Introduction:

To develop a good understanding of the physical and chemical properties of compounds it is extremely helpful to know the shape of the molecules or ions being considered.  In fact, current theories of bonding have depended, in part, upon experimental determination of the structures of molecules.  Bond length (distance between centers of atoms covalently bonded to each other) is used to indicate the presence of single or multiple bonds between atoms or, in fact, the presence of a covalent bond at all.  Bond angle (angle between covalent bonds from the same atom to two other atoms) is used to indicate which atomic orbital are overlapping.  A comparison of bond angles around certain atoms had led to an extension of Valence Bond Theory Called "hybridization theory".  According to the hybridization theory, the s, p, d, and f atomic orbitals can be mixed together in different combinations to form "hybridized atomic orbitals which have different orientation in space than the original "s", "p", "d", or "f" orbitals from which they were formed.  The use of hybridization to explain bonding in molecules or ions is, therefore, based on a consideration of shape.

The simplest and most convenient theory which explains many of the aspects of covalent bonding between atoms is the Valence Bond Theory.  According to VBT, the valence shell electrons on an atom are used in covalent bonding.  An atomic orbital on one atom is pictured as overlapping an atomic orbital on another atom as the two atoms move sufficiently close to each other.  The atomic orbital on each atom contains a single unpaired electron so that when overlapping of atomic orbitals occurs, the electrons pair up with each other in a region between the two atoms.  This shared pair of electrons is called the "covalent bond".

The atoms in the first three periods of the Periodic Table contain from one to eight electrons in their valence shells.  Based on the electronic configuration of the atom (and sometimes using the hybridization idea) it is possible to predict the number of unpaired electrons in the valence shell of the atom and hence the number of covalent bonds which that atom may form with other atoms.  THIS ABILITY TO PREDICT THE NUMBER OF COVALENT BONDS AND THE FORMULAS OF MOLECULES AND IONS CONTAINING COVALENT BONDS IS VALUABLE TO THE CHEMIST.

Based on the experimental observation that the rare gases (He, Ne, Ar, Kr, Xe) are chemically stable (for the most part resist reaction with other substances), it has been generally accepted that the s²p⁶ electronic configuration in the valence shell of an atom is an especially desirable situation from the standpoint of atomic stability.  One explanation for chemical bonding between atoms and ions is that by transferring electrons to form ions which attract each other or by sharing electrons between atoms, this desire on the part of atoms to achieve maximum stability (lowest energy dues to an s²p⁶ configuration) is achieved.  This idea is referred to as the "octet rule" or "rule of eight".  The octet rule is useful in explaining the formation of ionic and covalent compounds involving atoms in the 2^nd and 3^rd periods of the Periodic Table.  The majority of common substances are compounds of these atoms.  It is only fair to note that this generalization does not hold for atoms within two positions of He in the Periodic Table (where the "rule of two" applies) or for many atoms in the 4^th period and beyond where the larger atoms can actually accommodate six or more pairs of electrons when forming covalent bonds to other atoms.

Using the "octet rule" and the ideas of two electrons to an orbital and two electrons per covalent bond, G. N. Lewis drew electron dot formulas using a dot to represent each valence electron in the outermost energy level of an atom.  The correct formula of a compound and the number of covalent bonds between atoms could be predicted on the basis of valence shell electrons alone.  Because each atom attempts to have either two (remember H) or eight electrons surrounding it, by using all the valence shell electrons from all of the atoms in the molecule or ion, it is possible to draw a simple picture of the atom, ion, or molecule with each atom having a stable allotment of two or eight valence electrons (which has be achieved through electron transfer or sharing).  Such a picture is called a "Lewis dot structure" or "electron dot formula".

It is customary in such a picture to show electrons on four sides of the atom to symbolize the four orbitals which make up the octet - one s and three p orbitals.  In the final electron dot formula, electron pairs between atoms are called "bonding pairs" and electron pairs at other position around the atom are co called "nonbonding pairs" ("unshaired pair" and "lone pair" mean the same thing as "nonbonding pair").  The valence electrons are thought to occupy either the "s" or three "p" orbitals or four other atomic orbitals resulting from mixing of the "s" orbital with either one of the "p" orbitals, two of the "p" orbitals, or all three of the "p" orbitals.

Purpose:

To gain insight into the structural aspects of compounds, the importance of geometry in determining chemical and physical properties, and the role of atomic orbitals in determining the geometry of ions and molecules.

Procedure:

Draw Lewis structures for the atoms, ions, or molecules listed in the table below (YOU MUST DO THIS BEFORE COMING TO LAB).  Represent each valence electron as a dot.  Name the atom, ion, or molecule (using your textbook, if necessary).  Construct a model of the atom, ion, or molecule using a model kit.  Use a stick to represent each bonding or nonbonding pair of electrons.  If multiple bonds appear in your Lewis structure, you will need to use springs rather than sticks to hold those particular atoms together.  Have your instructor initial each model before proceeding to the next entry.
 
 

(Updated 6/4/07 by C.R. Snelling)