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.  The 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 orientations in space than the original "s", "p", "d", or "f" atomic orbitals from which they were formed. 

The simplest and most convenient theory which explains many of the aspects of covalent bonding between atoms is the Valence Bond Theory (VBT).  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.

Based on the experimental observation that the Noble gases (He, Ne, Ar, Kr, Xe) are chemically stable (do not react with other substances), it has been generally accepted that the s2p6 electronic configuration in the valence shell of an atom is particularly 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 due 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 2nd and 3rd periods of the Periodic Table.  The majority of common substances are compounds of these atoms.  However, 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 4th 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 "non-bonding pairs" ("unshared pair" and "lone pair" mean the same thing as "non bonding 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 the proper Lewis Dot structures for the atoms, ions, or molecules listed in the table below.  If your Lewis Dot structure shows resonance, draw all of the individual structures and then the average structure.  Make sure you calculate the formal charges on all of the atoms in your structure to make sure you have the best Lewis Dot structure.  NOTE:  YOU MUST COMPLETE THE FIRST FOUR COLUMNS OF THIS TABLE BEFORE COMING TO LAB.  Represent each covalent bond with a dash '-' and lone pairs with ':'.  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 the shorter gray plastic pieces to represent a single bond or non bonding pair of electrons.  If multiple bonds appear in your Lewis dot structure, you will need to use the longer gray plastic pieces to hold those particular atoms together.  

Make sure you use a central atom with the correct number of holes!  If you have four regions of electron density, use an atom with four holes.  It makes no sense to use a central atom with five holes if you only have four regions of electron density!!!!
 

Species
Name
Central
Atom
Lewis dot 
Structure
Bonding System
(AxByUz)
Electronic
Geometry
Molecular
Geometry
Valence Bond
Hybridization
Polar or
Non polar
PO43-
               
SO42-
               
CO32-
               
NO3-
               
ClO4-
               
NH3
               
CO2
               
H2O
               
Ne
               
Cl-
               
N2
               
HCl
               
CH4
               
CCl4
               
HNO3
               
H2SO4
               
CH2CH2
               
CH3CH3
               
O3
             
 
SF4
 
 
 
 
 
 
 
 
PF5
 
 
 
 
 
 
 
 
I3-
               
PF6-
               
BrF5
               
 IF4-
               

Note:  If you are having problems constructing the Lewis Dot structures for your pre-lab, or with completing the table after lab, I encourage you do use the internet and look up the answers.  I have also attached a PDF file I created from one of the best reports I have ever received for this lab (thank you Kady W.):  Molecular Modeling Key.

(Updated 9/30/13 by C.R. Snelling)