The combined effect of pH and temperature on smectite dissolution rate under acidic conditions

The effect of pH and temperature on smectite dissolution -1- 11/10/04

The combined effect of pH and temperature on smectite dissolution rate under acidic

conditions

Keren Amram and Jiwchar Ganor*

Department of Geological and Environmental Sciences, Ben-Gurion University of the Negev,

P. O. Box 653, Beer-Sheva 84105, Israel.

Phone ++972-8-6472-651 Fax ++972-8-6472-997 E-mail [email protected] WEB http://www.bgu.ac.il/geol/ganor/

Submitted to: Geochimica et Cosmochimica Acta, April, 2004

Revised and resubmitted, October, 2004

*Corresponding author.

Published in Geochimica et Cosmochimica Acta (2005) 69, 2535–2546


ABSTRACT

The main goal of this paper is to propose a new rate law describing the combined effect

of pH (1 to 4.5) and temperature (25°C to 70°C) on smectite dissolution rate, under far from

equilibrium conditions, as a step towards establishing the full rate law of smectite dissolution

under acidic conditions. Dissolution experiments were carried out using non-stirred flow-

through reactors fully immersed in a thermostatic water bath held at a constant temperature of

25.0°C, 50.0°C or 70.0°C ± 0.1°C. Smectite dissolution rates were obtained based on the

release of silicon and aluminum at steady state. The results show good agreement between

these two estimates of smectite dissolution rate. Low Al/Si ratios were obtained in

experiments that were conducted at pH≥ 4. These low Al/Si ratios are explained by

precipitation of gibbsite and/or diaspore.

Dissolution rate increases with temperature and decreases with increasing pH.

Dissolution rates of experiments in which ∆Gr≤-21kcal mol –1, are not affected by deviation

from equilibrium. Dissolution rates in most experiments are not affected by the addition of

up to 0.3 M NaNO3 to the input solution.

A simple model is used to describe the combined effect of pH and temperature on

smectite dissolution rate. According to this model, dissolution rate is linearly proportional to

the concentration of adsorbed protons on the mineral surface, and proton adsorption is

described using a Langmuir adsorption isotherm. All experimental results at pH<4 were fitted

to the model using a multiple non-linear regression. The resulting rate law is:

+

+

⋅⋅−⋅+

⋅⋅−⋅⋅−⋅=

H

HaRTe

aRTeRTeRate /1070061031

/107006103/17460220

where R is the gas constant, T is the temperature (K) and aH+ is the activity of protons in

solution.

According to the model, the dependence of dissolution rate on temperature is affected

by the activation energy and the adsorption enthalpy. The activation energy obtained from the

fitting (17±2 kcal/mol) is within error equal to the average value of 15 kcal/mol of apparent

activation energies for silicates dissolution rate (Lasaga et al., 1994). The obtained net


enthalpy of adsorption (-11±2 kcal/mol) is within the range of –7.9 to –23.1 kcal/mol,

experimentally obtained for oxides (Sverjensky and Sahai, 1998).

1 INTRODUCTION

Smectites are phyllosilicates with two tetrahedrally coordinated layers and one

octahedrally coordinated layer. Montmorillonites are the most abundant smectites in nature.

The structural formula of an ideal montmorillonite is:

M0.66[Mg0.66Al3.34][Si8]O20(OH)4·(H2O)n. M0.66 represents 0.66 monovalent cations or 0.33

divalent cations. In natural montmorillonites, the tetrahedrally coordinated Si is substituted

by up to 0.4 Al atoms, and octahedrally coordinated Al is considerably substituted by Mg,

FeIII or FeII. The structural charge resulting from the tetrahedral and octahedral substitution is

compensated by the presence of exchangeable interlayer cations, denoted by M in the

structural formula.

Smectites are generated in soils and sedimentary deposits by weathering, diagenesis or

hydrothermal processes, which can either involve degradation and transformation of

precursor phyllosilicates or precipitation from solution. Smectite is the most abundant mineral

in soils of temperate climates that are derived from basic igneous rocks, and from the clay

fraction of soils of arid climates derived from granitic pediment (Boettinger and Southard,

1995). Smectite, particularly montmorillonitic smectite, is the principal constituent of

bentonite clay deposits. These have been formed by the alteration of eruptive igneous rocks of

basic to intermediate composition.

Smectite-rich argillites and bentonites have been recognized as suitable clays to be used

as a sealant material in the multibarrier systems designed for storage of high level nuclear

waste in burial repositories (Chapman and McKinley, 1989). Due to its osmotic swelling

capacity (and consequently its plasticity and impermeability), smectite impedes groundwater

interaction with the metal canisters. The cation exchange reactions immobilize undesirable

cations from the radioactive waste, and retard its leakage from the canisters towards the local

groundwater. However, the durability of the smectite itself under confinement conditions is a

key parameter that must also be taken into consideration. For a repository in an argillaceous


formation, acidification may result from oxidation of pyrite, a common mineral phase in

claystone.

1.1 Previous studies on smectite dissolution rate and the motivation for the

present study

Only a few studies of smectite dissolution kinetics have been carried out in the last

decade (Furrer et al., 1993; Zysset and Schindler, 1996; Bauer and Berger, 1998; Cama et al.,

2000; Huertas et al., 2001; Metz, 2001; Metz et al., 2004a). All the experiment at 25°C and

the experiments of Metz (2001) and Metz et al. (2004a) at 50°C were conducted under acidic

conditions, whereas those at 20, 35, 40 and 60 to 80°C were conducted under basic (mostly

extremely basic) conditions. Furrer et al. (1993) and Zysset and Schindler (1996) conducted

both batch and flow-through dissolution experiments to study the proton-promoted

dissolution kinetics of K-montmorillonite. The dissolution experiments were conducted at

25°C in HCl / KCl solutions (pH = 1 to 5 and [KCl] = 0.03 to 1.0 M). The dissolution rate

increased with KCl concentrations and decreased with pH. At pH ≥ 3 the dissolution rate was

inhibited by aluminum (Furrer et al., 1993; Zysset and Schindler, 1996). Their observed

released rate ratio, RSi/RAl, depends on both the pH and KCl concentration. Bauer and Berger

(1998) conducted batch experiments at 35° and 80ºC in concentrated KOH solutions (0.1 to 4

M) to study the dissolution kinetics of industrial (Ibeco and Ceca) montmorillonite. They

found that under very basic conditions (11.5 ≤ pH ≤ 13.9) smectite dissolved independently of

aqueous silica or aluminum concentrations. They proposed a non-linear dependency of smectite dissolution rate, 06.0150 ±⋅= .

OH -a kRate . The apparent activation energy was found to

be 13 ± 1 kcal/mole. Huertas et al. (2001) measured bentonite dissolution in granitic

solutions (pH 7.6 to 8.5) in a semi-batch reactor at 20, 40 and 60°C. Sato et al. (2002; 2003)

conducted flow-through experiments on Na-montmorillonite dissolution at 30, 50 and 70°C at

pH 8.6 to 13.3 in NaOH-NaCl and KOH-KCl solutions. The effect of deviation from

equilibrium on the dissolution rate of smectite was studied by Cama et al. (2000) at pH 8.8

and 80°C, and by Metz (2001) at pH 3 and 50°C. Both studies show that smectite dissolution

rate is independent of deviation from equilibrium in the range of ∆Gr<-30 kcal mol –1. In

contrast to the observations of Furrer et al. (1993) and Zysset and Schindler (1996), Metz


(2001) found that far from equilibrium smectite dissolution rate is independent of both Al and

Si concentrations.

Besides dissolution kinetic studies on smectite powders employing batch and flow-

through reactors, a few papers about in-situ atomic force microscope (AFM) studies on

dissolution rates of single-cell smectite crystals have been recently published (Bickmore et

al., 1998; Bickmore et al., 1999; Bosbach et al., 2000; Bickmore et al., 2001). Smectite

dissolution rates determined in AFM studies appear to be faster than the rates determined in

flow-through or batch reactor studies at the same temperature and pH.

As part of our attempt to establish the full rate law of smectite dissolution under acidic

conditions, the present study introduces a new data set examining the effect of both

temperature (25°C to 70°C) and pH (1 to 4.5) on smectite dissolution rate. Cama et al. (2002)

demonstrated that in order to study the effect of temperature on kaolinite dissolution rate one

should distinguish between two effects: 1) the effect of temperature on the rate coefficient,

which may be modeled using the Arrhenius equation; and 2) the effect of temperature on the

adsorption of protons on the surface. The last effect cannot be studied independently of the

pH effect on the rate. Therefore, they presented a model describing the combined effect of pH

and temperature on kaolinite dissolution rate. The main objective of the present paper is to

model the experimental data using a rate law describing the combined effect of pH and

temperature, following the approach of Cama et al. (2002).

2 MATERIALS AND METHODS

2.1 Characterization and pretreatment of smectite

The smectite sample used in this study is the SAz-1, an international reference sample

of the Clay Mineral Society Source Clay Repository. Details about the origin of these

standard clay samples were published by Van Olphen and Fripiat (1979) and Moll (2001).

The sample is not a pure smectite, containing considerable amounts of silica phases (quartz,

cristobalite and/or amorphous silica) and minor quantities of alkali-feldspars, plagioclase and

carbonates. The structural formula of pure SAz-1 was determined by Metz et al. (2001;

2004a) to be K0.02Na0.05Ca0.41Mg0.18[Mg1.11Fe0.17Al2.77][Al0.30Si7.70]O20(OH)4. Using this structural


formula and the whole rock analysis, Metz et al. (2004a) calculated that the raw smectite

contains about 87% smectite, 8% excess SiO2 (probably mostly in amorphous silica). The

rest is mostly hydrated water.

The smectite was pretreated in 0.001 N HNO3 at 70°C for a few months, using the

procedure described in Ganor et al. (1995). The sample is composed of smectite aggregates

ranging in size from less than a micron to more than 100 micron (Metz et al., 2004a). BET

surface areas of raw SAz-1 sample determined in previous studies vary between 34 and 97 m2

g-1 (Metz et al., 2004b). Following dissolution experiments the BET surface area of SAz-1

increased to 127±13 m2 g-1. Using atomic force microscope, Metz et al. (2004b) determined

the edge surface area of SAz-1 to be 4.9 ± 0.7 m2 g-1.

2.2 Experimental setting

Dissolution experiments were carried out using non-stirred flow-through reactors (ca.

35 ml in volume) fully immersed in a thermostatic water-bath held at a constant temperature

of 25.0°C, 50.0°C or 70.0°C ± 0.1°C (Fig. 1). The reaction cells were composed of two

chambers, a lower chamber of 33-mm inner diameter and an upper chamber of 26-mm inner

diameter. The two chambers were separated by a fine (5 µm) nylon mesh, on which smectite

powder was placed. Some more details of the experimental procedure can be found in Metz

and Ganor (2001).

2.3 Solutions and Analyses

Input solutions were prepared at specific pH by diluting 1M HNO3 with double

deionized water. In experiments designed to study the effect of ionic strength, different

amounts of NaNO3 have been added into the input solution.

Input and output solutions were analyzed for Al, Si, and pH. Total Al and Si were

analyzed colorimetrically with a UV-visible spectrophotometer, using the Catechol violet

method (Dougan and Wilson, 1974) and Molybdate blue method (Koroleff, 1976),

respectively. The uncertainty in measured Al and Si was better than ± 5% for concentrations

above 4 µM. The precision dropped to ±15% and 33% for measurements at low

concentrations of 2 and 0.5 µM, respectively. The pH was measured at experimental


temperature on an unstirred aliquot of solution using a semi-micro 83-01 Orion Ross

combination electrode. The reported accuracy is ±0.02 pH units (±4.5% in H+ activities).

3 CALCULATIONS

The overall dissolution reaction of smectite sample SAz-1 under acidic conditions can

be expressed as:

(1) 44

332224207.73.077.217.011.118.041.005.002.0

7.707.317.029.141.005.002.0

8.62.13)())()((

SiOHAlFeMgCaNaK

OHHOHOSiAlAlFeMgMgCaNaK

++++++

⇔++++++++

+

The dissolution rate, Rate, (mol g-1 s-1) in steady state was based on the release of Al and Si

according to the expression:

(2) ),,( inpjCoutjC

mqRatej −−=⋅ν

where Cj,inp and Cj,out are the concentrations of component j (Al or Si) in the input and the

output solutions, respectively (mol m-3), νj is the stoichiometry coefficient of j in the

dissolution reaction, t is time (s), m is the sample mass (g) and q is the fluid volume flux

through the system (m3 s-1). Note that in our formalism, the rate is defined to be negative for

dissolution and positive for precipitation.

The common practice in experimental kinetics is to normalize the dissolution rate to the

total surface area of the pure mineral, which is measured by the Brunauer-Emmett-Teller

(BET) method, (Brunauer et al., 1938). In contrast to this common practice, Furrer et al.

(1993), Schlabach et al. (1999) and Zysset and Schindler (1996) normalized their smectite

dissolution rate data to the sample mass. They argued that normalization to surface area is

not appropriate, as long as it is not possible to measure exactly the extent of the edge surface

area of a smectite powder. Indeed, Metz et al. (2004b) showed that there is no correlation

between the total and the edge surface area of smectite, and as a result the BET surface area

cannot serve as a proxy for the reactive surface area of smectite. Therefore, we normalized

the dissolution rates by the mass of the smectite. The initial mass of the smectite in each

experiment was calculated from the product of the starting mass and the estimated percentage

of the smectite in SAz-1 (87%, Metz et al., 2004a). Following each stage (i.e., the time


between the replacements of two sequential output bottles), the remaining mass of each

mineral was updated based on the release rate of Al and the duration of the stage. The release

rate of Al and not of Si was used in the updating procedure following the conclusion of Metz

et al. (2004a) that the initial fast release of Si, before steady state, mainly reflects the

hydrolysis of a fast dissolving silica phase, while the initially slow release of Al reflects the

dissolution of the smectite itself. The error in the calculated rate is estimated using the

Gaussian error propagation method (Barrante, 1974) from the equation:

(3) ( )

2/12

2

∆

∂∂=∆ ∑

ii

ix

xPP

where ∆P is the uncertainty of the calculated parameter and ∆xi is the estimated uncertainty of

the measurements of the quantity xi.

The degree of saturation of the solution with respect to smectite dissolution is

calculated in terms of the Gibbs free energy of reaction ∆Gr

(4) )ln(

eqr K

IAPRTG =∆

where R is the gas constant, T is the absolute temperature, IAP is the ion activity product of

the solution, and Keq is the solubility constant. The activity coefficients and the activities of

the different species in solution were calculated using the EQ3NR module of the EQ3/6

software package (Wolery, 1992). Errors (∆P) in the above-calculated parameters (P), i.e.,

IAP and ∆Gr, were estimated according to the Gaussian error propagation equation (3). In the

absence of a specific Keq value for smectite sample SAz-1, the solubility constant of the

EQ3/6 thermodynamic data-set data0.com.R22a (Wolery, 1992) for Ca-endmember smectite

(Ca0.33[Mg0.66Al3.34][Si8]O20(OH)4) were used in the calculations. This proxy for the

solubility constant was selected so the calculation would be consistent with those of Metz

(2001).

4 RESULTS

The variations of output Al and Si concentrations in three representative flow-through

experiments as a function of time are shown in Fig. 2. Figures of all dissolution experiments


are presented in Amram (2002). Each of the experiments was composed of 1 to 5 stages,

where new stages were initiated by a change in the flow rate (e.g., 50-03.2, Fig. 2b) and/or in

the composition of the input solution (e.g., 70-15.2, Fig. 2c). The vertical lines in Fig. 2

delineate the different stages. Much of the noise in the non-steady state data results from

instabilities in flow rate. In several experiments, an established steady state was disturbed by

a long period of instabilities in flow, and thereafter a new steady state was established. In

such cases the new steady state was reported as a different stage (e.g., 50-03.3, Fig. 2b). The

experimental conditions of all the stages are compiled in Table 1. The first two digits in the

names of the experiments denote the experimental temperature. The last digit in the name

(following a dot) marked the consecutive stage number of the experiment. Al and Si analyses

used to calculate steady-state compositions are denoted by open symbols in Fig. 2. Si

concentrations were usually higher at the onset of the experiments (Fig. 2), after which Si

concentrations decreased until steady state was attained. A reversed trend was observed for

the release of Al, which increased usually from very low starting concentrations to higher

concentrations at steady state. The pH decreased continuously until steady state, where pH in

output solutions was close to the pH in the input solution (generally up to 0.2 pH units

higher).

Duration of experiments varies but mostly surpasses 1500 hours, and some last for more

than 10000 h. The time required to achieve the first steady state varied considerably. It was

usually shorter under conditions in which the dissolution rate was fast (high temperature and

low pH) and under higher ionic strength. In most of the experiments, steady state was easily

maintained for several hundred hours (up to 2500 h), as long as the flow rate was stable (e.g.,

Figs. 2a and b). The exceptions are experiments conducted under conditions in which

smectite dissolved very fast. In these experiments the mass of smectite decreased

significantly with time, and as a result the output Al and Si concentrations decreased as well

(e.g., 70-15.2, Fig. 2c). Several studies (e.g., Walther, 1996; Gautier et al., 2001) indicate

that the amount of time prior to steady state may influence the resulting steady-state

dissolution rate. Figure 2b shows the change in Al and Si concentrations in a multi-stage

experiment (50-03) at 50°C in which the flow rate varied between 0.013 and 0.04 ml/min.

The experiment attained the first steady state after about 6500 h. The dissolution rate at


steady state was 2.8±0.5x10-12 mol g-1 s-1. After about 2000 h at steady state the flow rate

increased to 0.04 and a change in concentration was observed. The dissolution rate at this

steady state was the same, 2.8±0.5x10-12 mol g-1 s-1. Following a period of instability in flow

rate a third steady state was obtained in which the rate was 2.6±0.5x10-12 mol g-1 s-1.

Regardless of the changes with time the same dissolution rate was observed in the three

stages.

Release of elements was highly incongruent during the first few hundreds to few

thousands hours of the experiments. Metz et al. (2004a) showed that the enhanced release of

Si reflects the hydrolysis of a silica phase which dissolves faster than smectite, while the

initially slow release of Al reflects the dissolution of smectite itself. As we thoroughly

discussed this initial non-congruent stage of smectite dissolution in Metz et al. (2004a), the

present study is discussing only the steady state dissolution rate.

Smectite dissolution rates at steady state (eq. (2)) were obtained based on the release of

silicon (RateSi), and aluminum (RateAl) at steady state for each stage (Table 1). Figure 3 plots

the dissolution rates evaluated based on the release of Si versus those obtained based on the

release of Al. The solid lines in Fig. 3 are the 1/1 diagonals. Taking into account the

appropriate errors, a good agreement between the two estimates of smectite dissolution rate is

observed in most of the experiments. In these experiments, the steady-state dissolution rates

are measured at conditions of under-saturation with respect to gibbsite, diaspore, kaolinite

and boehmite. Some experiments conducted at pH ≥4 in which equilibrium with respect to

these minerals is achieved and, consequently, incongruent dissolution occurs. In these

experiments the dissolution rate is based only on Si release.

In general, pretreated smectite samples were used in the flow-through experiments. Raw

samples of smectite SAz-1 underwent nine flow-through experiments (Table 1). The same

kinetic behavior was observed in experiments with pretreated and raw SAz-1.


5 DISCUSSION

5.1 Are the calculated steady-state dissolution rates influenced by the presence of

accessory phases in the smectite sample?

As the dissolved sample contains about 8 wt.% excess SiO2 it is important to verify that

the dissolution of the Si-rich accessory phases does not affect the release of Si at steady-state.

Metz et al. (2004a) discussed the general case in which a mixture of a major phase (smectite

in the case of SAz-1) and a minor phase (Si-rich phase) is dissolved. Following their

argumentations, there are three possible scenarios: 1) the half life of the smectite is

significantly shorter than that of the Si-rich phase (i.e., the smectite dissolves faster); 2) the

two phases have similar half lives; and 3) the half life of the smectite is significantly longer

than that of the Si-rich phase. In the first case, the Si-rich phase is both less abundance and

less reactive than the smectite and therefore the contribution of the Si-rich phase is expected

to be negligible. In the last case, the Si-rich phase would be extinct before steady state, due to

its shorter half life (e.g., Fig. 6 of Metz et al. (2004a)). Only in the second case, in which the

Si-rich phase has a half life similar to that of the smectite, the steady state would reflect the

dissolution rate of the two phases.

The following observations indicate that this is not the case: 1) Metz et al. (2004a)

showed that the initial fast release of Si during the dissolution of SAz-1 (Fig. 2) reflects the

hydrolysis of a silica phase which dissolves faster than smectite. This fast dissolving Si-rich

phase is accounted for more than 60% of the excess Si in SAz-1 (Metz et al., 2004a). 2) Metz

et al. (2004a) measured the composition of raw SAz-1 using XPS. Since this measurement

has an information depth of 6 nm, the measured composition is hardly obscured by

contaminants, and is therefore a good approximation for the composition of pure SAz-1. In

the present study, the average Al/Si ratio at steady state is within error equal to the Al/Si ratio

of the XPS measurement of the raw sample, indicating that the steady-state ratio is not

significantly influenced by the dissolution of phases with significantly different Al/Si ratio.

3) Metz et al. (2004a) recovered powder from their flow-through experiments with SAz-1 and

found that the recovered powder contained significantly less excess Si. The average steady

state Al/Si ratio in the present study is within error equal to the Al/Si ratio of in the recovered


powder, which was analyzed by both SEM-EDS and XPS. Therefore, we conclude that the

release of Al and Si at steady-state is due to the dissolution of the smectite itself, and can be

used to calculate the dissolution rate of the smectite.

5.2 Separating the direct effects of pH and temperature from effects of other

environmental variables

In order to model the effects of pH and temperature on smectite dissolution rate it is

important to separate between direct and indirect effects of the environmental variables

involved. By direct effect we mean an effect related to surface processes and therefore, one

that can be used to understand the reaction mechanism. In addition to temperature and pH,

other environmental variables such as output Al and Si concentrations, ionic strength, and the

degree of saturation vary between the experiments. Therefore, their possible effects on the

smectite dissolution rate are examined below.

5.2.1 The effect of degree of saturation

The Gibbs free energy of reaction (∆Gr) of the smectite dissolution reaction is a strong

function of pH and of Al and Si concentrations. If the dissolution rate varied due to changes

in ∆Gr in the different experiments, as the pH and the output concentrations varied, then the

calculated effect of pH would include a spurious contribution. This problem would be

minimal only in the "far-from-equilibrium" dissolution plateau region, which is defined as the

region in rate vs. ∆G space where there is no direct effect of the degree of saturation on

dissolution rate. Metz (2001) studied the effect of deviation from equilibrium on smectite

dissolution rate under acidic conditions. According to his study, near equilibrium (0 ≥ ∆Gr ≥

-20 kcal/mole) the rates increase gradually with increasing undersaturation. Far from

equilibrium, at ∆Gr ≤ -30 kcal/mole, the dissolution rate is much faster and is independent of

the degree of saturation. The transition between these two regions occurs somewhere in the

range of -20 ≥ ∆Gr ≥ -30 kcal/mole. The ∆Gr during the first steady state in experiment 50-

03 (Fig. 2b) was ∆Gr =-20.1 kcal/mol. In order to examine if this experiment is conducted

under the conditions of the dissolution plateau, we increased the flow rate by a factor of 3.2.

As a result the Al and Si concentration decreased and the average ∆Gr during the new steady


state was -29.8 kcal/mol. Figure 2b shows that the concentration of Si decreases by the same

factor (3.2) as the increase in flow rate, and as a result the dissolution rates during the two

steady states were the same. If the experiment was close to equilibrium during the first steady

state, the consequential decrease in ∆Gr would bring to an increase in smectite dissolution

rate. Therefore, we conclude that the dissolution rates in experiments with ∆Gr≤-20.1

kcal/mol were independent of the deviation from equilibrium. According to the results of

experiment 50-03 the dissolution plateau region for smectite is in the range of ∆Gr≤-20.1

kcal/mol, which was the ∆Gr range of most of the experiments (Table 1). The results of

Mogollon et al. (1996) showed that the dissolution plateau for gibbsite at 25°C was in very

good agreement with the results of Nagy and Lasaga (1992) at 80°C. Assuming that the

dissolution plateau for smectite is similarly independent of temperature, the dissolution rates

in most of the experiments were independent of the deviation from equilibrium. Some

experiments, however, were conducted under close to equilibrium conditions (Table 1). All

these experiments where conducted at pH>4. Therefore, these experiments are not used in the

fitting of the proposed model.

5.2.2 The effects of silicon and aluminum on smectite dissolution rate

Metz (2001) examined the effects of silicon and aluminum on smectite dissolution rate

at 50°C. He found that under far from equilibrium conditions (∆Gr ≤ -30 kcal mol-1) smectite

dissolution rate is independent both of Al concentration (ranging between 3 to 16 µM at pH 3

and between 21 and 139 µM at pH2) and of Si concentration (ranging between 9 and 41 µM

at pH 3 and between 56 and 329 µM at pH2). This rate independency is also supported by an

observation from the present study that few experiments that were conducted at the same pH

and temperature and with different steady-state Al and Si concentrations showed similar

dissolution rate. We suggest that the changes in far from equilibrium dissolution rates

observed in the present study (Fig. 4) are not significantly influenced by the variability in Al

and Si concentration. It is important to note that the present study ranges of Al (0.05-843

µM) and Si (2.8-2387 µM) concentrations are significantly larger than the ranges which were

thoroughly studied by Metz (2001). Therefore, we cannot rule out the possibility that some of

the experimental noise may be related to effects of Al or Si on dissolution rate.


5.2.3 The effects of ionic strength, Na+ and NO3-

To establish the effect of pH on dissolution rate, one should conduct a series of

experiments that are identical in all factors except pH. However, adding acid changes the

ionic strength and the concentration of the balancing anion. Therefore, studying the effect of

pH may be conducted using two possible experimental designs: 1) The input solution contains

only an acid and therefore the concentration of the balancing anion (NO3 in the present study)

varies in the different experiments and is equal to the H+ concentration; 2) The input solution

is composed of a mixture of an acid and a salt (NaNO3 in the present study), so the

concentration of the balancing anion and the ionic strength are the same in all the

experiments. In this design the concentration of the salt cation varies in the different

experiments. In order to assess the effects of ionic strength and Na and NO3 concentrations

on the determination of the observed pH dependence of smectite dissolution rate, we used

both experimental designs. About half of the experiments were conducted with input solution

composed solely of HNO3. The ionic strength in these experiments ranges from 0.000032 M

(pH 4.5) to 0.1 M (pH 1). In the rest of the experiments the ionic strength and NO3-

concentration of ~ 0.32 M were maintained by adding suitable amounts of NaNO3 into the

input solutions. By doing this, Na+ concentration increases as the pH increases whilst NO3-

concentration remains constant (~0.32 M). Figure 4 compares dissolution rates in

experiments conducted under constant ionic strength to those obtained without adding

NaNO3, i.e., where H+=NO3-. In most experiments at pH<4 the differences between the two

sets are small. The exceptions are the 25°C experiments at pH < 1.8 and the 50ºC at pH of 3

and 2, where the dissolution rate under constant ionic strength is significantly slower than that

in the experiments in which H+=NO3-. Above pH of 4 the dissolution rates under constant

ionic strength are significantly slower at 50 and 70°C and faster at 25°C than those in the

experiments in which H+=NO3-. Taking into account the experimental noise, it is hard either

to prove that the observed differences between the experiments represent a real effect of the

addition of salt, nor to rule out this possibility. For the purpose of modeling the effects of pH

and temperature on smectite dissolution rate we decided to use the results of all the

experiments conducted at pH<4, regardless of their ionic strength. As the effect on the rate is


small to insignificant, the results of the modeling will be only slightly influenced by this

decision.

5.3 Modeling the effect of pH and temperature on dissolution rate

The model proposed below is a simple version of the model proposed by Cama et al.

(2002). The model is based on two assumptions: 1) The proton promoted reaction mechanism

consists of fast adsorption of a proton on a surface site followed by a slow hydrolysis step;

and 2) The adsorption of the protons on the surface site may be described by a simple

Langmuir adsorption isotherm:

(5) +

+

⋅+⋅

=H

HadsH ab

abFX

1,

where F is the maximum surface coverage of protons on the reactive surface site, b is a

constant related to the energy of adsorption and aH+ is the activity of protons in solution.

Adsorption of a proton on a surface site close to the metal influences the bond strength and

thus affects the dissolution rate. If steady-state conditions are maintained, the rate of this

reaction path is (Lasaga, 1981):

(6) adsHXksRate

,⋅=⋅ ρ

where k (s-1) is the rate coefficient, Rate is the observed dissolution rate (mol g-1 s-1) , s is the

specific surface area (m2 g-1), ρ (mol m-2) is the density of reactive surface sites on the mineral

surface and XH,ads is the molar fraction of the surface site that is protonated. Substituting the

Langmuir adsorption isotherm (equation (5)) into equation (6), gives:

(7) +

+

⋅+⋅

⋅=⋅ H

Hab

abFk

sRate

1ρ

Both the rate coefficient, k, and the adsorption coefficient, b, in equation (7) depend on

temperature. The temperature dependence of the dissolution rate generally follows the

Arrhenius law:


(8) RTaEAek /−=

where A (s-1) is the pre-exponential factor, Ea is the apparent activation energy, R is the gas

constant and T is the temperature (K). The temperature dependence of the adsorption

coefficient may be evaluated recalling that the b constant in the Langmuir adsorption

isotherm is the equilibrium constant of the protonation reaction, and therefore its temperature

dependence may be written as:

(9) RTHRTHRS eKeeb /

0// 000 ∆−∆−∆ ⋅=⋅=

where ∆S0 (cal mol-1 K-1) is the entropy, and ∆H0 (kcal mol-1) is the net enthalpy of

adsorption. Following Sverjensky and Sahai (1998), the standard states for both surface and

aqueous species are assumed to reflect hypothetical 1 molal solutions referenced to infinite

dilution and a surface potential of zero at 25°C. The temperature dependence of the b

constant in the Langmuir adsorption isotherm may be described using equation (9) and

assuming that the heat capacity, ∆Cp, is equal to zero, and therefore ∆H0 is temperature

independent. This last assumption was examined and justified for kaolinite by Cama et al.

(2002).

The combined effect of pH and temperature on smectite dissolution rate may be

described by substituting equations. (8) and (9) into equation (7):

(10) +

+

⋅⋅+

⋅⋅⋅⋅=

⋅ ∆−

∆−−

HRTH

HRTH

RTEaaeK

aeKFeA

sRate

/0

/0/0

0

1ρ

The coefficients k'=A.F. s.ρ, Ea, K0 and ∆H0 were calculated from a multiple non-linear

regression of equation (10) using least squares. For the fitting we used all the experimental

results at 25, 50 and 70°C at pH<4. The resulting coefficients are k'=220±750 mol g-1 s-1,

Ea=17460±2000 cal mol-1, K0=3±14.10-6 and ∆H0=-10700±2500 cal mol-1. The regression

coefficient is R2=0.93. Substituting these values into equation (10) yields,

(11) +

+

⋅⋅−⋅+

⋅⋅−⋅⋅−⋅=

H

HaRTe

aRTeRTeRate /1070061031

/107006103/17460220


A comparison between the prediction of equation (11) and the experimental data at 25°C,

50°C and 70°C is shown in Fig. 4.

The obtained activation energy, 17±2 kcal/mol, is similar to the activation energy of the

dissolution of the edge site of kaolinite (22 kcal/mol, Cama et al., 2002) and is within error

equal to the average value of 15 kcal/mol of apparent activation energies for silicates

dissolution rate (Lasaga et al., 1994). The obtained net enthalpy of adsorption of -11±2

kcal/mol is within the range of –7.9 to –23.1 kcal/mol, experimentally obtained for oxides

(Sverjensky and Sahai, 1998). The reasonable activation energy and net enthalpy of

adsorption, although not proving the proposed model, provide support for its validity.

5.4 Comparing the prediction of the model with measured adsorption isotherms

The relative adsorption/desorption of protons on mineral surfaces is commonly

measured using potentiometric surface titration. The term "surface titration" is somewhat

misleading as the measurements of the adsorption of proton onto the mineral surface are

based on changes in the pH of the solution. The relative surface concentration of protons is

determined by mass balance between the proton (or hydroxide) added to solution and the

measured proton concentration in solution after equilibration using the so-called proton

consumption function

(12) AVOHHCCC BAs ⋅+−−=∆ −+ ])[][(

where ∆Cs is the change in surface concentration of protons (mol m-2), CA and CB are the

concentrations of the acid and base added (mol l-1), respectively, [H+] and [OH-] are the

solution concentrations of H+ and OH- after equilibration (mol l-1), V is the fluid volume (l)

and A is the total surface area (m2). It is important to note that the value of the proton surface

charge obtained by surface titration is arbitrary until a value of zero charge is established

(Davis and Kent, 1990; Ganor et al., 2003), i.e., the titration measures a relative change in

surface concentration (as defined by the proton consumption function, Eq. (12)) and not the

absolute concentrations (Schroth and Sposito, 1997). Therefore, the calculation of the

absolute proton adsorption density is based on an assumption regarding the pHPZNPC. Figure 5

shows the surface titration data obtained by Zysset and Schindler (1996) at 25°C between pH


4 and 1. According to the interpretation of Zysset and Schindler (1996) the surface

protonation at this pH range is controlled by both adsorption of H+ on aluminol site and ion

exchange reaction: at pH≥2.5 protonation of the aluminol is dominant (black dots in Fig. 5),

whereas at pH<2.5 ion exchange significantly contributes to the concentration of adsorbed H+

(squares in Fig. 5).

A byproduct of the fitting of the proposed model (eq. (10)) is that it predicts the molar

fraction of surface protonation. This prediction of the proposed model may be compared to

protonation data of Zysset and Schindler (1996). Such a comparison is not straightforward.

Zysset and Schindler (1996) measured the concentration of adsorbed H+ ions (mol g-1), which

results from several reactions including adsorption and cation exchange (Stadler and

Schindler, 1993; Zysset and Schindler, 1996). In contrast, the fitting of the proposed model

predicts the molar fraction (and not the total concentration) of a protonated edge site that

governs the dissolution rate under the examined pH range. It is important to note that the rate

law proposed in the present study assumes that the dissolution rate is proportional to the

concentration of protons that are adsorbed on the reactive edge surface site and is independent

of protonation due to ion exchange reaction. This is the major difference between the present

study rate law and that of Zysset and Schindler (1996), who proposed that the rate is linearly

proportional to the sum of the proton concentrations on the aluminol and the cation exchange

sites (see eq. 20 in Zysset and Schindler, 1996).

Based on the present study prediction, the total concentration of adsorbed H+ ions

({H+}, mol g-1) may be calculated by multiplying the predicted molar fraction at the reactive

edge site by the total concentration of this site (θ, mol g-1).

(13)

+

++∆−

∆−++

⋅⋅−⋅+

⋅⋅−⋅⋅+=

⋅⋅+

⋅⋅⋅+=

+

+

H

H

HRTH

HRTH

aRTe

aRTeH

aeK

aeKHH /1070061031

/107006103}{

1}{}{ 0/

0

/0

0 0

0

θθ

In the calculations we assume that the concentration of adsorbed H+ ions on sites that are not

influencing the rate is constant in the examined pH range and their sum equals {H+}0. The

term {H+}0 may include contribution from the permanent charge of the mineral surface as

well as the real surface charge at the pH that was assumed to be the pHPZNPC. The coefficients


{H+}0 and θ may be obtained by fitting Eq. (13) to adsorption isotherm data using least

squares. As the dissolution rate is assumed to be independent of the extent of ion exchange,

and as according to Zysset and Schindler (1996) ion exchange significantly contributes to the

observed {H+} at pH<2.5, only the seven data points at pH≥2.5 (black dots in Fig. 5) were

used for the fitting. The solid line in Fig. 5 is the best-fit curve obtained for the surface

titration of Zysset and Schindler (1996) at 25°C in the pH range of 2.5 to 4. The obtained

coefficients are {H+}0=8.3±0.2.10-5 mol g-1 and θ=2.42±0.08.10-4 mol g-1, and the regression

coefficient R2=0.994. Figure 5 shows that the model proposed in the present study predicts

the concentration of adsorbed H+, between pH 1.5 and 4, although only the data at pH≤2.5

were used for the fitting. Following Zysset and Schindler (1996), we suggest that the excess

concentration of adsorbed H+ below pH 1.5 is due to ion exchange. According to Zysset and

Schindler interpretations below pH 2.5 ion exchange significantly contributes to the

concentration of adsorbed H+. In contrast, the prediction of the present study indicates that

the contribution of ion exchange is significant only below pH 1.5.

Surface titration curves are commonly interpreted in terms of protonation and

deprotonation of surface sites (see for example Davis and Kent, 1990; Parks, 1990;

Lutzenkirchen and Kienzler, 2002, and references therein). This interpretation involves a

critical stage in which the amount of protonation and deprotonation reactions is assumed. In

most cases, this assumption is based on the minimum amount of sites that are required to

obtain an adequate agreement between the surface speciation model and the observations.

Mathematically, it is always possible to add more reactions while keeping (or improving) the

quality of the fitting. The comparison between the predictions of the present study and the

observations of Zysset and Schindler (1996) is based on the assumption that between pH 2.5

and 4 the protons consumption is controlled by the protonation of a single surface site.

Neither the present study, nor the study of Zysset and Schindler (1996) provided any evidence

supporting this assumption. Taking into account that smectite contained many different

surface and interlayer sites that may adsorb and exchange protons, this assumption may be

questioned. Moreover, Stadler and Schindler (1993) argued that smectite protonation below

pH 4 is possibly dominated by ion exchange. Therefore, although the predictions of the

proposed model show a very good agreement with the total concentration of adsorbed H+


between pH 1.5 and 4, this agreement cannot serve as a proof for the proposed model, and the

obtained value of the total concentration of the reactive edge site (θ) should be regarded with

caution.

5.5 Comparison of the results of the present study to those of previous studies

Figure 6 compares smectite dissolution rates that were obtained in the present study to

results of Furrer et al. (1993), Zysset and Schindler (1996) and Metz et al. (2004a). Furrer et

al. (1993) conducted both flow-through and batch dissolution experiments with K-

montmorillonite sample SWy-1. All the experiments were conducted at 25°C, using HCl/KCl

solutions. In most of their experiments the ionic strength was adjusted to 0.1M. The rates of

the flow-through dissolution experiments of Furrer et al. (1993) are generally slower than the

rates of their batch experiments (Fig. 6a). A comparison between the results of the present

study and those of Furrer et al. (1993) shows a similarity between the dissolution rates in the

flow-through experiments of the present study and those in the batch experiments of Furrer et

al. (1993). Zysset and Schindler (1996) conducted three sets of batch dissolution experiments

in which the solutions contained 0.03, 0.1 and 1M KCl, using the same sample as Furrer et al.

(1993). The experiments conducted at pH<2 and 0.03M KCl display stoichiometric

dissolution. The ionic strength in these experiments of Zysset and Schindler (1996) were

above the ionic strength of the present study experiments, which were conducted at the same

pH without salt and below those with salt. Their dissolution rates (Fig. 6a) were slightly

slower than those of both data sets of the present study. The experiments conducted by

Zysset and Schindler (1996) at pH>2 and 0.03M KCl display non stoichiometric dissolution.

The present study (stoichiometric) rates are located between the rates that were calculated by

Zysset and Schindler (1996) from the release of Si and those from the release of Al (Fig. 6b).

The rates obtained by Zysset and Schindler (1996) in experiments conducted with 0.1

and 1 M KCl show a significant enhancement of rate in comparison to the experiments

conducted under relatively low ionic strength. Adding NaNO3 in the present study hardly had

an effect on the dissolution rate, and in some experiments, adding NaNO3 even inhibited the

dissolution. Figure 6b also shows that the effect of pH on dissolution rate is smaller in the

experiments conducted by Zysset and Schindler (1996) with 0.1 and 1 M KCl than the effect


observed in the present study. We do not have a good explanation for the differences

between the observation of the present study and those of Zysset and Schindler (1996). It is

interesting to note that the dissolution rate observed by Furrer et al. (1993) at ionic strength of

0.1M resembled the rate that was calculated based on Si by Zysset and Schindler (1996) at

0.03M.

Metz et al. (2004a) measured the dissolution rate of smectite sample SAz-1 at 50°C and

pH 2 and 3. The input solution in the far-from equilibrium experiments of Metz et al. (2004a)

was composed of pure HClO4 (no salt was added). Figure 6c shows that the dissolution rate

of smectite is not significantly influenced by the type of balancing anion (ClO4- vs. NO3

-).

6 SUMMARY AND CONCLUSIONS

Steady-state smectite dissolution rates were examined using non-mixed flow-through

reactors. The experiments were conducted at 25°C, 50°C and 70°C in a pH range of 1 to 4.5.

Most of the experiments were conducted under far-from equilibrium conditions (∆Gr< -21

kcal mol-1). Some experiments at pH>4, were conducted under conditions closer to

equilibrium. These experiments were not used in the fitting of the proposed model. In

general, smectite dissolution rate increases with temperature and decreases with pH.

The experimental results were fitted to a model, which is a simple variation of the

model proposed by Cama et al. (2002) for the proton-promoted reaction of kaolinite.

According to our model, the dissolution rate is proportional to the concentration of protons

that are adsorbed on the reactive edge surface site, which may be expressed using the

Langmuir adsorption isotherm. We do not have any information regarding the identity of this

active site. Zysset and Schindler (1996) postulated that the active site that controls the rate

under acidic conditions is either Al-O-Si or Al-OH-Al group. The results of the present study

neither support nor contradict this suggestion. The dependence of the dissolution rate on

temperature is affected by the activation energy and the adsorption enthalpy. From fitting our

results to the proposed model we found activation energy of 17±2 kcal/mol and enthalpy of

-11±2 kcal/mol.


Acknowledgments. This research was supported by THE ISRAEL SCIENCE

FOUNDATION (grant No. 174/01). We gratefully acknowledge thorough reviews by the

associate editor, Jacques Schott, and by Stephan .J. Köhler and Andreas Bauer. We thank

Volker Metz for both fruitful and knowledgeable discussions and for conducting the

thermodynamic calculations. The technical assistance of Ester Shani, Nivi Kesler, Ruth Talby

and Gony Yagoda is greatly acknowledged.


1 Table 1: Experimental conditions and results Experiment sample NaNO3 mass Flow rate SS time p H [Al] [Si] ∆Gr

(M) (g) (mL min-1) (days) input output (µM) (µM) (kcal mol-1)

25°C25-12.2* treated 0.2162 0.1341 0.010 512 1.00 1.08 48.9 122.2 2.0E-11 ±5E-12 2.0E-11 ±5E-12 -61.125-17.2* treated 0.2162 0.2532 0.085 58 1.12 1.12 24.3 59.3 4.4E-11 ±1E-11 4.3E-11 ±1E-1125-17.1* treated 0.2162 0.2801 0.025 22 1.12 1.12 74.6 193.2 3.7E-11 ±1E-11 3.8E-11 ±1E-11 -59.625-03.3 raw 0 0.1696 0.026 837 1.18 1.13 105.8 263.8 8.8E-11 ±3E-11 8.7E-11 ±3E-11 -52.925-07.4 treated 0 0.1523 0.039 531 1.00 1.14 52.1 139.1 7.3E-11 ±3E-11 7.7E-11 ±3E-11 -57.525-07.2* treated 0.2846 0.2329 0.038 412 1.57 1.59 31.5 76.4 2.8E-11 ±6E-12 2.7E-11 ±6E-12 -56.125-07.3* treated 0.2846 0.2111 0.039 435 1.71 27.3 64.1 2.7E-11 ±6E-12 2.6E-11 ±6E-1225-06.2 treated 0 0.0880 0.038 410 1.50 1.75 22.8 54.7 5.4E-11 ±2E-11 5.2E-11 ±2E-11 -52.525-02.1 raw 0 0.2337 0.012 60 1.95 1.94 62.8 143.9 1.7E-11 ±4E-12 1.6E-11 ±3E-12 -42.025-02.2 raw 0 0.1926 0.011 123 2.01 2.01 62.2 148.7 2.0E-11 ±5E-12 1.9E-11 ±5E-12 -39.825-11.1* treated 0.3062 0.1652 0.027 56 2.06 2.02 18.1 42.8 1.6E-11 ±3E-12 1.5E-11 ±3E-1225-11.2* treated 0.3062 0.1502 0.028 145 2.11 2.12 12.1 27.7 1.2E-11 ±3E-12 1.1E-11 ±2E-1225-11.3* treated 0.3062 0.1335 0.029 230 2.00 2.17 9.3 20.7 1.1E-11 ±3E-12 9.8E-12 ±2E-12 -55.725-12.1* treated 0.3131 0.1595 0.042 74 2.49 2.49 7.7 16.3 1.1E-11 ±2E-12 9.4E-12 ±2E-12 -52.525-15.1 treated 0 0.2438 0.013 250 2.49 2.54 26.3 65.8 7.7E-12 ±1E-12 7.6E-12 ±1E-12 -36.725-06.1 treated 0 0.1508 0.039 218 2.56 2.55 11.4 25.9 1.6E-11 ±4E-12 1.4E-11 ±3E-12 -41.625-16.1 treated 0 0.3136 0.009 281 2.51 2.56 63.0 147.3 9.7E-12 ±2E-12 9.1E-12 ±2E-12 -28.925-08.1 treated 0 1.7001 0.012 253 2.56 2.58 109.2 238.2 4.2E-12 ±8E-13 3.7E-12 ±7E-1325-13.1* treated 0.3152 0.1650 0.039 77 3.00 3.01 5.1 10.5 6.5E-12 ±1E-12 5.4E-12 ±1E-12 -47.725-01.2 raw 0 0.1164 0.040 100 3.05 3.01 4.7 11.2 8.9E-12 ±2E-12 8.4E-12 ±1E-12 -40.925-01.1 raw 0 0.1076 0.044 260 3.07 3.02 2.9 6.3 6.5E-12 ±3E-12 5.6E-12 ±1E-12 -44.925-07.1 treated 0 0.2583 0.038 225 3.53 3.57 3.5 7.5 2.7E-12 ±1E-12 2.4E-12 ±4E-13 -33.425-14.2* treated 0.3159 0.2438 0.032 278 3.57 3.58 3.3 6.1 2.3E-12 ±9E-13 1.7E-12 ±3E-13 -39.525-14.1* treated 0.3159 0.2452 0.032 242 3.56 3.71 3.4 6.3 2.4E-12 ±9E-13 1.8E-12 ±3E-13 -39.525-11.4* treated 0.316 0.1276 0.026 303 3.50 3.80 1.9 4.0 2.1E-12 ±2E-12 1.8E-12 ±5E-13 -41.425-11.5* treated 0.316 0.1272 0.030 335 3.88 3.90 1.3 2.8 1.6E-12 ±1E-12 1.5E-12 ±7E-1325-03.1 raw 0 0.2583 0.013 491 4.08 4.12 1.2 9.7 3.3E-13 ±3E-13 1.1E-12 ±2E-1325-03.2 raw 0 0.2572 0.014 656 4.10 4.20 1.5 5.0 4.3E-13 ±3E-13 5.8E-13 ±1E-13 -27.525-15.2* treated 0.3161 0.2297 0.010 388 4.05 4.30 6.4 13.6 1.5E-12 ±3E-13 1.3E-12 ±3E-13 -24.225-08.2 treated 0 1.6916 0.016 638 4.00 4.36 2.0 14.2 1.1E-13 ±4E-14 2.9E-13 ±5E-1425-16.2* treated 0.3162 0.2945 0.013 428 4.60 4.51 2.5 7.7 6.2E-13 ±3E-13 7.6E-13 ±2E-13 -26.0

(mol g-1s-1)RateAl RateSi

(mol g-1s-1)

* Experiment that were conducted with NaNO3; SS time = time from the beginning of the experiment to steady state.

The effect of pH and temperature on smectite dissolution -24- 11/10/04 2 Table 1: Experimental conditions and results (continue)

Experiment sample NaNO3 mass Flow rate SS time p H [Al] [Si] ∆Gr


50°C50-17.1* treated 0.2162 0.1898 0.029 21 1.14 1.11 843.1 1979.2 6.9E-10 ±3E-10 6.5E-10 ±3E-10 -40.450-07.2 treated 0 0.1323 0.034 227 1.55 1.55 191.4 479.7 2.7E-10 ±1E-10 2.7E-10 ±1E-10 -40.850-10.1* treated 0.2846 0.3468 0.135 15 1.60 1.60 159.7 359.0 3.4E-10 ±1E-10 3.0E-10 ±1E-1050-02b.1 treated 0 0.1471 0.111 17 2.00 1.98 33.7 83.4 1.4E-10 ±3E-11 1.4E-10 ±3E-11 -44.050-02b.2 treated 0 0.1330 0.111 26 2.01 2.00 30.7 71.1 1.4E-10 ±4E-11 1.3E-10 ±3E-11 -45.150-02b.3 treated 0 0.0865 0.112 53 2.00 2.00 24.6 57.1 1.7E-10 ±7E-11 1.6E-10 ±7E-11 -46.750-12.4* treated 0.3062 0.0595 0.015 343 2.00 2.02 55.2 116.8 7.6E-11 ±5E-11 6.4E-11 ±4E-11 -46.550-12.3* treated 0.3062 0.1052 0.008 194 2.00 2.02 146.8 300.6 6.3E-11 ±2E-11 5.2E-11 ±2E-1150-02.2 raw 0 0.0736 0.109 12 2.01 2.03 23.0 57.0 1.9E-10 ±4E-11 1.8E-10 ±4E-11 -46.750-02.1 raw 0 0.0818 0.119 5 2.03 2.04 24.0 61.2 1.9E-10 ±4E-11 1.9E-10 ±4E-11 -46.750-12.2* treated 0.3131 0.1277 0.038 81 2.47 2.48 23.6 52.0 3.8E-11 ±1E-11 3.3E-11 ±9E-12 -44.150-12.1* treated 0.3131 0.1407 0.044 40 2.46 2.48 25.4 59.3 4.3E-11 ±1E-11 4.0E-11 ±9E-12 -43.350-06.1 treated 0 0.1424 0.035 66 2.47 2.49 29.8 68.1 4.0E-11 ±9E-12 3.6E-11 ±8E-12 -35.550-01.1 raw 0 0.1102 0.094 32 2.96 2.97 8.6 21.2 4.0E-11 ±8E-12 3.9E-11 ±7E-12 -35.550-01.2 raw 0 0.1074 0.088 45 2.96 3.00 8.7 21.1 3.8E-11 ±8E-12 3.7E-11 ±7E-12 -35.550-13.2* treated 0.3152 0.1272 0.026 278 3.02 3.05 9.6 21.1 1.1E-11 ±3E-12 9.4E-12 ±3E-1250-13.1* treated 0.3152 0.1304 0.025 237 3.04 3.06 15.3 32.6 1.6E-11 ±4E-12 1.4E-11 ±4E-1250-07.1 treated 0 0.2501 0.033 133 3.55 3.58 6.8 16.0 4.9E-12 ±9E-13 4.6E-12 ±8E-13 -26.850-14.1* treated 0.3159 0.1451 0.083 230 3.55 3.62 2.1 3.5 6.4E-12 ±3E-12 4.4E-12 ±2E-12 -45.550-03.3 raw 0 0.1725 0.039 428 4.07 4.07 0.6 5.3 7.8E-13 ±6E-13 2.6E-12 ±5E-13 -29.650-03.2 raw 0 0.1733 0.041 386 4.10 4.12 0.8 5.5 1.1E-12 ±8E-13 2.8E-12 ±5E-13 -29.650-03.1 raw 0 0.1738 0.013 273 4.02 4.15 1.5 17.4 5.9E-13 ±5E-13 2.8E-12 ±5E-13 -20.150-15.2* treated 0.3161 1.7362 0.015 388 4.20 5.9 28.9 2.8E-13 ±5E-14 5.5E-13 ±1E-1350-15.1* treated 0.3161 1.7379 0.015 334 4.72 4.21 5.1 31.6 2.4E-13 ±4E-14 6.0E-13 ±1E-13 -20.450-16.2* treated 0.3162 1.7302 0.015 414 4.74 4.59 2.1 21.4 1.0E-13 ±4E-14 4.1E-13 ±7E-1450-16.1* treated 0.3162 1.7308 0.015 313 4.54 4.65 2.2 25.5 1.1E-13 ±4E-14 4.9E-13 ±8E-14 -17.050-08.2 treated 0 1.7391 0.037 461 4.17 4.80 0.1 43.8 6.2E-15 ±1E-14 2.0E-12 ±3E-13 -12.450-08.1 treated 0 1.7390 0.037 505 4.00 5.46 0.1 35.5 1.5E-14 ±3E-14 1.6E-12 ±3E-13 -10.5

RateAl RateSi

(mol g-1s-1) (mol g-1s-1)


The effect of pH and temperature on smectite dissolution -25- 11/10/04 3 Table 1: Experimental conditions and results (continue)

Experiment sample NaNO3 mass Flow rate SS time p H [Al] [Si] ∆Gr


70°C70-04.1 treated 0 0.2631 0.062 25 1.17 1.14 413.8 1058.9 5.3E-10 ±2E-10 5.4E-10 ±2E-10 -43.470-17.1* treated 0.2162 0.2042 0.030 24 1.14 1.16 816.7 2387.5 6.6E-10 ±3E-10 7.7E-10 ±4E-10 -37.570-10.1* treated 0.2846 0.2925 0.107 10 1.60 1.62 166.0 412.5 3.3E-10 ±1E-10 3.3E-10 ±1E-1070-02.2 raw 0 0.0434 0.122 27 1.95 1.94 27.8 64.1 4.2E-10 ±2E-10 3.9E-10 ±2E-1070-02b.1 treated 0 0.1038 0.116 21 2.01 2.00 54.5 129.1 3.3E-10 ±1E-10 3.1E-10 ±1E-10 -40.170-02b.3 treated 0 0.0486 0.118 63 1.99 2.00 25.6 59.6 3.4E-10 ±3E-10 3.1E-10 ±3E-1070-02b.2 treated 0 0.0342 0.118 81 2.01 2.01 17.8 40.8 3.3E-10 ±4E-10 3.1E-10 ±4E-10 -49.770-02.1 raw 0 0.0586 0.121 18 2.06 2.02 30.4 72.0 3.4E-10 ±1E-10 3.2E-10 ±1E-1070-15.2* treated 0.3062 0.0731 0.018 557 2.24 2.29 69.6 176.6 9.1E-11 ±5E-11 9.2E-11 ±5E-1170-06.1 treated 0 0.1071 0.042 69 2.51 2.51 43.2 100.1 9.2E-11 ±3E-11 8.6E-11 ±3E-11 -31.870-06.2 treated 0 0.1520 0.042 19 2.49 2.52 71.2 163.6 1.1E-10 ±2E-11 9.8E-11 ±2E-11 -27.870-08b.3 treated 0 0.1447 0.040 49 2.54 2.55 54.3 132.9 8.1E-11 ±2E-11 7.9E-11 ±2E-1170-12.1* treated 0.3131 0.1025 0.027 48 2.50 2.55 43.9 101.6 6.3E-11 ±2E-11 5.8E-11 ±1E-11 -39.070-01.1 raw 0 0.0806 0.081 27 3.00 3.02 9.4 23.5 5.2E-11 ±1E-11 5.1E-11 ±1E-11 -34.270-13.1* treated 0.3152 0.1455 0.026 74 3.03 3.10 33.0 71.2 3.2E-11 ±8E-12 2.8E-11 ±7E-12 -30.470-07.1 treated 0 0.2625 0.034 123 3.58 3.57 7.3 30.4 5.1E-12 ±9E-13 8.5E-12 ±1E-12 -22.470-14.2* treated 0.3159 0.3261 0.083 80 3.60 3.59 7.7 16.7 1.1E-11 ±2E-12 9.2E-12 ±2E-12 -33.170-14.1* treated 0.3159 0.3454 0.084 27 3.55 3.61 10.9 26.9 1.4E-11 ±3E-12 1.4E-11 ±3E-1270-07.2 treated 0 0.2489 0.033 294 3.95 3.70 6.7 18.8 4.9E-12 ±9E-13 5.4E-12 ±1E-12 -23.870-15.1* treated 0.3161 0.1710 0.008 266 4.09 4.21 4.9 13.6 1.2E-12 ±2E-13 1.3E-12 ±2E-13 -24.970-16.3* treated 0.3162 0.2607 0.014 210 4.09 4.25 5.3 16.3 1.5E-12 ±3E-13 1.8E-12 ±3E-1370-03b.1 treated 0 0.2175 0.013 314 4.11 4.27 0.6 24.8 1.9E-13 ±1E-13 3.2E-12 ±5E-13 -19.770-08b.2 treated 0 0.1007 0.029 325 4.80 4.37 0.9 4.4 1.4E-12 ±1E-12 2.8E-12 ±1E-1270-16.2* treated 0.3162 0.2598 0.014 257 4.65 4.52 1.5 8.6 4.3E-13 ±3E-13 9.7E-13 ±2E-13 -26.370-08b.1 treated 0 0.1009 0.013 273 4.52 1.1 7.1 7.3E-13 ±7E-13 2.0E-12 ±7E-13 -25.470-16.1* treated 0.3162 0.2617 0.014 113 4.55 4.59 1.4 13.3 4.1E-13 ±3E-13 1.5E-12 ±3E-13 -22.770-03.1 raw 0 0.1736 0.010 156 4.05 4.61 0.6 29.5 1.9E-13 ±2E-13 3.8E-12 ±7E-13 -17.1

RateAl RateSi

(mol g-1s-1) (mol g-1s-1)



7 FIGURE CAPTIONS

1 Fig. 1: Experimental set-up: a) Schematic illustration of the flow-through system;

b) A detailed view of the reaction cell.

2 Fig. 2: Variation of concentrations of Si and Al in three representative

experiments. The vertical lines delineate different experimental stages.

Analyses used to calculate average compositions at steady state are denoted

by open symbols. The inserts in (c) show enlargements of parts of the

experiment.

3 Fig. 3: Comparison of dissolution rates evaluated based on the release of Al,

RateAl, with those obtained based on the release of Si, RateSi at (a) 25°C, (b)

50°C and (c) 70°C. The solid lines are the 1/1 diagonals. The inserts are

enlargements of the sections that show experiments with slow dissolution

rates.

4 Fig. 4: Variation of log dissolution rate with pH: (a) at 25 ºC, (b) at 50 ºC and (c) at

70 ºC. The figure compares experiments conducted under constant ionic

strength (I) and NO3- concentration of ~ 0.32 M with those obtained without

adding NaNO3, i.e., where H+=NO3- (see section 5.3). The dashed curves

are a result of multiple non-linear regression of equation (10) to all the

measured rate data at pH<4


5 Fig. 5 Comparing the prediction of the proposed model (solid line) to surface

protonation data of Zysset and Schindler (1996) (symbols). The white and

the black squares represent data points at pH<2.5, in which according to the

interpretation of Zysset and Schindler (1996) surface protonation includes

a significantly contribution from ion exchange. The solid line is the best-fit

curve of Eq. (13) to adsorption isotherm data between pH 2.5 and 4 (black

dots). The model proposed in the present study predicts the concentration

of adsorbed H+, between pH 1.5 and 4 (black dots and squares), although

only the data at pH≤2.5 (black dots) were used for the fitting (see section

5.4).

6 Fig. 6: Comparison of the present study smectite dissolution rate to those obtained

by (a) Furrer et al. (1993), (b) Zysset and Schindler (1996) and (c) Metz et

al. (2004a). The dashed curve in (b) represents the model fitting of all the

present study measured rate data at pH<4.

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The combined effect of pH and temperature on smectite dissolution rate under acidic conditions

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