NUREG/CR-6915

 LA-UR-05-4881

Aluminum Chemistry in a Prototypical

Post-Loss-of-Coolant-Accident, 

Pressurized-Water-Reactor Containment

Environment

Los Alamos National Laboratory

U.S. Nuclear Regulatory Commission

Office of Nuclear Regulatory Research

Washington, DC 20555-0001
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 NUREG/CR-6915

 LA-UR-05-4881

Aluminum Chemistry in a Prototypical

Post-Loss-of-Coolant-Accident,

Pressurized-Water-Reactor Containment

Environment 

Manuscript Completed: August 2006

Date Published: December 2006

Principal Investigator: M. Klasky

Prepared by

M. Klasky, J. Zhang, M. Ding, and B. Letellier (Los Alamos National Laboratory)

D. Chen and K. Howe (University of New Mexico) 

Los Alamos National Laboratory

Los Alamos, NM 87545

University of New Mexico

Department of Civil Engineering

Albuquerque, NM 87110

T.Y. Chang, NRC Project Manager

Prepared for

Division of Fuel, Engineering and Radiological Research 

Office of Nuclear Regulatory Research

U.S. Nuclear Regulatory Commission

Washington, DC 20555-0001

NRC Job Code N6285 
NUREG/CR-6915 has been reproduced 

from the best available copy. 
iii

Aluminum Chemistry in a Prototypical Post-Loss-of-CoolantAccident, Pressurized-Water-Reactor Containment Environment 

ABSTRACT 

An analysis of the integrated chemical effects tests (ICET) experiments has been 

performed by a comprehensive examination of both the test solutions and precipitates. In 

addition, a comprehensive review of the literature has been performed to assist in 

explaining the behavior of aluminum in alkaline solutions. The objective of this analysis 

was to elucidate the behavior of precipitate that formed when the ICET Tests 1 and 5 

solutions were allowed to cool so that the behavior of other solutions with different 

conditions, i.e., pH, temperature, etc., could be predicted throughout the pressurizedwater reactor following a loss-of-coolant accident (LOCA). This examination included 

supplemental analytical measurements using x-ray diffraction, 27Al and 11B nuclear 

magnetic resonance for both liquid and solid states, and quasi-elastic light-scattering 

measurements. Surrogate solutions were developed and compared with the analytical 

measurements of the ICET Tests 1 and 5 solutions. Finally, the characterization of the 

particle sizes and corrosion properties, including the corrosion mechanism and the 

corrosion rate of aluminum under LOCA conditions, has been elucidated. The current 

study should allow for the development of a head-loss correlation using the existing cake 

filtration theory, which could be used in conjunction with a corrosion model to predict 

system performance following a LOCA. 
iv

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v

FOREWORD 

The U.S. Nuclear Regulatory Commission (NRC) established Generic Safety Issue (GSI)-191, 

“Assessment of Debris Accumulation on PWR Sump Performance,” to assess the issue of 

potential blockage of pressurized-water reactor (PWR) sump screens during a loss-of-coolant 

accident (LOCA), and to determine the need for further regulatory actions. In this context, the 

definition of “debris” includes material (such as insulation and paint coatings) damaged by 

direct jet impingement from a LOCA, as well as latent dust and dirt that may be present in the 

containment buildings. During a review of GSI-191 in 2003, the NRC’s Advisory Committee 

on Reactor Safeguards (ACRS) raised a concern about “chemical effects.” Specifically, certain 

resultant products attributable to chemical interactions between emergency core cooling 

system (ECCS)/containment spray water and exposed materials (such as metal surfaces, paint 

chips, and insulation debris) could form another source of debris that has the potential to 

impede the performance of ECCS recirculation following a LOCA at a PWR plant. 

To address this concern, the NRC and the nuclear energy industry jointly sponsored an 

integrated chemical effects test (ICET) program, and the NRC is publishing the results of that 

program in the six-volume NUREG/CR-6914. As noted in that report, ICET Tests #1 and #5 

raised questions regarding the behavior of aluminum in sump pool solutions, and especially 

the behavior of aluminum-related chemical products that formed when ICET Tests #1 and #5 

solutions were allowed to cool. In order to answer these questions, the NRC sponsored a 

follow-on study, performed at Los Alamos National Laboratory, which included an analysis of 

both the test solutions and aluminum-related chemical products of ICET Tests #1 and #5, as 

well as a comprehensive review of the literature to improve our understanding of the behavior 

of aluminum in sump pool solutions. That follow-on study is the subject of this report. 

In particular, this report provides information about the characterization of particle sizes and 

corrosion properties of aluminum under LOCA conditions. It also provides information that 

will allow extrapolation of ICET behavior to predict the behavior of aluminum under the 

various pH and temperature conditions that might exist in PWR plants following a LOCA. 

This information will be useful to the nuclear utility industry and the NRC staff in assessing 

aluminum chemical effects on sump performance in a post-LOCA PWR containment 

environment. 

The findings and conclusions documented in this report were presented and discussed at 

several external Chemical Effects Peer-Review Group meetings and ACRS Committee 

meetings. As such, they have been tested for consistency with the professional experience of 

multiple individuals and compared to recent data generated at other national laboratories. 

_______________________________ 

Brian W. Sheron, Director 

Office of Nuclear Regulatory Research 

U.S. Nuclear Regulatory Commission 
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vii

CONTENTS

Page 

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

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

FIGURES ....................................................................................................................... viii

TABLES ......................................................................................................................... xi

EXECUTIVE SUMMARY ............................................................................................. xiii

ACKNOWLEDGMENTS................................................................................................ xv

ABBREVIATIONS ........................................................................................................ xvii

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

2. SUMMARY OF THE BEHAVIOR OF ALUMINUM OBSERVED 

DURING ICET TESTING.............................................................................. 3 

2.1. Analysis on Aluminum Concentration of ICET Tests 1 and 5................ 3 

2.2. Corrosion of ICET Metal Coupons ............................................................. 5 

2.3. ICET Turbidity and Total Suspended Solids (TSS) Measurements ..... 8 

2.4. ICET Hydrogen Concentration Measurements ...................................... 10 

2.5. ICET Viscosity Measurements.................................................................. 11 

2.6. ICET Calcium and Silica Concentrations ................................................ 12 

2.7. Analysis of Tests 1 and 5 Precipitates..................................................... 14 

2.8. TEM of Aluminum Precipitate.................................................................... 19 

2.9. Particle Size Measurements of Test 1 Solutions ................................... 21 

3. OVERVIEW OF OUTSTANDING ISSUES .............................................. 25

4. ALUMINUM-PHASE BEHAVIOR .............................................................. 27

4.1. Hydroxide Phases....................................................................................... 27 

4.2. Phase Transitions ....................................................................................... 28 

4.3. General Cation/Anion Effects on the Aluminum Phase ........................ 35 

5. ALUMINUM-BORON COMPLEX .............................................................. 41

5.1. Boron Behavior in the Solution ................................................................. 41 

5.2. Boron Adsorption ........................................................................................ 43 

6. PARTICLE SIZE DISTRIBUTION.............................................................. 49

7. ALUMINUM SOLUBILITY IN ALKALINE SOLUTION............................ 53

7.1. Predictions of Aluminum Solubility in Alkaline Solutions ...................... 53 

7.2. Effect of Organic Complexing Agents on Aluminum Solubility ............ 56 

8. ALUMINUM CORROSION IN ALKALINE SOLUTION .......................... 61

8.1. Corrosion Mechanism ................................................................................ 61 

8.2. Corrosion Rate ............................................................................................ 63 

8.3. Corrosion Inhibitor....................................................................................... 69 

9. BENCHMARK EXPERIMENTS................................................................. 73

9.1. Test 1 Redissolution Test .......................................................................... 73 

9.2. Solubility of Aluminum-Based Surrogates............................................... 74 

9.3. Effects of the Boron in the Aluminum-Alkaline Solution ....................... 74 

9.4. Particle Size Measurements of Surrogates at 60°C.............................. 76 

9.5. Inhibitor Passivation.................................................................................... 78 

10. CONCLUSIONS........................................................................................... 83

REFERENCES............................................................................................................... 85
viii

FIGURES 

Page 

Figure 1. Aluminum concentration versus time........................................................ 4 

Figure 2. ICET Benchtop pH comparison. ................................................................ 4 

Figure 3. Annotated SEM image magnified 1000 times for a Test 4, Day-30, 

submerged aluminum coupon sample...................................................... 6 

Figure 4. EDS counting spectrum for the deposits (EDS3) on the coupon 

surface shown in Figure 3........................................................................... 7 

Figure 5. SEM image magnified 1000 times for a Test 1, Day-30, submerged 

aluminum coupon sample........................................................................... 7 

Figure 6. EDS counting spectrum for the light spot (EDS1) on the coupon 

surface shown in Figure 5........................................................................... 8 

Figure 7. Turbidity versus time at 60°C..................................................................... 9 

Figure 8. Total suspended solids versus time.......................................................... 9 

Figure 9. Turbidity versus time at 23°C................................................................... 10 

Figure 10. Hydrogen concentration above tank versus time................................. 11 

Figure 11. ICET kinematic viscosity versus time at 60°C...................................... 12 

Figure 12. ICET kinematic viscosity at 25°C............................................................ 12 

Figure 13. Calcium concentration versus time. ....................................................... 13 

Figure 14. Silica concentration versus time. ............................................................ 14 

Figure 15. Day-30, high-volume sample counting spectrum (T1D30-EDS1) on 

the filtrate.................................................................................................... 15 

Figure 16. Day-15, Test 1 counting spectrum (EDS 1-2) for the surface on the 

film overlying the filter............................................................................... 16 

Figure 17. XRD of precipitate (damp, red; dried, blue). ......................................... 17 

Figure 18. XRD of precipitate sample (collected on 2/24/05, 1745 dried at 50°C).

 ...................................................................................................................... 17 

Figure 19. Test 1 separated precipitate collected on 2/24/05, 1745 bcl, dried at 

50°C. ............................................................................................................ 18 

Figure 20. Weight loss due to drying of ICET-separated precipitate (boehmite)... 

 ...................................................................................................................... 20 

Figure 21. Electron micrograph of the Day-15 filtered test sample magnified 

50,000 times. ............................................................................................. 20 

Figure 22. Electron micrograph of the Day-30 filtered test sample magnified 

50,000 times. ............................................................................................. 21 

Figure 23. TEM of the Day-15, Test 5 precipitate, magnified 4000 times........... 22 

Figure 24. TEM of the Day-30, Test 5 precipitate, magnified 8000 times........... 22 

Figure 25. Particle size (nanometer) distribution for the Test 1 “precipitate 

supernatant” solution at room temperature (aged 4 months). ........... 23 

Figure 26. Gibbsite and bayerite layer stacking patterns. ..................................... 28 

Figure 27. X-ray diffractograms of SDAHC after it has aged at 100% relative 

humidity (RH) and 25°C for 0 days (bottom), 62 days (second from 

bottom), 118 days (middle), 215 days (second from top), and 384 
ix

days (top). Diffraction bands associated with boehmite or bayerite are 

represented by dashed or solid lines, respectively.............................. 33 

Figure 28. X-ray diffractograms of SDAHC after it has aged at 100% RH and 

50°C for 0 days (bottom), 6 days (lower center), 36 days (upper 

center), and 78 days (top). Diffraction bands associated with 

boehmite are indicated by dashed lines.].............................................. 33 

Figure 29. Change in the O-H stretching frequency region of the IR spectrum of 

aluminum hydroxide gel during aging at 25°C for (A) 5 days, (B) 42 

days, (C) 54 days, (D) 61 days, (E) 69 days, and (F) 124 days........ 35 

Figure 30. Change in XRD pattern of aluminum hydroxide gel during aging at 

25o

C for (A) 57 days, (B) 65 days, (C) 71 days, and (D) 124 days... 35 

Figure 31. Possible pathways for polymerization of Al(OH)3. ............................... 36 

Figure 32. Electron transmission micrograph of aluminum adjuvant precipitated 

at pH = 4.5 in the presence of phosphate............................................. 38 

Figure 33. The 27Al nuclear magnetic resonance (NMR) of benchtop aluminum 

nitrate/boron/sodium hydroxide solution at pH = 9.5 at 60°C, 40°C, 

and 25°C after 4 weeks of aging. ........................................................... 42 

Figure 34. Spectra of (A) Al(OH)3 with boron (89 days old), (B) Al(OH)3 prepared 

without boron (13 days old), and (C) commercially prepared Al(OH)3.

...................................................................................................................... 44 

Figure 35. Absorbed boron as a function of pH at 25oC and 5oC......................... 45 

Figure 36. ATR-FTIR spectra of aqueous boric acid as a function of pH and total 

boron concentration. Each spectrum is the difference between the 

sample (boric acid in 0.1 M NaCl at each pH) spectrum and the 

reference (0.1-M NaCl solution) at the same pH. ................................ 46 

Figure 37. ATR-FTIR spectra of am-Al(OH)3 at (a) 25oC at pH = 6.8 and (b) 5oC 

at pH = 10.2. .............................................................................................. 47 

Figure 38. Particle size for aluminum in pH = 4.5 solution. ................................... 49 

Figure 39. Particle size versus pH for aluminum at ambient temperature.......... 50 

Figure 40. Particle growth dynamics in aluminum chloride solutions as a function 

of time for various OH/Al ratios at ambient temperature. ................... 51 

Figure 41. Solubility of (a) amorphous Al(OH)3, and (b) gibbsite [Al(OH)3] as a 

function of pH at 25°C. Also shown are lines indicating the solubility 

concentrations of Al3+ and individual hydroxyl complexes.................. 55 

Figure 42. Effect of KOAc on the potentiometric titration of 0.1 AlCl3 by 0.1 N 

NaOH. The molar ratio of acetate to aluminum was 0, 1.0, 2.0, and 

3.0................................................................................................................ 58 

Figure 43. Effect of K oxalate on the potentiometric titration of 0.1 AlCl3 by 0.1 N 

NaOH. The molar ratio of acetate to aluminum was 0, 0.5, 1.0, 1.5, 

2.0, and 3.0. ............................................................................................... 58 

Figure 44. Effect of K citrate on the potentiometric titration of 0.1 AlCl3 by 0.1 N 

NaOH. The molar ratio of acetate to aluminum was 0, 0.45, 1.0, 2.0, 

and 3.0........................................................................................................ 59 

Figure 45. Particle growth in the presence of citric acid. ....................................... 59 

Figure 46. Changes in film formation potential obtained from the pure aluminum 

rod specimen with time in 10-3 M NaOH solution at various applied 
x

anodic current densities of Ο , 5 mA cm-2; , 10mA cm-2; Δ, 20mA 

cm-2; • , 50mA cm-2; and , 100mA cm-2............................................... 62 

Figure 47. Corrosion rate as function of pH in systems (1) aluminum in 4 10− M 

KCl; (2) aluminum in 4 10− M KCl containing 4 10− M ethylenediamine; 

and (3) aluminum in 4 10− M KCl containing 3 10− M ethylenediamine. 

...................................................................................................................... 64 

Figure 48. The variation of weight loss of aluminum with exposure time at 30oC.. 

 ......................................................................................................... 65 

Figure 49. Corrosion rates based on weight loss shown in Figure 48................. 66 

Figure 50. Dependence of corrosion rate on the flow velocity.............................. 69 

Figure 51. Variation of inhibition efficiency with the concentration of some 

organic acides in 3 M NaOH. .................................................................. 71 

Figure 52. Potential-time curves for almasilium alloys immersed in 0.1 NaOH 

solution at 60oC, showing the effect of aluminate and “silicate” ions: 

curve a, 0.1 N NaOH alone; curve b, 0.03 M Si; curve c, 0.04 M Si; 

curve d, 0.04 M Si + 0.0003 M Al; curve e, 0.04 M Si + 0.0006 M Al; 

curve f, 0.04 M Si + 0.0012 M AI; and curve g, 0.04 M Si +0.0018 M 

AI.................................................................................................................. 71 

Figure 53. TGA for surrogate, where the red line is weight loss and the blue line 

is the derivative.......................................................................................... 75 

Figure 54. TGA for Test #1 precipitate. .................................................................... 75 

Figure 55. NMR solid-state spectrum of the precipitates....................................... 77 

Figure 56. Particle size distributions for aluminum/boron metal dissolution 

surrogate solution at 60°C after 8 hours................................................ 78 

Figure 57. Aluminum concentration (mg/L) versus time. ....................................... 79 

Figure 58. XPS spectra of Al 2p at 90 TOA on aluminum alloy coupons: (left) 

virgin aluminum, (right) aluminum soaked in Na2SiO3 solution. ........ 80 

Figure 59. The stability of a passivation layer of silicate with respect to high and 

low pH. ........................................................................................................ 81 
xi

TABLES 

Page 

Table 1. Experimental Conditions in the ICET Tests ............................................... 3 

Table 2. Mean Weight Gain/Loss Data for Submerged Coupons (g).................... 5 

Table 3. Elemental Composition (wt %) of Test 1, Day-30, High-Volume Filter 

Sample........................................................................................................... 15 

Table 4. Elemental Composition of Test 1 Film Material....................................... 16 

Table 5. Distribution of Main Elemental Components of Precipitates from Tests 

1 and 5 ........................................................................................................... 19 

Table 6. Structure of Precipitates, as Revealed by XRD ...................................... 29 

Table 7. Effect of Age on the pH and Crystalline Structure of Aluminum 

Hydroxide Prepared from 0.0378-M Aluminum Nitrate Solutions 

(Solution also 0.30 in Sodium Nitrate) ...................................................... 30 

Table 8. Characteristic Stretching and Bending Frequencies for Crystalline 

Aluminum Phases ........................................................................................ 34 

Table 9. Corrosion Rate of Aluminum in Static Alkaline Solution ........................ 65 

Table 10. Corrosion Rate on Various Aluminum Alloys in 4 KOH at 50oC ........... 68 

Table 11. Effect of Some Organic and Inorganic Additives on the Corrosion of 

Aluminum in NaOH Solutions .................................................................... 70 
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xiii

EXECUTIVE SUMMARY 

The Integrated Chemical Effects Test (ICET) program was conceived as a limited-scope 

exercise to simulate representative chemical conditions in pressurized-water-reactor 

(PWR) containment buildings following a postulated loss-of-coolant accident (LOCA). 

Principal objectives of the series included (1) monitoring long-term tests for the 

formation of chemical products that could adversely affect emergency-core-coolingsystem sump performance, (2) identifying the composition of adverse products, and 

(3) explaining the formation processes of adverse products. The enclosed stainless-steel 

test apparatus was designed to accommodate both submerged and suspended material 

samples that were introduced in proportion to representative inventories found in 

containment. Various debris sources included fiberglass and calcium-silicate insulation, 

crushed concrete, and common dirt. Heated 60°C solution was sprayed over suspended 

coupons for 4 hours, and the tank was circulated and monitored continuously for 30 days. 

While data were collected during Test 1, an extensive literature search was conducted to 

help explain the trends that were observed and additional bench-scale experiments were 

performed post-test to confirm the conclusions that were drawn regarding the formation 

of observed products. Test 5 shared many of the same attributes, including aluminum 

corrosion in a high-pH environment and complex interactions between sodium, 

aluminum, silicon, and boron. This report summarizes the findings of the literature search 

and benchscale experiments and synthesizes plausible explanations for the aluminum 

chemistry trends and chemical products that were observed in ICET Tests 1 and 5. 

A comprehensive examination of both the test solutions and precipitates from ICET Tests 

1 and 5 has been performed. In addition, a comprehensive review of literature and 

additional benchmark experiments of surrogates has been performed to assist in 

explaining the behavior of the aluminum present in the chemical solutions. The objective 

of this analysis was to elucidate of the behavior of precipitate that formed when ICET 

Tests 1 and 5 solutions were allowed to cool. This examination included supplemental 

analytical measurements using x-ray diffraction (XRD), 27Al and 11B nuclear magnetic 

resonance (NMR) (both liquid and solid state), and quasi-elastic light-scattering 

measurements. The results of this investigation permit extrapolation of the behavior of 

the ICET tests to predict the behavior of aluminum under different pH and temperature 

conditions that might exist throughout the PWR following a LOCA. Finally, 

characterization of the particle sizes and corrosion properties of aluminum under LOCA 

conditions have also been studied. These findings would facilitate development of a 

head-loss correlation using existing cake filtration theory, which could be used in 

conjunction with a surface corrosion model to predict system performance following a 

LOCA. 

Based on both inductively coupled plasma (ICP) analysis and examination of the ICET 

Tests 1 and 5 aluminum coupons, it was concluded that the aluminum weight loss from 

the metal coupons in these tests was consistent with the measured aluminum coupons. 

Furthermore, by examining the plateau of viscosity (at 23o

C), hydrogen production, and 

aluminum concentration, it was hypothesized that the plateau in the aluminum 
xiv

concentration was a result of passivation of the aluminum metal surface. The measured 

concentrations of aluminum in ICET Tests 1 and 5 solutions did not approach predicted 

solubility limits, and calculations confirmed that cooling of ICET Tests 1 and 5 solutions 

to ambient temperature would produce precipitation, as observed in practice. 

Examinations of the precipitates of ICET Tests 1 and 5 revealed that the precipitates, 

after being washed to ensure that sodium borates were not allowed to develop during the 

drying process, were largely composed of amorphous aluminum hydroxide, with a 

substantial quantity of boron adsorbed onto the surface. A survey of the literature 

suggested that the presence of the amorphous form of aluminum is expected because of 

the solution’s high concentration of anions, which have been shown to retard 

crystallization at temperatures below 60o

C. Thus, if amorphous-phase aluminum does 

form during cooling of the ICET Test 1 solution, as both XRD and TEM analyses would 

indicate, then it is not surprising that a significant amount of boron would also be found 

as a result of the adsorption of boron from solution. ICP results indicate that up to 35% of 

the boron from the initial solution may have been adsorbed onto the amorphous 

aluminum hydroxide precipitate. This degree of adsorption is certainly feasible. In fact, 

data from Su (1995) at a pH of ~9.5 predict a boron adsorption of ~35%. A review of the 

literature and NMR measurements, which revealed complexation between aluminum and 

boron, corroborate the hypothesis that complexation was responsible for impeding the 

crystallization of aluminum compounds. Furthermore, the measured weight percentage of 

boron on the aluminum precipitate was consistent with previous predictions based on 

adsorption models. 

Aluminate ions ( − Al(OH) ) are the only stable form of al 4 uminum in an alkaline solution. 

The solubility is a function of solid hydroxide phases and increases with pH. It is noted 

that the presence of some organics and inorganics can increase the aluminum solubility. 

The solubility can also be affected by the particle size presented in the solution. 

Nanometer-sized particles are found in the ICET Tests 1 and 5 solutions, and the particle 

size distribution is affected by the hydroxide phases and the addition of organic 

compounds. To determine if colloids exist in ICET Tests 1 and 5 solutions, dynamic light 

scattering was used. Results of these measurements revealed that colloids did exist at 

ambient temperature with a mean radius of 65 nm. 

Aluminum corrosion in alkaline solutions is an electrochemical process. In a static 

solution, the corrosion rate decreases with exposure time. In a dynamic solution, the 

corrosion rates increase with increasing flow velocity in mass-transfer-controlled regimes 

and do not depend on flow velocity in actively controlled regimes. The corrosion rate is 

influenced by adding organic and inorganic inhibitors to the alkaline solution and by the 

addition of some metal elements into the aluminum alloy. Sodium silicates have been 

found to be an effective inhibitor, with inhibition efficiency almost 100%. Thus, in ICET 

Test 4, the corrosion of aluminum was inhibited by the dissolution of calcium silicate. 
xv

ACKNOWLEDGMENTS 

The authors would like to thank several individuals for their insightful comments and 

generous assistance in the preparation of this report. John Apps from Berkeley 

Laboratory and Paul Klein from the NRC staff reviewed early drafts for technical 

accuracy. Their comments helped greatly to improve the practical content and 

applicability of the document. Steve Chipera from Los Alamos National Laboratory 

(LANL) graciously donated his time and x-ray diffraction equipment to help identify the 

chemical compounds present in ICET Test #1 precipitates. His assistance provided key 

evidence supporting the phenomenological explanation of trends observed in the ICET 

chemical systems. Debra Wrobleski from LANL also provided no-cost access to in situ, 

constant-temperature, particle-size diagnostics that helped confirm the presence of 

colloidal structures in the chemical solutions. The entire ICET investigation team 

provided a continual sounding board to exchange hypotheses and insights. LANL editor 

Lisa Rothrock again provided professional, cheerful assistance to transform our rambling 

collection of notes into a publication-quality document. Without her effort, we would not 

have obtained the permissions necessary to release this report. Finally, gratitude is 

offered to NRC project manager Rob Tregoning for recognizing the value of this analysis 

effort and providing the opportunity to share our work. 
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xvii

ABBREVIATIONS 

ATR-FTIR Attenuated Total Reflectance—Fourier Transform Infrared Spectroscopy 

BA Benzoin Acid 

CMP Chemical Mechanical Planarization 

DFB Desferriferrioxamine B 

DLS Dynamic Light Scattering 

EDS Energy Dispersive Spectroscopy 

EDTA Ethylenediaminetetraacetic 

FA Fulvic Acid 

ICET Integrated Chemical Effects Tests 

ICP-AES Inductively Coupled Plasma—Atomic Emission Spectroscopy 

IR Infrared 

LANL Los Alamos National Laboratory 

LOCA Loss-of-Coolant Accident 

NMR Nuclear Magnetic Resonance 

NPSH Net Positive Suction Head 

NTU Nephelometric Turbidity Unit 

NUREG Nuclear Regulatory Commission Report 

PZC Point of Zero Charge 

PWR Pressurized Water Reactor 

RH Relative Humidity 

SAXS Small-Angle X-Ray Scattering 

SDAHC Spray-Dried Aluminum Hydroxycarbonate 

SEM Scanning Electron Microscopy 

TEM Transmission Electron Microscopy 

TGA Thermogravimetric Analysis 

TOA Takeoff Angle 

TSP Trisodium Phosphate 

TSS Total Suspended Solids 

XPS X-Ray Photoelectron Spectroscopy 

XRD X-Ray Diffraction 

XRF X-Ray Fluorescence 
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1

1. INTRODUCTION 

During a loss-of-coolant accident (LOCA), the chemical environment generated by the 

injection of chemical agents into the emergency-core-cooling-system water, which is 

used to sequester the iodine, may create a chemical environment that facilitates the 

corrosion of metals and/or generation of precipitate. This corrosion may exacerbate the 

sump-screen head loss in a pressurized-water-reactor (PWR) containment sump pool. The 

integrated chemical effects tests (ICET) demonstrated that the interaction of sodium 

hydroxide with exposed aluminum surfaces that may be present in PWR containment 

may release substantial quantities of aluminum into solution via corrosion.1

Consequently, in this report we analyze the behavior of aluminum following a LOCA. 

Furthermore, previous small-scale testing, in which aluminum was artificially generated 

in solution by introducing aluminum nitrate, indicated that these chemical products may 

exacerbate sump-screen clogging and consequently lead to an unacceptable net-positivesuction-head (NPSH) margin.2

 Although ICET demonstrated the feasibility of generating 

substantial quantities of aluminum in solution, which on cooling could lead to the 

formation of gelatinous material, a detailed understanding of this material, including the 

conditions necessary for its production, was not provided in the ICET reports. 

Furthermore, additional analyses are needed to extrapolate these results to conditions 

outside the five ICET tests and for the ultimate development of filtration models that are 

needed to predict head loss. 

The objective of this report is to present the results of an extensive literature review and 

supplemental bench-scale testing of aluminum/sodium hydroxide solutions for a more 

detailed understanding of the generation of aluminum compounds (corrosion), their 

solubility properties, chemical speciation of these products, and particle size distributions 

in representative sodium hydroxide systems that are characteristic of those that may be 

encountered following a LOCA. Such an analysis may be used to allow for a 

determination of both the total quantity of chemical product formed under varying 

temperature and pH conditions. 
2

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3

2. SUMMARY OF THE BEHAVIOR OF ALUMINUM OBSERVED 

DURING ICET TESTING 

ICET testing was performed to investigate the possibility of chemical products being 

formed under conditions that might exist following a LOCA, which may compromise the 

filter screen and consequently prevent recirculation of cooling water to the reactor. Five 

tests were performed to address differences in both the chemical solutions and fiber 

insulation present throughout the PWR industry. Table 1 presents a summary of the five 

tests that were performed for the ICET project. Additional detail on the tests may be 

found in United States Nuclear Regulatory Commission report NUREG/CR-6914.1

Table 1. Experimental Conditions in the ICET Testsa,b 

Insulation Debris 

Volumef Test 

# 

Boronc

Added 

(mg/L) 

NaOH 

Added 

(mg/L) 

TSPd

Added 

(mg/L) 

Test pH 

Range 

pH 

Buffering 

Agente Fiberglass Calcium 

Silicate 

1 2800 7677 - 9.3–9.5 Borate 100% - 

2 2800 - 4000 7.1–7.4 Phosphate 100% - 

3 2800 - 4000 7.3–8.1 Phosphate 20% 80% 

4 2800 9600 - 9.5–9.9 Borate 20% 80% 

5 2400 - - 8.2–8.5 Borate 100% - 

a Temperature was maintained at 60°C (±3°C) in all tests. 

b The following chemicals were also added to the solution: LiOH = 0.7 mg/L and HCl = 100 mg/L in 

Tests 1 through 4. LiOH = 0.3 mg/L and HCl = 42.8 mg/L in Test 5. Also, 63.7 g of latent debris and 

21.21 g of concrete dust were added to the mixture prior to each test initiation. c The required boron concentration of 2800 mg/L was added as 16,000 mg/L H3BO3 in Tests 1 through 4. 

To attain the required boron concentration of 2400 mg/L for Test 5, 18,500 mg/L of Na2B4O7·10H2O 

(sodium tetraborate) in 143 gal. of water was added to 16,000 mg/L of H3BO3 in 107 gal. of water.

d Trisodium phosphate (TSP) = Na3PO4·12H2O. e Although other additives establish system pH on a wide scale, the buffering agent keeps pH relatively 

stable by resisting minor fluctuations from its pKa. f All tests consisted of 0.137 ft3

 of insulation represented per cubic foot of water. For each test, the 

insulation was present as fiberglass and/or calcium silicate material with the proportion of volume 

shown in the table.

2.1. Analysis on Aluminum Concentration of ICET Tests 1 and 5 

Throughout the ICET test series, measurements were taken of elemental solution 

concentrations, pH, hydrogen production, precipitate formation, turbidity, and total 

suspended solids. The results of the ICET tests have revealed that aluminum-based 

chemical products can be generated via the corrosion mechanism in solutions in which 

sodium hydroxide is added. During the first 25 days of the test, both filtered (pore size 

0.7 μm) and unfiltered samples were taken for analysis. Because of nondistinct 

differences between the filtered and unfiltered samples, only unfiltered samples were 

taken to be analyzed after Day 25 and for the duration of the test. From previous 
4

experimental work, it should be noted that filters of this size are incapable of removing 

the very small nanoparticles that are known to exist within aluminum solutions.3 Figure 1

presents the aluminum concentration found in the solutions for ICET Tests 1 and 5. From 

this figure we observe that the results from inductively coupled plasma—atomic emission 

spectroscopy (ICP-AES) analyses indicated that in Test 1, the aluminum concentration 

increased from nearly undetectable levels to ~350 mg/L after ~10 days, at which point the 

concentration reached a plateau. Likewise, we may observe that in Test 5, the aluminum 

concentration increased from nearly undetectable levels to ~50 mg/L after 20 days, at 

which point the concentration reached a plateau. Aluminum concentrations in Tests 2, 3, 

and 4 were below the levels of detection. Therefore, it may be concluded that aluminum 

was not present in solution to any appreciable degree in systems in which TSP was 

present. The pH is a major factor in determining the corrosion rate of aluminum, which 

will be discussed in Section 4. The changes of pH during ICET tests are shown in Figure 

2. The figure indicates that the pH remains almost constant during the tests. 

Figure 1. Aluminum concentration versus time. 

Figure 2. ICET Benchtop pH comparison. 
5

2.2. Corrosion of ICET Metal Coupons 

To determine if the measured concentrations of aluminum in Tests 1 and 5 were 

consistent with the observed weight loss of the aluminum coupons, the metal coupons 

were weighed after drying (no cleaning was performed before weighing). Weight 

loss/gain measurements, presented in Table 2, indicate that the submerged aluminum 

coupons experienced a significant weight change, which is consistent with the reduction 

and oxidation (redox) potentials of the metal coupons. The aluminum specimens in Test 1 

lost ~25% of their pre-test mass, whereas the carbon steel sample lost ~2% of its pre-test 

mass and the concrete sample gained <3% additional mass. The concrete weight gain 

may have resulted from retained water. Other samples either gained or lost much less 

than 1% of their pre-test mass. Likewise, the Test 5 coupons were weighed, and it was 

determined that only the submerged aluminum coupons experienced signification weight 

change: 3% of their weight. 

Table 2. Mean Weight Gain/Loss Data for Submerged Coupons (g) 

 Weight Changes (g) 

Coupon Type Initial Weight (g) #1 #2 #3 #4 #5 

Copper 1317.3 0.1 <0.1 0.3 0.2 -0.2 

Inorganic Zinc 1625.2 3.1 3.8 1.8 2.3 1.6 

Galvanized Steel 1054.83 0.0 28.6 15.0 0.3 0.1 

Aluminum 392.0 -98.6 -0.9 0.6 0.0 -11.2 

Uncoated Steel 1025.2 -23.3 1.4 -1.1 0.2 0.0 

Concrete 8586 233.0 240.7 180.5 239.6 225.9 

Based on the number of coupons and volume of water in the tank, the corrosion of the 

aluminum coupons would contribute 313 mg/L of aluminum to the tank solution in Test 1 

and 36 mg/L to the tank solution in Test 5. It should be noted that because of a buildup of 

corrosion products on the coupon surface, the weight loss of an aluminum coupon may 

underestimate the aluminum concentration in the solution. As a result, these values are 

fairly consistent with the measured aluminum concentrations, as shown in Figure 1. 

Therefore, it is unlikely that significant bulk precipitation occurred. Additional support 

for this conclusion is provided by supporting measurements, i.e., hydrogen generation, 

room temperature turbidity, and kinematic viscosity measurements presented later in this 

section. In addition, analysis of the passivation of aluminum surfaces, Section 4, and 

solubility of aluminum, Section 5, provide further substantiation of this hypothesis. 

As previously mentioned, the lack of aluminum present in solution in Test 4 (see Table 2) 

was initially surprising because the test conditions, with respect to pH, were similar to the 

initial test conditions of Test 1. That is, it was anticipated that the pH conditions in Test 4 

would facilitate corrosion of the aluminum, just as in Test 1, and that aluminum 

concentrations comparable to that of Test 1 would be observed. However, a primary 
6

difference between the two tests was the insulation material; the Test 1 insulation was 

100% fiberglass, and the Test 4 insulation was 20% fiberglass and 80% calcium silicate. 

To gain a preliminary qualitative understanding of the passivation of the aluminum 

surface, scanning electron microscopy (SEM) was performed on the Tests 1 and 4 

aluminum coupons. An SEM image of an aluminum coupon after 30 days of submersion 

in Test 4 is shown in Figure 3, and a corresponding energy dispersive spectroscopy 

(EDS) analysis associated with this coupon is shown in Figure 4. The EDS image 

indicates that the major components are aluminum, oxygen, silicon, sodium, and calcium, 

with small amounts of carbon and magnesium. Corresponding SEM and EDS spectrums 

of the Test 1 corrosion surface are provided in Figure 5 and Figure 6, respectively. A 

comparison of Figure 4 and Figure 6 indicates that silica is present to a much larger 

degree in Test 4 than in Test 1. An additional analysis of the effect of silica on 

passivating the aluminum surface is presented in Section 8.3. 

Figure 3. Annotated SEM image magnified 1000 times for a Test 4, Day-30, 

submerged aluminum coupon sample. 
7

Figure 4. EDS counting spectrum for the deposits (EDS3) on the coupon surface 

shown in Figure 3.

Figure 5. SEM image magnified 1000 times for a Test 1, Day-30, submerged 

aluminum coupon sample. 
8

Figure 6. EDS counting spectrum for the light spot (EDS1) on the coupon surface 

shown in Figure 5.

2.3. ICET Turbidity and Total Suspended Solids (TSS) 

Measurements 

To assess whether precipitate was being formed under test conditions, both turbidity 

measurements and total suspended solids were measured daily at test conditions. Results 

of these measurements are presented in Figure 7 and Figure 8, respectively. From these 

figures we observe that after an initial decline, presumably due to the settling of the 

debris, the values are fairly constant, after accounting for measurement uncertainty. 

Consequently, it would not appear that bulk aluminum precipitation occurred at test 

temperature. In fact, visual examination of the ICET solutions at 60o

C did not reveal any 

precipitate. (This observation does not address the possibility that colloids could be 

generated that would not be detectable via a turbidity measurement, nor would the TSS 

sample that contained solids via filtration through a 0.7-μm-sized filter identify colloids. 

Additional discussion on colloids is provided in Section 6.) 

Turbidity measurements were also performed at ambient temperature [23 (±2.0)°C], as 

presented in Figure 9. During the first day of Test 1, turbidity at ambient temperature was 

identical to turbidity at the test temperature. However, on Day 2 it was noted that 

turbidity at ambient temperature was higher than at the test temperature. It was also noted 

on Days 2 and 3 that the ambient-temperature turbidity was time dependent and increased 

as the holding time increased. Therefore, a procedure was implemented on Day 4 that 

required the ambient-temperature turbidity to be recorded after a cooling time of 10 

minutes. (The determination of turbidity at room temperature was extremely sensitive to 

the time at which the turbidity was measured and may also have been strongly affected 

by the rate at which the cooling takes place. In fact, the observance of visible precipitate 

from Test 5 was not seen for several days.) 
9

0

4

8

12

16

20

0 5 10 15 20 25 30

Time (Days)

Turbidity (NTU)

ICET1

ICET5

Figure 7. Turbidity versus time at 60°C. 

Figure 8. Total suspended solids versus time. 
10

0

20

40

60

80

100

120

140

160

0 5 10 15 20 25 30

Time (Days)

Turbidity (NTU)

ICET1

ICET5

Figure 9. Turbidity versus time at 23°C. 

From Figure 9, it may be observed that on cooling, the turbidity in Test 1 appeared to be 

rising relatively uniformly until approximately Day 18, when the turbidity appeared to 

reach an asymptotic limit of 133 nephelometric turbidity units (NTU). (From Day 18 on, 

the predominant behavior seemed to be determined primarily by variations due to 

measurement.) An explanation for a rapid rise in turbidity during cooling may be the 

appearance of a white, finely divided precipitate that gradually settled to the bottom of 

the sample storage bottles. The amount of visible room-temperature precipitation 

increased with time, and the precipitate formed more quickly at room temperature as 

testing progressed. A comparison of Figure 1 and Figure 9 suggests that the increasing 

turbidity as a function of time, at ambient temperature, may be related to the aluminum 

concentration. Furthermore, the plateau in the aluminum concentration and the turbidity 

appear to coincide, again suggesting a strong correlation. It should also be noted that the 

turbidity at room temperature for Test 5 increased from ~2 NTU after Day 2 to ~5 NTU 

after Day 17. Additional discussion on the nature of the white precipitate, the solubility of 

the aluminum, and the kinetics of aluminum precipitation is provided in Section 2.8, 

Section 5, and Section 6, respectively. 

2.4. ICET Hydrogen Concentration Measurements 

The hydrogen concentration was measured in the test tank atmosphere, although it should 

be noted that the tank was vented during the course of the experiment. The results of 

these measurements are presented in Figure 10. Although substantial variability can be 

found in the daily measurements, two points may be made: (1) measurable hydrogen, 

which is produced via the dissolution of aluminum, was observed in both Tests 1 and 5; 

and (2) the hydrogen concentrations for both Tests 1 and 5 appear to decline after 

approximately Days 20 and 18, respectively. The termination of hydrogen production, as 
11

noted by the absence of any hydrogen concentration in the tank headspace, is consistent 

with the plateaus in the aluminum concentration and the turbidity measurement at room 

temperature for Test 1. This observation again substantiates the hypothesis that the 

plateau in the aluminum concentration was a result of passivation of the submerged 

aluminum surfaces and not a consequence of reaching a solubility limit. 

0

0.1

0.2

0.3

0.4

0.5

0.6

0.7

0 5 10 15 20 25 30

Time [day]

Gerneration [%]

ICET1

ICET5

Figure 10. Hydrogen concentration above tank versus time. 

2.5. ICET Viscosity Measurements 

Kinematic viscosity measurements were also made to characterize the liquid phase at 

both 60°C and 23°C. Figure 11 and Figure 12 present the results of kinematic viscosity at 

6°C and 23°C, respectively. Measurements of Test 1 at 60°C indicate that the viscosity 

was 0.51 mm2

/s, as compared with a nominal value for water of 0.475 mm2/s (Ref. 4). 

Furthermore, from Figure 11 we observe that the kinematic viscosity for both Tests 1 and 

5 was relatively constant throughout the test. The lack of quantification of the uncertainty 

in the measurement precludes further analysis regarding the slightly elevated values 

observed during the first 20 days of the test. Uncertainty in the viscosity measurements 

was attributed variability in first few days of Test 1 because of technique. 

At room temperature we observed that the kinematic viscosity of Test 1 was substantially 

elevated relative to the starting value. However, substantial variability in the data was 

noted and is a consequence of the sensitivity of the measurement procedure because of 

potentially substantial variations resulting from small variations of temperature over time. 

Finally, the time behavior of the viscosity at 25°C for Test 1 appears to exhibit a strong 

positive correlation with the measured aluminum concentration. That is, the aluminum 

concentration and viscosity measured at 25°C both appear to be constant at 

approximately the same time. 
12

0.4

0.45

0.5

0.55

0.6

0 5 10 15 20 25 30

Time (Days)

Viscosity (mm2/s)

ICET 1

ICET 5

Figure 11. ICET kinematic viscosity versus time at 60°C. 

0.8

1

1.2

1.4

1.6

1.8

0 5 10 15 20 25 30

Time (Days)

Viscosity (mm2/s)

ICET 1

ICET 5

Figure 12. ICET kinematic viscosity at 25°C. 

2.6. ICET Calcium and Silica Concentrations 

Throughout the ICET test series, daily water samples were taken for ICP-AES analysis to 

determine the elemental constituents of the test solution. The concentrations of calcium 

and silicon are plotted in Figure 13 and Figure 14. In Figure 13 and Figure 14, we also 

have included the results of Test 4. It should be noted that 80% of the Test 4 insulation 

was calcium silicate. (In Section 4, we will examine the role of these species on the 

aluminum corrosion.) We note from Figure 14 that the silica concentrations of Tests 1 

and 5 are ~8 mg/L. This concentration of silica is significantly less than predicted by 

previous investigators.5

 In fact, as may be seen in SEM images of the fiberglass (which is 
13

presumably a source for silica and is presented later in the report), the surfaces of the 

fibers appear to be uniform and not to have disintegrated. A possible explanation for the 

apparently low silica concentrations is the fact that it has been reported that the presence 

of aluminum reduces the solubility of silicon, thereby reducing the driving force for the 

dissolution of the silica. 

Analysis of the ICP-AES results indicated negligible quantities of copper, magnesium, 

and zinc in the Tests 1, 4, and 5 solutions. 

As previously mentioned, precipitate was observed both in Tests 1 and 5 on cooling the 

solution to room temperature. To characterize the precipitates formed from the test 

solution, ICP spectroscopy, x-ray diffraction (XRD), and x-ray fluorescence (XRF) 

analyses were performed. 

0

10

20

30

40

50

60

70

80

0 5 10 15 20 25 30

Time [Days]

Concentration (mg/L)

ICET1

ICET4

ICET5

Figure 13. Calcium concentration versus time. 
14

0

50

100

150

200

250

0 5 10 15 20 25 30

Time (Days)

Concentration (mg/L)

ICET1

ICET4

ICET5

Figure 14. Silica concentration versus time. 

2.7. Analysis of Tests 1 and 5 Precipitates 

The precipitate examined from Test 1 was obtained from a high-volume water sample 

that was passed through filter paper (with a pore size of 0.7 μm) on Days 15 and 30. 

During this process, which naturally allowed for cooling of the solution, a white 

precipitate was observed to form on the filter paper. This material was the most 

physically homogeneous material extracted from the ICET experiment. Though 

consistent in appearance with a chemical-physical precipitate formed via precipitation, 

the white material may also have been formed by the aggregation of smaller particles that 

were not visible at the test temperature or by nucleation on small particles of other 

compounds that resided in solution at the test temperature. The precipitate was visually 

examined using both SEM and transmission electron microscopy (TEM). In addition, 

elemental analysis was obtained via EDS, ICP, XRD, and XRF. It should be noted that 

the precipitate was analyzed without washing. Figure 15 shows a typical EDS spectrum 

for a Day-30, high-volume filtered sample. It should be noted that gold (Au) and 

palladium (Pd) are present in almost all EDS spectra because of the sputtering technique 

used to prepare the SEM samples. The highly conductive sample surface provided by 

metallic sputtering prevented a charge from building up under the electron beam and 

destroying the image quality. Peak counting intensities were proportional to the elemental 

concentration but had to be compensated for by the energy-dependent detector response 

function and normalized to an assumed set of elemental constituents before proportional 

compositions could be inferred. Substantial variability in measured elemental 
15

concentrations was observed between samples analyzed with the same method and 

between diagnostic methods applied to the same sample. A survey of these measurements 

was performed to recommend suitably averaged mass proportions for the dominant 

constituents that were observed (see Table 3). 

Figure 15. Day-30, high-volume sample counting spectrum (T1D30-EDS1) on the 

filtrate. 

Table 3. Elemental Composition (wt %) of Test 1, Day-30, High-Volume Filter Sample 

Method O 

(%) 

Al (%) Ca (%) Na (%) B CO3

2- H2O + 

CO2

EDS 48 11 2 16 21 -- -- 

ICP-AES -- 11 0.4 43 18 28 -- 

XRF 45 7 0.8 24 -- -- 34 

In addition to the precipitate, a thin film was noted on the filter paper after passing the 

test solution through the filter paper. A representative EDS spectrum for the film is 

provided in Figure 16. Films also appeared to exist between fiber strands that were also 

examined at various stages during Test 1. Web-like structures between the fiber strands 

revealed an elemental composition that was very similar to that of the films observed on 

the filter paper. The elemental composition of the film-like material found on the fiber, as 

given by EDS, is provided in Table 4. It was found that the elements sodium, boron, and 

oxygen made up 97% of the measured composition. A comparison of the elemental 

composition of the compound found between the fiber strands/films with known sodium 

borate compounds has been performed. This comparison suggests that the compound 

likely is a sodium borate compound. If we consider that borax and tincalconite are both 

23% sodium, 21% boron, and 56% oxygen, it is reasonable to believe that the compounds 

are probably one of these or both. 
16

Figure 16. Day-15, Test 1 counting spectrum (EDS 1-2) for the surface on the film 

overlying the filter. 

Table 4. Elemental Composition of Test 1 Film Material 

Element Mass % 

O 56

Na 25

Al 0.6

B 16

Ca 0.17

Cl 1

When the information from Table 3 and Table 4 is examined, it can be seen that the 

elemental compositions of the precipitate and the thin film on the filter paper are 

different. That is, the precipitate appears to have a large concentration of aluminum, 

whereas the thin film does not have a significant aluminum concentration. 

After Test 1 was completed, the solution was drained from the tank. The resulting 

solution was then distributed into 5-gallon plastic buckets and plastic bottles, where a 

white precipitate formed as the solution cooled. The precipitate was somewhat denser 

than the bulk solution because it settled largely to the bottom of the bottle. Furthermore, 

it appeared to be gelatinous in nature. This material was examined through SEM, EDS, 

and ICP and appeared to be of an elemental composition similar to the material captured 

on the filter paper. 

To understand further the nature of the Test 1 precipitate formed on cooling, the 

precipitate captured following the completion of Test 1 was examined using XRD. The 

samples were obtained from the Test 1 precipitate, which formed on cooling the Test 1 

solution and was captured following the draining of the tank. XRD measurements are 

presented in Figure 17 and Figure 18. In Figure 17, the lower red pattern is the paste, still 

wet in the sample mount, whereas the top blue pattern is the precipitate after it was dried 

and powdered. Both of these samples denote a largely amorphous material. An additional 

sample was obtained for XRD examination (see Figure 18). In this case, the sample was 

dried on a 50°C hot plate and then powdered for the x-ray mount. Results indicate a 
17

significant amorphous component, along with tincalconite and borax at a ratio of ~90:10 

wt % of tincalconite to borax. In Figure 18, the red pattern is the T1 sample, the blue 

pattern is tincalconite, and the green pattern is borax. 

Figure 17. XRD of precipitate (damp, red; dried, blue). 

Figure 18. XRD of precipitate sample (collected on 2/24/05, 1745 dried at 50°C). 

From these figures we observe that the impact of drying appears to affect the materials 

formed. To understand the affect of the drying process further and also to gain a more 

quantitative understanding of the precipitate, a supplemental XRD analysis of the 

samples was performed. 
18

After the initial bulk analyses, the Test 1 post-test showed that the precipitate seen in 

Figure 18 contained both tincalconite and amorphous material. A closer examination of 

this material revealed that the sample was composed of two distinct materials, as may be 

observed from Figure 19. One material is a green-amber, translucent, harder material; the 

other is a white, softer, powdery material. XRD analyses conducted on the two separate 

splits revealed that the white material was relatively pure tincalconite and that the greenamber material was amorphous. 

A Rietveld refinement of the XRD analysis revealed that the material had crystal cell 

dimensions of a = 2.84440, b = 12.25996, and c = 3.68212 (Å) and a crystallite size of 

Figure 19. Test 1 separated precipitate collected on 2/24/05, 1745 bcl, dried at 50°C. 
19

20 (±1) Å. This material appears to be consistent with poorly crystalline boehmite,6

which is an aluminum hydroxide phase having a composition of AlOOH. 

If we consider that it has been found that the sodium borate is likely a manifestation of 

the drying process, then a reanalysis of the elemental composition is needed. 

Consequently, samples of the precipitate from Tests 1 and 5 were washed to eliminate the 

sodium borates and were analyzed with ICP to determine the elemental composition. 

Table 5 presents the results of this analysis. 

Table 5. Distribution of Main Elemental Components of Precipitates from Tests 1 

and 5 

Precipitate 

Sample 

Al (wt %) B (wt %) Ca (wt %) Na (wt %) Other 

Test 1

11/28 (Day 7) 62 24 7 7 0 

12/20 (Day 29) 41 46 2 10 1 

Average 51 35 5 8 0.5 

Test 5

08/02 (Day 7) 43 36 13 8 0 

08/24 (Day 29) 39 40 13 7 0 

Average 41 38 13 8 0 

Comparing the information in Table 3 and Table 5 reveals that the washing of the 

precipitate removed most of the sodium, which presumably was a constituent of the 

initial sodium borate additive. 

To gain an understanding of the amount of hydration of the pseudo boehmite, a drying 

investigation was conducted. Figure 20 presents the results of this investigation. From 

this figure we note that the precipitate is composed largely of water. 

2.8. TEM of Aluminum Precipitate 

To gain a further understanding of the details of the precipitate, TEM was used to 

examine the precipitate from both Tests 1 and 5. In Test 1, the precipitates formed early 

in the test, within a few minutes of cooling, and the quantity increased over the course of 

the test. However, in Test 5, the precipitates took longer to form (several days), and the 

quantities were much smaller. TEM images for both the Day-15 and Day-30, high-flow 

Test 1 samples are provided in Figure 21 and Figure 22. From these figures it may be 

seen that the precipitate appears to be composed of nanoparticles and that the size of the 

particles increases with time. 
20

IC E T -1 G e l

0.0

2 0.0

4 0.0

6 0.0

8 0.0

10 0.0

In itia l A ft e r 6 9 .5 h r s

at 23C

After 4 hrs at

60C

Aft er 4 hrs at

120C

Wt % of sample

Figure 20. Weight loss due to drying of ICET-separated precipitate (boehmite). 

Figure 21. Electron micrograph of the Day-15 filtered test sample magnified 50,000 

times. 
21

Figure 22. Electron micrograph of the Day-30 filtered test sample magnified 50,000 

times. 

Precipitate from Test 5 was also observed after several days at the bottom of the plastic 

bottles used to capture the Test 5 solution. Figure 23 and Figure 24 also present the 

results of Day-15 and Day-30 precipitates. 

From these TEM images, we again see that the basic building blocks of the precipitate 

appear to be nanosized particles and that the aggregates appear more fully developed 

after 30 days as compared with 15 days for both Tests 1 and 5. In addition, an 

examination of the precipitate with TEM did not reveal prevalent crystalline behavior. 

2.9. Particle Size Measurements of Test 1 Solutions 

Although precipitate was not observed during Tests 1 and 5 at 60°C, we cannot rule out 

the possibility that colloids may exist in solution. The determination of whether the 

aluminum species is truly soluble or colloidal may have a significant impact on the 

pressure drop across a filter during the course of a LOCA. The reason for this is that if 

colloids indeed are present in the solution, they may attach to the filter debris on the 

sump screen as they transverse the system. Thus, as a function of time, gel-like material 

may gradually build up as the colloids traverse the coolant circuit and lead to an 

increased pressure drop across the screen relative to the debris itself. To examine the 

possibility that colloids are present in the Test 1 solution, a particle size analysis was 

performed using dynamic light scattering (DLS). DLS provides information regarding 

particle sizes and size distributions in colloidal suspensions via the fluctuations in the 

intensity of light as a result of the Brownian motion of the particles. The rate of these 

fluctuations is a function of the density difference between the particle and the solvent. 

The diffusion coefficient and the particle size can both be estimated from these 
22

fluctuations. Figure 25 presents the particle size distribution of the Test 1 precipitate 

solution at ambient temperature. It should be noted that the range of sensitivity of the 

instrument was 50–1000 nm. Larger particles were removed by first filtering the solution. 

It should also be noted that all solutions were significantly diluted. 

Figure 23. TEM of the Day-15, Test 5 precipitate, magnified 4000 times. 

Figure 24. TEM of the Day-30, Test 5 precipitate, magnified 8000 times. 
23

Figure 25. Particle size (nanometer) distribution for the Test 1 “precipitate 

supernatant” solution at room temperature (aged 4 months). 

This measurement was made of the precipitate solution after approximately 4 months. 

From the measurement, it is seen that the particle distribution appears to be broad, in the 

50–90 nm region. Unfortunately, measurements at lower particle sizes are not possible 

with the laboratory equipment used. Furthermore, the measurements were conducted at 

room temperature in lieu of test conditions because of the limitations of the equipment. 

However, examination of the particle size distribution reveals that nanoparticles existed 

in the solution at room temperature. 

In conclusion, particle size measurements have been performed on room-temperature, 

Test 1 solutions using DLS. Results of these measurements indicate that colloids are 

indeed present. Additional investigations of the system at 60°C are not possible because 

all of the solutions were allowed to cool. However, surrogate solutions have been 

produced to address this issue (Section 9.4). 
24

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25

3. OVERVIEW OF OUTSTANDING ISSUES 

Although a significant understanding of ICET tests has been presented, significant issues 

still need to be resolved. Resolution of these issues is imperative to allow the information 

from the limited ICET tests to be used to assess the properties of systems with slightly 

different characteristics i.e., extrapolation/interpolation of the current test results. 

Furthermore, to ultimately address the issues of the resulting incremental head loss across 

the sump screen, additional information is needed to characterize the total quantity of 

chemical product that may be formed, its rate of production, and its physical 

characteristics for the development of a filtration model. Consequently, in the remainder 

of this report we present the results of supplemental bench-scale testing and analysis of 

the ICET tests, along with a comprehensive review of the literature. This supplemental 

analysis will allow the knowledge basis to be advanced to address these major objectives, 

along with the following outstanding issues: 

What led to plateau of the aluminum concentration in Tests 1 and 5? Did an oxide layer 

develop on the aluminum surface that ultimately inhibited corrosion? What was the 

reason for the abrupt passivation of the aluminum surface in Test 4? 

1. Was the aluminum fully soluble in the test solution at 60°C, or were nanoparticles 

present? 

2. Was the solution supersaturated with aluminum? Did nucleation take place on the 

nucleation sites on the fiberglass? 

3. How is the boron incorporated into the aluminum hydroxide gel that apparently 

forms when the solution cools? Is boron adsorbed onto the aluminum surface? 

What role does boron have in the phase behavior of aluminum and solubility? 

4. What is the role of the organic material within the ICET test? Was this material 

responsible for the increased solubility of aluminum during the test conditions? 

5. Can we predict the aluminum phase behavior and consequently the observed 

solubility behavior with respect to pH, temperature, and aging? What precluded 

the formation of crystalline aluminum hydroxide? 

6. At what temperature did the agglomeration occur that led to the formation of the 

gelatinous material when the solution was cooled? To what extent is the 

gelatinous material redissolved? 

7. What are the settling characteristics of the gel if it is formed, and what will the 

impact be on the aluminum corrosion if localized precipitation, away from the 

aluminum surfaces, does occur, i.e., the impact of nonisothermal conditions that 

are likely to be present in the actual containment? 

8. Can the formation of the gelatinous material and the deposition on the fiberglass 

be precluded through the modification of the solution chemistry? 

9. Can a surrogate solution/precipitate be developed for filtration testing of the 

aluminum systems? 
26

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27

4. ALUMINUM-PHASE BEHAVIOR 

It has been noted that all ICET tests were performed in alkaline environments, as 

indicated in Table 1. In this section, we examine the formation of the phase-controlling 

species in alkaline solutions and address some of the factors that determine the nature of 

the phase formed, including the kinetics of phase transitions (aging) and the experimental 

conditions. 

4.1 . Hydroxide Phases 

In an aluminum-alkaline solution system, a determination of the dissociation of the 

aluminum solid phase is needed to determine the solubility of aluminum in the solution. 

That is, the equilibrium of the following dissolving reaction must be considered: 

Al(OH)3(s) = Al3+ + 3OH- , (1) 

with an equilibrium constant: 

K1 = [Al3+][OH-

]

3

/[Al(OH)3] , (2) 

where [] is the species activity in the solution. If we consider that the aluminum 

hydroxide in Eq. (1) is solid, its activity is set to be unity. 

It has been noted that the equilibrium constant in Eq. (2) is a function of the aluminum 

hydroxide phase. The phase may exist in an amorphous form or as one of three crystalline 

forms: gibbsite, bayerite, or nordstrandite. The crystalline polymorphs differ only in the 

packing arrangement of the layers.7

 That is, the Al-O-Al layer framework is identical.8

Each aluminum atom has three neighboring Al atoms coordinated with pairs of hydroxyl 

groups. The coordination around aluminum is thus octahedral. The layers are held 

together by a network of hydrogen bounds. The ordering of these layers forms the 

essential difference between these structures. In gibbsite, the hydroxyl groups of one 

layer are stacked directly on top of the hydroxyl groups of the next layer, creating a 

crystal of hexagonal morphology, and its relationships are clearly of type AB-BA-AB, as 

shown in Figure 26. The hydroxyl groups of bayerite reside in the depressions of the 

layers below and above, creating a conically shaped crystal. The bayerite has an AB-ABtype lattice. Nordstrandite also has an AB-AB-type lattice and is composed of alternating 

layers of gibbsite and bayerite.8

 It is reported that the solubilities of gibbsite and bayerite 

have at times been shown to overlap, but some information suggests that gibbsite is 

slightly more insoluble than bayerite.9

 
28

 

 

Figure 26. Gibbsite and bayerite layer stacking patterns.7 [Used with permission of the 

publisher.*

] 

XRD has been used to characterize the phase of aluminum hydroxide solids. Gibbsite (α) 

is identified by the presence of three principal diffraction lines at 4.85, 4.37, and 4.32 Å, 

which correspond to scattering angles of sin 2θ of 7.17, 7.98, and 8.05, respectively. 

Bayerite(β) is distinguished by strong diffraction lines at 4.72 and 4.36 Å, whereas 

nordstrandite (γ) is detected by the presence of three major diffraction lines at 4.79, 4.33, 

and 4.22 Å. 

Another common crystal in aluminum-alkaline solution is boehmite, with a composition 

of AlOOH. Pseudoboehmite is built up out of the same fundamental unit as boehmite, a 

double sheet of octahedral crystals with aluminum ions at their center; however, the 

amount of water varies between double octahedral layers, which are identified by broad 

diffraction peaks at 6.6, 3.2 and 2.3 Å. The major distinction between pseudoboehmite 

and boehmite is that the former has broad diffraction peaks and that the spacing of the 

first peak is larger than the first peak of 6.11 Å for boehmite. The amorphous precipitate 

produces a diffraction pattern that is only a halo around the region of 3-4 Å.10

4.2 . Phase Transitions 

Numerous experiments have been conducted in alkaline systems to examine the nature of 

the aluminum species, solubility, and solid phase(s) present; this work has been 

summarized in Sposito.11 Precipitation studies in a basic medium have the distinct 

advantage that the hydrolysis is less complicated as compared with acid systems. In 

weakly aluminum solutions, the aluminate ion is the dominant species. Furthermore, 

experiments have revealed that solubility is very sensitive to the solid phase in 

 

*

 Reprinted from The Environmental Chemistry of Aluminum, S. Goldberg et al., “Chapter 7: The Surface 

Chemistry of Aluminum Oxides and Hydroxides,” 2nd Ed., CRC Press, Copyright 1996, with permission 

from Copyright Clearance Center, Inc. 
29

equilibrium with the solution and that the solid phase formed is a complex function of the 

experimental conditions and the nature of the materials used, i.e., particle size and surface 

area.

11

To address the sensitivity of the results to experimental conditions, an elaborate device 

was designed in which acidified aluminum nitrate solutions were titrated with alkali 

(NaOH or KOH).12 Considerable effort was devoted to ensuring vigorous mixing of the 

reactants so as to avoid developing local hydroxyl ion concentrations appreciably larger 

than the overall concentration because this would result in the formation of a precipitate 

that is most reluctant to redissolving, even though it is unstable. 

Results from this investigation12 indicate that visible precipitate in a system with an 

initial aluminum nitrate concentration of 0.072 M first occurred at a pH of ~6.1 at 25°C. 

Decreases in temperature led to a lowering of the onset pH of precipitation. It should be 

noted that this determination was also made with light scattering. The structure of the 

precipitate was determined after filtering, washing, and drying at 120°C using XRD. 

Table 6 provides the results of this investigation. 

Table 6. Structure of Precipitates, as Revealed by XRD [Used with permission of the 

publisher.*

] 

Temperature (°C) pH 24 28 40 60 90 

6 Aa A 

7 Microbayb A A A 

8 Microbay A A 

9 Microbay A A Gellboc

10 

Microbay 

Bayerite 

Bayerite A 

Bayerite 

Microbay 

Bayerite 

Gibbsite 

Gelbo 

(Gibbsite) 

(Bayerite) 

a

 Amorphous. 

b

 Amorphous with few diffuse lines of bayerite. 

c

 Gelatinous boehmite. 

In these studies neither precipitate nor colloids were found at OH/Al ratios of <2.5, even 

after aging for 2 months. (The value of OH/Al = 2.5 is very close to the vector 

percolation threshold of 2.4, which represents the minimum number of cross links 

required for a three-dimensional structure to attain rigidity.8

) However, in more common 

 

*

 Reprinted from Journal of Colloid and Interface Science, A. C. Vermeulen et al., “HydrolysisPrecipitation Studies of Aluminum (III),” 51:3, pp. 449–458, Copyright 1975, with permission from 

Elsevier. 
30

investigations where the hydroxyl ion concentration is nonuniformly added, a different 

behavior is observed. Depending on the rate of alkali addition, the solid phase can be 

formed at OH/Al ratios of <2.5. In addition, the nonuniform addition of the alkali yields a 

more-crystalline product than the amorphous product that is found in a more uniform 

setting.9,13

In another investigation of the phase behavior of aluminum hydroxide systems, Hayden 

found that the phase behavior was also dependent on the aging.14 Table 7 presents the 

results of this investigation. 

Table 7. Effect of Age on the pH and Crystalline Structure of Aluminum Hydroxide 

Prepared from 0.0378-M Aluminum Nitrate Solutions (Solution also 0.30 in 

Sodium Nitrate) [Used with permission of LANL legal counsel] 

OH/Al 

Ratio 

2.90 2.95 2.95+

 3.02 3.25 

Initial 

pH 

5.00 5.50 5.50 8.00 9.50 

1 Day 

(pH) 

4.85 

Amorphous 

5.48 

Amorphous 

5.54 

Amorphous 

7.85 

Weak Pseudoboehmite 

10.64 

Mixed Pseudoboehmite/ 

Bayerite 

1 Week 

(pH) 

4.82 

Amorphous 

5.18 

Amorphous 

5.41 

Amorphous 

7.60 

Mixed Pseudoboehmite/ 

Bayerite 

10.74 

Mixed Pseudoboehmite/ 

Bayerite/Gibbsite 

1 Month 

(pH) 

4.29 

Amorphous 

4.35 

Amorphous 

5.04 

Amorphous 

7.50 

Mixed 

Pseudoboehmite/Bayerite 

10.79 

Mixture Bayerite 

Gibbsite 

Heated 

(pH at 

65°C) 

4.05 

Gibbsite 

4.10 

Gibbsite 

4.42 

Amorphous 

7.25 

Pseudoboehmite 

9.47 

Gibbsite 

From these results it is clear that increasing the hydroxyl concentration and aging lead to 

the appearance of a more crystalline product. In addition, as may be seen at a low pH 

(i.e., <9.5), aging decreases the pH, whereas at pH = 9.5, aging increases the pH. 

Furthermore, an increase in temperature in all cases decreases the pH. The observed 

changes in pH illustrate that two distinct processes are occurring. The first process is the 

result of precipitation, as given by 

 Al(OH)3 (s) + Al(OH)4

- = 2 Al(OH)3 (s) + OH-

, (3) 

whereas the decrease in the pH is the result of the formation of double hydroxide bridges. 

The second process occurs by the sequential deprotonation and dehydration reactions 

shown in Eqs. (4) and (5) as 
31

Al(OH2)6

3+ =Al(OH)(OH2)5+2+H+

 (4) 

and 

2Al(OH)(OH2)5

+2 =Al2(OH)2(OH2)8+4+2H2O. (5) 

Results are also reported in the literature in terms of the OH/Al ratio. Because it is 

generally accepted that under slight alkaline conditions the rate-limiting step in the 

precipitation of aluminum hydroxide is the formation of Al-OH-Al bridges, the ratio 

OH/Al is of paramount importance.15 That is, below an OH/Al ratio of ~2.5, the pH at 

25°C shows a gradual increase from a pH of 1.2, at very low OH/Al ratios, to ~2.4 at an 

OH/Al ratio of 2.5. 

This pH behavior reveals two fundamentally different processes. On the first plateau, 

polynuclear hydroxo complexes are being formed. Based on acid kinetic and lightscattering studies, it has been concluded that these complexes are largely planar in 

structure, with OH bridges connecting the Al(III) cations. These units therefore resemble 

basal planes of the bayerite or gibbsite lattices. On the second plateau, separated from the 

first plateau by the rise in pH, the primary building blocks formed earlier are being linked 

into three-dimensional structures largely by the hydrogen bonding of additional OH-

 ions 

and Van der Waals interactions. This second stage in the precipitation process represents 

the actual nucleation and initial growth step. It should be noted that formation of the 

polynuclear species is not a necessity for precipitation because experiments have also 

been performed in which aluminum nitrate is titrated into a solution of base, as described 

previously.9

 In fact, ICET is such a configuration in that the aluminum is introduced into 

a very large excess of OH-

; therefore, polynuclear species probably will not be generated 

in such a configuration. 

In this investigation, the phase transition was found to proceed according to X = pH + 

pAT, where 

pH = -log[H+] (6) 

and 

pAT = -log[total aluminum concentration]. (7) 

In cases where X < 12, the transition of phases is amorphous-pseudoboehmite-bayerite, 

whereas if 12 <X <12.55, the phase transition will be pseudoboehmite-bayerite, and if X 

>12.55, bayerite will be immediately formed.9

 Thus, in the ICET configuration, the value 

of X is 11.3 and will result in the initial formation of amorphous aluminum hydroxide. 

Additional investigations have been performed on the transformation of aluminum 

hydroxycarbonate in aqueous solutions to address the mechanism of formation of 

crystalline aluminum species.16 Aluminum hydroxycarbonate is actually amorphous 

aluminum hydroxide in which some carbonate anions substitute for surface hydroxyl 
32

anions.17 The findings of this investigation16 indicate that when a suspension of 

aluminum hydroxide in water is aged at a pH of 7 or higher, it undergoes a two-step 

aging process: amorphous aluminum hydroxide transforms into poorly ordered boehmite 

(pseuoboehmite),18 which in turn transforms into bayerite, the stable polymorph. It has 

also been shown that supersaturated aluminate solutions form the most-soluble phase 

first, become saturated with that phase, and subsequently form the next soluble phase.9

This process precisely follows the Ostwald rule of stages. For example, a solution 

supersaturated with respect to amorphous Al(OH)3, pseudoboehmite, and bayerite 

(phases in order of decreasing solubility) will first form the amorphous Al(OH)3, then 

pseudoboehmite, and finally bayerite. Upon aging of the precipitate, amorphous Al(OH)3

will convert first to pseudoboehmite, which will subsequently convert to bayerite. 

The amorphous solid is initially high in both water and hydroxyl anion concentration. 

The rate of transformation to the more crystalline form is controlled by the rate at which 

hydroxyl anions replace water in amorphous solid.15 Boehmite is formed by solid-state 

interparticle and intraparticle condensation/aggregation.16 This reaction occurs as water is 

removed.19 It has been proposed that a hydroxyl anion reacts topochemically with 

another hydroxyl anion in the amorphous solid. This reaction forms a water molecule and 

under neutral or alkaline conditions produces the aluminum oxyhydroxide known as 

boehmite.9

 It has also been shown that the solubility of poorly ordered boehmite 

determines the rate of nucleation and growth of the crystalline aluminum polymorphs.20

Additional support for rapid precipitation and crystallization in neutral and alkaline 

media to produce bayerite has been provided.21 Finally, the solubility of the system also 

decreases through the transition process. 

Figure 27 and Figure 28 present the results of the investigation using commercial spraydried aluminum hydroxycarbonate (SDAHC) in water.16 The surface pH of the SDAHC 

was initially approximately 5. When the SDAHC was saturated with water, the surface 

pH was determined to be 10. This high surface pH is thought to cause SDAHC to age in 

the same fashion as aluminum hydroxide suspensions precipitated at a high pH or 

adjusted to a high pH.10 These figures clearly illustrate the role of aging in the 

transformation of the initially amorphous material into crystalline material. A comparison 

of Figure 27 and Figure 28 also illustrates the importance of temperature in the kinetics 

of the transformation. As may be seen from Figure 28, the phase transformation at 50°C 

proceeds quickly, and the stable form is crystalline boehmite. 

The transformation into bayerite is believed to occur at a pH of 10 because of the tight 

packing of the particles. The point of zero charge (PZC) of aluminum hydroxide is 9.6.16

Thus, the net surface charge of SDAHC disappears as the surface pH reaches 10. The 

absence of a substantial surface charge allows the particles to form bayerite, the most 

tightly packed polymorph of aluminum hydroxide. 
33

Figure 27. X-ray diffractograms of SDAHC after it has aged at 100% relative humidity 

(RH) and 25°C for 0 days (bottom), 62 days (second from bottom), 118 days 

(middle), 215 days (second from top), and 384 days (top). Diffraction bands 

associated with boehmite or bayerite are represented by dashed or solid 

lines, respectively. [Used with permission of the publisher.*

] 

. 

Figure 28. X-ray diffractograms of SDAHC after it has aged at 100% RH and 50°C for 0 

days (bottom), 6 days (lower center), 36 days (upper center), and 78 days 

(top). Diffraction bands associated with boehmite are indicated by dashed 

lines. [Used with permission of the publisher.*] 

Additional evidence of the phase transformation of amorphous aluminum hydroxide as a 

result of aging is provided by the infrared (IR) and XRD examination of amorphous 

aluminum hydroxide.22 In this investigation, both IR and XRD measurements, Figure 29 

 

*

 Reprinted from Journal of Colloid and Interface Science,, K. E. Hancock et al., “The Effect of Humidity 

on the Physical and Chemical Stability of Spray-Dried Aluminum Hydroxycarbonate,” 183:2, pp. 431–

440, Copyright 1996, with permission from Elsevier. 
34

and Figure 30, on aluminum hydroxide gel, produced by titrating ammonium into 

aluminum chloride solution to a pH of 7, were performed. In this pH regime, the 

experimentally observed phase is expected to be gibbsite. (By comparing the IR spectra 

results in Figure 29 to the O-H stretching and bending frequencies of crystalline 

aluminum phases in Table 8, transformation to the crystalline structure may be seen.) 

Further evidence of this transformation is also seen in the XRD diffraction lines in Figure 

30. 

Table 8. Characteristic Stretching and Bending Frequencies for Crystalline 

Aluminum Phases 

Frequency cm−1

Phase OH 

Stretching 

Region 

O-H 

Bending 

Region 

3612 1030 

3520 970 

3440 Gibbsite 

3395 

3650 1010 

Bayerite 3540 970 

3460 

3300 1150 Boehmite 3095 1070 

An examination of aluminum precipitation from supersaturated solutions showed that the 

rate of nucleation was determined largely by two parameters: the degree of 

supersaturation, Π, and the interfacial tension, σ.9

 Adsorption of lattice or foreign ions 

may (drastically) change the magnitude of the interfacial tension and thereby alter the 

rate of nucleation. Adsorption of surface-active species is also known to influence the 

rate and mechanism of crystal growth. The crystallization process may be retarded, or the 

crystal habit of the precipitate may change significantly23 and the polydispersity of the 

product may be controlled by adding surface-active species to the supersaturated solution. 

Consequently, in the next two subsections, we examine the role of the anion/cation and 

organics in the phase behavior of aluminum. 
35

 

Figure 29. Change in the O-H stretching frequency region of the IR spectrum of 

aluminum hydroxide gel during aging at 25°C for (A) 5 days, (B) 42 days, 

(C) 54 days, (D) 61 days, (E) 69 days, and (F) 124 days. [Used with 

permission of the publisher.*

] 

Figure 30. Change in XRD pattern of aluminum hydroxide gel during aging at 25oC for 

(A) 57 days, (B) 65 days, (C) 71 days, and (D) 124 days. [Used with 

permission of the publisher.*

] 

4.3. General Cation/Anion Effects on the Aluminum Phase 

In the previous section we examined the phase transformation behavior of aluminum in 

aqueous solutions due to aging. From this section we observed that the aging of the 

precipitate generally led to the formation of more crystalline material. However, the 

measured aluminum concentrations of the Tests 1 and 5 solutions, accompanied by the 

 

*

 Reprinted from Journal of Pharmaceutical Sciences, Nail, S. L., et al., “Comparison of IR Spectroscopic 

Analysis and X-Ray Diffraction of Aluminum Hydroxide Gel,” 64:7, pp. 1166-1169, Copyright 1975. 
36

XRD analysis of the wet precipitate obtained after draining the Test 1 tank solution while 

allowing it to cool, reveal a mostly amorphous form with a detectable crystalline structure 

of only sodium pentaborate (a sodium/boron compound). Consequently, in this section 

we seek to explain the persistence of the aluminum amorphous phase. 

When an aluminum hydroxide solid is prepared by precipitation from an aluminum salt 

solution, a highly random structure is formed. Anions present during precipitation are 

absorbed by the aluminum hydroxide gel and are believed to be important in stabilizing 

the colloidal system. Studies have shown that the crystallization of gibbsite from 

amorphous aluminum hydroxide can be greatly inhibited by the presence of sulfate,24

silic acid,25 and citric acid.26 Nitrate has a rather weak effect, and perchlorate has no 

effect at all.27 Figure 31 illustrates the mechanism of anion interference to 

crystallization.28

Figure 31. Possible pathways for polymerization of Al(OH)3. [Used with permission of 

the publisher.*

] 

Based on the adsorption capacity of aluminum hydroxide for several anions at different 

pH conditions, three types of anion adsorption have been proposed.29 (A relatively high 

degree of adsorption of anions occurs when aluminum hydroxide gel is precipitated at pH 

conditions <9.2, its PZC.) Nonspecific adsorption was suggested for anions such as 

nitrate, chlorate, and chloride, which are loosely held in the diffuse layer and therefore 

are adsorbed only by positively charged surfaces. Specific adsorption of anions of 

completely dissociated acids, such as sulfate and fluoride, is based on chemical 

adsorption and involves ligand exchange with surface water. A third type of adsorption is 

the specific adsorption of incompletely dissociated acids such as phosphate and silicate.29 

 

*

 Reprinted from Soil Sci Soc Am J, Beyrouty, C. A., et al., “Evidence Supporting Specific Adsorption of 

Boron on Synthetic Aluminum Hydroxides,” 48:2, 284–287, Copyright 1984, with permission from 

ASA-CSSA-SSSA. 
37

IR spectroscopy has proven useful in determining if the ligands are within the 

coordination sphere.30 IR spectroscopy studies have been performed using aluminum 

hydroxide gels prepared by stirring 1.9 M NaOH into the respective aluminum salt 

solution (0.25 M) until a pH of 6.5 was obtained. The influence of carbonate and 

bicarbonate also has been studied.31 The freshly precipitated aluminum hydroxide gels 

were air dried and examined by XRD and found to be amorphous. 

Gels produced with the nitrate anion show a nitrate band split in the aluminum hydroxide 

gel appearing at 1395 and 1350 cm-1, as compared with the undisturbed 1358 cm-1 IR 

band.28 The splitting of the nitrate band indicates that some perturbation occurs when the 

nitrate anion is incorporated into the aluminum hydroxide gel. It has been observed that 

the ν3 vibration splits into two new frequencies when the nitrate anion has a lowered 

symmetry due to the perturbation.32 The magnitude of the split appears to increase as the 

dissymmetry occurs. Based on previous observations, the relatively small degree of 

splitting observed for the nitrate anion in the aluminum gel, 45 cm-1, indicates an 

electrostatic interaction in which the nitrate anion is outside the coordination sphere of 

the aluminum cation.32 This observation is in good agreement with the nonspecific 

adsorption attributed to this anion.29 

The affinity of the carbonate anion for the aluminum cation is evidenced by the presence 

of carbonate in reactive aluminum gel.33 The IR spectra of the carbonate-containing 

aluminum hydroxide gel indicate that the carbonate anion is in a state of lower 

symmetry.31 The adsorption at 1090 cm-1 in the carbonate-containing aluminum 

hydroxide gel is active only in the Raman when the carbonate anion has full symmetry. In 

addition, the ν3

 vibration splits into two new adsorption bands at 1500 and 1435 cm-1.

31

The nature of the carbonate-to-aluminum bond in aluminum hydroxide gel can be 

inferred by comparison with carbonate complexes and carbonate surface oxides.34 The 

magnitude of the splitting, 115 cm-1, indicates that the carbonate anion is directly 

coordinated to the aluminum cation; therefore, specific adsorption is indicated.28

Figure 3235 is an electron micrograph illustrating the role of the anion PO3-in retarding 

the formation of crystalline growth. It appears that the primary particles aggregate to 

form larger secondary particles. This micrograph appears similar to the Test 1 

micrographs depicted in Figure 21 through Figure 24. Thus, it may be concluded that 

many anions interfere with Al(OH)3 crystallization. The interference is influenced by 

many factors, including the anion-aluminum stability constant, the anion ionic size and 

structure, and the anion concentration.21 

Studies also have been conducted to investigate the role of organics in the crystallization 

of aluminum. Experiments conducted to examine the effect of small concentrations of 

citric acid showed that the effect was to reduce the amount of precipitate at ambient 

temperature. Analysis of the precipitate revealed an amorphous material.27 It has been 

found that citric acid retards the crystallization of aluminum.27 The noncrystalline nature 

of the precipitate apparently was due to the occupation of coordination sites of aluminum 

ions by citric acid on dissociation of the proton, resulting in a distortion in the 

arrangement of the hexagonal ring units normally found in crystalline aluminum 
38

hydroxides. Crystalline behavior became apparent only with sufficient OH-

 and with 

aging. 

Figure 32. Electron transmission micrograph of aluminum adjuvant precipitated at pH 

= 4.5 in the presence of phosphate. [Used with permission of the 

publisher.*

] 

It has been suggested that the retardation of crystalline behavior was the result of activity 

at the solid/solution interface.36 Adsorption of the citrate also can lead to changes of the 

surface charge and interfacial tension and can influence the kinetics of numerous 

elementary steps involving the incorporation of the growth unit from the solution into the 

growing solid phase.31 However, the principal effect was suggested to be the occupation 

of certain surface sites, which caused the growth unit to be blocked.29 

Additional experiments have been conducted in which small quantities (0–50 mg) of 

fulvic acid (FA) were added to aluminum chloride solutions, 450 ml of a 10-3 M 

solution.10 (FA resembles citric acid in that it contains CO2H and aliphatic OH groups; it 

resembles quercetin in that it also contains phenolic hydroxyl and ketonic C=O groups. It 

is through these functional groups that FA can form stable complexes with aluminum.) 

The solutions were then titrated with 0.1 M NaOH to raise the pH to 6, 8, and 10. It was 

found that FA inhibited the formation of gibbsite at a pH of 6, leading to the formation of 

pseudoboehmite. In systems with a pH of 10, the addition of as little as 0.0125 mg/L of 

FA was sufficient to prevent the formation of any precipitate.10 A simple explanation may 

be that at a pH of 10, the surface of the aluminum and FA are both negative such that 

electrostatic repulsion between the two components appears responsible for the lack of 

precipitation. 

 

*

 Reprinted from Vaccine, L. Burrel, “Aluminum Phosphate Adjuvants Prepared by Precipitation at 

Constant pH: Part 2, Physicochemical Properties,” 19:2-3, p. 282, Copyright 2001, with permission from 

Elsevier. 
39

In summary, we have seen that the presence of certain organic and inorganic agents, even 

in small quantities, is sufficient to modify the solubility of aluminum and also preclude 

crystallization. In the next section we investigate the effect of boron. 
40

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41

5. ALUMINUM-BORON COMPLEX 

Boron is present within the Test 1 solution at concentrations approaching 2800 mg/L. We 

have previously noted that the precipitate that formed on cooling the ICET solution had a 

high weight percentage of boron. Consequently, an examination of the interaction of 

boron in solutions containing aluminum in conjunction with sodium hydroxide is 

warranted. To this end, research has been conducted on the role that boron may play in 

the formation of complexes, the determination of the solubility of aluminum, and the 

adsorption of boron onto aluminum surfaces. 

5.1. Boron Behavior in the Solution 

Boric acid is a Lewis acid and hydroxyl acceptor that through hydrolysis produces a 

change in coordination from planar to tetrahedral.37 The dissociation constant, pKa, for 

boric acid is 9.24, and therefore boric acid exists predominantly as undissociated boric 

acid [B(OH)3] in dilute aqueous solution below a pH of 7. Above a pH of 10, the 

metaborate anion B(OH)4 becomes the main species in solution. Between a pH of 6 and 

11 and at high concentration (>0.025 mol/L), highly water-soluble polyborate ions such 

as B3O3(OH)4

-

, B4O5(OH)4

2-, and B5O6(OH)4-

 may be formed.37

Calculation of the aqueous solution chemistry of boron is generally confined to 

hydrolysis because of the lack of thermodynamic data for aqueous cation complexes 

containing boron. Only one paper reports thermodynamic data for the aluminum boron 

complex.38 These data have drawn critical comments from several sources.39

Furthermore, the kinetic data have been unsuccessful in explaining the appearance of a 

white amorphous precipitate that appeared in testing an Al = 0.005 M and B = 0.04 M 

solution titrated with sodium hydroxide.40 This experimental finding is consistent with 

the white precipitate that was observed in the benchtop test performed. In addition, 27Al 

and 11B NMR measurements previously were unable to detect chemical complexes at 

either pH = 2 or pH = 12.37 However, 27Al and 11B solid-state NMR measurements 

performed at Los Alamos National Laboratory (LANL) on the Test 1 precipitate revealed 

the presence of increased coordination between aluminum and boron (as shown in Figure 

33) as the temperature of the solution was reduced from 60°C to room temperature. 

A previous investigation also revealed that aluminum borate complexes are present at 

25°C.41 The 27Al spectra performed at pH = 9 in Al-B solutions with B = 0.02 M show 

the presence of two peaks at 80.5 and 74.5 ppm, which correspond to Al(OH)4

-

 and a 

single substituted dimmer, Al(OH)3OB(OH)2. In 0.08- and 0.2-M borate solutions, a third 

peak appeared at 68.5 ppm, which could be assigned to the trimmer, 

Al(OH)2O2(B(OH)2)2

-

. Solubility studies were also performed that allowed for the 

determination of the equilibrium constant for the dimmer reaction 

Al(OH)4- + B(OH)3 aqo = Al(OH)3OB(OH)2

- + H2O . (8) 
42

Figure 33. The 27Al nuclear magnetic resonance (NMR) of benchtop aluminum 

nitrate/boron/sodium hydroxide solution at pH = 9.5 at 60°C, 40°C, and 25°C 

after 4 weeks of aging. 

Furthermore, the temperature dependence for the equilibrium constant has been found to 

be 

log K = 241.94 (+-130.7)/T(K) + .81(+-.36) . 

At 25°C, Log K was found to be 1.62. 

The impact of the aluminum borate complex on the equilibrium has been shown to lead 

to significant solubility in aluminum (up to factors of 6) for both gibbsite and boehmite.38

Future equilibrium calculations for the ICET configuration must include the borate 

complex to determine the important effects of the presence of borate in the solution on 

the aluminum solubility. 

 

Aluminum borates were synthesized by a precipitation process in which dilute solutions 

of aluminum nitrate nonahydrate Al(NO3)3:9H20 and boric acid were precipitated into a 

basic solution of ammonium carbonate; the resulting mixture was evaporated to dryness. 

It was found that until excessive heating was applied, the material was largely amorphous 

in nature.42,43,44 The amorphous nature of the precipitate from the Test 1 solution has also 

been confirmed via XRD at LANL. As previously discussed, the temperature of the 

precipitation may affect the quantity of boron absorbed. 
43

5.2. Boron Adsorption 

In a study of the adsorption of boron onto aluminum hydroxide surfaces, a solution was 

synthesized by the dropwise addition (~2 ml min-1) of 0.5 M NaOH (the base typically 

contained 0.75 M H3BO3) into 100 ml of 0.5 AlCl3 until the pH was 7.0. The solution 

was then analyzed using IR analysis,24 and the boron content, with respect to the 

Al(OH)3, was then determined. It was found that the boron present in the gel was held 

predominantly on the hydroxide surface as a specifically absorbed ion. This mechanism 

of anion exchange with hydroxyl ions led to specific surface adsorption.45 Specific 

adsorption of anions produced a shift in the ZPC of the mineral to a more-acid value. 

Furthermore, aging studies were conducted, and it was found that the adsorption of boron 

onto the aluminum surface precluded crystallization. 

Figure 34 presents the IR results obtained as a result of aging the aluminum hydroxide 

gels.28 As may be seen, the commercially prepared hydroxide shows three well-resolved 

peaks in the OH stretching region at 3605 cm-1, 3510 cm-1, and 3450 cm-1and a doublet at 

3380 cm-1and 3360 cm-1. These bands are characteristic of gibbsite, as may be seen from 

a comparison with Table 8. The 13-day-aged solution without boron also shows evidence 

of a high degree of structural order. These peaks indicate a gibbsite and bayerite mixture, 

as seen from Table 8. In contrast, the aged boron containing hydroxide exhibited a very 

broad adsorption band in the OH stretching region, centering at 3420 cm-1. The absence 

of any sharp peak confirms the hypothesis that boron was held directly on the hydroxide 

surface, thereby inhibiting the polymerization process. 

Identification of the boron absorbed onto the amorphous aluminum hydroxide was 

obtained by examining the IR spectra in the 800- to 1500-cm-1 region. Although the 1020-

cm-1and 965-cm-1 adsorption bands are not present in the gel with boron, two broader 

peaks appear at 1060 cm-1 and 900 cm-1. In addition, two new bands have also appeared 

at ~1415 cm-1and 1300 cm-1.

28 

Another means by which to evaluate the adsorption of an ion is the change in the PZC. 

The PZC is defined as the pH at which the net surface charge is zero or cation exchange 

capability minus anion exchange capacity equals zero. Specific adsorption of an anion 

makes the surface to which it is adsorbing more negatively charged. This specific 

adsorption produces a shift in the PZC to a more acidic value. Specific-adsorbed ions are 

ions held in inner-sphere surface complexes that contain no water between the adsorbed 

ion and the surface functional group. Kinetic experiments using pressure-jump-relaxation 

have also been used to confirm that boron adsorbs as an inner-sphere complex on 

aluminum oxide via a ligand exchange of borate with surface hydroxyl groups.46 The 

PZC for the Al(OH)3 free of any specifically adsorbed anions is 9.72, whereas for the 

boron-containing gel, the PZC was found to be 7.57–8.14.28 
44

Figure 34. Spectra of (A) Al(OH)3 with boron (89 days old), (B) Al(OH)3 prepared 

without boron (13 days old), and (C) commercially prepared Al(OH)3. [Used 

with permission of the publisher.*

] 

In another study on the adsorption of boron onto amorphous aluminum hydroxide, 

attenuated total reflectance Fourier transform infrared spectroscopy (ATR-FTIR) was 

used.47 The study of aqueous solutions of metal ligand systems by conventional IR 

spectroscopy is limited because water absorbs energy in several areas of the IR spectrum, 

resulting in very strong, broad bands in the regions of 3650–2930, 1720–1580, and 930–

400 cm-1. To minimize this, it is necessary to use solutions of high concentrations in cells 

with very short optical path lengths, which results in poor band resolution. ATR-FTIR

overcomes the conventional IR spectroscopy limitation and allows an IR spectrum of an 

aqueous solution of a metal ligand to be observed.48 It was found that trigonal B 

asymmetric bands shifted to higher frequencies and indicated that both B(OH)3 and 

B(OH)4

-

 were absorbed via a ligand exchange mechanism.47 

The adsorption of boron onto amorphous aluminum hydroxide surfaces has been 

previously studied and found to depend on both the pH and temperature.47 In Ref. 47, 

0.25 g of amorphous aluminum hydroxide was added to a 30-ml centrifuge tube and 

equilibrated with 12.5 ml of 0.1 M NaCl containing ~0.462 mmol/L of boron. The 

suspension was then pH adjusted with either 1.0 M HCl or NaOH. The addition of HCl or 

NaOH caused a <8% change in the final volume. The supernatant was separated by 

 

*

 Reprinted from Soil Sci Soc Am J, Beyrouty, C. A., et al., “Evidence Supporting Specific Adsorption of 

Boron on Synthetic Aluminum Hydroxides,” 48:2, 284–287, Copyright 1984, with permission from 

ASA-CSSA-SSSA. 
45

centrifugation for 30 minutes at 25°C and 5°C, analyzed for pH, and then filtered through 

a 0.1-mm filter. The boron concentration was determined using a Technicon autoanalyzer 

and the azomethine H method.49 The total boron adsorbed was the difference between the 

total boron added, indicating that the boron remained in solution. Results of this study are 

depicted in Figure 35.47 As may be seen in this figure, the decreasing temperature favors 

adsorption. This decrease occurs mainly because increasing temperature decreases the 

PZC.50 A decrease in PZC renders the surface more negative at a given pH and therefore 

decreases the adsorption of anions. Adsorption is maximized onto high-surface-area 

materials. The adsorption results at low temperature are also consistent with the 27Al 

measurements, which indicates that complexation between aluminum and boron is 

favored at low temperature. 

Figure 35. Absorbed boron as a function of pH at 25oC and 5oC. [Used with permission 

of the publisher.*

] 

ATR-FTIR spectra for boric acid as a function of pH and total boron concentration are 

given in Figure 36.47 The broad bands at 1410 and 1148 cm-1 are assigned to the B-O 

asymmetric stretching of triaganol boron and to B-OH in the plane bending of trigonal 

boron. These bands are seen to increase in intensity with increasing boron concentration. 

At a pH of 9, which is near the pKa for the monomeric boron species, roughly half of the 

total boron is in the form of B(OH)3 and half is in the form of the B(OH)4

-

 anion. The 

band at 955 cm-1 is assigned to the asymmetric stretching of tetrahedral boron, and the 

broad band at 1154 cm-1 is a mixture of B-OH bending of both trigonal and tetrahedral 

boron. Based on the above data, the most useful diagnostic bands are at 1410 cm-1 for 

 

*

 Reprinted from Environmental Science Technology, C. Su and D. L. Surez, “Coordination of Adsorbed 

Boron: A FTIR Spectroscopic Study,” 29, p. 302, Copyright 1995, with permission from The American 

Chemical Society. 
46

asymmetric stretching of triagonal boron and at 955 cm-1 for asymmetric stretching of 

tetrahedral boron. 

 

Figure 36. ATR-FTIR spectra of aqueous boric acid as a function of pH and total boron 

concentration. Each spectrum is the difference between the sample (boric 

acid in 0.1 M NaCl at each pH) spectrum and the reference (0.1-M NaCl 

solution) at the same pH. [Used with permission of the publisher.*

] 

. 

Figure 3747 presents the absorbance spectra for pH = 6.8 and 10.2 aluminum-boron 

solution for three initial boron concentrations: 4.62, 9.25, and 23.1 mmolL. The 

concentrations indicated in the figures are given for the equilibrium solution 

concentrations. Thus, boron is strongly absorbed, ~70% at 5°C. The spectra indicate the 

presence of triagonal boron, 1420 cm-1 and 1280 cm-1. The bands were shifted to higher 

frequencies than for the pure boric acid (1410 cm-1 and 1148 cm-1) at a pH of 7. This shift 

is probably due to the strengthening of the O-B and B-OH bonds in the surface complex 

Al-O-B(OH)2 when the boric acid molecule is complexed with surface functional groups 

such as Al-OH. Thus, at a pH of approximately 7, the trigonal boron is the primary 

species absorbed. 

 

*

 Reprinted from Environmental Science Technology, C. Su and D. L. Surez, “Coordination of Adsorbed 

Boron: A FTIR Spectroscopic Study,” 29, p. 302, Copyright 1995, with permission from The American 

Chemical Society. 
47

 

Figure 37. ATR-FTIR spectra of am-Al(OH)3 at (a) 25oC at pH = 6.8 and (b) 5oC at pH = 

10.2. [Used with permission of the publisher.*

] 

At a higher pH, the B-O asymmetric stretching of triagonal boron is not significantly 

shifted compared with 1410 cm-1 for pure boric acid. However, the B-OH bending of the 

triagonal boron band resulted in a narrower width and a higher wave number of 

1266 cm-1 from 1148 cm-1 due to the complexation on the amorphous aluminum 

hydroxide.47

Attempts to characterize the tetrahedral boron failed due to severe band interference in 

the range of 1000-900 cm-1 from the Al-O bond, which has a strong absorbance at 

969 cm-1. Although the tetrahedral boron is the dominant species at pH = 10.2 in aqueous 

solution, it need not be the dominant absorbed species on amorphous aluminum 

hydroxide due to the negative charge on the solid surface at this pH. The appearance of 

the triagonal boron on the surface of the amorphous aluminum hydroxide at a pH of 10.2 

suggests that the neutral B(OH)3 species could be preferred because its higher affinity for 

the negatively charged surface of the amorphous aluminum hydroxide is higher than it is 

for the borate ion, which would experience charge repulsion.47 (That is, at this pH, the 

surface charge on the amorphous aluminum hydroxide is negative.) 

Several surface complexation models are available to describe the adsorption of boron 

onto aluminum surfaces. These chemical models provide a molecular description of 

adsorption using an equilibrium approach with mass action and mass balance equations. 

In these models, boron adsorption occurs via complex formation with surface hydroxyl 

groups on aluminum. Surface complexation models define surface species, chemical 

 

*

 Reprinted from Environmental Science Technology, C. Su and D. L. Surez, “Coordination of Adsorbed 

Boron: A FTIR Spectroscopic Study,” 29, p. 302, Copyright 1995, with permission from The American 

Chemical Society. 
48

reactions, mass balance, and charge balances and calculate thermodynamic properties 

such as activity coefficients and equilibrium constants. Two models that are particularly 

well suited for studying boron’s adsorption onto amorphous aluminum hydroxide 

surfaces are the constant capacitance model and the triple-layer model. 

The constant capacitance model of the oxide aqueous solution interface has been applied 

to describe boron adsorption on aluminum.51 In the constant capacitance model, anion 

adsorption is assumed to occur via a ligand exchange mechanism with the reactive 

surface functional group, AlOH; no surface complexes are formed with ions in the 

background electrolyte. Specifically adsorbed ions reside in the surface plane of 

adsorption, along with protons and hydroxyl ions. (In some cases, chemically unrealistic 

values of the protonation-dissociation constants were obtained, potentially reducing the 

chemical significance of the model application.52) 

The triple-layer model has also been applied to describe boron adsorption on aluminum.52

In contrast with the constant-capacitance model, the triple-layer model allows anion 

adsorption to occur specifically via ligand exchange or nonspecifically through the 

formation of outer-sphere surface complexes with the reactive surface functional group, 

AlOH. Outer-sphere complexes contain at least one water molecule between the 

adsorbing ion and the surface functional group. The triple-layer model always includes 

outer-sphere complexation reactions for ions of the background electrolyte. Surface 

reactions, equilibrium constants, mass balance, and charge balances for the application of 

the triple-layer model are provided in Goldberg.53 However, it has been found that only 

the inner-sphere adsorption mechanism could successfully describe both equilibrium and 

pressure jump kinetic data using the triple-layer model.46 

Desorption of boron may be explained by various mechanisms, including ligand 

exchange, formation of bidentate surface complexes, and incorporation into lattices. 

Ligand exchange is reversible with respect to pH changes; however, anion desorption at 

constant pH exhibits varying degrees of irreversibility.54 

In conclusion, if amorphous-phase aluminum is indeed present, as both XRD and TEM 

analyses would indicate in Test 1, then it is not surprising that significant boron would be 

found in the precipitate as a result of the adsorption of boron from the solution during the 

cooling process. As indicated from the ICP results, up to 35% of the boron from the 

initial solution may have been adsorbed onto the amorphous aluminum hydroxide 

precipitate. This degree of adsorption is certainly feasible. In fact, if we use the data from 

Su and Surez47 at a pH of ~9.5, a boron adsorption of ~35% is predicted. It must be 

emphasized that we have shown that the complexation of aluminum and boron is 

dependent on the temperature. Consequently, if precipitate were to have formed during 

the course of Test 1, the appearance of boron would probably not have been found in the 

precipitate. Thus, it is reasonable to assume that precipitation of aluminum did not occur 

during Test 1 and was produced only via the cooling process. Finally, the inclusion of 

boron in the inner sphere of the tetrahedral aluminum has been shown at even very low 

Al to B ratios (~20) to preclude crystallization of aluminum, even when calcined to 

500°C.44 
49

6. PARTICLE SIZE DISTRIBUTION 

Aluminum behavior, such as solubility in an alkaline solution, is a function of the particle 

size distribution, which is affected by the aluminum hydroxide phase and the organic 

elements. In Section 2.9 we observed that colloids were present in aged Test 1 solution 

that had cooled to room temperature. In this section we briefly review the previous 

investigations of particle size measurements in alkaline aluminum systems at room 

temperature. 

Figure 38. Particle size for aluminum in pH = 4.5 solution. [Used with permission of 

the publisher.*

] 

Previous investigations into the particle size distribution of aluminum have been made 

using DLS, small-angle x-ray scattering (SAXS), and acoustic techniques.8,55,56 A typical 

particle size distribution, obtained using an acoustic technique for alumina used in 

chemical mechanical planarization (CMP) at a pH of 4, is given in Figure 38.55 The mean 

particle radius was determined to be ~65 nm at ambient temperature, as shown in the 

figure. 

The particle size distribution shown in Figure 38 is representative of particles that are 

very stable against aggregation because they repel one another. The stability of particle 

dispersion will depend on the balance of the repulsive and attractive forces that exist 

between particles as they approach one another. If all particles have a mutual repulsion, 

the dispersion will remain stable. However, if the particles have little or no repulsive 

force, some instability mechanism will eventually occur, e.g., flocculation or aggregation. 

The degree of particle repulsion can be measured using a zeta potential. The zeta 

potential of a particle is the overall charge that the particle acquires in a particular 

medium. The magnitude of the measured zeta potential is an indication of the repulsive 

 

*

 Reprinted from ACS Symposium Series, T. Oja et al., “Acoustic Analysis of Concentrated Colloidal 

Systems,” 881, pp. 231–248, Copyright 2002, with permission from The American Chemical Society. 
50

force that is present and can be used to predict the long-term stability of the product. If all 

particles in suspension have a large negative or positive zeta potential, they will tend to 

repel each other, and the particles will have no tendency to come together. However, if 

the particles have low zeta potential values, no force would prevent them from coming 

together and flocculating. 

At the pH at which the particle size distribution was measured in Figure 38, pH = 4.5, the 

zeta potential was large and thus the distribution was stable. As the pH was raised by 

adding base, such as NaOH, the electrical barrier was lowered from a large positive value 

to zero at around 9.0 and then to a negative potential at a high pH, i.e., >9. At a high pH, 

the system was again stable against further aggregation. Figure 39 demonstrates the 

dependence of particle size on the pH.55 

Figure 39. Particle size versus pH for aluminum at ambient temperature. [Used with 

permission of the publisher.*

] 

Dynamic studies of the particle size have also been conducted using quasi-elastic light 

scattering.57 In these investigations, it was found that in solutions of aluminum chloride, 

large particles that were quickly formed subsequently shrank until sufficient hydroxide 

was added, as shown in Figure 40. 

 

*

 Reprinted from ACS Symposium Series, T. Oja et al., “Acoustic Analysis of Concentrated Colloidal 

Systems,” 881, pp. 231–248, Copyright 2002, with permission from The American Chemical Society. 
51

Figure 40. Particle growth dynamics in aluminum chloride solutions as a function of 

time for various OH/Al ratios at ambient temperature. [Used with 

permission of the publisher.*

] 

Particle size investigations have also been conducted using SAXS.58 In nitrate and 

chloride solutions with concentrations of 0.25 M Al, with a pH of 7–9.8, the average 

radius and thickness of the particles were determined to be 880 Å and 50 Å, respectively. 

At reduced concentrations of 0.005 M Al, the radius and thickness of the particles were 

found to be ~200 Å and 40 Å, respectively. The “platelet” shape that was assumed in the 

analysis was found to be representative of aged, dried gels; gibbsite, as well as boehmite, 

has a tendency to form platelets. 

 

*

 Reprinted from Languir, D. Dabbs and A. Aksay, “Precipitation and Deposition of Aluminum-Containing 

Phases in Tank Wastes,” 21, pp. 11690–11695, Copyright 2005, with permission from The American 

Chemical Society. 
52

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53

7. ALUMINUM SOLUBILITY IN ALKALINE SOLUTION 

Previously, it was noted that aluminum at concentrations comparable to the Test 1 system 

formed on cooling but that no precipitates were observed at 60°C during the test. 

(However, it should be mentioned that a precipitate was observed after several weeks 

while the solution was being stored in an oven.) To assist in the understanding of this 

system, we examine the solubility of aluminum in alkaline systems to understand the 

sensitivity of the aluminum solubility with respect to pH and temperature. 

7.1. Predictions of Aluminum Solubility in Alkaline Solutions 

The aluminum solubility in alkaline solution is determined by the dissolving reaction of 

the aluminum hydroxide [Eq. (1)], where the equilibrium is a function of aluminum 

hydroxide phase. Therefore, the solubility product for Eq. (1) is very dependent on the 

aluminum hydroxide solid phase. For amorphous aluminum hydroxide, a solubility 

product of 10-31.2 at 25°C has been reported, whereas for gibbsite, a solubility product of 

10-33.9 has been reported.59 Differences in identifying the controlling phase of the 

precipitate have led to discrepancies in the predicted solubility of aluminum species. 

In alkaline or acidic solution, aluminum can exist as the following forms: Al3+, Al(OH)4

–

, 

AlOH2+, and Al(OH)2

+

. In an equilibrium solution, the following three reactions are in 

their equilibriums: 

Al3+ + H2O = AlOH2+ + H+ , (9) 

Al3+ + 2H2O = Al(OH)2

+

 + 2H+, (10) 

and 

Al3+ + 4 H2O = Al(OH)4

-

 + 4H+. (11) 

The equilibrium data for the above reactions can be found in Ref. 11. The equilibrium 

solubility of different forms and the corresponding total aluminum concentration in the 

solution are shown in Figure 41.59 The figure indicates that the main form is Al(OH)4

–

 in 

alkaline solution, whereas the aluminum ions are not stable. A comparison of Figure 41a 

and Figure 41b indicates that the amorphous aluminum hydroxide yielded soluble 

concentrations that were significantly larger than the concentrations for the crystalline 

gibbsite. The total aluminum concentration approached its smallest value near the neutral 

point and increased with an increasing pH in alkaline solution and decreased with pH 

increasing in acidic solution. The figure can be used to predict the aluminum 

concentration at 25o

C, whereas the solubility at other temperatures can be obtained by 

using a Van’t Hoff relationship. 

Experimental results have shown that the polynuclear aluminum complexes play no role 

unless the aluminum concentration is moderately high (>0.1 M).37 Confirmation that the 

aluminate ion is the only significant species in alkaline hydroxide solutions containing 
54

<~1.5 M total aluminum up to temperatures of 100°C has been provided via Raman and 

NMR studies.11

At 25°C an estimate of the solubility is 1.5 mg/L for gibbsite and 25 mg/L for amorphous 

aluminum hydroxide based on Figure 43.13 By comparison, the concentration of 

aluminum in Test 1 solution measured after 4 months was determined to be ~49 mg/L at 

room temperature. When we consider that boron existed in the Test 1 solution, aqueous 

boron could have been inhibiting precipitation, or the solution had not yet attained 

thermodynamic equilibrium. Furthermore, the concentration of boron in the solution 

indicated that a substantial quantity of boron had precipitated along with the aluminum 

(i.e., the concentration of boron was determined via ICP analysis to be 2000 mg/L, as 

compared with an initial 2800 mg/L). During the examination of the aged-4-months Test 

1 solution, it was observed that the pH had risen from 9.5 to 9.85. The deviation between 

the measured aluminum concentration and the theoretical prediction may be attributed to 

any number of factors: (1) the possible presence of complexing agents (organics), (2) the 

noted variability in the thermodynamic data, or (3) the complexation of aluminum with 

boron. An indication that the third factor may play a role is the fact that the pH had 

changed (increased) over the course of the aging process, which would indicate that 

further precipitation was occurring in the solution. 

Aluminum solubility may also be obtained at 60°C using the Van’t Hoff relationship. (It 

can be shown due to the weak dependency of the enthalpy on temperature over the small 

temperature range of interest, i.e., 25°C–60°C, that the Van’t Hoff approximation is 

appropriate.) For ICET 1 conditions, i.e., a pH of 9.6, an aluminum solubility of 48 mg/L 

is predicted if gibbsite is the controlling phase, whereas if the amorphous form of 

aluminum hydroxide is assumed, a solubility of ~1.3 gL is predicted. 

Based on the concentrations measured in Test 1, it would appear that the amorphous 

phase may be the phase controlling aluminum concentration. That is, the concentration of 

aluminum in Test 1 at 60°C was above the solubility limit of gibbsite and below the 

predicted solubility limit of the amorphous phase. Thus, the aluminum hydroxide may 

indeed be soluble under the isothermal Test 1 conditions. Additional support for this 

hypothesis is that the concentrations of measured aluminum were equivalent for both 

filtered and unfiltered water samples. Further support for the amorphous phase being the 

phase controlling species is given by the observation that the wet precipitate, containing 

~51 wt % Al, was found to be largely amorphous (Figure 19). Furthermore, the ICP 

analysis revealed that boron also was present and, as is discussed in Section 5, the 

presence of boron, even in very small quantities relative to aluminum, precludes 

crystallization. 
55

Figure 41. Solubility of (a) amorphous Al(OH)3, and (b) gibbsite [Al(OH)3] as a function 

of pH at 25°C. Also shown are lines indicating the solubility concentrations 

of Al3+ and individual hydroxyl complexes. [Used with permission of the 

publisher.*

] 

 

*

 Reprinted and electronically reproduced from Aqueous Environmental Geochemistry, Donald Langmuir, 

Copyright 1997, with permission from Pearson Education, Inc., Upper Saddle River, New Jersey. 
56

 

7.2. Effect of Organic Complexing Agents on Aluminum 

Solubility 

As previously noted in Section 7.1, the measured concentrations of aluminum were 

somewhat higher at room temperature than would be expected, even if the solubilitycontrolling aluminum hydroxide phase were amorphous. To explore this phenomenon, it 

has been noted that organics can increase the solubility.60 Consequently, in this 

subsection we review numerous studies that have been conducted to investigate the effect 

of organics on the solubility of aluminum. However, before doing so it is first instructive 

to review some basic theory on the action of the complexing agents. 

The stability of a complex is usually described by its stability constant, which is also 

known as its formation or association constant. The stability constant may refer to a 

stepwise or cumulative complex formation reaction. Kassoc applies to the simple reaction 

M + L = ML , (12) 

where M, L, and ML are the cation and ligand ML complex, respectively. The free 

energy of the reaction may be written as 

ΔG = ΔH-TΔS . (13) 

Thus, the stability of the complex depends on both the enthalpy of the reaction and the 

entropy. For most complexation reactions, the entropy stabilizes the reaction.59 The net 

positive entropy change that reflects structural changes in the solution consequent to 

complexation may be broken into many terms, as expressed in 

ΔSo = ΔSo

netchg + ΔSotr + ΔSovibr + ΔSodehydr , (14) 

In those cases in which there is an entropy reduction ΔSo

netchg because of the decrease in 

the number of charged particles, some of the translational entropy, ΔSo

tr, of the separate 

cation and ligand is converted to rotational and vibrational entropy of the complex. 

However, the chief contribution is usually the ΔSo

dehydr term that reflects partial 

dehydration of the cation and/or ligand. Complexation leads to a breakdown of the 

structured water of hydration, especially around the cation, with a resultant decrease in 

the order of the solution. Therefore, ΔSo

dehydr is proportional to the number of water 

molecules displaced by the ligand. Thus, ΔSo

dehydr is greatest for multivalent ion 

complexation because these species are initially the most hydrated. Because chelating 

agents bond a cation via more than one electron-donating atom, they tend to free several 

cation-coordinating waters on complexation. The result is an unusually large +ΔSo

, as 

opposed to that of monodentate complexes.59,61

Speciation calculations for typical hydroxyl organic ligands and their aluminum(III) 

complexes may be performed based on established protonation constants and reaction 
57

quotients.62 Citric acid has been investigated and found to establish three aluminum(III) 

complexes 1, 2, 3: the 1:1 complex and its protonated and deprotonated forms. 

 

Another possible chelating agent that has been studied is catechol. Catechol has been 

observed to have the highest stability. However, because of its high pKa’s, strong 

competition by hydrogen ions occurs, even in alkaline solutions. Ethylenediaminetetraacetic (EDTA) is another polydentate complexing agent that forms complexes with 

many metals. The anion of EDTA has a quadruple negative charge. The four carboxylate 

ion groups and also the two nitrogen atoms can form bonds with aluminum; the anion is 

accordingly a hexadentate complexing agent. Another promising chelating agent is 

desferriferrioxamine (Desferal R DFB). This agent is a microbial siderophore that has 

been used for the treatment of Colley’s anaemia and is now being used for the treatment 

of aluminum overload.62 

Experimental evidence for the effectiveness of chelating agents has been found.10 In these 

investigations, small quantities of FA (0–50 mg) were added to aluminum chloride 

solutions, 450 ml of a 10-3 M solution. (FA resembles citric acid in that it contains CO2H 

and resembles aliphatic OH groups and quercetin in that it also contains phenolic 

hydroxyl and ketonic C=O groups. It is through these functional groups that FA can form 

stable complexes with aluminum.) The solutions were then titrated with 0.1 M NaOH to 

raise the pH to 6, 8, and 10. All experimental solutions were allowed to age at 30o

C for 

70 days, with occasional shaking while maintaining pHs at the initial values. It was found 

that FA inhibited the formation of gibbsite at a pH of 6, leading to the formation of 

pseudoboehmite. In systems with a pH of 10, the addition of as little as 0.0125 mg/L of 

FA was sufficient to prevent the formation of any precipitate. A simple explanation may 

be that at a pH of 10, the surface of the aluminum and FA are both negative such that 

electrostatic repulsion between the two components appears responsible for the lack of 

precipitation. 

Polybasic acids have also been studied in conjunction with altering the solubility of 

aluminum hydroxide.13 In this study, 20 ml of 0.1 M AlCl3 was titrated with 0.1 M 

NaOH. Before titration, an appropriate quantity of potassium acetate, oxalate, or citrate 

was dissolved in AlCl3 solution. It was found that increasing the amounts of the organic 

acid delayed the onset of precipitation and, in the case of citric acid, precluded 

precipitation if a molar ratio of citric acid to aluminum were above 0.475. An infrared 

analysis of the air-dried 1:1 citrate-to-aluminum complex indicated that the complex may 

be similar to commercial aluminum citrate. From these studies it appears that the strength 

of the aluminum-anion bond is ordered as citrate>oxalate>acetate. Figure 42, Figure 43, 

and Figure 4413 present the results of this investigation. In the figures, the arrows denote 

the appearance of visible precipitate. 
58

Figure 42. Effect of KOAc on the potentiometric titration of 0.1 AlCl3 by 0.1 N NaOH. 

The molar ratio of acetate to aluminum was 0, 1.0, 2.0, and 3.0. [Used with 

permission of the publisher.*

] 

Figure 43. Effect of K oxalate on the potentiometric titration of 0.1 AlCl3 by 0.1 N 

NaOH. The molar ratio of acetate to aluminum was 0, 0.5, 1.0, 1.5, 2.0, and 

3.0. [Used with permission of the publisher.*] 

 

*

Reprinted from Clays and Clay Minerals, M. K. Wang et al., “Influence of Acetate, Oxalate, and Citrate 

Anions on Precipitation of Aluminum Hydroxide,” 31:1, pp. 65–68, Copyright 1983, with permission 

from Clays and Clay Minerals, The Macaulay Institute.. 
59

Figure 44. Effect of K citrate on the potentiometric titration of 0.1 AlCl3 by 0.1 N NaOH. 

The molar ratio of acetate to aluminum was 0, 0.45, 1.0, 2.0, and 3.0. [Used 

with permission of the publisher.*

] 

In a study of aluminum silicate scale formation and inhibition, it was found that chelating 

agents with carboxylate functionality, such as citric acid, acetic acid, and EDTA, may 

inhibit aluminum silicate scale formation.63 Nitrilotriacetic acid, diethylenetraminepentacetic acid, polycarboxylic acids, pyrocatechol, and humic acids may also be effective.62

Figure 45. Particle growth in the presence of citric acid. [Used with permission of the 

publisher.†

] 

 

*

Reprinted from Clays and Clay Minerals, M. K. Wang et al., “Influence of Acetate, Oxalate, and Citrate 

Anions on Precipitation of Aluminum Hydroxide,” 31:1, pp. 65–68, Copyright 1983, with permission 

from Clays and Clay Minerals, The Macaulay Institute. 

†

 Reprinted from Languir, D. Dabbs and A. Aksay, “Precipitation and Deposition of Aluminum-Containing 

Phases in Tank Wastes,” 21, pp. 11690–11695, Copyright 2005, with permission from The American 

Chemical Society. 
60

It is important to note that the dynamics of nucleation are also affected by the presence of 

complexing agents. In investigations using aluminum chloride with citric acid, using 

quasi elastic light scattering, it has been found that the addition of citric acid changed the 

dynamics of both solutions and suspensions.57 In solutions of very high hydrolysis 

(OH/Al = 3.29), particles did not form when the mole/mole ratio of citrate to aluminum 

equaled 0.8. Under the same hydrolysis conditions but at citrate/Al ratios <0.8, particles 

formed but nucleation and growth were markedly slowed. Under conditions of high 

hydrolysis (OH/Al = 2.46), particles formed in the presence of citric acid when base was 

added, but the rate of formation was slowed. Figure 45 presents the results of this 

investigation.57 
61

8. ALUMINUM CORROSION IN ALKALINE SOLUTION 

8.1. Corrosion Mechanism 

It is well known that aluminum metal is very inert in neutralized solution, whereas pure 

aluminum is too reactive to be used in concentrated alkaline solutions.64 Experimental 

results show that two competing processes occur at the aluminum metal surface: direct 

dissolution of the aluminum metal and electrochemically formation/dissolution of the 

aluminum hydroxide films. The first process is very intense, which leads to a high 

corrosion rate. With time, a film forms on the metal surface, which acts as a barrier for 

species transport. Therefore, the corrosion rate is significantly reduced after the film 

forms. 

When we consider that Al(OH)4

− is the only stable form of aluminum in alkaline solution, 

as shown in Figure 41, the direct dissolution of aluminum is due to 

− − + 2 + = H2 + Al(OH)4 2

3 Al . (15) 3H O OH

It is evident that the direct aluminum dissolution is accompanied by hydrogen production. 

Gas bubbles were observed in experiments at the aluminum metal surface. Therefore, the 

corrosion of pure aluminum in an alkaline solution proceeds mainly by water reduction, 

according to Eq. (15). The reaction indicates that the dissolution rate depends on the 

concentrations of OH−

 and aluminate ions Al(OH)4

−

 at the solid/liquid interface. Thus, 

the transportation of OH−

 and Al(OH)4

−

 through the solution to and from the interface, 

respectively, is expected to play an important role in the aluminum dissolution rate.64 If 

the reaction at the interface is much faster than the rate of mass transport provided by the 

flowing solution, then the dissolution rate is determined (or limited) by the mass transport 

efficiency. Thus, the aluminum corrosion rate will increase with increasing ability to 

remove the corrosion products, such as by increasing the solution velocity. However, if 

the transportation rate is greater than the reaction rate, the aluminum corrosion rate is 

determined by the reaction rate and depends little on the transportation of OH−

 and 

Al(OH)4

− through the solution, which has been reported by several researchers.65,66 

The electrochemical formation of hydroxide film on a pure aluminum surface in alkaline 

solution has been verified experimentally at an open circuit by the increases in the circuit 

potential obtained from the moment just after interrupting the abrading action on the 

specimen.67 This increase in potential has been considered by several other 

authors.64,67,68,69,70,71 The potential changes at the aluminum surface in alkaline solution 

as a function of time are shown in Figure 46.71 The potential increases with increasing 

time until it reaches a constant value. The increase in potential with time is known to 

arise from the growth of a surface oxide film.72
62

Figure 46. Changes in film formation potential obtained from the pure aluminum rod 

specimen with time in 10-3 M NaOH solution at various applied anodic 

current densities of Ο , 5 mA cm-2; , 10mA cm-2; Δ, 20mA cm-2; • , 50mA 

cm-2; and , 100mA cm-2. [Used with permission of the publisher.*

] 

As discussed in Ref. 64, the film formation is due to 

− − Al + 3OH = Al(OH) + 3e 3 . (16) 

The film forms because of the inward diffusion of − OH through the film. When a film 

with considerable thickness forms, the direct dissolution of aluminum metal ceases, as 

well as the hydrogen production. After that, the aluminum corrosion can be classified into 

a direct metal dissolution by the movement of aluminum ions through the film and an 

indirect metal dissolution by consecutive oxide film formation and dissolution.64 Because 

3+ Al is not thermodynamically stable in an alkaline solution, the direct ejection of 

aluminum ions from the film into the solution can never occur, which was shown in Refs. 

71 and 73. Thus, aluminum corrosion after the film formation occurs is due to 

− − Al(OH)3 + OH = Al(OH)4 , (17) 

which is indicative of an electrochemistry process. Thus, the corrosion of pure aluminum 

in an alkaline solution can be divided into two substeps of a partial anodic reaction 

comprising the electrochemical formation and chemical dissolution reactions of the film 

and a partial cathodic reaction of the water reduction reaction.64

In some applications, oxide layers are pre-formed on aluminum metal surfaces to protect 

the substrate. The oxide layer, composed of Al , is compact and protective in neutral 2O3

 

*

Reprinted from J. Solid State Electrochem., S. M. Moon and S. I. Pyun, “Growth Mechanism of Anodic 

Oxide Films on Pure Aluminum in Aqueous Acidic and Alkaline Solutions,” 2, pp. 156–161, Copyright 

1998, with kind permission of Springer Science and Business Media. 
63

solution. However, the film can be attached by OH−

 in an alkaline solution, resulting in 

film dissolution. If the oxide films were completely uniform, both physically and 

chemically, it would be expected that the thinning would be uniform over the entire 

surface. If not, it would be expected that the normal dissolution process would be a flowassisted process or a flow-centered process. This latter process was found to be mainly 

responsible in a series of experiments.74 The dissolution of the oxide film is due to 

− −

2 3 + + 2 = 4 Al O 2OH 3H O 2Al(OH) . (18) 

It should be noted that the dissolution of the aluminum oxide is much slower than the 

direct dissolution of aluminum metal [Eq. (15)]. On the other hand, the aluminum oxide 

can be converted into aluminum hydroxide if OH−

 ions take up the vacancies left by 

oxygen inward diffusion according to 

3

2

Al VO 3OH Al(OH) 2

3 Al (ox) + + = + − . (19)

Therefore, with time, the oxide layer may be converted to a hydroxide layer. 

8.2. Corrosion Rate 

The corrosion rate of aluminum in alkaline solution is a function of the operation 

conditions, such as the temperature, the pH, and the solution properties. Specifically, the 

corrosion rate depends on the time in a static solution. Because the aluminum 

concentration in the solution increases, the average corrosion rate decreases with time 

elapsing. Therefore, various authors obtained very different corrosion rates of aluminum 

in an alkaline solution because of the different exposure times. In this subsection, we will 

review experimental data on aluminum corrosion in an alkaline solution. 

It has been experimentally verified that the corrosion rate depends logarithmically on the 

pH,69 with different dependencies being noted on either side of the pH at which the 

corrosion rate is at its minimum (close to pH = 6). These results were reported by 

Pourbaix et al.75 Vujicic and Loverecek76 obtained a different logarithmic relation 

between the corrosion rate and pH value. The curves of corrosion rate as a function of 

pH76 are shown in Figure 47. It is evident that for the systems considered in the reference, 

the point of minimum corrosion rate is greater than a pH of 7.0, which is greater than the 

pH value, as shown in Ref. 75. Long-term corrosion test results of aluminum in sodium 

hydroxide solution over the pH range of 8–13 were carried out by McKee and Brown77

for up to 1 week and over the pH range of 7–12 were carried out by Tabrizi et al.69 up to 

80 days. A summary of corrosion rate data of aluminum in alkaline solutions is given in 

Table 9. 
64

Figure 47. Corrosion rate as function of pH in systems (1) aluminum in 4 10− M KCl; 

(2) aluminum in 4 10− M KCl containing 4 10− M ethylenediamine; and 

(3) aluminum in 4 10− M KCl containing 3 10− M ethylenediamine. [Used with 

permission of the publisher.*

] 

Table 9 indicates that the corrosion rate depends on the exposure time, which is clearly 

shown in Figure 48.69 Over 40 days of immersion, the weight loss increased with time 

elapsing. For immersion times beyond 40 days, in a solution having a pH of 10 and 11, 

the weight losses were similar to those measured at 40 days, indicating that no significant 

weight loss occurred during the last 20 days. Generally similar trends were evident for 

immersion at 60o

C,69 with initial high rates of weight loss that decreased with immersion 

time, particularly beyond 20 days. The weight losses were converted into corrosion rates, 

as shown in Figure 49. It is evident that the corrosion rate of aluminum in an alkaline 

solution depends strongly on the exposure time. It decreases with the exposure time 

increasing. The dependence on exposure time occurs because of the aluminate 

concentration increasing with time, where the rate is determined by the ratio of the 

sample surface to the solution volume. Therefore, when we discuss the corrosion rate of 

aluminum, we must consider the surface/volume ratio and the exposure time. This 

dependence can also explain why different authors obtained different corrosion rates at 

comparable conditions, e.g., temperature and pH, in the same solution at the same 

temperature. 

 

*

Reprinted from Surface Technology, V. Vujicic and B. Lovrecek, “A Study of the Influence of pH on the 

Corrosion Rate of Aluminum,” 25:1, pp. 49–57, Copyright 1985, with permission from Elsevier. 
65

Table 9. Corrosion Rate of Aluminum in Static Alkaline Solution 

pH Temperature Corrosion Rate Exposure Time Comments 

25oC 45 gm−2

y−1 1 hour Ref. 76 

25oC <30 gm−2

y−1 7 days Ref. 77 

9 

60o

C 242 gm−2

y−1 7days Ref. 78, alloy 

25oC 355 gm−2

y−1 1 hour Ref. 76 

25oC 55 gm−2

y−1 7 days Ref. 77 

30oC 110 gm−2

y−1 40 days Ref. 69, solution replenished 

30oC 130 gm−2

y−1 40 days Ref. 69, 1000 ppm Cl−

60oC 165 gm−2

y−1 40 days Ref. 69, solution replenished 

60oC 65 gm−2

y−1 40 days Ref. 69, 1000 ppm Cl−

10 

60o

C 670 gm−2

y−1 7 days Ref. 78, alloy 

30oC 270 gm−2

y−1 40 days Ref. 69, solution replenished 

30oC 285 gm−2

y−1 40 days Ref. 69, 1000 ppm Cl−

60oC 590 gm−2

y−1 40 days Ref. 69, solution replenished 

11 

60o

C 170 gm−2

y−1 40 days Ref. 69, 1000 ppm Cl−

Figure 48. The variation of weight loss of aluminum with exposure time at 30oC. [Used 

with permission of the publisher.*

] 

 

*

Reprinted from Corrosion Sci., M. R. Tabrizi, S. B. Lyon, G. E. Thompson, and J. M. Ferguson, “The 

Long-Term Corrosion of Aluminum in Alkaline Media,” 32, pp. 733–742, Copyright 1991, with 

permission from Elsevier. 
66

Figure 49. Corrosion rates based on weight loss shown in Figure 48. 

From Table 9, the corrosion rate is observed to be dependent on the temperature. Because 

the reaction and diffusion rates increase with temperature increasing, the corrosion rate of 

aluminum in alkaline solution becomes larger as the temperature increases, as shown in 

Table 9. By plotting the aluminum concentration in a 5-M KOH solution as a function of 

temperature, Chu and Savinell65 estimated the activation energy as ΔG ≈ 13.7kcal/mol

for the corrosion of aluminum in alkaline solution, which agrees very well with the value 

of 12.3 kcal/mol reported for aluminum near the open circuit.79 Chu and Savinell also 

reported that the film that formed on the metal surface became thin with increasing 

temperature, allowing easier OH−

 diffusion through the film and resulting in an increase 

of the aluminum dissolution rate. 

As discussed in Section 7.2, adding organic or inorganic ions into the solution has a 

significant effect on the aluminum concentration. As a result, these ions affect the 

corrosion rate. For example, the addition of 1000 ppm Cl−

 reduces the corrosion rate by 

two-thirds of the chloride-free solution, as shown in Table 9 at 60o

C and a pH of 10. The 

ions’ effect depends on the temperature. As indicated in the table, the addition of 1000 

ppm Cl−

 has little effect on the corrosion rate at 30oC. 

The corrosion rate can also be modified by changing the composition of the aluminum 

metal through adding small amounts of other elements, such as zinc, bismuth, tellurium, 

indium, gallium, lead, and titanium. The corrosion rate determination of various of 

aluminum alloys was conducted by Macdonald et al.;80 the results are presented in 
67

Table 10. The table indicates that the binary alloys containing <1% of the alloy elements 

exhibit corrosion rates in 4 M KOH at 50o

C that are higher than that for pure aluminum. 

Adding two or more alloying elements, such as aluminum alloys containing gallium, 

indium, tellurium, and phosphorus, can result in a sharp reduction in the corrosion rate to 

well below that for aluminum. The corrosion rate of Al-0.1%In alloy was also considered 

by Wilhelmsen et al.81 It was reported in Ref. 81 that Al-0.1%In exhibits excellent 

corrosion-resistant properties, which is very different from those in Ref. 80. 

In a static solution, the corrosion rate is a function of exposure time. With time, corrosion 

eventually can terminate when the corrosion product (aluminate ions) reaches its 

equilibrium concentration or solubility. The corrosion product does not reach an 

equilibrium concentration for a flowing system. Solution flow can carry corrosion 

products away from the metal and also can bring reactants to the metal surface, which 

accelerates the dissolution process. The dependence of corrosion rate on the flow velocity 

can be expressed simply, as shown in Figure 50.82 The figure indicates that at low 

velocities, the corrosion is controlled or partially controlled by mass transfer; in other 

words, the dissolution rate is greater than the mass transfer rate and the corrosion product 

interface concentration is saturated. In such cases, the thickness of the laminar mass 

transfer layer becomes thinner with increasing velocity and, as a result, the corrosion rate 

increases; when the velocity exceeds a critical value, the mass transfer rate becomes high 

enough to transport all corrosion products away from the interface. In this case, the 

corrosion rate is determined by the dissolution/reaction rate and is independent of the 

flow velocity. The corrosion is activation controlled; for very high velocities, the high 

shear stress at the interface can strip off the protective film on the surface of the structure. 

Some cavities appear at the interface, and the corrosion rate increases sharply with the 

flow velocity. Therefore, the dependence of the corrosion rate of aluminum in a dynamic 

alkaline solution on mass transfer is a function of the flow velocity. In different flow 

velocity ranges, the dependence may be much different. This conclusion can explain why 

several authors65,66 reported that the transport of OH−

 and Al(OH)4

−

 has little effect on the 

corrosion rate, whereas others64 found the corrosion rate increasing with the flow velocity 

increasing. Unlike the static solution, the corrosion rate is a constant in a flowing system 

when the flow is at a steady state. 
68

Table 10. Corrosion Rate on Various Aluminum Alloys in 4 KOH at 50oC80 [Used with 

permission of the publisher.*

] 

Element Composition (wt %) Corrosion Rate (mg/cm2/min) 

99.99 0.515 

Pure Al 99.999 0.876 

Alcoa1100-H14 2.735 

0.1 1.030 

0.5 1.074 

1.0 1.097 Zn 

5.0 1.190 

0.1 0.593 

0.5 0.461 

1.0 0.520 Bi 

5.0 0.507 

0.01 0.525 

0.05 0.584 

0.1 0.523 Te 

0.5 0.529 

0.01 1.910 

0.05 2.287 

0.1 1.965 In 

0.5 2.106 

0.01 8.740 

0.05 5.707 

0.1 5.845 Ga 

0.5 7.627 

Pb 0.5 0.736 

0.2 0.918 P 0.5 0.846 

0.1% P, 0.1% Ga 153.351 

0.1% P, 0.1% In, 0.2% Ga, 0.01% Ti 0.057 

0.07% In,0.2% Ga, 0.01% Ti 0.051 

0.05% In, 0.01% Ti 0.980 

0.05% In, 0.05% Ti 0.733 

0.2% In, 0.05% Ti 0.621 

0.25% In, 0.01% Ga, 0.1% Ti 0.041 

0.1% In, 0.2% Ga, 0.1% Ti 0.048 

 

*

 Reprinted from Corrosion, D. D. Macdonald, K. H. Lee, A. Moccari, and D. Harrington, “Evaluation of 

Alloy Anodes for Aluminum-Air Batteries: Corrosion Studies,” 44, pp. 652–657, Copyright 1988, with 

permission from NACE International, 2006. 
69

Figure 50. Dependence of corrosion rate on the flow velocity. [Used with permission 

of the publisher.*

] 

8.3. Corrosion Inhibitor 

Inhibition of the corrosion of aluminum and its alloys in alkaline solutions is well known, 

and many inhibitors have already been described and investigated;83 they may be 

organic,84 or inorganic.85 Using polarization methods as well as weight loss methods, AlSuhybani et al.86 studied in detail various organic and inorganic inhibitor effects on the 

corrosion rate of aluminum in alkaline solutions. It was reported that benzoin acids 

(BAs), as well as other organic acids, inhibit the corrosion but to different extents, 

depending on the structure, and the inhibition is due to the adsorption of these acids and 

not to the neutralizing effect. The effect of some organic and inorganic additives on the 

corrosion or aluminum in NaOH is presented in Table 11. It is evident that the citric acid 

shows the best inhibition at 5 M NaOH, as compared with other organic inhibitors. Some 

inorganic inhibitors in the table, such as chromate, metabisulphate, dihydrogen 

phosphate, disodium meta borate, and dihydrogenphosphate, are the best, with an 

inhibition efficiency of ~97.0%, whereas others, such as NaCl, NaNO3, Na2SO4, and 

NaIO3, are activators. 

 

*

Reprinted from Journal of Nuclear Materials, F. Balbaud-Celeier and F. Barbier, “Corrosion of Aluminum 

in Alkaline Solutions,” 289:1, pp. 227–242, Copyright 2001, with permission from Elsevier. 
70

Table 11. Effect of Some Organic and Inorganic Additives on the Corrosion of 

Aluminum in NaOH Solutions86 [Used with permission of the publisher.*

] 

(NaOH) Additive M Efficiency (%) 

Citric Acid 1.0 99.4 

2-Amino BA 1.0 15.6 

Tetraethylenediaminetetraactic Acid 1.0 37.3 

Chloramine-T 1.0 66.4 

Dithizone Saturated 78.5 

Sodium Chromate 1.0 45.5 

Sodium Metabisulphate 1.0 55.9 

Sodium Dihydrogen Phosphate 0.01 8.3 

Sodium Dihydrogen Phosphate 0.05 31.0 

5 M 

Sodium Dihydrogen Phosphate 1.0 96.9 

Disodium Metaborates 1.0 36.5 

NaCl, NaNO3, Na2SO4, NaIO3 1.0 Activators 

Dithizone 0.25 21.1 

Dithizone Saturated 47.5 

3M 

Polyvinylalcohol Saturated 12.6 

Al-Suhybani et al.86 reported that the inhibition efficiency depends on both the sodium 

hydroxide and the inhibitor concentrations. Increasing the sodium hydroxide 

concentration leads to a decrease in the inhibition efficiency, whereas the dependence on 

the inhibitor concentration is more complex. Figure 51 shows the dependence of 

inhibition efficiency on the concentration of several organic inhibitors. It is evident that 

the dependence is not linear. Specifically, the curve has a double-S shape for phthalic 

acid. The shape may indicate the formation of the second layer of adsorbed molecules. 

The figure also shows that phthalic acid and 3-hydroxy BA have stimulating effects at 

low concentrations, whereas citric acid inhibits the corrosion even at low concentrations. 

The effect of ethelencdiaminetetraacetic acid (EPTA) is also shown in the figure. The 

compound was found to be effective only at ~0.5 M. At low concentrations, it stimulates 

corrosion to the extent of ~90% at 0.01 M. 

 

*

 Reprinted from Materialwissenschaft und Werkstofftechnik, A. A. Al-Suhybani, Y. H. Sultan, and W. A. 

Hamid, “Corrosion of Aluminum in Alkaline Solutions,” 22:8, pp. 301–307, Copyright 1991, with 

permission from Wiley-VCH Verlag GmbH & Co KG. 
71

Figure 51. Variation of inhibition efficiency with the concentration of some organic 

acids in 3 M NaOH. [Used with permission of the publisher.*

] 

Figure 52. Potential-time curves for almasilium alloys immersed in 0.1 NaOH solution at 

60o

C, showing the effect of aluminate and “silicate” ions: curve a, 0.1 N NaOH 

alone; curve b, 0.03 M Si; curve c, 0.04 M Si; curve d, 0.04 M Si + 0.0003 M Al; 

curve e, 0.04 M Si + 0.0006 M Al; curve f, 0.04 M Si + 0.0012 M AI; and curve g, 

0.04 M Si +0.0018 M AI. [Used with permission of the publisher.†

] 

 

*

 Reprinted from Materialwissenschaft und Werkstofftechnik, A. A. Al-Suhybani, Y. H. Sultan, and W. A. 

Hamid, “Corrosion of Aluminium in Alkaline Solutions,” 22:8, pp. 301–307, Copyright 1991, with 

permission from Wiley-VCH Verlag GmbH & Co KG. 

†

 Reprinted from Thin Solid Films, J. P. Labbe and J. Pagetti, “Study of an Inhibiting Aluminosilicate 

Interface by Infrared Reflection Spectroscopy,” 82, pp. 113–119, Copyright 1981, with permission from 

Elsevier. 
72

The sodium silicates have been found to be an effective inhibitor of corrosion in 

aluminum-alkaline systems.83,87 The inhibition is due to the formation of amorphous 

aluminosilicate film on the metal surface. It has been reported that a total and almost 

instantaneous inhibition could be obtained if both SiIV and AlIII are present in the 

solution. The same phase could be built up more slowly with silicate ions alone. The 

solution potential for different silicate concentrations was also measured as a function of 

time. The results are shown in Figure 52. After an initial drop in the solution potential 

because of the dissolution of aluminum, the potential increased with time, followed by a 

series of oscillations that revealed a tendency for the inhibition film to become alternately 

partially or totally inhibiting.87 Finally, the potential reached a stabilization value. The 

larger the silicate concentration, the greater the stabilization, as shown in the figure. 

Another effective inorganic inhibitor of the corrosion of aluminum in an alkaline solution 

was calcium tartrate, which was studied by Shao et al.88 The authors reported an 

inhibition efficiency >98% for the corrosion of pure aluminum in a 1-M KOH solution. 

When the inhibitor concentration was low, its inhibition effect for the anodic process was 

much higher than that for the cathodic process. 
73

9. BENCHMARK EXPERIMENTS 

To assist in developing a surrogate for the Tests 1 and 5 solutions, a series of benchtop 

tests was performed. The objectives of these tests were to (1) reduce the complexity of 

the Test 1 system to facilitate the production of a surrogate that could be used to perform 

head-loss testing and (2) gain additional understanding of the chemical speciation and the 

mechanisms for their formation. That is, by understanding the mechanisms for the 

formation of these species, we may be able to add chelating agents to keep the 

detrimental species in solution and avoid additional head loss due to chemical products. 

In the previous benchtop tests, we addressed the solubility of the aluminum and the 

properties of the precipitate. Several small experiments were carried out at LANL. In this 

section, we report our experimental results. 

9.1. Test 1 Redissolution Test 

To investigate further the solubility properties and the kinetic aspects of the approach to 

equilibrium for the Test 1 solution, an experiment was conducted in which the solution 

after 4 months of storage was heated and held at 60o

C for 1 month. The pH during this 

test was lowered from 9.85 to 9.66, indicating that equilibrium had not been reached. 

(The initial pH before aging in ICET was 9.5.) Supernatant from the solution was then 

obtained, and aluminum nitrate at a concentration of ~0.02 M was titrated into the 

solution while the solution temperature was ~60°C. No noticeable precipitation occurred 

during this process. The lack of precipitation during the titration of the 0.02-M aluminum 

nitrate addition indicates that the solution may have been undersaturated with respect to 

aluminum at 60°C. The lack of precipitation can be explained by recognizing that the 

solution had previously precipitated while cooling to room temperature, where it was 

maintained for ~4 months. Following this precipitation process, the solution 

concentration was found to be ~48 mg/L. If the kinetics of the dissolution of the 

previously formed precipitate were slow, little precipitate would have redissolved, and 

the concentration of the solution at 60°C would not be much higher than the aluminum 

concentration at 25°C. Consequently, as may be seen from the solubility calculation in 

Section 5.1, the aluminum would have been undersaturated at 60°C. (The solubility at 

60°C is ~1.3 gL, assuming that the phase-controlling solid is amorphous aluminum 

hydroxide.) Therefore, the addition of a 0.02-M solution would not induce precipitation 

at a pH of ~9.6. 

After a small quantity of aluminum nitrate was added, the solution was allowed to cool at 

room temperature. No visible precipitation was observed after a few hours of aging. 

Subsequently, the solution was refrigerated to ~0°C and observable precipitation 

occurred. Finally, the solution was reheated and it was observed that the precipitate still 

remained 1 month after initial refrigeration. This result reinforces the point that either the 

reversibility is incomplete or the kinetics of redissolution is very slow, and equilibrium 

calculations may be of limited use in the 30-day period over which the ICET tests were 

operated. 
74

9.2. Solubility of Aluminum-Based Surrogates 

Benchtop tests have been performed to investigate the solubility limit of aluminum in a 

2800-mg/L boron, pH = 9.6 NaOH system at 60o

C. Two methods of introducing 

aluminum into solution have been examined. In the first method, aluminum metal (small 

pieces) initially was introduced at ambient temperature into the solution, followed by 

storage in an oven, where the temperature was maintained at 60o

C. It should be noted that 

the metal dissolved very rapidly into the solution because the metal was produced in an 

inert atmosphere and apparently did not have an oxide layer to limit the dissolution 

process. Thus, the initial dissolution was very rapid, resulting in vigorous bubbling 

during the first few hours of dissolution. Furthermore, it cannot be ensured that the 

solubility limit of ~25 mg/L at room temperature was not exceeded initially. After a few 

weeks, all of the aluminum metal appeared to have dissolved into the solution. A 

considerable amount of precipitate settled to the bottom of the bottle. Concentration 

measurements on material that was filtered through a 0.7-mm filter were made via ICP 

and indicated an aluminum concentration of 126 mg/L in the filtered solution. During the 

course of the dissolution of the metal, the pH dropped to 9.25. Normalizing to the ICET 

conditions, pH = 9.6, we found an aluminum solubility of 282 mg/L. The second method 

of introducing aluminum into solution consisted of introducing aluminum nitrate crystals 

into the base solution at 60o

C. In this system, the solubility at a measured pH = 8.69 was 

determined to be 80 mg/L. Converting to the Test 1 solution pH of 9.6 yielded a 

corresponding solubility of 650 mg/L. 

ICP measurements for both systems did not indicate any discernable quantity of boron 

precipitated from solution. It should be noted that in this test, the solution was not 

allowed to cool and form precipitate before a sample of the solution was obtained for ICP 

analysis. This boron deficiency is in contrast to the relatively high percentage of boron 

found in the precipitate from the Test 1 solution. 

9.3. Effects of the Boron in the Aluminum-Alkaline Solution 

To examine the boron effects, an aluminum coupon was added to a solution comprised of 

2800 mg/L boron and sodium hydroxide at a pH of 9.5. The solution was placed in a 

60o

C oven for 2 months and then hot filtered at the same temperature. The results of an 

XRD examination revealed that the precipitate was extremely disordered. A Rietveld 

refinement analysis revealed that the crystal cell dimensions were a = 2.83738, 

b = 12.26414, and c = 3.68170 (Å), with a crystallite size of 29 (±1) Å. Strain, if present, 

was very small. The coherent scattering domain was <10 unit cells, consequently 

resulting in very broad peaks. This phenomenon reveals similar cell dimensions for the 

amorphous precipitate. 

A comparison of the crystal cell size with the XRD analysis of the Test 1 precipitate 

revealed an extremely close match. Furthermore, a thermogravimetric analysis (TGA) of 

the surrogate precipitate was performed. Figure 53 and Figure 54 present the results of 

the TGA analysis for the surrogate and ICET precipitate, respectively. A comparison of 
75

Figure 53 and Figure 54 indicates that the surrogate precipitate and the Test 1 precipitate 

have comparable properties. 

Figure 53. TGA for surrogate, where the red line is weight loss and the blue line is the 

derivative. 

Letellier Experim ent

70

75

80

85

90

95

100

105

0 100 200 300 400 500 600 700 800 900

Temperature

Weight %

ICET 1 B ottom Sludge ICET-T1 G el Tinc alc onite

Figure 54. TGA for Test #1 precipitate. 

As has been previously reported, it would appear that on cooling, the adsorption of boron 

onto the surface of the amorphous aluminum hydroxide precluded crystallization. To 

further investigate the nature of the solid phase formed when the solution was cooled, an 

27Al NMR analysis was performed. To perform these analyses, the solution was first 

centrifuged to isolate the solid. The sample was then dried at 100°C overnight. The 27Al 

spectra are shown Figure 55a. The figure indicates that many chemical shifts exist; in 

addition to the broadening of the octahedrally coordinated aluminum, the shift at 60 ppm 
76

is similar in nature to the shift reported for tetrahedral aluminum with four boron nextnearest neighbors.89 In the reference, it was suggested that tetrahedrally and triagonally 

coordinated boron are both present. The relative amount depends on the B/Al ratio. To 

understand further the coordination between aluminum and boron solid-state 11B, NMR 

was performed, and the curve is shown in Figure 55b. From peaks in Figure 55b, we see 

that both tetrahedral and triagonal boron is coordinated with the aluminate ion. 

For further investigating the effects of boron on chemical speciation, sodium hydroxide, 

with a pH of 12.74, was titrated into a solution composed of ~2800 mg/L of boron and 

0.05 M of aluminum nitrate. A white precipitate formed as soon as the first drop of 

sodium hydroxide was added. The precipitate dissolved in ~1 hour. Another experiment 

was performed by adding sodium hydroxide into the aluminum/boron solution to raise the 

pH of the solution to 9.7. The solution became opaque and remained so for 24 hours, 

whereupon significant settling of the gel-like material led to the appearance of two 

distinct phases. The bottom phase had a milky appearance with apparent gel-like 

characteristics, and the top liquid phase was relatively clear in appearance.

After ~2 weeks, 27Al NMR analyses were performed, as shown in Figure 33. The figure 

shows that as the temperature was lowered from 60o

C, the appearance of a chemical shift 

at ~74 ppm was noted. The observed chemical shift was similar to that observed in the 

systems with organics, thereby indicating that the presence of boron had a more dominant 

effect. 

9.4. Particle Size Measurements of Surrogates at 60°C 

To examine the behavior of solutions similar to Tests 1 and 5 at 60oC, surrogates were 

used because of the inability to obtain ICET samples at the test temperature. The 

surrogates were made by dissolving aluminum coupons into a sodium hydroxide solution 

with 2800 mg of boron. The pH of this solution was initially 9.6, and the temperature of 

the solution was maintained at 60o

C. Distributions were obtained using a quasi-elastic 

light-scattering device with a range of particle detection of 2-nm–1-μm particle size. 

Figure 56 presents the particle distribution. From this figure it would appear that 

nanoparticles of ~30 and 500 nm were in solution. Furthermore, as was previously noted 

in the stored Test 1, 60o

C precipitate, precipitate was noted after several weeks at 60oC. 

To understand the bimodal particle size distribution, additional investigations on the 

properties of the surface charge are needed. 
77

(a) 27Al

(b) 11B

Figure 55. NMR solid-state spectrum of the precipitates. 
78

Figure 56. Particle size distributions for aluminum/boron metal dissolution surrogate 

solution at 60°C after 8 hours. 

Furthermore, additional measurements at times comparable to the duration of the ICET 

tests would be helpful to elucidate the evolution of the particle distribution dynamics. 

From examination of the surrogate system, it is clear that colloids exist at test 

temperatures. 

9.5. Inhibitor Passivation 

To examine the possibility of inhibiting aluminum corrosion in alkaline solution, several 

experiments were performed. In these experiments, aluminum coupons and aluminum 

alloy 3003 [having dimensions of 24 (l) × 13 (w) × 1.8 mm (h) and a weight of ~1.59 g] 

were ultrasonically cleaned with 95% ethanol, followed by drying in air before the test 

was conducted. The test coupons were then inserted into three 1.0-L solutions (the water 

used for these tests came from a laboratory water purification system): boric acid, boric 

acid and silicon, and boric acid and calcium. The concentration of boric acid was 

16,000 mg/L (i.e., 2800 mg/L of boron) in all of the three solutions. Na2SiO3·9H2O and 

CaCl2 were used to introduce silicon or calcium to the solution. For the solution 

containing silicon or calcium, its concentration was 88.7 mg/L or 50 mg/L, respectively. 

Each solution was put into a 1.0-L Nalgene® bottle. All of the solutions were preheated 

to 60°C, and the pH was adjusted to 9.5 using sodium hydroxide before putting 

aluminum coupons into the solutions. The temperature of the solution was maintained at 

60°C by a constant-temperature oven (a Yamato DKN400) throughout the 30-day test. 

The corrosion of the aluminum coupons was monitored by an ICP-AES and pH 

measurements. The aluminum coupons were analyzed by EDS. 
79

Figure 57. Aluminum concentration (mg/L) versus time. 

The measured aluminum concentrations in the three solutions are shown in Figure 57. It 

is evident that the concentration in the boron+ silicon solution was almost zero and did 

not change with time, indicating that the inhibition film was instantaneously formed on 

the coupon surface and that the inhibition efficiency was almost 100%. For the other two 

solutions (boron and boron + calcium), the concentrations increased with time until they 

reached an asymptotic value. As may be seen from Figure 57, it is evident that the 

introduction of calcium reduced the aluminum corrosion rate and the final aluminum 

concentration also could be reduced, but the inhibition efficiency was much lower than 

the efficiency obtained by adding silicate. From the SEM images of the sample, it was 

hard to observe the film, indicating that the film was very thin. Elemental compositions 

of the surface and the substrate have been examined using EDS; the analyses indicate the 

presence of small quantities of silica at the surface, whereas no silicate was detected in 

the bulk of the aluminum coupons; this result verifies the experimental results obtained in 

Ref. 87. According to the reference, the film was an amorphous aluminum-silicate phase. 

To provide direct evidence of the aluminum-silicate film formed on the aluminum alloy 

coupon surface, x-ray photoelectron spectroscopy (XPS) measurements were performed 

on a precleaned, virgin aluminum alloy coupon and on the silicate-passivated aluminum 

alloy coupon. The silicate passivated aluminum coupon was removed from the Na2SiO3

solution and gently rinsed with deionized water several times before storing it in a 

nitrogen-filled desiccator for drying. For the XPS measurements, a Kratos Axis Ultra 

spectrometer was used with an AlKa 1486-eV x-ray source with 300 W of power to 

generate x-rays. An electron flood gun for charge neutralization and a hemispherical 

analyzer with eight multichannel photomultiplier detectors were also used. 

As elucidated in Figure 58, the XPS examination was carried out with Al 2p at 90 

takeoff-angles (TOAs), which gives the chemical composition of the coupon subsurface 

up to a 10-nm depth. According to the result, the subsurface of the virgin aluminum 
80

coupon up to 10-nm depth was composed of Al2O3 (64%), Al(OH)3 (27%), and Al (9%). 

However, the subsurface of the silicon-passivated coupon was composed primarily of 

Al2OSiO4 (98%), with a small amount of Al2O3 (2%). This result indicates that silicon 

complexed significantly with aluminum at the coupon surface. The thickness of the 

complexation layer of silicon and aluminum is >10 nm. In addition, in contrast to the 

virgin aluminum coupon, no metal aluminum was detected up to 10 nm from the silicatepassivated coupon surface. This observation may be explained by the hypothesis that the 

complexation or passivation layer grew on top of the original coupon surface and 

consequently buried the aluminum metal in a deeper subsurface layer. 

To examine the passivation layer stability, the passivated aluminum coupons were treated 

in the following ways: two of the coupons were heated in an oven at 250°C for 2 hours; 

these coupons are referred to as the “preheated” coupons in Figure 59. When the coupons 

were cooled down, one unheated coupon and one preheated coupon were transferred to a 

12-ml HCl solution that was at a pH of 2.2, respectively. Similarly, one unheated and one 

preheated coupons were put into a 12-ml NaOH solution at a pH of 9.5, respectively. The 

concentration of aluminum in the solution with respect to the soaking time is shown in 

Figure 59. Based on the figure, the silicate passivation was stable below a pH of 9.5 

because little aluminum concentration was detected. However, a significant aluminum 

concentration (>100 mg/L) was found in the solution having a pH of 2.2. Little difference 

was noted between the preheated and unheated aluminum coupons. This result indicates 

that the passivation layer of silicate was stable below a pH of 9.5 and likely was 

dissolved at a pH of 2.2. 

Figure 58. XPS spectra of Al 2p at 90 TOA on aluminum alloy coupons: (left) virgin 

aluminum, (right) aluminum soaked in Na2SiO3 solution. 

 

Al(OH)3

Al2O3 Al 

AlOSiO4 

(a) (b)

Total Al
81

Figure 59. The stability of a passivation layer of silicate with respect to high and low 

pH. 
82

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83

10. CONCLUSIONS 

A comprehensive examination of both the test solution and precipitate from ICET Tests 

#1 and #5 has been performed. In addition, a comprehensive review of available literature 

and benchmark experiments of surrogates has been performed to assist in explaining the 

behavior of the aluminum in chemical solutions. The objective of this analysis was to 

elucidate the behavior of precipitate that formed when the Tests 1 and 5 solutions were 

allowed to cool. This examination included supplemental analytical measurements using 

XRD, 27Al, 11B NMR (both liquid and solid state), and quasi-elastic light-scattering 

measurements. The results of this investigation allow for the extrapolation of the 

behavior of ICET to prediction the behavior of aluminum under different pH and 

temperature conditions that might exist throughout the PWR following a LOCA. Finally, 

the characterization of the particle sizes and characterization of the corrosion properties 

of aluminum under LOCA conditions have been elucidated. These findings should allow 

for the development of a head-loss correlation using the existing cake filtration theory, 

which could be used in conjunction with a corrosion model to predict system 

performance following a LOCA. 

Based on both ICP analysis and examination of the Tests 1 and 5 aluminum coupons, it 

was concluded that the aluminum weight loss from the metal coupons in these tests was 

consistent with the measured aluminum coupons. Furthermore, by examining the plateau 

of viscosity (at 23o

C), hydrogen production, and aluminum concentration, it was 

hypothesized that the plateau in the aluminum concentration was a result of passivation 

of the aluminum metal surface. The measured aluminum concentrations of aluminum in 

Tests 1 and 5 solutions did not approach predicted solubility limits, and calculations 

confirmed that the cooling of Tests 1 and 5 solutions to ambient temperature would 

indeed produce precipitation. 

Examinations of the Tests 1 and 5 precipitates revealed that the precipitate after washing 

(performed to ensure that sodium borates were not allowed to develop during the drying 

process) was largely amorphous aluminum hydroxide, with a substantial quantity of 

boron adsorbed onto the surface. A survey of the literature suggests that the presence of 

the amorphous form of aluminum is indeed expected because of the solution’s high 

concentration of anions, which have been shown to retard crystallization at temperatures 

<60o

C. Thus, if amorphous-phase aluminum does indeed form during cooling of the Test 

1 solution, as both XRD and TEM analyses would indicate, then it is not surprising that 

significant boron would be found as a result of the adsorption of boron from the solution. 

As indicated from the ICP results, up to 35% of the boron from the initial solution may 

have been adsorbed onto the amorphous aluminum hydroxide precipitate. This degree of 

adsorption is certainly feasible. In fact, if we use the data from Su and Surez47 at a pH of 

~9.5, a boron adsorption of ~35% is predicted. A review of the literature and NMR 

measurements, which revealed complexation between aluminum and boron, was believed 

to be responsible for impeding the crystallization of aluminum. Furthermore, the 

measured weight percentage of boron on the aluminum precipitate was consistent with 

previous predictions based on adsorption models. 
84

The solid hydroxide equilibrium phases play an important role in aluminum solubility. 

Generally, the crystalline form has a smaller aluminum solubility than does the 

amorphous form. The phase transition is a complex function of operating conditions and 

the nature of the materials used. In an alkaline solution, it is likely that the initial 

precipitates are always amorphous hydroxide, which can be transformed into other forms 

of aluminum hydroxide by the action of water. At a specific operating condition, more 

crystalline hydroxide is produced with aging. The crystallization process may be 

retarded, or the crystal habit of the precipitates may be changed by adding some organic 

or inorganic cation/anion, such as citrate or boron. The retardation behavior is due to the 

activity change at the solid/solution interface because of adsorption of the adding 

anion/cation, which leads to changes in the surface charge and interfacial tension and can 

influence the kinetics of the incorporation of the growth unit from the solution into the 

growing solid phase. 

Aluminate ions ( − Al(OH) ) are the only stable form of al 4 uminum in an alkaline solution. 

The solubility is a function of solid hydroxide phases and increases with pH. It is noted 

that some organics and inorganics can increase the aluminum solubility. The solubility 

also can be affected by the particle size presented in the solution. Nanoparticles are found 

in the solution, and the particle size distribution is affected by the hydroxide phases and 

the adding organic elements. To determine if colloids exist in Tests 1 and 5 solutions, 

DLS was used. Results of these measurements revealed that colloids did indeed exist at 

ambient temperature with a mean radius of 65 nm. 

Aluminum corrosion in an alkaline solution is an electrochemical process. In a static 

solution, the corrosion rate decreases with exposure time. In a dynamic solution, the 

corrosion rates increase with increasing flow velocity in mass transfer-controlled regimes 

and does not depend on flow velocity in actively controlled regimes. The corrosion rate is 

influenced by adding organic and inorganic inhibitors into the alkaline solution or some 

other metal elements into the aluminum metal. The sodium silicates have been found to 

be effective inhibitors, with an inhibition efficiency of almost 100%. 
85

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