Bio-inspired Synthesis of Oxide Nanostructures

Nature is capable of building magnificently intricate and detailed structures out of otherwise boring materials such as calcium carbonate and silica. Anyone who has taken their children to see dinosaurs at a Natural History museum or visited natural wonders such as the Petrified Forest in Arizona are familiar with the natural process called fossilization by which the tissues of dead organisms are eventually replicated by objects of stone. Most living organisms (including humans) are critically dependent on more deliberate and controlled biomineralization phenomena that lead to the production of all hard tissues, including our teeth and bones, seashells and diatom skeletons, egg shells, and the magnetic nanoparticles that provide homing devices from bacteria to birds. All these processes are nothing more than specific examples of highly controlled nucleation and growth phenomena such as those described in generic terms in Chapter 7. At a molecular level, these processes are controlled by the same reaction mechanisms involving oxide surfaces, which were outlined in Chapter 6. However, biomineralization is orders of magnitude more sophisticated than standard nucleation and growth processes. The unique features of biomineralization involve the interplay between organic biomolecules and the nucleation and growth of inorganic phases such as oxides. This interplay is of critical importance in both biology and emerging nanotechnologies, providing specific examples that illustrate many of the concepts of oxide chemistry introduced in Chapters 5 through 7. In this chapter, we highlight the key concepts of biomineralization and provide examples of how researchers can now produce complex nanostructured oxides via biomimetic nucleation and growth strategies that replicate some of the key features used to make hard tissues in living systems. These strategies include the use of (1) molecular complexation and compartmentalization to control supersaturation levels, (2) specific ligands and surface structures to mediate nucleation phenomena, (3) hierarchical self-assembled organic architectures as templates for oxide formation, (4) functionalization to stimulate desired heterogeneous nucleation and growth processes on those templates, and (5) organic surfactants to manipulate both crystal-phase preferences and growth habits.

The Chemistry of Extended Oxide Surfaces

In Chapters 4 and 5, we demonstrated that local structures and charge distributions have an enormous impact on the equilibrium constants, trajectories, and kinetics of reactions involving soluble oxide precursors. In this chapter, we highlight those features that make reactions on extended oxide surfaces either similar to or dramatically different from the reactions documented in hydrolysis diagrams for each metal cation (see Chapter 5). We first describe oxide surface structures and then discuss how these structures impact both acid–base and ligand-exchange phenomena. In addition to dense oxides, we also introduce some of the chemistry associated with layered materials. Lamellar materials are important from both a fundamental and technological perspective, because water and ions can readily penetrate such structures and provide conditions under which almost every oxygen anion is at an oxide–water interface (see Chapter 10 and Chapter 11). This chapter focuses on oxides containing octahedral cations. The distinctive chemistry of oxides based on tetrahedral cations, including the clay minerals and the zeolites, are the focus of Part Five. The structures of bulk oxides were introduced in Chapter 2. However, for many oxides, the surface structures that interact with aqueous solutions are substantially different from structures found in the bulk. Here, we introduce the basic principles of oxide surfaces that make them chemically active. As a starting point, consider ideal oxide surfaces containing +2 octahedral cations. Pristine oxide surfaces can be created by cleaving perfect crystals in an ultrahigh-vacuum environment. The creation of new surfaces requires an expenditure of energy corresponding to the cohesive energy of the solid, which in turn represents the energy required to break every bond along a given fracture plane. For MgO, the Mg−O bond energy is 380 kJ/mole. Each surface created contains 1.4.1019 oxygen atoms/m2, or 2.4.10−5 moles of bonds. Because two surfaces are created in the fracture event, the initial interfacial energy of each resulting MgO surface is (1/2)(380 kJ/mole)/(2.4_10−5 mole/m2 )=4560 mJ/m2.

Solvated Ions in Water

In most undergraduate chemistry classes, students are taught to consider reactions in which cations and anions dissolved in water are depicted as isolated ions. For example, the magnesium ion is depicted as Mg2+, or at best Mg2+(aq). For anions, these descriptions may be adequate (if not accurate). However, for cations, these abbreviations almost always fail to describe the critical chemical attributes of the dissolved species. A much more meaningful description of Mg2+ dissolved in water is [Mg(H2O)6]2+, because Mg2+ in water does not behave like a bare Mg2+ ion, nor do the waters coordinated to the Mg2+ behave anything like water molecules in the bulk fluid. In many respects, the [Mg(H2O)6]2+ ion acts like a dissolved molecular species. In this chapter, we discuss the simple solvation of anions and cations as a prelude to exploring more complex reactions of soluble oxide precursors called hydrolysis products. The two key classes of water–oxide reactions introduced here are acid–base and ligand exchange. First, consider how simple anions modify the structure and properties of water. As discussed in Chapter 3, water is a dynamic and highly fluxional “oxide” containing transient rings and clusters based on tetrahedral oxygen anions held together by linear hydrogen bonds. Simple halide ions can insert into this structure by occupying sites that would normally be occupied by other water molecules because they have radii (ranging from 0.13 to 0.22 nm in the series from F− to I−) that are comparable to that of the O2− ion (0.14 nm). Such substitution is clearly seen in the structures of ionic clathrate hydrates, where the anion can replace one and sometimes even two water molecules. Larger anions can also replace water molecules within clathrate hydrate cages. For example, carboxylate hydrate structures incorporate the carboxylate group within the water framework whereas the hydrophobic hydrocarbon “tails” occupy a cavity within the water framework, as in methane hydrate (see Chapter 3). Water molecules form hydrogen bonds to dissolved halide ions just as they can to other water molecules, as designated by OH−Y−.

The Structure and Properties of Water

Water is one of the most complex fluids on Earth. Even after intense study, there are many aspects regarding the structure, properties, and chemistry of water that are not well understood. In this chapter, we highlight the attributes of water that dictate many of the reactions that take place between water and oxides. We start with a single water molecule and progress to water clusters, then finally to extended liquid and solid phases. This chapter provides a baseline for evaluating what happens when water encounters simple ions, soluble oxide complexes called hydrolysis products, and extended oxide phases. The primary phenomenon highlighted in this chapter is hydrogen bonding. Hydrogen bonding dominates the structure and properties of water and influences many water–oxide interactions. A single water molecule has eight valence electrons around a central oxygen anion. These electrons are contained in four sp3-hybridized molecular orbitals arranged as lobes that extend from the oxygen in a tetrahedral geometry. Each orbital is occupied by two electrons. Two of the lobes are bonded to protons; the other two lobes are referred to as lone pairs of electrons. The H–O–H bond angle of 104.5° is close to the tetrahedral angle of 109.5°. The O–H bond length in a single water molecule is 0.96 Ǻ. It is important to recognize that this bond length is really a measure of the electron density associated with the oxygen lone pair bonded to the proton. This is because a proton is so incredibly small (with an ionic radius of only 1.3·10−5 Ǻ) that it makes no contribution to the net bond length. The entire water molecule has a hard sphere diameter of 2.9 Ǻ, which is fairly typical for an oxygen anion. This means the unoccupied lone pairs are distended relative to the protonated lone pairs, extending out to roughly 1.9 Ǻ. The unequal distribution of charges introduces a dipole within the water molecule that facilitates electrostatic interactions with other molecules.

The Impact of Oxides on Environmental Chemistry

The ancient Greek philosopher Empedocles defined our environments using the four basic elements of fire, earth, wind, and water. Although we now know there are at least 118 elements, of which 98 are naturally occurring, these ancient descriptions aptly describe the habitats on Earth that are occupied by oxides and living things. Many oxides that comprise Earth’s surface are born by the fire represented by the massive heat of Earth’s interior as mediated by plate tectonics. This heat produces the igneous rocks found in volcanoes and our major mountain chains. Water weathers these pristine rocks, which are gradually broken down to form earth, which includes the wide diversity of other rock types, soils, and sediments covering the surfaces of our continents and ocean floors. Weathered oxides in the form of dust are blown by wind and enter the atmosphere, where they influence the chemistry of the air we breathe and the rainfall that supports continental life. The chemical transformations of oxides are strongly influenced by all the environmental conditions they encounter in their life cycle (see Chapter 17). Conversely, the interactions between oxides, water, and organisms help define many of the environments that allow life on Earth to thrive. These interactions form the basis for this final chapter of our book. Oxides are present in all our planet’s major environments. In this chapter, we explore each of the environments defined by the ancient Greeks in descending order based on their distance from Earth’s core. The chapter progresses from the stratosphere (air) to continental surfaces (earth) to our oceans (water) and finally to the subsurface environments of subduction zones such as the Marianas Trench (fire). In each section, we highlight reactions involving the two most important classes of oxides in terms of their environmental impact, both of which are weathering products: (1) the clay minerals and (2) the redox-active colloids of iron and manganese oxides. Clay mineral reactions impact colloidal interactions (Chapter 8), ion exchange (Chapter 10), and the sequestration of environmental nutrients and contaminants. Reactions of the redox-active oxidates of iron and manganese are dominant in terms of reversible and often complex electrochemical (Chapter 11) and photochemical (Chapter 13) processes that take place in natural environments.

Photochemistry and Excited-State Reactions of Oxides

The applied voltages that drive electrochemical processes (see Chapter 11) are only one of many energy sources that can be used to activate reactions in oxide molecules and materials. Another common energy source that drives many environmental and technological oxide reactions is light from the sun. Water plays a key role in many of these reactions. Imagine that you are on vacation floating in a warm ocean bathed by the sun. Many of the phenomena you experience, from your painful sunburn to the photosynthetic growth of the seaweed you see beneath you, are photoactivated processes. In this chapter, we highlight the roles that oxides play in photon-activated solar energy technologies. Also included are reactions stimulated by other nonthermal energy sources, including electrons in high-energy plasmas. Titanium oxide, found in common white paint, is the basis for much of the discussion, because this oxide is used in many photoelectrochemical energy storage technologies. The photochemistry of colloidal manganese- and iron-oxide particles suspended either in atmospheric droplets or in the upper photic zone of the ocean where the sunlight penetrates are discussed in Chapter 18. Such oxide reactions are important globally in the elimination of pollutants. Both industrial and environmental examples illustrate how oxides participate in a wide range of photoactivated chemical reactions, including the catalytic decomposition of water, photoelectrochemistry, and photoactivated dissolution and precipitation reactions. Before exploring excited-state reactions, we need to introduce the energy sources that provide such excitation. In most of this chapter, the excitation source of interest is light. Most of us are familiar with the electromagnetic spectrum, in which the energy of a photon is given by . . . E=hv=hc/λ=hcω (13.1). . . Here, h is Planck’s constant (h = 6.6 ·10 −34 J/second), c is the speed of light (3 ·1010cm/second), ν is the frequency of light (measured in Hertz or per second), λ is the wavelength of light (in centimeters), and ω is the wavelength expressed as wave number (measured per centimeter in infrared spectroscopy).

The Colloidal Chemistry of Oxides

Colloids are defined as suspensions of finely divided particles in a continuous medium that do not settle rapidly and are not readily filtered. To be more specific, the International Union of Pure and Applied Chemistry defines a colloid as any material for which one or more of its three dimensions lies within the size range of 1 to 1000 nm. As the nucleation and growth of oxides from aqueous solutions almost always produces suspensions containing submicron particles (see Chapter 7), typical oxide suspensions fall squarely within the colloidal domain. In this book, we consider colloidal particles to represent oxides or hydroxides that are small enough to stay in aqueous suspensions for more than a few hours, yet are larger and lacking in the specific molecular structures of typical hydrolysis products (see Chapter 5). Given the density range of most oxides (from around 2−10 g/cm3), the sizes of most colloidal oxides fall within the limits of the International Union of Pure and Applied Chemistry (see Section 8.4.5). Colloidal oxide particles suspended in water represent a complex chemical environment. At the molecular level, protons, ions, small molecules, and polymeric species interact with particle surfaces to create charged surface sites and promote adsorption and desorption phenomena (see Chapter 6). These modified surfaces perturb the adjacent liquid, creating ordered solvent layers and strong concentration gradients in ions and other dissolved species. These interfacial phenomena generate a range of forces called interaction potentials. Such forces determine whether particles repel each other (leading to stable suspensions) or are attracted to one another, resulting in agglomeration and sedimentation phenomena. The length scales of those components of the oxide–water interface that influence the interaction potentials to be discussed in this chapter are introduced in Figure 8.1. At the subatomic level, the correlated polarization of electron clouds gives rise to dispersion forces described by quantum mechanics that contribute to van der Waals interactions. At the atomic level, the inherent charge on each exposed oxygen anion that terminates the oxide surface is controlled by local chemical bonds to adjacent cations (see Chapter 6).

Glass Dissolution and Leaching

Oxide glasses represent some of the most important and prevalent materials that we encounter in our daily lives. The glass industry in the United States produces more than 75,000 glass products, with annual production estimated to be around 20,000,000 t. Roughly 50% of this production is for glass containers for food, beverages, and other liquids. Everyone relies on transparent glass windows for their homes, cars, and even their cell phones. Fiberglass provides insulation for our homes and businesses. We rely on glass for many optical systems, ranging from eyeglasses to microscope lenses to optical fiber communications. Glass is also an optically pleasing material found in many works of art, including stained glass windows. Glass even plays a role in energy transport and storage, being an important electrical insulator used in devices ranging from transformers to batteries. Glass compositions need to be optimized for specific applications, with important parameters being melting properties, thermal conductivity, thermal expansion, strength, dielectric properties, and, of course, optical properties. In most of these applications, glass objects encounter water, either to perform their basic functions or as a result of long-term environmental exposure. This means the chemical properties of many glasses also need to be optimized. Fortunately, borosilicate glasses, which represent the most widely used technological glass compositions, tend to exhibit a high level of resistance to aqueous attack. Understanding the kinetics and mechanisms of glass dissolution is critically important to the nuclear power and defense industries, which involves how to dispose of nuclear wastes safely. These wastes can be exceedingly complex, and contain almost every element found in the Periodic Table. The challenge is to incorporate these wastes into solids that encapsulate radionuclides safely for millions of years. Glass is an attractive option as a waste form because glass melts can accommodate almost all the constituents found in nuclear wastes. However, the deployment of glass waste forms requires the ability to predict the stability of the waste out to exceedingly long times based on science-based glass-dissolution models.

Nucleation and Growth of Solid Oxide and Hydroxide Phases

In this chapter, we consider what happens when solids begin to form from solution. To grow solids from solution, solution conditions are changed from a condition in which all species are completely soluble to a condition in which they are insoluble. In the context of hydrolysis diagrams, the solution composition moves in pH and total dissolved metal concentration from a regime below a solubility or saturation limit (given by the bold solid line in Figs. 5.2 and 5.3) to a regime above this limit where the solution is supersaturated. Supersaturated solutions are inherently unstable and have the potential to generate hydroxide or oxide solids. Sometimes these solutions can be maintained in a metastable state in which precipitation does not occur immediately. However, Mother Nature eventually reduces the energy of the solution by forming a stable mixture of solids plus solution species. As solids form, soluble complexes are removed from solution until concentrations drop back to the solubility limit. The precipitation of a solid from an aqueous solution is a surprisingly complex process, involving nucleation and growth phenomena that occur at nanometer-length scales. Nucleation involves reactions between oligomers to form new clusters or particles that are sufficiently large that they do not redissolve spontaneously via the reversible reactions denoted in hydrolysis diagrams. Homogeneous and heterogeneous nucleation processes represent events that occur within the bulk solution or at the interface of another phase, respectively. Growth involves the addition of monomers to clusters in solution or oligomers to existing particles or surfaces. The combination of nucleation and growth phenomena can lead to oxides exhibiting a bewildering range of sizes, shapes, and crystal structures. How do metal complexes decide whether to form a new particle or add to an existing particle? What determines the size, shape, and crystal structure of evolving particles? Do the particles aggregate with one another in an organized fashion? Because nucleation typically involves extremely rapid (<1 millisecond) events involving objects that are extremely small (on the order of a nanometer), it is difficult to probe such phenomena at a molecular level.

An Overview of Oxide Structures and Compositions

This entire book is devoted to exploring the chemistry of compounds that contain one simple anion: the O2− ion. Except under high-vacuum conditions (see Chapter 6), the species in oxides that interact with water and other environmental chemicals are O2− ions, because the charge-compensating cations are invariably buried beneath an oxide surface layer. One might wonder how we can fill an entire volume discussing the chemical interactions between this single anion and a single chemical: the water molecule. The single most important concept that must be appreciated to understand the contents of this book is that the chemistry and properties of O2− anions are critically dependent on all the cations to which the O2− ions are bound. Each bound cation modifies the electron distributions around the O2− site, changing its local charge, local bonding configurations, acid–base chemistry, ion exchange chemistry, electrochemical properties, chemical stability, and electrical and optical properties. None of these changes are subtle, and in fact most oxide properties are staggering in their diversity. Before considering the chemistry of oxides, it is important to gain an appreciation of just how diverse the structures of oxide materials really are. As this introduction makes clear, there is no such thing as a single, simple O2− ion. There are a myriad of different O2− sites found in the oxides that we encounter often in our daily lives, each of which exhibits its own unique properties. The purpose of this book is to provide a framework that can be used to predict, rationalize, and exploit the rich chemistry associated with those sites. The number of different structures and compositions that can be generated for oxides is almost limitless. The O2− ion forms compounds with more than 90 elements in the Periodic Table that are capable of losing electrons to form cations. The oxide anion combines with cations with charges that range from +1 to +7. Many elements exhibit more than one stable oxidation state, pushing the total number of chemically distinct cations with which O2− can interact to well more than 120.