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Chemical Geology 217

Ferric iron-bearing sediments as a mineral trap for CO2

sequestration: Iron reduction using sulfur-bearing waste gas

James L. PalandriT, Yousif K. Kharaka

U.S. Geological Survey, 345 Middlefield Rd., MS427, Menlo Park, California 94025, United States

Received 22 April 2003; accepted 10 December 2004

Abstract

We present a novel method for geologic sequestration of anthropogenic CO2 in ferrous carbonate, using ferric iron present in

widespread redbeds and other sediments. Iron can be reduced by SO2 that is commonly a component of flue gas produced by

combustion of fossil fuel, or by adding SO2 or H2S derived from other industrial processes to the injected waste gas stream.

Equilibrium and kinetically controlled geochemical simulations at 120 bar and 50 and 100 8C with SO2 or H2S show that iron

can be transformed almost entirely to siderite thereby trapping CO2, and simultaneously, that sulfur can be converted

predominantly to dissolved sulfate. If there is an insufficient amount of sulfur-bearing gas relative to CO2 as for typical flue gas,

then some of the iron is not reduced, and some of the CO2 is not sequestered. If there is an excess of sulfur-bearing gas, then

complete iron reduction is ensured, and some of the iron precipitates as pyrite or other solid iron sulfide, depending on their

relative precipitation kinetics. Gas mixtures with insufficient sulfur relative to CO2 can be used in sediments containing Ca, Mg,

or other divalent metals capable of precipitating carbonate minerals. For quartz arenite with an initial porosity of 21% and

containing 0.25 wt.% Fe2O3, approximately 0.7 g of CO2 is sequestered per kg of rock, and the porosity decrease is less than

0.03%. Sequestration of CO2 using ferric iron has the advantage of disposing of SO2 that may already be present in the

combustion gas.

D 2005 Published by Elsevier B.V.

1. Introduction

Disposal of anthropogenic CO2 captured from

fossil fuel-fired power plants into geologic formations

is a potential means of reducing the amount of

greenhouse gases released to the atmosphere. In

geologic sequestration, CO2 may be stored in deep

saline aquifers or structural traps such as depleted

0009-2541/$ - see front matter D 2005 Published by Elsevier B.V.

doi:10.1016/j.chemgeo.2004.12.018

* Corresponding author.

E-mail address: jlpaland@usgs.gov (J.L. Palandri).

petroleum or gas reservoirs, (1) as a supercritical fluid

(hydrodynamic trapping), (2) in formation water as a

dissolved constituent (solution trapping), or (3) as

carbonate minerals (mineral trapping), the latter being

the main theme of this study. Most studies of in situ

mineral trapping (e.g. Gunter et al., 2000; Gunter et

al., 1997; Johnson et al., 2001) consider disposal in

plagioclase-bearing (Ca-bearing), illitic (Mg-bearing),

or glauconitic (Fe II-bearing) sediments, to trap CO2

in calcite, magnesite, siderite, or ankerite. These

sediments are suitable repositories, but are generally

(2005) 351–364

J.L. Palandri, Y.K. Kharaka / Chemical Geology 217 (2005) 351–364352

of limited thickness and geographical occurrence, or

of limited alkaline earth or iron content. Other iron-

bearing sediments, including redbeds, have the

advantages of widespread geographic distribution,

generally greater thickness, and higher porosity and

permeability. Worldwide, sediments have average iron

contents of 4.8 and 2.4 wt.% for mudrocks and

sandstone, respectively (Blatt, 1982), and have an

enormous potential to trap CO2 in iron carbonate

minerals.

Ferric iron in these deposits must be reduced to

ferrous iron to precipitate in siderite. Possible reduc-

tants include organic matter, SO2 or H2S. At least one

study of CO2 sequestration (Pruess et al., 2001) makes

passing reference to reduction of ferric iron to

precipitate siderite, and refers to organic matter or

H2S as possible agents for the reduction of iron. In the

case of H2S, Pruess et al. (2001) state that even

relatively low H2S partial pressures favor the precip-

itation of ferrous iron in pyrite.

We show with geochemical simulations that SO2,

which is commonly a component of flue gas from

combustion of sulfur-bearing fossil fuel, can be used

to reduce ferric iron to ferrous iron, thereby allowing

CO2 to precipitate as iron carbonate. Fugacities of

H2S that is derived from the reaction of SO2 with

water, and of CO2 can be maintained so that

precipitation of iron carbonate is favored over iron

sulfide, and sulfur in the gas can be converted almost

entirely to dissolved sulfate.

2. Methods

Thermodynamic calculations are performed with

the computer programs CHILLER (Reed, 1982;

Reed, 1997; Reed, 1998) and GAMSPATH (Perkins

et al., 1997), which is the most recent version of

PATHARC (Perkins, 1980). These programs compute

the distribution of chemical components among

aqueous species, minerals, and gases, using equili-

brium constants for aqueous, solid, and gas reactions

at selected temperature and pressure. CHILLER is

used to compute reaction paths in geologic systems

by changing one of the system variables incremen-

tally, e.g. composition, temperature, pressure, or

enthalpy, and re-computing equilibrium at each step.

The thermodynamic database is identical to that

described by Palandri and Reed (2001), and is not

repeated here. CHILLER computes aqueous activity

coefficients using the extended Debye–Hqckel equa-tion of Helgeson et al. (1981), as modified by Tanger

and Helgeson (1988). CHILLER uses a virial

equation (Spycher and Reed, 1988) to compute the

fugacity coefficients for H2O, CO2, CH4, and

mixtures thereof, and for H2 alone; data are not

available for SO2 nor H2S which is one of the

reaction products of the reaction of SO2 with water,

and ideality is assumed. However, because SO2 and

derived H2S quantities are small and their solubilities

are large, the amounts of these two gases in any CO2-

rich phase that saturates is exceedingly small. A

further limitation is that the calculation of equili-

brium constants is limited to conditions along the

H2O liquid–vapor saturation curve. However varia-

tions in the equilibrium constants due to pressure

change are quite small compared to those due to

temperature change. In the CHILLER thermody-

namic database, mineral data are derived from the

data of Holland and Powell (1998), and aqueous

species data is derived from the SUPCRT92 database

of Johnson et al. (1992).

GAMSPATH is used to compute the effects of the

kinetics of mineral dissolution and precipitation,

which are important in low temperature systems. This

software simulates an infinitely well-stirred batch

reactor, and thus reaction kinetics are strictly surface

controlled. Concentration gradients near mineral

surfaces leading to diffusion controlled reactions are

not considered. Simulations therefore provide an

estimate of the minimum time for equilibration.

GAMSPATH computes reaction paths that are con-

strained by a general rate equation (modified from

Lasaga, 1995)

dm

dt¼ � Sd

Xj

�Aje

�Ej

RT jiani;ji;j ð1� XpÞq

�ð1Þ

where each term in the sum represents a single

reaction mechanism. The rate is dm/dt (mol/s), S is

the mineral surface area (m2), A is the Arrhenius pre-

exponential factor (mol/m2 s), E is the activation

energy (J/mol), T is the temperature (K), and R is the

gas constant. The dimensionless product of ai,jni,j is to

account for the activities of the ith species participat-

ing in the jth reaction mechanism, raised to a reaction

J.L. Palandri, Y.K. Kharaka / Chemical Geology 217 (2005) 351–364 353

order ni,j; the effects of surface adsorption are not

considered. The dimensionless function of the chem-

ical affinity (1�Xp)q where X is the mineral

saturation index, i.e. the activity product divided by

the equilibrium constant, accounts for slowing of

reactions as equilibrium is approached; the parame-

ters p and q are empirical and dimensionless. A

significant simplification is that nucleation is simu-

lated by precipitating 106 spherical shells with an

outer radius of 1 Am, and an inner radius that is

controlled by the amount of mineral that is computed

to precipitate in that time step. Once the shells are

filled, precipitation proceeds on the outer surface. A

default saturation index (Q/K) of 1.0 must be

attained for mineral precipitation to be initiated,

which we have used in the absence of experimental

data; other values can be specified if data are

available.

The GAMSPATH thermodynamic database is

derived from the SUPCRT92 database (Johnson et

al., 1992). Limitations of the GAMSPATH software

include restriction to isothermal and isobaric reaction

paths, a simplified heat capacity regression equation

for aqueous species containing a single variable

parameter, computation of gas fugacities based on

the assumption of the ideality of the gases and

mixtures thereof, and a simpler equation than

contained in CHILLER for computing aqueous

activity coefficients, the b-dot method of Helgeson

(1969). To facilitate comparison among the simula-

tions shown, all were constructed using a pressure of

120 bar, a temperature of 100 8C, and an initial 1.0

kg of H2O adjusted to 1.0 NaCl molality, and pH=5.1

using HCl.

3. Results and discussion

In the discussion below, we investigate several

geochemical reactions in order of increasing com-

plexity. First, we examine the reaction of SO2 gas

with water that produces sulfide, sulfate, and

extremely acidic conditions. Next, we investigate

the reaction of dissolved SO2 gas with a synthetic

redbed, a ferric iron-bearing quartz arenite, that

produces iron sulfide minerals and dissolved sulfate.

The SO2 reactions with water and redbed minerals are

considered in detail because they are essential to the

desired overall reaction with CO2. We then consider

the reaction of gas mixtures consisting of mostly CO2

and lesser amounts of SO2 with a synthetic redbed

that produces siderite and dissolved sulfate. Finally,

we consider the reaction of a CO2–SO2 mixture with

a rock containing both ferric iron and calcium that

produces the carbonate minerals siderite, calcite,

ankerite, and dawsonite.

3.1. The reaction of SO2 gas with water

Field observations of volcanic systems and ther-

modynamic modeling show (Getahun et al., 1996;

Symonds et al., 2001) that SO2 gas forms H2S and

H2SO4 in a disproportionation reaction:

4SO2ðgÞ þ 4H2OðlÞ ¼ H2SðaqÞ þ 3H2SO4ðaqÞ ð2Þ

Dissociation of H2S and H2SO4 to yield H+ further

leads to an extremely low fluid pH. SO2 will react

with water to produce H2SO4 and H2 (Symonds et al.,

2001) at temperatures as low as 100 8C, but only as a

gas (steam) phase reaction:

SO2ðgÞ þ 2H2OðgÞ ¼ H2ðgÞ þ H2SO4ðg; lÞ ð3Þ

Reaction (3) can be discounted because water

generally exists as a liquid in sedimentary systems.

The results of Reaction (2) are demonstrated in a

simulation at 100 8C, as shown in Fig. 1. The

reaction produces sulfate, dominantly as either SO42�

or HSO4� with lesser NaSO4

�,depending on pH, and

sulfide mostly as dissolved H2S (Fig. 1B) with

minor HS�. The pH (negative pH in Fig. 1) evolves

from near 6.0 initially, to less than zero at the

termination of the simulation where 0.78 kg of SO2

gas has been added, or �0.11 log10 kg of gas

added, hereafter referred to as (log) kga. Sulfur

reaches saturation at less than 1 g of SO2 added

(Fig. 1A), in response to increasing sulfate, sulfide,

and decreasing pH:

3H2Sþ H2SO4 ¼ 4SðnativeÞ þ 4H2O ð4Þ

If a sink exists to remove sulfide, e.g. by oxidation

to sulfate or by precipitation in iron sulfide minerals

(considered below), then sulfur will not precipitate.

Similar results are obtained at temperatures from 25 to

300 8C, albeit with slight differences because of the

Fig. 1. Reaction at 100 8C of SO2 with 1.0 kg of H2O at 1.0 NaCl molality. Equilibrium composition vs. log kg gas added: (A) minerals; (B)

selected aqueous species.

J.L. Palandri, Y.K. Kharaka / Chemical Geology 217 (2005) 351–364354

decreasing solubility of H2S and increasing solubility

of sulfur with increasing temperature.

3.2. The reaction of SO2 gas with ferric iron and

water

Within the last decade, research has been con-

ducted to investigate subsurface sequestration of H2S

(e.g. Palandri, 2000, and data of others therein), a

naturally occurring contaminant in natural gas, by

precipitation of iron sulfide minerals in iron-bearing

sediments. One possible reaction is:

4Fe2þ þ SO2�4 þ 7HS� þ Hþ

¼ 4FeS2ðpyriteÞ þ 4H2O ð5Þ

However, experimental results (Palandri, 2000)

show that the precipitation of amorphous iron sulfide

(FeS) is favored kinetically:

Fe2þ þ HS� ¼ FeSðamorphousÞ þ Hþ ð6Þ

Selected minerals may be suppressed from the

calculations, where a less stable phase is expected to

precipitate because of their relative kinetics; examples

include suppression of quartz where amorphous silica

is expected, or of pyrite and pyrrhotite where

amorphous FeS is expected. In the simulations

considered herein, pyrite is not suppressed to show

that even the most stable of the iron sulfide minerals is

a less stable repository for iron than are iron-bearing

carbonate minerals. Further, pyrite would likely

replace amorphous FeS if additional sulfide is

available after all iron precipitates in FeS as would

occur near an injection well, but the kinetics of the

reaction are uncertain.

Where sediments contain ferric iron, it must first be

reduced before precipitation as iron sulfide. The

reduction of ferric iron by sulfide is shown by the

following two reactions:

8Fe3þþ HS�þ 4H2O ¼ 8Fe2þþ SO2�4 þ 9Hþ ð7aÞ

4Fe2O3ðhematiteÞ þ 15Hþ þ HS�

¼ 8Fe2þ þ SO2�4 þ 8H2O ð7bÞ

We will consider only Fe2O3 as in hematite, rather

than Fe3+, because the primary source of ferric iron in

the environment is contained in minerals. Iron in

natural sedimentary systems may also occur in many

other minerals, especially phyllosilicates, but these are

not considered in the simulations because thermody-

namic data are lacking, or because they are metastable

and kinetic data are lacking. This limitation is a

significant source of uncertainty in our models,

especially in the time-dependent simulations. Com-

bining Reaction (5) or(6) with Eq. (7b) yields an

overall reaction for the sequestration of H2S in ferric

J.L. Palandri, Y.K. Kharaka / Chemical Geology 217 (2005) 351–364 355

iron-bearing (hematite-bearing) sediments as either

amorphous FeS or pyrite:

4Fe2O3ðhematiteÞ þ 15H2Sþ H2SO4

¼ 8FeS2ðpyriteÞ þ 16H2O ð8aÞ

4Fe2O3ðhematiteÞ þ 9H2S ¼ 8FeSðamorphousÞ

þH2SO4 þ 8H2O ð8bÞ

A simulation of the reaction of pure SO2 with

hematite in water at 100 8C is summarized in Fig. 2.

The simulated rock is a synthetic redbed composed of

99.75 wt.% quartz and 0.25 wt.% hematite. The

quantity of hematite is kept small to delay gas

saturation until much of the iron in hematite is

reduced, because data are not available to compute

accurately the fugacity coefficients for CO2–H2S

mixtures in the simulations. Before adding SO2 gas

in the simulation, 10.0 kg of rock is equilibrated with

1.0 kg of 1.0 m NaCl, yielding a porosity of 20.9%

based on densities of quartz (q=2.65 g/cm3), hematite

(q=5.28), and brine (q=1.04); the equilibration has

little affect on the mineral assemblage.

Where log kga for pure SO2 is less than �6.2 and

pyrite is absent (Fig. 2A), the pH increases initially

Fig. 2. Reaction at 100 8C of SO2 with 1.0 kg of H2O at 1.0 NaCl molality

0.25 wt.% hematite. Equilibrium composition vs. log kg gas added: (A) m

due to combined hematite dissolution and iron

reduction, reaction (7b). With further addition of

SO2, increasing concentrations of Fe2+, sulfate and

sulfide drive reaction (5) to the right, leading to pyrite

precipitation (Fig. 2A), and the pH begins to decrease.

Hematite replacement by pyrite proceeds over a wide

range of kga, accompanied by steadily increasing H+,

Fe2+ and sulfate; total sulfide decreases only slightly.

After hematite dissolves out, increasing H+, sulfate

and sulfide at near-constant Fe eventually drive the

replacement of pyrite by native sulfur. At the pyrite

maximum near �1.1 log kga, 55% of the iron is in the

aqueous phase in the ferrous state. Porosity change is

negligible because of the small amounts of reacting

minerals relative to quartz, increasing by ~0.03% near

�1.1 log kga at the pyrite (q=5.01 g/cm3) maximum.

Experiments using Na2S and redbed from Tempe, AZ

(Palandri, 2000), rather than SO2 used in the

simulation, show that the length of time for the

reaction to attain 90% completion ranges over time

scales of days to weeks, depending primarily on the

experimental temperatures of 28 to 71 8C. The most

significant result of this reaction is the steady increase

in the concentration of dissolved ferrous iron to over

12,000 ppm, despite its removal by pyrite precipita-

tion. This ferrous iron is therefore available for

precipitation in carbonate minerals.

, previously equilibrated with 10 kg of rock, 99.75 wt.% quartz and

inerals; (B) selected aqueous species.

J.L. Palandri, Y.K. Kharaka / Chemical Geology 217 (2005) 351–364356

3.3. The reaction of CO2–SO2 gas mixtures with ferric

iron and water

Here we consider the reaction of primary interest,

the reaction of CO2 and SO2 gas mixtures with ferric

iron in hematite and water. The overall reaction is:

Fe2O3ðhematiteÞ þ 2CO2ðgÞ þ SO2ðgÞ þ H2O

¼ 2FeCO3ðsideriteÞ þ H2SO4 ð9Þ

The first step in the reaction is the solution of the

gases and the conversion from their anhydrous to

aqueous forms, e.g. HCO3�, H2CO3, HS

� and H2S.

The coefficients in Reaction (9) show that the gas

composition required is 2/3 CO2 and 1/3 SO2 for

sediments containing ferric iron as the only sink for

CO2. Where there exist other divalent metals (e.g. Ca,

Mg), the amount of SO2 relative to CO2 required to

trap all of the CO2 can be much lower. In the

simulations considered in this section, the system

composition and conditions before adding gas are

identical to that discussed in Section 3.2.

Fig. 3 shows the results of a simulation using a 95/

5 wt.% mixture of CO2/SO2, a gas composition with a

higher concentration of SO2 than is representative of

Fig. 3. Reaction at 100 8C of a mixture of CO2, 95 wt.%, and SO2, 5 wt.%,

10 kg of rock, 99.75 wt.% quartz and 0.25 wt.% hematite. Equilibrium

species—total bicarbonate, dominantly HCO3�and H2CO3; total sulfate, do

Fe2+, FeCl+, and FeSO4.

the waste stream from burning low-sulfur fossil fuel

(0.01–0.1 wt.%). The desired reaction product, side-

rite, is stable over a wide range of total gas added,

from �3.3 to �0.34 log kga (0.5 to 460 g). Pyrite

eventually replaces siderite, but only after hematite

has dissolved completely. A supercritical CO2 domi-

nated gas phase saturates near �1.4 log kga (40 g),

and its H2S content remains less than 1.0 ppm until

siderite dissolves completely. Sulfur eventually repla-

ces pyrite at a higher gas–rock ratio, near 0.2 log kga

(1580 g).

At small values of log kga, less than that at which

siderite precipitates (Fig. 3A), the pH is nearly

constant because hematite dissolution consumes

copious H+, as in Reaction(7b), and addition of gas

causes the concentration of ferrous iron, total sulfide,

and total carbonate to increase steadily (Fig. 3B) due

to Reactions (2) and (7b) and dissolution of CO2.

Where siderite begins to precipitate the pH decreases

as the reaction produces additional H+:

H2CO3 þ Fe2þ ¼ FeCO3 þ 2Hþ ð10Þ

Total sulfide also begins to decrease because

decreasing pH displaces Reaction (7b) to the right.

A significant finding is that the total amount of sulfide

with 1.0 kg of H2O at 1.0 NaCl molality previously equilibrated with

composition vs. log kg of gas added: (A) minerals; (B) aqueous

minantly SO42�, FeSO4, NaSO4

�, and HSO4�; total iron, dominantly

J.L. Palandri, Y.K. Kharaka / Chemical Geology 217 (2005) 351–364 357

in the aqueous–gas system, as dissolved H2S, HS�,

and gaseous H2S remains small, less than 1.0 ppt,

until hematite dissolves out. Porosity change is again

negligible, increasing by ~0.002% near the siderite

(q=3.94 g/cm3) maximum. Approximately 0.7 g of

CO2 is sequestered in siderite per kg of rock, or about

1.9 kg of CO2 per m3 of rock. The pore fluid contains

~3 wt.% CO2, or about 4 times the amount in siderite

in a given volume of water and rock.

An additional simulation was constructed at 50 8Crather than 100 8C (not shown), with all other

parameters identical, to test the reaction at lower

temperatures. These results are quite similar to those

obtained at the higher temperature, except: (1) gas

saturation is delayed to a slightly larger amount of gas

added (�1.3 log kga) because of increased gas

solubility at lower temperature; (2) siderite precip-

itation is also delayed (�3.1 log kga) because siderite

stability decreases with decreasing temperature; and

(3) the total amount of sulfide in the aqueous and

gaseous phases is lower, because pyrite stability

increases with decreasing temperature.

Several additional simulations were constructed to

test the effect of varying the relative amounts of CO2

and SO2 in the gas added (Fig. 4). Siderite is the only

stable alteration mineral over a wide range of gas

added, at ratios of up to approximately 50/50 wt.%

CO2/SO2. At gas ratios less than ~80/20 wt.% CO2/

SO2, a small amount of pyrite precipitates before

siderite at small kga, because of the greater amount of

total dissolved sulfide relative to total carbonate; as

more gas is added, dissolved bicarbonate reaches a

concentration where siderite is preferred over pyrite.

Fig. 4D shows the reaction using CO2 and SO2 in a

molar ratio of 2 to 1, the composition required for the

stoichiometry shown in Reaction (9). The gas phase

saturates only after all of the hematite is consumed

because much CO2 remains dissolved in solution; the

gas composition (30% SO2) used in Fig. 4C is nearer

the composition where gas saturation is coincident

with exhaustion of hematite.

If the gas mixture contains less CO2 (Fig. 4E),

some of the iron that otherwise could trap CO2 in

siderite will precipitate in pyrite instead, but these

gas mixtures are not likely candidates for injection.

If the gas mixture contains more CO2 (Fig. 4A, B),

there will be insufficient SO2 to reduce the quantity

of iron required to fix all of the CO2 in siderite, and

the gas phase saturates before all of the hematite is

consumed. For the simulation shown in Fig. 3 (5%

SO2) at the siderite maximum, the amounts of CO2

in gas, brine, and siderite, respectively are approx-

imately 150 g, 30 g, and 10 g. It must therefore be

emphasized that to trap all of the CO2 in minerals

using gas mixtures representative of flue gas, i.e.

less than 1% SO2, the targeted rocks must contain

other metals capable of precipitating as carbonate

minerals (see below). If a rock type contains only

small amounts of other divalent metals, then it may

be possible to enhance the sulfur content of the

waste gas by adding H2S or SO2 derived from other

processes, e.g. H2S derived from sour natural gas

production. Further, it may actually be beneficial to

use flue gas from fossil fuel with high sulfur

content, which has generally been avoided in the

past where flue gas is released to the atmosphere, to

limit the production of acid rain. In any case, a rock

type targeted for CO2 sequestration needs to contain

only a small amount of ferric iron to effectively

oxidize all of the sulfur in flue gas derived from

low-sulfur fuel.

If H2S is used instead of SO2 as a reducing

agent, approximately the same results are obtained,

but at an H2S concentration approximately one

fourth of that for SO2, i.e. a mixture of 1.25 mol%

H2S and 98.75 mol% CO2 produces nearly the

same result as a mixture of 5 mol% SO2 and 95

mol% CO2 (see Reaction (1)). Differences in the

reaction using H2S rather than SO2 are minor, and

include a much lower sulfate concentration for a

given total mass of gas added, and because of the

lower sulfate, graphite precipitates instead of native

sulfur. However, carbonate or methane can coexist

metastably with sulfide for time scales up to

millions of years at temperatures less than ~140 8C(Heydari, 1997; Worden and Smalley, 1996; Wor-

den et al., 1995), and graphite would not likely

precipitate.

In subsurface sequestration of CO2–SO2 gas

mixtures with radial flow outward from an injection

well, it is likely that a roughly radial pattern would

develop around the well, centered horizontally and

skewed upward due to density effects. Near the well,

assemblages resulting from high gas–rock ratios

would form, including native S, iron sulfide, and

quartz. With increasing distance from the well, the

Fig. 4. Reaction at 100 8C of mixtures of CO2 and SO2, with 1.0 kg H2O at 1.0 NaCl molality previously equilibrated with 10 kg of rock, 99.75

wt.% quartz and 0.25 wt.% hematite. Equilibrium mineral assemblage vs. log kg gas added, wt.% CO2/SO2: (A) 99:1; (B) 90:10; (C) 70:30; (D)

58.2:41.8 (molar ratio of 2 to 1); (E) 30:70.

J.L. Palandri, Y.K. Kharaka / Chemical Geology 217 (2005) 351–364358

assemblage changes to pyrite, siderite, and quartz.

Distal from the well, the assemblage is hematite,

siderite, and quartz, and in the outermost zone,

unaltered rock.

3.3.1. Time dependence of the reaction

We constructed a time-dependent simulation of the

CO2–SO2–hematite–H2O reaction using program

GAMSPATH (Fig. 5). Because the simulations

Fig. 5. Time-dependent reaction at 100 8C of 10 kg of rock, 99.75 wt.% quartz and 0.25 wt.% hematite, with 1.0 kg H2O at 1.0 NaCl molality

previously equilibrated with 100 g of CO2/SO2 mixture, 95:5 wt.%. Composition vs. log time: (A) minerals; B) aqueous species—total

bicarbonate, dominantly HCO3� and H2CO3; total sulfate, dominantly SO4

2�, FeSO4, NaSO4�, and HSO4

�; total iron, dominantly Fe2+, FeCl+,

and FeSO4.

J.L. Palandri, Y.K. Kharaka / Chemical Geology 217 (2005) 351–364 359

considered herein are restricted to moderately to

extremely acidic conditions, only a single, H+-

catalyzed mechanism has a significant effect upon

reaction rates, and other mechanisms are negligible in

comparison.

Therefore, a much simpler equation

dm

dt¼ � Sd Ae

�ERT an1

Hþ d an2HCO�

3an3H2CO3

d an4Fe3þ

d 1� Xpð Þq

ð11Þ

yields rates nearly identical to the full Eq. (1), above.

The mineral dissolution rate parameters and grain

morphologies are summarized in Table 1. The rate

parameters are derived using surface areas determined

by BET, which are generally greater than geometric

surface areas. We use geometric surface area in the

simulations because the surface area of a hypothetical

mineral cannot be determined, and therefore our

computed rates may be slower by the factor of the

difference between BET and geometric surface area

for a particular mineral. This factor may exceed an

order of magnitude for poorly crystallized clay

minerals, but is generally less than 2.0 for the

crystalline minerals given in Table 1. The default

reaction order is zero if a particular species does not

participate in the reaction or if data are not available,

and default chemical affinity parameters p and q are

unity because data are absent. Mineral precipitation

rates were computed by dividing the dissolution rate

by the equilibrium constant, based on the principle of

microscopic reversibility (Lasaga, 1998).

The synthetic redbed rock is composed of 99.75

wt.% quartz and 0.25 wt.% hematite, and the gas is

composed of 90 wt.% CO2 and 10 wt.% SO2. The

amount of SO2 is slightly in excess of that required

to reduce all of the iron, and the amount of CO2 is

far in excess of the amount that can precipitate in

siderite, to ensure it is present as a separate phase

over the entire reaction. A GAMSPATH simulation

of this type is constructed by first equilibrating the

gas with the fluid, and then adding the rock. The

reaction path is then computed as constrained by free

energies and kinetics of mineral precipitation and

dissolution reactions. A significant limitation is that

there is no accommodation for the kinetics of

aqueous speciation, for which equilibrium is

assumed, and that the fluid must first be equilibrated

with the gas prior to reacting with the rock. Related

is that equilibration of fluid with SO2 gas leads to

saturation with respect to sulfur. In the current

simulation sulfur precipitation is suppressed, and

Table 1

Rate equation parameters

Mineral Morphology Rate constanta

(log mol/m2 s)

Activation energy

(kJ/mol)

Reaction orderb

w.r.t. activity of H+

Reaction orderb

w.r.t. given species

Quartzc 0.1 mm spheres �13.34 90.1 0.00 0.00

Hematited 0.01 mm spheres �9.39 66.2 1.00 0.23, HCO3�

Pyritee 0.1 mm spheres �7.52 56.9 �0.50 0.50, Fe3+

Sideritef 0.1 mm spheres �0.30 14.4 1.00 0.50, H2CO3

Albiteg 0.1 mm spheres �9.87 65.0 0.46 0.00

Anorthiteh 0.1 mm spheres �3.50 16.6 1.41 0.00

a At 25 8C, and a(H+)=a(Fe3+)=a(HCO3�)=a(H2CO3)=1.0.

b Default reaction order is 0.0 for species not in the reaction mechanism, or data are not available.c Data regressed from Tester et al. (1994).d Data regressed from Bruno et al. (1992) and Ruan and Gilkes (1995).e Data regressed from McKibben and Barnes (1986).f Rate set equal to calcite, data regressed from Talman et al. (1990).g Data regressed from Hellman (1994); p=0.76, q=90.0 in Eq. (11) from Alekseyev et al. (1997).h Data regressed from Hellman (1994), and Blum and Stillings (1995).

J.L. Palandri, Y.K. Kharaka / Chemical Geology 217 (2005) 351–364360

the concentrations of sulfate and sulfate are higher

than they would be if it were not. Fortunately, this

leads to fluid pH only slightly lower than is would

occur otherwise, 1.17 compared to 1.19.

The chemical reactions in this simulation are

similar to those in the equilibrium simulation dis-

cussed in Section 3.3, but with kinetic constraints

applied. Differences arise because the fluid is first

charged with gas before adding rock, rather than

equilibrating the fluid with the rock before adding gas.

Significant differences include a high initial concen-

tration of total sulfide, and a low pH before adding

rock, and the progression from a reducing, low pH

system to a more oxidizing system with slightly

higher pH, rather than the opposite computed in the

equilibrium simulation.

Quartz has the slowest dissolution rate, but it

equilibrates quickly because only a small amount

dissolves, and it does not participate in the reactions

of interest. Hematite has the next slowest dissolution

rate, and controls the length of time for the system to

equilibrate, because pyrite and siderite rates are faster.

The length of time to initial siderite precipitation is

about 2 h, and to complete system equilibration is about

18 h. Because the reactions are surface controlled, these

estimates are for the minimum length of time for the

reactions. However, given that the H2S sequestration

experiments discussed above were carried out in well

stirred reactors and required much longer to equilibrate,

the computed rates may be too fast for other reasons,

e.g. the hematite rate data are in error, or the iron in the

experimental redbed is contained in other minerals.

Given that published rates for dissolution of a single

mineral typically show uncertainties of up to one order

of magnitude within a single work, and up to several

orders of magnitude among different workers, the

reaction should be examined in experiments designed

to examine the reaction of CO2–SO2 gas mixtures with

iron-bearing rocks.

3.4. The reaction of CO2–SO2 with ferric iron in

calcium feldspar-bearing sediments

Here we consider the equilibrium reaction of a 95/

5 wt.% mixture of CO2/SO2 with a rock containing

94.5% quartz, 5.0% labradorite, and 0.5% hematite

by weight. The rock is equilibrated with the fluid

before adding gas, altering it to an assemblage of

quartz, stilbite, albite, hematite, and paragonite (Fig

6A). As in Section 3.3, iron-bearing carbonates are

stable over a wide range of total gas added, from

�3.6 to 0.4 log kga. However, ankerite replaces

siderite near �3.0 log kga, because of the increasing

calcium concentration due to the dissolution of

stilbite. The maximum amount of CO2 that is

sequestered in carbonate minerals occurs where

stilbite dissolves out near �1.2 log kga, and is

sequestered mostly in calcite, with minor ankerite.

The pH is buffered by the assemblage albite+par-

agonite+quartz until albite dissolve out, and then

shows a slight decrease until it is buffered by the

assemblage anhydrite+calcite together with the high

Fig. 6. Reaction of a mixture of CO2, 90 wt.%, and SO2, 105 wt.%, with 1.0 l of 1.0 molal NaCl at 100 8C previously equilibrated with 10 kg of

rock composed of 94.75 wt.% quartz, 5.0 wt.% labradorite, and 0.25 wt.% hematite, equilibrium composition vs. log kg of gas added: (A)

minerals, log mol per kg of initial fluid; (B) concentrations of selected aqueous species, molality. Total bicarbonate is composed dominantly of

HCO3� and H2CO3; total sulfate is composed dominantly of SO4

2�, FeSO4, NaSO4�, and HSO4

�; total iron is composed dominantly of Fe2+,

FeCl+, and FeSO4, total calcium is composed of Ca2+ and CaSO4 (aq.).

J.L. Palandri, Y.K. Kharaka / Chemical Geology 217 (2005) 351–364 361

concentrations of sulfate and bicarbonate. Anhydrite

replaces calcite and ankerite only at the highest gas–

rock ratios, only far after the initial reactants are

consumed; and is driven by the increasing sulfate

concentration. The porosity increases from 18.8%

initially to 20.6% near where dawsonite (q=2.43)precipitates and paragonite dissolves out, and then re-

mains nearly constant (qalbite=2.62 g/cm3, qanhydrite=

2.96, qankerite=3.27, qcalcite=2.71, qparagonite=2.89,

qstilbite=2.14,); however, in rocks containing unal-

tered primary plagioclase rather than hydrous zeolite,

it is expected that the porosity would decrease, as

feldspar is replaced with kaolinite (q=2.60). In this

reaction, approximately 5.6 g of CO2 is sequestered

per kg of rock, or about 14.8 kg of CO2 per m3 of

rock, at the calcite maximum where stilbite dissolves

out.

3.4.1. Time dependence of the reaction

We constructed a time-dependent simulation of

the CO2–SO2–hematite–labradorite–H2O reaction

using program GAMSPATH (Fig. 7). The parame-

ters are identical to those used in the time-depend-

ent simulation shown in Fig. 5 and discussed

above, except that the rock is composed of 94.75

wt.% quartz, 5.0 wt.% labradorite and 0.25 wt.%

hematite. Because of software and hardware limi-

tations, the simulation was computed using equili-

brium rather than kinetic mineral precipitation, and

overall reaction rates are controlled only by

dissolution rates of minerals that are present

initially. Therefore only reactant minerals initially

present are listed in Table 1. As above, quartz is

essentially inert once quartz saturation is reached

(Fig 7A).

Fig. 7. Time-dependent equilibrium reaction of 10 kg of rock composed of 94.75 wt.% quartz, 5.0 wt.% labradorite, and 0.25 wt.% hematite,

with 1.0 l of 1.0 molal NaCl at 100 8C previously equilibrated with a 100 g mixture CO2, 95 wt.%, and SO2, 5 wt.%, composition vs. log time:

(A) primary minerals, log mol per kg of initial fluid; (B) alteration minerals, log mol per kg of initial fluid; (C) concentrations of selected

aqueous species, molality. Total bicarbonate is composed dominantly of HCO3� and H2CO3; total sulfate is composed dominantly of SO4

2�,

FeSO4, NaSO4�, and HSO4

�; total iron composed dominantly of Fe2+, FeCl+, and FeSO4, total calcium is composed of Ca2+ and CaSO4 (aq.).

J.L. Palandri, Y.K. Kharaka / Chemical Geology 217 (2005) 351–364362

Thermodynamic data for ankerite is absent from

the GAMSPATH database and therefore ankerite is

absent from the mineral assemblages. Siderite begins

to precipitate after ~1 h, and reactions involving iron

are largely complete when hematite dissolves out after

~14 h (Fig. 7B). Subsequently, dawsonite and calcite

precipitate at 31 and 175 h, respectively. The total

time to system equilibration is controlled by the

dissolution rate of labradorite, which is treated as

separate endmembers albite and anorthite to bracket

the fastest and slowest possible rates for feldspar

dissolution. The length of time for labradorite to

dissolve would be intermediate to the two endmem-

bers. The times for anorthite and albite to dissolve out

are 0.57 and 22,400 years (off scale right), respec-

tively. However, the values for p and q in Eq. (11)

were obtained at a pH of 9 and at 300 8C (Alekseyev

et al., 1997), conditions which are substantially

different than in the simulation. If p and q are set

equal to their default value of unity, the time for albite

to dissolve out is 0.63 years. There is clearly a high

degree of uncertainty in these kinetic simulations, and

they provide only a very rough estimate of the length

of time for equilibration. This simulation provides a

better estimate of the expected porosity decrease,

because the initial mineral assemblage is composed of

anhydrous minerals; the initial porosity is 21.0%, and

the final porosity is 19.9% (qanorthite=2.60 g/cm3,

qkaolinite=2.60).

4. Conclusions

Theoretical calculations show that ferric iron in

sediments can be used to trap CO2 in ferrous iron-

bearing carbonate minerals, by including SO2 or H2S

J.L. Palandri, Y.K. Kharaka / Chemical Geology 217 (2005) 351–364 363

in the injected waste gas stream to reduce the iron.

The iron can be made to reside almost entirely in

siderite or ankerite, with sulfur existing predomi-

nantly as dissolved sulfate. In rocks that contain only

ferric iron as a potential trap for CO2, the appropriate

gas composition is 2 mol of CO2 to 1 mol of SO2.

Given that the waste gas from fossil fuel-fired power

plants has a higher ratio of CO2 to SO2, targeted

sediments must contain additional divalent metals

such as Ca or Mg, or additional sulfur-bearing gas

must be added to the waste gas stream to effectively

trap CO2 in carbonate minerals. Conversely, most

sediments contain ferric iron in quantities sufficient to

effectively scrub SO2 from waste gas produced by

combustion of fossil fuel, by converting gaseous

sulfur into dissolved sulfate. Therefore, CO2 and SO2

can be injected together without the need to first

separate SO2 from the waste gas stream. It may be

advantageous to make use of fossil fuels with large

amounts of sulfur, or to add waste SO2 or H2S

derived from other industrial processes to injected

flue gas, to maximize the amount of iron that is

reduced and made available for precipitation in

siderite, especially in sediments where Ca and Mg

are scarce. In rocks containing calcium in addition to

iron, anhydrite and gypsum are less stable than calcite

or ankerite except at very high sulfate concentrations,

and anhydrite precipitation that might plug the

formation should not occur.

Most of these conclusions are based on idealized

computer simulations. Although the thermodynamic

part of the calculations is built upon a solid frame-

work, there exist considerable uncertainties in the

kinetics of mineral precipitation and dissolution, and

of aqueous speciation. Our best time estimates for

equilibration of CO2–SO2 sequestration in redbeds

ranges from hours to weeks, based on the simulations

herein and previous experimental work investigating

H2S sequestration. For sediments containing Ca and

Mg in silicates, equilibration times are less certain

because of the absence of, or uncertainty in kinetic

data for mineral dissolution and precipitation; in any

case dissolution of preexisting minerals is likely the

rate limiting step in the overall reaction, rather than

precipitation of secondary carbonate minerals. Further

investigation of CO2 sequestration using ferric iron

should include rigorous experiments with detailed

mineral and solution chemistry.

The use of iron-rich sediments, e.g. redbeds,

introduces a CO2 mineral trap commonly with high

porosity and permeability, with widespread geo-

graphic distribution and great thickness, and there-

fore potentially great capacity. The benefits are

twofold—not only can ferric iron be used as a trap

for CO2, but SO2 gas that may already be part of

the flue gas can be injected with the CO2, so that it

does not have to be scrubbed from the waste gas

stream.

Acknowledgments

We thank Robert Michel, Keith Kirk, Robert

Rosenbauer and two anonymous reviewers for

comments and suggestions which has lead to

substantial improvement of this manuscript. We also

thank Eric Oelkers for editorial handling and his

efforts in organizing this special issue of Chemical

Geology. [EO]

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