THE COVALENT BOND – Introduction
Because the covalent bond is of central importance to organic chemistry, we begin with a review of bonding theory. Later, in Chapter 10, we shall return to develop certain aspects of the theory further in preparation for the discussion of pericyclic reactions.
MODELS OF CHEMICAL BONDING
Understanding and progress in natural science rest largely on models. A little reflection will make it clear that much of chemical thinking is in terms of models, and that the models useful in chemistry are of many kinds. Although we cannot see atoms, we have many excellent reasons for believing in them, and when we think about them we think in terms of models. For some purposes a very simple
model suffices. Understanding stoichiometry, for example, requires only the idea of atoms as small lumps of matter that combine with each other in definite proportions and that have definite weights. The mechanism by which the atoms are held together in compounds is not of central importance for this purpose. When thinking about stereochemistry, we are likely to use an actual physical model consisting of small balls of wood or plastic held together by springs or sticks. Now the relative weights of atoms are immaterial, and we do not bother to reproduce them in the model; instead we try to have the holes drilled carefully so that the model will show the geometrical properties of the molecules. Still other models are entirely mathematical. We think of chemical rate processes in terms of sets of differential equations, and the details of chemical bonding require still more abstract mathematical manipulations. The point to understand is that there may be many ways of building a model for a given phenomenon, none of which is complete but each of which serves its special purpose in helping us understand some
aspect of the physical reality.
The Electron Pair Bond-Lewis Structures
The familiar Lewis structure is the simplest bonding model in common use in organic chemistry. It is based on the idea that, at the simplest level, the ionic bonding force arises from the electrostatic attraction between ions of opposite charge, and the covalent bonding force arises from sharing of electron pairs between atoms. The starting point for the Lewis structure is a notation for an atom and its valence electrons. The element symbol represents the core, that is, the nucleus and all the inner-shell electrons. The core carries a number of positive charge equal to the number of valence electrons. This positive charge is called corecharge. Valence electrons are shown explicitly. For elements in the third and later rows ofthe periodic table, the d electrons in atoms of Main Groups 111, IV, V, VI, and VII are counted as part of the core. Thus :
Ions are obtained by adding or removing electrons. The charge on an ion is given by
charge = core charge – number of electrons shown exvlicidy
An ionic compound is indicated by writing the Lewis structures for the two ions. A covalent bond model is constructed by allowing atoms to share pairs of electrons. Ordinarily, a shared pair is designated by a line:
H-H
All valence electrons of all atoms in the structure must be shown explicitly. Those electrons not in shared covalent bonds are indicated as dots, for example:
If an ion contains two or more atoms covalently bonded to each other, the total charge on the ion must equal the total core charge less the total number of electrons, shared and unshared:
In order to write-correct Lewis structures, two more concepts are needed. First, consider the total number of electrons in the immediate neighborhood of each atom. This number is called the valence-shell occupancy of the atom, and to find it, all unshared electrons around the atom and all electrons in bonds leading to the atom must be counted. The valence-shell occupancy must not exceed 2 for hydrogen and must not exceed 8 for atoms of the first row of the periodic table. For elements of the second and later rows, the valence-shell occupancy may exceed 8. The structures
are acceptable.
The second idea is that of formal charge. For purposes of determining formal charge, partition all the electrons into groups as follows: Assign to each atom all of its unshared pair elec_tronsa nd half of all electrons in bonds leading it. Call the number of electrons assigned to the atom by this process its electron ownership. The formal charge of each atom is given by
formal charge = core charge – electron ownership
To illustrate formal charge, consider the hydroxide ion, OH-. The electron ownership of H is 1, its core charge is + 1, and its formal charge is therefore zero. The electron ownership of oxygen is 7, and the core charge is +6; therefore the formal charge is – 1. All nonzero formal charges must be shown explicitly in the
structure. The reader should verify the formal charges shown in the following examples
The algebraic sum of all formal charges in a structure is equal to the total charge. Formal charge is primarily useful as a bookkeeping device for electrons, but it also gives a rough guide to the charge distribution within a molecule. In writing Lewis structures, the following procedure is to be followed:
- Count the total number of valence electrons contributed by the electrically neutral atoms. If the species being considered is an ion, add one electron to the total for each negative charge; subtract one for each positive charge.
- Write the core symbols for the atoms and fill in the number of electrons determined in Step 1. The electrons should be added so as to make the valence shell occupancy of hydrogen 2 and the valence-shell occupancy of other atoms not less than 8 wherever possible.
- Valence-shell occupancy must not exceed 2 for hydrogen and 8 for a first-row atom; for a second-row atom it may be 10 or 12.
- Maximize the number of bonds, and minimize the number of unpaired erectrons, always taking care not to violate Rule 3.
- Find the formal charge on each atom.
We shall illustrate the procedure with two examples.
on February 24, 2009 at 11:28 am
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