“1.0 : : - :
ry

= 15 * 2 ote J

Z fF 5

E 20/4 bd °

= bd id 1

ra) e @

iH

Sash oe © OCuill) 4
e@pp

Se
-3.0 \ : : L
-7.0 -6.0 5.0 4.0 -3.0 -2.0
tog C(w) [M]

Figure 11. Adsorption isotherms of Cu(Il) (in the absence of dodecyipyridinium) and dodecylpyridinium (in the absence of Cu(Il)) on Lula aquifer
material (0.005 kg/L in 0.01 M NaCl). In the Cu(II) experiment, solution pH varied from about pH 5.6 at low concentrations of Cu(||) to about

pH 4.9 at high concentrations of Cu(!!).

3.5 T T T

sok S @ © e ee eT
2 88
a 2%
a
> 25 7 % 7
= © Cuil)

@oP
2.0 n n 1
75 65 55 4.5
log H [M]

Figure 12. The pH dependence of the distribution ratio of Cu({I) (in the absence of dodecylpyridinium) and dodecyipyridinium (in the absence of Cu(|!)),
for adsorption onto Lula aquifer material (0.005 kg/L in 0.01 M NaCl). Total Cu(!l) concentration was 40 pM in all experiments; total DP
concentration was 40 pM in all experiments.



0 r 1 ; —— SOS
© Montmorillonite °

<1 | @ EPAN2 °
3 @ Lula Ne
= A Kaolinite
2 2 ‘& Borden Sand
io
Oo 3
2
2

4

a wt
5 1. af L . n 1
9 8 7 4 5 4 3 2
log C(w) [M]

Figure 3.

kg/L; kaolinite Cq(w) =
kg/L. Solutions contained 0.01 M Nat.

of pH-dependent surface functional groups, or (ii) ion-exchange
of H+ for pyridinium. The data (Figure 4) indicate that both of these
effects are weak.

Effect of Major Cations on Adsorption of DP

The effect of concentration and type of electrolyte on adsorption
of DP was examined with two complementary types of data: (i)
adsorption isotherms of DP from solutions with various
concentrations of salts; and (ii) concentration distribution ratio of
DP as a function of concentration of NaCl or CaCl, with a
constant amount of DP in the system. In view of the nonlinearity
of the isotherms, the first method is preferable for development of
mechanistic models, while the second provides a quick estimate
of the magnitude of the effect.

A number of experiments were performed with the sorbents and
different inorganic salts at various concentrations. Here we
present data for one representative system. Figure 5 shows the
isotherms of DP on Lula aquifer material in the presence of 0.01
M NaCl, 0.1 M NaCl, and 0.001 M CaCl. The effect of the salt
concentration was generally much greater than that of pH.

A comparison of the isotherms in NaCl shows that, at low
concentrations of DP, an increase in the NaCl concentration
decreases the adsorption by a nearly constant factor. This
behavior is consistent with an ion-exchange mechanism. At high
concentrations of DP, the isotherms merge and cross at surface
concentrations that corresponds roughly tothe CEC. This behavior
is consistent with the predominance of adsorption of pyridinium-
chloride at high concentrations of DP. Similar isotherms (not
shown) were obtained for EPA-12, Borden sand, and kaolinite.

The effect of Ca2+ on the adsorption of DP on Lula aquifer material
is generally stronger than that of Na*, consistent with ion exchange.
A solution of 0.001 M Ca2+ impedes adsorption of DP more
effectively than 0.01 M Na+, but less than 0.1 M Nat.

Logarithmic adsorption isotherms of dodecylpyridinium on different materials. Montmorillonite, ratio of solids to liquids [C,(w)] was 0.0050
.025; soil EPA-12, C,(w) = 0.025; Lula aquifer material, C,(w) = 0.0025; and Borden sand, C,(w) = 0.025 or 0.050

Nonlinearity of isotherms

All of the adsorption isotherms determined in this study were
nonlinear, even at very low surface coverages of dodecylpyridinium.
While the Freundlich isotherm can be used to represent these data
empirically, it provides no mechanistic basis for evaluating
competition effects, saturation effects, and comparison of adsorption
energies for a homologous series of cationic surfactants. Thus the
usefulness of the Freundlich isotherm is limited.

In order to derive a more useful description of adsorption data, it
is necessary to consider the causes of nonlinearity: (i) non-
uniformity of energies of adsorption sites, with saturation of high-
energy sites; (ii) sorbate-sorbate interactions such as repulsive
electrostatic interactions or cooperative chain-chain interactions,
and (iii) experimental artifacts that stem from covariation of
experimental conditions, such as changes of pH, major ion
concentrations, or aggregation of particles, that accompany
changes in cationic surfactant concentrations.

We interpret the nonlinearity of isotherms in these experiments
primarily as the result of heterogeneity of sites. Particularly in the
low-concentration regions of the isotherms, the changes in pH,
concentrations of major ions in solution, and surface charge that
accompanied adsorption of the cationic surfactant were too small
toaccountfor theisotherm nonlinearity. Thus, atlow concentrations
of cationic surfactant, heterogeneity of sites is the preferred
explanation. Athigher concentrations it is more difficult to exclude
the influence of other processes on isotherm nonlinearity.

Model for DP Adsorption

The mechanistic models that were considered for these data were
(i) multisite models, (ii) electrostatic models, and (iii) some
combination of the two. Most of the phenomena observed in this
study can be explained readily through the multisite model, with


Westall, J. C. “FITEQL: A Computer Program for Determination
of Chemical Equilibrium Constants from Equilibrium Data Version
1.2”, Report 82-01; Oregon State University: Corvallis, OR,
1982a.

Westall, J.C. “FITEQL: A Computer Program for Determination
of Chemical Equilibrium Constants from Equilibrium Data Version
2.0", Report 82-02; Oregon State University: Corvallis, OR,
1982b.

Westall, J.C.; Chen, H. The Effect of Cationic Surfactants on the
Adsorption of Transition Metal lons on Aquifer Materials. To be
submitted for publication in Soil Sci, Soc, J, Am, 1994.

18



TABLE 4. DISTRIBUTION CONSTANTS, K, OF CHLOROBENZENES BETWEEN 0.01 M CaCi, AND LULA AQUIFER MATERIAL.
AMOUNTS OF DP ON THE AQUIFER MATERIAL ARE INDICATED.

Untreated Sorbent DP-treated Sorbent
batch K® column Kb batch K® column K>
(Lkg) (Lkgl) (Lkg') (Lkg")
ocB <0.1405 0.2 0.623 Vw
TCB 023203 0.7 321 3.5
TeCB 241 23 8+3 79
PeCB 743 9.0 2528 23

@ Determined from data in Figure 7.
> Determined from data in Figure 8.
© {DP} = 4.3 mmol kg"

1.2
1.0
0.8
0.6 ; A
0.4
0.2 DCB

0.0

C/iCo

0 4 8 12 16 20 0 10 20 30 40 50
Pore Volume Pore Volume

1.2 *
1.0

0.8
0.6
0.4
0.2
0.0

C/Co

PeCB

0 20 40 60 80 100 0 50 100 150 200 250
Pore Volume Pore Volume

Figure 8. The breakthrough curves of the chlorobenzenes on (0) untreated Lula aquifer material and (A) Lula aquifer material equilibrated with
10 uM DP. The solid lines were calculated from a finite difference approximation to the advection dispersion equation, with the parameters
in Table 3 and the partition constants of the chlorobenzenes defined by the data in Figure 7.

10


where [DP] represents the concentration of species DP in solution
(mol L:'), and{DPX} represents the concentration of species DP onthe
sorbent (mol kg"). The material balance equation for site iis:

Ti = (X} + {DPX} (5)

where T; is the total amount of site i per mass of sorbent (mol
kg"). Equations 4 and 5 exactly define a Langmuir isotherm.
Then the total amount of DP adsorbed is taken as the sum over
all types of sites i,

Cpp(s) = £ {DPX} ©)

According to the log K spectrum approach, the values of log Kj
were set to bracket the range of the data (i.e., log K; = 2, 3, 4, 5,
6, 7 to span the range of - log Cpp(w)), and values of T, were
determined by FITEQL (Westall, 1982a, 1982b). The log K
spectrum determined in this way for the DP-Lula-0.01 M CaCl
system is given in Table 2, and the isotherm calculated from this
model is given as the solid line in Figure 6. As can be seen, the
model yields an excellent representation of the experimental
data. The advantage of this approach is that it offers aconvenient
parameterization of nonlinear isotherms that is consistent with
general chemical equilibrium models.

Certainly this “discrete equilibrium-constant-spectrum” approach
is subject to the criticism that a model can be fit to any set of data
if enough adjustable parameters are used. Furthermore, there is
some covariance among the values of the adjustable parameters;

TABLE 2. MULTI-LANGMUIR MODEL FOR DP ADSORPTION?

T; (mol/g)

0
4.21 x 102
1.91 x 102
3.05 x 10°3
8.77 x 104

0

Nonaon

@ Model defined by Equations 4-6 in text.

the values are not unique, but just one of an infinite number of sets
that can be used to represent the data.

However, the advantages of this approach are significant, too.
The practical advantage is that the effects of multiple interactions
in complex systems can be represented in a way thatis completely
compatible with chemical equilibrium models used in environmental
chemistry. An adsorption isotherm that conforms closely to the
Freundlich equation over more than than six orders of magnitude
range in C(w) can be described very well by conventional mass
action and material balance equations; these equations have a

Figure 6.

0 T + T T
-~7f 7
<
Kj
o2F :
—

t
o 3 F 7
a
s
4 F 4
5 n 1 1 1
7 6 5 4 3 2
log C(w) [M]

Adsorption isotherm of dodecylpyridinium on Lula aquifer material (0.005 kg/L) in 0.01 M CaCl,. (A) experiments conducted with '4C-
labeled DP and direct determination of DP in solution and on the sorbent; (0) experiments conducted with unlabeled DP and determination
of DP in solution by UV spectrophotometry and DP on the sorbent by difference. The solid line is the isotherm calculated from the multi-
Langmuir model (Table 2), and the broken line is the empirically adjusted isotherm used in the transport model in Figure 9. The two symbols
(@) on the isotherm designate the DP concentration at which the Phase Il experiment (10 uM, Figure 8) was carried out and the initial DP
concentration for the Phase III experiment (3 mM, Figures 9 and 10).


4 — T r T
e ee e e e
3k Be" 8 ws . 4
&
D>
3 Aa Aa
2b a 4
@ EPA-12 an
TM LulaNé
4 Kaolinite
1 A
-11 7 3
log [H*]

Figure 4. _ Effectoflog[H*] on the concentration distribution ratio, D,, of dodecyipyridinium between aqueous solutions and selected sorbents. Sorbent.
concentrations (C,(w)) were as follows: EPA-12, C,(w) = 0.015 kg/L; Lula, C,(w) = 0.0025; and kaolinite, C,(w) = 0.0125. Solutions
contained 0.01 M Nat. The amounts of dodecylpyridinium adsorbed (Cpp(s)) were nearly constant for EPA-12 and Lula sorbents: Cpp(s)
= 0.20 and 29 ymol/g, respectively. Cpp(s) ranged from 7.6 to 15 mol/g for kaolinite.

0 T T + —— . .
0 0.01 M NaCl

-1 £ @ 0.1 MNaci 4
3 © 0.001 M CaCl
J
3 2+
&
s
o 3+
a
2

4t

$6

9 8 7 6 5 ~4 -3 -2
log C(w) [M)

Figure 5. Logarithmic adsorption isotherms of dodecyipyridinium on Lula aquifer material from 0.1 M NaCl, 0.01 M NaCl, and 0.001 M CaCl. The
ratio of solids to liquid was 0.0025 kg/L.

electroneutral reactions and no explicit electrostatic energy term. X; according to the reaction

While this approach is obviously an approximation, it is consistent

with a variety of observations, and it is conceptually very simple.

Here the multisite approach is illustrated for the adsorption of DP. DP + Xj = DPX, K (3)
(In this example, effects of Nat, Ca2+, and Cl are omitted for
simplicity; the full treatment is described by Brownawell et al.,
1990). The nonlinear isotherm was modeled as a sum of
Langmuir isotherms, with a discrete, regular distribution of log K’s
(Brownawell et al., 1990). DP is presumed to adsorb to each site [DP] {X} K; = {DPX} (4)

for which the mass action equation is


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Figure 7. The partition constant of chlorobenzenes between 0.01 M CaCl, and Lula aquifer material (0.025 kg/L solid/solution ratio) with different
loadings of DP.

TABLE 3. PHYSICAL PROPERTIES OF COLUMN PACKED WITH LULA AQUIFER MATERIAL

Symbol Description Experimental ® Model> Units

D ined Di

L column length 5.000 + 0.001 5.00 cm

dj internal diameter 1,000 + 0.001 1.00 cm

Ls sorbent mass. 6.84 + 0.01 6.84 g

Vin holdup volume 1.73 + 0.14 1.71 mt

Ve dead volume 0.14 + 0.01 0.14 mL

Fy flow rate 3.99 + 0.03 3.99 mint

D dispersion coefficient 0.0003 + 0.0001 0.0002¢ om? 51

Derived

Vi interstitial volume 1.59 + 0.14 1.57 mL
= Vin~ Vo

x empty column volume 3.927 + 0.008 3.927 me
= 3.1474) /2)2°L

Vs volume of sorbent 2.34 40.14 2.36 mie
=X-Vi

& Porosity 0.40 + 0.04 0.40 _
=V,IX

t residence time 1435 + 120 1417 s
= V;/ Fc*3600 s ht

u flow velocity 0.0035 + 0.0003 0.003 ems!
elit

08 bulk density 1.742 + 0.003 1.742 gmt
= Mg/X

Ps sorbent density 2.92 + 0.17 2.90 gmt?
=Mg/ Ve

a Estimated uncertainty in experimentally determined value is given.

> Values used in the models.

© As described in the text, several different values for D were used, D = 0.0002 for model of HOC
breakthrough on untreated aquifer material and for model of HOC breakthrough on aquifer material
equilibrated with 10 uM DP (Figure 8). D = 0.02 for DP transport (Figures 9 and 10). D = 0.002 for
HOC transport through column with time-varying DP (Figure 10).


United States
Environmental Protection
Agency

Robert S. Kerr Environmental
Research Laboratory
Ada, OK 74820

Research and Development

EPA/600/S-94/002 August 1994

EPA

ENVIRONMENTAL

RESEARCH BRIEF

The Use of Cationic Surfactants to Modify Aquifer Materials to Reduce the
Mobility of Hydrophobic Organic Compounds

John C. Westail@, Bruce J. Brownawell>
Julia Wagner? and Hua Chen

ABSTRACT

Cationic surfactants can be used to modify surfaces of soils and
subsurface materials to promote sorption of hydrophobic organic
compounds (HOC) and retard their migration. For example,
cationic surfactants could be injected into an aquifer downgradient
from a source of HOC contamination to provide a temporary
barrier against migration. A possible side effect of such a
treatment could be the mobilization of previously adsorbed metal
ions through cation exchange with the cationic surfactant.

Batch and column experiments were performed to investigate
these phenomena. The cationic surfactant was dodecylpyridinium
(DP); the HOC were chlorobenzene homologs; competing metal
ions were Pb2+, Cd2+, and Cu2+; and sorbents were low-organic-
carbon aquifer materials (Lula, EPA-12, Borden Sand) andpristine
minerals (kaolinite, montmorillonite). The investigation covered
three major topics: (i) adsorption of DP to particle surfaces; (ii)
sorption of chlorobenzenes to DP-modified surfaces; and (iii)
adsorption competition between DP and the metal ions.

The adsorption isotherms of DP were distinctly nonlinear, even at
very low surface concentrations. The distribution of DP was
strongly dependent on the solution concentrations of Na* and
Ca?+, but virtually independent of solution pH. Amuttisite adsorption
model was developed to describe adsorption over a wide range

8 Department of Chemistry, Oregon State University,
Corvallis, OR 97331-4003

> Waste Management Institute, Marine Sciences Research Center,
SUNY, Stony Brook, NY 11794-5000

> >

of DP, NaCl, and CaCl, concentrations. Two types of adsorption
reactions were found to be significant: exchange of DP with a
alkali-metal cation, and adsorption of pyridinium with chloride
counter-ion.

Distribution ratios of the chlorobenzenes varied nonlinearly with
DP loading of the surface. The elution of chlorobenzenes from
columns packed with DP-treated aquifer material showed
significant retardation with only moderate amounts of DP on
particle surfaces. A transport model based on results of the batch
experiments and the local equilibrium assumption yields an
acceptable approximation for the coelution of DP and HOC from
thecolumn. Itis concluded that treatment of surfaces with cationic
surfactants shows promise as ameans of promoting HOC sorption
in a variety of treatment processes.

Acomparison of the adsorption isotherms of DP, Pb2+, Cd2+, Cu2+
shows that DP is adsorbed more strongly than the metal ions from
solutions at pH 5.5 - 6.0. Competition experiments suggest that
DP and the metal ions adsorb to different types of sites, and that
competition becomes significant only when about half the CEC of
the material is occupied by DP. These results suggest that
adsorbed metal ions could not effectively be removed by washing
aquifer material with DP, but that aquifer materials could be
treated with DP to retard HOC without danger of mobilizing
adsorbed metal ions.

INTRODUCTION

Cationic surfactants that enter the subsurface environment will
modify the surfaces of soils and subsurface materials to promote
sorption of hydrophobic organic “Bm (HOC), and reduce

® Printed ‘on Recycled Paper



1.2 T T T
1.0
0.8
0.6

C/Cy

0.4
0.2

L 1 1

0.0
75

100
Pore Volume

125 150 175

1.2 T

1.0

0.8

0.6

CIC

0.4

0.2

0.0 +

200 400

600

800

Pore Volume

Figure 9.

The breakthrough curve and desorption curve of 3 mM DP in 0.01 M CaCl. The solid line was calculated from a finite difference

approximation to the advection dispersion model with the constants in Table 3 and the empirically adjusted adsorption isotherm depicted

in Figure 6.

adjustable parameters were not introduced. Thus, the model
used in Figure 10 was developed from the following data: (i) the
amount of DP on the column as a function of time and distance
was calculated from the advection-dispersion model with local
equilibrium according to the batch isotherm in Figure 6 and the
apparent dispersion coefficient determined from data in Figure 9;
(ii) the partition coefficients of the chlorobenzenes were calculated
for the amount of DP on the Lula and the third order polynomial
derived from the datain Figure 7; (iii) sorption ofthe chlorobenzenes
onthe sorbent were assumed to be linear; (iv) the hold-up volume
was determined from NO, breakthrough; and (v) the apparent
dispersion coefficient for the chlorobenzenes was set to 0.002
cm? s” to fit the model to the data. The actual breakthrough
curves were then calculated from the advection-dispersion model
based on these parameters.

An interesting feature of Figure 10 is noted: the ratio of effluent
concentration to influent concentration (C/Cy) for DCB is

12

appreciably greater than one. The model reproduces this feature
as well. This phenomenon is attributed to a “preconcentrate and
strip” mechanism: OCB is initially preconcentrated in a narrow
zone on the column, and then it is stripped off the column as the
OP itself elutes from the column.

We conclude that the simple local equilibrium model is successful
in describing retention volumes in Figures 9 and 10, but that
considerable liberty must be taken with the apparent dispersion
coefficients to achieve agreement of the model with the data.
However, the uncertainty about the treatment of dispersion does
not invalidate the general conclusions presented here, but points.
out an area for further experimentation.

COMPETITION BETWEEN DP AND METAL IONS

Adsorption of cationic surfactants can displace adsorbed metal
ions (Beveridge and Pickering, 1983; Bouchard et al., 1988). In


from this model (i.e., the solution to Equation 7 with parameters
in Tables 3 and 4).

This experiment has important implications for both applications
and theory. First note that the amount of DP adsorbed was only
4.3 mmol kg"; this amount is equivalent to a DP-fg¢ of about
0.09% compared to the natural fog of 0.02%. As seen in Figure
7, much higher loadings are possible. Even at this low amount of
DP, retardation factors for the HOC are approximately three times
the retardation factor for HOC on untreated sorbent, as shown in
Table 5.

TABLE 5. RETARDATION FACTORS, R®, FOR HOC ON
LULA AQUIFER MATERIAL AS CALCULATED
FROM COLUMN K GIVEN IN TABLE 2.

Untreated OP-treated

Sorbent Sorbent>
DCB 2 6
TCB 4 16
TeCB an) 35.
PeCB 40 101

@ R=1+K pg ((1-€) / &) where p, and ¢ are as.
.. defined in Table 3.
'b {DP} = 4.3 mmol kg?

Second, the model, which was based on the local equilibrium
assumption and values of K’s not significantly different from those
determined in batch experiments, agrees well with the experimental
data. This agreement indicates that a “particle concentration
effect” is negligible, both for the DP and the chlorobenzenes,
since the solid/liquid ratios in the batch experiments were 0.005
kg/L and 0.025 kg/L, and in the column experiment it was 4.35 kg/
L, that is, over a factor of 1000 difference. We attribute this good
agreement to the fact that we washed the materials used in the
batch experiments thoroughly to remove materials that cannot be
separated by centrifugation. Also the fact that the local equilibrium
assumption seems to hold indicates that sorption of the
chlorobenzenes is rapid compared to transport through the column.

Breakthrough of DP and Chlorobenzenes

Since the final phase of this part of the study deals with the
sorption of chlorobenzenes onto aquifer material with a time-
varying amount of DP, we discuss first our understanding of the
time variation of DP on the column.

The experimental data for the breakthrough of 3mM DP and the
elution of DP from a column containing Lula initially equilibrated
with 3 mM OP are presented in Figure 9. The retardation factor
calculated from integration of the breakthrough curve is 80
compared to 86 calculated from the batch isotherm. We feel that
this agreement within 10% is good, considering the uncertainty
introduced from factors such as the high concentrations of DP, the
effect of DP on the surface charge andthe packing of the particles,
and the uncertainty in the interstitial volume.

Whereas the retardation factor agrees with the simple batch-
based model, the shape of the breakthrough curve is impossible
to explain in terms of a simple focal-equilibrium model with the
dispersion coefficient from the NO; breakthrough curve. The
low-dispersion local equilibrium model would predict a sharp,
almoststep-like, breakthrough curve (due to the convex isotherm)
rather than the relatively broad wave that is observed. Moreover,
an understanding of the mechanism of broadening is important
since the distribution of DP along the length of the column
depends on the mechanism, and the transport of the
chlorobenzenes must depend to some degree on how the DP is.
distributed, since the curves in Figure 7 are not linear.

Among the many possibilities for explaining the breadth of the
wave is a slow adsorption-desorption step, which results in
greater dispersion. For solutes which follow a linear adsorption
isotherm, it can be shown how a slow adsorption-desorption step
leads to an increase in dispersion, which is directly proportional to
the flow velocity (Baetslé, 1969; Giddings, 1991; Bales and
Szecsody, 1985). Additional data, at different flow velocities,
would allow us to substantiate this mechanism andits applicability
to the nonlinear isotherms in this study. Atthis point, our objective
isto have aplausible explanation for the shape of the breakthrough
curve. This objective can be fulfilled by adjusting the apparent
dispersion coefficient to 0.02 cm? s', on the basis of slow
adsorption-desorption kinetics, and modeling adsorption by the
local equilibrium assumption according to the isotherm in Figure
6. itis worth noting that other more explicit kinetic models for
explaining the breadth of the wave were tested, but were not as
successful as the simple model presented here.

Although the desorption curve (Figure 9B) appears to approach
C/Cp = 0, integration of the elution curve reveals that, even after
approximately 1000 pore volumes, about 16% of the original DP
loading is still on the column. The nonlinearity of the isotherm
ensures that as the concentration of DP decreases, it adheres
more tenaciously to the column. The amount of DP remaining on
the column is almost twice the amount of DP on the column during
the experiment involving the breakthrough of chlorobenzenes on
column containing aquifer material with a uniform, time-invariant
amount of DP on the surface as shown in Figure 8. Thus after
1000 pore volumes there is still enough DP onthe column to retard
the chlorobenzenes. This extensive tailing is that which is
expected from the highly convex isotherm.

The DP adsorption isotherm in Figure 6, adjusted as shown for
concentrations above 170 uM, was incorporated into the transport
model previously described. The apparent dispersion coefficient
for DP in the model was set to 0.02 cm? s1. Other model
parameters are given in Table 3. The model so defined agrees
well with the experimental data, as seen in Figure 9.

Sorption-Desorption of Chlorobenzenes

After the column was equilibrated with 3 mM DP, a pulse of
chlorobenzenes in 0.01 MCaCl,, without DP, was passed through
the column, followed by 0.01 M CaClp. The elution histories are
shown in Figure 10, along with the elution histories from the
model, for which the parameters are given in Table 3. The model
for the chlorobenzene breakthrough curves in Figure 10 is based
on parameters that were determined in independent experiments
(Figures 6-9). The only adjustable parameter determined form
the data in Figure 10 was the dispersion coefficient of the
chlorobenzenes in the column with varying DP. Additional


mechanistic basis, which potentially could be used to understand
the factors that influence adsorption.

Another advantage of such a model is that it provides a framework
for evaluating quantitatively mechanisms that could be responsible
for the phenomena that are observed experimentally, such as (i)
reversal of effect of electrolyte concentration on isotherms near
the CEC; (ii) lateral shift of isotherms with salt concentration far
below the CEC; (iii) absence of pH effect; and (iv) relative effect
of Na vs. Ca on the isotherms.

SORPTION OF CHLOROBENZENES TO
DP-MODIFIED SURFACES

This work was carried out in three phases. First, in batch
experiments, we determined the adsorption isotherm of DP on
Lula aquifer material and investigated sorption of chlorobenzenes
to the DP-treated aquifer material.

Nextwe examined the breakthrough of chlorobenzenes on columns
containing the aquifer material with a uniform, time-invariant
amount of DP on the surface. These conditions are not those that
one would expect to encounter in field application, but they are
useful in evaluating the applicability of data generated in batch
experiments to column experiments and the validity of the local
equilibrium assumption in these systems. To maintain the constant
amount of DP on the sorbent, the solution flowing into the column
always contained a constant low concentration (10 1M) of DP.

Finally we investigated the breakthrough of chlorobenzenes on
columns that were initially uniformly coated with DP, but, when the
chlorobenzene pulse was applied to the column, the supply of DP
to the column was stopped, and the DP was allowed to desorb
slowly as the chlorobenzenes passed through the column. This
design is closer to that which would actually used in the field —a
pulse of the cationic surfactant would be applied to treat the
surfaces of aquifer particles, but the cationic surfactant would not
be injected into the aquifer over extended periods.

Adsorption Isotherm of DP

The adsorption isotherm of DP on Lula aquifer material in 0.01 M
CaCl, is shown in Figure 6 on a log-log scale. The isotherm is
nonlinear. The slope of the log-log plot is approximately 0.6 in the
low concentration range; on a linear scale (not shown), the
isotherm has an almost step-function appearance. In the high
concentration range, near the critical micelle concentration, the
isotherm levels off at approximately the cation exchange capacity
of the sorbent. The f,, of untreated Lula is about 0.02%; near the
plateau of the isotherm, the amount of DP on the surface is about
40 mmol/kg, or fo, = 1%. Thus, the fg, of the sorbent can be
increased by a factor of about 50 by loading with DP.

Asignificant feature of this isothermis its nonlinearity: at3 mM OP
in solution, the distribution ratio of DP is approximately 10 L kg",
while at 10 uM DP in solution, the distribution ratio is approximately
5000 L kg-1. Thus, the lower the concentration of DP in solution,
the more tenaciously it adheres to the sorbent. This behavior is
ideal for establishing absorbent barriers by injecting a pulse of
cationic surfactant into the ground. At high concentrations the
surfactant would be mobile and would disperse to cover a zone,
while at low concentrations it would be virtually immobile and
provide a coating to particles to retard migration of HOC.

Distribution of Chlorobenzenes on DP-treated
Aquifer Material

The distribution ratios of the chlorobenzenes between the DP-
treated aquifer material and 0.01 M CaCl, are shown in Figure 7
as a function of DP loading. The values of log K predicted for the
HOC from Koy (Miller et al., 1984) and f,, by the equation of
Schwarzenbach and Westall (1981) are lower than the values
found here. This result indicates that the DP accommodates the
chlorobenzenes much better than natural organic matter, per
carbon atom. The correlation developed by Schwarzenbach and
Westall (1981) included all of the compounds used in this study,
and the sorbents were a variety of environmental materials
including soils, sediments, aquifer materials and activated sludge.
Third order polynomials were fit to these data for incorporation in
the transport model of chlorobenzenes on DP-treated aquifer
material (Wagner et al., 1994). The chlorobenzene distribution
ratios presented in Figure 7 were determined from “one-point”
isotherms, rather than the slopes of multipoint isotherms, thatis,
the linearity of the chlorobenzene isotherms was not verified
experimentally in this study. Other studies (Smith and Jaffé,
1991; Smith et al., 1990; Boyd et al., 1988), have indicated that
isotherms of HOC are linear for sorbents treated with cationic
surfactants with long alkyl chains.

Breakthrough of Chlorobenzenes

The breakthrough curves of the chlorobenzenes through the
untreated and treated aquifer material are shown in Figure 8. In
these experiments, the DP on the surface was in equilibrium with
10 pM DP in the mobile phase, and the surface concentration did
not vary. The interstitial volume was determined from the NOz”
breakthrough curves. Physical properties of the column are given
in Table3. Curves calculated from a model as well as experimental
data are shown. Two approaches could be taken to determination
of the adjustable parameters in this model: (i) use independently
determined parameters from the batch experiments and the
column experiment with a nonsorbing tracer, and compute “ab
initio” breakthrough curves to be compared with the experimental
data; or (ii) adjustthe parameters in the model of the chlorobenzene
breakthrough curves to optimize agreement with the experimental
data, and then compare these parameter values with the
independently determined ones. Since the data from the
chlorobenzene breakthrough curves are as reliable as the data
from the batch experiments, we have chosen the second approach.

The model of the breakthrough curves is the one-dimensional,
advection-dispersion model for steady flow through a
homogeneous porous medium with local equilibrium:

aC (z,t)
ot

ac (z,t)
az

-2 #c (zt)
a2

u
R 7)

D

where C(z,t) is the concentration in the aqueous phase at time t
and distance along the column z, D is the dispersion coefficient,
R is the retardation factor, and u is the flow velocity. For all of
these experiments, the apparent dispersion coefficient, D, was
set to 0.0002 cm? s; the value of D estimated from the nitrate
breakthrough was 0.0003 + 0.0001 cm? s. The values of K
determined from the breakthrough curves can be compared with
those determined from batch experiments in Table 4; they are not
significantly different. The solid lines in Figure 8 were calculated



1000 T T T
[ (A) ]
fo) fe)
[ ° |
500 - 7
L, 4
L © Montmorillonite 4
0 nm 1 n 1 rn ]
200 T T T T T
3 r (B) e 1
= e °
3 . °
2 L 4
€ @ EPA-12 +
= @ LulaNé 1
2 = 8 4 Kaolinite 1
Pa 4
a A +
L A 1 A
r T T :
(c)
2F a 7
waa
be
t « Borden Sand
0 A 1 1 L 1.
Oo 2 4 6
C(w) [mM]

Figure 2.

Adsorption isotherms of dodecyipyridinium on different materials. Data are from the high concentration range of Figure 3, plotted here on

alinear scale. Note the correspondence between the plateaus of the adsorption isotherms and the cation-exchange capacities listed in
Table 1. Range of pH values for the data in the isotherms: (A) montmorillonite, pH = 7.13-5.46; (B) kaolinite, pH = 5.39-4.72; soil EPA-
12, pH =7.43-7.00; Lula aquifer material, pH = 6.81-4.49; (C) Borden sand, pH = 9.39-9.02. Solutions contained 0.01 MNat*; ratios of solids

to liquids are given in the legend of Figure 3.

These low values of n indicate nonlinearity in the isotherms, but
not the step-function that one perceives from Figure 2. (A linear
isotherm is one in which the surface concentration is directly
Proportional to the solution concentration, i.e.,n = 1 in Equation 1.)
The nonlinearity extends to surface concentrations that are as low
as 0.02% of the maximum surface coverage measured in thecase
of EPA-12 soil. As will be discussed, this nonlinearity is attributed
to heterogeneity of'adsorption sites.

The isotherms in Figure 3 exhibit no evidence of hemimicelle
formation, which we define operationally as a region in the log-log
isotherm with slope greater than one. Hemimicelle formation
occurs at surfaces when cooperative hydrophobic interactions
among adsorbed species dominates over repulsive and entropic

effects. The absence of an observable hemimicelle region in the
isotherms is attributed to the heterogeneity of the materials and
the electroneutrality of the ion-exchange reactions.

Effect of pH on Adsorption of DP

The effect of [H+] on the adsorption of DP is shown in Figure 4. A
very small dependence of D, on pH was observed for the
environmental sorbents; the slopes (A log D, /(A log [H*])
determined from linear regression of the data in Figure 4 wereless
than -0.05.

Two mechanisms through which pH can influence adsorption are
(i) a change in surface potential from protonation-deprotonation

acaba


Thus we expect that the data in Figure 11, which were obtained
in batch experiments without strict pH control, were influenced
slightly by this pH dependence of Cu adsorption. In these
experiments, the pH decreased from pH ~ 5.5 to pH ~ 5.0 as the
mildly acidic Cu(II) solution was added, presumably causing the
value of log D to decrease by up to about 0.10 -0.15 log unit. This.
change is large enough to be detected but not so large that it
would grossly affect our conclusions if it were not considered
explicitly in the data interpretation.

This weak pH dependence of Cu(II) resembles that expected for
ion exchange of Cu(Il) more than that expected for surface
complexation. Itis not unlikely that surface complexation of Cu(II)
sets in at slightly higher pH values. However, hydrolysis and
precipitation of Cu(!l) also set in at higher pH values.

Comparison of effect of DP on Cu(II) with effect of
Cu(II) on DP

The effects of DP on the adsorption isotherms of Cu(II) and the
effects of Cu(Il) on the adsorption isotherms of DP can be
compared in Figures 13a and 13b. At first glance, it might appear
that DP has a much stronger effect on Cu(II) than vice versa;
however, it should be noted that the total concentrations of the
competing DP are up to 1000 uM in Figure 13a, while the total
concentrations of the competing Cu(II) are up to only 100 pM in
Figure 13b. If one compares the isotherms with only up to 100 uM
total concentration of competing ion, the differences are not so
marked.

The series of Cu(II) isotherms obtained with various total
concentrations of DP in the system illustrate that large amounts
of DP in the system clearly suppress the adsorption of Cu(II). The
“nominal” amounts of DP on the surface are 0, 9, 17, 37, and 42,
mmol/kg, corresponding to Tpp in the system of 0, 50, 100, 400,
and 1000 uM. The solution pH varied from about pH 5.6 at low
concentrations of Cu(II) to about pH 4.9 at high concentrations of
Cull).

The isotherms show essentially no effect of DP until about 9
mmol/kg DP (about 20% of the DP exchange capacity) is on the
surface. Then the isotherms are flattened substantially by 17, 37,
and 42 mmol/kg DP on the surface. The flattening of the
isotherms seems to indicate some sort of saturation phenomenon:
either DP adsorbs to a certain extent, leaving a certain number of
sites free for Cu(II), or, DP adsorption alters the surface charge on
the surface, effectively causing “electrostatic saturation”.
Electrophoretic mobility measurements (data not shown) are
consistent with the latter explanation.

Total amount of DP and Cu(II) on the surface

Some insight into the nature of the competition between DP and
Cu(I!) can be gained from Figure 14, in which the amount of Cu(II)
adsorbed is plotted against amount of DP on the surface, for
several values of total Cu(I!) (20, 40, 70, and 100 1M) and total
DP (50, 100, 400, 1000 »M).

Atlow total concentration of both sorbates on the surface (Co, (s),
Cpp(s) < 10 mmol kg"), adsorption of the two species appears to
be independent (i.e., at constant total Cu(II), the amount of Cu(!!)
on the surface seems relatively independent of the DP surface
amount, and vice versa). Only at the point of 100 uM total Cu(I!)

15

and 50pM total DP are the values of Cpp(s) and Co,(s) shifted
from their no-competition values, with the Cu(|!) shifted more than
the DP.

At higher concentrations on the surface, both Cpp(s) and Cc,(s)
are shifted from their no competition values, with the Cu(II) always
shifted more than the DP. Ultimately the DP and Cu(II) on the
surface fall along the line corresponding to the saturation of the 45
mmol kg"! exchange capacity of the sorbent. It appears that the
DP affects the Cu(Il)-more than the Cu(II) affects the DP.

From the standpoint of environmental management, probably the
best way to summarize these data (Figures 11-14) is to say that
the capacity of the Lula aquifer material for Cu(II) is reduced to 25
% of the initial value in the presence of 42 mmol/kg DP, etc. The
nonlinearity of the isotherms makes it difficult to express the
effects in terms of log D. Furthermore, it appears that another
cationic surfactant that adsorbed more strongly than DP would
accomplish the same mission at lower aqueous-phase
concentrations.

From the standpoint of a mechanistic understanding of the
competition, one could consider two basic mechanisms: (i) OP
adsorbs much more intensely than Cu (.c., effectively irreversibly),
renders adsorption sites unavailable to Cu(I!), and causes an
effective reduction in capacity of the sorbent for Cu(II); or (ii) DP
and Cu(II) compete through a classic mass-action material-
balance mechanism, with or without electrostatic term. These
mechanisms are discussed in more detail by Westall and Chen
(1994).

SUMMARY

Large cationic surfactants, such as dodecylpyridinium, are strongly
adsorbed to aquifer materials. Adsorption depends strongly on
the concentration of major cations (e.g., Na+, Ca2+) in solution, but
is almost independent of solution pH. The capacity for adsorption
is close to the CEC of the sorbent; from a limited amount of data,
it rs that the organic carbon content of the sorbent has a
relatively minor effect on adsorption intensity and capacity.

The adsorption isotherms are distinctly nonlinear. Plotted on a
linear scale, theisotherms have almosta step function appearance,
suggesting that the adsorption reaction goes to completion and
that the surfactant is effectively irreversibly bound; however,
plotted on a logarithmic scale over several orders of magnitude,
the isotherms reflect the equilibrium distribution that is observed
inbatch experiments on adsorption-desorption kinetics andcolumn
experiments. The isotherm data conform to the Freundlich
equation C(w)" K = C(s), with a slope n = 0.6.

At high aqueous concentrations, the cationic surfactant adsorbs
rather weakly, the isotherm leveling off at approximately the
cation exchange capacity of the sorbent. At lower aqueous
concentrations, it adheres more tenaciously to the sorbent. This
behavior is consistent with the highly convex isotherm. This
behavior is ideal for establishing absorbent barriers by injecting a
pulse of cationic surfactant into the ground. Athigh concentrations
the surfactant would be mobile and would disperse to cover a
Zone, while at low concentrations it would be immobile and
provide a coating to particles to retard migration of HOC. The
cationic surfactant significantly promotes the sorption of HOC to
the aquifer material even with low amounts of surfactant on the
surface of the aquifer material.


A transport model based on the local equilibrium assumption and
on results of batch experiments can be used to predict HOC
column behavior reasonably well, although additional work on
dispersion of the surfactant breakthrough curveis warranted. The
agreement between batch and column experiments indicates that
“particle concentration effect” is negligible for the materials used
in this study. Also the fact that the local equilibrium assumption
‘seems to hold indicates that sorption of the HOC is rapid compared
to transport through the column.

It may be possible to design a system that allows for limited
dispersal of the cationic surfactant on injection and retardation of
HOC in the region treated with the dispersed surfactant. To a first
approximation, such a system can be designed by a model based
on reversible adsorption of the cationic surfactant and on the local
equilibrium assumption (for the materials used in this study).

There is relatively weak competition between DP and metal ions
(Cu(Il), Cd(!!). Pb(Il)) for adsorption sites on Lula aquifer material.
From the standpoint of environmental management, probably the
best way to summarize these results is to say that the capacity of
the Lula aquifer material for Cu(II) is reduced to 25 % of the initial
value in the presence of 42 mmol/kg DP, etc. The nonlinearity of
the isotherms makes it difficult to express the effects in terms of
Jog D. It appears that another cationic surfactant that adsorbed
more strongly than DP would accomplish the same mission at
lower aqueous-phase concentrations.

From the standpoint of a mechanistic understanding of the
competition, one could consider two basic mechanisms: (}) DP
adsorbs much moreintensely than Cu (i fectively irreversibly),
renders adsorption sites unavailable to Cu(II), and causes an
effective reduction in capacity of the sorbent for Cu(II); or (ii) OP
and Cu(II!) compete through a classic mass-action material-
balance mechanism, with or without electrostatic term.

g

QUALITY ASSURANCE STATEMENT

Allresearch projects funded by the U.S. Environmental Protection
Agency that make conclusions or recommendations based on
environmentally related measurements are required to participate
in the Agency Quality Assurance Program. This project was
conducted under an approved Quality Assurance Project Plan
and the procedures therein specified were used (with exceptions
noted). Information on the plan and documentation of the quality
assurance activities and results are available from the Principal
Investigator.

DISCLAIMER

The information in this document has been funded by the United
States Environmental Protection Agency under Cooperative
Agreement No. 816875 to Oregon State University. It has been
subjected to the Agency's peer and administrative review, and it
has been approved for publication as an EPA document. Mention
of trade names or commercial products does not constitute
endorsement or recommendation for use.

REFERENCES

Baetsié, L. H. In Progress in Nuclear Energy, Series XiL. Health
Physics; Duhamel, A. M. F., Ed.; Pergamon: Oxford, 1969; Vol.
2, pp 707-730.

17

Bales, R. C.; Szecsody, J. E. In Chemical Modeling of Aqueous

; Melchior, D. C.; Bassett, R. L. Eds.; ACS Symposium
Series 416; American Chemical Society, Washington, DC, 1985;
pp 526-538.

Beveridge, A.; Pickering, W. F. Water Res, 1983, 17, 215-225. -
Bijsterbosch, B. H. J, Colloid interface Sci, 1974, 47, 186-198.

Bouchard, D. C.; Powell, R. M.; Clark, D. A. J, Environ, Sci, Health
1988, A23, 585-601.

Boyd, S. A.; Mortland, M. M.; Chiou, C. T. Soil Sci, Soc, Am, J.
1988, 52, 652-657.

Brownawell, B. J.; Chen, H.; Collier, J. M.; Westall, J. C. Environ,
Sci, Technol, 1990, 24, 1234-1241.

Ford, , W. P. J.; Ottewill, R. H.; Parreira, H. C. J. Colloid Interface
Sci, 1966, 21, 522-533.

Fuller, C. C.; Davis, H. A. Geochim, Cosmochim, Acta 1987, 51,
1491-1502.

Giddings, J. C. Unified Separation Science; John Wiley & Sons:
New York, 1991.

Greenland, D. J.; Quirk, J. P. In. ir i
Conference on Clays and Clay Minerals; Swineford, A., Ed.;
Pergamon: New York, 1960; pp. 484-499.
Hassett, J. J.; Means, J. C.; Banwart, W. L.; Wood, S. G. Sorption
Properties of Sediments and Energy-Related Pollutants; U.S.
Environmental Protection Agency. National Technical Information
Service: Springfield, VA, 1980; EPA-600/3-80-041.

Malik, W. U.; Srivastava, S. K.; Gupta, D. Clay Minerals 1972, 9,
369-382.

Miller, M. M.; Ghodbane, S.; Wasik, S. P.; Terwarl, Y. B.; Martire,
DE. 1984, 29, 184-190.

Ralston, J.; Kitchener, J. A. J. Colloid Interface Sci, 1975, 50, 242-
249.

Schwarzenbach, R. P.; Westall J. Environ, Sci, Technol, 1981,
15, 1360-1367.

Smith, J. A.; Jaffé, P. R. Environ, Sci, Technol, 1991, 25, 2054-
2058.

Smith, J. A.; Jaffé, P. R.; Chiou, C. T. Environ, Sci, Technol, 1990,
24, 1167-1172.

Ter-Minassian-Saraga, L. J. Chim, Phys. 1966, 63, 1278-1280.
van Olphen, H.; Fripiat, J. J., Eds. |

; Permagon: New
York, 1979.

Wagner, J.; Chen, H.; Brownawell, B. J.; Westall, J.C. Environ,

Sci, Technol, 1994, 28, 231-237.


criteria would have to be considered if field studies were planned.
Labeled N-[1-14C]dodecylpyridinium bromide was obtained from
Pathfinder Laboratories, Inc., St. Louis, MO. The specific activity
was 10.6 Ci/mol, and the purity of the label was greater than 99%
as determined by HPLC with radioactivity detection. Solutions of
DP used in this study were checked periodically for purity; no
evidence of degradation was found. Unlabeled DP was purchased
as the chloride monohydrate from Aldrich Chemical and used as
received. The critical micelle concentrations (CMC) at 25°C in
0.01 and0.1 M NaCl solutions are estimated from the data of Ford
et al. (1966) as 13 and 6.3 mM, respectively.

The hydrophobic organic compounds were chlorobenzenes
(Figure 1): 1,2-dichlorobenzene (DCB), 1,2,3-trichlorobenzene
(TCB), 1,2,3,4-tetrachlorobenzene (TeCB), and pentachloro-
benzene (PeCB). These compounds are moderately hydrophobic,
with 3.4 < log K,,,< 5.0 (Miller et al., 1984).

Several low-organic-carbon aquifer materials and pristine
minerals were used as sorbents (Table 1). The sorbent for the
column experiments was Lula N6, a low-organic-carbon aquifer
material, obtained from the R. S. Kerr Environmental Research
Laboratory, U.S. EPA, Ada, Oklahoma. The materialis essentially
an iron oxide coated silica sand, with fraction of organic carbon,
foe: Of 0.2 Qo kg"! (0.02%). The fraction that passed through a 250
um sieve was used in all column experiments. Other sorbents
were examined for comparison. A 74-210 um fraction of aquifer
material Borden sand (Bouchard et al., 1988; Fuller et al., 1987),
isolated by dry-sieving, was used. EPA-12 soil (Hassett et al.,
1980) was provided by Professor John J. Hassett of the University
of Illinois. Georgia kaolin (KGa-1) and Wyoming Na-
montmorillonite (SWy-1) (van Olphen et al., 1979) were obtained
from Source Clays Repository at the University of Missouri.
Before use in batch sorption experiments, each sorbent was
washed several times to equilibrate the sorbent with the electrolyte
and to reduce the amounts of suspended material that are not
separated from the aqueous phase by centrifugation.

TABLE 1. SOME PROPERTIES OF THE SORBENTS USED
IN THIS STUDY?
organic CEC,
sorbent material carbon, % mmol/kg
kaolinite (KGa-1) 0.022 20
Na-montmorillonite (SWy-1) 0.020 7640
Borden sand 0.026, 0.020¢ 7.09, 23¢
Lula N6 aquifer material 0.021, 0.033¢ 90¢
EPA-12 soil 2.12, 2.33¢ 136¢

Organic carbon contents were determined on a LECO carbon
analyzer, unless otherwise referenced. Procedures for cation
exchange capacity (CEC) determination are given in the
reterences cited.

van Olphen et al., 1979.

Bouchard et al., 1988.

Fuller and Davis, 1987.

Hassett et al., 1980.

eacg

Procedures

Batch sorption experiments were carried out as described by
Brownawell et al. (1990), Wagner et al. (1994), and Westall and
Chen (1994). Theconcentration distribution ratio, D,, iscalculated
from

D, = Ci(s) / Cw) (1)

where C;(s) (umol/kg) and C,(w) (uM) are the concentrations of the
component i associated with the sorbent and in the water,
respectively.

Column experiments were carried out as described by Wagner et
al. (1994). The column, end fittings, tubing, and low dead-volume
unions were stainless steel, which was found to be superior to a
glass-teflon system in reducing sorption of these moderately
hydrophobic compounds to surfaces of the apparatus.

RESULTS
Adsorption of DP to Particle Surfaces
Adsorption Isotherms of DP

Adsorption isotherms of DP at relatively high (mM) concentrations,
on several different sorbents, are shown in Figure 2. The
equilibrium concentrations of DP in solution were well below the
critical micelle concentration (CMC) in all cases. The surface
concentrations reach plateau levels in almost all cases. The
plateau concentrations correspond closely, but not strictly, to the
cation exchange capacities reported for the five sorbents (Table
1).

For high energy surfaces, plateau adsorption can occur at
concentrations much below the CMC (Greenland and Quirk,
1962). When plateau adsorption occurs below the CMC, the
Possibility of multiple layer adsorption is more likely. Multiple layer
adsorption has been suggested to interpret adsorption of cationic
‘surfactants onto silica (Ter-Minassian-Saraga, 1966; Bijsterbosch,
1974), montmorillonite (Greenland and Quirk, 1962), and other
clays (Malik et al., 1972; Ralston and Kitchener, 1975). While it
is tempting to speculate about the orientation and structure of DP
at the surface, the heterogeneity of the surface and the absence
of additional data preclude any firm conclusions.

From the data in Figure 2, it appears that adsorption could be
explainedasa simple, virtually quantitative, ion-exchange reaction,
with step-function isotherms. However, if the data are considered
over a wider range of concentrations, it is apparent that the
situation is more complicated. The logarithmic adsorption
isotherms of DP on five different materials, over many orders of
magnitude, are shown in Figure 3. From these data itis apparent
that DP adsorption cannot be expressed as a simple reaction
going to completion. Over a wide range of concentrations (below
about 10-100 4M DP in solution), the isotherms conform to the
Freundlich equation:

log Cpp(s) = n log Cpp(w) + log K (2)

where Cpp(s) is the amount of DP adsorbed to the surface,
Cpp(w) is the concentration of DP in solution, and K and n are
constants. The values of n determined from the slopes of the
isotherms in the low concentration region are approximately 0.6.


the mobility of the HOC in the subsurface environment. While this
process will occur naturally as cationic surfactants enter the
environment, this process might also be the basis for remedial
action at contaminated sites. For example, cationic surfactants
could be injected into an aquifer downgradient from a source of
HOC contamination to provide a temporary barrier against
migration. Alternatively they could be used in conjunction with in-
situ bioremediation to eliminate the HOC contamination; the
surfactant could retard a contaminant such as TCE long enough
for a relatively slow process such as reductive dehalogenation to
be effective. Related applications might include the control of
vapor emissions from in-situ biofilters and the increase in retention
time of HOC in slurry bioreactors.

Apossible side effect of such a treatmentcouldbe the displacement
of previously adsorbed metal cations through ion exchange with
the cationic surfactant. Alternatively, the cationic surfactant could
be considered as a means of accelerating desorption of metal
cations as part of a soil cleaning treatment.

Thus, a wide range of questions on the environmental chemistry
ofcationic surfactants exist depending on whether the compounds
are introduced as waste or as part of a treatment program, and
whether the goal is to enhance or to retard the mobility of other
pollutants, and whether the other pollutants behave more like
metal ions or more like hydrophobic organic compounds. This
study was undertaken to elucidate some these questions as a
basis for management of cationic surfactants compounds in the
environment.

We focus on batch and column experiments designed to examine
the ability of cationic surfactants to modify the surface of a low
organic carbon aquifer material to promote the sorption of HOC.
Several questions were addressed: (i) will the cationic surfactant
adsorb strongly enough that it will not itself migrate? (ii) will the
cationic surfactant adsorb so strongly that it will be impossible to
disperse it in the subsurface environment? (iii) how do major

cations (Na+, Ca2+) and pH affect adsorption of the cationic
surfactant? (iv) will the cationic surfactants promote the sorption
of the HOC to a sufficient degree? (v) how well can column
behavior be predicted from batch behavior? (vi) is it possible to
engineer a system to determine the amount of cationic surfactant
that must be applied to obtain a preselected retardation of HOC?
(vil) to what extent will the treatment of particle surfaces with
cationic surfactants release adsorbed metal ions?

Although beyond the scope of this effort, several other issues
would have to be considered before field application of cationic
surfactants could be undertaken, including: the ultimate fate of
the surfactant, the effect of the surfactant on the permeability of
natural aquitards, the effect of the cationic surfactanton subsurface
microorganisms (particularly if bioremediation were being
considered), and the availability of cationic-surfactant-sorbed
HOC for biological degradation.

METHODS

The study was carried out in three major steps: (i) study of
adsorption of acationic surfactant to particle surfaces (Brownawell
et al., 1990); (ii) study of the sorption of chlorobenzenes to
surfactant-modified surfaces (Wagner et al., 1994); and (iii) study
of adsorption competition between the surfactant and the metal
ions Pb2+, Cd?+, and Cu2+ (Westall and Chen, 1994). The original
papers should be consulted for details about experimental
procedures.

Materials

The cationic surfactant used in this study was dodecylpyridinium
(DP), which was selected for the ease of analysis by UV
spectrophotometry andits resistance to biodegradation in slurries
of soils and subsurface materials; the structure of DP is shown in
Figure 1. In a limited number of experiments, we have found that
its behavior resembles that of other cationic surfactants. Additional

i
+ Cr
OCB
Cl
cl
OP TeCB

i
Cl Cl
ct
TCB
cl
Cl Cl
Cl cr Cl
cl
PeCB

Figure 1. Structures of dodecylpyridinium chloride and chlorobenzene homologs.


C/Co

TCB

600 800

200 400
Pore Volume

CiCo

tp

"200 400 600 800
Pore Volume

PeCB

0 300 600 900 12001500

Pore Volume

Figure 10. The breakthrough curves of the chlorobenzenes on a column of Lula aquifer material that was pre-equilibrated with 3 mM DP in 0.01 M
CaCly, but, before the pulse containing the chlorobenzenes was applied to the column, the supply of DP to the column was stopped. The
circles represent the concentration of the respective chlorobenzenes in the column effluent. The dashed line represents the concentration
of DP in the column effluent (not the amount remaining on the column!), the dotted line represents the concentration of chlorobenzenes
in the column influent, and the solid line represents the concentration of chlorobenzenes in the column effluent calculated from the finite
difference model. The model is defined by the constants in Table 3 for the column, the adjusted isotherm for local equilibrium of DP (Figure
6) with dispersion coefficient of 0.02 cm? s', the partition constants of the chlorobenzenes defined by the data in Figure 7.

this study we have investigated the effect of DP on adsorption of
Pb(Il), Cd(Il), and Cu(Il) on aquifer materials from Lula OK,
Elizabeth City NC, and Cape Cod MA. The results of the study of
Lula-DP-Cu(II) system, which illustrate the kinds of interactions
that are to be expected, are presented here.

Adsorption Isotherms of DP and Cu(II) on Lula

Adsorption isotherms of DP in the absence of Cu(I!), and of Cu(I!)
in the absence of DP, on Lula in 0.01 M NaCl are shown in Figure
11. Qualitatively it is seen that DP adsorbs more strongly than
Cu(II) under the solution conditions given, pH = 5.5-5.0 and 0.01
M NaCl.

The strongly nonlinear isotherms make it impossible to compare
adsorption energies of Cu(I!) and DP as single values of log K. As
an alternative, one could compare apparent distribution ratios of

the sorbates at equal concentrations on the surface, at equal
concentrations in solution, or at equal total concentrations in the
‘system, depending upon the question of interest. For example,
the distribution ratios for DP and Cu at concentrations of 10 uMin
solution are log Dpp = 3.2 L kg? and log Do, = 2.2 Lkg:!. Thus,
the DP and Cui!) adsorb “moderately strongly” under the conditions
of these experiments.

Dependence of adsorption on pH

The pH dependence of the distribution ratios of Cu(II) and DP are
illustrated in Figure 12. The distribution of DP on Lulais effectively
independent of pH over a relatively wide range, while that of Cu(II)
appears to decrease slightly with increasing log [H*]:

dlog D/d log [H+] ~ 0.25.


A T 1 1
& 20 | 7 8 Typist
a Toull)/yM 40 o-10 |
E os y ae 2B
aor 10 IE — babe |
3 me o> O= 1000
8 “
> 3.0 - ° 4
2 :

-3.5 1 n n

6.5 5.5 45 3.5
tog Coy (w) [M]

-1.0 T T T —
co] B 4091000
e 15 Tpp/M 4
2 Toull)/uM
z -2.0 - O=20 7
6 ren
$ 25 + A = 100 4
2

-3.0 1 1 A 1

7.0 6.0 -5.0 -4.0 -3.0 -2.0

log Cpp (w) [M]

Figure 13. (A) Logarithmic adsorption isotherms of Cu(l!) on Lula aquifer material (0.005 kg/L in 0.01 M NaC) in the presence of DP. The “nominal”
amounts of DP on the surface are 0, 9, 17, 37, and 42, mmol/kg, corresponding to Tpp in the system of 0, 50, 100, 400, and 1000 uM. The
solution pH varied from about pH 5.6 at low concentrations of Cu(II) to about pH 4.9 at high concentrations of Cu(ll). (B) Logarithmic
adsorption isotherms of DP on Lula aquifer material (0.005 kg/L in 0.01 M NaCl) in the presence of Cu(tl). The Tgya in the system is 20,
40,70, 100 uM, yielding “nominal” (i.e., in the absence of DP) amounts of Cu(II) on the surface of 2, 3, 6, 9, 11 mmol/kg. The solution pH
varied from about pH 5.6 at ow concentrations of Cu(II) to about pH 4.9 at high concentrations of Cu(II).

25
Toul
= 20
20 + 40 7
= 4=70
‘o a1 |
~ 164
3
z 0 gm
zg 10 } Aa? 4
38 4a
co) ° e 400
5 —_———_—_ 10 |
o——o ° S
0 ro 1 A
ie) 10 20 30 40 50

pp (s) [mmol kg"t}]

Figure 14. Amount of DP and Cu(II) on the surface of Lula aquifer material (0.005 kg/L in 0.01 M NaCl), with all possible combinations of Tpp in the
system 0, 10, 50, 100, 400, 1000 uM and Touuy in the system of 0, 20, 40, 70, 100 uM.