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I really don't understand why metals form metallic bonds. I mean, it makes no sense. It would make much more sense for them to form covalent bonds with themselves and have a 'pseudo-full' outer shell. How does freeing off electrons make them any more stable?

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6 Answers 6

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This is due to the low ionization energies of the metals. It's easier for them to release few electrons from the outer shell to obtain a noble gas configuration rather than consuming several ones. However, the difference between an ionic and a polar covalent bond is always fringe.

Again, it's not true that metals don't form covalent bonds at all. I guess you never heard of quadruple bond or δ-bond. There are several examples such as $\ce{K2[Re2Cl8]·2H2O}$ and Chromium(II) acetate hydrate.

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    $\begingroup$ >This is due to the low ionization energies of the metals || This is BS. silicon: ionisation energy 8 eV (nonmetallic bond); mercury: ionisation energy 10 eV(metallic bond). $\endgroup$
    – permeakra
    Commented Dec 1, 2015 at 8:17
  • $\begingroup$ Silicon is not entirely nonmetallic, it becomes a metal when melted. The point is, what makes things metallic is not as simple and brute force as ionization energy. $\endgroup$ Commented Feb 24, 2018 at 14:02
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    $\begingroup$ Atoms in molecules $\ce{Mo2}$ and $\ce{W2}$ in a gaseous phase form even a sextuple bond $\endgroup$
    – Poutnik
    Commented Mar 30, 2022 at 15:22
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  1. Bulk d-metals and especially intermetallic compounds often do have significant covalent bonding.
  2. Strictly speaking, metal bonding is a kind of covalent bonding in a sense. It is a common knowledge, that 3- or more atoms may be bound by one pair of electrons, like $\ce{H3+}$ ion. In bulk metals, similar bonding bond is working for valence electron.
  3. Even more strictly speaking, to fully understand the matter you have to consider the matter from position of theory of molecular orbitals.
    Let's assume we have a nanocrystal of 100 3d row element atoms. From 400 orbitals of said atoms 400 molecular orbitals are formed, with strict distribution of their energies depending on shape and size of the crystal and crystal cell. Then, the only difference between metals and non-metals is that metals have semi-full group of orbitals of same energy (so called conductivity zone, while non-metals have only full and vacant orbitals.
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  • $\begingroup$ 3 is rather misleading, and in this form does not explain why metals would have such orbitals or what is special about them $\endgroup$
    – Greg
    Commented Nov 12, 2016 at 9:08
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The difference between covalent and metallic bonding looks much clearer when written on paper than it does in the real world.

If you take the elements from group 15 from phosphorus downwards and analyse them, you have ‘true’ covalent bonding for one allotrope of phosphorus (the white one, $\ce{P4}$) and ‘true’ metallic bonding for bismuth. The entire transition downwards from phosphorus — technically already including phosphorus’ red and black allotropes — is a transition from covalent to metallic bonding.

If you look at an orbital picture of metallic bonding, you see the clear similarity to covalent bonding, except that you aren’t dealing with discreet molecules but a huge network.

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Arguably, the bonding in bulk metals is covalent bonding

The bonding in metals is often treated as a special type of bonding because it has unusual properties where many electrons are delocalised enough to make the bulk material conducting. So we don't describe the bonding the same ways we describe the bonding in, for example, diamond.

But, from a molecular orbital point of view the major difference between a metal and diamond is not which orbitals exist, but which bonds are occupied. To simplify a lot, in molecules when atoms join up we get molecular orbitals from combinations of the atomic orbitals. These fill up with the spare electrons. In bulk solids (as opposed to small molecules) the same process happens but there are far more orbitals to fill and there is a lot of duplication of their energy levels. In diamond the filled orbitals are very localised and the bonding orbitals confine most electrons between adjacent carbons. In metals, the highest level filled orbitals are delocalised over many atoms and the electrons can wander fairly freely over the bulk. WE could still describe them as "covalent" bonds but many of their properties are better understood by completely different tools than those used to describe small molecules.

That this is not a completely silly way to look at metallic bonding is illustrated by the bonding in graphite (or graphene). Hardly anyone would describe that as a solid held together by anything other than covalent bonds. But the pi-orbitals in the plane of the hexagonal rings are delocalised over the whole plane. Electrons can wander about fairly freely in those orbitals and the conductivity in the plane can reach metallic levels. Despite the 2D metal-like behaviour in the plane we still call these bonds "covalent". Bonding in actual metals isn't qualitatively different to this: it is just 3D not 2D in how electrons can move.

So from a chemistry point of view metallic bonds are just a special case of covalent bonds where the electron delocalisation is in 3D not 2D as it is in graphite. But we treat that special case very differently as we need a whole different set of tools to analyse and think about the properties of metallic bonds that matter and the tools for thinking about the bonding in small molecules are of little to no help.

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Certain metals under certain conditions form highly covalently bonded structures.

Certain allotropes (phases) of metals have no metal qualities such as alpha-tin. Formed at about 13.2 degrees celcius with pure tin- it has no metallic qualities and pure covalent bonds.

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  • $\begingroup$ Alpha-tin is a semi-metal, it certainly has metallic qualities (it still is conductive for one thing). The issue is that while the bottom of the conduction band overlaps with the top of the valance band in energy, those occur at different $k$ values. $\endgroup$
    – Jon Custer
    Commented Feb 24, 2022 at 21:10
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First of all! Metal do form covalent bond. It is very common in transition metal like platinum, palladium.

However, it is not the way you are talking about. Typically, when pure metal atoms bond together, they prefer metallic bond.

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