Shells of electrons, completed

It is not proposed to discuss here the stereochemistry of S, which is much more complex than that of 0 because of the availability of d orbitals, but merely to summarize some points of difference between 0 and S. From 0 to Te the atoms increase in size and we may associate the change in behaviour of the outer valency electrons with the increased screening of the nuclear charge by the intervening completed shells of electrons. This shows itself in a number of ways ... 

This simple case illustrates the idea of completed shells of electrons. The first two sets of quantum numbers remain the... 

It may be mentioned that the potential-energy function due to an s electron is spherically symmetrical, inasmuch as the probability distribution function w u , is independent of

potential-energy function due to a completed shell of electrons... 

In his 1916 paper The Atom and the Molecule, Lewis proposed that a chemical (covalent) bond between two atoms involves the sharing of electrons between the nuclei. Thus a single bond (for hydrogen, H-H) results when an electron from each atom forms an electron pair that is shared between the two nuclei (H H) a double bond involves two electrons from each atom (e.g., the carbon-carbon bond in (H )2C C( H)2) and a triple bond involves three electrons from each atom (e.g., the carbon-carbon bond in H C C H). Such representations are referred to as Lewis dot structures. Lewis further postulated that an electron octet (and in a few cases an electron pair) forms a complete shell of electrons with spatial rigidity and chemical inertness—hence a stable arrangement. 

Every chemist knows that atomic nuclei are surrounded by shells of electrons, which, when completed, contain 2 (inner shell), 8, 8, 18,. . . electrons, this being the explanation of periodicity in chemical properties. For reasons that will appear, these may be called (in the same order) K, L, M, N,. . . electrons on the basis of conclusive x-ray evidence. 

The interaction of two alkali metal atoms is to be expected to be similar to that of two hydrogen atoms, for the completed shells of the ions will produce forces similar to the van der Waals forces of a rare gas. The two valence electrons, combined symmetrically, will then be shared between the two ions, the resonance phenomenon producing a molecule-forming attractive force. This is, in fact, observed in band spectra. The normal state of the Na2 molecule, for example, has an energy of dissociation of 1 v.e. (44). The first two excited states are similar, as is to be expected they have dissociation energies of 1.25 and 0.6 v.e. respectively. 

Studies of the electron distributions around outer atoms consistently show that hydrogen is always associated with two electrons (one pair). All other outer atoms always have eight electrons (four pairs). The reason for this regularity is that each atom in a molecule is most stable when its valence shell of electrons is complete. For hydrogen, this requires a single pair of electrons, enough to make full use of the hydrogen 1 S orbital. Any other atom needs four pairs of electrons, the maximum number that can be accommodated by an .S p valence shell. Details of these features can be traced to the properties of atoms (Chapter 8) and are discussed further in Chapter 10. 

It must be emphasized that the octet rule does not describe the electron configuration of all compounds. The very existence of any compounds of the noble gases is evidence that the octet rule does not apply in all cases. Other examples of compounds that do not obey the octet rule are BF,. PF5, and SF6. But the octet rule does summarize, systematize, and explain the bonding in so many compounds that it is well worth learning and understanding. Compounds in which atoms attain the configuration of helium (the duet) are considered to obey the octet rule, despite the fact that they achieve only the duet characteristic of the complete first shell of electrons. 

Structures in which all of the atoms have a complete valence shell of electrons are especially stable and make large contributions to the hybrid. 

The outer shell of the helium atom is full and complete the shell can only accept two electrons and, indeed, is occupied by two electrons. Similarly, argon has a complete octet of electrons in its outer shell. Further reaction would increase the number of electrons if argon were to undergo a covalent bond or become an anion, or would decrease the number of electrons below the perfect eight if a cation were to form. There is no impetus for reaction because the monatomic argon is already at its position of lowest energy, and we recall that bonds form in order to decrease the energy. 

Each atom in both structures has a complete valence shell of electrons. There are no formal charges in the first structure, but in the second structure, the oxygen is formally positive and the carbon is formally negative. 

Gases form the group 8 of the periodic table. All contain a complete outer shell of electrons. They are helium, neon, argon and krypton. 

A system of chemical bonding whereby atoms try to obtain a stable outer electron arrangement of a noble gas by achieving a complete octet of electrons in their outer shell. See also covalent bonding and ionic bonding. 

The Representation of Phosphorus Compounds by Electronic Theories of Valency.—The compounds in which phosphorus is trivalent are saturated in the sense that all the covalent bonds on the element are made up, with the completion of the outer octet of electrons, the phosphorus atom thus assuming the argon type with three completed shells of 2, 8, 8 electrons of which only the outermost are shown by the formulas—1... 

Hybridization can also help explain the existence and structure of many inorganic molecular ions. Consider, for example, the zinc compounds shown here. At the top is shown the electron configuration of atomic zinc, and just below it, of the divalent zinc ion. Notice that this ion has no electrons at all in its 4-shell. In zinc chloride, shown in the third row, there are two equivalent chlorine atoms bonded to the zinc. The bonding orbitals are of sp character that is, they are hybrids of the 4s and one 4p orbital of the zinc atom. Since these orbitals are empty in the isolated zinc ion, the bonding electrons themselves are all contributed by the chlorine atoms, or rather, the chlor ide ions, for it is these that are the bonded species here. Each chloride ion possesses a complete octet of electrons, and two of these electrons occupy each sp bond orbital in the zinc chloride complex ion. This is an example of a coordinate covalent bond, in which the bonded atom contributes both of the electrons that make up the shared pair. 

I Adding a new shell of electrons considerably increases the size of the atom breaking into a complete octet to ionize an electron requires a huge quantity of energy. 

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