Period 1 element

All other periods in the periodic table contain at least eight elements, and it is often helpful to consider periodic trends across the period. However, period 1 contains only two elements, so this concept does not apply here.

In terms of vertical trends down groups, helium can be seen as a typical noble gas at the head of the IUPAC group 18, but as discussed below, hydrogen's chemistry is unique and it is not easily assigned to any group.[1]

The first electron shell, n = 1, consists of only one orbital, and the maximum number of valence electrons that a period 1 element can accommodate is two, both in the 1s orbital. The valence shell lacks "p" or any other kind of orbitals due to the general l < n constraint on the quantum numbers. Therefore, period 1 has exactly two elements. Although both hydrogen and helium are in the s-block, neither of them behaves similarly to other s-block elements. Their behaviour is so different from the other s-block elements that there is considerable disagreement over where these two elements should be placed in the periodic table.

Simply following electron configurations, hydrogen (electronic configuration 1s¹) and helium (1s²) should be placed in groups 1 and 2, above lithium (1s²2s¹) and beryllium (1s²2s²).[2] While such a placement is common for hydrogen, it is rarely used for helium outside of the context of illustrating the electron configurations. Usually, hydrogen is placed in group 1, and helium in group 18: this is the placement found on the IUPAC periodic table.[3] Some variation can be found on both these matters.[4]

Like the group 1 metals, hydrogen has one electron in its outermost shell[2] and typically loses its only electron in chemical reactions.[5] It has some metal-like chemical properties, being able to displace some metals from their salts.[5] But hydrogen forms a diatomic nonmetallic gas at standard conditions, unlike the alkali metals which are reactive solid metals. This and hydrogen's formation of hydrides, in which it gains an electron, brings it close to the properties of the halogens which do the same[5] (though it is rarer for hydrogen to form H⁻ than H⁺).[6] Moreover, the lightest two halogens (fluorine and chlorine) are gaseous like hydrogen at standard conditions.[5] Some properties of hydrogen are not a good fit for either group: hydrogen is neither highly oxidising nor highly reducing and is not reactive with water.[6] Hydrogen thus has properties corresponding to both those of the alkali metals and the halogens, but matches neither group perfectly, and is thus difficult to place by its chemistry.[5] Therefore, while the electronic placement of hydrogen in group 1 predominates, some rarer arrangements show either hydrogen in group 17,[7] duplicate hydrogen in both groups 1 and 17,[8][9] or float it separately from all groups.[9][10][4] The possibility of "floating" hydrogen has nonetheless been criticised by Eric Scerri, who points out that removing it from all groups suggests that it is being excluded from the periodic law, when all elements should be subject to that law.[11] A few authors have advocated more unusual placements for hydrogen, such as group 13 or group 14, on the grounds of trends in ionisation energy, electron affinity, and electronegativity.[6]

Helium is an unreactive noble gas at standard conditions, and has a full outer shell: these properties are like the noble gases in group 18, but not at all like the reactive alkaline earth metals of group 2. Therefore, helium is nearly universally placed in group 18[3] which its properties best match.[4] However, helium only has two outer electrons in its outer shell, whereas the other noble gases have eight; and it is an s-block element, whereas all other noble gases are p-block elements. Also, solid helium crystallises in a hexagonal close-packed structure, which matches beryllium and magnesium in group 2, but not the other noble gases in group 18.[12] In these ways helium better matches the alkaline earth metals.[2][4] Therefore, tables with both hydrogen and helium floating outside all groups may rarely be encountered.[10][4]

A few chemists, such as Henry Bent, have advocated that the electronic placement in group 2 be adopted for helium.[13][12][14][15][16] This assignment is also found in Charles Janet's left-step table. Arguments for this often rest on the first-row anomaly trend (s >> p > d > f), which states that the first element of each group often behaves quite differently from the succeeding ones: the difference is largest in the s-block (H and He), is moderate for the p-block (B to Ne), and is less pronounced for the d- and f-blocks.[13] Thus helium as the first s² element before the alkaline earth metals stands out as anomalous in a way that helium as the first noble gas does not.[13] The normalized ionization potentials and electron affinities show better trends with helium in group 2 than in group 18; helium is expected to be slightly more reactive than neon (which breaks the general trend of reactivity in the noble gases, where the heavier ones are more reactive); and predicted helium compounds often lack neon analogues even theoretically, but sometimes have beryllium analogues.[17][18][19]

• Bloch, D. R. (2006). Organic Chemistry Demystified. McGraw-Hill Professional. ISBN 0-07-145920-0.