IB Chemistry - Bonding

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In this section we examine the forces that hold a metal structure together. The metallic bond.

Delocalisation of electrons

Metal atoms, in common with all other metals apart from the noble gases, cannot exist for very long on their own. Metal atoms aggregate and attract one another in an attempt to stabilise themselves.

Metal atoms have very few electrons in the outer shell (valence electrons) and so cannot achieve a full outer shell by gaining electrons or sharing electrons. They tend to lose electrons, transferring them to non-metal atoms. However, in the absence of non-metal atoms, the only way that they can achieve stability is by sharing all of the outer shell electrons in giant delocalised orbitals. It is these delocalised orbital electrons that give metals their unique characteristics.


Formation of ions

Losing the outer electrons to a large delocalised orbital leaves the metal atoms as ions. These ions are then held in place by the attraction of the negative charge in the delocalised orbital. The ions themselves are arranged in a giant lattice (network).

The charge on the ions depends on the number of outer shell electrons. Group 1 metals provide one electron per atom to the delocalised orbital and the ions formed have a 1+ charge. Group 2 atoms have ions with a 2+ charge.

Transition metals also lose electrons forming ions, but the number of electrons cannot be predicted from the group number (as they are not arranged in groups). Generally transition metals form 2+ ions.


The metallic bond

The sea of electrons is a negative charge cloud that attracts all of the positive ions. It is rather like marbles stuck into blu-tack. The metal ions would repel each other without the electron charge cloud, however the force of electrostatic attraction between the electrons and the positive ions holds the whole structure together.

The strength of metallic bonding is a function of the number of electrons provided by the atoms and the consequent charge on the metal ions. The ionic radius also plays a part, as smaller ions exert a greater force of attraction on the negative charge cloud.

The effect of these two factors can be seen by comparing the melting points (the temperature needed to overcome the forces within the metal structure) down group 1 and across the third period.

Group 1 metals Li Na K Rb Cs
ionic radius / nm 0.068 0.098 0.133 0.148 0.167
melting point / K 454 371 337 312 302

It is clearly seen that as the ionic radius increases so the melting point decreases. Caesium would be a liquid on a warm summer's day.

Period 3 metals Na Mg Al
ionic radius / nm 0.098 0.065 0.045
ionic charge 1+ 2+ 3+
melting point / K 371 922 936

Although magnesium has a similar radius to lithium, the melting point is far higher, indicating that the effect of doubling the ionic charge is much more significant.

Aluminium has a higher melting point than magnesium although not such a diference as between lithium and magnesium. It is thought that the high charge density of the aluminium 3+ ion pulls electron density back onto the aluminium ions effectively decreasing their ionic charge.

Aluminium is known to do this in its compounds, giving them a high degree of covalent character, so it seems reasonable that similar effects apply to the metallic bond.


The metallic lattice

As the metal ions in a lattice of a metallic element are all the same radius they can easily pack together like marbles in a bucket.

The most common arrangement is called hexagonal close packing (HCP). It is the most efficient way for spheres to pack close together.

There are two main close packing systems, depending on how the third layer is placed in comparison to the other two. These two packing systems are called ABA and ABC. If the ions of third layer are directly above those of the first layer, it is called ABA. If the ions of the third layer sit in 'holes' that are not directly above any other ion the packing is called ABC. The best way to visualise this is using models.

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