracting
sulfur and iron compounds from pyritic waste materials, producing a final
refuse that is incapable of causing water pollution, and at the same time,
conserving useful mineral resources. The primary method under investigation
was the Chlorination of pyrites in non-aqueous solvents.
An attractive feature of the proposed method was its cyclic nature. Chlo-
rine would serve as the original driving force for the expected reactions, but
would be replaced as primary oxidant by two products of the Chlorination reac-
tion, ferric chloride and sulfur dichloride.
Test results indicated that the successful extraction of sulfur and iron
compounds depended upon the continuous addition of chlorine gas to the system.
The expected replacement of chlorine by ferric chloride and sulfur dichloride
was not achieved under conditions of this study. The reaction rates with
chlorine gas varied with the particle size and source of the pyritic material;
with the choide of solvent; and with the - means of providing intimate contact
between pyrites and chlorine-saturated solvents. Recycling solvents through
beds of pyrite significantly improved rates of reaction, while elevated temp-
eratures and treatment in the presence of ultra-violet radiation yielded.
slight improvements in reaction rate. This report contains 20 references.
Abstractor
Joseph A. Boros
Institution
Cyrus Wm. Rice Division - NUS Corporation
WR:I02 (REV. JULY 1969)
WRSI C
SEND TO. WATER RESOURCES SCIENTIFIC INFORMATION CENTER
U S DEPARTMENT OF THE INTERIOR
WASHINGTON. D C 20240
* GPO: 1969- 33 9-339
w U. S. GOVERNMENT PRINTING OFFICE : 1973—514-149/95
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