9 - 1 Molecular Geometry Molecules have different shapes (geometries) depending on the type of atoms making it up and the number of electron pairs present. Molecular shapes are described in terms of bond angles and bond lengths. The length of a chemical bond is measured between two chemically bonded atoms from the nucleus of one to the nucleus of the other.

9 - 2 The bond angle is the angle between two bonds that include a common atom. Molecular shapes determine the properties of molecules such as polarity and solubility. Lewis structures are explained in the Covalent Bonds and Ionic Bonds PowerPoints.

9 - 3 Lewis Structure of Water The Lewis structure only shows the bonding of each atom, not the shape of the water molecule. H O H The actual shape is with a bond angle of O 104.5° H

9 - 4 The previous slide shows that oxygen is the central atom having four electron pairs. The O shares two pairs of electrons with the H’s and two of the pairs are not shared. The unshared pair repel each other and they also strongly repel the shared pair. Instead of water being a linear molecule, it is said to be bent or angular.

9 - 5 VSEPR Theory The valence shell electron pair repulsion theory says that the shape of a molecule results from the repulsive interaction of electron pair in the valence shell of an atom. The most important atom for determining the geometry is the central atom. The geometry depends on the atoms having minimal interaction between the valence shell electron pairs.

9 - 6 VSEPR Theory The minimal interaction between the pairs of valence electrons maximizes the distance between the electron pairs and between the atoms making up the molecule. A multiple bond (double or triple) holds the multi-bonded atom in the same position as a single bond. A multiple bond is treated as a single bond for determining the molecular geometry.

9 - 7 VSEPR Theory The following VSEPR structures were made by Dr. Mark R. Leach who granted permission for their use. Check out all his graphics at

9 - 8 VSEPR for Molecules and Ions AX 2 Valence electron pairs: 2 Bonding electron pairs: 2 Nonbonding electron pairs: 0 Examples: BeCl 2, CO 2

9 - 9 VSEPR for Molecules and Ions AX 3 Valence electron pairs: 3 Bonding electron pairs: 3 Nonbonding electron pairs: 0 Examples: BF 3, In(CH 3 ) 3

VSEPR for Molecules and Ions AX 2 E Valence electron pairs: 3 Bonding electron pairs: 2 Nonbonding electron pairs: 1 Examples: SO 2, GeF 2

VSEPR for Molecules and Ions AX 4 Valence electron pairs: 4 Bonding electron pairs: 4 Nonbonding electron pairs: 0 Examples: CH 4, CCl 4

VSEPR for Molecules and Ions AX 3 E Valence electron pairs: 4 Bonding electron pairs: 3 Nonbonding electron pairs: 1 Examples: NH 3, H 3 O +

VSEPR for Molecules and Ions AX 2 E 2 Valence electron pairs: 4 Bonding electron pairs: 2 Nonbonding electron pairs: 2 Examples: OF 2, H 2 O

VSEPR for Molecules and Ions AX 5 Valence electron pairs: 5 Bonding electron pairs: 5 Nonbonding electron pairs: 0 Examples: PCl 5

VSEPR for Molecules and Ions AX 4 E Valence electron pairs: 5 Bonding electron pairs: 4 Nonbonding electron pairs: 1 Examples: SF 4

VSEPR for Molecules and Ions AX 3 E 2 Valence electron pairs: 5 Bonding electron pairs: 3 Nonbonding electron pairs: 2 Examples: ClF 3

VSEPR for Molecules and Ions AX 2 E 3 Valence electron pairs: 5 Bonding electron pairs: 2 Nonbonding electron pairs: 3 Examples: ICl 2 -, XeF 2

VSEPR for Molecules and Ions AX 6 Valence electron pairs: 6 Bonding electron pairs: 6 Nonbonding electron pairs: 0 Examples: SF 6

VSEPR for Molecules and Ions AX 5 E Valence electron pairs: 6 Bonding electron pairs: 5 Nonbonding electron pairs: 1 Examples: BrF 5

VSEPR for Molecules and Ions AX 4 E 2 Valence electron pairs: 6 Bonding electron pairs: 4 Nonbonding electron pairs: 2 Examples: XeF 4, ICl 4 -

Exceptions to the Octet Rule Three major exceptions to the octet rule:  Molecules or ions with more than eight electrons around the central atom.  Species with fewer than eight electrons around the central atom.  Species with an odd number of valence electrons.

Expanded Octets Starting with period three, atoms have the capability to accommodate d electrons (3d). AX 4 E molecules such as SF 4 are able to accommodate 4 bonding pairs of electrons and one nonbonding pair of electrons. This results in S being surrounded by 5 electron pairs.

The favored bonding scenario includes large central atoms (starting in the third period) and small terminal atoms such as fluorine, chlorine, and oxygen. As shown below, S also has the ability to accommodate six pairs of valence electrons as found in SF 6.

Less Than an Octet Molecules having either boron or beryllium as their central atom result in the central atom having only 2 or 3 valence pairs of electrons. These molecules are very reactive with a molecule having an unshared pair of electrons. BeCl 2 BF 3

Odd Number of Valence Electrons Most molecules have an even number of valence electrons. In rare cases, molecules such as NO and NO 2, there is one unpaired electron which is very reactive.

Odd Number of Valence Electrons nitrogen(II) oxide nitrogen(IV) oxide N OONO

Polar Molecules and Polar Bonds The molecular geometry of a molecule or ion determines if polar bonds in a species result in the species itself being polar. If all the bonds in a molecule are nonpolar, then the molecule itself is nonpolar regardless of geometry. A polar molecule has an asymmetrical distribution of charge.

The charge results from the atoms in the molecule having different electronegativities and their spatial arrangement. The polarity of the O-H bond contributes to the resultant polarity ( ) of the water molecule. H H O

Each O-H bond is polar and because of its bent or angular shape, the water molecule itself is polar. If the bond angle was 180° as it is in HCl, water would be a nonpolar molecule. The HCl molecule is polar because of the difference in electronegativities between H and Cl. HCl δ+δ+ δ-δ-

Molecular Geometry in Summary  Molecular geometry is determined by the position of the atoms, not by the position of electron pairs.  Lone pairs of electrons repel other lone pairs more strongly.  The electron cloud surrounding a lone pair of electrons is much bigger than the cloud surrounding a bonding pair of electrons.

 Bonding pairs of electrons have the smallest force of repulsion.  The order of electron pair repulsion is: lp-lp > lp-bp > bp-bp  Molecules or ions with lone pairs of electrons will have smaller bond angles than predicted.