A-level Chemistry/OCR (Salters)/Metallic bonding
Metallic bonding is what holds metal atoms together in the solid and liquid states.
Metals contain delocalised electrons 
The simplest description* of the bonding in metals consists of a lattice of metal cations in a sea of electrons.
A metal atom releases its valence electron(s) into the sea, but its core electrons remain localised around its nucleus. This extensive delocalisation of valence electrons is responsible for most of the properties of metals.
Physicists sometimes consider metals to consist of a 'gas' of electrons, free to flow around the regular arrangement of metal cations. In a metal, the valence electrons are free to roam throughout the whole structure, but cannot escape completely without getting extra energy from somewhere else.
Delocalised electrons conduct charge 
Because the valence electrons in a metal are completely delocalised, they can flow through the metal if an electric field or circuit encourages them to. Electrical conduction consists of electrons flowing in a specific direction, and generally decreases at higher temperatures as the electrons collide more frequently with vibrating atoms.
Delocalised electrons help conduct heat 
Electrons also help conduct heat through metals, supplementing the conduction that occurs in any solid as atoms bash into each other. Thermal conduction, in contrast to electrical conduction, consists of electrons moving in random directions and generally increases at higher temperatures as more frequent collisions help propagate thermal vibrations.
Metallic bonding is equally strong in all directions 
Metallic bonding is strong and flexible 
- Metals are malleable and ductile because metallic bonding is not directional
- Metallic bonding does not oppose deformation because a change in the positions of the metal nuclei does not increase the energy of the structure
Covalent bonding is strong but inflexible 
- Covalent bonds, on the other hand, are only strong in certain directions and oppose deformation
- Diamond is exceptionally hard because deforming it requires bond breaking, and it has many strong covalent bonds in all directions
- Molecular solids, like sulfur, S8, have strong intramolecular covalent bonding but weak intermolecular bonding, so they are soft
Ionic bonding is strong but ionic solids are brittle 
- Ionic bonds are not directional either, but deforming an ionic solid pushes ions of the same charge closer together
- The repulsion from these like charges either opposes any deformation
- This explains why ionic solids are brittle - they will resist deformation strongly up to a point, but beyond that point, they fail
The strength of metallic bonds 
The stronger the bonding in a metal, the higher its melting and boiling points will be. To melt a metal requires the metallic bonding to loosen, whereas boiling breaks almost all the metallic bonding (in the gas phase, metals exist either as single atoms or as small molecules or clusters, such as Li2).
Metallic bonding can be weak, but is more often strong. Atoms that contribute one electron to metallic bonding are weakly bonded and have low melting points. Atoms that contribute two electrons have much stronger metallic bonding and therefore much higher melting points
Strength of metallic bonding in Groups 1, 2 and 3 
|metal||mp / ° C||metal||mp / ° C||metal||mp / ° C|
The Group 1 metals, Li-Cs, all have low melting points because they contribute only one electron per atom to metallic bonding. The Group 2 metals, Be-Ba, have moderate melting points since they release two electrons per atom to metallic bonding. Group 3 metals have high melting points, as they can provide three electrons per atom for metallic bonding.
Group 13 metals are weird 
|metal||mp / ° C|
The Group 13 metals, B-Tl, do not seem to follow any pattern!
Boron is a metalloid and is covalently bonded — it is too electronegative to release its valence electrons for metallic bonding.
Al, Ga, In and Tl have different core electron configurations. Aluminium has a noble gas core, [Ne], but gallium and indium have "noble gas plus filled d subshell" cores, [Ar]3d10 and [Kr]4d10, respectively. Thallium has a "noble gas plus filled d and f subshells" core, [Kr]4f145d10.
The filled d and f subshells found in the cores of Ga, In and Tl are much less effective at shielding their valence electrons from their nuclei than filled s and p subshells are, because of the diffuse shapes of d and f orbitals. As a result, their valence electrons feel a stronger pull from the nucleus (a greater effective nuclear charge) and are less easily released for metallic bonding.
Strength of metallic bonding in transition metals 
The transition metals almost all have very high melting and boiling points, since they have many valence electrons to contribute to metallic bonding. Excluding the sort-of-transition-sort-of-not metals Zn, Cd and Hg, the minimum melting point is 961 °C (silver), while the maximum is 3407 °C (tungsten, higher than all elements except carbon). The mean melting point of the transition metals is around 1900 °C.
Although the melting points of transition metals are high, the increase in melting point on moving from Group 3 to Group 4 is not as great as the increases seen on going from Group 1 to Group 2 and Group 2 to Group 3. Also, the melting points of the transition metals increase from left to right until around the middle of the d block, where they begin to decrease again.
Metals are uniquely useful 
Metals and alloys that are strong are used as structural materials, most importantly iron (mostly as steel) and aluminium (alloyed with Mg, Si and/or Cu). While many materials are strong, metals and alloys are special because they are also elastic (can be deformed but will return to their original shape) and thus easy to shape (malleable, ductile and machinable). Other strong materials that are not metallically bonded are brittle rather than elastic, and may fracture too easily or be difficult to manufacture in the required shape.
* A more complicated, but more detailed and realistic, way of looking at metallic bonding is called band theory, which is introduced in A-level Physics and studied by undergraduates in physics and chemistry, but is not part of Salters Advanced Chemistry. It basically says that atomic orbitals from each metal atom interact to create an almost infinite number of great big delocalised orbitals, each orbital delocalised over every metal atom. The energy levels of these delocalised orbitals cover a range of energies called a band (hence the name band theory), which can be treated as continuous, so that there is an orbital for any energy, as long as that energy is within the band.