United St8teI

Environmental Protection

Agency

Environmental Researd'I

L8b0ratory

NarraglnSett RI 02882

EPA-600/3-80-011

January 1980

&EPA Adsorption of Trace

Metals by Hydrous

Ferric Oxide in

Seawater
This document IS available to the public through the National Technicallnformation Service, Springfield, Virginia 22161.
ADSORPTION OF TRACE METALS BY

HYDROUS FERRIC OXIDE IN SEAWATER

K. C. Swallow and Francois Morel

Ralph M. Parsons Laboratory

For Water Resources and Hydrodynamics

Department of Civil Engineering

Massachusetts Institute of Technology

Cambridge, Massachusetts 02139

Grant R-803738

Project Officer

Earl W. Davey

Environmental Research Laboratory

Narragansett, Rhode Island 02882

ENVIRONMENTAL RESEARCH LABORATORY

OFFICE OF RESEARCH AND DEVELOPMENT

U. S. ENVIRONMENTAL PROTECTION AGENCY

NARRAGANSETT, RHODE ISLAND 02882

EPA-600/3-80-011

January 1980
DISCLAIMER

This report has been reviewed by the Environmental Research Laboratory,

Narragansett, U. S. Environmental Protection Agency, and approved for

publication. Approval does not signify that the contents necessarily reflect

the views and policies of the U. S. Environmental Protection Agency, nor does

mention of trade names or commercial products constitute endorsement or

recommendation for use.

ii
FOREWORD

The Environmental Research Laboratory of the U.S. Environmental

Protection Agency is located on the shore of Narragansett Bay, Rhode Island.

In order to assure the protection of marine resources, the laboratory is

charged with providing a scientifically sound basis for Agency decisions on

the environmental safety of various uses of marine systems. To a great

extent, this requires research on the tolerance of marine organisms and their

life stages as well as of ecosystems to many forms of pollution stress.

In addition, a knowledge of pollutant transport and fate is needed.

This report describes a three-year project aimed at improving modelling

capabilities for the fate of metallic waste constituents in coastal waters.

The particular focus of the report is the study of metal adsorption onto

hydrous ferric oxide.

Donald K. Phelps

Acting Director, ERLN

iii
PREFACE

Assessment of the environmental impact of activities which alter the

chemical characteristics of a body of water is made especially difficult by

the great number and complexity of the possible reactions through which a

chemical can be mobilized or precipitated and can take biologically active--

sometimes toxic--or inactive forms. One approach to such a problem is to

build reasonable models of the systems under study by assembling all relevant

thermodynamic. kinetic and analytical information. In this fashion at

least "chemical boundary conditions" can be established that permit the

isolation of the essential processes and point the way to relevant research

on smaller and more manageable experimental systems.

This report describes a three year project aimed at improving our

modelling capabilities for the fate of metallic waste constituents in

coastal waters. The particular focus of the report is the study of metal

adsorption onto hydrous ferric oxide. This process was singled out as the

one whose ignorance was most responsible for the inadequacies of our present

thermodynamic modelling of the metal chemistry in waste fields.

As part of the same project. new computer programs for chemical

calculations have been developed. These (particularly the MINEQL series)

have been widely distributed among universities. industry and government

research centers and are used routinely for environmental studies. A

particular focus of the modelling effort has focussed on the question of

alternative models for adsorption of solutes at solid water interfaces.

The following reports and publications should be consulted for specifics

regarding these various models:

MINEQL. A Computer Program for the Calculation of Chemical Equilibrium

Composition of Aqueous Systems. J.C. Westall. J.L. Zachary. F.M.M.

Morel. 1976. Water Quality Laboratory. Ralph M. Parsons Laboratory

for Water Resources and Environmental Engineering. Department of

Civil Engineering. Massachusetts Institute of Technology. Cambridge.

MA 02138. Technical Note No. 18.

The Use and Abuse of MINEQL-II. J.L. Zachary. 1977. Ralph M. Parsons

Laboratory for Water Resources and Environmental Engineering.

Department of Civil Engineering. Massachusetts Institute of Technology.

Cambridge. MA 02138.

Morel. F.M.M•• J .G. Yeasted. J .C. Westall. "Adsorption Models: A

Mathematical Analysis in the Framework of General Equilibrium

Calculations". In: M.A. Anderson and A. Rubin [eds.l. Adsorption of

Inorganics at the Solid-Liquid Interface. Ann Arbor Science, Ann Arbo~

iv
Michigan. 1979 (in press).

Chemical Equilibrium Including Adsorption on Charged Surfaces. J.C.

Westall. 1979 (in press). Swiss Federal Institute of Technology,

EAWAG, CH-860~Duebendorf, Switzerland.

v
ABSTRACT

The adsorption of trace metals by amorphous hydrous ferric oxide in seawater is studied with reference to simple model systems designed to isolate

the factors which may have an effect on the isotherms. Results show that

the complex system beijaves in a remarkably simple way and that the data obtained under various conditions of total metal concentration and total oxide

concentration can be reduced to an apparent reaction constant, K, which is a

function of pH only. The high capacity of the oxide for trace metals renders

the concept of a surface reaction useless to explain the uptake of metals. A

physical picture of the oxide as a swollen hydrous gel permeable to hydrated

ions is presented.

This report was submitted in fulfillment of Grant Number R-803738 under

the sponsorship of the Environmental Protection Agency. The work cited

in this report was completed as of April 1978 and is included in the Ph.D

thesis of Kathleen C. Swallow (Massachusetts Institute of Technology, 1978).

This work has also been submitted (July 1979) for publication to the journal,

Environmental Science and Technology in a paper entitled, "Sorption of Copper

and Lead by Hydrous Ferric Oxide" by K.C. Swallow, D.N. Hume and F.M.M. MoreL

vi
CONTENTS

Foreword ----------------------------------------------------------- iii

Preface ----------------------------------------------------------- iv

Abstract ---------------------------------------------------------- vi

Figures

Tables

1. Background -------------------------------------------------

Introduction ------------------------------------------

2. Experimental Materials and Methods --------------------------

Materials ---------------------------------------------

Methods ----------------------------------------------

3. Experimental Results ----------------------------------------

Potentiometric Acid-Base Titrations -------------------

Effect of Ionic Strength ------------------------------

Effect of Composition of the Background Electrolyte ---

Effect of variation of metal and oxide concentrations--

Competition between metals for the hydrous ferric oxide

4. Discussion -------------------------------------------------

References ---------------------------------------------------------

vii

viii

xi

1

1

13

13

15

19

19

22

22

22

39

44

50
8

Number

1

2

3

4

5

6

FIGURES

Adsorption in the electrical double layer

Expertmental apparatus -------------------------------------------- 16

-3 Titration curve for 1 x 10 M amorphous hydrous ferric oxide

in 0.5 M NaC104 --------------------------------------------------- 20

+ -3 Surface excess H or OH- as a function of pH for 1 x 10 M

hydrous ferric oxide ---------------------------------------------- 21

-4 Changes with aging in the titration curve for 1 x 10 M

hydrous ferric oxide in S.O.W. ~----------------------------------- 23

2+ -5 2+ Effect of aging on the Cu isotherm for 1 x 10 M Cu on

-4 1 x 10 M hydrous ferric oxide in S.O.W. -------------------------- 24

7 Effect of

2+ Cu on 1

2+ -5 ionic strength on the Cu isotherm for 1 x 10 M

-3

x 10 M hydrous ferric oxide prepared in the

different background electrolytes --------------------------------- 25

8 Effect of the ionic strength on

2+ for 1 x 10- the Cu isotherm 5

M

2+ -4 Cu on 1 x 10 hydrous ferric oxide prepared in one large

9

10

batch and diluted with the appropriate background electrolyte ----- 26

2+ 5 Effect of ionic strength on the Pb isotherm for 1 x 10- M

for 1 x 10-5M Pb2+ on 1 x lO-4M hydrous ferric oxide prepared

in one large batch and diluted with the appropriate background

electrolyte ---------------------------------------------------____ 27

2+ Effect of the background electrolyte composition on the Cu

viii
-5 2+ -3 isotherm for 1 x 10 M Cu on 1 x 10 M hydrous ferric

oxide prepared in different background electrolytes ---____________ 22

11 Effect of background electrolyte composition on the Cu2+

-5 2+ -4 isotherm for 1 x 10 M Cu on 1 x 10 M hydrous ferric

12

oxide prepared in one large batch and diluted with the

appropriate background electrolyte -------------------------------- 29

2+ Effect of background electrolyte composition on the Pb

-5 2+ -4 isotherm for 1 x 10 M Pb on 1 x 10 M hydrous ferric oxide

13

prepared in one large batch and diluted with the appropriate

background electrolyte -------------------------------------------- 30

-5 2+ -3 Isotherms for 1 x 10 M Cu on 1 x 10 M hydrous ferric oxide

under various conditions of ionic strength and background

electrolyte composition obtained with different batches of

hydrous ferric oxide prepared in the different background

as a function of pH for variable

ferric oxide --------------------------- 35

electrolyte ------------------------------------------------------- 32

2+ 2+ The ratio of Cu d b d/Cuf as a function of pH for variable a sor e ree

CUT on 1 x 10-4M hydrous ferric oxide ----------------------------_ 33

2+ Moles of Cu adsorbed per mole hydrous ferric oxide as a function

-4

x 10 M hydrous ferric oxide ---------- 34

mole hydrous ferric oxide as a function

-5 of pH for variable CUT on 5.0 x 10 M hydrous ferric oxide -------_ 36

2+ 2+ The ratio of Pb d b d/Pbf as a function of pH for variable a sor e ree

-5 PbT on 5.0 x 10 M hydrous ferric oxide ---------------____________ 37

of pH for variable CUT on 1

16 The 2+ 2+ ratio Cu d b d/Cuf a sor e ree

CUT on 5.0 x

-5 10 M hydrous

17 Moles 2+ of Cu adsorbed per

15

14

18

ix
19

20

21

2+ Moles of Pb adsorbed per mole hydrous ferric oxide as a

-5 function of pH for variable PbT on 5.0 x 10 M hydrous

ferric oxide ----------------------------------------------------- 38

2+ Per cent Cu adsorbed as a function of pH for various

concentrations of hydrous ferric oxide in S.O.W. ----------------- 40

2+ Per cent Pb adsorbed as a function of pH for various

concentrations of hydrous ferric oxide in S.O.W. ---------------- 41

22 Reduction of data from isotherms for 1.0 x 10-5M Cu2

+ on

various concentrations of hydrous ferric oxide in S.O.W. to

2+ 2+

Cu d b d/Cuf FeT as a function of pH ------------------------ 42 a sor e ree

23 Reduction of data from isotherms for 1.0 x 10-5M Pb2+ on

45

24

25

various concentrations of hydrous ferric oxide in S.O.W. to

2+ 2+

Pb d b d/Pbf FeT as a function of pH ------------------------ 43 a Bor e ree

, -5 2+ -4 Depression of the isotherm for 1.0 x 10 M Cu on 4.0 x 10 M

-4 hydrous ferric oxide in S.Q.W. in the presence of 6.0 x 10 M

Fe2+

2+

Log K as a function of pH for all Cu data --------------------- 47

TABLES

Number Page

1 The Composition of S.Q.W. ~----------------------------------------- 14

x
SECTION 1

BACKGROUND

INTRODUCTION

A comparison of the measured concentrations of trace metals in seawater

with calculations of the solubilities of the least soluble compounds which

can be formed from trace metals and the major components of seawater has led

to the conclusion that seawater is undersaturated with respect to many trace

elements (1,2,3). Various mechanisms for removal of trace metals to the sediments have been proposed as the controlling factor for their concentra~ions in

seawater. Krauskopf (1) performed a series of simple experiments to test the

efficiency of a number of processes for removing trace metals to a solid phas~

precipitation as the least soluble species, precipitation by sulfides and adsorption by ferrous sulfide in reducing environments, and adsorption on

hydrous ferric oxide, hydrous manganese oxide, apatite, clay and organic

matter. He concluded that adsorption on one or more of these substrates could

account for the low aqueous concentrations of eight of thirteen metals studie~

In particular, hydrous ferric oxide removed 95% of the Zn, 98% of the Cu, 92%

of the V, 86% of the Pb, 89% of the Co, and 80% of the W from seawater solutions with initial concentrations of 0.1 to 10 ppm metal.

Although the adsorption of metals on hydrous ferric oxides has been investigated (4,5), the experimental data exist mainly for simple, carefully

controlled systems designed to allow for maximum insight into the process of

adsorption at the microscopic level. Isotherms were generated for one metal

adsorbing on a crystalline hydrous ferric oxide surface in the presence of an

inert background electrolyte, usually at low ionic strength. Adsorption in

seawater, on the other hand, involves many adsorbing metals adsorbing onto an

ill-defined, possibly largely amorphous surface in the presence of a complex

background electrolyte at high ionic strength. Any or all of these factors

may make it unfeasible to extrapolate data from the simpler systems to seawater.

Factors Which May Affect Adsorption of Trace Metals by Hydrous Ferric Oxide in

Seawater

The Nature of the Hydrous Ferric Oxide--

In a study of the effects of ionic strength, temperature and Fe(III) concentration on the hydrolysis and precipitation of ferric oxides from ferric

nitrate solution, Dousma and deBruyn (6) found that 6.25 x 10-3M Fe(III) at

24°C produced an amorphous ferric oxide. Higher temperature and Fe(III) concentration were necessary to produce geothite (a FeOOII). This effect of

1
temperature had earlier been reported by Kolthoff and Moscovitz (7) who also

noted that the surface area of the crystalline oxide was smaller than that of

the amorphous oxide and that the former adsorbed three to four times less

copper than the latter.

In general, slow hydrolysis and/or elevated temperature are required to

produce the crystalline ferric oxides. Hematite (Fe203) is usually prepared

by boiling solutions of ferric nitrate at low pH under reflux for 18 days

(8). Goethite is obtained by aging a solution of ferric nitrate at pH 12 for

24 hours at 60° (9). These methods produce particles of considerable uniformity in size and shape whose composition depend most strongly on the pH

and the anion present in the solution.

Gadde and Laitinen (10) prepared amorphous hydrous ferric oxide by adjusting the pH of a O.lM ferric nitrate solution to 6 and allowing the precipitate to settle for 2 hours at room temperature. After filtering, washing

and aging overnight in distilled water at pH 6, the suspension was used for

the experiments. X-ray powder diffraction confirmed that the oxide was

amorphous.

Matijevic reports that the background anions, although responsible for

the particle characteristics, are not found in the solid phase. Ellis et ale

(11), on the other hand, found that Cl- ion is specifically adsorbed by

B FeOOH and cannot be removed completely even with extensive washing with

distilled water or by ion exchange.

An extensive investigation of the slow hydrolysis of partially neutralized Fe(III) solution~ in the nitrate, chloride and perchlorate solutions

was carried out by Murphy, Posner and Quirk (12). They found that the initial

polymerization process yields spherical polycations independent of the anion

present. The anion does affect the subsequent aging and crystalline structure

of the precipitate as it progresses from the spherical polycations to rods to

rafts. Rods are formed from linear coalescence of spherical polycations and

addition of unpolymerized ferric species in solution. Both rods and spherical polycations are X-ray amorphous. Rafts result from lateral coalescence

of rods and show the X-ray pattern of geothite or BFeOOH.

In nitrate and perchlorate solutions a FeOOH (goethite) predominates,

while in chloride solutions BFeOOH is formed. The formation of BFeOOH in

chloride solution is thought to arise from the penetration of the Cl- ion

into the polycations resulting in a different internal structure when the

polycations coalesce. The Fe(III) concentration does not affect the crystalline structure· in chloride solutions and increased ionic strength only increases the rate of precipitate formation. In more dilute solutions both

nitrate and perchlorate preparations also contained traces of y FeOOH (lepidacrocite). High ionic strength or anion affinity for the Fe3+ ion inhibit the

formation of this crystalline form.

In a study of Fe(III) speciation in seawater, Byrne and Kester (13) note

that the X-ray amorphous solid phase formed initially in solutions supersaturated with respect to Fe(OH)3 is slowly transformed to a more stable crystalline solid phase. This transformation is extremely slow, however, and after

2
several months of aging a substantial portion of the solid phase still exists

as the amorphous oxide. Dousma and deBruyn had also noted this slow transformation, but found that when the X-ray amorphous solid initially formed in

nitrate solution was studied by infrared techniques, the presence of a FeOOH

(geothite) was suggested. It has been proposed that the iron oxide found in

seawater, while X-ray amorphous, actually has a largely geothite nature (14).

The ferric oxide in seawater may have originated from terrestrial sources

or may have been precipitated in situ. Since an in situ precipitation would

occur in the presence of the potentially adsorbing trace metals, the question

is raised whether coprecipitation will be the same as adsorption on a previously formed surface. Kolthoff and Moskovitz (7) found that if the oxide

was formed at room temperature (amorphous), there was only a slight increase

in the amount of copper coprecipitated versus copper adsorbed. With precipitate formed at 98°C, however (crystalline), the differences were more

pronounced. Kurbatov, Kulp and Mack (15) found that a higher percentage of

strontium and barium was coprecipitated than was adsorbed on an amorphous

oxide, but that after several days of standing at pH 8 the amounts of adsorbed

and coprecipitated metal were identical. Gadde and Laitinen (10) found that

more lead was adsorbed on hydrous ferric oxide if it was present in the solution during the precipitation than if it was added immediately after the precipitation. Aged hydrous ferric oxide adsorbed less lead than fresh, and

lead adsorbed on the aged oxide was more efficiently recovered, suggesting

that some occlusion occurs in the coprecipitation process. In a later paper,

however (16), under the same conditions they report that the effect of aging

of the hydrous ferric oxide, either in the presence of lead or before it was

added, on the amount of lead adsorbed was not pronounced.

The Complexity of the Background Electrolyte--

The inert background electrolyte used in adsorption experiments usually

consists of a dilute solution of a simple salt whose ions have no specific

chemical interaction with the surface. Seawater is a complex solution which

contains high concentrations of the divalent ions Ca2+ and Mg2+ which may

specifically adsorb on the hydrous ferric oxide and out-compete a less concentrated trace metal for the available surface sites. Alternatively,

ligands such as Cl- that form stable complexes with the adsorbable metal may

outcompete the surface for the metal. In either case, a reduction in adsorption of trace metal would be expected.

MCNau~hton and James (17) reported no exchange or competition between

Na+ or Mg2 and the hydrolyzed Hg(II) species adsorbed on aQuartz. On the

other hand, O'Connor and Kester (18) found that Mg in a concentration equal to

that in seawater suppresses C0 2+ adsorption onto illite, but not Cu2+ adsorption. It is possible that adsorption on the clay, illite, and the oxide

aQuartz occur through different mechanisms and that results from one surface

are not relevant to the other.

The effect of competing ligands has been interpreted in several different

ways. Forbes, Posner and Quirk (4) interpreted the failure of HgC12 to adsorb

on goethite as an indication that coordinated hydroxyl groups play a vital

role in adsorption. The same reasoning was used by McNaughton and James to

explain the same result for HgC12 on aQuartz. They had found that the abrupt

3
authors have expressed doubt as to the applicability of this method to adsorption on hydrous oxides.

There is no reason to expect that adsorption of a gas occurs via the

same mechanism as specific chemical adsorption of ions from solution. In

addition, removing the solid from suspension and drying it in order to perform a BET determination may drastically affect the physical configuration

of the solid and alter the surface.

Another technique which is sometimes used involves the adsorption of a

dye, such as wooly violet, followed by removal of the solid from the suspension and colorimetric measurement of the residual dye in solution. Again

the adsorption of large dye molecules is not necessarily the same as adsorption of metal cations, but this technique can be used on an aqueous suspension

of the solid.

When the surface charge is attributed to specific sites on the surface,

the available reaction sites may be determined by acid-base titration of the

surface. This assumes that all sites which react with acid and base will

also react with adsorbing metals.

Regardless of the physical picture of surface charge distribution, the

determination of the available surface on an amorphous hydrous ferric oxide

suspended in the medium in which it was precipitated is difficult to carry

out. B.E.T. determinations cannot be made on solution; dye methods, while

feasible, are not likely to be appropriate; and acid-base titrations are

difficult to interpret because of the unknown concentration of strong acid

or base already present in the suspension after precipitation of the oxide.

If the available surface is limited, it should be possible to demonstrate

that the amount of metal adsorbed on a given amount of oxide reaches a limiting or plateau level above which an increase in the equilibrium concentration

of the metal does not result in further adsorption. James and Healy (Zl)

found that for CoZ+ adsorption on SiOZ, isotherms plotted as adsorbed C02+

versus free CoZ+ did not reach a plateau as such. With increasing CoZ+ the

isotherms converged toward the region where free CoZ+ concentration was

limited by the solubility of Co(OH)Z. They estimated a maximum adsorption

density corresponding to bare CoZ+ ions being separated by two to four water

molecules.

It was pointed out by James (Z5), however, that the adsorption and precipitation processes are apparently independent since a smaller percent of

the total CoZ+ is adsorbed with increasing added metal, which is opposite to

the concentration dependence of precipitation and polymerization.

Further evidence that adsorption and precipitation are independent processes is found in very early work by Kurbatov (Z6) on the preparation of

"highly emanating preparates". These are preparations of hydrous ferric

oxide containing adsorbed Ra used to obtain nuclear emanation in exactly

measured quantities. He found that exceeding the Ksp of Ra or Ba salts resulted in preparates of diminished radioactivity. This was attributed to

the formation of microscopic crystals of the salts which resulted in less Ra

4
rise in the isotherm for Hg(II) adsorption on aQuartz corresponded to the pH

at which Hg(OH)2 predominates in the speciation diagram. The correlation

between adsorption and hydrolysis led Matijevic (20) to the conclusion that

the hydroxyl group is responsible for the adsorption of ions on colloid

particles. He postulated that hydrogen bonding may playa decisive role.

James and Healy (21) rejected the idea that adsorption and hydrolysis

were related because of the presence of the -OH group. Instead, they suggested that in lowering the charge on the central cation, hydrolysis made it

easier to disrupt the secondary hydration sheath of the metal cation because

the charge dipole interactions were weakened. The disruption of the secondary

hydration sheath is necessary if an adsorbing cation is to move to within one

water molecule of the surface, their definition of specific adsorption. To

test James and Healy's interpretation of the correlation between hydrolysis

and adsorption, Stanton and Burger (22) used acetate and phosphate ions to

reduce the charge of Zn2+. Acetate had little effect, but phosphate enhanced

the adsorption of Zn2+ on amorphous iron and aluminum oxides. They interpreted the different results for the two anions as indicative of bridging between

the zinc ions and the oxide surface by the phosphate ions, similar to the

bridging role played by hydroxyl ions. Acetate could not form oxo bridges,

so had no effect.

In ammoniacal media, Kolthoff and Moskovitz (7) found that increased NH4+

concentration decreased Cu2+ adsorption on hydrous ferric oxide and attributed

it to lowering of the O~ concentration by the NH4+. Increased NH3 concentration, while increasing O~ concentration, also reduced Cu2+ adsorption, however. This effect was due to the formation of the stable Cu2+ -NH3 complexes

which apparently did not adsorb.

With illite, O'Connor and Kester (18) found less copper adsorbed at a

given pH from seawater or 0.7M NaCl than from an artificial river water

solution. There was no difference if the system was CO2 free or contained

a carbonate buffering system, however. Vuceta (23) also found that carbonate buffering system neither enhanced nor depressed adsorption of Cu + and

Pb2+ on aQuartz. In both cases the HC03- concentrations were in the range

of the concentration in natural waters and the C02 was atmospheric. Under

these conditions, the carbonate complexes of the metals are not the predominant species.

The Presence of Many Potentially Adsorbing Trace Metals and Limited Surface

for Adsorption

Much of the existing data for adsorption of metals on hydrous oxides has

been obtained with an excess of adsorbing surface and one metal in solution.

The meaning of available surface is not clear and has been measured or defined

by various methods.

If the charge on the surface is pictured as evenly distributed over the

entire surface area, it is common to assume that the limit of adsorption

density is surface monolayer coverage. This value is usually obtained by

measuring available surface area per amount of solid with the B.E.T. gas

adsorption method (24) or a similar technique. Then from calculations based

on hydrated ionic radii, a maximum adsorption density is obtained. Various

5
or Ba being uniformly distributed throughout the hydroxide.

Gadde and Laitinen (16) were able to demonstrate a plateau value for

Pb2+ adsorption on amorphous hydrous ferric oxide at pH 6. Their value was

0.28 moles Pb2+ per mole of Fe(III). They could not demonstrate a plateau

at pH 5 or for Pb2+ on hydrous manganese oxide. Cd2+, Zn2+ and Th2+ reached

a limiting plateau on amorphous hydrous manganese oxide at about 0.2 molel

mole Mn at pH 6. The values for the plateau levels were found to be pHdependent, decreasing with decreasing pH.

With a hematite suspension, Breeuwsma and Lyklema (27) found that Ca2+,

Mg2+, Sr2+, and Ba2+ were specifically adsorbed and shifted the pH of the

suspension toward lower pH until it reached pH 6.5 in 10-3N salt solution.

No further shift was observed up to 10-lN salt solution, indicating no further

adsorption. The surface charge computed from maximum adsorption density was

lower than 1 micro coulomb per cm2 • Comparison of this value with the computed maximum surface charge obtained from potentiometric titration and B.E.T.

surface area data of 75 ~C cm-2 , indicates the presence of special groups or

adsorption sites on the surface.

The fact that a hydrous oxide surface can be saturated opens the possibility of competition between adsorbing metals as the total metal concentration exceeds the concentration of available sites. Gadde and Laitinen (16)

were able to demonstrate that Pb2+ outcompetes Cd2+, Zn2+, and Th2+ for

hydrous manganese oxide sites regardless of which metal is added first. This

demonstrates not only competition but the existence of a reversible adsorption

mechanism through which an adsorbed metal ion can be replaced by other competing metal ions. On hydrous ferric oxide, Zn2+, Ca2+, and all in large

excess over Pb2+ had no effect on the Pb2+ adsorption.

Kurbatov, Kulp and Mack (28) in a study of Ba and Sr adsorption on

hydrous ferric oxide found plateaux for both metals at about 10-2 mole of

metal per mole of ferric oxide, but could demonstrate competition only if one

metal was present in great excess over the other.

Theoretical Models of Adsorption

Predicting the possible effects of each of these factors on the adsorption isotherms on hydrous ferric oxide in seawater depends to a great extent

on the physical and chemical picture of adsorption that is considered. The

basic concepts on which most adsorption models have been built were expounded

in the Gouy-Chapman theory of the electrical double layer as modified by

Stern -Grahame. This theory, known as the electrical double layer or EDL

theory, was derived with reference to either mercury drops which become

charged as the result of the external application of a known potential difference, or solids such as AgI which become charged as the result of assymetrical

dissolution of the lattice ions and for which these ions are potential

determining. The double layer potential for these solids is gotten from the

Nernst Equation:

o

6

(1)
where 'l'

o

R

T

F

= the potential difference across the electrical double layer formed

by the reversible transfer of potential determining ions

- the gas constant

- the temperature in oK

• the Faraday constant

a - +

o

a - +

the activity of the positively charged potential determining ion in

solution

the activity of the positively charged potential determining ion

at the point of zero charge

For either surface, rigorous thermodynamic calculations relating the surface

charge and the potential of the electrical double layer are possible. The

adsorption of ions at either of these surfaces is a function of the coulombic

attraction or repulsion of the charged surface for the ions and a specific

chemical interaction energy.

6 Gads = ZF'l'o - ¢

where Z • the charge on the adsorbing ion

F - the Faraday constant

(2)

the potential at the distance of closest approach of the adsorbing

ions (0)

¢ - the specific chemical interaction energy

The physical picture of adsorption that results from this treatment is represented in Figure 1.

Application of EDL theory to hydrous oxides is not completely straightforward. With hydrous oxides, much higher surface charges are developed than

can be accounted for by the electrokinetically measured double layer potentials (zeta potentials). The surface charge arises from reaction of the surface with H+ and OH: ions, which are therefore the potential-determining ions

for oxides.

+ ::M-OH + H -+ ::M-OH +

2 (3)

The Nernst equation for oxide surfaces takes the form:

'l' = 2 3 RT ( H) o • F P-~ZC - P (4)

where P~ZC = the pH at which the surface contains equal numbers of adsorbed

H+ and OH- ions

7
Solution

Charoa = 0 - 0

+

+

+

+

+-

solid surface

r:-

nner Helmholtz Plane_specifi~ adsorptlon ~uter Helmholtz Plane

1"'

I

I

~drl '-outer Helmholtz Plane, Charge z cr

Inner :lelmholtz Plane. Charge z o. d

Surface, Charge = 0

0 0

+

+

+

I

I

I

I

I

~dlE-

~adSurface.

Solid

Distribution of

Charge in the

Electrical Double

Layer

00

POl:c:ntial :telative to ·!jIo

&ulk Solution ~~-----:~.----------

+

:10 S!"cc ifie .'\dsorpti on Wit~ S'ecific AJsorption

Tha Electricai Double Layer

Figure 1. Adsorption in the electrical double layer.
pH = any other pH

The use of the Nernst equation to calculate double layer potentials for oxides

is only considered valid near pHpZC where the number of adsorbed potential

determining ions is small compared to the number of surface lattice ions.

Lyk1ema (29) points out that in order for equation (4) to hold, the chemical

potential of g+ on the surface, Sg+ must be assumed to be independent of the

activity of g+ on the surface. He maintains that this is unlikely since g+

is not a constituent of the unhydrated solid, and thus adsorption of H+ will

cause the surface layer composition to be different from that of the bulk

solid.

This problem is also discussed by Berube and deBruyn (30) who observed

that the differential capacity of the double layer on hydrous oxides, given

by:

where 0

o

surface charge

c =

ao

o

Wo

(5)

'¥

o = surface potential relative to the bulk solution

is much larger than the differential capacities for mercury or silver iodide

surfaces.

Several models of the double layer on hydrous oxides which could account

for the existence of high surface charge despite low double layer potential

have been proposed. The porous double layer model of Breeuwsma and Lyk1ema

(31) is based on the ability of the potential determining ions and some

counter ions to penetrate into the surface layers of the solid. Thus the

specificity of a charged surfaced in adsorbing ions of the opposite charge

depends on the geometric penetrability of the ions as well as on their affinity for the solid. Their data suggests that of the two, porosity is the more

important factor. Since much of the surface charge would be neutralized

within the pores of the solid, the potential across the electrical double

layer would be lower than that expected from consideration of the total surface charge.

Berube and deBruyn in their structured water model place the locus of the

potential determining g+ and OH- ions not on the surface itself, but in a

layer at least one molecular layer of water away. This highly structured

water layer between the surface and the adsorbed potential determining ions

is dissociated, possibly to a greater extent than bulk water, because of its

chemisorption at the surface. Since g+ and OH- are now lattice ions, the

constraints on the Nernst equation are no longer a problem. Adsorbed counterions, especially strongly hydrated ones, can now approach the region occupied

by the surface charge very closely, again allowing high surface charges to

exist despite low potentials some distance from the surface.

In the site-binding model of Yates, Levine and Healy (32), the charged

9
(6)

=.M-O-Na + =.M-O + Na

surface sites arising from reaction of the surface with potential determining

ions, react further with ions of the indifferent (non-specifically adsorbed)

supporting electrolyte, e.g. NaCl.

+ - 1 =.M-OH2 + Cl + =.M-OH2C

+ All sites involving the potential determining Hand OH ions are in one plane

at potential ~o. The supporting electrolyte ions are situated in a different

common plane, the Inner Helmholtz Plane. The Outer Helmholtz Plane is the

plane of closest approach of ions in the diffuse part of the double layer.

Again high surface charge and differential capacity of the double layer can

exist despite low potential at the Outer Helmholtz Plane.

These models have been proposed to explain characteristics of the electrical double layer peculiar to hydrous oxide surfaces. They produce refinements in the isotherms for electrostatically bound adsorbants, but do nothing

to elucidate the nature of specific chemical adsorption. The fact that

charge-reversal due to super-equivalent adsorption of cations on a negatively

charged surface and the adsorption of cations on a positively charged surfaces

exist, indicates that in many cases the specific chemical interaction is more

important than the electrostatic interaction.

James and Healy modified the EDL treatment of adsorption by adding a

third energy term to equation (2). They observe that for each metal specifically adsorbing on a given surface there is a critical pH range, usually less

than two units wide, over which the fractional amount of metal adsorbed

increases from zero to almost unity. The onset of the range is related to

the onset of hydrolysis and tqe subsequent lowering of the charge of the

adsorbing ion. The added term opposes adsorption of highly charged free aquo

metal ions and decreases in magnitude as the charge on the metal decreases

with hydrolysis. Thus:

+ (7)

~GO solvation represents the energy required to disrupt the hydration sheath

of an adsorbing metal ion so that it can move to within one water molecule

of the surface. It depends on the dielectric constant of the adsorbant and

the charge and radius of the adsorbing metal ion. The more hydrolyzed the

adsorbing metal ion is, the less charged it is and the less unfavorable the

energy of solvation is because the water of hydration is less tightly bound.

This model represents a departure from the classical EDL treatment of adsorption in that it includes more information about the nature of the specific

chemical adsorption process.

A different approach to adsorption on oxides is taken in the surface complexation model (33). Adsorption is treated as a chemical reaction whose

equilibrium constants are affected by the presence of the charged surface.

First, since the development of the surface charge is dependent on the acidbase properties of the oxide, one can write the constants for the surface

10
reactions:

+ + ~-OH2 ~-OH + H

::M-OH

where { } indicates surface groups

[ ) indicates species in solution

Ka1

.. {= M-OH}[H+)

{=M-OH2

+}

_ {= M-O-}[H+)

- {= M-OH}

(8)

These constants can be evaluated experimentally from data for the potentiometric acid-base titration of the surface. The constants obtained in this

manner are "microscopic" constants, however, which are pH dependent. The

"microscopic" constants are analogous to the series of acidity constants obtained as protons are removed from proteins or other polymers. As each

successive proton is removed, the remaining protons become more tightly bound

because of the increasing number of negatively charged groups. Likewise,

for a charged surface, the acidity of a proton depends on the relative charge

on the rest of the surface which in turn depends on the pH. These microscopic

or apparent acidity constants can be converted to intrinsic constants, independent of pH, by extrapolating the apparent constant versus surface charge

curve to the zero charge condition. From EDL theory it follows that the

change in the acidity constants with surface charge is given by:

K

a exp {F '¥fRT} (9)

where K .. intrinsic acidity constant

a int

K = microscopic or apparent acidity constant

a

Similarly, for a specifically adsorbed metal ion (Me2+) , the surface

reactions and stability constants are given by:

K s

1

2+ + + ::M-OH + Me ~-O-Me + H

2+ + 2::M-OH + Me (~-O)2Me + 2H

= {= M-O-Me+}[H+)

{= M-OH} [Me2+)

+ 2 8/ = {(~-O)2Me}[H )

{= M-OH}2 [Me2+)

(10)

for mono and bidentate complexes. Again, these are microscopic constants and

can be extrapolated to the pH of zero charge to obtain the intrinsic constanta

This model presents adsorption as a chemical reactions which is influenced by electrostatic interactions. It does not give any direct insight into

the electrical properties of the surface or address itself to purely cou1ombic

interactions.

11
In order to determine what are the effects of the various factors discussed above on the adsorption of trace metals by hydrous ferric oxide in seawater, it is first necessary to isolate the factors. This2~an be done by

starting with a simple system of one metal, for example Cu ,adsorbing on

hydrous ferric oxide in a simple, non-specifically adsorbed background electr~

lyte, for example NaC104. Once a baseline result is established for the

simple system, the system can gradually be changed and/or made more complex

and any changes in the adsorption isotherms can be ascribed to the appropriate

factor. Finally, the isotherm for Cu2+ on hydrous ferric oxide in the

presence of the other adsorbable metals in seawater can be interpreted on the

basis of the known influence of each of the factors. Using the experimental

results, it should be possible to determine which, if any, of the theoretical

models of adsorption provides a satisfactory explanation for the process of

adsorption in seawater.

12
SECTION 2

EXPERIMENTAL MATERIALS AND METHODS

MATERIALS

Chemicals

All reagents were prepared from Analytical Reagent grade chemicals using

distilled, deionized water.

Sodium hydroxide solutions were prepared from 50% NaOH solution and

standardized against potassium acid phthalate with phenolphthalein indicator.

They were stored in polyethylene dispenser bottles and kept C02 free.

6 N hydrochloric acid was distilled from a quartz still. Solutions were

standardized against standard base with phenolphthalein indicator.

Synthetic Ocean Water (S.O.W.) was prepared according to the formulation

of the FWPCA (34). This is a complex solution of eleven salts containing the

major constituents of seawater. The composition is given in Table I. The

S.O.W. was passed through a Chelex 100 chelating resin column at pH 8.5 to

remove trace metals present as impurities in the reagent grade major salts

(35).

Copper and lead stock solutions were 0.01 and 0.1M Cu(N03 )2 and Pb(N03)t

50 ~l of the appropriate 0.01 M solution was added to 50 ml or sample gave

a final trace metal concentration of 1 x 10-5M. This concentration was used

for most experiments with 1 x 10-4M (as Fe) hydrous ferric oxide. In the

experiments where total trace metal concentration was varied, differing

amounts of the stock solutions in both concentrations were used to give the

desired final concentration. Stock solutions were naturally acidic as a

result of the hydrolysis of the Cu2+ and Pb2+. Periodically fresh stock

solutions were prepared and both old and fresh solutions were analyzed by

Atomic Absorption Spectrophotometry. Results showed that the stock solutions

remained stable indefinitely.

The ferric iron stock solution was 1.0 M Fe(N03)3. It had a pH of 2.7

immediately after preparation, so no acid was added to stabilize it. The

amber color of this solution increased in intensity with time, indicating

increasing polymerization of the iron hydrolysis species, but no cloudiness

or precipitate appeared. Periodic iron analysis by the orthophenanthroline

procedure indicated that the concentration was stable indefinitely.

Other trace metal solutions were prepared fresh as needed from reagent

13
salt

NaCl

KBr

NaF

KCl

Table 1

The Composition of S.O.W.

final concentration

4.20 x 10-~

1.05 x 10-~

8.40 x 10-4M

7.14 x 10-5M

9.39 x 10-3M

4.85 x 10-4M

-2 2.88 x 10 M

-3 2.38 x 10 M

-5 6.38 x 10 M

14
grade chemicals.

Apparatus

Before use, all glassware was cleaned, rinsed with distilled~ deionized

water, soaked in 3N HCL for at least 12 hours, and rinsed with distilled,

deionized water. All tubing and stirrers which contacted the sample solutions

were Teflon. Some experiments were done in acid-soaked glass beakers, others

in Teflon beakers. The apparatus used for most of the experiments is shown

in Figure 2. The sample was placed in a 100 ml beaker, either glass or Teflon

with a Teflon coated stirring bar. This beaker was placed inside a 250 m1

jacketed glass beaker which was connected to a Haake circulating constant

temperature bath. The temperature was maintained at 25°C for all experiments.

A #10 silicone rubber stopper bored to accommodate two electrodes, a N2 inlet

tube, a sample withdrawal tube and a reagent introduction inlet fit snugly

into the top of the beaker. The N2 was passed first through an Ascaritefilled drying tube to remove C02 and then through distilled, deionized water

to saturate it before it entered the reaction vessel through Teflon tubing.

The sample withdrawal tube was a Teflon tube ending in a Luer-loc fitting.

This end was kept immersed in a test tube of water to provide a pressure

outlet without allowing air to enter the system. The other end was only

immersed in the sample when an aliquot was being withdrawn. Reagents were

introduced into the system with Eppendorf pipettes. The reagent introduction

inlet was an Eppendorf pipette tip (5-100 ~l size), with about 2 cm cut off

from the tip inserted in the stopper until the cut off end was flush with the

stopper bottom. Except when in use, this inlet was stopped with another

Eppendorf pipette tip either filled with Ascarite or closed off with a medicine dropper bulb. When completely assembled this apparatus was a C02-free

system. The pH was measured with an Orion Model 91-01 glass electrode versus

an Orion Model 90-02 double junction reference electrode. The electrodes were

calibrated daily using Fisher Certified Buffer Solutions at pH 7 and 4.

Readout was obtained on either an Orion Model 701 digital pH meter of an Orion

Model BOlA digital ionalyzer. The output was recorded on a Cole-Parmer dual

channel strip chart recorder.

METHODS

Preparation of Hydrous Ferric Oxide

The hydrous ferric oxide sols used for these experiments were prepared

by filling a volumetric flask with the appropriate background electrolyte

solution, stirring with a magnetic stirring bar, and adding stock 1M Fe(N03)3

with an Eppendorf pipette. In NaCl04 and NaCl solutions a stoichiometric

amount of IN NaOH, calculated by assuming the formula Fe(OH)3 for the precipitate, was added to raise the pH and ensure complete, fast precipitation.

In synthetic ocean water (S.O.W.), the natural buffering made this unnecessar~

as the pH stayed around 7.5. All preparations were done open to the atmosphere. For experiments using freshly precipitated hydrous ferric oxide (HFO),

only enough suspension was prepared for one day's experiments, usually 100 or

250 ml. For experiments with aged BFa, 2 liters were prepared, which yielded

forty 50 ml samples.

15
:.

"I / 1 7.0 01 1 ... I / -0- "

/"" ..."", -J1

• :@~ / " ~==::t.-.JL_-l____ __-----',

Figure 2. Experimental apparatus.

, I ,. 0

-0-

" I ,
In experiments done on "fresh" precipitates, the suspension was stirred

for one hour before sampling. During stirring the oxide remained uniformly

suspended and no clumping of particles or accumulation of solid on the walls

of the flask occurred. Samples were withdrawn with 50 ml volumetric pipettes

while the suspension was stirred vigorously. The sample was transferred to

the beaker and placed in the apparatus shown in Figure 2. The experiment

was begun after pH equilibrium was reached. A signal drift below 0.002 pH

units/min was used as the criterion for equilibrium. The total time elapsed

between preparation and the start of an experiment on fresh hydrous iron

oxide was around' 3 hours.

In the first experiment done on "aged" precipitates, the preparation was

stirred continuously and sampled at appropriate times measured as time after

preparation. After 144 hours it was noticed that small shavings of Teflon

had flaked off the magnetic stirring bar and had become nuclei for agglomerations of the ferric hydroxide sol. This was extensive enough to significantly alter the concentration of iron in the samples withdrawn, since these

agglomerates were not taken into the sampling pipettes. Later experiments

on aged precipitates were done with sols that were quiescent between samplings and only stirred during sampling. The settling out of the precipitate

did not appear to affect its later dispersability or chemical characteristic&

Potentiometric Titrations

The titrations were performed in the C02-free system by adding a known

excess of acid--either HC104 for NaC104 solutions or HCl for NaCl or S.O.W.

solutions--to the sample in the beaker, to lower the pH below 4. The N2 was

turned on and the CO2 purged from the sample. After the pH equilibrated, but

not before 10 minutes, the pH was recorded and a known volume of standard

0.05N NaOH (C02-free) was added with an Eppendorf pipette. At the low and

high ends of the pH range, 50 ~l additions were made. At the break, additions were reduced to 10 ~l. The pH was allowed to re-equilibrate after

each addition and was recorded. A drift of less than 0.002 pH units/min was

used as a measure of equilibrium. This was continued until the sample reached pH 10. All titrations were kept free of CO2,

Generation of Isotherms

To generate an isotherm a 50 ml portion of the hydrous ferric oxide sol

was placed in the beaker in the C02-free apparatus shown in Figure 2. HCI04

or HCl was added to lower the sample pH below 4 and the sample was purged of

C02' After pH equilibration a known volume of a trace metal solution was

added with an Eppendorf pipette. An addition of 0.05N NaOH (C02-free) was

made to raise the pH. At low pH, 0.05N NaOH was added in 100 ~l portions.

Near the break of the titration the size of each addition was lowered to

50, 10, and finally 5 ~l. This made it possible to obtain samples distributed over the entire range from pH 4 to pH 7. In the low and high pH regions,

pH equilibrations occurred within one minute. Near the break, the time

necessary for equilibration became longer, reaching a maximum of about one

hour. Equilibration times of longer than one hour were indicative of a

clogged reference electrode, a pressure build-up in the system or inadequate

stirring. A pH drift back to lower pH indicated C02 contamination which

17
necessitated termination of the experiment.

After pH equilibration the sample withdrawal tube was lowered into the

suspension and a 5 cc disposable syringe was attached to the Luer-loc fitting.

About 3 m1 of the suspension and another addition of base was made to the

sample in the beaker. A Swinnex Millipore filter holder with a 0.025 ~m

Cellulose acetate filter was attached to the syringe containing the withdrawn sample. The sample was expressed through the filter into a 25 m1

Erlermeyer flask with a standard taper glass stopper. 10 ~l of 6N HCl were

added to the flask, it was stoppered, marked for identification and set

aside until all samples were collected in the same way.

When all the samples for an isotherm were collected they were analyzed

for the metal or metals of interest by flame atomic absorption spectrophotometry.

The efficiency of the 0.025 ~m Millipore filters was tested by analyzin~

the filtrate for iron by the orthophenanthroline procedure. 1.2 x 10-6M Fe3

was left in the filtrate from a 1 x 10-~ (as Fe) suspension of hydrous

ferric oxide. This corresponds to 99.8% removal of the iron.

Flame Atomic Absorption Spectrophotometry

Flame Atomic Absorption Spectrophotometry was done on a Perkin-Elmer

Model 360 double beam atomic absorption spectrophotometer. All of the metals

analyzed required an air-acetylene flame. Metal concentrations were always

within the detectable range of the instrument so no pre-concentration steps

were necessary. Standards were prepared in concentrations to bracket the

sample concentrations using the same background electrolyte as the sample.

The high salt content of standards and samples caused some variability in

absorption values obtained for different aspirations of the same sample.

Because of this, all standards and samples were aspirated in random order at

least three times. The average of the three values was taken as the absorbance for a sample. The relative standard deviation of the three readings

ranged from 1 to 3%. Standard curves prepared from the samples were linear

through all points, but had non-zero intercepts on the absorbance (y) axis

corresponding to the absorbance reading for background electrolyte with no

added metal. It was determined that this was due to the high salt content

and its interference with the flame and not to ambient metal concentrations

in the reagents by using the Cd lamp to analyze a solution of Cu. Cd has a

spectral line at 326.1 nm, very close to the Cu line at 324.8 nm. A distilled water solution of Cu has zero absorbance at this wavelength, but the

0.5 M NaCl04 0.5M NaCl and S.O.W. solutions all gave absorbance readings

corresponding to the non-zero intercept of their standard curves. The nonzero intercept was reproducible and stable from day to day.

18
SECTION 3

EXPERIMENTAL RESULTS

POTENTIOMETRIC ACID-BASE TITRATIONS

Existing models of adsorption of trace metals by hydrous oxides rely

heavily on data from acid-base titrations of the solid. These data are used

to compute the pHpZC and the number of reactive sites on the solid surface.

In this work, to avoid possible changes in the nature of the precipitate due

to filtration, the titrations were performed in the same medium from which

the hydrous ferric oxide had been precipitated. Reproducible titrations of

the amorphous hydrous ferric oxide used in this work were very difficult to

obtain. Similar difficulties were encountered by Davis (36). The experimental errors in the addition of Fe(N03)3 and NaOH to the medium resulted in

variability in the position of the titration curves. A representative titration curve is shown in Figure 3 for 1 x 10-~ amorphous hydrous ferric oxide

in 0.5 M NaC104• Similar curves were obtained in 0.1 and 0.25 M NaC104 .

The proton excess or deficiency in the solid phase (fg+ - fOH-) is calculated from the titration curves as the difference in the amount of acid or

base needed to reach a given pH in the sample and the blank. It is related

to surface charge by

cr - o

where cr = surface charge o

(1)

+ f

H

+ - f OH- = surface excess Hover OH

A = surface area

F = Faraday constant

Because the methods available for determining surface area are unsatisfactory for use with suspensions in the medium in which they have been precipitated, surface excess in eq/mole has been substituted for surface charge

in this work. If a constant specific surface area is assumed the two differ

by only a constant.

Plots of the surface excess shown in Figure 4 for the three ionic

strengths are the same within experimental error. They do not exhibit the

sigmoid characteristics obtained for other oxides (5, 30, 31). Therefore a

19
4 base 2

Sample _

0.5 M NaCI04

4

9

5

I

I

I

~8lank

I

I

I

I

I

I

I

I

3 '--__....I--__-L..__---& -=-, L..-__....1...-__--'-__---J

4 acid 2 eq x 104

8

10

7

pH

6 N

o

Figure 3. -3 Titration curve for 1 x 10 ~f anorphous hydrous ferric oxide

in 0.5 H NaCI04 .
-10 ,.--------------------------------,

-8

9 10

NaCI04

NaCI0 4

NaCI0 4

8

• .1 M

6. .25M

o .5M

5 6 7

pH

+ - -3 Figure 4. Surface excess H or OH as a function of pH for 1 x 10 M hydrous ferric oxide.

)(

-...

(f)

(f)

N Q)

t-' U

)(

W

Q)

u

0

lo...

:J

(f)

16

4
value for maximum reactive sites cannot be obtained, although it must be in

excess of 0.15 site/Fe. The p~Z appears to be in the range of 7.9 - 8.5.

This is in agreement with data re~orted for various crystalline and amorphous

iron oxides (4, 5, 8, 29).

Titration curves obtained as soon as possible after preparation of the

oxide (ca. 2 hours due to pH equilibration time) required more acid than

curves obtained for the same oxide after 24 hours aging. No further change

occurred after 24 hours. This shift in the titration curves, shown in Figure

5, can be interpreted as a release of acid or a consumptio~4of base by the

aging oxide. The two curves are roughly parallel 1.3 x 10 eq apart for

1 x 10-4M Fe. Copper isotherms obtained simultaneously with the same batch

of oxide showed no effect of aging. This is shown in Figure 6 which also

illustrates that the same isotherm was obtained when copper was coprecipitated

with the iron.

EFFECT OF IONIC STRENGTH

2+ The effect of ionic strength on the Cu isotherm was first studied using hydrous ferric oxide precipitated in media of varying ionic strength.

The results shown in Figure 7, exhibit a non-systematic variability presumably due to differences among the batches of ferric oxide. To eliminate

batch to batch variations in the oxide a large batch of 2 x 10-4M hydrous

ferric oxide was prepared in 0.01 M NaC104' By appropriate dilution (1:1)

with 0.99 M NaC104, 0.09 M NaC104, or deionized, distilled water, a background electrolyte concentration of 0.5 M NaC104' 0.05 M NaC104 or 0.005 M

NaC104 could be had with the same batch of oxide. As expected, the results

were less variable and demonstrate no effect of ionic strength either on the

Cu2+ isotherm (Figure 8) or on the Pb2+ isotherm (Figure 9).

EFFECT OF COMPOSITION OF THE BACKGROUND ELECTROLYTE

The effect of the composition of the background electrolyte was studied

in the same way. The Cu2+ isotherms obtained with hydrous ferric oxide precipitated in 0.5 M NaC104, 0.5 M NaCl and S.O.W. (Figure 10) again show a

small non-systematic variability. Diluting portions of the large batch of

2 x 10-4M hydrous ferric oxide in 0.01 M NaC104 with 2 x S.O.W. and 1.0 M NaCl

gave background electrolyte compositions of S.O.W. and 0.5 M NaCl for the

same batch of iron. While the Cu2+ isotherms were identical to the one

obtained in 0.5 M NaC104 (Figure 11), the Pb2+ isotherms were equally depressed relative to the Pb2+ isotherm in 0.5 M NaC104• The diffePence follows

from the relative strengths of the Cl- complexes of the two metals. While

the chloride complexes of Cu(II) are relatively unimportant, Pb(II) is known

to form important di and trichloro complexes at (Cl-) - 0.5 M. Apparently

the Cl- complexes are not adsorbed and the result of Cl- complex formation

is a decrease in available Pb2+ which depresses the isotherm. Neither the

Cu2+ nor the Pb2+ isotherm is affected b

2high concentrations of Mg2+ and

Ca2+, indicating either than Mg2+ and Ca + are not adsorbed or that they can

be displaced by Cu2+ and Pb2+.

EFFECT OF VARIATION OF METAL AND OXIDE CONCENTRATIONS

An important aspect of the amorphous hydrous ferric oxide system is the

22
10 --.--------

o

O _. _-x

--Xx Blank 24 hr

• 2 hrs

24 hrs

o 48 hrs

o 72 hrs

•

.------.

.,.---

/.,....... , ...X-

, x..... . /

./ /

I /X

/ x • x

I f

/ !

• I

./ //

.~ ~x •..----- x- .--.-.Arl.-JJ-'-"

5

4

8

9

7

pH

6

Figure 5. ferric oxide in S.O.W.

3 :-;1:---....L------:1~9--....J.---:-l=-----l----L---L---...L-----.L----l 17 15 13 II

meq acid x 100

-4 Changes with aging in the titration curve for 1 x 10 M hydrous
•

1

0

x ,

•

j ~.o

<> 0 'I 0 <>

-Jl,-----..-.O

•

<>

Fe T = I x 10 - 4 M

CUT = I X 10-5 M

HFO genp.rateci in presence of Cu2+

HFO prepared, Cu 2

+ added immediately

ofter pH equilibration

HFO prepared ,aged 24 hrs

HFO prepared, aged 48 hrs

x HFO prepared ,aged 72 hrs

•

o

•

<> 100

80

30

90

10

70

"tJ

60

o

CI)

"tJ

« 50

+

N

:::J

U 40

o

4 5

pH

6 7

Figure 6. 2+ -5 Effect of aging on the Cu isotherm for 1 x 10 M

2+ -4 Cu on 1 x 10 M hydrous ferric oxide in S.O.W.

24
100

90

80

70

0

w 60 CD

0::

0

en

0

« 50

+

N

:;)

U

0

40 • 0.5 M

• 0.05M

30 • 0.005M

Fer=10-3M

20 Cur= 10-5M

•

•

10

•

•

o 4 5

pH

6 7 8

2+ Figure 7. Effect of ionic strength on the Cu isotherm for

-5 2+ -3 1 x 10 M Cu on 1 x 10 M hydrous ferric oxide prepared

in the different background electrolytes.

25
100 r------------ ----.

90

80

70

u

Q) 60

.0

0

f/) • 0.5 M No C104

u 50

<t A 0.05 M No C104

+

N

::J 00.005 M NoCI04

u 40

0

30

20

10

4 5

pH

6 7

Figure 8. 2+ Effect of the ionic strength on the Cu isotherm for

-5 2+ -4 1 x 10 M Cu on 1 x 10 hydrous ferric oxide prepared

in one large batch and diluted with the appropriate

background electrolyte.

26
10

100

•

90

FeT=lxI0- 4 M

80 PbT = I x 10- 5 M

70

"tJ

• 0.5 M NaCI04

Q)

..c 60 6 0.05 M NaCI0 4

0

CJ) 0 0.005 M NaCI04

'U

<:t 50

+

N

.c

0.. 40

0

30

•

20

4 5 6 7

pH

Figure 9. Effect of ionic strength on the Pb2+ isotherm for 1 x lO-5M Pb2+ on 1 x lO-4M

hydrous ferric oxide prepared in one large batch and diluted with the

appropriate background electrolyte.

27
100

90

80

70

CI

w 60

CD

0::

o

(/)

CI

<t 50

+

N

:J

U

o 40

30

20

10

•

•

~.-.

• 0.5 M NoCIO~

• 0.5M NoCI

• SOW

FeT =10-3 M

CUT =10- 5 M

o 4 5 6 7 8

pH

Figure 10. Effect of the back~round electrolyte composition on the Cu2+

-5 2+ -3 isotherm for 1 x 10 M Cu on 1 x 10 M hydrous ferric oxide

prepared in the different background electrolytes.

28
100 ,-------------- --,

90

•

80

70

"0 60 (])

.Q

lo-

• 0.5 M NaCI04

e

V> 50 6 0.5 M NaCI "0

<!

0 SOW "'" N::J 4 a

u

0

30

20

10

•

o

4 5 6 7

Figure 11.

pH 2+ Effect of background electrolyte composition on the Cu isotherm

-5 2+ -4 for 1 x 10 M Cu on 1 x 10 M hydrous ferric oxide prepared in

one large batch and diluted with the appropriate background

electrolyte.

29
100 ,-------------------,

90

80

70 - •

"0 / <1J

.a 60 • 0.5M NaCI04

(/)

0 6 0.5M NaCI •

"0

« 50 0 SOW

+

N

..a

a.. 40 0

0

30

20

10 -6

•

/

•

4 5 6 7

Figure 12.

pH

2+ Effect of background electrolyte composition on the Ph isotherm

-5 2+ -4 for 1 x 10 M Pb on 1 x 10 M hydrous ferric oxide prepared in

one large batch and diluted with the appropriate background

electrolyte.

30
batch to batch variability under controlled laboratory conditions. Figure

13 which consolidates the Cu2+ data from seven different batches of iron

under various conditions of ionic strength and background electrolyte composition, gives a measure of this variability. Some of the low data at high

pH may reflect contamination of the filtrate from new apparatus and not true

batch to batch variations. Nevertheless there are significant differences

among the isotherms obtained with different batches which must be considered

when comparing the data of different researchers or applying the results to

natural systems.

To elucidate the nature of the reaction, metal and oxide concentrations

were systematically varied. The precipitation of the Cu2+ and Pb2+ hydroxides limited the metal concentrations that could be used to ca. 1 x 10-4M

Cu2+ and 4 x 10-4M Pb2+. The minimum oxide concentration yielding measurable

uptake of metals was 5 x 10-5M.

In the first set of experiments the hydrous ferric oxide concentration

was kept at 1.0 x 10-4M while the Cu~+ concentration was varied. From

consideration of a simple mass balance

- 2+ -+ + :=M-O + Me +- :=M-Q-Me (1)

it was expected that below saturation at a given pH the ratio of adsorbed to

free metal would remain constant while the total moles of metal adsorbed

increased as long as the available reaction sites were in large excess. As

saturation was approached, the ratio of adsorbed to free metal would begin

to decrease as the available reaction sites decreased. According to this

interpretation, the results shown in Figures 14 and 15 imply that saturation

is not approached up to a Cu2+/Fe ratio of 0.2. The ratio CUi!sorbed/CU~+ee remained constant while the moles of Cu2+ adsorbed increased with increasIng

Cu2+. Increasing the Cu2+ concentration above 1.0 x 10-4M would have led to

precipitation of Cu(OH)2 in the range in which the isotherms showed significant adsorption of Cu2+. The ferric oxide concentration was therefore decreased to 5.0 x 10-~ and the experiment repeated with the same Cu2+ concentrations. It appears that even under these conditions, an approach to

saturation cannot be reached before the Cu2+ precipitates as Cu(OH)2

(Figures 16-17). The distinction between adsorption and precipitation of

Cu2+ in these experiments was very clear. Before the onset of precipitation,

a gradual decrease in free Cu2+ accompanied a gradual increase in pH as

base was added. When precipitation of Cu(OH)2 began, a rapid decrease in

free Cu2+ occurred while the pH remained constant with the addition of base.

As previously mentioned! further reduction of the hydrous ferric oxide concentration to 1.0 x 10-JM resulted in negligible Cu2+ adsorption in the pH

range 4-7.

2+ Since the Cu hydrolysis and precipitation made it impossible to demonstrate saturation of the hydrous ferric oxide with Cu2+, the experiments

were repeated with Pb2+ which precipitates as Pb(OH)2 at a slightly higher

pH than Cu2+ precipitates as Cu(OH)2, but gives a similar isotherm on hydrous

ferric oxide. The results shown in Figures 18 and 19 are not as straightforward as those for Cu2+. During the generation of the isotherms, the

31
8

o

o~_......,

7

Fer =10-3 M

CUr =10-5M

6

10

o

80

90

30

20

100

70

00

Cl

UJ

m 60

0::

0

(/)

Cl

<f 50

+

N

::J

U

0

40

Figure 13.

pH

-5 2+ -3 Isotherms for 1 x 10 M eu on 1 x 10 M hydrous ferric oxide

under various conditions of ionic streneth and background electrolyte

composition obtained with different batch~s of h~drous ferric oxide

prepared in the different background electrolyte.

32
•

I

1.0 r---------------------

4 5 7

.1

.2

.9

.3

.8 -

0 CUr = 5 x 10- 6 M

.7 • CUT = Ix10- 5 M

6 Cu r = 5 x 10- 5 M

CII

CII .6 0 CUT = Ix 10-4M

+ ...

Fer = I x 10- 4 M (\J~

u .5 "-

II)

+'0

NO

:J

u .4

Figure 14.

pH 2+ 2+ The ratio Cu d b d/Cuf as a function of pH for variable Cua sor e ree T

on 1 x 10-~ hydrous ferric oxide.

33
A

.20

o

6 7

pH

5

e---

o CUT = 5 X 10- 6 M

e CUT'= 1X 10-S M

A CUT = 5 X 10-5 M 0

o CUT =1X 10- 4 M

FeT a r X 10-4 M

4

0

IJ.

::c .15

(U

-

0

:E

"-

(U

a.

"0

CP

.c

"- .10

0

en

"C

<:{

+

N

:J

U

(J)

ell

0.05

Figure 15. 2+ Moles of Cu adsorbed per mole hydrous ferric oxide as a function

-4 of pH for variable C~ on 1 x 10 M hydrouR ferric oxide.

34
1.0

.9

.8

.7

-.6 +-

Not

U

..........

.. +1:' .5 NO

u

.4

.3

.2

.1

4

• Cur =I x 10-5M

x CUr=4xI0-5M

• CUr =5xI0-5M

6 CUr =I x10-4 M

Fer =5xI0-5M

x------:7

•

5 6 7

Figure 16.

2+ 2+ pH The ratio Cu d b d/Cuf as a function of pH for variable Cua sor e ree T

on 5.0 x 10-~ hydrous ferric oxide.

35
.8

4 5

pH

• ......... ,....... _e-·

6

a

7

Figure 17. 2+ Moles of Cu adsorbed per mole hydrous ferric oxide as a function

-'l of pH for variable CUT on 5.0 x 10 -M hydrous ferric oxide.

36
1.0

.9

.8

.7

• Pbr =I x 10-4M

x Pbr=2x 10-4 M

• Pbr =3xI0-4 M

6 Pbr =4xI0-4 M

Fer =5 x 10-5 M

4 5 6 7

pH

Figure 18. The ratio of Pb2+ /Pb2+ as a function of pH for variable adsorbed free

PbT on 5.0 x -5 10 M hydrous ferric oxide.

37
1.0 ,--~--------

.8

0

lL.

I

Q)

-

0

.6

Q)

Q.

1J

Q)

1\

0

CI)

1J

<t

+ .4 N

.0

a..

CI)

Q)

-

0

0

.2

I x 10- 4 M Pb 2 0 +

A 2 X 10- 4 M Pb 2 +

• 3 X 10-4 M Pb 2

+

0 4 X 10- 4 M Pb 2

+

4 5

pH

6 7

Fi~ure 19. Moles of Pb2+ adsorbed per mole hydrous ferric oxide as a function of

-5 pH for variable Pb on 5.0 x 10 M hydrous ferric oxide.

T 38
Pb2+ solutions apparently became supersaturated with respect to Pb(OH)2 and

an addition of base actually lowered the pH if precipitation commenced. The

complex and variable hydrolysis of Pb2+ in these solutions reflects the known

polynuclear behavior of Pb2+ hydrolysis products. Britton (37) in a study

of the precipitation of trace metal hydroxides, found that Ksp for Pb(OH)2

increased with the amount of alkali added, ranging from 3 x 10-19 to

1.35 x 10-15 . A ~henomenon of this kind is indicated in the data presented

here. At high Pb + concentration (4.0 x 10-4M) the precipitation of Pb(OH)2

caused the hydrous ferric oxide to coagulate and a white solid was carried

down with the brick red ferric oxide. Before the appearance of the white

solid, however, the distinction between adsorption and precipitation of the

hydroxide was not as clear for Pb2+ as it was for Cu2+. Despite these 2+

difficulties, an approach to saturation is suggested with the highest Pb

concentration at a 0.8Pb/Fe ratio.

The expected increase in Cu2+ and Pb2+ adsorption with increasing FeT

at a given pH, is shown in Figures 20 and 21. The reduction of the data to

and

Cu2+

adsorbed

2+ CU f Fe

ree T

Pb2+

adsorbed

2+ Pbf Fer ree

(1)

(2)

presented in Figures 22 and 23, respectively shows that the increase is

approximately proportional to the concentration of the oxide. The high points

in the Cu2+ plot at low pH and the low points in the Pb2+ plot at high pH

reflect large indeterminant errors in the isotherms.

COMPETITION BETWEEN METALS FOR THE HYDROUS FERRIC OXIDE

. 2+ 2+ It was imposs1ble to demonstrate competition between Pb and Cu for

the iron. This was consistent with the difficulties encountered in demonstrating saturation for either of the metals. An apparent competition was

demonstrated between Cu2+ and Fe2+ when 1.0 x 10-5M Cu2+ and 5.0 x 10-4M Fe2+

were present with 4.0 x 10-4M hydrous ferric oxide in an 02-free system.*

This result is shown in Figure 24 as a large depression in the isotherm compared to the same system in the absence of Fe(II). This demonstration is of

academic interest only, as an 02-free system is not applicable to seawater

but it does indicate the possibility of competition if combined trace metal

concentrations are sufficiently high.

* The N2 used to exclude C02 from the system was also passed through a trap

containing Cu metal in NH40H solution to remove 02. a second trap containing

HCl to remove NH3 and a water trap to remove NH4Cl and HCl.

39
6 7 PH 4 5

10

20

100 CUT = 1X 10-sM ~.

•

• Fe T = I X 10-3 M

90 t::. FeT I X 10- = 11 M

0 FeT = 4x10- 4 M

80 x 5x10- FeT = 5 M

0 FeT = I X 10-5 M

70

"U

Q.l

.c 60

0

II)

"U

« 50

+

N

:J

U 40

0

30

Fi~re 20. 2+ Per cent eu adsorhed as a function of pH for various concentrations

of hydrous ferric oxide in S.O.P.

40
100.------------------------,

90

I X 10- • Fer = 3 M

6 Fer = I x 10-4 M

80 •

0 Fer =5x 10-5M .--

70

"0

(J)

.c 60

\""

0

(I)

"0

<! 50

+

C\I

a.. 40

0

30

20

10

4 5

pH

6 7

Figure 21. Per cent Pb 2+ adsorbed as a function of pH for various

concentrations of hydrous ferric oxide in S.O.W.

41
50

40

30

•

0-4 FeT =I x I

x

20

10

4 5

pH

6 7

Figure 22. S 2+ Reduction of data from isotherms for 1. 0 x 10- M Cu on various

concentrations of hydrous ferric oxide in S.Q.W. to r.u2+ I

2+ adsorbed

CUf Fe- as a function of pH. ree T

42
50,----------------------------,

x

• Fe = 5)( 10-5 M T

40 Fe = 1)( 10- 0

4 M

T

)( Fe = I )( 10-3 M T

PbT = I x 10-5 M

30

20

10

4 5

pH

6

o

7

Pb2+ Fe as a function of pH. free T

Figure 23. 5 2+ Reduction of data from isotherms for 1.0 x 10- ~! Ph on various

concentrations of hydrous ferric oxide in S.O.W. to P~:+ I adsorbed

43
SECTION 4

DISCUSSION

The mechanism by which hydrous metal oxides remove trace metals from

solution has been postulated by various workers to be adsorption, chemisorption, ion exchange, surface complexation and coprecipitation. Several simple,

well-defined experimental systems yielded data which could be fitted to

equations derived for models based on these mechanisms.

The electrical double layer theory and numerous modifications treat the

phenomenon of electrostatic adsorption of ions on a charged surface. James

and Healy added a term to the established double layer treatment to explain

the relationship between hydrolysis of metal ions and their specific adsorption (chemisorption). The surface complexation model proposes a chemical

reaction between surface sites and metal ions in solution for which the

equilibrium constants are modified by the surface charge. All of the mathematical models require that the data can be reduced to an apparent constant

for a given pH and a given background electrolyte according to a mass law

expression

(1)

As the reduction in Figs. 14, 16, 18, 22 and 23 demonstrates, the reaction of Cu2+ and Pb2+ with amorphous ferric oxide is, to a first approximation

proportional to the free metal ~oncentration in solution and to the amount

of oxid~ over the range 5 x 10-~ to 4 x 10-4M metal and 5 x 10-5M to

1 x 10-~ oxide.

To reduce the data further to test equation 1, requires an estimate of

the total number of reactive sites on the oxide in order to calculate

{=M-O-}. The acid-base titration data are of little help for this purpose as

not all of the acid-base sites appear to be titrated in the pH range of

interest. An estimate of total reactive sites based on such data would result

in orders of magnitude underestimation of the removal of Cu2+ and Pb2+ by

the iron. Upper and lower limits for the number of reactive sites can be

obtained directly from the adsorption data. The Pb2+ data (Figure 19)

provides the lower limit of 0.8 site/Fe; while an upper limit of 1.5 site/Fe

is implied from the Fe(II)-Cu competition experiment (Figure 24). This is

consistent with the estimate obtained by Davis (36) of 0.87 site/Fe on the

basis of Yates (38) data obtained from rapid tritium exchange experiments.

Fi~~re 25 shows the reduction of the data using a value of 1 site/Fe for the

Cu2+ isotherms in seawater for all values of Cut and FeT. It should be noted

44
7

o

6

pH

4 5

10

20

100 • X 10- 5 M Cu 2 +only

0 1 X 10- 5 M CU 2 + + 6x1O- 4 M FeC")

90 Fe = 4x1O- 4

T M /0_0

80

•

"0

Q,) 60 .c

0

(/)

"0

<I 50

+

N

u 40 •

0

30

:rigure 24. -5 2+ -4 Depression of the isother"1 for 1.0 x 10 rt eu on 4.0 x 10 M

-4 2+ hydrous fer~ic oxide i~ S.O.W. in the ~resence of 6.0 x 10 M Fe .

45
both that the data are consistent with the mass law expression and that since

all the isotherms were below saturation, this is not a sensitive test of the

value chosen for total reactive sites. Although the Pb2+ data would provide

a more sensitive test, their scatter renders such reduction meaningless.

The curve in Figure 25 provides a convenient means to predict Cu2+ removal by

the iron oxide for any combination of concentrations and pH.

The very complex system of amorphous ferric oxide in seawater behaves in

a remarkably simple way with respect to Cu2+ and Pb2+ removal. There is no

effect of ionic strength, no effect of background electrolyte (except for the

expected Cl- effect on Pb2+), no effect of aging, no difference whether the

iron is precipitated in the presence or absence of the metal. Such simplicity

was not expected on the basis of the various mechanisms that have been proposed.

The expected effect of ionic strength depends to some extent on the model

being considered. The ionic strength enters into electrical double layer

calculations and affects the coulombic interactions between the charged

surface and the adsorbing ions by changing the thickness of the double layer.

The parameter 11K, which can be conceived of as the distance between the

plates of a hypothetical electrical condenser representing the electrical

double layer is 10.0 nm in 1.0 x 10-~ NaCl and reduces to 0.36 nm in 0.7M

NaCl (39). The collapse of the electrical double layer at high ionic strength

decreases the diffuse layer potential. What effect this has on adsorption

of trace metal cations, which are specifically adsorbed against the electrostatic repulsion of a positively charged surface, is unclear.

In all cases, however, the coulombic interactions in the adsorption process can only play a role for reactions resulting in a net change of charge.

The remarkable consistency of the isotherms for various ionic strengths

invites a speculation that the reaction of iron oxide with Cu2+ and Pb2+

results in no net change in charge in the pH range of interest. For example,

two protons may exchange for one metal ion. This is reinforced by the aging

study which demonstrates a clear change in charge (evidenced by a change in

uptake of acid by the solid) with no concomitant change in the isotherm.

Britton (37) in extensive studies of the precipitation of hydroxides

found that in general the amount of base needed to precipitate a metal

hydroxide did not correspond to the computed stoichiometric amount. He concluded that other anions were incorporated into the solid forming what he

calls "basic salts". Such non-stoichiometric precipitation would lead to the

observed batch to batch variations, but it also implies a possible effect of

the background electrolyte on the nature of the iron oxide. The effect of

the background electrolyte on the formation of the crystalline ferric oxides

has been reported (12). Whatever effect the background electrolyte may have

on the nature of the amorphous iron oxide, however, it is not reflected in

the isotherms for Cu2+ and Pb2+

+ It is puzzling that the release of 1.3 eqlmole of H IFe over 24 hours of

aging is not accompanied by any change in the Cu2+ isotherms. There is a

strong implication that the protons lost in aging are not the exchangeable

protons on the reactive sites. This may be a consequence of different

46
CJl

o

5

4

3 0

o

0 0 0

o

o

o

0:10oo

o 0

o 00

o

o ~~ 00

o 00

o <X> 0;: 0

o 000

o Q) / 0

o o¥'%o 8

o 0

00 .t6 0

o /~ o () 0

0/0000

Q) /() 0

o~ 0

o 0

o

o o

o

6 7 pH 5

2L-- L..-__---JL..-__---JL..-__---JL..-__---'L..-__---'1---....I

4

Figure 25. Log K as a function of t>1I for all Cu:>'+ data where:

Cu2+

adsorbed K - 2+ - ----------2+-- --- - (Cu--Cu d b d)(~e--Cu d b d) a sor e a gor e

47
coordination of the oxygen atoms or of geometrical inaccessability of some

of the oxygen atoms. Dousma and de Bruyn (6) discuss the polymerization of

Fe(III) oxide and note that in addition to the olation processes involving

-OH addition to iron oxide polymers, oxolation also occurs. This process

is represented by

+ (n-l)+

(1)

The oxo bridges thus formed react very slowly with acid. Oxolation would

explain the release of acid over the first 24 hours. The oxygen atoms

involved in oxolation are already bound to two iron atoms and are not likely

to be the reactive sites. That the acid-base sites with rapidly exchangeable

protons and the coordination sites for Cu2+ are indeed the same is supported

by the dependence of the apparent adsorption constant (K) on pH. Note that

from pH 6 to 7 the apparent constant increases by an order of magnitude

(Figure 15).

The capacity of amorphous hydrous ferric oxide for metal ions makes the

meaning of the terms "surface" and "adsorption" somewhat arbitrary. It

becomes necessary to provide a more satisfactory image of the system than

that of an interface separating two semi-infinite phases. One is clearly

not concerned with the interface between the visible precipitate and the bulk

solution, but rather the interface between some microstructure and its

immediately surrounding water. Amorphous hydrous ferric oxide has been

described as "amorphous, randomly crosslinked aggregates containing large

and indefinite amounts of water" (40). This description calls to mind the

structure of a swollen ion exchange resin which is permeable to hydrated ions.

The ions are free to diffuse throughout the structure and are not restricted

to external "surface" sites.

The formation of hydrous ferric oxide proceeds from the hydrolysis of

the hexahydrated aquo ion

to formation of the dimer

which gives rise to longer straight chain polymers as the pH is increased.

The solid ferric oxide begins to precipitate around pH 3. In the time frame

of the experiments reported here, the hydrous ferric oxide reaches what

appear to be two equilibrium stages, one at about two hours and another at

about 24 hours. The first corresponds to separation of the solid in its most

hydrated state. The second occurs after the polymers have completed fast

olation and oxolation processes. It is likely that a third stage follows in

which there is very slow dehydration of the solid with increasing cross48
linking and crystallinity. Byrne and Kester (13) found evidence for increasing crystallinity after very long aging periods for amorphous ferric oxide.

The loose, highly hydrated structure of the ferric oxide used in this

work readily accommodates foreign ions which become incorporated into the

solid as they hydrolyze. A loss of H+ accompanies the formation of -OH and

-0- bridges. The observed effect of chloride is a consequence of the formation of complexes with the metal ions which are not incorporated into the

solid. Rudnev et al. (41) found a similar effect of NH3 on copper coprecipitation with ferric hydrous oxide.

In addition to accounting for the high capacity of the iron for metals,

this image also provides an explanation for the similarity between isotherms

obtained by precipitation of the iron in the presence and absence of Cu2+

It is also consistent with what is known regarding the formation of metal

ferrites (42). The amorphous complex oxides, such as those obtained in this

work, are precursors of the crystalline ferrites. Adamovich et al. (43) examined the solid resulting from the coprecipitation of Fe and Cu by X-ray

diffraction and thermogravimetry and found evidence for a single solid phase

rather than a mixture of iron and copper oxides. Heating the solid produced

a compound whose X-ray diffraction pattern was suggestive of copper ferrite.

Gmelin (44) documents the preparation of copper ferrite by heating a

coprecipitated iron and copper oxides. When the oxides are precipitated

separately, mixed, and heated only CuO and Fe203 are formed.

The results of this study show that despite its complexity, the amorphous

iron oxide in seawater system behaves rather simply. The removal of metal

can be approximately predicted under all concentration and electrolyte

conditions by a constant which is only a function of pH. This pH dependence

can presumably be accounted for in a numerical model by adjusting the acidbase chemistry of the solid which is difficult to quantitate experimentally.

Although the mathematics appears to be similar to the surface complexation

model for adsorption on crystalline metal oxides, the reaction between

amorphous iron oxide and metals should not be viewed as a macroscopic surface

phenomenon. In many ways, amphorous ferric oxide can be viewed as more

closely analogous to polyelectrolytes than to crystalline solids such as

quartz.

49
REFERENCES

1. Krauskopf, Konrad B. Geochim. et Cosmochim. Acta, 9:1, 1956.

2. Goldberg, E.D. J. Geology, 62:249, 1954.

3. Goldschmidt, V.M. J. Chem. Soc., 1937:655, 1937.

4. Forbes, E.A., A. Posner, and J.P. Quirk. J. Colloid Interface Sci.,

49(3), 1974.

5. Yates, D.E., and T.W. Healy. J. Colloid Interface Sci., 52(2):222, 1975.

6. Dousma, J., and P.L. deBruyn. J. Colloid Interface Sci., 56(3):527, 1976.

7. Ko1thoff, I.M., and B. Moskovitz. J. Phys. Chem., 41:629, 1937.

8. Parks, G.A., and P.L. deBruyn. J. Phys. Chem., 66:967, 1962.

9. Atkinson, R.J., Posner, A.M., and J.P. Quirk. J. Phys. Chem., 71(3),

1967.

10. Gadde, R.R., and H.A. Laitinen. Environ. Lett., 5(4):223, 1973.

11. Ellis, J., R. Giovano1i, and W. Stumm. Chimia, 30:194, 1976.

12. Murphy, P.J., A.M. Posner, and J.P. Quirk. J. Colloid Interface Sci.,

56(2):270, 1976.

13. Byrne, R.H., and D.R. Kester. Marine Chem., 4:255, 1976.

14. Stumm, W. Personal communication.

15. Kurbatov, J.D., J.L. Ku1p, and E. Mack, Jr. J. Am. Chem. Soc., 67:1923,

1945.

16. Gadde, R.R., and H.A. Laitinen. Anal. Chem., 46(13):2022, 1974.

17. McNaughton, M.G., and R.O. James. J. Colloid Interface Sci., 47(2):431,

1974.

18. O'Connor, T.P., and D.R. Kester. Geochim. et Cosmochim Acta, 39:1531,

1975.

19. Forbes, E.A., A. Posner, and J.P. Quirk. J. Colloid Interface Sci.,

50
49(3), 1974.

20. Matijevic, E. In: Principles and Applications of Water Chemistry,

Faust and Hunter.

21. James, R.O., and T.W. Healy. J. Colloid Interface Sci., 40(1):42, 1972.

22. Stanton, D.A., and R. duT. Burger. Geoderma, 1(1):13, 1967.

23. Vuceta, J. Ph.D Thesis, California Institute of Technology, 1976.

24. Brunauer, S., P.H. Emmett, and E. Teller. J. Am. Chem. Soc., 60:309,

1938.

25. James, R.O., P.J. Stig1ich, and T.W. Healy. Faraday Disc. Chem. Soc.,

59:142, 1976.

26. Kurbatov, J.D. J. Phys. Chem., 36:1241, 1932.

27. Breeuwsma, A., and J. Lyklema. J. Colloid Interface Sci., 43(2), 1973.

28. Kurbatov, J.D., J.L. Ku1p, and E. Mack, Jr. J. Am. Chem. Soc., 67:1923,

1945.

29. Lyk1ema, J. Croatica Chem. Acta, 43:249, 1971.

30. Berube, Y.G., and P.L. deBruyn. J. Colloid Interface Sci., 27(2):305,

1968.

31. Breeuwsma, A., and J. Lyk1ema. Disc. of Faraday Soc., 52:324, 1971.

32. Yates, D.E., S. Levine, and T.W. Healy. J. Chem. Soc., 70:1807, 1974.

33. Stumm, W., H. Ho1h, and F. Da1ang. Croatica Chem. Acta, 1976.

34. FWPCA Methods for Chemical Analysis of Water and Wastes, U.S. Dept. of

Interior, 1969. p. 240.

35. Morel, F.M.M., J.C. Westall, J.G. Rueter, and J.P. Chap1ick. Technical

Note #16, Ralph M. Parsons Laboratory, M.I.T., Cambridge, Massachusetts,

1975.

36. Davis, J.A. Ph.D Thesis, Stanford University, 1977.

37. Britton, H.T.S. Hydrogen Ions, Vol. II. D. Van Nostrand and Co., Inc.,

Princeton, New Jersey, 1956.

38. Yates, D.E. Ph.D Thesis, University of Melbourne, 1975.

39. Parks, G.A. Adsorption in the Marine Environment. In: Chemical

Oceanography, J.P. Riley and G. Skirrow, eds. Academic Press, London,

1975.

51
40. Walton, H.F. Principles and Methods of Chemical Analysis. PrenticeHall, Inc., Englewood Cliffs, New Jersey, 1964.

41. Rudnev, N.A., G.I. Malofeeva, N.P. Andreeva, and T.V. Tikhonova. Zhurnal

Anal. Khimmii, 26(4):697, 1971.

42. Hilpert, S., and A. Wille. Z. Physik. Chem., B-18:29l, 1932.

43. Adamovich, T.P., V.V. Sviridov, and A.D. Lobanok. Dokl. Akad. Nauk.

Belorussk. SSR, 8(5):312, 1964.

44. Gmelin's Handbuch der Anorganischen Chemie. Kupfer Teil B., Lief. 3,

System Num. 60:1276, Verlag Chem., 1965.

52
TECHNICAL REPORT DATA

(Please read IRr.lTUctions on the reverse before completing)

1. REPORT NO.

1

2

.

3. RECIPIENT'S ACCESSION-NO.

EPA-600/3-80-011

4. TITLE AND SUBTITLE 5. REPORT DATE

ADSORPTION OF TRACE METALS BY HYDROUS FERRIC January 1980 issuinQ date

OXIDE IN SEAWATER 6. PERFORMING ORGANIZATION CODE

7. AUTHOR(S) B. PERFORMING ORGANIZATION REPORT NO.

Francois Morel and K. C. Swallow

9. PERFORMING ORGANIZATION NAME AND ADDRESS 10. PROGRAM ELEMENT NO.

Ralph M. Parsons Laboratory for Water Resources and lBA8l9

Hydrodynamics, Department of Civil Engineering 11. CONTRACT/GRANT NO.

Massachusetts Institute of Technology R803738

Cambridge, Massachusetts 02139

12. SPONSORING AGENCY NAME AND ADDRESS 13. TYPE OF REPORT AND PERIOD COVERED

Environmental Research Laboratory - Narragansett, R.1. Final

Office of Research and Development 14. SPONSORING AGENCY CODE

U.S. Environmental Protection Agency EPA/600/05

Narragansett, Rhode Island 02882

15. SUPPLEMENTARY NOTES

----

16. ABSTRACT

The adsorption of trace metals by amorphous hydrous ferric oxide in seawater

is studied with reference to simple model systems designed to isolate the factors

which may have an effect on the isotherms. Results show that the complex system

behaves in a remarkably simple way and that the data obtained under various

conditions of total metal concentration and total oxide concentration can be reduced

to an apparent reaction constant, K, which is a function of pH only. The high capacity of the oxide for trace metals renders the concept of a surface reaction

useless to explain the uptake of metals. A physical picture of the oxide as a

swollen hydrous gel permeable to hydrated ions is presented.

17. KEY WORDS AND DOCUMENT ANALYSIS

a. DESCRIPTORS b.IDENTIFIERS/OPEN ENDED TERMS c. COSATI Field/Group

Trace elements Modelling 06/F

Metalic Wastes

18. DISTRIBUTION STATEMENT 19. SECURITY CLASS (This Report) 21. NO. OF PAGES

RELEASE TO PrySLIC UNCLASSIFIED 63

20. SECURITY CLASS (This page) 22. PRICE

UNCLASSIFIED

EPA Form 2220·1 (9·73)

53 "" us GOVERNMENT f'lINTING OfFICE. 111IO -651-146/5546

TD427.T7S93

Swallow, K.C. (Kathleen C.).

Adsorption of trace materials

by hydrous ferric oxide in

seawater.

14437687
UIiNIeII_

hIlir Agency ......."'~

OffIcial Business

Penalty for Private Use. $300

r...................'............... c....

·Cincinn.ti OH 452'88

Please melte all necessary changes on the above label,

detach or coPY. and return to the addreSt; ,n the upper

left-hand corner

If yOu do nol WISh 10 recelYtllhese reportl CHECK HEilE 0,

delach, or copy th.s cover. and return to the address 'n the

upper leh· hand corner

EPA-600/3-80-011

fee. Poaage·and ".ua _(lJ E'.1Vlronmental .' Protecllon

Agency >

EPA:,335.. . ..

Spec.., FOllflh·CI""1 Rate

Book