Editing Metallic bonding (section)

==The nature of metallic bonding==
The combination of two phenomena gives rise to metallic bonding: [[delocalized electron|delocalization of electrons]] and the availability of a far larger number of delocalized energy states than of delocalized electrons.{{clarify|date=January 2014}} The latter could be called [[electron deficiency]].

===In 2D===
[[Graphene]] is an example of two-dimensional metallic bonding. Its metallic bonds are similar to [[aromaticity|aromatic bonding]] in [[benzene]], [[naphthalene]], [[anthracene]], [[ovalene]], etc.

===In 3D===
[[Metal aromaticity]] in [[metal cluster]]s is another example of delocalization, this time often in three-dimensional arrangements. Metals take the delocalization principle to its extreme, and one could say that a crystal of a metal represents a single molecule over which all conduction electrons are delocalized in all three dimensions. This means that inside the metal one can generally not distinguish molecules, so that the metallic bonding is neither intra- nor inter-molecular. 'Nonmolecular' would perhaps be a better term. Metallic bonding is mostly non-polar, because even in [[alloys]] there is little difference among the [[Electronegativity|electronegativities]] of the [[atom]]s participating in the bonding interaction (and, in pure elemental metals, none at all). Thus, metallic bonding is an extremely delocalized communal form of covalent bonding. In a sense, metallic bonding is not a 'new' type of bonding at all. It describes the bonding only as present in a ''chunk'' of condensed matter: be it crystalline solid, liquid, or even glass. Metallic vapors, in contrast, are often atomic ([[mercury (element)|Hg]]) or at times contain molecules, such as [[sodium|Na<sub>2</sub>]], held together by a more conventional covalent bond. This is why it is not correct to speak of a single 'metallic bond'.{{clarify|date=January 2014}}

Delocalization is most pronounced for '''s'''- and '''p'''-electrons. Delocalization in [[caesium]] is so strong that the electrons are virtually freed from the caesium atoms to form a gas constrained only by the surface of the metal. For caesium, therefore, the picture of Cs<sup>+</sup> ions held together by a negatively charged [[nearly-free electron model|electron gas]] is very close to accurate (though not perfectly so).{{efn|If the electrons were truly ''free'', their energy would only depend on the magnitude of their [[wave vector]] '''k''', not its direction. That is, in [[momentum space|'''k'''-space]], the Fermi level should form a perfect [[sphere]]. The [[Fermi surface|shape of the Fermi level]] can be measured by [[Electron cyclotron resonance|cyclotron resonance]] and is never a sphere, not even for caesium.<ref>{{cite journal|title=The Fermi Surface of Caesium|author1=Okumura, K.  |author2=Templeton, I. M.  |name-list-style=amp |journal=Proceedings of the Royal Society of London A|issue=1408 |year=1965|pages=89–104|jstor=2415064|doi=10.1098/rspa.1965.0170|volume=287|bibcode = 1965RSPSA.287...89O|s2cid=123127614 }}</ref>}} For other elements the electrons are less free, in that they still experience the potential of the metal atoms, sometimes quite strongly. They require a more intricate quantum mechanical treatment (e.g., [[tight binding]]) in which the atoms are viewed as neutral, much like the carbon atoms in benzene. For '''d'''- and especially '''f'''-electrons the delocalization is not strong at all and this explains why these electrons are able to continue behaving as [[unpaired electron]]s that retain their spin, adding interesting [[magnetism|magnetic properties]] to these metals.

===Electron deficiency and mobility===
Metal [[atoms]] contain few [[electron]]s in their [[Electron shell#Valence shells|valence shells]] relative to their periods or [[energy level]]s. They are [[Electron deficiency|electron-deficient]] elements and the communal sharing does not change that. There remain far more available energy states than there are shared electrons. Both requirements for conductivity are therefore fulfilled: strong delocalization and partly filled energy bands. Such electrons can therefore easily change from one energy state to a slightly different one. Thus, not only do they become delocalized, forming a sea of electrons permeating the structure, but they are also able to migrate through the structure when an external electrical field is applied, leading to electrical conductivity. Without the field, there are electrons moving equally in all directions. Within such a field, some electrons will adjust their state slightly, adopting a different [[wave vector]]. Consequently, there will be more moving one way than another and a net current will result.

The freedom of electrons to migrate also gives metal atoms, or layers of them, the capacity to slide past each other. Locally, bonds can easily be broken and replaced by new ones after a deformation. This process does not affect the communal metallic bonding very much, which gives rise to metals' characteristic [[malleability]] and [[ductility]]. This is particularly true for pure elements. In the presence of dissolved impurities, the normally easily formed cleavages may be blocked and the material become harder. Gold, for example, is very soft in pure form (24-[[Carat (purity)|karat]]), which is why alloys are preferred in jewelry.

Metals are typically also good conductors of heat, but the conduction electrons only contribute partly to this phenomenon. Collective (i.e., delocalized) vibrations of the atoms, known as [[phonon]]s that travel through the solid as a wave, are bigger contributors.

However, a substance such as [[diamond]], which conducts heat quite well, is not an electrical conductor. This is not a consequence of delocalization being absent in diamond, but simply that carbon is not electron deficient.

Electron deficiency is important in distinguishing metallic from more conventional covalent bonding. Thus, we should amend the expression given above to: ''Metallic bonding is an extremely delocalized communal form of electron-deficient{{efn|Electron deficiency is a relative term: it means fewer than half of the electrons needed to complete the ''next'' noble gas configuration. For example, lithium is electron deficient with respect to [[neon]], but electron-''rich'' with respect to the previous noble gas, [[helium]].}} covalent bonding''.