The Octet Rule In Chemical Bonding
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The Octet Rule In Chemical Bonding

The octet rule, or rule of eight is the tendency of the Type II elements to enter into chemical combination through losing, gaining, or sharing electrons, in such a way to attain the electronic configuration of a noble gas.

The octet rule, or rule of eight is the tendency of the Type II elements to enter into chemical combination through losing, gaining, or sharing electrons, in such a way to attain the electronic configuration of a noble gas.

In the Type I elements called the noble or inert gases, the electron shells elements are complete since they have attained an electronic stability which cannot be increased either by electron transfer or electron sharing. As a result, they are highly indifferent to chemical change.

Electronic structure of the noble gases

Atoms containing total of 2, 10, 18, 36, 54, or 86 extranuclear electrons are remarkably stable. The elements with atoms containing electrons of these totals are known as the noble or rare gases, which are so inert that they form no chemical combinations. With the exception of helium, the atoms of these elements contain 8 electrons in their outermost shells. Apparently, electron totals of 2, 10, 18, 36, 54, and 86 attain electronic configurations that are more stable than any intermediate numbers.

Electronic Structure of Type II Elements

The chemical properties of these elements can be explained in terms of the electrons of their incomplete, outermost shell—often called the valence electrons. The number of electrons in this shell may lie between one and seven. Chemical behavior may be described as a tendency to gain, or to lose, or to share those electrons in chemical reactions, in such fashion that the outer shell may contain eight, as in one of the noble gases. (A few of the lightest elements tend toward an outer shell of two, as in helium.)

Type II atoms with one or two electrons in the outer shell appear to hold these electrons comparatively loosely and to lose them readily to form positive ions, e.g.,

              Na (2,8,1) → e + Na+                (2,8) isoelectronic with Ne (2,8)

              Ca (2,8,8,2) → 2e + Ca+2        (2,8,8) isoelectronic with A (2,8,8)

         Diagram illustrating electron transfer to form an ionic compound

 At the other extreme, atoms with seven electrons in the outer shell show a strong tendency to gain one electron, and thus form negative ions which are isoelectronic with the next higher noble gas, e.g.,

               Cl (2,8,7) + e → Cl-                 (2,8,8) isoelectronic with A (2,8,8)

The formation of an ionic compound such as sodium chloride ia a typical example of an electron transfer taking place in accordance with the octet rule. Each sodium atom loses its outer-shell electron, thus forming the sodium ion which is isoelectronic with neon; each chlorine atom gains one electron, forming a chloride ion with the electron configuration of argon.

For atoms of those elements which have more than two and fewer than seven valence electrons, they seldom react by electron transfer. They are more inclined to attain the noble gas structures by sharing electrons in the formation of covalent molecules. The sharing process presumably requires for the atoms of these particular elements a lower expenditure of energy than does electron transfer. A favorite illustration of the process of electron sharing is the electron distribution found in the carbon tetrachloride molecule. In this molecule, four chlorine atoms have attained the structure of argon by sharing four electron pairs with the carbon atom. At the same time, the sharing of the four electron pairs permits the carbon atom to become isoelectronic with the neon atom.

The rule of eight is not universal. It applies only to the Type II elements and, with well-known exceptions. For example, the compounds   have properties characteristic of covalent compounds, yet neither the beryllium nor the boron atom has attained eight electrons in the valence shell by forming these compounds. In short, the octet rule—which has its uses in an elementary discussion of valency—oversimplifies the factors involved in the formation of many chemical bonds. A better basis for predictions in regard to bond formation is one that utilizes the concept of uncompleted orbitals in one atom which may be filled by electrons from another atom

Adapted from: Qualitative Analysis by E. S. Gilreath, Philippines Copyright 1964

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