Chemical Bonding

Diatomic Elements

Elements found in their elemental form as diatomic molecules. In these molecules, two atoms are joined by one or more covalent bonds, forming a molecule with the general formula X2. They are:

H2 N2 O2 F2 Cl2 Br2 I2

HH− NN≡ OO=

F −F Cl−Cl BrBr− II−

Polyatomic Elements

These elements are found in their elemental form as molecules of more than two atoms. They are:

S8 Se8 P4

Molecular Shape Molecular geometry or molecular structure is the three- dimensional arrangement of the atoms that constitute a molecule. It determines several properties of a substance including its polarity, phase of matter, chemical reactivity, and various other things. The overall shape of a molecule is often more important than what atoms are involved in making it. Sometimes the same numbers and kinds of atoms make up molecules with different shapes and the shapes cause these molecules to have very different properties.

Three-dimensional shapes are best viewed with the aid of models. In order to represent such configurations on a two-dimensional surface (paper, blackboard or screen), we often use perspective drawings in which the direction of a bond is specified by the line connecting the bonded atoms. In most cases the focus of configuration is a carbon atom so the lines specifying bond directions will originate there.

As defined in the diagram below:

A simple straight line represents a bond lying approximately in the surface plane. The two bonds to substituents A in the structure on the left are of this kind.

A wedge shaped bond is directed in front of this plane (thick end toward the viewer), as shown by the bond to substituent B.

A hatched bond is directed in back of the plane (away from the viewer), as shown by the bond to substituent D. Some texts and other sources may use a dashed bond in the same manner as we have defined the hatched bond, but this can be confusing because the dashed bond is often used to represent a partial bond (i.e. a covalent bond that is partially formed or partially broken).

The following examples make use of this notation, and also illustrate the importance of including non-bonding valence shell electron pairs (colored blue) when viewing such configurations.

Methane Ammonia Water CH4 NH3 H2O

The above molecules must be memorized by shape, formula, and name.

Valence Shell Electron Pair Repulsion Theory

The acronym "VSEPR" is sometimes pronounced "vesper" for ease of pronunciation.

VSEPR theory proposes that the geometric shape of a molecule, is determined solely by the repulsions between electron pairs present in the valence shell of the central atom. The number of electron pairs around the central atom can be determined by writing the Lewis structure for the molecule. The geometry of the molecule depends on the number of bonding groups (pairs of electrons) and the number of nonbonding electrons on the central atom.

Electron pairs repel each other and will move as far apart as possible.

The repulsion caused by a lone pair is greater than the repulsion caused by a bonding pair.

The Highlighted shapes must be known.

# of lone # of bonding

pair groups (pair electrons electrons) Bond Molecular Geometry on 'central' on 'central' Angle atom atom 0 2 linear 180 0 3 trigonal planar 120 1 2 bent less than 120 0 4 tetrahedral 109.5 1 3 trigonal pyramidal less than 109.5 2 2 bent less than 109.5 0 5 trigonal bipyramidal 90, 120 and 180 1 4 seesaw 90, 120 and 180 2 3 T-shaped 90 and 180 3 2 linear 180 0 6 octrahedral 90 and 180 1 5 square pyramidal 90 and 180 2 4 square planar 90 and 180

VSEPR Theory at Oklahoma State Web Site

A represents the central atom. X represents the number of bonds between the central atoms and outside atoms. Multiple covalent bonds (double, triple, etc) count as one X. E represents the number of lone electron pairs surrounding the central atom.

VSEPR theory makes the following predictions:

Basic Geometry 1 lone pair 2 lone pairs 3 lone pairs 0 lone pair

linear

trigonal planar bent

bent tetrahedral trigonal pyramid

trigonal bipyramid seesaw T-shaped linear

octahedral square pyramid square planar

pentagonal bipyramid pentagonal pyramid

All of the links below connect to Wikipedia. This table originated on Wikipedia.

Electron arrangement including lone pairs, shown in pale yellow

Electron arrangement Geometry Molecule Observed geometry Examples Type Shape Lone pairs (excluding lone pairs) shown in yellow

AX1En Diatomic HF, O2

BeCl2, HgCl2, AX2E0 Linear CO2

− AX2E1 Bent NO2 , SO2, O3

AX2E2 Bent H2O, OF2

− AX2E3 Linear XeF2, I3

2− BF3, CO3 , AX3E0 Trigonal planar − NO3 , SO3

Trigonal AX3E1 NH3, PCl3 pyramidal

AX3E2 T-shaped ClF3, BrF3

3− CH4, PO4 , AX4E0 Tetrahedral 2− − SO4 , ClO4

AX4E1 Seesaw SF4

AX4E2 Square Planar XeF4

Trigonal AX5E0 PCl5 Bipyramidal

Square AX5E1 ClF5, BrF5 Pyramidal

AX6E0 Octahedral SF6

− 2− Pentagonal XeOF5 , IOF5 AX6E1 [6] pyramidal

Pentagonal AX7E0 IF7 bipyramidal

Covalent Bonding Theories

During the 1920s and 1930s while the electron orbitals were starting to be understood, two theories about how these orbitals are involved in bonding were developed. First came the Valence Bond (VB) Theory and then the Molecular Orbital (MO) Theory. Aspects of both are still used today and parts of both theories will be presented in this outline. One key difference: molecular orbital theory involves orbitals that cover the whole molecule and delocalize the shared electrons around the molecule. Valence bond theory considers the shared electrons to be concentrated in the bond region in a small overlapping area. When atoms interact to form a chemical bond, the atomic orbitals are said to mix in a process called orbital hybridization.

© Mr. D. Scott; CHS Hybridization is the merging or mixing of valence orbitals – usually s and p orbitals – during bonding producing new orbital shapes. These new shapes orient the valence electrons around the atom to achieve the lowest energy conditions (best balance of attractive and repulsive forces). These hybrid shapes give rise to the various molecular geometries (molecule shapes). Once the hybrid is formed, it then overlaps or connects with the valence orbitals of the adjoining atom thereby forming the covalent bond. EXAMPLE: sp3 hybridization

Sigma bonds occur when the orbitals of two shared electrons overlap head-to-head. Pi bonds occur when two orbitals overlap when they are parallel.

Sigma bonds: Overlapping occurs symmetrically along the internuclear axis. (End to End)

Pi bonds: Overlapping occurs parallel to the internuclear axis. (Side to side)

Single bond = 1 sigma bond Double bond = 1 sigma bond & 1 pi bond Triple bond = 1 sigma bond & 2 pi bonds

Sigma & Pi Bond Animation (requires flash player)

Sigma & Pi Bonding at the ChemEd Digital Library

Molecular Orbital = The space occupied by the shared electrons in a covalent bond. These orbital regions have unique shapes that are different from the atomic orbitals (s,p,d,f).