Atoms, Electrons, and Orbitals

3 Covalent Bonds, Lewis Formulas.and the Octet Rule 8 

5 Polar Covalent Bonds, Electronegativity, and Bond Dipoles 10 

9 Sulfur and Phosphorus-Containing Organic Compounds and the Octet Rule 23 

Before discussing stmcture and bonding in molecules, let s hrst review some fundamentals of atomic stmcture. Each element is characterized by a unique atomic number Z, which is equal to 

FIGURE 1.2 Cross sections of (a) a Is orbital and (b) a 2s orbital. The wave function has the same sign over the entire Is orbital. It is arbitrarily shown as +, but could just as well have been designated as -. The 2s orbital has a spherical node where the wave function changes sign. 

A hydrogen atom (Z = 1) has one electron a helium atom (Z = 2) has two. The single electron of hydrogen occupies a 1 orbital, as do the two electrons of helium. The respective electron configurations are described as  

The period (or row) of the periodic table in which an element appears corresponds to the principal quantum number of the highest numbered occupied orbital (n = 1 in the case of hydrogen and helium). Hydrogen and helium are first-row elements lithium (n = 2) is a second-row element. 

A complete periodic table of the elements is presented on the inside back cover. 

A complete perioctic table of the elements is presentee) on the insieJe back cover. 

Before discussing structure and bonding in molecules, let s first review some fundamentals of atomic structure. Each element is characterized by a unique atomic number Z, which is equal to the number of protons in its nucleus. A neutral atom has equal numbers of protons, which are positively charged, and electrons, which are negatively charged. 

Other methods are also used to contrast the regions of an orbital where the signs of the wave function are different. Some mark one lobe of a p orbital + and the other Others shade one lobe and leave the other blank. When this level of detail isn t necessary, no differentiation is made between the two lobes. 

Be careful, though. The electron cloud of a hydrogen atom, although drawn as a collection of many dots, represents only one electron. 

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When multi-electron atoms are combined to form a chemical bond they do not utilize all of their electrons. In general, one can separate the electrons of a given atom into inner-shell core electrons and the valence electrons which are available for chemical bonding. For example, the carbon atom has six electrons, two occupy the inner Is orbital, while the remaining four occupy the 2s and three 2p orbitals. These four can participate in the formation of chemical bonds. It is common practice in semi-empirical quantum mechanics to consider only the outer valence electrons and orbitals in the calculations and to replace the inner electrons + nuclear core with a screened nuclear charge. Thus, for carbon, we would only consider the 2s and 2p orbitals and the four electrons that occupy them and the +6 nuclear charge would be replaced with a +4 screened nuclear charge. 

Note In this table all metal ions are in high-spin states and liganding atoms are small O, N donors. S-donors favour lower co-ordination numbers. Ligand-field theory, that is polarisation of and binding by the core electrons and orbitals of the metal ion compounds, can explain the above observations see inorganic chemistry textbooks in Further Reading . 

Atoms are constructed from small particles known as protons and neutrons, housed in the atomic nucleus, and orbital electrons. 

In Section 1.5, we emphasised the pertinence of the question stable or unstable with respect to what . So far in this chapter, we have sought to rationalise the right to exist of known, stable covalent substances by devising plausible descriptions of their bonding, accounting properly for the available valence electrons and orbitals of the constituent atoms. We now turn to unstable species, with a view to understanding the factors which deny them a right to exist. 

In this book we are particularly interested in simple descriptions of structures that are easily visualized and providing information of the chemical environment of the ions and atoms involved. For metals, there is an obvious pattern of structures in the periodic table. The number of valence electrons and orbitals are important. These factors determine electron densities and compressibilities, and are essential for theoretical band calculations, etc. The first part of this book covers classical descriptions and notation for crystals, close packing, the PTOT system, and the structures of the elements. The latter and larger part of the book treats the structures of many crystals organized by the patterns of occupancies of close-packed layers in the PTOT system. 

The critical choice was made of a CASSCF(4,4)/6-31G calculation the active space is thus the degenerate filled 2s + 2s and 2s 2s pair of MOs, and the degenerate empty 2px + 2px and 2px — 2px pair of MOs. CASSCF(4,4) was chosen because it corresponds to the CASSCF(2,2) calculation on one beryllium atom in the sense that we are doubling up the number of electrons and orbitals in our noninteracting system. This calculation gave an energy of 29.1709451 hartree. We can compare this with twice the energy of one beryllium atom, 2 x — 14.5854725 hartree = —29.1709450 hartree. 

Typical theoretical concepts occurring in chemical and quantum chemical theories are the various corpuscles - that is atoms , molecules , ions , and electrons - and orbital", spin , chemical bond , and electric charge . The expressions atom, molecule, electron, and ion refer to particles that are thought to... 

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