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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.

Chemical Bonding I: Basic Concepts examines general ideas of chemical bonding between atoms and ions and how this bonding affects the chemical properties of the elements. An overview of Lewis symbols, Lewis structures and the octet rule is presented including the role of valence electrons in ionic and covalent bonding. The energy changes that accompany ionic bond formation are also discussed with emphasis on lattice energy. The chapter covers guidelines and general procedures for writing Lewis structures or electron dot formulas for molecular compounds and polyatomic ions. The concepts and applications of resonance, formal charge and exceptions to the octet rules are presented, along with coverage of the relationship between bond polarity and electronegativity.

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-.