BONDS BETWEEN ATOMS

Metallic bonding occurs between atoms which have similar low electronegativity values. In this form of bonding, each atom shares its valence electron(s) with every other atom in the structure, i.e., the electrons are pooled , free , or delocalized . The electrostatic attraction between the positive ions and the electron pool holds the structure together (Fig. 11.4). As with ionic bonding, each positive ion feels the influence of a large number of 

Covalent bonding occurs when the electronegativity values of the participating atoms are high, but similar in value. Therefore, covalent bonding is electron 

There is a wide range of molecules which use covalent bonding, and they can exhibit a wide range of physical and chemical properties. Many of these differences result from the scale of the bonding. Molecular covalent molecules 

---

The ROSDAL syntax is characterized by a simple coding of a chemical structure using alphanumeric symbols which can easily be learned by a chemist [14]. In the linear structure representation, each atom of the structure is arbitrarily assigned a unique number, except for the hydrogen atoms. Carbon atoms are shown in the notation only by digits. The other types of atoms carry, in addition, their atomic symbol. In order to describe the bonds between atoms, bond symbols are inserted between the atom numbers. Branches are marked and separated from the other parts of the code by commas [15, 16] (Figure 2-9). The ROSDAL linear notation is rmambiguous but not unique. 

I iiis can be helpful because it may enable more meaningful sets of orbitals to be generated from the original solutions. Molecular orbital calculations may give solutions that are nioared out throughout the entire molecule, whereas we may find orbitals that are Im alised in specific regions (e.g. in the bonds between atoms) to be more useful. 

The ring closure bond between atoms 1 and 5 would be strongly coupled to the other internal coordinates inless dummy atoms are used to define the Z-matrix (right). 

In the case of a triatomic molecule with two bonds (between atoms 1,2 and 2,3), twi constraint equations are obtained ... 

To understand the function of a protein at the molecular level, it is important to know its three-dimensional stmcture. The diversity in protein stmcture, as in many other macromolecules, results from the flexibiUty of rotation about single bonds between atoms. Each peptide unit is planar, ie, oJ = 180°, and has two rotational degrees of freedom, specified by the torsion angles ( ) and /, along the polypeptide backbone. The number of torsion angles associated with the side chains, R, varies from residue to residue. The allowed conformations of a protein are those that avoid atomic coUisions between nonbonded atoms. 

In this chapter, three methods for measuring the frequencies of the vibrations of chemical bonds between atoms in solids are discussed. Two of them, Fourier Transform Infrared Spectroscopy, FTIR, and Raman Spectroscopy, use infrared (IR) radiation as the probe. The third, High-Resolution Electron Enetgy-Loss Spectroscopy, HREELS, uses electron impact. The fourth technique. Nuclear Magnetic Resonance, NMR, is physically unrelated to the other three, involving transitions between different spin states of the atomic nucleus instead of bond vibrational states, but is included here because it provides somewhat similar information on the local bonding arrangement around an atom. 

For most combinations of atoms, a number of molecular structures that differ fk m each other in the sequence of bonding of the atoms are possible. Each individual molecular assembly is called an isomer, and the constitution of a compound is the particular combination of bonds between atoms (molecular connectivity) which is characteristic of that structure. Propanal, allyl alcohol, acetone, 2-methyloxinine, and cyclopropanol each correspond to the molecular formula CjH O, but differ in constitution and are isomers of one another. 

The theory of atoms in molecules defines chemical properties such as bonds between atoms and atomic charges on the basis of the topology of the electron density p, characterized in terms of p itself, its gradient Vp, and the Laplacian of the electron density V p. The theory defines an atom as the region of space enclosed by a zero-/lMx surface the surface such that Vp n=0, indicating that there is no component of the gradient of the electron density perpendicular to the surface (n is a normal vector). The nucleus within the atom is a local maximum of the electron density. 

The other syntax for supplying molecular structures to Gaussian 94 is the Z-matrix. A Z-matrix specifies the locations of and bonds between atoms using bond lengths, bond angles, and dihedral (torsion) angles. 

F. A. Cotton and R. A. Walton, Multiple Bonds between Atoms, 2nd edn., Oxford University Press, Oxford, 1993, 787 pp. 

Once the density matrix has been transformed to the NAO basis, bonds between atoms may be identified from the off-diagonal blocks. The procedure involves the following steps. 

Shortly after the tetravalent nature of carbon was proposed, extensions to the Kekule-Couper theory were made w7hen the possibility of multiple bonding between atoms was suggested. Emil Erlenmeyer proposed a carbon-carbon triple bond for acetylene, and Alexander Crum Brown proposed a carbon-carbon double bond for ethylene. In 1865, Kekule provided another major advance when he suggested that carbon chains can double back on themselves to form rings of atoms. 

How does electron sharing lead to bonding between atoms Two models have been developed to describe covalent bonding valence bond theory and molecular orbital theory. Each model has its strengths and weaknesses, and chemists tend... 

We saw in the last chapter how covalent bonds between atoms are described, and we looked at the valence bond model, which uses hybrid orbitals to account for the observed shapes of organic molecules. Before going on to a systematic study of organic chemistry, however, we still need to review a few fundamental topics. In particular, we need to look more closely at how electrons are distributed in covalent bonds and at some of the consequences that arise when the electrons in a bond are not shared equally between atoms. 

It is found experimentally that a bond between two atoms with very different ionization energies tends to be stronger than a bond between atoms with similar ionization energies. Since electric dipoles are caused by differences in the ionization energies of bonded atoms, we can conclude that strong bonds are expected in molecules with electric dipoles. 

Contrast the bonds between atoms in metals, in van der Waals solids, and in network solids in regard to ... 

Because nonmetals do not form monatomic cations, the nature of bonds between atoms of nonmetals puzzled scientists until 1916, when Lewis published his explanation. With brilliant insight, and before anyone knew about quantum mechanics or orbitals, Lewis proposed that a covalent bond is a pair of electrons shared between two atoms (3). The rest of this chapter and the next develop Lewis s vision of the covalent bond. In this chapter, we consider the types, numbers, and properties of bonds that can be formed by sharing pairs of electrons. In Chapter 3, we revisit Lewis s concept and see how to understand it in terms of orbitals. 

The charges on the atoms in HCI are called partial charges. We show the partial charges on the atoms by writing 8+l I—Cl8. A bond in which ionic contributions to the resonance result in partial charges is called a polar covalent bond. All bonds between atoms of different elements are polar to some extent. The bonds in homonuclear (same element) diatomic molecules and ions are nonpolar. 

The values in Table 2.4 show how resonance affects the strengths of bonds. For example, the strength of a carbon-carbon bond in benzene is intermediate between that of a single and that of a double bond. Resonance spreads multiple bond character over the bonds between atoms as a result, what were single bonds are strengthened and what were double bonds are weakened. The net effect overall is a stabilization of the molecule. 

We have to refine our atomic and molecular model of matter to see how bulk properties can be interpreted in terms of the properties of individual molecules, such as their size, shape, and polarity. We begin by exploring intermolecular forces, the forces between molecules, as distinct from the forces responsible for the formation of chemical bonds between atoms. Then we consider how intermolecular forces determine the physical properties of liquids and the structures and physical properties of solids. 

In the past, ionic radii have often been compared with observed interatomic distances without much regard to the nature of the crystal from which they were derived. Recently several investigators19 have concluded that in many crystals the bond between atoms does not consist of the electrostatic attraction of only slightly deformed ions. Goldschmidt in particular has divided crystals into two classes, ionic and atomic crystals, and has shown that ionic radii (using Wasastjema s set) do not account for the observed inter-atomic distances in atomic crystals. In the following pages our crystal radii will be compared with the experimental dis-... 

---