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Why do some people say that hydrogen should be above lithium in the periodic table and others argue it should be above fluorine?

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what have you found till now? –  tvamsisai Feb 16 '13 at 17:00
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This comes down to the fact that hydrogen has a $1s$ valence shell, capable of holding only two electrons.

In some sense, hydrogen is like the halogens, in that it can achieve the electronic configuration of a noble gas (namely, helium, which has a full $1s$ level) by gaining an electron. This reduction yields the hydride ion, $\ce{H-}$. It should be noted that hydride does not have the same stability as the halogen anions (e.g. hydrogen has an electron affinity of 0.75 eV vs. 3.40 eV for fluorine).

In another sense, hydrogen is like the alkali metals, as it only contains a single electron in its valence shell. Whilst the alkali metals can attain noble gas configurations by losing their valence electrons, it's a bit abstract to talk about hydrogen doing the same by losing its single electron to generate a proton ($\ce{H^{+}}$)(1). Nevertheless there is strong analogy between the electronic structure of hydrogen and the alkali metals which motivates hydrogen being often placed in this group.

Periodic table makers sometimes get away with duplicating hydrogen above the alkali metals and the halogens, or putting it in no group in particular.

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The simplest explanation is that hydrogen has a valence shell that can only accept two electrons. Hydrogen is thus in the unique position of being just as likely to want to gain an electron as to lose the one it has; either way, it reaches a stable valence state. For this reason, it defines the "center point" of the Pauling electronegativity scale; elements with a lower number are "electropositive" (they prefer to get rid of their valence electrons), while elements with a higher number are "electronegative" (they want more valence electrons than they have).

Elsewhere in the periodic table, alkali metals which have only one or two valence electrons want to get rid of them, to get back to a complete valence shell underneath. They'd need to gain 6, 7 or even more electrons to stabilize their existing valence shells, and that's unlikely (they don't have enough electrons to share, and their nucleus size won't support an additional 7 electrons at that energy level). Halogens, on the other side of the table, have a nearly-complete valence shell, and just need one electron to get there, so they'll take it from anything that'll give it up (and in the case of some of the more vigorous halogens, like fluorine, they'll even take it from things that won't give them up, like other halogens and even noble gases).

Hydrogen, however, can do either one. It can lose its electron and become a naked proton, the H+ cation, or it can gain an electron and become a H- anion. Which of these it is more likely to do depends on what it's reacting with. With alkalis, it'll happily accept the extra proton and form a hydride, such as LiH. With halogens, it'll happily give up the proton and form a hydrohalic acid like HF, HCl or HBr. With other elements, it'll share, generally forming "polar covalent" bonds where the other atom in the bond attracts the electrons more closely. Hydrogen's very "zen" about the whole thing; give one up, take one in, share, it's all good.

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The position of Hydrogen in the periodic table is an often cited problem Dr Eric Scerri. Its placement in most table presentations bows to the physicists' spdf model and is over Li where it sticks out like a sore thumb. Some tables place it over F, some over C and some try to avoid the issue by placing it mid-air, floating over the transition elements. Dr Mark Leach's database of periodic tables has many; also the common web search machines will show many "familiar periodic table" images as well as many novelties.

The 3rd International Conference on the Periodic Table occurred last year (2012). As long as the spdf model is the only acceptible model, this issue might never be resolved.

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