Why do metallic bonds have delocalized electrons, whereas covalent bonds only share electrons? Does it have something to with the atomic masses of the individual non/metals or something else? I'm really interested to know.

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    $\begingroup$ It is wrong to think of the covalent bond as a pair of shared electrons between two atoms. The only true difference between metallic and covalent bonds is that the metallic bond’s ‘molecule’ is magnitudes larger. $\endgroup$
    – Jan
    Oct 20, 2015 at 20:11
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    $\begingroup$ @Jan That is a really nice way of thinking about it. You should write it up into an answer. $\endgroup$
    – bon
    Oct 22, 2015 at 13:09
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    $\begingroup$ Indeed, the electrons in a covalent bond are also delocalized; it's just that they are delocalized over two atoms. Or maybe three, if we think of a three-centered bond like those in boranes. Or six atoms, if we think of benzene. Or even more, because larger conjugated structures are also known. $\endgroup$ Dec 21, 2015 at 7:56

2 Answers 2


Why do metallic bonds have delocalized electrons, whereas covalent bonds only share electrons?

To be somewhat flip this is like asking why are redbirds red and blue birds blue?

In order to pigeon-hole bonding ionic, covalent, aromatic, and metallic bonds are "artificial" constructs used to predict behavior. But there are always exceptions, and no bond is purely in any one of the categories.

So in general a sample which has metallic bonds has delocalized electrons and hence will conduct electricity. Such elements would be metals. If the sample only has covalent bonds then there would not be any delocalized electrons and the sample would be a nonmetal.

For some elements the nature of the bonding yields different results. For example graphite conducts electricity, but diamond won't.

To add another wrinkle in this we can think of gaseous ionic or covalent molecules but in general being metallic is a "bulk" property. So we have to consider not only the bonding in the "molecule", but how "bonding" is distributed through the solid matrix of the sample. For example there are discrete water molecules in ice, but in a diamond there are no discrete carbon "molecules."

  • $\begingroup$ You write "For example there are discrete water molecules in ice, but in a diamond there are no discrete carbon "molecules."" <-- what about silicon dioxide crystal? $\endgroup$
    – barlop
    May 13 at 10:25

The question’s premise is incorrect. While it is nice to think of clearly localised bonds and while it provides an easy access to chemistry at a high school level thinking of just moving single lines (a.k.a. bonding electron pairs) this is not much more than a very simplified model of reality.

The most well-known counterexample is the benzene molecule $\ce{C6H6}$ which is always introduced as not having localised double and single bonds but rather an ‘electron cloud’ in a delocalised π system above and below the molecular plane. But that is only the tip of the iceberg. In the end, all molecules that are not diatomic have all their bonding electrons delocalised across the entire molecule. Check out, for example, the molecular orbitals of water.[1]

It is only through mathematical methods (linear combination) that we can backtrack from molecular orbitals back to orbitals of a specific bond. For example, by taking $a_2 \Psi_\mathrm{mo2} + a_3 \Psi_\mathrm{mo3} + a_4 \Psi_\mathrm{mo4}$ and correctly choosing the coefficients $a_n$ we can construct the wavefunction $\Psi_{\unicode[Times]{x3c3}(\ce{O-H})}$ that corresponds to one of the two $\ce{O-H}$ bonds; by taking $a_3' = - a_3$ we will arrive at the other $\ce{O-H}$ bond. The physical reality is much closer to the entirely delocalised system, though.

That in turn means that the distinction between metals (large, three-dimensional networks of atoms, all somehow bonded to each other), ionic compounds (large, three-dimensional networks of atoms considered ions of opposite charge, somehow also bonded to each other) and molecules (small, three-dimensional discreet entities of atoms, somehow bonded to each other) lies only in the size of the network or in the types of particles involved. Indeed, graphite (the carbon modification) and black phosphorus both exhibit clear metal characteristics even though the elements are clearly non-metallic: absorption of all wavelengths (i.e. black colour), conductivity, high melting points and more.

[1]: Site by professor Zipse, LMU Munich; held in German. Click on one of the mos to access an image of the respective orbital.


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