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Cover Chemical Bonding

Simple bonding schemes  

This chapter discusses the Lewis description of bonding as described by the combination of Lewis atoms. The three principal chemical bond types are covalent, ionic, and metallic. A covalent bond is the interaction by which two atoms are held together by a pair of electrons located between the two atoms. Ionic compounds consist of regular arrays of cations and anions and ionic bonding consisting of electrostatic attractions experienced between the electric charges of the cations and anions. Meanwhile, metals may be regarded as a regular lattice array of cations, each of which consists of the metal atom less the valence electrons; the valence electrons are detached from any single metal cation and the spaces between the cations are filled with a ‘sea’ of detached electrons.


Cover Chemical Structure and Reactivity

Bonding between the elements  

This chapter describes what happens when different elements combine. The sizes of the orbitals involved and the number of valence electrons is important when considering the bonding in the elements, but when different elements combine there is the added complication that their orbital energies will be different. This leads to the possibilities of polar bonds and compounds in which electrons have essentially been transferred from one element to another, i.e. ionic compounds. Depending on the energy separation of the orbitals involved, the bonding can vary on a spectrum from purely covalent to ionic. Trying to predict the structure that a particular compound will adopt is very difficult. However, there are a few general concepts which can be applied to at least give some indication as to the likely structure. The concept of the radius ratio can help one to understand the different structures adopted by ionic solids.


Cover Foundations of Inorganic Chemistry


This chapter examines bonding, which is a consequence of electrostatic interactions between positively charged nuclei and negatively charged electrons. The type of bonding that holds atoms together in a metal, such as iron, is called metallic bonding. Compounds such as salt (sodium chloride) consist of lattices of anions and cations held together via electrostatic forces—this is ionic bonding. Meanwhile, compounds such as oxygen and nitrogen, in the air that we breathe, are held together by shared-electron bonds—this is covalent bonding. Most compounds are best described as being somewhere between purely covalent and purely ionic in nature. The chapter then looks at the Lewis model of bonding, van der Waals forces, and the concept of the oxidation state.


Cover Making the Transition to University Chemistry

Bonding and Molecular Shape  

This chapter discusses bonding and molecular shape of electrons. A covalent bond occurs when atoms share a pair of similar electrons. Modern theories on covalent bonding are dominated by molecular orbital theory. Polar covalent bonds occur when different atoms share the electron pair unequally due to electronegativity. With benzene being the most familiar molecule of the bonding, delocalization happens when more than two atoms are involved in the bonding. The valence-shell electron-pair repulsion (VSEPR) theory can help us to visualize the shapes of simple molecules. Additionally, ionic bonding occurs when an atom transfers an electron to another atom and the ions formed a crystal lattice electrostatically.


Cover Chemistry for the Biosciences

Compounds and chemical bonding: bringing atoms together  

This chapter examines compounds and chemical bonding. A compound is a substance that comprises atoms of more than one element. Chemical bonds hold the components of a compound together and are formed by the redistribution of valence electrons between atoms. According to the octet rule, valence electrons are redistributed so that atoms achieve full valence shells that typically contain eight electrons. The chapter then differentiates between the two types of chemical bond: ionic and covalent. An ionic bond forms when one or more electrons are totally transferred from one atom to another to generate an ionic compound. A covalent bond, however, forms when one or more pairs of electrons are shared between atoms to generate a covalent compound. The chapter looks at polarized bonds, in which electrons are shared unequally between two nuclei, and discusses how valency describes the number of chemical bonds an atom of a given element can participate in. It also describes chemical bonding in terms of atomic and molecular orbitals before discussing aromatic compounds and polyatomic ionic compounds.


Cover Chemical Structure and Reactivity

Molecules and molecular structures: an overview  

This chapter provides an overview of the key concepts that are commonly used in describing bonding and molecular structures, together with some related ideas, such as delocalization, oxidation state, and formal charge. It looks at the different types of bonding (covalent, ionic, and metallic) and their particular characteristics. The bonding in simple covalent molecules can be described using Lewis structures in which the focus is on bonding and non-bonding pairs of electrons. The shapes of simple molecules can often be predicted using the valence shell electron pair repulsion (VSEPR) model. Meanwhile, molecular structures are determined using X-ray diffraction. The chapter then considers the different types of solid materials, and the behaviour of the ideal gas.


Cover Chemistry3


This chapter looks closely at covalent, metallic, and ionic bonding in solid state structures, which is important as the properties of solid state materials depend on their structures and bonding. Some of the characteristic/properties of molecular solids, covalent network structures, metals, and ionic solids are summarized. The chapter describes the differences between cubic close packing (ccp) and hexagonal close packing (hcp). It demonstrates how to predict the limiting radius ratio for different geometries and how to use the radius ratio rule to predict the structures of ionic compounds. It also outlines how to calculate packing efficiencies and densities, lattice enthalpies using Born–Haber cycle compounds, and lattice energies using the Born–Landé equation and the Kapustinskii equation.


Cover f-Block Chemistry

Introducing the f-elements  

This chapter provides an overview of the f-elements: lanthanoids and actinoids. All of the lanthanoid elements (apart from Pm) occur naturally in greater abundance than either iodine or mercury, but they always occur as mixtures in their ores. This is due to the great similarity in their properties. All of the actinoids are radioactive, and only uranium and thorium occur naturally to any significant extent. The characteristic feature of lanthanoids and actinoids includes the progressive filling of the 4f and 5f-orbitals respectively. The f-orbitals of the lanthanoids and the later actinoids behave as core orbitals and do not take part in bonding; however, the greater radial extent of the 5f-orbitals of the early actinoids means that they can contribute to bonding. The chapter then looks at the metallic and ionic radii of both Ln and An.


Cover Chemical Structure and Reactivity

Bonding in solids  

This chapter focuses on bonding in solids, particularly metallic bonding. In solids, the overlap between orbitals on different atoms can give rise to crystal orbitals which extend throughout the material; these orbitals are analogous to delocalized molecular orbitals. The crystal orbitals which arise from a particular set of atomic orbitals form a band, which can hold a certain number of electrons. The electrical conductivity of metals is the result of partially filled bands; insulators have full bands, but in semiconductors there is a small gap between a filled and an empty band. The lattice enthalpy of an ionic solid can be estimated using a simple electrostatic model. It depends on the size (radii) of the ions and their three-dimensional arrangement in the crystal.


Cover The Electronic Structure and Chemistry of Solids

Electronic energy levels and chemical bonding  

This chapter examines a number of important features of the electronic structure of the solid, which include energies and widths of the various bands, the energy gaps between them, and the number of electrons that occupy them. It details how the electronic energy levels of simple solids are related to chemical models of bonding. It also highlights the ionic model, wherein bonding is accompanied by an electron transfer between one atom to another. The chapter highlights the use of the ionic model in predicting the heats of formation and the chemical and physical properties of compounds, such as halides and oxides. It considers the solids formed by non-metallic elements and looks at compounds where covalent bonding is combined with some degree of ionic character.


Cover Organometallics and Catalysis

Organometallic Compounds of Main Group Elements  

This chapter discusses the main group of organometallic compounds which contain metal–carbon σ-bonds generated by orbital overlap along the M–C axis. It reviews the possibility of interaction of the metal with the π-electron systems of unsaturated organic compounds, in particular, when the unsaturated moiety carries a negative charge. The chapter also looks at one important aspect that governs the reactivity of organometallic species: the polarity of the Mδ+ Cδ-bond. It explicates the nature of bond polarity, then studies the concept of electronegativity and ‘group electronegativity’. Towards the end, the chapter turns to the reactivity of main group metal alkyls. It also considers ionic organometallics and lithium ions.


Cover Periodicity and the s- and p-Block Elements

Compounds of the s- and p-block elements  

This chapter assesses s- and p-block element compounds and their properties. It begins by looking at halides, which is a large and diverse group of compounds, particularly in terms of the variety of structural types encountered. Halides include fluorides and chlorides. The chapter then considers ionic compounds, element oxides, and element hydrides. All the hydrides are molecular covalent species, but the trends in melting and boiling points are worthy of comment. In Group 14, one sees a progression to higher melting and boiling points as the group is descending resulting from the increasing van der Waals forces between the molecules. This is also largely the case for the Group 16 hydrides but with the obvious exception of H2O. The explanation for this feature is the extensive intermolecular hydrogen bonding between oxygen lone pairs and hydrogen, this being that much greater for the lighter element oxygen due to its high electronegativity.