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Brief Background: I was studying about the 'classical electronic configuration' of the first 20 elements of the periodic table today and was bewildered by the fact that there are an equal number of metals and non-metals.


Question(s)

How on earth is it possible for 92 of the 118 discovered elements to be metals!? I've read a little about it here, but if I agree with this, how is it that Calcium essentially 'forgets' about its penultimate shell but Scandium and many others don't? And still, how is it possible that some electrons 'jump' from the penultimate shell to the valence shell to result in variable valency in such a case?

Holy grail

I hope that a conceptual answer awaits for me, but I may be fine with others as well. Specifically, what is the exact physical cause behind this phenomenon?

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    $\begingroup$ I think it's possible to argue that metallicity (or more generally, delocalized bonding) is more natural, and it's the non-metals which are unusual - why do we have a bunch of those? $\endgroup$ Aug 28 at 8:26
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    $\begingroup$ Some non-metals have metallic states, see e. g. carnegiescience.edu/news/… $\endgroup$
    – Karsten
    Aug 29 at 1:11
  • $\begingroup$ Simultaneous posting across various sites is generally discouraged, but if you do not obtain answers (or satisfactory ones) on one site it can be justified to consider another site. $\endgroup$
    – Buck Thorn
    Aug 29 at 3:12
  • $\begingroup$ @NicolauSakerNeto Can you please explain 'delocalized bonding'? $\endgroup$ Aug 29 at 15:32
  • $\begingroup$ @Karsten Interesting! Never knew of this! $\endgroup$ Aug 29 at 15:32

3 Answers 3

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Very simply, metallicity is defined (*other than in astronomy, where everything beyond He is a metal) by behavior: a shared sea (conduction band) of electrons, causing conductivity, reflectivity, etc.

Elements such as the alkali metals are "glad" to be rid of their lone electron to have a complete outer shell, and similar effects are observed for dropping an electron or two to complete an orbital.

However, other important effects are shielding, average distance from the positive nucleus, and relativity effects, all of which make outer electrons easier to drop off. For example, look down the column of chalcogens:

  • Oxygen is very quick to add an electron to complete its outer shell, definitely nonmetallic.
  • Sulfur is a bit less likely to grab an electron, but is still nonmetallic.
  • Selenium is a metalloid, shiny in some allotropes, and conducting electricity with the help of a photon or two. Clearly, the outer electrons are not firmly attached.
  • Tellurium is also a semimetal, more conductive than Se.
  • Polonium is a metal, conductive (0.40 µΩ⋅m at 0 °C) and shiny (and incredibly dangerous).

In all these elements, the outer shells are somewhat similar, but the increasing size of that outer shell make them more metallic.

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    $\begingroup$ Well depending on who you ask,polonium is a semimetal. $\endgroup$ Aug 29 at 1:14
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    $\begingroup$ @BruhMoments Maybe. But there are few chemists indeed who will volunteer to do the work with polonium to double check that idea. $\endgroup$
    – matt_black
    Aug 29 at 13:59
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    $\begingroup$ @matt_black, Good point! See Derek Lowe on Po and other things he won't work with: chemistryworld.com/opinion/a-risk-not-worth-taking/… $\endgroup$ Aug 29 at 19:46
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I will use this excellent comment as a starting point:

[Nicolau Saker Neto] I think it's possible to argue that metallicity (or more generally, delocalized bonding) is more natural, and it's the non-metals which are unusual - why do we have a bunch of those?

Definition and structure of metals

Mostly, we think of metals as solids, although the liquid mercury is commonly defined as metal as well. While the OP's question is about elements, alloys are classified as metals as well, even when they contain some non-metalls like the carbon in steel.

So to be considered a metal, an element has to be in a condensed physical state. In a solid or liquid, the atoms will be close together; in an element, they will all be the same. Either the atoms will interact with each other in the same way (like in a metal), or they will form groups of strongly interacting atoms (called molecules) interacting less strongly with other groups (like liquid bromine or solid iodine). The former seems more "natural", given that the atoms are all of the same type. The absence of molecules and the abundance of nearest neighbors gives rise to the band structure of metals (and the delocalisation of electrons) that explains many of their properties.

Temperature and pressure

[OP] How on earth is it possible for 92 of the 118 discovered elements to be metals!?

On earth, indeed.

If an element is in the gas phase, we typically don't describe it as a metal. This includes hydrogen, oxygen, nitrogen, the halogens up to chlorine and the noble gases up to radon. Some of these elements (including hydrogen, oxygen, nitrogen, chlorine, and xenon) solidify and form metallic phases when under sufficient pressure. A paper on high-pressure forms of chlorine claims [[1]2]

Under sufficient compression, all molecular systems are expected to collapse into close-packed metals.

On the other hand, if we increase the temperature, most metals will vaporize (or even form a plasma). This shows that the number of elements we classify as metals depends on the ambient conditions we experience on the surface of our planet these days.

Allotropes

There are also elements that have allotropes, with one metallic and one non-metallic. A good example is tin, which has a metallic allotrope (white tin) and a non-metallic allotrope (gray tin), which is a powder and does not conduct electricity. Carbon comes in the form of diamond (non-conducting) and graphit (conducting). Boron has metal-like allotropes as well. Finally, metalloids are semiconductors, so their metallic character depends on temperature.

Chemical properties

This answer focused on physical properties. The key chemical property of metals is that they tend to form (monoatomic) cations. Taking away electrons from an isolated atom always costs energy (all ionization energies are positive). Interacting with a polar solvent (like water, made of nonmetals) or with anions (made of or at least containing nonmetals) offsets this cost in some cases, so we observe the formation of cations. This means that without non-metals, we would not be able to observed metals forming cations. To explain (or rationalize) which elements form cations or anions, we invoke the electronic structure of atoms, as you will find in any textbook. The introductory textbooks focus on main group metals and non-metals, yet the bulk of metals are transition metals and f-group metals, where things quickly become hand-waving (OP's "jumping" electrons).

Cause and effect

[OP] Specifically, what is the exact physical cause behind this phenomenon?

Sorry, other than pointing to experimental evidence showing that most elements have metallic properties under ambient conditions, and adding that this is well-described by theory (but the theory is complicated and the applications are subtle), I got nothing.

Reference

  1. Dalladay-Simpson, P.; Binns, J.; Peña-Alvarez, M.; Donnelly, M.-E.; Greenberg, E.; Prakapenka, V.; Chen, X.-J.; Gregoryanz, E.; Howie, R. T. Band Gap Closure, Incommensurability and Molecular Dissociation of Dense Chlorine. Nat Commun 2019, 10 (1), 1134. DOI:10.1038/s41467-019-09108-x.
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  • $\begingroup$ Can you please explain the theory behind it as well (in brief), mentioned in the subheading 'Cause and effect'? $\endgroup$ Aug 29 at 15:45
  • $\begingroup$ @ChinmayKrishna Sorry, I can't. If anyone else can, chemistry instructors would all lose their jobs. The results are easy to show (e.g. on periodic table), the theory takes a couple of years. $\endgroup$
    – Karsten
    Aug 29 at 15:51
  • $\begingroup$ The simplest view is if the bonding is due to overlap of the highest energy shells on the atoms then almost everything should be a metal as the resulting bands will be partially filled except when the atomic shell is full, such as the noble gases or alkaline earths. Thus the default position is nearly everything should be a metal and you have to explain why it isn't, not the other way around. $\endgroup$
    – Ian Bush
    Aug 29 at 20:13
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    $\begingroup$ And as for why Mg and Zn and similar are metals, well the problem is more than one band is involved, see chemistry.stackexchange.com/questions/164694/… $\endgroup$
    – Ian Bush
    Aug 29 at 20:13
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The elements are defined by their "molecular structure", do they form covalent bonds. Covalent bonds are formed by elements with two types of valence orbitals, first: orbitals that are close to the nucleus [1s,2s] second: degenerate orbitals of the same energy level [2p, 3p, d and f to a lesser extent] that are not shielded from the nucleus as electrons are added. They form covalent molecular bonds that are strong and directional maximizing attraction of electrons to the nuclei and minimizing electron repulsion. This explains the nonmetallic behavior of hydrogen and helium, the less nonmetallic behavior of lithium and beryllium and the nonmetallic behavior of boron thru neon. As we progress down the periodic table the valence electrons, while possibly having similar shielding, are farther from the nucleus so molecular bonds are farther from the nucleus and are weaker. Rather than saying metallic character increases nonmetallic behavior decreases so eventually the only nonmetals left are iodine, possibly tellurium, and the inert gases; and they are all suspect.

Nonmetals usually have large ionization potentials and electron affinities; loss of electrons is difficult and gaining one is usually exothermic. This affects chemical properties and nonmetals tend to be involved with negative ions, anions. Finally! metals, since their valence electrons are more shielded, have lower ionization potentials and less exothermic electron affinities. The individual electronic structure is what is important. This means that metals in their lower oxidation states tend to form positive ions, cations, and in their higher oxidation states start to assume nonmetal properties forming complex anions; the borderline is the +3 ions. A question here might be why do ionic compounds form? why not heteroatom metals? A quick answer is electronegativity differences. What gets those electrons closer to the nucleus is what is important.

With the exception of hydrogen and lithium covalent bonding is weak for elements that have lower nonmetal properties; lets call them metals! They still bond together sometimes with immensely strong bonds. Their bonds involve mobile electrons evidenced by electrical conductivity and no or little nuclear motion. The bonding possibly involves attractions of nuclei for the electrons of surrounding nuclei added to the attractions of the electrons in conduction bands for the array of nuclei. There must be an explanation for the immense range in physical properties from mercury to osmium This is beyond the scope of my understanding.

There are more metals than nonmetals because the electronic configurations that result in nonmetal properties are very limited to the elements in the first few periods and far right groups in the periodic table.

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  • $\begingroup$ In the array of nuclei of a metallic structure the better shielded the valence, especially the s, electrons are the stronger the attraction of nucleus A to the electrons of nuclei B,C etc. becomes than the repulsion of the nuclei. Someone must have worked out the QM implications for this. There must be a reason for the strong attraction between positive ions in a close packed array. The more shielded the valence electrons are the more metallic properties. $\endgroup$
    – jimchmst
    Sep 24 at 20:13

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