Why Hydrogen Could Sit Above the Alkali Metals or the Halogens

Hydrogen is the simplest atom — one proton and one electron — yet it refuses to fit neatly into any group of the periodic table. It fuels rockets, turns vegetable oil into solid fat, makes ammonia for fertilisers and combines with most elements to form hydrides of very different kinds.

This lesson covers hydrogen's position in the periodic table, the preparation, properties and uses of dihydrogen, and the three classes of hydrides.

Hydrogen has a single electron in its 1s orbital, so, like the alkali metals, it can lose one electron, and, like the halogens, it can gain one electron to reach a noble gas configuration — which is why it is placed on its own.

Resemblance to the alkali metals:

- Electronic configuration $1s^1$, like $n{s}^1$ for Group 1
- Forms a positive ion, ${\mathrm{H}}^{+}$, which exists in water as ${\mathrm{H}}_3{\mathrm{O}}^{+}$
- Shows the +1 oxidation state and forms HCl, ${\mathrm{H}}_2\mathrm{O}$ and ${\mathrm{H}}_2\mathrm{S}$, just as sodium forms NaCl, $\mathrm{N}{\mathrm{a}}_2\mathrm{O}$ and $\mathrm{N}{\mathrm{a}}_2\mathrm{S}$
- Acts as a reducing agent: $\mathrm{CuO}+{\mathrm{H}}_2\to \mathrm{Cu}+{\mathrm{H}}_2\mathrm{O}$

Resemblance to the halogens:

- One electron short of a noble gas configuration
- Forms a hydride ion, ${\mathrm{H}}^{-}$, in NaH, just as chlorine forms $\mathrm{C}{\mathrm{l}}^{-}$ in NaCl
- Exists as diatomic molecules, ${\mathrm{H}}_2$, like $\mathrm{C}{\mathrm{l}}_2$
- Shows the -1 oxidation state in metal hydrides

Differences from both:

- Its ionisation enthalpy, 1312 kJ mol${}^{-1}$, is far higher than that of sodium, 496 kJ mol${}^{-1}$, and closer to the halogens
- It is a non-metallic gas, unlike the alkali metals
- ${\mathrm{H}}^{+}$ is a bare proton, so tiny that it never exists free in solution

An everyday example. The periodic table chart on a classroom wall often shows hydrogen above Group 1 but set slightly apart — a visual reminder that it belongs fully to neither family.

The substance. Hydrogen is placed separately, not wrongly — its resemblances to Group 1 and Group 17 are real, but so are the differences that make it unique.

Dihydrogen is made in the laboratory from zinc and dilute acid, and on a large scale from steam and hydrocarbons or by electrolysis; it is a light, colourless gas that burns, reduces metal oxides and combines with many elements.

Laboratory preparation:

- Granulated zinc with dilute acid: $\mathrm{Zn}+2{\mathrm{H}}^{+}\to \mathrm{Z}{\mathrm{n}}^{2+}+{\mathrm{H}}_2$
- Zinc with aqueous sodium hydroxide: $\mathrm{Zn}+2NaOH\to \mathrm{N}{\mathrm{a}}_2\mathrm{Zn}{\mathrm{O}}_2+{\mathrm{H}}_2$

Commercial production:

- Electrolysis of acidified water, with hydrogen collected at the cathode
- Steam reforming: $\mathrm{C}{\mathrm{H}}_4+{\mathrm{H}}_2\mathrm{O}\to \mathrm{CO}+3{\mathrm{H}}_2$ at about 1270 K with a nickel catalyst; the mixture of CO and hydrogen is called syngas
- Water-gas shift reaction: $\mathrm{CO}+{\mathrm{H}}_2\mathrm{O}\to \mathrm{C}{\mathrm{O}}_2+{\mathrm{H}}_2$ at 673 K with a catalyst, raising the yield of hydrogen

Properties:

- Colourless, odourless, tasteless, the lightest gas, and only slightly soluble in water
- Its H-H bond enthalpy, 435.8 kJ mol${}^{-1}$, is very high, so it reacts slowly at room temperature
- Burns in oxygen with a pale blue flame: $2{\mathrm{H}}_2+{\mathrm{O}}_2\to 2{\mathrm{H}}_2\mathrm{O}$
- Combines with nitrogen to form ammonia and with halogens to form hydrogen halides

Uses:

- Making ammonia for fertilisers, and methanol: $\mathrm{CO}+2{\mathrm{H}}_2\to \mathrm{C}{\mathrm{H}}_3\mathrm{OH}$
- Hydrogenation of vegetable oils with a nickel catalyst
- Liquid hydrogen as rocket fuel, and hydrogen in fuel cells that generate electricity

An everyday example. Vanaspati, the solid cooking fat sold in tins across India, is made by passing hydrogen through vegetable oil over a nickel catalyst, turning some C=C double bonds into single bonds.

The substance. Very pure zinc reacts only slowly with dilute acid — the impurities in ordinary granulated zinc speed up the release of hydrogen.

Hydrides are binary compounds of hydrogen, classified as ionic hydrides of the s-block metals, covalent hydrides of the p-block elements, and metallic hydrides of the d- and f-block elements.

Ionic or saline hydrides:

- Formed by highly electropositive s-block metals: LiH, NaH, $\mathrm{Ca}{\mathrm{H}}_2$
- Crystalline, non-volatile solids with high melting points
- Conduct electricity when molten, releasing hydrogen at the anode, because hydrogen is present as ${\mathrm{H}}^{-}$
- React vigorously with water: $\mathrm{NaH}+{\mathrm{H}}_2\mathrm{O}\to \mathrm{NaOH}+{\mathrm{H}}_2$

Covalent or molecular hydrides:

- Formed by p-block elements, mostly as volatile molecules
- Electron-deficient — too few electrons for ordinary two-electron bonds, as in diborane, ${\mathrm{B}}_2{\mathrm{H}}_6$, of Group 13
- Electron-precise — exactly enough electrons for all bonds, as in $\mathrm{C}{\mathrm{H}}_4$ of Group 14
- Electron-rich — extra electrons held as lone pairs, as in $\mathrm{N}{\mathrm{H}}_3$, ${\mathrm{H}}_2\mathrm{O}$ and HF of Groups 15 to 17

Worked example. Count the lone pairs on the central atom: nitrogen in $\mathrm{N}{\mathrm{H}}_3$ has 1, oxygen in ${\mathrm{H}}_2\mathrm{O}$ has 2, and fluorine in HF has 3 — which is why these hydrides can form hydrogen bonds.

Metallic or interstitial hydrides:

- Formed by many d- and f-block metals, such as palladium, titanium and lanthanum
- Often non-stoichiometric, with formulas such as $\mathrm{La}{\mathrm{H}}_{2.87}$
- Palladium absorbs a large volume of hydrogen, so such hydrides are studied for hydrogen storage

An everyday example. Engineers designing hydrogen-powered buses consider metal hydrides as a way to store hydrogen safely in a compact tank instead of a high-pressure cylinder.

The substance. Hydrides change character across the periodic table — ionic on the left, covalent on the right and metallic in the middle, following the electronegativity of the element.