NUREG/CR-6977 

Redox and Sorption 

Reactions of Iodine and 

Cesium During Transport 

Through Aquifer Sediments 

Office of Nuclear Regulatory Research 
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NUREG/CR-6977 

Redox and Sorption 

Reactions of Iodine and 

Cesium During Transport 

Through Aquifer Sediments 

Manuscript Completed: May 2008 

Date Published: March 2009 

Prepared by 

J.A. Davis and P.M. Fox 

U.S. Geological Survey 

Menlo Park, Ca 94025 

M. Fuhrmann, NRC Project Manager 

NRC Job Code Y6462 

Office of Nuclear Regulatory Research
ABSTRACT 

Radioactive isotopes of iodine (131I and 129I) and cesium (137Cs) are important 

contaminants present in nuclear waste. These radioisotopes have been introduced into 

the environment through nuclear weapons tests as well as nuclear accidents such as 

Chernobyl. Although iodine is commonly found as iodide (I-), which is generally 

considered to behave conservatively, it has been proposed that iodide can be oxidized to 

elemental iodine (I2) or iodate (IO3

-

) by manganese oxides or nitrate, which may behave 

less conservatively in sediments due to uptake by organic matter or adsorption onto 

mineral surfaces. Cesium is generally present as a cation (Cs+

) and can be strongly 

adsorbed by sediments. 

In order to further our understanding of the chemical behavior of I and Cs in groundwater 

systems, a series of laboratory and field experiments were undertaken. The kinetics of Ioxidation by the manganese oxide, birnessite, was investigated under a variety of 

geochemical conditions. In order to determine Cs and I sorption and I oxidation, batch 

experiments with aquifer sediments and with binary sediment-Mn oxide systems were 

performed. Iodide transport was studied in a column filled with aquifer sediments. Three 

field tracer test experiments were performed to elucidate the redox chemistry and 

transport of I and Cs in an aquifer characterized by distinct geochemical zones: (1) 

injection of CsI into a well oxygenated zone of the aquifer, (2) injection of CsIO3 into a 

well oxygenated zone, and (3) injection of CsIO3 into a zone of the aquifer characterized 

by active Fe(III) reduction (but not sulfate reduction). 

In laboratory experiments, birnessite oxidized I-

 to I2 and IO3

-

 in a two-step process. The 

oxidation of I-

 proceeded according to first order kinetics with respect to initial Iconcentration, pH, and birnessite concentration. I2 sorption to birnessite was high (up to 

0.25 mmol/g), while IO3

-

 sorption to birnessite was an order of magnitude lower (up to 

0.024 mmol/g). Uptake of I- in batch experiments by sediments was fairly low at pH 4.8 

or above, as was I- retardation in column experiments at this pH. In column experiments 

at pH 4.50, the results suggested some oxidation of I-

 occurred due to a 7% loss of iodine 

mass exiting the column, presumably due to volatilization of elemental I2. IO3

-

 uptake in 
batch experiments with sediments was higher than that of I-

, reaching up to 12% 

adsorbed. Cs also adsorbed to aquifer sediments, with up to 22% removed from solution 

after 24 hours. 

Results from the field tracer tests show that I-

 was oxidized to both I2 (up to 46%) and 

IO3

-

 (up to 6%) in the oxic zone, with the extent of oxidation increasing with transport. A 

pulse of dissolved Mn was liberated from the sediments, providing evidence that Mn 

oxides were responsible for I-

 oxidation. Iodate was retarded relative to a conservative 

tracer (Br), arriving 5 days later at 3.9 m downgradient. In the Fe-reducing zone, IO3

-

was quickly reduced to I- without any observed production of I2 intermediate. About 60% 

of the iodate was reduced to I-

 in 1 m of transport, with complete reduction occurring 

after 3 m of transport. Cesium transport was retarded relative to the conservative tracer 

in all three tracer tests. In the oxic zone of the aquifer, the peak Cs concentration arrived 

3.9 m downgradient after 35 days of transport, whereas the Br peak arrived at 8 days. Cs 

was so attenuated during the first 3.9 m of transport that the maximum Cs concentration 

reached was only 6% of the injected concentration. 

While the Cs and I concentrations used in these experiments was much higher than would 

be relevant for concentrations of radioactive isotopes of these elements, the studies are 

relevant for revealing reaction mechanisms that affect the transport of these radionuclides 

in the environment. The results of these experiments demonstrate that not only can redox 

transformations of iodine easily occur in groundwater systems, but also that I2, IO3

-

, and 

Cs behave non-conservatively by adsorbing to sediments and minerals. The results 

indicate the importance of considering the complex redox and sorption chemistry of 

iodine when predicting its transport in waste plumes. 
Foreword

11

I

I

I

The speciation and related behavior of iodine is controlled by its oxidation state, which can change

according to different oxidation/reduction levels found in typical groundwater systems. This report,

prepared by staff of the U.S. Geological Survey for the U.S. Nuclear Regulatory Commission, describes

experiments examining transport and speciation changes of iodine in groundwater and the rates of

speciation change under different levels of oxygen content. These changes substantially affect

assumptions about transport of iodine in near surface groundwater and shoulq. be considered in

Perfonnance Assessment models of systems containing significant quantities of radionuclides of iodine,

such as waste forms generated incidental to nuclear fuel reprocessing.

Radioactive isotopes of iodine and cesium are important contaminants with widely differing chemical

properties. Cesium is monovalent and tends to strongly adsorb on many minerals. Iodine, in contrast,

sorbs poorly and can undergo several oxidation-reduction reactions; the products of these reactions may

have differing sorption properties. To further understand the chemical behavior of iodine in groundwater

systems, a series oflaboratory and field experiments were undertaken. Small-scale batch cesium and

iodine sorption experimentswere conducted with aquifer sediments and binary sediment-manganese

oxide systems.Iodide transport was studied in a column filled with aquifer sediments. Three field tracer

experimentswere performed to define the redox chemistry and transport of iodine and cesium:

. injection of CsI into a well-oxygenated zone of the aquifer

. injection of CsI03 into a well-oxygenated zone of the aquifer

. injection of CsI03 into a zone of the aquifer characterized by active iron (ill) reduction (but not

sulfate reduction).

These experimentswere conducted on a very well-characterized, heavily-instrumented site, allowing

observation of the differential transport of cesium, iodine, and the conservative tracer bromine.

Redox transformations of iodide easily occurred in oxygenated groundwater systems where it was

oxidized to both iodine and iodate. As iodide was oxidized, manganese appeared in the groundwater,

indicating that manganese oxides caused the iodine oxidation. In the reducing zone of the aquifer, iodate

was reduced to iodide and was transported at the same rate as nonreactive bromine tracer.. However, the

rate of iodate transport in oxic groundwater was retarded by a factor of about 1.4. In this report, smallscale experiments clarified processes controlling iodine speciation and identified analyses necessary to

confimi those processes in the field experiments. The field experiments demonstrated that very small

differences in location, associated with changed redox chemistry, can alter iodine speciation and transport

and mwt be considered in conceptual models assessing perfo~e g,fiodine-bearing systems.

.stiadaLuf, Director

Division of Risk Analysis

Office of Nuclear Regulatory Research

U.S. Nuclear Regulatory Commission

l

v

CONTENTS 

ABSTRACT .........................................................................................iii 

FOREWORD..........................................................................................v 

FIGURES...............................................................................................ix 

1 INTRODUCTION ............................................................................... 1 

2 LABORATORY STUDIES OF CESIUM AND IODINE ADSORPTION AND 

REDOX REACTIONS BY AQUIFER SEDIMENTS AND MINERALS .............3 

2.1 Materials and Methods ........................................................................3 

2.1.1 Synthesis of Birnessite .....................................................................3 

2.1.2 Aquifer Sediments ..........................................................................3 

2.1.3 Iodine-Birnessite Batch Experiments ....................................................3 

2.1.4 Cesium and Iodine Batch Experiments with Aquifer Sediments .....................4 

2.1.5 Column Experiment.........................................................................4 

2.1.5.1 Column Packing ............................................................................4 

2.1.5.2 Transport Experiment Apparatus.........................................................4 

2.1.5.3 Column Operation ..........................................................................5 

2.1.6 Iodine and Tritium Analysis ...............................................................5 

2.2 Results and Discussion .....................................................................6 

2.2.1 Kinetics of Iodide Oxidation by Birnessite .............................................6 

2.2.1.1 Effect of Ionic Strength ....................................................................6 

2.2.1.2 Effect of pH .................................................................................6 

2.2.1.3 Effect of Birnessite Concentration ........................................................6 

2.2.1.4 Elemental Iodine Oxidation by Birnessite ..............................................7 

2.2.1.5 Elemental Iodine and Iodate Adsorption ................................................8 

2.3 Rate Laws and Modeling Predictions ........................................................8 

2.4 Cesium and Iodine Reactions with Aquifer Sediments ..................................11 

2.4.1 Iodide Adsorption and Oxidation ........................................................11 

2.4.2 Iodate Adsorption ..........................................................................12 

2.4.3 Cesium Adsorption ........................................................................13 

2.4.4 Iodide Transport in Column Experiments with Aquifer Sediments ................14 

2.5 Summary of Laboratory Studies ............................................................15 
3 FIELD STUDIES OF CESIUM AND IODINE TRANSPORT UNDER VARIABLE 

ENVIRONMENTAL CONDITIONS .....................................................19 

3.1 Site Description ..............................................................................19 

3.2 Field Experiment Methodology ...........................................................19 

3.3 Results and Discussion ......................................................................21 

3.3.1 Iodine Oxidation in the Oxic Iodide Tracer Test .....................................21 

3.3.2 Iodate Adsorption in Oxic Iodate Tracer Test ........................................21 

3.3.3 Iodate Reduction in the Anoxic Iodate Tracer Test ..................................22 

3.3.4 Cesium Retardation and Attenuation ...................................................23 

3.4 Summary of Field Studies ....................................................................24 

4 SUMMARY ...................................................................................29 

5 REFERENCES ................................................................................31 
FIGURES 

Figure 1.1 pe-pH diagram for various redox couples under conditions relevant to the 

laboratory and field experiments described in this report.......................................2 

Figure 2.1. Experimental design for column experiment.......................................5 

Figure 2.2. (A) Iodide oxidation by birnessite (1g/L) at pH 5.00 in 0.1M NaClO4 solution. 

The initial I-

 concentration was 1.00 mM. (B) Plot of the same data for the first 8 hours of 

the experiment.........................................................................................7 

Figure 2.3 Plot of ln[I-

] versus time for 1 mM I- oxidation by birnessite (1 g/L) at pH 5.00 

in 0.1M NaClO4 solution............................................................................7 

Figure 2.4. Effect of ionic strength on I- oxidation by birnessite at pH 4.5-5.0, showing 

the disappearance of I-

 and the production of IO3

-

 and I2.......................................9 

Figure 2.5. Effect of pH on iodide oxidation by birnessite in 0.1M NaClO4 solution, 

showing the disappearance of I- and production of IO3

-

 and I2.................................9 

Figure 2.6. pH dependence of the log of the first order rate constant, k', describing the 

disappearance of iodide from solution, determined from reacting 1 mM I- with 1 g/L 

birnessite in 0.1M NaClO4 solution...............................................................10 

Figure 2.7. Oxidation of 0.1 mM I-

 at pH 5.00 in 0.1M NaClO4 solution for various 

concentrations of Birnessite, showing the disappearance of I-

 and the production of I2 and 

IO3

-

 versus time......................................................................................10 

Figure 2.8. Log-log plot of the pseudo-first order rate constant for iodide disappearance 

versus birnessite concentration for experiments performed with an initial I-

 concentration 

of 0.1 mM in 0.1M NaClO4 solution at pH 5.0.................................................11 

Figure 2.9. (A) I2 oxidation by birnessite (1 g/L) at pH 5.00 in 0.1M NaClO4 solution. 

The initial iodine concentration was 0.24 mM (as I). (B) Plot of the same data for the 

first 8 hours of the experiment....................................................................11 

Figure 2.10. Plot of ln[I2] over time for 0.24 mM I2 (as I) oxidation by birnessite (1 g/L) 

at pH 5.00 in 0.1M NaClO4 solution..............................................................12 

Figure 2.11. Log of k' for the disappearance of I2 from solution determined at 1 g/L 

birnessite in 0.1M NaClO4 solution at pH 5.0, plotted against the initial I2

concentration.........................................................................................12 

Figure 2.12. Iodine sorption by birnessite (1g/L) at pH 5.0 in 0.1M NaClO4 solution....12 
Figure 2.13. Iodate sorption by birnessite (1 g/L) at pH 5.0 in 0.1M NaClO4 solution. (A) 

Kinetics of IO3

-

 sorption for sample spiked with 0.50 mM IO3

-

. (B) Sorption isotherm 

showing adsorbed IO3

-

 versus IO3

-

 concentration in solution.................................13 

Figure 2.14. Comparison of actual (solid lines) and predicted (dashed lines) iodide 

oxidation by birnessite (1 g/L) using k1=1.5 and k2=0.0042 for (A) the first 10 hours, and 

(B) 72 hours..........................................................................................13 

Figure 2.15. Comparison of actual (solid lines) and predicted (dashed lines) iodide 

oxidation by birnessite (1 g/L) using k1=1.5 and k2=0.0473 for (A) the first 10 hours, and 

(B) 72 hours..........................................................................................14 

Figure 2.16. Comparison of actual (solid lines) and predicted (dashed lines) iodide 

oxidation by birnessite (1 g/L) using k1=1.5 and k2=0.1408..................................14 

Figure 2.17. Time dependence of iodide loss from solution via sorption or oxidation by 

sediments (1000g/L)................................................................................15 

Figure 2.18. Time dependence of iodate adsorption onto sediments (400g/L).............15 

Figure 2.19. Iodate adsorption onto sediments (400 g/L) at different pH values after 1 

week of reaction.....................................................................................16 

Figure 2.20. Time dependence of Cs adsorption onto aquifer sediments (400 g/L)........16 

Figure 2.21. Breakthrough curves for tritium (3

H) and iodide at pH 4.75 and 4.50........16 

Figure 3.1. Map of the Cape Cod field site showing the extent of the wastewater plume 

and the locations of the two study areas (wells F625 and F168)............................. 20 

Figure 3.2. Breakthrough curves showing Br-

, I-, I2, IO3

-

, and Mn concentrations at 

various distances downgradient of the injection well in the oxic I- tracer test..............22 

Figure 3.3. Iodate and Br breakthrough curves at various distances downgradient from the 

injection well in the oxic iodate tracer test......................................................23 

Figure 3.4. Br-

, IO3

-

 and I- breakthrough curves at various distances downgradient from 

the injection well for the anoxic iodate injection...............................................24 

Figure 3.5. Cs, Br, and K breakthrough curves at various distances downgradient from 

the injection well in the oxic iodide tracer test (well F168-M17-02)........................25 

Figure 3.6. Cs, Br, and K breakthrough curves at various distances downgradient from 

the injection well in oxic iodate tracer test (well F168-M17-08).............................26 

Figure 3.7. Cs, Br, and K during the anoxic iodate tracer test (well F625-M12-09).......27 
TABLES 

Table 2.1. Basic Characteristics of Aquifer Sediment Composite Samples..................4 

Table 2.2. Summary of Pseudo First Order Rate Constants Determined from the 

Disappearance of The Starting Compound (Iodide Or Iodine).................................8 

Table 2.3. Reaction Conditions and Speciation of Iodine after Reaction of Iodide with 

Aquifer Sediments..................................................................................15 

Table 3.1. Background Chemistry at Injection Wells Used In Field Experiments.........20 

Table 3.2. Tracer Test Characteristics and Injected Tracer Concentrations.................21 

1

1 INTRODUCTION 

Radioactive isotopes of I and Cs are important 

contaminants produced from nuclear fission. 

These radioisotopes are components of nuclear 

waste and have been introduced into the 

environment through fallout from nuclear 

weapons tests and nuclear accidents such as 

Chernobyl (Filipovic-Vincekovic et al., 1991; 

Blagoeva and Zikovsky, 1995). 131I has a very 

short half life (8 days) and high activity, whereas 

129I has a long half life (1.57 x 107

 years). While 

129I is a naturally occurring radioactive isotope 

of iodine, anthropogenic sources have greatly 

increased the pool of 129I in the environment 

(Rao and Fehn, 1999; Oktay et al., 2000; 

Reithmeier et al., 2006). 137Cs and 134Cs have 

half lives of 30.2 and 2 years, respectively. An 

understanding of the chemical behavior of 

iodine and cesium in natural systems is 

important for understanding the fate of their 

radioisotopes derived from nuclear wastes. 

 

Iodine can exist in oxidation states ranging from 

-1 to +7, but in natural waters is most commonly 

present as iodide (I-

), iodate (IO3

-

), and organic I 

(Luther III et al., 1991; Oktay et al., 2001; 

Schwehr et al., 2005). A pe-pH diagram for 

various redox couples is shown in Figure 1.1, 

calculated for conditions relevant to experiments 

described in this report. While in theory O2 may 

oxidize I-

, this reaction is very slow and thus the 

H2O/O2 redox couple is not shown in the 

diagram. Significant research into the redox 

behavior of iodine in seawater (e.g. Anschutz et 

al., 2000; Truesdale et al., 2001) has been 

performed. Anschutz et al. (2000) proposed 

manganese oxides (MnO2) and nitrate as 

possible electron acceptors in the oxidation of 

iodide to iodine and/or iodate (see Fig. 1.1): 

1⁄2MnO2 + I-

 +2H+ Æ 1⁄2Mn2+ + 1⁄2I2 + H2O 

(Eq. 1-1) 

I

-

 + 1/5NO3

-

 + 6/5H+ Æ 1/10N2 + 1⁄2I2 + 3/5H2O 

(Eq. 1-2) 

There is evidence from research in seawater that 

iodine redox chemistry is linked to that of 

nitrogen and manganese oxides, however, these 

reactions have not been directly proven. 

Both I-

 and IO3

-

 are considered to be highly 

mobile and unreactive with sediments. Sorption 

of these species to many minerals and sediments 

is low (Fuhrmann et al., 1998; Kaplan et al., 

2000). While most researchers have found that 

IO3

-

 sorption to minerals and sediments is 

slightly higher than I-

 sorption (Couture and 

Seitz, 1983; Hu et al., 2005), the opposite may 

be true for certain minerals such as magnetite 

(Fuhrmann et al., 1998). Elemental iodine (I2) 

can also be taken up by organic matter, 

including humic substances (Sheppard et al., 

1995; Warner et al., 2000; Schlegel et al., 2006), 

sorbed to mineral surfaces (Fuhrmann et al., 

1998) and volatilized to the atmosphere, 

providing additional pathways for the nonconservative behavior of iodine. 

Cesium is commonly found as the monovalent 

cation Cs+

, and sorbs strongly to sediments via 

cation exchange (Zachara et al., 2002; 

Ainsworth et al., 2005). Cs sorption can usually 

be described using a two site model including a 

strong site and a weak site corresponding to the 

frayed edge sites (FES) and planar sites of 

phyllosilicates, respectively (Poinssot et al., 

1999; Zachara et al., 2002). Several researchers 

have noted that a fraction of the Cs is 

irreversibly sorbed (Comans et al., 1991; 

Krumhansl et al., 2001; Zachara et al., 2002). 

In this study, a combination of field and 

laboratory experiments were conducted to 

further our understanding of Cs and I 

interactions with aquifer sediments, specifically 

(1) iodide oxidation, (2) sorption of I species, 

and (3) sorption of Cs species. Through the use 

of laboratory experiments, the kinetics of iodide 

oxidation by the manganese oxide, birnessite, 

was studied under a variety of conditions, 

including variable pH, ionic strength, initial I 

concentration, and solid concentration. Cs and I 

batch experiments with aquifer sediments and 

with binary sediment-Mn oxide systems were 

performed to determine the extent of Cs and I 

sorption and I oxidation. Iodide transport was 

also studied in a column filled with natural 

aquifer sediments. 
2

In addition to laboratory studies, the results of 

three field experiments are described in the 

report. Tracer tests were performed to elucidate 

the redox chemistry and transport of Cs and I in 

an aquifer characterized by distinct geochemical 

zones: (1) injection of CsI into a well 

oxygenated zone of the aquifer, (2) injection of 

CsIO3 into a well oxygenated zone, and (3) 

injection of CsIO3 into a zone of the aquifer 

characterized by active Fe(III) reduction (but not 

sulfate reduction). 

-5

0

5

10

15

20

4.0 4.5 5.0 5.5 6.0 6.5 7.0

pe

pH

FeOOH/Fe2+

SO

4

2-/H

2

S

AsO

4

3-/AsO

2

-

I

2

/IMnO

2

/Mn2+

IO

3

-

/INO

3

-

/N

2

O

2

/H

2

O

Figure 1.1 pe-pH diagram for various redox couples under conditions relevant to the 

laboratory and field experiments described in this report. [SO4

2-]=10-4M; [N2]=0.77atm; 

[O2]=0.21atm; [NO3

-

], [Mn2+], [IO3

-

], [I2], [Fe2+], [H2S]=10-6M; and [I-]=10-3M. The 

activities for all solid phases and water were taken as 1.0. 
3

2 Laboratory Studies of Cesium and Iodine Adsorption and Redox Reactions 

by Aquifer Sediments and Minerals 

2.1 Materials and Methods 

2.1.1 Synthesis of Birnessite 

Birnessite (δ-MnO2) was synthesized by adding 

concentrated HCl to a boiling solution of 0.4M 

KMnO4 (McKenzie, 1971). The resulting 

precipitate was washed several times with 

deionized water and then placed in dialysis 

tubing in a tub of MilliQ-deionized water to 

further remove salts. The deionized water was 

changed daily until the conductivity of the rinse 

water was close to that of MilliQ-deionized 

water. The solid was confirmed to be birnessite 

by XRD, and the surface area was determined to 

be 41 m2

/g by 5 point N2-BET. The average 

oxidation state of Mn in the solid was 

determined to be 4.09 by oxalate titration (Hem 

et al., 1982). The resulting birnessite solid 

contained 52.4% (w/w) Mn and 9.2% (w/w) K. 

2.1.2 Aquifer Sediments 

Several sediment composites were used for the 

batch and column experiments. The composites 

were created from aquifer materials collected 

from the oxic zone of a water table in a sand and 

gravel aquifer located on Cape Cod, 

Massachusetts. The site where the sediments 

were collected has been described in detail 

elsewhere (Davis et al., 2000). The shallow, 

unconfined aquifer consists of permeable, 

stratified glacial sand and gravel deposits, with 

mineralogy dominated by quartz with minor 

accessory minerals. Surfaces of sediment grains 

are heavily coated with hydroxypolymer and 

clay coatings containing Fe, Al and Mn (Coston 

et al., 1995; Fuller et al., 1996). A description of 

the collection and characteristics of these 

sediments is given in Amirbahman et al. (2006) 

and references therein. One composite sediment 

sample was collected at the location R23AW, 

while two other samples were collected near 

well F168 (F168-15; F168 2005). The sample, 

F168 2005, was a new composite created to 

replace the exhausted F168-15 composite. 

To prepare composite sample F168 2005, four 

cores were collected near well F168 and the 

sediments located 15-20 ft (below ground 

surface) were air-dried. The sediments were then 

sieved to remove materials greater than 2mm, 

and the <2mm fractions from each of the cores 

were combined to create the composite sample. 

An aliquot of the new composite was extracted 

with 0.25M hydroxylamine hydrochloride at 

50 ̊C for 96 hours in order to determine the 

reductively extractable Mn content. Surface 

area was measured using 5-point N2 BET 

analysis. The results are compared with the 

other composites used in this study in Table 2.1. 

Typical values of extractable Fe(III) can be 

found in Coston et al. (1995) and Fuller et al. 

(1996). 

2.1.3 Iodine-Birnessite Batch 

Experiments 

Batch kinetic experiments were performed in 

order to investigate the effect of numerous 

variables on iodide oxidation by birnessite. 

Experiments were performed in a gas tight, 4-

port 250-mL glass reaction vessel in the dark. 

Iodide oxidation by birnessite was determined 

over a range of conditions, including ionic 

strength (0.01-0.1 M), pH (4.50-6.25), solidliquid ratio (0.1-1.0 g/L), and initial iodide 

concentration (0.1-1.0 mM). The kinetics of I2

oxidation to IO3

-

, and IO3

-

 adsorption by 

birnessite were also measured. 

The birnessite was suspended in a background 

electrolyte solution of 0.01 or 0.10 M NaClO4

and the pH was adjusted to the desired value 

with 0.2N HCl or NaOH. The suspension was 

stirred on a stir plate overnight in order to 

equilibrate the birnessite prior to iodine addition. 

The suspension was then spiked with 0.1 M 

sodium iodide solution, 0.1 M sodium iodate 

solution, or ~2 mM I2 (as I) stock solution to 

achieve the desired concentration of I. The I2

stock solution was created by shaking several 

grams of doubly sublimed solid I2 in 0.1M 

NaClO4 in the dark for several hours to create a 
4

 Table 2.1. Basic Characteristics of Aquifer Sediment Composite Samples 

Composite Specific Surface Area Reductively Extracted Mn 

 m2/g μmol/m2

R23AW† 0.531 ± 0.115 0.59 ± 0.03 

F168-15† 0.608 ± 0.074 3.36 ± 0.20 

F168 2005 0.764 ± 0.023 1.29 ± 0.05 

 †Data from Amirbahman et al. (2006).

saturated solution, which was filtered and 

analyzed for I2 concentration prior to use. The 

pH of the birnessite suspension was maintained 

at the target value using a pH stat containing 

0.2N HCl throughout the experiment. Samples 

were periodically withdrawn using a syringe and 

filtered through a 0.2 μm nylon filter. The first 

5 mL of filtered sample were discarded, 

retaining the next 5-10 mL for analysis. The 

samples were immediately analyzed for the three 

iodine species, iodide (I-

), iodine (I2) and iodate 

(IO3

-

) by spectrophotometer. 

In order to investigate iodate sorption by 

birnessite, a series of samples were prepared at 

pH 5.0 in 0.1M NaClO4 solution with 1 g/L 

birnessite, and iodate concentrations ranging 

from 0.005 to 0.500 mM. The batch 

experiments were performed in duplicate in 

35mL polycarbonate centrifuge tubes. The 

birnessite suspensions were adjusted to pH 5.00 

and equilibrated overnight prior to spiking with 

sodium iodate. Samples were reacted for 22 

hours before filtering and analyzing in the same 

manner as in the kinetic experiments. 

2.1.4 Cesium and Iodine Batch 

Experiments with Aquifer Sediments 

Batch experiments were performed in order to 

investigate the uptake of IO3

-

, Cs+, and I- onto 

the natural sediments from the Cape Cod sand 

and gravel aquifer. An artificial groundwater 

(AGW) with a composition designed to match 

that of the oxic zone of the aquifer was created, 

containing 0.025 mM CaSO4, 0.045 mM 

MgSO4, and 0.25 mM NaCl. In addition, 5 mM 

MES buffer or 5 mM acetate/acetic acid was 

added as a pH buffer for experiments performed 

at pH 5.6 and 4.5-5.2, respectively. 

Uptake experiments were performed in 

polycarbonate centrifuge tubes at a sediment 

concentration of 400 or 1000 g/L. Sediment 

samples were mixed with AGW and the pH was 

adjusted to the target value. Sediments were 

equilibrated for 24 hours prior to spiking with I-

, 

IO3

-

, or Cs+ to achieve a final concentration of 

1.0 mM (I) or 0.5 mM (Cs). Samples were 

allowed to react for 2 to 168 hours, then 

centrifuged and filtered through a 0.45 μm nylon 

filter and analyzed for iodine species as 

described below, or Cs by ICP-MS. 

Experiments involving sediment amended with 

synthetic birnessite at rates of 1.0 and 2.7 mg 

birnessite per g sediment were also conducted as 

described above. 

2.1.5 Column Experiment 

2.1.5.1 Column Packing 

A glass adjustable length column from Millipore 

(2.2 cm inner diameter, 1 m length) was packed 

with the F168 2005 composite sediment. The 

column was packed by suspending the sediment 

in AGW and pouring the slurry into the column. 

The excess water was displaced by tightening 

the adjustable length column plungers. The final 

packed column length was 98.74 cm at a 

suspension density of 3849 g/L. 

2.1.5.2 Transport Experiment Apparatus 

A schematic of the column and pump system is 

shown in Figure 2.1. It consisted of a stainless 

steel piston pump (ISCO, 500 mL capacity), two 

Teflon reservoirs (650 mL), a Teflon tubing 

(3/16 inch inner diameter) loop, the packed 

column, an in-line vertical flow pH electrode, 

and a fraction collector. In order to avoid any 

contact between the feed solution and the 
5

Fraction 

Collector

Column

pH 

Electrode

Pump

Teflon Reservoirs

Valves

Fraction 

Collector

Column

pH 

Electrode

Pump

Teflon Reservoirs

Valves

Figure 2.1. Experimental design for column experiment (not to scale).

stainless steel pump, the feed solutions were 

contained in the Teflon reservoirs (AGW with or 

without iodide) or the Teflon tubing loop (AGW 

with tritium), and the pump was filled with a 

sucrose solution with a density of 1.2 g/mL. The 

sucrose solution displaced the AGW in the first 

Teflon reservoir, but was never allowed to reach 

the second reservoir. All tubing and fittings in 

the system were Teflon and the valves were 4-

way Hamilton valves. 

2.1.5.3 Column Operation 

The flow rate for the duration of the column 

experiment was 3.17 mL/hr. The column was 

first preconditioned with AGW at pH 4.75 until 

the pH of the effluent water stabilized at pH 4.75 

± 0.1 (19 days). Three separate pulses of AGW 

were pumped through the column: (1) AGW + 

3

HHO at pH 4.75, (2) AGW + NaI at pH 4.75, 

and (3) AGW + NaI at pH 4.50. In between 

each pulse, AGW was pumped through the 

column until all traces of the previous pulse 

were gone from the effluent water. A 5 mM 

acetate buffer was added to the pH 4.5 AGW in 

order to help buffer the system at this lower pH. 

Approximately 2 pore volumes (302.8 mL) of 

AGW spiked with 10 nCi/mL tritiated water 

(

3

HHO) was added to the column via the tubing 

loop. The I-

 pulses (about 3.7 pore volumes 

each) were pumped through the column from the 

Teflon reservoirs. During breakthrough, effluent 

samples were collected every 0.5 or 1.0 hour. 

After reaching a plateau, samples were collected 

every 2 or 3 hours. The tritium pulse was fitted 

using CXTFIT in order to determine average 

linear velocity and pore volume of the packed 

column. 

2.1.6 Iodine and Tritium Analysis 

Iodine species were measured using a UV-Vis 

spectrophotometer on 3 separate samples: (1) 

active iodine, (2) active iodine + iodide, and (3) 

iodate. Active iodine includes elemental I2 and 

its hydrolysis products HOI, OI-

, and in the 

presence of excess iodide, I3

-

. For the sake of 

simplicity, active iodine will be referred to as 

‘iodine’ or ‘I2’ hereafter. Active iodine and 

iodide (I-

) were measured using the leuco-crystal 

violet method (APHA, 1992). Briefly, active 

iodine (1) is determined by adding a citric acid 

buffer solution (pH 3.8) to each sample or 

standard, followed by the addition of the leuco 

crystal violet indicator solution and measuring 

absorbance at 592 nm. Active iodine + iodide 

(2) is determined by the same method with the 

addition of oxone (KHSO5) to oxidize the iodide 

to iodine. Iodide is then determined by 

subtraction of (1) from (2). Iodate was 
6

measured by adding an excess of iodide and 

sulfamic acid and measuring absorbance at 351 

nm (Truesdale, 1978). The iodate concentrations 

were corrected for the presence of I3

-

. Tritium 

(

3

H) was measured by liquid scintillation 

counting on a Beckman LS 6000 counter. 

2.2 Results and Discussion 

2.2.1 Kinetics of Iodide Oxidation by 

Birnessite 

Iodide is oxidized by birnessite to both 

elemental iodine (I2) and iodate (IO3

-

). Figure 

2.2 shows speciation data over time for a typical 

experiment. All of the experiments showed this 

same general trend, with iodide disappearing 

fairly quickly followed by a sharp increase in I2

and a slower production of iodate. It is clear 

from Figure 2.2 that a significant portion of the 

iodine is associated with the mineral phase for at 

least a short period of time. The rates at which 

these reactions occur vary significantly with 

factors such as pH and ionic strength. Each of 

these factors will be discussed further below. 

The disappearance of iodide over time follows a 

pseudo first order rate law, and the slope of the 

plot ln[I-

] versus time gives the pseudo first 

order rate constant (k'). An example of this is 

shown in Figure 2.3, where only the iodide data 

above the detection limits (the first 2 hours in 

this case) was used to determine k'. The pseudo 

first order rate constants for all experiments are 

summarized in Table 2.2. 

The initial I-

 concentration was varied from 0.1 

to 1.0 mM at pH 5.00. The reaction rate did not 

vary significantly with initial I-

 concentration. 

For example, with initial I-

 concentrations of 0.1 

and 1.0 mM, k' was determined to be 1.53±0.17 

and 1.39±0.03, respectively. Because the 

disappearance of I-

 appears to be first order, it is 

expected that varying the initial I-

 concentration 

will have no effect on the rate constant. 

2.2.1.1 Effect of Ionic Strength 

Iodide oxidation by birnessite was investigated 

in 0.01 and 0.10 M NaClO4 over the pH range 

4.50 to 5.00. Increasing the ionic strength 

slowed down the oxidation reaction (Figure 2.4). 

Both the disappearance of I-

 and the production 

of I2 and IO3

-

 are slower at higher ionic strength, 

particularly in the first 2 hours, presumably due 

to the lower I-

 activity in solution. However, 

after 8 hours the effect of ionic strength is only 

apparent in the production of IO3

-

. 

2.2.1.2 Effect of pH 

Iodide oxidation by birnessite was investigated 

for the pH range of 4.50 – 6.50. The loss of Iand production of I2 and IO3

-

 occurs more 

quickly at lower pH as shown in Figure 2.5. If 

the log of k' for each experiment is plotted 

against pH (Figure 2.6), the slope of the line 

represents the reaction order with respect to H+

activity (1.02±0.04). There are two possible 

explanations for this effect. One involves the 

pH dependence of iodide oxidation. As seen in 

equation 2-1, iodide oxidation by birnessite 

consumes protons, thus it would not be 

surprising for the rate of oxidation to depend 

upon the pH. 

1⁄2MnO2 + I-

 +2H+ Æ 1⁄2Mn2+ + 1⁄2I2 + H2O 

(Eq. 2-1) 

Acid must be added throughout the experiment 

in order to maintain a constant pH. However, 

this trend may also be a reflection of pH 

dependent adsorption of iodide onto the mineral 

surface. As the pH drops, the surface charge of 

the mineral becomes less negative, and the 

electrostatic repulsion that opposes the specific 

adsorption of anions decreases. The point of 

zero charge (pzc) of synthetic birnessite is low 

(typically around 1.5) and thus the mineral is 

negatively charged throughout the entire pH 

range studied (McKenzie, 1981). Several other 

researchers have demonstrated that I-

 may sorb 

to sediments and minerals such as illite, 

ferrihydrite, hematite, and imogolite, and the 

sorption is often inversely dependent on pH 

(Couture and Seitz, 1983; Yu et al., 1996; 

Kaplan et al., 2000). 

2.2.1.3 Effect of Birnessite Concentration 

The concentration of birnessite was varied from 

0.2 to 1.0 g/L in experiments containing 0.1 mM 
7

0.00

0.20

0.40

0.60

0.80

1.00

0 10 20 30 40 50 60 70 80

I Concentration (mM)

Time (hours)

0.00

0.20

0.40

0.60

0.80

1.00

0 2 4 6 8 10

Iodine

Iodide

Iodate

Total I

I Concentration (mM)

Time (hours)

A

B

Figure 2.2. (A) Iodide oxidation by 

birnessite (1g/L) at pH 5.00 in 0.1M 

NaClO4 solution. The initial Iconcentration was 1.00 mM. (B) Plot of 

the same data for the first 8 hours of the 

experiment.

I

-

 (Figure 2.7). In all three cases the available 

exchange sites on birnessite were in excess of 

the I-

 concentration. The surface exchange 

capacity of birnessite (synthesized by the same 

method used in this study) is reported in the 

literature to be 25-33 μmol/m2

 (McKenzie, 

1981; Scott and Morgan, 1995), corresponding 

to 1.0-1.4 mmol/g for the birnessite used in this 

study. The rate of I-

 oxidation was slower at 

lower birnessite concentrations, for example 

after 8 hours, 0.010, 0.020, and 0.044 mM of the 

I

-

 had been oxidized to IO3

-

 at 0.2, 0.5 and 1.0 

g/L birnessite concentrations, respectively. The 

log of k' is plotted against the log of birnessite 

-3.00

-2.50

-2.00

-1.50

-1.00

-0.50

0.00

0.50

0 0.5 1 1.5 2 2.5

y = 0.008 - 1.39x R2= 0.998 

ln[I-]

Time (hours)

Figure 2.3 Plot of ln[I-] versus time for 1 

mM I-

 oxidation by birnessite (1 g/L) at 

pH 5.00 in 0.1M NaClO4 solution. A 

linear regression of the data gives a 

pseudo first order rate constant (k'=1.39) 

for I- disappearance (the slope), with the 

initial I- concentration as the intercept. 

concentration in Figure 2.8, with the slope of the

line representing the reaction order with respect 

to birnessite concentration (0.98 ± 0.23). It 

appears that the reaction is first order with 

respect to birnessite concentration, however 

there is a large error associated with this 

determination because only three data points 

were used. 

2.2.1.4 Elemental Iodine Oxidation by 

Birnessite 

The oxidation of I2 by birnessite was 

investigated in order to separate out the two-step 

oxidation of I-

 to I2 and I2 to IO3

-

. An example is 

shown in Figure 2.9. There appears to be an 

initial drop in I2 concentration which occurs in 

the first few minutes, which is probably due to 

adsorption. This is followed by a slower 

decrease in I2 concentration and corresponding 

increase in IO3

-

. The data from the latter 

reaction period was used to determine the 

pseudo first order rate constant (k') for I2

disappearance (Figure 2.10). Although a plot of 

ln[I2] versus time is linear, k' varied with initial 

concentration (Table 2.2 and Figure 2.11), 

indicating that the reaction may not, in fact, 

follow first order kinetics. The reaction order 
8

Table 2.2. Summary of Pseudo First Order Rate Constants Determined from the 

Disappearance of The Starting Compound (Iodide Or Iodine) 

pH NaClO4

Concentration 

Initial I 

Species 

Initial I 

Concentration

Birnessite 

Concentration

k' †

 M mM g/L hr-1

4.50 0.01 I- 1.0 1.0 

4.75 0.01 I- 1.0 1.0 

5.00 0.01 I- 1.0 1.0 

4.50 0.10 I- 1.0 1.0 

4.75 0.10 I- 1.0 1.0 

5.00 0.10 I- 1.0 1.0 

5.25 0.10 I- 1.0 1.0 

5.50 0.10 I- 1.0 1.0 

5.75 0.10 I- 1.0 1.0 

6.00 0.10 I- 1.0 1.0 

6.25 0.10 I- 1.0 1.0 

5.00 0.10 I- 0.1 1.0 

5.00 0.10 I- 0.3 1.0 

5.00 0.10 I2 0.31 1.0 

5.00 0.10 I2 0.24 1.0 

5.00 0.10 I2 0.17 1.0 

5.00 0.10 I2 0.05 1.0 

5.00 0.10 I- 0.1 0.2 

5.00 0.10 I-

 0.1 0.5 

5.30 ± 0.17 

1.79 ± 0.06 

1.47 ± 0.04 

3.25 ± 0.30 

2.45 ± 0.06 

1.39 ± 0.03 

0.627 ± 0.010 

0.298 ± 0.013 

0.184 ± 0.004 

0.118 ± 0.004 

0.064 ± 0.003 

1.53 ± 0.17 

1.57 ± 0.05 

0.052 ± 0.002 

0.072 ± 0.003 

0.089 ± 0.005 

0.172 ± 0.009 

0.305 ± 0.035 

0.552 ± 0.050 

 † k' was determined using least squared regression analysis and is expressed along with the standard 

error. 

with respect to initial I2 concentration was 

determined to be 0.79±0.07. This indicates that 

the oxidation of I2 to IO3

-

 may be more complex 

than that of I-

 to I2; perhaps it is a multistep 

reaction or is limited by the number of available 

surface sites on birnessite. 

2.2.1.5 Elemental Iodine and Iodate 

Sorption 

From Figure 2.9 we can see that when I2 is 

added to a suspension of birnessite, there is an 

initial drop in aqueous I2 concentration in the 

first few minutes, followed by a slower decrease 

that occurs after about 1 hour. We hypothesize 

here that these two regions represent fast 

adsorption followed by slower oxidation. The 

adsorption of I2 to birnessite was determined by 

averaging the data points in the ‘adsorption’ 

kinetic region. Adsorbed versus aqueous I2 is 

plotted in Figure 2.12 below. Although there is 

a fair amount of scatter in the data, there is a 

clear trend especially at lower I2 concentrations. 

Iodate sorption onto birnessite was investigated 

and found to be rapid, reaching equilibrium 

within 30 minutes, and fairly low, reaching a 

maximum of 0.024 mmol/g as shown in Figure 

2.13. It is possible that all of the I species 

adsorb to birnessite, however, other researchers 

have found IO3

-

 sorption to be greater than Isorption on several minerals and sediments 

(Couture and Seitz, 1983; Hu et al., 2005). An 

estimate of the total (sorbed + dissolved) IO3

-

and I2 can be calculated from Figures 2.12 and 

2.13. 

2.3 Rate Laws and Modeling 

Predictions 

If we consider the simple case in which 

oxidation proceeds as a two step reaction from I-
9

0.0

0.2

0.4

0.6

0.8

1.0

0 2 4 6 8 10

Iodide (I-

)

pH 4.50, 0.01M

pH 4.50, 0.1M

pH 4.75, 0.01M

pH 4.75, 0.1M

pH 5.00, 0.01M

pH 5.00, 0.1M

I Concentration (mM)

Time (hours)

0.0

0.2

0.4

0.6

0.8

1.0

0 2 4 6 8 10

Iodine (I2

)

I Concentration (mM)

Time (hours)

0.00

0.05

0.10

0.15

0.20

0 2 4 6 8 10

Iodate (IO3

-

)

I Concentration (mM)

Time (hours)

Figure 2.4. Effect of ionic strength on Ioxidation by birnessite at pH 4.5-5.0, 

showing the disappearance of I-

 and the 

production of IO3

-

 and I2. Suspensions of 

1 g/L birnessite were spiked with 1.0 mM 

NaI solution. 

0.0

0.2

0.4

0.6

0.8

1.0

0 5 10 15 20 25 30

Iodide

pH 4.50

pH 5.00

pH 5.75

pH 6.25

I Concentration (mM)

Time (hours)

0.0

0.2

0.4

0.6

0.8

1.0

0 10 20 30 40 50 60 70 80

Iodine

I Concentration (mM)

Time (hours)

0.0

0.1

0.2

0.3

0.4

0.5

0 10 20 30 40 50 60 70 80

Iodate

I Concentration (mM)

Time (hours)

Figure 2.5. Effect of pH on iodide 

oxidation by birnessite in 0.1M NaClO4

solution, showing the disappearance of Iand production of IO3

-

 and I2. 

Suspensions of 1 g/L birnessite were 

spiked with 1.0 mM NaI solution. 
10

-1.50

-1.00

-0.50

0.00

0.50

1.00

4.0 4.5 5.0 5.5 6.0 6.5

y = 5.18 - 1.02x R2= 0.991 

log[k']

pH

Figure 2.6. pH dependence of the log of 

the first order rate constant, k', 

describing the disappearance of iodide 

from solution, determined from reacting 1 

mM I-

 with 1 g/L birnessite in 0.1M 

NaClO4 solution. The slope of the line 

represents the reaction order with respect 

to hydrogen ion activity. 

to I2 and then I2 to IO3

-

 (and both reactions are 

first order with respect to I), then the rate laws 

can be written as follows: 

d[I-

]/dt = -k1[I-] (Eq. 2-2) 

d[I2]/dt = k1[I-

] – k2[I2] (Eq. 2-3) 

d[IO3

-

]/dt = k2[I2], (Eq. 2-4) 

where k1 and k2 are the pseudo-first order rate 

constants for the oxidation of I-

 to I2 and I2 to 

IO3

-

 by birnessite, respectively. These equations 

can be solved for each I species: 

[I-

] = [I-]0e

-k1t (Eq. 2-5) 

[I2] = [I-

]0(k1/(k2-k1))( e-k1t - e-k2t) (Eq. 2-6) 

[IO3

-

] = [I-]0(1+((k1e

-k2t – k2e-k1t)/(k2 – k1))), 

(Eq. 2-7) 

where [I-

]0 is the initial iodide concentration. 

Under a given set of conditions, e.g. pH 5.00, 1 

g/L birnessite in 0.1M NaClO4 solution, we can 

determine k1 from the disappearance of iodide to 

be 1.5 (average of 3 values). As discussed 

above, in a given experiment the disappearance 

of I2 appears to follow first order kinetics 

0.00

0.02

0.04

0.06

0.08

0.10

0 2 4 6 8 10

Iodide

0.2g/L

0.5g/L

1.0g/L

I Concentration (mM)

Time (hours)

0.00

0.02

0.04

0.06

0.08

0.10

0 2 4 6 8 10

Iodine

I Concentration (mM)

Time (hours)

0.00

0.02

0.04

0.06

0.08

0.10

0 2 4 6 8 10

Iodate

I Concentration (mM)

Time (hours)

Figure 2.7. Oxidation of 0.1 mM I- at pH 

5.00 in 0.1M NaClO4 solution for various 

concentrations of Birnessite, showing the 

disappearance of I- and the production of 

I2 and IO3

-

 versus time. 
11

-1.50

-1.00

-0.50

0.00

0.50

-2 -1.5 -1 -0.5 0 0.5

y = 0.303 + 0.983x R2= 0.947 

ln[k']

ln [solid conc (g/L)]

Figure 2.8. Log-log plot of the pseudofirst order rate constant for iodide 

disappearance versus birnessite 

concentration for experiments performed 

with an initial I-

 concentration of 0.1 mM 

in 0.1M NaClO4 solution at pH 5.0. The 

slope of the line represents the reaction 

order with respect to birnessite 

concentration. 

 

(Figure 2.10). However, the pseudo-first order 

rate constant (k2) determined from these 

experiments ranges from 0.052 to 0.173 for 

initial I2 concentrations of 0.31 to 0.05 mM 

indicating that this reaction is complicated by a 

multi-step reaction or is limited by the available 

surface sites (Table 2.2 and Figure 2.11). As an 

approximation, we can treat the oxidation of I2

to IO3

-

 as a first order reaction by determining k2

for each individual reaction. Over longer time 

periods (after all of the I-

 has been oxidized to 

I2), k2 can be determined simply from the 

disappearance of I2. Equations 2-5, 2-6, and 2-7 

can then be used to predict the concentrations of 

all three I species over time as shown in Figures 

2.14 through 2.16. The model overpredicts the 

concentration of I2 during the beginning of the 

experiment, but predicts the overall shape of the 

data fairly well. The overprediction of I2

concentrations may be due to an inaccurate 

estimate of the amount of sorbed I2 during this 

time period or may reflect the presence of 

sorbed I-

. If there is significant sorbed I- present, 

then the reaction consists of fast I-

 sorption 

followed by a slower oxidation to I2. 

0.00

0.05

0.10

0.15

0.20

0.25

0.30

0 10 20 30 40 50 60 70 80

Iodine

Iodide

Iodate

Total I

I Concentration (mM)

Time (hours)

0.00

0.05

0.10

0.15

0.20

0.25

0.30

0 2 4 6 8 10

I Concentration (mM)

Time (hours)

A

B

Figure 2.9. (A) I2 oxidation by birnessite 

(1 g/L) at pH 5.00 in 0.1M NaClO4

solution. The initial iodine concentration 

was 0.24 mM (as I). (B) Plot of the same 

data for the first 8 hours of the 

experiment. 

2.4 Cesium and Iodine Reactions 

with Aquifer Sediments 

2.4.1 Iodide Adsorption and Oxidation 

All sediments were very unreactive with iodide 

under the conditions investigated (Table 2.3, 

Figure 2.17). Only under the most favorable 

conditions for I-

 oxidation, namely lower pH and 

higher Mn oxide content of F168-15, was there 

any appreciable uptake of I-

. In experiments with 

unamended sediments, oxidized I (IO3

-

 or I2) 
12

-6.00

-5.00

-4.00

-3.00

-2.00

-1.00

0 10 20 30 40 50 60

y = -1.81 - 0.072x R2= 0.993 

ln[I2

]

Time (hours)

Figure 2.10. Plot of ln[I2] over time for 

0.24 mM I2 (as I) oxidation by birnessite 

(1 g/L) at pH 5.00 in 0.1M NaClO4

solution. A linear regression of the data 

gives the pseudo first order rate constant 

(k'=0.072) for I2 disappearance (the 

slope), with the initial I2 concentration as 

the intercept. 

-3.00

-2.50

-2.00

-1.50

-1.00

-0.50

-4.50 -4.00 -3.50 -3.00 -2.50 -2.00 -1.50 -1.00

y = -3.847 - 0.787x R2= 0.982 

ln [k']

ln[I2]

initial

Figure 2.11. Log of k' for the 

disappearance of I2 from solution 

determined at 1 g/L birnessite in 0.1M 

NaClO4 solution at pH 5.0, plotted against 

the log of the initial I2 concentration. The 

slope of the line represents the reaction 

order with respect to the initial I2

concentration. 

0.00

0.05

0.10

0.15

0.20

0.25

0 0.1 0.2 0.3 0.4 0.5

[I2

]ads (mmol/g)

[I2]

aq (mM)

Figure 2.12. Iodine sorption by birnessite 

(1g/L) at pH 5.0 in 0.1M NaClO4 solution. 

Several data points in the first 10-30 

minutes of I2 reaction with birnessite were 

averaged (i.e. before the onset of 

oxidation to IO3

-

). Error bars represent 

the standard deviation for the averaged 

values. 

was below detection limits in most cases. 

However, the addition of a Mn oxide (birnessite, 

δ-MnO2) to the sediments significantly increased 

the loss of iodide from solution. It is interesting 

to note that less than 2% of the total I added to 

experiments with 1.0 mg/g MnO2 was detected, 

however at higher amendment levels of MnO2

(2.7 mg/g), close to 4 and 5% of the total I was 

detected as iodate and elemental iodine, 

respectively. These experiments were not 

performed in a gas-tight system, so I2 may be 

lost to volatilization or taken up by the solid 

phase, as was observed in experiments with pure 

birnessite. 

2.4.2 Iodate Adsorption 

As can be seen in Figures 2.18 and 2.19, 

sediment sample R23AW was fairly unreactive 

with iodate, with only 2% adsorption at pH 5.7 

(initial IO3

-

 concentration = 1.0 mM). However, 

a relatively slow adsorption of iodate onto F168 

sediments occurred, reaching 10-12% for F168-

15 sediments after 1 week and 9.2% for F168 
13

0.00

0.10

0.20

0.30

0.40

0.50

0 2 4 6 8 10

I Concentration (mM)

Time (hours)

0.00

0.01

0.01

0.02

0.02

0.03

0 0.1 0.2 0.3 0.4 0.5

[IO3

-

]ads (mmol/g)

[IO3

-

]

aq (mM)

A

B

Figure 2.13. Iodate sorption by birnessite 

(1 g/L) at pH 5.0 in 0.1M NaClO4

solution. (A) Kinetics of IO3

-

 sorption for 

sample spiked with 0.50 mM IO3

-

. (B) 

Sorption isotherm showing adsorbed IO3

-

versus IO3

-

 concentration in solution. 

2005 sediments after 72 hours, perhaps due to 

the lower pH of these samples. Samples were 

also analyzed for other iodine species, but none 

were detected indicating that reduction of iodate 

was not important in this system. Some other 

researchers have noted that uptake of IO3

-

 by 

soils and minerals was due to reduction of IO3

-

to I2, however, these studies involved Fe(II) 

minerals (pyrite) or anaerobic conditions 

(Fuhrmann et al., 1998; Yamaguchi et al., 2006). 

0.00

0.20

0.40

0.60

0.80

1.00

0 2 4 6 8 10

A

I

-

aq

I

2, total

IO3

-

,total

I Concentration (mM)

Time (hours)

0.00

0.20

0.40

0.60

0.80

1.00

0 10 20 30 40 50 60 70 80

B

I Concentration (mM)

Time (hours)

Figure 2.14. Comparison of actual (solid 

lines) and predicted (dashed lines) iodide 

oxidation by birnessite (1 g/L) using 

k1=1.5 and k2=0.0042 for (A) the first 10 

hours, and (B) 72 hours. [I-

]0=1.0 mM, 

pH=5.0, and [NaClO4]=0.1M. Actual IO3

-

and I2 are plotted as total (dissolved + 

adsorbed) concentration, and I-

 as 

dissolved only. 

2.4.3 Cesium Adsorption 

Unsurprisingly, Cs adsorption to aquifer 

sediments was much more extensive than I 

adsorption. After approximately 24 hours, 22% 

and 12% of the Cs was sorbed to sediment 

samples R23AW and F168 2005, respectively 

(Figure 2.20). Most of the Cs uptake by the 

sediments appears to occur quickly (within a few 

hours), although slight increases in Cs uptake 
14

0.00

0.05

0.10

0.15

0.20

0.25

0.30

0 2 4 6 8 10

A

I Concentration (mM)

Time (hours)

0.00

0.05

0.10

0.15

0.20

0.25

0.30

0 10 20 30 40 50 60 70 80

B

I

-

aq

I

2, total

IO3

-

,total

I Concentration (mM)

Time (hours)

Figure 2.15. Comparison of actual (solid 

lines) and predicted (dashed lines) iodide 

oxidation by birnessite (1 g/L) using 

k1=1.5 and k2=0.0473 for (A) the first 10 

hours, and (B) 72 hours. [I-

]0 = 0.3 mM, 

pH=5.0, and [NaClO4]=0.1M. Actual IO3

-

and I2 are plotted as total (dissolved + 

adsorbed) concentration, and I-

 as 

dissolved only. 

may continue over longer periods of several 

days. Other researchers have also observed 

strong Cs sorption by sediments (Zachara et al., 

2002; Ainsworth et al., 2005). Poinssot et al. 

(1999) noted that Cs sorption to the low affinity, 

planar sites on illite occurred within a few hours, 

while sorption onto the high affinity frayed-edge 

sites took up to 5 days to reach equilibrium. At 

the high Cs concentrations used in this study, 

adsorption should be dominated by the low0.00

0.02

0.04

0.06

0.08

0.10

0 2 4 6 8 10

I

-

aq

I

2, total

IO3

-

, total

I Concentration (mM)

Time (hours)

Figure 2.16. Comparison of actual (solid 

lines) and predicted (dashed lines) iodide 

oxidation by birnessite (1 g/L) using 

k1=1.5 and k2=0.1408. [I-

]0=0.1 mM, 

pH=5.0, and [NaClO4]=0.1M solution. 

Actual IO3

-

 and I2 are plotted as total 

(dissolved + adsorbed) concentration, and 

I

-

 as dissolved only. 

affinity planar sites. The presence of I- had no 

effect on Cs adsorption (data not shown). 

2.4.4 Iodide Transport in Column 

Experiments with Aquifer Sediments 

Figure 2.21 shows breakthrough curves for 

tritium and iodide in a column packed with the 

F168 2005 composite sediment sample. The 

tritium curve was fitted using CXTFIT, and the 

pore volume was found to be 140.6 mL; the 

average linear velocity was found to be 0.534 

m/day. When iodide was added to the column at 

pH 4.75, the concentrations increased during 

breakthrough to about 97% of the input 

concentration, and then slowly increased to 

100% over the next couple of days. There was 

no indication of iodide retardation during 

transport, with the breakthrough curve during 

the first pore volume being virtually identical to 

that observed for tritium. In the experiment at 

pH 4.50, however, iodide concentrations 

increased rapidly to 90% of the inlet 

concentration, and then reached a maximum of 

93% before declining as the pulse of introduced 

iodine passed. Because no elemental I2 or iodate 
15

Table 2.3. Reaction Conditions and Speciation of Iodine after Reaction of 1.0 mM Iodide

with Aquifer Sediments 

Sediment Amendment Solid 

conc 

Reaction 

Time 

Final 

pH‡ 

IO3

- I2 I-

 mg/g g/L hr —————mM—————

R23AW None 400 168 5.7 BDL† BDL 0.993 

R23AW None 400 168 5.1 BDL BDL 0.976 

F168-15 None 400 168 5.1 0.002 0.003 0.979 

F168-15 None 400 168 5.0 0.003 0.003 0.927 

F168 2005 None 1000 48 4.6 BDL BDL 0.940 

F168 2005 None 1000 24 4.9 BDL BDL 0.967 

F168 2005 1.0 (MnO2) 1000 48 5.0 0.002 BDL 0.017 

F168 2005 2.7 (MnO2) 1000 48 4.9 0.039 0.049 BDL 

†BDL= below detection limits of 0.001 mM (IO3

-

) and 0.002 mM (I2). 

‡ The pH drifted up from the initial values by up to 0.5 pH units. The pH value at the end of 

the experiment is shown. 

0.0

0.5

1.0

1.5

2.0

2.5

0 50 100 150

F168 2005, pH4.9

% Iodide Uptake

Time (hours)

Figure 2.17. Time dependence of iodide 

loss from solution via sorption or 

oxidation by sediments (1000g/L). The 

initial I concentration was 1.0 mM. 

was detected in the effluent, the results indicate 

a loss of iodine mass during the experiment. 

The most likely explanation for the loss is 

oxidation of iodide to elemental I2 with 

subsequent volatilization from samples open to 

the atmosphere in the effluent fraction collector. 

0.0

2.0

4.0

6.0

8.0

10.0

0 50 100 150 200

F168 2005, pH4.9

R23AW, pH5.7

% Iodate Sorbed

Time (hours)

Figure 2.18. Time dependence of iodate 

adsorption onto sediments (400g/L). The 

initial iodate concentration was 1.0 mM.

2.5 Summary of Laboratory 

Studies 

Birnessite, a Mn-oxide mineral commonly 

present in soils and sediments, was shown to 

oxidize I-

 to I2 and IO3

-

 in a two-step process. 

The oxidation of I-

 proceeded according to first 

order kinetics with respect to initial Iconcentration, pH, and birnessite concentration. 

While I2 sorption to birnessite was high (up to 

0.25 mmol/g), IO3

-

 sorption on birnessite was an 

order of magnitude lower (up to 0.024 mmol/g). 

Uptake of I-

 by sediments was near zero in batch 
16

0

2

4

6

8

10

12

R23AW pH 5.7 F168-15 pH 5.1 F168-15 pH 4.9

% IO3

- sorbed

Figure 2.19. Iodate adsorption onto 

sediments (400 g/L) at different pH values 

after 1 week of reaction. The initial 

iodate concentration was 1.0 mM. 

0

5

10

15

20

25

30

0 20 40 60 80 100

R23AW, pH5.5

F168 2005, pH4.8

% Cs sorbed

Time (hours)

Figure 2.20. Time dependence of Cs 

adsorption onto aquifer sediments (400 

g/L). The initial Cs concentration was 0.5 

mM. 

(<2.5%) and column experiments at pH 4.8 or 

above, but at pH 4.50, uptake of I-

 in column 

experiments was significantly higher (Iconcentrations reached a maximum of 93% of 

the input concentration). IO3

-

 uptake by 

sediments was higher than that of I-

, reaching 

close to 12% for the F168-15 sediment sample at 

pH 4.9 (initial I concentration of 1.0 mM). Cs 

uptake by sediments was higher than that of I, 

with 22 and 12% sorbed to the R23AW and 

F168 2005 sediment samples after 24 hours 

(initial Cs concentration of 0.5 mM). 

0.00

0.20

0.40

0.60

0.80

1.00

0.50 1.0 1.5 2.0 2.5 3.0 3.5 4.0

Tritium

C/C0

Pore Volume

0.00

0.20

0.40

0.60

0.80

1.00

0123456

1.0mM I-

, pH 4.75

C/C0

Pore Volume

0.00

0.20

0.40

0.60

0.80

1.00

0123456

1.0mM I-

, pH 4.50

C/C0

Pore Volume

Figure 2.21. Breakthrough curves for 

tritium (3

H) and iodide at pH 4.75 and 

4.50. Concentrations are expressed 

relative to the input concentrations. 
17

While the Cs and I concentrations used in these 

experiments are much higher than would be 

relevant for concentrations of radioactive 

isotopes of these elements, the studies are 

relevant for revealing reaction mechanisms that 

affect the transport of the radionuclides in the 

environment. 
18
19

3 Field Studies of Cesium and Iodine Transport under Variable 

Environmental Conditions 

3.1 Site Description 

Field experiments were performed in a shallow, 

unconfined sand and gravel aquifer located in 

western Cape Cod, MA. Discharge of effluent 

from a sewage treatment plant operated by the 

Massachusetts Military Reservation from 1938 

to 1995 resulted in a plume of sewagecontaminated groundwater characterized by high 

concentrations of dissolved salts, no dissolved 

O2, higher pH, and the presence of dissolved 

Fe(II) (Kent et al., 1994). The Fe-reducing 

sewage contaminated zone is overlain by 

pristine groundwater characterized by low 

dissolved salts, high dissolved O2, and lower pH. 

Ground-water flow velocities are approximately 

0.4 m/day. Further details on the site history and 

characteristics are described elsewhere 

(LeBlanc, 1984; LeBlanc et al., 1991; Repert et 

al., 2006). 

Two regions of the aquifer were chosen for the 

field experiments, one in the Fe-reducing zone 

of the aquifer resulting from sewage 

contamination (located near well F625, see Fig. 

3.1), and one in the pristine, oxic zone of the 

aquifer (located near well F168). In each area 

an array of multilevel samplers (MLS) exists 

that allows collection of breakthrough curves at 

several distances downgradient from an injection 

well. A summary of the background water 

chemistry at the injection wells used for the 

three field experiments are shown in Table 3.1. 

3.2 Field Experiment Methodology 

Three tracer tests were performed at the field 

site in August 2006: (1) injection of CsI and 

NaBr into the oxic zone of the aquifer at F168-

M17-02 (Oxic I-

), (2) injection of CsBr and 

NaIO3 into the oxic zone at F168-M17-08 (Oxic 

IO3

-

), and (3) injection of CsBr and NaIO3 into 

the Fe-reducing zone of the aquifer at F625-

M02-09 (Anoxic IO3

-

). A summary of the 

injections is shown in Table 3.2. For all three 

tracer tests, groundwater was collected from 

wells in the same area (within 5-10 m) and depth 

interval as the injection wells. Tracer salts were 

added to the groundwater and mixed well before 

injecting. In order to prevent oxygen from 

entering the test performed in the Fe-reducing 

zone, the injection water was collected in a N2

purged jet fuel bladder that was impervious to 

oxygen. The tracers were dissolved in a small 

amount of deionized water (boiled to remove 

oxygen) and pumped into the bladder, followed 

by several rinses with boiled, deionized water. 

However, some oxygen was inadvertently 

allowed into the bladder during the injection, 

resulting in low concentrations of dissolved 

oxygen in the injectate. 

Groundwater samples were collected from 

downgradient MLS at various times and 

distances in order to create breakthrough curves. 

Using a peristaltic pump, each MLS port was 

purged with 150 mL of groundwater before 

collecting water samples for in-field (pH, 

conductivity) and laboratory analyses. Lab 

samples were filtered through a 0.45 μm filter. 

Samples collected for iodine analysis were kept 

in the dark and refrigerated until they could be 

analyzed (within 4 hours) by spectrophotometer. 

Samples collected for Br and other anions were 

frozen. Br was measured using either ICP-MS 

(for samples containing I-

 or I2) or flow injection 

analysis (for samples containing only IO3

-

). A 

subset of Br samples was analyzed by both 

methods in order to verify agreement between 

the two methods. Selected samples were also 

analyzed by ion chromatography for nitrate, 

phosphate, chloride, sulfate, and ammonium. 

Samples collected for elemental analysis by 

ICP-AES (Na, K, Mg, Ca, Mn, Fe, Si, P) and 

ICP-MS (Cs, Br) were acidified to pH 2 with 

trace metal grade nitric acid. 

Iodine species were measured using a UV-Vis 

spectrophotometer on three separate samples as 

described in Section 2.1.6. 
20

Figure 3.1. Map of the Cape Cod field site showing the extent of the wastewater plume and 

the locations of the two study areas (wells F625 and F168).

Table 3.1. Background Chemistry at Injection Wells Used In Field Experiments 

Injection DO Ca K Mg Na Mn Fe 

Well mg/L —————————————mol/L————————————— 

F168-

M17-02 

4.34 9.22x10-06 1.31x10-05 2.94x10-05 2.43x10-04 5.32x10-07 bdl 

F168-

M17-08 

7.45 1.70x10-05 1.52x10-05 3.71x10-05 2.32x10-04 2.63x10-07 bdl 

F625-

M02-09 

bdl 1.98x10-04 8.45x10-05 5.91x10-05 8.99x10-04 2.59x10-06 2.07x10-04

pH Alk P Si Cl- SO4

2- NH4+ 

 meq/L ———————————mol/L——————————— 

F168-

M17-02 

4.71 bdl bdl 1.12x10-04 2.54x10-04 2.12x10-04 bdl 

F168-

M17-08 

5.06 0.005 bdl 1.36x10-04 2.04x10-04 2.13x10-04 bdl 

F625-

M02-09 

6.62 0.953 9.67x10-05 2.25x10-04 4.94x10-04 1.44x10-04 5.83x10-06
21

Table 3.2. Tracer Test Characteristics and Injected Tracer Concentrations 

 Oxic I- Oxic IO3

- Anoxic IO3- 

Injection Port F168 M17-02 F168 M17-08 F625 M02-09 

Injection Altitude, 

m to sea level 

8.21 6.75 1.02 

Volume, L 362 383 179 

Cs, mM 1.07 1.06 1.03 

I, mM 1.13 1.05 0.98 

Br, mM 1.13 1.02 1.01 

Na, mM 1.26 1.26 2.50 

pH 4.70 5.18 6.53 

3.3 Results and Discussion 

3.3.1 Iodine Oxidation in the Oxic Iodide 

Tracer Test 

Results from the oxic I- injection demonstrated 

that iodide was oxidized to both I2 and IO3

-

during the course of the tracer test (Figure 3.2). 

Very little oxidation occurred over the first 1 m 

of transport, but after 2 m more than 20% of the 

injected I-

 had been oxidized, primarily to I2, and 

after 4 m even more oxidation had occurred. 

This region of the aquifer was characterized by 

oxygenated groundwater (4.0-5.8 mg/L 

dissolved O2) and low pH (4.5-5.1), creating 

conditions conducive for the oxidation of iodide. 

There are several possible oxidants that may be 

responsible for iodide oxidation, including 

dissolved O2, NO3

-

, and mineral phases such as 

Mn and Fe oxides (Fig. 1.1). Mn oxides are 

clearly the most likely oxidant, however, 

because oxygen is slow to react with iodide and 

nitrate concentrations were extremely low (<1 

μM). 

Manganese oxides have been shown to 

abiotically oxidize a variety of inorganic 

contaminants such as As(III) and Cr(III) and are 

likely oxidants of I-

 as well (Fendorf and 

Zasoski, 1992; Scott and Morgan, 1995; 

Anschutz et al., 2000; Truesdale et al., 2001; 

Manning et al., 2002). Amirbahman et al. 

(2006) found that sediments collected from the 

same tracer test region oxidized As(III) to As(V) 

in laboratory batch experiments, most likely 

caused by Mn-oxides. Sediment samples 

collected closest to the tracer test location 

(F168-15) contained 2.04 μmol/g reductively 

extractable Mn (Amirbahman et al., 2006). If 

we consider an estimated bulk density of 4000 

g/L in the field, then there was an excess of Mnoxide available for I-

 oxidation (8.16 mM Mn). 

As shown in Figure 3.2, there was an increase in 

the concentration of dissolved Mn during this 

tracer test, arriving just after the Br 

breakthrough curve, thus providing evidence 

that Mn oxides were the likely electron acceptor 

responsible for I-

 oxidation. The observed 

dissolved Mn concentrations increased with 

transport distance from the injection well, 

reaching up to 7.6 x 10-5 M at 2 m 

downgradient. Mn2+ is known to adsorb to 

mineral surfaces, thus the observed dissolved 

Mn should be only a fraction of the total amount 

of Mn2+ produced during the reaction of Mn 

oxides with I-

. An increase in Mn 

concentrations was not observed in the oxic 

tracer test in which IO3

-

 was injected, and thus 

the increase in Mn concentrations cannot be 

attributed to ion exchange reactions releasing 

sorbed Mn2+. The mass balance for I indicates 

that little methyl-I was formed in the 

experiments. 

3.3.2 Iodate Adsorption in the Oxic 

Iodate Tracer Test 

Iodate was significantly retarded relative to the 

Br tracer (Figure 3.3) when iodate was injected 

into an oxic region of the aquifer. After 4 m of 

transport, the IO3

-

 pulse became very spread out, 

and reached a maximum of 29% of the injection 

concentration. The retardation and significant 

tailing indicates that iodate adsorbs to the 
22

0.00

0.20

0.40

0.60

0.80

1.00

0

1 10-5

2 10-5

3 10-5

4 10-5

5 10-5

6 10-5

7 10-5

8 10-5

0 50 100 150 200 250 300

0.97m

Normalized Concentration (C/C0

)

Mn Concentration (M)

Time (hours)

0.00

0.20

0.40

0.60

0.80

1.00

0

1 10-5

2 10-5

3 10-5

4 10-5

5 10-5

6 10-5

7 10-5

8 10-5

50 100 150 200 250 300 350

2.02m

Normalized Concentration (C/C0

)

Mn Concentration (M)

Time (hours)

0.00

0.20

0.40

0.60

0.80

1.00

0

1 10-5

2 10-5

3 10-5

4 10-5

5 10-5

6 10-5

7 10-5

8 10-5

100 150 200 250 300 350 400

3.98m

Normalized Concentration (C/C0

)

Mn Concentration (M)

Time (hours)

0.00

0.20

0.40

0.60

0.80

1.00

0

1 10-5

2 10-5

3 10-5

4 10-5

5 10-5

6 10-5

7 10-5

8 10-5

0 50 100 150 200 250 300

Injection Well

I

2

I

-

IO3

-

Br

Mn

Normalized Concentration (C/C0

)

Mn Concentration (M)

Time (hours)

Figure 3.2. Breakthrough curves showing Br-, I-, I2, IO3

-

, and Mn concentrations at various 

distances downgradient of the injection well in the oxic I-

 tracer test. All Br and I species 

concentrations are normalized to the injection concentration. 

sediments under these conditions. Other 

researchers have observed that iodate exhibits 

higher sorption and is more retarded than iodide 

(Couture and Seitz, 1983; Hu et al., 2005). Hu 

et al. (2005) observed both iodate sorption and 

reduction to I-

 during batch and column studies 

with a variety of soils and clay minerals. They 

postulated that structural Fe(II) present in the 

mineral phases was responsible for the iodate 

reduction. In our study, however, no iodide was 

detected during the tracer test. Nevertheless, 

iodate exhibited non-conservative transport in 

this experiment. 

3.3.3 Iodate Reduction in the Anoxic 

Iodate Tracer Test 

Iodate was injected into an iron-reducing region 

of the aquifer located near well F625. The 

iodate was quickly reduced to iodide, with more 

than 50% reduced after only 1 m of transport 

(Figure 3.4). Transport of iodine was 

conservative, matching Br transport very 

closely. No I2 was detected during the tracer test, 

suggesting that iodate was reduced directly to 

iodide. This is in contrast to the oxidation of 

iodide, which proceeded through I2 to IO3

-

. It is 

unclear by what mechanism iodate reduction 

occurs in this system. A number of inorganic 

compounds have been shown or proposed to be 
23

0.00

0.20

0.40

0.60

0.80

1.00

1.20

0 100 200 300 400

Injection Well

IO3

- BrNormalized Concentration (C/C0

)

Time (hours)

0.00

0.20

0.40

0.60

0.80

1.00

1.20

0 100 200 300 400

0.97m

Normalized Concentration (C/C0

)

Time (hours)

0.00

0.20

0.40

0.60

0.80

1.00

1.20

100 200 300 400 500

2.02m

Normalized Concentration (C/C0

)

Time (hours)

0.00

0.20

0.40

0.60

0.80

1.00

1.20

200 300 400 500 600 700 800 900 1000

3.98m

Normalized Concentration (C/C0

)

Time (hours)

Figure 3.3. Iodate and Br breakthrough curves at various distances downgradient from the 

injection well in the oxic iodate tracer test. All concentrations are normalized to the 

injection concentration. Note the different time scales for different wells.

reductants of iodate, including Fe2+ (both 

dissolved and solid phases), dissolved Mn2+, and 

dissolved S2- (Councell et al., 1997; Anschutz et 

al., 2000; Hu et al., 2005). Microbially mediated 

reduction by Desulfovibrio desulfuricans and 

Shewanella putrefaciens has been demonstrated 

(Councell et al., 1997; Farrenkopf et al., 1997). 

During this tracer test, there was a decrease in 

the concentrations of dissolved Fe that coincided 

with the I pulse, however, this may have been 

due to the fact that some dissolved oxygen 

contaminated the injectate solution, possibly 

causing oxidation of Fe(II) and subsequent 

precipitation of Fe(III). 

3.3.4 Cesium Retardation and 

Attenuation 

Cesium was significantly retarded relative to the 

conservative tracer (Br) in both tracer tests 

performed in the oxic zone of the aquifer near 

F168 (Figures 3.5 and 3.6). Cs is so attenuated 

that after only 4 m, the maximum Cs 

concentration reached was only 5% of the 

injected concentration. A pulse of dissolved K 

was liberated from the aquifer sediments during 

the tracer tests, an indicator that Cs uptake in the 

tracer tests was due to cation exchange. Cation 

exchange processes are generally considered to 

be independent of pH, and while this is generally 

true for Cs over a wide range of conditions, 
24

0.00

0.20

0.40

0.60

0.80

1.00

0 50 100 150 200 250

Injection Well

IO3

-

I

-

BrNormalized Concentration (C/C0

)

Time (hours)

0.00

0.20

0.40

0.60

0.80

1.00

0 50 100 150 200 250

0.99m

Normalized Concentration (C/C0

)

Time (hours)

0.00

0.20

0.40

0.60

0.80

1.00

100 150 200 250 300 350

2.96m

Normalized Concentration (C/C0

)

Time (hours)

Figure 3.4. Br-, IO3

-

 and I- breakthrough curves at various distances downgradient from 

the injection well for the anoxic iodate injection. Concentrations are normalized to the 

injection concentrations.

Poinssot et al. (1999) found Cs sorption to 

exhibit pH dependence (sorption increased with 

increasing pH for the pH range 3-6) at low ionic 

strength (in 0.01M NaClO4 solution). Because 

the oxic zone of the Cape Cod aquifer has a low 

ionic strength (Table 3.2), it is possible that the 

pH difference between the two tracer tests (5.18 

at M17-08 vs 4.70 at M17-02) affected Cs 

sorption. Reactive transport modeling of the 

data is necessary to determine if there was a real 

difference in Cs sorption between the two oxic 

tracer tests. The disappearance of Cs from the 

injection well during the anoxic tracer test is 

shown in Figure 3.7. Cs and Br disappearance 

appeared to be qualitatively similar in this tracer 

test to the behavior observed in the oxic tests 

(Figure 3.6). 

3.4 Summary of Field Experiments 

This study demonstrated that both iodine and 

cesium may behave non-conservatively in 

groundwater systems. In the oxic zone of the 

aquifer, iodide was oxidized to I2 and IO3

-

, and 

IO3

-

 transport was retarded. Iodide was most 

likely oxidized by Mn oxides present in the 

aquifer sediments. In the iron-reducing zone of 

the Cape Cod sand and gravel aquifer, IO3

-

 was 

reduced to I-

. Cs was even more retarded than 

IO3

-

 in both the oxic and reducing zones of the 

aquifer, due to cation exchange with K. These 

results demonstrate the importance of taking 

redox transformations into consideration when 

predicting iodine transport in the field. 
25

0.00

0.20

0.40

0.60

0.80

1.00

0

2 10-5

4 10-5

6 10-5

8 10-5

1 10-4

0 50 100 150 200 250 300 350

Injection Well

Br Cs

K

Normalized Concentration (C/C0

)

K Concentration (M)

Time (hours)

0.00

0.20

0.40

0.60

0.80

1.00

0

2 10-5

4 10-5

6 10-5

8 10-5

1 10-4

0 50 100 150 200 250 300 350

0.97m

Normalized Concentration (C/C0

)

K Concentration (M)

Time (hours)

0.00

0.20

0.40

0.60

0.80

1.00

0

2 10-5

4 10-5

6 10-5

8 10-5

1 10-4

50 100 150 200 250 300 350 400

2.02m

Normalized Concentration (C/C0

)

K Concentration (M)

Time (hours)

0.00

0.20

0.40

0.60

0.80

1.00

0

2 10-5

4 10-5

6 10-5

8 10-5

1 10-4

200 400 600 800 1000

3.98m

Normalized Concentration (C/C0

)

K Concentration (M)

Time (hours)

Figure 3.5. Cs, Br, and K breakthrough curves at various distances downgradient from the 

injection well in the oxic iodide tracer test (well F168-M17-02). Cs and Br concentrations 

are normalized to the concentrations in the injectate. Note the different time scales in each 

panel. 
26

0.00

0.20

0.40

0.60

0.80

1.00

1.20

0

5 10-5

1 10-4

1.5 10-4

2 10-4

0 100 200 300 400 500

Injection Well

Br Cs

K

Normalized Concentration (C/C0

)

K Concentration (M)

Time (hours)

0.00

0.20

0.40

0.60

0.80

1.00

1.20

0

5 10-5

1 10-4

1.5 10-4

2 10-4

0 100 200 300 400 500

0.97m

Normalized Concentration (C/C0

)

K Concentration (M)

Time (hours)

0.00

0.20

0.40

0.60

0.80

1.00

1.20

0

5 10-5

1 10-4

1.5 10-4

2 10-4

200 400 600 800 1000

2.02m

Normalized Concentration (C/C0

)

K Concentration (M)

Time (hours)

0.00

0.20

0.40

0.60

0.80

1.00

1.20

0

5 10-5

1 10-4

1.5 10-4

2 10-4

0 500 1000 1500 2000 2500

3.98m

Normalized Concentration (C/C0

)

K Concentration (M)

Time (hours)

Figure 3.6. Cs, Br, and K breakthrough curves at various distances downgradient from the 

injection well in oxic iodate tracer test (well F168-M17-08). Cs and Br concentrations are 

normalized to the concentrations in the injectate. Note the different time scales in each 

panel. 
27

0.00

0.20

0.40

0.60

0.80

1.00

0

2 10-5

4 10-5

6 10-5

8 10-5

1 10-4

1.2 10-4

1.4 10-4

1.6 10-4

0 100 200 300 400 500

Injection Well

Br Cs

K

Normalized Concentration (C/C0

)

K Concentration (M)

Time (hours)

0.00

0.20

0.40

0.60

0.80

1.00

0

2 10-5

4 10-5

6 10-5

8 10-5

1 10-4

1.2 10-4

1.4 10-4

1.6 10-4

0 500 1000 1500 2000

0.99m

Normalized Concentration (C/C0

)

K Concentration (M)

Time (hours)

Figure 3.7. Cs, Br, and K during the anoxic iodate tracer test (well F625-M12-09). No Cs 

was detected beyond 0.99 m downgradient. Cs and Br concentrations are normalized to 

the concentrations in the injectate. 
28
29

4 SUMMARY 

Radioactive isotopes of Cs (137Cs) and I (131I and 

129I) are produced during nuclear fission and 

have been released into the environment through 

nuclear fallout from weapons tests and nuclear 

accidents such as Chernobyl. Scientists and 

regulators need information on the geochemical 

behavior of Cs and I in order to accurately 

predict their behavior in subsurface 

environments. Prior to this study, very little 

information on the redox reactions of iodine in 

the subsurface was available. Cs sorption and 

transport has been extensively studied in 

laboratory batch and column experiments, but 

very little information on field studies in 

saturated environments is available. In this 

study iodide oxidation, sorption of I species, and 

sorption of Cs species were investigated through 

a combination of laboratory and field 

experiments. 

Through the use of batch kinetic experiments, 

we discovered that the manganese oxide, 

birnessite, could oxidize iodide (I-

) to elemental 

iodine (I2) and iodate (IO3

-

) in a two-step 

process, according to the following reactions: 

1⁄2MnO2 + I-

 + 2H+ Æ 1⁄2Mn2+ + 1⁄2I2 + H2O 

(Eq. 4-1) 

5/2MnO2 + 1⁄2I2 + 4H+ Æ 5/2Mn2+ + IO3

-

 + 2H2O 

(Eq. 4-2) 

The oxidation kinetics depend on a variety of 

factors including ionic strength, pH, and 

birnessite concentration. Both I2 and to a lesser 

extent IO3

-

 adsorb to birnessite, reaching 

concentrations of up to 0.25 and 0.024 mmol/g, 

respectively. I-

 may also adsorb to birnessite, 

but this probably occurs quickly relative to the 

oxidation of I-

 to I2. Reaction 4-1 appears to be 

first order with respect to initial I-

 concentration, 

H+

 concentration, and birnessite concentration, 

with pseudo-first order rate constants ranging 

from 0.064 to 5.30 hr-1. Reaction 4-2, however, 

does not appear to be strictly first order with 

respect to initial I2 concentration and may be 

complicated by I2 sorption. Despite this fact, a 

simple model which considers both Reaction 4-1 

and 4-2 to be first order with respect to I fits the 

data fairly well. 

In batch sediment experiments, uptake of I-

 was 

low (2-3% at an initial concentration of 1 mM). 

In binary sediment-birnessite experiments, Iwas oxidized to I2 and IO3

-

. In a 1-m long 

column filled with aquifer sediment, I-

 uptake 

was low at pH 4.8, but increased when the pH 

was lowered to 4.50. At pH 4.5, Iconcentrations increased during breakthrough to 

90% of the input concentration, then slowly 

increased to 93% over the next few days. No I2

or IO3

-

 was detected during the column 

experiments, suggesting that I-

 was oxidized to 

elemental I2, which was volatilized from the 

samples prior to analysis. During a field tracer 

test in which 1 mM I-

 was injected into the oxic 

zone of a sand and gravel aquifer on Cape Cod, 

MA, I-

 was oxidized to I2 (up to 46%) and IO3

-

(up to 6%) during 4 m of transport. Mn-oxides 

appeared to be responsible for the I-

 oxidation, 

as evidenced by a pulse of dissolved Mn that 

was released from the sediments during the 

experiment. The tracer tests allowed us to 

investigate iodide oxidation over much larger 

transport distances than is feasible in the 

laboratory. Over 1 m of transport, I-

 behaved 

similarly in the column experiment and tracer 

test, although low concentrations of I2 and IO3

-

(1-2% of the input concentration) were 

measured in the tracer test. 

Iodate sorption to aquifer sediments was higher 

than iodide, with up to 12% of the initial 

concentration (1 mM) taken up by sediments in 

batch experiments. In a tracer test in which 1 

mM IO3

-

 was injected into the oxic zone of the 

aquifer, IO3

-

 was retarded relative to bromide. 

After 4 m of transport, the IO3

-

 pulse became 

very spread out and reached a maximum of 29% 

of the injection concentration. No evidence of 

reduction of IO3

-

 in either the batch experiments 

or the oxic tracer test was found. However, 

when IO3

-

 was injected into the Fe-reducing 

zone of the aquifer, it was completely reduced to 

I

-

 during 3 m of transport. A number of abiotic 

and biotic pathways for IO3

-

 reduction are 
30

possible, including reduction by Fe2+ or Mn2+, 

and microbially mediated reduction. 

In batch experiments with aquifer sediment, up 

to 22% of the Cs (initial concentration of 0.5 

mM) was sorbed. During the tracer tests, Cs 

sorption was strong, with the breakthrough very 

attenuated. After 4 m of transport the Cs 

concentration peaked at only 5% of the injection 

concentration (1 mM). Following the long peak 

in Cs concentration, a very long tail of low Cs 

concentrations continued for several months. 

The Cs appeared to adsorb via cation exchange, 

as evidenced by a pulse of K released from the 

sediments during the tracer tests. 

While the Cs and I concentrations used in these 

experiments are much higher than would be 

relevant for concentrations of radioactive 

isotopes of these elements, the studies are 

relevant for revealing reaction mechanisms that 

affect the transport of the radionuclides in the 

environment. 
31

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