LEWIS THEORY - CHEMISTRY TUTORIALS

In the period from 1916 to 1919, two Americans, G. N. Lewis and Irving Langmuir, and a German, Walther Kossel, advanced an important proposal about chemical bonding:

Something unique in the electron configurations of noble gas atoms accounts for their inertness, and atoms of other elements combine with one another to acquire electron configurations like those of noble gas atoms.

The theory that grew out of this model has been most closely associated with G. N. Lewis and is called the Lewis theory. Some fundamental ideas associated with Lewis s theory follow:

1. Electrons, especially those of the outermost (valence) electronic shell, play a fundamental role in chemical bonding.

2. In some cases, electrons are transferred from one atom to another. Positive and negative ions are formed and attract each other through electrostatic forces called ionic bonds.

3. In other cases, one or more pairs of electrons are shared between atoms. A bond formed by the sharing of electrons between atoms is called a covalent bond.

4. Electrons are transferred or shared in such a way that each atom acquires an especially stable electron configuration. Usually this is a noble gas configuration, one with eight outer-shell electrons, or an octet.

Lewis developed a special set of symbols for his theory. A Lewis symbol consists of a chemical symbol to represent the nucleus and core (inner-shell) electrons of an atom, together with dots placed around the symbol to represent the valence (outer-shell) electrons.

A Lewis structure is a combination of Lewis symbols that represents either the transfer or the sharing of electrons in a chemical bond.

In Lewis theory, we use square brackets to identify ions, as we did in equation (10.1). The charge on the ion is given as a superscript.

Lewis s work dealt mostly with covalent bonding, which we will emphasize throughout this chapter. However, Lewis s ideas also apply to ionic bonding, and we briefly describe this application next.

ATOMIC RADIUS - BASIC INFORMATION AND TUTORIALS

Unfortunately, atomic radius is hard to define. The probability of finding an electron decreases with increasing distance from the nucleus, but nowhere does the probability fall to zero, so there is no precise outer boundary to an atom.

We might describe an effective atomic radius as, say, the distance from the nucleus within which 95% of all the electron charge density is found, but in fact, all that we can measure is the distance between the nuclei of adjacent atoms (internuclear distance).

Even though it varies, depending on whether atoms are chemically bonded or merely in contact without forming a bond, we define atomic radius in terms of internuclear distance.

Because we are primarily interested in bonded atoms, we will emphasize an atomic radius based on the distance between the nuclei of two atoms joined by a chemical bond. The covalent radius is one-half the distance between the nuclei of two identical atoms joined by a single covalent bond.

The ionic radius is based on the distance between the nuclei of ions joined by an ionic bond. Because the ions are not identical in size, this distance must be properly apportioned between the cation and anion. One way to apportion the electron density between the ions is to define the radius of one ion and then infer the
radius of the other ion.

The convention we have chosen to use is to assign an ionic radius of 140 pm. An alternative apportioning scheme is to use as the reference ionic radius.

When using ionic radii data, one should carefully note which convention is used and not mix radii from the different conventions. Starting with a radius of 140 pm for the radius of Mg2+ can be obtained from the internuclear distance in MgO, the radius of from the internuclear distance in and the radius of from the internuclear distance in NaCl.

For metals, we define a metallic radius as one-half the distance between the nuclei of two atoms in contact in the crystalline solid metal. Similarly in a solid sample of a noble gas the distance between the centers of neighboring atoms is called the van der Waals radius.

There is much debate about the values of the atomic radii of noble gases because the experimental determination of the van der Waals radii is difficult; consequently, the atomic radii of noble gases are left out of the discussion of trends in atomic radii.