Hydrogen Content from the guide to life, the universe and everything


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Overview of the Periodic Table
Hydrogen | Group 1 - Alkali Metals | Group 2 - Alkaline Earth Metals | Group 12


Hydrogen is the most abundant element in the universe and forms more compounds than any other element in the periodic table. It is also the third most abundant element in the Earth's crust, behind silicon and oxygen. It exists as three naturally occurring isotopes:

  1. Hydrogen (1H), which has nuclei each containing just one proton

  2. Deuterium (2D), which contains an extra neutron in each nucleus and constitutes 0.0156% of hydrogen

  3. Tritium (3T)

Tritium only occurs in minute quantities and is produced in the upper atmosphere by the action of cosmic rays. Each of its nuclei contains one proton and two neutrons, and is unstable having a radioactive half life of 12.4 years. Molecular hydrogen, or dihydrogen (H2), is a colourless and odourless gas and has a boiling point of just 20.4 K. An important feature in the chemistry of hydrogen is the role of the H+ ion in acid/base chemistry.

Preparation of Hydrogen

Elemental hydrogen can be produced by the action of water or dilute acids on the group 1 elements (the alkali metals) and the group 2 elements (the alkaline earth metals). It can also be produced by the electrolysis of aqueous salt solution and the hydrolysis of metal hydrides.

CaH2 + H20 → Ca(OH)2 + H2

On large scales it is produced in industry by the steam reforming reaction where methane and light hydrocarbons react with a nickel catalyst at about 750 °C to form a mixture of carbon monoxide and hydrogen. This mixture is called 'synthesis gas', or 'syngas'.

CH4 + H20 → CO + H2

It can also be produced by the water-gas shift reaction which can be catalysed by many transition metal complexes at relatively high temperatures. The carbon dioxide produced can then be used to produce dry ice.

CO + H20 → CO2 + H2

Hydrogen has many uses. An important process is the production of methanol, ethanol and higher alcohols from syngas. These are important chemical feed stocks and this reaction has been widely used in countries where oil imports have been blockaded, such as in South Africa in the times of apartheid.

2 CO + H2 → CH3CH2OH + H2O

It is also widely used for hydrogenation reactions of organic molecules, for example the hydrogenation of alkenes to alkanes.

R2C=CHR2 + H2 → R2HC-CHR2

It can be used in the formation of elemental metals from their oxides, for example:

MO2 + 2 H2 → M + 2 H2O

The Hydrogen Bond

The hydrogen bond is the strongest form of intermolecular attraction, the nature of which has been extensively studied. Hydrogen bonding between molecules is a key part of the chemistry of water, aqueous solutions, compounds containing hydroxyl (OH) groups and is of vital importance to biochemistry. It comes about due to the polarity of X-H bonds in molecules, where X is some electronegative element such as F, O, N or Cl. The electrons in the X-H bond are pulled mostly over to the atom X leaving the hydrogen atoms with a partial positive charge (δ+). Because hydrogen has no inner electron shells to shield the charge of its positively charged nucleus, it can then interact with a negatively charged, electron rich, atom Y on another molecule forming the hydrogen bond.


The H···Y interaction is considered to be essentially electrostatic and the distance between the two atoms is generally longer than a normal H-Y bond, while the X-H bond is usually slightly lengthened. In the strongest hydrogen bonds, the X-H and H-Y distances can become almost equal. Here, the H-Y interaction is closer to a covalent interaction.

The influence of hydrogen bonding is most evident in the behaviour of the XH3 compounds of the group 15 elements; the XH2 compounds of group 16; and the XH acids of group 17. The boiling points of NH3, H20 and HF are much higher than those of PH3, H2S and HCl respectively. This is because the electronegativities of N (3.07), O (3.50) and F (4.10) are much greater than that of H (2.20), whereas the eletronegativities of P (2.06), S (2.44) and Cl (2.83) are much more comparable to that of H.

The X-H bonds in these three first row compounds are much more polar than those of the second row compounds leading to stronger intermolecular associations through hydrogen bonding and higher boiling point. Conversely, the electronegativity of C (2.50) is much less than for N, O and F and comparable to H leading to very little polarity in the C-H bonds of methane, CH4. Because of this CH4 has a much lower boiling point than SiH4.

Hydrogen bonds are vital for life, which needed liquid water in which to start. Without hydrogen bonding interactions, water would be expected to boil at around -100 °C rather than +100 °C, and so there would be no liquid water on the Earth. They also play a major structural role in biochemical systems. Proteins are chains of amino acids containing N-H bonds that have complex and specific shapes to carry out their function. This framework is held together by the formation of intramolecular N-H···O and N-H···N hydrogen bonds between different parts of the chain. The structure of DNA is also dependent on hydrogen bonds. DNA molecules each consist of two chains of nucleotides containing bases commonly referred to as 'A', 'T', 'C' and 'G'. The two chains are held together in a double helix structure by forming hydrogen bonded AT and CG base pairs.


Almost any compound of any element containing hydrogen could be called a hydride, although specifically a hydridic compound is an H- ion donor, such as LiAlH4 or NaBH4. Hydrogen-containing compounds can be divided into three broad groups - ionic, covalent and acidic - where the hydrogen exists and reacts as the hydride ion H-, the radical H or the bear proton H+.

Hydrogen forms neutral compounds of the type XH4 with elements of group 14, the most well-known example of which is methane, CH4. Group 15 elements form XH3 compounds that act as bases, such as ammonia, NH3. The XH2 compounds of group 16, such as H2O and H2S, are weakly acidic or even amphoteric1. The HX compounds of group 17 are strongly acidic with the hydrogen atom tending toward the form of H+. Boron will form many different compounds with hydrogen called boranes, some with very complex structures. All these hydride compounds are covalently bonded. Although LiAlH4 and NaBH4 are H- donors they also contain covalent Al-H and B-H bonds. LiAlH4 can be formed from the reaction of LiH with Al2Cl6 and NaBH4 from the reaction of NaH with diborane.

8 LiH + Al2Cl6 → 2 LiAlH4 + 6 LiCl
2 NaH + B2H6 → 2 NaBH4

There are hydride compounds that are ionic containing H- and a positively charged metal ion. These are mainly confined to the elements of groups 1 and 2. Although the hydrides of Li, Na, Be and Mg have M-H covalent bonds, the structure of BeH2 can be thought of as a polymer containing Be-H-Be bridges. The hydrides of groups 1 and 2 are formed from the reaction of hydrogen gas with the elemental metals at high temperatures

2 M (l) + H2 (g) → 2 MH (s)
M (l) + H2 (g) → MH2(s)

These hydrides are very reactive with water, forming hydrogen gas and the metal hydroxide in their reactions with it. This reaction for CaH2 allows it to act as a drying agent to remove traces of water in chlorinated organic solvents, that may be later used for air and moisture sensitive reactions.

CaH2 + 2 H2O → H2 + Ca(OH)2

There are also many hydrides of the transition metal elements with a wide range of structures, reactivity and properties. There are complexes that contain covalent M-H bonds such as HMn(CO)5, anionic compounds such as FeH6- and a wide range of nonstoichiometric compounds where there is not a whole number ratio of metal and hydrogen atoms. They are formed from the reaction of hydrogen gas with the metals at high temperature and have complicated structures.

Dihydrogen itself can also act as a ligand in transition metal complexes under special circumstances. It binds to the metal side with a triangular arrangement of the two hydrogen atoms and the metal atom (figure 1). In these complexes we get overlap of the H-H σ-orbital with an empty atomic orbital on the metal. The metal can then use an orbital to overlap in a way similar to π bonding and donates electrons into the H-H σ*-orbital. This weakens the H-H bond and in almost all cases leads it breaking resulting in two MH2.

Figure 1: Generic structure of a metal dihydrogen complex and a metal dihydride.

H---H H H | \ / M M

A rare example of a complex where the H-H ligand is intact at least in equilibrium is (dppe)2(CO)Mo(H2), where dppe is a ligand with the structure (C6H5)2PCH2CH2P(C6H5)2. This ligand bonds to the metal through both P atoms.

1Amphoteric compounds can act as either acids or bases, depending on the conditions.

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