 1. VSEPR
2. Molecular shapes and bond angles
3. Determining molecular shapes

Valence shell electron-pair repulsion (VSEPR) theory is a set of rules which explains the geometry that molecules adopt. The rules are governed by finding the lowest energy conformation, holding both the bonding electrons and lone-pair electrons as far apart as possible.

Repulsion in an order of increasing strength is:
1. bonding pair - bonding pair,
2. bonding pair - lone pair,
3. lone pair - lone pair. To determine the shape of molecule, start from counting the number of atoms and lone pairs of electrons attached to the central atom. Then locate the lone pairs of electrons to positions that are further away from the other electron pairs and atoms. Lone pairs of electrons are usually ignored while naming the shape and the following shapes usually result from a particular number of bonds and lone pairs.

Assuming that all atoms bonded to the central atom are the same:
2 bonds → linear
2 bonds + some lone pairs → angular/bent
3 bonds → trigonal planar (120°)
3 bonds + 1 lone pair → trigonal pyramidal
4 bonds → tetrahedral (109.5°)
4 bonds + 1 lone pair → see-saw
4 bonds + 2 lone pairs → square planar (90°)
5 bonds → trigonal bipyramidal (90° and 120°)
6 bonds → octahedral (90°) The geometry of molecules can be determined using the VSEPR rules that require knowing the number of bonds as well as the number of lone pairs of electrons. The number of bonds in small molecules is usually the same as the number of atoms attached to the central atom. To find the number of lone pairs of electrons, add all of the electrons that the atoms in a molecule have, adjust for a possible charge, subtract electrons that are used in bonds, subtract non-bonding electrons of non-central atoms and divide the result by 2. Approximate values of angles between bonds can be found by adding 2° for each lone pairs of electrons pushing them closer to each other. # 1

What geometry do CO₂ molecules adopt?

number of valence electrons = 4 (carbon) + 2 × 6 (oxygen) = 16
There are two double C=O bonds in the molecule.
Each oxygen has 4 non-bonding electrons.
number of lone pairs = [16 − 2 × 4 (bonds) − 2 × 4 (oxygen atoms)] ÷ 2 = 0
There are no lone pairs to distort a linear structure.

linear # 2

Determine the shape of a ClF₃ molecule.

number of valence electrons = 7 (chlorine) + 3 × 7 (fluorine) = 28
There are 3 single Cl—F bonds in the molecule.
Each fluorine has 6 non-bonding electrons.
number of lone pairs = [28 − 3 × 2 (bonds) − 3 × 6 (fluorine atoms)] ÷ 2 = 2
There are 2 lone pairs that take the axial positions in a trigonal bipyramidal structure.
The fluorine atoms are in the planar positions and the molecular shape is trigonal planar.

trigonal planar # 3

Methane has 4 C—H bonds and no lone pairs of electrons. Sketch the molecule of methane.

Since all of the hydrogen atoms bonded to the central carbon atom are the same, tetrahedral geometry is formed (4 bonds → tetrahedral geometry).

image # 4

Which type of electrons exerts the largest repulsion, bonding electrons or lone-pair electrons?

Repulsion in an order of increasing strength is:
1. bonding pair - bonding pair,
2. bonding pair - lone pair,
3. lone pair - lone pair.

lone-pair electrons # 5

Use VSERP to predict the shape of a NH₄⁺ ion.

number of valence electrons = 5 (nitrogen) + 4 × 1 (hydrogen) = 9
There are four single N—H bonds in the molecule.
Each hydrogen has 0 non-bonding electrons.
The + charge means that one electron has to be subtracted.
number of lone pairs = [9 − 4 × 2 (bonds) − 2 × 0 (hydrogen atoms) − 1 (charge)] ÷ 2 = 0
There are no lone pairs to distort a tetrahedral shape.

tetrahedral End of page