3.1.3 Bonding - Shapes of simple molecules and ions


Students should:
  • understand the concept of bonding and lone (non bonding) pairs of electrons as charge clouds.
  • be able, in terms of electron pair repulsion, to predict the shapes of, and bond angles in, simple molecules and ions, limited to 2, 3, 4, 5 and 6 co-ordination
  • know that lone pair/lone pair repulsion is greater than lone pair/bonding pair repulsion, which is greater than bonding pair/bonding pair repulsion, and understand the resulting effect on bond angles

Electron pair repulsion

Like charges repel, negative charges repel other negative charges. The electrons are constrained around the nucleus in pairs. Each pair exerts a repulsion on other electron pairs. To ascertain the actual shape adopted by molecules we must consider these repulsions.

The electrons in a charge region are also attracted to the nucleus. Thus any repulsion will cause the electron pairs to move as far apart as possible, while still remaining attached to the central nucleus.

It is rather like having two magnetic north poles attached to two strings with the strings attached to a central pivot. The magnets will push as far away as possible while staying tied to the central pivot. The final electronic shape adopted depends on the number of regions of electronic charge around the nucleus.


Two regions of electron density (charge)

The two regions move as far apart as possible while remaining fixed to the nucleus. They adopt an arrangement in which they are at 180º to one another.

Here we can see the beryllium dichloride molecule that has only two pairs of electrons on the central atom. They move as far apart as possible and adopt an arrangement as shown.

The beryllium chloride molecule, showing how the  two lone pairs of electrons adopt a linear arrangement to minimise repulsion

The molecule is linear with bond angles of 180º


Three regions of electronic charge

Three regions move as far apart as possible while remaining attached to the central atom. They adopt a trigonal planar arrangement in which the angle formed between the electron pairs is 120º.

Here we can see the boron trichloride molecule. It has three pairs of electrons on the central atom. They move as far apart as possible and adopt an arrangement as shown:

the boron trichloride molecule, showing how the three pairs of bonding electrons arrange themselves to be at angles of 120º to one another to minimise repulsion.

The bond angle subtended by the Cl-B-Cl atoms is 120º in all cases.


Four regions of electronic charge

Four regions of charge adopt a tetrahedral arrangement. This is a three dimensional shape that can be considered formed from the peaks of a regular pyramid (tetrahedron) with the central atom in the very centre of the pyramid. The molecule methane has four regions of electronic charge around the central atom, which adopt a tetrahedral arrangement.

Methane, showing the three dimensional tetrahedral shape adopted by the four bonding pairs of electrons to minimise repulsion.

Methane has a tetrahedral 3-D arrangement


Electronic shape and molecular shape

The shape of a molecule is given by the relative positions of its constituent atoms. This may, or may not, be the same as the electronic shape of the molecule. If all of the electrons around the central atom are involved in bonding, then the molecular shape equals the electronic orientation. However, if there is one, or more, lone pairs (non-bonding pairs) of electrons, then the molecular shape will be different from the electronic orientation.

If, for example, the central atom has three regions of electron density (three charge centres), but is attached to only two atoms then the electronic shape is trigonal planar, but the molecular shape is angular (bent).

Trigonal planar electronic shape - three charge centres

Angular molecular shape - 2 out of 3 charge centres used for bonding.

The shape adopted by electron pairs around a central atom depends on the number of regions of electrical charge. Each region will move so as to minimise the repulsive forces experienced. The molecular shape is not necessarily the same as the electronic shape, as only the positions of the atoms themselves are used to describe the molecular shape.

To determine the molecular shape it is necessary to first determine the shape adopted by the electrons, and only then can the positions of the atoms be known.

The electronic shape is tetrahedral, but the shape of the water molecule considers only the H-O-H. It is angular, or bent.

Three dimensional structure of water.


Unequal repulsion - VSEPR

One of the successes of the Valence Shell Electron Pair Repulsion theory lies in its ability to predict, or explain, the bond angles of molecules. To do this, it considers that electron pairs that are shared by two atoms (bonding pairs) experience less repulsion than lone, or non-bonding pairs, of electrons. The logic of this assumption comes from the idea that shared pairs have their electron density relatively displaced away from the central atom by the bonded atom, when compared to a lone (non-bonding) pair.

This idea gives rise to an order or repulsive force experienced by electron pairs, in which lone pair - lone pair repulsion is greater than lone pair - bonding pair repulsion, which in turn is greater than bonding pair - bonding pair repulsion.

These different forces of repulsion distort the perfectly symmetrical shapes that would be adopted by the electron pairs under ideal conditions. We can appreciate this by studying examples such as the ammonia molecule, NH3.

The three dimensional arrangement of the ammonia molecule.



The ammonia molecule has a central nitrogen atom with four pairs of electrons. These electrons tend to adopt a tetrahedral orientation. A perfect tetrahedron has a bond angle of 109º 28'. However, there is a greater repulsion between the lone pair (non-bonding pair) of electrons and the bonding pairs, than there is between the bonding pairs themselves. This greater repulsive force has the effect of 'squeezing' the hydrogen atoms closer together, creating smaller H-N-H bond angles of 107º.

The ammonia molecule

The ammonia molecule - VSEPR



The water molecule has four pairs of electrons on the central oxygen atom. These electron pairs, or regions of electron density, adopt a tetrahedral arrangement. The theoretical bond angle of the electron pairs would be 109º 28'. However, only two of the electron pairs are actually used in bonding and the other electrons are lone (non-bonding) pairs. The lone pairs of electrons repel the bonding pairs more than the bonding pairs repel each other. This has the effect of closing the 109º 28' bond angle down to 104º 30'.

The water molecule

The water molecule - VSEPR


Dealing with ions

Ions are atoms or groups of atoms that have lost or gained electrons to attain a stable electronic configuration. In the case of ions involving more than one atom the atoms always have an octet of electrons.

When counting up the total number of electrons available, it is important to remember that a positive ion has lost an electron and a negative ion has gained one or more electrons with respect to the constituent atoms.

For example, the ammonium ion has the formula NH4+. It is comprised of one N atom (group V) and four hydrogen atoms. The total valence electrons from the atoms = 5 + 4 = 9. However, the ion has a positive charge, so it has lost a valence electron. It therefore has only 8 valence electrons.

The central nitrogen atom is attached to four hydrogen atoms by four shared pairs of electrons = 8 electrons. Thus, there are no electrons left over.

The three dimensional shape of the ammonium ion showing the tetrahedral arrangement of electon pairs that give rise to the tetrahedral shape of the ion.


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