Section # __________________
Date of Experiment __________
Covalent Bonding and Molecular Geometry
When atoms combine to form molecules (this also includes complex ions) by forming covalent bonds, the relative positions of the combined atoms vary greatly between the different molecules that are formed. The molecule is typically visualized as being compose of a central atom (A) with two or more other atoms (X), called ligands, bonded directly to the central atom but not bonded to each other. In the typical molecules we will consider, the ligands will be the same type of atom but the shapes of molecules in which the are not all the same may also be determined by the process detailed later. The relative arrangement or molecular shape is established by the number of electron pair bonds or the electron domains around the central atom (or atoms) in the molecule. The shape of the molecule is obtained by drawing lines to connect the nuclei of each pair of atoms that are bonded together within the molecule. The bond angles formed by these lines create a definite geometry or shape around the central atom. The name of this shape is determined by drawing planes through groups of three ligands at a time to create a standard geometric shape which is the name given to the molecular geometry. (Sometimes the plane(s) may be drawn through two ligands and the central atom, especially in simple molecules.
The arrangement of the ligands around the central atom is determined by the number of electron domains present on the central atom. In establishing the number of electron domains, it is necessary to draw the electron dot diagram or Lewis diagram to depict the number and types of covalent bonds plus the number of non-bonded electron pairs that are assigned to the central (A) atom. These electrons are then classified into one of two types of electron domains, bonding or non-bonding. If two atoms share more than one pair of electrons i.e. have a double or triple bond, all of the electrons in the bonds constitute a single electron domain. Single bonds and non-bonded pairs each represent a single electron domain.
For example water has two bonded pairs and two non-bonded pairs of electrons assigned to the oxygen in the correct electron dot diagram. This gives four electron domains around the oxygen which determine the electron domain geometry. In carbon dioxide the carbon also has four pairs of electrons but only two electron domains because each oxygen is double bonded to the carbon. The different electron domain geometries cause these two triatomic molecules to have different shapes. Sometimes, in larger molecules it may be necessary to consider a group of atoms such as a methyl group, (CH3− ) or amine group (NH2− ) as a single group to obtain the bulk shape of the compound and then add the known geometry of the grouped atoms to complete the structure. The typical molecular structures in the figures below show the different geometries that are used to describe the shapes of most molecules. These figures are representing the location of the central atom (A) and the ligands (x) only and does not try to locate any non-bonding electrons even though these electrons are very important in establishing this geometry. The electron domain structures that give rise to the various shapes are denoted by the AXnEm notation where E now represents a pair of non-bonding electrons and the subscripts n and m indicate the number of ligands and non-bonding electron pairs respectively.
The actual molecular structures that are displayed are one of the above regular geometric figures if all of the electron domains involve bonded electrons. Some variant or distortion of these geometries is displayed when one or more pairs of electrons of the central atom are non-bonding. The relationship between the shape of the molecule and the distribution of the electrons as predicted by the Lewis diagram is an important concept in understanding the chemical and physical properties of matter.
This exercise is designed to present the techniques and the concepts necessary to predict the geometry of most typical small molecules.
The process of determining the shape of molecules is known as the VSEPR –Valence Shell Electron Pair Repulsion Theory or the Electron Domain Theory. The concept can be created by using balloons and tying inflated balloons together by their stems. The shape of the distribution of the balloons is determined by the number of balloons tied together and is caused by the repulsion interaction between the balloons. When there are only two balloons tied together the shape with the minimum repulsion is a linear arrangement. Three balloons will take up a trigonal planer shape, four a tetrahedron, five form a trigonal bipyramid and six form an octahedron. The same concept applies to the repulsion of the electron domains around an atom. If all of the electron domains are used to form covalent bonds then the molecular shape is the same as the electron domain geometry. If one or more of the electron domains around the central atom house non-bonding electron pairs, then the molecular geometry is different from the electron domain geometry.
If the central atom has four electron domains and they are all used in bonding as in CH4, the molecular geometry is tetrahedral. In H3N:, there are three bonding domains and one non-bonding domain. This causes three of the triangular faces of the tetrahedral to be changed from equilateral triangles to isosceles triangles and the nitrogen is at the apex of the pyramid instead of in the center of the pyramid as is the case for CH4 . Thus CH4 is a tetrahedral molecule and the H3N: molecule is trigonal pyramidal even though both molecules have four electron pairs around the central atom of the molecule.
Although it is possible to predict the correct molecular geometry by the VSEPR model it does not explain why a bond exists between the atoms. One approach, the valence bond theory, considers that a bond is formed when the atomic orbitals, which contain the electrons of the bond, overlap or occupy some common space. This overlap can occur directly on the line separating the nuclei of the bonded atoms and form a sigma bond (σ - bond ). The overlap may also be above and below a plane containing this line by the sideways overlap of p orbitals. This type of bond is called a pi bond (π - bond). However, except in the case of a few very simple molecules the atomic orbitals are not directed in the proper spatial pattern to form the bonds in the direction where they actually occur.
Valence bond theory solves this problem by assuming the atomic orbitals are just the most simple set of electron orbitals that are available. Bonding orbitals can be created at the appropriate angles around the central atom by forming hybrid orbitals. These new orbitals are created by combining or hybridizing the s, p, and d atomic orbitals to form a new set of orbitals needed to produce bonds in the directions about the central atom where the ligands are located. The number of atomic orbitals needed is determined by the number of electron domains present in the electron dot diagram for the molecule. If there are two domains then it is necessary to use two atomic orbitals (an s and a p) and the two hybrid orbitals formed are called sp hybrids because they are constructed from a single s and a single p orbital. The new sp hybrid orbitals formed from the two atomic orbitals are directed 180°
from each other. This is just the angle required to minimize the electron-electron repulsion from two pairs of electrons in the bonding orbitals. Each of the two hybrid orbitals can form a bond by overlapping with an atomic orbital on the ligand atom. An example of this is Be F2.
In order to determine the electronic environment of each nucleus, you must first draw the Lewis diagram for the BeF2 molecule.
:F:Be:F:
Since there are two pairs of electrons around the Be atom and no non-bonded pairs, the molecule falls into the classification AX2 and should be linear. The hybridization required for a linear shape is sp . The F atoms bond by overlapping their p orbital with one electron to the Be sp hybrids which also contain one electron each giving rise to the correct molecular geometry.
The Lewis diagram of BF3 requires three bonds to be formed and the molecule is of the type AX3
since all of the valence electrons of the boron atom are used to form bonds. These three electron domains obtain a maximum separation when the system is planer and the bond angles are 120° . Thus BF3 is a trigonal planer molecule as shown below.
This picture only shows the sp2 hybrid orbitals of the boron atom and does not indicate what atomic orbital on the ligand fluorine that is overlapping with the boron hybrid orbitals to form the bond. This is the designation that will be used in following molecular models.
In both of the above examples all of the electrons around the central atom are bonding electrons and the molecular shape is the same as the electron domain shape. This is not the situation when not all of the electrons around the central atom are bonding electrons as is the case for SO2 . In this case there are three electron domains but only two are bonding and one is non-bonding. The three domains still require the sp2 hybridization so the electron domain geometry is trigonal planer but the molecular shape is simply bent because three points ( atomic nuclei) must be either on a line or in the same plane. The minimum energy configuration for three domains is trigonal planer but only two are bonding and the molecule must be bent.
This is the process you will use to establish the geometry of the many possible molecular geometries produced when different numbers of electron domains form molecules with the various possible combinations of bonding and non-bonding domains. To determine the molecular shape by the VSEPR theory a four step process will be followed.
1. Draw the Lewis diagram and determine the number of electron domains in the valence shell of the central atom.
2. Determine the electron arrangement that will give the minimum repulsion energy from this number of electron pairs.
3. If necessary, decide which pairs will be bonding pairs and which will be non- bonding in order to minimize the electron repulsion energy.
4. If necessary, determine if there is any distortion of the bond angles caused by the unsymmetrical electron repulsions.
EXPERIMENTAL PROCEDURE
Part 1. Determine the spatial arrangements that will minimize the electron repulsion energy in molecules with 2, 3, 4, 5, and 6 electron pairs in the valence shell and complete the following table.
The arrangements may be visualized by tying the appropriate number of balloons together and
viewing the geometrical shape that is produced naturally by the repulsion between these balloons.
You should make a sketch of the arrangement, name this shape and state the orbital hybridization needed to accommodate this number of electron domains.
2 electron pairs
3 electron pairs
4 electron pairs
5 electron pairs
6 electron pairs
Sketch of Arrangement of
Electron pairs
Name of the Arrangement
Orbital Hybridization Required
Part 2. The second part of this exercise is to determine the molecular geometry for all of the molecules it is possible to construct with the each of the electron domain configurations given above. The following combinations of bonded ligands and non-bonding electron pairs include all the possible combinations:
AX2E0; AX3E0, AX2E1; AX4E0, AX3E1, AX2E2; AX5E0 , AX4E1, AX3E2, , AX2E3;
AX6E0, AX5E1, AX4E2, AX3E3, AX2E4
The information should be organized in the following table which has columns labeled "AXE Notation",
"Arrangement of electron pairs", "Sketch of molecule", "Molecular geometry" and "Orbital hybridization"
AXE Notation
Arrangement of Electron
Sketch of Molecule
Molecular Geometry
Orbital
Hybridization
pairs
AX2E0
; AX3E0
AX2E1
AX4E0
AX3E1
AX2E2
AX5E0
AX4E1
AX3E2
AX2E3
AX6E0
AX5E1
AX4E2
AX3E3
AX2E4
Part 3. With the aid of the table constructed in Part 2 determine the molecular geometry for the molecules given below. Complete a table for each molecule similar to the one above having the following column headings: "molecule", " Lewis Diagram of Molecule", "Number of electron pairs" ,
"Arrangement of electron pairs", "Number of bonded pairs", "Number of non-bonded pairs", "AXE Notation"
"Hybridization", "Shape of Molecule", "Sketch of Molecule".
Molecule Lewis Diagram
Numbe r of Electro n pairs
Arrang e-ment of Electro n pairs
Number of Bonded pairs
Number of non- bonding pairs
AXE
notation Hybrid-
ization Shape of
Molecule Sketch of
Molecule
SnCl2
OF2
HgI2
I3-1
BrO2-1
AlCl3
AsF3
BrF3
SO3-2
BrO3-1
BeF4-2
TeCl4
XeF4
SeO4-2
IO4-1
GeCl4
BrF5
AsCl5
SiF6-2
SF6
Results and Conclusions:
