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Is Hydrogen metallic?

  1. Feb 4, 2013 #1
    I asked my chemistry professor this question, and he gave me a brief answer of "yes and no", and I never followed up to explain that.

    Hydrogen is in the same column as the alkali metals, and it does donate electrons. However, it is gaseous and found diatomically (making it similar to some halogens too).

    I heard hydrogen can be "metallic" under extreme extreme pressure, but then again, doesn't that apply to any other non-metal.

    Lastly, what makes it a non-metal, when it has the same outer ring configuration as the other metals in it's column.

  2. jcsd
  3. Feb 4, 2013 #2

    D H

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    Is this a question you are asking us or a rhetorical question you have answered for yourself?

    The placement of hydrogen in the periodic table has always been a bit controversial. It shares some characteristics with the alkali metals, others with the halogens, yet others the carbon group, and has some characteristics that don't jibe with any of those groups. Your professor's "yes and no" answer is just about right.
  4. Feb 4, 2013 #3
    No, it's not a rhetorical question.

    We both established that hydrogen is both a metal and a non-metal. But WHY?

    What makes it a non-metal, when it has the same outer ring configuration as the other metals in it's column?
  5. Feb 4, 2013 #4


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    The answer is that the metal/non-metal classification scheme we use to sort the elements is ultimately an arbitrary label that we have made up to simplify our understanding of chemistry. It is a useful label in many cases, but it is too simplistic to capture the full diversity of the elements, and there are numerous examples of metals acting like non-metals and vice versa.
  6. Feb 4, 2013 #5
    Good answer.

    I am curious as to what specifically makes hydrogen diatomically gaseous as opposed to naturally solid.
  7. Feb 4, 2013 #6


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    If you want to show off, know that under very extreme pressure , probably never realised on earth, hydrogen is confidently predicted to become metallic - conducting electricity, which is supposed to happen on Jupiter and explain the huge magnetic fields there. The explanation of why this happens is here http://cseligman.com/text/planets/metallichydrogen.htm, there is lots you can follow up, depending on your interests.
  8. Feb 8, 2013 #7
    A bit naive, but I hope intuitive explanation:

    What is special for alcali metallic atoms? It's readiness to give up electron, and smooth acceptance for extra 2nd, 3rd,... electrons. This attracts atoms to form big clusters, rather than 2-atomic molecules.

    What is special for halogens? It's gready for a 1 extra electron, and total indifference for more.

    Hydrogen has the same greedy/indifference step.
  9. Feb 9, 2013 #8
    There is a quite simple underlying pattern of nonmetal-semimetal-metal transition in the main groups.


    Oxygen - mainly covalent diatomic. Unstable three-member chain. No higher molecules, chains or networks.
    Going down: Sulphur does form covalent S2 molecules - in high temperature vapour.As the vapour cools, higher molecules S3 to S8 are formed.
    In condensed phases, sulphur has a fine balance between small cyclic molecules S8 and smaller, and long chains. At low temperatures, the balance is towards S8, forming molecular crystals and fluid yellow liquid; on heating, the balance goes towards viscous liquid of long chains; on further heating, it disintegrates back to smaller molecules in a fluid liquid, but now wider variety of them, with red-brown colour. Sudden cooling can preserve the long chains as a rubbery mass - in time, these revert to S8 molecules.
    Sulphur is a good insulator.

    Selenium - Molten selenium contains mostly long chains, like molten sulphur, with some smaller molecules. On rapid cooling, it forms glassy solid, from which some Se8 and Se7 molecules can be extracted with solvents and which is fairly metastable on standing; on slow cooling, the long chains will form a regular packing of spirals. Brittle, semiconducting.

    Tellurium - regular packing of spiral chains like selenium. Also a brittle semiconductor.
  10. Feb 11, 2013 #9
    This is a very good point.
  11. Feb 11, 2013 #10


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    Care to elaborate? Can't say I have heard a lot about Na- or Na2-.
  12. Feb 11, 2013 #11


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    What Graniar presumably wants to express is nothing more than Paulings definition of electronegativity, namely the difference between electron affinity and ionization energy. In a metal, this difference is small, in the metallic state atoms do not mind very much to be either anionic or cationic. This is a prerequisite for delocalized molecular orbitals forming a good ground state. In contrast, in an insulator only covalent bonds are possible in which ionic configurations play at best a minor role. Hence at normal pressure, the electronegativity of hydrogen is too high for it to form a metal. At higher pressures, the gain in kinetic energy by delocalization rises so that substances of higher and higher electronegativity will become metallic.
  13. Feb 11, 2013 #12
    DrDu, thank you for disclosure.

    Sure, Na2- has no meaning in a free state, (neither an O2-).
    Only in bond to positive charge.

    Please don't perceive my attempts to make things simpler as a sort of crackpottery.
    I realize, that uncorrect simplification may lead to the ground loosing.
    If such a style is undesirable at this forum, please let me know before banning.

    So, let's see, why H3 molecule is unstable in the free state, unlike Na3?

    (H3: http://en.wikipedia.org/wiki/Triatomic_hydrogen)

    H3+ is stable, because 3 H+ shares 2 e- along their 1s orbitals.
    Third e- may fill in only 2s orbitals.
    Energy levels of 1s and 2s differs much enough to make total molecule energy higher than for unbonded H2 and H.

    In case of Na3, third e- has to shift from 3s to 3p orbital, which energy levels are closer.
  14. Feb 14, 2013 #13
    Under extreme pressure such as giant gaseous planets some scientists have hypothesized that hydrogen becomes metallic
  15. Feb 15, 2013 #14
    Look to the other main groups:
    Nitrogen: like oxygen, the stable form is diatomic molecules, and unlike oxygen, no long-lived allotropes.
    Phosphorus: diphosphorus like disulphur occurs in hig temperature gases. At lower temperatures, P4 molecules are preferred, but unlike sulphur (and like selenium), it is metastable - on prolonged heating they get crosslinked. Random polymer red phosphorus is metastable - the stable form, reached on prolonged heating under high pressure, is puckered sheets of black phosphorus
    Arsenic: yellow arsenic consists of As4 molecules and can be condensed on rapid cooling, but is easily transformed into gray arsenic, again puckered sheets - brittle and semiconducting. Black arsenic, also metastable, reseembles red phosphorus.
    Antimony - stable form is puckeres sheets like arsenic. Likewise brittle. And again, metastable allotropes like black arsenic and red phosphorus are formed on rapid cooling of vapour.
    Bismuth - brittle, poorly conducive metal.

    So, again a regular transition from nonmetal to semimetal - but this time reaching metal, as bismuth is easier to examine in bulk than polonium.

    Carbon: dicarbon molecules are common in high temperature vapours, but unlike dinitrogen and dioxygen and like diphosphorus and disulphur readily react at lower temperatures.
    Last edited: Feb 16, 2013
  16. Feb 17, 2013 #15
    Carbon does form some larger molecules - beginning from C60. These, like P4 and S8, are reasonably metastable and form molecular solids - but in case of carbon, need special conditions to form. For carbon readily forms infinite networks on condensing.
    Graphite sheets are conductive; the metastable diamond network is insulating.

    Silicon is not prone to form pi bonds. Solid silicon has diamond structure, brittle crystals - but unlike diamond, semiconducting rather than insulating, and also melting at a much lower temperature.

    Germanium again has diamond structure - a semiconductor like silicon, but much more conductive. And lower melting again.

    Tin does have diamond structure, semiconductive and brittle form of gray tin. On slight warming, however, it converts into soft, conductive white tin.

    And lead is always soft and conductive metal - unlike tin and bismuth.

    Boron - not sure what molecules form boron vapour. Boron certainly forms several refractory, insulating forms of covalent bond networks.
    By contrast, aluminum is a good conductor, and ductile. Also reasonably high melting - gallium and lower are much softer and lower melting compared to aluminum.

    Beryllium is a metal, and conductive - but high melting and brittle, unlike magnesium and alkaline earths.

    Lithium also is a metal - conductive and soft, although higher melting than the other alkali metals.

    The systematic pattern is that in all main groups, the elements in earlier periods have a tendency to covalent bonds, and in second period in suitable groups to multiple bonds.
    Last edited: Feb 17, 2013
  17. Feb 17, 2013 #16


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    Any way your posts are related to the original question, or is it that you just want to type something, no matter if it is on-topic or off-topic?
  18. Feb 17, 2013 #17
    Yes, I believe they are. No one else, including you, has yet bothered to address the real approach to the original questions:
    What did you answer?

    The thing is, that besides the outer electron configurations, which is the same in each main group through all periods, there are systematic trends in what the same configurations in different periods/different electron shells will do.

    Hydrogen is a nonmetal because unlike alkali metals, and like the other nonmetals in the lower electron shells, hydrogen preferentially forms localized covalent bonds, and resists formation of delocalized metallic bonds.

    Hydrogen is diatomic because, unlike alkali metals, it cannot form metallic bonds and because, like halogens, it has only one covalent bond to form. Hydrogen is gaseous and not solid because it is diatomic as explained above and because unlike iodine and like and even more than chlorine, fluorine, oxygen and nitrogen, diatomic hydrogen molecules have low polarizability, form weak van der Waals bonds, and condense at extremely low temperatures.
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