APPROVED:

Paul S. Braterman, Major Professor Mohammad Omary, Committee Member Ruthanne D. Thomas, Chair of the Department of

Chemistry Sandra L. Terrell, Dean of the Robert B. Toulouse

School of Graduate Studies

LAYERED DOUBLE HYDROXIDES: SYNTHESIS, CHARACTERIZATION, AND

INTERACTION OF Mg-Al SYSTEMS WITH INTERCALATED

TETRACYANONICKELATE(II)

Fang Wei Brister, B.Sc.

Thesis Prepared for the Degree of

MASTER OF SCIENCE

UNIVERSITY OF NORTH TEXAS

August 2004

Brister, Fang Wei, Layered Double Hydroxides: Synthesis, Characterization, and

Interaction of Mg-Al Systems with Intercalated Tetracyanonickelate(II). Master of Science

(Inorganic Chemistry), August 2004, 37 pp., 7 tables, 8 illustrations, 115 references.

The square-planar tetracyanonickelate(II) anion was intercalated into 2:1 and 3:1 Mg-Al

layered double hydroxide systems (LDHs). In the 2:1 material, the anion holds itself at an angle

of about 30° to the layers, whereas in the 3:1 material it lies more or less parallel to the layers.

This is confirmed by orientation effects in the infrared spectra of the intercalated materials and

by X-ray diffraction (XRD) data. The measured basal spacings for the intercalated LDH hosts

are approximately 11 Å for the 2:1 and approximately 8 Å for the 3:1. The IR of the 2:1 material

shows a slight splitting in the ν(CN) peak, which is suppressed in that compound’s oriented IR

spectrum, indicating that at least some of the intercalated anion’s polarization is along the z-axis.

This effect is not seen in the 3:1 material.

A comparison between chloride LDHs and nitrate LDHs was made with respect to

intercalation of tetracyanonickelate(II) anions. Both XRD data and atomic absorption

spectroscopy (AAS) data of the LDH tetracyanonickelates confirms that there are no significant

differences between the products from the two types of starting materials. The presence of a

weak ν(NO) peak in the IR spectra of those samples made from nitrate parents indicates the

presence of small amounts of residual [NO3]- in those systems. Small amounts of Cl- present in

the chloride-derived samples, while perhaps detectable using AAS, would not be detectable in

this manner.

An attempted synthesis of Mg-Al LDH carbonates starting from reduced Mg and Al was

unsuccessful due to pH constraints on hydroxide solubility in the solvent system used (water).

The pH required to precipitate Al(OH)3 in the system was too high to allow precipitation of

Mg(OH)2. Consequently, we found it impossible to have both of the required metal hydroxides

present simultaneously in the system. An additional synthesis using a halogen as an oxidizing

agent also failed to produce material of any characterizable quality.

ii

ACKNOWLEDGEMENTS

I would like to thank Dr. Paul S. Braterman, my major advisor for his guidance. Thanks

also go to Dr. Faith Yarberry for her help and discussions. Mickey Richardson also contributed

to this work significantly.

This research was made possible by grants from NASA, the Robert A. Welch foundation

and the UNT Faculty Research Fund.

I would also like to especially thank my husband, Brian Brister, for his time, patience,

and thoughtful perspectives throughout the course of my research and while proofreading this

work. Finally, I would like to thank my daughter, Jike Wei, my son, William Brister, my mother,

Jinren Huang, and the rest of my family and friends for the strength and support they gave to me.

Through their love, patience, devotion and daily encouragement, this research was possible.

iii

TABLE OF CONTENTS

Page

ACKNOWLEDGEMENTS.......................................................................................................... ii LIST OF TABLES AND ILLUSTRATIONS ............................................................................. iv INTRODUCTION ........................................................................................................................ 1

Synthesis of Layered Double Hydroxides ........................................................................ 2

On Preparing LDH from Reduced Metals: A “Total Synthesis” Approach ..................... 4

Uses of LDHs.................................................................................................................... 5 ANALYTICAL TECHNIQUES................................................................................................... 7

Fourier Transform Infrared Spectroscopy (FTIR) ............................................................ 7

Powder X-Ray Diffraction.............................................................................................. 10

Atomic Absorption Spectroscopy (AAS) ....................................................................... 11

Carbon, Hydrogen, and Nitrogen (CHN) Elemental Analysis ....................................... 12 EXPERIMENTAL...................................................................................................................... 13 RESULTS AND DISCUSSION................................................................................................. 16

Synthesis of 2:1 Mg:Al LDH carbonate from Mg, Al, Na2CO3, and CO2 ..................... 24

Synthesis of Potassium Tetracyanonickelate(II)............................................................. 26 CONCLUSIONS......................................................................................................................... 29

Comments on the “Total Synthesis” Approach .............................................................. 30 REFERENCES ........................................................................................................................... 32

iv

LIST OF TABLES AND ILLUSTRATIONS

Page

Tables

1. Observed infrared frequencies for 2:1 Mg:Al LDHs containing nitrate and tetracyanonickelate (cm-1)............................................................................................... 19

2. Observed infrared frequencies for 2:1 Mg:Al LDH chloride and nickelocyanide (cm-1) ......................................................................................................................................... 19

3. Observed infrared frequencies for 3:1 Mg:Al LDH nitrate and nickelocyanide (cm-1) . 20

4. Observed infrared frequencies for 3:1 Mg:Al LDH chloride and nickelocyanide (cm-1) ......................................................................................................................................... 20

5. Observed powder XRD d-spacings of 2:1 Mg:Al LDH materials. Spacings given in angstroms ........................................................................................................................ 21

6. Observed powder XRD d-spacings of 3:1 Mg:Al LDH materials. Spacings given in angstroms ........................................................................................................................ 22

7. Elemental analysis results of 2:1 and 3:1 Mg:Al LDH –Ni(CN)4 materials (Cl-LDH and NO3-LDH as parent material) ......................................................................................... 24

Figures

1. Schematic of a layered double hydroxide......................................................................... 2

2. FTIR spectra of a) 3:1, b) 2:1 Mg:Al LDH [Ni(CN)4]2- from chloride precursors......... 16

3. FTIR spectra of a) 3:1, b) 2:1 Mg:Al LDH nitrate ......................................................... 17

4. FTIR spectra of a) 3:1, b) 2:1 Mg:Al LDH [Ni(CN)4]2- from nitrate precursors........... 17

5. FTIR spectra a) conventional, b) oriented, of 2:1 Mg:Al LDH [Ni(CN)4]2-. Note scale change ............................................................................................................................. 18

6. Powder XRD patterns for (a) 3:1 Mg-Al LDH-Ni(CN)4 and (b) 2:1 Mg-Al LDH-Ni(CN)4 from exchange with parent LDH-Cl ............................................................................... 22

7. Powder XRD patterns for (c) 3:1 Mg-Al LDH-Ni(CN)4 and (d) 2:1 Mg-Al LDH-Ni(CN)4 from exchange with LDH-NO3....................................................................................... 23

8. Infrared spectrum of material obtained from reaction of Mg and Al with I2.................. 31

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INTRODUCTION Layered double hydroxides (LDHs)1-3 are a family of natural and synthetic compounds

having the general formula of [M(II) 1-x M(III)x (OH)2 ](Yn-)x/n · yH2O, where M(II) and M(III)

represent divalent and trivalent metals, respectively, and Yn- is the anion between the layers.

(Figure 1 shows a general structure of a LDH.) In present-day LDH materials, divalent metals

can be Mg2+, Ca2+, Zn2+, Co2+, Cu2+, etc.; trivalent metals can be Al3+, Cr3+, Co3+, Fe3+, Mn3+,

etc. Many anions can be used, including Cl-, NO3-, CO3

2-, and many organic anions. LDHs are

also known as hydrotalcite-like compounds (due to their structural similarities to that mineral) or

anionic clays. They consist structurally of brucite-like [Mg(OH)2] layers, with a net positive

charge due to partial substitution of divalent metals (M2+) by trivalent metals (M3+); typically, the

substitution leads to values of x in the range of 0.2 to 0.4. This positive charge is effectively

dispersed uniformly across each layer, driven by repulsion of the positive charge centers. Anions

are attracted into the interlayer space along with water molecules, balancing the overall positive

charge with a negative one of equal magnitude.

LDH materials appear in nature and can be readily prepared in the laboratory. In nature they

are formed from the weathering of basalts4-5 or precipitation in saline water sources.5 All natural

LDH minerals have a structure similar to hydrotalcite, which has the formula

[Mg6Al2(OH)16]·CO3·4H2O. Unlike clays, however, layered double hydroxides are not

discovered in large, commercially exploitable deposits.6

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Figure 1. Schematic of a layered double hydroxide.

Synthesis of Layered Double Hydroxides There are many methods to synthesize LDH materials in the laboratory. Direct synthesis

is perhaps the most widely used technique of preparation.3, 7-17 In this method, the metal

hydroxides are co-precipitated by adding a base to a stirred aqueous solution containing M(II)

and M(III) salts and the desired anion. We have found that a 50 wt% solution sodium hydroxide

is especially useful, since any carbonate contamination precipitates out of solution as sodium

carbonate. LDH systems generally prefer carbonate to most other anions; for this reason,

carbonate and carbon dioxide must be rigorously excluded from the system whenever excess

water is present.6 The metal hydroxide layers nucleate and grow under stirring and are often

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refluxed to improve crystallinity. Titration at constant or varied pH and buffered precipitation

are included in this method. Other methods used to prepare LDH materials include co-

precipitation,18 anion exchange,19 glycerol-effected exchange,20-22 preparation from oxides and

hydroxides,23-25 preparation by sol-gel techniques,26 reaction of metal oxides/hydroxides with

salts,25, 27-28 so-called homogeneous precipitation,29-30 preparation via aluminate,31-36 preparation

by oxidation,40-50 and preparation from metals.51-55 Post-preparative treatments are usually

necessary to improve the crystallinity of the LDH materials prepared by conventional methods.

Improved crystallinity is desirable in LDHs because it allows for greater stabilization of

intercalated anions. Ageing the precipitation at or above room temperature is often used as a

post-treatment method. In addition, hydrothermal and microwave treatment techniques can be

employed as post treatments.

In preparation of LDHs, it is necessary to rigorously exclude carbon dioxide from the

system because of the well-known conversion of carbon dioxide to carbonate under mild to

strongly basic conditions, and ready uptake of the carbonate by the material. It is important to

note that carbonate contamination is more prevalent at higher pH. Mg-Al LDH forms around a

pH of 8-10, which allows for rapid conversion of CO2 to CO32- in solution. During the

intercalation process, however, the pH is generally measured as the pH of the anion in aqueous

solution. For aqueous [Ni(CN)4]2- we have a measured pH of 11, sufficient to merit the same

rigorous measures used in the synthesis procedures to exclude CO2. In cases where the layer

metal cations are easily oxidized under basic conditions, such as is the case with FeII-FeIII LDH

systems (the famous “Green Rusts”), the reaction should be performed in an oxygen-free as well

as carbon dioxide-free environment.50,56 However, since the magnesium-aluminum LDH system

is more robust than its iron-containing counterpart, less stringent precautions can be taken. It is

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possible to synthesize a 2:1 Mg-Al LDH with very low carbonate contamination using nothing

more than a blanket of high-purity (~99%) Nitrogen gas. This method is illustrated in greater

detail in the Experimental Section.

On Preparing LDH from Reduced Metals: A “Total Synthesis” Approach The preparation of LDH from metal starting materials could be very useful in certain

circumstances. Isotopic substitution is perhaps the most obvious advantage to this approach.

When carried out in D2O, layered double deuteroxides (LDDs) are expected to form. These

compounds are chemically identical to LDHs, but are more suited for special purposes such as

neutron diffraction studies, since hydrogen is more prone to randomly scatter incident neutrons

than deuterium. Since many of the metal salts that are normally used in syntheses are hydrated

(or worse, hygroscopic), their use should be avoided whenever a hydrogen-free LDD is desired.

Isotopes of Fe, Co, Ni, O, or other elements involved may enable characterization of other

aspects of the system such as the extent of exchange of LDH components with the surrounding

solution. A high degree of control is available using this “total synthesis” approach; however,

the method has also proven to be inherently difficult to execute. The primary difficulty lies in

the presence of excess electrons in the system. Since the metals have an initial oxidation state of

zero, they will only react with substances that will remove electrons from them. OH- (more

appropriately for our experiments, OD-) already holds one extra electron, and will not remove

another from a reduced metal. We must therefore proceed through the formation of intermediate

compounds such as metal salts of some form or another. Our experiments have shown, as

discussed later, that simply exposing the metals to aqueous solutions of weak acids is insufficient

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to cause appreciable ionization. Although some products were formed by this method, they were

highly admixed with other compounds and therefore unusable. After obtaining isotopically pure

forms of the precursors, the rest of the synthesis should proceed as described above, with the

added constraint of a more controlled environment and exclusion of water.

Uses of LDHs

LDHs are drawing more and more attention recently.2-3, 17, 22, 57-68 These materials are of

interest as catalysts and catalyst precursors and supports,69-71 as ceramics precursors,3 as hosts for

photoactivation and photocatalysis,72-75 as anion exchangers,76-77 as traps for anionic pollutants

including some kinds of nuclear waste,78-83 as antacids and delivery systems for

pharmaceuticals,84-98 as fire retardants99 and additives for polymers.100 The last of these is the

main commercial and industrial application. We consider that LDHs intercalated by cyanide-

containing anions, such as cobalticyanide, ferrocyanide, ferricyanide, could exhibit great promise

in the areas of catalysis, modified electrodes, and molecular adsorption. Additionally, much work

has been done in preparing noble metal containing catalysts by the calcination of LDHs carrying

the metal in question. For example, Mg-Al LDHs have been prepared with intercalated

hexacyanoruthenate(II) anions.101 In other cases, RuII was used as a layer metal to catalyze the

oxidation of alcohols to carbonyl compounds.102-104 In this work, our intercalation of

tetracyanonickelate(II) anions would produce Ni-bearing Layered Double Oxides (LDOs) upon

calcination. It is possible that intercalation of other Ni-containing anions could produce different

LDOs with different catalytic behaviors.

- 6 -

Layered double hydroxides have also been invoked in the study of the origins of life. 2:1

Mg-Al LDHs have been found to catalyze the self-addition of [CN]- in aqueous solution. The

principal product is diaminomaleonitirle, the cyanide tetramer. This was found to form in

preference to the hydrolysis product, formamide, even at fairly dilute (0.05M) concentrations.

Boclair et. al. published a detailed rate study of this reaction in 2001.106 Other, higher-order

products include glycine, triazinetrione, and possibly caffeine and adenine.6 Most notably, the

reaction produces a purple-pink oligomeric material, more or less firmly attached to the LDH.

Recent studies in our laboratory have focused on LDHs intercalated with [Fe(CN)5 (NO)]2-,

[Co(CN)6]3-, [Ru(CN)6]4-, [Pt(CN)4]2-, [Pt(CN)6]2-, and [Ni(CN)4]2-.101 In the present study, we

extend the work on LDH intercalated with [Ni(CN)4]2-, by characterizing and comparing

different ratios of M(II) to M(III) and by using different anion precursors to intercalate the anion.

The Infrared spectrum, X-ray diffraction, and elemental analysis of these materials are presented

here. Additional investigations on the “total synthesis” of LDD systems are also presented.

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ANALYTICAL TECHNIQUES

Fourier Transform Infrared Spectroscopy (FTIR)

We have made extensive use of Fourier transform infrared spectroscopy (FTIR)

throughout this work. FTIR has proven to be an extremely useful tool in the daily course of

research on all the projects mentioned here. It is used as a tool for the evaluation of a molecule’s

structure by studying the vibration change occurring from the excitation of the bonds. The

absorption peak shapes also indicate the environment that the vibration took place in; distortion

from an ideal symmetry causes splitting of the relevant peak in the infrared (IR). The hallmark of

IR spectroscopy is that certain types of bonds between atoms invariably appear in certain regions

of the infrared spectrum; we use this effect to determine the presence or absence of particular

intercalants in the interlayer, and to check that the layered double hydroxide (LDH) layer

material matches the expected spectrum. Since functional groups absorb at characteristic

frequencies of infrared radiation, many compounds can be identified using only their spectra. A

quick scan can be used to verify the purity and veracity of a sample before further experiments

are carried out; higher-quality scans can be performed that give detailed information about the

interlayer environment, the presence or absence of key functional groups in the system, the

relative orderliness of the hydroxide layers, and the extent of ordering of water in the sample.

We can also use IR to determine how much contamination (e.g. how much carbonate) is present

in the system; if the LDH parent was a nitrate, we can also semi-quantitatively determine the

extent of intercalation of our anion of interest.

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We can obtain detailed information about the extent of distortion that the intercalated

anions undergo by comparing the spectra of the original compounds with those of the

intercalated LDHs. This has provided valuable information on the ordering of these anions in the

interlayer. The relative intensity of the ν(CN) peak allows distinguishing between absorption

(intercalation) and adsorption of the metal-cyanide complex under study. Additional

comparisons between spectra of analogous compounds (for example, [FeIII(CN)6]3- vs.

[RuIII(CN)6]3- ) aid in characterizing the layers themselves, while comparing compounds in a

coordination series (such as [PtII(CN)4]2- vs. [PtII(CN)6]4-, where the former intercalates but the

latter does not) further characterizes the Hydrogen-bonding environment of the interlayer and

anions’ ability to intercalate in an ordered manner.

FTIR also has other uses. Many aspects of our system can be qualitatively determined by

the presence or absence of characteristic peaks in the compound’s spectrum. An example of this

is the degree of crystallinity in 2:1 Mg-Al LDH, where highly crystalline samples exhibit a

sharp, intense peak at 447 cm-1. In less ordered samples this peak will be less intense, to the

point where it becomes masked by other signals in the ν(M—OH) region. Certain systems (the

2:1 LDH ferrocyanides in particular) exhibit a noticeable splitting in the region 3300-3600 cm-1,

on the high-frequency side of the large, broad ν(OH) peak in that area. This peak is believed to

arise from water which has become trapped in between ferrocyanide anions in the interlayer, or

from pendant layer hydrogens similarly isolated. In some cases this splitting becomes quite

distinct and the usual ν(OH) is considerably more narrow than normal, indicating a high degree

of ordering amongst the intercalated anions. Yet another method that has frequently proven

useful is the use of oriented infrared spectroscopy. In this method, individual particles of LDH

are allowed to deposit themselves slowly onto the surface of a barium fluoride (BaF2) disc from

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aqueous suspension. This allows the particles time to settle into a roughly flat, even layer across

the disc’s surface. This causes one axis of polarization to be situated normal to the detection

plane, and it therefore becomes suppressed in the resulting spectrum. The spectroscopic effect

that this arrangement produces in LDH compounds is quite noticeable, particularly in the ν(CN)

region of interest here. Results from using this method have allowed us to better determine the

exact environments that our anions are subjected to in the layer. Boclair et al. used this method

extensively to show that distortion from Oh symmetry occurs along the z-axis whenever 6-

coordinate metal-cyanide complexes (such as [Fe(CN)6]3- and [Ru(CN)6]4-) undergo

intercalation.56 Careful analysis of the non-oriented spectra shows a splitting of the ν(CN) peak

into two distinct signals, one usually appearing as a shoulder on the other. The splitting of the

signal is attributed to the splitting of the T1u stretching mode into A2u and Eg modes in D3d

symmetry, indicating a distortion of the anion force-field in the LDH interlayer. Since the A2u

component is polarized perpendicularly to the plane of the layer, it becomes suppressed in the

oriented spectrum as the particles align with the layers normal to the incident beam thus causing

this mode to align with the beam’s direction. This suppression serves as positive evidence of the

force-field distortion hypothesis.

The FTIR instrument employed in our experiments was a Perkin Elmer ™ 1760 FTIR

spectrometer (Perkin Elmer Corp., Norwalk, CT). Data was processed using the Perkin Elmer ™

Spectrum 1700 software (Perkin Elmer Corp., Norwalk, CT) running on a Microsoft® Windows

95® operating system (Microsoft Corp., United States). All samples were obtained using a

pressed potassium bromide (KBr) pellet containing approximately 1% sample. The KBr pellet

were made using FTIR grade KBr as purchased from Aldrich. The KBr and the sample were

finely ground together to obtain a mixture in which the sample thoroughly dispersed throughout.

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This mixture was then placed in a die, the die was evacuated using a vacuum pump, and pressed

to form a pellet.

Powder X-Ray Diffraction (PXRD) Powder X-ray diffraction (PXRD) is a very useful tool to analyze the crystalline structure,

lattice parameters, crystal morphology, and crystal size107 of LDHs. This application is based on

the fact that an X- ray diffraction is unique for each crystalline substance.108 Thus if an XRD

pattern of an unknown sample is found to match the pattern of an authentic one, chemical or

structure identity can be presume.. The XRD pattern of a typical LDH chloride shows that the

compounds correspond to the space group 3R1. The peaks from the lower angle to the high angle

are the diffraction by planes (003), (006), (009), etc. thus the cell parameter c is usually

estimated as 3d003 or d003 + 2d006 + 3d009. And a can be easily calculated from the 110 reflection

(a = 2d110).

XRD is based on Bragg’s law:

nλ = 2d sin θ

where λ is the wavelength of the incident X-ray beam, d is the spacing between each lattice, θ is

the angle between the incident X-ray beam and the reflecting crystal plan, and n is an integer

representing the order of the reflection (in practice, taken to be 1). Bragg’s law indicates that the

diffraction occurs when two conditions are satisfied:

1) The angle of the incidence = the angle of the scattering.

2) The path length difference is equal to an integer number of the wavelength.

The research data discussed in this thesis were collected using a Siemens® D-500 X-ray

Diffractometer (Munich, Germany). The radiation used was Cu Kα at a wavelength of 1.541838

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Ǻ. The powder was finely ground with approximately 5 wt% of CaF2 as an internal standard and

placed in a sample holder.

Atomic Absorption Spectroscopy (AAS) AAS is a method by which relative metal ratios can be determined from analysis of a

sample. The raw data from this method is in the form of a concentration (ppm) with respect to a

standard solution. These standard solutions are very carefully calibrated with standard solutions

made by the National Institute for Standards in Technology (NIST). For this reason, our AAS

measurements can be used even when combined with measurements from another calibrated

AAS instrument if necessary.

The principles behind atomic spectroscopy are well-known. A solution of the sample

(usually as free ions) is vaporized in a flame, leaving the element of interest present as free

atoms. The vaporized solvent and analyte atoms pass through a beam from an element-specific

lamp, where electrons in the outer orbitals of the atoms can become excited to higher-energy

levels by absorbing characteristic “excitation wavelengths”. Either this absorption or the

corresponding emissions when the electrons return to their original orbitals can be measured by a

detector. Using element-specific lamps is essential as the lamps establish a base intensity and

provide light at precisely the needed wavelengths for excitation. More energy is absorbed at

higher analyte concentrations. The lamps used are normally hollow cathode lamps that excite

their respective elements by electric discharges; the atoms in the lamps emit the characteristic

wavelengths that the atoms in the sample absorb. Modern AAS is capable of detecting analyte

even at ppb levels, and is highly specific to the element being analyzed. All AAS data in this

work were gathered using a Perkin-Elmer™ AAnalyst 300 (Perkin-Elmer Corp., Norwalk, CT).

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The lamps used were hollow cathode lamps as provided from Perkin-Elmer, and the fuel source

for the flame was either an air/acetylene mix or a mixture of N2O and acetylene, depending on

the element analyzed. All samples analyzed were first digested in a 5 vol% solution of HNO3 to

allow the free metal ions to be analyzed. NIST-calibrated standard solutions were diluted to

follow the Perkin-Elmer recommended methods; the metal concentrations in the LDH samples

were treated in a similar fashion.

Carbon, Hydrogen, and Nitrogen (CHN) Elemental Analysis

Elemental Analyses for carbon, hydrogen, and nitrogen presented in this work were

performed by Atlantic Microlabs. The analyses were performed by combustion of the samples in

an oxygen-rich atmosphere and measurement of the resulting gases. However, the results

obtained by this method were irreproducible and inconsistent with data from other sources and so

were not used.

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EXPERIMENTAL Pure (18.2 MΩ/cm3) water, boiled and purged with UHP nitrogen to remove carbon

dioxide, was used throughout this work. Hydrated metal salts (A.C.S reagent grade),

AlCl3·6H2O, MgCl2·6H2O, and NaCl were used as supplied by Aldrich. Potassium cyanide

(KCN) and nickel (II) cyanide (Ni(CN)2) were used as supplied by Alfa Æsar. Potassium

tetracyanonickelate (II) (potassium nickelocyanide) was prepared by suspending nickel (II)

cyanide in water and adding KCN to a (CN-):Ni ratio of 4:1. Small amounts of dilute HCl were

added to decompose remove dissolved CO2. To avoid exposure to HCN, this procedure was done

in a fume hood with the sash drawn down almost entirely. Pure crystals were collected by

vigorous boiling to create a supersaturated solution, followed by cooling to room temperature to

allow for rapid formation of larger single crystals.

All parent layered double hydroxide (LDH) materials for this work were prepared using a

direct synthesis method. Mg2+ and Al3+ (chlorides or nitrates as appropriate) were dissolved in

water under a slow stream of nitrogen, along with enough of the appropriate sodium salt to bring

the anion concentration to 1 M. For a 2:1 Mg:Al LDH, a salt solution was prepared using the

hydrated salts of Mg and Al. An amount of water was calculated based on the total amount of

LDH desired from the synthesis, sufficient to make the final LDH concentration 0.1 M. The

required amounts of hydrated Mg and Al salts were calculated to make their concentrations 0.3M

and 0.1M, respectively. The appropriate amount of required Na salt was also calculated. The

required amounts of these salts were then added to a 3-necked, round-bottomed flask along with

most of the total water needed. The salt solution was placed under a blanket of ultrapure

nitrogen gas for 15 minutes to establish the atmosphere and allow for the solution to exchange

any dissolved gases. The amount of 50% NaOH needed to make a 6:1 OH-:Al ratio was diluted

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with the remaining water and added either directly or by titration to co-precipitate Mg(OH)2 and

Al(OH)3. In a typical 2:1 Mg-Al LDH synthesis, 85 mL of a solution containing 25 mmol Mg2+,

8 mmol Al3+, and 8 mmol Na+ (as NaCl, to make [Cl-] ≈1 M) was treated with approximately 4

mL of fresh 50 wt% NaOH. Freshness of the base is very important whenever a quantitative,

detailed analysis of synthetic procedures is required, since NaOH reacts with atmospheric CO2

and effectively lowers the OH- concentration in the reagent bottle. However, for syntheses that

follow well-established protocols, the buffering effect of excess Mg2+ in solution allows for some

ageing to take place without affecting the products. The resulting LDH suspension was refluxed

overnight under a slow stream of UHP nitrogen gas. After reflux, the solid was collected and

washed thoroughly via centrifuge to remove excess salts.

For a 3:1 LDH, the synthesis procedure is similar, except that amount of hydrated Mg salt

was increased to make a 0.4M final concentration, creating a final anion concentration of 1.1 M

and eliminating the need for the Na salt. Anion exchange was achieved by exposing a sample of

the parent LDH material (approximately 2 grams LDH) to a solution of containing a large excess

of K2Ni(CN)4 (approximately 0.2 M). The reaction was performed under a nitrogen blanket with

stirring for one hour. The intercalated LDH was then collected and washed via centrifuge.

Following this, the samples were placed in a vacuum dessicator until they appeared to be dry.

Preparation for the various analysis methods is described below.

Powder X-ray diffraction (XRD) data were obtained using a Siemens D500 Diffractometer

with Cu Kα radiation. All powdered XRD samples were prepared by grinding the samples with

approximately 5 wt% CaF2 as an internal measurement standard. Conventional infrared spectra

were gathered using a Perkin Elmer 1760-X FTIR spectrometer. The data was processed using

Spectrum for Windows™ data analysis software on a Microsoft Windows 95 operating system.

- 15 -

The spectra were obtained using KBr discs, containing approximately 1% of sample, mounted

perpendicular to the sample beam. Carbon, nitrogen and hydrogen contents were analyzed by

Atlantic Microlabs, Incorporated. Metal contents were determined using a Perkin Elmer

AAnalyst 300 Atomic Absorption Spectrometer connected to Perkin Elmer AA Winlab software.

The AAS standard metal solutions conforming to NIST standards requirements were obtained

from Perkin Elmer and diluted to concentration equivalent to those recommended by Perkin

Elmer to obtain optimum results. The AAS samples were digested in a 5% HNO3 acid, and the

concentration of the metals in the LDHs to be analyzed were adjusted by dilution to fall within

the linear regression obtain between 0 and the concentration of the calibration standard.

- 16 -

RESULTS AND DISCUSSION

Both 2:1 and 3:1 Mg:Al LDH nickelocyanides were successfully prepared by exposure of

the relevant layered double hydroxide (LDH) chlorides to a 0.2 M solution of K2Ni(CN)4. This

anion exchange is shown clearly in the compound’s infrared spectra, where a large ν(CN) peak

appears in the region 2000-2200 cm-1. The infrared spectra of both the 2:1 and 3:1 Mg-Al LDH

nickelocyanides derived from parent chlorides are shown in Figure 2. Exchange was also

achieved using the LDH nitrate precursors for both the 2:1 and the 3:1 Mg:Al LDH systems.

However, the exchange failed to remove 100% of the nitrate anions from the parent compounds.

This is clearly evidenced by the appearance of the characteristic ν(NO3) peak at 1384 cm-1 in the

compound’s infrared spectrum. Figure 3 shows the infrared spectra of the 2:1 and 3:1 LDH

nitrate parents; the nickelocyanide materials created from them are shown in Figure 4.

Figure 2. FTIR spectra of a) 3:1, b) 2:1 Mg:Al LDH [Ni(CN)4]2- from chloride precursors.

4000 3500 3000 2500 2000 1500 1000 500

Wavenumber (cm-1)

AB

S

- 17 -

Figure 3. FTIR spectra of a) 3:1, b) 2:1 Mg:Al LDH nitrate.

Figure 4. FTIR spectra of a) 3:1, b) 2:1 Mg:Al LDH [Ni(CN)4]2- from nitrate precursors.

4000 3500 3000 2500 2000 1500 1000 500Wavenumber (cm-1)

AB

S

4000 3500 3000 2500 2000 1500 1000 500Wavenumber (cm-1)

AB

S

- 18 -

Looking at the infrared spectrum of the 2:1 Mg:Al LDH nickelocyanide, we can see that

the ν(CN) vibration’s maximum is at 2124 cm-1, very close to the 2123 cm-1 maximum found109

for [Ni(CN)4]2- in aqueous solution. A.G. Sharpe reported the value of the ν(CN) peak at 2133

cm-1 in the solid potassium salt; this agrees with our observations of the purified material we

used in our experiments.110 A close-up of the ν(CN) region shows that this peak is actually split

into two peaks of near-equivalent intensity. We believe that this splitting is due to a lowering of

the anion’s symmetry from D4h to D2h upon intercalation. The oriented IR of this compound

shows suppression of the low-frequency splitting, indicating that the ν(CN) vibration involved

has at least some part of its polarization along the z-axis. This can only be achieved if the anion

is present at an appreciable angle to the layers. This phenomenon is shown in Figure 5.

Figure 5. FTIR spectra a) conventional, b) oriented, of 2:1 Mg:Al LDH [Ni(CN)4]2-. Note scale change.

Also of interest in the infrared spectrum of the 2:1 Mg:Al LDH nickelocyanide made

from the NO3- precursor (shown in Figure 4) is the existence of a well-resolved peak in the

ν(OH) region around 3640 cm-1. This peak is observed frequently in LDH ferrocyanides, where

- 19 -

it has been attributed to non-hydrogen-bonded OH groups present in the interlayer.56 These are a

natural geometric consequence of the interaction of these anions with the layers; this can be

assumed to be the case in the LDH nickelocyanide systems as well. Tables 1 and 2 list all of

these peaks for the 2:1 LDH system made from the nitrate and chloride precursors, respectively.

Table 1. Observed infrared frequencies for 2:1 Mg:Al LDHs containing nitrate and tetracyanonickelate (cm-1).

Anion ν (OH) δ (OH) ν, δ (anion) lattice vibrations

(NO3)- 3543, 3431 1634 1764, 1384, 826 672, 446

[Ni(CN)4]2- 3636, 3435 1645 2123 660, 447

Table 2. Observed infrared frequencies for 2:1 Mg:Al LDH chloride and nickelocyanide (cm-1).

Anion ν (OH) δ (OH) ν, δ (anion) lattice vibrations

Cl- 3542, 3467 1625 --- 670, 447

[Ni(CN)4]2- 3636, 3435 1640 2122 660, 448

The infrared spectrum of the 3:1 Mg:Al LDH nickelocyanide appears very similar to its

2:1 counterpart, with the exception that the splitting of the ν(CN) peak is no longer present. We

infer that the anion is aligned more parallel to the layers in that system. Another difference

between the two materials is the occasional appearance of a ν(CN) peak around 2170 cm-1 in the

3:1 but not the 2:1 materials. This peak is of varying intensity and frequently not present at all.

Similar behavior is seen in LDH ferrocyanide systems, where it is attributed to the formation of

cubic ferrocyanide species by reaction with cations from the hydroxide layers.56 We have further

found that exposure of potassium tetracyanonickelate to moderate concentrations of aluminum

chloride in aqueous solution creates a small amount of an insoluble, pale blue product. This

- 20 -

product has a complex infrared spectrum that includes a ν(CN) at 2170 cm-1 as seen in our 3:1

materials. This implies that the [Ni(CN)4]2- anion can, at times, remove Al3+ from the layer and

form this material in situ, although a plausible mechanism by which it could do this has yet to be

found. While interesting and perhaps worthy of future research, this information is not displayed

in Tables 3 or 4. Instead, Tables 3 and 4 show only the values that are consistently observed in

the infrared spectra of our 3:1 LDH systems.

Table 3. Observed infrared frequencies for 3:1 Mg:Al LDH nitrate and nickelocyanide (cm-1).

Anion ν (OH) δ (OH) ν, δ (anion) lattice vibrations

(NO3)- 3543, 3431 1635 1764, 1384, 826 626, 409

[Ni(CN)4]2- 3586, 3468 1625 2122 592, 414

Table 4. Observed infrared frequencies for 3:1 Mg:Al LDH chloride and nickelocyanide (cm-1).

Anion Ν (OH) δ (OH) ν, δ (anion) lattice vibrations

Cl- 3579, 3495 1625 --- 614, 409

[Ni(CN)4]2- 3586, 3468 1625 2121 611, 414

Figure 6 shows the observed XRD patterns for the 2:1 and 3:1 Mg:Al LDH

nickelocyanide systems as prepared from LDH chloride precursors. The relevant d-spacings

observed are shown in Tables 5 and 6. The observed d-spacings in each of the final products

indicate identical behavior between the LDH chlorides and their nitrate analogues. Sharper peaks

in the diffraction patterns of the 2:1 materials indicate a higher degree of crystallinity compared

to the 3:1 compounds. Using the crystal XRD data reported by A.G. Sharpe109 and the van der

- 21 -

Waals radius of Nitrogen111, the square-planar tetracyanonickelate(II) anion can be regarded as a

rounded square slab approximately 7.35 Å per side and about 3 Å thick. However, using the N—

H hydrogen-bond length of 2.84 Å reported by Taylor et. al.112 allows us to ignore the length

contribution from the van der Waals radius of Nitrogen. Thus we arrive at an estimated axis

length of 8.27 Å for the anion in the interlayer. This value, in conjunction with the commonly

accepted thickness of ~4.8 Å per Mg:Al LDH layer, allows an approximate angle of intercalation

for the [Ni(CN)4]2- anion to be calculated using the formula θ= sin-1 [(gallery height) ÷ (axis

length)]. With a measured d-spacing of 10.91 Å in the 2:1 and 8.02 Å in the 3:1 materials, the

angle of intercalation is calculated as approximately 50° in the 2:1 material and must be nearly

zero in the 3:1 material, since the calculated gallery height of 3.22Å is very close to the

calculated thickness of the anion itself. This hypothesis also explains the observed splitting of

the ν(CN) in the 2:1 infrared spectrum, as well as the suppression of the high-frequency

component in the oriented infrared of the 2:1 materials. We attribute this difference to the

change in charge density between the 2:1 and 3:1 materials. The observed d-spacings in the 3:1

Mg:Al LDH nickelocyanide system can be seen in Figure 7 and Tables 5 and 6.

Table 5. Observed Powder XRD d-spacings of 2:1 Mg:Al LDH materials. Spacings given in angstroms.

dnnn Cl- [Ni(CN)4]2-

Cl-LDH as precursor

NO3- [Ni(CN)4]2-

NO3-LDH as precursor

d003 7.75 10.91 8.58 11.05

d006 3.86 5.53 4.39 5.57

d009 2.82 3.69 2.60 3.73

- 22 -

10 20 30 40 50 60 700

100

200

300

400

500

600

700

800

b

a

CPS

2-Theta

Figure 6. Powder XRD patterns for (a) 3:1 Mg-Al LDH-Ni(CN)4 and (b) 2:1 Mg-Al LDH-Ni(CN)4 from exchange with parent LDH-Cl.

Table 6. Observed Powder XRD d-spacings of 3:1 Mg:Al LDH materials. Spacings given in angstroms.

dnnn Cl- [Ni(CN)4]2-

Cl-LDH as precursor NO3

- [Ni(CN)4]2-

NO3-LDH as precursor

d003 7.96 8.02 8.53 8.15

d006 3.97 3.93 4.42 4.11

d009 2.61 2.59 2.58 2.62

- 23 -

10 20 30 40 50 60 70

100

200

300

400

500

600

700

800

900

1000

d

cC

PS

2 -T he ta

Figure 7. Powder XRD patterns for (c) 3:1 Mg-Al LDH-Ni(CN)4 and (d) 2:1 Mg-Al LDH-Ni(CN)4 from exchange with LDH-NO3.

Table 7 lists the results of the elemental analysis of these samples. It is important to note

that the data for this table was compiled from very different analytical techniques, namely AAS

and microanalytical GC. The disagreement between the measured results and the values

predicted by theory is hopefully explained in light of this circumstance. While the metal ratios

we have found are within the bounds of reason, it seems that the ratios of carbon to nickel are

skewed somewhat, since they imply a CN to Ni ratio of only 3:1. We conclude that all the data

from the carbon, hydrogen, and nitrogen elemental analyses are unreliable, possibly due to the

stability of the M—CN bond. We therefore will make no inferences from these data.

- 24 -

Table 7. Elemental analysis results of 2:1 and 3:1 Mg:Al LDH –Ni(CN)4 materials (Cl-LDH and NO3-LDH as parent material).

Elemental Analysis for Nickelocyanide Project

Starting Mat’l. Mg:Al Ni:Al C:Ni N:Ni %Mg %Al %Ni %C %H %N 2:1 LDH-Cl 2.17 0.57 3.27 3.10 16.89 8.63 10.37 6.95 3.59 7.69 2:1 LDH-NO3 2.14 0.58 3.23 2.98 19.92 8.67 10.98 7.26 3.18 7.82 3:1 LDH-Cl 3.17 0.54 3.88 3.64 17.80 6.21 7.20 5.73 3.50 6.25 3:1 LDH-NO3 3.40 0.52 3.88 3.61 20.25 6.60 7.17 5.70 3.23 6.17

Synthesis of 2:1 Mg:Al LDH Carbonate from Mg, Al, Na2CO3, and CO2

Attempts were made to synthesize LDH carbonates directly from the metal carbonates in

order to develop a hydrogen-free method of LDH (more appropriately, Layered Double

Deuteroxide or LDD) synthesis, with a view toward analysis of our layered materials by neutron

diffraction. Hydrogen scatters neutrons in a wildly unpredictable manner; as a result, any

hydrogen in the samples used would vastly reduce the quality of the diffractograms obtained.

The simple substitution of D2O for H2O is ineffective; H competes favorably with D in most

respects because of an isotope effect. Heavier isotopes have naturally lower-energy vibration

states when they form bonds. This causes the lighter isotope to react faster and respond more

strongly to force-field effects than its heavier counterpart. This effect is most apparent between

hydrogen and tritium (hydrogen-3), and only slightly less apparent between hydrogen and

deuterium. H will out-compete D for layer locations near the anions (particularly carbonate

anions), and H2O will be preferred in the interlayer to D2O. It is therefore necessary that H2O be

rigorously excluded from LDD systems.

- 25 -

The problem that then arises is how to exclude H2O but not D2O from the synthesis

environments. We found that conventional water-removal methods were equally as effective on

both compounds. As a result, not only do they fail to remove all of the H2O, they also waste

expensive D2O in the process. In addition, the precursor compounds used in conventional LDH

synthesis incorporate H2O into their crystal structures; this must be removed before a successful

LDD synthesis can be performed.

We therefore sought a method by which the air-stable and moisture-insensitive metal

carbonates could be synthesized directly from the reduced metals as suspensions in D2O. This

would have allowed a total synthesis of LDD from precursor compounds to the final intercalated

products without introduction of H2O into the system. Our procedure for synthesis used the

reduced metals as metal shavings or powders, stirred and suspended in water, which were then

exposed to a carbon dioxide-rich environment by bubbling CO2 generated from a separate

container. Dry ice was used as the initial CO2 generator, since the first experiments were

designed to test the synthetic pathway rather than directly attempt the total deuterated synthesis.

The metals used were Mg and Al, as the Mg-Al LDH systems are by far the most robust and

commonly used systems.

In retrospect, this approach could not have worked because of constraints in the system.

Water, the only available source of OH- in the system, does not function as a good oxidizing

agent for these metals. The system could reach the pH required to dissolve Mg, but could not

dissolve Al at the same time; at the pH required to dissolve Al, the Mg would precipitate as

insoluble Mg(OH)2, creating a pale white material in solution. These experiments therefore

failed, even when allowed to proceed for weeks at a time.

- 26 -

Synthesis of Potassium Tetracyanonickelate(II)

As a related issue that may be of general interest to inorganic chemists, the compound

potassium tetracyanonickelate is fairly easily synthesized in rough (impure) form, and allows a

number of approaches to be taken in its synthesis. The most predictable, stable, and easily

accomplished tactic is to stir an aqueous solution of potassium cyanide with an excess of nickel

(II) cyanide until the solution turns a deep orange. Next, filter the solution to remove as much

leftover nickel (II) cyanide as possible. Boil the solution until its volume is reasonably reduced

or until it appears a deep cherry red color. Finally, cover the solution and allow it to cool and

slowly evaporate for 6-8 weeks. The resulting crystals will be large and almost entirely pure

with the exception of waters of hydration. The overall formula for these large crystals is

K2Ni(CN)4 • nH2O.

For those wishing to take less time, the solution can be boiled vigorously until an orange-

yellow sludge forms. This sludge can be dried by simply continuing to boil away the remaining

water but requires constant stirring and “scraping off” to prevent scorching the material. The

crystals formed by this method are reasonably pure (some potassium cyanide may still be present

in solution) and can be used right away. However, care must be taken to prevent material from

splattering out of the beaker when the sludge begins to dry. This is difficult since the solution

requires constant stirring; in our experience, it is best done by using a large (~4L) beaker and an

initial solution volume of no more than 1L.

Another approach allows stoichiometric addition of potassium cyanide: titration. Simply

titrate a solution of known KCN concentration into the stirred suspension of nickel (II) cyanide;

the titration is finished when the last of the solid disappears. Obtaining the crystals from this

solution is the same as before. One drawback to this approach is that the high pH of the solution

- 27 -

causes CO2 from the air to be drawn into solution as carbonate, [CO3]2-, which then competes

with cyanide for both potassium and nickel counterions. Although cyanide is comparatively high

on the spectrochemical series to carbonate, it has been our experience that some carbonate (in

one case as much as 40-50%) is drawn in due to the high pH (~11), particularly with respect to

potassium in solution.

It has been found that nearly all of this contamination can be removed if the pH of the

solution is sufficiently lowered by using HCl. For our purposes, this acid is an ideal source of

both protons and anions. The presence of Cl- in solution does not pose a problem since our

intent is to intercalate [Ni(CN)4]2- into our LDH. Cl- is the anion used in making the parent

LDH, and [Ni(CN)4]2- competes favorably with Cl-. Also, any excess Cl- is removed in the

washings after intercalation. Finally, any Cl- still present in the sample after washing does not

affect the spectroscopic quality of the resulting compounds.

There are some caveats to this method, however. During this process K2Ni(CN)4 is

converted into Ni(CN)2, and finally to NiCl2 due to the presence of excess chloride ion in

solution. The [CN]- in solution gets converted to HCN, which leaves as a (highly poisonous)

gas. We therefore recommend that anyone using this method perform it in a fume hood with the

sash drawn down, use dilute (10:1) HCl solution, and add this HCl dropwise to the solution

while monitoring the pH. An additional note for those wishing to repeat this work is that the

recrystallization and drying steps MUST be done without the use of ethanol. We have found that

the ethanol-water azeotropic mixture creates poorly characterizable crystals of K2Ni(CN)4 and its

use should therefore be avoided.

By now it should be apparent that the aqueous chemistry of nickel cyanides can become

quite involved. In any given solution, there exists a dynamic equilibrium between [Ni(CN)4]2-

- 28 -

and [Ni(CN)5]3-, which equilibrium constant has been determined to a fair degree of accuracy.

Fortunately, the pentacyanonickelate(II) anion does not appear in quantity until a large excess of

[CN]- is present in solution.113 This species is also thermodynamically unstable with respect to its

components, and decomposes rapidly. Pure crystals of K3Ni(CN)5, which is described as having

a deep red color, were obtained by Raymond and Basolo in 1966.114

We also performed a limited experiment using concentrated (37% wt.) HCl, adding acid

dropwise to a solution of fresh tetracyanonickelate(II) that we obtained from the “direct”

synthesis method. While adding the acid, a white precipitate was observed to form in small

quantities, which dissolved immediately on stirring. We have been unable to isolate this material

from solution as yet, however, based upon the compound’s color, it can not be the expected

Nickel (II) Chloride. It may be tetrachloronickelate(II) or finely powdered

tetracyanonickelate(II) acid salt, H2Ni(CN)4. More experiments are necessary before this

compound’s identity can be properly determined.

- 29 -

CONCLUSIONS

Tetracyanonickelate(II) was successfully intercalated into both 2:1 and 3:1 Mg:Al layered

double hydroxide (LDH) systems. Comparison of this intercalation between LDH chloride

precursors and LDH nitrate precursors revealed no significant differences between the two types

of precursor. A small amount of nitrate was found to remain in the system after intercalation, as

evidenced by the appearance of a ν(NO3) peak in the infrared spectra of the nitrate-derived LDH

nickelocyanides; it is reasonable to speculate that some chloride would remain in the other

systems as well. XRD data show no significant differences in post-intercalation gallery heights

between the nitrate-derived LDH nickelocyanides and those made from chloride parents. Based

on this data, the tetracyanonickelate(II) anion is believed to insert itself at an angle of

approximately 30° in the 2:1 systems, and is believed to lie nearly parallel to the layers in 3:1

systems. No conclusions are drawn from the results of the elemental analysis due to

inconsistency of the results.

The two-dimensional average charge density of a single [Ni(CN)4]2- anion can be

estimated by dividing the total anion charge by the calculated area of one square-planar anion.

By using the value of 7.35 Å calculated earlier, we arrive at area of 0.540 nm2; with a -2 charge,

the charge density is estimated as 2/0.540 = 3.70 e/nm2. This value can be compared to the value

for the charge densities in 2:1 and 3:1 Mg-Al LDH layers, which can be estimated using the

following formula:

°=

60sin C 2d a

x

- 30 -

where x is the charge per unit cell and a is the literature value of the a-spacing from AFM of

prepared samples.115 Using x=0.33 for 2:1 Mg-Al LDH (one excess positive charge for 3 metal

centers) and a=0.31 nm, we arrive at a value of 4.0 e/nm2. For the 3:1 Mg-Al LDH systems,

using x=0.25 and a=0.31 nm gives a value of 3.0 e/nm2. Thus we find that the charge density of

[Ni(CN)4]2- is expected to be between the densities of 2:1 and 3:1 Mg-Al LDHs. This would

mean that the anions in the 2:1 Mg-Al LDH system would need to tilt somewhat in order to

present a sufficient charge density to the layer. In contrast, the 3:1 Mg-Al LDH system receives

more than enough charge density to satisfy its positive layer charges, and so the anions can lie

parallel to the layers.

Comments on the “Total Synthesis” Approach

The most obvious difficulty in using reduced metals for the starting materials is the need

to oxidize them before any relevant chemistry can take place. While there are a variety of means

available to do this, most come with some severe drawbacks. Purchasing isotopically pure

substances is very expensive and those substances are often easily contaminated by chemically

identical but isotopically different compounds (e.g D2O is easily contaminated with H2O). Many

chemical oxidizing agents either draw water from the surrounding environment or are already

associated with significant amounts of water (or both) and are thus unfit for use in a hydrogen-

free system. One other method was attempted, which tried to use the halogen iodine (as I2) as

both an oxidizing agent and a source of anions for the resulting LDH. The suspension of reduced

metals was exposed to a large excess of I2 in hopes of directly forming a small amount of LDH

material. Infrared analysis of the resulting materials showed a poorly characterizeable system.

- 31 -

The sharp peak around 3650 cm-1 is indicative of Mg(OH)2 in the system. The spectrum of the

products is shown below in Figure 8.

Figure 8. Infrared spectrum of material obtained from reaction of Mg and Al with I2.

4000 3500 3000 2500 2000 1500 1000 500Wavenumber (cm-1)

AB

S

- 32 -

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